Nitrogen

Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in breaking the bond to convert the N₂ into useful compounds, but releasing large amounts of often useful energy, when these compounds burn, explode, or decay back into nitrogen gas.

The element nitrogen was discovered by Scottish physician Daniel Rutherford in 1772. Nitrogen occurs in all living organisms. It is a constituent element of amino acids and thus of proteins, and of nucleic acids (DNA and RNA). It resides in the chemical structure of almost all neurotransmitters, and is a defining component of alkaloids, biological molecules produced by many organisms.

1 History
2 Properties
3 Occurrence
4 Compounds
5 Production and applications
6 Biological role
7 Safety
8 See also
9 References
10 Further reading
11 External links

[edit] History

Nitrogen (Latin nitrogenium, where nitrum (from Greek nitron î�î��„ρî¿î�) means "saltpetre" (see nitre), and genes î�î�î�î�ς means "forming") is formally considered to have been discovered by Daniel Rutherford in 1772, who called it noxious air or fixed air.^[1] That there was a fraction of air that did not support combustion was well-known to the late 18th century chemist. Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as "mephitic air" or azote, from the Greek word î�î�ω�„î¿ς (azotos) meaning "lifeless".^[2] Animals died in it, and it was the principal component of air in which animals had suffocated and flames had burned to extinction. Lavoisier's name for nitrogen is used in many languages (French, Russian, etc.) and still remains in English in the common names of many compounds, such as hydrazine and compounds of the azide ion. Compounds of nitrogen were known in the Middle Ages.

The alchemists knew nitric acid as aqua fortis (strong water). The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold (the king of metals). The earliest military, industrial and agricultural applications of nitrogen compounds involved uses of saltpeter (sodium nitrate or potassium nitrate), notably in gunpowder, and later as fertilizer. In 1910, Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", an allotrope considered to be monatomic. The "whirling cloud of brilliant yellow light" produced by his apparatus reacted with quicksilver to produce explosive mercury nitride.^[3]

[edit] Properties

Nitrogen is a nonmetal, with an electronegativity of 3.04. It has five electrons in its outer shell and is therefore trivalent in most compounds. The triple bond in molecular nitrogen (N₂) is the strongest. The resulting difficulty of converting N₂ into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N₂, have dominated the role of nitrogen in both nature and human economic activities.

At atmospheric pressure molecular nitrogen condenses (liquefies) at 77 K ('195.8 –C) and freezes at 63 K ('210.0 –C) into the beta hexagonal close-packed crystal allotropic form. Below 35.4 K ('237.6 –C) nitrogen assumes the alpha cubic crystal allotropic form. Liquid nitrogen, a fluid resembling water in appearance, but with 80.8% of the density (the density of liquid nitrogen at its boiling point is 0.808 g/mL), is a common cryogen.

Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like N₃ and N₄.^[4] Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced using a diamond anvil cell, nitrogen polymerizes into the single-bonded cubic gauche crystal structure. This structure is similar to that of diamond, and both have extremely strong covalent bonds. N₄ is nicknamed "nitrogen diamond."^[5]

[edit] Isotopes

[edit] Electromagnetic spectrum

A 1×5 cm vial of glowing ultrapure nitrogen

Nitrogen discharge (spectrum) tube

Molecular nitrogen (¹⁴N₂) is largely transparent to infrared and visible radiation because it is a homonuclear molecule and thus has no dipole moment to couple to electromagnetic radiation at these wavelengths. Significant absorption occurs at extreme ultraviolet wavelengths, beginning around 100 nanometers. This is associated with electronic transitions in the molecule to states in which charge is not distributed evenly between nitrogen atoms. Nitrogen absorption leads to significant absorption of ultraviolet radiation in the Earth's upper atmosphere and the atmospheres of other planetary bodies. For similar reasons, pure molecular nitrogen lasers typically emit light in the ultraviolet range.

Nitrogen also makes a contribution to visible air glow from the Earth's upper atmosphere, through electron impact excitation followed by emission. This visible blue air glow (seen in the polar aurora and in the re-entry glow of returning spacecraft) typically results not from molecular nitrogen, but rather from free nitrogen atoms combining with oxygen to form nitric oxide (NO).

