Key Concepts of Solubility Product Constants to Know for General Chemistry II

Understanding solubility product constants (Ksp) is key in General Chemistry II. Ksp helps us quantify how well sparingly soluble ionic compounds dissolve in water, linking solubility to various factors like temperature, pH, and the presence of common ions.

1. Definition of solubility product constant (Ksp)

   • Ksp is an equilibrium constant that quantifies the solubility of a sparingly soluble ionic compound in water.
   • It represents the product of the molar concentrations of the ions, each raised to the power of their coefficients in the balanced dissolution equation.
   • A higher Ksp value indicates greater solubility of the compound in solution.

2. Relationship between Ksp and solubility

   • Solubility (S) is the maximum amount of solute that can dissolve in a given amount of solvent at equilibrium.
   • Ksp can be expressed in terms of solubility for a given ionic compound, allowing for the calculation of solubility from Ksp values.
   • The relationship is particularly useful for determining the solubility of salts in different conditions.

3. Writing Ksp expressions for ionic compounds

   • The Ksp expression is derived from the balanced dissolution equation of the ionic compound.
   • For a compound like AB → A⁺ + B⁻, the Ksp expression is Ksp = [A⁺][B⁻].
   • Each ion's concentration is raised to the power of its stoichiometric coefficient in the dissolution equation.

4. Calculating Ksp from solubility data

   • To find Ksp, first determine the solubility (S) of the compound in mol/L.
   • Substitute the solubility into the Ksp expression, using the stoichiometry of the dissolution reaction.
   • For example, if S = 0.1 M for AB, then Ksp = (S)(S) for a 1:1 salt, resulting in Ksp = 0.01.

5. Predicting precipitation using Ksp

   • Precipitation occurs when the product of the ion concentrations exceeds the Ksp value.
   • Calculate the ion product (Q) using the concentrations of the ions in solution.
   • If Q > Ksp, a precipitate will form; if Q < Ksp, no precipitation occurs.

6. Common ion effect on solubility

   • The presence of a common ion decreases the solubility of a salt due to Le Chatelier's principle.
   • Adding a common ion shifts the equilibrium to the left, reducing the concentration of dissolved ions.
   • This effect is significant in applications like water treatment and analytical chemistry.

7. pH effects on solubility

   • The solubility of certain salts, especially those containing basic anions, can be affected by pH changes.
   • Lowering the pH (increasing acidity) can increase solubility by converting the anion to a less soluble form.
   • For example, the solubility of calcium carbonate increases in acidic solutions due to the formation of soluble calcium ions.

8. Temperature effects on Ksp

   • Ksp values can change with temperature, affecting the solubility of salts.
   • For most salts, Ksp increases with temperature, leading to higher solubility at elevated temperatures.
   • However, some salts may exhibit decreased solubility with increasing temperature, so it is essential to consider the specific salt.

9. Selective precipitation of ions

   • Selective precipitation involves adding a reagent to precipitate one ion while keeping others in solution.
   • This technique relies on differences in Ksp values; ions with lower Ksp will precipitate first.
   • It is commonly used in qualitative analysis to separate and identify ions in a mixture.

10. Comparing solubilities using Ksp values

    • Ksp values can be used to compare the solubility of different ionic compounds.
    • A higher Ksp indicates a more soluble compound, while a lower Ksp suggests lower solubility.
    • This comparison is useful in predicting which salts will dissolve more readily in a given solution.