How Fast? The Rate of Chemical Change (DP IB Chemistry: SL): Exam Questions

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.

[2]

ii) Sketch a distribution for this same sample of gas, at a higher temperature, and label it as T₂.

[2]

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.

[1]

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

[2]

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

A student investigates the reaction between solid A and aqueous solution B:

A (s) + B (aq) → C (g)

In the first experiment, B was the limiting reagent. The volume of gas C was collected over time, as shown by the solid line in the graph below.

Define the term rate of reaction.

Explain what is indicated by the horizontal section of the graph, labelled R.

On the graph, sketch the curve that would be obtained if the experiment were repeated using double the concentration of solution B, with all other conditions kept the same.

Explain why the gradient of the original curve decreases over time.

Explain, using collision theory, why a small increase in temperature has a large effect on the rate of reaction.

The rate of the slow decomposition of aqueous hydrogen peroxide can be investigated by measuring the volume of oxygen gas produced over time. A catalyst, such as manganese(IV) oxide, is often used.

Write the balanced chemical equation for the decomposition of hydrogen peroxide.

Draw a labelled diagram of the assembled apparatus suitable for collecting and measuring the volume of oxygen produced at regular time intervals. 

Explain how the initial rate of reaction could be determined from the experimental results

The graph shown below represents the decomposition of hydrogen peroxide.

Figure 1

Explain why the rate of reaction decreases as the reaction proceeds.

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. 

[1]

ii) On your graph, label the point at which an equilibrium is first established the letter E. 

[1]

In the reaction in part (a), large pieces 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 large pieces.

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. 

i) 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.

[3]

Figure 2

ii) Explain what your distribution shows.

[3]

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 probable 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.

[2] 

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

[2] 

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

[2]

A student carried out a metal displacement reaction between zinc powder and copper(II) sulfate solution. The equation for the reaction is

Zn (s) + CuSO₄ (aq) → ZnSO₄ (aq) + Cu (s)

3.78 g of zinc powder was added to 50.0 cm³ of 0.250 moldm^-3 copper(II) sulfate solution.

Determine the limiting reagent showing your working.

The reaction between the zinc and copper sulfate was carried out in a polystyrene cup and the temperature change was measured using a temperature probe. The maximum temperature rise the student recorded was 8.5 ^oC. 

Using sections 1 and 2 of the data booklet, calculate the enthalpy change, ∆H, for the reaction, in kJ.

Assume that all the heat evolved was absorbed by the solution, and that the density and specific heat capacity of the copper(II) sulfate solution are the same as pure water.

State two further assumptions made in the calculation of ∆H.

On the axes below, sketch a graph of the concentration of zinc sulfate, ZnSO₄ (aq), versus time and show how the graph may be used to find the initial rate of reaction.

A student investigated the rate of decomposition of hydrogen peroxide, H₂O₂, at a temperature of 45 ^o The decomposition reaction occurs in the presence of a catalyst, MnO₂.

2H₂O₂ (aq)  O₂ (g) + 2H₂O (l) 

The results she obtained are shown in the below.

 Plot a graph on the axes below in and determine the rate of reaction after 60 s.

i) On the same graph, sketch the shape obtained if the student had carried out the same reaction at 60 ^oC.

[1]

ii) Explain the shape of the graph at 60 ^oC.

[2]

The decomposition of hydrogen peroxide can be investigated by measuring the volume of oxygen given off using the apparatus shown below.

i) Explain why the volume of oxygen given off can be used as a measure of the concentration of hydrogen peroxide. 

[1]

ii) Suggest one limitation of using the apparatus used.

[1]

iii) Suggest an alternative method of measuring the rate of reaction.

[1]

Two students decide to measure the rate of decomposition for H₂O₂ using the change in mass as oxygen escapes from the reaction container.

One student says that they should use a three decimal place rather than two decimal place balance because it will make their results more accurate. The second student disagrees and says it will make their results more precise, but not more accurate.

