Compounds of Nitrogen: We know that nitrogen is the most abundant gas in the earth’s atmosphere. The element nitrogen is a group \(15\) elements. It is non-metal and exists in the gaseous state. The gaseous nitrogen \(\left( {{{\rm{N}}_2}} \right)\) comprises about \(78\% \) by the volume of our atmosphere. Nitrogen combines with other elements to form several compounds. In this article, we are going to discuss the different compounds of nitrogen in detail.

What are the Compounds of Nitrogen?

The common compounds of nitrogen include ammonia, nitric acid and different oxides of nitrogen. We will discuss it one by one in detail.

Ammonia \(\left( {{\rm{N}}{{\rm{H}}_3}} \right)\)

Ammonia is a compound containing nitrogen and hydrogen atoms. Its chemical formula is \({\rm{N}}{{\rm{H}}_3}.\) Ammonia is a covalent compound that is formed by the sharing of electrons between nitrogen and hydrogen atoms. The structure of ammonia can be given as,

Ammonia has a trigonal planar structure with the three hydrogen atoms and a lone pair of electrons attached to the central nitrogen atom.

Preparation of Ammonia

Ammonia can be prepared in the laboratory using the following processes:

Preparation of Ammonia through Haber’s Process

Haber’s process is the cheapest and the common method used for the preparation of ammonia. In this method, nitrogen combines with hydrogen to form ammonia. That is,

\({{\rm{N}}_2} + 3{{\rm{H}}_2} \leftrightarrow 2{\mkern 1mu} {\rm{N}}{{\rm{H}}_3};{\mkern 1mu} \,{\Delta _{\rm{f}}}{{\rm{H}}^0} = \, – 46.1{\mkern 1mu} {\rm{kJ}}{\mkern 1mu} {\rm{mo}}{{\rm{l}}^{ – 1}}\)

The above reaction is highly exothermic.

During the process, the favourable conditions for the maximum ammonia yield are high concentration of the reactants, high pressure and low temperature, according to Le- Chatelier’s principle. Hence, an optimum temperature of \( \sim 700\,{\rm{K}}\) and a pressure of \( \sim 200\) atm is used for the maximum yield. The use of a catalyst such as iron oxide with small amounts of \({{\rm{K}}_2}{\rm{O}}\) and \({\rm{A}}{{\rm{l}}_2}{{\rm{O}}_3}\) increases the rate of attainment of equilibrium.

Preparation of Ammonia through Laboratory Process

Ammonia can be prepared by heating ammonium chloride with slaked lime or quicklime. That is,

\(2\,{\rm{N}}{{\rm{H}}_4}{\rm{Cl}} + {\rm{Ca}}{\left( {{\rm{OH}}} \right)_2} \to {\rm{CaC}}{{\rm{l}}_2} + 2\,{\rm{N}}{{\rm{H}}_3} + 2\,{{\rm{H}}_2}{\rm{O}}\)

\(2\,{\rm{N}}{{\rm{H}}_4}{\rm{Cl}} + {\rm{CaO}} \to {\rm{CaC}}{{\rm{l}}_2} + 2\,{\rm{N}}{{\rm{H}}_3} + {{\rm{H}}_2}{\rm{O}}\)

Ammonia is present in air and soil in very small quantities, which is formed by the decay of nitrogenous organic matter. For example, ammonia can be obtained from urea as,

\({\rm{N}}{{\rm{H}}_2}{\rm{CON}}{{\rm{H}}_2} + 2\,{{\rm{H}}_2}{\rm{O}} \to {\left( {{\rm{N}}{{\rm{H}}_4}} \right)_2}{\rm{C}}{{\rm{O}}_3} \leftrightarrow 2\,{\rm{N}}{{\rm{H}}_3} + {{\rm{H}}_2}{\rm{O}} + {\rm{C}}{{\rm{O}}_2}\)

Properties of Ammonia

Let us discuss the physical and chemical properties of ammonia in detail.

