ICSE Class 10 Chemistry Chapter 10 Study of Compounds Nitric Acid

Study of Compounds - Nitric Acid

Syllabus

Nitric Acid: One laboratory method of preparation of nitric acid from potassium nitrate or sodium nitrate; Nitric acid as an oxidising agent.

Nitric Acid: Laboratory method of preparation of nitric acid from potassium nitrate or sodium nitrate; the laboratory method can be studied in terms of reactant, product, condition, equation, setting, diagram, precaution, collection, identification. As an oxidising agent: its reaction with copper, carbon, sulphur.

Molecular Formula and Structure

Molecular formula: HNO3

Relative molecular mass: 63

Structure of nitric acid (Nitrogen shows valency of 5):

H - O - N=O with another O below

Formation of Nitric Acid in Atmosphere

1. During lightning discharge, the nitrogen present in the atmosphere reacts with the oxygen to form nitric oxide.

\[N_2 + O_2 \rightleftharpoons 2NO\]

2. Nitric oxide is further oxidised to nitrogen dioxide.

\[2NO + O_2 \rightleftharpoons 2NO_2\]

3. This nitrogen dioxide dissolves in atmospheric moisture or rain water in the presence of oxygen of the air and forms nitric acid in the free state, which is washed down by the rain and combines with the salt present on the surface of the earth.

\[4NO_2 + 2H_2O + O_2 \rightarrow 4HNO_3\]

The conversion of free atmospheric nitrogen into useful nitrogenous compounds in the soil is known as fixation of atmospheric nitrogen.

10.1 Introduction

The acid was formerly known as aqua fortis meaning strong water. It is so called because it reacts with nearly all metals. It can even dissolve silver which does not dissolve in other acids.

10.2 Discovery

Glauber, in 1658, obtained nitric acid by distilling nitre, (potassium nitrate, KNO3) with sulphuric acid.

Lavoisier, in 1776, proved that nitric acid contains oxygen.

Cavendish, in 1784, proved that it also contains hydrogen and nitrogen in addition to oxygen.

10.3 Occurrence

In the free state, nitric acid is found in rain water, where it occurs in traces after lightning.

In the combined state, it is found in the form of metallic nitrates such as Chile saltpetre (NaNO3), nitre (KNO3), or calcium nitrate [Ca(NO3)2].

10.4 Laboratory Preparation of Nitric Acid

Reactants: It is obtained by distilling conc. sulphuric acid with nitrates of potassium KNO3 (nitre) or sodium NaNO3 (Chile saltpetre).

Procedure: A mixture of equal parts, by weight, of potassium/sodium nitrate and concentrated sulphuric acid is heated gently to 180-200 degrees Celsius in a glass retort (Fig 10.1). Sulphuric acid is a non-volatile acid and produces volatile nitric acid on reacting with potassium nitrate or sodium nitrate.

Conc. sulphuric acid and sodium or potassium nitrate are heated at a temperature that should not exceed 200 degrees Celsius, because sodium sulphate formed at higher temperature forms a hard crust which sticks to the walls of the retort and is difficult to remove, although the yield of HNO3 is higher.

\[NaNO_3 + NaHSO_4 > 200°C \rightarrow Na_2SO_4 + HNO_3\]

The higher temperature (i) may damage the glass apparatus, (ii) decomposition of nitric acid can also occur, (iii) Wastage of fuel.

Reactions

Collection: The vapours of nitric acid are condensed to a light yellow liquid by chilling the receiver with running cold water.

Note: Pure acid is colourless but the acid obtained in the laboratory is slightly yellow. The yellow colour is due to dissolution of reddish brown coloured nitrogen dioxide gas in the acid. This gas is produced due to the thermal decomposition of a portion of nitric acid.

\[4HNO_3 \rightarrow 2H_2O + 4NO_2 + O_2\]

The Yellow Colour of the Acid is Removed

(i) If dry air or CO2 is bubbled through the yellow acid, the latter turns colourless because it drives out NO2 from warm acid which is further oxidised to nitric acid.

