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CHM 233 Review
Basic Principles Part A
1 Energies of Electrons in Atoms and in Covalent Bonds
• Bonding, structure, shape and reactions of organic molecules are determined PRIMARILY by the Energies
of the electrons in ATOMIC and MOLECULAR ORBITALS
Electrons in ATOMS (energies in eV, i.e. electron Volts)
energy of an electron that is infinitely far from any nucleus (not stabilized by any nucleus)
IP ~ 5.4
Energy
IP ~ 13.6
IP ~ 11.3
x
2s
x
x
1s
1s
2s
Li
1s22s
IP ~ 14.5
x
2px
H
1s
IP ~ 13.6
2py
x
x x
x
2px
2pz
x x
1s
Р "core" electrons, not involved in reactions/bonding
Р "valence" electrons, are involved in reactions/bonding
x x
x
2py
2px
2s
C
1s22s22p2
x
2py
2s
1s
x
1s
N
1s22s22p3
x x
x
2pz
O
1s22s22p4
• Electrons that are held "less tightly" by the nucleus are HIGH in energy, and thus require less energy to remove
from an atom, and thus have a low IP
• I.P. decreases (electron energy increases) with increasing orbital size (e.g. down the periodic table)
• I.P. increases (electron energy decreases) with increasing electronegativity (e.g. left to right in the table)
increasing electronegativity
hydrogen
1
H
beryllium
3
4
Li Be
6.941
9.012
sodium
magnesium
11
12
24.306
potassium
calcium
19
20
K Ca
39.098
40.078
rubidium
strontium
37
38
Rb Sr
4.0026
boron
decreasing electron energy
increasing IP
Na Mg
22.990
2
He
GENERALLY
1.0079
lithium
helium
scandium
titanium
vanadium
21
22
23
Sc Ti V
44.956
47.867
50.942
yttrium
zirconium
40
39
Y
carbon
nitrogen
oxygen
fluorine
6
7
8
9
10
C
N
O
F
Ne
argon
10.811
12.0107
14.007
15.999
18.998
aluminium
silicon
phosphorus
sulfur
chlorine
13
14
15
16
17
18
Si P
S
Cl
Ar
Al
26.912
28.086
30.974
32.067
35.453
chromium
manganese
iron
cobalt
nickel
copper
zinc
gallium
germanium
arsenic
selenium
bromine
24
25
26
27
28
29
30
31
32
33
34
35
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
58.693
63.546
65.39
69.723
72.61
74.922
78.96
79.904
niobium
molybdenum
technetium
ruthenium
rhodium
palladium
silver
cadmium
indium
tin
antimony
tellurium
iodine
41
42
43
44
45
46
47
48
49
50
51
52
53
54
I
Xe
radon
In Sn Sb Te
87.62
91.224
92.906
95.94
[98.91]
101.07
85.468
106.42
107.87
112.41
114.818
118.71
121.760
127.60
126.904
barium
lutetium
hafnium
tantalum
tungsten
rhenium
osmium
iridium
platinum
gold
mercury
thallium
lead
bismuth
polonium
astatine
56
71
72
73
74
75
76
77
78
79
80
81
82
83
84
85
137.32
Lu Hf Ta W Re Os Ir
174.97
178.49
36
39.098
55
Cs Ba
krypton
Kr
55.845
Zr Nb Mo Tc Ru Rh Pd Ag Cd
180.95 183.84
Bonding and Structure
186.21
190.23
132.905
Pt Au Hg
195.08
196.97
1
200.59
Tl Pb Bi Po At
204.383
207.2
208.980
[208.98]
[209.99]
increasing
electronegativity
GENERALLY
39.948
54.938
cesium
132.905
20.180
51.996
85.468
88.906
neon
5
B
83.80
xenon
decreasing energy
increasing IP
131.29
86
Rn
[222.08]
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Electrons in a simple MOLECULE
infinity
Energy
IP ~ 13.6
1s atomic orbital
H atom
IP ~ 15.4
lower energy IN A BOND
 molecular orbital
H2 molecule
• IP in MOLECULAR hydrogen is LARGER than in atomic hydrogen
• the electrons in MOLECULAR hydrogen are thus LOWER in energy
Question. Why are the electrons lower in energy in molecular hydrogen compared to H atom?
