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Ozone
General
Systematic name	Trioxygen
Molecular formula	O3
Molar mass	47.998 g·mol
Appearance	bluish colored gas
CAS number
Properties
Density and phase	2.144 g·L (0 °C), gas
Solubility in water	0.105 g·100mL (0 °C)
Melting point	80.7 K, −192.5 °C
Boiling point	161.3 K, −111.9 °C
Thermodynamic data
Standard enthalpy of formation ΔfH°solid	+142.3 kJ·mol
Standard molar entropy S°solid	237.7 J·K.mol
Hazards
EU classification	not listed
NFPA 704
Supplementary data page
Structure and properties	n, εr, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Regulatory data	Flash point, RTECS number, etc.
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references

Ozone (O₃) is a triatomic molecule, consisting of three oxygen atoms. It is an allotrope of oxygen that is much less stable than the diatomic species O₂. Ground-level ozone is an air pollutant with harmful effects on our respiratory system. On the other hand, ozone in the upper atmosphere protects living organisms by preventing damaging ultraviolet light from reaching the Earth's surface. It is present in low concentrations throughout the Earth's atmosphere. It has many industrial and consumer applications as well as being used in ozone therapy.

Ozone, the first allotrope of a chemical element to be described by science, was discovered by Christian Friedrich Schönbein in 1840, who named it after the Greek word for smell (ozein), from the peculiar odour in lightning storms. The odour from a lightning strike is from ions produced during the rapid chemical changes, not the ozone itself.

Physical properties

Undiluted ozone is a pale blue gas at standard temperature and pressure; it forms a dark blue liquid below −112 °C and a violet-black solid below −193 °C . At concentrations found in the atmosphere it is colorless . The concentration above which it can be smelt (odor threshold) is between 0.0076 and 0.036 ppm.

Structure

The structure of ozone, according to experimental evidence from microwave spectroscopy, is bent, with C_2v symmetry (similar to the water molecule), O–O distance of 127.2 pm and O–O–O angle of 116.78°. Ozone is a polar molecule with a dipole moment of 0.5337 D.

Cyclic ozone

An alternative molecular structure for ozone is the cyclic structure with D_3h symmetry, similar to cyclopropane. The cyclic structure has been studied theoretically using ab initio quantum chemistry methods, and most calculations agree that it is less stable than the bent structure by about 30 kcal/mol, with a ring-opening barrier of about 20 kcal/mol. This structure would look much like a equilateral triangle with all the 3 oxygen atoms bonded to one another. Unfortunately, this structure has never been observed in real life, let alone synthesized. Currently, research is being done to synthesize cyclic ozone molecules by hitting it with ultrafast lasers (http://www.physorg.com/news2904.html). Creation and isolation of such cyclic molecules could allow more energy to be packed into rockets and hence may allow for farther space travel.

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Chemistry

Ozone is a powerful oxidizing agent. It is also unstable at high concentrations, decaying to ordinary diatomic oxygen:

2 O₃ → 3 O₂.

This reaction proceeds more rapidly with increasing temperature and decreasing pressure. Ozone will oxidize metals (except gold, platinum, and iridium) to oxides of the metals in their highest oxidation state:

2 Cu + 2 H + O₃ → 2 Cu + H₂O + O₂

Ozone converts oxides to peroxides:

SO₂ + O₃ → SO₃ + O₂

It also increases the oxidation number of oxides:

NO + O₃ → NO₂ + O₂

The above 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:

NO₂ + NO₃ → N₂O₅

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

C + 2 O₃ → CO₂ + 2 O₂

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

2 NH₃ + 4 O₃ → NH₄NO₃ + 4 O₂ + H₂O

Ozone reacts with sulfides to make sulfates:

PbS + 4 O₃ → PbSO₄ + 4 O₂

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

S + H₂O + O₃ → H₂SO₄

3 SO₂ + 3 H₂O + O₃ → 3 H₂SO₄

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

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

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

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

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

2 NO₂ + 2 ClO₂ 2 O₃ → 2 NO₂ClO₄ + O₂

Ozone can be used for combustion reactions and combusting gases in ozone provides higher temperatures than combusting in dioxygen (O₂). Following is a reaction for the combustion of carbon subnitride:

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:

H + O₃ → HO₂ + O

2 HO₂ → H₂O₄

Ozonides can be formed, which contain the ozonide anion, O₃. These compounds are explosive and must be stored at cryogenic temperatures. Ozonides for all the alkali metals are known. KO₃, RbO₃, and CsO₃ can 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:

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:

CsO₃ + Na → Cs + NaO₃

Treatment with ozone of calcium dissolved in ammonia leads to ammonium ozonide and not calcium ozonide:

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 the water, forming a precipitate which can be filtered:

2 Mn + 2 O₃ + 4 H₂O → 2 MnO(OH)_{2 (s)} + 2 O₂ + 4 H

Ozone will also turn cyanides to the one thousand times less toxic cyanates:

CN + O₃ → CNO + O₂

Finally, ozone will also completely decompose urea:

(NH₂)₂CO + O₃ → N₂ + CO₂ + 2 H₂O

Ozone in Earth's atmosphere

The standard way to express total ozone levels (the volume of ozone in a vertical column) in the atmosphere is by using Dobson units. Concentrations at a point are measured in parts per billion (ppb) or in μg/m³.

Ozone layer

The highest levels of ozone in the atmosphere are in the stratosphere, in a region also known as the ozone layer between about 10 km and 50 km above the surface. Here it filters out the shorter wavelengths (less than 320 nm) of ultraviolet light (270 to 400 nm) from the Sun that 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, which is essential for human health. Ozone in the stratosphere is mostly produced from ultraviolet rays reacting with oxygen:

O₂ + (radiation < 240 nm) → 2 O

O + O₂ → O₃

It is destroyed by the reaction with atomic oxygen:

O₃ + O → 2 O₂

(See Ozone-oxygen cycle for more detail.)

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 due to emissions of CFCs and similar chlorinated and brominated organic molecules, which have increased the concentration of ozone-depleting catalysts above the natural background. See ozone depletion for more information. For more information on stratospheric ozone see Seinfeld and Pandis (1999).

Low level ozone

Low level ozone (or tropospheric ozone) is regarded as a pollutant by the World Health Organization. It is not emitted directly by car engines or by industrial operations. It is formed by the reaction of sunlight on air containing hydrocarbons and nitrogen oxides that react to form ozone directly at the source of the pollution or many kilometers down wind. For more details of the complex chemical reactions that produce low level ozone see tropospheric ozone or Seinfled and Pandis (1998).

Ozone reacts directly with some hydrocarbons such as aldehydes and thus begins their removal from the air, but the products are themselves key components of smog. Ozone photolysis by UV light leads to production of the hydroxyl radical and this plays a part in the removal of hydrocarbons from the air, but is also the first step in the creation of components of smog such as peroxyacyl nitrates which can be powerful eye irritants. The atmospheric lifetime of tropospheric ozone is about 22 days and its main removal mechanisms are being deposited to the ground, the above mentioned reaction giving OH, and by reactions with OH and the peroxy radical HO₂· (Stevenson et al, 2006) .

As well as having an impact on human health (see below) there is also evidence of significant reduction in agricultural yields due to increased ground-level ozone and pollution which interferes with photosynthesis and stunts overall growth of some plant species.

Ozone as a greenhouse gas

Although ozone was present at ground level before the industrial revolution, peak concentrations are far higher than the pre-industrial levels and even background concentrations well away from sources of pollution are substantially higher. This increase in ozone is of further concern as ozone present in the upper troposphere acts as a greenhouse gas, absorbing some of the infrared energy emitted by the earth. Quantifying the greenhouse gas potency of ozone is difficult as it is not present in uniform concentrations across the globe. However, the most recent scientific review on the climate change (the IPCC Third Assessment Report) suggests that the radiative forcing of tropospheric ozone is about 25% that of carbon dioxide.

Ozone and health

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Ozone in air pollution

There is a great deal of evidence to show that high concentrations (ppm) of ozone, created by high concentrations of pollution and daylight UV rays at the earth's surface, can harm lung function and irritate the respiratory system . There has also been shown to be a connection between increased ozone caused by thunderstorms and hospital admissions of asthma sufferers . Air quality guidelines such as those from the World Health Organization are based on detailed studies of what levels can cause measurable health effects.

