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= Nitrogen =
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Introduction
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Nitrogen is a chemical element; it has symbol N and atomic number 7.
Nitrogen is a nonmetal and the lightest member of group 15 of the
periodic table, often called the pnictogens. It is a common element in
the universe, estimated at seventh in total abundance in the Milky Way
and the Solar System. At standard temperature and pressure, two atoms
of the element bond to form N2, a colourless and odourless diatomic
gas. N2 forms about 78% of Earth's atmosphere, making it the most
abundant chemical species in air. Because of the volatility of
nitrogen compounds, nitrogen is relatively rare in the solid parts of
the Earth.
It was first discovered and isolated by Scottish physician Daniel
Rutherford in 1772 and independently by Carl Wilhelm Scheele and Henry
Cavendish at about the same time. The name was suggested by French
chemist Jean-Antoine-Claude Chaptal in 1790 when it was found that
nitrogen was present in nitric acid and nitrates. Antoine Lavoisier
suggested instead the name 'azote', from the "no life", as it is an
asphyxiant gas; this name is used in a number of languages, and
appears in the English names of some nitrogen compounds such as
hydrazine, azides and azo compounds.
Elemental nitrogen is usually produced from air by pressure swing
adsorption technology. About 2/3 of commercially produced elemental
nitrogen is used as an inert (oxygen-free) gas for commercial uses
such as food packaging, and much of the rest is used as liquid
nitrogen in cryogenic applications. Many industrially important
compounds, such as ammonia, nitric acid, organic nitrates (propellants
and explosives), and cyanides, contain nitrogen. The extremely strong
triple bond in elemental nitrogen (N≡N), the second strongest bond in
any diatomic molecule after carbon monoxide (CO), dominates nitrogen
chemistry. This causes difficulty for both organisms and industry in
converting N2 into useful compounds, but at the same time it means
that burning, exploding, or decomposing nitrogen compounds to form
nitrogen gas releases large amounts of often useful energy.
Synthetically produced ammonia and nitrates are key industrial
fertilisers, and fertiliser nitrates are key pollutants in the
eutrophication of water systems. Apart from its use in fertilisers
and energy stores, nitrogen is a constituent of organic compounds as
diverse as aramids used in high-strength fabric and cyanoacrylate used
in superglue.
Nitrogen occurs in all organisms, primarily in amino acids (and thus
proteins), in the nucleic acids (DNA and RNA) and in the energy
transfer molecule adenosine triphosphate. The human body contains
about 3% nitrogen by mass, the fourth most abundant element in the
body after oxygen, carbon, and hydrogen. The nitrogen cycle describes
the movement of the element from the air, into the biosphere and
organic compounds, then back into the atmosphere. Nitrogen is a
constituent of every major pharmacological drug class, including
antibiotics. Many drugs are mimics or prodrugs of natural
nitrogen-containing signal molecules: for example, the organic
nitrates nitroglycerin and nitroprusside control blood pressure by
metabolising into nitric oxide. Many notable nitrogen-containing
drugs, such as the natural caffeine and morphine or the synthetic
amphetamines, act on receptors of animal neurotransmitters.
History
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Nitrogen compounds have a very long history, ammonium chloride having
been known to Herodotus. They were well-known by the Middle Ages.
Alchemists knew nitric acid as 'aqua fortis' (strong water), as well
as other nitrogen compounds such as ammonium salts and nitrate salts.
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 discovery of nitrogen is attributed to the Scottish physician
Daniel Rutherford in 1772, who called it 'noxious air'. Though he did
not recognise it as an entirely different chemical substance, he
clearly distinguished it from Joseph Black's "fixed air", or carbon
dioxide. The fact that there was a component of air that does not
support combustion was clear to Rutherford, although he was not aware
that it was an element. 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'. French
chemist Antoine Lavoisier referred to nitrogen gas as "mephitic air"
or 'azote', from the Greek word (azotikos), "no life", because it is
asphyxiant. In an atmosphere of pure nitrogen, animals died and flames
were extinguished. Though Lavoisier's name was not accepted in English
since it was pointed out that all gases but oxygen are either
asphyxiant or outright toxic, it is used in many languages (French,
Italian, Portuguese, Polish, Russian, Albanian, Turkish, etc.; the
German 'Stickstoff' similarly refers to the same characteristic, viz.
'ersticken' "to choke or suffocate") and still remains in English in
the common names of many nitrogen compounds, such as hydrazine and
compounds of the azide ion. Finally, it led to the name "pnictogens"
for the group headed by nitrogen, from the Greek πνίγειν "to choke".
The English word nitrogen (1794) entered the language from the French
, coined in 1790 by French chemist Jean-Antoine Chaptal (1756-1832),
from the French 'nitre' (potassium nitrate, also called saltpetre) and
the French suffix '-gène', "producing", from the Greek -γενής (-genes,
"begotten"). Chaptal's meaning was that nitrogen is the essential part
of nitric acid, which in turn was produced from nitre. In earlier
times, nitre had been confused with Egyptian "natron" (sodium
carbonate) - called νίτρον (nitron) in Greek - which, despite the
name, contained no nitrate.
The earliest military, industrial, and agricultural applications of
nitrogen compounds used saltpetre (sodium nitrate or potassium
nitrate), most notably in gunpowder, and later as fertiliser. In 1910,
Lord Rayleigh discovered that an electrical discharge in nitrogen gas
produced "active nitrogen", a monatomic allotrope of nitrogen. The
"whirling cloud of brilliant yellow light" produced by his apparatus
reacted with mercury to produce explosive mercury nitride.
For a long time, sources of nitrogen compounds were limited. Natural
sources originated either from biology or deposits of nitrates
produced by atmospheric reactions. Nitrogen fixation by industrial
processes like the Frank-Caro process (1895-1899) and Haber-Bosch
process (1908-1913) eased this shortage of nitrogen compounds, to the
extent that half of global food production now relies on synthetic
nitrogen fertilisers. At the same time, use of the Ostwald process
(1902) to produce nitrates from industrial nitrogen fixation allowed
the large-scale industrial production of nitrates as feedstock in the
manufacture of explosives in the World Wars of the 20th century.
Atomic
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A nitrogen atom has seven electrons. In the ground state, they are
arranged in the electron configuration 1s2s2p2p2p. It, therefore, has
five valence electrons in the 2s and 2p orbitals, three of which (the
p-electrons) are unpaired. It has one of the highest
electronegativities among the elements (3.04 on the Pauling scale),
exceeded only by chlorine (3.16), oxygen (3.44), and fluorine (3.98).
