Redox Titration of Iron (II) with Potassium Permanganate

Introduction

• We aim to determine the concentration of iron (II) ions using a redox titration.
• Initial unknown concentration of iron (II) in acidic solution (10 mL).

Procedure

• Add potassium permanganate (KMnO4) to the solution.
  • KMnO4 concentration = 0.02 M.
  • KMnO4 provides permanganate anions (MnO4-).
• Drip KMnO4 into the solution and observe redox reaction.

Redox Reaction Details

• Balanced Redox Reaction:
  • MnO4- (purple) + Fe²⁺ -> Mn²⁺ (colorless) + Fe³⁺ (colorless)
• Oxidation States:
  • Manganese (Mn) in MnO4-: +7 to Mn²⁺: +2 (reduction).
  • Iron (Fe) from Fe²⁺: +2 to Fe³⁺: +3 (oxidation).

Endpoint Identification

• As permanganate is added, the solution becomes colorless.
• Endpoint when a light purple color persists after a drop, indicating excess permanganate.

Calculations

• Volume of KMnO4 used: 20 mL (0.02 L).
• Moles of Permanganate Needed:
  • Molarity (M) = moles/L.
  • Moles of MnO4- = 0.02 M * 0.02 L = 0.0004 moles.
• Mole Ratio from Balanced Equation:
  • 1 MnO4- : 5 Fe²⁺.
• Moles of Iron (II) Calculated:
  • 5 * 0.0004 moles = 0.002 moles of Fe²⁺.
• Concentration of Iron (II):
  • Molarity (M) = moles/volume (L).
  • 0.002 moles / 0.01 L = 0.2 M.

Conclusion

• The concentration of iron (II) ions in the original solution is 0.2 M.
• Alternative method: Use the MV = MV equation with adjustments for non-1:1 mole ratio.
  • Preferred method: Step-by-step calculation to understand the process.