|Name, Symbol, Number||nitrogen, N, 7|
|Group, Period, Block||15, 2, p|
|Atomic mass||14.0067(2) g/mol|
|Electron configuration||1s2 2s2 2p3|
|Electrons per shell||2, 5|
|Density||(0 °C, 101.325 kPa)
|Melting point||63.15 K
(-210.00 °C, -346.00 °F)
|Boiling point||77.36 K
(-195.79 °C, -320.42 °F)
|Critical point||126.21 K, 3.39 MPa|
|Heat of fusion||(N2) 0.720 kJ/mol|
|Heat of vaporization||(N2) 5.57 kJ/mol|
|Heat capacity||(25 °C) (N2)
|Oxidation states||±3, 5, 4, 2
(strongly acidic oxide)
|Electronegativity||3.04 (Pauling scale)|
|1st: 1402.3 kJ/mol|
|2nd: 2856 kJ/mol|
|3rd: 4578.1 kJ/mol|
|Atomic radius||65 pm|
|Atomic radius (calc.)||56 pm|
|Covalent radius||75 pm|
|Van der Waals radius||155 pm|
|Magnetic ordering||no data|
|Thermal conductivity||(300 K) 25.83 mW/(m·K)|
|Speed of sound||(gas, 27 °C) 353 m/s|
|CAS registry number||7727-37-9|
Nitrogen (symbol N, atomic number 7) is the chief constituent of the Earth's atmosphere and a vital element in all known forms of life. At ordinary temperatures and pressures, free nitrogen (unbound to any other element) is a colorless, odorless, and tasteless gas. As an inert gas, it reduces the amount of oxygen available for the oxidation of natural materials, thus restricting spontaneous combustion of flammable materials and the corrosion of metals. It also protects living organisms from the toxic effects of breathing pure (or highly concentrated) oxygen. The Earth's nitrogen continually cycles through the atmosphere, biosphere, and lithosphere, effected by such processes as nitrogen fixation by bacteria, metabolic processing in living things, and decomposition of dead organic matter.
In living organisms, nitrogen atoms are part of the molecular structures of such key substances as amino acids, proteins, and nucleic acids. In industry, nitrogen gas is used as an inert replacement for air in the packaging of foods and the manufacture of steel and electronic components. Liquid nitrogen is a cryogen (low-temperature refrigerant) used for freezing and transport of food and other perishable products. Ammonia, a significant compound of nitrogen, is useful for fertilizers and for the synthesis of nitric acid and other valuable compounds. Nitric acid is an oxidizing agent used in liquid-fueled rockets, potassium nitrate is used in gunpowder, and trinitrotoluene (TNT) is a significant explosive. In addition, nitrogen is a constituent element in every major class of drugs.
Nitrogen makes up 78.084 percent of the volume (and 75.5 percent of the mass) of air. Its most common isotope, nitrogen-14 (14N), appears to be created through nuclear fusion processes in stars.
Compounds that contain this element have been observed by astronomers, and molecular nitrogen has been detected in interstellar space by David Knauth and coworkers using the Far Ultraviolet Spectroscopic Explorer. Molecular nitrogen occurs in trace amounts in the atmospheres of various planets, but it is a major constituent of Titan, the planet Saturn's largest moon.
Nitrogen is present in all living organisms, as part of the molecular structures of proteins, nucleic acids, and other important substances. It is a large component of animal waste, usually in the form of urea, uric acid, and their derivatives.
Nitrogen (Latin nitrum, Greek Nitron meaning "native soda"; genes meaning "forming") is formally considered to have been discovered in 1772 by the chemist Daniel Rutherford, who knew that there was a fraction of air that did not support combustion. He called it "noxious air" or "fixed air." 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 (unreactive) enough that Antoine Lavoisier referred to it as "azote," from the Greek word αζωτος meaning "lifeless." It was the principle component of air in which animals had suffocated and flames had burned to extinction. This term became the French word for nitrogen and later spread to many other languages.
The alchemists of the Middle Ages experimented with various compounds of nitrogen. For instance, they knew nitric acid as aqua fortis (strong water). The mixture of nitric acid and hydrochloric acid was called aqua regia (royal water), celebrated for its ability to dissolve gold (the "king" of metals). In the earliest industrial and agricultural applications of nitrogen compounds, saltpeter (sodium nitrate or potassium nitrate) was used in gunpowder, much later as fertilizer, and later still as a chemical feedstock.