(The light noble gases, helium, neon, and argon, would presumably also
be more electronegative, and in fact are on the Allen scale.)
Following periodic trends, its single-bond covalent radius of 71 pm is
smaller than those of boron (84 pm) and carbon (76 pm), while it is
larger than those of oxygen (66 pm) and fluorine (57 pm). The nitride
anion, N3−, is much larger at 146 pm, similar to that of the oxide
(O2−: 140 pm) and fluoride (F−: 133 pm) anions. The first three
ionisation energies of nitrogen are 1.402, 2.856, and 4.577 MJ·mol−1,
and the sum of the fourth and fifth is . Due to these very high
figures, nitrogen has no simple cationic chemistry.
The lack of radial nodes in the 2p subshell is directly responsible
for many of the anomalous properties of the first row of the p-block,
especially in nitrogen, oxygen, and fluorine. The 2p subshell is very
small and has a very similar radius to the 2s shell, facilitating
orbital hybridisation. It also results in very large electrostatic
forces of attraction between the nucleus and the valence electrons in
the 2s and 2p shells, resulting in very high electronegativities.
Hypervalency is almost unknown in the 2p elements for the same reason,
because the high electronegativity makes it difficult for a small
nitrogen atom to be a central atom in an electron-rich three-center
four-electron bond since it would tend to attract the electrons
strongly to itself. Thus, despite nitrogen's position at the head of
group 15 in the periodic table, its chemistry shows huge differences
from that of its heavier congeners phosphorus, arsenic, antimony, and
bismuth.
Nitrogen may be usefully compared to its horizontal neighbours' carbon
and oxygen as well as its vertical neighbours in the pnictogen column,
phosphorus, arsenic, antimony, and bismuth. Although each period 2
element from lithium to oxygen shows some similarities to the period 3
element in the next group (from magnesium to chlorine; these are known
as diagonal relationships), their degree drops off abruptly past the
boron-silicon pair. The similarities of nitrogen to sulfur are mostly
limited to sulfur nitride ring compounds when both elements are the
only ones present.
Nitrogen does not share the proclivity of carbon for catenation. Like
carbon, nitrogen tends to form ionic or metallic compounds with
metals. Nitrogen forms an extensive series of nitrides with carbon,
including those with chain-, graphitic-, and fullerenic-like
structures.
It resembles oxygen with its high electronegativity and concomitant
capability for hydrogen bonding and the ability to form coordination
complexes by donating its lone pairs of electrons. There are some
parallels between the chemistry of ammonia NH3 and water H2O. For
example, the capacity of both compounds to be protonated to give NH4+
and H3O+ or deprotonated to give NH2− and OH−, with all of these able
to be isolated in solid compounds.
Nitrogen shares with both its horizontal neighbours a preference for
forming multiple bonds, typically with carbon, oxygen, or other
nitrogen atoms, through p'π'-p'π' interactions. Thus, for example,
nitrogen occurs as diatomic molecules and therefore has very much
lower melting (−210 °C) and boiling points (−196 °C) than the rest of
its group, as the N2 molecules are only held together by weak van der
Waals interactions and there are very few electrons available to
create significant instantaneous dipoles. This is not possible for its
vertical neighbours; thus, the nitrogen oxides, nitrites, nitrates,
nitro-, nitroso-, azo-, and diazo-compounds, azides, cyanates,
thiocyanates, and imino-derivatives find no echo with phosphorus,
arsenic, antimony, or bismuth. By the same token, however, the
complexity of the phosphorus oxoacids finds no echo with nitrogen.
Setting aside their differences, nitrogen and phosphorus form an
extensive series of compounds with one another; these have chain,
ring, and cage structures.
Table of thermal and physical properties of nitrogen (N2) at
atmospheric pressure:
|Temperature (K) |Density (kg m−3) |Specific heat (kJ kg−1 °C−1)
|Dynamic viscosity (kg m−1 s−1) |Kinematic viscosity (m2 s−1) |Thermal
conductivity (W m−1 °C−1) |Thermal diffusivity (m2 s−1) |Prandtl
number
|100 |3.4388 |1.07 | | | | |0.768
|150 |2.2594 |1.05 | | | | |0.759
|200 |1.7108 |1.0429 | | | | |0.747
|300 |1.1421 |1.0408 | | | | |0.713
|400 |0.8538 |1.0459 | | | | |0.691
|500 |0.6824 |1.0555 | | | | |0.684
|600 |0.5687 |1.0756 | | | | |0.686
|700 |0.4934 |1.0969 | | | | |0.691
|800 |0.4277 |1.1225 | | | | |0.7
|900 |0.3796 |1.1464 | | | | |0.711
|1000 |0.3412 |1.1677 | | | | |0.724
|1100 |0.3108 |1.1857 | | | | |0.736
|1200 |0.2851 |1.2037 | | | | |0.748
| |0.2591 |1.219 | | | | |0.701
Isotopes
==========
Nitrogen has two stable isotopes: 14N and 15N. The first is much more
common, making up 99.634% of natural nitrogen, and the second (which
is slightly heavier) makes up the remaining 0.366%. This leads to an
atomic weight of around 14.007 u. Both of these stable isotopes are
produced in the CNO cycle in stars, but 14N is more common as its
proton capture is the rate-limiting step. 14N is one of the five
stable odd-odd nuclides (a nuclide having an odd number of protons and
neutrons); the other four are 2H, 6Li, 10B, and 180mTa.
The relative abundance of 14N and 15N is practically constant in the
atmosphere but can vary elsewhere, due to natural isotopic
fractionation from biological redox reactions and the evaporation of
natural ammonia or nitric acid. 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.
