Ozone

Ozone, also called trioxygen, is an inorganic molecule with the chemical formula . It is a pale-blue gas with a distinctively pungent odour. It is an allotrope of oxygen that is much less stable than the diatomic allotrope, breaking down in the lower atmosphere to . Ozone is formed from dioxygen by the action of ultraviolet light and electrical discharges within the Earth's atmosphere. It is present in very low concentrations throughout the atmosphere, with its highest concentration high in the ozone layer of the stratosphere, which absorbs most of the Sun's ultraviolet radiation.
Ozone's odour is reminiscent of chlorine, and detectable by many people at concentrations of as little as in air. Ozone's O₃ structure was determined in 1865. The molecule was later proven to have a bent structure and to be weakly diamagnetic. At standard temperature and pressure, ozone is a pale blue gas that condenses at cryogenic temperatures to a dark blue liquid and finally a violet-black solid. Ozone's instability with regard to more common dioxygen is such that both concentrated gas and liquid ozone may decompose explosively at elevated temperatures, physical shock, or fast warming to the boiling point. It is therefore used commercially only in low concentrations.
Ozone is a powerful oxidising agent and has many industrial and consumer applications related to oxidation. This same high oxidising potential, however, causes ozone to damage mucous and respiratory tissues in animals, and also tissues in plants, above concentrations of about. While this makes ozone a potent respiratory hazard and pollutant near ground level, a higher concentration in the ozone layer is beneficial, preventing damaging UV light from reaching the Earth's surface.

Nomenclature

The trivial name ozone is the most commonly used and preferred IUPAC name. The systematic names 2λ⁴-trioxidiene and catena-trioxygen, valid IUPAC names, are constructed according to the substitutive and additive nomenclatures, respectively. The name ozone derives from ozon, the Greek neuter present participle of ozein "to smell", referring to ozone's distinctive smell.
In appropriate contexts, ozone can be viewed as trioxidane with two hydrogen atoms removed, and as such, trioxidanylidene may be used as a systematic name, according to substitutive nomenclature. By default, these names pay no regard to the radicality of the ozone molecule. In an even more specific context, this can also name the non-radical singlet ground state, whereas the diradical state is named trioxidanediyl.
Trioxidanediyl is used, non-systematically, to refer to the substituent group. Care should be taken to avoid confusing the name of the group for the context-specific name for the ozone given above.

History

In 1785, Dutch chemist Martinus van Marum was conducting experiments involving electrical sparking above water when he noticed an unusual smell, which he attributed to the electrical reactions, failing to realize that he had in fact produced ozone.
A half century later, Christian Friedrich Schönbein noticed the same pungent odour and recognized it as the smell often following a bolt of lightning. In 1839, he succeeded in isolating the gaseous chemical and named it "ozone", from the Greek word ozein meaning "to smell".
For this reason, Schönbein is generally credited with the discovery of ozone. He also noted the similarity of ozone smell to the smell of phosphorus, and in 1844 proved that the product of reaction of white phosphorus with air is identical. A subsequent effort to call ozone "electrified oxygen" he ridiculed by proposing to call the ozone from white phosphorus "phosphorized oxygen". The chemical formula for ozone, O₃, was not determined until 1865 by Jacques-Louis Soret and confirmed by Schönbein in 1867.
For much of the second half of the 19th century and well into the 20th, ozone was considered a healthy component of the environment by naturalists and health-seekers. Beaumont, California, had as its official slogan "Beaumont: Zone of Ozone", as evidenced on postcards and Chamber of Commerce letterhead. Naturalists working outdoors often considered the higher elevations beneficial because of their ozone content which was readily monitored. "There is quite a different atmosphere with enough ozone to sustain the necessary energy ", wrote naturalist Henry Henshaw, working in Hawaii. Seaside air was considered to be healthy because of its believed ozone content. The smell giving rise to this belief is in fact that of halogenated seaweed metabolites and dimethyl sulphide.
Much of ozone's appeal seems to have resulted from its "fresh" smell, which evoked associations with purifying properties. Scientists noted its harmful effects. In 1873 James Dewar and John Gray McKendrick documented that frogs grew sluggish, birds gasped for breath, and rabbits' blood showed decreased levels of oxygen after exposure to "ozonized air", which "exercised a destructive action". Schönbein himself reported that chest pains, irritation of the mucous membranes, and difficulty breathing occurred as a result of inhaling ozone, and small mammals died. In 1911, Leonard Hill and Martin Flack stated in the Proceedings of the Royal Society B that ozone's healthful effects "have, by mere iteration, become part and parcel of common belief; and yet exact physiological evidence in favour of its good effects has been hitherto almost entirely wanting... The only thoroughly well-ascertained knowledge concerning the physiological effect of ozone, so far attained, is that it causes irritation and œdema of the lungs, and death if inhaled in relatively strong concentration for any time."
During World War I, ozone was tested at Queen Alexandra Military Hospital in London as a possible disinfectant for wounds. The gas was applied directly to wounds for as long as 15 minutes. This resulted in damage to both bacterial cells and human tissue. Other sanitizing techniques, such as irrigation with antiseptics, were found preferable.
Until the 1920s, it was not certain whether small amounts of oxozone,, were also present in ozone samples due to the difficulty of applying analytical chemistry techniques to the explosive concentrated chemical. In 1923, Georg-Maria Schwab was the first to successfully solidify ozone and perform accurate analysis which conclusively refuted the oxozone hypothesis. Further hitherto unmeasured physical properties of pure concentrated ozone were determined by the Riesenfeld group in the 1920s.

