Ozone, the first allotrope of a 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.The formula for ozone, O3, was not determined until 1865 byJacques-Louis Soret and confirmed by Schönbein in 1867.

Physical properties

Ozone is a pale blue gas, slightly soluble in water and much more soluble in inert non-polar solvents such ascarbon tetrachloride or fluorocarbons, where it forms a blue solution. At 161 K (–112 °C), 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 °C), it forms a violet-black solid.

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. 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.


According to experimental evidence from microwave spectroscopy, ozone is a bent molecule, with C2vsymmetry (similar to the water molecule). The O – O distances are 127.2 pm. The O – O – O angle is 116.78°.[9] The central atom is sp² hybridized with one lone pair. Ozone is a polar molecule with adipole moment of 0.53 D.[10] The bonding can be expressed as a resonance hybrid with a single bondon 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


Ozone is a powerful oxidizing agent, far stronger than O2. It is also unstable at high concentrations, decaying to ordinary diatomic oxygen (with a half-life of about half an hour in atmospheric conditions):[11]

    2 O3 → 3 O2

This reaction proceeds more rapidly with increasing temperature and increased pressure. Deflagrationof ozone can be triggered by a spark, and can occur in ozone concentrations of 10 wt% or higher.[12]

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 H3O+ + O3 → 2 Cu2+ + 3 H2O + O2

With nitrogen and carbon compound

Ozone also oxidizes nitric oxide to nitrogen dioxide:

    NO + O3 → NO2 + O2

This reaction is accompanied by chemiluminescence. The NO2 can be further oxidized:

    NO2 + O3 → NO3 + O2

The NO3 formed can react with NO2 to form N2O5:

Solid nitryl perchlorate can be made from NO2, ClO2, and O3 gases:

    2 NO2 + 2 ClO2 + 2 O3 → 2 NO2ClO4 + O2

Ozone does not react with ammonium salts but it oxidizes with ammonia to ammonium nitrate:

    2 NH3 + 4 O3 → NH4NO3 + 4 O2 + H2O

Ozone reacts with carbon to form carbon dioxide, even at room temperature:

    C + 2 O3 → CO2 + 2 O2

With sulfur compounds

Ozone oxidizes sulfides to sulfates. For example, lead(II) sulfide is oxidised to lead(II) sulfate:

    PbS + 4 O3 → PbSO4 + 4 O2

Sulfuric acid can be produced from ozone, water and either elemental sulfur or sulfur dioxide:

    S + H2O + O3 → H2SO4
    3 SO2 + 3 H2O + O3 → 3 H2SO4

In the gas phase, ozone reacts with hydrogen sulfide to form sulfur dioxide:

    H2S + O3 → SO2 + H2O

In an aqueous solution, however, two competing simultaneous reactions occur, one to produce elemental sulfur, and one to produce sulfuric acid:

    H2S + O3 → S + O2 + H2O
    3 H2S + 4 O3 → 3 H2SO4

Other substratesAll three atoms of ozone may also react, as in the reaction of tin(II) chloride with hydrochloric acid and ozone:

    3 SnCl2 + 6 HCl + O3 → 3 SnCl4 + 3 H2O

Iodine perchlorate can be made by treating iodine dissolved in cold anhydrous perchloric acid with ozone:

    I2 + 6 HClO4 + O3 → 2 I(ClO4)3 + 3 H2O


Ozone can be used for combustion reactions and combusting gases; ozone provides higher temperatures than combusting in dioxygen (O2). The following is a reaction for the combustion ofcarbon subnitride which can also cause higher temperatures:

    3 C4N2 + 4 O3 → 12 CO + 3 N2

Ozone can react at cryogenic temperatures. At 77 K (−196 °C), atomic hydrogen reacts with liquid ozone to form a hydrogen superoxide radical, which dimerizes:[13]

    H + O3 → HO2 + O
    2 HO2 → H2O4

Reduction to ozonidesReduction of ozone gives the ozonide anion, O3– . Derivatives of this anion are explosive and must be stored at cryogenic temperatures. Ozonides for all the alkali metals are known. KO3, RbO3, and CsO3can be prepared from their respective superoxides:

    KO2 + O3 → KO3 + O2

Although KO3 can be formed as above, it can also be formed from potassium hydroxide and ozone:[14]

    2 KOH + 5 O3 → 2 KO3 + 5 O2 + H2O

NaO3 and LiO3 must be prepared by action of CsO3 in liquid NH3 on an ion exchange resin containing Na+ or Li+ ions:[15]

    CsO3 + Na+ → Cs+ + NaO3

A solution of calcium in ammonia reacts with ozone to give to ammonium ozonide and not calcium ozonide:[13]

    3 Ca + 10 NH3 + 6 O3 → Ca·6NH3 + Ca(OH)2 + Ca(NO3)2 + 2 NH4O3 + 2 O2 + H2


Ozone can be used to remove manganese from water, forming a precipitate which can be filtered:

    2 Mn2+ + 2 O3 + 4 H2O → 2 MnO(OH)2 (s) + 2 O2 + 4 H+

Ozone will also detoxify cyanides by converting it to cyanate, which is a thousand times less toxic.

    CN- + O3 → CNO− + O2

Ozone will also completely decompose urea:[16]

    (NH2)2CO + O3 → N2 + CO2 + 2 H2O

Ozone will cleave alkenes to form carbonyl compounds in the ozonolysis process.

    A generalized scheme of ozonolysis

Ozone in Earth's atmosphereThe standard way to express total ozone levels (the amount of ozone in a vertical column) in the atmosphere is by usingDobson units. Point measurements are reported as mole fractions in nmol/mol (parts per billion, ppb) or asconcentrations in μg/m3.

Ozone layer
Main article: Ozone layer

The highest levels of ozone in the atmosphere are in thestratosphere, in a region also known as the ozone layerbetween about 10 km and 50 km above the surface (or between about 6 and 31 miles). Here it filters out photonswith shorter wavelengths (less than 320 nm) of ultraviolet light, also called UV rays, (270 to 400 nm) from the Sunthat would be harmful to most forms of life in large doses. These same wavelengths are also among those responsible for the production of vitamin D in humans. Ozone in the stratosphere is mostly produced from ultraviolet rays reacting with oxygen:

    O2 + photon (radiation < 240 nm) → 2 O

    O + O2 + M → O3 + M

It is dsetroyed by the reaction with atomic oxygen:

    O3 + O → 2 O2

The latter reaction is catalysed by the presence of certain free radicals, of which the most important are hydroxyl (OH), nitric oxide (NO) and atomic chlorine (Cl) and bromine (Br). In recent decades the amount of ozone in the stratosphere has been declining mostly because of emissions of CFCs and similar chlorinated and brominated organic molecules, which have increased the concentration of ozone-depleting catalysts above the natural background. Ozone only makes up 0.00006% of the atmosphere.

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