How do you know a reaction has happened?
In a chemical reaction, substances (the reactants) change into new substances with different properties (the products). You can usually tell a reaction has occurred from one or more of these observations:
a change in colour,
a change in temperature (heat released or absorbed),
evolution of a gas (bubbles, fizzing),
formation of a precipitate (an insoluble solid),
a change of state.
Example. When a clean magnesium ribbon is burnt in air, it burns with a dazzling white flame and leaves a white powder: magnesium oxide.
Writing a chemical equation
A word equation names the substances:
magnesium + oxygen → magnesium oxide
A chemical equation uses formulae instead: Mg + O₂ → MgO. This is a skeletal equation. It is not balanced yet, because there are two oxygen atoms on the left and only one on the right.
States are shown in brackets: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution in water.
Balancing an equation
Atoms are neither created nor destroyed in a chemical reaction. This is the law of conservation of mass. So the number of atoms of each element must be the same on both sides.
To balance, change only the coefficients (the numbers in front of formulae). Never change the subscripts inside a formula, because that would change the substance itself.
Example. Balance Fe + H₂O → Fe₃O₄ + H₂.
Iron: 3 on the right, so write 3Fe.
Oxygen: 4 on the right, so write 4H₂O.
Hydrogen: now 8 on the left, so write 4H₂.
3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)
Check: Fe 3 = 3, O 4 = 4, H 8 = 8. ✓
Types of chemical reactions
Type | General pattern | Example |
|---|---|---|
Combination | A + B → AB | CaO + H₂O → Ca(OH)₂ |
Decomposition | AB → A + B | CaCO₃ → CaO + CO₂ (on heating) |
Displacement | A + BC → AC + B | Fe + CuSO₄ → FeSO₄ + Cu |
Double displacement | AB + CD → AD + CB | Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl |
Combination. Quicklime (CaO) reacts vigorously with water to form slaked lime, Ca(OH)₂, and releases a lot of heat. Reactions that release heat are exothermic. Respiration is also exothermic.
Decomposition. One substance breaks into simpler ones. It needs energy, supplied as heat (thermal decomposition), light (silver chloride turning grey in sunlight) or electricity (electrolysis of water). Reactions that absorb energy are endothermic.
Displacement. A more reactive metal pushes a less reactive metal out of its compound. An iron nail placed in blue copper sulphate solution becomes coated with brown copper, and the solution turns pale green as iron sulphate forms.
Double displacement. Two compounds exchange ions. When sodium sulphate and barium chloride solutions are mixed, a white precipitate of barium sulphate forms. Such reactions are also called precipitation reactions.
Oxidation and reduction (redox)
Oxidation: gain of oxygen or loss of hydrogen.
Reduction: loss of oxygen or gain of hydrogen.
In CuO + H₂ → Cu + H₂O, copper oxide loses oxygen (it is reduced) and hydrogen gains oxygen (it is oxidised). When both happen in the same reaction it is a redox reaction. The substance that gets oxidised (here H₂) is the reducing agent.
Oxidation in everyday life
Corrosion: metals are slowly attacked by air and moisture. Rusting of iron is the best-known example. It is prevented by painting, oiling, galvanising (coating with zinc) or chrome plating.
Rancidity: fats and oils in food are oxidised and develop a bad smell and taste. It is slowed by adding antioxidants, by refrigeration, and by packing chips in nitrogen instead of air.
Quick summary
Balance equations by changing coefficients, never subscripts.
Combination joins, decomposition splits, displacement swaps one element, double displacement swaps ions.
Oxidation: gain of O or loss of H. Reduction: the reverse. Both together form a redox reaction.