In order to understand why many ionic compounds are soluble we first need to consider our model of an ionic lattice in a little more detail.
Lattice energy or lattice enthalpy is both the energy needed to from gaseous ions from one mole of a solid ionic compound / ionic lattice AND the energy released when one mole of a solid ionic compound / ionic lattice is formed from its ions in the gaseous state.
And herein lies much confusion … text books can use the term to mean both!
Clearly forming a lattice is highly exothermic as millions of strong electrostatic attractions pull the gaseous ions into a regular crystalline arrangement. Breaking up a lattice to from gaseous ions must be an highly endothermic process.
Na+(g) + Cl–(g) ⇾ NaCl(s) ΔH⦵ = -787 kJ mol-1
NaCl(s) ⇾ Na+(g) + Cl– (g) ΔH⦵ = +787 kJ mol-1
How do we tell which version of lattice energy a question or text book is talking about?
- Check the sign – if it is negative then we are forming a lattice and if it is positive we are breaking it apart
- Text books and exam questions often talk in terms of ‘enthalpy of formation of a lattice’ which equates to the exothermic process
To be clear, I make the assumption that lattice energy or enthalpy is an endothermic process to convert one mole of an ionic solid into its constituent ions which is the most common definition outside of A level.
It doesn’t actually make any difference to our explanations or our calculations as long as we all agree on a definition before beginning. Just be aware!
A note on the nature of ionic lattices
A quick note on solvents
