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{{Infobox oxygen}}
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==Overview==
'''Oxygen ''' is the [[chemical element|element]] with [[atomic number]] 8 and represented by the symbol '''O'''. It is a member of the [[chalcogen]] [[Group (periodic table)|group]] on the [[periodic table]], and is a highly reactive [[nonmetal]]lic [[period 2 element]] that readily forms [[chemical compound|compounds]] (notably [[oxide]]s) with almost all other elements. At [[standard temperature and pressure]] two atoms of the element [[chemical bond|bind]] to form dioxygen, a colorless, odorless, tasteless [[diatomic molecule|diatomic]] [[gas]] with the formula {{chem|O|2}}. Oxygen is the [[Abundance of the chemical elements|third most abundant]] element in the universe by mass after [[hydrogen]] and [[helium]]<ref name="NBB297">[[#Reference-idEmsley2001|Emsley 2001]], p.297</ref> and the [[abundances of the elements (data page)|most abundant]] element by mass in the Earth's crust.<ref name="lanl"/> Oxygen constitutes 88.8% of the mass of water and 20.9% of the volume of air.<ref name="ECE500"/>

All major classes of structural molecules in living organisms, such as [[protein]]s, [[carbohydrate]]s, and [[fat]]s, contain oxygen, as do the major [[inorganic compound]]s that comprise animal shells, teeth, and bone. Oxygen in the form of {{chem|O|2}} is produced from water by [[cyanobacteria]], [[algae]] and plants during [[photosynthesis]] and is used in [[cellular respiration]] for all complex life. Oxygen is toxic to [[anaerobic organism]]s, which were the dominant form of early life on Earth until {{chem|O|2}} began to accumulate in the atmosphere 2.5&nbsp;billion years ago.<ref>{{cite press release|title=NASA Research Indicates Oxygen on Earth 2.5 Billion Years Ago|url=http://www.nasa.gov/lb/centers/ames/news/releases/2007/07_70AR.html|publisher=[[NASA]]|date=2007-09-27|accessdate=2008-03-13}}</ref> Another form ([[allotropes of oxygen|allotrope]]) of oxygen, [[ozone]] ({{chem|O|3}}), helps protect the biosphere from [[ultraviolet radiation]] with the high-altitude [[ozone layer]], but is a pollutant near the surface where it is a by-product of smog. 

Oxygen was independently discovered by [[Joseph Priestley]] and [[Carl Wilhelm Scheele]] in the 1770s, but Priestley is usually given priority because he published his findings first. The name ''oxygen'' was coined in 1777 by [[Antoine Lavoisier]],<ref name=mellor>[[#Reference-idMellor1939|Mellor 1939]]</ref> whose experiments with oxygen helped to discredit the then-popular phlogiston theory of [[combustion]] and [[corrosion]]. Oxygen is produced industrially by fractional distillation of liquefied air, use of [[zeolite]]s to remove [[carbon dioxide]] and [[nitrogen]] from air, [[electrolysis of water]] and other means. Uses of oxygen include the production of steel, plastics and textiles; rocket propellant; [[oxygen therapy]]; and life support in aircraft, submarines, spaceflight and diving.

== Characteristics ==
===Structure===
[[Image:Electron shell 008 Oxygen.svg|thumb|left|Electron shell diagram of oxygen]]
At [[standard temperature and pressure]], oxygen is a colorless, odorless gas with the [[molecule| molecular]] formula {{chem|O|2}}, in which the two oxygen atoms are [[chemical bond|chemically bonded]] to each other with a spin triplet [[electron configuration]]. This bond has a [[bond order]] of two, and is often over-simplified in description as a [[double bond]].<ref>{{citeweb|url=http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/mo.html#bond|title=Molecular Orbital Theory|publisher = Purdue University | accessdate =2008-01-28}}</ref>

[[Triplet oxygen]] is the [[ground state]] of the {{chem|O|2}} molecule.<ref name="BiochemOnline">{{cite web | title = Biochemistry Online | url = http://employees.csbsju.edu/hjakubowski/classes/ch331/bcintro/default.html| chapterurl=http://employees.csbsju.edu/hjakubowski/classes/ch331/oxphos/oldioxygenchem.html |chapter=Chapter 8: Oxidation-Phosphorylation, the Chemistry of Di-Oxygen|first=Henry|last=Jakubowski|accessdate=2008-01-28|publisher=Saint John's University}}</ref> The electron configuration of the molecule has two unpaired electrons occupying two [[degenerate orbitals|degenerate]] [[molecular orbital]]s.<ref>An orbital is a concept from [[quantum mechanics]] that models an electron as a [[Wave–particle duality|wave-like particle]] that has a spacial distribution about an atom or molecule.</ref> These orbitals are classified as [[antibonding]] (weakening the bond order from three to two), so the diatomic oxygen bond is weaker than the diatomic [[nitrogen]] triple bond in which all bonding molecular orbitals are filled, but some antibonding orbitals are not.<ref name="BiochemOnline"/> 

In normal triplet form, {{chem|O|2}} molecules are [[paramagnetism| paramagnetic]]&mdash;they form a magnet in the presence of a magnetic field&mdash;because of the [[Spin (physics)|spin]] [[magnetic moment]]s of the unpaired electrons in the molecule, and the negative exchange energy between neighboring {{chem|O|2}} molecules.<ref name="NBB303"/> Liquid oxygen is attracted to a magnet to a sufficient extent that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.<ref>{{cite web | url = http://genchem.chem.wisc.edu/demonstrations/Gen_Chem_Pages/0809bondingpage/liquid_oxygen.htm | title = Demonstration of a bridge of liquid oxygen supported against its own weight between the poles of a powerful magnet | publisher = University of Wisconsin-Madison Chemistry Department DEMONSTRATION LAB | accessdate = 2007-12-15 }}</ref><ref>Oxygen's paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine the purity of gaseous oxygen. ({{cite web | url = http://www.servomex.com/oxygen_gas_analyser.html | title = Company literature of Oxygen analyzers (triplet) | publisher = Servomex | accessdate = 2007-12-15 }})</ref>

[[Singlet oxygen]], a name given to several higher-energy species of molecular {{chem|O|2}} in which all the electron spins are paired, is much more reactive towards common [[organic compound|organic molecules]]. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.<ref>[[#Reference-idKrieger-Liszkay2005|Krieger-Liszkay 2005]], 337-46</ref> It is also produced in the troposphere by the photolysis of ozone by light of short wavelength,<ref name=harrison>[[#Reference-idHarrison1990|Harrison 1990]]</ref> and by the immune system as a source of active oxygen.<ref name=immune-ozone>[[#Reference-idWentworth2002|Wentworth 2002]]</ref> [[Carotenoid]]s in photosynthetic organisms (and possibly also in animals) play a major role in absorbing energy from singlet oxygen and converting it to the unexcited ground state before it can cause harm to tissues.<ref>[[#Reference-idHirayama1994|Hirayama 1994]], 149-150</ref>

=== Allotropes ===
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{{main|Allotropes of oxygen}}
[[Image:Ozone-montage.png|thumb||left|upright|Ozone is a rare gas on Earth found mostly in the stratosphere.]]
The common [[Allotropy|allotrope]] of elemental oxygen on Earth is called dioxygen, {{Chem|O|2}}. It has a bond length of 121&nbsp;[[Picometre|pm]] and a bond energy of 498&nbsp;[[joule per mole|kJ·mol<sup>-1</sup>]].<ref>{{citeweb|last=Chieh|first=Chung|title=Bond Lengths and Energies|url=http://www.science.uwaterloo.ca/~cchieh/cact/c120/bondel.html|publisher= University of Waterloo|accessdate=2007-12-16}}</ref> This is the form that is used by complex forms of life, such as animals, in [[cellular respiration]] (see [[#Biological role|Biological role]]) and is the form that is a major part of the Earth's atmosphere (see [[#Occurrence|Occurrence]]). Other aspects of {{Chem|O|2}} are covered in the remainder of this article. 

