Kinetics (AQA A Level Chemistry)

Exam Questions

3 hours45 questions
1a1 mark

A student investigated the rate of a reaction between calcium carbonate and hydrochloric acid.

State the meaning of the term rate of reaction.

1b1 mark

The equation below shows the reaction the student is investigating.

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

State which of the following will not affect the rate of the reaction the student is investigating.

A - Surface Area

B - Temperature

C - Pressure 

D - Concentration

1c3 marks

Figure 1 shows the collection of carbon dioxide gas over time from the reaction mentioned in part (b). 

Figure 1

ZY39vjT3_1

i)
Describe the general trend as shown in Figure 1.

ii)
State and explain what has happened to the rate of the reaction at point A in Figure 1.
1d1 mark

The student repeats the experiment but this time at a fixed temperature. The student doubles the concentration of hydrochloric acid used instead.

State what would happen to the rate of the reaction if the concentration of acid is doubled.

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2a2 marks

The Maxwell-Boltzman distribution curve shown in Figure 1 shows the distribution of molecular energies in a sample of gas.

Figure 1

h9t9An5R_2

i)
Sketch a second curve on Figure 1 to show the distribution of molecular energies in this mixture at a higher temperature.

ii)
State what the area underneath the curve represents.
2b2 marks

Some reactions which are slow at room temperature involve the use of a catalyst to speed up the rate of the reaction.

i)
Suggest one reason why a solid catalyst for a gas phase reaction is often in powdered form.

ii)
State the name of the type of catalyst used in this reaction.
2c2 marks

A dynamic equilibrium is established when gas X is mixed with gas W at a given temperature.

X (g) + W (g) rightwards harpoon over leftwards harpoon for blank of Y (g) + Z (g)

State the meaning of the term dynamic equilibrium.

2d3 marks

Catalysts have both economic and sustainability importance when it comes to the chemical industry.

i)
State the effect, if any, of a catalyst on the activation energy of a reaction.

ii)
State the effect, if any, of a catalyst on the amount of products produced at equilibrium.

iii)
State the effect, if any, of a catalyst on the time taken for equilibrium to be established in a reaction. 

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3a2 marks

Figure 1 below shows the Maxwell-Boltzmann distribution of molecular energies for a catalysed reaction.

Figure 1

SE3~UKfY_3

i)
State what would happen to the shape of the curve if the temperature in this reaction was lowered. 

ii)
State which other factor would cause the activation energy, Ea, to shift to the left in this
3b2 marks

A solid catalyst, vanadium (V) oxide, V2O5, is used in industry to increase the rate of the production of sulfur trioxide, SO3, in this reaction.

2SO2 (g) + O2 (g) → 2SO3 (g)   ΔH = −196 kJ mol−1

i)
State and explain whether V2O5 is a homogeneous or heterogeneous catalyst.

ii)
The use of catalysts in industrial processes can be beneficial to the environment.
State one reason for this.
3c1 mark

Figure 2 below shows the Maxwell-Boltzmann distribution of molecular energies for a reaction at a fixed temperature.

Figure 2

36

State what the energy marked X in Figure 2 shows.

3d1 mark

In the reaction in part (c), the pressure of the original reaction mixture is doubled.

State the effect, if any, of this change on the value of X.

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4a2 marks

Figure 1 shows a sketch of the Maxwell-Boltzmann curve for the distribution of molecular energies for a gas at a given temperature.

Ea represents the activation energy of the uncatalysed reaction.

Figure 1

jdu8J7EF_5

i)
Draw a vertical line on Figure 1, to represent the activation energy of the catalysed reaction. 
Label this line Ea (with catalyst).

ii)
State the effect, if any, of increasing temperature on the activation energy of a reaction.
4b2 marks

This question concerns the Maxwell-Boltzmann energy distribution shown below in Figure 2.

Figure 2

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i)
State how does the curve in Figure 2 change when the temperature is increased.

ii)
State what would be the effect on Figure 2 if the reactant concentrations were increased.
4c2 marks

Figure 3 shows an incomplete sketch of the Maxwell-Boltzmann curve for the distribution of molecular energies for a gas at a given temperature.

Figure 3

1-3

i)
State why the curve starts at the origin.

ii)
Label the missing axis on Figure 3.
4d2 marks

Figure 4 shows the Maxwell–Boltzmann distribution shown for a sample of a gas, Z, at two different temperatures.

Figure 4

tw6R7qzR_8

i)
Label both curves, T1, for low temperature and T2, for high temperature.

ii)
State which line, A-D, represents the mean energy of the molecules at a low temperature.

