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{{Use British English|date=October 2024}}
{{Infobox nitrogen|engvar=en-GB}}
'''Nitrogen''' is a [[chemical element]]; it has [[Symbol (chemistry)|symbol]] '''N''' and [[atomic number]] 7. Nitrogen is a [[Nonmetal (chemistry)|nonmetal]] and the lightest member of [[pnictogen|group 15]] of the [[periodic table]], often called the [[Pnictogen|pnictogens]]. It is a common element in the [[universe]], estimated at [[Abundance of the chemical elements|seventh in total abundance]] in the [[Milky Way]] and the [[Solar System]]. At [[standard temperature and pressure]], two atoms of the element [[chemical bond|bond]] to form N<sub>2</sub>, a colourless and odourless [[diatomic molecule|diatomic]] [[gas]]. N<sub>2</sub> forms about 78% of [[Atmosphere of Earth|Earth's atmosphere]], making it the most abundant chemical species in air. Because of the volatility of nitrogen compounds, nitrogen is relatively rare in the solid parts of the Earth.
It was first discovered and isolated by Scottish physician [[Daniel Rutherford]] in 1772 and independently by [[Carl Wilhelm Scheele]] and [[Henry Cavendish]] at about the same time. The name {{Lang|fr|nitrogène}} was suggested by French chemist [[Jean-Antoine-Claude Chaptal]] in 1790 when it was found that nitrogen was present in [[nitric acid]] and [[nitrate]]s. [[Antoine Lavoisier]] suggested instead the name ''azote'', from the {{langx|grc|ἀζωτικός}} "no life", as it is an [[asphyxiant gas]]; this name is used in a number of languages, and appears in the English names of some nitrogen compounds such as [[hydrazine]], [[azide]]s and [[azo compound]]s.
Elemental nitrogen is usually produced from air by [[pressure swing adsorption]] technology. About 2/3 of commercially produced elemental nitrogen is used as an inert (oxygen-free) gas for commercial uses such as food packaging, and much of the rest is used as [[liquid nitrogen]] in [[cryogenic]] applications. Many industrially important compounds, such as [[ammonia]], nitric acid, organic nitrates ([[propellant]]s and [[explosive]]s), and [[cyanide]]s, contain nitrogen. The extremely strong [[triple bond]] in elemental nitrogen (N≡N), the second strongest bond in any [[diatomic molecule]] after [[carbon monoxide]] (CO),<ref>[http://www.wiredchemist.com/chemistry/data/bond_energies_lengths.html Common Bond Energies (D) and Bond Lengths (r)] {{Webarchive|url=https://web.archive.org/web/20100515215439/http://www.wiredchemist.com/chemistry/data/bond_energies_lengths.html |date=2010-05-15 }}. wiredchemist.com</ref> dominates nitrogen chemistry. This causes difficulty for both organisms and industry in converting N<sub>2</sub> into useful [[Chemical compound|compounds]], but at the same time it means that burning, exploding, or decomposing nitrogen compounds to form nitrogen gas releases large amounts of often useful energy. Synthetically produced ammonia and nitrates are key industrial [[fertiliser]]s, and fertiliser nitrates are key [[pollutant]]s in the [[eutrophication]] of water systems.  Apart from its use in fertilisers and energy stores, nitrogen is a constituent of organic compounds as diverse as [[aramid]]s used in high-strength fabric and [[cyanoacrylate]] used in [[superglue]].
Nitrogen occurs in all organisms, primarily in [[amino acid]]s (and thus [[protein]]s), in the [[nucleic acid]]s ([[DNA]] and [[RNA]]) and in the energy transfer molecule [[adenosine triphosphate]]. The [[Composition of the human body|human body contains]] about 3% nitrogen by mass, the fourth most abundant element in the body after oxygen, carbon, and hydrogen. The [[nitrogen cycle]] describes the movement of the element from the air, into the [[biosphere]] and organic compounds, then back into the atmosphere. Nitrogen is a constituent of every major pharmacological drug class, including [[antibiotic]]s. Many drugs are mimics or [[prodrug]]s of natural nitrogen-containing [[Cell signaling|signal molecules]]: for example, the organic nitrates [[nitroglycerin]] and [[nitroprusside]] control [[blood pressure]] by metabolising into [[nitric oxide]]. Many notable nitrogen-containing drugs, such as the natural [[caffeine]] and [[morphine]] or the synthetic [[amphetamine]]s, act on receptors of animal [[neurotransmitter]]s.
==History==
[[File:Rutherford Daniel.jpg|thumb|left|[[Daniel Rutherford]], discoverer of nitrogen]]
Nitrogen compounds have a very long history, [[ammonium chloride]] having been known to [[Herodotus]]. They were well-known by the Middle Ages. [[Alchemy|Alchemists]] knew nitric acid as ''[[aqua fortis]]'' (strong water), as well as other nitrogen compounds such as [[ammonium]] salts and [[nitrate]] salts. 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.<ref name="Greenwood406" />
The discovery of nitrogen is attributed to the Scottish physician [[Daniel Rutherford]] in 1772, who called it ''noxious air''.<ref>Rutherford, Daniel (1772) "[https://books.google.com/books?id=JxUUAAAAQAAJ Dissertatio Inauguralis de aere fixo, aut mephitico] {{Webarchive|url=https://web.archive.org/web/20200806192616/https://books.google.com/books?id=JxUUAAAAQAAJ&printsec=frontcover&hl=en&source=gbs_ge_summary_r&cad=0 |date=2020-08-06 }}" (Inaugural dissertation on the air [called] fixed or mephitic), M.D. dissertation, University of Edinburgh, Scotland. English translation: {{cite journal | first = Leonard | last = Dobbin | title = Daniel Rutherford's inaugural dissertation |year = 1935 | journal = Journal of Chemical Education | volume = 12 | issue = 8 | pages = 370–75 |doi = 10.1021/ed012p370| bibcode = 1935JChEd..12..370D }}</ref><ref name="Weeks">{{cite journal | doi = 10.1021/ed009p215 | title = The discovery of the elements. IV. Three important gases |year = 1932 | last1 = Weeks | first1 = Mary Elvira |author-link1=Mary Elvira Weeks| journal = Journal of Chemical Education | volume = 9 | issue = 2 | page = 215|bibcode = 1932JChEd...9..215W }}</ref> Though he did not recognise it as an entirely different chemical substance, he clearly distinguished it from Joseph Black's [[Joseph Black#Carbon dioxide|"fixed air"]], or carbon dioxide.<ref>{{cite book |first=Aaron J. |last=Ihde |title=The Development of Modern Chemistry |location=New York | publisher=Harper & Row |year=1964}}</ref> The fact that there was a component of air that does not support [[combustion]] was clear to Rutherford, although he was not aware that it was an element. Nitrogen was also studied at about the same time by [[Carl Wilhelm Scheele]],<ref>Carl Wilhelm Scheele, [https://archive.org/stream/CarlWilhelmSche00Sche#page/n1/mode/2up ''Chemische Abhandlung von der Luft und dem Feuer''] [Chemical treatise on air and fire] (Upsala, Sweden: Magnus Swederus, 1777; and Leipzig, (Germany): Siegfried Lebrecht Crusius, 1777). In the section titled "Die Luft muß aus elastischen Flüßigkeiten von zweyerley Art, zusammengesetzet seyn." (The air must be composed of elastic fluids of two sorts), pp. 6–14, Scheele presents the results of eight experiments in which air was reacted with various substances. He concluded ([https://archive.org/stream/CarlWilhelmSche00Sche#page/12/mode/2up p. 13]): ''"So viel sehe ich aus angeführten Versuchen, daß die Luft aus 2 von einander unterschiedenen Flußigkeiten bestehe, von welchen die eine die Eigenschaft das Phlogiston anzuziehen gar nicht äussere, die andere aber zur solchen Attraction eigentlich aufgeleget ist und welche zwischen dem 3:ten und 4:ten Theil von der ganzen Luftmasse aus machet."'' (So I see [this] much from the experiments [that were] conducted: that the air consists of two fluids [that] differ from one another, of which the one doesn't express at all the property of attracting phlogiston; the other, however, is capable of such attraction and which makes up between 1/3 and 1/4 part of the entire mass of the air.)</ref> [[Henry Cavendish]],<ref>{{cite journal | first = Joseph | last = Priestley | title = Observations on different kinds of air |year = 1772 | journal = Philosophical Transactions of the Royal Society of London | volume = 62 | pages = 147–256|doi =10.1098/rstl.1772.0021 | s2cid = 186210131 }} ; [https://books.google.com/books?id=aBxWAAAAYAAJ&pg=PA225 see p. 225.] {{Webarchive|url=https://web.archive.org/web/20160903135622/https://books.google.com/books?id=aBxWAAAAYAAJ&pg=PA225 |date=2016-09-03 }}</ref> and [[Joseph Priestley]],<ref>{{cite journal | first = Joseph | last = Priestley | title = Observations on different kinds of air |year = 1772 | journal = Philosophical Transactions of the Royal Society of London | volume = 62 | pages = 147–256 | doi=10.1098/rstl.1772.0021| s2cid = 186210131 }} ; see: [https://books.google.com/books?id=aBxWAAAAYAAJ&pg=PA225 "VII. Of air infected with the fumes of burning charcoal." pp. 225–28.] {{Webarchive|url=https://web.archive.org/web/20160903135622/https://books.google.com/books?id=aBxWAAAAYAAJ&pg=PA225 |date=2016-09-03 }}</ref> who referred to it as ''burnt air'' or ''[[Phlogiston theory|phlogisticated air]]''. French chemist [[Antoine Lavoisier]] referred to nitrogen gas as "[[wikt:mephitic air|mephitic air]]" or ''azote'', from the [[Greek language|Greek]] word {{lang|el|άζωτικός}} (azotikos), "no life", due to it being [[Asphyxiant gas|asphyxiant]].<ref>Lavoisier, Antoine with Robert Kerr, trans., ''Elements of Chemistry'', 4th ed. (Edinburgh, Scotland: William Creech, 1799), pp. 85–86. [https://books.google.com/books?id=adYKAAAAIAAJ&pg=PA85 [p. 85]:] {{Webarchive|url=https://web.archive.org/web/20200806230924/https://books.google.com/books?id=adYKAAAAIAAJ&pg=PA85 |date=2020-08-06 }} "In reflecting upon the circumstances of this experiment, we readily perceive; that the mercury, during its calcination [i.e., roasting in air], absorbs the salubrious and respirable part of the air, or, to speak more strictly, the base of this respirable part; that the remaining air is a species of mephitis [i.e., a poisonous gas emitted from the earth], incapable of supporting combustion or respiration; ... " [https://books.google.com/books?id=adYKAAAAIAAJ&pg=PA86 [p. 86]:] {{Webarchive|url=https://web.archive.org/web/20200806192501/https://books.google.com/books?id=adYKAAAAIAAJ&pg=PA86 |date=2020-08-06 }} "I shall afterwards shew, that at least in our climate, the atmospheric air is composed of respirable and mephitic airs, in the proportion of 27 and 73; ... "</ref><ref>Lavoisier, Antoine with Robert Kerr, trans., ''Elements of Chemistry'', 4th ed. (Edinburgh, Scotland: William Creech, 1799), [https://books.google.com/books?id=adYKAAAAIAAJ&pg=PA101 p. 101:] "The chemical properties of the noxious portion of the atmospheric air being hitherto but little known, we have been satisfied to derive the name of its base from its known quality of killing such animals as are forced to breathe it, giving it the name of ''azot'', from the Greek privitive particle α and ξωη, ''vita''; hence the name of the noxious part of atmospheric air is ''azotic gas''."</ref> In an atmosphere of pure nitrogen, animals died and flames were extinguished. Though Lavoisier's name was not accepted in English since it was pointed out that all gases but oxygen are either asphyxiant or outright toxic, it is used in many languages (French, Italian, Portuguese, Polish, Russian, Albanian, Turkish, etc.; the German ''Stickstoff'' similarly refers to the same characteristic, viz. ''ersticken'' "to choke or suffocate") and still remains in English in the common names of many nitrogen compounds, such as [[hydrazine]] and compounds of the [[azide]] ion. Finally, it led to the name "[[pnictogen]]s" for the group headed by nitrogen, from the Greek πνίγειν "to choke".<ref name="Greenwood406">Greenwood and Earnshaw, pp. 406–07</ref>
The English word nitrogen (1794) entered the language from the French {{Lang|fr|nitrogène}}, coined in 1790 by French chemist [[Jean-Antoine Chaptal]] (1756–1832),<ref>Chaptal, J. A. and Nicholson, William trans. (1800) ''Elements of Chemistry'', 3rd ed. London, England: C.C. and J. Robinson, vol. 1. [https://archive.org/stream/elementsofchemis01chap#page/n47/mode/2up pp. xxxv–xxxvi:] "In order to correct the Nomenclature on this head [i.e., in this regard], nothing more is necessary than to substitute to [i.e., for] this word a denomination which is derived from the general system made use of; and I have presumed to propose that of Nitrogene Gas. In the first place, it is deduced from the characteristic and exclusive property of this gas, which forms the radical of the nitric acid. By this means we shall preserve to the combinations [i.e., compounds] of this substance the received [i.e., prevailing] denominations, such as those of the Nitric Acid, Nitrates, Nitrites, &c."</ref> from the French ''nitre'' ([[potassium nitrate]], also called [[saltpetre]]) and the French suffix ''-gène'', "producing", from the [[Greek language|Greek]] -γενής (-genes, "begotten"). Chaptal's meaning was that nitrogen is the essential part of [[nitric acid]], which in turn was produced from [[nitre]]. In earlier times, nitre had been confused with Egyptian "natron" ([[sodium carbonate]])&nbsp;– called νίτρον (nitron) in Greek&nbsp;– which, despite the name, contained no nitrate.<ref>[http://www.etymonline.com/index.php?term=nitrogen nitrogen] {{Webarchive|url=https://web.archive.org/web/20170702134749/http://www.etymonline.com/index.php?term=nitrogen |date=2017-07-02 }}. Etymonline.com. Retrieved 2011-10-26.</ref>
The earliest military, industrial, and agricultural applications of nitrogen compounds used saltpetre ([[sodium nitrate]] or potassium nitrate), most notably in [[gunpowder]], and later as [[fertiliser]]. In 1910, [[John William Strutt, 3rd Baron Rayleigh|Lord Rayleigh]] discovered that an electrical discharge in nitrogen gas produced "active nitrogen", a [[monatomic]] [[allotrope]] of nitrogen.<ref>Strutt, R. J. (1911) [http://rspa.royalsocietypublishing.org/content/royprsa/85/577/219.full.pdf "Bakerian Lecture. A chemically active modification of nitrogen, produced by the electric discharge"], {{Webarchive|url=https://web.archive.org/web/20161220041712/http://rspa.royalsocietypublishing.org/content/royprsa/85/577/219.full.pdf |date=2016-12-20 }} ''Proceedings of the Royal Society A'', '''85''' (577): 219–29.</ref> The "whirling cloud of brilliant yellow light" produced by his apparatus reacted with [[mercury (element)|mercury]] to produce explosive [[mercury nitride]].<ref>[http://www.lateralscience.co.uk/activen/index.html Lord Rayleigh's Active Nitrogen] {{webarchive|url=https://web.archive.org/web/20121101100510/http://www.lateralscience.co.uk/activen/index.html |date=2012-11-01 }}. Lateralscience.co.uk. Retrieved 2011-10-26.</ref>
For a long time, sources of nitrogen compounds were limited. Natural sources originated either from biology or deposits of nitrates produced by atmospheric reactions. [[Nitrogen fixation]] by industrial processes like the [[Frank–Caro process]] (1895–1899) and [[Haber–Bosch process]] (1908–1913) eased this shortage of nitrogen compounds, to the extent that half of global [[food production]] now relies on synthetic nitrogen fertilisers.<ref name="Haber100">{{cite journal | doi = 10.1038/ngeo325 | title = How a century of ammonia synthesis changed the world |year = 2008 | last1 = Erisman | first1 = Jan Willem | last2 = Sutton | first2 = Mark A. | last3 = Galloway | first3 = James | last4 = Klimont | first4 = Zbigniew | last5 = Winiwarter | first5 = Wilfried | journal = Nature Geoscience | volume = 1 | issue = 10 | page = 636|bibcode = 2008NatGe...1..636E | s2cid = 94880859 }}</ref> At the same time, use of the [[Ostwald process]] (1902) to produce nitrates from industrial nitrogen fixation allowed the large-scale industrial production of nitrates as [[feedstock]] in the manufacture of [[explosives]] in the [[World war|World Wars]] of the 20th century.<ref>{{cite patent |inventor1-last=Ostwald |inventor1-first=Wilhelm |inventor1-link=Wilhelm Ostwald |title=Improvements in the Manufacture of Nitric Acid and Nitrogen Oxides |country=GB |number=190200698 |pubdate=1902-03-20}}</ref><ref>{{cite patent |inventor-last=Ostwald |inventor-first=Wilhelm |inventor-link=Wilhelm Ostwald |title=Improvements in and relating to the Manufacture of Nitric Acid and Oxides of Nitrogen |country=GB |number=190208300 |pubdate= 1903-02-26 |fdate=1902-12-18}}</ref>
==Properties==
===Atomic===
[[File:Neon orbitals.png|thumb|upright=1.6|right|The shapes of the five orbitals occupied in nitrogen. The two colours show the phase or sign of the wave function in each region. From left to right: 1s, 2s (cutaway to show internal structure), 2p<sub>''x''</sub>, 2p<sub>''y''</sub>, 2p<sub>''z''</sub>.]]
