Nitrogen

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Nitrogen, 7N
A transparent liquid, with visible evaporation, being poured
Liquid nitrogen (N2 at below −196 °C)
Nitrogen
Allotropessee § Allotropes
Appearancecolorless gas, liquid or solid
Standard atomic weight Ar°(N)
Nitrogen in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
–
↑
N
↓
P
carbon ← nitrogen → oxygen
kJ/mol
Heat of vaporisation(N2) 5.57 kJ/mol
Molar heat capacity(N2) 29.124 J/(mol·K)
Vapour pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 37 41 46 53 62 77
Atomic properties
Discovery
Daniel Rutherford (1772)
Named byJean-Antoine Chaptal (1790)
Isotopes of nitrogen
Main isotopes Decay
abun­dance half-life (t1/2) mode pro­duct
13N trace 9.965 min
β+
 		13C
14N 99.6%
stable
15N 0.4% stable
 Category: Nitrogen
| references

Nitrogen is a

nonmetal and the lightest member of group 15 of the periodic table, often called the pnictogens. It is a common element in the universe, estimated at seventh in total abundance in the Milky Way and the Solar System. At standard temperature and pressure, two atoms of the element bond to form N2, a colorless and odorless diatomic gas. N2 forms about 78% of Earth's atmosphere
, making it the most abundant uncombined element 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

Ancient Greek: 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, azides and azo compounds
.

Elemental nitrogen is usually produced from air by

superglue
.

Nitrogen occurs in all organisms, primarily in

nitroprusside control blood pressure by metabolizing into nitric oxide. Many notable nitrogen-containing drugs, such as the natural caffeine and morphine or the synthetic amphetamines, act on receptors of animal neurotransmitters
.

History

Daniel Rutherford, discoverer of nitrogen

Nitrogen compounds have a very long history,

aqua fortis (strong water), as well as other nitrogen compounds such as ammonium salts and nitrate salts. The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold, the king of metals.[7]

The discovery of nitrogen is attributed to the Scottish physician Daniel Rutherford in 1772, who called it noxious air.[8][9] 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.[10] 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,[11] Henry Cavendish,[12] and Joseph Priestley,[13] who referred to it as burnt air or phlogisticated air. French chemist Antoine Lavoisier referred to nitrogen gas as "mephitic air" or azote, from the Greek word άζωτικός (azotikos), "no life", due to it being asphyxiant.[14][15] 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 "pnictogens" for the group headed by nitrogen, from the Greek πνίγειν "to choke".[7]

The English word nitrogen (1794) entered the language from the French nitrogène, coined in 1790 by French chemist

nitre. In earlier times, niter had been confused with Egyptian "natron" (sodium carbonate) – called νίτρον (nitron) in Greek – which, despite the name, contained no nitrate.[17]

The earliest military, industrial, and agricultural applications of nitrogen compounds used saltpeter (

allotrope of nitrogen.[18] The "whirling cloud of brilliant yellow light" produced by his apparatus reacted with mercury to produce explosive mercury nitride.[19]

For a long time, sources of nitrogen compounds were limited. Natural sources originated either from biology or deposits of nitrates produced by atmospheric reactions.

explosives in the World Wars of the 20th century.[21][22]

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 five valence electrons in the 2s and 2p orbitals, three of which (the p-electrons) are unpaired. It has one of the highest 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 gases, helium, neon, and argon, would presumably also be more electronegative, and in fact are on the Allen scale.)[23] Following periodic trends, its single-bond covalent radius of 71 pm is smaller than those of boron (84 pm) and carbon (76 pm), while it is larger than those of oxygen (66 pm) and fluorine (57 pm). The nitride anion, N3−, is much larger at 146 pm, similar to that of the oxide (O2−: 140 pm) and fluoride (F−: 133 pm) anions.[23] 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 is 16.920 MJ·mol−1. Due to these very high figures, nitrogen has no simple cationic chemistry.[24] The lack of radial nodes in the 2p subshell is directly responsible for many of the anomalous properties of the first row of the

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. 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.[25]

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 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.[26]

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-, graphitic-, and fullerenic-like structures.[27]

It resembles oxygen with its high electronegativity and concomitant capability for hydrogen bonding and the ability to form coordination complexes by donating its lone 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.[28]

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.

