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8 nitrogen ← oxygen → fluorine - ↑ O ↓ S Periodic table - Extended periodic table
General
Name, symbol, number	oxygen, O, 8
Chemical series	nonmetals, chalcogens
Group, period, block	16, 2, p
Appearance	colorless gas above light blue liquid
Standard atomic weight	15.9994(3) g·mol−1
Electron configuration	1s2 2s2 2p4
Electrons per shell	2, 6
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 1.429 g/L
Melting point	54.36 K (-218.79 °C, -361.82 °F)
Boiling point	90.20 K (-182.95 °C, -297.31 °F)
Critical point	154.59 K, 5.043 MPa
Heat of fusion	(O2) 0.444 kJ·mol−1
Heat of vaporization	(O2) 6.82 kJ·mol−1
Heat capacity	(25 °C) (O2) 29.378 J·mol−1·K−1
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K       61 73 90
Atomic properties
Crystal structure	cubic
Oxidation states	2, 1, −1, −2 (neutral oxide)
Electronegativity	3.44 (Pauling scale)
Ionization energies (more)	1st: 1313.9 kJ·mol−1
	2nd: 3388.3 kJ·mol−1
	3rd: 5300.5 kJ·mol−1
Atomic radius	60 pm
Atomic radius (calc.)	48 pm
Covalent radius	73 pm
Van der Waals radius	152 pm
Miscellaneous
Magnetic ordering	paramagnetic
Thermal conductivity	(300 K) 26.58x10-3  W·m−1·K−1
Speed of sound	(gas, 27 °C) 330 m/s
CAS registry number	7782-44-7
Selected isotopes
Main article: Isotopes of oxygen iso NA half-life DM DE (MeV) DP 16O 99.76% O is stable with 8 neutrons 17O 0.038% O is stable with 9 neutrons 18O 0.21% O is stable with 10 neutrons
References
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Oxygen (pronounced /ˈɒksɪdʒən/) is a colorless, odorless, tasteless, gaseous chemical element with the chemical symbol O and atomic number 8. It is a chalcogen, period 2, nonmetallic element that can form binary compounds (known as oxides) with almost all the other elements. The valency of oxygen is 2 and the most common oxidation state is -2.^[1] On Earth it is usually bonded to other elements covalently or ionically. Oxygen is the third most abundant element in the universe by mass (after hydrogen and helium),^[2] the most abundant element by mass in the Earth's crust,^[3] and the most abundant element by mass in the human body.^[4]

The name oxygen was coined in 1777 by Antoine Lavoisier from the Greek roots οξύς (oxys) (acid, lit. "sharp," from the taste of acids) and -γενής (-genēs) (producer, lit. begetter), because he mistakenly thought that it was a constituent of all acids.^[5]

Free Diatomic oxygen or dioxygen (O₂) is, together with nitrogen, one of the two major components of air, constituting about a fifth of its volume.^[6]

Oxygen is highly reactive, and readily forms compounds with most other elements. Its compounds with silicon and metals are abundant in the earth's crustal rocks and with hydrogen in water (H₂O). The nucleic acids and all of the major classes of structural molecules in living organisms, proteins, polysaccharides and fats contain oxygen, as do the major inorganic compounds that comprise animal shell (calcium carbonate), tooth and bone (calcium phosphate). Dioxygen is produced from water by cyanobacteria, algae and plants during photosynthesis, and the energy required to sustain life in aerobically-respiring organisms is provided by enzyme-mediated oxidation of sugars and carboxylic acids, which are themselves already oxygen-containing compounds. Without oxygen, most organisms with aerobic respiration die within minutes.^[7] However, free oxygen is toxic to obligate anaerobic organisms and was a poisonous waste product for early life on Earth.

Main articles: Triplet oxygen and Singlet oxygen

At standard temperature and pressure, oxygen is a colorless, odorless gas with the molecular formula O₂, in which the two oxygen atoms are chemically bonded to each other with a triplet electron configuration. This bond has a bond order of two, and is often simplified in description as a double bond.^[8]

The molecular orbital diagram of dioxygen (middle) in the ground triplet state.

Triplet oxygen is the ground state of the oxygen molecule. The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals. These orbitals are classified as antibonding (weakening the bond order from three to two), so the diatomic oxygen bond is weaker than the diatomic nitrogen triple bond in which all bonding molecular orbitals are filled, but fewer antibonding ones are. Though unpaired electrons are commonly associated with high reactivity in chemical compounds, triplet oxygen is relatively (and fortunately) nonreactive by comparison with most radicals.

