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Oxygen



8 nitrogenoxygenfluorine
-

O

S
General
Name, symbol, number oxygen, O, 8
Chemical seriesnonmetals, chalcogens
Group, period, block 162, p
Appearance

colorless gas above
light blue liquid
Standard atomic weight 15.9994(3) g·mol−1
Electron configuration 1s2 2s2 2p4
Electrons per shell 2, 6
Physical properties
Phasegas
Density(0 °C, 101.325 kPa)
1.429 g/L
Melting point54.36 K
(-218.79 °C, -361.82 °F)
Boiling point90.20 K
(-182.95 °C, -297.31 °F)
Critical point154.59 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 structurecubic
Oxidation states2, 1, −1, −2
(neutral oxide)
Electronegativity3.44 (Pauling scale)
Ionization energies
(more)
1st: 1313.9 kJ·mol−1
2nd: 3388.3 kJ·mol−1
3rd: 5300.5 kJ·mol−1
Atomic radius60 pm
Atomic radius (calc.)48 pm
Covalent radius73 pm
Van der Waals radius152 pm
Miscellaneous
Magnetic orderingparamagnetic
Thermal conductivity(300 K) 26.58x10-3  W·m−1·K−1
Speed of sound(gas, 27 °C) 330 m/s
CAS registry number7782-44-7
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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Oxygen (pronounced /ˈɒksɪdʒən/) is a colorless, odorless, tasteless, gaseous chemical element with the chemical symbol O and atomic number 8. It is a chalcogen, period 2, nonmetallic element that can form binary compounds (known as oxides) with almost all the other elements. The valency of oxygen is 2 and the most common oxidation state is -2.[1] On Earth, oxygen is usually bonded to other elements covalently or ionically. Oxygen is the third-most-abundant element in the universe by mass (most abundant after hydrogen and helium),[2] the most abundant element by mass in the Earth's crust,[3] and the most abundant element by mass in the human body.[4]

The name oxygen was coined in 1777 by Antoine Lavoisier from the Greek roots οξύς (oxys) (acid, lit. "sharp," from the taste of acids) and -γενής (-genēs) (producer, lit. begetter), because he mistook oxygen to be a constituent of all acids.[5]

Free diatomic oxygen or dioxygen (O2) is, together with nitrogen, one of the two major components of air, constituting about a fifth of the volume of air.[6]

Oxygen is highly reactive, and readily forms compounds with most other elements. Its compounds with silicon and metals are abundant in the earth's crustal rocks and with hydrogen in water (H2O). The nucleic acids and all of the major classes of structural molecules in living organisms, proteins, polysaccharides, and fats contain oxygen, as do the major inorganic compounds that comprise animal shell (calcium carbonate), and tooth and bone (calcium phosphate). Dioxygen is produced from water by cyanobacteria, algae, and plants during photosynthesis, and the energy required to sustain life in aerobically-respiring organisms is provided by enzyme-mediated oxidation of sugars and carboxylic acids, which are themselves already oxygen-containing compounds. Without oxygen, most organisms with aerobic respiration die within minutes.[7] However, free oxygen is toxic to obligate anaerobic organisms and was a poisonous waste product for early life on Earth.

Additional recommended knowledge

Contents

Characteristics

Structure

Main articles: Triplet oxygen and Singlet oxygen

At standard temperature and pressure, oxygen is a colorless, odorless gas with the molecular formula O2, in which the two oxygen atoms are chemically bonded to each other with a triplet electron configuration. This bond has a bond order of two, and is often simplified in description as a double bond.[8]  

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 (weakening the bond order from three to two), so the diatomic oxygen bond is weaker than the diatomic nitrogen triple bond in which all bonding molecular orbitals are filled, but fewer antibonding ones are. Though unpaired electrons are commonly associated with high reactivity in chemical compounds, triplet oxygen is relatively nonreactive by comparison with most radicals.

In normal triplet form, oxygen molecules are paramagnetic due to the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O2 molecules.[9] Liquid oxygen is attracted to a magnet to a sufficient extent that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.[10] Oxygen's paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine the purity of gaseous oxygen.[11]

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.[12] It is also produced in the troposphere by the photolysis of ozone by light of short wavelength,[13] and by the immune system as a source of active oxygen.[14] 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.[15]

Physical properties

Oxygen is more soluble in water than nitrogen, water containing approximately 1 part of oxygen to 2 of nitrogen, compared with a ratio in the atmosphere 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).[16][17] At 25 °C at 1 atm of air, freshwater contains about 6.04 ml (8.63 mg, 0.27 mmol) of oxygen per liter, whereas sea water contains about 4.9 ml (7.0 mg, 0.22 mmol) per liter. At 0 °C the solubilities increase to 10.3 ml (14.7 mg, 0.46 mmol) per liter for water and 8.0 ml (11.4 mg, 0.36 mmol) per liter for sea water.

