dpsl dpsl

Organic ChemistrySecond Edition

The INSTANT NOTES series

Series Editor BD Hames School of Biochemistry and Molecular BiologyUniversity of Leeds Leeds UK

Animal Biology 2nd editionBiochemistry 2nd editionBioinformaticsChemistry for Biologists 2nd editionDevelopmental BiologyEcology 2nd editionImmunology 2nd editionGenetics 2nd editionMicrobiology 2nd editionMolecular Biology 2nd editionNeurosciencePlant Biology

Chemistry seriesConsulting Editor Howard Stanbury

Analytical ChemistryInorganic Chemistry 2nd editionMedicinal ChemistryOrganic Chemistry 2nd editionPhysical Chemistry

Psychology seriesSub-series Editor Hugh Wagner Dept of Psychology University of CentralLancashire Preston UK

Psychology

Forthcoming titlesCognitive PsychologyPhysiological Psychology

Organic Chemistry

Second Edition

G L Patrick

Department of Chemistry and Chemical Engineering Paisley University Paisley Scotland

copy Garland ScienceBIOS Scientific Publishers 2004

First published 2000Second edition published 2004

All rights reserved No part of this book may be reproduced or transmitted in any form or by any meanswithout permission

A CIP catalogue record for this book is available from the British Library

ISBN 1 85996 264 5

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Library of Congress Cataloging-in-Publication Data

Patrick Graham LOrganic chemistry GL Patrickmdash2nd ed

p cm mdash (Instant notes series)Includes bibliographical references and indexISBN 1-85996-264-5 (alk paper)1 Chemistry OrganicmdashOutlines syllabi etc I Title Instant notes organic chemistry

II Title III Series

QD2565P37 2003547mdashdc22

Production Editor Andrew Watts

This edition published in the Taylor amp Francis e-Library 2005

To purchase your own copy of this or any of Taylor amp Francis or Routledges

collection of thousands of eBooks please go to wwweBookstore

ISBN 0-203-42761-0 Master e-book ISBN

ISBN 0-203-44168-0 (Adobe eReader Format)(Print Edition)

tandfcoukrdquo

Preface ix

Section A ndash Structure and bonding 1A1 Atomic structure of carbon 1A2 Covalent bonding and hybridization 3A3 sp3 Hybridization 5A4 sp2 Hybridization 8A5 sp Hybridization 14A6 Bonds and hybridized centers 17

Section B ndash Alkanes and cycloalkanes 19B1 Definition 19B2 Drawing structures 20B3 Nomenclature 22

Section C ndash Functional groups 27C1 Recognition of functional groups 27C2 Aliphatic and aromatic functional groups 29C3 Intermolecular bonding 30C4 Properties and reactions 33C5 Nomenclature of compounds with functional groups 35C6 Primary secondary tertiary and quaternary

nomenclature 43

Section D ndash Stereochemistry 45D1 Constitutional isomers 45D2 Configurational isomers ndash alkenes and cycloalkanes 46D3 Configurational isomers ndash optical isomers 49D4 Conformational isomers 56

Section E ndash Nucleophiles and electrophiles 63E1 Definition 63E2 Charged species 64E3 Neutral inorganic species 66E4 Organic structures 69

Section F ndash Reactions and mechanisms 73F1 Reactions 73F2 Mechanisms 75

Section G ndash Acidndashbase reactions 79G1 BroslashnstedndashLowry acids and bases 79G2 Acid strength 82G3 Base strength 88G4 Lewis acids and bases 94G5 Enolates 95

CONTENTS

Section H ndash Alkenes and alkynes 99H1 Preparation of alkenes and alkynes 99H2 Properties of alkenes and alkynes 101H3 Electrophilic addition to symmetrical alkenes 105H4 Electrophilic addition to unsymmetrical alkenes 112H5 Carbocation stabilization 115H6 Reduction and oxidation of alkenes 117H7 Hydroboration of alkenes 121H8 Electrophilic additions to alkynes 124H9 Reduction of alkynes 127H10 Alkylation of terminal alkynes 129H11 Conjugated dienes 131

Section I ndash Aromatic chemistry 135I1 Aromaticity 135I2 Preparation and properties 137I3 Electrophilic substitutions of benzene 139I4 Synthesis of mono-substituted benzenes 147I5 Electrophilic substitutions of mono-substituted

aromatic rings 150I6 Synthesis of di- and tri-substituted benzenes 160I7 Oxidation and reduction 164

Section J ndash Aldehydes and ketones 167J1 Preparation 167J2 Properties 169J3 Nucleophilic addition 173J4 Nucleophilic addition ndash charged nucleophiles 175J5 Electronic and steric effects 181J6 Nucleophilic addition ndash nitrogen nucleophiles 184J7 Nucleophilic addition ndash oxygen and sulfur

nucleophiles 187J8 Reactions of enolate ions 191J9 α-Halogenation 198J10 Reduction and oxidation 200J11 αβ-Unsaturated aldehydes and ketones 202

Section K ndash Carboxylic acids and carboxylic acid derivatives 205K1 Structure and properties 205K2 Nucleophilic substitution 209K3 Reactivity 213K4 Preparations of carboxylic acids 217K5 Preparations of carboxylic acid derivatives 219K6 Reactions 224K7 Enolate reactions 234

Section L ndash Alkyl halides 239L1 Preparation and physical properties of alkyl halides 239L2 Nucleophilic substitution 242L3 Factors affecting SN2 versus SN1 reactions 247L4 Elimination 252L5 Elimination versus substitution 256L6 Reactions of alkyl halides 258L7 Organometallic reactions 261

vi Contents

Section M ndash Alcohols phenols and thiols 263M1 Preparation of alcohols 263M2 Preparation of phenols 264M3 Properties of alcohols and phenols 266M4 Reactions of alcohols 270M5 Reactions of phenols 277M6 Chemistry of thiols 281

Section N ndash Ethers epoxides and thioethers 283N1 Preparation of ethers epoxides and thioethers 283N2 Properties of ethers epoxides and thioethers 286N3 Reactions of ethers epoxides and thioethers 289

Section O ndash Amines and nitriles 295O1 Preparation of amines 295O2 Properties of amines 299O3 Reactions of amines 305O4 Chemistry of nitriles 311

Section P ndash Organic spectroscopy and analysis 315P1 Spectroscopy 315P2 Visible and ultra violet spectroscopy 317P3 Infra-red spectroscopy 322P4 Proton nuclear magnetic resonance spectroscopy 324P5 13C nuclear magnetic resonance spectroscopy 339P6 Mass spectroscopy 342

Further reading 347Index 349

Contents vii

This textbook aims to provide a comprehensive set of basic notes in organicchemistry which will be suitable for undergraduate students taking chemistrychemistry-related courses or courses which involve organic chemistry as anancillary subject The book concentrates on core topics which are most likelyto be common to those organic chemistry courses which follow on from a foun-dation or introductory general chemistry course

Organic chemistry is a subject which can lead some students to the heightsof ecstasy yet drive others up the wall Some students lsquoswitch onrsquo to it imme-diately while others can make neither head nor tail of it no matter how hardthey try Certainly one of the major problems in studying the subject is thevast amount of material which often has to be covered Many students blancheat the prospect of having to learn a seemingly endless number of reactionsand when it comes to drawing mechanisms and curly arrows they see onlya confusing maze of squiggly lines going everywhere yet nowhere Theconcepts of organic reaction mechanisms are often the most difficult to masterThese difficulties are often compounded by the fact that current textbooks inorganic chemistry are typically over 1200 pages long and can be quite expen-sive to buy

This book attempts to condense the essentials of organic chemistry into amanageable text of 310 pages which is student friendly and which does notcost an arm and a leg It does this by concentrating purely on the basics of thesubject without going into exhaustive detail or repetitive examplesFurthermore key notes at the start of each topic summarize the essential factscovered and help focus the mind on the essentials

Organic chemistry is a peculiar subject in that it becomes easier as you goalong This might seem an outrageous statement to make especially to a first-year student who is struggling to come to terms with the rules of nomenclaturetrying to memorize a couple of dozen reactions and making sense of mecha-nisms at the same time However these topics are the basics of the subject andonce they have been grasped the overall picture becomes clear

Understanding the mechanism of how a reaction takes place is particularlycrucial in this It brings a logic to the reactions of the different functional groupsThis in turn transforms a list of apparently unrelated facts into a sensible themewhich makes remembering the reactions a lsquopiece of cakersquo (well nearly)

Once this happy state of affairs has been reached the relevance of organicchemistry to other subjects such as genetics and biochemistry suddenly leapsoff the page Understanding organic chemistry leads to a better understandingof life chemistry and how the body works at the molecular level It also helpsin the understanding of the molecular mechanisms involved in disease andbodily malfunction leading in turn to an understanding of how drugs can bedesigned to cure these disease states ndash the science of medicinal chemistry

And thatrsquos not all An understanding of organic chemistry will help the indus-trial chemist or chemical engineer faced with unexpected side-reactions in achemical process and the agro-scientist trying to understand the molecularprocesses taking place within plants and crops and it will assist in the designand synthesis of new herbicides and fungicides which will be eco-friendly It

PREFACE

will aid the forensic scientist wishing to analyze a nondescript white powderndash is it heroin or flour

The list of scientific subject areas involving organic chemistry is endless ndashdesigning spacesuits developing new photographic dyes inventing new mole-cular technology in microelectronics ndash one could go on and on Organicchemistry is an exciting subject since it leads to an essential understanding ofmolecules and their properties

The order in which the early topics of this book are presented is importantThe first two sections cover structure and bonding which are crucial to latersections Just why does carbon form four bonds What is hybridization

The third section on functional groups is equally crucial if students are to becapable of categorizing the apparent maze of reactions which organiccompounds can undergo It is followed by section D on stereochemistry thensections E and F in which the basic theory of reactions and mechanisms iscovered What are nucleophiles and electrophiles What does a mechanismrepresent What does a curly arrow mean

The remaining sections can be used in any order These look at the reactionsand mechanisms of the common functional groups which are important inchemistry and biochemistry

It is hoped that students will find this textbook useful in their studies andthat once they have grasped what organic chemistry is all about they will readmore widely and enter a truly exciting world of molecular science

x Preface

Section A ndash Structure and bonding

A1 ATOMIC STRUCTURE OF CARBON

Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table This means that thereare two shells of atomic orbitals available for these electrons The first shell closestto the nucleus has a single s orbital ndash the 1s orbital The second shell has a single sorbital (the 2s orbital) and three p orbitals (3 2p) Therefore there are a total offive atomic orbitals into which these six electrons can fit The s orbitals are sphericalin shape with the 2s orbital being much larger then the 1s orbital The p orbitalsare dumbbell-shaped and are aligned along the x y and z axes Therefore they areassigned the 2px 2py and 2pz atomic orbitals (Fig 1)

Key Notes

The atomic orbitals available for the six electrons of carbon are the s orbitalin the first shell the s orbital in the second shell and the three p orbitals inthe second shell The 1s and 2s orbitals are spherical in shape The 2porbitals are dumbbell in shape and can be assigned 2px 2py or 2pz depend-ing on the axis along which they are aligned

The 1s orbital has a lower energy than the 2s orbital which has a lowerenergy than the 2p orbitals The 2p orbitals have equal energy (ie theyare degenerate)

Carbon is in the second row of the periodic table and has six electrons whichwill fill up lower energy atomic orbitals before entering higher energyorbitals (aufbau principle) Each orbital is allowed a maximum of two elec-trons of opposite spin (Pauli exclusion principle) When orbitals of equalenergy are available electrons will occupy separate orbitals before pairingup (Hundrsquos rule) Thus the electronic configuration of a carbon atom is 1s2

2s2 2px1 2py

1

Related topic Covalent bonding andhybridization (A2)

Electronicconfiguration

Atomic orbitals

Energy levels

1s 2s 2px 2py 2pz

y

xz

y

xz

y

xz

y

xz

y

xz

Fig 1 Atomic orbitals

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Organic ChemistrySecond Edition

The INSTANT NOTES series

Series Editor BD Hames School of Biochemistry and Molecular BiologyUniversity of Leeds Leeds UK

Animal Biology 2nd editionBiochemistry 2nd editionBioinformaticsChemistry for Biologists 2nd editionDevelopmental BiologyEcology 2nd editionImmunology 2nd editionGenetics 2nd editionMicrobiology 2nd editionMolecular Biology 2nd editionNeurosciencePlant Biology

Chemistry seriesConsulting Editor Howard Stanbury

Analytical ChemistryInorganic Chemistry 2nd editionMedicinal ChemistryOrganic Chemistry 2nd editionPhysical Chemistry

Psychology seriesSub-series Editor Hugh Wagner Dept of Psychology University of CentralLancashire Preston UK

Psychology

Forthcoming titlesCognitive PsychologyPhysiological Psychology

Organic Chemistry

Second Edition

G L Patrick

Department of Chemistry and Chemical Engineering Paisley University Paisley Scotland

copy Garland ScienceBIOS Scientific Publishers 2004

First published 2000Second edition published 2004

All rights reserved No part of this book may be reproduced or transmitted in any form or by any meanswithout permission

A CIP catalogue record for this book is available from the British Library

ISBN 1 85996 264 5

Garland ScienceBIOS Scientific Publishers4 Park Square Milton Park Abingdon Oxon OX14 4RN UK and29 West 35th Street New York NY 10001ndash2299 USAWorld Wide Web home page wwwbioscouk

Garland ScienceBIOS Scientific Publishers is a member of the Taylor amp Francis Group

Distributed in the USA byFulfilment CenterTaylor amp Francis10650 Toebben DriveIndependence KY 41051 USAToll Free Tel +1 800 634 7064 E-mail taylorandfrancisthomsonlearningcom

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Library of Congress Cataloging-in-Publication Data

Patrick Graham LOrganic chemistry GL Patrickmdash2nd ed

p cm mdash (Instant notes series)Includes bibliographical references and indexISBN 1-85996-264-5 (alk paper)1 Chemistry OrganicmdashOutlines syllabi etc I Title Instant notes organic chemistry

II Title III Series

QD2565P37 2003547mdashdc22

Production Editor Andrew Watts

This edition published in the Taylor amp Francis e-Library 2005

To purchase your own copy of this or any of Taylor amp Francis or Routledges

collection of thousands of eBooks please go to wwweBookstore

ISBN 0-203-42761-0 Master e-book ISBN

ISBN 0-203-44168-0 (Adobe eReader Format)(Print Edition)

tandfcoukrdquo

Preface ix

Section A ndash Structure and bonding 1A1 Atomic structure of carbon 1A2 Covalent bonding and hybridization 3A3 sp3 Hybridization 5A4 sp2 Hybridization 8A5 sp Hybridization 14A6 Bonds and hybridized centers 17

Section B ndash Alkanes and cycloalkanes 19B1 Definition 19B2 Drawing structures 20B3 Nomenclature 22

Section C ndash Functional groups 27C1 Recognition of functional groups 27C2 Aliphatic and aromatic functional groups 29C3 Intermolecular bonding 30C4 Properties and reactions 33C5 Nomenclature of compounds with functional groups 35C6 Primary secondary tertiary and quaternary

nomenclature 43

Section D ndash Stereochemistry 45D1 Constitutional isomers 45D2 Configurational isomers ndash alkenes and cycloalkanes 46D3 Configurational isomers ndash optical isomers 49D4 Conformational isomers 56

Section E ndash Nucleophiles and electrophiles 63E1 Definition 63E2 Charged species 64E3 Neutral inorganic species 66E4 Organic structures 69

Section F ndash Reactions and mechanisms 73F1 Reactions 73F2 Mechanisms 75

Section G ndash Acidndashbase reactions 79G1 BroslashnstedndashLowry acids and bases 79G2 Acid strength 82G3 Base strength 88G4 Lewis acids and bases 94G5 Enolates 95

CONTENTS

Section H ndash Alkenes and alkynes 99H1 Preparation of alkenes and alkynes 99H2 Properties of alkenes and alkynes 101H3 Electrophilic addition to symmetrical alkenes 105H4 Electrophilic addition to unsymmetrical alkenes 112H5 Carbocation stabilization 115H6 Reduction and oxidation of alkenes 117H7 Hydroboration of alkenes 121H8 Electrophilic additions to alkynes 124H9 Reduction of alkynes 127H10 Alkylation of terminal alkynes 129H11 Conjugated dienes 131