Physical Properties of Ammonia

Ammonia is a colourless gas with a characteristic pungent smell. It is lighter than air and is highly soluble in water. Ammonia has high melting and boiling point than expected based on its molecular mass. This is because the molecules are associated with each other by intermolecular hydrogen bonding.

Chemical Properties of Ammonia

The chemical properties of ammonia are discussed below:

1. Basic Properties of Ammonia

The aqueous solution of ammonia is basic due to the formation of \({\rm{O}}{{\rm{H}}^ – }\) ions.

\({\rm{N}}{{\rm{H}}_3}\left( {\rm{g}} \right) + {{\rm{H}}_2}{\rm{O}}\left( {\rm{I}} \right) \leftrightarrow {\rm{NH}}_4^ + + {\rm{O}}{{\rm{H}}^ – }\left( {{\rm{aq}}} \right)\)

It turns red litmus blue, and it forms ammonium salts with acids. For example,

\({\rm{N}}{{\rm{H}}_3} + {\rm{HCl}} \to {\rm{N}}{{\rm{H}}_4}{\rm{Cl}}\)

2. Reducing Properties of Ammonia

Ammonia is a good reducing agent. That is, the nitrogen atom in ammonia has lone pair of electrons; hence it can donate electrons and can get oxidized itself. For example,

When ammonia is passed over heated copper oxide, metallic copper, nitrogen, and water vapour are formed. That is,

\(3\,{\rm{CuO}} + 2\,{\rm{N}}{{\rm{H}}_3} \to 3\,{\rm{Cu}} + {{\rm{N}}_2} + 3\,{{\rm{H}}_2}{\rm{O}}\)

Here, ammonia reduces cupric oxide to copper metal.

3. Reaction of Ammonia with Air or Oxygen (Oxidation)

Ammonia burns in atmospheric air or oxygen with a blue flame. That is,

\(4\,{\rm{N}}{{\rm{H}}_3} + 3\,{{\rm{O}}_2} \to 2\,{{\rm{N}}_2} + 6\,{{\rm{H}}_2}{\rm{O}}\)

If the same reaction is done using platinum gauze, at \(500{\,^{\rm{o}}}{\rm{C}},\) nitric oxide is formed. That is,

\(4{\mkern 1mu} {\rm{N}}{{\rm{H}}_3} + 5{\mkern 1mu} {{\rm{O}}_{2\,{\rm{Pt}}}} \to 4{\mkern 1mu} {\rm{NO}} + 6{\mkern 1mu} {{\rm{H}}_2}{\rm{O}}\)

4. Reaction of Ammonia with Halogens

Ammonia, with an insufficient amount of chlorine, produces nitrogen and hydrochloric acid. The hydrochloric acid formed reacts with ammonia producing ammonium chloride.

\(\begin{array}{l} 2\,{\rm{N}}{{\rm{H}}_3} + 3\,{\rm{C}}{{\rm{l}}_2} \to {{\rm{N}}_2} + 6\,{\rm{HCl}}\\ 6\,{\rm{N}}{{\rm{H}}_3} + 6\,{\rm{HCl}} \to 6\,{\rm{N}}{{\rm{H}}_3}{\rm{Cl}}\\ \_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\\ 8\,{\rm{N}}{{\rm{H}}_3} + 3\,{\rm{C}}{{\rm{l}}_2} \to 6\,{\rm{N}}{{\rm{H}}_4}{\rm{Cl}} + {{\rm{N}}_2} \end{array}\)

With an excess of chlorine, \({\rm{NC}}{{\rm{l}}_3}\) is produced. That is,

\({\rm{N}}{{\rm{H}}_3} + 3\,{\rm{C}}{{\rm{l}}_2} \to {\rm{NC}}{{\rm{l}}_3} + 3\,{\rm{HCl}}\)

An explosive compound nitrogen tri-iodide can be prepared by reacting \({\rm{N}}{{\rm{H}}_3}\) into excess of iodine.

