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Oxygen

January 26, 2023

This article is about the chemical element. For other uses, see Oxygen (disambiguation).

Oxygen is the chemical element with the symbol O and atomic number 8. It is a member of the chalcogen group in the periodic table, a highly reactive nonmetal, and an oxidizing agent that readily forms oxides with most elements as well as with other compounds. Oxygen is Earth's most abundant element, and after hydrogen and helium, it is the third-most abundant element in the universe. At standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O
2. Diatomic oxygen gas currently constitutes 20.95% of the Earth's atmosphere, though this has changed considerably over long periods of time. Oxygen makes up almost half of the Earth's crust in the form of oxides.[3]

Oxygen, 8O
Liquid oxygen boiling (O2)
Oxygen
AllotropesO2, O3 (ozone) and more (see Allotropes of oxygen)
Appearancegas: colorless
liquid and solid: pale blue
Standard atomic weight Ar°(O)
  • [15.99903, 15.99977]
  • 15.999±0.001 (abridged)[1]
Abundance
in the Earth's crust461000 ppm
Oxygen in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
–
↑
O
↓
S
nitrogen ← oxygen → fluorine
Atomic number (Z)8
Groupgroup 16 (chalcogens)
Periodperiod 2
Block  p-block
Electron configuration[He] 2s2 2p4
Electrons per shell2, 6
Physical properties
Phase at STPgas
Melting point(O2) 54.36 K ​(−218.79 °C, ​−361.82 °F)
Boiling point(O2) 90.188 K ​(−182.962 °C, ​−297.332 °F)
Density (at STP)1.429 g/L
when liquid (at b.p.)1.141 g/cm3
Triple point54.361 K, ​0.1463 kPa
Critical point154.581 K, 5.043 MPa
Heat of fusion(O2) 0.444 kJ/mol
Heat of vaporization(O2) 6.82 kJ/mol
Molar heat capacity(O2) 29.378 J/(mol·K)
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K)       61 73 90
Atomic properties
Oxidation states−2, −1, 0, +1, +2
ElectronegativityPauling scale: 3.44
Ionization energies
  • 1st: 1313.9 kJ/mol
  • 2nd: 3388.3 kJ/mol
  • 3rd: 5300.5 kJ/mol
  • (more)
Covalent radius66±2 pm
Van der Waals radius152 pm
Spectral lines of oxygen
Other properties
Natural occurrenceprimordial
Crystal structure ​cubic
Speed of sound330 m/s (gas, at 27 °C)
Thermal conductivity26.58×10−3  W/(m⋅K)
Magnetic orderingparamagnetic
Molar magnetic susceptibility+3449.0×10−6 cm3/mol (293 K)[2]
CAS Number7782-44-7
History
DiscoveryCarl Wilhelm Scheele (1771)
Named byAntoine Lavoisier (1777)
Main isotopes of oxygen
Iso­tope Decay
abun­dance half-life (t1/2) mode pro­duct
15O syn 122.266 s β+ 15N
16O 99.757% stable
17O 0.038% stable
18O 0.205% stable
 Category: Oxygen
| references

Many major classes of organic molecules in living organisms contain oxygen atoms, such as proteins, nucleic acids, carbohydrates, and fats, as do the major constituent inorganic compounds 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 is continuously replenished in Earth's atmosphere by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide. Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form (allotrope) of oxygen, ozone (O
3), strongly absorbs ultraviolet UVB radiation and the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation. However, ozone present at the surface is a byproduct of smog and thus a pollutant.

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. The name oxygen was coined in 1777 by Antoine Lavoisier, who first recognized oxygen as a chemical element and correctly characterized the role it plays in combustion.

Common uses of oxygen include production of steel, plastics and textiles, brazing, welding and cutting of steels and other metals, rocket propellant, oxygen therapy, and life support systems in aircraft, submarines, spaceflight and diving.

History of study

Early experiments

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.[4] 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.[5]

In the late 17th century, Robert Boyle proved that air is necessary for combustion. English chemist John Mayow (1641–1679) refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus.[6] 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.[7] 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.[6] 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.[6] Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione".[7]

Phlogiston theory

Main article: 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.[8] 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.[9]

Established in 1667 by the German alchemist J. J. Becher, and modified by the chemist Georg Ernst Stahl by 1731,[10] 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.[5]

Highly combustible materials that leave little 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.[5]

Discovery

 
Joseph Priestley is usually given priority in the discovery.

Polish alchemist, philosopher, and physician Michael Sendivogius (Michał Sędziwój) in his work De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti (1604) described a substance contained in air, referring to it as 'cibus vitae' (food of life,[11]) and according to Polish historian Roman Bugaj, this substance is identical with oxygen.[12] Sendivogius, during his experiments performed between 1598 and 1604, properly recognized that the substance is equivalent to the gaseous byproduct released by the thermal decomposition of potassium nitrate. In Bugaj's view, the isolation of oxygen and the proper association of the substance to that part of air which is required for life, provides sufficient evidence for the discovery of oxygen by Sendivogius.[12] This discovery of Sendivogius was however frequently denied by the generations of scientists and chemists which succeeded him.[11]

It is also commonly claimed that oxygen was first discovered by Swedish pharmacist Carl Wilhelm Scheele. He had produced oxygen gas by heating mercuric oxide (HgO) and various nitrates in 1771–72.[13][14][5] Scheele called the gas "fire air" because it was then the only known agent to support combustion. He wrote an account of this discovery in a manuscript titled Treatise on Air and Fire, which he sent to his publisher in 1775. That document was published in 1777.[15]

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".[14] 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 common air, but I fancied that my breast felt peculiarly light and easy for some time afterwards."[8] 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.[5][16] Because he published his findings first, Priestley is usually given priority in the discovery.

The French chemist Antoine Laurent 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.)[15]

Lavoisier's contribution

 
Antoine Lavoisier discredited the phlogiston theory.

Lavoisier conducted the first adequate quantitative experiments on oxidation and gave the first correct explanation of how combustion works.[14] He used these and similar experiments, all started in 1774, to discredit the phlogiston theory and to prove that the substance discovered by Priestley and Scheele was a chemical element.

In one experiment, Lavoisier observed that there was no overall increase in weight when tin and air were heated in a closed container.[14] He noted that air rushed in when he opened the container, which indicated that part of the trapped air had been consumed. He also noted that the tin had increased in weight and that increase was the same as the weight of the air that rushed back in. This and other experiments on combustion were documented in his book Sur la combustion en général, which was published in 1777.[14] In that work, he proved that air is a mixture of two gases; 'vital air', which is essential to combustion and respiration, and azote (Gk. ἄζωτον "lifeless"), which did not support either. Azote later became nitrogen in English, although it has kept the earlier name in French and several other European languages.[14]

Etymology

Lavoisier renamed 'vital air' to oxygène in 1777 from the Greek roots ὀξύς (oxys) (acid, literally "sharp", from the taste of acids) and -γενής (-genēs) (producer, literally begetter), because he mistakenly believed that oxygen was a constituent of all acids.[17] Chemists (such as Sir Humphry Davy in 1812) eventually determined that Lavoisier was wrong in this regard, but by then the name was too well established.[18]

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.[15]

Later history

 
Robert H. Goddard and a liquid oxygen-gasoline rocket

John Dalton's original 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.[19] 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.[20][a]

The first commercial method of producing oxygen was chemical, the so-called Brin process involving a reversible reaction of barium oxide. It was invented in 1852 and commercialized in 1884, but was displaced by newer methods in early 20th century.

By the late 19th century scientists realized that air could be liquefied and its components isolated by compressing and cooling it. Using a cascade method, Swiss chemist and physicist Raoul Pierre Pictet 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.[21] Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen.[21] 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.[22]

 
An experiment setup for preparation of oxygen in academic laboratories

In 1891 Scottish chemist James Dewar was able to produce enough liquid oxygen for study.[23] 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 distilled the component gases by boiling them off one at a time and capturing them separately.[24] Later, in 1901, oxyacetylene welding was demonstrated for the first time by burning a mixture of acetylene and compressed O
2. This method of welding and cutting metal later became common.[24]

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 m at 97 km/h on March 16, 1926, in Auburn, Massachusetts, US.[24][25]

In academic laboratories, oxygen can be prepared by heating together potassium chlorate mixed with a small proportion of manganese dioxide.[26]

Oxygen levels in the atmosphere are trending slightly downward globally, possibly because of fossil-fuel burning.[27]

Characteristics

Properties and molecular structure

 
Orbital diagram, after Barrett (2002),[28] showing the participating atomic orbitals from each oxygen atom, the molecular orbitals that result from their overlap, and the 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* and π and π* orbital pairs).

