Class 10 › Science › Chapter 1

Chemical Reactions and Equations

How to recognise a chemical reaction, write and balance chemical equations, the five types of reactions, energy changes, oxidation and reduction, and their effects in daily life.

1. Chemical Reactions and How to Recognise Them
Chemical reaction: a change in which the nature and identity of the initial substances (reactants) change and new substances (products) are formed.
Observations

Signs of a reaction

Change in state, change in colour, evolution of a gas, change in temperature.

Examples

From the activities

Mg burns with a dazzling white flame to give white MgO (Act. 1.1). Pb(NO3)2 + KI gives a yellow precipitate (Act. 1.2). Zn + dilute acid gives H2 bubbles and warmth (Act. 1.3).

  • Daily-life reactions: milk turning sour, rusting of iron, fermentation of grapes, cooking, digestion and respiration.
  • Why clean Mg ribbon? Sandpaper removes the protective oxide layer so Mg can burn in air. Never look directly at the flame.
2. Chemical Equations
  • Word-equation: reactants (LHS) and products (RHS) in words, joined by +, with an arrow pointing towards the products.
Magnesium + Oxygen → Magnesium oxide
  • Chemical equation: the same reaction written with chemical formulae.
  • Skeletal equation: an equation with correct formulae but not balanced. Example: Mg + O2 → MgO.
  • Balanced equation: the number of atoms of each element is the same on both sides.
Why balance? Law of conservation of mass: mass is neither created nor destroyed in a reaction, so atoms of each element must be equal before and after.
3. Balancing a Chemical Equation (Hit-and-Trial)
Worked example: Fe + H2O → Fe3O4 + H2
Step 1: Draw boxes around each formula. Never change anything inside a box; only add coefficients in front.
Step 2: Count atoms on both sides. Fe: 1 vs 3. H: 2 vs 2. O: 1 vs 4.
Step 3: Start with the compound having the most atoms (Fe3O4) and its element with the highest count (O). Put 4 before H2O: Fe + 4H2O → Fe3O4 + H2.
Step 4: H is now 8 on the left, so put 4 before H2: Fe + 4H2O → Fe3O4 + 4H2.
Step 5: Fe is 1 vs 3, so put 3 before Fe.
Step 6: Check every element on both sides.
3Fe + 4H2O → Fe3O4 + 4H2Balanced: Fe 3 = 3, H 8 = 8, O 4 = 4
  • Rule: to balance O in water, write 4H2O, not H2O4 or (H2O)4. Changing a subscript changes the substance.
  • Use the smallest whole-number coefficients.
  • Practice: Mg + O2 → MgO balances to 2Mg + O2 → 2MgO.
4. State Symbols and Reaction Conditions
SymbolMeaning
(s)Solid
(l)Liquid
(g)Gas
(aq)Aqueous: dissolved in water
3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g)(g) with H2O shows water is used as steam.
  • State symbols are written only when it is necessary to specify them.
  • Conditions (temperature, pressure, catalyst, sunlight) go above or below the arrow. Example: CO(g) + 2H2(g) 340 atm → CH3OH(l).
  • Photosynthesis: 6CO2(aq) + 12H2O(l) sunlight, chlorophyll → C6H12O6(aq) + 6O2(aq) + 6H2O(l).
5. Combination Reaction
Combination reaction: two or more substances (elements or compounds) combine to form a single product.
CaO(s) + H2O(l) → Ca(OH)2(aq) + HeatQuick lime + water gives slaked lime; the beaker becomes hot (Act. 1.4).
C(s) + O2(g) → CO2(g)
2H2(g) + O2(g) → 2H2O(l)
2Mg + O2 → 2MgOBurning of Mg (Act. 1.1) is also a combination reaction.
  • Whitewashing: slaked lime reacts slowly with CO2 in air: Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l). After 2 to 3 days CaCO3 (same formula as marble) gives a shiny finish.
6. Decomposition Reaction
Decomposition reaction: a single reactant breaks down into two or more simpler products. It is the opposite of a combination reaction.
Thermal

Energy = heat

2FeSO4(s) → Fe2O3(s) + SO2(g) + SO3(g)
Green crystals turn to a different colour; smell of burning sulphur.

CaCO3(s) → CaO(s) + CO2(g)
Used in making cement.

2Pb(NO3)2(s) → 2PbO(s) + 4NO2(g) + O2(g)
Brown fumes of NO2.

Electrolytic

Energy = electricity

2H2O(l) → 2H2(g) + O2(g) (Act. 1.7)

Gas at the cathode is H2 and at the anode is O2. The volume of H2 is double that of O2, since water has H : O = 2 : 1. H2 burns with a pop; O2 makes the candle burn brighter.

