Nitrogen

Nitrogen,  ₇N

Liquid nitrogen
Spectral lines of nitrogen
General properties
Name, symbol	nitrogen, N
Appearance	colorless gas, liquid or solid
Pronunciation	/ˈnaɪtrədʒən/ NY-trə-jən
Nitrogen in the periodic table
– ↑ N ↓ P carbon ← nitrogen → oxygen
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] 2s^2 2p^3
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/cm^3
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 vaporization	(N2) 5.56 kJ/mol
Molar heat capacity	(N2) 29.124 J/(mol·K)
vapor 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
Ionization 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
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 99.634% ^14N is stable with 7 neutrons ^15N 0.366% ^15N is stable with 8 neutrons

Nitrogen is a chemical element with symbol N and atomic number 7. It is the lightest pnictogen and at room temperature, it is a transparent, odorless diatomic gas. Nitrogen is a common element in the universe, estimated at about seventh in total abundance in the Milky Way and the Solar System. On Earth, the element forms about 78% of Earth's atmosphere and is the most abundant uncombined element. The element nitrogen was discovered as a separable component of air by Scottish physician Daniel Rutherford in 1772.

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) dominates nitrogen chemistry, causing difficulty for both organisms and industry in converting the 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 fertilizers, and fertilizer nitrates are key pollutants in the eutrophication of water systems.

Apart from its use in fertilizers and energy-stores, nitrogen is a constituent of organic compounds as diverse as Kevlar fabric and cyanoacrylate "super" glue. 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 proteins), 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 and etymology

Nitrogen is formally considered to have been discovered by Scottish physician Daniel Rutherford in 1772, who called it noxious air.^[3][4][5] Though he did not recognize it as an entirely different chemical substance, he clearly distinguished it from Joseph Black's "fixed air", or carbon dioxide.^[6] The fact that there was a component of air that does not support combustion was clear to Rutherford. Nitrogen was also studied at about the same time by Carl Wilhelm Scheele,^[7] Henry Cavendish,^[8] and Joseph Priestley,^[9] 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 ἄζωτος azotos, "lifeless".^[10] 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.

Lavoisier's name for nitrogen is used in many languages (French, Italian, Portuguese, Polish, Russian, Albanian, Turkish, etc.) and still remains in English in the common names of many compounds, such as hydrazine and compounds of the azide ion. The English word nitrogen (1794) entered the language from the French nitrogène, coined in 1790 by French chemist Jean-Antoine Chaptal (1756–1832),^[11] from the French nitre (potassium nitrate, also called saltpeter) and the French -gène, "producing", from the Greek -γενής (-genes, "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 nitron (νίτρον) in Greek — which contained no nitrate.^[12]

Nitrogen compounds were well known by the Middle Ages. Alchemists knew nitric acid as aqua fortis (strong water). 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). The earliest military, industrial, and agricultural applications of nitrogen compounds used saltpeter (sodium nitrate or potassium nitrate), most notably in gunpowder, and later as fertilizer. 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 quicksilver to produce explosive mercury nitride.^[13]

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 fertilizers.^[14] 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.

Production

Nitrogen gas is an industrial gas produced by the fractional distillation of liquid air, or by mechanical means using gaseous air (i.e., pressurized reverse osmosis membrane or pressure swing adsorption). Nitrogen gas generators using membranes or PSA are typically more cost and energy efficient than bulk delivered nitrogen.^[15] 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).^[16]

In a chemical laboratory, it is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite.^[17]

NH₄Cl(aq) + NaNO₂(aq) → N₂(g) + NaCl(aq) + 2 H₂O (l)

Small amounts of 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.^[17] Very pure nitrogen can be prepared by the thermal decomposition of barium azide or sodium azide.^[18]

2 NaN₃ → 2 Na + 3 N₂

Properties

Nitrogen is a nonmetal, with an electronegativity of 3.04.^[19] It has five electrons in its outer shell and is, therefore, trivalent in most compounds. The triple bond in molecular nitrogen (N
2) is one of the strongest. The resulting difficulty of converting N
2 into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N
2, have dominated the role of nitrogen in both nature and human economic activities.

