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In science, oxygen (International Phonetic Alphabet: ) is a chemical element with the chemical symbol O and atomic number 8. The word oxygen derives from two roots in Ancient Greek, οξύς (oxys) (acid, lit. sharp) and -γενής (-genēs) (producer, lit. begetter). It was recognized in 1777 by Antoine Lavoisier, who coined the name oxygen from the Greek roots mentioned above because he erroneously thought that it was a constituent of all acids. (The acid#Definitions of acids and bases has since been revised). Oxygen has a valency (chemistry) of 2. On Earth it is usually bonded to other elements Covalent bond or ionic bond. Examples for common oxygen-containing compounds include water (H2O), sand (silica, SiO2), and rust (iron oxide, Fe2O3).

Diatomic oxygen (O2) is one of the two major components of air (20.95%). It is produced by plants during photosynthesis, and is necessary for aerobic respiration in animals. It is toxic to Anaerobic organism and was a poisonous waste product for early life on Earth.

Triatomic oxygen (ozone, O3) forms through radiation in the upper layers of the atmosphere and acts as a shield against UV radiation.

Characteristics

At standard temperature and pressure, oxygen exists as a diatomic molecule with the formula O2, in which the two oxygen atoms are chemical bond to each other with the electron configuration of triplet oxygen. This bond has a bond order of two, and is thus often grossly simplified in description as a double bond. Triplet oxygen is the ground state of the oxygen molecule. The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals. These orbitals are classified as antibonding, so the diatomic oxygen bond is weaker than the diatomic nitrogen bond, where all bonding molecular orbitals are filled. Though unpaired electrons are commonly associated with high reactivity in chemical compounds, triplet oxygen is relatively (and fortunately) nonreactive by comparison with most radicals.

Singlet oxygen, a name given to several higher energy species of molecular oxygen in which all the electron spins are paired, is much more reactive towards common organic compound molecules. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight. It is also produced by the immune system as a source of active oxygen. Carotenoids in photosynthetic organisms and possibly also in 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.

Liquid oxygen and solid O2 are clear substances with a light diffuse sky radiation. The phenomena are not related; the color of the sky is due to Rayleigh scattering. In normal triplet form they are paramagnetism due to the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O2 molecules. Liquid oxygen is attracted to a magnet to a sufficient extent that a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet, in laboratory demonstrations. Liquid O2 is usually obtained by the fractional distillation of liquid air or by the condensation out of air. It is a highly reactive substance and should be handled extremely carefully.

Oxygen is slightly soluble in water, but naturally occurring dissolved amounts are enough to support animal life (see below).

O2 has a bond length of 121 pm and a bond energy of 498 kJ/mol.

=== Allotropes ===The common allotrope of elemental oxygen on Earth, O2, is known as dioxygen. Elemental oxygen is most commonly encountered in this form, as 21% of Earth's atmosphere.

.Ozone (O3), the less common triatomic allotrope of oxygen, is a poisonous gas with a distinct, sharp odor. It is thermodynamically unstable toward the more common dioxygen form. It is formed continuously in the upper atmosphere of the Earth by short-wave UV radiation, and also functions as a shield against UV radiation reaching the ground. Ozone has recently been found to be produced by the immune system as an antimicrobial (see below). Liquid and solid O3 (ozone) have a deeper blue color than ordinary oxygen, and they are unstable and explosive. Traces of it can be detected sometimes as a sharp smell coming from electromotors.

A newly discovered allotropy of oxygen, tetraoxygen (O4), is a deep red solid that is created by pressurizing O2 to the order of 20 GPa. Its properties are being studied for use in rocket fuels and similar applications, as it is a much more powerful oxidizing agent than either O2 or O3.