Photolytic

Energy = light

2AgCl(s) → 2Ag(s) + Cl2(g) (Act. 1.8)

2AgBr(s) → 2Ag(s) + Br2(g)

White silver chloride turns grey in sunlight. Used in black and white photography.

  • Decomposition reactions need energy (heat, light or electricity) to break the reactant.
  • Handle heating tubes carefully: do not point the mouth at anyone; waft gas gently towards the nose.
7. Displacement Reaction
Displacement reaction: a more reactive element displaces a less reactive element from its compound.
Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
Pb(s) + CuCl2(aq) → PbCl2(aq) + Cu(s)
  • Act. 1.9 observations: the blue colour of CuSO4 fades (light green FeSO4 forms) and the iron nail gets a brownish coating of copper.
  • Fe, Zn and Pb are more reactive than Cu, so they displace it.
8. Double Displacement Reaction
Double displacement reaction: two compounds exchange ions to form two new compounds.
Na2SO4(aq) + BaCl2(aq) → BaSO4(s) + 2NaCl(aq)White precipitate of BaSO4 forms from SO42− and Ba2+.
Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3(aq)Yellow precipitate (Act. 1.2); also a double displacement reaction.
  • Precipitate: an insoluble solid formed in a reaction. A reaction that forms one is a precipitation reaction.
  • Displacement vs double displacement: in displacement, one element replaces another in a compound; in double displacement, two compounds swap ions.
9. Exothermic and Endothermic Reactions
Exothermic

Heat is released

CaO + H2O, burning of natural gas (CH4 + 2O2 → CO2 + 2H2O), respiration, and decomposition of vegetable matter into compost.

Endothermic

Energy is absorbed

Decomposition reactions (heat, light or electricity needed). Ba(OH)2 + NH4Cl makes the test tube feel cold.

  • Respiration is exothermic: glucose + oxygen gives CO2 + H2O + energy, C6H12O6(aq) + 6O2(aq) → 6CO2(aq) + 6H2O(l) + energy.
10. Oxidation and Reduction (Redox)
Oxidation

Gain of oxygen or loss of hydrogen

Reduction

Loss of oxygen or gain of hydrogen

2Cu + O2 → 2CuOCu gains oxygen: Cu is oxidised (black copper(II) oxide coating, Act. 1.11).
CuO + H2 → Cu + H2OCuO loses oxygen: reduced. H2 gains oxygen: oxidised. The black coating turns brown.
  • Redox reaction: oxidation and reduction occur together. One reactant is oxidised while the other is reduced.
  • ZnO + C → Zn + CO: C is oxidised, ZnO is reduced.
  • MnO2 + 4HCl → MnCl2 + 2H2O + Cl2: HCl is oxidised (to Cl2), MnO2 is reduced.
  • Mg burning in air: Mg gains oxygen, so it is oxidised.
  • The substance that is oxidised and the one that is reduced are always different substances.
11. Effects of Oxidation in Everyday Life
Corrosion

Metal attacked by moisture, acids, etc.

Rusting of iron gives a reddish-brown coating. Silver turns black; copper gets a green coating. It damages car bodies, bridges, railings and ships and costs huge sums to repair. Prevention: painting iron articles.

Rancidity

Oxidation of fats and oils

Smell and taste of food change. Prevention: add antioxidants, use air-tight containers, flush chip bags with nitrogen.

12. Quick Revision
TypeKey ideaExample
CombinationA + B → ABC + O2 → CO2
DecompositionAB → A + BCaCO3 → CaO + CO2
DisplacementA + BC → AC + BFe + CuSO4 → FeSO4 + Cu
Double displacementAB + CD → AD + CBNa2SO4 + BaCl2 → BaSO4 + 2NaCl
RedoxOxidation + reduction togetherCuO + H2 → Cu + H2O
Check

Balanced equation test

Count each element on both sides after adding coefficients. Every element must match.

Remember

Energy and precipitate

Exothermic releases heat; endothermic absorbs energy. A precipitate is an insoluble solid formed in solution.

Common mistakes:
• Changing subscripts (H2O4) instead of adding coefficients.
• Forgetting to count atoms inside brackets, e.g. in Ca(OH)2.
• Calling a reaction decomposition when it has more than one reactant.
• Mixing up oxidised and reduced substances.
• Writing H2 : O2 volumes as 1 : 2 in electrolysis; it is 2 : 1.
• Leaving out conditions or state symbols when the question asks for them.