Chapter 2 Structure and Properties of Organic …faculty.sdmiramar.edu/choeger/Chap 2_Part1.pdf1...

1

Chapter 2Structure and Properties

of Organic Molecules

Chemistry 231Organic Chemistry I

Fall 2007

Organic Molecules Slide 2-2

Advanced Bonding: Review• Atomic Quantum Mechanics cannot explain

how molecules like CH4 form:• Valence Bond Theory (localized)

and Molecular Orbital Theory (delocalized)• Linear Combination of Atomic Orbitals

on different atoms produce molecular orbitals between partners on the same atom give hybrid orbitals which combine with

others.• Conservation of orbitals.• Orbitals that are in phase add together: amplitude increases.• Orbitals that are out of phase cancel out.• Bonding regions vs. antibonding regions.

H

C

H

H

H

109.5

2


Bonding Region• Electrons are between both nuclei, either on or off axis.

z axis


Sigma Bonding• Electron density lies between the nuclei on the z axis (by def).• A bond may be formed by s-s, pz-pz, s-pz, or hybridized orbital

overlaps along same axis (later).

• Anything attached to another atom must use one (and only one0sigma bond per group to do so.

• The bonding MO is lower in energy than the original atomicorbitals.

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

AO1AO2

!*

!

3


Bonding Molecular OrbitalTwo hydrogens, 1s constructive overlap


Anti-Bonding Molecular Orbital

Two hydrogens, destructive overlap.

Antibonding MO’s havenodes between nucleicenters

4


H2: s-s overlap


Cl2: p-p overlap

Constructive overlap along the sameaxis forms a sigma bond.

5


HCl: s-p overlap

What is the predicted shape for the bonding MO and the antibondingMO of the HCl molecule?


Pi (π) Bonding• Pi bonds form after sigma bonds; do not form from hybrid

orbitals.• Sideways overlap of parallel p orbitals; electron density above

and below nuclear axis.

6


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.

C C

H

H

H

H

C CH H

HH

C C HH


Molecular Shapes

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

X

A

X

X

X

109.5

X A

X

X

X

A

X

X

120

A XX

180

120

tetrahedral trigonal planar linear

X

A

XX

XX

trigonal bipyramid(transition state)

7


sp Hybrid Orbitals

• 2 VSEPR pairs• Linear electron pair

geometry• 180° bond angle• 2 atomic orbitals =

2 hybrid orbitals• 1 hybrid orbital per

group attached tocentral atom

• Orbitals not used forhybrid can be used forpi bonding


sp2 Hybrid Orbitals• 3 VSEPR pairs; Trigonal planar e- pair geometry• 120° bond angle• 3 atomic orbitals = 3 hybrid orbitals• 1 hybrid orbital per group attached to central atom• Orbital not used for hybrid can be used for pi bonding

8


sp3 Hybrid Orbitals• 4 VSEPR pairs• Tetrahedral e- pair geometry• 109.5° bond angle• 4 atomic orbitals = 4 hybrid orbitals• 1 hybrid orbital per group attached to central atom• NO Orbital available for pi bonding in Period 2 elements


• Single bonds in acyclic connections freely rotate

• Double bonds cannot rotate unless the bond is broken(restricted rotation).

Rotation around Bonds

X

C C

H

HH

H X

X

C C

H H

HH

X

rotate

Called conformations

C CX H

XH

C CX X

HH

X

Called configurations

9


Isomerism• Same molecular formula, but different arrangement of atoms:

isomers.• Two major types:

Constitutional (or structural) isomers: differ in theirbonding sequence.

Stereoisomers: differ only in the arrangement of the atomsin 3D-space.


Structural Isomers

CH3 O CH3 and CH3 CH2 OH

CH3

CH3

and

10


Stereoisomers

C C

Br

CH3

Br

H3C

C C

CH3

Br

Br

H3C

and

Cis - same side Trans - across

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

C C

H3C

H H

HNo cis-trans isomers possible


Molecular Dipole Moments• Depend on bond polarity (due to electronegativity differences)

and bond angles (VSEPR structure).• Vector sum of all the bond dipole moments in the molecule.• All polar molecules have polar bonds but not vice versa

11


Effect of Lone PairsLone pairs of electrons contribute strongly to the dipole moment.


Intermolecular Forces• Strength of attractions between individual molecules

influence m.p., b.p., and solubility, esp. for solids and liquids.• Classification depends on structure.

Dipole-dipole interactionsLondon dispersionsHydrogen bonding

12


Dipole-Dipole Forces• Between polar molecules; requires fixed and permanent

dipole.• Partial positive end of one molecule aligns with partial

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.


Dipole-Dipole

13


London Dispersions• Occurs for all molecules; most important IMF between

nonpolar molecules• Temporary dipole-dipole interactions caused by

instantaneous, but fleeting, molecular dipoles due toelectron cloud distortions.

• Larger atoms are more polarizable, as are larger molecules.• Branching lowers b.p. because of decreased surface contact

between molecules.


Dispersions

14


Hydrogen Bonding

• Not a formal bond; rather, a strong dipole-dipole attraction.• For this to occur, one of the organic molecules undergoing

H-bonding must have N-H or O-H while the other must havean N or O.

• The hydrogen attached to an O or N on one molecule isstrongly attracted to the lone pair of electrons belonging toan O or N on the other molecule.

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


H Bonds

15


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


Solubility• Like dissolves like.• Polar solutes dissolve in polar solvents.• Nonpolar solutes dissolve in nonpolar solvents.• Molecules with similar intermolecular forces will mix freely.• “Hoeger’s Rules”

16


Ionic Solute withPolar Solvent

Hydration releases energy.Entropy increases.


Ionic Solute withNonpolar Solvent

=>

17


Nonpolar Solute withNonpolar Solvent

=>


Nonpolar Solutewith Polar Solvent

=>

18


Classes of Compounds

• Classification based on functional group.• Three general classes

HydrocarbonsCompounds containing oxygenCompounds containing nitrogen.


Classification of CompoundsIn addition to classifying compounds according to functional group, we also use a

broader scheme to describe/discuss compounds and atoms in reactionschemes:

• Classification by hybridization: the hybridization of the atom at whichreaction occurs is quite often important to know. Classes are sp3, sp2, and spand are used for all atoms.

• Classification by carbon attachment: this classification is only used for sp3

carbons; any atom or functional group will by default carry that carbon’sclassification. Classes are primary (1°), secondary (2°), tertiary (3°), andquaternary (4°).

C C OH

H

H1o carbon

1o hydrogen

1o alcohol

C C Br

C

H2o carbon

2o hydrogen

2o halide

C C H

C

C3o carbon

3o hydrogen

19


Hydrocarbons• Alkane: single bonds, sp3 carbons

Cycloalkane: sp3 carbons form a ring• Alkene: double bond, sp2 carbons

Cycloalkene: double bond in ring• Alkyne: triple bond, sp carbons; not usually found in a ring• Aromatic: simple: contains a benzene ring; more

complicated: follow Huckel’s rules (later)

Into this category we also include the hydrocarbon derivativessuch as the nitro- and the halogen- containing compounds thatdo not contain other functional groups.

1


Chapter 2 Homework

23, 24, 28, 29, 35, 37, 39, 40, 42, 44

Chapter 2 Structure and Properties of Organic …faculty.sdmiramar.edu/choeger/Chap 2_Part1.pdf1...

Documents

Transcript of Chapter 2 Structure and Properties of Organic …faculty.sdmiramar.edu/choeger/Chap 2_Part1.pdf1...