Solubility Calculator

Pick your salt type, enter a Ksp or a molar solubility, and this converts between them with the whole derivation written out: the dissolution equation, the s-expression, and your numbers substituted. Add a molar mass and you get grams per liter too. A salt dissolving is one specific equilibrium worked so often it earns its own tool; for any other reaction, our equilibrium constant calculator does the general job.

Ksp to molar solubility Molar solubility to Ksp

AB, like AgCl AB₂ or A₂B, like CaF₂ AB₃ or A₃B, like Fe(OH)₃ A₂B₃ or A₃B₂, like Ca₃(PO₄)₂

Or pick a salt (fills in the Ksp, the type and the molar mass for you)

Ksp (e-notation works: 1.77e-10)

Molar mass (g/mol) (optional: adds grams per liter; our molar mass calculator computes it from a formula)

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How this solubility calculator works

When a sparingly soluble salt sits in water, a little of it dissolves and the rest stays solid, and the balance point is an equilibrium like any other. Its constant is called the solubility product, Ksp, and the question this page answers is the conversion every general chemistry course drills: given Ksp, how much actually dissolves, and given how much dissolves, what is Ksp? Pick the salt's type, because the type decides the algebra, and the whole derivation is written out with your numbers: the dissolution equation, the s-expression, the substitution.

This is one specific case of the general machinery on our equilibrium constant calculator: a salt dissolving is worked so often, with its own conventions and its own classic mistakes, that it earns its own tool. Any other reaction belongs on the general page; the pure solid dropping out of the expression is the same rule on both.

The formula

AB: K_sp = s²
AB₂ or A₂B: K_sp = 4s³
AB₃ or A₃B: K_sp = 27s⁴
A₂B₃ or A₃B₂: K_sp = 108s⁵

s is the molar solubility, the moles of salt that dissolve per liter of saturated solution. The multipliers are not magic numbers: each one is the stoichiometry counted honestly. For AB₂, every s of dissolved salt makes s of A and 2s of B, so K_sp = (s)(2s)² = 4s³. The coefficient 2 appears twice, once inside the concentration and once as the exponent, and forgetting one of those appearances is the single most common lost point on this topic. The solid salt itself never appears in the expression, because a pure solid has an activity of 1.

Worked example

How soluble is silver chloride? AgCl is a 1 to 1 salt with K_sp = 1.77e-10 at 25 C.

AgCl(s) ⇌ Ag⁺(aq) + Cl^-(aq). Let s be the molar solubility: then [Ag⁺] = s and [Cl^-] = s, so K_sp = s × s = s², and s = √(1.77e-10) = 1.33e-5 mol/L.

With AgCl's molar mass of 143.32 g/mol, that is 1.33e-5 × 143.32 = 0.0019 g/L: about two milligrams of silver chloride in a liter of water, which is why the salt is the textbook example of "insoluble" while still, strictly, dissolving. Run the same page the other way and a measured solubility hands back the Ksp: 1.33e-5 squared is 1.77e-10 again.

Ksp does not rank solubility

Here is the trap this page most wants you to walk around. It feels obvious that a bigger Ksp means a more soluble salt, and across different salt types it is simply not true. CaF₂ has K_sp = 3.45e-11, five times smaller than AgCl's 1.77e-10. Yet its molar solubility is 2.05e-4 mol/L against AgCl's 1.33e-5: about 15 times more soluble, on a smaller Ksp. The reason is the algebra above. AgCl's solubility is a square root of a tiny number and CaF₂'s is a cube root of one, and a cube root lifts a tiny number harder than a square root does. The rule to keep: compare molar solubilities, never raw Ksp values, across different salt types. Within one type the ranking works fine.

The table below is sortable: click a column heading. Sort by Ksp and then by solubility and watch the order shuffle, which is the whole lesson in one click. The same inversion sits right at the top: PbCl₂'s Ksp of 1.70e-5 is smaller than CaSO₄'s 4.93e-5, and PbCl₂ still out-dissolves it better than two to one. Values are the standard compilation figures at 25 C.

