Ozone

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Ozone
Ozone-CRC-MW-3D-balls.png
Ozone-1,3-dipole.png
Ozone-CRC-MW-3D-vdW.png
Ozone-elpot-3D-vdW.png
IUPAC name Trioxygen
Identifiers
InChI InChI=1/O3/c1-3-2
InChIKey CBENFWSGALASAD-UHFFFAOYAY
Standard InChI InChI=1S/O3/c1-3-2
Standard InChIKey CBENFWSGALASAD-UHFFFAOYSA-N
CAS number [10028-15-6]
EC number 233-069-2
RTECS RS8225000
ChemSpider 23208
PubChem 24823
Properties
Chemical formula O3
Molar mass 47.998 g·mol−1
Appearance bluish colored gas
Density 2.144 g/L (0 °C), gas
Melting point

80.7 K, −192.5 °C

Boiling point

161.3 K, −111.9 °C

Solubility in water 0.105 g/100mL (0 °C)
Refractive index (nD) 1.2226 (liquid)
Thermochemistry
Std enthalpy of formation ΔfHo298 +142.3 kJ·mol−1
Standard molar entropy So298 237.7 J·K−1.mol−1
Hazards
EU classification Oxidant (O)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)

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. The ozone layer 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 verb ozein (ὄζειν, "to smell"), from the peculiar odor in lightning storms.[1][2] The formula for ozone, O3, was not determined until 1865 by Jacques-Louis Soret[3] and confirmed by Schönbein in 1867.[1][4]

Physical properties

Ozone is a pale blue gas, 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 -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 -193 °C, it forms a violet-black solid.[5]

Most people can detect about 0.01 ppm of ozone in air where it has a very specific sharp odor somewhat resembling chlorine bleach. Exposure of 0.1 to 1 ppm produces headaches, burning eyes, and irritation to the respiratory passages.[6] Even low concentrations of ozone in air are very destructive to organic materials such as latex, plastics, and lungs.

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°.[7] The central atom forms an sp² hybridization with one lone pair. Ozone is a polar molecule with a dipole moment of 0.5337 D.[8] 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

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[9]):

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.[10]

Metals

Ozone will oxidize metals (except gold, platinum, and iridium) to oxides of the metals in their highest oxidation state:

2 Cu+ (aq) + 2 H3O+ (aq) + O3 (g) → 2 Cu2+ (aq) + 3 H2O (l) + O2 (g)

Non-metals

Ozone also increases the oxidation number of oxides, such as the oxidation of 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:

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. For example, lead(II) sulfide is oxidised to lead(II) sulfate:

PbS + 4 O3 → PbSO4 + 4 O2

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 of tin(II) chloride with hydrochloric acid and NaCl:

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

Combustion

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 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:[11]

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:[12]

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:[13]

CsO3 + Na+ → Cs+ + NaO3

Treatment with ozone of calcium dissolved in ammonia leads to ammonium ozonide and not calcium ozonide:[11]

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 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:[14]

(NH2)2CO + O3 → N2 + CO2 + 2 H2O

Ozone in Earth's atmosphere

The distribution of atmospheric ozone in partial pressure as a function of altitude.
Concentration of ozone as measured by the Nimbus-7 satellite.
Total ozone concentration in June 2000 as measured by EP-TOMS satellite instrument.

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

Main article: 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.

See also: Ozone-oxygen cycle and Ozone depletion

Low level ozone

Main articles: Tropospheric ozone and Photochemical smog

Low level ozone (or tropospheric ozone) is regarded as a pollutant by the World Health Organization[15] and the United States Environmental Protection Agency (EPA). It is not emitted directly by car engines or by industrial operations, but 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).[16]

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.[17][18]

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.[18]

Ozone cracking

Ozone cracking in Natural rubber tubing

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. Storing rubber products in close proximity to DC electric motors can accelerate the rate at which ozone cracking occurs. The commutator of the motor creates sparks which in turn produce ozone.

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.[19][20] 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[21]) suggests that the radiative forcing of tropospheric ozone is about 25% that of carbon dioxide.