In normal triplet form oxygen molecules are paramagnetic due to the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O₂ molecules.^[9] Liquid oxygen is attracted to a magnet to a sufficient extent that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.^[10] Oxygen's paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine the purity of gaseous oxygen.^[11]

Singlet oxygen, a name given to several higher energy species of molecular oxygen in which all the electron spins are paired, is much more reactive towards common organic molecules. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.^[12] It is also produced in the troposphere by the photolysis of ozone by light of short wavelength,^[13] and by the immune system as a source of active oxygen.^[14] Carotenoids in photosynthetic organisms (and possibly also in animals) play a major role in absorbing energy from singlet oxygen and converting it to the unexcited ground state before it can cause harm to tissues.^[15]

Oxygen is more soluble in water than nitrogen, water containing approximately 1 part of oxygen to 2 of nitrogen, compared with a ratio in the atmosphere of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg l^-1)dissolves at 0^oC than at 20^oC(7.6 mg l^-1.^[16][17] At 25°C at 1 atm of air, freshwater contains about 6.04 cc (8.63 mg, 0.27 mmol) of oxygen per litre, whereas sea water contains about 4.9 cc (7.0 mg, 0.22 mmol) per litre. At 0°C the solubilities increase to 10.3 cc (14.7 mg, 0.46 mmol) per litre for water and 8.0 cc (11.4 mg, 0.36 mmol) per litre for sea water.

Oxygen condenses at 90.20 K (-182.95 °C, -297.31 °F), and freezes at 54.36 K (-218.79 °C, -361.82 °F). Both liquid and solid O₂ are clear substances with a light sky blue color caused by absorption in the red (by contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light). High purity liquid O₂ is usually obtained by the fractional distillation of liquified air;^[18] Liquid oxygen, may also be produced by condensation out of air, using liquid nitrogen as a coolant. It is a highly reactive substance and must be segregated from combustible materials.^[19]

Main article: Allotropes of oxygen

The simplified representation of Dioxygen.

Ozone is a rare gas on Earth found mostly in the stratosphere.

The common allotrope of elemental oxygen on Earth, O₂, is known as dioxygen. Elemental oxygen is most commonly encountered in this form, as about 21% (by volume) of the Earth's atmosphere. O₂ has a bond length of 121 pm and a bond energy of 498 kJ/mol.^[20]

Triatomic oxygen (Ozone, O₃), the less common triatomic allotrope of oxygen, is a poisonous gas with a distinctive odor. Traces of it can be detected as a sharp, chlorine-like smell coming from electric motors, laser printers and photocopiers. It was named "ozone" by Christian Friedrich Schönbein, in 1840, from the Greek word ÖĮώ (ozo) for smell. ^[5] Ozone is thermodynamically unstable toward the more common dioxygen form, and is formed by reaction of O₂ with atomic oxygen produced by splitting of O₂ by UV radiation in the upper atmosphere.^[5] Ozone absorbs strongly in the ultraviolet and functions as a shield for the biosphere against the mutagenic and other damaging effects of solar UV radiation (see ozone layer).^[5] Ozone is also formed by electrostatic discharge in the presence of dioxygen. The immune system produces ozone as an antimicrobial (see below).^[14] Liquid and solid O₃ have a deeper blue color than ordinary oxygen and they are unstable and explosive.^[21][5]

A newly discovered allotrope of oxygen, tetraoxygen (O₄), is a deep red solid that is created by pressurizing O₂ to the order of 20 GPa. Its properties are being studied for use in rocket fuels and similar applications, as it is a much more powerful oxidizer than either O₂ or O₃.^[22][23] When tetraoxygen is subjected to a pressure of 96 GPa, it becomes metallic, similarly to hydrogen,^[24] and becomes more similar to the heavier chalcogens, such as tellurium and polonium, both of which show significant metallic character.

Water (H₂O) is the most familiar oxygen compound

Main article: Oxygen compounds

In almost all known compounds of oxygen, the oxidation state of oxygen is -2. The oxidation state -1 is found in a few compounds such as peroxides. Compounds containing oxygen in other oxidation states are very uncommon: -1/2 (superoxides),-1/3 (ozonides), 0 (elemental, hypofluorous acid), +1/2 (dioxygenyl), +1 (dioxygen difluoride) and +2 (oxygen difluoride).

Arguably, the most familiar oxygen compound is water, the oxide of hydrogen, H₂O. However the earth's crustal rock is predominantly composed of oxides of silicon as silica, SiO₂ (found in granite and sand), silicates (found in feldspars), and oxygen compounds of metals such as calcium (as calcium carbonate in limestone), aluminium (as silicates in feldspars and as aluminium oxide in bauxite and corundum), iron (as iron (III) oxide Fe₂O₃ in hematite and rust), etc.