A nitrogen atom has seven electrons. In the ground state, they are arranged in the electron configuration 1s{{su|p=2}}2s{{su|p=2}}2p{{su|p=1|b=''x''}}2p{{su|p=1|b=''y''}}2p{{su|p=1|b=''z''}}. It, therefore, has five [[valence electron]]s in the 2s and 2p orbitals, three of which (the p-electrons) are unpaired. It has one of the highest [[electronegativity|electronegativities]] among the elements (3.04 on the Pauling scale), exceeded only by [[chlorine]] (3.16), [[oxygen]] (3.44), and [[fluorine]] (3.98). (The light [[noble gas]]es, [[helium]], [[neon]], and [[argon]], would presumably also be more electronegative, and in fact are on the Allen scale.)<ref name="Greenwood411" /> Following periodic trends, its single-bond [[covalent radius]] of 71&nbsp;pm is smaller than those of [[boron]] (84&nbsp;pm) and [[carbon]] (76&nbsp;pm), while it is larger than those of oxygen (66&nbsp;pm) and fluorine (57&nbsp;pm). The [[nitride]] anion, N<sup>3−</sup>, is much larger at 146&nbsp;pm, similar to that of the [[oxide]] (O<sup>2−</sup>: 140&nbsp;pm) and [[fluoride]] (F<sup>−</sup>: 133&nbsp;pm) anions.<ref name="Greenwood411" /> The first three ionisation energies of nitrogen are 1.402, 2.856, and 4.577&nbsp;MJ·mol<sup>−1</sup>, and the sum of the fourth and fifth is {{val|16.920|u=MJ·mol<sup>−1</sup>}}. Due to these very high figures, nitrogen has no simple cationic chemistry.<ref name="Greenwood550">Greenwood and Earnshaw, p. 550</ref>
The lack of radial nodes in the 2p subshell is directly responsible for many of the anomalous properties of the first row of the [[p-block]], especially in nitrogen, oxygen, and fluorine. The 2p subshell is very small and has a very similar radius to the 2s shell, facilitating [[orbital hybridisation]]. It also results in very large electrostatic forces of attraction between the nucleus and the valence electrons in the 2s and 2p shells, resulting in very high electronegativities. [[Hypervalent molecule|Hypervalency]] is almost unknown in the 2p elements for the same reason, because the high electronegativity makes it difficult for a small nitrogen atom to be a central atom in an electron-rich [[three-center four-electron bond]] since it would tend to attract the electrons strongly to itself. Thus, despite nitrogen's position at the head of group 15 in the periodic table, its chemistry shows huge differences from that of its heavier congeners [[phosphorus]], [[arsenic]], [[antimony]], and [[bismuth]].<ref name="Kaupp">{{cite journal |last=Kaupp |first=Martin |date=1 December 2006 |title=The role of radial nodes of atomic orbitals for chemical bonding and the periodic table |journal=Journal of Computational Chemistry |volume=28 |issue=1 |pages=320–25 |doi=10.1002/jcc.20522 |pmid=17143872 |s2cid=12677737 |doi-access=free }}</ref>
Nitrogen may be usefully compared to its horizontal neighbours' carbon and oxygen as well as its vertical neighbours in the pnictogen column, phosphorus, arsenic, antimony, and bismuth. Although each period 2 element from lithium to oxygen shows some similarities to the period 3 element in the next group (from magnesium to chlorine; these are known as [[diagonal relationship]]s), their degree drops off abruptly past the boron–silicon pair. The similarities of nitrogen to sulfur are mostly limited to sulfur nitride ring compounds when both elements are the only ones present.<ref name="Greenwood412" />
Nitrogen does not share the proclivity of carbon for [[catenation]]. Like carbon, nitrogen tends to form ionic or metallic compounds with metals. Nitrogen forms an extensive series of nitrides with carbon, including those with chain-, [[graphite|graphitic-]], and [[fullerene|fullerenic]]-like structures.<ref>{{cite journal |last1=Miller |first1=T. S. |last2=Belen |first2= A.|last3= Suter|first3= T. M.|last4= Sella|first4= A.|last5= Corà|first5= A.|last6= McMillan|first6= P. F.|date=2017 |title= Carbon nitrides: synthesis and characterization of a new class of functional materials |journal=Physical Chemistry Chemical Physics |volume= 19|issue= 24|pages=15613–15638 |doi=10.1039/C7CP02711G|pmid=28594419 |bibcode=2017PCCP...1915613M |doi-access= free}}</ref>
It resembles oxygen with its high electronegativity and concomitant capability for [[hydrogen bond]]ing and the ability to form [[coordination complex]]es by donating its [[lone pair]]s of electrons. There are some parallels between the chemistry of ammonia NH<sub>3</sub> and water H<sub>2</sub>O. For example, the capacity of both compounds to be protonated to give NH<sub>4</sub><sup>+</sup> and H<sub>3</sub>O<sup>+</sup> or deprotonated to give NH<sub>2</sub><sup>−</sup> and OH<sup>−</sup>, with all of these able to be isolated in solid compounds.<ref>{{cite book |last1= House|first1=J. E. |last2=House |first2= K. A.|date=2016 |title=Descriptive Inorganic Chemistry |location= Amsterdam|publisher= Elsevier|page= 198|isbn=978-0-12-804697-5 }}</ref>
Nitrogen shares with both its horizontal neighbours a preference for forming multiple bonds, typically with carbon, oxygen, or other nitrogen atoms, through p<sub>''π''</sub>–p<sub>''π''</sub> interactions.<ref name="Greenwood412" /> Thus, for example, nitrogen occurs as diatomic molecules and therefore has very much lower [[melting point|melting]] (−210&nbsp;°C) and [[boiling point]]s (−196&nbsp;°C) than the rest of its group, as the N<sub>2</sub> molecules are only held together by weak [[van der Waals interaction]]s and there are very few electrons available to create significant instantaneous dipoles. This is not possible for its vertical neighbours; thus, the [[nitrogen oxide]]s, [[nitrite]]s, [[nitrate]]s, [[nitro compound|nitro-]], [[nitroso]]-, [[azo compound|azo]]-, and [[diazo]]-compounds, [[azide]]s, [[cyanate]]s, [[thiocyanate]]s, and [[imino]]-derivatives find no echo with phosphorus, arsenic, antimony, or bismuth. By the same token, however, the complexity of the phosphorus oxoacids finds no echo with nitrogen.<ref name="Greenwood412" /> Setting aside their differences, nitrogen and phosphorus form an extensive series of compounds with one another; these have chain, ring, and cage structures.<ref>{{cite book |last1=Roy  |first1= A. K.|last2=Burns  |first2= G. T.|last3=Grigora  |first3= S.|last4=Lie  |first4= G. C.|editor-last1= Wisian-Neilson|editor-first1= P.|editor-last2= Alcock|editor-first2= H. R.|editor-last3= Wynne|editor-first3= K. J.|title=Inorganic and organometallic polymers II: advanced materials and intermediates |publisher=American Chemical Society |date=1994 |pages=344–357 |chapter=Poly(alkyl/aryloxothiazenes), [N=S(O)R]''<sub>n</sub>'' : New direction in inorganic polymers |doi=10.1021/bk-1994-0572.ch026}}</ref>
Table of thermal and physical properties of nitrogen (N<sub>2</sub>) at atmospheric pressure:<ref>{{Cite book |last=Holman |first=Jack P. |url=https://www.worldcat.org/oclc/46959719 |title=Heat transfer |publisher=McGraw-Hill Companies, Inc. |year=2002 |isbn=9780072406559 |edition=9th |location=New York, NY |pages=600–606 |language=English |oclc=46959719}}</ref><ref>{{Cite book |last1=Incropera |last2=Dewitt |last3=Bergman |last4=Lavigne |first1=Frank P. |first2=David P. |first3=Theodore L. |first4=Adrienne S. |url=https://www.worldcat.org/oclc/62532755 |title=Fundamentals of heat and mass transfer. |publisher=John Wiley and Sons, Inc. |year=2007 |isbn=9780471457282 |edition=6th |location=Hoboken, NJ |pages=941–950 |language=English |oclc=62532755}}</ref>
{|class="wikitable mw-collapsible mw-collapsed", style="text-align: right"
|style="text-align: center"|Temperature (K)
|style="text-align: center"|Density (kg&nbsp;m<sup>−3</sup>)
|style="text-align: center"|Specific heat (kJ&nbsp;kg<sup>−1</sup>&nbsp;°C<sup>−1</sup>)
|style="text-align: center"|Dynamic viscosity (kg&nbsp;m<sup>−1</sup>&nbsp;s<sup>−1</sup>)
|style="text-align: center"|Kinematic viscosity (m<sup>2</sup>&nbsp;s<sup>−1</sup>)
|style="text-align: center"|Thermal conductivity (W&nbsp;m<sup>−1</sup>&nbsp;°C<sup>−1</sup>)
|style="text-align: center"|Thermal diffusivity (m<sup>2</sup>&nbsp;s<sup>−1</sup>)
|style="text-align: center"|[[Prandtl number]]
|-
|100
|3.4388
|1.07{{figure space}}{{figure space}}
|{{val|6.88e-6}}
|{{val|2.00e-6}}
|{{gaps|0.009|58}}
|{{val|2.60e-6}}
|0.768
|-
|150
|2.2594
|1.05{{figure space}}{{figure space}}
|{{val|1.01e-5}}
|{{val|4.45e-6}}
|{{gaps|0.013|9{{figure space}}}}
|{{val|5.86e-6}}
|0.759
|-
|200
|1.7108
|1.0429
|{{val|1.29e-5}}
|{{val|7.57e-6}}
|{{gaps|0.018|24}}
|{{val|1.02e-5}}
|0.747
|-
|300
|1.1421
|1.0408
|{{val|1.78e-5}}
|{{val|1.56e-5}}
|{{gaps|0.026|2{{figure space}}}}
|{{val|2.20e-5}}
|0.713
|-
|400
|0.8538
|1.0459
|{{val|2.20e-5}}
|{{val|2.57e-5}}
|{{gaps|0.033|35}}
|{{val|3.73e-5}}
|0.691
|-
|500
|0.6824
|1.0555
|{{val|2.57e-5}}
|{{val|3.77e-5}}
|{{gaps|0.039|84}}
|{{val|5.53e-5}}
|0.684
|-
|600
|0.5687
|1.0756
|{{val|2.91e-5}}
|{{val|5.12e-5}}
|{{gaps|0.045|8{{figure space}}}}
|{{val|7.49e-5}}
|0.686
|-
|700
|0.4934
|1.0969
|{{val|3.21e-5}}
|{{val|6.67e-5}}
|{{gaps|0.051|23}}
|{{val|9.47e-5}}
|0.691
|-
|800
|0.4277
|1.1225
|{{val|3.48e-5}}
|{{val|8.15e-5}}
|{{gaps|0.056|09}}
|{{val|1.17e-4}}
|0.7{{figure space}}{{figure space}}
|-
|900
|0.3796
|1.1464
|{{val|3.75e-5}}
|{{val|9.11e-5}}
|{{gaps|0.060|7{{figure space}}}}
|{{val|1.39e-4}}
|0.711
|-
|1000
|0.3412
|1.1677
|{{val|4.00e-5}}
|{{val|1.19e-4}}
|{{gaps|0.064|75}}
|{{val|1.63e-4}}
|0.724
|-
|1100
|0.3108
|1.1857
|{{val|4.23e-5}}
|{{val|1.36e-4}}
|{{gaps|0.068|5{{figure space}}}}
|{{val|1.86e-4}}
|0.736
|-
|1200
|0.2851
|1.2037
|{{val|4.45e-5}}
|{{val|1.56e-4}}
|{{gaps|0.071|84}}
|{{val|2.09e-4}}
|0.748
|-
|{{val|1300}}
|0.2591
|1.219{{figure space}}
|{{val|4.66e-5}}
|{{val|1.80e-4}}
|{{gaps|0.081|{{figure space}}{{figure space}}}}
|{{val|2.56e-4}}
|0.701
|}
===Isotopes===
{{main|Isotopes of nitrogen}}
[[File:NuclideMap C-F.png|thumb|right|upright=2.3|Table of nuclides (Segrè chart) from carbon to fluorine (including nitrogen). Orange indicates [[proton emission]] (nuclides outside the proton drip line); pink for [[positron emission]] (inverse beta decay); black for [[stable nuclide|stable]] nuclides; blue for [[electron emission]] (beta decay); and violet for [[neutron emission]] (nuclides outside the neutron drip line). Proton number increases going up the vertical axis and neutron number going to the right on the horizontal axis.]]
