|colorless gas, liquid or
|Name, symbol, number||nitrogen, N, 7|
|Group, period, block||15, 2, p|
|Standard atomic weight||14.0067(2) g·mol−1|
|Electron configuration||1s2 2s2 2p3|
|Electrons per shell||2, 5 (Image)|
|Density||(0 °C, 101.325 kPa)
|Melting point||63.153 K, -210.00 °C, -346.00 °F|
|Boiling point||77.36 K, -195.79 °C, -320.3342 °F|
|Triple point||63.1526 K (-210°C), 12.53 kPa|
|Critical point||126.19 K, 3.3978 MPa|
|Heat of fusion||(N2) 0.72 kJ·mol−1|
|Heat of vaporization||(N2) 5.56 kJ·mol−1|
|Specific heat capacity||(25 °C) (N2)
|Oxidation states||5, 4, 3, 2, 1, -1, -2, -3
(strongly acidic oxide)
|Electronegativity||3.04 (Pauling scale)|
|1st: 1402.3 kJ·mol−1|
|2nd: 2856 kJ·mol−1|
|3rd: 4578.1 kJ·mol−1|
|Covalent radius||71±1 pm|
|Van der Waals radius||155 pm|
|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|
|Most stable isotopes|
|Main article: Isotopes of nitrogen|
Nitrogen is a chemical element that has the symbol N, the atomic number of 7 and an atomic mass 14.00674 u. Elemental nitrogen is a colorless, odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78% by volume of Earth's atmosphere.
Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in converting the N2 into useful compounds, and releasing large amounts of energy when these compounds burn or decay back into nitrogen gas.
The element nitrogen was discovered by Daniel Rutherford, a Scottish physician, in 1772. Nitrogen occurs in all living organisms. It is a constituent element of amino acids and thus of proteins, and of nucleic acids (DNA and RNA). It resides in the chemical structure of almost all neurotransmitters, and is a defining component of alkaloids, biological molecules produced by many organisms.
Nitrogen (Latin nitrogenium, where nitrum (from Greek nitron νιτρον) means "saltpetre" (see nitre), and genes γενης 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 "mephitic air" or azote, from the Greek word άζωτος (azotos) 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. Lavoisier's name for nitrogen is used in many languages (French, Russian, etc.) and still remains in English in the common names of many compounds, such as hydrazine and compounds of the azide ion. 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 military, industrial and agricultural applications of nitrogen compounds involved uses of saltpeter (sodium nitrate or potassium nitrate), notably in gunpowder, and much later, as fertilizer.
Nitrogen is a nonmetal, with an electronegativity of 3.04. It has five electrons in its outer shell and is therefore trivalent in most compounds. The triple bond in molecular nitrogen (N2) is 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 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.
Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like N3 and N4. Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced using 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. N4 is nicknamed "nitrogen diamond."
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. 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 the isotopologue 14N15N and almost all the rest is 14N2.
Radioisotope 16N is the dominant radionuclide in the coolant of pressurized 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, but during its decay back to 16O produces high-energy gamma radiation (5 to 7 MeV). Because of this, the access to the primary coolant piping must be restricted during reactor power operation. 16N is one of the main means used to immediately detect even small leaks from the primary coolant to the secondary steam cycle.
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).
Nitrogen is generally 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.
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(Ph2CH2CH2PPh 2)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. 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. (see nitrogen fixation)
The starting point for industrial production of nitrogen compounds is the Haber process, in which nitrogen is fixed by reacting N2 and H2 over an iron(III) oxide (Fe3O4) 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 which contains Fe and Mo atoms, using energy derived from hydrolysis of adenosine triphosphate (ATP) into adenosine diphosphate and inorganic phosphate (−20.5 kJ/mol).
Nitrogen is the largest single constituent of the Earth's atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air). It is created by fusion processes in stars, and is estimated to be the 7th most abundant chemical element by mass in the universe.
Molecular nitrogen and nitrogen compounds have been detected in interstellar space by astronomers using the Far Ultraviolet Spectroscopic Explorer. Molecular nitrogen is a major constituent of the Saturnian moon Titan's thick atmosphere, and occurs in trace amounts in other planetary atmospheres.
Nitrogen is present in all 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 are unable to fix atmospheric nitrogen.
Nitrogen occurs naturally in a number of minerals, such as saltpetre (potassium nitrate), Chile saltpetre (sodium nitrate) and sal ammoniac (ammonium chloride). Most of these are relatively uncommon, partly because of the minerals' ready solubility in water. See also Nitrate minerals and Ammonium minerals.
