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nitric acid

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Nitric Acid and Nitrates 1. Concentrated nitric acid is commonly called aqua fortis. 2. In the laboratory, nitric acid is prepared by heating potassium nitrate or sodium nitrate with conc. sulphuric acid in a retort, when nitric acid vapours distil over in a receiver to form nitric acid in liquid state. less than KHSO4 + HNO3 KNO3 + H2SO4 200 C more than KNO3+ KHSO4 K2SO4 + HNO3 200 C 3. The nitric acid prepared in the laboratory is generally light yellow in colour. It is because nitric acid vapours partially decompose to form nitrogen dioxide gas, which redissolves in acid and gives it a yellow colour. 4. Industrially, nitric acid is prepared by Ostwald s process. (i) A mixture of two parts of pure oxygen and one part of pure ammonia is passed over heated platinum gauze at 900 C, when the following exothermic reaction takes place with the formation of nitric oxide gas. pt 900 6C 4NH3 + 5O2 4NO + 6H2O (ii) The mixture of gases so formed, is cooled and then led into oxidation chamber, where excess of oxygen reacts with nitric oxide to form nitrogen dioxide. 4NO + 2O2 4NO2 (iii) The mixture of nitrogen dioxide and excess of oxygen is then led to absorption chamber, in which fine spray of water trickles from top. The following reaction takes place with the formation of dilute nitric acid, which collects on the bottom of chamber. 4NO2 + 2H2O + O2 4HNO3 (iv) The dilute nitric acid is concentrated under reduced pressure. 1 Key Points 5. Physical properties of nitric acid : (i) Pure nitric acid is a colourless, fuming liquid, hygroscopic in nature. (ii) It has a choking smell and sour taste. (iii) It stains skin yellow and causes very painful burns. (iv) Its b.p is 86 C. It is highly soluble in water. The 68% nitric acid forms a constant boiling liquid. 6. Chemical properties of nitric acid : (i) It decomposes on heating to form nitrogen dioxide and oxygen. heat 2H2O + 4NO2 + O2 4HNO3 (ii) It reacts with metallic oxides and hydroxides to form their respective nitrates and water. CuO + 2HNO3 (dil.) Cu(NO3)2 + H2O Ca(OH)2 + 2HNO3 (dil.) Ca(NO3)2 + 2H2O (iii) It reacts with metallic carbonates to form their respective nitrates, carbon dioxide gas and water. MgCO3 + 2HNO3 (dil.) Mg(NO3)2 + H2O + CO2 (iv) Conc. nitric acid reacts with carbon, sulphur and phosphorus to form oxides or oxyacids, nitrogen dioxide and water. boiling C + 4HNO3 (conc.) CO2 (g) + 4NO2 (g) + 2H2O boiling S + 6HNO3 (conc.) H2SO4 + 2H2O + 6NO2 (g) boiling P + 5HNO3 (conc.) H3PO4 + H2O + 5NO2 (g) (v) 1% nitric acid reacts with magnesium and manganese to liberate hydrogen. Mg + 2HNO3 (v. dil.) Mg(NO3)2 + H2 (vi) 1% nitric acid reacts with zinc to form ammonium nitrate. 4Zn + 10HNO3 (v. dil.) 4Zn(NO3)2 + 3H2O + NH4NO3 (vii) Dilute nitric acid (20%) reacts with metals to form metallic nitrate, nitric oxide and water. However, in case of zinc, nitrous oxide is formed. 3Mg + 8HNO3 (dil.) 3Mg(NO3)2 + 4H2O + 2NO (g) 4Zn + 10HNO3 (dil.) 4Zn (NO3)2 + 5H2O + N2O 2 Key Points (viii) Concentrated nitric acid (60%) reacts with metals to form metallic nitrates, nitrogen dioxide and water. Cu + 4HNO3 (conc.) Cu(NO3)2 + 2H2O + 2NO2 (ix) Iron, cobalt and nickel become passive when treated with conc. nitric acid. (x) Pure nitric acid does not react with gold or platinum. However, the above metals dissolve in aqua regia. (xi) Moderately, concentrated nitric acid oxidises : (i) Hydrogen sulphide gas to sulphur. H2S + 2HNO3 2H2O + 2NO2 + S (ii) Potassium iodide to iodine. 6KI + 8HNO3 6KNO3 + 4H2O + 2NO + 3I2 (iii) Sulphur dioxide to sulphuric acid. 2HNO3 + SO2 H2SO4 + 2NO2 (iv) Acidified ferrous sulphate solution to ferric sulphate. 6FeSO4 + 3H2SO4 + 2HNO3 3Fe2(SO4)3 + 4H2O + 2NO (v) Hydrochloric acid to chlorine. 3HCl + HNO3 2H2O + NOCl + Cl2 (xii) It reacts with organic compounds by replacing hydrogen atom with NO2 radical. C6H6 + HNO3 C6H5NO2 + H2O 7. The salts of nitric acid with metals are called nitrates. Only nonmetallic nitrate is ammonium nitrate. 8. The nitrates can be prepared by the action of dilute nitric acid on (i) metallic oxide (ii) metallic hydroxide (iii) metallic carbonates (iv) metallic sulphites (v) metals. 9. All nitrates are crystalline in nature, do not contain water of crystallisation and are soluble in water. All metallic nitrates are white in colour, except iron nitrate and copper nitrate. 3 Key Points 10. All metallic nitrates decompose on heating. (i) Ammonium nitrates decompose to form nitrous oxide and steam and leave behind no residue. heat N2O + 2H2O NH4NO3 (ii) Sodium nitrate and potassium nitrate melt on heating and decompose to form their respective nitrites and oxygen. 2NaNO3 2NaNO2 + O2 (iii) Nitrates of calcium, magnesium, aluminium, zinc, lead, tin and copper on heating decompose to form their oxides, nitrogen dioxide and oxygen. 2Pb(NO3)2 2PbO + 4NO2 + O2 (iv) Nitrates of silver and mercury decompose on heating to form a metal, nitrogen dioxide and oxygen. Hg(NO3)2 Hg + 2NO2 + O2 11. Uses of nitric acid : 1. It is used in making silver nitrate, which is required in the photographic industry. 2. It is used in the manufacture of fertilisers. 3. It is used in the manufacture of explosives like T.N.T and T.N.G. 4. It is used in making dyes and drugs. 5. It is used for etching designs in metals. 6. It is used as a rocket fuel. 12. Tests of nitric acid and nitrates : 1. All nitrates on heating with conc. sulphuric acid and small amount of copper turnings liberate reddish-brown gas nitrogen dioxide. 2. All nitrates when treated with freshly prepared ferrous sulphate solution and conc. sulphuric acid form a brown ring at the junction of liquids. 4 Key Points

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