Solubility and Precipitation Notes

Solubility and Insoluble Compounds

  • Insoluble compounds do not dissolve in water; they don't dissociate into ions.
  • Soluble compounds dissolve and dissociate into ions in water. See solubility rules for specific ions.

Solubility Rules

  • Soluble Ionic Compounds: Generally contain Group I cations (Li+, Na+, K+, Rb+, Cs+), Cl-, Br-, I- (except with Ag+, Hg22+, Pb2+), F- (except with Group 2 metal cations, Pb2+, Fe3+), C2H3O2-, HCO3-, NO3-, ClO3-, and SO42- (except with Ag+, Ba2+, Ca2+, Hg22+, Pb2+, Sr2+).
  • Insoluble Ionic Compounds: Generally contain CO32-, CrO42- (except with Group 1 cations and NH4+), PO43-, S2- (except with Group 1 cations and Ba2+), and OH- (except with Group 1 cations and Ba2+).

Solubility Equilibrium

  • Even insoluble compounds dissolve to a small extent, establishing an equilibrium between the solid and its ions.
  • Example: AgCl(s) \rightleftharpoons Ag^+(aq) + Cl^–(aq); equilibrium constant K << 1.

Solubility Product Constant (Ksp)

  • K_{sp} is the solubility product constant; it represents the product of ion concentrations at equilibrium in a saturated solution.
  • For SrCrO4(s) \rightleftharpoons Sr^{2+}(aq) + CrO4^{2–}(aq), K{sp} = [Sr^{2+}][CrO4^{2–}] = 3.6 × 10^{-5} at 25°C.
  • K_{sp} has a fixed value at a given temperature.

Calculating Ksp and Molar Solubility

  • Fluorite example: CaF2(s) \rightleftharpoons Ca^{2+}(aq) + 2F^–(aq). If [Ca^{2+}] = 2.15 × 10^{–4} M, then K{sp} can be calculated.
  • Molar solubility is the solubility in moles per liter (M), calculated from K_{sp}.
  • Copper(I) bromide example: If the K_{sp} of CuBr is 6.3 × 10^{–9}, the molar solubility of copper bromide can be determined.
  • Calcium hydroxide example: If the K{sp} of Ca(OH)2 is 1.3 × 10^{–6}, the molar solubility of calcium hydroxide can be calculated.

Ksp and Gram Solubility

  • Chrome yellow (PbCrO4) has a solubility of 4.6 × 10^{–6} g/L. Convert to molar solubility (M) before calculating K_{sp}.

Ksp and Precipitate Formation

  • Predict precipitate formation using the reaction quotient, Q.
  • If Q > K_{sp}, a precipitate forms.
  • If Q < K_{sp}, the solution is unsaturated.
  • If Q = K_{sp}, the solution is saturated.

Precipitation Examples

  • Magnesium hydroxide: Given [Mg^{2+}] = 0.0537 M in seawater, determine if Mg(OH)2 precipitates when [OH^–] = 0.0010 M, given K{sp} = 8.9 × 10^{–12}.
  • Silver chloride: Determine if AgCl precipitates when equal volumes of 2.0 × 10^{–4} M AgNO_3 and NaCl are mixed. When calculating Q, remember to account for the dilution that occurs when the two solutions are mixed.

Selective Precipitation

  • Use differences in K_{sp} values to selectively precipitate ions.
  • Example: Separating Ba^{2+} and Mg^{2+} using CO3^{2–} ions. BaCO3 will precipitate first.

Precipitation of Silver Halides

  • Given a solution containing KBr and KCl, determine which precipitates first, AgBr or AgCl, upon adding AgNO_3.

Common Ion Effect

  • The presence of a common ion reduces the solubility of a salt.
  • Example: Barium sulfate is less soluble in 0.10 M Na2SO4 than in pure water due to the common ion SO_4^{2–}.
  • Calculate the molar solubility of cadmium sulfide (CdS) in a 0.010M solution of cadmium bromide (CdBr2). The Ksp of CdS is 1.0 × 10–28.

Lewis Acids and Bases

  • Lewis acid: Accepts an electron pair.
  • Lewis base: Donates an electron pair.
  • Lewis acid-base reaction forms a coordinate covalent bond.

Coordinate Covalent Bond

  • Formed when one atom provides both bonding electrons.

Lewis Acid-Base Adduct

  • Compound containing a coordinate covalent bond.

Complex Ions

  • Polyatomic ions with a central atom (usually a transition metal) surrounded by ligands (Lewis bases).
  • The formula must be determined experimentally.

Complex Ion Equilibria

  • Formation constant (K_f): Equilibrium constant for complex ion formation from its components.
  • Example: Cu^+(aq) + 2CN^–(aq) \rightleftharpoons Cu(CN)_2^–(aq).

Dissociation of a Complex Ion

  • Calculate the concentration of the silver ion in a solution that initially is 0.10 M with respect to Ag(NH3)2^+.

Dissolving Precipitates in Acid

  • Strong acids can dissolve compounds with low water solubility if H3O^+ reacts with the anion (or NH3).
  • Compounds containing hydroxide, weak base anions, or NH_3 are more soluble in acid.

Dissolving Hydroxide Containing Compounds with Strong Acids

  • Zn(OH)2(s) dissolving in strong acid: Zn(OH)2(s) + 2H^+(aq) \rightarrow Zn^{2+}(aq) + 2H_2O(l).

Dissolving Compounds with Conjugate Base of weak acid

  • Salts that dissociate to give the conjugate base of a weak acid dissolve more readily in acidic solution.

Complex Formation

  • Lewis bases can dissolve ionic compounds containing metal cations that form complex ions.
  • Overall: Zn(OH)2 (s) + 4NH3 (aq) \rightleftharpoons Zn(NH3)4^{2+} (aq) + 2OH^– (aq).

Calcium Carbonate

  • The dissolution of calcium carbonate in acidic solution is : CaCO3(s) + H3O^+(aq) = Ca^{2+}(aq) + HCO3^–(aq) + H2O(l). K = K{sp}/K{a2} = 180