# Ksp Calculator - Molar Solubility, Common-Ion Effect & Precipitation

> Free solubility product calculator with 80 Ksp values at 25 °C: Ksp to molar solubility, the common-ion effect, Q vs Ksp precipitation and hydroxide pH.

- URL: https://qilingtools.com/en/math/ksp
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- Last updated: 2026-10-26
- Price: Free, no sign-up
- Privacy: Runs locally in the browser; files are never uploaded

The solubility product Ksp is the equilibrium constant for a sparingly soluble salt dissolving. For MₓAᵧ(s) ⇌ x Mᶻ⁺ + y Aᶻ⁻ with molar solubility s, Ksp = (xs)ˣ(ys)ʸ = xˣyʸ·s^(x+y). Silver chloride has Ksp = 1.77×10⁻¹⁰, so s = √Ksp = 1.33×10⁻⁵ mol/L; silver chromate has a smaller Ksp of 1.12×10⁻¹², yet Ksp = 4s³ gives s = 6.54×10⁻⁵ mol/L, so it is actually more soluble — which is why Ksp values can only be compared directly for salts of the same type. This calculator includes about 80 common sparingly soluble compounds (silver, barium, calcium, lead and mercury(I) salts, metal hydroxides and sulfides) with 25 °C values mainly from the CRC Handbook. It converts Ksp to and from solubility in mol/L, g/L or g/100 mL; solves the common-ion effect exactly rather than only with the usual approximation (AgCl in 0.1 M NaCl dissolves to just 1.77×10⁻⁹ mol/L); mixes two solutions and compares the ion product Q with Ksp to decide whether a precipitate forms and how much; works out selective precipitation, including the classic Mohr titration of chloride with chromate indicator; and charts the pH at which each metal hydroxide starts to precipitate and is complete (10⁻⁵ M), so you can design a pH separation.

## How to use

1. Pick a mode: Ksp ↔ solubility, common-ion effect, will it precipitate, selective precipitation, or hydroxide pH
2. Choose a salt from the built-in table (or enter your own ions and Ksp), fill in concentrations and volumes, or click an example
3. Read the result with the dissolution equation, Ksp expression and step-by-step working, and copy any value

## Features

- About 80 Ksp values at 25 °C (CRC and Lange’s Handbook)
- Two-way Ksp ↔ molar solubility in mol/L, g/L and g/100 mL
- Exact common-ion solubility with the approximation error shown
- Mixing two solutions: Q vs Ksp, mass of precipitate and ions left
- Selective precipitation and the Mohr titration: which salt comes first and whether they separate
- Bar chart of hydroxide precipitation pH ranges and pH separation windows
- Same-family comparison table that flags where ranking by Ksp and by s disagree
- Worked steps and textbook examples in every mode

## FAQ

### How do I calculate molar solubility from Ksp?

For a salt MₓAᵧ with molar solubility s, [M] = xs and [A] = ys, so Ksp = xˣyʸ·s^(x+y). For 1:1 salts such as AgCl or BaSO₄, s = √Ksp; for 1:2 or 2:1 salts such as CaF₂ or Ag₂CrO₄, Ksp = 4s³ and s = ∛(Ksp/4); for Fe(OH)₃, Ksp = 27s⁴. Multiply s by the molar mass to get g/L.

### Does a smaller Ksp always mean a less soluble salt?

Only when the salts have the same ion ratio. AgCl has a larger Ksp (1.77×10⁻¹⁰) than Ag₂CrO₄ (1.12×10⁻¹²), but its molar solubility (1.33×10⁻⁵ M) is lower than that of Ag₂CrO₄ (6.54×10⁻⁵ M). Convert to s before comparing salts of different types; the comparison table highlights these reversals.

### How do I know if a precipitate will form when two solutions are mixed?

Work out each ion’s concentration after mixing (allowing for the combined volume), then the ion product Q = [M]ˣ[A]ʸ. If Q > Ksp a precipitate forms, if Q = Ksp the solution is just saturated, and if Q < Ksp nothing precipitates. Mixing equal volumes of 0.01 M Pb(NO₃)₂ and 0.01 M KI gives Q = 0.005 × 0.005² = 1.25×10⁻⁷, well above Ksp = 9.8×10⁻⁹, so PbI₂ precipitates.

### What counts as complete precipitation?

An ion is usually considered completely precipitated when less than 1.0×10⁻⁵ mol/L remains (quantitative analysis often uses less than 0.1% remaining). Using this, 0.01 M Fe³⁺ starts to precipitate as Fe(OH)₃ at about pH 1.8 and is complete by pH 2.8, while Cu²⁺ does not start until about pH 5.2, so adjusting to pH 3–4 removes iron from a copper solution.

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