Section I ndash Aromatic chemistry 135I1 Aromaticity 135I2 Preparation and properties 137I3 Electrophilic substitutions of benzene 139I4 Synthesis of mono-substituted benzenes 147I5 Electrophilic substitutions of mono-substituted

aromatic rings 150I6 Synthesis of di- and tri-substituted benzenes 160I7 Oxidation and reduction 164

Section J ndash Aldehydes and ketones 167J1 Preparation 167J2 Properties 169J3 Nucleophilic addition 173J4 Nucleophilic addition ndash charged nucleophiles 175J5 Electronic and steric effects 181J6 Nucleophilic addition ndash nitrogen nucleophiles 184J7 Nucleophilic addition ndash oxygen and sulfur

nucleophiles 187J8 Reactions of enolate ions 191J9 α-Halogenation 198J10 Reduction and oxidation 200J11 αβ-Unsaturated aldehydes and ketones 202

Section K ndash Carboxylic acids and carboxylic acid derivatives 205K1 Structure and properties 205K2 Nucleophilic substitution 209K3 Reactivity 213K4 Preparations of carboxylic acids 217K5 Preparations of carboxylic acid derivatives 219K6 Reactions 224K7 Enolate reactions 234

Section L ndash Alkyl halides 239L1 Preparation and physical properties of alkyl halides 239L2 Nucleophilic substitution 242L3 Factors affecting SN2 versus SN1 reactions 247L4 Elimination 252L5 Elimination versus substitution 256L6 Reactions of alkyl halides 258L7 Organometallic reactions 261

vi Contents

Section M ndash Alcohols phenols and thiols 263M1 Preparation of alcohols 263M2 Preparation of phenols 264M3 Properties of alcohols and phenols 266M4 Reactions of alcohols 270M5 Reactions of phenols 277M6 Chemistry of thiols 281

Section N ndash Ethers epoxides and thioethers 283N1 Preparation of ethers epoxides and thioethers 283N2 Properties of ethers epoxides and thioethers 286N3 Reactions of ethers epoxides and thioethers 289

Section O ndash Amines and nitriles 295O1 Preparation of amines 295O2 Properties of amines 299O3 Reactions of amines 305O4 Chemistry of nitriles 311

Section P ndash Organic spectroscopy and analysis 315P1 Spectroscopy 315P2 Visible and ultra violet spectroscopy 317P3 Infra-red spectroscopy 322P4 Proton nuclear magnetic resonance spectroscopy 324P5 13C nuclear magnetic resonance spectroscopy 339P6 Mass spectroscopy 342

Further reading 347Index 349

Contents vii

This textbook aims to provide a comprehensive set of basic notes in organicchemistry which will be suitable for undergraduate students taking chemistrychemistry-related courses or courses which involve organic chemistry as anancillary subject The book concentrates on core topics which are most likelyto be common to those organic chemistry courses which follow on from a foun-dation or introductory general chemistry course

Organic chemistry is a subject which can lead some students to the heightsof ecstasy yet drive others up the wall Some students lsquoswitch onrsquo to it imme-diately while others can make neither head nor tail of it no matter how hardthey try Certainly one of the major problems in studying the subject is thevast amount of material which often has to be covered Many students blancheat the prospect of having to learn a seemingly endless number of reactionsand when it comes to drawing mechanisms and curly arrows they see onlya confusing maze of squiggly lines going everywhere yet nowhere Theconcepts of organic reaction mechanisms are often the most difficult to masterThese difficulties are often compounded by the fact that current textbooks inorganic chemistry are typically over 1200 pages long and can be quite expen-sive to buy

This book attempts to condense the essentials of organic chemistry into amanageable text of 310 pages which is student friendly and which does notcost an arm and a leg It does this by concentrating purely on the basics of thesubject without going into exhaustive detail or repetitive examplesFurthermore key notes at the start of each topic summarize the essential factscovered and help focus the mind on the essentials

Organic chemistry is a peculiar subject in that it becomes easier as you goalong This might seem an outrageous statement to make especially to a first-year student who is struggling to come to terms with the rules of nomenclaturetrying to memorize a couple of dozen reactions and making sense of mecha-nisms at the same time However these topics are the basics of the subject andonce they have been grasped the overall picture becomes clear

Understanding the mechanism of how a reaction takes place is particularlycrucial in this It brings a logic to the reactions of the different functional groupsThis in turn transforms a list of apparently unrelated facts into a sensible themewhich makes remembering the reactions a lsquopiece of cakersquo (well nearly)

Once this happy state of affairs has been reached the relevance of organicchemistry to other subjects such as genetics and biochemistry suddenly leapsoff the page Understanding organic chemistry leads to a better understandingof life chemistry and how the body works at the molecular level It also helpsin the understanding of the molecular mechanisms involved in disease andbodily malfunction leading in turn to an understanding of how drugs can bedesigned to cure these disease states ndash the science of medicinal chemistry

And thatrsquos not all An understanding of organic chemistry will help the indus-trial chemist or chemical engineer faced with unexpected side-reactions in achemical process and the agro-scientist trying to understand the molecularprocesses taking place within plants and crops and it will assist in the designand synthesis of new herbicides and fungicides which will be eco-friendly It

PREFACE

will aid the forensic scientist wishing to analyze a nondescript white powderndash is it heroin or flour

The list of scientific subject areas involving organic chemistry is endless ndashdesigning spacesuits developing new photographic dyes inventing new mole-cular technology in microelectronics ndash one could go on and on Organicchemistry is an exciting subject since it leads to an essential understanding ofmolecules and their properties

The order in which the early topics of this book are presented is importantThe first two sections cover structure and bonding which are crucial to latersections Just why does carbon form four bonds What is hybridization

The third section on functional groups is equally crucial if students are to becapable of categorizing the apparent maze of reactions which organiccompounds can undergo It is followed by section D on stereochemistry thensections E and F in which the basic theory of reactions and mechanisms iscovered What are nucleophiles and electrophiles What does a mechanismrepresent What does a curly arrow mean

The remaining sections can be used in any order These look at the reactionsand mechanisms of the common functional groups which are important inchemistry and biochemistry

It is hoped that students will find this textbook useful in their studies andthat once they have grasped what organic chemistry is all about they will readmore widely and enter a truly exciting world of molecular science

x Preface

Section A ndash Structure and bonding

A1 ATOMIC STRUCTURE OF CARBON

Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table This means that thereare two shells of atomic orbitals available for these electrons The first shell closestto the nucleus has a single s orbital ndash the 1s orbital The second shell has a single sorbital (the 2s orbital) and three p orbitals (3 2p) Therefore there are a total offive atomic orbitals into which these six electrons can fit The s orbitals are sphericalin shape with the 2s orbital being much larger then the 1s orbital The p orbitalsare dumbbell-shaped and are aligned along the x y and z axes Therefore they areassigned the 2px 2py and 2pz atomic orbitals (Fig 1)

Key Notes

The atomic orbitals available for the six electrons of carbon are the s orbitalin the first shell the s orbital in the second shell and the three p orbitals inthe second shell The 1s and 2s orbitals are spherical in shape The 2porbitals are dumbbell in shape and can be assigned 2px 2py or 2pz depend-ing on the axis along which they are aligned

The 1s orbital has a lower energy than the 2s orbital which has a lowerenergy than the 2p orbitals The 2p orbitals have equal energy (ie theyare degenerate)

Carbon is in the second row of the periodic table and has six electrons whichwill fill up lower energy atomic orbitals before entering higher energyorbitals (aufbau principle) Each orbital is allowed a maximum of two elec-trons of opposite spin (Pauli exclusion principle) When orbitals of equalenergy are available electrons will occupy separate orbitals before pairingup (Hundrsquos rule) Thus the electronic configuration of a carbon atom is 1s2

2s2 2px1 2py

1

Related topic Covalent bonding andhybridization (A2)

Electronicconfiguration

Atomic orbitals

Energy levels

1s 2s 2px 2py 2pz

y

xz

y

xz

y

xz

y

xz

y

xz

Fig 1 Atomic orbitals

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

The INSTANT NOTES series

Series Editor BD Hames School of Biochemistry and Molecular BiologyUniversity of Leeds Leeds UK

Animal Biology 2nd editionBiochemistry 2nd editionBioinformaticsChemistry for Biologists 2nd editionDevelopmental BiologyEcology 2nd editionImmunology 2nd editionGenetics 2nd editionMicrobiology 2nd editionMolecular Biology 2nd editionNeurosciencePlant Biology

Chemistry seriesConsulting Editor Howard Stanbury

Analytical ChemistryInorganic Chemistry 2nd editionMedicinal ChemistryOrganic Chemistry 2nd editionPhysical Chemistry

Psychology seriesSub-series Editor Hugh Wagner Dept of Psychology University of CentralLancashire Preston UK

Psychology

Forthcoming titlesCognitive PsychologyPhysiological Psychology

Organic Chemistry

Second Edition

G L Patrick

Department of Chemistry and Chemical Engineering Paisley University Paisley Scotland

copy Garland ScienceBIOS Scientific Publishers 2004

First published 2000Second edition published 2004

All rights reserved No part of this book may be reproduced or transmitted in any form or by any meanswithout permission

A CIP catalogue record for this book is available from the British Library

ISBN 1 85996 264 5

Garland ScienceBIOS Scientific Publishers4 Park Square Milton Park Abingdon Oxon OX14 4RN UK and29 West 35th Street New York NY 10001ndash2299 USAWorld Wide Web home page wwwbioscouk

Garland ScienceBIOS Scientific Publishers is a member of the Taylor amp Francis Group

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Library of Congress Cataloging-in-Publication Data

Patrick Graham LOrganic chemistry GL Patrickmdash2nd ed

p cm mdash (Instant notes series)Includes bibliographical references and indexISBN 1-85996-264-5 (alk paper)1 Chemistry OrganicmdashOutlines syllabi etc I Title Instant notes organic chemistry

II Title III Series

QD2565P37 2003547mdashdc22

Production Editor Andrew Watts

This edition published in the Taylor amp Francis e-Library 2005

To purchase your own copy of this or any of Taylor amp Francis or Routledges

collection of thousands of eBooks please go to wwweBookstore

ISBN 0-203-42761-0 Master e-book ISBN

ISBN 0-203-44168-0 (Adobe eReader Format)(Print Edition)

tandfcoukrdquo

Preface ix

Section A ndash Structure and bonding 1A1 Atomic structure of carbon 1A2 Covalent bonding and hybridization 3A3 sp3 Hybridization 5A4 sp2 Hybridization 8A5 sp Hybridization 14A6 Bonds and hybridized centers 17

Section B ndash Alkanes and cycloalkanes 19B1 Definition 19B2 Drawing structures 20B3 Nomenclature 22

Section C ndash Functional groups 27C1 Recognition of functional groups 27C2 Aliphatic and aromatic functional groups 29C3 Intermolecular bonding 30C4 Properties and reactions 33C5 Nomenclature of compounds with functional groups 35C6 Primary secondary tertiary and quaternary

nomenclature 43

Section D ndash Stereochemistry 45D1 Constitutional isomers 45D2 Configurational isomers ndash alkenes and cycloalkanes 46D3 Configurational isomers ndash optical isomers 49D4 Conformational isomers 56

Section E ndash Nucleophiles and electrophiles 63E1 Definition 63E2 Charged species 64E3 Neutral inorganic species 66E4 Organic structures 69

Section F ndash Reactions and mechanisms 73F1 Reactions 73F2 Mechanisms 75

Section G ndash Acidndashbase reactions 79G1 BroslashnstedndashLowry acids and bases 79G2 Acid strength 82G3 Base strength 88G4 Lewis acids and bases 94G5 Enolates 95

CONTENTS

Section H ndash Alkenes and alkynes 99H1 Preparation of alkenes and alkynes 99H2 Properties of alkenes and alkynes 101H3 Electrophilic addition to symmetrical alkenes 105H4 Electrophilic addition to unsymmetrical alkenes 112H5 Carbocation stabilization 115H6 Reduction and oxidation of alkenes 117H7 Hydroboration of alkenes 121H8 Electrophilic additions to alkynes 124H9 Reduction of alkynes 127H10 Alkylation of terminal alkynes 129H11 Conjugated dienes 131

Section I ndash Aromatic chemistry 135I1 Aromaticity 135I2 Preparation and properties 137I3 Electrophilic substitutions of benzene 139I4 Synthesis of mono-substituted benzenes 147I5 Electrophilic substitutions of mono-substituted

aromatic rings 150I6 Synthesis of di- and tri-substituted benzenes 160I7 Oxidation and reduction 164

Section J ndash Aldehydes and ketones 167J1 Preparation 167J2 Properties 169J3 Nucleophilic addition 173J4 Nucleophilic addition ndash charged nucleophiles 175J5 Electronic and steric effects 181J6 Nucleophilic addition ndash nitrogen nucleophiles 184J7 Nucleophilic addition ndash oxygen and sulfur

nucleophiles 187J8 Reactions of enolate ions 191J9 α-Halogenation 198J10 Reduction and oxidation 200J11 αβ-Unsaturated aldehydes and ketones 202

Section K ndash Carboxylic acids and carboxylic acid derivatives 205K1 Structure and properties 205K2 Nucleophilic substitution 209K3 Reactivity 213K4 Preparations of carboxylic acids 217K5 Preparations of carboxylic acid derivatives 219K6 Reactions 224K7 Enolate reactions 234

Section L ndash Alkyl halides 239L1 Preparation and physical properties of alkyl halides 239L2 Nucleophilic substitution 242L3 Factors affecting SN2 versus SN1 reactions 247L4 Elimination 252L5 Elimination versus substitution 256L6 Reactions of alkyl halides 258L7 Organometallic reactions 261

vi Contents

Section M ndash Alcohols phenols and thiols 263M1 Preparation of alcohols 263M2 Preparation of phenols 264M3 Properties of alcohols and phenols 266M4 Reactions of alcohols 270M5 Reactions of phenols 277M6 Chemistry of thiols 281

Section N ndash Ethers epoxides and thioethers 283N1 Preparation of ethers epoxides and thioethers 283N2 Properties of ethers epoxides and thioethers 286N3 Reactions of ethers epoxides and thioethers 289

Section O ndash Amines and nitriles 295O1 Preparation of amines 295O2 Properties of amines 299O3 Reactions of amines 305O4 Chemistry of nitriles 311

Section P ndash Organic spectroscopy and analysis 315P1 Spectroscopy 315P2 Visible and ultra violet spectroscopy 317P3 Infra-red spectroscopy 322P4 Proton nuclear magnetic resonance spectroscopy 324P5 13C nuclear magnetic resonance spectroscopy 339P6 Mass spectroscopy 342

Further reading 347Index 349

Contents vii

This textbook aims to provide a comprehensive set of basic notes in organicchemistry which will be suitable for undergraduate students taking chemistrychemistry-related courses or courses which involve organic chemistry as anancillary subject The book concentrates on core topics which are most likelyto be common to those organic chemistry courses which follow on from a foun-dation or introductory general chemistry course

Organic chemistry is a subject which can lead some students to the heightsof ecstasy yet drive others up the wall Some students lsquoswitch onrsquo to it imme-diately while others can make neither head nor tail of it no matter how hardthey try Certainly one of the major problems in studying the subject is thevast amount of material which often has to be covered Many students blancheat the prospect of having to learn a seemingly endless number of reactionsand when it comes to drawing mechanisms and curly arrows they see onlya confusing maze of squiggly lines going everywhere yet nowhere Theconcepts of organic reaction mechanisms are often the most difficult to masterThese difficulties are often compounded by the fact that current textbooks inorganic chemistry are typically over 1200 pages long and can be quite expen-sive to buy

This book attempts to condense the essentials of organic chemistry into amanageable text of 310 pages which is student friendly and which does notcost an arm and a leg It does this by concentrating purely on the basics of thesubject without going into exhaustive detail or repetitive examplesFurthermore key notes at the start of each topic summarize the essential factscovered and help focus the mind on the essentials

