Rates of Reaction (Edexcel IGCSE Chemistry)

The copper(II) carbonate in the mineral, malachite, reacts with hydrochloric acid according to this equation.

CuCO₃ (s) + 2HCl (aq) → CuCl₂ (aq) + H₂O (l) + CO₂ (g)

Some students investigate the effect of changing the concentration of acid on the rate of this reaction. The diagram shows the apparatus they use.

This is the method they use:

• set the balance to zero

• add an excess of malachite lumps to the conical flask and replace the cotton wool

• start a timer and record the balance reading after one minute

The experiment is repeated using different concentrations of hydrochloric acid. The mass and number of malachite lumps are kept the same in each experiment.

The table shows the results obtained in one series of experiments.

State why the balance readings have negative values.

The graph shows the results of this series of experiments.

The circled point indicates an anomalous result.

i) Suggest one mistake the students could have made to produce this result.

(1) 

ii) State the relationship shown by the graph.

(1)

Explain why an increase in the concentration of the acid causes an increase in the rate of the reaction. You should use the particle collision theory in your answer.

A student uses this apparatus to investigate the rate of reaction between marble chips and dilute hydrochloric acid.

i)

Explain how using the cotton wool increases the accuracy of this investigation.

(2)

ii)

Why is mass lost from the flask?

(1)

This is the equation for the reaction between marble chips and dilute hydrochloric acid. Complete the equation by adding the state symbols.

CaCO₃ (.........) + 2HCl (.........) → CaCl₂ (.........) + H₂O (.........) + CO₂ (.........)

The student uses large marble chips in the investigation. This is a graph of his results.

The student repeats the experiment using the same total mass of smaller marble chips. On the graph, draw the curve that would be obtained. [assume the marble chips are in excess]

The rate of this reaction can be altered by increasing the temperature or by increasing the concentration of the hydrochloric acid.

i) Explain, using the particle collision theory, how increasing the concentration of the hydrochloric acid would affect the rate of this reaction.

(3)

ii) Explain, using the particle collision theory, how increasing the temperature would affect the rate of this reaction.

(3)

A solution of hydrogen peroxide decomposes when a catalyst of manganese(IV) oxide is added.

The products of the reaction are water and oxygen.

Complete the chemical equation for this reaction.

............ H₂O₂ → ........... H₂O + ........... O₂

Give a test for oxygen.

State the reason for adding a catalyst.

A student investigates how changing the concentration of the hydrogen peroxide solution affects the rate of this reaction.

She uses this apparatus.

The student records the volume of oxygen that collects every 2 minutes for 16 minutes.

The table shows her results.

i) Plot the student’s results on the grid.

(1)

ii) Draw a circle on the grid around the anomalous result.

(1)

iii) Draw a curve of best fit through the points, ignoring the anomalous result.

(1)

iv) Suggest a mistake that the student might have made to cause the anomalous result.

(1)

v) Determine the volume of oxygen collected during the first 3 minutes. Show on your graph how you obtain your answer.

(2)

volume of oxygen = ............................................... cm³

The student repeats the experiment using hydrogen peroxide solution of half the concentration of the original solution. She keeps the volume of the hydrogen peroxide solution and all other conditions the same.

i) Draw on the grid the curve you would expect the student to obtain.

(2)

ii) Explain how using hydrogen peroxide solution of half the concentration affects the rate of the reaction.

Refer to particle collision theory in your answer.

(3)

Zinc reacts with dilute hydrochloric acid to form hydrogen.

 i) Give a chemical equation for this reaction.

(2)

ii) Give a test for hydrogen gas.

(1)

A student investigates the reaction between pieces of zinc and dilute hydrochloric acid.

In each experiment, he uses the same mass of zinc but a different volume of the acid.

He collects the hydrogen and measures its volume in each experiment.

The graph shows the student’s results.

i) Use the graph to find the minimum volume of acid needed to react with all of the zinc.

(1)

ii) The student repeats the investigation, using hydrochloric acid of double the original concentration.

Determine the volume of hydrogen that would be collected using 15 cm³ of this acid.

Show your working on the graph.

(2)

volume = ............................................... cm³

Explain how increasing the concentration of the hydrochloric acid affects the rate of reaction.

