Study of Compounds

Na₂CO₃ + 2HCl [dil.] ⟶ 2NaCl + H₂O + CO₂ [g.]

Hydrogen chloride gas is denser than NH₃.
Reason — V.D. of NH₃ = 8.5 and that of HCl = 18.25

HCl gas is collected by upward displacement of air.

NH₃ [g] + HCl [g] ⟶ NH₄Cl (s-solid)

When a gas jar containing HCl gas in inverted over a jar full of ammonia gas, dense white fumes are formed. These fumes consist of particles of solid ammonium chloride suspended in air.

(i) Fe + 2HCl ⟶ FeCl₂ + H₂

(ii) NaHCO₃ + HCl ⟶ NaCl + H₂O + CO₂

(iii) FeS + 2HCl ⟶ FeCl₂ + H₂S

(iv) Na₂SO₃ + 2HCl ⟶ 2NaCl + H₂O + SO₂ [g]

Hydrogen chloride is 1.28 times heavier than air, hence it is collected by downward delivery [upward displacement]. Hydrogen chloride is not collected over water since it is highly soluble in water.

(i) Na₂S₂O₃ + 2HCl ⟶ 2NaCl + SO₂ + S + H₂O

(ii) Pb(NO₃)₂ + 2HCl ⟶ PbCl₂ ↓ + 2HNO₃

Hydrogen chloride (HCl)

$\text{NaCl} + \text{H}_2\text{SO}_4 \xrightarrow{\lt 200 ^{\circ}\text{C}} \text{NaHSO}_\text{4} + \text{HCl [g]}$

(i) [C] Hydrogen chloride
(ii) [B] Silver chloride

(i) $\text{Pb}_3\text{O}_4 + 8\text{HCl} \xrightarrow{\Delta} 3\text{PbCl}_2 + 4\text{H}_2\text{O} + \text{Cl}_2$

(ii) Mg + 2HCl ⟶ MgCl₂ + H₂

HCl acid is prepared in the lab. by passing hydrogen chloride gas in water using a special funnel arrangement.

D : 3 parts of conc. hydrochloric acid and 1 part of conc. nitric acid

Dense white fumes of ammonium chloride are formed when a glass rod dipped in NH₄OH soln. is brought near an open bottle containing conc. HCl

$\underset{\text{[aq.]}}{\text{NH}_3} + \underset{\text{[vapour]}}{\text{HCl}} \longrightarrow \underset{\text{[dense white fumes]}}{\text{NH}_4\text{Cl}}$

(i) Salt — Sodium Chloride [NaCl], Acid — Sulphuric acid [H₂SO₄]

(ii) $\text{NaCl} + \text{H}_2\text{SO}_4 \xrightarrow{\lt 200 ^{\circ}\text{C}} \text{NaHSO}_\text{4} + \text{HCl}$

(iii) A glass rod dipped in NH₄OH soln. is brought near the mouth of the gas jar 'Y'. If the gas jar is full with HCl gas then dense white fumes of ammonium chloride are formed.

$\underset{\text{[aq.]}}{\text{NH}_3} + \underset{\text{[vapour]}}{\text{HCl}} \longrightarrow \underset{\text{[dense white fumes]}}{\text{NH}_4\text{Cl}}$

(iv) This method indicates that HCl gas is heavier than air.

When silver nitrate soln. is added to dil. HCl, curdy white ppt. of silver chloride is formed. Whereas, there is no reaction when silver nitrate soln. is added to dil. HNO₃.

AgNO₃ + HCl ⟶ AgCl (white ppt.) + HNO₃

AgNO₃ + HNO₃ ⟶ no reaction

Dry hydrogen chloride gas (HCl) and Ammonia (NH₃) gas can be used in the study of the fountain experiment.

High solubility of HCl gas in water is the common property demonstrated by the fountain experiment.

Concentrated sulphuric acid.

Reason — Drying agent used for drying should only remove the moisture and not react with it, hence, conc. sulphuric acid is used as the drying agent.

Ca(HCO₃)₂ + 2HCl ⟶ CaCl₂ + 2H₂O + 2CO₂

(i) Below diagrams show the special funnel arrangement used for the absorption of HCl gas in water:

(ii) The reasons are :

1. Prevents or minimizes back-suction of water.
2. Provides a large surface area for the absorption of the HCl gas.

(iii) The equations are:

(A) $\text{NaCl} + \text{H}_2\text{SO}_4 \xrightarrow{\lt 200 ^{\circ}\text{C}} \text{NaHSO}_4 + \text{HCl [g]}$

(B) $\text{2NaCl} + \text{H}_2\text{SO}_4 \xrightarrow{\gt 200 ^{\circ}\text{C}} \text{Na}_2\text{SO}_4 + \text{2HCl [g}]$

Aqua regia contains one part by volume of conc. nitric acid and three parts by volume of conc. hydrochloric acid.

Quicklime [CaO] is alkaline in nature and it reacts with HCl forming the respective chloride. Hence, it can't be used as a drying agent.

$2\text{KMnO}_4 + 16\text{HCl} \xrightarrow{\Delta} 2\text{KCl} + 2\text{MnCl}_2 + 8\text{H}_2\text{O}+ 5\text{Cl}_2$

A: $\text{NaCl} + \text{H}_2\text{SO}_4 \xrightarrow{\lt 200 ^{\circ}\text{C}} \text{NaHSO}_4 + \text{HCl [g]}$

B: Fe + 2HCl (dil.) ⟶ FeCl₂ + H₂ [g]

C: NH₄OH + HCl (dil.)⟶ NH₄Cl + H₂O

D: $\text{PbO}_2 + 4\text{HCl} \xrightarrow{\Delta} \text{PbCl}_2 + 2\text{H}_2\text{O} + \text{Cl}_2$

(i) Sulphur dioxide gas

(ii) Chlorine gas
$\text{MnO}_2 + 4\text{HCl} \xrightarrow{\Delta} \text{MnCl}_2 + 2\text{H}_2\text{O} + \text{Cl}_2$

(i) A colourless gas having smell of rotten eggs is given off. The gas evolved is H₂S

CuS + 2HCl ⟶ CuCl₂ + H₂S

(ii) Curdy white precipitate of silver chloride [AgCl] is obtained, which is soluble in excess of NH₄OH.

AgNO₃ + HCl ⟶ AgCl ↓ + HNO₃

AgCl + 2NH₄OH ⟶ [Ag(NH₃)₂Cl)] + 2H₂O

Quicklime is not used to dry HCl gas because CaO is alkaline.

Action of dilute hydrochloric acid on sodium sulphide.

Na₂S + 2HCl ⟶ 2NaCl + H₂S [g.]

Na₂CO₃ + 2HCl ⟶ 2NaCl + H₂O + CO₂ ↑

Effervescence of CO₂ seen which turns lime water milky

(i) Dry HCl (Hydrogen chloride) gas is the gas Y

(ii) High solubility of HCl in water

(iii) Ammonia (NH₃) gas also demonstrates high solubility in water.

Hydrogen chloride

NH₃ [g] + HCl [g] ⟶ NH₄Cl (s-solid)

When a gas jar containing HCl gas in inverted over a jar full of ammonia gas, dense white fumes are formed. These fumes consists of particles of solid ammonium chloride suspended in air.

Dilute hydrochloric acid

AgNO₃ + HCl ⟶ AgCl (white ppt.) + HNO₃

(i) The equation for the laboratory preparation of hydrogen chloride gas :

$\text{NaCl} + \text{H}_2\text{SO}_4 \xrightarrow{\lt 200 ^{\circ}\text{C}} \text{NaHSO}_4 + \text{HCl [g]}$

(ii) Concentrated sulphuric acid is used as the drying agent used in the above preparation.
Reason — Drying agent used for drying should only remove the moisture and not react with it, hence, conc. sulphuric acid is used as the drying agent.

(iii) The precautions taken are:

1. The reaction mixture is initially heated very slowly to control the evolution of HCl gas.
2. The lower end of the thistle funnel should dip below the acid in the flask otherwise the HCl gas obtained will escape out from the thistle funnel.

(C) HCl is highly soluble in water

AgCl, a white precipitate is soluble in excess NH₄OH.

