What%20Do%20Molecules%20Look%20Like?

1
What Do Molecules Look Like?
The Lewis Dot Structure approach provides some
insight into molecular structure in terms of
bonding, but what about 3D geometry?

• Recall that we have two types of electron pairs
  bonding and lone.

Valence-Shell Electron-Pair Repulsion (VSEPR).
3D structure is determined by minimizing
repulsion of electron pairs.
2
Electron pairs (both bonding and lone) are
distributed around a central atom such that
electron-electron repulsions are minimized.
3
Electron pairs (both bonding and lone) are
distributed around a central atom such that
electron-electron repulsions are minimized.
2 electron pairs
3 electron pairs
4 electron pairs
Period 1, 2
5 electron pairs
6 electron pairs
Period 3 beyond
4
Arranging Electron Pairs
Must consider both bonding and lone pairs
when minimizing repulsion.

• Example CH4 (bonding pairs only)

Lewis Structure
VSEPR Structure
5
Arranging Electron Pairs (cont.)
Example NH3 (both bonding and lone pairs).
Lewis Structure
VSEPR Structure Noteelectron pair geometry
vs. molecular shape
6
VSEPR Structure Guidelines

• The previous examples illustrate the strategy for
  applying VSEPR to predict molecular structure
• Construct the Lewis Dot Structure
• Arrange bonding/lone electron pairs in space such
  that repulsions are minimized (electron pair
  geometry).
• Name the molecular shape from the position of the
  atoms.

VSEPR Shorthand 1. Refer to central atom as
A 2. Attached atoms are referred to as X 3.
Lone pair are referred to as E
Examples CH4 AX4 NH3 AX3E H2O
AX2E2 BF3 AX3
7
VSEPR 2 electron pairs
Experiments show that molecules with multiple
bonds can also be linear.
Linear (AX2) angle between bonds is 180
Example BeF2
Multiple bonds are treated as a single effective
electron group.
180
More than one central atom? Determine shape
around each.
8
VSEPR 3 electron pairs
Trigonal Planar (AX3) angle between bonds is
120
Multiple bond is treated as a single effective
electron group.
Example BF3
120
9
VSEPR 4 electron pairs (cont.)
Tetrahedral (AX4) angle between bonds is 109.5
Example CH4
109.5
tetrahedral e- pair geometry AND tetrahedral
molecular shape
10
Bonding vs. Lone pairs
Bond angle in a tetrahedral arrangement of
electron pairs may vary from 109.5 due to size
differences between bonding and lone pair
electron densities.
bonding pair is constrained by two nuclear
potentials more localized in space.
lone pair is constrained by only one nuclear
potential less localized (needs more room).
11
VSEPR 4 electron pairs
Trigonal pyramidal (AX3E) Bond angles are
lt109.5, and structure is nonplanar due to
repulsion of lone pair.
Example NH3
107
tetrahedral e- pair geometry trigonal pyramidal
molecular shape
12
VSEPR 4 electron pairs (cont.)
Classic example of tetrahedral angle shift from
109.5 is water (AX2E2)
104.5o
bent
tetrahedral e- pair geometry bent molecular shape
13
VSEPR 4 electron pairs (cont.)
Comparison of CH4 (AX4), NH3 (AX3E), and H2O
(AX2E2)
14
1. Refer to central atom as A 2. Attached
atoms are referred to as X 3. Lone pair are
referred to as E
15
Molecular vs. Electron-Pair Geometry
C
N
O
F
Central Atom
Compound
Electron-Pair Geometry
Molecular Shape
Carbon, C
CH4
tetrahedral
tetrahedral
Nitrogen, N
NH3
tetrahedral
trigonal pyramidal
Oxygen, O
H2O
tetrahedral
bent
Fluorine, F
HF
tetrahedral
linear
16

• What is the electron-pair geometry and the
  molecular shape for HCFS?

1. trigonal planar, bent
2. trigonal planar, trigonal planar
3. tetrahedral, trigonal planar
4. tetrahedral, tetrahedral

17
VSEPR Beyond the Octet
Systems with expanded valence shells will have
five or six electron pairs around a central atom.
90
90
90
120
18
VSEPR 5 electron pairs

• Consider the structure of SF4 (34 e-, AX4E)
• What is the optimum arrangement of electron
  pairs around S?