Salt	Type	Ksp (25 C)	Molar solubility (mol/L)	vs AgCl
Lead(II) chloride, PbCl₂	AB₂ (4s³)	0.000017 1.70e-5	0.016198 1.62e-2	1,218
Calcium hydroxide, Ca(OH)₂	AB₂ (4s³)	0.00000502 5.02e-6	0.010787 1.08e-2	811
Calcium sulfate, CaSO₄	AB (s²)	0.0000493 4.93e-5	0.0070214 7.02e-3	528
Lead(II) iodide, PbI₂	AB₂ (4s³)	0.0000000098 9.80e-9	0.0013481 1.35e-3	101
Calcium fluoride, CaF₂	AB₂ (4s³)	0.0000000000345 3.45e-11	0.00020508 2.05e-4	15.4
Silver carbonate, Ag₂CO₃	A₂B (4s³)	0.00000000000846 8.46e-12	0.00012836 1.28e-4	9.6
Magnesium hydroxide, Mg(OH)₂	AB₂ (4s³)	0.00000000000561 5.61e-12	0.00011194 1.12e-4	8.4
Silver chromate, Ag₂CrO₄	A₂B (4s³)	0.00000000000112 1.12e-12	0.000065421 6.54e-5	4.9
Calcium carbonate, CaCO₃	AB (s²)	0.00000000336 3.36e-9	0.000057966 5.80e-5	4.4
Silver chloride, AgCl	AB (s²)	0.000000000177 1.77e-10	0.000013304 1.33e-5	1.0
Barium sulfate, BaSO₄	AB (s²)	0.000000000108 1.08e-10	0.000010392 1.04e-5	0.78
Silver bromide, AgBr	AB (s²)	0.000000000000535 5.35e-13	0.00000073136 7.31e-7	0.055
Calcium phosphate, Ca₃(PO₄)₂	A₃B₂ (108s⁵)	0.00000000000000000000000000000000207 2.07e-33	0.00000011390 1.14e-7	0.0086
Silver iodide, AgI	AB (s²)	0.0000000000000000852 8.52e-17	0.0000000092304 9.23e-9	0.00069
Aluminum hydroxide, Al(OH)₃	AB₃ (27s⁴)	0.0000000000000000000000000000000003 3.00e-34	0.0000000018257 1.83e-9	0.00014
Iron(III) hydroxide, Fe(OH)₃	AB₃ (27s⁴)	2.79e-39	0.00000000010084 1.01e-10	0.0000076

The common ion effect

A salt dissolves dramatically less in a solution that already contains one of its ions, and the arithmetic is short enough to do at a glance. Drop AgCl into 0.10 M NaCl. The chloride concentration is already 0.10 M, and the tiny extra chloride the AgCl adds changes it by so little that we can call it 0.10 exactly. The Ksp expression still has to hold: [Ag⁺] × [Cl^-] = 1.77e-10, so [Ag⁺] = 1.77e-10 ÷ 0.10 = 1.77e-9 mol/L. That is roughly 7,500 times less soluble than the 1.33e-5 mol/L in pure water. Same salt, same constant, wildly different solubility, because one factor of the product arrived pre-filled.

The shortcut leans on one assumption worth naming: that s is much smaller than 0.10, so the dissolved silver chloride's own chloride contribution can be ignored. Here s comes out at 1.77e-9, about fifty million times smaller than 0.10, so the approximation is magnificently legal. When the common ion concentration is not so dominant, the honest route is the full quadratic, but for a sparingly soluble salt in a deliberate excess of common ion, this shortcut is the standard and correct move.

What the number leaves out

This page computes the Ksp equilibrium alone, which is the exam's question and the first-order truth, and real beakers have three more things going on. Temperature: every value here is a 25 C figure, and Ksp moves with temperature like any equilibrium constant. pH: hydroxides and carbonates dissolve more in acid, because acid consumes their anion and Le Chatelier pulls more solid apart; Fe(OH)₃'s spectacularly tiny Ksp is really a statement about neutral water. Complex ions: AgCl in a large excess of chloride starts dissolving more again, as AgCl₂^- complexes form, so the common ion curve is not monotonic forever.