The heavy isotope 15N was first discovered by S. M. Naudé in 1929, and
soon after heavy isotopes of the neighbouring elements oxygen and
carbon were discovered. It presents one of the lowest thermal neutron
capture cross-sections of all isotopes. It is frequently used in
nuclear magnetic resonance (NMR) spectroscopy to determine the
structures of nitrogen-containing molecules, due to its fractional
nuclear spin of one-half, which offers advantages for NMR such as
narrower line width. 14N, though also theoretically usable, has an
integer nuclear spin of one and thus has a quadrupole moment that
leads to wider and less useful spectra. 15N NMR nevertheless has
complications not encountered in the more common 1H and 13C NMR
spectroscopy. The low natural abundance of 15N (0.36%) significantly
reduces sensitivity, a problem which is only exacerbated by its low
gyromagnetic ratio, (only 10.14% that of 1H). As a result, the
signal-to-noise ratio for 1H is about 300 times as much as that for
15N at the same magnetic field strength. This may be somewhat
alleviated by isotopic enrichment of 15N by chemical exchange or
fractional distillation. 15N-enriched compounds have the advantage
that under standard conditions, they do not undergo chemical exchange
of their nitrogen atoms with atmospheric nitrogen, unlike compounds
with labelled hydrogen, carbon, and oxygen isotopes that must be kept
away from the atmosphere. The 15N:14N ratio is commonly used in stable
isotope analysis in the fields of geochemistry, hydrology,
paleoclimatology and paleoceanography, where it is called 'δ'15N.
Of the thirteen other isotopes produced synthetically, ranging from 9N
to 23N, 13N has a half-life of ten minutes and the remaining isotopes
have half-lives less than eight seconds. Given the half-life
difference, 13N is the most important nitrogen radioisotope, being
relatively long-lived enough to use in positron emission tomography
(PET), although its half-life is still short and thus it must be
produced at the venue of the PET, for example in a cyclotron via
proton bombardment of 16O producing 13N and an alpha particle.
The radioisotope 16N is the dominant radionuclide in the coolant of
pressurised water reactors or boiling water reactors during normal
operation. It is produced from 16O (in water) via an (n,p) reaction,
in which the 16O atom captures a neutron and expels a proton. It has a
short half-life of about 7.1 s, but its decay back to 16O produces
high-energy gamma radiation (5 to 7 MeV). Because of this, access to
the primary coolant piping in a pressurised water reactor must be
restricted during reactor power operation. It is a sensitive and
immediate indicator of leaks from the primary coolant system to the
secondary steam cycle and is the primary means of detection for such
leaks.
Allotropes
============
Atomic nitrogen, also known as active nitrogen, is highly reactive,
being a triradical with three unpaired electrons. Free nitrogen atoms
easily react with most elements to form nitrides, and even when two
free nitrogen atoms collide to produce an excited N2 molecule, they
may release so much energy on collision with even such stable
molecules as carbon dioxide and water to cause homolytic fission into
radicals such as CO and O or OH and H. Atomic nitrogen is prepared by
passing an electric discharge through nitrogen gas at 0.1-2 mmHg,
which produces atomic nitrogen along with a peach-yellow emission that
fades slowly as an afterglow for several minutes even after the
discharge terminates.
Given the great reactivity of atomic nitrogen, elemental nitrogen
usually occurs as molecular N2, dinitrogen. This molecule is a
colourless, odourless, and tasteless diamagnetic gas at standard
conditions: it melts at −210 °C and boils at −196 °C. Dinitrogen is
mostly unreactive at room temperature, but it will nevertheless react
with lithium metal and some transition metal complexes. This is due to
its bonding, which is unique among the diatomic elements at standard
conditions in that it has an N≡N triple bond. Triple bonds have short
bond lengths (in this case, 109.76 pm) and high dissociation energies
(in this case, 945.41 kJ/mol), and are thus very strong, explaining
dinitrogen's low level of chemical reactivity.
Other nitrogen oligomers and polymers may be possible. If they could
be synthesised, they may have potential applications as materials with
a very high energy density, that could be used as powerful propellants
or explosives. Under extremely high pressures (1.1 million atm) and
high temperatures (2000 K), as produced in a diamond anvil cell,
nitrogen polymerises into the single-bonded cubic gauche crystal
structure. This structure is similar to that of diamond, and both have
extremely strong covalent bonds, resulting in its nickname "nitrogen
diamond".
At atmospheric pressure, molecular nitrogen condenses (liquefies) at
77 K (−195.79 °C) and freezes at 63 K (−210.01 °C) into the beta
hexagonal close-packed crystal allotropic form. Below 35.4 K (−237.6
°C) nitrogen assumes the cubic crystal allotropic form (called the
alpha phase). Liquid nitrogen, a colourless 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.
Solid nitrogen has many crystalline modifications. It forms a
significant dynamic surface coverage on Pluto and outer moons of the
Solar System such as Triton. Even at the low temperatures of solid
nitrogen it is fairly volatile and can sublime to form an atmosphere,
or condense back into nitrogen frost. It is very weak and flows in the
form of glaciers, and on Triton geysers of nitrogen gas come from the
polar ice cap region.
Dinitrogen complexes
======================
The first example of a dinitrogen complex to be discovered was
[Ru(NH3)5(N2)]2+ (see figure at right), and soon many other such
complexes were discovered. These complexes, in which a nitrogen
molecule donates at least one lone pair of electrons to a central
metal cation, illustrate how N2 might bind to the metal(s) in
nitrogenase and the catalyst for the Haber process: these processes
involving dinitrogen activation are vitally important in biology and
in the production of fertilisers.
Dinitrogen is able to coordinate to metals in five different ways. The
more well-characterised ways are the end-on M←N≡N ('η'1) and M←N≡N→M
('μ', bis-'η'1), in which the lone pairs on the nitrogen atoms are
donated to the metal cation. The less well-characterised ways involve
dinitrogen donating electron pairs from the triple bond, either as a
bridging ligand to two metal cations ('μ', bis-'η'2) or to just one
('η'2). The fifth and unique method involves triple-coordination as a
bridging ligand, donating all three electron pairs from the triple
bond ('μ'3-N2). A few complexes feature multiple N2 ligands and some
feature N2 bonded in multiple ways. Since N2 is isoelectronic with
carbon monoxide (CO) and acetylene (C2H2), the bonding in dinitrogen
complexes is closely allied to that in carbonyl compounds, although N2
is a weaker 'σ'-donor and 'π'-acceptor than CO. Theoretical studies
show that 'σ' donation is a more important factor allowing the
formation of the M-N bond than 'π' back-donation, which mostly only
weakens the N-N bond, and end-on ('η'1) donation is more readily
accomplished than side-on ('η'2) donation.
Today, dinitrogen complexes are known for almost all the transition
metals, accounting for several hundred compounds. They are normally
prepared by three methods:
# Replacing labile ligands such as H2O, H−, or CO directly by
nitrogen: these are often reversible reactions that proceed at mild
conditions.
