Elemental nitrogen is usually produced from air bypressure 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 asliquid nitrogen incryogenic applications. Many industrially important compounds, such asammonia, nitric acid, organic nitrates (propellants andexplosives), andcyanides, contain nitrogen. The extremely strongtriple bond in elemental nitrogen (N≡N), the second strongest bond in anydiatomic molecule aftercarbon monoxide (CO),[8] dominates nitrogen chemistry. This causes difficulty for both organisms and industry in converting N2 into usefulcompounds, 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 industrialfertilisers, and fertiliser nitrates are keypollutants in theeutrophication of water systems. Apart from its use in fertilisers and energy stores, nitrogen is a constituent of organic compounds as diverse asaramids used in high-strength fabric andcyanoacrylate used insuperglue.
Nitrogen compounds have a very long history,ammonium chloride having been known toHerodotus. They were well-known by the Middle Ages.Alchemists knew nitric acid asaqua fortis (strong water), as well as other nitrogen compounds such asammonium salts andnitrate salts. The mixture of nitric andhydrochloric acids was known asaqua regia (royal water), celebrated for its ability to dissolvegold, the king of metals.[9]
The discovery of nitrogen is attributed to the Scottish physicianDaniel Rutherford in 1772, who called itnoxious air.[10][11] Though he did not recognise it as an entirely different chemical substance, he clearly distinguished it from Joseph Black's"fixed air", or carbon dioxide.[12] The fact that there was a component of air that does not supportcombustion was clear to Rutherford, although he was not aware that it was an element. Nitrogen was also studied at about the same time byCarl Wilhelm Scheele,[13]Henry Cavendish,[14] andJoseph Priestley,[15] who referred to it asburnt air orphlogisticated air. French chemistAntoine Lavoisier referred to nitrogen gas as "mephitic air" orazote, from theGreek wordάζωτικός (azotikos), "no life", because it isasphyxiant.[16][17] 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 GermanStickstoff 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 ashydrazine and compounds of theazide ion. Finally, it led to the name "pnictogens" for the group headed by nitrogen, from the Greek πνίγειν "to choke".[9]
The English word nitrogen (1794) entered the language from the Frenchnitrogène, coined in 1790 by French chemistJean-Antoine Chaptal (1756–1832),[18] from the Frenchnitre (potassium nitrate, also calledsaltpetre) and the French suffix-gène, "producing", from theGreek -γενής (-genes, "begotten"). Chaptal's meaning was that nitrogen is the essential part ofnitric acid, which in turn was produced fromnitre. In earlier times, nitre had been confused with Egyptian "natron" (sodium carbonate) – called νίτρον (nitron) in Greek – which, despite the name, contained no nitrate.[19]
The earliest military, industrial, and agricultural applications of nitrogen compounds used saltpetre (sodium nitrate or potassium nitrate), most notably ingunpowder, and later asfertiliser. In 1910,Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", amonatomicallotrope of nitrogen.[20] The "whirling cloud of brilliant yellow light" produced by his apparatus reacted withmercury to produce explosivemercury nitride.[21]
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 theFrank–Caro process (1895–1899) andHaber–Bosch process (1908–1913) eased this shortage of nitrogen compounds, to the extent that half of globalfood production now relies on synthetic nitrogen fertilisers.[22] At the same time, use of theOstwald process (1902) to produce nitrates from industrial nitrogen fixation allowed the large-scale industrial production of nitrates asfeedstock in the manufacture ofexplosives in theWorld Wars of the 20th century.[23][24]
Properties
Atomic
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), 2px, 2py, 2pz.