Health effects

Air pollution

Red Alder leaf, showing the typical discolouration caused by ozone pollution.[22]

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.[15][23] A connection has also been known to exist between increased ozone caused by thunderstorms and hospital admissions of asthma sufferers.[24] Air quality guidelines such as those from the World Health Organization are based on detailed studies of what levels can cause measurable health effects. Exposure to ozone and the pollutants that produce it has been linked to premature death, asthma, bronchitis, heart attack, and other cardiopulmonary problems. According to scientists with the United States Environmental Protection Agency (EPA), susceptible people can be adversely effected by ozone levels as low as 40 ppb.[25]

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 standards 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 lowering 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.[25] 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 ppb are described as "Unhealthy for Sensitive Groups", 105 ppb to 124 ppb as "unhealthy" and 125 ppb to 404 ppb as "very unhealthy".[26]

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.[27] 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.[28]

Long-term exposure to ozone has been shown to increase risk of death from respiratory illness. A study of 450,000 people living in United States cities showed a significant correlation between ozone levels and respiratory illness over the 18-year follow-up period. The study revealed that people living in cities with high ozone levels such as Houston or Los Angeles had an over 30% increased risk of dying from lung disease.[29][30]

Physiology

See also: trioxidane

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.[31]

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[32] as well as a secondary condensation product via aldolization.[33]

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."[34]

Safety regulations

Due to the strongly oxidizing properties of ozone, ozone is a primary irritant, affecting especially the eyes and respiratory systems and can be hazardous at even low concentrations. The Canadian Center for Occupation Safety and Health reports that:

"Even very low concentrations of ozone can be harmful to the upper respiratory tract and the lungs. The severity of injury depends on both by the concentration of ozone and the duration of exposure. Severe and permanent lung injury or death could result from even a very short-term exposure to relatively low concentrations." [35]

To protect workers potentially exposed to ozone, OSHA has established a permissible exposure limit (PEL) of 0.1 ppm (29 CFR 1910.1000 table Z-1), calculated as an 8 hour time weighted average. Higher concentrations are especially hazardous and NIOSH has established an Immediately Dangerous to Life and Health Limit (IDLH) of 5 ppm. [36] Work environments where ozone is used or where it is likely to be produced should have adequate ventilation and it is prudent to have a monitor for ozone that will alarm if the concentration exceeds the OSHA PEL. Continuous monitors for ozone are available from several suppliers.

Production

Ozone often forms in nature under conditions where O2 will not react.[6] 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.[37] 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.[38] 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.VUV Ozone generators are used in swimming pool and spa applications ranging to millions of gallons of water. VUV Ozone generators, unlike Corona Discharge generators) do not produce harmful nitrogen by-products and also unlike Corona Discharge systems, VUV Ozone generators work extremely well in humid air environments. There is also not normally a need for expensive off-gas mechanisms, and no need for air driers or oxygen concentrators which require extra costs and maintenance.

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. Ozone is similarly formed in the Catatumbo lightning storms phenomenon on the Catatumbo River in Venezuela, which helps to replenish ozone in the upper troposhere. It is the world's largest single natural generator of ozone, lending calls for it to be designated a UNESCO World Heritage Site.[39]

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 3 molar sulfuric acid electrolyte.[40] The half cell reactions taking place are:

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)

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 out spigots at the top and bottom of the outer tube. The inner core should have a length of metal foil inserted into it connected to one side of the power source. The other side of the power source should be connected to another piece of foil wrapped around the outer tube. Dry O2 should be run through the tube in one spigot. As the O2 is run through one spigot into the apparatus and 10,000-20,000 volts DC are applied to the foil leads, electricity will discharge between the dry dioxygen in the middle and form O3 and O2 out the other spigot. The reaction can be summarized as follows:[6]

3 O2 — electricity → 2 O3

Ionic air purifiers

Some air purifiers create ozone.[41]

Applications

Industry

The largest use of ozone is in the preparation of pharmaceuticals, synthetic lubricants, as well as many other commercially useful organic compounds, where it is used to sever carbon-carbon bonds.[6] It can also be used for bleaching substances and for killing microorganisms in air and water sources.[42] Many municipal drinking water systems kill bacteria with ozone instead of the more common chlorine.[43] Ozone has a very high oxidation potential.[44] Ozone does not form organochlorine compounds, nor does it remain in the water after treatment. The Safe Drinking Water Act mandate that these systems introduce an amount of chlorine to maintain a minimum of 0.2 ppm residual Free Chlorine in the pipes, based on results of regular testing. Where electrical power is abundant, ozone is a cost-effective method of treating water, since 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.

Although low levels of ozone have been advertised to be of some disinfectant use in residential homes, the concentration of ozone in dry air required to have a rapid, substantial effect on airborne pathogens exceeds safe levels recommended by the U.S. Occupational Safety and Health Administration and Environmental Protection Agency. Humidity control can vastly improve both the killing power of the ozone and the rate at which it decays back to oxygen (more humidity allows more effectiveness). Spore forms of most pathogens are very tolerant of atmospheric ozone in concentrations where asthma patients start to have issues.

