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{{About|the chemical element}}
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'''Oxygen''' is a [[chemical element]]; it has the [[chemical symbol|symbol]]{{nbsp}}'''O''' and its [[atomic number]] is{{nbsp}}8. It is a member of the [[chalcogen]] [[group (periodic table)|group]] in the [[periodic table]], a highly [[reactivity (chemistry)|reactive]] [[nonmetal (chemistry)|nonmetal]], and a potent [[oxidizing agent]] that readily forms [[oxide]]s with most elements as well as with other [[chemical compound|compounds]]. Oxygen is [[abundance of elements in Earth's crust|the most abundant element in Earth's crust]], making up almost half of the [[Earth's crust]] in the form of various oxides such as [[water]], [[carbon dioxide]], [[iron oxide]]s and [[silicate]]s.<ref name="Atkins7th">Atkins, P.; Jones, L.; Laverman, L. (2016).''Chemical Principles'', 7th edition. Freeman. {{ISBN|978-1-4641-8395-9}}</ref> It is also [[abundance of chemical elements|the third-most abundant element in the universe]] after [[hydrogen]] and [[helium]].

At [[standard temperature and pressure]], two oxygen atoms will [[chemical bond|bind]] [[covalent bond|covalently]] to form [[dioxygen]], a colorless and odorless [[diatomic]] gas with the [[chemical formula]] {{chem|O|2}}. Dioxygen gas currently constitutes approximately 20.95% [[molar fraction]] of the [[Earth's atmosphere]], though this has [[geological history of oxygen|changed considerably]] over long periods of time in [[Earth's history]]. A much rarer [[triatomic]] [[allotropes of oxygen|allotrope of oxygen]], [[ozone]] ({{chem|O|3}}), strongly absorbs the [[UVB]] and [[ultraviolet C|UVC]] wavelengths and forms a protective [[ozone layer]] at the lower [[stratosphere]], which shields the [[biosphere]] from [[ionizing radiation|ionizing]] [[ultraviolet radiation]]. However, ozone present at the surface is a [[corrosive]] byproduct of [[smog]] and thus an [[air pollutant]].

All [[eukaryotic]] [[organism]]s, including plants, animals, fungi, algae and most [[protist]]s, need oxygen for [[cellular respiration]], a process that extracts [[chemical energy]] by the [[dioxygen in biological reactions|reaction of oxygen]] with [[organic molecule]]s derived from food and releases carbon dioxide as a waste product. 
Many major classes of organic molecules in living organisms contain oxygen atoms, such as [[protein]]s, [[nucleic acid]]s, [[carbohydrate]]s and [[fat]]s, as do the major constituent [[inorganic compound]]s of animal shells, teeth, and bone. Most of the mass of living organisms is oxygen as a component of water, the major constituent of lifeforms. Oxygen in Earth's atmosphere is produced by biotic [[photosynthesis]], in which [[photon energy]] in sunlight is captured by [[chlorophyll]] to split water molecules and then react with carbon dioxide to produce [[carbohydrate]]s and oxygen is released as a [[byproduct]]. Oxygen is too chemically reactive to remain a free element in air without being [[oxygen cycle|continuously replenished]] by the photosynthetic activities of [[autotroph]] such as [[cyanobacteria]], [[chloroplast]]-bearing algae and plants.

Oxygen was isolated by [[Michael Sendivogius]] before 1604, but it is commonly believed that the element was discovered independently by [[Carl Wilhelm Scheele]], in [[Uppsala]], in 1773 or earlier, and [[Joseph Priestley]] in [[Wiltshire]], in 1774. Priority is often given for Priestley because his work was published first. Priestley, however, called oxygen "dephlogisticated air", and did not recognize it as a chemical element. In 1777, [[Antoine Lavoisier]] first recognized oxygen as a chemical element and correctly characterized the role it plays in combustion.

Common industrial uses of oxygen include production of steel, plastics and textiles, [[oxy-fuel welding and cutting|brazing, welding and cutting]] of steels and other metals, [[rocket propellant]], [[oxygen therapy]], and [[life-support system]]s in aircraft, submarines, spaceflight and diving.

==History of study==
The modern concept of the element oxygen developed over five centuries and included many related discoveries and unsuccessful theories. Multiple people made different contributions to the concept: no one person discovered oxygen.<ref name=Severinghaus-2016>{{Cite journal |last=Severinghaus |first=John W. |date=September 2016 |title=Eight sages over five centuries share oxygen's discovery |url=https://journals.physiology.org/doi/full/10.1152/advan.00076.2016 |journal=Advances in Physiology Education |volume=40 |issue=3 |pages=370–376 |doi=10.1152/advan.00076.2016 |pmid=27458241 |issn=1043-4046}}</ref><ref>{{Cite journal |last=Hancock |first=John T. |date=2022-03-15 |title=A Brief History of Oxygen: 250 Years on |journal=Oxygen |language=en |volume=2 |issue=1 |pages=31–39 |doi=10.3390/oxygen2010004 |doi-access=free |issn=2673-9801}}</ref>

===Early experiments===
One of the first known experiments on the relationship between [[combustion]] and air was conducted by the 2nd-century&nbsp;BCE Greek writer on mechanics, [[Philo of Byzantium]]. In his work ''{{tlit|grc|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>{{cite book|title = Story of Human Error|first = Joseph|last = Jastrow|url = https://books.google.com/books?id=tRUO45YfCHwC&pg=PA171|page = 171|date = 1936|publisher = Ayer |isbn = 978-0-8369-0568-7|access-date = August 23, 2020|archive-date = October 1, 2021|archive-url = https://web.archive.org/web/20211001032137/https://books.google.com/books?id=tRUO45YfCHwC&pg=PA171|url-status = live}}</ref> Philo incorrectly surmised that parts of the air in the vessel were converted into the [[Fire (classical element)|classical element fire]] and thus were able to escape through pores in the glass.  Many centuries later [[Ibn al-Nafis]], writing in 1250&nbsp;CE, correctly described oxygenation of blood in the circulatory system; [[Michael Servetus]] rediscovered this concept in 1553 but his books were systematically destroyed.<ref name=Severinghaus-2016/> A scientifically based and influential description was published by [[William Harvey]] in 1628.<ref>{{Cite journal |last=Izquierdo |first=J. J. |date=1937 |title=A New and More Correct Version of the Views of Servetus on the Circulation of the Blood |url=https://www.jstor.org/stable/44438194 |journal=Bulletin of the Institute of the History of Medicine |volume=5 |issue=10 |pages=914–932 |jstor=44438194 |issn=2576-4810}}</ref>

[[Leonardo da Vinci]] observed that a portion of air is consumed during combustion and [[respiration (physiology)|respiration]].<ref name="ECE499">[[#Reference-idCook1968|Cook & Lauer 1968]], p. 499.</ref>
Polish alchemist, philosopher, and physician [[Michael Sendivogius]] (Michał Sędziwój), writing in 1604,<ref>''{{lang|la|De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti}}'' (''Twelve Treatises on the Philosopher's Stone drawn from the source of nature and manual experience''; 1604)</ref> described a substance contained in air, referring to it as {{lang|la|cibus vitae}} ('food of life');<ref name=Severinghaus-2016/> this substance is identical with oxygen.<ref name="Bugaj">{{cite journal |last1=Bugaj |first1=Roman |title=Michał Sędziwój – Traktat o Kamieniu Filozoficznym |journal=Biblioteka Problemów |date=1971 |volume=164 |pages=83–84 |url=https://books.google.com/books?id=d0gaAQAAMAAJ |language=pl |issn=0137-5032 |access-date=August 23, 2020 |archive-date=October 1, 2021 |archive-url=https://web.archive.org/web/20211001032100/https://books.google.com/books?id=d0gaAQAAMAAJ |url-status=live }}</ref> During his experiments, performed between 1598 and 1604, Sendivogius properly recognized that the substance is equivalent to the gaseous byproduct released by the [[thermal decomposition]] of [[potassium nitrate]]. However, this important connection was not understood by contemporary scientists like Robert Boyle.<ref name=Severinghaus-2016/><ref name="Bugaj"/>

Unaware of Sendivogius's work, [[John Mayow]] wrote about a portion of air that provided heat in a fire and the human body. This work was ignored because it failed to align with the prevailing phlogiston theory of air and fire. Mayow observed that [[antimony]] increased in weight when heated, and inferred that the {{lang|la|nitroaereus}} must have combined with it.<ref name="EB1911" /> He also thought that the lungs separate {{lang|la|nitroaereus}} from air and pass it into the blood and that animal heat and muscle movement result from the reaction of {{lang|la|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 ''{{lang|la|Tractatus duo}}'' in the tract "{{langr|la|De respiratione}}".<ref name="WoC" />

After [[Robert Boyle]] proved that air is necessary for combustion in the late 17th&nbsp;century, English chemist [[John Mayow]] (1641–1679) refined this work by showing that fire requires only a part of air that he called {{lang|la|spiritus nitroaereus}}.<ref name="EB1911">{{cite EB1911|wstitle=Mayow, John|volume=17|pages=938–939}}</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">{{cite book|title=World of Chemistry|chapter=John Mayow|date=2005|publisher=Thomson Gale|chapter-url=http://www.bookrags.com/John_Mayow|access-date=December 16, 2007|isbn=978-0-669-32727-4|archive-date=April 17, 2020|archive-url=https://web.archive.org/web/20200417002720/http://www.bookrags.com/John_Mayow/|url-status=live}}</ref> From this, he surmised that {{lang|la|nitroaereus}} is consumed in both respiration and combustion.<ref>{{cite book |last=Lagerkvist |first=Ulf |title=The Enigma of Ferment |publisher=World Scientific |date=2005 |isbn=978-981-256-421-4 |page=58}}</ref>

===Phlogiston theory===
{{Main|Phlogiston theory}}

[[Robert Hooke]], [[Ole Borch]], [[Mikhail Lomonosov]], and [[Pierre Bayen]] all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a [[chemical 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.<ref>{{cite journal | last1 = Best | first1 = Nicholas W. | year = 2015 | title = Lavoisier's 'Reflections on Phlogiston' I: Against Phlogiston Theory | journal = [[Foundations of Chemistry]] | volume = 17 | issue = 2| pages = 137–51 | doi=10.1007/s10698-015-9220-5| s2cid = 170422925 }}</ref>

Established in 1667 by the German alchemist [[J. J. Becher]], and modified by the chemist [[Georg Ernst Stahl]] by 1731,<ref name="morris">{{cite book| title = The last sorcerers: The path from alchemy to the periodic table
| url = https://archive.org/details/lastsorcererspat0000morr
| url-access = registration
| last = Morris| first = Richard| date = 2003|publisher = Joseph Henry Press|location = Washington, D.C.|isbn = 978-0-309-08905-0}}</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" />

Highly combustible materials that leave little [[residue (chemistry)|residue]], such as wood or coal, were thought to be made mostly of phlogiston; 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" />

===Scientific era===
[[File:PriestleyFuseli.jpg|thumb|upright=0.8|left|Among several contemporaries who had made discoveries independently from one another, [[Joseph Priestley]] was the first to publish his findings on oxygen.|alt=A drawing of an elderly man sitting by a table and facing parallel to the drawing. His left arm rests on a notebook, legs crossed.]]

Swedish pharmacist [[Carl Wilhelm Scheele]] produced and described some properties of oxygen sometime around 1770–1775, but did not publish his work until a few years later<ref>{{Cite journal |last=West |first=John B. |date=December 2014 |title=Carl Wilhelm Scheele, the discoverer of oxygen, and a very productive chemist |url=https://journals.physiology.org/doi/full/10.1152/ajplung.00223.2014 |journal=American Journal of Physiology. Lung Cellular and Molecular Physiology |volume=307 |issue=11 |pages=L811–L816 |doi=10.1152/ajplung.00223.2014 |pmid=25281638 |issn=1040-0605|url-access=subscription }}</ref> because he was unable to interpret his work in the framework of the phlogiston theory.<ref name=Severinghaus-2016/> Scheele had produced oxygen gas by heating [[mercuric oxide]] (HgO) and various [[nitrate]]s in 1771–1772.<ref>{{cite web |url=http://www.rsc.org/periodic-table/element/8/oxygen |title=Oxygen |publisher=RSC.org |access-date=2016-12-12 |archive-date=January 28, 2017 |archive-url=https://web.archive.org/web/20170128145051/http://www.rsc.org/periodic-table/element/8/Oxygen |url-status=live }}</ref><ref name="ECE500" /><ref name="ECE499" /> After reading about Priestley's work in 1775, Scheele published in 1777, calling the gas "fire air" because it was then the only known [[Oxidizing agent|agent]] to support combustion.<ref name="NBB300">[[#Reference-idEmsley2001|Emsley 2001]], p. 300</ref>

In the meantime, on August 1, 1774, an experiment conducted by the British clergyman [[Joseph Priestley]] focused sunlight on mercuric oxide contained in 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, Priestley wrote: "The feeling of it to my lungs was not sensibly different from that of [[Atmosphere of Earth|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 name="ECE499" /><ref>{{cite journal|title = An Account of Further Discoveries in Air|first = Joseph |last = Priestley |journal = Philosophical Transactions |date = 1775 |volume = 65 |pages = 384–394 |doi = 10.1098/rstl.1775.0039|doi-access = free }}</ref>

The French chemist [[Antoine Lavoisier]] later claimed to have discovered the new substance independently. Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele had also dispatched a letter to Lavoisier on September 30, 1774, which described his 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" />

=== Discrediting Philogiston theory ===
[[File:Antoine lavoisier.jpg|thumb|upright|left|[[Antoine Lavoisier]] discredited the phlogiston theory.|alt=A drawing of a young man facing towards the viewer, but looking on the side. He wear a white curly wig, dark suit and white scarf.]]