Oxygen condenses at 90.20 K (-182.95 °C, -297.31 °F), and freezes at 54.36 K (-218.79 °C, -361.82 °F). Both liquid and solid O2 are clear substances with a light sky-blue color caused by absorption in the red (by contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light). High-purity liquid O2 is usually obtained by the fractional distillation of liquified air;[18] Liquid oxygen may also be produced by condensation out of air, using liquid nitrogen as a coolant. It is a highly-reactive substance and must be segregated from combustible materials.[19]

Allotropes

Main article: Allotropes of oxygen

 

 

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% (by volume) of Earth's atmosphere. O2 has a bond length of 121 pm and a bond energy of 498 kJ/mol.[20]

Triatomic oxygen (Ozone, O3), the less common triatomic allotrope of oxygen, is a poisonous gas with a distinctive odor. Traces of it can be detected as a sharp, chlorine-like smell coming from electric motors, laser printers, and photocopiers. It was named "ozone" by Christian Friedrich Schönbein, in 1840, from the Greek word ÖĮώ (ozo) for smell. [5] Ozone is thermodynamically unstable toward the more common dioxygen form, and is formed by reaction of O2 with atomic oxygen produced by splitting of O2 by UV radiation in the upper atmosphere.[5] Ozone absorbs strongly in the ultraviolet and functions as a shield for the biosphere against the mutagenic and other damaging effects of solar UV radiation (see ozone layer).[5] Ozone is also formed by electrostatic discharge in the presence of dioxygen. The immune system produces ozone as an antimicrobial (see below).[14] Liquid and solid O3 have a deeper-blue color than ordinary oxygen and they are unstable and explosive.[21][5]

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.[22][23] When tetraoxygen is subjected to a pressure of 96 GPa, it becomes metallic, in a similar manner as hydrogen,[24] and becomes more similar to the heavier chalcogens, such as tellurium and polonium, both of which show significant metallic character.

Compounds

 

Main article: Oxygen compounds

Oxidation states

In almost all known compounds of oxygen, the oxidation state of oxygen is -2. The oxidation state -1 is found in a few compounds such as peroxides. 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).

Oxygen compounds as minerals

The most familiar oxygen compound is water, the oxide of hydrogen, H2O. Oxygen as a compound is also present in the atmosphere in trace quantities in the form of carbon dioxide (CO2). However the earth's crustal rock is composed predominantly of oxides of silicon as silica, SiO2 (found in granite and sand), silicates (found in feldspars), and oxygen compounds of metals such as calcium (as calcium carbonate in limestone), aluminium (as silicates in feldspars and as aluminium oxide in bauxite and corundum), iron (as iron (III) oxide Fe2O3 in hematite and rust), etc.

Organic compounds

Other important examples of oxygen compounds include the compounds of carbon and oxygen, such as alcohols (R-OH where "R" is an organic group), carbonyls (R-CO-H or R-CO-R) such as formaldehyde and acetone, and carboxylic acids (R-COOH) such as acetic acid and palmitic acid, a common fatty acid found in animals and plants and the food products derived from them, such as palm oil and milk products.

Inorganic compounds

Oxygenated radicals such as chlorates (ClO3),perchlorates (ClO4), chromates (CrO42−), dichromates (Cr2O72−), permanganates (MnO4), and nitrates (NO3) are strong oxidizing agents. Phosphorus is important in its oxygenated form as the phosphate (PO43−) ion, present in bone as calcium phosphate and as the backbone of RNA and DNA.

Oxygen forms heteropoly acids and polyoxometalate ions with tungsten, molybdenum and some other transition metals. Phosphotungstic acid (PTA), or dodecatungstophosphoric acid, has the chemical formula H3PW12O40, while octadecamolybdophosphoric acid is H6P2Mo18O62.

Oxygen forms compounds with almost all of the other known elements, including some of the rarest: technetium (TcO4), promethium (Pm2O3) and neptunium (NpO2); and also with some of the least reactive elements such as xenon (XeO3), gold (Au2O3) and platinum (PtO2). Synthetic elements that have known oxides include plutonium (PuO2), americium (AmO2), curium (CuO2), berkelium (BkO3), californium (Cf2O3) and einsteinium (Es2O3).

Unexpected compounds

One unexpected oxygen compound is dioxygen hexafluoroplatinate, O2+PtF6, discovered when Neil Bartlett was studying the properties of platinum hexafluoride (PtF6).[25] He noticed a change in color when this compound was exposed to atmospheric air and reasoned that xenon should be oxidized by PtF6. This led him to the discovery of xenon hexafluoroplatinate Xe+PtF6.

The cation O22+ in O2F2 is only formed in the presence of stronger oxidants than oxygen, which limits it to oxygen fluorides, e.g. oxygen fluoride.[21]

When dissolved in water, many metallic oxides form alkaline solutions while many oxides of nonmetals form acidic solutions. For example, sodium oxide in solution forms the strong base sodium hydroxide while phosphorus pentoxide in solution forms phosphoric acid.[26]

Oxides and peroxides

The alkali metals in groups 1 and 2 of the periodic table - lithium, sodium, potassium, rubidium, cesium, magnesium, calcium, strontium and barium, all react spontaneously with oxygen when exposed to air to form oxides and form hydroxides in the presence of water. None of these elements are found in nature as free metals. Cesium is so reactive with oxygen that it is used as a getter in vacuum tubes. The surface of aluminium is always oxidised in the presence of air, coated with a thin film of aluminium oxide that passivates the metal and slows further corrosion. The aluminium oxide layer can be built to greater thickness by the process of electrolytic anodising. Although solid magnesium and aluminium react slowly with oxygen at STP, they are both capable of burning in air, generating very high temperatures, and the metal powders may form explosive mixtures with air.