At standard temperature and pressure, oxygen is a colorless, odorless, and tasteless gas with the molecular formula O
2, referred to as dioxygen.[29]

As dioxygen, two oxygen atoms are 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 orbitals 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, filling of orbitals, and the resulting cancellation of contributions from the 2s electrons, after sequential filling of the low σ and σ* 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* orbitals.[28]

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—is a configuration termed a spin triplet state. Hence, the ground state of the O
2 molecule is referred to as triplet oxygen.[30][b] 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.[31]

 
Liquid oxygen, temporarily suspended in a magnet owing to its paramagnetism

In the triplet form, O
2 molecules are paramagnetic. That is, they impart magnetic character to oxygen when it is in the presence of a magnetic field, because of the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O
2 molecules.[23] Liquid oxygen is so magnetic that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.[32][c]

Singlet oxygen is a name given to several higher-energy species of molecular O
2 in which all the electron spins are paired. It is much more reactive with common organic molecules than is normal (triplet) molecular oxygen. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.[33] It is also produced in the troposphere by the photolysis of ozone by light of short wavelength[34] and by the immune system as a source of active oxygen.[35] Carotenoids 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.[36]

Allotropes

Main article: Allotropes of oxygen
 
Space-filling model representation of dioxygen (O2) molecule

The common allotrope of elemental oxygen on Earth is called dioxygen, O
2, the major part of the Earth's atmospheric oxygen (see Occurrence). O2 has a bond length of 121 pm and a bond energy of 498 kJ/mol.[37] O2 is used by complex forms of life, such as animals, in cellular respiration. Other aspects of O
2 are covered in the remainder of this article.

Trioxygen (O
3) is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue.[38] Ozone is produced in the upper atmosphere when O
2 combines with atomic oxygen made by the splitting of O
2 by ultraviolet (UV) radiation.[17] Since ozone absorbs strongly in the UV region of the spectrum, the ozone layer of the upper atmosphere functions as a protective radiation shield for the planet.[17] Near the Earth's surface, it is a pollutant formed as a by-product of automobile exhaust.[38] At low earth orbit altitudes, sufficient atomic oxygen is present to cause corrosion of spacecraft.[39]

The metastable molecule tetraoxygen (O
4) was discovered in 2001,[40][41] and was assumed to exist in one of the six phases of solid oxygen. It was proven in 2006 that this phase, created by pressurizing O
2 to 20 GPa, is in fact a rhombohedral O
8 cluster.[42] This cluster has the potential to be a much more powerful oxidizer than either O
2 or O
3 and may therefore be used in rocket fuel.[40][41] A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa[43] and it was shown in 1998 that at very low temperatures, this phase becomes superconducting.[44]

Physical properties

 
Oxygen discharge (spectrum) tube
See also: Liquid oxygen and solid oxygen

Oxygen dissolves more readily in water than in nitrogen, and in freshwater more readily than in seawater. Water in equilibrium with air contains approximately 1 molecule of dissolved O
2 for every 2 molecules of 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 (14.6 mg/L) dissolves at 0 °C than at 20 °C (7.6 mg/L).[8][45] At 25 °C and 1 standard atmosphere (101.3 kPa) of air, freshwater can dissolve about 6.04 milliliters (mL) of oxygen per liter, and seawater contains about 4.95 mL per liter.[46] At 5 °C the solubility increases to 9.0 mL (50% more than at 25 °C) per liter for freshwater and 7.2 mL (45% more) per liter for sea water.

Oxygen gas dissolved in water at sea-level
(milliliters per liter)
5 °C 25 °C
Freshwater 9.00 6.04
Seawater 7.20 4.95

Oxygen condenses at 90.20 K (−182.95 °C, −297.31 °F) and freezes at 54.36 K (−218.79 °C, −361.82 °F).[47] Both liquid and solid O
2 are clear substances with a light sky-blue color caused by absorption in the red (in contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light). High-purity liquid O
2 is usually obtained by the fractional distillation of liquefied air.[48] Liquid oxygen may also be condensed from air using liquid nitrogen as a coolant.[49]

Liquid oxygen is a highly reactive substance and must be segregated from combustible materials.[49]

The spectroscopy of molecular oxygen is associated with the atmospheric processes of aurora and airglow.[50] 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.[51] Excited-state singlet molecular oxygen is responsible for red chemiluminescence in solution.[52]

Table of thermal and physical properties of oxygen (O2) at atmospheric pressure:[53][54]

Temperature (K) Density (kg/m^3) Specific heat (kJ/kg °C) Dynamic viscosity (kg/m s) Kinematic viscosity (m^2/s) Thermal conductivity (W/m °C) Thermal diffusivity (m^2/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

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 three stable isotopes, 16O, 17O, and 18O, with 16O being the most abundant (99.762% natural abundance).[55]

Most 16O is synthesized at the end of the helium fusion process in massive stars but some is made in the neon burning process.[56] 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.[56] Most 18O is produced when 14N (made abundant from CNO burning) captures a 4He nucleus, making 18O common in the helium-rich zones of evolved, massive stars.[56]

Thirteen radioisotopes have been characterized, ranging from 11O to 26O.[57] The most stable are 15O with a half-life of 122.24 seconds and 14O with a half-life of 70.606 seconds.[55] 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.[55] The most common decay mode of the isotopes lighter than 16O is β+ decay[58][59][60] to yield nitrogen, and the most common mode for the isotopes heavier than 18O is beta decay to yield fluorine.[55]

Occurrence

See also: Silicate minerals, Category:Oxide minerals, Stellar population, Cosmochemistry, and Astrochemistry
Ten most common elements in the Milky Way Galaxy estimated spectroscopically[61]
Z Element Mass fraction in parts per million
1 Hydrogen 739,000 71 × mass of oxygen (red bar)
2 Helium 240,000 23 × mass of oxygen (red bar)
8 Oxygen 10,400 10400
 
6 Carbon 4,600 4600
 
10 Neon 1,340 1340
 
26 Iron 1,090 1090
 
7 Nitrogen 960 960
 
14 Silicon 650 650
 
12 Magnesium 580 580
 
16 Sulfur 440 440
 

Oxygen is the most abundant chemical element by mass in the Earth's biosphere, air, sea and land. Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.[62] About 0.9% of the Sun's mass is oxygen.[14] Oxygen constitutes 49.2% of the Earth's crust by mass[63] as part of oxide compounds such as silicon dioxide and is the most abundant element by mass in the Earth's crust. It is also the major component of the world's oceans (88.8% by mass).[14] 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 1015 tonnes).[14][64][d] 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% O
2 by volume) and Venus have much less. The O
2 surrounding those planets is produced solely by the action of ultraviolet radiation on oxygen-containing molecules such as carbon dioxide.

 
Cold water holds more dissolved O
2.

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 respiration, decay, and combustion remove it from the atmosphere. In the present equilibrium, production and consumption occur at the same rate.[65]

Free oxygen also occurs in solution in the world's water bodies. The increased solubility of O
2 at lower 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.[66] Water polluted with plant nutrients such as nitrates or phosphates may stimulate growth of algae by a process called eutrophication and the decay of these organisms and other biomaterials may reduce the O
2 content in eutrophic water bodies. Scientists assess this aspect of water quality by measuring the water's biochemical oxygen demand, or the amount of O
2 needed to restore it to a normal concentration.[67]

Analysis

 
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 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.[68] 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.[68] Paleoclimatologists also directly measure this ratio in the water molecules of ice core samples as old as hundreds of thousands of years.

Planetary geologists have measured the relative quantities of oxygen isotopes in samples from the Earth, the Moon, Mars, and meteorites, but were long unable to obtain reference values for the isotope ratios in the Sun, believed to be the same as those of the primordial solar nebula. Analysis of a silicon wafer exposed to the solar wind in space and returned by the crashed Genesis spacecraft has shown that the Sun has a higher proportion of oxygen-16 than does the Earth. The measurement implies that an unknown process depleted oxygen-16 from the Sun's disk of protoplanetary material prior to the coalescence of dust grains that formed the Earth.[69]

Oxygen presents two spectrophotometric absorption bands peaking at the wavelengths 687 and 760 nm. Some remote sensing scientists have proposed using the measurement of the radiance coming from vegetation canopies in those bands to characterize plant health status from a satellite platform.[70] 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.

Biological production and role of O2

Main article: Dioxygen in biological reactions

Photosynthesis and respiration

 
Photosynthesis splits water to liberate O
2 and fixes CO
2 into sugar in what is called a Calvin cycle.

