What counts as a chemical reaction
A chemical reaction makes new substances by breaking bonds and forming new ones. A physical change (melting, dissolving sugar) changes only the state or mixture, not the substances.
Signs a reaction happened: a gas forms, a solid (precipitate) appears in a solution, the colour changes, heat or light is given out or taken in, or a smell appears. None of these alone is proof (boiling also gives bubbles), so chemists check whether a new substance with new properties formed.
Law of conservation of mass: atoms are not created or destroyed, only rearranged, so the total mass of reactants equals the total mass of products in a closed system.
Every reaction also has an energy change: exothermic reactions release heat (ΔH < 0), endothermic reactions absorb it (ΔH > 0).
Representations of reactions: from words to net ionic equations
One reaction, four ways to write it:
- Word equation: silver nitrate + sodium chloride → silver chloride + sodium nitrate.
- Balanced formula equation with state symbols (s, l, g, aq): AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq).
- Full (complete) ionic equation: write dissolved strong electrolytes as ions: Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq).
- Net ionic equation: cross out ions that appear unchanged on both sides (spectator ions): Ag⁺(aq) + Cl⁻(aq) → AgCl(s).
Rules: keep solids, liquids, gases, weak acids and water as whole formulas; split only soluble ionic compounds and strong acids/bases. A net ionic equation must balance in atoms and in charge.
Particle diagrams
Chemists also draw reactions as particle pictures, like the 3D scene. A correct diagram shows the same number of each kind of atom before and after.
Balancing by inspection
Change only the big numbers in front (coefficients), never the small subscripts. Balance elements that appear in one substance on each side first, then H and O, and finish with the smallest whole numbers.
Types of chemical reactions
| Type | Pattern | Example |
|---|---|---|
| Synthesis (combination) | A + B → AB | 2Mg + O₂ → 2MgO |
| Decomposition | AB → A + B | CaCO₃ → CaO + CO₂ (heat) |
| Single displacement | A + BC → AC + B | Zn + CuSO₄ → ZnSO₄ + Cu |
| Double displacement | AB + CD → AD + CB | AgNO₃ + NaCl → AgCl↓ + NaNO₃ |
| Combustion | fuel + O₂ → CO₂ + H₂O | CH₄ + 2O₂ → CO₂ + 2H₂O |
Double displacement includes precipitation (an insoluble solid forms), neutralisation (acid + base → salt + water; net H⁺ + OH⁻ → H₂O) and gas-forming reactions (carbonate + acid → CO₂).
Sorting by what is transferred
- Redox: electrons move. The species that loses electrons is oxidised (oxidation number goes up); the one that gains is reduced. Synthesis, combustion and single displacement are almost always redox.
- Acid–base (Brønsted): a proton (H⁺) moves from acid to base.
- Precipitation: ions join into an insoluble solid; no electrons move.
A reactivity series tells you whether single displacement happens: a metal displaces a less reactive metal from its salt (Zn pushes out Cu, but Cu cannot push out Zn).
Predicting simple reactions from outer electrons
Atoms react to reach a stable outer shell, usually 8 electrons (2 for H and He). The number of outer (valence) electrons tells you what an element will do:
- Group 1 metals (1 outer e⁻) lose 1 → +1 ions; Group 2 lose 2 → +2; aluminium loses 3 → +3.
- Group 17 (7 outer e⁻) gain 1 → −1; Group 16 gain 2 → −2; Group 15 gain 3 → −3.
- Two non-metals share electrons instead (covalent bonds), e.g. H₂O, CH₄.
Formula rule: electrons lost by the metal must equal electrons gained by the non-metal. For Mg and Cl: Mg gives 2, each Cl takes 1, so MgCl₂. For Al and O: lowest common multiple of 3 and 2 is 6, so 2 Al and 3 O: Al₂O₃.
You can also predict how vigorously: going down Group 1, the outer electron is farther from the nucleus and lost more easily, so K reacts faster with water than Na. Going down Group 17, gaining an electron gets harder, so F₂ is the most reactive halogen, and a more reactive halogen displaces a less reactive one: Cl₂ + 2KBr → 2KCl + Br₂.
Coupled reactions
Whether a reaction can go by itself depends on the Gibbs free energy change, ΔG. If ΔG < 0 the reaction is spontaneous (thermodynamically favourable); if ΔG > 0 it will not go on its own.
