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Nitrogen

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Nitrogen,  7N
Liquidnitrogen.jpg
Nitrogen Spectra.jpg
Spectral lines of nitrogen
General properties
Name, symbol nitrogen, N
Pronunciation /ˈntrəən/
NY-trə-jən
Appearance colorless gas, liquid or solid
Nitrogen in the periodic table
Hydrogen (diatomic nonmetal)
Helium (noble gas)
Lithium (alkali metal)
Beryllium (alkaline earth metal)
Boron (metalloid)
Carbon (polyatomic nonmetal)
Nitrogen (diatomic nonmetal)
Oxygen (diatomic nonmetal)
Fluorine (diatomic nonmetal)
Neon (noble gas)
Sodium (alkali metal)
Magnesium (alkaline earth metal)
Aluminium (post-transition metal)
Silicon (metalloid)
Phosphorus (polyatomic nonmetal)
Sulfur (polyatomic nonmetal)
Chlorine (diatomic nonmetal)
Argon (noble gas)
Potassium (alkali metal)
Calcium (alkaline earth metal)
Scandium (transition metal)
Titanium (transition metal)
Vanadium (transition metal)
Chromium (transition metal)
Manganese (transition metal)
Iron (transition metal)
Cobalt (transition metal)
Nickel (transition metal)
Copper (transition metal)
Zinc (transition metal)
Gallium (post-transition metal)
Germanium (metalloid)
Arsenic (metalloid)
Selenium (polyatomic nonmetal)
Bromine (diatomic nonmetal)
Krypton (noble gas)
Rubidium (alkali metal)
Strontium (alkaline earth metal)
Yttrium (transition metal)
Zirconium (transition metal)
Niobium (transition metal)
Molybdenum (transition metal)
Technetium (transition metal)
Ruthenium (transition metal)
Rhodium (transition metal)
Palladium (transition metal)
Silver (transition metal)
Cadmium (transition metal)
Indium (post-transition metal)
Tin (post-transition metal)
Antimony (metalloid)
Tellurium (metalloid)
Iodine (diatomic nonmetal)
Xenon (noble gas)
Caesium (alkali metal)
Barium (alkaline earth metal)
Lanthanum (lanthanide)
Cerium (lanthanide)
Praseodymium (lanthanide)
Neodymium (lanthanide)
Promethium (lanthanide)
Samarium (lanthanide)
Europium (lanthanide)
Gadolinium (lanthanide)
Terbium (lanthanide)
Dysprosium (lanthanide)
Holmium (lanthanide)
Erbium (lanthanide)
Thulium (lanthanide)
Ytterbium (lanthanide)
Lutetium (lanthanide)
Hafnium (transition metal)
Tantalum (transition metal)
Tungsten (transition metal)
Rhenium (transition metal)
Osmium (transition metal)
Iridium (transition metal)
Platinum (transition metal)
Gold (transition metal)
Mercury (transition metal)
Thallium (post-transition metal)
Lead (post-transition metal)
Bismuth (post-transition metal)
Polonium (post-transition metal)
Astatine (metalloid)
Radon (noble gas)
Francium (alkali metal)
Radium (alkaline earth metal)
Actinium (actinide)
Thorium (actinide)
Protactinium (actinide)
Uranium (actinide)
Neptunium (actinide)
Plutonium (actinide)
Americium (actinide)
Curium (actinide)
Berkelium (actinide)
Californium (actinide)
Einsteinium (actinide)
Fermium (actinide)
Mendelevium (actinide)
Nobelium (actinide)
Lawrencium (actinide)
Rutherfordium (transition metal)
Dubnium (transition metal)
Seaborgium (transition metal)
Bohrium (transition metal)
Hassium (transition metal)
Meitnerium (unknown chemical properties)
Darmstadtium (unknown chemical properties)
Roentgenium (unknown chemical properties)
Copernicium (transition metal)
Nihonium (unknown chemical properties)
Flerovium (post-transition metal)
Moscovium (unknown chemical properties)
Livermorium (unknown chemical properties)
Tennessine (unknown chemical properties)
Oganesson (unknown chemical properties)


N

P
carbonnitrogenoxygen
Atomic number (Z) 7
Group, block group 15 (pnictogens), p-block
Period period 2
Element category   diatomic nonmetal
Standard atomic weight (Ar) 14.007[1] (14.00643–14.00728)[2]
Electron configuration [He] 2s2 2p3
per shell
2, 5
Physical properties
Phase gas
Melting point 63.15 K ​(−210.00 °C, ​−346.00 °F)
Boiling point 77.355 K ​(−195.795 °C, ​−320.431 °F)
Density at stp (0 °C and 101.325 kPa) 1.251 g/L
when liquid, at b.p. 0.808 g/cm3
Triple point 63.151 K, ​12.52 kPa
Critical point 126.192 K, 3.3958 MPa
Heat of fusion (N2) 0.72 kJ/mol
Heat of vaporisation (N2) 5.56 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
Oxidation states 5, 4, 3, 2, 1, −1, −2, −3 ​(a strongly acidic oxide)
Electronegativity Pauling scale: 3.04
Ionisation energies 1st: 1402.3 kJ/mol
2nd: 2856 kJ/mol
3rd: 4578.1 kJ/mol
(more)
Covalent radius 71±1 pm
Van der Waals radius 155 pm
Miscellanea
Crystal structure hexagonal
Hexagonal crystal structure for nitrogen
Speed of sound 353 m/s (gas, at 27 °C)
Thermal conductivity 25.83×10−3 W/(m·K)
Magnetic ordering diamagnetic
CAS Number 7727-37-9
History
Discovery Daniel Rutherford (1772)
Named by Jean-Antoine Chaptal (1790)
Most stable isotopes of nitrogen
iso NA half-life DM DE (MeV) DP
13N syn 9.965 min ε 2.220 13C
14N 99.6% is stable with 7 neutrons
15N 0.4% is stable with 8 neutrons
| references | in Wikidata

Nitrogen is a chemical element with symbol N and atomic number 7. It is the lightest member of group 15 of the periodic table, the heavier members of which are the pnictogens phosphorus, arsenic, antimony, and bismuth. Nevertheless, nitrogen is sometimes not considered a pnictogen, due to its distinct chemistry arising from its small size. Nitrogen is a common element in the universe, estimated at about seventh in total abundance in the Milky Way and the Solar System. At standard temperature and pressure, two atoms of the element bind to form dinitrogen, a colourless and odorless diatomic gas with the formula N2. Dinitrogen forms about 78% of Earth's atmosphere, making it the most abundant uncombined element. Additionally, nitrogen compounds, such as proteins, are essential to all known life forms.

