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8 nitrogen ← oxygen → fluorine - ↑ O ↓ S Periodic table - Extended periodic table
General
Name, symbol, number	oxygen, O, 8
Chemical series	nonmetals, chalcogens
Group, period, block	16, 2, p
Appearance	colorless (gas) pale blue (liquid)
Standard atomic weight	15.9994(3)?g·mol?1
Electron configuration	1s2 2s2 2p4
Electrons per shell	2, 6
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 1.429 g/L
Melting point	54.36 K (-218.79 °C, -361.82 °F)
Boiling point	90.20 K (-182.95 °C, -297.31 °F)
Critical point	154.59 K, 5.043 MPa
Heat of fusion	(O2) 0.444 kJ·mol?1
Heat of vaporization	(O2) 6.82 kJ·mol?1
Heat capacity	(25 °C) (O2) 29.378 J·mol?1·K?1
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K       61 73 90
Atomic properties
Crystal structure	cubic
Oxidation states	?2, ?1 (neutral oxide)
Electronegativity	3.44 (Pauling scale)
Ionization energies (more)	1st: 1313.9 kJ·mol?1
	2nd: 3388.3 kJ·mol?1
	3rd: 5300.5 kJ·mol?1
Atomic radius	60 pm
Atomic radius (calc.)	48 pm
Covalent radius	73 pm
Van der Waals radius	152 pm
Miscellaneous
Magnetic ordering	paramagnetic
Thermal conductivity	(300 K) 26.58 m W·m?1·K?1
Speed of sound	(gas, 27 °C) 330 m/s
CAS registry number	7782-44-7
Selected isotopes
Main article: Isotopes of oxygen iso NA half-life DM DE (MeV) DP 16O 99.76% O is stable with 8 neutrons 17O 0.038% O is stable with 9 neutrons 18O 0.21% O is stable with 10 neutrons
References
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In science, oxygen (IPA: /??k?s??d??n/) is a chemical element with the chemical symbol O and atomic number 8. The word oxygen derives from two roots in Greek, οξ?? (oxys) (acid, lit. sharp) and -γεν?? (-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

At standard temperature and pressure, oxygen exists as a colorless, odorless 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 grossly simplified in description as a double bond.^[2]

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 sky blue coloration. The phenomena are not related; the color of the sky is due to Rayleigh scattering. 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.^[3] 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; 20 cc of the gas can dissolve in 1 l of water.^[4] O₂ has a bond length of 121 pm and a bond energy of 498 kJ/mol.^[5]

[edit] Allotropes

The common allotrope of elemental oxygen on Earth, O₂, is known as dioxygen. Elemental oxygen is most commonly encountered in this form, as about 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 ultraviolate (UV) radiation, and also functions as a shield against UV radiation reaching the ground. Ozone has also been found to be produced by the immune system as an antimicrobial (see below). Liquid and solid O₃ have a deeper blue color than ordinary oxygen and they are unstable and explosive. Traces of it can be detected as a sharp smell coming from electromotors.

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.^{[6] [7]} When tetraoxygen is subjected to a pressure of 96 GPa, it becomes metallic, similarly to hydrogen, and becomes more similar to the heavier chalcogens, such as tellurium and polonium, both of which show significant metallic character.

[edit] Applications

[edit] Breathing gas and supplement

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 home oxygen concentrator in situ in an emphysema patient's house. The model shown is the DeVILBISS LT 4000.

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, 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] Industrial

Smelting of iron ore into steel consumes 55% of commercially produced oxygen.^[8] In this process, oxygen is injected through a high-pressure lance into molten iron, which removes sulfur and carbon impurities. The reaction is exothermic, so the temperature increases to 1700 ° C. ^[8]

Another 25% of commercially produced oxygen is used by the chemical industry.^[8] Ethylene is reacted with oxygen 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). ^[8]

Most of the remaining 20% of commercially produced oxygen is used in medical applications, metal cutting, as an oxidizer in rocket fuel, and in water treatment.^[8] Oxygen is used in welding (such as the oxyacetylene torch).

[edit] Scientific

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.^[9] 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.

Paleoclimatologists measure the ratio of oxygen-18 and oxygen-16 in the shells and skeletons of marine organisms to determine what the climate was like millions of years ago. During periods of lower global temperatures, sea water molecules that contain the lighter isotope, oxygen-16, evaporate at a slightly faster rate than water molecules containing the heavier oxygen-18.^[10] Snow and rain from that evaporated water tends to be enriched in oxygen-16 and the seawater left behind tends to be enriched in oxygen-18. Marine organisms then incorporate more oxygen-18 into their skeletons and shells than they would in a warmer climate.^[10] Paleoclimatologists also directly measure this ratio in air trapped in ice core samples that are up to several hundreds of thousand years old.

[edit] History

[edit] 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 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.^[11] 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. ^[12]

Oxygen's discovery as a separate element was delayed by a philosophy of combustion and corrosion called the phlogiston theory. Established in 1667 by German alchemist J. J. Becher and modified by chemist Georg Ernst Stahl by 1731,^[13] 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, its calx. ^[12] 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.^[12]

In the late 16th century, Polish alchemist and philosopher Micha? S?dziwój 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.^[4]

[edit] Discovery by Priestley and Scheele

Joseph Priestley is usually given priority in the discovery

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'.^[14] 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."^[4] Priestley published his findings in 1775 in a work titled Experiments and Observations on Different Kinds of Air.^[12] Because he published first, Priestley 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.^{[12] [14]} 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.^[15] Scheele called the gas 'fire air' because it was the only known supporter of combustion.

