Nitrogen – facts and properties
- Symbol
- N
- Atomic number
- 7
- Atomic mass
- 14.007 u
- Period
- 2
- Group
- 15
- Category
- diatomic nonmetal
- Electron configuration
- [He] 2s² 2p³
- Electronegativity
- 3.04
- Atomic radius
- 56 pm
- Valence electrons
- 5
- State at standard conditions
- gas
- Melting point
- -210 °C
- Boiling point
- -195.79 °C
- Density
- 1.25 g/L
- Discovered
- 1772
Nitrogen is the gas that makes up most of the air we breathe, around 78 per cent by volume. It has atomic number 7 and is essential to every living thing, since proteins and DNA contain nitrogen. Yet the gas itself is so unreactive that plants and animals cannot use it directly. Learning how to turn nitrogen from the air into fertiliser was one of the most consequential achievements in the history of chemistry.
Position in the periodic table
Occurrence and the nitrogen cycle
The atmosphere is by far the largest reservoir of nitrogen, where it exists as the diatomic molecule N₂. The Earth's crust contains comparatively little. The main mineral deposits are nitrates, above all Chile saltpetre (sodium nitrate), which accumulated in the rainless Atacama Desert and was a vital raw material for fertiliser and explosives in the nineteenth century. In living matter nitrogen is bound in amino acids, proteins, nucleic acids and urea; in soil it occurs as ammonium and nitrate ions that plants absorb through their roots.
In the nitrogen cycle, certain bacteria convert atmospheric nitrogen into ammonium, a process called fixation. Some of these bacteria live in the root nodules of legumes such as peas, beans and clover. Other bacteria oxidise ammonium to nitrate, and yet others convert nitrate back into nitrogen gas, returning it to the air. Lightning adds a small share by forming nitrogen oxides that rain washes into the soil. Industrially, nitrogen is obtained by separating air. Air is liquefied and fractionally distilled: nitrogen, with the lower boiling point, evaporates first and leaves the oxygen behind. Smaller plants use membranes or molecular sieves instead. Nitrogen was identified in 1772 by Daniel Rutherford in Edinburgh, who isolated the part of air that supports neither combustion nor life. The name means nitre-former, after the mineral saltpetre.
Properties and reactions
Nitrogen has the electron configuration [He] 2s² 2p³ with five valence electrons and sits in period 2 and group 15. Two atoms join through a triple bond to form the N₂ molecule. This bond is one of the strongest known, and it explains why the gas is so inert: nitrogen does not burn, does not support combustion and reacts with almost nothing at room temperature. Breaking the triple bond requires high temperatures, high pressures or a catalyst.
The gas is colourless and odourless, and with a density of 1.25 g/L it is slightly lighter than air. It liquefies at -195.79 °C and solidifies at -210 °C; liquid nitrogen is one of the most widely used coolants. Nitrogen's electronegativity of 3.04 is high, exceeded only by oxygen, fluorine and chlorine. In compounds it shows a wide range of oxidation states, from minus three in ammonia to plus five in nitric acid. With hydrogen it forms ammonia, a base that reacts with acids to give ammonium salts. With oxygen it forms nitrogen oxides, which are produced in engines and power stations and contribute to smog and acid rain. Nitric acid is both a strong acid and a powerful oxidising agent. Many nitrogen compounds, such as nitroglycerine and TNT, are explosives because their decomposition releases extremely stable N₂ along with a great deal of energy. Heated with lithium or magnesium, nitrogen forms nitrides directly.
Uses in everyday life and technology
The most important use of nitrogen is the manufacture of ammonia by the Haber-Bosch process, in which nitrogen and hydrogen react over an iron catalyst at high pressure and temperature. Ammonia is converted into fertilisers such as ammonium nitrate and urea; without them the world could not feed its present population. Ammonia is also the starting point for nitric acid, for polymers such as nylon, and for dyes and explosives.
Nitrogen gas serves as an inert blanket wherever oxygen would cause trouble: in bags of crisps and packets of coffee, in wine tanks, in welding and in electronics manufacturing. Aircraft tyres and racing tyres are inflated with nitrogen because the dry gas contains no water vapour, which would amplify pressure changes, and does not oxidise the inside of the tyre. Liquid nitrogen flash-freezes food, preserves biological samples, sperm and egg cells for decades, and freezes off warts in the doctor's surgery. Liquid nitrogen also makes exceptionally smooth ice cream. Nitrous oxide, laughing gas, is an anaesthetic and the propellant in whipped-cream dispensers. In daily life nitrogen appears mainly as fertiliser in gardens and fields. Too much of it leaches into groundwater as nitrate and escapes into the air as ammonia and nitrous oxide, which is why the use of nitrogen fertiliser is now strictly regulated in many countries.
Molecules containing this element
- Nitrogen N2
- Ammonia NH3
- Ammonium bromide NH4Br
Quick question
Which pair of elements has the same number of valence electrons?
FAQ
Is nitrogen toxic?
Nitrogen gas is not toxic; we inhale it with every breath. The hazard is suffocation when it displaces oxygen, for example in tanks or in poorly ventilated rooms where liquid nitrogen has evaporated. Because the body senses rising carbon dioxide rather than falling oxygen, a person in pure nitrogen can lose consciousness without warning. Liquid nitrogen also causes severe cold burns.
Why can't plants use nitrogen from the air?
The triple bond in the nitrogen molecule is too strong for plants to break. Only certain bacteria possess the enzyme nitrogenase, which converts atmospheric nitrogen into ammonium. Plants therefore take up nitrogen from the soil as nitrate or ammonium, supplied by fertiliser, by decaying organic matter or by nitrogen-fixing bacteria living in their roots.
What is liquid nitrogen used for?
Liquid nitrogen is cheap, non-toxic and extremely cold. It cools superconducting materials, stores cells and tissue samples for decades, freezes food rapidly and is used in medicine to freeze skin lesions. In laboratories it fills cold traps that condense vapours before they reach vacuum pumps.