The element nitrogen was first discovered and isolated by Scottish physician Daniel Rutherford in 1772. Although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is generally accorded the credit because his work was published first. The name nitrogen was suggested by Jean-Antoine-Claude Chaptal in 1790, when it was found that nitrogen was present in nitric acid and nitrates; this name derives from the Greek roots νἰτρον "nitre" and -γεννᾶν "to form". Antoine Lavoisier suggested instead the name azote, from the Greek άζωτικός "no life", as it is an asphyxiant gas; his name is instead used in many languages, such as French, Russian, and Turkish, and appears in the English names of some nitrogen compounds such as hydrazine, azides and azo compounds.

Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong triple bond in elemental nitrogen (N≡N), the strongest bond in any diatomic molecule, dominates nitrogen chemistry, causing difficulty for both organisms and industry in converting N2 into useful compounds, but at the same time causing release of large amounts of often useful energy when the compounds burn, explode, or decay back into nitrogen gas. Synthetically produced ammonia and nitrates are key industrial fertilisers, and fertiliser nitrates are key pollutants in the eutrophication of water systems.

Apart from its use in fertilisers and energy-stores, nitrogen is a constituent of organic compounds as diverse as Kevlar used in high-strength fabric and cyanoacrylate used in superglue. Nitrogen is a constituent of every major pharmacological drug class, including antibiotics. Many drugs are mimics or prodrugs of natural nitrogen-containing signal molecules: for example, the organic nitrates nitroglycerin and nitroprusside control blood pressure by metabolizing into nitric oxide. Plant alkaloids (often defense chemicals) contain nitrogen by definition, and many notable nitrogen-containing drugs, such as caffeine and morphine, are either alkaloids or synthetic mimics that act (as many plant alkaloids do) on receptors of animal neurotransmitters (for example, synthetic amphetamines).

Nitrogen occurs in all organisms, primarily in amino acids (and thus proteinhttps://en.wikipedia.org/w/index.php?title=Nitrogen&action=edits), in the nucleic acids (DNA and RNA) and in the energy transfer molecule adenosine triphosphate. The human body contains about 3% by mass of nitrogen, the fourth most abundant element in the body after oxygen, carbon, and hydrogen. The nitrogen cycle describes movement of the element from the air, into the biosphere and organic compounds, then back into the atmosphere.

History

Daniel Rutherford, discoverer of nitrogen

The discovery of nitrogen compounds has a very long history, ammonium chloride having been known to Herodotus. They were well known by the Middle Ages. Alchemists knew nitric acid as aqua fortis (strong water), as well as other nitrogen compounds such as ammonium salts and nitrate salts. The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold, the king of metals.[3]

The existence of nitrogen is formally considered to have been discovered by Scottish physician Daniel Rutherford in 1772, who called it noxious air.[4][5][6][7] 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.[8] 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,[9] Henry Cavendish,[10] and Joseph Priestley,[11] who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as "mephitic air" or azote, from the Greek word άζωτικός, "no life".[12] In an atmosphere of pure nitrogen, animals died and flames were extinguished. This "mephitic air" consisted mostly of N2, but might have included more than 1% argon. Though Lavoisier's name was not accepted in English, since it was pointed out that almost all gases (indeed, with the sole exception of oxygen) are mephitic, it is used in many languages (French, Italian, Portuguese, Polish, Russian, Albanian, Turkish, etc.; the German Stickstoff similarly refers to the same characteristic, viz. sticken "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".[3]

The English word nitrogen (1794) entered the language from the French nitrogène, coined in 1790 by French chemist Jean-Antoine Chaptal (1756–1832),[13] from the French nitre (potassium nitrate, also called saltpeter) and the French -gène, "producing", from the Greek -γενής ("producer, begetter"). Chaptal's meaning was that nitrogen gas is the essential part of nitric acid, which in turn was produced from niter. In earlier times, niter had been confused with Egyptian "natron" (sodium carbonate) — called νίτρον (nitron) in Greek — which, despite the name, contained no nitrate.[14]

The earliest military, industrial, and agricultural applications of nitrogen compounds used saltpeter (sodium nitrate or potassium nitrate), most notably in gunpowder, and later as fertiliser. In 1910, Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", a monatomic allotrope of nitrogen. The "whirling cloud of brilliant yellow light" produced by his apparatus reacted with mercury to produce explosive mercury nitride.[15]

For a long time, sources of nitrogen compounds were limited. Natural sources originated either from biology or deposits of nitrates produced by atmospheric reactions. Nitrogen fixation by industrial processes like the Frank–Caro process (1895–1899) and Haber–Bosch process (1908–1913) eased this shortage of nitrogen compounds, to the extent that half of global food production (see Applications) now relies on synthetic nitrogen fertilisers.[16] At the same time, use of the Ostwald process (1902) to produce nitrates from industrial nitrogen fixation allowed the large-scale industrial production of nitrates as feedstock in the manufacture of explosives in the World Wars of the 20th century.[17][18]

Properties

Atomic

A nitrogen atom has seven electrons. In the ground state, they are arranged in the electron configuration 1s2
2s2
2p1
x
2p1
y
2p1
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 oxygen (3.44), fluorine (3.98), and chlorine (3.16).[19] 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 (75 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.[19] Ionic radii of 16 pm and 13 pm respectively have been published for the N3+ and N5+ cations, but these must be taken as purely notional figures as simple cationic chemistry is unknown for nitrogen due to its high ionisation energies: 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.[20]

The lack of radial nodes in the 2p subshell is directly responsible for many of the anomalous properties of the first row of the p-block, especially in nitrogen, oxygen, and fluorine. The 2p subshell is very small and has a very similar radius to the 2s shell, facilitating orbital hybridisation. It also results in very large electrostatic forces of attraction between the nucleus and the valence electrons in the 2s and 2p shells, resulting in very high electronegativities. Hypervalency is almost unknown in these 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.[21]

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). This far to the right of the periodic table, the diagonal relationship with sulfur has with only a few exceptions faded completely into insignificance. Nitrogen resembles oxygen far more than it does carbon with its high electronegativity and concomitant capability for hydrogen bonding and the ability to form coordination complexes by donating its lone pairs of electrons. It does not share carbon's proclivity for catenation, with the longest chain of nitrogen yet discovered being composed of only eight nitrogen atoms (PhN=N–N(Ph)–N=N–N(Ph)–N=NPh). One property nitrogen does share with both its horizontal neighbours is its preferentially forming multiple bonds, typically with carbon, nitrogen, or oxygen atoms, through pπ–pπ interactions. This is not possible for its vertical neighbours; thus, the nitrogen oxides, nitrites, nitrates, nitro-, nitroso-, azo-, and diazo-compounds, azides, cyanates, thiocyanates, and imino-derivatives find no echo with phosphorus, arsenic, antimony, or bismuth. By the same token, however, the complexity of the phosphorus oxoacids finds no echo with nitrogen.[22]