[edit] Reactions

Structure of [Ru(NH₃)₅(N₂)]²⁺.

Nitrogen is generally unreactive at standard temperature and pressure. N₂ reacts spontaneously with few reagents, being resilient to acids and bases as well as oxidants and most reductants. When nitrogen reacts spontaneously with a reagent, the net transformation is often called nitrogen fixation.

Nitrogen reacts with elemental lithium.^[7] Lithium burns in an atmosphere of N₂ to give lithium nitride:

6 Li + N₂ ' 2 Li₃N

Magnesium also burns in nitrogen, forming magnesium nitride.

3 Mg + N₂ ' Mg₃N₂

N₂ forms a variety of adducts with transition metals. The first example of a dinitrogen complex is [Ru(NH₃)₅(N₂)]²⁺ (see figure at right). Such compounds are now numerous, other examples include IrCl(N₂)(PPh₃)₂, W(N₂)₂(Ph₂CH₂CH₂PPh₂)₂, and [(î�⁵-C₅Me₄H)₂Zr]₂(î�₂, î�–,î�–-N₂). These complexes illustrate how N₂ might bind to the metal(s) in nitrogenase and the catalyst for the Haber process.^[8] A catalytic process to reduce N₂ to ammonia with the use of a molybdenum complex in the presence of a proton source was published in 2005.^[7] (see nitrogen fixation)

The starting point for industrial production of nitrogen compounds is the Haber process, in which nitrogen is fixed by reacting N₂ and H₂ over an iron(III) oxide (Fe₃O₄) catalyst at about 500 –C and 200 atmospheres pressure. Biological nitrogen fixation in free-living cyanobacteria and in the root nodules of plants also produces ammonia from molecular nitrogen. The reaction, which is the source of the bulk of nitrogen in the biosphere, is catalyzed by the nitrogenase enzyme complex which contains Fe and Mo atoms, using energy derived from hydrolysis of adenosine triphosphate (ATP) into adenosine diphosphate and inorganic phosphate ('20.5 kJ/mol).

[edit] Occurrence

Nitrogen is the largest single constituent of the Earth's atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air). It is created by fusion processes in stars, and is estimated to be the 7th most abundant chemical element by mass in the universe.^[9]

Molecular nitrogen and nitrogen compounds have been detected in interstellar space by astronomers using the Far Ultraviolet Spectroscopic Explorer.^[10] Molecular nitrogen is a major constituent of the Saturnian moon Titan's thick atmosphere, and occurs in trace amounts in other planetary atmospheres.^[11]

Nitrogen is present in all living organisms, in proteins, nucleic acids and other molecules. It typically makes up around 4% of the dry weight of plant matter, and around 3% of the weight of the human body. It is a large component of animal waste (for example, guano), usually in the form of urea, uric acid, ammonium compounds and derivatives of these nitrogenous products, which are essential nutrients for all plants that cannot fix atmospheric nitrogen.

Nitrogen occurs naturally in many minerals, such as saltpetre (potassium nitrate), Chile saltpetre (sodium nitrate) and sal ammoniac (ammonium chloride). Most of these are uncommon, partly because of the minerals' ready solubility in water. See also Nitrate minerals and Ammonium minerals.

[edit] Compounds

[edit] Production and applications

A computer rendering of the nitrogen molecule, N₂.

Nitrogen gas is an industrial gas produced by the fractional distillation of liquid air, or by mechanical means using gaseous air (i.e. pressurized reverse osmosis membrane or Pressure swing adsorption). Commercial nitrogen is often a byproduct of air-processing for industrial concentration of oxygen for steelmaking and other purposes. When supplied compressed in cylinders it is often called OFN (oxygen-free nitrogen).^[12]

Nitrogen gas has a variety of applications, including serving as an inert replacement for air where oxidation is undesirable;