NaHCO₃ + HCl ⟶ NaCl + H₂O + CO₂ ↑

(i) White precipitate of PbCl₂ is formed which is soluble in hot water.

Pb(NO₃)₂ + 2HCl ⟶ PbCl₂ ↓ + 2HNO₃

(ii) Effervescence of CO₂ seen which turns lime water milky

CuCO₃ + 2HCl ⟶ CuCl₂ + H₂O + CO₂

(iii) Yellow particles of S precipitate.

Na₂S₂O₃ + 2HCl ⟶ 2NaCl + SO₂ + S + H₂O

(i) SO₂ is evolved when dilute hydrochloric acid reacts with sodium sulphite.

Na₂SO₃ + 2HCl ⟶ 2NaCl + H₂O + SO₂ [g]

(ii) H₂S gas is evolved when dilute hydrochloric acid reacts with Iron [II] sulphide.

FeS + 2HCl ⟶ FeCl₂ + H₂S

Potassium sulphite on reacting with hydrochloric acid, releases SO₂ gas.

Ammonium chloride.

NH₃ [g] + HCl [g] ⟶ NH₄Cl (s-solid)

H₂S gas is evolved. It has a foul smell like rotten eggs.

FeS + 2HCl [g] ⟶ FeCl₂ + H₂S

A : The products are sodium hydrogen sulphate (NaHSO₄) and HCl gas.

B : HCl gas is collected by the upward displacement of air method.

MgSO₃ + 2HCl ⟶ MgCl₂ + H₂O + SO₂ [g]

White precipitate of PbCl₂ is formed which is soluble in hot water.

Pb(NO₃)₂ + 2HCl ⟶ PbCl₂ + 2HNO₃

H₂S gas is evolved.

Na₂S + 2HCl ⟶ 2NaCl + H₂S [g.]

Dry hydrogen chloride gas can be collected by upward displacement of air.

Conc. H₂SO₄ is used for the preparation of hydrogen chloride gas in the laboratory.

As conc. H₂SO₄ is non-volatile and has a high boiling point, therefore, it displaces the volatile hydrogen chloride from the salt sodium chloride. Hence, conc. H₂SO₄ is used as a reactant in the laboratory preparation of HCl from sodium chloride

$\text{NaCl} + \text{H}_2\text{SO}_4 \xrightarrow{\lt 200 ^{\circ}\text{C}} \text{NaHSO}_4 + \text{HCl [g.]}$

(i) Direct absorption of hydrogen chloride gas in water is not feasible because it causes back suction in the delivery tube.

1. Hydrogen chloride gas is extremely soluble in water.
2. If a delivery tube through which HCl gas is passed is directly immersed in water, the rate of absorption of HCl gas is high and hence a partial vacuum is created in the tube.
3. The pressure outside being higher causes the water to be pushed up into the delivery tube and damages the apparatus. This is called back-suction.

(ii) Special funnel arrangement is used for avoiding back suction. It (a) Prevents or minimizes back-suction of water (b) Provides a large surface area for absorption of HCl gas.

1. The rim of the funnel is placed so that it just touches the trough containing water. If back-suction occurs, the water rises up the funnel and the level outside the funnel falls, creating an air gap between the rim of the funnel and the surface of water.
2. The pressure outside and inside equalize and the water which had risen in the funnel falls down again. This process continues till the water in the trough is saturated with hydrogen chloride gas resulting in the formation of hydrochloric acid.
3. Hence, hydrochloric acid is not prepared in the laboratory by passing hydrogen chloride gas directly through water, but prepared using a special funnel arrangement.

Conc. H₂SO₄

Reason — Drying agent used for drying should only remove the moisture and not react with it, hence, conc. sulphuric acid is used as the drying agent.

When sodium chloride is heated with concentrated sulphuric acid below 200°C, one of the products formed is sodium hydrogen sulphate

Effervescence of H₂ gas can be seen.

Zn + 2HCl ⟶ ZnCl₂ + H₂

Due to its high solubility, HCl gas fumes in moist air and forms a mist of droplets of HCl acid.

(i) The indicator which does not change colour on passage of HCl gas is phenolphthalein.

(ii) The acid which cannot act as an oxidizing agent is conc. HCl

(i) Balanced chemical equation for the reaction are :

$\text{NaCl} + \text{H}_2\text{SO}_4 \xrightarrow{\lt 200 ^{\circ}\text{C}} \text{NaHSO}_4 + \text{HCl [g.]}$

(ii) As conc. H₂SO₄ is non-volatile and has a high boiling point, therefore, it displaces the volatile hydrogen chloride from the salt sodium chloride. Hence, conc. H₂SO₄ is used as a reactant in the laboratory preparation of HCl from sodium chloride.
Conc. nitric acid cannot be used in place of conc. H₂SO₄ since it is volatile and may volatize out along with the hydrogen chloride.

(iii) The gas is collected by upward displacement of air method.

(iv) Quicklime (CaO) and phosphorous pentoxide [P₂O₅] are not used as drying agent, since they react with HCl forming their respective chlorides.

1. Hydrogen chloride gas is not dried using CaO.
2. Hydrogen chloride gas on heating above 500°C gives hydrogen and chlorine.The reaction is an example of thermal dissociation.
3. Iron reacts with hydrogen chloride gas forming iron (II) chloride and hydrogen. The reaction is an example of simple displacement.
4. Hydrogen chloride and water are examples of polar covalent compounds and a solution of hydrogen chloride in water contains free ions.
5. Addition of silver nitrate to hydrochloric acid, gives an insoluble precipitate of the respective chloride. This precipitate is soluble in ammonium hydroxide and insoluble in dilute nitric acid.
6. Addition of iron (II) sulphide to dilute hydrochloric acid results in liberation of hydrogen sulphide gas.
7. Aqua regia is a mixture of one part/s of concentrated nitric acid and three part/s of concentrated hydrochloric acid by volume. In aqua regia, nitric acid oxidises hydrochloric acid to chlorine.
8. Hydrochloric acid can be converted into chlorine by heating with lead (IV) oxide which acts as an oxidising agent.

(a) HCl is present in gastric juices of mammals, hence the gastric juices of mammals are acidic. It helps in digestion.

(b) Due to difference in electronegativities of H and Cl the bond in HCl is polar covalent.

Moist hydrogen gas combines with chlorine in the presence of diffused sunlight.

$\text{H}_2 + \text{Cl}_2 \xrightarrow[\text{sunlight}]{\text{Diffused}} 2\text{HCl [g]}$

1. The reaction is explosive in direct sunlight and negligible in the dark. Moisture acts as a catalyst in the above synthesis reaction.
2. Burning jet of hydrogen also burns in chlorine forming hydrogen chloride.

Below is the balanced equation for preparation of HCl gas in the laboratory from sodium chloride:

$\text{NaCl} + \text{H}_2\text{SO}_4 \xrightarrow{\lt 200 ^{\circ}\text{C}} \text{NaHSO}_4 + \text{HCl [g]}$

(i) As conc. H₂SO₄ is non-volatile and has a high boiling point, therefore, it displaces the volatile hydrogen chloride from the salt sodium chloride. Hence, conc. H₂SO₄ is used as a reactant in the laboratory preparation of HCl from sodium chloride

(ii) The fuel is not wasted, the glass apparatus does not crack and sodium sulphate does not form hard crust, which in turn may stick to the glass and will be difficult to remove. Therefore, temperature below 200°C is used.

(iii) Drying agent used for drying should only remove the moisture and not react with it, hence, conc. sulphuric acid is used as the drying agent.

Dry HCl gas is collected by the upward displacement of air because it is 1.28 times heavier than air. It is highly soluble in water, hence not collected over water.

V.D of HCl = 18.25 and air = 14.4. HCl gas is highly soluble in water.

Due to high solubility, HCl gas fumes in moist air and forms a mist of droplets of HCl acid.

The high solubility of HCl gas in water is demonstrated by the Fountain experiment.

