pH of Strong Acids & Bases SL
Strong acids and bases fully dissociate, so the initial concentration directly gives you [H⁺] or [OH⁻].
Key Equations
- pH = −log[H⁺] | pOH = −log[OH⁻] | pH + pOH = 14.00 (at 25°C)
Worked Example: Strong Acid
Calculate the pH of 0.10 M HCl.
HCl fully ionises → [H⁺] = 0.10 M
pH = −log(0.10) = 1.00
Worked Example: Strong Base
Calculate the pH of 0.0044 M Ca(OH)₂.
Ca(OH)₂ gives 2 OH⁻ per unit → [OH⁻] = 0.0044 × 2 = 0.0088 M
pOH = −log(0.0088) = 2.06
pH = 14.00 − 2.06 = 11.94
HLpH of Weak Acids (Using Ka)
Weak acids partially dissociate. Use an ICE table to find [H⁺].
Worked Example: Weak Acid
Find the pH of 0.534 M formic acid (HCO₂H), Ka = 1.8 × 10⁻⁴
Ka = x² / (0.534 − x) ≈ x² / 0.534 (assuming x ≪ 0.534)
x² = 1.8 × 10⁻⁴ × 0.534 = 9.6 × 10⁻⁵
x = [H⁺] = 0.0098 M
pH = −log(0.0098) = 2.01
HLpH of Weak Bases (Using Kb)
Worked Example: Weak Base
Find the pH of 0.0325 M NH₃, Kb = 1.76 × 10⁻⁵
Kb = x² / 0.0325 → x = √(0.0325 × 1.76 × 10⁻⁵) = 7.56 × 10⁻⁴ M
pOH = −log(7.56 × 10⁻⁴) = 3.12
pH = 14.00 − 3.12 = 10.88
Titration Curves
Strong Acid + Strong Base Titration
HLAll Four Titration Curve Types
| Combination | Initial pH | Equiv. PH | Key Features |
|---|---|---|---|
| SA + SB | ~1 | 7 | Steep symmetric jump; no buffer region |
| WA + SB | ~3–4 | > 7 | Buffer region before equiv.; half-equiv: pH = pKa |
| SA + WB | ~11 | < 7 | Buffer region; half-equiv: pOH = pKb |
| WA + WB | Moderate | Varies | Very shallow curve; no steep section; pH meter needed |
HLChoosing an Indicator
An indicator is a weak acid (HInd) whose conjugate base (Ind⁻) has a different colour. It changes colour when pH ≈ pKa of the indicator (±1).
The Rule
The indicator's colour-change range must fall within the steep vertical section of the titration curve.
- Phenolphthalein (range 8.3–10.0) → good for WA+SB and SA+SB
- Methyl orange (range 3.1–4.4) → good for SA+WB and SA+SB
- WA+WB → no indicator works (no steep section). Use a pH meter
HLSalt Hydrolysis
When a salt dissolves, its ions may react with water (hydrolyse). The pH depends on the parent acid and base:
| Salt Type | Made From | Solution pH | Example |
|---|---|---|---|
| Neutral | SA + SB | = 7 | NaCl |
| Basic | WA + SB | > 7 | CH₃COONa |
| Acidic | SA + WB | < 7 | NH₄Cl |
HLBuffer Solutions
A buffer resists changes in pH when small amounts of acid or base are added. Made from comparable amounts of a conjugate pair.
Types of Buffer
- Acidic buffer: weak acid + its conjugate base (e.g. CH₃COOH + CH₃COO⁻)
- Basic buffer: weak base + its conjugate acid (e.g. NH₃ + NH₄⁺)
How It Works
- If H⁺ added → absorbed by the conjugate base (A⁻ + H⁺ → HA)
- If OH⁻ added → neutralised by the weak acid (HA + OH⁻ → A⁻ + H₂O)
Henderson-Hasselbalch Equation
pH = pKa + log([A⁻] / [HA])
At the half-equivalence point, [A⁻] = [HA], so log(1) = 0 → pH = pKa. This is how you experimentally determine Ka from a titration curve.