As a modified atmosphere, pure or mixed with carbon dioxide, to preserve the freshness of packaged or bulk foods (by delaying rancidity and other forms of oxidative damage)
In ordinary incandescent light bulbs as an inexpensive alternative to argon.^[13]
The production of electronic parts such as transistors, diodes, and integrated circuits
Dried and pressurized, as a dielectric gas for high voltage equipment
The manufacturing of stainless steel^[14]
Use in military aircraft fuel systems to reduce fire hazard, (see inerting system)
On top of liquid explosives as a safety measure
Filling automotive and aircraft tires^[15] due to its inertness and lack of moisture or oxidative qualities, as opposed to air, though this is not necessary for consumer automobiles.^[16]^[17]
Use as a propellant for draught wine, and as an alternative to or together with carbon dioxide for other beverages.

Nitrogen is commonly used during sample preparation procedures for chemical analysis. Specifically, it is used to concentrate and reduce the volume of liquid samples. Directing a pressurized stream of nitrogen gas perpendicular to the surface of the liquid allows the solvent to evaporate while leaving the solute(s) and un-evaporated solvent behind.^[18]

Nitrogen tanks are also replacing carbon dioxide as the main power source for paintball guns. The downside is that nitrogen must be kept at higher pressure than CO₂, making N₂ tanks heavier and more expensive.

[edit] Nitrogenated beer

A further example of its versatility is its use as a preferred alternative to carbon dioxide to pressurize kegs of some beers, particularly stouts and British ales, due to the smaller bubbles it produces, which make the dispensed beer smoother and headier. A modern application of a pressure sensitive nitrogen capsule known commonly as a "widget" now allows nitrogen charged beers to be packaged in cans and bottles.^[19]

A mixture of nitrogen and carbon dioxide can be used for this purpose as well, to maintain the saturation of beer with carbon dioxide.^[20]

[edit] Liquid nitrogen

Air balloon submerged into liquid nitrogen

Main article: Liquid nitrogen

Liquid nitrogen is a cryogenic liquid. At atmospheric pressure, it boils at '195.8 –C. When insulated in proper containers such as Dewar flasks, it can be transported without much evaporative loss.^[21]

Like dry ice, the main use of liquid nitrogen is as a refrigerant. Among other things, it is used in the cryopreservation of blood, reproductive cells (sperm and egg), and other biological samples and materials. It is used medically in cryotherapy to remove cysts and warts on the skin^[22]. It is used in cold traps for certain laboratory equipment and to cool X-ray detectors. It has also been used to cool central processing units and other devices in computers which are overclocked, and which produce more heat than during normal operation.^[23]

[edit] Applications of nitrogen compounds

Molecular nitrogen (N₂) in the atmosphere is relatively non-reactive due to its strong bond, and N₂ plays an inert role in the human body, being neither produced nor destroyed. In nature, nitrogen is converted into biologically (and industrially) useful compounds by lightning, and by some living organisms, notably certain bacteria (i.e. nitrogen fixing bacteria ' see Biological role below). Molecular nitrogen is released into the atmosphere in the process of decay, in dead plant and animal tissues.

The ability to combine or fix molecular nitrogen is a key feature of modern industrial chemistry, where nitrogen and natural gas are converted into ammonia via the Haber process. Ammonia, in turn, can be used directly (primarily as a fertilizer, and in the synthesis of nitrated fertilizers), or as a precursor of many other important materials including explosives, largely via the production of nitric acid by the Ostwald process.

The organic and inorganic salts of nitric acid have been important historically as convenient stores of chemical energy. They include important compounds such as potassium nitrate (or saltpeter used in gunpowder) and ammonium nitrate, an important fertilizer and explosive (see ANFO). Various other nitrated organic compounds, such as nitroglycerin and trinitrotoluene, and nitrocellulose, are used as explosives and propellants for modern firearms. Nitric acid is used as an oxidizing agent in liquid fueled rockets. Hydrazine and hydrazine derivatives find use as rocket fuels and monopropellants. In most of these compounds, the basic instability and tendency to burn or explode is derived from the fact that nitrogen is present as an oxide, and not as the far more stable nitrogen molecule (N₂) which is a product of the compounds' thermal decomposition. When nitrates burn or explode, the formation of the powerful triple bond in the N₂ produces most of the energy of the reaction.