Lavoisier conducted the first adequate quantitative experiments on [[oxidation]] and gave the first correct explanation of how combustion works. 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]].<ref name="ECE500" />

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 ''{{lang|fr|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 {{lang|fr|azote}} (from Greek {{lang|grc|ἄζωτον}} 'lifeless'), which did not support either. {{lang|fr|Azote}} later became ''[[nitrogen]]'' in English, although it has kept the earlier name in French and several other European languages.<ref name="ECE500" />
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====Etymology====
Lavoisier renamed "vital air" to {{tlit|grc|oxygène}} in 1777 from the [[Ancient Greek|Greek]] roots {{tlit|grc|oxys}} ({{lang|grc|ὀξύς}}; "[[acid]]", literally 'sharp', from the taste of acids) and {{tlit|grc|-genēs}} ({{lang|grc|-γενής}}; "producer", literally 'begetter'), because he mistakenly believed that oxygen was a constituent of all acids.<ref name="mellor">{{cite book|last1=Parks|first1=G. D.|last2=Mellor|first2=J. W.|date=1939|title=Mellor's Modern Inorganic Chemistry|edition=6th |publisher=Longmans, Green and Co.|location=London}}</ref> Chemists (such as Sir [[Humphry Davy]] in 1812) eventually determined that Lavoisier was wrong in this regard (e.g. [[Hydrogen chloride]] (HCl) is a strong acid that does not contain oxygen), but by then the name was too well established.<!-- ref defined in Template:Element-symbol--to-oxidation-state-data --><ref name=GE28/>{{rp|793}}

''Oxygen'' entered the English language despite opposition by English scientists and the fact that the Englishman Priestley had first isolated the gas and written about it. 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" />
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===Later history===
[[John Dalton]]'s original [[History of atomic theory#Dalton|atomic hypothesis]] presumed that all elements were monatomic 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, leading to the conclusion that the [[atomic mass]] of oxygen was 8 times that of hydrogen, instead of the modern value of about 16.<ref>{{cite book| title = The Interactive Textbook of PFP96 |chapter= Do We Take Atoms for Granted?|chapter-url=http://www.physics.upenn.edu/courses/gladney/mathphys/subsubsection1_1_3_2.html |url=http://www.physics.upenn.edu/courses/gladney/mathphys/Contents.html |first1=Dennis |last1=DeTurck |last2=Gladney|first2=Larry|last3=Pietrovito|first3=Anthony| publisher=University of Pennsylvania|date=1997|access-date=January 28, 2008|archive-url = https://web.archive.org/web/20080117230939/http://www.physics.upenn.edu/courses/gladney/mathphys/subsubsection1_1_3_2.html |archive-date = January 17, 2008|url-status=dead}}</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 diatomic elemental molecules in those gases.<ref>{{cite book|title=A Treatise on Chemistry|first1=Henry Enfield |last1=Roscoe |last2=Schorlemmer|first2=Carl|page=38|date=1883|publisher=D. Appleton and Co.}}</ref>

In 1879 the French brothers Quentin and Arthur
Brin discovered a [[Brin process|commercially viable reaction]] to create oxygen. They realized that the known reversible reaction
{{chem2|2BaO}}(s) + {{chem2|O2}}(g) ↔ {{chem2|2BaO2}}(s) was deactivated by the formation of barium carbonate from carbon dioxide in the air; treating air to remove the carbon dioxide allowed the reaction be reversed indefinitely. Their company used the process between 1886 and 1906 when more economical [[fractional distillation]] began to be used.<ref>{{Cite journal |last=Jensen |first=William B. |date=November 2009 |title=The Origin of the Brin Process for the Manufacture of Oxygen |url=https://pubs.acs.org/doi/abs/10.1021/ed086p1266 |journal=Journal of Chemical Education |language=en |volume=86 |issue=11 |pages=1266 |doi=10.1021/ed086p1266 |bibcode=2009JChEd..86.1266J |issn=0021-9584|url-access=subscription }}</ref>

[[File:Goddard and Rocket.jpg|thumb|left|upright|[[Robert H. Goddard]] and a liquid-oxygen–gasoline [[rocket]]|alt=A metal frame structure stands on the snow near a tree. A middle-aged man wearing a coat, boots, leather gloves and a cap stands by the structure and holds it with his right hand.]]

By the late 19th century scientists realized that air could be liquefied and its components isolated by compressing and cooling it. Using a [[Cascade (chemical engineering)|cascade]] method, Swiss chemist and physicist [[Raoul Pictet|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">{{cite book|title=Biographical Encyclopedia of Scientists|last=Daintith|first=John|date=1994|publisher=CRC Press|isbn=978-0-7503-0287-6|page=707}}</ref> 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 each case and no meaningful analysis could be conducted. Oxygen was liquefied in a stable state for the first time on March 29, 1883, by Polish scientists from [[Jagiellonian University]], [[Zygmunt Wróblewski]] and [[Karol Olszewski]].<ref>{{cite journal|title = Louis Paul Cailletet: The liquefaction of oxygen and the emergence of low-temperature research |first =Faidra |last = Papanelopoulou |journal =Notes and Records of the Royal Society of London |date = 2013|volume = 67 |issue=4|pages = 355–73|doi = 10.1098/rsnr.2013.0047 |pmc=3826198}}</ref>

In 1891 Scottish chemist [[James Dewar]] was able to produce enough liquid oxygen for 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 separately.<ref name="HPAM">{{cite book|title=How Products are Made|chapter=Oxygen|publisher=The Gale Group, Inc.|date=2002|chapter-url=http://www.answers.com/topic/oxygen|access-date=December 16, 2007|archive-date=April 3, 2019|archive-url=https://web.archive.org/web/20190403220006/http://www.answers.com/topic/oxygen|url-status=live}}</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]] that burned liquid fuel; 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]], US.<ref name="HPAM" /><ref>{{cite web|title=Goddard-1926 |url=http://grin.hq.nasa.gov/ABSTRACTS/GPN-2002-000132.html |publisher=NASA |access-date=November 18, 2007 |url-status=dead |archive-url=https://web.archive.org/web/20071108225824/http://grin.hq.nasa.gov/ABSTRACTS/GPN-2002-000132.html |archive-date=November 8, 2007 }}</ref>
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==Characteristics==
===Properties and molecular structure===
[[File:Oxygen molecule orbitals diagram-en.svg|thumb|left|upright=1.2|Orbital diagram, after Barrett (2002),<ref name="Barrett2002" /> showing the participating atomic orbitals from each oxygen atom, the molecular orbitals that result from their overlap, and the [[Aufbau principle|aufbau]] filling of the orbitals with the 12 electrons, 6 from each O atom, beginning from the lowest-energy orbitals, and resulting in covalent double-bond character from filled orbitals (and cancellation of the contributions of the pairs of σ and σ<sup>*</sup> and π and π<sup>*</sup> orbital pairs).]]
At [[standard temperature and pressure]], oxygen is a colorless, odorless, and tasteless gas with the [[molecular formula]] {{chem|O|2}}, referred to as dioxygen.<ref>{{cite web |url=http://www.sciencekids.co.nz/sciencefacts/chemistry/oxygen.html |title=Oxygen Facts |publisher=Science Kids |date=February 6, 2015 |access-date=November 14, 2015 |archive-date=May 7, 2020 |archive-url=https://web.archive.org/web/20200507223541/https://www.sciencekids.co.nz/sciencefacts/chemistry/oxygen.html |url-status=live}}</ref>

As ''dioxygen'', two oxygen atoms are [[chemical bond|chemically bound]] to each other. The bond can be variously described based on level of theory, but is reasonably and simply described as a covalent [[double bond]] that results from the filling of [[molecular orbitals]] formed from the [[atomic orbital]]s of the individual oxygen atoms, the filling of which results in a [[bond order]] of two. More specifically, the double bond is the result of sequential, low-to-high energy, or [[Aufbau principle|Aufbau]], filling of orbitals, and the resulting cancellation of contributions from the 2s electrons, after sequential filling of the low σ and σ<sup>*</sup> orbitals; σ overlap of the two atomic 2p orbitals that lie along the O–O molecular axis and π overlap of two pairs of atomic 2p orbitals perpendicular to the O–O molecular axis, and then cancellation of contributions from the remaining two 2p electrons after their partial filling of the π<sup>*</sup> orbitals.<ref name="Barrett2002">Jack Barrett, 2002, "Atomic Structure and Periodicity", (Basic concepts in chemistry, Vol.&nbsp;9 of Tutorial chemistry texts), Cambridge, UK: Royal Society of Chemistry, p.&nbsp;153, {{ISBN|0854046577}}. See [https://books.google.com/books?isbn=0854046577 Google Books]. {{Webarchive|url=https://web.archive.org/web/20200530044101/https://books.google.com/books?isbn=0854046577%2F |date=May 30, 2020 }} accessed January 31, 2015.</ref>

This combination of cancellations and σ and π overlaps results in dioxygen's double-bond character and reactivity, and a triplet electronic [[ground state]]. An [[electron configuration]] with two unpaired electrons, as is found in dioxygen orbitals (see the filled π* orbitals in the diagram) that are of equal energy—i.e., [[degenerate orbitals|degenerate]]—is a configuration termed a [[spin triplet]] state. Hence, the ground state of the {{chem|O|2}} molecule is referred to as [[triplet oxygen]].<ref name="BiochemOnline">{{cite web |work=Biochemistry Online |url=http://employees.csbsju.edu/hjakubowski/classes/ch331/oxphos/oldioxygenchem.html |title=Chapter 8: Oxidation-Phosphorylation, the Chemistry of Di-Oxygen |first=Henry |last=Jakubowski |access-date=January 28, 2008 |publisher=Saint John's University |archive-date=October 5, 2018 |archive-url=https://web.archive.org/web/20181005032115/http://employees.csbsju.edu/hjakubowski/classes/ch331/oxphos/oldioxygenchem.html |url-status=live}}</ref> The highest-energy, partially filled orbitals are [[antibonding]], and so their filling weakens the bond order from three to two. Because of its unpaired electrons, triplet oxygen reacts only slowly with most organic molecules, which have paired electron spins; this prevents spontaneous combustion.<ref name="astm-tpt">{{cite conference|editor1-last=Werley|editor1-first=Barry L.|date=1991|work=Fire Hazards in Oxygen Systems|title=ASTM Technical Professional training|publisher=[[ASTM International]] Subcommittee G-4.05|location=Philadelphia}}</ref>

[[File:Liquid oxygen in a magnet 2.jpg|thumb|left|upright|Liquid oxygen, temporarily suspended in a magnet owing to its paramagnetism]]
In the triplet form, {{chem|O|2}} molecules are [[paramagnetism|paramagnetic]]. That is, they impart magnetic character to oxygen when it is in the presence of a magnetic field, 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 so [[magnet]]ic 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 |access-date = December 15, 2007 |archive-url = https://web.archive.org/web/20071217064218/http://genchem.chem.wisc.edu/demonstrations/Gen_Chem_Pages/0809bondingpage/liquid_oxygen.htm |archive-date = December 17, 2007 |url-status=dead}}</ref> Oxygen's paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine gaseous oxygen concentration, especially in industrial process control and medicine.<ref>{{Cite journal |last1=Jasek |first1=Krzysztof |last2=Pasternak |first2=Mateusz |last3=Grabka |first3=Michał |date=2022-11-25 |title=Paramagnetic Sensors for the Determination of Oxygen Concentration in Gas Mixtures |url=https://doi.org/10.1021/acssensors.2c00938 |journal=ACS Sensors |volume=7 |issue=11 |pages=3228–3242 |doi=10.1021/acssensors.2c00938 |pmc=9706814 |pmid=36300934 |bibcode=2022ACSSe...7.3228J }}</ref><ref>{{cite web |url=http://www.servomex.com/oxygen_gas_analyser.html |title=Company literature of Oxygen analyzers (triplet) |publisher=Servomex |access-date=December 15, 2007 |url-status=dead |archive-url=https://web.archive.org/web/20080308213517/http://www.servomex.com/oxygen_gas_analyser.html |archive-date=March 8, 2008 }}</ref>

[[Singlet oxygen]] is a name given to several higher-energy species of molecular {{chem|O|2}} in which all the electron spins are paired. It is much more reactive with common [[organic compound|organic molecules]] than is normal (triplet) molecular oxygen. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.<ref>{{cite journal |first=Anja |last=Krieger-Liszkay |journal=Journal of Experimental Botany |volume=56 |pages=337–346 |date=October 13, 2004 |title=Singlet oxygen production in photosynthesis |doi=10.1093/jxb/erh237 |pmid=15310815 |issue=411 |doi-access=free}}</ref> It is also produced in the [[troposphere]] by the photolysis of ozone by light of short wavelength<ref name="harrison">{{cite book |last=Harrison |first=Roy M. |author-link=Roy M. Harrison |date=1990 |title=Pollution: Causes, Effects & Control |edition=2nd |location=Cambridge |publisher=[[Royal Society of Chemistry]] |isbn=978-0-85186-283-5 |url-access=registration |url=https://archive.org/details/pollutioncausese0000unse}}</ref> and by the [[immune system]] as a source of active oxygen.<ref name="immune-ozone">{{cite journal |journal=Science |title=Evidence for Antibody-Catalyzed Ozone Formation in Bacterial Killing and Inflammation |date=December 13, 2002 |volume=298 |pages=2195–2219 |doi=10.1126/science.1077642 |pmid=12434011 |last1=Wentworth |first1=Paul |last2=McDunn |first2=J. E. |last3=Wentworth |first3=A. D. |last4=Takeuchi |first4=C. |last5=Nieva |first5=J. |last6=Jones |first6=T. |last7=Bautista |first7=C. |last8=Ruedi |first8=J. M. |last9=Gutierrez |first9=A. |last10=Janda |first10=K. D. |last11=Babior |first11=B. M. |last12=Eschenmoser |first12=A. |last13=Lerner |first13=R. A. |issue=5601 |bibcode=2002Sci...298.2195W |s2cid=36537588 |doi-access=free }}</ref> [[Carotenoid]]s in photosynthetic organisms (and possibly 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>{{cite journal |title=Singlet oxygen quenching ability of naturally occurring carotenoids |journal=Lipids |first1=Osamu |last1=Hirayama |last2=Nakamura |first2=Kyoko |last3=Hamada |first3=Syoko |last4=Kobayasi |first4=Yoko |volume=29 |issue=2 |date=1994 |doi=10.1007/BF02537155 |pages=149–150 |pmid=8152349 |bibcode=1994Lipid..29..149H |s2cid=3965039}}</ref>

