Oxygen is a chemical element with symbol O and atomic number 8.
It is a member of the chalcogen group on the periodic table, a highly
reactive nonmetal, and an oxidizing agent that readily forms oxides
with most elements as well as with other compounds. By mass, oxygen is
the third-most abundant element in the universe, after hydrogen and
helium. At standard temperature and pressure, two atoms of the element
bind to form dioxygen, a colorless and odorless diatomic gas with the
2. Diatomic oxygen gas constitutes 20.8% of the Earth's atmosphere. As
compounds including oxides, the element makes up almost half of the
Dioxygen is used in cellular respiration and many major classes of
organic molecules in living organisms contain oxygen, 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
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
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. 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
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.
2.1 Early experiments
2.2 Phlogiston theory
2.4 Lavoisier's contribution
2.5 Later history
3.1 Properties and molecular structure
3.3 Physical properties
3.4 Isotopes and stellar origin
4 Biological role of O2
Photosynthesis and respiration
4.2 Living organisms
4.3 Build-up in the atmosphere
5 Industrial production
7.2 Life support and recreational use
8.1 Oxides and other inorganic compounds
8.2 Organic compounds
9 Safety and precautions
Combustion and other hazards
10 See also
14 External links
The name oxygen was coined in 1777 by Antoine Lavoisier, 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.
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. 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
Leonardo da Vinci built on Philo's work by observing that a portion of
air is consumed during combustion and respiration.
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. 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. 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. 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.
Accounts of these and other experiments and ideas were published in
1668 in his work Tractatus duo in the tract "De respiratione".
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. 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.
Established in 1667 by the German alchemist J. J. Becher, and modified
by the chemist
Georg Ernst Stahl
Georg Ernst Stahl by 1731, 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.
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
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 in 1771–2. 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. That document was
published in 1777.
In the meantime, on August 1, 1774, an experiment conducted by the
Joseph Priestley focused sunlight on mercuric oxide
(HgO) inside a glass tube, which liberated a gas he named
"dephlogisticated air". 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." 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. Because he published
his findings first, Priestley is usually given priority in the
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 also posted a letter to Lavoisier on
September 30, 1774 that 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
Lavoisier conducted the first adequate quantitative experiments on
oxidation and gave the first correct explanation of how combustion
works. He used these and similar experiments, all started in 1774,
to discredit the phlogiston theory and to prove that the substance
discovered by Priestley and Scheele was a chemical element.
Antoine Lavoisier discredited the phlogiston theory.
In one experiment, Lavoisier observed that there was no overall
increase in weight when tin and air were heated in a closed
container. 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. 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.
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.
Chemists (such as Sir
Humphry Davy in 1812) eventually determined that
Lavoisier was wrong in this regard (hydrogen forms the basis for acid
chemistry), but by then the name was too well established.
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.
Robert H. Goddard
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, giving the atomic mass of
oxygen was 8 times that of hydrogen, instead of the modern value of
about 16. In 1805,
Joseph Louis Gay-Lussac
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
Avogadro's law and the diatomic elemental molecules in those
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. Just two days later, French physicist Louis Paul Cailletet
announced his own method of liquefying molecular oxygen. Only a
few drops of the liquid were produced in each case and no meaningful
analysis could be conducted.
Oxygen was liquified in a stable state
for the first time on March 29, 1883 by Polish scientists from
Zygmunt Wróblewski and Karol Olszewski.
In 1891 Scottish chemist
James Dewar was able to produce enough liquid
oxygen for study. The first commercially viable process for
producing liquid oxygen was independently developed in 1895 by German
Carl von Linde
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. Later, in 1901, oxyacetylene welding was
demonstrated for the first time by burning a mixture of acetylene and
2. This method of welding and cutting metal later became common.
In 1923, the American scientist
Robert H. Goddard
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.
Oxygen levels in the atmosphere are trending slightly downward
globally, possibly because of fossil-fuel burning.
Properties and molecular structure
Orbital diagram, after Barrett (2002), 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.
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 of the six 2p electrons after their partial filling of
the lowest π and π* orbitals.
