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When gaseous ions come together to build an ionic lattice, strong electrostatic attractions between the oppositely charged ions release a large amount of energy. This energy release is what holds the solid crystal together and is quantified by the lattice enthalpy.
OCR uses the formation convention for lattice enthalpy. The OCR definition is:
Lattice enthalpy (ΔH°_LE or ΔH°_lat) is the enthalpy change when one mole of an ionic compound is formed from its gaseous ions under standard conditions.
The key features of this definition are:
CRITICAL - OCR CONVENTION: OCR defines lattice enthalpy as an energy release (formation of lattice from gaseous ions). It is always negative. This is DIFFERENT from the AQA convention, which defines lattice enthalpy as a dissociation (gaseous ions formed from the lattice, positive value). Under OCR you should ALWAYS write lattice enthalpy as a negative number.
For sodium chloride:
Na+(g) + Cl-(g) -> NaCl(s) ΔH°_LE = -787 kJ mol^-1
For magnesium oxide:
Mg^2+(g) + O^2-(g) -> MgO(s) ΔH°_LE = -3791 kJ mol^-1
Notice the state symbols: the ions on the left are (g) to show they are isolated gaseous ions, and the product is (s) because ionic compounds are solids at standard temperature.
For calcium chloride (where stoichiometry matters):
Ca^2+(g) + 2Cl-(g) -> CaCl2(s) ΔH°_LE = -2258 kJ mol^-1
You must balance the equation so that exactly one mole of the ionic compound is formed. For CaCl2 that means 2 moles of Cl- ions are needed on the left, but the enthalpy change is per mole of CaCl2, not per mole of ions.
Under the OCR convention, gaseous ions (well separated, in the gas phase) come together to form a regular array of ions tightly bound by electrostatic attraction. As the ions approach, potential energy is released. The ionic lattice sits at a much lower energy than the separated gaseous ions, and this energy difference is released to the surroundings as heat.
The magnitude of the lattice enthalpy is a direct measure of how strongly the ions attract each other in the lattice. A more negative (more exothermic) lattice enthalpy means a stronger ionic bond.
| Compound | Lattice enthalpy / kJ mol^-1 | Comment |
|---|---|---|
| NaCl | -787 | +1, -1 ions |
| NaBr | -751 | larger anion, slightly weaker |
| MgCl2 | -2526 | doubly charged cation |
| MgO | -3791 | both ions doubly charged, small radii |
| CaO | -3401 | larger cation than Mg^2+ |
| LiF | -1031 | small ions |
Lattice enthalpies for simple 1+/1- salts are around -700 to -1000 kJ mol^-1; for 2+/2- compounds they are typically -3000 to -4000 kJ mol^-1.
Unlike enthalpy of combustion, lattice enthalpy cannot be measured in a calorimeter. You simply cannot take isolated gaseous Na+ and Cl- ions, bring them together, and measure the heat given out. The starting state (gaseous ions) is not experimentally accessible.
Instead, lattice enthalpy is found indirectly using a Born-Haber cycle - an application of Hess's law that links the lattice enthalpy to experimentally measurable quantities such as enthalpies of formation, ionisation energies, electron affinities and atomisation energies. Lessons 2-4 of this course develop the Born-Haber cycle in detail.
The lattice enthalpy governs many of the properties that make ionic compounds distinctive:
Melting and boiling points: breaking the lattice requires supplying energy equal in magnitude to the lattice enthalpy (approximately). MgO, with its very exothermic lattice enthalpy, melts at 2852 °C and is used to line furnaces. NaCl melts at 801 °C, much lower than MgO because its lattice enthalpy is much smaller.
Hardness and brittleness: strong ionic attractions give hard crystals, but the same alignment of charges means a small shift of one layer can bring like charges face-to-face, splitting the crystal along a cleavage plane.
Solubility: whether an ionic compound dissolves depends on whether the energy released in hydrating the ions can compensate for the energy required to break the lattice (see lesson 5).
Electrical conductivity: solid ionic compounds do not conduct because the ions are locked in place, but molten or dissolved ionic compounds conduct because the ions are free to move.
graph TD
A[Gaseous ions<br/>Na+ g plus Cl- g<br/>High energy] -->|Lattice enthalpy<br/>exothermic| B[Ionic solid<br/>NaCl s<br/>Low energy]
The arrow points downward in energy: releasing energy to go from the scattered gaseous ions to the ordered solid. The length of the arrow is proportional to the magnitude of the lattice enthalpy.
Question: Write an equation, with state symbols, to represent the lattice enthalpy of aluminium oxide, Al2O3.
Answer: Al2O3 contains 2 Al^3+ and 3 O^2-. To form 1 mole of Al2O3 we need 2 mol Al^3+(g) and 3 mol O^2-(g):
2Al^3+(g) + 3O^2-(g) -> Al2O3(s) ΔH°_LE
The lattice enthalpy here is very large (about -15900 kJ mol^-1) because of the high charges and small ionic radii.
Question: The lattice enthalpy of CaO is -3401 kJ mol^-1. Interpret this value.
Answer: 3401 kJ of energy is released when 1 mole of solid CaO is formed from 1 mole of gaseous Ca^2+ ions and 1 mole of gaseous O^2- ions. The negative sign shows the process is exothermic. The large magnitude reflects the strong electrostatic attraction between the doubly charged ions.
OCR lattice enthalpy is the exothermic enthalpy change accompanying the formation of one mole of an ionic compound from its gaseous ions under standard conditions. It is always negative and its magnitude is a direct measure of ionic bonding strength. Because gaseous ions are not a practical starting material, lattice enthalpy is determined indirectly through a Born-Haber cycle, which is the subject of the next lesson.