Acids, Bases & Buffers (OCR A Level Chemistry A): Flashcards

pKa is calculated using the expression pKa = -log Ka. A lower pKa value corresponds to a stronger weak acid.

pH = -log[H⁺], where [H⁺] is the concentration of hydrogen ions in mol dm^-3. The scale is logarithmic with base 10.

True.

Kw = [H⁺][OH^-] and always equals 1 × 10^-14 mol² dm^-6 at 298 K, reflecting the self-ionisation of water.

Water undergoes self-ionisation, where a small proportion of molecules donate a proton to another water molecule, producing equal concentrations of H⁺ and OH^- so pH = 7 at 298 K.

A strong acid is one that is completely ionised in aqueous solution, so the concentration of H⁺ equals the initial concentration of the acid.

Strong acids are fully ionised, so every mole of acid produces one mole of H⁺. The H⁺ contribution from water self-ionisation is negligible in comparison.

To find the pH of a strong base, first calculate [H⁺] using [H⁺] = Kw ÷ [OH^-], then apply pH = -log [H⁺].

A weak acid only partially dissociates, so [H⁺] is much smaller than [HA]. Setting [H⁺] = [A^-] and combining terms in the Ka expression gives [H⁺]² = Ka × [HA], hence the square root.

True.

Because the degree of dissociation is very small, the change in [HA] is negligible, so the initial concentration is used as the equilibrium concentration in the Ka expression.

First, [H⁺] = [A^-] at equilibrium (H⁺ from water is negligible). Second, the initial concentration of the undissociated acid equals its equilibrium concentration because dissociation is so slight.

True.

A larger Ka means greater dissociation, giving more H⁺. Since pKa = -log Ka, a larger Ka produces a lower pKa.

As the acid becomes stronger, dissociation increases and the difference between initial [HA] and equilibrium [HA] grows, making the assumption that initial [HA] is unchanged less valid and introducing greater error.

A buffer solution is one that resists changes in pH when small amounts of acid or alkali are added, by maintaining a near-constant concentration of H⁺ ions.

The mixture provides a large reservoir of both the undissociated weak acid (to react with added OH^-) and the conjugate base (to react with added H⁺), keeping [H⁺] nearly constant.

The Henderson-Hasselbalch equation states: pH = pKa + log₁₀ ([salt ] / [acid]).

The extra H⁺ reacts with CH₃COO^- to form CH₃COOH, shifting the equilibrium left and consuming the added acid so pH remains almost constant.

False.

For a buffer to form, the weak acid must be in excess so that both the acid and its conjugate salt are present. If the alkali is in excess, the weak acid is fully consumed and no buffer forms.

You must state that the weak acid is in excess (partial neutralisation has occurred) and that both the weak acid and its conjugate salt are present in the resulting solution.

True.

OH^- reacts with H⁺ to form water, lowering [H⁺]. The equilibrium shifts right as CH₃COOH dissociates further to replenish H⁺ and maintain a constant pH.

Excess H⁺ reacts with HCO₃^- to shift the equilibrium CO₂ + H₂O ⇌ H⁺ + HCO₃^- to the left, removing H⁺ and preventing acidosis.

An equivalence point (stoichiometric point) is the stage in a titration where the acid and base have reacted in exact molar ratio, located at the midpoint of the vertical section of the pH curve.

There is no sharp, vertical section on the pH curve, so no indicator has a colour-change range narrow enough to fall within the gradual pH change at the equivalence point.

An acid-base indicator is a weak acid (HIn) whose undissociated form and conjugate base (In^-) have different colours, allowing it to signal the endpoint of a titration by a colour change.

In a strong acid-strong base titration the pH changes from 4 to 10 at the end-point, so both methyl red and phenolphthalein are suitable indicators.

The colour depends on the ratio of [HIn] to [In^-], which shifts continuously as pH changes. The transition occurs over a range centred on the pKa of the indicator, not at a single abrupt point.

True.

For weak acid-strong base the pH range at the end-point is 7-10, which matches phenolphthalein (pKa ≈ 9.4). For strong acid-weak base the range is 4-7, so methyl red is required instead.

The colour-change pH range of the indicator must fall within the steep, near-vertical section of the titration curve at the equivalence point. The indicator pKa should match the pH at the equivalence point.