Chemical Reactions

Exercise I

Question 1

(a) Define a chemical reaction.

(b) What happens during a chemical reaction?

(c) What do you understand by a chemical bond?

Answer

(a) Any chemical change in matter which involves transformation into one or more substances with entirely different properties is called a chemical reaction.

(b) A chemical reaction involves breaking of chemical bonds between the atoms or groups of atoms of reacting substances and rearrangement of atoms making new bonds to form new substances with absorption or release of energy in the form of heat and light normally.

(c) A chemical bond is the attractive force that holds the atoms of a molecule together in a compound.

Question 2

Give one example each that illustrates the following characteristics of a chemical reaction:

(a) evolution of a gas

(b) change of colour

(c) change in state

Answer

(a) In many chemical reactions one of the product is a gas.

When zinc reacts with dilute sulphuric acid, hydrogen gas is evolved with an effervescence.

Zn + H₂SO₄ (dil) ⟶ ZnSO₄ + H₂ (g)

(b) Certain chemical reactions are characterised by a change in the colour of the reactants.

When a few pieces of iron are dropped into a blue coloured copper sulphate solution, the blue colour of the solution fades and turns into light green.

Fe + CuSO₄ (aq) ⟶ FeSO₄ + Cu

(c) In many chemical reactions a change of state is observed.

Ammonia and hydrogen chloride gases react to produce a white solid of ammonium chloride.

NH₃ (g) + HCl (g) ⟶ NH₄Cl (s)

Question 3

How do the following help in bringing about a chemical change? Explain each with a suitable example.

(a) Pressure

(b) Light

(c) Catalyst

(d) Heat

Answer

(a) Some chemical reactions take place when the reactants are subjected to high pressure.

When nitrogen and hydrogen gas are subjected to high pressure, ammonia gas is produced.

$\text{N}_2 + 3\text{H}_2 \xrightleftharpoons[\underset{\text{pressure}}{\text{200 atmos.}}]{\text{450 \degree C }} \text{2NH}_3$

(b) Some chemical reactions can take place only in the presence of light and are called photochemical reactions.

Photosynthesis is a chemical reaction in which glucose is prepared by the green leaves of a plant but light is necessary for the reaction to take place.

$\underset{\text{[carbon dioxide]}}{6\text{CO}_2} + 6\text{H}_2\text{O} \xrightarrow[\text{Sunlight}]{\text{Chlorophyll}} \underset{\text{[glucose]}}{\text{C}_6\text{H}_{12}\text{O}_6} + \underset{\text{[oxygen]}}{6\text{O}_2}$

(c) Some chemical reactions need a catalyst to change the rate of reaction, if the reaction is too slow or too fast. Positive catalyst increases the rate of chemical reaction whereas negative catalyst decreases the rate of reaction.

Finely divided iron is used as a positive catalyst in the manufacture of ammonia from hydrogen and oxygen.

     $\text{N}_2 + 3\text{H}_2 \xrightleftharpoons[\text{200-900 atmos.}]{\text{Fe catalyst/450\degree C }} \text{2NH}_3$

(d) Some chemical reactions take place only in the presence of heat.

If iron powder and sulphur powder are mixed, they do not react. But they are heated, they react to form iron sulphide.

Fe + S $\xrightarrow{\enspace\Delta\enspace}$ FeS

Question 4

(a) Define catalyst.

(b) What are (i) positive catalysts and (ii) negative catalysts? Support your answer with one example for each of them.

(c) Name three biochemical catalysts found in the human body.

Answer

(a) A catalyst is a substance that either increases or decreases the rate of a chemical reaction without itself undergoing any chemical change during the reaction.

(b) Positive Catalyst — When a catalyst increases the rate of a chemical reaction, it is known as a positive catalyst.
For example, Finely divided iron is used as a positive catalyst in the manufacture of ammonia from hydrogen and oxygen.

$\text{N}_2 + 3\text{H}_2 \xrightleftharpoons[\text{200-900 atmos.}]{\text{Fe catalyst/450\degree C }} \text{2NH}_3$

Negative Catalyst — When a catalyst decreases the rate of a chemical reaction, it is known as a negative catalyst.
For example, Phosphoric acid acts as a negative catalyst to decrease the rate of decomposition of hydrogen peroxide.

