Benzene.

The electronic structure of benzene was, and still is, a matter of debate. The two possible structures according to Lewis-theory are shown in the figure above. These cannot be correct, however, for the following reasons. First, they imply that there are two different C-C bonds. In particular, double bonds are generally shorter than single bonds. But all C-C bonds in benzene are the same. Second, double bonds show characteristic reactions, which are not found for benzene. Third, the heat of formation of benzene is too large. If we look at the MO-diagram on the left in the next figure that shows the formation of a single p bond, we see that it contributes 2|b| to the heat of formation. The value of b can be determined from the heat of formation of alkenes. For the Lewis-structures we would then expect a contribution of 6|b|, which turns out to be too small.

The solution of these discrepancies is given by the MO's that are formed from the p orbitals of the C atoms perpendicular to the plane of the molecule. These do not mix with the other AO's, which form the s skeleton. We number the carbon atoms from 1 to 6 so that n+1 is a neighbor of n and 1 is a neighbor of 6. We write the orbitals of the p system as a linear combination

The following figure shows the MO's. We see that none of these MO's look like some bonding orbital for two C atoms. We might think that it is possible to transform to new MO's that do describe C-C bonds; as is possible for methane. That is not the case here, however. Such a transformation is not allowed to change the Slater determinant. Consequently, we can only transform the occupied MO's. We have three of them, which can only give three transformed MO's. This is, of course, not enough to describe six C-C bonds. So as we cannot make MO's form individual bonds, we talk about delocalization , delocalized bonds and delocalization energy . Such bonds are typical for aromatic molecules, and give them their specific properties.