Everything in oxygen chemistry seemed more or less in place:
Up there, in the stratosphere, there were oxygen atoms,
O2 molecules and ozone, O3, as well as
ions derived from these and a bit of active OH, all in a dance
of creation and destruction.
Meanwhile, within our
bodies, normal O2 served us well. There was even a
place, under enzyme supervision, for the somewhat nasty
relatives hydrogen peroxide (H2O2 and its
deprotonated form, O2
2-) and superoxide (O2
- and its protonated alter ego, HOO�), whose chemistry may generate the harmful
hydroxyl radical �OH. Of course, there's
also water everywhere. And here and there singlet dioxygen, a
more reactive and excited state of normal diatomic oxygen.
A nice, neatly compartmentalized world: ozone for
atmospheric chemists, but not biologists, who had plenty of more
complicated molecules to worry about.
So they thought.
This complacent state has now changed—dramatically
so—with a series of remarkable discoveries. There is new
evidence for the occurrence of ozone in living cells. Hydrogen
peroxide is being made by molecules thought incapable of doing
so. Even a metastable laboratory curiosity, the unusual HOOOH
molecule (which sounds like a holler; call it
dihydrogen-trioxide), may be in living systems. Meanwhile, we
are still puzzling out the state of oxygen in high-temperature
superconductors. And the menagerie of alternative forms of
elemental oxygen continues to expand—there are strong
theoretical arguments for the existence of a cyclic ozone
isomer, and there may even be ways of stabilizing it.