===Allotropes===
{{Main|Allotropes of oxygen}}
[[File:Oxygen molecule.png|thumb|right|upright=0.9|[[Space-filling model]] representation of dioxygen (O<sub>2</sub>) molecule]]
The common [[Allotropy|allotrope]] of elemental oxygen on Earth is called [[Allotropes of oxygen|dioxygen]], {{chem|O|2}}, the major part of the Earth's atmospheric oxygen (see [[#Occurrence|Occurrence]]). O<sub>2</sub> has a bond length of 121&nbsp;[[Picometre|pm]] and a bond energy of 498&nbsp;[[joule per mole|kJ/mol]].<ref>{{cite web|last=Chieh|first=Chung|title=Bond Lengths and Energies|url=http://www.science.uwaterloo.ca/~cchieh/cact/c120/bondel.html|publisher=University of Waterloo|access-date=December 16, 2007|archive-url=https://web.archive.org/web/20071214215455/http://www.science.uwaterloo.ca/~cchieh/cact/c120/bondel.html|archive-date=December 14, 2007|url-status=dead}}</ref> 
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">{{cite book|title=Guide to the Elements|url=https://archive.org/details/guidetoelements00stwe|url-access=registration|edition=Revised |first=Albert|last=Stwertka|publisher=Oxford University Press|date=1998|isbn=978-0-19-508083-4|pages=[https://archive.org/details/guidetoelements00stwe/page/48 48–49]}}</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 [[Electromagnetic spectrum|spectrum]], the [[ozone layer]] of the upper atmosphere functions as a protective radiation shield for the planet.<ref name="mellor" /> Near the Earth's surface, it is a [[air pollution|pollutant]] formed as a by-product of [[exhaust system|automobile exhaust]].<ref name="GuideElem48" /> At [[low earth orbit]] altitudes, sufficient atomic oxygen is present to cause [[corrosion in space|corrosion of spacecraft]].<ref>{{cite web|access-date=August 8, 2009|url=http://www.spenvis.oma.be/spenvis/help/background/atmosphere/erosion.html|title=Atomic oxygen erosion|archive-url = https://web.archive.org/web/20070613121048/http://www.spenvis.oma.be/spenvis/help/background/atmosphere/erosion.html |archive-date = June 13, 2007|url-status=dead}}</ref>

The [[Metastability in molecules|metastable]] molecule [[tetraoxygen]] ({{chem|O|4}}) was discovered in 2001,<ref name="o4">{{cite journal|last1=Cacace|first1=Fulvio|last2=de Petris|first2=Giulia|last3=Troiani|first3=Anna |date=2001|title=Experimental Detection of Tetraoxygen|journal=Angewandte Chemie International Edition|volume=40|issue=21|pages=4062–65|doi = 10.1002/1521-3773(20011105)40:21<4062::AID-ANIE4062>3.0.CO;2-X|pmid=12404493}}</ref><ref name="newform">{{cite news|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=September 16, 2001|access-date=January 9, 2008|archive-date=October 21, 2013|archive-url=https://web.archive.org/web/20131021083801/http://www.nature.com/news/2001/011122/pf/011122-3_pf.html|url-status=live}}</ref> and was assumed to exist in one of the six phases of [[solid oxygen]]. In 2006 this phase, created by pressurizing {{chem|O|2}} to 20&nbsp;[[Pascal (unit)|GPa]], was shown to form a [[rhombohedral]] {{chem|O|8}} [[Cluster chemistry|cluster]].<ref>{{cite journal| title=Observation of an{{chem|O|8}} molecular lattice in the phase of solid oxygen|journal=Nature|volume=443|issue=7108|pages=201–04|doi=10.1038/nature05174|first1=Lars F. |last1=Lundegaard|pmid=16971946| display-authors=4| last2=Weck|first2=Gunnar|last3=McMahon|first3=Malcolm I.|last4=Desgreniers|first4=Serge|last5=Loubeyre|first5=Paul|date=2006|bibcode = 2006Natur.443..201L|s2cid=4384225}}</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 propellant|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>{{cite journal|last1=Desgreniers |first1=S. |last2=Vohra|first2=Y. K.|last3=Ruoff|first3=A. L.|title=Optical response of very high density solid oxygen to 132 GPa|journal=J. Phys. Chem.|volume=94|pages=1117–22|date=1990|doi=10.1021/j100366a020|issue=3 |bibcode=1990JPhCh..94.1117D }}</ref> and it was shown in 1998 that at very low temperatures, this phase becomes [[superconductivity|superconducting]].<ref>{{cite journal|last1=Shimizu|first1=K.|display-authors=4|last2=Suhara|first2=K.|last3=Ikumo|first3=M.|last4=Eremets|first4=M. I.|last5= Amaya|first5=K.|title=Superconductivity in oxygen|journal=Nature|volume=393|pages=767–69|date=1998|doi=10.1038/31656|issue=6687|bibcode = 1998Natur.393..767S |s2cid=205001394|author4-link=Mikhail Eremets}}</ref>

===Physical properties===
[[File:Liquid oxygen in a beaker 4.jpg|thumb|Liquid oxygen boiling (O<sub>2</sub>)|alt=A transparent beaker containing a light blue fluid with gas bubbles.]]
{{see also|Liquid oxygen|solid oxygen}}
Oxygen [[Solubility|dissolves]] more readily in water than nitrogen does. Water in equilibrium with air contains approximately 1 molecule of dissolved {{chem|O|2}} for every 2 molecules of {{chem|N|2}} (1:2), compared with an atmospheric ratio of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much ({{val|14.6|u=mg/L}}) dissolves at {{Convert|0|°C|abbr=on}} than at {{Convert|20|°C|abbr=on}} ({{val|7.6|u=mg/L}}).<ref name="NBB299" /><ref>{{cite web |url=http://www.engineeringtoolbox.com/air-solubility-water-d_639.html |title=Air solubility in water |access-date=December 21, 2007 |publisher=The Engineering Toolbox |archive-date=April 4, 2019 |archive-url=https://web.archive.org/web/20190404044017/https://www.engineeringtoolbox.com/air-solubility-water-d_639.html |url-status=live}}</ref>

At {{Convert|25|°C|abbr=on}} and {{convert|1|atm|lk=on|sigfig=6}} of air, freshwater can dissolve about 6.04&nbsp;[[Litre|milliliters]]&nbsp;(mL) of oxygen per [[liter]], and [[seawater]] contains about 4.95&nbsp;mL per liter. 
At {{Convert|5|°C|abbr=on}} the solubility increases to 9.0&nbsp;mL (50% more than at {{Convert|25|°C|abbr=on}}) per liter for freshwater and 7.2&nbsp;mL (45% more) per liter for sea water.<ref>{{cite book |title = The Physiology of Fishes |url=https://books.google.com/books?id=lBltoKDaBVEC&pg=PA88 |first1=David Hudson |last1=Evans |last2=Claiborne |first2=James B. |page=88 |date=2005 |edition=3rd |publisher=CRC Press |isbn=978-0-8493-2022-4}}</ref>
{| class="wikitable" style="float:left; margin-right:2em"
|+Oxygen gas dissolved in water at sea-level<br />(milliliters per liter)
!
!{{Convert|5|°C|abbr=on}}
!{{Convert|25|°C|abbr=on}}
|-
|Freshwater
|9.00
|6.04
|-
|Seawater
|7.20
|4.95
|}
Oxygen condenses at 90.20&nbsp;[[kelvin|K]] (−182.95&nbsp;°C, −297.31&nbsp;°F) and freezes at 54.36&nbsp;K (−218.79&nbsp;°C, −361.82&nbsp;°F).<ref>{{cite book |first=David R. |last=Lide |title=CRC Handbook of Chemistry and Physics |edition=84th |publisher=[[CRC Press]] |location=Boca Raton, Florida |date=2003 |chapter=Section 4, Properties of the Elements and Inorganic Compounds; Melting, boiling, and critical temperatures of the elements |isbn=978-0-8493-0595-5 |url=https://archive.org/details/crchandbookofche0000unse_p1y5}}</ref> Both [[liquid oxygen|liquid]] and [[solid oxygen|solid]] {{chem|O|2}} are clear substances with a light [[diffuse sky radiation|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. |access-date = December 15, 2007 |archive-date = October 21, 2018 |archive-url = https://web.archive.org/web/20181021010346/http://www.uigi.com/cryodist.html |url-status = live}}</ref> Liquid oxygen may also be condensed from air using [[liquid nitrogen]] as a coolant.<ref name="LOX MSDS">{{cite web |url=https://www.mathesontrigas.com/pdfs/msds/00225011.pdf |title=Liquid Oxygen Material Safety Data Sheet |publisher=Matheson Tri Gas |access-date=December 15, 2007 |url-status=dead |archive-url=https://web.archive.org/web/20080227014309/https://www.mathesontrigas.com/pdfs/msds/00225011.pdf |archive-date=February 27, 2008 }}</ref> Liquid oxygen is a highly reactive substance and must be segregated from combustible materials.<ref name="LOX MSDS" />

The spectroscopy of molecular oxygen is associated with the atmospheric processes of [[aurora]] and [[airglow]].<ref name="Krupenie1972">{{cite journal |last1=Krupenie |first1=Paul H. |title=The Spectrum of Molecular Oxygen |journal=Journal of Physical and Chemical Reference Data |volume=1 |issue=2 |year=1972 |pages=423–534 |doi=10.1063/1.3253101 |bibcode=1972JPCRD...1..423K |s2cid=96242703 }}</ref> The absorption in the [[Herzberg continuum]] and [[Schumann–Runge bands]] in the ultraviolet produces atomic oxygen that is important in the chemistry of the middle atmosphere.<ref name="BrasseurSolomon2006">{{cite book |author1=Guy P. Brasseur |author2=Susan Solomon |title=Aeronomy of the Middle Atmosphere: Chemistry and Physics of the Stratosphere and Mesosphere |url=https://books.google.com/books?id=Z5OtlDjfXkkC&pg=PA220 |date=January 15, 2006 |publisher=Springer Science & Business Media |isbn=978-1-4020-3824-2 |pages=220– |access-date=July 2, 2015 |archive-date=February 2, 2017 |archive-url=https://web.archive.org/web/20170202143926/https://books.google.com/books?id=Z5OtlDjfXkkC&pg=PA220 |url-status=live}}</ref> Excited-state singlet molecular oxygen is responsible for red chemiluminescence in solution.<ref name="Kearns1971">{{cite journal |last1=Kearns |first1=David R. |title=Physical and chemical properties of singlet molecular oxygen |journal=Chemical Reviews |volume=71 |issue=4 |year=1971 |pages=395–427 |doi=10.1021/cr60272a004}}</ref>

Table of thermal and physical properties of oxygen (O<sub>2</sub>) at atmospheric pressure:<ref>{{Cite book |last=Holman |first=Jack P. |title=Heat transfer |publisher=McGraw-Hill Companies, Inc. |year=2002 |isbn=9780072406559 |edition=9th |location=New York, NY |pages=600–606 |language=English |oclc=46959719}}</ref><ref>{{Cite book |last=Incropera 1 Dewitt 2 Bergman 3 Lavigne 4 |first=Frank P. 1 David P. 2 Theodore L. 3 Adrienne S. 4 |title=Fundamentals of heat and mass transfer. |publisher=John Wiley and Sons, Inc. |year=2007 |isbn=9780471457282 |edition=6th |location=Hoboken, NJ |pages=941–950 |language=English |oclc=62532755}}</ref>
{|class="wikitable mw-collapsible mw-collapsed"
|[[Temperature]] (K)
|[[Density]] (kg/m<sup>3</sup>)
|[[Specific heat]] (kJ/(kg·K))
|[[Dynamic viscosity]] (kg/(m·s))
|[[Kinematic viscosity]] (m<sup>2</sup>/s)
|[[Thermal conductivity]] (W/(m·K))
|[[Thermal diffusivity]] (m<sup>2</sup>/s)
|[[Prandtl Number]]
|-
|100
|3.945
|0.962
|7.64E-06
|1.94E-06
|0.00925
|2.44E-06
|0.796
|-
|150
|2.585
|0.921
|1.15E-05
|4.44E-06
|0.0138
|5.80E-06
|0.766
|-
|200
|1.93
|0.915
|1.48E-05
|7.64E-06
|0.0183
|1.04E-05
|0.737
|-
|250
|1.542
|0.915
|1.79E-05
|1.16E-05
|0.0226
|1.60E-05
|0.723
|-
|300
|1.284
|0.92
|2.07E-05
|1.61E-05
|0.0268
|2.27E-05
|0.711
|-
|350
|1.1
|0.929
|2.34E-05
|2.12E-05
|0.0296
|2.90E-05
|0.733
|-
|400
|0.962
|1.0408
|2.58E-05
|2.68E-05
|0.033
|3.64E-05
|0.737
|-
|450
|0.8554
|0.956
|2.81E-05
|3.29E-05
|0.0363
|4.44E-05
|0.741
|-
|500
|0.7698
|0.972
|3.03E-05
|3.94E-05
|0.0412
|5.51E-05
|0.716
|-
|550
|0.6998
|0.988
|3.24E-05
|4.63E-05
|0.0441
|6.38E-05
|0.726
|-
|600
|0.6414
|1.003
|3.44E-05
|5.36E-05
|0.0473
|7.35E-05
|0.729
|-
|700
|0.5498
|1.031
|3.81E-05
|6.93E-05
|0.0528
|9.31E-05
|0.744
|-
|800
|0.481
|1.054
|4.15E-05
|8.63E-05
|0.0589
|1.16E-04
|0.743
|-
|900
|0.4275
|1.074
|4.47E-05
|1.05E-04
|0.0649
|1.41E-04
|0.74
|-
|1000
|0.3848
|1.09
|4.77E-05
|1.24E-04
|0.071
|1.69E-04
|0.733
|-
|1100
|0.3498
|1.103
|5.06E-05
|1.45E-04
|0.0758
|1.96E-04
|0.736
|-
|1200
|0.3206
|1.0408
|5.33E-05
|1.661E-04
|0.0819
|2.29E-04
|0.725
|-
|1300
|0.296
|1.125
|5.88E-05
|1.99E-04
|0.0871
|2.62E-04
|0.721
|}

