Chemical Change & Rate of Reaction (Cambridge (CIE) IGCSE Co-ordinated Sciences (Double Award): Chemistry): Exam Questions

Suggest two changes to experiment C which would increase the speed of the reaction and explain why the speed would increase. The volume of the acid, the concentration of the acid and the mass of magnesium used were kept the same.

change 1 ....................................................................................................

explanation ................................................................................................

change 2 ....................................................................................................

explanation .................................................................................................

Manganese is a transition element. It has more than one valency and the metal and its compounds are catalysts.

Predict three other properties of manganese that are typical of transition elements.

Extended tier only

It has several oxides, three of which are shown below. Manganese(II) oxide, which is basic. Manganese(III) oxide, which is amphoteric. Manganese(IV) oxide, which is acidic.

i) Complete the word equation.

manganese(II) oxide + hydrochloric acid → .....................  ..................... + .....................

[2]

ii) Which, if any, of these oxides will react with sodium hydroxide?

[1]

Aqueous hydrogen peroxide decomposes to form water and oxygen.

2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)

This reaction is catalysed by manganese(IV) oxide.

The following experiments were carried out to investigate the rate of this reaction.

A 0.1 g sample of manganese(IV) oxide was added to 20 cm³ of 0.2 mol / dm³ hydrogen peroxide solution. The volume of oxygen produced was measured every minute. The results of this experiment are shown on the graph.

i) How does the rate of reaction vary with time? Explain why the rate varies.

[3]

ii) The following experiment was carried out at the same temperature. 0.1 g of manganese(IV) oxide and 20 cm³ of 0.4 mol / dm³ hydrogen peroxide. Sketch the curve for this experiment on the same grid.

[2]

iii) How would the shape of the graph differ if only half the mass of catalyst had been used in these experiments?

[2]

The equation for the reaction between sodium thiosulfate and hydrochloric acid is given below.

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

The speed of this reaction was investigated using the following experiment. A beaker containing 50 cm³ of 0.2 mol/dm³ sodium thiosulfate was placed on a black cross.

5.0 cm³ of 2.0 mol/dm³ hydrochloric acid was added and the clock was started.

Initially, the cross was clearly visible. When the solution became cloudy and the cross could no longer be seen, the clock was stopped and the time recorded.

The experiment was repeated with 25 cm³ of 0.2 mol/dm³ sodium thiosulfate and 25 cm³ of water. Typical results for this experiment and a further two experiments are given in the table.

i) Explain why it is necessary to keep the total volume the same in all the experiments.

[2]

ii) Complete the table.

[1]

iii) How and why does the speed of the reaction vary from experiment 1 to 4?

[3]

Extended tier only

The idea of collisions between reacting particles is used to explain changes in the speed of reactions. Use this idea to explain the following results.

The rate of the reaction between iron and aqueous bromine can be investigated using the apparatus shown below.

A piece of iron was weighed and placed in the apparatus. It was removed at regular intervals and the clock was paused. The piece of iron was washed, dried, weighed and replaced. The clock was restarted. This was continued until the solution was colourless. The mass of iron was plotted against time. The graph shows the results obtained.

i) Suggest an explanation for the shape of the graph.

[3]

ii) Predict the shape of the graph if a similar piece of iron with a much rougher surface had been used. Explain your answer.

[2]

iii) Describe how you could find out if the rate of this reaction depended on the speed of stirring.

[2]

Extended tier only

Iron has two oxidation states +2 and +3. There are two possible equations for the redox reaction between iron and bromine.

Fe + Br₂ → Fe²⁺ + 2Br^– 2Fe + 3Br₂ → 2Fe³⁺ + 6Br^–

i) Indicate, on the first equation, the change which is oxidation. Give a reason for your choice.

[2]

ii) Which substance in the first equation is the reductant (reducing agent)?

[1]

Describe how you could test the solution to find out which ion, Fe²⁺ or Fe³⁺, is present.

Some of the factors that can determine the rate of a reaction are concentration, temperature and light intensity.

A small piece of calcium carbonate was added to an excess of hydrochloric acid. The time taken for the carbonate to react completely was measured.