Industrially, ozone is used to:

  • Disinfect laundry in hospitals, food factories, care homes etc;[45]
  • Disinfect water in place of chlorine[6]
  • Deodorize air and objects, such as after a fire. This process is extensively used in Fabric Restoration
  • Kill bacteria on food or on contact surfaces;[46]
  • Sanitize swimming pools and spas
  • Kill insects in stored grain[47]
  • Scrub yeast and mold spores from the air in food processing plants;
  • Wash fresh fruits and vegetables to kill yeast, mold and bacteria;[46]
  • Chemically attack contaminants in water (iron, arsenic, hydrogen sulfide, nitrites, and complex organics lumped together as "colour");
  • 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[48]
  • Clean and bleach fabrics (the former use is utilized in Fabric Restoration; 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;
  • Eradicate water borne parasites such as Giardia lamblia and Cryptosporidium in surface water treatment plants.

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.[49]

Ozone is used as an alternative to chlorine or chlorine dioxide in the bleaching of wood pulp[50] . It is often used in conjunction with oxygen and hydrogen peroxide to eliminate the need for chlorine-containing compounds in the manufacture of high-quality, white paper[51]

Ozone can be used to detoxify cyanide wastes (for example from gold and silver mining) by oxidizing cyanide to cyanate and eventually to carbon dioxide.[52]

Consumers

Devices generating high levels of ozone, some of which use ionization, are used to sanitize and deodorize uninhabited buildings, rooms, ductwork, woodsheds, and boats and other vehicles.

In the U.S., air purifiers emitting lower levels of ozone have been sold. This kind of air purifier is sometimes claimed to imitate nature's way of purifying the air[53] without filters and to sanitize both it and household surfaces. The United States Environmental Protection Agency (EPA) has declared that there is "evidence to show that at concentrations that do not exceed public health standards, ozone is not effective at removing many odor-causing chemicals" or "viruses, bacteria, mold, or other biological pollutants." Furthermore, its report states that "results of some controlled studies show that concentrations of ozone considerably higher than these [human safety] standards are possible even when a user follows the manufacturer’s operating instructions."[54] The government successfully sued one company in 1995, ordering it to stop repeating health claims without supporting scientific studies.

Ozonated water is used to launder clothes and to sanitize food, drinking water, and surfaces in the home. According to the U.S. Food and Drug Administration (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 demonstrated that 0.3 ppm levels of ozone dissolved in filtered tapwater can produce a reduction of more than 99.99% in such food-borne microorganisms as salmonella, E. coli 0157:H7, and Campylobacter. This quantity exceeds 20,000 times the WHO recommended limits stated above.[46][55] Ozone can be used to remove pesticide residues from fruits and vegetables.[56][57]

Ozone is used in homes and hot tubs to kill bacteria in the water and to reduce the amount of chlorine or bromine required by reactivating them to their free state. Since ozone does not remain in the water long enough, ozone by itself is ineffective at preventing cross-contamination among bathers and must be used in conjunction with these halogens. Gaseous ozone created by ultraviolet light or by corona discharge is injected into the water.[58]

Ozone is also widely used in treatment of water in aquariums and fish ponds. Its use can minimize bacterial growth, control parasites, eliminate transmission of some diseases, and reduce or eliminate "yellowing" of the water. Ozone must not come in contact with fish's gill structures. Natural salt water (with life forms) provides enough "instantaneous demand" that controlled amounts of ozone activate bromide ion to hypobromous acid, and the ozone entirely decays in a few seconds to minutes. If oxygen fed ozone is used, the water will be higher in dissolved oxygen, fish's gill structures will atrophy and they will become dependent on higher dissolved oxygen levels.

