Oxygen is the chemical element with the 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 formula O
. Diatomic oxygen gas constitutes 20.8% of the Earth's atmosphere. As compounds including oxides, the element makes up almost half of the Earth's crust.

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 replenished in Earth's atmosphere by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide. Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form (allotrope) of oxygen, ozone (O
), strongly absorbs ultraviolet UVB radiation and the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation. However, ozone present at the surface is a byproduct of smog and thus a pollutant.

Oxygen was isolated by Michael Sendivogius before 1604, but it is commonly believed that the element was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph Priestley in Wiltshire, in 1774. Priority is often given for Priestley because his work was published first. Priestley, however, called oxygen "dephlogisticated air", and did not recognize it as a chemical element. The name oxygen was coined in 1777 by Antoine Lavoisier, who first recognized oxygen as a chemical element and correctly characterized the role it plays in combustion.

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

Oxygen,  8O
A transparent beaker containing a light blue fluid with gas bubbles
Liquid oxygen boiling
AllotropesO2, O3 (Ozone)
Appearancegas: colorless
liquid and solid: pale blue
Standard atomic weight Ar, std(O)[15.9990315.99977] conventional: 15.999
Oxygen in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson


Atomic number (Z)8
Groupgroup 16 (chalcogens)
Periodperiod 2
Element category  reactive nonmetal
Electron configuration[He] 2s2 2p4
Electrons per shell
2, 6
Physical properties
Phase at STPgas
Melting point54.36 K ​(−218.79 °C, ​−361.82 °F)
Boiling point90.188 K ​(−182.962 °C, ​−297.332 °F)
Density (at STP)1.429 g/L
when liquid (at b.p.)1.141 g/cm3
Triple point54.361 K, ​0.1463 kPa
Critical point154.581 K, 5.043 MPa
Heat of fusion(O2) 0.444 kJ/mol
Heat of vaporization(O2) 6.82 kJ/mol
Molar heat capacity(O2) 29.378 J/(mol·K)
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K)       61 73 90
Atomic properties
Oxidation states−2, −1, +1, +2
ElectronegativityPauling scale: 3.44
Ionization energies
  • 1st: 1313.9 kJ/mol
  • 2nd: 3388.3 kJ/mol
  • 3rd: 5300.5 kJ/mol
  • (more)
Covalent radius66±2 pm
Van der Waals radius152 pm
Color lines in a spectral range
Spectral lines of oxygen
Other properties
Natural occurrenceprimordial
Crystal structurecubic
Cubic crystal structure for oxygen
Speed of sound330 m/s (gas, at 27 °C)
Thermal conductivity26.58×10−3  W/(m·K)
Magnetic orderingparamagnetic
Magnetic susceptibility+3449.0·10−6 cm3/mol (293 K)[1]
CAS Number7782-44-7
DiscoveryCarl Wilhelm Scheele (1771)
Named byAntoine Lavoisier (1777)
Main isotopes of oxygen
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct
16O 99.76% stable
17O 0.04% stable
18O 0.20% stable


Early experiments

One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck.[2] 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.[3]

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.[4] 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.[5] 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.[4] 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.[4] Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione".[5]

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.[6] 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.[7]

Established in 1667 by the German alchemist J. J. Becher, and modified by the chemist Georg Ernst Stahl by 1731,[8] 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.[3]

Highly combustible materials that leave little residue, such as wood or coal, were thought to be made mostly of phlogiston; non-combustible substances that corrode, such as iron, contained very little. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea; instead, it was based on observations of what happens when something burns, that most common objects appear to become lighter and seem to lose something in the process.[3]


Joseph Priestley is usually given priority in the discovery.

Polish alchemist, philosopher, and physician Michael Sendivogius in his work De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti (1604) described a substance contained in air, referring to it as 'cibus vitae' (food of life[9]), and this substance is identical with oxygen.[10] Sendivogius, during his experiments performed between 1598 and 1604, properly recognized that the substance is equivalent to the gaseous byproduct released by the thermal decomposition of potassium nitrate. In Bugaj’s view, the isolation of oxygen and the proper association of the substance to that part of air which is required for life, lends sufficient weight to the discovery of oxygen by Sendivogius.[10] This discovery of Sendivogius was however frequently denied by the generations of scientists and chemists which succeeded him.[9]

It is also commonly claimed that oxygen was first discovered by Swedish pharmacist Carl Wilhelm Scheele. He had produced oxygen gas by heating mercuric oxide and various nitrates in 1771–2.[11][12][3] Scheele called the gas "fire air" because it was then the only known agent to support combustion. He wrote an account of this discovery in a manuscript titled Treatise on Air and Fire, which he sent to his publisher in 1775. That document was published in 1777.[13]

In the meantime, on August 1, 1774, an experiment conducted by the British clergyman Joseph Priestley focused sunlight on mercuric oxide (HgO) contained in a glass tube, which liberated a gas he named "dephlogisticated air".[12] He noted that candles burned brighter in the gas and that a mouse was more active and lived longer while breathing it. After breathing the gas himself, Priestley wrote: "The feeling of it to my lungs was not sensibly different from that of common air, but I fancied that my breast felt peculiarly light and easy for some time afterwards."[6] 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.[3][14] Because he published his findings first, Priestley is usually given priority in the discovery.

The French chemist Antoine Laurent Lavoisier later claimed to have discovered the new substance independently. Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele also dispatched 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 death).[13]

Lavoisier's contribution

Lavoisier conducted the first adequate quantitative experiments on oxidation and gave the first correct explanation of how combustion works.[12] 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
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.[12] 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.[12] 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.[12]

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.[15] 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.[16]

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.[13]

Later history

Goddard and Rocket
Robert H. Goddard and a liquid oxygen-gasoline rocket

John Dalton's original atomic hypothesis presumed that all elements were monatomic and that the atoms in compounds would normally have the simplest atomic ratios with respect to one another. For example, Dalton assumed that water's formula was HO, leading to the conclusion that the atomic mass of oxygen was 8 times that of hydrogen, instead of the modern value of about 16.[17] In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen; and by 1811 Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called Avogadro's law and the diatomic elemental molecules in those gases.[18][a]

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.[19] Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen.[19] Only a few drops of the liquid were produced in each case and no meaningful analysis could be conducted. Oxygen was liquefied in a stable state for the first time on March 29, 1883 by Polish scientists from Jagiellonian University, Zygmunt Wróblewski and Karol Olszewski.[20]

In 1891 Scottish chemist James Dewar was able to produce enough liquid oxygen for study.[21] The first commercially viable process for producing liquid oxygen was independently developed in 1895 by German engineer Carl von Linde and British engineer William Hampson. Both men lowered the temperature of air until it liquefied and then distilled the component gases by boiling them off one at a time and capturing them separately.[22] 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.[22]