??
S
S
S
Compare e pair angles
lone-pair / bond-pair
two at 90o, two at 120o
three at 90o
three at 90o, three at 120o
four at 90o, one at 120o
bond-pair / bond-pair
Repulsive forces (strongest to weakest) lone-pair
/lone-pair gt lone-pair/bond-pair gt
bond-pair/bond-pair
19
VSEPR 5 electron pairs
The optimum structure maximizes the angular
separation of the lone pairs. I3- (AX2E3)
20
5-electron-pair geometries
our previous example
21
VSEPR 6 electron pairs
Which of these is the more likely structure?
See-saw
Square Planar
22
6-electron-pair geometries
our previous example
23
Molecular Dipole Moments
We can use VSEPR to determine the polarity of a
whole molecule.

1. Draw Lewis structures to determine 3D arrangement
   of atoms.

2. If one side of the molecule has more EN
atoms than the other, the molecule has a net
dipole.
Shortcut completely symmetric molecules will not
have a dipole regardless of the polarity of the
bonds.
24
Molecular Dipoles
The CO bonds have dipoles of equal magnitude but
opposite direction, so there is no net dipole
moment.
The O-H bonds have dipoles of equal magnitude
that do not cancel each other, so water has a net
dipole moment.
25
Molecular Dipoles (cont.)
symmetric
symmetric
asymmetric
26
Molecular Dipole Example

• Write the Lewis dot and VESPR structures for
  CF2Cl2. Does it have a dipole moment?

32 e-
Tetrahedral
27
Advanced VSEPR Application

• Molecules with more than one central atom
  methanol (CH3OH)

tetrahedral e- pairs tetrahedral shape
tetrahedral e- pairs bent shape
28
The VSEPR Table
e- pairs
e- Geom.
Molec. Geom.
2 AX2 BeF2 linear linear
3 AX3 BF3 trigonal planar trigonal planar
AX2E O3 trigonal planar bent
4 AX4 CH4 tetrahedral tetrahedral
AX3E NH3 tetrahedral pyramidal
AX2E2 H2O tetrahedral bent
29
The VSEPR Table
e- pairs
e- Geom.
Molec. Geom.
5 AX5 PF5 trigonal bipyramidal trigonal bipyramidal
AX4E SF4 trigonal bipyramidal see saw
AX3E2 ClF3 trigonal bipyramidal T-shaped
AX2E3 I3- trigonal bipyramidal linear
6 AX6 SF6 octahedral octahedral
AX4E2 XeF4 octahedral square planar
30
What is the expected shape of ICl2?
AX2E2
20 e-
A. linear
C. tetrahedral
D. square planar
B. bent
31
Valence Bond Theory
Basic Principle of Localized Electron Model A
covalent bond forms when the orbitals from two
atoms overlap and a pair of electrons occupies
the region between the two nuclei. Rule 1
Maximum overlap. The bond strength depends on the
attraction of nuclei to the shared electrons,
so The greater the orbital overlap, the
stronger the bond.
32
Valence Bond Theory
Basic Principle of Localized Electron Model A
covalent bond forms when the orbitals from two
atoms overlap and a pair of electrons occupies
the region between the two nuclei. Rule 2
Spins pair. The two electrons in the overlap
region occupy the same space and therefore must
have opposite spins. There may be no more than 2
electrons in a molecular orbital.
33
Valence Bond Theory
Basic Principle of Localized Electron Model A
covalent bond forms when the orbitals from two
atoms overlap and a pair of electrons occupies
the region between the two nuclei. Rule 3
Hybridization. To explain experimental
observations, Pauling proposed that the valence
atomic orbitals in a molecule are different from
those in the isolated atoms. We call this concept
Hybridization
34
What is hybridization?

• Atoms adjust to meet the needs of the molecule.
• In a molecule, electrons rearrange in an attempt
  to give each atom a noble gas configuration and
  to minimize electron repulsion.
• Atoms in a molecule adjust their orbitals through
  hybridization in order for the molecule to have a
  structure with minimum energy.
• The source of the valence electrons is not as
  important as where they are needed in the
  molecule to achieve a maximum stability.

