Nitrogen

In 1772 Daniel Rutherford discovered Nitrogen. Nitrogen is the first member of group 15 of the periodic table. Nitrogen exists as a diatomic molecule (N₂) with a triple bond between the nitrogen atoms (NN). So nitrogen gas is also called dinitrogen. Its electronic configuration is 1s² 2s² 2p³.

Occurrence

Nitrogen occurs in free state in atmosphere, which contains 78% of nitrogen by volume. It is present in volcanic gases and gases evolved by burning coal. In combined state, nitrogen occurs in animals and plants as proteins, and as nitrates, such as NaNO₃ (chile salt peter), KNO₃ (nitre), nitrites and ammonium salts in the soil.

Preparation

Nitrogen can be easily prepared from its compounds.

By heating ammonium nitrite

Ammonium nitrite on heating decomposes to form nitrogen.

Since ammonium nitrite is very unstable, it cannot be stored and handled as such. Hence nitrogen is usually prepared by heating a mixture of ammonium chloride and sodium nitrite.

Laboratory preparation of nitrogen

In laboratory, dinitrogen is prepared by heating an aqueous solution containing equivalent amounts of ammonium chloride and sodium nitrite.

During the preparation, small amounts of NO, ammonia and HNO₃ are also formed. The gas is freed from these by passing through a solution of potassium dichromate acidified with sulphuric acid and collected over water.

Other Methods

By heating ammonium dichromate

Ammonium dichromate on heating decomposes to give nitrogen gas.

This reaction is very vigorous and inconvenient for the preparation of nitrogen.

By the oxidation of ammonia

When ammonia vapours are passed over heated cupric oxide (CuO), nitrogen is formed.

By the action of nitrous acid on urea

Urea when treated with nitrous acid gives nitrogen gas

Nitrous acid is unstable and is always prepared in the solution itself by the action of dilute HCl on sodium nitrite.

The gas is bubbled through NaOH solution to remove CO2 and then collected over water.

From air

On commercial scale, nitrogen is obtained by the fractional distillation of

liquefied air. The compressed carbon dioxide free air is liquefied by allowing it to expand in a region of low pressure. The liquefied air is then subjected to fractional distillation. Nitrogen having lower boiling point (78 K) distils over first leaving behind oxygen. Nitrogen obtained from air contains traces of oxygen. To remove these, the nitrogen so obtained is passed over heated copper turnings when copper reacts with oxygen to form CuO.

Physical properties of nitrogen

• Nitrogen gas (dinitrogen) is a colorless, odorless and tasteless gas.

• Nitrogen gas is slightly lighter than air.

• Nitrogen gas is very sparingly soluble (almost insoluble) in water: 23 mL per litre of water at 0°C and 1 atm pressure.

• Nitrogen gas liquefies at 77.2 K (-195.9°C) and solidifies at 63.2 K (209.9°C).

• It is not poisonous, but animals die in nitrogen atmosphere for want of oxygen.

  

• Nitrogen is neither combustible nor a supporter of combustion.

Chemical Properties of Nitrogen

Reactivity of nitrogen

Nitrogen gas (dinitrogen) is chemically non-reactive at ordinary temperature due to:

• Its small atomic size.

• Its high bond order. The N-N bond in the N₂ molecule has a formal triple bond character.

The NN bond distance is 109.8 pm and is very stable. Its bond dissociation energy is 946kJ mol^-1.

The dissociation constant is extremely small and even at 3000 K there is no appreciable dissociation. This indicates the strength of the N-N bond.

Reaction with highly electropositive metals

Nitrogen readily combines with highly electropositive metals (Group 1), even at room temperature, to give nitrides. For example,

These nitrides are ionic compounds in which nitrogen is present as a trinegative anion, N^3-. Thus, lithium nitride may be represented as (Li+)₃N^3-. Ionic nitrides are crystalline compounds and have high melting points.

Reaction with less electropositive metals

Nitrogen reacts with less electropositive such as magnesium, calcium etc., at higher temperatures, with boron and aluminium at red heat, and with silicon at white heat (very high temperature) to form the corresponding nitrides.

The nitrides of alkaline Earth metals (group 2) are ionic, while those of groups 13 and 14 are covalent in nature.

Reaction with transition metals

Transition metals such as iron, manganese form nitrides when heated with nitrogen at very high temperatures. The nitrides of the transition metals are interstitial compounds in which nitrogen atoms occupy the interstices (due to its small size). These nitrides are extremely hard and have high melting points.

Reactions with non-metals

• Nitrogen reacts with many non-metals under suitable conditions. For example with hydrogen:

• With oxygen, it forms nitric oxide, which being unstable reacts with more oxygen to form nitrogen dioxide.

This reaction between N₂ and O₂ also takes place during lightening.

Combination with compounds

Nitrogen combines with certain compounds on strong heating. For example,

With calcium carbide:

With alumina in presence of carbon:

Calcium cyanamide, aluminium nitride and magnesium nitride hydrolyze on boiling with water to give ammonia.