 

Some substances need to be heated before they will react with oxygen in bulk but some, such as iron, readily forms iron oxide, or rust, Fe2O3. The production of free oxygen by photosynthetic bacteria some 3.5 billion years ago precipitated iron out of solution in the oceans as Fe2O3 in the economically-important iron ore hematite.

Due to its electronegativity, oxygen forms chemical bonds with almost all other free elements at elevated temperatures to give corresponding oxides. So-called noble metals (common examples: gold, platinum) resist direct chemical combination with oxygen, and substances like gold(III) oxide must be formed by an indirect route.

Peroxides retain some of oxygen's original molecular structure(-O-O-). White or light yellow sodium peroxide (Na2O2) is formed when metallic sodium (Na) is burned in oxygen. Each oxygen atom in its peroxide ion may have a full octet of 4 pairs of electrons.[27] Superoxides are a class of compounds that are very similar to peroxides, but with just one unpaired electron for each pair of oxygen atoms (O2-).[27]. These compounds form from oxidation of alkali metals with larger ionic radii (K, Rb, Cs). For example, potassium superoxide ( KO2) is an orange-yellow solid formed when potassium (K) reacts with oxygen.

Hydrogen peroxide (H2O2) can be produced by passing a volume of 96% to 98% hydrogen and 2 to 4% oxygen through an electric discharge.[26] A more commercially-viable method is allow autoxidation of an organic intermediate; 2-ethylanthrahydroquinone dissolved in an organic solvent is oxidized to H2O2 and 2-ethylanthraquinone.[26] The 2-ethylanthraquinone is then reduced and recycled back into the process.

Silicates and silica

 

Silica is the common name for the compound silicon dioxide (SiO2). Quartz (illustrated) is the naturally-occurring crystalline mineral form of silica, and common deposits of quartz are in igneous rocks such as granite, sandstones, and sand.

Most chemically-combined oxygen is locked in a class of minerals called silicates (which, in turn, are the major component of rocks and clays). The basic structure of silicates consists of two parts; units of silicon surrounded by four oxygen anions in a tetrahedral arrangement and units of metal-oxygen polyhedra that contain metal cations (examples: aluminium, calcium, iron, and sodium).[27] Both units are linked together by sharing oxygen anions, forming complex polymers in the process.

Water- soluble silicates in the form of Na4SiO4, Na2SiO3, and Na2Si2O5 are used as detergents and adhesives.[27] NaxSixOx with a higher ratio of SiO2 to Na2O has a greater molecular weight and a lower solubility.

In organic compounds

 

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); amides (R-C(O)-NR2). There are many important organic solvents that contain oxygen, among which: acetone, methanol, ethanol, isopropanol, furan, THF, diethyl ether, dioxane, ethylacetate, DMF, DMSO, acetic acid, formic acid. Acetone ((CH3)2CO) and phenol (C6H5OH) 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, acetamide, etc. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.

Oxygen reacts spontaneously with many organic compounds at or below room temperature in a process called autoxidation.[26] Most of the organic compounds that contain oxygen are not made by direct action of oxygen. Organic compounds important in industry and commerce are made by direct oxidation of a precursor include[27]:

- Ethylene oxide (used to make the antifreeze ethylene glycol) is obtained by direct oxidation of ethylene: C2H4 + ½ O2 +catalyst→ C2H4O
- Peracetic acid (feeder material for various epoxy compounds) is obtained from acetaldehyde: CH3CHO + O2 +catalyst→ CH3C(O)-OOH

Of the organic compounds with biological relevance, carbohydrates (such as glucose) contain a large amount of oxygen. All fatty acids (such as oleic acid) and aminoacids contain oxygen (due to the presence of carboxyl group). Furthermore, seven of the amino acids incorporate oxygen in the side-chain too: serine, tyrosine, threonine, glutamic acid, glutamine, aspartic acid, and asparagine. Oxygen also occurs in phosphate groups in the biologically important energy-carrying molecules ATP and ADP and in the backbone of RNA and DNA.

Occurrence

 

See also: Silicate minerals and Category:Oxide minerals

Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.[2] About 0.87% of the Sun's mass is in the form of oxygen.[28] Oxygen constitutes 49.2% of the Earth's crust by mass[3] and is the most common component of the world's oceans (88.81% by mass).[28] It is also the second-most-common component of the Earth's atmosphere, taking up 20.947% of its volume and 23.14% of its mass (some million billion tonnes).[29][28][30] Earth is unusual in having such a high concentration of free oxygen in its atmosphere. With 0.15% oxygen by volume, the atmosphere of Mars has the second-most-abundant concentration by volume of any planet in the solar system, while Venus comes in third place.[2] However, their oxygen is solely produced by ultraviolet radiation impacting oxygen-containing molecules such as carbon dioxide.

