Introduction to Solubility Equilibria
Quick Notes
- Ksp is the solubility product constant — it shows the equilibrium position of a dissolving ionic compound.
- A saturated solution is in equilibrium between the undissolved salt and its dissociated ions.
- We can use Ksp to:
- Calculate molar solubility.
- Compare the relative solubility of different salts.
- Predict precipitation from ion concentrations.
Full Notes
What Is Ksp?
When an ionic salt dissolves, it breaks into its ions. For example:
AB(s) ⇌ A+(aq) + B−(aq)
When a salt is partially soluble, only a small portion of it dissolves in water — meaning some, but not all, of its ions break away into solution. The process is reversible: as ions dissolve from the solid (dissolution), some also recombine to re-form the solid salt (precipitation).
Eventually, the solution reaches a point where the rate of dissolution equals the rate of precipitation. At this stage, the system is at equilibrium and the concentration of dissolved ions stays constant, even though the solid is still present.
We describe this equilibrium using the solubility product constant, Ksp. It reflects how soluble a salt is — a higher Ksp means greater solubility, while a lower Ksp indicates that only a small amount dissolves.
Writing the Ksp Expression
For the 1:1 salt above, the expression is:
AB(s) ⇌ A+(aq) + B−(aq)
Ksp = [A+][B−]
![AP Chemistry Dissolution at equilibrium with Ksp expression Ksp = [A+][B−]](images/kspconstant.png)
For salts with multiple ions (and stoichiometry), we need to include coefficients as powers.
For Example:
CaF2(s) ⇌ Ca2+(aq) + 2F−(aq) → Ksp = [Ca2+][F−]2
Molar Solubility
Molar solubility refers to the number of moles of a substance that dissolve in one liter of solution to form a saturated solution. It is a measure of how much solute can dissolve under equilibrium conditions.
We can relate molar solubility x to Ksp with an ICE table.
Problem: What is the molar solubility of PbCl2 if Ksp = 1.6 × 10−5?
Dissociation: PbCl2(s) ⇌ Pb2+(aq) + 2Cl−(aq)
| Species | Initial | Change | Equilibrium |
|---|---|---|---|
| Pb2+ | 0 | + x | x |
| Cl− | 0 | + 2x | 2x |
Ksp expression: Ksp = [Pb2+][Cl−]2 = x(2x)2 = 4x3
Solve: 4x3 = 1.6 × 10−5 → x3 = 4.0 × 10−6 → x ≈ 1.6 × 10−2 M.
Molar solubility of PbCl2 ≈ 1.6 × 10−2 mol dm−3.
Comparing Solubilities Using Ksp
The solubility product constant, Ksp, tells us how soluble a salt is in water – but care must be taken when comparing values:
- Same stoichiometry (1:1 vs 1:1): you may compare Ksp values directly.
Examples: AgCl ⇌ Ag+ + Cl−vs NaBr ⇌ Na+ + Br−.
A higher Ksp directly means greater solubility. - Different stoichiometries: compare molar solubilities, not Ksp directly.
Examples: CaF2 ⇌ Ca2+ + 2F− (1:2) vs AgCl (1:1).
Here, even if CaF₂ has a larger Ksp than AgCl, that doesn’t automatically mean it’s more soluble. You must calculate the molar solubility (the number of moles that dissolve per liter) for each salt based on their dissociation equations.
Key rule to remember: Same ion ratio then compare Ksp directly. Different ion ratio then calculate and compare the molar solubility.
Solubility Rules and Ksp
Common solubility rules (e.g., all sodium, potassium, ammonium, and nitrate salts are soluble in water) correspond to high Ksp values (≫ 1). Very low Ksp values (e.g., < 10−6) indicate low solubility, so such compounds are often called “insoluble” in practice.
Summary
- Ksp gives a quantitative measure of how much salt dissolves.
- Use ICE tables with the Ksp expression to calculate molar solubility.
- Compare Ksp values directly only when salts share the same stoichiometric ion ratio.