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5a2 marks

Many chemical industries use catalysts to increase the rate of the reaction, as catalysts work by lowering the activation energy required.

Define the term ‘activation energy’

5b2 marks

Figure 1 shows a Maxwell-Boltzmann distribution of molecular energies for gaseous molecules.

Figure 1

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i)
State what area of the graph corresponds to the number of molecules with sufficient energy to react when a catalyst is present.
  1. Y
  2. Y - Z
  3. Y + Z
  4. Z
ii)
State what factor would always result in a decrease in the number of molecules contained within area Y.
5c1 mark

A student mixes hydrogen and iodine at room temperature and pressure and allows the mixture to reach dynamic equilibrium.

H2 (g) + I2 (g) rightwards harpoon over leftwards harpoon with blank on top 2 HI (g)    ΔH = -9 kJ mol-1

State why this reaction is slow at room temperature.

5d1 mark

The Maxwell-Boltzmann distribution of energies for a gas shown in Figure 2.

Figure 2

RJ0B7QD6_10

State what the shaded area of Figure 2 represents.

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1a2 marks

Sodium thiosulfate and hydrochloric acid will react together readily, as shown by the equation below:

Na2S2O3 + 2HCl  →  2NaCl + S + SO2 + H2O

This reaction is often referred to as the ‘disappearing cross’ experiment, where students use different concentrations of hydrochloric acid to determine the effect that concentration has on the rate of a chemical reaction. 

The cross disappears because the solution turns cloudy as the reaction progresses. 

State the type of chemical reaction taking place and identify the product responsible for turning the solution cloudy. 

1b3 marks

A student completed the disappearing cross experiment from part (a) and recorded the results shown in Table 1

Table 1

Expt. 

Concentration of Na2S2O3
(mol dm-3)

Concentration of HCl
(mol dm-3)

Time for X to disappear
(s)

1

0.050

0.25

113

2

0.050

0.50

99

3

0.050

1.0

68

4

0.050

1.5

43

Use the results shown in Table 1 and collision theory, to state and explain the effect that increasing the concentration of a reactant has on the rate of a chemical reaction. 

1c2 marks

The Maxwell-Boltzmann distribution of energies for a gas shown in Figure 1.

Figure 1

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A student stated that when you increase the temperature:

  • The curve will be lower and shift to the right.
  • The shaded area on the Maxwell-Boltzmann curve will be smaller.

Explain, with reasons, whether you agree with this students’ statement.

1d2 marks

Explain how catalysts increase the rate of a chemical reaction. 

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2a4 marks

In any chemical reaction, the particles will all be moving around in different directions, at different speeds, with different amounts of energy. 

A Maxwell-Boltzmann distribution is a graph which shows the distribution of energy amongst particles within a chemical reaction. 

Figure 1 below shows the Maxwell-Boltzmann distribution in a sample of a gas at a fixed temperature, T1

Figure 1

1-3

i)
Label the x and y axes of the graph. 

ii)
Sketch a distribution for this same sample of gas, at a higher temperature, T2
2b2 marks

State why a Maxwell-Boltzmann distribution curve always starts at the origin and what the area under the curve represents. 

2c3 marks

Chemical reactions take place at different speeds. For a chemical reaction to take place, particles must collide with each other in the correct orientation and with sufficient energy. 

i)
Explain why most collisions between particles in the gas phase do not result in a reaction taking place.

ii)
State and explain one way that the rate of reaction could be increased, other than by increasing the temperature. 
2d1 mark

Give one reason why a reaction may be slow at room temperature.

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3a1 mark

State the meaning of the term rate of reaction.

3b3 marks

A group of students were completing a practical, investigating the factors which affect the rate of the chemical reaction shown below. 

A (s) + B (aq) → C (g)

The students collected the gas produced and plotted the graph shown in Figure 1

Figure 1

2-3

i)
State and explain what the letter R represents on the students graph in Figure 1

ii)
In the original reaction above, the students used 0.5 g of A and 50 cm3 of 1.0 mol dm-3 B.
Sketch a curve on the graph to show how the total volume of gas collected would change if the students still used 0.5 g of A, but used 50 cm3 of 2.0 mol dm-3 of B.
3c2 marks

Explain why the gradient of the curve in part (b) decreases as the time of the reaction progresses. 

3d2 marks

Another way to increase the rate of reaction is to increase the temperature. 

Explain why a small increase in temperature has a large effect on the initial rate of a chemical reaction.

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4a1 mark

The decomposition of hydrogen peroxide into water and oxygen is a very slow chemical reaction. 

Write the equation for the decomposition of hydrogen peroxide. 