Action on indicators:

(i) $2\text{HCl} \xrightleftharpoons{\gt 500 ^{\circ}\text{C}} \text{H}_\text{2} + \text{Cl}_\text{2}$

(ii) $\text{NH}_4\text{Cl} \xrightarrow{\Delta} \text{NH}_3 + \text{HCl}$

NH₃ [g] + HCl [g] ⟶ NH₄Cl (s-solid)

When a gas jar containing HCl gas in inverted over a jar full of ammonia gas dense white fumes are formed. These fumes consists of particles of solid ammonium chloride suspended in air.

Fe + 2HCl ⟶ FeCl₂ + H₂

Hydrochloric acid is prepared by dissolving hydrogen chloride gas in water using a special funnel arrangement.

(i) Direct absorption of HCl gas in water using a delivery tube causes back suction. Hence, it is not used for preparing HCl acid from HCl gas.

1. Hydrogen chloride gas is extremely soluble in water.
2. If a delivery tube through which HCl gas is passed is directly immersed in water, the rate of absorption of HCl gas is high and hence a partial vacuum is created in the tube.
3. The pressure outside being higher causes the water to be pushed up into the delivery tube and damages the apparatus. This is called back-suction.

(ii) Special funnel arrangement is used for converting HCl gas into HCl acid as it avoids back suction. It:
    (a) Prevents or minimizes back-suction of water.
    (b) Provides a large surface area for absorption of HCl gas.

1. The rim of the funnel is placed so that it just touches the trough containing water. If back-suction occurs, the water rises up the funnel and the level outside the funnel falls, creating an air gap between the rim of the funnel and the surface of water.
2. The pressure outside and inside equalize and the water which had risen in the funnel falls down again. This process continues till the water in the trough is saturated with hydrogen chloride gas resulting in the formation of hydrochloric acid.
3. Hence, hydrochloric acid is not prepared in the laboratory - by passing hydrogen chloride gas directly through water, but prepared using a special funnel arrangement.

A solution which boils without any change in it's composition is known as the constant boiling mixture. A solution of HCl in water forms a constant boiling mixture at 110° C.

On boiling, the mixture evolves out the vapours of both acid and water in the same proportion as in the liquid. Hence, dilute HCl cannot be concentrated beyond a certain concentration by boiling.

The ions obtained are hydronium ions (H₃O⁺) and chloride ions (Cl^-)

HCl + H₂O ⟶ H₃O⁺ + Cl^-

The presence of hydrogen ions [H⁺] in HCl imparts acidic properties to an aqueous solution of hydrochloric acid.

Water being a polar solvent, ionizes the polar covalent HCl molecule into free ions. Hence, the solution of HCl in water exhibits acidic properties and is an electrolyte whereas toluene not being a polar solvent does not dissociate HCl molecule into free ions. Hence, the solution of HCl in toluene does not exhibit acidic properties and is a non electrolyte.

1. Magnesium + Hydrochloric acid ⟶ Magnesium chloride + Hydrogen
2. Copper oxide + Hydrochloric acid ⟶ Copper chloride + Water
3. Sodium carbonate + Hydrochloric acid ⟶ Sodium chloride + Water + Carbon dioxide gas
4. Magnesium sulphite + Hydrochloric acid ⟶ Magnesium chloride + Water + sulphur dioxide gas

(i) Mg + 2HCl ⟶ MgCl₂ + H₂

(ii) NaHCO₃ + HCl ⟶ NaCl + H₂O + CO₂ [g]

(iii) Na₂SO₃ + 2HCl ⟶ 2NaCl + H₂O + SO₂ [g]

(iv) FeS + 2HCl ⟶ FeCl₂ + H₂S

(i) AgNO₃ + HCl ⟶ AgCl ↓ + HNO₃

(ii) Pb(NO₃)₂ + 2HCl ⟶ PbCl₂ ↓ + 2HNO₃

(i) Fe + 2HCl ⟶ FeCl₂ + H₂

(ii) MnO₂ + 4HCl ⟶ MnCl₂ + 2H₂O + Cl₂

Aqua regia is a mixture of 1 part of conc. HNO₃ and 3 parts of conc. HCl [by vol.]

Nitric acid is the oxidizing agent in aqua regia.

HNO₃ + 3HCl ⟶ 2H₂O + NOCl + 2[Cl] [nascent chlorine]

Aqua regia : 1 part of conc. HNO₃ and 3 parts of conc. HCl

Nitric acid present in aqua regia oxidizes HCl to chlorine.

HNO₃ + 3HCl ⟶ NOCl + 2H₂O + 2[Cl] [nascent chlorine]

This nascent chlorine reacts with noble metals like gold forming the corresponding soluble chloride.

Au + 3[Cl] ⟶ AuCl₃ [Gold (III) Chloride]

Hence, aqua regia dissolves noble metals like gold and platinum which are insoluble in all other acids.

(i) When glass rod dipped in ammonia solution (NH₄OH) is brought near vapours of hydrochloric acid, dense white fumes of ammonium chloride are formed.

$\underset{\text{[aq.]}}{\text{NH}_3} + \underset{\text{[vapour]}}{\text{HCl}} \longrightarrow \underset{\text{[dense white fumes]}}{\text{NH}_4\text{Cl}}$

(ii) On addition of silver nitrate soln. to dil. HCl (acidified with dil. HNO₃), curdy white precipitate of silver chloride is obtained.

AgNO₃ + HCl ⟶ AgCl ↓ + HNO₃

(iii) On heating a mixture of manganese dioxide and conc. HCl, Greenish yellow gas (chlorine) is evolved.

MnO₂ + 4HCl ⟶ MnCl₂ ↓ + 2H₂O + Cl₂

Industrial products manufactured from HCl acid are :

1. Dyes
2. Paint

General uses of hydrochloric acid are :

1. Hydrochloric acid dissolves the calcium phosphate present in bones. Hence, it is used in the extraction of glue from bones.
2. Hydrochloric acid dissolves the metallic oxide coating [impurities] on the surface of the metal, thereby cleans the surface. Hence, it is used in pickling of metals before galvanizing.

1. A — $4\text{HCl(conc.)} + \text{PbO}_2 \xrightarrow{\Delta} \text{PbCl}_2 + 2\text{H}_2\text{O} + \text{Cl}_2$
   B — $2\text{HCl(dil.)} + \text{Pb(NO}_3)_2 \xrightarrow{\Delta} \text{PbCl}_2 \downarrow + 2\text{HNO}_3$
2. C — NaHCO₃ + HCl ⟶ NaCl + H₂O + CO₂
   D — $\text{NaCl} + \text{H}_2\text{SO}_4 \xrightarrow{\lt 200 ^{\circ}\text{C}} \text{NaHSO}_\text{4} + \text{HCl [g]}$
3. E₁ — 2Fe + 3Cl₂ ⟶ 2FeCl₃
   E₂ — Fe + 2HCl [g] ⟶ FeCl₂ + H₂ [g]

At temperatures above 200°C, Sodium sulphate is formed. It forms a hard crust, sticks to the glass and is difficult to remove.

When a gas jar containing HCl gas in inverted over a jar full of ammonia gas, dense white fumes are formed. These fumes consists of particles of solid ammonium chloride suspended in air.

NH₃ [g] + HCl [g] ⟶ NH₄Cl (s-solid)

HCl gas present in the flask dissolves in water due to high solubility, creating a partial vacuum in the flask. The outside pressure being higher pushes the blue litmus solution up the jet tube which emerges out at the end of the tube as red fountain. Hence, the inference is that the HCl gas is highly soluble in water and acidic in nature.

Hydrochloric acid dissolves the metallic oxide coating [impurities] on the surface of the metal, thereby cleans the metallic surface before galvanizing.

Due to the high solubility of HCl gas, it fumes in moist air and forms a mist of droplets of HCl acid. On the other hand, Hydrogen Sulphide does not do so as it is slightly soluble in water.

1. An aqueous solution of HCl gas is named muriatic acid.
2. The salt obtained when rock salt reacts with conc. H₂SO₄ at temperatures below 200°C is an acid salt.
3. In the preparation of HCl acid from HCl gas, a funnel arrangement provides more surface area for absorption of the gas.
4. The ions which impart acidic properties to an aqueous solution of hydrogen chloride are hydrogen.
5. The indicator which does not change colour on passage of hydrogen chloride gas is phenolphthalein.