Nitrogen has two stable [[isotope]]s: <sup>14</sup>N and <sup>15</sup>N. The first is much more common, making up 99.634% of natural nitrogen, and the second (which is slightly heavier) makes up the remaining 0.366%. This leads to an atomic weight of around 14.007&nbsp;u.<ref name="Greenwood411">Greenwood and Earnshaw, pp. 411–12</ref> Both of these stable isotopes are produced in the [[CNO cycle]] in [[star]]s, but <sup>14</sup>N is more common as its proton capture is the rate-limiting step. <sup>14</sup>N is one of the five stable [[even and odd atomic nuclei|odd–odd nuclides]] (a nuclide having an odd number of protons and neutrons); the other four are [[deuterium|<sup>2</sup>H]], <sup>6</sup>Li, <sup>10</sup>B, and <sup>180m</sup>Ta.<ref name="BetheBible">{{cite journal|last=Bethe |first=H. A.|date=1939|title=Energy Production in Stars|journal=[[Physical Review]]|volume=55 |issue=5 |pages=434–56|bibcode=1939PhRv...55..434B|doi= 10.1103/PhysRev.55.434|pmid=17835673|doi-access=free}}</ref>
The relative abundance of <sup>14</sup>N and <sup>15</sup>N is practically constant in the atmosphere but can vary elsewhere, due to natural isotopic fractionation from biological [[redox]] reactions and the evaporation of natural [[ammonia]] or [[nitric acid]].<ref>{{cite web |url=http://www.ciaaw.org/nitrogen.htm |title=Atomic Weight of Nitrogen |author=CIAAW |date=2003 |website=ciaaw.org |publisher=CIAAW |access-date=13 October 2016 |archive-date=14 October 2016 |archive-url=https://web.archive.org/web/20161014112435/http://www.ciaaw.org/nitrogen.htm |url-status=live }}</ref> 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]].<ref name="enrich">{{cite book | url = https://books.google.com/books?id=U9y3whFC2DIC&pg=PA74 | pages = 74–75 | title = Stable Isotopes and Biosphere – Atmosphere Interactions: Processes and Biological Controls | isbn = 978-0-08-052528-0 | last1 = Flanagan | first1 = Lawrence B. | last2 = Ehleringer | first2 = James R. | last3 = Pataki | first3 = Diane E. | date = 15 December 2004 | publisher = Elsevier | access-date = 20 December 2015 | archive-date = 5 February 2016 | archive-url = https://web.archive.org/web/20160205191759/https://books.google.com/books?id=U9y3whFC2DIC&pg=PA74 | url-status = live }}</ref>
The heavy isotope <sup>15</sup>N was first discovered by S. M. Naudé in 1929, and soon after heavy isotopes of the neighbouring elements [[oxygen]] and [[carbon]] were discovered.<ref name="Greenwood408">Greenwood and Earnshaw, p. 408</ref> It presents one of the lowest thermal neutron capture cross-sections of all isotopes.<ref>{{cite web |url=http://www.nndc.bnl.gov/sigma/index.jsp?as=15&lib=endfb7.1&nsub=10 |title=Evaluated Nuclear Data File (ENDF) Retrieval & Plotting |publisher=National Nuclear Data Center |access-date=2016-11-23 |archive-date=2020-08-09 |archive-url=https://web.archive.org/web/20200809124433/https://www.nndc.bnl.gov/sigma/index.jsp?as=15&lib=endfb7.1&nsub=10 |url-status=live }}</ref> It is frequently used in [[nuclear magnetic resonance]] (NMR) spectroscopy to determine the structures of nitrogen-containing molecules, due to its fractional [[nuclear spin]] of one-half, which offers advantages for NMR such as narrower line width. <sup>14</sup>N, though also theoretically usable, has an integer nuclear spin of one and thus has a [[quadrupole moment]] that leads to wider and less useful spectra.<ref name="Greenwood411" /> <sup>15</sup>N NMR nevertheless has complications not encountered in the more common <sup>1</sup>H and <sup>13</sup>C NMR spectroscopy. The low natural abundance of <sup>15</sup>N (0.36%) significantly reduces sensitivity, a problem which is only exacerbated by its low [[gyromagnetic ratio]], (only 10.14% that of <sup>1</sup>H). As a result, the signal-to-noise ratio for <sup>1</sup>H is about 300 times as much as that for <sup>15</sup>N at the same magnetic field strength.<ref name="autogenerated2007">{{cite book| author=Arthur G Palmer| title=Protein NMR Spectroscopy| publisher=Elsevier Academic Press| date=2007 | isbn = 978-0-12-164491-8}}</ref> This may be somewhat alleviated by isotopic enrichment of <sup>15</sup>N by chemical exchange or fractional distillation. <sup>15</sup>N-enriched compounds have the advantage that under standard conditions, they do not undergo chemical exchange of their nitrogen atoms with atmospheric nitrogen, unlike compounds with labelled [[hydrogen]], carbon, and oxygen isotopes that must be kept away from the atmosphere.<ref name="Greenwood411" /> The <sup>15</sup>N:<sup>14</sup>N ratio is commonly used in stable isotope analysis in the fields of [[geochemistry]], [[hydrology]], [[paleoclimatology]] and [[paleoceanography]], where it is called [[δ15N|''δ''<sup>15</sup>N]].<ref>{{cite book | last=Katzenberg | first=M. A. | title=Biological Anthropology of the Human Skeleton | chapter=Chapter 13: Stable Isotope Analysis: A Tool for Studying Past Diet, Demography, and Life History | year=2008 | publisher=Wiley | edition=2nd | isbn=978-0-471-79372-4 }}</ref>
Of the thirteen other isotopes produced synthetically, ranging from <sup>9</sup>N to <sup>23</sup>N, [[nitrogen-13|<sup>13</sup>N]] has a [[half-life]] of ten minutes and the remaining isotopes have half-lives less than eight seconds.<ref name=N9sci>{{cite news |url=https://www.science.org/content/article/fleeting-form-nitrogen-stretches-nuclear-theory-its-limits |title=Fleeting form of nitrogen stretches nuclear theory to its limits |last=Cho |first=Adrian |website=[[science.org]] |date=25 September 2023 |access-date=27 September 2023}}</ref><ref name="NUBASE">{{NUBASE 2003}}</ref> Given the half-life difference, <sup>13</sup>N is the most important nitrogen radioisotope, being relatively long-lived enough to use in [[positron emission tomography]] (PET), although its half-life is still short and thus it must be produced at the venue of the PET, for example in a [[cyclotron]] via proton bombardment of <sup>16</sup>O producing <sup>13</sup>N and an [[alpha particle]].<ref name="Carlson 151">{{cite book | last = Carlson | first = Neil | title = Physiology of Behavior | publisher = Pearson | series = Methods and Strategies of Research | volume = 11th edition | date = January 22, 2012 | page = 151 | isbn = 978-0-205-23939-9}}</ref>
The [[radioisotope]] <sup>16</sup>N is the dominant [[radionuclide]] in the coolant of [[pressurised water reactor]]s or [[boiling water reactor]]s during normal operation. It is produced from <sup>16</sup>O (in water) via an [[Np reaction|(n,p) reaction]], in which the <sup>16</sup>O atom captures a neutron and expels a proton. It has a short half-life of about 7.1&nbsp;s,<ref name="NUBASE" /> but its decay back to <sup>16</sup>O produces high-energy [[gamma radiation]] (5 to 7&nbsp;MeV).<ref name="NUBASE" /><ref name="Neeb">{{Cite book|last=Neeb|first=Karl Heinz|title=The Radiochemistry of Nuclear Power Plants with Light Water Reactors|publisher=Walter de Gruyter|location=Berlin-New York|date=1997|isbn=978-3-11-013242-7|url=https://books.google.com/books?id=SJOE00whg44C&pg=PA227|page=227|access-date=2015-12-20|archive-date=2016-02-05|archive-url=https://web.archive.org/web/20160205191759/https://books.google.com/books?id=SJOE00whg44C&pg=PA227|url-status=live}}</ref> Because of this, access to the primary coolant piping in a pressurised water reactor must be restricted during [[Nuclear reactor|reactor]] power operation.<ref name="Neeb" /> It is a sensitive and immediate indicator of leaks from the primary coolant system to the secondary steam cycle and is the primary means of detection for such leaks.<ref name="Neeb"/>
===Allotropes===
{{see also|Solid nitrogen}}
[[File:N2MolecularDiagramCR.jpg|thumb|right|upright=1.8|[[Molecular orbital diagram]] of dinitrogen molecule, N<sub>2</sub>. There are five bonding orbitals and two antibonding orbitals (marked with an asterisk; orbitals involving the inner 1s electrons not shown), giving a total bond order of three.]]
Atomic nitrogen, also known as active nitrogen, is highly reactive, being a [[radical (chemistry)|triradical]] with three unpaired electrons. Free nitrogen atoms easily react with most elements to form nitrides, and even when two free nitrogen atoms collide to produce an excited N<sub>2</sub> molecule, they may release so much energy on collision with even such stable molecules as [[carbon dioxide]] and [[water]] to cause homolytic fission into radicals such as CO and O or OH and H. Atomic nitrogen is prepared by passing an electric discharge through nitrogen gas at 0.1–2&nbsp;mmHg, which produces atomic nitrogen along with a peach-yellow emission that fades slowly as an afterglow for several minutes even after the discharge terminates.<ref name="Greenwood412">Greenwood and Earnshaw, pp. 412–16</ref>
Given the great reactivity of atomic nitrogen, elemental nitrogen usually occurs as molecular N<sub>2</sub>, dinitrogen. This molecule is a colourless, odourless, and tasteless [[diamagnetic]] gas at standard conditions: it melts at −210&nbsp;°C and boils at −196&nbsp;°C.<ref name="Greenwood412" /> Dinitrogen is mostly unreactive at room temperature, but it will nevertheless react with [[lithium]] metal and some [[transition metal]] complexes. This is due to its bonding, which is unique among the diatomic elements at standard conditions in that it has an N≡N [[triple bond]]. Triple bonds have short bond lengths (in this case, 109.76&nbsp;pm) and high dissociation energies (in this case, 945.41&nbsp;kJ/mol), and are thus very strong, explaining dinitrogen's low level of chemical reactivity.<ref name="Greenwood412" /><ref>{{cite web | url=http://www.uigi.com/nitrogen.html | title=Universal Industrial Gases, Inc...Nitrogen N2 Properties, Uses, Applications - Gas and Liquid }}</ref>
Other nitrogen [[oligomers]] and polymers may be possible. If they could be synthesised, they may have potential applications as materials with a very high energy density, that could be used as powerful propellants or explosives.<ref name="Lewars">{{cite book |title=Modeling Marvels: Computational Anticipation of Novel molecules |last=Lewars |first=Errol G. |year=2008 |publisher=[[Springer Science+Business Media]] |isbn=978-1-4020-6972-7 |doi=10.1007/978-1-4020-6973-4 |pages=141–63 }}</ref>  Under extremely high pressures (1.1&nbsp;million&nbsp;[[Atmosphere (unit)|atm]]) and high temperatures (2000&nbsp;K), as produced in a [[diamond anvil cell]], nitrogen polymerises into the single-bonded [[cubic gauche]] crystal structure. This structure is similar to that of [[diamond]], and both have extremely strong [[covalent bond]]s, resulting in its nickname "nitrogen diamond".<ref>{{Cite news|url=http://www.physorg.com/news693.html|title=Polymeric nitrogen synthesized|publisher=physorg.com|date=5 August 2004|access-date=2009-06-22|archive-date=2012-01-24|archive-url=https://web.archive.org/web/20120124231419/http://www.physorg.com/news693.html|url-status=live}}</ref>
[[File:Mountainous Shoreline of Sputnik Planum (PIA20198).png|thumb|right|upright=1.1|[[Solid nitrogen]] on the plains of [[Sputnik Planitia]] (on the bottom-right side of the image) on [[Pluto]] next to water ice mountains (on the up-left side of the image)]]
