Nitrogen

7 carbon ← nitrogen → oxygen - ↑ N ↓ P periodic table
General
Name, Symbol, Number	nitrogen, N, 7
Chemical series	nonmetals
Group, Period, Block	15, 2, p
Appearance	colorless
Atomic mass	14.0067(2) g/mol
Electron configuration	1s2 2s2 2p3
Electrons per shell	2, 5
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 1.251 g/L
Melting point	63.15 K (-210.00 °C, -346.00 °F)
Boiling point	77.36 K (-195.79 °C, -320.42 °F)
Critical point	126.21 K, 3.39 MPa
Heat of fusion	(N2) 0.720 kJ/mol
Heat of vaporization	(N2) 5.57 kJ/mol
Heat capacity	(25 °C) (N2) 29.124 J/(mol·K)
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K 37 41 46 53 62 77
Atomic properties
Crystal structure	hexagonal
Oxidation states	±3, 5, 4, 2 (strongly acidic oxide)
Electronegativity	3.04 (Pauling scale)
Ionization energies (more)	1st: 1402.3 kJ/mol
	2nd: 2856 kJ/mol
	3rd: 4578.1 kJ/mol
Atomic radius	65 pm
Atomic radius (calc.)	56 pm
Covalent radius	75 pm
Van der Waals radius	155 pm
Miscellaneous
Magnetic ordering	no data
Thermal conductivity	(300 K) 25.83 mW/(m·K)
Speed of sound	(gas, 27 °C) 353 m/s
CAS registry number	7727-37-9
Notable isotopes
Main article: Isotopes of nitrogen iso NA half-life DM DE (MeV) DP 13N syn 9.965 m ε 2.220 13C 14N 99.634% N is stable with 7 neutrons 15N 0.366% N is stable with 8 neutrons

Nitrogen is a chemical element which has the symbol N and atomic number 7 in the periodic table. Elemental nitrogen is a colorless, odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78.08% percent of Earth's atmosphere. Nitrogen is a constituent element of all living tissues and amino acids. Many industrially important compounds, such as ammonia, nitric acid, and cyanides, contain nitrogen.

Occurrence

Nitrogen is the largest single component of the Earth's atmosphere—78.084% by volume and 75.5% by weight. It appears that the most common isotope, nitrogen-14 (¹⁴N) is created as part of the nuclear fusion processes in stars.

Compounds that contain this element have been observed by astronomers, and molecular nitrogen has been detected in interstellar space by David Knauth and coworkers using the Far Ultraviolet Spectroscopic Explorer. Molecular nitrogen occurs in trace amounts in the atmospheres of various planets, but it is a major constituent of Titan, the planet Saturn's largest moon.

Nitrogen is present in all living organisms, as part of the molecular structures of proteins, nucleic acids, and other important substances. It is a large component of animal waste, usually in the form of urea, uric acid, and their derivatives.

Discovery and etymology

Nitrogen (Latin nitrum, Greek Nitron meaning "native soda," "genes," "forming") is formally considered to have been discovered in 1772 by the chemist Daniel Rutherford, who knew that there was a fraction of air that did not support combustion. He called it "noxious air" or "fixed air." 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 (unreactive) 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 became the French word for nitrogen and later spread to many other languages.

The alchemists of the Middle Ages experimented with various compounds of nitrogen. For instance, they knew nitric acid as aqua fortis (strong water). The mixture of nitric acid and hydrochloric acid was called aqua regia (royal water), celebrated for its ability to dissolve gold (the "king" of metals). In the earliest industrial and agricultural applications of nitrogen compounds, saltpeter (sodium nitrate or potassium nitrate) was used in gunpowder, much later as fertilizer, and later still as a chemical feedstock.

Notable characteristics

Nitrogen is a chemical element in the periodic table, situated at the head of group 15 (former group 5A), just above phosphorus. In addition, it lies in period 2, flanked by carbon and oxygen. Classified as a nonmetal, it has an electronegativity of 3.0. Each atom of nitrogen has five electrons in its outer shell, and it forms three covalent bonds in most compounds.

A computer rendering of the nitrogen molecule, N₂.

Nitrogen gas consists of diatomic molecules, each of which has the chemical formula N₂. This gas condenses to the liquid form at 77 Kelvin (K) at atmospheric pressure and freezes at 63 K. Liquid nitrogen is a common cryogen (an extremely low-temperature refrigerant) that can cause instant frostbite on direct contact with living tissue.

Isotopes

There are two stable isotopes of nitrogen: ¹⁴N and ¹⁵N. By far the most common is ¹⁴N (99.634%), which is thought to be produced in stars by a set of nuclear fusion reactions called the "carbon-nitrogen-oxygen cycle" (CNO cycle). Of the 10 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.

