Ozone: Preparation, Structure and Properties

Ozone, $\text{O}_3$, is the triatomic allotrope of oxygen and, along with ordinary diatomic

dioxygen, one of only two allotropes of the element studied at this level. It is prepared in the

laboratory using a Siemens ozoniser: dry oxygen gas is passed between two coaxial electrodes

separated by a narrow gap, across which a high-voltage silent (non-sparking) electric discharge is maintained. A small fraction (5-10%) of the oxygen passing through is converted to

ozone:

A silent discharge, rather than an ordinary spark, is essential, because ozone itself is

thermodynamically unstable relative to $\text{O}_2$ and decomposes readily on heating or in a

localised hot spark; the diffuse, low-heat silent discharge produces ozone from oxygen without

immediately destroying it again.

Structurally, ozone is a bent (angular) molecule, with a central oxygen atom that is

$sp^2$ hybridised and bears one lone pair, giving an $\text{O-O-O}$ bond angle of about

$116.8^\circ$. Simple valence-bond structures suggest one $\text{O=O}$ double bond and one

$\text{O-O}$ single bond, but the two are found experimentally to be of equal length

(about 128 pm, intermediate between a typical O-O single bond, about 148 pm, and a

typical O=O double bond, about 121 pm). This is explained by resonance: two equally

valid canonical structures can be drawn, differing only in which of the two terminal oxygens

carries the formal double bond (with the central oxygen carrying a formal positive charge and the

single-bonded terminal oxygen a formal negative charge in each), and the true molecule is a

resonance hybrid of the two, with the extra $\pi$-electron density effectively delocalised over

both O-O linkages equally.

Ozone is a powerful oxidising agent, considerably stronger than dioxygen itself, because its

decomposition to ordinary, stable $\text{O}_2$ releases one highly reactive atomic (nascent) oxygen as a by-product: $\text{O}_3 \rightarrow \text{O}_2 + [\text{O}]$. This nascent oxygen

is the true oxidising species in most of ozone's reactions. The classic laboratory test for

ozone exploits exactly this: ozone oxidises iodide ion in neutral potassium iodide solution to

free iodine (which turns starch solution blue-black), while itself being reduced to oxygen and

hydroxide:

Ordinary dioxygen does not oxidise iodide under these mild conditions at all, which is why this …