===Isotopes and stellar origin===
<!--COPYEDITS AND CORRECTIONS ONLY: DIRECT EXPANSION OF THIS SUBTOPIC TO [[Isotopes of oxygen]] -->
{{Main|Isotopes of oxygen}}
[[File:Evolved star fusion shells.svg|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 <sup>18</sup>O in the He-shell.|alt=A concentric-sphere diagram, showing, from the core to the outer shell, iron, silicon, oxygen, neon, carbon, helium and hydrogen layers.]]
Naturally occurring oxygen is composed of three stable [[isotope]]s, [[oxygen-16|<sup>16</sup>O]], [[oxygen-17|<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|access-date=December 17, 2007|archive-date=July 12, 2012|archive-url=https://archive.today/20120712195516/http://environmentalchemistry.com/yogi/periodic/O-pg2.html|url-status=live}}</ref>

<sup>16</sup>O is the one of the dominant fusion products in massive [[star]]s. It is  [[nucleosynthesis|synthesized]] at the end of the [[helium fusion]] process with some synthesis in the [[neon burning process]].<ref name="Meyer2005">{{cite conference|first=B. S.|last=Meyer|title=Nucleosynthesis and Galactic Chemical Evolution of the Isotopes of Oxygen|conference=Workgroup on Oxygen in the Earliest Solar System|date=September 19–21, 2005|location=Gatlinburg, Tennessee|url=http://www.lpi.usra.edu/meetings/ess2005/pdf/9022.pdf|access-date=January 22, 2007|work=Proceedings of the NASA Cosmochemistry Program and the Lunar and Planetary Institute|conference-url=http://www.lpi.usra.edu/meetings/ess2005/|id=9022|archive-date=December 29, 2010|archive-url=https://web.archive.org/web/20101229194925/http://www.lpi.usra.edu/meetings/ess2005/pdf/9022.pdf|url-status=live}}</ref> Both <sup>17</sup>O and <sup>18</sup>O require seed nuclei. <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 [[Nitrogen-14|<sup>14</sup>N]] (made abundant from CNO burning) captures a [[Helium-4|<sup>4</sup>He]] nucleus, making <sup>18</sup>O common in the helium-rich zones of [[Stellar evolution#Massive stars|evolved, massive stars]].<ref name="Meyer2005" />

Fifteen [[radioisotope]]s have been characterized, ranging from <sup>11</sup>O to <sup>28</sup>O.{{NUBASE2020|ref}}<ref name=O-28-SA>{{cite news |url=https://www.sciencealert.com/scientists-have-observed-a-never-before-seen-form-of-oxygen |first=Michelle |last=Starr |date=30 August 2023 |title=Scientists Have Observed A Never-Before-Seen Form of Oxygen |work=ScienceAlert |access-date=30 August 2023}}</ref> The most stable are <sup>15</sup>O with a [[half-life]] of 122.24&nbsp;seconds and <sup>14</sup>O with a half-life of 70.606&nbsp;seconds.<ref name="EnvChem-Iso" /> All of the remaining [[Radioactive decay|radioactive]] isotopes have half-lives that are less than 27&nbsp;seconds 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 [[positron emission|β<sup>+</sup> decay]]<ref name="NUDAT-13O">{{cite web|url=http://www.nndc.bnl.gov/nudat2/decaysearchdirect.jsp?nuc=13O&unc=nds|title=NUDAT 13O|access-date=July 6, 2009|archive-date=June 9, 2022|archive-url=https://web.archive.org/web/20220609000104/http://www.nndc.bnl.gov/nudat2/decaysearchdirect.jsp?nuc=13O|url-status=live}}</ref><ref name="NUDAT-14O">{{cite web|url=http://www.nndc.bnl.gov/nudat2/decaysearchdirect.jsp?nuc=14O&unc=nds|title=NUDAT 14O|access-date=July 6, 2009|archive-date=June 7, 2022|archive-url=https://web.archive.org/web/20220607045357/http://www.nndc.bnl.gov/nudat2/decaysearchdirect.jsp?nuc=14O|url-status=live}}</ref><ref name="NUDAT-15O">{{cite web|url=http://www.nndc.bnl.gov/nudat2/decaysearchdirect.jsp?nuc=15O&unc=nds|title=NUDAT 15O|access-date=July 6, 2009|archive-date=June 7, 2022|archive-url=https://web.archive.org/web/20220607045434/http://www.nndc.bnl.gov/nudat2/decaysearchdirect.jsp?nuc=15O|url-status=live}}</ref> 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===
{{see also|Silicate minerals|Category:Oxide minerals|Stellar population|Cosmochemistry|Astrochemistry}}
{| class="wikitable sortable" style="float:left; margin-right: 20px"
|+Ten most common elements in the [[Milky Way Galaxy]] estimated spectroscopically<ref name="croswell">{{cite book | last = Croswell | first = Ken | title = Alchemy of the Heavens | publisher = Anchor | year = 1996 | url = http://kencroswell.com/alchemy.html | isbn = 978-0-385-47214-2 | access-date = December 2, 2011 | archive-date = May 13, 2011 | archive-url = https://web.archive.org/web/20110513233910/http://www.kencroswell.com/alchemy.html | url-status = live }}</ref>
|-
![[Atomic number|Z]] !! Element !! colspan="2"|Mass fraction in parts per million
|-
| 1 || [[Hydrogen]] || align="right"|{{bartable| 739,000||0.001}}
|-
| 2 || [[Helium]] || align="right"|{{bartable| 240,000||0.001}}
|-
| 8 || Oxygen || align="right"|{{bartable| 10,400||0.001||background:red;}}
|-
| 6 || [[Carbon]] || align="right"|{{bartable| 4,600||0.001}}
|-
| 10 || [[Neon]] || align="right"|{{bartable| 1,340||0.001}}
|-
| 26 || [[Iron]] || align="right"|{{bartable| 1,090||0.001}}
|-
| 7 || [[Nitrogen]] || align="right"|{{bartable| 960||0.001}}
|-
| 14 || [[Silicon]] || align="right"|{{bartable| 650||0.001}}
|-
| 12 || [[Magnesium]] || align="right"|{{bartable| 580||0.001}}
|-
| 16 || [[Sulfur]] || align="right"|{{bartable| 440||0.001}}
|}

Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.<ref name="NBB297">[[#Reference-idEmsley2001|Emsley 2001]], p. 297</ref> 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">{{cite web |url=http://periodic.lanl.gov/elements/8.html|publisher=Los Alamos National Laboratory|title=Oxygen|access-date=December 16, 2007|archive-url=https://web.archive.org/web/20071026034224/http://periodic.lanl.gov/elements/8.html|archive-date=October 26, 2007}}</ref> as part of oxide compounds such as [[silicon dioxide]] and is the most abundant element by mass in the [[crust (geology)#Earth's crust and mantle|Earth's crust]]. It is also the major component of the world's oceans (88.8% by mass).<ref name="ECE500" /> Oxygen gas is the second most common component of the [[Earth's atmosphere]], taking up 20.8% of its volume and 23.1% of its mass (some 10<sup>15</sup> tonnes).<ref name="ECE500" /><ref name="NBB298">[[#Reference-idEmsley2001|Emsley 2001]], p. 298</ref><ref group="lower-alpha">Figures given are for values up to {{convert|80|km|mi|abbr=on}} 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 much less. The {{chem|O|2}} surrounding those planets is produced solely by the action of ultraviolet radiation on oxygen-containing molecules such as carbon dioxide.<ref>{{cite book |author1=Richard Peer Wayne |title=Chemistry of Atmospheres |date=2006 |publisher=Oxford University Press |isbn=9780198503750 |language=en |pages=562–584}}</ref>
The unusually high concentration of oxygen gas 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. Photosynthesis releases oxygen into the atmosphere, while [[Cellular respiration|respiration]], [[Decomposition|decay]], and combustion remove it from the atmosphere. In the present equilibrium, production and consumption occur at the same rate.<ref name=GE28/>{{rp|602}} 

Oxygen levels in the atmosphere are trending slightly downward globally, possibly because of fossil-fuel burning.<ref>{{cite web|url=http://scrippso2.ucsd.edu/|title=Atmospheric Oxygen Research|author=Scripps Institute|access-date=October 8, 2011|archive-date=July 25, 2017|archive-url=https://web.archive.org/web/20170725074925/http://scrippso2.ucsd.edu/|url-status=live}}</ref>

[[File:WOA09 sea-surf O2 AYool.png|thumb|right|Cold water holds more dissolved {{chem|O|2}}.|alt=World map showing that the sea-surface oxygen is depleted around the equator and increases towards the poles.]]
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 notes that according to later articles in ''Nature'', the values appear to be about 3% too high.</ref> 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>
Significant deoxygenation has been observed in tropical oceans. Warming oceans waters are expected to lose oxygen over the next century and into the future for a thousand years; the possible consequences include minimal oxygen zones which are unable to support [[macrofauna]].<ref>{{Cite journal |last1=Keeling |first1=Ralph F. |last2=Körtzinger |first2=Arne |last3=Gruber |first3=Nicolas |date=2010-01-15 |title=Ocean Deoxygenation in a Warming World |url=https://www.annualreviews.org/content/journals/10.1146/annurev.marine.010908.163855 |journal=Annual Review of Marine Science |language=en |volume=2 |pages=199–229 |doi=10.1146/annurev.marine.010908.163855 |pmid=21141663 |bibcode=2010ARMS....2..199K |issn=1941-1405|url-access=subscription }}</ref>

===Analysis===
[[File:Phanerozoic Climate Change.png|thumb|left|upright=1.15|500 million years of [[Climate variability and change|climate change]] vs. <sup>18</sup>O|alt=Time evolution of oxygen-18 concentration on the scale of 500 million years showing many local peaks.]]

[[Paleoclimatology|Paleoclimatologists]] measure the ratio of oxygen-18 and oxygen-16 in the [[Exoskeleton|shells]] and [[skeleton]]s of marine organisms to determine the climate 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, and 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 as old as hundreds of thousands of years.<ref>{{cite journal|last = Hand|first = Eric|title = The Solar System's first breath|journal = Nature|volume = 452|page = 259|date = March 13, 2008|doi = 10.1038/452259a|pmid = 18354437|issue = 7185|bibcode = 2008Natur.452..259H |s2cid = 789382|doi-access = free}}</ref><ref>{{Cite book|title=Reconstructing Quaternary Environments|last1=Lowe|first1=J. John|last2=Walker|first2=Mike|publisher=Routledge|year=2014|isbn=978-0-415-74075-3|edition=3rd|location=Abingdon, UK|pages=165–170}}</ref>

[[Geology of solar terrestrial planets|Planetary geologists]] have measured the relative quantities of oxygen isotopes in samples from the [[Earth]], the [[Moon]], [[Mars]], and [[meteorite]]s, but were long unable to obtain reference values for the isotope ratios in the [[Sun]], believed to be the same as those of the [[Nebular hypothesis|primordial solar nebula]]. Analysis of a [[silicon]] wafer exposed to the [[solar wind]] in space and returned by the crashed [[Genesis (spacecraft)|''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 [[Protoplanetary disk|disk of protoplanetary material]] prior to the coalescence of dust grains that formed the Earth.<ref>{{cite journal | last1=McKeegan | first1=K. D. | last2=Kallio | first2=A. P. A. | last3=Heber | first3=V. S. | last4=Jarzebinski | first4=G. | last5=Mao | first5=P. H. | last6=Coath | first6=C. D. | last7=Kunihiro | first7=T. | last8=Wiens | first8=R. C. | last9=Nordholt | first9=J. E. | last10=Moses | first10=R. W. | last11=Reisenfeld | first11=D. B. | last12=Jurewicz | first12=A. J. G. | last13=Burnett | first13=D. S. | title=The Oxygen Isotopic Composition of the Sun Inferred from Captured Solar Wind | journal=Science | volume=332 | issue=6037 | date=2011-06-24 | issn=0036-8075 | doi=10.1126/science.1204636 | pages=1528–1532 | pmid=21700868 | bibcode=2011Sci...332.1528M | url=https://www.science.org/doi/10.1126/science.1204636 | access-date=2025-05-27| url-access=subscription }}</ref><ref>{{cite journal | last1=Rumble | first1=D. | last2=Bowring | first2=S. | last3=Iizuka | first3=T. | last4=Komiya | first4=T. | last5=Lepland | first5=A. | last6=Rosing | first6=M. T. | last7=Ueno | first7=Y. | title=The oxygen isotope composition of earth's oldest rocks and evidence of a terrestrial magma ocean | journal=Geochemistry, Geophysics, Geosystems | volume=14 | issue=6 | date=2013 | issn=1525-2027 | doi=10.1002/ggge.20128 | pages=1929–1939 | bibcode=2013GGG....14.1929R | url=https://agupubs.onlinelibrary.wiley.com/doi/10.1002/ggge.20128 | access-date=2025-05-27| url-access=subscription }}</ref>

Oxygen presents two spectrophotometric [[absorption band]]s 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 [[Earth observation satellite|satellite]] platform.<ref>{{cite conference|title=Progress on the development of an integrated canopy fluorescence model|last1=Miller|first1=J. R.|display-authors=4|author2=Berger, M.|author3=Alonso, L.|author4=Cerovic, Z.|author5=Goulas, Y.|author6=Jacquemoud, S.|author7=Louis, J.|author8=Mohammed, G.|author9=Moya, I.|author10=Pedros, R.|author11=Moreno, J.F.|author12=Verhoef, W.|author13=Zarco-Tejada, P.J.|work=Geoscience and Remote Sensing Symposium, 2003. IGARSS '03. Proceedings. 2003 IEEE International|year=2003 |volume=1 |pages=601–603 |doi=10.1109/IGARSS.2003.1293855|isbn=0-7803-7929-2 |citeseerx=10.1.1.473.9500}}</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.<ref>{{cite web | title=Paleoclimatology: The Oxygen Balance | publisher=NASA Earth Observatory | date=2011-03-11 | url=https://earthobservatory.nasa.gov/features/Paleoclimatology_OxygenBalance | access-date=2025-05-27}}</ref>
{{Clear}}

==Biological production and role of O<sub>2</sub>==
{{Main|Dioxygen in biological reactions}}
<!-- CopyEdits Only&nbsp;— DIRECT ALL FUTURE EXPANSION to [[dioxygen in biological reactions]] -->