# Reducing metal complexes in the presence of a suitable co-ligand in
excess under nitrogen gas. A common choice includes replacing chloride
ligands with dimethylphenylphosphine (PMe2Ph) to make up for the
smaller number of nitrogen ligands attached to the original chlorine
ligands.
# Converting a ligand with N-N bonds, such as hydrazine or azide,
directly into a dinitrogen ligand.
Occasionally the N≡N bond may be formed directly within a metal
complex, for example by directly reacting coordinated ammonia (NH3)
with nitrous acid (HNO2), but this is not generally applicable. Most
dinitrogen complexes have colours within the range
white-yellow-orange-red-brown; a few exceptions are known, such as the
blue [{Ti('η'5-C5H5)2}2-(N2)].
Nitrides, azides, and nitrido complexes
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Nitrogen bonds to almost all the elements in the periodic table except
the first two noble gases, helium and neon, and some of the very
short-lived elements after bismuth, creating an immense variety of
binary compounds with varying properties and applications. Many binary
compounds are known: with the exception of the nitrogen hydrides,
oxides, and fluorides, these are typically called nitrides. Many
stoichiometric phases are usually present for most elements (e.g. MnN,
Mn6N5, Mn3N2, Mn2N, Mn4N, and Mn'x'N for 9.2 < 'x' < 25.3). They
may be classified as "salt-like" (mostly ionic), covalent,
"diamond-like", and metallic (or interstitial), although this
classification has limitations generally stemming from the continuity
of bonding types instead of the discrete and separate types that it
implies. They are normally prepared by directly reacting a metal with
nitrogen or ammonia (sometimes after heating), or by thermal
decomposition of metal amides:
:3 Ca + N2 → Ca3N2
:3 Mg + 2 NH3 → Mg3N2 + 3 H2 (at 900 °C)
:3 Zn(NH2)2 → Zn3N2 + 4 NH3
Many variants on these processes are possible. The most ionic of these
nitrides are those of the alkali metals and alkaline earth metals,
Li3N (Na, K, Rb, and Cs do not form stable nitrides for steric
reasons) and M3N2 (M = Be, Mg, Ca, Sr, Ba). These can formally be
thought of as salts of the N3− anion, although charge separation is
not actually complete even for these highly electropositive elements.
However, the alkali metal azides NaN3 and KN3, featuring the linear
anion, are well-known, as are Sr(N3)2 and Ba(N3)2. Azides of the
B-subgroup metals (those in groups 11 through 16) are much less ionic,
have more complicated structures, and detonate readily when shocked.
Many covalent binary nitrides are known. Examples include cyanogen
((CN)2), triphosphorus pentanitride (P3N5), disulfur dinitride (S2N2),
and tetrasulfur tetranitride (S4N4). The essentially covalent silicon
nitride (Si3N4) and germanium nitride (Ge3N4) are also known: silicon
nitride, in particular, would make a promising ceramic if not for the
difficulty of working with and sintering it. In particular, the group
13 nitrides, most of which are promising semiconductors, are
isoelectronic with graphite, diamond, and silicon carbide and have
similar structures: their bonding changes from covalent to partially
ionic to metallic as the group is descended. In particular, since the
B-N unit is isoelectronic to C-C, and carbon is essentially
intermediate in size between boron and nitrogen, much of organic
chemistry finds an echo in boron-nitrogen chemistry, such as in
borazine ("inorganic benzene"). Nevertheless, the analogy is not exact
due to the ease of nucleophilic attack at boron due to its deficiency
in electrons, which is not possible in a wholly carbon-containing
ring.
The largest category of nitrides are the interstitial nitrides of
formulae MN, M2N, and M4N (although variable composition is perfectly
possible), where the small nitrogen atoms are positioned in the gaps
in a metallic cubic or hexagonal close-packed lattice. They are
opaque, very hard, and chemically inert, melting only at very high
temperatures (generally over 2500 °C). They have a metallic lustre and
conduct electricity as do metals. They hydrolyse only very slowly to
give ammonia or nitrogen.
The nitride anion (N3−) is the strongest 'π' donor known among ligands
(the second-strongest is O2−). Nitrido complexes are generally made by
the thermal decomposition of azides or by deprotonating ammonia, and
they usually involve a terminal {≡N}3− group. The linear azide anion
(), being isoelectronic with nitrous oxide, carbon dioxide, and
cyanate, forms many coordination complexes. Further catenation is
rare, although (isoelectronic with carbonate and nitrate) is known.
Hydrides
==========
Industrially, ammonia (NH3) is the most important compound of nitrogen
and is prepared in larger amounts than any other compound because it
contributes significantly to the nutritional needs of terrestrial
organisms by serving as a precursor to food and fertilisers. It is a
colourless alkaline gas with a characteristic pungent smell. The
presence of hydrogen bonding has very significant effects on ammonia,
conferring on it its high melting (−78 °C) and boiling (−33 °C)
points. As a liquid, it is a very good solvent with a high heat of
vaporisation (enabling it to be used in vacuum flasks), that also has
a low viscosity and electrical conductivity and high dielectric
constant, and is less dense than water. However, the hydrogen bonding
in NH3 is weaker than that in H2O due to the lower electronegativity
of nitrogen compared to oxygen and the presence of only one lone pair
in NH3 rather than two in H2O. It is a weak base in aqueous solution
(p'K''b' 4.74); its conjugate acid is ammonium, . It can also act as
an extremely weak acid, losing a proton to produce the amide anion, .
It thus undergoes self-dissociation, similar to water, to produce
ammonium and amide. Ammonia burns in air or oxygen, though not
readily, to produce nitrogen gas; it burns in fluorine with a
greenish-yellow flame to give nitrogen trifluoride. Reactions with the
other nonmetals are very complex and tend to lead to a mixture of
products. Ammonia reacts on heating with metals to give nitrides.
Many other binary nitrogen hydrides are known, but the most important
are hydrazine (N2H4) and hydrogen azide (HN3). Although it is not a
nitrogen hydride, hydroxylamine (NH2OH) is similar in properties and
structure to ammonia and hydrazine as well. Hydrazine is a fuming,
colourless liquid that smells similar to ammonia. Its physical
properties are very similar to those of water (melting point 2.0 °C,
boiling point 113.5 °C, density 1.00 g/cm3). Despite it being an
endothermic compound, it is kinetically stable. It burns quickly and
completely in air very exothermically to give nitrogen and water
vapour. It is a very useful and versatile reducing agent and is a
weaker base than ammonia. It is also commonly used as a rocket fuel.
