Summary Revision Notes

Acid–Base Equilibria Revision
12 October 2025 11:10


1. Definitions
• Arrhenius Acid: Produces H+ in water
• Arrhenius Base: Produces OH- in water
• Bronsted–Lowry Acid: Proton donor
• Bronsted–Lowry Base: Proton acceptor
• Conjugate Pair: Two species that differ by one proton
Example: HCl / Cl-, NH4+ / NH3

2. Strong vs Weak
Type Dissociation
Strong Acid Fully (HCl, HNO3)
Weak Acid Partially (CH3COOH)
Strong Base Fully (NaOH, KOH)
Weak Base Partially (NH3)


3. pH Calculations
Strong Acid:
[H+] = concentration
pH = -log[H+]
Strong Base:
[OH-] = concentration
pOH = -log[OH-]
pH = 14 - pOH
Weak Acid:
HA ⇌ H+ + A-
Ka = [H+][A-] / [HA]
If weak: [H+] = sqrt(Ka × HA concentration)

4. Kw (Ionisation of Water)
H2O ⇌ H+ + OH-
Kw = [H+][OH-]
At 25°C, Kw = 1.00 × 10^-14
Neutral pH = when [H+] = [OH-], not always pH 7 (changes with temperature)

5. Buffer Solutions
Definition: Solution that resists pH change when small amounts of acid or base are added.
Made from:
• Weak acid + salt of that acid (CH3COOH + CH3COONa)
• Weak base + its salt (NH3 + NH4Cl)
Henderson–Hasselbalch:
pH = pKa + log ( [A-] / [HA] )

6. How Buffers Work
Add acid (H+):
A- + H+ → HA
Add base (OH-):
HA + OH- → A- + H2O

7. Titration Curves

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