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Nitrogen

2007 Schools Wikipedia Selection. Related subjects: Chemical elements


                 7               carbon ← nitrogen → oxygen
                 -
                ↑
                N
                ↓
                P

                                  Periodic Table - Extended Periodic Table

                                                                   General
                                       Name, Symbol, Number nitrogen, N, 7
                                                 Chemical series nonmetals
                                             Group, Period, Block 15, 2, p
                                                      Appearance colorless
                                             Atomic mass 14.0067 (2) g/mol
                                     Electron configuration 1s^2 2s^2 2p^3
                                                  Electrons per shell 2, 5
                                                       Physical properties
                                                                 Phase gas
                                              Density (0 °C, 101.325 kPa)
                                                                 1.251 g/L
                                                    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 (N[2]) 0.720 kJ·mol^−1
                             Heat of vaporization (N[2]) 5.57 kJ·mol^−1
                                             Heat capacity (25 °C) (N[2])
                                                29.124 J·mol^−1·K^−1

   CAPTION: Vapor pressure

                                            P/Pa  1  10 100 1 k 10 k 100 k
                                           at T/K 37 41 46  53   62   77

                                                         Atomic properties
                                               Crystal structure hexagonal
                                             Oxidation states ±3, 5, 4, 2
                                                   (strongly acidic oxide)
                                    Electronegativity 3.04 (Pauling scale)
                                                       Ionization energies
                                          ( more) 1st: 1402.3 kJ·mol^−1
                                                    2nd: 2856 kJ·mol^−1
                                                  3rd: 4578.1 kJ·mol^−1
                                                       Atomic radius 65 pm
                                               Atomic radius (calc.) 56 pm
                                                     Covalent radius 75 pm
                                               Van der Waals radius 155 pm
                                                             Miscellaneous
                                                 Magnetic ordering no data
                     Thermal conductivity (300 K) 25.83 mW·m^−1·K^−1
                                      Speed of sound (gas, 27 °C) 353 m/s
                                             CAS registry number 7727-37-9
                                                         Selected isotopes

                 CAPTION: Main article: Isotopes of nitrogen

                                  iso    NA    half-life DM DE ( MeV)  DP
                                  ^13N syn     9.965 m   ε  2.220     ^13C
                                  ^14N 99.634% N is stable with 7 neutrons
                                  ^15N 0.366%  N is stable with 8 neutrons

                                                                References

   Nitrogen ( IPA: /ˈnʌɪtrə(ʊ)dʒən/) is a chemical element which has the
   symbol N and atomic number 7 in the periodic table. Elemental nitrogen
   is a colorless, odorless, tasteless and mostly inert diatomic gas at
   standard conditions, constituting 78.08% percent 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.

Notable characteristics of elemental nitrogen

   Nitrogen is a non-metal, 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 (N[2]) is one of the
   strongest in nature. The resulting difficulty of converting (N[2]) into
   other compounds, and the relative ease (and associated high energy
   release) of converting nitrogen compounds into elemental N[2], have
   dominated the role of nitrogen in both nature and human economic
   activities.

   Molecular nitrogen condenses at 77 K at atmospheric pressure and
   freezes at 63 K. Liquid nitrogen, a fluid resembling water, but with
   81% of the density, is a common cryogen.

Occurrence

   Nitrogen is the largest single component of the Earth's atmosphere
   (78.084% by volume, 75.5% by weight).

   ^14Nitrogen is created as part of the 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 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.

Isotopes

   See also Isotopes of nitrogen

   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 almost always 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 ^14N^15N and almost all the rest is ^14N[2].

Electromagnetic spectrum

   Molecular nitrogen (^14N[2]) 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 far 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 aurora 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).

History

   Nitrogen (Latin nitrogenium, where nitrum (from Greek nitron) means
   "native soda" (see niter), 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 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, from the Greek word
   αζωτος 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 languages.

   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
   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.

Biological role

   See also nitrogen cycle

   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 and nitrate, 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 (see nitrogen fixation) into
   a form (ammonium ion) 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 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 (Jahn et al. 2005).

   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 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 odour in unfresh saltwater fish: PMID 15186102). 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 odour 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.

