Factors affecting the rate of reaction

A student investigated the rate of the reaction of sodium thiosulfate solution with hydrochloric acid.

When sodium thiosulfate solution reacts with hydrochloric acid, the mixture becomes cloudy.

Figure 6 shows the apparatus.

This is the method used.

1. Put 75 cm³ of sodium thiosulfate solution in a conical flask.

2. Draw a pencil cross on paper.

3. Put the conical flask on the pencil cross.

4. Add 15 cm³ of hydrochloric acid to the contents of the conical flask.

5. Swirl the conical flask to mix the contents and immediately start a timer.

6. Stop the timer when the pencil cross is no longer visible through the reaction mixture.

7. Repeat steps 1 to 6 using different concentrations of sodium thiosulfate solution.

Explain why a 50 cm³ conical flask is unsuitable to use in this method.

Name a piece of equipment suitable for measuring the volume of sodium thiosulfate solution.

The student measured the time taken for the pencil cross to be no longer visible for different concentrations of sodium thiosulfate solution.

Draw one line from each type of variable to the variable in this investigation.

Type of variable
Variable in this investigation

What effect will using a darker pencil cross have on the time taken for the cross to be no longer visible?

Tick (✓) one box.

Table 2 shows the results.

Which concentration of sodium thiosulfate solution had the highest rate of reaction?

Tick (✓) one box.

Increasing the concentration of sodium thiosulfate solution changes the rate of the reaction with hydrochloric acid.

Which two statements explain the effect of increasing the concentration?

Tick (✓) two boxes.

The effect on the time taken for the cross to disappear can also be investigated by:

changing the temperature of the hydrochloric acid

changing the concentration of the hydrochloric acid.

Complete the sentences.

Choose the answers from the box.

If the temperature of the hydrochloric acid is increased, the time taken for the cross

If the concentration of the hydrochloric acid is decreased, the time taken for the

A student investigated the rate of the reaction between zinc and sulfuric acid.

The equation for the reaction is

Figure 10 shows the apparatus.

This is the method used.

1. Pour 50 cm³ of sulfuric acid into the conical flask.

2. Add excess zinc to the conical flask.

3. Insert the stopper and start a timer.

4. Measure the volume of hydrogen collected in the 50 cm³ measuring cylinder every 20 seconds for 180 seconds.

Explain why the volume of hydrogen collected in the 50 cm³ measuring cylinder is less than the volume of hydrogen produced.

Figure 11 shows the volumes of hydrogen collected in the 50 cm³ measuring cylinder after 40 seconds and after 100 seconds.

Determine the number of moles of hydrogen collected between 40 seconds and 100 seconds.

The volume of one mole of any gas at room temperature and pressure is 24 dm³.

A different student investigated how the concentration of sulfuric acid affected the rate of the reaction.

The student did a different experiment using sulfuric acid of

The student calculated the number of moles of hydrogen collected after every 20 seconds.

Figure 12 shows the results.

Determine the rate of reaction at 45 seconds.

You should draw a tangent on Figure 12.

Give your answer in standard form.

Figure 13 shows the results for 0.40 mol/dm³ sulfuric acid.

The student repeated the experiment using 0.20 mol/dm³ sulfuric acid instead of 0.40 mol/dm³ sulfuric acid.

Excess zinc was used in each experiment.

Sketch a line on Figure 13 to show the results you would expect.

Explain how increasing the temperature would affect the rate of reaction between zinc and sulfuric acid.

This question is about oxygen.

Scientists think that there was little or no oxygen in the Earth’s early atmosphere.

Which planet today has an atmosphere that is similar to the Earth’s early atmosphere?

Tick (✓) one box.

Which is the approximate percentage of oxygen in the Earth’s atmosphere today?

Tick (✓) one box.

Which two of the following increased the percentage of oxygen in the Earth’s atmosphere?

Tick (✓) two boxes.

Some scientists think that 1100 million years ago the Earth’s atmosphere contained:

• 16% oxygen

• 4% carbon dioxide.

Complete Figure 1.

You should:

• complete the y-axis scale

• plot the percentage of oxygen in the Earth’s atmosphere 1100 million years ago.

Oxygen is produced when manganese dioxide is added to hydrogen peroxide solution.

The equation for the reaction is:

hydrogen peroxide → water + oxygen

A student investigated the effect of changing the temperature on the decomposition of hydrogen peroxide.

This is the method used.

