odd-electron molecules

The octet rule and Lewis structures quietly assume electrons come in pairs. But some perfectly real molecules have an odd total number of electrons, so at least one electron is left unpaired no matter how you draw them. These odd-electron molecules, or free radicals, simply cannot give every atom a neat octet, and they show that pairing is a tendency, not an iron law.

Nitric oxide, NO, is the textbook case: nitrogen contributes five valence electrons and oxygen six, eleven in all, an odd number. Draw the best Lewis structure you like and one electron will be left over, unpaired, sitting on the nitrogen. Nitrogen dioxide NO2 (17 valence electrons) and chlorine dioxide ClO2 are other examples. Because of that lone unpaired electron, odd-electron molecules are paramagnetic — they are attracted by a magnetic field — and they are usually very reactive, eager to pair up their stray electron by reacting with another species or dimerising (NO2 partly dimerises to N2O4, pairing the electrons in an N-N bond).

Odd-electron species are central to a great deal of real chemistry: NO is a signalling molecule in the body and a key player in atmospheric and combustion chemistry, while NO2 and other radicals drive air pollution and smog. The simple Lewis picture struggles with them, but molecular-orbital theory handles the unpaired electron naturally — it just places that one electron in the appropriate molecular orbital.