nitrogen

 

(nī'trəjən) , gaseous chemical element; symbol N; at. no. 7; at. wt. 14.0067; m.p. −209.86°C; b.p. −195.8°C; density 1.25 grams per liter at STP; valence principally −3, +3, or +5. Nitrogen is a colorless, odorless, tasteless diatomic gas. It is found in Group 15 of the periodic table. It does not burn, does not support combustion, and is only slightly soluble in water. It is relatively inactive chemically, but many of its compounds display marked reactivity. At high temperatures it reacts with some of the other elements to form nitrides.

Nitrogen has several oxides. Nitrous oxide, N₂O, is a gas used as an anesthetic; it is often called laughing gas. Nitric oxide, NO, is a gas used in the manufacture of sulfuric acid; in air it forms nitrogen dioxide, NO₂, a poisonous reddish brown gas. Nitrogen trioxide, N₂O₃, is unstable at ordinary temperatures. Nitrogen pentoxide, N₂O₅, forms nitric acid when dissolved in water. Important compounds of nitrogen include nitric acid, ammonia, many explosives, cyanides, fertilizers, and the proteins. Many organic compounds contain nitrogen.

Nitrogen for industrial use is produced largely by the fractional distillation of liquid air. Nitrogen is used to some extent for filling light bulbs, in thermometers, and generally anywhere a relatively inert atmosphere is needed, as in the production of electronic parts such as transistors, diodes, and integrated circuits, and in food storage packaging to prevent spoilage. It is used in the manufacture of stainless steel and as a coolant for the immersion freezing of food products, for the transportation of foods, for the preservation of bodies and reproductive cells (sperm and eggs), and for the storage of biological samples. However, the chief importance of the element lies in its compounds, among them ammonia, nitric acid, and cyanide.

The expression “nitrogen fixation” refers to the extraction of the element from the atmosphere by its combination with other elements to form compounds. This is accomplished commercially in several ways. In the Haber process, nitrogen is reacted with hydrogen to form ammonia; in the cyanamide process, nitrogen is reacted with calcium carbide at high temperatures to form calcium cyanamide; in the arc process, nitrogen is reacted with oxygen in an electric arc to form nitrogen oxides.

Nitrogen is abundant in the atmosphere; it is about 78% (by volume) of dry air. Nitrogen is present in living things; it and its compounds are necessary for the continuation of life (see nitrogen cycle). Nitrogen also is found in foods and is important in the human diet.

Nitrogen compounds were known to alchemists as early as the Middle Ages, but nitrogen is formally considered to have been discovered by Daniel Rutherford in 1772, who called it noxious air or phlogisticated air (air from which the oxygen had been removed, usually by combustion). Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as burnt air or dephlogisticated air. It was well known to these late 18th century chemists that there was a fraction of air that did not support combustion. Antoine Lavoisier was the first to treat oxygenless air as a separate element, which he called azote, meaning without life. The term nitrogen was first used by J. A. Chaptal in 1790. This early “nitrogen” was later shown by John Strutt (Lord Rayleigh), and William Ramsay to contain argon; Henry Cavendish had shown in 1785 that there was an unreactive gas other than nitrogen present in air.

7 carbon ← nitrogen → oxygen - ↑ N ↓ P Periodic table - Extended periodic table
General
Name, symbol, number	nitrogen, N, 7
Chemical series	nonmetals
Group, period, block	15, 2, p
Appearance	colorless gas
Standard atomic weight	14.0067(2) g·mol−1
Electron configuration	1s2 2s2 2p3
Electrons per shell	2, 5
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 1.251 g/L
Melting point	63.15 K (-210.00 °C, -346.00 °F)
Boiling point	77.36 K (-195.79 °C, -320.42 °F)
Critical point	126.21 K, 3.39 MPa
Heat of fusion	(N2) 0.720 kJ·mol−1
Heat of vaporization	(N2) 5.57 kJ·mol−1
Heat capacity	(25 °C) (N2) 29.124 J·mol−1·K−1
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K 37 41 46 53 62 77
Atomic properties
Crystal structure	hexagonal
Oxidation states	±3, 5, 4, 2 (strongly acidic oxide)
Electronegativity	3.04 (Pauling scale)
Ionization energies (more)	1st: 1402.3 kJ·mol−1
	2nd: 2856 kJ·mol−1
	3rd: 4578.1 kJ·mol−1
Atomic radius	65 pm
Atomic radius (calc.)	56 pm
Covalent radius	75 pm
Van der Waals radius	155 pm
Miscellaneous
Magnetic ordering	diamagnetic
Thermal conductivity	(300 K) 25.83 × 10−3 W·m−1·K−1
Speed of sound	(gas, 27 °C) 353 m/s
CAS registry number	7727-37-9
Selected isotopes
Main article: Isotopes of nitrogen iso NA half-life DM DE (MeV) DP 13N syn 9.965 min ε 2.220 13C 14N 99.634% N is stable with 7 neutrons 15N 0.366% N is stable with 8 neutrons
References