History Early experiments and Phlogiston theory One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Ancient Greece 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.{{cite book] 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 (physiology).Cook (1969). The Encyclopedia of the Chemical Elements, "Oxygen", page 499

Oxygen's discovery as a separate element was delayed by a philosophy of combustion and corrosion called the phlogiston theory. Established in 1667 by Holy Roman Empire alchemist J. J. Becher and modified by chemist Georg Ernst Stahl by 1731,{{cite book]. Highly combustible materials, such as coal, were made mostly of phlogiston while non-combustible substances, such as iron, contained very little. Air did not play a role in phlogiston theory and no initial quantitative experiments were conducted to test the idea; instead it was based on observations of what happened when something burned.

In the late 16th century, Polish-Lithuanian Commonwealth alchemist and philosopher Michał Sędziwój thought of the gas given off by warm niter (saltpeter) as "the elixir of life".

Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all also produced oxygen in experiments in the 17th century but none of them recognized it as an element.Emsley (2001). Nature's Building Blocks, page 299

Discovery by Priestley and Scheele An experiment conducted by Kingdom of Great Britain clergyman Joseph Priestley on August 1 1774 focused sunlight on mercury(II) oxide (mercury (element)O) inside a glass tube, which liberated a gas he named 'dephlogisticated air'.Cook (1969). The Encyclopedia of the Chemical Elements, "Oxygen", page 500 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 work titled Experiments and Observations on Different Kinds of Air. Because he published first, Priestley is usually given priority in the discovery.

Unknown to Priestley, Sweden pharmacist Carl Wilhelm Scheele had already produced oxygen by heating mercuric oxide and various nitrates some time around 1772. Scheele wrote an account of this discovery in a manuscript he titled Treatise on Air and Fire, which he sent to his publisher in 1775. However, that document was not published until 1777.Emsley (2001). Nature's Building Blocks, page 300 Scheele called the gas 'fire air' because it was the only known supporter of combustion.

Noted France chemist Antoine Lavoisier later claimed to have independently discovered the new substance. However, Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele posted a letter to Lavoisier on September 30 1774 that described the 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).

Lavoisier's contribution What Lavoisier did indisputably do was to conduct the first adequate quantitative experiments on oxidation and give 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.

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 general, 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', which did not support either.

Lavoisier later renamed 'vital air' to oxygène after the Greek language roots meaning "acid-former" while 'azote' was renamed nitrogen. Oxygen entered the English language despite opposition by English scientists and the fact that Priestley had priority. 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.



Biological role and oceans in reaction to the evolution of oxygenic photosynthesis: A) no oxygen produced by biosphere, B) oxygen produced, but absorbed in oceans and by seabed rock, C) oxygen starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer.) -2) First land plants -3) Ordovician-Silurian extinction events -4) Huge forests form on land, first land animals and seed plants -5) Coal formation, first conifers, insect and amphibian giantism -6) Low ocean levels, supercontinent Pangaea forms -7) Permian-Triassic extinction event -8) First primitive flowering plants and dinosaurs -9) Triassic-Jurassic extinction event -10) Age of dinosaurs -11) Radiation of flowering plants -12) Cretaceous-Tertiary extinction event -13) Radiation of mammalsMolecular oxygen, O2, is essential for cellular respiration in all aerobic organisms. It is used as electron acceptor in the mitochondria to generate chemical energy in the form of adenosine triphosphate (ATP) during oxidative phosphorylation. During this reaction, oxygen is reduced to water. Conversely, free oxygen is produced in the biosphere through photolysis (light-driven oxidation and splitting) of water during photosynthesis in cyanobacteria, green algae and plants, thus closing the biological water-oxygen redox cycle.

Before the evolution of water oxidation in photosynthetic bacteria, oxygen was almost nonexistent in earth's atmosphere. Free oxygen first appeared in significant quantities during the Paleoproterozoic era (between 2.5 billion years ago and 1.6 billion years ago) as a product of the metabolic action of early anaerobes (archaea and bacteria). These organisms developed the mechanism of oxygen evolution between 3.5 and 2.7 billion years ago. At first, the produced oxygen dissolved in the oceans and reacted with iron. It started to "gas out" of the oxygen-saturated waters about 2.7 billion years ago as evident in the rusting of iron-rich terrestrial rocks starting around that time. The amount of oxygen in the atmosphere increased gradually at first and shot up rapidly around 2.2 to 1.7 billion years ago to about 10% of its present level.