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.[22][23]

Oxygen levels in the atmosphere are trending slightly downward globally, possibly because of fossil-fuel burning.[24]


Properties and molecular structure

Oxygen molecule orbitals diagram
Orbital diagram, after Barrett (2002),[25] 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
, referred to as dioxygen.[26]

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.[25]

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
molecule is referred to as triplet oxygen.[27][b] The highest energy, partially filled orbitals are antibonding, and so their filling weakens the bond order from three to two. Because of its unpaired electrons, triplet oxygen reacts only slowly with most organic molecules, which have paired electron spins; this prevents spontaneous combustion.[28]

Liquid oxygen in a magnet 2
Liquid oxygen, temporarily suspended in a magnet owing to its paramagnetism

In the triplet form, O
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
molecules.[21] 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.[29][c]

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 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.[30] It is also produced in the troposphere by the photolysis of ozone by light of short wavelength,[31] and by the immune system as a source of active oxygen.[32] Carotenoids in photosynthetic organisms (and possibly animals) play a major role in absorbing energy from singlet oxygen and converting it to the unexcited ground state before it can cause harm to tissues.[33]


Oxygen molecule
Space-filling model representation of dioxygen (O2) molecule

The common allotrope of elemental oxygen on Earth is called dioxygen, O
, 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,[34] 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.[28][35] Due to its energy content, O2 is used by complex forms of life, such as animals, in cellular respiration. Other aspects of O
are covered in the remainder of this article.

Trioxygen (O
) is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue.[36] Ozone is produced in the upper atmosphere when O
combines with atomic oxygen made by the splitting of O
by ultraviolet (UV) radiation.[15] 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.[15] Near the Earth's surface, it is a pollutant formed as a by-product of automobile exhaust.[36] At low earth orbit altitudes, sufficient atomic oxygen is present to cause corrosion of spacecraft.[37]

The metastable molecule tetraoxygen (O
) was discovered in 2001,[38][39] 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.[40] 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.[38][39] A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa[41] and it was shown in 1998 that at very low temperatures, this phase becomes superconducting.[42]

Physical properties

Oxygen discharge tube
Oxygen discharge (spectrum) tube

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
for every 2 molecules of N
(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).[6][43] 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.[44] 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
5 °C 25 °C
Freshwater 9.0 mL 6.04 mL
Seawater 7.2 mL 4.95 mL

Oxygen condenses at 90.20 K (−182.95 °C, −297.31 °F), and freezes at 54.36 K (−218.79 °C, −361.82 °F).[45] 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.[46] Liquid oxygen may also be condensed from air using liquid nitrogen as a coolant.[47]

Oxygen is a highly reactive substance and must be segregated from combustible materials.[47]

The spectroscopy of molecular oxygen is associated with the atmospheric processes of aurora and airglow.[48] 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.[49] Excited state singlet molecular oxygen is responsible for red chemiluminescence in solution.[50]

Isotopes and stellar origin

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

Naturally occurring oxygen is composed of three stable isotopes, 16O, 17O, and 18O, with 16O being the most abundant (99.762% natural abundance).[51]

Most 16O is synthesized at the end of the helium fusion process in massive stars but some is made in the neon burning process.[52] 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.[52] 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.[52]

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.[51] 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.[51] The most common decay mode of the isotopes lighter than 16O is β+ decay[53][54][55] to yield nitrogen, and the most common mode for the isotopes heavier than 18O is beta decay to yield fluorine.[51]


Ten most common elements in the Milky Way Galaxy estimated spectroscopically[56]
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
6 Carbon 4,600
10 Neon 1,340
26 Iron 1,090
7 Nitrogen 960
14 Silicon 650
12 Magnesium 580
16 Sulfur 440

Oxygen is the most abundant chemical element by mass in the Earth's biosphere, air, sea and land. Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.[57] About 0.9% of the Sun's mass is oxygen.[12] Oxygen constitutes 49.2% of the Earth's crust by mass[58] as part of oxide compounds such as silicon dioxide and is the most abundant element by mass in the Earth's crust. It is also the major component of the world's oceans (88.8% by mass).[12] 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).[12][59][d] Earth is unusual among the planets of the Solar System in having such a high concentration of oxygen gas in its atmosphere: Mars (with 0.1% O
by volume) and Venus have much less. The O
surrounding those planets is produced solely by the action of ultraviolet radiation on oxygen-containing molecules such as carbon dioxide.

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.[60]

WOA09 sea-surf O2 AYool
Cold water holds more dissolved O

Free oxygen also occurs in solution in the world's water bodies. The increased solubility of O
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.[61] Water polluted with plant nutrients such as nitrates or phosphates may stimulate growth of algae by a process called eutrophication and the decay of these organisms and other biomaterials may reduce the O
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
needed to restore it to a normal concentration.[62]


Phanerozoic Climate Change
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.[63] 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.[63] 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.[64]

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.[65] 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

Photosynthesis and respiration

Simple photosynthesis overview
Photosynthesis splits water to liberate O
and fixes CO
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.[66] 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.[67]

A simplified overall formula for photosynthesis is:[68]

6 CO2 + 6 H
+ photonsC
+ 6 O

or simply

carbon dioxide + water + sunlight → glucose + dioxygen

Photolytic oxygen evolution occurs in the thylakoid membranes of photosynthetic organisms and requires the energy of four photons.[e] Many steps are involved, but the result is the formation of a proton gradient across the thylakoid membrane, which is used to synthesize adenosine triphosphate (ATP) via photophosphorylation.[69] The O
remaining (after production of the water molecule) is released into the atmosphere.[f]

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 O
→ 6 CO2 + 6 H
+ 2880 kJ/mol

In vertebrates, O
diffuses through membranes in the lungs and into red blood cells. Hemoglobin binds O
, changing color from bluish red to bright red[36] (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).[59] A liter of blood can dissolve 200 cm3 of O

Until the discovery of anaerobic metazoa,[70] oxygen was thought to be a requirement for all complex life.[71]

Reactive oxygen species, such as superoxide ion (O
) and hydrogen peroxide (H
), are reactive by-products of oxygen use in organisms.[59] 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.[69] Oxygen is damaging to obligately anaerobic organisms, which were the dominant form of early life on Earth until O
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.[72][73]

An adult human at rest inhales 1.8 to 2.4 grams of oxygen per minute.[74] This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year.[g]

Living organisms

Partial pressures of oxygen in the human body (PO2)
Unit Alveolar pulmonary
gas pressures
Arterial blood oxygen Venous blood gas
kPa 14.2 11[75]-13[75] 4.0[75]-5.3[75]
mmHg 107 75[76]-100[76] 30[77]-40[77]

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.[78]

Build-up in the atmosphere

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

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).[79] 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.[79][80]

The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the extant anaerobic organisms to extinction during the Great Oxygenation Event (oxygen catastrophe) about 2.4 billion years ago. Cellular respiration using O
enables aerobic organisms to produce much more ATP than anaerobic organisms.[81] Cellular respiration of O
occurs in all eukaryotes, including all complex multicellular organisms such as plants and animals.

