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GES166/266, Soil Chemistry
Winter 2000
Lecture Supplement 4
Oxidation and Reduction (Redox) Reactions
The hydrogen ion (H+) and the availability of electrons (e-) are the master
variables of soil chemical reactions. Unlike H+, electrons are not 'free forming'; they are
contained within atoms or molecules. Thus, electrons are only transferred between
species. Redox, short for reduction-oxidation, is the termed used to denote the transfer
of electrons. Oxidation and reduction are defined as:
Reduction: The gain of electrons by a compound
Oxidation: The loss of electrons by a compound
Remember, oxidation and reduction MUST occur together! An electron cannot exist as
an isolated entity; it is transferred from one species (the reductant) to another (the
oxidant). Individual oxidation or reduction reactions are termed half-reactions; half
reactions must be coupled to give an overall reaction. When an oxidation or reduction
reaction is written independently; for example, the reduction of CO2
CO2 + 4e- + 4H+ = CH2O + H2O
(photosynthesis)
[V.1]
or for the oxidation of H2O
2H2O = O2 + 4e- + 4H+
[V.2].
a ‘free’ electron (e-) written in the equation. An overall redox reactions will never have
an e- shown. You can couple [1] and [2] together for an overall reaction:
CO2 + H2O = CH2O + O2
[V.3]
In soils, micro-organisms play a key role in driving redox reactions. They drive
redox reactions by using a species as an electron donor (usually reduced carbon) or as an
electron acceptor (usually oxygen in aerobic environments). We can, however, classify
compounds as electron donors or acceptors
Electron Donors
By far the most prevalent electron donor is organic carbon. Carbohydrate oxidation
serves as an example of organic carbon oxidation to carbon dioxide.
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CH2O + H2O = CO2 + 4eThis half reaction is what supplies energy to microorganisms within soils (we will speak
of the electron acceptor they use in a moment). Other electron donors that occur in soils
and participate in chemical redox couples are compounds found in reduced (oxygen
limited) soils. Examples include:
Mn(II) = Mn(IV) + 2eFe(II) = Fe(III) + eS2- = SO42- + 8eAs(III) = As(V) + 2eall of which are used much less frequently than carbon sources. Bacterial oxidation of
Fe(II) and sulfide are common in mine spoils and reduced sediment; sulfide is also used
as an electron donor by bacteria in hydrothermal marine vents.
Electron Acceptors
The most prevalent electron acceptor is O2 (g), and is the sole electron acceptor in aerated
systems (aerobic systems).
O2 + 4 e- + 4H+ = 2H2O
[V.4]
Living compounds create energy by oxidizing carbohydrates and reducing oxygen. That
is, organisms take electrons from the organic carbon, run it through their metabolic
cycles, and then dump the electron to an electron acceptor. When oxygen is not present,
microorganisms must seek alternate electron acceptors. The energy gain for the
organisms is the energy difference between reduced carbon and the electron acceptor. In
order of favorability, electron acceptors are:
O2 > NO3- > Mn(VI) > Fe(III) > AsO43- >SO42A lack of oxygen leads to an anaerobic condition and results in the build up of
reduced species: Mn(II), N2, Fe(II), As(III), and S2-. When reduced species are build up
in the system, it is termed "REDUCED". When O2 is present with reduced species such
as S2-, Fe(II), Mn(II), or As(III) it is energetically favorable for oxygen to oxidizing the
previous species and thus become reduced to water. Thus, if oxygen is introduced to a
‘reduced’ system, the reduced species are oxidizes in opposite order:
SO42- > AsO43- > Fe(III) > Mn(VI) > NO3We will specify the exact energies of these overall reactions shortly.
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A Hypothetical Thought: Electron Activity (pe)
Although this is not truly correct, we can think of the availability of electrons in terms
of an electron activity. In this way, electrons are treated just like H+: They have an
artificial ACTIVITY (mol / l). Usually, electron activity is written as pe
- log (e-) = pe
Given this notation, we can treat electrons just as any other reactant. Let’s consider a
generalized half reaction to illustrate this convention (and its utility).
OX + n e- = RED
K = (RED) / (OX) (e-)n
log K = log (RED)/OX - n log e= log (RED)/OX + npe
log K - npe = log (RED)/(OX)
- note: we could measure pe and know the ratio of RED/OX
Electrode Potential
While pe is a convenient way to 'view' redox reactions, it is not real. Since electrons do
not exist in solution we need a separate way of measuring their reactivity in a system.