Nitrogen is a constituent of molecules in every major drug class in pharmacology and medicine. Nitrous oxide (N₂O) was discovered early in the 19th century to be a partial anesthetic, though it was not used as a surgical anesthetic until later. Called "laughing gas", it was found capable of inducing a state of social disinhibition resembling drunkenness. Other notable nitrogen-containing drugs are drugs derived from plant alkaloids, such as morphine (there exist many alkaloids known to have pharmacological effects; in some cases they appear natural chemical defenses of plants against predation). Drugs that contain nitrogen include all major classes of antibiotics, and organic nitrate drugs like nitroglycerin and nitroprusside that regulate blood pressure and heart action by mimicking the action of nitric oxide.

[edit] Biological role

Nitrogen is an essential building block of amino and nucleic acids, essential to life on Earth.

Elemental nitrogen in the atmosphere cannot be used directly by either plants or animals, and must be converted to a reduced (or 'fixed') state in order to be useful for higher plants and animals. Precipitation often contains substantial quantities of ammonium and nitrate, thought to result from nitrogen fixation by lightning and other atmospheric electric phenomena.^[24] This was first proposed by Liebig in 1827 and later confirmed.^[24] However, because ammonium is preferentially retained by the forest canopy relative to atmospheric nitrate, most fixed nitrogen reaches the soil surface under trees as nitrate. Soil nitrate is preferentially assimilated by these tree roots relative to soil ammonium.

Specific bacteria (e.g. Rhizobium trifolium) possess nitrogenase enzymes which can fix atmospheric nitrogen (see nitrogen fixation) into a form (ammonium ion) that is chemically useful to higher organisms. This process requires a large amount of energy and anoxic conditions. Such bacteria may live freely in soil (e.g. Azotobacter) but normally exist in a symbiotic relationship in the root nodules of leguminous plants (e.g. clover, Trifolium, or soybean plant, Glycine max). Nitrogen-fixing bacteria are also symbiotic with a number of unrelated plant species such as alders (Alnus) spp., lichens (Casuarina), Myrica, liverworts, and Gunnera.^[25]

As part of the symbiotic relationship, the plant converts the 'fixed' ammonium ion to nitrogen oxides and amino acids to form proteins and other molecules, (e.g. alkaloids). In return for the 'fixed' nitrogen, the plant secretes sugars to the symbiotic bacteria.^[25]

Plants are able to assimilate nitrogen directly in the form of nitrates which may be present in soil from natural mineral deposits, artificial fertilizers, animal waste, or organic decay (as the product of bacteria, but not bacteria specifically associated with the plant). Nitrates absorbed in this fashion are converted to nitrites by the enzyme nitrate reductase, and then converted to ammonia by another enzyme called nitrite reductase.^[25]

Nitrogen compounds are basic building blocks in animal biology as well. Animals use nitrogen-containing amino acids from plant sources, as starting materials for all nitrogen-compound animal biochemistry, including the manufacture of proteins and nucleic acids. Plant-feeding insects are dependent on nitrogen in their diet, such that varying the amount of nitrogen fertilizer applied to a plant can affect the reproduction rate of insects feeding on fertilized plants.^[26]

Soluble nitrate is an important limiting factor in the growth of certain bacteria in ocean waters.^[27] In many places in the world, artificial fertilizers applied to crop-lands to increase yields result in run-off delivery of soluble nitrogen to oceans at river mouths. This process can result in eutrophication of the water, as nitrogen-driven bacterial growth depletes water oxygen to the point that all higher organisms die. Well-known "dead zone" areas in the U.S. Gulf Coast and the Black Sea are due to this important polluting process.

