Nitrogen
The main chemical characteristic of nitrogen is the contrast between the unreactive dinitrogen molecule and the very large number of reactive nitrogen compounds.
Nitrogen is a chemical element with the symbol N and atomic number 7. It is a nonmetal and the lightest member of group 15 of the periodic table, the pnictogens. At standard temperature and pressure it exists as the colorless, odorless diatomic gas dinitrogen (N2), which makes up about 78 percent of the atmosphere by volume (78.08 percent of dry air, or about 75.5 percent by mass). Nitrogen is essential to life, because it is a component of amino acids, proteins, nucleic acids, chlorophyll, and many other biomolecules. It is also the basis of fertilizers and explosives. Nitrogen is the seventh most abundant element in the Milky Way and the solar system, and it was produced mainly in stars by hydrogen burning in the CNO cycle, not by carbon fusion.
The main chemical characteristic of nitrogen is the contrast between the unreactive dinitrogen molecule and the very large number of reactive nitrogen compounds. N2 is chemically inert under ordinary conditions, because it has a very strong triple bond, but nitrogen compounds in which the element has an oxidation state from −3 to +5 take part in nearly all biological and industrial chemistry.
History: How Was Nitrogen Discovered and Fixed Industrially?
Pneumatic Chemistry and the Haber-Bosch Breakthrough
Nitrogen was identified in 1772 by the Scottish physician Daniel Rutherford, who found that the gas left after combustion and removal of carbon dioxide from air does not support combustion or life. Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley studied the gas at about the same time, and Antoine Lavoisier called it azote ("without life"). The name "nitrogen", from the Greek for "niter-forming", was proposed by Jean-Antoine Chaptal in 1790, from its presence in potassium nitrate. The industrial fixation of nitrogen by the Haber-Bosch process was developed by Fritz Haber and Carl Bosch in 1909 to 1913. Both received Nobel Prizes in Chemistry for this work (Haber in 1918 and Bosch in 1931).
Atomic Structure and Isotopes: How Do Valence Orbitals and Nuclear Spin Define Nitrogen?
Electron Configuration (1s2 2s2 2p3) and Oxidation States
A nitrogen atom has 7 protons and 7 electrons, with the electron configuration 1s2 2s2 2p3. The valence shell has five electrons, three of them unpaired, so that the atom needs three more electrons to complete an octet. Nitrogen has a high electronegativity (3.04 on the Pauling scale), exceeded among the elements only by oxygen, fluorine, and (nearly) chlorine. It forms covalent bonds with most elements, and it can take the oxidation states −3 (as in ammonia and nitrides), −2 (hydrazine), −1 (hydroxylamine), 0 (N2), +1 (nitrous oxide), +2 (nitric oxide), +3 (nitrites), +4 (nitrogen dioxide), and +5 (nitrates and nitric acid).
Stable (14N, 15N) and Radiogenic Isotopes
Naturally occurring nitrogen consists of two stable isotopes:
- Nitrogen-14 (14N) makes up 99.636 percent of natural nitrogen. It has 7 neutrons, and a nuclear spin of 1.
- Nitrogen-15 (15N) makes up 0.364 percent. It has 8 neutrons and a nuclear spin of 1/2, which makes it useful in nuclear magnetic resonance spectroscopy of biomolecules, usually in enriched form.
Radioactive isotopes are also known. Nitrogen-13 has a half-life of about 10 minutes and is used in positron emission tomography (for example as 13N-ammonia for imaging of the myocardium). In the upper atmosphere, cosmic-ray neutrons convert 14N to carbon-14 by the reaction 14N + n → 14C + p, which is the origin of radiocarbon dating.
Isotopic Fractionation
Biological and chemical processes discriminate slightly between the two stable isotopes, because bonds of the lighter isotope are broken somewhat faster, so that they are enriched in the reaction products. The nitrogen isotope composition of a sample is given as δ15N, in parts per thousand (per mil) relative to atmospheric N2 as the standard:
Atmospheric N2 has δ15N = 0 by definition, with a ratio of about 0.00368. Consumers are typically enriched in 15N relative to their food by about 3 to 4 per mil at each trophic level, because the light isotope is preferentially excreted. A consumer three levels above a plant with δ15N of 0 would therefore have a value of about +10 per mil, which is the basis for the use of nitrogen isotopes in studies of food webs and of the diet of animals and of ancient people. The isotope ratio of nitrates and of sediments is also used to study the cycle of nitrogen in the oceans and its changes in the past, for example from ice cores and marine sediments, and it is used to trace sources of nitrogen pollution. Fractionation depends on the process (for example, denitrification discriminates strongly), so that interpretation requires care.