At atmospheric pressure, molecular nitrogen condenses (liquefies) at 77 K (−195.79 °C) and freezes at 63 K (−210.01 °C)^[20] 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).^[21] Liquid nitrogen, a 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.^[22]

Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like N
3 and N
4.^[23] Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced in a diamond anvil cell, nitrogen polymerizes into the single-bonded cubic gauche crystal structure. This structure is similar to that of diamond, and both have extremely strong covalent bonds. N
4 is nicknamed "nitrogen diamond".^[24]

Other (theoretical but not yet synthesized) allotropes include hexazine (N
6, a benzene analog)^[25] and octaazacubane (N
8, a cubane analog).^[26] The former is predicted to be highly unstable, while the latter is predicted to be kinetically stable, by reason of its orbital symmetry.^[27]

Isotopes

There are two stable isotopes of nitrogen: ¹⁴N and ¹⁵N. By far the most common is ¹⁴N (99.634%), which is produced in the CNO cycle in stars.^[28] Of the ten isotopes produced synthetically, ¹³N has a half-life of ten minutes and the remaining isotopes have half-lives on the order of seconds or less.^[29]

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.^[30]

A small part (0.73%) of the molecular nitrogen in Earth's atmosphere is the isotopologue ¹⁴N¹⁵N, and almost all the rest is ¹⁴N₂.^[31]

The radioisotope ¹⁶N is the dominant radionuclide in the coolant of pressurized water reactors or boiling water reactors during normal operation. It is produced from ¹⁶O (in water) via (n,p) reaction. 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][32]

Because of this, access to the primary coolant piping in a pressurized 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.^[32]

Similarly, access to any of the steam cycle components in a boiling water reactor nuclear power plant must be restricted during operation. Condensate is typically retained for 10 minutes to allow for decay of ¹⁶N to radioactive depletion.

Electromagnetic spectrum

Nitrogen discharge (spectrum) tube

Molecular nitrogen (¹⁴N₂) is largely transparent to infrared and visible radiation because it is a homonuclear molecule and has no dipole moment to couple to electromagnetic radiation at these wavelengths. Significant absorption occurs at extreme ultraviolet wavelengths,^[33] beginning around 100 nanometres. This is associated with electronic transitions in the molecule to states in which charge is not distributed evenly between nitrogen atoms. Nitrogen absorption leads to significant absorption of ultraviolet radiation in the Earth's upper atmosphere and the atmospheres of other planetary bodies. For similar reasons, pure molecular nitrogen lasers typically emit light in the ultraviolet range.

Nitrogen also makes a contribution to visible air glow from the Earth's upper atmosphere, through electron impact excitation followed by emission. This visible blue air glow (seen in the polar aurora and in the re-entry glow of returning spacecraft) typically results not from molecular nitrogen but rather from free nitrogen atoms combining with oxygen to form nitric oxide (NO).

Nitrogen gas also exhibits scintillation.

Reactions

Structure of dinitrogen, N₂

Structure of [Ru(NH₃)₅(N₂)]²⁺

In general, nitrogen is unreactive at standard temperature and pressure. N₂ reacts spontaneously with few reagents, being resilient to acids and bases as well as oxidants and most reductants. When nitrogen reacts spontaneously with a reagent, the net transformation is often called nitrogen fixation.

Elemental lithium reacts ("burns") in an atmosphere of N₂ to give lithium nitride:^[34]

6 Li + N₂ → 2 Li₃N

Magnesium also burns in nitrogen, forming magnesium nitride.

3 Mg + N₂ → Mg₃N₂

N₂ forms a variety of adducts with transition metals. The first example of a dinitrogen complex is [Ru(NH₃)₅(N₂)]²⁺ (see figure at right). Notably, N₂ ligand was obtained by the decomposition of hydrazine, 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.^[35] 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.^[34]

The starting point for industrial production of nitrogen compounds is the Haber process, in which nitrogen is fixed by reacting N
2 and H
2 over an iron(II, III) oxide (Fe
3O
4) catalyst at about 500 °C and 200 atmospheres pressure. Biological nitrogen fixation in free-living cyanobacteria and in the root nodules of plants also produces ammonia from molecular nitrogen. The reaction, which is the source of the bulk of nitrogen in the biosphere, is catalyzed by the nitrogenase enzyme complex that contains Fe and Mo atoms, using energy derived from hydrolysis of adenosine triphosphate (ATP) into adenosine diphosphate and inorganic phosphate (−20.5 kJ/mol).