  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 modern Earth's atmosphere. Because of the vast amounts of oxygen in the atmosphere, even if all photosynthesis were to cease, it would take at least 5,000 years to strip out more or less all oxygen.[31]

Free elemental dioxygen also occurs in solution in the world's water bodies. The higher solubility of O2 at low 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.[32] Polluted water may have reduced amounts of oxygen in it, depleted by decaying algae and other biomaterials (see eutrophication). Scientists assess this aspect of water quality by measuring the water's biochemical oxygen demand (BOD), or the amount of oxygen needed to restore a normal oxygen concentration.[6]

Isotopes and stellar origin

Main article: Isotopes of oxygen

  Naturally occurring oxygen is composed of 3 stable isotopes, 16O, 17O, and 18O, with 16O being the most abundant (99.762% natural abundance).[33] Oxygen isotopes range in mass number from 12 to 28.[33]

Relative and absolute abundance of 16O is due to it being a principal product of stellar evolution and the fact that it is a primary isotope, meaning it can be made by stars that were initially made exclusively of hydrogen.[34] Most 16O is synthesized at the end of the helium fusion process in stars; the triple-alpha reaction creates 12C, which captures an additional 4He to make 16O. The neon burning process creates additional 16O.[34]

Both 17O and 18O are secondary isotopes, meaning that their nucleosynthesis requires seed nuclei. 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.[34] Most 18O is produced when 14N (made abundant from CNO burning) captures a 4He nuclei, making 18O common in the helium-rich zones of stars.[34] A billion degrees Celsius are required for two oxygen nuclei to undergo nuclear fusion to form the heavier nucleus of sulfur.[2]

Fourteen radioisotopes have been characterized, with the most stable being 15O with a half-life of 122.24 s and 14O with a half-life of 70.606 s.[33] 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.[33] The most common decay mode before the stable isotopes is electron capture and the most common mode after is beta decay. The decay products before the stable isotopes are element 7 (nitrogen) isotopes and the products after are element 9 (fluorine) isotopes.[33]

An atomic mass of 16 was assigned to oxygen prior to the definition of the unified atomic mass unit based upon 12C.[35] Since physicists referred to 16O only, while chemists meant the naturally-abundant mixture of isotopes, this led to slightly different atomic mass scales.

The isotopic composition of oxygen atoms in the earth's atmosphere is 99.759% 16O, 0.037% 17O and 0.204% 18O.[28] Because water molecules containing the lighter isotope are slightly more likely to evaporate and fall as precipitation[36], fresh water and polar ice on earth contains slightly less (0.1981%) of the heavy isotope 18O than air (0.204%) or seawater containing (0.1995%). Since the 18O/16O isotope ratio in marine calcium carbonate equilibrates with that in the atmosphere, fluctuations in the oxygen isotope ratio in foraminifera can be used as a climate proxy, increasing during accumulation of polar ice and decreasing during warmer periods.

Biological role

Photosynthesis

  In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis in cyanobacteria, green algae and plants.[37] Algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on earth.[38] The remainder is produced by terrestrial plants, although almost all oxygen produced in tropical forests is consumed by organisms in those forests.[39]

The overall formula for photosynthesis is:

6CO2 + 6H2O + sunlight \longrightarrow C6H12O6 + 6O2

Or simply: carbon dioxide + water + sunlight \longrightarrow glucose + oxygen

Photolytic oxygen evolution part of photosynthesis occurs via the light-dependent oxidation of water to molecular oxygen and can be written as the following simplified chemical reaction:

2H2O \longrightarrow 4e- + 4H+ + O2

The reaction requires the energy of four photons. The electrons from the oxidized water molecules replace electrons in the P680 component of photosystem II, which have been removed into an electron transport chain via light-dependent excitation and resonance energy transfer onto plastoquinone.[40] Photosytem II therefore has also been referred to as water-plastoquinone oxido-reductase.[41] The protons are released into the thylakoid lumen, thus contributing to the generation of a proton gradient across the thylakoid membrane. This proton gradient is the driving force for ATP synthesis via photophosphorylation and coupling the absorption of light energy and photolysis of water to the creation of chemical energy during photosynthesis.[40]

Water oxidation is catalyzed by a manganese-containing enzyme complex associated with thylakoid membranes known as the oxygen evolving complex (OEC) or water-splitting complex. Manganese is an important cofactor, and calcium and chloride are also required for the reaction to occur.[40]

Cellular oxidations

 

See also: Aerobic respiration

DNA and proteins contain 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. Vertebrate animals use hemoglobin in their blood to transport oxygen from their lungs to their tissues, but other animals use hemocyanin (molluscs and some arthropods) or hemerythrin (spiders and lobsters).[29] A liter of blood can dissolve 200 cc of oxygen gas, which is much more than water can dissolve (see Physical Properties).[29]

In vertebrates, oxygen uptake is carried out by the following processes:

Oxygen diffuses through membranes and into red blood cells after inhalation into the lungs. The heme group of hemoglobin binds oxygen when it is present, and CO2 is released from another part of the hemoglobin molecule, as is acid which causes CO2 to be released from bicarbonate, its major reservoir in blood plasma (Bohr effect). After being carried in blood to a body tissue in need of oxygen, O2 is handed-off from the heme group to monooxygenase, an enzyme that also has an active site with an atom of iron.[29] Monooxygenase uses oxygen to catalyze many oxidation reactions in the body. Oxygen is also used as an electron acceptor in mitochondria to generate chemical energy in the form of adenosine triphosphate (ATP) during oxidative phosphorylation. Carbon dioxide, a waste product, is released from the cell and into the blood, where it combines with bicarbonate and hemoglobin for transport to the lungs. Blood circulates back to the lungs and the process repeats.[42] [9]

