1. Hybridization of Orbitals
Creating Bonds that are neither pure s nor p nor d!

2. Valence Bond Model
Linus Pauling developed valence bond theory/model in the 1930’s and won the Nobel Prize for his work in 1954
A covalent bond is a pair of electrons with opposite spins
1s 2s 2px 2py 2pz F
(in HF)

3. So to form a covalent bond, an atom needs an unpaired electron
However, most atoms don’t have unpaired electrons
1s 2s 2px 2py 2pz C
We assume that as the bonding atoms approach one another, the orbitals undergo a dramatic change: they hybridize

7. How do I predict the hybridization of the central atom?
Count the number of lone pairs AND the number
of atoms bonded to the central atom
# of Lone Pairs +
# of Bonded Atoms
Hybridization
Examples 2 sp
BeCl2 3
sp2
BF3 4
sp3
CH4, NH3, H2O 5
sp3d
PCl5 6
sp3d2
SF6

9. Practice Problems PCl5 SO3 Protocol:
Lewis dot structure, count pairs around central atom
Atoms bond hybrid orbitals repulsions
Take a picture shape bond angles polarity

10. Multiple Bonds
Geometry is not effected by multiple bonds (double, triple)
This is because the extra pair of electrons to make a multiple bond are not hybridized!
Ethylene acetylene
H H H—C C—H
C C
H H

11. Sigma and Pi bonds Single bonds
In the valence bond theory, single bonds are made up of atomic orbitals that are cylindrically symmetric about the line joining the bonded atoms (the inter-nuclear axis). Such bonds are called sigma (s) bonds.
Overlap of 2 s orbitals in H2
Overlap of an s and a p orbital in HF
Overlap of 2 p orbitals in F2
Overlap of an s or p orbital with an spy hybrid orbital - BeCl2, CH4, PCl5, etc.
Multiple bonds
The second (or third) bond in a multiple bond involves overlap of two p orbitals that are perpendicular to the internuclear axis.
The inter-nuclear axis contains a nodal plane.
Electron density is above and below the nodal plane.
These bonds are called pi (p) bonds.

12. Bonding in Ethylene
Sigma bonds on carbon result from sp2 hybrid orbitals. The resulting geometry is trigonal planar for each carbon.
The unhybridized p orbital on each carbon form the second C–C bond, which is a pi bond.

14. Free Rotation
Single Bonds (sigma bonds) allow free rotation about the bond.
Multiple Bonds do not allow the molecule to rotate about the bond, since the degree of overlap of the pi-bonded orbitals would be changed.

16. Sigma (s) and Pi (p) Bonds
1 sigma bond
Single bond
Double bond
1 sigma bond and 1 pi bond
Triple bond
1 sigma bond and 2 pi bonds
How many s and p bonds are in the acetic acid
(vinegar) molecule CH3COOH? C H O
s bonds = 6
+ 1 = 7
p bonds = 1

17. cis and trans Isomerism
Rotation about the internuclear axis of C-C single bonded atoms is relatively easy at room temperature.
Rotation about the internuclear axis of C=C double bonded atoms is not very easy at room temperature: p-bond energy ~ 260 kJ/mol
The result can be that two structural forms are “frozen out”
Each form is actually a different compound because the atoms have different arrangements in 3-space
Each can have different physical properties such as boiling temperature, molecular polarity or color
Isomers are compounds that have the same molecular formula but different arrangements of the atoms in 3-space
Example: cis- and trans- 1,2-dichloroethene

18. Valence Bond Theory
Delocalized p bonds: When two or more resonance structures can be written involving p bonds, the p bonds can be represented as being spread out over the molecule where the p bonds occur.
The electrons and the bonds containing them are delocalized.
Example: Benzene
The s framework of C-C and C-H bonds is based on sp2 hybridized C atoms.
a. After accounting for s bonding, an unhybridized p orbital remains on each C atom, with one electron per p orbital
b. The p orbitals overlap to form p bonds that form a continuous p electron cloud above and below the plane of the ring
Delocalized p bonds

19. Summary of Valence Bond Results
Bonded atoms share one or more pairs of electrons.
At least one sigma bond exists between each pair of bonded electrons.
Sigma bonds are cylindrically symmetric along the inter-nuclear axis and electrons are concentrated - localized - between the bonded atoms.
An appropriate set of hybrid atomic orbitals is formed to form s bonds.
The set of hybrid orbitals depends on the number of s bonds to be formed, the number of non-bonded electron pairs and the geometry of the overall molecule.
Atoms sharing more than one pair of electrons form pi bonds by sideways overlap of p atomic orbitals.
pi bonds have a nodal plane containing the inter-nuclear axis
In pi bonds, electron density is concentrated above and below the nodal plane.
Molecules with two or more resonance structures can have pi bonds delocalized over more than two atoms.

20. Practice Problem
What is the hybridization of each indicated atom in the following
molecule? How many sigma and pi bonds are in the molecule?