The concentration of a solution of iron $($II$)$ sulfate, $FeS{O}_4,$ can be determined through a redox titration. A $50\mathrm{.}00mL$ sample of the solution is diluted to $250\mathrm{.}00mL$ with deionized water. A $25\mathrm{.}00mL$ aliquot is then pipetted into an Erlenmeyer where it is acidified with sulfuric acid and then titrated with a standard solution of potassium dichromate, $K_{2}C{r}_2{O}_7.$ The reactants and products of the reaction $($unbalanced$)$ are:

$C{r}_2{O}_7^{2-}(aq)+F{e}^{2+}(aq)C{r}^{3+}(aq)+F{e}^{3+}(aq),\{in$ acidic solution $\}$

Question $($a$)$ Write the balanced half$-$reactions $($state whether oxidation or reduction halfreaction$),$ and the overall balanced redox equation for the reaction between dichromate ion and iron$($II$)$ ion in acidic solution.

Question $($b$)$ From the data given above and the following data, calculate the molarity of the original iron sulfate solution. Show your reasoning$/$calculations$.$

DATA:

Volume of original sample $=$ Volume of diluted solution $=$ Volume of diluted sample titrated $=$ Standard dichromate solution: Initial burette reading $=$ Final burette reading $=$

$50\mathrm{.}00mL$

$250\mathrm{.}00mL$

$25\mathrm{.}00mL$

$0\mathrm{.}04204M$

$0\mathrm{.}21mL$

$21\mathrm{.}10mL$