At [[atmospheric pressure]], molecular nitrogen [[condensation|condenses]] ([[liquid|liquefies]]) at 77&nbsp;[[Kelvin|K]] (−195.79&nbsp;°[[Celsius|C]]) and [[freezing|freezes]] at 63&nbsp;K (−210.01&nbsp;°C)<ref name="Gray">{{cite book|last=Gray|first=Theodore|title=The Elements: A Visual Exploration of Every Known Atom in the Universe|date=2009|publisher=Black Dog & Leventhal Publishers|location=New York|isbn=978-1-57912-814-2|url-access=registration|url=https://archive.org/details/elementsvisualex0000gray}}</ref> into the beta [[hexagonal close-packed]] crystal [[Allotropy|allotropic]] form. Below 35.4&nbsp;K (−237.6&nbsp;°C) nitrogen assumes the [[Cubic crystal system|cubic]] crystal allotropic form (called the alpha phase).<ref name="schu">{{cite journal|last1=Schuch|first1=A. F.|last2=Mills|first2=R. L.|title=Crystal Structures of the Three Modifications of Nitrogen 14 and Nitrogen 15 at High Pressure|journal=The Journal of Chemical Physics|date=1970|volume=52|issue=12|pages=6000–08|doi=10.1063/1.1672899|bibcode=1970JChPh..52.6000S}}</ref> [[Liquid nitrogen]], a colourless fluid resembling water in appearance, but with 80.8% of the density (the density of liquid nitrogen at its boiling point is 0.808&nbsp;g/mL), is a common [[cryogen]].<ref>{{Cite journal | last1 = Iancu | first1 = C. V. | last2 = Wright | first2 = E. R. | last3 = Heymann | first3 = J. B. | last4 = Jensen | first4 = G. J. | title = A comparison of liquid nitrogen and liquid helium as cryogens for electron cryotomography | doi = 10.1016/j.jsb.2005.12.004 | journal = Journal of Structural Biology | volume = 153 | issue = 3 | pages = 231–40 | year = 2006 | pmid = 16427786}}</ref> [[Solid nitrogen]] has many crystalline modifications. It forms a significant dynamic surface coverage on Pluto<ref>{{cite news|title=Flowing nitrogen ice glaciers seen on surface of Pluto after New Horizons flyby|url=http://www.abc.net.au/news/2015-07-25/flowing-nitrogen-ice-glaciers-seen-on-surface-of-pluto/6647636|newspaper=ABC News|access-date=6 October 2015|date=25 July 2015|archive-date=29 September 2015|archive-url=https://web.archive.org/web/20150929044226/http://www.abc.net.au/news/2015-07-25/flowing-nitrogen-ice-glaciers-seen-on-surface-of-pluto/6647636|url-status=live}}</ref> and outer moons of the Solar System such as [[Triton (moon)|Triton]].<ref>{{cite encyclopedia|title = Encyclopedia of the Solar System|chapter = Triton|last1 = McKinnon|first1 = William B.|last2 = Kirk|first2 = Randolph L.|publisher = [[Elsevier]]|date = 2014|editor1-first = Tilman|editor1-last = Spohn|editor2-first = Doris|editor2-last = Breuer|editor3-first = Torrence|editor3-last = Johnson|edition = 3rd|location = Amsterdam; Boston|isbn = 978-0-12-416034-7|pages = 861–82|chapter-url = https://books.google.com/books?id=0bEMAwAAQBAJ&pg=PA861|access-date = 2016-04-30|archive-date = 2016-09-03|archive-url = https://web.archive.org/web/20160903233037/https://books.google.com/books?id=0bEMAwAAQBAJ&pg=PA861|url-status = live}}</ref> Even at the low temperatures of solid nitrogen it is fairly volatile and can [[sublimation (phase transition)|sublime]] to form an atmosphere, or condense back into nitrogen frost. It is very weak and flows in the form of glaciers, and on Triton [[geyser]]s of nitrogen gas come from the polar ice cap region.<ref>{{cite web |url=http://solarsystem.nasa.gov/planets/profile.cfm?Object=Triton |publisher=[[NASA]] |access-date=September 21, 2007 |title=Neptune: Moons: Triton |archive-url=https://web.archive.org/web/20111015074425/http://solarsystem.nasa.gov/planets/profile.cfm?Object=Triton |archive-date=October 15, 2011 |url-status=dead }}</ref>
==Chemistry and compounds==
===Dinitrogen complexes===
{{main|Dinitrogen complex}}
[[File:RuA5N2.png|thumb|right|Structure of [Ru(NH<sub>3</sub>)<sub>5</sub>(N<sub>2</sub>)]<sup>2+</sup> ([[pentaamine(dinitrogen)ruthenium(II) chloride|pentaamine(dinitrogen)ruthenium(II)]]), the first dinitrogen complex to be discovered]]
The first example of a [[dinitrogen complex]] to be discovered was [Ru(NH<sub>3</sub>)<sub>5</sub>(N<sub>2</sub>)]<sup>2+</sup> (see figure at right), and soon many other such complexes were discovered. These [[Complex (chemistry)|complexes]], in which a nitrogen molecule donates at least one lone pair of electrons to a central metal cation, illustrate how N<sub>2</sub> might bind to the metal(s) in [[nitrogenase]] and the [[Catalysis|catalyst]] for the [[Haber process]]: these processes involving dinitrogen activation are vitally important in biology and in the production of fertilisers.<ref>{{Cite journal|author=Fryzuk, M. D.|author2=Johnson, S. A.|name-list-style=amp|title=The continuing story of dinitrogen activation|journal=Coordination Chemistry Reviews|volume=200–202|page=379|year=2000|doi=10.1016/S0010-8545(00)00264-2}}</ref><ref name="Schrock">{{Cite journal|author=Schrock, R. R.|title=Catalytic Reduction of Dinitrogen to Ammonia at a Single Molybdenum Center|journal=[[Acc. Chem. Res.]]|year=2005|volume=38|issue=12|pmid=16359167|pages=955–62|pmc=2551323|doi=10.1021/ar0501121|author-link=Richard R. Schrock}}</ref>
Dinitrogen is able to coordinate to metals in five different ways. The more well-characterised ways are the end-on M←N≡N (''[[hapticity|η]]''<sup>1</sup>) and M←N≡N→M (''[[bridging ligand|μ]]'', bis-''η''<sup>1</sup>), in which the lone pairs on the nitrogen atoms are donated to the metal cation. The less well-characterised ways involve dinitrogen donating electron pairs from the triple bond, either as a [[bridging ligand]] to two metal cations (''μ'', bis-''η''<sup>2</sup>) or to just one (''η''<sup>2</sup>). The fifth and unique method involves triple-coordination as a bridging ligand, donating all three electron pairs from the triple bond (''μ''<sub>3</sub>-N<sub>2</sub>). A few complexes feature multiple N<sub>2</sub> ligands and some feature N<sub>2</sub> bonded in multiple ways. Since N<sub>2</sub> is isoelectronic with [[carbon monoxide]] (CO) and [[acetylene]] (C<sub>2</sub>H<sub>2</sub>), the bonding in dinitrogen complexes is closely allied to that in [[carbonyl]] compounds, although N<sub>2</sub> is a weaker ''σ''-donor and ''π''-acceptor than CO. Theoretical studies show that ''σ'' donation is a more important factor allowing the formation of the M–N bond than ''π'' back-donation, which mostly only weakens the N–N bond, and end-on (''η''<sup>1</sup>) donation is more readily accomplished than side-on (''η''<sup>2</sup>) donation.<ref name="Greenwood412" />
Today, dinitrogen complexes are known for almost all the [[transition metal]]s, accounting for several hundred compounds. They are normally prepared by three methods:<ref name="Greenwood412" />
# Replacing labile ligands such as [[water|H<sub>2</sub>O]], [[hydride|H<sup>−</sup>]], or [[carbon monoxide|CO]] directly by nitrogen: these are often reversible reactions that proceed at mild conditions.
# Reducing metal complexes in the presence of a suitable co-ligand in excess under nitrogen gas. A common choice includes replacing chloride ligands with [[dimethylphenylphosphine]] (PMe<sub>2</sub>Ph) to make up for the smaller number of nitrogen ligands attached to the original chlorine ligands.
# Converting a ligand with N–N bonds, such as hydrazine or azide, directly into a dinitrogen ligand.
Occasionally the N≡N bond may be formed directly within a metal complex, for example by directly reacting coordinated [[ammonia]] (NH<sub>3</sub>) with [[nitrous acid]] (HNO<sub>2</sub>), but this is not generally applicable. Most dinitrogen complexes have colours within the range white-yellow-orange-red-brown; a few exceptions are known, such as the blue [{Ti(''η''<sup>5</sup>-C<sub>5</sub>H<sub>5</sub>)<sub>2</sub>}<sub>2</sub>-(N<sub>2</sub>)].<ref name="Greenwood412" />
===Nitrides, azides, and nitrido complexes===
Nitrogen bonds to almost all the elements in the periodic table except the first two [[noble gas]]es, [[helium]] and [[neon]], and some of the very short-lived elements after [[bismuth]], creating an immense variety of binary compounds with varying properties and applications.<ref name="Greenwood412" /> Many binary compounds are known: with the exception of the nitrogen hydrides, oxides, and fluorides, these are typically called [[nitride]]s. Many stoichiometric phases are usually present for most elements (e.g. MnN, Mn<sub>6</sub>N<sub>5</sub>, Mn<sub>3</sub>N<sub>2</sub>, Mn<sub>2</sub>N, Mn<sub>4</sub>N, and Mn<sub>''x''</sub>N for 9.2 < ''x'' < 25.3). They may be classified as "salt-like" (mostly ionic), covalent, "diamond-like", and metallic (or [[interstitial compound|interstitial]]), although this classification has limitations generally stemming from the continuity of bonding types instead of the discrete and separate types that it implies. They are normally prepared by directly reacting a metal with nitrogen or ammonia (sometimes after heating), or by [[thermal decomposition]] of metal amides:<ref name="Greenwood417">Greenwood and Earnshaw, pp. 417–20</ref>
:3 Ca + N<sub>2</sub> → Ca<sub>3</sub>N<sub>2</sub>
:3 Mg + 2 NH<sub>3</sub> → Mg<sub>3</sub>N<sub>2</sub> + 3 H<sub>2</sub> (at 900&nbsp;°C)
:3 Zn(NH<sub>2</sub>)<sub>2</sub> → Zn<sub>3</sub>N<sub>2</sub> + 4 NH<sub>3</sub>
Many variants on these processes are possible. The most ionic of these nitrides are those of the [[alkali metal]]s and [[alkaline earth metal]]s, Li<sub>3</sub>N (Na, K, Rb, and Cs do not form stable nitrides for steric reasons) and M<sub>3</sub>N<sub>2</sub> (M = Be, Mg, Ca, Sr, Ba). These can formally be thought of as salts of the N<sup>3−</sup> anion, although charge separation is not actually complete even for these highly electropositive elements. However, the alkali metal [[azide]]s NaN<sub>3</sub> and KN<sub>3</sub>, featuring the linear {{chem|N|3|-}} anion, are well-known, as are Sr(N<sub>3</sub>)<sub>2</sub> and Ba(N<sub>3</sub>)<sub>2</sub>. Azides of the B-subgroup metals (those in [[group 11 element|groups 11]] through [[chalcogen|16]]) are much less ionic, have more complicated structures, and detonate readily when shocked.<ref name="Greenwood417" />
[[File:Borazin Mesomers1.svg|thumb|right|upright=1.8|Mesomeric structures of borazine, (–BH–NH–)<sub>3</sub>]]
Many covalent binary nitrides are known. Examples include [[cyanogen]] ((CN)<sub>2</sub>), [[triphosphorus pentanitride]] (P<sub>3</sub>N<sub>5</sub>), [[disulfur dinitride]] (S<sub>2</sub>N<sub>2</sub>), and [[tetrasulfur tetranitride]] (S<sub>4</sub>N<sub>4</sub>). The essentially covalent [[silicon nitride]] (Si<sub>3</sub>N<sub>4</sub>) and [[germanium nitride]] (Ge<sub>3</sub>N<sub>4</sub>) are also known: silicon nitride, in particular, would make a promising [[ceramic]] if not for the difficulty of working with and sintering it. In particular, the [[boron group|group 13]] nitrides, most of which are promising [[semiconductor]]s, are isoelectronic with graphite, diamond, and [[silicon carbide]] and have similar structures: their bonding changes from covalent to partially ionic to metallic as the group is descended. In particular, since the B–N unit is isoelectronic to C–C, and carbon is essentially intermediate in size between boron and nitrogen, much of [[organic chemistry]] finds an echo in boron–nitrogen chemistry, such as in [[borazine]] ("inorganic [[benzene]]"). Nevertheless, the analogy is not exact due to the ease of [[nucleophile|nucleophilic]] attack at boron due to its deficiency in electrons, which is not possible in a wholly carbon-containing ring.<ref name="Greenwood417" />
The largest category of nitrides are the interstitial nitrides of formulae MN, M<sub>2</sub>N, and M<sub>4</sub>N (although variable composition is perfectly possible), where the small nitrogen atoms are positioned in the gaps in a metallic cubic or [[hexagonal close-packed]] lattice. They are opaque, very hard, and chemically inert, melting only at very high temperatures (generally over 2500&nbsp;°C). They have a metallic lustre and conduct electricity as do metals. They hydrolyse only very slowly to give ammonia or nitrogen.<ref name="Greenwood417" />
The nitride anion (N<sup>3−</sup>) is the strongest ''π'' donor known among ligands (the second-strongest is O<sup>2−</sup>). Nitrido complexes are generally made by the thermal decomposition of azides or by deprotonating ammonia, and they usually involve a terminal {≡N}<sup>3−</sup> group. The linear azide anion ({{chem|N|3|-}}), being isoelectronic with [[nitrous oxide]], [[carbon dioxide]], and [[cyanate]], forms many coordination complexes. Further catenation is rare, although {{chem|N|4|4-}} (isoelectronic with [[carbonate]] and [[nitrate]]) is known.<ref name="Greenwood417" />
===Hydrides===
[[File:Nitrogen electrode potentials.svg|thumb|right|upright=2.3|Standard reduction potentials for nitrogen-containing species. Top diagram shows potentials at pH&nbsp;0; bottom diagram shows potentials at pH&nbsp;14.<ref name="Greenwood434">Greenwood and Earnshaw, pp. 434–38</ref>]]
Industrially, [[ammonia]] (NH<sub>3</sub>) is the most important compound of nitrogen and is prepared in larger amounts than any other compound because it contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to food and fertilisers. It is a colourless alkaline gas with a characteristic pungent smell. The presence of [[hydrogen bonding]] has very significant effects on ammonia, conferring on it its high melting (−78&nbsp;°C) and boiling (−33&nbsp;°C) points. As a liquid, it is a very good solvent with a high heat of vaporisation (enabling it to be used in vacuum flasks), that also has a low viscosity and electrical conductivity and high [[dielectric constant]], and is less dense than water. However, the hydrogen bonding in NH<sub>3</sub> is weaker than that in H<sub>2</sub>O due to the lower electronegativity of nitrogen compared to oxygen and the presence of only one lone pair in NH<sub>3</sub> rather than two in H<sub>2</sub>O. It is a weak base in aqueous solution ([[acid dissociation constant|p''K''<sub>''b''</sub>]] 4.74); its conjugate acid is [[ammonium]], {{chem|NH|4|+}}. It can also act as an extremely weak acid, losing a proton to produce the amide anion, {{chem|NH|2|-}}. It thus undergoes self-dissociation, similar to water, to produce ammonium and amide. Ammonia burns in air or oxygen, though not readily, to produce nitrogen gas; it burns in fluorine with a greenish-yellow flame to give [[nitrogen trifluoride]]. Reactions with the other nonmetals are very complex and tend to lead to a mixture of products. Ammonia reacts on heating with metals to give nitrides.<ref name="Greenwood420">Greenwood and Earnshaw, pp. 420–26</ref>
