N2, a colorless, tasteless, odorless gas which makes up 78.1% of the atmosphere. Atmospheric nitrogen is converted by nitrogen fixation and nitrification into compounds used by plants and animals. In the far upper atmosphere, N2 is broken down when large numbers of energetic secondary electrons are produced and available to react with the N2. This leads to the eventual production of NO in that part of the atmosphere and is not--by definition--anthropogenic in nature.

(Crutzen, Paul J. and T.E. Graedel, Atmospheric Change. W.H. Freeman and Company, New York, 147.) (Webster's New World Dictionary. Prentice Hall, New York, 918.)

Source: Atmospheric Chemistry Glossary

Atomic Number: 7
Atomic Symbol: N
Atomic Weight: 14.00674
Electron Configuration: [He]2s22p3

History
(L. nitrum, Gr. Nitron, native soda; genes, forming) Nitrogen was discovered by chemist and physician Daniel Rutherford in 1772. He removed oxygen and carbon dioxide from air and showed that the residual gas would not support combustion or living organisms. At the same time there were other noted scientists working on the problem of nitrogen. These included Scheele, Cavendish, Priestley, and others. They called it "burnt or dephlogisticated air," which meant air without oxygen.

Sources
Nitrogen gas (N2) makes up 78.1% of the Earth’s air, by volume. The atmosphere of Mars, by comparison, is only 2.6% nitrogen. From an exhaustible source in our atmosphere, nitrogen gas can be obtained by liquefaction and fractional distillation. Nitrogen is found in all living systems as part of the makeup of biological compounds.

The Element
The French chemist Antoine Laurent Lavoisier named nitrogen azote, meaning without life. However, nitrogen compounds are found in foods, fertilizers, poisons, and explosives. Nitrogen, as a gas is colorless, odorless, and generally considered an inert element. As a liquid (boiling point = -195.8° C), it is also colorless and odorless, and is similar in appearance to water. Nitrogen gas can be prepared by heating a water solution of ammonium nitrite (NH4NO3).

Nitrogen Compounds and Nitrogen in Nature
Sodium nitrate (NaNO3) and potassium nitrate (KNO3) are formed by the decomposition of organic matter with compounds of these metals present. In certain dry areas of the world these saltpeters are found in quantity and are used as fertilizers. Other inorganic nitrogen compounds are nitric acid (HNO3), ammonia (NH3), the oxides (NO, NO2, N2O4, N2O), cyanides (CN-), etc.

The nitrogen cycle is one of the most important processes in nature for living organisms. Although nitrogen gas is relatively inert, bacteria in the soil are capable of "fixing" the nitrogen into a usable form (as a fertilizer) for plants. In other words, Nature has provided a method to produce nitrogen for plants to grow. Animals eat the plant material where the nitrogen has been incorporated into their system, primarily as protein. The cycle is completed when other bacterial convert the waste nitrogen compounds back to nitrogen gas. Nitrogen has become crucial to life being a component of all proteins.

Ammonia
Ammonia (NH3) is the most important commercial compound of nitrogen. It is produced by the Haber Process. Natural gas (methane, CH4) is reacted with steam to produce carbon dioxide and hydrogen gas (H2) in a two step process. Hydrogen gas and nitrogen gas are then reacted in the Haber Process to produce ammonia. This colorless gas with a pungent odor is easily liquefied. In fact, the liquid is used as a nitrogen fertilizer. Ammonia is also used in the production of urea, NH2CONH2, which is used as a fertilizer, in the plastic industry, and in the livestock industry as a feed supplement. Ammonia is often the starting compound for many other nitrogen compounds.

Sources: CRC Handbook of Chemistry and Physics and the American Chemical Society.












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