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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

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Redox Geochemistry

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  1. Redox Geochemistry

  2. 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

  3. 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.

  4. 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)

  5. 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

  6. 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

  7. 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-

  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 - the solution is reducing. • If pe is high - 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. THE pe OF A HALF REACTION - II Taking the logarithm of both sides of the above equation and multiplying by -1 we obtain: or

  11. THE pe OF A HALF REACTION - III We can calculate K from: so

  12. 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.

  13. 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.

  14. STANDARD HYDROGEN ELECTRODE ½H2(g)  H+ + e-

  15. ELECTROCHEMICAL CELL Fe3++ e- Fe2+ ½H2(g)  H+ + e-

  16. 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.

  17. 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

  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+/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

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

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

  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. 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.

  24. 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

  25. 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

  26. 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

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

  28. 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!

  29. 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

  30. 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.

  31. As for the pe-pH diagram, we assume that pO2 = 1 atm. This results in This yields a line with slope of -0.0592.

  32. 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

  33. O2/H2O C2HO

  34. 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

  35. 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?

  36. 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???)

  37. 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)

  38. Redox titration II • Let’s modify a bjerrum plot to reflect pe changes

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