Advanced Solubility Product Calculator

Calculator Inputs

Use the responsive calculator grid below. Large screens show three columns, medium screens show two, and mobile shows one.

Calculation mode

Select the chemistry task you need.

Compound name

Example: Silver chloride or Calcium fluoride.

Temperature (°C)

Stored for reference. This version does not temperature-correct Ksp.

Cation label

Example: Ag+, Ca2+, Pb2+.

Anion label

Example: Cl-, F-, SO4^2-.

Ion charges

Charges are displayed for user reference.

Stoichiometric coefficients

For CaF₂, enter 1 and 2.

Molar solubility, s (mol/L)

Used when calculating Ksp from solubility.

Ksp value

Used for solubility and Qsp comparisons.

Starting cation concentration (mol/L)

Enter a common-ion concentration or current cation concentration.

Starting anion concentration (mol/L)

Enter a common-ion concentration or current anion concentration.

Formula Used

General dissolution reaction:
M_aX_b(s) ⇌ aMⁿ⁺ + bX^m−

Solubility product:
Ksp = [Mⁿ⁺]^a[X^m−]^b

Pure-water molar solubility:
If the molar solubility is s, then [Mⁿ⁺] = as and [X^m−] = bs.
Therefore, Ksp = (as)^a(bs)^b.

Common-ion condition:
With starting concentrations C₀ and A₀, the equilibrium expression becomes:
Ksp = (C₀ + as)^a(A₀ + bs)^b

Ionic product:
Qsp = [Mⁿ⁺]^a[X^m−]^b
If Qsp < Ksp, the solution is undersaturated.
If Qsp = Ksp, the solution is at equilibrium.
If Qsp > Ksp, precipitation is favored.

Saturation index:
SI = log₁₀(Qsp / Ksp)

How to Use This Calculator

Choose a calculation mode that matches your chemistry problem.
Enter the compound name and ion labels for clear reporting.
Set the stoichiometric coefficients from the balanced dissolution reaction.
Enter molar solubility if you want the calculator to determine Ksp.
Enter Ksp when solving molar solubility or checking precipitation risk.
Add starting ion concentrations to model common-ion effects or an existing solution.
Press the calculation button to show the result below the header and above the form.
Use the export buttons to save the result table as CSV or PDF.

Example Data Table

Compound	Dissolution	Approx. Ksp	Approx. molar solubility	Common-ion example
AgCl	AgCl(s) ⇌ Ag⁺ + Cl⁻	1.8 × 10⁻¹⁰	1.34 × 10⁻⁵ mol/L	0.010 mol/L Cl⁻ greatly lowers solubility.
CaF₂	CaF₂(s) ⇌ Ca²⁺ + 2F⁻	3.9 × 10⁻¹¹	2.14 × 10⁻⁴ mol/L	Added fluoride suppresses dissolution strongly.
BaSO₄	BaSO₄(s) ⇌ Ba²⁺ + SO₄²⁻	1.1 × 10⁻¹⁰	1.05 × 10⁻⁵ mol/L	Sulfate-rich water reduces further dissolution.

Frequently Asked Questions

1) What does Ksp represent?

Ksp is the equilibrium constant for the dissolution of a sparingly soluble ionic solid. It connects equilibrium ion concentrations to the salt’s dissolution stoichiometry.

2) What is the difference between Ksp and Qsp?

Ksp is the equilibrium benchmark. Qsp is calculated from the solution’s current ion concentrations. Comparing them shows whether dissolution, equilibrium, or precipitation is favored.

3) Why do stoichiometric coefficients matter?

The coefficients become exponents in the solubility product expression. Even a small change in stoichiometry can change the relationship between molar solubility and Ksp significantly.

4) How does the common-ion effect change solubility?

A common ion raises one equilibrium concentration before the salt dissolves. That shifts the system toward the solid phase and lowers the extra amount that can dissolve.

5) Why can a salt with low Ksp still have noticeable solubility?

Ksp depends on both ion concentrations and stoichiometric exponents. A salt producing multiple ions can show a solubility pattern that is not obvious from Ksp alone.

6) Does temperature affect Ksp?

Yes. Ksp usually changes with temperature because dissolution has an enthalpy change. This calculator stores temperature for reference, but it does not apply a temperature correction automatically.

7) What does a positive saturation index mean?

A positive saturation index means Qsp is greater than Ksp. The solution is supersaturated, so precipitation is thermodynamically favored until equilibrium is restored.

8) When should I use the ionic product mode?

Use it when you already know the present ion concentrations and need to predict whether a precipitate can form, whether the system is balanced, or whether more solid may dissolve.