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.
Strong vs Weak Acids and Bases
| Property | Strong Acid / Base | Weak Acid / Base |
|---|---|---|
| Dissociation | 100% complete (\(\rightarrow\)) | Partial equilibrium (\(\rightleftharpoons\)) |
| [H⁺] or [OH⁻] | Equal to initial concentration (monoprotic) | Much less than initial concentration |
| Electrical Conductivity | High (high mobile ion concentration) | Low (low mobile ion concentration) |
| Reaction Rate | Fast (with metals, carbonates) | Slow (with same equimolar reagents) |
IB Required Lists
Memorise these. Any acid/base not on this list should be treated as weak.
Required Strong Acids & Bases
Strong Acids (Must Memorise)
- Hydrochloric acid: \(\text{HCl}\)
- Nitric acid: \(\text{HNO}_3\)
- Sulfuric acid: \(\text{H}_2\text{SO}_4\)
Strong Bases (Must Memorise)
- Group 1 hydroxides: \(\text{LiOH}, \text{NaOH}, \text{KOH}\)
- Group 2 soluble hydroxides: \(\text{Ba(OH)}_2\)
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
pH and Water Auto-ionisation Formulas
\[\text{pH} = -\log_{10}[\text{H}^+] \quad\Longleftrightarrow\quad [\text{H}^+] = 10^{-\text{pH}}\] \[K_{\text{w}} = [\text{H}^+][\text{OH}^-] = 1.00 \times 10^{-14} \text{ at } 298\text{ K}\] \[\text{pH} + \text{pOH} = 14.00 \text{ (at } 298\text{ K)}\]
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.
Temperature Dependence of Kw and Water Neutrality
The auto-ionisation of water (\(\text{H}_2\text{O} \rightleftharpoons \text{H}^+ + \text{OH}^-\)) is endothermic (\(\Delta H > 0\)).
- As temperature rises, equilibrium shifts right → \(K_{\text{w}}\) increases → \([\text{H}^+]\) increases → pH decreases below 7.00.
- However, water remains strictly neutral because \([\text{H}^+] = [\text{OH}^-]\) always holds for pure water!
HLKa, pKa, Kb, pKb
Acid/Base Dissociation Constants (Ka, Kb)
For weak acid \(\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^-\) and weak base \(\text{B} + \text{H}_2\text{O} \rightleftharpoons \text{BH}^+ + \text{OH}^-\):
Higher \(K_{\text{a}}\) (lower \(\text{p}K_{\text{a}}\)) = stronger acid.
For conjugate pair: \(K_{\text{a}} \times K_{\text{b}} = K_{\text{w}}\) and \(\text{p}K_{\text{a}} + \text{p}K_{\text{b}} = 14\).
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.
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