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ION–ION INTERACTIONS 265

Fig. 3.19. Determina-
tion of individual ionic
activity from the cell
potential.

where is the standard potential of the electrode reaction, at 25 °C and a unit activity
of the ion i in solution; the Faraday, or charge on 1 gram ion.
It follows that if one has two electrodes, each in contact with the same ions but at
different activities, and the reactions are in thermodynamic equilibrium, then neglect-
ing for a moment any potential that might exist at the contact between the two solutions,
i.e., any liquid-junction potential, from Eq. (3.103),

and a schematic representation of this idea is shown in Fig. 3.19.
Thus, one could argue as follows: if one of the solutions in the cell has a
sufficiently low concentration, e.g., for 1:1 electrolyte, then the
Debye–Hückel limiting law applies excellently. Hence, if one of the solutions has a
concentration of we know the activity of the ion i in that cell,
so that Eq. (3.104) would give at once the activity of species i in the more
concentrated cell.
Moreover—and still keeping the question of the liquidjunction rigidly suppressed
in one’s mind—the answer would have one big advantage, it would give an individual
ionic activity coefficient,
A method of such virtue must indeed have a compensating complication, and the
truth is that the neglected liquid junction potential (LJP) may not be negligible at all. 14

14 It turns out that where and are, respectively, the cationic and anionic transport
numbers. There are cases (e.g., for junctions of solutions of KC1) where and are almost the same and
hence the and the correction due to the liquid junction is negligible. In some cases the difference
may be quite considerable.