Many other binary nitrogen hydrides are known, but the most important are [[hydrazine]] (N<sub>2</sub>H<sub>4</sub>) and [[hydrogen azide]] (HN<sub>3</sub>). Although it is not a nitrogen hydride, [[hydroxylamine]] (NH<sub>2</sub>OH) is similar in properties and structure to ammonia and hydrazine as well. Hydrazine is a fuming, colourless liquid that smells similar to ammonia. Its physical properties are very similar to those of water (melting point 2.0&nbsp;°C, boiling point 113.5&nbsp;°C, density 1.00&nbsp;g/cm<sup>3</sup>). Despite it being an endothermic compound, it is kinetically stable. It burns quickly and completely in air very exothermically to give nitrogen and water vapour. It is a very useful and versatile reducing agent and is a weaker base than ammonia.<ref name="Greenwood426" /> It is also commonly used as a rocket fuel.<ref name="Vieira">{{cite journal | last = Vieira | first = R. |author2=C. Pham-Huu |author3=N. Keller |author4=M. J. Ledoux | year = 2002 | title = New carbon nanofiber/graphite felt composite for use as a catalyst for hydrazine catalytic decomposition | journal = [[Chemical Communications]] | issue = 9 | pages = 954–55 | doi = 10.1039/b202032g| pmid = 12123065 }}</ref>
Hydrazine is generally made by reaction of ammonia with alkaline [[sodium hypochlorite]] in the presence of gelatin or glue:<ref name="Greenwood426">Greenwood and Earnshaw, pp. 426–33</ref>
:NH<sub>3</sub> + OCl<sup>−</sup> → NH<sub>2</sub>Cl + OH<sup>−</sup>
:NH<sub>2</sub>Cl + NH<sub>3</sub> → {{chem|N|2|H|5|+}} + Cl<sup>−</sup> (slow)
:{{chem|N|2|H|5|+}} + OH<sup>−</sup> → N<sub>2</sub>H<sub>4</sub> + H<sub>2</sub>O (fast)
(The attacks by hydroxide and ammonia may be reversed, thus passing through the intermediate NHCl<sup>−</sup> instead.) The reason for adding gelatin is that it removes metal ions such as Cu<sup>2+</sup> that catalyses the destruction of hydrazine by reaction with [[monochloramine]] (NH<sub>2</sub>Cl) to produce [[ammonium chloride]] and nitrogen.<ref name="Greenwood426" />
[[Hydrogen azide]] (HN<sub>3</sub>) was first produced in 1890 by the oxidation of aqueous hydrazine by nitrous acid. It is very explosive and even dilute solutions can be dangerous. It has a disagreeable and irritating smell and is a potentially lethal (but not cumulative) poison. It may be considered the conjugate acid of the azide anion, and is similarly analogous to the [[hydrohalic acid]]s.<ref name="Greenwood426" />
===Halides and oxohalides===
[[File:Nitrogen trichloride.JPG|thumb|right|[[Nitrogen trichloride]]]]
All four simple nitrogen trihalides are known. A few mixed halides and hydrohalides are known, but are mostly unstable; examples include NClF<sub>2</sub>, NCl<sub>2</sub>F, NBrF<sub>2</sub>, NF<sub>2</sub>H, [[fluoroamine|NFH<sub>2</sub>]], [[Dichloramine|NCl<sub>2</sub>H]], and [[Monochloramine|NClH<sub>2</sub>]].<ref name="Greenwood438">Greenwood and Earnshaw, pp. 438–42</ref>
[[Nitrogen trifluoride]] (NF<sub>3</sub>, first prepared in 1928) is a colourless and odourless gas that is thermodynamically stable, and most readily produced by the [[electrolysis]] of molten [[ammonium fluoride]] dissolved in anhydrous [[hydrogen fluoride]]. Like [[carbon tetrafluoride]], it is not at all reactive and is stable in water or dilute aqueous acids or alkalis. Only when heated does it act as a fluorinating agent, and it reacts with [[copper]], arsenic, antimony, and bismuth on contact at high temperatures to give [[tetrafluorohydrazine]] (N<sub>2</sub>F<sub>4</sub>). The cations {{chem|NF|4|+}} and {{chem|N|2|F|3|+}} are also known (the latter from reacting tetrafluorohydrazine with strong fluoride-acceptors such as [[arsenic pentafluoride]]), as is ONF<sub>3</sub>, which has aroused interest due to the short N–O distance implying partial double bonding and the highly polar and long N–F bond. Tetrafluorohydrazine, unlike hydrazine itself, can dissociate at room temperature and above to give the radical NF<sub>2</sub>•. [[Fluorine azide]] (FN<sub>3</sub>) is very explosive and thermally unstable. [[Dinitrogen difluoride]] (N<sub>2</sub>F<sub>2</sub>) exists as thermally interconvertible ''cis'' and ''trans'' isomers, and was first found as a product of the thermal decomposition of FN<sub>3</sub>.<ref name="Greenwood438" />
[[Nitrogen trichloride]] (NCl<sub>3</sub>) is a dense, volatile, and explosive liquid whose physical properties are similar to those of [[carbon tetrachloride]], although one difference is that NCl<sub>3</sub> is easily hydrolysed by water while CCl<sub>4</sub> is not. It was first synthesised in 1811 by [[Pierre Louis Dulong]], who lost three fingers and an eye to its explosive tendencies. As a dilute gas it is less dangerous and is thus used industrially to bleach and sterilise flour. [[Nitrogen tribromide]] (NBr<sub>3</sub>), first prepared in 1975, is a deep red, temperature-sensitive, volatile solid that is explosive even at −100&nbsp;°C. [[Nitrogen triiodide]] (NI<sub>3</sub>) is still more unstable and was only prepared in 1990. Its adduct with ammonia, which was known earlier, is very shock-sensitive: it can be set off by the touch of a feather, shifting air currents, or even [[alpha particle]]s.<ref name="Greenwood438" /><ref>{{ cite journal | author = Bowden, F. P. | title = Initiation of Explosion by Neutrons, α-Particles, and Fission Products | journal = Proceedings of the Royal Society of London A | year = 1958 | volume = 246 | issue = 1245 | pages = 216–19 | doi = 10.1098/rspa.1958.0123 | bibcode = 1958RSPSA.246..216B | s2cid = 137728239 }}</ref> For this reason, small amounts of nitrogen triiodide are sometimes synthesised as a demonstration to high school chemistry students or as an act of "chemical magic".<ref>{{cite book | author1 = Ford, L. A. | author2 = Grundmeier, E. W. | title = Chemical Magic | publisher = Dover | year = 1993 | page = [https://archive.org/details/chemicalmagic00ford_0/page/76 76] | isbn = 978-0-486-67628-9 | url-access = registration | url = https://archive.org/details/chemicalmagic00ford_0/page/76 }}</ref> [[Chlorine azide]] (ClN<sub>3</sub>) and [[bromine azide]] (BrN<sub>3</sub>) are extremely sensitive and explosive.<ref>{{ cite journal |author1=Frierson, W. J. |author2=Kronrad, J. |author3=Browne, A. W. | title = Chlorine Azide, ClN<sub>3</sub>. I | journal = [[Journal of the American Chemical Society]] | year = 1943 | volume = 65 | issue = 9 | pages = 1696–1698 | doi = 10.1021/ja01249a012 }}</ref><ref name="solid">{{cite journal|last=Lyhs|first=Benjamin|author2=Bläser, Dieter|author3=Wölper, Christoph|author4=Schulz, Stephan|author5=Jansen, Georg|title=Solid-State Structure of Bromine Azide|journal=Angewandte Chemie International Edition|date=20 February 2012|volume=51|issue=8|pages=1970–1974|doi=10.1002/anie.201108092|pmid=22250068|url=https://duepublico2.uni-due.de/servlets/MCRFileNodeServlet/duepublico_derivate_00073133/Accepted_Manuscript_Angew_Chem_Int_Ed_2012_51_1970.pdf|access-date=25 August 2021|archive-date=25 August 2021|archive-url=https://web.archive.org/web/20210825213503/https://duepublico2.uni-due.de/servlets/MCRFileNodeServlet/duepublico_derivate_00073133/Accepted_Manuscript_Angew_Chem_Int_Ed_2012_51_1970.pdf|url-status=live}}</ref>
Two series of nitrogen oxohalides are known: the nitrosyl halides (XNO) and the nitryl halides (XNO<sub>2</sub>). The first is very reactive gases that can be made by directly halogenating nitrous oxide. [[Nitrosyl fluoride]] (NOF) is colourless and a vigorous fluorinating agent. [[Nitrosyl chloride]] (NOCl) behaves in much the same way and has often been used as an ionising solvent. [[Nitrosyl bromide]] (NOBr) is red. The reactions of the nitryl halides are mostly similar: [[nitryl fluoride]] (FNO<sub>2</sub>) and [[nitryl chloride]] (ClNO<sub>2</sub>) are likewise reactive gases and vigorous halogenating agents.<ref name="Greenwood438" />
===Oxides===
{{main|Nitrogen oxide}}
[[File:Nitrogen dioxide at different temperatures.jpg|thumb|upright=1.36|right|Nitrogen dioxide at −196&nbsp;°C, 0&nbsp;°C, 23&nbsp;°C, 35&nbsp;°C, and 50&nbsp;°C. {{chem|NO|2}} converts to colourless dinitrogen tetroxide ({{chem|N|2|O|4}}) at low temperatures, and reverts to {{chem|NO|2}} at higher temperatures.]]
Nitrogen forms nine molecular oxides, some of which were the first gases to be identified: N<sub>2</sub>O ([[nitrous oxide]]), NO ([[nitric oxide]]), N<sub>2</sub>O<sub>3</sub> ([[dinitrogen trioxide]]), NO<sub>2</sub> ([[nitrogen dioxide]]), N<sub>2</sub>O<sub>4</sub> ([[dinitrogen tetroxide]]), N<sub>2</sub>O<sub>5</sub> ([[dinitrogen pentoxide]]), N<sub>4</sub>O ([[nitrosylazide]]),<ref name="Greenwood443">Greenwood and Earnshaw, pp. 443–58</ref> and N(NO<sub>2</sub>)<sub>3</sub> ([[trinitramide]]).<ref>{{cite journal |last1=Rahm |first1=Martin |last2=Dvinskikh |first2=Sergey V. |last3=Furó |first3=István |last4=Brinck |first4=Tore |date=23 December 2010 |title=Experimental Detection of Trinitramide, N(NO<sub>2</sub>)<sub>3</sub> |journal=Angewandte Chemie International Edition |volume=50 |issue=5 |pages=1145–48 |doi=10.1002/anie.201007047|pmid=21268214 |s2cid=32952729 }}</ref> All are thermally unstable towards decomposition to their elements. One other possible oxide that has not yet been synthesised is [[oxatetrazole]] (N<sub>4</sub>O), an aromatic ring.<ref name="Greenwood443" />
Nitrous oxide (N<sub>2</sub>O), better known as laughing gas, is made by thermal decomposition of molten [[ammonium nitrate]] at 250&nbsp;°C. This is a redox reaction and thus nitric oxide and nitrogen are also produced as byproducts. It is mostly used as a propellant and aerating agent for [[cream|sprayed canned whipped cream]], and was formerly commonly used as an anaesthetic. Despite appearances, it cannot be considered to be the [[inorganic anhydride|anhydride]] of [[hyponitrous acid]] (H<sub>2</sub>N<sub>2</sub>O<sub>2</sub>) because that acid is not produced by the dissolution of nitrous oxide in water. It is rather unreactive (not reacting with the halogens, the alkali metals, or [[ozone]] at room temperature, although reactivity increases upon heating) and has the unsymmetrical structure N–N–O (N≡N<sup>+</sup>O<sup>−</sup>↔<sup>−</sup>N=N<sup>+</sup>=O): above 600&nbsp;°C it dissociates by breaking the weaker N–O bond.<ref name="Greenwood443" />
Nitric oxide (NO) is the simplest stable molecule with an odd number of electrons. In mammals, including humans, it is an important cellular [[signalling molecule]] involved in many physiological and pathological processes.<ref>{{cite journal|pmid=10390607|year=1999|last1=Hou|first1=Y. C.|last2=Janczuk|first2=A.|last3=Wang|first3=P. G.|title=Current trends in the development of nitric oxide donors|volume=5|issue=6|pages=417–41|journal=Current Pharmaceutical Design|doi=10.2174/138161280506230110111042 }}</ref> It is formed by catalytic oxidation of ammonia. It is a colourless paramagnetic gas that, being thermodynamically unstable, decomposes to nitrogen and oxygen gas at 1100–1200&nbsp;°C. Its bonding is similar to that in nitrogen, but one extra electron is added to a ''π''* antibonding orbital and thus the bond order has been reduced to approximately 2.5; hence dimerisation to O=N–N=O is unfavourable except below the boiling point (where the ''cis'' isomer is more stable) because it does not actually increase the total bond order and because the unpaired electron is delocalised across the NO molecule, granting it stability. There is also evidence for the asymmetric red dimer O=N–O=N when nitric oxide is condensed with polar molecules. It reacts with oxygen to give brown nitrogen dioxide and with halogens to give nitrosyl halides. It also reacts with transition metal compounds to give nitrosyl complexes, most of which are deeply coloured.<ref name="Greenwood443" />
Blue dinitrogen trioxide (N<sub>2</sub>O<sub>3</sub>) is only available as a solid because it rapidly dissociates above its melting point to give nitric oxide, nitrogen dioxide (NO<sub>2</sub>), and dinitrogen tetroxide (N<sub>2</sub>O<sub>4</sub>). The latter two compounds are somewhat difficult to study individually because of the equilibrium between them, although sometimes dinitrogen tetroxide can react by heterolytic fission to [[nitrosonium]] and [[nitrate]] in a medium with high dielectric constant. Nitrogen dioxide is an acrid, corrosive brown gas. Both compounds may be easily prepared by decomposing a dry metal nitrate. Both react with water to form [[nitric acid]]. Dinitrogen tetroxide is very useful for the preparation of anhydrous metal nitrates and nitrato complexes, and it became the storable oxidiser of choice for many rockets in both the United States and [[USSR]] by the late 1950s. This is because it is a [[hypergolic propellant]] in combination with a [[hydrazine]]-based [[Rocket propellant|rocket fuel]] and can be easily stored since it is liquid at room temperature.<ref name="Greenwood443" />
