Lattice energy

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The lattice energy of an ionic solid is a measure of the strength of bonds in that ionic compound. It is given the symbol U and is equivalent to the amount of energy required to separate a solid ionic compound into gaseous ions. Lattice energy can also be considered as the energy given off when gaseous ions form an ionic solid. It is dependent on ionic charge and the ionic radius: as the charge of the ions increases the magnitude of the lattice energy increases (becomes more negative), and as the radius decreases (the ions in the ionic solid move closer together) the lattice energy increases.

Lattice energy is used primarily in respect to inorganic, crystalline solids. It cannot be determined directly but can be calculated by using experimental data in the Born Haber cycle or from theoretical calculations.

Lattice energy is inversely related to the distances of ionic bonds. Barium oxide (BaO), for instance, has a bond radius of 275 picometers and a lattice energy of -3054 kJ/mol, while sodium chloride (NaCl) has a bond radius of 283 picometers and a lattice energy of -786 kJ/mol.

[edit] Calculation

Lattice energy is the energy change when an ionic compound is separated into its constituent ions at infinite separation. It is given in units of kilojoules per mole. For example, when solid sodium chloride is separated into gaseous sodium and chlorine ions

the lattice energy is 786 kJ/mol.

Lattice energy is usually calculated by using the Born-Haber cycle. The Kapustinskii equation can also be used.

Lattice energy, the potential energy of two interacting ions (taken as two point charges), can also be calculated by using a modified version of Coulomb's law: , where , r is the distance between the centers of the ions, and q₁,q₂ are the charges on the ions.^[1]

[edit] Lattice enthalpy

The lattice enthalpy is the enthalpy change involved in formation of the ionic compound from gaseous ions. Some older textbooks define it as the energy required to break the ionic compound into gaseous ions. The former definition is invariably exothermic and the latter is endothermic. The existence of both definitions is in part due to inconsistencies in defining Standard Enthalpy Changes of Bond Dissociation, which it is often associated with.

[edit] References

1. ^ Brown, Theodore L, H Eugene Lemay, and Bruce E. Bursten. Chemistry: The Central Science. pp265. New Jersey: Prentice Hall, 2000.

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