Since the beginning of the Cambrian period 540 million years ago, atmospheric O
levels have fluctuated between 15% and 30% by volume.[82] Towards the end of the Carboniferous period (about 300 million years ago) atmospheric O
levels reached a maximum of 35% by volume,[82] which may have contributed to the large size of insects and amphibians at this time.[83]

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.[84]

At the current rate of photosynthesis it would take about 2,000 years to regenerate the entire O
in the present atmosphere.[85]

Industrial production

Hofmann voltameter fr
Hofmann electrolysis apparatus used in electrolysis of water.

One hundred million tonnes of O
are extracted from air for industrial uses annually by two primary methods.[13] The most common method is fractional distillation of liquefied air, with N
distilling as a vapor while O
is left as a liquid.[13]

The other primary method of producing O
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
.[13] 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).[86]

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
evolution from oxides and oxoacids. Chemical catalysts can be used as well, such as in chemical oxygen generators or oxygen candles that are used as part of the life-support equipment on submarines, and are still part of standard equipment on commercial airliners in case of depressurization emergencies. Another air separation method is forcing air to dissolve through ceramic membranes based on zirconium dioxide by either high pressure or an electric current, to produce nearly pure O


Compressed gas cylinders.mapp and oxygen.triddle
Oxygen and MAPP gas compressed gas cylinders with regulators

Oxygen storage methods include high pressure oxygen tanks, cryogenics and chemical compounds. For reasons of economy, oxygen is often transported in bulk as a liquid in specially insulated tankers, since one liter of liquefied oxygen is equivalent to 840 liters of gaseous oxygen at atmospheric pressure and 20 °C (68 °F).[13] 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.[13]



Home oxygen concentrator
An oxygen concentrator in an emphysema patient's house

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.[87]

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.[88]

Hyperbaric (high-pressure) medicine uses special oxygen chambers to increase the partial pressure of O
around the patient and, when needed, the medical staff.[89] Carbon monoxide poisoning, gas gangrene, and decompression sickness (the 'bends') are sometimes addressed with this therapy.[90] Increased O
concentration in the lungs helps to displace carbon monoxide from the heme group of hemoglobin.[91][92] Oxygen gas is poisonous to the anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them.[93][94] 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
as soon as possible helps to redissolve the bubbles back into the blood so that these excess gasses can be exhaled naturally through the lungs.[87][95][96] Normobaric oxygen administration at the highest available concentration is frequently used as first aid for any diving injury that may involve inert gas bubble formation in the tissues. There is epidemiological support for its use from a statistical study of cases recorded in a long term database.[97][98][99]

Life support and recreational use

Wisoff on the Arm - GPN-2000-001069
Low pressure pure O
is used in space suits.

An application of O
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
. This trade-off of higher oxygen concentration for lower pressure is needed to maintain suit flexibility.[100][101]

Scuba and surface-supplied underwater divers and submariners also rely on artificially delivered O
. 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
use in diving at pressures higher than atmospheric is usually limited to rebreathers, or decompression at relatively shallow depths (~6 meters depth, or less),[102][103] 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
with other gases, such as nitrogen or helium, to prevent oxygen toxicity.[102]

People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental O
supplies.[h] Pressurized commercial airplanes have an emergency supply of O
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.[62] A steady stream of oxygen gas is then produced by the exothermic reaction.

Oxygen, as a mild euphoric, has a history of recreational use in oxygen bars and in sports. Oxygen bars are establishments found in the United States since the late 1990s that offer higher than normal O
exposure for a minimal fee.[104] 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.[104] Available studies support a performance boost from oxygen enriched mixtures only if it is breathed during aerobic exercise.[105]

Other recreational uses that do not involve breathing include pyrotechnic applications, such as George Goble's five-second ignition of barbecue grills.[106]


Clabecq JPG01
Most commercially produced O
is used to smelt and/or decarburize iron.

Smelting of iron ore into steel consumes 55% of commercially produced oxygen.[62] 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.[62]

Another 25% of commercially produced oxygen is used by the chemical industry.[62] 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).[62]

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.[62] Oxygen is used in oxyacetylene welding burning acetylene with O
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


Stilles Mineralwasser
Water (H
) 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.[108] 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).[109]

Oxides and other inorganic compounds

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 per hydrogen atom) to an adjacent oxygen atom in a separate molecule.[110] 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.[111][i]

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

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements to give corresponding oxides. The surface of most metals, such as aluminium and titanium, are oxidized in the presence of air and become coated with a thin film of oxide that passivates the metal and slows further corrosion. Many oxides of the transition metals are non-stoichiometric compounds, with slightly less metal than the chemical formula would show. For example, the mineral FeO (wüstite) is written as Fe
1 − x
, where x is usually around 0.05.[112]

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 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.

Water-soluble silicates in the form of Na
, Na
, and Na
are used as detergents and adhesives.[113]

Oxygen also acts as a ligand for transition metals, forming transition metal dioxygen complexes, which feature metal–O
. This class of compounds includes the heme proteins hemoglobin and myoglobin.[114] An exotic and unusual reaction occurs with PtF
, which oxidizes oxygen to give O2+PtF6, dioxygenyl hexafluoroplatinate.[115]

Organic compounds

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 (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. 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.[116] Most of the organic compounds that contain oxygen are not made by direct action of O
. Organic compounds important in industry and commerce that are made by direct oxidation of a precursor include ethylene oxide and peracetic acid.[113]

Safety and precautions

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.[117]


Symptoms of oxygen toxicity
Main symptoms of oxygen toxicity[118]

Oxygen gas (O
) can be toxic at elevated partial pressures, leading to convulsions and other health problems.[102][j][119] 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).[6]

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.[6]

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.[100][120] 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.[121]

Oxygen toxicity to the lungs and central nervous system can also occur in deep scuba diving and surface supplied diving.[6][102] Prolonged breathing of an air mixture with an O
partial pressure more than 60 kPa can eventually lead to permanent pulmonary fibrosis.[122] 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 (217 ft) or more of depth; the same thing can occur by breathing 100% O
at only 6 m (20 ft).[122][123][124][125]

Combustion and other hazards

Apollo 1 fire
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.

