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Oxygen,  8O
A glass bottle half-filled with a bluish bubbling liquid
Liquid oxygen, oxygen bubbles
Spectral lines of oxygen
General properties
Name, symbol oxygen, O
Appearance gas: colorless
liquid: pale blue
Oxygen in the periodic table


Atomic number 8
Standard atomic weight 15.999(4)
Element category diatomic nonmetal
Group, block group 16 (chalcogens), p-block
Period period 2
Electron configuration [He] 2s2 2p4
per shell 2, 6
Physical properties
Phase gas
Melting point 54.36 K ​(−218.79 °C, ​−361.82 °F)
Boiling point 90.188 K ​(−182.962 °C, ​−297.332 °F)
Density at stp (0 °C and 101.325 kPa) 1.429 g·L−1
liquid, at b.p. 1.141 g·cm−3
Triple point 54.361 K, ​0.1463 kPa
Critical point 154.581 K, 5.043 MPa
Heat of fusion (O2) 0.444 kJ·mol−1
Heat of vaporization (O2) 6.82 kJ·mol−1
Molar heat capacity (O2) 29.378 J·mol−1·K−1
at T (K)       61 73 90
Atomic properties
Oxidation states 2, 1, −1, −2
Electronegativity Pauling scale: 3.44
Ionization energies 1st: 1313.9 kJ·mol−1
2nd: 3388.3 kJ·mol−1
3rd: 5300.5 kJ·mol−1
Covalent radius 66±2 pm
Van der Waals radius 152 pm
Crystal structure cubic
Cubic crystal structure for oxygen
Speed of sound 330 m·s−1 (gas, at 27 °C)
Thermal conductivity 26.58×10−3  W·m−1·K−1
Magnetic ordering paramagnetic
CAS number 7782-44-7
Discovery Carl Wilhelm Scheele (1772)
Named by Antoine Lavoisier (1777)
Most stable isotopes
iso NA half-life DM DE (MeV) DP
16O 99.76% 16O is stable with 8 neutrons
17O 0.039% 17O is stable with 9 neutrons
18O 0.201% 18O is stable with 10 neutrons
Blue white glow from an oxygen discharge tube.

Oxygen is a chemical element with symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table and is a highly reactive nonmetallic element and oxidizing agent that readily forms compounds (notably oxides) with most elements.[1] By mass, oxygen is the third-most abundant element in the universe, after hydrogen and helium.[2] At STP, two atoms of the element bind to form dioxygen, a diatomic gas that is colorless, odorless, and tasteless, with the formula O

Many major classes of organic molecules in living organisms, such as early life on Earth until O
began to accumulate in the atmosphere. Free elemental O
only began to accumulate in the atmosphere about 2.5 billion years ago during the

  • Oxygen at The Periodic Table of Videos (University of Nottingham)
  • Oxidizing Agents > Oxygen
  • ) Properties, Uses, Applications2Oxygen (O
  • Roald Hoffmann article on "The Story of O"
  • – Oxygen
  • Chemistry in its element podcast (MP3) from the Royal Society of Chemistry's Chemistry World: Oxygen
  • Oxygen on In Our Time at the BBC. (listen now)
  • Scripps Institute: Atmospheric Oxygen has been dropping for 20 years

External links

  • Walker, J. (1980). "The oxygen cycle". In Hutzinger O. Handbook of Environmental Chemistry. Volume 1. Part A: The natural environment and the biogeochemical cycles. Berlin; Heidelberg; New York: Springer-Verlag. p. 258.  

