Aqueous iron(II) ions, Fe2+(aq), are usually kept in acidic conditions to prevent them readily oxidising to aqueous iron(III) ions, Fe3+(aq). Fe2+(aq) ions react with Ag+(aq) ions in a redox reaction. The following equilibrium is established. Fe2+(aq) + Ag+(aq) ⇋ Fe3+(aq) + Ag(s) The concentration of Fe2+(aq) at equilibrium can be found by titration with a standard solution of aqueous potassium manganate(VII), KMnO4(aq). KMnO4(aq) is deep purple in colour. The equilibrium constant for the reaction can be found using the following equation. Kc = [Fe3+(aq)]eqm / ([Fe2+(aq)]eqm × [Ag+(aq)]eqm) A student carries out the experiment using the following instructions. step 1 Add 100.0 cm³ of 0.200 moldm⁻³ AgNO3(aq) to 100.0 cm³ of 0.200 moldm⁻³ Fe(NO3)2(aq) in a 500 cm³ conical flask and stopper the flask. Label the conical flask A. step 2 Leave conical flask A for four hours, shaking intermittently. Then leave conical flask A untouched for one hour. step 3 Use a pipette to transfer 25.00cm³ of the solution from conical flask A into a clean 250 cm³ conical flask. Label this conical flask B. step 4 Add 5 cm³ of 1.00 moldm⁻³ NaCl(aq) to the solution in conical flask B. A white precipitate of silver chloride forms. step 5 Use a measuring cylinder to add 20cm³ of 1.00moldm⁻³ sulfuric acid to conical flask B. step 6 Rinse a burette and fill it with a standard solution of KMnO4(aq). step 7 Add KMnO4(aq) to the mixture in conical flask B until an end-point is reached. step 8 Empty conical flask B and rinse it with distilled water ready for the next titration. The student repeats the titration until concordant readings are achieved.

1. (a)The student records their results in Table 1.1.
2. (a)(i)Complete Table 1.1.[1]
3. (a)(ii)Calculate a suitable mean titre to be used in the student's calculations. Show clearly how you obtain the mean titre. mean titre = .......................... cm³[1]
4. (b)State what is meant by a standard solution in step 6.[1]
5. (c)(i)Suggest why conical flask A is left for four hours in step 2.[1]
6. (c)(ii)Suggest why conical flask A is not shaken during the final hour in step 2.[1]
7. (d)Suggest why a measuring cylinder is the most appropriate apparatus to use for measuring sulfuric acid in step 5.[1]
8. (e)State what the burette should be rinsed with in step 6.[1]
9. (f)State the change of colour seen in the mixture in conical flask B at the end-point in step 7. from .......................... to ..........................[1]
10. (g)The student repeats the experiment using KMnO4(aq) at a lower concentration. The student obtains a larger mean titre. Suggest one reason why a larger titre is better than a smaller titre.[1]
11. (h)The equilibrium is shown. Fe2+(aq) + Ag+(aq) ⇋ Fe3+(aq) + Ag(s) When another student carries out the titration with 0.0200moldm⁻³ KMnO4(aq), the mean titre is 21.10cm³. The ionic equation for the reaction between MnO4⁻(aq) and Fe2+(aq) is shown. MnO4⁻(aq) + 8H+(aq) + 5Fe2+(aq) → 5Fe3+(aq) + 4H2O(l) + Mn2+(aq)
12. (h)(i)Calculate the concentration of Fe2+(aq) in the equilibrium mixture. [Fe2+(aq)]eqm = .......................... mol dm⁻³[2]
13. (h)(ii)Suggest why it is not necessary to measure the concentration of Ag+(aq) ions in the equilibrium mixture experimentally.[1]
14. (h)(iii)Determine the decrease in concentration of Fe2+(aq) from the initial solution. Hence, determine the concentration of Fe3+(aq) in the equilibrium mixture. If you were unable to obtain an answer to (h)(i), use [Fe2+(aq)]eqm = 0.0804 moldm⁻³. This is not the correct answer. [Fe3+(aq)]eqm = .......................... mol dm⁻³[1]
15. (h)(iv)Determine the value of Kc. Include units in your answer. Kc = .......................... units = ..........................[2]