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California State Polytechnic University, Pomona Dr. Laurie S. Starkey, Organic...
Transcript of California State Polytechnic University, Pomona Dr. Laurie S. Starkey, Organic...
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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