|
|
| Appearance |
colorless gas, liquid or solid
Liquid nitrogen
Spectral lines of nitrogen |
| General properties |
| Name, symbol, number |
nitrogen, N, 7 |
| Pronunciation |
/ˈnaɪtrɵdʒən/ NYE-trə-jən |
| Element category |
nonmetal |
| Group, period, block |
15, 2, p |
| Standard atomic weight |
14.0067(2) |
| Electron configuration |
1s2 2s2 2p3 |
| Electrons per shell |
2, 5 (Image) |
| Physical properties |
| Phase |
gas |
| Density |
(0 °C, 101.325 kPa)
1.251 g/L |
| Liquid density at b.p. |
0.808 g·cm−3 |
| Melting point |
63.15 K, -210.00 °C, -346.00 °F |
| Boiling point |
77.36 K, -195.79 °C, -320.33 °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 |
| Molar heat capacity |
(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 |
| Oxidation states |
5, 4, 3, 2, 1, -1, -2, -3
(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 |
| Covalent radius |
71±1 pm |
| Van der Waals radius |
155 pm |
| Miscellanea |
| Crystal structure |
hexagonal |
| Magnetic ordering |
diamagnetic |
| Thermal conductivity |
25.83 × 10−3 W·m−1·K−1 |
| Speed of sound |
(gas, 27 °C) 353 m·s−1 |
| CAS registry number |
7727-37-9 |
| Most stable isotopes |
| Main article: Isotopes of nitrogen |
|
|
|
· r |
Nitrogen (
/ˈnaɪtrɵdʒən/ NY-trə-jən) is a chemical element that has the symbol N, atomic number of 7 and atomic mass 14.00674 u. Elemental nitrogen is a colorless, odorless, tasteless, and mostly inert diatomic gas at standard conditions, constituting 78.09% by volume of Earth's atmosphere.[1] The element nitrogen was discovered as a separable component of air, by Scottish physician Daniel Rutherford, in 1772.
Nitrogen is a common element in the universe, estimated at about seventh in total abundance in our galaxy and the solar system. Its occurrence there is thought to be entirely due to synthesis by fusion from carbon and hydrogen in supernovas. Due to the volatility of elemental nitrogen and its common compounds with hydrogen and oxygen, nitrogen is far less common on the rocky planets of the inner solar system, and it is a relatively rare element on Earth as a whole. However, as is the case on Earth, nitrogen and its compounds occur commonly as gases in the atmospheres of planets and moons that have atmospheres.
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 breaking the bond to convert the N2 into useful compounds, but at the same time causing release of large amounts of often useful energy when the compounds burn, explode, or decay back into nitrogen gas.
Nitrogen occurs in all living organisms, primarily in amino acids and thus proteins and in the nucleic acids (DNA and RNA). The human body contains about 3% by weight of nitrogen, the fourth most abundant element after oxygen, carbon, and hydrogen. Nitrogen resides in the chemical structure of almost all neurotransmitters, and is a defining component of alkaloids, biological molecules produced as secondary metabolites by many organisms. The nitrogen cycle describes movement of the element from the air into the biosphere and organic compounds, then back into the atmosphere. Synthetically produced nitrates are key ingredients of industrial fertilizers,[1] and key pollutants in causing the eutrophication of water systems.
Nitrogen is formally considered to have been discovered by Scottish physician Daniel Rutherford in 1772, who called it noxious air or fixed air.[2] The fact that there was an element of air that does not support combustion was clear to Rutherford. 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 ἄζωτος (azotos) meaning "lifeless".[3] In it, animals died and flames were extinguished. Lavoisier's name for nitrogen is used in many languages (French, Polish, Russian, etc.) and still remains in English in the common names of many compounds, such as hydrazine and compounds of the azide ion.
The English word nitrogen (1794) entered the language [4] from the French nitrogène, coined in 1790 by French chemist Jean-Antoine Chaptal (1756–1832), from "nitre" + Fr. gène "producing" (from Gk. -γενής means "forming" or "giving birth to."). The gas had been found in nitric acid. Chaptal's meaning was that nitrogen gas is the essential part of nitric acid, in turn formed from saltpetre (potassium nitrate), then known as nitre. This word in the more ancient world originally described sodium salts that did not contain nitrate, and is a cognate of natron.
Nitrogen compounds were well known during the Middle Ages. Alchemists knew nitric acid as aqua fortis (strong water). The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold (the king of metals). The earliest military, industrial, and agricultural applications of nitrogen compounds used saltpetre (sodium nitrate or potassium nitrate), most notably in gunpowder, and later as fertilizer. In 1910, Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", an allotrope considered to be monatomic. The "whirling cloud of brilliant yellow light" produced by his apparatus reacted with quicksilver to produce explosive mercury nitride.[5]
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 called OFN (oxygen-free nitrogen).[6]
In a chemical laboratory it is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite.
- NH4Cl(aq) + NaNO2(aq) → N2(g) + NaCl(aq) + 2 H2O (l)
Small amounts of impurities NO and HNO3 are also formed in this reaction. The impurities can be removed by passing the gas through aqueous sulfuric acid containing potassium dichromate. Very pure nitrogen can be prepared by the thermal decomposition of barium or sodium azide.
- 2 NaN3 → 2 Na + 3 N2
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 one of the strongest. 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 (liquefies) at 77 K (−195.79 °C) and freezes at 63 K (−210.01 °C)[1] into the beta hexagonal close-packed crystal allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the cubic crystal allotropic form (called the alpha-phase). 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.[7] 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."[8]
Other (as yet unsynthesized) allotropes include hexazine (N6, a benzene analog)[9] and octaazacubane (N8, a cubane analog).[10] The former is predicted to be highly unstable, while the latter is predicted to be kinetically stable, for reasons of orbital symmetry.[11]
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.
A small part (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 or boiling water reactors during normal operation. It is produced from 16O (in water) via (n,p) reaction. It has a short half-life of about 7.1 s, 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 in a pressurized water reactor must be restricted during reactor power operation.[12] 16N is one of the main means used to immediately detect even small leaks from the primary coolant to the secondary steam cycle.
In similar fashion, access to any of the steam cycle components in a boiling water reactor nuclear power plant must be restricted during operation. Condensate from the condenser is typically retained for 10 minutes to allow for decay of the 16N. This eliminates the need to shield and restrict access to any of the feed water piping or pumps.
Nitrogen discharge (spectrum) tube
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 and the atmospheres of other planetary bodies. For similar reasons, pure molecular nitrogen lasers typically emit light in the ultraviolet range.
Nitrogen also makes a contribution to visible air glow from the Earth's upper atmosphere, through electron impact excitation followed by emission. This visible blue air glow (seen in the polar aurora and in the re-entry glow of returning spacecraft) typically results not from molecular nitrogen but rather from free nitrogen atoms combining with oxygen to form nitric oxide (NO).
Nitrogen gas also exhibits scintillation.
Structure of dinitrogen, N
2
Structure of [Ru(NH
3)
5(N
2)]
2+
In general, nitrogen is unreactive at standard temperature and pressure. N2 reacts spontaneously with few reagents, being resilient to acids and bases as well as oxidants and most reductants. When nitrogen reacts spontaneously with a reagent, the net transformation is often called nitrogen fixation.
Nitrogen reacts with elemental lithium.[13] Lithium burns in an atmosphere of N2 to give lithium nitride:
- 6 Li + N2 → 2 Li3N
Magnesium also burns in nitrogen, forming magnesium nitride.
- 3 Mg + N2 → Mg3N2
N2 forms a variety of adducts with transition metals. The first example of a dinitrogen complex is [Ru(NH3)5(N2)]2+ (see figure at right). However, it is interesting to note that the N2 ligand was obtained by the decomposition of hydrazine, and not coordination of free dinitrogen. Such compounds are now numerous, other examples include IrCl(N2)(PPh3)2, W(N2)2(Ph2PCH2CH2PPh2)2, and [(η5-C5Me4H)2Zr]2(μ2, η2,η2-N2). These complexes illustrate how N2 might bind to the metal(s) in nitrogenase and the catalyst for the Haber process.[14] 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.[13]
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(II, 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 that contains Fe and Mo atoms, using energy derived from hydrolysis of adenosine triphosphate (ATP) into adenosine diphosphate and inorganic phosphate (−20.5 kJ/mol).
Nitrogen is the largest constituent of the Earth's atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air). However, this high concentration does not reflect nitrogen's overall low abundance in the makeup of the Earth, from which most of the element escaped by solar evaporation, early in the planet's formation.
Nitrogen is a common element in the universe, and is estimated to be approximately seventh most abundant chemical element by mass in the universe, our galaxy and the solar system. Its occurrence there is thought to be entirely due to synthesis by fusion from carbon and hydrogen in supernovas. In these places it was originally created by fusion processes from carbon and hydrogen in supernovas.[15]Molecular nitrogen and nitrogen compounds have been detected in interstellar space by astronomers using the Far Ultraviolet Spectroscopic Explorer.[16]
Due to the volatility of elemental nitrogen and also its common compounds with hydrogen and oxygen, nitrogen and its compounds were driven out of the planetesimals in the early solar system by the heat of the Sun, and in the form of gases, were lost to the rocky planets of the inner solar system. Nitrogen is therefore a relatively rare element on these inner planets, including Earth, as a whole. In this, nitrogen resembles neon, which has a similar abundance in the universe, but is also rare in the inner solar system. Nitrogen is estimated at 30th of the elements in crustal abundance. There exist some relatively uncommon nitrogen minerals, such as saltpetre (potassium nitrate), Chile saltpetre (sodium nitrate) and sal ammoniac (ammonium chloride). Even these are known mainly as concentrated from evaporative ocean beds, on account of their ready solubility of most naturally-occurring nitrogen compounds in water. A similar pattern occurs with the water solubility of the uncommon light element boron.
However, nitrogen and its compounds occur far more commonly as gases in the atmospheres of planets and moons that are large enough to have atmospheres.[17] For example, molecular nitrogen is a major constituent of not only Earth's atmosphere, but also the Saturnian moon Titan's thick atmosphere. Also, due to retension by gravity at colder temperatures, nitrogen and its compounds occur in appreciable to trace amounts in planetary atmospheres of the gas giant planets.[18]
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 cannot fix atmospheric nitrogen.
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.
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 a tendency to dimerize (thus pairing the electrons), and are, in general, 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 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 less powerful explosive trinitrotoluene (TNT) is the standard explosive against which the power of nuclear explosions are measured.[19]
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.
A computer rendering of the nitrogen molecule, N
2
Nitrogen gas has a variety of applications, including serving as an inert replacement for air where oxidation is undesirable;
Nitrogen is commonly used during sample preparation procedures for chemical analysis. It is used to concentrate and reduce the volume of liquid samples. Directing a pressurized stream of nitrogen gas perpendicular to the surface of the liquid allows the solvent to evaporate while leaving the solute(s) and un-evaporated solvent behind.[25]
Nitrogen tanks are also replacing carbon dioxide as the main power source for paintball guns. Nitrogen must be kept at higher pressure than CO2, making N2 tanks heavier and more expensive.
Nitrogen can be used instead of carbon dioxide to pressurize kegs of some beers, in particular, stouts and British ales, due to the smaller bubbles it produces, which makes the dispensed beer smoother and headier. A pressure sensitive nitrogen capsule known commonly as a "widget" allows nitrogen charged beers to be packaged in cans and bottles.[26]
A mixture of nitrogen and carbon dioxide can be used for this purpose as well, to maintain the saturation of beer with carbon dioxide.[27]
Air balloon submerged in liquid nitrogen
Liquid nitrogen is a cryogenic liquid. At atmospheric pressure, it boils at −195.8 °C. When insulated in proper containers such as Dewar flasks, it can be transported without much evaporative loss.[28]
Like dry ice, the main use of liquid nitrogen is as a refrigerant. Among other things, it is used in the cryopreservation of blood, reproductive cells (sperm and egg), and other biological samples and materials. It is used in the clinical setting in cryotherapy to remove cysts and warts on the skin.[29] It is used in cold traps for certain laboratory equipment and to cool infrared detectors or X-ray detectors. It has also been used to cool central processing units and other devices in computers that are overclocked, and that produce more heat than during normal operation.[30]
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),[1] 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, 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 as natural chemical defenses of plants against predation). Drugs that contain nitrogen include all major classes of antibiotics and organic nitrate drugs like nitroglycerin and nitroprusside that regulate blood pressure and heart action by mimicking the action of nitric oxide.
Nitrogen is an essential building block of amino and nucleic acids, essential to life on Earth.
Elemental nitrogen in the atmosphere cannot be used directly by either plants or animals, and must be converted to a reduced (or 'fixed') state to be useful for higher plants and animals. Precipitation often contains substantial quantities of ammonium and nitrate, thought to result from nitrogen fixation by lightning and other atmospheric electric phenomena.[31] This was first proposed by Liebig in 1827 and later confirmed.[31] However, because ammonium is preferentially retained by the forest canopy relative to atmospheric nitrate, most fixed nitrogen reaches the soil surface under trees as nitrate. Soil nitrate is preferentially assimilated by tree roots relative to soil ammonium[citation needed].
Specific bacteria (e.g., Rhizobium trifolium) possess nitrogenase enzymes that can fix atmospheric nitrogen (see nitrogen fixation) into a form (ammonium ion) that is chemically useful to higher organisms. This process requires a large amount of energy and anoxic conditions. Such bacteria may live freely in soil (e.g., Azotobacter) but normally exist in a symbiotic relationship in the root nodules of leguminous plants (e.g. clover, Trifolium, or soybean plant, Glycine max). Nitrogen-fixing bacteria are also symbiotic with a number of unrelated plant species such as alders (Alnus) spp., lichens, Casuarina, Myrica, liverworts, and Gunnera.[32]
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.[32] Legumes maintain an anaerobic (oxygen free) environment for their nitrogen-fixing bacteria.
Plants are able to assimilate nitrogen directly in the form of nitrates that may be present in soil from natural mineral deposits, artificial fertilizers, animal waste, or organic decay (as the product of bacteria, but not bacteria specifically associated with the plant). Nitrates absorbed in this fashion are converted to nitrites by the enzyme nitrate reductase, and then converted to ammonia by another enzyme called nitrite reductase.[32]
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.[33]
Soluble nitrate is an important limiting factor in the growth of certain bacteria in ocean waters.[34] 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.[35] In animals, free radical nitric oxide (NO) (derived from an amino acid), serves as an important regulatory molecule for circulation.[34]
Animal metabolism of NO results in production of nitrite. Animal metabolism of nitrogen in proteins, in general, results in excretion of urea, while animal metabolism of nucleic acids results in excretion of urea and uric acid. The characteristic odor of animal flesh decay is caused by the creation of long-chain, nitrogen-containing amines, such as putrescine and cadaverine, which are breakdown products of the amino acids ornithine and lysine, respectively, in decaying proteins.[36]
Decay of organisms and their waste products may produce small amounts of nitrate, but most decay eventually returns nitrogen content to the atmosphere, as molecular nitrogen. The circulation of nitrogen from atmosphere, to organic compounds, then back to the atmosphere, is referred to as the nitrogen cycle.[32]
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.[37] An example occurred shortly before the launch of the first Space Shuttle mission in 1981, when two technicians lost consciousness (and one of them 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.[38][39]
Nitrogen also dissolves in the bloodstream and body fats. Rapid decompression (in particular, 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 the bends), when nitrogen bubbles form in the bloodstream, nerves, joints, and other sensitive or vital areas.[40][41] 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.[42]
Direct skin contact with liquid nitrogen will cause severe frostbite (cryogenic "burns"). This may happen almost instantly on contact, or after a second or more, depending on the form of liquid nitrogen. Bulk liquid nitrogen causes less rapid freezing than a spray of nitrogen mist (such as is used to freeze certain skin growths in the practice of dermatology). The extra surface area provided by nitrogen-soaked materials is also important, with soaked clothing or cotton causing far more rapid damage than a spill of direct liquid to skin. Full "contact" between naked skin and large collected-droplets or pools of liquid nitrogen may be prevented for a second or two, by a layer of insulating gas from the Leidenfrost effect. This may give the skin a second of protection from nitrogen bulk liquid. However, liquid nitrogen applied to skin in mists, and on fabrics, bypasses this effect, and causes local frostbite immediately.
Oxygen sensors are sometimes used as a safety precaution when working with liquid nitrogen to alert workers of gas spills into a confined space.[42]
- ^ a b c d Gray, Theodore (2009). The Elements: A Visual Exploration of Every Known Atom in the Universe. New York: Black Dog & Leventhal Publishers. ISBN 978-1-57912-814-2.
- ^ Lavoisier, Antoine Laurent (1965). Elements of chemistry, in a new systematic order: containing all the modern discoveries. Courier Dover Publications. p. 15. ISBN 0-486-64624-6. http://books.google.com/?id=yS_m3PrVbpgC&pg=PR15.
- ^ Elements of Chemistry, trans. Robert Kerr (Edinburgh, 1790; New York: Dover, 1965), 52.
- ^ nitrogen. Etymonline.com. Retrieved on 2011-10-26.
- ^ Lord Rayleigh's Active Nitrogen. Lateralscience.co.uk. Retrieved on 2011-10-26.
- ^ Reich, Murray; Kapenekas, Harry (1957). "Nitrogen Purfication. Pilot Plant Removal of Oxygen". Industrial & Engineering Chemistry 49 (5): 869–873. DOI:10.1021/ie50569a032.
- ^ "A new molecule and a new signature – Chemistry – tetranitrogen". Science News. February 16, 2002. http://www.findarticles.com/p/articles/mi_m1200/is_7_161/ai_83477565. Retrieved 2007-08-18.
- ^ "Polymeric nitrogen synthesized". physorg.com. 2004-08-05. http://www.physorg.com/news693.html. Retrieved 2009-06-22.
- ^ Fabian, J. and Lewars, E. (2004). "Azabenzenes (azines) — The nitrogen derivatives of benzene with one to six N atoms: Stability, homodesmotic stabilization energy, electron distribution, and magnetic ring current; a computational study". Canadian Journal of Chemistry 82 (1): 50–69. DOI:10.1139/v03-178. http://pubs.nrc-cnrc.gc.ca/rp/rppdf/v03-178.pdf.
- ^ Muir, B.. "Cubane". http://www.ch.ic.ac.uk/local/projects/b_muir/Cubane/Cubanepro/Start.html(See "further topics" section.)
- ^ Ujwala N. Patil, Nilesh R. Dhumal and Shridhar P. Gejji. "Theoretical studies on the molecular electron densities and electrostatic potentials in azacubanes". Theoretical Chemistry Accounts: Theory, Computation, and Modeling (Theoretica Chimica Acta) 112: p. 27-32. http://www.springerlink.com/content/w3pap8xmmju00j3e/.
- ^ Karl Heinz Neeb (1997). The Radiochemistry of Nuclear Power Plants with Light Water Reactors. Berlin-New York: Walter de Gruyter. ISBN 3-11-013242-7.
- ^ a b Schrock, R. R. (2005). "Catalytic Reduction of Dinitrogen to Ammonia at a Single Molybdenum Center". Acc. Chem. Res. 38 (12): 955–962. DOI:10.1021/ar0501121. PMC 2551323. PMID 16359167. //www.pubmedcentral.nih.gov/articlerender.fcgi?tool=pmcentrez&artid=2551323.
- ^ Fryzuk, M. D. and Johnson, S. A. (2000). "The continuing story of dinitrogen activation". Coordination Chemistry Reviews 200–202: 379. DOI:10.1016/S0010-8545(00)00264-2.
- ^ Croswell, Ken (February 1996). Alchemy of the Heavens. Anchor. ISBN 0-385-47214-5. http://kencroswell.com/alchemy.html.
- ^ Meyer, Daved M.; Cardelli, Jason A.; Sofia, Ulysses J. (1997). "Abundance of Interstellar Nitrogen". The Astrophysical Journal 490: L103–L106. arXiv:astro-ph/9710162. Bibcode 1997ApJ...490L.103M. DOI:10.1086/311023.
- ^ Note: nitrogen and its compounds are far less common in atmospheres of smaller rocky moons and planets than neon, due to being less volatile than neon.
- ^ Hamilton, Calvin J.. "Titan (Saturn VI)". Solarviews.com. http://www.solarviews.com/eng/titan.htm. Retrieved 2007-12-24.
- ^ Häring, Heinz-Wolfgang (2008). Industrial Gases Processing. John Wiley & Sons. pp. 243–. ISBN 978-3-527-62125-5. http://books.google.com/books?id=4XNBJIrdK9wC&pg=PA243. Retrieved 2 January 2012.
- ^ Harding, Charlie, ed. (2002). Elements of the p Block. Cambridge: Royal Society of Chemistry. ISBN 978-0-85404-690-4. http://books.google.com/?id=W0HW8wgmQQsC&pg=PA90.
- ^ Gavriliuk, V. G.; Berns, Hans (1999). High nitrogen steels: structure, properties, manufacture, applications. Springer. ISBN 3-540-66411-4. http://books.google.com/?id=6eF4AfEwF4YC&pg=PA338.
- ^ "Why don't they use normal air in race car tires?". Howstuffworks. http://auto.howstuffworks.com/question594.htm. Retrieved 2006-07-22.
- ^ "Why Nitrogen?". http://www.getnitrogen.org/savebillions/index.php. Retrieved 2011-06-20.
- ^ "Passenger Vehicle Studies. Nitrogen in Tires: Information about Nitrogen Tire Inflation News and Benefits". http://www.getnitrogen.org/why/index.php. Retrieved 2011-06-20.
- ^ Kemmochi, Y; Tsutsumi, K; Arikawa, A; Nakazawa, H (2002). "Centrifugal concentrator for the substitution of nitrogen blow-down micro-concentration in dioxin/polychlorinated biphenyl sample preparation". Journal of Chromatography A 943 (2): 295–297. DOI:10.1016/S0021-9673(01)01466-2. PMID 11833649.
- ^ "How does the widget in a beer can work?". Howstuffworks. http://recipes.howstuffworks.com/question446.htm.
- ^ NORBERTOVY BOMBY – rozvoz technických plynů. Bomby.cz. Retrieved on 2011-10-26.
- ^ Kaganer, M. G. et al. (1967). "Vessels for the storage and transport of liquid oxygen and nitrogen". Chemical and Petroleum Engineering 3 (12): 918–922. DOI:10.1007/BF01136404.
- ^ Ahmed I, Agarwal S, Ilchyshyn A, Charles-Holmes S, Berth-Jones J (May 2001). "Liquid nitrogen cryotherapy of common warts: cryo-spray vs. cotton wool bud". Br. J. Dermatol. 144 (5): 1006–9. PMID 11359389.
- ^ Kent, Allen; Williams, James G. (1994). Encyclopedia of Computer Science and Technology, Volume 30. CRC Press. p. 318. ISBN 0-8247-2283-3. http://books.google.com/?id=9fYFiZ7QeEcC&pg=PA318.
- ^ a b Rakov, Vladimir A.; Uman, Martin A. (2007). Lightning: Physics and Effects. Cambridge University Press. p. 508. ISBN 978-0-521-03541-5. http://books.google.com/?id=TuMa5lAa3RAC&pg=PA508.
- ^ a b c d Bothe, Hermann; Ferguson, Stuart John; Newton, William Edward (2007). Biology of the nitrogen cycle. Elsevier. p. 283. ISBN 0-444-52857-1. http://books.google.com/?id=qmZDpnV-sYYC&pg=PA283.
- ^ Jahn, G.C. et al. (2005). "Effect of nitrogen fertilizer on the intrinsic rate of increase of the rusty plum aphid, Hysteroneura setariae (Thomas) (Homoptera: Aphididae) on rice (Oryza sativa L.)". Environmental Entomology 34 (4): 938–943. DOI:10.1603/0046-225X-34.4.938. http://puck.esa.catchword.org/vl=33435372/cl=21/nw=1/rpsv/cw/esa/0046225x/v34n4/s26/p938.
- ^ a b Knox, G. A. (2007). Biology of the Southern Ocean. CRC Press. p. 392. ISBN 0-8493-3394-6. http://books.google.com/?id=qPuCVvAv2GwC&pg=PA392.
- ^ Nielsen, M. K.; Jørgensen, B. M. (Jun 2004). "Quantitative relationship between trimethylamine oxide aldolase activity and formaldehyde accumulation in white muscle from gadiform fish during frozen storage". Journal of Agricultural and Food Chemistry 52 (12): 3814–3822. DOI:10.1021/jf035169l. PMID 15186102.
- ^ Vickerstaff Joneja, Janice M. (2004). Digestion, diet, and disease: irritable bowel syndrome and gastrointestinal function. Rutgers University Press. p. 121. ISBN 0-8135-3387-2. http://books.google.com/?id=RMPis9OnJxkC&pg=PT121.
- ^ "Biology Safety – Cryogenic materials. The risks posed by them". University of Bath. http://www.bath.ac.uk/internal/bio-sci/bbsafe/asphyx.htm. Retrieved 2007-01-03.
- ^ Fowler, B; Ackles, KN; Porlier, G (1985). "Effects of inert gas narcosis on behavior—a critical review". Undersea Biomed. Res. 12 (4): 369–402. PMID 4082343. http://archive.rubicon-foundation.org/3019. Retrieved 2008-09-21.
- ^ Rogers, W. H.; Moeller, G. (1989). "Effect of brief, repeated hyperbaric exposures on susceptibility to nitrogen narcosis". Undersea Biomed. Res. 16 (3): 227–32. OCLC 2068005. PMID 2741255. http://archive.rubicon-foundation.org/2522. Retrieved 2008-09-21.
- ^ Acott, C. (1999). "A brief history of diving and decompression illness". South Pacific Underwater Medicine Society Journal 29 (2). OCLC 16986801. http://archive.rubicon-foundation.org/6004. Retrieved 2008-09-21.
- ^ Kindwall, E. P.; A. Baz; E. N. Lightfoot; E. H. Lanphier; A. Seireg. (1975). "Nitrogen elimination in man during decompression". Undersea Biomed. Res. 2 (4): 285–97. OCLC 2068005. PMID 1226586. http://archive.rubicon-foundation.org/2741. Retrieved 2008-09-21.
- ^ a b US Navy Diving Manual, 6th revision. United States: US Naval Sea Systems Command. 2006. http://www.supsalv.org/00c3_publications.asp?destPage=00c3&pageID=3.9. Retrieved 2008-04-24.
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