Enthalpy of solution and hydrationEdexcel A-Level Chemistry: Revision notes
Section 1
Enthalpy change of solution
The enthalpy change of solution, ΔH(sol), is the enthalpy change when one mole of an ionic solid dissolves in water to form a solution at infinite dilution (so that no further enthalpy change occurs on adding more water).
Example: NaCl(s) + aq → Na⁺(aq) + Cl⁻(aq). It can be exothermic or endothermic.
Section 2
Enthalpy change of hydration
The enthalpy change of hydration, ΔH(hyd), is the enthalpy change when one mole of gaseous ions dissolves in water to form aqueous ions at infinite dilution, for example Na⁺(g) + aq → Na⁺(aq).
It is always exothermic, because the ions attract the polar water molecules (ion-dipole attractions) and energy is released.
Hydration starts from gaseous ions, solution starts from the solid. Do not mix the two definitions.
Section 3
Energy cycles and calculations
To dissolve a solid, the lattice is first broken into gaseous ions (the reverse of lattice energy, so the sign changes), and then the gaseous ions are hydrated:
ΔH(sol) = −ΔH(latt) + ΔH(hyd)(cation) + ΔH(hyd)(anion)
For a compound with two anions, count the anion hydration twice. Worked example (KCl): ΔH(sol) = +717 − 322 − 364 = +31 kJ mol⁻¹.
The same relationship can be shown on an energy level diagram: gaseous ions at the top, solid and aqueous ions below, with arrows for −ΔH(latt) up, ΔH(hyd) down and ΔH(sol) between solid and aqueous ions.
Write down the cycle first, then substitute. Check the sign of the lattice energy and count each ion in the formula.
Section 4
Effect of ionic charge and radius
Lattice energy becomes more exothermic as ionic charge increases and ionic radius decreases (stronger ion-ion attraction). MgO is much more exothermic than NaCl.
Enthalpy change of hydration becomes more exothermic with higher charge and smaller radius, as the charge density is greater and water molecules are attracted more strongly (Mg²⁺ ≫ Na⁺; F⁻ more exothermic than Cl⁻).
Whether a solid dissolves exothermically depends on the balance of the two: ΔH(sol) is small compared with the large terms that are being added and subtracted.
That's the notes covered.
Carry on to the next subtopic.
Exam questions on Enthalpy of solution and hydration
- A student investigates an instant cold pack in which potassium chloride dissolves in water and the temperature of the water falls. In energy cycles, lattice energy is taken as the enthalpy change when one mole of the solid is formed from its gaseous ions. For potassium chloride the lattice energy is −717 kJ mol⁻¹, and the enthalpy changes of hydration are −322 kJ mol⁻¹ for K⁺ and −364 kJ mol⁻¹ for Cl⁻.Calculate the enthalpy change of solution of potassium chloride and use your answer to explain why the temperature of the water falls.2 marks
- A chemist compares how strongly different ions are attracted to water molecules when they dissolve. She is looking at the enthalpy change of hydration of a series of gaseous ions, and wants to explain the differences in terms of ionic charge and radius.Explain why the enthalpy change of hydration of Mg²⁺ is much more exothermic than that of Na⁺.2 marks
- Energy data (in kJ mol⁻¹) are given for sodium chloride and magnesium chloride. Lattice energy (formation of the solid from gaseous ions): NaCl −787; MgCl₂ −2526. Enthalpy change of hydration: Na⁺ −406; Mg²⁺ −1920; Cl⁻ −364. The enthalpy change of solution of NaCl is +17 kJ mol⁻¹.Calculate the enthalpy change of solution of magnesium chloride.3 marks
Written by the Exaim team, led by Shaun Daswani (Head of Upper Secondary, Improve ME Institute; MSc Financial Mathematics, Imperial College London; BSc, UCL) and Jason Daswani (operational lead, Improve ME Institute; LSE).