Oxygen

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In science, oxygen (IPA: /ˈɒkˑsəˑdʒɪn/) is a chemical element with the chemical symbol O and atomic number 8. The word oxygen derives from two roots in 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 definition of acid has since been revised). Oxygen has a valency of 2. On Earth it is usually bonded to other elements covalently or ionically. 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 obligate anaerobic organisms 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.

[edit] Characteristics

The common allotrope of elemental oxygen on Earth, O₂, is known as dioxygen.

At standard temperature and pressure, oxygen exists as a diatomic molecule with the formula O₂, in which the two oxygen atoms are bonded to each other with the electron configuration of triplet oxygen. This bond has a bond order of two, and is thus often very grossly simplified in description as a double bond.^[1] 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 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 O₂ and solid O₂ are clear substances with a light sea-blue color. In normal triplet form they are paramagnetic 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.

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.^[2]

[edit] Allotropes

Ozone, 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.

A newly discovered allotrope 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 oxidizer than either O₂ or O₃.^[3][4]

[edit] Applications

Uptake of oxygen from the air is the essential purpose of respiration, so oxygen supplementation has found use in medicine (as oxygen therapy). People who climb mountains or fly in non-pressurized aeroplanes sometimes have supplemental oxygen supplies; the reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspired oxygen partial pressure nearer to that found at sea-level.

A notable application of oxygen as a very low-pressure breathing gas, is in modern spacesuits, where use of nearly pure oxygen at a total ambient pressure of about one third normal, results in normal blood partial pressures of oxygen. This trade-off of breathing gas content and needed pressure is important for space applications, because the issue of flexible spacesuits working at Earth sea-level pressures remains a technological challenge of aerospace technology. Oxygen is used in welding (such as the oxyacetylene torch), and in the industrial production of steel and methanol. Also, liquid oxygen finds use as a classic oxidizer in rocket propulsion.

Oxygen presents two spectrophotometric absorption bands peaking at the wavelengths 687 and 760 nanometers. Some scientists have proposed to use the measurement of the radiance coming from vegetation canopies in those oxygen bands to characterize plant health status from a satellite platform.^{[citation needed]} This is because in those bands, it is possible to discriminate the vegetation's reflectance from the vegetation's fluorescence, which is much weaker. The measurement presents several technical difficulties due to the low signal to noise ratio and due to the vegetation's architecture, but it has been proposed as a possibility to monitor the carbon cycle from satellites on a global scale.

Oxygen, as a supposed mild euphoric, has a history of recreational use (see oxygen bar). However, the reality of a pharmacological effect is doubtful, a metabolic boost being the most plausible explanation. Controlled tests of high oxygen mixtures in diving (see nitrox) and other activities, even at higher than normal pressures, demonstrated no particular effects on humans other than promotion of an increased tolerance to aerobic exercise.

In the 19th century, oxygen was often mixed with nitrous oxide to temper its analgesic effect. A stable 50% gaseous mixture (Entonox) is commonly used in medicine today as an analgesic. However, the common basic anaesthetic mixture is 30% oxygen with 70% nitrous oxide; the pain-suppressing effects, obviously, are due to the nitrous oxide and not to oxygen.

[edit] History

Oxygen was first described by Michał Sędziwój, a Polish alchemist and philosopher in the late 16th century. Sędziwój thought of the gas given off by warm niter (saltpeter) as "the elixir of life".^[5] Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all produced oxygen in experiments in the 17th century but none of them recognized it as an element.^[6]

An experiment conducted by British clergyman Joseph Priestley on August 1, 1774 focused sunlight on mercuric oxide (HgO) inside a glass tube, which liberated a gas he named dephlogisticated air (see phlogiston theory). 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 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 and consequently is usually given priority in the discovery.

Unknown to Priestley, Swedish 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 sent to his publisher in 1775. However, that document wasn't published until 1777.^[7] Scheele called the gas 'fire air' because it was the only known supporter of combustion. It was later called 'vital air' because it was and is vital for the existence of animal life.

Noted French chemist Antoine Laurent Lavoisier later claimed to have independently discovered the new substance. However, Priestley visited Lavoisier in October 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).^[7]

Lavoisier named the new element oxygène from Greek roots meaning "acid-former". As noted, the name reflects the then-common incorrect belief that all acids contain oxygen. 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.^[7]

Oxygen was condensed for the first time in 1883 by Zygmunt Wróblewski and Karol Olszewski, both professors of Jagiellonian University.

[edit] Biological role

Molecular 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.^[8]

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.^[9] 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.

[edit] Occurrence

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 Sun's mass, ¹⁶O nuclei are produced during the Carbon burning process. ¹⁶O can also be produced in stars with at least 8 times the Sun's mass as a result of photodisintegration during the Neon burning process.^[11]

Oxygen is the most common component of the Earth's crust (49% by mass),^[12] 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 atm of air, a litre of water will dissolve about 6.04 cc (8.63 mg, 0.270 mmol) 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. ^[13]

See also Silicate minerals, Oxide minerals.

[edit] Production

Main article: Oxygen evolution

In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis in cyanobacteria, green algae and plants.^[14] Algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth.^[15] The remainder is produced by terrestrial plants, although almost all oxygen produced in tropical forests is consumed by organisms in those forests.^[16]

In the laboratory and industrially, oxygen can 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.

Industrially, oxygen is typically produced in bulk quantity as a liquid produced by distillation from atmospheric air. In large quantities, the price of liquid oxygen (2001) is approximately $0.21/kg ^[17]. Since the primary cost of production is the energy cost of liquifying the air, the production cost will change as energy cost varies.

In the modern era, oxygen is increasingly obtained by non-cryogenic technolgies such as pressure swing adsorption (PSA) and vacuum-pressure swing adsorption (VPSA) technolgies.^[18]

[edit] 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, 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 radicals 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 dioxygen hexafluoroplatinate O₂⁺PtF₆⁻. It was discovered when Neil Bartlett was studying the properties of PtF₆. 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⁻.

See also: Category:Oxygen compounds

[edit] Isotopes

Main article: isotopes of oxygen

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 unified 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.

[edit] Precautions

[edit] Toxicity of O₂

Main article: oxygen toxicity

Oxygen can be toxic 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.^[19] 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.

[edit] 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 very quickly disproportionates 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 disproportionates 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.^[20]

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.

[edit] 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 burns.