Further Physical Chemistry Practicals (OCR A Level Chemistry)

The reaction of calcium carbonate with dilute hydrochloric acid results in the formation of a gas.

The effect of increasing the concentration of the acid on the rate of reaction can be determined by measuring the amount of gas produced at regular intervals for several minutes.

The following procedure was used to measure the rate of the reaction of CaCO₃ with HCl:

Procedure

1. 1. Support a gas syringe with a stand, boss and clamp
   2. Measure 50 cm³ of 0.050 mol dm^-3 dilute HCl in a beaker and then add it to a conical flask
   3. Add about 2.5 g of CaCO₃ to the flask using a spatula, immediately connect the gas syringe and start counting.
   4. Record the volume of gas produced every 20 seconds for 2 minutes.
   5. Repeat Steps 1-4 for different concentrations of HCl.

The results for the reaction of 2.5 g of CaCO₃ with 50 cm³ of 0.050 mol dm^-3 dilute HCl were recorded in the table below.

Plot a graph of time (s) on the x-axis and volume of CO₂ (cm³) on the y-axis. Draw a line of best fit on the graph.

Use the results and your graph to calculate the initial rate of reaction. Give your answer to 2 decimal places and state the correct units.

A student investigates the rate of reaction between copper sulfate solution and dilute sodium hydroxide solution.

75 cm³ of copper sulfate solution was poured into a 250 cm³ conical flask using a measuring cylinder.

The conical flask was then placed on a cross drawn on a piece of paper. 15 cm³ of dilute sodium hydroxide was added to the conical flask as shown.

The timer was started, and then was stopped when the cross could not be seen anymore. The student then recorded the time. 

The student investigates the effect of temperature on the rate of the reaction, described in part (a), using the following procedure.

Procedure

1. 1. Measure out copper sulfate using a beaker and pour it into a conical flask.
   2. Record the initial temperature of the copper sulfate in the conical flask using a thermometer.
   3. Place the flask with the solution onto a cross drawn on a piece of paper.
   4. Heat the sodium hydroxide to a specific temperature.
   5. Start the stopwatch and then add the dilute sodium hydroxide.
   6. Stop the stopwatch when the cross can no longer be seen.
   7. Repeat the experiment at different temperatures.

The student concludes that the rate of reaction increases with increasing temperature.

Explain whether the student’s conclusion is correct using the collision theory.

Peroxodisulafte(VI) ions, S₂O₈^2-, reacts with iodide ions, I^–, in solution to form sulfate ions, SO₄^2-. and iodine, I_2. 

S₂O₈^2- (aq) + 2I^– (aq) → 2SO₄^2- (aq) + I₂ (aq)

If the reaction is carried out in the presence of a fixed amount of aqueous thiosulfate ions, S₂O₃^2- (aq), the iodine produced is reduced back to iodine.

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

When all of the thiosulfate ions have been used up, the iodine remaining will react with starch solution to produce a blue black colour.

Using appropriate apparatus and the chemicals listed below, plan an experiment to find the order of this reaction with respect to iodide, I^– (aq). Your answer should quote suitable volumes of solutions for the experiment

Chemicals

• • 1.00 mol dm^-3 aqueous potassium iodide, KI (aq)
  • 0.040 mol dm^-3 aqueous dipotassium peroxodisulfate(VI), K₂S₂O₈ (aq)
  • 0.010 mol dm^-3 aqueous sodium thiosulfate, Na₂S₂O₃ (aq)
  • 1% aqueous starch solution
  • Distilled water

Explain how the results can be used to show that the reaction is first order with respect to iodide ions, I^– (aq).

In a separate experiment, a student reacts together:

• • 7.2 x 10^-3 mol dm^-3 I^– (aq)
  • 3.8 x 10^-3 mol dm^-3 S₂O₈^2- (aq)

The initial rate is 8.95 x 10^-3 mol dm^-3 s^-1. The reaction is first order with respect to both I^– and S₂O₈^2-.

Calculate the rate constant k for this reaction. State the units, if any.

The following pH curve was produced when an acid and an alkali were reacted together.

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. 

A student performs a titration with a weak acid and weak base to produce the following pH curve by titration.

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 limited amount of thiosulfate ion as shown.

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

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

Explain how the colour change in the indicator is caused.

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.

A second student completes the same experiment. Using their results, the second student plots a rate-concentration graph. 

Describe how the second student calculates the rate for their graph.

The graph of the second student is shown.

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

Evaluate the second student’s conclusion.

The reaction between methanoic acid and bromine water, in the presence of an acid catalyst, can be monitored using the following equipment.

Write a balanced chemical equation for the reaction between methanoic acid and bromine water.

