Geoscience Reference
In-Depth Information
7.2 Dominance diagrams
The purpose of this section is to introduce some graphic methods used to assess which
species dominate solutions in conditions of given acidity (pH) and oxidation state (pe). Let
us consider the following acid-base equilibrium:
HSO
4
SO
2
4
H
+
⇔
+
(log
K
=−
2.0)
(7.12)
We follow the conventional presentation of the reaction as a dissociation. We can write the
constant of the equilibrium as:
SO
2
4
H
+
HSO
4
=
K
(7.13)
Taking the logarithm, we obtain:
SO
2
4
pH
=−
log
K
+
log
HSO
4
(7.14)
log
K
,wehave
SO
2
4
>
HSO
4
, while the opposite is true
Clearly, for pH
>
p
K
=−
<
for pH
p
K
.
The same principle holds for redox reactions, for example:
Fe
3
+
+
e
−
⇔
Fe
2
+
=+
or
E
H
=
(log
K
13.0
0.77 V)
(7.15)
(note that this time we consider that the electron combines with the oxidant). Writing the
equilibrium constant and taking the logarithm, we obtain:
log
Fe
3
+
Fe
2
+
pe
=
log
K
+
(7.16)
(note the ratio in the order of oxidant/reductant on the right-hand side). We note pe
0
=log
K
0.77 V), we conclude that Fe
3
+
>
Fe
2
+
, and Fe
3
+
<
=
13. When pe
>
pe
0
(or
E
H
>
Fe
2
+
otherwise.
Figure 7.1
shows these two examples of dominance diagrams. When the two types of
dominance diagrams are combined, the useful pe-pH diagram (or, equivalently, the
E
H
-
pH diagram) is obtained. This is normally complicated by the fact that oxo-anions, such as
SO
2
4
or NO
3
, require a proton to achieve reduction, which makes redox reactions and the
redox dominance diagrams pH-dependent as in:
SO
2
4
9H
+
+
8e
−
⇔
HS
−
+
+
4H
2
O g
K
=
34)
(7.17)
The geochemically important example of the sulfur system, with H
2
S, HS
−
,S
2
(solid
sulfur), SO
2
4
, and HSO
4