The development of an oxygen-rich atmosphere was one of the most important events in the history of life on earth. The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the anaerobic organisms then living to extinction during the oxygen catastrophe about 2.4 billion years ago. However, the high electronegativity of O2 creates a large potential energy drop for cellular respiration, thus enabling organisms using aerobic respiration to produce much more ATP than anaerobic organisms. This makes them so efficient that they have come to dominate earth's biosphere. Photosynthesis and cellular respiration of oxygen allowed for the evolution of eukaryotic cells and ultimately complex multicellular organisms such as plants and animals.

The atmospheric abundance of free oxygen in later geological epochs and its gradual increase up to the present has been largely due to synthesis by photosynthetic organisms. Over the past 500 million years, oxygen levels fluctuated between 15 and 35% per volume. Towards the end of the Carboniferous era (coal age) about 300 million years ago, atmospheric oxygen levels reached a maximum of 35% by volume, allowing insects and amphibians with limiting respiratory systems to grow much larger than today's species. Today, oxygen is the second most common component of the earth's atmosphere (about 21% by volume) after nitrogen.

Occurrence . Note more oxygen in cold water near the poles. Data from the World Ocean Atlas 2001.

Oxygen is the third most abundant chemical element in the universe by mass, after hydrogen and helium (see chemical element). Some of this oxygen was produced during stellar nucleosynthesis as a step in the CNO-II branch of the CNO cycle. However oxygen is primarily produced in massive stars. In stars with at least four times the solar mass, 16O nuclei are produced during the Carbon burning process. 16O can also be produced in stars with at least 8 times the solar mass as a result of photodisintegration during the Neon burning process.

Oxygen is the most common component of the Earth's crust (49% by mass), Los Alamos National Laboratory – Oxygen the second most common component of the Earth as a whole (28% by mass), the most common component of the world's oceans (86% by mass), and the second most common component of the Earth's atmosphere (20.947% by volume), second to nitrogen.

Elemental oxygen occurs not only in the atmosphere, but also as solution in the world's water bodies. At 25° C under 1 atmosphere (unit) of air, a litre of water will dissolve about 6.04 cubic centimetre (8.63 milligram, 0.270 millimole) of oxygen, whereas sea water will dissolve about 4.9 cc (7.0 mg, 0.22 mmol). At 0° C the solubilities increase to 10.29 cc (14.7 mg, 0.460 mmol) for water and 8.0 cc (11.4 mg, 0.36 mmol) for sea water. This difference has important implications for ocean life, as polar oceans support a much higher density of life due to their oxygen content. From The Chemistry and Fertility of Sea Waters by H.W. Harvey, 1955, citing C.J.J. Fox, "On the coefficients of absorption of atmospheric gases in sea water", Publ. Circ. Cons. Explor. Mer, no. 41, 1907. Harvey however notes that according to later articles in Nature the values appear to be about 3% too high.

See also :Category:Silicate minerals, :Category:Oxide minerals.

Production In nature, free oxygen is produced by the light-driven photolysis during oxygenic photosynthesis in cyanobacteria, green algae and plants. Algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth.{{cite book ] liquefied air into its various components, with nitrogen distillation as a vapor while oxygen is left as a liquid. The other major method of producing oxygen involves passing a stream of clean, dry air through a bed of zeolite molecular sieves, which absorb the nitrogen and leave a gas stream that is 90 to 93% oxygen. Nitrogen is released from saturated zeolite by diverting air flow to another zeolite bed and reducing the chamber's air pressure. This allows for a continuous supply of gaseous oxygen to be pumped through a pipeline.