Further reading

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


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  1. ^ An orbital is a concept from quantum mechanics that models an electron as a wave-like particle that has a spacial distribution about an atom or molecule.
  2. ^ 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. Retrieved December 15, 2007. )
  3. ^ Figures given are for values up to 50 miles (80 km) above the surface
  4. ^ Thylakoid membranes are part of chloroplasts in algae and plants while they simply are one of many membrane structures in cyanobacteria. In fact, chloroplasts are thought to have evolved from cyanobacteria that were once symbiotic partners with the progenerators of plants and algae.
  5. ^ 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)
  6. ^ (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
  7. ^ However, these results were mostly ignored until 1860. Part of this rejection was due to the belief that atoms of one element would have no chemical affinity towards atoms of the same element, and part was due to apparent exceptions to Avogadro's law that were not explained until later in terms of dissociating molecules.
  8. ^ The reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspired O
    partial pressure nearer to that found at sea-level.
  9. ^ 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.
  10. ^ Since O
    's partial pressure is the fraction of O
    times the total pressure, elevated partial pressures can occur either from high O
    fraction in breathing gas or from high breathing gas pressure, or a combination of both.
  11. ^ No single ignition source of the fire was conclusively identified, although some evidence points to arc from an electrical spark.[107]


See also

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.[106] As with other cryogenic liquids, on contact with the human body it can cause frostbites to the skin and the eyes.

Concentrated O
will allow combustion to proceed rapidly and energetically.[106] 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
systems requires special training to ensure that ignition sources are minimized.[106] The fire that killed the Apollo 1 crew in a launch pad test spread so rapidly because the capsule was pressurized with pure O
but at slightly more than atmospheric pressure, instead of the 13 normal pressure that would be used in a mission.[11][108]

Highly concentrated sources of oxygen promote rapid combustion. Fire and explosion hazards exist when concentrated oxidants and fuels are brought into close proximity; however, an ignition event, such as heat or a spark, is needed to trigger combustion.[106] Oxygen itself is not the fuel, but the oxidant. Combustion hazards also apply to compounds of oxygen with a high oxidative potential, such as peroxides, chlorates, nitrates, perchlorates, and dichromates because they can donate oxygen to a fire.

An inside of some device, charred and apparently destroyed.
The interior of the Apollo 1 Command Module. Pure O
at higher than normal pressure and a spark led to a fire and the loss of the Apollo 1 crew.

Combustion and other hazards

Oxygen toxicity to the lungs and central nervous system can also occur in deep scuba diving and surface supplied diving.[25][85] Prolonged breathing of an air mixture with an O
partial pressure more than 60 kPa can eventually lead to permanent pulmonary fibrosis.[102] Exposure to a O
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
at 66 m or more of depth; the same thing can occur by breathing 100% O
at only 6 m.[102][103][104][105]

Breathing pure O
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.[83][101] In the case of spacesuits, the O
partial pressure in the breathing gas is, in general, about 30 kPa (1.4 times normal), and the resulting O
partial pressure in the astronaut's arterial blood is only marginally more than normal sea-level O
partial pressure (for more information on this, see space suit and arterial blood gas).

At one time, premature babies were placed in incubators containing O
-rich air, but this practice was discontinued after some babies were blinded by the oxygen content being too high.[25]

Oxygen gas (O
) can be toxic at elevated partial pressures, leading to convulsions and other health problems.[85][10][100] 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
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
by volume (about 30 kPa at standard pressure).[25] (although this figure also is subject to wide variation, depending on type of mask).

Four divers, equipped with oxygen cylinders, at the see bottom.
Oxygen toxicity occurs when the lungs take in 2 12 higher than normal O
partial pressure, which can occur in deep scuba diving.

A diagraph showing a man torso and listing symptoms of oxygen toxicity: Eyes – visual field loss, near)sightedness, cataract formation, bleeding, fibrosis; Head – seizures; Muscles – twitching; Respiratory system – jerky breathing, irritation, coughing, pain, shortness of breath, tracheobronchitis, acute respiratory distress syndrome.
Main symptoms of oxygen toxicity[99]


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.

Safety and precautions

The element is found in almost all carbohydrates contain the largest proportion by mass of oxygen. All fats, fatty acids, amino acids, and proteins contain oxygen (due to the presence of carbonyl groups in these acids and their ester residues). Oxygen also occurs in phosphate (PO3−
) groups in the biologically important energy-carrying molecules ATP and ADP, in the backbone and the purines (except adenine) and pyrimidines of RNA and DNA, and in bones as calcium phosphate and hydroxylapatite.

