Organic Chemistry ICHM 201

William A. Price, Ph.D.

Introduction and Review: Structure and Bonding

Atomic structureLewis Structures

ResonanceStructural Formulas

Acids and Bases

Chapter 1 5

Electronic Structure of the Atom

• An atom has a dense, positively charged nucleus surrounded by a cloud of electrons.

• The electron density is highest at the nucleus and drops off exponentially with increasing distance from the nucleus in any direction.

Orbitals are Probabilities

2s Orbital Has a Node

The p Orbital

The 2p Orbitals

• There are three 2p orbitals, oriented at right angles to each other.

• Each p orbital consists of two lobes.

• Each is labeled according to its orientation along the x, y, or z axis.

Chapter 1 9

px, py, pz

Electronic Configurations

• The aufbau principle states to fill the lowest energy orbitals first.

• Hund’s rule states that when there are two or more orbitals of the same energy (degenerate), electrons will go into different orbitals rather than pairing up in the same orbital.

Chapter 1 11

Electronic Configurations of Atoms• Valence electrons are electrons on the

outermost shell of the atom.

Chapter 1 12

Covalent Bonding• Electrons are shared between the atoms to complete

the octet.• When the electrons are shared evenly, the bond is

said to be nonpolar covalent, or pure covalent.• When electrons are not shared evenly between the

atoms, the resulting bond will be polar covalent.

Chapter 1 13

Lewis Dot Structure of Methane

C...

.

carbon - 4 valence e hydrogen - 1 valence e

H.

1s22s22p2 1s

Tetrahderal Geometry

CH4 NH3

H2O Cl2

Lewis Structures

C

H

H

H

HN H

H

H

O HH ClCl

Carbon: 4 e4 H@1 e ea: 4 e 8 e

Nitrogen: 5 e3 H@1 e ea: 3 e 8 e

Oxygen: 6 e2 H@1 e ea: 2 e 8 e

2 Cl @7 e ea: 14 e

Chapter 1 16

Bonding Patterns

Valence electrons (group #)

# Bonds # Lone Pair Electrons

C

N

O

Halides(F, Cl, Br, I)

4 4 0

5 3 1

6 2 2

7 1 3

Chapter 1 17

Bonding Characteristics of Period 2 Elements

Lewis structures are the way wewrite organic chemistry.

Learning now to draw themquickly and correctly will helpyou throughout this course.

HINT

Chapter 1 20

Multiple Bonding

• Sharing two pairs of electrons is called a double bond.• Sharing three pairs of electrons is called a triple bond.

Chapter 1 21

Convert Formula into Lewis Structure

• HCN• HNO2

• CHOCl• C2H3Cl

• N2H2

• O3

• HCO3-

• C3H4

Formal Charges

H3O+ NO+

OH H

H

Formal charge = [group number ] – [nonbonding electrons ] – ½ [shared electrons]

N O

6 – 2 – ½ (6) = +1 6 – 2 – ½ (6) = +1

5 – 2 – ½ (6) = 0+

+

• Formal charges are a way of keeping track of electrons.• They may or may not correspond to actual charges in the

molecule.

Chapter 1 23

Common Bonding Patterns

Chapter 1 24

Work enough problems tobecome familiar with thesebonding patterns so you canrecognize other patterns as

being either unusual or wrong.

HINT

Chapter 1 25

Electronegativity Trends Ability to Attract the Electrons in a Covalent Bond

Dipole Moment

• Dipole moment is defined to be the amount of charge separation (d) multiplied by the bond length (m).

• Charge separation is shown by an electrostatic potential map (EPM), where red indicates a partially negative region and blue indicates a partially positive region.