Occurrence

Nitrogen gas (N₂) is the largest constituent of Earth's atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air).^[36] However, this high concentration does not reflect nitrogen's overall low abundance in the makeup of the Earth, from which most of the element escaped by solar evaporation, early in the planet's formation.

Nitrogen is a common element in the universe, and is considered the seventh most abundant chemical element by mass in the universe, the Milky Way, and the Solar System. According to physical cosmology theory, it was originally created by fusion processes from carbon and hydrogen in supernovas.^[37] Molecular nitrogen and nitrogen compounds have been detected in interstellar space by astronomers using the Far Ultraviolet Spectroscopic Explorer.^[38]

Due to the volatility of elemental nitrogen and its common compounds with hydrogen and oxygen, nitrogen was driven out of the planetesimals in the early Solar System by the heat of the Sun and lost to the rocky planets of the inner Solar System. Nitrogen is consequently a relatively rare element on these inner planets, including Earth. Neon, which occurs in equal abundance in the universe, is rare in the inner Solar System for the same reason. Nitrogen is estimated at 31st element in crustal abundance. Nitrogen also occurs in some relatively uncommon minerals, such as saltpeter (potassium nitrate), Chile saltpeter (sodium nitrate) and sal ammoniac (ammonium chloride). Because of the ready solubility of naturally-occurring nitrogen compounds, these minerals were concentrated from the crust by the leaching of rivers and collected in ancient seas which then evaporated, leaving the deposits of present day. A similar process occurred with the water-soluble compounds of the uncommon light element boron.

Nitrogen and its compounds occur far more commonly as gases in the atmospheres of planets and moons that are large enough to have atmospheres.^{[lower-alpha 1]} Molecular nitrogen is a major constituent of not only Earth's atmosphere, but also the Saturnian moon Titan's thick atmosphere. Also, held by gravity at colder temperatures, nitrogen and its compounds occur in variable amounts in the planetary atmospheres of the gas giant planets.^[39]

Nitrogen is present in all known living organisms, in proteins, nucleic acids, and other molecules. It typically makes up around 4% of the dry weight of plant matter, and around 3% of the weight of the human body. It is a large component of animal waste (for example, guano), usually in the form of urea, uric acid, ammonium compounds, and derivatives of these nitrogenous products, which are essential nutrients for all plants that cannot fix atmospheric nitrogen.

Compounds

The main neutral hydride of nitrogen is ammonia (NH
3), although hydrazine (N
2H
4) is also commonly used. Ammonia is more basic than water by 6 orders of magnitude. In solution, ammonia forms the ammonium ion (NH+
4). Liquid ammonia (boiling point 240 K) is amphiprotic (displaying either Brønsted–Lowry acidic or basic character) and forms ammonium and the less common amide ions (NH−
2); both amides and nitride (N3−
) salts are known, Many of these salts decompose in water, but there are exceptions like gallium nitride or boron nitride that have strong polar covalent network structures. Singly, doubly, triply and quadruply substituted alkyl compounds of ammonia are called amines (four substitutions, to form commercially and biologically important quaternary amines, results in a positively charged nitrogen, and thus a water-soluble, or at least amphiphilic, compound). Larger chains, rings and structures of nitrogen hydrides are also known, but are generally unstable.

Other classes of nitrogen anions (negatively charged ions) are the poisonous azides (N−
3), which are linear and isoelectronic to carbon dioxide, but which bind to important iron-containing enzymes in the body in a manner more resembling cyanide. Another molecule of the same structure is the colorless and relatively inert anesthetic gas Nitrous oxide (dinitrogen monoxide, N
2O), also known as laughing gas. This is one of a variety of nitrogen oxides that form a family often abbreviated as NOx. Nitric oxide (nitrogen monoxide, NO), is a natural free radical used in signal transduction in both plants and animals. It is used, for example, in vasodilation, causing the smooth muscle of blood vessels to relax. The reddish and poisonous nitrogen dioxide NO
2 contains an unpaired electron and is an important component of smog. Nitrogen molecules containing unpaired electrons show a tendency to dimerize (thus pairing the electrons), and are, in general, highly reactive. The corresponding acids are nitrous HNO
2 and nitric acid HNO
3, with the corresponding salts called nitrites and nitrates.

