Ksp & Solubility Calculator
Turn a solubility product into molar solubility and back. Pick a salt from a 93-entry Ksp table or enter your own, find its solubility in mol/L and g/L, see how a common ion lowers it (solved exactly, with the usual approximation beside it), and check whether mixing two solutions makes a precipitate.
How to Use
- Choose what to work out: solubility from Ksp, Ksp from a measured solubility, the effect of a common ion, or whether a precipitate forms.
- Pick the salt. Its formula sets the ion counts a and b, and its Ksp at 25 °C is filled in from the table; you can overwrite it, or choose Custom salt and enter a, b and a molar mass.
- Enter the solubility (in mol/L, mmol/L, g/L or mg/L), the common-ion concentration, or the two solutions to mix, depending on the mode.
- Read the molar and mass solubility, the ion concentrations or Q against Ksp. The picture shows the saturated solution and a log scale comparing the numbers.
- Show Work gives the Ksp expression, each step, and for a common ion the exact answer beside the approximation and its error.
Worked Example
Silver chloride in pure water. AgCl(s) ⇌ Ag⁺ + Cl⁻, so Ksp = [Ag⁺][Cl⁻] = s². With Ksp = 1.6 × 10⁻¹⁰, s = √(1.6 × 10⁻¹⁰) = 1.265 × 10⁻⁵ mol/L. At 143.32 g/mol that is 1.81 mg in a litre.
The same salt in 0.10 M NaCl. Now Ksp = (s)(s + 0.10). Solving exactly gives s = 1.6 × 10⁻⁹ mol/L, 7,906 times less than in pure water. Here the shortcut s ≈ Ksp ÷ 0.10 is fine, because s is tiny next to 0.10 M.
The common mistake: forgetting the coefficients. For Ag₂CrO₄ the silver concentration is 2s, so Ksp = (2s)²(s) = 4s³. Writing Ksp = s³ gives s = ∛(9.0 × 10⁻¹²) = 2.08 × 10⁻⁴ M, which is wrong; s = (9.0 × 10⁻¹² ÷ 4)^(1/3) = 1.31 × 10⁻⁴ mol/L, and [Ag⁺] = 2.62 × 10⁻⁴ M.
Show Work
Formulas
Where Solubility Products Come From
Cato Guldberg and Peter Waage stated the law of mass action in 1864, and Walther Nernst applied it to sparingly soluble salts in 1889, showing that a salt dissolves until the product of its ion concentrations reaches a fixed value at a given temperature. That product, the solubility product, explained why adding a salt with a shared ion makes the first one less soluble, the common ion effect used throughout classical qualitative analysis.
The Ksp values here are from OpenStax Chemistry 2e, Appendix J, “Solubility Products” (CC BY 4.0), at 25 °C. Published tables disagree, sometimes by a lot, because these values are hard to measure: OpenStax lists AgCl as 1.6 × 10⁻¹⁰ and BaSO₄ as 2.3 × 10⁻⁸, and other tables give different numbers for both. The simple formulas also ignore activity coefficients, ion pairing and reactions of ions such as S²⁻ and CO₃²⁻ with water, so real solubilities can be higher.
About This Tool
This calculator works with any salt MₐXᵦ, including ions like Hg₂²⁺ and [Fe(CN)₆]⁴⁻: it turns Ksp into molar and mass solubility, a measured solubility into Ksp, and finds the solubility with a common ion by solving the full equation numerically, with the shortcut and its percentage error shown beside it. The precipitation mode mixes two solutions, compares Q with Ksp and works out how much solid forms and what is left in solution.
Everything runs in your browser; nothing is sent anywhere.
Related tools: Chemical Equilibrium (ICE Table) Calculator, Molarity Calculator, and Nernst Equation Calculator.
Frequently Asked Questions
How do you calculate molar solubility from Ksp?
Write Ksp in terms of s: for MₐXᵦ, [M] = as and [X] = bs, so Ksp = aᵃbᵇ s^(a+b). For Ag₂CrO₄, Ksp = (2s)²(s) = 4s³, so s = (9.0 × 10⁻¹² ÷ 4)^(1/3) = 1.31 × 10⁻⁴ mol/L, which is 0.0435 g/L at 331.73 g/mol.
Does a smaller Ksp always mean a less soluble salt?
Only for salts with the same ion ratio. AgCl has the larger Ksp, 1.6 × 10⁻¹⁰, yet dissolves to 1.265 × 10⁻⁵ M, while Ag₂CrO₄, Ksp 9.0 × 10⁻¹², dissolves to 1.31 × 10⁻⁴ M, 10.4 times more, because its Ksp contains s³ rather than s².
What is the common ion effect?
An ion already in solution pushes the equilibrium back towards the solid. AgCl in 0.10 M NaCl dissolves to only 1.6 × 10⁻⁹ M, 7,906 times less than in pure water. The usual shortcut, ignoring s next to the added ion, can fail: PbCl₂ in 0.010 M Cl⁻ gives 0.16 M by the shortcut but 0.01273 M exactly.
How do you tell if a precipitate will form?
Work out the ion concentrations after mixing and the ion product Q, then compare it with Ksp. Mixing 50 mL of 1.0 × 10⁻⁴ M Ag⁺ with 50 mL of 1.0 × 10⁻⁴ M Cl⁻ gives 5.0 × 10⁻⁵ M of each, Q = 2.5 × 10⁻⁹ > 1.6 × 10⁻¹⁰, so AgCl precipitates: about 0.535 mg of it.
How do you find Ksp from solubility?
Convert the solubility to mol/L, find each ion’s concentration and multiply them with their powers. PbI₂ dissolving to 0.70 g/L is 0.70 ÷ 461.0 = 0.001518 M, with [I⁻] = 0.003037 M, so Ksp = 0.001518 × 0.003037² = 1.4 × 10⁻⁸.
How do I use the Ksp & Solubility Calculator?
Just type your numbers. The answer shows up right away — there is no button to press. Change anything and it updates by itself.
Is it free? Does it work without internet?
Yes to both. It is free with no sign-up, and once the page has loaded it keeps working even with no internet.
Where does my data go?
Nowhere — every calculation runs on your own device. Nothing you enter is uploaded, logged, or stored.
Common Use Cases
Hard water and scale
CaCO₃ with Ksp 8.7 × 10⁻⁹ dissolves to 9.33 × 10⁻⁵ M in pure water, about 9.3 mg/L.
Mohr titration
With 0.010 M Cl⁻, AgCl starts to form at 1.6 × 10⁻⁸ M Ag⁺; Ag₂CrO₄ with 0.010 M CrO₄²⁻ needs 3.0 × 10⁻⁵ M, so the red chromate appears only after the chloride is used up.
Fluoride and teeth
CaF₂ (Ksp 4.0 × 10⁻¹¹) dissolves to 2.15 × 10⁻⁴ M, 16.8 mg/L, and 0.010 M Ca²⁺ cuts that 6.8-fold.
Antacids
Mg(OH)₂, Ksp 8.9 × 10⁻¹², dissolves to only 7.6 mg/L, which is why milk of magnesia is a suspension.
Golden rain demonstration
Mixing 100 mL each of 0.010 M Pb²⁺ and I⁻ gives Q = 1.25 × 10⁻⁷, nine times Ksp, and 138.3 mg of yellow PbI₂.
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