2-1

2-2

California State Polytechnic University, PomonaDr. Laurie S. Starkey, Organic Chemistry CHM 314,

Wade Chapter 2: Structure and Physical Properties of Organic Molecules

Molecular Orbitals (MO) - formed by overlap of Atomic Orbitals (AO) to make covalent bonds - TWO AO's combine to give TWO MO's (there are TWO possible combinations)

Example 1 Consider the formation of the sigma bond in H2 by combining two H atoms:

H H H Ha b

two electrons shared in a σ bond (a σ MO)

AO's(s orbitals) sa sb

(same sign)

sa - sb(out of phase)

sa + sb(in phase) σ MO - bonding molecular orbital

(favorable overlap, low E)

σ* MO - antibonding "sigma star"(high E, usually empty)

no electron density holdingatoms together - "anti"bonding

σ*

σ

AO AO

antibonding orbital is empty

bonding orbital contains twoelectrons = a sigma bond!

Ener

gy

sa sb

PLEASE NOTEan increase ↑ in # of nodes

results in anincrease ↑ in Energy

(the orbital with MORE nodes is LESS stable)

1) Orbitals & Bonding (2-1 to 2-3)

Chapter Outline1) Orbitals and Bonding (2-1 to 2-3)2) Hybrid Orbitals (2-4)3) 3-D Sketches (2-5 to 2-7)4) Physical Properties (2-9 to 2-11)5) Isomers (2-8)6) Introduction to Functional Groups (FG) (2-12 to 2-14)

Atomic Orbitals (AO) - a region with a high probability of finding electron (e-) density - defined by mathematical equations called wave functions - mathematical sign of the wave function changes at a "node" - electron density = 0 at any node

x

y

z

s

x

y

z

p

x

y

z

p

x

y

z

p

2-2Example 2 Consider the formation of a pi bond, by overlapping two p orbitals

pa pb

C C

pa + pbπ bond

(bonding MO)

C C

pa - pbπ* "pi star"

(antibonding MO)

and

pi bond = cloud of electrondensity above and below

Ener

gy

Overall E levels of MO's

σπ

nπ∗

σ∗

(n = nonbonding "lone pairs")antibonding orbitalshigh E = less stable

(usually empty)electrons in these MO's are less stable than σ electrons

and are more reactivemost stable, strongest bond, least reactive

twopossible

combinations

resulting intwo new MOs

FYI: Electronic Transitions (Wade 15-13B, UV-Vis Spectroscopy)

π*

πground state

pi bond(low E)

π*

πexcited state

pi bond(high E)

hν(light E,usually

UV light)

increase ↑ # of conjugated pi bondsincrease ↑ resonance stabilizationdecrease ↓ E needed for π → π*

if... then...and...

lower Energyvisible lightis absorbed

i.e., COLOR!

this E gapgets smallerif conjugated

π−σ−π

two AO's

π*

π

AO AO

Ener

gy

pa pb

2) Hybridization (2-4)

How are the bonds in methane, CH4, formed?

carbon's atomic orbitals (AOs) contain _________ valence electrons

2p ____ ____ ____

2s ____

px py pzs

carbon's AOs

2p ____ ____ ____

2s ____

add Energy

promoteelectron

But CH4 has 4 identical bonds. How can that be?

mixing of AOs to give new hybrid orbitals

type of hybridization (sp, sp2, sp3) depends on the number of groups around the carbon"regions of electron density"

2-3Determining Hybridization

Examplemolecule

Regions ofe– density

Hybrid-ization

s p p p Result Geometry(VSEPR)

C CH

HH

HH

H

C CH

H

H

H

C C HH

practice: assign hybridizations on given molecule

1) complete Lewis structure2) hybridization is for each atom3) count "regions" on each atom

a "region of electron density" is a lone pair or single bond or double bond or triple bond

3) 3-D Sketches of Molecules (2-5 to 2-7)

CH3 CH3

CH2 CH2

HC CH

note: can rotate about σ bond(many drawings are possible)

note: CANNOT rotate about π bond

(aligned p orbitals)

C C HH

For the indicated bonds, describe the type of bond and determine which orbitals overlap to form them.

3 bondsN C C

O

C H

H

H

2-4practice: provide 3D sketch of given molecule

C C

O

CH3

1) complete Lewis structure

2) assign atom hybridizations

3) sketch with maximum number of atoms in the plane of the pageN

Hybridization and Resonance (see Wade problem 2-7 and solved problem 2-8)

try 3-D sketch of allene H2CCCH2 (problem 2-6)

CH3 C

O

NH2

4) Physical Properties (2.9 to 2.11)Physical properties, such as water solubility and boiling point (bp) are based onintermolecular forces/attractions.

methanol liquid

heatΔ

(bp)

methanol gas

if molecules are strongly attracted to one another, then- requires a lot of energy to separate them from each other

A Dipole-Dipole

B Hydrogen Bonding

C van der Waal's/London Dispersion

Types of "nonbonding" interactions

CH3

OH

CH3

OH

CH3

OHCH3

OH

CH3

OH

CH3

OH

CH3

OHCH3

OHCH3

OHCH3

OH

CH3

OH CH3

OH

CH3

OHCH3

OH

CH3

OH

CH3

OH

CH3

OH

- will have a high/low boiling point

An allylic lone pair must be in a p orbital in order to have resonance delocalization!Atoms do not move in resonance, so hybridization is sp2 throughout all resonance forms.