The higher oxides dinitrogen trioxide N
2O
3, dinitrogen tetroxide N
2O
4 and dinitrogen pentoxide N
2O
5, are unstable and explosive, a consequence of the low energy state of N
2. Nearly every hypergolic rocket engine uses N
2O
4 as the oxidizer; the fuels, one of the various forms of hydrazine, are also nitrogen compounds. These engines are extensively used on spacecraft such as the space shuttle and those of the Apollo Program because their propellants are liquids at room temperature and ignition occurs spontaneously without an ignition system, allowing many precisely controlled burns. Some launch vehicles such as the Titan II and Ariane 1 through 4 also use hypergolic fuels, although the trend is away from such engines for cost and safety reasons. N
2O
4 is an intermediate in the manufacture of nitric acid HNO
3, a strong acid and a fairly strong oxidizing agent.

Nitrogen is notable for the range of explosively unstable compounds. Nitrogen triiodide NI
3 is an extremely sensitive contact explosive. Nitrocellulose, produced by nitration of cellulose with nitric acid, is also known as guncotton. Nitroglycerin, made by nitration of glycerin, is the dangerously unstable explosive ingredient of dynamite. The comparatively stable, but less powerful explosive trinitrotoluene (TNT) is the standard explosive against which the power of nuclear explosions are measured.^[40]

Nitrogen can also be found in many organic compounds. Common nitrogen functional groups include: amines, amides, nitro groups, imines, and enamines. The amount of nitrogen in a chemical substance can be determined by the Kjeldahl method.

Applications

Nitrogen gas

Nitrogen gas has a variety of applications, including serving as an inert replacement for air where oxidation is undesirable;^[41]

• As a modified atmosphere, pure or mixed with carbon dioxide, to nitrogenate and preserve the freshness of packaged or bulk foods (by delaying rancidity and other forms of oxidative damage). Pure nitrogen as food additive is labeled in the European Union with the E number E941.^[42]
• In incandescent light bulbs as an inexpensive alternative to argon.^[43]
• In fire suppression systems for Information technology (IT) equipment
• In photolithography in deep ultraviolet, nitrogen is used to avoid the strong oxygen absorption of UV at these wavelengths.
• Dried and pressurized, it is used as an dielectric gas for high voltage equipment.
• In the manufacture of stainless steel.^[44]
• In some aircraft fuel systems to reduce fire hazard, (see inerting system).^[45]
• In packaging liquid explosives as a safety measure.
• To inflate race car and aircraft tires,^[46] reducing the problems caused by moisture and oxygen in natural air.

Nitrogen is commonly used during sample preparation in chemical analysis. It is used to concentrate and reduce the volume of liquid samples. Directing a pressurized 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.^[47]

Nitrogen can be used as a replacement, or in combination with, carbon dioxide to pressurize kegs of some beers, particularly stouts and British ales, due to the smaller bubbles it produces, which makes the dispensed beer smoother and headier.^[48] A pressure-sensitive nitrogen capsule known commonly as a "widget" allows nitrogen-charged beers to be packaged in cans and bottles.^[49][50]

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.^[51]

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.^[52] Nitrogen is promoted by euthanasia advocate Philip Nitschke to end life in a "peaceful, reliable [and] totally legal" manner.^[53]

Liquid nitrogen

Air balloon submerged in liquid nitrogen

Liquid nitrogen is a cryogenic liquid. At atmospheric pressure, it boils at −195.8 °C (−320.4 °F). When insulated in proper containers such as Dewar flasks, it can be transported without much evaporative loss.^[54]

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.^[55] 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.^[56]

Solid nitrogen

The solid form of the element nitrogen forms a significant dynamic surface coverage on Pluto^[57] and outer moons of the Solar System such as Triton.^[58] 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.^[59]

There are a variety of crystalline forms of solid nitrogen known, not all with discrete dinitrogen molecular structures. Such non-molecular forms of solid nitrogen are of interest as they possess a higher energy density than any other non-nuclear material.

Nitrogen compounds

Molecular nitrogen (N₂) in the atmosphere is relatively non-reactive due to its strong triple bond, N≡N, and molecular nitrogen plays an inert role in the human body, being neither produced nor destroyed. In nature, nitrogen is converted into biologically (and industrially) useful compounds by lightning, and by some living organisms, notably certain bacteria (i.e., nitrogen-fixing bacteria—see Biological role below). Molecular nitrogen is released into the atmosphere in the process of decay, in dead plant and animal tissues.