Nitrogen is a chemical element in the periodic table, situated at the head of group 15 (former group 5A), just above phosphorus. In addition, it lies in period 2, flanked by carbon and oxygen. Classified as a nonmetal, it has an electronegativity of 3.0. Each atom of nitrogen has five electrons in its outer shell, and it forms three covalent bonds in most compounds.
Nitrogen gas consists of diatomic molecules, each of which has the chemical formula N2. The two nitrogen atoms in each molecule are attached to each other by a strong, triple covalent bond. For this reason, nitrogen gas is extremely stable and inert.
The gas condenses to the liquid form at 77 Kelvin (K) at atmospheric pressure and freezes at 63 K. Liquid nitrogen is a common cryogen (an extremely low-temperature refrigerant) that can cause instant frostbite on direct contact with living tissue.
There are two stable isotopes of nitrogen: 14N and 15N. By far the most common is 14N (99.634 percent), which is thought to be produced in stars by a set of nuclear fusion reactions called the "carbon-nitrogen-oxygen cycle" (CNO cycle). Of the 10 isotopes produced synthetically, 13N has a half life of nine minutes, and the remaining isotopes have half lives on the order of seconds or less. In the Earth's atmosphere, 0.73 percent of molecular nitrogen consists of 14N15N, and almost all the rest is 14N2.
Nitrogen is an essential element in the molecules of amino acids, proteins, nucleic acids, and other substances vital to life. Specific bacteria (such as those of the genus Rhizobium) possess certain enzymes ("nitrogenases") that can fix atmospheric nitrogen (see nitrogen fixation) into a form (ammonium ion) that is chemically useful for higher organisms. This process requires a large amount of energy and anoxic (oxygen-free) conditions. Usually, these bacteria exist in a symbiotic relationship in the root nodules of leguminous plants such as clover or the soya bean plant. Nitrogen-fixing bacteria can also be symbiotic with other plant species, such as alders, lichens, casuarina, myrica, liverwort, and gunnera.
As part of the symbiotic relationship, the plant converts the ammonium ions to nitrogen oxides and amino acids to form proteins and other biologically useful molecules, such as alkaloids. In return, the plant secretes sugars that the bacteria can use.
Some plants can assimilate nitrogen directly in the form of nitrates, which may be present in the soil from natural mineral deposits, artificial fertilizers, animal waste, or organic decay (as the product of bacteria that are not specifically associated with the plant). Nitrates absorbed in this fashion are converted to nitrites by the enzyme nitrate reductase, and then converted to ammonia by another enzyme called nitrite reductase.
Nitrogen compounds are basic building blocks in animal biology. Animals use nitrogen-containing amino acids from plant sources as starting materials for all nitrogen-compound animal biochemistry, including the manufacture of proteins and nucleic acids. Many saltwater fish manufacture large amounts of trimethylamine oxide to protect them from the high osmotic effects of their environment. In animals, nitric oxide (NO) is derived from an amino acid and serves as an important regulatory molecule for circulation.
Animal metabolism of nitrogen in proteins generally results in the excretion of urea, while animal metabolism of nucleic acids results in the excretion of urea and uric acid. The characteristic odor of animal flesh decay is caused by nitrogen-containing long-chain amines, such as putrescene and cadaverine. The decay of organisms and their waste products may produce small amounts of nitrate, but most decay processes eventually return nitrogen content to the atmosphere, as molecular nitrogen.
Nitrogen is produced industrially in large quantities by a method known as the fractional distillation of liquefied air. Isolated in the liquid form, this nitrogen is often referred to by the quasi-formula LN2, but it is more accurately written as N2(l). Technologies that isolate nitrogen from gaseous air include methods known as pressure swing adsorption and membrane separation. In addition, commercial nitrogen is often obtained as a byproduct of air processing during the industrial concentration of oxygen for steelmaking and other purposes. The formula for gaseous nitrogen is N2(g).
Nitrogen gas has a wide variety of applications, including serving as an inert replacement for air where oxidation is undesirable. Some examples are as follows.
When appropriately insulated from ambient heat, liquid nitrogen serves as a compact, readily transportable source of nitrogen gas without pressurization. Furthermore, its ability to maintain temperatures far below the freezing point of water as it boils at (77 K, -196 °C or -320 °F) makes it extremely useful in a wide range of applications as an open-cycle refrigerant. Some uses of liquid nitrogen include:
Nitrogen forms part of various inorganic compounds, some of which are noted below.