Organic chemistry is a peculiar subject in that it becomes easier as you goalong This might seem an outrageous statement to make especially to a first-year student who is struggling to come to terms with the rules of nomenclaturetrying to memorize a couple of dozen reactions and making sense of mecha-nisms at the same time However these topics are the basics of the subject andonce they have been grasped the overall picture becomes clear

Understanding the mechanism of how a reaction takes place is particularlycrucial in this It brings a logic to the reactions of the different functional groupsThis in turn transforms a list of apparently unrelated facts into a sensible themewhich makes remembering the reactions a lsquopiece of cakersquo (well nearly)

Once this happy state of affairs has been reached the relevance of organicchemistry to other subjects such as genetics and biochemistry suddenly leapsoff the page Understanding organic chemistry leads to a better understandingof life chemistry and how the body works at the molecular level It also helpsin the understanding of the molecular mechanisms involved in disease andbodily malfunction leading in turn to an understanding of how drugs can bedesigned to cure these disease states ndash the science of medicinal chemistry

And thatrsquos not all An understanding of organic chemistry will help the indus-trial chemist or chemical engineer faced with unexpected side-reactions in achemical process and the agro-scientist trying to understand the molecularprocesses taking place within plants and crops and it will assist in the designand synthesis of new herbicides and fungicides which will be eco-friendly It

PREFACE

will aid the forensic scientist wishing to analyze a nondescript white powderndash is it heroin or flour

The list of scientific subject areas involving organic chemistry is endless ndashdesigning spacesuits developing new photographic dyes inventing new mole-cular technology in microelectronics ndash one could go on and on Organicchemistry is an exciting subject since it leads to an essential understanding ofmolecules and their properties

The order in which the early topics of this book are presented is importantThe first two sections cover structure and bonding which are crucial to latersections Just why does carbon form four bonds What is hybridization

The third section on functional groups is equally crucial if students are to becapable of categorizing the apparent maze of reactions which organiccompounds can undergo It is followed by section D on stereochemistry thensections E and F in which the basic theory of reactions and mechanisms iscovered What are nucleophiles and electrophiles What does a mechanismrepresent What does a curly arrow mean

The remaining sections can be used in any order These look at the reactionsand mechanisms of the common functional groups which are important inchemistry and biochemistry

It is hoped that students will find this textbook useful in their studies andthat once they have grasped what organic chemistry is all about they will readmore widely and enter a truly exciting world of molecular science

x Preface

Section A ndash Structure and bonding

A1 ATOMIC STRUCTURE OF CARBON

Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table This means that thereare two shells of atomic orbitals available for these electrons The first shell closestto the nucleus has a single s orbital ndash the 1s orbital The second shell has a single sorbital (the 2s orbital) and three p orbitals (3 2p) Therefore there are a total offive atomic orbitals into which these six electrons can fit The s orbitals are sphericalin shape with the 2s orbital being much larger then the 1s orbital The p orbitalsare dumbbell-shaped and are aligned along the x y and z axes Therefore they areassigned the 2px 2py and 2pz atomic orbitals (Fig 1)

Key Notes

The atomic orbitals available for the six electrons of carbon are the s orbitalin the first shell the s orbital in the second shell and the three p orbitals inthe second shell The 1s and 2s orbitals are spherical in shape The 2porbitals are dumbbell in shape and can be assigned 2px 2py or 2pz depend-ing on the axis along which they are aligned

The 1s orbital has a lower energy than the 2s orbital which has a lowerenergy than the 2p orbitals The 2p orbitals have equal energy (ie theyare degenerate)

Carbon is in the second row of the periodic table and has six electrons whichwill fill up lower energy atomic orbitals before entering higher energyorbitals (aufbau principle) Each orbital is allowed a maximum of two elec-trons of opposite spin (Pauli exclusion principle) When orbitals of equalenergy are available electrons will occupy separate orbitals before pairingup (Hundrsquos rule) Thus the electronic configuration of a carbon atom is 1s2

2s2 2px1 2py

1

Related topic Covalent bonding andhybridization (A2)

Electronicconfiguration

Atomic orbitals

Energy levels

1s 2s 2px 2py 2pz

y

xz

y

xz

y

xz

y

xz

y

xz

Fig 1 Atomic orbitals

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Organic Chemistry

Second Edition

G L Patrick

Department of Chemistry and Chemical Engineering Paisley University Paisley Scotland

copy Garland ScienceBIOS Scientific Publishers 2004

First published 2000Second edition published 2004

All rights reserved No part of this book may be reproduced or transmitted in any form or by any meanswithout permission

A CIP catalogue record for this book is available from the British Library

ISBN 1 85996 264 5

Garland ScienceBIOS Scientific Publishers4 Park Square Milton Park Abingdon Oxon OX14 4RN UK and29 West 35th Street New York NY 10001ndash2299 USAWorld Wide Web home page wwwbioscouk

Garland ScienceBIOS Scientific Publishers is a member of the Taylor amp Francis Group

Distributed in the USA byFulfilment CenterTaylor amp Francis10650 Toebben DriveIndependence KY 41051 USAToll Free Tel +1 800 634 7064 E-mail taylorandfrancisthomsonlearningcom

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Library of Congress Cataloging-in-Publication Data

Patrick Graham LOrganic chemistry GL Patrickmdash2nd ed

p cm mdash (Instant notes series)Includes bibliographical references and indexISBN 1-85996-264-5 (alk paper)1 Chemistry OrganicmdashOutlines syllabi etc I Title Instant notes organic chemistry

II Title III Series

QD2565P37 2003547mdashdc22

Production Editor Andrew Watts

This edition published in the Taylor amp Francis e-Library 2005

To purchase your own copy of this or any of Taylor amp Francis or Routledges

collection of thousands of eBooks please go to wwweBookstore

ISBN 0-203-42761-0 Master e-book ISBN

ISBN 0-203-44168-0 (Adobe eReader Format)(Print Edition)

tandfcoukrdquo

Preface ix

Section A ndash Structure and bonding 1A1 Atomic structure of carbon 1A2 Covalent bonding and hybridization 3A3 sp3 Hybridization 5A4 sp2 Hybridization 8A5 sp Hybridization 14A6 Bonds and hybridized centers 17

Section B ndash Alkanes and cycloalkanes 19B1 Definition 19B2 Drawing structures 20B3 Nomenclature 22

Section C ndash Functional groups 27C1 Recognition of functional groups 27C2 Aliphatic and aromatic functional groups 29C3 Intermolecular bonding 30C4 Properties and reactions 33C5 Nomenclature of compounds with functional groups 35C6 Primary secondary tertiary and quaternary

nomenclature 43

Section D ndash Stereochemistry 45D1 Constitutional isomers 45D2 Configurational isomers ndash alkenes and cycloalkanes 46D3 Configurational isomers ndash optical isomers 49D4 Conformational isomers 56

Section E ndash Nucleophiles and electrophiles 63E1 Definition 63E2 Charged species 64E3 Neutral inorganic species 66E4 Organic structures 69

Section F ndash Reactions and mechanisms 73F1 Reactions 73F2 Mechanisms 75

Section G ndash Acidndashbase reactions 79G1 BroslashnstedndashLowry acids and bases 79G2 Acid strength 82G3 Base strength 88G4 Lewis acids and bases 94G5 Enolates 95

CONTENTS

Section H ndash Alkenes and alkynes 99H1 Preparation of alkenes and alkynes 99H2 Properties of alkenes and alkynes 101H3 Electrophilic addition to symmetrical alkenes 105H4 Electrophilic addition to unsymmetrical alkenes 112H5 Carbocation stabilization 115H6 Reduction and oxidation of alkenes 117H7 Hydroboration of alkenes 121H8 Electrophilic additions to alkynes 124H9 Reduction of alkynes 127H10 Alkylation of terminal alkynes 129H11 Conjugated dienes 131

Section I ndash Aromatic chemistry 135I1 Aromaticity 135I2 Preparation and properties 137I3 Electrophilic substitutions of benzene 139I4 Synthesis of mono-substituted benzenes 147I5 Electrophilic substitutions of mono-substituted

aromatic rings 150I6 Synthesis of di- and tri-substituted benzenes 160I7 Oxidation and reduction 164

Section J ndash Aldehydes and ketones 167J1 Preparation 167J2 Properties 169J3 Nucleophilic addition 173J4 Nucleophilic addition ndash charged nucleophiles 175J5 Electronic and steric effects 181J6 Nucleophilic addition ndash nitrogen nucleophiles 184J7 Nucleophilic addition ndash oxygen and sulfur

nucleophiles 187J8 Reactions of enolate ions 191J9 α-Halogenation 198J10 Reduction and oxidation 200J11 αβ-Unsaturated aldehydes and ketones 202

Section K ndash Carboxylic acids and carboxylic acid derivatives 205K1 Structure and properties 205K2 Nucleophilic substitution 209K3 Reactivity 213K4 Preparations of carboxylic acids 217K5 Preparations of carboxylic acid derivatives 219K6 Reactions 224K7 Enolate reactions 234

Section L ndash Alkyl halides 239L1 Preparation and physical properties of alkyl halides 239L2 Nucleophilic substitution 242L3 Factors affecting SN2 versus SN1 reactions 247L4 Elimination 252L5 Elimination versus substitution 256L6 Reactions of alkyl halides 258L7 Organometallic reactions 261

vi Contents

Section M ndash Alcohols phenols and thiols 263M1 Preparation of alcohols 263M2 Preparation of phenols 264M3 Properties of alcohols and phenols 266M4 Reactions of alcohols 270M5 Reactions of phenols 277M6 Chemistry of thiols 281

Section N ndash Ethers epoxides and thioethers 283N1 Preparation of ethers epoxides and thioethers 283N2 Properties of ethers epoxides and thioethers 286N3 Reactions of ethers epoxides and thioethers 289

Section O ndash Amines and nitriles 295O1 Preparation of amines 295O2 Properties of amines 299O3 Reactions of amines 305O4 Chemistry of nitriles 311

Section P ndash Organic spectroscopy and analysis 315P1 Spectroscopy 315P2 Visible and ultra violet spectroscopy 317P3 Infra-red spectroscopy 322P4 Proton nuclear magnetic resonance spectroscopy 324P5 13C nuclear magnetic resonance spectroscopy 339P6 Mass spectroscopy 342

Further reading 347Index 349

Contents vii

This textbook aims to provide a comprehensive set of basic notes in organicchemistry which will be suitable for undergraduate students taking chemistrychemistry-related courses or courses which involve organic chemistry as anancillary subject The book concentrates on core topics which are most likelyto be common to those organic chemistry courses which follow on from a foun-dation or introductory general chemistry course

Organic chemistry is a subject which can lead some students to the heightsof ecstasy yet drive others up the wall Some students lsquoswitch onrsquo to it imme-diately while others can make neither head nor tail of it no matter how hardthey try Certainly one of the major problems in studying the subject is thevast amount of material which often has to be covered Many students blancheat the prospect of having to learn a seemingly endless number of reactionsand when it comes to drawing mechanisms and curly arrows they see onlya confusing maze of squiggly lines going everywhere yet nowhere Theconcepts of organic reaction mechanisms are often the most difficult to masterThese difficulties are often compounded by the fact that current textbooks inorganic chemistry are typically over 1200 pages long and can be quite expen-sive to buy

This book attempts to condense the essentials of organic chemistry into amanageable text of 310 pages which is student friendly and which does notcost an arm and a leg It does this by concentrating purely on the basics of thesubject without going into exhaustive detail or repetitive examplesFurthermore key notes at the start of each topic summarize the essential factscovered and help focus the mind on the essentials

Organic chemistry is a peculiar subject in that it becomes easier as you goalong This might seem an outrageous statement to make especially to a first-year student who is struggling to come to terms with the rules of nomenclaturetrying to memorize a couple of dozen reactions and making sense of mecha-nisms at the same time However these topics are the basics of the subject andonce they have been grasped the overall picture becomes clear

Understanding the mechanism of how a reaction takes place is particularlycrucial in this It brings a logic to the reactions of the different functional groupsThis in turn transforms a list of apparently unrelated facts into a sensible themewhich makes remembering the reactions a lsquopiece of cakersquo (well nearly)

Once this happy state of affairs has been reached the relevance of organicchemistry to other subjects such as genetics and biochemistry suddenly leapsoff the page Understanding organic chemistry leads to a better understandingof life chemistry and how the body works at the molecular level It also helpsin the understanding of the molecular mechanisms involved in disease andbodily malfunction leading in turn to an understanding of how drugs can bedesigned to cure these disease states ndash the science of medicinal chemistry

And thatrsquos not all An understanding of organic chemistry will help the indus-trial chemist or chemical engineer faced with unexpected side-reactions in achemical process and the agro-scientist trying to understand the molecularprocesses taking place within plants and crops and it will assist in the designand synthesis of new herbicides and fungicides which will be eco-friendly It

PREFACE

will aid the forensic scientist wishing to analyze a nondescript white powderndash is it heroin or flour

The list of scientific subject areas involving organic chemistry is endless ndashdesigning spacesuits developing new photographic dyes inventing new mole-cular technology in microelectronics ndash one could go on and on Organicchemistry is an exciting subject since it leads to an essential understanding ofmolecules and their properties

The order in which the early topics of this book are presented is importantThe first two sections cover structure and bonding which are crucial to latersections Just why does carbon form four bonds What is hybridization

The third section on functional groups is equally crucial if students are to becapable of categorizing the apparent maze of reactions which organiccompounds can undergo It is followed by section D on stereochemistry thensections E and F in which the basic theory of reactions and mechanisms iscovered What are nucleophiles and electrophiles What does a mechanismrepresent What does a curly arrow mean

The remaining sections can be used in any order These look at the reactionsand mechanisms of the common functional groups which are important inchemistry and biochemistry

It is hoped that students will find this textbook useful in their studies andthat once they have grasped what organic chemistry is all about they will readmore widely and enter a truly exciting world of molecular science

x Preface

Section A ndash Structure and bonding

A1 ATOMIC STRUCTURE OF CARBON

Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table This means that thereare two shells of atomic orbitals available for these electrons The first shell closestto the nucleus has a single s orbital ndash the 1s orbital The second shell has a single sorbital (the 2s orbital) and three p orbitals (3 2p) Therefore there are a total offive atomic orbitals into which these six electrons can fit The s orbitals are sphericalin shape with the 2s orbital being much larger then the 1s orbital The p orbitalsare dumbbell-shaped and are aligned along the x y and z axes Therefore they areassigned the 2px 2py and 2pz atomic orbitals (Fig 1)

Key Notes

The atomic orbitals available for the six electrons of carbon are the s orbitalin the first shell the s orbital in the second shell and the three p orbitals inthe second shell The 1s and 2s orbitals are spherical in shape The 2porbitals are dumbbell in shape and can be assigned 2px 2py or 2pz depend-ing on the axis along which they are aligned

The 1s orbital has a lower energy than the 2s orbital which has a lowerenergy than the 2p orbitals The 2p orbitals have equal energy (ie theyare degenerate)

Carbon is in the second row of the periodic table and has six electrons whichwill fill up lower energy atomic orbitals before entering higher energyorbitals (aufbau principle) Each orbital is allowed a maximum of two elec-trons of opposite spin (Pauli exclusion principle) When orbitals of equalenergy are available electrons will occupy separate orbitals before pairingup (Hundrsquos rule) Thus the electronic configuration of a carbon atom is 1s2

2s2 2px1 2py

1

Related topic Covalent bonding andhybridization (A2)

Electronicconfiguration

Atomic orbitals

Energy levels

1s 2s 2px 2py 2pz

y

xz

y

xz

y

xz

y

xz

y

xz

Fig 1 Atomic orbitals

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

copy Garland ScienceBIOS Scientific Publishers 2004

First published 2000Second edition published 2004

All rights reserved No part of this book may be reproduced or transmitted in any form or by any meanswithout permission

A CIP catalogue record for this book is available from the British Library

ISBN 1 85996 264 5

Garland ScienceBIOS Scientific Publishers4 Park Square Milton Park Abingdon Oxon OX14 4RN UK and29 West 35th Street New York NY 10001ndash2299 USAWorld Wide Web home page wwwbioscouk

Garland ScienceBIOS Scientific Publishers is a member of the Taylor amp Francis Group

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Library of Congress Cataloging-in-Publication Data