1. A soluble salt obtained on reaction of a metallic chloride with liquor ammonia — E : [Ag(NH₃)₂Cl]
2. A salt which is insoluble in dilute nitric acid but soluble in ammonium hydroxide. — B : AgCl
3. A salt obtained on reaction of an active metal with hydrogen chloride gas. — D : FeCl₂
4. A salt obtained when a basic gas reacts with hydrogen chloride gas. — A : NH₄Cl
5. A salt soluble in hot water but not in cold, obtained on heating an oxidising agent with conc. HCl. — C : PbCl₂

1. The substance reacted with conc. HCl and heated to prove that conc. HCl contains Cl₂. — PbO₂
2. The metal reacted with dil. HCl to prove that dil. HCl contains hydrogen. — Fe
3. The gases which are heavier than air and highly soluble in water. — HCl and NH₃
4. The acid which is not an oxidizing agent. — Conc. HCl
5. The acid which is not a monobasic acid. — Sulphurous acid

1. Active metal + hydrochloric acid (dil) ⟶ salt + hydrogen
2. Metallic oxide, hydroxide + hydrochloric acid (dil) ⟶ salt + water
3. Metallic carbonate / bicarbonate + hydrochloric acid (dil) ⟶ salt + water + carbon dioxide
4. Metallic bisulphite + hydrochloric acid (dil) ⟶ salt + water + sulphur dioxide
5. Metallic sulphide + hydrochloric acid (dil) ⟶ salt + hydrogen sulphide

(i) Ammonia (NH₃)

CuSO₄ + 2NH₄OH [ammonia soln.]⟶ (NH₄)₂SO₄ + Cu(OH)₂
Cu(OH)₂ + (NH₄)₂SO₄ + 2NH₄OH ⟶ [Cu(NH₃)₄]SO₄ + 4H₂O

(ii) Ammonia (NH₃)

4NH₃ + 3O₂ ⟶ 2N₂ + 6H₂O

N₂ + 3H₂ ⇌ 2NH₃ + Δ

Liquefaction is a technique used for separating ammonia from the uncombined N₂ and H₂ reactant mixture during Haber's process.

8NH₃ [excess] + 3Cl₂ ⟶ 6NH₄Cl + N₂

Hydroxyl ion [OH^-1]

NH₃ + H₂O ⟶ NH₄OH

NH₄OH ⇌ NH₄⁺ + OH^-1

(i) 2NH₄Cl + Ca(OH)₂ ⟶ CaCl₂ + 2H₂O + 2NH₃

(ii) AlN + 3H₂O ⟶ Al(OH)₃ + NH₃ [g]

Ammonia.

Reason — The aq. soln. of ammonia [NH₄OH] is alkaline in nature due to the presence of hydroxyl ions [OH^-]

NH₃ + H₂O ⟶ NH₄OH

NH₄OH ⇌ NH₄⁺ + OH^-1

Ammonium sulphate.

Reason — Ammonia gas is produced when ammonium sulphate reacts with NaOH. Ammonia gas forms dense white fumes with HCl.

(NH₄)₂SO₄ + 2NaOH ⟶ Na₂SO₄ + 2H₂O + 2NH₃[g]

Lead hydroxide is precipitated, chalky white in colour and is insoluble in excess of ammonium hydroxide soln.

Pb(NO₃)₂ + 2NH₄OH ⟶ 2NH₄NO₃ + Pb(OH)₂ ↓

Quicklime (CaO) is used for drying ammonia.

Ammonia gas is a strong reducing agent. Below equation shows reduction of Black Copper [II] oxide to brown Copper by Ammonia.

2NH₃ + 3CuO ⟶ 3Cu + 3H₂O + N₂ [g]

Nitrogen reacts with hydrogen under specific conditions liberating ammonia

N₂ + 3H₂ ⇌ 2NH₃ + Δ

Conditions :

Temperature : 450-500°C [Optimum temperature]

Pressure : 200 to 900 atmospheres [Optimum pressure]

Catalyst : Finely divided iron [Fe]

Promotor : Molybdenum [Mo]

[Catalyst - iron [III] oxide [Fe₂O₃] may also be used containing promoters about 1% K₂O and 3% Al₂O₃]

Precautions : Impurities like CO, CO₂, and H₂S may poison the catalyst hence, the nitrogen-hydrogen reactant mixture must be free from them.

2NH₃ + H₂SO₄ (dil.) ⟶ (NH₄)₂SO₄

(i) 2NH₃ + 3CuO ⟶ 3Cu + 3H₂O + N₂ [g]

(ii) 8NH₃ [excess] + 3Cl₂ ⟶ 6NH₄Cl + N₂

If dry ammonia gas is present in the round bottom flask, the red litmus soln. kept in the trough below emerges out as a blue fountain since ammonia gas is basic in nature. On the other hand, if dry HCl was used in the flask instead of ammonia, the blue litmus soln. kept in the trough below, emerges out as a red fountain since HCl is acidic in nature.

Magnesium nitride
Mg₃N₂ + 6H₂O ⟶ 3Mg(OH)₂ + 2NH₃ [g]

NH₃ (Ammonia)
NH₃ + HCl ⟶ NH₄Cl

AlN + 3H₂O ⟶ Al(OH)₃ + NH₃ [g]

Completed table is shown below:

Ammonia (NH₃)
4NH₃ + 3O₂ ⟶ 2N₂ + 6H₂O

Mg₃N₂ + 6H₂O ⟶ 3Mg(OH)₂ + 2NH₃ [g]

Substances 'Q' are Ammonium salts like ammonium chloride, ammonium sulphate. They are white solids, soluble in water and liberate pungent smelling ammonia gas on heating with alkalis like NaOH.

(a) Ammonia

(b) Alkaline

(c) Ammonium

(d) Hydroxyl

(e) Dirty green

Green or greenish yellow flame is seen when ammonia gas is burnt in an atmosphere of oxygen. The reaction of combustion of Ammonia is:

4NH₃ + 3O₂ ⟶ 2N₂ + 6H₂O

NH₄Cl + NaOH ⟶ NaCl + H₂O + NH₃

(i) Haber's Process

(ii) Nitrogen (N₂) and hydrogen (H₂) in the ratio 1:3 by volume

(iii) Finely divided iron (Fe)

(iv) N₂ + 3H₂ ⇌ 2NH₃ + Δ

Ammonia gas is a strong reducing agent. Below equation shows reduction of Black Copper [II] oxide to brown Copper by Ammonia.

2NH₃ + 3CuO ⟶ 3Cu + 3H₂O + N₂ [g]

Hydrogen chloride gas (HCl) and Ammonia (NH₃) are the two gases which can be used to study the fountain experiment.

The common property demonstrated by the fountain experiment is the solubility of gases.

Ammonium hydroxide if first added in small quantity and then in excess to a solution of copper sulphate, a pale blue ppt. of copper hydroxide is formed which dissolves in excess of ammonium hydroxide forming a soluble complex salt.

CuSO₄ + 2NH₄OH ⟶ (NH₄)₂SO₄ + Cu(OH)₂ ↓

Cu(OH)₂ + (NH₄)₂SO₄ + 2NH₄OH ⟶ [Cu(NH₃)₄]SO₄ + 4H₂O

(i) Ammonia (NH₃) gas is collected.

(ii) 2NH₄Cl + Ca(OH)₂ ⟶ CaCl₂ + 2H₂O + 2NH₃

(iii) Downward displacement of air is the method used to collect the gas.

(iv) Quicklime (CaO) is the drying agent.

(v) A glass rod dipped in conc. HCl acid is brought near the mouth of the jar. If the jar is full-dense white fumes of ammonium chloride are formed.

8NH₃ [excess] + 3Cl₂ ⟶ 6NH₄Cl + N₂

When Al metal is burnt in a jar of nitrogen gas, its nitride i.e., AlN is formed. When warm water is added to AlN, it undergoes hydrolysis and is decomposed by warm water to give pungent smelling ammonia gas. The corresponding insoluble metal hydroxide (i.e., Al(OH)₃) is precipitated out. The reactions are shown below:

2Al + N₂ ⟶ 2AlN

AlN + 3H₂O ⟶ Al(OH)₃ + NH₃ [g]

Nitrogen

8NH₃ [excess] + 3Cl₂ ⟶ 6NH₄Cl + N₂

Magnesium nitride reacts with warm water to liberate ammonia along with magnesium hydroxide.