Trioxygen ({{chem|O|3}}) is usually known as [[ozone]] and is a very reactive allotrope of oxygen that is damaging to lung tissue.<ref name="GuideElem48">[[#Reference-idStwertka1998|Stwertka 1998]], p.48</ref> Ozone is produced in the upper atmosphere when {{chem|O|2}} combines with atomic oxygen made by the splitting of {{chem|O|2}} by [[ultraviolet]] (UV) radiation.<ref name=mellor/> Since ozone absorbs strongly in the UV region of the spectrum, it functions as a protective radiation shield for the planet (see [[ozone layer]]).<ref name=mellor/>  Near the earth's surface, however, it is a [[air pollution| pollutant]] formed as a by-product of automobile exhaust.<ref name="GuideElem49">[[#Reference-idStwertka1998|Stwertka 1998]], p.49</ref>  

The [[Metastability in molecules|metastable]] molecule tetraoxygen ({{chem|O|4}}) was discovered in 2001,<ref name=o4>[[#Reference-idCacace2001|Cacace 2001]], 4062</ref><ref name=newform>{{citenews|first=Phillip|last=Ball|url=http://www.nature.com/news/2001/011122/pf/011122-3_pf.html|title=New form of oxygen found|work = Nature News|date=2001-09-16|accessdate=2008-01-09}}</ref> and was assumed to exist in one of the six phases of solid oxygen. It was proven in 2006 that that phase, created by pressurizing {{chem|O|2}} to 20 [[Pascal (unit)|GPa]], is in fact a rhombohedral {{chem|O|8}} [[Cluster chemistry|cluster]].<ref>[[#Reference-idLundegaard2006|Lundegaard 2006]], 201–04</ref> This cluster has the potential to be a much more powerful [[oxidizing agent|oxidizer]] than either {{chem|O|2}} or {{chem|O|3}} and may therefore be used in rocket fuel.<ref name=o4/><ref name=newform/> A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa<ref>[[#Reference-idDesgreniers1990|Desgreniers 1990]], 1117–22</ref> and it was shown in 1998 that at very low temperatures, this phase becomes superconducting.<ref>[[#Reference-idShimizu1998|Shimizu 1998]], 767–69</ref>

===Physical properties===
Oxygen is more [[Solubility|soluble]] in water than nitrogen; water contains approximately 1 molecule of {{chem|O|2}} for every 2 molecules of {{chem|N|2}}, compared to an atmospheric ratio of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg·L<sup>−1</sup>) dissolves at 0 °C than at 20 °C (7.6 mg·L<sup>−1</sup>).<ref name="NBB299"/><ref>{{citeweb|url=http://www.engineeringtoolbox.com/air-solubility-water-d_639.html|title=Air solubility in water|accessdate=2007-12-21|publisher=The Engineering Toolbox}}</ref> At 25 °C and 1 [[atmosphere (unit)|atm]] of air, freshwater contains about 6.04&nbsp;[[Litre|milliliter]]s&nbsp;(mL) of oxygen per [[liter]], whereas seawater contains about 4.95&nbsp;mL per liter.<ref>[[#Reference-idEvansClaiborne2006|Evans & Claiborne 2006]], 88</ref> At 5 °C the solubility increases to 9.0&nbsp;mL (50% more than at 25 °C) per liter for water and 7.2&nbsp;mL (45% more) per liter for sea water. 

Oxygen condenses at 90.20&nbsp;[[Kelvin|K]] (−182.95 °C, −297.31 °F), and freezes at 54.36&nbsp;K (−218.79 °C, −361.82 °F).<ref>[[#Reference-idLide2003|Lide 2003]], Section 4</ref> Both [[liquid oxygen|liquid]] and [[solid oxygen|solid]] {{chem|O|2}} are clear substances with a light sky-blue color caused by absorption in the red (in contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light). High-purity liquid {{chem|O|2}} is usually obtained by the fractional distillation of liquefied air;<ref>{{cite web | url = http://www.uigi.com/cryodist.html | title = Overview of Cryogenic Air Separation and Liquefier Systems | publisher = Universal Industrial Gases, Inc. | accessdate = 2007-12-15 }}</ref>  Liquid oxygen may also be produced by condensation out of air, using liquid nitrogen as a coolant. It is a highly-reactive substance and must be segregated from combustible materials.<ref>{{cite web | url = https://www.mathesontrigas.com/pdfs/msds/00225011.pdf |format=PDF| title = Liquid Oxygen Material Safety Data Sheet|publisher = Matheson Tri Gas | accessdate = 2007-12-15 }}</ref>

=== Isotopes and stellar origin ===
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[[Image:Evolved star fusion shells.svg|200px|left|thumb|Late in a massive star's life, <sup>16</sup>O concentrates in the O-shell, <sup>17</sup>O in the H-shell and [[oxygen-18|<sup>18</sup>O]] in the He-shell.]]

Naturally occurring oxygen is composed of three stable [[isotope]]s, <sup>16</sup>O, <sup>17</sup>O, and [[oxygen-18|<sup>18</sup>O]], with <sup>16</sup>O being the most abundant (99.762% [[natural abundance]]).<ref name="EnvChem-Iso">{{cite web|url=http://environmentalchemistry.com/yogi/periodic/O-pg2.html|title=Oxygen Nuclides / Isotopes|publisher=EnvironmentalChemistry.com|accessdate=2007-12-17}}</ref> Oxygen isotopes range in [[mass number]] from 12 to 28.<ref name="EnvChem-Iso"/> 

Most <sup>16</sup>O is synthesized at the end of the helium fusion process in stars but some is made in the  neon burning process.<ref name="Meyer2005">[[#Reference-idMeyer2005|Meyer 2005]], 9022</ref> <sup>17</sup>O is primarily made by the burning of hydrogen into helium during the CNO cycle, making it a common isotope in the hydrogen burning zones of stars.<ref name="Meyer2005"/> Most <sup>18</sup>O is produced when <sup>14</sup>N (made abundant from CNO burning) captures a <sup>4</sup>He nucleus, making <sup>18</sup>O common in the helium-rich zones of stars.<ref name="Meyer2005"/> 

Fourteen [[radioisotope]]s have been characterized, the most stable being <sup>15</sup>O with a [[half-life]] of 122.24&nbsp;seconds&nbsp;(s) and <sup>14</sup>O with a half-life of 70.606&nbsp;s.<ref name="EnvChem-Iso"/>  All of the remaining [[Radioactive decay|radioactive]] isotopes have half-lives that are less than 27&nbsp;s and the majority of these have half-lives that are less than 83&nbsp;milliseconds.<ref name="EnvChem-Iso"/> The most common decay mode of the isotopes lighter than <sup>16</sup>O is [[electron capture]] to yield nitrogen, and the most common mode for the isotopes heavier than <sup>18</sup>O is [[beta decay]] to yield [[fluorine]].<ref name="EnvChem-Iso"/>

=== Occurrence ===
Oxygen is the most abundant chemical element, by mass, in our biosphere, air, sea and land.
Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.<ref name="NBB297"/> About 0.9% of the Sun's mass is oxygen.<ref name="ECE500"/> Oxygen constitutes 49.2% of the Earth's crust by mass<ref name="lanl">
{{citeweb|url=http://periodic.lanl.gov/elements/8.html|title=Oxygen|publisher=Los Alamos National Laboratory|title=Oxygen|accessdate=2007-12-16}}
</ref> and is the major component of the world's oceans (88.8% by mass).<ref name="ECE500"/> It is the second most common component of the Earth's atmosphere, taking up 21.0% of its volume and 23.1% of its mass (some 10<sup>15</sup> tonnes).<ref name="NBB298">[[#Reference-idEmsley2001|Emsley 2001]], p.298</ref><ref name="ECE500"/><ref>Figures given are for values up to {{convert|50|mi|km}} above the surface</ref> Earth is unusual among the planets of the Solar System in having such a high concentration of oxygen gas in its atmosphere: Mars (with 0.1% {{chem|O|2}} by volume) and Venus have far lower concentrations.  However, the {{chem|O|2}} surrounding these other planets is produced solely by ultraviolet radiation impacting oxygen-containing molecules such as [[carbon dioxide]].  

