Quick Summary
Chemical reactions convert reactants into products. This chapter is important for identifying reaction types, balancing equations, writing observations, and understanding oxidation-reduction.
Key Definitions
- Reactants: Substances that take part in a chemical reaction.
- Products: New substances formed in a chemical reaction.
- Combination reaction: Two or more reactants form one product.
- Decomposition reaction: One compound breaks into simpler substances.
- Displacement reaction: A more reactive element displaces a less reactive element.
- Double displacement reaction: Exchange of ions between two compounds.
- Oxidation: Addition of oxygen, removal of hydrogen, or loss of electrons.
- Reduction: Removal of oxygen, addition of hydrogen, or gain of electrons.
Major Reaction Types
| Reaction type | General pattern | Example idea |
|---|---|---|
| Combination | A + B -> AB | Formation of one product |
| Decomposition | AB -> A + B | Breaking a compound |
| Displacement | A + BC -> AC + B | More reactive element replaces another |
| Double displacement | AB + CD -> AD + CB | Ion exchange |
| Neutralisation | Acid + Base -> Salt + Water | Acid-base reaction |
Balancing Chemical Equations
A balanced equation has the same number of atoms of each element on both sides. This follows the law of conservation of mass.
Reactants -> Products
Atoms before reaction = Atoms after reaction
Oxidation and Reduction
Oxidation and reduction often happen together in the same reaction. A substance that gets oxidised loses electrons, while a substance that gets reduced gains electrons.
Worked Examples (with real equations)
Example 1 — Combination
2Mg + O₂ → 2MgO (magnesium burns with a dazzling white flame)
CaO + H₂O → Ca(OH)₂ (quick lime + water, exothermic)
Example 2 — Decomposition
Thermal: CaCO₃ → CaO + CO₂↑ (heating limestone)
Electrolytic: 2H₂O → 2H₂ + O₂ (electrolysis of water)
Photolytic: 2AgCl → 2Ag + Cl₂ (silver chloride in sunlight)
Example 3 — Displacement
Fe + CuSO₄ → FeSO₄ + Cu
(blue colour of CuSO₄ fades; iron is more reactive than copper)
Example 4 — Double displacement (precipitation)
Na₂SO₄ + BaCl₂ → BaSO₄↓ + 2NaCl
(white precipitate of barium sulphate forms)
Example 5 — Redox (identify oxidised/reduced)
ZnO + C → Zn + CO
C is oxidised (gains O) ; ZnO is reduced (loses O)
Example 6 — Balancing step by step
Unbalanced: Fe + H₂O → Fe₃O₄ + H₂
Balance Fe and O: 3Fe + 4H₂O → Fe₃O₄ + 4H₂
Check: Fe 3=3, O 4=4, H 8=8 ✓
Textbook Solution Strategy
1. Write the reactants and products correctly. 2. Count atoms of each element on both sides. 3. Balance one element at a time using coefficients. 4. Identify the reaction pattern. 5. Mention observations such as colour change, gas evolution, precipitate, or heat.
Important Questions
One Mark
- Define chemical reaction.
- What is a decomposition reaction?
- What is oxidation?
Two Marks
- Differentiate between combination and decomposition reactions.
- What is a displacement reaction? Give the general form.
- Why should chemical equations be balanced?
Five Marks
- Explain any four types of chemical reactions with examples.
- Balance a given chemical equation and identify the reaction type.
- Explain oxidation and reduction with examples.
Common Mistakes
- Changing subscripts while balancing equations.
- Forgetting that coefficients are used for balancing.
- Confusing displacement and double displacement reactions.
- Treating oxidation and reduction as unrelated reactions.
Revision Checklist
- Major reaction types revised
- General patterns memorised
- Equation balancing practised
- Oxidation and reduction definitions revised
- Examples written for each reaction type
Frequently Asked Questions
A chemical reaction is a process in which reactants are converted into new substances called products.