Oxygen can also be produced through electrolysis of water into oxygen and hydrogen. A similar method is the electrocatalytic oxygen 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 spacecraft and submarines. Oxygen is increasingly obtained by non-cryogenic technologies such as pressure swing adsorption (PSA), vacuum-pressure swing adsorption (VPSA) Non-Cryogenic Air Separation Processes 2003, or vacuum swing adsorption (VSA) technolgies. Air can be forced to dissolve through ceramic membranes based on zirconium oxide by either high pressure or an electric current to produce nearly pure oxygen.Emsley (2001). Nature's Building Blocks, page 301

In large quantities, the price of liquid oxygen (2001) is approximately $0.21/kg. NASAFacts FS-2001-09-015-KSC, Space Shuttle Use of Propellants and Fluids, National Aeronautics and Space Administration, September 2001 (postscript file here Since the primary cost of production is the energy cost of liquefying the air, the production cost will change as energy cost varies.

Oxygen is often transported in bulk as a liquid in specially insulated tankers because one liter of liquefied oxygen is equivalent to 840 liters of the gas. Oxygen is also stored and shipped in cylinders containing the compressed gas; a form that is useful in medical applications and Oxy-fuel welding and cutting.

Compounds Due to its electronegativity, oxygen forms chemical bonds with almost all other elements hence the original definition of oxidation. The only elements known to escape the possibility of oxidation are a few of the noble gases, and fluorine. However, many noble metals (common examples: gold, platinum) resist direct chemical combination with oxygen, and substances like gold oxide must be formed by an indirect route.

The most familiar oxygen compound is water. Other well-known examples include silica (found in sand, glass, Rock (geology), etc.), and the compounds of carbon and oxygen, such as carbon dioxide (CO2), alcohols (R-OH), carbonyls, (R-CO-H or R-CO-R), and carboxylic acids (R-COOH). Oxygenated radical (chemistry)s such as chlorates (ClO3−), perchlorates (ClO4−), chromates (CrO42−), dichromates (Cr2O72−), permanganates (MnO4−), and nitrates (NO3−) are strong oxidizing agents in and of themselves. Phosphorus is biologically important in its oxygenated form as the phosphate (PO43−) ion.Many metals bond with oxygen atoms, such as iron in iron(III) oxide (Fe2O3), commonly called rust.

Ozone (O3) is formed by electrostatic discharge in the presence of molecular oxygen. A double oxygen molecule (O2)2 is known and is found as a minor component of liquid oxygen. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.

One unexpected oxygen compound is Platinum hexafluoride#Reactions O2+PtF6−. It was discovered when Neil Bartlett was studying the properties of platinum hexafluoride. He noticed a change in color when this compound was exposed to atmospheric air. Bartlett reasoned that xenon should be oxidized by PtF6. This led him to the discovery of xenon hexafluoroplatinate Xe+PtF6−.

Isotopes Oxygen has seventeen known isotopes with atomic masses ranging from 12.03 u to 28.06 u. Three are stable, 16O, 17O, and 18O, of which 16O is the most abundant (over 99.7%). The radioisotopes all have half-lives of less than three minutes. Nonetheless, 15O is used in positron emission tomography.

An atomic weight of 16 was assigned to oxygen prior to the definition of the atomic mass unit based upon 12C. Since physicists referred to 16O only, while chemists meant the naturally abundant mixture of isotopes, this led to slightly different atomic weight scales.

