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nitrogen (nīˈtrəjən), gaseous chemical element; symbol N; at. no. 7; interval in which at. wt. ranges 14.00643–14.00728; m.p. −209.86℃; b.p. −195.8℃; 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.
(in Russian, azot), N; a chemical element in Group V of Mendeleev’s periodic table, with atomic number 7 and atomic weight 14.0067; a colorless, odorless, and tasteless ga
History Nitrogen compounds—saltpeter, nitric acid, ammonia—were known long before nitrogen was prepared in a free state. In 1772, Rutherford burned phosphorus and other substances in a glass bell jar and showed that the gas remaining after combustion, which he called “suffocating air,” does not support respiration or combustion. In 1787, A. Lavoisier established that the “vital” and “suffocating” gases that are constituents of the air are elements and proposed the name “azote” for the “suffocating” gas. In 1784, H. Cavendish demonstrated that nitrogen is a component of saltpeter; this is the origin of the name “nitrogen” (from Late Latin nitrum, “niter” or “saltpeter,” and Greek gennao, “I give birth, I produce”), which was proposed in 1790 by J. A. Chaptal. By the beginning of the 19th century, the chemical inertness of free nitrogen and its extraordinary role in the form of fixed nitrogen in compounds with other elements had become known. From that time on, the “fixation” of atmospheric nitrogen has become one of the most important technological problems in chemistry.
Natural occurrence Nitrogen is one of the most abundant elements on earth; its basic mass (about 4 ξ 1015 tons) is found in the free state in the atmosphere. In the air free nitrogen (in the form of N2 molecules) constitutes 78.09 percent by volume (or 75.6 percent by weight), ignoring negligible admixtures in the form of ammonia and oxides. The average nitrogen content in the lithosphere is 1.9 x 10−3 percent by weight. Naturally occurring nitrogen compounds are ammonium chloride (NH4C1) and various nitrates. Large nitrate deposits are characteristic of dry desert climates (Chile, Middle Asia). Nitrates were long the main source of nitrogen for industrial use (at present, the commercial synthesis of ammonia from hydrogen and atmospheric nitrogen is of primary importance in nitrogen fixation). Small amounts of bound nitrogen occur in coal (1–2.5 percent) and petroleum (0.02–1.5 percent), as well as in river, sea, and ocean water. Nitrogen accumulates in well as in river, sea and ocean water. Nitrogen accumulates in soils (0.1 percent) and organisms (0.3 percent).
Although the name azote means “not supporting life,” nitrogen is actually an element essential for life processes. Animal and human protein contains 16–17 percent nitrogen. In the bodies of carnivores, protein is formed from the proteins in the bodies of herbivores and plants they consume. Plants synthesize protein by assimilating nitrogenous substances, primarily inorganic ones, occurring in the soil. Considerable amounts of nitrogen enter the soil by means of nitrogen-fixing microorganisms capable of converting free atmospheric nitrogen into nitrogen compounds.
In nature there is a nitrogen cycle in which nitrifying, denitrifying, nitrogen-fixing, and other microorganisms play the main role. However, as a result of the extraction of an enormous amount of bound nitrogen from the soil by plants (especially in cases of intensive cultivation), soils show nitrogen depletion. A nitrogen deficiency characterizes agriculture in almost all countries and is observed in animal usbandry as well (“protein starvation”). Plants develop poorly in soils deficient in accessible nitrogen. Nitrogen fertilizers and supplementary protein feeding of animals are highly important means of raising agricultural production. Man’s economic activity disturbs the nitrogen cycle. Thus, the combustion of fuel enriches the atmosphere in nitrogen, but fertilizer factories bind atmostpheric nitrogen. The shipment of fertilizers and agricultural products redistributes the nitrogen on the surface of the earth.
Nitrogen is the fourth most abundant element in the solar system after hydrogen, helium, and oxygen.
Isotopes, atom, and molecule Naturally occurring nitrogen consists of two stable isotopes, 14N (99.635 percent) and 15 N (0.365 percent). The 15 N isotope is used in chemical and biochemical research as a tracer atom. Among the artificial radioactive isotopes of nitrogen, 13N has the longest half-life (T½ = 10.08 min.), whereas the others are very short-lived. In the upper layers of the atmosphere, 14N is converted by the action of cosmic-ray neutrons into the radioactive carbon isotope 14C. This process is used in nuclear reactions to produce 14C. The outer electron shell of the nitrogen atom consists of five electrons (one unshared electron pair and three unpaired electrons, to give the configuration 2s22p3). Most commonly, nitrogen has a covalence of 3 in its compounds by virtue of its unpaired electrons (as in ammonia NH3). The existence of the unshared electron pair can lead to the formation of one more covalent bond, and nitrogen then assumes a covalence of 4 (as in the ammonium ion NH4+). The oxidation states of nitrogen vary from +5 (in N205) to -3 (in NH3). Under ordinary conditions, free nitrogen forms the molecule N2, in which the N atoms are connected by three covalent bonds. The nitrogen molecule is very stable: its energy of dissociation into atoms is 942.9 kilojoules per mole (225.2 kilocalories per mole), and therefore even at approximately 3300°C the degree of dissociation of nitrogen is only about 0.1 percent.
