A-Level Chemistry: Required Practicals and Practical Skills
A-level Chemistry required practicals and practical skills is the practical-skills strand AQA examines directly in the written papers, worth a substantial share of the overall marks, even though no single required practical is assessed as a written lab report. It covers what the required practicals actually investigate: making up a standard solution and carrying out an acid-base titration, measuring an enthalpy change by calorimetry, following a pH titration curve for a weak acid against a strong base, simple test-tube reactions to identify cations and anions, and the preparation and purification of an organic solid or liquid. On top of naming apparatus and technique, the exam questions test calculating and combining uncertainties, evaluating a method to suggest a genuine improvement, and processing titration and calorimetry data correctly.
Before you start
Make sure you're comfortable with these topics first:
Method
- Read a burette to the nearest 0.05 cm3 (half of the smallest graduation, usually 0.10 cm3), and remember that every titre is found from two readings (the initial and final burette readings), so the absolute uncertainty in a single titre is the sum of the uncertainties in these two readings, not just the uncertainty in one reading alone.
- When calculating a mean titre, use only concordant results, defined as titres within 0.10 cm3 of each other, and discard, rather than average in, any anomalous result before calculating the mean; state which titres were used and why.
- For a temperature change (delta T) measured in a calorimetry experiment, apply the same two-readings rule as for a titre: the absolute uncertainty in delta T is the sum of the absolute uncertainties in the initial and the final temperature readings, since delta T is found by subtracting one reading from another.
- Calculate percentage uncertainty using (absolute uncertainty / measured value) x 100, and combine percentage uncertainties by adding them whenever quantities are multiplied or divided together, such as when a concentration and a titre volume are combined to calculate moles.
- For test-tube tests on ions, learn the reagent, the observation and the specific colour for each: acidified barium chloride gives a white precipitate with a sulfate; silver nitrate, followed by ammonia to test solubility, gives a white, cream or yellow precipitate with a chloride, bromide or iodide respectively; dilute acid followed by limewater tests for a carbonate; and warming with NaOH tests for an ammonium ion, releasing ammonia gas that turns damp red litmus paper blue.
- For a pH titration curve, identify the equivalence point as the midpoint of the steep, near-vertical part of the curve, and choose an indicator whose colour-change range lies entirely within this steep part; for a weak acid and a strong base, the steep part lies entirely in the alkaline range, so phenolphthalein is suitable but methyl orange is not, since methyl orange would change colour gradually over a range of volumes rather than sharply at the equivalence point.
- When evaluating a practical method, give a specific weakness tied to its effect on the result, not a vague 'human error', followed by a specific, practicable improvement that addresses it, such as using a lidded, insulated calorimeter and extrapolating the cooling curve back to the moment of mixing to correct for heat loss, rather than simply reading the highest temperature recorded.
Worked example
A student reacts 25.0 cm3 of 2.00 mol/dm3 hydrochloric acid with 25.0 cm3 of 2.00 mol/dm3 sodium hydroxide solution in an insulated polystyrene cup, using a thermometer with a resolution of 1 degree C. The initial temperature of both solutions is 18.5 degrees C; the highest temperature reached after mixing is 25.0 degrees C. Assume the density of the mixture is 1.00 g/cm3 and its specific heat capacity is 4.18 J/(g degreesC), the same as water. (a) Calculate the temperature change, delta T, and its percentage uncertainty. (b) Calculate the enthalpy change of neutralisation, in kJ/mol.
- Calculate the temperature change: delta T = 25.0 - 18.5 = 6.5 degrees C.
- Find the absolute uncertainty in delta T: since delta T is found from two thermometer readings, each with an absolute uncertainty of half the resolution (0.5 degrees C), the uncertainty in delta T is the sum of the two: 0.5 + 0.5 = 1.0 degrees C.
- Calculate the percentage uncertainty in delta T: (1.0 / 6.5) x 100 = 15.4% (3 s.f.). This is a precision (resolution) issue, separate from any heat lost to the surroundings.
- Calculate the heat released, using the total mass of solution (25.0 + 25.0 = 50.0 g, since density = 1.00 g/cm3): Q = m x c x delta T = 50.0 x 4.18 x 6.5 = 1358.5 J = 1.36 kJ (3 s.f.).
- Calculate the moles of the limiting reagent, here exactly stoichiometric since equal volumes and equal concentrations of a 1:1 acid-base reaction are used: moles = concentration x volume = 2.00 x (25.0/1000) = 0.0500 mol.
