Oxygen is the chemical element with atomic number 8, symbol O. It is the leader of the chalcogen group, often referred to the oxygen group. Discovered independently in 1772 by the Swede Carl Wilhelm Scheele in Uppsala, and in 1774 by Pierre Bayen in Châlons-en-Champagne as well as by the British Joseph Priestley in Wiltshire, oxygen was named in 1777 by the Frenchman Antoine Lavoisier and his wife in Paris from the ancient Greek ὀξύς / oxús (“acute”, that is to say here “acid”), and γενής / genḗs (“generator”), because Lavoisier mistakenly thought — oxidation and acidification being related — that:

“We have given the base of the breathable portion of the air the name of oxygen, deriving it from two Greek words ὀξύς, acid and γείνομαι, I beget, because indeed one of the most general properties of this base [Lavoisier speaks of oxygen] is to form acids by combining with most substances. We will therefore call oxygen gas the union of this base with the caloric.»

Atomic number8
Period2nd period
BlockBlock p
Element familyNon-metal
Electronic configuration[He] 2s 2 2 p4
Electrons by energy level2, 6
Atomic mass15.999 4 ± 0.000 3 u (atom O)
Atomic radius (calc)60 pm (48 pm)
Covalence radius66 ± 2 pm
Van der Waals radius140 pm
Oxidation state-2, -1
Electronegativity (Pauling)3,44
1re: 13,618 05 eV2nd: 35.121 1 eV
3rd: 54.935 5 eV4th: 77.413 53 eV
5th: 113.899 0 eV6th: 138.119 7 eV
7th: 739.29 eV8th: 871,410 1 eV
Ordinary stateParamagnetic gas
Allotrope in the standard stateDioxygen O2
Other allotropesOzone O 3, singlet oxygen O 2*, cyclic ozone O 3, tetraoxygen O 4, octaoxygen O8
Density1.427 63 kg·m-3 T.P.N. (molecule O2)
Crystal systemCubic
Melting point−218.79 °C
Boiling point−182.95 °C
Fusion energy0.222 59 kJ·mol-1
Vaporization energy3.409 9 kJ·mol-1
Critical temperature−118.56 °C
Critical pressure5,043 MPa
Triple point−218.79 °C
Molar volume22.414×10-3 m3·mol-1
Speed of sound317 m·s-1 at 20 °C,5
Mass heat920 J·kg-1· K-1
Thermal conductivity0.026 74 W·m-1· K-1
CAS No.17778-80-2
SI & CNTP units, unless otherwise stated.

A molecule with the chemical formula O2, commonly called “oxygen” but “dioxygen” by chemists, consists of two oxygen atoms connected by covalent bonding: at normal conditions of temperature and pressure, dioxygen is a gas, which constitutes 20.8% of the volume of the Earth’s atmosphere at sea level.

Oxygen is a non-metal that very easily forms compounds, including oxides, with virtually all other chemical elements. This ease results in high formation energies but, kinetically, dioxygen is often not very reactive at room temperature. Thus a mixture of dioxygen and dihydrogen, iron or sulfur, etc., evolves only extremely slowly.

It is, by mass, the third most abundant element in the Universe after hydrogen and helium, and the most abundant of the elements of the Earth’s crust; oxygen on Earth is thus:

  • 86% of the mass of the oceans, in the form of water;
  • 46.4% of the mass of the earth’s crust, in particular in the form of oxides and silicates;
  • 23.1% of the mass of the air, in the form of dioxygen or ozone, or 1.2 × 1015 tons, or nearly 21% of the total volume of the atmosphere;
  • 62.5% of the mass of the human body;
  • Up to 88% of the mass of some marine animals.

The Earth was originally devoid of dioxygen. It was formed through photosynthesis by plants, algae and cyanobacteria, the latter having appeared perhaps 2.8 billion years ago. Dioxygen O2 is toxic to anaerobic organisms, which included the first forms of life on Earth, but is essential for the respiration of aerobic organisms, which constitute the vast majority of living species today. Cellular respiration is the set of metabolic pathways, such as the Krebs cycle and the respiratory chain, fueled for example by glycolysis and β-oxidation, by which a cell produces energy in the form of ATP and reducing power in the form of NADH + H + and FADH2.