===Photosynthesis and respiration===
<!-- CopyEdits Only&nbsp;— DIRECT ALL FUTURE EXPANSION to [[dioxygen in biological reactions]] -->
[[File:Simple photosynthesis overview.svg|thumb|Photosynthesis splits water to liberate {{chem|O|2}} and fixes {{chem|CO|2}} into sugar in what is called a [[Calvin cycle]].|alt=A diagram of photosynthesis processes, including income of water and carbon dioxide, illumination and release of oxygen. Reactions produce ATP and NADPH in a Calvin cycle with a sugar as a by product.]]
In nature, [[free oxygen]] is produced as a [[byproduct]] of [[photolysis|light-driven splitting]] of water during [[chlorophyll]]ic [[photosynthesis]]. According to some estimates, marine [[photoautotroph]]s such as [[red algae|red]]/[[green algae]] and [[cyanobacteria]] provide about 70% of the free oxygen produced on Earth, and the rest is produced in terrestrial environments by plants.<ref>{{cite book|chapter-url=https://books.google.com/books?id=g6RfkqCUQyQC&pg=PA147|title=Plants: the potentials for extracting protein, medicines, and other useful chemicals (workshop proceedings)|date=September 1983|chapter=Marine Plants: A Unique and Unexplored Resource|last=Fenical|first=William|page=147|isbn=978-1-4289-2397-3|publisher=DianePublishing|access-date=August 23, 2020|archive-date=March 25, 2015|archive-url=https://web.archive.org/web/20150325221600/http://books.google.com/books?id=g6RfkqCUQyQC&pg=PA147|url-status=live}}</ref> Other estimates of the oceanic contribution to atmospheric oxygen are higher, while some estimates are lower, suggesting oceans produce ~45% of Earth's atmospheric oxygen each year.<ref>{{cite book|last=Walker|first=J. C. G.|date=1980|title=The oxygen cycle in the natural environment and the biogeochemical cycles|publisher=Springer-Verlag|location=Berlin}}</ref>

A simplified overall formula for photosynthesis is<ref>{{cite book|last1=Brown|first1=Theodore L. |last2=LeMay|first2=Burslen|title=Chemistry: The Central Science|url=https://archive.org/details/studentlectureno00theo|url-access=registration|isbn=978-0-13-048450-5|page=958|date=2003|publisher=Prentice Hall/Pearson Education}}</ref>

: 6 {{CO2}} + 6 {{chem|H|2|O}} + photons → {{chem|C|6|H|12|O|6}} + 6 {{chem|O|2}}

or simply

: carbon dioxide + water + sunlight → [[glucose]] + dioxygen

Photolytic [[oxygen evolution]] occurs in the [[thylakoid membrane]]s of photosynthetic organisms and requires the energy of four [[photon]]s.<ref group=lower-alpha>Thylakoid membranes are part of [[chloroplast]]s 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 progenitors 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 production of the water molecule) is released into the atmosphere.<ref group=lower-alpha>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>

Oxygen is used in [[mitochondria]] of [[eukaryote]]s to generate 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 {{CO2}} + 6 {{chem|H|2|O}} + 2880 kJ/mol

In [[aquatic animal]]s, [[dissolved oxygen]] in water is absorbed by [[gill]]s, [[cutaneous respiration|through the skin]] or [[enteral respiration|via the gut]]; in terrestrial animals such as [[tetrapod]]s, oxygen in air is [[breathing|actively taken into the body]] via [[lung]]s, where [[gas exchange]] takes place to [[diffusion|diffuse]] oxygen into the blood and carbon dioxide out, and the body's [[circulatory system]] then transports the oxygen to other tissues where cellular respiration takes place.<ref>{{cite book|last1=Hall|first1=John|title=Guyton and Hall textbook of medical physiology|date=2011|publisher=Saunders/Elsevier|location=Philadelphia, Pa.|isbn=978-1-4160-4574-8|page=5|edition=12th}}</ref><ref name="Pocock2">{{cite book|last1=Pocock|first1=Gillian|last2=Richards|first2=Christopher D.|title=Human physiology : the basis of medicine|date=2006|publisher=Oxford University Press|location=Oxford|isbn=978-0-19-856878-0|page=311|edition=3rd}}</ref> However, in insects, the most successful and [[biodiverse]] terrestrial [[clade]], oxygen is directly conducted to the internal tissues via [[respiratory system of insects|a deep network of airways]]. [[Hemoglobin]] in [[red blood cell]]s binds {{chem|O|2}}, changing color from bluish red to bright red<ref name="GuideElem48" /> ({{chem|CO|2}} is released from another part of hemoglobin through the [[Bohr effect]]). Other terrestrial [[invertebrate]]s use [[hemocyanin]] ([[mollusc]]s and some [[arthropod]]s) or [[hemerythrin]] (spiders and lobsters) instead.<ref name="NBB298" /> A liter of blood can dissolve up to 200&nbsp;cm<sup>3</sup> of {{chem|O|2}}.<ref name="NBB298" />

Until the discovery of [[anaerobic organism|anaerobic]] [[animal|metazoa]],<ref name="pmid20370908">{{cite journal |display-authors=4 |author=Danovaro R |author2=Dell'anno A |author3=Pusceddu A|author4=Gambi C |author5=Heiner I|author6=Kristensen RM |title=The first metazoa living in permanently anoxic conditions |journal=BMC Biology |volume=8 |issue=1 |pages=30 |date=April 2010 |pmid=20370908 |pmc=2907586 |doi=10.1186/1741-7007-8-30 |doi-access=free}}</ref> oxygen was thought to be a requirement for all complex life.<ref>{{cite book |last1=Ward |first1=Peter D. |last2=Brownlee |first2=Donald |title=Rare Earth: Why Complex Life is Uncommon in the Universe |publisher=Copernicus Books (Springer Verlag) |date=2000 |isbn=978-0-387-98701-9 |page=217}}</ref>

[[Reactive oxygen species]], such as [[superoxide]] ion ({{chem|O|2|-}}) and [[hydrogen peroxide]] ({{chem|H|2|O|2}}), are reactive by-products of oxygen use in organisms.<ref name="NBB298" /> Parts of the [[immune system]] of higher organisms 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" /> Oxygen is damaging to [[Obligate anaerobe|obligately anaerobic organisms]], which were the dominant form of [[Evolutionary history of life|early life]] on Earth until {{chem|O|2}} began to accumulate in the atmosphere about 2.5 billion years ago during the [[Great Oxygenation Event]], about a billion years after the first appearance of these organisms.<ref>{{cite press release |title=NASA Research Indicates Oxygen on Earth 2.5 Billion Years ago |url=http://www.nasa.gov/home/hqnews/2007/sep/HQ_07215_Timeline_of_Oxygen_on_Earth.html |publisher=[[NASA]] |date=September 27, 2007 |access-date=March 13, 2008 |archive-date=March 13, 2008 |archive-url=https://web.archive.org/web/20080313063940/http://www.nasa.gov/home/hqnews/2007/sep/HQ_07215_Timeline_of_Oxygen_on_Earth.html |url-status=live }}</ref><ref name="NYT-20131003">{{cite news |last=Zimmer |first=Carl |author-link=Carl Zimmer |title=Earth's Oxygen: A Mystery Easy to Take for Granted |url=https://www.nytimes.com/2013/10/03/science/earths-oxygen-a-mystery-easy-to-take-for-granted.html |date=October 3, 2013 |work=[[The New York Times]] |access-date=October 3, 2013 |archive-date=May 16, 2020 |archive-url=https://web.archive.org/web/20200516083101/https://www.nytimes.com/2013/10/03/science/earths-oxygen-a-mystery-easy-to-take-for-granted.html |url-status=live }}</ref>

An adult human at rest inhales<!--simply inhales (most is exhaled again) or takes up and respires?--> 1.8 to 2.4&nbsp;grams of oxygen per minute.<ref>{{Cite web|url=https://patents.google.com/patent/US6224560B1/en|title=Flow restrictor for measuring respiratory parameters|access-date=August 4, 2019|archive-date=May 8, 2020|archive-url=https://web.archive.org/web/20200508103811/https://patents.google.com/patent/US6224560B1/en|url-status=live}}</ref> This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year.<ref group=lower-alpha>(1.8 grams/min/person)×(60 min/h)×(24 h/day)×(365 days/year)×(6.6 billion people)/1,000,000 g/t=6.24 billion tonnes</ref>

===Living organisms===
{{anchor|partial pressure}}
{|class="wikitable" style="float:right; margin-left:25px"
|+Partial pressures of oxygen in the human body (PO<sub>2</sub>)
|-
! Unit !! Alveolar pulmonary<br /> gas pressures !! Arterial blood oxygen !! Venous blood gas
|-
| [[kPa]] || 14.2 || 11{{efn|name=mmHg|Derived from mmHg values using 0.133322 kPa/mmHg}}-13{{efn|name=mmHg}} || 4.0{{efn|name=mmHg}}-5.3{{efn|name=mmHg}}
|-
| [[mmHg]] || 107 || 75<ref name="southwest">
[http://pathcuric1.swmed.edu/PathDemo/nrrt.htm Normal Reference Range Table] {{Webarchive|url=https://web.archive.org/web/20111225185659/http://pathcuric1.swmed.edu/PathDemo/nrrt.htm |date=December 25, 2011 }} from The University of Texas Southwestern Medical Center at Dallas. Used in Interactive Case Study Companion to Pathologic basis of disease.
</ref>-100<ref name="southwest" /> || 30<ref name="brookside" />-40<ref name="brookside">[http://www.brooksidepress.org/Products/OperationalMedicine/DATA/operationalmed/Lab/ABG_ArterialBloodGas.htm The Medical Education Division of the Brookside Associates--> ABG (Arterial Blood Gas)] {{Webarchive|url=https://web.archive.org/web/20170812201558/http://www.brooksidepress.org/Products/OperationalMedicine/DATA/operationalmed/Lab/ABG_ArterialBloodGas.htm |date=August 12, 2017 }} Retrieved on December 6, 2009</ref>
|-
|}
The free oxygen [[partial pressure]] in the body of a living vertebrate organism is highest in the [[respiratory system]], and decreases along any [[arterial system]], peripheral tissues, and [[venous system]], respectively. Partial pressure is the pressure that oxygen would have if it alone occupied the volume.<ref>{{cite book|author=Charles Henrickson|title=Chemistry|publisher=Cliffs Notes|date=2005|isbn=978-0-7645-7419-1|url=https://archive.org/details/chemistry00henr}}</ref>

===Build-up in the atmosphere===
{{Main|Geological history of oxygen}}
<!-- CopyEdits Only&nbsp;— DIRECT ALL FUTURE EXPANSION to [[Geological history of oxygen]] or [[dioxygen in biological reactions]] -->

[[File:Oxygenation-atm.svg|thumb|left|upright=1.35|{{chem|O|2}} build-up in Earth's atmosphere: 1) no {{chem|O|2}} produced; 2) {{chem|O|2}} produced, but absorbed in oceans & seabed rock; 3) {{chem|O|2}} starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer; 4–5) {{chem|O|2}} sinks filled and the gas accumulates<ref>{{Cite journal |last=Holland |first=Heinrich D |date=2006-06-29 |title=The oxygenation of the atmosphere and oceans |journal=Philosophical Transactions of the Royal Society B: Biological Sciences |language=en |volume=361 |issue=1470 |pages=903–915 |doi=10.1098/rstb.2006.1838 |issn=0962-8436 |pmc=1578726 |pmid=16754606}}</ref>|alt=A graph showing time evolution of oxygen pressure on Earth; the pressure increases from zero to 0.2 atmospheres.]]

Before [[photosynthesis]] evolved, [[Earth's atmosphere]] had little free [[diatomic]] [[chemical element|element]]al [[oxygen]] (O<sub>2</sub>).<ref name=Krump-2008>{{Cite journal |last=Kump |first=Lee R. |date=2008-01-17 |title=The rise of atmospheric oxygen |url=https://www.nature.com/articles/nature06587 |journal=Nature |language=en |volume=451 |issue=7176 |pages=277–278 |doi=10.1038/nature06587 |pmid=18202642 |bibcode=2008Natur.451..277K |issn=0028-0836}}</ref><ref name="NYT-20131003"/>
Oxygen began building up in the [[prebiotic atmosphere]] at approximately 2.45 Ga during the [[Neoarchean]]-[[Paleoproterozoic]] boundary, a [[paleogeological]] event known as the [[Great Oxygenation Event]] (GOE).<ref name=Canfield-2005>{{Cite journal |journal=Annual Review of Earth and Planetary Sciences |volume=33 |issue=1 |language=en |issn=0084-6597 |date=May 31, 2005 |pages=1–36 |doi=10.1146/annurev.earth.33.092203.122711 |url=https://www.annualreviews.org/doi/10.1146/annurev.earth.33.092203.122711 |title=The Early History of Atmospheric Oxygen: Homage to Robert M. Garrels |first=D.E. |last=Canfield |bibcode=2005AREPS..33....1C }}</ref> The concentrations of O<sub>2</sub> attained were less than 10% of today's and probably fluctuated greatly. Around 500Mya a second event known as the [[Neoproterozoic Oxygenation Event]] lead to oxygen levels similar or even higher than the present.<ref name=Mills-2023/> 

