1.4.7Development of Practical Skills in Chemistry

pH – Problem Solving

5 min read 5 parts Lesson 9 of 11

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Section 1 of 5Ways to Measure pH

Different tests give different levels of precision, so a solution is usually identified by combining several of them.

Universal indicator

A mixture of dyes giving a range of colours;
- it is quick but only gives an approximate whole-number pH by comparison with a chart.

pH probe

A meter gives readings to two decimal places, but it must be calibrated with buffers of known pH before use.

Phenolphthalein

A single indicator that is colourless in acid and pink in alkaline solution, so it only distinguishes alkalis from acids.

Data logger

A probe linked to a logger records pH continuously, which is useful during a titration or other changing system.

For a strong monoprotic acid, [H⁺] equals the acid concentration, so pH = −log₁₀[H⁺].

For a strong base, find [OH⁻] first, then use [H⁺] = Kw ÷ [OH⁻], where Kw = 1.00 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K.

For a weak acid, use Ka = [H⁺]² ÷ [HA], assuming negligible dissociation and that [H⁺] equals [A⁻].

Rearrange to [H⁺] = √(Ka × [HA]) and then take the negative logarithm to obtain pH.

Compare the calculated pH with the measured pH to decide which solution is which in an unknown set.

Identifying Unknown Solutions

Strong acid

Very low pH such as 1 for 0.1 mol dm⁻³ HCl, and no colour change with phenolphthalein.

Weak acid

Higher pH than a strong acid of the same concentration, for example about 2.9 for 0.10 mol dm⁻³ ethanoic acid.

Strong base

Buffer

Has a pH close to its pKa, for example near 4.76 for an equimolar ethanoic acid and ethanoate mixture.

Question Calculate the pH of 0.100 mol dm⁻³ ethanoic acid, Ka = 1.74 × 10⁻⁵ mol dm⁻³.

Assume [H⁺] = [CH₃COO⁻] and that [CH₃COOH] stays 0.100 mol dm⁻³.
[H⁺]² = Ka × [HA] = 1.74 × 10⁻⁵ × 0.100 = 1.74 × 10⁻⁶, so [H⁺] = 1.32 × 10⁻³ mol dm⁻³.
pH = −log₁₀(1.32 × 10⁻³) = 2.88.

Answer pH = 2.88.

pH The negative logarithm to base 10 of the hydrogen ion concentration, pH = −log₁₀[H⁺].
Kw The ionic product of water, [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K.
Ka The acid dissociation constant, Ka = [H⁺][A⁻] ÷ [HA].
Calibration Adjusting a pH probe against buffers of known pH so that its readings are reliable.

Using a probe straight from storage without calibrating it, which gives systematic errors in every reading.
Treating a weak acid as fully dissociated and equating [H⁺] to the acid concentration.
Reading a universal indicator colour as an exact pH instead of an approximate value.
Forgetting to rinse the probe in distilled water between solutions, contaminating the next reading.