[95].peracetic acid and ethylene oxide. Organic compounds important in industry and commerce that are made by direct oxidation of a precursor include
O Oxygen reacts spontaneously with many

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
). 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
) and phenol (C
) 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.

Skeletal chemical structure with a linear chain of O-P-O bonds connected to three different carbon-nitrogen rings.
Oxygen represents more than 40% of the molecular mass of the ATP molecule.

A ball structure of a molecule. Its backbone is a zig-zag chain of three carbon atoms connected in the center to an oxygen atom and on the end to 6 hydrogens.
Acetone is an important feeder material in the chemical industry.

Organic compounds and biomolecules

Oxygen also acts as a ligand for transition metals, forming transition metal dioxygen complexes, which feature metal–O
. Notable members of this class of compounds are the heme proteins hemoglobin and myoglobin.[96] An exotic and unusual reaction occcurs with PtF
, which oxidizes oxygen to give O2+PtF6-.[97]

Water-soluble silicates in the form of Na
, Na
, and Na
are used as detergents and adhesives.[95]

Oxygen is present in the atmosphere in trace quantities in the form of carbon dioxide (CO
). The Earth's crustal rock is composed in large part of oxides of silicon (silica SiO
, as found found in granite and quartz), aluminium (aluminium oxide Al
, in bauxite and corundum), iron (iron(III) oxide Fe
, 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.

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements to give corresponding oxides. The surface of most metals, notably 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 actually written as Fe
1 − x
, where x is usually around 0.05.[94]

A rusty piece of a bolt.
Oxides, such as iron oxide or rust, form when oxygen combines with other elements.

Water (H
) 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−1 per hydrogen atom) to an adjacent oxygen atom in a separate molecule.[92] 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.[93][9]

Oxides and other inorganic compounds

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.[91] 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).

Water flowing from a bottle into a glass.
Water (H
) is the most familiar oxygen compound.


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

Another 25% of commercially produced oxygen is used by the chemical industry.[39] Ethylene is reacted with O
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).[39]

Smelting of iron ore into steel consumes 55% of commercially produced oxygen.[39] In this process, O
is injected through a high-pressure lance into molten iron, which removes sulfur impurities and excess carbon as the respective oxides, SO
and CO
. The reactions are exothermic, so the temperature increases to 1,700 °C.[39]

An elderly worker in a helmet is facing his side to the viewer in an industrial hall. The hall is dark but is illuminated yellow glowing splashes of a melted substance.
Most commercially produced O
is used to smelt iron into steel.


Other recreational uses that do not involve breathing the gas include barbecue grills.[89]

Oxygen, as a supposed mild euphoric, has a history of recreational use in oxygen bars and in sports. Oxygen bars are establishments, found in Japan, California, and Las Vegas, Nevada since the late 1990s that offer higher than normal O
exposure for a fee.[87] Professional athletes, especially in American football, also sometimes go off field between plays to wear oxygen masks in order to get a "boost" in performance. The pharmacological effect is doubtful; a placebo effect is a more likely explanation.[87] Available studies support a performance boost from enriched O
mixtures only if they are breathed during aerobic exercise.[88]

People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental O
supplies.[8] Passengers traveling in (pressurized) commercial airplanes have an emergency supply of O
automatically supplied to them in case of cabin depressurization. Sudden cabin pressure loss activates chemical oxygen generators above each seat, causing oxygen masks to drop. 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.[39] A steady stream of oxygen gas is then produced by the exothermic reaction.

