Electrolytes and ionisation
An electrolyte is a substance that conducts electricity when melted or dissolved in water, because it forms free ions (charged particles). Breaking into ions in water is called dissociation or ionisation.
- Strong electrolytes split almost completely: soluble salts (NaCl), strong acids (HCl, HNO₃, H₂SO₄) and strong alkalis (NaOH, KOH). NaCl → Na⁺ + Cl⁻.
- Weak electrolytes split only a little: weak acids like vinegar (CH₃COOH), and ammonia solution. We write them as molecules.
- Non-electrolytes do not form ions: sugar, alcohol.
Acids give H⁺ ions in water; alkalis give OH⁻ ions; salts give metal ions and acid-radical ions.
Full and net ionic equations
Take silver nitrate + sodium chloride.
1. Molecular equation: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)
2. Full ionic equation: split every strong electrolyte that is dissolved (aq) into its ions. Keep solids, gases, water and weak electrolytes as formulas.
Ag⁺ + NO₃⁻ + Na⁺ + Cl⁻ → AgCl(s) + Na⁺ + NO₃⁻
3. Cross out spectator ions (the same on both sides): Na⁺ and NO₃⁻.
4. Net (short) ionic equation: Ag⁺ + Cl⁻ → AgCl↓
5. Check: atoms balance and total charge is the same on both sides (here 0 = 0).
The net equation shows the real change. Any soluble silver salt with any soluble chloride gives the same net equation.
When does an ion-exchange reaction happen?
An ion-exchange reaction goes to completion only if ions are removed from the solution as:
- a precipitate (an insoluble solid): Ba²⁺ + SO₄²⁻ → BaSO₄↓ (white); Cu²⁺ + 2OH⁻ → Cu(OH)₂↓ (blue)
- a gas: CO₃²⁻ + 2H⁺ → H₂O + CO₂↑; NH₄⁺ + OH⁻ → NH₃↑ + H₂O (on warming)
- water (or another weak electrolyte): H⁺ + OH⁻ → H₂O. This is the net equation for every strong acid + strong alkali neutralisation.
If none of these forms, for example KCl + NaNO₃, all ions are spectators and there is no reaction.
Ion coexistence
Ions can stay together in one solution only if they do not form a precipitate, gas or water. Ba²⁺ and SO₄²⁻ cannot coexist; H⁺ and OH⁻ cannot; H⁺ and CO₃²⁻ cannot. Na⁺, K⁺, NO₃⁻ and Cl⁻ can coexist with almost anything. Use a solubility table to check.
Tests for common ions (qualitative reactions)
- Chloride Cl⁻: add dilute nitric acid, then silver nitrate → white precipitate. Ag⁺ + Cl⁻ → AgCl↓
- Sulfate SO₄²⁻: add dilute hydrochloric acid, then barium chloride → white precipitate. Ba²⁺ + SO₄²⁻ → BaSO₄↓
- Carbonate CO₃²⁻: add dilute acid → fizzing; the gas turns limewater milky. CO₃²⁻ + 2H⁺ → H₂O + CO₂↑
- Ammonium NH₄⁺: warm with sodium hydroxide → gas that turns damp red litmus blue. NH₄⁺ + OH⁻ → NH₃↑ + H₂O
- Metal ions with NaOH: Cu²⁺ blue precipitate, Fe²⁺ green, Fe³⁺ red-brown. Fe³⁺ + 3OH⁻ → Fe(OH)₃↓
Hydrolysis, amphoteric compounds and water treatment
Hydrolysis (introduction)
Some salt ions react a little with water itself. Sodium carbonate solution is alkaline because CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻. Ammonium chloride solution is slightly acidic because NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺. A salt of a strong acid and a strong alkali (NaCl) gives a neutral solution.
Amphoteric compounds (idea)
Some hydroxides, like Zn(OH)₂ and Al(OH)₃, react with both acids and alkalis. Zn(OH)₂ + 2H⁺ → Zn²⁺ + 2H₂O, and with extra OH⁻ it dissolves again.
Ion exchange for water
Hard water contains Ca²⁺ and Mg²⁺. An ion-exchange resin holds Na⁺ ions and swaps them: Ca²⁺ goes onto the resin, Na⁺ goes into the water. Mixed resins can swap all ions for H⁺ and OH⁻ to give very pure (deionised) water.