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.[26] Setting aside their differences, nitrogen and phosphorus form an extensive series of compounds with one another; these have chain, ring, and cage structures.[29]

Table of thermal and physical properties of nitrogen (N2) at atmospheric pressure:[30][31]

Temperature (K) Density (kg m−3) Specific heat (kJ kg−1 °C−1) Dynamic viscosity (kg m−1 s−1) Kinematic viscosity (m2 s−1) Thermal conductivity (W m−1 °C−1) Thermal diffusivity (m2 s−1) Prandtl number
100 3.4388 1.07   6.88×10−6 2.00×10−6 0.00958 2.60×10−6 0.768
150 2.2594 1.05   1.01×10−5 4.45×10−6 0.0139  5.86×10−6 0.759
200 1.7108 1.0429 1.29×10−5 7.57×10−6 0.01824 1.02×10−5 0.747
300 1.1421 1.0408 1.78×10−5 1.56×10−5 0.0262  2.20×10−5 0.713
400 0.8538 1.0459 2.20×10−5 2.57×10−5 0.03335 3.73×10−5 0.691
500 0.6824 1.0555 2.57×10−5 3.77×10−5 0.03984 5.53×10−5 0.684
600 0.5687 1.0756 2.91×10−5 5.12×10−5 0.0458  7.49×10−5 0.686
700 0.4934 1.0969 3.21×10−5 6.67×10−5 0.05123 9.47×10−5 0.691
800 0.4277 1.1225 3.48×10−5 8.15×10−5 0.05609 1.17×10−4 0.7  
900 0.3796 1.1464 3.75×10−5 9.11×10−5 0.0607  1.39×10−4 0.711
1000 0.3412 1.1677 4.00×10−5 1.19×10−4 0.06475 1.63×10−4 0.724
1100 0.3108 1.1857 4.23×10−5 1.36×10−4 0.0685  1.86×10−4 0.736
1200 0.2851 1.2037 4.45×10−5 1.56×10−4 0.07184 2.09×10−4 0.748
1300 0.2591 1.219  4.66×10−5 1.80×10−4 0.081   2.56×10−4 0.701

Isotopes

Main article: Isotopes of nitrogen
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 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 isotopes: 14N and 15N. 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.[23] Both of these stable isotopes are produced in the CNO cycle in stars, but 14N is more common as its proton capture is the rate-limiting step. 14N is one of the five stable odd–odd nuclides (a nuclide having an odd number of protons and neutrons); the other four are 2H, 6Li, 10B, and 180mTa.[32]

The relative abundance of 14N and 15N 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.[33] Biologically mediated reactions (e.g., assimilation, nitrification, and denitrification) strongly control nitrogen dynamics in the soil. These reactions typically result in 15N enrichment of the substrate and depletion of the product.[34]

The heavy isotope 15N was first discovered by S. M. Naudé in 1929, and soon after heavy isotopes of the neighbouring elements

quadrupole moment that leads to wider and less useful spectra.[23] 15N NMR nevertheless has complications not encountered in the more common 1H and 13C NMR spectroscopy. The low natural abundance of 15N (0.36%) significantly reduces sensitivity, a problem which is only exacerbated by its low gyromagnetic ratio, (only 10.14% that of 1H). As a result, the signal-to-noise ratio for 1H is about 300 times as much as that for 15N at the same magnetic field strength.[37] This may be somewhat alleviated by isotopic enrichment of 15N 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 labelled hydrogen, carbon, and oxygen isotopes that must be kept away from the atmosphere.[23] The 15N:14N ratio is commonly used in stable isotope analysis in the fields of geochemistry, hydrology, paleoclimatology and paleoceanography, where it is called δ15N.[38]

Of the thirteen other isotopes produced synthetically, ranging from 9N to 23N, 13N has a half-life of ten minutes and the remaining isotopes have half-lives less than eight seconds.[39][40] Given the half-life difference, 13N 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 16O producing 13N and an alpha particle.[41]

The

gamma radiation (5 to 7 MeV).[40][42] Because of this, access to the primary coolant piping in a pressurised water reactor must be restricted during reactor power operation.[42] 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.[42]

Allotropes

See also: Solid nitrogen
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 a 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 N2 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 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.[26]

Given the great reactivity of atomic nitrogen, elemental nitrogen usually occurs as molecular N2, dinitrogen. This molecule is a colourless, odourless, and tasteless

diamagnetic gas at standard conditions: it melts at −210 °C and boils at −196 °C.[26] 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 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.[26][43]

Other nitrogen

atm) and high temperatures (2000 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 bonds, resulting in its nickname "nitrogen diamond".[45]

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

cryogen.[48] Solid nitrogen has many crystalline modifications. It forms a significant dynamic surface coverage on Pluto[49] and outer moons of the Solar System such as Triton.[50] Even at the low temperatures of solid nitrogen it is fairly volatile and can 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 geysers of nitrogen gas come from the polar ice cap region.[51]

Chemistry and compounds

Dinitrogen complexes

Main article:
Dinitrogen complex
Structure of [Ru(NH3)5(N2)]2+ (pentaamine(dinitrogen)ruthenium(II)), the first dinitrogen complex to be discovered

The first example of a

complexes, 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) in nitrogenase and the catalyst for the Haber process: these processes involving dinitrogen activation are vitally important in biology and in the production of fertilisers.[52][53]

