Solubility Product Calculator

Free Solubility Product Calculator - calculate solubility product for chemistry problems. Accurate results using standard formulas.

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What it is and when to use it

The solubility product calculator compares the ion product of a sparingly soluble salt with its solubility product constant, Ksp. For a salt that dissolves into a cation and an anion, the ion product Q is built from the concentrations of those ions currently in solution, each raised to its stoichiometric coefficient. If Q is larger than Ksp the solution holds more dissolved ions than equilibrium allows, so solid should precipitate.

Use it to predict whether mixing two solutions will form a precipitate, to check whether a solution is saturated, or to find how many moles of a salt dissolve per litre in pure water. It is a standard tool in general chemistry, water-quality work and qualitative analysis, where selective precipitation is used to separate ions.

The formula and how it works

For a salt with the dissolution reaction AxBy(s) ⇌ xAy+(aq) + yBx-(aq), the calculator uses:

  • Q = [A]x [B]y, where [A] and [B] are the entered molar concentrations and x and y the coefficients.
  • Ksp is the equilibrium value of the same expression, supplied by you.
  • Q / Ksp above 1 means supersaturated, below 1 means unsaturated, and 1 means saturated.
  • s = (Ksp / (xx yy))1/(x+y) is the molar solubility in pure water, in mol/L.

For a one-to-one salt this simplifies to s = sqrt(Ksp).

Worked example

Take silver chloride, AgCl, with the commonly tabulated Ksp of 1.8 x 10-10 at 25 C. Suppose [Ag+] = 1 x 10-5 M and [Cl-] = 1 x 10-4 M, both coefficients 1.

Q = (1 x 10-5)(1 x 10-4) = 1 x 10-9. Dividing, Q / Ksp = 1 x 10-9 / 1.8 x 10-10 = 5.56. Since Q exceeds Ksp, AgCl should precipitate until the ion product falls to 1.8 x 10-10. The molar solubility in pure water is the square root of 1.8 x 10-10, about 1.342 x 10-5 mol/L. The calculator reports these same values.

Common mistakes and how to interpret the result

  • Using initial concentrations of separate solutions without accounting for dilution: mixing equal volumes halves each concentration before you compute Q.
  • Forgetting the exponents: for salts like CaF2 or Ag2CrO4 the coefficient must be entered, otherwise Q is wrong by orders of magnitude.
  • Applying the pure-water molar solubility when a common ion is present: the real solubility is lower in that case.
  • Assuming Q slightly above Ksp guarantees visible precipitate: supersaturation can persist without nucleation, and very small excesses may not produce visible solid.

Frequently Asked Questions

What is the difference between Q and Ksp?
Ksp is the equilibrium constant for dissolving a sparingly soluble salt, fixed at a given temperature. Q, the ion product, has the same form but uses the concentrations actually present at this moment. Comparing them tells you the direction the system will move: Q above Ksp drives precipitation, Q below Ksp allows more solid to dissolve.
Why do I enter coefficients?
The exponents in the Ksp expression come from the balanced dissolution equation. For calcium fluoride, CaF2 dissolves to one Ca2+ and two F-, so Ksp = [Ca2+][F-]^2 and the anion coefficient is 2. Leave both at 1 for salts such as AgCl or BaSO4 that dissociate one-to-one.
Does a common ion change the result?
Yes, and this calculator handles it automatically if you enter the total concentration of each ion actually in solution. Adding a soluble chloride raises [Cl-], which raises Q for AgCl and pushes it toward precipitating. The molar solubility shown is for pure water only, so it overstates solubility when a common ion is present.
How accurate is this for real solutions?
It treats concentrations as activities, which is a good approximation for dilute solutions. At higher ionic strength, activity coefficients fall below 1 and salts appear somewhat more soluble than the simple calculation predicts. Ksp values also vary with temperature, so use a value measured near your conditions.

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