CIE Nov 2015 v1 Paper 3 Q1

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In this experiment you will determine the ionic equation for the reaction of acidified potassium manganate(VII) with potassium iodide. Excess potassium iodide is used and the reaction produces iodine. The amount of iodine produced is measured by titration with sodium thiosulfate. PMT FA 1 is 0.0180 mol dm3 potassium manganate(VII), KMnO4. FA 2 is 1.00 mol dm3 sulfuric acid, H2SO4. FA 3 is 0.500 mol dm3 potassium iodide, KI. FA 4 is 0.100 mol dm3 sodium thiosulfate, Na2S2O3. starch indicator (a) MethodPipette 25.0 cm3 of FA 1 into a conical flask. Use the measuring cylinder to add 25 cm3 of FA 2 to the conical flask. Use the measuring cylinder to add 20 cm3 of FA 3 to the conical flask. Fill the burette with FA 4. Carry out a rough titration. When the colour of the mixture becomes yellow/orange, add a few drops of starch indicator. Then titrate until the mixture goes colourless. Record all your burette readings in the space below. The rough titre iscm3.Carry out as many accurate titrations as you think necessary to obtain consistent results. Make sure any recorded results show the precision of your practical work. Record in a suitable form below all of your burette readings and the volume of FA 4 added in each accurate titration. Keep FA 1 and FA 2 for use in Question 3 and FA 4 for use in Question 2. II III IV VI VII [7] (b) From your accurate titration results, obtain a suitable value for the volume of FA 4 to be used in your calculations. Show clearly how you have obtained this value. Volume of FA 4 required iscm3. [1]UCLES 2015 9701/31/O/N/15<br /> PMT (c) Calculations Show your working and appropriate significant figures in the final answer to each step of your calculations. (i) Calculate the number of moles of sodium thiosulfate in the volume of FA 4 calculated in (b). (ii) Use the equation below to calculate the number of moles of iodine that reacted with the sodium thiosulfate in the titration. moles of Na2S2O3 =mol I2 + 2Na2S2O3Na2S4O6 + 2NaI (iii) Use information on page 2 to calculate the number of moles of potassium manganate(VII) in FA 1 used in the titration. moles of I2 =mol (iv) From your answers to (ii) and (iii), calculate the number of moles of iodine produced by the reaction of 2.00 moles of potassium manganate(VII) with excess potassium iodide. moles of KMnO4 =mol (v) Using your answer to (iv), put a tick next to the ionic equation that represents the reaction between FA 1 and FA 3. moles I2 =mol 2MnO4+ 2I + 16H+I2 + 2Mn6+ + 8H2O2MnO4+ 4I + 16H+2I2 + 2Mn5+ + 8H2O2MnO4+ 6I + 16H+3I2 + 2Mn4+ + 8H2O2MnO4+ 8I + 16H+4I2 + 2Mn3+ + 8H2O2MnO4+ 10I + 16H+5I2 + 2Mn2+ + 8H2O2MnO4+ 12I + 16H+6I2 + 2Mn+ + 8H2OUCLES 2015 9701/31/O/N/15 [Turn over<br /> PMT (vi) Prove that the iodide ion has been oxidised in the equation that you selected in (v).[5] (d) (i) The error in calibration of the pipette you used is 0.06 cm3. Calculate the percentage error when measuring FA 1, using the pipette. percentage error =% (ii) A student suggested that the experiment would be more accurate if a pipette was used to measure solution FA 3. State and explain whether you agree with the student.[2] [Total: 15]UCLES 2015 9701/31/O/N/15<br />

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