[[Image:AYool WOA surf O2.png|left|thumb|Cold water holds more dissolved {{chem|O|2}}.]]
The unusually high concentration of oxygen on Earth is the result of the oxygen cycle. This biogeochemical cycle describes the movement of oxygen within and between its three main reservoirs on Earth: the atmosphere, the biosphere, and the lithosphere. The main driving factor of the oxygen cycle is photosynthesis, which is responsible for modern Earth's atmosphere. Because of the vast amounts of oxygen gas available in the atmosphere, even if all photosynthesis were to cease completely, it would take all the oxygen-consuming processes at the present rate at least another 5,000 years to strip all the {{chem|O|2}} from the atmosphere.<ref>[[#Reference-idWalker1980|Walker 1980]]</ref><ref>This is calculated by dividing all the free {{chem|O|2}} in the atmosphere to the rate it is used for respiration by the entire biosphere. This is obviously an extreme calculation since most organisms would die well before the pressure of {{chem|O|2}} fell to zero, and therefore the rate of consumption would decrease significantly from the present rate.</ref>

Free oxygen also occurs in solution in the world's water bodies. The increased solubility of {{chem|O|2}} at lower temperatures (see [[#Physical properties|Physical properties]]) has important implications for ocean life, as polar oceans support a much higher density of life due to their higher oxygen content.<ref>From The Chemistry and Fertility of Sea Waters by H.W. Harvey, 1955, citing C.J.J. Fox, "On the coefficients of absorption of atmospheric gases in sea water", Publ. Circ. Cons. Explor. Mer, no. 41, 1907. Harvey however notes that according to later articles in Nature the values appear to be about 3% too high.</ref> [[Water pollution|Polluted water]] may have reduced amounts of {{chem|O|2}} in it, depleted by decaying algae and other biomaterials (see eutrophication). Scientists assess this aspect of water quality by measuring the water's biochemical oxygen demand, or the amount of {{chem|O|2}} needed to restore it to a normal concentration.<ref name="NBB301">[[#Reference-idEmsley2001|Emsley 2001]], p.301</ref>

== Biological role ==
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===Photosynthesis and respiration===
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[[Image:Oxygen evolving complex.png|left|thumb|Oxygen evolution by water oxidation during photosynthesis. The jagged lines represent four photons oxidizing the central cluster of the oxygen evolving complex by exciting and removing four electrons through a cycle of ''S-states''.]]

In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis. Green algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth and the rest is produced by terrestrial plants.<ref>[[#Reference-idFenical1983|Fenical 1983]], "Marine Plants"</ref>

A simplified overall formula for photosynthesis is:<ref>[[#Reference-idBrown2003|Brown 2003]], 958</ref> 
::6{{chem|C||O|2}} + 6{{chem|H|2|O}} + [[photon]]s &rarr; {{chem|C|6|H|12|O|6}} + 6{{chem|O|2}}  <small>(or simply carbon dioxide + water + sunlight → glucose + dioxygen)</small>

Photolytic oxygen evolution occurs in the thylakoid membranes of photosynthetic organisms and requires the energy of four photons.<ref>Thylakoid membranes are part of chloroplasts in algae and plants while they simply are one of many membrane structures in cyanobacteria. In fact, chloroplasts are thought to have evolved from cyanobacteria that were once symbiotic partners with the progenerators of plants and algae.</ref> Many steps are involved, but the result is the formation of a proton gradient across the thylakoid membrane, which is used to synthesize [[Adenosine triphosphate|ATP]] via photophosphorylation.<ref name="Raven">[[#Reference-idRaven2005|Raven 2005]], 115–27</ref> The {{chem|O|2}} remaining after oxidation of the water molecule is released into the atmosphere.<ref>Water oxidation is catalyzed by a [[manganese]]-containing [[enzyme]] complex known as the oxygen evolving complex (OEC) or water-splitting complex found associated with the lumenal side of thylakoid membranes. Manganese is an important [[Cofactor (biochemistry)|cofactor]], and [[calcium]] and [[chloride]] are also required for the reaction to occur.(Raven 2005)</ref>

Molecular dioxygen, {{chem|O|2}}, is essential for [[cellular respiration]] in all [[aerobic organism]]s. Oxygen is used in [[Mitochondrion|mitochondria]] to help generate [[adenosine triphosphate]] (ATP) during [[oxidative phosphorylation]]. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as:
::{{chem|C|6|H|12|O|6}} + 6{{chem|O|2}} → 6{{chem|C||O|2}} + 6{{chem|H|2|O}} + 2880 kJ·mol<sup>-1</sup>

In [[vertebrate]]s, {{chem|O|2}} is [[diffusion| diffused]] through membranes in the lungs and into [[red blood cell]]s. [[Hemoglobin]] binds {{chem|O|2}}, changing its color from bluish red to bright red.<ref>CO<sub>2</sub> is released from another part of hemoglobin (see [[Bohr effect]])</ref><ref name="GuideElem48"/> Other animals use [[hemocyanin]] (molluscs and some arthropods) or hemerythrin (spiders and lobsters).<ref name="NBB298"/> A liter of blood can dissolve 200&nbsp;cc of {{chem|O|2}}.<ref name="NBB298"/>

[[Reactive oxygen species]], such as [[superoxide]] ion (O<sub>2</sub><sup>−</sup>) and [[hydrogen peroxide]] ({{chem|H|2|O|2}}), are dangerous by-products of oxygen use in organisms.<ref name="NBB298"/> Parts of the [[immune system]] of higher organisms, however, create peroxide, superoxide, and singlet oxygen to destroy invading microbes. Reactive oxygen species also play an important role in the [[hypersensitive response]] of plants against pathogen attack.<ref name="Raven"/>

An adult human in rest [[breathing|inhales]] 1.8 to 2.4 grams of oxygen per minute.<ref> "For humans, the normal volume is 6-8 liters per minute." [http://www.patentstorm.us/patents/6224560-description.html]</ref> This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year. <ref>(1.8 grams)*(60 minutes)*(24 hours)*(365 days)*(6.6 billion people)/1,000,000=6.24 billion tonnes</ref>

===Build-up in the atmosphere===
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[[Image:Oxygenation-atm.png|thumb|left|300px|O<sub>2</sub> build-up in Earth's atmosphere: 1) no O<sub>2</sub> produced; 2) O<sub>2</sub> produced, but absorbed in oceans & seabed rock; 3) O<sub>2</sub> starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer; 4-5) O<sub>2</sub> sinks filled and the gas accumulates]]

Free oxygen gas was almost nonexistent in Earth's atmosphere before photosynthetic archaea and bacteria evolved. Free oxygen first appeared in significant quantities during the Paleoproterozoic era (between 2.5 and 1.6&nbsp;billion years ago). At first, the oxygen combined with dissolved [[iron]] in the oceans to form banded iron formations. Free oxygen started to gas out of the oceans 2.7&nbsp;billion years ago, reaching 10% of its present level around 1.7&nbsp;billion years ago.<ref name="Campbell">[[#Reference-idCampbell2005|Campbell 2005]], 522–23</ref>

The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the anaerobic organisms then living to extinction during the oxygen catastrophe about 2.4&nbsp;billion years ago. However, [[cellular respiration]] using O<sub>2</sub> enables [[aerobic organism]]s to produce much more ATP than anaerobic organisms, helping the former to dominate Earth's biosphere.<ref name="Freeman">[[#Reference-idFreeman2005|Freeman 2005]], 214, 586</ref> Photosynthesis and cellular respiration of {{chem|O|2}} allowed for the evolution of eukaryotic cells and ultimately complex multicellular organisms such as plants and animals.

Since the beginning of the Cambrian era 540&nbsp;million years ago, {{chem|O|2}} levels have fluctuated between 15% and 30% per volume.<ref name="geologic">[[#Reference-idBerner1999|Berner 1999]], 10955–57</ref> Towards the end of the Carboniferous era (about 300&nbsp;million years ago) atmospheric {{chem|O|2}} levels reached a maximum of 35% by volume,<ref name=geologic/> allowing insects and amphibians to grow much larger than today's species. Human activities, including the burning of 7&nbsp;billion tonnes of fossil fuels each year have had very little effect on the amount of free oxygen in the atmosphere.<ref name="NBB303"/> At the current rate of photosynthesis it would take about 2,000&nbsp;years to regenerate the entire {{chem|O|2}} in the present atmosphere.<ref>[[#Reference-idDole1965|Dole 1965]], 5–27</ref>

== History ==
===Early experiments===
[[Image:Philos experiment of the burning candle.PNG|100px||thumb|left|Philo's experiment inspired later investigators.]]
One of the first known experiments on the relationship between [[combustion]] and air was conducted by the second century&nbsp;BCE Greek writer on mechanics, Philo of Byzantium. In his work ''Pneumatica'', Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck.<ref>[[#Reference-idJastrow1936|Jastrow 1936]], 171</ref>
Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries later [[Leonardo da Vinci]] built on Philo's work by observing that a portion of air is consumed during combustion and [[respiration (physiology)|respiration]].<ref name="ECE499">[[#Reference-idCook1968|Cook & Lauer 1968]], p.499.</ref>

In the late 17th&nbsp;century, Robert Boyle proved that air is necessary for combustion. English chemist [[John Mayow]] refined this work by showing that fire requires only a part of air that he called ''spiritus nitroaereus'' or just ''nitroaereus''.<ref name="EB1911">[[#Reference-idEB1911|''Britannica'' contributors 1911]], "John Mayow"</ref> 
In one experiment he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.<ref name="WoC">[[#Reference-idWoC2005|''World of Chemistry'' contributors 2005]], "John Mayow"</ref> 
From this he surmised that nitroaereus is consumed in both [[Respiration (physiology)|respiration]] and combustion. 

