ASR 2020 J2Prelim H2Chem P4 MS

Anderson Serangoon Junior College 2020 JC2 H2 Chemistry Paper 4 Suggested Mark Scheme 1 Determination of the oxidation number of iodine in a compound In this experiment you will determine the oxidation number of iodine in one of its compounds by titration. You are provided with the following: FA 1 is a 0.0197 mol dm–3 solution of the iodine–containing compound. FA 2 is dilute sulfuric acid, H2SO4. FA 3 is aqueous potassium iodide, KI. FA 4 is 0.105 mol dm–3 sodium thiosulfate, Na2S2O3. starch indicator FA 1 reacts with excess acidified potassium iodide to produce iodine, I2. This iodine is then titrated with aqueous sodium thiosulfate using starch indicator. (a) Titration of iodine formed against FA 4 1. Fill the burette with FA 4. (You will use the same burette for Question 2.) 2. Pipette 25.0 cm3 of FA 1 into a conical flask. 3. Using a 25 cm3 measuring cylinder, add 10 cm3 of FA 2 to the same conical flask. 4. Using the same measuring cylinder, add 20 cm3 of FA 3 to the mixture in the conical flask. The mixture will now be a red–brown colour, due to iodine produced. 5. Add FA 4 from the burette until the mixture becomes light brown. 6. Then add 10 drops of starch indicator. The mixture will change to a dark blue colour. 7. Continue titrating until the mixture becomes colourless. This is the end–point. 8. Carry out as many accurate titrations as you think necessary to obtain consistent results. 9. Record in a suitable form below all of your burette readings and the volume of FA 4 added in each accurate titration. (i) Results 1 2 initial burette reading / cm3 0.00 10.00 final burette reading / cm3 28.40 38.25 volume of FA 4 / cm3 28.40 28.35 [5]

(ii) 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 obtained this value. 28.35 28.40 2  = 28.38 cm3 The iodine produced required 28.38 cm3 of FA 4. [1] (b) 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 (a)(ii). n(S2O32–) used = 28.380.105 x mol1000 = 2.98 x 10–3 mol moles of Na2S2O3 = ............................ [1] (ii) The equation for the reaction of iodine with sodium thiosulfate is shown. I2(aq) + 2Na2S2O3(aq)  Na2S4O6(aq) + 2NaI(aq) Calculate the number of moles of iodine that reacted with the sodium thiosulfate calculated in (i). n(I2) reacted = ½ x 2.98 x 10–3 = 1.49 x 10–3 mol moles of I2 = ............................ [1] (iii) Use the information on page 2 to calculate the number of moles of iodine – containing compound in the 25 cm3 of FA 1 used in each titration. FA1 25n = x 0.01971000 = 0.000493 mol moles of iodine–containing compound FA 1 = ............................ [1] (iv) Use your answers to (b)(ii) and (b)(iii) to calculate the number of moles of iodine produced when 1 mole of the iodine –containing compound in FA 1 reacts with excess FA 3. Give your answer as an integer. 3 4 1.49 x 10no. of moles = 4.93 x 10   = 3 moles of I2 = ............................ [1]

(v) The anion in FA 1 is IOx– where x is the number of oxygen atoms present in the formula. Use your answer to (iv) to balance the ionic equation for the reaction between FA 1 and FA 3 under acidic conditions. Hence deduce the value of x in the formula IOx – . IO...– + ...... I– + ......H+  ...... I2 + ......H2O By observation: IO…– + …I– + …H+  3I2 + …H2O (use ans from (iv)) IO…– + 5I– + …H+  3I2 + …H2O (balance I) IO…– + 5I– + 6H+  3I2 + …H2O (balance charge) IO…– + 5I– + 6H+  3I2 + 3H2O (balance H) IO3– + 5I– + 6H+  3I2 + 3H2O (balance O) therefore x = 3 x = ............................[2] (vi) Calculate the oxidation state of iodine in FA 1. (If you were unable to calculate x in part (v), assume that x = 4.) Let the O.S. of iodine in FA 1 be a. a + 3(–2) = –1 a = +5 oxidation state of iodine = ............................[2] [Total: 14]

