How to solve Concentration problem:

Question

A 50.00 mL solution containing an unknown amount of UO2−4 was passed through a Jones reductor, converting UO2−4− to U3+. Upon exposure to atmospheric oxygen, the U3+ was reoxidized to U4+.U4+. The solution was then diluted to 100.0 mL using 1 M H2SO4.  and titrated with 19.62 mL of 0.0806 M KMnO4 to reach the purple endpoint. A blank solution subjected to the same procedure required 0.14 mL.

 Balance the titration reaction and determine the concentration of UO2−4 in the original solution.

How would I solve for concentration?

Answer 1

The U^4+ that is in solution is oxidized to UO2^2+ (U^6+) by the permanganate. The first step in determining the concentration is to balance the redox reaction between in acidic solution using these to unbalanced half-reactions

oxidation: U^4+ ------> UO2^2+ + 2e-

reduction: MnO4^- + 5e- -----> Mn^2+

2 H2O (l) + 5 U^4+ (aq) + 2 MnO4^- (aq) ----> 5 UO2^2+ (aq) + 2 Mn^2+ (aq) + 4 H+ (aq)

Taking the blank titration into consideration the volume of MnO4^- solution required for the titration of the U^4+ is (19.62 - 0.14) = 19.48 mL

so the concentration of U^4+ in the 100 mL diluted solution was

19.84 mL MnO4^- x (0.0806 mmol MnO4^-/ 1 mL MnO4-) x 5 mmol U^4+ / 2 mmol MnO4^-) x (1/ 100.0 mL U^4+ )

dimensional analysis will show the result is mmol U4^+ / mL U^4+ or U^4+ M so 0.0400 M U^4+

Since the dilution was a two-fold dilution the concentration of UO2^2- in the original solutions is 0.0800 M.

3 significant figures because of the concentration of the permanganate.