page_title: Nitrogen
section_title: Isotopes
hierarchical_section_title: Nitrogen / Properties / Isotopes
caption_attribution_description: English: Chart of nuclides for Carbon to Fluorine &#160;&#160; proton emission &#160;&#160; beta+/EC decay &#160;&#160; stable nuclide &#160;&#160; beta- decay &#160;&#160; neutron emission
context_page_description: Nitrogen is the chemical element with the symbol N and atomic number 7. It was first discovered and isolated by Scottish physician Daniel Rutherford in 1772. Although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is generally accorded the credit because his work was published first. 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 Greek ἀζωτικός "no life", as it is an asphyxiant gas; this name is instead used in many languages, such as French, Italian, Russian, Romanian and Turkish, and appears in the English names of some nitrogen compounds such as hydrazine, azides and azo compounds.
  Nitrogen is the lightest member of group 15 of the periodic table, often called the pnictogens. It is a common element in the universe, estimated at about seventh in total abundance in the Milky Way and the Solar System. At standard temperature and pressure, two atoms of the element bind to form dinitrogen, a colourless and odorless diatomic gas with the formula N₂.
context_section_description: Nitrogen has two stable isotopes: ¹⁴N and ¹⁵N. The first is much more common, making up 99.634% of natural nitrogen, and the second (which is slightly heavier) makes up the remaining 0.366%. This leads to an atomic weight of around 14.007 u. Both of these stable isotopes are produced in the CNO cycle in stars, but ¹⁴N is more common as its neutron capture is the rate-limiting step. ¹⁴N is one of the five stable odd–odd nuclides (a nuclide having an odd number of protons and neutrons); the other four are ²H, ⁶Li, ¹⁰B, and ¹⁸⁰ᵐTa.
  The relative abundance of ¹⁴N and ¹⁵N is practically constant in the atmosphere but can vary elsewhere, due to natural isotopic fractionation from biological redox reactions and the evaporation of natural ammonia or nitric acid. 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.
  The heavy isotope ¹⁵N was first discovered by S. M. Naudé in 1929, soon after heavy isotopes of the neighbouring elements oxygen and carbon were discovered. It presents one of the lowest thermal neutron capture cross-sections of all isotopes. It is frequently used in nuclear magnetic resonance (NMR) spectroscopy to determine the structures of nitrogen-containing molecules, due to its fractional nuclear spin of one-half, which offers advantages for NMR such as narrower line width. ¹⁴N, though also theoretically usable, has an integer nuclear spin of one and thus has a quadrupole moment that leads to wider and less useful spectra. ¹⁵N NMR nevertheless has complications not encountered in the more common ¹H and ¹³C NMR spectroscopy. The low natural abundance of ¹⁵N (0.36%) significantly reduces sensitivity, a problem which is only exacerbated by its low gyromagnetic ratio, (only 10.14% that of ¹H). As a result, the signal-to-noise ratio for ¹H is about 300 times as much as that for ¹⁵N at the same magnetic field strength. This may be somewhat alleviated by isotopic enrichment of ¹⁵N by chemical exchange or fractional distillation. ¹⁵N-enriched compounds have the advantage that under standard conditions, they do not undergo chemical exchange of their nitrogen atoms with atmospheric nitrogen, unlike compounds with labelled hydrogen, carbon, and oxygen isotopes that must be kept away from the atmosphere. The ¹⁵N:¹⁴N ratio is commonly used in stable isotope analysis in the fields of geochemistry, hydrology, paleoclimatology and paleoceanography, where it is called δ¹⁵N.
  Of the ten other isotopes produced synthetically, ranging from ¹²N to ²³N, ¹³N has a half-life of ten minutes and the remaining isotopes have half-lives on the order of seconds (¹⁶N and ¹⁷N) or milliseconds. No other nitrogen isotopes are possible as they would fall outside the nuclear drip lines, leaking out a proton or neutron. Given the half-life difference, ¹³N is the most important nitrogen radioisotope, being relatively long-lived enough to use in positron emission tomography (PET), although its half-life is still short and thus it must be produced at the venue of the PET, for example in a cyclotron via proton bombardment of ¹⁶O producing ¹³N and an alpha particle.
  The radioisotope ¹⁶N is the dominant radionuclide in the coolant of pressurised water reactors or boiling water reactors during normal operation, and thus it is a sensitive and immediate indicator of leaks from the primary coolant system to the secondary steam cycle, and is the primary means of detection for such leaks. It is produced from ¹⁶O (in water) via an (n,p) reaction in which the ¹⁶O atom captures a neutron and expels a proton. It has a short half-life of about 7.1 s, but during its decay back to ¹⁶O produces high-energy gamma radiation (5 to 7 MeV). Because of this, access to the primary coolant piping in a pressurised water reactor must be restricted during reactor power operation.
caption: Table of nuclides (Segrè chart) from carbon to fluorine (including nitrogen). Orange indicates proton emission (nuclides outside the proton drip line); pink for positron emission (inverse beta decay); black for stable nuclides; blue for electron emission (beta decay); and violet for neutron emission (nuclides outside the neutron drip line). Proton number increases going up the vertical axis and neutron number going to the right on the horizontal axis.
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