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).[29] The body uses superoxide dismutase to reduce superoxide radicals to hydrogen peroxide. Glutathione peroxidase and similar enzymes then convert the H2O2 to water and dioxygen.[29]

Parts of the immune system of higher organisms, 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.[43]

Biosynthesis: geologic timeline

  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 photosynthetic action of early anaerobes (archaea and bacteria). These organisms, fossil evidence for which occurs in the form of stromatolites and oncolites, developed the mechanism of oxygen evolution in the Archean era, between 3.5 and 2.7 billion years ago. At first, the produced oxygen dissolved in the oceans, where it was reduced by dissolved iron compounds, precipitating iron oxide (Fe2O3) and creating banded iron formations that are now a valuable resource of iron ore, hematite. Oxygen started to gas out of the oxygen-saturated waters from about 2.7 billion years ago, as is evident from the rusting of iron-rich terrestrial rocks starting around that time. The amount of oxygen in the atmosphere increased gradually at first and then more rapidly around 2.2 to 1.7 billion years ago to about 10% of its present level, as available reducing agents in the oceans and crustal rocks became oxidized.[44]

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.[45] 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 30% per volume.[46] Towards the end of the Carboniferous era (coal age) about 300 million years ago, atmospheric oxygen levels reached a maximum of 35% by volume,[46] 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), the most-common being 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.[9] It was estimated that, at the current rate of photosynthesis, it would take about 2,000 years to regenerate the entire oxygen in the present atmosphere.[47]

Anthropogenic Production

See also: Oxygen evolution and fractional distillation

Fractional distillation

Two major methods are employed to produce the 100 million tonnes of oxygen extracted from air for industrial uses annually.[48] 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.[48]

  The other major method of producing oxygen involves 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% oxygen.[48] Simultaneously, nitrogen is released from the other nitrogen-saturated zeolite bed, by reducing the chamber operating pressure and diverting part of the oxygen from the producer bed through it, in the reverse direction of flow. After a set cycle time, the operation of the two beds is switched, so that the producer bed is reverse purged and the purged bed becomes the producer bed. This allows for a continuous supply of gaseous oxygen to be pumped through a pipeline. It is known as pressure swing adsorption (PSA). Oxygen is increasingly obtained by these non-cryogenic technologies (see also the related vacuum-pressure swing adsorption (VPSA),[49] or vacuum swing adsorption (VSA) technolgies).

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 submarines, which are still part of standard equipment on commercial airliners in case of depressurization emergencies.

Another air separation technology involves forcing air to dissolve through ceramic membranes based on zirconium oxide by either high pressure or an electric current, to produce nearly pure oxygen.[6]

In large quantities, the price of liquid oxygen (2001) is approximately $0.21/kg.[50] Since the primary cost of production is the energy cost of liquefying the air, the production cost will change as energy cost varies.

Transportation

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 gaseous oxygen, at atmospheric pressure and 20°C.[48] Such tankers are used to refill bulk liquid oxygen storage containers, which stand outside hospitals and other institutions with a need for large volumes of pure oxygen. 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.[48]

Applications

See also: Breathing gas, Redox, and Combustion

Medical

  Uptake of oxygen from the air is the essential purpose of respiration, so oxygen supplementation is used in medicine. Oxygen therapy is used to treat emphysema, pneumonia, some heart disorders, and any disease that impairs the body's ability to take up and use oxygen.[51] 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. Hyperbaric medicine uses hyperbaric oxygen chambers to increase the partial pressure of oxygen around the patient and, when needed, the medical staff.

Carbon monoxide poisoning, gas gangrene, and decompression sickness (the "bends") are sometimes treated using these devices. Increased oxygen concentration in the lungs helps to displace carbon monoxide from the heme group of hemoglobin. Oxygen is poisonous to the anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them. Decompression sickness occurs in divers who decompress too quickly after a dive, resulting in bubbles of inert gas, mostly nitrogen and argon, forming in their blood. Increasing the pressure of oxygen as soon as possible is part of the treatment.[51]

Life support and recreational use

  A notable application of oxygen as a low-pressure breathing gas is in modern spacesuits, which surround their occupant's body with pressurized air. These devices use nearly pure oxygen at about one third normal pressure, resulting in a normal blood partial pressures of oxygen. This trade-off of higher oxygen concentration for lower pressure is needed to maintain flexible spacesuits.

Scuba divers and submariners also rely on artificially-delivered oxygen, but most often use normal pressure, and/or mixtures of oxygen and air. Pure or nearly pure oxygen use in diving at higher-than-sea-level pressures, is usually limited to rebreather, decompression, or emergency treatment use at relatively shallow depths (~ 6 meters depth, or less). Deeper diving requires significant dilution of oxygen with other gases, such as nitrogen or helium, to help prevent oxygen toxicity.