2-5A Dipole-Dipole - attraction between polar molecules (consider geometry! Is CCl4 polar?)

a polar molecule:

NaCl H

bp ˚C 1413 76 36

O

Overall trend:

polarity

bpB Hydrogen Bonding - strongest known dipole due to H on N or O

H N H O both are extremely polar bonds, can cause H-bond formation

OH

H

δ-δ+

δ+

hydrogen-bonding in water:

hydrogen-bonding in DNA base pairs:

NN

O

O

CH3

NN

N

NNHH

H

thymine (T) adenine (A)

CH3CH2CH2CH3

bp ˚C -1 10 36

C Van der Waal's/London Dispersion Forces - induced (temporary) dipoles in nonpolar molecules

C31H64

> 300

CH3CH2CH2CH2CH3CH3 C CH3

CH3

CH3

temporary attraction because of unevendistribution of electrons

- the greater the surface area, the greater the VDW/Dispersion forces (think "Velcro")- the higher the MW, the higher the bp (if all polarity is equal)

H2O

bp ˚C 100 78 -24

CH3OCH3

-42

CH3CH2CH3CH3CH2OH

2-6

straight-chain vs. branched

to predict boiling points

1) H-bonding (OH or NH)2) polar vs. nonpolar

3) MW, bp4) branching (least important!)

bp 36˚CCH3CH2CH2CH2CH3

bp 10˚C

CH3 C CH3

CH3

CH3

Water Solubility (2-11) - "like dissolves like" (see Figures 2-26 to 2-29) - water is polar and can form hydrogen bonds (H-bonds)

CH3CH2OH CH3CH2CH2OH OH

CH3

C

O

CH3

acetone

- miscible with water

- polar

- H-bond acceptor

5) Isomerism (2-8) Isomers are different compounds that have the same molecular formula.

Constitutional (Structural) Isomers: same formula, different connectivity

Stereoisomers: same formula AND same connectivity, but different spatial arrangement (3D)

2-7 California State Polytechnic University, Pomona Organic Chemistry, CHM 314, Dr. Laurie S. Starkey

Chapter 2 Summary (Wade textbook)

I. Atomic Orbitals (AO's) combine to give Molecular Orbitals (MO's) (2-1 to 2-3) A) Bonding MO's (σ, π) contain electrons in covalent bonds B) Antibonding MO's (σ∗, π∗) are usually empty, can contain excited electrons C) Relative energies, stabilities of MO'sII. Hybrid Orbitals (2-4) A) sp3 hybridization i) 4 regions of electron density ii) tetrahedral geometry B) sp2 hybridization i) 3 regions of electron density ii) trigonal planar geometry iii) contains an unhybridized p orbital C) sp hybridization i) 2 regions of electron density ii) linear geometry iii) contains two unhybridized p orbitalsIII. 3-D sketches (2-5 to 2-7) A) determine hybridization to learn geometry about each atom B) draw aligned p orbitals to show π bonds C) apply sp2 hybridization for atoms involved in resonanceIV. Physical Properties (2-9 to 2-11) A) Nonbonding (intermolecular) Interactions affect bp, mp i) dipole-dipole for polar molecules (δ+, δ-) ii) hydrogen bonding for molecules containing NH, OH or HF (STRONG dipole) iii) van der Waal's (London dispersion) temporary dipole moments a) explains why bp varies by MW (higher MW, higher bp) b) straight vs. branched molecules (greater surface area, higher bp) B) mp increases for molecules that can pack tighter (more spherical, higher mp) C) water solubility increases with polarity, hydrogen-bondingV. Isomerism (2-8) A) structural (constitutional): same molecular formula, different connectivity B) cis-trans (stereoisomers): structures vary only by orientation in spaceVI. Intro to Functional Groups (FG) (2-12 – 2-14)

Suggested Textbook Problems: see syllabus www.csupomona.edu/~lsstarkey/courses/CHM314

2-8

• hybrid orbitals (sp, sp2, sp3) contain lone pairs or are used to form σ bonds• p orbitals are used to form π bonds, may contain lone pairs (for resonance) or may be empty

Functional Group Example Abbreviation Namealkane methane

alkyl halide RX or RCl chloromethane(methyl chloride)

alkene ethene(ethylene)

alkyne

alcohol methanol(methyl alcohol)

ROR or R2O methoxymethane(dimethyl ether)

amine methanamine(methyl amine)

aldehyde

ketone RCOR or R2CO

2-propanone(acetone)

ethanoic acid(acetic acid)

acid chloride(acyl halide)

ethanoyl chloride(acetyl chloride)

ester methyl ethanoate(methyl acetate)

amide RCONR2 ethanamide(acetamide)

anhydride RCO2COR or(RCO)2O

ethanoic anhydride(acetic anhydride)

nitrile ethanenitrile(acetonitrile)

CH4

CH3Cl

H2C CH2

HC CH

CH3OH

CH3OCH3

CH3NH2

CO

HCH3

CO

CH3CH3

CO

OHCH3

CO

ClCH3

CO

OCH3CH3

CO

NH2CH3

CO

CH3 CO

O CH3

CH3CN

carboxylic acid

ethanal(acetaldehyde)

R3N

ether

ethyne(acetylene)

aromatic benzeneArH

RH

RCCR

ROH

RCHO

RCO2H

RCOCl

RCO2R

RCN

R2CCR2

CHM

314

CH

M 3

15C

HM 3

16