|Boiling Point: 77.5°K, -195.65°C,
Melting Point: 63.29°K, -209.86°C, -345.75°F
Electrons Energy Level: 2, 5
Isotopes: 11 + 2 Stable
Heat of Vaporization: 2.7928 kJ/mol
Heat of Fusion: 0.3604 kJ/mol
Density: 1.2506 g/L @ 273°K & 1atm
Specific Heat: 1.04 J/g°K
Atomic Radius: 0.75Å
Ionic Radius: 0.13Å
Electronegativity: 3.04 (Pauling); 3.07 (Allrod Rochow)
nitrogenium, where nitrum (from Greek nitron) means "native
soda" and genes 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 Canvendish, 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 azote, from a Greek word 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. This term has
become the French word for "nitrogen" and later spread out to many other
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 industrial and agricultural applications of nitrogen compounds involved uses in the form of saltpeter (Sodium or Potassium Nitrate), notably in gunpowder, and much later, as fertilizer, and later still, as a chemical feedstock.
Nitrogen forms a surprising number of compounds with oxygen, exhibiting a wide range of oxidation states. Nitrogen is a nonmetal, with an electronegativity of 3.0. 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.
Molecular nitrogen condenses at 77oK at atmospheric pressure and freezes at 60oK. Liquid nitrogen, a fluid resembling water, but with 88% of the density, is a common cryogen.
Elemental nitrogen is a colorless, odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78.1% by volume of Earth's atmosphere. Nitrogen is a constituent element of all living tissues and Amino Acids. Many industrially important compounds, such as Ammonia, Nitric Acid, and Cyanides, contain nitrogen.
Nitrogen is the largest single component of the Earth's atmosphere (78.082% by volume, 75.3% by weight).
14N is created as part of the fusion processes in stars, and is estimated to be the 7th most abundant chemical element (by mass) in our universe.
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 is a major constituent of Titan's thick atmosphere, and occurs in trace amounts of other planetary atmospheres.
Nitrogen is present in all living tissues as Proteins, Nucleic Acids and other molecules. It is a large component of animal waste (for example, Guano), usually in the form of Urea, Uric Acid, and compounds of these nitrogenous products.
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 auroa 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 an essential part of Amino Acids and Nucleic Acids, both of which are essential to all life.
Molecular Nitrogen in the atmosphere cannot be used directly by either plants or animals, and needs to be converted to other compounds, or "fixed," in order to be used by life. Precipitation often contains substantial quantities of Ammonium, NH4+, and Nitrate, NO3-, both thought to be a result of Nitrogen fixation by lightning and other atmospheric electric phenomena. However, because Ammonium is preferentially retained by the forest canopy relative to atmospheric Nitrate, most of the fixed Nitrogen that reaches the soil surface under trees is in the form of nitrate. Soil Nitrate is preferentially assimilated by tree roots relative to soil Ammonium.
Specific bacteria (e.g. Rhizobium trifolium) possess Nitrogenase enzymes which can fix atmospheric Nitrogen into a form (Ammonium Ion, NH4+) which is chemically useful to higher organisms. This process requires a large amount of energy and anoxic conditions. Such bacteria may be free in the soil (e.g. azotobacter) but normally exist in a symbiotic relationship in the root nodules of leguminous plants (e.g. clover or the soya bean plant). Nitrogen fixating bacteria can be symbiotic with a number of unrelated plant species. Common examples are legumes, alders, lichens, casuarina, myrica, liverwort, and gunnera.
As part of the symbiotic relationship, the plant subsequently converts the Ammonium ion to Nitrogen Oxides and Amino Acids to form proteins and other biologically useful molecules, such as alkaloids. In return for the usable (fixed) Nitrogen, the plant secretes sugars to the symbiotic bacteria.
Some 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, NH3, 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. Some plant-feeding insects are so dependent on Nitrogen in their diet, that varying the amount of Nitrogen fertilizer applied to a plant can affect the birth rate of the insects feeding on it.
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 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). In animals, the free radical molecule Nitric Oxide (NO), which is 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 Nitrogen-containing long-chain 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.
Nitrogen gas is acquired for industrial purposes by the fractional distillation of liquid air, or by mechanical means using gaseous air. Commercial Nitrogen is often a byproduct of air-processing for industrial concentration of Oxygen for steelmaking and other purposes.
Molecular Nitrogen (Gas & Liquid)
A Computer Rendering of the Nitrogen Molecule N2.
Nitrogen gas has a wide variety of applications, including serving as an inert replacement for air where oxidation is undesirable;
Nitrogen molecules are less likely to escape from the inside of a tire compared to the traditional air mixture used. Air consists mostly of Nitrogen and Oxygen. Nitrogen molecules have a larger effective diameter than Oxygen molecules and therefore diffuse through porous substances more slowly
A further example of its versatility is its use as a preferred alternative to Carbon Dioxide, CO2, to pressurize kegs of some beers, particularly thicker stouts and Scottish and English ales, due to the smaller bubbles it produces, which make the dispensed beers 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.
Molecular Nitrogen, a diatomic gas, is apt to dimerize into a linear four Nitrogen long polymer. This is an important phenomenon for understanding high voltage Nitrogen dielectric switches because the process of polymerization can continue to lengthen the molecule to still longer lengths in the presence of an intense electric field. A Nitrogen polymer fog is thereby created. The second virial coefficient of nitrogen also shows this effect as the compressibility of Nitrogen Gas is changed by the dimerization process at moderate and low temperatures.
