The three acid–base theories, the pH scale, acid/base strength (Ka, Kb), neutralisation, salts & hydrolysis, and buffers — exam-focused for the BIEK / Sindh Board paper. Read it, or open the interactive lecture and run the pH / titration simulator.
1 — Properties of acids & bases
| Acids | Bases |
|---|
| sour taste; turn litmus red | bitter taste, soapy feel; turn litmus blue |
| release H⁺ (H₃O⁺) in water | release OH⁻ in water |
| react with metals → H₂ gas | react with acids → salt + water |
| react with carbonates → CO₂ | conduct electricity (aqueous) |
2 — Theories of acids & bases
| Theory | Acid | Base |
|---|
| Arrhenius | gives H⁺ in water | gives OH⁻ in water |
| Brønsted–Lowry | proton (H⁺) donor | proton acceptor |
| Lewis | electron-pair acceptor | electron-pair donor |
Lewis is the broadest — e.g. BF₃ (no H) is a Lewis acid; NH₃ (lone pair) is a Lewis base.
3 — Conjugate acid–base pairs
In the Brønsted theory, when an acid donates a proton it becomes its conjugate base; when a base accepts a proton it becomes its conjugate acid. They differ by one H⁺.
ExampleHCl + H₂O → H₃O⁺ + Cl⁻
acid₁ base₂ acid₂ base₁ (HCl/Cl⁻ and H₂O/H₃O⁺ are the pairs)
A strong acid has a weak conjugate base, and vice versa.
4 — Strong vs weak acids & bases
- Strong acid/base — almost completely ionised in water (HCl, H₂SO₄, HNO₃; NaOH, KOH).
- Weak acid/base — only partially ionised (CH₃COOH, H₂CO₃; NH₃).
Strength ≠ concentration. A dilute strong acid can have a lower H⁺ than a concentrated weak acid. Strength is the degree of ionisation.
5 — Ionic product of water (Kw)
Water itself ionises slightly: H₂O ⇌ H⁺ + OH⁻.
Kw at 25 °CKw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴
in pure water [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ mol/dm³
6 — The pH & pOH scale
DefinitionspH = −log[H⁺] · pOH = −log[OH⁻]
pH + pOH = 14 (at 25 °C)
| pH | Nature |
|---|
| < 7 | acidic ([H⁺] > [OH⁻]) |
| = 7 | neutral |
| > 7 | basic / alkaline ([OH⁻] > [H⁺]) |
pH from [H⁺]
Find the pH of 0.001 M HCl.
[H⁺] = 10⁻³ → pH = −log(10⁻³) = 3
7 — Ka, Kb and pKa
For a weak acid HA ⇌ H⁺ + A⁻Ka = [H⁺][A⁻] / [HA] · pKa = −log Ka
smaller Ka (larger pKa) → weaker acid
Ka × Kb = Kw for a conjugate acid–base pair. A weak base has Kb similarly defined.
8 — Neutralisation & salts
- Neutralisation — acid + base → salt + water. The net ionic equation is H⁺ + OH⁻ → H₂O (ΔH ≈ −57 kJ/mol).
- Salt — an ionic compound formed when the H⁺ of an acid is replaced by a metal (or NH₄⁺) ion.
ExampleHCl + NaOH → NaCl + H₂O
9 — Salt hydrolysis
The ions of a salt can react with water, making the solution acidic or basic.
| Salt from | Example | Solution |
|---|
| strong acid + strong base | NaCl | neutral (pH 7) |
| strong acid + weak base | NH₄Cl | acidic (pH < 7) |
| weak acid + strong base | CH₃COONa | basic (pH > 7) |
10 — Buffer solutions
- Buffer — a solution that resists a change in pH when small amounts of acid or base are added. Made from a weak acid + its salt (acidic buffer) or a weak base + its salt (basic buffer).
Henderson–HasselbalchpH = pKa + log( [salt] / [acid] )
Blood is buffered near pH 7.4 by the H₂CO₃ / HCO₃⁻ system.
11 — Indicators & titration
- Indicator — a weak acid/base whose colour depends on pH (litmus, phenolphthalein, methyl orange).
In a titration, a measured volume of one solution is added to another until the end point. At the equivalence point, moles of acid = moles of base.
titration
25 cm³ of HCl is neutralised by 20 cm³ of 0.1 M NaOH. Find the HCl concentration.
moles NaOH = 0.1 × 20/1000 = 0.002 = moles HCl
[HCl] = 0.002 / (25/1000) = 0.08 M
12 — Exam recap
- Properties of acids & bases; the three theories (Arrhenius, Brønsted, Lewis).
- Conjugate acid–base pairs; strong vs weak (degree of ionisation).
- Kw; pH = −log[H⁺]; pH + pOH = 14.
- Ka, Kb, pKa; Ka × Kb = Kw.
- Neutralisation, salts & hydrolysis; buffers (Henderson–Hasselbalch).
- Indicators & titration calculations.