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.[28][126] Oxygen is the oxidant, not the fuel, but nevertheless the source of most of the chemical energy released in combustion.[28][35]

Concentrated O
will allow combustion to proceed rapidly and energetically.[126] 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.[126] 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.[k][128]

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.[126]

See also


  1. ^ These results were mostly ignored until 1860. Part of this rejection was due to the belief that atoms of one element would have no chemical affinity towards atoms of the same element, and part was due to apparent exceptions to Avogadro's law that were not explained until later in terms of dissociating molecules.
  2. ^ An orbital is a concept from quantum mechanics that models an electron as a wave-like particle that has a spatial distribution about an atom or molecule.
  3. ^ Oxygen's paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine the purity of gaseous oxygen. ("Company literature of Oxygen analyzers (triplet)". Servomex. Archived from the original on March 8, 2008. Retrieved December 15, 2007.)
  4. ^ Figures given are for values up to 50 miles (80 km) above the surface
  5. ^ Thylakoid membranes are part of chloroplasts in algae and plants while they simply are one of many membrane structures in cyanobacteria. In fact, chloroplasts are thought to have evolved from cyanobacteria that were once symbiotic partners with the progenitors of plants and algae.
  6. ^ Water oxidation is catalyzed by a manganese-containing enzyme complex known as the oxygen evolving complex (OEC) or water-splitting complex found associated with the lumenal side of thylakoid membranes. Manganese is an important cofactor, and calcium and chloride are also required for the reaction to occur. (Raven 2005)
  7. ^ (1.8 grams/min/person)×(60 min/h)×(24 h/day)×(365 days/year)×(6.6 billion people)/1,000,000 g/t=6.24 billion tonnes
  8. ^ 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 an arc from an electrical spark.[127]


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General references

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

External links

Anaerobic organism

An anaerobic organism or anaerobe is any organism that does not require oxygen for growth. It may react negatively or even die if free oxygen is present. (In contrast, an aerobic organism (aerobe) is an organism that requires an oxygenated environment.)

An anaerobic organism may be unicellular (e.g. protozoans, bacteria) or multicellular. For practical purposes, there are three categories of anaerobe: obligate anaerobes, which are harmed by the presence of oxygen; aerotolerant organisms, which cannot use oxygen for growth but tolerate its presence; and facultative anaerobes, which can grow without oxygen but use oxygen if it is present.

Apollo 13

Apollo 13 was the seventh manned mission in the Apollo space program and the third intended to land on the Moon. The craft was launched on April 11, 1970 from the Kennedy Space Center, Florida, but the lunar landing was aborted after an oxygen tank exploded two days later, crippling the service module (SM) upon which the command module (CM) had depended. Despite great hardship caused by limited power, loss of cabin heat, shortage of potable water, and the critical need to make makeshift repairs to the carbon dioxide removal system, the crew returned safely to Earth on April 17, 1970, six days after launch.

The flight passed the far side of the Moon at an altitude of 254 kilometers (137 nautical miles) above the lunar surface, and 400,171 km (248,655 mi) from Earth, a spaceflight record marking the farthest humans have ever traveled from Earth. The mission was commanded by James A. Lovell with John L. "Jack" Swigert as Command Module Pilot and Fred W. Haise as Lunar Module Pilot. Swigert was a late replacement for the original CM pilot Ken Mattingly, who was grounded by the flight surgeon after exposure to German measles.

The story of the Apollo 13 mission has been dramatized multiple times, most notably in the 1995 film Apollo 13.


Asphyxia or asphyxiation is a condition of severely deficient supply of oxygen to the body that arises from abnormal breathing. An example of asphyxia is choking. Asphyxia causes generalized hypoxia, which affects primarily the tissues and organs. There are many circumstances that can induce asphyxia, all of which are characterized by an inability of an individual to acquire sufficient oxygen through breathing for an extended period of time. Asphyxia can cause coma or death.

In 2015 about 9.8 million cases of unintentional suffocation occurred which resulted in 35,600 deaths. The word asphyxia is from Ancient Greek α- "without" and σφύξις sphyxis, "squeeze" (throb of heart).


Blood is a body fluid in humans and other animals that delivers necessary substances such as nutrients and oxygen to the cells and transports metabolic waste products away from those same cells.In vertebrates, it is composed of blood cells suspended in blood plasma. Plasma, which constitutes 55% of blood fluid, is mostly water (92% by volume), and contains proteins, glucose, mineral ions, hormones, carbon dioxide (plasma being the main medium for excretory product transportation), and blood cells themselves. Albumin is the main protein in plasma, and it functions to regulate the colloidal osmotic pressure of blood. The blood cells are mainly red blood cells (also called RBCs or erythrocytes), white blood cells (also called WBCs or leukocytes) and platelets (also called thrombocytes). The most abundant cells in vertebrate blood are red blood cells. These contain hemoglobin, an iron-containing protein, which facilitates oxygen transport by reversibly binding to this respiratory gas and greatly increasing its solubility in blood. In contrast, carbon dioxide is mostly transported extracellularly as bicarbonate ion transported in plasma.

Vertebrate blood is bright red when its hemoglobin is oxygenated and dark red when it is deoxygenated. Some animals, such as crustaceans and mollusks, use hemocyanin to carry oxygen, instead of hemoglobin. Insects and some mollusks use a fluid called hemolymph instead of blood, the difference being that hemolymph is not contained in a closed circulatory system. In most insects, this "blood" does not contain oxygen-carrying molecules such as hemoglobin because their bodies are small enough for their tracheal system to suffice for supplying oxygen.

Jawed vertebrates have an adaptive immune system, based largely on white blood cells. White blood cells help to resist infections and parasites. Platelets are important in the clotting of blood. Arthropods, using hemolymph, have hemocytes as part of their immune system.

Blood is circulated around the body through blood vessels by the pumping action of the heart. In animals with lungs, arterial blood carries oxygen from inhaled air to the tissues of the body, and venous blood carries carbon dioxide, a waste product of metabolism produced by cells, from the tissues to the lungs to be exhaled.

Medical terms related to blood often begin with hemo- or hemato- (also spelled haemo- and haemato-) from the Greek word αἷμα (haima) for "blood". In terms of anatomy and histology, blood is considered a specialized form of connective tissue, given its origin in the bones and the presence of potential molecular fibers in the form of fibrinogen.

Cellular respiration

Cellular respiration is a set of metabolic reactions and processes that take place in the cells of organisms to convert biochemical energy from nutrients into adenosine triphosphate (ATP), and then release waste products. The reactions involved in respiration are catabolic reactions, which break large molecules into smaller ones, releasing energy in the process, as weak so-called "high-energy" bonds are replaced by stronger bonds in the products. Respiration is one of the key ways a cell releases chemical energy to fuel cellular activity. Cellular respiration is considered an exothermic redox reaction which releases heat. The overall reaction occurs in a series of biochemical steps, most of which are redox reactions themselves. Although cellular respiration is technically a combustion reaction, it clearly does not resemble one when it occurs in a living cell because of the slow release of energy from the series of reactions.

