Nitrogen and it compounds (O-level chemistry)

Nitrogen and its compounds (O-level)

Source: Air

Industrial preparation of nitrogen

By distillation of air from which carbon dioxide and water vapor have been removed. Nitrogen distills of leaving oxygen as the residue. (oxygen has a higher boiling point than nitrogen)

Laboratory preparation

By oxidation of ammonia with copper oxide

Set up

Equation

2NH₃(g) + 3CuO (s)  →   N₂ (g) + 3H₂O(l) + 3Cu(s)

Physical properties on nitrogen

• It is colorless
• Odorless
• Nonreactive

Chemical properties of nitrogen

Nitrogen does not react easily because the nitrogen atoms are bonded by strong triple covalent bonds

However, under drastic conditions, it reacts to form compounds. For instance, nitrogen reacts with magnesium and lithium to form nitrides. This is because the formation of nitrides liberates a lot of heat that breaks the strong nitrogen-nitrogen triple bond

3Mg(s) + N₂(g)   →    Mg₃N₂ (s)

6Li(s) + N₂ (g)     →   2Li₃N(s)

The nitrides hydrolyze in water to give ammonia

Mg₃N₂(s) + 6H₂O(l) →  3Mg(OH)₂ (s) + NH₃ (g)

Ammonia is an alkaline gas that turns litmus solution blue or damp red litmus paper blue.

Note that magnesium burns in air to produce magnesium oxide as well.

Mg(s) + O₂(g)     →   2MgO(s)

Magnesium oxide is a white solid that is insoluble in water.

Exercise 1

Magnesium burns in air to produce two solids A and B. A is insoluble in water while B reacts with water to form a gas that turns damp red litmus paper blue

• Identify Solids A and B.
• Write equations leading to the formation of A and B.
• Write an equation between B and water.

Uses of nitrogen

1. The raw material for synthesis of ammonia and nitric acid
2. Coolant

Compounds of nitrogen

Ammonia, NH₃

Industrial preparation

By reaction nitrogen and hydrogen

N₂(g) + 3H₂(g)    ↔   2NH₃(g) + heat

Conditions for the production of ammonia.

1. Pressure:

Production of ammonia is accompanied with a decrease in the number of moles of a gas and as such high yield of ammonia is favored by high pressure.

2. Temperature

Production of ammonia produces heat; high yields are thus expected at low temperatures.

However, at low temperature, the reaction is slow that a compromise temperature of 450⁰-500⁰C is used.

3. Catalyst is iron speed up the reaction

Laboratory preparation of ammonia

By reacting calcium hydroxide (Ca(OH)₂) with ammonium chloride (NH₄Cl)

Equation

Ca(OH)₂ (s) + 2NH₄Cl (s)    →   CaCl₂(aq) + 2H₂O (l) + 2NH₃ (g)

Properties of ammonia

(i) Has chocking smell

(ii) Turns damp red litmus paper from red to blue.

(iii) It is very soluble in water

    This solubility is demonstrated by a Fountain experiment.

(a) The hard glass flask is filled with ammonia and taps A and B are closed

(b) The flask is inverted in water as shown above.

(c) Tap B is open water dissolves ammonia to create a vacuum inside the flask

(d) Immediately tap A is open, water splashes into the flask like a fountain.

(iv) Ammonia as a reducing agent

It reduces black copper oxide to brown copper and orange lead II oxide to grey lead metal.

(v) Oxidation of ammonia

Ammonia burns with a blue flame in oxygen.

Equation

4NH₃ (s)  + 3O₂ (s)  →    2N₂(g) + 6H₂O (l)

In presence of excess O₂ and platinum or copper wire catalyst, ammonia is oxidized to nitrogen monoxide and the brown fumes on nitrogen dioxide.

4NH₃ (s)  + 5O₂ (s)   →    4NO(g) + 6H₂O (l)

2NO₂ (s)  + O₂ (s)     →    2NO₂(g)

(vi) The reaction of ammonia solution with common cations

(a) Form a white precipitate with lead (II) and aluminium salts that are insoluble in excess

NH₃ (aq) + H₂O (l)     →       NH₄⁺(aq) + OH^– (aq)

Then

Pb²⁺(aq) + 2OH^–(aq)        →           Pb(OH)₂(s)

Al³⁺(aq) + 3OH^–(aq)          →           Al(OH)₃(s)

(b) Forms white precipitate with zinc II salts soluble in excess due to formation of a complex ion

Zn²⁺(aq) + 2OH^–(aq)         →                           Zn(OH)₂ (s)

Then

Zn(OH)₂ (s) + 4NH₃ (aq)  (excess)        →          Zn(NH₃)₄²⁺(aq)

 

(c) Forms blue precipitate with copper II salts soluble in excess to give a deep blue solution due to formation of a complex ion

Cu²⁺(aq) + 2OH_–(aq)        →                           Cu(OH)₂ (s)

Then

Cu(OH)₂ (s) + 4NH₃ (aq) (Excess)  →         Cu(NH₃)₄²⁺(aq)

Uses of ammonia

• Disinfectant
• For the preparation of liquid soap
• Fertilizers

Note that prolonged use of ammonia nitrate fertilizer may make the soil acidic because ammonium nitrate in soil reacts with water forming ammonium hydroxide and nitric acid plants absorb ammonium ions leaving nitric acid in solution.