By accumulating in the Earth’s atmosphere, oxygen O2 from photosynthesis formed an ozone layer at the base of the stratosphere under the effect of solar radiation. Ozone is an allotrope of oxygen with the chemical formula O3 even more oxidizing than dioxygen — which makes it an undesirable pollutant when present in the troposphere at ground level — but which has the particularity of absorbing ultraviolet rays from the Sun and thus protecting the biosphere from this harmful radiation: the ozone layer was the shield that allowed the first terrestrial plants to leave the oceans a long time ago. is nearly 475 million years old.

The oxygen content of the oceans has been falling significantly for several years. This deoxygenation of the ocean – due to global warming and agricultural fertilizer discharges – affects marine biodiversity. The oceans have lost 77 billion tons of oxygen over the past fifty years.

In industry, it has enormous importance as an oxidizer. In power plants, fuel is burned either with air or with pure oxygen (oxy-fuel process). Oxy-cracking of heavy petroleum fractions produces valuable compounds. For example, the chemical industry uses it for the production of acrylic acid, a very important monomer. Heterogeneous catalytic oxidation holds promise for benzoic acid production. It is also a promising raw material for the electrochemical synthesis of hydrogen peroxide. Air oxidation plays a very important role in the conversion of hazardous gases (CO, methane) into less harmful CO2.


Isotopes and stellar origin

Oxygen has seventeen isotopes with mass numbers ranging from 12 to 28. Naturally occurring oxygen is composed of three stable isotopes: oxygen-16-16O, oxygen-17-17O, and oxygen-18-18O . Oxygen is also assigned a standard atomic mass of 15.999 4 u. Oxygen-16 is the most abundant, with a natural abundance of 99.762%.

The majority of oxygen-16 is synthesized at the end of the helium fusion process in massive stars, but some is also produced during neon fusion reactions. Oxygen-17 is mainly derived from the fusion of hydrogen into helium during the NOC cycle. It is therefore a common isotope of the hydrogen combustion zones of stars. The majority of oxygen-18 is produced when the nitrogen-14-14-N made abundant by the CNO cycle captures a helium-4-4 He nucleus. Oxygen-18 is therefore commonly present in helium-rich areas of massive evolved stars.

Fourteen radioisotopes were identified. The most stable are oxygen-15-15 O with the longest half-life (122.24 s) and oxygen-14-14 O with a half-life of 70.606 s. All other radioactive isotopes have half-lives of less than 27 s and the majority have half-lives of less than 83 ms. Oxygen 12 12O has the shortest lifetime (580 × 10−24 s). The most common type of radioactive decay in isotopes lighter than oxygen-16 is nitrogen-producing positron emission. The most common type of decay for isotopes heavier than oxygen-18 is radioactivity β giving rise to fluorine.

Use of oxygen-18

Oxygen-18 is a paleoclimatic indicator used to determine the temperature in a region at a given time: the higher the 18 O/16 O isotopic ratio, the lower the corresponding temperature. This ratio can be determined from ice cores, as well as aragonite or calcite from some fossils.

This process is very useful for confirming or disproving a theory about terrestrial natural climate change such as Milanković parameters.

As a stable isotopic marker, it has been used to measure the unidirectional flow of oxygen absorbed during photosynthesis by the phenomenon of photorespiration. It has been shown that, before the increase in CO2 of the industrial era, half of the oxygen emitted by the leaves was reabsorbed. This reduced the yield of photosynthesis by half (Gerbaud and André, 1979-1980).

Importance of the presence of oxygen

List of the ten most common elements of the Milky Way (spectroscopic estimate) ZElementMass fractionin ppm
16Mass fraction in ppm440

Oxygen is the most mass-abundant chemical element in the biosphere, air, water and terrestrial rocks. It is also the third most abundant element in the universe after hydrogen and helium and accounts for about 0.9% of the Sun’s mass. It constitutes 49.2% of the mass of the Earth’s crust and is the main constituent of our oceans (88.8% of their mass). Dioxygen is the second most important component of the Earth’s atmosphere, accounting for 20.8% of its volume and 23.1% of its mass (some 10-15 tonnes). The Earth, by having such a high level of gaseous oxygen in its atmosphere, is an exception within the planets of the solar system: oxygen from the neighboring planets Mars (which represents only 0.1% of the volume of its atmosphere) and Venus have much lower concentrations. However, the dioxygen surrounding these other planets is only produced by ultraviolet rays acting on oxygen-containing molecules such as carbon dioxide.