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.[19] Both of these stable isotopes are produced in the CNO cycle in stars, but 14N is more common as its neutron 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 1H, 6Li, 10B, and 180mTa.[23]

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.[24] 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.[25]

The heavy isotope 15N was first discovered by S. M. Naudé in 1929, soon after heavy isotopes of the neighbouring elements oxygen and carbon were discovered.[26] It presents one of the lowest thermal neutron capture cross sections of all isotopes.[27] It is frequently used in nuclear magnetic resonance (NMR) spectroscopy, due to its fractional nuclear spin of one-half, which offers advantages for NMR such as narrower line width. In principle, 14N could also be used, but it has an integer nuclear spin of one and thus has a quadrupole moment that leads to wider spectra. Nevertheless, the first chemical shift observed in 1950 was from 14N in an aqueous solution of ammonium nitrate.[19]

15N NMR nevertheless has complications not encountered in 1H and 13C NMR spectroscopy. The 0.36% natural abundance of 15N results in a major sensitivity penalty. Sensitivity is made worse by its low gyromagnetic ratio, which is 10.14% that of 1H. The signal to noise ratio for 1H is about 300 times greater than 15N at the same magnetic field strength.[28] This may be somewhat alleviated by isotopic enrichment of 15N, either by chemical exchange of nitrogen atoms in equilibrium reactions (e.g. NO and NO
3
, NH3 and NH+
4
, or NO and NO2), or by fractional distillation of nitric oxide (NO), which results in enrichment of the doubly-labelled 15N18O. 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.[19]

Of the ten other isotopes produced synthetically, ranging from 12N to 23N, 13N has a half-life of ten minutes and the remaining isotopes have half-lives on the order of seconds or less. No other nitrogen isotopes are possible as they would fall outside the nuclear drip lines, leaking out a proton or neutron.[29] The radioisotope 16N is the dominant radionuclide in the coolant of pressurised water reactors or boiling water reactors during normal operation. It is produced from 16O (in water) via an (n,p) reaction in which the 16O atom captures a neutron and expels a proton. It has a short half-life of about 7.1 s,[29] but during its decay back to 16O produces high-energy gamma radiation (5 to 7 MeV).[29][30] Because of this, access to the primary coolant piping in a pressurised water reactor must be restricted during reactor power operation. 16N 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.[30]

Chemistry and compounds

Standard reduction potentials for nitrogen-containing species. Top diagram shows potentials at pH 0; bottom diagram shows potentials at pH 14.[31]

Allotropes

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. It 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.[22]

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.[22] The bonding in N2 is characterised by overlap between the 2s orbitals to give occupied σ2s (bonding) and σ*2s (antibonding) orbitals (which therefore cancel out and have no net contribution to the bond order), a head-on overlap between one of the p-orbitals on each nitrogen atom to give the σ2p bonding and σ*2p antibonding orbitals, and side-on overlap between the other two p-orbitals of each nitrogen atom to form the two π2p bonding and two π*2p antibonding orbitals.[21] All three bonding orbitals from 2p overlaps are occupied and none of the three antibonding ones are, for an overall bond order of 3, corroborated by short bond length (109.76 pm) and high dissociation energy (945.41 kJ/mol). Nitrogen is mostly unreactive at room temperature, but will react with lithium metal and some transition metal complexes.[22]

There are some theoretical indications that other nitrogen oligomers and polymers may be possible. If they could be synthesised, they may have potential applications as materials with a very high energy density, that could be used as powerful propellants or explosives.[32] This is because they should all decompose to dinitrogen, whose N≡N triple bond (bond energy 946 kJ⋅mol−1) is much stronger than those of the N=N double bond (418 kJ⋅mol−1) or the N–N single bond (160 kJ⋅mol−1): indeed, triple bond has more than thrice the energy of the single bond. (The opposite is true for the heavier pnictogens, which prefer polyatomic allotropes.)[33] Tetranitrogen (N4) has been discovered, but is poorly characterised due to its having a mean lifetime on the order of microseconds: its structure is probably triplet azidonitrene, N3–N.[34] The most hopeful other nitrogen allotropes for synthesis are hexanitrogen (N6) as diazide, and tetraazatetrahedrane (N4), although the latter is expected to be more challenging to synthesise than even the long-sought tetrahedrane, (CH)4. Other neutral polynitrogens, some of the more symmetric of which are hexazine (N6; an analogue of benzene) and octaazacubane (N8), are not expected to have a large barrier towards decomposition. This stands in contrast to the well-characterised cationic and anionic polynitrogens azide (N
3
), pentazenium (N+
5
), and pentazolide (cyclic aromatic N
5
).[32] Under extremely high pressures (1.1 million 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".[35]

At atmospheric pressure, molecular nitrogen condenses (liquefies) at 77 K (−195.79 °C) and freezes at 63 K (−210.01 °C)[36] into the beta hexagonal close-packed crystal allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the cubic crystal allotropic form (called the alpha phase).[37] 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 common cryogen.[38] Solid nitrogen has many crystalline modifications. It forms a significant dynamic surface coverage on Pluto[39] and outer moons of the Solar System such as Triton.[40] 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.[41]

Dinitrogen complexes

Main article: Dinitrogen complex

Dinitrogen is able to coordinate to metals in five different ways. The more well-characterised ways are the end-on ways 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 a bridging 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. In the M←N≡N structure, the M–N–N structure is linear. Since N2 is isoelectronic with carbon monoxide (CO) and acetylene (C2H2), the bonding in dinitrogen complexes is closely allied to that in 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.[22]

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

The first example of a dinitrogen complex to be discovered was [Ru(NH3)5(N2)]2+ (see figure at right). Notably, N2 ligand was obtained by the decomposition of hydrazine (N2H4), and not coordination of free dinitrogen. Many such compounds are now known, including IrCl(N2)(PPh3)2, W(N2)2(Ph2PCH2CH2PPh2)2, and [(η5-C5Me4H)2Zr]2(μ2,η2,η2-N2). These complexes illustrate how N2 might bind to the metal(s) in nitrogenase and the catalyst for the Haber process.[42] A catalytic process to reduce N2 to ammonia with the use of a molybdenum complex in the presence of a proton source was published in 2005.[43]

