A student titrates a \(25.00 \text{ mL}\) sample of an acidified \(\text{Fe}^{2+}\text{(aq)}\) solution of unknown concentration with a standardized \(0.0500 \text{ M } \text{KMnO}_4\text{(aq)}\) solution. The balanced equation for the reaction that occurs during the titration is shown below.

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

If the faint pink endpoint is reached after exactly \(10.00 \text{ mL}\) of the \(0.0500 \text{ M } \text{KMnO}_4\text{(aq)}\) titrant is added, what is the concentration of \(\text{Fe}^{2+}\text{(aq)}\) in the original sample?