Acid-base Equilibria (Edexcel A Level Chemistry): Flashcards

[H⁺] is found by rearranging pH = −log₁₀[H⁺]:

[H⁺] = 10−pH^

For example, pH 3.10 gives [H⁺] = 10−3.10 ^{= 7.94 × 10}−4 ^{mol dm}−3^.

True.

The pH scale is logarithmic (base 10). Each unit decrease in pH means [H⁺] increases by a factor of 10, and each unit increase means [H⁺] decreases by a factor of 10.

pH = −log₁₀[H⁺] = −log₁₀(1.60 × 10⁻⁴) = 3.80

pH values are given to 2 decimal places; the digits before the decimal reflect the order of magnitude.

pH = −log₁₀ [H⁺]

[H⁺] = 10 ^−pH

The pH of the final solution is 3.

[H⁺] is reduced by a factor of 100 (= 10²). On a logarithmic scale, this corresponds to an increase of 2 pH units: pH 1.0 + 2 = 3.0.

False.

pH can be negative for very concentrated strong acid solutions where [H⁺] > 1 mol dm⁻³, since −log(value > 1) gives a negative result.

pH values are normally given to 2 decimal places.

The digits before the decimal point reflect the order of magnitude of [H⁺], while the two decimal places carry the significant figures of the concentration measurement.

A strong acid is an acid that dissociates completely in aqueous solution, so the equilibrium lies entirely to the right.

Examples: HCl, HNO₃, H₂SO₄.

Two assumptions are made:

1. The contribution of H⁺ from the ionisation of water is negligible.

2. The equilibrium concentrations of H⁺ and A⁻ are equal (since both come from the same dissociation step).

This allows the simplification [H⁺]² = K_a × [HA].

A weak acid is an acid that only partially dissociates in aqueous solution, establishing an equilibrium between the undissociated acid and its ions.

Examples: CH₃COOH, HCN, H₂CO₃.

True.

A higher K_a indicates greater dissociation at equilibrium, meaning more H⁺ ions are produced. The acid is therefore stronger. For example, K_a of HCl ≫ K_a of CH₃COOH.

HCl is a strong acid that fully dissociates:

[H⁺] = [HCl] = 0.01 mol dm⁻³

pH = −log(0.01) = 2.00

K_a is the acid dissociation constant — the equilibrium constant for the dissociation of a weak acid HA in water:

K_a = [H⁺][A⁻] / [HA]

Units: mol dm⁻³. A larger K_a means a stronger acid.

For a weak acid HA, the K_a expression is:

K_a = [H⁺][A⁻] ÷ [HA]

[H⁺]² = K_a × [HA] = 1.74 × 10⁻⁵ × 0.100 = 1.74 × 10⁻⁶

[H⁺] = √(1.74 × 10⁻⁶) = 1.32 × 10⁻³ mol dm⁻³

pH = −log(1.32 × 10⁻³) = 2.88

True.

Weak acids only partially ionise, so energy must be used to complete their ionisation before neutralisation can occur. This makes the overall enthalpy change less negative (less exothermic) compared to the −57 to −58 kJ mol⁻¹ typical of strong acid–strong base reactions.

K_w is the ionic product of water — the equilibrium constant for the self-ionisation of water:

K_w = [H⁺][OH⁻]

At 298 K, K_w = 1.0 × 10⁻¹⁴ mol² dm⁻⁶.

True.

In pure water, self-ionisation produces equal concentrations of H⁺ and OH⁻. Since K_w = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴, each concentration equals √(1.0 × 10⁻¹⁴) = 1.0 × 10⁻⁷ mol dm⁻³.

K_w = 1.0 × 10⁻¹⁴ at 298 K, where the units are mol² dm⁻⁶.

For a strong base, [OH⁻] = [base].

Rearrange K_w = [H⁺][OH⁻] to find:

[H⁺] = K_w / [OH⁻]

Then: pH = −log[H⁺]

For example, 0.15 mol dm⁻³ NaOH gives pH = 13.18.

pK_w is defined as:

pK_w = −log K_w

At 298 K, pK_w = 14. It is the negative logarithm of the ionic product of water, analogous to the relationship between [H⁺] and pH.

True.

Even in alkaline solutions, the self-ionisation of water still produces a small concentration of H⁺. K_w = [H⁺][OH⁻] applies to all aqueous solutions; in alkalis, [OH⁻] > [H⁺], so [H⁺] is small but non-zero.

NaOH is a strong base: [OH⁻] = 0.15 mol dm⁻³

[H⁺] = K_w / [OH⁻] = (1.0 × 10⁻¹⁴) / 0.15 = 6.67 × 10⁻¹⁴ mol dm⁻³

pH = −log(6.67 × 10⁻¹⁴) = 13.18

[H⁺] = 10−10.50 ^{= 3.16 × 10}−11 ^{mol dm}−3^

[OH⁻] = K_w / [H⁺] = (1.0 × 10⁻¹⁴) / (3.16 × 10⁻¹¹)

[OH⁻] = 3.16 × 10⁻⁴ mol dm⁻³

To find [H⁺] in a solution of a strong base, use:

[H⁺] = K_w ÷ [OH⁻]

CH₃COONa is formed from a weak acid (CH₃COOH) and a strong base (NaOH). The ethanoate ion accepts protons from water, producing an alkaline solution.