Hydrazine is generally made by reaction of ammonia with alkaline
sodium hypochlorite in the presence of gelatin or glue:
:NH3 + OCl− → NH2Cl + OH−
:NH2Cl + NH3 → + Cl− (slow)
: + OH− → N2H4 + H2O (fast)
(The attacks by hydroxide and ammonia may be reversed, thus passing
through the intermediate NHCl− instead.) The reason for adding gelatin
is that it removes metal ions such as Cu2+ that catalyses the
destruction of hydrazine by reaction with monochloramine (NH2Cl) to
produce ammonium chloride and nitrogen.
Hydrogen azide (HN3) was first produced in 1890 by the oxidation of
aqueous hydrazine by nitrous acid. It is very explosive and even
dilute solutions can be dangerous. It has a disagreeable and
irritating smell and is a potentially lethal (but not cumulative)
poison. It may be considered the conjugate acid of the azide anion,
and is similarly analogous to the hydrohalic acids.
Halides and oxohalides
========================
All four simple nitrogen trihalides are known. A few mixed halides and
hydrohalides are known, but are mostly unstable; examples include
NClF2, NCl2F, NBrF2, NF2H, NFH2, NCl2H, and NClH2.
Nitrogen trifluoride (NF3, first prepared in 1928) is a colourless and
odourless gas that is thermodynamically stable, and most readily
produced by the electrolysis of molten ammonium fluoride dissolved in
anhydrous hydrogen fluoride. Like carbon tetrafluoride, it is not at
all reactive and is stable in water or dilute aqueous acids or
alkalis. Only when heated does it act as a fluorinating agent, and it
reacts with copper, arsenic, antimony, and bismuth on contact at high
temperatures to give tetrafluorohydrazine (N2F4). The cations and
are also known (the latter from reacting tetrafluorohydrazine with
strong fluoride-acceptors such as arsenic pentafluoride), as is ONF3,
which has aroused interest due to the short N-O distance implying
partial double bonding and the highly polar and long N-F bond.
Tetrafluorohydrazine, unlike hydrazine itself, can dissociate at room
temperature and above to give the radical NF2•. Fluorine azide (FN3)
is very explosive and thermally unstable. Dinitrogen difluoride (N2F2)
exists as thermally interconvertible 'cis' and 'trans' isomers, and
was first found as a product of the thermal decomposition of FN3.
Nitrogen trichloride (NCl3) is a dense, volatile, and explosive liquid
whose physical properties are similar to those of carbon
tetrachloride, although one difference is that NCl3 is easily
hydrolysed by water while CCl4 is not. It was first synthesised in
1811 by Pierre Louis Dulong, who lost three fingers and an eye to its
explosive tendencies. As a dilute gas it is less dangerous and is thus
used industrially to bleach and sterilise flour. Nitrogen tribromide
(NBr3), first prepared in 1975, is a deep red, temperature-sensitive,
volatile solid that is explosive even at −100 °C. Nitrogen triiodide
(NI3) is still more unstable and was only prepared in 1990. Its adduct
with ammonia, which was known earlier, is very shock-sensitive: it can
be set off by the touch of a feather, shifting air currents, or even
alpha particles. For this reason, small amounts of nitrogen triiodide
are sometimes synthesised as a demonstration to high school chemistry
students or as an act of "chemical magic". Chlorine azide (ClN3) and
bromine azide (BrN3) are extremely sensitive and explosive.
Two series of nitrogen oxohalides are known: the nitrosyl halides
(XNO) and the nitryl halides (XNO2). The first is very reactive gases
that can be made by directly halogenating nitrous oxide. Nitrosyl
fluoride (NOF) is colourless and a vigorous fluorinating agent.
Nitrosyl chloride (NOCl) behaves in much the same way and has often
been used as an ionising solvent. Nitrosyl bromide (NOBr) is red. The
reactions of the nitryl halides are mostly similar: nitryl fluoride
(FNO2) and nitryl chloride (ClNO2) are likewise reactive gases and
vigorous halogenating agents.
Oxides
========
Nitrogen forms nine molecular oxides, some of which were the first
gases to be identified: N2O (nitrous oxide), NO (nitric oxide), N2O3
(dinitrogen trioxide), NO2 (nitrogen dioxide), N2O4 (dinitrogen
tetroxide), N2O5 (dinitrogen pentoxide), N4O (nitrosylazide), and
N(NO2)3 (trinitramide). All are thermally unstable towards
decomposition to their elements. One other possible oxide that has not
yet been synthesised is oxatetrazole (N4O), an aromatic ring.
Nitrous oxide (N2O), better known as laughing gas, is made by thermal
decomposition of molten ammonium nitrate at 250 °C. This is a redox
reaction and thus nitric oxide and nitrogen are also produced as
byproducts. It is mostly used as a propellant and aerating agent for
sprayed canned whipped cream, and was formerly commonly used as an
anaesthetic. Despite appearances, it cannot be considered to be the
anhydride of hyponitrous acid (H2N2O2) because that acid is not
produced by the dissolution of nitrous oxide in water. It is rather
unreactive (not reacting with the halogens, the alkali metals, or
ozone at room temperature, although reactivity increases upon heating)
and has the unsymmetrical structure N-N-O (N≡N+O−↔−N=N+=O): above 600
°C it dissociates by breaking the weaker N-O bond.
Nitric oxide (NO) is the simplest stable molecule with an odd number
of electrons. In mammals, including humans, it is an important
cellular signalling molecule involved in many physiological and
pathological processes. It is formed by catalytic oxidation of
ammonia. It is a colourless paramagnetic gas that, being
thermodynamically unstable, decomposes to nitrogen and oxygen gas at
1100-1200 °C. Its bonding is similar to that in nitrogen, but one
extra electron is added to a 'π'* antibonding orbital and thus the
bond order has been reduced to approximately 2.5; hence dimerisation
to O=N-N=O is unfavourable except below the boiling point (where the
'cis' isomer is more stable) because it does not actually increase the
total bond order and because the unpaired electron is delocalised
across the NO molecule, granting it stability. There is also evidence
for the asymmetric red dimer O=N-O=N when nitric oxide is condensed
with polar molecules. It reacts with oxygen to give brown nitrogen
dioxide and with halogens to give nitrosyl halides. It also reacts
with transition metal compounds to give nitrosyl complexes, most of
which are deeply coloured.