Below is the equation for the catalytic oxidation of Ammonia:

$4\text{NH}_3 + 5\text{O}_2 \xrightarrow[800 ^{\circ} \text{C}]{\text{Pt.}} \text{4NO} + \text{6H}_\text{2}\text{O} + \Delta$

(i) As ammonium nitrate is explosive in nature and may itself decompose forming nitrous oxide and water vapour, hence it is not used in the preparation of ammonia.

(ii) Quick lime [CaO] is used as a drying agent during the process.

(iii) Downward displacement of air is the method used for the collection of the ammonia gas. As ammonia gas is highly soluble in water, therefore, it is not collected over water.

Colourless ammonia gas reacts with greenish yellow excess chlorine giving a yellow explosive liquid (Nitrogen trichloride).

NH₃ + 3Cl₂ [excess] ⟶ 3HCl + NCl₃

Ammonium nitrite

NH₄NO₂ ⟶ 2H₂O + N₂

White gelatinous precipitate of zinc hydroxide is formed which dissolves in excess of NH₄OH solution.

Zn(NO₃)₂ + 2NH₄OH ⟶ 2NH₄NO₃ + Zn(OH)₂ ↓

2NH₃ + 3CuO ⟶ 3Cu + 3H₂O + N₂ [g]

(i) Temperature : 450 to 500°C [Optimum temperature]

(ii) Catalyst : Finely divided iron (Fe)

(iii) Equation : N₂ + 3H₂ ⇌ 2NH₃ + Δ

Ammonia gas.

NH₃ + HCl ⟶ NH₄Cl

Ammonia gas is collected by downward displacement of air.

Mg₃N₂ + 6H₂O ⟶ 3Mg(OH)₂ + 2NH₃ [g]

(i) On reaction with ammonium hydroxide, Iron [II] sulphate forms a dirty green ppt whereas iron [III] sulphate forms a reddish brown ppt. Hence, the two can be distinguished easily.

FeSO₄ + 2NH₄OH ⟶ (NH₄)₂SO₄ + Fe(OH)₂ ↓

Fe₂(SO₄)₃ + 6NH₄OH ⟶ 3(NH₄)₂SO₄ + 2Fe(OH)₃ ↓

(ii) A lead salt gives a chalky white ppt. on reaction with ammonium hydroxide that is insoluble in excess of ammonium hydroxide. For example:

Pb(NO₃)₂ + 2NH₄OH ⟶ 2NH₄NO₃ + Pb(OH)₂ ↓

On the other hand, zinc salt forms a white gelatinous ppt. which dissolves when excess of ammonium hydroxide is added. Hence, the two can be distinguished.

ZnSO₄ + 2NH₄OH ⟶ (NH₄)₂SO₄ + Zn(OH)₂ ↓

[Zn(OH)₂] + (NH₄)₂SO₄ + 2NH₄OH ⟶ [Zn(NH₃)₄]SO₄ + 4H₂O

Pungent smelling gas (ammonia) is given out.

2NH₄Cl + Ca(OH)₂ ⟶ CaCl₂ + 2H₂O + 2NH₃

Hydroxyl ion [OH^-1] and ammonium ions [NH₄⁺] are formed.

NH₃ + H₂O ⟶ NH₄OH

NH₄OH ⇌ NH₄⁺ + OH^-1

The alkaline NH₄OH, due to the presence of hydroxyl ions [OH^-] turns red litmus blue and phenolphthalein soln. pink

Conditions for catalytic oxidation of ammonia to nitric oxide are :

1. Temperature : 800°C
2. Catalyst : Platinum (Pt)

$4\text{NH}_3 + 5\text{O}_2 \xrightarrow[800 ^{\circ} \text{C}]{\text{Pt.}} \text{4NO} + \text{6H}_\text{2}\text{O} + \Delta$

Ammonia

Reducing nature of ammonia:
2NH₃ + 3CuO ⟶ 3Cu + 3H₂O + N₂ [g]

4NH₃ + 3O₂ ⟶ 2N₂ + 6H₂O

Ammonia burns in the atmosphere of excess oxygen with a green or greenish yellow flame, forming nitrogen and water vapour.

X (metal) has a valency 2, hence,

X - 2e^- ⟶ X²⁺

Y (non metal) has a valency 3. hence,

Y + 3e^- ⟶ X^3-

Therefore, the two combine as :

3X + 2Y ⟶ X₃Y₂

(i) 2NH₄Cl + Ca(OH)₂ ⟶ CaCl₂ + 2H₂O + 2NH₃

(ii) NH₃ + 3Cl₂ [excess] ⟶ 3HCl + NCl₃

(iii) 2NH₃ + H₂SO₄ (dil.) ⟶ (NH₄)₂SO₄

(i) AlN + 3H₂O ⟶ Al(OH)₃ + NH₃ [g]

(ii) 8NH₃ [excess] + 3Cl₂ ⟶ 6NH₄Cl + N₂

(iii) 2NH₃ + 3CuO ⟶ 3Cu + 3H₂O + N₂ [g]

(i) Ammonia gas

2NH₄Cl + Ca(OH)₂ ⟶ CaCl₂ + 2H₂O + 2NH₃

(ii) Nitrogen gas

NH₄Cl + NaNO₂ ⟶ NaCl + NH₄NO₂

NH₄NO₂ ⟶ 2H₂O + N₂

(i) 2NH₃ + 3CuO ⟶ 3Cu + 3H₂O + N₂ [g]

(ii) 2NH₄Cl + Ca(OH)₂ ⟶ CaCl₂ + 2H₂O + 2NH₃

Ammonia burns in the atmosphere of excess oxygen with a green or greenish yellow flame, forming nitrogen and water.

4NH₃ + 3O₂ ⟶ 2N₂ + 6H₂O

C : Magnesium nitride (Mg₃N₂) and water (H₂O).

D : Quick lime [CaO]

E : Downward displacement of air.

(i) $4\text{NH}_3 + 5\text{O}_2 \xrightarrow[800 ^{\circ} \text{C}]{\text{Pt.}} \text{4NO} + \text{6H}_\text{2}\text{O} + \Delta$

(ii) NH₃ + HNO₃ ⟶ NH₄NO₃

2NH₄Cl + Ca(OH)₂ ⟶ CaCl₂ + 2H₂O + 2NH₃ [g]

(i) As sulphuric acid reacts chemically with ammonia to form ammonium sulphate hence, it is not used as a drying agent for drying ammonia.

2NH₃ + H₂SO₄ ⟶ (NH₄)₂SO₄

(ii) As ammonia gas is highly soluble in water hence, it is not collected over water.

Ammonia reacts with excess chlorine to form nitrogen trichloride

Black copper [II] oxide is reduced to brown copper.

2NH₃ + 3CuO ⟶ 3Cu + 3H₂O + N₂ [g]

Platinum

$4\text{NH}_3 + 5\text{O}_2 \xrightarrow[800 ^{\circ} \text{C}]{\text{Pt.}} \text{4NO} + \text{6H}_\text{2}\text{O} + \Delta$

Nitrogen

2NH₃ + 3CuO ⟶ 3Cu + 3H₂O + N₂ [g]

A : Mg₃N₂ + 6H₂O ⟶ 3Mg(OH)₂ + 2NH₃ [g]

B : NH₃ + HCl ⟶ NH₄Cl

C: NH₄Cl + NaOH ⟶ NaCl + H₂O + NH₃ [g]

Name of the compound — Ammonia

Name of the process — Haber's process

Catalytic equation [with the catalyst] — $\text{N}_2 + 3\text{H}_2 \xrightleftharpoons[\text{450-500 ^{\circ} C}]{\text{Fe}} \text{2NH}_3 + \Delta$

(i) 8NH₃ [excess] + 3Cl₂ ⟶ 6NH₄Cl + N₂

(ii) Pb(NO₃)₂ + 2NH₄OH ⟶ 2NH₄NO₃ + Pb(OH)₂ ↓

Ammonium salts react with alkalis on heating to liberate Ammonia gas. Hence, we can use NaOH as reagent to distinguish between Ammonium Sulphate crystals and Sodium Sulphate crystals. Below is the reaction for the same:

(NH₄)₂SO₄ + 2NaOH ⟶ Na₂SO₄ + 2H₂O + 2NH₃ [g]

Platinum

$4\text{NH}_3 + 5\text{O}_2 \xrightarrow[800 ^{\circ} \text{C}]{\text{Pt.}} \text{4NO} + \text{6H}_\text{2}\text{O} + \Delta$

Ammonia is produced when nitrogenous matter decays in the absence of air. Putrefying bacteria on organic matter in the soil and the ammonifying bacteria on nitrogenous organic matter bring about decomposition and produces ammonia.

