Ozone
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Ozone or trioxygen (O3) is a triatomic molecule, consisting of three oxygen atoms. It is an allotrope of oxygen that is much less stable than the diatomic O2. Ground-level ozone is an air pollutant with harmful effects on the respiratory systems of animals and humans. Ozone in the upper atmosphere filters potentially 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.
Ozone, the first allotrope of a chemical element to be recognized by science, was proposed as a distinct chemical compound by Christian Friedrich Schönbein in 1840, who named it after the Greek word for smell (ozein), from the peculiar odor in lightning storms.<ref name="ozo">{{#if:Rubin
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Contents |
Physical properties
Most people can detect about 0.01 ppm in air. Exposure of 0.1 to 1 ppm produces headaches, burning eyes, and irritation to the respiratory passages.<ref name=brown>Template:Cite book</ref>
Ozone is 1.5 times as dense as oxygen. At -112 °C, it forms a dark blue liquid. At temperatures below -193 °C, it forms a violet-black solid.<ref>Template:Cite web</ref>
Ozone is diamagnetic, meaning that it will resist formation of a magnetic field and will decrease the energy stored in the field once the field is established.
Structure
The structure of ozone, according to experimental evidence from microwave spectroscopy, is bent, with C2v symmetry (similar to the water molecule), O – O distance of 127.2 pm and O – O – O angle of 116.78°.<ref>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.</ref> The central atom forms an sp² hybridization with one lone pair. Ozone is a polar molecule with a dipole moment of 0.5337 D.<ref>Kenneth M. Mack; J. S. Muenter. Stark and Zeeman properties of ozone from molecular beam spectroscopy. Journal of Chemical Physics 1977, 66, 5278–5283. Template:Doi</ref> The bonding is single bond on one side and double bond on the other side, and these bonds are blended to become known as resonance structures. The bond order is 1.5 for each side.
Chemistry
Ozone is a powerful oxidizing agent, far better than dioxygen. It is also unstable at high concentrations, decaying to ordinary diatomic oxygen (in about half an hour in atmospheric conditions<ref>Earth Science FAQ: Where can I find information about the ozone hole and ozone depletion? Goddard Space Flight Center, National Aeronautics and Space Administration, March 2008.</ref>):
- 2 O3 → 3 O2
This reaction proceeds more rapidly with increasing temperature and decreasing pressure. Deflagration of ozone can be triggered by a spark, and can occur in ozone concentrations of 10 wt% or higher.
Metals
Ozone will oxidize metals (except gold, platinum, and iridium) to oxides of the metals in their highest oxidation state:
- 2 Cu1+(aq) + 2 H3O+(aq) + O3(g) → 2 Cu2+(aq) + 3 H2O(l) + O2(g)
Non-metals
Ozone also increases the oxidation number of oxides:
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:
Sulfuric acid can be produced from ozone, starting either 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 and NaCl along with Ammonium Nitrate:
- 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:
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
Combustion
Ozone can be used for combustion reactions and combusting gases; ozone provides higher temperatures than combusting in dioxygen (O2). Following is a reaction for the combustion of carbon subnitride which can also cause lower 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:<ref>Horvath M., Bilitzky L., & Huttner J., 1985. "Ozone." pg 44–49</ref>
- H + O3 → HO2 + O
- 2 HO2 → H2O4
Ozonides
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:<ref>Housecroft & Sharpe, 2005. "Inorganic Chemistry." pg 439</ref>
- 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:<ref>Housecroft & Sharpe, 2005. "Inorganic Chemistry." pg 265</ref>
- CsO3 + Na+ → Cs+ + NaO3
Treatment with ozone of calcium dissolved in ammonia leads to ammonium ozonide and not calcium ozonide:<ref>Horvath M., Bilitzky L., & Huttner J., 1985. "Ozone." pg 44–49</ref>
- 3 Ca + 10 NH3 + 6 O3 → Ca•6NH3 + Ca(OH)2 + Ca(NO3)2 + 2 NH4O3 + 2 O2 + H2
Applications
Ozone can be used to remove manganese from the 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 turn cyanides to the one thousand times less toxic cyanates:
- CN- + O3 → CNO- + O2
Finally, ozone will also completely decompose urea:<ref>Horvath M., Bilitzky L., & Huttner J., 1985. "Ozone." pg 259, 269–270</ref>
- (NH2)2CO + O3 → N2 + CO2 + 2 H2O
Ozone in Earth's atmosphere
The standard way to express total ozone levels (the amount 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 (or between about 6 and 31 miles). Here it filters out photons with shorter wavelengths (less than 320 nm) of ultraviolet light, also called UV rays, (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, a vitamin also produced by the human body. Ozone in the stratosphere is mostly produced from ultraviolet rays reacting with oxygen:
- O2 + photon(radiation< 240 nm) → 2 O
- O + O2 → O3
It is destroyed 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.