Other important examples of oxygen compounds include the compounds of carbon and oxygen, such as carbon dioxide (CO₂), alcohols (R-OH where "R" is an organic group), carbonyls (R-CO-H or R-CO-R) such as formaldehyde and acetone and carboxylic acids (R-COOH) such as acetic acid and palmitic acid, a common fatty acid found in animals and plants and the food products derived from them, such as palm oil and milk products.

Oxygenated radicals such as chlorates (ClO₃⁻),perchlorates (ClO₄⁻), chromates (CrO₄²⁻), dichromates (Cr₂O₇²⁻), permanganates (MnO₄⁻) and nitrates (NO₃⁻) are strong oxidizing agents. Phosphorus is biologically important in its oxygenated form as the phosphate (PO₄³⁻) ion, present in bone as calcium phosphate and as the backbone of RNA and DNA.

Oxygen forms compounds with almost all of the other known elements, including some of the rarest elements: technetium (TcO₄⁻), promethium (Pm₂O₃) and neptunium (NpO₂); and also with some of the least reactive elements such as xenon (XeO₃), gold (Au₂O₃) and platinum (PtO₂). Synthetic elements that have known oxides include plutonium (PuO₂), americium (AmO₂), curium (CuO₂), berkelium (BkO₃), californium (Cf₂O₃) and einsteinium (Es₂O₃).

One unexpected oxygen compound is dioxygen hexafluoroplatinate, O₂⁺PtF₆⁻, discovered when Neil Bartlett was studying the properties of platinum hexafluoride (PtF₆).^[25] He noticed a change in color when this compound was exposed to atmospheric air and reasoned that xenon should be oxidized by PtF₆. This led him to the discovery of xenon hexafluoroplatinate Xe⁺PtF₆⁻.

The cation O₂²⁺ in O₂F₂ is only formed in the presence of stronger oxidants than oxygen, which limits it to oxygen fluorides, e.g. oxygen fluoride.^[21]

When dissolved in water, many metallic oxides form alkaline solutions while many oxides of nonmetals form acidic solutions. For example, sodium oxide in solution forms the strong base sodium hydroxide while phosphorus pentoxide in solution forms phosphoric acid.^[26]

The alkali metals in groups 1 and 2 of the periodic table - lithium, sodium, potassium, rubidium, cesium, magnesium, calcium, strontium and barium, all react spontaneously with oxygen when exposed to air to form oxides and form hydroxides in the presence of water. None of these elements are found in nature as free metals. Cesium is so reactive with oxygen that it is used as a getter in vacuum tubes. The surface of aluminium is always oxidised in the presence of air, coated with a thin film of aluminium oxide that passivates the metal and slows further corrosion. The aluminium oxide layer can be built to greater thickness by the process of electrolytic anodising. Although solid magnesium and aluminium react slowly with oxygen at STP, they are both capable of burning in air, generating very high temperatures, and the metal powders may form explosive mixtures with air.

Oxides, such iron oxide or rust, Fe₂O₃, form when oxygen combines with other elements

Some substances need to be heated before they will react with oxygen in bulk but some, such as iron, readily forms iron oxide, or rust, Fe₂O₃. The production of free oxygen by photosynthetic bacteria some 3.5 billion years ago precipitated iron out of solution in the oceans as Fe₂O₃ in the economically-important iron ore hematite.

Due to its electronegativity, oxygen forms chemical bonds with almost all other free elements at elevated temperatures to give corresponding oxides. So-called noble metals (common examples: gold, platinum) resist direct chemical combination with oxygen, and substances like gold(III) oxide must be formed by an indirect route.

Peroxides retain some of oxygen's original molecular structure(^-O-O^-). White or light yellow sodium peroxide (Na₂O₂) is formed when metallic sodium (Na) is burned in oxygen. Each oxygen atom in its peroxide ion may have a full octet of 4 pairs of electrons.^[27] Superoxides are a class of compounds that are very similar to peroxides, but with just one unpaired electron for each pair of oxygen atoms (O₂^-).^[27]. These compounds form from oxidation of alkali metals with larger ionic radii (K, Rb, Cs). For example, potassium superoxide ( KO₂) is an orange-yellow solid formed when potassium (K) reacts with oxygen.

Hydrogen peroxide (H₂O₂) can be produced by passing a volume of 96 to 98% hydrogen and 2 to 4% oxygen through an electric discharge.^[26] A more commercially viable method is allow autoxidation of an organic intermediate; 2-ethylanthrahydroquinone dissolved in an organic solvent is oxidized to H₂O₂ and 2-ethylanthraquinone.^[26] The 2-ethylanthraquinone is then reduced and recycled back into the process.

