Nitrogen

Nitrogen element

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.

It is placed along with p-block elements. The group members of nitrogen are phosphorus, arsenic, antimony, and bismuth.

Isotopes

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.

In 1919, a British physicist, Ernest Rutherford first introduced artificial facts of radioactive nuclear reaction by the bombardment of the alpha ray on ¹⁴N isotope to form ¹⁷O and protons particles.

Properties of nitrogen

Nitrogen
Symbol	N
Discovery	Daniel Rutherford in 1772
Name derived from	The Greek word nitron and genes mean nitre forming
Allotropes	Dinitrogen (N2)
Common isotope	7N14
Crystal structure	Hexagonal crystal lattice
Periodic properties
Atomic number	7
Electron per shell	2, 5
Atomic weight	14.007
Electronic configuration	[He] 2s2 2p3
Group	15
Period	2
Block	p-block
Physical properties
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
Chemical properties
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
	1402.33	2856.09	5300.47
Electron affinity	Unknown
Electronegativity	3.04 (Pauling scale)
Molar heat capacity	29.124 J mol−1 K−1
CAS number	7727-37-9

Chemical properties

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.

Oxidation number

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.

The formation of nitride ion (N⁻³) requires 2130 kJ mol⁻¹ energy. Therefore, it occurs only in nitride salts of strongly electropositive metals like lithium, beryllium, magnesium, calcium, etc.

Chemical reactivity

Dinitrogen is generally unreactive at room temperature due to the following reasons,

1. It has high N≡N bond energy.
2. The molecular orbital description of dinitrogen is 1σ_g¹ 2σ_u² 1π_u⁴ 3σ_g². A large difference between the HOMO (3σ_g) and LUMO (2π_u).
3. 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.

It can also be produced on a large scale by burning carbon or hydrocarbon in the air. The produced carbon dioxide and water are separated from the gas mixture.

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,

1. 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.
2. It is prepared by heating ammonium dichromate or sodium azide.
   (NH₄)₂Cr₂O₇ → N₂ + Cr₂O₃ + 4 H₂O
   2 NaN₃ → 2 Na + N₂
3. The reaction of ammonia with bromine water produced dinitrogen in the laboratory.
   8 NH₃ + 3 Br₂ → N₂ + 6 NH₄Br
4. 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 N₂ 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

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.