And one honest note about the constants themselves: published Ksp compilations differ by small factors for many salts, because measuring an equilibrium this dilute is genuinely hard and activity corrections nibble at the result. Your textbook's 1.8e-10 and this page's 1.77e-10 are the same physical claim at homework precision. If your concentrations started life as grams in a flask, our molarity calculator gets you to mol/L, and our molar mass calculator turns any formula into the g/mol figure the grams-per-liter conversion wants.

Sources

Where the numbers on this page come from. We go to the body that publishes the figure, not to another calculator. See how we verify.

Solubility Product Constants near 25 degrees C, CRC Handbook of Chemistry and Physics (standard compilation values). The Ksp values in the reference table are the widely taught compilation figures at 25 C. Compilations differ from one another by small factors for many salts, because the underlying measurements are hard; the page says so out loud, and the arithmetic works identically with whichever value your textbook prints.

Frequently asked questions

How do you calculate solubility from Ksp?

Write the dissolution equation, call the molar solubility s, express each ion's concentration in terms of s, substitute into the Ksp expression, and solve. For a 1 to 1 salt like AgCl, Ksp = s squared, so s is the square root of Ksp: the square root of 1.77e-10 is 1.33e-5 mol/L. For a salt like CaF2 that releases two fluorides per formula unit, Ksp = 4s cubed, so s is the cube root of Ksp over 4. The salt's formula decides the equation, which is why the pattern buttons come first on this page.

What is Ksp?

The solubility product constant: the equilibrium constant for a solid salt dissolving into its ions. Like every equilibrium constant it is a product of concentrations raised to their coefficients, and like every heterogeneous equilibrium the pure solid itself does not appear in the expression. A saturated solution sitting over undissolved solid is the physical situation Ksp describes, and it is fixed at a given temperature.

Why does the salt type change the formula?

Because the stoichiometry puts the solubility into the expression more than once. When CaF2 dissolves, every s mol/L of dissolved salt makes s of calcium and 2s of fluoride, so Ksp = (s)(2s) squared = 4s cubed. The 2 appears twice, once inside the concentration and once as the exponent, and dropping either one is the classic exam slip. A 1 to 1 salt has no such doubling, so its expression is just s squared.

Can a salt with a smaller Ksp be more soluble?

Yes, when the two salts have different patterns. CaF2 has a Ksp of 3.45e-11, five times smaller than AgCl at 1.77e-10, yet CaF2 is about 15 times more soluble: 2.05e-4 mol/L against 1.33e-5. The cube root in the CaF2 formula lifts a tiny number more than the square root does. The rule: compare molar solubilities, not Ksp values, whenever the salts are different types. Within one type, bigger Ksp does mean more soluble.

What is the common ion effect?

A salt dissolves less in a solution that already contains one of its ions. Put AgCl into 0.10 M NaCl and the chloride is already at 0.10 M, so the silver concentration only needs to be Ksp divided by 0.10, which is 1.77e-9 mol/L, roughly 7,500 times less than in pure water. It is Le Chatelier in action: the product of the two ion concentrations must not exceed Ksp, and one factor is already large.

Why does my textbook's Ksp differ slightly from this table?

Because compilations genuinely disagree by small factors. Solubility equilibria are measured at low concentrations where activity corrections, ion pairing and hydrolysis all nibble at the result, so one handbook prints 1.77e-10 for AgCl and another prints 1.8e-10. Neither is wrong for homework purposes. Use whichever value your course provides; the arithmetic on this page is identical, and the answer moves by the same small factor.

Does Ksp change with temperature?

Yes. Every equilibrium constant is a function of temperature and nothing else, and the values here are the standard 25 C figures. Most salts dissolve more readily when warm, so their Ksp rises with temperature, but not all: calcium carbonate and calcium sulfate go the other way, which is why kettles scale up with hot water rather than cold.