NH₄NO₃(s) + H₂O(l)         →           NH₄OH (aq) + HNO₃ (aq)

Ammonium salts

Common ammonium salt include ammonium chloride (NH₄Cl), ammonium sulphate ((NH₄)₂SO₄) ammonium carbonate ((NH₄)₂CO₃) and ammonium nitrate (NH₄NO₃)

1. Ammonium chloride, ammonium carbonate, and ammonium sulphate decompose to give alkaline gas ammonia (turns damp red litmus paper blue). Ammonium chloride sublimes the rest do not.

2. Ammonium salts react with sodium hydroxide to liberate ammonia

NH₄⁺(aq) + OH^–(aq)          →           NH₃(g) + H₂O(l)

Ammonia is an alkaline gas that turns damp red litmus paper blue.

Nitric acid

Industrial preparation

Ammonia is oxidized to nitrogen dioxide in the presence of platinum catalyst.

4NH₃ (s)  + 7O₂ (s)            →           4NO₂(g) + 6H₂O (l)

The nitrogen dioxide is dissolved in water in presence of oxygen to produce nitric acid

4NO₂ (g) + 2H₂O (l)  + O₂ (g)         →          4HNO₃ (aq)

Laboratory preparation

By heating potassium nitrate with concentrated sulphuric acid

Equation

H₂SO₄ (aq) + 2KNO₃ (s)  →           HNO₃(aq) + KHSO₄(aq)

Properties of nitric acid

1. Turns blue litmus red
2. Liberate carbon dioxide from carbonates and hydrogen carbonates

HCO₃^–  + H⁺ (aq)    →       CO₂(g) + H₂O(l)

CO₃^2-  + 2H⁺ (aq)    →       CO₂(g) + H₂O(l)

3. It reacts with metals to produce hydrogen

Mg (s) + 2H⁺(aq)    →      Mg²⁺(aq) + H₂(g)

4. It oxidizes nonmetals

S(s) + 6HNO₃(aq)   →      H₂SO₄ (aq) + 6NO₂(aq) + 2H₂O (l)

C(s) + 4HNO₃(aq)  →       CO₂ (s) + 6NO₂(aq) + 2H₂O (l)

5. It decomposes on heating to form nitrogen dioxide and oxygen

4HNO₃(l)      (Heat) →     4NO₂ (g) + O₂ (g) + 2H₂O(l)

Nitrogen dioxide is soluble in water where oxygen is insoluble in water. Therefore, when fuming nitric acid is heated, the gas collected over water is oxygen.

6. It oxidizes metals

It oxidizes copper and lead to copper II nitrate and lead II nitrate respectively.

3Cu(s) + 8HNO₃(aq) (dil) →  3Cu(NO₃)₂(aq) + 2NO(g) + 4H₂O (l)

3Pb(s) + 8HNO₃(aq) (dil)  →   3Pb(NO₃)₂(aq) + 2NO(g) + 4H₂O (l)

Cu(s) + 4HNO₃(aq)(conc.)  →   Cu(NO₃)₂(aq) + 2NO₂(g) + 2H₂O (l)

Pb(s) + 4HNO₃(aq) (conc) →   Pb(NO₃)₂(aq) + 2NO₂(g) + 2H₂O (l)

Nitrates

These are salt of nitric acid

Decomposition of nitrates

1. Nitrates of group 1 elements decompose on heating to give nitrites and oxygen

2NaNO₃ (s)     →        2NaNO₂ (s) + O₂ (g)

2. Nitrates of other elements decompose to liberate oxide, nitrogen dioxide, and oxygen.

2Cu(NO₃)₂ (s)     →      2CuO(s) + 2NO₂(g) + O₂ (g)

2Pb(NO₃)₂ (s)     →      2PbO(s) + 2NO₂(g) + O₂ (g)

3. Silver nitrate decomposes on heating  to liberate silver metal, nitrogen dioxide, and oxygen

2Ag(NO₃)₂ (s)         →                 2Ag (s) + 2NO₂(g) + O₂ (g)

4. Ammonium nitrate decomposes on heating  into dinitrogen oxide and water

NH₄NO₃   →   N₂O(g) + 2H₂O(l)

                Testing for nitrates

1. Nitrates are identified from the brown fumes when nitrates are heated strongly.
2. Brown ring test

To 1 cm³ of a nitrate solution, add freshly prepared iron II sulphate, followed sulphuric acid slowly over the wall of a slanted test tube.

Observation

A brown ring forms at the junction of the two solutions.

Uses of nitric acid

For fertilizers, dyes, explosives.

For sample questions and answers download PDF below

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Compiled By Dr. Bbosa Science