Strong vs Weak: The Ionisation Concept
The strength of an acid or base refers to its extent of ionisation (dissociation) in aqueous solution. not its concentration.
Comparison Table
| Property | Strong | Weak |
|---|---|---|
| Ionisation | Complete (~100%) → use → | Partial (<5%) → use ⇌ |
| pH (same conc) | Lower (more acidic) | Higher |
| Conductivity | Higher (more ions) | Lower |
| Rate of reaction | Faster | Slower |
IB Required Lists
Memorise these. Any acid/base not on this list should be treated as weak.
Strong Acids
HCl, HBr, HI, HNO₃, H₂SO₄ (first proton only)
Strong Bases
Group 1 hydroxides: LiOH, NaOH, KOH, RbOH, CsOH
Heavy Group 2 hydroxides: Ca(OH)₂, Sr(OH)₂, Ba(OH)₂
⚠️ Exam Trap: Strength ≠ Concentration
Strength = proportion of molecules that ionise (100% vs partial).
Concentration = amount of solute per unit volume (mol dm⁻³).
It is entirely possible to have a dilute strong acid (e.g. 0.0001 M HCl) and a concentrated weak acid (e.g. 5.0 M CH₃COOH). A dilute strong acid can have a lower pH than a concentrated weak acid because it fully ionises.
The pH Scale
Key Equations
- pH = −log₁₀[H⁺] and [H⁺] = 10⁻ᵖᴴ
- pOH = −log₁₀[OH⁻] and [OH⁻] = 10⁻ᵖᴼᴴ
- At 25°C: pH + pOH = 14.00
Self-Ionisation of Water & Kw
Water is amphiprotic and undergoes auto-ionisation:
H₂O(l) ⇌ H⁺(aq) + OH⁻(aq)
The equilibrium constant for this is the ionic product of water:
Kw = [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ at 25°C
In pure water: [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ mol dm⁻³ → pH = 7.00
HLTemperature Dependence of Kw
Auto-ionisation of water is endothermic (ΔH > 0). By Le Chatelier's principle, increasing temperature shifts equilibrium to the right → more H⁺ and OH⁻ → Kw increases → pH of pure water decreases.
Key Insight
Even though pH drops below 7 at higher temperatures, the water remains perfectly neutral because [H⁺] = [OH⁻] at all times. Neutral ≠ pH 7; neutral = [H⁺] = [OH⁻].
Example: At 37°C (body temp), Kw = 2.4 × 10⁻¹⁴. For pure water: [H⁺]² = 2.4 × 10⁻¹⁴ → [H⁺] = 1.55 × 10⁻⁷ M → pH = 6.81 (still neutral!)
HLKa, pKa, Kb, pKb
Definitions
- For weak acid HA: Ka = [H⁺][A⁻] / [HA] → pKa = −log Ka
- For weak base B: Kb = [BH⁺][OH⁻] / [B] → pKb = −log Kb
- Stronger acid → larger Ka, smaller pKa
- Stronger base → larger Kb, smaller pKb
The Conjugate Relationship
For any conjugate acid-base pair:
Ka × Kb = Kw and pKa + pKb = pKw = 14.00 (at 25°C)
This means: the stronger an acid, the weaker its conjugate base (and vice versa). A strong acid like HCl has a conjugate base (Cl⁻) so weak it doesn't act as a base in water.