Quartz is a common crystalline mineral made of silica, or silicon dioxide (SiO₂)

Silica is the common name for the compound silicon dioxide (SiO₂). Quartz (illustrated) is the naturally-occurring crystalline mineral form of silica, and common deposits of quartz are in igneous rocks such as granite, sandstones and sand.

Most chemically combined oxygen is locked in a class of minerals called silicates (which in turn are the major component of rocks and clays). The basic structure of silicates consists of two parts; units of silicon surrounded by four oxygen anions in a tetrahedral arrangement and units of metal-oxygen polyhedra that contain metal cations (examples: aluminium, calcium, iron and sodium).^[27] Both units are linked together by sharing oxygen anions, forming complex polymers in the process.

Water- soluble silicates in the form of Na₄SiO₄, Na₂SiO₃, and Na₂Si₂O₅ are used as detergents and adhesives.^[27] Na_xSi_xO_x with a higher ratio of SiO₂ to Na₂O has a greater molecular weight and a lower solubility.

Acetone is an important feeder material in the chemical industry.
Oxygen is in red, carbon in black and hydrogen in white.

Among the most important classes of organic compounds that contain oxygen are (where "R" is an organic group): alcohols (R-OH); ethers (R-O-R); ketones (R-CO-R); aldehydes (R-CO-H); carboxylic acids (R-COOH); esters (R-COO-R); acid anhydrides (R-CO-O-CO-R); amides (R-C(O)-NR₂). There are many important organic solvents that contain oxygen, among which: acetone, methanol, ethanol, isopropanol, furan, THF, diethyl ether, dioxane, ethylacetate, DMF, DMSO, acetic acid, formic acid. Acetone ((CH₃)₂CO) and phenol (C₆H₅OH) 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, acetamide, etc. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.

Oxygen reacts spontaneously with many organic compounds at or below room temperature in a process called autoxidation.^[26] Most of the organic compounds that contain oxygen are not made by direct action of oxygen. Commercially important organic compounds that are made by direct oxidation of a precursor include^[27]:

- Ethylene oxide (used to make the antifreeze ethylene glycol) is obtained by direct oxidation of ethylene: C₂H₄ + ½ O₂ +catalyst→ C₂H₄O

- Peracetic acid (feeder material for various epoxy compounds) is obtained from acetaldehyde: CH₃CHO + O₂ +catalyst→ CH₃C(O)-OOH

Of the organic compounds with biological relevance, carbohydrates (such as glucose) contain a large amount of oxygen. All fatty acids (such as oleic acid) and aminoacids contain oxygen (due to the presence of carboxyl group). Furthermore, seven of the amino acids incorporate oxygen in the side-chain too: serine, tyrosine, threonine, glutamic acid, glutamine, aspartic acid and asparagine. Oxygen also occurs in phosphate groups in the biologically important energy-carrying molecules ATP and ADP and in the backbone of RNA and DNA.

The oxygen cycle maintains the ~21% concentration of O₂ in the earth's atmosphere

See also: Silicate minerals and Category:Oxide minerals

Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.^[2] About 0.87% of the Sun's mass is in the form of oxygen.^[28] Oxygen constitutes 49.2% of the Earth's crust by mass^[3] and is the most common component of the world's oceans (88.81% by mass).^[28] It is also the second most common component of the Earth's atmosphere, taking up 20.947% of its volume and 23.14% of its mass (some million billion tonnes).^[29][28][30] The Earth is unusual in having such a high concentration of free oxygen in its atmosphere. With 0.15% oxygen by volume, the atmosphere of Mars has the second most abundant concentration by volume of any planet in the solar system while Venus comes in third place.^[2] However, their oxygen is solely produced by ultraviolet radiation impacting oxygen-containing molecules such as carbon dioxide.

Cold water holds more dissolved oxygen.

