Nitrogen ( ) is a
chemical element that has the symbol
N and
atomic number 7 and
atomic mass 14.00674 u. Elemental nitrogen is a colorless, odorless, tasteless and mostly
inert diatomic gas at
standard conditions, constituting 78% by volume of
Earth's atmosphere.
Many industrially important compounds, such as
ammonia,
nitric acid, organic nitrates (
propellants and
explosives), and
cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in converting the into useful compounds, and releasing large amounts of energy when these compounds burn or decay back into nitrogen gas.
The element nitrogen was discovered by
Daniel Rutherford, a Scottish physician, in 1772. Nitrogen occurs in all living organisms. It is a constituent element of
amino acids and thus of
proteins, and of
nucleic acids (
DNA and
RNA). It resides in the
chemical structure of almost all
neurotransmitters, and is a defining component of
alkaloids, biological molecules produced by many organisms.
History
Nitrogen (
Latin nitrogenium, where
nitrum (from
Greek nitron νιτρον) means "
saltpetre" (see
nitre), and
genes γενης 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 "mephitic air" or
azote, from the
Greek word
(
azotos) 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.
Lavoisier's name for nitrogen is used in many languages (French, Russian, etc.) and still remains in English in the common names of many compounds, such as hydrazine and compounds of the azide ion. 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 military, industrial and
agricultural applications of nitrogen compounds involved uses of
saltpeter (
sodium nitrate or
potassium nitrate), notably in
gunpowder, and much later, as
fertilizer.
Properties
Nitrogen is a
nonmetal, with an
electronegativity of 3.04. It has five
electrons in its
outer shell and is therefore
trivalent in most compounds. The
triple bond in molecular nitrogen () is the strongest in nature. The resulting difficulty of converting into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental , have dominated the role of nitrogen in both nature and human economic activities.
At
atmospheric pressure molecular nitrogen
condenses (
liquifies) at 77
K (−195.8 °
C) and
freezes at 63 K (−210.0 °C) into the beta
hexagonal close-packed crystal
allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the alpha
cubic crystal allotropic form.
Liquid nitrogen, a fluid resembling water in appearance, but with 80.8% of the density (the density of liquid nitrogen at its boiling point is 0.808 g/mL), is a common
cryogen.
Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like and
. Under extremely high pressures (1.1 million
atm) and high temperatures (2000 K), as produced using a
diamond anvil cell, nitrogen polymerizes into the single-bonded cubic gauche crystal structure. This structure is similar to that
diamond, and both have extremely strong
covalent bonds. is nicknamed "nitrogen diamond."
Isotopes
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. Of the ten isotopes produced synthetically,
13N has a
half-life of ten 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.
0.73% of the molecular nitrogen in Earth's atmosphere is comprised of the
isotopologue 14N
15N and almost all the rest is
14N
2.
Radioisotope
16N is the dominant radionuclide in the coolant of
pressurized water reactors during normal operation. It is produced from
16O (in water) via (n,p) reaction. It has a short half-life of about 7.1 s, but during its decay back to
16O produces high-energy gamma radiation (5 to 7 MeV). Because of this, the access to the primary coolant piping must be restricted during reactor power operation.
16N is one of the main means used to immediately detect even small leaks from the primary coolant to the secondary steam cycle.
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 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).
Reactions
[[Image:RuA5N2.png|thumb|right|Structure of [Ru(NH
3)
5(N
2)]
2+.]]
Nitrogen is generally unreactive at standard temperature and pressure. N
2 reacts spontaneously with few
reagents, being resilient to
acids and
bases as well as oxidants and most reductants. When nitrogen reacts spontaneously with a reagent, the net transformation is often called
nitrogen fixation.
Nitrogen reacts with elemental
lithium at
STP.
Lithium burns in an atmosphere of N
2 to give
lithium nitride:
6 Li + N2 → 2 Li3N
Magnesium also burns in nitrogen, forming
magnesium nitride.
3 Mg + N2 → Mg3N2
N
2 forms a variety of
adducts with transition metals. The first example of a
dinitrogen complex is [Ru(NH
3)
5(N
2)]
2+ (see figure at right). Such compounds are now numerous, other examples include IrCl(N
2)(PPh
3)
2, W(N
2)
2(
Ph2CH2CH2PPh2)
2, and [(η
5-C
5Me
4H)
2Zr]
2(
μ2,
η²,η²-N
2). These
complexes illustrate how N
2 might bind to the metal(s) in
nitrogenase and the
catalyst for the
Haber process. A catalytic process to
reduce N
2 to ammonia with the use of a
molybdenum complex in the presence of a proton source was published in 2005.