(c) The three biochemical catalysts found in the human body are amylase, trypsin, pepsin.

Question 5

What do you observe when:

(a) dilute sulphuric acid is added to granulated zinc?

(b) a few pieces of iron are dropped in a blue solution of copper sulphate?

(c) silver nitrate is added to a solution of sodium chloride?

(d) ferrous sulphate solution is added to an aqueous solution of sodium hydroxide?

(e) solid lead nitrate is heated?

(f) when dilute sulphuric acid is added to barium chloride solution?

Answer

(a) When dilute sulphuric acid is added to granulated zinc, hydrogen gas is evolved with an effervescence.

Zn + H₂SO₄ (dil) ⟶ ZnSO₄ + H₂ (g)

(b) When a few pieces of iron are dropped into a blue coloured copper sulphate solution, the blue colour of the solution fades and turns into light green.

Fe + CuSO₄ (aq) ⟶ FeSO₄ + Cu

(c) When a solution of silver nitrate is added to a solution of sodium chloride a white insoluble precipitate of silver chloride is formed.

AgNO₃ (aq) + NaCl (aq) ⟶ NaNO₃ (aq) + AgCl ↓

(d) When ferrous sulphate solution is added to sodium hydroxide solution, a dirty green precipitate of ferrous hydroxide is formed.

FeSO₄ (aq) + 2NaOH (aq) ⟶ Fe(OH)₂ ↓ + Na₂SO₄ (aq)

(e) When solid lead nitrate is heated strongly, it decomposes to produce light yellow solid lead monoxide, reddish brown nitrogen dioxide gas and colourless oxygen gas.

2Pb(NO₃)₂ ⟶ 2PbO + 4NO₂ ↑ + O₂ ↑

(f) When dilute sulphuric acid is added to barium chloride solution, a white precipitate of barium sulphate is formed.

BaCl₂ + H₂SO₄ ⟶ BaSO₄ ↓ + 2HCl

Question 6

Complete and balance the following chemical equations:

(a) N₂ + O₂ ⟶

(b) H₂S + Cl₂ ⟶

(c) Na + H₂O ⟶

(d) NaCl + AgNO₃ ⟶

(e) Zn + H₂SO₄ (dil.) ⟶

(f) FeSO₄ (aq) + NaOH (aq) ⟶

(g) Pb(NO₃)₂ $\xrightarrow{\enspace\Delta\enspace}$

(h) BaCl₂ (aq) + H₂SO₄ (aq) ⟶

Answer

(a) N₂ + O₂ ⟶ 2NO ↑

(b) H₂S + Cl₂ ⟶ 2HCl ↑ + S

(c) 2Na + 2H₂O ⟶ 2NaOH + H₂ ↑

(d) NaCl + AgNO₃ ⟶ NaNO₃ + AgCl ↓

(e) Zn + H₂SO₄ (dil.) ⟶ ZnSO₄ + H₂ ↑

(f) FeSO₄ (aq) + 2NaOH (aq) ⟶ Fe(OH)₂ ↓ + Na₂SO₄ (aq)

(g) 2Pb(NO₃)₂ $\xrightarrow{\enspace\Delta\enspace}$ 2PbO + 4NO₂ ↑ + O₂ ↑

(h) BaCl₂ (aq) + H₂SO₄ (aq) ⟶ BaSO₄ ↓ + 2HCl

Exercise II

Question 1

Fill in the blanks:

(a) A reaction in which two or more substances combine to form a single substance is called a ............... reaction.

(b) A ............... is a substance which changes the rate of a chemical reaction without undergoing a chemical change itself.

(c) The formation of gas bubbles in a liquid during a reaction is called ............... .

(d) The reaction between an acid and a base is called ............... .

(e) Soluble bases are called ............... .

(f) The chemical change involving iron and hydrochloric acid illustrates a ............... reaction.

(g) In the type of reaction called ............... two compounds exchange their positive an negative radicals respectively.