Physical and chemical properties Nitrogen is somewhat lighter than air; its density at 0°C and 101,325 newtons per m2 or 760 mm Hg is 1.2506 kg/m3. It melts at −209.86°C and boils at −195.8°C. Nitrogen is difficult to liquefy; its critical temperature is quite low (−147.1°C) and its critical pressure is high (3.39 meganewtons per m2 [MN/m2] or 34.6 kilogram force per cm2). The density of liquid nitrogen is 808 kg/m3. Nitrogen is less soluble in water than oxygen—at 0C C, 23.3 g of nitrogen dissolve in 1 m3 of H20. Nitrogen is more soluble in some hydrocarbons than in water.
Only with such active metals as lithium, calcium, and magnesium does nitrogen react during moderate heating. With most other elements nitrogen reacts at high temperature and in the presence of catalysts. The compounds of nitrogen and oxygen N20, NO, N203, N02, and N205 have been thoroughly studied. Of these, the oxide NO forms upon direct interaction of the elements at 4000°C; upon cooling, it readily oxidizes further to the dioxide NO2. In the air, oxides of nitrogen are formed during atmospheric discharges. They can also be produced by the action of ionizing radiation on a mixture of nitrogen with oxygen. When nitrous anhydride (N203) and nitric anhydride (N2O5) dissolve in water, nitrous acid (HN02) and nitric acid (HN03) respectively are formed; they form salts called nitrites and nitrates. With hydrogen, nitrogen combines only at high temperature and in the presence of catalysts to give ammonia (NH3). In addition to ammonia, numerous other compounds of nitrogen with hydrogen are known—for example, hydrazine (H2N—NH2), diimide (HN ═ NH), hy-drazoic acid [HN3 (—N═N≡N)], and octazone (N8H14); most of the compounds of nitrogen with hydrogen have been isolated only in the form of organic derivatives. Nitrogen does not react directly with halogens, and therefore all nitrogen halides are prepared only indirectly; for example, nitrogen trifluoride (NF3) is prepared by the interaction of fluorine and ammonia. As a rule, the nitrogen halides are unstable compounds (except for NF3); the nitrogen oxy-halides (NOF, NOCI, NOBr, N02F, and N02C1) are more stable. Nitrogen also does not react directly with sulfur; sulfur nitride (N4S4) is formed by the reaction of liquid sulfur with ammonia. Cyanogen (CN)2 is formed by the reaction of incandescent coke with nitrogen. Hydrogen cyanide (HCN) can be prepared by heating nitrogen with acetylene (C2H2) to 1500°C. The interaction of nitrogen with metals at high temperatures leads to the formation of nitrides, such as Mg3N2.
During the action of electric discharges on ordinary nitrogen (at a pressure of 130–270 N/m2 or 1–2 mm Hg) or during the decomposition of the nitrides of boron, titanium, magnesium, and calcium, as well as during electric discharges in the air, active nitrogen may form; this is a mixture of molecules and atoms of nitrogen that have an increased energy content. In contrast with molecular nitrogen, active nitrogen reacts very vigorously with oxygen, hydrogen, sulfur vapors, phosphorus, and some metals.
Nitrogen is a constituent of a great many very important organic compounds, including amines, amino acids, and nitro compounds.
Preparation and use Nitrogen can be easily prepared in the laboratory by heating a concentrated solution of ammonium nitrite: NH4N02 = N2 + 2H20. The commercial method of preparing nitrogen is based on the separation of preliquefied air, which is then distilled.
The bulk of the free nitrogen produced is used in the commercial production of ammonia, considerable amounts of which are then converted into nitric acid, fertilizers, explosives, and so on. In addition to the direct synthesis of ammonia from the elements, the cyanamide method, developed in 1905, is of commercial importance in the fixation of atmospheric nitrogen. Calcium carbide (prepared by heating a mixture of lime and coal in an electric furnace) reacts at 1000°C with free nitrogen: CaC2 + N2 = CaCN 2+ C. The calcium cynamide formed is decomposed by superheated steam with the release of ammonia CaCN2 + 3H20 = CaCO3 + 2NH3
Free nitrogen is used in many branches of industry—as an inert medium in various chemical and metallurgical processes, for filling the empty space in mercury thermometers, in the pumping of flammable liquids, and so on. Liquid nitrogen is used in various cooling devices. It is stored and shipped in steel Dewar flasks, while cylinders are used for compressed gaseous nitrogen. Many nitrogen compounds are widely used. The production of fixed nitrogen began to develop rapidly after World War I and has now reached enormous proportions.
REFERENCESNekrasov, B. V. Osnovy obshchei khimii, vol. 1. Moscow, 1965.
Remy, G. Kurs neorganicheskoi khimii, vol. 1. Moscow, 1963.
(Translated from German.)
Khimiia i tekhnologiia sviazannogo azota. [Moscow-Leningrad,]1934.
Kratkaia khimicheskaia entsiklopediia, vol. 1. Moscow, 1961.