A non-spontaneous reaction can be driven by coupling it to a strongly favourable reaction that shares an intermediate. Add the equations and add their ΔG values: if the total is negative, the overall process goes.
Example (metal extraction): Fe₂O₃ → 2Fe + 3/2 O₂ has a large positive ΔG. Burning carbon, 3/2 C + 3/2 O₂ → 3/2 CO₂, has a large negative ΔG. Together, Fe₂O₃ + 3/2 C → 2Fe + 3/2 CO₂ has a negative ΔG at furnace temperature.
Example (living cells): joining glucose and fructose has ΔG ≈ +27 kJ/mol; ATP hydrolysis has ΔG ≈ −30 kJ/mol. Coupled, the total is about −3 kJ/mol, so the cell can build the sugar.
Electrochemical cells use the same idea: a favourable redox reaction can push an unfavourable one (as in charging a battery from an outside source).
Try it: a safe kitchen precipitation
Dissolve a spoon of washing soda (sodium carbonate) in one cup of water and a spoon of Epsom salt or hard tap water in another. Pour them together: a white cloudiness appears. Predict first: which ions are spectators? (Na⁺ and SO₄²⁻.) The net ionic equation is Mg²⁺(aq) + CO₃²⁻(aq) → MgCO₃(s). Then use the free-play step to predict formulas such as CaO and AlF₃.
Key formulas and definitions
- Mass of reactants = mass of products (closed system)
- Net ionic equation = full ionic equation − spectator ions
- Net ionic equation must balance atoms and total charge
- Oxidation = loss of electrons; Reduction = gain of electrons (OIL RIG)
- Ionic formula: (electrons lost by metal) × (number of metal atoms) = (electrons gained per non-metal) × (number of non-metal atoms)
- Neutralisation (net): H⁺(aq) + OH⁻(aq) → H₂O(l)
- Spontaneous if ΔG < 0; coupled reactions: ΔG(total) = ΔG₁ + ΔG₂
- Combustion of hydrocarbon: CxHy + O₂ → CO₂ + H₂O
Worked examples
1. Balance: Al + O₂ → Al₂O₃.
O: 2 on left, 3 on right; LCM 6, so 3 O₂ and 2 Al₂O₃. Then Al: 4 on right, so 4 Al. Answer: 4Al + 3O₂ → 2Al₂O₃.
2. Write the net ionic equation for BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq).
Full ionic: Ba²⁺ + 2Cl⁻ + 2Na⁺ + SO₄²⁻ → BaSO₄(s) + 2Na⁺ + 2Cl⁻. Cross out Na⁺ and Cl⁻: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s). Charge: +2 − 2 = 0 on both sides.
3. Classify: (a) 2H₂O₂ → 2H₂O + O₂ (b) HCl + NaOH → NaCl + H₂O (c) C₃H₈ + 5O₂ → 3CO₂ + 4H₂O.
(a) decomposition (also redox), (b) double displacement / neutralisation (acid–base), (c) combustion (redox).
4. Predict the formula of the compound of calcium (2 outer e⁻) and nitrogen (needs 3).
LCM of 2 and 3 is 6. Ca: 6 ÷ 2 = 3 atoms; N: 6 ÷ 3 = 2 atoms. Formula Ca₃N₂.
5. Will Cu + ZnSO₄ react? Will Mg + CuSO₄ react?
Cu is less reactive than Zn, so no reaction. Mg is more reactive than Cu, so Mg + CuSO₄ → MgSO₄ + Cu (Mg oxidised, Cu²⁺ reduced).
6. Reaction X has ΔG = +18 kJ/mol. Reaction Y has ΔG = −31 kJ/mol. They are coupled. Does the overall process go?
ΔG total = +18 + (−31) = −13 kJ/mol. It is negative, so yes, the coupled process is spontaneous.
Common mistakes
- Changing subscripts to balance (writing H₂O₂ instead of 2H₂O). Only coefficients may change.
- Splitting solids, gases or water into ions in an ionic equation. Only dissolved strong electrolytes are split.
- Forgetting to balance charge in a net ionic equation, e.g. Fe³⁺ + Cu → Fe²⁺ + Cu²⁺ looks balanced in atoms but not charge (needs 2Fe³⁺).
- Thinking a coupled reaction makes an unfavourable step favourable on its own. Only the total ΔG is negative; the shared intermediate links them.