The unusually high concentration of elemental oxygen on earth is the result of the oxygen cycle. This biogeochemical cycle describes the movement of oxygen within and between its three main reservoirs on earth: the atmosphere, the biosphere, and the lithosphere. The main driving factor of the oxygen cycle is photosynthesis, which is responsible for the modern Earth's atmosphere. Because of the vast amounts of oxygen in the atmosphere, even if all photosynthesis were to cease it would take at least 5,000 years to strip out more or less all oxygen.^[31]

Free elemental dioxygen also occurs in solution in the world's water bodies. The higher solubility of O₂ at low temperatures (see Physical Properties) has important implications for ocean life, as polar oceans support a much higher density of life due to their higher oxygen content.^[32] Polluted water may have reduced amounts of oxygen in it, depleted by decaying algae and other biomaterials (see eutrophication). Scientists assess this aspect of water quality by measuring the water's biochemical oxygen demand (BOD), or the amount of oxygen needed to restore a normal oxygen concentration.^[6]

Main article: Isotopes of oxygen

Late in a massive star's life, ¹⁶O concentrates in the O-shell, ¹⁷O in the H-shell and ¹⁸O in the He-shell

Naturally occurring oxygen is composed of 3 stable isotopes, ¹⁶O, ¹⁷O, and ¹⁸O, with ¹⁶O being the most abundant (99.762% natural abundance).^[33] Oxygen isotopes range in mass number from 12 to 28.^[33]

Relative and absolute abundance of ¹⁶O is due to it being a principal product of stellar evolution and the fact that it is a primary isotope, meaning it can be made by stars that were initially made exclusively of hydrogen.^[34] Most ¹⁶O is synthesized at the end of the helium fusion process in stars; the triple-alpha reaction creates ¹²C, which captures an additional ⁴He to make ¹⁶O. The neon burning process creates additional ¹⁶O.^[34]

Both ¹⁷O and ¹⁸O are secondary isotopes, meaning that their nucleosynthesis requires seed nuclei. ¹⁷O 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.^[34] Most ¹⁸O is produced when ¹⁴N (made abundant from CNO burning) captures a ⁴He nuclei, making ¹⁸O common in the helium-rich zones of stars.^[34] A billion degrees Celsius are required for two oxygen nuclei to undergo nuclear fusion to form the heavier nucleus of sulfur.^[2]

Fourteen radioisotopes have been characterized, with the most stable being ¹⁵O with a half-life of 122.24 s and ¹⁴O with a half-life of 70.606 s.^[33] All of the remaining radioactive isotopes have half-lives that are less than 27 s and the majority of these have half lives that are less than 83 milliseconds.^[33] The most common decay mode before the stable isotopes is electron capture and the most common mode after is beta decay. The decay products before the stable isotopes are element 7 (nitrogen) isotopes and the products after are element 9 (fluorine) isotopes.^[33]

An atomic mass of 16 was assigned to oxygen prior to the definition of the unified atomic mass unit based upon ¹²C.^[35] Since physicists referred to ¹⁶O only, while chemists meant the naturally abundant mixture of isotopes, this led to slightly different atomic mass scales.

The isotopic composition of oxygen atoms in the earth's atmosphere is 99.759% ¹⁶O, 0.037% ¹⁷O and 0.204% ¹⁸O.^[28] Because water molecules containing the lighter isotope are slightly more likely to evaporate and fall as precipitation^[36], fresh water and polar ice on earth contains slightly less (0.1981%) of the heavy isotope ¹⁸O than air (0.204%) or seawater containing (0.1995%). Since the ¹⁸O/¹⁶O isotope ratio in marine calcium carbonate equilibrates with that in the atmosphere, fluctuations in the oxygen isotope ratio in foraminifera can be used as a climate proxy, increasing during accumulation of polar ice and decreasing during warmer periods.

Oxygen evolution by water oxidation during photosynthesis. The jagged lines represent four photons oxidizing the central cluster of the oxygen evolving complex by exciting and removing four electrons through a cycle of S-states.

In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis in cyanobacteria, green algae and plants.^[37] Algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth.^[38] The remainder is produced by terrestrial plants, although almost all oxygen produced in tropical forests is consumed by organisms in those forests.^[39]

The overall formula for photosynthesis is:

6CO₂ + 6H₂O + sunlight C₆H₁₂O₆ + 6O₂

Or simply: carbon dioxide + water + sunlight glucose + oxygen

Photolytic oxygen evolution part of photosynthesis occurs via the light-dependent oxidation of water to molecular oxygen and can be written as the following simplified chemical reaction:

2H₂O 4e^- + 4H⁺ + O₂

The reaction requires the energy of four photons. The electrons from the oxidized water molecules replace electrons in the P₆₈₀ component of photosystem II which have been removed into an electron transport chain via light-dependent excitation and resonance energy transfer onto plastoquinone.^[40] Photosytem II therefore has also been referred to as water-plastoquinone oxido-reductase.^[41] The protons are released into the thylakoid lumen, thus contributing to the generation of a proton gradient across the thylakoid membrane. This proton gradient is the driving force for ATP synthesis via photophosphorylation and coupling the absorption of light energy and photolysis of water to the creation of chemical energy during photosynthesis.^[40]

Water oxidation is catalyzed by a manganese-containing enzyme complex associated with thylakoid membranes known as the oxygen evolving complex (OEC) or water-splitting complex. Manganese is an important cofactor, and calcium and chloride are also required for the reaction to occur.^[40]

In all vertebrates, the heme group of hemoglobin binds the oxygen dissolved in the blood.