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 (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, 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 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. N 2+2 is another polyatomic cation as in hydrazine.
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, N2O), 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, for example in vasodilation by causing the smooth muscle of blood vessels to relax. The reddish and poisonous nitrogen dioxide NO2 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 corresponding acids are nitrous HNO2 and nitric acid HNO3, with the corresponding salts called nitrites and nitrates.
The higher oxides dinitrogen trioxide N2O3, dinitrogen tetroxide N2O4 and dinitrogen pentoxide N2O5, are fairly unstable and explosive, a consequence of the chemical stability of N2. Nearly every hypergolic rocket engine uses N2O4 as the oxidizer; their fuels, 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 on contact 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. N2O4 is an intermediate in the manufacture of nitric acid HNO3, one of the few acids stronger than hydronium and a fairly strong oxidizing agent.
Nitrogen is notable for the range of explosively unstable compounds that it can produce. Nitrogen triiodide NI3 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 more powerful explosive trinitrotoluene (TNT) is the standard explosive against which the power of nuclear explosions are measured.
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 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). 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 referred to as OFN (oxygen-free nitrogen).
Nitrogen is commonly used during sample preparation procedures for chemical analysis. Specifically, it is used as a means of concentrating and reducing the volume of liquid samples. Directing a pressurized stream of nitrogen gas perpendicular to the surface of the liquid allows the solvent to evaporate while leaving the solute(s) and un-evaporated solvent behind.
Nitrogen tanks are also replacing carbon dioxide as the main power source for paintball guns. The downside is that nitrogen must be kept at higher pressure than CO2, making N2 tanks heavier and more expensive.
A further example of its versatility is its use as a preferred alternative to carbon dioxide to pressurize kegs of some beers, particularly stouts and British 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.
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 cold traps for certain laboratory equipment and to cool x-ray detectors. 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.
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 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, 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 important historically as convenient stores of chemical energy. They include important compounds such as potassium nitrate (or saltpeter used 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 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 natural chemical defenses 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.
Elemental nitrogen in the atmosphere cannot be used directly by either plants or animals, and must be converted to a reduced (or 'fixed') state in order 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. This was first proposed by Liebig in 1827 and later confirmed. 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 these 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) 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). Nitrogen-fixing bacteria are also symbiotic with a number of unrelated plant species such as alders (Alnus) spp., lichens (Casuarina), Myrica, liverworts, and Gunnera.
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.
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 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.
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 not fresh saltwater fish . In animals, free radical nitric oxide (NO) (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 the creation of long-chain, nitrogen-containing 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. The circulation of nitrogen from atmosphere to organic compounds and back is referred to as the nitrogen cycle.
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 a poor low-oxygen (hypoxia) sensing system. 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 inhaled at high partial pressures (more than about 4 bar, encountered at depths below about 30 m in scuba diving) nitrogen begins to act as an anesthetic agent. It can cause nitrogen narcosis, a temporary semi-anesthetized state of mental impairment similar to that caused by nitrous oxide.
Nitrogen also dissolves in the bloodstream and body fats. Rapid decompression (particularly in the case of divers ascending too quickly, or astronauts decompressing too quickly from cabin pressure to spacesuit pressure) can lead to a potentially fatal condition called decompression sickness (formerly known as caisson sickness or more commonly, the "bends"), when nitrogen bubbles form in the bloodstream, nerves, joints, and other sensitive or vital areas. Other "inert" gases (those gases other than carbon dioxide and oxygen) cause the same effects from bubbles composed of them, so replacement of nitrogen in breathing gases may prevent nitrogen narcosis, but does not prevent decompression sickness.
Direct skin contact with liquid nitrogen will eventually cause severe frostbite (cryogenic burns). This may happen almost instantly on contact, 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 droplets or pools of undisturbed liquid nitrogen may be prevented for a few seconds by a layer of insulating gas from the Leidenfrost effect. However, liquid nitrogen applied to skin in mists, and on fabrics, bypasses this effect.
NITROGEN [[[symbol]] N., atomic weight 14. 01, 0 =16]. A non-metallic chemical element, first isolated in 1772 by D. Rutherford, who showed that on removing oxygen from air a gas remained, which was incapable of supporting combustion or respiration. Nitrogen forms approximately 79% by volume (or 77% by weight) of the atmosphere; actual values are: % by volume-79.07 (Regnault), 79.20 (Dumas); % by weight76.87 (Regnault), 77.00 (Dumas), 77.002 (Lewy), 76.900 (Stas), 77.010 (Marignac). No absolutely accurate determinations appear to have been made recently. Free nitrogen is also found in some natural waters and has been recognized in certain nebulae. In the combined state nitrogen is fairly widely distributed, being found in nitre, Chile saltpetre, ammonium salts and in various animal and vegetable tissues and liquids. It is invariably present in soils, where compounds are formed by nitrifying bacteria.