Low level ozone
Low level ozone (or tropospheric ozone) is regarded as a pollutant by the World Health Organization<ref name=who-Europe>WHO-Europe reports: Health Aspects of Air Pollution (2003) (PDF)</ref> and the United States Environmental Protection Agency (EPA). 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.
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 OH 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; 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).<ref>Template:Cite web</ref>
There is evidence of significant reduction in agricultural yields because of increased ground-level ozone and pollution which interferes with photosynthesis and stunts overall growth of some plant species.<ref>Template:Cite web</ref><ref>Template:Cite web</ref>
Certain examples of cities with elevated ozone readings are Houston, Texas, and Mexico City, Mexico. Houston has a reading of around 41 ppb, while Mexico City is far more hazardous, with a reading of about 125 ppb.<ref>Template:Cite web</ref>
Ozone as a greenhouse gas
Although ozone was present at ground level before the Industrial Revolution, peak concentrations are now far higher than the pre-industrial levels, and even background concentrations well away from sources of pollution are substantially higher.<ref>Template:Cite web</ref><ref>Template:Cite web</ref> This increase in ozone is of further concern because 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 because it is not present in uniform concentrations across the globe. However, the scientific review on the climate change (the IPCC Third Assessment Report<ref>Template:Cite web</ref>) suggests that the radiative forcing of tropospheric ozone is about 25% that of carbon dioxide.
Ozone cracking
Ozone gas attacks any polymer possessing olefinic or double bonds within its chain structure, such materials including natural rubber, nitrile rubber, and Styrene-butadiene rubber. Products made using these polymers are especially susceptible to attack, which causes cracks to grow longer and deeper with time, the rate of crack growth depending on the load carried by the product and the concentration of ozone in the atmosphere. Such materials can be protected by adding antiozonants, such as waxes, which bond to the surface to create a protective film or blend with the material and provide long term protection. Ozone cracking used to be a serious problem in car tires for example, but the problem is now seen only in very old tires. On the other hand, many critical products like gaskets and O-rings may be attacked by ozone produced within compressed air systems. Fuel lines are often made from reinforced rubber tubing and may also be susceptible to attack, especially within engine compartments where low levels of ozone are produced from electrical equipment.
Health effects
Air pollution
There is a great deal of evidence to show that high concentrations 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.<ref name="who-Europe"/><ref>Answer to follow-up questions from CAFE (2004) (PDF)</ref> A connection has also been shown to exist between increased ozone caused by thunderstorms and hospital admissions of asthma sufferers.<ref> {{#if:Anderson
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The Clean Air Act directs the EPA to set National Ambient Air Quality Standards for several pollutants, including ground-level ozone, and counties out of compliance with these standard are required to take steps to reduce their levels. In May 2008, the EPA lowered its ozone standard from 80 ppb to 75 ppb. This proved controversial, since the Agency's own scientists and advisory board had recommended lowing the standard to 60 ppb, and the World Health Organization recommends 51 ppb. Many public health and environmental groups also supported the 60 ppb standard. On the other hand, the EPA had already designated over 300 mostly urban counties as out of compliance, and lowering the standard to 75 ppb put hundreds more in non-compliance. Lowering it further to 60 ppb would likely have left most of the US in non-compliance. Manufacturers, employers, and others argued that the cost of compliance with the lower standard would be prohibitive.<ref name="pmid18629332"/> The EPA has also developed an Air Quality Index to help explain air pollution levels to the general public. Eight-hour average ozone concentrations of 85 to 104 ppbv are described as "Unhealthy for Sensitive Groups", 105 ppbv to 124 ppbv as "unhealthy" and 125 ppb to 404 ppb as "very unhealthy".<ref>Template:Cite web</ref>
Ozone can also be present in indoor air pollution.
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 benefits.<ref>Ashfield District Council: Monitored Air Pollutants, downloaded February 2, 2007</ref> 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.<ref>University of East Anglia press release, Cloning the smell of the seaside, February 2, 2007</ref>
Physiology
Template:Seealso Ozone, along with reactive forms of oxygen such as superoxide, singlet 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.<ref> {{#if:Hoffmann
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}}
}}
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| .
}}{{#if:Hoffmann2004
|
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}}{{#if:American Scientist
|. American Scientist
}}{{#if:92
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}}{{#if:1
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}}{{#if:23
|: 23
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When inhaled, ozone reacts with compounds lining the lungs 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<ref>{{#if:Smith
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|{{#ifeq:Oxygen, oxysterols, ouabain, and ozone: a cautionary tale||Citation is missing a title.
