The Story of O
Oxygen is the most abundant element in the crust of the Earth. It is mostly tied up in carbonates and phosphates, and in a wider range of silicates, from clays to zeolites to quartz.
Under ambient conditions, diatomic oxygen is the stable form of the element. O2 is a strange molecule, for all its ubiquity. Its problem (no, our problem as we try to think about it) is that oxygen's two least strongly held electrons, responsible for most of its chemistry, have available to them two orbitals. One electron goes into each orbital, so the ground state of O2 has two unpaired electrons.
From here on I will refer to this ground state simply as O2. The first excited state of O2, usually simply called singlet oxygen or 1O2, lies 95 kilojoules per mole above the ground state. Singlet oxygen is an energetic, still more reactive form of oxygen that is relatively easy to make, either chemically or with light.
Normal ground state O2 is not an inert molecule. A single unpaired electron on oxygen (as in ●OH) or on an organic fragment (as in methyl, ●CH3) makes such a "radical" very reactive. Radicals rip hydrogens from previously stable molecules; they also start polymerization chains. With its two unpaired electrons, O2 is a "diradical." It enters many organic and inorganic reactions—as we'll soon see, this changed the course of evolution.
But compared to other diradicals, O2 is surprisingly unreactive. "Otherwise California would be burning permanently, and not just from time to time all the time," as a La Jolla-based colleague remarked in the midst of the recent fires. Many organic molecules have large barriers to reaction with O2. The reasons for O2's attenuated reactivity are still being debated, but there is no ambiguity about the destructive nature of singlet oxygen and the ●OH and ●OOH radicals. Or of the metastable allotrope of oxygen, ozone. These attack, with alacrity, nearly anything organic.