Redox Geochemistry

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Redox Geochemistry
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WHY?

• Redox gradients drive life processes!
• The transfer of electrons between oxidants and
  reactants is harnessed as the battery, the source
  of metabolic energy for organisms
• Metal mobility ? redox state of metals and
  ligands that may complex them is the critical
  factor in the solubility of many metals
• Contaminant transport
• Ore deposit formation

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Oxidation Reduction Reactions

• Oxidation - a process involving loss of
  electrons.
• Reduction - a process involving gain of
  electrons.
• Reductant - a species that loses electrons.
• Oxidant - a species that gains electrons.
• Free electrons do not exist in solution. Any
  electron lost from one species in solution must
  be immediately gained by another.
• Ox1 Red2 ? Red1 Ox2

LEO says GER
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Fundamental electromagnetic relations

• Electric charge (q) is measured in coulombs (C).
• The magnitude of the charge of a single electron
  is 1.602 x 10-19 C. 1 mole of electrons has a
  charge of 9.649 x 104 C which is called the
  Faraday constant (F)
• qnF
• The quantity of charge flowing each second
  through a circuit is called the current (i). The
  unit of current is the ampere (A) 1 A 1 C/sec
• The difference in electric potential (E) between
  two points is a measure of the work that is
  needed when an electric charge moves from one
  point to another. Potential difference is
  measured in volts (V) 1 V 1 J/C
• The greater the potential difference between two
  points, the stronger will be the "push" on a
  charged particle traveling between those points.
  A 12 V battery will push electrons through a
  circuit 8 times harder than a 1.5 V battery.
• Ohms Law V I R ? potential is equal to
  current resistance

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Half Reactions

• Often split redox reactions in two
• oxidation half rxn ?
• Fe2 ? Fe3 e-
• Reduction half rxn ?
• O2 4 e- 4 H ? 2 H2O
• SUM of the half reactions yields the total redox
  reaction
• 4 Fe2 ? 4 Fe3 4 e-
• O2 4 e- 4 H? 2 H2O
• 4 Fe2 O2 4 H ? 4 Fe3 2 H2O

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Redox Couples

• For any half reaction, the oxidized/reduced pair
  is the redox couple
• Fe2 ? Fe3 e-
• Couple Fe2/Fe3
• H2S 4 H2O ? SO42- 10 H 8 e-
• Couple H2S/SO42-

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Half-reaction vocabulary part II

• Anodic Reaction an oxidation reaction
• Cathodic Reaction a reduction reaction
• Relates the direction of the half reaction
• A ? A e- anodic
• B e- ? B- cathodic

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ELECTRON ACTIVITY

• Although no free electrons exist in solution, it
  is useful to define a quantity called the
  electron activity
• The pe indicates the tendency of a solution to
  donate or accept a proton.
• If pe is low, there is a strong tendency for the
  solution to donate protons - the solution is
  reducing.
• If pe is high, there is a strong tendency for the
  solution to accept protons - the solution is
  oxidizing.

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THE pe OF A HALF REACTION - I

• Consider the half reaction
• MnO2(s) 4H 2e- ? Mn2 2H2O(l)
• The equilibrium constant is
• Solving for the electron activity

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THE pe OF A HALF REACTION - II

• Taking the logarithm of both sides of the above
  equation and multiplying by -1 we obtain
• or

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THE pe OF A HALF REACTION - III

• We can calculate K from
• so

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WE NEED A REFERENCE POINT!

• Values of pe are meaningless without a point of
  reference with which to compare. Such a point is
  provided by the following reaction
• ½H2(g) ? H e-
• By convention
• so K 1.

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THE STANDARD HYDROGEN ELECTRODE

• If a cell were set up in the laboratory based on
  the half reaction
• ½H2(g) ? H e-
• and the conditions a H 1 (pH 0) and p H2
  1, it would be called the standard hydrogen
  electrode (SHE).
• If conditions are constant in the SHE, no
  reaction occurs, but if we connect it to another
  cell containing a different solution, electrons
  may flow and a reaction may occur.

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STANDARD HYDROGEN ELECTRODE
½H2(g) ? H e-
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ELECTROCHEMICAL CELL
Fe3 e- ? Fe2
½H2(g) ? H e-
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ELECTROCHEMICAL CELL

• We can calculate the pe of the cell on the right
  with respect to SHE using
• If the activities of both iron species are equal,
  pe 12.8. If a Fe2/a Fe3 0.05, then
• The electrochemical cell shown gives us a method
  of measuring the redox potential of an unknown
  solution vs. SHE.

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DEFINITION OF Eh

• Eh - the potential of a solution relative to the
  SHE.
• Both pe and Eh measure essentially the same
  thing. They may be converted via the
  relationship
• Where ? 96.42 kJ volt-1 eq-1 (Faradays
  constant).
• At 25C, this becomes
• or

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Eh Measurement and meaning

• Eh is the driving force for a redox reaction
• No exposed live wires in natural systems
  (usually) ? where does Eh come from?
• From Nernst ? redox couples exist at some Eh
  (Fe2/Fe31, Eh 0.77V)
• When two redox species (like Fe2 and O2) come
  together, they should react towards equilibrium
• Total Eh of a solution is measure of that
  equilibrium

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FIELD APPARATUS FOR Eh MEASUREMENTS
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CALIBRATION OF ELECTRODES

• The indicator electrode is usually platinum.
• In practice, the SHE is not a convenient field
  reference electrode.
• More convenient reference electrodes include
  saturated calomel (SCE - mercury in mercurous
  chloride solution) or silver-silver chloride
  electrodes.
• A standard solution is employed to calibrate the
  electrode.
• Zobells solution - solution of potassium
  ferric-ferro cyanide of known Eh.

