9.5. Hybridization and Molecular Orbitals http://www.ck12.org
According to the description of valence bond theory so far, carbon would be expected to form only two bonds,
corresponding to its two unpaired electrons. However, methane is a common and stable molecule that contains four
equivalent C−H bonds. One way to account for this might be to promote one of the 2s electrons to the empty 2p
orbital.
Now, four bonds are possible. The promotion of the electron “costs” a small amount of energy, but recall that the
process of bond formation is accompanied by a decrease in energy. The two extra bonds that can now be formed
result in a lower overall energy, and thus a greater stability, for the CH 4 molecule. Carbon normally forms four
bonds in most of its compounds.
The number of bonds is now correct, but the geometry is wrong. The three p orbitals (px, py, pz) are oriented at 90°
angles relative to one another. However, as we saw in our discussion of VSEPR theory, the observed H−C−H bond
angle in the tetrahedral CH 4 molecule is actually 109.5°. Therefore, the methane molecule cannot be adequately
represented by simple overlap of the 2s and 2p orbitals of carbon with the 1s orbitals of each hydrogen atom.
To explain the bonding in methane, it is necessary to introduce the concept of hybridization and hybrid atomic
orbitals. Hybridizationis the mixing of the atomic orbitals in an atom to produce a set of hybrid orbitals. When
hybridization occurs, it must do so as a result of the mixing of nonequivalent orbitals. For example, s and p orbitals
can hybridize, but p orbitals cannot hybridize only with other p orbitals. Hybrid orbitalsare the atomic orbitals
obtained when two or more nonequivalent orbitals from the same atom combine in preparation for bond formation.
In the current case of carbon, the single 2s orbital hybridizes with the three 2p orbitals to form a set of four hybrid
orbitals, called sp^3 hybrids.
The sp^3 hybrids are all equivalent to one another. Spatially, the hybrid orbitals point towards the four corners of a
tetrahedron (Figure9.23):
The large lobe from each of the sp^3 hybrid orbitals then overlaps with normal unhybridized 1s orbitals on each
hydrogen atom to form the tetrahedral methane molecule.
Another example of sp^3 hybridization occurs in the ammonia (NH 3 ) molecule. The electron domain geometry
of ammonia is also tetrahedral, meaning that there are four groups of electrons around the central nitrogen atom.
Unlike methane, however, one of those electron groups is a lone pair. The resulting molecular geometry is trigonal
pyramidal. Just as in the carbon atom, the hybridization process starts as a promotion of a 2s electron to a 2p orbital,
followed by hybridization to form a set of four sp^3 hybrids. In this case, one of the hybrid orbitals already contains a
pair of electrons, leaving only three half-filled orbitals available to form covalent bonds with three hydrogen atoms.