Nitrogen may be obtained from the atmosphere by the removal of the oxygen with which it is there mixed. This may be effected by burning phosphorus in a confined volume of air, by the action of an alkaline solution of pyrogallol on air, by passing air over heated copper, or by the action of copper on air in the presence of ammoniacal solutions.
It is also prepared by heating ammonium nitrite (or a mixture of sodium nitrite and ammonium chloride): NH 4 NO 2 =2H20+N21 by heating a mixture of ammonium nitrate and chloride (the chlorine which is simultaneously produced being absorbed by milk of lime or by a solution of sodium hydroxide): 4NH4N03+2NH4C1=5N2 +C1 2 +12H 2 O; by heating ammonium dichromate (or a mixture of ammonium chloride and potassium dichromate): (NH4)2Cr207 = Cr203+4H20+ N2; by passing chlorine into a concentrated solution of ammonia (which should be present in considerable excess): 8NH3+3C12=6NH4C1-F-N2; by the action of hypochlorites or hypobromites on ammonia: 3NaOBr-+2NH 3 =3NaBr+3H 2 OH-N 2; and by the action of manganese dioxide on ammonium nitrate at 180-20o° C. It is also formed by the reduction of nitric and nitrous oxides with hydrogen in the presence of platinized asbestos at a red heat (G. v. Knorre and K. Arndt, Ber., 1899, 32, p. 2136); by the oxidation of hydroxylamine (ibid., 1900, 33, p. 30); and by the electrolysis of hydrazine and its salts (E. Ch. Szarvasy, Jour. Chem. Soc., 1900, 77, p. 603).
The chief importance of nitrogenous compounds depends upon their assimilation by living plants, which, in their development, absorb these compounds from the soil, wherein they are formed mainly by the action of nitrifying bacteria. Since these compounds are essential to plant life, it becomes necessary to replace the amount abstracted from the soil, and hence a demand for nitrogenous manures was created. This was met in a very large measure by deposits of natural nitre and the products of artificial nitrieres, whilst additional supplies are available in the ammoniacal liquors of the gas-manufacturer, &c. The possible failure of the nitre deposits led to attempts to convert atmospheric nitrogen into manures by processes permitting economic success. Combination can be made in five directions, viz. to form (1) oxides and nitric acids, (2) ammonia, (3) readily decomposable nitrides, (4) cyanides, (5) cyanamides. The first three will he treated here; for the others see Prussic Acid and Cyanamide.
The combination of nitrogen with oxygen was first effected by Cavendish in 1785, who employed a spark discharge. The process was developed by Madame Lefebre in 1859; by Meissner in 1863, who found that moist gases gave a better result; and by Prim in 1882, who sparked the gases under pressure; it was also used by Lord Rayleigh in his isolation of argon. It was not, however, a commercial success, and the same result attended Siemens and Halske's application of the silent discharge. More effective was the electric arc. In 1892 Sir W. Crookes showed that the arc brought about combination; and in 1897 Lord Rayleigh went into the process more fully. But the first careful working-out of the conditions was made in 1900 by A. McDougall and F. Howles, who, employing a high tension alternating arc, showed that the effectiveness depended upon the temperature. The commercial manufacture of nitric acid was attempted by C. S. Bradley and D. R. Lovejoy at Niagara Falls, who passed atmospheric air, or air enriched with oxygen, about a high tension arc made as long as possible; but the company (the Atmospheric Products Company) was a failure. Better results have attended the process of K. Birkeland and S. Eyde, which is being worked on a large scale at Notodden, Norway. The arc is produced by leading a current of about 5000 volts equatorially between the poles of an electromagnet; this produces what is practically a disk of flame, 62 ft. in diameter and having a temperature of about 3000°. The disk really consists of a series of successive arcs which increase in size until they burst. The first product of the reaction is nitric oxide, which on cooling with the residual gases produces nitrogen peroxide. The cooled gases are then led into towers where they meet a stream of water coming in the contrary direction. Nitric acid (up to 50%) is formed in the first tower, and weaker acids in the successive ones; the last tower contains milk of lime which combines with the gases to form calcium nitrite and nitrate (this product, being unsuitable as a manure, is decomposed with the acid, and the evolved gases sent back). It was found advantageous not to work for acid but for a basic calcium nitrate (normal calcium nitrate being very deliquescent); for this purpose the acid is treated with the requisite amount of milk of lime. In the process of the Badische Anilinand Soda-Fabrik, the arc, which is said to be 30 to 50 ft. long or more, is formed in a long tube, and the gases are sent round the arc by obliquely injecting them. A 30% acid is said to be formed. Moscicki and J. von Kowalski have patented a process wherein the arc is formed at two vertical concentric copper electrodes and rotated by an electromagnet; it is worked at Vevey, Switzerland. The Rankin process, of which very little is known, produces the arc with much lower current.
