INTRODUCTION TO OXIDATIONREDUCTION CONCEPTS
• Oxidation-reduction reactions control solubility of
minerals or solid phases, influence adsorption processes,
and control fate and transport of organic and inorganic
species.
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APPLICATIONS:
Landfills
Acid Mine Drainage
Hydrocarbon Plumes
Chlorinated Solvent Plumes
Sewage Treatment Plant Discharges
Radionuclide Leachate Plumes
OXIDATION-REDUCTION REACTIONS
• A reduction-oxidation (or redox) reaction
is a chemical reaction in which electrons
are transferred completely from one
species to another.
– The chemical species that loses electrons
(reducing agent, electron donor) in this
charge transfer process is called oxidized.
– The chemical species receiving electrons
(oxidizing agent, electron acceptor) is called
reduced.
OXIDATION-REDUCTION REACTIONS
• EXAMPLE: Benzene Oxidation Half Reaction
– C6H6 + 12H2O = 6CO2(g) + 30H+ + 30e• EXAMPLE: Goethite Reduction Half Reaction
– 30FeOOH(s) + 90H+ + 30e- = 30Fe2+ + 60H2O
• Coupled Redox Reaction
– 1/6C6H6 + 5FeOOH(s) + 10H+ = CO2(g) +
5Fe2+ + 8H2O
NERNST EQUATION
• Eh = Eo + RT/nF ln (activity of oxidized species)
(activity of reduced species)
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R = 1.987 x 10-3 kcal/mol K
T = temp (K)
n = number of electrons
F = 23.06 kcal/V g eq
at 25°C and converting natural ln to log10
• Eh = Eo + 0.0592 log (activity of oxidized species)
n
(activity of reduced species)
NERNST EQUATION: Example
• Fe2+ = Fe3+ + e- Eo = 0.77 volt
• Eh = 0.77 + 0.0592 log a Fe3+
1
a Fe2+
– If aFe3+ = aFe2+, then Eh = 0.77 volt = Eo
– If aFe3+ /Fe2+ = 10-2, then
– Eh = 0.77 + 0.0592log[10-2] = 0.65 volt
BALANCING REDOX REACTIONS
1. Write for each half reaction the oxidized and reduced
species into the two equations and balance the
elements, except H and O at left and right.
2. Balance the number of oxygen atoms by adding H2O.
3. Balance the number of hydrogen atoms by adding H+.
4. Balance electroneutrality by adding electrons.
5. Subtract the two half reactions, canceling electrons, to
obtain the redox reaction.
Useful oxidation state information: O(-II), Fe(II), Fe(III),
O2(g)(0), H(I), OH(-I) Mn(II), Mn(III), Mn(IV), Cl(-I),
e(-I), S(VI)O4 2-
BALANCING REDOX REACTIONS
Step 1. Fe2+ = Fe3+
oxidation reaction
O2(aq) = H2O
reduction reaction
Step 2. Fe2+ = Fe3+
O2(aq) = 2H2O
balance oxygens
Step 3. Fe2+ = Fe3+
4H+ + O2(aq) = 2H2O
balance hydrogens
Step 4. Fe2+ = Fe3+ + e4e- + 4H+ + O2(aq) = 2H2O
balance electrons
Step 5. 4(Fe2+ = Fe3+ + e-)
multiply and add
4e- + 4H+ + O2(aq) = 2H2O
4H+ + 4Fe2+ + O2(aq) = 2H2O + 4Fe3+
This reaction may occur in low pH environments characteristic of
acid mine drainage and metal processing sites.
REDOX REACTION SEQUENCES AND REDOX LADDERS IN NATURAL SYSTEMS
The above figure shows the Eh value at pH 7 and 25°C for important redox couples in natural waters under
specified conditions (Langmuir, 1997). The reduction reactions and Eh-pH equations that correspond to these
couples at 25°C are provided in a table in order of decreasing Eh. Langmuir D., 1997, Aqueous Environmental
Geochemistry, Prentice Hall, New Jersey, 600 p.
IMPORTANT VARIABLES THAT CONTROL REDOX REACTIONS
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The oxygen content of recharge water.
– Near surface groundwater with dissolved oxygen present tend to be
relatively oxidizing.
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The distribution and reactivity of organic matter and other
potential reductants.
– Surface waters with low organic carbon content tend to be relatively
oxidizing.
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The distribution of potential redox buffers (MnO2, Fe2O3, Fe(OH)3).
– The Fe(III)/Fe(II) redox couple controls dissolved iron concentrations and is in
equilibrium with ferric hydroxide.
– This redox couple is important under oxidizing, acidic and reducing, near
neutral pH conditions typical of acid-mine drainage and organic-rich
hydrocarbon and chlorinated solvent) contaminated waters, respectively.
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The circulation rate (residence time) of groundwater.
Bacterial processes control the rates of redox reactions involving
oxygen, iron, manganese, nitrogen, sulfur, carbon, and other redox
sensitive elements.
Master
chemical
variables
Eh
pH
Eh-pH DIAGRAM
FOR PART OF
THE IRON
SYSTEM
(Brookins, 1988)
Reduced Zone [Fe(II) Stable, Fe(III)
Solids Dissolve]
Oxidized Zone [Fe(III) Solids
Precipitate]
BEI of Fe(OH)3 in a sandstone uranium rollfront
Uranium is
dissolved and
mobile
Eh
pH
for
UO2
Uranium is
solid and not
mobile
REDOX REACTIONS OF ORGANIC COMPOUNDS
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REDUCTION REACTIONS WITH ORGANIC COMPOUNDS
FREQUENTLY INVOLVE A LOSS OF OXYGEN AND HALOGEN
(Cl, Br) ATOMS AND A GAIN OF HYDROGEN ATOMS.
EXAMPLE: REDUCTIVE DEHALOGENATION REACTIONS WITH
TETRACHLOROETHENE
C2Cl4 → C2Cl3H → C2Cl2H2 → C2ClH3
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LOSS OF CHLORINE ATOMS
GAIN OF HYDROGEN ATOMS
OXIDATION REACTIONS WITH ORGANIC COMPOUNDS
FREQUENTLY INVOLVE A GAIN OF OXYGEN ATOMS AND A
LOSS OF HYDROGEN ATOMS (PROTONS).
Methanogenesis
Immobilization
Reactions
CH3COOH + H2O = CH4 + HCO3- + H+
Sulfate Reducing
SO4 + 9H+ + 8e- = HS- + 4H2O
2-
Iron Reducing
Fe(OH)3(s) + 3H+ + e- = Fe2+ + 3H2O
Manganese Reducing
Denitrifying
Aerobic
Cu2+ + 2[=FeOH](S) = 2[=FeO]Cu(S) + 2H+
U(VI)OH+ + 2e- = U(IV)O2(S) + OHxNi2+ + (1-x)Fe2+ + HS− = Fe(1−x )NixS(S) + H+
Pb2+ + HCO3- = PbCO3(s) + H+
REDUCTION OF ORGANIC COMPOUNDS
• Biotransformations of organic compounds are
controlled by oxidation and reduction reactions.
• Redox reactions generally are controlled by
microorganisms.
• Aromatic hydrocarbons (BTEX) biodegrade through
oxidation reactions, eventually forming carbon dioxide
and water.
• Chlorinated aliphatic hydrocarbons biodegrade
through reductive dehalogenation reactions by addition
of hydrogen atoms.
• Important electron acceptors (species undergoing
reduction) include oxygen, nitrate, ferric hydroxide,
sulfate, and carbon dioxide.
Expression for Conservation of electrons
Stoichiometric coefficients of electrons in species
Element stoichiometry
Element valence