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Ch 1Chemical Reactions and Equations

Unit I: Chemical Substances — Nature and Behaviour (25 marks)

Chapter 1: Chemical Reactions and Equations

1. Chemical Reactions

A chemical reaction is a process in which one or more substances (reactants) are converted into new substances (products) with different properties. A chemical change is indicated by one or more of the following observations:

  • Change in state
  • Change in colour
  • Evolution of a gas
  • Change in temperature
  • Formation of a precipitate

2. Chemical Equations

A chemical equation is the symbolic representation of a chemical reaction using formulae of reactants and products. Example: burning of magnesium ribbon in air produces magnesium oxide (dazzling white flame; white powder MgO).

2Mg(s) + O2(g) → 2MgO(s)

A skeletal equation is an unbalanced equation. A balanced equation has an equal number of atoms of each element on both sides, satisfying the Law of Conservation of Mass (mass can neither be created nor destroyed in a chemical reaction).

Balancing (Hit and Trial Method)

  1. Write the skeletal equation.
  2. Count atoms of each element on both sides.
  3. Balance the compound with the maximum atoms first; balance elements one by one by adjusting coefficients (never change formulae).
  4. Finally add physical states: (s) solid, (l) liquid, (g) gas, (aq) aqueous.
Fe(s) + H2O(g) → Fe3O4(s) + H2(g)   (skeletal)
3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g)   (balanced)
Exam tip: Always write physical states and conditions (heat Δ, sunlight, catalyst) above/below the arrow — often 1 mark is reserved for states and balancing.

3. Types of Chemical Reactions

(a) Combination Reaction

Two or more reactants combine to form a single product.

CaO(s) + H2O(l) → Ca(OH)2(aq) + heat   (quick lime → slaked lime)
C(s) + O2(g) → CO2(g)
2H2(g) + O2(g) → 2H2O(l)

Whitewashing: Slaked lime Ca(OH)2 reacts slowly with CO2 of air to form a shiny layer of calcium carbonate (same formula as marble):

Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)

(b) Decomposition Reaction

A single reactant breaks down into two or more simpler products. Energy (heat, light or electricity) is required, so decomposition reactions are generally endothermic.

  • Thermal decomposition (by heat):
    2FeSO4(s) → Fe2O3(s) + SO2(g) + SO3(g)   (on heating, Δ)

    Green ferrous sulphate crystals lose water, colour changes to white then reddish-brown Fe2O3; smell of burning sulphur due to SO2 and SO3.

    2Pb(NO3)2(s) → 2PbO(s) + 4NO2(g) + O2(g)

    Lead nitrate on heating gives brown fumes of nitrogen dioxide (NO2) and yellow lead oxide.

    CaCO3(s) → CaO(s) + CO2(g)   (limestone → quick lime; used in cement industry)
  • Electrolytic decomposition (by electricity): Electrolysis of water gives hydrogen at cathode and oxygen at anode in the ratio 2:1 by volume.
    2H2O(l) → 2H2(g) + O2(g)
  • Photolytic decomposition (by sunlight):
    2AgCl(s) → 2Ag(s) + Cl2(g)   (white AgCl turns grey)
    2AgBr(s) → 2Ag(s) + Br2(g)

    These reactions are used in black and white photography.

(c) Displacement Reaction

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

Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

An iron nail dipped in copper sulphate solution becomes brownish (copper deposit) and the blue colour of the solution fades to light green (FeSO4).

Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
Pb(s) + CuCl2(aq) → PbCl2(aq) + Cu(s)

Zinc and lead are more reactive than copper, so they displace copper from its salt solutions.

(d) Double Displacement Reaction

Two compounds exchange their ions to form two new compounds. Usually a precipitate (insoluble solid) is formed — then it is also called a precipitation reaction.

Na2SO4(aq) + BaCl2(aq) → BaSO4(s)↓ + 2NaCl(aq)

White precipitate of barium sulphate is formed by the reaction of Ba2+ and SO42− ions.

(e) Oxidation, Reduction and Redox Reactions

  • Oxidation: gain of oxygen or loss of hydrogen.
  • Reduction: loss of oxygen or gain of hydrogen.
  • A reaction in which oxidation and reduction occur together is a redox reaction.
2Cu(s) + O2(g) → 2CuO(s)   (copper oxidised; black CuO coating)
CuO(s) + H2(g) → Cu(s) + H2O(l)   (CuO reduced to Cu, H2 oxidised to H2O — redox)
ZnO + C → Zn + CO   (ZnO reduced, C oxidised)
MnO2 + 4HCl → MnCl2 + 2H2O + Cl2   (HCl oxidised to Cl2, MnO2 reduced to MnCl2)
Trick: "Oxidising agent gets reduced; reducing agent gets oxidised." In CuO + H2 → Cu + H2O, CuO is the oxidising agent and H2 is the reducing agent.

4. Exothermic and Endothermic Reactions

ExothermicEndothermic
Heat is released along with products.Heat/energy is absorbed.
Examples: burning of natural gas CH4 + 2O2 → CO2 + 2H2O + heat; CaO + H2O; respiration; decomposition of vegetable matter into compost.Examples: most decomposition reactions (heating of CaCO3, FeSO4); photosynthesis; electrolysis of water.

Respiration is exothermic: glucose obtained from digested food is oxidised in cells to provide energy.

C6H12O6(aq) + 6O2(aq) → 6CO2(aq) + 6H2O(l) + energy

5. Effects of Oxidation in Daily Life

(a) Corrosion

When a metal is attacked by moisture, air, acids etc. present around it, it is said to corrode. Examples:

  • Rusting of iron: reddish-brown coating of hydrated iron(III) oxide, Fe2O3·xH2O, forms in presence of both air (O2) and moisture. Causes huge damage to bridges, ships, railings, car bodies.
  • Black coating on silver (Ag2S with H2S of air); green coating on copper (basic copper carbonate, CuCO3·Cu(OH)2, with moist CO2).

(b) Rancidity

When fats and oils in food are oxidised, their smell and taste change — the food becomes rancid. Prevention:

  • Adding antioxidants to foods containing fats and oils.
  • Storing food in airtight containers, refrigeration.
  • Flushing packets of chips with nitrogen gas (an inert atmosphere prevents oxidation).
Frequently asked 3-markers: (i) Why is respiration an exothermic reaction? (ii) Why are chips packets flushed with nitrogen? (iii) Why does the colour of CuSO4 solution change when an iron nail is dipped in it? Prepare these with equations.
Rationalisation note: The core syllabus of this chapter is fully retained for 2026-27. Detailed treatment of oxidation number method of balancing is not in NCERT Class 10; stick to the hit-and-trial method and gain/loss of oxygen-hydrogen definitions of redox.
Ch 2: Acids, Bases and Salts →