People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental oxygen supplies.[52] Passengers traveling in commercial airplanes have an emergency supply of oxygen automatically supplied to them in case of cabin depressurization. Sudden cabin pressure loss activates chemical oxygen generators above each seat, causing oxygen masks to drop and forcing iron fillings into the sodium chlorate inside the canister.[6] A steady stream of oxygen gas is produced by the exothermic reaction.[53]

Oxygen, as a supposed mild euphoric, has a history of recreational use in oxygen bars and in sports. Oxygen bars are establishments, found in Japan, California and Las Vegas, Nevada since the late 1990s that offer higher than normal oxygen exposure for a fee.[54] Professional athletes, especially in American football, also sometimes go off field between plays to wear oxygen masks in order to get a supposed "boost" in performance. However, the reality of a pharmacological effect is doubtful; a placebo or psychological boost being the most plausible explanation.[54]. Available studies support a performance boost from enriched oxygen mixtures only if they are breathed during actual aerobic exercise. [55]

Industrial

  Smelting of iron ore into steel consumes 55% of commercially-produced oxygen.[6] In this process, oxygen is injected through a high-pressure lance into molten iron, which removes sulfur impurities and excess carbon as the respective oxides, SO2 and CO2. The reactions are exothermic, so the temperature increases to 1700° C.[6]

Another 25% of commercially-produced oxygen is used by the chemical industry.[6] 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).[6]

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.[6] Oxygen is used in oxyacetylene welding burning acetylene with oxygen to produce a very hot flame. In this process, metal up to 60  cm thick is first heated with a small oxy-acetylene flame and then quickly cut by a large stream of oxygen.[56] Rocket propulsion requires a fuel and an oxidizer. Larger rockets use liquid oxygen as their oxidizer, which is mixed and ignited with the fuel for propulsion.

Scientific

  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 12% heavier oxygen-18.[57] 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.[57] Paleoclimatologists also directly measure this ratio in the water molecules of ice core samples that are up to several hundreds of thousands of years old.

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.[58] 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.

In human history

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.[59] 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.[60]

In the late 17th century, Robert Boyle proved that air is necessary for combustion. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called 'spiritus nitroaereus' or just 'nitroaereus'.[61] 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.[62] 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.[61] 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.[61] Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione".[62]

Phlogiston theory

  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.[16] This was largely due to the prevalence of a philosophy of combustion and corrosion called the phlogiston theory, which was then the favored explanation of how those processes worked.

Established in 1667 by German alchemist J. J. Becher, and modified by chemist Georg Ernst Stahl by 1731,[63] 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.[60]

Highly-combustible materials that leave little residuum, such as wood or coal, were thought of as made mostly of phlogiston, whereas non-combustible substances that corrode, 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 burns: that most common objects appear to become lighter and seem to lose something in the process.[60] The fact that a substance like wood actually gains overall weight in burning was hidden by the buoyancy of the gaseous combustion products. That metals actually gain weight in rusting (when they were supposed to be losing phlogiston) was one of the first clues that the phlogiston theory is incorrect.

Discovery

  An experiment conducted by the 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'.[28] 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."[16] 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.[64][60] 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 by about 1772.[60][28] 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.[48] 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 discovered the new substance independently. However, Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele also posted a letter to Lavoisier on September 30 1774 that described his own 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).[48]

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.[28] 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.[28] 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.[28] 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. "no life"), which did not support either.

Lavoisier later renamed 'vital air' to oxygène after the Greek roots meaning "acid-former" while 'azote' was renamed nitrogen.[28] 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.[48]

Later history

  Scientists realized by the late 19th century that compressing and cooling air could be used to liquefy and isolate its components. 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 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.[65] Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying oxygen.[65] Only a few drops of liquid oxygen were produced in either case so no meaningful analysis could be conducted.

In 1891, Scottish chemist James Dewar was able to produce enough liquid oxygen to study.[9] 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.[66] Later, in 1901, oxyacetylene welding was demonstrated for the first time by burning a mixture of acetylene and compressed oxygen. This method of welding and cutting metal later became common.[66]

In 1923, American scientist Robert H. Goddard became the first person to develop a rocket engine; the engine used gasoline for fuel and liquid oxygen as the oxidizer. Goddard successfully flew a small liquid-fueled rocket 56 m at around 97 kph on March 16 1926 in Auburn, Massachusetts.[66][67]

Precautions

Toxicity

 

Main articles: Oxygen toxicity and Oxygen intoxication

Oxygen can be toxic at elevated partial pressures, leading to convulsions and other health problems.[68] 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.[16] (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.)[16]

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.[69] In the case of spacesuits, the oxygen partial pressure in the breathing gas is, in general, about 0.30 bar (1.4 times normal), and the resulting oxygen partial pressure in the astronaut's arterial blood (due to downward adjustments due to water vapor and CO2 in the alveoli) is only marginally more than the normal sea-level oxygen partial pressure of 0.13 bar (see arterial blood gas).

In deep scuba diving and surface supplied diving, and when using equipment that 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 that produces lung damage may be considerably less than 50%. More important, 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.[16] This form of oxygen toxicity usually occurs after several hours exposure to oxygen partial pressures over about 1.4 atmospheres (bars) (i.e., seven 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 forms of oxygen such as ozone (O3), singlet oxygen, and some derivatives of oxygen such as hydrogen peroxide, hydroxyl radicals, and superoxide are also highly toxic.

Combustion hazards

0
0
0
OX

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.[70] Oxygen itself is not the fuel, but an oxidant.