Chapter 1 27

Methanol

Dipole Moment (m) is sum of the Bond Moments

Nonpolar CompoundsBond Moments Cancel Out

Nitromethane

Nitromethane has 2 Formal Charges

CH3NO2 C N

O

O

H

HH

Formal Charge = [Group #] - [# bonds] - [# non-bonding electrons]

N = 5-4-0 = +1

O = 6-1-6 = -1

Both Resonance Structures Contribute to the Actual Structure

CH3NO2

C N

O

O

H

H

HHH

H

N

O

O

C

2 Equivalent Resonance Structures

Dipole Moment reflects Both Resonance Structures

C N

O

O

H

H

HHH

H

N

O

O

C

Resonance Hybrid

C

H

HH

O

O

N

Resonance Rules

• Cannot break single (sigma) bonds• Only electrons move, not atoms

3 possibilities:– Lone pair of e- to adjacent bond position

• Forms p bond

- p bond to adjacent atom- p bond to adjacent bond position

Curved Arrow Formalism Shows flow of electrons

C N

O

O

H

H

HHH

H

N

O

O

C

Arrows depict electron pairs moving

Resonance Forms• The structures of some compounds are not

adequately represented by a single Lewis structure.• Resonance forms are Lewis structures that can be

interconverted by moving electrons only. • The true structure will be a hybrid between the

contributing resonance forms.

Chapter 1 37

Resonance Forms

Resonance forms can be compared using the following criteria, beginning with the most important:1. Has as many octets as possible.2. Has as many bonds as possible.3. Has the negative charge on the most

electronegative atom.4. Has as little charge separation as possible.

Chapter 1 38

Two Nonequivalent Resonance Structures in Formaldehyde

Major and Minor Contributors

• When both resonance forms obey the octet rule, the major contributor is the one with the negative charge on the most electronegative atom.

N C O N C O

MAJOR MINOR

The oxygen is more electronegative,so it should have more of the negativecharge.

Chapter 1 40

Resonance Stabilization of IonsPos. charge is “delocalized”

CH

HC

C

H

H

H

C

C

H

H

H

H

HC

H

HC

H

H

H

CC

resonance hybrid

Solved Problem 2

Draw the important resonance forms for [CH3OCH2]+. Indicate which structure is the major and minor contributor or whether they would have the same energy.

The first (minor) structure has a carbon atom with only six electrons around it. The second (major) structure has octets on all atoms and an additional bond.

Solution

Chapter 1 42

Solved Problem 3Draw the resonance structures of the compound below. Indicate which structure is the major and minor contributor or whether they would have the same energy.

Both of these structures have octets on oxygen and both carbon atoms, and they have the same number of bonds. The first structure has the negative charge on carbon, the second on oxygen. Oxygen is the more electronegative element, so the second structure is the major contributor.

Solution

Chapter 1 43

Resonance Forms for the Acetate Ion

• When acetic acid loses a proton, the resulting acetate ion has a negative charge delocalized over both oxygen atoms.

• Each oxygen atom bears half of the negative charge, and this delocalization stabilizes the ion.

• Each of the carbon–oxygen bonds is halfway between a single bond and a double bond and is said to have a bond order of 1½.

Chapter 1 44

Condensed Structural Formulas Lewis Condensed

CH3CH3H C C H

H

H

H

H

1 2

• Condensed forms are written without showing all the individual bonds.

• Atoms bonded to the central atom are listed after the central atom (CH3CH3, not H3CCH3).

• If there are two or more identical groups, parentheses and a subscript may be used to represent them.

Chapter 1 45

Drawing Structures

CC

CC

H

H

H

H

HH H

HH

H

Butane, C4H10

CC

C

C

H

H H

HH

H H

HHH

Methylpropane, C4H10

CH3CH2CH2CH3 CH3CH(CH3)CH3

Octane Representations

CC

CC

CC

CC

H

HH

H H

H H

H H

H H

H H

HHH

HH

CH3CH2CH2CH2CH2CH2CH2CH3

CH3(CH2)6CH3

Lewis structure condensed structural formula

C8H18 is molecular formula but there are 18 possible structural isomers

Line-Angle Structures are Often Used as a Short-hand

Line-angle or "zig-zag" structures

Lewis structure

CC

CC

CC

CC

H

HH

H H

H H

H H

H H

H H

HHH

HH

Line-Angle Structures

CC

CC

CC

C

CC

C

CC

HH

H

H

H

H

H HH

H H

H

H

H

H

H

HH

HH

H

H H

HH

H

HH

H

H

H

H

H HH

H H

H

H

H

H

H

HH

HH

H

H H

HH

H

C12H26

HH

H

H

H

H

H HH

H H

H

H

H

H

H

HH

HH

H

H H

HH

H

At the end of every line and at the intersection of any lines there is a carbon atom with 4 bonds. Hydrogen atoms

are mentally supplied to fill the valency to 4.

