Solubility Equilibria Ksp and Qsp

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.

Meaning of the Ksp Ksp for AgCl is 1.56 x 10-10 . 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 10-10 . 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

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 0.010 M. Ksp for AgI is 1.5 x 10-16

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 0.004 moles/L. Determine Ksp.

Type 3: Find Ksp given solubility in g/L. Example: the solubility of CaF2 is 0.0180 g/L. Find the Ksp of CaF2.

Type 4: Given Ksp, find the solubility in moles/L or g/L. Example: The Ksp for Ag2S is 1.6 x 10-49. Find its solubility in a) moles/L b) g/L

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