Nitrogen
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2008-07-15T09:05:01Z
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{{Infobox nitrogen}}
'''Nitrogen''' ({{pronEng|ˈnaɪtɹəʤɪn}}) is a [[chemical element]] that has the symbol '''N''' and [[atomic number]] 7 and [[atomic weight]] 14.0067. Elemental nitrogen is a colorless, odorless, tasteless and mostly [[inert]] [[diatomic]] gas at [[standard conditions]], constituting 78.08% by volume of [[Earth's atmosphere]].
Many industrially important compounds, such as [[ammonia]], [[nitric acid]], organic nitrates ([[propellant]]s and [[explosive]]s), and [[cyanide]]s, contain nitrogen. The very strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in converting the {{chem|N|2}} into useful compounds, and releasing large amounts of energy when these compounds burn or decay back into nitrogen gas.
The element nitrogen was discovered by [[Daniel Rutherford]] in [[1772]]. Nitrogen occurs in all living organisms — it is a constituent element of [[amino acids]] and thus of [[protein]]s, and of [[nucleic acids]] ([[DNA]] and [[RNA]]); resides in the [[chemical structure]] of almost all [[neurotransmitter]]s; and is a defining component of [[alkaloid]]s, biological molecules produced by many organisms.
== Properties ==
Nitrogen is a [[nonmetal]], with an [[electronegativity]] of 3.0. It has five [[electrons]] in its outer shell and is therefore [[valence (chemistry)|trivalent]] in most compounds. The triple bond in molecular nitrogen ({{chem|N|2}}) is one of the strongest in nature. The resulting difficulty of converting ({{chem|N|2}}) into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental {{chem|N|2}}, have dominated the role of nitrogen in both nature and human economic activities.
At [[atmospheric pressure]] molecular nitrogen [[condensation|condenses]] ([[liquid|liquifies]]) at 77 [[Kelvin|K]] (−195.8 °[[Celsius|C]]) and [[freezing|freezes]] at 63 K (−210.0 °C) into the beta [[Close-packing|hexagonal close-packed]] crystal [[allotropic]] form. Below 35.4 K (−237.6 °C) nitrogen assumes the alpha [[Cubic crystal system|cubic]] crystal allotropic form. [[Liquid nitrogen]], a fluid resembling water, but with 80.8% of the density, is a common [[cryogen]].
Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like {{chem|N|3}} and [[tetranitrogen|{{chem|N|4}}]].<ref>{{cite web |url=http://www.findarticles.com/p/articles/mi_m1200/is_7_161/ai_83477565 |title=A new molecule and a new signature - Chemistry - tetranitrogen |work=Science News |date=[[February 16]][[2002]] |accessdate=2007-08-18}}</ref> Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced under [[diamond anvil]] conditions, nitrogen polymerizes into the single bonded [[diamond]] crystal structure, an allotrope nicknamed "nitrogen diamond."<ref>{{cite web |url=http://www.physorg.com/news693.html |title=Polymeric nitrogen synthesized |publisher=physorg.com |date=[[August 5]][[2004]] |accessdate=2007-08-18}}</ref>
== Occurrence ==
Nitrogen is the largest single constituent of the [[Earth|Earth's]] [[Earth's atmosphere|atmosphere]] (78.082% by volume of dry air, 75.3% by weight in dry air). It is created by [[stellar nucleosynthesis|fusion]] processes in [[star]]s, and is estimated to be the 7th most abundant [[chemical element]] by mass in the universe. {{Fact|date=February 2008}}
[[Molecule|Molecular]] nitrogen and nitrogen [[compound]]s have been detected in [[interstellar space]] by astronomers using the [[Far Ultraviolet Spectroscopic Explorer]].<ref>{{cite web|url=http://arxiv.org/abs/astro-ph/9710162v1|title=Abundance of Interstellar Nitorgen|author=Daved M. Meyer, Jason A. Cardelli, and Ulysses J. Sofia|publisher=arXiv|accessdate=2007-12-24|date=1997}}</ref> Molecular nitrogen is a major constituent of the [[Saturn]]ian moon [[Titan (moon)|Titan]]'s thick atmosphere, and occurs in trace amounts in other planetary atmospheres.<ref>{{cite web|url=http://www.solarviews.com/eng/titan.htm|title=Titan (Saturn VI)|accessdate=2007-12-24|publisher=Solarviews.com|author=Calvin J. Hamilton}}</ref>
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 are unable to [[nitrogen fixation|fix atmospheric nitrogen]].
