Solubility vs Temperature Calculator

Calculate how solubility changes with temperature using van't Hoff equation for dissolution.

20°C=293K, 25°C=298K
Positive=endothermic, NaCl=+3.9, KNO3=+34.9
Please check your inputs and try again.

Why Solubility Changes with Temperature

Dissolving requires breaking solute–solute and solvent–solvent attractions and forming solute–solvent ones. The net heat absorbed or released is the enthalpy of solution, ΔHsol, and its sign determines everything about the temperature behaviour.

Treat dissolution as an equilibrium and Le Chatelier applies directly. If dissolving absorbs heat (ΔHsol positive, endothermic), heat behaves as a reactant — adding it drives more solute into solution. If dissolving releases heat, the reverse holds and warming reduces solubility.

ΔHsolProcessEffect of heatingExamples
Large positiveStrongly endothermicSolubility rises steeplyKNO3, NH4NO3, KCl
Near zeroThermally neutralAlmost flatNaCl (+3.9 kJ/mol)
NegativeExothermicSolubility fallsCe2(SO4)3, Ca(OH)2, most gases

NaCl is the textbook case of near-independence. Its ΔHsol of only +3.9 kJ/mol means solubility rises from about 35.7 g per 100 mL at 0°C to just 39.1 g at 100°C — which is precisely why you cannot recrystallise salt from water by cooling, but can recrystallise KNO3 beautifully.

Why Gases Behave Oppositely

Gas dissolution is essentially always exothermic, because no energy is needed to separate gas molecules from each other — only the favourable solvation step remains. Consequently gas solubility falls as temperature rises. This is why a warm fizzy drink goes flat faster, why boiling water releases bubbles before it boils, and why thermal pollution of rivers reduces dissolved oxygen available to fish.

The same van’t Hoff relationship applies throughout, with solubility taking the place of the equilibrium constant. A large ΔHsol gives a steep solubility curve; a small one gives a flat curve.

Worked Examples

Example 1: KNO3: S1=31.6g/100mL at 20°C, ΔH=+34.9kJ
S2 at 100°C (373K)
Result: S2=247g/100mL — huge increase!
Hot water dissolves much more KNO3
Example 2: NaCl: ΔH=+3.9kJ/mol — nearly flat
Small ΔH = weak T dependence
Result: S increases only slightly with T
Explains why salting out requires other methods
Example 3: Exothermic dissolution
Ca(OH)2, ΔHsol negative, heated from 20°C to 80°C
Result: Solubility decreases
Lime water becomes cloudier when warmed as dissolved Ca(OH)2 comes out of solution — the opposite of intuition and a classic demonstration.
Example 4: Why recrystallisation works
KNO3: 31.6 g/100 mL at 20°C, about 245 g/100 mL at 100°C
Result: Cooling a saturated hot solution drops most of it out
The steep curve from a large ΔHsol is exactly what makes recrystallisation an effective purification. A flat curve like NaCl gives almost no recovery.
Example 5: Dissolved oxygen and temperature
O2 in water: about 14.6 mg/L at 0°C, 7.6 mg/L at 30°C
Result: Solubility roughly halves
Exothermic gas dissolution. This is why warm-water rivers hold far less oxygen and why thermal discharge stresses aquatic life.

Common Mistakes

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Assuming all solids become more soluble when heated

Most do, but not all. Ce2(SO4)3 and Ca(OH)2 have exothermic dissolution and become less soluble as temperature rises.

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Applying solid behaviour to gases

Gas dissolution is exothermic, so gas solubility decreases with temperature — the opposite of most salts. Confusing the two leads to wrong predictions in environmental and beverage chemistry.

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Extrapolating far beyond measured data

The calculation assumes ΔHsol is constant with temperature. Over a wide range it is not, and predictions well outside the measured region become unreliable.

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Confusing solubility with dissolution rate

Temperature affects both, but they are different properties. Heating speeds dissolution even for a salt whose equilibrium solubility barely changes.

Frequently Asked Questions

Exothermic dissolution and temperature?
When ΔHsol<0 (exothermic), Le Chatelier: adding heat drives dissolution backward → solubility decreases. Examples: Ce2(SO4)3, Na2SO4 (above 32°C). Most salts are endothermic (solubility increases with T). Gases are exothermic (solubility decreases with T).
Supersaturation and crystallization?
Heat solution to dissolve maximum solute. Cool slowly → supersaturation → crystallization. This is how large pure crystals form. Rapid cooling → many small crystals. Used in sugar refining, pharmaceutical crystallization.
Why does heating usually increase solubility?
Because dissolution is endothermic for most salts. Heat acts as a reactant in the dissolution equilibrium, so adding it shifts the balance toward more dissolved solute.
Which substances become less soluble when heated?
Those with exothermic dissolution — calcium hydroxide, cerium sulfate, and essentially all gases. For these, heat is effectively a product and warming drives it out.
Why is NaCl solubility almost flat with temperature?
Its enthalpy of solution is only about +3.9 kJ/mol, so the temperature term in the van’t Hoff expression is tiny. Solubility rises from roughly 35.7 to 39.1 g/100 mL between 0 and 100°C.
Why do gases become less soluble in warm water?
Gas dissolution releases heat, since no energy is required to separate gas molecules from one another. Warming therefore shifts the equilibrium back toward the gas phase.
How does this relate to recrystallisation?
Purification by recrystallisation depends on a steep solubility curve. Compounds with large positive ΔHsol dissolve readily when hot and crystallise on cooling; flat-curve compounds recover poorly.

Formula Explorer connections

Interpretation: This formula describes how reactants, products, ions or phases distribute when opposing processes reach equilibrium. Assumption: Use equilibrium rather than initial concentrations, correct stoichiometric exponents, and the specified temperature; activities may replace concentrations in nonideal systems.

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