Question #44650

A $3.000 \text{ g}$ sample containing a mixture of \ce{Fe3O4}, \ce{FeC2O4}, and inert impurities is dissolved in dilute \ce{H2SO4} to prepare exactly $250.0 \text{ mL}$ of stock solution. A $20.00 \text{ mL}$ aliquot of this solution is heated to $70^\circ\text{C}$ and requires $42.35 \text{ mL}$ of $0.02000 \text{ M}$ \ce{KMnO4} for complete titration. Another $20.00 \text{ mL}$ aliquot is treated with Zn dust and dilute \ce{H2SO4} to completely reduce all iron to \ce{Fe^{2+}}. After filtering off the unreacted zinc, the solution is heated to $70^\circ\text{C}$ and requires $42.74 \text{ mL}$ of the same \ce{KMnO4} solution for complete titration. Calculate the mass of \ce{Fe3O4} in the original sample. (Molar mass of \ce{Fe3O4} = $231.53 \text{ g/mol}$). Apply strict rules of significant figures at all stages of your calculation.