Today, dinitrogen complexes are known for almost all the transition metals, accounting for several hundred compounds. They are normally prepared by three methods:[22]

  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 coligand in excess under nitrogen gas. A common choice include replacing chloride ligands by dimethylphenylphosphine (PMe2Ph) to make up for the smaller number of nitrogen ligands attached than 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)].[22]

Nitrides, azides, and nitrido complexes

Nitrogen bonds to almost all the elements in the periodic table except the first three noble gases, helium, neon, and argon.[22] Many binary compounds are known: with the exception of the nitrogen hydrides, oxides, and fluorides, these are typically called nitrides. 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 (or 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:[44]

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. 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.[44]

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). In particular, the group 13 nitrides 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"), although the analogy is not exact due to the ease of nucleophilic attack at boron. 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.[44]

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.[44]

The nitride anion (N3−) is the strongest π donor known amongst ligands (the second-strongest is O2−). Nitrido complexes are generally made by thermal decomposition of azides or by deprotonating ammonia, and they usually involve a terminal {≡N}3− group. Linear symmetrical bridging, trigonal planar μ3 bridging, and tetrahedral coordination are rarer modes. The linear azide anion (N
3
), being isoelectronic with nitrous oxide, carbon dioxide, and cyanate, forms many coordination complexes, involving end-on η1 coordination, bridging μ,η1, and bridging μ,η1:η1 configuration. Further catenation is rare, although planar bridging N4−
4
(isoelectronic with carbonate and nitrate) is known.[44]

Hydrides

Structure of the ammonia molecule

Industrially, ammonia (NH3) is the most important compound of nitrogen and is prepared in larger amounts than any other compound. It is a colourless alkaline gas with a characteristic pungent smell. The presence of hydrogen bonding has very significant effects on ammonia, conferring on it its high melting (−78 °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 high 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.[45]

Hydrazine hydrate

Many other binary nitrogen hydrides are known, but the most important are hydrazine (N2H4) and 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 similarly 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·cm−3). It is generally made by reaction of ammonia with alkaline sodium hypochlorite in the presence of gelatin or glue:[46]

NH3 + OCl → NH2Cl + OH
NH2Cl + NH3N
2
H+
5
+ Cl (slow)
N
2
H+
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 chloramine (NH2Cl) to produce ammonium chloride and nitrogen. 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.[46] It is also commonly used as a rocket fuel.[47]

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 the hydrohalic acids.[46]

Halides and oxohalides

All four simple nitrogen trihalides are known. A few mixed halides and hydrohalides are known, but are mostly unstable and uninteresting: examples include NClF2, NCl2F, NBrF2, NF2H, NCl2H, and NClH2.[48]

Five nitrogen fluorides are known. 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
2
F+
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.[48]

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.[48][49]

Two series of nitrogen oxohalides are known: the nitrosyl halides (XNO) and the nitryl halides (XNO2). The first are 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.[48]

Oxides

Main article: Nitrogen oxide

Nitrogen forms nine molecular oxides, some of which were the first gases to be identified: N2O (nitrous oxide), NO (nitric oxide), N2O3 (dinitrogen trioxide), NO2 (nitrogen dioxide), N2O4 (dinitrogen tetroxide), N2O5 (dinitrogen pentoxide), NO3 (nitrogen trioxide), N4O (nitrosylazide),[50] and N(NO2)3 (trinitramide).[51] All are thermally unstable towards decomposition to their elements. One other possible oxide that has not yet been synthesised is oxatetrazole (N4O), an aromatic ring.[50]

Nitrous oxide (N2O), better known as laughing gas, is made by thermal decomposition of molten ammonium nitrate at 250 °C. This is a redox reaction and thus nitric oxide and nitrogen are also produced as byproducts. Despite appearances, it cannot be considered to be the 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+ON=N+=O): above 600 °C it dissociates by breaking the weaker N–O bond. It reacts with molten alkali metal amides at 200 °C to produce the ligands and, like nitrogen itself, can act as a ligand in [Ru(NH3)5(N2O)]2+. It is mostly used as a propellant and aerating agent for whipped ice cream, and was formerly commonly used as an anaesthetic.[50]

Nitric oxide (NO) is the least complicated molecule with an odd number of electrons that is stable. 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.[50]

Nitrogen dioxide at −196 °C, 0 °C, 23 °C, 35 °C, and 50 °C. (NO
2
) converts to the colorless dinitrogen tetroxide (N
2
O
4
) at low temperatures, and reverts to NO
2
at higher temperatures.

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 nitrosonium and nitrate in a medium with high dielectric constant. Nitrogen dioxide is an acrid, corrosive brown gas. Both compounds may be easily prepared by decomposing a dry metal nitrate. Both react with water to form nitric acid. Dinitrogen tetroxide is very useful for the preparation of anhydrous metal nitrates and nitrato complexes.[50]

The thermally unstable and very reactive dinitrogen pentoxide (N2O5) is the anhydride of nitric acid, and can be made from it by dehydration with phosphorus pentoxide. It is a deliquescent, 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 with hydrogen peroxide giving peroxonitric acid (HOONO2). It is a violent oxidising agent. Gaseous dinitrogen pentoxide decomposes as follows:[50]

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 solution 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
2
O2−
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.[52]

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
2
H+
5
react with nitrous acid to produce azides which further react to give nitrous oxide and nitrogen. Sodium nitrite is mildly toxic in concentrations above 100 mg/kg, but small amounts are often used to cure meat and as a presarvative to avoid bacterial spoilage. It is also used to synthesise hydroxylamine and to diazotise primary aromatic amines as follows:[52]

ArNH2 + HNO2 → [ArNN]Cl + 2 H2O

The reaction is conducted in acid solution. The reaction is different with aliphatic amines. Primary aliphatic amines result in alcohols and nitrogen:[52]

RNH2 + HNO2 −H2O  RNHNO → RN=NOH −OH  RN+
2
→ N2 + R+ → ROH

Secondary aliphatic amines react to form nitrosamines and no nitrogen:[52]

R2NH + HONO → R2NNO + H2O

Tertiary aliphatic amines react at low tepmeratures to give nitrite salts that decompose when warmed, giving the nitrosamine and the alcohol:[52]

R3N + HNO2 → [R3NH]+[NO2] → R2NNO + ROH

Nitrite is also a common ligand that can coordinate in five ways: nitro (M–N<⦂); nitrito (M–O–N=O); chelating (M<⦂>N); unsymmetrical bridging(N,O) (O=N–O, with M bonded to the single-bonded N and O), and η1-O bridging (with the single bonded O bonded to two M atoms). Nitro-nitrito isomerism is common, where the nitrito form is usually less stable.[52]