The rate of reaction could also be affected by changing the temperature of the hydrochloric acid, or by using a catalyst. Explain one other way in which the rate of reaction between zinc and hydrochloric acid can be affected.

Hydrogen peroxide solution decomposes to give water and oxygen gas.

The equation for this reaction is:

2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)

Three different solids are catalysts for the decomposition of hydrogen peroxide solution.

 A student is given hydrogen peroxide solution and a sample of each of the solid catalysts.

 The student has a timer, a measuring cylinder, a balance and the apparatus shown in the diagram.

 Describe a method the student could use to find which of the three solids is the most effective catalyst for the decomposition of hydrogen peroxide solution. 

The diagram shows the reaction profile for the decomposition of hydrogen peroxide without a catalyst.

 i) Label the diagram to show the activation energy (E_a) and the enthalpy change (ΔH) for this reaction.

 (2)

 ii) On the diagram, draw a curve to show the reaction profile for the same reaction when a catalyst is used.

 (1)

Hydrogen peroxide, H₂O₂, quickly decomposes to produce water and oxygen when a catalyst is added to the solution:

2H₂O₂ (aq) → 2H₂O (l)  + O₂ (g)

The volume of gas produced was measured and is shown in the graph below:

Which statement is not true for this reaction?

This question is about ethane and ethene.

Ethane can be obtained from crude oil. Describe the industrial process used to separate crude oil into fractions.

The equation for the reaction between ethene gas and hydrogen gas is

C₂H₄ + H₂ →  C₂H₆

i) The rate of this reaction can be increased by increasing the pressure. Explain why increasing the pressure increases the rate of this reaction.

(2)

ii) The rate of this reaction can also be increased by using a catalyst. Explain how using a catalyst increases the rate of a reaction.

(2)

iii) Give one other way that the rate of reaction between ethene gas and hydrogen gas can be increased.

(1)

iv) The reaction between ethene and hydrogen is exothermic. Complete the reaction profile diagram, including labels for the activation energy and the enthalpy change, ΔH.

(3)

The reaction between ethene and hydrogen can be represented using displayed formulae.

Use the bond energies in the table to calculate the enthalpy change, ΔH, in kJ/mol for this reaction.

ΔH = .............................................................. kJ/mol

A student uses this apparatus to investigate the rate of reaction between magnesium and an excess of dilute hydrochloric acid.

She uses this method.

• use a graduated beaker to pour 50 cm³ of dilute hydrochloric acid of concentration 2.00 mol / dm³ into the conical flask

• add a piece of magnesium ribbon of mass 0.086 g to the acid and put the bung into the neck of the flask

• measure the total volume of gas collected every ten seconds until the reaction stops

The table shows the student's results.

 i) Plot the student's results on the grid.

 (1)

 ii) Draw a curve of best fit 

 (1)

i) The student repeats the experiment using

• 0.043 g of magnesium ribbon

• 50 cm³ of 2.00 mol / dm³ hydrochloric acid

 Draw, on the grid in part (a), the curve you would expect in this experiment.

 Label this curve Y.

(2)

 ii) The student repeats the experiment again, using

• 0.086 g of magnesium ribbon

• 50 cm³ of 2.00 mol / dm³ hydrochloric acid

• a slightly higher temperature than the first experiment

 Draw, on the grid in part (a), the curve you would expect in this experiment.

 Label this curve Z.

 (2)

The expected volume of gas produced in the first experiment is 86 cm³.

Suggest why the volume collected is less than the expected volume.

The student uses a graduated beaker to measure the volume of dilute hydrochloric acid. Explain why it is not necessary to use a measuring cylinder in this experiment.

The ionic equation for the reaction between magnesium and hydrochloric acid is

Mg (s) + 2H⁺ (aq) → Mg²⁺ (aq) + H₂ (g)

Use the information in this equation, and the particle collision theory, to explain why the rate of reaction decreases during each of the experiments.

Dilute hydrochloric acid reacts with a solution of sodium thiosulfate (Na₂S₂O₃) to form a precipitate.

The equation for the reaction is

Na₂S₂O₃ (aq) + 2HCl (aq) → 2NaCl (aq) + S (s) + H₂O (l) + SO₂ (g)

State the name of the precipitate that forms.

The reaction is often used to investigate rates of reaction.

The diagram shows the apparatus a student uses to investigate the effect of temperature on the rate of the reaction.