Blue dinitrogen trioxide (N2O3) is only available as a solid because
it rapidly dissociates above its melting point to give nitric oxide,
nitrogen dioxide (NO2), and dinitrogen tetroxide (N2O4). The latter
two compounds are somewhat difficult to study individually because of
the equilibrium between them, although sometimes dinitrogen tetroxide
can react by heterolytic fission to nitrosonium and nitrate in a
medium with high dielectric constant. Nitrogen dioxide is an acrid,
corrosive brown gas. Both compounds may be easily prepared by
decomposing a dry metal nitrate. Both react with water to form nitric
acid. Dinitrogen tetroxide is very useful for the preparation of
anhydrous metal nitrates and nitrato complexes, and it became the
storable oxidiser of choice for many rockets in both the United States
and USSR by the late 1950s. This is because it is a hypergolic
propellant in combination with a hydrazine-based rocket fuel and can
be easily stored since it is liquid at room temperature.
The thermally unstable and very reactive dinitrogen pentoxide (N2O5)
is the anhydride of nitric acid, and can be made from it by
dehydration with phosphorus pentoxide. It is of interest for the
preparation of explosives. It is a deliquescent, colourless
crystalline solid that is sensitive to light. In the solid state it is
ionic with structure [NO2]+[NO3]−; as a gas and in solution it is
molecular O2N-O-NO2. Hydration to nitric acid comes readily, as does
analogous reaction with hydrogen peroxide giving peroxonitric acid
(HOONO2). It is a violent oxidising agent. Gaseous dinitrogen
pentoxide decomposes as follows:
:N2O5 NO2 + NO3 → NO2 + O2 + NO
:N2O5 + NO 3 NO2
Oxoacids, oxoanions, and oxoacid salts
========================================
Many nitrogen oxoacids are known, though most of them are unstable as
pure compounds and are known only as aqueous solutions or as salts.
Hyponitrous acid (H2N2O2) is a weak diprotic acid with the structure
HON=NOH (p'K'a1 6.9, p'K'a2 11.6). Acidic solutions are quite stable
but above pH 4 base-catalysed decomposition occurs via [HONNO]− to
nitrous oxide and the hydroxide anion. Hyponitrites (involving the
anion) are stable to reducing agents and more commonly act as reducing
agents themselves. They are an intermediate step in the oxidation of
ammonia to nitrite, which occurs in the nitrogen cycle. Hyponitrite
can act as a bridging or chelating bidentate ligand.
Nitrous acid (HNO2) is not known as a pure compound, but is a common
component in gaseous equilibria and is an important aqueous reagent:
its aqueous solutions may be made from acidifying cool aqueous nitrite
(, bent) solutions, although already at room temperature
disproportionation to nitrate and nitric oxide is significant. It is a
weak acid with p'K''a' 3.35 at 18 °C. They may be titrimetrically
analysed by their oxidation to nitrate by permanganate. They are
readily reduced to nitrous oxide and nitric oxide by sulfur dioxide,
to hyponitrous acid with tin(II), and to ammonia with hydrogen
sulfide. Salts of hydrazinium react with nitrous acid to produce
azides which further react to give nitrous oxide and nitrogen. Sodium
nitrite is mildly toxic in concentrations above 100 mg/kg, but small
amounts are often used to cure meat and as a preservative to avoid
bacterial spoilage. It is also used to synthesise hydroxylamine and to
diazotise primary aromatic amines as follows:
:ArNH2 + HNO2 → [ArNN]Cl + 2 H2O
Nitrite is also a common ligand that can coordinate in five ways. The
most common are nitro (bonded from the nitrogen) and nitrito (bonded
from an oxygen). Nitro-nitrito isomerism is common, where the nitrito
form is usually less stable.
Nitric acid (HNO3) is by far the most important and the most stable of
the nitrogen oxoacids. It is one of the three most used acids (the
other two being sulfuric acid and hydrochloric acid) and was first
discovered by alchemists in the 13th century. It is made by the
catalytic oxidation of ammonia to nitric oxide, which is oxidised to
nitrogen dioxide, and then dissolved in water to give concentrated
nitric acid. In the United States of America, over seven million
tonnes of nitric acid are produced every year, most of which is used
for nitrate production for fertilisers and explosives, among other
uses. Anhydrous nitric acid may be made by distilling concentrated
nitric acid with phosphorus pentoxide at low pressure in glass
apparatus in the dark. It can only be made in the solid state, because
upon melting it spontaneously decomposes to nitrogen dioxide, and
liquid nitric acid undergoes self-ionisation to a larger extent than
any other covalent liquid as follows:
:2 HNO3 + H2O + [NO2]+ + [NO3]−
Two hydrates, HNO3·H2O and HNO3·3H2O, are known that can be
crystallised. It is a strong acid and concentrated solutions are
strong oxidising agents, though gold, platinum, rhodium, and iridium
are immune to attack. A 3:1 mixture of concentrated hydrochloric acid
and nitric acid, called 'aqua regia', is still stronger and
successfully dissolves gold and platinum, because free chlorine and
nitrosyl chloride are formed and chloride anions can form strong
complexes. In concentrated sulfuric acid, nitric acid is protonated to
form nitronium, which can act as an electrophile for aromatic
nitration:
:HNO3 + 2 H2SO4 + H3O+ + 2
The thermal stabilities of nitrates (involving the trigonal planar
anion) depends on the basicity of the metal, and so do the products of
decomposition (thermolysis), which can vary between the nitrite (for
example, sodium), the oxide (potassium and lead), or even the metal
itself (silver) depending on their relative stabilities. Nitrate is
also a common ligand with many modes of coordination.
Finally, although orthonitric acid (H3NO4), which would be analogous
to orthophosphoric acid, does not exist, the tetrahedral orthonitrate
anion is known in its sodium and potassium salts:
:NaNO3{} + Na2O ->[\ce{Ag~crucible}][\ce{300^\circ C~for~7 days}]
Na3NO4
These white crystalline salts are very sensitive to water vapour and
carbon dioxide in the air:
:Na3NO4 + H2O + CO2 → NaNO3 + NaOH + NaHCO3
Despite its limited chemistry, the orthonitrate anion is interesting
from a structural point of view due to its regular tetrahedral shape
and the short N-O bond lengths, implying significant polar character
to the bonding.
