Simple Rate Equations, Orders of Reaction & Rate Constants (Cambridge (CIE) A Level Chemistry): Exam Questions

The initial rate of the reaction of chlorine dioxide, ClO₂, and fluorine, F₂, is measured in a series of experiments at a constant temperature.

2 ClO₂ + F₂ → 2ClO₂F

The results obtained are shown in Table 1.1.

Table 1.1

The rate equation is rate = k[ClO₂][F₂].

i) Define the term order of reaction with respect to a particular reagent.

[1]

ii) Use the results of experiment 1 to calculate the rate constant, k, for this reaction. State the units of k. Show your working.

k = ............................. units .............................

[2]

iii) Use the data in Table 1.1 to calculate the initial rate in experiment 2.

initial rate = ...................................... mol dm^-3 s^-1

[1]

iv) Use the data in Table 1.1 to calculate [ClO₂] in experiment 3.

[ClO₂] = ............................................ mol dm^-3

[1]

i) Define the term rate-determining step.

[1]

ii) The mechanism of the reaction between ClO₂ and F₂ has two steps.

Suggest equations for the two steps of this mechanism.

step 1 ..............................................................................

step 2 ...............................................................................

[1]

iii) State and explain which of the two steps is the rate-determining step.

rate-determining step = ..........................

[1]

Describe the effect of temperature change on the rate of a reaction and the rate constant.

This question is about the reactions of different unknown compounds.

Two compounds, X and Y, were reacted together.

The initial rate of reaction was measured when compound X and compound Y were reacted together. The temperature was kept constant and the results of the experiments are shown in Table 2.1.

Table 2.1

i) Use the data in Table 2.1 to deduce the order of reaction with respect to X.

[1]

ii) State the order of the reaction with respect to Y.

[1]

iii) State the overall order of the reaction.

[1]

iv) Construct the rate equation for the reaction.

[1]

Three separate experiments were carried out at 400K with different starting concentrations of A and B. The results are shown in Table 2.2.

Table 2.2

i) Deduce the order of reaction with respect to each reactant. Explain your reasoning.

[2]

ii) Construct the rate equation for this reaction. Calculate the rate constant, k.

State the units of k. Show your working.

[4]

The overall rate equation for a reaction is rate = k[P]²[Q]

 i) State what the units would be of the rate constant, k, in this reaction. 

[1]

 ii) State the overall order of the reaction above. 

[1]

Chemists measured the rate of a chemical reaction in a series of experiments between compounds C and D at a fixed temperature as shown in Table 2.3.

Table 2.3

i) Deduce the order of reaction with respect to C and D.

[2]

ii) Construct the rate equation for this reaction.

[1]

The initial rate of reaction for iodine and propanone in acid solution is measured in a series of experiments at constant temperature.

CH₃COCH₃ + I₂ CH₃COCH₂I + HI

The rate equation was determined experimentally to be as shown.

rate = k [CH₃COCH₃][H⁺]

i) State the order of reaction with respect to CH₃COCH₃, I₂ and H⁺.

[1]

ii) State the overall order of the reaction.

[1]

iii) State and explain the role of the acid.

[2]

The rate of this reaction is 5.40 × 10^–3 mol dm^–3 s^–1 when

• the concentration of CH₃COCH₃ is 1.50 × 10^–2 mol dm^–3

• the concentration of I₂ is 1.25 × 10^–2 mol dm^–3

• the concentration of H⁺ is 7.75 × 10^–1 mol dm^–3.

i) Calculate the rate constant, k, for this reaction. State the units of k.

k = .............................. 

units = ..............................

[2]

ii) Complete the table by placing one tick (✔) in each row to describe the effect of increasing the temperature on the rate constant and on the rate of reaction.

 [1]

 Fig. 3.1 is produced from the results, which shows how the concentration of I₂ changes during the reaction. 

 Fig. 3.1

 Describe how Fig 3.1 could be used to determine the initial rate of the reaction.

On Fig 3.2, sketch a graph to show how the initial rate changes with different initial concentrations of H⁺ in this reaction.

 Fig. 3.2 

The equation for the decomposition of hydrogen peroxide without a catalyst is shown.

2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)

With this information, a student suggests that the rate equation for this reaction is as follows.

rate = k [H₂O₂]²

State whether the student is correct. Explain your answer.

