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For other uses, see Oxygen (disambiguation).

8 nitrogen ← oxygen → fluorine - ↑ O ↓ S periodic table
General
Name, Symbol, Number	oxygen, O, 8
Chemical series	Chalcogens
Group, Period, Block	16, 2, p
Appearance	colorless
Atomic mass	15.9994(3) g/mol
Electron configuration	1s2 2s2 2p4
Electrons per shell	2, 6
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 1.429 g/L
Melting point	54.36 K (-218.79 °C, -361.82 °F)
Boiling point	90.20 K (-182.95 °C, -297.31 °F)
Heat of fusion	(O2) 0.444 kJ/mol
Heat of vaporization	(O2) 6.82 kJ/mol
Heat capacity	(25 °C) (O2) 29.378 J/(mol·K)
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K       61 73 90
Atomic properties
Crystal structure	cubic
Oxidation states	−2, −1 (neutral oxide)
Electronegativity	3.44 (Pauling scale)
Ionization energies (more)	1st: 1313.9 kJ/mol
	2nd: 3388.3 kJ/mol
	3rd: 5300.5 kJ/mol
Atomic radius	60 pm
Atomic radius (calc.)	48 pm
Covalent radius	73 pm
Van der Waals radius	152 pm
Miscellaneous
Magnetic ordering	paramagnetic
Thermal conductivity	(300 K) 26.58 mW/(m·K)
Speed of sound	(gas, 27 °C) 330 m/s
CAS registry number	7782-44-7
Notable isotopes
Main article: Isotopes of oxygen iso NA half-life DM DE (MeV) DP 16O 99.762% O is stable with 8 neutrons 17O 0.038% O is stable with 9 neutrons 18O 0.2% O is stable with 10 neutrons
References

Oxygen is a chemical element in the periodic table. It has the symbol O and atomic number 8. The element is very common, found not only on Earth but throughout the universe, usually covalently bonded with other elements. Unbound oxygen (usually called molecular oxygen, O₂, a diatomic molecule) first appeared on Earth during the Paleoproterozoic era (between 2500 million years ago and 1600 million years ago) and as a product of the metabolic action of early anaerobes (archaea and bacteria). The presence of free oxygen drove most of the organisms then living to extinction. The atmospheric abundance of free oxygen in later geological epochs and up to the present has been largely driven by photosynthetic organisms, roughly three quarters by phytoplankton and algae in the oceans and one quarter from terrestrial plants.

Contents 1 Characteristics 2 Applications 3 History 4 Occurrence 5 Compounds 6 Isotopes 7 Precautions 8 See also 9 References 10 External links

Characteristics

At standard temperature and pressure, oxygen exists as a diatomic molecule with the formula O₂, in which the two oxygen atoms are doubly bonded to each other. In its most stable form, oxygen exists as a diradical (triplet oxygen). Though radicals are commonly associated with highly reactive compounds, triplet oxygen is surprisingly (and fortunately) unreactive towards most compounds. Singlet oxygen, a name given to several higher energy species in which all the electron spins are paired, is much more reactive towards common organic molecules.

Oxygen is a major component of air, produced by plants during photosynthesis, and is necessary for aerobic respiration in animals. The word oxygen derives from two words in Greek, οξυς (oxys) (acid, sharp) and γεινομαι (geinomai) (engender). The name "oxygen" was chosen because, at the time it was discovered in the late 18th century, it was believed that all acids contained oxygen. The definition of acid has since been revised to not require oxygen in the molecular structure.

Liquid O₂ and solid O₂ have a light blue color and both are highly paramagnetic. Liquid O₂ is usually obtained by the fractional distillation of liquid air.

Liquid and solid O₃ (ozone) have a deeper color of blue.

A recently discovered allotrope of oxygen, tetraoxygen (O₄), is a deep red solid that is created by pressurizing O₂ to the order of 20 GPa. Its properties are being studied for use in rocket fuels and similar applications, as it is a much more powerful oxidizer than either O₂ or O₃.

Applications

Liquid oxygen finds use as an oxidizer in rocket propulsion. Oxygen is essential to respiration, so oxygen supplementation has found use in medicine (as oxygen therapy). People who climb mountains or fly in airplanes sometimes have supplemental oxygen supplies (as air). Oxygen is used in welding (such as the oxyacetylene torch), and in the making of steel and methanol.

