Solubility Equilibrium

Beaker View: Saturated Solution

A⁺ (cation) B⁻ (anion) AₐBᵦ(s) solid

Ion Product Q vs Solubility Product Ksp

Current Q: 0.00

Ksp: 0.00

Status: Saturated

Solubility (S) vs Temperature

Current S: 0.00 M

Temperature: 25 °C

Common Ion Effect: S vs [Common Ion]

With Common Ion: 0.00 M

Without Common Ion: 0.00 M

Reduction: 0%

Solubility Equilibrium Controls

Compound Parameters

Compound Type: Solubility Product Ksp: 1.00e-10

Value is log₁₀(Ksp), range: 10⁻¹⁵ to 10⁻²

Temperature: 25 °C ΔH° (kJ/mol): 20

Positive = endothermic, Negative = exothermic

Ion Concentrations (Test for Precipitation)

[A⁺] Concentration: 0.010 M [B⁻] Concentration: 0.010 M Common Ion [Common]: 0.00 M

Simulates common ion effect on solubility

Q Calculated: 0.00

Will Precipitate? No

Solubility S: 0.00 M

Common Compounds

Simulation Control

Solubility Equilibrium Equations

Dissociation Reaction: AₐBᵦ(s) ⇌ aA^(b+) + bB^(a-)

Solubility Product: Ksp = [A]^a[B]^b

Ion Product (Q): Q = [A]^a[B]^b (compared to Ksp)

For AB (1:1): Ksp = S², S = √Ksp

For AB₂ (1:2): Ksp = 4S³, S = ³√(Ksp/4)

van't Hoff Equation: ln(Ksp₂/Ksp₁) = -ΔH°/R (1/T₂ - 1/T₁)

Common Ion [Common] S decreases when common ion present

What is Solubility Equilibrium?

Solubility equilibrium is a dynamic equilibrium between a solid salt and its dissolved ions in a saturated solution. At equilibrium, the rate of dissolution equals the rate of precipitation. The solubility product constant (Ksp) quantifies this equilibrium and is temperature-dependent. When the ion product Q exceeds Ksp, precipitation occurs. When Q is less than Ksp, more solid can dissolve. The common ion effect reduces solubility when one of the ions is already present in the solution.

Solubility Product Constant (Ksp)

Definition: Ksp is the equilibrium constant for dissolution of a sparingly soluble salt.
Expression: For AₐBᵦ(s) ⇌ aA + bB, Ksp = [A]^a[B]^b (solid not included).
Temperature Dependence: Ksp increases with temperature for endothermic dissolution, decreases for exothermic.
Precipitation Criterion: If Q > Ksp, precipitation occurs; if Q = Ksp, at equilibrium; if Q < Ksp, unsaturated.
Practical Use: Predicts precipitate formation and calculates maximum solubility.

Common Ion Effect

Definition: Decreased solubility when one ion is already present.
Le Chatelier's Principle: Adding common ion shifts equilibrium toward solid.
Calculation: For AB with [A]₀: Ksp = (S + [A]₀)(S) ≈ [A]₀ × S, so S = Ksp/[A]₀.
Applications: Qualitative analysis, water purification, pharmaceutical formulation.

Temperature Effect on Solubility

Endothermic: Most salts (KNO₃, NH₄Cl) absorb heat. Solubility increases with T.
Exothermic: Some salts (Ce₂(SO₄)₃, Li₂CO₃) release heat. Solubility decreases with T.
van't Hoff: ln(Ksp₂/Ksp₁) = -ΔH°/R (1/T₂ - 1/T₁).
Industrial: Temperature-controlled crystallization for purification.

Precipitation Dynamics

Supersaturation: Q > Ksp causes precipitation, but may be delayed (metastable).
Nucleation: Initial crystal formation requires energy, may need seed crystal.
Crystal Growth: Crystals grow by adding more ions from solution.
Kinetics: Rate depends on mixing, temperature, nucleation sites.
Applications: Removing interfering ions, heavy metal removal from wastewater.

Real-World Applications

Qualitative Analysis: Selective precipitation identifies ions (e.g., Group I cations).
Water Treatment: Lime softening precipitates Ca²⁺ and Mg²⁺ as carbonates.
Biological: Kidney stones form when calcium salts exceed solubility.
Geology: Stalactites form from CaCO₃ precipitation.
Pharmaceuticals: Drug solubility affects bioavailability.

Factors Affecting Solubility

Temperature: Most salts more soluble at higher T.
pH: Affects salts with basic anions (carbonates, phosphates).
Common Ion: Presence decreases solubility.
Complex Formation: Ligands (NH₃, CN⁻) increase solubility via complexes.
Particle Size: Smaller particles have higher solubility.