(see
nitrogen fixation)
The starting point for industrial production of nitrogen compounds is the Haber process, in which nitrogen is fixed by reacting and over an
iron(III) oxide () catalyst at about 500 °C and 200 atmospheres pressure. Biological nitrogen fixation in free-living
cyanobacteria and in the
root nodules of plants also produces ammonia from molecular nitrogen. The reaction, which is the source of the bulk of nitrogen in the
biosphere, is catalyzed by the
nitrogenase enzyme complex which contains Fe and Mo atoms, using energy derived from hydrolysis of
adenosine triphosphate (ATP) into
adenosine diphosphate and
inorganic phosphate (−20.5 kJ/mol).
Occurrence
Nitrogen is the largest single constituent of the
Earth's
atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air). It is created by
fusion processes in
stars, and is estimated to be the 7th most abundant
chemical element by mass in the universe.
nitrogen and nitrogen
compounds have been detected in
interstellar space by astronomers using the
Far Ultraviolet Spectroscopic Explorer. Molecular nitrogen is a major constituent of the
Saturnian moon
Titan's thick atmosphere, and occurs in trace amounts in other planetary atmospheres.
Nitrogen is present in all living organisms, in proteins, nucleic acids and other molecules. It typically makes up around 4% of the dry weight of plant matter, and around 3% of the weight of the human body. It is a large component of animal waste (for example,
guano), usually in the form of
urea,
uric acid,
ammonium compounds and derivatives of these nitrogenous products, which are essential nutrients for all plants that are unable to
fix atmospheric nitrogen.
Nitrogen occurs naturally in a number of minerals, such as
saltpetre (potassium nitrate),
Chile saltpetre (sodium nitrate) and
sal ammoniac (ammonium chloride). Most of these are relatively uncommon, partly because of the minerals' ready solubility in water. See also
Nitrate minerals and
Ammonium minerals.
Compounds
The main neutral
hydride of nitrogen is
ammonia (), although
hydrazine () is also commonly used. Ammonia is more
basic than
water by 6 orders of magnitude. In
solution ammonia forms the
ammonium ion (). Liquid ammonia (boiling point 240 K) is
amphiprotic (displaying either
Brønsted-Lowry acidic or basic character) and forms ammonium and the less common
amide ions (); both amides and
nitride ()
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 quaternary 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. is another polyatomic cation as in hydrazine.
Other classes of nitrogen
anions (negatively charged ions) are the poisonous
azides (), 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
Nitrous oxide (dinitrogen monoxide, ), also known as laughing gas. This is one of a variety of nitrogen
oxides that form a family often abbreviated as
NOx.
Nitric oxide (
nitrogen monoxide, NO), is a natural
free radical used in
signal transduction in both plants and animals, for example in
vasodilation by causing the smooth muscle of blood vessels to relax. The reddish and poisonous
nitrogen dioxide 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 corresponding acids are
nitrous and
nitric acid , with the corresponding salts called
nitrites and
nitrates.
The higher oxides
dinitrogen trioxide ,
dinitrogen tetroxide and
dinitrogen pentoxide , are fairly unstable and explosive, a consequence of the chemical stability of . Nearly every
hypergolic rocket engine uses as the oxidizer; their fuels, various forms of
hydrazine, are also nitrogen compounds. These engines are extensively used on spacecraft such as the
space shuttle and those of the
Apollo Program because their propellants are liquids at room temperature and ignition occurs on contact without an ignition system, allowing many precisely controlled burns. Some launch vehicles, such as the
Titan II and
Ariane 1 through 4 also use hypergolic fuels, although the trend is away from such engines for cost and safety reasons. is an intermediate in the manufacture of nitric acid , one of the few acids stronger than
hydronium and a fairly strong
oxidizing agent.
Nitrogen is notable for the range of explosively unstable compounds that it can produce. Nitrogen triiodide is an extremely sensitive
contact explosive.
Nitrocellulose, produced by nitration of cellulose with nitric acid, is also known as guncotton.
Nitroglycerin, made by nitration of
glycerin, is the dangerously unstable explosive ingredient of
dynamite. The comparatively stable, but more powerful explosive
trinitrotoluene (TNT) is the standard explosive against which the power of nuclear explosions are measured.
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.
Applications
A computer rendering of the nitrogen molecule, N2.
Nitrogen gas is an
industrial gas produced by the fractional
distillation of liquid
air, or by mechanical means using gaseous air (i.e. pressurized 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. When supplied compressed in cylinders it is often referred to as OFN (oxygen-free nitrogen).
Nitrogen gas has a wide variety of applications, including serving as an
inert replacement for
air where
oxidation is undesirable;
Nitrogen is commonly used during sample preparation procedures for chemical analysis. Specifically, it is used as a means of concentrating and reducing the volume of liquid samples. Directing a pressurized stream of nitrogen gas perpendicular to the surface of the liquid allows the solvent to evaporate while leaving the solute(s) and un-evaporated solvent behind.
