04/18/2025 Chem Notes

Understanding Ksp and Ionic Compounds

• We start with solid ionic compounds in equilibrium with aqueous ions.
• Essential to split ionic compounds into individual ions for balanced reactions.
• Ksp (Solubility Product Constant) is the equilibrium constant for these specific reactions.
• Ksp expressions only include ions from the product side, excluding solids.

Ksp Values of Slightly Soluble Ionic Compounds

• Many ionic compounds considered insoluble are actually slightly soluble with low Ksp values.
  • Example: Silver sulfide has a Ksp of $8 imes 10^{-48}$, indicating very low solubility.
  • Example: Zinc iodate has a higher Ksp of $3.9 imes 10^{-6}$, indicating more solubility than silver sulfide.

Comparing Ksp Values

• Ksp values can be compared to infer relative solubility among ionic compounds.
• For compounds with the same number of ions in dissolution, a higher Ksp correlates with higher solubility.
  • Example:
  • Magnesium carbonate - Ksp = $x$
  • Lead sulfate - Ksp = $y$
  • Barium chromate - Ksp = $z$
  • As Ksp decreases, solubility decreases.

Dissociation and Solubility

• Ksp can be affected by the stoichiometry of the ions.
• For example:
  • Calcium hydroxide: Ca(OH)_2
    ightleftharpoons Ca^{2+} + 2OH^- produces 1 Ca and 2 OH.
  • Silver chromate: Ag2CrO4
    ightleftharpoons 2Ag^+ + CrO_4^{2-} produces 2 Ag and 1 CrO4.
• Cannot directly compare solubility of compounds with different numbers of ions.

Using ICE Tables

• Set up ICE (Initial, Change, Equilibrium) tables to calculate Ksp or solubility:
  • Start with solid and 0 aqueous ions initially.
  • Increase ions by amounts corresponding to stoichiometry from dissolving.
• Example - Magnesium hydroxide dissolving:
  • $[Mg^{2+}]$ equals themolirity found from solubility information, and $[OH^-] = 2 imes [Mg^{2+}]$.
• If $[Mg^{2+}] = 1.31 imes 10^{-4}$ M,
  • So $[OH^-] = 2.62 imes 10^{-4}$ M, from which Ksp can be calculated.

Calculating Ksp from Solubility

• Ksp calculation:
  • From earlier values, Ksp = $[Mg^{2+}][OH^-]^2$,
  • Resulting Ksp = $8.9951 imes 10^{-12}$.

Finding Molar Solubility from Ksp

• Reverse calculation of known Ksp to find molar solubility.
• Example with lead iodide:
  • $Ksp = [Pb^{2+}][I^-]^2 = 1.4 imes 10^{-8}$
  • Set up ICE table with conditions for dissolution.
  • Solve for $x$, where $x$ = molar solubility of lead ions and $2x$ = iodide ions.

Precipitate Formation

• Predicting whether a precipitate will form involves comparing Q (reaction quotient) to Ksp:
  • If Q > Ksp, a precipitate will form.
  • If Q < Ksp, no precipitate forms.
• Example of Strontium fluoride:
  • $Ksp = 4.3 imes 10^{-9}$,
  • Given concentrations of strontium and fluoride, calculate Q and compare with Ksp to predict precipitation.

Common Ion Effect

• Adding a common ion can reduce solubility due to Le Chatelier's principle.
• Example with lead chromate in sodium chromate solution:
  • Increased chromate ions push equilibrium left, decreasing solubility of lead chromate.

Example Calculations of Solubility

• Solubility in pure water:
  • $Ksp = 6.5 imes 10^{-6}$ for calcium hydroxide indicates solubility in pure water is $1.2 imes 10^{-2}$ M.
• Solubility in a 0.1 M calcium nitrate solution indicates reduced solubility, approx. $4 imes 10^{-3}$ M due to common ion effect.

Conclusion

• Mastery of Ksp and solubility principles is vital in solubility predictions and computations.
• Fundamental stoichiometry and equilibrium principles guide these calculations.