Nitrogen: Difference between revisions

Nitrogen, ₇N
	Liquid nitrogen (N2 at below −196 °C)
Nitrogen
Allotropes	see § Allotropes
Appearance	colorless gas, liquid or solid
Standard atomic weight Ar°(N)
	[14.00643, 14.00728]; 14.007±0.001 (abridged);
Nitrogen in the periodic table
	carbon ← nitrogen → oxygen
Atomic number (Z)	7
Group	group 15 (pnictogens)
Period	period 2
Block	p-block
Electron configuration	[He] 2s2 2p3
Electrons per shell	2, 5
Physical properties
Phase at STP	gas
Melting point	(N2) 63.23 K (−209.86 °C, −345.75 °F)
Boiling point	(N2) 77.355 K (−195.795 °C, −320.431 °F)
Density (at STP)	1.2506 g/L at 0 °C, 1013 mbar
when liquid (at b.p.)	0.808 g/cm3
Triple point	63.151 K, 12.52 kPa
Critical point	126.21 K, 3.39 MPa
Heat of fusion	(N2) 0.72 kJ/mol
Heat of vaporization	(N2) 5.57 kJ/mol
Molar heat capacity	(N2) 29.124 J/(mol·K)
Atomic properties
Oxidation states	common: −3, +3, +5; −2, −1, 0, +1, +2, +4
Electronegativity	Pauling scale: 3.04
Ionization energies	1st: 1402.3 kJ/mol ; 2nd: 2856 kJ/mol ; 3rd: 4578.1 kJ/mol ; (more) ;
Covalent radius	71±1 pm
Van der Waals radius	155 pm
	Spectral lines of nitrogen
Other properties
Natural occurrence	primordial
Crystal structure	hexagonal (hP4)
Lattice constants	a = 411.6 pm; c = 673.4 pm (at t.p.)
Thermal conductivity	25.83×10−3 W/(m⋅K)
Magnetic ordering	diamagnetic
Speed of sound	353 m/s (gas, at 27 °C)
CAS Number	17778-88-0 ; 7727-37-9 (N2)
History
Discovery	Daniel Rutherford (1772)
Named by	Jean-Antoine Chaptal (1790)
Isotopes of nitrogen v; e;
β−α<0.01%	12C
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Nitrogen (IPA: /ˈnaɪtrədʒən/) is a chemical element which has the symbol N and atomic number 7. Elemental nitrogen is a colourless, odourless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78.1% by volume 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 nonmetal, 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 (N₂) is one of the strongest in nature. The resulting difficulty of converting (N₂) into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N₂, 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, but with 80.8% of the density, is a common cryogen.

Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like N₃ and N₄.[1] Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced under diamond anvil conditions, nitrogen polymerizes into the single bonded diamond crystal structure, an allotrope nicknamed "nitrogen diamond." [2].

Occurrence

Nitrogen is the largest single component of the Earth's atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air).

¹⁴N is created as part of the fusion processes in stars, and is estimated to be the 7th most abundant chemical element (by mass) in our universe.

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 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 ¹⁴N¹⁵N and almost all the rest is ¹⁴N₂.

Electromagnetic spectrum

Molecular nitrogen (¹⁴N₂) 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).

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 involved uses 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

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 odor 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 odor 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.

Nitrogen gas has a wide variety of applications, including serving as an 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)
In ordinary incandescent light bulbs as an inexpensive alternative to argon
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 manufacturing of stainless steel
Use in military aircraft fuel systems to reduce fire hazard, see inerting system
Filling automotive and aircraft tires^[8] due to its inertness and lack of moisture or oxidative qualities, as opposed to air, though this is not necessary for consumer automobiles.^[9]^[10]

Nitrogen molecules are less likely to escape from the inside of a tire compared with the traditional air mixture used. Air consists mostly of nitrogen and oxygen. Nitrogen molecules have a larger effective diameter than oxygen molecules and therefore diffuse through porous substances more slowly.^[11]

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.

Molecular nitrogen, a diatomic gas, is apt to dimerize into a linear four nitrogen long polymer. This is an important phenomenon for understanding high voltage nitrogen dielectric switches because the process of polymerization can continue to lengthen the molecule to still longer lengths in the presence of an intense electric field. A nitrogen polymer fog is thereby created. The second virial coefficient of nitrogen also shows this effect as the compressibility of nitrogen gas is changed by the dimerization process at moderate and low temperatures.

File:Solid nitrogen.jpg

Solid nitrogen ice in a small plastic beaker with melting liquid flowing off. The nitrogen has been frozen by immersion in liquid helium^[12]

Liquid nitrogen

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 abbreviation, LN₂. 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 below)

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.
as a method of freezing water pipes in order to work on them in situations where a tap is not available to block water flow to the work area.

Nitrogen compounds in industry

Simple compounds

See also the category Nitrogen compounds.

