Physical Chemistry Practicals (A Level only) (AQA A Level Chemistry)

A standard hydrogen electrode can be used to measure the standard electrode potential for aqueous iron (III) ions being reduced to aqueous iron (II) ions. This reaction can be used to provide electrical energy in a cell.

Draw a clearly labelled diagram to show the components, reagents, including their concentrations, in this Fe³⁺/Fe²⁺ electrode.

You do not need to include the standard hydrogen electrode or salt bridge in your diagram.

In order to obtain a pH curve, a student is provided with a conical flask containing 50.0 cm³ of a 0.500 mol dm^–3 ethanoic acid solution and a burette filled with 0.250 mol dm^–3 potassium hydroxide solution. They are also provided with a pH meter.

Briefly describe how the student would ensure that the reading on the pH meter was accurate.

Describe how the student would carry out the titration for this reaction in order to obtain results to plot a pH curve.

An experiment was carried out to determine the e.m.f of this cell.

Fe(s) |  Fe³⁺(aq)  ||  H⁺(aq) |  H₂(g)  |  Pt(s)

In a Daniell cell, zinc is a more reactive metal than copper and this type of cell can be used to provide electrical energy. The conventional representation for the Daniell cell is;

Zn(s) | Zn²⁺(aq)  ||  Cu²⁺(aq) | Cu(s)

When an acid and alkali reacted together, the following pH curve was produced.

Figure 1

A salt bridge is used to complete an electrochemical cell. This could be prepared using filter paper and potassium chloride solution.

A student performs a titration with a weak acid and weak base. The pH curve produced from the titration is shown below in Figure 2.

Figure 2

The equation for magnesium reacting with hydrochloric acid is below:

Mg + 2HCl → MgCl₂ + H₂

The rate of reaction can be determined by collecting the hydrogen formed and measuring its volume at regular intervals.

A general equation for a reaction is shown below;

V(aq) + W(aq) + X(aq) → Y(aq) + Z(aq)

V, W, X and Y are all colourless aqueous solutions and solution Z is a dark blue aqueous solution.

A reagent (A) reacts rapidly with Z. If a small amount of reagent A is added to the initial reaction mixture, it will react with any solution Z that is formed until all of the reagent A has been used up.

Explain how you could use a series of experiments to determine the order of this reaction with respect to V. In each experiment you should produce a measure of the initial rate of reaction, t = 0.

A mixture containing iodine and an acid catalyst reacts slowly with butanone. The equation for the reaction is shown.

CH₃COCH₂CH₃ + I₂ ⟶ CH₃COCH₂CH₂I + HI

The rate of this reaction can be determined by changing the initial concentration of butanone. At appropriate time intervals, small samples of the mixture are removed and each sample is immediately added to an excess of sodium hydrogen carbonate solution before being titrated with sodium thiosulfate solution.

State the purpose of adding each sample to an excess of sodium hydrogen carbonate. Explain your answer.

Iodide ions are oxidised to iodine in acidic conditions by hydrogen peroxide, as shown in the equation below.

H₂O₂(aq) + 2H⁺(aq) + 2I^–(aq) → I₂(aq) + 2H₂O(l)

In the reaction mixture, a large excess of both H₂O₂ and I^– is used at a fixed temperature.

Figure 1 shows a graph of the results of the reaction.

Figure 1

Figure 2 below shows the results of a titration of NaOH and an acid. pH was plotted against the volume of sodium hydroxide solution. An indicator shows the pH range over which it changes colour.

Figure 2

A standard hydrogen electrode can be used to determine the standard electrode potential of a Cu(s)/Cu²⁺(aq) electrode.

Draw a labelled diagram of the apparatus which would be used to determine the standard electrode potential of Cu(s)/Cu²⁺(aq).

State the necessary conditions for this cell, which would allow the standard electrode potential to be measured.

Figure 1 below shows four pH curves representing different titrations performed with aqueous solutions of a variety of acids and bases. The concentrations of each are
0.5 mol dm^-3.

Figure 1

Which of these pH curves W, X, Y and Z shows the addition of:

For this question you will need to use Figure 1 from part (b) and Table 1 below, which shows indicators that can be used for acid-base reactions.

Table 1

Table 2 shows the half equations and standard electrode potentials of some half cells.