See also

References

  1. ↑ 1.0 1.1 Rubin, Mordecai B. (2001). "The History of Ozone. The Schönbein Period, 1839-1868" (PDF). Bull. Hist. Chem. 26 (1). Retrieved on 2008-02-28. 
  2. ↑ Today in Science History. Retrieved on 2006-05-10.
  3. ↑ Jacques-Louis Soret (1865). "Recherches sur la densité de l'ozone". Comptes rendus de l'Académie des sciences 61: 941. 
  4. ↑ Ozone FAQ. Global Change Master Directory. Retrieved on 2006-05-10.
  5. ↑ Oxygen. WebElements. Retrieved on 2006-09-23.
  6. ↑ 6.0 6.1 6.2 6.3 6.4 [1977] (2003) "22", in Nicole Folchetti: Chemistry: The Central Science, 9th, Pearson Education, 882–883. ISBN 0-13-066997-0.  Cite error: Invalid <ref> tag; name "brown" defined multiple times with different content
  7. ↑ 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.
  8. ↑ 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
  9. ↑ 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.
  10. ↑ http://www.iitk.ac.in/che/jpg/papersb/full%20papers/K-106.pdf
  11. ↑ 11.0 11.1 Horvath M., Bilitzky L., & Huttner J., 1985. "Ozone." pg 44–49
  12. ↑ Housecroft & Sharpe, 2005. "Inorganic Chemistry." pg 439
  13. ↑ Housecroft & Sharpe, 2005. "Inorganic Chemistry." pg 265
  14. ↑ Horvath M., Bilitzky L., & Huttner J., 1985. "Ozone." pg 259, 269–270
  15. ↑ 15.0 15.1 WHO-Europe reports: Health Aspects of Air Pollution (2003) (PDF)
  16. ↑ Stevenson et al. (2006). Multimodel ensemble simulations of present-day and near-future tropospheric ozone. American Geophysical Union. Retrieved on 2006-09-16.
  17. ↑ Rising Ozone Levels Pose Challenge to U.S. Soybean Production, Scientists Say. NASA Earth Observatory (2003-07-31). Retrieved on 2006-05-10.
  18. ↑ 18.0 18.1 Mutters, Randall (March 1999). Statewide Potential Crop Yield Losses From Ozone Exposure. California Air Resources Board. Retrieved on 2006-05-10.
  19. ↑ Tropospheric Ozone in EU - The consolidated report. European Environmental Agency (1998). Retrieved on 2006-05-10.
  20. ↑ Atmospheric Chemistry and Greenhouse Gases. Intergovernmental Panel on Climate Change. Retrieved on 2006-05-10.
  21. ↑ Climate Change 2001. Intergovernmental Panel on Climate Change (2001). Retrieved on 2006-09-12.
  22. ↑ Jeannie Allen (2003-08-22). Watching Our Ozone Weather. NASA Earth Observatory. Retrieved on 2008-10-11.
  23. ↑ Answer to follow-up questions from CAFE (2004) (PDF)
  24. ↑ Anderson, W.; G.J. Prescott, S. Packham, J. Mullins, M. Brookes, and A. Seaton (August 2001). "Asthma admissions and thunderstorms: a study of pollen, fungal spores, rainfall, and ozone". QJM: an International Journal of Medicine 94 (8): 429–433. Oxford Journals. doi:10.1093/qjmed/94.8.429. PMID 11493720. 
  25. ↑ 25.0 25.1 Weinhold B (July 2008). "Ozone nation: EPA standard panned by the people". Environ. Health Perspect. 116 (7): A302–A305. PMID 18629332. 
  26. ↑ Smog - Who does it hurt? What You Need to Know About Ozone and Your Health. AIRNow.gov. Retrieved on 2007-07-10.
  27. ↑ Ashfield District Council: Monitored Air Pollutants, downloaded February 2, 2007
  28. ↑ University of East Anglia press release, Cloning the smell of the seaside, February 2, 2007
  29. ↑ Jerrett, Michael; Burnett, Richard T. and Pope, C. Arden, III and Ito, Kazuhiko and Thurston, George and Krewski, Daniel and Shi, Yuanli and Calle, Eugenia and Thun, Michael (March 12, 2009). "Long-Term Ozone Exposure and Mortality". N. Engl. J. Med. 360 (11): 1085–1095. doi:10.1056/NEJMoa0803894. PMID 19279340. 
  30. ↑ Wilson, Elizabeth K. (March 16, 2009). "Ozone's Health Impact". Chemical & Engineering News 87 (11): 9. American Chemical Society Publications. 
  31. ↑ Hoffmann, Roald (January 2004). "The Story of O". American Scientist 92 (1): 23. doi:10.1511/2004.1.23. Retrieved on 2006-10-11. 
  32. ↑ Smith, LL (1 August 2004). "Oxygen, oxysterols, ouabain, and ozone: a cautionary tale". Free radical biology & medicine 37 (3): 318–24. Elsevier. doi:10.1016/j.freeradbiomed.2004.04.024. Retrieved on 2008-05-17. 