Scuba divers and submariners also rely on artificially delivered O
, but most often use normal pressure, and/or mixtures of oxygen and air. Pure or nearly pure O
use in diving at higher-than-sea-level pressures is usually limited to rebreather, decompression, or emergency treatment use at relatively shallow depths (~6 meters depth, or less).[85][86] Deeper diving requires significant dilution of O
with other gases, such as nitrogen or helium, to help prevent oxygen toxicity.[85]

A notable application of O
as a low-pressure breathing gas is in modern space suits, which surround their occupant's body with pressurized air. These devices use nearly pure oxygen at about one third normal pressure, resulting in a normal blood partial pressure of O
.[83][84] This trade-off of higher oxygen concentration for lower pressure is needed to maintain flexible spacesuits.

Low pressure pure O
is used in space suits.

Life support and recreational use

Oxygen is also used medically for patients who require mechanical ventilation, often at concentrations above the 21% found in ambient air.

Hyperbaric (high-pressure) medicine uses special oxygen chambers to increase the partial pressure of O
around the patient and, when needed, the medical staff.[75] Carbon monoxide poisoning, gas gangrene, and decompression sickness (the 'bends') are sometimes treated using these devices.[76] Increased O
concentration in the lungs helps to displace carbon monoxide from the heme group of hemoglobin.[77][78] Oxygen gas is poisonous to the anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them.[79][80] Decompression sickness occurs in divers who decompress too quickly after a dive, resulting in bubbles of inert gas, mostly nitrogen and helium, forming in their blood. Increasing the pressure of O
as soon as possible is part of the treatment.[73][81][82]

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.[74]

Uptake of O
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.[73]

A gray device with a label DeVILBISS LT4000 and some text on the front panel. A green plastic pipe is running from the device.
An oxygen concentrator in an emphysema patient's house



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).[63] Such tankers are used to refill bulk liquid oxygen storage containers, which stand outside hospitals and other institutions with a need for large volumes of pure oxygen gas. Liquid oxygen is passed through heat exchangers, which convert the cryogenic liquid into gas before it enters the building. Oxygen is also stored and shipped in smaller cylinders containing the compressed gas; a form that is useful in certain portable medical applications and oxy-fuel welding and cutting.[63]


In large quantities, the price of liquid oxygen in 2001 was approximately $0.21/kg.[72] Since the primary cost of production is the energy cost of liquefying the air, the production cost will change as energy cost varies.

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. Contrary to popular belief, the 2:1 ratio observed in the DC electrolysis of acidified water does not prove that the empirical formula of water is H2O unless certain assumptions are made about the molecular formulae of hydrogen and oxygen themselves. A similar method is the electrocatalytic O
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 technology involves forcing air to dissolve through ceramic membranes based on zirconium dioxide by either high pressure or an electric current, to produce nearly pure O

The other major method of producing O
gas involves passing a stream of clean, dry air through one bed of a pair of identical zeolite molecular sieves, which absorbs the nitrogen and delivers a gas stream that is 90% to 93% O
.[63] 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).[71]

A drawing of three vertical pipes connected at the bottom and filled with oxygen (left pipe), water (middle) and hydrogen (right). Anode and cathode electrodes are inserted into the left and right pipes and externally connected to a battery.
Hofmann electrolysis apparatus used in electrolysis of water.

Two major methods are employed to produce 100 million tonnes of O
extracted from air for industrial uses annually.[63] The most common method is to fractionally distill liquefied air into its various components, with N
distilling as a vapor while O
is left as a liquid.[63]

Industrial production

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.[69][70]

In 1891 Scottish chemist James Dewar was able to produce enough liquid oxygen to study.[12] 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.[69] Later, in 1901, oxyacetylene welding was demonstrated for the first time by burning a mixture of acetylene and compressed O
. This method of welding and cutting metal later became common.[69]

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.[67] Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen.[67] Only a few drops of the liquid were produced in either case so no meaningful analysis could be conducted. Oxygen was liquified in stable state for the first time on March 29, 1883 by Polish scientists from Jagiellonian University, Zygmunt Wróblewski and Karol Olszewski.[68]

John Dalton's original atomic hypothesis assumed 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, giving the atomic mass of oxygen as 8 times that of hydrogen, instead of the modern value of about 16.[65] In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen; and by 1811 Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called Avogadro's law and the assumption of diatomic elemental molecules.[66][7]

A metal frame structure stands on the snow near a tree. A middle-aged man wearing a coat, boots, leather gloves and a cap stands by the structure and holds it with his right hand.
Robert H. Goddard and a liquid oxygen-gasoline rocket

Later history

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.[63]

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.[8] Chemists (notably Sir Humphry Davy in 1812) eventually determined that Lavoisier was wrong in this regard (it is in fact hydrogen that forms the basis for acid chemistry), but by that time it was too late; the name had taken.

