8. Reaction kinetics
A section of Chemistry, 9701
Listing 10 of 220 questions
Aqueous acidified iodate(ions, IO3 –, react with iodide ions, as shown. IO3 – + 6H+ + 5I– 3I2 + 3H2O The initial rate of this reaction is investigated. Table 3.1 shows the results obtained. Table 3.1 experiment [IO3 –] / mol dm–3 [H+] / mol dm–3 [I–] / mol dm–3 initial rate / mol dm–3 min–1 0.0400 0.0150 0.0250 4.20 × 10–2 0.120 to be calculated 0.0125 7.09 × 10–2 The rate equation for this reaction is rate = k [IO3 –][H+]2[I–]2. Explain what is meant by order of reaction. Complete Table 3.2. Table 3.2 the order of reaction with respect to [IO3 –] the order of reaction with respect to [H+] the order of reaction with respect to [I–] the overall order of reaction Use your answer to to sketch lines in to show the relationship between the initial rates and the concentrations of [IO3 –] and [I–]. initial rate [IO3 –] initial rate [I–] Use data from Table 3.1 to calculate the rate constant, k, for this reaction. Include the units of k. k = units Use data from Table 3.1 to calculate the concentration of hydrogen ions, [H+], in experiment 2. [H+] = mol dm–3 This reaction is repeated in two separate experiments. The experiments are carried out at the same temperature and with the same concentrations of I– and IO3 –. One experiment takes place at pH 1.0 and the other experiment takes place at pH 2.0. Calculate the value of rate at pH 1.0 rate at pH 2.0 . value of rate at pH 1.0 rate at pH 2.0 = In aqueous solution, iron(ions react with iodide ions, as shown. 2Fe3+ + 2I– 2Fe2+ + I2 The initial rate of reaction is first order with respect to Fe3+ and second order with respect to I–. The mechanism for this reaction has three steps. Each step involves only two ions reacting together. Suggest equations for the three steps of this mechanism. Identify the rate‑determining step. step 1 step 2 step 3 rate‑determining step =
9701_s23_qp_41
THEORY
2023
Paper 4, Variant 1
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In aqueous solution, iron(ions react with iodide ions, as shown. 2Fe3+ + 2I– 2Fe2+ + I2 A series of experiments is carried out using different concentrations of Fe3+ and I–, as shown in Table 4.1. Table 4.1 experiment [Fe3+] / mol dm–3 [I–] / mol dm–3 initial rate / mol dm–3 s–1 0.0400 0.0200 2.64 × 10–4 0.1200 0.0200 7.92 × 10–4 0.0800 0.0400 2.11 × 10–3 Explain what is meant by overall order of reaction. Use the data in Table 4.1 to deduce the order of reaction with respect to Fe3+ and with respect to I–. Explain your reasoning. Use your answer to to construct the rate equation for this reaction. rate = Use your answer to and the data from experiment 1 to calculate the rate constant, k, for this reaction. Include the units of k. k = units Describe qualitatively the effect of an increase in temperature on the rate constant and on the rate of this reaction. In aqueous solution, iodide ions react with acidified hydrogen peroxide, as shown. 2I– + H2O2 + 2H+ I2 + 2H2O The initial rate of reaction is found to be first order with respect to I–, first order with respect to H2O2 and zero order with respect to H+. shows a possible four-step mechanism for this reaction. step 1 H2O2 + I– IO– + H2O step 2 H+ + IO– HIO step 3 HIO + I– I2 + OH– step 4 OH– + H+ H2O Suggest which of the steps, 1, 2, 3 or 4, in this mechanism is the rate-determining step. Explain your answer. Identify a step in that involves a redox reaction. Explain your answer in terms of oxidation numbers. Suggest the role of HIO in this mechanism. Explain your reasoning.
9701_s23_qp_42
THEORY
2023
Paper 4, Variant 2
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Aqueous acidified iodate(ions, IO3 –, react with iodide ions, as shown. IO3 – + 6H+ + 5I– 3I2 + 3H2O The initial rate of this reaction is investigated. Table 3.1 shows the results obtained. Table 3.1 experiment [IO3 –] / mol dm–3 [H+] / mol dm–3 [I–] / mol dm–3 initial rate / mol dm–3 min–1 0.0400 0.0150 0.0250 4.20 × 10–2 0.120 to be calculated 0.0125 7.09 × 10–2 The rate equation for this reaction is rate = k [IO3 –][H+]2[I–]2. Explain what is meant by order of reaction. Complete Table 3.2. Table 3.2 the order of reaction with respect to [IO3 –] the order of reaction with respect to [H+] the order of reaction with respect to [I–] the overall order of reaction Use your answer to to sketch lines in to show the relationship between the initial rates and the concentrations of [IO3 –] and [I–]. initial rate [IO3 –] initial rate [I–] Use data from Table 3.1 to calculate the rate constant, k, for this reaction. Include the units of k. k = units Use data from Table 3.1 to calculate the concentration of hydrogen ions, [H+], in experiment 2. [H+] = mol dm–3 This reaction is repeated in two separate experiments. The experiments are carried out at the same temperature and with the same concentrations of I– and IO3 –. One experiment takes place at pH 1.0 and the other experiment takes place at pH 2.0. Calculate the value of rate at pH 1.0 rate at pH 2.0 . value of rate at pH 1.0 rate at pH 2.0 = In aqueous solution, iron(ions react with iodide ions, as shown. 2Fe3+ + 2I– 2Fe2+ + I2 The initial rate of reaction is first order with respect to Fe3+ and second order with respect to I–. The mechanism for this reaction has three steps. Each step involves only two ions reacting together. Suggest equations for the three steps of this mechanism. Identify the rate‑determining step. step 1 step 2 step 3 rate‑determining step =
