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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.)
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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