Issue Date: July 17, 2017
Dioxygen is a diradical, so why is it stable?
Oxygen as O2 is stable enough to be abundant in the environment and is required for many forms of life. But from the standpoint of theory, dioxygen’s stability is curious: Its highest occupied molecular orbitals contain two unpaired electrons, making it a diradical. Instead of wafting around as O2, the molecule should be busy abstracting hydrogen atoms or forming oligomers; isoelectronic sulfur, for example, is most stable as S8. The key to dioxygen’s reactivity lies in resonance. That finding comes from experimental and computational analysis by Weston T. Borden of the University of North Texas, Roald Hoffmann of Cornell University, and their colleagues (J. Am. Chem. Soc. 2017, DOI: 10.1021/jacs.7b04232). Confirming a 1931 proposal by Linus Pauling, the researchers determined that O2’s π bond can be thought of as a pair of two-center, three-electron bonds, with resonance contributing a net stabilization energy of 418 kJ/mol. S2 is stabilized by only about 213 kJ/mol. The consequence is that trimerization of O2 is endothermic, while S2 is exothermic. Meanwhile, O2’s σ bond is relatively weak, so oxidation reactions are ultimately exothermic when they do occur.
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