In the Earth's atmosphere, 0.73% of molecular nitrogen consists of ¹⁴N¹⁵N, and almost all the rest is ¹⁴N₂.

Biological role

Nitrogen is an essential element in the molecules of amino acids, proteins, nucleic acids, and other substances vital to life. Specific bacteria (such as those of the genus Rhizobium) possess certain enzymes ("nitrogenases") that can fix atmospheric nitrogen (see nitrogen fixation) into a form (ammonium ion) that is chemically useful for higher organisms. This process requires a large amount of energy and anoxic (oxygen-free) conditions. Usually, these bacteria exist in a symbiotic relationship in the root nodules of leguminous plants such as clover or the soya bean plant. Nitrogen-fixing bacteria can also be symbiotic with other plant species, such as alders, lichens, casuarina, myrica, liverwort, and gunnera.

As part of the symbiotic relationship, the plant converts the ammonium ions to nitrogen oxides and amino acids to form proteins and other biologically useful molecules, such as alkaloids. In return, the plant secretes sugars that the bacteria can use.

Some plants can assimilate nitrogen directly in the form of nitrates, which may be present in the soil from natural mineral deposits, artificial fertilizers, animal waste, or organic decay (as the product of bacteria that are not 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. Many saltwater fish manufacture large amounts of trimethylamine oxide to protect them from the high osmotic effects of their environment. In animals, nitric oxide (NO) is derived from an amino acid and serves as an important regulatory molecule for circulation.

Animal metabolism of nitrogen in proteins generally results in the excretion of urea, while animal metabolism of nucleic acids results in the excretion of urea and uric acid. The characteristic odor of animal flesh decay is caused by nitrogen-containing long-chain amines, such as putrescene and cadaverine. The decay of organisms and their waste products may produce small amounts of nitrate, but most decay processes eventually return nitrogen content to the atmosphere, as molecular nitrogen.

Industrial production

Nitrogen is produced industrially in large quantities by a method known as the fractional distillation of liquefied air. Isolated in the liquid form, this nitrogen is often referred to by the quasi-formula LN₂, but it is more accurately written as N₂(l). Technologies that isolate nitrogen from gaseous air include methods known as pressure swing adsorption and membrane separation. In addition, commercial nitrogen is often obtained as a byproduct of air-processing during the industrial concentration of oxygen for steelmaking and other purposes. The formula for gaseous nitrogen is N₂(g).

Modern applications of molecular nitrogen

Nitrogen gas has a wide variety of applications, including serving as a more inert replacement for air where oxidation is undesirable. Some examples are as follows.

• It helps preserve the freshness of packaged or bulk foods by delaying the onset of rancidity and other forms of oxidative damage.
• It is placed on top of liquid explosives for safety.
• It is used in the production of electronic parts such as transistors, diodes, and integrated circuits.

A tank of liquid nitrogen that supplies a cryogenic freezer (for storing laboratory samples at a temperature of about -150 °C).

• Dried and pressurized nitrogen is used as a dielectric gas for high-voltage equipment.
• It is used in the manufacture of stainless steel.
• Given its inertness and lack of moisture (as opposed to air), nitrogen gas is used to fill race-car and aircraft tires, though this is not necessary for consumer automobiles.

When appropriately insulated from ambient heat, liquid nitrogen serves as a compact, readily transportable source of nitrogen gas without pressurization. Furthermore, its ability to maintain temperatures far below the freezing point of water as it boils at (77 K, -196 °C or -320 °F) makes it extremely useful in a wide range of applications as an open-cycle refrigerant. Some uses of liquid nitrogen include:

• The immersion freezing and transportation of food products.
• The cryopreservation of blood, reproductive cells (sperm and egg), and other biological samples and materials.
• The cryonic preservation of humans and pets in the hope of future revival with molecular repair technology.
• The study of cryogenics.
• Demonstrations in science education.
• As a coolant, it is used for for highly sensitive sensors and low-noise amplifiers.
• In dermatology, it is used for removing unsightly or potentially malignant skin lesions such as warts and actinic keratosis.
• It is a cooling medium during machining of high strength materials.
• It is a supplement for cooling computer hardware such as a central processing unit or a graphics processining unit.

Compounds of nitrogen

Inorganic compounds

Ammonia: The main neutral hydride of nitrogen is ammonia (NH₃), although hydrazine (N₂H₄) is also common. Ammonia is a chemical base— more basic than water by 6 orders of magnitude. In solution, ammonia combines with protons (H⁺ ions) to form ammonium ions (NH₄⁺). Liquid ammonia (boiling point 240 K) is "amphiprotic"—that is, it can behave as an acid as well as a base. As an acid, it donates a proton (H⁺) to another molecule to form the amide ion (NH₂⁻); as a base, it receives a proton (H⁺) from another molecule to form the ammonium ion (NH₄⁺).