In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis. According to some estimates, green algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on Earth, and the rest is produced by terrestrial plants.[71] 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.[72]

A simplified overall formula for photosynthesis is[73]

6 CO2 + 6 H
2O + photons → C
6H
12O
6 + 6 O
2

or simply

carbon dioxide + water + sunlight → glucose + dioxygen

Photolytic oxygen evolution occurs in the thylakoid membranes of photosynthetic organisms and requires the energy of four photons.[e] 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.[74] The O
2 remaining (after production of the water molecule) is released into the atmosphere.[f]

Oxygen is used in mitochondria in the generation of ATP during oxidative phosphorylation. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as

C
6H
12O
6 + 6 O
2 → 6 CO2 + 6 H
2O + 2880 kJ/mol

In vertebrates, O
2 diffuses through membranes in the lungs and into red blood cells. Hemoglobin binds O
2, changing color from bluish red to bright red[38] (CO
2 is released from another part of hemoglobin through the Bohr effect). Other animals use hemocyanin (molluscs and some arthropods) or hemerythrin (spiders and lobsters).[64] A liter of blood can dissolve 200 cm3 of O
2.[64]

Until the discovery of anaerobic metazoa,[75] oxygen was thought to be a requirement for all complex life.[76]

Reactive oxygen species, such as superoxide ion (O−
2) and hydrogen peroxide (H
2O
2), are reactive by-products of oxygen use in organisms.[64] 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.[74] Oxygen is damaging to obligately anaerobic organisms, which were the dominant form of early life on Earth until 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.[77][78]

An adult human at rest inhales 1.8 to 2.4 grams of oxygen per minute.[79] This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year.[g]

Living organisms

Partial pressures of oxygen in the human body (PO2)
Unit Alveolar pulmonary
gas pressures		 Arterial blood oxygen Venous blood gas
kPa 14.2 11[h]-13[h] 4.0[h]-5.3[h]
mmHg 107 75[80]-100[80] 30[81]-40[81]

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.[82]

Build-up in the atmosphere

Main article: Geological history of oxygen
 
O
2 build-up in Earth's atmosphere: 1) no O
2 produced; 2) O
2 produced, but absorbed in oceans & seabed rock; 3) O
2 starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer; 4–5) O
2 sinks filled and the gas accumulates

Free oxygen gas was almost nonexistent in Earth's atmosphere before photosynthetic archaea and bacteria evolved, probably about 3.5 billion years ago. Free oxygen first appeared in significant quantities during the Paleoproterozoic eon (between 3.0 and 2.3 billion years ago).[83] Even if there was much dissolved iron in the oceans when oxygenic photosynthesis was getting more common, it appears the banded iron formations were created by anoxyenic or micro-aerophilic iron-oxidizing bacteria which dominated the deeper areas of the photic zone, while oxygen-producing cyanobacteria covered the shallows.[84] Free oxygen began to outgas from the oceans 3–2.7 billion years ago, reaching 10% of its present level around 1.7 billion years ago.[83][85]

The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the extant anaerobic organisms to extinction during the Great Oxygenation Event (oxygen catastrophe) about 2.4 billion years ago. Cellular respiration using O
2 enables aerobic organisms to produce much more ATP than anaerobic organisms.[86] Cellular respiration of O
2 occurs in all eukaryotes, including all complex multicellular organisms such as plants and animals.

Since the beginning of the Cambrian period 540 million years ago, atmospheric O
2 levels have fluctuated between 15% and 30% by volume.[87] Towards the end of the Carboniferous period (about 300 million years ago) atmospheric O
2 levels reached a maximum of 35% by volume,[87] which may have contributed to the large size of insects and amphibians at this time.[88]

Variations in atmospheric oxygen concentration have shaped past climates. When oxygen declined, atmospheric density dropped, which in turn increased surface evaporation, causing precipitation increases and warmer temperatures.[89]

At the current rate of photosynthesis it would take about 2,000 years to regenerate the entire O
2 in the present atmosphere.[90]

Extraterrestrial free oxygen

Main article: Extraterrestrial atmosphere

In the field of astrobiology and in the search for extraterrestrial life oxygen is a strong biosignature. That said it might not be a definite biosignature, being possibly produced abiotically on celestial bodies with processes and conditions (such as a peculiar hydrosphere) which allow free oxygen,[91][92][93] like with Europa's and Ganymede's thin oxygen atmospheres.[94]

Industrial production

See also: Air separation, Oxygen evolution, and Fractional distillation
 
Hofmann electrolysis apparatus used in electrolysis of water.

One hundred million tonnes of O
2 are extracted from air for industrial uses annually by two primary methods.[15] The most common method is fractional distillation of liquefied air, with N
2 distilling as a vapor while O
2 is left as a liquid.[15]

The other primary method of producing 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% to 93% O
2.[15] Simultaneously, nitrogen gas is released from the other nitrogen-saturated zeolite bed, by reducing the chamber operating pressure and diverting part of the oxygen gas from the producer bed through it, in the reverse direction of flow. After a set cycle time the operation of the two beds is interchanged, thereby allowing for a continuous supply of gaseous oxygen to be pumped through a pipeline. This is known as pressure swing adsorption. Oxygen gas is increasingly obtained by these non-cryogenic technologies (see also the related vacuum swing adsorption).[95]

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 O
2 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 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 O
2 gas.[67]

Storage

 
Oxygen and MAPP gas compressed-gas cylinders with regulators

Oxygen storage methods include high-pressure oxygen tanks, cryogenics and chemical compounds. For reasons of economy, oxygen is often transported in bulk as a liquid in specially insulated tankers, since one liter of liquefied oxygen is equivalent to 840 liters of gaseous oxygen at atmospheric pressure and 20 °C (68 °F).[15] 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 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.[15]

Applications

See also: Breathing gas, Redox, and Combustion

Medical

 
An oxygen concentrator in an emphysema patient's house
Main article: Oxygen therapy

Uptake of O
2 from the air is the essential purpose of 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.[96]

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 replaced mostly by the use of oxygen masks or nasal cannulas.[97]

Hyperbaric (high-pressure) medicine uses special oxygen chambers to increase the partial pressure of O
2 around the patient and, when needed, the medical staff.[98] Carbon monoxide poisoning, gas gangrene, and decompression sickness (the 'bends') are sometimes addressed with this therapy.[99] Increased O
2 concentration in the lungs helps to displace carbon monoxide from the heme group of hemoglobin.[100][101] Oxygen gas is poisonous to the anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them.[102][103] 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 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.[96][104][105] 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.[106][107][108]

Life support and recreational use

 
Low-pressure pure O
2 is used in space suits.

An application of O
2 as a low-pressure breathing gas is in modern space suits, which surround their occupant's body with the breathing gas. These devices use nearly pure oxygen at about one-third normal pressure, resulting in a normal blood partial pressure of O
2. This trade-off of higher oxygen concentration for lower pressure is needed to maintain suit flexibility.[109][110]

Scuba and surface-supplied underwater divers and submariners also rely on artificially delivered 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 O
2 use in diving at pressures higher than atmospheric is usually limited to rebreathers, or decompression at relatively shallow depths (~6 meters depth, or less),[111][112] or 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 O
2 with other gases, such as nitrogen or helium, to prevent oxygen toxicity.[111]

People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental O
2 supplies.[i] Pressurized commercial airplanes have an emergency supply of O
2 automatically supplied to the passengers in case of cabin depressurization. Sudden cabin pressure loss activates chemical oxygen generators above each seat, causing oxygen masks 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.[67] A steady stream of oxygen gas is then produced by the exothermic reaction.

Oxygen, as a mild euphoric, has a history of recreational use in oxygen bars and in sports. Oxygen bars are establishments found in the United States since the late 1990s that offer higher than normal O
2 exposure for a minimal fee.[113] 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.[113] Available studies support a performance boost from oxygen enriched mixtures only if it is inhaled during aerobic exercise.[114]

Other recreational uses that do not involve breathing include pyrotechnic applications, such as George Goble's five-second ignition of barbecue grills.[115]

Industrial

 
Most commercially produced O
2 is used to smelt and/or decarburize iron.

Smelting of iron ore into steel consumes 55% of commercially produced oxygen.[67] In this process, O
2 is injected through a high-pressure lance into molten iron, which removes sulfur impurities and excess carbon as the respective oxides, SO
2 and CO
2. The reactions are exothermic, so the temperature increases to 1,700 °C.[67]

Another 25% of commercially produced oxygen is used by the chemical industry.[67] Ethylene is reacted with 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 plastics and fabrics).[67]

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.[67] Oxygen is used in oxyacetylene welding, burning acetylene with O
2 to produce a very hot flame. In this process, metal up to 60 cm (24 in) thick is first heated with a small oxy-acetylene flame and then quickly cut by a large stream of O
2.[116]

Compounds

Main article: Compounds of oxygen
 
Water (H
2O) is the most familiar oxygen compound.

