The concentration of a solution of iron (II) sulfate, FeSO_4, can be determined
ID: 473827 • Letter: T
Question
The concentration of a solution of iron (II) sulfate, FeSO_4, can be determined through a redox titration. A 50.00 mL sample of the solution is diluted to 250.00 mL with deionized water. A 25.00 mL 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_2Cr_2O_7. The reactants and products of the reaction (unbalanced) are: Cr_2O_7^2- (aq) + Fe^2+ (aq) rightarrow Cr^3+ (aq) + Fe^3+ (aq) {in acidic solution} 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. From the data given above and the following data, calculate the molarity of the original iron sulfate solution. Show your reasoning/calculations.Explanation / Answer
For the given data
(a) balanced half reactions,
Fe2+ ---> Fe3+ + e-
Cr2O7^2- + 14H+ + 6e- ---> 2Cr3+ + 7H2O
Complete balaned reaction
Cr2O7^2- + 6Fe2+ + 14H+ ---> 2Cr3+ + 6Fer3+ + 7H2O
(b) moles of Cr2O7^2- used = molarity x volume
= 0.04326 M x (19.97 - 0.51) ml
= 0.842 mmol
moles of FeSO4 present = 0.842 mmol x 6 = 5.052 mmol
Volume of sample titrated = 25 ml
molarity of dichromate solution = 5.052 mmol/25 ml
= 0.2021 M
molarity of original FeSO4 solution = 0.2021 M x 250/50
= 1.0105 M
Related Questions
drjack9650@gmail.com
Navigate
Integrity-first tutoring: explanations and feedback only — we do not complete graded work. Learn more.