(h) A catalyst either ............... or ............... the rate of a chemical change but itself remains ............... at the end of the reaction.

(i) The chemical reaction between hydrogen and chlorine is a ............... reaction.

(j) When a piece of copper is added to silver nitrate solution, it turns ............... in colour.

Answer

(a) combination

(b) catalyst

(c) effervescence

(d) neutralization reaction

(e) alkalis

(f) displacement

(g) double displacement

(h) increases, decreases, unchanged

(i) combination

(j) blue

Question 2

Classify the following reactions as combination, decomposition, displacement, precipitation and neutralization. Also balance the equations.

(a) CaCO₃ (s) $\xrightarrow{\enspace\Delta\enspace}$ CaO (s) + CO₂ (g)

(b) Zn(s) + H₂SO₄ ⟶ ZnSO₄ (s) + H₂ (g)

(c) AgNO₃ (aq) + NaCl (aq) ⟶ AgCl(s) + NaNO₃

(d) NH₃ (g) + HCl (g) ⟶ NH₄Cl (s)

(e) CuSO₄ (aq) + H₂S (g) ⟶ CuS (s) + H₂SO₄ (l)

(f) Zn (s) + CuSO₄ (aq) ⟶ ZnSO₄ (aq) + Cu(s)

(g) Ca(s) + O₂ (g) ⟶ CaO(s)

(h) NaOH + HCl ⟶ NaCl + H₂O

(i) KOH + H₂SO₄ ⟶ K₂SO₄ + H₂O

Answer

(a) Decomposition reaction

(b) Displacement reaction

(c) Precipitation reaction

(d) Combination reaction

(e) Precipitation reaction

(f) Displacement reaction

(g) Combination reaction
      Balanced Equation:
      2Ca(s) + O₂ (g) ⟶ 2CaO(s)

(h) Neutralization reaction

(i) Neutralization reaction
     Balanced Equation:
     2KOH + H₂SO₄ ⟶ K₂SO₄ + 2H₂O

Question 3

Define:

(a) Precipitation

(b) Neutralization

(c) Catalyst

Answer

(a) A chemical reaction in which two compounds in their aqueous state react to form an insoluble solid (a precipitate) as one of the products is known as a precipitation reaction.

(b) A chemical reaction in which a base or an alkali reacts with an acid to produce a salt and water only is known as a neutralization reaction.

(c) A catalyst is a substance that either increases or decreases the rate of a chemical reaction without itself undergoing any chemical change during the reaction.

Question 4

Explain the following types of chemical reactions giving two examples for each of them:

(a) combination reaction

(b) decomposition reaction

(c) displacement reaction

(d) double displacement reaction

Answer

(a) Combination Reaction — A reaction in which two or more substances combine to form a single substance is called a combination reaction. It is also called a synthesis reaction.

Examples:

(i) Two elements combine to form a compound.

When iron and sulphur (both elements) are heated together they combine to form a compound iron sulphide.

Fe (s) + S (s) $\xrightarrow{\enspace\Delta\enspace}$ FeS (s)

(ii) An element and a compound can also combine to form a product.

Carbon monoxide, a compound, burns in presence of oxygen, an element, to form a single product, carbon dioxide.

2CO (g) + O₂ (g) $\xrightarrow{\enspace\Delta\enspace}$ 2CO₂ (g)

(b) Decomposition Reaction — A reaction in which a compound breaks up on heating into two or more simpler substances is called a decomposition reaction.

Examples:

(i) Calcium carbonate decomposes on strong heating to form two compounds, calcium oxide and carbon dioxide.

CaCO₃ (s) $\xrightarrow{\enspace\Delta\enspace}$ CaO (s) + CO₂ (g)

(ii) Mercuric oxide when heated decomposes to form mercury and oxygen.

2HgO (s) $\xrightarrow{\enspace\Delta\enspace}$ 2Hg (s) + O₂ (g)

(c) Displacement Reaction — A reaction in which a more active element displaces a less active element from a compound is called a displacement reaction.

Examples:

(i) Zinc is more reactive than copper so it displaces copper from copper sulphate solution to form zinc sulphate.