The ability to combine, or fix, molecular nitrogen is a key feature of modern industrial chemistry. Previously to the 20th century, access to nitrogen compounds for fertilizers and gunpowder had been through deposits of natural nitrates, such as Chilean saltpeter. However, first the Frank–Caro process for producing cyanamide, and then the Haber–Bosch process for producing ammonia from air and natural gas^[20] (developed just before the first world war) eased this shortage of nitrogen compounds, to the extent that half of global food production now relies on synthetic nitrogen fertilizers.

The Ostwald process, developed a few years before the Haber process, allowed large-scale production of nitric acid and nitrate from ammonia, thus freeing large-scale industrial production of nitrate explosives and weapons propellants from the need to mine nitrate salt deposits.^[41] The organic and inorganic salts of nitric acid have been important historically as convenient stores of chemical energy for warfare and rocket fuels. Historically, such compounds included important compounds such as potassium nitrate, used in gunpowder^[60] which was often produced by biological means (bacterial fermentation) before natural mineral sources were discovered. Later, all such sources were displaced by industrial production, in the early 1900s.

Table tennis ball made from nitrocellulose.

Ammonium nitrate has been used as both fertilizer^[36] and explosive (see ANFO). Various other nitrated organic compounds, such as nitroglycerin, trinitrotoluene,^[60] and nitrocellulose,^[61] are used as explosives and propellants for modern firearms. Nitric acid is used as an oxidizing agent in liquid fueled rockets. Hydrazine and hydrazine derivatives find use as rocket fuels and monopropellants. In most of these compounds, the basic instability and tendency to burn or explode is derived from the fact that nitrogen is present as an oxide, and not as the far more stable nitrogen molecule (N₂), which is a product of the compounds' thermal decomposition. When nitrates burn or explode, the formation of the powerful triple bond in the N₂ produces most of the energy of the reaction.

Nitrogen is a constituent of molecules in every major drug class in pharmacology and medicine. Nitrous oxide (N₂O) was discovered early in the 19th century to be a partial anesthetic, though it was not used as a surgical anesthetic until later. Called "laughing gas", it was found capable of inducing a state of social disinhibition resembling drunkenness. Other notable nitrogen-containing drugs are drugs derived from plant alkaloids, such as morphine (there exist many alkaloids known to have pharmacological effects; in some cases, they appear as natural chemical defenses of plants against predation). Drugs that contain nitrogen include all major classes of antibiotics and organic nitrate drugs like nitroglycerin and nitroprusside that regulate blood pressure and heart action by mimicking the action of nitric oxide.

Biological role

Nitrogen is an essential building block of amino and nucleic acids, essential to life on Earth.^[36]

Elemental nitrogen in the atmosphere cannot be used directly by either plants or animals, and must be converted to a reduced (or 'fixed') state to be useful for higher plants and animals. Precipitation often contains substantial quantities of ammonium and nitrate, thought to result from nitrogen fixation by lightning and other atmospheric electric phenomena.^[62] This was first proposed by Liebig in 1827 and later confirmed.^[62] However, because ammonium is preferentially retained by the forest canopy relative to atmospheric nitrate, most fixed nitrogen reaches the soil surface under trees as nitrate. Soil nitrate is preferentially assimilated by tree roots relative to soil ammonium, however this is not the case for all species.^[63][64]

Specific bacteria (e.g., Rhizobium trifolium) possess nitrogenase enzymes that can fix atmospheric nitrogen (see nitrogen fixation) into a form (ammonium ion) that is chemically useful to higher organisms. This process requires a large amount of energy and anoxic conditions. Such bacteria may live freely in soil (e.g., Azotobacter) but normally exist in a symbiotic relationship in the root nodules of leguminous plants (e.g. clover, Trifolium, or soybean plant, Glycine max) and fertilizer trees. Nitrogen-fixing bacteria are also symbiotic with a number of unrelated plant species such as alders (Alnus) spp., lichens, Casuarina, Myrica, liverworts, and Gunnera.^[65]

As part of the symbiotic relationship, the plant converts the 'fixed' ammonium ion to nitrogen oxides and amino acids to form proteins and other molecules, (e.g., alkaloids). In return for the 'fixed' nitrogen, the plant secretes sugars to the symbiotic bacteria.^[65] Legumes maintain an anaerobic (oxygen free) environment for their nitrogen-fixing bacteria.

