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5 Redox Processes.doc
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University of Toronto St. George
Earth Sciences
ESS102H1
John Ferris
Winter
Description
5 Redox Processes
5.0 Introduction
The importance of redox (i.e., oxidationreduction) 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 oxidationreduction
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 oxidationreduction 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 OxidationReduction Reactions
In redox reactions, electrons are transferred from an electrondonor 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 oxidationreduction reaction in which reduced species red is 1
oxidized when it donates one electron per molecule to oxidized species ox . The 2verall
reaction can be taken to be the result of two halfcell reactions as follows
Oxidation: red1=conjugateox + e1 − (1)
Reduction: ox2+ e = conjugatered 2 (2)
Overall redox reaction: red1+ ox 2conjugateox +conj1gatered 2 (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 oxidants and reductants. The conventional approach is to use the H /H ha2fcell reaction
H
(Eqn. 4) in which the corresponding equilibrium constant K is assigned a standard value
of unity.
+ 
2H + 2e = H 2 (4)
K H = {H }2 =1
{H }{e }− 2 (5)
In accordance with IUPAC guidelines, all half cell reactions are written as
n+ 
reduction reactions with oxidized species A and n electrons eon the left and reduced
species B on the right
n+ 
A + ne = B (6)
The corresponding mass action relationship for the halfcell reaction in which oxidized
species A is reduced to species B is
A/ B {B}
K = n+ − n (7)
{A }{e }
A/B +
where K is the equilibrium constant for the reduction reaction. Combining the H /H 2
halfcell reaction with Eqn. 6 gives the overall reaction
An+ + n H = B + n H +
2 2 2 (8)
with a mass action expression of
R {B}{H }+ n/ 2 K A/ B A/ B
K = n+ n/2 = H = K (9)
{A }pH 2 K
This relationship shows that the standard H /H h2lfcell reaction makes the equilibrium
constant for the overall oxidationreduction reaction K equivalent to the equilibrium
n+ A/B
constant of the A /B halfcell K .
n+
As the mass action expresn+on for the A /B halfcell (Eqn. 7) indicates that the
greater the capacity of oxidant A to serve as an electron acceptor, the greater the value
of K . Taking common logarithms of Eqn. 7 and rearranging gives
1 {B}
−log{e }= pe = logK A/ B−log (10)
n {A }+
1
International Union of Pure and Applied Chemistry
2 where, as the negative common logarithm of electron activity, pe is by analogy similar to
the definition of pH = log{H }. When {B} = {A }+
{B} =1 (11)
{A }
such that
1 A/ B 0
pe = logK = pe (12)
n
This gives from Eqns. 3 and 5
0 1 {B}
pe = pe − log n+ (13)
n {A }
Considering that
− ΔG 0
lnK = 2.3logK A/ = (14)
RT
then from Eqn. 5
0 − ΔG 0
pe = (15)
2.303nRT
and
pe = − ΔG (16)
2.303nRT
Combining Eqns. 13, 15 and 16 gives the change in free energy for the halfcell reaction
illustrated by Eqn. 1.
ΔG = ΔG + 2.303RT log {B} (17)
{A }
In most aqueous systems, electron transfer in a complete oxidationreduction
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 halfcell are
conveyed to the electron acceptor at another electrode through an electrical circuit. The
3 free energy difference associated with the transfer of electrons between the two
electrodes is given by the fundamental electrochemical relationship
ΔG = −nFEm (18)
and at standard state
ΔG = −nFEm 0 (19)
Where F is the Faraday constant and Em is the potential difference in volts between the
two electrodes of the electrical cell
Because the absolute value of Em depends fundamentally on two different half
cell reactions at separate electrodes, it is advantageous to adopt a standard reference
electrode against which to measure the tendency of redox active substances to donate or
accept electrons. This is done by according the standard hydrogen (halfcell) electrode
(SHE) with a potential of zero with unit electron, H and H ac2ivities. Relative to the
SHE, Em values are expressed more explicitly in terms of Eh, the potential difference of
the electrical halfcell relative to the SHE, which gives
ΔG = −nFEh (20)
and additionally
ΔG = −nFEh 0 (21)
Another important outcome from Eqn. 21 is that the standard potential of a halfcell
0 0
reaction Eh can be calculated directly from ΔG , or indirectly from standard free energy
of formation values (Table 5.1).
Substitution of Eqn. 20 and Eqn. 21 into Eqn. 17 and rearranging yields the
Nernst Equation
2.303RT {A }+
Eh = Eh +0 log (22)
nF {B}
This can be generalized for any half cell reaction
2.303RT activityproduct of oxidizedspecies
Eh = Eh +0 log (23)
nF activityproduct of reducedspecies
4 Table 5.1: Examples of common halfcell reactions with corresponding Eh , ΔG ,0
and log K values.
Conjugate 0 ΔG 0
Oxidant Reductant Eh (V) (kJ/mol) log K
O 2 H 2 1.231 118.7 20.8

NO 3 N 2 1.248 120.5 21.1
NO  NO  0.840 81.1 14.2
3 2
NO 3 NH 4 0.882 85.1 14.9
3+ 2+
Fe Fe 0.770 74.3 13.0
Fe(OH) Fe 2+ 1.065 102.8 18.0
3
SO 42 H 2 0.311 30.0 5.3
2 
SO 4 HS 0.251 24.3 4.3
CO 2 CH 4 0.170 16.4 2.9
CO 2 CH O2 0.071 6.9 1.2
+
H H 2 0.000 0.0 0.0
5.2 Redox Stability Diagrams
As illustrated by Eqn. 8, redox reactions often involve consumption or production
of protons. This means that Eh can be expressed as a function of pH on the basis of the
Nernst Equation (Eqn. 15). Theoretical expressions developed from this approach
provide a convenient way to display stability relationships between oxidized and reduced
species in redox reactions; however, an important point to remember is t
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