- Final answer: (a) delta T = 6.5 degrees C, with a percentage uncertainty of 15.4%; (b) enthalpy change = -Q / moles = -1358.5 / 0.0500 = -27170 J/mol = -27.2 kJ/mol (3 s.f.). The negative sign shows the reaction is exothermic, consistent with the temperature rise; heat lost to the surroundings during the experiment, a systematic error distinct from the 15.4% resolution-based uncertainty, means this measured value is smaller in magnitude than it would be if no heat had been lost.
Practice questions
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Q1A burette is read to the nearest 0.05 cm3. State the absolute uncertainty in a single titre, given that a titre is calculated from two burette readings (an initial and a final reading).Show answer
Answer: 0.10 cm3: the uncertainty in each individual reading, 0.05 cm3, is added for the two readings used to find the titre (0.05 + 0.05 = 0.10 cm3).
Q2A titration gives a titre of 23.40 cm3, with an absolute uncertainty of 0.10 cm3. Calculate the percentage uncertainty in this titre.Show answer
Answer: (0.10 / 23.40) x 100 = 0.43% (2 s.f.).
Q3A student weighs out a solid on a balance that reads to 2 decimal places (resolution 0.01 g), recording a mass of 4.85 g. Calculate the percentage uncertainty in this mass.Show answer
Answer: Absolute uncertainty = half the resolution = 0.005 g. Percentage uncertainty = (0.005 / 4.85) x 100 = 0.10% (2 s.f.).
Q4Describe the test and positive result used to confirm the presence of sulfate ions in a solution.Show answer
Answer: Add dilute hydrochloric acid, then aqueous barium chloride; a white precipitate (of barium sulfate) confirms sulfate ions are present.
Q5A student adds aqueous silver nitrate to a solution, followed by dilute ammonia, then concentrated ammonia if needed. A cream precipitate forms that dissolves only in concentrated ammonia. Identify the halide ion present.Show answer
Answer: Bromide (Br-): the cream precipitate is silver bromide, which is insoluble in dilute ammonia but dissolves in concentrated ammonia (a white precipitate, silver chloride, would dissolve in dilute ammonia; a yellow precipitate, silver iodide, would not dissolve even in concentrated ammonia).
Q6Explain why phenolphthalein is a suitable indicator for a titration between a weak acid and a strong base, but methyl orange is not.Show answer
Answer: For a weak acid titrated with a strong base, the steep (near-vertical) part of the pH curve at the equivalence point lies entirely within phenolphthalein's colour-change range, in the alkaline region, so phenolphthalein changes colour sharply at the equivalence point; methyl orange's colour-change range lies in the acidic region, where this curve is not steep, so it would change colour gradually over a range of volumes rather than sharply at the equivalence point.
Q7State why a burette should be rinsed with the solution it is about to contain, rather than with distilled water only, before a titration.Show answer
Answer: Rinsing with distilled water only would leave a thin film of water inside the burette, diluting the titrant and making its concentration, and therefore the calculated titre and any result derived from it, inaccurate; rinsing with the solution itself removes this risk.
Q8In a titration, a student obtains titres of 24.05 cm3, 24.10 cm3, 24.08 cm3 and 24.12 cm3. Explain why none of these titres needs to be excluded before calculating the mean, and calculate this mean.Show answer
Answer: All four titres are within 0.10 cm3 of each other (the range is 24.12 - 24.05 = 0.07 cm3), so all four are concordant and none needs to be excluded. Mean = (24.05 + 24.10 + 24.08 + 24.12) / 4 = 96.35 / 4 = 24.09 cm3 (2 d.p.).
Exam-style questions
Written in the style of a A Level Science exam paper, with a full mark scheme.
A student standardises a solution of sodium hydroxide by titrating it against 25.0 cm3 of a 0.0500 mol/dm3 standard solution of hydrochloric acid, using phenolphthalein indicator. The titres obtained were: 24.60 cm3, 24.15 cm3, 24.12 cm3 and 24.18 cm3. (a) Identify the anomalous titre and calculate the mean of the concordant titres. (b) Calculate the concentration of the sodium hydroxide solution, in mol/dm3.
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A student measures the enthalpy of combustion of ethanol by burning it under a metal calorimeter containing a known mass of water, and recording the maximum temperature reached. The measured enthalpy of combustion obtained is considerably less exothermic (smaller in magnitude) than the accepted data book value. (a) Suggest two reasons, other than incomplete combustion, why the measured value differs from the data book value in this way. (b) Describe one specific improvement to the method that would reduce the effect of one of the reasons you have given, and explain how it works.
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A student prepares a sample of a solid organic product by recrystallisation and determines its melting point using a melting point apparatus, finding that the sample melts gradually over the range 128 to 133 degrees C. A data book gives the melting point of the pure compound as a sharp 135 degrees C. (a) State what the range, rather than a single sharp value, observed for the student's sample suggests about its purity. (b) Suggest one further step the student could take to improve the purity of their sample.
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