Chapter 2Structure and Properties

of Organic Molecules

Organic Chemistry, 6th EditionL. G. Wade, Jr.

Jo BlackburnRichland College, Dallas, TX

Dallas County Community College District2006,Prentice Hall

Chapter 2 2

Wave Properties of Electrons

• Standing wave vibrates in fixed location.

• Wave function, , mathematical description of size, shape, orientation.

• Amplitude may be positive or negative.

• Node: amplitude is zero.

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Chapter 2 3

Wave Interactions• Linear combination of atomic orbitals

on different atoms produce molecular orbitals

on the same atom give hybrid orbitals.

• Conservation of orbitals.• Waves that are in phase add together.

Amplitude increases.• Waves that are out of phase cancel out.

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Chapter 2 4

Bonding Region

• Electrons are close to both nuclei.

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Chapter 2 5

Sigma Bonding

• Electron density lies between the nuclei.• A bond may be formed by s-s, p-p, s-p,

or hybridized orbital overlaps.• The bonding MO is lower in energy than

the original atomic orbitals.• The antibonding MO is higher in energy

than the atomic orbitals. =>

Chapter 2 6

Bonding Molecular Orbital

Two hydrogens, 1s constructive overlap

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Chapter 2 7

Anti-Bonding Molecular Orbital

Two hydrogens, destructive overlap.

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Chapter 2 8

H2: s-s overlap

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Chapter 2 9

Cl2: p-p overlap

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Constructive overlap along the sameaxis forms a sigma bond.

Chapter 2 10

HCl: s-p overlap

Question: What is the predicted shape for the bonding MO and the antibonding MO of the HCl molecule?

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Chapter 2 11

Pi Bonding• Pi bonds form after sigma bonds.

• Sideways overlap of parallel p orbitals.

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Chapter 2 12

Multiple Bonds

• A double bond (2 pairs of shared electrons) consists of a sigma bond and a pi bond.

• A triple bond (3 pairs of shared electrons) consists of a sigma bond and two pi bonds.

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Chapter 2 13

Molecular Shapes

• Bond angles cannot be explained with simple s and p orbitals. Use VSEPR theory.

• Hybridized orbitals are lower in energy because electron pairs are farther apart.

• Hybridization is LCAO within one atom, just prior to bonding. =>

Chapter 2 14

sp Hybrid Orbitals

• 2 VSEPR pairs

• Linear electron pair geometry

• 180° bond angle

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Chapter 2 15

sp2 Hybrid Orbitals• 3 VSEPR pairs

• Trigonal planar e- pair geometry

• 120° bond angle

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Chapter 2 16

sp3 Hybrid Orbitals

• 4 VSEPR pairs

• Tetrahedral e- pair geometry

• 109.5° bond angle

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Chapter 2 17

Sample Problems• Predict the hybridization, geometry,

and bond angle for each atom in the following molecules:

• Caution! You must start with a good Lewis structure!

• NH2NH2

• CH3-CC-CHO

CH3 C

O

CH2

_

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Chapter 2 18=>

• Single bonds freely rotate.

• Double bonds cannot rotate unless the bond is broken.

Rotation around Bonds

Chapter 2 19

Isomerism

• Same molecular formula, but different arrangement of atoms: isomers.

• Constitutional (or structural) isomers differ in their bonding sequence.

• Stereoisomers differ only in the arrangement of the atoms in space. =>

Chapter 2 20

Structural Isomers

CH3 O CH3 and CH3 CH2 OH

CH3

CH3

and

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Chapter 2 21

Stereoisomers

C CBr

CH3

Br

H3CC C

CH3

Br

Br

H3Cand

Cis - same side Trans - across

Cis-trans isomers are also called geometric isomers.There must be two different groups on the sp2 carbon.

