hydrogen peroxide

Hydrogen peroxide, H2O2, is the pale, slightly syrupy liquid sold dilute as an antiseptic and bleach — the bottle that fizzes when you pour it on a cut. It looks like water with one extra oxygen, and that is exactly what it is, but the extra oxygen changes everything: H2O2 is reactive, unstable, and one of inorganic chemistry's most versatile oxidising and reducing agents.

The key feature is the oxygen-oxygen single bond at its heart: H-O-O-H, with oxygen in the unusual -1 oxidation state (rather than the -2 it shows almost everywhere else). That O-O bond is weak, and oxygen would much rather be in water (-2) or O2 gas (0), so H2O2 is poised to change. It readily disproportionates — the same -1 oxygen splits, some going down to -2 (water) and some up to 0 (oxygen gas) in the reaction 2 H2O2 to 2 H2O + O2 — which is why a bottle slowly loses strength and why it foams vigorously when a catalyst such as manganese dioxide, or the enzyme catalase in your blood and in a potato, speeds the breakdown. The molecule itself is not flat; it has a skewed, open-book shape with the two O-H bonds twisted out of plane.

Because its oxygen sits in the middle of its possible range, H2O2 can act either way: it is usually an oxidising agent (bleaching hair and paper, disinfecting, treating wastewater), but against a stronger oxidiser it acts as a reducing agent, itself being oxidised to O2. Industrially it is valued as a comparatively clean oxidant, because its by-product is just water and oxygen. The honest caveat: the household 3% solution is mild, but concentrated hydrogen peroxide is dangerous — a strong oxidiser that can burn skin and, in high concentration, decompose explosively or feed rocket and weapon propellants.