Noted French chemist Antoine Laurent 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).^[15]

[edit] Lavoisier's contribution

Antoine Lavoisier discredited the Phlogiston theory

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.^[14] 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.^[14] 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.^[14] 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 roots meaning "acid-former" while 'azote' was renamed nitrogen. ^[14] 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.^[15]

[edit] Later history

Oxygen was liquefied for the first time in 1877 by Frenchman Louis Cailletet.^[3] Cailletet was only able to produce a few drops of liquid oxygen so no meaningful analysis of the substance could be conducted. In 1891 James Dewar was able to produce enough liquid oxygen to study.^[3]

[edit] Biological role

[edit] Respiration and use in biomolecules

Parts of DNA are made of oxygen and the element is found in almost all molecules that are important to life. Molecular oxygen, O₂, is essential for cellular respiration in all aerobic organisms. Almost all animals use hemoglobin in their blood to transport oxygen from their lungs to their tissues, but some, such as spiders and lobsters, use hemerythrin. ^[16] Hemoglobin uses iron at its active sites to attract oxygen while hemerythrin uses copper. A liter of blood can dissolve 200 cc of oxygen gas, which is much more than water can dissolve.^[16]

Oxygen is handed-off from hemoglobin to monooxygenase, an enzyme that also has an active site with an atom of iron.^[16] Monooxygenase uses oxygen to catalyze many oxidation reactions in the body. 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. On average, a oxygen atom is used in respiration once every 3,000 years.^[3]

Reactive oxygen species are dangerous by-products that sometimes result from the use of oxygen in organisms. Important examples include; oxygen free radicals such as the highly dangerous superoxide O₂^-, and the less harmful hydrogen peroxide ( H₂O₂). ^[16] The body uses superoxide dismutase to reduces superoxide radicals to hydrogen peroxide. Glutathione peroxidase and similar enzymes, then convent the H₂O₂ to water and dioxygen.^[16]

Parts of the immune system of higher organims, however, create peroxide, superoxide and singlet oxygen to destroy invading microbes. 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.^[17]

[edit] Life brings oxygen to the atmosphere

Oxygen was almost nonexistent in earth's atmosphere before the evolution of water oxidation in photosynthetic bacteria. 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, creating banded iron formations. 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.^[18]

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.^[19] 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. Human activities, including the burning of 7 billion tonnes of fossil fuels each year have had very little effect on the amount of free oxygen in the atmosphere.^[3]

[edit] Occurrence

Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.^[21] 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. ¹⁶O nuclei are produced during the carbon burning process in stars with at least four times the Sun's mass. 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.^[22]

Oxygen and its compounds constitute 49.2% of the Earth's crust by mass,^[23] 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. The Earth is unusual in having such a high concentration of free oxygen in its atmosphere. With only 0.15% oxygen by volume, the atmosphere of Mars has the second most abundant concentration of oxygen of any planet in the solar system (Venus comes in third place).^[21] However, their oxygen is only produced by ultraviolate radiation impacting oxygen-containing molecules such as carbon dioxide.

Elemental oxygen also occurs in 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.^[24]

The unusually high concentration of elemental oxygen 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 the modern Earth's atmosphere and life as we know it. Because of the vast amounts of oxygen in the atmosphere, even if all photosynthesis were to cease it would take between 5,000^[25] to 2.5 million years to strip out more or less all oxygen.

See also Silicate minerals, Oxide minerals.

[edit] Production

Hoffman electrolysis apparatus used in electrolysis of water

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. ^[26] Algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth.^[27] The remainder is produced by terrestrial plants, although almost all oxygen produced in tropical forests is consumed by organisms in those forests.^[28]

Two major methods are employed to produce the 100 million tonnes of oxygen extracted from air for industrial uses annually.^[15] The most common method is to fractionally distill liquefied air into its various components, with nitrogen distilling as a vapor while oxygen is left as a liquid.^[15] 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.^[15] 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) ^[29], or vacuum swing adsorption (VSA) [1] 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.^[8]

In large quantities, the price of liquid oxygen (2001) is approximately $0.21/kg. ^[30] 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.^[15] 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.^[15]

[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 (CrO₄^2?), dichromates (Cr₂O₇^2?), permanganates (MnO₄^?), and nitrates (NO₃^?) are strong oxidizing agents in and of themselves. Phosphorus is biologically important in its oxygenated form as the phosphate (PO₄^3?) 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.

A billion degrees are required for two oxygen nuclei to undergo nuclear fusion to form heavier the nuclei of silicon, phosphorus and sulfur. ^[21]

[edit] Precautions

[edit] Toxicity

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. Therefore, air supplied through oxygen masks in medical applications is typically composed of 30% oxygen by volume.^[4] (At one time, premature babies were placed in incubators containing oxygen-rich air, but this practice was discontinued after some babies were blinded by it.)^[4]

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.^[31] 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 or convulsions. ^[4] 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.

Certain derivatives of oxygen, such as ozone (O₃), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic.

[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 ? 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.

[edit] See also

[edit] References

• Nist atomic spectra database
• Nuclides and Isotopes Fourteenth Edition: Chart of the Nuclides, General Electric, 1989
• A. Szydlo, "Water which does not wet hands - The Alchemy of Michael Sendivogius", Warsaw Institute for the History of Sciences, Polish Academy of Sciences, 1994.

[edit] External links

• Priestley Society, Dedicated to Joseph Priestley the man who discovered oxygen
• Los Alamos National Laboratory – Oxygen
• WebElements.com – Oxygen
• Molecular Oxygen Site
• It's Elemental – Oxygen
• Oxygen (O2) Properties, Uses, Applications
• Computational Chemistry Wiki
• Oxidizing Agents > Oxygen
• Recent Roald Hoffman article on "The Story of O"