Calculate solubility product constants (Ksp) and ion concentrations for sparingly soluble ionic compounds. Pre-loaded with 10+ common compounds for quick analysis.
Select a compound and choose a calculation mode. Mode 1: Calculate Ksp from given ion concentrations. Mode 2: Calculate molar solubility and ion concentrations from Ksp.
Silver chloride is a classic example of a sparingly soluble salt. If the concentration of Ag⁺ is 1.34 × 10⁻⁵ M and Cl⁻ is 1.34 × 10⁻⁵ M, what is the Ksp?
Ksp = [Ag⁺][Cl⁻] = (1.34×10⁻⁵)(1.34×10⁻⁵) = 1.80 × 10⁻¹⁰
Barium sulfate is used in medical imaging. Given its Ksp = 1.1 × 10⁻¹⁰, what is the molar solubility and the concentration of each ion in a saturated solution?
For a 1:1 salt, s = √Ksp = √(1.1×10⁻¹⁰) = 1.05 × 10⁻⁵ M
[Ba²⁺] = [SO₄²⁻] = 1.05 × 10⁻⁵ M
Calcium fluoride is important in dental health. With Ksp = 3.9 × 10⁻¹¹, what are the ion concentrations in a saturated solution?
For a 1:2 salt, s = ∛(Ksp/4) = ∛(3.9×10⁻¹¹/4) = 2.14 × 10⁻⁴ M
Iron(III) hydroxide has a very low Ksp. Given [Fe³⁺] = 1.0 × 10⁻¹⁰ M and [OH⁻] = 3.0 × 10⁻¹⁰ M, does a precipitate form?
Q = [Fe³⁺][OH⁻]³ = (1.0×10⁻¹⁰)(3.0×10⁻¹⁰)³ = 2.7×10⁻³⁹
Ksp = 2.6×10⁻³⁹, so Q > Ksp → Precipitate forms!
Silver chromate is a red-brown precipitate. With Ksp = 1.1 × 10⁻¹², what is the molar solubility?
For a 2:1 salt, s = ∛(Ksp/4) = ∛(1.1×10⁻¹²/4) = 6.50 × 10⁻⁵ M
The solubility product constant (Ksp) is an equilibrium constant that describes the extent to which a sparingly soluble ionic compound dissolves in water. For a general salt AₘBₙ(s) ⇌ mAⁿ⁺(aq) + nBᵐ⁻(aq), the Ksp expression is Ksp = [Aⁿ⁺]ᵐ[Bᵐ⁻]ⁿ.
When a sparingly soluble salt dissolves, it establishes an equilibrium between the solid phase and its constituent ions in solution. The Ksp value indicates how soluble the compound is — smaller Ksp values mean lower solubility. Molar solubility (s) represents the moles of compound that dissolve per liter of solution to reach saturation.
Where s = molar solubility, m and n are stoichiometric coefficients
General formula for calculating molar solubility from Ksp
The reaction quotient Q is calculated the same way as Ksp but using the actual (non-equilibrium) ion concentrations. Comparing Q to Ksp tells you whether a precipitate will form:
The solubility product constant, Ksp, is an equilibrium constant that quantifies the solubility of a sparingly soluble ionic compound in water. It represents the product of the concentrations of its constituent ions, each raised to the power of its stoichiometric coefficient in the dissolution equation, at equilibrium with excess solid present.
For example, for the dissolution of calcium fluoride: CaF₂(s) ⇌ Ca²⁺(aq) + 2F⁻(aq), the Ksp expression is Ksp = [Ca²⁺][F⁻]². The value of Ksp is temperature-dependent and is typically determined experimentally. A smaller Ksp indicates that the compound is less soluble — it reaches saturation at lower ion concentrations.
Solubility calculations are fundamental in chemistry and related fields. They help predict whether a precipitate will form when solutions are mixed, understand mineral formation in geology, design water treatment processes, formulate pharmaceuticals, and analyze environmental samples. In analytical chemistry, solubility equilibria are crucial for gravimetric analysis, qualitative analysis of ions, and controlling precipitation reactions.
Our Solubility Calculator is valuable for a wide range of scenarios:
Solubility refers to the maximum amount of a compound that can dissolve in a solvent (usually expressed as molar solubility in mol/L or g/L). Ksp (solubility product constant) is the equilibrium constant for the dissolution reaction. While Ksp is directly related to solubility, they are different quantities. For a simple 1:1 salt like AgCl, Ksp = s², so s = √Ksp. But for salts with different stoichiometries (like CaF₂ or Ag₂CrO₄), the relationship is more complex.
Most dissolution reactions are endothermic (absorb heat), so increasing temperature generally increases solubility and Ksp values. The exact relationship is given by the van't Hoff equation. Ksp values are typically reported at 25°C (298 K). This calculator uses standard 25°C values unless otherwise specified.
The common ion effect refers to the decrease in solubility of a sparingly soluble salt when one of its constituent ions is already present in the solution. For example, AgCl is less soluble in a NaCl solution than in pure water because the additional Cl⁻ ions shift the equilibrium toward the solid phase, according to Le Chatelier's principle. This calculator assumes pure water as the solvent unless you enter non-zero initial ion concentrations.
The formula depends on the stoichiometry. For a salt AₘBₙ: s = (Ksp / (mᵐ × nⁿ))^(1/(m+n)). For 1:1 salts (AgCl, BaSO₄): s = √Ksp. For 1:2 salts (CaF₂, PbCl₂): s = ∛(Ksp/4). For 2:1 salts (Ag₂CrO₄): s = ∛(Ksp/4). For 1:3 salts (Fe(OH)₃, Al(OH)₃): s = ∜(Ksp/27). This calculator automatically applies the correct formula based on the compound stoichiometry.
Yes! You can enter any custom compound formula in the text field (e.g., PbI₂, SrSO₄, Mn(OH)₂). The calculator will parse the formula to determine the stoichiometry and let you perform calculations. You'll need to provide the Ksp value if using Mode 2 (calculate concentrations from Ksp). For Mode 1, you enter the ion concentrations directly.
When the reaction quotient Q exceeds Ksp, the solution is supersaturated with respect to the compound. This means the ion concentrations are higher than the equilibrium values, and a precipitate will form until Q equals Ksp. The calculator will show a warning when Q > Ksp to indicate that precipitation is thermodynamically favored under the given conditions.
⚠️ Important Note: This calculator provides approximate results based on ideal solution behavior and standard Ksp values at 25°C. Real-world solubility may differ due to temperature variations, ionic strength effects, complex ion formation, and other factors. The values shown are suitable for educational purposes and preliminary analysis. Always verify critical calculations with experimental data or authoritative references.