5. Nitric acid

NITRIC ACID (HNO₃)

Preparation :

(i)    In the laboratory, nitric acid is prepared by heating KNO₃ or NaNO₃ and concentrated H₂SO₄ in a glass retort.

NaNO₃ + H₂SO₄  —® NaHSO₄ + HNO₃ 

(ii)   On a large scale it is prepared mainly by Ostwald’s process.

This method is based upon catalytic oxidation of NH₃ by atmospheric oxygen.

4 NH₃ (g) + 5 O₂ (g) (from air)  4 NO (g) + 6 H₂O (g)

  Nitric oxide thus formed combines with oxygen giving NO₂.

 2 NO (g) + O₂ (g)  2 NO₂ (g)

Nitrogen dioxide so formed, dissolves in water to give HNO3.

3 NO₂ (g) + H₂O (l) —® 2 HNO₃ (aq) + NO (g)

NO thus formed is recycled and the aqueous HNO₃ can be  concentrated by distillation upto ~ 68% by mass. Further concentration to 98% can be achieved by dehydration with concentrated H₂SO₄.

Properties :

Physical properties :

It is a colourless liquid. Freezing point is 231.4 K and boiling point is 355.6 K. Laboratory grade nitric acid contains ~ 68% of the HNO₃ by mass and has a specific gravity of 1.504.

In the gaseous state, HNO₃ exists as a planar molecule.

In aqueous solution, nitric acid behaves as a strong acid giving hydronium and nitrate ions.

HNO₃ (aq) + H₂O (l) —® H₃O⁺ (aq) + NO₃^– (aq)

(i) Concentrated nitric acid is a strong oxidising agent and attacks most metals except noble metals such as gold and platinum. The products of oxidation depend upon the concentration of the acid , temperature and the nature of the material undergoing oxidation.

Some metals (e.g., Fe, Cr , Al) do not dissolve in concentrated nitric acid because of the formation of a passive film of oxide on the surface.

1.  4Zn + 10HNO₃ (dilute) —® 4Zn(NO₃)₂ + N₂O  + 5H₂O

 Zn + 4HNO₃ (concentrated) —® Zn(NO₃)2 + 2NO₂ + 2H₂O

2.  3Cu + 8HNO₃ (dilute) —® 2NO + Cu(NO₃)₂ + 4H₂O

Cu + 4HNO₃ (concentrated)  —® 2NO₂ + Cu(NO₃)₂ + 2H₂O

(ii)   Concentrated nitric acid also oxidises non–metals and their compounds. Iodine is oxidised to iodic acid , carbon to carbon dioxide , sulphur to H2SO4 and phosphorus to phosphoric acid.

I₂ + 10 HNO₃ —® 2 HIO₃ + 10 NO₂ + 4 H₂O

C + 4 HNO₃ —® CO₂ + 2 H₂O + 4 NO₂ 

S₈ + 48 HNO₃ (concentrated) —® 8 H₂SO₄ + 48 NO₂ + 16 H₂O

P₄ + 20 HNO₃ (concentrated) —® 4 H₃PO₄ + 20 NO₂ + 4 H₂O

Brown Ring Test :

The familiar brown ring test for nitrates depends on the ability of Fe²⁺ to reduce nitrates to nitric oxide, which reacts with Fe²⁺ to form a brown coloured complex. The test is usually carried out by adding dilute ferrous sulphate solution to an aqueous solution containing nitrate ion , and then carefully adding concentrated sulphuric acid along the sides of the test tube. A brown ring at the interface between the solution and sulphuric acid layers indicate the presence of nitrate ion in solution.

 NO₃^– + 3 Fe²⁺ + 4H⁺ —® NO + 3Fe³⁺ + 2 H₂O

[Fe (H2O)₆]²⁺ + NO  —® [Fe (H₂O)₅ (NO)]²⁺ + H₂O  

Uses :

The major use of nitric acid is in the manufacture of ammonium nitrate for fertilizers and other nitrates for use in explosives and pyrotechnics. It is also used for the preparation of nitroglycerin, trinitrotoluene and other organic nitro compounds. Other major uses are in the pickling of stainless steel, etching of metals and as an oxidiser in rocket fuels.