Solubility Equilibrium
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
Magnitude of Ksp
Ksp >> 1
Product favoured
Mostly ions
Soluble ionic compounds
Ksp << 1
Reactant favoured
Very low [ions]
Ionic compounds with low
solubility
E.g. 1: Calculate Ksp for magnesium fluoride at
25◦C, given a solubility of 1.72 X 10-3 g/100mL
E.g. 2: Calculate the solubility of zinc hydroxide, in
mol/L, given a Ksp of 7.7 X10-17 at 25◦C.
Practice!
• P. 486 #1-4
• Ksp worksheet (1st page only)
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
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
The Common Ion Effect
E.g. What is the solubility of PbCl2(s) in a 0.20 mol/L
NaCl(aq) solution at SATP
Practice!
•
•
•
•
P. 489 #5-6
P. 492 #7-12
P. 493 #1-11
Second page of worksheet