Physical Chemistry Practicals (OCR AS Chemistry A): Exam Questions

A student is asked to plan and carry out a practical to determine the molar mass of magnesium by reacting it with dilute hydrochloric acid at room temperature and pressure (RTP).

Describe how the student would carry out this experiment and the measurements they should record. In your answer, identify the apparatus the student should use.

Write an equation for this reaction. Include state symbols.

The student reacts 0.11 g of magnesium ribbon with an excess of 1.0 mol dm^-3 hydrochloric acid. The reaction produces 110 cm³ of hydrogen gas at room temperature and pressure (RTP).

Calculate the molar mass of magnesium to 1 decimal place. Show your working.

Another student carries out the same procedure using 0.11 g of magnesium powder but only records a volume of 92.0 cm³.

Suggest one reason why the student records a volume lower than expected. Explain your answer.

Lactic acid, C₃H₆O₃, is a chemical by-product of anaerobic respiration. Bacteria in yoghurt produce it, as do bacteria in our stomach. Bacteria in beer also ferment glucose, producing lactic acid by anaerobic respiration. The increase in lactic acid decreases the pH of the beer and gives it its characteristic sour taste.

A student devises an experiment to investigate the concentration of lactic acid in a sample of beer.

1. Transfer 15.0 cm³ of beer to a 250.0 cm³ volumetric flask using a pipette.

2. Make up the volume to 250.0 cm³ with distilled water.

3. Use a pipette to take a 25.0 cm³ portion and add it to a conical flask.

4. Swirl the conical flask and place on the bench.

5. Add a few drops of phenolphthalein indicator to the solution in the conical flask.

6. Titrate 25.0 cm³ portions of this solution with 0.0750 mol dm^-3 sodium hydroxide solution.

Complete the following table, stating how each improvement to the procedure could be achieved:

The student's incomplete results table is shown below:

i) Complete the student's results table.

[1]

ii) Calculate the mean titre to 2 decimal places.

[2]

The burette used by the student has an error of ±0.05 cm³. It was used to determine the initial reading, final reading, and the end point.

Calculate the maximum percentage uncertainty in using this burette, using your mean titre from part (b).

Another student is required to prepare 250.0 cm³ of an aqueous solution containing a known mass of lactic acid. The student is provided with a sample bottle containing the lactic acid.

Describe the method, including apparatus and practical details, that the student should use to prepare this solution.

The enthalpy of combustion of an unknown liquid fuel can be determined by carrying out a calorimetry experiment.

i) Draw a labelled diagram of the apparatus required to carry out this experiment.

[3]

ii) State one way to reduce uncertainties in this experiment.

[1]

A student used the following procedure to carry out a calorimetry experiment.

Procedure

1. Measure 100 cm³ of water into a calorimeter using a measuring cylinder.

2. Record the initial temperature of the water using a thermometer.

3. Heat the water using the flame from the burning fuel and record the final temperature.

4. Measure the final mass of the spirit burner with the liquid fuel.

The student has missed a step in the method. State the missing step and explain why it is necessary.

The student obtained the following results:

• Mass of water in calorimeter = 350 g

• Initial temperature of water = 9 °C

• Final temperature of water = 36 °C

• Mass of liquid fuel burned = 3.15 g

i) Calculate the energy transferred to the water, q, in joules.

[1]

ii) Calculate the amount, in mol, of liquid fuel burned. The molar mass of the liquid fuel is 63.2 g mol^-1.

[1]

iii) Calculate the enthalpy of combustion, ΔH_c, of the liquid fuel. Give your answer in kJ mol^-1.

[2]

The actual value of ΔH_c of the unknown liquid is greater in magnitude than the experimental value calculated above.

Give two reasons for this.

A student wanted to determine the formula of hydrated magnesium sulfate, MgSO₄·xH₂O.

The student used the following procedure to determine the water of crystallisation.