5. Reaction of Ammonia with Alkali Metals

When ammonia reacts with alkali metals, metal amide and hydrogen are formed. That is,

\(2\,{\rm{Na}} + 2\,{\rm{N}}{{\rm{H}}_3} \to 2\,{\rm{NaN}}{{\rm{H}}_2} + {{\rm{H}}_2}\)

\(2\,{\rm{K}} + 2\,{\rm{N}}{{\rm{H}}_3} \to 2\,{\rm{KN}}{{\rm{H}}_2} + {{\rm{H}}_2}\)

6. Ammonia as a Lewis base

In ammonia, the presence of a lone pair of electrons on the central nitrogen atom makes it behave as a Lewis base. Hence, it forms a dative bond with electron-deficient molecules and transition metal cations having vacant d-orbitals, and as a result, it forms complex compounds. Many such reactions are used for the detection of metal ions such as \({\rm{C}}{{\rm{u}}^{2 + }},\,{\rm{A}}{{\rm{g}}^ + },\)etc.

7. Reaction of Ammonia with Nessler’s reagent

Ammonia reacts with Nessler’s reagent (alkaline solution of \({{\rm{K}}_2}{\rm{Hg}}{{\rm{l}}_4}\)) to give a brown precipitate due to the formation of iodide of Million’s base. This test can be used for the detection of ammonium ions in qualitative analysis.

Uses of Ammonia

Ammonia has got many uses. They are:

1. Ammonia is used for the preparation of various compounds such as nitric acid and sodium carbonate.
2. Ammonia is mainly used in the production of fertilizers like ammonium sulphate, urea, ammonium phosphate and ammonium nitrate.
3. Liquid ammonia is used as a refrigerant.
4. Ammonia is used as a laboratory reagent.
5. Ammonia is used for the purification of water supplies.

Nitric acid \(\left( {{\rm{HN}}{{\rm{O}}_3}} \right)\)

Nitric acid is an inorganic acid. It is one of the most important acids formed from nitrogen. It is also known as aqua fortis or the spirit of nitre. The chemical formula of nitric acid is \({\rm{HN}}{{\rm{O}}_3}.\) In nitric acid, nitrogen is the central atom. In \({\rm{HN}}{{\rm{O}}_3},\) one oxygen atom is doubly bonded to the nitrogen atom, another oxygen atom is singly bonded to the nitrogen atom, and one oxygen atom is bonded to the nitrogen atom as well as with one hydrogen atom. \({\rm{HN}}{{\rm{O}}_3}\) is a covalent compound. The structure of nitric acid can be given as,

Let us discuss more about nitric acid.

Preparation of Nitric acid \(\left( {{\rm{HN}}{{\rm{O}}_3}} \right)\)

1. Nitric acid is generally prepared on a large scale in the laboratory by Ostwald’s process. In this process, a mixture of ammonia and is heated in the air in the presence of a platinum catalyst to about \(500\,{\rm{K}}{\rm{.}}\)

\({\text{2NO}}\left( {\text{g}} \right){\text{ + }}{{\text{O}}_{\text{2}}}\left( {\text{g}} \right) \rightleftharpoons {\text{2N}}{{\text{O}}_{\text{2}}}\left( {\text{g}} \right)\)

\(3{\mkern 1mu} {\rm{N}}{{\rm{O}}_2}\left( {\rm{g}} \right) + {{\rm{H}}_2}{\rm{O}}\left( {\rm{l}} \right) \to 2{\mkern 1mu} {\rm{HN}}{{\rm{O}}_3}\left( {{\rm{aq}}} \right) + {\rm{NO}}\left( {\rm{g}} \right)\)

The \({\rm{NO}}\) thus formed is recycled. The aqueous \({\rm{HN}}{{\rm{O}}_3}\) can be concentrated by distillation up to \( \sim 68\,\% \) by mass. The further concentration to \(98\,\% \) can be achieved by dehydration with concentrated sulphuric acid.