4b5 marks

The rate of decomposition of hydrogen peroxide can be ascertained by collecting and measuring the volume of gas formed at specific time intervals.

i)
Draw a labelled diagram to show the apparatus that you would use to collect and measure the volume of gas formed during this reaction. 

ii)
Explain how you would use the results to determine the initial rate of the reaction.
4c3 marks

Two students set up the practical apparatus from part (b) to measure the rate of decomposition of hydrogen peroxide. Student A states that one way to increase the rate of the decomposition, would be to increase the concentration of hydrogen peroxide. 

Is student A correct? Explain your answer. 

4d1 mark

The decomposition of hydrogen peroxide is a slow reaction, so a catalyst is often added to speed up the rate of the reaction. Catalysts are used in many chemical reactions to increase the rate. 


The following shows a two-step reaction mechanism of a chemical reaction, where a catalyst, X is used. 

STEP 1:                               W + X  →  Y + Z

STEP 2:                               Y + W → Z + A + X

OVERALL REACTION:       2W → 2Z + A

Give a reason, other than the rate of reaction increasing, why it can be deduced from the three equations above that X is a catalyst. 

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5a2 marks

During the following reaction, A and B react together to produce C

A  +  2B  →  C

Figure 1 shows the production of C over time.  

Figure 1

KAFiI~XW_3-2

i)
Sketch a graph to show what happens to A and B during the progress of the reaction. 

ii)
On your graph, write the letter E at the point at which an equilibrium is first established. 
5b3 marks

In the reaction in part (a), lumps of A were used.
Use collision theory to explain what would happen to the rate of the reaction if powdered A was used instead of lumps. 

5c6 marks

In a different reaction, gaseous W and X were added together to produce Y and Z as shown in the equation below:

2W (g) + X (g) rightwards arrow Y (g) + 2Z (g)

A catalyst was added to speed up the rate of reaction.

Sketch a Maxwell-Boltzmann distribution on the axes below in Figure 2 to show the distribution of molecular energies at a constant temperature with and without a catalyst.

Use Ea to label the activation energy without a catalyst and Ec to label the activation energy with a catalyst.

Explain what your distribution shows.

Figure 2

UFb70ERf_4-2

5d6 marks

Some changes were made individually to the experiment completed in part (c).

Consider your Maxwell-Boltzmann distribution curve from part (c). For each of the changes in parts (i), (ii) and (iii) below, state and explain the effect that the change would have on:

  • The area under the curve
  • The value of the most probably energy of the molecules (Emp)
  • The proportion of molecules with energy greater than or equal to Ea
i)
The temperature of the original reaction is increased, but no other changes are made.

ii)
The number of molecules in the original reaction mixture is increased, but no other changes are made.

iii)
A catalyst is added to the original reaction mixture, but no other changes are made.
    .

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1a4 marks

Reaction rates can be affected by a range of factors including changes in pressure and temperature.

Figure 1

1-2

On Figure 1, sketch one Maxwell-Boltzmann distribution labelled T1 and a second Maxwell-Boltzmann distribution at a higher temperature labelled T2.

 State how the mean energy of the molecules would be at T2 compared to T1.

1b6 marks

Hydrogen iodide can be used in the manufacturing of pharmaceuticals and can be broken down back into its elements in standard form, iodine and hydrogen. 

2HI (g) → H2 (g) + I2 (g)     ΔH = - 52 kJ mol-1

The activation energy when uncatalysed is +183 kJ mol-1 and when catalysed with gold it is +105 kJ mol-1

i)
Sketch an energy profile diagram for the reaction, including the curves for the activation energies for both the catalysed and uncatalysed reactions.

ii)
Calculate the activation energy for the reverse reaction in both the uncatalysed and catalysed reactions.

iii)
Explain why increasing the concentration of hydrogen iodide gas results in a faster reaction rate.
1c3 marks

Catalysts are often used in industrial processes and can be used in a variety of forms. 

i)
Explain why it is likely that the solid gold catalyst was used in powder form to catalyse the reaction mentioned in part (c). 

ii)
Gold is a heterogeneous catalyst used in the formation of hydrogen iodide. State the difference between a homogenous and heterogenous catalyst.

ii)
State how, if at all, the area under the curve of a Maxwell-Boltzmann distribution curve, changes as a catalyst is introduced without changing the temperature or the total number of molecules.
1d3 marks

The Contact process is an important industrial process, contributing to the production of sulfuric acid. In the Contact process, solid vanadium (V) oxide, a heterogeneous catalyst, is used to make sulfur trioxide from sulfur dioxide and oxygen. This process is reversible.

i)
Write a balanced symbol equation for this reaction. Include state symbols in your answer.

ii)
Explain why the use of the catalyst in the Contact process, reduces energy demand and benefits the environment.

