9.4 Nitrogen

Reason for the inertness of nitrogen

The nitrogen molecule contains a triple bond (N≡N) — one sigma bond and two pi bonds — making it one of the strongest bonds in chemistry. Its key data:

This inertness explains why nitrogen is abundant in the atmosphere but scarce in inorganic minerals — it simply refuses to react and form compounds unless forced by extreme conditions (high temperature, high pressure, or electrical discharge).

Active nitrogen

When an electrical discharge is passed through low-pressure N₂ gas, the molecules absorb energy and dissociate into individual nitrogen atoms. These recombine in excited electronic states, producing a characteristic yellow-orange glow that persists even after the gas moves away from the discharge region. This highly reactive form is called active nitrogen.

Active nitrogen is far more reactive than ordinary N₂ because the atoms carry excess energy. Other natural sources of active nitrogen include:

• Lightning strikes in the atmosphere
• Reaction of nitrous oxide (N₂O) with excited oxygen atoms (O*) in the stratosphere

Ammonia (NH₃) is a colourless gas with a sharp, pungent smell.

Its central nitrogen atom carries a lone pair of electrons, which is the key to understanding almost all of ammonia’s chemistry — it makes NH₃ a Lewis base and gives it the ability to donate electrons, dissolve readily in water, and form complexes with metal ions.

1. Basic Nature

• Ammonia dissolves in water, giving OH⁻ ions.

Thus, the aqueous solution of ammonia (aq. NH₃), ammonium hydroxide, and ammonia solution mean the same thing.

• Ammonia changes the colour of moist litmus blue.
• Ammonia reacts with acids, forming salts.

2. Oxidation reactions

Ammonia can be oxidised — its nitrogen is in the −3 oxidation state, which can be raised.

a) Combustion

Ammonia burns in an atmosphere of oxygen with a greenish-yellow flame, producing N₂ and H₂O.

b) Formation of nitric oxide

When a mixture of ammonia and oxygen gas is passed over heated platinum gauze (800 °C), it is oxidised to nitric oxide.

[Practice writing the ONs of all the atoms in the above equations.]

3. Complex formations

Salts of d-block elements like copper, silver, nickel, cobalt, cadmium, etc., form soluble complexes with ammonium hydroxide.

e.g., white precipitate of AgCl dissolves in ammonia solution, giving a soluble complex.

CuSO₄ solution gives a bluish-white ppt. on treatment with ammonia. On adding excess ammonia, it dissolves, giving a deep-blue solution of tetraamine copper sulphate (a complex compound).

4. Precipitation reactions

Salts of heavy metals like Iron, Chromium, Aluminium, zinc, etc., are precipitated as their hydroxides from their solutions.

These reactions serve as confirmatory tests for the presence of respective metal ions in the solution.

5. Reaction with Mercurous nitrate paper

Ammonia forms a black precipitate with mercurous salts.

This blackening test is a quick confirmatory test for ammonia gas.

Applications of Ammonia

Ammonia is one of the most widely produced industrial chemicals in the world. It is used:

• To produce urea, ammonium sulphate & other nitrogen fertilizers.
• To produce nitric acid, sodium carbonate, etc.
• As a laboratory reagent in the form of ammonium hydroxide.
• As a refrigerant in industrial cooling systems.

Despite its industrial importance, ammonia is highly toxic and must be handled with care.

In the laboratory and industry, ammonia must always be handled in well-ventilated areas or under a fume hood. In case of exposure, flush immediately with large quantities of water and seek medical attention.

There are two oxyacids of nitrogen, viz. nitrous acid (HNO₂) and nitric acid (HNO₃).

The oxidation numbers of nitrogen in these two acids are +3 and +5, respectively. N₂O₃ and N₂O₅ are the anhydrides of nitrous acid and nitric acid, respectively.

Nitrous acid is relatively unstable and cannot be stored. It is prepared fresh by mixing solutions of NaNO₂ and HCl as

It easily disproportionates into HNO₃ and NO as

Nitric acid (HNO₃) is a colourless, fuming liquid that turns yellow on prolonged exposure to light (due to partial decomposition into NO₂). It is both a strong acid and a powerful oxidising agent.

[A] Acidic Nature

Being a very strong acid, it dissociates completely in an aqueous solution.

Action on bases: HNO₃ neutralises bases, forming nitrate salts.

Action with carbonates and bicarbonates: HNO₃ liberates CO₂ gas in reaction with metal carbonates and bicarbonates.

Action with highly electropositive metals: Very dilute HNO₃ reacts with only Mg and Mn to produce salt and hydrogen gas.

The hydrogen gas produced is not pure; it is accompanied to some extent by gaseous reduction products of acids (i.e., NO, NO₂, etc.).

[B] Oxidising Nature

Nitric acid is a very good oxidising agent.

1. Action on metals

This reacts with almost all metals, even with some that lie below hydrogen in the electrochemical series, except noble metals like Pt and Au.

The reactions with active metals are balanced by the nascent hydrogen formation method in the partial equation method.

a. Action with Zn

With conc. HNO₃

With moderately conc. HNO₃ (1:1)

With dilute HNO₃

With very dilute HNO₃

b. Action with Fe

• Highly concentrated nitric acid renders iron passive due to the formation of a layer of ferroso-ferric oxide Fe₃O₄. This phenomenon is used to prevent the rusting of iron.

With moderately conc. HNO₃ (1:1)

With very dilute HNO₃

c. Action with Cu

Conc. HNO₃ gives NO₂ gas

(1:1) HNO₃ gives NO gas.

Summary table to remember easily:

Metal	HNO₃ conc.	Reduction product
Zn	Conc.	NO₂
Zn	Mod. conc.	NO
Zn	Dilute	N₂O
Zn	Very dilute	NH₄NO₃
Cu	Conc.	NO₂
Cu	Dilute	NO

2. Action on Non-metals

Hot and conc. Nitric acid oxidises many nonmetals, such as carbon, sulfur, phosphorus, iodine, etc., to form their oxides or oxyacids.

a. With carbon (C is oxidised to carbonic acid)

b. With sulphur (S is oxidised to sulphuric acid)

3. Action on reducing compounds

a. It oxidises H₂S to sulphur

b. It oxidises SO₂ to H₂SO₄

(Practice writing all these reactions several times. These are very important)

For the test of nitric acid or nitrate ion in a salt, the solution is first mixed with an equal volume of conc. H₂SO₄ in a test tube. The solution is cooled under tap water, and a freshly prepared FeSO₄ solution is added slowly from the wall of the test tube, tilting the test tube. The formation of a brown ring confirms the presence of nitrate ion or nitric acid in the solution.