Ozone

Ozone: The Protective Allotrope

Ozone, $O_3$, is an allotropic form of dioxygen that is far more reactive than $O_2$ and therefore does not persist for long once formed at ground level. High in the atmosphere — around 20 km altitude — sunlight continuously converts ordinary atmospheric oxygen into ozone, building up the stratospheric ozone layer, which shields the Earth's surface from an excessive concentration of harmful ultraviolet (UV) radiation.

Preparation

Ozone is generated by passing a slow, dry stream of oxygen through a silent electrical discharge; under these conditions about 10% of the oxygen converts to ozone, giving what is known as ozonised oxygen:

Because the conversion of oxygen to ozone is endothermic, a silent discharge (rather than a sparking one) is essential — a violent discharge would simply decompose the ozone as fast as it forms. If a concentration higher than 10% is wanted, several ozonisers can be connected together, and the pure gas — which boils at 101.1 K — can then be condensed by surrounding the collection vessel with liquid oxygen.

Physical Properties

Pure ozone exists as a pale blue gas, condenses to a dark blue liquid, and freezes to a violet-black solid. It has a sharp, characteristic odour and is harmless at very low concentrations, but once the concentration exceeds about 100 parts per million, breathing becomes uncomfortable and causes headache and nausea.

Thermodynamically, ozone is unstable relative to oxygen: its decomposition to $O_2$ releases heat ($\Delta H$ negative) and also increases entropy ($\Delta S$ positive), so the two effects reinforce each other, making $\Delta G$ for the conversion strongly negative. This is exactly why concentrated ozone can be dangerously explosive.

Chemical Behaviour — A Powerful Oxidising Agent

Because ozone readily liberates nascent oxygen atoms ($O_3 \rightarrow O_2 + O$), it behaves as a powerful oxidant. Two characteristic examples are its oxidation of lead sulphide to lead sulphate, and of iodide ion to iodine:

This iodine-liberating reaction is put to practical use: when ozone is passed into excess potassium iodide solution buffered at pH 9.2 with a borate buffer, the liberated iodine can be titrated against standard sodium thiosulphate — a quantitative method for estimating how much $O_3$ gas is present.

Ozone is also consumed by reaction with nitrogen oxides, particularly nitrogen monoxide, and this is why nitrogen oxides emitted by supersonic jet exhausts are suspected of slowly depleting the ozone layer in the upper atmosphere:

Freons, widely used in aerosol sprays and as refrigerants, pose a further threat to the ozone layer.

Structure …