Either specify one, or click here and a bot will try to complete the citation details for you. }}[{{#if:http://toxnet.nlm.nih.gov/cgi-bin/sis/search/r?dbs+toxline:@term+@mh+%22+5,6-secosterol+%22+@OR+@na+%22+5,6-secosterol+%22+@OR+@ab+%22+5,6-secosterol+%22+@OR+@kw+%22+5,6-secosterol+%22|http://toxnet.nlm.nih.gov/cgi-bin/sis/search/r?dbs+toxline:@term+@mh+%22+5,6-secosterol+%22+@OR+@na+%22+5,6-secosterol+%22+@OR+@ab+%22+5,6-secosterol+%22+@OR+@kw+%22+5,6-secosterol+%22|http://www.pubmedcentral.gov/articlerender.fcgi?tool=pmcentrez&artid={{{pmc}}}}} Oxygen, oxysterols, ouabain, and ozone: a cautionary tale]
|Oxygen, oxysterols, ouabain, and ozone: a cautionary tale
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}}{{#if:Free radical biology & medicine
|. Free radical biology & medicine
}}{{#if:37
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}}{{#if:3
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}}{{#if:318–24
|: 318–24
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|. Bibcode: {{{bibcode}}}
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}}</ref> as well as a secondary condensation product via aldolization.<ref>Template:Cite web</ref>
Ozone has been implicated to have an adverse effect on plant growth, "...Ozone reduced total chlorophylls, carotenoid and carbohydrate concentration, and increased 1-aminocyclopropane-1-carboxylic acid (ACC) content and ethylene production. In treated plants, the ascorbate leaf pool was decreased, while lipid peroxidation and solute leakage were significantly higher than in ozone-free controls. The data indicated that ozone triggered protective mechanisms against oxidative stress in citrus."<ref>{{#if:Iglesias
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|{{#if:Ángeles Calatayuda, Eva Barrenob, Eduardo Primo-Milloa and Manuel Talon
| ; Ángeles Calatayuda, Eva Barrenob, Eduardo Primo-Milloa and Manuel Talon
}}{{#if:February-March 2006
| (February-March 2006)
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|{{#ifeq:Responses of citrus plants to ozone: leaf biochemistry, antioxidant mechanisms and lipid peroxidation||Citation is missing a title.
Either specify one, or click here and a bot will try to complete the citation details for you. }}[{{#if:http://www.sciencedirect.com/science?_ob=ArticleURL&_udi=B6VRD-4JSF4VY-1&_user=10&_rdoc=1&_fmt=&_orig=search&_sort=d&view=c&_acct=C000050221&_version=1&_urlVersion=0&_userid=10&md5=088005a9f3bd83a2de1605ca93f9461a]|http://www.sciencedirect.com/science?_ob=ArticleURL&_udi=B6VRD-4JSF4VY-1&_user=10&_rdoc=1&_fmt=&_orig=search&_sort=d&view=c&_acct=C000050221&_version=1&_urlVersion=0&_userid=10&md5=088005a9f3bd83a2de1605ca93f9461a]|http://www.pubmedcentral.gov/articlerender.fcgi?tool=pmcentrez&artid={{{pmc}}}}} Responses of citrus plants to ozone: leaf biochemistry, antioxidant mechanisms and lipid peroxidation]
|Responses of citrus plants to ozone: leaf biochemistry, antioxidant mechanisms and lipid peroxidation
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}}{{#if:Plant Physiology and Biochemistry
|. Plant Physiology and Biochemistry
}}{{#if:44
| 44
}}{{#if:2-3
| (2-3)
}}{{#if:125–131
|: 125–131
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|. doi:{{#if: | {{#tag:nowiki|10.1016/j.plaphy.2006.03.007}} (inactive [[]]) {{#ifeq: | | [[Category:Pages with DOIs broken since {{#time: Y | }}]] }} | }}
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Production
Ozone often forms in nature under conditions where O2 will not react.<ref name=brown /> Ozone used in industry is measured in g/Nm³ or weight percent. The regime of applied concentrations ranges from 1 to 5 weight percent in air and from 6 to 14 weight percent in oxygen.
Corona discharge method
This is the most popular type of ozone generator for most industrial and personal uses. While variations of the "hot spark" coronal discharge method of ozone production exist, including medical grade and industrial grade ozone generators, these units usually work by means of a corona discharge tube.<ref>Organic Syntheses, Coll. Vol. 3, p.673 (1955); Vol. 26, p.63 (1946). (Article)</ref> They are typically very cost-effective and do not require an oxygen source other than the ambient air. However, they also produce nitrogen oxides as a by-product. Use of an air dryer can reduce or eliminate nitric acid formation by removing water vapor and increase ozone production. Use of an oxygen concentrator can further increase the ozone production and further reduce the risk of nitric acid formation by removing not only the water vapor, but also the bulk of the nitrogen.