2. Nitric acid can also be prepared by heating \({\rm{KN}}{{\rm{O}}_3}\) or \({\rm{NaN}}{{\rm{O}}_3}\) and concentrated \({{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4}\) in a glass retort.
\({\rm{NaN}}{{\rm{O}}_3}{\rm{ + }}{{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4} \to {\rm{NaHS}}{{\rm{O}}_4} + {\rm{HN}}{{\rm{O}}_3}\)

Properties of Nitric acid \(\left( {{\rm{HN}}{{\rm{O}}_3}} \right)\)

Let us discuss the physical and chemical properties of nitric acid.

1. Physical Properties of Nitric acid \(\left( {{\rm{HN}}{{\rm{O}}_3}} \right)\)

Nitric acid is a colourless liquid. We find it as yellow in bottles because it is impure due to the presence of dissolved \({\rm{N}}{{\rm{O}}_2}\) gas in it. The specific gravity of nitric acid is \(1.504.\) It is highly corrosive in nature; hence it produces yellow blisters on skin. It is hygroscopic and fumes in the air. It is soluble in water too.

2. Chemical Properties of Nitric acid \(\left( {{\rm{HN}}{{\rm{O}}_3}} \right)\)

The chemical properties of nitric acid are disscused below

(i) Acidic Properties of Nitric acid \(\left( {{\rm{HN}}{{\rm{O}}_3}} \right)\)

Nitric acid is a strong monobasic acid. It reacts with oxides, hydroxides and carbonates and results in the formation of nitrates.

\({\rm{CuO}} + 2\,{\rm{HN}}{{\rm{O}}_3} \to {\rm{Cu}}{\left( {{\rm{N}}{{\rm{O}}_3}} \right)_2} + {{\rm{H}}_2}{\rm{O}}\)

\({\rm{NaOH}} + {\rm{HN}}{{\rm{O}}_3} \to {\rm{NaN}}{{\rm{O}}_3} + {{\rm{H}}_2}{\rm{O}}\)

\({\rm{N}}{{\rm{a}}_2}{\rm{C}}{{\rm{O}}_3} + 2\,{\rm{HN}}{{\rm{O}}_3} \to 2\,{\rm{NaN}}{{\rm{O}}_3} + {{\rm{H}}_2}{\rm{O}} + {\rm{C}}{{\rm{O}}_2}\)

(ii) Oxidising Properties of Nitric acid \(\left( {{\rm{HN}}{{\rm{O}}_3}} \right)\)

Nitric acid is a strong oxidizing agent as it can release nascent (atomic) oxygen in the following manner,

\(2\,{\rm{HN}}{{\rm{O}}_3} \to {{\rm{H}}_2}{\rm{O}} + 2\,{\rm{NO}} + 3\,{\rm{O}}\)

Examples are:

(a) Nitric acid oxidizes hydrogen sulphide to sulphur. That is,

\(\begin{array}{l} 2\,{\rm{HN}}{{\rm{O}}_3} \to {{\rm{H}}_2}{\rm{O}} + 2\,{\rm{NO}} + 3\,{\rm{O}}\\ \left( {{{\rm{H}}_2}{\rm{S}} + {\rm{O}} \to {{\rm{H}}_2}{\rm{O}} + {\rm{S}}} \right) \times 3\\ \_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\\ 2\,{\rm{HN}}{{\rm{O}}_3} + 3\,{{\rm{H}}_2}{\rm{S}} \to 2\,{\rm{NO}} + 4\,{{\rm{H}}_2}{\rm{O}} + 3\,{\rm{S}} \end{array}\)