A common British folk myth dating back to the Victorian era holds that the smell of the sea is caused by ozone, and that this smell has "bracing" health-giving effects . Neither of these is true. The characteristic "smell of the sea" is not caused by ozone, but by the presence of dimethyl sulfide generated by phytoplankton, and dimethyl sulfide, like ozone, is toxic in high concentrations.

Physiology of ozone

Ozone, along with reactive forms of oxygen such as superoxide, singlet oxygen (see oxygen), hydrogen peroxide, and hypochlorite ions, is naturally produced by white blood cells and other biological systems (such as the roots of marigolds) as a means of destroying foreign bodies. Ozone reacts directly with organic double bonds. Also, when ozone breaks down to dioxygen it gives rise to oxygen free radicals, which are highly reactive and capable of damaging many organic molecules. Ozone has been found to convert cholesterol in the blood stream to plaque (which causes hardening and narrowing of arteries). Moreover, it is believed that the powerful oxidizing properties of ozone may be a contributing factor of inflammation. The cause-and-effect relationship of how the ozone is created in the body and what it does is still under consideration and still subject to various interpretations, since other body chemical processes can trigger some of the same reactions. A team headed by Dr. Paul Wentworth Jr. of the Department of Chemistry at the Scripps Research Institute has shown evidence linking the antibody-catalyzed water-oxidation pathway of the human immune response to the production of ozone. In this system, ozone is produced by antibody-catalyzed production of trioxidane from water and neutrophil-produced singlet oxygen. . See also trioxidane for more on this biological ozone-producing reaction.

Ozone has also been proven to form specific, cholesterol-derived metabolites that are thought to facilitate the build-up and pathogenesis of atherosclerotic plaques (A form of heart disease). These metabolites have been confirmed as naturally occurring in human atherosclerotic arteries and are categorized into a class of secosterols termed “Atheronals”, generated by ozonolysis of cholesterol's double bond to form a 5,6 secosterol as well as a secondary condensation product via aldolization. Volume: Number: Page: 23 DOI:

Safety

Artificial production

Ozone may be formed from O₂ by electrical discharges and by action of high energy electromagnetic radiation. Certain electrical equipment generate significant levels of ozone. This is especially true of devices using high voltages, such as ionic air purifiers, laser printers, photocopiers, and arc welders. Electric motors using brushes can generate ozone from repeated sparking inside the unit. Large motors that use brushes, such as those used by elevators (most elevator motors don't have brushes) or hydraulic pumps, will generate more ozone than smaller motors.

Industrial production

Ozone used in industry is measured in g/Nm or weight percent. The regime of applied concentrations ranges from 1 to 5 weight percent in air and from 6 to 13 weight percent in oxygen.

Formation and enrichment of ozone is obtained by exposure of an oxygen carrying gas to plasma, which is made of so-called silent or dielectric barrier discharges (DBD). Basically, molecular oxygen is dissociated into atomic oxygen, which subsequently recombines to ozone. The discharges manifest as filamentary transfer of electrons (micro discharges) in a gap between two electrodes. In order to evenly distribute the micro discharges, a dielectric insulator must be used to separate the metallic electrodes and to prevent arcing.

Ozone cannot be stored and transported like other industrial gases and must therefore be produced on site. Available ozone generators vary in the arrangement and design of the high-voltage electrodes. At production capacities higher than 20kg per hour, a gas/water tube heat-exchanger is utilized as ground electrode and assembled with tubular high-voltage electrodes on the gas-side. The regime of typical gas pressures is around 2bar absolute in oxygen and 3bar absolute in air. Several megawatt of electrical power may be installed in large facilities, applied as one phase AC current at 600 to 2000Hz and peak voltages between 3000 and 20000 volts.

The dominating parameter influencing ozone generation efficiency is the gas temperature, which is controlled by the cooling water temperature. The cooler the water, the better the ozone synthesis. At typical industrial conditions, almost 90 percent of the effective power is dissipated as heat and needs to be removed by a sufficient cooling water flow.

Due to the high reactivity of ozone, only few materials may be used like stainless steel (quality 316L), glass, polytetrafluorethylene, or polyvinylidene fluoride. Viton may be used with the restriction of constant mechanical forces and absence of humidity.