The thermally unstable and very reactive dinitrogen pentoxide (N<sub>2</sub>O<sub>5</sub>) is the anhydride of [[nitric acid]], and can be made from it by dehydration with [[phosphorus pentoxide]]. It is of interest for the preparation of explosives.<ref>{{cite journal|author=Talawar, M. B.|title=Establishment of Process Technology for the Manufacture of Dinitrogen Pentoxide and its Utility for the Synthesis of Most Powerful Explosive of Today – CL-20|journal=Journal of Hazardous Materials|year= 2005| volume =124|issue=1–3| pages =153–64|doi=10.1016/j.jhazmat.2005.04.021|pmid=15979786|bibcode=2005JHzM..124..153T |display-authors=etal}}</ref> It is a [[deliquescent]], colourless crystalline solid that is sensitive to light. In the solid state it is ionic with structure [NO<sub>2</sub>]<sup>+</sup>[NO<sub>3</sub>]<sup>−</sup>; as a gas and in solution it is molecular O<sub>2</sub>N–O–NO<sub>2</sub>. Hydration to nitric acid comes readily, as does analogous reaction with [[hydrogen peroxide]] giving [[peroxonitric acid]] (HOONO<sub>2</sub>). It is a violent oxidising agent. Gaseous dinitrogen pentoxide decomposes as follows:<ref name="Greenwood443" />
:N<sub>2</sub>O<sub>5</sub> {{eqm}} NO<sub>2</sub> + NO<sub>3</sub> → NO<sub>2</sub> + O<sub>2</sub> + NO
:N<sub>2</sub>O<sub>5</sub> + NO {{eqm}} 3 NO<sub>2</sub>
===Oxoacids, oxoanions, and oxoacid salts===
Many nitrogen [[oxoacid]]s are known, though most of them are unstable as pure compounds and are known only as aqueous solutions or as salts. [[Hyponitrous acid]] (H<sub>2</sub>N<sub>2</sub>O<sub>2</sub>) is a weak diprotic acid with the structure HON=NOH (p''K''<sub>a1</sub> 6.9, p''K''<sub>a2</sub> 11.6). Acidic solutions are quite stable but above pH 4 base-catalysed decomposition occurs via [HONNO]<sup>−</sup> to nitrous oxide and the hydroxide anion. [[Hyponitrite]]s (involving the {{chem|N|2|O|2|2-}} anion) are stable to reducing agents and more commonly act as reducing agents themselves. They are an intermediate step in the oxidation of ammonia to nitrite, which occurs in the [[nitrogen cycle]]. Hyponitrite can act as a bridging or chelating bidentate ligand.<ref name="Greenwood459">Greenwood and Earnshaw, pp. 459–72</ref>
[[Nitrous acid]] (HNO<sub>2</sub>) is not known as a pure compound, but is a common component in gaseous equilibria and is an important aqueous reagent: its aqueous solutions may be made from acidifying cool aqueous [[nitrite]] ({{chem|NO|2|-}}, bent) solutions, although already at room temperature disproportionation to [[nitrate]] and nitric oxide is significant. It is a weak acid with p''K''<sub>''a''</sub> 3.35 at 18&nbsp;°C. They may be [[titration|titrimetrically]] analysed by their oxidation to nitrate by [[permanganate]]. They are readily reduced to nitrous oxide and nitric oxide by [[sulfur dioxide]], to hyponitrous acid with [[tin]](II), and to ammonia with [[hydrogen sulfide]]. Salts of [[hydrazinium]] {{chem|N|2|H|5|+}} react with nitrous acid to produce azides which further react to give nitrous oxide and nitrogen. [[Sodium nitrite]] is mildly toxic in concentrations above 100&nbsp;mg/kg, but small amounts are often used to cure meat and as a preservative to avoid bacterial spoilage. It is also used to synthesise hydroxylamine and to [[Diazonium compound#Preparation|diazotise]] primary aromatic amines as follows:<ref name="Greenwood459" />
:ArNH<sub>2</sub> + HNO<sub>2</sub> → [ArNN]Cl + 2 H<sub>2</sub>O
Nitrite is also a common ligand that can coordinate in five ways. The most common are nitro (bonded from the nitrogen) and nitrito (bonded from an oxygen). Nitro-nitrito isomerism is common, where the nitrito form is usually less stable.<ref name="Greenwood459" />
[[File:Fuming nitric acid 40ml.jpg|thumb|right|Fuming nitric acid contaminated with yellow nitrogen dioxide]]
[[Nitric acid]] (HNO<sub>3</sub>) is by far the most important and the most stable of the nitrogen oxoacids. It is one of the three most used acids (the other two being [[sulfuric acid]] and [[hydrochloric acid]]) and was first discovered by alchemists in the 13th century. It is made by the catalytic oxidation of ammonia to nitric oxide, which is oxidised to nitrogen dioxide, and then dissolved in water to give concentrated nitric acid. In the [[United States|United States of America]], over seven million tonnes of nitric acid are produced every year, most of which is used for nitrate production for fertilisers and explosives, among other uses. Anhydrous nitric acid may be made by distilling concentrated nitric acid with phosphorus pentoxide at low pressure in glass apparatus in the dark. It can only be made in the solid state, because upon melting it spontaneously decomposes to nitrogen dioxide, and liquid nitric acid undergoes [[Molecular autoionization|self-ionisation]] to a larger extent than any other covalent liquid as follows:<ref name="Greenwood459" />
:2 HNO<sub>3</sub> {{eqm}} {{chem|H|2|NO|3|+}} + {{chem|NO|3|-}} {{eqm}} H<sub>2</sub>O + [NO<sub>2</sub>]<sup>+</sup> + [NO<sub>3</sub>]<sup>−</sup>
Two hydrates, HNO<sub>3</sub>·H<sub>2</sub>O and HNO<sub>3</sub>·3H<sub>2</sub>O, are known that can be crystallised. It is a strong acid and concentrated solutions are strong oxidising agents, though [[gold]], [[platinum]], [[rhodium]], and [[iridium]] are immune to attack. A 3:1 mixture of concentrated hydrochloric acid and nitric acid, called ''[[aqua regia]]'', is still stronger and successfully dissolves gold and platinum, because free chlorine and nitrosyl chloride are formed and chloride anions can form strong complexes. In concentrated sulfuric acid, nitric acid is protonated to form [[nitronium]], which can act as an electrophile for aromatic nitration:<ref name="Greenwood459" />
:HNO<sub>3</sub> + 2 H<sub>2</sub>SO<sub>4</sub> {{eqm}} {{chem|NO|2|+}} + H<sub>3</sub>O<sup>+</sup> + 2 {{chem|HSO|4|-}}
The thermal stabilities of [[nitrate]]s (involving the trigonal planar {{chem|NO|3|-}} anion) depends on the basicity of the metal, and so do the products of decomposition (thermolysis), which can vary between the nitrite (for example, sodium), the oxide (potassium and [[lead]]), or even the metal itself ([[silver]]) depending on their relative stabilities. Nitrate is also a common ligand with many modes of coordination.<ref name="Greenwood459" />
Finally, although orthonitric acid (H<sub>3</sub>NO<sub>4</sub>), which would be analogous to [[orthophosphoric acid]], does not exist, the tetrahedral [[orthonitrate]] anion {{chem|NO|4|3-}} is known in its sodium and potassium salts:<ref name="Greenwood459" />
:<chem>NaNO3{} + Na2O ->[\ce{Ag~crucible}][\ce{300^\circ C~for~7 days}] Na3NO4</chem>
These white crystalline salts are very sensitive to water vapour and carbon dioxide in the air:<ref name="Greenwood459" />
:Na<sub>3</sub>NO<sub>4</sub> + H<sub>2</sub>O + CO<sub>2</sub> → NaNO<sub>3</sub> + NaOH + NaHCO<sub>3</sub>
Despite its limited chemistry, the orthonitrate anion is interesting from a structural point of view due to its regular tetrahedral shape and the short N–O bond lengths, implying significant polar character to the bonding.<ref name="Greenwood459" />
===Organic nitrogen compounds===
Nitrogen is one of the most important elements in [[organic chemistry]]. Many organic [[functional group]]s involve a [[carbon–nitrogen bond]], such as [[amide]]s (RCONR<sub>2</sub>), [[amine]]s (R<sub>3</sub>N), [[imine]]s (RC(=NR)R), [[imide]]s (RCO)<sub>2</sub>NR, [[azide]]s (RN<sub>3</sub>), [[azo compound]]s (RN<sub>2</sub>R), [[Cyanate ester|cyanate]]s (ROCN), [[isocyanate]]s (RNCO), [[nitrate]]s (RONO<sub>2</sub>), [[nitrile]]s (RCN), [[isonitrile]]s (RNC), [[Nitrite#Organic nitrites|nitrite]]s (RONO), [[nitro compound]]s (RNO<sub>2</sub>), [[nitroso compound]]s (RNO), [[oxime]]s (RC(=NOH)R), and [[pyridine]] derivatives. C–N bonds are strongly polarised towards nitrogen. In these compounds, nitrogen is usually trivalent (though it can be tetravalent in [[quaternary ammonium salt]]s, R<sub>4</sub>N<sup>+</sup>), with a lone pair that can confer basicity on the compound by being coordinated to a proton. This may be offset by other factors: for example, amides are not basic because the lone pair is delocalised into a double bond (though they may act as bases at very low pH, being protonated at the oxygen), and [[pyrrole]] is not basic because the lone pair is delocalised as part of an [[aromaticity|aromatic]] ring.<ref name="Jerry">{{JerryMarch}}</ref> The amount of nitrogen in a [[chemical substance]] can be determined by the [[Kjeldahl method]].<ref>{{cite book|year=2008|doi=10.1007/978-1-4020-6754-9_9066|isbn=978-1-4020-6753-2|title=Encyclopedia of Genetics, Genomics, Proteomics and Informatics|page=1063|chapter=Kjeldahl Method|last1=Rédei|first1=George P}}</ref> In particular, nitrogen is an essential component of [[nucleic acids]], [[amino acid]]s and thus [[protein]]s, and the energy-carrying molecule [[adenosine triphosphate]] and is thus vital to all life on Earth.<ref name="Jerry" />
==Occurrence==
[[File:Nitrogen Cycle.svg|thumb|right|upright=1.8|Schematic representation of the [[Nitrogen cycle|flow of nitrogen compounds]] through a land environment]]
{{see also|Nitrogen cycle}}
Nitrogen is the most common pure element in the earth, making up 78.1% of the volume of the [[Atmosphere of Earth|atmosphere]]<ref name="Greenwood406" /> (75.5% by mass), around 3.89 million [[gigatonnes]] ({{val|3.89e18|u=kg}}). Despite this, it is [[Abundance of elements in Earth's crust|not very abundant]] in Earth's crust, making up somewhere around 19&nbsp;[[parts per million]] of this, on par with [[niobium]], [[gallium]], and [[lithium]]. (This represents 300,000 to a million gigatonnes of nitrogen, depending on the mass of the crust.<ref>Depending on the average thickness which is somewhere between 10 and 30 km, the mass of the earth's crust is between {{val|15|e=18}} and {{val|45|e=18}} tonnes.</ref>) The only important nitrogen minerals are [[nitre]] ([[potassium nitrate]], saltpetre) and [[sodanitre|soda nitre]] ([[sodium nitrate]], Chilean saltpetre). However, these have not been an important source of nitrates since the 1920s, when the industrial synthesis of ammonia and nitric acid became common.<ref name="Greenwood407">Greenwood and Earnshaw, pp. 407–09</ref>
Nitrogen compounds constantly interchange between the atmosphere and living organisms. Nitrogen must first be processed, or "[[nitrogen fixation|fixed]]", into a plant-usable form, usually ammonia. Some nitrogen fixation is done by lightning strikes producing the nitrogen oxides, but most is done by [[diazotroph]]ic bacteria through enzymes known as [[nitrogenase]]s (although today industrial nitrogen fixation to ammonia is also significant). When the ammonia is taken up by plants, it is used to synthesise proteins. These plants are then digested by animals who use the nitrogen compounds to synthesise their proteins and excrete nitrogen-bearing waste. Finally, these organisms die and decompose, undergoing bacterial and environmental oxidation and [[denitrification]], returning free dinitrogen to the atmosphere. Industrial nitrogen fixation by the [[Haber process]] is mostly used as fertiliser, although excess nitrogen–bearing waste, when leached, leads to [[eutrophication]] of freshwater and the creation of marine [[Dead zone (ecology)|dead zones]], as nitrogen-driven bacterial growth depletes water oxygen to the point that all higher organisms die. Furthermore, nitrous oxide, which is produced during denitrification, attacks the atmospheric [[ozone layer]].<ref name="Greenwood407" />
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 odour in unfresh saltwater fish.<ref>{{Cite journal|pmid=15186102|date=Jun 2004|author=Nielsen, M. K.|author2=Jørgensen, B. M.|title=Quantitative relationship between trimethylamine oxide aldolase activity and formaldehyde accumulation in white muscle from gadiform fish during frozen storage|volume=52|issue=12|pages=3814–22|doi=10.1021/jf035169l|journal=Journal of Agricultural and Food Chemistry}}</ref> In animals, [[free radical]] [[nitric oxide]] (derived from an [[amino acid]]), serves as an important regulatory molecule for circulation.<ref name="ocean">{{Cite book|url=https://books.google.com/books?id=qPuCVvAv2GwC&pg=PA392|page=392|title=Biology of the Southern Ocean|author=Knox, G. A.|publisher=CRC Press|date=2007|isbn=978-0-8493-3394-1|access-date=2020-08-24|archive-date=2021-10-01|archive-url=https://web.archive.org/web/20211001035610/https://books.google.com/books?id=qPuCVvAv2GwC&pg=PA392|url-status=live}}</ref>
Nitric oxide's rapid reaction with water in animals results in the production of its metabolite [[nitrite]]. Animal [[metabolism]] of nitrogen in proteins, in general, results in the [[excretion]] of [[urea]], while animal metabolism of [[nucleic acids]] results in the excretion of [[urea]] and [[uric acid]]. The characteristic odour of animal flesh decay is caused by the creation of long-chain, nitrogen-containing [[amine]]s, such as [[putrescine]] and [[cadaverine]], which are breakdown products of the amino acids [[ornithine]] and [[lysine]], respectively, in decaying proteins.<ref>{{Cite book|url=https://books.google.com/books?id=RMPis9OnJxkC&pg=PT121|page=121|title=Digestion, diet, and disease: irritable bowel syndrome and gastrointestinal function|author=Vickerstaff Joneja|author2=Janice M.|publisher=Rutgers University Press|date=2004|isbn=978-0-8135-3387-2|access-date=2020-08-24|archive-date=2021-10-01|archive-url=https://web.archive.org/web/20211001035605/https://books.google.com/books?id=RMPis9OnJxkC&pg=PT121|url-status=live}}</ref>
