Redox Reactions

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Redox Reactions

• Oxidation loss of electrons
• Reduction gain of electrons
• Oxidation-reduction reactions are the most
  energetic (greatest enthalpies) in chemistry and
  geochemistry.
• Examples
• Photosynthesis (reduction)
• Respiration (oxidation)
• Corrosion
• Pyrite weathering
• Nitrogen cycle reactions

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Redox Reactions (Cont.)

• Voltage and Free Energy
• The free energy change (? G) associated with the
  transfer of n electrons is equivalent to a
  voltage drop (E) experienced by the electrons
• ? G -NFE
• Where F is the faraday constant which is the
  charge (in coulombs) of a mole of electrons (F
  9.64846 x 104 C/mol).
• N is the number of chemical equivalents
• Charge ? Voltage Energy

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Redox Reactions (Cont.)

• The Nerst Equation
• By convention, we write half-reactions as
  reductions
• aA bB ne- cC dD
• The voltage of this reaction is
• E E RT/NF ln AaBb /CcDd
• Since R and F are constants, and changing from ln
  to log, we can use (at 25C)
• E E 0.059/N log AaBb /CcDd

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Redox Reactions (Cont.)

• We can measure the potential of a natural
  solution witrh respect to a standard hydrogen
  half-cell
• When this is done, the symbol Eh is used.
• We assign standard potential (E) of 0.00 volt
  for the reaction
• H(aq) e- ½ H2(g)
• Oxidation Potential (also Redox Potential)
• Eh -0.059 log P(H2)1/2 0.059pH

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Redox Reactions (Cont.)

• The Standard Hydrogen Electrode
• This cell would allow us to measure the E for the
  Cu2-Cu half reaction relative to the S.H.E.
• The salt bridge allows ions but not electrons to
  pass.

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Redox Reactions (Cont.)

• Two half reactions
• H2 ? 2H 2e-
• Cu2 2e- ? Cu
• Cu2 H2 Cu 2H
• Eh E 0.0592/N log AaBb /CcDd
• Where aA bB ne- cC dD
• Eh E 0.0592/N log H2/Cu2

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Calculating Eh

• Eh E 0.0592/N log AaBb /CcDd
• Where aA bB ne- cC dD
• Calculate the Eh of a groundwater in which Fe2
  10-4m, pH 7 and the solution is saturated in
  Fe(OH)3
• The Eh is determined from the couple
• Fe(OH)3 3H e- Fe2 H2O E 0.975 volts
• Eh 0.975 0.0592 log H3/Fe2
• Eh 0.975 0.0592 log 10-73/10-4
• Eh -0.0314 volts

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Redox Reactions (Cont.)

• Copper Zinc Battery
• Two half reactions
• Zn ? Zn2 2e-
• Cu2 2e- ? Cu
• Electron flow gives a voltage
• Depends on relative concentrations of the ions
• Electromotive force
• Tendency of the reaction to take place

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Redox Reactions (Cont.)

• Half-reactions are determined to have electrode
  potentials.
• The potential for any reaction may be calculated
  from published values.
• Zn ? Zn2 2e- -0.76 volt
• Cu2 2e- ? Cu 0.34 volt
• Zn Cu2 ? Zn2 Cu - 1.10 volt
• Any reaction with a negative potential will go to
  the right
• Any reaction with a positive potential will go to
  the left

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Redox Reactions (Cont.)

• Zinc is a reducing agent
• Gives up electrons
• We say it is oxidized
• Copper is an oxidizing agent
• Accepts electrons
• We say it is reduced

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Redox Reactions (Cont.)

• Lower Bounds on Eh
• Aqueous solutions cannot be at an Eh lower than
  that required to reduce water
• 2H2O 2e- H2 2OH- or 2H 2e- H2 E
  0.00 volt
• From the Nerst equation, we have
• Eh Eo 0.059/N log (H2 / H2)
• The standard potential Eo, when pH 0.0 and
  P(H2) 1, is defined to be 0.0 by definition at
  298oK. Now, at the Earth's surface, the highest
  possible partial pressure of hydrogen is simply
  1.0. Hence from the Nerst equation, we can write
• Eh -0.059pH

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Redox Reactions (Cont.)

