Nitrogen is a nonmetallic chemical element with the symbol N and atomic number 7. Under ordinary conditions, elemental nitrogen exists as a colorless, odorless gas composed of N₂ molecules. It constitutes approximately 78% of the Earth’s atmosphere by volume. Nitrogen compounds are essential constituents of organisms and major materials in agriculture and chemical manufacturing. (pubchem.ncbi.nlm.nih.gov)
Atomic structure and physical properties
Nitrogen occupies group 15 and period 2 of the periodic table. A neutral nitrogen atom has seven electrons, with the ground-state electron configuration 1s²2s²2p³. Its five outer-shell electrons account for its diverse bonding chemistry. The commonly quoted atomic weight is 14.007, although the standard atomic-weight interval is [14.00643, 14.00728], reflecting natural differences in isotopic composition. (pubchem.ncbi.nlm.nih.gov)
At atmospheric pressure, nitrogen melts at approximately −210°C and boils at approximately −195.8°C, or 77.4 K. Its low boiling point makes liquid nitrogen useful in cryogenics. Liquid nitrogen requires insulated storage because heat entering from the surroundings continuously causes evaporation; industrial storage systems commonly use vacuum-insulated vessels. (periodic-table.rsc.org)
Natural nitrogen has two stable isotopes, nitrogen-14 and nitrogen-15. Nitrogen-14 accounts for roughly 99.6% of natural nitrogen. Differences in isotope ratios arise during biological and chemical transformations, making those ratios useful for investigating nitrogen sources and movement through environmental systems. Atmospheric N₂ serves as the reference standard for nitrogen-isotope measurements. (ciaaw.org)
Chemical behavior
In the N₂ molecule, two nitrogen atoms share a triple covalent bond. This strong bond, together with the molecule’s nonpolarity, makes molecular nitrogen relatively unreactive under ordinary conditions. Its stability presents substantial thermodynamic and kinetic obstacles to converting atmospheric nitrogen into other compounds. Nitrogen gas is therefore useful as a protective atmosphere, although “inert” describes its behavior in particular applications rather than an absolute inability to react. (pubmed.ncbi.nlm.nih.gov)
Nitrogen exhibits oxidation states ranging from −3 to +5 in its compounds. Familiar examples include −3 in ammonia, NH₃, and +5 in nitrate, NO₃⁻. Other environmentally important forms include ammonium, NH₄⁺, nitrite, NO₂⁻, and nitrogen oxides. The distinction between elemental nitrogen and these combined forms is fundamental: the abundance of atmospheric N₂ does not mean that nitrogen is readily available for biological use. (pubchem.ncbi.nlm.nih.gov)
Discovery and naming
The discovery of nitrogen is conventionally credited to Daniel Rutherford, who described it in his doctoral thesis in September 1772. His experiments distinguished the residual gas remaining after oxygen and carbon dioxide had been removed from air. Henry Cavendish and Joseph Priestley had also obtained nitrogen earlier, but Rutherford’s published account established his association with its discovery. (periodic-table.rsc.org)
Antoine Lavoisier used the name azote, derived from Greek words meaning “without life,” because the gas did not support respiration. The name nitrogen means “nitre-former,” referring to its presence in nitrate compounds such as saltpetre. Forms of azote remain in use in several languages. (periodic-table.rsc.org)
Biological importance and the nitrogen cycle
Nitrogen is present in amino acids, the building blocks of proteins, and in the nitrogen-containing bases of DNA and RNA. Plants incorporate available nitrogen compounds into these biological materials; animals obtain nitrogen by consuming other organisms or their products. Nitrogen availability consequently influences biological growth even where atmospheric nitrogen is abundant. (periodic-table.rsc.org)
The nitrogen cycle transfers nitrogen among the atmosphere, organisms, soils, and water. Nitrogen fixation converts N₂ into chemically combined forms. Specialized bacteria perform biological fixation, producing ammonium that can enter biological pathways. Lightning also converts atmospheric nitrogen into compounds that subsequently reach the surface. Most plants cannot use N₂ directly and instead absorb nitrogen chiefly as nitrate or ammonium. (gml.noaa.gov)
Decomposers return nitrogen from dead organisms and wastes to ammonium. During nitrification, microorganisms convert ammonium through nitrite to nitrate. Denitrification converts nitrate into gaseous products, returning nitrogen to the atmosphere. These transformations link nutrient availability to microbial activity and local environmental conditions. (gml.noaa.gov)
Industrial production and uses
Industrial nitrogen is separated from air rather than extracted from mineral deposits. Large plants employ cryogenic distillation, separating the components of liquefied air. Other systems use pressure-swing adsorption, in which carbon molecular sieves preferentially adsorb oxygen under pressure, or membranes that separate gases through differences in permeation rates. The appropriate production method depends on required purity and flow. (airproducts.com)
A major chemical use is ammonia manufacture by the Haber process, which reacts nitrogen with hydrogen. Ammonia supplies nitrogen for fertilizers and provides feedstock for other industrial chemicals. Nitrogen gas also displaces oxygen in food packaging, electronics production, and metal processing. Liquid nitrogen supplies rapid cooling for food freezing, biological-sample preservation, and materials processing. (periodic-table.rsc.org)
Environmental effects and hazards
Agriculture, wastewater discharges, and fuel combustion alter the quantity and distribution of reactive nitrogen. Excess nutrients entering waters can cause eutrophication, promoting algal growth and oxygen depletion. Agricultural soils and other nitrogen-cycle processes also release nitrous oxide, N₂O, a greenhouse gas distinct from atmospheric N₂. (pubs.usgs.gov)
Nitrogen gas can cause asphyxiation by displacing oxygen, particularly in enclosed spaces. Liquid nitrogen introduces additional hazards: extreme cold can cause freezing injuries, and evaporation in a sealed vessel can generate dangerous pressure. These physical hazards exist despite nitrogen’s comparatively low chemical reactivity. (cdc.gov)