Oxygen



8 fluorine
-
↑
O
↓
S
General
number oxygen, O, 8
chalcogens
block p
Appearance

colorless gas above
light blue liquid
(3) g·mol−1
Electron configuration 1s2 2s2 2p4
shell 2, 6
Physical properties
PhasekJ·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
Electronegativity3.44 (Pauling scale)
more) 1st: 1313.9 kJ·mol−1
2nd: 3388.3 kJ·mol−1
3rd: 5300.5 kJ·mol−1
Van der Waals radius152 pm
Miscellaneous
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 neutrons
17O 0.038% O is neutrons
18O 0.21% O is neutrons
References
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Oxygen (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 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 obligate anaerobic organisms and was a poisonous waste product for early life on Earth.

Characteristics

Structure

Main articles: Singlet oxygen

At double bond.[8]  

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 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]

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 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 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 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 oxidizer than either O2 or O3.[22][23] When tetraoxygen is subjected to a pressure of 96 GPa, it becomes polonium, both of which show significant metallic character.

Compounds

 

Main article: Oxygen compounds

Oxidation states

In almost all known compounds of oxygen, the 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 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), fatty acid found in animals and plants and the food products derived from them, such as palm oil and milk products.

Inorganic compounds

Oxygenated DNA.

Oxygen forms heteropoly acids and 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: einsteinium (Es2O3).

Unexpected compounds

One unexpected oxygen compound is dioxygen hexafluoroplatinate, O2+PtF6−, discovered when 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 phosphoric acid.[26]

Oxides and peroxides

The 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 hematite.

Due to its platinum) resist direct chemical combination with oxygen, and substances like gold(III) oxide must be formed by an indirect route.

potassium (K) reacts with oxygen.

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

 

sandstones, and sand.

Most chemically-combined oxygen is locked in a class of polymers in the process.

Water-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); acetamide, etc. ethers in which the oxygen atom is part of a ring of three atoms.

Oxygen reacts spontaneously with many 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]:

- catalyst→ C2H4O
- acetaldehyde: CH3CHO + O2 +catalyst→ CH3C(O)-OOH

Of the organic compounds with biological relevance, carbohydrates (such as DNA.

Occurrence

 

See also: 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 carbon dioxide.

  The unusually high concentration of elemental oxygen on Earth is the result of the 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] 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 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 He 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 sulfur.[2]

Fourteen 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 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 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 plastoquinone.[40] Photosytem II therefore has also been referred to as water-plastoquinone oxido-reductase.[41] The protons are released into the 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 chloride are also required for the reaction to occur.[40]

Cellular oxidations

 

See also: Aerobic respiration

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 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]

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 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 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 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 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: 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 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 vacuum swing adsorption (VSA) technolgies).

Oxygen can also be produced through 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 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°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. partial pressure of oxygen around the patient and, when needed, the medical staff.

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 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 exothermic reaction.[53]

Oxygen, as a supposed mild euphoric, has a history of recreational use in 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

  C.[6]

Another 25% of commercially-produced oxygen is used by the chemical industry.[6] Ethylene is reacted with oxygen to create fabrics).[6]

Most of the remaining 20% of commercially-produced oxygen is used in medical applications, 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, ice core samples that are up to several hundreds of thousands of years old.

Oxygen presents two spectrophotometric 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 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, 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, phlogiston theory, which was then the favored explanation of how those processes worked.

Established in 1667 by German alchemist 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 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 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 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 "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 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 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 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 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 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] 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 cryogenic burns to the skin and the eyes.

Combustion hazards also apply to compounds of oxygen with a high oxidative potential, such as dichromates because they can donate oxygen to a fire.