Oxygen presents two absorption bands centered in the wavelengths 687 and 760 nanometers. Some scientists have proposed to use the measurement of the radiance coming from vegetation canopies in those oxygen bands to characterize plant health status from a satellite platform. This is because in those bands, it is possible to discriminate the vegetation's reflectance from the vegetation's fluorescence, which is much weaker. The measurement presents several technical difficulties due to the low signal to noise ratio and due to the vegetation's architecture, but it has been proposed as possibility to monitor the carbon cycle from satellite, thus in a global scale.

Oxygen, as a mild euphoric, has a history of recreational use that extends into modern times. Oxygen bars can be seen at parties to this day. In the 19th century, oxygen was often mixed with nitrous oxide to promote an analgesic effect; indeed, such a mixture (Entonox) is commonly used in medicine today.

History

Oxygen was first discovered by Michał Sędziwój, Polish alchemist and philosopher in late 16th century. Sędziwój assumed the existence of oxygen by warming nitre (saltpeter). He thought of the gas given off as "the elixir of life".

Oxygen was again discovered by the Swedish pharmacist Carl Wilhelm Scheele sometime before 1773, but the discovery was not published until after the independent discovery by Joseph Priestley on August 1, 1774, who called the gas dephlogisticated air (see phlogiston theory). Priestley published his discoveries in 1775 and Scheele in 1777; consequently Priestley is usually given the credit. It was named by Antoine Laurent Lavoisier after Priestley's publication in 1775.

Occurrence

Oxygen is the second most common component of the earth's atmosphere (20.947% by volume).

Compounds

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements (which is the origin of the original definition of oxidation). The only elements to escape the possibility of oxidation are a few of the noble gases. The most famous of these oxides is dihydrogen monoxide, or water (H₂O). Other well known examples include compounds of carbon and oxygen, such as carbon dioxide (CO₂), alcohols (R-OH), aldehydes, (R-CHO), and carboxylic acids (R-COOH). Oxygenated radicals such as chlorates (ClO₃⁻), perchlorates (ClO₄⁻), chromates (CrO₄²⁻), dichromates (Cr₂O₇²⁻), permanganates (MnO₄⁻), and nitrates (NO₃⁻) are strong oxidizing agents in and of themselves. Many metals such as iron bond with oxygen atoms, iron (III) oxide (Fe₂O₃). Ozone (O₃) is formed by electrostatic discharge in the presence of molecular oxygen. A double oxygen molecule (O₂)₂ is known and is found as a minor component of liquid oxygen. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.

Isotopes

Oxygen has fifteen known isotopes with atomic masses ranging from 12 to 26. Three of them are stable and twelve are radioactive. The radioisotopes all have half lives of less than three minutes. The stable isotopes have mass numbers of 16, 17 and 18, of which oxygen-16 is the most common (over 99%).

Precautions

Oxygen can be toxic at elevated partial pressures (i.e. high relative concentrations). This is important in some forms of scuba diving, such as with a rebreather.

Certain derivatives of oxygen, such as ozone (O₃), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic. The body has developed mechanisms to protect against these toxic species. For instance, the naturally-occurring glutathione can act as an antioxidant, as can bilirubin which is normally a breakdown product of hemoglobin. To protect against the destructive nature of peroxides, nearly every organism on earth has developed some form of the enzyme catalase, which very quickly dispropoitionates peroxide into water and dioxygen. Highly concentrated sources of oxygen promote rapid combustion and therefore are fire and explosion hazards in the presence of fuels. This is true as well of compounds of oxygen such as chlorates, perchlorates, dichromates, etc. Compounds with a high oxidative potential can often cause chemical burns.

The fire that killed the Apollo 1 crew on a test launchpad spread so rapidly because the pure oxygen atmosphere was at normal atmospheric pressure instead of the one third pressure that would be used during an actual launch. (See partial pressure.)

Oxygen derivatives are prone to form free radicals, especially in metabolic processes. Because they can cause severe damage to cells and their DNA, they are thought to be related to cancer and aging.

