Solubility Product (Ksp) Calculator
Calculate Ksp from molar solubility, or find molar solubility from Ksp.
Supports ionic compounds with presets for AgCl, BaSO4, CaF2, and more.
The solubility product constant (Ksp) describes the equilibrium between an ionic solid and its dissolved ions.
For the dissolution: MₓN_y(s) ⇌ xM^(y+)(aq) + yN^(x-)(aq)
Ksp = [M^(y+)]^x × [N^(x-)]^y
From molar solubility (s mol/L):
[M^(y+)] = x × s
[N^(x-)] = y × s
Ksp = (xs)^x × (ys)^y
From Ksp to solubility (solving for s):
For 1:1 salt (MX): Ksp = s² → s = √Ksp
For 1:2 salt (MX₂): Ksp = (s)(2s)² = 4s³ → s = (Ksp/4)^(1/3)
For 1:3 salt (MX₃): Ksp = (s)(3s)³ = 27s⁴ → s = (Ksp/27)^(1/4)
Common Ksp values at 25°C:
| Compound | Formula | Ksp |
|---|---|---|
| Silver chloride | AgCl | 1.8 × 10⁻¹⁰ |
| Barium sulfate | BaSO₄ | 1.1 × 10⁻¹⁰ |
| Calcium fluoride | CaF₂ | 3.9 × 10⁻¹¹ |
| Lead iodide | PbI₂ | 9.8 × 10⁻⁹ |
| Calcium carbonate | CaCO₃ | 3.4 × 10⁻⁹ |
| Iron(II) hydroxide | Fe(OH)₂ | 4.9 × 10⁻¹⁷ |
| Aluminum hydroxide | Al(OH)₃ | 1.9 × 10⁻³³ |
Predicting precipitation: Compare the ion product (Q_sp) to Ksp:
- Q_sp > Ksp → precipitation occurs
- Q_sp < Ksp → solution is unsaturated (no precipitate)
- Q_sp = Ksp → saturated (at the edge of precipitation)
Worked example
Silver chloride, AgCl, has Ksp = 1.8 × 10⁻¹⁰. It dissolves one ion of each, so Ksp = s²:
s = √(1.8 × 10⁻¹⁰) = 1.34 × 10⁻⁵ mol/L
Multiply by the molar mass of 143.32 g/mol and that is about 1.9 mg per litre. Silver chloride is the classic “insoluble” salt, and this is what insoluble actually means: not zero, just small enough to ignore in most chemistry.
Why you cannot rank solubility by Ksp
This is the trap the table above sets, and almost everyone falls into it once. Barium sulfate has Ksp = 1.1 × 10⁻¹⁰ and calcium fluoride has Ksp = 3.9 × 10⁻¹¹, so barium sulfate looks the more soluble of the two. It is not. Work the solubilities out:
- BaSO₄ dissolves as one-to-one, so s = √(1.1 × 10⁻¹⁰) = 1.05 × 10⁻⁵ mol/L
- CaF₂ dissolves as one-to-two, so s = ∛(3.9 × 10⁻¹¹ / 4) = 2.14 × 10⁻⁴ mol/L
Calcium fluoride is about twenty times more soluble despite the smaller Ksp, because its exponent is different. Comparing Ksp values only tells you anything when the two salts share the same dissolution stoichiometry. Otherwise you have to convert both to solubility first.
Two caveats worth knowing
Ksp is temperature-specific, and the values in the table are for 25 °C. Most salts get more soluble as they warm, though a few, calcium sulfate among them, go the other way.
And solubility drops sharply if one of the ions is already present in the solution. Adding sodium chloride to a saturated silver chloride solution pushes chloride up, so silver has to come down to keep the product constant, and solid silver chloride precipitates. That is the common-ion effect, and it is the whole basis of washing a precipitate with a dilute solution of one of its own ions rather than with pure water.
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