Ozone

Ozone is a significant inorganic molecule within the Earth’s atmosphere, notable for its characteristic odour, strong oxidising power and crucial role in protecting the biosphere from harmful ultraviolet radiation. Although chemically related to the more stable dioxygen, it possesses distinct structural and reactive properties that define its environmental relevance and industrial uses. As a naturally occurring gas found in trace concentrations throughout the atmosphere, its behaviour and effects vary substantially with altitude, source and concentration.

Physical and Chemical Characteristics

Ozone is a pale blue gas at standard temperature and pressure, with a distinctly sharp and chlorine-like smell detectable by many individuals at extremely low concentrations. Its solubility is relatively limited in water but it dissolves more readily in non-polar solvents, such as fluorocarbons and carbon tetrachloride, where it imparts a deep blue colour. When cooled to cryogenic temperatures, it condenses to a dark blue liquid and eventually solidifies into a violet-black solid. These condensed phases are extremely unstable, as liquid and solid ozone can detonate with abrupt temperature changes, physical shock or rapid warming.
Chemically, the molecule is composed of three oxygen atoms arranged in a bent configuration exhibiting C₂v symmetry. Microwave spectroscopy confirms an O–O–O bond angle of approximately 117°, with both O–O bonds possessing an intermediate bond order of 1.5 due to resonance structures. The central oxygen atom displays sp²-like hybridisation with a lone electron pair, contributing to the molecule’s polarity and a dipole moment of roughly 0.53 Debye. Ozone is weakly diamagnetic and is isoelectronic with the nitrite ion, sharing comparable electron distributions. Although theoretical models predict the possibility of a cyclic form of ozone, such a structure has not been observed experimentally.
The molecule is inherently unstable, decomposing to dioxygen at a rate highly dependent on temperature, humidity and air movement. Its atmospheric half-life can be minutes or hours, but under controlled laboratory conditions it typically averages around 15 minutes.

Formation and Occurrence in the Atmosphere

Ozone forms naturally in the atmosphere through the interaction of ultraviolet radiation or electrical discharges with dioxygen. In the stratosphere, its concentration reaches a maximum within the ozone layer, where it absorbs large amounts of ultraviolet-B and ultraviolet-C radiation, preventing these wavelengths from penetrating to the Earth’s surface. This protective layer generally contains between two and eight parts per million by volume of ozone, levels vital for shielding biological systems from genetic damage.
Closer to the Earth’s surface, ozone occurs in much lower concentrations and behaves very differently. Ground-level ozone is not emitted directly but is formed as a secondary pollutant through photochemical reactions involving nitrogen oxides and volatile organic compounds. In the troposphere it acts as a harmful pollutant and respiratory irritant, damaging plant tissues and reducing crop yields. Even brief exposure to concentrations around 0.01 mol per mol of air can cause headaches, eye irritation and a burning sensation in the respiratory tract.

Nomenclature and Terminology

The commonly used name “ozone” is the trivial and preferred IUPAC designation for the molecule, derived from the Greek ozein, meaning “to smell”, alluding to its easily recognisable odour. IUPAC nomenclature also permits the systematic names 2λ⁴-trioxidiene and catenatrioxygen, which describe its bonding and sequential oxygen arrangement. In contexts where ozone is considered a hydrogen-deficient derivative of trioxidane, the term trioxidanylidene may be used, although this description does not inherently distinguish between its radical and non-radical states. The diradical form is specifically referred to as trioxidanediyl. The associated OOO substituent group is sometimes informally labelled “ozonide”, but this must not be confused with ozonide salts formed in organic and inorganic reactions.

Historical Development

The discovery of ozone traces back to late-eighteenth-century experiments. In 1785, the Dutch chemist Martinus van Marum observed an unusual odour when producing electrical sparks over water, but he did not identify the substance. It was the German–Swiss chemist Christian Friedrich Schönbein who, in 1839, isolated the gas and recognised the same odour that follows lightning discharges. He named the new substance “ozone”, linking its scent to earlier observations.
Throughout the nineteenth century, ozone attracted scientific interest for both its perceived health benefits and harmful physiological effects. In 1865, Jacques-Louis Soret determined its chemical formula as O₃, a finding later confirmed by Schönbein in 1867. Early naturalists associated fresh, high-altitude or coastal air with ozone, believing it contributed to vitality. However, subsequent investigations clarified that the characteristic seaside smell often attributed to ozone was actually produced by halogenated organic compounds released by marine organisms.
Scientific studies soon established its toxicity. In 1873, James Dewar and John Gray McKendrick demonstrated that animals exposed to ozonised air exhibited respiratory distress, reduced blood oxygenation and lethargy. Schönbein himself suffered chest discomfort and mucous membrane irritation during experiments. By the early twentieth century, physiological research concluded that ozone inhalation, even at moderate concentrations, could cause lung inflammation, oedema and potentially fatal respiratory injury.
During the First World War, attempts were made to employ ozone as a wound disinfectant at Queen Alexandra Military Hospital in London. While the gas succeeded in killing bacteria, it also damaged living tissues, leading to its rejection in favour of liquid antiseptics. Analytical challenges persisted into the 1920s due to the explosive instability of concentrated ozone. Georg-Maria Schwab’s doctoral work under Ernst Riesenfeld finally enabled reliable solidification and analysis of pure ozone, debunking theories suggesting the presence of tetraoxygen impurities and providing precise measurements of its physical properties.

Industrial Applications and Behaviour

Ozone’s technical significance arises from its exceptional oxidising strength, which surpasses that of dioxygen and many other common oxidants. As a result, it is applied across several industrial contexts, typically in low concentrations to minimise hazards. Key applications include water purification, bleaching of textiles and paper pulp, deodorisation, and sterilisation of air and surfaces. In water treatment, ozone is valued for its ability to degrade organic pollutants, inactivate bacteria and viruses, and remove odours without leaving harmful residues.
In chemical synthesis, ozone is used in ozonolysis reactions to cleave carbon–carbon double bonds in alkenes, producing aldehydes, ketones or carboxylic acids depending on reaction conditions. These reactions must be handled carefully due to the instability of ozonide intermediates, which require controlled reductive or oxidative work-up steps.
Its instability also restricts storage and transport. Ozone is generally produced on-site using electrical corona discharge or ultraviolet radiation, immediately applied and allowed to decompose back to dioxygen. Even under controlled industrial conditions, its oxidative potential poses corrosion risks to materials such as rubber, plastics and metal surfaces.

Environmental and Biological Implications

Ozone’s environmental effects depend critically on altitude. In the stratosphere, ozone is indispensable for life, preventing harmful ultraviolet radiation from reaching the surface and playing a key role in atmospheric heat balance. Depletion of this layer—through reactions with anthropogenic halogenated compounds—has been a central environmental concern, leading to international regulatory efforts to reduce emissions of ozone-depleting substances.
In contrast, at ground level it is considered a primary component of photochemical smog. Exposure can impair lung function, aggravate asthma, and reduce plant productivity by damaging leaf tissues and inhibiting photosynthesis. Sensitive materials such as natural rubber deteriorate rapidly when exposed, highlighting ozone’s ability to attack organic structures even at low concentrations.
Despite these hazards, its atmospheric role is complex. Tropospheric ozone also participates in oxidising processes that influence climate and air quality, serving as both a greenhouse gas and a reactive species shaping the lifetimes of other atmospheric constituents.