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.