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 the alchemists in the 13th century. It is made by 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:[52]

2 HNO3H
2
NO+
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 gold, platinum, rhodium, and iridium are immune to attack. A 3:1 mixture of concentrated hydrochloric acid and nitric acid 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 concetrated sulfuric acid, nitric acid is protonated to form nitronium, which can act as an electrophile for aromatic nitration:[52]

HNO3 + 2 H2SO4NO+
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, including bidentate (M<⦂>N–O, symmetrical or unsymmetrical), unidentate (M–O–N<⦂), bridging syn-syn or anti-anti (with two of the oxygens bicoordinated to metal cations), μ2-O-bridging (with one of the oxygens bicoordinated), and possibly μ3-O-bridging (with one of the oxygens tricoordinated).[52]

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:[52]

NaNO3 + Na2O Ag crucible300 °C for 7 days Na3NO4

These white crystalline salts are very sensitive to water vapour and carbon dioxide in the air:[52]

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.[52]

Organic nitrogen compounds

Nitrogen is one of the most important elements in organic chemistry. Many organic functional groups involve a carbon–nitrogen bond, such as amides (RCONR2), amines (R3N), imines (RC(=NR)R), imides (RCO)2NR, azides (RN3), azo compounds (RN2R), cyanates and isocyanates (ROCN or RCNO), nitrates (RONO2), nitriles and isonitriles (RCN or RNC), nitrites (RONO), nitro compounds (RNO2), nitroso compounds (RNO), oximes (RCR=NOH), and pyridine derivatives. C–N bonds are strongly polarised towards nitrogen. In these compounds, nitrogen is usually trivalent (though it can be tetravalent in quaternary ammonium salts, R4N+), with a lone pair that can confer basicity on the compound by being coordinated to a proton. This may be offset by other factors: for example, amides are not basic because the lone pair is delocalised into a double bond (though they may act as acids at very low pH, being protonated at the oxygen), and pyrrole is not acidic because the lone pair is delocalised as part of an aromatic ring.[53] The amount of nitrogen in a chemical substance can be determined by the Kjeldahl method.[54] In particular, nitrogen is an essential component of nucleic acids, amino acids and thus proteins, and the energy-carrying molecule adenosine triphosphate and is thus vital to all life on Earth.[53]

Occurrence

Schematic representation of the flow of nitrogen compounds through a land environment

Nitrogen is the most common pure element in the earth, making up 78.1% of the entire volume of the atmosphere.[3] Despite this, it is not very abundant in Earth's crust, making up only 19 parts per million of this, on par with niobium, gallium, and lithium. The only important nitrogen minerals are nitre (potassium nitrate, saltpetre) and sodanitre (sodium nitrate, Chilean saltpetre). However, these are not an important source of nitrates any more since the 1920s, when the industrial synthesis of ammonia and nitric acid became common.[55]

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 own 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.[55]

Many saltwater fish manufacture large amounts of trimethylamine oxide to protect them from the high osmotic effects of their environment; conversion of this compound to dimethylamine is responsible for the early odour in unfresh saltwater fish.[56] In animals, free radical nitric oxide (derived from an amino acid), serves as an important regulatory molecule for circulation.[57]

Nitric oxide's rapid reaction with water in animals results in production of its metabolite nitrite. Animal metabolism of nitrogen in proteins, in general, results in excretion of urea, while animal metabolism of nucleic acids results in 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.[58]

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.[59] 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).[60] 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.[61]

In a chemical laboratory, it is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite.[62]

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.[62] Very pure nitrogen can be prepared by the thermal decomposition of barium azide or sodium azide.[63]

2 NaN3 → 2 Na + 3 N2

Applications

Gas

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 of nitrogen produced by industry is sold as the gas and the remaining one-third as the liquid. The gas is mostly used as an inert amtmosphere whenever the oxygen in the air would pose a fire, explosion, or oxidising hazard. Some examples include:[61]

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.[69]

Nitrogen can be used as a replacement, or in combination with, carbon dioxide to pressurise kegs of some beers, particularly stouts and British ales, due to the smaller bubbles it produces, which makes the dispensed beer smoother and headier.[70] A pressure-sensitive nitrogen capsule known commonly as a "widget" allows nitrogen-charged beers to be packaged in cans and bottles.[71][72] Nitrogen tanks are also replacing carbon dioxide as the main power source for paintball guns. Nitrogen must be kept at higher pressure than CO2, making N2 tanks heavier and more expensive.[73] Nitrogen gas has become the inert gas of choice for inert gas asphyxiation, and is under consideration as a replacement for lethal injection in Oklahoma.[74][75] Nitrogen gas, formed from the decomposition of sodium azide, is used for the inflation of airbags.[76]

Liquid

Air balloon submerged in liquid nitrogen

Liquid nitrogen is a cryogenic liquid. When insulated in proper containers such as Dewar flasks, it can be transported without much evaporative loss.[77]

Like dry ice, the main use of liquid nitrogen is as a refrigerant. Among other things, it is used in the cryopreservation of blood, reproductive cells (sperm and egg), and other biological samples and materials. It is used in the clinical setting in cryotherapy to remove cysts and warts on the skin.[78] It is used in cold traps for certain laboratory equipment and to cool infrared detectors or X-ray detectors. It has also been used to cool central processing units and other devices in computers that are overclocked, and that produce more heat than during normal operation.[79] 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 whenever necessary (around −200 °C). Because of its low cost, liquid nitrogen is also often used 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.[61]

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.[80] An example occurred shortly before the launch of the first Space Shuttle mission in 1981, when two technicians died from asphyxiation after they walked into a space located in the Shuttle's Mobile Launcher Platform that was pressurised with pure nitrogen as a precaution against fire.[81]

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.[82][83]

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.[84][85] 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.[86]

Liquid

As a cryogenic liquid, liquid nitrogen can be dangerous by causing cold burns on contact, although the Leidenfrost effect provides protection for very short exposure (about one second).[87] 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.[88]

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.[89]

As liquid nitrogen evaporates it reduces the oxygen concentration in the air and can act as an asphyxiant, especially in confined spaces. Nitrogen is odorless, colorless, and tasteless and may produce asphyxia without any sensation or prior warning.[90][91][92] For this reason, oxygen sensors are sometimes used as a safety precaution when working with liquid nitrogen to alert workers of gas spills into a confined space.[93]

