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Lecture

ESS102H1 Lecture Notes - Aquifer, Hydronium, Bromine


Department
Earth Sciences
Course Code
ESS102H1
Professor
John Ferris

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5 Redox Processes
5.0 Introduction
The importance of redox (i.e., oxidation-reduction) processes in biogeochemistry
cannot be over emphasized. These are reactions that involve the transfer of electrons
between one or more chemical species. Life would not exist without oxidation-reduction
reactions, and the intense biological and geochemical cycling of redox sensitive elements
that essentially define the uniqueness of the Earth system would not occur. Everything
we appreciate about our planet depends on oxidation-reduction reactions.
Many elements occur in nature in more than one oxidation state. The major
elements with such redox behavior include H, O,C,S,N, and Fe. In some systems, Mn is
also a major redox element.
As a general rule, most reactions that involve electrons also involve protons.
Oxidation usually releases protons or acidity. This is a basic cause of acid mine drainage.
Conversely, reduction generally consumes protons, and the pH rises.
5.1 Oxidation-Reduction Reactions
In redox reactions, electrons are transferred from an electron-donor or reductant
to an electron acceptor or oxidant. The surrender of electrons from a compound is
referred to as oxidation, and the acceptance of electrons by a different compound is
reduction. In a complete redox reaction, the reductant is transformed into a conjugate
oxidized species and the oxidant is transformed into a conjugate reduced species.
Consider a general oxidation-reduction reaction in which reduced species red1 is
oxidized when it donates one electron per molecule to oxidized species ox2. The overall
reaction can be taken to be the result of two half-cell reactions as follows
Oxidation:
+= eoxred 11 conjugate
(1)
Reduction:
22 conjugate redeox =+
(2)
Overall redox reaction:
2121 conjugate conjugate redoxoxred +=+
(3)
The strength of an oxidant or reductant is determined electrochemically by their
relative capacities to accept or donate electrons, respectively from a reductant. As such, a
strong oxidant is a potent electron acceptor, whereas a weak oxidant has a poor capacity
to accept electrons. Similarly, a strong reductant donates electrons readily, whereas
electrons are not given up easily by a weak reductant. These sweeping generalizations
emphasize the importance of adopting a reference scale to measure the actual strength of
1

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oxidants and reductants. The conventional approach is to use the H+/H2 half-cell reaction
(Eqn. 4) in which the corresponding equilibrium constant KH is assigned a standard value
of unity.
2H+ + 2e- = H2 (4)
1
}}{{
}{
2
2== + eH
H
KH
(5)
In accordance with IUPAC1 guidelines, all half cell reactions are written as
reduction reactions with oxidized species An+ and n electrons e-on the left and reduced
species B on the right
An+ + ne- = B(6)
The corresponding mass action relationship for the half-cell reaction in which oxidized
species An+ is reduced to species B is
nn
BA
eA
B
K}}{{
}{
/
+
=
(7)
where KA/B is the equilibrium constant for the reduction reaction. Combining the H+/H2
half-cell reaction with Eqn. 6 gives the overall reaction
++ +=+ H
n
BH
n
An
22 2
(8)
with a mass action expression of
BA
H
BA
n
n
n
RK
K
K
pHA
HB
K/
/
2/
2
2/
}{
}}{{ === +
+
(9)
This relationship shows that the standard H+/H2 half-cell reaction makes the equilibrium
constant for the overall oxidation-reduction reaction KR equivalent to the equilibrium
constant of the An+/B half-cell KA/B.
As the mass action expression for the An+/B half-cell (Eqn. 7) indicates that the
greater the capacity of oxidant An+ to serve as an electron acceptor, the greater the value
of KA/B. Taking common logarithms of Eqn. 7 and rearranging gives
==+
}{
}{
loglog
1
}log{ /
n
BA
A
B
K
n
pee
(10)
1 International Union of Pure and Applied Chemistry
2

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where, as the negative common logarithm of electron activity, pe is by analogy similar to
the definition of pH = -log{H+}. When {B} = {An+}
1
}{
}{ =
+n
A
B
(11)
such that
(12)
This gives from Eqns. 3 and 5
}{
}{
log
1
0
+
= n
A
B
n
pepe
(13)
Considering that
RT
G
KK BAR
0
/
log3.2ln
==
(14)
then from Eqn. 5
nRT
G
pe 303.2
0
0
=
(15)
and
nRT
G
pe 303.2
=
(16)
Combining Eqns. 13, 15 and 16 gives the change in free energy for the half-cell reaction
illustrated by Eqn. 1.
}{
}{
log303.2
0
+
+=n
A
B
RTGG
(17)
In most aqueous systems, electron transfer in a complete oxidation-reduction
reaction takes place on a molecular level, and there is no practical way of measuring the
actual magnitude of electron activity; however, oxidation and reduction steps can be
carried out at separate electrodes if the electrons yielded by the oxidation half-cell are
conveyed to the electron acceptor at another electrode through an electrical circuit. The
3
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