A nitrogen atom has seven electrons. In the ground state, they are arranged in the electron configuration 1s2 2s2 2p1 x2p1 y2p1 z. It, therefore, has fivevalence electrons in the 2s and 2p orbitals, three of which (the p-electrons) are unpaired. It has one of the highestelectronegativities among the elements (3.04 on the Pauling scale), exceeded only bychlorine (3.16),oxygen (3.44), andfluorine (3.98). (The lightnoble gases,helium,neon, andargon, would presumably also be more electronegative, and in fact are on the Allen scale.)[25] Following periodic trends, its single-bondcovalent radius of 71 pm is smaller than those ofboron (84 pm) andcarbon (76 pm), while it is larger than those of oxygen (66 pm) and fluorine (57 pm). Thenitride anion, N3−, is much larger at 146 pm, similar to that of theoxide (O2−: 140 pm) andfluoride (F−: 133 pm) anions.[25] The first three ionisation energies of nitrogen are 1.402, 2.856, and 4.577 MJ·mol−1, and the sum of the fourth and fifth is16.920 MJ·mol−1. Due to these very high figures, nitrogen has no simple cationic chemistry.[26]The lack of radial nodes in the 2p subshell is directly responsible for many of the anomalous properties of the first row of thep-block, especially in nitrogen, oxygen, and fluorine. The 2p subshell is very small and has a very similar radius to the 2s shell, facilitatingorbital 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.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-richthree-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 congenersphosphorus,arsenic,antimony, andbismuth.[27]
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 asdiagonal relationships), 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.[28]
Nitrogen does not share the proclivity of carbon forcatenation. 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-,graphitic-, andfullerenic-like structures.[29]
It resembles oxygen with its high electronegativity and concomitant capability forhydrogen bonding and the ability to formcoordination complexes by donating itslone pairs of electrons. There are some parallels between the chemistry of ammonia NH3 and water H2O. For example, the capacity of both compounds to be protonated to give NH4+ and H3O+ or deprotonated to give NH2− and OH−, with all of these able to be isolated in solid compounds.[30]
Nitrogen shares with both its horizontal neighbours a preference for forming multiple bonds, typically with carbon, oxygen, or other nitrogen atoms, through pπ–pπ interactions.[28] Thus, for example, nitrogen occurs as diatomic molecules and therefore has very much lowermelting (−210 °C) andboiling points (−196 °C) than the rest of its group, as the N2 molecules are only held together by weakvan der Waals interactions and there are very few electrons available to create significant instantaneous dipoles. This is not possible for its vertical neighbours; thus, thenitrogen oxides,nitrites,nitrates,nitro-,nitroso-,azo-, anddiazo-compounds,azides,cyanates,thiocyanates, andimino-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.[28] Setting aside their differences, nitrogen and phosphorus form an extensive series of compounds with one another; these have chain, ring, and cage structures.[31]
Table of thermal and physical properties of nitrogen (N2) at atmospheric pressure:[32][33]
Table of nuclides (Segrè chart) from carbon to fluorine (including nitrogen). Orange indicatesproton emission (nuclides outside the proton drip line); pink forpositron emission (inverse beta decay); black forstable nuclides; blue forelectron emission (beta decay); and violet forneutron 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 stableisotopes:14N and15N. 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 u.[25] Both of these stable isotopes are produced in theCNO cycle instars, but14N is more common as its proton capture is the rate-limiting step.14N is one of the five stableodd–odd nuclides (a nuclide having an odd number of protons and neutrons); the other four are2H,6Li,10B, and180mTa.[34]
The relative abundance of14N and15N is practically constant in the atmosphere but can vary elsewhere, due to natural isotopic fractionation from biologicalredox reactions and the evaporation of naturalammonia ornitric acid.[35] Biologically mediated reactions (e.g.,assimilation,nitrification, anddenitrification) strongly control nitrogen dynamics in the soil. These reactions typically result in15N enrichment of thesubstrate and depletion of theproduct.[36]
