Geoscience Reference
In-Depth Information
6.1 Solubility Equilibrium
The solubility equilibrium, subject to natural processes in the subsurface matrix,
molecular properties of the compound, the composition of the aqueous solution, and
the ambient temperature. Here, we focus our discussion on pollutant behavior.
The subsurface system contains a dynamic and complex array of inorganic and
organic constituents. The fate of pollutants in the aqueous phase therefore is
governed by a variety of reactions, including acid-base equilibria, oxidation-
reduction equilibria, complexation with organic and inorganic ligands, precipita-
tion and dissolution, ion exchange, and adsorption. The rate at which these reac-
tions occur, together with the rate of volatilization and biologically or chemically
induced degradation, controls contaminant concentrations in subsurface water.
Figure
6.1
shows the relationships among these reactions.
Solubility equilibrium is the final state to be reached by a chemical and the
subsurface aqueous phase under specific environmental conditions. Equilibrium
provides a valuable reference point for characterizing chemical reactions. Equi-
librium constants can be expressed on a concentration basis (K
C
), on an activity
basis (K
o
), or as mixed constants (K
m
) in which all parameters are given in terms of
concentration, except for H
+
,OH
-
, and e
-
(electron) which are given as activities.
Next, we provide a short description of the reactions involving transfer of
protons and electrons that affect the solubility equilibria.
Acid-base equilibria are described by a group of reactions covering the transfer
of protons, in which the proton donor is an acid and the acceptor is a base,
acid base
þ
proton, with the equilibrium constant given by
K
A
¼½H
þ
½base
=
½acid
ð
6
:
1
Þ
where K
o
is an acidity constant, and concentrations are denoted by []. From
Eq.
6.1
, it follows that, for a given ratio of activities of a particular acid and its
conjugate base, the proton activity has a fixed value. In a system with a complex of
two bases and two acids, the relation becomes acid
1
þ
base
2
base
1
þ
acid
2
and
the corresponding equilibrium becomes
K
1
;
2
¼ K
A1
=
K
A2
¼½base
1
½acid
2
=
½acid
1
½base
2
;
ð
6
:
2
Þ
o
and K
A
o
denote the acidity constants of the two acids.
Water is the ever-present proton acceptor in the subsurface. During the disso-
ciation of an acid in subsurface water, H
3
O
+
is one of the dissociation products
and the acid strength is a measurable parameter. In a dilute solution, the activity of
the hydrated protons equals that of H
3
O
+
and the pH value characterizes the H-ion
activity. Substituting for pH in Eq.
6.1
, we obtain
where K
A1