The conversion of nitrogen into ammonia by electricity has received much attention, but the commercial aspect appears to have been first worked out by de Hemptinne in 1900, who used both the spark and silent discharge on mixtures of hydrogen and nitrogen, and found that the pressure and temperature must be kept low and the spark gap narrow. J. Schlutius in 1903 employed Dowson gas as a source of hydrogen, and induced combination by means of platinum and the silent discharge. Several non-electrical processes have been devised. In 1862 Fleck passed a mixture of steam, nitrogen and carbon monoxide over red-hot lime, whilst in 1904 Woltereck induced combination by passing steam and air over red-hot iron oxide (peat is used in practice). In de Lambilly's process air and steam is led over white-hot coke, and carbon dioxide or monoxide removed from the escaping gases according as ammonium formate or carbonate is wanted. The residual gas is then passed through a tube containing porous materials, such as woodor bone-charcoal, platinized pumice or spongy platinum, then mixed with steam and again forced through the tube. The reactions are represented as (I) N2+3H2+2C0 -1-2H 2 0=2H CO 2 NH 4 (Ammonium formate). (2) N2+3 H 2 +2CO 2 +2H 2 0 =2 HO. CO 2 NH 4 (Ammonium carbonate). The best temperature for the first reaction is between 80°C. and 130°C. and for the second between 40° C. and 60° C. In another process, which originated with C. Kaiser (Abst. J.C.S., 1907, ii. p. 862), calcium is heated in a current of hydrogen, and nitrogen passed over the hydride so formed; this gives ammonia and calcium nitride, the latter of which gives up its nitrogen as ammonia and reforms the hydride when heated in a current of hydrogen.
The fixation of nitrogen as a nitride has not been attended with commercial success. H. Mehner patented heating the oxides of silicon, boron or magnesium with coal or coke in an electric furnace, and then passing in nitrogen, which forms, with the metal liberated by the action of the carbon, a readily decomposable nitride.
Nitrogen is a colourless, tasteless and odourless gas, which is only very slightly soluble in water. It is slightly lighter than air. Lord Rayleigh in 1894 found that the density of atmospheric nitrogen was about 2% higher than that of chemically prepared nitrogen, a discovery which led to the isolation of the rare gases of the atmosphere (see Argon). The values obtained are shown below.
097209 (see D. L. Chapman and L. Vodden, Jour. Chem. Soc., 1909, 95, p. 138). Chlorine azide, C1 N 31 was discovered by F. Raschig in 1908 (see Azoimide); the corresponding iodine compound had been obtained in 1900 by A. Hantzsch (Ber., 33, p. 522). For the so-called nitrogen iodide see Ammonia.