The heavy isotope15N was first discovered by S. M. Naudé in 1929, and soon after heavy isotopes of the neighbouring elementsoxygen andcarbon were discovered.[37] It presents one of the lowest thermal neutron capture cross-sections of all isotopes.[38] It is frequently used innuclear magnetic resonance (NMR) spectroscopy to determine the structures of nitrogen-containing molecules, due to its fractionalnuclear spin of one-half, which offers advantages for NMR such as narrower line width.14N, though also theoretically usable, has an integer nuclear spin of one and thus has aquadrupole moment that leads to wider and less useful spectra.[25]15N NMR nevertheless has complications not encountered in the more common1H and13C NMR spectroscopy. The low natural abundance of15N (0.36%) significantly reduces sensitivity, a problem which is only exacerbated by its lowgyromagnetic ratio, (only 10.14% that of1H). As a result, the signal-to-noise ratio for1H is about 300 times as much as that for15N at the same magnetic field strength.[39] This may be somewhat alleviated by isotopic enrichment of15N by chemical exchange or fractional distillation.15N-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 labelledhydrogen, carbon, and oxygen isotopes that must be kept away from the atmosphere.[25] The15N:14N ratio is commonly used in stable isotope analysis in the fields ofgeochemistry,hydrology,paleoclimatology andpaleoceanography, where it is calledδ15N.[40]
Of the thirteen other isotopes produced synthetically, ranging from9N to23N,13N has ahalf-life of ten minutes and the remaining isotopes have half-lives less than eight seconds.[41][42] Given the half-life difference,13N is the most important nitrogen radioisotope, being relatively long-lived enough to use inpositron 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 acyclotron via proton bombardment of16O producing13N and analpha particle.[43]
Theradioisotope16N is the dominantradionuclide in the coolant ofpressurised water reactors orboiling water reactors during normal operation. It is produced from16O (in water) via an(n,p) reaction, in which the16O atom captures a neutron and expels a proton. It has a short half-life of about 7.1 s,[42] but its decay back to16O produces high-energygamma radiation (5 to 7 MeV).[42][44] Because of this, access to the primary coolant piping in a pressurised water reactor must be restricted duringreactor power operation.[44] 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.[44]
Molecular orbital diagram of dinitrogen molecule, N2. 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 atriradical 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 N2 molecule, they may release so much energy on collision with even such stable molecules ascarbon dioxide andwater 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 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.[28]
Given the great reactivity of atomic nitrogen, elemental nitrogen usually occurs as molecular N2, dinitrogen. This molecule is a colourless, odourless, and tastelessdiamagnetic gas at standard conditions: it melts at −210 °C and boils at −196 °C.[28] Dinitrogen is mostly unreactive at room temperature, but it will nevertheless react withlithium metal and sometransition metal complexes. This is due to its bonding, which is unique among the diatomic elements at standard conditions in that it has an N≡Ntriple bond. Triple bonds have short bond lengths (in this case, 109.76 pm) and high dissociation energies (in this case, 945.41 kJ/mol), and are thus very strong, explaining dinitrogen's low level of chemical reactivity.[28][45]
Other nitrogenoligomers 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.[46] Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced in adiamond anvil cell, nitrogen polymerises into the single-bondedcubic gauche crystal structure. This structure is similar to that ofdiamond, and both have extremely strongcovalent bonds, resulting in its nickname "nitrogen diamond".[47]
Solid nitrogen on the plains ofSputnik Planitia (on the bottom-right side of the image) onPluto next to water ice mountains (on the up-left side of the image)
Atatmospheric pressure, molecular nitrogencondenses (liquefies) at 77 K (−195.79 °C) andfreezes at 63 K (−210.01 °C)[48] into the betahexagonal close-packed crystalallotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes thecubic crystal allotropic form (called the alpha phase).[49]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 g/mL), is a commoncryogen.[50]Solid nitrogen has many crystalline modifications. It forms a significant dynamic surface coverage on Pluto[51] and outer moons of the Solar System such asTriton.[52] Even at the low temperatures of solid nitrogen it is fairly volatile and cansublime to form an atmosphere, or condense back into nitrogen frost. It is very weak and flows in the form of glaciers, and on Tritongeysers of nitrogen gas come from the polar ice cap region.[53]
