Chemical reactions are very varied, so they are classified in several ways: by the number and composition of reactants and products, heat change, reversibility, change in oxidation states, phases and the involvement of a catalyst. Tests often ask you to give one reaction several "labels". For example, Zn + 2HCl → ZnCl₂ + H₂ is at the same time a displacement, redox, irreversible, heterogeneous and exothermic reaction. This topic organises these classifications and shows how to balance equations correctly and do thermochemical calculations. It follows on from "Basic concepts" and introduces "Redox reactions", "Electrolytic dissociation" and "Reaction rate".
§1. By number and composition of substances
Combination (synthesis) — two or more substances form one compound: 2H₂ + O₂ → 2H₂O, CaO + H₂O → Ca(OH)₂, NH₃ + HCl → NH₄Cl.
Decomposition — one compound gives two or more substances: CaCO₃ → CaO + CO₂ (heat), 2KNO₃ → 2KNO₂ + O₂, 2KMnO₄ → K₂MnO₄ + MnO₂ + O₂.
Displacement (single replacement) — atoms of an element replace atoms of one element in a compound: Fe + CuSO₄ → FeSO₄ + Cu, Zn + 2HCl → ZnCl₂ + H₂.
Exchange (double replacement) — two compounds swap parts (ions): NaOH + HCl → NaCl + H₂O.
An exchange reaction in solution goes to completion if one product is a precipitate, a gas or a weakly dissociated substance (water, a weak acid) — Berthollet’s rule. Exchange between an acid and a base is neutralisation. One reaction can be viewed in more than one way: for example, CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ is an exchange reaction because the unstable intermediate H₂CO₃ breaks down into water and gas.
Worked example
Solutions of BaCl₂ + Na₂SO₄ and of KCl + NaNO₃ are mixed. In which pair does a reaction occur?
BaCl₂ + Na₂SO₄ → BaSO₄↓ + 2NaCl — exchange; a BaSO₄ precipitate forms, so the reaction goes to completion.
KCl + NaNO₃ — the possible products KNO₃ and NaCl are soluble; there is no precipitate, gas or water.
Answer: only the first pair reacts.
§2. By heat change
Exothermic (+Q, ΔH < 0) — heat is released; endothermic (−Q, ΔH > 0) — heat is absorbed
Combustion, neutralisation and most combination reactions are exothermic; most decompositions, N₂ + O₂ → 2NO and photosynthesis are endothermic. The heat change is the difference between the energy needed to break bonds and the energy released when bonds form. A thermochemical equation shows the heat change: C + O₂ → CO₂ + 394 kJ. The heat refers to exactly the amounts written, so calculations use a simple proportion.
Worked example
2H₂ + O₂ → 2H₂O + 572 kJ. How much heat is released when 11.2 L of H₂ (STP) burns?
n(H₂) = 11.2/22.4 = 0.5 mol.
From the equation, 2 mol of H₂ → 572 kJ.
Q = 572·0.5/2 = 143 kJ.
§3. By reversibility and change in oxidation state
Irreversible reactions go to completion (a precipitate, gas or water forms, or combustion); reversible reactions go both ways at once and reach equilibrium (⇌): N₂ + 3H₂ ⇌ 2NH₃, esterification, 2SO₂ + O₂ ⇌ 2SO₃.
In redox reactions oxidation states change: all displacement reactions, combination and decomposition reactions involving elements, and combustion. Exchange reactions are never redox. Combination and decomposition reactions without elements (CaO + H₂O, CaCO₃ → CaO + CO₂) usually do not change oxidation states, but there are exceptions: in NH₄NO₃ → N₂O + 2H₂O nitrogen goes from −3 and +5 to +1. So the safest approach is always to assign oxidation states.
Worked example
Classify the reaction 2KMnO₄ → K₂MnO₄ + MnO₂ + O₂.
One substance gives three — decomposition.
Mn⁺⁷ → Mn⁺⁶ and Mn⁺⁴ (reduced), O⁻² → O⁰ (oxidised) — a redox reaction (intramolecular).
It needs heating and absorbs heat — endothermic and irreversible.
§4. By other features
By phase: homogeneous (in one phase — a gas mixture or a solution: H₂ + Cl₂ → 2HCl) and heterogeneous (at a boundary between phases: Zn(s) + HCl(aq), burning coal).
By catalyst: catalysed (2SO₂ + O₂ ⇌ 2SO₃, V₂O₅; enzyme reactions) and uncatalysed.
By energy source: photochemical (light: photosynthesis, the H₂ + Cl₂ explosion), electrochemical (electrolysis), thermal (heating).
In organic chemistry: addition (hydrogenation, halogenation), substitution, elimination (dehydration), polymerisation, isomerisation; by mechanism — radical and ionic.
§5. Balancing equations
Write the formulae correctly; when balancing never change subscripts — only add coefficients.
Balance metals first, then non-metals (treat acid residues as groups), then hydrogen and finally oxygen.
Check that each element has the same number of atoms on both sides; the coefficients should be the smallest whole numbers.
For combustion of organic compounds, balance carbon first, then hydrogen and finally oxygen; if oxygen comes out odd, double all the coefficients.
Worked example
Balance Fe₂O₃ + H₂SO₄ → Fe₂(SO₄)₃ + H₂O.
Iron: 2 = 2. Three sulfate groups on the right → 3H₂SO₄.
External signs show that a reaction is happening: a colour change, gas bubbles, a precipitate forming or dissolving, a new smell, heat or light given out. Reactions often need conditions to start: contact between substances (grinding, dissolving), heating (to start combustion), light, a catalyst, pressure. In an equation, conditions are written above the arrow, a precipitate is marked ↓ and a gas ↑. Some reactions keep going once started (combustion — the heat released warms the next molecules), while others need a continuous supply of energy (decomposition of CaCO₃, electrolysis). Unlike physical changes, chemical reactions produce new substances, but the number of atoms is conserved — which is why equations must be balanced.
Check yourself
What type of reaction is CuO + H₂ → Cu + H₂O?
Which pair reacts to completion: a) NaCl + KNO₃; b) Na₂CO₃ + HCl?
Is CaCO₃ → CaO + CO₂ − 178 kJ exothermic or endothermic?
Balance Al + O₂ → Al₂O₃.
C + O₂ → CO₂ + 394 kJ. What mass of carbon must burn to give 1970 kJ?
What is the net ionic equation for neutralisation?
Balance C₃H₈ + O₂ → CO₂ + H₂O and find the sum of the coefficients.
Is NH₄Cl → NH₃ + HCl a redox reaction?
Common mistakes
Changing subscripts when balancing (for example turning H₂O into H₂O₂).
Assuming every exchange reaction happens: check Berthollet’s rule.
Calling exchange reactions redox reactions.
Ignoring the equation’s coefficient in a thermochemical calculation.
Treating every combination and decomposition reaction as redox: assign oxidation states and check.