Liquid Nitrogen (liquid density at the triple point is 0.807 g/mL) is produced industrially in large quantities by fractional distillation of liquid air and is often referred to by the abbreviation, LN2. It is a cryogenic fluid which is potentially capable of causing instant frosbite on contact with living tissue. When appropriately insulated from ambient heat, liquid nitrogen serves as a compact and readily transported source of Nitrogen Gas without pressurization. Further, its ability to maintain temperatures far below the freezing point of water (it boils at 77°K, which equals -196°C or -320°F) makes it extremely useful in a wide range of applications as an open-cycle refrigerant, including;
Nitrogen Compounds in Industry
Small amounts of Nitrogen (for laboratory study) can be prepared by heating solid Ammonium Chloride (NH4Cl) with solid Ammonium Nitrite (NH4NO2). Sodium Chloride and water are by-products.
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 (NH4+). Liquid ammonia (b.p. 240°K) is amphiprotic (displaying either Bronsted-Lowry acidic or basic character) and forms ammonium and the less common Amide Ions (NH2-); 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 quarternary 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 (N3-), 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 Dinitrogen Monoxide (N2O), also known as laughing gas. This is one of a variety of Oxides, the most prominent of which are Nitrogen Monoxide (NO) (known more commonly as Nitric Oxide in biology), a natural free radical molecule used by the body as a signal for short-term control of smooth muscle in the circulation. Another notable nitrogen oxide compound (a family often abbreviated NOx) is the reddish and poisonous Nitrogen Dioxide (NO2), which also 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 more standard oxides, Dinitrogen Trioxide (N2O3) and Dinitrogen Pentoxide (N2O5), are actually fairly unstable and explosive-- a tendency which is driven by the stability of N2 as a product. The corresponding acids are Nitrous (HNO2) and Nitric Acid (HNO3), with the corresponding salts called Nitrites and Nitrates. Nitric Acid is one of the few acids stronger than Hydronium, and is a fairly strong oxidizing agent.
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 Compounds of Notable Economic Importance
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 or destroyed. In nature, nitrogen is slowly converted into biologically (and industrially) useful compounds by some living organisms, notably certain bacteria (i.e. Nitrogen fixing bacteria). Molecular nitrogen is also 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 historically important as stores of chemical energy. They include important compounds such as Potassium Nitrate, KNO3, (or saltpeter, important historically for its use in gunpowder) and Ammonium Nitrate, NH4NO3, an important fertilizer and explosive. 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. 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 which results, 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 defences 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 mimicing the action of Nitric Oxide.
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 and the remaining is 15N. Of the ten 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. 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.
The molecular Nitrogen in Earth's atmosphere is 0.73% comprised of the isotopologue 14N15N and almost all the rest is 14N2.
|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 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 Launch 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 breathed at high partial pressures (more than about 3 atmospheres, encountered at depths below about 30 m in scuba diving) 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, and 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.
Direct skin contact with liquid nitrogen causes severe frostbite (cryogenic burns) within moments to seconds, though not instantly on contact, depending on form of liquid nitrogen (liquid vs. mist) and surface area of the nitrogen-soaked material (soaked clothing or cotton causing more rapid damage than a spill of direct liquid to skin, which for a few seconds is protected by the Leidenfrost Effect).
Atomic Radius (Å): 0.75Å
Electrochemical Equivalents: 0.10452 g/amp-hr
Atomic Mass Average: 14.00674
(L. nitrum, Gr. nitron, native soda; genes, forming) Discovered by Daniel Rutherford in 1772, but Scheele, Cavendish, Priestley, and others at about the same time studied "burnt or dephlogisticated air," as air without oxygen was then called. Nitrogen makes up 78% of the air, by volume. The atmosphere of Mars, by comparison, is 2.6% nitrogen. The estimated amount of this element in our atmosphere is more than 4000 trillion (?). From this inexhaustible source it can be obtained by liquifaction and fractional distillation. Nitrogen molecules give the orange-red, blue-green, blue-violet, and deep violet shades to the aurora. The element is so inert that Lavoisier named it azote, meaning without life, yet its compounds are so active as to be most important in foods, poisons, fertilizers, and explosives. Nitrogen can also be easily prepared by heating a water solution of ammonium nitrite. Nitrogen, as a gas, is colorless, odorless, and a generally inert element. As a liquid it is also colorless and odorless, and is similar in appearance to water. Two allotropic forms of solid nitrogen exist, with the transition from the alpha to the beta form taking place at -237oC. When nitrogen is heated, it combines directly with magnesium, lithium, or calcium; when mixed with oxygen and subjected to electric sparks, it forms first nitric acid (NO) and then the dioxide (NO2); when heated under pressure with a catalyst with hydrogen, ammonia is formed (Haber process). The ammonia that is formed is of the utmost importance as it is used in fertilizers, and can be oxidized to nitric acid (Ostwald process). The ammonia industry is the largest consumer of nitrogen. Large amounts of gas are also used by the electronics industry, which uses the gas as a blanketing medium during production of such componenets as transistors, diodes, etc. Large quantities of nitrogen are used in annealing stainless steel and other steel mill products. The drug industry also uses large quantities. Nitrogen is used as a refrigerant both for the immersion freezing of food products and for transportation of foods. Liquid nitrogen is also used in missile work as a purge for components, insulators for space chambers, etc., and by the oil industry to build up great pressures in wells to force crude oil upward. Sodium and potassium nitrates are formed by the decomposition of organic matter with compounds of the metals present. In certain dry areas of the world these saltpeters are found in quantity. Ammonia, nitric acid, the nitrates, the five oxides, TNT, the cyanides, etc. are but a few of the important compounds. Nitrogen gas prices vary from 2 cents to $2.75 per 100 ft3 depending on purity, etc. Production of elemental nitrogen in the U.S. is more than 9 million short tons per year.
Source: CRC Handbook of Chemistry and Physics, 1913-1995. David R. Lide, Editor in Chief. Author: C.R. Hammond