C CH3C

H H

H No cis-trans isomers possible

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Chapter 2 22

Bond Dipole Moments• are due to differences in electronegativity.• depend on the amount of charge and

distance of separation.• In debyes,

x (electron charge) x d(angstroms)

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Chapter 2 23

Molecular Dipole Moments

• Depend on bond polarity and bond angles.

• Vector sum of the bond dipole moments.

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Chapter 2 24

Effect of Lone Pairs

Lone pairs of electrons contribute to the dipole moment.

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Chapter 2 25

Intermolecular Forces

• Strength of attractions between molecules influence m.p., b.p., and solubility, esp. for solids and liquids.

• Classification depends on structure.Dipole-dipole interactionsLondon dispersionsHydrogen bonding

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Chapter 2 26

Dipole-Dipole Forces• Between polar molecules.

• Positive end of one molecule aligns with negative end of another molecule.

• Lower energy than repulsions, so net force is attractive.

• Larger dipoles cause higher boiling points and higher heats of vaporization. =>

Chapter 2 27

Dipole-Dipole

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Chapter 2 28

London Dispersions• Between nonpolar molecules• Temporary dipole-dipole interactions• Larger atoms are more polarizable.• Branching lowers b.p. because of

decreased surface contact between molecules.

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Chapter 2 29

Dispersions

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Chapter 2 30

Hydrogen Bonding

• Strong dipole-dipole attraction.

• Organic molecule must have N-H or O-H.

• The hydrogen from one molecule is strongly attracted to a lone pair of electrons on the other molecule.

• O-H more polar than N-H, so stronger hydrogen bonding. =>

Chapter 2 31

H Bonds

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Chapter 2 32

Boiling Points and Intermolecular Forces

CH3 CH2 OH

ethanol, b.p. = 78°C

CH3 O CH3

dimethyl ether, b.p. = -25°C

trimethylamine, b.p. 3.5°C

N CH3H3C

CH3

propylamine, b.p. 49°C

CH3CH2CH2 N

H

H

ethylmethylamine, b.p. 37°C

N CH3CH3CH2

H

CH3 CH2 OH CH3 CH2 NH2

ethanol, b.p. = 78° C ethyl amine, b.p. = 17 ° C

Chapter 2 33

Solubility

• Like dissolves like.

• Polar solutes dissolve in polar solvents.

• Nonpolar solutes dissolve in nonpolar solvents.

• Molecules with similar intermolecular forces will mix freely. =>

Chapter 2 34

Ionic Solute with Polar Solvent

Hydration releases energy.Entropy increases. =>

Chapter 2 35

Ionic Solute withNonpolar Solvent

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Chapter 2 36

Nonpolar Solute withNonpolar Solvent

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Chapter 2 37

Nonpolar Solute with Polar Solvent

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Chapter 2 38

Classes of Compounds

• Classification based on functional group.

• Three broad classesHydrocarbonsCompounds containing oxygenCompounds containing nitrogen.

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Chapter 2 39

Hydrocarbons

• Alkane: single bonds, sp3 carbons

• Cycloalkane: carbons form a ring

• Alkene: double bond, sp2 carbons

• Cycloalkene: double bond in ring

• Alkyne: triple bond, sp carbons

• Aromatic: contains a benzene ring =>

Chapter 2 40

Compounds Containing Oxygen

• Alcohol: R-OH

• Ether: R-O-R'

• Aldehyde: RCHO

• Ketone: RCOR'

CH3CH2 C

O

H

CH3 C

O

CH3

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Chapter 2 41

Carboxylic Acids and Their Derivatives

• Carboxylic Acid: RCOOH

• Acid Chloride: RCOCl

• Ester: RCOOR'

• Amide: RCONH2

C

O

OH

C

O

Cl

C

O

OCH3C

O

NH2

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Chapter 2 42

Compounds Containing Nitrogen

• Amines: RNH2, RNHR', or R3N

• Amides: RCONH2, RCONHR, RCONR2

• Nitrile: RCN

N

O

CH3

CH3 C N

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Chapter 2 43

End of Chapter 2