Answer Because they are IN A BOND - this is really important
HҐ
H
H atom
H
H2 molecule, covalent bond
each electron
electron stabilized
stabilized by two nuclei
by one nucleus
1s atomic orbital
 molecular orbital
• In the molecule, the nuclei are shielded from each other by the two electrons
• In the molecule there is an electrostatically stable configuration for the two negatively and two positively
charged particles (the electrons and the protons)
The following represent the TWO MOST CRITICAL CONCEPTS for UNDERSTANDING organic chemistry:
1. Forming bonds stabilizes (lowers the energies of) electrons
2. Higher energy electrons are MORE CHEMICALLY REACTIVE
2 Energy Diagram for Bond Formation: Simplest Reaction (more...)
When we allow two H atoms to make a bond, what happens to the overall energy?
half arrow means: "this electron moves here to make a
new bond (when the atoms get close enough!)"
H H
Lewis "dot" structure
two
bring them
or
hydrogen
Lewis or Kekule structure
together
H
H
H•
•H atoms
hydrogen molecule
Energy
energy
increases
again when
nuclei get
too close
together
HҐ
ҐH
(bond broken)
two atoms, higher in energy
electron energy DECREASE as bond FORMS
bond
dissociation
energy (BDE)
electron energy INCREASE as bond BREAKS
~104 kcal/mol
for H2
H H
molecule, electrons lower in energy
equilibrium bond length
~0.8Å for H2
Bonding and Structure
2
H-H separation distance
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3 Energies of Electrons in Molecules
Example Problem Give the relative energies of the indicated electron pairs
H A non-bonding electrons
H
B -bonding electrons
O
C
H
H
H
C
C non-bonding electrons
C
C
C
C
H
H H
H
O
Energy
D -bonding electrons
C. non-bonding electrons on atom with formal negative charge clearly highest
A. non-bonding electrons higher than any bonding electrons
B. -bonding orbital "further" from nucleus than s-bonding orbital
D. -bond "made from" 2 sp3 A.O.s, larger more "p-like" orbital, higher energy electrons
4 Bond Strengths
Bond Dissociation Energy (BDE) defined as energy required to break a bond HOMOLYTICALLY
A
A
B
B
homolysis
The energy required to Break a Bond HETEROLYTICALLY a bond is very different, and NOT = BDE
A
A
B
BH2
heterolysis
Some HOMOLYTIC Bond Dissociation Energies and Bond Lengths
(do not memorize these numbers, but be able to recognize the trends)
• as usual, think about the energies of the electrons BEFORE, i.e. in the bond, and AFTER, i.e. in the radicals that
are formed upon homolytic cleavage
• IT COSTS ENERGY TO BREAK A BOND: after the bond is broken, the ELECTRONS are NOT IN A BOND!
1. Across the periodic table : Effect of electronegativity
Length (Ѓ )
1.09
H3C
H
H3C
H
B.D.E. (kcal/mol)
104
H2N
H
H2N
H
103
1.01
HO
H
HO
H
~119
0.96
F
H
F
H
~135
0.95
• bonds to more electronegative elements are stronger and shorter
2. Down the periodic table : Effect of atomic size
Length (Ѓ )
0.92
F
H
F
H
B.D.E. (kcal/mol)
134
Cl
H
Cl
H
103
1.27
Br
H
Br
H
88
1.41
I
H
I
H
71
1.61
• bonds to larger atoms are weaker and longer
• F is very electronegative and small, the electrons in the H-F bond are very low in energy, this is a
strong short bond. With increasing atomic size and decreasing electronegativity, there is poorer A.O.
Bonding and Structure
3
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overlap in bond, larger orbitals, less electrostatic stabilization, the energy of the electrons in the bonds
goes up, the bonds get weaker and longer.