The first example of adinitrogen complex to be discovered was [Ru(NH3)5(N2)]2+ (see figure at right), and soon many other such complexes were discovered. Thesecomplexes, in which a nitrogen molecule donates at least one lone pair of electrons to a central metal cation, illustrate how N2 might bind to the metal(s) innitrogenase and thecatalyst for theHaber process: these processes involving dinitrogen activation are vitally important in biology and in the production of fertilisers.[54][55]
Dinitrogen is able to coordinate to metals in five different ways. The more well-characterised ways are the end-on M←N≡N (η1) and M←N≡N→M (μ, bis-η1), 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 abridging ligand to two metal cations (μ, bis-η2) or to just one (η2). The fifth and unique method involves triple-coordination as a bridging ligand, donating all three electron pairs from the triple bond (μ3-N2). A few complexes feature multiple N2 ligands and some feature N2 bonded in multiple ways. Since N2 is isoelectronic withcarbon monoxide (CO) andacetylene (C2H2), the bonding in dinitrogen complexes is closely allied to that incarbonyl compounds, although N2 is a weakerσ-donor andπ-acceptor than CO. Theoretical studies show thatσ donation is a more important factor allowing the formation of the M–N bond thanπ back-donation, which mostly only weakens the N–N bond, and end-on (η1) donation is more readily accomplished than side-on (η2) donation.[28]
Today, dinitrogen complexes are known for almost all thetransition metals, accounting for several hundred compounds. They are normally prepared by three methods:[28]
Replacing labile ligands such asH2O,H−, orCO 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 withdimethylphenylphosphine (PMe2Ph) 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 coordinatedammonia (NH3) withnitrous acid (HNO2), 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(η5-C5H5)2}2-(N2)].[28]
Nitrides, azides, and nitrido complexes
Nitrogen bonds to almost all the elements in the periodic table except the first twonoble gases,helium andneon, and some of the very short-lived elements afterbismuth, creating an immense variety of binary compounds with varying properties and applications.[28] Many binary compounds are known: with the exception of the nitrogen hydrides, oxides, and fluorides, these are typically callednitrides. Many stoichiometric phases are usually present for most elements (e.g. MnN, Mn6N5, Mn3N2, Mn2N, Mn4N, and MnxN for 9.2 <x < 25.3). They may be classified as "salt-like" (mostly ionic), covalent, "diamond-like", and metallic (orinterstitial), 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 bythermal decomposition of metal amides:[56]
3 Ca + N2 → Ca3N2
3 Mg + 2 NH3 → Mg3N2 + 3 H2 (at 900 °C)
3 Zn(NH2)2 → Zn3N2 + 4 NH3
Many variants on these processes are possible. The most ionic of these nitrides are those of thealkali metals andalkaline earth metals, Li3N (Na, K, Rb, and Cs do not form stable nitrides for steric reasons) and M3N2 (M = Be, Mg, Ca, Sr, Ba). These can formally be thought of as salts of the N3− anion, although charge separation is not actually complete even for these highly electropositive elements. However, the alkali metalazides NaN3 and KN3, featuring the linearN− 3 anion, are well-known, as are Sr(N3)2 and Ba(N3)2. Azides of the B-subgroup metals (those ingroups 11 through16) are much less ionic, have more complicated structures, and detonate readily when shocked.[56]
Mesomeric structures of borazine, (–BH–NH–)3
Many covalent binary nitrides are known. Examples includecyanogen ((CN)2),triphosphorus pentanitride (P3N5),disulfur dinitride (S2N2), andtetrasulfur tetranitride (S4N4). The essentially covalentsilicon nitride (Si3N4) andgermanium nitride (Ge3N4) are also known: silicon nitride, in particular, would make a promisingceramic if not for the difficulty of working with and sintering it. In particular, thegroup 13 nitrides, most of which are promisingsemiconductors, are isoelectronic with graphite, diamond, andsilicon 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 oforganic chemistry finds an echo in boron–nitrogen chemistry, such as inborazine ("inorganicbenzene"). Nevertheless, the analogy is not exact due to the ease ofnucleophilic attack at boron due to its deficiency in electrons, which is not possible in a wholly carbon-containing ring.[56]
The largest category of nitrides are the interstitial nitrides of formulae MN, M2N, and M4N (although variable composition is perfectly possible), where the small nitrogen atoms are positioned in the gaps in a metallic cubic orhexagonal close-packed lattice. They are opaque, very hard, and chemically inert, melting only at very high temperatures (generally over 2500 °C). They have a metallic lustre and conduct electricity as do metals. They hydrolyse only very slowly to give ammonia or nitrogen.[56]