Nitrogen (IPA: /ˈnaɪtrədʒən/) is a chemical element which has the symbol N and atomic number 7. Elemental nitrogen is a colorless, odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78.1% by volume of Earth's atmosphere. Nitrogen is a constituent element of all living tissues and amino acids. Many industrially important compounds, such as ammonia, nitric acid, and cyanides, contain nitrogen.

Properties

Nitrogen is a nonmetal, with an electronegativity of 3.0. It has five electrons in its outer shell and is therefore trivalent in most compounds. The triple bond in molecular nitrogen (N₂) is one of the strongest in nature. The resulting difficulty of converting (N₂) into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N₂, have dominated the role of nitrogen in both nature and human economic activities.

At atmospheric pressure molecular nitrogen condenses (liquifies) at 77 K (−195.8 °C) and freezes at 63 K (−210.0 °C) into the beta hexagonal close-packed crystal allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the alpha cubic crystal allotropic form. Liquid nitrogen, a fluid resembling water, but with 80.8% of the density, is a common cryogen.

Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like N₃ and N₄.^[1] Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced under diamond anvil conditions, nitrogen polymerizes into the single bonded diamond crystal structure, an allotrope nicknamed "nitrogen diamond."^[2]

Occurrence

Nitrogen is the largest single component of the Earth's atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air).

¹⁴N is created as part of the fusion processes in stars, and is estimated to be the 7th most abundant chemical element (by mass) in our universe.

Compounds that contain this element have been observed by astronomers, and molecular nitrogen has been detected in interstellar space by David Knauth and coworkers using the Far Ultraviolet Spectroscopic Explorer. Molecular nitrogen is a major constituent of Titan's thick atmosphere, and occurs in trace amounts of other planetary atmospheres.

Nitrogen is present in all living tissues as proteins, nucleic acids and other molecules. It is a large component of animal waste (for example, guano), usually in the form of urea, uric acid, and compounds of these nitrogenous products.

See also: Nitrate minerals and Ammonium minerals

Isotopes

See also: Isotopes of nitrogen

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 and the remaining is ¹⁵N. 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. 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.

0.73% of the molecular nitrogen in Earth's atmosphere is comprised of the isotopologue ¹⁴N¹⁵N and almost all the rest is ¹⁴N₂.

Electromagnetic spectrum

Molecular nitrogen (¹⁴N₂) is largely transparent to infrared and visible radiation because it is a homonuclear molecule and thus has no dipole moment to couple to electromagnetic radiation at these wavelengths. Significant absorption occurs at extreme ultraviolet wavelengths, beginning around 100 nanometers. 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 as well as in 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).

History

Nitrogen (Latin nitrogenium, where nitrum (from Greek nitron) means "native soda" (see niter), and genes means "forming") is formally considered to have been discovered by Daniel Rutherford in 1772, who called it noxious air or fixed air. That there was a fraction of air that did not support combustion was well known to the late 18th century chemist. Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as azote, from the Greek word αζωτος meaning "lifeless". Animals died in it, and it was the principal component of air in which animals had suffocated and flames had burned to extinction. This term has become the French word for "nitrogen" and later spread out to many other languages.

Argon was discovered when it was noticed that nitrogen from air is not identical to nitrogen from chemical reactions.

Compounds of nitrogen were known in the Middle Ages. The 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 industrial and agricultural applications of nitrogen compounds involved uses in the form of saltpeter (sodium- or potassium nitrate), notably in gunpowder, and much later, as fertilizer,

Applications

Nitrogen gas is acquired for industrial purposes by the fractional distillation of liquid air, or by mechanical means using gaseous air (i.e. pressurised reverse osmosis membrane or pressure swing adsorption). Commercial nitrogen is often a byproduct of air-processing for industrial concentration of oxygen for steelmaking and other purposes.