CaCO₃ (s) + 2HCl (aq) → CaCl₂ (aq) + CO₂ (g) + H₂O (l)

The experiment was repeated at the same temperature, using pieces of calcium carbonate of the same size but with acid of a different concentration. In all the experiments an excess of acid was used. 

 i) Complete the table (assume the rate is proportional to both the acid concentration and the number of pieces of calcium carbonate).

[3]

 ii) Explain why the reaction rate would increase if the temperature was increased.

[2]

 iii) Explain why the rate of this reaction increases if the piece of carbonate is crushed to a powder.

[1]

Sodium chlorate(I) decomposes to form oxygen and sodium chloride. This is an example of a photochemical reaction. The rate of reaction depends on the intensity of the light.

2NaClO (aq) → 2NaCl (aq) + O₂ (g)

i) Describe how the rate of this reaction could be measured.

[2]

ii) How could you show that this reaction is photochemical?

[1]

The decomposition of hydrogen peroxide is catalysed by manganese(IV) oxide.

2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)

To 50 cm³ of aqueous hydrogen peroxide, 0.50 g of manganese(IV) oxide was added. The volume of oxygen formed was measured every 20 seconds. The average reaction rate was calculated for each 20 second interval.

Explain how the average reaction rate, 2.4 cm³/s, was calculated for the first 20 seconds.

Complete the table.

Explain why the average reaction rate decreases with time.

The experiment was repeated but 1.0 g of manganese(IV) oxide was added.

What effect, if any, would this have on the reaction rate and on the final volume of oxygen?

Give a reason for each answer.

i) Effect on rate.

[1]

ii) Reason

[2]

iii) Effect on final volume of oxygen.

[1]

iv) Reason

[2]

The rate of a reaction depends on the concentration of reactants, temperature and possibly a catalyst or light.

A piece of magnesium ribbon was added to 100 cm³ of 1.0 mol/dm³ hydrochloric acid. The hydrogen evolved was collected in a gas syringe and its volume was measured every 30 seconds.

In all the experiments mentioned in this question, the acid was in excess. The results were plotted to give a graph.

i) The experiment was repeated. Two pieces of magnesium ribbon were added to 100 cm³ of 1.0 mol/dm³ hydrochloric acid. Sketch this graph on the same grid and label it X.

[2]

ii) The experiment was repeated using one piece of magnesium ribbon and 100 cm³ of 1.0 mol/dm³ ethanoic acid. Describe how the shape of this graph would differ from the one given on the grid.

[2]

Extended tier only

Reaction rate increases when concentration or temperature is increased.

Using the idea of reacting particles, explain why;

i) Increasing concentration increases the reaction rate,

[2]

ii) Increasing temperature increases reaction rate.

[2]

Extended tier only

The following apparatus was used to measure the rate of the reaction between zinc and iodine.

The mass of the zinc plate was measured every minute until the reaction was complete.

Write an ionic equation for the redox reaction that occurred between zinc atoms and iodine molecules.

Describe how you could show by adding aqueous sodium hydroxide and aqueous ammonia that a solution contained zinc ions.

Result with sodium hydroxide .....................................................................................

Excess sodium hydroxide .....................................................................................

Result with aqueous ammonia .....................................................................................

Excess aqueous ammonia .....................................................................................

From the results of this experiment two graphs were plotted.

i) Which reagent iodine or zinc was in excess? Give a reason for your choice.

[1]

ii) Describe how the shape of graph 1 would change if 100 cm³ of 0.05 mol/dm³ iodine had been used.

[2]

iii) On graph 2, sketch the shape if the reaction had been carried out using 100 cm³ of 0.1 mol/dm³ iodine at 35 °C instead of at 25 °C.

[2]

A student investigates the progress of the reaction between dilute hydrochloric acid, HC, and an excess of large pieces of marble, CaCO₃, using the apparatus shown in Fig. 5.1.

Fig. 5.1

A graph of the volume of gas produced against time is shown in Fig. 5.2.

Fig. 5.2

State how the shape of the graph shows that the rate of reaction decreases as the reaction progresses.

Suggest why the rate of reaction decreases as the reaction progresses.

Deduce the time at which the reaction finishes.

