Question

A student followed the procedure of this experiment to determine the solubility product of zinc(II) iodate, Zn(IO3)2. Solutions of ZN(NO3)2 of known initial concentrations were titrated with 0.200 M KIO3 solutions to the first appearance of a white precipitate. The following data were collected.
Complete the table below and determine the solubility product constant.

[Zn(NO3)2]0, M Initial 0.226, 0.101 0.0452, 0.0118
[KIO3]0, M Titrant 0.200, 0.200, 0.200, 0.200
V0, mL of Zn(NO3)2 100.0, 100.0, 100.0, 100.0
V, mL of KIO3 titrant 12.9, 12.4, 13.0, 18.3

Answer 1

Answer:

Explanation:

Complete the table below and determine the solubility product constant.

[Zn(NO₃)₂] 0, M Initial 0.226, 0.101 0.0452, 0.0118

[KIO₃] 0, M Titrant 0.200, 0.200, 0.200, 0.200

V₀, mL of Zn(NO₃)₂100.0, 100.0, 100.0, 100.0

V, mL of KIO₃ titrant 12.9, 12.4, 13.0, 18.3

What needs to be determined are the following for each column/sample

V0 + V, mL

[Zn²⁺]  + [IO³⁻]    ⇄  [Zn²⁺][IO³⁻]2

log [Zn²⁺][IO³⁻]2

$\frac{\sqrt{Zn^{2+}}}{1 + \sqrt{Zn^{2+}}}$

Then there's the determination of Ksp itself

log Ksp =  Ksp of Zn(IO3)2 =

Considering this, which of the ion products is closest to Ksp, and why?

Finally, the titration volumes for the first three samples don't vary greatly. However the last sample is considerably larger. Why is this to be expected?

The attached figures shed more light on the solution to this problem