Ammonia: The main neutral hydride of nitrogen is ammonia (NH3), although hydrazine (N2H4) is also common. Ammonia is a chemical base— more basic than water by 6 orders of magnitude. In solution, ammonia combines with protons (H+ ions) to form ammonium ions (NH4+). Liquid ammonia (boiling point 240 K) is "amphiprotic"—that is, it can behave as an acid as well as a base. As an acid, it donates a proton (H+) to another molecule to form the amide ion (NH2−); as a base, it receives a proton (H+) from another molecule to form the ammonium ion (NH4+).
Amides: An inorganic amide is a compound in which a metal cation is combined with the amide anion (NH2-) mentioned above. An example is sodium amide (NaNH2). An inorganic amide is an extremely strong base and decomposes in water. Note that inorganic amide salts are distinctly different from organic amide compounds mentioned below.
Nitrides: In the molecule of a nitride compound, a nitrogen atom is attached to an atom of a more electropositive element. Some nitrides, such as lithium nitride (Li3N), are salt-like, in which the nitrogen exists as an ion with three negative charges (N3−). The salt-like nitrides are strong bases and readily decompose in water. Other nitrides, such as boron nitride (BN), are inert.
Azides: An inorganic azide is a salt in which a metal cation is combined with one or more azide anions (N3−). Each azide anion has a linear structure, N−=N+=N−, with a net charge of -1. Sodium azide is used in airbags, but the azide anion is toxic. Organic azides are noted below.
Oxides and Oxoacids: Nitrogen forms a variety of oxides. The most prominent ones are nitrogen monoxide (NO) (known more commonly as nitric oxide in biology) and nitrogen dioxide (NO2). Both types of molecules contain an unpaired electron. The latter shows some tendency to form dimers and is a significant component of smog. In addition, nitrogen forms dinitrogen monoxide (or nitrous oxide) (N2O), which is also known as laughing gas. The more standard oxides, dinitrogen trioxide (N2O3) and dinitrogen pentoxide (N2O5), are fairly unstable and explosive. The corresponding acids (oxoacids, or oxygen-containing acids) are nitrous acid (HNO2) and nitric acid (HNO3), and the corresponding salts are called nitrites and nitrates.
Nitrogen is also an important element in many organic compounds, in which it is directly bound to one or more carbon atoms. Some examples are given below.
The ability to combine or "fix" molecular nitrogen is a key feature of modern industrial chemistry, where nitrogen and hydrogen are made to react to form ammonia, by what is called the Haber process. Ammonia, in turn, can be used directly as a fertilizer and in the synthesis of nitrated fertilizers, or it can be used as a precursor of many other important materials, largely through the production of nitric acid by the Ostwald process.
The salts of nitric acid include important compounds such as potassium nitrate (or saltpeter, important historically for its use in gunpowder) and ammonium nitrate, an important fertilizer and explosive. 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 its derivatives find use as rocket fuels. For all 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). When nitrates burn or explode, the formation of the powerful triple bond in the resultant (N2) generates most of the energy of the reaction.
In medicine, nitrogen is a constituent of molecules in every major drug class. Nitrous oxide (N20) was discovered early in the nineteenth 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, such as morphine, are derived from plant alkaloids. Many alkaloids are known to have pharmacological effects, and some of them may play a role in the plants' natural defenses against predation. Organic nitro drugs, such as nitroglycerin and nitroprusside, help reduce blood pressure by widening blood vessels.
Rapid release of nitrogen gas into an enclosed space can displace oxygen, thus representing an asphyxiation hazard. Shortly before the launch of the first Space Shuttle mission in 1981, two technicians were killed in an area in the Shuttle's Mobile Launch Platform that was pressurized with pure nitrogen as a precaution against fire.
When breathed at high partial pressures (higher than 3 atmospheres, encountered at diving depths below about 30 meters) nitrogen begins to act as an anesthetic agent. As such, it can cause nitrogen narcosis—a temporary, semi-anesthetized condition of mental impairment similar to that caused by nitrous oxide.
Nitrogen also dissolves in the bloodstream. When divers ascend too quickly, or astronauts decompress too quickly from cabin pressure to spacesuit pressure, the rapid decompression 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.
Direct skin contact with liquid nitrogen causes severe frostbite (cryogenic burns) within seconds, depending on the form of liquid nitrogen (liquid versus mist) and surface area of the nitrogen-soaked material.
All links retrieved August 23, 2014.
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