The oxidation state of oxygen is −2 in almost all known compounds of oxygen. The oxidation state −1 is found in a few compounds such as peroxides.[117] 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).[118]

Oxides and other inorganic compounds

Water (H
2O) is an oxide of hydrogen and the most familiar oxygen compound. Hydrogen atoms are covalently bonded to oxygen in a water molecule but also have an additional attraction (about 23.3 kJ/mol per hydrogen atom) to an adjacent oxygen atom in a separate molecule.[119] These hydrogen bonds between water molecules hold them approximately 15% closer than what would be expected in a simple liquid with just van der Waals forces.[120][j]

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

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements to give corresponding oxides. 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 passivates the metal and slows further corrosion. Many oxides of the transition metals are non-stoichiometric compounds, with slightly less metal than the chemical formula would show. For example, the mineral FeO (wüstite) is written as  , where x is usually around 0.05.[121]

Oxygen is present in the atmosphere in trace quantities in the form of carbon dioxide (CO
2). The Earth's crustal rock is composed in large part of oxides of silicon (silica SiO
2, as found in granite and quartz), aluminium (aluminium oxide Al
2O
3, in bauxite and corundum), iron (iron(III) oxide Fe
2O
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 silicates (in silicate minerals). The Earth's mantle, of much larger mass than the crust, is largely composed of silicates of magnesium and iron.

Water-soluble silicates in the form of Na
4SiO
4, Na
2SiO
3, and Na
2Si
2O
5 are used as detergents and adhesives.[122]

Oxygen also acts as a ligand for transition metals, forming transition metal dioxygen complexes, which feature metal–O
2. This class of compounds includes the heme proteins hemoglobin and myoglobin.[123] An exotic and unusual reaction occurs with PtF
6, which oxidizes oxygen to give O2+PtF6−, dioxygenyl hexafluoroplatinate.[124]

Organic compounds

 
Acetone is an important feeder material in the chemical industry.
  Oxygen
  Carbon
  Hydrogen

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); and amides (R-C(O)-NR
2). There are many important organic solvents that contain oxygen, including: acetone, methanol, ethanol, isopropanol, furan, THF, diethyl ether, dioxane, ethyl acetate, DMF, DMSO, acetic acid, and formic acid. Acetone ((CH
3)
2CO) and phenol (C
6H
5OH) 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. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms. The element is similarly found in almost all biomolecules that are important to (or generated by) life.

Oxygen reacts spontaneously with many organic compounds at or below room temperature in a process called autoxidation.[125] Most of the organic compounds that contain oxygen are not made by direct action of O
2. Organic compounds important in industry and commerce that are made by direct oxidation of a precursor include ethylene oxide and peracetic acid.[122]

Safety and precautions

Oxygen
Hazards
GHS labelling:
 
H272
P220, P244, P370+P376, P403
NFPA 704 (fire diamond)
0
0
1

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.[126]

Toxicity

Main article: Oxygen toxicity
 
Main symptoms of oxygen toxicity[127]

Oxygen gas (O
2) can be toxic at elevated partial pressures, leading to convulsions and other health problems.[111][k][128] Oxygen toxicity usually begins to occur at partial pressures more than 50 kilopascals (kPa), equal to about 50% oxygen composition at standard pressure or 2.5 times the normal sea-level O
2 partial pressure of about 21 kPa. This is not a problem except for patients on mechanical ventilators, since gas supplied through oxygen masks in medical applications is typically composed of only 30–50% O
2 by volume (about 30 kPa at standard pressure).[8]

At one time, premature babies were placed in incubators containing O
2-rich air, but this practice was discontinued after some babies were blinded by the oxygen content being too high.[8]

Breathing pure O
2 in space applications, such as in some modern space suits, or in early spacecraft such as Apollo, causes no damage due to the low total pressures used.[109][129] In the case of spacesuits, the O
2 partial pressure in the breathing gas is, in general, about 30 kPa (1.4 times normal), and the resulting O
2 partial pressure in the astronaut's arterial blood is only marginally more than normal sea-level O
2 partial pressure.[130]

Oxygen toxicity to the lungs and central nervous system can also occur in deep scuba diving and surface supplied diving.[8][111] Prolonged breathing of an air mixture with an O
2 partial pressure more than 60 kPa can eventually lead to permanent pulmonary fibrosis.[131] Exposure to an O
2 partial pressures greater than 160 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% O
2 at 66 m (217 ft) or more of depth; the same thing can occur by breathing 100% O
2 at only 6 m (20 ft).[131][132][133][134]

Combustion and other hazards

 
The interior of the Apollo 1 Command Module. Pure O
2 at higher than normal pressure and a spark led to a fire and the loss of the Apollo 1 crew.

Highly concentrated sources of oxygen promote rapid combustion. Fire and explosion hazards exist when concentrated oxidants and fuels are brought into close proximity; an ignition event, such as heat or a spark, is needed to trigger combustion.[31] Oxygen is the oxidant, not the fuel.

Concentrated O
2 will allow combustion to proceed rapidly and energetically.[31] 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 O
2 systems requires special training to ensure that ignition sources are minimized.[31] The fire that killed the Apollo 1 crew in a launch pad test spread so rapidly because the capsule was pressurized with pure O
2 but at slightly more than atmospheric pressure, instead of the 1⁄3 normal pressure that would be used in a mission.[l][136]

Liquid oxygen spills, if allowed to soak into organic matter, such as wood, petrochemicals, and asphalt can cause these materials to detonate unpredictably on subsequent mechanical impact.[31]

See also

Notes

  1. ^ These results were mostly ignored until 1860. Part of this rejection was due to the belief that atoms of one element would have no chemical affinity towards atoms of the same element, and part was due to apparent exceptions to Avogadro's law that were not explained until later in terms of dissociating molecules.
  2. ^ An orbital is a concept from quantum mechanics that models an electron as a wave-like particle that has a spatial distribution about an atom or molecule.
  3. ^ Oxygen's paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine the purity of gaseous oxygen. (. Servomex. Archived from the original on March 8, 2008. Retrieved December 15, 2007.)
  4. ^ Figures given are for values up to 80 km (50 mi) above the surface
  5. ^ Thylakoid membranes are part of chloroplasts in algae and plants while they simply are one of many membrane structures in cyanobacteria. In fact, chloroplasts are thought to have evolved from cyanobacteria that were once symbiotic partners with the progenitors of plants and algae.
  6. ^ 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, and calcium and chloride are also required for the reaction to occur. (Raven 2005)
  7. ^ (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
  8. ^ a b c d Derived from mmHg values using 0.133322 kPa/mmHg
  9. ^ The reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspired O
    2     partial pressure nearer to that found at sea-level.
  10. ^ Also, since oxygen has a higher electronegativity than hydrogen, the charge difference makes it a polar molecule. The interactions between the different dipoles of each molecule cause a net attraction force.
  11. ^ Since O
    2    's partial pressure is the fraction of O
    2     times the total pressure, elevated partial pressures can occur either from high O
    2     fraction in breathing gas or from high breathing gas pressure, or a combination of both.
  12. ^ No single ignition source of the fire was conclusively identified, although some evidence points to an arc from an electrical spark.[135]

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General references

  • Cook, Gerhard A.; Lauer, Carol M. (1968). "Oxygen". In Clifford A. Hampel (ed.). The Encyclopedia of the Chemical Elements. New York: Reinhold Book Corporation. pp. 499–512. LCCN 68-29938.
  • Emsley, John (2001). "Oxygen". Nature's Building Blocks: An A–Z Guide to the Elements. Oxford, England: Oxford University Press. pp. 297–304. ISBN 978-0-19-850340-8.
  • Raven, Peter H.; Evert, Ray F.; Eichhorn, Susan E. (2005). Biology of Plants (7th ed.). New York: W. H. Freeman and Company Publishers. pp. 115–27. ISBN 978-0-7167-1007-3.

