A student analyzes a \(1.12\text{ g}\) specimen of iron ore. The specimen is dissolved in acid, and all the iron in the resulting solution is converted to \(\text{Fe}^{2+}\text{(aq)}\). After the excess reducing agent is removed, the solution is diluted to \(250.0\text{ mL}\). A \(25.00\text{ mL}\) aliquot requires \(10.00\text{ mL}\) of \(0.0200\text{ M}\) \(\text{KMnO}_4\text{(aq)}\) to reach the endpoint. The reaction in the acidic solution is represented by the following equation.

\[ \text{MnO}_4^{-}\text{(aq)}+8\text{H}^{+}\text{(aq)}+5\text{Fe}^{2+}\text{(aq)}\rightarrow \text{Mn}^{2+}\text{(aq)}+4\text{H}_2\text{O(l)}+5\text{Fe}^{3+}\text{(aq)} \]

Assume that only \(\text{Fe}^{2+}\text{(aq)}\) reacts with the permanganate and that the molar mass of iron is \(56.0\text{ g/mol}\). What is the mass percent of iron in the ore specimen?