Physical Properties: What Are the Phase Transitions and Cryogenic Characteristics of Dinitrogen?
Boiling Point, Freezing Point, and Critical Parameters
N2 is a nonpolar molecule with a very low solubility in water. It condenses at 77.4 K (−195.8 °C, −320.4 °F) at atmospheric pressure and freezes at 63.2 K (−210 °C). The critical temperature is 126.2 K and the critical pressure is about 3.4 MPa. Liquid nitrogen is a colorless liquid with a density of about 0.81 g·cm−3 at its boiling point.
Chemical Bonding and the Inertness of Dinitrogen: Why Is the N<sub>2</sub> Triple Bond So Stable?
Bond Order Anomaly and HOMO-LUMO Gap
The N2 molecule is held together by a triple bond, with one σ bond and two π bonds, and the bond length is about 109.8 pm. The bond dissociation energy is 945 kJ·mol−1, one of the highest of any diatomic molecule (comparable with that of carbon monoxide, 1,072 kJ·mol−1). Nitrogen forms bonds of different orders with itself, whose strengths are approximately 160 kJ·mol−1 for a single bond (N–N), 418 kJ·mol−1 for a double bond (N=N), and 945 kJ·mol−1 for the triple bond. The triple bond is much more than three times as strong as the single bond (3 × 160 = 480 kJ·mol−1), which explains two apparently opposite properties: N2 is very stable, and compounds with N–N and N–O single bonds are often unstable and energy-rich.
Kinetic Metastability of the Atmosphere
The inertness of N2 has several causes in addition to the bond strength. The molecule has no dipole moment and is not easily polarized, its highest occupied orbital is low in energy (the first ionization energy is 15.6 eV, much higher than that of O2, 12.1 eV), and its lowest unoccupied orbital is high, so that it is a poor donor and acceptor of electrons, and reactions have high activation energies. The reaction of N2 + O2 → 2 NO, is endothermic (+180 kJ per mole of the reaction as written) and requires high temperatures, such as in lightning, in internal combustion engines, and in furnaces. This is why air does not burn, and why NOx forms in combustion. The reactions of N2 with oxygen and water to form nitrate are thermodynamically favorable, which shows that the air and the ocean are not in chemical equilibrium and that the stability of N2 is kinetic. N2 can be activated by transition metal complexes, and the first dinitrogen complex was reported in 1965.
Compounds: How Does Nitrogen Span Oxidation States from −3 to +5?
| Oxidation State | Representative Chemical Species | Formula & Molecular Geometry | Standard Δ_G_f° (kJ/mol, 298 K) | Primary Biogeochemical or Industrial Role |
|---|---|---|---|---|
| −3 | Ammonia / Ammonium ion | NH3 (Trigonal pyramidal, 107.8°) / NH4+ | −16.4 (NH3(g)); −79.3 (NH4+(aq)) | Primary product of nitrogen fixation and fertilizer feedstock |
| −2 | Hydrazine | N2H4 (Skewed gauche, N–N single bond) | +149.3 (liquid) | Hypergolic rocket propellant and anammox intermediate |
| 0 | Dinitrogen gas | N≡N (Linear diatomic, d = 109.8 pm) | 0.0 (Reference element) | 78.08% of dry atmosphere (~3.9 × 1018 kg reservoir) |
| +1 | Nitrous oxide | N2O (Linear asymmetric, N≡N+–O−) | +104.2 (gas) | Potent greenhouse gas (GWP100 ≈ 273) and ozone-depleting agent |
| +2 | Nitric oxide | •NO (Free radical diatomic, bond order 2.5) | +87.6 (gas) | Mammalian vasodilator signaling molecule and combustion NOx |
| +3 | Nitrite ion / Nitrous acid | NO2− (Bent, 115.4°, sp2 resonance) | −32.2 (NO2−(aq)) | Intermediate in soil nitrification (AOB/AOA) and denitrification |
| +5 | Nitrate ion / Nitric acid | NO3− (Trigonal planar, 120°) / HNO3 | −111.3 (NO3−(aq)); −80.7 (HNO3(l)) | Dominant oxidized bioavailable nitrogen pool in soils and oceans |
Hydrides, Oxides, Oxyacids, and Organic Nitrogen
- Hydrides. Ammonia (NH3) is a colorless, pungent gas that is the main source of fixed nitrogen, and hydrazine (N2H4) is used as a rocket propellant.