1. Add 5 cm³ of hydrogen peroxide solution to three test tubes labelled A, B and C.

2. Place each test tube in a water bath at a different temperature.

3. Add 0.2 g of manganese dioxide to each test tube.

Figure 2 shows the results.

Which test tube contained hydrogen peroxide solution at the highest temperature?

Tick (✓) one box.

The student tested the gas produced.

What is used to prove that the gas is oxygen?

Tick (✓) one box.

Manganese dioxide does not appear in the chemical equation for this reaction.

Which is a correct statement about manganese dioxide in this reaction?

Tick (✓) one box.

A student investigated the rate of the reaction between zinc and sulfuric acid.

Hydrogen gas is produced during this reaction.

Figure 7 shows the apparatus.

This is the method used.

1. Add 50 cm³ of sulfuric acid to a conical flask.

2. Add 2.0 g of zinc to the conical flask.

3. Quickly put a stopper in the conical flask and start a timer.

4. Measure the time taken to collect 20 cm³ of gas.

5. Repeat steps 1 to 4 three more times.

Suggest why the stopper must be put in the conical flask as quickly as possible in step 3.

The student calculated the rate of the reaction for each trial.

Table 7 shows the results of the calculations.

Determine the mean time taken to collect 20 cm³ of gas.

Do not include any anomalous results.

Use the equation:

The student changed the investigation so that the mean time taken to collect 20 cm³ of gas was greater.

Which two changes would increase the mean time taken to collect 20 cm³ of gas?

Tick (✓) two boxes.

Hydrogen gas is produced during this reaction.

Describe the test for hydrogen gas.

Give the result of the test.

A student investigated the rate of the reaction between zinc and sulfuric acid. Hydrogen gas is produced during this reaction. Figure 1 shows the apparatus.

This is the method used.

1. Add 50 cm³ of sulfuric acid to a conical flask.

2. Add 2.0 g of zinc to the conical flask.

3. Quickly put a stopper in the conical flask and start a timer.

4. Measure the time taken to collect 20 cm³ of gas.

5. Repeat steps 1 to 4 three more times.

Suggest why the stopper must be put in the conical flask as quickly as possible in step 3.

The student calculated the rate of the reaction for each trial. Table 1 shows the results of the calculations.

Determine the mean time taken to collect 20 cm³ of gas. Do not include any anomalous results. Use the equation:

The student changed the investigation so that the mean time taken to collect 20 cm³ of gas was greater. Which two changes would increase the mean time taken to collect 20 cm³ of gas? Tick (✓) two boxes.

Hydrogen gas is produced during this reaction. Describe the test for hydrogen gas. Give the result of the test.

Manganese dioxide catalyses the decomposition of hydrogen peroxide solution. Oxygen and water are produced.

Explain how a manganese dioxide catalyst increases the rate of decomposition of hydrogen peroxide.

A student investigated the rate of this reaction. This is the method used.

1. Add 50 cm³ of 2.0 mol/dm³ hydrogen peroxide solution to a conical flask.

2. Add 1.0 g of manganese dioxide to the conical flask.

3. Place the conical flask on a balance and start a timer.

4. Record the total mass lost from the conical flask every 20 seconds for 180 seconds.

Explain why the mass of the conical flask and contents decreased.

Figure 4 shows the results for 50 cm³ of 2.0 mol/dm³ hydrogen peroxide solution and 1.0 g of manganese dioxide. A tangent to the line has been drawn at 75 seconds.

Determine the rate of reaction when the time was 75 seconds. Give your answer to 2 significant figures.

The results for 50 cm³ of 2.0 mol/dm³ hydrogen peroxide solution and 1.0 g of manganese dioxide are shown again on Figure 5.

The student repeated the investigation using 50 cm³ of 1.0 mol/dm³ hydrogen peroxide solution and 1.0 g of manganese dioxide. Sketch the expected results for 1.0 mol/dm³ hydrogen peroxide solution on Figure 5.

A student investigated the rate of reaction between sodium thiosulfate solution and hydrochloric acid. Figure 7 shows the apparatus used.

When hydrochloric acid is added to sodium thiosulfate solution, the mixture gradually becomes cloudy. A smaller percentage of light from the light source reaches the light sensor as the mixture becomes more cloudy.

This is the method used.

1. Measure 50 cm³ of sodium thiosulfate solution into the beaker.

2. Add 10 cm³ of hydrochloric acid to the sodium thiosulfate solution.

3. Immediately start a timer.

4. Record the percentage of light from the light source that reaches the light sensor every 20 seconds for 120 seconds.

Balance the equation for this reaction.

