The solubility of a sparingly soluble salt primarily depends on its solubility product (Ksp). This fundamental constant quantifies the extent to which a salt dissolves in a solvent, typically water, at a given temperature.
The Pivotal Role of Solubility Product (Ksp)
The solubility product (Ksp) is a specific equilibrium constant that represents the product of the concentrations of the ions in a saturated solution, each raised to the power of their stoichiometric coefficients in the balanced dissolution equation. For a sparingly soluble salt, a smaller Ksp value indicates lower solubility, while a larger Ksp value suggests higher solubility.
As stated in the reference, the "Solubility of sparingly soluble salt depend upon solubility product." This means that the intrinsic ability of a sparingly soluble salt to dissolve is directly tied to its Ksp value.
Consider the examples provided:
- MY and NY3 are two nearly insoluble salts.
- They both share the same Ksp value of 6.2×10⁻¹³ at room temperature.
This common Ksp value indicates that despite having different chemical formulas (MY vs. NY3), their solubility product is the same under these conditions. While their actual molar solubilities might differ based on their stoichiometry (e.g., Ksp = s² for MY, Ksp = 27s⁴ for NY3), the Ksp is the direct measure of the ionic product at saturation.
Here's a quick comparison:
| Salt | Formula Type | Ksp at Room Temperature |
|---|---|---|
| MY | 1:1 Salt | 6.2 × 10⁻¹³ |
| NY3 | 1:3 Salt | 6.2 × 10⁻¹³ |
This demonstrates that the Ksp is a characteristic property for a given salt at a specific temperature, defining its maximum ion concentration in solution.
Other Influencing Factors
While Ksp defines the intrinsic solubility, several other factors can significantly influence the observed or effective solubility of a sparingly soluble salt in a given system:
1. Temperature
Solubility is highly temperature-dependent.
- For most ionic solids, solubility increases with increasing temperature because dissolution is often an endothermic process (absorbs heat).
- However, for some salts, solubility may decrease with increasing temperature if the dissolution process is exothermic.
2. Common Ion Effect
The presence of a common ion (an ion already present in the solution that is also a component of the sparingly soluble salt) will decrease the solubility of the sparingly soluble salt. This is a direct consequence of Le Chatelier's Principle, where the equilibrium shifts to relieve the stress of the added common ion, leading to more precipitation and less dissolution.
- Example: Adding sodium chloride (NaCl) to a saturated solution of silver chloride (AgCl) will decrease the solubility of AgCl because both salts share the common chloride ion (Cl⁻).
3. pH of the Solution
The pH of the solution can affect the solubility of sparingly soluble salts, particularly if one of the ions produced upon dissolution is a conjugate base or acid.
- If the anion of the salt is a conjugate base of a weak acid (e.g., CO₃²⁻, S²⁻, PO₄³⁻), its solubility will increase in acidic solutions. The acid will react with the anion, removing it from solution and shifting the dissolution equilibrium to the right.
- If the cation is a conjugate acid (e.g., NH₄⁺), its solubility may be affected by pH, though less commonly for sparingly soluble salts.
4. Complex Ion Formation
The formation of stable complex ions can significantly increase the solubility of a sparingly soluble salt. If a metal cation from the salt can react with a ligand (e.g., ammonia, cyanide, thiosulfate) to form a soluble complex ion, it removes the free metal cation from the solution. This shifts the dissolution equilibrium of the sparingly soluble salt to the right, causing more of it to dissolve.
- Example: Silver chloride (AgCl), which is sparingly soluble, becomes much more soluble in the presence of ammonia (NH₃) due to the formation of the soluble diamminesilver(I) complex ion, [Ag(NH₃)₂]⁺.
In summary, while the solubility product (Ksp) provides the fundamental measure of a sparingly soluble salt's intrinsic solubility, external factors like temperature, common ions, pH, and complex ion formation can profoundly alter its observed solubility in a real-world scenario.