Laboratory production

In the laboratory ozone can be produced by electrolysis using a 9 volt battery, a pencil graphite rod cathode, a platinum wire anode and a 3M sulfuric acid electrolyte. The half cell reactions taking place are:

3 H₂O → O₃ + 6 H + 6 e ΔE = − 1.53 V

6 H + 6 e → 3 H₂ ΔE = 0 V

2 H₂O → O₂ + 4 H + 4 e ΔE = −1. 23 V

So that in the net reaction three equivalents of water are converted into one equivalent of ozone and three equivalents of hydrogen. Oxygen formation is a competing reaction.

Applications

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Industrial applications

Ozone can be used for bleaching substances and for killing bacteria. Many municipal drinking water systems kill bacteria with ozone instead of the more common chlorine. Ozone has a very high oxidation potential. Ozone does not form organochlorine compounds, but it also does not remain in the water after treatment, so some systems introduce a small amount of chlorine to prevent bacterial growth in the pipes, or may use chlorine intermittently, based on results of periodic testing. Where electrical power is abundant, ozone is a cost-effective method of treating water, as it is produced on demand and does not require transportation and storage of hazardous chemicals. Once it has decayed, it leaves no taste or odor in drinking water. Low level of Ozone is helpful to purify air inside the house.

Industrially, ozone or ozonated water is used to

• Disinfect laundry in hospitals, food factories, care homes etc;
• disinfect water before it is bottled;
• deodorize air and objects, such as after a fire;
• kill bacteria on food or on contact surfaces;
• scrub yeast and mold spores from the air in food processing plants;
• wash fresh fruits and vegetables to kill yeast, mold and bacteria;
• chemically attack contaminants in water (iron, arsenic, hydrogen sulfide, nitrites, and complex organics lumped together as "color");
• provide an aid to flocculation (agglomeration of molecules, which aids in filtration, where the iron and arsenic are removed);
• manufacture chemical compounds via chemical synthesis
• clean and bleach fabrics (the latter use is patented);
• assist in processing plastics to allow adhesion of inks;
• age rubber samples to determine the useful life of a batch of rubber;
• hospital operating rooms where air needs to be sterile;
• eradicate water borne parasites such as Giardia and Cryptosporidiumin surface water treatment plants. This process is known as ozonation.

Ozone is a reagent in many organic reactions in the laboratory and in industry. Ozonolysis is the cleavage of an alkene to carbonyl compounds.

Many hospitals in the U.S. and around the world use large ozone generators to decontaminate operating rooms between surgeries. The rooms are cleaned and then sealed airtight before being filled with ozone which effectively kills or neutralizes all remaining bacteria.

Consumer applications

Ozone machines, with or without ionization, are currently used to sanitize (high ozone output) and deodorize non-inhabited rooms, ductwork, vehicles, boats, woodsheds, and buildings.

Some models of air purifiers that also emit low levels of ozone have been sold in the US. These type of air purifiers claim to imitate nature's "filterless" air purifying mechanisms and claim to "sanitize" the air and/or household surfaces. The government successfully sued one company in 1995, ordering them to stop repeating health claims without supporting scientific studies. Ozone generators are typically used in hospital operating rooms where the air needs to be sterile.

Ozonated water is used to launder clothes, sanitize food, drinking water, and surfaces in the home. According to the FDA, it is "amending the food additive regulations to provide for the safe use of ozone in gaseous and aqueous phases as an antimicrobial agent on food, including meat and poultry." Studies at California Polytechnic University, have proven that low levels of ozone dissolved in filtered tapwater can produce more than a four-log (99.99%) reduction in such food-borne microorganisms as salmonella, e. Coli 0157:H7, campylobacter and others. Ironically, while ozone is considered an atmospheric pollutant, pollution and smog by the US government, it can actually reduce pollutants like pesticides in fruits and vegetables.

Ozone is used in spas or hot tubs with reduced levels of chlorine or bromine for keeping the water free of bacteria. As it does not remain in the water after treatment, it is ineffective at preventing bather cross-contamination, and must be used in conjunction with another sanitizer. Ozone gas is created by an ultraviolet light bulb or corona discharge chip and injected into the plumbing system.

Ozone is also widely used in treatment of water in aquaria and fish ponds. Its use can minimise bacterial growth control parasites and removes or reduce "yellowing" of the water. As the Ozone rapidly decomposes, at correctly controlled levels the application has no effect on the fish.

