Now we can begin to put all the pieces together and make sense of what we know so far. We’ll compare the way that the electrons fi t into the energy level diagrams for benzene and planar cyclooctatetraene. We are not concerned with the actual shapes of the molecular orbitals involved, just their energies. Benzene has six π electrons, which means that all its three bonding molecular orbitals are fully occupied, giving what we can call a ‘closed shell’ structure. Cyclooctatetraene’s eight electrons, on the other hand, do not fi t so neatly into its orbitals. Six of these fill up the bond ing molecular orbitals but there are two electrons left. These must go into the degenerate pair of non-bonding orbitals. Hund’s rule (Chapter 4) would suggest one in each. Planar cycloocta tetraene would not have the closed shell structure that benzene has—to get one it must either lose or gain two electrons. This is exactly what we have already seen—both the dianion and dication from cyclooctatetraene are planar, allowing delocalization all over the ring, whereas neutral cyclooctatetraene avoids the unfavorable arrangement of electrons shown below by adopting a tub shape with localized bonds.