Mayow observed that [[antimony]] increased in weight when heated, and inferred that the nitroaereus must have combined with it.<ref name="EB1911"/> He also thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body.<ref name="EB1911"/> Accounts of these and other experiments and ideas were published in 1668 in his work ''Tractatus duo'' in the tract "De respiratione".<ref name="WoC"/>

===Phlogiston theory===
[[Image:Georg Ernst Stahl.png|left|thumb|170px|[[Georg Ernst Stahl]] helped develop and popularize the phlogiston theory.]]

Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all produced oxygen in experiments in the 17th century but none of them recognized it as an element.<ref name="NBB299">[[#Reference-idEmsley2001|Emsley 2001]], p.299</ref> This may have been in part due to the prevalence of the philosophy of [[combustion]] and [[corrosion]] called the phlogiston theory, which was then the favored explanation of those processes.

Established in 1667 by the German alchemist [[J. J. Becher]], and modified by the chemist [[Georg Ernst Stahl]] by 1731,<ref name="morris">[[#Reference-idMorris2003|Morris 2003]]</ref> 
phlogiston theory stated that all combustible materials were made of two parts. One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or [[calx]].<ref name="ECE499">[[#Reference-idCook1968|Cook & Lauer 1968]], p.499</ref>

Highly combustible materials that leave little residuum, such as wood or coal, were thought to be made mostly of phlogiston; whereas non-combustible substances that corrode, such as iron, contained very little. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea; instead, it was based on observations of what happens when something burns, that most common objects appear to become lighter and seem to lose something in the process.<ref name="ECE499"/> The fact that a substance like wood actually ''gains'' overall weight in burning was hidden by the buoyancy of the gaseous combustion products. Indeed one of the first clues that the phlogiston theory was incorrect was that metals, too, gain weight in rusting (when they were supposedly losing phlogiston).

===Discovery===
[[Image:Carl Wilhelm Scheele from Familj-Journalen1874.png|thumb|170px|Carl Wilhelm Scheele beat Priestley to the discovery but published afterwards.]]
Oxygen was first discovered by Swedish pharmacist [[Carl Wilhelm Scheele]]. He had produced oxygen gas by heating mercuric oxide and various [[nitrate]]s by about 1772.<ref name="ECE499"/><ref name="ECE500"/>   Scheele called the gas 'fire air' because it was the only known supporter of combustion. He wrote an account of this discovery in a manuscript he titled ''Treatise on Air and Fire'', which he sent to his publisher in 1775. However, that document was not published until 1777.<ref name="NBB300">[[#Reference-idEmsley2001|Emsley 2001]], p.300</ref>   

[[Image:PriestleyFuseli.jpg|thumb|170px|left|Joseph Priestley is usually given priority in the discovery.]]
In the meantime, an experiment was conducted by the British clergyman [[Joseph Priestley]] on August 1 1774 focused sunlight on [[mercury(II) oxide|mercuric oxide]] (HgO) inside a glass tube, which liberated a gas he named 'dephlogisticated air'.<ref name="ECE500">[[#Reference-idCook1968|Cook & Lauer 1968]], p.500</ref> He noted that candles burned brighter in the gas and that a mouse was more active and lived longer while breathing it. After breathing the gas himself, he wrote: "The feeling of it to my lungs was not sensibly different from that of common air, but I fancied that my breast felt peculiarly light and easy for some time afterwards."<ref name="NBB299"/> Priestley published his findings in 1775 in a paper titled "An Account of Further Discoveries in Air" which was included in the second volume of his book titled ''Experiments and Observations on Different Kinds of Air''.<ref>[[#Reference-idPriestley1775|Priestley 1775]], 384–94</ref><ref name="ECE499"/> Because he had published his findings first, Priestley is usually given priority in the discovery.

The noted French chemist [[Antoine Lavoisier|Antoine Laurent Lavoisier]] later claimed to have discovered the new substance independently. However, Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele also posted a letter to Lavoisier on September 30 1774 that described his own discovery of the previously-unknown substance, but Lavoisier never acknowledged receiving it (a copy of the letter was found in Scheele's belongings after his death).<ref name="NBB300"/>

===Lavoisier's contribution===
[[Image:Antoine lavoisier.jpg|thumb|left|170px|Antoine Lavoisier discredited the Phlogiston theory.]]
What Lavoisier did indisputably do (although this was disputed at the time) was to conduct the first adequate quantitative experiments on [[oxidation]] and give the first correct explanation of how combustion works.<ref name="ECE500"/> He used these and similar experiments, all started in 1774, to discredit the phlogiston theory and to prove that the substance discovered by Priestley and Scheele was a [[chemical element]].  

In one experiment, Lavoisier observed that there was no overall increase in weight when [[tin]] and air were heated in a closed container.<ref name="ECE500"/> He noted that air rushed in when he opened the container, which indicated that part of the trapped air had been consumed. He also noted that the tin had increased in weight and that increase was the same as the weight of the air that rushed back in. This and other experiments on combustion were documented in his book ''Sur la combustion en général'', which was published in 1777.<ref name="ECE500"/> In that work, he proved that air is a mixture of two gases; 'vital air', which is essential to combustion and respiration, and ''azote'' (Gk. ''{{Polytonic|ἄζωτον}}'' "lifeless"), which did not support either.   

Lavoisier renamed 'vital air' to ''oxygène'' in 1777 from the [[Ancient Greek|Greek]] roots ''{{Polytonic|ὀξύς}} (oxys)'' ([[acid]], literally "sharp," from the taste of acids) and ''-γενής (-genēs)'' (producer, literally begetter), because he mistook oxygen to be a constituent of all acids.<ref name=mellor>[[#Reference-idMellor1939|Mellor 1939]]</ref> ''Azote'' later became  ''[[nitrogen]]'' in English, although it has kept the name in French and several other European languages.<ref name="ECE500"/> 

''Oxygen'' entered the English language despite opposition by English scientists and the fact that Priestley had priority.  This is partly due to a poem praising the gas titled "Oxygen" in the popular book ''The Botanic Garden'' (1791) by Erasmus Darwin, grandfather of [[Charles Darwin]].<ref name="NBB300"/>

===Later history===
[[Image:Goddard and Rocket.jpg|thumb|170px|Robert H. Goddard and a liquid oxygen-gasoline rocket]]
John Dalton's original atomic hypothesis assumed that all elements were monoatomic and that the atoms in compounds would normally have the simplest atomic ratios with respect to one another. For example, Dalton assumed that water's formula was HO, giving the [[atomic mass]] of oxygen as 8 times that of hydrogen, instead of the modern value of about 16.<ref>{{cite web| title = The Interactive Textbook of PFP96 |chapter= Do We Take Atoms for Granted?| chapterurl=http://www.physics.upenn.edu/courses/gladney/mathphys/subsubsection1_1_3_2.html |url=http://www.physics.upenn.edu/courses/gladney/mathphys/Contents.html |first=Dennis |last=DeTurck |coauthors=Gladney, Larry  and Pietrovito, Anthony| publisher=University of Pennsylvania|year=1997|accessdate=2008-01-28}}</ref>  In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen; and by 1811 Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called [[Avogadro's law]] and the assumption of diatomic elemental molecules.<ref>[[#Reference-idRoscoe1883|Roscoe 1883]], 38</ref><ref>However, these results were mostly ignored until 1860.  Part of this rejection was due to the belief that atoms of one element would have no [[chemical affinity]] towards atoms of the same element, and part was due to apparent exceptions to Avogadro's law that were not explained until later in terms of dissociating molecules.</ref>

By the late 19th century scientists realized that air could be liquefied, and its components isolated, by compressing and cooling it. Using a cascade method, Swiss chemist and physicist Raoul Pierre Pictet [[evaporation| evaporated]] liquid [[sulfur dioxide]] in order to liquefy [[carbon dioxide]], which in turn was evaporated to cool oxygen gas enough to liquefy it. He sent a telegram on December 22 1877 to the French Academy of Sciences in Paris announcing his discovery of [[liquid oxygen]].<ref name="BES707">[[#Reference-idDaintith1994|Daintith 1994]], p.707</ref> <!-- NEEDS TO BE CHECKED W/ GOOD CITE The telegram read  "Oxygen liquefied to-day under 320 atmospheres and 140 degrees of cold by combined use of sulfurous and carbonic acid." /NEEDS TO BE CHECKED /W GOOD CITE --> Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen.<ref name="BES707"/> Only a few drops of the liquid were produced in either case so no meaningful analysis could be conducted. 

