1. Chapter 9 Notes
AP CHEMISTRY
Galster

2. Molecular Geometry VESPR Theory
Abbreviation for valence shell electron pair repulsion.
Valence electrons will arrange themselves to reduce repulsion
i.e. they will be as far apart as possible
2 e-sets
Bond angle:
Shape name:
AZx designation:
3 e- sets

3. 4 e-sets
Bond angle:
Shape name:
AZx designation:
5 e- sets
6 e-sets

4. Electron Set Geometry
Name
Molecular Geometry
Example
3-D Shape
AZ2
(2 e- domains)
Linear
AX2
CO2
AZ3
(3 e- domains)
Trigonal planar
AX3
BF3
AX2E
Bent
SO2

5. Electron Set Geometry
Name
Molecular Geometry
Example
3-D Shape
AZ4
(4 e- domains)
Tetrahedral
AX4
CH4
AX3E
Trigonal pyramidal
NH3
AX2E2
Bent
H2O
AXE3
Linear
HCl

6. Electron Set Geometry
Name
Molecular Geometry
Example
3-D Shape
AZ5
(5 e- domains)
Trigonal bipyramidal
AX5
PCl5
AX4E
See-Saw
SF4
AX3E2
T-shaped
ClF3
AX2E3
Linear
XeF2

7. Electron Set Geometry
Name
Molecular Geometry
Example
3-D Shape
AZ6
(6 e- domains)
Octahedral
AX6
SF6
AX5E
Square pyramidal
BrF5
AX4E2
Square planar
XeF4

8. Effect of non-bonding electrons on bond angles
Order of Repulsion
Nonbonding-nonbonding > nonbonding-bonding > bonding-bonding
Trends
Angle between nonbonding pairs increase
Angles between bonding pairs decrease
Examples
1. CH NH Water, H2O
tetrahedral tetrahedral tetrahedral
tetrahedral trig. Pyramidal bent

9. 4. Sulfur tetrafluoride, SF4
5. COCl2

10. Polarity of Molecules Review of polar bonds Molecular polarity
Polar bonds - ∆EN is less than or equal to 0.5 (but >1.7, ionic)
Nonpolar bonds ∆EN is less than 0.5
Molecular polarity
Nonpolar bonds
Diatomic molecules or others with nonpolar bonds
Ex: H2, Cl2, I2
No dipole moment
So… molecules with all nonpolar bonds are nonpolar
Cancellation of dipoles
Depends on geometry “tug of war”
Polar bonds – yes
Dipole moment – no
Ex: CCl4, CO2

11. Examples of Nonpolar Molecules
CO2
SiCl5
AsCl5
ICl4-1

12. Examples of Polar Molecules
Polar bonds: yes
Dipole Moment? Yes
H2O SF4 BrF5

13. Which of these are polar?
SO2
I3-1
CHCl3
PCl3
XeCl4
SeBr4

14. Atomic Orbitals: Hybridization
Atomic Orbitals – covalent bonds result when orbitals in valence level overlap
**VESPR theory says that the number of unpaired electrons tells you the number of bonds the atom will form.
Ex: H2
HCl
Cl2

15. Hybrid Orbitals Mixing of orbitals
Pure orbitals cannot always be used to explain bonding
Basics
Atomic orbitals are combined to create new hybrid orbitals
The number of hybrid orbitals formed is equal to the number of atomic orbitals combined
* We are not increasing the number of orbitals, just rearranging them.
The names of the hybrid orbitals are based on the type and number of atomic orbitals combined.
The hybrid orbitals are degenerate and have the same shape.

16. Process of hybridization
A. methane, CH4
B. BF3
C. BeCl2

17. D. H2O
E. NH3
F. XeF4

18. Hybridization and Electron Set Geometry
sp = one s & one p = AZ2 linear
sp2= one s & two p = AZ3 trigonal planar
sp3= one s & three p = AZ4 tetrahedral
sp3d = one s, three p, and one d = AZ5 trigonal bipyramidal
sp3d2 = one s, three p, and two d = AZ6 octahedral

20. Multiple Bonds
Multiple bonds do not affect the electron set geometry of a molecule
Multiple bonds make attractions (i.e. bond) stronger and shorter
Types of bonds
Single bond = σ (sigma) bond
Double bond = σ (sigma) bond – first bond
π (pi) bond – 2nd or subsequent bonds
Ex: C2H4

21. Triple bond
Ex: C2H2 (Acetylene)

22. 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

23. Delocalized Electrons: Resonance
In reality, each of the four
atoms in the nitrate ion has
a p orbital.
The p orbitals on all three
oxygens overlap with the p
orbital on the central nitrogen.
This means the electrons are
not localized between the
nitrogen and one of the
oxygens, but rather are
delocalized throughout the
ion.

24. Resonance
The organic molecule benzene has six σ bonds and a p orbital on each carbon atom.
In reality the π electrons in benzene are not localized but delocalized
The even distribution of the π electrons in benzene makes the molecule unusually stable.

25. Molecular-Orbital (MO) Theory
Through valence bond
theory effectively conveys
most observed properties of
ions and molecules, there
are some concepts better
represented by molecular
orbitals.
In MO theory, we invoke the
wave nature of electrons
If waves interact constructively
the resulting orbital is lower in
energy: a bond

26. Molecular-Orbital (MO) Theory
If waves interact destructively, the resulting orbital is higher in energy: an antibonding molecular orbital.

27. MO Theory
In H2 the two electrons go into the bonding molecular orbital
The bond order is one half the difference between the number of bonding and antibonding electrons

28. MO Theory
For hydrogen, with two electrons in the bonding MO and none in the antibonding MO, the bond order is
½ (2 – 0) = 1

29. MO Theory In the case of He2, the bond order would be ½ (2 - 2) = 0
Therefore, He2 does not exist.

30. MO Theory For atoms with both s and p, there are two
types of interactions:
The s and the p orbitals
that face each other
overlap in σ fashion.
The other two sets of p
orbitals overlap in π
fashion.

31. MO Theory The resulting MO diagram looks like this :
There are both σ and π
bonding molecular orbitals
and σ* and π* antibonding
molecular orbitals

32. MO Theory
The smaller p-block elements in the second period have a sizable interaction between the s and p orbitals.
This flips the order of the σ and π molecular orbitals in these elements

33. Second-Row MO Diagrams