Precautions Toxicity of O2 Oxygen can be Oxygen toxicity at elevated partial pressures. Since oxygen partial pressure is the fraction of oxygen times the total pressure, elevated partial pressures can occur either from high oxygen fraction in breathing gas, or from high breathing gas pressure, or a combination of both. Oxygen toxicity usually begins to occur at partial pressures more than 0.5 atmospheres, or 2.5 times the normal sea-level oxygen partial pressure of about 0.2 atmospheres or bars. This means that at sea-level pressures, mixtures containing less than 50% oxygen are essentially non-toxic. However in medical applications (such as in ventilation gas mixtures in hospital applications) mixtures containing more than 50% oxygen can be expected to show lung toxicity, causing slow damage to the lungs over periods of days, with the rate of damage rising rapidly from mixtures between 50% and 100% oxygen. On the other hand, breathing 100% oxygen in space applications (such as in some modern spacesuits, or in early spacecraft such as the Apollo spacecraft), causes no damage due to the low total pressures (30% to 33% sea-level) used. In the case of spacesuits, oxygen partial pressure in the breathing gas is typically about 0.30 bar (1.4 times normal), and oxygen partial pressure in the astronaut's blood (due to downward adjustments due to water vapor and CO2 in the alveoli) is close to sea-level normal of 0.2 bar.

In deep scuba diving and surface supplied diving and when using equipment which can provide high partial pressures of oxygen, such as rebreathers, oxygen toxicity to the lungs can occur, just as in medical applications. Due to the higher total pressures in these applications, the fraction of oxygen which produces lung damage may be considerably less than 50%. More importantly, under pressures higher than normal sea-level, a far more serious form of oxygen toxicity in the central nervous system may lead to generalized seizures. This form of oxygen toxicity usually occurs after several hours exposure to oxygen partial pressures over about 1.4 atmospheres (bars) (i.e. 7 times normal), with the time decreasing for higher pressures above this, and with great variation from person to person. At over three bars of oxygen partial pressure (15 times normal), seizures typically occur within minutes.

Toxicity and antibacterial use of other chemical oxygen forms Certain derivatives of oxygen, such as ozone (O3), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic. Cells have developed various mechanisms to protect against all of these toxic compounds. For instance, the naturally-occurring glutathione can act as an antioxidant, as can bilirubin which is normally a breakdown product of hemoglobin. To protect against the destructive nature of peroxides, nearly every organism on earth has developed some form of the enzyme catalase, which quickly disproportionation hydrogen peroxide into water and dioxygen. Another nearly universally present enzyme in living organisms (except for a few species of bacteria which use Mn2+ ions directly for the job) is superoxide dismutase. This family of enzymes disproportionation superoxide to oxygen and peroxide, which is then in turn dealt with, by catalase.

Immune systems of higher organisms have long made use of reactive forms of oxygen which they produce. Not only do antibodies catalyze production of peroxide from oxygen, it is now known that immune cells produce peroxide, superoxide, and singlet oxygen in the course of an immune response. Recently, singlet oxygen has been found to be a source of biologically-produced ozone: this reaction proceeds through an unusual compound dihydrogen trioxide, also known as trioxidane, (HOOOH) which is an antibody-catalyzed product of singlet oxygen and water. This compound in turn disproportionates to ozone and peroxide, providing two powerful antibacterials. The body's range of defense against all of these active oxidizing agents is hardly surprising, then, given their "deliberate" employment as antimicrobial agents in the immune response.

Oxygen derivatives are prone to form free radicals, especially in metabolic processes. Because they can cause severe damage to cells and their DNA before they are dealt with, they form part of many theories of carcinogenesis and aging.

Combustion hazard Highly concentrated sources of oxygen promote rapid combustion and therefore are fire and explosion hazards in the presence of fuels. Oxygen itself is not the fuel, but as a reactant, concentrated oxygen may allow combustion to proceed dangerously rapidly. The fire that killed the Apollo 1 crew on a test launchpad spread so rapidly because the capsule was pressurized with pure oxygen as would be usual in an actual flight, but to maintain positive pressure in the capsule, this was at slightly more than atmospheric pressure instead of the ⅓ normal pressure that would be used in flight. (See partial pressure.)

Hazards also apply to compounds of oxygen with a high oxidative potential, such as high concentration peroxides, chlorates, perchlorates, and dichromates; they also can often cause Chemical burn.

See also

References

External links



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