Plants are able to assimilate nitrogen directly in the form of nitrates that may be present in soil from natural mineral deposits, artificial fertilizers, animal waste, or organic decay (as the product of bacteria, but not bacteria specifically associated with the plant). Nitrates absorbed in this fashion are converted to nitrites by the enzyme nitrate reductase, and then converted to ammonia by another enzyme called nitrite reductase.^[65]

Nitrogen compounds are basic building blocks in animal biology as well. Animals use nitrogen-containing amino acids from plant sources as starting materials for all nitrogen-compound animal biochemistry, including the manufacture of proteins and nucleic acids. Plant-feeding insects are dependent on nitrogen in their diet, such that varying the amount of nitrogen fertilizer applied to a plant can affect the reproduction rate of insects feeding on fertilized plants.^[66]

Soluble nitrate is an important limiting factor in the growth of certain bacteria in ocean waters.^[67] In many places in the world, artificial fertilizers applied to crop-lands to increase yields result in run-off delivery of soluble nitrogen to oceans at river mouths. This process can result in eutrophication of the water, as nitrogen-driven bacterial growth depletes water oxygen to the point that all higher organisms die. Well-known "dead zone" areas in the U.S. Gulf Coast and the Black Sea are due to this important polluting process.

Many saltwater fish manufacture large amounts of trimethylamine oxide to protect them from the high osmotic effects of their environment; conversion of this compound to dimethylamine is responsible for the early odor in unfresh saltwater fish.^[68] In animals, free radical nitric oxide (NO) (derived from an amino acid), serves as an important regulatory molecule for circulation.^[67]

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 odor 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.^[69]

Decay of organisms and their waste products may produce small amounts of nitrate, but most decay eventually returns nitrogen content to the atmosphere, as molecular nitrogen. The circulation of nitrogen from atmosphere, to organic compounds, then back to the atmosphere, is referred to as the nitrogen cycle.^[65]

Safety

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.^[70] 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 pressurized with pure nitrogen as a precaution against fire.^[71]

A closed space, initially filled with air, will ultimately contain only nitrogen, carbon dioxide, and trace gases if some oxidative process, such as the rusting of iron, consumes all the oxygen.

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.^[72][73]

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.^[74][75] 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.^[76]

Direct skin contact with liquid nitrogen will cause severe frostbite (cryogenic "burns"). This may happen almost instantly on contact, or after a second or more, depending on the form of liquid nitrogen. Bulk liquid nitrogen causes less rapid freezing than a spray of nitrogen mist (such as is used to freeze certain skin growths in the practice of dermatology). The extra surface area provided by nitrogen-soaked materials is also important, with soaked clothing or cotton causing far more rapid damage than a spill of direct liquid to skin. Full "contact" between naked skin and large collected-droplets or pools of liquid nitrogen may be prevented for a second or two, by a layer of insulating gas from the Leidenfrost effect. This may give the skin a second of protection from nitrogen bulk liquid. However, liquid nitrogen applied to skin in mists, and on fabrics, bypasses this effect, and causes local frostbite immediately.

Oxygen sensors are sometimes used as a safety precaution when working with liquid nitrogen to alert workers of gas spills into a confined space.^[77] Nitrogen that evaporates from liquid is still colder and heavier than (oxygenated) air, therefore any spill of liquid nitrogen tends to stay at the bottom of a closed space, expelling oxygen.

See also

• Reactive nitrogen species
• Yeast assimilable nitrogen
• Nitrogen fixation
• Nitrogen deficiency in plants
• nitrogen execution

Notes

1. Nitrogen and its compounds are far more common in atmospheres of smaller rocky moons and planets than neon, due to nitrogen being less volatile than neon.