Ultraviolet light
UV ozone generators employ a light source that generates a narrow-band ultraviolet light, a subset of that produced by the Sun. The Sun's UV sustains the ozone layer in the stratosphere of Earth.<ref>{{#if:Dohan
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| ; W. J. Masschelein
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|{{#if:
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| (1987)
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|
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|Photochemical Generation of Ozone: Present State-of-the-Art
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|{{#if:
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| (1987)
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}}{{#if:Ozone Sci. Eng.
|. Ozone Sci. Eng.
}}{{#if:9
| 9
}}{{#if:
| ()
}}{{#if:315–334
|: 315–334
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}}</ref> While standard UV ozone generators tend to be less expensive, they usually produce ozone with a concentration of about 0.5% or lower. Another disadvantage of this method is that it requires the air (oxygen) to be exposed to the UV source for a longer amount of time, and any gas that is not exposed to the UV source will not be treated. This makes UV generators impractical for use in situations that deal with rapidly moving air or water streams (in-duct air sterilization, for example). Production of ozone is one of the potential dangers of ultraviolet germicidal irradiation.
Cold plasma
In the cold plasma method, pure oxygen gas is exposed to a plasma created by dielectric barrier discharge. The diatomic oxygen is split into single atoms, which then recombine in triplets to form ozone.
Cold plasma machines utilize pure oxygen as the input source and produce a maximum concentration of about 5% ozone. They produce far greater quantities of ozone in a given space of time compared to ultraviolet production. However, because cold plasma ozone generators are very expensive, they are found less frequently than the previous two types.
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.
Some cold plasma units also have the capability of producing short-lived allotropes of oxygen which include O4, O5, O6, O7, etc. These anions are even more reactive than ordinary O3.
Special considerations
Ozone cannot be stored and transported like other industrial gases (because it quickly decays into diatomic oxygen) 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 20 kg per hour, a gas/water tube heat-exchanger may be utilized as ground electrode and assembled with tubular high-voltage electrodes on the gas-side. The regime of typical gas pressures is around 2 bar absolute in oxygen and 3 bar absolute in air. Several megawatts of electrical power may be installed in large facilities, applied as one phase AC current at 50 to 8000 Hz and peak voltages between 3,000 and 20,000 volts. Applied voltage is usually inversely related to the applied frequency.
The dominating parameter influencing ozone generation efficiency is the gas temperature, which is controlled by cooling water temperature and/or gas velocity. The cooler the water, the better the ozone synthesis. The lower the gas velocity, the higher the concentration (but the lower the net ozone produced). At typical industrial conditions, almost 90% of the effective power is dissipated as heat and needs to be removed by a sufficient cooling water flow.
Because of the high reactivity of ozone, only few materials may be used like stainless steel (quality 316L), titanium, aluminium (as long as no moisture is present), glass, polytetrafluorethylene, or polyvinylidene fluoride. Viton may be used with the restriction of constant mechanical forces and absence of humidity (humidity limitations apply depending on the formulation). Hypalon may be used with the restriction that no water come in contact with it, except for normal atmospheric levels. Embrittlement or shrinkage is the common mode of failure of elastomers with exposure to ozone. Ozone cracking is the common mode of failure of elastomer seals like O-rings.
Silicone rubbers are usually adequate for use as gaskets in ozone concentrations below 1 wt%, such as in equipment for accelerated ageing of rubber samples.
Incidental 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, tasers 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 or hydraulic pumps, will generate more ozone than smaller motors.
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.<ref>{{#if:Ibanez
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|{{#if:Rodrigo Mayen-Mondragon and M. T. Moran-Moran
| ; Rodrigo Mayen-Mondragon and M. T. Moran-Moran
}}{{#if:
| ({{{date}}})
|{{#if:2005
|{{#if:October
| (October 2005)
| (2005)
}}
}}
}}
}}{{#if:Ibanez
| .
}}{{#if:Ibanez2005
|
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| no
|
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|Laboratory Experiments on the Electrochemical Remediation of the Environment. Part 7: Microscale Production of Ozone
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}}{{#if:Journal of Chemical Education
|. Journal of Chemical Education
}}{{#if:82
| 82
}}{{#if:
| ({{{issue}}})
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|: 1546
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|. {{#if:
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|. Bibcode: {{{bibcode}}}
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|. Retrieved on 2006-05-10{{#if: | , [[{{{accessyear}}}]] }}
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- 3 H2O → O3 + 6 H+ + 6 e−; ΔEo = −1.53 V;
- 6 H+ + 6 e− → 3 H2; ΔEo = 0 V;
- 2 H2O → O2 + 4 H+ + 4 e−; ΔEo = −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.
It can also be prepared by passing 10,000-20,000 volts DC through dry O2. This can be done with an apparatus consisting of two concentric glass tubes sealed together at the top, with in and