The nitride anion (N3−) is the strongestπ donor known among ligands (the second-strongest is O2−). Nitrido complexes are generally made by the thermal decomposition of azides or by deprotonating ammonia, and they usually involve a terminal {≡N}3− group. The linear azide anion (N− 3), being isoelectronic withnitrous oxide,carbon dioxide, andcyanate, forms many coordination complexes. Further catenation is rare, althoughN4− 4 (isoelectronic withcarbonate andnitrate) is known.[56]
Hydrides
Standard reduction potentials for nitrogen-containing species. Top diagram shows potentials at pH 0; bottom diagram shows potentials at pH 14.[57]
Industrially,ammonia (NH3) 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 ofhydrogen bonding has very significant effects on ammonia, conferring on it its high melting (−78 °C) and boiling (−33 °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 highdielectric constant, and is less dense than water. However, the hydrogen bonding in NH3 is weaker than that in H2O due to the lower electronegativity of nitrogen compared to oxygen and the presence of only one lone pair in NH3 rather than two in H2O. It is a weak base in aqueous solution (pKb 4.74); its conjugate acid isammonium,NH+ 4. It can also act as an extremely weak acid, losing a proton to produce the amide anion,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 givenitrogen 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.[58]
Many other binary nitrogen hydrides are known, but the most important arehydrazine (N2H4) andhydrogen azide (HN3). Although it is not a nitrogen hydride,hydroxylamine (NH2OH) 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 °C, boiling point 113.5 °C, density 1.00 g/cm3). 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.[59] It is also commonly used as a rocket fuel.[60]
Hydrazine is generally made by reaction of ammonia with alkalinesodium hypochlorite in the presence of gelatin or glue:[59]
NH3 + OCl− → NH2Cl + OH−
NH2Cl + NH3 →N 2H+ 5 + Cl− (slow)
N 2H+ 5 + OH− → N2H4 + H2O (fast)
(The attacks by hydroxide and ammonia may be reversed, thus passing through the intermediate NHCl− instead.) The reason for adding gelatin is that it removes metal ions such as Cu2+ that catalyses the destruction of hydrazine by reaction withmonochloramine (NH2Cl) to produceammonium chloride and nitrogen.[59]
Hydrogen azide (HN3) 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 thehydrohalic acids.[59]
All four simple nitrogen trihalides are known. A few mixed halides and hydrohalides are known, but are mostly unstable; examples include NClF2, NCl2F, NBrF2, NF2H,NFH2,NCl2H, andNClH2.[61]
Nitrogen trifluoride (NF3, first prepared in 1928) is a colourless and odourless gas that is thermodynamically stable, and most readily produced by theelectrolysis of moltenammonium fluoride dissolved in anhydroushydrogen fluoride. Likecarbon 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 withcopper, arsenic, antimony, and bismuth on contact at high temperatures to givetetrafluorohydrazine (N2F4). The cationsNF+ 4 andN 2F+ 3 are also known (the latter from reacting tetrafluorohydrazine with strong fluoride-acceptors such asarsenic pentafluoride), as is ONF3, 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 NF2•.Fluorine azide (FN3) is very explosive and thermally unstable.Dinitrogen difluoride (N2F2) exists as thermally interconvertiblecis andtrans isomers, and was first found as a product of the thermal decomposition of FN3.[61]
Nitrogen trichloride (NCl3) is a dense, volatile, and explosive liquid whose physical properties are similar to those ofcarbon tetrachloride, although one difference is that NCl3 is easily hydrolysed by water while CCl4 is not. It was first synthesised in 1811 byPierre 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 (NBr3), first prepared in 1975, is a deep red, temperature-sensitive, volatile solid that is explosive even at −100 °C.Nitrogen triiodide (NI3) 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 evenalpha particles.[61][62] 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".[63]Chlorine azide (ClN3) andbromine azide (BrN3) are extremely sensitive and explosive.[64][65]
Two series of nitrogen oxohalides are known: the nitrosyl halides (XNO) and the nitryl halides (XNO2). 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 (FNO2) andnitryl chloride (ClNO2) are likewise reactive gases and vigorous halogenating agents.[61]
Nitrogen dioxide at −196 °C, 0 °C, 23 °C, 35 °C, and 50 °C.NO 2 converts to colourless dinitrogen tetroxide (N 2O 4) at low temperatures, and reverts toNO 2 at higher temperatures.
