Enthalpy of hydration and solution

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  1. Enthalpy terms in the dissolving process
  2. Enthalpy change of hydration.
  3. Enthalpy change of solution.
  4. Using solution enthalpies to draw Born-Haber cycles.
  5. Factors affecting hydration enthalpies.
Topic Notes

In this lesson, we will learn:

  • To recall the definitions of enthalpy of hydration and enthalpy of solution.
  • To understand the dissolving process as a combination of several energetic steps.
  • To represent these energetic steps in a Born-Haber cycle.

  • We saw in Enthalpy: Lattice energy, atomisation and electron affinity that lattice enthalpy measures the strength of the ionic bonds in an ionic compound. Recall that many ionic compounds are soluble in water. Dissolving involves more than one step in terms of bonding interactions. Dissolving therefore has two different enthalpy terms too:
    • The ionic solid is ‘pulled apart’ by the solvent water molecules. This breaking of the lattice costs energy – the lattice energy, H \triangle H lat, which we saw in the previous lesson.
      Being pulled apart, we now have dispersed gaseous ions. In terms of a normal chemical reaction, this is the bond breaking phase.
    • These dispersed ions are surrounded by water molecules which are favourably interacting with them. These new ion-water bonds release energy – the hydration enthalpy, H \triangle H hyd. In terms of a normal chemical reaction, this is the bond-making phase that gives us the product; an aqueous solution.

    Combined, these steps form the dissolution process. In the same way, these two energetic steps combined will form the enthalpy change of solution, H \triangle H sol. The definitions are below:
    • The enthalpy change of hydration, H \triangle H hyd, is the enthalpy change when one mole of gaseous ions is dissolved in water, resulting in an infinitely dilute solution.
      • This is only the bond-forming step after the lattice is broken up (where we spent the lattice energy).
      • Because ionic compounds contain at least two oppositely charged ions, an ionic lattice (for example NaCl) that has been broken up will have two different hydration enthalpies; one for the cation (Na+) and one for the anion (Cl-). They are both attractive forces so the enthalpies are negative.
    • The enthalpy change of solution, H \triangle H sol, is the enthalpy change when 1 mole of an ionic solid is dissolved in water, resulting in an infinitely dilute solution.
      • This total enthalpy change includes the lattice energy combined with the enthalpy change of hydration.

  • From the description and definitions above, you should be able to see that lattice energy, solution enthalpy and hydration enthalpy can be written in a Born-Haber cycle by applying Hess’s law.

  • The diagram above shows how using data1 on lattice energy and hydration enthalpy, the enthalpy change of solution can be worked out.

  • Just like lattice energy, there are factors that affect the hydration enthalpy. It is just a measure of the attractive forces between water molecules and whatever ions the solid ionic compound is made of. Just like with lattice enthalpy, ionic radius and charge affect hydration enthalpy for both cations and anions:
    • Hydration enthalpy becomes smaller (ion-water bonding interactions get weaker) going down the group, where ions are larger.
    • Larger charge increases the hydration enthalpy (ion-water interactions get stronger).

    1 Source for lattice enthalpies and hydration enthalpies:ATKINS, P. W., & DE PAULA, J. (2006).?Atkins' Physical chemistry. Oxford, Oxford University Press.