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CONVERTING ELECTRODE READING TO Eh

• Once a stable potential has been obtained, the
  reading can be converted to Eh using the equation
• Ehsys Eobs EhZobell - EhZobell-observed
• Ehsys the Eh of the water sample.
• Eobs the measured potential of the water sample
  relative to the reference electrode.
• EhZobell the theoretical Eh of the Zobell
  solution
• EhZobell 0.428 - 0.0022 (t - 25)
• EhZobell-observed the measured potential of the
  Zobell solution relative to the reference
  electrode.

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PROBLEMS WITH Eh MEASUREMENTS

• Natural waters contain many redox couples NOT at
  equilibrium it is not always clear to which
  couple (if any) the Eh electrode is responding.
• Eh values calculated from redox couples often do
  not correlate with each other or directly
  measured Eh values.
• Eh can change during sampling and measurement if
  caution is not exercised.
• Electrode material (Pt usually used, others also
  used)
• Many species are not electroactive (do NOT react
  electrode)
• Many species of O, N, C, As, Se, and S are not
  electroactive at Pt
• electrode can become poisoned by sulfide, etc.

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Figure 5-6 from Kehew (2001). Plot of Eh values
computed from the Nernst equation vs.
field-measured Eh values.
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Other methods of determining the redox state of
natural systems

• For some, we can directly measure the redox
  couple (such as Fe2 and Fe3)
• Techniques to directly measure redox SPECIES
• Amperometry (ion specific electrodes)
• Voltammetry
• Chromatography
• Spectrophotometry/ colorimetry
• EPR, NMR
• Synchrotron based XANES, EXAFS, etc.

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Free Energy and Electropotential

• Talked about electropotential (aka emf, Eh) ?
  driving force for e- transfer
• How does this relate to driving force for any
  reaction defined by DGr ??
• DGr n?DE or DG0r n?DE0
• Where n is the of e-s in the rxn, ? is
  Faradays constant (23.06 cal V-1), and E is
  electropotential (V)
• pe for an electron transfer between a redox
  couple analagous to pK between conjugate
  acid-base pair

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Electromotive Series

• When we put two redox species together, they will
  react towards equilibrium, i.e., e- will move ?
  which ones move electrons from others better is
  the electromotive series
• Measurement of this is through the
  electropotential for half-reactions of any redox
  couple (like Fe2 and Fe3)
• Because DGr n?DE, combining two half reactions
  in a certain way will yield either a or
  electropotential (additive, remember to switch
  sign when reversing a rxn)
• -E ? - DGr, therefore ? spontaneous
• In order of decreasing strength as a reducing
  agent ? strong reducing agents are better e-
  donors

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Biologys view ? upside down?
Reaction directions for 2 different redox couples
brought together?? More negative potential ?
reductant // More positive potential ? oxidant
Example O2/H2O vs. Fe3/Fe2 ? O2 oxidizes
Fe2 is spontaneous!
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Nernst Equation

• Consider the half reaction
• NO3- 10H 8e- ? NH4 3H2O(l)
• We can calculate the Eh if the activities of H,
  NO3-, and NH4 are known. The general Nernst
  equation is
• The Nernst equation for this reaction at 25C is

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• Lets assume that the concentrations of NO3- and
  NH4 have been measured to be 10-5 M and 3?10-7
  M, respectively, and pH 5. What are the Eh and
  pe of this water?
• First, we must make use of the relationship
• For the reaction of interest
• ?rG 3(-237.1) (-79.4) - (-110.8)
• -679.9 kJ mol-1

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• The Nernst equation now becomes
• substituting the known concentrations (neglecting
  activity coefficients)
• and

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REDOX CLASSIFICATION OF NATURAL WATERS

• Oxic waters - waters that contain measurable
  dissolved oxygen.
• Suboxic waters - waters that lack measurable
  oxygen or sulfide, but do contain significant
  dissolved iron (gt 0.1 mg L-1).
• Reducing waters (anoxic) - waters that contain
  both dissolved iron and sulfide.

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Redox titrations

• Imagine an oxic water being reduced to become an
  anoxic water
• We can change the Eh of a solution by adding
  reductant or oxidant just like we can change pH
  by adding an acid or base
• Just as pK determined which conjugate acid-base
  pair would buffer pH, pe determines what redox
  pair will buffer Eh (and thus be reduced/oxidized
  themselves)

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Redox titration II

• Lets modify a bjerrum plot to reflect pe changes

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Redox titrations in complex solutions

• For redox couples not directly related, there is
  a ladder of changing activity
• Couple with highest potential reduced first,
  oxidized last
• Thermodynamics drives this!!

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The Redox ladder
The redox-couples are shown on each stair-step,
where the most energy is gained at the top step
and the least at the bottom step. (Gibbs free
energy becomes more positive going down the
steps)