Patrick Graham LOrganic chemistry GL Patrickmdash2nd ed

p cm mdash (Instant notes series)Includes bibliographical references and indexISBN 1-85996-264-5 (alk paper)1 Chemistry OrganicmdashOutlines syllabi etc I Title Instant notes organic chemistry

II Title III Series

QD2565P37 2003547mdashdc22

Production Editor Andrew Watts

This edition published in the Taylor amp Francis e-Library 2005

To purchase your own copy of this or any of Taylor amp Francis or Routledges

collection of thousands of eBooks please go to wwweBookstore

ISBN 0-203-42761-0 Master e-book ISBN

ISBN 0-203-44168-0 (Adobe eReader Format)(Print Edition)

tandfcoukrdquo

Preface ix

Section A ndash Structure and bonding 1A1 Atomic structure of carbon 1A2 Covalent bonding and hybridization 3A3 sp3 Hybridization 5A4 sp2 Hybridization 8A5 sp Hybridization 14A6 Bonds and hybridized centers 17

Section B ndash Alkanes and cycloalkanes 19B1 Definition 19B2 Drawing structures 20B3 Nomenclature 22

Section C ndash Functional groups 27C1 Recognition of functional groups 27C2 Aliphatic and aromatic functional groups 29C3 Intermolecular bonding 30C4 Properties and reactions 33C5 Nomenclature of compounds with functional groups 35C6 Primary secondary tertiary and quaternary

nomenclature 43

Section D ndash Stereochemistry 45D1 Constitutional isomers 45D2 Configurational isomers ndash alkenes and cycloalkanes 46D3 Configurational isomers ndash optical isomers 49D4 Conformational isomers 56

Section E ndash Nucleophiles and electrophiles 63E1 Definition 63E2 Charged species 64E3 Neutral inorganic species 66E4 Organic structures 69

Section F ndash Reactions and mechanisms 73F1 Reactions 73F2 Mechanisms 75

Section G ndash Acidndashbase reactions 79G1 BroslashnstedndashLowry acids and bases 79G2 Acid strength 82G3 Base strength 88G4 Lewis acids and bases 94G5 Enolates 95

CONTENTS

Section H ndash Alkenes and alkynes 99H1 Preparation of alkenes and alkynes 99H2 Properties of alkenes and alkynes 101H3 Electrophilic addition to symmetrical alkenes 105H4 Electrophilic addition to unsymmetrical alkenes 112H5 Carbocation stabilization 115H6 Reduction and oxidation of alkenes 117H7 Hydroboration of alkenes 121H8 Electrophilic additions to alkynes 124H9 Reduction of alkynes 127H10 Alkylation of terminal alkynes 129H11 Conjugated dienes 131

Section I ndash Aromatic chemistry 135I1 Aromaticity 135I2 Preparation and properties 137I3 Electrophilic substitutions of benzene 139I4 Synthesis of mono-substituted benzenes 147I5 Electrophilic substitutions of mono-substituted

aromatic rings 150I6 Synthesis of di- and tri-substituted benzenes 160I7 Oxidation and reduction 164

Section J ndash Aldehydes and ketones 167J1 Preparation 167J2 Properties 169J3 Nucleophilic addition 173J4 Nucleophilic addition ndash charged nucleophiles 175J5 Electronic and steric effects 181J6 Nucleophilic addition ndash nitrogen nucleophiles 184J7 Nucleophilic addition ndash oxygen and sulfur

nucleophiles 187J8 Reactions of enolate ions 191J9 α-Halogenation 198J10 Reduction and oxidation 200J11 αβ-Unsaturated aldehydes and ketones 202

Section K ndash Carboxylic acids and carboxylic acid derivatives 205K1 Structure and properties 205K2 Nucleophilic substitution 209K3 Reactivity 213K4 Preparations of carboxylic acids 217K5 Preparations of carboxylic acid derivatives 219K6 Reactions 224K7 Enolate reactions 234

Section L ndash Alkyl halides 239L1 Preparation and physical properties of alkyl halides 239L2 Nucleophilic substitution 242L3 Factors affecting SN2 versus SN1 reactions 247L4 Elimination 252L5 Elimination versus substitution 256L6 Reactions of alkyl halides 258L7 Organometallic reactions 261

vi Contents

Section M ndash Alcohols phenols and thiols 263M1 Preparation of alcohols 263M2 Preparation of phenols 264M3 Properties of alcohols and phenols 266M4 Reactions of alcohols 270M5 Reactions of phenols 277M6 Chemistry of thiols 281

Section N ndash Ethers epoxides and thioethers 283N1 Preparation of ethers epoxides and thioethers 283N2 Properties of ethers epoxides and thioethers 286N3 Reactions of ethers epoxides and thioethers 289

Section O ndash Amines and nitriles 295O1 Preparation of amines 295O2 Properties of amines 299O3 Reactions of amines 305O4 Chemistry of nitriles 311

Section P ndash Organic spectroscopy and analysis 315P1 Spectroscopy 315P2 Visible and ultra violet spectroscopy 317P3 Infra-red spectroscopy 322P4 Proton nuclear magnetic resonance spectroscopy 324P5 13C nuclear magnetic resonance spectroscopy 339P6 Mass spectroscopy 342

Further reading 347Index 349

Contents vii

This textbook aims to provide a comprehensive set of basic notes in organicchemistry which will be suitable for undergraduate students taking chemistrychemistry-related courses or courses which involve organic chemistry as anancillary subject The book concentrates on core topics which are most likelyto be common to those organic chemistry courses which follow on from a foun-dation or introductory general chemistry course

Organic chemistry is a subject which can lead some students to the heightsof ecstasy yet drive others up the wall Some students lsquoswitch onrsquo to it imme-diately while others can make neither head nor tail of it no matter how hardthey try Certainly one of the major problems in studying the subject is thevast amount of material which often has to be covered Many students blancheat the prospect of having to learn a seemingly endless number of reactionsand when it comes to drawing mechanisms and curly arrows they see onlya confusing maze of squiggly lines going everywhere yet nowhere Theconcepts of organic reaction mechanisms are often the most difficult to masterThese difficulties are often compounded by the fact that current textbooks inorganic chemistry are typically over 1200 pages long and can be quite expen-sive to buy

This book attempts to condense the essentials of organic chemistry into amanageable text of 310 pages which is student friendly and which does notcost an arm and a leg It does this by concentrating purely on the basics of thesubject without going into exhaustive detail or repetitive examplesFurthermore key notes at the start of each topic summarize the essential factscovered and help focus the mind on the essentials

Organic chemistry is a peculiar subject in that it becomes easier as you goalong This might seem an outrageous statement to make especially to a first-year student who is struggling to come to terms with the rules of nomenclaturetrying to memorize a couple of dozen reactions and making sense of mecha-nisms at the same time However these topics are the basics of the subject andonce they have been grasped the overall picture becomes clear

Understanding the mechanism of how a reaction takes place is particularlycrucial in this It brings a logic to the reactions of the different functional groupsThis in turn transforms a list of apparently unrelated facts into a sensible themewhich makes remembering the reactions a lsquopiece of cakersquo (well nearly)

Once this happy state of affairs has been reached the relevance of organicchemistry to other subjects such as genetics and biochemistry suddenly leapsoff the page Understanding organic chemistry leads to a better understandingof life chemistry and how the body works at the molecular level It also helpsin the understanding of the molecular mechanisms involved in disease andbodily malfunction leading in turn to an understanding of how drugs can bedesigned to cure these disease states ndash the science of medicinal chemistry

And thatrsquos not all An understanding of organic chemistry will help the indus-trial chemist or chemical engineer faced with unexpected side-reactions in achemical process and the agro-scientist trying to understand the molecularprocesses taking place within plants and crops and it will assist in the designand synthesis of new herbicides and fungicides which will be eco-friendly It

PREFACE

will aid the forensic scientist wishing to analyze a nondescript white powderndash is it heroin or flour

The list of scientific subject areas involving organic chemistry is endless ndashdesigning spacesuits developing new photographic dyes inventing new mole-cular technology in microelectronics ndash one could go on and on Organicchemistry is an exciting subject since it leads to an essential understanding ofmolecules and their properties

The order in which the early topics of this book are presented is importantThe first two sections cover structure and bonding which are crucial to latersections Just why does carbon form four bonds What is hybridization

The third section on functional groups is equally crucial if students are to becapable of categorizing the apparent maze of reactions which organiccompounds can undergo It is followed by section D on stereochemistry thensections E and F in which the basic theory of reactions and mechanisms iscovered What are nucleophiles and electrophiles What does a mechanismrepresent What does a curly arrow mean

The remaining sections can be used in any order These look at the reactionsand mechanisms of the common functional groups which are important inchemistry and biochemistry

It is hoped that students will find this textbook useful in their studies andthat once they have grasped what organic chemistry is all about they will readmore widely and enter a truly exciting world of molecular science

x Preface

Section A ndash Structure and bonding

A1 ATOMIC STRUCTURE OF CARBON

Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table This means that thereare two shells of atomic orbitals available for these electrons The first shell closestto the nucleus has a single s orbital ndash the 1s orbital The second shell has a single sorbital (the 2s orbital) and three p orbitals (3 2p) Therefore there are a total offive atomic orbitals into which these six electrons can fit The s orbitals are sphericalin shape with the 2s orbital being much larger then the 1s orbital The p orbitalsare dumbbell-shaped and are aligned along the x y and z axes Therefore they areassigned the 2px 2py and 2pz atomic orbitals (Fig 1)

Key Notes

The atomic orbitals available for the six electrons of carbon are the s orbitalin the first shell the s orbital in the second shell and the three p orbitals inthe second shell The 1s and 2s orbitals are spherical in shape The 2porbitals are dumbbell in shape and can be assigned 2px 2py or 2pz depend-ing on the axis along which they are aligned

The 1s orbital has a lower energy than the 2s orbital which has a lowerenergy than the 2p orbitals The 2p orbitals have equal energy (ie theyare degenerate)

Carbon is in the second row of the periodic table and has six electrons whichwill fill up lower energy atomic orbitals before entering higher energyorbitals (aufbau principle) Each orbital is allowed a maximum of two elec-trons of opposite spin (Pauli exclusion principle) When orbitals of equalenergy are available electrons will occupy separate orbitals before pairingup (Hundrsquos rule) Thus the electronic configuration of a carbon atom is 1s2

2s2 2px1 2py

1

Related topic Covalent bonding andhybridization (A2)

Electronicconfiguration

Atomic orbitals

Energy levels

1s 2s 2px 2py 2pz

y

xz

y

xz

y

xz

y

xz

y

xz

Fig 1 Atomic orbitals

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Preface ix

Section A ndash Structure and bonding 1A1 Atomic structure of carbon 1A2 Covalent bonding and hybridization 3A3 sp3 Hybridization 5A4 sp2 Hybridization 8A5 sp Hybridization 14A6 Bonds and hybridized centers 17

Section B ndash Alkanes and cycloalkanes 19B1 Definition 19B2 Drawing structures 20B3 Nomenclature 22

Section C ndash Functional groups 27C1 Recognition of functional groups 27C2 Aliphatic and aromatic functional groups 29C3 Intermolecular bonding 30C4 Properties and reactions 33C5 Nomenclature of compounds with functional groups 35C6 Primary secondary tertiary and quaternary

nomenclature 43

Section D ndash Stereochemistry 45D1 Constitutional isomers 45D2 Configurational isomers ndash alkenes and cycloalkanes 46D3 Configurational isomers ndash optical isomers 49D4 Conformational isomers 56

Section E ndash Nucleophiles and electrophiles 63E1 Definition 63E2 Charged species 64E3 Neutral inorganic species 66E4 Organic structures 69

Section F ndash Reactions and mechanisms 73F1 Reactions 73F2 Mechanisms 75

Section G ndash Acidndashbase reactions 79G1 BroslashnstedndashLowry acids and bases 79G2 Acid strength 82G3 Base strength 88G4 Lewis acids and bases 94G5 Enolates 95

CONTENTS

Section H ndash Alkenes and alkynes 99H1 Preparation of alkenes and alkynes 99H2 Properties of alkenes and alkynes 101H3 Electrophilic addition to symmetrical alkenes 105H4 Electrophilic addition to unsymmetrical alkenes 112H5 Carbocation stabilization 115H6 Reduction and oxidation of alkenes 117H7 Hydroboration of alkenes 121H8 Electrophilic additions to alkynes 124H9 Reduction of alkynes 127H10 Alkylation of terminal alkynes 129H11 Conjugated dienes 131

Section I ndash Aromatic chemistry 135I1 Aromaticity 135I2 Preparation and properties 137I3 Electrophilic substitutions of benzene 139I4 Synthesis of mono-substituted benzenes 147I5 Electrophilic substitutions of mono-substituted

aromatic rings 150I6 Synthesis of di- and tri-substituted benzenes 160I7 Oxidation and reduction 164

Section J ndash Aldehydes and ketones 167J1 Preparation 167J2 Properties 169J3 Nucleophilic addition 173J4 Nucleophilic addition ndash charged nucleophiles 175J5 Electronic and steric effects 181J6 Nucleophilic addition ndash nitrogen nucleophiles 184J7 Nucleophilic addition ndash oxygen and sulfur

nucleophiles 187J8 Reactions of enolate ions 191J9 α-Halogenation 198J10 Reduction and oxidation 200J11 αβ-Unsaturated aldehydes and ketones 202

Section K ndash Carboxylic acids and carboxylic acid derivatives 205K1 Structure and properties 205K2 Nucleophilic substitution 209K3 Reactivity 213K4 Preparations of carboxylic acids 217K5 Preparations of carboxylic acid derivatives 219K6 Reactions 224K7 Enolate reactions 234

Section L ndash Alkyl halides 239L1 Preparation and physical properties of alkyl halides 239L2 Nucleophilic substitution 242L3 Factors affecting SN2 versus SN1 reactions 247L4 Elimination 252L5 Elimination versus substitution 256L6 Reactions of alkyl halides 258L7 Organometallic reactions 261

vi Contents

Section M ndash Alcohols phenols and thiols 263M1 Preparation of alcohols 263M2 Preparation of phenols 264M3 Properties of alcohols and phenols 266M4 Reactions of alcohols 270M5 Reactions of phenols 277M6 Chemistry of thiols 281

Section N ndash Ethers epoxides and thioethers 283N1 Preparation of ethers epoxides and thioethers 283N2 Properties of ethers epoxides and thioethers 286N3 Reactions of ethers epoxides and thioethers 289

Section O ndash Amines and nitriles 295O1 Preparation of amines 295O2 Properties of amines 299O3 Reactions of amines 305O4 Chemistry of nitriles 311

Section P ndash Organic spectroscopy and analysis 315P1 Spectroscopy 315P2 Visible and ultra violet spectroscopy 317P3 Infra-red spectroscopy 322P4 Proton nuclear magnetic resonance spectroscopy 324P5 13C nuclear magnetic resonance spectroscopy 339P6 Mass spectroscopy 342

Further reading 347Index 349

Contents vii

This textbook aims to provide a comprehensive set of basic notes in organicchemistry which will be suitable for undergraduate students taking chemistrychemistry-related courses or courses which involve organic chemistry as anancillary subject The book concentrates on core topics which are most likelyto be common to those organic chemistry courses which follow on from a foun-dation or introductory general chemistry course

Organic chemistry is a subject which can lead some students to the heightsof ecstasy yet drive others up the wall Some students lsquoswitch onrsquo to it imme-diately while others can make neither head nor tail of it no matter how hardthey try Certainly one of the major problems in studying the subject is thevast amount of material which often has to be covered Many students blancheat the prospect of having to learn a seemingly endless number of reactionsand when it comes to drawing mechanisms and curly arrows they see onlya confusing maze of squiggly lines going everywhere yet nowhere Theconcepts of organic reaction mechanisms are often the most difficult to masterThese difficulties are often compounded by the fact that current textbooks inorganic chemistry are typically over 1200 pages long and can be quite expen-sive to buy

This book attempts to condense the essentials of organic chemistry into amanageable text of 310 pages which is student friendly and which does notcost an arm and a leg It does this by concentrating purely on the basics of thesubject without going into exhaustive detail or repetitive examplesFurthermore key notes at the start of each topic summarize the essential factscovered and help focus the mind on the essentials

