Nitrogen

Nitrogen is the chemical element in the periodic table that has the symbol N and atomic number 7. Commonly a colorless, odorless, tasteless and mostly inert diatomic non-metal gas, nitrogen constitutes 78 percent of Earth's atmosphere and is a constituent of all living tissues. Nitrogen forms many important compounds such as amino acids, ammonia, nitric acid, and cyanides.

1 Notable characteristics
2 Applications
- 2.1 Nitrogen Compounds
- 2.2 Molecular nitrogen (gas and liquid)
3 History
4 Occurrence
5 Compounds
6 Biological role
7 Isotopes
8 Precautions
9 See also
10 References
11 External links

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Notable characteristics

Nitrogen is a non-metal, with an electronegativity of 3.0. It has five electrons in its outer shell, so is trivalent in most compounds. Pure nitrogen is an unreactive colorless diatomic gas at room temperature, and comprises about 78.08% of the Earth's atmosphere. It condenses at 77 K and freezes at 63 K. Liquid nitrogen is a common cryogen.

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Applications

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Nitrogen Compounds

Molecular nitrogen in the atmosphere is relatively non-reactive, but in nature it is slowly converted into biologically (and industrially) useful compounds by some living organisms, notably certain bacteria (see Biological role below). The ability to combine or fix nitrogen is a key feature of modern industrial chemisty, where nitrogen (along with natural gas) is converted into ammonia (via the Haber process). Ammonia, in turn, can be used directly (primarily as a fertilizer), or as a precursor of many other important materials including explosives, largely via the production of nitric acid by the Ostwald process.

The salts of nitric acid include important compounds like potassium nitrate (or saltpeter, important historically for its use in gunpowder) and ammonium nitrate, an important fertilizer. Various other nitrated organic compounds, such as nitroglycerin and trinitrotoluene, are used as explosives. Nitric acid is used as an oxidizer in liquid fueled rockets. Hydrazine and hydrazine derivatives find use as rocket fuels.

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Molecular nitrogen (gas and liquid)

Nitrogen gas is readily produced by allowing liquid nitrogen (see below) to warm up and evaporate. It has a wide variety of applications, including serving as a more inert replacement for air where oxidation is undesireable;

to preserve the freshness of packaged or bulk foods (by delaying rancidity and other forms of oxidative damage)
on top of liquid explosives for safety

It is also used in:

the production of electronic parts such as transistors, diodes, and integrated circuits
the manufacture of stainless steel
filling automotive tires due to its inertness and lack of moisture or oxidative qualities, as opposed to air.

A further example of its versitility is its use (as a preferred alternative to carbon dioxide) 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 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. A very popular example of this is Guinness Draught.

Liquid nitrogen is produced industrially in large quantities by distillation from liquid air and is often referred to by the quasi-formula LN₂. It is a cryogenic (extremely cold) fluid which can cause instant frostbite on direct contact with living tissue. When appropriately insulated from ambient heat it serves as a compact and readily transported source of nitrogen gas without pressurization. Further, its ability to maintain an unearthly temperature as it evaporates (77 K, -196 °C or -320 °F) makes it extremely useful in a wide range of applications as an open-cycle refrigerant, including;

the immersion freezing and transportation of food products
the preservation of bodies, reproductive cells (sperm and egg), and biological samples and materials
in the study of cryogenics
for demonstrations in science education
in dermatology for removing unsightly or potentially malignant skin lesions,e.g., warts, actinic keratosis, etc.

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History

Nitrogen (Latin nitrum, Greek Nitron meaning "native soda", "genes", "forming") is formally considered to have been discovered by Daniel Rutherford in 1772, who called it noxious air or fixed air. That there was a fraction of air that did not support combustion was well known to the late 18th century chemist. Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as azote, which stands for without life; this term has become the French word for "nitrogen" and later spread out to many other languages.

Compounds of nitrogen were known in the Middle Ages. The alchemists knew nitric acid as aqua fortis. The mixture of nitric and hydrochloric acids was known as aqua regia, celebrated for its ability to dissolve gold. The earliest industrial and agricultural applications of nitrogen compounds used it in the form of saltpeter (sodium- or potassium nitrate), notably in gunpowder, and much later, as fertilizer, and later still, as a chemical feedstock.

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Occurrence

Nitrogen is the largest single component of the Earth's atmosphere (78.084% by volume, 75.5% by weight) and is acquired for industrial purposes by the fractional distillation of liquid air or by mechanical means of gaseous air (i.e. pressurised reverse osmosis membrane or PSA (Pressure Swing Adsorption). Compounds that contain this element have been observed in outer space. Nitrogen-14 is created as part of the fusion processes in stars. Nitrogen 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 has been known to occur in Titan's atmosphere for some time, and has now been detected in interstellar space by David Knauth and coworkers using the Far Ultraviolet Spectroscopic Explorer.

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Compounds

The main hydride of nitrogen is ammonia (NH₃) although hydrazine (N₂H₄) is also well known. Ammonia is somewhat more basic than water, and in solution forms ammonium ions (NH₄⁺). Liquid ammonia is in fact slightly amphiprotic and forms ammonium and amide ions (NH₂^-); both amides and nitride (N^3-) salts are known, but decompose in water. Singly and doubly substituted compounds of ammonia are called amines. Larger chains, rings and structures of nitrogen hydrides are also known but virtually unstable.

Other classes of nitrogen anions are azides (N₃^-), which are linear and isoelectronic to carbon dioxide. Another molecule of the same structure is dinitrogen monoxide (N₂O), or laughing gas. This is one of a variety of oxides, the most prominent of which are nitrogen monoxide (NO) and nitrogen dioxide (NO₂), which both contain an unpaired electron. The latter shows some tendency to dimerize and is an important component of smog.

The more standard oxides, dinitrogen trioxide (N₂O₃) and dinitrogen pentoxide (N₂O₅), are actually fairly unstable and explosive. 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.

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Biological role

Nitrogen is an essential part of amino and nucleic acids which makes nitrogen vital to all life. Legumes like the soybean plant, can recover nitrogen directly from the atmosphere because their roots have nodules harboring microbes that do the actual conversion to ammonia in a process known as nitrogen fixation. The legume subsequently converts ammonia to nitrogen oxides and amino acids to form proteins.

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Isotopes

There are two stable isotopes: N-14 and N-15. By far the most common is N-14 (99.634%), which is produced in the CNO cycle in stars. The rest is N-15. Of the ten isotopes produced synthetically, one 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 almost always result in N-15 enrichment of the substrate and depletion of the product. Although precipitation often contains subequal quantities of ammonium and nitrate, because ammonium is preferentially retained by the canopy relative to atmospheric nitrate, most of the atmospheric nitrogen that reaches the soil surface is in the form of nitrate. Soil nitrate is preferentially assimilated by tree roots relative to soil ammonium.

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Precautions

Nitrate fertilizer washoff is a major source of ground water and river pollution. Cyano (-CN) containing compounds form extremely poisonous salts and are deadly to many animals and all mammals.

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