Nutrients that are commonly used by animal and plant cells in respiration include sugar, amino acids and fatty acids, and the most common oxidizing agent (electron acceptor) is molecular oxygen (O2). The chemical energy stored in ATP (its third phosphate group is weakly bonded to the rest of the molecule and is cheaply broken allowing stronger bonds to form, thereby transferring energy for use by the cell) can then be used to drive processes requiring energy, including biosynthesis, locomotion or transportation of molecules across cell membranes.

Circulatory system

The circulatory system, also called the cardiovascular system or the vascular system, is an organ system that permits blood to circulate and transport nutrients (such as amino acids and electrolytes), oxygen, carbon dioxide, hormones, and blood cells to and from the cells in the body to provide nourishment and help in fighting diseases, stabilize temperature and pH, and maintain homeostasis.

The circulatory system includes the lymphatic system, which circulates lymph. The passage of lymph for example takes much longer than that of blood. Blood is a fluid consisting of plasma, red blood cells, white blood cells, and platelets that is circulated by the heart through the vertebrate vascular system, carrying oxygen and nutrients to and waste materials away from all body tissues. Lymph is essentially recycled excess blood plasma after it has been filtered from the interstitial fluid (between cells) and returned to the lymphatic system. The cardiovascular (from Latin words meaning "heart" and "vessel") system comprises the blood, heart, and blood vessels. The lymph, lymph nodes, and lymph vessels form the lymphatic system, which returns filtered blood plasma from the interstitial fluid (between cells) as lymph.

The circulatory system of the blood is seen as having two components, a systemic circulation and a pulmonary circulation.While humans, as well as other vertebrates, have a closed cardiovascular system (meaning that the blood never leaves the network of arteries, veins and capillaries), some invertebrate groups have an open cardiovascular system. The lymphatic system, on the other hand, is an open system providing an accessory route for excess interstitial fluid to be returned to the blood. The more primitive, diploblastic animal phyla lack circulatory systems.

Many diseases affect the circulatory system. This includes cardiovascular disease, affecting the cardiovascular system, and lymphatic disease affecting the lymphatic system. Cardiologists are medical professionals which specialise in the heart, and cardiothoracic surgeons specialise in operating on the heart and its surrounding areas. Vascular surgeons focus on other parts of the circulatory system.


Combustion, or burning, is a high-temperature exothermic redox chemical reaction between a fuel (the reductant) and an oxidant, usually atmospheric oxygen, that produces oxidized, often gaseous products, in a mixture termed as smoke. Combustion in a fire produces a flame, and the heat produced can make combustion self-sustaining. Combustion is often a complicated sequence of elementary radical reactions. Solid fuels, such as wood and coal, first undergo endothermic pyrolysis to produce gaseous fuels whose combustion then supplies the heat required to produce more of them. Combustion is often hot enough that incandescent light in the form of either glowing or a flame is produced. A simple example can be seen in the combustion of hydrogen and oxygen into water vapor, a reaction commonly used to fuel rocket engines. This reaction releases 242 kJ/mol of heat and reduces the enthalpy accordingly (at constant temperature and pressure):

2H2(g) + O2(g) → 2 H2O(g)Combustion of an organic fuel in air is always exothermic because the double bond in O2 is much weaker than other double bonds or pairs of single bonds, and therefore the formation of the stronger bonds in the combustion products CO2 and  H2O results in the release of energy. The bond energies in the fuel play only a minor role, since they are similar to those in the combustion products; e.g., the sum of the bond energies of CH4 is nearly the same as that of CO2. The heat of combustion is approximately -418 kJ per mole of O2 used up in the combustion reaction, and can be estimated from the elemental composition of the fuel.Uncatalyzed combustion in air requires fairly high temperatures. Complete combustion is stoichiometric with respect to the fuel, where there is no remaining fuel, and ideally, no remaining oxidant. Thermodynamically, the chemical equilibrium of combustion in air is overwhelmingly on the side of the products. However, complete combustion is almost impossible to achieve, since the chemical equilibrium is not necessarily reached, or may contain unburnt products such as carbon monoxide, hydrogen and even carbon (soot or ash). Thus, the produced smoke is usually toxic and contains unburned or partially oxidized products. Any combustion at high temperatures in atmospheric air, which is 78 percent nitrogen, will also create small amounts of several nitrogen oxides, commonly referred to as NOx, since the combustion of nitrogen is thermodynamically favored at high, but not low temperatures. Since combustion is rarely clean, flue gas cleaning or catalytic converters may be required by law.

Fires occur naturally, ignited by lightning strikes or by volcanic products. Combustion (fire) was the first controlled chemical reaction discovered by humans, in the form of campfires and bonfires, and continues to be the main method to produce energy for humanity. Usually, the fuel is carbon, hydrocarbons or more complicated mixtures such as wood that contains partially oxidized hydrocarbons. The thermal energy produced from combustion of either fossil fuels such as coal or oil, or from renewable fuels such as firewood, is harvested for diverse uses such as cooking, production of electricity or industrial or domestic heating. Combustion is also currently the only reaction used to power rockets. Combustion is also used to destroy (incinerate) waste, both nonhazardous and hazardous.

Oxidants for combustion have high oxidation potential and include atmospheric or pure oxygen, chlorine, fluorine, chlorine trifluoride, nitrous oxide and nitric acid. For instance, hydrogen burns in chlorine to form hydrogen chloride with the liberation of heat and light characteristic of combustion. Although usually not catalyzed, combustion can be catalyzed by platinum or vanadium, as in the contact process.

Great Oxygenation Event

The Great Oxygenation Event, the beginning of which is commonly known in scientific media as the Great Oxidation Event (GOE, also called the Oxygen Catastrophe, Oxygen Crisis, Oxygen Holocaust, Oxygen Revolution, or Great Oxidation) was the biologically induced appearance of dioxygen (O2) in Earth's atmosphere. Geological, isotopic, and chemical evidence suggests a start of around 2.45 billion years ago (2.45 Ga), during the Siderian period, at the beginning of the Proterozoic eon. The causes of the event remain unclear. As of 2016, the geochemical and biomarker evidence for the development of oxygenic photosynthesis before the Great Oxidation Event is inconclusive.

