Understanding Buffers in A-Level Chemistry

How Buffer Solutions Work

Ever wondered why your blood doesn't become dangerously acidic when you drink a fizzy drink? That's buffers in action! A buffer solution consists of a weak acid and one of its salts working together to maintain a stable pH.

Take ethanoic acid (CH₃COOH) mixed with sodium ethanoate (CH₃COONa) as a classic example. The weak acid only partially dissociates, whilst the salt completely dissociates to flood the solution with ethanoate ions (CH₃COO⁻).

When you add acid (H⁺ ions), these react with the abundant ethanoate ions to form more ethanoic acid molecules. This shifts the equilibrium left and mops up the extra H⁺. When you add alkali (OH⁻ ions), they react with H⁺ to make water, so the ethanoic acid dissociates more to replace the lost H⁺ ions.

Key Insight: The salt provides a "reservoir" of conjugate base ions that can neutralise added acid, whilst the weak acid can neutralise added base.

Calculating Buffer pH

Calculating buffer pH might seem tricky, but it's actually straightforward once you know the formula. You'll use the expression: [H⁺] = Ka × [weak acid]/[conjugate base], where Ka is the acid dissociation constant.

Let's work through a typical exam question: 100 cm³ of 0.75 mol dm⁻³ NaOH added to 150 cm³ of 1.5 mol dm⁻³ HCOOH. First, find the moles of each reactant, then determine what's left after the neutralisation reaction.

The calculation shows 0.075 mol of salt (HCOONa) is produced, leaving 0.15 mol of unreacted acid. In the final 250 cm³ solution, you get concentrations of 0.6 mol dm⁻³ for the acid and 0.3 mol dm⁻³ for the salt.

Exam Tip: Always remember to use the total volume when calculating final concentrations - it's the sum of both solutions mixed together!

Plugging into the formula: [H⁺] = 1.78×10⁻⁴ × $0.6/0.3$ = 3.56×10⁻⁴, giving pH = 3.45.