Oxygen ( /ˈɒksɨdʒɨn/ ok-si-jin) is the element with atomic number 8 and represented by the symbol O. Its name derives from the Greek roots ὀξύς (oxys) ("acid", literally "sharp", referring to the sour taste of acids) and -γενής (-genēs) ("producer", literally "begetter"), because at the time of naming, it was mistakenly thought that all acids required oxygen in their composition. At standard temperature and pressure, two atoms of the element bind to form dioxygen, a very pale blue, odorless, tasteless diatomic gas with the formula O2.

Oxygen is a member of the chalcogen group on the periodic table and is a highly reactive nonmetallic element that readily forms compounds (notably oxides) with almost all other elements. Oxygen is a strong oxidizing agent and has the second highest electronegativity of all the elements (only fluorine has a higher electronegativity). By mass, oxygen is the third most abundant element in the universe after hydrogen and helium and the most abundant element by mass in the Earth's crust, making up almost half of the crust's mass. Free oxygen is too chemically reactive to appear on Earth without the photosynthetic action of living organisms, which use the energy of sunlight to produce elemental oxygen from water. Elemental O2 only began to accumulate in the atmosphere after the evolutionary appearance of these organisms, roughly 2.5 billion years ago. Diatomic oxygen gas constitutes 20.8% of the volume of air.

Because it comprises most of the mass in water, oxygen comprises most of the mass of living organisms (for example, about two-thirds of the human body's mass). All major classes of structural molecules in living organisms, such as proteins, carbohydrates, and fats, contain oxygen, as do the major inorganic compounds that comprise animal shells, teeth, and bone. Elemental oxygen is produced by cyanobacteria, algae and plants, and is used in cellular respiration for all complex life. Oxygen is toxic to obligately anaerobic organisms, which were the dominant form of early life on Earth until O2 began to accumulate in the atmosphere. Another form (allotrope) of oxygen, ozone (O3), helps protect the biosphere from ultraviolet radiation with the high-altitude ozone layer, but is a pollutant near the surface where it is a by-product of smog. At even higher low earth orbit altitudes atomic oxygen is a significant presence and a cause of erosion for spacecraft.

Oxygen was independently discovered by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph Priestley in Wiltshire, in 1774, but Priestley is often given priority because his work was published first. The name oxygen was coined in 1777 by Antoine Lavoisier, whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion.

Oxygen is produced industrially by fractional distillation of liquefied air, use of zeolites with pressure-cycling to concentrate oxygen from air, electrolysis of water and other means. Uses of oxygen include the production of steel, plastics and textiles; rocket propellant; oxygen therapy; and life support in aircraft, submarines, spaceflight and diving.

Contents 1 Safety and precautions 1.1 Toxicity 1.2 Combustion and other hazards 2 Notes and citations 3 References 4 Further reading

Safety and precautionsEdit

ToxicityEdit

Main article: Oxygen toxicity

Oxygen gas (O₂) can be toxic at elevated partial pressures, leading to convulsions and other health problems.^[1][2][3] Oxygen toxicity usually begins to occur at partial pressures more than 50 kilopascals (kPa), or 2.5 times the normal sea-level O₂ partial pressure of about 21 kPa (equal to about 50% oxygen composition at standard pressure). 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% O₂ by volume (about 30 kPa at standard pressure).^[4] (although this figure also is subject to wide variation, depending on type of mask).

At one time, premature babies were placed in incubators containing O₂-rich air, but this practice was discontinued after some babies were blinded by it.^[4][5]

Breathing pure O₂ 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.^[6][7] In the case of spacesuits, the O₂ partial pressure in the breathing gas is, in general, about 30 kPa (1.4 times normal), and the resulting O₂ partial pressure in the astronaut's arterial blood is only marginally more than normal sea-level O₂ partial pressure (for more information on this, see space suit and arterial blood gas).

Oxygen toxicity to the lungs and central nervous system can also occur in deep scuba diving and surface supplied diving.^[4][1] Prolonged breathing of an air mixture with an O₂ partial pressure more than 60 kPa can eventually lead to permanent pulmonary fibrosis.^[8] Exposure to a O₂ 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% O₂ at 66 m or more of depth; the same thing can occur by breathing 100% O₂ at only 6 m.^{[8][9][10][11]}

Combustion and other hazardsEdit

Highly concentrated sources of oxygen promote rapid combustion. Fire and explosion hazards exist when concentrated oxidants and fuels are brought into close proximity; however, an ignition event, such as heat or a spark, is needed to trigger combustion.^[12] Oxygen itself is not the fuel, but the oxidant. Combustion hazards also apply to compounds of oxygen with a high oxidative potential, such as peroxides, chlorates, nitrates, perchlorates, and dichromates because they can donate oxygen to a fire.

Concentrated O₂ will allow combustion to proceed rapidly and energetically.^[12] 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 O₂ systems requires special training to ensure that ignition sources are minimized.^[12] The fire that killed the Apollo 1 crew in a launch pad test spread so rapidly because the capsule was pressurized with pure O₂ but at slightly more than atmospheric pressure, instead of the ¹⁄₃ normal pressure that would be used in a mission.^[13][14]

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.^[12] As with other cryogenic liquids, on contact with the human body it can cause frostbites to the skin and the eyes.

Notes and citationsEdit

ReferencesEdit

• 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, UK: Oxford University Press. pp. 297–304. ISBN 0-19-850340-7. 
• Raven, Peter H.; Ray F. Evert, Susan E. Eichhorn (2005). Biology of Plants, 7th Edition. New York: W.H. Freeman and Company Publishers. pp. 115–27. ISBN 0-7167-1007-2. 

Further readingEdit

• Walker, J. (1980). "The oxygen cycle". In Hutzinger O.. Handbook of Environmental Chemistry. Volume 1. Part A: The natural environment and the biogeochemical cycles. Berlin; Heidelberg; New York: Springer-Verlag. p. 258. ISBN 0-387-09688-4.