Nitrogen occurs naturally in a number of minerals, such as [[potassium nitrate|saltpetre]] (potassium nitrate), [[sodium nitrate|Chile saltpetre]] (sodium nitrate) and [[sal ammoniac]] (ammonium chloride). Most of these are relatively uncommon, partly because of the minerals' ready solubility in water. See also [[:category:Nitrate minerals|Nitrate minerals]] and [[:category:Ammonium minerals|Ammonium minerals]].
==Isotopes==
{{seealso|Isotopes of nitrogen}}
There are two stable [[isotope]]s of nitrogen: <sup>14</sup>N and <sup>15</sup>N. By far the most common is <sup>14</sup>N (99.634%), which is produced in the [[CNO cycle]] in [[star]]s. Of the ten isotopes produced synthetically, <sup>13</sup>N has a [[half life]] of ten minutes and the remaining isotopes have half lives on the order of seconds or less.
Biologically-mediated reactions (e.g., [[Assimilation (biology)|assimilation]], [[nitrification]], and [[denitrification]]) strongly control nitrogen dynamics in the soil. These reactions typically result in <sup>15</sup>N enrichment of the [[Substrate (chemistry)|substrate]] and depletion of the [[Product (chemistry)|product]].
0.73% of the molecular nitrogen in Earth's atmosphere is comprised of the [[isotopologue]] <sup>14</sup>N<sup>15</sup>N and almost all the rest is <sup>14</sup>N<sub>2</sub>.
== Electromagnetic spectrum ==
Molecular nitrogen (<sup>14</sup>N<sub>2</sub>) 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 as well as in the atmospheres of other planetary bodies. For similar reasons, pure molecular [[nitrogen laser]]s 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 (astronomy)|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).
== History ==
Nitrogen ([[Latin]] ''nitrogenium'', where ''nitrum'' (from [[Greek language|Greek]] ''nitron'') means "saltpetre" (see [[niter]]), and ''genes'' means "forming") is formally considered to have been discovered by [[Daniel Rutherford]] in [[1772]], who called it ''noxious air'' or ''fixed air''. 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 ''[[Phlogiston theory|phlogisticated air]]''. Nitrogen gas was [[inert]] enough that [[Antoine Lavoisier]] referred to it as ''azote'', from the [[Greek language|Greek]] word αζωτος meaning "lifeless". Animals died in it, and it was the principal component of air in which animals had suffocated and flames had burned to extinction. This term has become the [[French language|French]] word for "nitrogen" and later spread out to many other languages.
[[Argon]] was discovered when it was noticed that nitrogen from air is not identical to nitrogen from chemical reactions.
Compounds of nitrogen were known in the [[Middle Ages]]. The [[alchemy|alchemists]] knew [[nitric acid]] as ''aqua fortis'' (strong water). The mixture of nitric and [[hydrochloric acid]]s was known as ''[[aqua regia]]'' (royal water), celebrated for its ability to dissolve [[gold]] (the ''king'' of metals). The earliest industrial and [[Agriculture|agricultural]] applications of nitrogen compounds involved uses in the form of [[Potassium nitrate|saltpeter]] ([[sodium nitrate|sodium-]] or [[potassium nitrate]]), notably in [[gunpowder]], and much later, as [[fertilizer]].
==Biological role==
{{See also|nitrogen cycle}}
Nitrogen is an essential part of [[amino acids]] and [[nucleic acid]]s, both of which are essential to all life on Earth.
Molecular nitrogen in the atmosphere cannot be used directly by either plants or animals, and needs to be converted into nitrogen compounds, or "fixed," in order to be used by life. [[Precipitation (meteorology)|Precipitation]] often contains substantial quantities of [[ammonium]] and nitrate, both thought to be a result of [[nitrogen fixation]] by [[lightning]] and other atmospheric electric phenomena. However, because [[ammonium]] is preferentially retained by the [[forest canopy]] relative to atmospheric nitrate, most of the fixed nitrogen that reaches the [[soil]] surface under trees is in the form of nitrate. Soil nitrate is preferentially assimilated by tree [[root]]s relative to soil ammonium.
Specific [[bacteria]] (e.g. ''[[Rhizobia|Rhizobium]] trifolium'') possess [[nitrogenase]] [[enzyme]]s which can fix atmospheric nitrogen (see [[nitrogen fixation]]) into a form (ammonium ion) which is chemically useful to higher organisms. This process requires a large amount of energy and [[anoxia|anoxic]] conditions. Such bacteria may be free in the soil (e.g. ''[[Azotobacter]]'') but normally exist in a [[symbiosis|symbiotic]] relationship in the [[root nodule]]s of leguminous plants (e.g. [[clover]], ''[[Trifolium]]'' species, or the [[soya bean]] plant, ''[[Glycine max]]''). Nitrogen-fixing bacteria can be symbiotic with a number of unrelated plant species. Common examples are legumes, alders (''[[Alnus]]'') spp., lichens, ''[[Casuarina]]'', ''[[Myrica]]'', liverworts, and ''[[Gunnera]]''.
