Short Answer Questions — Chapter 10: Types of Chemical Reactions
Combination (Synthesis) reaction: A chemical reaction in which two or more reactants combine to form a single product.
General form: A + B → AB
Examples:
1. Formation of water: 2H₂ + O₂ → 2H₂O (Hydrogen combines with oxygen)
2. Burning of magnesium: 2Mg + O₂ → 2MgO (Magnesium ribbon burns with bright white light)
3. Formation of sulphur dioxide: S + O₂ → SO₂
4. Quicklime + water: CaO + H₂O → Ca(OH)₂ (Slaking of lime — highly exothermic)
Combination reactions are usually exothermic (they release energy).
Decomposition reaction: A single compound breaks down into two or more simpler substances.
General form: AB → A + B
Types:
Single Displacement (Substitution) Reaction: A more reactive element displaces a less reactive element from its compound in solution.
General form: A + BC → AC + B (A is more reactive than B)
Examples:
- Fe + CuSO₄ → FeSO₄ + Cu (iron displaces copper — iron more reactive)
- Zn + H₂SO₄ → ZnSO₄ + H₂↑ (zinc displaces hydrogen from acid)
- Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag (copper displaces silver)
Activity series (reactivity): K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > H > Cu > Ag > Au
Double Displacement (Metathesis) Reaction: The cations and anions of two different ionic compounds exchange partners.
General form: AB + CD → AD + CB
Usually results in: formation of a precipitate (↓), a gas (↑), or water
Examples:
- AgNO₃ + NaCl → AgCl↓ + NaNO₃ (white precipitate of AgCl)
- BaCl₂ + Na₂SO₄ → BaSO₄↓ + 2NaCl (white precipitate of BaSO₄)
- Na₂CO₃ + HCl → 2NaCl + H₂O + CO₂↑ (gas evolution)
- HCl + NaOH → NaCl + H₂O (neutralisation — a special case)
Definitions of Oxidation: 1. Addition of oxygen to a substance 2. Removal of hydrogen from a substance 3. Loss of electrons (most fundamental definition) 4. Increase in oxidation number
Definitions of Reduction: 1. Removal of oxygen from a substance 2. Addition of hydrogen to a substance 3. Gain of electrons 4. Decrease in oxidation number
Examples:
| Reaction | Oxidised | Reduced |
|---|---|---|
| 2Mg + O₂ → 2MgO | Mg (gains O) | O₂ (gains e⁻ from Mg) |
| CuO + H₂ → Cu + H₂O | H₂ (gains O) | CuO (loses O) |
| Fe + CuSO₄ → FeSO₄ + Cu | Fe (Fe → Fe²⁺, loses e⁻) | Cu²⁺ (gains e⁻ → Cu) |
Oxidising Agent: Substance that oxidises another substance. It itself gets reduced (gains electrons).
Reducing Agent: Substance that reduces another substance. It itself gets oxidised (loses electrons).
Precipitation reaction: A chemical reaction between two aqueous solutions in which an insoluble product (precipitate) is formed. The precipitate settles at the bottom or is seen as a cloudiness/turbidity.
Example 1: Lead nitrate + potassium iodide
Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂↓ + 2KNO₃(aq)
Observation: Bright yellow precipitate of lead iodide (PbI₂) forms.
Example 2: Barium chloride + sodium sulphate
BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄↓ + 2NaCl(aq)
Observation: White precipitate of barium sulphate forms. (This is the test for SO₄²⁻ ions)
Example 3: Silver nitrate + sodium chloride
AgNO₃(aq) + NaCl(aq) → AgCl↓ + NaNO₃(aq)
Observation: White precipitate of silver chloride forms. (This is the test for Cl⁻ ions)
Exothermic reactions — release heat (ΔH < 0, negative):
| Reaction | Observation |
|---|---|
| Combustion of fuels (coal, petrol) | Burns, produces heat and light |
| Neutralisation (HCl + NaOH) | Solution becomes warm |
| Slaking of lime (CaO + H₂O) | Vessel gets hot |
| Cellular respiration | Body warmth |
| Dissolving H₂SO₄ in water | Extreme heat — always add acid to water |
Endothermic reactions — absorb heat (ΔH > 0, positive):
| Reaction | Observation |
|---|---|
| Photosynthesis | Requires sunlight |
| Thermal decomposition of CaCO₃ | Requires continuous heating |
| Dissolving ammonium nitrate in water | Solution becomes cold |
| Evaporation of sweat | Skin cools |
| Cooking food | Requires heat |
How to identify: Measure temperature change of surroundings. Temperature increase → exothermic; Temperature decrease → endothermic.
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