Ozone
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For other uses, see Ozone (disambiguation).
Ozone
IUPAC name[hide]
Trioxygen
Identifiers
CAS number 10028-15-6 [Yes]
PubChem 24823
ChemSpider 23208 [Yes]
UNII 66H7ZZK23N [Yes]
EC number 233–069–2
MeSH Ozone
ChEBI CHEBI:25812 [Yes]
RTECS number RS8225000
Gmelin Reference 1101
Jmol-3D images Image 1
Image 2
SMILES
[show]
InChI
[show]
Properties
Molecular formula O3
Molar mass 48.00 g mol−1
Appearance Pale blue gas
Density 2.144 mg cm−3 (at 0 °C)
Melting point −192.2 °C; −313.9 °F; 81.0 K
Boiling point −112 °C; −170 °F; 161 K
Solubility in water 1.05 g L−1 (at 0 °C)
Solubility very soluble in CCl4,sulfuric acid
Refractive index (nD) 1.2226 (liquid)
Structure
Space group C2v
Coordination
geometry Digonal
Molecular shape Dihedral
Hybridisation sp2 for O1
Dipole moment 0.53 D
Thermochemistry
Std enthalpy of
formation ΔfHo298 142.67 kJ mol−1
Standard molar
entropy So298 238.92 J K−1 mol−1
Hazards
EU classification [Oxidising agent] O [Irritant] Xi
NFPA 704
[NFPA 704.svg]
0
4
4
OX
Related compounds
Related compounds Sulfur dioxide
Trisulfur
Disulfur monoxide
Cyclic ozone
[Yes] (verify) (what is: [Yes] / ?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
Infobox references
Ozone /ˈoʊzoʊn/ (systematically named 1λ1,3λ1-trioxidane and μ-oxidodioxygen), or trioxygen, is an inorganic compound with the chemical formula O
3(μ-O) (also written [O(μ-O)O] or 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 normal dioxygen. Ozone is formed from dioxygen by the action of ultraviolet light and also atmospheric electrical discharges, and is present in low concentrations throughout the Earth's atmosphere. In total, ozone makes up only 0.6 ppm of the atmosphere.
Ozone's odor is sharp, reminiscent of chlorine, and detectable by many people at concentrations of as little as 10 ppb in air. Ozone's O3 formula was determined in 1865. The molecule was later proven to have a bent structure and to be diamagnetic. In standard conditions, ozone is a pale blue gas that condenses at progressively cryogenic temperatures to a dark blue liquid and finally a violet-black solid.[1] Ozone's instability with regard to more common dioxygen is such that both concentrated gas and liquid ozone may decompose explosively.[1] 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 mucus and respiratory tissues in animals, and also tissues in plants, above concentrations of about 100 ppb. This makes ozone a potent respiratory hazard and pollutant near ground level. However, the so-called ozone layer (a portion of the stratosphere with a higher concentration of ozone, from two to eight ppm) is beneficial, preventing damaging ultraviolet light from reaching the Earth's surface, to the benefit of both plants and animals.
Contents
[hide]
1 Nomenclature
2 History
3 Physical properties
4 Structure
5 Reactions
Nomenclature
The trivial name ozone is the most commonly used and preferred IUPAC name. The systematic names 1λ1,3λ1-trioxidane and μ-oxidodioxygen, valid IUPAC names, are constructed according to the substitutive and additive nomenclatures, respectively. The nameozone derives from ozein (ὄζειν), the Greek word for smell (verb), 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 context-specific systematic name, according to substitutive nomenclature. By default, this name pay no regard to the radicality of the ozone molecule. In 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 ozone given above.
History[edit]
ozonometer, 1865
Ozone, the first allotrope of any chemical element to be recognized, was proposed as a distinct chemical substance by Christian Friedrich Schönbein in 1840, who named it after the Greek verb ozein (ὄζειν, "to smell"), from the peculiar odor in lightning storms.[2][3] The formula for ozone, O3, was not determined until 1865 by Jacques-Louis Soret[4] and confirmed by Schönbein in 1867.[2][5]
For much of the second half of the nineteenth century and well into the twentieth, ozone was considered a healthy component of the environment by naturalists and health-seekers. The City of Beaumont in California had as its official slogan "Beaumont: Zone of Ozone," as evidenced on postcards and Chamber of Commerce letterhead.[6] 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.[7] Seaside air was considered to be healthy because of its "ozone" content but the smell giving rise to this belief is in reality that of rotting seaweed.[8]
Physical properties[edit]
Ozone is colourless gas but bluish when liquified, slightly soluble in water and much more soluble in inert non-polar solvents such as carbon tetrachloride or fluorocarbons, where 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.[9]
Most people can detect about 0.01 μmol/mol of ozone in air where it has a very specific sharp odor somewhat resembling chlorine bleach. Exposure of 0.1 to 1 μmol/mol produces headaches, burning eyes and irritation to the respiratory passages.[10] Even low concentrations of ozone in air are very destructive to organic materials such as latex, plastics and animal lung tissue.
Ozone is diamagnetic, which means that its electrons are all paired. In contrast, O2 is paramagnetic, containing two unpaired electrons.
Structure[edit]
According to experimental evidence from microwave spectroscopy, ozone is a bent molecule, with C2v symmetry (similar to the water molecule). The O – O distances are 127.2 pm (1.272 Å). The O – O – O angle is 116.78°.[11] The central atom is sp² hybridized with one lone pair. Ozone is a polar molecule with a dipole moment of 0.53 D.[12] The bonding can be expressed as a resonance hybrid with a single bond on one side and double bond on the other producing an overall bond order of 1.5 for each side.
[Resonance Lewis structures of the ozone molecule]
Reactions[edit]
Ozone is a powerful oxidizing agent, far stronger than O2. It is also unstable at high concentrations, decaying to ordinary diatomic oxygen. It has a varying length half-life (meaning half as concentrated, or half-depleted), depending upon atmospheric conditions (temperature, humidity, and air movement). In a sealed chamber, with fan moving the gas, ozone has a half-life of approximately a day at room temperature[13] Some claims have been stated that ozone can have a half life as short as a half an hour in atmospheric conditions, although this claim is not verified by this reference:[14]
2 O
3 → 3 O
2
This reaction proceeds more rapidly with increasing temperature and increased pressure. Deflagration of ozone can be triggered by a spark, and can occur in ozone concentrations of 10 wt% or higher.[15]
With metals[edit]
Ozone will oxidize most metals (except gold, platinum, and iridium) to oxides of the metals in their highest oxidation state. For example:
2 Cu+
+ 2 H
3O+
+ O
3 → 2 Cu2+
+ 3 H
2O + O
2