As part of the symbiotic relationship, the plant subsequently converts the ammonium ion to nitrogen oxides and amino acids to form [[protein]]s and other biologically useful molecules, such as [[alkaloids]]. In return for the usable (fixed) nitrogen, the plant secretes sugars to the symbiotic bacteria.
Some 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.
Nitrogen compounds are basic building blocks in animal biology. 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]]. Some plant-feeding insects are so dependent on nitrogen in their diet, that varying the amount of nitrogen fertilizer applied to a plant can affect the rate of reproduction of the insects feeding on it.<ref>{{cite journal |author=Jahn, GC, LP Almazan, and J Pacia |year=2005 |title=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.) |journal=Environmental Entomology |volume=34 |issue=4 |pages=938–943 |url=http://puck.esa.catchword.org/vl=33435372/cl=21/nw=1/rpsv/cw/esa/0046225x/v34n4/s26/p938}}</ref>
Soluble nitrate is an important limiting factor in the growth of certain bacteria in ocean waters. In many places in the world, artificial [[fertilizer]]s 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 (ecology)|"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 [[osmosis|osmotic]] effects of their environment (conversion of this compound to [[dimethylamine]] is responsible for the early odor in unfresh saltwater fish: PMID 15186102). In animals, the [[free radical]] molecule [[nitric oxide]] (NO), which is derived from an [[amino acid]], serves as an important regulatory molecule for circulation.
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 nitrogen-containing long-chain [[amines]], such as [[putrescine]] and [[cadaverine]].
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.
== Reactions==
[[Image:RuA5N2.png|thumb|right|Structure of [Ru(NH<sub>3</sub>)<sub>5</sub>(N<sub>2</sub>)]<sup>2+</sup>.]]
Nitrogen is generally unreactive at standard temperature and pressure. N<sub>2</sub> 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 at [[Standard conditions for temperature and pressure|STP]].<ref name=Schrock>{{cite journal | author = [[Richard R. Schrock]] | title = Catalytic Reduction of Dinitrogen to Ammonia at a Single Molybdenum Center | journal = [[Acc. Chem. Res.]] | year = 2005 | volume = 38 | pages = 955–962 | doi =10.1021/ar0501121}}</ref> Lithium burns in an atmosphere of N<sub>2</sub> to give [[lithium nitride]]:
:6 Li + N<sub>2</sub> → 2 Li<sub>3</sub>N
Magnesium also burns in nitrogen, forming [[magnesium nitride]].
:3 Mg + N<sub>2</sub> → Mg<sub>3</sub>N<sub>2</sub>
N<sub>2</sub> forms a variety of [[adduct]]s with transition metals. The first example of a [[dinitrogen complex]] is [Ru(NH<sub>3</sub>)<sub>5</sub>(N<sub>2</sub>)]<sup>2+</sup> (see figure at right). Such compounds are now numerous, other examples include IrCl(N<sub>2</sub>)(PPh<sub>3</sub>)<sub>2</sub>, W(N<sub>2</sub>)<sub>2</sub>([[Dppe|Ph<sub>2</sub>CH<sub>2</sub>CH<sub>2</sub>PPh<sub>2</sub>]])<sub>2</sub>, and [(η<sup>5</sup>-C<sub>5</sub>Me<sub>4</sub>H)<sub>2</sub>Zr]<sub>2</sub>([[bridging ligand|μ]]<sub>2</sub>,[[hapticity|η]]²,η²-N<sub>2</sub>). These complexes illustrate how N<sub>2</sub> might bind to the metal(s) in [[nitrogenase]] and the catalyst for the [[Haber-Bosch Process]].<ref>{{cite journal |author=Fryzuk, M. D. and Johnson, S. A. |title=The continuing story of dinitrogen activation |journal=[[Coordination Chemistry Reviews]] |volume=200–202 |pages=379 |year=2000 |doi=10.1016/S0010-8545(00)00264-2}}</ref> A catalytic process to reduce N<sub>2</sub> to ammonia with the use of a [[molybdenum]] complex in the presence of a proton source was published in 2005.<ref name=Schrock/> (see [[nitrogen fixation]])
The starting point for industrial production of nitrogen compounds is the [[Haber-Bosch]] process, in which nitrogen is fixed by reacting {{chem|N|2}} and {{chem|H|2}} over a [[ferric oxide]] ({{chem|Fe|3|O|4}}) 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 catalysed 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).