Oceanic cyanobacteria, were the first microbes to produce oxygen by photosynthesis. They evolved into tufted microbial mats more than 2.3 billion years ago, approximately 200 million years before the GOE. The free oxygen produced during this time was chemically captured by dissolved iron, converting iron and to magnetite () which is insoluable in water, and sank to the bottom of the shallow seas to create massive, large scale, banded iron formations. Some of the oxygen was captured by organic matter. The GOE started after these oxygen sinks were filled.

The increased production of oxygen set Earth's original atmosphere off-balance. Free oxygen is toxic to obligate anaerobic organisms, and the rising concentrations may have destroyed most such organisms.

A spike in chromium contained in ancient rock-deposits formed underwater shows the accumulation washed off from the continental shelves. Chromium is not easily dissolved and its release from rocks requires the presence of a powerful acid. One such acid, sulfuric acid (H2SO4), might have formed through bacterial reactions with pyrite. Mats of oxygen-producing cyanobacteria can produce a thin layer, one or two millimeters thick, of oxygenated water in an otherwise anoxic environment even under thick ice; before oxygen started accumulating in the atmosphere, these organisms would already have adapted to oxygen. Additionally, the free oxygen would have reacted with atmospheric methane, a greenhouse gas, greatly reducing its concentration and triggering the Huronian glaciation, possibly the longest episode of glaciation in Earth's history and called "snowball Earth".

Eventually, the evolution of aerobic organisms that consumed oxygen established an equilibrium in its availability. Free oxygen has been an important constituent of the atmosphere ever since.


Hemoglobin (American) or haemoglobin (British) (), abbreviated Hb or Hgb, is the iron-containing oxygen-transport metalloprotein in the red blood cells (erythrocytes) of almost all vertebrates (the exception being the fish family Channichthyidae) as well as the tissues of some invertebrates. Haemoglobin in the blood carries oxygen from the lungs or gills to the rest of the body (i.e. the tissues). There it releases the oxygen to permit aerobic respiration to provide energy to power the functions of the organism in the process called metabolism. A healthy individual has 12 to 16 grams of haemoglobin in every 100 ml of blood.

In mammals, the protein makes up about 96% of the red blood cells' dry content (by weight), and around 35% of the total content (including water). Haemoglobin has an oxygen-binding capacity of 1.34 mL O2 per gram, which increases the total blood oxygen capacity seventy-fold compared to dissolved oxygen in blood. The mammalian hemoglobin molecule can bind (carry) up to four oxygen molecules.Hemoglobin is involved in the transport of other gases: It carries some of the body's respiratory carbon dioxide (about 20–25% of the total) as carbaminohemoglobin, in which CO2 is bound to the heme protein. The molecule also carries the important regulatory molecule nitric oxide bound to a globin protein thiol group, releasing it at the same time as oxygen.Haemoglobin is also found outside red blood cells and their progenitor lines. Other cells that contain haemoglobin include the A9 dopaminergic neurons in the substantia nigra, macrophages, alveolar cells, lungs, retinal pigment epithelium, hepatocytes, mesangial cells in the kidney, endometrial cells, cervical cells and vaginal epithelial cells. In these tissues, haemoglobin has a non-oxygen-carrying function as an antioxidant and a regulator of iron metabolism.Haemoglobin and haemoglobin-like molecules are also found in many invertebrates, fungi, and plants. In these organisms, haemoglobins may carry oxygen, or they may act to transport and regulate other small molecules and ions such as carbon dioxide, nitric oxide, hydrogen sulfide and sulfide. A variant of the molecule, called leghaemoglobin, is used to scavenge oxygen away from anaerobic systems, such as the nitrogen-fixing nodules of leguminous plants, before the oxygen can poison (deactivate) the system.

Hyperbaric medicine

Hyperbaric medicine is medical treatment in which an ambient pressure greater than sea level atmospheric pressure is a necessary component. The treatment comprises hyperbaric oxygen therapy (HBOT), the medical use of oxygen at an ambient pressure higher than atmospheric pressure, and therapeutic recompression for decompression illness, intended to reduce the injurious effects of systemic gas bubbles by physically reducing their size and providing improved conditions for elimination of bubbles and excess dissolved gas.

The equipment required for hyperbaric oxygen treatment consists of a pressure chamber, which may be of rigid or flexible construction, and a means of delivering 100% oxygen. Operation is performed to a predetermined schedule by trained personnel who monitor the patient and may adjust the schedule as required. HBOT found early use in the treatment of decompression sickness, and has also shown great effectiveness in treating conditions such as gas gangrene and carbon monoxide poisoning. More recent research has examined the possibility that it may also have value for other conditions such as cerebral palsy and multiple sclerosis, but no significant evidence has been found.

Therapeutic recompression is usually also provided in a hyperbaric chamber. It is the definitive treatment for decompression sickness and may also be used to treat arterial gas embolism caused by pulmonary barotrauma of ascent. In emergencies divers may sometimes be treated by in-water recompression if a chamber is not available and suitable diving equipment to reasonably secure the airway is available.

A number of hyperbaric treatment schedules have been published over the years for both therapeutic recompression and hyperbaric oxygen therapy for other conditions.

Hypoxia (medical)

Hypoxia is a condition in which the body or a region of the body is deprived of adequate oxygen supply at the tissue level. Hypoxia may be classified as either generalized, affecting the whole body, or local, affecting a region of the body. Although hypoxia is often a pathological condition, variations in arterial oxygen concentrations can be part of the normal physiology, for example, during hypoventilation training or strenuous physical exercise.

Hypoxia differs from hypoxemia and anoxemia in that hypoxia refers to a state in which oxygen supply is insufficient, whereas hypoxemia and anoxemia refer specifically to states that have low or zero arterial oxygen supply. Hypoxia in which there is complete deprivation of oxygen supply is referred to as anoxia.

Generalized hypoxia occurs in healthy people when they ascend to high altitude, where it causes altitude sickness leading to potentially fatal complications: high altitude pulmonary edema (HAPE) and high altitude cerebral edema (HACE). Hypoxia also occurs in healthy individuals when breathing mixtures of gases with a low oxygen content, e.g. while diving underwater especially when using closed-circuit rebreather systems that control the amount of oxygen in the supplied air. Mild, non-damaging intermittent hypoxia is used intentionally during altitude training to develop an athletic performance adaptation at both the systemic and cellular level.Hypoxia is a common complication of preterm birth in newborn infants. Because the lungs develop late in pregnancy, premature infants frequently possess underdeveloped lungs. To improve lung function, doctors frequently place infants at risk of hypoxia inside incubators (also known as humidicribs) that provide continuous positive airway pressure.

Mount Everest

Mount Everest, known in Nepali as Sagarmatha (सगरमाथा) and in Tibetan as Chomolungma (ཇོ་མོ་གླང་མ), is Earth's highest mountain above sea level, located in the Mahalangur Himal sub-range of the Himalayas. The international border between Nepal (Province No. 1) and China (Tibet Autonomous Region) runs across its summit point.