The experiment is repeated using the same mass of smaller pieces of marble. All other conditions are kept the same. Draw a line on the grid in Fig. 5.2 to show the progress of the reaction using the smaller pieces of marble. 

Extended tier only

The original experiment is repeated at a higher temperature. All other conditions are kept the same. The resulting increase in rate of reaction can be explained in terms of activation energy and collisions between particles.

Define the term activation energy. 

Extended tier only

Explain why the rate of a reaction increases when temperature increases, in terms of activation energy and collisions between particles.

Sulfur dioxide, SO₂ , is used in the manufacture of sulfuric acid.

Why is a catalyst used?

Extended tier only

Explain, in terms of particles, why a high temperature increases the rate of this reaction.

Oxygen is produced by the decomposition of hydrogen peroxide. Manganese(IV) oxide is the catalyst for this reaction.

What is meant by the term catalyst?

Extended tier only

A student measures the volume of oxygen produced at regular time intervals using the apparatus shown. Large lumps of manganese(IV) oxide are used.

A graph of the results is shown.

What happens to the rate of this reaction as time increases? In your answer, explain why the rate changes in this way.

The experiment is repeated using the same mass of manganese(IV) oxide.

Powdered manganese(IV) oxide is used instead of large lumps. All other conditions stay the same.

Sketch a graph on the axes in (b) to show how the volume of oxygen changes with time.

Extended tier only

In terms of particles, explain what happens to the rate of this reaction when the temperature is increased.

Extended tier only

Ammonia is manufactured by the Haber process.

Explain, in terms of particles, what happens to the rate of this reaction when the temperature is increased.

A student investigates the rate of reaction of small pieces of calcium carbonate with an excess of hydrochloric acid of concentration 1mol/dm³.

Name the salt formed when calcium carbonate reacts with hydrochloric acid.

The graph shows how the mass of the reaction mixture changes with time.

State why the reaction mixture decreases in mass.

Calculate the loss in mass during the first 40 seconds of the experiment.

The experiment is repeated using hydrochloric acid of concentration 2 mol/dm³ . 

All other conditions are kept the same.

Draw a line on the grid for the experiment using hydrochloric acid of concentration 2 mol/dm³.

In the experiment, when 2.00 g of calcium carbonate is used, the loss in mass of the reaction mixture is 0.88 g. All other conditions are kept the same.

Calculate the loss in mass when 0.50 g of calcium carbonate is used.

The experiment is repeated using the same mass of different-sized pieces of calcium carbonate. All other conditions are kept the same. The sizes of the pieces of calcium carbonate are:

• powder

• small pieces

• large pieces.

Complete the table by writing the sizes of the pieces of calcium carbonate in the first column.

The graph shows how the volume of hydrogen produced changes with time.

Describe how the rate of reaction changes with time. Use the graph to explain your answer.

How many seconds did it take to collect the first 25 cm³ of hydrogen?

The experiment is repeated at a higher temperature. All other conditions are kept the same.

Draw a line on the grid for the experiment using a higher temperature.

If 2.4 g of magnesium is used, 0.2 g of hydrogen is produced. Calculate the mass of magnesium needed to produce 0.8 g of hydrogen using an excess of dilute hydrochloric acid.

mass of magnesium = .............................. g

Oxides of nitrogen are pollutants in the air.

Oxides of nitrogen act as catalysts.

What is meant by the term catalyst?

Hydrogen peroxide decomposes to form water and oxygen. This reaction is catalysed by manganese(IV) oxide.

2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)

The rate of this reaction can be investigated using the following apparatus.

40 cm³ of aqueous hydrogen peroxide was put in the flask and 0.1 g of small lumps of manganese(IV) oxide was added. The volume of oxygen collected was measured every 30 seconds. The results were plotted to give the graph shown below.

i) How do the rates at times t₁, t₂ and t₃ differ?

[2]

ii) Explain the trend in reaction rate that you described in a) i).

[2]

The experiment was repeated using 0.1 g of finely powdered manganese(IV) oxide. All the other variables were kept the same.

i) On the same axes as the graph in part a), sketch the graph that would be expected for this change.

[2]

ii) Explain the shape of this graph.

[2]

Describe how you could show that the catalyst, manganese(IV) oxide, was not used up in the reaction.