 This is the student’s method.

• Pour 50 cm³ of cold sodium thiosulfate solution into a conical flask and heat it to 20 °C

• Draw a cross (X) on a piece of paper and place it under the flask

• Add 5 cm³ of dilute hydrochloric acid to the flask

• Look at the cross from above and record the time taken until the cross cannot be seen

 The student repeats the experiment four times, using sodium thiosulfate solution at a different temperature each time.

 He keeps the volumes of sodium thiosulfate solution and hydrochloric acid constant in each experiment.

 Give two other factors that the student should keep constant.

 1 ....................................................................................................

 2 ....................................................................................................

The table shows the student’s results.

 The highest temperature the student uses is 60 °C because he thinks the results might not be as accurate at temperatures higher than 60 °C.

 Suggest a reason why the results might not be as accurate at temperatures higher than 60 °C.

The student wants to compare the rates of the reaction at the different temperatures.

 He uses this formula to obtain a value for each rate of reaction

 rate = 

 The table shows the value of the rate of reaction at each temperature. 

 Plot the values of temperature and rate of reaction on the grid.

 Draw a curve of best fit through the points.

i) Use the graph to determine a value for the rate of the reaction at 45 °C.

 Show on the graph how you obtained your answer.

 (2)

 rate of reaction = .............................................................. s⁻¹

 ii) Calculate the time that it would take for the cross not to be seen at 45 °C.

 (2)

 time = .............................................................. s

 iii) Describe the relationship between rate of reaction and temperature shown by the graph.

 (1)

Explain, in terms of particle collision theory, the effect that increasing the temperature has on the rate of a reaction.

A student uses this apparatus to investigate the rate of reaction between marble chips and dilute hydrochloric acid.

The equation for the reaction is

CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂

During the reaction the mass of the contents of the flask decreases.

 i) State why the mass of the contents of the flask decreases.

 (1)

 ii) State the purpose of the cotton wool.

 (1)

 iii) Explain why sulfuric acid is not a suitable acid to use in this investigation.

 (2)

The graph shows the student’s results.

  i) In the investigation the marble chips are in excess.

 Explain the shape of the graph.

 (4)

 ii) The student repeats the experiment using the same volume of hydrochloric acid but of half the concentration of the original acid. All other conditions are kept the same.

 On the grid, draw the curve the student would obtain.

 (2)

Explain, using particle collision theory, how increasing the temperature affects the rate of a reaction.

This question is about reactions that form gases.

Hydrogen peroxide decomposes to form water and oxygen. The equation for the reaction is

2H₂O₂ → 2H₂O + O₂

25.0 cm³ of hydrogen peroxide solution are poured into a conical flask and 1.00 g of solid manganese(IV) oxide is added. Bubbles of oxygen gas are formed.

i) Give the test for oxygen gas.

(1)

ii) Describe a method to show that solid manganese(IV) oxide is a catalyst in this reaction and not a reactant.

(3)

A student uses this apparatus to investigate the rate of the reaction between zinc and an excess of dilute hydrochloric acid.

This is the student’s method.

• pour 50 cm³ of dilute hydrochloric acid into a conical flask

• add about 1.2 g of zinc lumps

• record the volume of hydrogen gas collected every 30 s until no more hydrogen is collected

The graph shows the student’s results.

i) Calculate the mean (average) rate of reaction, in cm³ / s, in the first 120 s.

(2)

mean rate = .............................................................. cm³ / s

ii) The equation for the reaction between zinc and hydrochloric acid is

Zn (s) + 2HCl (aq) → ZnCl₂ (aq) + H₂ (g)

Use this equation and the particle collision theory to explain why the rate of reaction is greatest at the start of the reaction.

(3)

iii) The student repeats the experiment at a higher temperature but keeps all other conditions the same. On the grid, draw the curve you would expect to see in this experiment.

(2)

iv) Explain why the rate of reaction is greater if the same mass of zinc powder is used instead of zinc lumps. All other conditions are kept the same.

(2)

In another experiment, the student adds 0.55 g of zinc to a solution containing 2.50 × 10⁻² moles of hydrochloric acid. Use the equation to show that hydrochloric acid is in excess.