Nitrogen forms five oxides, viz. nitrous oxide, N 2 0, nitric oxide, NO, nitrogen trioxide, N203, nitrogen peroxide, N02, and nitrogen pentoxide, N205, whilst three oxyacids of nitrogen are known: hyponitrous acid, H2N202, nitrous acid, HN02, and nitric acid, HNO 3 (q.v.). The first four oxides are gases, the fifth is a solid. Nitrous oxide, N 2 O, isolated in 1776 by J. Priestley, who obtained it by reducing nitrogen peroxide with iron, may be prepared by heating ammonium nitrate at 170-260° C., or by reducing a mixture of nitric and sulphuric acid with zinc. It is a colourless gas, which is practically odourless, but possesses a sweetish taste. It is somewhat soluble in water. When liquefied it boils at -89.8° C., and by further cooling may be solidified, the solid melting at -102.3° C. (W. Ramsay, Chem. News, 1893, 67, p. 140). It does not burn, but supports the combustion of heated substances almost as well as oxygen. It is used as an anaesthetic, principally in dentistry, producing when inhaled a condition of hysterical excitement often accompanied by loud laughter, whence it is sometimes called "laughing gas." Nitric oxide, NO, first obtained by Van Helmont, is usually prepared by the action of dilute nitric acid (sp. gr. 1.2) on copper. This method does not give a pure gas, varying amounts of nitrous oxide and nitrogen being present (see Nitric Acid). In a purer condition it may be obtained by the action of sulphuric acid on a mixture of potassium nitrate and ferrous sulphate, or of hydrochloric acid on a mixture of potassium nitrate and ferric chloride. It is also formed by the action of concentrated sulphuric acid on sodium nitrite in the presence of mercury. It is a colourless gas which is only sparingly soluble in water. It may be liquefied, its critical temperature being -93, 5°, and the liquid boils at -153.6° C. It is not a supporter of combustion, unless the sustance introduced is at a sufficiently high temperature to decompose the gas, when combustion will continue at the expense of the liberated oxygen. If the gas be mixed with the vapour of carbon disulphide, the mixture burns with a vivid lavender-coloured flame Nitric oxide is soluble in solutions of ferrous salts, a dark brown solution being formed, which is readily decomposed by heat, with evolution of nitric oxide. It combines with oxygen to form nitrogen peroxide. Nascent hydrogen reduces it to hydroxylamine (q.v.), whilst solutions of hypochlorites oxidize it to nitric acid. In some instances it reacts as a reducing agent, e.g. silver oxide is reduced to metallic silver at 170° C., lead dioxide to the monoxide and manganese dioxide to sesquioxide.
Nitrogen trioxide, N203, was first mentioned by J. R. Glauber in 1648 as a product of the reaction between nitric acid and arsenious oxide. Sir W. Ramsay (Jour. Chem. Soc., 18 9 0, 5, p. 59 o), by distilling arsenious oxide with nitric acid and cooling the distillate, obtained a green liquid which consisted of nitrogen trioxide and peroxide in varying proportions, and concluded that the trioxide could not be obtained pure. He then tried the direct combination of nitric oxide with liquid nitrogen peroxide. A dark blue liquid is produced, and the first portions of gas boiling off from the mixture correspond fairly closely in composition with nitrogen trioxide. H. B. Baker (Jour. Chem. Soc., 1907, 91, p. 1862) obtained nitrogen trioxide in the gaseous form by volatilizing the liquid under special conditions. L. Francesconi and N. Sciacca (Gazz., 1904, 34 (i.), P. 447) have shown that liquid nitric oxide and oxygen, or gaseous nitric oxide a.nd liquid oxygen, mixed in all proportions and yielded nitrogen trioxide, whilst gaseous nitric oxide mixed with excess, of oxygen always gave the trioxide if the mixture was kept below -110° C. They also state that nitrogen trioxide is stable at ordinary pressure up to -21° C. N. M. v. Wittorf (Zeit. anorg. Chem., 1904, 41, p. 85) obtained blue crystals of the trioxide (melting at -103° C.) on saturating liquid nitrogen peroxide with nitric oxide and cooling the mixture. The liquid prepared by Baker is green in colour, and has a specific gravity III at ordinary temperature, but below -2° C. becomes of a deep indigo blue colour. It forms a mass of deep blue crystals at the temperature of liquid air. It is exceedingly soluble in concentrated sulphuric acid.
Nitrogen peroxide, NO 2 or N204, may be obtained by mixing oxygen with nitric oxide and passing the red gas so obtained through a freezing mixture. The production of this red gas when air is mixed with nitric oxide was mentioned by R. Boyle in 1671. Nitrogen peroxide is also prepared by heating lead nitrate and passing the products of decomposition through a tube surrounded by a freezing mixture, when the gas liquefies. At low temperatures it is a colourless crystalline solid which melts at -10.14° C. (W. Ramsay, Chem. News, 1900, 61, p. 91). As the temperature increases the liquid becomes yellowish, the colour deepening with rise of temperature until at +15° C. it has a deep orange tint. The liquid boils at about 22° C. This change of colour is accompanied by a change in the vapour density, and is explained by the fact that nitrogen peroxide consists of a mixture of a colourless compound N204, and a redbrown gas N02, the latter increasing in amount at the expense of the former as the temperature is raised (G. Salet, Comptes rendus, 1868, 67, p. 488; see also E. and L. Natanson, Wied. Ann., 1885, 24, Chemical Nitrogen.
o 96727 Lord Rayleigh, Chem. News, 1897, 76, p. 315.0. 9720 0.9671 A. Leduc, Comptes rendus, 1896, 123, p. 805.
