Nitrogen is a nonmetallic chemical element of group 15 of the periodic table with symbol N and chemical formula N2. It is a colourless, odorless, tasteless gas, and an essential element of the living matter. Dinitrogen is the most plentiful substance in the earth’s atmosphere which is present in air 78% by volume and 75% by weight.
The abundance of nitrogen in rocks and soil of the earth’s environment is remarkably low (about 19 ppm) comparable to gallium, niobium, and lithium. Only nitrogenous minerals like potassium nitrate (saltpeter) and sodium nitrate (Chile saltpeter) are isolated via nitric acid by the action of nitrifying or nitrogen fixation bacteria.
Who discovered nitrogen?
Carl Wilhelm Scheele and Henry Cavendish had independently discovered and isolated nitrogen but the credit for discovery was given to Scottish physician Daniel Rutherford in 1772. Rutherford has given this credit because his work was published first in a science journal.
The name nitrogen was given by French chemist Jean-Antoine-Claude Chaptal in 1790. He may derive the name from the Greek word ‘nitron’ and ‘genes’ mean nitre forming.
Where is nitrogen found?
Major deposits of nitrogen compounds like nitrate are found in Bolivia, Italy, Spain, Russia, and some regions of India. The major human interference occurs through the artificial fixation of N2 by fertilizers to meet the increasing demand for foods in our growing human family.
Nitrogen in the periodic table
The electronic configuration of nitrogen is 1s2 2s2 2p3. Therefore, it is placed in group 15 and period 2 of the periodic table.
In learning chemistry, naturally occurring nitrogen exists mainly in two stable isotopes, 14N and 15N. The relative abundance of 14N:15N ≈ 272:1. The radioactive isotope, 7N15 may be prepared or separated by an exchange reaction or by thermal diffusion.
Properties of nitrogen
|Discovery||Daniel Rutherford in 1772|
|Name derived from||The Greek word nitron and genes mean nitre forming|
|Crystal structure||Hexagonal crystal lattice|
|Electron per shell||2, 5|
|Electronic configuration||[He] 2s2 2p3|
|State at 20 °C||Gas|
|Melting point||−210 °C, −346 °F, 63.2 K|
|Boiling Point||−195.79 °C, −320.43 °F, 77.35 K|
|Critical temperature||126.21 K|
|Density||0.001145 g cm−3|
|Atomic radius (non-bonded)||1.55 Å|
|Covalent radius||0.71 Å|
|Oxidation number or states||5, 4, 3, 2, -3|
|Ionization energy (kJ mol−1)||1st||2nd||3rd|
|Electronegativity||3.04 (Pauling scale)|
|Molar heat capacity
||29.124 J mol−1 K−1|
The outer quantum shell electronic configuration of nitrogen is s2p3 with two paired s-electrons and three unpaired p-electrons. The electronic configuration suggests that the element is closer to the next noble gas than the preceding noble gas.
Assuming the +5 cationic configurations by losing all the five outer electrons is just impossible. Therefore, most of the chemical compounds of nitrogen are covalent compounds with an oxidation state of +3 or +5. The oxidation number of the element covers a wide range from −3 (in NH3) to +5 (HNO3). Due to the non-availability of energetically vacant d-orbitals, it has no scope for valency expansion.
The chemistry of nitrogen describes that the atom may attain the next noble gas electronic configuration by gaining three electrons to form an N−3 ion (oxidation number = −3). However, such electron attachment involves a high positive enthalpy change.
Dinitrogen is generally unreactive at room temperature due to the following reasons,
- It has high N≡N bond energy.
- The molecular orbital description of dinitrogen is 1σg1 2σu2 1πu4 3σg2. A large difference between the HOMO (3σg) and LUMO (2πu).
- Symmetrical electron distribution. It makes nitrogen molecules nonpolar.
Production of nitrogen
Production from air
Industrially or commercially large amounts of dinitrogen are obtained during the isolation of oxygen by fractional distillation of liquid air. The N2 molecule boils off before oxygen because its boiling point is lower (N2: −195.8 °C; O2: −183.1 °C). The product usually contains a fraction of argon.
Laboratory preparation of nitrogen
Nitrogen may be obtained in the laboratory by a variety of chemical reactions. Some of these production processes are given below,
- It is prepared in the laboratory by heating ammonium nitrate.
NH4NO2 → N2 + 2 H2O
Some NO and HNO3 are formed by this process. These may be removed by absorbents like potassium dichromate in presence of sulfuric acid.
- It is prepared by heating ammonium dichromate or sodium azide.
(NH4)2Cr2O7 → N2 + Cr2O3 + 4 H2O
2 NaN3 → 2 Na + N2
- The reaction of ammonia with bromine water produced dinitrogen in the laboratory.
8 NH3 + 3 Br2 → N2 + 6 NH4Br
- The high-temperature reaction of ammonia with cupric oxide.
3 CuO + 2 NH3 → 3 Cu + N2 + 3 H2O
What is nitrogen used for?
- Elemental nitrogen is largely used to provide an inert atmosphere in metallurgy.
- It is also used in various chemical industries such as the iron and steel industry for welding, soldering, brazing, and petrochemical industry.
- A large scale of nitrogen is used in the chemical industry for the manufacture of ammonia and calcium cyanamide
- It is a useful chemical for cryogenic research.
- Dinitrogen is used as an inhibitor for fire or explosions.
Uses of liquid nitrogen
- The liquid nitrogen is used in low-temperature machining.
- In grinding rubbers and rubber-like substances, we also used dinitrogen.
- Liquid nitrogen is also used for the preservation of biological specimens.
- In medicine, liquid nitrogen is used as a refrigerant for rapid freezing to preserve blood, bone marrow, tissue, bacteria, and semen.
What is the nitrogen cycle?
Nitrogen is an essential element for the growth of plant and animal life in nature. It continuously interchanges between the atmosphere and biosphere. To balance the nitrogen cycle in nature, it returns to the atmosphere by the following major steps,
- The action of nitrogen fixation bacteria
- Artificial fixation of dinitrogen
- Lighting in the upper atmosphere
- Death and decay of plants and animals
- Burning of wood, coal, and petroleum
- Drainage of surface water
Nitrogen fixation by bacteria
About 60 percent of N2 is input into the soil by the action of nitrifying bacteria (Rhizobium, Anabaena, Nostoc, Azotobactor, and Clostridium pastorium). These nitrifying bacteria directly convert dinitrogen into nitrates or ammonium salts.
Rhizobium is the most important bacteria in this category. It lives symbiotically in the nodules or roots of certain plants like Leguminosae, pea, beans, etc.
The nitrogen-fixing enzyme nitrogenase in fixed nitrogen into the soil. The fixation is subsequently greater than the artificial fixing through the Haber process (150 million tonnes vs 120 million tonnes).
Industrial fixation is used to grow nitrogen into the soil for the production of large amounts of food for the ever-increasing population of the world. The main procedure involves the Haber process for the manufacture of ammonia which converted nitric acid and other fertilizers.
Other processes to fix nitrogen
- Lighting in the upper atmosphere leads to the formation of NO and NO2, which are carried by nitric acid rain to the soil.
- The death or decay of plants and animals returns most of the nitrogen to the soil.
- The combustion of wood, coal, and petroleum releases a small amount of NO2 into the atmosphere.
- The drainage of surface water carries some nitrogen gas into the sea which uses to support the common marine life.