1. Problems with Valence Bond Theory
VB theory predicts many properties better than Lewis Theory
bonding schemes, bond strengths, bond lengths, bond rigidity
however, there are still many properties of molecules it doesn’t predict perfectly
magnetic behavior of O2 1 1

2. Aurora Borealis
Chapter 9 | Slide 2
Chapter 9 | Slide 2

3. Valence Bond Theory
Valence Bond Model of covalent bonding is easy to visualize, but it does have some problems:
Incorrectly assumes that electrons are localized and so we have to use resonance to describe some molecules.
Does not do a good job of describing molecules containing unpaired electrons
Does not indicate bond energies 3 3

4. Valence Bond Theory
Valence Bond Model does not explain why O2 is attracted to a magnetic field while N2 is slightly repelled nor accounts for the emission of light by molecules in an aurora.
The need to explain the magnetic behavior seen for O2 led to the development of another bonding theory called the Molecular Orbital (MO) Theory. 4 4

5. Molecular Orbital Theory
in MO theory, we apply Schrödinger’s wave equation to the molecule to calculate a set of molecular orbitals
in practice, the equation solution is estimated
we start with good guesses from our experience as to what the orbital should look like
then test and tweak the estimate until the energy of the orbital is minimized
in this treatment, the electrons belong to the whole molecule – so the orbitals belong to the whole molecule
unlike VB Theory where the atomic orbitals still exist in the molecule 5 5

6. LCAO
the simplest guess starts with the atomic orbitals of the atoms adding together to make molecular orbitals – this is called the Linear Combination of Atomic Orbitals (LCAO) method
weighted sum
because the orbitals are wave functions, the waves can combine either constructively (additive) or destructively (subtractive) 6 6

7. Bonding in H2
9.2
Chapter 9 | Slide 7

8. Molecular Orbitals
when the wave functions combine constructively, the resulting molecular orbital has less energy than the original atomic orbitals – it is called a Bonding Molecular Orbital
s, p
most of the electron density between the nuclei 8 8

9. Molecular Orbitals
when the wave functions combine destructively, the resulting molecular orbital has more energy than the original atomic orbitals – it is called a Antibonding Molecular Orbital
s*, p*
most of the electron density outside the nuclei
nodes between nuclei 9 9

10. Molecular Orbital Model
Molecular electron configurations can be written similar to atomic electron configurations.
Each molecular orbital can hold 2 electrons with opposite spins.
Orbitals are conserved.
Chapter 9 | Slide 10

11. Sigma Bonding and Antibonding Orbitals
Chapter 9 | Slide 11

12. Molecular Orbital Theory
Electrons in bonding MOs are stabilizing
Lower energy than the atomic orbitals
Electrons in anti-bonding MOs are destabilizing
Higher in energy than atomic orbitals
Electron density located outside the internuclear axis
Electrons in anti-bonding orbitals cancel stability gained by electrons in bonding orbitals 12 12

13. Since more electrons are in
Dihydrogen, H2
Molecular
Orbitals
Hydrogen
Atomic
Orbital
Hydrogen
Atomic
Orbital s* 1s 1s s
Since more electrons are in
bonding orbitals than are in antibonding orbitals,
net bonding interaction 13 13

14. Since there are as many electrons in
Dihelium, He2
Molecular
Orbitals
Helium
Atomic
Orbital
Helium
Atomic
Orbital s* 1s 1s s
Since there are as many electrons in
antibonding orbitals as in bonding orbitals,
there is no net bonding interaction 14 14

15. MO and Properties
Bond Order = difference between number of electrons in bonding and antibonding orbitals
only need to consider valence electrons
may be a fraction (partial bond order)
higher bond order = stronger and shorter bonds
if bond order = 0, then bond is unstable compared to individual atoms - no bond will form. 15 15

16. Since more electrons are in
Dilithium, Li2
Molecular
Orbitals
Lithium
Atomic
Orbitals
Lithium
Atomic
Orbitals s* 2s 2s s s*
BO = ½(4-2) = 1 1s 1s s
Since more electrons are in
bonding orbitals than are in antibonding orbitals, net bonding interaction 16 16

17. Diatomic O2 dioxygen is paramagnetic
paramagnetic material have unpaired electrons
neither Lewis Theory nor Valence Bond Theory predict this result
Paramagnetism – substance is attracted into the inducing magnetic field.
Diamagnetism – substance is repelled from the inducing magnetic fiel 17 17

18. Diatomic Oxygen, O2 Dioxygen is attracted to a magnetic field!
Neither Lewis Theory nor Valence Bond Theory predict this result.
Paramagnetism – substance is attracted into the inducing magnetic field.
- Unpaired electrons (O2)
Diamagnetism – substance is repelled from the inducing magnetic field.
- Paired electrons (N2) 18 18

19. Magnetic Properties of Liquid Nitrogen and Oxygen
Chapter 9 | Slide 19

20. Pi Bonding and Antibonding Orbitals
Chapter 9 | Slide 20

21. p Atomic Orbitals and the Formation of Molecular Orbitals

22. Molecular Orbital Diagram for B2

23. Molecular Orbital Diagram for O2

24. No s-p mixing s-p mixing
s and p Orbital Mixing
No s-p mixing
s-p mixing
B2s = 14 eV vs B2p = 8.3 eV O2s = 32.3eV vs O2p = 15.9 eV

25. s-p mixing
No s-p mixing

26. 1 electron volt (eV) = 1.60217646 × 10-19 joules

27. Factors that Affect the Formation of Molecular Orbitals
Symmetry
s and s
pz and pz (pz is along the bonding axis)
s and pz
Energy
Orbitals must have similar energy in order to overlap

28. Heteronuclear Diatomic Molecules
the more electronegative atom has lower energy orbitals
when the combining atomic orbitals are identical and equal energy, the weight of each atomic orbital in the molecular orbital are equal
when the combining atomic orbitals are different kinds and energies, the atomic orbital closest in energy to the molecular orbital contributes more to the molecular orbital
lower energy atomic orbitals contribute more to the bonding MO
higher energy atomic orbitals contribute more to the antibonding MO
nonbonding MOs remain localized on the atom donating its atomic orbitals 28 28

29. Molecular Orbital Diagram of Hydrogen Fluoride
H1s = 13.6 eV
F2s = 46.4 eV
F2p = 18.7 eV
32.8 eV
n.b. - nonbonding orbitals

30. Molecular Orbital Diagram of Carbon Monoxide
C2s = 19.5 eV
O2s = 32.3 eV
O2p = 15.9 eV

31. Polyatomic Molecules
when many atoms are combined together, the atomic orbitals of all the atoms are combined to make a set of molecular orbitals which are delocalized over the entire molecule
gives results that better match real molecule properties than either Lewis or Valence Bond theories 31 31

32. Valence Bond Model of Ozone

33. Molecular Orbital Model of Ozone

34. Resonance Structures of Benzene

35. Molecular Orbital Model of Benzene