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2a4 marks

Methanol, for use as a fuel, can be produced by the reaction of carbon monoxide with hydrogen. The reaction is typically carried out at 310 degreeC and 3.25 × 107 Pa as shown by sample 1 in Figure 1.

CO (g) + 2 H2 (g)  ⇌  CH3OH (g)    ΔH = – 90 kJ mol–1

Figure 1

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i)
A change was made to the reaction conditions as shown by sample 2 in Figure 1. Deduce the change that was made to the reaction conditions.

ii)
Explain the effect this change has had on the rate of the reaction.

iii)
State the effect, if any, this change has had on the most probable value for the energy of the molecules (Emp).
2b6 marks

Consider the Maxwell-Boltzmann distribution curve below in Figure 2

Figure 2

83

For each of the changes in parts (i), (ii) and (iii) below, state and explain the effect that the change would have on:

  • The area under the curve 
  • The value of the most probably energy of the molecules (Emp
  • The proportion of molecules with energy greater than or equal to Ea

i)
The temperature of the original reaction is increased, but no other changes are made. 

ii)
The number of molecules in the original reaction mixture is increased, but no other changes are made. 

iii)
A catalyst is added to the original reaction mixture, but no other changes are made. 
2c4 marks

A chemist performed a reaction at three different temperatures, 100K, 300K and 700K as shown by the Maxwell-Boltzmann distribution graph in Figure 3.

Figure 3

7-1

i)
Label each curve in Figure 3 with the correct temperature values, 100K, 300K and 700K.

ii)
Consider the following statement, ‘All reacting molecules have higher kinetic energy at 700K than they do at 300K’. State whether you agree this statement is correct and justify your reasons.
2d5 marks

Hydrogen will react with chlorine to form the hydrogen halide, hydrogen chloride, a colourless gas.

H2 (g) + Cl2 (g) → 2HCl (g)

i)
Give one reason why most collisions between hydrogen and chlorine molecules do not lead to the formation of hydrogen chloride. 

ii)
Apart from changing the temperature, state and explain two ways of speeding up the formation of hydrogen chloride.

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3a5 marks

Students investigated the effect of increasing temperature on the reaction between sodium thiosulfate solution and dilute hydrochloric acid as shown in Figure 1

The students reacted 20.0 cm3 of 1.00 mol dm-3 sodium thiosulfate solution with 5.0 cm3 of hydrochloric acid in a 100 cm3 conical flask at 20 degreeC. The reactants are mixed and the time taken for a cross on paper placed under the conical flask to disappear was measured. The students rinsed the conical flask with hydrochloric acid after every trial. They then repeated the experiment at different temperatures of sodium thiosulfate.

Figure 1

9-1

i)
Explain why the students rinsing the conical flask with acid after each experiment will reduce the accuracy of the experiment.

ii)
Explain why it is necessary to use separate measuring cylinders for sodium thiosulfate solution and for dilute hydrochloric acid.

iii)
Describe how the students could use their experimental results to determine the rate of the reaction.
3b4 marks

Some changes were made to the experiment completed in part (a). The student repeated the experiment but kept the temperature of the solutions constant but they changed the concentration of sodium thiosulfate to 2.00 mol dm-3.

i)
State two variables, other than temperature, the student would need to keep the same to make sure their results were valid.

ii)
The student stated that if you use a catalyst in this reaction the amount of product produced will double. They also predicted that doubling the concentration from their original experiment would result in the rate quadrupling. Justify and explain if the student is correct with their ideas.
3c4 marks

Aqueous hydrogen peroxide decomposes according to the following equation.

2H2O2 (aq)  →  2H2O (l) + O2 (g)

The decomposition is catalysed by adding 0.25 g manganese (IV) oxide. Manganese (IV) oxide acts as a heterogeneous catalyst.

This can be investigated by measuring the volume of oxygen produced at various times as the reaction proceeds.

Catalysts are not used up during a reaction. 

Describe and outline a method to show that the manganese (IV) oxide is not used up in the decomposition of hydrogen peroxide and that it still functions as a catalyst.

3d3 marks

A student performed the experiment mentioned in part (c) to investigate how the volume of oxygen produced varied over time under different conditions. The same mass of catalyst was used in each experiment.