Procedure

1. Using a two-decimal place balance, record the mass of an empty crucible.

2. Add a small amount of hydrated magnesium sulfate to the crucible, re-weigh and record the mass.

3. Gently heat the crucible for 2 minutes.

4. Allow the crucible to cool, re-weigh and record the mass.

Suggest two modifications to the method that would reduce the percentage uncertainty in the mass of the water removed.

Explain how heating until a constant mass is reached will improve the accuracy of the student's results.

Two other students carry out the same experiment.

Student A uses 50.0 g of MgSO₄·xH₂O.

Student B uses 0.01 g of MgSO₄·xH₂O.

Explain why both of these masses are likely to lead to an inaccurate value of x.

The student's results are shown in the table below.

Calculate the value of x in MgSO₄·xH₂O. Show your working.

x = ...............

When copper oxides in air, it forms copper(II) carbonate basic, CuCO₃·Cu(OH)₂ (s), which is a dull green colour. The percentage by mass of CuCO₃ in a sample of copper(II) carbonate basic can be found by reacting copper(II) carbonate basic with acid and measuring the volume of carbon dioxide evolved.

A student put exactly 0.50 g of copper(II) carbonate basic to the flask and added 50.0 cm³ of 1.00 mol dm^-3 sulfuric acid, which was in excess in the flask. They replaced the bung and collected the gas using the apparatus shown above. They collected 48.0 cm³ of carbon dioxide gas in the 250 cm³ measuring cylinder, which had graduations every 2 cm³.

Another student repeated the experiment but used an inverted burette graduated to 0.1 cm³ instead of the measuring cylinder.

Which method of gas collection has the lower percentage uncertainty associated with it?

Assume the same volume of gas is collected using the inverted burette.

Without changing the apparatus used, explain how the student could reduce the percentage uncertainty in using the measuring cylinder to measure the volume of gas collected.

The student realises that the percentage by mass of copper(II) carbonate in copper(II) carbonate basic should be given to two significant figures.

i) Why should the results be given to two significant figures?

[2]

ii) Suggest how the student might vary the method, without altering the apparatus, which would give a result that could be accurately reported to three significant figures. Explain your answer.

[1]

Suggest two errors in the method that could cause the calculated value for percentage by mass of copper(II) carbonate in copper(II) carbonate basic to be inaccurate.

For each error, state and explain the effect on the calculated value of percentage by mass and how the method could be altered to reduce the inaccuracy in the value.

A student performed a titration to identify an unknown Group 1 metal carbonate.

The student prepared 250 cm³ of a standard solution using 3.00 g of the unknown metal carbonate. They then used 25.0 cm³ of this solution and titrated it with 0.250 mol dm^-3 of hydrochloric acid.

Describe how the student prepared the standard solution of the unknown metal carbonate.

The student records their measurements and chemicals used. They write the concentration of hydrochloric acid used as 0.25 mol dm^-3 instead of 0.250 mol dm^-3.

Explain why the two concentrations do not mean the same.

The student's results are shown in the table below.

Identify any errors in their results table.

The student used the following steps to perform the titration:

1. Measure 25.0 cm³ of the unknown metal carbonate solution using a volumetric pipette and transfer into a conical flask

2. Add 3 to 4 drops of methyl orange

3. Place the conical flask onto a white tile

4. Rinse the burette with distilled water

5. Using a funnel, fill the burette with 0.250 mol dm^-3 hydrochloric acid

6. Titrate the metal carbonate solution with the hydrochloric acid, constantly swirl until the solution turns from yellow to orange — this is the rough titration

7. Repeat the titration until two results are concordant, but add the hydrochloric acid drop-by-drop as the end point is approached

Identify any errors in the method.

State and explain their effect on the unknown metal's calculated relative atomic mass.

The enthalpy change for the following reaction is known as the enthalpy change of hydration of anhydrous copper(II) sulfate. It is an exothermic reaction.