Hydrogen bonding

Nitrogen in its compounds with hydrogen shows hydrogen bonding though to a lesser extent than oxygen. For example, ammonia (NH₃) shows intermolecular hydrogen bonding.

Catenation

The tendency to establish self-linkage of the type M-M-M- leading to the formation of a chain or ring is called catenation. Due to p - p overlap nitrogen forms multiple bonds with itself as well as with carbon and oxygen e.g.,

Thus, nitrogen shows catenation tendency though weaker than that shown by carbon. The following compounds show the catenation property of nitrogen.

R - N = N - NR₂ R₂N - N = N - NR₂

RN = N - N(R) - N = NR

RN = N- N(R) - N = N - N(R) - N = NR

where R is some organic alkyl or aryl group. In some compounds R may also be H.

Uses

• Nitrogen gas is mainly used for the manufacture of compounds such as, ammonia, nitric acid, calcium cyanamide etc.

• Nitrogen gas is used for providing inert atmosphere in many metallurgical operations.

• Liquid nitrogen is used for producing low temperatures required in cryosurgery, preservation of biological materials etc.

Nitrogen Fixation

Nitrogen is present in the elemental state in the atmosphere to the extent of about 78% by volume. Nitrogen is important for life processes in all living organisms. But, free nitrogen cannot be directly assimilated by the living cells. Therefore, it is desirable to convert the free atmospheric nitrogen into useful nitrogen compounds. The process of conversion of free atmospheric nitrogen into useful nitrogen compounds is called fixation of nitrogen or nitrogen fixation.

Types of fixation

Natural fixation of nitrogen

Atmospheric nitrogen can be fixed into usable form by natural processes. Two ways of natural fixation of free nitrogen are:

Fixation of atmospheric nitrogen by electric discharge

During rains, when lightning strikes, nitrogen and oxygen of the atmosphere combine to form nitric oxide. This is further oxidized to nitrogen dioxide in the presence of excess of air. Nitrogen dioxide combines with rainwater in presence of oxygen to produce nitric acid. Nitric acid reacts with some alkalies in the soil, e.g., limestone (CaCO₃), to form nitrates. These nitrates serve as plant food and are responsible for their growth.

Fixation of atmospheric nitrogen by symbiotic bacteria

Certain plants such as peas, beans, etc., have some special kind of nitrogen fixing bacteria (symbiotic bacteria) in the nodules of their roots. These bacteria fix free nitrogen of the air as nitrogen compounds. These compounds spread into the soil after the decay of such plants.

Artificial fixation of nitrogen

Large amounts of free atmospheric nitrogen is however, fixed artificially by a number of methods. Three important methods are:

The cyanamide process

Nitrogen is also fixed as calcium cyanamide on heating it with calcium carbide at 1000°C in an electric furnace.

The mixture of calcium cyanamide and carbon (trade name nitrolim) is an important fertilizer. Calcium cyanamide is also be used as a source of ammonia. The ammonia so produced can be converted into useful fertilizers. Calcium cyanamide is decomposed by water to give ammonia.

Haber's process

When a mixture of nitrogen (obtained from the liquefied air) and hydrogen in the ratio 1:3, under a pressure of 200 - 800 atmospheres is passed over iron and molybdenum catalyst at about 450°C, ammonia is produced.

Ammonia is a basic raw material for preparing many salts and fertilizers.

Ostwald's process

A part of ammonia produced from the fixation of nitrogen is used for the manufacture of nitric acid by Ostwald's process.

Ammonia and nitric acid so produced are used for the production of many ammonium salts and nitrates, which act as fertilizers, e.g., calcium ammonium nitrate (CAN).

Birkland - Eyde's process

When nitrogen and oxygen are made to combine by passing air through an electric arc (~ 3000°C), nitric oxide is formed. It combines with more oxygen to form nitrogen dioxide which when absorbed in water, produces nitric acid.

Problems

7. How do you account for the inert character of dinitrogen?

Solution

The non-reactivity or inertness of nitrogen is due to:

Its small atomic size. Its high bond order. The N-N bond in the N₂ molecule has a formal triple bond character

The NN bond distance is 109.8 pm and is very stable. Its bond dissociation energy is 946kJ mol^-1.

The dissociation constant is extremely small and even at 3000 K there is no appreciable dissociation. This indicates the strength of the N-N bond and accounts for its inert character.

8. When a heap of ammonium dichromate is ignited, it erupts like a volcano, why?

Solution

When a heap of ammonium dichromate is ignited the chemical reaction that occurs is as follows:

Ammonium dichromate on heating decomposes to give nitrogen gas. But this reaction is so very vigorous that it is accompanied by the production of heat, light and lot of gases. As ammonium dichromate decomposes violently to form chromium oxide, this method is inconvenient for the preparation of nitrogen. The whole process looks like the eruption of a volcano and so this reaction is also referred to as chemical volcano.