  Concentrated oxygen will allow combustion to proceed rapidly and energetically.[70] 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 oxygen systems requires special training to ensure that ignition sources are minimized.[70] The fire that killed the Apollo 1 crew on a test launch pad spread so rapidly because the capsule was pressurized with pure oxygen, as would be usual in the beginning of an actual flight, at slightly more than atmospheric pressure, instead of the ⅓ normal pressure that would be used in the upper atmosphere and in space. (No single ignition source of the fire was conclusively identified, although some evidence points to arc from an electrical spark). [71][72]

Liquid oxygen spills, if allowed to soaked into organic matter, such as, wood, petrochemicals, and asphalt, can cause these materials to detonate unpredictably on subsequent mechanical impact.[70] On contact with the human body, it can also cause cryogenic burns to the skin and the eyes.

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.

See also

Notes

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  2. ^ a b c d Emsley 2001, p.297
  3. ^ a b Oxygen. Los Alamos National Laboratory. Retrieved on 2007-12-16.
  4. ^ Distribution of elements in the human body (by weight). The Internet Encyclopedia of Science. Retrieved on 2007-12-06.
  5. ^ a b c d e Mellor 1939
  6. ^ a b c d e f g h i Emsley 2001, p.301
  7. ^ wiseGeek.com on Oxygen. Retrieved on 2007-12-07.
  8. ^ Structure of Oxygen Molecule (triplet). Glasser Group, University of Missouri-Columbia. Retrieved on 2007-03-03.
  9. ^ a b c d Emsley 2001, p.303
  10. ^ Demonstration of a bridge of liquid oxygen supported against its own weight between the poles of a powerful magnet. University of Wisconsin-Madison Chemistry Department DEMONSTRATION LAB. Retrieved on 2007-12-15.
  11. ^ Company literature of Oxygen analyzers (triplet). Servomex. Retrieved on 2007-12-15.
  12. ^ Anja Krieger-Liszkay* (2005). "Singlet oxygen production in photosynthesis.". Journal of Experimental Botanics 56: 337-46. Oxford Journals. Retrieved on 2007-12-16.
  13. ^ Harrison, Roy M. (1990). Pollution: Causes, Effects & Control. (2nd Edition). Cambridge: Royal Society of Chemistry. ISBN 0-85186-283-7.
  14. ^ a b Paul Wentworth Jr., Jonathan E. McDunn, Anita D. Wentworth, Cindy Takeuchi, Jorge Nieva, Teresa Jones, Cristina Bautista, Julie M. Ruedi, Abel Gutierrez, Kim D. Janda, Bernard M. Babior, Albert Eschenmoser, Richard A. Lerner (2002-12-13). "Evidence for Antibody-Catalyzed Ozone Formation in Bacterial Killing and Inflammation". Science 298: 2195 - 219. doi:10.1126/science.1077642. Retrieved on 2007-12-16.
  15. ^ Osamu Hirayama, Kyoko Nakamura, Syoko Hamada and Yoko Kobayasi (1994-02-). "Singlet oxygen quenching ability of naturally occurring carotenoids". Lipids 29 (2): 149-150. Springer Berlin / Heidelberg. doi:10.1007/BF02537155. Retrieved on 2007-12-15.
  16. ^ a b c d e f Emsley 2001, p.299
  17. ^ Air solubility in water. The Engineering Toolbox. Retrieved on 2007-12-21.
  18. ^ Overview of Cryogenic Air Separation and Liquefier Systems. Universal Industrial Gases, Inc.. Retrieved on 2007-12-15.
  19. ^ Liquid Oxygen Material Safety Data Sheet. Matheson Tri Gas. Retrieved on 2007-12-15.
  20. ^ Chieh, Chung. Bond Lengths and Energies. University of Waterloo. Retrieved on 2007-12-16.
  21. ^ a b Cotton, F. Albert and Wilkinson, Geoffrey (1972). Advanced Inorganic Chemistry: A comprehensive Text. (3rd Edition). New York, London, Sydney, Toronto: Interscience Publications. ISBN 0-471-17560-9.
  22. ^ Ball, Philip. "New form of oxygen found", news@nature.com, November 16, 2001. Retrieved on 2007-03-03. 
  23. ^ F. Cacace, G. de Petris, A. Troiani, (2001). "Experimental Detection of Tetraoxygen". Angewandte Chemie International Edition 40 (21): 4062 - 4065.
  24. ^ Peter P. Edwards and Friedrich Hensel (2002-01-14). "Metallic Oxygen". ChemPhysChem 3 (1): 53 - 56. Retrieved on 2007-12-16.
  25. ^ Cook 1968, p.505
  26. ^ a b c d Cook 1968, p.506
  27. ^ a b c d e Cook 1968, p.507
  28. ^ a b c d e f g h i j Cook 1968, p.500
  29. ^ a b c d e f Emsley 2001, p.298
  30. ^ Figures given are for values up to 50 miles above the surface
  31. ^ Walker, J. C. G. (1980) The oxygen cycle in the natural environment and the biogeochemical cycles, Springer-Verlag, Berlin, Federal Republic of Germany (DEU)
  32. ^ 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.
  33. ^ a b c d e Oxygen Nuclides / Isotopes. EnvironmentalChemistry.com. Retrieved on 2007-12-17.