Manganese(IV) oxide is insoluble in water.

Extended tier only

The speed (rate) of a chemical reaction depends on a number of factors, including temperature and the presence of a catalyst.

Reaction speed increases as the temperature increases.

Explain why reaction speed increases with temperature.

Extended tier only

An organic compound decomposes to give off nitrogen.

C₆H₅N₂Cl (aq) → C₆H₅Cl (l) + N₂ (g)

The speed of this reaction can be determined by measuring the volume of nitrogen formed at regular intervals. Typical results are shown in the graph below.

The reaction is catalysed by copper.

i) Sketch the graph for the catalysed reaction on the diagram above.

[2]

ii) How does the speed of this reaction vary with time?

[1]

iii) Why does the speed of reaction vary with time?

[3]

Extended tier only

Catalytic converters reduce the pollution from motor vehicles.

i) Describe how carbon monoxide and the oxides of nitrogen are formed in car engines. 

[4]

ii) Describe the reaction(s) inside the catalytic converter which change these pollutants into less harmful gases.

Include at least one equation in your description.

[3]

Hydrogen peroxide, H₂O₂, decomposes into water and oxygen in the presence of a catalyst, manganese(IV) oxide.

2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)

What is meant by the term catalyst?

A student studies the rate of decomposition of hydrogen peroxide using the apparatus shown.

The student uses 20 cm³ of 0.1 mol/dm³ hydrogen peroxide and 1.0 g of manganese(IV) oxide.

The student measures the volume of oxygen given off at regular time intervals until the reaction stops. A graph of the results is shown.

i) When is the rate of reaction highest?

[1]

ii) Suggest one method of increasing the rate of reaction using the same amounts of hydrogen peroxide and manganese(IV) oxide.

[1]

Extended tier only

The reaction can be examined quantitatively.

i) Calculate the number of moles of hydrogen peroxide used in this experiment.

[1]

ii) Use your answer to c) i) and the equation to calculate the number of moles of oxygen produced in the reaction.

2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)

[1]

iii) Calculate the volume (at r.t.p.) of oxygen produced in dm³

[1]

iv) What would be the effect on the volume of oxygen produced if the mass of the catalyst was increased?

[1]

v) Deduce the volume of oxygen, in dm³, that would be produced if 20 cm³ of 0.2 mol/dm³ hydrogen peroxide was used instead of 20 cm³ of 0.1 mol/dm³ hydrogen peroxide

[1]

The student carries out a second experiment to investigate whether another substance, copper(II) oxide, is a better catalyst than manganese(IV) oxide.

Describe how the second experiment is carried out. You should state clearly how you would make sure that the catalyst is the only variable.

20.0 g of small lumps of calcium carbonate and 40 cm³ of hydrochloric acid, concentration 2.0 mol / dm³, were placed in a flask on a top pan balance. The mass of the flask and contents was recorded every minute.

The mass of carbon dioxide given off was plotted against time.

CaCO₃ (s) + 2HCl (aq) → CaCl₂ (aq) + H₂O (l) + CO₂ (g)

In all the experiments mentioned in this question, the calcium carbonate was in excess.

i) Explain how you could determine the mass of carbon dioxide given off in the first five minutes.

[1]

ii) Explain how the shape of the graph shows where the rate is fastest, where it is slowing down and where the rate is zero.

[2]

Sketch on the same graph, the line which would have been obtained if 20.0 g of small lumps of calcium carbonate and 80 cm³ of hydrochloric acid, concentration 1.0 mol / dm³, had been used. 

Extended tier only

Explain in terms of collisions between reacting particles each of the following.

i) The reaction rate would be slower if 20.0 g of larger lumps of calcium carbonate and 40 cm³ of hydrochloric acid, concentration 2.0 mol / dm³, were used.

[2]

ii) The reaction rate would be faster if the experiment was carried out at a higher temperature.

[2]

Extended tier only

Calculate the maximum mass of carbon dioxide given off when 20.0 g of small lumps of calcium carbonate react with 40 cm³ of hydrochloric acid, concentration 2.0 mol / dm³.