  33. ↑ Paul Wentworth (November 2003). Evidence for Ozone Formation in Human Atherosclerotic Arteries. Retrieved on 2006-08-03.
  34. ↑ Iglesias, Domingo J.; Ángeles Calatayuda, Eva Barrenob, Eduardo Primo-Milloa and Manuel Talon (February-March 2006). "Responses of citrus plants to ozone: leaf biochemistry, antioxidant mechanisms and lipid peroxidation". Plant Physiology and Biochemistry 44 (2-3): 125–131. Elsevier. doi:10.1016/j.plaphy.2006.03.007. Retrieved on 2008-05-17. 
  35. ↑ [1]
  36. ↑ http://www.cdc.gov/niosh/idlh/intridl4.html
  37. ↑ Organic Syntheses, Coll. Vol. 3, p.673 (1955); Vol. 26, p.63 (1946). (Article)
  38. ↑ Dohan, J. M.; W. J. Masschelein (1987). "Photochemical Generation of Ozone: Present State-of-the-Art". Ozone Sci. Eng. 9: 315–334. 
  39. ↑ Fire in the Sky. Retrieved on 2008-08-16.
  40. ↑ Ibanez, Jorge G.; Rodrigo Mayen-Mondragon and M. T. Moran-Moran (October 2005). "Laboratory Experiments on the Electrochemical Remediation of the Environment. Part 7: Microscale Production of Ozone". Journal of Chemical Education 82: 1546. Retrieved on 2006-05-10. 
  41. ↑ Phillips, TJ; Bloudoff DP, Jenkins PL, Stroud KR. (1999 Nov-Dec). "Ozone emissions from a "personal air purifier".". J Expo Anal Environ Epidemiol.: 594–601. 6: 9. Retrieved on 29 May 2009. 
  42. ↑ Ozone and Color Removal. Ozone Information. Retrieved on 2009-01-09.
  43. ↑ Hoigné, J. (1998). Handbook of Environmental Chemistry, Vol. 5 part C, p83-141. Berlin: Springer-Verlag. 
  44. ↑ Oxidation Potential of Ozone. Ozone-Information.com. Retrieved on 2008-05-17.
  45. ↑ Decontamination: Ozone scores on spores. Hospital Development. Wilmington Media Ltd. (2007-04-01). Retrieved on 2007-05-30.
  46. ↑ 46.0 46.1 46.2 Montecalvo, Joseph; Doug Williams. Application of Ozonation in Sanitizing Vegetable Process Washwaters (PDF). California Polytechnic State University. Retrieved on 2008-03-24.
  47. ↑ http://news.uns.purdue.edu/UNS/html4ever/030130.Mason.ozone.html
  48. ↑ Chemical Synthesis with Ozone. Ozone-Information.com. Retrieved on 2008-05-17.
  49. ↑ de Boer, Hero E. L.; Carla M. van Elzelingen-Dekker, BSc; Cora M. F. van Rheenen-Verberg, BSc; Lodewijk Spanjaard, MD, PhD (October 2006). "Use of Gaseous Ozone for Eradication of Methicillin-Resistant Staphylococcus aureus From the Home Environment of a Colonized Hospital Employee". Infection Control and Hospital Epidemiology 27 (10): 1120–1122. Chicago: The University of Chicago Press. doi:10.1086/507966. Retrieved on 2008-05-17. 
  50. ↑ Sjöström, Eero (1993). Wood Chemistry: Fundamentals and Applications. San Diego, CA: Academic Press, Inc.. ISBN 0-12-64748. 
  51. ↑ Su, Yu-Chang (2001). "Enzone Bleaching Sequence and Color Reversion of Ozone-Bleached Pulps". Taiwan Journal of Forest Science 16 (2): 93–102. 
  52. ↑ Bollyky, L. J. (1977). Ozone Treatment of Cyanide-Bearing Wastes, EPA Report 600/2-77-104. Research Triangle Park, N.C.: U.S. Environmental Protection Agency. 
  53. ↑ The Unknown Truth Regarding Ozone!. Retrieved on 2006-09-16.
  54. ↑ EPA report on consumer ozone air purifiers
  55. ↑ Long, Ron (2008). POU Ozone Food Sanitation: A Viable Option for Consumers & the Food Service Industry (pdf). (report also shows tapwater removes 99.95% of pathogens from lettuce; samples were first inoculated with pathogens before treatment)
  56. ↑ Tersano Inc (2007). lotus Sanitises Food without Chemicals. Retrieved on 2007-02-11.
  57. ↑ Jongen, W (2005). Improving the Safety of Fresh Fruit and Vegetables. Boca Raton: Woodhead Publishing Ltd. ISBN 1855739569. 
  58. ↑ Alternative Disinfectants and Oxidant Guidance Manual (PDF). United States Environmental Protection Agency. Retrieved on 2008-01-14.

Further reading

  • Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements, 2nd ed.; Butterworth-Heinemann: Oxford, 1997. ISBN 0-7506-3365-4
  • 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

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