In 1891 Scottish chemist James Dewar was able to produce enough liquid oxygen to study.<ref name="NBB303">[[#Reference-idEmsley2001|Emsley 2001]], p.303</ref> The first commercially-viable process for producing liquid oxygen was independently developed in 1895 by German engineer Carl von Linde and British engineer William Hampson. Both men lowered the temperature of air until it liquefied and then [[distillation|distilled]] the component gases by boiling them off one at a time and capturing them.<ref name="HPAM">[[#Reference-idHPaM2005|''How Products are Made'' contributors]], "Oxygen"</ref> Later, in 1901, oxyacetylene welding was demonstrated for the first time by burning a mixture of [[acetylene]] and compressed {{chem|O|2}}. This method of welding and cutting metal later became common.<ref name="HPAM"/>

In 1923 the American scientist Robert H. Goddard became the first person to develop a rocket engine; the engine used gasoline for fuel and liquid oxygen as the oxidizer. Goddard successfully flew a small liquid-fueled rocket 56&nbsp;m at 97&nbsp;km/h on March 16 1926 in Auburn, Massachusetts, USA.<ref name="HPAM"/><ref>{{cite web|title=Goddard-1926|url=http://grin.hq.nasa.gov/ABSTRACTS/GPN-2002-000132.html|publisher=NASA|accessdate=2007-11-18}}</ref>

== Industrial production ==
Two major methods are employed to produce the 100&nbsp;million&nbsp;tonnes of {{chem|O|2}} extracted from air for industrial uses annually.<ref name="NBB300"/> The most common method is to fractionally-distill liquefied air into its various components, with nitrogen {{chem|N|2}} [[distillation| distilling]] as a vapor while oxygen {{chem|O|2}} is left as a liquid.<ref name="NBB300"/> 

The other major method of producing {{chem|O|2}} gas involves passing a stream of clean, dry air through one bed of a pair of identical [[zeolite]] molecular sieves, which absorbs the nitrogen and delivers a gas stream that is 90%&nbsp;to&nbsp;93% {{chem|O|2}}.<ref name="NBB300"/> Simultaneously, nitrogen gas is released from the other nitrogen-saturated zeolite bed, by reducing the chamber operating pressure and diverting part of the oxygen gas from the producer bed through it, in the reverse direction of flow. After a set cycle time the operation of the two beds is interchanged, thereby allowing for a continuous supply of gaseous oxygen to be pumped through a pipeline. This is known as pressure swing adsorption. Oxygen gas is increasingly obtained by these non-cryogenic technologies (see also the related vacuum swing adsorption).<ref>{{citeweb|url=http://www.uigi.com/noncryo.html|title=Non-Cryogenic Air Separation Processes|year=2003|accessdate=2007-12-16|publisher=UIG Inc.}}</ref>

Oxygen gas can also be produced through [[electrolysis of water]] into molecular oxygen and hydrogen. A similar method is the electrocatalytic {{chem|O|2}} evolution from [[oxide]]s and [[oxoacid]]s. Chemical catalysts can be used as well, such as in [[chemical oxygen generator]]s or oxygen candles that are used as part of the life-support equipment on submarines, and are still part of standard equipment on commercial airliners in case of depressurization emergencies. Another air separation technology involves forcing air to dissolve through ceramic membranes based on [[zirconium dioxide]] by either high pressure or an electric current, to produce nearly pure {{chem|O|2}} gas.<ref name="NBB301"/>

In large quantities, the price of liquid oxygen in 2001 was approximately $0.21/kg.<ref>{{citation|quote=NASAFacts FS-2001-09-015-KSC|title=Space Shuttle Use of Propellants and Fluids|publisher=National Aeronautics and Space Administration|date=2001=09|url=http://www-pao.ksc.nasa.gov/kscpao/nasafact/ps/SSP.ps|accessdate=2007-12-16}}</ref>  Since the primary cost of production is the energy cost of liquefying the air, the production cost will change as energy cost varies.

For reasons of economy oxygen is often transported in bulk as a liquid in specially-insulated tankers, since one [[litre]] of liquefied oxygen is equivalent to 840&nbsp;liters of gaseous oxygen at atmospheric pressure and 20 °[[Celsius|C]].<ref name="NBB300"/> Such tankers are used to refill bulk liquid oxygen storage containers, which stand outside hospitals and other institutions with a need for large volumes of pure oxygen gas. Liquid oxygen is passed through heat exchangers, which convert the cryogenic liquid into gas before it enters the building. Oxygen is also stored and shipped in smaller cylinders containing the compressed gas; a form that is useful in certain portable medical applications and oxy-fuel welding and cutting.<ref name="NBB300"/>

== Applications ==
{{seealso|Breathing gas|Redox|Combustion}}
===Medical===
[[Image:Home oxygen concentrator.jpg|thumb|right|An [[oxygen concentrator]] in an [[emphysema]] patient's house]]
Uptake of {{chem|O|2}} from the air is the essential purpose of [[Respiration (physiology)|respiration]], so oxygen supplementation is used in [[medicine]]. [[Oxygen therapy]] is used to treat [[emphysema]], [[pneumonia]], some heart disorders, and any [[disease]] that impairs the body's ability to take up and use gaseous oxygen.<ref name="ECE510">[[#Reference-idCook1968|Cook & Lauer 1968]], p.510</ref> Treatments are flexible enough to be used in hospitals, the patient's home, or increasingly by portable devices. [[Oxygen tent]]s were once commonly used in oxygen supplementation, but have since been replaced mostly by the use of [[oxygen mask]]s or [[nasal cannula]]s. 

[[Hyperbaric medicine|Hyperbaric]] (high-pressure) medicine uses special oxygen chambers to increase the [[partial pressure]] of {{chem|O|2}} around the patient and, when needed, the medical staff.  [[Carbon monoxide poisoning]], [[gas gangrene]], and [[decompression sickness]] (the 'bends') are sometimes treated using these devices. Increased {{chem|O|2}} concentration in the lungs helps to displace [[carbon monoxide]] from the heme group of [[hemoglobin]]. Oxygen gas is poisonous to the [[anaerobic bacteria]] that cause gas gangrene, so increasing its partial pressure helps kill them. Decompression sickness occurs in divers who decompress too quickly after a dive, resulting in bubbles of inert gas, mostly nitrogen and argon, forming in their blood. Increasing the pressure of {{chem|O|2}} as soon as possible is part of the treatment.<ref name="ECE510"/> 

Oxygen  is also used medically for patients who require [[mechanical ventilation]], often at concentrations above the 21% found in ambient air. 
<!-- NEEDS TO BE EXPANDED  The isotope <sup>15</sup>O was experimentally used in [[positron emission tomography]].<ref>[[#Reference-idAgostini1995|Agostini et al. 1968}}, 69-72</ref>   /NEEDS TO BE EXPANDED -->

===Life support and recreational use===
[[Image:GPN-2000-001069.jpg|thumb|left|Low pressure pure {{chem|O|2}} is used in space suits]]

A notable application of {{chem|O|2}} as a low-pressure [[breathing gas]] is in modern space suits, which surround their occupant's body with pressurized air. These devices use nearly pure oxygen at about one third normal pressure, resulting in a normal blood [[partial pressure]] of {{chem|O|2}}. This trade-off of higher oxygen concentration for lower pressure is needed to maintain flexible spacesuits.

Scuba divers and submariners also rely on artificially-delivered {{chem|O|2}}, but most often use normal pressure, and/or mixtures of oxygen and air. Pure or nearly pure {{chem|O|2}} use in diving at higher-than-sea-level pressures is usually limited to rebreather, decompression, or emergency treatment use at relatively shallow depths (~ 6 meters depth, or less). Deeper diving requires significant dilution of {{chem|O|2}} with other gases, such as nitrogen or helium, to help prevent [[oxygen toxicity]].

