In aqueous solution, iron(III) 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]

1. (a)(i)Explain what is meant by overall order of reaction.[1]
2. (a)(ii)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.[2]
3. (a)(iii)Use your answer to (a)(ii) to construct the rate equation for this reaction. rate[1]
4. (a)(iv)Use your answer to (a)(iii) and the data from experiment 1 to calculate the rate constant, k, for this reaction. Include the units of k. k = units[2]
5. (a)(v)Describe qualitatively the effect of an increase in temperature on the rate constant and on the rate of this reaction.[1]
6. (b)(i)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+. Fig. 4.1 shows a possible four-step mechanism for this reaction. [Figure 4.1] Suggest which of the steps, 1, 2, 3 or 4, in this mechanism is the rate-determining step. Explain your answer.[1]
7. (b)(ii)Identify a step in Fig. 4.1 that involves a redox reaction. Explain your answer in terms of oxidation numbers.[1]
8. (b)(iii)Suggest the role of HIO in this mechanism. Explain your reasoning.[1]

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