The mystery deepens when we look at what happens when we treat cyclooctatetraene with powerful oxidizing or reducing agents. If 1,3,5,7-tetramethylcyclooctatetraene is treated at low temperature (–78 °C) with SbF₅/SO₂ClF (strongly oxidizing conditions) a dication is formed. This cation, unlike the neutral compound, is planar and all the C–C bond lengths are the same.

It is also possible to add electrons to cyclooctatetraene by treating it with alkali metals and a dianion results. X-ray structures reveal this dianion to be planar, again with all C–C bond lengths the same (140.7 pm). The difference between the anion and cation of cycloocta tetraene on the one hand and cyclooctatetraene on the other is the number of electrons in the π system. Neutral, non-planar, cyclooctatetraene has eight π electrons, the planar dication has six π electrons (as does benzene), and the planar anion has ten. Can you see a pattern forming? The important point is not the number of conjugated atoms but the number of electrons in the π system.

● When they have four or eight π electrons, both six- and eight-membered rings adopt non-planar structures; when they have six or ten π electrons, a planar structure is preferred.

If you made a model of cyclooctatetraene, you might have tried to force it to be flat. If you managed this you probably found that it didn’t stay like this for long and that it popped back into the tub shape. The strain in planar cyclooctatetraene can be overcome by the molecule adopting the tub conformation. The strain is due to the numbers of atoms and double bonds in the ring—it has nothing to do with the number of electrons. The planar dication and dianion of cyclooctatetraene still have this strain. The fact that these ions do adopt planar structures must mean there is some other form of stabilization that outweighs the strain of being planar. This extra stabilization is called aromaticity.