Lattice enthalpy is a fundamental concept in IB Chemistry Topic 5 (Energetics) and Topic 4 (Bonding). 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.
Factors Affecting Lattice Enthalpy
Lattice enthalpy depends on two main factors: charge and ionic radius.
1. Charge of the Ions
Greater charge → stronger electrostatic attraction → higher lattice enthalpy.
For example:
- Mg²⁺ and O²⁻ attract more strongly than Na⁺ and Cl⁻
- Therefore:
ΔHlatt( MgO ) > ΔHlatt( NaCl )
Charge has a very large effect because ionic attraction is proportional to the product of the charges.
2. Ionic Radius
Smaller ions → shorter distance between ions → stronger attraction → higher lattice enthalpy.
For example:
- LiF has a higher lattice enthalpy than KF because Li⁺ is smaller than K⁺
- Similarly, F⁻ forms stronger lattices than I⁻ because it is smaller
This explains many periodic trends in bonding strength.
Lattice Enthalpy and the Born–Haber Cycle
Lattice enthalpy cannot be measured directly.
Instead, it is calculated using a Born–Haber cycle, which applies Hess’s Law.
A Born–Haber cycle includes:
- Enthalpy of atomization
- Ionization energy
- Electron affinity
- Enthalpy of formation
- Lattice enthalpy
By constructing a cycle, you can solve for the unknown value, usually the lattice enthalpy itself.
Born–Haber cycles are common in IB Paper 2 problems.
Lattice Enthalpy and Solubility
Whether an ionic compound dissolves depends on two competing energies:
1. Lattice enthalpy (opposes dissolving)
Energy is needed to separate ions in the lattice.
2. Enthalpy of hydration (supports dissolving)
Energy is released when ions interact with water.
A salt dissolves when:
|ΔH hydration| > |ΔH lattice|
For example:
- NaCl dissolves easily
- MgO does not dissolve in water because its lattice enthalpy is extremely high
This relationship is central in solubility predictions.
Lattice Enthalpy and Melting Points
Higher lattice enthalpy usually means:
- Higher melting point
- Greater hardness
- Stronger ionic attraction
Ionic compounds with large charge and small radii (e.g., Al₂O₃, MgO) have very high melting points because their ionic lattices are exceptionally strong.
FAQs
Why is lattice enthalpy exothermic in IB definitions?
IB defines lattice enthalpy as the formation of an ionic lattice from gaseous ions, a process that releases energy.
Is lattice enthalpy the same as lattice energy?
Yes, the terms are often used interchangeably, but definitions may differ between formation and dissociation. Always follow the IB version.
Why can’t lattice enthalpy be measured directly?
Gaseous ions are extremely difficult to isolate experimentally, making direct measurement impractical. Hess’s Law provides an indirect method.
Conclusion
Lattice enthalpy measures the strength of attraction between ions in an ionic solid. It depends on ionic charge and ionic size, and it influences melting point, solubility, and stability. Using Born–Haber cycles, IB Chemistry students can calculate lattice enthalpy and understand how energy changes determine the properties of ionic compounds.