(ii) By addition of excess of water, nitrogen dioxide gas dissolves in water and thus the yellow colour of the acid is removed. The reaction is the reverse of the decomposition of nitric acid.

Precautions

(1) All glass apparatus is used because nitric acid vapours attack rubber and cork.

(2) Conc. HCl is not used in place of conc. H2SO4 because HCl is volatile and hence nitric acid vapours will carry HCl vapours.

(3) The temperature of the reaction should not exceed 200 degrees Celsius because sodium sulphate formed at higher temperature forms a hard crust which sticks to the walls of the retort and is difficult to remove, although the yield of HNO3 is higher.

10.5 Properties of Nitric Acid

A. Physical Properties

(1) Colour: It is a colourless liquid (98% concentration) but the commercial nitric acid (68% concentration) is yellowish brown in colour.

(2) Odour: Suffocating smell.

(3) Taste: Acidic (sour taste).

(4) Nature: Hygroscopic and fumes in air. Hence, the bottle containing nitric acid should always be stoppered.

(5) Density: 1.54 g/cm3 (98%) but the commercial nitric acid (68%) has lower density 1.42 g/cm3.

(6) Boiling point and melting point: It boils at 86 degrees Celsius and melts at -42 degrees Celsius.

An aqueous solution of nitric acid (68% concentration) forms a constant boiling mixture at 121 degrees Celsius.

(7) Solubility: Soluble in water in all proportions.

(8) Physiological action: Non-poisonous. It has a corrosive action on the skin and causes painful blisters. It stains the skin yellow as it reacts with protein of the skin and forms xanthoproteic acid.

B. Chemical Properties

(1) Stability: Pure nitric acid is unstable to heat or sunlight. It decomposes to give yellow solution due to the formation of nitrogen dioxide. Thus, 100% nitric acid is not generally used.

\[4HNO_3 \xrightarrow{\text{Sunlight}} 4NO_2 + 2H_2O + O_2\]

In the presence of sunlight, nitric acid decomposes even at room temperature. The nitric acid stored in a bottle turns yellow. This colour is due to dissolved NO2 in HNO3. To avoid the decomposition, nitric acid is normally stored in coloured bottles.

(2) Acidic properties: Nitric acid is a very strong monobasic acid.

HNO3 ionises almost completely in aqueous solution to produce hydrogen ions and nitrate ions.

\[HNO_3(aq) \rightleftharpoons H^+ + NO_3^-\]

(i) It turns:

blue litmus red, methyl orange pink, phenolphthalein remains colourless.

(ii) Reaction with alkalis:

It neutralises alkalis to form salt and water. All metallic oxides and hydroxides react with dilute nitric acid to form their respective soluble metallic nitrates and water only.

(iii) Reaction with carbonates and bicarbonates:

It reacts with carbonates and bicarbonates to give salt, water and carbon dioxide.

(iv) Action with metallic sulphites and bisulphites:

Metallic sulphites and bisulphites react with dilute nitric acid to form their soluble metallic nitrates, water and sulphur dioxide gas.

(3) Oxidising properties: Nitric acid vigorously oxidises non-metals, metals, inorganic compounds and organic substances.

Nitric acid itself undergoes reduction to form different reduction products like NO, NO2, N2O, etc. These products depend on the acid concentration and the temperature of the reaction.

Oxidising properties are due to nascent oxygen which it gives on decomposition.

\[2HNO_3 \text{(conc.)} \rightarrow 2NO_2 + H_2O + [O]\]

\[2HNO_3 \text{(dil.)} \rightarrow 2NO + H_2O + 3[O]\]

Nitric acid is a powerful oxidising agent and the nascent oxygen formed on decomposition oxidises hydrogen to water.

Teacher's Note

Nitric acid is commonly used in laboratories and industries for cleaning metal surfaces and etching designs. Understanding its properties helps in appreciating why it must be handled with care and stored properly.

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