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. See:
   • Daniel Rutherford (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.
   • See also: James R. Marshall and Virginia L. Marshall (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.
4. 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.
5. 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.
6. Aaron J. Ihde, The Development of Modern Chemistry, New York 1964.
7. 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.)
8. 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.
9. Priestley, Joseph (1772). "Observations on different kinds of air". Philosophical Transactions of the Royal Society of London 62: 147–256. ; see: "VII. Of air infected with the fumes of burning charcoal." pp. 225–228.
10. Elements of Chemistry, trans. Robert Kerr (Edinburgh, 1790; New York: Dover, 1965), 52.
11. J. A. Chaptal with William Nicholson, trans., Elements of Chemistry, 3rd ed. (London, England: C.C. and J. Robinson, 1800), 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."
12. nitrogen. Etymonline.com. Retrieved 2011-10-26.
13. Lord Rayleigh's Active Nitrogen. Lateralscience.co.uk. Retrieved 2011-10-26.
14. 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.
15. Froehlich, Peter (May 2013). "A Sustainable Approach to the Supply of Nitrogen". www.parker.com. Parker Hannifin Corporation. Retrieved April 2016.
16. Reich, Murray.; Kapenekas, Harry. (1957). "Nitrogen Purfication. Pilot Plant Removal of Oxygen". Industrial & Engineering Chemistry 49 (5): 869–873. doi:10.1021/ie50569a032.
17. 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.
18. 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.
19. Lide, D. R., ed. (2003). CRC Handbook of Chemistry and Physics (84th ed.). Boca Raton, FL: CRC Press.
20. 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.
21. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
22. 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.
23. "A new molecule and a new signature – Chemistry – tetranitrogen". Science News. 16 February 2002. Retrieved 2007-08-18.
24. "Polymeric nitrogen synthesized". physorg.com. 5 August 2004. Retrieved 2009-06-22.
25. Fabian, J. & Lewars, E. (2004). "Azabenzenes (azines)—The nitrogen derivatives of benzene with one to six N atoms: Stability, homodesmotic stabilization energy, electron distribution, and magnetic ring current; a computational study" (PDF). Canadian Journal of Chemistry 82 (1): 50–69. doi:10.1139/v03-178.
26. Muir, B. Cubane. (See "further topics" section.)
27. Patil, Ujwala N.; Dhumal, Nilesh R. & Gejji, Shridhar P. (2004). "Theoretical studies on the molecular electron densities and electrostatic potentials in azacubanes". Theoretica Chimica Acta 112: 27–32. doi:10.1007/s00214-004-0551-2.
28. Bethe, H. A. (1939). "Energy Production in Stars". Physical Review 55 (5): 434–56. Bibcode:1939PhRv...55..434B. doi:10.1103/PhysRev.55.434.
29. 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.
30. 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 9780080525280.
31. "Atomic Weights and Isotopic Compositions for Nitrogen". NIST. Retrieved 2013-05-22.
32. 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.
33. Worley, R. (1943). "Absorption Spectrum of N2 in the Extreme Ultraviolet". Physical Review 64 (7–8): 207. Bibcode:1943PhRv...64..207W. doi:10.1103/PhysRev.64.207.
34. 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 2551323. PMID 16359167.
35. 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.
36. Emsley, p. 360
37. Croswell, Ken (February 1996). Alchemy of the Heavens. Anchor. ISBN 0-385-47214-5.
38. Meyer, Daved M.; Cardelli, Jason A.; Sofia, Ulysses J. (1997). "Abundance of Interstellar Nitrogen". The Astrophysical Journal 490: L103–L106. arXiv:astro-ph/9710162. Bibcode:1997ApJ...490L.103M. doi:10.1086/311023.
39. Hamilton, Calvin J. "Titan (Saturn VI)". Solarviews.com. Retrieved 2007-12-24.
40. Häring, Heinz-Wolfgang (2008). Industrial Gases Processing. John Wiley & Sons. pp. 243–. ISBN 978-3-527-62125-5. Retrieved 2012-01-02.
41. Emsley, p. 364
42. Ministers, Nordic Council of (2002). "Food Additives in Europe 2000": 591. ISBN 9789289308298.
43. Harding, Charlie, ed. (2002). Elements of the p Block. Cambridge: Royal Society of Chemistry. ISBN 978-0-85404-690-4.