9701_s23_qp_43
THEORY
2023
Paper 4, Variant 3
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In aqueous solution, persulfate ions, S2O8 2–, react with iodide ions, as shown in reaction 1. reaction 1 S2O8 2– + 2I– 2SO4 2– + I2 The rate of reaction 1 is investigated. A sample of S2O8 2– is mixed with a large excess of iodide ions of known concentration. The graph in shows the results obtained. 0.000 0.100 0.200 0.300 0.400 0.500 0.600 0.700 0.800 0.900 time / min [S2O8 2–] / mol dm–3 Use to determine the initial rate of reaction 1. Show your working. rate = mol dm–3 min–1 The rate equation for reaction 1 is rate = k [S2O8 2–] [I–]. Suggest why a large excess of iodide ions allows the rate constant to be determined from the half-life in this investigation. The reaction of persulfate ions, S2O8 2–, with iodide ions is catalysed by Fe2+ ions. Write two equations to show how Fe2+ catalyses reaction 1. equation 1 equation 2 Describe the effect of an increase in temperature on the rate constant and the rate of reaction 1. In aqueous solution, thiosulfate ions, S2O3 2–, react with hydrogen ions, as shown in reaction 2. reaction 2 S2O3 2– + 2H+ SO2 + S + H2O The rate of reaction is first order with respect to [S2O3 2–] and zero order with respect to [H+] under certain conditions. The rate constant, k, for this reaction is 1.58 × 10–2 s–1. Calculate the half-life, t 1 2, for reaction 2. t 1 2 = s The compound nitrosyl bromide, NOBr, can be formed as shown in reaction 3. reaction 3 2NO+ Br22NOBrThe rate is first order with respect to and first order with respect to . The reaction mechanism has two steps. Suggest equations for the two steps of this mechanism. State which is the rate-determining step. step 1 step 2 rate-determining step =
9701_s24_qp_41
THEORY
2024
Paper 4, Variant 1
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In aqueous solution, persulfate ions, S2O8 2–, react with iodide ions, as shown in reaction 1. reaction 1 S2O8 2– + 2I– 2SO4 2– + I2 The rate of reaction 1 is investigated. A sample of S2O8 2– is mixed with a large excess of iodide ions of known concentration. The graph in shows the results obtained. 0.000 0.100 0.200 0.300 0.400 0.500 0.600 0.700 0.800 0.900 time / min [S2O8 2–] / mol dm–3 Use to determine the initial rate of reaction 1. Show your working. rate = mol dm–3 min–1 The rate equation for reaction 1 is rate = k [S2O8 2–] [I–]. Suggest why a large excess of iodide ions allows the rate constant to be determined from the half-life in this investigation. The reaction of persulfate ions, S2O8 2–, with iodide ions is catalysed by Fe2+ ions. Write two equations to show how Fe2+ catalyses reaction 1. equation 1 equation 2 Describe the effect of an increase in temperature on the rate constant and the rate of reaction 1. In aqueous solution, thiosulfate ions, S2O3 2–, react with hydrogen ions, as shown in reaction 2. reaction 2 S2O3 2– + 2H+ SO2 + S + H2O The rate of reaction is first order with respect to [S2O3 2–] and zero order with respect to [H+] under certain conditions. The rate constant, k, for this reaction is 1.58 × 10–2 s–1. Calculate the half-life, t 1 2, for reaction 2. t 1 2 = s The compound nitrosyl bromide, NOBr, can be formed as shown in reaction 3. reaction 3 2NO+ Br22NOBrThe rate is first order with respect to and first order with respect to . The reaction mechanism has two steps. Suggest equations for the two steps of this mechanism. State which is the rate-determining step. step 1 step 2 rate-determining step =
9701_s24_qp_43
THEORY
2024
Paper 4, Variant 3
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Use In the late 19th century the two pioneers of the study of reaction kinetics, Vernon Harcourt and William Esson, studied the rate of the reaction between hydrogen peroxide and iodide ions in acidic solution. H2O2 + 2I– + 2H+ 2H2O + I2 This reaction is considered to go by the following steps. step 1 H2O2 + I– IO– + H2O step 2 IO– + H+ HOI step 3 HOI + H+ + I– I2 + H2O The general form of the rate equation is as follows. rate = ka[I–]b[H+]c Suggest how the appearance of the solution might change as the reaction takes place. Suggest values for the orders a, b and c in the rate equation for each of the following cases. case numerical value a b c step 1 is the slowest overall step 2 is the slowest overall step 3 is the slowest overall A study was carried out in which both and [H+] were kept constant at 0.05 mol dm–3, and [I–] was plotted against time. The following curve was obtained. 0.001 0.0009 0.0008 0.0007 0.0006 0.0005 0.0004 0.0003 0.0002 0.0001 120 150 180 210 240 270 300 time / s [I– ion] / mol dm–3 Examiner’s Use To gain full marks for the following answers you will need to draw relevant construction lines on the graph opposite to show your working. Draw them using a pencil and ruler. Calculate the initial rate of this reaction and state its units. rate = units Use half-life data calculated from the graph to show that the reaction is first order with respect to [I–]. Use the following data to deduce the orders with respect to and [H+], explaining your reasoning. / mol dm–3 [H+] / mol dm–3 relative rate 0.05 0.05 1.0 0.07 0.05 1.4 0.09 0.07 1.8 order with respect to = order with respect to [H+] = From your results, deduce which of the three steps is the slowest (rate determining) step.
9701_w08_qp_4
THEORY
2008
Paper 4, Variant 0
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Questions Discovered
220