Lattice Energy: Strengths of Ionic Bonds

Although there are several important factors which must be considered in determining the energy associated with forming ionic bonds, the lattice energy is where we will focus. The lattice energy is defined as the energy required in the following reaction;

M_mX_n(s) ---> Mⁿ⁺_(g) + X^m-_(g)

Lattice Energies of Alkali Metal Halides (kJ mol^-1)

Lattice Energies of Salts of F^- (kJ mol^-1)

The lattice energy is the energy liberated when oppositely charged ions in the gas phase come together to form a solid. So for sodium chloride the lattice energy is 787 kJ mol^–1. This is the energy liberated when Na⁺ and Cl^– ions in the gas phase come together to form the lattice of alternating Na⁺ and Cl^– ions in a NaCl crystal.

The charges and sizes of the ions are the critical factors which determine the magnitude of the lattice energy. The lattice energy is related to the potential energy of two interacting charges which is given in the mathematical equation;

Q₁ and Q₂ are the magnitudes of the charge on the particles in coulombs and d is the distance between centers in meters. The constant k has a value of 8.99 x 10⁹J m C^–2. From this relationship the magnitude of the lattice energy is directly related to the charge on the ions and inversely related to the ionic radii of the ions. The combination of a small anion and a small cation liberates more energy compared to the combination of large ions. In general holding one of the ions constant the trend in lattice energy is reflected in the change in size of the counterion. The strength of the ionic bond also depends on the magnitude of the charge on the cation or anion. In the case of the fluoride compounds of sodium, magnesium and aluminum there is a dramatic difference in the lattice energy with cation charge.

It is the formation of the ionic bond which is the driving force for this reaction. In the solid sodium chloride considerable energy is liberated as a result of the arrangement of the sodium and chloride ions.

To generalize this electron lose and electron gain behavior we should note that Group IA metals, with their single valence electron will easily lose the electron to form a +1 cation. A behavior that continues down the Group. Group IIA metals will lose their two valence electrons to form M²⁺ cations. When this occurs the cations are isoelectronic with the preceeding noble gas. Group IIIA metals form M³⁺ species, although as we go down the Group we do see M⁺ species appearing. Removing 3 electrons requires considerably more energy, and as a result we see fewer ionic compounds in Group IIIA compared to Group IA and Group IIA.

Group IVA and VA exhibit little tendency to lose or gain electrons to form ionic bonds. Members of these two groups prefer to share electrons.

In Groups VIA and VIIA, elements with high electron affinities, we see gain of electrons to become isoelectronic with the next noble gas. Group VIA gaining two electrons and Group VIIA gaining one electron.