==Production==
Nitrogen gas is an [[industrial gas]] produced by the [[fractional distillation]] of [[liquid air]], or by mechanical means using gaseous air (pressurised reverse [[Osmotic pressure|osmosis membrane]] or [[pressure swing adsorption]]). Nitrogen gas generators using membranes or pressure swing adsorption (PSA) are typically more cost and energy efficient than bulk-delivered nitrogen.<ref>{{Cite web|url=http://www.parkern2.com/_literature_176619/A_Sustainable_Approach_to_the_Supply_of_Nitrogen|title=A Sustainable Approach to the Supply of Nitrogen|last=Froehlich|first=Peter|date=May 2013|website=www.parker.com|publisher=Parker Hannifin Corporation|access-date=24 November 2016|archive-date=16 March 2016|archive-url=https://web.archive.org/web/20160316190914/http://www.parkern2.com/_literature_176619/A_Sustainable_Approach_to_the_Supply_of_Nitrogen|url-status=live}}</ref> 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).<ref>{{Cite journal|doi= 10.1021/ie50569a032|title=Nitrogen Purfication. Pilot Plant Removal of Oxygen|year=1957|journal=Industrial & Engineering Chemistry|volume= 49|pages= 869–73|issue= 5|last1= Reich|first1= Murray|last2= Kapenekas|first2= Harry}}</ref> Commercial-grade nitrogen already contains at most 20&nbsp;ppm oxygen, and specially purified grades containing at most 2&nbsp;ppm oxygen and 10&nbsp;ppm [[argon]] are also available.<ref name="Greenwood409">Greenwood and Earnshaw, pp. 409–11</ref>
In a chemical laboratory, it is prepared by treating an aqueous solution of [[ammonium chloride]] with [[sodium nitrite]].<ref name="labProduction">{{Cite journal| last1 = Bartlett |first1 = J. K.| title = Analysis for nitrite by evolution of nitrogen: A general chemistry laboratory experiment | doi = 10.1021/ed044p475 | journal = Journal of Chemical Education | volume = 44 | issue = 8 | page = 475 | year = 1967 | bibcode = 1967JChEd..44..475B}}</ref>
:NH<sub>4</sub>Cl + NaNO<sub>2</sub> → N<sub>2</sub> + NaCl + 2 H<sub>2</sub>O
Small amounts of the impurities NO and HNO<sub>3</sub> are also formed in this reaction. The impurities can be removed by passing the gas through aqueous sulfuric acid containing [[potassium dichromate]].<ref name="labProduction" />
It can also be obtained by the thermal decomposition of [[ammonium dichromate]].<ref>{{cite journal | last1=Mahieu | first1=B. | last2=Apers | first2=D.J. | last3=Capron | first3=P.C. | title=Thermal decomposition of ammonium dichromate | journal=Journal of Inorganic and Nuclear Chemistry | publisher=Elsevier BV | volume=33 | issue=9 | year=1971 | issn=0022-1902 | doi=10.1016/0022-1902(71)80047-7 | pages=2857–2866}}</ref>
:3(NH<sub>4</sub>)<sub>2</sub>Cr<sub>2</sub>O<sub>7</sub> → 2N<sub>2</sub> + 9H<sub>2</sub>O + 3Cr<sub>2</sub>O<sub>3</sub> + 2NH<sub>3</sub> + 32O<sub>2</sub>
Very pure nitrogen can be prepared by the thermal decomposition of [[barium azide]] or [[sodium azide]].<ref>{{Cite journal | last1 = Eremets | first1 = M. I. | last2 = Popov | first2 = M. Y. | last3 = Trojan | first3 = I. A. | last4 = Denisov | first4 = V. N. | last5 = Boehler | first5 = R. | last6 = Hemley | first6 = R. J. | doi = 10.1063/1.1718250 | title = Polymerization of nitrogen in sodium azide | journal = The Journal of Chemical Physics | volume = 120 | issue = 22 | pages = 10618–23 | year = 2004 | pmid = 15268087|bibcode = 2004JChPh.12010618E }}</ref>
:2 NaN<sub>3</sub> → 2 Na + 3 N<sub>2</sub>
==Applications==
The applications of nitrogen compounds are naturally extremely widely varied due to the huge size of this class: hence, only applications of pure nitrogen itself will be considered here. Two-thirds (2/3) of nitrogen produced by industry is sold as gas and the remaining one-third (1/3) as a liquid.
===Gas{{anchor|Nitrogen_gas}}===
The gas is mostly used as a low reactivity safe atmosphere wherever the oxygen in the air would pose a fire, explosion, or oxidising hazard. Some examples include:<ref name="Greenwood409" />
* As a [[modified atmosphere]], pure or mixed with [[carbon dioxide]], to nitrogenate and preserve the freshness of packaged or bulk foods (by delaying [[Rancidification|rancidity]] and other forms of [[Redox|oxidative damage]]). Pure nitrogen as food additive is labelled in the [[European Union]] with the [[E number]] E941.<ref>{{Cite book | url = https://books.google.com/books?id=Fvm-sqd90-oC&pg=PA591 | page = 591 | title = Food Additives in Europe 2000 | isbn = 978-92-893-0829-8 | year = 2002 | last1 = Ministers | first1 = Nordic Council of | publisher = Nordic Council of Ministers | access-date = 2015-12-20 | archive-date = 2016-02-05 | archive-url = https://web.archive.org/web/20160205191759/https://books.google.com/books?id=Fvm-sqd90-oC&pg=PA591 | url-status = live }}</ref>
* In [[incandescent light bulb]]s as an inexpensive alternative to [[argon]].<ref>{{Cite book|url=https://books.google.com/books?id=W0HW8wgmQQsC&pg=PA90|isbn=978-0-85404-690-4|editor=Harding, Charlie|date=2002|publisher=Royal Society of Chemistry|location=Cambridge|title=Elements of the p Block|access-date=2020-08-24|archive-date=2021-10-01|archive-url=https://web.archive.org/web/20211001035604/https://books.google.com/books?id=W0HW8wgmQQsC&pg=PA90|url-status=live}}</ref>
* In [[Gaseous fire suppression|fire suppression systems for Information technology (IT) equipment]].<ref name="Greenwood409" />
* In the manufacture of [[stainless steel]].<ref>{{Cite book|title=High nitrogen steels: structure, properties, manufacture, applications|author=Gavriliuk, V. G.|author2=Berns, Hans|url=https://books.google.com/books?id=6eF4AfEwF4YC&pg=PA338|publisher=Springer|date=1999|isbn=978-3-540-66411-6|access-date=2020-08-24|archive-date=2021-10-01|archive-url=https://web.archive.org/web/20211001035605/https://books.google.com/books?id=6eF4AfEwF4YC&pg=PA338|url-status=live}}</ref>
* In the [[case-hardening]] of steel by [[nitriding]].<ref name="Duplex">{{cite journal|doi=10.1179/1743284715Y.0000000098|last1=Meka|first1=S. R.|last2=Chauhan|first2=A.|last3=Steiner|first3=T.|last4=Bischoff|first4=E.|last5=Ghosh|first5=P. K.|last6=Mittemeijer|first6=E. J.|year=2015|title=Generating duplex microstructures by nitriding; nitriding of iron based Fe–Mn alloy|page=1743284715Y.000|journal=Materials Science and Technology|volume=32 |issue=9 |doi-access=free}}</ref>
* In some aircraft fuel systems to reduce fire hazard (see [[inerting system]]).
* To inflate race car and aircraft [[tire]]s,<ref>{{cite web|url=http://auto.howstuffworks.com/question594.htm|title=Why don't they use normal air in race car tires?|publisher=Howstuffworks|access-date=2006-07-22|date=2001-03-16|archive-date=2011-07-12|archive-url=https://web.archive.org/web/20110712225925/http://auto.howstuffworks.com/question594.htm|url-status=live}}</ref> reducing the problems of inconsistent expansion and contraction caused by moisture and [[Redox|oxygen]] in natural air.<ref name="Greenwood409" />
Nitrogen is commonly used during sample preparation in [[chemical analysis]]. It is used to concentrate and reduce the volume of liquid samples. Directing a pressurised stream of nitrogen gas perpendicular to the surface of the liquid causes the solvent to evaporate while leaving the solute(s) and un-evaporated solvent behind.<ref>{{Cite journal|doi=10.1016/S0021-9673(01)01466-2|title=Centrifugal concentrator for the substitution of nitrogen blow-down micro-concentration in dioxin/polychlorinated biphenyl sample preparation|year=2002|author=Kemmochi, Y|journal=Journal of Chromatography A|volume=943|pages=295–97|pmid=11833649|last2=Tsutsumi|first2=K.|last3=Arikawa|first3=A.|last4=Nakazawa|first4=H.|issue=2}}</ref>
Nitrogen can be used as a replacement, or in combination with, [[carbon dioxide]] to pressurise kegs of some [[beer]]s, particularly [[stout]]s and British [[ale]]s, due to the smaller [[liquid bubble|bubbles]] it produces, which makes the dispensed beer smoother and [[beer head|headier]].<ref>{{cite book|last1=Baxter|first1=E. Denise|last2=Hughes|first2=Paul S.|title=Beer: Quality, Safety and Nutritional Aspects|date=2001|publisher=Royal Society of Chemistry|isbn=978-0-85404-588-4|page=22|url=https://books.google.com/books?id=X95O-1_x-a0C&pg=PA22|access-date=2015-06-20|archive-date=2020-03-21|archive-url=https://web.archive.org/web/20200321190342/https://books.google.com/books?id=X95O-1_x-a0C&pg=PA22|url-status=live}}</ref> A pressure-sensitive nitrogen capsule known commonly as a "[[widget (beer)|widget]]" allows nitrogen-charged beers to be packaged in [[Beverage can|cans]] and [[bottle]]s.<ref>{{cite web|url=http://recipes.howstuffworks.com/question446.htm|title=How does the widget in a beer can work?|publisher=Howstuffworks|date=2000-08-16|access-date=2008-07-30|archive-date=2007-11-02|archive-url=https://web.archive.org/web/20071102214918/http://recipes.howstuffworks.com/question446.htm|url-status=live}}</ref><ref>{{cite book| url = https://books.google.com/books?id=ogXc9xTTCVsC&pg=PA131| page = 131| title = Froth!: The Science of Beer| isbn = 978-0-8018-9569-2| last1 = Denny| first1 = Mark| date = 1 November 2009| publisher = Johns Hopkins University Press+ORM| access-date = 20 December 2015| archive-date = 5 February 2016| archive-url = https://web.archive.org/web/20160205191759/https://books.google.com/books?id=ogXc9xTTCVsC&pg=PA131| url-status = live}}</ref> Nitrogen tanks are also replacing carbon dioxide as the main power source for [[paintball gun]]s. Nitrogen must be kept at a higher pressure than CO<sub>2</sub>, making N<sub>2</sub> tanks heavier and more expensive.<ref>{{cite thesis| hdl = 1721.1/45820| last=Kennett |first=Andrew J. |title =Design of a pneumatically assisted shifting system for Formula SAE racing applications|year=2008|publisher=Dept. of Mechanical Engineering, Massachusetts Institute of Technology. | type=Thesis }}</ref>
==== Equipment ====
Some construction equipment uses pressurised nitrogen gas to help [[Hydraulics|hydraulic system]] to provide extra power to devices such as [[Breaker (hydraulic)|hydraulic hammer]].  Nitrogen gas, formed from the decomposition of [[sodium azide]], is used for the inflation of [[airbag]]s.<ref>{{cite journal | author = Betterton, E. A. | title = Environmental Fate of Sodium Azide Derived from Automobile Airbags | journal = Critical Reviews in Environmental Science and Technology | year = 2003 | volume = 33 | issue = 4 | pages = 423–58 | doi = 10.1080/10643380390245002 | bibcode = 2003CREST..33..423B | s2cid = 96404307 }}</ref>
==== Execution ====
{{see|Inert gas asphyxiation}}
As nitrogen is an asphyxiant gas in itself, some jurisdictions have considered asphyxiation by inhalation of pure nitrogen as a means of [[capital punishment]] (as a substitute for [[lethal injection]]).<ref name="Time-apr2015">{{cite news |last=Sanburn |first=Josh |url=https://time.com/3749879/nitrogen-gas-execution-oklahoma-lethal-injection/ |title=The Dawn of a New Form of Capital Punishment |magazine=[[Time (magazine)|Time]] |date=2015-04-10 |access-date=2015-04-11 |archive-date=2015-04-11 |archive-url=https://web.archive.org/web/20150411201357/http://time.com/3749879/nitrogen-gas-execution-oklahoma-lethal-injection/ |url-status=live }}</ref><ref>{{cite news |url=http://www.abc.net.au/7.30/content/2012/s3657550.htm |title=Euthanasia campaigner under scrutiny |author=Sexton, Mike |date=18 December 2012 |work=[[ABC News (Australia)|ABC]] |access-date=6 May 2013 |archive-date=7 July 2013 |archive-url=https://web.archive.org/web/20130707143023/http://www.abc.net.au/7.30/content/2012/s3657550.htm |url-status=live }}</ref><ref>{{Cite news|url=https://www.washingtonpost.com/news/post-nation/wp/2015/04/17/oklahoma-says-it-will-now-use-nitrogen-gas-as-its-backup-method-of-execution/|title=Oklahoma says it will now use nitrogen gas as its backup method of execution|last=Berman|first=Mark|date=April 17, 2015|newspaper=The Washington Post|access-date=June 22, 2019|archive-date=June 23, 2019|archive-url=https://web.archive.org/web/20190623024008/https://www.washingtonpost.com/news/post-nation/wp/2015/04/17/oklahoma-says-it-will-now-use-nitrogen-gas-as-its-backup-method-of-execution/|url-status=live}}</ref> In January 2024, [[Execution of Kenneth Eugene Smith|Kenneth Eugene Smith]] became the first person executed by nitrogen asphyxiation.<ref>{{cite web |title=Kenneth Eugene Smith executed by nitrogen hypoxia in Alabama, marking a first for the death penalty |url=https://www.cbsnews.com/news/kenneth-eugene-smith-executed-by-nitrogen-hypoxia-in-alabama/ |publisher=[[CBS News]] |date=26 January 2024}}</ref>
===Liquid===
{{main |Liquid nitrogen}}
[[File:Nitrogen.ogv|thumb|Air balloon submerged in liquid nitrogen]]