CuSO₄ (s) + 5H₂O (l) → CuSO₄⋅5H₂O (s)

The enthalpy change cannot be measured directly. Suggest why not.

The enthalpy change of the hydration of copper(II) sulfate can be determined experimentally using Hess' law. This involves recording the temperature change that occurs when both anhydrous copper(II) sulfate and hydrated copper(II) sulfate are dissolved in an excess of water.

The equations for the reactions are shown below, where nH₂O represents an excess of water.

CuSO₄ (s) + nH₂O → CuSO₄ (aq)

CuSO₄·5H₂O (s) + nH₂O → CuSO₄ (aq)

Explain why an excess of water is used.

The method below is used to determine the temperature change when anhydrous and hydrated copper(II) sulfate is added to water. The measured temperatures are then used to determine the enthalpy change of the reactions.

1. Measure 50 cm³ of water and add to a polystyrene cup placed inside a 250 cm³ glass beaker

2. Accurately weigh a weighing bottle containing approximately 5.0 g of anhydrous copper(II) sulfate

3. Measure and record the temperature of the water

4. Add the anhydrous copper(II) sulfate to the water, and continually stir. Measure and record the highest or lowest temperature

5. Reweigh the weighing bottle

6. Repeat steps 1 to 5 using 7.8 g hydrated copper(II) sulfate

A student suggests that the experiment should be altered to record the mass of the empty polystyrene cup before adding the water and then again at the end of the reaction.

Discuss if this alteration could have any effect on the calculated enthalpy changes.

The reaction scheme for the formation of hydrated copper sulfate from its anhydrous form is shown below.

Use the reaction scheme, Hess' law and the data provided in the table below to calculate the enthalpy of hydration for anhydrous copper(II) sulfate.

A student investigated the enthalpy change of combustion of different liquid fuels using a spirit burner and the apparatus shown below.

They used the following method:

1. Weigh the spirit burner and methanol

2. Measure 150 cm³ of water and pour into the beaker

3. Record the initial temperature of the water

4. Place the spirit burner under the beaker and light

5. Use the thermometer to stir the water

6. After about 5 minutes, extinguish the flame and record the maximum temperature reached by the water

7. Reweigh the spirit burner and methanol

The student repeated the method using ethanol and propanol.

When analysing the results, the student noticed that when the mass of the three fuels burnt was the same, the resulting temperature increase was different. Suggest why.

The measured temperature increase is used to calculate the energy transferred in the experiment. The student thought that if they used the same initial temperature and final temperature in each experiment, the energy transferred in each experiment would be the same.

State the assumption that the student must make for this to be the case.

After extinguishing the flame after the final temperature was recorded, the student tidied their equipment away before reweighing the spirit burner.

State and explain the effect, if any, on the calculated value of the enthalpy change of combustion.

When performing the experiment using another fuel, the student accidentally spilt some of the water when transferring it to the beaker.

State and explain the effect, if any, on the calculated value of the enthalpy change of combustion.

This question is about hydrated aluminium sulfate, Al₂(SO₄)₃·xH₂O (s).

Describe how the amount of water of crystallisation, x, in the formula of hydrated aluminium sulfate can be determined experimentally.

In one experiment, the student did not remove all of the water of crystallisation.

Explain how their calculated value of the number of moles of water of crystallisation, x, would differ from the actual value of x.

In another experiment, the student heated the hydrated aluminium sulfate very strongly for 30 minutes.

i) Give a reason why this may cause their calculated value of x to be incorrect.

[1]

ii) State if their calculated value of x would be greater or smaller than the actual value of x. Explain your answer.

[2]

The student decided to start with a larger mass of hydrated aluminium sulfate as they said that this would reduce the percentage uncertainty in their calculated value of x.

i) Explain why this would reduce the percentage uncertainty.

[2]

ii) Give a reason why using a larger initial mass of hydrated aluminium sulfate may produce anomalous results.

[1]