(b) Nitric acid reacts with potassium iodide to liberate iodine. That is,

\(\begin{array}{l} 2\,{\rm{HN}}{{\rm{O}}_3} \to {{\rm{H}}_2}{\rm{O}} + 2\,{\rm{NO}} + 3\,{\rm{O}}\\ \left( {2\,{\rm{Kl}} + {\rm{O}} \to {{\rm{K}}_2}{\rm{O}} + {{\rm{I}}_2}} \right) \times 3\\ \left( {{{\rm{K}}_2}{\rm{O}} + 2\,{\rm{HN}}{{\rm{O}}_3} \to 2\,{\rm{KN}}{{\rm{O}}_3} + {{\rm{H}}_2}{\rm{O}}} \right) \times 3\\ \_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\\ 8\,{\rm{HN}}{{\rm{O}}_3} + 6\,{\rm{Kl}} \to 6\,{\rm{KN}}{{\rm{O}}_3} + 2\,{\rm{NO}} + 3\,{{\rm{I}}_2} + 4\,{{\rm{H}}_2}{\rm{O}} \end{array}\)

(iii) Reaction of Nitric acid with Non-metals of Nitric acid \(\left( {{\rm{HN}}{{\rm{O}}_3}} \right)\)

In the presence of non-metals, nitric acid dissociates to give \({\rm{N}}{{\rm{O}}_2}.\) That is,

\(2\,{\rm{HN}}{{\rm{O}}_3} \to {{\rm{H}}_2}{\rm{O}} + 2\,{\rm{N}}{{\rm{O}}_2} + {\rm{O}}\)
For example, with carbon, it reacts as,

\(\begin{array}{l} \left( {2\,{\rm{HN}}{{\rm{O}}_3} \to {{\rm{H}}_2}{\rm{O}} + 2\,{\rm{N}}{{\rm{O}}_2} + {\rm{O}}} \right) \times 2\\ {\rm{C}} + 2\,{\rm{O}} \to {\rm{C}}{{\rm{O}}_2}\\ \_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\_\\ 4\,{\rm{HN}}{{\rm{O}}_3} + {\rm{C}} \to {\rm{C}}{{\rm{O}}_2} + 4\,{\rm{N}}{{\rm{O}}_2} + 2\,{{\rm{H}}_2}{\rm{O}} \end{array}\)

(iv) Reaction of Nitric acid with Metals of Nitric acid \(\left( {{\rm{HN}}{{\rm{O}}_3}} \right)\)

All metals react with nitric acid except gold and platinum. For example, zinc reacts with dilute nitric acid to give \({\rm{N}}{{\rm{O}}_2}\) and it reacts with concentrated acid to give \({\rm{N}}{{\rm{O}}_2}.\)That is,

\(4\,{\rm{Zn}} + 10\,{\rm{HN}}{{\rm{O}}_3}\left( {{\rm{dilute}}} \right) \to 4\,{\rm{Zn}}{\left( {{\rm{N}}{{\rm{O}}_3}} \right)_2} + 5\,{{\rm{H}}_2}{\rm{O}} + {{\rm{N}}_2}{\rm{O}}\)

\({\rm{Zn}} + 4\,{\rm{HN}}{{\rm{O}}_3}\left( {{\rm{conc}}.} \right) \to {\rm{Zn}}{\left( {{\rm{N}}{{\rm{O}}_3}} \right)_2} + 2\,{{\rm{H}}_2}{\rm{O}} + 2\,{\rm{N}}{{\rm{O}}_2}\)

Some metals will become passive when it is treated with concentrated nitric acid because of the formation of a passive film of oxide on their surface.

3. Brown ring test for Nitrates \(\left( {{\rm{NO}}_3^ – } \right)\)

Brown ring test is used as a confirmatory test for nitrates. In this method, about \(1\,{\rm{mL}}\) of the salt solution is treated with an equal volume of concentrated sulphuric acid. The resulting solution is cooled under the tap, and then a freshly prepared solution of \({\rm{FeS}}{{\rm{O}}_4}\) is added to it without disturbing the solution.