35
Example Methane

• 4 equivalent C-H covalent bonds
• VSEPR predicts a tetrahedral geometry

36
The Valence Orbitals of a Carbon Atom
Carbon 2s22p2
37
Hybridization Mixing of Atomic Orbitals to form
New Orbitals for Bonding













38
Other Representations of Hybridization
y1 1/2(2s) (2px) (2py) (2pz) y2
1/2(2s) (2px) - (2py) - (2pz) y3 1/2(2s) -
(2px) (2py) - (2pz) y4 1/2(2s) - (2px) -
(2py) (2pz)
39
Hybridization is related to the number of valence
electron pairs determined from VSEPR Methane
(CH4) VSEPR AB4 ? tetrahedral ?
sp3 hybridized
Electron pair geometry determines hybridization,
not vice versa!!
40
Hybridization is related to the number of valence
electron pairs determined from VSEPR
Ammonia (NH3) VSEPR AB3E ? tetrahedral
? sp3 hybridized
41
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42
s bonding and p bonding

• Two modes of bonding are important for
• 1st and 2nd row elements s bonding and p bonding
• These two differ in their relationship to the
  internuclear axis
• s bonds have electron density ALONG the axis

p bonds have electron density ABOVE AND BELOW the
axis
43
Problem Describe the hybridization and bonding
of the carbon orbitals in ethylene (C2H4)
VSEPR AB3 ? trigonal planar ? sp2
hybridized orbitals for s bonding
sp2 hybridized orbitals used for s
bonding remaining p orbital used for p bonding
44
Bonding in ethylene (C2H4)
45

• Problem Describe the hybridization and bonding
  of the carbon orbitals in Carbon Dioxide (CO2)
• VSEPR AB2
• linear
• sp hybridized orbitals for s bonding

46
Bonding in Carbon Dioxide (CO2)
47
Atoms of the same kind can have different
hybridizations
C2 AB4
lone pair
48
What have we learned so far?

• Molecular orbitals are combinations of atomic
  orbitals
• Atomic orbitals are hybridized to satisfy
  bonding in molecules
• Hybridization follows simple rules that can be
  deduced from the number of chemical bonds in the
  molecule and the VSEPR model for electron pair
  geometry

49
Hybridization

• sp3 Hybridization (CH4)
• This is the sum of one s and three p orbitals on
  the carbon atom
• We use just the valence orbitals to make bonds
• sp3 hybridization gives rise to the tetrahedral
  nature of the carbon atom

50
Hybridization

• sp2 Hybridization (H2CCH2)
• This is the sum of one s and two p orbitals on
  the carbon atom
• Leaves one p orbital uninvolved this is free to
  form a p bond (the second bond in a double bond)

51
Hybridization

• sp Hybridization (OCO)
• This is the sum of one s and one p orbital on the
  carbon atom
• Leaves two p orbitals free to bond with other
  atoms (such a O in CO2), or with each other as in
  HCCH

52
General Notes

• This is a model and only goes so far, but it is
  especially helpful in understanding geometry and
  expanding Lewis dot structures.
• Orbitals are waves. Hybridized orbitals are just
  the sums of waves constructive and destructive
  interference.

53
What is important to know about hybridization?

• You should be able to give the hybridization of
  an atom in a molecule based on the formula given.
• Example CH3-CH2-CHO
• Step 1 Draw the Lewis Dot Structure

54
What is important to know about hybridization?

• Step 2 What is the electron pair geometry and
  molecular shape?

AXE2
Trigonal Planar
AX3
AX4
Trigonal Planar
AX4
Tetrahedral
Tetrahedral
55
What is important to know about hybridization?

• Step 3 Use the molecular shape to determine the
  hybridization.

sp2
sp3
sp2
sp3
56
The Localized Electron Model is very powerful for
explaining geometries and basic features
of bonding in molecules, but it is just a model.

• Assumes electrons are highly localized between
  the nuclei (sometimes requires resonance
  structures)
• Doesnt easily deal with unpaired electrons
  (incorrectly predicts physical properties in some
  cases)
• Doesnt provide direct information about bond
  energies

57
The Molecular Orbital Model
Basic premise When atomic orbitals interact to
form a bond, the result is the formation of new
molecular orbitals
HY EY
Important features of molecular orbitals 1.
Atomic Orbitals are solutions of the Schrödinger
equation for atoms. Molecular orbitals are the
solutions of the same Schrödinger equation
applied to the molecule.
58
Molecular Orbital Theory

• 2. Atomic Orbitals can hold 2 electrons with
  opposite spins.
• Molecular Orbitals can hold 2 electrons with
  opposite spins.
• 3. The electron probability for the Atomic
  Orbital is given by Y2.
• The electron probability for the Molecular
  Orbital is given by Y2.
• 4. Orbitals are conserved - in bringing together
  2 atomic orbitals, we have to end up with 2
  molecular orbitals!
• How does this work?