The main neutral hydride of nitrogen is ammonia (NH₃), although hydrazine (N₂H₄) is also commonly used. Ammonia is more basic than water by 6 orders of magnitude. In solution ammonia forms the ammonium ion (NH₄⁺). 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 (NH₂^-); 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 (N₃^-), 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 (N₂O), 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 (NO₂), 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 (N₂O₃) and dinitrogen pentoxide (N₂O₅), are actually fairly unstable and explosive-- a tendency which is driven by the stability of N₂ as a product. The corresponding acids are nitrous (HNO₂) and nitric acid (HNO₃), 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 (N₂) in the atmosphere is relatively non-reactive due to its strong bond, and N₂ 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 (N₂) which is a product of the compounds' thermal decomposition. When nitrates burn or explode, the formation of the powerful triple bond in the N₂ 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 (N₂O) 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 mimicking 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.^[13] 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 body fats, 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, nerves, joints, and other sensitive or vital areas.

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).

References

^ "Standard Atomic Weights: Nitrogen". CIAAW. 2009.
^ Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.
^ ^a ^b ^c Lide, David R. (1990–1991). CRC Handbook of Physics and Chemistry (71st ed.). Boca Raton, Ann Arbor, Boston: CRC Press, inc. pp. 4-22 (one page).
^ "Gases - Density". The Engineering Toolbox. Retrieved 27 January 2019.
^ ^a ^b ^c ^d ^e Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 28. ISBN 978-0-08-037941-8.
^ Tetrazoles contain a pair of double-bonded nitrogen atoms with oxidation state 0 in the ring. A Synthesis of the parent 1H-tetrazole, CH₂N₄ (two atoms N(0)) is given in Henry, Ronald A.; Finnegan, William G. (1954). "An Improved Procedure for the Deamination of 5-Aminotetrazole". Journal of the American Chemical Society. 76 (1): 290–291. doi:10.1021/ja01630a086. ISSN 0002-7863.
^ Arblaster, John W. (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.
^ "Howstuffworks "Why don't they use normal air in race car tires". Retrieved 2006-07-22.
^ "Car Talk: Diffusion, moisture] and tyre expansion". Retrieved 2006-07-22.
^ "The Straight Dope : Is it better to fill your tires with nitrogen instead of air?". Retrieved 2007-02-16.
^ G. J. Van Amerongen. The Permeability of Different Rubbers to Gases and Its Relation to Diffusivity and Solubility. Journal of Applied Physics 1946, 17, 972-985. doi:10.1063/1.1707667.
^ Full video of the solid nitrogen experiment.
^ Biology Safety - Cryogenic materials. The risks posed by them. Accessed 03 Jan 07.

Los Alamos National Laboratory – Nitrogen
Chemistry of the Elements, N. N. Greenwood and A. Earnshaw. ISBN 0-08-022057-6
Biochemistry, R.H. Garrett and C.M. Grisham. 2nd edition, 1999. ISBN 0-03-022318-0
Jahn, GC, LP Almazan, and J Pacia. 2005. Effect of nitrogen fertilizer on the intrinsic rate of increase of the rusty plum aphid, Hysteroneura setariae (Thomas) (Homoptera: Aphididae) on rice (Oryza sativa L.). Environmental Entomology 34 (4): 938-943. [3]

External links

Template:ChemicalSources

Template:Link FA

[1] "Standard Atomic Weights: Nitrogen". CIAAW. 2009.

[CIAAW2021-2] Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.

[CRC71-3] Lide, David R. (1990–1991). CRC Handbook of Physics and Chemistry (71st ed.). Boca Raton, Ann Arbor, Boston: CRC Press, inc. pp. 4-22 (one page).

[4] "Gases - Density". The Engineering Toolbox. Retrieved 27 January 2019.

[GE28-5] Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 28. ISBN 978-0-08-037941-8.

[6] Tetrazoles contain a pair of double-bonded nitrogen atoms with oxidation state 0 in the ring. A Synthesis of the parent 1H-tetrazole, CH₂N₄ (two atoms N(0)) is given in Henry, Ronald A.; Finnegan, William G. (1954). "An Improved Procedure for the Deamination of 5-Aminotetrazole". Journal of the American Chemical Society. 76 (1): 290–291. doi:10.1021/ja01630a086. ISSN 0002-7863.

[Arblaster_2018-7] Arblaster, John W. (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.

[8] "Howstuffworks "Why don't they use normal air in race car tires". Retrieved 2006-07-22.

[9] "Car Talk: Diffusion, moisture] and tyre expansion". Retrieved 2006-07-22.

[10] "The Straight Dope : Is it better to fill your tires with nitrogen instead of air?". Retrieved 2007-02-16.

[11] G. J. Van Amerongen. The Permeability of Different Rubbers to Gases and Its Relation to Diffusivity and Solubility. Journal of Applied Physics 1946, 17, 972-985. doi:10.1063/1.1707667.

[12] Full video of the solid nitrogen experiment.

[13] Biology Safety - Cryogenic materials. The risks posed by them. Accessed 03 Jan 07.

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