External links

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  • Oxygen at The Periodic Table of Videos (University of Nottingham)
  • Oxidizing Agents > Oxygen
  • Oxygen (O2) Properties, Uses, Applications
  • Roald Hoffmann article on "The Story of O"
  • WebElements.com – Oxygen
  • Oxygen on In Our Time at the BBC
  • Scripps Institute: Atmospheric Oxygen has been dropping for 20 years

January 26, 2023
oxygen, this, article, about, chemical, element, other, uses, disambiguation, chemical, element, with, symbol, atomic, number, member, chalcogen, group, periodic, table, highly, reactive, nonmetal, oxidizing, agent, that, readily, forms, oxides, with, most, el. This article is about the chemical element For other uses see Oxygen disambiguation Oxygen is the chemical element with the symbol O and atomic number 8 It is a member of the chalcogen group in the periodic table a highly reactive nonmetal and an oxidizing agent that readily forms oxides with most elements as well as with other compounds Oxygen is Earth s most abundant element and after hydrogen and helium it is the third most abundant element in the universe At standard temperature and pressure two atoms of the element bind to form dioxygen a colorless and odorless diatomic gas with the formula O2 Diatomic oxygen gas currently constitutes 20 95 of the Earth s atmosphere though this has changed considerably over long periods of time Oxygen makes up almost half of the Earth s crust in the form of oxides 3 Oxygen 8OLiquid oxygen boiling O2 OxygenAllotropesO2 O3 ozone and more see Allotropes of oxygen Appearancegas colorlessliquid and solid pale blueStandard atomic weight Ar O 15 99903 15 99977 15 999 0 001 abridged 1 Abundancein the Earth s crust461000 ppmOxygen in the periodic tableHydrogen HeliumLithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine NeonSodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine ArgonPotassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine KryptonRubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine XenonCaesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury element Thallium Lead Bismuth Polonium Astatine RadonFrancium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson O Snitrogen oxygen fluorineAtomic number Z 8Groupgroup 16 chalcogens Periodperiod 2Block p blockElectron configuration He 2s2 2p4Electrons per shell2 6Physical propertiesPhase at STPgasMelting point O2 54 36 K 218 79 C 361 82 F Boiling point O2 90 188 K 182 962 C 297 332 F Density at STP 1 429 g Lwhen liquid at b p 1 141 g cm3Triple point54 361 K 0 1463 kPaCritical point154 581 K 5 043 MPaHeat of fusion O2 0 444 kJ molHeat of vaporization O2 6 82 kJ molMolar heat capacity O2 29 378 J mol K Vapor pressureP Pa 1 10 100 1 k 10 k 100 kat T K 61 73 90Atomic propertiesOxidation states 2 1 0 1 2ElectronegativityPauling scale 3 44Ionization energies1st 1313 9 kJ mol2nd 3388 3 kJ mol3rd 5300 5 kJ mol more Covalent radius66 2 pmVan der Waals radius152 pmSpectral lines of oxygenOther propertiesNatural occurrenceprimordialCrystal structure cubicSpeed of sound330 m s gas at 27 C Thermal conductivity26 58 10 3 W m K Magnetic orderingparamagneticMolar magnetic susceptibility 3449 0 10 6 cm3 mol 293 K 2 CAS Number7782 44 7HistoryDiscoveryCarl Wilhelm Scheele 1771 Named byAntoine Lavoisier 1777 Main isotopes of oxygenveIso tope Decayabun dance half life t1 2 mode pro duct15O syn 122 266 s b 15N16O 99 757 stable17O 0 038 stable18O 0 205 stable Category Oxygenviewtalkedit referencesMany major classes of organic molecules in living organisms contain oxygen atoms such as proteins nucleic acids carbohydrates and fats as do the major constituent inorganic compounds 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 is continuously replenished in Earth s atmosphere by photosynthesis which uses the energy of sunlight to produce oxygen from water and carbon dioxide Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms Another form allotrope of oxygen ozone O3 strongly absorbs ultraviolet UVB radiation and the high altitude ozone layer helps protect the biosphere from ultraviolet radiation However ozone present at the surface is a byproduct of smog and thus a pollutant 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 The name oxygen was coined in 1777 by Antoine Lavoisier who first recognized oxygen as a chemical element and correctly characterized the role it plays in combustion Common uses of oxygen include production of steel plastics and textiles brazing welding and cutting of steels and other metals rocket propellant oxygen therapy and life support systems in aircraft submarines spaceflight and diving Contents 1 History of study 1 1 Early experiments 1 2 Phlogiston theory 1 3 Discovery 1 4 Lavoisier s contribution 1 4 1 Etymology 1 5 Later history 2 Characteristics 2 1 Properties and molecular structure 2 2 Allotropes 2 3 Physical properties 2 4 Isotopes and stellar origin 2 5 Occurrence 2 6 Analysis 3 Biological production and role of O2 3 1 Photosynthesis and respiration 3 2 Living organisms 3 3 Build up in the atmosphere 3 4 Extraterrestrial free oxygen 4 Industrial production 5 Storage 6 Applications 6 1 Medical 6 2 Life support and recreational use 6 3 Industrial 7 Compounds 7 1 Oxides and other inorganic compounds 7 2 Organic compounds 8 Safety and precautions 8 1 Toxicity 8 2 Combustion and other hazards 9 See also 10 Notes 11 References 11 1 General references 12 External linksHistory of studyEarly experiments 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 4 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 5 In the late 17th century Robert Boyle proved that air is necessary for combustion English chemist John Mayow 1641 1679 refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus 6 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 7 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 6 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 6 Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract De respiratione 7 Phlogiston theory Main article 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 8 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 9 Established in 1667 by the German alchemist J J Becher and modified by the chemist Georg Ernst Stahl by 1731 10 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 5 Highly combustible materials that leave little 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 5 Discovery Joseph Priestley is usually given priority in the discovery Polish alchemist philosopher and physician Michael Sendivogius Michal Sedziwoj in his work De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti 1604 described a substance contained in air referring to it as cibus vitae food of life 11 and according to Polish historian Roman Bugaj this substance is identical with oxygen 12 Sendivogius during his experiments performed between 1598 and 1604 properly recognized that the substance is equivalent to the gaseous byproduct released by the thermal decomposition of potassium nitrate In Bugaj s view the isolation of oxygen and the proper association of the substance to that part of air which is required for life provides sufficient evidence for the discovery of oxygen by Sendivogius 12 This discovery of Sendivogius was however frequently denied by the generations of scientists and chemists which succeeded him 11 It is also commonly claimed that oxygen was first discovered by Swedish pharmacist Carl Wilhelm Scheele He had produced oxygen gas by heating mercuric oxide HgO and various nitrates in 1771 72 13 14 5 Scheele called the gas fire air because it was then the only known agent to support combustion He wrote an account of this discovery in a manuscript titled Treatise on Air and Fire which he sent to his publisher in 1775 That document was published in 1777 15 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 14 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 common air but I fancied that my breast felt peculiarly light and easy for some time afterwards 8 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 5 16 Because he published his findings first Priestley is usually given priority in the discovery The French chemist Antoine Laurent 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 15 Lavoisier s contribution Antoine Lavoisier discredited the phlogiston theory Lavoisier conducted the first adequate quantitative experiments on oxidation and gave the first correct explanation of how combustion works 14 He used these and similar experiments all started in 1774 to discredit the phlogiston theory and to prove that the substance discovered by Priestley and Scheele was a chemical element In one experiment Lavoisier observed that there was no overall increase in weight when tin and air were heated in a closed container 14 He noted that air rushed in when he opened the container which indicated that part of the trapped air had been consumed He also noted that the tin had increased in weight and that increase was the same as the weight of the air that rushed back in This and other experiments on combustion were documented in his book Sur la combustion en general which was published in 1777 14 In that work he proved that air is a mixture of two gases vital air which is essential to combustion and respiration and azote Gk ἄzwton lifeless which did not support either Azote later became nitrogen in English although it has kept the earlier name in French and several other European languages 14 Etymology Lavoisier renamed vital air to oxygene in 1777 from the Greek roots ὀ3ys oxys acid literally sharp from the taste of acids and genhs genes producer literally begetter because he mistakenly believed that oxygen was a constituent of all acids 17 Chemists such as Sir Humphry Davy in 1812 eventually determined that Lavoisier was wrong in this regard but by then the name was too well established 18 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 15 Later history Robert H Goddard and a liquid oxygen gasoline rocket John Dalton s original 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 19 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 20 a The first commercial method of producing oxygen was chemical the so called Brin process