People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental {{chem|O|2}} supplies.<ref>The reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspired {{chem|O|2}} [[partial pressure]] nearer to that found at sea-level.</ref> Passengers traveling in (pressurized) commercial airplanes have an emergency supply of {{chem|O|2}} automatically supplied to them in case of cabin depressurization. Sudden cabin pressure loss activates [[chemical oxygen generator]]s above each seat, causing [[oxygen mask]]s to drop and forcing iron filings into the sodium chlorate inside the canister.<ref name="NBB301"/> A steady stream of oxygen gas is produced by the [[exothermic]] reaction. However, even this may pose a danger if inappropriately triggered: a ValuJet airplane crashed after use-date-expired {{chem|O|2}} canisters, which were being shipped in the cargo hold, activated and caused fire. The canisters were mis-labeled as ''empty'', and carried against dangerous goods regulations.<ref>{{citeweb |url=http://www.ntsb.gov/NTSB/brief.asp?ev_id=20001208X05743&key=1 |title=NTSB Summary report |publisher=National Transportation Safety Board |accessdate=2007-12-16}})</ref>

Oxygen, as a supposed mild [[euphoria (emotion)|euphoric]], has a history of recreational use in [[oxygen bar]]s and in sports. Oxygen bars are establishments, found in Japan, California, and Las Vegas, Nevada since the late 1990s that offer higher than normal {{chem|O|2}} exposure for a fee.<ref name="FDA-O2Bars">{{cite web |url=http://www.fda.gov/Fdac/features/2002/602_air.html |title=Oxygen Bars: Is a Breath of Fresh Air Worth It? |last=Bren |first=Linda |work=FDA Consumer magazine |publisher=U.S. Food and Drug Administration |date=November–December 2002 |accessdate=2007-12-23}}</ref> Professional athletes, especially in American football, also sometimes go off field between plays to wear oxygen masks in order to get a supposed "boost" in performance. However, the reality of a pharmacological effect is doubtful; a [[placebo]] or psychological boost being the most plausible explanation.<ref name="FDA-O2Bars"/> Available studies support a performance boost from enriched {{chem|O|2}} mixtures only if they are breathed ''during'' actual aerobic exercise.<ref>{{citeweb|url=http://www.pponline.co.uk/encyc/1008.htm|title=Ergogenic Aids|accessdate=2008-01-04|publisher=Peak Performance Online}}</ref> Other recreational uses include pyrotechnic applications, such as George Goble's five-second ignition of barbecue grills.<ref>{{cite web |url=http://www.bkinzel.de/misc/ghg/index.html |title=George Goble's extended home page (mirror)}}</ref><!--- Primary source; many secondary sources exist but they only provide less information and more ads --->

===Industrial===
[[Image:Clabecq JPG01.jpg|thumb|Most commercially-produced {{chem|O|2}} is used to smelt iron into steel.]]
Smelting of iron ore into steel consumes 55% of commercially-produced oxygen.<ref name="NBB301"/> In this process, {{chem|O|2}} is injected through a high-pressure lance into molten iron, which removes [[sulfur]] impurities and excess [[carbon]] as the respective oxides, SO<sub>2</sub> and CO<sub>2</sub>. The reactions are [[exothermic reaction| exothermic]], so the temperature increases to 1700 °[[Celsius|C]].<ref name="NBB301"/> 

Another 25% of commercially-produced oxygen is used by the chemical industry.<ref name="NBB301"/> [[Ethylene]] is reacted with {{chem|O|2}} to create [[ethylene oxide]], which, in turn, is converted into [[ethylene glycol]]; the primary feeder material used to manufacture a host of products, including antifreeze and [[polyester]] polymers (the precursors of many [[plastic]]s and fabrics).<ref name="NBB301"/> 

Most of the remaining 20% of commercially-produced oxygen is used in medical applications, metal cutting and welding, as an oxidizer in rocket fuel, and in [[water treatment]].<ref name="NBB301"/> Oxygen is used in oxyacetylene welding burning [[acetylene]] with {{chem|O|2}} to produce a very hot flame. In this process, metal up to 60&nbsp;[[centimetre|cm]] thick is first heated with a small oxy-acetylene flame and then quickly cut by a large stream of {{chem|O|2}}.<ref name="ECE508">[[#Reference-idCook1968|Cook & Lauer 1968]], p.508</ref> Rocket propulsion requires a fuel and an oxidizer. Larger rockets use liquid oxygen as their oxidizer, which is mixed and ignited with the fuel for propulsion.

===Scientific===
[[Image:Phanerozoic Climate Change.png|thumb|left|300px|500 million years of climate change vs [[oxygen-18|<sup>18</sup>O]]]]
Paleoclimatologists measure the ratio of [[oxygen-18]] and oxygen-16 in the shells and [[skeleton]]s of marine organisms to determine what the climate was like millions of years ago (see oxygen isotope ratio cycle). Seawater molecules that contain the lighter [[isotope]], oxygen-16, evaporate at a slightly faster rate than water molecules containing the 12% heavier oxygen-18; this disparity increases at lower temperatures.<ref name="NBB304">[[#Reference-idEmsley2001|Emsley 2001]], p.304</ref> During periods of lower global temperatures, snow and rain from that evaporated water tends to be higher in oxygen-16, and the seawater left behind tends to be higher in oxygen-18. Marine organisms then incorporate more oxygen-18 into their skeletons and shells than they would in a warmer climate.<ref name="NBB304"/> Paleoclimatologists  also directly measure this ratio in the water molecules of ice core samples that are up to several hundreds of thousands of years old.

Planetary geologists have measured different abundances of oxygen isotopes in samples from the Earth, the Moon, Mars, and meteorites, but were long unable to obtain reference values for the isotope ratios in the Sun, believed to be the same as those of the primordial solar nebula. However, analysis of a [[silicon]] wafer exposed to the solar wind in space and returned by the crashed Genesis spacecraft has shown that the Sun has a higher proportion of oxygen-16 than does the Earth. The measurement implies that an unknown process depleted oxygen-16 from the Sun's disk of protoplanetary material prior to the coalescence of dust grains that formed the Earth.<ref>{{cite journal
  | last = Hand
  | first = Eric
  | authorlink = 
  | coauthors = 
  | title = The Solar System's first breath
  | journal = Nature
  | volume = 452
  | issue = 
  | pages = 259
  | publisher = 
  | location = 
  | date = 2008-03-13
  | url = http://www.nature.com/news/2008/080313/full/452259a.html
  | doi = 10.1038/452259a
  | id = 
  | accessdate = }}</ref>

Oxygen presents two spectrophotometric absorption bands peaking at the wavelengths 687 and 760&nbsp;[[Nanometre|nm]]. Some remote sensing scientists have proposed using the measurement of the radiance coming from vegetation canopies in those bands to characterize plant health status from a satellite platform.<ref>[[#Reference-idMiller2003|Miller et al. 2003]]</ref> This approach exploits the fact that in those bands it is possible to discriminate the vegetation's reflectance from its [[fluorescence]], which is much weaker. The measurement is technically difficult owing to the low signal-to-noise ratio and the physical structure of vegetation; but it has been proposed as a possible method of monitoring the [[carbon cycle]] from satellites on a global scale.
{{-}}

== Compounds ==
<!-- DIRECT ALL FUTURE EXPANSION to Compounds of oxygen -->
[[Image:Stilles Mineralwasser.jpg|thumb|[[Water]] (H<sub>2</sub>O) is the most familiar oxygen compound.]]

The [[oxidation state]] of oxygen is −2 in almost all known compounds of oxygen. The oxidation state −1 is found in a few compounds such as [[peroxide]]s.<ref>[[#Reference-idGreenwood1997|Greenwood & Earnshaw 1997]], 28</ref> Compounds containing oxygen in other oxidation states are very uncommon: −1/2 ([[superoxide]]s), −1/3 ([[ozonide]]s), 0 ([[Allotropes of oxygen|elemental]], hypofluorous acid), +1/2 (dioxygenyl), +1 ([[dioxygen difluoride]]), and +2 ([[oxygen difluoride]]). 

