Chemistry, Third edition

13 · ENERGY CHANGES IN CHEMICAL REACTIONS

Lattice enthalpy as a measure of the strength of ionic

bonding

The lattice enthalpy of NaCl(s) is an index of the strength of the bonding between Na

and Clions in the sodium chloride crystal. However, ions in a sodium chloride crys-

tal lattice are not paired, and each sodium (or chloride) ion is attracted by six oppos-

itely charged neighbouring ions and (less strongly) by more distant ions. This means

thatH—L^ is a measure of the totalattractive force between sodium and chloride ions

within the lattice, and does not simply reflect the force of attraction between isolated

pairsof Na,Cl.

The ‘MX’-type lattices NaCl(s), NaI(s) and NaBr(s) possess the following lattice

enthalpies:771,684 and 731 kJ mol^1 respectively. Of these crystals, NaCl

possesses the most positive H—L^ and is said to be the ‘most stable’ of the series.

234

Lattice enthalpies

Given that

Ca^2 (g)2Cl(g) Ca^2 ,2Cl(s) H^ —L 2237 kJ mol–1

how much energy is required to break up exactly 0.1 mol of calcium chloride crystal into the

separate gaseous ions at 298 K?

Exercise 13L

Use of Hess’s law in calculating the lattice enthalpy of an

ionic crystal

The thermochemical equation showing the reaction in which sodium chloride is

formed from its elements at 298 K,

Na(s)^1 ⁄ 2 Cl 2 (g) Na,Cl(s) H—^ f  411 kJ mol–1 (13.16)

may be split up into several imaginary steps known as the Born–Haber cycle(Fig.

13.8). The enthalpy change for each step has been defined in Table 13.1. As always,

Progress of reaction

Enthalpy/kJ

Na+,Cl

• (s)

Na(s) +^1 / 2 Cl 2 (g)

atomization

Na(g) +^1 / 2 Cl 2 (g)

Na+(g) + e– +^1 / 2 Cl 2 (g)

ionization

bond dissociation

Na+(g) + Cl•(g) + e–

ΔH 2 –

ΔH 1 –

ΔH 3 – ΔH 4

ΔH 5 –

ΔHf^ –

lattice formation

electron gain

Na+(g) + Cl

• (g)

Fig. 13.8The Born–Haber cycle,

showing the enthalpy changes

involved in the formation of NaCl(s)

from its elements. The enthalpy

axis is not to scale.