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