Dinitrogen is able to coordinate to metals in five different ways. The more well-characterised ways are the end-on M←N≡N (

carbonyl compounds, although N2 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 (η1) donation is more readily accomplished than side-on (η2) donation.[26]

Today, dinitrogen complexes are known for almost all the transition metals, accounting for several hundred compounds. They are normally prepared by three methods:[26]

  1. Replacing labile ligands such as H2O, H−, or CO directly by nitrogen: these are often reversible reactions that proceed at mild conditions.
  2. 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 (PMe2Ph) to make up for the smaller number of nitrogen ligands attached to the original chlorine ligands.
  3. 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 (NH3) with nitrous 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)].[26]

Nitrides, azides, and nitrido complexes

Nitrogen bonds to almost all the elements in the periodic table except the first two

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:[54]

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 the alkali metals and alkaline 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 metal azides NaN3 and KN3, featuring the linear N−
3 anion, are well-known, as are Sr(N3)2 and Ba(N3)2. Azides of the B-subgroup metals (those in groups 11 through 16) are much less ionic, have more complicated structures, and detonate readily when shocked.[54]

Mesomeric structures of borazine, (–BH–NH–)3

Many covalent binary nitrides are known. Examples include cyanogen ((CN)2), triphosphorus pentanitride (P3N5), disulfur dinitride (S2N2), and tetrasulfur tetranitride (S4N4). The essentially covalent silicon nitride (Si3N4) and germanium nitride (Ge3N4) 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 group 13 nitrides, most of which are promising semiconductors, 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 nucleophilic attack at boron due to its deficiency in electrons, which is not possible in a wholly carbon-containing ring.[54]

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 or

hexagonal 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.[54]

The nitride anion (N3−) is the strongest π 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 with nitrous oxide, carbon dioxide, and cyanate, forms many coordination complexes. Further catenation is rare, although N4−
4 (isoelectronic with carbonate and nitrate) is known.[54]

Hydrides

Standard reduction potentials for nitrogen-containing species. Top diagram shows potentials at pH 0; bottom diagram shows potentials at pH 14.[55]

Industrially,

dielectric 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 is ammonium, 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 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.[56]

Many other binary nitrogen hydrides are known, but the most important are

hydrogen 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.[57] It is also commonly used as a rocket fuel.[58]

Hydrazine is generally made by reaction of ammonia with alkaline sodium hypochlorite in the presence of gelatin or glue:[57]

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 with monochloramine (NH2Cl) to produce ammonium chloride and nitrogen.[57]

hydrohalic acids.[57]

Halides and oxohalides

Nitrogen trichloride

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, and NClH2.[59]

Nitrogen trifluoride (NF3, 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 (N2F4). The cations NF+
4 and N
2F+
3 are also known (the latter from reacting tetrafluorohydrazine with strong fluoride-acceptors such as arsenic 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 interconvertible cis and trans isomers, and was first found as a product of the thermal decomposition of FN3.[59]

Nitrogen trichloride (NCl3) is a dense, volatile, and explosive liquid whose physical properties are similar to those of carbon tetrachloride, although one difference is that NCl3 is easily hydrolysed by water while CCl4 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 (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 even alpha particles.[59][60] 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".[61] Chlorine azide (ClN3) and bromine azide (BrN3) are extremely sensitive and explosive.[62][63]

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) and nitryl chloride (ClNO2) are likewise reactive gases and vigorous halogenating agents.[59]

Oxides

Main article: Nitrogen oxide
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 to NO
2 at higher temperatures.

Nitrogen forms nine molecular oxides, some of which were the first gases to be identified: N2O (

oxatetrazole (N4O), an aromatic ring.[64]

Nitrous oxide (N2O), better known as laughing gas, is made by thermal decomposition of molten

anhydride of hyponitrous 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, or ozone 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.[64]
Nitric oxide (NO) is the simplest stable molecule with an odd number of electrons. In mammals, including humans, it is an important cellular
signaling molecule involved in many physiological and pathological processes.[66] 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 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.[64]

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 to

USSR by the late 1950s. This is because it is a hypergolic propellant in combination with a hydrazine-based rocket fuel and can be easily stored since it is liquid at room temperature.[64]

The thermally unstable and very reactive dinitrogen pentoxide (N2O5) is the anhydride of

peroxonitric acid (HOONO2). It is a violent oxidising agent. Gaseous dinitrogen pentoxide decomposes as follows:[64]

N2O5 ⇌ NO2 + NO3 → NO2 + O2 + NO
N2O5 + NO ⇌ 3 NO2

Oxoacids, oxoanions, and oxoacid salts

Many nitrogen

oxoacids 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 the N
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 the nitrogen cycle. Hyponitrite can act as a bridging or chelating bidentate ligand.[68]