Vessels containing liquid nitrogen can condense oxygen from air. The liquid in such a vessel becomes increasingly enriched in oxygen (boiling point −183 °C) as the nitrogen evaporates, and can cause violent oxidation of organic material.[94]

See also

References

  1. ^ Conventional Atomic Weights 2013. Commission on Isotopic Abundances and Atomic Weights
  2. ^ Standard Atomic Weights 2013. Commission on Isotopic Abundances and Atomic Weights
  3. ^ a b c Greenwood and Earnshaw, pp. 406–7
  4. ^ Rutherford, Daniel (1772) "Dissertatio Inauguralis de aere fixo, aut mephitico" (Inaugural dissertation on the air [called] fixed or mephitic), M.D. dissertation, University of Edinburgh, Scotland. English translation: Dobbin, Leonard (1935). "Daniel Rutherford's inaugural dissertation". Journal of Chemical Education. 12 (8): 370–375. doi:10.1021/ed012p370. 
  5. ^ Marshall, James R. and Marshall, Virginia L. (Spring 2015) "Rediscovery of the Elements: Daniel Rutherford, nitrogen, and the demise of phlogiston," The Hexagon (of Alpha Chi Sigma), 106 (1) : 4–8. Available on-line at: University of North Texas.
  6. ^ Lavoisier, Antoine Laurent (1965). Elements of chemistry, in a new systematic order: containing all the modern discoveries. Courier Dover Publications. p. 15. ISBN 0-486-64624-6. 
  7. ^ Weeks, Mary Elvira (1932). "The discovery of the elements. IV. Three important gases". Journal of Chemical Education. 9 (2): 215. Bibcode:1932JChEd...9..215W. doi:10.1021/ed009p215. 
  8. ^ Aaron J. Ihde, The Development of Modern Chemistry, New York 1964.
  9. ^ Carl Wilhelm Scheele, Chemische Abhandlung von der Luft und dem Feuer [Chemical treatise on air and fire] (Upsala, Sweden: Magnus Swederus, 1777 ; and Leipzig, (Germany): Siegfried Lebrecht Crusius, 1777). In the section titled "Die Luft muß aus elastischen Flüßigkeiten von zweyerley Art, zusammengesetzet seyn." (The air must be composed of elastic fluids of two sorts), pp. 6–14, Scheele presents the results of eight experiments in which air was reacted with various substances. He concluded (p. 13): "So viel sehe ich aus angeführten Versuchen, daß die Luft aus 2 von einander unterschiedenen Flußigkeiten bestehe, von welchen die eine die Eigenschaft das Phlogiston anzuziehen gar nicht äussere, die andere aber zur solchen Attraction eigentlich aufgeleget ist und welche zwischen dem 3:ten und 4:ten Theil von der ganzen Luftmasse aus machet." (So I see [this] much from the experiments [that were] conducted: that the air consists of two fluids [that] differ from one another, of which the one doesn't express at all the property of attracting phlogiston; the other, however, is capable of such attraction and which makes up between 1/3 and 1/4 part of the entire mass of the air.)
  10. ^ Priestley, Joseph (1772). "Observations on different kinds of air". Philosophical Transactions of the Royal Society of London. 62: 147–256. doi:10.1098/rstl.1772.0021.  ; see p. 225.
  11. ^ Priestley, Joseph (1772). "Observations on different kinds of air". Philosophical Transactions of the Royal Society of London. 62: 147–256. doi:10.1098/rstl.1772.0021.  ; see: "VII. Of air infected with the fumes of burning charcoal." pp. 225–228.
  12. ^ Elements of Chemistry, trans. Robert Kerr (Edinburgh, 1790; New York: Dover, 1965), p. 52.
  13. ^ Chaptal, J. A. and Nicholson, William trans. (1800) Elements of Chemistry, 3rd ed. London, England: C.C. and J. Robinson, vol. 1, pp. xxxv-xxxvi. From pp. xxxv-xxxvi: "In order to correct the Nomenclature on this head [i.e., in this regard], nothing more is necessary than to substitute to [i.e., for] this word a denomination which is derived from the general system made use of; and I have presumed to propose that of Nitrogene Gas. In the first place, it is deduced from the characteristic and exclusive property of this gas, which forms the radical of the nitric acid. By this means we shall preserve to the combinations [i.e., compounds] of this substance the received [i.e., prevailing] denominations, such as those of the Nitric Acid, Nitrates, Nitrites, &c."
  14. ^ nitrogen. Etymonline.com. Retrieved 2011-10-26.
  15. ^ Lord Rayleigh's Active Nitrogen. Lateralscience.co.uk. Retrieved 2011-10-26.
  16. ^ Erisman, Jan Willem; Sutton, Mark A.; Galloway, James; Klimont, Zbigniew; Winiwarter, Wilfried (2008). "How a century of ammonia synthesis changed the world". Nature Geoscience. 1 (10): 636. Bibcode:2008NatGe...1..636E. doi:10.1038/ngeo325. 
  17. ^ GB 190200698, Ostwald, Wilhelm, "Improvements in the Manufacture of Nitric Acid and Nitrogen Oxides", published January 9, 1902, issued March 20, 1902 
  18. ^ GB 190208300, Ostwald, Wilhelm, "Improvements in and relating to the Manufacture of Nitric Acid and Oxides of Nitrogen", published December 18, 1902, issued February 26, 1903 
  19. ^ a b c d e Greenwood and Earnshaw, pp. 411–2
  20. ^ Greenwood and Earnshaw, p. 550
  21. ^ a b Kaupp, Martin (1 December 2006). "The role of radial nodes of atomic orbitals for chemical bonding and the periodic table" (PDF). Journal of Computational Chemistry. 28 (1): 320–5. doi:10.1002/jcc.20522. Retrieved 14 October 2016. 
  22. ^ a b c d e f g h Greenwood and Earnshaw, pp. 412–6