44. Gavriliuk, V. G.; Berns, Hans (1999). High nitrogen steels: structure, properties, manufacture, applications. Springer. ISBN 3-540-66411-4.
45. "Centre Fuel Tank Inerting". B737.org.uk. Retrieved 2013-08-21.
46. "Why don't they use normal air in race car tires?". Howstuffworks. Retrieved 2006-07-22.
47. 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.
48. Baxter, E. Denise; Hughes, Paul S. (2001). Beer: Quality, Safety and Nutritional Aspects. Royal Society of Chemistry. p. 22. ISBN 9780854045884.
49. "How does the widget in a beer can work?". Howstuffworks.
50. Denny, Mark (1 November 2009). Froth!: The Science of Beer. p. 131. ISBN 9780801895692.
51. Kennett, Andrew J. "Design of a pneumatically assisted shifting system for Formula SAE® racing applications". hdl:1721.1/45820.
52. Sanburn, Josh (2015-04-10). "The Dawn of a New Form of Capital Punishment". Time. Retrieved 2015-04-11.
53. Sexton, Mike (18 December 2012). "Euthanasia campaigner under scrutiny". ABC. Retrieved 6 May 2013.
54. 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.
55. 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.
56. Kent, Allen; Williams, James G. (1994). Encyclopedia of Computer Science and Technology 30. CRC Press. p. 318. ISBN 0-8247-2283-3.
57. "Flowing nitrogen ice glaciers seen on surface of Pluto after New Horizons flyby". ABC. 25 July 2015. Retrieved 6 October 2015.
58. 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.
59. "Neptune: Moons: Triton". NASA. Archived from the original on October 5, 2011. Retrieved September 21, 2007.
60. Emsley, p. 362
61. Cleveland, David (2002) "Don't Try This at Home: Some Thoughts on Nitrate Film, With Particular Reference to Home Movie Systems" in Roger Smither and Catherine Surowiec (eds.), This Film is Dangerous: A Celebration of Nitrate Film, Brussels, FIAF, ISBN 978-2-9600296-0-4, p. 196
62. Rakov, Vladimir A.; Uman, Martin A. (2007). Lightning: Physics and Effects. Cambridge University Press. p. 508. ISBN 978-0-521-03541-5.
63. Jama, Bashir; Ndufa, J. K.; Buresh, R. J.; Shepherd, K. D. "Vertical Distribution of Roots and Soil Nitrate: Tree Species and Phosphorus Effects" 62 (1). Soil Science Society of America Journal: 280–286. Retrieved 2013-01-02.
64. McKane, Robert B.; Johnson, Loretta C.; Shaver, Gaius R.; Nadelhoffer, Knute J.; Rastetter, Edward B.; Fry, Brian; Giblin, Anne E.; Kielland, Knut; Kwiatkowski, Bonnie L. "Resource-based niches provide a basis for plant species diversity and dominance in arctic tundra". Nature 415 (6867): 68–71. doi:10.1038/415068a.
65. Bothe, Hermann; Ferguson, Stuart John; Newton, William Edward (2007). Biology of the nitrogen cycle. Elsevier. p. 283. ISBN 0-444-52857-1.
66. Jahn, G.C.; Almazan, Liberty P.; Pacia, Jocelyn B. (2005). "Effect of nitrogen fertilizer on the intrinsic rate of increase of the rusty plum aphid, Hysteroneura setariae (Thomas) (Homoptera: Aphididae) on rice (Oryza sativa L.)". Environmental Entomology 34 (4): 938–943. doi:10.1603/0046-225X-34.4.938.
67. Knox, G. A. (2007). Biology of the Southern Ocean. CRC Press. p. 392. ISBN 0-8493-3394-6.
68. 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.
69. 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.
70. "Biology Safety – Cryogenic materials. The risks posed by them.". University of Bath. Archived from the original on February 6, 2007. Retrieved 2007-01-03.
71. "Space Shuttle Columbia Fast Facts". CNN. September 30, 2013.
72. 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.
73. 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.
74. 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.
75. 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.
76. US Navy Diving Manual, 6th revision. United States: US Naval Sea Systems Command. 2006. Retrieved 2008-04-24.
77. Liquid Nitrogen – Code of practice for handling. United Kingdom: Birkbeck, University of London. 2007. Retrieved 2012-02-08.

Bibliography

• Emsley, John (2011). Nature's Building Blocks: An A-Z Guide to the Elements (New ed.). New York, NY: Oxford University Press. ISBN 978-0-19-960563-7.

Further reading

• Garrett, Reginald H.; Grisham, Charles M. (1999). Biochemistry (2nd ed.). Fort Worth: Saunders College Publ. ISBN 0-03-022318-0.

External links

• Etymology of Nitrogen
• Nitrogen at The Periodic Table of Videos (University of Nottingham)
• Nitrogen podcast from the Royal Society of Chemistry's Chemistry World