| Mineral | Ksp (25°C) | s (mol/L) | Solubility (g/L) |
|---|---|---|---|
| Silver chloride (AgCl) | 1.8e-10 | 1.34e-5 | 0.00192 |
| Barium sulfate (BaSO4) | 1.1e-10 | 1.05e-5 | 0.00245 |
| Calcium fluoride (CaF2) | 3.9e-11 | 0.000214 | 0.0167 |
| Calcite (CaCO3) | 3.3e-9 | 0.000674 | 0.0675 |
| Magnesium hydroxide (Mg(OH)2) | 5.6e-12 | 0.000112 | 0.00652 |
For a sparingly soluble salt MaXb:
- Equilibrium: MaXb(s) ⇌ aM + bX
- Solubility product (ideal): Ksp = [M]a[X]b
- With molar solubility s in pure water: [M]=a·s, [X]=b·s
- So: Ksp = (a·s)a(b·s)b and s = (Ksp/(aabb1/(a+b)
This calculator also supports:
- Common-ion effect: adds an initial cation or anion concentration and solves for s numerically.
- Activities: uses Ksp = (γM[M])a(γX[X]free)b, with γ estimated by the Davies equation from ionic strength.
- Temperature option: adjusts Ksp using a van ’t Hoff form when ΔH is provided.
- Speciation option: for carbonate, phosphate, and sulfide, the free anion is [X]free=α·[X]total, where α depends on pH.
- Select a mineral preset, or choose Custom and enter your values.
- Confirm Ksp at 25°C, stoichiometry (a, b), charges, and molar mass.
- Optional: add a common ion concentration to model suppression of solubility.
- Optional: set ionic strength to estimate activity coefficients (γ).
- Optional: choose a speciation system and pH to estimate α for the free anion.
- Press Calculate Solubility. Results appear above the form.
- Use the export buttons to download CSV or PDF.
Ksp-to-solubility conversion in practice
Ksp links a solid to dissolved ions at equilibrium. For MaXb, the calculator solves for molar solubility s, then reports g/L and mg/L using molar mass. When a=b=1, s≈√Ksp in ideal water; for CaF2 (a=1, b=2), s scales with Ksp1/3. Use consistent mol/L inputs because Ksp is defined by activities of ions, not by the solid amount added.
Ionic strength and activity correction
Natural waters rarely behave ideally. The Davies equation estimates activity coefficients (γ) from ionic strength I, helping convert measured concentrations into activities. Fresh surface water often ranges from I≈0.001–0.01, groundwater 0.01–0.1, and seawater near 0.7. At I=0.10, γ for a monovalent ion is about 0.78, while γ for a divalent ion is about 0.37, which can shift calculated solubility by orders of magnitude for high-charge salts.
Common-ion suppression and mixing scenarios
Adding a common ion reduces solubility because the ion product approaches Ksp sooner. The calculator lets you specify an initial cation or anion concentration (for example, background Ca2+ or SO42−). This is useful for blend calculations: mixing two waters can raise a common ion above its original level and trigger precipitation. Track the saturation index (SI); SI>0 indicates supersaturation and likely scaling, while SI<0 indicates undersaturation.
pH-driven speciation for polyprotic anions
Carbonate, phosphate, and sulfide systems distribute across protonated forms, so only a fraction α exists as the fully deprotonated anion that enters Ksp. At pH 8.2, α for CO32− is only about 0.007, meaning bicarbonate dominates and free carbonate is limited. Raising pH increases α dramatically (near 0.98 by pH 12), often lowering apparent solubility for carbonate minerals and affecting metal sulfide stability in alkaline conditions.
Temperature option and reporting outputs
If you provide ΔH, the tool applies a van ’t Hoff adjustment to estimate Ksp at your temperature. Positive ΔH typically increases Ksp with temperature, while negative ΔH decreases it. Report both the “Ksp used” and the chosen temperature to keep results clearly reproducible. For documentation, export CSV for spreadsheets or PDF for lab notes, including inputs, γ values, α, equilibrium concentrations, and calculation timestamp when needed.
1) My reference shows pKsp. What should I enter?
Convert using Ksp = 10−pKsp. Enter the value at 25°C when possible, then apply temperature and ΔH options if you need a different condition.
2) Why does a common ion reduce solubility?
A pre‑existing cation or anion raises the ion product, so equilibrium is reached with a smaller added s. This is the classic common‑ion effect predicted by Le Châtelier’s principle.
3) When should I use ionic strength?
Use it for waters that are not very dilute, especially I above about 0.01 mol/L (groundwaters, brackish water, brines). Activity corrections become important for divalent and trivalent ions.
4) What does “supersaturated” mean in the result?
It means the initial ion product exceeds Ksp under your inputs. The solution is expected to precipitate until SI approaches zero, assuming precipitation kinetics and inhibitors are not limiting.
5) How do I choose the speciation system and pH?
Select carbonate, phosphate, or sulfide when the controlling anion can protonate. Enter pH to estimate α for the fully deprotonated form used in Ksp. Choose None for simple anions like Cl−.
6) Does this tool model complexation or multiple minerals?
No. It is a focused Ksp‑based estimate with optional activity and speciation factors. For full geochemical speciation, buffering, and multi‑mineral equilibria, use dedicated models and compare with measured alkalinity and ions.