1. Solubility Equilibria
Ksp and Qsp

2. AgCl (s)  Ag+ (aq) + Cl- (aq) Ksp = [Ag+] [Cl-]
The dissociation equation is written as a reversible reaction. You need to use the table of ions and always remember to balance.
example: silver chloride
AgCl (s)  Ag+ (aq) + Cl- (aq)
Ksp = [Ag+] [Cl-]
Since the equilibrium constant of such a dissolving reaction is written as the product of the ion concentrations, it is usually called the “solubility product” & written as Ksp instead of Keq.

3. Meaning of the Ksp
Ksp for AgCl is 1.56 x This really means that in a saturated solution containing Ag+ and Cl- ions, the product of the concentrations of Ag+ and Cl- must always equal 1.56 x
Thus if [Ag+] [Cl-] > Ksp, the solution is “supersaturated” and some of the Ag+ & Cl- ions will recombine & precipitate out. If [Ag+] [Cl-] < Ksp, the solution is not saturated and more AgCl will dissolve.
Example: Show the Ksp expression for the dissolving of copper (II) iodate.
Iron (III) hydroxide Ksp values p.578 in text

4. Ksp Problem Types
Type 1: Finding an equilibrium concentration given the Ksp & other conditions
Example: What will be the [I-] in a saturated solution of AgI where [Ag+] is M. Ksp for AgI is 1.5 x 10-16

5. Type 2: Finding Ksp, given solubility in moles/L (M).
Example 1: The solubility of BaSO4 is 3.9 x 10-5 moles/L. Determine Ksp.
Example 2: Solubility of Al(OH)3 is moles/L. Determine Ksp.

6. Type 3: Find Ksp given solubility in g/L.
Example: the solubility of CaF2 is g/L. Find the Ksp of CaF2.

7. Type 4: Given Ksp, find the solubility in moles/L or g/L.
Example: The Ksp for Ag2S is 1.6 x Find its solubility in a) moles/L b) g/L

8. Type 5: Predicting if precipitation will occur (Qsp).
Example: Will there be a precipitate of AgCl formed when these are mixed: 250 mL of 0.01 M Ag2SO4 & 250 mL of 0.02 M FeCl3. Ksp of AgCl = 1.8 x 10-10