== Nitrogen compounds in industry ==
===Simple compounds ===
''See also the category [[:category:Nitrogen compounds|Nitrogen compounds]].''
The main neutral [[hydride]] of nitrogen is [[ammonia]] (N[[hydrogen|H]]<sub>3</sub>), although [[hydrazine]] (N<sub>2</sub>H<sub>4</sub>) is also commonly used. Ammonia is more [[Basic (chemistry)|basic]] than [[water]] by 6 orders of magnitude. In [[solution]] ammonia forms the [[ammonium]] [[ion]] (NH<sub>4</sub><sup>+</sup>). Liquid ammonia (b.p. 240 K) is [[amphiprotic]] (displaying either [[Brønsted-Lowry]] acidic or basic character) and forms ammonium and the less common [[amide]] ions (NH<sub>2</sub><sup>-</sup>); both amides and [[nitride]] (N<sup>3-</sup>) [[salt]]s are known, but [[Chemical decomposition|decompose]] in water. Singly, doubly, triply and quadruply substituted alkyl compounds of ammonia are called [[amine]]s (four substitutions, to form commercially and biologically important quarternary amines, results in a positively charged nitrogen, and thus a water-soluble, or at least [[amphiphilic]], compound). Larger chains, rings and structures of nitrogen hydrides are also known, but are generally unstable. N<sub>2</sub><sup>2+</sup> is another polyatomic cation as in hydrazine.
Other classes of nitrogen [[anion]]s (negatively charged ions) are the poisonous [[azide]]s (N<sub>3</sub><sup>-</sup>), which are linear and [[isoelectronic]] to [[carbon dioxide]], but which bind to important iron-containing enzymes in the body in a manner more resembling [[cyanide]]. Another [[molecule]] of the same structure is the colorless and relatively inert anesthetic gas [[Nitrous oxide|dinitrogen monoxide]] {{chem|N|2|O}}, also known as laughing gas. This is one of a variety of [[oxide]]s, the most prominent of which are [[nitrogen monoxide]] (NO) (known more commonly as [[nitric oxide]] in biology), a natural free radical molecule used by the body as a signal for short-term control of smooth muscle in the circulation. Another notable nitrogen oxide compound (a family often abbreviated '''NOx''') is the reddish and poisonous [[nitrogen dioxide]] {{chem|NO|2}}, which also contains an unpaired [[electron]] and is an important component of [[smog]]. Nitrogen molecules containing unpaired electrons show an understandable tendency to [[dimerize]] (thus pairing the electrons), and are generally highly reactive.
The more standard oxides, [[dinitrogen trioxide]] {{chem|N|2|O|3}} and [[dinitrogen pentoxide]] {{chem|N|2|O|5}}, are actually fairly unstable and explosive-- a tendency which is driven by the stability of {{chem|N|2}} as a product. The corresponding acids are [[nitrous acid|nitrous]] {{chem|HNO|2}} and [[nitric acid]] {{chem|HNO|3}}, with the corresponding salts called [[nitrite]]s and [[nitrate]]s. [[Dinitrogen tetroxide]] {{chem|N|2|O|4}} (DTO) is one of the most important oxidisers of rocket fuels, used to oxidise [[hydrazine]] in the [[Titan rocket]] and in the recent [[NASA]] [[MESSENGER]] probe to [[Mercury (planet)|Mercury]]. DTO is an intermediate in the manufacture of nitric acid {{chem|HNO|3}}, one of the few acids stronger than [[hydronium]] and a fairly strong [[oxidizing agent]].
Nitrogen is notable for the range of explosively unstable compounds that it can produce. Nitrogen triiodide {{chem|NI|3}} is an extremely sensitive [[contact explosive]]. [[Nitrocellulose]], produced by nitration of cellulose with nitric acid, is also known as guncotton. [[Nitroglycerin]], made by nitration of [[glycerin]], is the dangerously unstable explosive ingredient of [[dynamite]]. The comparatively stable, but more powerful explosive [[trinitrotoluene]] (TNT) is the standard explosive against which the power of nuclear explosions are measured.
Nitrogen can also be found in [[organic compound]]s. Common nitrogen [[functional group]]s include: [[amines]], [[amides]], [[nitro]] groups, [[imine]]s, and [[enamine]]s. The amount of nitrogen in a [[chemical substance]] can be determined by the [[Kjeldahl method]].