This is done with an electron potential: The potential for an electron to participate in a
reaction.
Derivation of the System Potential (Eh)
∆Grxn = ∆Gro + RT ln Q (Lewis Eqn)
[V.5]
where Q is the reaction quotient and is analogous to the ion activity product (IAP) we
discussed in the “Precipitation” section.
∆Gr = -nFE or
∆Gro = -nFEo
where n = # of e- (in mols)
and F = 96,500 C / mol esubstituting [2] into [1] yields
or
[V.6]
23,061 Cal / V • mol e-
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-nFE = -nFEo + RT ln Q
then
E = Eo - RT/nF ln Q
For a reaction:
Oxid + e- = Red
This then simplifies to the Nernst Equation:
E = Eo - 0.059/n log [ Red / Oxid]
or we could simply write a more generic, complete reaction as
aA + bB = cC + dD
and
E = Eo - 0.059/n log {[A]a [B]b / [C]c [D]d }
If we are at equilibrium, then Q (IAP) = Keq and
Eo (V) = (0.059 / n) log K
• Converting between Eh and pe:
Eh(mV) = 59.2 pe
Application of pe and Eh
Most soils maintain a pH between 3.5 and 9; a pe of -6 to +12 (Eh: -355 to 710 mV);
and a pe + pH of 2 - 18.
pe and Eh Ranges
Soils have been broadly classified into 3 redox states depending on their pe values:
Oxic:
Suboxic
Anoxic
pe > 7
pe 2-7
pe < 2
Eh > 400 mV
100 < Eh < 400 mV
Eh < 100 mV
but this classification does not include a necessary consideration--pH. A better way to
classify the redox status is to simply state that suboxic zones are characterized by the
reduction of Fe and Mn-oxides, but not the reduction of sulfate. Anoxic systems have
been lowered to where sulfate becomes an active electron acceptor, yielding S2- or other
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reduced sulfur species. Methane (CH3) is also produced under highly anoxic conditions.
In soils, realize that there may be microenvironments that are anoxic very close to areas
that are oxic.
How do soils become oxidized or reduced, and how are they oxidizing or reducing?
The redox potential is governed by two primary factors:
1. Microbial activity
2. Availability of electron acceptors.
In the oxygen supply is the dominant factor controlling the redox status of a soil; when
oxygen is limited other terminal electron acceptors become important.
Chemical species also influence the redox status of systems, but in the presence of
microbial activity they are a small consideration compared to the microorganisms
themselves. Let’s ask another important question regarding the redox status of a soil.
Why does the presence of O2 increase the redox potential of the system and make it
more oxidized? The answer is that coupling oxygen’s reduction half-reaction with any
redox couple having a lower Eh value will result in an energetically favor redox
reaction. This leads to the reduction of O2 and oxidation of the originally reduced
compound.
For example, consider Fe(II) coming in contact with molecular oxygen (O2):
H+ + 1/4 O2 + e- = 1/2 H2O
1.229
Fe2+ = Fe3+ + e-0.711 V
so
Fe2+ + 1/2O2 + H+ = Fe3+ + 1/2 H2O
0.518
A positive ∆Eh, or Ehrxn, means that the reaction is thermodynamically favorable. Thus,
if kinetic constraints are not limiting oxygen will oxidize everything with a lower E(h)
value and cause the system to have a 'low electron activity' (or high redox potential).
Microorganisms
Microorganisms are probably the most important factor controlling the redox status of
soils:
-they reduce carbon to store energy
-they oxidize it to release energy
-they use oxygen or other molecules as electron acceptors, thus forming reduced
species under oxygen limiting conditions.
Micro-organisms control the redox potential along with the redox couple of the
electron acceptor. Oxygen is the preferred acceptor because it is most easily reduced to
water of the available acceptors. This leaves the greatest energy from respiration. In
its absence, other redox couples must be used to put electrons liberated from the
oxidation of carbon compounds.