The high and unusual concentration of dioxygen on Earth is the result of oxygen cycles. This biogeochemical cycle describes the movements of dioxygen within and between its three main reservoirs on Earth: the atmosphere, the biosphere and the lithosphere. The main factor in the realization of these cycles is photosynthesis which is the main responsible for the current dioxygen content on Earth. Dioxygen is essential to any ecosystem: photosynthetic living beings release dioxygen into the atmosphere while respiration and the decomposition of animals and plants consumes it. In the current equilibrium, production and consumption are carried out in the same proportions: each of these transfers corresponds to about 1/2000 of all atmospheric oxygen each year. Finally, oxygen is an essential component of the molecules found in any living being: amino acids, sugars, etc.

Oxygen also plays an important role in the aquatic environment. Increasing oxygen solubility at low temperatures has a noticeable impact on life in the oceans. For example, the density of living species is higher in polar waters due to the higher concentration of oxygen. Polluted waters containing plant nutrients such as nitrates or phosphates can stimulate algae growth through a process called eutrophication, and the breakdown of these organisms and other biomaterials can reduce the amount of dioxygen in eutrophic waters. Scientists assess this aspect of water quality by measuring the biological oxygen demand of the water or the amount of oxygen needed to return to a normal O2 concentration.

Single body

Structure of oxygen

molecule dioxygen
Compact molecular model of dioxygen

Under normal conditions of temperature and pressure, oxygen is in the form of an odorless and colorless gas, dioxygen has the chemical formula O2. Within this molecule, the two oxygen atoms are chemically bonded to each other in a triplet state. This bond, having an order of 2, is often represented in a simplified way by a double bond or by the association of a bond with two electrons and two bonds with three electrons. The triplet state of oxygen is the ground state of the dioxygen molecule. The electron configuration of the molecule presents two unpaired electrons occupying two degenerate molecular orbitals. These orbitals are called “antibonds” and lower the bond order from three to two, so that the bond of oxygen is weaker than the triple bond of dinitrogen for which all bonding atomic orbitals are filled but several antibonding orbitals are not.

In its normal triplet state, the dioxygen molecule is paramagnetic, that is, it acquires magnetization under the effect of a magnetic field. This is due to the magnetic spin moment of the unpaired electrons of the molecule as well as the negative exchange interaction between neighboring molecules of O2. Liquid oxygen can be attracted to a magnet so that in laboratory experiments, liquid oxygen can be kept in equilibrium against its own weight between the two poles of a powerful magnet.

Singlet oxygen is the name given to several excited species of the dioxygen molecule in which all spins are paired. In nature, it is commonly formed from water, during photosynthesis, using the energy of the sun’s rays. It is also produced in the troposphere through the photolysis of ozone by short-wavelength light rays and by the immune system as an active oxygen source. Carotenoids from photosynthetic organisms (but also sometimes animals) play a major role in absorbing energy from singlet oxygen and converting it to its deexcited ground state before it harms tissue.

Oxygen is very electronegative. It easily forms many ionic compounds with metals (oxides, hydroxides). It also forms ionocovalent compounds with nonmetals (examples: carbon dioxide, sulfur trioxide) and is used in many classes of organic molecules, for example, alcohols (R-OH), carbonyls R-CHO or R2CO and carboxylic acids (R-COOH).


The ordinary allotrope of oxygen on Earth is called “dioxygen”, with the chemical formula O2. It has a bond length of 121 pm and a binding energy of 498 kJ/mol. This is the form used by the most complex life forms, such as animals, during cellular respiration and the form that makes up most of the Earth’s atmosphere.

Trioxygen O3, usually called “ozone”, is a highly reactive allotrope of oxygen that is harmful to lung tissue. Ozone is a metastable gas produced in the upper layers of the atmosphere when dioxygen combines with atomic oxygen itself from the fragmentation of dioxygen by ultraviolet rays. Since ozone absorbs heavily in the ultraviolet range of the electromagnetic spectrum, the ozone layer contributes to the filtration of ultraviolet rays that strike the Earth. However, near the Earth’s surface, it is a pollutant produced by the decomposition on hot days of nitrogen oxides from the combustion of fossil fuels under the effect of ultraviolet solar rays. Since the 1970s, the concentration of ozone in the air at ground level has increased as a result of human activities.