===Oxides and other inorganic compounds===
<!-- DIRECT ALL FUTURE EXPANSION to [[Compounds of oxygen]] -->
[[Water]] (H<sub>2</sub>O) is the [[oxide]] of [[hydrogen]] and the most familiar oxygen compound. Hydrogen atoms are [[covalent bonding|covalently bonded]] to oxygen in a water molecule but also have an additional attraction (about 23.3&nbsp;kJ·mol<sup>−1</sup> per hydrogen atom) to an adjacent oxygen atom in a separate molecule.<ref>[[#Reference-idMaksyutenko2006|Maksyutenko et al. 2006]]</ref> These [[hydrogen bond]]s between water molecules hold them approximately 15% closer than what would be expected in a simple liquid with just [[Van der Waals force]]s.<ref>{{cite web|title=Water Hydrogen Bonding|last=Chaplin|first=Martin|url=http://www.lsbu.ac.uk/water/hbond.html|accessdate=2008-01-06|date=2008-01-04}}</ref><ref> Also, since oxygen has a higher electronegativity than hydrogen, the charge difference makes it a [[polar molecule]]. The interactions between the different [[dipole]]s of each molecule cause a net attraction force.</ref>

[[Image:Rust screw.jpg|thumb|left|Oxides, such as iron oxide or [[rust]], Fe<sub>2</sub>O<sub>3</sub>, form when oxygen combines with other elements.]]
Due to its [[electronegativity]], oxygen forms [[chemical bond]]s with almost all other elements at elevated temperatures to give corresponding [[oxide]]s. However, some elements readily form oxides at [[standard conditions for temperature and pressure]]; the [[rust]]ing of [[iron]] is an example. The surface of metals like [[aluminium]] and [[titanium]] are oxidized in the presence of air and become coated with a thin film of oxide that passivates the metal and slows further [[corrosion]]. Some of the transition metal oxides are found in nature as non-stoichiometric compounds, with a slightly less metal than the [[chemical formula]] would show. For example, the natural occurring [[Iron(II) oxide|FeO]] (wüstite) is actually written as {{chem|Fe|''1−x''|O}}, where ''x'' is usually around 0.05.<ref>[[#Reference-idSmart2005|Smart 2005]], 214</ref>

Oxygen as a compound is present in the atmosphere in trace quantities in the form of [[carbon dioxide]] ({{chem|C||O|2}}). The earth's crustal rock is composed in large part of oxides of [[silicon]] ([[Silicon dioxide|silica]] {{chem|Si||O|2}}, found in granite and sand), [[aluminium]] (aluminium oxide {{chem|Al|2|O|3}}, in bauxite and corundum), iron (iron(III) oxide {{chem|Fe|2|O|3}}, in hematite and [[rust]]) and other [[metal]]s.

The rest of the Earth's crust is also made of oxygen compounds, in particular [[calcium carbonate]] (in limestone) and [[silicate]]s (in feldspars).  Water-[[solubility|soluble]] silicates in the form of {{chem|Na|4|Si||O|4}}, {{chem|Na|2|Si||O|3}}, and {{chem|Na|2|Si|2|O|5}} are used as [[detergent]]s and adhesives.<ref name="ECE507">[[#Reference-idCook1968|Cook & Lauer 1968]], p.507</ref>

Oxygen also acts as a ligand for transition metals, forming metal–O<sub>2</sub> bonds with the [[iridium]] atom in [[Vaska's complex]],<ref>[[#Reference-idCrabtree2001|Crabtree 2001]], 152</ref> with the [[platinum]] in {{chem|Pt||F|6}},<ref name="ECE505">[[#Reference-Cook1968|Cook & Lauer 1968]], p.505</ref> and with the iron center of the [[heme]] group of [[hemoglobin]].

===Organic compounds and biomolecules===
<!-- DIRECT ALL FUTURE EXPANSION to [[Compounds of oxygen]] -->
[[Image:Acetone-3D-vdW.png|thumb|left|[[Acetone]] is an important feeder material in the chemical industry<br><small>(oxygen is in red, carbon in black and hydrogen in white)</small>.]] 
Among the most important classes of organic compounds that contain oxygen are (where "R" is an organic group): [[alcohol]]s (R-OH); [[ether]]s (R-O-R); [[ketone]]s (R-CO-R); [[aldehyde]]s (R-CO-H); [[carboxylic acid]]s (R-COOH); [[ester]]s (R-COO-R); [[acid anhydride]]s (R-CO-O-CO-R); and [[amide]]s (R-C(O)-NR<sub>2</sub>). There are many important organic [[solvent]]s that contain oxygen, including: [[acetone]], [[methanol]], [[ethanol]], [[Isopropyl alcohol|isopropanol]], [[furan]], [[tetrahydrofuran|THF]], [[diethyl ether]], [[1,4-Dioxane|dioxane]], [[ethyl acetate]], [[dimethylformamide|DMF]], [[dimethyl sulfoxide|DMSO]], [[acetic acid]], and [[formic acid]]. [[Acetone]] ((CH<sub>3</sub>)<sub>2</sub>CO) and [[phenol]] (C<sub>6</sub>H<sub>5</sub>OH) are used as feeder materials in the synthesis of many different substances. Other important organic compounds that contain oxygen are: [[glycerol]], [[formaldehyde]], [[glutaraldehyde]], [[citric acid]], [[acetic anhydride]], and [[acetamide]]. [[Epoxide]]s are [[ether]]s in which the oxygen atom is part of a ring of three atoms. 

Oxygen reacts spontaneously with many [[organic chemistry| organic]] compounds at or below room temperature in a process called autoxidation.<ref name="ECE506">[[#Reference-idCook1968|Cook & Lauer 1968]], p.506</ref> Most of the [[organic compound]]s that contain oxygen are not made by direct action of {{chem|O|2}}. Organic compounds important in industry and commerce that are made by direct oxidation of a precursor include [[ethylene oxide]] and [[peracetic acid]].<ref name="ECE507"/>

[[Image:ATP structure.svg|thumb|Oxygen represents more than 40% of the [[molecular mass]] of the [[Adenosine triphosphate|ATP]] molecule.]]
The element is found in almost all [[biomolecule]]s that are important to (or generated by) life. Only a few common complex biomolecules, such as [[squalene]] and the [[carotene]]s, contain no oxygen. Of the organic compounds with biological relevance, [[carbohydrate]]s contain the largest proportion by mass of oxygen. All [[fat]]s, [[fatty acid]]s, [[amino acid]]s, and [[protein]]s contain oxygen (due to the presence of [[carbonyl]] groups in these acids and their [[ester]] residues). Oxygen also occurs in [[phosphate]] (PO<sub>4</sub><sup>3&minus;</sup>) groups in the biologically important energy-carrying molecules [[Adenosine triphosphate|ATP]] and [[Adenosine diphosphate|ADP]], in the backbone and the [[purine]]s (except [[adenine]]) and [[pyrimidine]]s of [[RNA]] and [[DNA]], and in bones as [[calcium phosphate]] and [[hydroxylapatite]].

==Precautions==
===Toxicity===
[[Image:Scuba-diving.jpg|thumb|Oxygen toxicity occurs when lungs take in a higher than normal O<sub>2</sub> [[partial pressure]], which can occur in deep scuba diving.]] 
{{main|Oxygen toxicity}}
Oxygen gas ({{chem|O|2}}) can be [[Oxygen toxicity|toxic]] at elevated [[partial pressure]]s, leading to [[convulsion]]s and other health problems.<ref>Since {{chem|O|2}}'s partial pressure is the fraction of {{chem|O|2}} times the total pressure, elevated partial pressures can occur either from high {{chem|O|2}} fraction in breathing gas or from high breathing gas pressure, or a combination of both.</ref><ref name="ECE511">[[#Reference-idCook1968|Cook & Lauer 1968]], p.511</ref> Oxygen toxicity usually begins to occur at partial pressures more than 50&nbsp;kilo[[Pascal (unit)|pascal]]s&nbsp;(kPa), or 2.5 times the normal sea-level {{chem|O|2}} partial pressure of about 21&nbsp;kPa. Therefore, air supplied through [[oxygen mask]]s in medical applications is typically composed of 30% {{chem|O|2}} by volume (about 30&nbsp;kPa at standard pressure).<ref name="NBB299"/> At one time, [[Premature birth|premature babies]] were placed in incubators containing {{chem|O|2}}-rich air, but this practice was discontinued after some babies were blinded by it.<ref name="NBB299"/>

Breathing pure {{chem|O|2}} in space applications, such as in some modern space suits, or in early spacecraft such as Apollo, causes no damage due to the low total pressures used.<ref>
{{cite web | last = Wade | first = Mark | year = 2007 | url = http://www.astronautix.com/craftfam/spasuits.htm | title = Space Suits | publisher = Encyclopedia Astronautica |accessdate=2007-12-16}}</ref> In the case of spacesuits, the {{chem|O|2}} partial pressure in the breathing gas is, in general, about 30&nbsp;kPa (1.4 times normal), and the resulting {{chem|O|2}} partial pressure in the astronaut's arterial blood is only marginally more than normal sea-level {{chem|O|2}} partial pressure (see [[arterial blood gas]]).