See also

Notes

  1. ^ (1997) IUPAC Compendium of Chemical Terminology, 2, IUPAC. Retrieved on 2007-12-15. 
  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 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
  58. ^ Zarco-Tejada, P.J., Miller, J.R.; Berger, M., Alonso, L., Cerovic, Z., Goulas, Y., Jacquemoud, S., Louis, J., Mohammed, G. Moya, I., Pedros, R., Moreno, J.F., Verhoef, W.. Progress on the development of an integrated canopy fluorescence model. Geoscience and Remote Sensing Symposium, 2003. IGARSS '03. Proceedings. 2003 IEEE International. Retrieved on 2007-12-16.
  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. 
  64. ^ Joseph Priestley (1775). "An Account of Further Discoveries in Air". Philosophical Transactions 65: 384-394. Retrieved on 2007-12-16.
  65. ^ a b Daintith 1994, p.707
  66. ^ a b c (2002) "Oxygen", How Products are Made. The Gale Group, Inc. Retrieved on 2007-12-16. 
  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.

References

  • Agostini, D.; H. Iida, and A. Takahashi (1995). "Positron emission tomography with oxygen-15 of stunned myocardium caused by coronary artery vasospasm after recovery.". British Heart Journal 73 (1): 69–72. Retrieved on 2007-12-16.
  • Balachandran, S. C. (October 9-11, 1995). "Carbon and Oxygen Nucleosynthesis in the Galaxy: Problems and Prospects". Proceedings of the sixth (6th) annual October Astrophysics Conference: 188-195, College Park; Maryland: Astronomical Society of the Pacific. Retrieved on 2007-01-08. 
  • Berner, Robert A. (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.
  • Braker, William; Mossman, Allen L. (1971). "Oxygen", Matheson Gas Data Book. East Rutherford: Matheson Gas Products, Pp 441-445. 
  • Dole, Malcom (1965). "The Natural History of Oxygen". The Journal of General Physiology 49: 5-27. Retrieved on 2007-12-16.
  • Cook, Gerhard A.; Lauer, Carol M. (1968). "Oxygen", in Clifford A. Hampel: The Encyclopedia of the Chemical Elements. New York: Reinhold Book Corporation, 499-512. LCCN 68-29938. 
  • Cotton, F. Albert; Wilkinson, Geoffrey (1972). "Chapter 14: Oxygen", Advanced Inorganic Chemistry: A Comprehensive Text (3rd Edition). New York, London, Sidney, Toronto: Interscience Publishers, Pp 403-420. 
  • Daintith, John (1994). Biographical Encyclopedia of Scientists. CRC Press. ISBN 0750302879. 
  • Edwards, Peter P.; Friedrich Hensel (2002-01-14). "Metallic Oxygen". ChemPhysChem 3 (1): 53 - 56. Retrieved on 2007-12-16.
  • Emsley, John (2001). "Oxygen", Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press, 297-304. ISBN 0198503407. 
  • Gribbin, John (2002). The Scientists:A History of Science Told Through the Lives of Its Greatest Inventors. Random House. ISBN 0307290719. 
  • Hirayama, Osamu; 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.
  • Krieger-Liszkay*, Anja (2005). "Singlet oxygen production in photosynthesis.". Journal of Experimental Botanics 56: 337-46. Oxford Journals. Retrieved on 2007-12-16.
  • Parks, G. D.; Mellor, J. W. (1939). Mellor's Modern Inorganic Chemistry, 6th edition, London: Longmans, Green and Co. 
  • Ventura1, P.; F. D'Antona and I. Mazzitelli (2002). "Yields from low metallicity, intermediate mass AGB stars: Their role for the CNO and lithium abundances in Globular Cluster stars". Astronomy&Astrophysics 393: 215-223. doi:10.1051/0004-6361:20021001. Retrieved on 2007-12-16.
  • Wentworth Jr., Paul; 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.
 v • d • e chemical elements

Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2

v • d • e
E numbers

pH regulators & anti-caking agents (E500–599) • Flavour enhancers (E600–699) • Miscellaneous (E900–999) • Additional chemicals (E1100–1599)

Foaming agents (E990–999)

Calcium peroxide (E930) • Hydrogen (E949)

 
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.