• Upper Bound on Eh
• Aqueous solutions cannot be at an Eh higher than
  that required to oxidize water
• 2H2O 4e- O2 4H E 1.23
• From the Nerst equation, we have
• Eh Eo 0.059/N log (O2H4/H2O2)
• Eh 1.22 0.059/4 logH4 pO2
• We can assume a(H2O) 1. In the Earth's
  atmosphere, the partial pressure of oxygen is
  0.2. Therefore, we can write
• Eh 1.23 - 0.059 pH 0.015 log (0.2)
• Eh 1.22 0.059 pH

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• In nature, Eh varies from about 0.8 to 0.5 V,
  and pH varies between about 1 and 10
• The Range of Eh and pH values in the environment
  is shown to the right

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Eh and pH

• We can think of Eh as reflecting the abundance of
  electrons in the environment
• We can think of pH as reflecting the abundance of
  protons in the environment.
• Eh-pH diagrams illustrate reaction stability
  fields under natural conditions

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Redox Reactions (Cont.)

• In general, environments in contact with the
  atmosphere (e.g. rainwater, oceans, streams) have
  higher Eh and / or pH than environments isolated
  from the atmosphere (e.g. peat bogs).
• The potential for oxygen reactions is not
  reached, and the oxygen acts as a weak oxidizing
  agent.

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Controls on Eh and pH

• The major controls on Eh and pH are
• (i) organic processes photosynthesis,
  respiration, decay
• (ii) redox reactions involving Fe, S, N, C
• (iii) balance between dissolved CO2 and CaCO3 in
  natural waters

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Controls on Eh and pH

• Selected environments

Modified from Brownlow, 1996
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FIELD APPARATUS FOR Eh MEASUREMENTS
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Eh and pH

• Limitations of Eh-pH daigrams
• Eh very difficult to measure
• Typically a Pt electrode is used, but most redox
  reactions dont react on a Pt surface at near
  neutral pH
• The actual process of sampling will locally alter
  the geochemical environment
• Natural systems are rarely at equilibrium
• Many different redox equilibria govern pairs of
  half-reactions
• No single control on system Eh

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Difficult to get meaningful Eh results
21
Eh-pH

• However
• Eh-pH diagrams can still show us what is likely
  in a particular system if it could reach
  equilibrium

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pe

• Calculations - an easier way
• Ignore Eh and E by using pe
• Define pe as
• pe - loge-
• similar to pH
• Converting Eh to pe
• Eh 2.303RT/F pe
• Or Eh 0.059 pe

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pe

• Pe-pH diagrams
• Show stability fields for solid species and ions

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Construction of a pe-pH Diagram

• Need to know species of interest.
• Iron in water Fe O C System
• Ferrous iron
• Fe2
• Fe(OH) or Fe(OH)2
• FeCO3
• Ferric iron
• Fe3
• Fe(OH)2- or Fe(OH)3
• Assumptions
• Boundaries with aqueous species are picked at
  equal activity ratios
• Boundaries with solids require a choice of
  aqueous activity

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Construction of a pe-pH Diagram

• Start with the redox diagram for water

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Construction of a pe-pH Diagram

• Add redox diagram for ferrous/ferric iron
• For Fe2 ? Fe3 e-
• K Fe3e-/Fe2 10-13.05
• logFe3 pe logFe2 -13.05
• pe 13.05

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Construction of a pe-pH Diagram

• The boundary for Fe(OH)2 - Fe3
• Fe(OH)2 ? Fe3 OH-
• K Fe3OH-/Fe(OH)2 10-11.6
• logFe3 logFe(OH)2 pH 2.4
• pH 2.4

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Construction of a pe-pH Diagram

• Consider the solid Fe(OH)3
• Fe(OH)3 3H e- ?
• Fe2 3H2O
• Pe 23.9 3pH

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Construction of a pe-pH Diagram

• Final diagram after several additional
  calculations

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Eh-pH Diagram for Fe-O-S System

• Iron sulfate ion-pairs at low pH
• Iron sulfides (pyrite and pyrrhotite at low Eh

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Eh-pH Diagrams and Speciation

• As(III) species are much more toxic than As(V)
  species.
• AsO43- will strongly sorb to FeOOH