59
Molecular Orbitals are simply Linear Combinations
of Atomic Orbitals
Example H2
s anti-bonding (s)
Molecular Orbitals have phases ( or -)
s bonding
Next Question Why does this work?
60
Constructive and Destructive Interference
Destructive interference between two orbitals of
opposite sign leads to an anti-bonding orbital.
Constructive interference between two overlapping
orbitals leads to a bonding orbital.
61
Bonding is driven by stabilization of electrons

• Electrons are negatively charged
• Nuclei are positively charged

The bonding combination puts electron density
between the two nuclei - stabilization The
anti-bonding combination moves electron density
away from region between the nuclei -
destabilization
62
MO Diagrams

• We can depict the relative energies of molecular
  orbitals with a molecular orbital diagram

The new molecular orbital is lower in energy than
the atomic orbitals
63
s M.O. is raised in energy
s M.O. is lowered in energy
H atom (1s)1 electron configuration H2 molecule
(s1s)2 electron configuration

64
Same as previous description of bonding
65
Review of Orbital Filling

• Pauli Exclusion Principle
• No more than 2 e- in an orbital, spins must be
  paired (??)
• Aufbau Principle (a.k.a. Building-Up)
• Fill the lowest energy levels with electrons
  first
• 1s 2s 2p 3s 3p 4s 3d 4p
• Hunds Rule
• When more than one orbital has the same energy,
  electrons occupy separate orbitals with parallel
  spins

No
No
Yes
66
Filling Molecular Orbitals with Electrons
1) Orbitals are filled in order of increasing
Energy (Aufbau principle)
67
Filling Molecular Orbitals with Electrons
2) An orbital has a maximum capacity of two
electrons with opposite spins (Pauli
exclusion principle)
68
Filling Molecular Orbitals with Electrons
3) Orbitals of equal energy (degenerate orbitals)
are half filled, with spins parallel,
before any is filled completely (Hunds rule)
69
Bond Order
70
Bond Order Examples
Bond order (2-0)/2 1 Single bond Stable
molecule (436 kJ/mol bond)
Bond order (2-2)/2 0 No bond! Unstable
molecule (0 kJ/mol bond)
71
He2
Bond order (2-1)/2 1/2 Half of a single
bond Can be made, but its not very stable (250
kJ/mol bond)
Fractional bond orders are okay!
H2
Bond order (1-0)/2 1/2 Half of a single
bond Can be made, but its not very stable (255
kJ/mol bond)
72
Forming Bonds

• A s bond can be formed a number of ways
• s, s overlap
• s, p overlap
• p, p overlap

Only orbitals of the same phase (, ) can form
bonds
73
Anti-bonding Orbitals

• For every bonding orbital we form, we also form
  an anti-bonding orbital

74
MO Theory in Bonding

• Homonuclear atoms (H2, O2, F2, N2)

H2 (Only 1s orbitals available for bonding)
75
Covalent Bonding in Homonuclear Diatomics

• Atomic orbitals must overlap in space in order to
  participate in molecular orbitals
• Covalent bonding is dominated by the valence
  orbitals (only valence orbitals are shown in the
  MO diagrams)

76
Covalent Bonding in Homonuclear Diatomics
Region of shared e- density
77
Valence configurations of the 2nd row atoms
Li Be B C N O
F 2s1 2s2 2s22p1 2s22p2 2s22p3
2s22p4 2s22p5
So far we have focused on bonding involving the s
orbitals. What happens when we have to consider
the p orbitals?
78
For diatomic molecules containing atoms with
valence electrons in the p orbitals, we must
consider three possible bonding interactions
nucleus
p-type
p-type
s-type
79
() destructive mixing
() constructive mixing
80
2) Doesnt easily deal with unpaired electrons
(incorrectly predicts physical properties in some
cases)
Experiments show O2 is paramagnetic
81
A quick note on magnetism
Paramagnetic The molecule contains unpaired
electrons and is attracted to (has a positive
susceptibility to) an applied magnetic field
Diamagnetic The molecule contains only paired
electrons and is not attracted to (has a negative
susceptibility to) an applied magnetic field
82
Example the O2 Diatomic
Oxygen atom has a 2s22p4 valence configuration
M.O. O2
O atom
O atom
____ ?2p
___ ___ ?2p
___ ___ ___2p ___ ___ ___ 2p
___ ___ ?2p
Energy ____ ?2p

