Question

Need step by step solution to following: Suppose that you want to determine the concentration of an aqueous solution of iron(II) sulfate using titration. You decide to use the oxidation of iron(II) to iron(III) with permanganate, 〖MnO〗_4^-, as the oxidizing agent in acidic solution. You take advantage of the fact that the reaction is quite fast, and the permanganate exhibits an intense purple color, while the other reactants and products are essentially colorless Consequently, if you add a solution of potassium permanganate to a solution iron(II) sulfate in acidic aqueous solution, the purple color of the potassium permanganate rapidly disappears until all of the iron (II) has been oxidized. You find that it takes 38.2 mL of 0.0100 M 〖KMnO〗_4 to titrate 25.0 mL of the 〖FeSO〗_4 solution. What is the concentration of the 〖FeSO〗_4 solution?

Answer 1

The balanced reaction for oxidation of Fe (11) to felii) by permanganate ion is - Mnoj +8H7+ 5 F2t 5 & 3 tt Ma 2+ + 4H2O. 1 mol 5 mol so one mol of KMnog Now we find the moles Moles - onidises 5 moles of Felely to Fe (111) of knnou required for titration. Molarity & volume in L 0.0100 MX 38.2 m2 x 14 1000mL = 3.828164 mores. . Moles of Fe (W) present in sovition = 5x 3.82210-moles. 21.91810?mo? Now, Molarity of Fe (504) = . Moles volume in L 1.91810-3 mol 25.0m2 x IL 10com L = 0.0764M. which is the required molarity of resoy solution.