involving a reversible reaction of barium oxide It was invented in 1852 and commercialized in 1884 but was displaced by newer methods in early 20th century By the late 19th century scientists realized that air could be liquefied and its components isolated by compressing and cooling it Using a cascade method Swiss chemist and physicist Raoul Pierre Pictet 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 21 Just two days later French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen 21 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 Wroblewski and Karol Olszewski 22 An experiment setup for preparation of oxygen in academic laboratories In 1891 Scottish chemist James Dewar was able to produce enough liquid oxygen for study 23 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 distilled the component gases by boiling them off one at a time and capturing them separately 24 Later in 1901 oxyacetylene welding was demonstrated for the first time by burning a mixture of acetylene and compressed O2 This method of welding and cutting metal later became common 24 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 m at 97 km h on March 16 1926 in Auburn Massachusetts US 24 25 In academic laboratories oxygen can be prepared by heating together potassium chlorate mixed with a small proportion of manganese dioxide 26 Oxygen levels in the atmosphere are trending slightly downward globally possibly because of fossil fuel burning 27 CharacteristicsProperties and molecular structure Orbital diagram after Barrett 2002 28 showing the participating atomic orbitals from each oxygen atom the molecular orbitals that result from their overlap and the 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 s and s and p and p orbital pairs At standard temperature and pressure oxygen is a colorless odorless and tasteless gas with the molecular formula O2 referred to as dioxygen 29 As dioxygen two oxygen atoms are 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 orbitals 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 filling of orbitals and the resulting cancellation of contributions from the 2s electrons after sequential filling of the low s and s orbitals s overlap of the two atomic 2p orbitals that lie along the O O molecular axis and p 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 p orbitals 28 This combination of cancellations and s and p 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 p orbitals in the diagram that are of equal energy i e degenerate is a configuration termed a spin triplet state Hence the ground state of the O2 molecule is referred to as triplet oxygen 30 b 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 31 Liquid oxygen temporarily suspended in a magnet owing to its paramagnetism In the triplet form O2 molecules are paramagnetic That is they impart magnetic character to oxygen when it is in the presence of a magnetic field because of the spin magnetic moments of the unpaired electrons in the molecule and the negative exchange energy between neighboring O2 molecules 23 Liquid oxygen is so magnetic that in laboratory demonstrations a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet 32 c Singlet oxygen is a name given to several higher energy species of molecular O2 in which all the electron spins are paired It is much more reactive with common organic molecules than is normal triplet molecular oxygen In nature singlet oxygen is commonly formed from water during photosynthesis using the energy of sunlight 33 It is also produced in the troposphere by the photolysis of ozone by light of short wavelength 34 and by the immune system as a source of active oxygen 35 Carotenoids 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 36 Allotropes Main article Allotropes of oxygen Space filling model representation of dioxygen O2 molecule The common allotrope of elemental oxygen on Earth is called dioxygen O2 the major part of the Earth s atmospheric oxygen see Occurrence O2 has a bond length of 121 pm and a bond energy of 498 kJ mol 37 O2 is used by complex forms of life such as animals in cellular respiration Other aspects of O2 are covered in the remainder of this article Trioxygen O3 is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue 38 Ozone is produced in the upper atmosphere when O2 combines with atomic oxygen made by the splitting of O2 by ultraviolet UV radiation 17 Since ozone absorbs strongly in the UV region of the spectrum the ozone layer of the upper atmosphere functions as a protective radiation shield for the planet 17 Near the Earth s surface it is a pollutant formed as a by product of automobile exhaust 38 At low earth orbit altitudes sufficient atomic oxygen is present to cause corrosion of spacecraft 39 The metastable molecule tetraoxygen O4 was discovered in 2001 40 41 and was assumed to exist in one of the six phases of solid oxygen It was proven in 2006 that this phase created by pressurizing O2 to 20 GPa is in fact a rhombohedral O8 cluster 42 This cluster has the potential to be a much more powerful oxidizer than either O2 or O3 and may therefore be used in rocket fuel 40 41 A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa 43 and it was shown in 1998 that at very low temperatures this phase becomes superconducting 44 Physical properties Oxygen discharge spectrum tube See also Liquid oxygen and solid oxygen Oxygen dissolves more readily in water than in nitrogen and in freshwater more readily than in seawater Water in equilibrium with air contains approximately 1 molecule of dissolved O2 for every 2 molecules of N2 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 14 6 mg L dissolves at 0 C than at 20 C 7 6 mg L 8 45 At 25 C and 1 standard atmosphere 101 3 kPa of air freshwater can dissolve about 6 04 milliliters mL of oxygen per liter and seawater contains about 4 95 mL per liter 46 At 5 C the solubility increases to 9 0 mL 50 more than at 25 C per liter for freshwater and 7 2 mL 45 more per liter for sea water Oxygen gas dissolved in water at sea level milliliters per liter 5 C 25 CFreshwater 9 00 6 04Seawater 7 20 4 95Oxygen condenses at 90 20 K 182 95 C 297 31 F and freezes at 54 36 K 218 79 C 361 82 F 47 Both liquid and solid O2 are clear substances with a light sky blue color caused by absorption in the red in contrast with the blue color of the sky which is due to Rayleigh scattering of blue light High purity liquid O2 is usually obtained by the fractional distillation of liquefied air 48 Liquid oxygen may also be condensed from air using liquid nitrogen as a coolant 49 Liquid oxygen is a highly reactive substance and must be segregated from combustible materials 49 The spectroscopy of molecular oxygen is associated with the atmospheric processes of aurora and airglow 50 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 51 Excited state singlet molecular oxygen is responsible for red chemiluminescence in solution 52 Table of thermal and physical properties of oxygen O2 at atmospheric pressure 53 54 Temperature K Density kg m 3 Specific heat kJ kg C Dynamic viscosity kg m s Kinematic viscosity m 2 s Thermal conductivity W m C Thermal diffusivity m 2 s Prandtl Number100 3 945 0 962 7 64E 06 1 94E 06 0 00925 2 44E 06 0 796150 2 585 0 921 1 15E 05 4 44E 06 0 0138 5 80E 06 0 766200 1 93 0 915 1 48E 05 7 64E 06 0 0183 1 04E 05 0 737250 1 542 0 915 1 79E 05 1 16E 05 0 0226 1 60E 05 0 723300 1 284 0 92 2 07E 05 1 61E 05 0 0268 2 27E 05 0 711350 1 1 0 929 2 34E 05 2 12E 05 0 0296 2 90E 05 0 733400 0 962 1 0408 2 58E 05 2 68E 05 0 033 3 64E 05 0 737450 0 8554 0 956 2 81E 05 3 29E 05 0 0363 4 44E 05 0 741500 0 7698 0 972 3 03E 05 3 94E 05 0 0412 5 51E 05 0 716550 0 6998 0 988 3 24E 05 4 63E 05 0 0441 6 38E 05 0 726600 0 6414 1 003 3 44E 05 5 36E 05 0 0473 7 35E 05 0 729700 0 5498 1 031 3 81E 05 6 93E 05 0 0528 9 31E 05 0 744800 0 481 1 054 4 15E 05 8 63E 05 0 0589 1 16E 04 0 743900 0 4275 1 074 4 47E 05 1 05E 04 0 0649 1 41E 04 0 741000 0 3848 1 09 4 77E 05 1 24E 04 0 071 1 69E 04 0 7331100 0 3498 1 103 5 06E 05 1 45E 04 0 0758 1 96E 04 0 7361200 0 3206 1 0408 5 33E 05 1 661E 04 0 0819 2 29E 04 0 7251300 0 296 1 125 5 88E 05 1 99E 04 0 0871 2 62E 04 0 721Isotopes and stellar origin 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 three stable isotopes 16O 17O and 18O with 16O being the most abundant 99 762 natural abundance 55 Most 16O is synthesized at the end of the helium fusion process in massive stars but some is made in the neon burning process 56 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 56 Most 18O is produced when 14N made abundant from CNO burning captures a 4He nucleus making 18O common in the helium rich zones of evolved massive stars 56 Thirteen radioisotopes have been characterized ranging from 11O to 26O 57 The most stable are 15O with a half life of 122 24 seconds and 14O with a half life of 70 606 seconds 55 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 55 The most common decay mode of the isotopes lighter than 16O is b decay 58 59 60 to yield nitrogen and the most common mode for the isotopes heavier than 18O is beta decay to yield fluorine 55 Occurrence See also Silicate minerals Category Oxide minerals Stellar population Cosmochemistry and Astrochemistry Ten most common elements in the Milky Way Galaxy estimated spectroscopically 61 Z Element Mass fraction in parts per million1 Hydrogen 739 000 71 mass of oxygen red bar 2 Helium 240 000 23 mass of oxygen red bar 8 Oxygen 10 400 10400 6 Carbon 4 600 4600 10 Neon 1 340 1340 26 Iron 1 090 1090 7 Nitrogen 960 960 14 Silicon 650 650 12 Magnesium 580 580 16 Sulfur 440 440 Oxygen is the most abundant chemical element by mass in the Earth s biosphere air sea and land Oxygen is the third most abundant chemical element in the universe after hydrogen and helium 62 About 0 9 of the Sun s mass is oxygen 14 Oxygen constitutes 49 2 of the Earth s crust by mass 63 as part of oxide compounds such as silicon dioxide and is the most abundant element by mass in the Earth s crust It is also the major component of the world s oceans 88 8 by mass 14 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 1015 tonnes 14 64 d 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 O2 by volume and Venus have much less The O2 surrounding those