As a result, we can see the formation of a brown ring at the junction of two liquids that can disappear and reappear on further addition of \({\rm{FeS}}{{\rm{O}}_4}.\)This confirms the presence of \({\rm{NO}}_3^ – \) ions. The brown ring is formed due to the formation of pentaaquanitroso iron (II) sulphate.

Uses of Nitric acid

Nitric acid has got so many uses. They are:

1. Nitric acid is used in the manufacture of artificial silk, fertilizers and the explosives like nitroglycerine, picric acid, etc.
2. It is used for the preparation of aqua regia (A mixture of concentrated \({\rm{HCl}}\) and concentrated \({\rm{HN}}{{\rm{O}}_3}\) in ratio \(3:1\)), which is used for dissolving noble metals.
3. It is used for the purification of silver and gold.
4. It is used as a laboratory reagent.
5. It is used as an oxidizer in rocket fuel.

Oxides of Nitrogen

Nitrogen forms various oxides in different oxidation states. Let us discuss them.

(i) Dinitrogen Oxide \(\left( {{{\rm{N}}_2}{\rm{O}}} \right)\)

Dinitrogen oxide has a chemical formula \({{\rm{N}}_2}{\rm{O}}.\) In this compound, the oxidation state of nitrogen is \(+1.\)

Dinitrogen oxide has a linear shape. It is a covalent compound.

Dinitrogen oxide is a colourless gas, and is neutral in nature. It can be prepared by the action of heat on ammonium nitrate. That is,

\({\text{N}}{{\text{H}}_{\text{4}}}{\text{N}}{{\text{O}}_{\text{3}}}\xrightarrow{{{\text{Heat}}}}{{\text{N}}_{\text{2}}}{\text{O + 2}}{{\text{H}}_{\text{2}}}{\text{O}}\)

(ii) Nitrogen Monoxide \(\left( {{\rm{NO}}} \right):\)

The chemical formula of nitrogen monoxide is \({\rm{NO}}.\) The oxidation state of nitrogen in this compound is \( + 2.\)

Nitrogen monoxide is a colourless gas. It is neutral in nature. It can be prepared by the reaction of sodium nitrite, ferrous sulphate and sulphuric acid. That is,

\(2\,{\rm{NaN}}{{\rm{O}}_2} + 2\,{\rm{FeS}}{{\rm{O}}_4} + 3\,{{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4} \to {\rm{F}}{{\rm{e}}_2}{\left( {{\rm{S}}{{\rm{O}}_4}} \right)_3} + 2\,{\rm{NaHS}}{{\rm{O}}_4} + 2\,{{\rm{H}}_2}{\rm{O}} + 2\,{\rm{NO}}\)

(iii) Dinitrogen Trioxide \(\left( {{{\rm{N}}_2}{{\rm{O}}_3}} \right)\)

The chemical formula of dinitrogen trioxide is \({{\rm{N}}_2}{{\rm{O}}_3}.\) The oxidation state of nitrogen in \({{\rm{N}}_2}{{\rm{O}}_3}\) is \( + 3.\)

Dinitrogen trioxide is a blue solid. It is acidic in nature. It can be prepared by heating nitrogen monoxide with dinitrogen tetroxide.

\({\text{2NO + }}{{\text{N}}_{\text{2}}}{{\text{O}}_{\text{4}}}\xrightarrow{{250\,{\text{K}}}}{\text{2}}{{\text{N}}_{\text{2}}}{{\text{O}}_{\text{3}}}\)

(iv) Nitrogen Dioxide \(\left( {{\rm{N}}{{\rm{O}}_2}} \right)\)

The chemical formula of nitrogen dioxide is \({\rm{N}}{{\rm{O}}_2}.\) The oxidation state of nitrogen in \({\rm{N}}{{\rm{O}}_2}\) is \( + 4.\)

Nitrogen dioxide is a brown gas. It is acidic in nature. It can be prepared by the thermal decomposition of lead nitrate.
\({\text{2Pb}}{\left( {{\text{N}}{{\text{O}}_{\text{3}}}} \right)_{\text{2}}}\xrightarrow{{{\text{673}}\,{\text{K}}}}{\text{4N}}{{\text{O}}_{\text{2}}}{\text{ + 2PbO + }}{{\text{O}}_{\text{2}}}\)