In one experiment, Lavoisier observed that there was no overall increase in weight when tin and air were heated in a closed container.[5] 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.[5] 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 name in French and several other European languages.[5]

A drawing of a young man facing towards the viewer, but looking on the side. He wear a white curly wig, dark suit and white scarf.
Antoine Lavoisier discredited the Phlogiston theory.

What Lavoisier did indisputably do (although this was disputed at the time) was to conduct the first adequate quantitative experiments on oxidation and give the first correct explanation of how combustion works.[5] 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.

Lavoisier's contribution

The noted French chemist Antoine Laurent Lavoisier later claimed to have discovered the new substance independently. However, Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele also posted a letter to Lavoisier on September 30, 1774 that described his own discovery of the previously unknown substance, but Lavoisier never acknowledged receiving it (a copy of the letter was found in Scheele's belongings after his death).[63]

In the meantime, on August 1, 1774, an experiment conducted by the British clergyman Joseph Priestley focused sunlight on mercuric oxide (HgO) inside a glass tube, which liberated a gas he named "dephlogisticated air".[5] He noted that candles burned brighter in the gas and that a mouse was more active and lived longer while breathing it. After breathing the gas himself, he wrote: "The feeling of it to my lungs was not sensibly different from that of common air, but I fancied that my breast felt peculiarly light and easy for some time afterwards."[25] 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.[59][64] Because he published his findings first, Priestley is usually given priority in the discovery.

A drawing of an elderly man sitting by the table and facing parallel to the drawing. His left arm rests on a notebook, legs crossed
Joseph Priestley is usually given priority in the discovery.

Oxygen was first discovered by Swedish pharmacist Carl Wilhelm Scheele. He had produced oxygen gas by heating mercuric oxide and various nitrates by about 1772.[5][59] Scheele called the gas "fire air" because it was the only known supporter of combustion, and wrote an account of this discovery in a manuscript he titled Treatise on Air and Fire, which he sent to his publisher in 1775. However, that document was not published until 1777.[63]

Profile drawing of a young men's head in an oval frame.
Carl Wilhelm Scheele beat Priestley to the discovery but published afterwards.


Highly combustible materials that leave little residue, such as wood or coal, were thought to be made mostly of phlogiston; whereas non-combustible substances that corrode, such as iron, contained very little. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea; instead, it was based on observations of what happens when something burns, that most common objects appear to become lighter and seem to lose something in the process.[59] The fact that a substance like wood actually gains overall weight in burning was hidden by the buoyancy of the gaseous combustion products. Indeed one of the first clues that the phlogiston theory was incorrect was that metals, too, gain weight in rusting (when they were supposedly losing phlogiston).

[59].calx 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 [62] Established in 1667 by the German alchemist

Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen (fr) all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element.[25] 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.

Stahl helped develop and popularize the phlogiston theory.

Phlogiston theory

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

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 or just nitroaereus.[60] In one experiment he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.[61] From this he surmised that nitroaereus is consumed in both respiration and combustion.

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

Drawing of a burning candle enclosed in a glass bulb.
Philo's experiment inspired later investigators.