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.[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
Liquid oxygen, temporarily suspended in a magnet owing to its
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
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.[c]
Singlet oxygen is a name given to several higher-energy species of
2 in which all the electron spins are paired. It is much more reactive
with common organic molecules than is molecular oxygen per se. In
nature, singlet oxygen is commonly formed from water during
photosynthesis, using the energy of sunlight. It is also produced
in the troposphere by the photolysis of ozone by light of short
wavelength, and by the immune system as a source of active
oxygen. 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
Main article: Allotropes of oxygen
Space-filling model representation of dioxygen (O2) molecule
The common allotrope of elemental oxygen on
Earth is called dioxygen,
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−1, which is smaller than the energy of other
double bonds or pairs of single bonds in the biosphere and responsible
for the exothermic reaction of O2 with any organic molecule.
Due to its energy content, 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.
3) is usually known as ozone and is a very reactive allotrope of
oxygen that is damaging to lung tissue.
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. 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. Near
the Earth's surface, it is a pollutant formed as a by-product of
automobile exhaust. At low earth orbit altitudes, sufficient
atomic oxygen is present to cause corrosion of spacecraft.
The metastable molecule tetraoxygen (O
4) was discovered in 2001, 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. 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. A metallic phase
was discovered in 1990 when solid oxygen is subjected to a pressure of
above 96 GPa and it was shown in 1998 that at very low
temperatures, this phase becomes superconducting.
Oxygen discharge (spectrum) tube
Liquid oxygen and solid oxygen
Oxygen dissolves more readily in water than nitrogen, and in
freshwater more readily than 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−1) dissolves at 0 °C than at
20 °C (7.6 mg·L−1). At 25 °C and 1 standard
atmosphere (101.3 kPa) of air, freshwater contains about
6.04 milliliters (mL) of oxygen per liter, and seawater
contains about 4.95 mL per liter. 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.
Oxygen gas dissolved in water at sea-level
Oxygen condenses at 90.20 K (−182.95 °C,
−297.31 °F), and freezes at 54.36 K (−218.79 °C,
−361.82 °F). Both liquid and solid 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
2 is usually obtained by the fractional distillation of liquefied
Liquid oxygen may also be condensed from air using liquid
nitrogen as a coolant.
Oxygen is a highly reactive substance and must be segregated from
The spectroscopy of molecular oxygen is associated with the
atmospheric processes of aurora and airglow. 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. Excited state singlet molecular oxygen is
responsible for red chemiluminescence in solution.
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
Most 16O is synthesized at the end of the helium fusion process in
massive stars but some is made in the neon burning process. 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. 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.
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. 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. The most
common decay mode of the isotopes lighter than 16O is β+
decay to yield nitrogen, and the most common mode for the
isotopes heavier than 18O is beta decay to yield fluorine.
Silicate minerals, Category:
Oxide minerals, Stellar
population, Cosmochemistry, and Astrochemistry
Ten most common elements in the
Milky Way Galaxy
Milky Way Galaxy estimated
Mass fraction in parts per million
71 × mass of oxygen (red bar)
23 × mass of oxygen (red bar)
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. About
0.9% of the Sun's mass is oxygen.
Oxygen constitutes 49.2% of the
Earth's crust by mass 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
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).[d]
Earth is unusual among the planets
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
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.
Cold water holds more dissolved O
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.
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.
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.
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. 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.
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. 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 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. 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
A simplified overall formula for photosynthesis is:
6 CO2 + 6 H
2O + photons → C
6 + 6 O
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. The O
2 remaining (after production of the water molecule) is released into
Oxygen is used in mitochondria to generate ATP during oxidative
phosphorylation. The reaction for aerobic respiration is essentially
the reverse of photosynthesis and is simplified as:
6 + 6 O
2 → 6 CO2 + 6 H
2O + 2880 kJ·mol−1
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 (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). A liter of blood can
dissolve 200 cm3 of O
Until the discovery of anaerobic metazoa, oxygen was thought to be
a requirement for all complex life.
Reactive oxygen species, such as superoxide ion (O−
2) and hydrogen peroxide (H
2), are reactive by-products of oxygen use in organisms. 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.
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.
An adult human at rest inhales 1.8 to 2.4 grams of oxygen per
minute. This amounts to more than 6 billion tonnes of oxygen
inhaled by humanity per year.[g]
Partial pressures of oxygen in the human body (PO2)
Arterial blood oxygen
Venous blood gas
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.
Build-up in the atmosphere
Main article: Geological history of oxygen
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). For the first billion years, any free oxygen
produced by these organisms combined with dissolved 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.
