Acids, Bases, pH and Buffers

Acids and bases (Brønsted-Lowry)

• An acid is a proton (H⁺) donor; a base is a proton acceptor.
• In water, an acid donates H⁺ to form H₃O⁺ (often written H⁺).
• Conjugate pairs: an acid becomes its conjugate base after losing H⁺.

Strong vs weak

• Strong acids (HCl, HNO₃, H₂SO₄) fully dissociate in water.
• Weak acids (e.g. ethanoic acid) only partially dissociate — an equilibrium.

pH

pH = −log₁₀[H⁺]        and        [H⁺] = 10^(−pH)

For a strong monoprotic acid, [H⁺] = the acid concentration.

The acid dissociation constant, Ka

For a weak acid HA ⇌ H⁺ + A⁻:

Ka = [H⁺][A⁻] ÷ [HA]         pKa = −log₁₀ Ka

• A larger Ka (smaller pKa) = stronger acid.
• For a weak acid, assume [H⁺] ≈ [A⁻], so [H⁺] = √(Ka × [HA]).

The ionic product of water, Kw

Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ (at 298 K)

Use this to find the pH of strong bases: [H⁺] = Kw ÷ [OH⁻].

Buffers

A buffer resists changes in pH when small amounts of acid or base are added. An acidic buffer is made from a weak acid + its salt (e.g. ethanoic acid + sodium ethanoate), providing a reservoir of both HA and A⁻:

• Add acid → A⁻ removes the extra H⁺.
• Add base → HA releases H⁺ to neutralise the OH⁻.

Buffer pH is found by rearranging Ka:

[H⁺] = Ka × ([HA] ÷ [A⁻])

Titration curves

Plotting pH against volume of base added shows a characteristic curve with a steep vertical section at the equivalence point. The correct indicator must change colour within that steep range (e.g. phenolphthalein for strong base–weak acid).

Worked example

Find the pH of 0.10 mol dm⁻³ ethanoic acid, Ka = 1.8 × 10⁻⁵.

• [H⁺] = √(Ka × [HA]) = √(1.8 × 10⁻⁵ × 0.10) = √(1.8 × 10⁻⁶) = 1.34 × 10⁻³.
• pH = −log(1.34 × 10⁻³) = 2.87. ✓

Common mistakes

• Treating a weak acid like a strong one ([H⁺] ≠ acid concentration for weak acids).
• Forgetting to use Kw to get [H⁺] from [OH⁻] for bases.
• Choosing an indicator that doesn't change within the steep part of the titration curve.

Exam tips

• Learn pH = −log[H⁺], Ka/pKa, Kw and the buffer expression.
• For weak acids, use [H⁺] = √(Ka[HA]); for strong bases, use Kw.
• Explain buffer action with the equilibrium HA ⇌ H⁺ + A⁻ shifting.

Key facts to remember

• Acid = H⁺ donor; pH = −log[H⁺]; strong acids fully dissociate, weak acids partially (Ka, small pKa = stronger).
• Kw = [H⁺][OH⁻] = 10⁻¹⁴; use it for strong-base pH.
• A buffer (weak acid + its salt) resists pH change; [H⁺] = Ka × [HA]/[A⁻].