Urine is the liquid source of ammonia.

Ammonium Chloride + Calcium Hydroxide ⟶ Calcium Chloride + Water + Ammonia

(NH₄)₂SO₄ + 2NaOH ⟶ Na₂SO₄ + 2H₂O + 2NH₃

(NH₄)₂SO₄ + Ca(OH)₂ ⟶ CaSO₄ + 2H₂O + 2NH₃

As ammonium nitrate is explosive in nature and may itself decompose forming nitrous oxide and water vapour, hence a mixture of ammonium nitrate and NaOH are not used in the lab. preparation of ammonia gas.

The ammonia gas is passed through a drying tower before collection of the gas. The drying agent used is quicklime [CaO]
Reason — Quicklime being basic in nature does not react with basic ammonia gas. The drying agents not used are conc. H₂SO₄, P₂O₅ and fused CaCl₂ since they react chemically with ammonia.

Dry ammonia gas is collected by the downward displacement of air as it is lighter than air. Ammonia gas is not collected over water since it is highly soluble in water.

When metals like Mg, Ca and Al are burnt in Nitrogen gas, they form their respective metal nitrides i.e., Magnesium Nitride, Calcium Nitride and Aluminium Nitride as the products.
When warm water is added to these metal nitrides, they undergo hydrolysis decomposing to give pungent ammonia gas. The chemical reactions for each metal is given below:

1. Mg to NH₃
   3Mg + N₂ ⟶ Mg₃N₂
   Mg₃N₂ + 6H₂O ⟶ 3Mg(OH)₂ + 2NH₃ [g]

2. Ca to NH₃
   3Ca + N₂ ⟶ Ca₃N₂
   Ca₃N₂ + 6H₂O ⟶ 3Ca(OH)₂ + 2NH₃ [g]

3. Al to NH₃
   2Al + N₂ ⟶ 2AlN
   AlN + 3H₂O ⟶ Al(OH)₃ + NH₃ [g]

Nitrogen reacts with hydrogen under specific conditions liberating ammonia.

Below balanced equation represents this reaction :

N₂ + 3H₂ ⇌ 2NH₃ + Δ

Conditions :

Temperature : 450-500°C [Optimum temperature]

Pressure : 200-900 atmospheres [Optimum pressure]

Catalyst : Finely divided iron [Fe]

Promotor : Molybdenum [Mo]

[Catalyst - iron [III] oxide [Fe₂O₃] may also be used containing promoters about 1% K₂O and 3% Al₂O₃]

Ammonia is separated and recovered from the uncombined nitrogen and hydrogen by technique based on the physical properties of ammonia.

1. Ammonia is easily liquefiable — It condenses to a liquid at around -33°C at 8 atmosphere pressure. N₂ and H₂ are difficult to liquify. They condense at temperatures around -196°C and -253°C respectively and at comparatively higher pressures.
2. Ammonia is highly soluble in water — 1 vol of water dissolves about 702 vols. of ammonia at 20°C and 1 atmosphere pressure. N₂ and H₂ are almost insoluble in water.

Ammonia is lighter than air. The vapour density of ammonia is 8.5 while that of air is 14.4.

The two gases lighter than ammonia are :

1. Hydrogen
2. Helium.

Other gases highly soluble in water are :

1. Hydrogen chloride
2. Sulphur dioxide

Fountain Experiment demonstrates the high solubility of ammonia in water.

Apparatus:

• Dry round bottom flask filled with ammonia gas.
• Mouth of the flask has a rubber stopper with two holes for (a) Jet tube, (b) Dropper containing water.
• Red litmus solution is placed in the trough below.

Procedure:

• The dropper containing water is squeezed and water enters the flask.

Observation:

• Ammonia gas present in the flask dissolves in water due to it's high solubility, thereby creating a partial vacuum in the flask.
• The outside pressure being higher pushes the red litmus solution up the jet tube which emerges out at the end of the tube as a blue fountain.

Inference:

• Ammonia gas is highly soluble in water and basic in nature.

4NH₃ + 3O₂ ⟶ 2N₂ + 6H₂O

Ammonia burns in the atmosphere of excess oxygen with a green or greenish yellow flame, forming nitrogen and water vapour.

Reaction:

$4\text{NH}_3 + 5\text{O}_2 \xrightarrow[800 ^{\circ} \text{C}]{\text{Pt}.} 4\text{NO} + 6\text{H}_2\text{O} + \Delta$

Conditions:
Temperature : 800°C
Catalyst : Platinum (Pt)

Below diagram shows the setup for catalytic oxidation of ammonia:

Following are the observations with reasons:

• The colourless nitric oxide formed, on oxidation gives reddish brown nitrogen dioxide
  2NO + O₂ ⟶ 2NO₂
• The platinum [catalyst] continues to glow even after the heating is discontinued since the catalytic oxidation of ammonia is an exothermic reaction.

Ostwald’s process for preparing nitric acid.

An aqueous solution of NH₃ is called Liquor Ammonia.

Preparation:
By dissolving ammonia gas in water using a funnel arrangement.

Reasons:

• Prevents back suction of water and
• Provides a large surface area for absorption of ammonia gas.

NH₃ + H₂O ⟶ NH₄OH [liquor ammonia]

(i) The alkaline behaviour of aq. soln. of ammonia [NH₄OH] is due to the presence of hydroxyl ions [OH^-]. Hence, the aq. soln. of ammonia turns red litmus blue.

(ii) The aq. soln. of ammonia [NH₄OH] is alkaline. It undergoes partial dissociation in aq. soln. to give hydroxyl ions [OH^-] in low concentration. Therefore, it is a weak base and a weak electrolyte.

Below are the two methods of preparing NH₄Cl using hydrochloric acid:

1. NH₃ + HCl (dil.) ⟶ NH₄Cl
2. NH₄OH + HCl (dil.) ⟶ NH₄Cl + H₂O

1. Ammonia to ammonium salt:
   (a) NH₃ + HNO₃ (dil.) ⟶ NH₄NO₃
   (b) 2NH₃ + H₂SO₄ (dil.) ⟶ (NH₄)₂SO₄
2. Ammonium hydroxide to ammonium salt:
   (a) NH₄OH + HNO₃ (dil.) ⟶ NH₄NO₃ + H₂O
   (b) 2NH₄OH + H₂SO₄ (dil.) ⟶ (NH₄)₂SO₄ + 2H₂O

As liquor ammonia is an alkali due to the presence of hydroxyl ions.

NH₃ + H₂O ⟶ NH₄OH

NH₄OH ⇌ NH₄⁺ + OH^-1

When an acid (nitric acid) and an alkali or base combine completely to form salt and water then the reaction is called a neutralization reaction.

NH₄OH + HNO₃ ⟶ NH₄NO₃ + H₂O

Hence, the reaction between liquor ammonia and nitric acid is a neutralization reaction.

The precipitate which are hydroxides of metallic radicals vary in :

• colour
• solubility in excess of NH₄OH

Hence, ammonium hydroxide is used in qualitative analysis for identifying positive radicals or cations.