Oxygen toxicity to the lungs and [[central nervous system]] can also occur in deep scuba diving and surface supplied diving.<ref name="NBB299"/> Prolonged breathing of an air mixture with an {{chem|O|2}} partial pressure more than 60&nbsp;kPa can eventually lead to permanent   [[pulmonary fibrosis]].<ref name="BMJ">{{cite web|url=http://www.pubmedcentral.nih.gov/articlerender.fcgi?artid=1114047|title=ABC of oxygen: Diving and oxygen|last=Wilmshurst|first=Peter|publisher=British Medical Journal|year=1998|accessdate=2008-01-06}}</ref> Exposure to a {{chem|O|2}} partial pressures greater than 160&nbsp;kPa may lead to convulsions (normally fatal for divers). Acute oxygen toxicity can occur by breathing an air mixture with 21% {{chem|O|2}} at 66&nbsp;m or more of depth while the same thing can occur by breathing 100% {{chem|O|2}} at only 6&nbsp;m.<ref name="BMJ"/><ref name="Donald">[[#Reference-idDonald1992|Donald 1992]]</ref>

===Combustion and other hazards===
<div style="float: left; margin: 5px;">
{{NFPA 704|Health=0|Flammability=0|Reactivity=0|Other=OX}}
</div>
Highly-concentrated sources of oxygen promote rapid [[combustion]]. Fire and explosion hazards exist when concentrated oxidants and fuels are brought into close proximity; however, an ignition event, such as heat or a spark, is needed to trigger combustion.<ref name=astm-tpt>[[#Reference-idWerley1991|Werley 1991]]</ref> Oxygen itself is not the fuel, but the oxidant. Combustion hazards also apply to compounds of oxygen with a high oxidative potential, such as  [[peroxide]]s, [[chlorate]]s, [[nitrate]]s, [[perchlorate]]s, and [[dichromate]]s because they can donate oxygen to a fire.

[[Image:Apollo 1 fire.jpg|thumb|right|Pure {{chem|O|2}} at higher than normal pressure and a spark led to a fire and the loss of the Apollo 1 crew.]]
Concentrated {{chem|O|2}} will allow combustion to proceed rapidly and energetically.<ref name=astm-tpt/> Steel pipes and storage vessels used to store and transmit both gaseous and [[liquid oxygen]] will act as a fuel; and therefore the design and manufacture of {{chem|O|2}} systems requires special training to ensure that ignition sources are minimized.<ref name=astm-tpt/> The fire that killed the Apollo 1 crew on a test launch pad spread so rapidly because the capsule was pressurized with pure {{chem|O|2}} but at slightly more than atmospheric pressure, instead of the ⅓ normal pressure that would be used in a mission.<ref>No single ignition source of the fire was conclusively identified, although some evidence points to arc from an electrical spark). (Report of Apollo 204 Review Board NASA Historical Reference Collection, NASA History Office, NASA HQ, Washington, DC)</ref><ref name=chiles>[[#Reference-idChiles2001|Chiles 2001]]</ref>

Liquid oxygen spills, if allowed to soak into organic matter, such as wood, [[petrochemical]]s, and asphalt can cause these materials to detonate unpredictably on subsequent mechanical impact.<ref name=astm-tpt/> On contact with the human body, it can also cause cryogenic burns to the skin and the eyes.

==See also==
<div class= style="-moz-column-count:2; -webkit-column-count:2; column-count:2;">
*[[Hypoxia (medical)|Hypoxia]], a lack of oxygen
*Hypoxia (environmental) for {{chem|O|2}} depletion in aquatic ecology
*Winkler test for dissolved oxygen for instructions on how to determine the amount of {{chem|O|2}} dissolved in fresh water.
*Optode for a method of measuring {{chem|O|2}} concentration in solution
*Oxygen Catastrophe in geology
*Oxygen isotope ratio cycle
</div>

==Notes and citations== 
<!-- Full reference information for Cook, Daintith, and Emsley given in the references section -->
{{Reflist|2}}

==References==
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{{refend}}

==External links==
* [http://periodic.lanl.gov/elements/8.html Los Alamos National Laboratory &ndash; Oxygen]
* [http://www.webelements.com/webelements/elements/text/O/index.html WebElements.com &ndash; Oxygen]
* [http://www.uigi.com/oxygen.html Oxygen (O2) Properties, Uses, Applications]
* [http://www.organic-chemistry.org/chemicals/oxidations/oxygen.shtm Oxidizing Agents > Oxygen]
* [http://www.americanscientist.org/template/AssetDetail/assetid/29647/page/1 Roald Hoffmann article on "The Story of O"]

{{diatomicelements}}
{{E number infobox 930-949}}
{{Compact periodic table}}


[[Category:Chemical elements]]
[[Category:Oxygen| ]]

[[af:Suurstof]]
[[als:Sauerstoff]]
[[ar:أكسجين]]
[[ast:Oxíxenu]]
[[az:Oksigen]]
[[bn:অক্সিজেন]]
[[zh-min-nan:O (goân-sò͘)]]
[[be:Кісларод]]
[[bar:Sauastoff]]
[[bs:Kiseonik]]
[[br:Oksigen]]
[[bg:Кислород]]
[[ca:Oxigen]]
[[cv:Йӳçлĕк]]
[[cs:Kyslík]]
[[co:Ossigenu]]
[[cy:Ocsigen]]
[[da:Ilt]]
[[de:Sauerstoff]]
[[et:Hapnik]]
[[el:Οξυγόνο]]
[[es:Oxígeno]]
[[eo:Oksigeno]]
[[eu:Oxigeno]]
[[fa:اکسیژن]]
[[fr:Oxygène]]
[[fy:Soerstof]]
[[fur:Ossigjen]]
[[ga:Ocsaigin]]
[[gv:Ocsygien]]
[[gd:Àile-beatha]]
[[gl:Osíxeno (elemento)]]
[[gu:ઑક્સીજન]]
[[ko:산소]]
[[hy:Թթվածին]]
[[hi:आक्सीजन]]
[[hsb:Kislik]]
[[hr:Kisik]]
[[io:Oxo]]
[[id:Oksigen]]
[[ia:Oxygeno]]
[[is:Súrefni]]
[[it:Ossigeno]]
[[he:חמצן]]
[[jv:Oksigen]]
[[pam:Oxygen]]
[[kn:ಆಮ್ಲಜನಕ]]
[[ka:ჟანგბადი]]
[[sw:Oksijeni]]
[[ht:Oksijèn]]
[[ku:Oksîjen]]
[[la:Oxygenium]]
[[lv:Skābeklis]]
[[lb:Sauerstoff]]
[[lt:Deguonis]]
[[li:Zuurstof]]
[[ln:Oksijɛ́ní]]
[[jbo:kijno]]
[[hu:Oxigén]]
[[mk:Кислород]]
[[ml:ഓക്സിജന്‍]]
[[mi:Hāora]]
[[mr:ऑक्सिजन]]
[[ms:Oksigen]]
[[mn:Хүчилтөрөгч]]
[[nah:Ehēcayoh]]
[[nl:Zuurstof]]
[[new:अक्सिजन]]
[[ja:酸素]]
[[no:Oksygen]]
[[nn:Oksygen]]
[[oc:Oxigèn]]
[[om:Oxygen]]
[[uz:Kislorod]]
[[nds:Suerstoff]]
[[pl:Tlen]]
[[pt:Oxigénio]]
[[ksh:Sauerstoff]]
[[ro:Oxigen]]
[[qu:Muksichaq]]
[[ru:Кислород]]
[[sq:Oksigjeni]]
[[scn:Ossìgginu]]
[[si:ඔක්සිජන්]]
[[simple:Oxygen]]
[[sk:Kyslík]]
[[sl:Kisik]]
[[sr:Кисеоник]]
[[sh:Kiseonik]]
[[su:Oksigén]]
[[fi:Happi]]
[[sv:Syre]]
[[tl:Oksiheno]]
[[ta:ஆக்ஸிஜன்]]
[[te:ఆక్సిజన్]]
[[th:ออกซิเจน]]
[[vi:Ôxy]]
[[tg:Оксиген]]
[[tr:Oksijen]]
[[uk:Кисень]]
[[yi:זויערשטאף]]
[[zh-yue:氧]]
[[bat-smg:Degounis]]
[[zh:氧]]

{{WH}}
{{WS}}