Nitrogen is the fifth most abundant element in the observable universe (after hydrogen, helium, oxygen, and carbon) and plays a crucial 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 mainly produced in intermediate-mass stars (2-8 solar masses) during the AGB phase (asymptotic giant branch), where it is synthesized from carbon via the CN cycle. These stars then enrich the interstellar medium with nitrogen through their powerful stellar winds.
In the interstellar medium, nitrogen exists in several forms: atomic (N, N⁺), molecular (N₂, CN, HCN, NH₃, and many other complex nitrogen 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 indirectly deduced via other nitrogen species.
The isotopic ratio ¹⁴N/¹⁵N varies significantly 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 makes up 78% of the atmosphere and is essential for life. On Titan (Saturn's moon), the atmosphere is 98% nitrogen. The study of atmospheric nitrogen 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.
Nitrogen was independently discovered by several chemists in the late 18th century. In 1772, Scottish physician and chemist Daniel Rutherford (1749-1819) isolated this gas by removing oxygen and carbon dioxide from the air, leaving a gaseous residue he called "vitiated 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, French chemist Jean-Antoine Chaptal (1756-1832) proposed the name azote (from Greek a = without and zoe = life), emphasizing that this gas could not support life or combustion. The English name "nitrogen" (niter generator) was introduced in 1790 by Chaptal, referring to saltpeter (potassium nitrate).
N.B.:
The Haber-Bosch process (early 20th century) enables the industrial fixation of atmospheric nitrogen into ammonia. It has allowed the mass production of fertilizers, supporting global food security—more than half of the nitrogen in our proteins comes from it. However, this artificial fixation (∼150 Mt/year) now exceeds natural fixation, causing water pollution (nitrates), N₂O emissions (greenhouse gas), and a major disruption of the natural nitrogen cycle. The current challenge is to reconcile food production with the restoration of a sustainable nitrogen cycle.
Nitrogen (symbol N, atomic number 7) is a non-metal of group 15 (pnictogens) of 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), making it particularly stable and unreactive under normal conditions. This stability explains why nitrogen gas makes up 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 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 | Not natural | 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 | Not natural | 7.13 s | Radioactive β\(^-\) decay to \(\,^{16}\mathrm{O}\) ; produced in nuclear reactors. |
| Nitrogen-17 — \(\,^{17}\mathrm{N}\,\) | 7 | 10 | 17.008450 u | Not natural | 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) | Not natural | \(10^{-22}\) — 0.63 s | Very unstable states observed in nuclear physics; decay by particle emission or β radioactivity. |
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Electron shells: How electrons organize around the nucleus.
Nitrogen has 7 electrons distributed in two electron shells. Its full electronic configuration is: 1s² 2s² 2p³, or simplified: [He] 2s² 2p³. This configuration can also be written as: K(2) L(5).
K shell (n=1): contains 2 electrons in the 1s subshell. This inner shell is complete and very 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).
Nitrogen, from group 15 (pnictogens), has 5 valence electrons (2s² 2p³). This configuration explains its versatility: it forms three covalent bonds (state -3 as in NH₃), but can also reach positive oxidation states up to +5 (HNO₃). The N₂ molecule, with its exceptionally stable triple bond, corresponds to the 0 state.
Dinitrogen (N₂) makes up 78% of Earth's atmosphere. Its inertness at room temperature makes it useful for inert atmospheres. However, once activated, nitrogen becomes a key element of life (proteins, DNA), agriculture (fertilizers via the Haber-Bosch process), and industry (explosives, nitric acid). Its liquefaction also enables cryogenic 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 industrially exploited 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, especially 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), as well as 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).
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