Organic chemistry is a peculiar subject in that it becomes easier as you goalong This might seem an outrageous statement to make especially to a first-year student who is struggling to come to terms with the rules of nomenclaturetrying to memorize a couple of dozen reactions and making sense of mecha-nisms at the same time However these topics are the basics of the subject andonce they have been grasped the overall picture becomes clear

Understanding the mechanism of how a reaction takes place is particularlycrucial in this It brings a logic to the reactions of the different functional groupsThis in turn transforms a list of apparently unrelated facts into a sensible themewhich makes remembering the reactions a lsquopiece of cakersquo (well nearly)

Once this happy state of affairs has been reached the relevance of organicchemistry to other subjects such as genetics and biochemistry suddenly leapsoff the page Understanding organic chemistry leads to a better understandingof life chemistry and how the body works at the molecular level It also helpsin the understanding of the molecular mechanisms involved in disease andbodily malfunction leading in turn to an understanding of how drugs can bedesigned to cure these disease states ndash the science of medicinal chemistry

And thatrsquos not all An understanding of organic chemistry will help the indus-trial chemist or chemical engineer faced with unexpected side-reactions in achemical process and the agro-scientist trying to understand the molecularprocesses taking place within plants and crops and it will assist in the designand synthesis of new herbicides and fungicides which will be eco-friendly It

PREFACE

will aid the forensic scientist wishing to analyze a nondescript white powderndash is it heroin or flour

The list of scientific subject areas involving organic chemistry is endless ndashdesigning spacesuits developing new photographic dyes inventing new mole-cular technology in microelectronics ndash one could go on and on Organicchemistry is an exciting subject since it leads to an essential understanding ofmolecules and their properties

The order in which the early topics of this book are presented is importantThe first two sections cover structure and bonding which are crucial to latersections Just why does carbon form four bonds What is hybridization

The third section on functional groups is equally crucial if students are to becapable of categorizing the apparent maze of reactions which organiccompounds can undergo It is followed by section D on stereochemistry thensections E and F in which the basic theory of reactions and mechanisms iscovered What are nucleophiles and electrophiles What does a mechanismrepresent What does a curly arrow mean

The remaining sections can be used in any order These look at the reactionsand mechanisms of the common functional groups which are important inchemistry and biochemistry

It is hoped that students will find this textbook useful in their studies andthat once they have grasped what organic chemistry is all about they will readmore widely and enter a truly exciting world of molecular science

x Preface

Section A ndash Structure and bonding

A1 ATOMIC STRUCTURE OF CARBON

Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table This means that thereare two shells of atomic orbitals available for these electrons The first shell closestto the nucleus has a single s orbital ndash the 1s orbital The second shell has a single sorbital (the 2s orbital) and three p orbitals (3 2p) Therefore there are a total offive atomic orbitals into which these six electrons can fit The s orbitals are sphericalin shape with the 2s orbital being much larger then the 1s orbital The p orbitalsare dumbbell-shaped and are aligned along the x y and z axes Therefore they areassigned the 2px 2py and 2pz atomic orbitals (Fig 1)

Key Notes

The atomic orbitals available for the six electrons of carbon are the s orbitalin the first shell the s orbital in the second shell and the three p orbitals inthe second shell The 1s and 2s orbitals are spherical in shape The 2porbitals are dumbbell in shape and can be assigned 2px 2py or 2pz depend-ing on the axis along which they are aligned

The 1s orbital has a lower energy than the 2s orbital which has a lowerenergy than the 2p orbitals The 2p orbitals have equal energy (ie theyare degenerate)

Carbon is in the second row of the periodic table and has six electrons whichwill fill up lower energy atomic orbitals before entering higher energyorbitals (aufbau principle) Each orbital is allowed a maximum of two elec-trons of opposite spin (Pauli exclusion principle) When orbitals of equalenergy are available electrons will occupy separate orbitals before pairingup (Hundrsquos rule) Thus the electronic configuration of a carbon atom is 1s2

2s2 2px1 2py

1

Related topic Covalent bonding andhybridization (A2)

Electronicconfiguration

Atomic orbitals

Energy levels

1s 2s 2px 2py 2pz

y

xz

y

xz

y

xz

y

xz

y

xz

Fig 1 Atomic orbitals

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Section H ndash Alkenes and alkynes 99H1 Preparation of alkenes and alkynes 99H2 Properties of alkenes and alkynes 101H3 Electrophilic addition to symmetrical alkenes 105H4 Electrophilic addition to unsymmetrical alkenes 112H5 Carbocation stabilization 115H6 Reduction and oxidation of alkenes 117H7 Hydroboration of alkenes 121H8 Electrophilic additions to alkynes 124H9 Reduction of alkynes 127H10 Alkylation of terminal alkynes 129H11 Conjugated dienes 131

Section I ndash Aromatic chemistry 135I1 Aromaticity 135I2 Preparation and properties 137I3 Electrophilic substitutions of benzene 139I4 Synthesis of mono-substituted benzenes 147I5 Electrophilic substitutions of mono-substituted

aromatic rings 150I6 Synthesis of di- and tri-substituted benzenes 160I7 Oxidation and reduction 164

Section J ndash Aldehydes and ketones 167J1 Preparation 167J2 Properties 169J3 Nucleophilic addition 173J4 Nucleophilic addition ndash charged nucleophiles 175J5 Electronic and steric effects 181J6 Nucleophilic addition ndash nitrogen nucleophiles 184J7 Nucleophilic addition ndash oxygen and sulfur

nucleophiles 187J8 Reactions of enolate ions 191J9 α-Halogenation 198J10 Reduction and oxidation 200J11 αβ-Unsaturated aldehydes and ketones 202

Section K ndash Carboxylic acids and carboxylic acid derivatives 205K1 Structure and properties 205K2 Nucleophilic substitution 209K3 Reactivity 213K4 Preparations of carboxylic acids 217K5 Preparations of carboxylic acid derivatives 219K6 Reactions 224K7 Enolate reactions 234

Section L ndash Alkyl halides 239L1 Preparation and physical properties of alkyl halides 239L2 Nucleophilic substitution 242L3 Factors affecting SN2 versus SN1 reactions 247L4 Elimination 252L5 Elimination versus substitution 256L6 Reactions of alkyl halides 258L7 Organometallic reactions 261

vi Contents

Section M ndash Alcohols phenols and thiols 263M1 Preparation of alcohols 263M2 Preparation of phenols 264M3 Properties of alcohols and phenols 266M4 Reactions of alcohols 270M5 Reactions of phenols 277M6 Chemistry of thiols 281

Section N ndash Ethers epoxides and thioethers 283N1 Preparation of ethers epoxides and thioethers 283N2 Properties of ethers epoxides and thioethers 286N3 Reactions of ethers epoxides and thioethers 289

Section O ndash Amines and nitriles 295O1 Preparation of amines 295O2 Properties of amines 299O3 Reactions of amines 305O4 Chemistry of nitriles 311

Section P ndash Organic spectroscopy and analysis 315P1 Spectroscopy 315P2 Visible and ultra violet spectroscopy 317P3 Infra-red spectroscopy 322P4 Proton nuclear magnetic resonance spectroscopy 324P5 13C nuclear magnetic resonance spectroscopy 339P6 Mass spectroscopy 342

Further reading 347Index 349

Contents vii

This textbook aims to provide a comprehensive set of basic notes in organicchemistry which will be suitable for undergraduate students taking chemistrychemistry-related courses or courses which involve organic chemistry as anancillary subject The book concentrates on core topics which are most likelyto be common to those organic chemistry courses which follow on from a foun-dation or introductory general chemistry course

Organic chemistry is a subject which can lead some students to the heightsof ecstasy yet drive others up the wall Some students lsquoswitch onrsquo to it imme-diately while others can make neither head nor tail of it no matter how hardthey try Certainly one of the major problems in studying the subject is thevast amount of material which often has to be covered Many students blancheat the prospect of having to learn a seemingly endless number of reactionsand when it comes to drawing mechanisms and curly arrows they see onlya confusing maze of squiggly lines going everywhere yet nowhere Theconcepts of organic reaction mechanisms are often the most difficult to masterThese difficulties are often compounded by the fact that current textbooks inorganic chemistry are typically over 1200 pages long and can be quite expen-sive to buy

This book attempts to condense the essentials of organic chemistry into amanageable text of 310 pages which is student friendly and which does notcost an arm and a leg It does this by concentrating purely on the basics of thesubject without going into exhaustive detail or repetitive examplesFurthermore key notes at the start of each topic summarize the essential factscovered and help focus the mind on the essentials

Organic chemistry is a peculiar subject in that it becomes easier as you goalong This might seem an outrageous statement to make especially to a first-year student who is struggling to come to terms with the rules of nomenclaturetrying to memorize a couple of dozen reactions and making sense of mecha-nisms at the same time However these topics are the basics of the subject andonce they have been grasped the overall picture becomes clear

Understanding the mechanism of how a reaction takes place is particularlycrucial in this It brings a logic to the reactions of the different functional groupsThis in turn transforms a list of apparently unrelated facts into a sensible themewhich makes remembering the reactions a lsquopiece of cakersquo (well nearly)

Once this happy state of affairs has been reached the relevance of organicchemistry to other subjects such as genetics and biochemistry suddenly leapsoff the page Understanding organic chemistry leads to a better understandingof life chemistry and how the body works at the molecular level It also helpsin the understanding of the molecular mechanisms involved in disease andbodily malfunction leading in turn to an understanding of how drugs can bedesigned to cure these disease states ndash the science of medicinal chemistry

And thatrsquos not all An understanding of organic chemistry will help the indus-trial chemist or chemical engineer faced with unexpected side-reactions in achemical process and the agro-scientist trying to understand the molecularprocesses taking place within plants and crops and it will assist in the designand synthesis of new herbicides and fungicides which will be eco-friendly It

PREFACE

will aid the forensic scientist wishing to analyze a nondescript white powderndash is it heroin or flour

The list of scientific subject areas involving organic chemistry is endless ndashdesigning spacesuits developing new photographic dyes inventing new mole-cular technology in microelectronics ndash one could go on and on Organicchemistry is an exciting subject since it leads to an essential understanding ofmolecules and their properties

The order in which the early topics of this book are presented is importantThe first two sections cover structure and bonding which are crucial to latersections Just why does carbon form four bonds What is hybridization

The third section on functional groups is equally crucial if students are to becapable of categorizing the apparent maze of reactions which organiccompounds can undergo It is followed by section D on stereochemistry thensections E and F in which the basic theory of reactions and mechanisms iscovered What are nucleophiles and electrophiles What does a mechanismrepresent What does a curly arrow mean

The remaining sections can be used in any order These look at the reactionsand mechanisms of the common functional groups which are important inchemistry and biochemistry

It is hoped that students will find this textbook useful in their studies andthat once they have grasped what organic chemistry is all about they will readmore widely and enter a truly exciting world of molecular science

x Preface

Section A ndash Structure and bonding

A1 ATOMIC STRUCTURE OF CARBON

Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table This means that thereare two shells of atomic orbitals available for these electrons The first shell closestto the nucleus has a single s orbital ndash the 1s orbital The second shell has a single sorbital (the 2s orbital) and three p orbitals (3 2p) Therefore there are a total offive atomic orbitals into which these six electrons can fit The s orbitals are sphericalin shape with the 2s orbital being much larger then the 1s orbital The p orbitalsare dumbbell-shaped and are aligned along the x y and z axes Therefore they areassigned the 2px 2py and 2pz atomic orbitals (Fig 1)

Key Notes

The atomic orbitals available for the six electrons of carbon are the s orbitalin the first shell the s orbital in the second shell and the three p orbitals inthe second shell The 1s and 2s orbitals are spherical in shape The 2porbitals are dumbbell in shape and can be assigned 2px 2py or 2pz depend-ing on the axis along which they are aligned

The 1s orbital has a lower energy than the 2s orbital which has a lowerenergy than the 2p orbitals The 2p orbitals have equal energy (ie theyare degenerate)

Carbon is in the second row of the periodic table and has six electrons whichwill fill up lower energy atomic orbitals before entering higher energyorbitals (aufbau principle) Each orbital is allowed a maximum of two elec-trons of opposite spin (Pauli exclusion principle) When orbitals of equalenergy are available electrons will occupy separate orbitals before pairingup (Hundrsquos rule) Thus the electronic configuration of a carbon atom is 1s2

2s2 2px1 2py

1

Related topic Covalent bonding andhybridization (A2)

Electronicconfiguration

Atomic orbitals

Energy levels

1s 2s 2px 2py 2pz

y

xz

y

xz

y

xz

y

xz

y

xz

Fig 1 Atomic orbitals

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Section M ndash Alcohols phenols and thiols 263M1 Preparation of alcohols 263M2 Preparation of phenols 264M3 Properties of alcohols and phenols 266M4 Reactions of alcohols 270M5 Reactions of phenols 277M6 Chemistry of thiols 281

Section N ndash Ethers epoxides and thioethers 283N1 Preparation of ethers epoxides and thioethers 283N2 Properties of ethers epoxides and thioethers 286N3 Reactions of ethers epoxides and thioethers 289

Section O ndash Amines and nitriles 295O1 Preparation of amines 295O2 Properties of amines 299O3 Reactions of amines 305O4 Chemistry of nitriles 311

Section P ndash Organic spectroscopy and analysis 315P1 Spectroscopy 315P2 Visible and ultra violet spectroscopy 317P3 Infra-red spectroscopy 322P4 Proton nuclear magnetic resonance spectroscopy 324P5 13C nuclear magnetic resonance spectroscopy 339P6 Mass spectroscopy 342

Further reading 347Index 349

Contents vii

This textbook aims to provide a comprehensive set of basic notes in organicchemistry which will be suitable for undergraduate students taking chemistrychemistry-related courses or courses which involve organic chemistry as anancillary subject The book concentrates on core topics which are most likelyto be common to those organic chemistry courses which follow on from a foun-dation or introductory general chemistry course

Organic chemistry is a subject which can lead some students to the heightsof ecstasy yet drive others up the wall Some students lsquoswitch onrsquo to it imme-diately while others can make neither head nor tail of it no matter how hardthey try Certainly one of the major problems in studying the subject is thevast amount of material which often has to be covered Many students blancheat the prospect of having to learn a seemingly endless number of reactionsand when it comes to drawing mechanisms and curly arrows they see onlya confusing maze of squiggly lines going everywhere yet nowhere Theconcepts of organic reaction mechanisms are often the most difficult to masterThese difficulties are often compounded by the fact that current textbooks inorganic chemistry are typically over 1200 pages long and can be quite expen-sive to buy

This book attempts to condense the essentials of organic chemistry into amanageable text of 310 pages which is student friendly and which does notcost an arm and a leg It does this by concentrating purely on the basics of thesubject without going into exhaustive detail or repetitive examplesFurthermore key notes at the start of each topic summarize the essential factscovered and help focus the mind on the essentials

Organic chemistry is a peculiar subject in that it becomes easier as you goalong This might seem an outrageous statement to make especially to a first-year student who is struggling to come to terms with the rules of nomenclaturetrying to memorize a couple of dozen reactions and making sense of mecha-nisms at the same time However these topics are the basics of the subject andonce they have been grasped the overall picture becomes clear

Understanding the mechanism of how a reaction takes place is particularlycrucial in this It brings a logic to the reactions of the different functional groupsThis in turn transforms a list of apparently unrelated facts into a sensible themewhich makes remembering the reactions a lsquopiece of cakersquo (well nearly)

Once this happy state of affairs has been reached the relevance of organicchemistry to other subjects such as genetics and biochemistry suddenly leapsoff the page Understanding organic chemistry leads to a better understandingof life chemistry and how the body works at the molecular level It also helpsin the understanding of the molecular mechanisms involved in disease andbodily malfunction leading in turn to an understanding of how drugs can bedesigned to cure these disease states ndash the science of medicinal chemistry

And thatrsquos not all An understanding of organic chemistry will help the indus-trial chemist or chemical engineer faced with unexpected side-reactions in achemical process and the agro-scientist trying to understand the molecularprocesses taking place within plants and crops and it will assist in the designand synthesis of new herbicides and fungicides which will be eco-friendly It

