The valance bond theory looks at the question of how the electrons of different atoms interact as they approach one another and eventually bond.
This theory focuses on the central atom of the structure and how its bonding electrons behave during a bond. In doing so, the theory looks at areas of electron location , which can be either a pair of electrons or a lone electron going into a bond. Again, the focus is on the central atom only.
For example, in the molecule of water, the oxygen atom has four areas of electron location. Two are pairs of electrons are two are lone electrons bonding with the hydrogen.
The valance bond theory asserts that, as the orbitals of the terminal atoms approach the orbitals of the central atom, an overlap takes place. It is this overlap that creates the bond. The greater the overlap, the stronger the bond energy. High overlap occurs when the atoms approach the bond in an end-to-end fashion.
Hybridization of Atomic Orbitals:
This is the crux of the valance bond theory. Remember, we are focusing only on the behavior of the central atom.
Case in point: Carbon
Carbon’s electron configuration:
However, as we have seen, carbon bonds in four directions. This can be explained by the following change in the electron configuration:
In this situation, the bonding electrons include the one electron in the 2s sublevel and the three electrons in the 2p orbitals.

If this were the case, you would expect the bond arising from the s sublevel to be different from the bonds arising from the p sublevels, because the overlap patterns would be different.
However, this is not the case. In reality, all four bonds are identical as a result of hybridization.
Hybridization is the interaction that occurs between the regions of high electron density surrounding a central atom. In the case of CH
4
, the one s orbital and the three p orbitals hybridize to create four equivalent orbitals, called sp
3
orbitals.
You must be able to predict the type of hybridization orbitals in a molecule. This can be most easily identified by identifying the number of areas of high electron density.
Areas of high electron density
2
3
4
5
6
Hybrid orbitals utilized sp sp 2 sp 3 sp 3 d sp 3 d 2
Remember, a region of high electron density can be either a pair of non-bonding electrons or a bonding area. A double or triple bond count as only one area of electron density because they exist in the same region of space.
Ex: NCl
3
Nitrogen has one unshared pair of electrons and three regions of bonding.
Therefore, there are four regions of electron density so this molecule is utilizing sp
3 orbitals.

If you consider the electron configuration of nitrogen, 1s
2
2s
2
2p
3
, this makes sense. The unshared pair is in the 2s sublevel and the three bonding electrons are in the 2p orbitals.
Therefore, the hybridizing orbitals are sp
3
.
Ex: OF
2
Oxygen is single bonded to each of the two fluorine’s and has two unshared pairs of electrons. Therefore, there are four regions of electron density and this indicates sp 3 hybridization. In light of oxygen’s electron configuration, 1s 2
2s
2
2p
4
, this makes sense.
The electrons of the 2s and the first 2p orbitals represent the paired electrons. The single electrons in the last 2p orbitals represent the bonding sites.
Ex: CO
2
Carbon is double bonded to the oxygens on each side. Therefore, there are two areas of electron density so the hybridized orbitals are sp.
However, when you consider the electron configuration of carbon shown above, how does this make sense? If only the s and one of the p orbitals hybridizes, what about the other two p orbitals? This introduces the concept of sigma and pi bonds.
Sigma bonds are the first bond between two atoms. They utilize hybrid orbitals and have a greater overlap than pi bonds. Because they have greater overlap, sigma bonds have high bond energies.
Pi bonds are the second and third bonds between two atoms. Because they have side by side overlap, they are not as strong as the first bond between two atoms. However, they do have some strength and so add to the bond energy.
Count sigma and pi bonds in HCN, CO
3
-2
, and C
2
H
4
.

Ex: SF
6
The sketch of SF
6
shows a sulfur bonded to six individual fluorines. Therefore, there are six regions of electron density and it utilizes sp 3 d 2 orbitals.
When you consider the electron configuration of sulfur, it is
How can the sp
3 d
2
hybridization be explained?
Why do you suppose PCl
5
exists but NCl
5
does not?