  23. ^ Bethe, H. A. (1939). "Energy Production in Stars". Physical Review. 55 (5): 434–56. Bibcode:1939PhRv...55..434B. doi:10.1103/PhysRev.55.434. 
  24. ^ CIAAW (2003). "Atomic Weight of Nitrogen". ciaaw.org. CIAAW. Retrieved 13 October 2016. 
  25. ^ Flanagan, Lawrence B.; Ehleringer, James R; Pataki, Diane E. (15 December 2004). Stable Isotopes and Biosphere - Atmosphere Interactions: Processes and Biological Controls. pp. 74–75. ISBN 978-0-08-052528-0. 
  26. ^ Greenwood and Earnshaw, p. 408
  27. ^ "Evaluated Nuclear Data File (ENDF) Retrieval & Plotting". National Nuclear Data Center. 
  28. ^ Arthur G Palmer (2007). Protein NMR Spectroscopy. Elsevier Academic Press. ISBN 0-12-164491-X. 
  29. ^ a b c Audi, G.; Wapstra, A. H.; Thibault, C.; Blachot, J. & Bersillon, O. (2003). "The NUBASE evaluation of nuclear and decay properties" (PDF). Nuclear Physics A. 729: 3–128. Bibcode:2003NuPhA.729....3A. doi:10.1016/j.nuclphysa.2003.11.001. [dead link]
  30. ^ a b Neeb, Karl Heinz (1997). The Radiochemistry of Nuclear Power Plants with Light Water Reactors. Berlin-New York: Walter de Gruyter. p. 227. ISBN 3-11-013242-7. 
  31. ^ Greenwood and Earnshaw, pp. 434–8
  32. ^ a b Lewars, Errol G. (2008). Modeling Marvels: Computational Anticipation of Novel molecules. Springer Science+Business Media. pp. 141–163. doi:10.1007/978-1-4020-6973. ISBN 978-1-4020-6972-7. 
  33. ^ Greenwood and Earnshaw, p. 483
  34. ^ "A new molecule and a new signature – Chemistry – tetranitrogen". Science News. 16 February 2002. Retrieved 2016-09-30. 
  35. ^ "Polymeric nitrogen synthesized". physorg.com. 5 August 2004. Retrieved 2009-06-22. 
  36. ^ Gray, Theodore (2009). The Elements: A Visual Exploration of Every Known Atom in the Universe. New York: Black Dog & Leventhal Publishers. ISBN 978-1-57912-814-2. 
  37. ^ Schuch, A. F.; Mills, R. L. (1970). "Crystal Structures of the Three Modifications of Nitrogen 14 and Nitrogen 15 at High Pressure". The Journal of Chemical Physics. 52 (12): 6000–6008. Bibcode:1970JChPh..52.6000S. doi:10.1063/1.1672899. 
  38. ^ Iancu, C. V.; Wright, E. R.; Heymann, J. B.; Jensen, G. J. (2006). "A comparison of liquid nitrogen and liquid helium as cryogens for electron cryotomography". Journal of Structural Biology. 153 (3): 231–240. doi:10.1016/j.jsb.2005.12.004. PMID 16427786. 
  39. ^ "Flowing nitrogen ice glaciers seen on surface of Pluto after New Horizons flyby". ABC. 25 July 2015. Retrieved 6 October 2015. 
  40. ^ McKinnon, William B.; Kirk, Randolph L. (2014). "Triton". In Spohn, Tilman; Breuer, Doris; Johnson, Torrence. Encyclopedia of the Solar System (3rd ed.). Amsterdam; Boston: Elsevier. pp. 861–882. ISBN 978-0-12-416034-7. 
  41. ^ "Neptune: Moons: Triton". NASA. Archived from the original on October 5, 2011. Retrieved September 21, 2007. 
  42. ^ Fryzuk, M. D. & Johnson, S. A. (2000). "The continuing story of dinitrogen activation". Coordination Chemistry Reviews. 200–202: 379. doi:10.1016/S0010-8545(00)00264-2. 
  43. ^ Schrock, R. R. (2005). "Catalytic Reduction of Dinitrogen to Ammonia at a Single Molybdenum Center". Acc. Chem. Res. 38 (12): 955–962. doi:10.1021/ar0501121. PMC 2551323Freely accessible. PMID 16359167. 
  44. ^ a b c d e Greenwood and Earnshaw, pp. 417–20
  45. ^ Greenwood and Earnshaw, pp. 420–6
  46. ^ a b c Greenwood and Earnshaw, pp. 426–33
  47. ^ Vieira, R.; C. Pham-Huu; N. Keller; M. J. Ledoux (2002). "New carbon nanofiber/graphite felt composite for use as a catalyst support for hydrazine catalytic decomposition". Chemical Communications (9): 954–955. doi:10.1039/b202032g. 
  48. ^ a b c d Greenwood and Earnshaw, pp. 438–42
  49. ^ Bowden, F. P. (1958). "Initiation of Explosion by Neutrons, α-Particles, and Fission Products". Proceedings of the Royal Society of London A. 246 (1245): 216–219. doi:10.1098/rspa.1958.0123. 
  50. ^ a b c d e f Greenwood and Earnshaw, pp. 443–58
  51. ^ Rahm, Martin; Dvinskikh, Sergey V.; Furó, István; Brinck, Tore (23 December 2010). "Experimental Detection of Trinitramide, N(NO2)3". Angewandte Chemie International Edition. 50 (5): 1145–8. doi:10.1002/anie.201007047. 
  52. ^ a b c d e f g h i j k l Greenwood and Earnshaw, pp. 459–72
  53. ^ a b March, Jerry (1985), Advanced Organic Chemistry: Reactions, Mechanisms, and Structure (3rd ed.), New York: Wiley, ISBN 0-471-85472-7 
  54. ^ "Kjeldahl Method". Encyclopedia of Genetics, Genomics, Proteomics and Informatics. 2008. p. 1063. doi:10.1007/978-1-4020-6754-9_9066. ISBN 978-1-4020-6753-2. 
  55. ^ a b Greenwood and Earnshaw, pp. 407–9
  56. ^ Nielsen, M. K.; Jørgensen, B. M. (Jun 2004). "Quantitative relationship between trimethylamine oxide aldolase activity and formaldehyde accumulation in white muscle from gadiform fish during frozen storage". Journal of Agricultural and Food Chemistry. 52 (12): 3814–3822. doi:10.1021/jf035169l. PMID 15186102. 
  57. ^ Knox, G. A. (2007). Biology of the Southern Ocean. CRC Press. p. 392. ISBN 0-8493-3394-6. 
  58. ^ Vickerstaff Joneja; Janice M. (2004). Digestion, diet, and disease: irritable bowel syndrome and gastrointestinal function. Rutgers University Press. p. 121. ISBN 0-8135-3387-2. 
  59. ^ Froehlich, Peter (May 2013). "A Sustainable Approach to the Supply of Nitrogen". www.parker.com. Parker Hannifin Corporation. Retrieved 24 November 2016. 