Oxygen is both a result of biological activity and a key enabler. [[Photosynthesis]] produces oxygen while plants and animals using [[aerobic respiration]] consume it. Consequently the [[evolution of life]] is closely related to the concentration of available oxygen. Understanding the relationship between oxygen and evolution would aid in seeking evidence of [[extraterrestrial life]] in [[exoplanet]] data.<ref name=Mills-2023>{{Cite journal |last1=Mills |first1=Benjamin J. W. |last2=Krause |first2=Alexander J. |last3=Jarvis |first3=Ian |last4=Cramer |first4=Bradley D. |date=2023-05-31 |title=Evolution of Atmospheric O2 Through the Phanerozoic, Revisited |url=https://www.annualreviews.org/content/journals/10.1146/annurev-earth-032320-095425 |journal=Annual Review of Earth and Planetary Sciences |language=en |volume=51 |pages=253–276 |doi=10.1146/annurev-earth-032320-095425 |issn=0084-6597}}</ref>{{rp|252}} Oxygen concentration plays a key role in the geochemical composition of sedimentary rocks, making oxygen concentration important for geology and sedimentary rocks important for understanding oxygen concentration over geologic time.<ref>{{Cite journal |last=Kendall |first=Brian |date=2021-05-30 |title=Recent Advances in Geochemical Paleo-Oxybarometers |url=https://www.annualreviews.org/doi/10.1146/annurev-earth-071520-051637 |journal=Annual Review of Earth and Planetary Sciences |language=en |volume=49 |issue=1 |pages=399–433 |doi=10.1146/annurev-earth-071520-051637 |bibcode=2021AREPS..49..399K |issn=0084-6597}}</ref> The increase in oxygen concentrations had wide-ranging and significant impacts on Earth's [[geochemistry]] and [[biosphere]]. However, detailed connections between oxygen and evolution remain elusive.<ref name=Lyons-2014>{{Cite journal |last1=Lyons |first1=Timothy W. |last2=Reinhard |first2=Christopher T. |last3=Planavsky |first3=Noah J. |date=2014 |title=The rise of oxygen in Earth's early ocean and atmosphere |url=https://www.nature.com/articles/nature13068 |journal=Nature |language=en |volume=506 |issue=7488 |pages=307–315 |doi=10.1038/nature13068 |pmid=24553238 |bibcode=2014Natur.506..307L |issn=0028-0836}}</ref>

Variations in atmospheric oxygen concentration may have shaped past climates. When oxygen declined, atmospheric density dropped, which in turn increased surface evaporation, causing precipitation increases and warmer temperatures.<ref>{{cite journal|url=http://ns.umich.edu/new/releases/22942-variations-in-atmospheric-oxygen-levels-shaped-earth-s-climate-through-the-ages|doi=10.1126/science.1260670|pmid=26068848|journal=Science|title=Long-term climate forcing by atmospheric oxygen concentrations|author1=Poulsen, Christopher J.|author2=Tabor, Clay|author3=White, Joseph D.|volume=348|issue=6240|pages=1238–41|bibcode=2015Sci...348.1238P|year=2015|s2cid=206562386|access-date=June 12, 2015|archive-date=July 13, 2017|archive-url=https://web.archive.org/web/20170713125418/http://ns.umich.edu/new/releases/22942-variations-in-atmospheric-oxygen-levels-shaped-earth-s-climate-through-the-ages|url-status=live|url-access=subscription}}</ref>
{{clear}}

===Extraterrestrial free oxygen===
{{Main|Extraterrestrial atmosphere}}
In the field of [[astrobiology]] and in the search for [[extraterrestrial life]] oxygen is a strong [[biosignature]]. 
The value of this signature dependents on the stability of the oxygen content on Earth. Simulations of Earth's oxygen balance have  
estimated that oxygen on Earth will last for about one billion years.<ref>{{Cite journal|url=https://www.nature.com/articles/s41561-021-00693-5|title=The future lifespan of Earth's oxygenated atmosphere|first1=Kazumi|last1=Ozaki|first2=Christopher T.|last2=Reinhard|date=March 9, 2021|journal=Nature Geoscience|volume=14|issue=3|pages=138–142|via=www.nature.com|doi=10.1038/s41561-021-00693-5|arxiv=2103.02694|bibcode=2021NatGe..14..138O |s2cid=232083548 }}</ref><ref>{{Cite web|url=https://www.eurekalert.org/news-releases/825455|title=How much longer will the oxygen-rich atmosphere be sustained on Earth?|website=EurekAlert!}}</ref>
Another potential issue is that oxygen [[Extraterrestrial atmosphere#Abiotic oxygen|may be produced abiotically]] on [[celestial bodies]] with processes and conditions (such as a peculiar [[hydrosphere]]) which allow free oxygen,<ref>{{cite web|url=https://earthsky.org/space/oxygen-exoplanets-not-always-indicator-of-life|title=Oxygen and life: a cautionary tale|date=3 January 2019|author=Paul Scott Anderson|access-date=29 December 2020|archive-date=January 22, 2021|archive-url=https://web.archive.org/web/20210122134654/https://earthsky.org/space/oxygen-exoplanets-not-always-indicator-of-life|url-status=live}}</ref><ref>{{cite journal | vauthors = Luger R, Barnes R | title = Extreme water loss and abiotic O2 buildup on planets throughout the habitable zones of M dwarfs | journal = Astrobiology | volume = 15 | issue = 2 | pages = 119–43 | date = February 2015 | pmid = 25629240 | pmc = 4323125 | doi = 10.1089/ast.2014.1231 | bibcode = 2015AsBio..15..119L | arxiv = 1411.7412 }}</ref><ref>{{cite journal |last1=Wordsworth |first1=Robin |last2=Pierrehumbert |first2=Raymond |title=Abiotic oxygen-dominated atmospheres on terrestrial habitable zone planets |journal=The Astrophysical Journal |date=1 April 2014 |volume=785 |issue=2 |pages=L20 |doi=10.1088/2041-8205/785/2/L20 |bibcode=2014ApJ...785L..20W |arxiv=1403.2713 |s2cid=17414970 }}</ref> like with [[Europa (moon)|Europa]]'s and [[Ganymede (moon)|Ganymede]]'s thin oxygen atmospheres.<ref name="Hall1998">{{cite journal |last1=Hall |first1=D.T. |last2=Feldman |first2=P.D. |last3=McGrath |first3=M.A. |last4=Strobel |first4=D. F. |display-authors=2 |title=The Far-Ultraviolet Oxygen Airglow of Europa and Ganymede |journal=The Astrophysical Journal |date=1998 |volume=499 |issue=1 |pages=475–81 |doi=10.1086/305604 |bibcode=1998ApJ...499..475H |doi-access=free }}</ref>

==Industrial production==
{{see also|Air separation|Oxygen evolution|Fractional distillation}}
[[File:Hofmann voltameter fr.svg|thumb|upright|[[Hofmann voltameter|Hofmann electrolysis apparatus]] used in electrolysis of water|alt=A drawing of three vertical pipes connected at the bottom and filled with oxygen (left pipe), water (middle) and hydrogen (right). Anode and cathode electrodes are inserted into the left and right pipes and externally connected to a battery.]]

Every year one hundred&nbsp;million&nbsp;tonnes of {{chem|O|2}} are extracted from air for industrial uses.<ref name="NBB300" /> The most common method of extraction is [[fractional distillation]] of liquefied air, with {{chem|N|2}} [[distillation|distilling]] as a vapor while {{chem|O|2}} is left as a liquid.<ref name="NBB300" /> The other primary method of producing {{chem|O|2}} is 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-[[cryogenics|cryogenic]] technologies (see also the related [[vacuum swing adsorption]]).<ref>{{cite web|url=http://www.uigi.com/noncryo.html|title=Non-Cryogenic Air Separation Processes|date=2003|access-date=December 16, 2007|publisher=UIG Inc.|archive-date=October 3, 2018|archive-url=https://web.archive.org/web/20181003082121/http://www.uigi.com/noncryo.html|url-status=live}}</ref>

[[File:A setup for preparation of Oxygen.jpg|alt=An experiment setup with test tubes to prepare oxygen|left|thumb|280x280px|An experiment setup for preparation of oxygen in academic laboratories]]
In academic laboratories, oxygen can be prepared by heating together potassium chlorate mixed with a small proportion of manganese dioxide.<ref>{{Cite book|url=https://archive.org/details/flescscho_1114918|title=A school chemistry|last=Flecker|first=Oriel Joyce|publisher=Oxford, Clarendon press|others=MIT Libraries|year=1924|page=[https://archive.org/details/flescscho_1114918/page/n41 30]}}</ref>

Oxygen gas can also be produced through [[electrolysis of water]] into molecular oxygen and hydrogen. DC electricity must be used: if AC is used, the gases in each limb consist of hydrogen and oxygen in the explosive ratio 2:1. A similar method is the electrocatalytic {{chem|O|2}} evolution from oxides 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 method is 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" />

===Storage===
[[File:Compressed gas cylinders.mapp and oxygen.triddle.jpg|thumb|Oxygen and [[MAPP gas]] compressed-gas cylinders with regulators]]

[[Oxygen storage]] methods include high-pressure [[oxygen tank]]s, cryogenics and chemical compounds. For reasons of economy, oxygen is often transported in bulk as a liquid in specially insulated tankers, since one [[litre|liter]] of liquefied oxygen is equivalent to 840&nbsp;liters of gaseous oxygen at atmospheric pressure and {{convert|20|C|F}}.<ref name="NBB300" /> Such tankers are used to refill bulk liquid-oxygen storage containers, which stand outside hospitals and other institutions that need large volumes of pure oxygen gas. Liquid oxygen is passed through [[heat exchanger]]s, 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==
{{see also|Breathing gas|Redox|Combustion}}

===Medical===
[[File:Home oxygen concentrator.jpg|thumb|upright|left|An [[oxygen concentrator]] in an [[emphysema]] patient's house|alt=A gray device with a label DeVILBISS LT4000 and some text on the front panel. A green plastic pipe is running from the device.]]

{{Main|Oxygen therapy}}

Uptake of {{chem|O|2}} from the air is the essential purpose of [[Respiration (physiology)|respiration]], so oxygen supplementation is used in [[medicine]]. Treatment not only increases oxygen levels in the patient's blood, but has the secondary effect of decreasing resistance to blood flow in many types of diseased lungs, easing work load on the heart. [[Oxygen therapy]] is used to treat [[emphysema]], [[pneumonia]], some heart disorders ([[congestive heart failure]]), some disorders that cause increased [[pulmonary artery pressure]], 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.<ref name="pmid18540928">{{cite journal |author=Sim MA |display-authors=4 |author2=Dean P |author3=Kinsella J |author4= Black R |author5=Carter R|author6=Hughes M |title=Performance of oxygen delivery devices when the breathing pattern of respiratory failure is simulated |journal=Anaesthesia |volume=63 |issue=9 |pages=938–40 |date=2008 |pmid=18540928 |doi=10.1111/j.1365-2044.2008.05536.x|s2cid=205248111 |doi-access=free }}</ref>

[[Hyperbaric medicine|Hyperbaric]] (high-pressure) medicine uses special [[hyperbaric oxygen chamber|oxygen chambers]] to increase the [[partial pressure]] of {{chem|O|2}} around the patient and, when needed, the medical staff.<ref name="pmid8931286">{{cite journal |author=Stephenson RN |author2=Mackenzie I |author3=Watt SJ |author4=Ross JA |title=Measurement of oxygen concentration in delivery systems used for hyperbaric oxygen therapy |journal=Undersea Hyperb Med |volume=23 |issue=3 |pages=185–88 |date=1996 |pmid=8931286 |url=http://archive.rubicon-foundation.org/2245 |access-date=September 22, 2008 |archive-date=August 11, 2011 |archive-url=https://web.archive.org/web/20110811175247/http://archive.rubicon-foundation.org/2245 |url-status=usurped }}</ref> [[Carbon monoxide poisoning]], [[gas gangrene]], and [[decompression sickness]] (the 'bends') are sometimes addressed with this therapy.<ref>{{cite web|url=http://www.uhms.org/Default.aspx?tabid=270 |title=Indications for hyperbaric oxygen therapy |author=Undersea and Hyperbaric Medical Society |access-date=September 22, 2008 |author-link=Undersea and Hyperbaric Medical Society |url-status=dead |archive-url=https://web.archive.org/web/20080912184905/http://www.uhms.org/Default.aspx?tabid=270 |archive-date=September 12, 2008 }}</ref> Increased {{chem|O|2}} concentration in the lungs helps to displace [[carbon monoxide]] from the heme group of [[hemoglobin]].<ref>{{cite web |url=http://www.uhms.org/ResourceLibrary/Indications/CarbonMonoxidePoisoning/tabid/272/Default.aspx |title=Carbon Monoxide |author=Undersea and Hyperbaric Medical Society |access-date=September 22, 2008 |archive-url=https://web.archive.org/web/20080725005744/http://www.uhms.org/ResourceLibrary/Indications/CarbonMonoxidePoisoning/tabid/272/Default.aspx <!--Added by H3llBot--> |archive-date=July 25, 2008}}</ref><ref name="pmid15233173">{{cite journal |author=Piantadosi CA |title=Carbon monoxide poisoning |journal=Undersea Hyperb Med |volume=31 |issue=1 |pages=167–77 |date=2004 |pmid=15233173 |url=http://archive.rubicon-foundation.org/4002 |access-date=September 22, 2008 |archive-date=February 3, 2011 |archive-url=https://web.archive.org/web/20110203090807/http://archive.rubicon-foundation.org/4002 |url-status=usurped }}</ref> Oxygen gas is poisonous to the [[anaerobic bacteria]] that cause gas gangrene, so increasing its partial pressure helps kill them.<ref>{{cite journal |author=Hart GB |author2=Strauss MB |title=Gas Gangrene&nbsp;– Clostridial Myonecrosis: A Review |journal=J. Hyperbaric Med |volume=5 |issue=2 |pages=125–44 |date=1990 |url=http://archive.rubicon-foundation.org/4428 |access-date=September 22, 2008 |archive-date=February 3, 2011 |archive-url=https://web.archive.org/web/20110203090838/http://archive.rubicon-foundation.org/4428 |url-status=usurped }}</ref><ref>{{cite journal |author=Zamboni WA |author2=Riseman JA |author3=Kucan JO |title=Management of Fournier's Gangrene and the role of Hyperbaric Oxygen |journal=J. Hyperbaric Med |volume=5 |issue=3 |pages=177–86 |date=1990 |url=http://archive.rubicon-foundation.org/4431 |access-date=September 22, 2008 |archive-date=February 3, 2011 |archive-url=https://web.archive.org/web/20110203090958/http://archive.rubicon-foundation.org/4431 |url-status=usurped }}</ref> Decompression sickness occurs in divers who decompress too quickly after a dive, resulting in bubbles of inert gas, mostly nitrogen and helium, forming in the blood. Increasing the pressure of {{chem|O|2}} as soon as possible helps to redissolve the bubbles back into the blood so that these excess gasses can be exhaled naturally through the lungs.<ref name="ECE510" /><ref>{{cite web |url=http://www.uhms.org/ResourceLibrary/Indications/DecompressionSickness/tabid/275/Default.aspx |title=Decompression Sickness or Illness and Arterial Gas Embolism |author=Undersea and Hyperbaric Medical Society |access-date=September 22, 2008 |archive-url=https://web.archive.org/web/20080705210353/http://www.uhms.org/ResourceLibrary/Indications/DecompressionSickness/tabid/275/Default.aspx <!--Added by H3llBot--> |archive-date=July 5, 2008}}</ref><ref>{{cite journal |last=Acott |first=C. |title=A brief history of diving and decompression illness |journal=South Pacific Underwater Medicine Society Journal |volume=29 |issue=2 |date=1999 |url=http://archive.rubicon-foundation.org/6004 |access-date=September 22, 2008 |archive-date=September 5, 2011 |archive-url=https://web.archive.org/web/20110905152645/http://archive.rubicon-foundation.org/6004 |url-status=usurped }}</ref> Normobaric oxygen administration at the highest available concentration is frequently used as first aid for any diving injury that may involve inert gas bubble formation in the tissues. There is epidemiological support for its use from a statistical study of cases recorded in a long term database.<ref name="Longphre et al 2007">{{cite journal|title=First aid normobaric oxygen for the treatment of recreational diving injuries |last1=Longphre |first1=JM |last2=Denoble |first2=PJ |last3=Moon |first3=RE |last4=Vann |first4=RD |last5=Freiberger |first5=JJ |journal=Undersea & Hyperbaric Medicine |date=2007 |volume=34 |issue=1 |pages=43–49|url=https://pdfs.semanticscholar.org/3c96/eec9b2ae3f25ffc0569f26b7329d5b05e213.pdf |archive-url=https://web.archive.org/web/20181001104203/https://pdfs.semanticscholar.org/3c96/eec9b2ae3f25ffc0569f26b7329d5b05e213.pdf |url-status=dead |archive-date=2018-10-01 |via=Rubicon Research Repository |pmid=17393938 |s2cid=3236557 }}</ref><ref name="Emergency O2 for scuba">{{cite web |url=https://www.diversalertnetwork.org/training/courses/course_eo2 |title=Emergency Oxygen for Scuba Diving Injuries |publisher=Divers Alert Network |author=<!--not specified--> |access-date=October 1, 2018 |archive-date=April 20, 2020 |archive-url=https://web.archive.org/web/20200420114653/https://www.diversalertnetwork.org/training/courses/course_eo2 |url-status=live }}</ref><ref name="DAN Europe">{{cite web |url=https://daneurope.org/web/guest/readarticle;jsessionid=F8EB8916CD93E6A793F9F875BF5FC782?p_p_id=web_content_reading&p_p_lifecycle=0&p_p_mode=view&p_r_p_-1523133153_groupId=10103&p_r_p_-1523133153_articleId=11601&p_r_p_-1523133153_articleVersion=1.0&p_r_p_-1523133153_commaCategories=&p_r_p_-1523133153_commaTags= |title=Oxygen First Aid for Scuba Diving Injuries |publisher=Divers Alert Network Europe |author=<!--not specified--> |access-date=October 1, 2018 |archive-date=June 10, 2020 |archive-url=https://web.archive.org/web/20200610202203/https://daneurope.org/web/guest/readarticle;jsessionid=F8EB8916CD93E6A793F9F875BF5FC782?p_p_id=web_content_reading&p_p_lifecycle=0&p_p_mode=view&p_r_p_-1523133153_groupId=10103&p_r_p_-1523133153_articleId=11601&p_r_p_-1523133153_articleVersion=1.0&p_r_p_-1523133153_commaCategories=&p_r_p_-1523133153_commaTags= |url-status=live }}</ref>
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===Life support and recreational use===
[[File:STS057-89-067 - Wisoff on the Arm (Retouched).jpg|thumb|Low-pressure pure {{chem|O|2}} is used in [[space suit]]s.]]

