8. Reaction kinetics
A section of Chemistry, 9701
Listing 10 of 220 questions
The initial rate of reaction for propanone and iodine in acid solution is measured in a series of experiments at a constant temperature. H+ catalyst CH3COCH3 + I2 CH3COCH2I + HI The rate equation was determined experimentally to be as shown. rate = k [H+] State the order of reaction with respect to ● CH3COCH3 ● I2 ● H+ and state the overall order of this reaction.  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. Calculate the rate constant, k, for this reaction. State the units of k.  k =  units =  Complete the table by placing one tick () in each row to describe the effect of decreasing the temperature on the rate constant and on the rate of reaction. decreases no change increases rate constant rate of reaction  From the results, a graph is produced which shows how the concentration of I2 changes during the reaction. time Describe how this graph could be used to determine the initial rate of the reaction. On the axes below, sketch a graph to show how the initial rate changes with different initial concentrations of CH3COCH3 in this reaction. rate  The rate of a reaction between metal ions was studied. The following three-step mechanism has been suggested for this reaction. Step1 is the rate-determining step. step 1 Ce4+ + Mn2+ Ce3+ + Mn3+ step 2 Ce4+ + Mn3+ Ce3+ + Mn4+ step 3 Mn4+ + Tl + Tl 3+ + Mn2+ Explain the meaning of the term rate-determining step. Use this mechanism to ● determine the overall equation for this reaction ● suggest the role of Mn2+ ions in this mechanism. Explain your answer.  
9701_s19_qp_42
THEORY
2019
Paper 4, Variant 2
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Chlorate(ions undergo the following reaction under aqueous conditions. 2NH3 + Cl O– N2H4 + Cl – + H2O A series of experiments was carried out at different concentrations of Cl O– and NH3. The table shows the results obtained. experiment [Cl O–] / mol dm–3 / mol dm–3 initial rate / mol dm–3 s–1 0.200 0.100 0.256 0.400 0.200 2.05 0.400 0.400 8.20 Use the data in the table to determine the order with respect to each reactant, Cl O– and NH3. Show your reasoning. Write the rate equation for this reaction. rate = Use the results of experiment1 to calculate the rate constant, k, for this reaction. Include the units of k.  k =  units =  On the axes sketch a graph to show how the value of k changes as temperature is increased. rate constant, k temperature  In another experiment, the reaction between chlorate(ions and iodideions in aqueous alkali was investigated. A solution of iodide ions in aqueous alkali was added to a large excess of chlorate(ions and [I–] was measured at regular intervals. Describe how the results of this experiment can be used to confirm that the reaction is first-order with respect to [I–]. A three-step mechanism for this reaction is shown. step 1 Cl O– + H2O HCl O + OH– step 2 I– + HCl O HIO + Cl – step 3 HIO + OH– H2O + IO– Use this mechanism to deduce the overall equation for this reaction. Identify a step that involves a redox reaction. Explain your answer.  
9701_s19_qp_43
THEORY
2019
Paper 4, Variant 3
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Aqueous bromine reacts with methanoicacid to form hydrogenbromide and carbondioxide gas. Br2+ HCO2H→ 2HBr+ CO2The table shows the oxidation numbers of bromine and carbon in the species involved in this reaction. Br in Br2 C in HCO2H Br in HBr C in CO2 oxidation number +2 –1 +4 Identify the oxidising agent in this reaction. Explain your reasoning with reference to oxidation numbers. Suggest one change you would observe, ignoring temperature changes, when bromine reacts with methanoicacid. This reaction can be followed by measuring the concentration of bromine present in the mixture at regular time intervals. The graph shows the change in concentration of bromine against time in a reaction carried out at 20 °C. time / s 105 / mol dm–3 Use the graph to calculate the average rate of reaction at 20 °C during the first 600 s. State the units of this rate of reaction. average rate of reaction units  The experiment is repeated at a temperature of 40 °C. This relatively small increase in temperature produces a large increase in reaction rate. Sketch a graph, on the same axes, to show the expected results when repeating the experiment at 40 °C. The rate of reaction increases when the frequency of successful collisions between reactant particles increases. Explain why an increase in temperature produces this effect. Complete the ‘dot-and-cross’ diagram, showing outer electrons only, to show the bonding in methanoic acid, HCO2H. C O O H H  
9701_s21_qp_21
THEORY
2021
Paper 2, Variant 1
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Dinitrogenpentoxide, N2O5, is dissolved in an inert solvent and the rate of decomposition of N2O5 is investigated. This reaction produces nitrogendioxide, which remains in solution, and oxygen gas. N2O5→ 2NO2+ 1 2O2Suggest what measurements could be used to follow the rate of this reaction from the given information. In a separate experiment, the rate of the decomposition of N2O5is investigated. N2O5→ 2NO2+ 1 2O2The graph shows the results obtained. 0.30 0.25 0.20 0.15 0.10 0.05 time / s / mol dm–3 The reaction is first order with respect to N2O5. This can be confirmed from the graph using half‑lives. Explain the term half-life of a reaction. Determine the half-life of this reaction. Show your working on the graph.  half-life = s Suggest the effect on the half-life of this reaction if the initial concentration of N2O5 is halved. Use the graph in 5to determine the rate of reaction at 200 s. Show your working.  rate =  units =  The rate equation for this reaction is shown. rate = k Use your answer to to calculate the value of the rate constant, k, for this reaction and state its units.  k = units Nitrogendioxide reacts with ozone, O3, as shown. 2NO2 + O3 → N2O5 + O2 The rate equation for this reaction is rate = k . Suggest a possible two-step mechanism for this reaction. 