planets is produced solely by the action of ultraviolet radiation on oxygen containing molecules such as carbon dioxide Cold water holds more dissolved O2 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 respiration decay and combustion remove it from the atmosphere In the present equilibrium production and consumption occur at the same rate 65 Free oxygen also occurs in solution in the world s water bodies The increased solubility of O2 at lower 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 66 Water polluted with plant nutrients such as nitrates or phosphates may stimulate growth of algae by a process called eutrophication and the decay of these organisms and other biomaterials may reduce the O2 content in eutrophic water bodies Scientists assess this aspect of water quality by measuring the water s biochemical oxygen demand or the amount of O2 needed to restore it to a normal concentration 67 Analysis 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 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 68 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 68 Paleoclimatologists also directly measure this ratio in the water molecules of ice core samples as old as hundreds of thousands of years Planetary geologists have measured the relative quantities of oxygen isotopes in samples from the Earth the Moon Mars and meteorites but were long unable to obtain reference values for the isotope ratios in the Sun believed to be the same as those of the primordial solar nebula Analysis of a silicon wafer exposed to the solar wind in space and returned by the crashed Genesis spacecraft has shown that the Sun has a higher proportion of oxygen 16 than does the Earth The measurement implies that an unknown process depleted oxygen 16 from the Sun s disk of protoplanetary material prior to the coalescence of dust grains that formed the Earth 69 Oxygen presents two spectrophotometric absorption bands peaking at the wavelengths 687 and 760 nm Some remote sensing scientists have proposed using the measurement of the radiance coming from vegetation canopies in those bands to characterize plant health status from a satellite platform 70 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 Biological production and role of O2Main article Dioxygen in biological reactions Photosynthesis and respiration Photosynthesis splits water to liberate O2 and fixes CO2 into sugar in what is called a Calvin cycle In nature free oxygen is produced by the light driven splitting of water during oxygenic photosynthesis According to some estimates green algae and cyanobacteria in marine environments provide about 70 of the free oxygen produced on Earth and the rest is produced by terrestrial plants 71 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 72 A simplified overall formula for photosynthesis is 73 6 CO2 6 H2 O photons C6 H12 O6 6 O2or simply carbon dioxide water sunlight glucose dioxygenPhotolytic oxygen evolution occurs in the thylakoid membranes of photosynthetic organisms and requires the energy of four photons e 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 74 The O2 remaining after production of the water molecule is released into the atmosphere f Oxygen is used in mitochondria in the generation of ATP during oxidative phosphorylation The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as C6 H12 O6 6 O2 6 CO2 6 H2 O 2880 kJ molIn vertebrates O2 diffuses through membranes in the lungs and into red blood cells Hemoglobin binds O2 changing color from bluish red to bright red 38 CO2 is released from another part of hemoglobin through the Bohr effect Other animals use hemocyanin molluscs and some arthropods or hemerythrin spiders and lobsters 64 A liter of blood can dissolve 200 cm3 of O2 64 Until the discovery of anaerobic metazoa 75 oxygen was thought to be a requirement for all complex life 76 Reactive oxygen species such as superoxide ion O 2 and hydrogen peroxide H2 O2 are reactive by products of oxygen use in organisms 64 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 74 Oxygen is damaging to obligately anaerobic organisms which were the dominant form of early life on Earth until O2 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 77 78 An adult human at rest inhales 1 8 to 2 4 grams of oxygen per minute 79 This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year g Living organisms Partial pressures of oxygen in the human body PO2 Unit Alveolar pulmonary gas pressures Arterial blood oxygen Venous blood gaskPa 14 2 11 h 13 h 4 0 h 5 3 h mmHg 107 75 80 100 80 30 81 40 81 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 82 Build up in the atmosphere Main article Geological history of oxygen O2 build up in Earth s atmosphere 1 no O2 produced 2 O2 produced but absorbed in oceans amp 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 Free oxygen gas was almost nonexistent in Earth s atmosphere before photosynthetic archaea and bacteria evolved probably about 3 5 billion years ago Free oxygen first appeared in significant quantities during the Paleoproterozoic eon between 3 0 and 2 3 billion years ago 83 Even if there was much dissolved iron in the oceans when oxygenic photosynthesis was getting more common it appears the banded iron formations were created by anoxyenic or micro aerophilic iron oxidizing bacteria which dominated the deeper areas of the photic zone while oxygen producing cyanobacteria covered the shallows 84 Free oxygen began to outgas from the oceans 3 2 7 billion years ago reaching 10 of its present level around 1 7 billion years ago 83 85 The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the extant anaerobic organisms to extinction during the Great Oxygenation Event oxygen catastrophe about 2 4 billion years ago Cellular respiration using O2 enables aerobic organisms to produce much more ATP than anaerobic organisms 86 Cellular respiration of O2 occurs in all eukaryotes including all complex multicellular organisms such as plants and animals Since the beginning of the Cambrian period 540 million years ago atmospheric O2 levels have fluctuated between 15 and 30 by volume 87 Towards the end of the Carboniferous period about 300 million years ago atmospheric O2 levels reached a maximum of 35 by volume 87 which may have contributed to the large size of insects and amphibians at this time 88 Variations in atmospheric oxygen concentration have shaped past climates When oxygen declined atmospheric density dropped which in turn increased surface evaporation causing precipitation increases and warmer temperatures 89 At the current rate of photosynthesis it would take about 2 000 years to regenerate the entire O2 in the present atmosphere 90 Extraterrestrial free oxygen Main article Extraterrestrial atmosphere In the field of astrobiology and in the search for extraterrestrial life oxygen is a strong biosignature That said it might not be a definite biosignature being possibly produced abiotically on celestial bodies with processes and conditions such as a peculiar hydrosphere which allow free oxygen 91 92 93 like with Europa s and Ganymede s thin oxygen atmospheres 94 Industrial productionSee also Air separation Oxygen evolution and Fractional distillation Hofmann electrolysis apparatus used in electrolysis of water One hundred million tonnes of O2 are extracted from air for industrial uses annually by two primary methods 15 The most common method is fractional distillation of liquefied air with N2 distilling as a vapor while O2 is left as a liquid 15 The other primary method of producing O2 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 to 93 O2 15 Simultaneously nitrogen gas is released from the other nitrogen saturated zeolite bed by reducing the chamber operating pressure and diverting part of the oxygen gas from the producer bed through it in the reverse direction of flow After a set cycle time the operation of the two beds is interchanged thereby allowing for a continuous supply of gaseous oxygen to be pumped through a pipeline This is known as pressure swing adsorption Oxygen gas is increasingly obtained by these non cryogenic technologies see also the related vacuum swing adsorption 95 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 O2 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 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 O2 gas 67 Storage Oxygen and MAPP gas compressed gas cylinders with regulators Oxygen storage methods include high pressure oxygen tanks cryogenics and chemical compounds For reasons of economy oxygen is often transported in bulk as a liquid in specially insulated tankers since one liter of liquefied oxygen is equivalent to 840 liters of gaseous oxygen at atmospheric pressure and 20 C 68 F 15 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 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 15 ApplicationsSee also Breathing gas Redox and Combustion Medical An oxygen concentrator in an emphysema patient s house Main article Oxygen therapy Uptake of O2 from the air is the essential purpose of 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 96 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 replaced mostly by the use of oxygen masks or nasal cannulas 97 Hyperbaric high pressure medicine uses special oxygen chambers to increase the partial pressure of O2 around the patient and when needed the medical staff 98 Carbon monoxide poisoning gas gangrene and decompression sickness the bends are sometimes addressed with this therapy 99 Increased O2 concentration in the lungs helps to displace carbon monoxide from the heme group of hemoglobin 100 101 Oxygen gas is poisonous to the anaerobic bacteria that cause gas gangrene so increasing its partial pressure helps kill them 102 103 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 O2 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 96 104 105 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 106 107 108 Life support and recreational use Low pressure pure O2 is used in space suits An application of O2 as a low pressure breathing gas is in modern space suits which surround their occupant s body with the breathing gas These devices use nearly pure oxygen at about one third normal pressure resulting in a normal blood