exp 7_prelablecture_permanganate

Purpose of Experiment 13

  • Illustrate a redox reaction.

  • Reuse a titration technique.

  • Determine the formula mass of an iron compound.

Standardization of KMnO4

  • KMnO4 solutions cannot be prepared directly by dissolving the solid in a given volume; they must be standardized using sodium oxalate (Na2C2O4).

  • In acidic solutions, oxalate ion (C2O4^2-) changes to oxalic acid (H2C2O4), which is oxidized by MnO4^- to produce CO2.

  • Oxidation half-reaction: H2C2O4 → 2CO2 + 2H+ + 2e−

  • Reaction slow at room temperature but rapid above 70°C.

Redox Reactions and Stoichiometry

  • In redox reactions, the loss of electrons (oxidation) must equal the gain of electrons (reduction).

  • MnO4^- reduction: gains 5 electrons.

  • H2C2O4 oxidation: loses 2 electrons.

  • To balance, multiply oxidation by 5 and reduction by 2:

    • 2MnO4^- + 5H2C2O4 → 10CO2 + 10H+ + 2Mn2+

  • Cancel identical electrons to find the overall balanced reaction.

Color Changes in Titration

  • MnO4^- ion is intensely colored, while Mn2+ is colorless in dilute solutions.

  • As KMnO4 is added during titration, its color disappears until the reducing agent is fully reacted.

  • Upon reaching the endpoint, excess MnO4^- gives a faint pink color that persists for about 30 seconds, indicating the end of titration.

Determining Percent of Iron in Unknown Compound

  • Iron is analyzed in the form of Fe2+ which is oxidized to Fe3+ during titration.

  • Fe2+ oxidation half-reaction: Fe2+ → Fe3+ + e−

  • Multiply iron half-reaction by 5 to balance with the permanganate reaction.

Standardization Procedure

  • Dissolve 0.09 to 0.1 g of Na2C2O4 in ~50 mL of 1 M H2SO4.

  • Heat the solution to 80-85°C.

  • Titrate with KMnO4 until the endpoint persists for 30 seconds.

  • Record initial and final burette readings to determine the volume of KMnO4 used.

Calculating Molarity of KMnO4

  • Convert the mass of sodium oxalate to moles:

    • moles = mass (g) / molar mass (134 g/mol)

  • Convert moles to millimoles by multiplying by 1000.

  • Use the balanced redox equation (2 moles of KMnO4 react with 5 moles of Na2C2O4) to find millimoles of KMnO4.

  • Calculate the molarity: Molarity = millimoles of KMnO4 / mL of KMnO4 used.

  • Average molarity calculated from three trials.

Titration of Iron Compound

  • Measure between 0.6 to 0.7 g of the unknown iron compound and dissolve in H2SO4 (no heating).

  • Titrate with KMnO4 to light pink endpoint.

  • Record KMnO4 volume from burette readings (final - initial).

Finding Mass of Iron in Sample

  • Use average molarity of KMnO4 from the first three titrations to convert volume used into millimoles of KMnO4 and then to millimoles of Fe2+ using the ratio from the balanced equation (1 mole of KMnO4 reacts with 5 moles of Fe2+).

  • Convert millimoles of Fe2+ to mass:

    • grams of Fe = moles of Fe × molar mass (55.85 g/mol).

Calculating Percentage by Mass of Iron

  • Percentage by mass of iron = (mass of Fe / mass of compound) × 100

  • Average percentage calculated from three trials.

Finding Formula Mass of Iron Compound

  • Average mass of iron compound obtained from the three trials.

  • Average moles of Fe2+ = total millimoles of Fe2+ for trials / 3 (convert to moles).

  • Formula mass = average mass of the iron compound / average moles of Fe2+.