Solubility Product (Ksp) Calculator

Solubility Product Guide

What Is the Solubility Product?

For a sparingly soluble salt AB dissolving: AB(s) ⇌ A⁺(aq) + B⁻(aq). Ksp = [A⁺][B⁻]. If molar solubility = s mol/L, then [A⁺] = s and [B⁻] = s, so Ksp = s². For AB₂: PbCl₂ ⇌ Pb²⁺ + 2Cl⁻. [Pb²⁺] = s, [Cl⁻] = 2s. Ksp = s × (2s)² = 4s³. Large Ksp = relatively more soluble. Ksp only applies to saturated solutions of sparingly soluble salts — it does not apply to freely soluble salts like NaCl.

What should I know about Common Ksp Values?

Very sparingly soluble (Ksp < 10⁻¹⁵): AgI (8.5×10⁻¹⁷), PbS (3.0×10⁻²⁸), BaSO₄ (1.1×10⁻¹⁰). Sparingly soluble (Ksp 10⁻⁸ to 10⁻¹²): AgCl (1.8×10⁻¹⁰), Ag₂CrO₄ (1.1×10⁻¹²), CaF₂ (3.9×10⁻¹¹). Slightly soluble (Ksp 10⁻⁴ to 10⁻⁸): PbCl₂ (1.6×10⁻⁵), CaSO₄ (4.9×10⁻⁵). Applications: Ksp controls the precipitation of sparingly soluble salts in qualitative analysis and water treatment.

What should I know about The Common Ion Effect?

Adding a common ion reduces solubility — Le Chatelier's principle. If AgCl is dissolved in 0.1 mol/L NaCl instead of pure water: [Cl⁻] = 0.1 + s ≈ 0.1 (s is tiny). Ksp = [Ag⁺] × 0.1 = 1.8×10⁻¹⁰. [Ag⁺] = 1.8×10⁻⁹ mol/L (compared to 1.34×10⁻⁵ mol/L in pure water). The common ion suppresses solubility by a factor of approximately 7,000 in this example. Applications: water softening adds calcium hydroxide to precipitate calcium carbonate; qualitative analysis exploits the common ion effect.

Will Precipitation Occur?

Compare the ion product Q to Ksp: Q = [A⁺]initial × [B⁻]initial. If Q > Ksp: solution is supersaturated — precipitation will occur until Q = Ksp. If Q = Ksp: solution is exactly saturated — at equilibrium. If Q < Ksp: solution is unsaturated — no precipitation, dissolving can continue. Example: mix 50 mL of 2×10⁻⁴ mol/L AgNO₃ with 50 mL of 2×10⁻⁴ mol/L NaCl. After mixing (both halved): [Ag⁺] = [Cl⁻] = 10⁻⁴. Q = 10⁻⁸. Ksp(AgCl) = 1.8×10⁻¹⁰. Q > Ksp → precipitation occurs.