Modern inorganic chemistry

ENERGETICS 79

the enthalpy changes involved in dissolving them in water are as

follows:

NaF NaCl NaBr Nal

A/i 5 MX(s)-»M+(g) +X~(g) +919 +787 +752 +703

SAfchvd:M + (g) + X-(g) -» M + (aq)

4- X~(aq) -921 -787 -753 -711

A/zs MX(s) -> M+(aq) 4- X"(aq) -2 0 -1 -8

AgF AgCl AgBr Agl

Afcs MX(s) -> M+(g) + X~(g) +966 +917 +905 +891

£Ahyd: M+(g) + X-(g)->M+(aq)

+ X^(aq) -986 -851 -820 -778

A/is MX(s)-> M+(aq) 4- X~(aq) -20 +66 +85 +113

Although the data for the silver halides suggest that silver(I) fluoride
is likely to be more soluble than the other silver halides (which is in
fact the case), the hydration enthalpies for the sodium halides almost
exactly balance the lattice energies. What then is the driving force
which makes these salts soluble, and which indeed must be res-
ponsible for the solution process where this is endothermic? We
have seen on p. 66 the relationship AG^ = A//*^9 " — TAS^ and
noted that for a reaction to be spontaneous AG^ must be negative.
The driving force, then, is to be found in the entropy term TAS^.
When a crystal dissolves the orderly arrangement of ions in the
lattice is destroyed, but since each ion becomes solvated order is
brought into the areas of solvent around each ion. Generally, how-
ever, despite this 'ordering of the solvent' there is an overall increase
in entropy and AS^ is positive. Hence, negative values of AG^ can
be produced even for endothermic reactions, and since TAS"^0 "
increases with temperature, it is not surprising to find that the
solubility of nearly all simple ionic substances increases as the
temperature is increased.
Prediction of solubility for simple ionic compounds is difficult
since we need to know not only values of hydration and lattice
enthalpies but also entropy changes on solution before any informed
prediction can be given. Even then kinetic factors must be considered.
This problem does not become easier when considering ionic
compounds of Group II elements since with the increase in ionic
charge and decrease in ionic radius of the Group II ions not only
does hydration energy increase but also the lattice energy of the
compound itself, and again the value of the enthalpy of solution is
the difference between two large (indeed, in the case of Group II,
very large) quantities.