The mixture becomes cloudy because the sulfur produced is a solid. What is the state symbol for a solid? Tick (✓) one box.

The student monitored the cloudiness of the reaction mixture using a light sensor. What other piece of equipment could be used to monitor the cloudiness of the reaction mixture? Tick (✓) one box.

Figure 8 shows the results.

What happened to the rate of reaction between 40 and 60 seconds? Use Figure 8. Tick (✓) one box.

The student stopped taking measurements after 120 seconds because the percentage of light reaching the sensor stayed constant. Why did the percentage of light reaching the sensor stay constant? Tick (✓) one box.

The student repeated the experiment using sodium thiosulfate solution of a higher concentration. How would the line of best fit for sodium thiosulfate solution of a higher concentration compare with the line of best fit on Figure 8 ? Tick (✓) one box.

The student then investigated the effect of changing the temperature on the rate of reaction. The student used sodium thiosulfate solution and hydrochloric acid which had been kept in an ice bath. Which are two effects of using reactants kept in an ice bath rather than at room temperature? Tick (✓) two boxes.

This question is about the reaction between sodium thiosulfate solution and hydrochloric acid. When hydrochloric acid is added to sodium thiosulfate solution, the mixture gradually becomes cloudy. The equation for the reaction is:

Sulfur is produced in the reaction. Why does the mixture become cloudy?

A student investigated the effect of changing the concentration of sodium thiosulfate solution on the rate of the reaction. Figure 6 shows the apparatus used.

A smaller percentage of light from the light source reaches the light sensor as the mixture becomes more cloudy. This is the method used.

1. Measure 50 cm³ of 0.10 mol/dm³ sodium thiosulfate solution into the beaker.

2. Add 10 cm³ of hydrochloric acid to the sodium thiosulfate solution.

3. Immediately start a timer.

4. Record the percentage of light from the light source that reaches the light sensor every 20 seconds for 120 seconds.

5. Repeat steps 1 to 4 using 0.20 mol/dm³ sodium thiosulfate solution.

Figure 7 shows the results for 0.10 mol/dm³ sodium thiosulfate solution.

The percentage of light reaching the light sensor decreases by 1% when 7.1 × 10⁻⁵ moles of sulfur is produced. Determine the rate of reaction in mol/s for the production of sulfur at 30 seconds. You should draw a tangent on Figure 7.

Explain why the rate of reaction changes between 0 and 60 seconds. Answer in terms of concentration. Use Figure 7.

Figure 8 is a repeat of Figure 7.

Figure 8 shows the results for 0.10 mol/dm³ sodium thiosulfate solution. Sodium thiosulfate solution was in excess in the investigation.

The line of best fit on Figure 8 is horizontal between 80 and 120 seconds because the reaction stopped. Why did the reaction stop?

Sketch a line on Figure 8 to show the results you would predict for 0.20 mol/dm³ sodium thiosulfate solution.

The same student did the investigation again the next day. The student found that the same method produced different results for the percentage of light reaching the light sensor.

How could the student improve the method so that the same percentages of light reached the light sensor?

Tick (✓) one box.

The student improved the method so that similar results were obtained on different days. What name is given to similar results obtained on different days under the same conditions by the same student? Tick (✓) one box.

Figure 9 shows the volumes of:

• sodium thiosulfate solution of concentration 0.10 mol/dm³

• hydrochloric acid of concentration 0.05 mol/dm³

which completely react to produce different masses of sulfur.

Which expression represents the relationship between the volume (V) of sodium thiosulfate solution used and the mass (m) of sulfur produced? Use Figure 9. Tick (✓) one box.

Determine the simplest whole number ratio of the volumes of sodium thiosulfate solution: hydrochloric acid which completely react with each other. Use Figure 9.

This question is about reactions between gases.

When hydrogen gas is heated with iodine gas, hydrogen iodide gas is produced.

The equation for this reversible reaction is:

hydrogen + iodine hydrogen iodide

This reversible reaction reaches equilibrium in a sealed container.

How does the equation show that the reaction is reversible?

Which two statements are correct when the reaction reaches equilibrium?

Tick (✓) two boxes.

The initial mixture of hydrogen and iodine in the sealed container is purple.

Hydrogen iodide is colourless.

How will the colour of the mixture in the sealed container have changed when equilibrium is reached?

Tick (✓) one box.

The rate of reaction between gases is affected by changing the pressure.

Complete the sentences.

When the pressure of the reacting gases is increased,

This is because at higher pressures the distance

Give one other way of changing the rate of reaction between gases.

