As soon as we try to make a localized MO model of benzene, we run into trouble. The planar hexagonal geometry of the molecule, with 120 bond angles, suggests sp2 hybridization around the carbons, with one spl orbital from each C pointed toward an H, and the other two directed toward the neighboring carbon atoms in the ring. This skeleton of s bonds uses 24 of the 30 bonding electrons (6 X 4 from carbons plus 6 x 1 from hydrogens), and all of the outer orbitals except the six p orbitals perpendicular to the plane of the hexagon. This s framework is shown. What should be done with the six unused electrons and six remaining p orbitals? These are shown in perspective at the bottom left. Adjacent p orbitals could be combined in pairs around the ring to make every other carbon-carbon bond a double bond. There are two ways of doing this, represented schematically below. These are known as the Kekule structures after the man who first proposed them, but they cannot be correct because we know that all of the carbon-carbon bonds are the same length. A somewhat less plausible way of pairing the p orbitals would be to connect two across the ring, and then pair the remaining two at either side, as in the three Dewar structures shown below the Kekule rings.