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 aqueous nitrite (NO−
2, bent) solutions, although already at room temperature disproportionation to nitrate and nitric oxide is significant. It is a weak acid with pKa 3.35 at 18 °C. They may be 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 N
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 to diazotise primary aromatic amines as follows:[68]

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.[68]

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 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 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 self-ionisation to a larger extent than any other covalent liquid as follows:[68]

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, though

nitronium, which can act as an electrophile for aromatic nitration:[68]

HNO3 + 2 H2SO4 ⇌ NO+
2 + H3O+ + 2 HSO−
4

The thermal stabilities of nitrates (involving the trigonal planar 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.[68]

Finally, although orthonitric acid (H3NO4), which would be analogous to

orthophosphoric acid, does not exist, the tetrahedral orthonitrate anion NO3−
4 is known in its sodium and potassium salts:[68]

These white crystalline salts are very sensitive to water vapour and carbon dioxide in the air:[68]

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.[68]

Organic nitrogen compounds

Nitrogen is one of the most important elements in

nucleic acids, amino acids and thus proteins, and the energy-carrying molecule adenosine triphosphate and is thus vital to all life on Earth.[69]

Occurrence

Schematic representation of the 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

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.[72]

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 by diazotrophic bacteria through enzymes known as nitrogenases (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 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.[72]

Many saltwater fish manufacture large amounts of

free radical nitric oxide (derived from an amino acid), serves as an important regulatory molecule for circulation.[74]

Nitric oxide's rapid reaction with water in animals results in the production of its metabolite

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 amines, such as putrescine and cadaverine, which are breakdown products of the amino acids ornithine and lysine, respectively, in decaying proteins.[75]

Production

Nitrogen gas is an industrial gas produced by the fractional distillation of liquid air, or by mechanical means using gaseous air (pressurised reverse 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.[76] 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).[77] Commercial-grade nitrogen already contains at most 20 ppm oxygen, and specially purified grades containing at most 2 ppm oxygen and 10 ppm argon are also available.[78]

In a chemical laboratory, it is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite.[79]

NH4Cl + NaNO2 → N2 + NaCl + 2 H2O

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 containing potassium dichromate.[79]

It can also be obtained by the thermal decomposition of ammonium dichromate. [80]

3(NH4)2Cr2O7 → 2N2 + 9H2O + 3Cr2O3 + 2NH3 + 32O2

Very pure nitrogen can be prepared by the thermal decomposition of barium azide or sodium azide.[81]

2 NaN3 → 2 Na + 3 N2

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

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:[78]

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.[87]

Nitrogen can be used as a replacement, or in combination with,

paintball guns. Nitrogen must be kept at a higher pressure than CO2, making N2 tanks heavier and more expensive.[91]

Equipment

Some construction equipment uses pressurized nitrogen gas to help hydraulic system to provide extra power to devices such as hydraulic hammer. Nitrogen gas, formed from the decomposition of sodium azide, is used for the inflation of airbags.[92]

Execution

Further information: 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).[93][94][95] In January 2024, Kenneth Eugene Smith became the first person executed by nitrogen asphyxiation.[96]

Liquid

Main article: Liquid nitrogen
Air balloon submerged in liquid nitrogen

Liquid nitrogen is a

dewar flasks, it can be transported and stored with a low rate of evaporative loss.[97]

A container vehicle carrying liquid nitrogen

Like

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.[78]

Safety

Gas

Although nitrogen is non-toxic, when released into an enclosed space it can displace oxygen, and therefore presents an

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.[99] An example occurred shortly before the launch of the first Space Shuttle mission on March 19, 1981, when two technicians died from asphyxiation after they walked into a space located in the Space Shuttle's mobile launcher platform that was pressurised with pure nitrogen as a precaution against fire.[100]

When inhaled at high partial pressures (more than about 4 bar, encountered at depths below about 30 m in scuba diving), nitrogen is an anesthetic agent, causing nitrogen narcosis, a temporary state of mental impairment similar to nitrous oxide intoxication.[101][102]

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.[103][104] Bubbles from other "inert" gases (gases other than carbon dioxide and oxygen) cause the same effects, so replacement of nitrogen in breathing gases may prevent nitrogen narcosis, but does not prevent decompression sickness.[105]

Liquid

As a

cold burns on contact, although the Leidenfrost effect provides protection for very short exposure (about one second).[106] 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.[107]

Because the liquid-to-gas expansion 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, 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 m off their foundations.[108]

Liquid nitrogen readily evaporates to form gaseous nitrogen, and hence the precautions associated with gaseous nitrogen also apply to liquid nitrogen.[109][110][111] 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.[112]

Vessels containing liquid nitrogen can condense 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.[113]

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%.[114] 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

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Bibliography

External links
Retrieved from "https://en.wikipedia.org/w/index.php?title=Nitrogen&oldid=1205496181"