8.1. Rate of reaction
A subsection of Chemistry, 9701, through 8. Reaction kinetics
Listing 10 of 86 questions
Nitrogenmonoxide, NO, reacts with hydrogen, H2, under certain conditions. 2NO+ 2H2N2+ 2H2ODefine the term rate of reaction. Identify a change in the reaction mixture that would enable the rate of this reaction to be studied. The rate equation for this reaction is given. rate = k 2 The result of an experiment in which NO reacted with H2 is shown in the table. initial / mol dm–3 initial / mol dm–3 initial rate of reaction / mol dm–3 s–1 2.50 × 10–3 2.50 × 10–3 1.27 × 10–3 Use the data and the rate equation to calculate a value for the rate constant k. Give the units of k. k = units = A second experiment is performed at the same temperature. The initial concentration of H2is 4.60×10–3 mol dm–3. The initial rate of the reaction is 2.31×10–3 mol dm–3 s–1. Calculate the initial concentration of NO. initial concentration of NO= mol dm–3 State the order of the reaction with respect to NOand with respect to H2, and the overall order of the reaction. overall order The reaction is believed to proceed in three steps. 2NO N2O2 N2O2 + H2 N2O + H2O N2O + H2 N2 + H2O Deduce which of the three steps is the rate-determining step. Explain your answer to . A third experiment is performed under different conditions. A small amount of H2of concentration 0.0200 mol dm–3 is mixed with a large excess of NO. The concentration of H2is found to have a constant half-life of 2.00seconds under the conditions used. Define the term half-life. Use the axes below to construct a graph of the variation in the concentration of H2during the first 6seconds under the conditions used. 0.02 0.01 time / s / mol dm–3 NOacts as a catalyst in the oxidation of atmospheric sulfurdioxide. Give two equations to describe how NOacts as a catalyst in this process. equation 1 equation 2 Explain why NOcan be described as a catalyst in this reaction. Describe, with the aid of an equation, an environmental consequence of the oxidation of atmospheric sulfurdioxide.
9701_s18_qp_41
THEORY
2018
Paper 4, Variant 1
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Iodine monochloride, ICl, is a yellow-brown gas. It reacts with hydrogen gas under certain conditions as shown. 2ICl + H22HCl + I2Experiments are performed using different starting concentrations of ICl and H2. The initial rate of each reaction is measured. The following results are obtained. experiment / mol dm–3 / mol dm–3 relative rate of reaction 4.00 × 10–3 4.00 × 10–3 1.00 4.00 × 10–3 7.00 × 10–3 1.75 4.00 × 10–3 1.00 × 10–2 2.50 5.00 × 10–3 8.00 × 10–3 2.50 7.00 × 10–3 8.00 × 10–3 3.50 Identify a change, taking place in the reaction mixture, that would enable measurements of the rate of this reaction to be made. Use the data in the table to show that the reaction is first order with respect to H2. Use the data in the table to show that the reaction is first order with respect to ICl . Complete the rate equation for the reaction between ICl and H2. rate = Use experiment3 to calculate a numerical value for the rate constant, k. k = The reaction 2ICl + H22HCl + I2is first order with respect to ICl and first order with respect to H2. Suggest a mechanism for this reaction. You should assume ● the mechanism has two steps, ● the first step is much slower than the second step. first step second step An alternative method is used to show that the reaction is first order with respect to H2. This method uses a large excess of ICl and measures how the concentration of H2varies with time. Describe two ways of using these results to show the reaction is first order with respect to H2concentration. Explain the reason for using a large excess of ICl . A chemical reaction may be speeded up by the presence of a catalyst. Explain why a catalyst increases the rate of a chemical reaction.
9701_s18_qp_42
THEORY
2018
Paper 4, Variant 2
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Nitrogenmonoxide, NO, reacts with hydrogen, H2, under certain conditions. 2NO+ 2H2N2+ 2H2ODefine the term rate of reaction. Identify a change in the reaction mixture that would enable the rate of this reaction to be studied. The rate equation for this reaction is given. rate = k 2 The result of an experiment in which NO reacted with H2 is shown in the table. initial / mol dm–3 initial / mol dm–3 initial rate of reaction / mol dm–3 s–1 2.50 × 10–3 2.50 × 10–3 1.27 × 10–3 Use the data and the rate equation to calculate a value for the rate constant k. Give the units of k. k = units = A second experiment is performed at the same temperature. The initial concentration of H2is 4.60×10–3 mol dm–3. The initial rate of the reaction is 2.31×10–3 mol dm–3 s–1. Calculate the initial concentration of NO. initial concentration of NO= mol dm–3 State the order of the reaction with respect to NOand with respect to H2, and the overall order of the reaction. overall order The reaction is believed to proceed in three steps. 2NO N2O2 N2O2 + H2 N2O + H2O N2O + H2 N2 + H2O Deduce which of the three steps is the rate-determining step. Explain your answer to . A third experiment is performed under different conditions. A small amount of H2of concentration 0.0200 mol dm–3 is mixed with a large excess of NO. The concentration of H2is found to have a constant half-life of 2.00seconds under the conditions used. Define the term half-life. Use the axes below to construct a graph of the variation in the concentration of H2during the first 6seconds under the conditions used. 0.02 0.01 time / s / mol dm–3 NOacts as a catalyst in the oxidation of atmospheric sulfurdioxide. Give two equations to describe how NOacts as a catalyst in this process. equation 1 equation 2 Explain why NOcan be described as a catalyst in this reaction. Describe, with the aid of an equation, an environmental consequence of the oxidation of atmospheric sulfurdioxide.
9701_s18_qp_43
THEORY
2018
Paper 4, Variant 3
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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_41
THEORY
2019
Paper 4, Variant 1
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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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