Class 10 Science - Chapter 1 Chemical Reactions and Equations (Notes)

Introduction to Chemical Reactions

A chemical reaction involves a change in the nature and identity of substances, resulting in the formation of new substances. This chapter explores how to identify, represent, and classify chemical reactions.

Examples of Chemical Reactions in Daily Life:

• Milk spoiling at room temperature (chemical change).
• Iron rusting in humid conditions.
• Fermentation of grapes.
• Cooking food, digestion, and respiration.

Indicators of a Chemical Reaction:

• Change in state.
• Change in colour.
• Evolution of a gas.
• Change in temperature.

Chemical Equations

Writing Chemical Equations

A chemical equation represents reactants (starting substances) and products (resulting substances) using chemical formulae. For example:

• Word Equation: Magnesium + Oxygen → Magnesium oxide.
• Chemical Equation: Mg + O₂ → MgO (skeletal equation).

Reactants: Written on the left-hand side (LHS) with a plus sign (+).

Products: Written on the right-hand side (RHS) with a plus sign (+).

Arrow (→): Indicates the direction of the reaction.

Balancing Chemical Equations

The law of conservation of mass states that mass is neither created nor destroyed in a chemical reaction. Thus, the number of atoms of each element must be equal on both sides of the equation.

Example: Balancing Fe + H₂O → Fe₃O₄ + H₂

Steps to Balance:

1. Draw boxes around each formula: [Fe] + [H₂O] → [Fe₃O₄] + [H₂].
2. List atoms:

   Element	LHS	RHS
   Fe	1	3
   H	2	2
   O	1	4

3. Start with the compound with the most atoms (e.g., Fe₃O₄). Balance oxygen by adjusting H₂O: Fe + 4H₂O → Fe₃O₄ + H₂.
4. Balance hydrogen: Fe + 4H₂O → Fe₃O₄ + 4H₂.
5. Balance iron: 3Fe + 4H₂O → Fe₃O₄ + 4H₂.
6. Verify:

   Element	LHS	RHS
   Fe	3	3
   H	8	8
   O	4	4

7. Add physical states: 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g).

Physical States: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution.

Note: Reaction conditions (e.g., temperature, catalyst) may be indicated above/below the arrow, e.g., CO(g) + 2H₂(g) →[340 atm] CH₃OH(l).

Types of Chemical Reactions

Combination Reaction

Two or more reactants combine to form a single product.

• Example: CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat (Quick lime + Water → Slaked lime).
• Others:
  • C(s) + O₂(g) → CO₂(g) (Burning of coal).
  • 2H₂(g) + O₂(g) → 2H₂O(l) (Formation of water).

Exothermic Reactions: Release heat, e.g., Activity 1.1 (Mg → MgO, heat released).

Other Exothermic Examples:

• Burning of natural gas: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g).
• Respiration: C₆H₁₂O₆(aq) + 6O₂(aq) → 6CO₂(aq) + 6H₂O(l) + energy.
• Decomposition of vegetable matter into compost.

Decomposition Reaction

A single reactant breaks down into two or more products (opposite of combination).

• Thermal Decomposition (Heat):
  • CaCO₃(s) →[Heat] CaO(s) + CO₂(g) (Limestone → Quick lime).
  • Activity 1.5: 2FeSO₄·7H₂O(s) →[Heat] Fe₂O₃(s) + SO₂(g) + SO₃(g) (Green crystals turn brown, gases evolve).
  • Activity 1.6: 2Pb(NO₃)₂(s) →[Heat] 2PbO(s) + 4NO₂(g) + O₂(g) (Brown fumes of NO₂).
• Photolytic Decomposition (Light):
  • Activity 1.8: 2AgCl(s) →[Sunlight] 2Ag(s) + Cl₂(g) (White to grey, used in photography).
  • 2AgBr(s) →[Sunlight] 2Ag(s) + Br₂(g).
• Electrolytic Decomposition (Electricity):
  • Activity 1.7: 2H₂O(l) →[Electricity] 2H₂(g) + O₂(g) (Hydrogen:Oxygen ratio 2:1 by volume).

Endothermic Reactions: Absorb energy (heat, light, electricity), e.g., decomposition reactions.

Displacement Reaction

A more reactive element displaces a less reactive element from its compound.

• Example: Activity 1.9: Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s) (Iron displaces copper, blue colour fades).
• Others:
  • Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s).
  • Pb(s) + CuCl₂(aq) → PbCl₂(aq) + Cu(s).

Double Displacement Reaction

Exchange of ions between two reactants, often forming a precipitate.

• Example: Activity 1.10: Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq) (White precipitate of BaSO₄).
• Activity 1.2: Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq) (Yellow precipitate of PbI₂).

Precipitation Reaction: Produces an insoluble product (precipitate).

Oxidation and Reduction (Redox) Reactions

Involve gain/loss of oxygen or hydrogen.

• Oxidation: Gain of oxygen or loss of hydrogen.
  • 2Cu + O₂ →[Heat] 2CuO (Copper oxidised to black CuO).
  • C + O₂ → CO₂ (Carbon oxidised).
• Reduction: Loss of oxygen or gain of hydrogen.
  • CuO + H₂ →[Heat] Cu + H₂O (CuO reduced to Cu).
  • ZnO + C → Zn + CO (ZnO reduced to Zn).

Redox Reaction: One substance is oxidised, another reduced, e.g., CuO + H₂ → Cu + H₂O.

Effects of Oxidation in Everyday Life

Corrosion

Metals react with substances (moisture, acids) and deteriorate.