____ ?2s
___ 2s ___ 2s
____ ?2s
Bond Order (8-4)/2 2 O2 is stable (498
kJ/mol bond strength)
(s2s)2(s2s)2(s2p)2(p2p)4(p2p)2
83
A prediction from the M.O. diagram of O2
The Lewis dot structure predicts O2 should be
diamagnetic-all electrons are paired.
The unpaired electrons predicted by the M.O.
diagram should behave as small magnets- O2
should be magnetic!
84
N2 Video
O2 Video
85
What have we learned so far?
1. Molecular orbitals (MO) are linear
combinations of atomic orbitals 2. Both s and p
atomic orbitals can be mixed to form MOs 3.
Molecular orbitals are bonding and
anti-bonding 4. Bonding and anti-bonding MOs
lead to the definition of the bond order 5. Bond
order is related to the bond strength (bond
dissociation energy)

86
MO Diagram for H2 vs. N2
Atomic orbital overlap sometimes forms both ? and
? bonds. Examples N2, O2, F2
87
M.O. Diagram for N2
88
A Complication
M.O. Diagram for O2 (similar for F2 and Xe2)
M.O. Diagram for B2 (similar for C2 and N2)
O
O
O2
89
A Complication
M.O. Diagram for O2 (similar for F2 and Ne2)
M.O. Diagram for B2 (similar for C2 and N2)
No s-p mixing
s-p mixing
90
Why does s-p mixing occur?
Electron repulsion!!
s2s and s2p both have significant e- probability
between the nuclei, so e- in s2s will repel e- in
s2p
Effect will decrease as you move across the
Periodic Table ? increased nuclear charge pulls
the s2s e- closer, making the s2s orbital smaller
and decreasing the s2s and s2p interaction
91
Molecular Orbitals of X2 Molecules

• s?p orbital mixing (a little hybridization)
• lowers the energy of the ?2s orbitals and
• raises the energy of the ?2p orbitals.
• As a result, E(?2p) gt E(? 2p) for B2, C2, and
  N2.
• As one moves right in Row 2, 2s and 2p get
  further apart in energy, decreasing sp mixing
  ? E(?2p) lt E(?2p) for O2, F2, and Ne2. See
  text pages 680-681.
• Note that sp mixing does not affect bond order
  or magnetism in the common diatomics (N2, O2, and
  F2). Hence it is not of much practical
  importance.

92
No s-p mixing
s-p mixing
93
When does s-p mixing occur?
B, C, and N all have ? 1/2 filled 2p orbitals
O, F, and Xe all have gt 1/2 filled 2p orbitals

• If 2 electrons are forced to be in the same
  orbital, their energies go up.
• Electrons repel each other because they are
  negatively charged.
• Having gt 1/2 filled 2p orbitals raises the
  energies of these orbitals due to e- - e-
  repulsion

94
Relating the M.O. Diagrams to Physical Properties
95
Sample Problem
Using MO Theory to Explain Bond Properties
Problem Consider the following data for these
homonuclear diatomic species
N2
N2 O2 O2 Bond
energy (kJ/mol) 945 841
498 623 Bond length (pm)
110 112 121
112 No. of valence electrons 10
9 12 11
Removing an electron from N2 decreases the bond
energy of the resulting ion, whereas removing an
electron from O2 increases the bond energy of the
resulting ion. Explain these facts using M.O.
diagrams.
96
Sample Problem
Using MO Theory to Explain Bond Properties
Problem Consider the following data for these
homonuclear diatomic species
N2
N2 O2 O2 Bond
energy (kJ/mol) 945 841
498 623 Bond length (pm)
110 112 121
112 No. of valence electrons 10
9 12 11
97
Sample Problem - Continued
Solution The MO energy levels are
N2
??p
?2p
?2p
?2p
?2s
?2s
Bond Orders
(8-2)/2 3
(7-2)/2 2.5
(8-4)/2 2
(8-3)/2 2.5
98
Sample Problem
Using MO Theory to Explain Bond Properties
Problem Consider the following data for these
homonuclear diatomic species
N2
N2 O2 O2 Bond
energy (kJ/mol) 945 841
498 623 Bond length (pm)
110 112 121
112 No. of valence electrons 10
9 12 11 Bond Order
3 2.5 2 2.5
99
What have we learned so far?
1. Molecular orbitals (MO) explain the properties
of valence electrons in molecules (Example
O2) 2. s and p atomic orbitals can be mixed to
form s, s, p, and p molecular orbitals 3.
Electrons in p or p molecular orbitals can have
the same energies Degenerate orbitals 4. The
ordering of s2p and p2p molecular orbitals
depends on the electron occupancy s-p mixing
100
Bonding in Diatomic Molecules
Covalent
Ionic
Covalent
Ionic
101
Homonuclear H2
Nonpolar covalent bond (450 kJ/mol bond)
102
Electrons are not equally shared in heteronuclear
bonds
HF
Because F (EN 4.0) is more electronegative than
H (EN 2.2), the electrons move closer to F.
This gives rise to a polar bond
Electronegativity
H F
Figure 14.45
103
M.O.s of a Polar Covalent Bond HF
s Antibonding (s) Mostly H(1s)
H F
H F
This approach simplifies model and only considers
electrons involved in bond.
s Bonding Mostly F(2p)
104

