What is a chemical reaction?
A chemical reaction is a process in which one or more substances (the reactants) change into new substances (the products) with different properties.
You can tell a chemical reaction has happened if you see any of these:
- Change in state
- Change in colour
- Evolution of a gas
- Change in temperature
- Formation of a precipitate
Example: When magnesium ribbon is burnt in air, it burns with a dazzling white flame and forms a white powder, magnesium oxide.
2Mg + O₂ → 2MgO
Chemical equations
A chemical equation uses symbols and formulae to show a reaction.
- Word equation: Magnesium + Oxygen → Magnesium oxide
- Skeletal equation: Mg + O₂ → MgO (not balanced)
- Balanced equation: 2Mg + O₂ → 2MgO
Why balance an equation?
The Law of Conservation of Mass says mass can neither be created nor destroyed in a chemical reaction. So the number of atoms of each element must be the same on both sides.
Steps to balance (hit-and-trial method)
- Write the skeletal equation.
- Count the atoms of each element on both sides.
- Start with the compound that has the most atoms. Balance its elements using coefficients, never by changing formulae.
- Balance the remaining elements one by one. Leave H and O for last.
- Check every element again.
- Add state symbols: (s), (l), (g), (aq).
Example: Fe(s) + H₂O(g) → Fe₃O₄(s) + H₂(g)
Balanced: 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)
Types of chemical reactions
| Type | What happens | Example |
|---|---|---|
| Combination | Two or more reactants form a single product | CaO + H₂O → Ca(OH)₂ |
| Decomposition | A single reactant breaks into simpler products | 2FeSO₄ → Fe₂O₃ + SO₂ + SO₃ |
| Displacement | A more reactive element displaces a less reactive one | Fe + CuSO₄ → FeSO₄ + Cu |
| Double displacement | Two compounds exchange ions | Na₂SO₄ + BaCl₂ → BaSO₄↓ + 2NaCl |
| Redox | Oxidation and reduction happen together | CuO + H₂ → Cu + H₂O |
Decomposition: three kinds
- Thermal (by heat): CaCO₃ → CaO + CO₂
- Electrolytic (by electricity): 2H₂O → 2H₂ + O₂
- Photolytic (by light): 2AgCl → 2Ag + Cl₂. This is why silver chloride is stored in dark bottles. It is also used in black-and-white photography.
Exothermic and endothermic reactions
- Exothermic: heat is released. Examples: burning of natural gas, respiration, and slaking of lime (CaO + H₂O).
- Endothermic: heat is absorbed. Example: decomposition of calcium carbonate.
Oxidation and reduction
- Oxidation: gain of oxygen or loss of hydrogen
- Reduction: loss of oxygen or gain of hydrogen
In CuO + H₂ → Cu + H₂O:
- CuO is reduced to Cu (it loses oxygen).
- H₂ is oxidised to H₂O (it gains oxygen).
- CuO is the oxidising agent. H₂ is the reducing agent.
Effects of oxidation in daily life
Corrosion
When a metal is attacked by moisture, air or acids, it slowly gets eaten away. Rusting of iron is the most common example. Silver turns black and copper turns green for the same reason.
Prevention: painting, oiling, greasing, galvanisation (coating with zinc) and chrome plating.
Rancidity
Fats and oils are oxidised when kept for a long time, and they develop a bad smell and taste.
Prevention: adding antioxidants, storing in air-tight containers, and flushing chip packets with nitrogen gas.
Quick revision: important questions
- Why should a magnesium ribbon be cleaned before burning? (To remove the protective layer of magnesium oxide.)
- Why is respiration considered an exothermic reaction? (Glucose is oxidised and energy is released.)
- Why does the colour of copper sulphate solution change when an iron nail is dipped in it? (Iron displaces copper, so the solution turns from blue to light green, which is FeSO₄.)
- What is a precipitation reaction? Give an example. (A reaction that forms an insoluble solid, for example BaSO₄.)
- Define the oxidising and reducing agent in ZnO + C → Zn + CO.
Exam tip: In MP Board papers, “balance the equation” and “identify the type of reaction” questions appear almost every year. Practise at least 10 equations.