Redox Geochemistry

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

23. 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.

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. 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

26. 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

27. Fe3+ + e- Fe2+ E0= +0.77 • 2 H2O  4 H+ + O2 + 4 e- E0= +1.23 ----------------------------------------------------------- 4 Fe3+ + H2O  4 Fe2+ + 4 H+ + O2 +2.00 V 4 Fe2+ + 4 H+ + O2  4 Fe3+ + H2O -2.00 V Reaction direction!!