In modern [[space suit]]s, which surround their occupant's body, oxygen gas is used as a low-pressure [[breathing gas]]. 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 suit flexibility.<ref name="pmid11541018">{{cite journal|author=Morgenthaler GW|author2=Fester DA|author3=Cooley CG| title=As assessment of habitat pressure, oxygen fraction, and EVA suit design for space operations|journal=Acta Astronautica |volume= 32|issue=1|pages=39–49|date=1994|pmid=11541018|doi=10.1016/0094-5765(94)90146-5|bibcode = 1994AcAau..32...39M }}</ref><ref name="pmid2730484">{{cite journal|author=Webb JT|author2= Olson RM|author3=Krutz RW|author4=Dixon G|author5=Barnicott PT|title=Human tolerance to 100% oxygen at 9.5 psia during five daily simulated 8-hour EVA exposures|journal=Aviat Space Environ Med|volume=60|issue=5|pages=415–21|date=1989|pmid=2730484|doi=10.4271/881071}}</ref>

[[Scuba diving|Scuba]] and [[Surface-supplied diving|surface-supplied]] [[underwater diving|underwater diver]]s and [[submarine]]s also rely on artificially delivered {{chem|O|2}}. Submarines, submersibles and [[atmospheric diving suits]] usually operate at normal atmospheric pressure. Breathing air is scrubbed of carbon dioxide by chemical extraction and oxygen is replaced to maintain a constant partial pressure. [[Ambient pressure]] divers breathe air or gas mixtures with an oxygen fraction suited to the operating depth. Pure or nearly pure {{chem|O|2}} use in diving at pressures higher than atmospheric is usually limited to [[rebreathers]], or [[Decompression (diving)|decompression]] at relatively shallow depths (~6 meters depth, or less),<ref name="Acott">{{cite journal|last=Acott|first=C.|title=Oxygen toxicity: A brief history of oxygen in diving|journal=South Pacific Underwater Medicine Society Journal|volume=29|issue=3|date=1999|url=http://archive.rubicon-foundation.org/6014|access-date=September 21, 2008|archive-date=December 25, 2010|archive-url=https://web.archive.org/web/20101225073221/http://archive.rubicon-foundation.org/6014|url-status=usurped}}</ref><ref name="Longphre">{{cite journal|last1=Longphre|first1=J. M.|title=First aid normobaric oxygen for the treatment of recreational diving injuries|journal=Undersea Hyperb. Med.|volume=34|issue=1|pages=43–49|date=2007|pmid=17393938|url=http://archive.rubicon-foundation.org/5514|access-date=September 21, 2008|display-authors=4|last2=Denoble|first2=P. J.|last3=Moon|first3=R. E.|last4=Vann|first4=R. D.|last5=Freiberger|first5=J. J.|archive-url=https://web.archive.org/web/20080613163501/http://archive.rubicon-foundation.org/5514|archive-date=June 13, 2008|url-status=usurped}}</ref> or [[Hyperbaric treatment schedules|medical treatment in recompression chambers]] at pressures up to 2.8 bar, where acute oxygen toxicity can be managed without the risk of drowning. Deeper diving requires significant dilution of {{chem|O|2}} with other gases, such as nitrogen or helium, to prevent [[oxygen toxicity]].<ref name="Acott" />

People who climb mountains or fly in non-pressurized [[fixed-wing aircraft]] sometimes have supplemental {{chem|O|2}} supplies.<ref group=lower-alpha>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> Pressurized commercial airplanes have an emergency supply of {{chem|O|2}} automatically supplied to the passengers in case of cabin depressurization. Sudden cabin pressure loss activates [[chemical oxygen generator]]s above each seat, causing [[oxygen mask]]s to drop. Pulling on the masks "to start the flow of oxygen" as cabin safety instructions dictate, forces iron filings into the [[sodium chlorate]] inside the canister.<ref name="NBB301" /> A steady stream of oxygen gas is then produced by the [[exothermic]] reaction.<ref name=GE28/>

Oxygen, as a mild [[euphoria|euphoric]], has a history of recreational use in [[oxygen bar]]s and in [[sport]]s. Oxygen bars are establishments found in the United States since the late 1990s that offer higher than normal {{chem|O|2}} exposure for a minimal fee.<ref name="FDA-O2Bars">{{cite journal|url=https://www.fda.gov/Fdac/features/2002/602_air.html| title=Oxygen Bars: Is a Breath of Fresh Air Worth It?|last=Bren|first=Linda|journal=FDA Consumer Magazine| volume=36| issue=6| pages=9–11|publisher=U.S. Food and Drug Administration|date=November–December 2002|access-date=December 23, 2007|archive-url=https://web.archive.org/web/20071018041754/https://www.fda.gov/Fdac/features/2002/602_air.html|archive-date=October 18, 2007|url-status=dead| pmid=12523293}}</ref> Professional athletes, especially in [[American football]], sometimes go off-field between plays to don oxygen masks to boost performance. The pharmacological effect is doubted; a [[placebo]] effect is a more likely explanation.<ref name="FDA-O2Bars" /> Available studies support a performance boost from oxygen enriched mixtures only if it is inhaled ''during'' [[aerobic exercise]].<ref>{{cite web|url=http://www.pponline.co.uk/encyc/1008.htm|title= Ergogenic Aids|access-date=January 4, 2008|publisher=Peak Performance Online |archive-url = https://web.archive.org/web/20070928051412/http://www.pponline.co.uk/encyc/1008.htm <!--Added by H3llBot--> |archive-date = September 28, 2007}}</ref>

Other recreational uses that do not involve breathing 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)|access-date=March 14, 2008|archive-url=https://web.archive.org/web/20090211213613/http://www.bkinzel.de/misc/ghg/index.html|archive-date=February 11, 2009|url-status=dead}}</ref><!-- - Primary source; many secondary sources exist but they only provide less information and more ads - -->

===Industrial===
[[File:Clabecq JPG01.jpg|thumb|Most commercially produced {{chem|O|2}} is used to [[smelting|smelt]] and/or [[Decarburization|decarburize]] [[iron]].|alt=An elderly worker in a helmet is facing his side to the viewer in an industrial hall. The hall is dark but is illuminated yellow glowing splashes of a melted substance.]]

[[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, {{chem|SO|2}} and {{chem|CO|2}}. The reactions are [[exothermic reaction|exothermic]], so the temperature increases to 1,700&nbsp;°[[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 [[fabric]]s).<ref name="NBB301" />

Most of the remaining 20% of commercially produced oxygen is used in medical applications, [[gas welding|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 {{convert|60|cm|abbr=on}} 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>

==Compounds==
{{Main|Oxygen compounds}}
<!-- DIRECT ALL FUTURE EXPANSION to [[Oxygen compounds]] -->
[[File:Stilles Mineralwasser.jpg|thumb|upright|[[Water]] ({{chem|H|2|O}}) is the most familiar oxygen compound.|alt=Water flowing from a bottle into a glass.]]
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 name=GE28/>{{rp|28}} 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]]).<ref>[[International Union of Pure and Applied Chemistry|IUPAC]]: [http://old.iupac.org/publications/books/rbook/Red_Book_2005.pdf ''Red Book.''] {{Webarchive|url=https://web.archive.org/web/20180709210050/http://old.iupac.org/publications/books/rbook/Red_Book_2005.pdf |date=July 9, 2018 }} pp.&nbsp;73, 320.</ref>

===Oxides and other inorganic compounds===
<!-- DIRECT ALL FUTURE EXPANSION to [[Compounds of oxygen]] -->
[[Water]] ({{chem|H|2|O}}) is an 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 per hydrogen atom) to an adjacent oxygen atom in a separate molecule.<ref>{{cite journal|first1=P.|last1=Maksyutenko|first2=T. R.|last2=Rizzo|first3=O. V.|last3=Boyarkin|date=2006|title=A direct measurement of the dissociation energy of water|pmid=17115729|journal=J. Chem. Phys.|page=181101 |doi=10.1063/1.2387163|issue=18|volume=125|bibcode = 2006JChPh.125r1101M |url=http://infoscience.epfl.ch/record/90816 }}</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|access-date=January 6, 2008|date=January 4, 2008|archive-date=October 10, 2007|archive-url=https://web.archive.org/web/20071010055658/http://www.lsbu.ac.uk/water/hbond.html|url-status=live}}</ref><ref group=lower-alpha>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>

[[File:Rust screw.jpg|thumb|left|Oxides, such as [[iron oxide]] or [[rust]], form when oxygen combines with other elements.|alt=A rusty piece of a bolt.]]
Due to its [[electronegativity]], oxygen forms [[chemical bond]]s with almost all other elements to give corresponding [[oxide]]s. The surface of most metals, such as [[aluminium]] and [[titanium]], are oxidized in the presence of air and become coated with a thin film of oxide that [[Passivation (chemistry)|passivates]] the metal and slows further [[corrosion]]. Many oxides of the [[transition metal]]s are [[non-stoichiometric compound]]s, with slightly less metal than the [[chemical formula]] would show. For example, the mineral [[Iron(II) oxide|FeO]] ([[wüstite]]) is written as <math chem>\ce{Fe}_{1-x}\ce{O}</math>, where ''x'' is usually around 0.05.<ref>{{cite book|first1=Lesley E.|last1=Smart|last2=Moore|first2=Elaine A. |title=Solid State Chemistry: An Introduction|edition=3rd |publisher=CRC Press|date=2005|page=214|isbn=978-0-7487-7516-3}}</ref>

Oxygen is present in the atmosphere in trace quantities in the form of [[carbon dioxide]] ({{chem|CO|2}}). The [[Earth's crust]]al [[Rock (geology)|rock]] is composed in large part of oxides of [[silicon]] ([[Silicon dioxide|silica]] {{chem|SiO|2}}, as found in [[granite]] and [[quartz]]), 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 [[calcium carbonate]] (in [[limestone]]). The rest of the Earth's crust is also made of oxygen compounds, in particular various complex [[silicate]]s (in [[silicate minerals]]). The Earth's mantle, of much larger mass than the crust, is largely composed of silicates of magnesium and iron.<ref>{{cite book | title = Inorganic Chemistry | last1 = Housecroft | first1 = C. E. | last2 = Sharpe | first2 = A. G. | year = 2018 | publisher = Prentice Hall | edition = 5th | isbn = 978-1292134147}}</ref>

Water-[[solubility|soluble]] silicates in the form of {{chem|Na|4|SiO|4}}, {{chem|Na|2|SiO|3}}, and {{chem|Na|2|Si|2|O|5}} are used as [[detergent]]s and [[adhesive]]s.<ref name="ECE507">[[#Reference-idCook1968|Cook & Lauer 1968]], p. 507</ref>

Oxygen also acts as a [[ligand]] for transition metals, forming [[transition metal dioxygen complexes]], which feature metal–{{chem|O|2}}. This class of compounds includes the [[heme]] proteins [[hemoglobin]] and [[myoglobin]].<ref>{{cite book|last=Crabtree|first=R.|title=The Organometallic Chemistry of the Transition Metals|edition=3rd |publisher=John Wiley & Sons|date=2001|page=152|isbn=978-0-471-18423-2}}</ref> An exotic and unusual reaction occurs with [[platinum hexafluoride|{{chem|PtF|6}}]], which oxidizes oxygen to give {{chem2|O2+PtF6−}}, [[dioxygenyl hexafluoroplatinate]].<ref name="ECE505">[[#Reference-idCook1968|Cook & Lauer 1968]], p.505</ref>

===Organic compounds===
<!-- DIRECT ALL FUTURE EXPANSION to [[Compounds of oxygen]] -->
[[File:Acetone-3D-vdW.png|thumb|[[Acetone]] is an important feeder material in the chemical industry.
{{legend|red|Oxygen}}
{{legend|black|Carbon}}
{{legend|white|Hydrogen|outline=silver}}
|alt=A ball structure of a molecule. Its backbone is a zig-zag chain of three carbon atoms connected in the center to an oxygen atom and on the end to 6 hydrogens.]]

