Acids and bases — common mistakes

When asked to explain how a buffer works, you must write two separate equations: one showing the reaction with added H⁺ and one showing the reaction with added OH⁻. State that there are large reservoirs of both the weak acid and its conjugate base.

When sketching titration curves, always label your axes (pH on y-axis, Volume of titrant/cm³ on x-axis). Mark the volume and pH at the equivalence point. For weak acid/base titrations, also indicate the approximate pH at the half-equivalence point, where $pH = pK_a$.

Strength and concentration are different concepts. 'Strength' refers to the degree of dissociation of an acid in water; a strong acid like HCl dissociates completely. 'Concentration' refers to the amount of acid substance dissolved in a unit volume of solution. It is possible to have a dilute solution of a strong acid (e.g., 0.01 M HCl) or a concentrated solution of a weak acid (e.g., 17 M ethanoic acid).

The pH at the equivalence point is 7 only for a strong acid-strong base titration. For a weak acid-strong base titration, the salt formed at equivalence contains the conjugate base of the weak acid, which hydrolyses water to produce OH⁻ ions, making the solution alkaline (pH > 7). Conversely, for a strong acid-weak base titration, the conjugate acid formed makes the solution acidic (pH < 7).

The indicator's pKin value (the pH at which it changes colour) should be as close as possible to the pH at the equivalence point of the titration. More accurately, the entire pH range of the indicator's colour change must lie on the steep, vertical part of the titration curve. For example, phenolphthalein (pH range 8.3-10.0) is suitable for weak acid-strong base titrations, while methyl orange (pH range 3.1-4.4) is suitable for strong acid-weak base titrations.