- Oxides. Nitrous oxide (N2O) is used as an anesthetic and is a greenhouse gas and an ozone-depleting substance. Nitric oxide (NO) is a signaling molecule in physiology and an intermediate in combustion. Nitrogen dioxide (NO2) is a brown gas and an air pollutant. Together, NO and NO2 are called NOx.
- Oxyacids and salts. Nitric acid (HNO3) is made by the Ostwald process, in which ammonia is oxidized to NO (4 NH3 + 5 O2 → 4 NO + 6 H2O) and then to nitric acid. Nitrates are used as fertilizers and as oxidizers, and nitrites as food preservatives.
- Organic compounds. Amines, amides, nitriles, nitro compounds, and the heterocycles of nucleic acids and many drugs.
- Nitrides and azides. Metal nitrides, such as titanium nitride and boron nitride, are hard and refractory. Azides, such as sodium azide, decompose rapidly and release N2, which is the basis of their use in some airbag inflators.
Nitrogen in Biology and the Nitrogen Cycle: How Is Atmospheric N<sub>2</sub> Fixed and Recycled?
Most organisms cannot use N2, because they lack nitrogenase. It is used only by the diazotrophs, which are certain bacteria and archaea, and the rest of the biosphere depends on reduced nitrogen that they supply or that is made by other processes. Nitrogen is a limiting nutrient in many ecosystems.
Fixation
Dinitrogen is converted to ammonia by nitrogen fixation, which occurs in three main ways:
- Biological fixation by diazotrophs, which include free-living bacteria, cyanobacteria (such as Trichodesmium in the oceans), and symbionts such as Rhizobium in the root nodules of legumes. The enzyme nitrogenase has a metal cofactor (an iron-molybdenum cofactor in the most common form) and requires 16 ATP per N2 in the idealized reaction, and it is damaged by oxygen:
- Abiotic fixation by lightning and by other high-energy events produces NO and then nitrate in small amounts, of the order of a few Tg of nitrogen per year.
- Industrial fixation by the Haber-Bosch process, which combines N2 and hydrogen at high temperatures and pressures over an iron-based catalyst:
The Biogeochemical Cycle
In the nitrogen cycle, fixed nitrogen is assimilated into organic matter, returned to ammonium by decomposition (ammonification), oxidized to nitrite and nitrate by nitrifying bacteria and archaea (nitrification, which includes the "comammox" bacteria that perform both steps), and finally converted back to N2 by denitrification (in which nitrate is used as an electron acceptor under low oxygen) and by anammox (the anaerobic oxidation of ammonium with nitrite). The nitrogen in the atmosphere (about 3.9 × 1018 kg) is a huge reservoir compared with the other pools, so that the lifetime of an N2 molecule in the atmosphere, with respect to the natural fixation of about 200 Tg per year, is of the order of 107 years.
Environmental Effects: How Does Anthropogenic Nitrogen Fixation Perturb Planetary Boundaries?
Eutrophication, Nitrous Oxide Radiative Forcing, and Air Quality
Humans now create reactive nitrogen at a rate of the order of 200 Tg of nitrogen per year (including Haber-Bosch, the cultivation of legumes, and combustion), which exceeds natural terrestrial fixation, although the estimates of natural fluxes are uncertain. Part of the fertilizer that is applied is not taken up by crops. The excess has several effects:
- Eutrophication. Runoff of nitrate and ammonium stimulates the growth of algae in rivers, lakes, and coastal waters, and the decomposition of the algae depletes oxygen and forms "dead zones", as in the Gulf of Mexico and the Baltic Sea.
- Greenhouse gas emissions. Nitrous oxide is produced by nitrification and denitrification in soils and waters. It has a lifetime of about 110 years, a global warming potential about 270 times that of CO2 over 100 years, and a concentration of about 337 ppb, up from about 270 ppb before the industrial period. It is also the largest emission that is currently depleting stratospheric ozone.
- Air pollution. NOx from combustion contributes to tropospheric ozone, acid rain, and fine particles, and ammonia from agriculture forms particulate ammonium salts.
- Groundwater. High nitrate in drinking water can cause methemoglobinemia in infants and is regulated.
Nitrogen is one of the processes in the planetary boundaries framework, and a large share of the nitrogen use in agriculture could be reduced by better management of fertilizer, crops, and livestock manure.
Production and Uses: How Is Elemental Nitrogen Separated and Applied in Industry?
Production of Nitrogen Gas
Nitrogen gas is made on a large scale by fractional distillation of liquid air, in which nitrogen, with a lower boiling point (77.4 K) than oxygen (90.2 K) and argon (87.3 K), is separated, and it can be made by pressure swing adsorption or by membranes, which are used for on-site generation at lower purity.