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photosynthesis stores energy by reducing C
Red. of C
balancer
CO2 + 4 e- + 4 H+ => CH2O + H2O
2 H2O = 2O2 + 4e- + 4 H+
0.21 V
overall
CO2 + H2O = CH2O + O2
-1.00 V
-1.21 V
To acquire the stored energy the C source is oxidized
=> CO2 + 4 e- + 4 H+
Oxid of C
CH2O + H2O
Red of O2
2O2 + 4e- + 4 H+ => 2 H2O
Overall
CH2O + O2 => CO2 + H2O
+1.00
Alternate electron acceptors, in order of preference (based on E° values):
NO3 -> N2 (or other reduced N forms)
Mn(IV) --> Mn(II)
Fe(III) --> Fe(II)
SO42- --> S2- (or H2S)
The order of their preference is due to their redox potentials for the half-reactions (see
Table 7.1 in Environmental Chemistry of Soils by M.B. McBride). Note that O2 has
the highest redox potential, followed by NO3>Mn(IV)>Fe(III)>SO42-. The reverse
order will apply upon oxidation. Sulfide will go first, followed by Fe(II), and then
Mn(II). Also note that if a reduced species with a low redox potential is present with a
species in the oxidized form which has a higher redox potential, then there is a high
probability for oxidation of the reduced species and reduction of the oxidized species.
Example, if S2- and MnO2 were present together.
1/8H2S + 1/2 H2O = 1/8 SO42- + 5/4H+ + e1/2 MnO2 (s) + 2H+ + e- = 1/2 Mn2+ + H2O
1/8H2S + 1/2MnO2 (s) + 1/2H2O + 3/4H+ = 1/8SO42- + 1/2 Mn2+ + 1/2H2O
-0.303
1.23
0.927
The selectivity of the terminal electron acceptor also leads to the soil being 'Poised' at a
certain redox potential. This means that as a soil becomes reduced it does not have a
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smoothly decreasing potential. Rather the redox potential moves in steps from the
potential of one redox potential to another along the sequence discussed above.
pH Changes
As soils or sediments become reduced, their pH tends to move toward neutrality.
That is, if the pH of the system is acidic then the pH will increase while soils with pH
values above 7 tend to have decreasing pH values.
Reasons:
• When the pH is initially low, H+ consumption in the reduction reactions
increases the pH.
For example: MnO2 (s) + 4 H+ + 2e- = Mn2+ + 2H2O
• If the pH is initially basic, then the liberation and reprecipitation of metal ions
such as Fe and Mn as hydroxides, carbonates, or sulfides tends to lower the pH.
For example: Fe2+ + 2 H2O = Fe(OH)2 + 2 H+
Fe2+ + HCO3 = FeCO3 + H+
Soils that undergo seasonal flooding and drying tend to become more acidic. Based on
reversible reaction such as that for sulfate-sulfide you would not expect this.
1/8 SO42- + 5/4H+ + e- = 1/8H2S + 1/2 H2O
However, if sulfate is leached from the system then the acidity generated in the oxidation
direction will not be removed by reduction. This then leads to a gradual acidification of
the soil.
Specific Cycles
Iron Cycles:
Iron goes through pronounced changes with redox potential changes.
-Fe(III)-oxides are dominant in oxidized systems
-suboxic soils result in high Fe(II) concentrations
* Fe(II) can precipitate as FeCO3 (siderite)
-in anoxic systems, Fe(II) may precipitate as FeS / FeS2 (pyrite)
-upon aeration, the reduced Fe forms are oxidized quickly
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Manganese Cycles:
Like Fe, Mn goes through extensive changes with redox potential changes.
-Mn-oxides dominate in aerated systems
-Mn(II) is formed upon reduction
-MnCO3 forms in reduced (suboxic) systems, less likely are MnS but they do
form
-oxidation is much slower than for Fe(II) except at mineral surfaces
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O2
Oxic
zone
Mn2+
oxidation
precipitation
2+
Mn
+ O 2 = MnO 2
diffusion
Suboxic
zone
Anoxic
zone
Mn 2+
MnO2 + 4H + = Mn 2+ + 2H 2 O
reductive dissolution
Mn2+ + CO3 2- = MnCO 3
Fe-carbonate precipitation
Nitrogen Cycles
-Oxidation of N2 does not occur except in a few organisms
-Reduction of NO3- to reduced-N forms accounts for denitrification in soils
-NO3- to NH4+ is termed mineralization and only occurs in living organisms
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Sulfur Cycles
• Sulfur goes through dramatic changes in redox cycling. It covers a full 8 electron
transfer sequence with most intermediates being possible.
• We usually only consider sulfate and sulfide, but we should look at the other
possible oxidation states.
-Sulfate dominates in oxidized environments
-anoxic conditions leads to the formation of sulfides and other reduced S forms
-plant tissue degradation leaves many intermediate S oxidation states
-sulfide tends to form strong, insoluble complexes with soft to intermediate acids:
Fe(II), Cd, Cu, Pb, Hg, ...
-oxidation of sulfide leads to acidification of many systems (acid mine drainage)