Solubility and Complex-Ion Equilibria Notes

WARNING

  • These are abbreviated sets of class notes provided to help you master the course material. They do not replace lectures as they lack key information shown on class slides, comments, examples, and illustrations presented in class.

Solubility Equilibria

  • The precipitation of ionic substances or the dissolving of slightly soluble salts are important in natural and industrial processes.
  • All salts, considered insoluble, dissociate to some degree in solution.
  • Solubility Rules for Ionic Compounds: Familiarize yourself with solubility rules including exceptions.

The Solubility Product Constant, Ksp

  • A slightly soluble ionic compound establishes an equilibrium between solid and its respective ions when mixed with water.
  • Example: For barium sulfate, BaSO4 (s) \rightleftharpoons Ba^{2+}(aq) + SO4^{2-}(aq)
  • The equilibrium constant for this process is called the solubility product constant, Ksp, defined as:
    Ksp = [Ba^{2+}][SO_4^{2-}]
  • For calcium phosphate, the expression is: Ksp = [Ca^{2+}]^3[PO_4^{3-}]^2
  • If you found the maximum solubility of strontium carbonate, SrCO_3, as 2.37 \times 10^{-5} mol/L in water at 25ºC, you can calculate its Ksp:
    • Ksp = [Sr^{2+}][CO_3^{2-}]
    • Given molar solubility x = 2.37 \times 10^{-5}, therefore:
    • Ksp = (2.37 \times 10^{-5})(2.37 \times 10^{-5}) = 5.6 \times 10^{-10}

Calculating Ksp from the Solubility

  • Setting up an ICE table is essential:
    • Initial/Equilibrium change for SrCO_3:
    • SrCO3 (s) \rightleftharpoons Sr^{2+}(aq) + CO3^{2-}(aq)
    • Initial concentrations: 0 and 0
    • Change: +x for each aqueous component
    • You determine the Ksp for various salts like PbCl_2:
      • Ksp = [Pb^{2+}][Cl^-]^2
      • Determined molar solubility from mass (7.95 g in 2.0 L) leads to:
      • Ksp = (0.0143)(0.0286)^2 = 1.2 \times 10^{-5}

Criteria for Precipitation

  • Use reaction quotient Q_c to determine if precipitation will occur by comparing it with Ksp.
  • Example with Fe(OH)_3:
    • Qc = [Fe^{3+}][OH^-]^3
    • If Qc > Ksp, precipitation occurs.

Factors Affecting Solubility

  • Common-Ion Effect: The presence of a common ion reduces solubility.
  • Example: Calculate the solubility of lead(II) sulfate in a sodium sulfate solution (0.25 M):
    • Ksp = [Pb^{2+}SO_4^{2-}] and solved to yield 7.3 \times 10^{-8} M in 0.25 M sodium sulfate.
  • Effect of pH on Solubility: Certain substances become more soluble in acidic conditions.
  • Fractional Precipitation: The process of selectively precipitating ions from a solution using a reagent.

Complex-Ion Equilibria

  • Transition metals can form complex ions via coordinate covalent bonds with ligands.
  • Formation constant K_f measures stability.
    • For example, for Ni^{2+} reacting with form [Ni(NH3)6]^{2+} is represented as:
      Ni^{2+}(aq) + 6NH3(aq) \rightleftharpoons [Ni(NH3)_6]^{2+}
  • Large K_f values indicate stable complexes formed readily.

Conclusion

  • Understand the principles of solubility equilibria, Ksp calculations, factors affecting solubility, and how complex ions form and dissolve in solutions to excel in this area of chemistry.