Nitrous oxide (N2O), better known as laughing gas, is made by thermal decomposition of moltenammonium nitrate at 250 °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 forsprayed canned whipped cream, and was formerly commonly used as an anaesthetic. Despite appearances, it cannot be considered to be theanhydride ofhyponitrous acid (H2N2O2) 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, orozone at room temperature, although reactivity increases upon heating) and has the unsymmetrical structure N–N–O (N≡N+O−↔−N=N+=O): above 600 °C it dissociates by breaking the weaker N–O bond.[66]Nitric oxide (NO) is the simplest stable molecule with an odd number of electrons. In mammals, including humans, it is an important cellularsignalling molecule involved in many physiological and pathological processes.[68] 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 °C. Its bonding is similar to that in nitrogen, but one extra electron is added to aπ* 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 thecis 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.[66]
Blue dinitrogen trioxide (N2O3) is only available as a solid because it rapidly dissociates above its melting point to give nitric oxide, nitrogen dioxide (NO2), and dinitrogen tetroxide (N2O4). 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 tonitrosonium andnitrate 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 formnitric 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 andUSSR by the late 1950s. This is because it is ahypergolic propellant in combination with ahydrazine-basedrocket fuel and can be easily stored since it is liquid at room temperature.[66]
The thermally unstable and very reactive dinitrogen pentoxide (N2O5) is the anhydride ofnitric acid, and can be made from it by dehydration withphosphorus pentoxide. It is of interest for the preparation of explosives.[69] It is adeliquescent, colourless crystalline solid that is sensitive to light. In the solid state it is ionic with structure [NO2]+[NO3]−; as a gas and in solution it is molecular O2N–O–NO2. Hydration to nitric acid comes readily, as does analogous reaction withhydrogen peroxide givingperoxonitric acid (HOONO2). It is a violent oxidising agent. Gaseous dinitrogen pentoxide decomposes as follows:[66]
N2O5 ⇌ NO2 + NO3 → NO2 + O2 + NO
N2O5 + NO ⇌ 3 NO2
Oxoacids, oxoanions, and oxoacid salts
Many nitrogenoxoacids are known, though most of them are unstable as pure compounds and are known only as aqueous solutions or as salts.Hyponitrous acid (H2N2O2) is a weak diprotic acid with the structure HON=NOH (pKa1 6.9, pKa2 11.6). Acidic solutions are quite stable but above pH 4 base-catalysed decomposition occurs via [HONNO]− to nitrous oxide and the hydroxide anion.Hyponitrites (involving theN 2O2− 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 thenitrogen cycle. Hyponitrite can act as a bridging or chelating bidentate ligand.[70]
Nitrous acid (HNO2) 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 aqueousnitrite (NO− 2, bent) solutions, although already at room temperature disproportionation tonitrate and nitric oxide is significant. It is a weak acid with pKa 3.35 at 18 °C. They may betitrimetrically analysed by their oxidation to nitrate bypermanganate. They are readily reduced to nitrous oxide and nitric oxide bysulfur dioxide, to hyponitrous acid withtin(II), and to ammonia withhydrogen sulfide. Salts ofhydraziniumN 2H+ 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 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 todiazotise primary aromatic amines as follows:[70]
ArNH2 + HNO2 → [ArNN]Cl + 2 H2O
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.[70]
Fuming nitric acid contaminated with yellow nitrogen dioxide
Nitric acid (HNO3) 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 beingsulfuric acid andhydrochloric 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 theUnited 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 undergoesself-ionisation to a larger extent than any other covalent liquid as follows:[70]
2 HNO3 ⇌H 2NO+ 3 +NO− 3 ⇌ H2O + [NO2]+ + [NO3]−
