Acids and bases (A Level) flashcards

Topic 12: Acids and bases

Brønsted–Lowry acid and base

An acid is a proton donor; a base is a proton acceptor.

pH

−log₁₀[H⁺].

Kw: expression, value at 298 K and units

Kw = [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K.

Pure water at a higher temperature has a pH below 7. Is it acidic?

No, it is still neutral, because [H⁺] = [OH⁻]. Kw increases with temperature because the dissociation of water is endothermic.

pH of 0.0100 mol dm⁻³ NaOH at 298 K

[H⁺] = Kw ÷ [OH⁻] = 1.00 × 10⁻¹⁴ ÷ 0.0100 = 1.00 × 10⁻¹² mol dm⁻³, so pH = 12.00.

pH of a weak acid: the two assumptions

[H⁺] = [A⁻], and [HA] at equilibrium equals the starting concentration (only slight dissociation). So [H⁺] = √(Ka × [HA]).

pH of a buffer made from HA and its salt

[H⁺] = Ka × [HA] ÷ [A⁻] (moles can be used, as the volume cancels), then pH = −log₁₀[H⁺]. At half-neutralisation pH = pKa.

Buffer

A solution that keeps an approximately constant pH despite dilution or the addition of small amounts of acid or base.

How does an acidic buffer (CH₃COOH / CH₃COO⁻) resist added acid and added alkali?

Added H⁺: H⁺ + A⁻ → HA. Added OH⁻: OH⁻ + HA → A⁻ + H₂O. The large reservoirs of HA and A⁻ mean the ratio [HA]/[A⁻] barely changes.

Indicator for a weak acid–strong base titration, and why

Phenolphthalein (range 8.3–10.0), because the vertical section and the equivalence point are in the alkaline range.

Ka vs pKa: which way means a stronger acid?

Larger Ka means stronger acid; larger pKa means weaker acid.