Many saltwater fish manufacture large amounts of trimethylamine oxide to protect them from the high osmotic effects of their environment (conversion of this compound to dimethylamine is responsible for the early odor in unfresh saltwater fish ^[28]. In animals, free radical nitric oxide (NO) (derived from an amino acid), serves as an important regulatory molecule for circulation.^[27]

Animal metabolism of NO results in production of nitrite. Animal metabolism of nitrogen in proteins generally results in excretion of urea, while animal metabolism of nucleic acids results in excretion of urea and uric acid. The characteristic odor of animal flesh decay is caused by the creation of long-chain, nitrogen-containing amines, such as putrescine and cadaverine which are (respectively) breakdown products of the amino acids ornithine and lysine in decaying proteins.^[29]

Decay of organisms and their waste products may produce small amounts of nitrate, but most decay eventually returns nitrogen content to the atmosphere, as molecular nitrogen. The circulation of nitrogen from atmosphere, to organic compounds, then back to the atmosphere, is referred to as the nitrogen cycle.^[25]

[edit] Safety

Rapid release of nitrogen gas into an enclosed space can displace oxygen, and therefore represents an asphyxiation hazard. This may happen with few warning symptoms, since the human carotid body is a relatively slow and a poor low-oxygen (hypoxia) sensing system.^[30] An example occurred shortly before the launch of the first Space Shuttle mission in 1981, when two technicians lost consciousness and died after they walked into a space located in the Shuttle's Mobile Launcher Platform that was pressurized with pure nitrogen as a precaution against fire. The technicians would have been able to exit the room if they had experienced early symptoms from nitrogen-breathing.

When inhaled at high partial pressures (more than about 4 bar, encountered at depths below about 30 m in scuba diving) nitrogen begins to act as an anesthetic agent. It can cause nitrogen narcosis, a temporary semi-anesthetized state of mental impairment similar to that caused by nitrous oxide.^[31]^[32]

Nitrogen also dissolves in the bloodstream and body fats. Rapid decompression (particularly in the case of divers ascending too quickly, or astronauts decompressing too quickly from cabin pressure to spacesuit pressure) can lead to a potentially fatal condition called decompression sickness (formerly known as caisson sickness or more commonly, the "bends"), when nitrogen bubbles form in the bloodstream, nerves, joints, and other sensitive or vital areas.^[33]^[34] Other "inert" gases (those gases other than carbon dioxide and oxygen) cause the same effects from bubbles composed of them, so replacement of nitrogen in breathing gases may prevent nitrogen narcosis, but does not prevent decompression sickness.^[35]

Direct skin contact with liquid nitrogen will eventually cause severe frostbite (cryogenic burns). This may happen almost instantly on contact, depending on the form of liquid nitrogen. Bulk liquid nitrogen causes less rapid freezing than a spray of nitrogen mist (such as is used to freeze certain skin growths in the practice of dermatology). The extra surface area provided by nitrogen-soaked materials is also important, with soaked clothing or cotton causing far more rapid damage than a spill of direct liquid to skin. Full "contact" between naked skin and large droplets or pools of undisturbed liquid nitrogen may be prevented for a few seconds by a layer of insulating gas from the Leidenfrost effect. However, liquid nitrogen applied to skin in mists, and on fabrics, bypasses this effect.