The current official elevation of 8,848 m (29,029 ft), recognized by China and Nepal, was established by a 1955 Indian survey and subsequently confirmed by a Chinese survey in 1975. In 2005, China remeasured the rock height of the mountain, with a result of 8844.43 m (29,017 ft). There followed an argument between China and Nepal as to whether the official height should be the rock height (8,844 m, China) or the snow height (8,848 m, Nepal). In 2010, an agreement was reached by both sides that the height of Everest is 8,848 m, and Nepal recognizes China's claim that the rock height of Everest is 8,844 m.In 1865, Everest was given its official English name by the Royal Geographical Society, upon a recommendation by Andrew Waugh, the British Surveyor General of India. As there appeared to be several different local names, Waugh chose to name the mountain after his predecessor in the post, Sir George Everest, despite Everest's objections.Mount Everest attracts many climbers, some of them highly experienced mountaineers. There are two main climbing routes, one approaching the summit from the southeast in Nepal (known as the "standard route") and the other from the north in Tibet. While not posing substantial technical climbing challenges on the standard route, Everest presents dangers such as altitude sickness, weather, and wind, as well as significant hazards from avalanches and the Khumbu Icefall. As of 2017, nearly 300 people have died on Everest, many of whose bodies remain on the mountain.The first recorded efforts to reach Everest's summit were made by British mountaineers. As Nepal did not allow foreigners into the country at the time, the British made several attempts on the north ridge route from the Tibetan side. After the first reconnaissance expedition by the British in 1921 reached 7,000 m (22,970 ft) on the North Col, the 1922 expedition pushed the north ridge route up to 8,320 m (27,300 ft), marking the first time a human had climbed above 8,000 m (26,247 ft). Seven porters were killed in an avalanche on the descent from the North Col. The 1924 expedition resulted in one of the greatest mysteries on Everest to this day: George Mallory and Andrew Irvine made a final summit attempt on 8 June but never returned, sparking debate as to whether or not they were the first to reach the top. They had been spotted high on the mountain that day but disappeared in the clouds, never to be seen again, until Mallory's body was found in 1999 at 8,155 m (26,755 ft) on the north face. Tenzing Norgay and Edmund Hillary made the first official ascent of Everest in 1953, using the southeast ridge route. Norgay had reached 8,595 m (28,199 ft) the previous year as a member of the 1952 Swiss expedition. The Chinese mountaineering team of Wang Fuzhou, Gonpo, and Qu Yinhua made the first reported ascent of the peak from the north ridge on 25 May 1960.

Oxygen (TV channel)

Oxygen is an American pay television channel that is owned by NBCUniversal, which is owned by Comcast. The channel primarily airs true crime programming targeted towards women.

The network was founded by Geraldine Laybourne, and carried a format focused on lifestyle and entertainment programming oriented towards women, similar to competing channels such as Lifetime. NBCUniversal acquired the network in 2007; under NBC ownership, the network increasingly produced reality shows aimed at the demographic, and was relaunched in 2014 to target a "modern", younger female audience. After the network experienced ratings successes with a programming block dedicated to such programming, Oxygen was relaunched in mid-2017 to focus primarily on true-crime programs.

As of February 2015, approximately 77.5 million American households (66.5% of households with television) receive Oxygen. Under its current format, the network primarily competes with Investigation Discovery and HLN.

Oxygen toxicity

Oxygen toxicity is a condition resulting from the harmful effects of breathing molecular oxygen (O2) at increased partial pressures. Severe cases can result in cell damage and death, with effects most often seen in the central nervous system, lungs, and eyes. Historically, the central nervous system condition was called the Paul Bert effect, and the pulmonary condition the Lorrain Smith effect, after the researchers who pioneered the discoveries and descriptions in the late 19th century. Oxygen toxicity is a concern for underwater divers, those on high concentrations of supplemental oxygen (particularly premature babies), and those undergoing hyperbaric oxygen therapy.

The result of breathing increased partial pressures of oxygen is hyperoxia, an excess of oxygen in body tissues. The body is affected in different ways depending on the type of exposure. Central nervous system toxicity is caused by short exposure to high partial pressures of oxygen at greater than atmospheric pressure. Pulmonary and ocular toxicity result from longer exposure to increased oxygen levels at normal pressure. Symptoms may include disorientation, breathing problems, and vision changes such as myopia. Prolonged exposure to above-normal oxygen partial pressures, or shorter exposures to very high partial pressures, can cause oxidative damage to cell membranes, collapse of the alveoli in the lungs, retinal detachment, and seizures. Oxygen toxicity is managed by reducing the exposure to increased oxygen levels. Studies show that, in the long term, a robust recovery from most types of oxygen toxicity is possible.

Protocols for avoidance of the effects of hyperoxia exist in fields where oxygen is breathed at higher-than-normal partial pressures, including underwater diving using compressed breathing gases, hyperbaric medicine, neonatal care and human spaceflight. These protocols have resulted in the increasing rarity of seizures due to oxygen toxicity, with pulmonary and ocular damage being mainly confined to the problems of managing premature infants.

In recent years, oxygen has become available for recreational use in oxygen bars. The US Food and Drug Administration has warned those suffering from problems such as heart or lung disease not to use oxygen bars. Scuba divers use breathing gases containing up to 100% oxygen, and should have specific training in using such gases.


Ozone , or trioxygen, is an inorganic molecule with the chemical formula O3. It is a pale blue gas with a distinctively pungent smell. It is an allotrope of oxygen that is much less stable than the diatomic allotrope O2, breaking down in the lower atmosphere to O2 (dioxygen). Ozone is formed from dioxygen by the action of ultraviolet light (UV) and electrical discharges within the Earth's atmosphere. It is present in very low concentrations throughout the latter, with its highest concentration high in the ozone layer of the stratosphere, which absorbs most of the Sun's ultraviolet (UV) radiation.

Ozone's odour is reminiscent of chlorine, and detectable by many people at concentrations of as little as 0.1 ppm in air. Ozone's O3 structure was determined in 1865. The molecule was later proven to have a bent structure and to be diamagnetic. In standard conditions, ozone is a pale blue gas that condenses at progressively cryogenic temperatures to a dark blue liquid and finally a violet-black solid. Ozone's instability with regard to more common dioxygen is such that both concentrated gas and liquid ozone may decompose explosively at elevated temperatures or fast warming to the boiling point.

It is therefore used commercially only in low concentrations.