Nitrogen is a very inert gas: it will neither burn nor support the combustion of ordinary combustibles. It combines directly with lithium, calcium and magnesium when heated, whilst nitrides of the rare earth metals are also produced when their oxides are mixed with magnesium and heated in a current of nitrogen (C. Matignon, Comptes rendus, 1900, 131, p. 837). Nitrogen has been liquefied, the critical temperature being -149° C. and the critical pressure 27.54 atmospheres. The liquefied gas boils at -195.5° C., and its specific gravity at its boiling point is 0 8103 (E. C. C. Baly and F. G. Donnan, Jour. Chem. Soc., 1902, 81, p. 912).
Compounds. Nitrogen combines with hydrogen to form ammonia, NH 3, hydrazine, N 2 H 4, and azoimide, N 3 H (qq.v.); the other known hydrides, N 4 H 4 and N5H5, are salts of azoimide, viz. NH 4 N3 and N2H4 N3H.
Nitrogen trichloride, NC1 3, discovered by P. L. Dulong in 1811 (Schweigg. Journ., 1811, 8, p. 302), and obtained by the action of chlorine or sodium hypochlorite on ammonium chloride, or by the electrolysis of ammonium chloride solution, is a very volatile yellow oil. It possesses an extremely pungent smell, and its vapour is extremely irritating to the eyes. It is a most dangerous explosive P. 454; 1886, 27, p. 606). M. Berthelot and J. Ogier (Bull. Soc. Chem., 1882 , 37, P. 434; 38, p. 60) have also shown that the specific heat of the gas decreases with increase of temperature until it reaches a minimum at about 198-253° C. Cryoscopic determinations of the molecular weight of nitrogen peroxide dissolved in glacial acetic acid show that it corresponds to the molecular formula N204 at low temperatures (W. Ramsay, Jour. Chem. Soc., 1888, 53, p. 621). Nitrogen peroxide is the most stable oxide of nitrogen. It is decomposed by water, giving at o° C. a mixture of nitric and nitrous acids: 2N02+H20=HN03+HN02. It combines with sulphuric acid to form nitro-sulphonic acid, SO 2 (OH) (N02). It does not support the combustion of a taper, but burning phosphorus and red-hot carbon will continue to burn in the gas. It converts many metallic oxides into mixtures of nitrates and nitrites, and attacks many metals, forming nitrates and being itself reduced to nitric oxide. It is an energetic oxidizing agent.
Nitrogen pentoxide, N 2 O 5, was first obtained in 1849 by H. SainteClaire-Deville (Ann. Chim. Phys., 1850 , 28, p. 241) by the action of dry chlorine on silver nitrate: 4AgN03+2C12=4AgC1+2N205 +02. It may also be obtained by distilling nitric acid over phosphorus pentoxide. It crystallizes in large prisms which melt at 29-30° C. to a yellowish liquid, which boils at 45-50° C. with rapid decomposition. It is very unstable, decomposing slowly even at ordinary temperatures. It dissolves in water, forming nitric acid.
Hyponitrous acid, H2N202, was first obtained in the form of its salts by E. Divers in 1871 (Chem. News, 23, p. 206) by reducing a solution of potassium nitrite with sodium amalgam, and subsequent precipitation as silver salt. Hyponitrites also result when hydroxyamido-sulphonates, e.g. HO NH SO 3 Na, are hydrolysed by caustic alkalis (E. Divers and T. Haga, Jour. Chem. Soc., 188 9, 55, p. 760), or when benzsulphohydroxamic acid, C 6 H 5 SO 2 NH OH, is treated in the same manner (0. Piloty, Ber., 1896, 29, p. 1560). They may also be prepared by the action of mercuric or cupric oxides on alkaline solutions of hydroxylamine (A. Hantzsch, Ann., 1896, 2 9 2, p. 3 1 7); by the action of hydroxylamine sulphate on alkaline nitrites in the presence of lime or calcium carbonate, the mixture being rapidly heated to 60° C.; or by the hydrolysis of dimethyl nitroso-oxyurea, (CH 3) 2 N CO N(NO) OH (A. Hantzsch, Ber., 18 97, 3 0, p. 2356). The free acid, which crystallizes in brilliant scales, is best prepared by decomposing the silver salt with an ethereal solution of hydrochloric acid. It is very explosive, dissolves readily in water and behaves as a dibasic acid. It does not liberate iodine from potassium iodide, neither does it decolorize iodine solution. Bromine oxidizes it to nitric acid, but the reaction is not quantitative. In acid solution, potassium permanganate oxidizes it to nitric acid, but in alkaline solution only to nitrous acid. It decomposes slowly on standing, yielding water and nitrous oxide. The silver salt is a bright yellow solid, soluble in dilute sulphuric and nitric acids, and may be crystallized from concentrated solutions of ammonia. It slowly decomposes on exposure or on heating. The calcium salt, CaN 2 O 2.4H 2 O, formed by the action of calcium chloride on the silver salt in the presence of a small quantity of nitric acid, is a lustrous crystalline powder, almost insoluble in water but readily soluble in dilute acids. It is decomposed by sulphuric acid, with evolution of nitrous oxide.