Early experiments


Since the beginning of the Cambrian period 540 million years ago, O
levels have fluctuated between 15% and 30% by volume.[55] Towards the end of the Carboniferous period (about 300 million years ago) atmospheric O
levels reached a maximum of 35% by volume,[55] which may have contributed to the large size of insects and amphibians at this time.[56] Human activities, including the burning of 7 billion tonnes of fossil fuels each year have had very little effect on the amount of free oxygen in the atmosphere.[12] At the current rate of photosynthesis it would take about 2,000 years to regenerate the entire O
in the present atmosphere.[57]

The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the extinction during the Great Oxygenation Event (oxygen catastrophe) about 2.4 billion years ago. However, cellular respiration using O
enables biosphere.[54] Cellular respiration of O
occurs in all eukaryotes, including all complex multicellular organisms such as plants and animals.

Free oxygen gas was almost nonexistent in iron in the oceans to form banded iron formations. When such oxygen sinks became saturated, 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.[52][53]

A graph showing time evolution of oxygen pressure on Earth; the pressure increases from zero to 0.2 atmospheres.
build-up in Earth's atmosphere: 1) no O
produced; 2) O
produced, but absorbed in oceans & seabed rock; 3) O
starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer; 4–5) O
sinks filled and the gas accumulates

Build-up in the atmosphere

The free oxygen respiratory system, and decreases along any arterial system, peripheral tissues and venous system, respectively. Partial pressure is the pressure which oxygen would have if it alone occupied the volume.[51]

Partial pressures of oxygen in the human body (PO2)
Unit Alveolar pulmonary
gas pressures
Arterial blood oxygen Venous blood gas
kPa 14.2 11[48]-13[48] 4.0[48]-5.3[48]
mmHg 107 75[49]-100[49] 30[50]-40[50]

Content in body

An adult human in rest inhales 1.8 to 2.4 grams of oxygen per minute.[47] This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year.[6]

Reactive oxygen species, such as superoxide ion (O
) and hydrogen peroxide (H
), are dangerous by-products of oxygen use in organisms.[37] Parts of the hypersensitive response of plants against pathogen attack.[46]

In vertebrates, O
diffuses through membranes in the lungs and into red blood cells. Hemoglobin binds O
, changing its color from bluish red to bright red[19] (CO
is released from another part of hemoglobin through the Bohr effect). Other animals use hemocyanin (molluscs and some arthropods) or hemerythrin (spiders and lobsters).[37] A liter of blood can dissolve 200 cm3 of O

+ 6 O
→ 6 CO
+ 6 H
+ 2880 kJ·mol−1

Molecular dioxygen, O
, is essential for cellular respiration in all mitochondria to help generate adenosine triphosphate (ATP) during oxidative phosphorylation. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as:

Photolytic photons.[4] Many steps are involved, but the result is the formation of a proton gradient across the thylakoid membrane, which is used to synthesize ATP via photophosphorylation.[46] The O
remaining after oxidation of the water molecule is released into the atmosphere.[5]

carbon dioxide + water + sunlight → glucose + dioxygen

or simply

6 CO
+ 6 H
+ photonsC
+ 6 O

A simplified overall formula for photosynthesis is:[45]

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.[43] 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.[44]

A diagram of photosynthesis processes, including income of water and carbon dioxide, illumination and release of oxygen. Reactions produce ATP and NADPH in a Calvin cycle with a sugar as a by product.
Photosynthesis splits water to liberate O
and fixes CO
into sugar in what is called a Calvin cycle.

Photosynthesis and respiration

Biological role of O2

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.[42] 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.

Planetary geologists have measured different abundances of oxygen isotopes in samples from the Earth, the Moon, Mars, and meteorites, but were long unable to obtain reference values for the isotope ratios in the Sun, believed to be the same as those of the primordial solar nebula. However, analysis of a silicon wafer exposed to the solar wind in space and returned by the crashed Genesis spacecraft has shown that the Sun has a higher proportion of oxygen-16 than does the Earth. The measurement implies that an unknown process depleted oxygen-16 from the Sun's disk of protoplanetary material prior to the coalescence of dust grains that formed the Earth.[41]

samples that are up to several hundreds of thousands of years old. ice core Paleoclimatologists also directly measure this ratio in the water molecules of [40]