Zn (s) + CuSO₄ (aq) ⟶ ZnSO₄ (aq) + Cu (s)

(ii) Chlorine is more reactive than bromine, so it displaces bromine from potassium bromide solution.

2KBr + Cl₂ ⟶ 2KCl + Br₂

(d) Double Displacement Reaction — A chemical reaction in which two compounds in their aqueous state exchange their ions or radicals to form new compounds is called a double decomposition reaction.

AB + CD ⟶ CB + AD

Here AB and CD are reactants. They exchange their ions to form CB and AD which are the products.

Examples:

(i) When a solution of silver nitrate is added to a solution of sodium chloride a precipitate of silver chloride is formed.

AgNO₃ (aq) + NaCl (aq) ⟶ NaNO₃ (aq) + AgCl ↓

(ii) When ferrous sulphate solution is added to sodium hydroxide solution, a dirty green precipitate of ferrous hydroxide is formed.

FeSO₄ (aq) + 2NaOH (aq) ⟶ Fe(OH)₂ ↓ + Na₂SO₄ (aq)

Question 5

Write the missing reactants and products and balance the equations:

(a) NaOH + ............... ⟶ NaCl + ................

(b) KClO₃ $\xrightarrow{\text{heat}}$ ................. + ...................

(c) ............. + HCl ⟶ NaCl + H₂O + .................

Answer

(a) NaOH + HCl ⟶ NaCl + H₂O

(b) 2KClO₃ $\xrightarrow{\text{heat}}$ 2KCl + 3O₂

(c) Na₂SO₃ + 2HCl ⟶ 2NaCl + H₂O + SO₂

Question 6

How will you obtain?

(a) Magnesium oxide from magnesium.

(b) Silver chloride from silver nitrate.

(c) Nitrogen dioxide from lead nitrate.

(d) Zinc chloride from zinc.

(e) Ammonia from nitrogen.

Also give balanced equations for the reactions.

Answer

(a) When magnesium reacts with oxygen it produces magnesium oxide.

2Mg + O₂ ⟶ 2MgO

(b) When silver nitrate reacts with hydrochloric acid silver chloride is formed.

AgNO₃ (aq) + HCl ⟶ AgCl↓ + HNO₃

(c) When lead nitrate is heated strongly, a reddish brown gas nitrogen dioxide is produced.

2Pb(NO₃)₂ $\xrightarrow{\enspace\Delta\enspace}$ 2PbO + 4NO₂ (g) + O₂ (g)

(d) Zinc reacts with dilute HCl to produce zinc chloride.

Zn + 2HCl (dil.) ⟶ ZnCl₂ + H₂ (g)

(e) When nitrogen gas reacts with hydrogen gas in presence of finely divided iron as catalyst, subjected to pressure of 200-900 atm and temperature of about 450°C, ammonia gas is produced.

     $\text{N}_2 + 3\text{H}_2 \xrightleftharpoons[\text{200-900 atmos.}]{\text{Fe catalyst/450\degree C }} \text{2NH}_3$

Question 7

What do you observe when:

(a) Iron nail is kept in copper sulphate solution for sometime?

(b) Phenolphthalein is added to sodium hydroxide solution?

(c) Blue litmus paper is dipped in dilute hydrochloric acid?

(d) Lead nitrate is heated?

(e) Magnesium ribbon is burnt in oxygen?

(f) Ammonia is brought in contact with hydrogen chloride gas?

Answer

(a) When iron nail is kept in a blue coloured copper sulphate solution for sometime, a reddish brown coating is seen on the iron nail and the colour of the solution changes gradually from blue to light green. The reason for this observation is that iron being more reactive than copper, displaces copper from the solution and copper is deposited on the iron nail.
Fe + CuSO₄ (aq) ⟶ FeSO₄ + Cu

(b) When a few drops of phenolphthalein is added to sodium hydroxide solution, it turns pink indicating that sodium hydroxide solution is basic in nature.

(c) Blue litmus paper changes into red colour when dipped in dilute hydrochloric acid confirming its acidic nature.