Pharmaceutical applications

Ozone has been used in alternative medicine as a medical treatment in a number of different countries. Its use however is controversial.

See also

• Ozone depletion, including the phenomenon known as the ozone hole.
• Ozone layer
• Tropospheric ozone
• Tetraoxygen (O₄)

References

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2. "Ozone FAQ". Global Change Master Directory. Retrieved 2006-05-10.
3. "Oxygen". WebElements. Retrieved 2006-09-23.
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5. "Ozone". Haz-MAP (occupational health database}. Retrieved 2006-09-23.
6. Takehiko Tanaka; Yonezo Morino. Coriolis interaction and anharmonic potential function of ozone from the microwave spectra in the excited vibrational states Journal of Molecular Spectroscopy 1970, 33, 538-551.
7. Kenneth M. Mack; J. S. Muenter. Stark and Zeeman properties of ozone from molecular beam spectroscopy. Journal of Chemical Physics 1977, 66, 5278-5283. doi:10.1063/1.433909
8. Sotiris S. Xantheas; Gregory J. Atchity; Stephen T. Elbert; Klaus Ruedenberg. Potential energy surfaces of ozone. Journal of Chemical Physics 1991, 94, 8054-8069. doi:10.1063/1.460140
9. Horvath M., Bilitzky L., & Huttner J., 1985. "Ozone." pg 44-49
10. Housecroft & Sharpe, 2005. "Inorganic Chemistry." pg 439
11. Housecroft & Sharpe, 2005. "Inorganic Chemistry." pg 265
12. Horvath M., Bilitzky L., & Huttner J., 1985. "Ozone." pg 44-49
13. Horvath M., Bilitzky L., & Huttner J., 1985. "Ozone." pg 259, 269-270
14. ^ WHO-Europe reports: Health Aspects of Air Pollution (2003) (PDF)
15. Stevenson; et al. (2006). "Multimodel ensemble simulations of present-day and near-future tropospheric ozone". American Geophysical Union. Retrieved 2006-09-16. {{cite web}}: Explicit use of et al. in: |author= (help)
16. "Rising Ozone Levels Pose Challenge to U.S. Soybean Production, Scientists Say". NASA Earth Observatory. 2003-07-31. Retrieved 2006-05-10. {{cite web}}: Check date values in: |date= (help)
17. Mutters, Randall (1999). "Statewide Potential Crop Yield Losses From Ozone Exposure". California Air Resources Board. Retrieved 2006-05-10. {{cite web}}: Unknown parameter |month= ignored (help)
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23. Ashfield District Council: Monitored Air Pollutants, downloaded February 2, 2007
24. University of East Anglia press release, Cloning the smell of the seaside, February 2, 2007
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• Seinfeld, John H.; Pandis, Spyros N (1998). Atmospheric Chemistry and Physics - From Air Pollution to Climate Change. John Wiley and Sons, Inc. ISBN 0-471-17816-0
• Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
• Series in Plasma Physics: Non-Equilibrium Air Plasmas at Atmospheric Pressure. Edited by K.H. Becker, U. Kogelschatz, K.H. Schoenbach, R.J. Barker; Bristol and Philadelphia: Institute of Physics Publishing Ltd; ISBN 0-7503-0962-8; 2005

External links

• NASA's Ozone Resource Page
• Paul Crutzen Interview Freeview video of Paul Crutzen Nobel Laureate for his work on decomposition of ozone talking to Harry Kroto Nobel Laureate by the Vega Science Trust.
• NASA's Earth Observatory article on Ozone
• International Day for the Preservation of the Ozone Layer
• International Chemical Safety Card 0068
• NIOSH Pocket Guide to Chemical Hazards
• National Institute of Environmental Health Sciences Ozone Alerts
• Ground-level Ozone Air Pollution — A summary for non specialists by GreenFacts of the above WHO reports.
• NASA Study Links "Smog" to Arctic Warming— NASA Goddard Institute for Space Studies (GISS) study shows the warming effect of ozone in the Arctic during winter and spring.
• EPA Assessment of Effectiveness and Health Consequences of Ozone Generators that are Sold as Air Cleaners

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• Articles lacking sources from October 2006
• Environmental chemistry
• Oxygen compounds
• Ozone depletion
• Reagents for organic chemistry
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