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Organic Chemistry

Organic Chemistry (10 lectures)

Book: Essentials in Organic Chemistry by P.M.

Dewick

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Organic Chemistry

• 1st Lecture: Chapter 1 (Molecular representations and nomenclature) + Chapter 2 (Atomic structure and bonding)• 2nd Lecture: Chapter 4 (Acids and bases) and Chapter 3 (Stereochemistry)• 3rd Lecture: Chapter 5 (Reaction mechanism) + Chapter 6 (Nucleophilic

substitution) • 4th Lecture: Chapter 7 (Nucleophilic reaction of carbonyl groups – part 1)• 5th Lecture: Chapter 7 (Nucleophilic reaction of carbonyl groups – part 2)• 6th Lecture: Chapter 8 (Electrophilic reaction – part 1) + Chapter 9 (Radical• reactions: Catalytic hydration)• 7th Lecture: Chapter 9 (Radical reactions) + Chpater 8 (Electrophilic reactions: • Aromatic substitution)• 8th Lecture: Chapter 10 (Nucleophilic reaction involving enolate anions –part 1)• 9th Lecture: Chapter 10 (Nucleophilic reaction involving enolate anions –part 2)• 10th Lecture: Chapter 12 (Carbohydrates)

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Organic Chemistry

• The study of the compounds of carbon• Over 10 million organic compounds have been

identified– about 1000 new ones are discovered or

synthesized and identified each day!• C is a small atom

– it forms single, double, and triple bonds– it is intermediate in electronegativity (2.5)

– it forms strong bonds with C, H, O, N, and some metals

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Chapter 1 - Molecular representations

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Molecular representations

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Partial Structure

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Functional Groups

• Functional groupFunctional group: an atom or group of atoms within a molecule that shows a characteristic set of physical and chemical properties

• Functional groups are important for three reasons, they are– the units by which we divide organic compounds into

classes– the sites of characteristic chemical reactions– the basis for naming organic compounds

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Functional Groups

• Hydroxyl (OH) group of alcohols

An alcohol (Ethanol)

H-C-C-O-H

HH

H H

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Functional Groups

• Amino group of 1°, 2°, and 3° amines

CH3 N HH

CH3 N H

CH3

CH3 N CH3

CH3

(a 1° amine) (a 2° amine) (a 3° amine)

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Functional Groups

• Carbonyl group of aldehydes and ketones

O

CH3-C-H

O

CH3-C-CH3

An aldehyde A ketone

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Functional Groups

• Carboxyl group of carboxylic acids

O

CH3-C-O-H CH3COOH CH3CO2Hor or

Acetic acid

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Functional groups

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High priority

Lower priority

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Higher priority

Lowest priority

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Systematic Nomenclature (IUPAC)

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Common groups and abbreviations

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Chapter 2: Atomic structure and

bonding

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Electronic Structure of Atoms

• Shells are divided into subshells called orbitals, which are designated by the letters s, p, d,........– s (one per shell) – p (set of three per shell 2 and higher) – d (set of five per shell 3 and higher) .....

Shell Orbitals Contained in That Shell

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2

1 1s

2s, 2px, 2py, 2pz

3s, 3px, 3py, 3pz, plus five 3d orbitals

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Electronic configuration

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Lewis Structures

• Gilbert N. Lewis• Valence shellValence shell: the outermost electron shell of an

atom• Valence electronsValence electrons: electrons in the valence shell

of an atom; these electrons are used in forming chemical bonds

• Lewis structureLewis structure – the symbol of the atom represents the nucleus

and all inner shell electrons– dots represent valence electrons

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Lewis Structures

• Lewis structures for elements 1-18 of the Periodic Table

N OB

H

Li Be

Na Mg

He

1A 2A 5A 6A 7A 8A3A 4A

Cl

F

S

Ne

Ar

C

SiAl P

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Lewis Model of Bonding

• Atoms bond together so that each atom in the bond acquires the electron configuration of the noble gas nearest it in atomic number– an atom that gains electrons becomes an anionanion– an atom that loses electrons becomes a cationcation– Ionic bondIonic bond: a chemical bond resulting from the

electrostatic attraction of an anion and a cation– Covalent bondCovalent bond: a chemical bond resulting from two

atoms sharing one or more pairs of electrons

• We classify chemical bonds as ionic, polar covalent, and nonpolar covalent based on the difference in electronegativity between the atoms

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Electronegativity

• ElectronegativityElectronegativity: a measure of the force of an atom’s attraction for the electrons it shares in a chemical bond with another atom

• Pauling scale– increases from left to right within a period– increases from bottom to top in a group

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Electronegativity

• Electronegativity of atoms (Pauling scale)

insert Table 1.5page 7

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Electronegativity

• Electronegativity and chemical bonding

Difference in ElectronegativityBetween Bonded Atoms Type of Bond

less than 0.5

0.5 to 1.9greater than 1.9

nonpolar covalentpolar covalent

ionic

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Polar Covalent Bonds

• In a polar covalent bond– the more electronegative atom has a partial negative charge,

indicated by the symbol δδ--– the less electronegative atom has a partial positive charge,

indicated by the symbol δδ++• in an electron density model

– red indicates a region of high electron density – blue indicates a region of low electron density

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Drawing Lewis Structures

• To draw a Lewis structure:– determine the number of valence electrons in the

molecule or ion– determine the arrangement of atoms– connect the atoms by single bonds– arrange the remaining electrons so that each atom has a

complete valence shell– show bonding electrons as a single bond– show nonbonding electrons as a pair of dots– atoms share 1 pair of electrons in a single bond, 2 pairs

in a double bond, and 3 pairs in a triple bond

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Lewis Structures

Hydrogen chloride

Methane Ammonia

Water

H O

H

H

H NH C

H

H

H Cl

H

H

H2O (8)