Table 1

Experiment

Concentration of H2O2 (aq) / mol dm-3

Volume of HCl (aq) / dm3

Temperature / degreeC

Catalyst

1

1.0

50

15

lumps

2

1.0

50

25

lumps

3

0.5

50

15

lumps

The student plotted a graph for their first set of results as shown in Figure 2.

Figure 2

10-1

i)
Draw two lines on Figure 2 to show how time affected the volume of oxygen collected in experiments 2 and 3. Label each line with the experiment number.

ii)
State one other variable the student could change in their experiment, to fully understand the factors that affect the rate of a reaction. 

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4a3 marks

Aqueous solutions of hydrogen peroxide, H2O2 (aq), decompose slowly into water and oxygen.

2 H2O2 (aq)  →  2 H2O (l) + O2 (g)

A student investigates the decomposition of H2O2 by measuring the volume of oxygen gas produced over time. All gas volumes are measured at room temperature and pressure.

The student uses 20.0 cm3 of 2.45 mol dm−3 H2O2.

From the results, the student determines the concentration of H2O2 at each time. The student then plots a concentration–time graph as shown in Figure 1.

Figure 1

5MCLWver_11

i)
Determine the total volume of oxygen in cm3, measured at room temperature and pressure, that the student should be prepared to collect in this investigation.

ii)
Suggest a piece of apparatus with an appropriate scale that would allow this gas volume to be collected.
4b4 marks

A student reacted excess zinc with aqueous dilute hydrochloric acid and measured the volume of hydrogen gas produced at regular intervals. The temperature of the acid was changed between experiments.

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

i)
Draw a labelled diagram to show how the gas can be collected and its volume measured, labelling the apparatus used.

ii)
A student suggested that an appropriate method was placing the reactants in a conical flask on the top of a top pan balance, to measure the rate of the reaction. Justify why this experimental procedure is not suitable for the above reaction.
4c4 marks

When 35.0 cm3 of 0.500 mol dm−3 aqueous ethanoic acid reacts with 410 mg calcium, the following reaction occurs and hydrogen gas is produced.

The volume of hydrogen produced can be measured using the equipment shown in Figure 1.

Figure 1

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i)
Write a balanced symbol equation for this reaction. 

ii)
With the aid of calculations, show that the calcium is in excess in this reaction.
4d5 marks

The student repeated the experiment and used 0.250 mol dm−3 of ethanoic acid solution with all other conditions the same. The calcium was still in excess.

Figure 2 shows the original experimental results from part (c).

Figure 2

nV_lhW4G_13

i)
Sketch a line on Figure 2 to show how the volume of hydrogen produced varies with time in this second experiment.

ii)
Describe how you would find a numerical value for the initial rate of reaction and for the maximum rate of reaction in this experiment from the graph. 

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5a2 marks

In a series of experiments, a student reacted 50.0 cm3 of 1.00 mol dm–3 hydrochloric acid to small pieces of calcium carbonate in a conical flask placed on an electronic balance. The loss in mass of the flask and its contents was recorded for 15 minutes.

A graph was plotted of the students’ results in Figure 1.

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

Figure 1
18-1

Describe and explain the shape of the curve drawn for Experiment 1. Link your answer to the mass lost in the experiment.

5b3 marks

The student carried out three further experiments and varied the reaction conditions as shown in Table 1.

Table 1

Experiment

CaCO3 size

HCl

Temperature / degreeC

2

Small pieces

50.0 cm3 and 1.00 mol dm-3

80

3

One large piece

50.0 cm3 and 1.00 mol dm-3

20

4

Small pieces

50.0 cm3 and 2.00 mol dm-3

20

Using the graph from part (a), draw and label three curves to represent experiments 2, 3 and 4.

5c6 marks

In a different reaction, gaseous W and X were added together to produce Y and Z as shown in the equation below:

2 W (g) + X (g) → Y (g) + 2 Z (g)

A catalyst was added to speed up the rate of reaction. 

Sketch a Maxwell-Boltzmann distribution on the axes below in Figure 2 to show the distribution of molecular energies at a constant temperature with and without a catalyst.

Use Ea to label the activation energy without a catalyst and Ec to label the activation energy with a catalyst. 

Explain what your distribution shows. 

Figure 2

4-2

5d6 marks

Ammonia, NH3, is manufactured by the chemical industry from nitrogen and hydrogen gases. 

 N2 (g) + 3 H2 (g) ⇌ 2 NH3 (g)   ΔH = − 92 kJ mol−1

An iron catalyst is used which provides several benefits for sustainability.

The chemical industry uses operational conditions that are different from the conditions predicted to give a maximum equilibrium yield.

Justify the above statements using your knowledge of kinetics and equilibrium. 

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