Kinetics (AQA A Level Chemistry)

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. 

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

Table 1

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. 

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

Figure 1

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.

Explain how catalysts increase the rate of a chemical reaction. 

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

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. 

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

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. 

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

State the meaning of the term rate of reaction.

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

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.

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

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.

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

Write the equation for the decomposition of hydrogen peroxide. 

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.

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. 

Explain whether student A is correct. 

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

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

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

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. 

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. 

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)  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 E_a to label the activation energy without a catalyst and E_c to label the activation energy with a catalyst.

Explain what your distribution shows.

Figure 2

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 (E_mp)

• The proportion of molecules with energy greater than or equal to E_a

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.     .

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

Figure 1

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.

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.

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 (b). 

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

iii) 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.

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.

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 C 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

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 (E_mp).

Consider the Maxwell-Boltzmann distribution curve below in Figure 2. 

Figure 2

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 (E_mp) 

• The proportion of molecules with energy greater than or equal to E_a

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. 

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

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.

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.

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 C. 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

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.

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.

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 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.

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

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

Figure 2

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. 

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

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.

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.

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

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.

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

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. 

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

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

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

Table 1

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

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 E_a to label the activation energy without a catalyst and E_c to label the activation energy with a catalyst. 

Explain what your distribution shows. 

Figure 2

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.