Oxygen

Oxygen is the most abundant element in the Earth's crust, accounting for close to half its mass. It also makes up about 21 % of the atmosphere by volume and about 65 % of the mass of a human body, mostly as water.

Almost all the energy that living things use comes from oxidation — which is to say, in the end, from oxygen. And the free oxygen in the atmosphere is itself biological in origin: photosynthesising organisms made it.

Bound, oxygen is in water, in the overwhelming majority of crustal minerals — silicates, carbonates, oxides — and in organic matter.

Free, it exists in two forms. Ordinary oxygen, O₂, is a colourless, odourless gas. Ozone, O₃, is built from the same atoms but a different molecule, and behaves differently: it has a sharp smell, is far more reactive, is poisonous at ground level, and in the stratosphere forms the layer that absorbs hard ultraviolet.

This is a good illustration of why the properties of a substance cannot be attributed to the element. Oxygen and ozone are made of identical atoms, yet they boil at different temperatures and behave unlike each other.

Oxygen was obtained independently and almost simultaneously. Carl Wilhelm Scheele isolated it around 1772 but published late. Joseph Priestley heated mercuric oxide in 1774 and got a gas in which a candle burned more brightly; he explained it within phlogiston theory and called it "dephlogisticated air".

The decisive step was Lavoisier's. He repeated the experiments, weighed everything before and after, and showed that combustion is the addition of part of the air to a substance rather than the release of phlogiston. Phlogiston theory ended there, and with it alchemical chemistry. The name oxygène, "acid former", is also Lavoisier's — and there he was wrong: there are acids without oxygen.

Oxygen is the second most electronegative element after fluorine, and nearly all its behaviour follows from that. It oxidises the great majority of simple substances, and in compounds it usually shows oxidation state −2.

Reactions with oxygen almost always release heat, but at room temperature they are often stalled: paper and petrol sit quietly in air until something lights them. The reason is the strength of the O₂ molecule — the reaction cannot start until it is broken.

The exceptions to −2 are worth remembering: in peroxides such as H₂O₂ oxygen is −1, and in the fluoride OF₂ it is +2, because fluorine is the more electronegative of the two.

Metallurgy consumes the most. Blowing oxygen through a converter burns the impurities out of pig iron and turns it into steel.

Then come the chemical industry, welding and cutting, rocketry — liquid oxygen as an oxidiser — and medicine. Water treatment is a field of its own, where oxygen and ozone break down organic contaminants.

Respiration is the oxidation of organic matter by oxygen, broken into many small steps so that the energy is released in portions a cell can store rather than all at once.

The reverse process is photosynthesis: plants and cyanobacteria split water, take the hydrogen to build organic matter, and discard oxygen as waste. All the free oxygen in the atmosphere accumulated that way, and it took some two billion years.