A student carries out separate experiments using different initial concentrations of H₂O₂. The initial rate of each reaction is measured.

 The table shows the results that are obtained.

 i) Plot a graph on the grid of H₂O₂ concentration against rate of reaction. 

Draw a line of best fit through the plotted points.  

[2]

ii) State the rate equation for this reaction.

[1]

Another student monitors the concentration of H₂O₂ over time. The graph obtained is shown.

Using appropriate calculations, describe how the student could use this graph to show the order of reaction with respect to H₂O₂.

Under certain conditions, the decomposition of hydrogen peroxide is found to be first order with respect to hydrogen peroxide, with a rate constant, k, of 2.0 × 10^–6 s^–1.

i) Calculate the initial rate of decomposition of a 0.65 mol dm^–3 hydrogen peroxide solution.

initial rate = .............................. mol dm^–3 s^–1

[1]

ii) Explain the effect (if any) of decreasing the temperature on the rate constant and the rate of reaction. 

[2]

This question is about nitrogen oxides.

Nitrogen monoxide is produced by combustion in car engines and released into the atmosphere.

In the lower atmosphere, nitrogen monoxide is responsible for the formation of ozone in a series of reactions shown below.

NO (g) + ½O₂ (g) → NO₂ (g)

NO₂ (g) → NO (g) + O (g)

O₂ (g) + O (g) → O₃ (g)

i) Construct the overall equation for this series of reactions.

[1]

ii) Explain why NO is acting as a catalyst in this reaction.

[1]

The rate equation for the reaction between nitrogen monoxide (NO) and oxygen (O₂) under certain conditions is given.

Rate = k [NO]²[O₂]

The result of an experiment in which NO reacted with O₂ is shown in the table.

Use the data and the rate equation to calculate a value for the rate constant k.

State the units of k. Show your working.

Nitrogen dioxide reacts with carbon monoxide at 100 ℃ to form nitrogen monoxide and carbon dioxide.

2NO₂ (g) + CO (g) → NO (g) + CO₂ (g)

The rate of this reaction was measured at different initial concentrations of the two reagents.

The rate equation was determined to be

Rate = k [NO₂]²

The table shows an incomplete set of results.

i) Use the data from Experiment 1 to calculate a value for the rate constant, k, at this temperature. State the units of k. Show your working.

[2]

ii) Use your value of k from (i) to complete the table for the reaction between NO₂ and CO.

[2]

Nitrogen monoxide, NO (g), reacts with hydrogen, H₂ (g), under certain conditions. 

2NO (g) + 2H₂ (g) → N₂ (g) + 2H₂O (g) 

The rate equation for this reaction is given. 

rate = k [NO]² [H₂] 

One proposed mechanism for this reaction occurs in two steps.

1. 2NO + H₂ → N₂ + H₂O₂ 

2. H₂O₂ + H₂ → 2H₂O

 Explain which of the steps is the rate-determining step.

In the presence of acid, H⁺(aq), chlorine and propanone react together:

CH₃COCH₃ (aq) + Cl₂ (aq) → CH₃COCH₂Cl (aq) + HCl (aq)

A student carried out an investigation into the kinetics of this reaction.

The student investigated how different concentrations of chlorine affect the initial rate of the reaction. A graph of [Cl₂ (aq)] against time is shown in Fig. 1.1.

Fig. 1.1

The student then investigated how different concentrations of propanone and H⁺(aq) affect the initial rate of reaction.

Their results are shown below.

Use the student's results to deduce the reaction orders. Explain your answer.

i) Deduce the rate equation for this reaction.

Rate equation ...........................................................

[1]

ii) Calculate the rate constant, k, for this reaction. State the units of k. Show your working.

[3]

The student varied the concentrations of the reactants in another experiment and they found the rate to be 0.26 x 10^-5 mol dm^-3 s^-1. They also used a pH probe and found that the reaction mixture had a pH of 2.

Calculate the initial rate of the reaction mixture if the amount of acid added was altered to give a pH of 1. Show your working.

Assume the temperature and the initial concentrations of the other reactants remained the same.

The experiment was repeated at a lower temperature. State and explain the effect on the rate and the rate constant of the reaction.