• Examples:
  • Iron rusting (reddish-brown coating).
  • Copper forming green patina.
  • Silver developing black tarnish.
• Impact: Damages cars, bridges, ships; costs billions to repair.
• Prevention: Painting iron articles to block moisture.

Rancidity

Oxidation of fats/oils in food changes smell and taste.

• Prevention:
  • Add antioxidants to foods.
  • Store in air-tight containers.
  • Flush chips packets with nitrogen gas.

Key Terms

• Chemical Reaction: Process where substances (reactants) form new substances (products).
• Balanced Equation: Equal number of atoms of each element on both sides.
• Combination Reaction: Two or more reactants form one product.
• Decomposition Reaction: One reactant breaks into multiple products.
• Exothermic Reaction: Releases heat.
• Endothermic Reaction: Absorbs energy.
• Displacement Reaction: More reactive element displaces less reactive one.
• Double Displacement Reaction: Ion exchange between reactants.
• Redox Reaction: Involves oxidation and reduction.
• Corrosion: Metal deterioration due to environmental reactions.
• Rancidity: Oxidation of fats/oils in food.

Exercises and Solutions

Questions (Page 6)

• 1. Magnesium ribbon is cleaned to remove oxide coating for faster burning.
• 2. Balanced equations:
  • (i) H₂ + Cl₂ → 2HCl
  • (ii) 3BaCl₂ + Al₂(SO₄)₃ → 3BaSO₄ + 2AlCl₃
  • (iii) 2Na + 2H₂O → 2NaOH + H₂
• 3. Balanced equations with states:
  • (i) BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)
  • (ii) NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)

Questions (Page 10)

• 1.
  • (i) Substance X: Calcium oxide (CaO).
  • (ii) CaO(s) + H₂O(l) → Ca(OH)₂(aq)
• 2. Hydrogen gas (H₂) is collected in double volume due to 2:1 ratio in water (2H₂O → 2H₂ + O₂).

Questions (Page 13)

• 1. Iron displaces copper: Fe + CuSO₄ → FeSO₄ + Cu, fading blue colour.
• 2. Example: Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq).
• 3. Oxidised/Reduced:
  • (i) Na oxidised (gains O), O₂ reduced.
  • (ii) H₂ oxidised (gains O), CuO reduced.

Exercises (Pages 14-15)

• 1. Incorrect: (i) (a) and (b). PbO is reduced, C is oxidised.
• 2. (d) Displacement reaction.
• 3. (a) Hydrogen gas and iron chloride produced.
• 4. Balanced equation: Equal atoms on both sides, required by conservation of mass.
• 5. Balanced equations:
  • (a) 3H₂ + N₂ → 2NH₃
  • (b) 2H₂S + 3O₂ → 2H₂O + 2SO₂
  • (c) 3BaCl₂ + Al₂(SO₄)₃ → 3BaSO₄ + 2AlCl₃
  • (d) 2K + 2H₂O → 2KOH + H₂
• 6. Balanced:
  • (a) 2HNO₃ + Ca(OH)₂ → Ca(NO₃)₂ + 2H₂O
  • (b) 2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O
  • (c) NaCl + AgNO₃ → AgCl + NaNO₃
  • (d) BaCl₂ + H₂SO₄ → BaSO₄ + 2HCl
• 7. Balanced:
  • (a) Ca(OH)₂ + CO₂ → CaCO₃ + H₂O
  • (b) Zn + 2AgNO₃ → Zn(NO₃)₂ + 2Ag
  • (c) 2Al + 3CuCl₂ → 2AlCl₃ + 3Cu
  • (d) BaCl₂ + K₂SO₄ → BaSO₄ + 2KCl
• 8. Balanced and Type:
  • (a) 2KBr(aq) + BaI₂(aq) → 2KI(aq) + BaBr₂(s) (Double displacement).
  • (b) ZnCO₃(s) → ZnO(s) + CO₂(g) (Decomposition).
  • (c) H₂(g) + Cl₂(g) → 2HCl(g) (Combination).
  • (d) Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g) (Displacement).
• 9. Exothermic: Releases heat (e.g., CH₄ + 2O₂ → CO₂ + 2H₂O). Endothermic: Absorbs energy (e.g., CaCO₃ → CaO + CO₂).
• 10. Respiration is exothermic as glucose oxidation releases energy.
• 11. Decomposition splits one reactant into multiple products, opposite of combination. E.g., CaCO₃ → CaO + CO₂ vs. CaO + CO₂ → CaCO₃.
• 12. Decomposition examples:
  • Heat: CaCO₃ → CaO + CO₂
  • Light: 2AgCl → 2Ag + Cl₂
  • Electricity: 2H₂O → 2H₂ + O₂
• 13. Displacement: One element displaces another (e.g., Fe + CuSO₄ → FeSO₄ + Cu). Double displacement: Ion exchange (e.g., Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl).
• 14. Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag.
• 15. Precipitation: Forms insoluble product, e.g., Na₂SO₄ + BaCl₂ → BaSO₄(s) + 2NaCl.
• 16. Oxidation: 2Cu + O₂ → 2CuO, C + O₂ → CO₂. Reduction: CuO + H₂ → Cu + H₂O, ZnO + C → Zn + CO.
• 17. Element X: Copper (Cu), Compound: Copper(II) oxide (CuO).
• 18. Paint prevents corrosion by blocking moisture.
• 19. Nitrogen prevents rancidity by reducing oxidation.
• 20.
  • Corrosion: Iron rusting due to moisture.
  • Rancidity: Spoilage of oily food due to oxidation.