Valence Bond Theory

The orbitals shown at the left of the figure are for a nonbonded carbon atom, but once the carbon atom begins to bond with other atoms (in this case hydrogen), the atomic orbitals hybridize, and this changes their shape considerably. Notice how the first set of figures form the sp³ atomic orbital, the hybrid, and this leads to further hybridization.

Ammonia also has sp³ hybridization, even though it has a lone pair.

Now let’s look more closely at pi bonds. As we mentioned earlier in this chapter, pi (p) bonds result from the sideways overlap of p orbitals, and pi orbitals are defined by the region above and below an imaginary line connecting the nuclei of the two atoms. Keep in mind that pi bonds never occur unless a sigma bond has formed first, and they may form only if unhybridized p orbitals remain on the bonded atoms. Also, they occur when sp or sp² hybridization is present on central atom but not sp³ hybridization.

Below, we show the formation of a set of sp² orbitals. This molecule would contain a double bond, like ethene. Notice again how the first set of figures form the sp² atomic orbital, the hybrid, and the last figure shows full hybridization:

The set of p orbitals that are unhybridized are not shown in this depiction:

A different view, which doesn’t show the hydrogens and centers on the C atoms, shows the unhybridized p orbitals that create the sideways overlap that’s necessary to create the double pi bond:

Here’s how it looks with all the pieces put together:

Here is a table summarizing hybridization and structure:

effective pairs	hybridization	geometry
2	sp	Linear
3	sp2	trigonal planar
4	sp3	tetrahedral
5	dsp3	trigonal bipyramidal
6	d2sp3	octahedral