1. Chapter 7.4 – Molecular Orbital (MO) Theory Continued
CHM1111 Section 04
Instructor: Dr. Jules Carlson
Class Time: M/W/F 1:30-2:20
Monday, November 7th

2. Housekeeping Final Exam is Monday, December 9th
The fall session lectures finish Tuesday November 29th, but we have Lectures on Wednesday, November 30th and Thursday December 1st (1:30 – 2:20 PM) to make up for the holidays (Thanksgiving Day + Remembrance Day).
At least half the lecture on Thursday December 1st will be for review but I would like to schedule another review with you. We can talk more later about when to do this.

3. Different Models for Describing Different Properties
Started with Orbital Overlap Model – described how wavefunctions overlapped and in some cases what molecular structures looked like.
Then added VSEPR theory to show what molecules looked like when simple orbital overlap model didn’t work (although sometimes it does not apply i.e. H2S).
In order to explain why molecular shapes in VSEPR make sense, we introduced hybridization theory.
Now to describe bond strength, and magnetism, we need MO Theory.

4. Filling and Energy of Molecular Orbitals for H2
H2 forms σ1s bonding and σ1s* antibonding orbitals
Aufbau principle, Pauli Exclusion principle and Hund’s rule all still apply for electron filling.
Remember antibonding orbitals are higher in energy than bonding orbitals.

5. Rules for Building and Filling Molecular Orbitals
The total number of molecular orbitals produced by a set of interacting atomic orbitals is always equal to the number of interacting atomic orbitals.
The bonding molecular orbital is lower in energy than the parent orbitals and the antibonding orbital is higher in energy.
Electrons of the molecule are assigned to orbitals of successively higher energy according to the aufbau principle and Hund’s Rule.

6. Bond Order
Bond order – describes the relative effect of bonding and antibonding orbitals.
Bond Order = ½(number of electrons in bonding MOs – number of electrons in antibonding MOs).
Bonding is stable if Bond Order > 0.
Antibonding orbitals indicated with a *.
He2 Bond Order = 0
H2 Bond Order = 1

7. MO Diagrams for Diatomics – Li2, Be2

8. Adding p Orbitals to MO Diagrams
Note: x, y and z directions are not defined the same way in Chapter 4, now z is along the bond axis.
There are three p orbitals in directions x, y, z.
The z direction is defined as along the bond axis – the MO formed with two pz orbitals is σ as it lies along the bond axis.
Both x and y directions are at 90⁰ to the bond axis, the MO formed with two px or two py orbitals is π as it lies above and below the bond axis.

9. MO Diagram for B2 Bond Order = 1 Bond Energy = 280 kJ/mol σ2pz* π2px*
Electron Configuration:
(σ1s)2 (σ1s*)2 (σ2s)2 (σ2s*)2 (π2px)1 (π2py)1
σ2pz*
π2px*
π2py*
2px
2py
2pz
σ2pz
2px
2py
2pz
π2px
π2py
σ2s*
Energy
σ2s 2s 2s
σ1s*
σ1s 1s 1s

10. Paramagnetism and Diamagnetism
Paramagnetism – Exhibited when a compound attracted to a strong magnetic field. Occurs when a molecule has unpaired electrons. B2 is paramagnetic.
Diamagnetism – Exhibited when a compound is repelled by a strong magnetic field. Occurs when a molecule has all paired electrons, H2 is diamagnetic.

11. MO Diagram for C2 Bond Order = 2 Bond Energy = 602 kJ/mol σ2pz* π2px*
Electron Configuration:
(σ1s)2 (σ1s*)2 (σ2s)2 (σ2s*)2 (π2px)2 (π2py)2
σ2pz*
π2px*
π2py*
2px
2py
2pz
σ2pz
2px
2py
2pz
π2px
π2py
σ2s*
Energy
σ2s 2s 2s
σ1s*
σ1s 1s 1s

12. MO Diagram for N2 Bond Order = 3 Bond Energy = 945 kJ/mol σ2pz* π2px*
Electron Configuration:
(σ1s)2 (σ1s*)2 (σ2s)2 (σ2s*)2 (π2px)2 (π2py)2 (σ2pz)2
σ2pz*
π2px*
π2py*
2px
2py
2pz
σ2pz
2px
2py
2pz
π2px
π2py
σ2s*
Energy
σ2s 2s 2s
σ1s*
σ1s 1s 1s

13. I Clicker Question Which of the following statements is INCORRECT?
Li2 is diamagnetic.
N2+ has a bond order of 1.5.
The Bond Energy of Li2 is larger than for Be2.
The sum of energy levels of molecular orbitals must equal the sum of energy of atomic orbitals for the atoms combining to form the molecule.
More than one statement is incorrect.

14. Antibonding p orbitals
As we saw with s orbitals, we can have additive or subtractive overlap with p orbitals
Remember pz orbitals are aligned along the bond axis.
Remember px and py orbitals are aligned at 90 angles to the bond axis.

15. Review Slide of MO shapes
Additive overlap of p σ bonds
Subtractive overlap of p σ bonds
Additive overlap of p π bonds
Subtractive overlap of p π bonds

16. Orbital Mixing Effects
2s and 2p orbitals do not act independently of each other.
2s and 2p orbitals have similar radii so the 2s and 2pz orbitals of one atom overlap with both the 2s and 2pz orbitals of the other atom.
This destabilizes σ2pz (high electron density in small region) and stabilzes σ2s
Larger atoms have larger nuclear charge, reduces mixing as higher electron density can be accepted from higher nucleus-electron attraction.
Therefore we need a different orbital filling diagram for Z > 7.

17. MO Diagram for O2 Bond Order = 2 Bond Energy = 495 kJ/mol σ2pz* π2px*
Electron Configuration:
(σ1s)2 (σ1s*)2 (σ2s)2 (σ2s*)2 (σ2pz)2
(π2px)1 (π2py)1
σ2pz*
π2px*
π2py*
2px
2py
2pz
π2px
π2py
2px
2py
2pz
O2 is paramagnetic, this could not be explained without antibonding orbitals.
σ2pz
σ2s*
Energy
σ2s 2s 2s
σ1s*
σ1s 1s 1s

18. MO Diagram for F2 Bond Order = 1 Bond Energy = 155 kJ/mol σ2pz* π2px*
Electron Configuration:
(σ1s)2 (σ1s*)2 (σ2s)2 (σ2s*)2 (σ2pz)2
(π2px)2 (π2py)2
σ2pz*
π2px*
π2py*
2px
2py
2pz
π2px
π2py
2px
2py
2pz
σ2pz
σ2s*
Energy
σ2s 2s 2s
σ1s*
σ1s 1s 1s

19. MO Diagram for Ne2 Bond Order = 0 Bond Energy ~ 0 kJ/mol σ2pz* π2px*
Electron Configuration:
(σ1s)2 (σ1s*)2 (σ2s)2 (σ2s*)2 (σ2pz)2
(π2px)2 (π2py)2 (σ2pz*)2
σ2pz*
π2px*
π2py*
2px
2py
2pz
π2px
π2py
2px
2py
2pz
σ2pz
σ2s*
Energy
σ2s 2s 2s
σ1s*
σ1s 1s 1s

20. MO Diagram for Heteronuclear Species – Carbon Monoxide
Let’s look at CO:

21. MO Diagram for Heteronuclear Species – Nitrogen Oxide
Let’s look at NO: