Ozone
Allotrope of oxygen (O₃) present in Earth's atmosphere / From Wikipedia, the free encyclopedia
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Ozone (/ˈoʊzoʊn/) (or trioxygen) is an inorganic molecule with the chemical formula O
3. It is a pale blue gas with a distinctively pungent smell. It is an allotrope of oxygen that is much less stable than the diatomic allotrope O
2, breaking down in the lower atmosphere to O
2 (dioxygen). Ozone is formed from dioxygen by the action of ultraviolet (UV) 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 (UV) radiation.
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Names | |||
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IUPAC name
Ozone | |||
Systematic IUPAC name
Trioxygen | |||
Other names
2λ4-trioxidiene; catena-trioxygen | |||
Identifiers | |||
3D model (JSmol) |
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ChEBI | |||
ChemSpider |
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ECHA InfoCard | 100.030.051 | ||
EC Number |
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1101 | |||
MeSH | Ozone | ||
PubChem CID |
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RTECS number |
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UNII | |||
CompTox Dashboard (EPA) |
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Properties | |||
O3 | |||
Molar mass | 47.997 g·mol−1 | ||
Appearance | Colourless to pale blue gas[1] | ||
Odor | Pungent[1] | ||
Density | 2.144 g/L (at 0 °C) | ||
Melting point | −192.2 °C; −313.9 °F; 81.0 K | ||
Boiling point | −112 °C; −170 °F; 161 K | ||
1.05 g L−1 (at 0 °C) | |||
Solubility in other solvents | Very soluble in CCl4, sulfuric acid | ||
Vapor pressure | 55.7 atm[2] (−12.15 °C or 10.13 °F or 261.00 K)[lower-alpha 1] | ||
Conjugate acid | Protonated ozone | ||
+6.7·10−6 cm3/mol | |||
Refractive index (nD) |
1.2226 (liquid), 1.00052 (gas, STP, 546 nm—note high dispersion)[3] | ||
Structure | |||
C2v | |||
Digonal | |||
Dihedral | |||
Hybridisation | sp2 for O1 | ||
0.53 D | |||
Thermochemistry | |||
Std molar entropy (S⦵298) |
238.92 J K−1 mol−1 | ||
Std enthalpy of formation (ΔfH⦵298) |
142.67 kJ mol−1 | ||
Hazards | |||
GHS labelling: | |||
Danger | |||
H270, H314 | |||
NFPA 704 (fire diamond) | |||
Lethal dose or concentration (LD, LC): | |||
LCLo (lowest published) |
12.6 ppm (mouse, 3 hr) 50 ppm (human, 30 min) 36 ppm (rabbit, 3 hr) 21 ppm (mouse, 3 hr) 21.8 ppm (rat, 3 hr) 24.8 ppm (guinea pig, 3 hr) 4.8 ppm (rat, 4 hr)[4] | ||
NIOSH (US health exposure limits): | |||
PEL (Permissible) |
TWA 0.1 ppm (0.2 mg/m3)[1] | ||
REL (Recommended) |
C 0.1 ppm (0.2 mg/m3)[1] | ||
IDLH (Immediate danger) |
5 ppm[1] | ||
Related compounds | |||
Related compounds |
Sulfur dioxide Trisulfur Disulfur monoxide Cyclic ozone | ||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Ozone's odor is reminiscent of chlorine, and detectable by many people at concentrations of as little as 0.1 ppm 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. In standard conditions, 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.[5][6] It is therefore used commercially only in low concentrations.
Ozone is a powerful oxidant (far more so than dioxygen) and has many industrial and consumer applications related to oxidation. This same high oxidizing potential, however, causes ozone to damage mucous and respiratory tissues in animals, and also tissues in plants, above concentrations of about 0.1 ppm. While this makes ozone a potent respiratory hazard and pollutant near ground level, a higher concentration in the ozone layer (from two to eight ppm) is beneficial, preventing damaging UV light from reaching the Earth's surface.
The trivial name ozone is the most commonly used and preferred IUPAC name. The systematic names 2λ4-trioxidiene[dubious – discuss] and catena-trioxygen, valid IUPAC names, are constructed according to the substitutive and additive nomenclatures, respectively. The name ozone derives from ozein (ὄζειν), the Greek neuter present participle for smell,[7] 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 (or ozonide) is used, non-systematically, to refer to the substituent group (-OOO-). Care should be taken to avoid confusing the name of the group for the context-specific name for the ozone given above.
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 created ozone.[8][9]
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 (ὄζειν) meaning "to smell".[10][11] For this reason, Schönbein is generally credited with the discovery of ozone.[12][13][14][8] 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.[10] A subsequent effort to call ozone "electrified oxygen" he ridiculed by proposing to call the ozone from white phosphorus "phosphorized oxygen".[10] The formula for ozone, O3, was not determined until 1865 by Jacques-Louis Soret[15] and confirmed by Schönbein in 1867.[10][16]
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.[17] Naturalists working outdoors often considered the higher elevations beneficial because of their ozone content. "There is quite a different atmosphere [at higher elevation] with enough ozone to sustain the necessary energy [to work]", wrote naturalist Henry Henshaw, working in Hawaii.[18] 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[19] and dimethyl sulfide.[20]
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".[21][12] 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.[22] 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."[12][23]
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.[12][24]
Until the 1920s, it was not certain whether small amounts of oxozone, O
4, were also present in ozone samples due to the difficulty of applying analytical chemistry techniques to the explosive concentrated chemical.[25][26] In 1923, Georg-Maria Schwab (working for his doctoral thesis under Ernst Hermann Riesenfeld) was the first to successfully solidify ozone and perform accurate analysis which conclusively refuted the oxozone hypothesis.[25][26] Further hitherto unmeasured physical properties of pure concentrated ozone were determined by the Riesenfeld group in the 1920s.[25]
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 161 K (−112 °C; −170 °F), 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 80 K (−193.2 °C; −315.7 °F), it forms a violet-black solid.[27]
Most people can detect about 0.01 μmol/mol of ozone in air where it has a very specific sharp odour somewhat resembling chlorine bleach. Exposure of 0.1 to 1 μmol/mol produces headaches, burning eyes and irritation to the respiratory passages.[28] Even low concentrations of ozone in air are very destructive to organic materials such as latex, plastics and animal lung tissue.
Ozone is weakly diamagnetic.[29]
According to experimental evidence from microwave spectroscopy, ozone is a bent molecule, with C2v symmetry (similar to the water molecule).[30] The O–O distances are 127.2 pm (1.272 Å). The O–O–O angle is 116.78°.[31] The central atom is sp² hybridized with one lone pair. Ozone is a polar molecule with a dipole moment of 0.53 D.[32] 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 (16O, 17O, 18O). A cyclic form has been predicted but not observed.