Redox geochemistry
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Redox Geochemistry. 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

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Redox Geochemistry

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Redox geochemistry

Redox Geochemistry


Redox geochemistry

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


Redox classification of natural waters

REDOX CLASSIFICATION OF NATURAL WATERS

Oxicwaters - waters that contain measurable dissolved oxygen.

Suboxic waters - waters that lack measurable oxygen or sulfide, but do contain significant dissolved iron (> ~0.1 mg L-1).

Reducing waters(anoxic) - waters that contain both dissolved iron and sulfide.


The redox ladder

O2

Aerobes

Oxic

H2O

Dinitrofiers

NO3-

N2

Maganese reducers

Post - oxic

MnO2

Mn2+

Iron reducers

Fe(OH)3

Fe2+

SO42-

Sulfate reducers

Sulfidic

H2S

CO2

Methanogens

CH4

Methanic

H2O

H2

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. (Gibb’s free energy becomes more positive going down the steps)


Oxidation reduction reactions

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


Half reactions

Half Reactions

  • Often split redox reactions in two:

    • oxidation half rxn 

      • Fe2+ Fe3+ + e-

    • Reduction half rxn 

      • O2 + 4 e- + 4H+  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


Redox couples

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-


Electron activity

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 - the solution is reducing.

  • If pe is high - the solution is oxidizing.


The pe of a half reaction i

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


The pe of a half reaction ii

THE pe OF A HALF REACTION - II

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

or


The pe of a half reaction iii

THE pe OF A HALF REACTION - III

We can calculate K from:

so


We need a reference point

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.


The standard hydrogen electrode

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 aH+ = 1 (pH = 0) and pH2 = 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.


Standard hydrogen electrode

STANDARD HYDROGEN ELECTRODE

½H2(g)  H+ + e-


Electrochemical cell

ELECTROCHEMICAL CELL

Fe3++ e- Fe2+

½H2(g)  H+ + e-


Electrochemical cell1

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.


Definition of eh

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 (Faraday’s constant).

At 25°C, this becomes

or


Eh measurement and meaning

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+/Fe3+=1, 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


Field apparatus for eh measurements

FIELD APPARATUS FOR Eh MEASUREMENTS


Calibration of electrodes

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.

  • Zobell’s solution - solution of potassium ferric-ferro cyanide of known Eh.


Redox geochemistry

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


Problems with eh measurements

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.


Other methods of determining the redox state of natural systems

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.


Free energy and electropotential

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 Faraday’s 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


Nernst equation

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 25°C is


Redox geochemistry

Let’s 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


Redox geochemistry

The Nernst equation now becomes

substituting the known concentrations (neglecting activity coefficients)

and


Redox geochemistry

Biology’s 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!


Stability limits of water

Stability Limits of Water

  • H2O  2 H+ + ½ O2(g) + 2e-

    Using the Nernst Equation:

  • Must assign 1 value to plot in x-y space (PO2)

  • Then define a line in pH – Eh space


Upper stability limit of water eh ph

UPPER STABILITY LIMIT OF WATER (Eh-pH)

To determine the upper limit on an Eh-pH diagram, we start with the same reaction

1/2O2(g) + 2e- + 2H+ H2O

but now we employ the Nernst eq.


Redox geochemistry

As for the pe-pH diagram, we assume that pO2 = 1 atm. This results in

This yields a line with slope of -0.0592.


Lower stability limit of water eh ph

LOWER STABILITY LIMIT OF WATER (Eh-pH)

Starting with

H+ + e- 1/2H2(g)

we write the Nernst equation

We set pH2 = 1 atm. Also, Gr° = 0, so E0 = 0. Thus, we have


Redox geochemistry

O2/H2O

C2HO


Making stability diagrams

Making stability diagrams

  • For any reaction we wish to consider, we can write a mass action equation for that reaction

  • We make 2-axis diagrams to represent how several reactions change with respect to 2 variables (the axes)

  • Common examples: Eh-pH, PO2-pH, T-[x], [x]-[y], [x]/[y]-[z], etc


Construction of these diagrams

Construction of these diagrams

  • For selected reactions:

    Fe2+ + 2 H2O  FeOOH + e- + 3 H+

    How would we describe this reaction on a 2-D diagram? What would we need to define or assume?


Redox geochemistry

  • How about:

  • Fe3+ + 2 H2O  FeOOH(ferrihydrite) + 3 H+

    Ksp=[H+]3/[Fe3+]

    log K=3 pH – log[Fe3+]

    How would one put this on an Eh-pH diagram, could it go into any other type of diagram (what other factors affect this equilibrium description???)


Redox titrations

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)


Redox titration ii

Redox titration II

  • Let’s modify a bjerrum plot to reflect pe changes


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