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Redox Geochemistry

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- 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

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.

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-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 - 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

- Often split redox reactions in two:
- oxidation half rxn
- Fe2+ Fe3+ + e-

- Reduction half rxn
- O2 + 4 e- + 4H+ 2 H2O

- oxidation half rxn
- 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

- 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-

- 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.

Consider the half reaction

MnO2(s) + 4H+ + 2e- Mn2+ + 2H2O(l)

The equilibrium constant is

Solving for the electron activity

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

or

We can calculate K from:

so

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.

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.

½H2(g) H+ + e-

Fe3++ e- Fe2+

½H2(g) H+ + e-

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.

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 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

- 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.

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

- 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.

- Many species are not electroactive (do NOT react electrode)

- 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.

- 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 = nDE or DG0r = nDE0

- 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

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

Let’s assume that the concentrations of NO3- and NH4+ have been measured to be 10-5 M and 310-7 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

The Nernst equation now becomes

substituting the known concentrations (neglecting activity coefficients)

and

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!

- 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

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.

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

This yields a line with slope of -0.0592.

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

O2/H2O

C2HO

- 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

- 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?

- 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???)

- 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)

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