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Section 17.4: Solubility Equilibria

  • Polyprotic Acid Titration:

    • The pH at halfway to the equivalence point of a titration for a polyprotic acid is equal to the corresponding pKa value.
    • Example:
      • For H3PO3, pKa1 can be estimated from the titration curve.
      • First equivalence point: 50 mL of NaOH (Titrant) added.
      • Halfway (25 mL) to equivalence point gives pH ≈ 1.5, thus pKa1 = 1.5 (estimated).
      • Second equivalence point at 100 mL NaOH added, halfway at 75 mL gives pH ≈ 6.5, so pKa2 ≈ 6.5 (estimated).
      • Actual pKa values: pKa1 = 1.3, pKa2 = 6.7 (
        close to estimates).

  • Dominant Species in Solution:

    • At pH = 4: Dominant species could be H3PO3 or H2PO3^−, etc.
    • At pH = 11: Dominant species change, possibly to HPO3^2− or PO3^3−.

  • Solubility Equilibria Overview:

    • Equilibria involving the dissolution or precipitation of ionic compounds are heterogeneous.
    • Example in Nature: Tooth enamel dissolves in acid, creating tooth decay; kidney stones form from precipitation of salts.

  • Solubility Product Constant (Ksp):

    • A saturated solution has a solid in equilibrium with its dissolved ions.
    • Example: Ksp for BaSO4 is derived from the equilibrium:
      BaSO4(s) ⇌ Ba²⁺(aq) + SO4²⁻(aq)
      => Ksp = [Ba²⁺] [SO4²⁻]
    • Ksp values vary with temperature, indicative of how soluble a solid is in water.
    • Example Values: Ksp(BaSO4) = 1.1 * 10^(-10).

  • Distinguishing Solubility vs. Ksp:

    • Solubility (g/L): Amount of solute that can be dissolved in solvent.
    • Molar Solubility (mol/L): Number of moles that dissolve in 1 L of saturated solution.
    • Ksp: Indicates extent of dissolving in equilibrium.

  • Factors Influencing Solubility:

    • Solubility can be affected by the presence of common ions (Common-Ion Effect).
      • Example: The solubility of CaF2 decreases in the presence of Ca²⁺ or F⁻ ions in solution.
    • Changes in pH can also affect solubility, especially for hydroxide salts.
    • Example: Mg(OH)2 is more soluble in acidic solution due to hydrolysis effect.

  • Dissolution Calculations:

    • Ksp can be used to predict solubility and vice versa, but care is needed in calculations under various ionic concentrations in solution.
Practice Problems for Solubility:
  • Exercise: Write Ksp expression for CaF2.
  • Exercise: Find Ksp for various compounds from Appendix D and compare their molar solubility based on Ksp values.
  • Exercise: Predict the impact of adding a common ion on the solubility of given salts.
Example Calculation of Ksp and Solubility:
  • BaSO4 Example: Given Ba²⁺ concentration from a saturated solution to compute Ksp.
  • Hydroxide Solubility: Influence of solution pH on solubility of hydroxides.
    • Example: Calculate Ksp from solubility concentration of a saturated solution or given ion concentration.
17.5: Factors Affecting Solubility

  • Common Ion Effect:

    • Influences solubility of ionic compounds, reduces solubility when common ions are increased.

  • pH Effect:

    • Basic anions increase solubility in acidic conditions.

  • Complex Formation:

    • Strong Lewis bases increase solubility of metal salts by forming complex ions (higher solubility when interacting with metal ions).
    • Examples include interactions with NH3 or OH−.
Ksp and Solubility Summary:
  • High Ksp values indicate soluble salts, while low Ksp values suggest insolubility.
    • Practical applications in environmental chemistry (tooth decay, kidney stones, stalactites).