Physical properties

Ozone is a colourless or pale blue gas, slightly soluble in water, and much more soluble in inert non-polar solvents such as carbon tetrachloride or fluorocarbons, in which it forms a blue solution. At, it condenses to form a dark blue liquid. It is dangerous to allow this liquid to warm to its boiling point, because both concentrated gaseous ozone and liquid ozone can detonate. At temperatures below, it forms a violet-black solid.
Ozone has a very specific sharp odour somewhat resembling chlorine bleach. Most people can detect it at the 0.01 μmol/mol level in air. Exposure of 0.1 to 1 μmol/mol produces headaches and burning eyes and irritates the respiratory passages.
Even low concentrations of ozone in air are very destructive to organic materials such as latex, plastics, and animal lung tissue.
The ozone molecule is weakly diamagnetic.

Structure

According to experimental evidence from microwave spectroscopy, ozone is a bent molecule, with C_2v symmetry. The O–O distances are. The O–O–O angle is 116.78°. The central atom is sp² hybridized with one lone pair. Ozone is a polar molecule with a dipole moment of 0.53 D. The molecule can be represented as a resonance hybrid with two contributing structures, each with a single bond on one side and double bond on the other. The arrangement possesses an overall bond order of 1.5 for both sides. It is isoelectronic with the nitrite anion. Naturally occurring ozone can be composed of substituted isotopes. A cyclic form has been predicted but not observed.

Reactions

Ozone is among the most powerful oxidising agents known, far stronger than dioxygen|. It is also unstable at high concentrations, decaying into ordinary diatomic oxygen. Its half-life varies with atmospheric conditions such as temperature, humidity, and air movement. Under laboratory conditions, the half-life will average ~1500 minutes in still air at room temperature, zero humidity with zero air changes per hour.
This reaction proceeds more rapidly with increasing temperature. Deflagration of ozone can be triggered by a spark and can occur in ozone concentrations of 10 wt% or higher.
Ozone can also be produced from oxygen at the anode of an electrochemical cell. This reaction can create smaller quantities of ozone for research purposes.
This can be observed as an unwanted reaction in a Hoffman apparatus during the electrolysis of water when the voltage is set above the necessary voltage.

With metals

Ozone oxidises most metals into oxides of the metals in their highest oxidation state. For example:

With nitrogen and carbon compounds

Ozone oxidises nitric oxide to nitrogen dioxide:
This reaction is accompanied by chemiluminescence.
The can be further oxidised to nitrate radical:
The formed can react with to form dinitrogen pentoxide.
Solid nitronium perchlorate can be made from, and gases:
Ozone does not react with ammonium salts, but it oxidises ammonia to ammonium nitrate:
Ozone reacts with carbon to form carbon dioxide, even at room temperature:

With sulphur compounds

Ozone oxidizes sulphides to sulphates. For example, lead sulphide is oxidized to lead sulphate:
Sulphuric acid can be produced from ozone, water and either elemental sulphur or sulphur dioxide:
In the gas phase, ozone reacts with hydrogen sulphide to form sulphur dioxide:
In an aqueous solution, however, two competing simultaneous reactions occur, one to produce elemental sulphur, and one to produce sulphuric acid: