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{{Infobox nitrogen}}
{{Infobox nitrogen}}


'''Nitrogen''' ({{IPAc-en|icon|ˈ|n|aɪ|t|r|ɵ|dʒ|ɨ|n}} {{respell|NYE|tro-jin}}) 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 78.08% by volume of [[Earth's atmosphere]]. The element nitrogen was discovered as a separable component of air, by Scottish physician [[Daniel Rutherford]], in 1772.
'''Nitrogen''' ({{IPAc-en|icon|ˈ|n|aɪ|t|r|ɵ|dʒ|ɨ|n}} {{respell|NYE|tro-jin}}) is a [[A fat guy eating to much]] that has the symbol no symbol at all, [[atomic number]] of 9 and [[atomic mass]] 14.00674 u. Elemental nitrogen is a colorless, odorless, tasteless and mostly [[inert]] [[diatomic]] gas at [[standard conditions]], constituting 78.08% by volume of [[Earth's atmosphere]]. The element nitrogen was discovered as a separable component of air, by Scottish physician [[Daniel Rutherford]], in 1772.


Many industrially important compounds, such as [[ammonia]], [[nitric acid]], organic nitrates ([[propellant]]s and [[explosive]]s), and [[cyanide]]s, 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 {{chem|N|2}} into useful [[Chemical compound|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.
Many industrially important compounds, such as [[ammonia]], [[nitric acid]], organic nitrates ([[propellant]]s and [[explosive]]s), and [[cyanide]]s, 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 {{chem|N|2}} into useful [[Chemical compound|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.

Revision as of 18:35, 22 October 2010

Nitrogen, 7N
A transparent liquid, with visible evaporation, being poured
Liquid nitrogen (N2 at below −196 °C)
Nitrogen
Allotropessee § Allotropes
Appearancecolorless gas, liquid or solid
Standard atomic weight Ar°(N)
Nitrogen in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
–
↑
N
↓
P
carbon ← nitrogen → oxygen
Atomic number (Z)7
Groupgroup 15 (pnictogens)
Periodperiod 2
Block  p-block
Electron configuration[He] 2s2 2p3
Electrons per shell2, 5
Physical properties
Phase at STPgas
Melting point(N2) 63.23[3] K ​(−209.86[3] °C, ​−345.75[3] °F)
Boiling point(N2) 77.355 K ​(−195.795 °C, ​−320.431 °F)
Density (at STP)1.2506 g/L[4] at 0 °C, 1013 mbar
when liquid (at b.p.)0.808 g/cm3
Triple point63.151 K, ​12.52 kPa
Critical point126.21 K, 3.39 MPa
Heat of fusion(N2) 0.72 kJ/mol
Heat of vaporization(N2) 5.57 kJ/mol
Molar heat capacity(N2) 29.124 J/(mol·K)
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 37 41 46 53 62 77
Atomic properties
Oxidation statescommon: −3, +3, +5
−2,[5] −1,[5] 0,[6] +1,[5] +2,[5] +4[5]
ElectronegativityPauling scale: 3.04
Ionization energies
  • 1st: 1402.3 kJ/mol
  • 2nd: 2856 kJ/mol
  • 3rd: 4578.1 kJ/mol
  • (more)
Covalent radius71±1 pm
Van der Waals radius155 pm
Color lines in a spectral range
Spectral lines of nitrogen
Other properties
Natural occurrenceprimordial
Crystal structure ​hexagonal (hP4)
Lattice constants
Hexagonal crystal structure for nitrogen
a = 411.6 pm
c = 673.4 pm (at t.p.)[7]
Thermal conductivity25.83×10−3 W/(m⋅K)
Magnetic orderingdiamagnetic
Speed of sound353 m/s (gas, at 27 °C)
CAS Number17778-88-0
7727-37-9 (N2)
History
DiscoveryDaniel Rutherford (1772)
Named byJean-Antoine Chaptal (1790)
Isotopes of nitrogen
Main isotopes Decay
abun­dance half-life (t1/2) mode pro­duct
13N trace 9.965 min β+ 13C
14N 99.6% stable
15N 0.4% stable
16N synth 7.13 s β− 16O
β−α<0.01% 12C
 Category: Nitrogen
| references

Nitrogen (/[invalid input: 'icon']ˈnaɪtr[invalid input: 'ɵ']dʒ[invalid input: 'ɨ']n/ NYE-tro-jin) is a A fat guy eating to much that has the symbol no symbol at all, atomic number of 9 and atomic mass 14.00674 u. Elemental nitrogen is a colorless, odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78.08% by volume of Earth's atmosphere. The element nitrogen was discovered as a separable component of air, by Scottish physician Daniel Rutherford, in 1772.

Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in breaking the bond to convert the N
2 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, and the nitrogen cycle describes movement of the element from air into the biosphere and organic compounds, then back into the atmosphere. Synthetically-produced nitrates are key ingredients of industrial fertilizers, and also key pollutants in causing the eutrophication of water systems. Nitrogen 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.

History

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.[8] The fact that there was an element of air which did 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 άζωτος (azotos) meaning "lifeless".[9] In it animals died and flames were extinguished. 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.

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 involved uses of saltpeter (sodium nitrate or potassium nitrate), 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.[10]

Properties

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 (N
2) is the strongest. The resulting difficulty of converting N
2 into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N
2, have dominated the role of nitrogen in both nature and human economic activities.

At atmospheric pressure molecular nitrogen condenses (liquefies) at 77 K (−195.8 °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 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 N
3 and N
4.[11] 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. N
4 is nicknamed "nitrogen diamond."[12]

Isotopes

See also: Isotopes of nitrogen

There are two stable isotopes of nitrogen: 14N and 15N. By far the most common is 14N (99.634%), which is produced in the CNO cycle in stars. 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 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.[13] 16N is one of the main means used to immediately detect even small leaks from the primary coolant to the secondary steam cycle.

Electromagnetic spectrum

A 1×5 cm vial of glowing ultrapure nitrogen
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).

Reactions

Structure of [Ru(NH3)5(N2)]2+.

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.

Nitrogen reacts with elemental lithium.[14] 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). Such compounds are now numerous, other examples include IrCl(N2)(PPh3)2, W(N2)2(Ph2CH2CH2PPh2)2, and [(η5-C5Me4H)2Zr]2(μ2, η²,η²-N2). These complexes illustrate how N2 might bind to the metal(s) in nitrogenase and the catalyst for the Haber process.[15] 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.[14]

The starting point for industrial production of nitrogen compounds is the Haber process, in which nitrogen is fixed by reacting N
2 and H
2 over an iron(III) oxide (Fe
3O
4) catalyst at about 500 °C and 200 atmospheres pressure. Biological nitrogen fixation in free-living cyanobacteria and in the root nodules of plants also produces ammonia from molecular nitrogen. The reaction, which is the source of the bulk of nitrogen in the biosphere, is catalyzed by the nitrogenase enzyme complex 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).

Occurrence

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.[16]

Molecular nitrogen and nitrogen compounds have been detected in interstellar space by astronomers using the Far Ultraviolet Spectroscopic Explorer.[17] Molecular nitrogen is a major constituent of the Saturnian moon Titan's thick atmosphere, and occurs in slightly appreciable to trace amounts in other planetary atmospheres.[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.

Nitrogen occurs naturally in many minerals, such as saltpetre (potassium nitrate), Chile saltpetre (sodium nitrate) and sal ammoniac (ammonium chloride). Most of these are uncommon, partly because of the minerals' ready solubility in water. See also Nitrate minerals and Ammonium minerals.

Compounds

See also: Category:Nitrogen compounds

The main neutral hydride of nitrogen is ammonia (NH
3), although hydrazine (N
2H
4) is also commonly used. Ammonia is more basic than water by 6 orders of magnitude. In solution ammonia forms the ammonium ion (NH+
4). Liquid ammonia (boiling point 240 K) is amphiprotic (displaying either Brønsted-Lowry acidic or basic character) and forms ammonium and the less common amide ions (NH−
2); both amides and nitride (N3−
) salts are known, 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, N
2O), also known as laughing gas. This is one of a variety of nitrogen oxides that form a family often abbreviated as NOx. Nitric oxide (nitrogen monoxide, NO), is a natural free radical used in signal transduction in both plants and animals, for example in vasodilation by causing the smooth muscle of blood vessels to relax. The reddish and poisonous nitrogen dioxide NO
2 contains an unpaired electron and is an important component of smog. Nitrogen molecules containing unpaired electrons show an understandable tendency to dimerize (thus pairing the electrons), and are generally highly reactive. The corresponding acids are nitrous HNO
2 and nitric acid HNO
3, with the corresponding salts called nitrites and nitrates.

The higher oxides dinitrogen trioxide N
2O
3, dinitrogen tetroxide N
2O
4 and dinitrogen pentoxide N
2O
5, are unstable and explosive, a consequence of the chemical stability of N
2. Nearly every hypergolic rocket engine uses N
2O
4 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. N
2O
4 is an intermediate in the manufacture of nitric acid HNO
3, 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 NI
3 is an extremely sensitive contact explosive. Nitrocellulose, produced by nitration of cellulose with nitric acid, is also known as guncotton. Nitroglycerin, made by nitration of glycerin, is the dangerously unstable explosive ingredient of dynamite. The comparatively stable, but 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.

Production and applications

A computer rendering of the nitrogen molecule, N2.

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).[19]

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. Specifically, 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. The downside is that nitrogen must be kept at higher pressure than CO2, making N2 tanks heavier and more expensive.

Nitrogenated beer

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.[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]

Liquid nitrogen

Air balloon submerged into liquid nitrogen
Main article: 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 medically in cryotherapy to remove cysts and warts on the skin.[29] 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.[30]

Applications of nitrogen compounds

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). Drugs that contain nitrogen include all major classes of antibiotics, and organic nitrate drugs like nitroglycerin and nitroprusside that regulate blood pressure and heart action by mimicking the action of nitric oxide.

Biological role

See also: Nitrogen cycle and Human impacts on the nitrogen cycle

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 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.[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 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.[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 anerobic (oxygen free) environment for their nitrogen-fixing 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.[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 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 which are (respectively) breakdown products of the amino acids ornithine and lysine 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]

Safety

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 (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.[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 eventually 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 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.

See also

References

  1. ^ "Standard Atomic Weights: Nitrogen". CIAAW. 2009.
  2. ^ Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.
  3. ^ a b c Lide, David R. (1990–1991). CRC Handbook of Physics and Chemistry (71st ed.). Boca Raton, Ann Arbor, Boston: CRC Press, inc. pp. 4-22 (one page).
  4. ^ "Gases - Density". The Engineering Toolbox. Retrieved 27 January 2019.
  5. ^ a b c d e Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 28. ISBN 978-0-08-037941-8.
  6. ^ Tetrazoles contain a pair of double-bonded nitrogen atoms with oxidation state 0 in the ring. A Synthesis of the parent 1H-tetrazole, CH2N4 (two atoms N(0)) is given in Henry, Ronald A.; Finnegan, William G. (1954). "An Improved Procedure for the Deamination of 5-Aminotetrazole". Journal of the American Chemical Society. 76 (1): 290–291. doi:10.1021/ja01630a086. ISSN 0002-7863.
  7. ^ Arblaster, John W. (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.
  8. ^ Lavoisier, Antoine Laurent (1965). Elements of chemistry, in a new systematic order: containing all the modern discoveries. Courier Dover Publications. p. 15. ISBN 0486646246.
  9. ^ Elements of Chemistry, trans. Robert Kerr (Edinburgh, 1790; New York: Dover, 1965), 52.
  10. ^ Lord Rayleigh's Active Nitrogen
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  13. ^ Karl Heinz Neeb (1997). The Radiochemistry of Nuclear Power Plants with Light Water Reactors. Berlin-New York: Walter de Gruyter. ISBN 3110132427.
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  15. ^ 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.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  16. ^ Croswell, Ken (1996). Alchemy of the Heavens. Anchor. ISBN 0-385-47214-5. {{cite book}}: Cite has empty unknown parameter: |coauthors= (help); Unknown parameter |month= ignored (help)
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  22. ^ "Why don't they use normal air in race car tires?". Howstuffworks. Retrieved 2006-07-22.
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  26. ^ "How does the widget in a beer can work?". Howstuffworks.
  27. ^ http://www.bomby.cz/plneni.htm
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  29. ^ PMID 11359389
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  31. ^ a b Rakov, Vladimir A. (2007). Lightning: Physics and Effects. Cambridge University Press. p. 508. ISBN 9780521035415. {{cite book}}: Unknown parameter |coauthor= ignored (|author= suggested) (help)
  32. ^ a b c d Bothe, Hermann; Ferguson, Stuart John; Newton, William Edward (2007). Biology of the nitrogen cycle. Elsevier. p. 283. ISBN 0444528571.{{cite book}}: CS1 maint: multiple names: authors list (link)
  33. ^ 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. {{cite journal}}: Explicit use of et al. in: |author= (help)
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  36. ^ Vickerstaff Joneja, Janice M. (2004). Digestion, diet, and disease: irritable bowel syndrome and gastrointestinal function. Rutgers University Press. p. 121. ISBN 0813533872.
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  39. ^ Rogers, W. H.; Moeller, G. (1989). "Effect of brief, repeated hyperbaric exposures on susceptibility to nitrogen narcosis". Undersea Biomed. Res. 16 (3): 227–32. ISSN 0093-5387. OCLC 2068005. PMID 2741255. Retrieved 2008-09-21.{{cite journal}}: CS1 maint: multiple names: authors list (link)
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Further reading

  • Garrett, Reginald H. (1999). Biochemistry (2nd ed.). Fort Worth: Saunders College Publ. ISBN 0030223180. {{cite book}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
  • Greenwood, Norman N. (1984). Chemistry of the Elements. Oxford: Pergamon Press. ISBN 0080220576. {{cite book}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
  • "Nitrogen". Los Alamos National Laboratory. 2003-10-20.
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