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The question in question (no pun intended) asks:
The Ksp of PbCl2 is 1.6 x 10^-5 at 25 degrees C. What is the solubility of PbCl2 in 0.01M KCl?
First, let's find the molar solubility of PbCl2 with no common ion effect:
Ksp = 1.6 x 10^-5 = x(2x)^2
1.6 x 10^-5 = 4x^3
x = 0.0158
Now, if we solve the question asked using 0.01M Cl- as the common ion:
Ksp = 1.6 x 10^-5 = x(0.01)^2
1.6 x 10^-5 = 0.001x
x = 0.16
0.16 is the correct answer given in DAT Destroyer. My question is: why is the molar solubility of PbCl2 so much higher when you start with Cl- in solution? Is that not supposed to deter PbCl2 from dissolving? Doesn't this problem contradict what you would expect if you understand the common ion effect?
The Ksp of PbCl2 is 1.6 x 10^-5 at 25 degrees C. What is the solubility of PbCl2 in 0.01M KCl?
First, let's find the molar solubility of PbCl2 with no common ion effect:
Ksp = 1.6 x 10^-5 = x(2x)^2
1.6 x 10^-5 = 4x^3
x = 0.0158
Now, if we solve the question asked using 0.01M Cl- as the common ion:
Ksp = 1.6 x 10^-5 = x(0.01)^2
1.6 x 10^-5 = 0.001x
x = 0.16
0.16 is the correct answer given in DAT Destroyer. My question is: why is the molar solubility of PbCl2 so much higher when you start with Cl- in solution? Is that not supposed to deter PbCl2 from dissolving? Doesn't this problem contradict what you would expect if you understand the common ion effect?