Time evolution of oxygen-18 concentration on the scale of 500 million years showing many local peaks.
500 million years of climate change vs 18O


[39] needed to restore it to a normal concentration.O
, or the amount of biochemical oxygen demand in eutrophic water bodies. Scientists assess this aspect of water quality by measuring the water's
O [38]) has important implications for ocean life, as polar oceans support a much higher density of life due to their higher oxygen content.Physical properties at lower temperatures (see O
Free oxygen also occurs in solution in the world's water bodies. The increased solubility of

World map showing that the sea-surface oxygen is depleted around the equator and increases towards the poles.
Cold water holds more dissolved O

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 and decay remove it from the atmosphere. In the present equilibrium, production and consumption occur at the same rate of roughly 1/2000th of the entire atmospheric oxygen per year.

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.[2] About 0.9% of the Sun's mass is oxygen.[5] Oxygen constitutes 49.2% of the Earth's crust by mass[6] and is the major component of the world's oceans (88.8% by mass).[5] 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).[5][37][3] 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
by volume) and Venus have far lower concentrations. However, the O
surrounding these other planets is produced solely by ultraviolet radiation impacting oxygen-containing molecules such as carbon dioxide.

Ten most common elements in the Milky Way Galaxy estimated spectroscopically[36]
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


Fourteen radioisotopes have been characterized. The most stable are 15O with a half-life of 122.24 seconds and 14O with a half-life of 70.606 seconds.[31] 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.[31] The most common decay mode of the isotopes lighter than 16O is β+ decay[33][34][35] to yield nitrogen, and the most common mode for the isotopes heavier than 18O is beta decay to yield fluorine.[31]

Most 16O is synthesized at the end of the helium fusion process in massive stars but some is made in the neon burning process.[32] 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.[32] 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.[32]

Naturally occurring oxygen is composed of three stable isotopes, 16O, 17O, and 18O, with 16O being the most abundant (99.762% natural abundance).[31]

A concentric-sphere diagram, showing, from the core to the outer shell, iron, silicon, oxygen, neon, carbon, helium and hydrogen layers.
Late in a massive star's life, 16O concentrates in the O-shell, 17O in the H-shell and 18O in the He-shell.

Isotopes and stellar origin

Oxygen condenses at 90.20 K (−182.95 °C, −297.31 °F), and freezes at 54.36 K (−218.79 °C, −361.82 °F).[28] Both liquid and solid O
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
is usually obtained by the fractional distillation of liquefied air.[29] Liquid oxygen may also be produced by condensation out of air, using liquid nitrogen as a coolant. It is a highly reactive substance and must be segregated from combustible materials.[30]

Oxygen is more soluble in water than nitrogen is. Water in equilibrium with air contains approximately 1 molecule of dissolved O
for every 2 molecules of N
, compared to an atmospheric ratio of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg·L−1) dissolves at 0 °C than at 20 °C (7.6 mg·L−1).[25][26] At 25 °C and 1 standard atmosphere (101.3 kPa) of air, freshwater contains about 6.04 milliliters (mL) of oxygen per liter, whereas seawater contains about 4.95 mL per liter.[27] At 5 °C the solubility increases to 9.0 mL (50% more than at 25 °C) per liter for water and 7.2 mL (45% more) per liter for sea water.

Physical properties

Trioxygen (O
) is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue.[19] Ozone is produced in the upper atmosphere when O
combines with atomic oxygen made by the splitting of O
by ultraviolet (UV) radiation.[8] 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.[8] Near the Earth's surface, however, it is a pollutant formed as a by-product of automobile exhaust.[19] The metastable molecule tetraoxygen (O
) was discovered in 2001,[20][21] 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
to 20 GPa, is in fact a rhombohedral O
cluster.[22] This cluster has the potential to be a much more powerful oxidizer than either O
or O
and may therefore be used in rocket fuel.[20][21] A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa[23] and it was shown in 1998 that at very low temperatures, this phase becomes superconducting.[24]

The common allotrope of elemental oxygen on Earth is called dioxygen, O
. It has a bond length of 121 pm and a bond energy of 498 kJ·mol−1.[18] This is the form that is used by complex forms of life, such as animals, in cellular respiration (see Biological role) and is the form that is a major part of the Earth's atmosphere (see Occurrence). Other aspects of O
are covered in the remainder of this article.

