Nitrogen was discovered independently by several chemists in the late 18th century. In 1772, the Scottish physician and chemist Daniel Rutherford (1749-1819) isolated this gas by removing oxygen and carbon dioxide from air, leaving a gaseous residue he called "noxious air" or "phlogisticated air". Around the same time, Carl Wilhelm Scheele (1742-1786) in Sweden, Henry Cavendish (1731-1810) in England, and Joseph Priestley (1733-1804) conducted similar experiments. In 1790, the French chemist Jean-Antoine Chaptal (1756-1832) proposed the name azote (from Greek a = without and zoe = life), emphasizing that this gas could not sustain life or combustion. The English name "nitrogen" (nitre generator) was introduced in 1790 by Chaptal as well, referring to saltpeter (potassium nitrate).
Nitrogen (symbol N, atomic number 7) is a non-metal of group 15 (pnictogens) in the periodic table, consisting of seven protons, usually seven neutrons (for the most common isotope), and seven electrons. The two stable isotopes are nitrogen-14 \(\,^{14}\mathrm{N}\) (≈ 99.636%) and nitrogen-15 \(\,^{15}\mathrm{N}\) (≈ 0.364%).
At room temperature, nitrogen exists as a diatomic gas (N₂), colorless, odorless, and relatively chemically inert. The N₂ molecule has a very strong triple bond (N≡N) that makes it particularly stable and unreactive under normal conditions. This stability explains why nitrogen gas constitutes about 78% of Earth's atmosphere by volume. N₂ gas has a density of about 1.251 g/L at standard temperature and pressure. The temperature at which the liquid and solid states can coexist (melting point): 63.15 K (−210.00 °C). The temperature at which it transitions from liquid to gas (boiling point): 77.355 K (−195.795 °C).
| Isotope / Notation | Protons (Z) | Neutrons (N) | Atomic mass (u) | Natural abundance | Half-life / Stability | Decay / Remarks |
|---|---|---|---|---|---|---|
| Nitrogen-13 — \(\,^{13}\mathrm{N}\,\) | 7 | 6 | 13.005739 u | Unnatural | 9.965 minutes | Radioactive β\(^+\) decay to \(\,^{13}\mathrm{C}\); used in positron emission tomography (PET). |
| Nitrogen-14 — \(\,^{14}\mathrm{N}\,\) | 7 | 7 | 14.003074 u | ≈ 99.636 % | Stable | Major isotope; basis of all proteins and nucleic acids of terrestrial life. |
| Nitrogen-15 — \(\,^{15}\mathrm{N}\,\) | 7 | 8 | 15.000109 u | ≈ 0.364 % | Stable | Used in NMR spectroscopy, as a tracer in biology, and to study the nitrogen cycle. |
| Nitrogen-16 — \(\,^{16}\mathrm{N}\,\) | 7 | 9 | 16.006102 u | Unnatural | 7.13 s | Radioactive β\(^-\) decay to \(\,^{16}\mathrm{O}\); produced in nuclear reactors. |
| Nitrogen-17 — \(\,^{17}\mathrm{N}\,\) | 7 | 10 | 17.008450 u | Unnatural | 4.173 s | Radioactive β\(^-\); used in nuclear research. |
| Other isotopes — \(\,^{10}\mathrm{N}-\,^{12}\mathrm{N},\,^{18}\mathrm{N}-\,^{25}\mathrm{N}\) | 7 | 3-5, 11-18 | — (resonances) | Unnatural | \(10^{-22}\) — 0.63 s | Very unstable states observed in nuclear physics; decay by particle emission or β radioactivity. |
N.B. :
Electron shells: How electrons organize around the nucleus.
Nitrogen has 7 electrons distributed across two electron shells. Its full electronic configuration is: 1s² 2s² 2p³, or simplified as: [He] 2s² 2p³. This configuration can also be written as: K(2) L(5).
K Shell (n=1): Contains 2 electrons in the 1s sub-shell. This inner shell is complete and highly stable.
L Shell (n=2): Contains 5 electrons distributed as 2s² 2p³. The 2s orbitals are complete, while the 2p orbitals contain only 3 electrons out of 6 possible, with one electron in each of the three 2p orbitals according to Hund's rule. Thus, 3 electrons are missing to reach the stable neon configuration with 8 electrons (octet).
The 5 electrons in the outer shell (2s² 2p³) are the valence electrons of nitrogen. This configuration explains its chemical properties:
By gaining 3 electrons, nitrogen forms the N³⁻ ion (oxidation state -3), present in metal nitrides and ammonia (NH₃), adopting the stable neon configuration [Ne].
Nitrogen can exhibit several positive oxidation states: +1, +2, +3, +4, and +5, observed in its oxides and oxoacids (e.g., nitric acid HNO₃ at +5).
The oxidation state 0 corresponds to dinitrogen N₂, its natural molecular form, where two nitrogen atoms are bonded by an extremely stable triple bond.
Nitrogen's electronic configuration, with 5 valence electrons, classifies it among the pnictogens (Group 15) and gives it intermediate properties between metals and non-metals. This structure grants it characteristic properties: the ability to form three covalent bonds by sharing its three unpaired 2p electrons, the high stability of the N₂ molecule due to its triple bond (very high dissociation energy), and chemical versatility with oxidation states ranging from -3 to +5. Nitrogen forms the nitride ion N³⁻ in some ionic compounds, but this ion is rare and easily hydrolyzed. More commonly, nitrogen shares its electrons in covalent bonds. Dinitrogen N₂ is a colorless, odorless gas, chemically very unreactive at room temperature due to the high stability of its triple bond. This inertness makes dinitrogen an ideal gas for creating protective inert atmospheres.
