Lattice enthalpy

CIE A-Level Chemistry· Unit 17: Further chemical energetics· 40 min read

1. Definition of Lattice Enthalpy★☆☆☆☆⏱ 10 min

📘 Definition

Lattice Enthalpy

\Delta H_{latt}

There are two standard definitions: lattice enthalpy of formation describes the formation of 1 mole of solid ionic compound from gaseous ions under standard conditions (298 K, 1 atm). Lattice enthalpy of dissociation is the reverse process, breaking 1 mole of solid ionic compound into gaseous ions.

Example:

For NaCl: $Δ H_{l a tt, formation} = - 787$ kJ mol⁻¹, $Δ H_{l a tt, dissociation} = + 787$ kJ mol⁻¹

📐 Worked Example

State whether the process $M g^{2 +} (g) + O^{2 -} (g) \to M g O (s)$ corresponds to lattice enthalpy of formation or dissociation, and predict the sign of $Δ H$ .

1
Step 1: Check the state and direction of the process. The reactants are isolated gaseous ions, and the product is 1 mole of solid ionic compound.
2
This matches the definition of lattice enthalpy of formation.
3
Step 2: Forming an ionic lattice requires forming new ionic bonds, which releases energy, so the process is exothermic.
4
Conclusion: This is lattice enthalpy of formation, with a negative $Δ H$ .

2. Factors Affecting Lattice Enthalpy★★☆☆☆⏱ 15 min

The magnitude of lattice enthalpy (how negative it is for formation, how positive for dissociation) depends on the strength of electrostatic attraction between oppositely charged ions in the lattice, which is controlled by two key factors:

Ionic charge: Higher ion charge increases electrostatic attraction between ions, leading to a larger magnitude lattice enthalpy.
Ionic radius: Smaller ionic radius allows ions to sit closer together, increasing attraction, leading to a larger magnitude lattice enthalpy.

📐 Worked Example

Explain why the magnitude of lattice enthalpy of MgO (-3795 kJ mol⁻¹) is much larger than that of NaF (-894 kJ mol⁻¹).

1
Step 1: Compare ionic charges first. MgO contains $M g^{2 +}$ (+2) and $O^{2 -}$ (-2), while NaF contains $N a^{+}$ (+1) and $F^{-}$ (-1). Higher charge creates much stronger electrostatic attraction.
2
Step 2: Compare ionic radii. $M g^{2 +}$ (72 pm) is smaller than $N a^{+}$ (102 pm), and $O^{2 -}$ (140 pm) is similar in size to $F^{-}$ (133 pm), so the distance between ions in MgO is smaller.
3
Step 3: Stronger attraction from higher charge and smaller ion size means more energy is released when the lattice forms, so magnitude is much larger for MgO.

✓ Quick check

Test your understanding:

Which of the following has the largest magnitude lattice enthalpy of formation?
- KF
- KCl
- KBr
- KI
Reveal answer
KF —
Correct! All have the same +1 potassium cation. $F^{-}$ is the smallest halide ion, so strongest attraction, largest magnitude lattice enthalpy.

3. Lattice Enthalpy and Ionic Compound Properties★★★☆☆⏱ 15 min

Lattice enthalpy magnitude is a direct measure of ionic bond strength in a solid ionic compound. Higher magnitude lattice enthalpy means stronger ionic bonds, which directly impacts physical properties like melting point and hardness. More energy is required to break a stronger ionic lattice, so melting point increases with lattice enthalpy magnitude.

📐 Worked Example

Predict which compound has a higher melting point: CaS or CaO. Explain your answer in terms of lattice enthalpy.

1
Step 1: Compare ionic charges. Both compounds have +2 calcium cations and -2 anions, so ion charges are equal.
2
Step 2: Compare anionic radii. $O^{2 -}$ (140 pm) is much smaller than $S^{2 -}$ (184 pm), so the average distance between ions in CaO is smaller than in CaS.
3
Step 3: Smaller ion distance increases electrostatic attraction, so CaO has a larger magnitude lattice enthalpy than CaS.
4
Step 4: More energy is required to break the stronger lattice of CaO, so CaO has a higher melting point.

4. Common Pitfalls

Wrong move:

Defining lattice enthalpy starting from solid elements rather than gaseous ions

Why:

This describes enthalpy of formation of the compound, not lattice enthalpy, and will lose all marks in exams

Correct move:

Always explicitly state that the ions involved are gaseous in your definition

Wrong move:

Mixing up the sign of lattice enthalpy for formation vs dissociation

Why:

This leads to incorrect values in Born-Haber cycle calculations, costing multiple marks

Correct move:

Formation = solid from gaseous ions, bond formation is exothermic → ΔH negative; dissociation = gaseous ions from solid, bond breaking is endothermic → ΔH positive

Wrong move:

Claiming larger ionic radius gives larger magnitude lattice enthalpy

Why:

Larger ions are further apart, so electrostatic attraction is weaker, not stronger

Correct move:

Smaller ionic radius leads to stronger attraction and larger magnitude lattice enthalpy

Wrong move:

Comparing ionic radius before ionic charge when ranking lattice enthalpy

Why:

Ionic charge has a much larger effect on lattice enthalpy than radius, leading to incorrect rankings

Correct move:

Always compare ion charges first, then compare ionic radii only when charges are equal

5. Quick Reference Cheatsheet

Concept	Key Detail	ΔHlatt (formation) magnitude
Lattice formation	1 mole solid from gaseous ions	Always negative
Lattice dissociation	1 mole solid to gaseous ions	Always positive
Higher ion charge	Stronger electrostatic attraction	Larger (more negative)
Smaller ion radius	Ions closer, stronger attraction	Larger (more negative)
Charge vs radius	Charge effect dominates over radius	Large change for +2/+2 vs +1/-1

When this came up on past exams

AI-estimated based on syllabus patterns — cross-check with official past papers for accuracy. Use only as revision-focus signals.

2022 · 2
Compare lattice enthalpy values
2023 · 1
Definition multiple choice
2021 · 4
Factors affecting lattice enthalpy

Going deeper

internalBorn-Haber CyclesNext sub-topic in Unit 17

What's Next

Understanding lattice enthalpy is the core foundation for all other topics in the further chemical energetics unit. Lattice enthalpy values are used extensively in Born-Haber cycle calculations to find unknown enthalpy changes like electron affinity, and to compare theoretical and experimental values to identify the degree of covalent character in nominally ionic compounds. The same principles of electrostatic attraction that govern lattice enthalpy also apply to enthalpy of hydration and solubility trends for ionic compounds, so the skills you learn here transfer directly to those topics. Mastery of the definitions and factors affecting lattice enthalpy is essential for exam success, as this sub-topic is almost always tested in both multiple choice and extended response questions.