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
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
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,
2 levels have fluctuated between 15% and 30% by volume. Towards
the end of the
Carboniferous period (about 300 million years ago)
2 levels reached a maximum of 35% by volume, which may have
contributed to the large size of insects and amphibians at this
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.
At the current rate of photosynthesis it would take about
2,000 years to regenerate the entire O
2 in the present atmosphere.
Hofmann electrolysis apparatus used in electrolysis of water.
See also: Air separation,
Oxygen evolution, and Fractional
One hundred million tonnes of O
2 are extracted from air for industrial uses annually by two primary
methods. 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.
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. 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).
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
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). 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.
Breathing gas, Redox, and Combustion
An oxygen concentrator in an emphysema patient's house
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
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.
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. Carbon
monoxide poisoning, gas gangrene, and decompression sickness (the
'bends') are sometimes addressed with this therapy. Increased O
2 concentration in the lungs helps to displace carbon monoxide from
the heme group of hemoglobin.
Oxygen gas is poisonous to the
anaerobic bacteria that cause gas gangrene, so increasing its partial
pressure helps kill them.
Decompression sickness occurs in
divers who decompress too quickly after a dive, resulting in bubbles
of inert gas, mostly nitrogen and 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
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.
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), 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
People who climb mountains or fly in non-pressurized fixed-wing
aircraft sometimes have supplemental O
2 supplies.[h] 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. A steady stream of oxygen gas is then produced by the
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
2 exposure for a fee. 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. Available studies
support a performance boost from oxygen enriched mixtures only if it
is breathed during aerobic exercise.
Other recreational uses that do not involve breathing include
pyrotechnic applications, such as George Goble's five-second ignition
of barbecue grills.
Most commercially produced O
2 is used to smelt iron into steel.
Smelting of iron ore into steel consumes 55% of commercially produced
oxygen. 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,
2 and CO
2. The reactions are exothermic, so the temperature increases to
Another 25% of commercially produced oxygen is used by the chemical
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).
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.
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
Main article: Compounds of oxygen
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. 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).
Oxides and other inorganic compounds
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−1 per
hydrogen atom) to an adjacent oxygen atom in a separate molecule.
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.[i]
Oxides, such as iron oxide or rust, form when oxygen combines with
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 Fe
1 − xO, where x is usually around 0.05.
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
3, in bauxite and corundum), iron (iron(III) oxide Fe
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
3, and Na
5 are used as detergents and adhesives.
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. An exotic and unusual reaction occurs with PtF
6, which oxidizes oxygen to give O2+PtF6−, dioxygenyl
Acetone is an important feeder material in the chemical industry.
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
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
2CO) and phenol (C
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. Most of the
organic compounds that contain oxygen are not made by direct action of
2. Organic compounds important in industry and commerce that are made
by direct oxidation of a precursor include ethylene oxide and
Safety and precautions
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.
Main symptoms of oxygen toxicity
Oxygen gas (O
2) can be toxic at elevated partial pressures, leading to convulsions
and other health problems.[j]
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
2 by volume (about 30 kPa at standard pressure).
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.
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. 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.
Oxygen toxicity to the lungs and central nervous system can also occur
in deep scuba diving and surface supplied diving. Prolonged
breathing of an air mixture with an O
2 partial pressure more than 60 kPa can eventually lead to
permanent pulmonary fibrosis. Exposure to a 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).
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.
Oxygen is the
oxidant, not the fuel, but nevertheless the source of most of the
chemical energy released in combustion.
2 will allow combustion to proceed rapidly and energetically.
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. The fire that killed the
Apollo 1 crew in a launch
pad test spread so rapidly because the capsule was pressurized with
2 but at slightly more than atmospheric pressure, instead of the
1⁄3 normal pressure that would be used in a mission.[k]
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.
Geological history of oxygen
Hypoxia (environmental) for O
2 depletion in aquatic ecology
Hypoxia (medical), a lack of oxygen
Limiting oxygen concentration
View or order collections of articles
Period 2 elements
Chemical elements (sorted alphabetically)
Chemical elements (sorted by number)
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Find out more on's
^ 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
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^ An orbital is a concept from quantum mechanics that models an
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cyanobacteria. In fact, chloroplasts are thought to have evolved from
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of plants and algae.
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