1. A soln. of Fe²⁺ ions gives dirty green ppt. which is insoluble in excess of NH₄OH.
   FeSO₄ + 2NH₄OH ⟶ (NH₄)₂SO₄+ Fe(OH)₂ ↓

2. A soln. of Fe³⁺ ions gives reddish brown ppt. which is insoluble in excess of NH₄OH
   FeCl₃ + 3NH₄OH ⟶ 3NH₄Cl + Fe(OH)₃ ↓

3. A soln. of Pb²⁺ ions gives chalky white ppt. which is insoluble in excess of NH₄OH
   Pb(NO₃)₂ + 2NH₄OH ⟶ 2NH₄NO₃ + Pb(OH)₂ ↓

4. A soln. of Zn²⁺ ions gives gelatinous white ppt. which is soluble in excess of NH₄OH
   ZnSO₄ + 2NH₄OH ⟶ (NH₄)₂SO₄ + Zn(OH)₂ ↓
   Zn(OH)₂ + (NH₄)₂SO₄ + 2NH₄OH ⟶ [Zn(NH₃)₄]SO₄ + 4H₂O

5. A soln. of Cu²⁺ ions gives pale blue ppt. which is soluble in excess of NH₄OH
   CuSO₄ + 2NH₄OH ⟶ (NH₄)₂SO₄ + Cu(OH)₂ ↓
   Cu(OH)₂ + (NH₄)₂SO₄ + 2NH₄OH ⟶ [Cu(NH₃)₄]SO₄ + 4H₂O

Ammonium hydroxide if first added in small quantity and then in excess to a solution of copper sulphate, a pale blue ppt. of copper hydroxide is formed which dissolves in excess of ammonium hydroxide forming a soluble complex salt.

CuSO₄ + 2NH₄OH ⟶ (NH₄)₂SO₄ + Cu(OH)₂ ↓

Cu(OH)₂ + (NH₄)₂SO₄ + 2NH₄OH ⟶ [Cu(NH₃)₄]SO₄ + 4H₂O

The pale blue ppt. of copper hydroxide is soluble in excess of a NH₄OH, since a soluble complex [Cu(NH₃)₄]SO₄ [tetrammine copper [II] sulphate] is formed.

Zinc salt forms a white gelatinous ppt. of Zn(OH)₂, which dissolves when excess of ammonium hydroxide is added due to the formation of [Zn(NH₃)₄]SO₄

ZnSO₄ + 2NH₄OH ⟶ (NH₄)₂SO₄ + Zn(OH)₂ ↓

Zn(OH)₂ + (NH₄)₂SO₄ + 2NH₄OH ⟶ [Zn(NH₃)₄]SO₄ + 4H₂O

(i) 2NH₃ + 3CuO ⟶ 3Cu + 3H₂O + N₂ [g]

(ii) 2NH₃ + 3PbO ⟶ 3Pb + 3H₂O + N₂ [g]

(iii) (a) 8NH₃ [excess] + 3Cl₂ ⟶ 6NH₄Cl + N₂

(iii) (b) NH₃ + 3Cl₂ [excess] ⟶ 3HCl + NCl₃

Five tests for ammonia

Ammonia turns —

1. Red litmus solution to blue.
2. Methyl orange solution to yellow.
3. Colourless phenolphthalein solution to pink.
4. Pale blue precipitate of CuSO₄ into deep blue coloured solution
5. Colourless Nesseler's reagent to pale brown.

(i) Hydrogen

(ii) Nitric acid

(iii) Ammonium nitrate (NH₄NO₃)

(iv) Ammonium sulphate [(NH₄)₂SO₄]

(a) Ammonium Carbonate, [(NH₄)₂CO₃] is used as smelling salt. It dissociates evolving pungent ammonia gas and hence is used as a constituent for smelling salts.

(NH₄)₂CO₃ ⟶ 2NH₃ + CO₂ + H₂O

The pungent smelling NH₃ gas revives a fainted person.

(b) Ammonium Chloride (NH₄Cl) is used in dry cells. It oxidizes Zn to Zn²⁺ ions. The electrons so produced constitute the electric current.

(i) An aqueous solution of NH₃ is called Liquor Ammonia. It is used as a cleaning agent to remove oil or grease stains from clothes, cleaning window panes, porcelain articles, etc as it emulsifies or dissolves fats, grease etc.

(ii) Liquefied NH₃ is used as refrigerant because (a) it is highly volatile and has a high latent heat of evaporation. (b) it easily liquefies under pressure, at low temperature.

Chlorofluorocarbons (CFC’s) are compounds of carbon with chlorine or fluorine. CFC are chemicals which like liquefied ammonia gas are used in refrigeration gas.

They enter the atmosphere due to their use as coolants in refrigeration and A/C plants and aerosol sprays and cause ozone depletion in the atmosphere.

Chlorofluorocarbons (CFC’s) are decomposed by U.V. rays to give free Cl radicals.

CFCl₃ ⟶ Cl atom

This Cl atom breaks the ozone layer causing ozone depletion.

Cl + O₃ [ozone] ⟶ ClO + O₂

Hence, ozone layer located in the stratosphere above the earth, which prevents harmful U.V. rays from reaching the earth, gets depleted.

Suitable alternatives to chlorofluorocarbons which are non-ozone depleting include Hydrochlorofluorocarbons [HCFC]

1. Dry ammonia and dry oxygen gas are ignited together. — C : N₂
2. Ammonia is passed over heated litharge. — C : N₂
3. A greenish yellow gas reacts with excess ammonia. — C : N₂
4. (a) Dry NH₃ and O₂ are passed over heated Pt. — B : NO
   (b) The gaseous product obtained is further oxidized. — A : NO₂
5. Ammonium nitrite undergoes thermal decomposition. — C : N₂

1. Pink colour
2. Deep blue coloured solution
3. Green or greenish yellow flame
4. White gelatinous precipitate soluble in excess of ammonium hydroxide
5. The colourless nitric oxide formed, on oxidation gives reddish brown nitrogen dioxide.

• A : FeSO₄ + 2NH₄OH ⟶ (NH₄)₂SO₄ + Fe(OH)₂ ↓

• B : (NH₄)₂SO₄ + 2NaOH ⟶ Na₂SO₄ + 2H₂O + 2NH₃ [g]

• C : NH₃ + HCl ⟶ NH₄Cl

• D : 8NH₃ [excess] + 3Cl₂ ⟶ 6NH₄Cl + N₂

• E : 4NH₃ + 3O₂ ⟶ 2N₂ + 6H₂O

The alkaline behaviour of NH₄OH is due to the presence of hydroxyl ions [OH^-]. It acts as a weak base, since it undergoes partial dissociation in aq. soln. to give hydroxyl ions [OH^-] in low concentration.

As ammonium nitrate is explosive in nature and may itself decompose forming nitrous oxide and water vapour, hence a mixture of ammonium nitrate and slaked lime are not used in the lab. preparation of ammonia gas.

The catalyst only accelerates the reaction and does not form a part of the reaction and hence it does not affect the percentage yield of ammonia in Haber’s process.

Ammonia gas or aqueous solution of Ammonia is basic in nature. Hence, reaction between dilute acids and ammonia is a neutralization reaction between an acid and a base forming salt and water as the end products. Therefore, ammonium salts are formed when ammonia reacts with dilute acids in the gaseous or aq. medium.

A lead salt gives a chalky white ppt. on reaction with ammonium hydroxide which does not dissolve when excess of ammonium hydroxide is added.

Pb(NO₃)₂ + 2NH₄OH ⟶ 2NH₄NO₃ + Pb(OH)₂ ↓

On the other hand, zinc salt forms a white gelatinous ppt. which dissolves when excess of ammonium hydroxide is added. Hence, the two can be easily distinguished using an aqueous solution of ammonia.

1. The salt solution which does not give an insoluble precipitate on addition of ammonium hydroxide in small amount is NaNO₃
2. The alkaline behaviour of liquor ammonia is due to the presence of hydroxyl ions.
3. Ammonia in the liquefied form is basic
4. Ammonia reduces chlorine to hydrogen chloride
5. The chemical not responsible for ozone depletion is ammonia

1. NH₄Cl ⇌ NH₃ + HCl
   This reaction is an example of thermal dissociation.
2. NH₄NO₃ ⟶ N₂O + 2H₂O
   This reaction is an example of thermal decomposition.

