Not sure what it is asking or how to complete

Question

Assuming that no equilibria other than dissolution are involved, calculate the concentration of all solute species in each of the following solution of salt in contact with a solution containing a common ion. Show that changes in the initial concentrations:

PbCl2(s) in 0.631 M MgCl2 (MgCl2 is strong electrolyte) Ksp PbCl2 = 1.6 × 10−5

Answer 1

This appears to be a problem dealing with solubility and Ksp. It appears they are asking you to find the concentration of Pb²⁺ and Cl^- in the presence of the common ion Cl^- which is present in the form of MgCl₂. The way the question is written, it isn't clear what is meant by "Show that changes in the initial concentrations:"

PbCl₂(s) <==> Pb²⁺(aq) + 2Cl^-(aq) ... this is the dissolution equilibria

Ksp = 1.6x10^-5 = [Pb²⁺][Cl^-]²

The [Cl^-] can be taken as 2 x 0.631 M = 1.26 M since MgCl₂ ==> Mg²⁺ + 2Cl^-

1.6x10^-5 = [Pb²⁺][1.26]²

1.6x10^-5 = 1.59 [Pb²⁺]

[Pb²⁺] = 1.0x10^-5 M

[Mg²⁺] = 0.631 M

[Cl^-] = 2 x 0.631 M = 1.26 M