Collision theory & activation energy

Alkanes and alkenes are hydrocarbons.

Hydrocarbon molecules are made up of hydrogen and carbon atoms only.

An alkene has the formula C₃H₆

Which alkene has the formula C₃H₆?

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How many atoms are there in a molecule of C₃H₆?

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Which is a use of C₃H₆?

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The percentage (%) by mass of each element in a molecule of C₃H₆ is:

hydrogen 14%

carbon 86%.

Figure 1 shows the percentage by mass of hydrogen in a molecule of C₃H₆

Complete Figure 1.

You should:

• complete the y-axis scale

• plot the percentage by mass of carbon as a bar

• label the bar plotted for carbon.

Alkanes can be cracked to produce alkenes and smaller alkane molecules.

Describe a test for alkenes.

Give the result of the test.

C₁₀H₂₂ can be cracked to produce C₃H₆ and an alkane.

Complete the equation for the reaction.

Complete the sentence.

Choose the answer from the box.

ceramics fertilisers fuels

High temperatures between 450 °C and 750 °C are used in cracking reactions.

Which statement describes the molecules at 750 °C compared with the molecules at 450 °C?

Tick () one box.

How does the rate of reaction at 750 °C compare with the rate of reaction at 450 °C?

A student analysed three solutions, A, B and C.

Each solution contained one ionic compound.

The student added a few drops of sodium hydroxide solution to solution A.

Solution A produced a white precipitate.

Aluminium ions produce a white precipitate with a few drops of sodium hydroxide solution.

Name one other metal ion which produces a white precipitate with a few drops of sodium hydroxide solution.

Describe how the student could confirm the presence of aluminium ions in solution A

Give the result of the test.

The student tested solution B to identify the ions present.

The student did a flame test on solution B.

An orange-red flame was produced.

Write the symbol for the metal ion in solution B.

The student added dilute nitric acid followed by silver nitrate solution to solution B.

A white precipitate was produced.

Write the symbol for the non-metal ion in solution B.

The student added 1.0 mol/dm³ nitric acid to solution C.

Bubbles of gas were produced.

Describe a test on the bubbles of gas that confirms that solution C contained carbonate ions.

Give the result of the test.

The student repeated the test on solution C using 1.0 mol/dm³ ethanoic acid instead of 1.0 mol/dm³ nitric acid.

Bubbles of gas were produced more slowly using ethanoic acid.

Explain why.

Hydrogen peroxide decomposes into water and oxygen.

The equation for the reaction is:

A student investigated the effect of catalyst A and of catalyst B on the rate of this reaction.

Figure 6 shows the apparatus.

This is the method used.

1. Put the conical flask on the balance.

2. Add 50 cm³ of 1.5 mol/dm³ hydrogen peroxide solution to the conical flask.

3. Add 1.00 g of catalyst A to the conical flask.

4. Start a timer.

5. Record the loss in mass of the conical flask and contents every 30 seconds for 3 minutes.

6. Repeat steps 1 to 5 using catalyst B.

Explain why the conical flask and contents lost mass.

Figure 7 shows the reaction profile for the decomposition of hydrogen peroxide using:

catalyst A

catalyst B.

Explain why catalyst A and catalyst B give different rates of reaction for the decomposition of hydrogen peroxide.

Figure 8 shows the results for catalyst A.

Determine the rate of the reaction when the time was 60 seconds.

Give your answer in standard form.

Explain how the rate of decomposition of hydrogen peroxide depends on the temperature of the solution.

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 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.

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.

Some students investigated the rate of decomposition of hydrogen peroxide.

The equation for the reaction is:

hydrogen peroxide → water + oxygen

Complete the sentence.

Choose an answer from the box.

The students tested the gas produced to show that it was oxygen.

The students used.

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

This is the method used.

1. Measure 25 cm³ hydrogen peroxide solution into a conical flask.

2. Add some fine manganese dioxide powder to the conical flask.

3. Measure the volume of oxygen produced every 30 seconds for 10 minutes.

4. Repeat steps 1 to 3 two more times.

5. Repeat steps 1 to 4 with coarse manganese dioxide lumps.

The method student A used did not give repeatable results.

How could student A make the results repeatable?

Tick (✓) one box.

Student B used a method which gave repeatable results.

How could student B improve the accuracy of these results?

Tick (✓) one box.

Figure 5 shows student B ’s results for coarse manganese dioxide lumps.

Calculate the mean rate of reaction between 30 and 250 seconds for coarse manganese dioxide lumps.

Use Figure 5 and the equation:

Give your answer to 3 significant figures.

Volume of oxygen formed

Mean rate of reaction = cm³/s

Fine manganese dioxide powder produces a higher rate of reaction than coarse manganese dioxide lumps.

Sketch on Figure 5 the results you would expect for student B ’s experiment with fine manganese dioxide powder.

Hydrogen peroxide molecules collide with manganese dioxide particles during the reaction.

Why does fine manganese dioxide powder produce a higher rate of reaction than coarse manganese dioxide lumps?

Tick (✓) one box.

This question is about rate of reaction.

A student investigated the rate of the reaction between magnesium and dilute hydrochloric acid.

The equation for the reaction is:

Which state symbol in the equation for the reaction does not represent one of the three states of matter?

The student determined the rate of production of hydrogen gas.

What two pieces of measuring apparatus could the student use to find the rate of production of hydrogen gas?

Table 6 shows the results of the investigation.

Plot the data from Table 6 on Figure 8.

You should draw a line of best fit.

Give three conclusions that can be drawn about the rate of reaction between magnesium and dilute hydrochloric acid in this investigation.

Use data from Figure 8 and Table 6.

The student repeated the investigation using dilute hydrochloric acid at a higher temperature.

All the other variables were kept the same.

Which two statements are correct?

Tick (✓) two boxes.

This question is about rate of reaction.

A student investigated the rate of the reaction between magnesium and dilute hydrochloric acid.

The equation for the reaction is:

Which state symbol in the equation for the reaction does not represent one of the three states of matter?

The student determined the rate of production of hydrogen gas.

What two pieces of measuring apparatus could the student use to find the rate of production of hydrogen gas?

Table 2 shows the results of the investigation.

Plot the data from Table 2 on Figure 3.

You should draw a line of best fit.

Give three conclusions that can be drawn about the rate of reaction between magnesium and dilute hydrochloric acid in this investigation.

Use data from Figure 3 and Table 2.

The student repeated the investigation using dilute hydrochloric acid at a higher temperature.

All the other variables were kept the same.

Which two statements are correct?

Tick (✓) two boxes.

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.

This question is about methanol.

Methanol is broken down in the body during digestion. What type of substance acts as a catalyst in this process?

Tick one box.

In industry, methanol is produced by reacting carbon monoxide with hydrogen. The equation for the reaction is:

How many moles of carbon monoxide react completely with 4.0 10³ moles of hydrogen?

Tick one box.

The reaction is carried out at a temperature of 250 ºC and a pressure of 100 atmospheres. The forward reaction is exothermic. Explain what happens to the yield of methanol if a temperature higher than 250 ºC is used.

A pressure of 100 atmospheres is used instead of atmospheric pressure. The higher pressure gives a greater yield of methanol and an increased rate of reaction. Explain why.

A catalyst is used in the reaction to produce methanol from carbon monoxide and hydrogen.

Explain how a catalyst increases the rate of a reaction.

Suggest why a catalyst is used in this industrial process. Do not give answers in terms of increasing the rate of reaction.

Suggest the effect of using the catalyst on the equilibrium yield of methanol.