2 To investigate the effect of concentration changes on the rate of a reaction. FA 5 is 0.0500 mol dm–3 acidified iron(III) chloride, FeCl3. FA 6 is 0.0500 mol dm–3 potassium iodide, KI. Iron(III) ions oxidise iodide ions, I–, to iodine, I2 as shown in equation 1. equation 1 2Fe3+(aq) + 2I–(aq)  2Fe2+(aq) + I2(aq) In this experiment you will investigate how the rate of this reaction is affected by the concentration of Fe3+ ions. To do this you will add thiosulfate ions, S 2O32–, and starch indicator to a mixture of Fe3+(aq) and I–(aq). The iodine produced by the reaction reacts immediately with the thiosu lfate ions and is reduced back to iodide as shown in equation 2. equation 2 I2(aq) + 2S2O32–(aq)  2I–(aq) + S4O62–(aq) When all the thiosulfate has reacted, the iodine remaining in solution turns the starch indicator blue–black. The rate of reaction can be determined by timing how long it takes for the reaction mixture to turn blue–black. (a) Preparation of FA 7 by dilution of FA 4 FA 7 which contains 0.00500 mol dm–3 of sodium thiosulfate, Na2S2O3, can be prepared by diluting FA 4 from Question 1 in a 250 cm3 graduated flask. (i) The concentration of sodium thiosulfate, Na2S2O3, in FA 4 is 0.105 mol dm3. Show, by calculation, that the volume of FA 4 needed to prepare this diluted solution in the 250 cm3 graduated flask is 11.90 cm3. [1] Let V1 cm3 be the volume of FA 4 needed for dilution. c1V1 = c2V2 0.105 x V1 = 0.005 x 250 V1 = 11.90 cm3 (ii) Using the burette from Question 1 , measure 11.90 cm 3 of FA 4 into a 250 cm 3 graduated flask labelled FA 7. Make up to the mark with distilled water. Record the volume of FA 4 added to the flask in the space provided. volume of FA 4 added = …………………………….. cm3

(b) Experiment 1 1. Fill the burette labelled FA 5 with FA 5. 2. Run 20.00 cm3 of FA 5 into a 100 cm3 beaker. 3. Using the same measuring cylinder, add the following to the second 100 cm3 beaker:  10 cm3 of FA 6  20 cm3 of FA 7  10 cm3 of starch solution 4. Add the contents of the first beaker to the second beaker and start timing immediately. 5. Stir the mixture once and place the beaker on a white tile. 6. Stop timing as soon as the solution turns blue –black. Ignore any colour changes that occur before the intense blue–black colouration. 7. Record this reaction time , to the nearest second , in the space provided on page 8. 8. Rinse both beakers and shake dry. Rinse and dry the glass rod. Experiment 2 1. Run 10.00 cm3 of FA 5 into a 100 cm3 beaker. 2. Using another measuring cylinder, add 10. 0 cm 3 of distilled water into the beaker containing FA 5. 3. Using the measuring cylinder from Experiment 1, add the following to the second 100 cm3 beaker:  10 cm3 of FA 6  20 cm3 of FA 7  10 cm3 of starch solution 4. Add the contents of the first beaker to the second beaker and start timing immediately. 5. Stir the mixture once and place the beaker on a white tile. 6. Stop timing as soon as the solution turns blue –black. Ignore any colour changes that occur before the intense blue–black colouration. 7. Record this reaction time , to the nearest second , in the space provided on page 8. 1. Rinse both beakers and shake dry. Rinse and dry the glass rod.

Experiment 3 to 5 Carry out three further experiments to investigate how the reaction time changes with different volumes of FA 5. Remember that the combined volume of FA 5 and distilled water must always be 20.00 cm3. Do not carry out an experiment using 15.00 cm3 of FA 5. Do not use a volume of FA 5 that is