PREFACE

will aid the forensic scientist wishing to analyze a nondescript white powderndash is it heroin or flour

The list of scientific subject areas involving organic chemistry is endless ndashdesigning spacesuits developing new photographic dyes inventing new mole-cular technology in microelectronics ndash one could go on and on Organicchemistry is an exciting subject since it leads to an essential understanding ofmolecules and their properties

The order in which the early topics of this book are presented is importantThe first two sections cover structure and bonding which are crucial to latersections Just why does carbon form four bonds What is hybridization

The third section on functional groups is equally crucial if students are to becapable of categorizing the apparent maze of reactions which organiccompounds can undergo It is followed by section D on stereochemistry thensections E and F in which the basic theory of reactions and mechanisms iscovered What are nucleophiles and electrophiles What does a mechanismrepresent What does a curly arrow mean

The remaining sections can be used in any order These look at the reactionsand mechanisms of the common functional groups which are important inchemistry and biochemistry

It is hoped that students will find this textbook useful in their studies andthat once they have grasped what organic chemistry is all about they will readmore widely and enter a truly exciting world of molecular science

x Preface

Section A ndash Structure and bonding

A1 ATOMIC STRUCTURE OF CARBON

Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table This means that thereare two shells of atomic orbitals available for these electrons The first shell closestto the nucleus has a single s orbital ndash the 1s orbital The second shell has a single sorbital (the 2s orbital) and three p orbitals (3 2p) Therefore there are a total offive atomic orbitals into which these six electrons can fit The s orbitals are sphericalin shape with the 2s orbital being much larger then the 1s orbital The p orbitalsare dumbbell-shaped and are aligned along the x y and z axes Therefore they areassigned the 2px 2py and 2pz atomic orbitals (Fig 1)

Key Notes

The atomic orbitals available for the six electrons of carbon are the s orbitalin the first shell the s orbital in the second shell and the three p orbitals inthe second shell The 1s and 2s orbitals are spherical in shape The 2porbitals are dumbbell in shape and can be assigned 2px 2py or 2pz depend-ing on the axis along which they are aligned

The 1s orbital has a lower energy than the 2s orbital which has a lowerenergy than the 2p orbitals The 2p orbitals have equal energy (ie theyare degenerate)

Carbon is in the second row of the periodic table and has six electrons whichwill fill up lower energy atomic orbitals before entering higher energyorbitals (aufbau principle) Each orbital is allowed a maximum of two elec-trons of opposite spin (Pauli exclusion principle) When orbitals of equalenergy are available electrons will occupy separate orbitals before pairingup (Hundrsquos rule) Thus the electronic configuration of a carbon atom is 1s2

2s2 2px1 2py

1

Related topic Covalent bonding andhybridization (A2)

Electronicconfiguration

Atomic orbitals

Energy levels

1s 2s 2px 2py 2pz

y

xz

y

xz

y

xz

y

xz

y

xz

Fig 1 Atomic orbitals

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

This textbook aims to provide a comprehensive set of basic notes in organicchemistry which will be suitable for undergraduate students taking chemistrychemistry-related courses or courses which involve organic chemistry as anancillary subject The book concentrates on core topics which are most likelyto be common to those organic chemistry courses which follow on from a foun-dation or introductory general chemistry course

Organic chemistry is a subject which can lead some students to the heightsof ecstasy yet drive others up the wall Some students lsquoswitch onrsquo to it imme-diately while others can make neither head nor tail of it no matter how hardthey try Certainly one of the major problems in studying the subject is thevast amount of material which often has to be covered Many students blancheat the prospect of having to learn a seemingly endless number of reactionsand when it comes to drawing mechanisms and curly arrows they see onlya confusing maze of squiggly lines going everywhere yet nowhere Theconcepts of organic reaction mechanisms are often the most difficult to masterThese difficulties are often compounded by the fact that current textbooks inorganic chemistry are typically over 1200 pages long and can be quite expen-sive to buy

This book attempts to condense the essentials of organic chemistry into amanageable text of 310 pages which is student friendly and which does notcost an arm and a leg It does this by concentrating purely on the basics of thesubject without going into exhaustive detail or repetitive examplesFurthermore key notes at the start of each topic summarize the essential factscovered and help focus the mind on the essentials

Organic chemistry is a peculiar subject in that it becomes easier as you goalong This might seem an outrageous statement to make especially to a first-year student who is struggling to come to terms with the rules of nomenclaturetrying to memorize a couple of dozen reactions and making sense of mecha-nisms at the same time However these topics are the basics of the subject andonce they have been grasped the overall picture becomes clear

Understanding the mechanism of how a reaction takes place is particularlycrucial in this It brings a logic to the reactions of the different functional groupsThis in turn transforms a list of apparently unrelated facts into a sensible themewhich makes remembering the reactions a lsquopiece of cakersquo (well nearly)

Once this happy state of affairs has been reached the relevance of organicchemistry to other subjects such as genetics and biochemistry suddenly leapsoff the page Understanding organic chemistry leads to a better understandingof life chemistry and how the body works at the molecular level It also helpsin the understanding of the molecular mechanisms involved in disease andbodily malfunction leading in turn to an understanding of how drugs can bedesigned to cure these disease states ndash the science of medicinal chemistry

And thatrsquos not all An understanding of organic chemistry will help the indus-trial chemist or chemical engineer faced with unexpected side-reactions in achemical process and the agro-scientist trying to understand the molecularprocesses taking place within plants and crops and it will assist in the designand synthesis of new herbicides and fungicides which will be eco-friendly It

PREFACE

will aid the forensic scientist wishing to analyze a nondescript white powderndash is it heroin or flour

The list of scientific subject areas involving organic chemistry is endless ndashdesigning spacesuits developing new photographic dyes inventing new mole-cular technology in microelectronics ndash one could go on and on Organicchemistry is an exciting subject since it leads to an essential understanding ofmolecules and their properties

The order in which the early topics of this book are presented is importantThe first two sections cover structure and bonding which are crucial to latersections Just why does carbon form four bonds What is hybridization

The third section on functional groups is equally crucial if students are to becapable of categorizing the apparent maze of reactions which organiccompounds can undergo It is followed by section D on stereochemistry thensections E and F in which the basic theory of reactions and mechanisms iscovered What are nucleophiles and electrophiles What does a mechanismrepresent What does a curly arrow mean

The remaining sections can be used in any order These look at the reactionsand mechanisms of the common functional groups which are important inchemistry and biochemistry

It is hoped that students will find this textbook useful in their studies andthat once they have grasped what organic chemistry is all about they will readmore widely and enter a truly exciting world of molecular science

x Preface

Section A ndash Structure and bonding

A1 ATOMIC STRUCTURE OF CARBON

Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table This means that thereare two shells of atomic orbitals available for these electrons The first shell closestto the nucleus has a single s orbital ndash the 1s orbital The second shell has a single sorbital (the 2s orbital) and three p orbitals (3 2p) Therefore there are a total offive atomic orbitals into which these six electrons can fit The s orbitals are sphericalin shape with the 2s orbital being much larger then the 1s orbital The p orbitalsare dumbbell-shaped and are aligned along the x y and z axes Therefore they areassigned the 2px 2py and 2pz atomic orbitals (Fig 1)

Key Notes

The atomic orbitals available for the six electrons of carbon are the s orbitalin the first shell the s orbital in the second shell and the three p orbitals inthe second shell The 1s and 2s orbitals are spherical in shape The 2porbitals are dumbbell in shape and can be assigned 2px 2py or 2pz depend-ing on the axis along which they are aligned

The 1s orbital has a lower energy than the 2s orbital which has a lowerenergy than the 2p orbitals The 2p orbitals have equal energy (ie theyare degenerate)

Carbon is in the second row of the periodic table and has six electrons whichwill fill up lower energy atomic orbitals before entering higher energyorbitals (aufbau principle) Each orbital is allowed a maximum of two elec-trons of opposite spin (Pauli exclusion principle) When orbitals of equalenergy are available electrons will occupy separate orbitals before pairingup (Hundrsquos rule) Thus the electronic configuration of a carbon atom is 1s2

2s2 2px1 2py

1

Related topic Covalent bonding andhybridization (A2)

Electronicconfiguration

Atomic orbitals

Energy levels

1s 2s 2px 2py 2pz

y

xz

y

xz

y

xz

y

xz

y

xz

Fig 1 Atomic orbitals

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

will aid the forensic scientist wishing to analyze a nondescript white powderndash is it heroin or flour

The list of scientific subject areas involving organic chemistry is endless ndashdesigning spacesuits developing new photographic dyes inventing new mole-cular technology in microelectronics ndash one could go on and on Organicchemistry is an exciting subject since it leads to an essential understanding ofmolecules and their properties

The order in which the early topics of this book are presented is importantThe first two sections cover structure and bonding which are crucial to latersections Just why does carbon form four bonds What is hybridization

The third section on functional groups is equally crucial if students are to becapable of categorizing the apparent maze of reactions which organiccompounds can undergo It is followed by section D on stereochemistry thensections E and F in which the basic theory of reactions and mechanisms iscovered What are nucleophiles and electrophiles What does a mechanismrepresent What does a curly arrow mean

The remaining sections can be used in any order These look at the reactionsand mechanisms of the common functional groups which are important inchemistry and biochemistry

It is hoped that students will find this textbook useful in their studies andthat once they have grasped what organic chemistry is all about they will readmore widely and enter a truly exciting world of molecular science

x Preface

Section A ndash Structure and bonding

A1 ATOMIC STRUCTURE OF CARBON

Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table This means that thereare two shells of atomic orbitals available for these electrons The first shell closestto the nucleus has a single s orbital ndash the 1s orbital The second shell has a single sorbital (the 2s orbital) and three p orbitals (3 2p) Therefore there are a total offive atomic orbitals into which these six electrons can fit The s orbitals are sphericalin shape with the 2s orbital being much larger then the 1s orbital The p orbitalsare dumbbell-shaped and are aligned along the x y and z axes Therefore they areassigned the 2px 2py and 2pz atomic orbitals (Fig 1)

Key Notes

The atomic orbitals available for the six electrons of carbon are the s orbitalin the first shell the s orbital in the second shell and the three p orbitals inthe second shell The 1s and 2s orbitals are spherical in shape The 2porbitals are dumbbell in shape and can be assigned 2px 2py or 2pz depend-ing on the axis along which they are aligned

The 1s orbital has a lower energy than the 2s orbital which has a lowerenergy than the 2p orbitals The 2p orbitals have equal energy (ie theyare degenerate)

Carbon is in the second row of the periodic table and has six electrons whichwill fill up lower energy atomic orbitals before entering higher energyorbitals (aufbau principle) Each orbital is allowed a maximum of two elec-trons of opposite spin (Pauli exclusion principle) When orbitals of equalenergy are available electrons will occupy separate orbitals before pairingup (Hundrsquos rule) Thus the electronic configuration of a carbon atom is 1s2

2s2 2px1 2py

1

Related topic Covalent bonding andhybridization (A2)

Electronicconfiguration

Atomic orbitals

Energy levels

1s 2s 2px 2py 2pz

y

xz

y

xz

y

xz

y

xz

y

xz

Fig 1 Atomic orbitals

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Section A ndash Structure and bonding

A1 ATOMIC STRUCTURE OF CARBON

Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table This means that thereare two shells of atomic orbitals available for these electrons The first shell closestto the nucleus has a single s orbital ndash the 1s orbital The second shell has a single sorbital (the 2s orbital) and three p orbitals (3 2p) Therefore there are a total offive atomic orbitals into which these six electrons can fit The s orbitals are sphericalin shape with the 2s orbital being much larger then the 1s orbital The p orbitalsare dumbbell-shaped and are aligned along the x y and z axes Therefore they areassigned the 2px 2py and 2pz atomic orbitals (Fig 1)

Key Notes

The atomic orbitals available for the six electrons of carbon are the s orbitalin the first shell the s orbital in the second shell and the three p orbitals inthe second shell The 1s and 2s orbitals are spherical in shape The 2porbitals are dumbbell in shape and can be assigned 2px 2py or 2pz depend-ing on the axis along which they are aligned

The 1s orbital has a lower energy than the 2s orbital which has a lowerenergy than the 2p orbitals The 2p orbitals have equal energy (ie theyare degenerate)

Carbon is in the second row of the periodic table and has six electrons whichwill fill up lower energy atomic orbitals before entering higher energyorbitals (aufbau principle) Each orbital is allowed a maximum of two elec-trons of opposite spin (Pauli exclusion principle) When orbitals of equalenergy are available electrons will occupy separate orbitals before pairingup (Hundrsquos rule) Thus the electronic configuration of a carbon atom is 1s2

2s2 2px1 2py

1

Related topic Covalent bonding andhybridization (A2)

Electronicconfiguration

Atomic orbitals

Energy levels

1s 2s 2px 2py 2pz

y

xz

y

xz

y

xz

y

xz

y

xz

Fig 1 Atomic orbitals

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

2 Section A ndash Structure and bonding

Energy levels The atomic orbitals described above are not of equal energy (Fig 2) The 1s orbitalhas the lowest energy The 2s orbital is next in energy and the 2p orbitals have thehighest energies The three 2p orbitals have the same energy meaning that theyare degenerate

Electronic Carbon is in the second row of the periodic table and has six electrons which will configuration fill up the lower energy atomic orbitals first This is known as the aufbau princi-

ple The 1s orbital is filled up before the 2s orbital which is filled up before the 2porbitals The Pauli exclusion principle states that each orbital is allowed a maxi-mum of two electrons and that these electrons must have opposite spins There-fore the first four electrons fill up the 1s and 2s orbitals The electrons in eachorbital have opposite spins and this is represented in Fig 3 by drawing the arrowspointing up or down There are two electrons left to fit into the remaining 2porbitals These go into separate orbitals such that there are two half-filled orbitalsand one empty orbital Whenever there are orbitals of equal energy electrons willonly start to pair up once all the degenerate orbitals are half filled This is knownas Hundrsquos rule

The electronic configuration for carbon is 1s2 2s2 2px1 2py

1 The numbers insuperscript refer to the numbers of electrons in each orbital The letters refer to thetypes of atomic orbital involved and the numbers in front refer to which shell theorbital belongs

Energy

1s

2s

2px 2py 2pz

Fig 2 Energy levels of atomic orbitals

Energy

1s

2s

2px 2py 2pz

Fig 3 Electronic configuration for carbon

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Section A ndash Structure and bonding

A2 COVALENT BONDING ANDHYBRIDIZATION

Covalent bonding A covalent bond binds two atoms together in a molecular structure and is formedwhen atomic orbitals overlap to produce a molecular orbital ndash so called becausethe orbital belongs to the molecule as a whole rather than to one specific atom Asimple example is the formation of a hydrogen molecule (H2) from two hydrogenatoms Each hydrogen atom has a half-filled 1s atomic orbital and when the atomsapproach each other the atomic orbitals interact to produce two MOs (the numberof resulting MOs must equal the number of original atomic orbitals Fig 1)

The MOs are of different energies One is more stable than the original atomicorbitals and is called the bonding MO The other is less stable and is called theantibonding MO The bonding MO is shaped like a rugby ball and results from

Key Notes

When two hydrogen atoms approach each other their 1s atomic orbitalsinteract to form a bonding and an antibonding molecular orbital (MO) Astable covalent bond is formed when the bonding MO is filled with a pairof electrons and the antibonding MO is empty

Sigma (σ) bonds are strong bonds with a circular cross-section formed bythe head-on overlap of two atomic orbitals

The electronic configuration of atomic carbon implies that carbon shouldform two bonds However it is known that carbon forms four bonds Whencarbon is part of an organic structure it can lsquomixrsquo the 2s and 2p orbitals ofthe valence shell in a process known as hybridization There are threepossible types of hybridization ndash sp3 sp2 and sp hybridization

Related topics Atomic structure of carbon (A1) sp3 Hybridization (A3)

sp2 Hybridization (A4)sp Hybridization (A5)

Covalentbonding

Sigma bonds

Hybridization

+

1s atomicorbital

1s atomicorbital

= Bonding molecular orbital(full)

Antibonding molecular orbital(empty)Energy

HH

H H

H H H H

Fig 1 Molecular orbitals for hydrogen (H2 )

