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Redox Geochemistry

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H2O H2 O2 H2O NO3- N2 MnO2 Mn2+ Fe(OH)3 Fe2+ SO42- H2S CO2 CH4 Oxic Post - oxic Sulfidic Methanic The redox-couples are shown on each stair-step, ... – PowerPoint PPT presentation

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Title: Redox Geochemistry


1
Redox Geochemistry
2
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
3
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

4
Half Reactions
  • Often split redox reactions in two
  • oxidation half rxn ? e- leaves left, goes right
  • Fe2 ? Fe3 e-
  • Reduction half rxn ? e- leaves left, goes right
  • O2 4 e- ? 2 H2O
  • SUM of the half reactions yields the total redox
    reaction
  • 4 Fe2 ? 4 Fe3 4 e-
  • O2 4 e- ? 2 H2O
  • 4 Fe2 O2 ? 4 Fe3 2 H2O

5
Examples
  • Balance these and write the half reactions
  • Mn(IV) H2S ? Mn2 S0 H
  • CH2O O2 ? CO2 H2O
  • H2S O2 ? S8 H2O

6
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-

7
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

8
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.

9
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

10
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.

11
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.

12
STANDARD HYDROGEN ELECTRODE
½H2(g) ? H e-
13
ELECTROCHEMICAL CELL
Fe3 e- ? Fe2
½H2(g) ? H e-
14
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.

15
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

16
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?E
  • 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

17
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

18
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

19
FIELD APPARATUS FOR Eh MEASUREMENTS
20
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.

21
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.

22
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.

23
Figure 5-6 from Kehew (2001). Plot of Eh values
computed from the Nernst equation vs.
field-measured Eh values.
24
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.

25
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.

26
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)

27
Redox titration II
  • Lets modify a bjerrum plot to reflect pe changes

28
The Redox ladder
O2
Oxic
H2O
NO3-
N2
MnO2
Post - oxic
Mn2
Fe(OH)3
Fe2
SO42-
Sulfidic
H2S
CO2
CH4
H2O
Methanic
H2
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)
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