[edit] See also

[edit] References

^ Lavoisier, Antoine Laurent (1965). Elements of chemistry, in a new systematic order: containing all the modern discoveries. Courier Dover Publications. p. 15. ISBN 0486646246. http://books.google.com/?id=yS_m3PrVbpgC&pg=PR15.
^ Elements of Chemistry, trans. Robert Kerr (Edinburgh, 1790; New York: Dover, 1965), 52.
^ Lord Rayleigh's Active Nitrogen
^ "A new molecule and a new signature ' Chemistry ' tetranitrogen". Science News. February 16, 2002. http://www.findarticles.com/p/articles/mi_m1200/is_7_161/ai_83477565. Retrieved 2007-08-18.
^ "Polymeric nitrogen synthesized". physorg.com. 2004-08-05. http://www.physorg.com/news693.html. Retrieved 2009-06-22.
^ Karl Heinz Neeb (1997). The Radiochemistry of Nuclear Power Plants with Light Water Reactors. Berlin-New York: Walter de Gruyter. ISBN 3110132427.
^ ^a ^b Schrock, R. R. (2005). "Catalytic Reduction of Dinitrogen to Ammonia at a Single Molybdenum Center". Acc. Chem. Res. 38 (12): 955'962. doi:10.1021/ar0501121. PMID 16359167.
^ Fryzuk, M. D. and Johnson, S. A. (2000). "The continuing story of dinitrogen activation". Coordination Chemistry Reviews 200'202: 379. doi:10.1016/S0010-8545(00)00264-2.
^ Croswell, Ken (February 1996). Alchemy of the Heavens. Anchor. ISBN 0-385-47214-5. http://kencroswell.com/alchemy.html.
^ Meyer, Daved M.; Cardelli, Jason A.; Sofia, Ulysses J. (1997). "Abundance of Interstellar Nitrogen". arXiv. http://arxiv.org/abs/astro-ph/9710162v1. Retrieved 2007-12-24.
^ Hamilton, Calvin J.. "Titan (Saturn VI)". Solarviews.com. http://www.solarviews.com/eng/titan.htm. Retrieved 2007-12-24.
^ Reich, Murray; Kapenekas, Harry (1957). "Nitrogen Purfication. Pilot Plant Removal of Oxygen". Industrial & Engineering Chemistry 49: 869. doi:10.1021/ie50569a032.
^ Harding, Charlie, ed (2002). Elements of the p Block. Cambridge: Royal Society of Chemistry. ISBN 9780854046904. http://books.google.com/?id=W0HW8wgmQQsC&pg=PA90.
^ Gavriliuk, V. G.; Berns, Hans (1999). High nitrogen steels: structure, properties, manufacture, applications. Springer. ISBN 3540664114. http://books.google.com/?id=6eF4AfEwF4YC&pg=PA338.
^ "Why don't they use normal air in race car tires?". Howstuffworks. http://auto.howstuffworks.com/question594.htm. Retrieved 2006-07-22.
^ "Diffusion, moisture and tyre expansion". Car Talk. http://www.cartalk.com/content/columns/Archive/1997/September/05.html. Retrieved 2006-07-22.
^ "Is it better to fill your tires with nitrogen instead of air?". The Straight Dope. http://www.straightdope.com/columns/070216.html. Retrieved 2007-02-16.
^ Kemmochi, Y; Tsutsumi, K; Arikawa, A; Nakazawa, H (2002). "Centrifugal concentrator for the substitution of nitrogen blow-down micro-concentration in dioxin/polychlorinated biphenyl sample preparation". Journal of Chromatography A 943 (2): 295. doi:10.1016/S0021-9673(01)01466-2. PMID 11833649.
^ "How does the widget in a beer can work?". Howstuffworks. http://recipes.howstuffworks.com/question446.htm.
^ http://www.bomby.cz/plneni.htm
^ Kaganer, M. G. et al. (1967). "Vessels for the storage and transport of liquid oxygen and nitrogen". Chemical and Petroleum Engineering 3 (12): 918. doi:10.1007/BF01136404.
^ PubMed:11359389
^ Kent, Allen; Williams, James G. (1994). Encyclopedia of Computer Science and Technology, Volume 30. CRC Press. p. 318. ISBN 0824722833. http://books.google.com/?id=9fYFiZ7QeEcC&pg=PA318.
^ ^a ^b Rakov, Vladimir A.; Uman, Martin A. (2007). Lightning: Physics and Effects. Cambridge University Press. p. 508. ISBN 9780521035415. http://books.google.com/?id=TuMa5lAa3RAC&pg=PA508.
^ ^a ^b ^c ^d Bothe, Hermann; Ferguson, Stuart John; Newton, William Edward (2007). Biology of the nitrogen cycle. Elsevier. p. 283. ISBN 0444528571. http://books.google.com/?id=qmZDpnV-sYYC&pg=PA283.
^ Jahn, G.C. et al. (2005). "Effect of nitrogen fertilizer on the intrinsic rate of increase of the rusty plum aphid, Hysteroneura setariae (Thomas) (Homoptera: Aphididae) on rice (Oryza sativa L.)". Environmental Entomology 34 (4): 938'943. doi:10.1603/0046-225X-34.4.938. http://puck.esa.catchword.org/vl=33435372/cl=21/nw=1/rpsv/cw/esa/0046225x/v34n4/s26/p938.
^ ^a ^b Knox, G. A. (2007). Biology of the Southern Ocean. CRC Press. p. 392. ISBN 0849333946. http://books.google.com/?id=qPuCVvAv2GwC&pg=PA392.
^ Nielsen, Mk; Jørgensen, Bm (Jun 2004). "Quantitative relationship between trimethylamine oxide aldolase activity and formaldehyde accumulation in white muscle from gadiform fish during frozen storage.". Journal of agricultural and food chemistry 52 (12): 3814'22. doi:10.1021/jf035169l. ISSN 0021-8561. PMID 15186102.
^ Vickerstaff Joneja, Janice M. (2004). Digestion, diet, and disease: irritable bowel syndrome and gastrointestinal function. Rutgers University Press. p. 121. ISBN 0813533872. http://books.google.com/?id=RMPis9OnJxkC&pg=PT121.
^ "Biology Safety ' Cryogenic materials. The risks posed by them". University of Bath. http://www.bath.ac.uk/internal/bio-sci/bbsafe/asphyx.htm. Retrieved 2007-01-03.
^ Fowler, B; Ackles, KN; Porlier, G (1985). "Effects of inert gas narcosis on behavior'a critical review.". Undersea Biomed. Res. 12 (4): 369'402. PMID 4082343. http://archive.rubicon-foundation.org/3019. Retrieved 2008-09-21.
^ Rogers, W. H.; Moeller, G. (1989). "Effect of brief, repeated hyperbaric exposures on susceptibility to nitrogen narcosis". Undersea Biomed. Res. 16 (3): 227'32. ISSN 0093-5387. OCLC 2068005. PMID 2741255. http://archive.rubicon-foundation.org/2522. Retrieved 2008-09-21.
^ Acott, C. (1999). "A brief history of diving and decompression illness.". South Pacific Underwater Medicine Society journal 29 (2). ISSN 0813-1988. OCLC 16986801. http://archive.rubicon-foundation.org/6004. Retrieved 2008-09-21.
^ Kindwall, E. P.; A. Baz; E. N. Lightfoot; E. H. Lanphier; A. Seireg. (1975). "Nitrogen elimination in man during decompression.". Undersea Biomed. Res. 2 (4): 285'97. ISSN 0093-5387. OCLC 2068005. PMID 1226586. http://archive.rubicon-foundation.org/2741. Retrieved 2008-09-21.
^ US Navy Diving Manual, 6th revision. United States: US Naval Sea Systems Command. 2006. http://www.supsalv.org/00c3_publications.asp?destPage=00c3&pageID=3.9. Retrieved 2008-04-24.