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

4 Section A ndash Structure and bonding

the combination of the 1s atomic orbitals Since this is the more stable MO thevalence electrons (one from each hydrogen) enter this orbital and pair up Theantibonding MO is of higher energy and consists of two deformed spheres Thisremains empty Since the electrons end up in a bonding MO which is more stablethan the original atomic orbitals energy is released and bond formation isfavored In the subsequent discussions we shall concentrate solely on the bond-ing MOs to describe bonding and molecular shape but it is important to realizethat antibonding molecular orbitals also exist

Sigma bonds The bonding molecular orbital of hydrogen is an example of a sigma (σ) bond σbonds have a circular cross-section and are formed by the head-on overlap of twoatomic orbitals This is a strong interaction and so sigma bonds are strong bondsIn future discussions we shall see other examples of σ bonds formed by theinteraction of atomic orbitals other than the 1s orbital

Hybridization Atoms can form bonds with each other by sharing unpaired electrons such thateach bond contains two electrons In Topic A1 we identified that a carbon atomhas two unpaired electrons and so we would expect carbon to form two bondsHowever carbon forms four bonds How does a carbon atom form four bondswith only two unpaired electrons

So far we have described the electronic configuration of an isolated carbonatom However when a carbon atom forms bonds and is part of a molecular struc-ture it can lsquomixrsquo the s and p orbitals of its second shell (the valence shell) This isknown as hybridization and it allows carbon to form the four bonds which weobserve in reality

There are three ways in which this mixing process can take place

the 2s orbital is mixed with all three 2p orbitals This is known as sp3 hybridization the 2s orbital is mixed with two of the 2p orbitals This is known as sp2

hybridization the 2s orbital is mixed with one of the 2p orbitals This is known as sp

hybridization

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Section A ndash Structure and bonding

A3 SP3 HYBRIDIZATION

Definition In sp3 hybridization the 2s orbital is mixed with all three of the 2p orbitals to give aset of four sp3 hybrid orbitals (The number of hybrid orbitals must equal thenumber of original atomic orbitals used for mixing) The hybrid orbitals willeach have the same energy but will be different in energy from the originalatomic orbitals That energy difference will reflect the mixing of the respect-ive atomic orbitals The energy of each hybrid orbital is greater than the original sorbital but less than the original p orbitals (Fig 1)

Electronic The valence electrons for carbon can now be fitted into the sp3 hybridized orbitalsconfiguration (Fig 1) There was a total of four electrons in the original 2s and 2p orbitals The s

orbital was filled and two of the p orbitals were half filled After hybridizationthere is a total of four hybridized sp3 orbitals all of equal energy By Hundrsquos rule

Key Notes

In sp3 hybridization the s and the p orbitals of the second shell are lsquomixedrsquoto form four hybridized sp3 orbitals of equal energy

Each hybridized orbital contains a single unpaired electron and so fourbonds are possible

Each sp3 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The hybridized orbitals arrange themselves as farapart from each other as possible such that the major lobes point to the cor-ners of a tetrahedron sp3 Hybridization explains the tetrahedral carbon insaturated hydrocarbon structures

Sigma (σ) bonds are strong bonds formed between two sp3 hybridized car-bons or between an sp3 hybridized carbon and a hydrogen atom A σ bondformed between two sp3 hybridized carbon atoms involves the overlap ofhalf filled sp3 hybridized orbitals from each carbon atom A σ bond formedbetween an sp3 hybridized carbon and a hydrogen atom involves a half-filled sp3 orbital from carbon and a half-filled 1s orbital from hydrogen

Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicmolecules This means that nitrogen has three half-filled sp3 orbitals and canform three bonds which are pyramidal in shape Oxygen has two half-filledsp3 orbitals and can form two bonds which are angled with respect to eachother Chlorine has a single half-filled sp3 orbital and can only form a singlebond All the bonds which are formed are σ bonds

Related topics Covalent bonding andhybridization (A2)

Bonds and hybridized centers (A6)

Electronicconfiguration

Nitrogen oxygenand chlorine

Sigma bonds

Definition

Geometry

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

they are all half filled with electrons which means that there are four unpairedelectrons Four bonds are now possible

Geometry Each of the sp3 hybridized orbitals has the same shape ndash a rather deformed lookingdumbbell (Fig 2) This deformed dumbbell looks more like a p orbital than an sorbital since more p orbitals were involved in the mixing process

Each sp3 orbital will occupy a space as far apart from each other as possible bypointing to the corners of a tetrahedron (Fig 3) Here only the major lobe of eachhybridized orbital has been shown and the angle between each of these lobes is1095 This is what is meant by the expression tetrahedral carbon The three-dimensional shape of the tetrahedral carbon can be represented by drawing a nor-mal line for bonds in the plane of the page Bonds going behind the page arerepresented by a hatched wedge and bonds coming out the page are representedby a solid wedge

Sigma bonds A half-filled sp3 hybridized orbital from one carbon atom can be used to form abond with a half-filled sp3 hybridized orbital from another carbon atom In Fig 4athe major lobes of the two sp3 orbitals overlap directly leading to a strong σ bondIt is the ability of hybridized orbitals to form strong σ bonds that explains whyhybridization takes place in the first place The deformed dumbbell shapes allowa much better orbital overlap than would be obtained from a pure s orbital or apure p orbital A σ bond between an sp3 hybridized carbon atom and a hydrogenatom involves the carbon atom using one of its half-filled sp3 orbitals and thehydrogen atom using its half-filled 1s orbital (Fig 4b)

6 Section A ndash Structure and bonding

2py 2pz

2s

2px

Original atomic orbitals

Energy

sp3 hybridized orbitals

Fig 1 sp3 Hybridization

Minor lobe Major lobe

Fig 2 sp3 Hybridized orbital

H

H

HCC

H

Bond going behind the page

Bond coming out of the pageTetrahedral shape

Bond in the plane of the page109 o5

Fig 3 Tetrahedral shape of an sp3 hybridized carbon

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

A3 ndash sp3 Hybridization 7

Nitrogen oxygen Nitrogen oxygen and chlorine atoms can also be sp3 hybridized in organicand chlorine structures Nitrogen has five valence electrons in its second shell After

hybridization it will have three half-filled sp3 orbitals and can form threebonds Oxygen has six valence electrons After hybridization it will have twohalf-filled sp3 orbitals and will form two bonds Chlorine has seven valenceelectrons After hybridization it will have one half-filled sp3 orbital and willform one bond

The four sp3 orbitals for these three atoms form a tetrahedral arrangement withone or more of the orbitals occupied by a lone pair of electrons Considering theatoms alone nitrogen forms a pyramidal shape where the bond angles are slightlyless than 1095 (c 107) (Fig 5a) This compression of the bond angles is due to theorbital containing the lone pair of electrons which demands a slightly greateramount of space than a bond Oxygen forms an angled or bent shape where twolone pairs of electrons compress the bond angle from 1095 to c 104 (Fig 5b)

Alcohols amines alkyl halides and ethers all contain sigma bonds involvingnitrogen oxygen or chlorine Bonds between these atoms and carbon are formedby the overlap of half-filled sp3 hybridized orbitals from each atom Bonds involv-ing hydrogen atoms (eg OndashH and NndashH) are formed by the overlap of the half-filled 1s orbital from hydrogen and a half-filled sp3 orbital from oxygen ornitrogen

C+

sp3sigma bondsp3

C Ca)

C

Fig 4 (a) σ Bond between two sp3 hybridized carbons (b) σ bond between an sp3

hybridized carbon and hydrogen

+

sp3

CH

1s sigma bond

HC

b)

N

CH3H

H

N

CH3H

H

107o

=

Pyramidal

a)

Fig 5 (a) Geometry of sp3 hybridized nitrogen (b) geometry of sp3 hybridized oxygen

O

H3C

HH

O

CH3

104o

Angled molecule

=

b)

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Section A ndash Structure and bonding

A4 SP2 HYBRIDIZATION

Key Notes

In sp2 hybridization a 2s orbital is lsquomixedrsquo with two of the 2p orbitals toform three hybridized sp2 orbitals of equal energy A single 2p orbital is leftover which has a slightly higher energy than the hybridized orbitals

For carbon each sp2 hybridized orbital contains a single unpaired electronThere is also a half-filled 2p orbital Therefore four bonds are possible

Each sp2 orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbital is a symmetrical dumbbellThe major lobes of the three sp2 hybridized orbitals point to the corners ofa triangle with the 2p orbital perpendicular to the plane

Each sp2 hybridized carbon forms three σ bonds using three sp2 hybridizedorbitals The remaining 2p orbital overlaps lsquoside onrsquo with a neighboring 2porbital to form a pi (π) bond The π bond is weaker than the σ bond but isstrong enough to prevent rotation of the CC bond Therefore alkenes areplanar with each carbon being trigonal planar

The oxygen and carbon atoms are both sp2 hybridized The carbon has threesp2 hybridized orbitals and can form three σ bonds one of which is to theoxygen The oxygen has one sp2 orbital which is used in the σ bond withcarbon The p orbitals on carbon and oxygen are used to form a π bond

Aromatic rings are made up of six sp2 hybridized carbons Each carbonforms three σ bonds which results in a planar ring The remaining 2p orbitalon each carbon is perpendicular to the plane and can overlap with a neigh-boring 2p orbital on either side This means that a molecular orbital isformed round the whole ring such that the six π electrons are delocalizedaround the ring This results in increased stability such that aromatic ringsare less reactive than alkenes

Conjugated systems such as conjugated alkenes and αβ-unsaturatedcarbonyl compounds involve alternating single and double bonds In suchsystems the p lobes of one π bond are able to overlap with the p lobes of aneighboring π bond and thus give a small level of double bond character tothe connecting bond This partial delocalization gives increased stability tothe conjugated system

Related topics Properties of alkenes and alkynes(H2)

Conjugated dienes (H11)Aromaticity (I1)

Properties (J2)αβ-Unsaturated aldehydes and

ketones (J11)Structureandproperties (K1)

Electronicconfiguration

Geometry

Definition

Alkenes

Carbonyl groups

Aromatic rings

Conjugated systems

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

A4 ndash sp2 Hybridization 9

Definition In sp2 hybridization the s orbital is mixed with two of the 2p orbitals (eg 2px and2pz) to give three sp2 hybridized orbitals of equal energy The remaining 2py orbitalis unaffected The energy of each hybridized orbital is greater than the original sorbital but less than the original p orbitals The remaining 2p orbital (in this casethe 2py orbital) remains at its original energy level (Fig 1)

Electronic For carbon there are four valence electrons to fit into the three hybridized sp2

configuration orbitals and the remaining 2p orbital The first three electrons are fitted into eachof the hybridized orbitals according to Hundrsquos rule such that they are all half-filled This leaves one electron still to place There is a choice between pairing itup in a half-filled sp2 orbital or placing it into the vacant 2py orbital The usual prin-ciple is to fill up orbitals of equal energy before moving to an orbital of higherenergy However if the energy difference between orbitals is small (as here) it iseasier for the electron to fit into the higher energy 2py orbital resulting in threehalf-filled sp2 orbitals and one half-filled p orbital (Fig 1) Four bonds are possible

Geometry The 2py orbital has the usual dumbbell shape Each of the sp2 hybridized orbitalshas a deformed dumbbell shape similar to an sp3 hybridized orbital Howeverthe difference between the sizes of the major and minor lobes is larger for the sp2

hybridized orbitalThe hybridized orbitals and the 2py orbital occupy spaces as far apart from each

other as possible The lobes of the 2py orbital occupy the space above and belowthe plane of the x and z axes (Fig 2a) The three sp2 orbitals (major lobes shownonly) will then occupy the remaining space such that they are as far apart from the2py orbital and from each other as possible As a result they are all placed in thexndashz plane pointing toward the corner of a triangle (trigonal planar shape Fig 2b)The angle between each of these lobes is 120 We are now ready to look at thebonding of an sp2 hybridized carbon

sp2

2px 2py

2s

2pz

Original atomic orbitals

2py

sp2 hybridized orbitals

Energy

Fig 1 sp2 Hybridization

x

y

z z

b)

120o

a)y

x

Fig 2 (a) Geometry of the 2py orbital (b) geometry of the 2 py orbital and the sp2

hybridized orbitals

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

10 Section A ndash Structure and bonding

Alkenes sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form atrigonal planar shape The use of these three orbitals in bonding explains theshape of an alkene for example ethene (H2CCH2) As far as the CndashH bonds areconcerned the hydrogen atom uses a half-filled 1s orbital to form a strong σ bondwith a half filled sp2 orbital from carbon (Fig 3a) A strong σ bond is also possiblebetween the two carbon atoms of ethene due to the overlap of sp2 hybridizedorbitals from each carbon (Fig 3b)

The full σ bonding diagram for ethene is shown in Fig 4a and can be simplified asshown in Fig 4b Ethene is a flat rigid molecule where each carbon is trigonal pla-nar We have seen how sp2 hybridization explains the trigonal planar carbons butwe have not explained why the molecule is rigid and planar If the σ bonds werethe only bonds present in ethene the molecule would not remain planar sincerotation could occur round the CndashC σ bond (Fig 5) Therefore there must be fur-ther bonding which lsquolocksrsquo the alkene into this planar shape This bond involves

the remaining half-filled 2py orbitals on each carbon which overlap side-on to pro-duce a pi (p) bond) with one lobe above and one lobe below the plane of the mol-ecule (Fig 6) This π bond prevents rotation round the CndashC bond since the π bondwould have to be broken to allow rotation A π bond is weaker than a σ bond sincethe 2py orbitals overlap side-on resulting in a weaker overlap The presence of a πbond also explains why alkenes are more reactive than alkanes since a π bond ismore easily broken and is more likely to take part in reactions

H

sp2 1s

C+

CH

Fig 3 (a) Formation of a CndashH σ bond (b) formation of a CndashC σ bond

+C C

sp2

C

sp2

C

Fig 4 (a) σ Bonding diagram for ethene (b) simple representation of σ bonds for ethene

C CH H

H HC C

H

H

H

HBond rotation

Fig 5 Bond rotation around a σ bond

C CH H

H HH

C CHHH

H

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

A4 ndash sp2 Hybridization 11

Carbonyl groups The same theory explains the bonding within a carbonyl group (CO) whereboth the carbon and oxygen atoms are sp2 hybridized The following energy leveldiagram (Fig 7) shows how the valence electrons of oxygen are arranged after sp2

hybridization Two of the sp2 hybridized orbitals are filled with lone pairs ofelectrons which leaves two half-filled orbitals available for bonding The sp2

orbital can be used to form a strong σ bond while the 2py orbital can be used forthe weaker π bond Figure 8 shows how the σ and π bonds are formed in thecarbonyl group and explains why carbonyl groups are planar with the carbonatom having a trigonal planar shape It also explains the reactivity of carbonylgroups since the π bond is weaker than the σ bond and is more likely to beinvolved in reactions

Fig 6 Formation of a π bond

sp2 hybridized orbitals

2px 2py 2pz

2s

Energy

2py

sp2

Original atomic orbitals

Fig 7 Energy level diagram for sp2 hybridized oxygen

C OH

HC O

H

H

sp2

a)

sigma bondsp2

Fig 8 (a) Formation of the carbonyl σ bond (b) formation of the carbonyl π bond

C O

H

H

2py2py

b)

C O

H

H

H

H

H

H

2py2py

π Bond

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Aromatic rings All the carbons in an aromatic ring are sp2 hybridized which means that eachcarbon can form three σ bonds and one π bond In Fig 9a all the single bonds areσ while each double bond consists of one σ bond and one π bond However thisis an oversimplification of the aromatic ring For example double bonds areshorter than single bonds and if benzene had this exact structure the ring wouldbe deformed with longer single bonds than double bonds (Fig 9b)

