Bronsted-Lowry acids and bases and pHEdexcel International A Level Chemistry: Revision notes
Section 1
Brønsted–Lowry acids and bases
A Brønsted–Lowry acid is a proton donor and a Brønsted–Lowry base is a proton acceptor. An acid–base reaction is therefore a proton transfer: the acid gives H⁺ and the base takes it, usually by using a lone pair of electrons.
Example: HCl + NH₃ → NH₄⁺ + Cl⁻. HCl donates a proton and NH₃ accepts it. In water, an acid transfers its proton to H₂O to form the oxonium ion, H₃O⁺ (often written H⁺(aq)).
Water is amphoteric: it accepts protons (base) in the presence of acids and donates them (acid) in the presence of bases.
Section 2
Conjugate acid–base pairs
When an acid donates a proton, what remains is its conjugate base. When a base accepts a proton, it forms its conjugate acid. A conjugate pair differs by exactly one proton.
In CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺ there are two pairs:
- CH₃COOH (acid) / CH₃COO⁻ (conjugate base)
- H₂O (base) / H₃O⁺ (conjugate acid)
To find a pair, look for two species on opposite sides of the equation that differ by H⁺.
Pairing species from the same side of the equation (for example CH₃COOH with H₂O). A conjugate pair always has one species on each side.
Section 3
Defining pH
pH is a measure of the hydrogen ion concentration: pH = −log₁₀[H⁺], where [H⁺] is in mol dm⁻³. Rearranged, [H⁺] = 10⁻ᵖᴴ.
A low pH means a high [H⁺]. Because the scale is logarithmic, a change of one pH unit is a tenfold change in [H⁺]. Give pH to two decimal places, matching the precision of the data.
Example: pH 3.20 gives [H⁺] = 10⁻³·²⁰ = 6.31 × 10⁻⁴ mol dm⁻³.
On a calculator use 10ˣ with the negative of the pH to get back to [H⁺]. Check by taking the log again.
Section 4
Strong and weak acids
A strong acid is fully dissociated in aqueous solution: HCl → H⁺ + Cl⁻, with a single arrow. A weak acid is only partially dissociated, so an equilibrium is set up, for example CH₃COOH ⇌ CH₃COO⁻ + H⁺, with the position of equilibrium well to the left.
The degree of dissociation, not the concentration, is what makes an acid strong or weak. Two acids of the same concentration therefore have different pH: 0.100 mol dm⁻³ HCl has pH 1.00, while 0.100 mol dm⁻³ ethanoic acid has pH 2.88.
Confusing strong with concentrated. Strength is about dissociation; concentration is about amount per dm³.
Section 5
Calculating the pH of a strong acid
For a monoprotic strong acid [H⁺] equals the acid concentration, so pH = −log₁₀(c).
Worked example: 25.0 cm³ of 0.0800 mol dm⁻³ HCl is made up to 1.00 dm³ with water.
- Amount of HCl = 0.0800 × 0.0250 = 2.00 × 10⁻³ mol
- [H⁺] = 2.00 × 10⁻³ ÷ 1.00 = 2.00 × 10⁻³ mol dm⁻³
- pH = −log₁₀(2.00 × 10⁻³) = 2.70
When two strong acids are mixed, add the amounts of H⁺ and divide by the total volume. For a weak acid, you cannot assume [H⁺] equals the concentration; you must be given the pH or Ka.
Averaging the two pH values when mixing acids. Always work with amounts of H⁺ and convert to pH at the end.
That's the notes covered.
Carry on to the next subtopic.
Exam questions on Bronsted-Lowry acids and bases and pH
- Ethanoic acid dissolves in water and reaches the equilibrium CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺.Identify the two conjugate acid–base pairs in this equilibrium.2 marks
- A swimming-pool technician lowers the pH of pool water by dosing it with dilute hydrochloric acid. Treat hydrochloric acid as fully dissociated, and assume all solutions are at 25 °C.A technician dilutes 25.0 cm³ of 0.0800 mol dm⁻³ hydrochloric acid with water to a total volume of 1.00 dm³. Calculate the pH of the diluted solution.2 marks
- A student compares two acid solutions at 25 °C, each of concentration 0.100 mol dm⁻³: hydrochloric acid and ethanoic acid. The pH of the ethanoic acid solution is measured as 2.88.Calculate the pH of the hydrochloric acid and explain why the ethanoic acid has a higher pH at the same concentration.3 marks
Written by the Exaim team, led by Shaun Daswani (Head of Upper Secondary, Improve ME Institute; MSc Financial Mathematics, Imperial College London; BSc, UCL) and Jason Daswani (operational lead, Improve ME Institute; LSE).