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The concentration of iron (II) sulfate, FeSO4, can be determined through a redox titration. A 25.

The concentration of iron (II) sulfate, FeSO4, can be determined through a redox titration. A 25.00 mL aliquot is pipetted into an Erlenmeyer where it is acidified with sulfuric acid and then titrated with a standard solution of potassium dichromate, K2Cr2O7. . The reactants and products of the unbalanced reaction are:

Cr2O72-(aq) + Fe2+(aq) Cr3+(aq) + Fe3+(aq) (in acidic solution)

Question (a):

Write the balanced half-reactions and the overall balanced redox equation forthe reaction between dichromate ion and iron (II) ion in acidic solution.
Here is what I got:

Reduction half-reaction: Cr2O7^2- + 14H+ +6e^- ---> 2Cr^3 + 7H2O

Oxidation half-reaction: 6Fe^2+ ---> 6Fe^3+ + 6e^-

Redox equation: 6Fe^2+ (aq) + Cr2O7^2- + 14H+ ---> 6Fe^3+ (aq) +2Cr^3+ +7H2O
How do I calculate part 2?
From the data given above and the following data, calculate the molarity of the original iron sulfate solution.

DATA:
Volume of sample = 25.00 mL
Concentration of standard dichromate, Cr2O72-, solution = 0.04343 M
Initial burette reading = 0.47 mL
Final burette reading = 21.10 mL

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