9701_s21_qp_41
THEORY
2021
Paper 4, Variant 1
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In aqueous solution, chlorinedioxide, Cl O2, reacts with hydroxide ions as shown. 2Cl O2 + 2OH– → Cl O3 – + Cl O2 – + H2O A series of experiments is carried out using different concentrations of Cl O2 and OH–. The table shows the results obtained. experiment [Cl O2] / mol dm–3 [OH–] / mol dm–3 initial rate / mol dm–3 min–1 0.020 0.030 7.20 × 10–4 0.020 0.120 2.88 × 10–3 0.050 0.030 4.50 × 10–3 Explain the term order of reaction. Use the data in the table to determine the order of reaction with respect to each reactant, Cl O2 and OH–. Explain your reasoning. Use your answer to to construct the rate equation for this reaction. rate = Use your rate equation and the data from experiment1 to calculate the rate constant, k, for this reaction. Include the units of k.  k = units The decomposition of benzenediazonium ions, C6H5N2 +, using a large excess of water, is a first‑order reaction. The graph shows the results obtained. 0.025 0.020 0.015 0.010 0.005 time / s [C6H5N2 +] / mol dm–3 Draw the structure of the organic product formed in this reaction.  Use the graph to determine the rate of reaction at 100 s. Show your working.  rate = mol dm–3 s–1 Sketch a concentration–time graph for a zero-order reaction. Use your graph to suggest how successive half-lives for a zero-order reaction vary as the concentration of a reactant decreases. Indicate this by placing a tick ( ) in the appropriate box in the table. concentration time successive half-lives decrease no change in successive half-lives successive half-lives increase  
9701_s21_qp_42
THEORY
2021
Paper 4, Variant 2
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Dinitrogenpentoxide, N2O5, is dissolved in an inert solvent and the rate of decomposition of N2O5 is investigated. This reaction produces nitrogendioxide, which remains in solution, and oxygen gas. N2O5→ 2NO2+ 1 2O2Suggest what measurements could be used to follow the rate of this reaction from the given information. In a separate experiment, the rate of the decomposition of N2O5is investigated. N2O5→ 2NO2+ 1 2O2The graph shows the results obtained. 0.30 0.25 0.20 0.15 0.10 0.05 time / s / mol dm–3 The reaction is first order with respect to N2O5. This can be confirmed from the graph using half‑lives. Explain the term half-life of a reaction. Determine the half-life of this reaction. Show your working on the graph.  half-life = s Suggest the effect on the half-life of this reaction if the initial concentration of N2O5 is halved. Use the graph in 5to determine the rate of reaction at 200 s. Show your working.  rate =  units =  The rate equation for this reaction is shown. rate = k Use your answer to to calculate the value of the rate constant, k, for this reaction and state its units.  k = units Nitrogendioxide reacts with ozone, O3, as shown. 2NO2 + O3 → N2O5 + O2 The rate equation for this reaction is rate = k . Suggest a possible two-step mechanism for this reaction. 
9701_s21_qp_43
THEORY
2021
Paper 4, Variant 3
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A large amount of N2Odecomposes into nitrogen gas and oxygen gas in the presence of a tiny amount of a gold foil catalyst. The gold foil provides a solid surface on which the catalysed reaction takes place. The graph shows the concentration of N2Oagainst time as it decomposes. The graph is a straight line. concentration of N2O0 time / min Which row describes: ● the change in rate of reaction as N2Odecomposes from 0 to 10 minutes ● the effect of adding more gold foil catalyst on the rate of decomposition of the same amount and concentration of N2O?
9701_s22_qp_12
MCQ
2022
Paper 1, Variant 2
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Questions Discovered
220