This solubility product calculator converts between Ksp and molar solubility for any salt AmBn — Ksp = (ms)m(ns)n — including the gram-per-litre solubility bench chemists actually weigh.
| Molar solubility s | – |
| Ksp | – |
| Solubility in g/L | – |
Defaults are AgCl (m 1, n 1, Ksp 1.8 × 10⁻¹⁰, MW 143.32). For CaF₂ use m 1, n 2 (Ksp 3.9 × 10⁻¹¹, MW 78.07); for Ag₂CrO₄, m 2, n 1. Pure-water solubility only — a common ion already in solution suppresses s (the common-ion effect), and pH matters for salts of basic anions.
Ksp: An Equilibrium Constant Wearing a Solubility Badge
The general machinery is our Equilibrium Constant Calculator. Part of our Chemistry Calculators collection.
Worked Examples (Verified)
- AgCl (defaults): s = √Ksp = √(1.8 × 10⁻¹⁰) = 1.34 × 10⁻⁵ M = 0.00192 g/L — a litre of saturated solution holds under 2 mg of dissolved silver chloride
- CaF₂ (m 1, n 2, Ksp 3.9 × 10⁻¹¹): s = (Ksp/4)^⅓ = 2.14 × 10⁻⁴ M — sixteen times AgCl’s solubility despite a SMALLER Ksp: cross-stoichiometry Ksp comparisons mislead, convert to s first
- The exponent structure is the lesson: Ksp = (ms)^m (ns)^n — each ion’s concentration (s × its subscript) raised to that subscript, because the dissolution equation’s coefficients become exponents
Solubility Discipline
- Never rank different-stoichiometry salts by raw Ksp: AgCl (1.8e-10) vs Ag₂CrO₄ (1.1e-12) — the “smaller” Ksp belongs to the MORE soluble salt in mol/L (s 6.5e-5). Convert to s, then compare.
- Common-ion effect: AgCl in 0.1 M NaCl dissolves to s ≈ Ksp/0.1 = 1.8e-9 M — the shared Cl⁻ suppresses it 10,000-fold (Le Chatelier with numbers).
- Salts of basic anions (F⁻, CO₃²⁻, OH⁻) dissolve more in acid — anion protonation pulls the equilibrium; pure-water Ksp math understates their real-world solubility.
- Q > Ksp predicts precipitation — the same quotient logic as the equilibrium page, applied to ion products.
Related Tools
- Equilibrium Constant Calculator — Kc, Kp & Q
- pOH Calculator — hydroxide solubility partner
- Chemistry Calculators — the full collection
Frequently Asked Questions
How do you find molar solubility from Ksp?
Write the dissolution (AmBn → m A + n B), set ion concentrations to ms and ns, and solve Ksp = (ms)^m (ns)^n for s: s = [Ksp / (m^m · n^n)]^(1/(m+n)). AgCl (1:1): s = √Ksp = 1.34 × 10⁻⁵ M. CaF₂ (1:2): s = (Ksp/4)^⅓. This calculator runs the formula both directions.
Why can’t you compare Ksp values directly?
Because the exponent pattern differs by stoichiometry. A 2:1 salt’s Ksp is a product of three concentration terms; a 1:1 salt’s is two. Identical Ksp numbers mean different solubilities across types, and a smaller Ksp can hide a larger s (Ag₂CrO₄ vs AgCl, worked in the notes above). Convert to molar solubility to compare fairly.
What is the common-ion effect?
Dissolving a salt into a solution already containing one of its ions suppresses its solubility: s ≈ Ksp/[common ion] for 1:1 salts. AgCl in 0.1 M chloride drops from 1.34e-5 M to ~1.8e-9 M. It’s Le Chatelier’s principle applied to the dissolution equilibrium.
Does pH affect solubility?
Yes, for salts whose anion is basic (fluorides, carbonates, hydroxides, sulfides): acid protonates the anion, depleting it and pulling more salt into solution. Salts of strong-acid anions (Cl⁻, NO₃⁻, ClO₄⁻) are essentially pH-proof. Ksp alone — this calculator’s scope — describes pure water.
How do you know if a precipitate forms?
Compute the ion product Q with the same formula as Ksp but with actual (mixed) concentrations: Q > Ksp → precipitate forms until Q = Ksp; Q < Ksp → solution is unsaturated and any solid dissolves. It’s the solubility dialect of the Q-vs-K test on our equilibrium page.
Is this solubility product calculator free?
Yes — free, instant, no sign-up.