A pH meter is an instrument used to measure the pH of a solution accurately by detecting the hydrogen ion concentration.

By measuring the pH of equimolar aqueous solutions at the same temperature — a higher pH indicates a weaker acid.

False.

A higher pH means a lower hydrogen ion concentration. A higher pH for an equimolar acid solution indicates a weaker acid, not a stronger one.

At very low acid concentrations, the hydrogen ions from water dissociation become significant. The actual pH is close to 7, as the contribution from water cannot be ignored.

For a strong acid, as concentration increases by a factor of 10, the pH decreases by one unit.

For a weak acid, a tenfold increase in concentration decreases the pH by approximately 0.5 units, compared to 1 unit for a strong acid, because weak acids only partially dissociate.

True.

Both salts are formed from a strong acid and a strong base, so neither ion reacts significantly with water. The resulting solution has a pH of 7.00.

NH₄Cl is formed from a strong acid (HCl) and a weak base (NH₃). The ammonium ion partially donates protons to water, producing an acidic solution.

The equivalence point is the point in a titration at which the acid and base have completely reacted in stoichiometric proportions, with no excess of either.

1. Strong acid + strong base

2. Weak acid + strong base

3. Strong acid + weak base

4. Weak acid + weak base

True.

Phenolphthalein changes colour in the pH range 8.2–10.0. This falls within the rapid pH change at the equivalence point for both titration types.

An indicator is a weak acid whose undissociated form (HIn) and conjugate base (In^-) are different colours, so its colour changes with the pH of the solution.

There is no sharp pH change at the equivalence point, so no common indicator changes colour within the gradual pH transition. The equivalence point cannot be reliably detected.

For an indicator, the pKa of the indicator equals the pH at its endpoint, because at that point [HIn] = [In^-].

Methyl red is most suitable. Its colour change occurs in the pH range 4–6, which falls within the rapid pH change (pH 4–7) at the equivalence point of a strong acid-weak base titration.

False.

The end point is when the indicator changes colour; the equivalence point is when the acid and base have reacted in exact stoichiometric proportions. They are close but not identical.

From a pH titration curve you can read the initial pH of the acid, the pH at the equivalence point, the volume of base at the equivalence point, and the pH range of the near-vertical section.

A buffer solution is a solution that resists changes in pH when small amounts of acid or alkali are added.

The equilibrium shifts left: added H⁺ ions react with CH₃COO^- to form more CH₃COOH. The large reserve of ethanoate ions keeps the pH nearly constant.

The Henderson-Hasselbalch equation states: pH = pKa + log([A^-] / [HA]).

OH^- reacts with H⁺ to form water. The equilibrium shifts right: CH₃COOH dissociates to replace the H⁺ consumed. The large reserve of ethanoic acid keeps the pH nearly constant.

True.

Buffers can consist of either a weak acid and its conjugate base or a weak base and its conjugate acid. Both systems resist pH change.

Sodium ethanoate is a salt that fully ionises in solution, releasing a large reserve of CH₃COO^- (conjugate base) that can react with any added H⁺ ions.

pKa is the negative logarithm (base 10) of the acid dissociation constant Ka: pKa = -log₁₀ Ka. It is a measure of acid strength; a lower pKa indicates a stronger acid.

HCO₃^- ions act as a conjugate base in the equilibrium CO₂ + H₂O ⇌ H⁺ + HCO₃^-. This equilibrium shifts left or right to resist pH changes and keeps blood pH between 7.35 and 7.45.

False.

A blood pH below 7.35 is called acidosis (too much acid). Alkalosis refers to a blood pH that is too high (above 7.45).

At the equivalence point, the weak acid has been fully converted to its conjugate base (e.g. CH₃COO^-). This ion is a weak base that accepts protons from water, raising the pH above 7.

The half-equivalence point is the stage in a weak acid-strong base titration at which exactly half the weak acid has been neutralised, so [acid] = [conjugate base] and pKa = pH.

False.

A weak acid partially dissociates, so its H⁺ concentration is lower than that of an equimolar strong acid. Its starting pH is therefore higher (e.g. _{3 for ethanoic acid, vs} 1 for hydrochloric acid).

The buffer region is the area of gradual pH rise before the equivalence point. It forms because the added NaOH partially neutralises the weak acid, producing its conjugate base. The resulting weak acid / conjugate base mixture resists pH change.

Read off the pH at the half-equivalence point on the titration curve. At this point, pKa = pH, because [weak acid] = [conjugate base] and the log term in the Henderson-Hasselbalch equation equals zero.

At the half-equivalence point of a weak acid-strong base titration, pH = pKa because [weak acid] = [conjugate base].

True.

Phenolphthalein changes colour between pH 8.2 and 10.0. This range falls within the rapid pH change at the equivalence point (above pH 7) of a weak acid-strong base titration.

NaOH (aq) + CH₃COOH (aq) → CH₃COONa (aq) + H₂O (l)

The strong base neutralises the weak acid to form a salt and water.