1.8.4Physical Chemistry

Enthalpy of Hydration

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Section 1 of 6Definitions

Enthalpy of solution

The enthalpy change when one mole of an ionic solid dissolves in enough water to form an effectively infinitely dilute solution.

Enthalpy of hydration

The enthalpy change when one mole of a gaseous ion dissolves in enough water to form an infinitely dilute solution.

Example

Mg²⁺(g) + aq → Mg²⁺(aq) is the hydration of magnesium ions, and this process is always exothermic.

Meaning of aq

The symbol aq represents excess water, so the ions are surrounded by water and no longer interact with each other.

Section 2 of 6Why Hydration Is Exothermic

Why Hydration Is Exothermic

Polar water

Water has a partial negative oxygen and partial positive hydrogens, so it can attract both cations and anions.

Cations

The oxygen lone pairs of water point towards positive ions and form ion–dipole attractions with them.

Anions

The hydrogen atoms of water point towards negative ions, forming ion–dipole attractions that release energy.

Energy

Forming these attractions releases energy, which is why all hydration enthalpies are negative.

Section 3 of 6Factors and the Energy Cycle

Charge density

Ions with a higher charge and smaller radius attract water more strongly, giving a more exothermic hydration enthalpy.

Group 1 and 2

Mg²⁺ is more strongly hydrated than Na⁺ because it is smaller and carries twice the charge.

Cycle

Solid → gaseous ions (lattice dissociation, endothermic) → aqueous ions (hydration, exothermic) gives the enthalpy of solution.

Equation

ΔHsol = ΔHlatt(dissociation) + ΣΔHhyd, or equivalently −ΔHlatt(formation) + ΣΔHhyd.

Quantity	Sign	Process
Lattice dissociation	Positive	Solid to gaseous ions
Hydration of an ion	Negative	Gaseous ion to aqueous ion
Solution	Either	Solid to aqueous ions

Section 4 of 6Interpreting Results

Interpreting Results

Exothermic solution

The hydration energy released outweighs the lattice energy needed, so the solution warms.

Endothermic solution

The lattice energy is greater than the hydration energy released, so the solution cools.

Solubility

Solubility also depends on entropy change, so an endothermic solution can still occur if disorder increases enough.

Section 5 of 6Worked Example

Question

Calculate the enthalpy of solution of NaCl using: lattice dissociation = +787;
- hydration of Na⁺ = −406;
- hydration of Cl⁻ = −364 (all kJ mol⁻¹).

EquationΔHsol = ΔHlatt(dissociation) + ΔHhyd(Na⁺) + ΔHhyd(Cl⁻).
Substitute+787 + (−406) + (−364).
Add the hydration terms−770.
787 − 770 = +17.

Answer ΔHsol = +17 kJ mol⁻¹ (slightly endothermic)

Section 6 of 6Exam Checklist & Tips

1.8.4 – Key Points for the Exam

Enthalpy of solution Enthalpy change when one mole of solute dissolves in water to give an infinitely dilute solution.
Enthalpy of hydration Enthalpy change when one mole of gaseous ions dissolves in water to give an infinitely dilute solution.
Ion–dipole attraction Electrostatic attraction between an ion and the partial charge on a polar molecule.
Charge density Ionic charge divided by ionic radius, a measure of how concentrated the charge is.

Common Mistakes

Writing hydration for the solid instead of for gaseous ions.
Using the lattice formation value without changing its sign in the solution equation.
Saying hydration enthalpies can be positive; they are always negative.
Assuming a solid dissolves only if ΔHsol is negative; entropy matters too.