What a learner can do afterwards
- Writes the Ka expression for a named weak acid
- Calculates the pH of a weak acid solution and names the two approximations the calculation makes
- Converts between Ka and pKa and ranks acids by strength from either
- Explains why a weak acid and a strong acid at the same concentration differ in pH yet need the same volume of alkali to neutralise
1 · Read
Vinegar holds ethanoic acid and lemon holds citric acid. Both are weak acids that only partly split into ions. For a weak acid HA splitting into H+ and A-, the equilibrium constant is Ka = [H+][A-] over [HA]. At the same concentration and temperature, a bigger Ka means a bigger share has split, so bigger Ka means a stronger weak acid.
Take ethanoic acid, the acid in vinegar, as a 0.10 mol per cubic decimetre solution with Ka 1.7 x 10-5. The working gives [H+] near 1.3 x 10-3, so the pH is about 2.9. It leans on two quiet guesses: the water gives almost no H+ next to the acid, and 0.10 stays near 0.10 since so little splits.
pKa = -log Ka puts strength on a pH-like scale: lower pKa wins, and each single unit is a factor of ten in Ka. Ranking by pKa and ranking by Ka always agree.
Take equal volumes of 0.10 mol per cubic decimetre ethanoic acid and 0.10 mol per cubic decimetre hydrochloric acid. The strong acid shows a lower pH because it is fully split. Yet both need the same volume of the same alkali to neutralise. A monoprotic acid can give one proton per molecule, so both flasks hold the same total neutralisable amount, and the titre cannot tell them apart.
Ka measures the split, pKa rescales it, approximations unlock the pH, and titre counts total neutralisable amount, not strength.
The rewritten readings were written for this dataset. The retained reading '14.3: Relative Strengths of Acids and Bases' comes from the supplied library record for OpenStax Chemistry; its license terms are not established in this packet.
More to read
Comparing acids with Ka and pKa
In water, strong acids ionise essentially completely. Weak acids, including ethanoic acid, ionise only partly and reach equilibrium. For HA ⇌ H+ + A-, the constant is Ka = [H+][A-]/[HA]. Under matching conditions, a larger Ka indicates stronger ionisation and greater acid strength.
Weak-acid Ka values are often small. Chemists express them using pKa = -log10(Ka). The ranking runs in the opposite direction: smaller pKa indicates a stronger acid. Converting between the two forms lets you compare acids using either kind of data.
14.3: Relative Strengths of Acids and Bases
To find the pH of a weak acid you let x stand for [H+], then notice that [A-] equals x too, since each split makes one of each. That turns Ka into x squared over the starting concentration, which solves in one step. For 0.10 mol/dm3 ethanoic acid with Ka 1.7 x 10-5, x comes out near 1.3 x 10-3, giving pH about 2.9. The working quietly leans on two approximations. It assumes the starting concentration barely drops, so [HA] stays at its initial value, and it ignores the H+ that water itself supplies. Both are safe here because x is tiny next to 0.10 and huge next to the ions from pure water.
Why pH and neutralisation differ
Compare matching volumes and concentrations of monoprotic acids, one strong and one weak. When water's own ionisation is negligible, the strong acid produces an H+ concentration approximately equal to its acid concentration. Its pH is therefore lower. The weak acid ionises less completely under these conditions, giving a higher pH, sometimes by three or more units.
Complete neutralisation still requires equal volumes of the same alkali. OH- removes H+, prompting further weak-acid ionisation until all the acid has reacted. pH reflects the H+ present in solution; the neutralisation titre measures the total amount of acid that can react.
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