Rate of Reaction (Cambridge O Level Chemistry)

A length of magnesium ribbon was added to 50 cm³ of sulfuric acid, concentration 1.0 mol/dm³. The time taken for the magnesium to react was measured. The experiment was repeated with the same volume of different acids. In all these experiments, the acid was in excess and the same length of magnesium ribbon was used.

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.

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.

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.

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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.

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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.

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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^–

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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.

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.

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.

Photosynthesis is another example of a photochemical reaction. Glucose and more complex carbohydrates are made from carbon dioxide and water.

Complete the equation.

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.

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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.

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Reaction rate increases when concentration or temperature is increased.

Using the idea of reacting particles, explain why;

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The rate of a photochemical reaction is affected by light. A reaction, in plants, between carbon dioxide and water is photochemical.

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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.

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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.

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. 

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. 

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?

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?

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.

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

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?

Aqueous ammonium nitrite, NH₄NO₂(aq), decomposes when heated, as shown.

NH₄NO₂(aq) → N₂(g) + 2H₂O(l)

A 25.0 cm³ sample of 0.150 mol/dm³ NH₄NO₂(aq) is heated.

Calculate the maximum volume, in dm³, of nitrogen formed, measured at room temperature and pressure. 

The concentration of NH₄NO₂(aq) is decreased.

The temperature of the reaction remains constant.

State and explain how the rate of reaction changes

NH₄NO₂ contains the ammonium ion, NH₄⁺, and the nitrite ion.

A mixture of aqueous calcium hydroxide and NH₄NO₂(s) is warmed.

Calcium nitrite, water and a gas are formed. The gas turns damp red litmus paper blue. Construct the equation for this reaction. 

NH₄NO₂(aq) is added to a sample of aqueous potassium iodide.

A brown solution is formed.

Peroxodisulfate ions, S₂O₈^2–, react with iodide ions in aqueous solution.

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

Iodide ions are oxidised in this reaction.

State how the equation shows this.

Table 17.1 shows how the relative rate of this reaction changes when different concentrations of peroxodisulfate ions and iodide ions are used.

Table 17.1

Using the information in Table 17.1, describe how increasing the concentration of each of these ions affects the relative rate of reaction:

peroxodisulfate ions

iodide ions

Iron(III) ions, Fe³⁺, catalyse this reaction.

Explain how catalysts increase the rate of a reaction.

Three of the factors that can influence the rate of a chemical reaction are:

• physical state of the reactants
• light
• the presence of a catalyst

The first recorded dust explosion was in a flour mill in Italy in 1785. Flour contains carbohydrates. Explosions are very fast exothermic reactions.

The decomposition of silver(I) bromide is the basis of film photography. The equation for this decomposition is:

This reaction is photochemical.

A piece of white paper was coated with silver(I) bromide and the following experiment was carried out.

Explain the results.

The fermentation of glucose is catalysed by enzymes from yeast. Yeast is added to aqueous glucose, the solution starts to bubble and becomes cloudy as more yeast cells are formed.

C₆H₁₂O₆ (aq) → 2C₂H₅OH (aq) + 2CO₂ (g)

The reaction is exothermic.

Eventually the fermentation stops when the concentration of ethanol is about 12%.

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.

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

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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.

In the first experiment, the maximum volume of oxygen produced was 96 cm³ measured at r.t.p.

Calculate the concentration of the aqueous hydrogen peroxide in mol/dm³ by completing the following steps.

2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)

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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.

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Suggest an explanation why initially the reaction speed increases then above a certain temperature the speed decreases.

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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.

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Catalytic converters reduce the pollution from motor vehicles.

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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.

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The reaction can be examined quantitatively.

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2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)

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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.

The reactions between metals and acids are redox reactions.

Zn + 2H⁺ → Zn²⁺ + H₂

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The rate of reaction between a metal and an acid can be investigated using the apparatus shown below.

A piece of zinc foil was added to 50 cm³ of hydrochloric acid, of concentration 2.0 mol/dm³. The acid was in excess. The hydrogen evolved was collected in the gas syringe and its volume was measured every minute.

The results were plotted and labelled as graph 1.

The experiment was repeated to show that the reaction between zinc metal and hydrochloric acid is catalysed by copper.

A small volume of aqueous copper(II) chloride was added to the acid before the zinc was added.

The results of this experiment were plotted on the same grid and labelled as graph 2.

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If the first experiment was repeated using ethanoic acid, CH₃COOH, instead of hydrochloric acid, how and why would the graph be different from graph 1?

Calculate the maximum mass of zinc which will react with 50 cm³ of hydrochloric acid, of concentration 2.0 mol/dm³.

Zn + 2HCl → ZnCl₂ + H₂

Show your working.

Sodium chlorate(I) decomposes to form sodium chloride and oxygen. The rate of this reaction is very slow at room temperature provided the sodium chlorate(I) is stored in a dark bottle to prevent exposure to light.

2NaClO → 2NaCl + O₂

The rate of this decomposition can be studied using the following experiment.

Sodium chlorate(I) is placed in the flask and 0.2 g of copper(II) oxide is added. This catalyses the decomposition of the sodium chlorate(I) and the volume of oxygen collected is measured every minute. The results are plotted to give a graph of the type shown below.

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The ions present in aqueous sodium chloride are Na⁺ (aq), Cl^– (aq), H⁺(aq) and OH^– (aq).
The electrolysis of concentrated aqueous sodium chloride forms three products. They are hydrogen, chlorine and sodium hydroxide.

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Cl₂ + ...NaOH → ..................... + ..................... + .....................

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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.

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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. 

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

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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.

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

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The experiment was repeated using 1.0 mol/dm³ propanoic acid instead of 1.0 mol/dm³ hydrochloric acid. Propanoic acid is a weak acid.

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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.

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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) → ............ + ............ + ............

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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.

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

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Many organic compounds which contain a halogen have chloro, bromo or iodo in their name.

The following diagram shows the structure of 1-chloropropane.

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

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.