Nitrous acid, HN02, is found to some extent in the form of its salts in the atmosphere and in rain water. The pure acid has not yet been obtained, since in the presence of water it decomposes with formation of nitric acid and liberation of nitric oxide: 3HN02 =HNO 3 +2N0+H20. Its salts may be obtained in some cases by heating the corresponding nitrates, but the method does not give good results. Sodium nitrite, the most commonly used salt of the acid, is generally obtained by heating the nitrate with metallic lead; by heating sodium nitrate with sulphur and sodium hydroxide, the product then being fractionally crystallized;(Read, Holliday & Sons): 3NaNO 3 +S+2NaOH = Na2S04+3NaN02+H20; by oxidizing atmospheric nitrogen in an electric arc, keeping the gases above 300° C., until absorption in alkaline hydroxide solution is effected (German Pat. 188188); or by passing air, or a mixture of oxygen and ammonia, over heated metallic oxides (ibid., 168272). The salts of the acid are colourless or faintly yellow. In aqueous solution the free acid acts as an oxidizing agent, bleaching indigo and liberating iodine from potassium iodide, or it may act as a reducing agent since it readily tends to pass into nitric acid: consequently it discharges the colour of acid solutions of permanganates and chromates. The acid finds considerable use in organic chemistry, being employed to discriminate between the different types of alcohols and of amines, and also in the production of diazo, azo and diazo-amino compounds. It may be recognized by the blue colour it gives with diphenylamine sulphate and by its reaction with potassium iodide-starch paper.
Nitrosyl chloride, NOC1, is obtained by the direct union of nitric oxide with chlorine; or by distilling a mixture of concentrated nitric and hydrochloric acids, passing the resulting gases into concentrated sulphuric acid and heating the so-formed nitrosyl hydrogen sulphate with dry salt: HN03+3HCl=NOC1+C12 +H 2 O; NOC1 + H2S04 = HCl + NO SO 4 H; NO SO 4 H + NaC1 = Noci+NaHS04 (W. A. Tilden, Jour. Chem. Soc., 1860, p. 630). It is also prepared by the action of phosphorus pentachloride on potassium nitrite or on nitrogen peroxide. It is an orange-coloured gas which may be readily liquefied and by further cooling may be solidified. The liquid boils at -5° C. and the solid melts at -65° C. It forms double compounds with many metallic chlorides, and finds considerable application as a means of separating various members of the terpene group of compounds. It is readily decomposed by water and alkaline hydroxides, yielding a mixture of nitrite and chloride. On treatment with silver fluoride it yields nitrosyl fluoride, NOF (0. Ruff, Zeit. anorg. Chem., 1905 47, p. 190). Nitroxyl fluoride, NO 2 F, is formed by the action of fluorine on nitric oxide at the temperature of liquid oxygen (H. Moissan and P. Lebeau, Comptes rendus, 1905, 140, pp. 1573, 1621). It is a gas at ordinary temperature; when liquefied it boils at -63.5° C. and on solidification melts at -139° C. Water decomposes it into nitric and hydrofluoric acids. Nitramide, NH 2 NO 2, is obtained by the action of sulphuric and nitric acids on potassium imidosulphonate, or by the action of ice-cold sulphuric acid on potassium nitro-carbamate (J. Thiele and A. Lachmann, Ann., 1895, 288, p. 297): N02 NK C02K+H2S04 = NH2N02+K2S04+C02. It crystallizes in prisms or leaflets which melt at 72-75°C. and are readily soluble in water and in all organic solvents except ligroin. It is somewhat volatile at ordinary temperature, and its aqueous solution possesses a strongly acid reaction. It is very unstable, decomposing into nitrous oxide and water when mixed with copper oxide, lead chromate or even powdered glass. On reduction it gives a strongly reducing substance, probably hydrazine. According to A. Hantzsch (Ann., 1896, 2 9 2, pp. 34 0 et seq.) hyponitrous acid and nitramide are to be regarded as stereoisomers, being the anti-and synforms of the same compound. Thiele, however, regards nitramide as imidonitric acid, HN :NO(OH). Nitrogen sulphide, N4S4, first obtained by W. Gregory (Jour. pharm., 1835, 21, p. 315) by the action of ammonia on sulphur chloride, has been investigated by 0. Ruff and E. Geisel (Ber., 1904, 37, p. 1 573; 1905, 38, p. 2659), who also obtained it by dissolving sulphur in liquid ammonia. It is a reddish-yellow crystalline solid, insoluble in water and melting at 178° C. It explodes readily when melted or subjected to shock. Dry hydrochloric acid gives ammonia but no nitrogen; with ammonia it gives N:SNH 2 and S :S(NH 2) 2; and with secondary amines it forms thiodiamines, S(NR2)2, nitrogen and ammonia being liberated. When heated with CS 2 to 1 00° C. under pressure, it forms liquid nitrogen sulphide, N 2 S 5, a mobile red liquid which solidifies to an iodine-like mass of crystals which melt at Io-I I° C. Water, alkalis and acids decompose it into sulphur and ammonia (W. Muthmann, Zeit. anorg. Chem., 18 97, 13, 200).