Industrial, Cryogenic, and Energetic Applications
- Inert atmospheres. Nitrogen is used to displace oxygen and moisture in chemical, pharmaceutical, and electronics manufacturing, in the blanketing of tanks to prevent fires and explosions, in the packaging of foods (modified atmospheres, which slow oxidation and the growth of aerobic microbes, but do not prevent the growth of anaerobic microbes), and in the welding and heat treating of metals.
- Cryogenics. Liquid nitrogen is used as a refrigerant for the storage of biological samples (cells, semen, embryos, and blood products), in cryosurgery, for the shrink-fitting of metal parts, and in physics laboratories. One liter of liquid nitrogen gives about 694 liters of gas at room temperature, so that a closed container of it can burst.
- Fertilizers and chemicals. Ammonia and its derivatives (urea, ammonium nitrate, and ammonium phosphates) are the largest use of fixed nitrogen, and ammonia is also the starting material for nitric acid, nylon, and many other chemicals.
- Explosives. Many explosives contain nitrogen, including nitroglycerin, TNT, and ammonium nitrate. They release much energy not because the bonds in the starting compounds are particularly weak (the N–N and N–O single bonds are fairly weak, and the compounds are metastable), but because the products include N2, with its very strong triple bond, and CO2 and H2O, so that the reaction is strongly exothermic, and it produces a large volume of hot gas.
- Other uses. Nitrous oxide is used for anesthesia and as an oxidizer in rocketry, and in whipped cream dispensers. Nitrogen is used in tires, in lasers and in other equipment, and compounds are used in medicine (such as nitroglycerin for angina, which acts by releasing nitric oxide, and many drugs contain nitrogen).
Health and Safety: What Are the Physiological Risks of Asphyxiation and Hyperbaric Narcosis?
Oxygen Displacement, Nitrogen Narcosis, and Decompression Sickness
Nitrogen is not toxic, but it can cause asphyxiation in confined spaces, where it displaces oxygen. The danger is greater because there is no warning odor, and a person may lose consciousness without noticing. At high partial pressures, as in deep diving with compressed air, nitrogen causes nitrogen narcosis, an alteration of mental state that increases with depth (noticeable at about 30 m, or a pressure of 4 bar), and a rapid decompression can cause decompression sickness (the bends), as dissolved nitrogen forms bubbles in tissues. Contact with liquid nitrogen causes cold burns and frostbite.
Frequently Asked Scientific Questions
Why is the N≡N triple bond (945 kJ/mol) nearly six times stronger than the N–N single bond (160 kJ/mol), whereas the C≡C triple bond (835 kJ/mol) is less than 2.5 times stronger than the C–C single bond (348 kJ/mol)?
In a single N–N bond (such as hydrazine, H2N–NH2), the small atomic radius of nitrogen forces the non-bonding lone electron pairs on adjacent nitrogen atoms into close proximity, generating severe inter-electronic lone-pair repulsion that weakens the σ bond to 160 kJ/mol. In N≡N, compact 2_p_ orbitals achieve exceptional sideways π overlap while the lone pairs point 180° away from the internuclear axis, yielding 945 kJ/mol; carbon lacks lone pairs in C–C bonds.
Why is Earth's atmosphere of 78% N2 and 21% O2 thermodynamically unstable with respect to the oceans?
As noted by Gilbert N. Lewis, the overall reaction of atmospheric dinitrogen, dioxygen, and liquid water to form aqueous nitric acid (N2(g) + 5/2 O2(g) + H2O(l) → 2 H+(aq) + 2 NO3−(aq)) is thermodynamically spontaneous at 298 K (Δ_G_° < 0 at neutral oceanic pH). Earth's atmosphere and oceans do not convert into dilute nitric acid solely because the endothermic initial step (N2 + O2 → 2 NO, Δ_H_° = +180.5 kJ/mol) imposes an enormous kinetic activation barrier.
Why does trophic enrichment increase δ15N by +3‰ to +4‰ per step in a food web?
During amino acid catabolism and transamination in consumer organisms, enzymes cleave C–14N bonds slightly faster than C–15N bonds due to the lower zero-point vibrational energy of the heavier 15N isotope (the kinetic isotope effect). Consequently, metabolic waste products excreted as urea, uric acid, or ammonia are isotopically depleted in 15N (enriched in 14N), leaving the retained body proteins of the consumer enriched by +3‰ to +4‰ relative to its diet.