Flashcards · 9620 · A Level Topic 10

Acids and bases (A Level) flashcards

28 cards on A Level Topic 10, subtopics 3.1.10.1 to 3.1.10.6. Pick a subtopic or study them all.

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Brønsted–Lowry acid and base: Acid: proton donor. Base: proton acceptor.
Acid–base equilibrium: One in which protons (H⁺) are transferred. Check both directions: the species that loses H⁺ in the reverse reaction is also an acid.
In HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O, which species is the Brønsted–Lowry acid?: HCO₃⁻: it donates a proton to OH⁻.
Strong vs concentrated (acid): “Strong” describes the extent of dissociation (complete); “concentrated” describes how many moles are in each dm³.

Definition of pH: pH = −log₁₀[H⁺], so [H⁺] = 10^(−pH).
pH of a strong monoprotic acid: [H⁺] = concentration of the acid; pH = −log₁₀[H⁺]. Example: 0.0072 mol dm⁻³ HI has pH = −log₁₀(0.0072) = 2.14.
Calculate [H⁺] in a solution of pH 10.46: [H⁺] = 10^(−10.46) = 3.47 × 10⁻¹¹ mol dm⁻³
Method: pH after mixing a strong acid with a strong base: Find n(H⁺) and n(OH⁻), subtract the smaller from the larger, divide the excess by the combined volume. Excess H⁺ gives pH directly; excess OH⁻ needs Kw.

Ionic product of water, Kw: Kw = [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ mol² dm⁻⁶ at 25 °C. It comes from Kc for H₂O ⇌ H⁺ + OH⁻, with [H₂O] effectively constant.
pH of a strong base: [OH⁻] = concentration × OH⁻ per formula unit; [H⁺] = Kw ÷ [OH⁻]; pH = −log₁₀[H⁺]. Example: 0.0028 mol dm⁻³ NaOH → [H⁺] = 3.57 × 10⁻¹² → pH 11.45.
Neutral solution: [H⁺] = [OH⁻]. This is pH 7 only at 25 °C.
What happens to Kw and the pH of pure water as temperature rises?: Water’s dissociation is endothermic, so Kw increases and the pH falls below 7. The water is still neutral because [H⁺] = [OH⁻].

Weak acid or base: Dissociates only slightly in water.
Ka expression and pKa: Ka = [H⁺][A⁻] ÷ [HA] (mol dm⁻³). pKa = −log₁₀ Ka. Larger Ka = stronger acid; larger pKa = weaker acid.
pH of a weak acid from Ka: Assume [H⁺] = [A⁻] and [HA] ≈ starting concentration, so [H⁺] = √(Ka × [HA]). Example: Ka = 2.0 × 10⁻⁶, [HA] = 0.500 → [H⁺] = 1.0 × 10⁻³, pH 3.00.
Ka from the pH of a weak acid solution: [H⁺] = 10^(−pH), then Ka = [H⁺]² ÷ [HA].

Indicator colour-change ranges: methyl orange and phenolphthalein: Methyl orange: 3.1–4.4 (red to yellow). Phenolphthalein: 8.3–10.0 (colourless to pink).
When is an indicator suitable for a titration?: When its entire colour-change range lies within the vertical (steep) section of the pH curve.
Indicator for weak acid–strong base and for strong acid–weak base titrations: Weak acid–strong base (vertical about 7–11): phenolphthalein. Strong acid–weak base (vertical about 3–7): methyl orange. Weak–weak: no sharp change; use a pH meter.
How can you find pKa from a weak acid–strong base pH curve?: Read the pH at half the equivalence volume. There [HA] = [A⁻], so pH = pKa.
End point vs equivalence point: The end point is where the indicator changes; the equivalence point is where reacting amounts match.
Required practical 7: key steps for a pH curve: Calibrate the pH meter with buffers. Add base from a burette to the acid, stir, record pH and total volume, using smaller additions near the end point. Plot pH against volume.

Buffer: A solution whose pH stays nearly constant when it is diluted or when a small amount of acid or base is added.
What does an acidic buffer contain? A basic buffer?: Acidic: a weak acid and its salt (or a weak acid partly neutralised by NaOH). Basic: a weak base and its salt, e.g. methylamine with methylammonium chloride.
How does an acidic buffer resist added acid and added base?: Added H⁺ reacts with A⁻ to form HA (HA ⇌ H⁺ + A⁻ shifts left). Added OH⁻ removes H⁺, and HA dissociates to replace it (shifts right). Large amounts of HA and A⁻ keep [HA]/[A⁻] almost constant.
[H⁺] in an acidic buffer: [H⁺] = Ka × [HA] ÷ [A⁻] (or the ratio of moles). Do not use √(Ka[HA]) for a buffer.
Buffer: 0.024 mol HA and 0.016 mol A⁻, Ka = 3.0 × 10⁻⁵ mol dm⁻³. pH?: [H⁺] = 3.0 × 10⁻⁵ × 0.024 ÷ 0.016 = 4.5 × 10⁻⁵ mol dm⁻³, so pH = 4.35.
Applications of buffers: Calibrating pH meters, controlling the pH of blood, controlling acidity in foods and drinks, and any process (such as with enzymes) that needs a steady pH.