Numerous determinations of the atomic weight of nitrogen have been made by different observers, the values obtained varying somewhat according to the methods used. These methods have been purely chemical (either gravimetric or volumetric), physical (determinations of the density of nitrogen, nitric oxide, &c.) or physicochemical. P. A. Guye has given a critical discussion of the relative accuracy of the gravimetric and physico-chemical methods, and favours the latter, giving for the atomic weight a value less than 14.01. The more important papers dealing with the subject are: J. Stas, Ouvres completes, i. pp. 342 et seq.; Lord Rayleigh, Proc. Roy. Soc. (1894), 55, p. 34 0; (1904) 73, p. 1 53; G. Dean, Jour. Chem. Soc. (1901), 79, p. 1 47; R. W. Gray, Jour. Chem. Soc. (1906), 88, p. 1174; A. Scott, Proc. Chem. Soc. (1905), 21, p. 309; P. A. Guye, Chem. News (1905), 92, pp. 261 et seq.; (1906) 93, p. 13 et seq.; D. Berthelot, Comptes rendus (1907), 1 44, P z69
|N||Previous: carbon (C)|
|Next: oxygen (O)|
For etymology and more information refer to: http://elements.vanderkrogt.net/elem/be.html (A lot of the translations were taken from that site with permission from the author)
From French nitrogène
nitrogen m. (uncountable)
Nitrogen is a nonmetal chemical element. It has the chemical symbol N and atomic number 7. Its stable nuclei typically contains 14 nucleons (7 protons and 7 neutrons). It has 5 electrons in its outermost shell.
Nitrogen is a colorless odorless gas at normal temperature. It is normally bonded to another nitrogen atom, making a nitrogen molecule (N2). The bond is very strong. That is why many explosives contain nitrogen. The bond is broken when the explosive is made. When it explodes the bond forms, releasing a lot of energy.
It turns into a liquid at -195.8°C and turns into a solid at -210°C. If it is compressed, it can be turned into a liquid without making it cold.
It is very unreactive because of its strong bond. The bond prevents it from reacting. Lithium is one of the only chemical elements that react with nitrogen without being heated. Magnesium can burn in nitrogen. Nitrogen also makes blue electric sparks. The blue color is caused by the atoms being excited. When they get normal again, they release light. When nitrogen is excited, it reacts with many things that it does not normally react with. [[File:|thumb|A nice electric spark through a tube filled with nitrogen]]
Pure liquid nitrogen can be made by cooling air. The nitrogen turns into a liquid at a different temperature than the oxygen. It can also be made by heating certain chemical compounds, such as sodium azide.
Nitrogen as an element is used to prevent things from reacting with the oxygen in the air. It can be used to fill chip bags and incandescent bulbs. It is also used to fill some tires. It can be used to make electric components like transistors. Liquid nitrogen can be used to freeze things.
Nitrogen was discovered by Daniel Rutherford in 1772, who called it noxious gas or fixed gas. They discovered that part of air did not burn. It was found that animals died in it. It was known as "azote". Many nitrogen compounds also contain the "azide" or "azine" letters, such as hydrazine.
In 1910, Lord Rayleigh found out that when a spark was passed through nitrogen, it made a reactive form of nitrogen. This nitrogen reacted with many metals and compounds.
Nitrogen is not poisonous. We can safely breathe it when it is a part of air, but we cannot breathe pure nitrogen by itself, because it does not have the oxygen that we need to live. Someone that breaths in pure nitrogen will just fall asleep and die.