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

• To preserve the freshness of packaged or bulk foods (by delaying rancidity and other forms of oxidative damage)
• In ordinary incandescent light bulbs as an inexpensive alternative to argon
• On top of liquid explosives for safety
• The production of electronic parts such as transistors, diodes, and integrated circuits
• Dried and pressurized, as a dielectric gas for high voltage equipment
• The manufacturing of stainless steel
• Use in military aircraft fuel systems to reduce fire hazard, see inerting system
• Filling automotive and aircraft tires^[4] due to its inertness and lack of moisture or oxidative qualities, as opposed to air, though this is not necessary for consumer automobiles.^[5][6]

Nitrogen molecules are less likely to escape from the inside of a tire compared with the traditional air mixture used. Air consists mostly of nitrogen and oxygen. Nitrogen molecules have a larger effective diameter than oxygen molecules and therefore diffuse through porous substances more slowly.^[7]

A further example of its versatility is its use as a preferred alternative to carbon dioxide to pressurize kegs of some beers, particularly thicker stouts and Scottish and English ales, due to the smaller bubbles it produces, which make the dispensed beer smoother and headier. A modern application of a pressure sensitive nitrogen capsule known commonly as a "widget" now allows nitrogen charged beers to be packaged in cans and bottles.

Molecular nitrogen, a diatomic gas, is apt to dimerize into a linear four nitrogen long polymer. This is an important phenomenon for understanding high voltage nitrogen dielectric switches because the process of polymerization can continue to lengthen the molecule to still longer lengths in the presence of an intense electric field. A nitrogen polymer fog is thereby created. The second virial coefficient of nitrogen also shows this effect as the compressibility of nitrogen gas is changed by the dimerization process at moderate and low temperatures.^{[citation needed]}

Nitrogen tanks are also replacing carbon dioxide as the main power source for paintball markers. The downside is that nitrogen must be kept at higher pressure than CO2, making N2 tanks heavier and more expensive.

Liquid nitrogen

Main article: Liquid nitrogen

Liquid nitrogen is a cryogenic liquid. At atmospheric pressure, it boils at −196 °C. When insulated in proper containers such as dewar flasks, it can be transported without much evaporative losses.

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 also used in cold traps for certain laboratory equipment. It has also been used to cool central processing units and other devices in computers which are overclocked, and which produce more heat than during normal operation.

Biological role

See also: nitrogen cycle

Nitrogen is an essential part of amino acids and nucleic acids, both of which are essential to all life on Earth.

Molecular nitrogen in the atmosphere cannot be used directly by either plants or animals, and needs to be converted to other compounds, or "fixed," in order to be used by life. Precipitation often contains substantial quantities of ammonium and nitrate, both thought to be a result of nitrogen fixation by lightning and other atmospheric electric phenomena. However, because ammonium is preferentially retained by the forest canopy relative to atmospheric nitrate, most of the fixed nitrogen that reaches the soil surface under trees is in the form of nitrate. Soil nitrate is preferentially assimilated by tree roots relative to soil ammonium.

Specific bacteria (e.g. Rhizobium trifolium) possess nitrogenase enzymes which can fix atmospheric nitrogen (see nitrogen fixation) into a form (ammonium ion) which is chemically useful to higher organisms. This process requires a large amount of energy and anoxic conditions. Such bacteria may be free in the soil (e.g. azotobacter) but normally exist in a symbiotic relationship in the root nodules of leguminous plants (e.g. clover or the soya bean plant). Nitrogen fixating bacteria can be symbiotic with a number of unrelated plant species. Common examples are legumes, alders, lichens, casuarina, myrica, liverwort, and gunnera.

As part of the symbiotic relationship, the plant subsequently converts the ammonium ion to nitrogen oxides and amino acids to form proteins and other biologically useful molecules, such as alkaloids. In return for the usable (fixed) nitrogen, the plant secretes sugars to the symbiotic bacteria.

Some plants are able to assimilate nitrogen directly in the form of nitrates which 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.

Nitrogen compounds are basic building blocks in animal biology. 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. Some plant-feeding insects are so dependent on nitrogen in their diet, that varying the amount of nitrogen fertilizer applied to a plant can affect the birth rate of the insects feeding on it.^[8]

Soluble nitrate is an important limiting factor in the growth of certain bacteria in ocean waters. 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: PMID 15186102). In animals, the free radical molecule nitric oxide (NO), which is derived from an amino acid, serves as an important regulatory molecule for circulation.