See also: Aerobic respiration

DNA and proteins contain oxygen and the element is found in almost all molecules that are important to life. Molecular oxygen, O₂, is essential for cellular respiration in all aerobic organisms. Vertebrate animals use hemoglobin in their blood to transport oxygen from their lungs to their tissues, but other animals use hemocyanin (molluscs and some arthropods) or hemerythrin (spiders and lobsters).^[29] A liter of blood can dissolve 200 cc of oxygen gas, which is much more than water can dissolve (see Physical Properties).^[29]

In vertebratess, oxygen uptake is carried out by the following processes:

Oxygen diffuses through membranes and into red blood cells after inhalation into the lungs. The Heme group of hemoglobin by now already has carbon dioxide in its active site, but releases it for exhalation when oxygen is present. After being carried in blood to a body tissue in need of oxygen, it is handed-off from the Heme group to monooxygenase, an enzyme that also has an active site with an atom of iron.^[29] Monooxygenase uses oxygen to catalyze many oxidation reactions in the body. Oxygen is also used as an electron acceptor in mitochondria to generate chemical energy in the form of adenosine triphosphate (ATP) during oxidative phosphorylation. Carbon dioxide, one of the waste products produced, is released from the cell and into the blood, where it combines with empty Heme groups. Blood circulates back to the lungs and the process repeats.^[42] On average, a given oxygen atom cycles through the respiration pathway once every 3,000 years.^[9]

Reactive oxygen species are dangerous by-products that sometimes result from the use of oxygen in organisms. Important examples include; oxygen free radicals such as the highly dangerous superoxide O₂^-, and the less harmful hydrogen peroxide ( H₂O₂).^[29] The body uses superoxide dismutase to reduce superoxide radicals to hydrogen peroxide. Glutathione peroxidase and similar enzymes, then convent the H₂O₂ to water and dioxygen.^[29]

Parts of the immune system of higher organisms, however, create peroxide, superoxide and singlet oxygen to destroy invading microbes. Recently, singlet oxygen has been found to be a source of biologically-produced ozone: this reaction proceeds through an unusual compound dihydrogen trioxide, also known as trioxidane, (HOOOH) which is an antibody-catalyzed product of singlet oxygen and water. This compound in turn disproportionates to ozone and peroxide, providing two powerful antibacterials. The body's range of defense against all of these active oxidizing agents is hardly surprising, then, given their "deliberate" employment as antimicrobial agents in the immune response.^[43]

O₂ build-up in earth's atmosphere: 1) no O₂ produced, 2) O₂ produced, but absorbed in oceans & seabed rock, 3) O₂ starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer, 4-5) O₂ sinks filled and the gas accumulates

Oxygen was almost nonexistent in earth's atmosphere before the evolution of water oxidation in photosynthetic bacteria. Free oxygen first appeared in significant quantities during the Paleoproterozoic era (between 2.5 billion years ago and 1.6 billion years ago) as a product of the metabolic action of early anaerobes (archaea and bacteria). These organisms developed the mechanism of oxygen evolution between 3.5 and 2.7 billion years ago. At first, the produced oxygen dissolved in the oceans and reacted with iron, creating banded iron formations. It started to gas out of the oxygen-saturated waters about 2.7 billion years ago as evident in the rusting of iron-rich terrestrial rocks starting around that time. The amount of oxygen in the atmosphere increased gradually at first and shot up rapidly around 2.2 to 1.7 billion years ago to about 10% of its present level.^[44]

The development of an oxygen-rich atmosphere was one of the most important events in the history of life on earth. The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the anaerobic organisms then living to extinction during the oxygen catastrophe about 2.4 billion years ago. However, the high electronegativity of O₂ creates a large potential energy drop for cellular respiration, thus enabling organisms using aerobic respiration to produce much more ATP than anaerobic organisms. This makes them so efficient that they have come to dominate earth's biosphere.^[45] Photosynthesis and cellular respiration of oxygen allowed for the evolution of eukaryotic cells and ultimately complex multicellular organisms such as plants and animals.