  34. ^ a b c d Meyer, B.S. (September 19-21, 2005). "NUCLEOSYNTHESIS AND GALACTIC CHEMICAL EVOLUTION OF THE ISOTOPES OF OXYGEN" in Workgroup on Oxygen in the Earliest Solar System. Proceedings of the NASA Cosmochemistry Program and the Lunar and Planetary Institute. 9022. Retrieved on 2007-12-23. 
  35. ^ Mellor 1939, Chapter VI, Section 7
  36. ^ Dansgaard, W (1964) Stable isotopes in precipitation. Tellus 16, 436-468
  37. ^ Raven, Peter H.; Ray F. Evert, Susan E. Eichhorn (2005). Biology of Plants, 7th Edition. New York: W.H. Freeman and Company Publishers, 115-127. ISBN 0-7167-1007-2. 
  38. ^ Fenical, William (September 1983). "Marine Plants: A Unique and Unexplored Resource", Plants: the potentials for extracting protein, medicines, and other useful chemicals (workshop proceedings). DIANE Publishing, 147. ISBN 1428923977. 
  39. ^ Broeker, W.S. (2006). Breathing easy, Et tu, O2. Columbia University. Retrieved on 2007-10-21.
  40. ^ a b c Raven, Peter H.; Ray F. Evert, Susan E. Eichhorn (2005). Biology of Plants, 7th Edition. New York: W.H. Freeman and Company Publishers, 115-127. ISBN 0-7167-1007-2. 
  41. ^ Raval M, Biswal B, Biswal U (2005). "The mystery of oxygen evolution: analysis of structure and function of photosystem II, the water-plastoquinone oxido-reductase". Photosynthesis Research 85 (3): 267-93. doi:10.1007/s11120-005-8163-4. PMID 16170631.
  42. ^ During oxidative phosphorylation, oxygen is reduced to water. In contrast, 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.
  43. ^ Hoffmann, Roald (2004). "The Story of O". American Scientist 92 (1): 23. Retrieved on 2007-03-03.
  44. ^ Campbell, Neil A.; Reece, Jane B. (2005). Biology, 7th Edition. San Francisco: Pearson - Benjamin Cummings, 522-523. ISBN 0-8053-7171-0. 
  45. ^ Freeman, Scott (2005). Biological Science, 2nd Edition. Upper Saddle River, NJ: Pearson - Prentice Hall, 214, 586. ISBN 0-13-140941-7. 
  46. ^ a b Robert A. Berner (1999-09-18). "Atmospheric oxygen over Phanerozoic time". Proceedings of the National Academy of Sciences of the USA (20): 10955-10957. Retrieved on 2007-12-16.
  47. ^ Malcom Dole (1965). "The Natural History of Oxygen". The Journal of General Physiology 49: 5-27. Retrieved on 2007-12-16.
  48. ^ a b c d e f g h Emsley 2001, p.300
  49. ^ Non-Cryogenic Air Separation Processes. UIG Inc. (2003). Retrieved on 2007-12-16.
  50. ^ , National Aeronautics and Space Administration, 2001=09, . Retrieved on 2007-12-16
  51. ^ a b Cook 1968, p.510
  52. ^ 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.
  53. ^ Even this may pose a danger if inappropriately triggered: a ValuJet airplane crashed after use-date-expired oxygen canisters, which were being shipped in the cargo hold, activated and caused fire. (They were mis-labeled as empty, and carried against dangerous goods regulations). (NTSB Summary report. National Transportation Safety Board. Retrieved on 2007-12-16.)
  54. ^ a b Bren, Linda (November-December 2002). Oxygen Bars: Is a Breath of Fresh Air Worth It?. FDA Consumer magazine. U.S. Food and Drug Administration. Retrieved on 2007-12-23.
  55. ^ [1] Accessed Jan 4, 2008
  56. ^ Cook 1968, p.508
  57. ^ a b Emsley 2001, p.304
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  59. ^ Jastrow, Joseph (1936). Story of Human Error. Ayer Publishing, 171. ISBN 0836905687. Retrieved on 2007-12-16. 
  60. ^ a b c d e Cook 1968, p.499.
  61. ^ a b c Britannica contributors (1911). "John Mayow", Encyclopaedia Britannica, 11th edition. Retrieved on 2007-12-16. 
  62. ^ a b World of Chemistry contributors (2005). "John Mayow", World of Chemistry. Thomson Gale. Retrieved on 2007-12-16. 
  63. ^ Morris, Richard (2003). The last sorcerers: The path from alchemy to the periodic table (Hardback), Washington, D.C.: Joseph Henry Press. ISBN 0309089050. 
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  65. ^ a b Daintith 1994, p.707
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  67. ^ Goddard-1926. NASA. Retrieved on 2007-11-18.
  68. ^ Cook 1968, p.511
  69. ^ Wade, Mark (2007). Space Suits. Encyclopedia Astronautica. Retrieved on 2007-12-16.
  70. ^ a b c d Werley, Barry L. (Edtr.) (1991). Fire Hazards in Oxygen Systems. (ASTM Technical Professional training: ASTM Subcommittee G-4.05) Philadelphia: ASTM International.
  71. ^ Report of Apollo 204 Review Board NASA Historical Reference Collection, NASA History Office, NASA HQ, Washington, DC
  72. ^ Chiles, James R. (2001). Inviting Disaster: Lessons from the edge of Technology: An inside look at catastrophes and why they happen. New York: HarperCollins Publishers Inc. ISBN 0-06-662082-1.

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This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Oxygen". A list of authors is available in Wikipedia.
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