Liquid nitrogen is a [[cryogen|cryogenic liquid]] which looks like [[water]]. When insulated in proper containers such as [[dewar flask]]s, it can be transported and stored with a low rate of [[evaporation|evaporative loss]].<ref>{{Cite journal|doi=10.1007/BF01136404|title=Vessels for the storage and transport of liquid oxygen and nitrogen|author=Kaganer, M. G.|author2=Kozheurov, V.|author3=Levina, Zh. L.|name-list-style=amp|journal=Chemical and Petroleum Engineering|volume=3|issue=12|year=1967|pages=918–22|bibcode=1967CPE.....3..918K |s2cid=96762552}}</ref>
[[File:A cylinder container, containing liquid nitrogen.jpg|thumb|A container vehicle carrying liquid nitrogen]]
Like [[dry ice]], the main use of liquid nitrogen is for cooling to low temperatures. It is used in the [[cryopreservation]] of biological materials such as blood and reproductive cells ([[sperm]] and [[ovum|eggs]]). It is used in [[cryotherapy]] to remove cysts and warts on the skin by freezing them.<ref>{{cite journal|pmid=11359389|volume=144|issue=5|title=Liquid nitrogen cryotherapy of common warts: cryo-spray vs. cotton wool bud|date=May 2001|journal=Br. J. Dermatol.|pages=1006–09|author=Ahmed I|author2=Agarwal S|author3=Ilchyshyn A|author4=Charles-Holmes S|author5=Berth-Jones J|doi=10.1046/j.1365-2133.2001.04190.x|s2cid=221325640}}</ref> It is used in laboratory [[cold trap]]s, and in [[cryopump]]s to obtain lower pressures in [[vacuum pump]]ed systems. It is used to cool heat-sensitive electronics such as [[infrared detector]]s and [[X-ray detector]]s. Other uses include freeze-grinding and machining materials that are soft or rubbery at room temperature, shrink-fitting and assembling engineering components, and more generally to attain very low temperatures where necessary. Because of its low cost, liquid nitrogen is often used for cooling even when such low temperatures are not strictly necessary, such as refrigeration of food, [[freeze-branding]] livestock, freezing pipes to halt flow when valves are not present, and consolidating unstable soil by freezing whenever excavation is going on underneath.<ref name="Greenwood409" />
==Safety==
===Gas===
Although nitrogen is non-toxic, when released into an enclosed space it can displace oxygen, and therefore presents an [[Nitrogen asphyxiation|asphyxiation]] hazard. This may happen with few warning symptoms, since the human [[carotid body]] is a relatively poor and slow low-oxygen (hypoxia) sensing system.<ref>{{cite web|url=http://www.bath.ac.uk/internal/bio-sci/bbsafe/asphyx.htm |title=Biology Safety&nbsp;– Cryogenic materials. The risks posed by them |publisher=University of Bath |access-date=2007-01-03 |url-status=dead |archive-url=https://web.archive.org/web/20070206095504/http://www.bath.ac.uk/internal/bio-sci/bbsafe/asphyx.htm |archive-date=February 6, 2007 }}</ref> An example occurred shortly before the launch of the [[STS-1|first Space Shuttle mission]] on March 19, 1981, when two technicians died from asphyxiation after they walked into a space located in the [[Mobile launcher platform#Saturn V and Space Shuttle|Space Shuttle's mobile launcher platform]] that was pressurised with pure nitrogen as a precaution against fire.<ref>{{cite news| title = Space Shuttle Columbia Fast Facts| url = http://www.cnn.com/2013/09/30/us/space-shuttle-columbia-fast-facts| date = September 30, 2013| publisher = CNN| access-date = January 20, 2016| archive-date = February 2, 2016| archive-url = https://web.archive.org/web/20160202181327/http://www.cnn.com/2013/09/30/us/space-shuttle-columbia-fast-facts| url-status = live}}</ref>
When inhaled at high [[partial pressure]]s (more than about 4 bar, encountered at depths below about 30&nbsp;m in [[scuba diving]]), nitrogen is an anaesthetic agent, causing [[nitrogen narcosis]], a temporary state of mental impairment similar to [[nitrous oxide]] intoxication.<ref>{{Cite journal|last1=Fowler|first1=B.|last2=Ackles|first2=K. N.|last3=Porlier|first3=G.|title=Effects of inert gas narcosis on behavior – a critical review|journal=Undersea Biomed. Res.|volume=12|issue=4|pages=369–402|year=1985|pmid=4082343|url=http://archive.rubicon-foundation.org/3019|access-date=2008-09-21|archive-url=https://web.archive.org/web/20101225052236/http://archive.rubicon-foundation.org/3019|archive-date=2010-12-25|url-status=usurped}}</ref><ref>{{Cite journal|author=Rogers, W. H.|author2=Moeller, G.|title=Effect of brief, repeated hyperbaric exposures on susceptibility to nitrogen narcosis|journal=Undersea Biomed. Res.|volume=16|issue=3|pages=227–32|year=1989|oclc=2068005|pmid=2741255|url=http://archive.rubicon-foundation.org/2522|access-date=2008-09-21|archive-url=https://web.archive.org/web/20090901020853/http://archive.rubicon-foundation.org/2522|archive-date=2009-09-01|url-status=usurped}}</ref>
Nitrogen dissolves in the [[blood]] and body fats. Rapid decompression (as when divers ascend too quickly or astronauts decompress too quickly from cabin pressure to spacesuit pressure) can lead to a potentially fatal condition called [[decompression sickness]] (formerly known as caisson sickness or ''the bends''), when nitrogen bubbles form in the bloodstream, nerves, joints, and other sensitive or vital areas.<ref name="DCShx">{{Cite journal|last=Acott|first=C.|title=A brief history of diving and decompression illness|journal=South Pacific Underwater Medicine Society Journal|volume=29|issue=2|year=1999|oclc=16986801|url=http://archive.rubicon-foundation.org/6004|access-date=2008-09-21|archive-date=2011-09-05|archive-url=https://web.archive.org/web/20110905152645/http://archive.rubicon-foundation.org/6004|url-status=usurped}}</ref><ref name="Kindwall">{{Cite journal|last1=Kindwall|first1=E. P.|last2=Baz|first2=A.|last3=Lightfoot|first3=E. N.|last4=Lanphier|first4=E. H.|last5=Seireg|first5=A.|title=Nitrogen elimination in man during decompression|journal=Undersea Biomed. Res.|volume=2|issue=4|pages=285–97|year=1975|oclc=2068005|pmid=1226586|url=http://archive.rubicon-foundation.org/2741|access-date=2008-09-21|archive-url=https://web.archive.org/web/20110727224419/http://archive.rubicon-foundation.org/2741|archive-date=2011-07-27|url-status=usurped}}</ref> Bubbles from other "inert" gases (gases other than carbon dioxide and oxygen) cause the same effects, so replacement of nitrogen in [[breathing gas]]es may prevent nitrogen narcosis, but does not prevent decompression sickness.<ref name="usn1">{{Cite book|title=US Navy Diving Manual, 6th revision|date=2006|publisher=US Naval Sea Systems Command|location=United States|url=http://www.supsalv.org/00c3_publications.asp?destPage=00c3&pageID=3.9|access-date=2008-04-24|archive-date=2008-05-02|archive-url=https://web.archive.org/web/20080502023541/http://www.supsalv.org/00c3_publications.asp?destPage=00c3&pageId=3.9|url-status=live}}</ref>
===Liquid===
As a [[cryogenic]] liquid, liquid nitrogen can be dangerous by causing [[cold burn]]s on contact, although the [[Leidenfrost effect]] provides protection for very short exposure (about one second).<ref>{{cite journal|last1=Walker|first1=Jearl|title=Boiling and the Leidenfrost Effect|journal=Fundamentals of Physics|pages=1–4|url=http://www.wiley.com/college/phy/halliday320005/pdf/leidenfrost_essay.pdf|access-date=11 October 2014|archive-date=13 December 2019|archive-url=https://web.archive.org/web/20191213010916/https://www.wiley.com/college/phy/halliday320005/pdf/leidenfrost_essay.pdf|url-status=live}}</ref> Ingestion of liquid nitrogen can cause severe internal damage. For example, in 2012, a young woman in England had to have her stomach removed after ingesting a cocktail made with liquid nitrogen.<ref>{{cite web|url=https://www.bbc.co.uk/news/uk-19878511 |title=Liquid nitrogen cocktail leaves teen in hospital|date=2012-10-08|publisher=BBC News|archiveurl=https://web.archive.org/web/20170412232656/http://www.bbc.co.uk/news/uk-19878511|archivedate=2017-04-12 }}</ref>
Because the liquid-to-gas [[expansion ratio]] of nitrogen is 1:694 at 20&nbsp;°C, a tremendous amount of force can be generated if liquid nitrogen is rapidly vaporised in an enclosed space. In an incident on January 12, 2006, at [[Texas A&M University]], the pressure-relief devices of a tank of liquid nitrogen were malfunctioning and later sealed. As a result of the subsequent pressure buildup, the tank failed catastrophically. The force of the explosion was sufficient to propel the tank through the ceiling immediately above it, shatter a reinforced concrete beam immediately below it, and blow the walls of the laboratory 0.1–0.2&nbsp;m off their foundations.<ref>{{cite web |title=Investigative Report on Chemistry 301A Cylinder Explosion |author=Mattox, Brent S. |publisher=Texas A&M University |format=reprint |url=http://ucih.ucdavis.edu/docs/chemistry_301a.pdf |url-status=dead |archive-url=https://web.archive.org/web/20140430205659/http://ucih.ucdavis.edu/docs/chemistry_301a.pdf |archive-date=2014-04-30 }}</ref>
Liquid nitrogen readily evaporates to form gaseous nitrogen, and hence the precautions associated with gaseous nitrogen also apply to liquid nitrogen.<ref name="BCGACOPCP30">British Compressed Gases Association (2000) BCGA Code of Practice CP30. {{cite web|url=http://www.bcga.co.uk/preview/products.php?g1=3ff921&n=2 |title=The Safe Use of Liquid nitrogen Dewars up to 50 litres.|archive-url=https://web.archive.org/web/20070718050900/http://www.bcga.co.uk/preview/products.php?g1=3ff921&n=2 |archive-date=2007-07-18 |issn=0260-4809}}.</ref><ref>{{cite web|url=http://www.csb.gov/assets/1/19/Valero_Case_Study.pdf |title=Confined Space Entry – Worker and Would-be Rescuer Asphyxiated|archive-url=https://web.archive.org/web/20150922022123/http://www.csb.gov/assets/1/19/Valero_Case_Study.pdf |archive-date=2015-09-22 |publisher= Valero Refinery Asphyxiation Incident Case Study.}}</ref><ref>{{cite web|url=http://news.bbc.co.uk/2/hi/uk_news/scotland/484813.stm |title=Inquiry after man dies in chemical leak|archive-url=https://web.archive.org/web/20170107063950/http://news.bbc.co.uk/2/hi/uk_news/scotland/484813.stm |archive-date=2017-01-07|publisher=BBC News|date=1999-10-25}}</ref> For example, [[oxygen sensor]]s are sometimes used as a safety precaution when working with liquid nitrogen to alert workers of gas spills into a confined space.<ref name="usn">{{Cite book |title=Liquid Nitrogen – Code of practice for handling |year=2007 |publisher=Birkbeck, University of London |location=United Kingdom |url=http://www.bbk.ac.uk/so/policies/liqn2 |access-date=2012-02-08 |archive-date=2018-06-12 |archive-url=https://web.archive.org/web/20180612141434/http://www.bbk.ac.uk/so/policies/liqn2 |url-status=live }}</ref>
Vessels containing liquid nitrogen can [[liquid oxygen|condense oxygen]] from air. The liquid in such a vessel becomes increasingly enriched in oxygen (boiling point −183&nbsp;°C, higher than that of nitrogen) as the nitrogen evaporates, and can cause violent oxidation of organic material.<ref>{{Cite web|title = Liquid Nitrogen Safety|author = Levey, Christopher G.|publisher = Thayer School of Engineering at Dartmouth|url = http://engineering.dartmouth.edu/microeng/ln2.html|access-date = 2016-11-23|archive-date = 2016-03-05|archive-url = https://web.archive.org/web/20160305020126/http://engineering.dartmouth.edu/microeng/ln2.html|url-status = live}}</ref>
===Oxygen deficiency monitors===
Oxygen deficiency monitors are used to measure levels of oxygen in confined spaces and any place where nitrogen gas or liquid are stored or used. In the event of a nitrogen leak, and a decrease in oxygen to a pre-set alarm level, an oxygen deficiency monitor can be programmed to set off audible and visual alarms, thereby providing notification of the possible impending danger. Most commonly the oxygen range to alert personnel is when oxygen levels get below 19.5%. OSHA specifies that a hazardous atmosphere may include one where the oxygen concentration is below 19.5% or above 23.5%.<ref>National Institutes of Health. Protocol for Use and Maintenance of Oxygen Monitoring Devices. February 2014, at 1:35 UTC. Available at: https://www.ors.od.nih.gov/sr/dohs/documents/protocoloxygenmonitoring.pdf {{Webarchive|url=https://web.archive.org/web/20201205185802/https://www.ors.od.nih.gov/sr/dohs/documents/protocoloxygenmonitoring.pdf |date=2020-12-05 }}. Accessed June 23, 2020</ref>
Oxygen deficiency monitors can either be fixed, mounted to the wall and hard-wired into the building's power supply or simply plugged into a power outlet, or a portable hand-held or wearable monitor.
==See also==
* [[Reactive nitrogen species]]
* [[Soil gas]]
==References==
{{Reflist}}
==Bibliography==
* {{Greenwood&Earnshaw2nd}}
==External links==
* [http://www.balashon.com/2008/07/neter-and-nitrogen.html Etymology of Nitrogen]
* [http://www.periodicvideos.com/videos/007.htm Nitrogen] at ''[[The Periodic Table of Videos]]'' (University of Nottingham)
* [http://www.rsc.org/periodic-table/podcast/7/nitrogen Nitrogen podcast] from the Royal Society of Chemistry's ''[[Chemistry World]]''
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