(v) Dinitrogen Tetroxide \(\left( {{{\rm{N}}_2}{{\rm{O}}_4}} \right)\)

The chemical formula of dinitrogen tetroxide is \({{\rm{N}}_2}{{\rm{O}}_4}.\) In this compound, the oxidation state of nitrogen is \( + 4.\)

It can exist as a colourless solid or liquid. It is acidic in nature. It can be prepared by heating nitrogen dioxide. That is,
\({\text{2N}}{{\text{O}}_{\text{2}}}\mathop \rightleftharpoons \limits_{{\text{Heat}}}^{{\text{Cool}}} {{\text{N}}_{\text{2}}}{{\text{O}}_{\text{4}}}\)

(vi) Dinitrogen Pentoxide \(\left( {{{\rm{N}}_2}{{\rm{O}}_5}} \right)\)

The chemical formula of dinitrogen pentoxide is \({{\rm{N}}_2}{{\rm{O}}_5}.\) In this compound, the oxidation state of nitrogen is \( + 5.\)

Dinitrogen peroxide appears as a colourless solid. It is acidic in nature. It can be prepared by the reaction of nitric acid with phosphorous pentoxide.
\(4\,{\rm{HN}}{{\rm{O}}_3} + {{\rm{P}}_4}{{\rm{O}}_{10}} \to 4\,{\rm{HP}}{{\rm{O}}_3} + 2\,{{\rm{N}}_2}{{\rm{O}}_5}\)

Compounds of Nitrogen and Sulphur

Sulphur and nitrogen are non-metals. They combine together to form sulphur nitrides. They include sulphur monoxide \(\left( {{\rm{SN}}} \right),\) disulphur monoxide \(\left( {{{\rm{N}}_2}{\rm{S}}} \right),\) monosulphur dinitride \(\left( {{\rm{S}}{{\rm{N}}_2}} \right).\) These compounds are generally unstable in nature. The other examples of nitrides include pentasulphur hexanitride \(\left( {{{\rm{S}}_5}{{\rm{N}}_6}} \right),\) tetrasulphur tetranitride \(\left( {{{\rm{S}}_4}{{\rm{N}}_4}} \right),\) tetrasulphur dinitride \(\left( {{{\rm{S}}_4}{{\rm{N}}_2}} \right),\) disulphur dinitride, \(\left( {{{\rm{S}}_2}{{\rm{N}}_2}} \right),\) polythiazyl \(\left[ {{{\left( {{\rm{SN}}} \right)}_{\rm{x}}}} \right]\) and thiatetrazole \(\left( {{\rm{S}}{{\rm{N}}_4}} \right).\)

Compounds of Nitrogen and Hydrogen

One of the common compounds formed between hydrogen and nitrogen is ammonia. In this article, we have discussed ammonia in detail. The other compounds include nitrogen hydrides. They are hydroxylamine \(\left( {{\rm{N}}{{\rm{H}}_2}{\rm{OH}}} \right),\) hydrazine \(\left( {{{\rm{N}}_2}{{\rm{H}}_4}} \right),\) hydrogen azide \(\left( {{\rm{H}}{{\rm{N}}_3}} \right)\) and diimide \(\left( {{{\rm{H}}_2}{{\rm{N}}_2}} \right).\)

Compounds of Nitrogen and Phosphorus

The chemical compounds formed between phosphorous and nitrogen include phosphorous nitrides. They include phosphorous mononitride \(\left( {{\rm{PN}}} \right),\) Tetraphosphorous hexanitride \(\left( {{{\rm{P}}_4}{{\rm{N}}_6}} \right)\) and Triphosphorus pentanitride \(\left( {{{\rm{P}}_3}{{\rm{N}}_5}} \right).\)

Compounds of Nitrogen Group Elements

Nitrogen is a group \(15\) element. The other elements in this group include phosphorus, arsenic, antimony, bismuth and moscovium (a highly radioactive element). They form a wide range of compounds when combined with other elements.