Modern applications

   Nitrogen gas is acquired for industrial purposes by the fractional
   distillation of liquid air, or by mechanical means using gaseous air
   (i.e. pressurised reverse osmosis membrane or pressure swing
   adsorption). Commercial nitrogen is often a byproduct of air-processing
   for industrial concentration of oxygen for steelmaking and other
   purposes.

Molecular nitrogen (gas and liquid)

   A computer rendering of the Nitrogen Molecule, N2.
   Enlarge
   A computer rendering of the Nitrogen Molecule, N[2].

   Nitrogen gas has a wide variety of applications, including serving as a
   more inert replacement for air where oxidation is undesirable;
     * 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
     * The production of electronic parts such as transistors, diodes, and
       integrated circuits
     * dried and pressurized, as a dielectric gas for high voltage
       equipment
     * The manufacture of stainless steel
     * Use in military aircraft fuel systems to reduce fire hazard, see
       inerting system
     * Filling automotive and aircraft tires due to its inertness and lack
       of moisture or oxidative qualities, as opposed to air, though this
       is not necessary for consumer automobiles.

   Contrary to some claims that nitrogen will diffuse more rapidly through
   rubber tires than air (and oxygen), 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 are larger than oxygen molecules and therefore, all else
   being equal, larger molecules diffuse through porous substances slower
   than smaller molecules.

   A further example of its versatility 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.
   Liquid nitrogen may be used to prepare "home-made" ice cream, as these
   students are doing.
   Enlarge
   Liquid nitrogen may be used to prepare "home-made" ice cream, as these
   students are doing.

   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 quasi-formula LN[2] (but is
   more accurately written N[2](l) ). It is a cryogenic fluid which is
   potentially capable of causing instant frostbite on contact with living
   tissue (see precautions). 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;
     * the immersion freezing and transportation of food products
     * the cryopreservation of blood, reproductive cells ( sperm and egg),
       and other biological samples and materials (see image at right)
     * the cryonic preservation of humans and pets in the unproven hope of
       future reanimation.
     * in the study of cryogenics
     * for demonstrations in science education
     * as a coolant for highly sensitive sensors and low-noise amplifiers
     * in dermatology for removing unsightly or potentially malignant skin
       lesions such as warts and actinic keratosis
     * as a cooling supplement for overclocking a central processing unit,
       a graphics processing unit, or another type of computer hardware
     * as a cooling medium during machining of high strength materials.
     * as the working fluid in a binary engine
     * as a means of final disposition of the dead, known as promession.

   A tank of liquid nitrogen, used to supply a cryogenic freezer (for
   storing laboratory samples at a temperature of about -150 Celsius).
   Enlarge
   A tank of liquid nitrogen, used to supply a cryogenic freezer (for
   storing laboratory samples at a temperature of about -150 Celsius).

Nitrogen compounds in industry

Simple compounds

   The main neutral hydride of nitrogen is ammonia (NH[3]), although
   hydrazine (N[2]H[4]) is also commonly used. Ammonia is more basic than
   water by 6 orders of magnitude. In solution ammonia forms the ammonium
   ion (NH[4]^+). Liquid ammonia (b.p. 240 K) is amphiprotic (displaying
   either Brønsted-Lowry acidic or basic character) and forms ammonium and
   the less common amide ions (NH[2]^-); both amides and nitride (N^3-)
   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 (N[3]^-), 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 (N[2]O), 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 (NO[2]), 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 (N[2]O[3]) and dinitrogen
   pentoxide (N[2]O[5]), are actually fairly unstable and explosive-- a
   tendency which is driven by the stability of N[2] as a product. The
   corresponding acids are nitrous (HNO[2]) and nitric acid (HNO[3]), 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 (N[2]) in the atmosphere is relatively non-reactive
   due to its strong bond, and N[2] 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 - see Biological role above). 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 (or saltpeter, important
   historically for its use in gunpowder) and ammonium nitrate, an
   important fertilizer and explosive (see ANFO). 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 (N[2]) which is
   a product of the compounds' thermal decomposition. When nitrates burn
   or explode, the formation of the powerful triple bond in the N[2] 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 (N[2]0) 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.

Dangers

   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).

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