Organic nitrogen compounds
============================
Nitrogen is one of the most important elements in organic chemistry.
Many organic functional groups involve a carbon-nitrogen bond, such as
amides (RCONR2), amines (R3N), imines (RC(=NR)R), imides (RCO)2NR,
azides (RN3), azo compounds (RN2R), cyanates (ROCN), isocyanates
(RNCO), nitrates (RONO2), nitriles (RCN), isonitriles (RNC), nitrites
(RONO), nitro compounds (RNO2), nitroso compounds (RNO), oximes
(RC(=NOH)R), and pyridine derivatives. C-N bonds are strongly
polarised towards nitrogen. In these compounds, nitrogen is usually
trivalent (though it can be tetravalent in quaternary ammonium salts,
R4N+), with a lone pair that can confer basicity on the compound by
being coordinated to a proton. This may be offset by other factors:
for example, amides are not basic because the lone pair is delocalised
into a double bond (though they may act as bases at very low pH, being
protonated at the oxygen), and pyrrole is not basic because the lone
pair is delocalised as part of an aromatic ring. The amount of
nitrogen in a chemical substance can be determined by the Kjeldahl
method. In particular, nitrogen is an essential component of nucleic
acids, amino acids and thus proteins, and the energy-carrying molecule
adenosine triphosphate and is thus vital to all life on Earth.
Occurrence
======================================================================
Nitrogen is the most common pure element in the earth, making up 78.1%
of the volume of the atmosphere (75.5% by mass), around 3.89 million
gigatonnes (). Despite this, it is not very abundant in Earth's crust,
making up somewhere around 19 parts per million of this, on par with
niobium, gallium, and lithium. (This represents 300,000 to a million
gigatonnes of nitrogen, depending on the mass of the crust.) The only
important nitrogen minerals are nitre (potassium nitrate, saltpetre)
and soda nitre (sodium nitrate, Chilean saltpetre). However, these
have not been an important source of nitrates since the 1920s, when
the industrial synthesis of ammonia and nitric acid became common.
Nitrogen compounds constantly interchange between the atmosphere and
living organisms. Nitrogen must first be processed, or "fixed", into a
plant-usable form, usually ammonia. Some nitrogen fixation is done by
lightning strikes producing the nitrogen oxides, but most is done by
diazotrophic bacteria through enzymes known as nitrogenases (although
today industrial nitrogen fixation to ammonia is also significant).
When the ammonia is taken up by plants, it is used to synthesise
proteins. These plants are then digested by animals who use the
nitrogen compounds to synthesise their proteins and excrete
nitrogen-bearing waste. Finally, these organisms die and decompose,
undergoing bacterial and environmental oxidation and denitrification,
returning free dinitrogen to the atmosphere. Industrial nitrogen
fixation by the Haber process is mostly used as fertiliser, although
excess nitrogen-bearing waste, when leached, leads to eutrophication
of freshwater and the creation of marine dead zones, as
nitrogen-driven bacterial growth depletes water oxygen to the point
that all higher organisms die. Furthermore, nitrous oxide, which is
produced during denitrification, attacks the atmospheric ozone layer.
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. In animals, free radical nitric
oxide (derived from an amino acid), serves as an important regulatory
molecule for circulation.
Nitric oxide's rapid reaction with water in animals results in the
production of its metabolite nitrite. Animal metabolism of nitrogen in
proteins, in general, results in the excretion of urea, while animal
metabolism of nucleic acids results in the excretion of urea and uric
acid. The characteristic odour of animal flesh decay is caused by the
creation of long-chain, nitrogen-containing amines, such as putrescine
and cadaverine, which are breakdown products of the amino acids
ornithine and lysine, respectively, in decaying proteins.
Production
======================================================================
Nitrogen gas is an industrial gas produced by the fractional
distillation of liquid air, or by mechanical means using gaseous air
(pressurised reverse osmosis membrane or pressure swing adsorption).
Nitrogen gas generators using membranes or pressure swing adsorption
(PSA) are typically more cost and energy efficient than bulk-delivered
nitrogen. 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 called OFN
(oxygen-free nitrogen). Commercial-grade nitrogen already contains at
most 20 ppm oxygen, and specially purified grades containing at most 2
ppm oxygen and 10 ppm argon are also available.
In a chemical laboratory, it is prepared by treating an aqueous
solution of ammonium chloride with sodium nitrite.
:NH4Cl + NaNO2 → N2 + NaCl + 2 H2O
Small amounts of the impurities NO and HNO3 are also formed in this
reaction. The impurities can be removed by passing the gas through
aqueous sulfuric acid containing potassium dichromate.
It can also be obtained by the thermal decomposition of ammonium
dichromate.
:3(NH4)2Cr2O7 → 2N2 + 9H2O + 3Cr2O3 + 2NH3 + 32O2
Very pure nitrogen can be prepared by the thermal decomposition of
barium azide or sodium azide.
:2 NaN3 → 2 Na + 3 N2
Applications
======================================================================
The applications of nitrogen compounds are naturally extremely widely
varied due to the huge size of this class: hence, only applications of
pure nitrogen itself will be considered here. Two-thirds (2/3) of
nitrogen produced by industry is sold as gas and the remaining
one-third (1/3) as a liquid.
Gas{{anchor|Nitrogen_gas}}
============================
The gas is mostly used as a low reactivity safe atmosphere wherever
the oxygen in the air would pose a fire, explosion, or oxidising
hazard. Some examples include:
* As a modified atmosphere, pure or mixed with carbon dioxide, to
nitrogenate and preserve the freshness of packaged or bulk foods (by
delaying rancidity and other forms of oxidative damage). Pure nitrogen
as food additive is labelled in the European Union with the E number
E941.
* In incandescent light bulbs as an inexpensive alternative to argon.
* In fire suppression systems for Information technology (IT)
equipment.
* In the manufacture of stainless steel.
* In the case-hardening of steel by nitriding.
* In some aircraft fuel systems to reduce fire hazard (see inerting
system).
* To inflate race car and aircraft tires, reducing the problems of
inconsistent expansion and contraction caused by moisture and oxygen
in natural air.
Nitrogen is commonly used during sample preparation in chemical
analysis. It is used to concentrate and reduce the volume of liquid
samples. Directing a pressurised stream of nitrogen gas perpendicular
to the surface of the liquid causes the solvent to evaporate while
leaving the solute(s) and un-evaporated solvent behind.