For example, nitrogen forms common compounds such as ammonia \(\left( {{\rm{N}}{{\rm{H}}_3}} \right),\) nitric acid \(\left( {{\rm{HN}}{{\rm{O}}_3}} \right),\) oxides of nitrogen, nitrogen hydrides, etc. Similarly, common compounds of phosphorous include phosphine, phosphorous halides such as phosphorous trichloride \(\left( {{\rm{PC}}{{\rm{l}}_3}} \right),\) phosphorous pentachloride \(\left( {{\rm{PC}}{{\rm{l}}_5}} \right)\) and oxoacids of phosphorus such as orthophosphorous acid \(\left( {{{\rm{H}}_3}{\rm{P}}{{\rm{O}}_3}} \right),\) orthophosphoric acid \(\left( {{{\rm{H}}_3}{\rm{P}}{{\rm{O}}_4}} \right),\) etc.

The compounds formed by arsenic include arsenic trioxide \(\left( {{\rm{A}}{{\rm{s}}_2}{{\rm{O}}_3}} \right),\) sodium arsenite \(\left( {{\rm{NaAs}}{{\rm{O}}_2}} \right),\) arsenic acid \(\left( {{{\rm{H}}_3}{\rm{As}}{{\rm{O}}_4}} \right),\) arsenic trichloride \(\left( {{\rm{AsC}}{{\rm{l}}_3}} \right),\) etc. Bismuth generally forms trivalent and pentavalent compounds. The compounds made of bismuth include bismuth oxide \(\left( {{\rm{B}}{{\rm{i}}_2}{{\rm{O}}_3}} \right),\) bismuth chloride \(\left( {{\rm{BiC}}{{\rm{l}}_3}} \right),\) Bismuth oxychloride \(\left( {{\rm{BiOCl}}} \right),\) bismuthine \(\left( {{\rm{Bi}}{{\rm{H}}_3}} \right),\) etc.

Summary

Nitrogen forms a large number of compounds. In this article, we have learnt about the compounds of nitrogen in detail. Their physical and chemical properties and their uses in industrial fields and daily life. We have also learnt about the compounds formed between nitrogen and sulphur, nitrogen and phosphorus and the compounds of group \(15\) elements.

FAQs on Compounds of Nitrogen

Q.1. What are examples of nitrogen compounds?
Ans: Nitrogen is a group \(15\) element. It forms many compounds. Some examples of nitrogen compounds are ammonia, nitric acid, oxides of nitrogen, etc.

Q.2. What are nitrogenous organic compounds?
Ans: Organic compounds that contain nitrogen are termed nitrogenous organic compounds. The nitrogenous organic compounds include amines, amides, alkyl nitrates, nitrosamines, nitroarenes, etc.

Q.3. Why are nitrogen compounds important?
Ans: Nitrogen forms some important compounds that play a major role in different aspects of life. That is, they are involved in the industrial production of fertilizers, and some of the forms the building blocks of life. Therefore, nitrogen compounds are important.

Q.4. What substances contain nitrogen?
Ans: Nitrogen is the most abundant gas in the earth’s atmosphere. It naturally occurs as sodium nitrate (called Chile saltpetre) and potassium nitrate (Indian saltpetre). In plants and animals, it is found in the form of proteins.

Q.5. Where are nitrogen compounds found?
Ans: Nitrogen compounds are found in plants and animals. They include amino acids, amides, amino acids, proteins, polyamines, etc.

Q.6. Why is nitrogen gas called Azota?
Ans: Nitrogen is one of the components of air. It does not support life on its own. Based on this, Antoine Lavoisier called nitrogen ‘azota’, which has the meaning ‘lifeless’.