A | B | C | D | E | F | G | H | CH | I | J | K | L | M | N | O | P | Q | R | S | T | U | V | W | X | Y | Z | 0 | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9
 Liquid nitrogen (N2 at below −196 °C) | |||||||||||||||||||||||||||||||||
| Nitrogen | |||||||||||||||||||||||||||||||||
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| Allotropes | see § Allotropes | ||||||||||||||||||||||||||||||||
| Appearance | colorless gas, liquid or solid | ||||||||||||||||||||||||||||||||
| Standard atomic weight Ar°(N) | |||||||||||||||||||||||||||||||||
| Nitrogen in the periodic table | |||||||||||||||||||||||||||||||||
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| 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[3] K (−209.86[3] °C, −345.75[3] °F) | ||||||||||||||||||||||||||||||||
| Boiling point | (N2) 77.355 K (−195.795 °C, −320.431 °F) | ||||||||||||||||||||||||||||||||
| Density (at STP) | 1.2506 g/L[4] 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 vaporisation | (N2) 5.57 kJ/mol | ||||||||||||||||||||||||||||||||
| Molar heat capacity | (N2) 29.124 J/(mol·K) | ||||||||||||||||||||||||||||||||
Vapour pressure
| |||||||||||||||||||||||||||||||||
| Atomic properties | |||||||||||||||||||||||||||||||||
| Oxidation states | −3, −2, −1, 0,[5] +1, +2, +3, +4, +5 (a strongly acidic oxide) | ||||||||||||||||||||||||||||||||
| Electronegativity | Pauling scale: 3.04 | ||||||||||||||||||||||||||||||||
| Ionisation energies |
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| Covalent radius | 71±1 pm | ||||||||||||||||||||||||||||||||
| Van der Waals radius | 155 pm | ||||||||||||||||||||||||||||||||
 | |||||||||||||||||||||||||||||||||
| Other properties | |||||||||||||||||||||||||||||||||
| Natural occurrence | primordial | ||||||||||||||||||||||||||||||||
| Crystal structure | hexagonal | ||||||||||||||||||||||||||||||||
| 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 | |||||||||||||||||||||||||||||||||
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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 colorless and odorless diatomic gas. N2 forms about 78% of Earth's atmosphere, making it the most abundant uncombined element 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 nitrogène 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 Ancient Greek: ἀζωτικός "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),[6] 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 metabolizing 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
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.[7]
The discovery of nitrogen is attributed to the Scottish physician Daniel Rutherford in 1772, who called it noxious air.[8][9] 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.[10] 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,[11] Henry Cavendish,[12] and Joseph Priestley,[13] 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 άζωτικός (azotikos), "no life", due to it being asphyxiant.[14][15] 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".[7]
The English word nitrogen (1794) entered the language from the French nitrogène, coined in 1790 by French chemist Jean-Antoine Chaptal (1756–1832),[16] from the French nitre (potassium nitrate, also called saltpeter) 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, niter had been confused with Egyptian "natron" (sodium carbonate) – called νίτρον (nitron) in Greek – which, despite the name, contained no nitrate.[17]
The earliest military, industrial, and agricultural applications of nitrogen compounds used saltpeter (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.[18] The "whirling cloud of brilliant yellow light" produced by his apparatus reacted with mercury to produce explosive mercury nitride.[19]
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.[20] 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.[21][22]
Properties
Atomic
A nitrogen atom has seven electrons. In the ground state, they are arranged in the electron configuration 1s2
2s2
2p1
x2p1
y2p1
z. 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.)[23] 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.[23] 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 16.920 MJ·mol−1. Due to these very high figures, nitrogen has no simple cationic chemistry.[24]
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.[25]
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.[26]
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.[27]
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.[28]
Nitrogen shares with both its horizontal neighbours a preference for forming multiple bonds, typically with carbon, oxygen, or other nitrogen atoms, through pπ–pπ interactions.[26] 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.[26] Setting aside their differences, nitrogen and phosphorus form an extensive series of compounds with one another; these have chain, ring, and cage structures.[29]
Table of thermal and physical properties of nitrogen (N2) at atmospheric pressure:[30][31]
| 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 | 6.88×10−6 | 2.00×10−6 | 0.00958 | 2.60×10−6 | 0.768 |
| 150 | 2.2594 | 1.05 | 1.01×10−5 | 4.45×10−6 | 0.0139 | 5.86×10−6 | 0.759 |
| 200 | 1.7108 | 1.0429 | 1.29×10−5 | 7.57×10−6 | 0.01824 | 1.02×10−5 | 0.747 |
| 300 | 1.1421 | 1.0408 | 1.78×10−5 | 1.56×10−5 | 0.0262 | 2.20×10−5 | 0.713 |
| 400 | 0.8538 | 1.0459 | 2.20×10−5 | 2.57×10−5 | 0.03335 | 3.73×10−5 | 0.691 |
| 500 | 0.6824 | 1.0555 | 2.57×10−5 | 3.77×10−5 | 0.03984 | 5.53×10−5 | 0.684 |
| 600 | 0.5687 | 1.0756 | 2.91×10−5 | 5.12×10−5 | 0.0458 | 7.49×10−5 | 0.686 |
| 700 | 0.4934 | 1.0969 | 3.21×10−5 | 6.67×10−5 | 0.05123 | 9.47×10−5 | 0.691 |
| 800 | 0.4277 | 1.1225 | 3.48×10−5 | 8.15×10−5 | 0.05609 | 1.17×10−4 | 0.7 |
| 900 | 0.3796 | 1.1464 | 3.75×10−5 | 9.11×10−5 | 0.0607 | 1.39×10−4 | 0.711 |