Table 2

A student was asked to make up a standard solution of ethanedioic acid. The student used the method below to make up a solution of ethanedioic acid;

• 750 g of an impure sample of ethanedioic acid was weighed out into a weighing boat
• The contents of the weighing boat were tipped into a beaker and 100 cm³ of distilled water was added
• The ethanedioic acid does not dissolve easily in cold water so the beaker was heated gently until all of the solid ethanedioic acid had dissolved
• The solution was poured into a 250 cm³ graduated cylinder and made up to the mark with distilled water

The value of K_a for a weak acid, HA, at 25 °C is 2.54 × 10^–5 mol dm^–3.

A student needs to obtain a pH curve, and they are provided with a conical flask containing 25.0 cm³ of a 0.500 mol dm^–3 ethanoic acid solution and 0.500 mol dm^–3 lithium hydroxide solution. They are also provided with a calibrated pH meter.

A student made up a buffer solution made from a weak acid and its salt. They combined 150 cm³ of 1.00 mol dm^-3 HCOOH mixed with 100 cm³ 0.750 mol dm^-3 HCOONa.

Calculate the pH of the buffer solution they produced.

K_a(HCOOH) = 1.78 x 10^-4 mol dm^-3.

A student performed an experiment to determine the acid dissociation constant by reacting a strong base with a weak acid and drawing a pH curve from their results.

In order to do this the student wrote down their method before beginning the experiment

Step 1      Pour a fixed volume of acid in a beaker

Step 2      Add alkali in small portions from a burette

Step 3      Use a pH meter to record the pH after each portion of alkali was added

Suggest two improvements to the student’s method

The student’s data was recorded in Table 1. Plot the graphs on the graph paper shown in Figure 1. You should start the y-axis at pH 4.00.

Table 1

Figure 1

Using your graph and the data in Table 1 in part (b), calculate the value for K_a of this acid and state the units.

Suggest what change the student could make to the procedure in order to give a more precise end point.

A buffer solution is made from a mixture of propanoic acid and its salt. A small amount of potassium hydroxide is added to the buffer.

Write an equation, including state symbols, to show how this buffer can resist the change in pH.

A separate buffer is prepared by the following steps:

Step 1             50.00 cm³ of distilled water was added to a 100 cm³ beaker

Step 2             5.00 cm³ of 0.30 mol dm^-3 of ethanoic acid was added to the beaker

Step 3             0.30 g of sodium ethanoate was added to the beaker

Step 4             A calibrated pH probe was used to measure the pH of the solution

Step 5             Sodium ethanoate was added in small quantities while stirring the solution
                        with a glass rod

Step 6             The pH was recorded until it reached the desired pH

Step 7             This solution was then transferred into a 100 cm³ volumetric flask and
                        distilled water was added up to the mark

K_a ethanoic acid = 1.74 x 10^-5 mol dm^-3

Explain why it is important to use a calibrated pH probe and describe how they are calibrated

Calculate the pH of the solution after the addition of 0.30 g of sodium ethanoate. 

An example of a basic buffer solution contains ammonia and ammonium chloride.

Explain why a mixture of 100 cm³ of 0.16 mol dm^-3 ammonia and 100 cm³ of 0.2 mol dm^-3 ammonium chloride can act as a basic buffer. You do not need to calculate the pH of the buffer solution.

A student is asked to complete the following experiment to determine the entropy change for the vaporisation of water.

1. Using a 500 cm³ measuring cylinder (± 5.0 cm³), place 500 cm³ of water into a 2 kW kettle
2. Boil the kettle
3. Measure the volume of water that remains after the kettle is boiled
4. Re-boil the water for 100 seconds by keeping the automatic cut-off switch depressed
5. Measure the volume of water that remains after the kettle is re-boiled for 100 seconds

Another student completes the entropy of vaporisation practical outlined in part (a).

After re-boiling a 2.7 kW kettle for 60 seconds, the student determines that 120 cm³ of water has evaporated.

Use the student’s results to calculate the entropy of vaporisation for their experiment.

A student is given the following method to determine the value of an equilibrium constant by preparing the propyl ethanoate ester.