Central atom is positively charged and end atoms are negatively charged.
Ozone is a rare gas on Earth found mostly in the stratosphere.


Singlet oxygen is a name given to several higher-energy species of molecular O
in which all the electron spins are paired. It is much more reactive towards common singlet oxygen and converting it to the unexcited ground state before it can cause harm to tissues.[17]

In normal triplet form, O
molecules are paramagnetic. That is, they form a magnet 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
molecules.[12] Liquid oxygen is attracted to a magnet to a sufficient extent that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.[13][2]

A trickle of liquid oxygen is deflected by a magnetic field, illustrating its paramagnetic property

Triplet oxygen (not to be confused with ozone, O
) is the ground state of the O
molecule.[11] The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals.[1] These orbitals are classified as antibonding (weakening the bond order from three to two), so the diatomic oxygen bond is weaker than the diatomic nitrogen triple bond in which all bonding molecular orbitals are filled, but some antibonding orbitals are not.[11]

At standard temperature and pressure, oxygen is a colorless, odorless gas with the molecular formula O
, in which the two oxygen atoms are chemically bonded to each other with a spin triplet electron configuration. This bond has a bond order of two, and is often simplified in description as a double bond[9] or as a combination of one two-electron bond and two three-electron bonds.[10]

Oxygen O2 molecule.




  • Characteristics 1
    • Structure 1.1
    • Allotropes 1.2
    • Physical properties 1.3
    • Isotopes and stellar origin 1.4
    • Occurrence 1.5
    • Analysis 1.6
  • Biological role of O2 2
    • Photosynthesis and respiration 2.1
    • Content in body 2.2
    • Build-up in the atmosphere 2.3
  • History 3
    • Early experiments 3.1
    • Phlogiston theory 3.2
    • Discovery 3.3
    • Lavoisier's contribution 3.4
    • Later history 3.5
  • Industrial production 4
  • Storage 5
  • Applications 6
    • Medical 6.1
    • Life support and recreational use 6.2
    • Industrial 6.3
  • Compounds 7
    • Oxides and other inorganic compounds 7.1
    • Organic compounds and biomolecules 7.2
  • Safety and precautions 8
    • Toxicity 8.1
    • Combustion and other hazards 8.2
  • See also 9
  • Notes 10
  • Citations 11
  • References 12
  • Further reading 13
  • External links 14

Oxygen was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph Priestley in Wiltshire, in 1774, but Priestley is often given priority because his work was published first. The name oxygen was coined in 1777 by Antoine Lavoisier,[8] whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. Its name derives from the Greek roots ὀξύς oxys, "acid", literally "sharp", referring to the sour taste of acids and -γενής -genes, "producer", literally "begetter", because at the time of naming, it was mistakenly thought that all acids required oxygen in their composition.

Oxygen is an important part of the atmosphere, and is necessary to sustain most terrestrial life as it is used in respiration. However, it is too chemically reactive to remain a free element in allotrope) of oxygen, ozone (O
), strongly absorbs UVB radiation and consequently the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation, but is a pollutant near the surface where it is a by-product of smog. At even higher low earth orbit altitudes, atomic oxygen is a significant presence and a cause of erosion for spacecraft.[7] Oxygen is produced industrially by fractional distillation of liquefied air, use of zeolites with pressure-cycling to concentrate oxygen from air, electrolysis of water and other means. Uses of elemental oxygen include the 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.

[6], making up almost half of the crust's mass.silicon dioxide as part of oxide compounds such as Earth's crust element by mass in the most abundant Oxygen is the [5] Diatomic oxygen gas constitutes 20.8% of the volume of air.[4][3]