CaCO₃ (s) + 2HCl( aq) → CaCl₂ (aq) + H₂O (l) + CO₂ (g)

A diagram of the apparatus which could be used to investigate the rate of reaction between magnesium and an excess of an acid is drawn below.

The magnesium kept rising to the surface. In one experiment, this was prevented by twisting the magnesium around a piece of copper. In a second experiment, the magnesium was held down by a plastic net fastened to the beaker.

i) Suggest a reason why magnesium, which is denser than water, floated to the surface.

[1]

ii) Iron, zinc and copper have similar densities. Why was copper a better choice than iron or zinc to weigh down the magnesium?

[1]

The only difference between the two experiments was the method used to hold down the magnesium. The results are shown below. 

i) In which experiment did the magnesium react faster?

[1]

ii) Suggest a reason why the experiment chosen in part i) had the faster rate.

[1]

Extended tier only

Give two factors which would alter the rate of this reaction. For each factor explain why it alters the rate.

Factor ............................................................................................................................

Explanation ....................................................................................................................

Factor ............................................................................................................................

Explanation ....................................................................................................................

A small piece of marble, CaCO₃, was added to 5.0 cm³ of hydrochloric acid, concentration 1.0 mol / dm³, at 25°C. The time taken for the reaction to stop was measured. The experiment was repeated using 5.0 cm³ of different solutions of acids. The acid was in excess in all of the experiments.

Typical results are given in the table.

.

i) Explain why it is important that the pieces of marble are the same size and the same shape.

[2]

ii) How would you know when the reaction had stopped?

[1]

Extended tier only

The equation for the reaction in experiment 1 is:

CaCO₃ (s) + 2HCl (aq) → CaCl₂ (aq) + CO₂ (g) + H₂O (l)

Complete the following ionic equation.

CaCO₃ (s) + 2H⁺ (aq) → ............ + ............ + ............

Extended tier only

A small piece of marble, calcium carbonate, was added to 5 cm³ of hydrochloric acid at 25 °C. The time taken for the reaction to stop was measured.

CaCO₃ (s) + 2HCl (aq) → CaCl₂ (aq) + CO₂ (g) + H₂O (l)

Similar experiments were performed always using 5cm³ of hydrochloric acid.

Explain each of the following in terms of collisions between reacting particles.

i) Why is the rate in experiment 2 slower than in experiment 1?

[2]

ii) Why is the rate in experiment 3 faster than in experiment 1?

[2]

iii) Why is the rate in experiment 4 faster than in experiment 1?

[2]

An alternative method of measuring the rate of this reaction would be to measure the volume of carbon dioxide produced at regular intervals.

i) Sketch this graph.

[2]

ii) One piece of marble, 0.3 g, was added to 5 cm³ of hydrochloric acid, concentration 1.00 mol / dm³. Which reagent is in excess? Give a reason for your choice.

Mass of one mole of CaCO₃ = 100g
Number of moles of CaCO₃ =
Number of moles of HCl =
Reagent in excess is:
Reason:

[4]

iii) Use your answer to (ii) to calculate the maximum volume of carbon dioxide produced measured at r.t.p.

[1]

Extended tier only

Organic halides react with water to form an alcohol and a halide ion.

CH₃–CH₂–I + H₂O → CH₃–CH₂–OH + I^– 

i) Describe how you could show that the reaction mixture contained an iodide ion.

[2]

ii) Name the alcohol formed when 1-chloropropane reacts with water.

[1]

Extended tier only

The speed (rate) of reaction between an organic halide and water can be measured by the following method.

A mixture of 10 cm³ of aqueous silver nitrate and 10 cm³ of ethanol is warmed to 60 °C. Drops of the organic halide are added and the time taken for a precipitate to form is measured.

Silver ions react with the halide ions to form a precipitate of the silver halide.

Ag⁺ (aq) + X^– (aq) → AgX (s)

Typical results for four experiments, A, B, C and D, are given in the table.

i) Explain why it takes longer to produce a precipitate in experiment A than in B.

 [2]

ii) How does the order of reactivity of the organic halides compare with the order of reactivity of the halogens?

 [2]

iii) Explain why the time taken to produce a precipitate would increase if the experiments were repeated at 50 °C.

 [3]