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EVSC 590

Oxidation -Reduction (Redox)

Soil Microbiology:

An Exploratory ApproachChapter 2 and 13

Pages 17-18 and 163-165

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Redox Reactions

An Oxidation-Reduction (redox) reaction is

a chemical reaction in which electrons aretransferred completely from one species to

another.

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Redox Reactions

The chemical species which loses electrons

in this transfer process is called oxidized

whereas the one receiving electrons is

called reduced.

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Oxidation is the loss of electrons and

Reduction is the gain of electrons.

All chemical elements can accept or donate

electrons under appropriate conditions.

Redox Reactions

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Redox Reactions

Oxidation and Reduction always occur

together because a substance can only

donate electrons if another substance canaccept it.

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Redox Reactions

The chemical conditions of soils and nature

limit the number of elements involved in

natural electron exchange.

Many inorganic chemical reactions and

virtually all biological reactions of C, N,and S are oxidation-reduction (redox)

reactions

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Redox Potentials

The redox reaction below has two half

cells

H2+ 1/2O2 H20 Eq. 1an oxidation

H2 2H+ + 2e- Eq. 2and a reduction

1/2 O2 + 2H+ +2e- H20 Eq. 3

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Redox Potentials

The tendency of reactions to occur is based

on the relative tendencies of reactions 2 and

3 to proceed in the directions indicated.

This tendency is determined by the

electromotive force or potential(E) of the

reaction.

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Redox Potentials

The electromotive force for a half reactionis determined by measuring the electricpotential generated relative to a referenceelectrode(normally a hydrogen electrode).

By convention, the H electrode is assigned ahalf reaction potential (Eh) of 0.00v under

standard state conditions, i.e.H+ (1M) + e- 1/2H2 (1atm) Eh

o=0.00V

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Redox Potentials

The Eho is corrected to pH 7 and designated

Eho to reflect metabolic activity at neutral

conditions.

To facilitate comparisons, half reactions arewritten as reductions and the sign of the

corresponding electrical potential isadjusted accordingly.

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Redox Couples

In redox couples oxidized form is always

placed on the left.

e.g. 2H +/H2 or O2/ H2O

H+ and O2 are the oxidized forms

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Redox Couples

When constructing complete oxidation

reduction reactions from constituent half

reactions, the reduced substance of a redoxcouple whose reduction potential is the

more negative donates electron to the

oxidized substance whose potential is more

positive.

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Redox Potentials

The standardized values are referred to as

reduction potentials , for example,

2H+ + 2e- H2 Eo= -0.42V Eq.41/2 O2 + 2H

++2e-H20 Eo=+0.82V Eq.5

Since Eq 4 is more negative than Eq 5, Eq 5has a greater tendency to proceed as written.

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FREE ENERGY

For the reaction; H2 +1/2 O2H20

Eo = +0.82 (-0.42) = +1.2

Go = -(2) (23.1) (1.24)

= -239 KJ mol-1 ofH20 produced

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FREE ENERGY

One very useful measure of metabolic

reaction is the free energy change also

known as the Gibbs free energy.

Gro is the intrinsic energy contained in a

given substance.

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FREE ENERGY

For a reaction X Y, the free energychange DGo is the difference between the

free energy of the product Y and thereactant X

DGo = Gy -Gx

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FREE ENERGY

Negative DG = Exergonic,

Positive DG = Endogonic,

Zero DG = Equlibrium

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FREE ENERGY

DGo for redox reaction can be calculated

using the following equation.

DGo = -nFDEo

N = # of electrons transferred

F = Faradays constant (96.5KJV-1 mol-1)

= 23.1 V-1 mol-1

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FREE ENERGY

DEo= difference in reduction potentials for

the two half reactions comprising the

combined reaction. For the reaction; H2 + 1/2 O2 H20DEo = +0.82 (-0.42) = +1.2

DGo = -(2) (23.1) (1.24)

= -239 KJ mol-1 of H20 produced

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FREE ENERGY

For the reaction

Ox + mH+ + ne Red DGr = DGro + RTln(Red)/(Ox)(H+)m

DGro= -nFEo

DGr= -nFE

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Nerst Equation

Nernst Equation relates activities of to Eh

-nFE = -nFEo + RT ln(Red)/(Ox)(H+)m

Eh= Eho- RT/nFln(red)/(Ox) - mRT/nF lnH+

Reduces to Eh = 2.303RT/F x pe

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Concept of pe

Just as pH is based on moles L-1, redox

potential can also be expressed in terms of

pe (-log of electron activity) which iscompatible with units of moles per liter.

