Chapter 11
Acid-Base Equilibria in Saline Media:
Application of the Mean Spherical Approximation
M. E. Sastre de Vicente . T. Vilarifio
11.1
Introduction
The simplest mathematical formulation of the hypothesis that chemical equilibrium
exists in the bulk of a saline solution is based on the equilibrium constant, KT = K* Q{ }';),
where KT is the thermodynamic equilibrium constant, K* the stoichiometric constant
and Q{}';) the ratio of activity coefficients (}';) associated with the equilibrium. Studies
on the effect of ionic strength on stoichiometric constants are based on modelling of
the Q{}';) term (Sastre de Vicente 1997; Daniele et al. 1997).
With regard to this equation, the following points can be mentioned:
• Most treatments rely on the Debye-Hiickellimiting law, although their final purpose
is their possible application to moderate and high ionic strengths, as shown through
many studies devoted to the analysis of changes in activity and osmotic coefficients
for strong electrolytes.
• The vast majority of the treatments lead to explicit expressions for pK* as a function
of I, which is clearly an advantage; however, interpreting the parameters associated
to the fitting of a pK* vs. I curve is often far from easy.
• This kind of study is carried out in a constant ionic medium, which has been traditionally employed as a synonym for "background electrolyte;" the principal weakness of the method arises from the assumption that the added salt is inert to the
equilibria involved. If the background electrolyte cannot be considered as inert, ion
association must be included and the equilibria redefined in the system.
11.2
Acid-Base Equilibria in Saline Media
For the case of acid-base equilibrium reactions in saline media, we should consider
the most significant acid-base reactions that take place in solution, which can be classified as follows (Edsall 1943; Izatt et al. 1995):
• Type I:
AH=A- +H+
(n.l)
where the charge separation involved usually has a strong effect on H20 structure.
Acid-Base Equilibria in Saline Media:
Application of the Mean Spherical Approximation
M. E. Sastre de Vicente . T. Vilarifio
11.1
Introduction
The simplest mathematical formulation of the hypothesis that chemical equilibrium
exists in the bulk of a saline solution is based on the equilibrium constant, KT = K* Q{ }';),
where KT is the thermodynamic equilibrium constant, K* the stoichiometric constant
and Q{}';) the ratio of activity coefficients (}';) associated with the equilibrium. Studies
on the effect of ionic strength on stoichiometric constants are based on modelling of
the Q{}';) term (Sastre de Vicente 1997; Daniele et al. 1997).
With regard to this equation, the following points can be mentioned:
• Most treatments rely on the Debye-Hiickellimiting law, although their final purpose
is their possible application to moderate and high ionic strengths, as shown through
many studies devoted to the analysis of changes in activity and osmotic coefficients
for strong electrolytes.
• The vast majority of the treatments lead to explicit expressions for pK* as a function
of I, which is clearly an advantage; however, interpreting the parameters associated
to the fitting of a pK* vs. I curve is often far from easy.
• This kind of study is carried out in a constant ionic medium, which has been traditionally employed as a synonym for "background electrolyte;" the principal weakness of the method arises from the assumption that the added salt is inert to the
equilibria involved. If the background electrolyte cannot be considered as inert, ion
association must be included and the equilibria redefined in the system.
11.2
Acid-Base Equilibria in Saline Media
For the case of acid-base equilibrium reactions in saline media, we should consider
the most significant acid-base reactions that take place in solution, which can be classified as follows (Edsall 1943; Izatt et al. 1995):
• Type I:
AH=A- +H+
(n.l)
where the charge separation involved usually has a strong effect on H20 structure.