NH3 (8)CH4 (8)

HCl (8)

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Lewis Structures

Carbonic acidMethanal

Acetylene

Ethylene

H

C C

H

C C HH

HO

CC

HO

H

O

H

HH

O

C2H4 (12)

C2H2 (10)

CH2O (12) H2CO3 (24)

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Lewis Structures

• In neutral molecules containing C, H, N, O, and halogen– hydrogen has one bond– carbon has 4 bonds and no unshared electrons– nitrogen has 3 bonds and 1 unshared pair of electrons– oxygen has 2 bonds and 2 unshared pairs of electrons– halogen has 1 bond and 3 unshared pairs of electrons

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Formal Charge

• Formal chargeFormal charge: the charge on an atom in a molecule or polyatomic ion– write a Lewis structure for the molecule or ion– assign each atom all its unshared (nonbonding) electrons and

one-half its shared (bonding) electrons– compare this number with the number of valence electrons in

the neutral, unbonded atom– if the number is less than that assigned to the unbonded atom,

the atom has a formal positive charge– if the number is greater, the atom has a formal negative

charge # of valence electrons in

unbonded atom

all unsharedelectrons

one-half of all shared

electrons+Formal

charge=

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Formal Charge

• Example:Example:

NH4+ CH3NH3

+

HCO3- CO3

2-OH-

CH3CO2-CH3

-

CH3OH2+

BF4-

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Polar and Nonpolar Molecules• A molecule is polar if

– it has polar bonds and– its center of positive charge is at a different place than

its center of negative charge

O C O-- +

Carbon dioxide(nonpolar)

-

+

Water(polar)

Insert elpot ofcarbon dioxide(page 19) Insert elpot

of water(page 19)H H

O

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Polar and Nonpolar Molecules

– ammonia and formaldehyde are polar molecules– acetylene is a nonpolar molecule

-

+

Insert elpot ofammonia(page 19)

Insert elpot offormaldehyde(page 20)

H HN

-

+H HC

O

Formaldehyde(polar)

HAmmonia

(polar)

H C C H

Insert elpot ofacetylene(page 20)

Acetylene(nonpolar)

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Resonance• For many molecules and ions, no single Lewis structure

provides a truly accurate representation

andO

CO

CH3 -

OC

OCH3

-

• Linus Pauling - 1930s– many molecules and ions are best described by writing

two or more Lewis structures– individual Lewis structures are called contributing

structures– connect individual contributing structures by a double-

headed (resonance) arrow– the molecule or ion is a hybrid of the various

contributing structures

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Resonance

N

O

O

N

O

O

Ethanoate ion(equivalent contributing structures)

C

O

O

CH3 C

O

O

CH3

Nitrite ion(equivalent contributing structures)

Curved arrowCurved arrow: a symbol used to show the redistribution of valence electrons

In using curved arrows, there are only two allowed types of electron redistribution:

—› from a bond to an adjacent atom

—› from an atom to an adjacent bond

Electron pushing by the use of curved arrows is a survival skill in organic chemistry

—› learn it well!

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Resonance

• All acceptable contributing structures must:

1. have the same number of valence electrons2. obey the rules of covalent bonding

– no more than 2 electrons in the valence shell of H – no more than 8 electrons in the valence shell of a 2nd

period element– 3rd period elements may have up to 12 electrons in

their valence shells3. differ only in distribution of valence electrons4. have the same number of paired and unpaired electrons

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VSEPR (Valence shell electron pair repulsion) Model

• This model is based on two concepts– atoms are surrounded by regions of electron density– regions of electron density repel each other

HC C

H

O C

C

H

NH

H

C

HH

OH

CH C H

H

O

4 regions of e- density(tetrahedral, 109.5°)

3 regions of e- density(trigonal planar, 120°)

2 regions of e- density(linear, 180°)

H

CH H

HN

H HH

::

::

::

::

:H O

H

:

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A covalent bond - Orbital Overlap Model

• A covalent bond forms when a portion of an atomic orbital of one atom overlaps a portion of an atomic orbital of another atom– in forming the covalent bond in H-H, for example, there

is overlap of the 1s orbitals of each hydrogen

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Hybrid Orbitals

• The Problem:– overlap of 2s and 2p atomic orbitals would give bond

angles of approximately 90°– instead we observe bond angles of approximately

109.5°, 120°, and 180°

• A Solution– hybridization of atomic orbitals– 2nd row elements use sp3, sp2, and sp hybrid orbitals for

bonding

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Hybrid Orbitals

• We study three types of hybrid atomic orbitals

spsp33 (one s orbital + three p orbitals give four sp3 orbitals)

spsp22 (one s orbital + two p orbitals give three sp2 orbitals)

spsp (one s orbital + one p orbital give two sp orbitals)

• Overlap of hybrid orbitals can form two types of bonds, depending on the geometry of the overlap

(sigma) bonds(sigma) bonds are formed by “direct” overlap

(pi) bonds(pi) bonds are formed by “parallel” overlap

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sp3 Hybrid Orbitals

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sp2 Hybrid Orbitals

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sp Hybrid Orbitals

– a carbon-carbon triple bond consists of one sigma (s) bond and two pi (p) bonds

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Hybrid Orbitals

• Summary of orbitals and bond types

H-C-C-HH

H

H

H

H

C C

H

H H

H-C C-H

Hybrid-ization

Types of Bonds to Carbon Example

sp3 four sigma bonds

sp2 three sigma bondsand one pi bond

sp two sigma bondsand two pi bonds

Ethane

Ethylene

Acetylene

Name