A 0.5510-g sample consisting of a mixture of iron and iron(III) oxide was dissolved completely in acid to give a solution containing iron(II) and iron(III) ions. A reducing agent was added to convert all of the iron to iron(II) ions, and the solution was then titrated with the standardized KMnO₄ (0.04240 M); 37.50 mL of the KMnO₄ solution was required. Calculate the mass percent of Fe and Fe₂O₃ in the 0.5510-g sample. (Example 4.13 gives the equation for the reaction of iron(II) ions and KMnO₄.)

Data given in Example 4.13

Transcribed Image Text:

EXAMPLE 4.13 Using an Oxidation-Reduction Reaction in a Titration Problem The iron in a sample of an iron ore can be converted quantitatively to the iron (II) ion, Fe²+, in aqueous solution, and this solution can then be titrated with aqueous potassium permanganate, KMnO4. The balanced, net ionic equation for the reaction oc- curring in this titration is MnO4 (aq) + 5 Fe²+ (aq) + 8 H30+ (aq) →Mn²+ (aq) + 5 Fe³+ (aq) + 12 H₂O(l) purple colorless colorless pale yellow A 1.026-g sample of iron-containing ore requires 24.35 mL of 0.0195 M KMnO4 to reach the equivalence point. What is the mass percent of iron in the ore?