You should not refer to pressure in your answer.

A student investigated how a change in concentration affects the rate of the reaction between zinc powder and sulfuric acid. The equation for the reaction is:

This is the method used.

1. Pour 50 cm³ of sulfuric acid of concentration 0.05 mol/dm³ into a conical flask.

2. Add 0.2 g of zinc powder to the conical flask.

3. Put the stopper in the conical flask.

4. Measure the volume of gas collected every 30 seconds for 5 minutes.

5. Repeat steps 1 to 4 with sulfuric acid of concentration 0.10 mol/dm³

Figure 8 shows the apparatus used.

The student made an error in setting up the apparatus in Figure 8. What error did the student make?

The student corrected the error. Figure 9 shows the student’s results.

Explain why the lines of best fit on Figure 9 become horizontal.

How does Figure 9 show that zinc powder reacts more slowly with 0.05 mol/dm³ sulfuric acid than with 0.10 mol/dm³ sulfuric acid?

Determine the rate of the reaction for 0.05 mol/dm³ sulfuric acid at 80 seconds. Show your working on Figure 9. Give your answer to 2 significant figures.

The activation energy for the reaction between zinc and sulfuric acid is lowered if a solution containing metal ions is added. What is the most likely formula of the metal ions added? Tick (✓) one box.

This question is about the rate of the reaction between hydrochloric acid and calcium carbonate.

A student investigated the effect of changing the size of calcium carbonate lumps on the rate of this reaction.

This is the method used.

1. Pour hydrochloric acid into a conical flask up to the 50 cm³ line.

2. Add 10.0 g of small calcium carbonate lumps to the conical flask.

3. Attach a gas syringe to the conical flask.

4. Measure the volume of gas produced every 20 seconds for 100 seconds.

5. Repeat steps 1 to 4 using 10.0 g of large calcium carbonate lumps.

The student used the 50 cm³ line on the conical flask to measure the volume of hydrochloric acid.

Suggest a piece of equipment the student could use to make the measurement of volume more accurate.

Carbon dioxide gas is produced in the reaction between hydrochloric acid and calcium carbonate.

Which test is used to identify carbon dioxide gas?

Tick (✓) one box.

Table 2 shows the student’s results for large calcium carbonate lumps.

Figure 4 shows the student’s results for small calcium carbonate lumps.

Complete Figure 4.

You should:

• plot the data for large calcium carbonate lumps from Table 2 on Figure 4

• draw a line of best fit for large calcium carbonate lumps.

Determine the mean rate of reaction using small calcium carbonate lumps between 0 seconds and 60 seconds.

Use the equation:

Use Figure 4.

Describe what happens to the volume of gas collected using small calcium carbonate lumps:

• between 0 and 20 seconds

• between 80 and 100 seconds.

Use Figure 4.

Between 0 and 20 seconds

Between 80 and 100 seconds

The balance used to weigh 10.0 g of calcium carbonate lumps caused an error.

The balance always read 0.2 g before being used.

What type of error was caused by the balance?

Tick (✓) one box.

Figure 5 shows the dimensions of two cubes of calcium carbonate.

A cube of calcium carbonate has six faces.

Calculate the total surface area of the large cube of calcium carbonate.

Use Figure 5.

The large cube of calcium carbonate was divided into eight smaller cubes.

The eight smaller cubes have a greater total surface area than the one large cube.

Compare the rate of reaction when using the eight smaller cubes with the rate of reaction when using the large cube.

Complete the sentence.

Choose the answer from the box.

This question is about carboxylic acids. Carboxylic acids belong to a homologous series. Table 3 shows information about the first three carboxylic acids in this homologous series.

Complete Table 3.

Ethanoic acid ionises in water. The equation for the reaction is:

Explain how the equation shows that ethanoic acid is a weak acid.

A student adds a solution of ethanoic acid to zinc carbonate in an open flask on a balance. Explain what happens to the mass of the flask and its contents during the reaction.

The student compares the rates of the reaction of zinc carbonate with:

• 0.01 mol/dm³ methanoic acid

• 0.01 mol/dm³ ethanoic acid.

The rate of the reaction with methanoic acid is greater than the rate of the reaction with ethanoic acid. Explain why. You should refer to ions in your answer. Use Table 3.

Ethanoic acid reacts with ethanol to produce an ester.

Give the name of the ester produced when ethanoic acid reacts with ethanol.