Among the most important classes of organic compounds that contain oxygen are (where "R" is an organic group): [[Alcohol (chemistry)|alcohol]]s ({{chem|R-OH}}); [[ether]]s ({{chem|R-O-R}}); [[ketone]]s ({{chem|R-CO-R}}); [[aldehyde]]s ({{chem|R-CO-H}}); [[carboxylic acid]]s ({{chem|R-COOH}}); [[ester]]s ({{chem|R-COO-R}}); [[Organic acid anhydride|acid anhydrides]] ({{chem|R-CO-O-CO-R}}); and [[amide]]s ({{chem|R-CO-NR|2}}). 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 ({{chem|(CH|3|)|2|CO}}) and [[phenol]] ({{chem|C|6|H|5|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 ethers in which the oxygen atom is part of a ring of three atoms. The element is similarly found in almost all [[biomolecule]]s that are important to (or generated by) life.<ref name="Claydentext">{{Clayden}}</ref>
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" />

==Safety and precautions==
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The [[NFPA 704]] standard rates compressed oxygen gas as nonhazardous to health, nonflammable and nonreactive, but an oxidizer. Refrigerated liquid oxygen (LOX) is given a health hazard rating of 3 (for increased risk of [[hyperoxia]] from condensed vapors, and for hazards common to cryogenic liquids such as frostbite), and all other ratings are the same as the compressed gas form.<ref name="nfpa">{{cite web|url = http://www.rivcoeh.org/Portals/0/documents/guidance/hazmat/bep_nfparatings.pdf|publisher = Riverside County Department of Environmental Health|access-date = August 22, 2017|title = NFPA 704 ratings and id numbers for common hazardous materials|archive-date = July 11, 2019|archive-url = https://web.archive.org/web/20190711171240/http://www.rivcoeh.org/Portals/0/documents/guidance/hazmat/bep_nfparatings.pdf|url-status = live}}</ref>

===Toxicity===
{{Main|Oxygen toxicity}}
[[File:Symptoms of oxygen toxicity.png|thumb|left|upright=1.35|Main symptoms of oxygen toxicity<ref>{{cite journal |author=Dharmeshkumar N Patel |display-authors=4 |author2=Ashish Goel |author3=SB Agarwal |author4=Praveenkumar Garg |author5=Krishna K Lakhani |title=Oxygen Toxicity |journal=Indian Academy of Clinical Medicine |volume=4 |issue=3 |page=234 |date=2003 |url=http://medind.nic.in/jac/t03/i3/jact03i3p234.pdf |access-date=April 26, 2009 |archive-date=September 22, 2015 |archive-url=https://web.archive.org/web/20150922093352/http://medind.nic.in/jac/t03/i3/jact03i3p234.pdf |url-status=dead }}</ref>|alt=A diagram showing a male torso and listing symptoms of oxygen toxicity: Eyes&nbsp;– visual field loss, nearsightedness, cataract formation, bleeding, fibrosis; Head&nbsp;– seizures; Muscles&nbsp;– twitching; Respiratory system&nbsp;– jerky breathing, irritation, coughing, pain, shortness of breath, tracheobronchitis, acute respiratory distress syndrome.]]

Oxygen gas ({{chem|O|2}}) can be [[Oxygen toxicity|toxic]] at elevated [[partial pressure]]s, leading to [[convulsion]]s and other health problems.<ref name="Acott" /><ref group=lower-alpha>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 kilo[[Pascal (unit)|pascals]]&nbsp;(kPa), equal to about 50% oxygen composition at standard pressure or 2.5 times the normal sea-level {{chem|O|2}} partial pressure of about 21&nbsp;kPa. This is not a problem except for patients on [[mechanical ventilator]]s, since gas supplied through [[oxygen mask]]s in medical applications is typically composed of only 30–50% {{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 the oxygen content being too high.<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 spacecraft|Apollo]], causes no damage due to the low total pressures used.<ref name="pmid11541018" /><ref>{{cite web|last = Wade|first = Mark|date = 2007|url = http://www.astronautix.com/craftfam/spasuits.htm|title = Space Suits|publisher = Encyclopedia Astronautica |access-date=December 16, 2007 |url-status = dead|archive-url = https://web.archive.org/web/20071213122134/http://www.astronautix.com/craftfam/spasuits.htm |archive-date = December 13, 2007}}</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.<ref>{{cite web |url=http://www.globalrph.com/martin_4_most2.htm |title=The Four Most Important Equations In Clinical Practice |last=Martin |first=Lawrence |website=GlobalRPh |publisher=David McAuley |access-date=June 19, 2013 |archive-date=September 5, 2018 |archive-url=https://web.archive.org/web/20180905215615/http://www.globalrph.com/martin_4_most2.htm |url-status=live }}</ref>

Oxygen toxicity to the lungs and [[central nervous system]] can also occur in deep [[scuba diving]] and [[surface-supplied diving]].<ref name="NBB299" /><ref name="Acott" /> 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 journal |author=Wilmshurst P |title=Diving and oxygen |journal=BMJ |volume=317 |issue=7164 |pages=996–99 |date=1998 |pmid=9765173 |pmc=1114047 |doi=10.1136/bmj.317.7164.996}}</ref> Exposure to an {{chem|O|2}} partial pressure greater than 160&nbsp;kPa (about 1.6 atm) may lead to convulsions (normally fatal for divers). Acute oxygen toxicity (causing seizures, its most feared effect for divers) can occur by breathing an air mixture with 21% {{chem|O|2}} at {{convert|66|m|abbr=on}} or more of depth; the same thing can occur by breathing 100% {{chem|O|2}} at only {{convert|6|m|abbr=on}}.<ref name="BMJ" /><ref name="Donald">{{cite book |last=Donald |first=Kenneth |title=Oxygen and the Diver |isbn = 978-1-85421-176-7|date=1992 |publisher=SPA in conjunction with K. Donald |location=England}}</ref><ref name="Donald1">{{cite journal |author=Donald K. W. |title=Oxygen Poisoning in Man: Part I |journal=Br Med J |volume=1 |issue=4506 |pages=667–72 |date=1947 |pmc=2053251 |doi=10.1136/bmj.1.4506.667 |pmid=20248086}}</ref><ref name="Donald2">{{cite journal |author=Donald K. W. |title=Oxygen Poisoning in Man: Part II |journal=Br Med J |volume=1 |pages=712–17 |date=1947 |pmc=2053400|issue=4507 |doi=10.1136/bmj.1.4507.712 |pmid=20248096}}</ref>

===Combustion and other hazards===
[[File:Apollo 1 fire.jpg|thumb|right|The interior of the [[Apollo 1]] Command Module. Pure {{chem|O|2}} at higher than normal pressure and a spark led to a fire and the loss of the Apollo 1 crew.|alt=The inside of a small spaceship, charred and apparently destroyed.]]
Unless non-flammable containers are used or all sources of ignition are eliminated, oxygen rich environments are extremely hazardous. Many materials including most metals burn faster in oxygen rich environments and ignite at lower temperatures.<ref>{{Cite report |url=https://ntrs.nasa.gov/citations/19970001787 |title=Guide for Oxygen Hazards Analyses on Components and Systems |last1=Stoltzfus |first1=Joel M. |last2=Dees |first2=Jesse |last3=Poe |first3=Robert F. |date=1996-10-01 |language=en}}</ref> 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 1967 fire that killed the [[Apollo 1]] crew in a launch pad test spread so rapidly because the capsule was pressurized with pure {{chem|O|2}} at slightly more than atmospheric pressure to force the doors against their seals. Once in space the exterior pressure would be minimal and the oxygen pressure would have been reduced to 5psi to reduce fire hazard. The ignition source for the fire was traced to an electrical surge under the seat of one of the astronauts, probably due to chafed wires.<ref>{{Cite journal |last1=DiLisi |first1=Greg |last2=McLean |first2=Stella |date=2019-04-01 |title=The Apollo 1 Fire: A Case Study in the Flammability of Fabrics |url=https://pubs.aip.org/pte/article/57/4/236/278421/The-Apollo-1-Fire-A-Case-Study-in-the-Flammability |journal=The Physics Teacher |language=en |volume=57 |issue=4 |pages=236–239 |doi=10.1119/1.5095379 |bibcode=2019PhTea..57..236D |issn=0031-921X|url-access=subscription }}</ref>

Liquid oxygen spills, if allowed to soak into organic matter, such as [[wood]], [[petrochemical]]s, and [[Bitumen|asphalt]] can cause these materials to [[Detonation|detonate]] unpredictably on subsequent mechanical impact.<ref name="astm-tpt" />
{{clear}}

==See also==
{{div col|colwidth=20em}}
* [[Geological history of oxygen]]
* [[Hypoxia (environmental)]] for {{chem|O|2}} depletion in aquatic ecology
* [[Ocean deoxygenation]]
* [[Hypoxia (medical)]], a lack of oxygen
* [[Limiting oxygen concentration]]
* [[:Category:Oxygen compounds|Oxygen compounds]]
* [[Oxygen plant]]
* [[Oxygen sensor]]
* [[Dark oxygen]]
{{div col end}}
{{Subject bar
|book1=Oxygen
|book2=Period 2 elements
|book3=Chalcogens
|book4=Chemical elements (sorted&nbsp;alphabetically)
|book5=Chemical elements (sorted by number)
|portal1=Chemistry
|portal2=Medicine
|commons=y
|wikt=y
|v=y
|v-search=Oxygen atom
|b=y
|b-search=Wikijunior:The Elements/Oxygen
}}

==Notes==
{{reflist|30em|group=lower-alpha}}

==References==
<!-- Full reference information for Cook, Daintith, and Emsley given in the "General references" subsection -->
{{reflist | 30em}}

===General references===
<!-- Please do not list cite web references here unless it is cited more than once -->
* <!-- Co -->{{cite book| ref=Reference-idCook1968|title=The Encyclopedia of the Chemical Elements| chapter-url=https://archive.org/details/encyclopediaofch00hamp| chapter-url-access=registration|last1=Cook|first1=Gerhard A.|last2=Lauer|first2=Carol M.|publisher=Reinhold Book Corporation|location=New York|date=1968|pages=[https://archive.org/details/encyclopediaofch00hamp/page/499 499–512]|editor=Clifford A. Hampel|chapter=Oxygen| lccn=68-29938}}
* <!-- Em -->{{cite book|ref=Reference-idEmsley2001|title=Nature's Building Blocks: An A–Z Guide to the Elements|last=Emsley|first=John|publisher=Oxford University Press|date=2001|location=Oxford, England|isbn=978-0-19-850340-8|chapter=Oxygen|pages=[https://archive.org/details/naturesbuildingb0000emsl/page/297 297–304]|chapter-url=https://archive.org/details/naturesbuildingb0000emsl/page/297}}
* <!-- Ra -->{{cite book| ref=Reference-idRaven2005 |last1=Raven|first1=Peter H.|first2=Ray F.|last2=Evert|first3=Susan E.|last3=Eichhorn|title=Biology of Plants| url=https://archive.org/details/biologyofplants00rave_0 | url-access=registration |edition=7th|publisher=W. H. Freeman and Company Publishers|date=2005|location = New York|pages=[https://archive.org/details/biologyofplants00rave_0/page/115 115–27]|isbn = 978-0-7167-1007-3}}

==External links==
{{Spoken Wikipedia|En-oxygen-article.ogg|date=2008-06-23}}
* [https://www.periodicvideos.com/videos/008.htm Oxygen] at ''[[The Periodic Table of Videos]]'' (University of Nottingham)
* [https://www.organic-chemistry.org/chemicals/oxidations/oxygen.shtm Oxidizing Agents > Oxygen]
* [https://www.uigi.com/oxygen.html Oxygen (O<sub>2</sub>) Properties, Uses, Applications]
* [https://www.americanscientist.org/issues/pub/the-story-of-o Roald Hoffmann article on "The Story of O"] {{Webarchive|url=https://web.archive.org/web/20170607115422/http://www.americanscientist.org/issues/pub/the-story-of-o |date=June 7, 2017 }}
* [https://www.webelements.com/webelements/elements/text/O/index.html WebElements.com&nbsp;– Oxygen]
* {{In Our Time|Oxygen|b0088nql|Oxygen}}
* [https://scrippso2.ucsd.edu/ Scripps Institute: Atmospheric Oxygen has been dropping for 20 years]

{{Periodic table (navbox)}}
{{diatomicelements}}
{{Oxygen compounds}}
{{Authority control}}

[[Category:Oxygen| ]]
[[Category:Chemical elements]]
[[Category:Diatomic nonmetals]]
[[Category:Reactive nonmetals]]
[[Category:Chalcogens]]
[[Category:Chemical substances for emergency medicine]]
[[Category:Breathing gases]]
[[Category:E-number additives]]
[[Category:Oxidizing agents]]