Zn (s) + 2HCl (aq) → ZnCl₂ (aq) + H₂ (g)

[A_r of Zn = 65]

A student uses this apparatus to investigate the rate of reaction between calcium carbonate chips and dilute hydrochloric acid.

Every 20 seconds the student records the reading on the balance.

Explain why using a cotton wool plug increases the accuracy of the student’s results.

Complete the equation for the reaction by adding the state symbols.

CaCO₃ (..........) + 2HCl (..........) → CaCl₂ (aq) + H₂O (..........) + CO₂ (..........)

The student uses the balance readings to find the decrease in mass of the flask and contents.

The graph shows the student’s results.

 i) Give a reason why there are some calcium carbonate chips remaining in the flask when the reaction stops.

 (1)

 ii) State how the student would know when the reaction has stopped.

 (1)

 iii) Use the graph to determine the amount, in moles, of carbon dioxide produced during the reaction.

 [M_r of CO₂ = 44]

 (2)

 amount = .............................. mol

 iv) Use the graph to calculate the rate of reaction, in grams per second, at time 60 seconds.

 Show your working on the graph.

 (3)

 rate of reaction = .............................. g / s

The student repeats the investigation by diluting the original hydrochloric acid. 

The student then determines the initial rate of reaction at different percentage concentrations of the original hydrochloric acid. 

The graph shows the student’s results.

 i) Describe the relationship between the initial rate of reaction and percentage concentration of the original hydrochloric acid.

 (2)

 ii) Explain why changing the concentration of hydrochloric acid has an effect on the initial rate of reaction.

 (2)

A student uses this apparatus in an experiment to study the rate of the reaction between zinc and dilute sulfuric acid.

This is the student’s method.

• Add a few zinc granules to a conical flask on a balance

• Add 100 cm³ of dilute sulfuric acid to the flask, start a timer and immediately record the mass of the flask and contents

• Record the mass of the flask and contents every minute until the mass remains constant

The mass of the flask and contents decreases because hydrogen gas is produced and leaves the flask.

The student uses the mass readings to calculate the total mass of hydrogen produced.

Complete the equation for the reaction by adding the state symbols. 

 Zn (................) + H₂SO₄ (................) → ZnSO₄ (................) + H₂ (................)

The table shows the student’s results.

 i) Plot the student’s results. The first three have been done for you.

 (1)

 ii) Draw a circle around the anomalous result.

 (1)

 iii) Draw a curve of best fit.

 (1)

 iv) Give a possible reason for the anomalous result.

 (1)

 v) Determine a more likely value for this result.

 (1)

i) Explain how the shape of the curve shows how the rate of the reaction changes as time increases.

 (2)

 ii) At the end of the experiment there is no zinc left in the flask.

 Give a conclusion the student could make from this observation.

 (1)

The student does another experiment using

• the same amount of similarly sized magnesium granules instead of zinc

• the same volume of sulfuric acid, but of a lower concentration

Explain why it is difficult to predict how the rate of reaction in this experiment compares with the rate of reaction in the first experiment.

Explain, in terms of particle collision theory, how increasing the temperature affects the rate of a reaction.

Hydrogen peroxide solution decomposes slowly at room temperature to form water and oxygen. The equation for the reaction is

2H₂O₂ → 2H₂O + O₂

A catalyst increases the rate of this reaction. State one other property of a catalyst.

A student has samples of three solids, X, Y and Z. The student uses this apparatus to find out which solids act as catalysts in the decomposition of hydrogen peroxide solution.

Describe the method that the student should use to find out which solids act as catalysts.

The decomposition of hydrogen peroxide solution is exothermic.

On the diagram, draw and label the reaction profiles for the reaction

• without a catalyst

• with a catalyst

The equation for the reaction can be shown using displayed formulae.

2H-O-O-H → 2H-O-H + O=O    ΔH = –204kJ

The table gives the bond energies for the two of the bonds

i)Use this information to calculate the total amount of energy needed to break all the bonds in two moles of H₂O₂.

(1)

energy needed = .............................................................. kJ

ii) Use this information to calculate the total amount of energy released when all the bonds in two moles of H₂O are formed.

(1)

energy released = .............................................................. kJ

iii) Use the value of ΔH and your answers for (i) and (ii) to calculate the bond energy, in kJ/mol, for the O=O bond.

(2)

bond energy = .............................................................. kJ/mol