Nitrogen can be used as a replacement, or in combination with, carbon
dioxide to pressurise kegs of some beers, particularly stouts and
British ales, due to the smaller bubbles it produces, which makes the
dispensed beer smoother and headier. A pressure-sensitive nitrogen
capsule known commonly as a "widget" allows nitrogen-charged beers to
be packaged in cans and bottles. Nitrogen tanks are also replacing
carbon dioxide as the main power source for paintball guns. Nitrogen
must be kept at a higher pressure than CO2, making N2 tanks heavier
and more expensive.
Equipment
===========
Some construction equipment uses pressurised nitrogen gas to help
hydraulic system to provide extra power to devices such as hydraulic
hammer. Nitrogen gas, formed from the decomposition of sodium azide,
is used for the inflation of airbags.
Execution
===========
As nitrogen is an asphyxiant gas in itself, some jurisdictions have
considered asphyxiation by inhalation of pure nitrogen as a means of
capital punishment (as a substitute for lethal injection). In January
2024, Kenneth Eugene Smith became the first person executed by
nitrogen asphyxiation.
Liquid
========
Liquid nitrogen is a cryogenic liquid which looks like water. When
insulated in proper containers such as dewar flasks, it can be
transported and stored with a low rate of evaporative loss.
Like dry ice, the main use of liquid nitrogen is for cooling to low
temperatures. It is used in the cryopreservation of biological
materials such as blood and reproductive cells (sperm and eggs). It is
used in cryotherapy to remove cysts and warts on the skin by freezing
them. It is used in laboratory cold traps, and in cryopumps to obtain
lower pressures in vacuum pumped systems. It is used to cool
heat-sensitive electronics such as infrared detectors and X-ray
detectors. Other uses include freeze-grinding and machining materials
that are soft or rubbery at room temperature, shrink-fitting and
assembling engineering components, and more generally to attain very
low temperatures where necessary. Because of its low cost, liquid
nitrogen is often used for cooling even when such low temperatures are
not strictly necessary, such as refrigeration of food, freeze-branding
livestock, freezing pipes to halt flow when valves are not present,
and consolidating unstable soil by freezing whenever excavation is
going on underneath.
As a cryogenic liquid, liquid nitrogen can be dangerous by causing
cold burns on contact, although the Leidenfrost effect provides
protection for very short exposure (about one second). Ingestion of
liquid nitrogen can cause severe internal damage. For example, in
2012, a young woman in England had to have her stomach removed after
ingesting a cocktail made with liquid nitrogen.
Because the liquid-to-gas expansion ratio of nitrogen is 1:694 at 20
°C, a tremendous amount of force can be generated if liquid nitrogen
is rapidly vaporised in an enclosed space. In an incident on January
12, 2006, at Texas A&M University, the pressure-relief devices of
a tank of liquid nitrogen were malfunctioning and later sealed. As a
result of the subsequent pressure buildup, the tank failed
catastrophically. The force of the explosion was sufficient to propel
the tank through the ceiling immediately above it, shatter a
reinforced concrete beam immediately below it, and blow the walls of
the laboratory 0.1-0.2 m off their foundations.
Liquid nitrogen readily evaporates to form gaseous nitrogen, and hence
the precautions associated with gaseous nitrogen also apply to liquid
nitrogen. For example, oxygen sensors are sometimes used as a safety
precaution when working with liquid nitrogen to alert workers of gas
spills into a confined space.
Vessels containing liquid nitrogen can condense oxygen from air. The
liquid in such a vessel becomes increasingly enriched in oxygen
(boiling point −183 °C, higher than that of nitrogen) as the nitrogen
evaporates, and can cause violent oxidation of organic material.
Gas
=====
Although nitrogen is non-toxic, when released into an enclosed space
it can displace oxygen, and therefore presents an asphyxiation hazard.
This may happen with few warning symptoms, since the human carotid
body is a relatively poor and slow low-oxygen (hypoxia) sensing
system. An example occurred shortly before the launch of the first
Space Shuttle mission on March 19, 1981, when two technicians died
from asphyxiation after they walked into a space located in the Space
Shuttle's mobile launcher platform that was pressurised with pure
nitrogen as a precaution against fire.
When inhaled at high partial pressures (more than about 4 bar,
encountered at depths below about 30 m in scuba diving), nitrogen is
an anaesthetic agent, causing nitrogen narcosis, a temporary state of
mental impairment similar to nitrous oxide intoxication.
Nitrogen dissolves in the blood and body fats. Rapid decompression (as
when divers ascend too quickly or astronauts decompress too quickly
from cabin pressure to spacesuit pressure) can lead to a potentially
fatal condition called decompression sickness (formerly known as
caisson sickness or 'the bends'), when nitrogen bubbles form in the
bloodstream, nerves, joints, and other sensitive or vital areas.
Bubbles from other "inert" gases (gases other than carbon dioxide and
oxygen) cause the same effects, so replacement of nitrogen in
breathing gases may prevent nitrogen narcosis, but does not prevent
decompression sickness.
Oxygen deficiency monitors
============================
Oxygen deficiency monitors are used to measure levels of oxygen in
confined spaces and any place where nitrogen gas or liquid are stored
or used. In the event of a nitrogen leak, and a decrease in oxygen to
a pre-set alarm level, an oxygen deficiency monitor can be programmed
to set off audible and visual alarms, thereby providing notification
of the possible impending danger. Most commonly, the oxygen range to
alert personnel is when oxygen levels get below 19.5%. OSHA specifies
that a hazardous atmosphere may include one where the oxygen
concentration is below 19.5% or above 23.5%.
Oxygen deficiency monitors can either be fixed, mounted to the wall
and hard-wired into the building's power supply or simply plugged into
a power outlet, or a portable hand-held or wearable monitor.
See also
======================================================================
* Reactive nitrogen species
* Soil gas
External links
======================================================================
* [
http://www.balashon.com/2008/07/neter-and-nitrogen.html Etymology
of Nitrogen]
* [
http://www.periodicvideos.com/videos/007.htm Nitrogen] at 'The
Periodic Table of Videos' (University of Nottingham)
* [
http://www.rsc.org/periodic-table/podcast/7/nitrogen Nitrogen
podcast] from the Royal Society of Chemistry's 'Chemistry World'
License
=========
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Original Article:
http://en.wikipedia.org/wiki/Nitrogen
