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
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