(d) When solid lead nitrate is heated strongly, it decomposes to produce light yellow solid lead monoxide, reddish brown nitrogen dioxide gas and colourless oxygen gas.

2Pb(NO₃)₂ $\xrightarrow{\enspace\Delta\enspace}$ 2PbO + 4NO₂ (g) + O₂ (g)

(e) When magnesium ribbon is burnt in oxygen gas, it burns with a dazzling white light and produces a white powder of magnesium oxide.

2Mg + O₂ ⟶ 2MgO

(f) When ammonia is brought in contact with hydrogen chloride gas, dense white fumes of ammonium chloride are produced.

NH₃ + HCl ⟶ NH₄Cl

Question 8a

Give reason:

A person suffering from acidity is advised to take an antacid.

Answer

Antacids are alkaline in nature. So it neutralizes the acid produced in the stomach. Hence, antacids are advised for acidity.

Question 8b

Give reason:

Acidic soil is treated with quicklime.

Answer

Quicklime is a base which neutralizes the acidic soil. So acidic soil is treated with base like quicklime.

Question 8c

Give reason:

Wasp sting is treated with vinegar.

Answer

Wasp sting is alkaline in nature which is neutralized by vinegar which is a weak acid.

Question 9

What is meant by metal reactivity series? State its importance (any two points).

Answer

A list in which the metals are arranged in a decreasing order of their chemical reactivity is called the metal reactivity series.

The importance of metal reactivity series are:

1. The series facilitates the comparative study of metals in terms of the degree of their reactivity.
2. The compounds of the metals like oxides, carbonates, nitrates and hydroxides too can be easily compared.

Question 10

What are oxides? Give two examples of each of the following oxides.

(a) Basic oxide

(b) Acidic oxide

(c) Amphoteric oxide

(d) Neutral oxide

Answer

Oxides are compounds which essentially contains oxygen in its molecule, chemically combined with a metal or a non-metal.

(a) Calcium oxide (CaO) and magnesium oxide (MgO) are basic oxides.

(b) Carbon dioxide (CO₂) and sulphur dioxide (SO₂) are acidic oxides.

(c) Zinc oxide (ZnO) and aluminium oxide (Al₂O₃) are amphoteric oxides.

(d) Carbon monoxide (CO) and nitric oxide (NO) are neutral oxides.

Question 11

Define exothermic and endothermic reactions. Give two examples of each.

Answer

Exothermic Reaction

A chemical reaction in which heat is given out is called an exothermic reaction.

Examples:

1. When carbon burns in oxygen, carbon dioxide is formed and a lot of heat is produced.
   C + O₂ ⟶ CO₂ + Heat
2. When water is added to quicklime, slaked lime is formed and along with it a large amount of heat energy is produced.
   CaO + H₂O ⟶ Ca(OH)₂ + Heat

Endothermic Reaction

A chemical reaction in which heat is absorbed is called an endothermic reaction.

Examples:

1. When nitrogen and oxygen gas are heated together to a temperature of about 3000°C, nitric oxide gas is formed.
   N₂ + O₂ $\xrightarrow[3000\degree\text{C}]{\text{heat}}$ 2NO
2. When calcium carbonate is heated at 1000°C, it decomposes to give calcium oxide and carbon dioxide.
   CaCO₃ $\xrightarrow[1000\degree\text{C}]{\text{heat}}$ CaO + CO₂

Question 12

State the effect of the following on the surroundings:

(a) an endothermic reaction

(b) an exothermic reaction

Answer

(a) An endothermic reaction causes fall in temperature. These reactions absorb heat from the surroundings leading to cooling effect of the surroundings.

(b) An exothermic reaction causes rise in temperature. These reactions release heat into surroundings leading to increase in temperature of the surroundings.

Question 13

What do you observe when:

(a) an acid is added to a basic solution?

(b) ammonium chloride is dissolved in water?

Answer

(a) When an acid is added to a basic solution, neutralization reaction takes place and salt and water are produced. For example:
NaOH + HCl ⟶ NaCl + H₂O

(b) When ammonium chloride is dissolved in water, heat is absorbed leading to fall in temperature. So the apparatus in which this reaction occurs becomes colder than earlier.