3. Carbon-Hydrogen Bonds : Stabilization of the RADICAL by Resonance
H3C
H2C
H3C
H
H
CH
CH2
H2C
H2C
CH
CH
B.D.E.
kcal/mol
104
H
CH2
H
~87
Length
Ѓ
1.09
~1.07
CH2
• resonance stabilized radical is easier to form, the non-bonding electron is lower in energy (it is
stabilized be delocalization, resonance), thus the energy required to break the bond is lower, the B.D.E.
is lower
higher energy electron
higher BDE
H3CҐ + ҐH
lower energy electron
Relative
smaller BDE
Energy
CH
CH
+ H
H2C
CH2
H2C
CH2
costs this
much energy
to break
RРH
these bonds
energies normalized here
rRРH
•IMPORTANT: it is resonance stabilization in the RADICAL that is formed upon bond fragmentation,
not resonance in the bond (there is none!) that lowers the B.D.E.
4. Carbon-Hydrogen Bonds : Stabilization of the RADICAL by HYPERCONJUGATION
H3C
H
H3CРH2C
H
H3C
HC
H
H3C
H3CРC
H3C
H
H3C
H3C
H
H3C–H2C
H
1° radical
H3C
H
H3C HC
2° radical
H3C
H
H3C–C
3°
radical
H3C
B.D.E.
kcal/mol
104
Length
Ѓ
1.09
~98
~1.09
~95
~1.08
~91
~1.07
• increasing radical stability upon substitution at carbon due to hyperconjugation results in lower energy
electrons in the radical, thus lower B.D.E.
• costs less energy to form radicals that are more stable
Bonding and Structure
4
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• What is HYPERCONJUGATION?
H
methyl
radical
ethyl
radical
C
CH3
half-filled p orbital
H
sp2 hybridized
H
H
CH2 CH3
C
H
C
H
H
H
H
H
H
C
C
H
H
H
H
H
C
C H
hyperconjugation
H
increasing stability
(decreasing energy of the non-bonding electron)
H
R
R
H 2C C
R C
< CH2–R < CH
<
~
CH2
R ~ same
R
primary
secondary
resonance stabilized allyl
tertiary
CH3
methyl
• radicals are stabilized by alkyl substituents on the carbon bearing the non-bonding electron by hyperconjugation
• in a tertiary radical, the stabilization can be as good a real conjugation
5. Carbon-Hydrogen Bonds : Effect of C Hybridization
H3C
H
H3C
H
H
H
H
C C
H
C
H
sp2
H
C C
H
sp3
C
H
H
H
H
C
C
Length
Ѓ
1.09
B.D.E.
kcal/mol
104
108
sp
H
133
~1.08
~1.055
• the electrons in a bond from H to a sp2 and sp hybridized carbons are lower in energy than those in a
bond to a sp3 hybridized carbon, due to smaller orbitals, it is thus harder to break the bond, thus larger
B.D.E. and stronger bond means shorter bond
6. Carbon-Carbon Bonds
H3C
H3C
H
H3C
CH3
H3C
B.D.E. (kcal/mol) Length (Ѓ )
1.09
104
~88
~1.54
CH3
H
• Overlap of a sp3 A.O. with another sp3 A.O. in a C-C bond is poorer that overlap between a sp3 and a
1s A.O. in a C-H bond, poorer overlap means weaker and longer bonds
7. Weaker Bonds
repulsion HO
OH
HO
repulsion Br
Br
HO
OH
B.D.E. (kcal/mol)
~51
OH
46
Length (Ѓ )
1.48
2.28
• Oxygen lone pairs repel each other in the peroxide, electrons are higher in energy in the bonded
state, thus a weak and long bond, even more so for the larger bromine atoms
Bonding and Structure
5
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8. Multiple Bonds
B.D.E. (kcal/mol) Length (Ѓ )
~88
~1.54
H3C
CH3
H3C
CH3
H2C
CH2
H2C
CH2
~160
~1.33
H2C
O
H2C
O
~175
~1.21
~230
HC
CH
HC CH
• Multiple bonds are obviously stronger (and shorter) than single bonds
• Electronegativity effects are still important
~1.20
Bonding and Structure
6
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