The metastable molecule called “tetraoxygen” (O4) was discovered in 2001, and was previously thought to exist in one of the six phases of solid oxygen. It is proven in 2006 that this phase, obtained by pressurizing dioxygen at 20 GPa is in fact consisting of a rhombohedral cluster O8. This cluster is potentially a more potent oxidizer than dioxygen or ozone and could therefore be used in rocket propellants. A metallic phase, discovered in 1990, appears when solid oxygen is subjected to a pressure greater than 96 GPa and it was shown in 1998 that at very low temperatures, this phase becomes superconducting.

Physical properties of oxygen

Oxygen is more soluble in water than nitrogen. Water in equilibrium with air contains approximately one molecule of dissolved dioxygen for two molecules of dinitrogen, whereas in the atmosphere the ratio is approximately one molecule of dioxygen for four of dinitrogen.

The solubility of oxygen in water depends on temperature: about twice as much (14.6 mg/L) is dissolved at 0°C as at 20°C (7.6 mg/L). At 25°C and an air pressure of 1 atmosphere, fresh water contains about 6.04 mL of oxygen per liter while seawater contains about 4.95 mL/L. At 5°C, solubility increases to 9.0 mL/L of fresh water, 50% more than at 25°C, and to 7.2 mL/L of seawater, or 45% more.

Oxygen condenses at −182.95 °C and solidifies at −218.79 °C. The liquid and solid phases of dioxygen are both transparent with a slight coloration reminiscent of the blue color of the sky caused by absorption into red. High-purity liquid oxygen is usually obtained by fractional distillation of liquid air. Liquid oxygen can also be produced by condensation of air using liquid nitrogen as a coolant. It is an extremely reactive substance that must stay away from combustible materials.

Although oxygen-17 is stable, oxygen, composed mainly of oxygen-16, has a particularly low thermal neutron capture cross-section: 0.267 mb (weighted average over the three stable isotopes), which allows its use in nuclear reactors as an oxide in fuel, and in water as a coolant and moderator.

Nevertheless, the activation of oxygen by the neutrons of the core causes the formation of nitrogen 16 emitting a gamma radiation specially energetic (10.419 MeV), but whose period is only 7.13 s, which means that this radiation is extinguished quickly after the shutdown of the reactor.

Historical of Oxygen

First experiences

One of the earliest known experiments concerning the relationship between combustion and air was conducted by Philo of Byzantium, a Greek writer of the second century BC. In his book entitled Pneumatics, Philo observes that by burning a candle in an overturned container whose opening is immersed in water, it causes an elevation of water in the neck of the container containing the candle. Philo makes an incorrect conjecture, claiming that some of the air in the container turned into one of the four elements, fire, which may have escaped from the container because of the porosity of the glass. Many centuries later, Leonardo da Vinci drew on the work of Philo of Byzantium and observed that some of the air is consumed during combustion and respiration.

At the end of the seventeenth century, Robert Boyle proved that air is necessary for combustion. The English chemist John Mayow refined Boyle’s work by showing that combustion only needed a part of the air that he named spiritus nitroaereus or simply nitroaereus. In an experiment, he found that when he placed a lit mouse or candle in a closed container whose opening was immersed in water, the water level in the container rose and replaced a fourteenth of the volume of air before the subjects were extinguished. Therefore, he conjectures that nitroaereus is consumed both by combustion and respiration.

Mayow observes that antimony increases in mass when heated and deduces that nitroaereus must be associated with it. He also believes that the lungs separate nitroaereus from the air and pass it into the bloodstream and that animal heat and muscle movements result from nitroaereus’ reaction with certain substances in the body. Accounts of these and other experiments and Mayow’s ideas were published in 1668 in Tractatus Duo from De Respiratione.


Robert Hooke, Ole Borch, Mikhail Lomonosov and Pierre Bayen all managed to produce oxygen in experiments in the seventeenth and eighteenth centuries, but none of them recognized it as a chemical element. This is probably due in part to the scientific theory concerning combustion and corrosion and called “phlogistic” which was then the most widespread explanation for these phenomena.