See also

• Winkler test for dissolved oxygen for instructions on how to determine the amount of oxygen dissolved in fresh water.
• Combustion
• Oxidation
• Oxygen Catastrophe in geology
• The role of oxygen as a diving breathing gas
• Oxygen depletion aquatic ecology
• Ozone layer

References

• Los Alamos National Laboratory – Oxygen
• Nist atomic spectra database
• Nuclides and Isotopes Fourteenth Edition: Chart of the Nuclides, General Electric Company, 1989

External links

• Priestley Society, Dedicated to Joseph Priestley the man who discovered oxygen
• Joseph Priestley Information Website, about the man who discovered oxygen
• Los Alamos National Laboratory – Oxygen
• WebElements.com – Oxygen
• It's Elemental – Oxygen
• Oxygen Toxicity
• Oxygen (O2) Properties, Uses, Applications

• Computational Chemistry Wiki
• Tests with liquid oxygen :-)

This article is based on the article "Oxygen" from Wikipedia - the free encyclopedia created and edited by online user community. This article is distributed under the terms of GNU Free Documentation License. Here you find the list of authors of this article. The article can only edited within Wikipedia. Edit this article in Wikipedia.

Ozone
O=O-O
General
Systematic name	Trioxygen
Molecular formula	O3
Molar mass	47.998 g/mol
Appearance	colorless gas
CAS number	[10028-15-6]
Properties
Density and phase	2.144 g/l (0 °C), gas
Solubility in water	0.105 g/100 ml (0 °C)
Melting point	−197.2 °C
Boiling point	−111.9 °C
Thermodynamic data
Standard enthalpy of formation ΔfH°solid	+142.3 kJ/mol
Standard molar entropy S°solid	237.7 J.K−1.mol−1
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

For the Moldavian pop group see O-Zone

Ozone (O₃) is an allotrope of oxygen, the molecule consisting of three oxygen atoms instead of the more stable diatomic O₂.

Ozone is a colorless gas at standard temperature and pressure. It forms a dark blue liquid below -112 °C and a dark blue solid below -193 °C. Ozone is a powerful oxidizing agent. It is also unstable, decaying to ordinary oxygen through the reaction:

2O₃ → 3O₂. This reaction proceeds more rapidly with increasing temperature and decreasing pressure.

Ozone is a highly corrosive, poisonous substance and a common pollutant. It has a sharp, pungent odor. It is present in low concentrations throughout the Earth's atmosphere. It is also formed from O₂ by electrical discharges such as lightning, and by action of high energy electromagnetic radiation.

Some kinds of electrical equipment generate levels of ozone that a human can easily smell. This is especially true of devices using high voltages, such as television sets, laser printers, and photocopiers. Electric motors using brushes can generate ozone from repeated sparking inside the unit. Large motors, such as those used by elevators or hydraulic pumps, will generate more ozone than smaller motors.

Contents 1 Ozone layer 2 Discovery of ozone 3 Industrial production 4 Use in industry 5 Use in medicine 6 Air pollution 7 Other uses 8 See also 9 External links 10 References

Ozone layer

See 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. 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 responsible for the production of vitamin D, which is essential for human health. The standard way to express total ozone amounts in the atmosphere is by using Dobson units. Ozone used in industry is measured in ppm (OSHA exposure limits for example), and percent by mass or weight.

Discovery of ozone

Ozone was discovered by Christian Friedrich Schönbein in 1840, who named it after the Greek word for smell (ozein), from the peculiar odor in lightning storms. [1]. Even so, the odor from a lightning strike is usually from electrons freed during the rapid chemical changes, not the ozone itself. [2].

Industrial production

Industrially, ozone is produced with short wavelength ultraviolet radiation from a mercury vapor lamp or the application of a high voltage electrical field in a process called cold discharge. The cold discharge apparatus consists of two metal plates separated by an air gap and a high dielectric strength electrical insulator such as borosilicate glass or mica. A high voltage alternating current is applied to the plates and the ozone is formed in the air gap when O₂ molecules disassociate and recombine into O₃. A faint corona may be present in the air gap, but the voltage is maintained below that which would cause punch-through of the insulator with subsequent arcing and plasma formation. 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 ^[3]. The half cell reactions taking place are:

3H₂O → O₃ + 6H⁺ + 6e^- ΔE^o = - 1.53 V

6H⁺ + 6e^- → 3H₂ ΔE^o = 0 V

2H₂O → O₂ + 4H⁺ + 4e^- ΔE^o = -1. 23 V

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

Use in industry

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 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.