Potassium nitrate [nitre] or sodium nitrate [chile salt petre] and conc. sulphuric acid are the products required for the laboratory preparation of nitric acid.

Pure nitric acid [HNO₃] is colourless and unstable and decomposes slightly even at ordinary temperatures and in the presence of sunlight.

The decomposition results in formation of reddish brown nitrogen dioxide [NO₂] which remains dissolved in the acid thus imparting a slight yellowish brown colour.

Cu + 4HNO₃ (conc.) ⟶ Cu(NO₃)₂ + 2H₂O + 2NO₂.

Platinum gauze

Concentrated nitric acid (HNO₃) gives nitrogen dioxide with copper

Cu + 4HNO₃ (conc.) ⟶ Cu(NO₃)₂ + 2H₂O + 2NO₂.

All glass apparatus is used in the laboratory preparation of nitric acid since the vapours of nitric acid being highly corrosive attack rubber, cork, etc.

Below is the equation for the lab. preparation of HNO₃ from potassium nitrate and conc. H₂SO₄:

$\underset{[conc.]}{\text{KNO}_3} + \underset{[conc.]}{\text{H}_2\text{SO}_4} \xrightarrow{\lt 200 ^{\circ}\text{C}} \underset{[acid \space salt]}{\text{KHSO}_4} + \text{HNO}_3$

Neutralization reaction.

KOH + HNO₃ [dil.] ⟶ KNO₃ + H₂O

Hot and conc. nitric acid.

S + 6HNO₃ ⟶ H₂SO₄ + 2H₂O + 6NO₂

A reddish-brown gas is obtained when X, Y and Z are separately warmed with concentrated sulphuric acid and copper turnings added to the mixture. The common anion is the [NO^–] Nitrate ion.

Cu + 4HNO₃ ⟶ Cu(NO₃)₂ + 2H₂O + 2NO₂

S + 6HNO₃ ⟶ H₂SO₄ + 2H₂O + 6NO₂

Dilute nitric acid is generally considered a typical acid except for it's reaction with metals since it generally does not liberate hydrogen on reaction with metals. Nitric oxide on decomposition forms nascent oxygen which oxidizes the hydrogen to water.

3Cu + 8HNO₃ [cold and dil.] ⟶ 3Cu(NO₃)₂ + 4H₂O + 2NO

Pure nitric acid [HNO₃] is colourless and unstable and decomposes slightly even at ordinary temperatures and in the presence of sunlight. The decomposition results in formation of reddish brown nitrogen dioxide [NO₂] which remains dissolved in the acid thus imparting a slight yellowish brown colour.

Copper nitrate

When nitric acid is kept in a reagent bottle for a long time, it turns dark yellowish brown in colour.

This is because nitric acid is unstable hence, it decomposes slightly even at ordinary temperatures and in the presence of sunlight. When it is kept for a long time, the decomposition is complete resulting in the formation of reddish brown nitrogen dioxide [NO₂] that dissolves in the acid giving it a darker yellowish brown colour.

All glass apparatus is used in the laboratory preparation of nitric acid since the vapours of nitric acid being highly corrosive attack rubber, cork, etc.

(A) Liquid → Conc. sulphuric acid

(B) Solid → Potassium Nitrate (Nitre)

4HNO₃ ⟶ 4NO₂ + 2H₂O + O₂

Cu + 4HNO₃ [conc.] ⟶ Cu(NO₃)₂ + 2H₂O + 2NO₂

HNO₃ (Nitric acid)

When nitric acid undergoes decomposition it yields nascent oxygen. This nascent oxygen oxidises copper allowing it to react with nitric acid. Hence, the oxidising nature of nitric acid helps it to react with copper.

3Cu + 8HNO₃ [cold & dil.] ⟶ 3Cu(NO₃)₂ + 4H₂O + 2NO

Nitric oxide (NO)

3Cu + 8HNO₃ [cold and dil.] ⟶ 3Cu(NO₃)₂ + 4H₂O + 2NO

(A) Sulphur

S + 6HNO₃ ⟶ H₂SO₄ + 2H₂O + 6NO₂

Copper reacts with concentrated nitric acid to produce nitrogen dioxide.

Cu + 4HNO₃ [conc.] ⟶ Cu(NO₃)₂ + 2H₂O + 2NO₂

(A) Nitroso Iron [II] sulphate

X is Copper nitrate [Cu(NO₃)₂]

Y is Nitrogen dioxide gas (NO₂)

$2\text{Cu(NO}_3)_2 \xrightarrow{\Delta} 2\text{CuO} + \text{O}_2 + 4\text{NO}_2$

Platinum

Dense brown fumes of nitrogen dioxide can be seen.

$\text{Cu} + 4\text{HNO}_3 \xrightarrow{\Delta} \text{Cu(NO}_3)_2 + 2\text{H}_2\text{O} + 2\text{NO}_2$

NO₃^1–

6FeSO₄ + 3H₂SO₄ [conc.] + 2HNO₃ [dil.] ⟶ 3Fe₂(SO₄)₃ + 4H₂O + 2NO

FeSO₄ + NO ⟶ FeSO₄.NO [Nitroso Iron [II] Sulphate] - brown compound

(i) All glass apparatus is used since, the vapours of nitric acid being highly corrosive attack rubber, cork. etc.

(ii) Temperature above 200°C may cause:

1. Damage to the glass apparatus.
2. Further decomposition of nitric acid
3. Formation of hard residual crust of the corresponding sulphate [Na₂SO₄ or K₂SO₄] which being a poor conductor of heat, sticks to the glass and cannot be easily removed from the apparatus.

Fe(OH)₃ + 3HNO₃ [dil.] ⟶ Fe(NO₃)₃ + 3H₂O

Nitrogen dioxide.

Cu + 4HNO₃ ⟶ Cu(NO₃)₂ + 2H₂O + 2NO₂

Reddish brown nitrogen dioxide gas is evolved on heating zinc nitrate crystals.

$2\text{Zn(NO}_3)_2 \xrightarrow{\Delta} 2\text{ZnO} + \text{O}_2 + 4\text{NO}_2$

Magnesium reacts with cold, very dilute nitric acid to liberate hydrogen gas.

Iron or aluminium are rendered passive on reaction with fuming HNO₃ due to formation of a thin oxide coating on the surface of the metal which prevents further reaction

CuCO₃ + 2HNO₃ [dil.] ⟶ Cu(NO₃)₂ + H₂O + CO₂

(i) Nitrogen dioxide gas

S + 6HNO₃ ⟶ H₂SO₄ + 2H₂O + 6NO₂

(ii) Oxygen gas

$2\text{KNO}_3 \xrightarrow{\Delta} 2\text{KNO}_2 + \text{O}_2$

Reddish brown nitrogen dioxide gas is evolved.

Buff coloured residue of PbO is obtained in the test tube.

$\text{Pb(NO}_3)_2 \xrightarrow{\Delta} 2\text{PbO} + \text{O}_2 + 4\text{NO}_2$

C + 4HNO₃ ⟶ CO₂ + 2H₂O + 4NO₂

Cold, dilute nitric acid reacts with copper to form nitric oxide.

(i) Balanced equation for laboratory preparation of nitric acid:

$\underset{[conc.]}{\text{KNO}_3} + \underset{[conc.]}{\text{H}_2\text{SO}_4} \xrightarrow{\lt 200 ^{\circ}\text{C}} \underset{[acid \space salt]}{\text{KHSO}_4} + \text{HNO}_3$

(ii) Balanced equation for action of heat on a mixture of copper and concentrated nitric acid:

$\text{Cu} + 4\underset{[conc.]}{\text{HNO}_3} \longrightarrow \text{Cu(NO}_3 \text{)}_2 + 2\text{H}_2 \text{O} + 2\text{NO}_2$

Reddish brown nitrogen dioxide gas is evolved.

Black coloured residue of CuO is obtained in the test tube.

$2\text{Cu(NO}_3)_2 \xrightarrow{\Delta} 2\text{CuO} + \text{O}_2 + 4\text{NO}_2$

Nitric acid