(1) The Ksp of PbCl2 is 1.6x10E-5 at 25C. what is the solubility of PbCl2 in 0.01 M KCl?
PbCl2--> Pb2+ +2Cl-
Ksp=(Pb2+)(Cl-)^2 -->1.6x10^-5=x(2x+.01)^2-->1.6x10^-5=x(0.01)^2
-->x=.16M
I know that the addition of an common ion will push the equilibrium in a direction which tends to reduce the concentration of that ion. That means the equilibrium will move to the left, and the solubility of the compound will reduce. But please check my calculation below. I will just calculate the solubility of PbCl2 without a common ion.
(2) Ksp=(Pb2+)(Cl-)^2-->1.6x10^-5=x(2x)^2-->1.6x10^-5=4X^3
-->x=0.016M
when I compared the 2 solubilities, the one with an common ion is more soluble, did I do anything wrong? Shouldn't (2) be more soluble than (1)?
PbCl2--> Pb2+ +2Cl-
Ksp=(Pb2+)(Cl-)^2 -->1.6x10^-5=x(2x+.01)^2-->1.6x10^-5=x(0.01)^2
-->x=.16M
I know that the addition of an common ion will push the equilibrium in a direction which tends to reduce the concentration of that ion. That means the equilibrium will move to the left, and the solubility of the compound will reduce. But please check my calculation below. I will just calculate the solubility of PbCl2 without a common ion.
(2) Ksp=(Pb2+)(Cl-)^2-->1.6x10^-5=x(2x)^2-->1.6x10^-5=4X^3
-->x=0.016M
when I compared the 2 solubilities, the one with an common ion is more soluble, did I do anything wrong? Shouldn't (2) be more soluble than (1)?