The atmospheric abundance of free oxygen in later geological epochs and its gradual increase up to the present has been largely due to synthesis by photosynthetic organisms. Over the past 500 million years, oxygen levels fluctuated between 15 and 30% per volume.^[46] Towards the end of the Carboniferous era (coal age) about 300 million years ago, atmospheric oxygen levels reached a maximum of 35% by volume,^[46] allowing insects and amphibians with limiting respiratory systems to grow much larger than today's species. Today, oxygen is the second most common component of the earth's atmosphere (about 21% by volume) after nitrogen. Human activities, including the burning of 7 billion tonnes of fossil fuels each year have had very little effect on the amount of free oxygen in the atmosphere.^[9] It was estimated that at the current rate of photosynthesis, it would take about 2,000 years to regenerate the entire oxygen in the present atmosphere.^[47]

See also: Oxygen evolution and fractional distillation

Two major methods are employed to produce the 100 million tonnes of oxygen extracted from air for industrial uses annually.^[48] The most common method is to fractionally distill liquefied air into its various components, with nitrogen distilling as a vapor while oxygen is left as a liquid.^[48]

Hoffman electrolysis apparatus used in electrolysis of water

The other major method of producing oxygen involves passing a stream of clean, dry air through one bed of a pair of identical zeolite molecular sieves, which absorbs the nitrogen and delivers a gas stream that is 90 to 93% oxygen.^[48] Simultaneously, nitrogen is released from the other nitrogen-saturated zeolite bed, by reducing the chamber operating pressure and diverting part of the oxygen from the producer bed through it, in the reverse direction of flow. After a set cycle time, the operation of the two beds is switched, so that the producer bed is reverse purged and the purged bed becomes the producer bed. This allows for a continuous supply of gaseous oxygen to be pumped through a pipeline. It is known as pressure swing adsorption (PSA). Oxygen is increasingly obtained by these non-cryogenic technologies (see also the related vacuum-pressure swing adsorption (VPSA),^[49] or vacuum swing adsorption (VSA) technolgies).

Oxygen can also be produced through electrolysis of water into oxygen and hydrogen. A similar method is the electrocatalytic oxygen evolution from oxides and oxoacids. Chemical catalysts can be used as well, such as in chemical oxygen generators or oxygen candles that are used as part of the life support equipment on submarines, and which are still part of standard equipment on commercial airliners in case of depressurization emergencies.

Another air separation technology involves forcing air to dissolve through ceramic membranes based on zirconium oxide by either high pressure or an electric current, to produce nearly pure oxygen.^[6]

In large quantities, the price of liquid oxygen (2001) is approximately $0.21/kg.^[50] Since the primary cost of production is the energy cost of liquefying the air, the production cost will change as energy cost varies.

Oxygen is often transported in bulk as a liquid in specially insulated tankers because one liter of liquefied oxygen is equivalent to 840 liters of gaseous oxygen, at atmospheric pressure and 20°C.^[48] Such tankers are used to refill bulk liquid oxygen storage containers, which stand outside hospitals and other institutions with a need for large volumes of pure oxygen. Liquid oxygen is passed through heat-exchangers, which convert the cryogenic liquid into gas before it enters the building. Oxygen is also stored and shipped in smaller cylinders containing the compressed gas; a form that is useful in certain portable medical applications and Oxy-fuel welding and cutting.^[48]

See also: Breathing gas, Redox, and Combustion

An oxygen concentrator in an emphysema patient's house

Uptake of oxygen from the air is the essential purpose of respiration, so oxygen supplementation is used in medicine. Oxygen therapy is used to treat emphysema, pneumonia, some heart disorders and any disease that impairs the body's ability to take up and use oxygen.^[51] Treatments are flexible enough to be used in hospitals, the patient's home, or increasingly by portable devices. Oxygen tents were once commonly used in oxygen supplementation, but have since been mostly replaced by the use of oxygen masks or nasal cannulas. Hyperbaric medicine uses hyperbaric oxygen chambers to increase the partial pressure of oxygen around the patient and, when needed, the medical staff.

Carbon monoxide poisoning, gas gangrene and decompression sickness (the "bends") are sometimes treated using these devices. Increased oxygen concentration in the lungs helps to displace carbon monoxide from the heme group of hemoglobin. Oxygen is poisonous to the anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them. Decompression sickness occurs in divers who decompress too quickly after a dive, resulting in bubbles of inert gas, mostly nitrogen and argon, forming in their blood. Increasing the pressure of oxygen as soon as possible is part of the treatment.^[51]

Low pressure pure oxygen is used in spacesuits

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

Scuba divers and submariners also rely on artificially-delivered oxygen, but most often use normal pressure and air mixtures. Deep sea diving, however, requires alternate mixtures of oxygen with other gases, such as helium, to help prevent oxygen toxicity, nitrogen narcosis and other diving hazards.

