Lattice enthalpy is a fundamental concept in IB Chemistry. It helps explain why certain ionic compounds are more stable than others, why some dissolve easily in water, and why melting points vary across ionic solids. Although students often find lattice enthalpy confusing, the idea becomes much clearer when you understand how lattice formation and separation relate to energy changes.
What Is Lattice Enthalpy?
Lattice enthalpy is the enthalpy change when one mole of an ionic solid is formed from its gaseous ions.
This is the definition used in the IB syllabus:
M⁺(g) + X⁻(g) → MX(s) ΔHlatt = negative
This is an exothermic process because forming an ionic lattice releases energy as oppositely charged ions attract each other strongly.
However, some textbooks use the reverse definition:
The enthalpy required to separate one mole of an ionic solid into its gaseous ions.
MX(s) → M⁺(g) + X⁻(g) ΔHlatt = positive
Both definitions describe the same magnitude of energy change.
IB uses the formation definition (exothermic), so always check for consistency in exam questions.
Why Lattice Enthalpy Is Important
Lattice enthalpy is a direct measure of:
- The strength of ionic bonding
- The stability of an ionic solid
- Trends in melting points and solubility
A large magnitude lattice enthalpy means ions are strongly held together.
Examples:
- MgO has a very high lattice enthalpy → extremely high melting point
- NaCl has a moderate lattice enthalpy → melts more easily than MgO
Understanding lattice enthalpy helps predict the properties of ionic substances.