[edit] Further reading

Garrett, Reginald H.; Grisham, Charles M. (1999). Biochemistry (2nd ed.). Fort Worth: Saunders College Publ.. ISBN 0030223180.
Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. ISBN 0080220576.
"Nitrogen". Los Alamos National Laboratory. 2003-10-20. http://periodic.lanl.gov/elements/7.html.

[edit] External links

Wikimedia Commons has media related to: Nitrogen

Look up nitrogen in Wiktionary, the free dictionary.

v ' d ' e Diatomic chemical elements

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

E numbers

Colors (E100'199) – Preservatives (E200'299) – Antioxidants & acidity regulators (E300'399) – Thickeners, stabilisers & emulsifiers (E400'499) – pH regulators & anti-caking agents (E500'599) – Flavour enhancers (E600'699) – Miscellaneous (E900'999) – Additional chemicals (E1100'1599)

Waxes (E900'909) – Synthetic glazes (E910'919) – Improving agents (E920'929) – Packaging gases (E930'949) – Sweeteners (E950'969) – Foaming agents (E990'999)

Calcium peroxide (E930) – Argon (E938) – Helium (E939) – Dichlorodifluoromethane (E940) – Nitrogen (E941) – Nitrous oxide (E942) – Butane (E943a) – Isobutane (E943b) – Propane (E944) – Oxygen (E948) – Hydrogen (E949)

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