Established in 1667 by the German chemist Johann Joachim Becher and modified by the chemist Georg Ernst Stahl in 1731, the phlogiston theory states that all combustible materials consist of two parts: a part called “phlogiston” that escapes when the substance containing it burns while the dephlogisticated part constitutes the true form of the substance.

Highly combustible materials that leave very little residue, such as wood or coal, are considered to contain predominantly phlogiston, while non-combustible substances that corrode such as metal, contain very little. Air plays no role in phlogiston theory, nor do the first experiments originally conducted to test the idea. On the contrary, the theory is based on the observation of what happens when an object burns and on the fact that the majority of objects appear lighter and seem to have lost something during the combustion process. To justify the fact that a material such as wood actually sees its mass increase when burning, Stahl claims that phlogiston has a negative mass. Indeed, the fact that metals also see their mass increase by rusting while they are supposed to lose phlogiston is one of the first clues to invalidate the theory of phlogiston.


Oxygen was first discovered by Swedish chemist Carl Wilhelm Scheele. He produced dioxygen by heating mercury oxide and various nitrates around 1772. Scheele named this gas “Feuerluf” (fire air) because it was the only known oxidizer and wrote an account of its discovery in a manuscript he entitled Traité Chimique de l’Air et du feu which he sent to his publisher in 1775 but which was not published until 1777.

At the same time, on August 1, 1774, an experiment led the British pastor Joseph Priestley to converge the Sun’s rays on a glass tube containing mercury oxide (HgO). This causes the release of gas he calls “dephlogisticated air”. He finds that the flame of the candles is brighter in this gas and that a mouse is more active and lives longer breathing it. After breathing the gas himself he wrote: “The sensation of [this gas] in my lungs was not appreciably different from that of ordinary air, but I had the impression that my breathing was particularly light and easy for some time afterward.” Priestley published his findings in 1775 in an article entitled An Account of Further Discoveries in Air included in the second volume of his book, Experiments and Observations on Different Kinds of Air.

The French chemist Antoine Laurent Lavoisier later claimed to have discovered this new substance independently of Priestley. However, Priestley visited Lavoisier in October 1774, told him about his experience and how he had released the gas. Scheele also sent a letter to Lavoisier on September 30, 1774, in which he described his own discovery of the hitherto unknown substance, but Lavoisier stated that he had never received it (a copy of the letter was found in Scheele’s belongings after his death).

Lavoisier’s contribution

Even if this is disputed in his time, Lavoisier’s contribution is undoubtedly to have carried out the first satisfactory quantitative experiments on oxidation and to have given the first correct explanation of how combustion takes place. His experiments, all begun in 1774, led to discredit the phlogiston theory and prove that the substance discovered by Priestley and Scheele was a chemical element.

In an experiment, Lavoisier observed that there is generally no increase in mass when tin and air are heated in a closed enclosure. He notices that the ambient air rushes into the enclosure when he opens it, which proves that some of the trapped air has been consumed. He also notes that the mass of tin has increased and that this increase corresponds to the same mass of air that rushed into the enclosure when it was opened. Other experiments as well as this one are detailed in his book On Combustion in General, published in 1777. In this work, he proves that air is a mixture of two gases: “vital air” which is essential for respiration and combustion and nitrogen (from the Greek ἄζωτον, “deprived of life”) which is useless to them.

Lavoisier renamed the “vital air” oxygen in 1777 from the Greek root ὀξύς  (oxys)  (acid, literally “bitter” after the taste of acids and -γενής (-genēs) (producer, literally “which generates”) because he mistakenly believed that oxygen is a constituent of all acids. Chemists, notably Sir Humphry Davy in 1812, finally proved that Lavoisier was wrong in this regard (it was actually hydrogen that was the basis of acid chemistry) but the name stuck.

Nineteenth-century and after

John Dalton’s atomic theory assumes that all elements are monoatomic and that atoms in compound bodies are in simple ratios. For example, Dalton assumes that the chemical formula for water is HO, giving oxygen an atomic mass eight times that of hydrogen as opposed to the current value which is about sixteen times that of hydrogen. 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 in 1811 Amedeo Avogadro managed to correctly interpret the composition of water on the basis of what is now called Avogadro’s law and the hypothesis of elementary diatomic molecules.

In December 1877, Louis Paul Cailletet in France and Raoul Pictet in Switzerland succeeded in producing, by two different processes and independently of each other, the first drops of liquid oxygen.

References (sources)