Amide: An inorganic amide is a compound in which a metal cation is combined with the amide anion (NH₂^-) mentioned above. An example is sodium amide (NaNH₂). An inorganic amide is an extremely strong base and decomposes in water. Note that inorganic amide salts are distinctly different from organic amide compounds mentioned below.

Nitride: In a nitride compound, nitrogen is attached to a more electropositive element. Some nitrides, such as lithium nitride (Li₃N), are salt-like, in which the nitrogen exists as an ion with three negative charges (N³⁻). The salt-like nitrides are strong bases and readily decompose in water. Other nitrides, such as boron nitride (BN), are inert.

Another class of nitrogen anions (negatively charged ions) are azides (N₃^-). The molecule of an azide has a linear structure.

Nitrogen forms a variety of oxides. The most prominent ones are nitrogen monoxide (NO) (known more commonly as nitric oxide in biology) and nitrogen dioxide (NO₂). Both types of molecules contain an unpaired electron. The latter shows some tendency to form dimers and is an important component of smog. In addition, nitrogen forms dinitrogen monoxide (or nitrous oxide) (N₂O), which is also known as laughing gas.

The more standard oxides, dinitrogen trioxide (N₂O₃) and dinitrogen pentoxide (N₂O₅), are fairly unstable and explosive. The corresponding acids are nitrous acid (HNO₂) and nitric acid (HNO₃), and 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.

Organic compounds

Nitrogen is also an important element in many organic compounds, in which it is directly bound to one or more carbon atoms. Some examples are given below.

• amines: Between one and four carbon-containing groups ("alkyl" or "aryl" groups) are attached to a nitrogen atom, and the corresponding products are known as primary, secondary, tertiary, and quarternary amines. Many amines are commercially and biologically important compounds.
• amides: In an organic amide, a "carbonyl" group (C=O) is directly attached to a nitrogen atom, and the bond is called an "amide bond." The general chemical formula of an organic amide is R₁(CO)NR₂R₃, where R₂ or R₃ or both may represent a hydrogen atom (each). A peptide or protein is a chain of amino acids that are linked to one another through amide bonds.
• nitro:
• imines:
• enamines:

Common nitrogen functional groups include: amides, nitro groups, imines, and enamines.

Nitrogen compounds of notable economic importance

Molecular nitrogen in the atmosphere is relatively non-reactive due to its strong bond, and the (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 salts of nitric acid 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 all 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 compound's 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 (N₂0) was discovered early in the 19th century to be a partial anesthetic, though it was not used as a surgical anesthetic until later. Called "laughing gas", it was found capable of inducing a state of social disinhibition resembling drunkenness. Other notable nitrogen-containing drugs are drugs derived from plant alkaloids, such as morphine (there exist many alkaloids known to have pharmacological effects; in some cases they appear natural chemical defences of plants against predation). Nitrogen containing drugs include all of the major classes of antibiotics, and organic nitrate drugs like nitroglycerin and nitroprusside which regulate blood pressure and heart action by mimicing the action of nitric oxide.

Precautions

Rapid release of nitrogen gas into an enclosed space can displace oxygen, and therefore represents an asphyxiation hazard. An example occurred shortly before the launch of the first Space Shuttle mission in 1981, when two technicians were killed in a space located in the Shuttle's Mobile Launch Platform that was pressurized with pure nitrogen as a precaution against fire.

When breathed at high partial pressures (more than about 3 atmospheres, encountered a depths below about 30 m in diving) nitrogen begins to act as an anesthetic agent. As such, it can cause nitrogen narcosis, a temporary semi-anesthetized condition of mental impairment similar to that caused by nitrous oxide.

Nitrogen also dissolves in the bloodstream, and rapid decompression (particularly in the case of divers ascending too quickly, or astronauts decompressing too quickly from cabin pressure to spacesuit pressure) can lead to a potentially fatal condition called decompression sickness (formerly known as caisson sickness or more commonly, the "bends"), when nitrogen bubbles form in the bloodstream.

Direct skin contact with liquid nitrogen causes severe frostbite (cryogenic burns) within moments to seconds, 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 ).

See also

• Nutrient
• Nitrogen cycle
• NOx
• Nitrous oxide

References

ISBN links support NWE through referral fees

• 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

External links

• Why high nitrogen density in explosives?
• WebElements.com – Nitrogen
• It's Elemental – Nitrogen
• Schenectady County Community College – Nitrogen
• Nitrogen N2 Properties, Uses, Applications
• Computational Chemistry Wiki
• Handling procedures for liquid nitrogen

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