Using the equipment shown in part (a), describe the steps required to collect the results and determine the order of reaction with respect to bromine.

Another method to determine the order of reaction with respect to bromine involves changing the concentration of the bromine solution and using a white tile.

State the dependent variable in this experiment and suggest how the use of a white tile may improve the reliability of the results.

The order of reaction with respect to bromine can be investigated experimentally by using a colorimeter to measure the change in concentration of the bromine solution over time. 

The results of such an investigation are shown.

Plot a graph of the results.

The reaction between sodium thiosulfate solution, Na₂S₂O₃, and hydrochloric acid is commonly investigated by measuring the time taken for a cross placed underneath the reaction mixture to disappear, as shown.

Write a balanced symbol equation, including state symbols, for the reaction between sodium thiosulfate and hydrochloric acid.

A group of students are asked to investigate the effect of hydrochloric acid concentration on the reaction between sodium thiosulfate solution and hydrochloric acid.

One student identifies the following control variables:

• Volume of chemicals used
• Size of conical flask
• Person judging the end of the reaction 

State the assumption that links the size of the conical flask to the amount of sulfur produced.

Another student identifies the production of sulfur dioxide as a potential chemical hazard in this reaction.

Explain two associated precautions that can be implemented with respect to the sulfur dioxide produced. You do not need to discuss the use of personal protective equipment or fume cupboards.

A third student suggests that the initial rate of the reaction between sodium thiosulfate and hydrochloric acid could be measured by collecting the sulfur dioxide produced in a 100 cm³ gas syringe. 

The student’s suggested method is as follows:

1. Place 100 cm³ of 0.075 mol dm^-3 of sodium thiosulfate solution into a conical flask
2. Add 10 cm³ of 3.0 mol dm^-3 hydrochloric acid into the conical flask
3. Connect the conical flask to the gas syringe with a delivery tube and start the timer
4. Record the volume of gas produced every 30 seconds until the volume of gas has remained constant for 2 minutes
5. Plot and use a graph of the results to estimate the initial rate of reaction

 Using your equation from part (a), justify whether or not this method would be effective in producing valid results in a school laboratory.

The reaction between potassium manganate(VII) solution and ethanedioic acid, in the presence of sulfuric acid, can be used to determine the activation energy of a reaction.

2KMnO₄ (aq) + 5H₂C₂O₄ (aq) + 3H₂SO₄ (aq) → 2MnSO₄ (aq) + K₂SO₄ (aq) + 8H₂O (l)
                        + 10CO₂ (g)

Identify any spectator ion(s) and the reducing agent in this reaction, showing all your reasoning.

The following table provides information about the three reactants in this experiment.

The method for this experiment is as follows. 

1. Measure 10.0 cm³ of each of the following into a separate boiling tube:
   • 0.1 mol dm^-3 potassium manganate(VII) solution, KMnO₄ (aq)
   • 2.0 mol dm^-3 sulfuric acid, H₂SO₄ (aq)
   • 0.25 mol dm^-3 ethanedioic acid, H₂C₂O₄ (aq)
2. Heat all three boiling tubes in a water bath
3. Place a test tube in a test tube rack
4. Measure and record the temperature of all three solutions
5. Add 0.5 cm³ of the potassium manganate(VII) solution to the test tube
6. Add 0.5 cm³ of the sulfuric acid to the test tube
7. Then add 1.0 cm³ of the ethanedioic acid to the test tube
8. Immediately start the stopwatch and swirl the mixture

Explain why a measuring cylinder can be used, in step 1, to measure the reagents into the boiling tubes but a pipette should be used, in steps 5, 6 and 7, to measure the reagents into the test tube.

The temperature of each reagent is recorded before the experiment is performed.  

Explain why three thermometers may be required, in step 4, to measure the temperature of each reagent.

The purpose of heating the ethanedioic acid, potassium manganate(VII) solution and sulfuric acid in a water bath is to provide the variety of temperatures required. The results from the various temperatures are used to calculate the activation energy of the reaction.

Explain why the temperature recorded, in step 4, for each experiment, is only an approximation.

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. 

Table 1

Figure 1

Plot the graphs on the graph paper shown in Figure 1. You should start the y-axis at pH 4.00. 

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

The student repeated the experiment using a pH probe. The pH probe was calibrated using buffer solutions of pH 5.00 and 9.00. 

Explain how a buffer solution of pH 5.00 and a volume of 100 cm³ could be made from 10 cm³ of 0.25 mol dm^-3 ethanoic acid and sodium ethanoate. Your answer should Include relevant calculations to determine the mass of sodium ethanoate required.