Nitrogen tanks are also replacing carbon dioxide as the main power source for paintball guns. The downside is that nitrogen must be kept at higher pressure than CO
2, making N
2 tanks heavier and more expensive.
Nitrogenated beer
A further example of its versatility is its use as a preferred alternative to
carbon dioxide to
pressurize kegs of some
beers, particularly
stouts and British
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
Liquid nitrogen is a
cryogenic liquid. At atmospheric pressure, it boils at −195.8 °C. When insulated in proper containers such as
Dewar flasks, it can be transported without much
evaporative loss.
Like
dry ice, the main use of liquid nitrogen is as a
refrigerant. Among other things, it is used in the
cryopreservation of
blood, reproductive cells (
sperm and
egg), and other biological samples and materials. It is used in
cold traps for certain laboratory equipment and to cool
x-ray detectors. It has also been used to cool
central processing units and other devices in computers which are
overclocked, and which produce more heat than during normal operation.
Applications of nitrogen compounds
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 nor destroyed. In nature, nitrogen is converted into biologically (and industrially) useful compounds by lightning, and by some living organisms, notably certain
bacteria (i.e.
nitrogen fixing bacteria – see
Biological role below). Molecular nitrogen is 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 important historically as convenient stores of chemical energy. They include important compounds such as
potassium nitrate (or
saltpeter used 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 and
monopropellants. 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 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
2O) 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 defenses 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 mimicking the action of
nitric oxide.
Biological role
Nitrogen is an essential building block of
amino and
nucleic acids, essential to life on Earth.
Elemental nitrogen in the atmosphere cannot be used directly by either plants or animals, and must be converted to a reduced (or 'fixed') state in order to be useful for higher plants and animals.
Precipitation often contains substantial quantities of
ammonium and
nitrate, thought to result from
nitrogen fixation by
lightning and other atmospheric electric phenomena.
This was first proposed by
Liebig in 1827 and later confirmed.
However, because
ammonium is preferentially retained by the
forest canopy relative to atmospheric nitrate, most fixed nitrogen reaches the
soil surface under trees as nitrate. Soil nitrate is preferentially assimilated by these 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) that is chemically useful to higher organisms. This process requires a large amount of energy and
anoxic conditions. Such bacteria may live freely in soil (e.g.
Azotobacter) but normally exist in a
symbiotic relationship in the
root nodules of
leguminous plants (e.g.
clover,
Trifolium, or
soybean plant,
Glycine max). Nitrogen-fixing bacteria are also symbiotic with a number of unrelated plant species such as alders (
Alnus) spp., lichens (
Casuarina),
Myrica,
liverworts, and
Gunnera.
As part of the symbiotic relationship, the plant converts the 'fixed' ammonium ion to nitrogen oxides and amino acids to form
proteins and other molecules, (e.g.
alkaloids). In return for the 'fixed' nitrogen, the plant secretes sugars to the symbiotic bacteria.
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 as well. 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. Plant-feeding insects are dependent on nitrogen in their diet, such that varying the amount of nitrogen fertilizer applied to a plant can affect the reproduction rate of insects feeding on fertilized plants.
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 odor in not fresh saltwater fish . In animals,
free radical nitric oxide (
NO) (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 the creation of long-chain, nitrogen-containing
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. The circulation of nitrogen from atmosphere to organic compounds and back is referred to as the
nitrogen cycle.
Safety
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 a 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 Launcher 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 inhaled at high
partial pressures (more than about 4 bar, encountered at depths below about 30 m in
scuba diving) nitrogen begins to act as an anesthetic agent. It can cause
nitrogen narcosis, a temporary semi-anesthetized state of mental impairment similar to that caused by
nitrous oxide.
Nitrogen also dissolves in the
bloodstream and body fats. 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, nerves, joints, and other sensitive or vital areas.
Other "inert" gases (those gases other than carbon dioxide and oxygen) cause the same effects from bubbles composed of them, so replacement of nitrogen in
breathing gases may prevent nitrogen narcosis, but does not prevent decompression sickness.
Direct skin contact with
liquid nitrogen will eventually cause severe
frostbite (cryogenic burns). This may happen almost instantly on contact, depending on the form of liquid nitrogen. Bulk liquid nitrogen causes less rapid freezing than a spray of nitrogen mist (such as is used to freeze certain skin growths in the practice of
dermatology). The extra surface area provided by nitrogen-soaked materials is also important, with soaked clothing or cotton causing far more rapid damage than a spill of direct liquid to skin. Full "contact" between naked skin and large droplets or pools of undisturbed liquid nitrogen may be prevented for a few seconds by a layer of insulating gas from the
Leidenfrost effect. However, liquid nitrogen applied to skin in mists, and on fabrics, bypasses this effect.
See also
