Carbon Hybridization

In this excited state (having an unmatched 2s1 and 2p1 orbital), Beryllium could certainly bind with hydrogen in two unmatched orbitals according to this theory. Research, however, has shown that these two bonds are identical. How can this be? Basically, the only way this makes sense is through the process of hybridization. What this involves is combining orbitals that are not equivalent but that are properly oriented to form bonds. These new combinations are called hybrid orbitals because they are made by hybridizing two or more atomic orbitals from the same atom. This leads to a unique hybrid orbital in beryllium that has this energy pattern, as shown in figure 5:

Figure 5.

This obviously leads to the ability of the beryllium to bind with the hydrogen atom in order to make BeH2. This will produce a linear BeH2 molecule. Both promotion and hybridization require an input of energy; however, this is made up for when beryllium bonds with 2 hydrogen atoms. It means that, in situations not associated with beryllium, some compounds are so unstable because they have such a high energy necessary to make the hybrid orbital. This energy is not made up for by the energy recouped in the bonding effort.

CARBON HYBRIDIZATION

Carbon has six electrons, with a 2s22p2 configuration. Based on this, one would expect that it would be likely to bond with just two other atoms; we know, on the other hand, that it bonds with four other atoms to make covalent bonds. On the second level, there is one 2s orbital and 3 2p orbitals, which can be hybridized to make four degenerate sp3

hybrid orbitals that contain a single electron available for bonding. Figure 6 shows what this hybridization looks like:

Figure 6.

The energy of the hybridized orbitals is somewhere in the middle of the ground state and the excited state of carbon. This leads to orbitals that are in the shape of a tetrahedron with 109.5-degree angles between them. These are equivalent in energy and allow for a tetrahedral molecule, which is the exact shape of methane.

The amount of energy released in any hybrid orbital situation increases with the number of bonds formed. In the case of carbon, more energy is released in the formation of four bonds versus just two, making CH4 so much more stable than CH2 or CF2. Intermediates like CH2 are so reactive that they only form temporarily under certain experimental circumstances.

Valence bond theory also explains why molecules such as NH3 and H2O are formed. With nitrogen, the 2s22p3 electron configuration becomes hybridized to make four sp3 hybrid orbitals. One of the orbitals is completely occupied with a pair of electrons, while the remaining three are single and are available for bonding.