  60. ^ Reich, Murray; Kapenekas, Harry (1957). "Nitrogen Purfication. Pilot Plant Removal of Oxygen". Industrial & Engineering Chemistry. 49 (5): 869–873. doi:10.1021/ie50569a032. 
  61. ^ a b c d e Greenwood and Earnshaw, pp. 409–11
  62. ^ a b Bartlett, J. K. (1967). "Analysis for nitrite by evolution of nitrogen: A general chemistry laboratory experiment". Journal of Chemical Education. 44 (8): 475. Bibcode:1967JChEd..44..475B. doi:10.1021/ed044p475. 
  63. ^ Eremets, M. I.; Popov, M. Y.; Trojan, I. A.; Denisov, V. N.; Boehler, R.; Hemley, R. J. (2004). "Polymerization of nitrogen in sodium azide". The Journal of Chemical Physics. 120 (22): 10618–10623. Bibcode:2004JChPh.12010618E. doi:10.1063/1.1718250. PMID 15268087. 
  64. ^ Ministers, Nordic Council of (2002). "Food Additives in Europe 2000": 591. ISBN 978-92-893-0829-8. 
  65. ^ Harding, Charlie, ed. (2002). Elements of the p Block. Cambridge: Royal Society of Chemistry. ISBN 978-0-85404-690-4. 
  66. ^ Gavriliuk, V. G.; Berns, Hans (1999). High nitrogen steels: structure, properties, manufacture, applications. Springer. ISBN 3-540-66411-4. 
  67. ^ "Centre Fuel Tank Inerting". B737.org.uk. Retrieved 2013-08-21. 
  68. ^ "Why don't they use normal air in race car tires?". Howstuffworks. Retrieved 2006-07-22. 
  69. ^ Kemmochi, Y; Tsutsumi, K; Arikawa, A; Nakazawa, H (2002). "Centrifugal concentrator for the substitution of nitrogen blow-down micro-concentration in dioxin/polychlorinated biphenyl sample preparation". Journal of Chromatography A. 943 (2): 295–297. doi:10.1016/S0021-9673(01)01466-2. PMID 11833649. 
  70. ^ Baxter, E. Denise; Hughes, Paul S. (2001). Beer: Quality, Safety and Nutritional Aspects. Royal Society of Chemistry. p. 22. ISBN 978-0-85404-588-4. 
  71. ^ "How does the widget in a beer can work?". Howstuffworks. 
  72. ^ Denny, Mark (1 November 2009). Froth!: The Science of Beer. p. 131. ISBN 978-0-8018-9569-2. 
  73. ^ Kennett, Andrew J. (2008). "Design of a pneumatically assisted shifting system for Formula SAE® racing applications". Dept. of Mechanical Engineering, Massachusetts Institute of Technology. hdl:1721.1/45820. 
  74. ^ Sanburn, Josh (2015-04-10). "The Dawn of a New Form of Capital Punishment". Time. Retrieved 2015-04-11. 
  75. ^ Sexton, Mike (18 December 2012). "Euthanasia campaigner under scrutiny". ABC. Retrieved 6 May 2013. 
  76. ^ Betterton, E. A. (2003). "Environmental Fate of Sodium Azide Derived from Automobile Airbags". Critical Reviews in Environmental Science and Technology. 33 (4): 423–458. doi:10.1080/10643380390245002. 
  77. ^ Kaganer, M. G.; Kozheurov, V. & Levina, Zh. L. (1967). "Vessels for the storage and transport of liquid oxygen and nitrogen". Chemical and Petroleum Engineering. 3 (12): 918–922. doi:10.1007/BF01136404. 
  78. ^ Ahmed I; Agarwal S; Ilchyshyn A; Charles-Holmes S; Berth-Jones J (May 2001). "Liquid nitrogen cryotherapy of common warts: cryo-spray vs. cotton wool bud". Br. J. Dermatol. 144 (5): 1006–9. doi:10.1046/j.1365-2133.2001.04190.x. PMID 11359389. 
  79. ^ Kent, Allen; Williams, James G. (1994). Encyclopedia of Computer Science and Technology. 30. CRC Press. p. 318. ISBN 0-8247-2283-3. 
  80. ^ "Biology Safety – Cryogenic materials. The risks posed by them.". University of Bath. Archived from the original on February 6, 2007. Retrieved 2007-01-03. 
  81. ^ "Space Shuttle Columbia Fast Facts". CNN. September 30, 2013. 
  82. ^ Fowler, B.; Ackles, K.N.; Porlier, G. (1985). "Effects of inert gas narcosis on behavior—a critical review". Undersea Biomed. Res. 12 (4): 369–402. PMID 4082343. Retrieved 2008-09-21. 
  83. ^ Rogers, W. H.; Moeller, G. (1989). "Effect of brief, repeated hyperbaric exposures on susceptibility to nitrogen narcosis". Undersea Biomed. Res. 16 (3): 227–32. OCLC 2068005. PMID 2741255. Retrieved 2008-09-21. 
  84. ^ Acott, C. (1999). "A brief history of diving and decompression illness". South Pacific Underwater Medicine Society Journal. 29 (2). OCLC 16986801. Retrieved 2008-09-21. 
  85. ^ Kindwall, E. P.; Baz, A.; Lightfoot, E. N.; Lanphier, E. H.; Seireg, A. (1975). "Nitrogen elimination in man during decompression". Undersea Biomed. Res. 2 (4): 285–97. OCLC 2068005. PMID 1226586. Retrieved 2008-09-21. 
  86. ^ US Navy Diving Manual, 6th revision. United States: US Naval Sea Systems Command. 2006. Retrieved 2008-04-24. 
  87. ^ Walker, Jearl. "Boiling and the Leidenfrost Effect" (PDF). Fundamentals of Physics: 1–4. Retrieved 11 October 2014. 
  88. ^ Liquid nitrogen cocktail leaves teen in hospital, BBC News, October 8, 2012.
  89. ^ Mattox, Brent S. "Investigative Report on Chemistry 301A Cylinder Explosion" (reprint). Texas A&M University. 
  90. ^ British Compressed Gases Association (2000) BCGA Code of Practice CP30. The Safe Use of Liquid nitrogen Dewars up to 50 litres. ISSN 0260-4809.
  91. ^ Confined Space Entry - Worker and Would-be Rescuer Asphyxiated, Valero Refinery Asphyxiation Incident Case Study.
  92. ^ Inquiry after man dies in chemical leak, BBC News, October 25, 1999.
  93. ^ Liquid Nitrogen – Code of practice for handling. United Kingdom: Birkbeck, University of London. 2007. Retrieved 2012-02-08. 
  94. ^ Levey, Christopher G. "Liquid Nitrogen Safety". Thayer School of Engineering at Dartmouth. 

Bibliography

External links