Animal metabolism of NO results in production of nitrite. Animal metabolism of nitrogen in proteins generally 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 nitrogen-containing long-chain amines, such as putrescine and cadaverine.

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.

Reactions

Nitrogen is generally considered unreactive. N₂ reacts spontaneously with few reagents, being resilient to acids and bases as well as oxidants and most reductants. It does however react with lithium metal. Lithium burns in an atmosphere of N₂ to give lithium nitride:

6 Li + N₂ → 2 Li₃N

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). Such compounds are now numerous, other examples include IrCl(N₂)(PPh₃)₂, W(N₂)₂(Ph₂CH₂CH₂PPh₂)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-Bosch Process.^[9]

Nitrogen compounds in industry

Simple compounds

See also the category Nitrogen compounds.

The main neutral hydride of nitrogen is ammonia (NH₃), although hydrazine (N₂H₄) is also commonly used. Ammonia is more basic than water by 6 orders of magnitude. In solution ammonia forms the ammonium ion (NH₄⁺). Liquid ammonia (b.p. 240 K) is amphiprotic (displaying either Brønsted-Lowry acidic or basic character) and forms ammonium and the less common amide ions (NH₂^-); both amides and nitride (N^3-) salts are known, but decompose in water. Singly, doubly, triply and quadruply substituted alkyl compounds of ammonia are called amines (four substitutions, to form commercially and biologically important quarternary 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 (N₃^-), 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 dinitrogen monoxide (N₂O), also known as laughing gas. This is one of a variety of oxides, the most prominent of which are nitrogen monoxide (NO) (known more commonly as nitric oxide in biology), a natural free radical molecule used by the body as a signal for short-term control of smooth muscle in the circulation. Another notable nitrogen oxide compound (a family often abbreviated NOx) is the reddish and poisonous nitrogen dioxide (NO₂), which also contains an unpaired electron and is an important component of smog. Nitrogen molecules containing unpaired electrons show an understandable tendency to dimerize (thus pairing the electrons), and are generally highly reactive.

The more standard oxides, dinitrogen trioxide (N₂O₃) and dinitrogen pentoxide (N₂O₅), are actually fairly unstable and explosive-- a tendency which is driven by the stability of N₂ as a product. The corresponding acids are nitrous (HNO₂) and nitric acid (HNO₃), with the corresponding salts called nitrites and nitrates. Nitric acid is one of the few acids stronger than hydronium, and is a fairly strong oxidizing agent.

Nitrogen can also be found in 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.

Nitrogen compounds of notable economic importance

Molecular nitrogen (N₂) in the atmosphere is relatively non-reactive due to its strong bond, and N₂ plays an inert role in the human body, being neither produced or destroyed. In nature, nitrogen is converted into biologically (and industrially) useful compounds by some living organisms, notably certain bacteria (i.e. nitrogen fixing bacteria – see Biological role above). Molecular nitrogen is also 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, where nitrogen and natural gas are converted into ammonia via the Haber process. Ammonia, in turn, can be used directly (primarily as a fertilizer, and in the synthesis of nitrated fertilizers), or as a precursor of many other important materials including explosives, largely via the production of nitric acid by the Ostwald process.

The organic and inorganic salts of nitric acid have been historically important as stores of chemical energy. They include important compounds such as potassium nitrate (or saltpeter, important historically for its use in gunpowder) and ammonium nitrate, an important fertilizer and explosive (see ANFO). Various other nitrated organic compounds, such as nitroglycerin and trinitrotoluene, and nitrocellulose, 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. 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 which results, 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 natural chemical defences of plants against predation). Nitrogen containing drugs include all of the major classes of antibiotics, and organic nitrate drugs like nitroglycerin and nitroprusside which regulate blood pressure and heart action by mimicking the action of nitric oxide.

Dangers

Rapid release of nitrogen gas into an enclosed space can displace oxygen, and therefore represents an asphyxiation hazard. This may happen with few warning symptoms, since the human carotid body is a relatively slow and poor low-oxygen (hypoxia) sensing system.^[10] An example occurred shortly before the launch of the first Space Shuttle mission in 1981, when two technicians lost consciousness and died 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. The technicians would have been able to exit the room if they had experienced early symptoms from nitrogen-breathing.

When breathed at high