AIM To estimate the amount of iron present in a given solution by potentiometric method. You are provided with a standard KMnO4 solution. PRINCIPLE This is a typical redox titratio... AIM To estimate the amount of iron present in a given solution by potentiometric method. You are provided with a standard KMnO4 solution. PRINCIPLE This is a typical redox titration using a potentiometer experiment. For a redox titration, an electrochemical cell has to be set by taking a known quantity of an aqueous solution of the reducing agent (here FAS) in a beaker provided with an indicator electrode – Platinum electrode and a reference electrode – calomel electrode connected to a potentiometer. The potential developed by the cell is measured. The electrode potential of the reference electrode is constant, and the potential of the indicator electrode depends on the concentration of Fe2+ to Fe3+. When a ferrous ion is titrated against potassium permanganate in an acidic medium, it gets oxidized into a ferric ion. From the Nernst equation, the electrode potential of any redox system is given by, E= Eo + 2.303 RT/nF log [Reactant] / [Product] For Fe2+/Fe3+ system, E=E Fe2+ / Fe3+ + 0.0591/n log [Fe3+]/[Fe2+] at 25 0C. To determine the emf of the above system, the following cell may be constructed, Pt/ Fe2+ / Fe3+ // Hg2Cl2; KCl (satd)/ Hg/Hg2Cl2. KCl (saturated) represents a reference electrode (saturated calomel electrode) whose single electrode potential is constant i.e., +0.2422V. Emf of the above cell is given by E cell = E RHS- E LHS = E cathode – E anode = [E Fe2+ / Fe3+ + 0.0591/n log [Fe3+]/[Fe2+] – [0.2422]. As the titration proceeds, i.e. when the KMnO4 is added to ferrous ammonium sulfate, Fe3+ (ferric) concentration increases and Fe2+ (ferrous) concentration decreases. Therefore, the half-cell potential and cell potential increase by the above Nernst equation. Near the endpoint, the rate of change in potential will be maximum due to the [Fe3+]/[ Fe2+] ratio changes significantly. On crossing the equivalence point, EMF changes in small increments and finally, it reaches saturation. This is due to the reason that, after an equivalence point, only Fe3+ ions are present since no Fe2+ ions are present in the solution. After the endpoint, the cell potential is governed by another redox couple [Mn7+]/[Mn2+]. The electrode potential depends on the concentration of H+ ions besides the concentration of Fe2+ and Fe3+. Therefore, to avoid the effect of change in the [H+] on the electrode potential, titration is carried out in the presence of large amounts of [H+]. Hence the cell potential is made to depend on concentrations of the Fe2+ and Fe3+ ratio. PROCEDURE Prepare 100 mL of 0.1N KMnO4 solution by taking 0.316 g of KMnO4 into a clean 100 mL standard flask and make up the solution to the mark with distilled water. Fill the burette with 0.1N KMnO4 solution. Transfer the given Fe2+ ion solution into a clean 100 mL standard flask. Make up the solution to the mark with distilled water. Pipette out 20 mL of the made-up solution into a 100 mL beaker and add 1 test tube of dilute sulphuric acid. Keep the platinum and reference electrodes in the solution and connect them to the potentiometer. Note the emf. Add 1 mL of std. KMnO4 from the burette and record the emf. During the emf determination, the solution should be under stirring conditions. Titrate by adding a standard KMnO4 solution in 1 mL increment from the burette. Note down the potential each time. Continue the addition until there is a sudden change in potential and take a few readings beyond that level. Plot a graph between the volume of KMnO4 and the corresponding potential. The midpoint of the steeper raising portion of the graph is the endpoint. Take a range of a few mL from the volume corresponding to the midpoint volume in fractions of 0.2 mL. Further, plot ΔE /ΔV against the volume of the titrant. The peak point of the graph is the endpoint. Read more