Try it at home
Put a spoon of baking soda in a glass and add vinegar. The fizz is CO₂ gas leaving the solution: HCO₃⁻ + CH₃COOH → CH₃COO⁻ + H₂O + CO₂↑ (vinegar is a weak acid, so we keep it as a molecule). Then in the 3D, choose "Carbonate + acid" and watch the bubbles leave the beaker.
Key formulas and definitions
- Ag⁺ + Cl⁻ → AgCl↓ (white)
- Ba²⁺ + SO₄²⁻ → BaSO₄↓ (white)
- Cu²⁺ + 2OH⁻ → Cu(OH)₂↓ (blue)
- Fe³⁺ + 3OH⁻ → Fe(OH)₃↓ (red-brown)
- H⁺ + OH⁻ → H₂O
- CO₃²⁻ + 2H⁺ → H₂O + CO₂↑
- NH₄⁺ + OH⁻ → NH₃↑ + H₂O
- Rule: split only dissolved strong electrolytes; keep (s), (g), H₂O and weak electrolytes whole
Worked examples
1. Write the full and net ionic equations for barium chloride + sodium sulfate.
Molecular: BaCl₂ + Na₂SO₄ → BaSO₄↓ + 2NaCl. Full ionic: Ba²⁺ + 2Cl⁻ + 2Na⁺ + SO₄²⁻ → BaSO₄↓ + 2Na⁺ + 2Cl⁻. Spectators: Na⁺, Cl⁻. Net: Ba²⁺ + SO₄²⁻ → BaSO₄↓.
2. Write the net ionic equation for hydrochloric acid + sodium hydroxide.
Full: H⁺ + Cl⁻ + Na⁺ + OH⁻ → Na⁺ + Cl⁻ + H₂O. Spectators Na⁺, Cl⁻. Net: H⁺ + OH⁻ → H₂O.
3. Copper(II) sulfate solution is mixed with sodium hydroxide. Write the net ionic equation and say what you see.
Cu²⁺ + 2OH⁻ → Cu(OH)₂↓. A blue jelly-like precipitate forms. Na⁺ and SO₄²⁻ are spectators.
4. Write the net ionic equation for sodium carbonate + hydrochloric acid.
Full: 2Na⁺ + CO₃²⁻ + 2H⁺ + 2Cl⁻ → 2Na⁺ + 2Cl⁻ + H₂O + CO₂↑. Net: CO₃²⁻ + 2H⁺ → H₂O + CO₂↑. Check charge: (−2) + (+2) = 0 on the left, 0 on the right.
5. Will a reaction happen when potassium nitrate and sodium chloride solutions are mixed?
Ions present: K⁺, NO₃⁻, Na⁺, Cl⁻. Possible new pairs KCl and NaNO₃ are both soluble; no gas, no water. All ions are spectators: no reaction.
6. Write the molecular equation that matches the net equation Fe³⁺ + 3OH⁻ → Fe(OH)₃↓.
Choose a soluble iron(III) salt and a soluble hydroxide, e.g. FeCl₃ + 3NaOH → Fe(OH)₃↓ + 3NaCl. Full ionic: Fe³⁺ + 3Cl⁻ + 3Na⁺ + 3OH⁻ → Fe(OH)₃↓ + 3Na⁺ + 3Cl⁻.
7. Which pair of ions cannot be present together in a solution: (a) K⁺ and SO₄²⁻, (b) Ag⁺ and Cl⁻, (c) Na⁺ and NO₃⁻?
(b) Ag⁺ and Cl⁻, because they form the insoluble precipitate AgCl. The others form soluble salts and can coexist.
Common mistakes
- Splitting a precipitate, gas or water into ions. Write AgCl(s), CO₂(g) and H₂O as whole formulas.
- Splitting a weak acid like CH₃COOH into ions in the equation. Weak electrolytes stay as molecules.
- Forgetting to balance charge: Cu²⁺ needs 2OH⁻, not one.
- Cancelling ions that are not exactly the same on both sides (e.g. cancelling Cl⁻ when it ends up in a precipitate).