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Redox Ladders

• Assume that redox conditions are relatively
  constant
• Factors that control the abundance of electron
  donors and acceptors (along with kinetics)
  determine the Eh.
• Usually catalyzed by micro-oganisms
• Organic carbon is the most common electron donor.
• Near the surface, electron acceptors are
  dissolved oxygen and nitrate.
• Other electron acceptors originate in the
  weathered rock material
• Sulfate
• Ferric iron
• Mn
• CO2

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Redox Ladder

• General sequence of reaction for neutral pH
• Reaction Eh (pH 7)
• ½O2 2H 2e- H2O 0.816 V
• NO3 6H 5e- 1/2N2 3H2O 0.713 V
• MnO2 4H 2e- Mn2 2H2O 0.544 V
• NO3 2H 2e- NO2- H2O 0.431 V
• NO2 8H 6e- NH4 2H2O 0.340 V
• Fe(OH)3 3H e- Fe2 3H2O 0.014 V
• Fe2 SO42 16H 14e- FeS2 8H2O -0.156
  V
• S 2H 2e- H2S -0.181 V
• SO42 10H 8e- H2S 4H2O -0.217 V
• HCO3- 9H 8e- CH4 3H2O -0.260 V
• H e- ½ H2 -0.414 V
• HCO3 5H 4e- CH2O H2O -0.454 V

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Redox Ladder

• Sequence of changes in concentration, either with
  time or distance, in a flow system

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Geochemical Environments

• Based on measured dissolved oxygen and H2S
• Easy to measure
• Strongly tied to redox reactions
• Concentrations have a major effect on organisms
• Control the formation of authigenic minerals
• Cannot coexist
• Bacteria divide into aerobes (killed by H2S) and
  anaerobes (killed by oxygen)

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Geochemical Environments (Berner 1981)

• Oxic Environments
• Dissolved oxygen gt 30 µM
• Mn2 below detection
• Suboxic Environments
• Dissolved oxygen gt 1 µM, lt 30 µM
• Fe2 below detection
• Mn2 detectable
• Anoxic Environments
• Dissolved oxygen lt 1 µM)
• Anoxic Sulfidic
• H2S gt 1 µM
• Anoxic Non-sulfidic
• H2S lt 1 µM
• Post-Oxic
• Methanic

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Geochemical Environments
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Redox Buffering

• Sequential utilization of electron acceptors
  maintains the redox potential at specific
  intervals
• Until that acceptor is used up

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Redox Examples

• Aquifer in Illinois (from Barcelona et al. 1989)

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Redox Examples

• Changes in composition along distance in
  groundwater flow, Lincolnshire, England (Champ
  and Gulens, 1979)

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Redox Examples

• Profile of dissolved species in a sediment core
  from a NC estuary (Martens and Goldhaber, 1978)

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• SEQUENCE OF REDUCTION REACTIONS
• If a system is effectively a closed system and
  there is abundant carbon to support microbial
  activity, then there is a well defined
• sequence of reduction reactions affecting the
  inorganic species present in the system.
• For each of the major redox elements there is a
  specific range of pE for the initiation of the
  reduction of the element.
• These half reactions are coupled with the
  oxidation of carbon. In this way, the redox
  elements act as electron acceptors in C
• oxidation.
• e- Acceptor pE Range Reactions
• O(-II) 5.0-11.0 Aerobic organisms. Oxidized forms
  of Fe, S, and Mn
• N(-III) 3.4-8.5 Nitrate respiration below pE 8.
  Products include NO2-, N2, N2O,
• or NH4. Denitrification produces N gases.
• Mn(IV) 3.4-6.8 Solid phase Mn reduction, can
  occur in the presence of NO3-
• Fe(III) 1.7-5.0 Solid phase Fe reduction, does
  not occur in the presence of NO3-
• or O2. Fe and Mn reduction characteristics of
  suboxic system.
• S(VIII) -2.5-0.0 Anaerobic organisms. Products
  include H2S, HS-, S2O3-
• Changes in pE induce both chemical reduction
  sequences as well as sequences in microbial
  ecology.
• Aerobic organisms that utilize O2 do not
  function below pE of 5
• Denitrifying bacteria function in the pE range
  of 10 to 0
• Sulfate-reducing bacteria live at pE's below 2

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