In fact the CndashC bonds in benzene are all the same length In order to understandthis we need to look more closely at the bonding which takes place Figure 10ashows benzene with all its σ bonds and is drawn such that we are looking into theplane of the benzene ring Since all the carbons are sp2 hybridized there is a 2py

orbital left over on each carbon which can overlap with a 2py orbital on either sideof it (Fig 10b) From this it is clear that each 2py orbital can overlap with its neigh-bors right round the ring This leads to a molecular orbital which involves all the 2py

orbitals where the upper and lower lobes merge to give two doughnut-like lobesabove and below the plane of the ring (Fig 11a) The molecular orbital is symmetri-cal and the six π electrons are said to be delocalized around the aromatic ring sincethey are not localized between any two particular carbon atoms The aromatic ringis often represented as shown in Fig 11b to represent this delocalization of the π

12 Section A ndash Structure and bonding

H

H

H

H

H

H

b)

=

a)

Fig 9 (a) Representation of the aromatic ring (b) lsquodeformedrsquo structure resulting from fixedbonds

H

H H

HHH

b)a)

C C

C

CC

C= =

a) b)

H

H

H

H

H

H

Fig 11 Bonding molecular orbital for benzene (b) representation of benzene to illustratedelocalization

Fig 10 (a) σ Bonding diagram for benzene (b) π Bonding diagram for benzene

a) b)

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

A4 ndash sp2 Hybridization 13

electrons Delocalization increases the stability of aromatic rings such that they areless reactive than alkenes (ie it requires more energy to disrupt the delocalized πsystem of an aromatic ring than it does to break the isolated π bond of an alkene)

Conjugated Aromatic rings are not the only structures where delocalization of π electrons cansystems take place Delocalization occurs in conjugated systems where there are alternat-

ing single and double bonds (eg 13-butadiene) All four carbons in 13-butadieneare sp2 hybridized and so each of these carbons has a half-filled p orbital which caninteract to give two π bonds (Fig 12a) However a certain amount of overlap isalso possible between the p orbitals of the middle two carbon atoms and so thebond connecting the two alkenes has some double bond character (Fig 12b) ndashborne out by the observation that this bond is shorter in length than a typicalsingle bond This delocalization also results in increased stability However it isimportant to realize that the conjugation in a conjugated alkene is not as greatas in the aromatic system In the latter system the π electrons are completelydelocalized round the ring and all the bonds are equal in length In 13-butadienethe π electrons are not fully delocalized and are more likely to be found in the ter-minal CndashC bonds Although there is a certain amount of π character in the middlebond the latter is more like a single bond than a double bond

Other examples of conjugated systems include αβ-unsaturated ketones and αβ-unsaturated esters (Fig 13) These too have increased stability due to conjugation

C CH

H HC C

H HHC C

HH H

C CH H

H

2py 2py 2py 2py

13-Butadiene

a)

Fig 12 (a) π Bonding in 13-butadiene (b) delocalization in 13-butadiene

CC

H

H3C H

C

O

CH3

CC

H

H3C H

C

O

OCH3

a) b)

Fig 13 (a) αβ-Unsaturated ketone (b) αβ-unsaturated ester

C CH

H HC

HHC

H

2py2py

2py2py

b)

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Section A ndash Structure and bonding

A5 SP HYBRIDIZATION

Definition In sp hybridization the 2s orbital is mixed with one of the 2p orbitals (eg 2px) togive two sp hybrid orbitals of equal energy This leaves two 2p orbitals unaffected(2py and 2pz) with slightly higher energy than the hybridized orbitals (Fig 1)

Key Notes

In sp hybridization the 2s orbital and one of the three 2p orbitals are lsquomixedrsquoto form two hybridized sp orbitals of equal energy Two 2p orbitals are leftover and have slightly higher energy than the unhybridized orbitals

For carbon each sp hybridized orbital contains a single unpaired electronThere are also two half-filled 2p orbitals Therefore four bonds are possible

Each sp orbital is shaped like a deformed dumbbell with one lobe muchlarger than the other The remaining 2p orbitals are symmetrical dumbbellsIf we define the 2p orbitals as being aligned along the y and the z axes thetwo sp hybridized orbitals point in opposite directions along the x axis

Each sp hybridized carbon of an alkyne can form two σ bonds using sphybridized orbitals The remaining 2p orbitals can overlap lsquoside-onrsquo to formtwo π bonds Alkynes are linear molecules and are reactive due to the πbonds

The nitrogen and carbon atoms of a nitrile group (CN) are both sphybridized The carbon has two sp hybridized orbitals and can form two σbonds one of which is to nitrogen The nitrogen has one sp orbital which isused in the σ bond with carbon Both the carbon and the nitrogen have two2p orbitals which can be used to form two π bonds

Related topics Properties of alkenes andalkynes (H2)

Chemistry of nitriles (O4)

Electronicconfiguration

Alkynes

Definition

Geometry

Nitrile groups

2px 2pz

Original atomic orbitals sp hybridized orbitals

sp

2py 2pz

Energy

2s

2py

Fig 1 sp Hybridization of carbon

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Electronic For carbon the first two electrons fit into each sp orbital according to Hundrsquos ruleconfiguration such that each orbital has a single unpaired electron This leaves two electrons

which can be paired up in the half-filled sp orbitals or placed in the vacant 2py and2pz orbitals The energy difference between the orbitals is small and so it is easierfor the electrons to fit into the higher energy orbitals than to pair up This leads totwo half-filled sp orbitals and two half-filled 2p orbitals (Fig 1) and so four bondsare possible

Geometry The 2p orbitals are dumbbell in shape while the sp hybridized orbitals aredeformed dumbbells with one lobe much larger than the other The 2py and 2pz

orbitals are at right angles to each other (Fig 2a) The sp hybridized orbitalsoccupy the space left over and are in the x axis pointing in opposite directions(only the major lobe of the sp orbitals are shown in black Fig 2b)

A molecule using the two sp orbitals for bonding will be linear in shape Thereare two common functional groups where such bonding takes place ndash alkynes andnitriles

Alkynes Let us consider the bonding in ethyne (Fig 3) where each carbon is sp hybridizedThe CndashH bonds are strong σ bonds where each hydrogen atom uses its half-filled

1s orbital to bond with a half-filled sp orbital on carbon The remaining sp orbitalon each carbon is used to form a strong σ carbonndashcarbon bond The full σ bondingdiagram for ethyne is linear (Fig 4a) and can be simplified as shown (Fig 4b)

Further bonding is possible since each carbon has half-filled p orbitals Thus the2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds(Fig 5) The π bond formed by the overlap of the 2py orbitals is represented in dark

A5 ndash sp Hybridization 15

sp

spz

b)a)

xx

y

2py2py

z

y

z2p z2p

Fig 2 (a) 2py and 2pz orbitals of an sp hybridized carbon (b) 2py 2pz and sp hybridizedorbitals of an sp hybridized carbon

H C C H

Fig 3 Ethyne

C CH HHa) b)

C C H

Fig 4 (a) σ Bonding for ethyne (b) representation of σ bonding

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

gray The π bond resulting from the overlap of the 2pz orbitals is represented inlight gray Alkynes are linear molecules and are reactive due to the relatively weakπ bonds

Nitrile groups Exactly the same theory can be used to explain the bonding within a nitrile group(CN) where both the carbon and the nitrogen are sp hybridized The energylevel diagram in Fig 6 shows how the valence electrons of nitrogen are arrangedafter sp hybridization A lone pair of electrons occupies one of the sp orbitals butthe other sp orbital can be used for a strong σ bond The 2py and 2pz orbitals can beused for two π bonds Figure 7 represents the σ bonds of HCN as lines and howthe remaining 2p orbitals are used to form two π bonds

16 Section A ndash Structure and bonding

CC

H

H

2py

2pz

2pz

2py

C

Pi Bond (π)

C

Fig 5 π-Bonding in ethyne

2py

sp

2s

2px 2py 2pz

Original atomic orbitals sp hybridized orbitals

2pz

Energy

Fig 6 sp Hybridization of nitrogen

CN

H

2py

2pz

2py

2pz

CN

C N

Pi Bond (π)

Fig 7 π-Bonding in HCN

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Section A ndash Structure and bonding

A6 BONDS AND HYBRIDIZED CENTERS

σ and π bonds Identifying σ and π bonds in a molecule (Fig 1) is quite easy as long as youremember the following rules

all bonds in organic structures are either sigma (σ) or pi (π) bonds all single bonds are σ bonds all double bonds are made up of one σ bond and one π bond all triple bonds are made up of one σ bond and two π bonds

Hybridized centers All the atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized (Fig 2)

Key Notes

Every bond in an organic structure is a σ bond or a π bond Every atom ina structure is linked to another by a single σ bond If there is more than onebond between any two atoms the remaining bonds are π bonds

All atoms in an organic structure (except hydrogen) are either sp sp2 or sp3

hybridized Atoms linked by single bonds are sp3 hybridized atoms linkedby double bonds are sp2 hybridized and atoms linked by triple bonds aresp hybridized

sp3 Hybridized centers are tetrahedral sp2 hybridized centers are trigonalplanar and sp centers are linear This determines the shape of functionalgroups Functional groups containing sp2 hybridized centers are planarwhile functional groups containing sp hybridized centers are linear

Functional groups containing π bonds tend to be reactive since the π bondis weaker than a σ bond and is more easily broken

Related topics

( with the exception of allenes R2CCCR2)

sp3 Hybridization (A3)sp2 Hybridization (A4)

sp Hybridization (A5)

Hybridized centers

σ and π bonds

Shape

Reactivity

C

H

C

H

CC

C

OC

O

CH

Cl H

OH3C

CH3H3CCH2

CH3H3CCH2

CH3H3C π

ππππ

Fig 1 Examples ndash all the bonds shown are σ bonds except those labelled as

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

The identification of sp sp2 and sp3 centers is simple if you remember the followingrules

all atoms linked by a single bond are sp3 hybridized (except hydrogen) both carbon atoms involved in the double bond of an alkene (CC) must be sp2

hybridized both the carbon and the oxygen of a carbonyl group (CO) must be sp2

hybridized all aromatic carbons must be sp2 hybridized both atoms involved in a triple bond must be sp hybridized hydrogen uses a 1s orbital for bonding and is not hybridized

Hydrogen atoms cannot be hybridized They can only bond by using an s orbitalsince there are no p orbitals in the first electron shell It is therefore impossible fora hydrogen to take part in π bonding Oxygen nitrogen and halogens on the otherhand can form hybridized orbitals which are either involved in bonding or inholding lone pairs of electrons

Shape The shape of organic molecules and the functional groups within them isdetermined by the hybridization of the atoms present For example functionalgroups containing trigonal planar sp2 centers are planar while functional groupscontaining sp centers are linear

planar functional groups ndash aldehyde ketone alkene carboxylic acid acidchloride acid anhydride ester amide aromatic

linear functional groups ndash alkyne nitrile functional groups with tetrahedral carbons ndash alcohol ether alkyl halide

Reactivity Functional groups which contain π bonds are reactive since the π bond is weakerthan a σ bond and can be broken more easily Common functional groups whichcontain π bonds are aromatic rings alkenes alkynes aldehydes ketonescarboxylic acids esters amides acid chlorides acid anhydrides and nitriles

18 Section A ndash Structure and bonding

H3CCH

CH2

CH3H3C

CCH2

CH3

H3CC

CCH3

Cl O H

H

sp2sp3

sp3

sp3 sp3

sp3

sp3 sp3sp3

sp3

sp3sp2

sp2sp2

Fig 2 Examples of sp sp2 and sp3 hybridized centers

H3CC

OC

CO

sp

sp

sp3 sp3sp2

sp2 H

Functional groups known as allenes (R2CCCR2) have an sp hybridized carbon locatedat the center of two double bonds but these functional groups are beyond the scope of thistext

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading

Section B ndash Alkanes and cycloalkanes

B1 DEFINITION

Alkanes Alkanes are organic molecules with the general formula CnH2n2 which consist ofcarbon and hydrogen atoms linked together by CndashC and CndashH single bonds Theyare often referred to as saturated hydrocarbons ndash saturated because all the bondsare single bonds hydrocarbons because the only atoms present are carbon andhydrogen All the carbon atoms in an alkane are sp3 hybridized and tetrahedral inshape The CndashC and CndashH bonds are strong σ bonds and so alkanes are unreactiveto most chemical reagents

Alkanes are sometimes referred to as straight chain or acyclic alkanes todistinguish them from cycloalkanes or alicyclic compounds

Cycloalkanes Cycloalkanes are cyclic alkanes (alicyclic compounds) having the general formulaCnH2n where the carbon atoms have been linked together to form a ring All sizesof ring are possible However the most commonly encountered cycloalkane inorganic chemistry is the six-membered ring (cyclohexane) Most cycloalkanes areunreactive to chemical reagents However small three- and four-membered ringsare reactive and behave like alkenes Such cyclic structures are highly strainedsince it is impossible for the carbon atoms to adopt their preferred tetrahedralshape

Key Notes

Alkanes are organic molecules consisting solely of carbon and hydrogenatoms linked by single σ bonds All the carbon atoms are tetrahedral and sp3

hybridized Alkanes are stable molecules and unreactive to most chemicalreagents They have the general formula CnH2n2

Cycloalkanes are cyclic alkane structures They have the general formulaCnH2n Most cycloalkanes are unreactive to chemical reagents Howeverthree- and four-membered rings are reactive due to ring strain and behavelike alkenes

Related topics sp3 Hybridization (A3) Conformational isomers (D4)

Alkanes

Cycloalkanes

• Book Cover
• Title
• Contents
• Preface
• - Structure and bonding
• Covalent bonding and hybridization
• sp3 Hybridization
• sp2 Hybridization
• sp Hybridization
• Bonds and hybridized centers
• - Alkanes and cycloalkanes
• Drawing structures
• Nomenclature
• - Functional groups
• Aliphatic and aromatic functional groups
• Intermolecular bonding
• Properties and reactions
• Nomenclature of compounds with functional groups
• Primary secondary tertiary and quaternary nomenclature
• - Stereochemistry
• Configurational isomers - alkenes and cycloalkanes
• Configurational isomers - optical isomers
• Conformational isomers
• - Nucleophiles and electrophiles
• Charged species
• Neutral inorganic species
• Organic structures
• - Reactions and mechanisms
• Mechanisms
• - Acid-base reactions
• Acid strength
• Base strength
• Lewis acids and bases
• Enolates
• - Alkenes and alkynes
• Properties of alkenes and alkynes
• Electrophilic addition to symmetrical alkenes
• Electrophilic addition to unsymmetrical alkenes
• Carbocation stabilization
• Reduction and oxidation of alkenes
• Hydroboration of alkenes
• Electrophilic additions to alkynes
• Reduction of alkynes
• Alkylation of terminal alkynes
• Conjugated dienes
• - Aromatic chemistry
• Preparation and properties
• Electrophilic substitutions of benzene
• Synthesis of mono-substituted benzenes
• Electrophilic substitutions of mono-substituted aromatic rings
• Synthesis of di- and tri-substituted benzenes
• Oxidation and reduction
• - Aldehydes and ketones
• Properties
• Nucleophilic addition
• Nucleophilic addition - charged nucleophiles
• Electronic and steric effects
• Nucleophilic addition - nitrogen nucleophiles
• Nucleophilic addition - oxygen and sulfur nucleophiles
• Reactions of enolate ions
• -Halogenation
• Reduction and oxidation
• -Unsaturated aldehydes and ketones
• - Carboxylic acids and carboxylic acid derivatives
• Nucleophilic substitution
• Reactivity
• Preparations of carboxylic acids
• Preparations of carboxylic acid derivatives
• Reactions
• Enolate reactions
• - Alkyl halides
• Nucleophilic substitution
• Factors affecting SN2 versus SN1 reactions
• Elimination
• Elimination versus substitution
• Reactions of alkyl halides
• Organometallic reactions
• - Alcohols phenols and thiols
• Preparation of phenols
• Properties of alcohols and phenols
• Reactions of alcohols
• Reactions of phenols
• Chemistry of thiols
• - Ethers epoxides and thioethers
• Properties of ethers epoxides and thioethers
• Reactions of ethers epoxides and thioethers
• - Amines and nitriles
• Properties of amines
• Reactions of amines
• Chemistry of nitriles
• - Organic spectroscopy and analysis
• Visible and ultra violet spectroscopy
• Infra-red spectroscopy
• Proton nuclear magnetic resonance spectroscopy
• 13C nuclear magnetic resonance spectroscopy
• Mass spectroscopy
• Further reading