• MOs OF XY MOLECULES
• Equal or unequal e? sharing between 2 atoms is
  reflected in the composition of the MOs
• When 2 atoms X and Y have the same
  electronegativity (purely covalent bond), their
  overlapping AOs have the same energy, and the
  bonding and anti-bonding MOs are each half X and
  half Y AO. All electrons spend equal time near X
  and Y. Examples N2, O2, F2.
• If EN(Y) gt EN(X) (polar covalent X??Y??), the Y
  AO has lower energy than the X AO. The bonding
  MO is more like the Y AO and the anti-bonding MO
  more like the X AO. Bonding e? spend more time
  near Y than X vice versa for anti-bonding e?.
  Example CO.

105
MOs OF XY MOLECULES
Electronegativity
C Atom (4e) CdOd (10e) O Atom
(6e)

• CO Bond Order 3.0 (same as N2).
• CO Bond Energy 1,076 kJ/mol (N2 945
  kJ/mol).
• Isoelectronic to CO and N2 CN, NO.
• NO has 1e in ? ? bond order 2.5 this
  e is more on N than O NO ? NO easy

106
Bonding in NO

• Two possible Lewis dot structures for NO
• The simplest structure minimizes formal charges
  and places the lone (unpaired) electron on the
  nitrogen.
• The Lewis structure predicts a bond order of 2,
  but experimental evidence suggests a bond order
  between 2 and 3.
• How does MO theory help us understand bonding in
  NO?

107
When the electronegativities of the 2 atoms are
more similar, the bonding becomes less polar.
2p
2p
Electronegativity
EN(N) 3.0 EN(O) 3.4
2s
2s
NO
N
O
Bond order 2.5, unpaired electron is in a
N-like orbital
108
NO is easily oxidized to form NO. Why? What
changes can we predict in the bonding and
magnetism of the molecule?
NO
NO
oxidation
Bond Order (8-3)/2 2.5 Paramagnetic
Bond Order (8-2)/2 3 Diamagnetic
109
M.O. diagram for NO
p2p
p2p (empty)
-597
p2p
p2p
-1444
-1374
s2s
-1835
s2s
-3320
110
Key Points of MO Theory Heteronuclear Molecules

• The more electronegative atom has orbitals lower
  in energy than the more positive atom.
• Electrons in bonding orbitals are closer to the
  more electronegative atom, anti-bonding electrons
  are closer to the more positive atom.
• For most diatomic molecules, s-p mixing changes
  the orbital energy levels, but since these
  orbitals are almost always fully occupied, their
  order is less important to us.

111
Combining the Localized Electron and Molecular
Orbital Models (into a convenient working model)
Figure 14.47
Only the p bonding changes between these
resonance structures - The M.O. model describes
this p bonding more effectively.
112
Atomic Orbitals Molecular Orbitals
Figure 14.51
113
Another example Benzene
s bonding
p bonding
114
MO Theory Expectations

• You should be able to
• predict which atomic orbitals are higher or lower
  in energy (based on electronegativity
  differences).
• correctly fill a molecular orbital diagram.
• correctly calculate bond order.
• predict molecular magnetic properties based on
  orbital occupation.
• understand how molecular properties change upon
  ionization (oxidation or reduction) of molecules.