Two hydrates, HNO3·H2O and HNO3·3H2O, are known that can be crystallised. It is a strong acid and concentrated solutions are strong oxidising agents, thoughgold,platinum,rhodium, andiridium are immune to attack. A 3:1 mixture of concentrated hydrochloric acid and nitric acid, calledaqua 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 formnitronium, which can act as an electrophile for aromatic nitration:[70]
HNO3 + 2 H2SO4 ⇌NO+ 2 + H3O+ + 2HSO− 4
The thermal stabilities ofnitrates (involving the trigonal planarNO− 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 andlead), or even the metal itself (silver) depending on their relative stabilities. Nitrate is also a common ligand with many modes of coordination.[70]
Finally, although orthonitric acid (H3NO4), which would be analogous toorthophosphoric acid, does not exist, the tetrahedralorthonitrate anionNO3− 4 is known in its sodium and potassium salts:[70]
These white crystalline salts are very sensitive to water vapour and carbon dioxide in the air:[70]
Na3NO4 + H2O + CO2 → NaNO3 + NaOH + NaHCO3
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.[70]
Nitrogen is the most common pure element in the earth, making up 78.1% of the volume of theatmosphere[9] (75.5% by mass), around 3.89 milliongigatonnes (3.89×1018 kg). Despite this, it isnot very abundant in Earth's crust, making up somewhere around 19 parts per million of this, on par withniobium,gallium, andlithium. (This represents 300,000 to a million gigatonnes of nitrogen, depending on the mass of the crust.[73]) The only important nitrogen minerals arenitre (potassium nitrate, saltpetre) andsoda 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.[74]
Nitrogen compounds constantly interchange between the atmosphere and living organisms. Nitrogen must first be processed, or "fixed", into a plant-usable form, usually ammonia. Some nitrogen fixation is done by lightning strikes producing the nitrogen oxides, but most is done bydiazotrophic bacteria through enzymes known asnitrogenases (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 anddenitrification, returning free dinitrogen to the atmosphere. Industrial nitrogen fixation by theHaber process is mostly used as fertiliser, although excess nitrogen–bearing waste, when leached, leads toeutrophication of freshwater and the creation of marinedead 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 atmosphericozone layer.[74]
Many saltwater fish manufacture large amounts oftrimethylamine oxide to protect them from the highosmotic effects of their environment; conversion of this compound todimethylamine is responsible for the early odour in unfresh saltwater fish.[75] In animals,free radicalnitric oxide (derived from anamino acid), serves as an important regulatory molecule for circulation.[76]
Nitric oxide's rapid reaction with water in animals results in the production of its metabolitenitrite. Animalmetabolism of nitrogen in proteins, in general, results in theexcretion ofurea, while animal metabolism ofnucleic acids results in the excretion ofurea anduric acid. The characteristic odour of animal flesh decay is caused by the creation of long-chain, nitrogen-containingamines, such asputrescine andcadaverine, which are breakdown products of the amino acidsornithine andlysine, respectively, in decaying proteins.[77]
Production
Nitrogen gas is anindustrial gas produced by thefractional distillation ofliquid air, or by mechanical means using gaseous air (pressurised reverseosmosis membrane orpressure swing adsorption). Nitrogen gas generators using membranes or pressure swing adsorption (PSA) are typically more cost and energy efficient than bulk-delivered nitrogen.[78] Commercial nitrogen is often a byproduct of air-processing for industrial concentration ofoxygen for steelmaking and other purposes. When supplied compressed in cylinders it is often called OFN (oxygen-free nitrogen).[79] Commercial-grade nitrogen already contains at most 20 ppm oxygen, and specially purified grades containing at most 2 ppm oxygen and 10 ppmargon are also available.[80]
Small amounts of the impurities NO and HNO3 are also formed in this reaction. The impurities can be removed by passing the gas through aqueous sulfuric acid containingpotassium dichromate.[81]
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
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:[80]
In some aircraft fuel systems to reduce fire hazard (seeinerting system).