Which student is correct?

Explain why the reaction represented below is a redox reaction. 

2ClO₂ (aq) + 2NaOH (aq) → NaClO₃ (aq) + NaClO₂ (aq) + H₂O (l) 

A group of students planned how to investigate the effect of changing the concentration of H₂SO₄ on the initial rate of reaction with magnesium: 

Mg (s) + H₂SO₄ (aq) MgSO₄ (aq) + H₂ (g) 

They decided to measure how long the reaction took to complete when similar masses of magnesium were added to acid. 

Two methods were suggested:

Method 1 - Use small pieces of magnesium ribbon, an excess of acid and record the time taken for the magnesium ribbon to disappear

Method 2 - Use large strips of magnesium ribbon, an excess of magnesium and record the time taken for bubbles to stop forming 

Deduce, giving a reason, which of method 1 and method 2 would be the least affected if the masses of magnesium ribbon used varied slightly between each experiment.

Neither method in part a) actually allows the initial rate to be calculated. Outline a method that would allow the calculation of initial rate.

The reaction is to be conducted across a few weeks.

State a factor that has a significant effect on reaction rate, which could vary between experiments across the weeks and therefore needs to be controlled.

One group collected the following data using 1.50 mol dm^-3 acid: 

i) Comment on the use of uncertainty when calculating the mean.

[2]

ii) Calculate the mean time for the set of results.

[2]

When investigating the reaction between sulfuric acid and calcium carbonate, it was observed that a small increase of temperature of around 10 ^oC caused a doubling in the rate of the reaction.

Sketch and label Maxwell-Boltzmann curve for the two temperatures T and T+10, and use this diagram to help to explain this effect of temperature on rate.

Why do some collisions at high temperatures still not result in the formation of the product?

Identify and explain another factor that affects the number of particles present in a solution with sufficient energy to react.

Some groups investigating the effect of temperature on rate stirred their reactions, some did not.

Explain the effect of stirring upon the rate of the reaction.

0.5 g of magnesium reacts with 50 cm³ of 0.01 moldm^-3 nitric acid. Magnesium is in excess.

A graph monitoring the volume of hydrogen gas produced is shown below:

i) Calculate the mean rate of reaction over the first 15 seconds of the reaction

[1]

ii) Calculate the actual rate of reaction at 15 seconds

[3]

iii) Explain the difference in values for rate

[1]

Compare the expected rate and progress of the reaction if 25 cm³ of 0.2 mol dm^-3 nitric acid was used instead of 50 cm³ of 0.1 mol dm^-3 nitric acid.

Suggest one change to the reaction that could be made to produce more hydrogen gas in total and explain your choice.

Suggest why it is often better to study a slower reaction instead of a faster one.

The following energy profile diagram shows the pathways for both a catalysed and uncatalysed reversible reaction:

Identify the letter(s) representing the activation energy for the catalysed reverse reaction.

State and explain the effect that this catalyst will have on the equilibrium yield.

Vehicles with combustion engines usually have catalytic convertors added to catalyse the oxidation of carbon monoxide into carbon dioxide and to catalyse the reduction of nitrogen oxides to nitrogen. These catalysts are usually rhodium or platinum.

Leaded fuels were phased out as they were found to poison these catalysts, binding irreversibly to the metal surface.

Explain the problems for drivers of the catalysts being poisoned.

Suggest a situation in which using a catalyst would not be appropriate.

The reaction between iodide ions and persulfate ions is a 'clock' reaction and often used to study reaction kinetics.

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

Deduce the redox changes taking place in the reaction.

The rate constant for a reaction doubles when the temperature is increased from 25.0 °C to 35 °C.

Calculate the activation energy, E_a, in kJ mol⁻¹ for the reaction using section 1 and 2 of the data booklet.

A different reaction route is used which reduces the activation energy of the reaction. 

Explain how the rate constant calculated in part(b) would differ.