partial pressure of O2 This trade off of higher oxygen concentration for lower pressure is needed to maintain suit flexibility 109 110 Scuba and surface supplied underwater divers and submariners also rely on artificially delivered O2 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 O2 use in diving at pressures higher than atmospheric is usually limited to rebreathers or decompression at relatively shallow depths 6 meters depth or less 111 112 or 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 O2 with other gases such as nitrogen or helium to prevent oxygen toxicity 111 People who climb mountains or fly in non pressurized fixed wing aircraft sometimes have supplemental O2 supplies i Pressurized commercial airplanes have an emergency supply of O2 automatically supplied to the passengers in case of cabin depressurization Sudden cabin pressure loss activates chemical oxygen generators above each seat causing oxygen masks 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 67 A steady stream of oxygen gas is then produced by the exothermic reaction Oxygen as a mild euphoric has a history of recreational use in oxygen bars and in sports Oxygen bars are establishments found in the United States since the late 1990s that offer higher than normal O2 exposure for a minimal fee 113 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 113 Available studies support a performance boost from oxygen enriched mixtures only if it is inhaled during aerobic exercise 114 Other recreational uses that do not involve breathing include pyrotechnic applications such as George Goble s five second ignition of barbecue grills 115 Industrial Most commercially produced O2 is used to smelt and or decarburize iron Smelting of iron ore into steel consumes 55 of commercially produced oxygen 67 In this process O2 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 1 700 C 67 Another 25 of commercially produced oxygen is used by the chemical industry 67 Ethylene is reacted with O2 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 67 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 67 Oxygen is used in oxyacetylene welding burning acetylene with O2 to produce a very hot flame In this process metal up to 60 cm 24 in thick is first heated with a small oxy acetylene flame and then quickly cut by a large stream of O2 116 CompoundsMain article Compounds of oxygen Water H2 O is the most familiar oxygen compound The oxidation state of oxygen is 2 in almost all known compounds of oxygen The oxidation state 1 is found in a few compounds such as peroxides 117 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 118 Oxides and other inorganic compounds Water H2 O is an oxide of hydrogen and the most familiar oxygen compound Hydrogen atoms are covalently bonded to oxygen in a water molecule but also have an additional attraction about 23 3 kJ mol per hydrogen atom to an adjacent oxygen atom in a separate molecule 119 These hydrogen bonds between water molecules hold them approximately 15 closer than what would be expected in a simple liquid with just van der Waals forces 120 j Oxides such as iron oxide or rust form when oxygen combines with other elements Due to its electronegativity oxygen forms chemical bonds with almost all other elements to give corresponding oxides 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 passivates the metal and slows further corrosion Many oxides of the transition metals are non stoichiometric compounds with slightly less metal than the chemical formula would show For example the mineral FeO wustite is written as Fe 1 x O displaystyle ce Fe 1 x ce O where x is usually around 0 05 121 Oxygen is present in the atmosphere in trace quantities in the form of carbon dioxide CO2 The Earth s crustal rock is composed in large part of oxides of silicon silica SiO2 as found in granite and quartz aluminium aluminium oxide Al2 O3 in bauxite and corundum iron iron III oxide Fe2 O3 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 silicates in silicate minerals The Earth s mantle of much larger mass than the crust is largely composed of silicates of magnesium and iron Water soluble silicates in the form of Na4 SiO4 Na2 SiO3 and Na2 Si2 O5 are used as detergents and adhesives 122 Oxygen also acts as a ligand for transition metals forming transition metal dioxygen complexes which feature metal O2 This class of compounds includes the heme proteins hemoglobin and myoglobin 123 An exotic and unusual reaction occurs with PtF6 which oxidizes oxygen to give O2 PtF6 dioxygenyl hexafluoroplatinate 124 Organic compounds Acetone is an important feeder material in the chemical industry Oxygen Carbon Hydrogen 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 and amides R C O NR2 There are many important organic solvents that contain oxygen including acetone methanol ethanol isopropanol furan THF diethyl ether dioxane ethyl acetate DMF DMSO acetic acid and formic acid Acetone CH3 2 CO and phenol C6 H5 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 Epoxides are ethers in which the oxygen atom is part of a ring of three atoms The element is similarly found in almost all biomolecules that are important to or generated by life Oxygen reacts spontaneously with many organic compounds at or below room temperature in a process called autoxidation 125 Most of the organic compounds that contain oxygen are not made by direct action of O2 Organic compounds important in industry and commerce that are made by direct oxidation of a precursor include ethylene oxide and peracetic acid 122 Safety and precautionsOxygen HazardsGHS labelling Pictograms Hazard statements H272Precautionary statements P220 P244 P370 P376 P403NFPA 704 fire diamond 001OX 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 126 Toxicity Main article Oxygen toxicity Main symptoms of oxygen toxicity 127 Oxygen gas O2 can be toxic at elevated partial pressures leading to convulsions and other health problems 111 k 128 Oxygen toxicity usually begins to occur at partial pressures more than 50 kilopascals kPa equal to about 50 oxygen composition at standard pressure or 2 5 times the normal sea level O2 partial pressure of about 21 kPa This is not a problem except for patients on mechanical ventilators since gas supplied through oxygen masks in medical applications is typically composed of only 30 50 O2 by volume about 30 kPa at standard pressure 8 At one time premature babies were placed in incubators containing O2 rich air but this practice was discontinued after some babies were blinded by the oxygen content being too high 8 Breathing pure O2 in space applications such as in some modern space suits or in early spacecraft such as Apollo causes no damage due to the low total pressures used 109 129 In the case of spacesuits the O2 partial pressure in the breathing gas is in general about 30 kPa 1 4 times normal and the resulting O2 partial pressure in the astronaut s arterial blood is only marginally more than normal sea level O2 partial pressure 130 Oxygen toxicity to the lungs and central nervous system can also occur in deep scuba diving and surface supplied diving 8 111 Prolonged breathing of an air mixture with an O2 partial pressure more than 60 kPa can eventually lead to permanent pulmonary fibrosis 131 Exposure to an O2 partial pressures greater than 160 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 O2 at 66 m 217 ft or more of depth the same thing can occur by breathing 100 O2 at only 6 m 20 ft 131 132 133 134 Combustion and other hazards The interior of the Apollo 1 Command Module Pure O2 at higher than normal pressure and a spark led to a fire and the loss of the Apollo 1 crew Highly concentrated sources of oxygen promote rapid combustion Fire and explosion hazards exist when concentrated oxidants and fuels are brought into close proximity an ignition event such as heat or a spark is needed to trigger combustion 31 Oxygen is the oxidant not the fuel Concentrated O2 will allow combustion to proceed rapidly and energetically 31 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 O2 systems requires special training to ensure that ignition sources are minimized 31 The fire that killed the Apollo 1 crew in a launch pad test spread so rapidly because the capsule was pressurized with pure O2 but at slightly more than atmospheric pressure instead of the 1 3 normal pressure that would be used in a mission l 136 Liquid oxygen spills if allowed to soak into organic matter such as wood petrochemicals and asphalt can cause these materials to detonate unpredictably on subsequent mechanical impact 31 See alsoGeological history of oxygen Hypoxia environmental for O2 depletion in aquatic ecology Ocean deoxygenation Hypoxia medical a lack of oxygen Limiting oxygen concentration Oxygen compounds Oxygen plant Oxygen sensor Portals Chemistry MedicineOxygen at Wikipedia s sister projects Definitions from Wiktionary Media from Commons Quotations from Wikiquote Textbooks from Wikibooks Resources from WikiversityNotes These results were mostly ignored until 1860 Part of this rejection was due to the belief that atoms of one element would have no chemical affinity towards atoms of the same element and part was due to apparent exceptions to Avogadro s law that were not explained until later in terms of dissociating molecules An orbital is a concept from quantum mechanics that models an electron as a wave like particle that has a spatial distribution about an atom or molecule Oxygen s paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine the purity of gaseous oxygen Company literature of Oxygen analyzers triplet Servomex Archived from the original on March 8 2008 Retrieved December 15 2007 Figures given are for values up to 80 km 50 mi above the surface Thylakoid membranes are part of chloroplasts in algae and plants while they simply are one of many membrane structures in cyanobacteria In fact chloroplasts are thought to have evolved from cyanobacteria that were once symbiotic partners with the progenitors of plants and algae Water oxidation is catalyzed by a manganese containing enzyme complex known as the oxygen evolving complex OEC or water splitting complex found associated with 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