1. Solubility Equilibrium

2. The Solubility Product Constant
Most salts (ionic compounds) readily dissociate in water
E.g. Ca3(PO4)2(s) ↔ 3 Ca2+(aq) + 2PO43-(aq)
For the dissociation equilibrium equation:
AB(s) ↔ bB+(aq) + cC-(aq)
Ksp = [B+(aq)]b[C-(aq)]c
Recall: solids are not included in Ksp
Common Ksp values at SATP are listed on p. 802
Remember: Ksp is temperature dependent

3. Magnitude of Ksp Ksp >> 1 Product favoured Mostly ions
Soluble ionic compounds
Ksp << 1 Reactant favoured
Very low [ions]
Ionic compounds with low solubility

4. E.g. 1: Calculate Ksp for magnesium fluoride at 25◦C, given a solubility of 1.72 X 10-3 g/100mL

5. E.g. 2: Calculate the solubility of zinc hydroxide, in mol/L, given a Ksp of 7.7 X10-17 at 25◦C.

6. Practice!
P. 486 #1-4
Ksp worksheet (1st page only)

7. Predicting Precipitation
Trial ion product, Q
If Q = Ksp the system is at equilibrium, saturated solution
If Q < Ksp the system will shift right, unsaturated solution
If Q > Ksp the system will shift left, supersaturated solution
No precipitate
Precipitate

8. The Common Ion Effect
Equilibrium can be shifted by dissolving a salt with a common ion or a compound that reacts with one of the ions in solution (Le Chatelier’s Principle)
E.g. NaCl(s) ↔ Na+(aq) + Cl-(aq)
Explain the effect on the equilibrium system above when:
HCl(aq) is added
NaNO3(s) is added
Ca(NO3)2(aq) is added
Pb(NO3)2(s) is added

9. The Common Ion Effect
E.g. What is the solubility of PbCl2(s) in a 0.20 mol/L NaCl(aq) solution at SATP

10. Practice!
P. 489 #5-6
P. 492 #7-12
P. 493 #1-11
Second page of worksheet