Flask 1:

1. 0 g of concentrated ethanoic acid and 6.2 g of ethanol are added to a conical flask

2. A drop of concentrated sulfuric acid is added and the flask is stoppered and shaken

Flask 2:

3. A second ‘blank’ flask is made up containing a drop of sulfuric acid and 20 cm³ of distilled water

4. The reaction mixtures are allowed to reach equilibrium over the course of a week

Analysis

5. The reaction mixtures are titrated against 0.2 mol dm^-3 sodium hydroxide solution using phenolphthalein as an indicator

The differences in the titres can then be used to calculate the equilibrium amount of ethanoic acid and the value of K_c.

Explain the changes that the student needs to make to the method in order to make propyl ethanoate. Your answer should include relevant calculations.

Using an adapted method from part (c), a student placed 0.15 moles of ethanoic acid and 0.15 moles of butan-2-ol in a conical flask with 1 cm³ of concentrated sulfuric acid. The flask was left for a week to reach equilibrium. 

After this time, 25.0 cm³ of the reaction was titrated against 1.0 mol dm^-3 sodium hydroxide solution using phenolphthalein as an indicator. The average titre was found to be 23.60 cm³.

Using the average titre, the student calculated that there was 2.36 x 10^-2 moles of ethanoic acid remaining at equilibrium.

The student then completed the following K_c calculation to determine the value of the equilibrium constant.

K_c =  = 3.49 x 10^-2

The actual value of the equilibrium constant for the student’s experiment should be 3.23 x 10^-2.

Identify the single error that the student made and its consequent effects.

A student is asked to investigate the effect of changing the concentration of iodide ions on the iodine clock reaction between hydrogen peroxide and iodide ions.

Hydrogen peroxide reacts with iodide ions to form iodine which then immediately reacts with the thiosulfate ion as shown.

H₂O₂ (aq) + 2H⁺ (aq) + 2I^– (aq) → I₂ (aq) + 2H₂O (l)

2S₂O₃^2– (aq) + I₂ (aq) → 2I^– (aq) + S₄O₆^2– (aq)

The amount of thiosulfate ions present is limited. When the iodine produced has reacted with all of the thiosulfate ions present, an excess of iodine remains which turns the starch a dark blue-black colour.

The student is given 100 cm³ of 0.10 mol dm^-3 potassium iodide solution.

Describe how they could produce a minimum of five different concentrations of potassium iodide solution. Your answer should consider possible control variables.

The results of another student for the experiment in part (a) are shown in Figure 1.

Figure 1

From the graph, the student concludes that the reaction is first order with respect to iodide ions.

Evaluate the student’s conclusion.

In the presence of an acid catalyst, methanoic acid reacts with bromine according to the following equation.

HCOOH (aq) + Br₂ (aq) → 2Br^- (aq) + 2H⁺ (aq) + CO₂ (g)

Describe the basic steps to determine the order of reaction with respect to bromine using colorimetry.

A student suggests that the order of reaction with respect to bromine for the reaction described in part (c) can be determined by measuring the volume of gas produced.

Outline an initial rates method that the student could use to determine the order of reaction with respect to bromine by measuring the volume of gas produced.

Table 1 shows some standard electrode potential data.

Table 1

Give the conventional representation, including state symbols, of the cell that is used to measure the electromotive force between Cu²⁺ (aq) ions and Zn²⁺ (aq) ions.

A student sets up the electrochemical cell shown in Figure 1.

Figure 1

The voltmeter reading that the student gets for this electrochemical cell is -0.59 V.

Explain the student’s error and propose a suitable correction. Use Table 1, from part (a), to help you answer.

Identify the ionic half equations in Table 1 that have the largest difference in reactivity.

Use this information to give the conventional representation, including state symbols, of this cell.

A student hypothesised that increasing the temperature of the Zn / Zn²⁺ and Cu / Cu²⁺ electrochemical cell would result in an increase in electromotive force.

The student used 1.00 mol dm^-3 solutions of zinc(II) sulfate solution and copper(II) sulfate solution connected using a potassium nitrate salt bridge and a high resistance voltmeter. 

They had planned to heat both solutions to 25 ^oC and record the electromotive force. Then they would increase the temperature by 5 ^oC for each subsequent experiment.

The student’s results are shown in Figure 2.

Figure 2

Suggest why the student was practically unable to record the electromotive force every 5^oC.