$>=$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:

Question $($a$)$ Write the balanced half$-$reactions $($state whether oxidation or reduction half$-$

reaction$),$ 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 $=,50\mathrm{.}00mL$

Volume of diluted solution $=,250\mathrm{.}00mL$

Volume of diluted sample titrated $=,25\mathrm{.}00mL$

Standard dichromate solution: $[C{r}_2{O}_7^{2-}]=0\mathrm{.}04240M$

Initial burette reading $=,0\mathrm{.}83mL$

Final burette reading $=,24\mathrm{.}28mL$