dioxygen and ozone allotropes

Oxygen comes in two everyday forms. Ordinary dioxygen, O2, is the colourless gas that keeps you alive and makes up about a fifth of the air. Ozone, O3, is its three-atom cousin: a pale blue, sharp-smelling gas (you notice it near sparking electrical equipment or after a thunderstorm) that is both a shield high in the atmosphere and a pollutant near the ground. Same element, two allotropes, two very different molecules.

Dioxygen hides a famous surprise. A simple Lewis structure draws O=O with a tidy double bond and all electrons paired, but that prediction is wrong: liquid oxygen is paramagnetic — it sticks to a magnet — which means it has unpaired electrons. Molecular orbital theory explains this beautifully: when you fill the MO diagram for O2, the last two electrons go one each into a pair of equal-energy antibonding pi orbitals (following Hund's rule), leaving two unpaired electrons. So O2 is a diradical with a net double bond, and MO theory's success here is one of the strongest arguments for the whole model. Ozone, O3, is a bent molecule, with the central oxygen bonded to the other two and the bonding delocalised by resonance over both O-O links, so its bonds are identical and intermediate between single and double.

Both forms matter enormously. O2 is the oxidant that powers respiration, combustion, and corrosion. Ozone is a Jekyll-and-Hyde molecule: high in the stratosphere a thin layer of it absorbs the Sun's hard ultraviolet light and protects life, which is why chlorofluorocarbon damage to that ozone layer was so alarming; but down at ground level ozone is a toxic, irritating component of photochemical smog and a respiratory hazard. The honest summary often heard is "good up high, bad nearby" — the same molecule, helpful or harmful depending only on where it is.