Line-Angle structure Superimposed on Lewis Structure

CC

CC

CC

C

CC

C

CC

HH

H

H

H

H

H HH

H H

H

H

H

H

H

HH

HH

H

H H

HH

H

Line-Angle Drawings

H C C C

H

H

H

H

C

H

C

H

H

H H

H

C

O

H

O

H

1 2 3 4 5 61 2 3 4 5 6

• Atoms other than carbon must be shown.• Double and triple bonds must also be shown.

Chapter 1 52

For Cyclic Structures, Draw the Corresponding Polygon

Cyclohexane

CC

CC

C

C

HHH

HH

HHH

H

HH

H

Some Steroids

HOCholesterol

C27H44O

OH

OTestosterone

C19H26O2

Definitions of Acids/Bases

Arrhenius acid - forms H3O+ in H2O

Bronsted-Lowry acid - donates a H+ (proton)

Lewis acid - accepts an electron pair to form a new bond

Arrhenius base - forms OH- in H2O

Bronsted-Lowry base - accepts a H+ (proton)

Lewis base - donates an electron pair to form a new bond

Dissociation in H2OArrhenius Acid forms H3O+

Bronsted-Lowry Acid donates a H+

Brønsted-Lowry Acids and BasesBrønsted-Lowry acids are any species that donate a proton.

Brønsted-Lowry bases are any species that can accept a proton.

Chapter 1 57

Conjugate Acids and Bases

• Conjugate acid: when a base accepts a proton, it becomes an acid capable of returning that proton.

• Conjugate base: when an acid donates its proton, it becomes capable of accepting that proton back.

Chapter 1 58

Acid Strength defined by pKa

HCl + H2O H3O + Cl

Keq = [H3O ][Cl ]

[HCl][H2O]

Ka = Keq[H2O] =[H3O ][Cl ]

[HCl]= 10

7

pKa = -log(Ka) = -7

Stronger Acid Controls Equilibrium

HCl + H2O H3O + Clacid base conjugate conjugate

acid base

pKa = -7 -1.7 stronger weaker

Reaction Described with Arrows

H Cl + OHH O

H

HH+ Cl

Equilibrium Reactions

Identify the Acid and Base

CH3OH + OCOH

O

CH3O + HOCOH

O

Equilibrium Favors Reactants

CH3OH + OCOH

O

CH3O + HOCOH

O

acid base conj. base conj. acid

pKa 15.5 6.5

The Effect of Resonance on pKa

Effect of Electronegativity on pKa

• As the bond to H becomes more polarized, H becomes more positive and the bond is easier to break.

Chapter 1 66

Effect of Size on pKa

• As size increases, the H is more loosely held and the bond is easier to break.

• A larger size also stabilizes the anion.

Chapter 1 67

Lewis Acids and Lewis Bases

• Lewis bases are species with available electrons than can be donated to form a new bond.

• Lewis acids are species that can accept these electrons to form new bonds.

• Since a Lewis acid accepts a pair of electrons, it is called an electrophile.

Chapter 1 68

Nucleophiles and Electrophiles

• Nucleophile: Donates electrons to a nucleus with an empty orbital (same as Lewis Base)

• Electrophile: Accepts a pair of electrons (same as Lewis Acid)

• When forming a bond, the nucleophile attacks the electrophile, so the arrow goes from negative to positive.

• When breaking a bond, the more electronegative atom receives the electrons.

Chapter 1 69

Nucleophiles and Electrophiles

Chapter 1 70