Industrially, ozone or ozonated water is used to:

• disinfect water before it is bottled,
• kill bacteria on food-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 (a process of agglomeration of molecules, which aids in filtration... this is where the iron and arsenic are removed),
• 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.

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

Use in medicine

Ozone, along with hypochlorite ions, is naturally produced by white blood cells and the roots of marigolds as a means of destroying foreign bodies. When ozone breaks down it gives rise to oxygen free radicals, which are highly reactive and damage or destroy most organic molecules.

Ozone has a number of medical uses. It can be used to affect the body's antioxidant-prooxidant balance, since the body usually reacts to its presence by producing antioxidant enzymes. 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 bacterium.

Ozone therapy has blossomed into a thriving field of alternative medicine, and there are a host of claimed applications above and beyond what has actually been verified by studies. Ozone treatments are dangerous, however, since ozone is highly corrosive.

In the United States ozone therapy is illegal, as the Food and Drug Administration (FDA) has not approved its use on humans. At least one death has been attributed to application of ozone through insufflation in the U.S. "Air cleaners" which produce "activated oxygen", i.e., ozone, are often sold in the U.S. nonetheless. See Air ioniser.

Air pollution

See main articles: Tropospheric ozone and Air pollution.

Ozone is not directly emitted by car engines or by industial operations themselves. These sources emit hydrocarbons and nitrogen oxides that react with sunlight to form ozone directly at the source of the pollution being emmitted and in the atmosphere's boundary layer (1 to 3 km altitude). The mix of hydrocarbons, nitrogen oxides, and ozone are the major components of smog that frequently occurs in urban and suburban areas. Recent satellite maps of Nitrogen Dioxide (NO₂) clearly show the worldwide distribution of polluted regions associated with industrial activity (automobiles, factories, and fossil fuel power generation).

There is a great deal of evidence to show that ozone at the earth's surface can harm lung function and irritate the respiriatory system (WHO Europe reports, cited below). Ozone has been found to convert cholesterol in the blood stream to plaque (which causes hardening and narrowing of arteries). This cholesterol product has also been implicated in Alzheimer's disease, suggesting a link between the inflammatory response associated with head injury and Alzheimer's. Air quality guidlines such as those from the World Health Organization are based on detailed studies of what levels can cause measurable health effects.

Although ozone was present at ground level before the industrial revolution, peak concentrations are far higher than the pre-industrial levels [4] and even background concentrations well away from sources of pollution are substantially higher [5].

Ozone reacts directly with some hydrocarbons such as aldehydes and thus begins their removal from the air but the products of ozonolysis are themselves key components of smog. Ozone photolysis by UV light leads to production of the hydroxyl radical and this plays a key part in the removal of hydrocarbons from the air but is again a key step in the creation of components of smog such as Peroxyacyl nitrates which are powerful eye irritants. Ultimately, ozone is one component of smog which is harmful in itself and contributes both to the production and ultimate removal of other air pollutants.

Other uses

During the 1992 U.S. Presidential election, George H.W. Bush referred to his opponents Bill Clinton and Al Gore as "Bozo and Ozone", respectively, the latter in connection with Gore's well known stance on environmental issues.

See also

• Ozone depletion, including the phenomenon known as the Ozone Hole.
• Ozone layer
• Tropospheric ozone

External links

• 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
• WHO-Europe reports: Health Aspects of Air Pollution (2003) (PDF) and "Answer to follow-up questions from CAFE (2004) (PDF)
• Ground-level Ozone Air Pollution — A summary for non specialists by GreenFacts of the above WHO reports.

References

• ^  Laboratory Experiments on the Electrochemical Remediation of the Environment. Part 7: Microscale Production of Ozone Jorge G. Ibanez, Rodrigo Mayen-Mondragon, and M. T. Moran-Moran J. Chem. Ed. October 2005 Vol. 82 No. 10 p. 1546 Abstract
• 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

This article is based on the article "Ozone" from Wikipedia - the free encyclopedia created and edited by online user community. This article is distributed under the terms of GNU Free Documentation License. Here you find the list of authors of this article. The article can only edited within Wikipedia. Edit this article in Wikipedia.