Solubility Equilibria 16.6 AgCl (s) Ag + (aq) + Cl - (aq) K sp = [Ag + ][Cl - ]K sp is the solubility product constant MgF 2 (s) Mg 2+ (aq) + 2F - (aq) K sp = [Mg 2+ ][F - ] 2 Ag 2 CO 3 (s) 2Ag + (aq) + CO (aq) K sp = [Ag + ] 2 [CO ] Ca 3 (PO 4 ) 2 (s) 3Ca 2+ (aq) + 2PO (aq) K sp = [Ca 2+ ] 3 [PO ] 2 Dissolution of an ionic solid in aqueous solution: (Q= reaction quotient) Q = K sp Saturated solution Q < K sp Unsaturated solution No precipitate Q > K sp Supersaturated solution Precipitate will form Note: We are assuming ideal behavior. In reality, there may be hydrolysis or ion pairs that decrease solubility.

16.6

Molar solubility (mol/L) is the number of moles of solute dissolved in 1 L of a saturated solution. Solubility (g/L) is the number of grams of solute dissolved in 1 L of a saturated solution. 16.6

What is the solubility of silver chloride in g/L ? AgCl (s) Ag + (aq) + Cl - (aq) K sp = [Ag + ][Cl - ] Initial (M) Change (M) Equilibrium (M) s+s +s+s ss K sp = s 2 s = K sp  s = 1.3 x [Ag + ] = 1.3 x M [Cl - ] = 1.3 x M Solubility of AgCl = 1.3 x mol AgCl 1 L soln g AgCl 1 mol AgCl x = 1.9 x g/L K sp = 1.6 x

If 2.00 mL of M NaOH are added to 1.00 L of M CaCl 2, will a precipitate form? 16.6 The ions present in solution are Na +, OH -, Ca 2+, Cl -. Only possible precipitate is Ca(OH) 2 (solubility rules). Is Q > K sp for Ca(OH) 2 ? [Ca 2+ ] 0 = M [OH - ] 0 = 4.0 x M K sp = [Ca 2+ ][OH - ] 2 = 8.0 x Q = [Ca 2+ ] 0 [OH - ] 0 2 = 0.10 x (4.0 x ) 2 = 1.6 x Q < K sp No precipitate will form

What concentration of Ag is required to precipitate ONLY AgBr in a solution that contains both Br - and Cl - at a concentration of 0.02 M? AgCl (s) Ag + (aq) + Cl - (aq) K sp = [Ag + ][Cl - ] K sp = 1.6 x AgBr (s) Ag + (aq) + Br - (aq) K sp = 7.7 x K sp = [Ag + ][Br - ] [Ag + ] = K sp [Br - ] 7.7 x = = 3.9 x M [Ag + ] = K sp [Cl - ] 1.6 x = = 8.0 x M 3.9 x M < [Ag + ] < 8.0 x M What [Br - ] remains just before AgCl begins to precipitate? [Br - ] = K sp [Ag + ]

The Common Ion Effect and Solubility The presence of a common ion decreases the solubility of the salt. What is the molar solubility of AgBr in (a) pure water and (b) M NaBr? AgBr (s) Ag + (aq) + Br - (aq) K sp = 7.7 x s 2 = K sp s = 8.8 x NaBr (s) Na + (aq) + Br - (aq) [Br - ] = M AgBr (s) Ag + (aq) + Br - (aq) [Ag + ] = s [Br - ] = s  K sp = s x s = 7.7 x Ksp = [Ag + ] [Br - ]

pH and Solubility The presence of a common ion decreases the solubility. Insoluble bases dissolve in acidic solutions Additional of H + ions uses up OH - ions Insoluble acids dissolve in basic solutions Additional OH - ions use up H + ions Mg(OH) 2 (s) Mg 2+ (aq) + 2OH - (aq) K sp = [Mg 2+ ][OH - ] 2 = 1.2 x K sp = (s)(2s) 2 = 4s 3 4s 3 = 1.2 x s = 1.4 x M [OH - ] = 2s = 2.8 x M pOH = 3.55 pH = At pH less than Lower [OH - ] OH - (aq) + H + (aq) H 2 O (l) remove Increase solubility of Mg(OH) 2 At pH greater than Raise [OH - ] add Decrease solubility of Mg(OH)

Complex Ion Equilibria and Solubility A complex ion is an ion containing a central metal cation bonded to one or more molecules or ions. Co 2+ (aq) + 4Cl - (aq) CoCl 4 (aq) 2- K f = [CoCl 4 ] [Co 2+ ][Cl - ] 4 2- The formation constant or stability constant (K f ) is the equilibrium constant for the complex ion formation. Co(H 2 O) 6 2+ CoCl KfKf stability of complex

16.10

16.11

Qualitative Analysis of Cations 16.11

Flame Test for Cations lithium sodium potassiumcopper 16.11

Chemistry In Action: How an Eggshell is Formed Ca 2+ (aq) + CO 3 2- (aq) CaCO 3 (s) H 2 CO 3 (aq) H + (aq) + HCO 3 - (aq) HCO 3 - (aq) H + (aq) + CO 3 2- (aq) CO 2 (g) + H 2 O (l) H 2 CO 3 (aq) carbonic anhydrase