Hexanedioic acid and ethanediol join together to produce a polyester. Ethanoic acid and ethanol join together in the same way to produce an ester. Which is the displayed structural formula of the ester produced when ethanoic acid reacts with ethanol? Tick (✓) one box.

This question is about the rate of the reaction between hydrochloric acid and calcium carbonate. A student investigated the effect of changing the size of calcium carbonate lumps on the rate of this reaction. This is the method used.

1. Pour 40 cm³ of hydrochloric acid into a conical flask.

2. Add 10.0 g of small calcium carbonate lumps to the conical flask.

3. Attach a gas syringe to the conical flask.

4. Measure the volume of gas produced every 30 seconds for 180 seconds.

5. Repeat steps 1 to 4 using 10.0 g of large calcium carbonate lumps.

The student calculated the number of moles of gas from each volume of gas measured. Table 4 shows the student’s results for large calcium carbonate lumps.

The student plotted the results for small calcium carbonate lumps on Figure 4.

Complete Figure 4. You should:

• plot the data for large calcium carbonate lumps from Table 4

• draw a line of best fit.

Determine the mean rate of reaction for small calcium carbonate lumps between 20 seconds and 105 seconds. Give the unit. Use Figure 4.

The student concluded that the large calcium carbonate lumps reacted more slowly than the small calcium carbonate lumps. How do the student’s results show that this conclusion is correct?

The difference in the rates of reaction of large lumps and of small lumps of calcium carbonate depends on the surface area to volume ratios of the lumps. Figure 5 shows a cube of calcium carbonate.

Calculate the surface area to volume ratio of the cube in Figure 5. Give your answer as the simplest whole number ratio.

A larger cube of calcium carbonate has sides of 5 cm Describe how the surface area to volume ratio of this larger cube differs from that of the cube shown in Figure 5.

Some students investigated the rate of decomposition of hydrogen peroxide, H₂O₂

The equation for the reaction is:

The catalyst for the reaction is manganese dioxide.

Describe a test to identify the gas produced in the reaction.

Give the result of the test.

Student A investigated the effect of the particle size of manganese dioxide on the rate of the reaction.

This is the method used.

1. Measure 25 cm³ of 0.3 mol/dm³ hydrogen peroxide solution into a conical flask.

2. Add a spatula of fine manganese dioxide powder to the conical flask.

3. Measure the volume of gas produced every minute for 10 minutes.

4. Repeat steps 1 to 3 with some coarse manganese dioxide lumps.

The method student A used did not give valid results.

What two improvements could student A make to the method to give valid results?

Tick (✓) two boxes.

Student B used a method which gave valid results.

Figure 9 shows student B ’s results.

Determine the mean rate of reaction in cm³/s between 2 and 4 minutes for coarse manganese dioxide lumps.

Give your answer to 2 significant figures.

Use data from Figure 9.

Hydrogen peroxide molecules must collide with manganese dioxide particles for catalysis to take place.

Student B repeated the experiment with coarse lumps of manganese dioxide.

Student B used the same volume of 0.2 mol/dm³ hydrogen peroxide instead of 0.3 mol/dm³ hydrogen peroxide.

Sketch on Figure 9 the curve you would expect to see.

Assume that the reaction is complete after 9 minutes.

The rate of reaction is different when manganese dioxide is used as a fine powder rather than coarse lumps.

Explain why.

You should answer in terms of collision theory.

Sodium thiosulfate solution reacts with dilute hydrochloric acid. The solution becomes cloudy as the reaction takes place.

The equation for the reaction is:

Explain why the solution becomes cloudy.

Plan an investigation to show how the concentration of the sodium thiosulfate solution affects the rate of the reaction with dilute hydrochloric acid. Your plan should give valid results.

A student investigated how temperature affects the rate of reaction between magnesium carbonate and dilute hydrochloric acid. This is the method used.

1. Heat hydrochloric acid to 30 ºC in a conical flask.

2. Add magnesium carbonate powder to the conical flask.

3. Measure the loss in mass of the flask and contents every 20 seconds for 140 seconds.

4. Repeat steps 1-3 with hydrochloric acid heated to 50 ºC

Explain why the contents of the conical flask lose mass.

Table 5 shows the student’s results for hydrochloric acid at 30 ºC

Plot the data from Table 5 on Figure 6. Draw a line of best fit.

Figure 7 shows the student’s results for hydrochloric acid at 50 ºC

Determine the rate of reaction at 50 °C when the loss of mass is 0.95 g Show your working on Figure 7. Give your answer to 2 significant figures.