People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental oxygen supplies.^[52] Passengers traveling in commercial airplanes have an emergency supply of oxygen automatically supplied to them in case of cabin depressurization. Sudden cabin pressure loss activates chemical oxygen generators above each seat, causing oxygen masks to drop and forcing iron fillings into the sodium chlorate inside the canister.^[6] A steady stream of oxygen gas is produced by the exothermic reaction.^[53]

Oxygen, as a supposed mild euphoric, has a history of recreational use in oxygen bars and in sports. Oxygen bars are establishments, found in Japan, California and Las Vegas, Nevada since the late 1990s that offer higher than normal oxygen exposure for a fee.^[54] Professional athletes, especially in American football, also sometimes go off field between plays to wear oxygen masks in order to get a supposed "boost" in performance. However, the reality of a pharmacological effect is doubtful; a placebo or psychological boost being the most plausible explanation.^[54]

Most commercially produced oxygen is used to smelt iron into steel

Smelting of iron ore into steel consumes 55% of commercially produced oxygen.^[6] In this process, oxygen is injected through a high-pressure lance into molten iron, which removes sulfur impurities and excess carbon as the respective oxides, SO₂ and CO₂. The reactions are exothermic, so the temperature increases to 1700° C.^[6]

Another 25% of commercially produced oxygen is used by the chemical industry.^[6] Ethylene is reacted with oxygen 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 plastics and fabrics).^[6]

Most of the remaining 20% of commercially produced oxygen is used in medical applications, metal cutting and welding, as an oxidizer in rocket fuel, and in water treatment.^[6] Oxygen is used in oxyacetylene welding burning acetylene with oxygen to produce a very hot flame. In this process, metal up to 60  cm thick is first heated with a small oxy-acetylene flame and then quickly cut by a large stream of oxygen.^[55] Rocket propulsion requires a fuel and an oxidizer. Larger rockets use liquid oxygen as their oxidizer, which is mixed and ignited with the fuel for propulsion.

500 million years of climate change vs ¹⁸O

Paleoclimatologists measure the ratio of oxygen-18 and oxygen-16 in the shells and skeletons of marine organisms to determine what the climate was like millions of years ago. During periods of lower global temperatures, sea water molecules that contain the lighter isotope, oxygen-16, evaporate at a slightly faster rate than water molecules containing the 12% heavier oxygen-18.^[56] Snow and rain from that evaporated water tends to be enriched in oxygen-16 and the seawater left behind tends to be enriched in oxygen-18. Marine organisms then incorporate more oxygen-18 into their skeletons and shells than they would in a warmer climate.^[56] Paleoclimatologists also directly measure this ratio in the water molecules of ice core samples that are up to several hundreds of thousands of years old.

Oxygen presents two spectrophotometric absorption bands peaking at the wavelengths 687 and 760 nanometers. Some scientists have proposed to use the measurement of the radiance coming from vegetation canopies in those oxygen bands to characterize plant health status from a satellite platform.^[57] This is because in those bands, it is possible to discriminate the vegetation's reflectance from the vegetation's fluorescence, which is much weaker. The measurement presents several technical difficulties due to the low signal to noise ratio and due to the vegetation's architecture, but it has been proposed as a possibility to monitor the carbon cycle from satellites on a global scale.

Philo's experiment inspired later investigators

One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck.^[58] Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries later Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration.^[59]

In the late 17th century Robert Boyle proved that air is necessary for combustion. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called 'spiritus nitroaereus' or just 'nitroaereus'.^[60] 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.^[61] From this he surmised that nitroaereus is consumed in both respiration and combustion.

Mayow observed that antimony increased in weight when heated and inferred that the nitroaereus must have combined with it.^[60] He also thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body.^[60] Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione".^[61]

Georg Ernst Stahl helped develop and popularize the phlogiston theory

Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all also produced oxygen in experiments in the 17th century but none of them recognized it as an element.^[16] This was largely due to the prevalence of a philosophy of combustion and corrosion called the phlogiston theory, which was then the favored explanation of how those processes worked.

Established in 1667 by German alchemist J. J. Becher, and modified by chemist Georg Ernst Stahl by 1731,^[62] 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, its calx.^[59]

Highly combustible materials which left little residuum, such as wood or coal, were thought of as made mostly of phlogiston, while non-combustible substances which corroded, such as iron, contained very little. Air did not play a role in phlogiston theory and no initial quantitative experiments were conducted to test the idea; instead it was based on observations of what happened when something burned: that most common objects appeared to become lighter and seemed to lose