Electrochemistry Question

Question

Manganese is used in the steelmaking process to produce a tough steel for use in earth-moving machinery. It is derived from pyrolusite ore, an impure form of manganese dioxide. To analyze an ore sample for its manganese dioxide content the following procedure is used. A 0.533g samle is treated with 1.651g of oxalic acid * dihydrate in an acidic medium. Following this procedure the excess oxalic acid is titrated with 0.1000M potassium permanganate. The titration requires 30.06mL of the permanganate solution. What is the precent manganese dioxide, by mass, in the ore?

Answer 1

This problem involves 2 separate chemical reaction/equations.  One for the redox involving MnO2 and oxalic acid, and a second for oxidation of excess oxalic acid by permanganate.

 

(1)  MnO2 + H2SO4 + H2C2O4 ===> MnSO4 + 2CO2 + 2H2O

  Mn goes from 4+ in MnO2 to 2+ in MnSO4, so it has been reduced and C goes from 3+ in O.A. to 4+ in CO2 (oxidized)

moles oxalic acid used = 1.651 g x 1 mole/126 g = 0.01310 moles O.A. dihydrate used

 

(2) 2MnO4^- + 5H2C2O4 + 6H^+ ===> 2Mn^2+ + 10 CO2 + 8H2O

moles MnO4^- used = 0.03006 L x 0.100 mol/L = 0.003006 moles 

excess moles H2C2O4 = 0.003006 x 5/2 = 0.007515 moles xs O.A.

 

0.01310 mole - 0.007515 mole = 0.005585 moles  O.A. to titrate MnO2

moles MnO2 = 0.005585

mass MnO2 = 0.005585 moles x 86.94 g/mole = 0.4856 g

% MnO2 by mass = 0.4856 g/0.533 g (x100%) = 91.1% by mass