Solubility Product (Ksp) Calculator

Precipitation occurs when the ion product Q = [cation]^m × [anion]^n exceeds Ksp. If Q < Ksp, the solution is unsaturated and more solid can dissolve. If Q = Ksp, the solution is exactly saturated.

Only for compounds with the same stoichiometry (e.g., both 1:1 or both 1:2). For different stoichiometries, you must compare molar solubilities, as the mathematical relationship between Ksp and solubility changes.

Most Ksp values increase with temperature since dissolution is typically endothermic. Exceptions include some calcium salts (like CaSO₄) where solubility decreases at higher temperatures.

Adding a common ion (one already present in the equilibrium) shifts the equilibrium left by Le Chatelier's principle, decreasing solubility. Ksp itself does not change, but the molar solubility does.

Yes. For salts of weak acids (like CaCO₃), lower pH increases solubility because H⁺ reacts with the anion. For salts of strong acids, pH has minimal effect on solubility.

Ksp values range enormously: from ~10⁻² for slightly soluble salts (like CaSO₄) to ~10⁻⁵⁰ for extremely insoluble compounds (like HgS). Most common Ksp values are between 10⁻⁵ and 10⁻³⁰.

Ksp can be determined by measuring the concentration of dissolved ions in a saturated solution using techniques like AAS, ICP-OES, conductimetry, or gravimetric analysis.

No. Ksp specifically applies to ionic compounds that dissociate into ions upon dissolution. Molecular compounds (like sugar in water) do not have Ksp values.

pKsp = −log₁₀(Ksp). It converts the very small Ksp numbers to a more manageable scale. Higher pKsp means lower solubility.