The importance of nitrogen is paramount: it constitutes about 78% of Earth's atmosphere, making it the most abundant atmospheric gas. It is essential for life as it is a fundamental component of proteins, nucleic acids (DNA, RNA), and many other biomolecules. The nitrogen cycle, which enables its transformation into various chemical forms, is fundamental to ecosystems. Industrially, nitrogen is used for ammonia synthesis via the Haber-Bosch process (basis of nitrogen fertilizers), nitric acid production, explosives manufacturing, and as an inert atmosphere in the food and electronics industries. Liquid nitrogen, obtained by air distillation, serves as a cryogenic fluid for biological preservation and various industrial applications.
Nitrogen has five valence electrons and typically forms three covalent bonds (oxidation state −3 in ammonia NH₃) or can lose its electrons to reach various oxidation states from −3 to +5. The triple bond N≡N in the diatomic N₂ molecule is one of the strongest known chemical bonds (dissociation energy ≈ 945 kJ/mol), making molecular nitrogen very unreactive at room temperature. This inertness is exploited industrially to create protective inert atmospheres.
However, once the triple bond is broken (requiring high temperature, high pressure, or catalysts), nitrogen becomes very reactive. It forms compounds with almost all elements, notably hydrogen (ammonia NH₃, hydrazine N₂H₄), oxygen (nitrogen oxides: NO, NO₂, N₂O, N₂O₃, N₂O₅), halogens (nitrogen trihalides), and many metals (nitrides). Nitrogen compounds exhibit an extraordinary range of properties, from essential fertilizers (nitrates, ammonia) to powerful explosives (TNT, nitroglycerin) to proteins and nucleic acids of life.
The nitrogen cycle is one of the most important biogeochemical cycles on Earth. Although N₂ is abundant in the atmosphere, most organisms cannot use it directly. Biological nitrogen fixation by certain bacteria (symbiotic or free-living) converts N₂ into ammonia, which can then be assimilated by plants. Other bacteria carry out nitrification (conversion to nitrites then nitrates) and denitrification (return of nitrogen to the atmosphere). Humanity has profoundly disrupted this natural cycle with the massive industrial production of nitrogen fertilizers (Haber-Bosch process).
Nitrogen is the fifth most abundant element in the observable universe (after hydrogen, helium, oxygen, and carbon) and plays an important role in the chemical evolution of galaxies. Unlike primordial elements, nitrogen is entirely produced by stellar nucleosynthesis.
The main pathway for nitrogen production in stars is the CNO cycle (carbon-nitrogen-oxygen), where nitrogen appears as a catalytic intermediate in the fusion of hydrogen into helium. In massive stars, this cycle dominates energy production. Nitrogen-14 is produced mainly in intermediate-mass stars (2-8 solar masses) during the AGB (asymptotic giant branch) phase, where it is synthesized from carbon via the CN cycle. These stars then enrich the interstellar medium with nitrogen through their powerful stellar winds.
Nitrogen can also be produced by spallation reactions in the interstellar medium (fragmentation of heavier atoms by cosmic rays), although this contribution is minor compared to stellar nucleosynthesis.
In the interstellar medium, nitrogen exists in several forms: atomic (N, N⁺), molecular (N₂, CN, HCN, NH₃, and many other complex nitrogen-containing molecules). Nitrogen-containing molecules are important tracers of physical and chemical conditions in dense molecular clouds where stars form. Dinitrogen (N₂) is difficult to detect directly in space due to the lack of a permanent dipole moment, but its abundance can be inferred indirectly via other nitrogen species.
The isotopic ratio ¹⁴N/¹⁵N varies considerably in the universe and provides valuable information about nucleosynthesis and mixing processes in stars. This ratio, measured in meteorites, comets, planetary atmospheres, and the interstellar medium, reveals the complex history of matter recycling in our galaxy. The solar system has a ¹⁴N/¹⁵N ratio of about 272, but this ratio can vary significantly depending on the sources and objects observed.
In planetary atmospheres, nitrogen plays a major role. On Earth, it constitutes 78% of the atmosphere and is essential to life. On Titan (Saturn's moon), the atmosphere is 98% nitrogen. The study of nitrogen in atmospheres and its chemical cycle on different bodies in the solar system and on potentially habitable exoplanets is crucial for understanding planetary evolution and the search for extraterrestrial life.
N.B.:
The Haber-Bosch process, developed in the early 20th century, revolutionized global agriculture by enabling the industrial synthesis of ammonia from atmospheric nitrogen and hydrogen under high pressure and temperature, with an iron catalyst. This innovation allowed the mass production of nitrogen fertilizers, significantly increasing agricultural yields and feeding an exponentially growing world population. It is estimated that more than half of the nitrogen in human proteins today comes from nitrogen artificially fixed by this process. However, this massive industrial nitrogen fixation (about 150 million tons per year) now far exceeds natural biological fixation and has created major environmental problems: water pollution by nitrates, eutrophication of aquatic ecosystems, emissions of nitrous oxide (N₂O, a potent greenhouse gas), and profound disruption of the natural nitrogen cycle on a planetary scale. The challenge of the 21st century is to maintain food production while restoring a more balanced and sustainable nitrogen cycle.