===Nitrogen compounds of notable economic importance===
Molecular nitrogen (N<sub>2</sub>) in the atmosphere is relatively non-reactive due to its strong bond, and N<sub>2</sub> plays an inert role in the human body, being neither produced or destroyed. In nature, nitrogen is converted into biologically (and industrially) useful compounds by some living organisms, notably certain [[bacteria]] (i.e. [[nitrogen fixing bacteria]] – see ''[[#Biological role|Biological role]]'' above). Molecular nitrogen is also 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 [[salt]]s of nitric acid have been important historically as stores of chemical energy. They include important compounds such as [[potassium nitrate]] (or [[Potassium nitrate|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 [[rocket]]s. [[Hydrazine]] and hydrazine derivatives find use as rocket [[fuel]]s and [[monopropellant]]s. 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 (N<sub>2</sub>) which is a product of the compounds' thermal decomposition. When nitrates burn or explode, the formation of the powerful triple bond in the N<sub>2</sub> which results, 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]] (N<sub>2</sub>O) 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 defences of plants against predation). Nitrogen containing drugs include all of the major classes of antibiotics, and organic nitrate drugs like [[nitroglycerin]] and [[nitroprusside]] which regulate blood pressure and heart action by mimicking the action of [[nitric oxide]].
== Dangers ==
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.<ref>{{cite web |url=http://www.bath.ac.uk/internal/bio-sci/bbsafe/asphyx.htm |title=Biology Safety - Cryogenic materials. The risks posed by them |publisher=University of Bath |accessdate=2007-01-03}}</ref> 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]].
Nitrogen also dissolves in the [[Cardiovascular system|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.
Direct skin contact with [[liquid nitrogen]] causes severe [[frostbite]] (cryogenic burns) within seconds, though not instantly on contact, depending on form of liquid nitrogen (liquid vs. mist) and surface area of the nitrogen-soaked material (soaked clothing or cotton causing more rapid damage than a spill of direct liquid to skin, which for a few seconds is protected by the [[Leidenfrost effect]]).
== See also ==
* [[Nutrient]]
* [[Nitrogenomics]]
* [[Nitrogen oxide]]
* [[TKN]]
* [[Tetranitrogen]]
==References==
{{Reflist|2}}
==Further reading==
*{{cite book |title=Biochemistry |last=Garrett |first=Reginald H. |authorlink= |coauthors=Grisham, Charles M. |year=1999 |edition=2nd edition |publisher=Saunders College Publ. |location=Fort Worth |isbn=0030223180 |pages= }}
*{{cite book |title=Chemistry of the Elements |last=Greenwood |first=Norman N. |authorlink= |coauthors=Earnshaw, Alan |year=1984 |publisher=Pergamon Press |location=Oxford |isbn=0080220576 |pages= }}
*{{cite web |url=http://periodic.lanl.gov/elements/7.html |title=Nitrogen |accessdate= |work= |publisher=Los Alamos National Laboratory |date=2003-10-20 }}
== External links ==
{{Commons|Nitrogen}}
{{wiktionary|nitrogen}}
* [http://www.balashon.com/2008/07/neter-and-nitrogen.html Etymology of Nitrogen]
* [http://www.newton.dep.anl.gov/askasci/chem99/chem99306.htm Why high nitrogen density in explosives?]
* [http://www.webelements.com/nitrogen/ WebElements.com – Nitrogen]
* [http://education.jlab.org/itselemental/ele007.html It's Elemental – Nitrogen]
* [http://www.sunysccc.edu/academic/mst/ptable/n.html Schenectady County Community College – Nitrogen]
* [http://www.uigi.com/nitrogen.html Nitrogen N2 Properties, Uses, Applications]
* [http://box27.bluehost.com/~edsanvil/wiki/index.php?title=Nitrogen_gas Computational Chemistry Wiki]
* [http://www.2spi.com/catalog/instruments/nitrodew-supp.html Handling procedures for liquid nitrogen]
* [http://www.safety.vanderbilt.edu/pdf/hcs_msds/NitrogenCryo_G103_06_04.pdf Material Safety Data Sheet]
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[[ru:Азот]]
[[sq:Azoti]]
[[scn:Azzotu]]
[[simple:Nitrogen]]
[[sk:Dusík]]
[[sl:Dušik]]
[[sr:Азот]]
[[sh:Dušik]]
[[fi:Typpi]]
[[sv:Kväve]]
[[ta:நைட்ரஜன்]]
[[te:నత్రజని]]
[[th:ไนโตรเจน]]
[[vi:Nitơ]]
[[tg:Азот]]
[[tr:Azot]]
[[uk:Азот]]
[[zh-yue:氮]]
[[bat-smg:Azuots]]
[[zh:氮]]
[[mt:Ażotu]]