Ozone is a powerful oxidant (far more so than dioxygen) and has many industrial and consumer applications related to oxidation. This same high oxidising potential, however, causes ozone to damage mucous and respiratory tissues in animals, and also tissues in plants, above concentrations of about 0.1 ppm. While this makes ozone a potent respiratory hazard and pollutant near ground level, a higher concentration in the ozone layer (from two to eight ppm) is beneficial, preventing damaging UV light from reaching the Earth's surface.


Photosynthesis is a process used by plants and other organisms to convert light energy into chemical energy that can later be released to fuel the organisms' activities. This chemical energy is stored in carbohydrate molecules, such as sugars, which are synthesized from carbon dioxide and water – hence the name photosynthesis, from the Greek φῶς, phōs, "light", and σύνθεσις, synthesis, "putting together". In most cases, oxygen is also released as a waste product. Most plants, most algae, and cyanobacteria perform photosynthesis; such organisms are called photoautotrophs. Photosynthesis is largely responsible for producing and maintaining the oxygen content of the Earth's atmosphere, and supplies all of the organic compounds and most of the energy necessary for life on Earth.Although photosynthesis is performed differently by different species, the process always begins when energy from light is absorbed by proteins called reaction centres that contain green chlorophyll pigments. In plants, these proteins are held inside organelles called chloroplasts, which are most abundant in leaf cells, while in bacteria they are embedded in the plasma membrane. In these light-dependent reactions, some energy is used to strip electrons from suitable substances, such as water, producing oxygen gas. The hydrogen freed by the splitting of water is used in the creation of two further compounds that serve as short-term stores of energy, enabling its transfer to drive other reactions: these compounds are reduced nicotinamide adenine dinucleotide phosphate (NADPH) and adenosine triphosphate (ATP), the "energy currency" of cells.

In plants, algae and cyanobacteria, long-term energy storage in the form of sugars is produced by a subsequent sequence of light-independent reactions called the Calvin cycle; some bacteria use different mechanisms, such as the reverse Krebs cycle, to achieve the same end. In the Calvin cycle, atmospheric carbon dioxide is incorporated into already existing organic carbon compounds, such as ribulose bisphosphate (RuBP). Using the ATP and NADPH produced by the light-dependent reactions, the resulting compounds are then reduced and removed to form further carbohydrates, such as glucose.

The first photosynthetic organisms probably evolved early in the evolutionary history of life and most likely used reducing agents such as hydrogen or hydrogen sulfide, rather than water, as sources of electrons. Cyanobacteria appeared later; the excess oxygen they produced contributed directly to the oxygenation of the Earth, which rendered the evolution of complex life possible. Today, the average rate of energy capture by photosynthesis globally is approximately 130 terawatts, which is about eight times the current power consumption of human civilization.

Photosynthetic organisms also convert around 100–115 billion tonnes (91-104 petagrams) of carbon into biomass per year.

Radical (chemistry)

In chemistry, a radical is an atom, molecule, or ion that has an unpaired valence electron.

With some exceptions, these unpaired electrons make radicals highly chemically reactive. Many radicals spontaneously dimerize. Most organic radicals have short lifetimes.

A notable example of a radical is the hydroxyl radical (HO•), a molecule that has one unpaired electron on the oxygen atom. Two other examples are triplet oxygen and triplet carbene (:CH2) which have two unpaired electrons.

Radicals may be generated in a number of ways, but typical methods involve redox reactions. Ionizing radiation, heat, electrical discharges, and electrolysis are known to produce radicals. Radicals are intermediates in many chemical reactions, more so than is apparent from the balanced equations.

Radicals are important in combustion, atmospheric chemistry, polymerization, plasma chemistry, biochemistry, and many other chemical processes. A large fraction of natural products is generated by radical-generating enzymes. In living organisms, the radicals superoxide and nitric oxide and their reaction products regulate many processes, such as control of vascular tone and thus blood pressure. They also play a key role in the intermediary metabolism of various biological compounds. Such radicals can even be messengers in a process dubbed redox signaling. A radical may be trapped within a solvent cage or be otherwise bound.


Redox (short for reduction–oxidation reaction) (pronunciation: redoks or reedoks) is a chemical reaction in which the oxidation states of atoms are changed. Any such reaction involves both a reduction process and a complementary oxidation process, two key concepts involved with electron transfer processes. Redox reactions include all chemical reactions in which atoms have their oxidation state changed; in general, redox reactions involve the transfer of electrons between chemical species. The chemical species from which the electron is stripped is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. It can be explained in simple terms:

Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion.

Reduction is the gain of electrons or a decrease in oxidation state by a molecule, atom, or ion.As an example, during the combustion of wood, oxygen from the air is reduced, gaining electrons from carbon which is oxidized. Although oxidation reactions are commonly associated with the formation of oxides from oxygen molecules, oxygen is not necessarily included in such reactions, as other chemical species can serve the same function.The reaction can occur relatively slowly, as with the formation of rust, or more quickly, in the case of fire. There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide (CO2) or the reduction of carbon by hydrogen to yield methane (CH4), and more complex processes such as the oxidation of glucose (C6H12O6) in the human body.

Respiratory system

The respiratory system (also respiratory apparatus, ventilatory system) is a biological system consisting of specific organs and structures used for gas exchange in animals and plants. The anatomy and physiology that make this happen varies greatly, depending on the size of the organism, the environment in which it lives and its evolutionary history. In land animals the respiratory surface is internalized as linings of the lungs. Gas exchange in the lungs occurs in millions of small air sacs called alveoli in mammals and reptiles, but atria in birds. These microscopic air sacs have a very rich blood supply, thus bringing the air into close contact with the blood. These air sacs communicate with the external environment via a system of airways, or hollow tubes, of which the largest is the trachea, which branches in the middle of the chest into the two main bronchi. These enter the lungs where they branch into progressively narrower secondary and tertiary bronchi that branch into numerous smaller tubes, the bronchioles. In birds the bronchioles are termed parabronchi. It is the bronchioles, or parabronchi that generally open into the microscopic alveoli in mammals and atria in birds. Air has to be pumped from the environment into the alveoli or atria by the process of breathing which involves the muscles of respiration.

In most fish, and a number of other aquatic animals (both vertebrates and invertebrates) the respiratory system consists of gills, which are either partially or completely external organs, bathed in the watery environment. This water flows over the gills by a variety of active or passive means. Gas exchange takes place in the gills which consist of thin or very flat filaments and lammelae which expose a very large surface area of highly vascularized tissue to the water.

Other animals, such as insects, have respiratory systems with very simple anatomical features, and in amphibians even the skin plays a vital role in gas exchange. Plants also have respiratory systems but the directionality of gas exchange can be opposite to that in animals. The respiratory system in plants includes anatomical features such as stomata, that are found in various parts of the plant.


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