To inflate race car and aircrafttires,[88] reducing the problems of inconsistent expansion and contraction caused by moisture andoxygen in natural air.[80]
Nitrogen is commonly used during sample preparation inchemical 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.[89]
Nitrogen can be used as a replacement, or in combination with,carbon dioxide to pressurise kegs of somebeers, particularlystouts and Britishales, due to the smallerbubbles it produces, which makes the dispensed beer smoother andheadier.[90] A pressure-sensitive nitrogen capsule known commonly as a "widget" allows nitrogen-charged beers to be packaged incans andbottles.[91][92] Nitrogen tanks are also replacing carbon dioxide as the main power source forpaintball guns. Nitrogen must be kept at a higher pressure than CO2, making N2 tanks heavier and more expensive.[93]
Equipment
Some construction equipment uses pressurised nitrogen gas to helphydraulic system to provide extra power to devices such ashydraulic hammer. Nitrogen gas, formed from the decomposition ofsodium azide, is used for the inflation ofairbags.[94]
As nitrogen is an asphyxiant gas in itself, some jurisdictions have considered asphyxiation by inhalation of pure nitrogen as a means ofcapital punishment (as a substitute forlethal injection).[95][96][97] In January 2024,Kenneth Eugene Smith became the first person executed by nitrogen asphyxiation.[98]
Likedry ice, the main use of liquid nitrogen is for cooling to low temperatures. It is used in thecryopreservation of biological materials such as blood and reproductive cells (sperm andeggs). It is used incryotherapy to remove cysts and warts on the skin by freezing them.[100] It is used in laboratorycold traps, and incryopumps to obtain lower pressures invacuum pumped systems. It is used to cool heat-sensitive electronics such asinfrared detectors andX-ray detectors. 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.[80]
Safety
Gas
Although nitrogen is non-toxic, when released into an enclosed space it can displace oxygen, and therefore presents anasphyxiation hazard. This may happen with few warning symptoms, since the humancarotid body is a relatively poor and slow low-oxygen (hypoxia) sensing system.[101] An example occurred shortly before the launch of thefirst Space Shuttle mission on March 19, 1981, when two technicians died from asphyxiation after they walked into a space located in theSpace Shuttle's mobile launcher platform that was pressurised with pure nitrogen as a precaution against fire.[102]
Nitrogen dissolves in theblood 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 calleddecompression sickness (formerly known as caisson sickness orthe bends), when nitrogen bubbles form in the bloodstream, nerves, joints, and other sensitive or vital areas.[105][106] Bubbles from other "inert" gases (gases other than carbon dioxide and oxygen) cause the same effects, so replacement of nitrogen inbreathing gases may prevent nitrogen narcosis, but does not prevent decompression sickness.[107]
Liquid
As acryogenic liquid, liquid nitrogen can be dangerous by causingcold burns on contact, although theLeidenfrost effect provides protection for very short exposure (about one second).[108] 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.[109]
Because the liquid-to-gasexpansion ratio of nitrogen is 1:694 at 20 °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, atTexas 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 m off their foundations.[110]
Liquid nitrogen readily evaporates to form gaseous nitrogen, and hence the precautions associated with gaseous nitrogen also apply to liquid nitrogen.[111][112][113] For example,oxygen sensors are sometimes used as a safety precaution when working with liquid nitrogen to alert workers of gas spills into a confined space.[114]
Vessels containing liquid nitrogen cancondense oxygen from air. The liquid in such a vessel becomes increasingly enriched in oxygen (boiling point −183 °C, higher than that of nitrogen) as the nitrogen evaporates, and can cause violent oxidation of organic material.[115]
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%.[116]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.
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