Enthalpy of Solution Calculator
Two huge numbers nearly cancelling — which is the point.
Work out Enthalpy of Solution. Two huge numbers nearly cancelling — which is the point. Written for the problem set you are stuck on.
As the enthalpy of lattice FORMATION, so negative. NaCl is −787.
Both ions added together. For NaCl, −406 for Na⁺ plus −378 for Cl⁻.
Enthalpy of solution
+3 kJ/mol
Endothermic — the solution gets colder as it dissolves
Dissolving an ionic solid is two competing processes: pulling the lattice apart costs energy, and hydrating the freed ions releases it. The enthalpy of solution is the small difference between two very large numbers, which is why salts that look almost identical can behave oppositely. Sodium chloride comes out at about +3 kJ/mol from terms near 800 — so salt water cools imperceptibly. Ammonium nitrate is strongly endothermic and is what makes an instant cold pack work; calcium chloride is strongly exothermic and heats a hand warmer. A positive enthalpy of solution does not prevent dissolving. Entropy also drives the process, and for most salts the entropy gained by dispersing the ions is more than enough to make a mildly endothermic dissolution spontaneous.
How the Enthalpy of Solution Calculator works
Enter the lattice energy and total hydration enthalpy to get the enthalpy of solution. The two terms are large and nearly cancel, which is why chemically similar salts behave oppositely — ammonium nitrate cools an instant cold pack while calcium chloride heats a hand warmer.
Also known as: heat of solution calculator · hydration enthalpy calculator · why does ammonium nitrate get cold · endothermic dissolving calculator
Frequently asked questions
How do I calculate enthalpy of solution?
ΔH_sol = −(lattice energy) + (hydration enthalpy), taking lattice energy in the negative lattice-formation convention. It is the energy cost of breaking the lattice apart set against the energy released by hydrating the ions.
Why does ammonium nitrate get cold when it dissolves?
Because its lattice energy exceeds its hydration enthalpy, so breaking the crystal apart costs more than hydrating the ions returns. The difference is drawn from the surroundings, which is exactly how an instant cold pack works.
Why can a salt dissolve if the process is endothermic?
Because spontaneity depends on Gibbs energy, not enthalpy alone. Dispersing ions through a solvent raises entropy substantially, and for most salts that is more than enough to outweigh a mildly positive ΔH.
Why do similar salts behave so differently?
Because ΔH_sol is a small difference between two numbers near 800 kJ/mol. A few percent shift in either term flips the sign, so chemically similar compounds can warm or cool a solution with no obvious pattern.
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