In this way, electrons can be treated as other

reactants and products so that both can beexpressed by a single equilibrium constant.

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Concept of pe

pe = -log (e-)

It measures the relative tendency of a

solution to accept electrons.

Reducing solutions have low pe and tend to

donate electrons to species placed in thesolution.

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Concept of pe

Oxidizing solutions have high pe and tend

to accept electrons from species placed in

the solution. Large values of pe favor the existence of

electron- poor (i.e) oxidized species just as

large values of pH favor the existence ofproton poor species (i.e bases)

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Concept of pe

Small values of pe favor electron- rich, or

reduced, species, just as small values of pH

favor proton rich species, acids.

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Concept of pe

Unlike pH, however, pe can take on

negative values.

The largest pe value is +13.0 and the

smallest is near -6.0.

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pe vs Eh

Eh (millivolts) = 59.2pe

Eh (Volts) = 0.059 pe

e.g. For (CO2/ CH2O), -0.43V

pe = 0.059 x0. 43 =

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Table 1. Redox Pairs Arranged in Order from the

Strongest Reductants to the Strongest Oxidantsa

(Paul and Clark, 1989).

Redox Pair pe Eh (volts)

CO2/CH2O -7.3 -0.43

SO4

2-/H2S -3.7 -0.22

NO3-/N2 7.11 0.42

MnO2/Mn2+ 8.1 0.48

NO3-/N2 12.4 0.74

Fe3+/Fe2+ 12.9 0.761/2 O2/H2O 13.9 0.82

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Concept of pe

In soils the pe range can be divided into 3

parts that corresponds to:

1) oxic soils (pe > + 7 at pH 7)

2) Sub oxic soils (2 < pe < + 7 pe at pH 7)

3) Anoxic soils ( pe < + 2 at pH 7)

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pe + pH

Most soil systems consist of aqueous

environments in which the dissociation of

water into H2(g) or O2(g) imposes redoxlimit on soils

On the reduced side the redox limit is given

by the reaction;H+ + e- 1/2 H2 g Rxn.1

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pe + pH

H+ + e- 1/2 H2 g K = (H2)g

1/2/(H+) (e-) ,

K = 1, thus log K = 0

or log K = 1/2 log H2 (g) - log (H+) - log(e-)

At H2

=1 atm, pe + pH = O

This represents the most reduced equilibrium

conditions in natural aqueous environments.

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pe + pH

On the oxidized side the redox limit is given

by the reaction

H+

+ e-

+ 1/4 O2(g) = 1/2 H2O Rxn.2K =1/2 H2O /(H

+)(e-)(O2)1/4 =1020.78.

Log K =Log (H+) -log(e-) 1/4 log O2(g)

= 20.78

pe + pH = 20.78 + 1/4 O2(g)

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pe + pH

Thus when O2 is 1 atm,

pe + pH = 20.78.

This represents the most oxidized

equilibrium conditions in natural aqueousenvironments

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pe + pH

When the redox limits of natural aqueous

environments defined by reactions 1 and 2

are plotted we get a graph which is knownas a pE-pH diagram, and shows the domain

of electron and proton activity that has been

observed in soil environment worldwide.

Both pe and pH are needed to specify the

redox status of aqueous systems.

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Redox in soils

This pe range can be divided into 3 parts

that corresponds to:

1) Oxic soils (pE > + 7 ) at pH 7

2) Sub-oxic soils (2 < pE < + 7 ) at pH 7

3) Anoxic soils (pE < + 2 ) at pH 7