Chemistry Reference
In-Depth Information
the relative oxidation and reduction capabilities of
the particular species. Redox systems with positive
potentials will oxidise hydrogen to protons and thus
increasing positive potentials correspond to increas-
ing oxidation conditions. The converse is true for
increasing negative potentials: reductants (electron
donors) will reduce protons to hydrogen. For ex-
ample, increasing positive potentials equate to oxi-
dation of metals to solution species or to passivation,
e.g. oxide films, whereas increasing negative poten-
tials equate to stable metal species. Thus, generally
for two redox couples (denoted as 1 and 2) with
values of standard potentials
E
1 and
E
2 the reduced
form of couple 1 can be oxidised by the oxidised
form of couple 2 when
E
2 >
E
1. Thus the thermo-
dynamic driving force is positive and the reaction
is spontaneous with species O
2
as the oxidant. In
aqueous solution the pH can be a significant factor
and reaction equilibria commonly are expressed
in terms of potential-pH (Pourbaix) diagrams [3]
that incorporate both chemical and electrochemical
(redox) reactions. A typical Pourbaix diagram is
shown in Fig. 19.1 for copper. The broken lines are
used to signify the equilibrium between solution
species and the solid lines define equilibria between
solid phases or between solid and solution phases.
Figure 19.1 shows only the predominant phases
under specified conditions of temperature, pressure
and solution species activity and also the region of
stability of water. Equations 19.1-19.4 are the basis
of two important technological applications of elec-
trochemistry, i.e. the production of hydrogen by
electrolysis and the generation of electrical energy
from fuel cells:
Fig. 19.1
Potential pH diagram for copper.
3.1 Electrode potential, kinetics and
mass transport
Electrochemical processes are driven by the applica-
tion of a potential field, the magnitude of which gen-
erally will determine the rate of the relevent process:
charge transfer and ionic flux. Electrochemical reac-
tions are surface processes that are instigated by a
suitable charge transfer at a fluid/solid interface.
When two electrodes are placed in an ionic con-
ducting solution and are connected externally they
become charged. Thus, locally, at the solution/elec-
trode interface there is a large potential difference
over a molecular scale of a few nanometres. A simple
model of this situation [4] consists of a double layer
comprising a plane of closest approach (ihp) and a
diffuse layer or outer layer. The equilibrium estab-
lished at the interface is electrostatic and somewhat
analogous to that in a capacitor. The electrostatic
interactions determine the distribution of the poten-
O
2
+ 4H
+
+ 4e
-
Æ 2H
2
O
(19.1)
E
=
1 2291
.
-
0 0591
.
pH + 0.0148 log
(
p
)
(19.2)
OHO
O
22
2
2H
+
+ 2e
-
= H
2
(19.3)
E
=
0 0591
.
pH
-
0
.
0296
log
(
p
)
(19.4)
HH
+
H
2
2
In the case of the copper system, with
E
o
= 0.34 V,
corrosion of the copper can occur only by the re-
duction of dissolved oxygen and not by proton
reduction.
Although Pourbaix diagrams give useful informa-
tion on the behaviour of many electrochemical
systems from purely thermodynamic considerations,
in practice electrolytic systems are driven by an
overpotential.
