1. Chemistry 141 Monday, November 27, 2017 Lecture 33 Pi bonding and MOs
Chemistry 11 - Lecture 11
9/30/2009
Chemistry 141
Monday, November 27, 2017
Lecture 33
Pi bonding and MOs

2. Questions for today:
How can we understand multiple bonds in terms of orbitals?
How can we explain more complex molecular properties such as light absorption and magnetism?

3. Bonding in ethylene Ethane Ethylene Experimental data on ethylene:
Bond length is shorter than in ethane
Bond energy is larger than in ethane (346 kJ/mol), but by less than a factor of 2
Molecule is planar and rotation about the C-C bond is not allowed
Ethylene
Ethane
bond bond bond
length energy angles
(Å) (kJ/mol)
C-C °
C-H
bond bond bond
length energy angles
(Å) (kJ/mol)
C-C °
C-H
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4. Sigma () and Pi () Bonds
Sigma bonds are characterized by
head-to-head overlap.
cylindrical symmetry of electron density about the internuclear axis.
Pi bonds are characterized by
side-to-side overlap.
electron density above and below the internuclear axis.

5. Bonding in Ethylene

6. Bonding in Acetylene

7. Benzene

8. Lewis/VSEPR/Hybridization
Successes:
Simplicity!
Lewis structures rationalize atomic connectivity, bond order
VSEPR predicts molecular shape, bond polarity
Hybridization rationalizes bond length, rotation
Limitations:
Resonance structures
Excited states and transitions (colors, etc.)
Magnetism

9. Magnetism Diamagnetic – unaffected by magnetic field (no unpaired e–)
Chemistry 11 - Lecture 28
9/8/2018
Magnetism
Diamagnetic – unaffected by magnetic field (no unpaired e–)
Paramagnetic – attracted to magnetic field (has unpaired e–)
Prediction:
N2 predict diamagnetic observe diamagnetic
O2 predict diamagnetic observe paramagnetic N N O O
Liquid O2 is poured between the poles of a strong magnet
Paramagnetic O2 is held in place by the magnet
Liquid O2 will condense from the air in the presence of liquid N2 (Tbp = –195.8 °C).
Don’t try this at home… liquid O2 is strongly oxidizing and thus quite dangerous!

10. Molecular Orbital Theory
Two models of bonding
Hybridization
Simple
Hybridization of valence shell s,p, and d orbitals on one atom to form new atomic orbitals with the correct geometry
Atomic orbitals on adjacent pairs of atoms overlap to form bonds
Electrons are shared by only 2 atoms
Molecular Orbital Theory
Simple only for diatomic molecules
Atomic orbitals on all atoms combine to form molecular orbitals
Electrons fill molecular orbitals; bonds form when there are more e- in bonding than antibonding orbitals
Electrons delocalized over molecule
Doesn’t handle resonance
excited states O2 paramagnetism
Does handle resonance
excited states O2 paramagnetism

11. Forming molecular orbitals

12. Bond order in molecular orbitals
A bond is defined as a configuration where the bond order is greater than zero.
We define the bond order as follows:

13. Bonding in H2 and He2 and their ions (1s MO)
Configuration
# bonding e–
# anti-bonding e–
Bond order
Bond energy (kJ/mol)
Bond length (Å)
H2+
σ1s(1) 1
1/2
255
1.06 H2
σ1s(2) 2
432
0.76
He2+
σ1s(2)σ*1s(1)
322
1.08
He2
σ1s(2)σ*1s(2)
Note that any bond order greater than 0 implies that a bond can form!

14. MOs from p orbitals • • • • • • • • • • • • • • • • + – + – + – – + +
Choose the z-axis to be the bond axis
bonding combinations
antibonding combinations
non-bonding combinations • + – + • –
s-type interaction • + – – • + • + – • pz – + pz pz pz pz py • + – • + – • + – –
p-type interaction • • + – • + – + py py py py px py • + – • + – • + – –
p-type interaction • + px px px px

15. MOs formed from p orbitals
What do the MOs look like? • + –
σ*2p
anti-bonding
out of phase: • + – • + –
2pz
2pz • + –
σ2p
bonding
in phase: • + – • + –
2pz
2pz • + –
anti-bonding
π*2p • + – • + –
out of phase:
2py
2py • + –
bonding
π2p • + – • + –
in phase:
2py
2py
There is an identical set of orbitals from the 2px perpendicular to the 2py

16. Energy levels for 2p MOs σ*2p π*2p 2p 2p π2p σ2p σ*2s 2s 2s σ2s XA X2XB

17. s-p Orbital Interactions

18. MO Diagrams for 2nd Period Diatomics

19. Heteronuclear diatomics: NO
σ*2p
Example: NO (5 + 6 = 11 valence e–)
Electron configuration:
(σ2s)2(σ*2s)2(σ2p)2(π2p)4(π*2p)1
Where is the unpaired electron?
MO scheme suggests more density on N
This is in agreement with the Lewis structure:
π*2p 2p
π2p 2p
σ2p
σ*2s 2s 2s
σ2s N NO O
N O : .

20. Molecular Orbital Theory
Two models of bonding
Hybridization
Simple
Hybridization of valence shell s,p, and d orbitals on one atom to form new atomic orbitals with the correct geometry
Atomic orbitals on adjacent pairs of atoms overlap to form bonds
Electrons are shared by only 2 atoms
Molecular Orbital Theory
Simple only for diatomic molecules
Atomic orbitals on all atoms combine to form molecular orbitals
Electrons fill molecular orbitals; bonds form when there are more e- in bonding than antibonding orbitals
Electrons delocalized over molecule
Doesn’t handle resonance
excited states O2 paramagnetism
Does handle resonance
excited states O2 paramagnetism