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. J. Willard Gibbs • Gibbs realized that for a reaction, a certain amount of energy goes to an increase in entropy of a system. • G = H –TS or DG0R = DH0R – TDS0R • Gibbs Free Energy (G) is a state variable, measured in KJ/mol or Cal/mol • Tabulated values of DG0R available…

4. Equilibrium Constant • for aA + bB  cC + dD: • Restate the equation as: DGR = DG0R + RT ln Q • DGR= available metabolic energy (when negative = exergonic process as opposed to endergonic process for + energy) for a particular reaction whose components exist in a particular concentration

5. Activity • Activity, a, is the term which relates Gibbs Free Energy to chemical potential: mi-G0i = RT ln ai • Why is there now a correction term you might ask… • Has to do with how things mix together • Relates an ideal solution to a non-ideal solution

6. Ions in solution • Ions in solutions are obviously nonideal states! • Use activities (ai) to apply thermodynamics and law of mass action ai = gimi • The activity coefficient, gi, is found via some empirical foundations

7. Activity Coefficients • Extended Debye-Huckel approximation (valid for I up to 0.5 M): • Where A and B are constants (tabulated), and a is a measure of the effective diameter of the ion (tabulated)

8. Speciation • Any element exists in a solution, solid, or gas as 1 to n ions, molecules, or solids • Example: Ca2+ can exist in solution as: Ca++ CaCl+ CaNO3+ Ca(H3SiO4)2 CaF+ CaOH+ Ca(O-phth) CaH2SiO4 CaPO4- CaB(OH)4+ CaH3SiO4+ CaSO4 CaCH3COO+ CaHCO3+ CaHPO40 CaCO30 • Plus more species  gases and minerals!!

9. Mass Action & Mass Balance • mCa2+=mCa2++MCaCl+ + mCaCl20 + CaCL3- + CaHCO3+ + CaCO30 + CaF+ + CaSO40 + CaHSO4+ + CaOH+ +… • Final equation to solve the problem sees the mass action for each complex substituted into the mass balance equation

10. Geochemical models • Hundreds of equations solved iteratively for speciation, solve for DGR • All programs work on same concept for speciation thermodynamics and calculations of mineral equilibrium – lots of variation in output, specific info…

11. 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

12. 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

13. 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

14. 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.

15. 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

16. 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

17. 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 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

18. Electropotentials • E0 is standard electropotential, also standard reduction potential (write rxn as a reduction ½ rxn) – EH is relative to SHE (Std Hydrogen Electrode) At non-standard conditions: At 25° C:

19. 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E, 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

20. Redox reactions with more negative reduction potentials will donate electrons to redox reactions with more positive potentials. NADP+ + 2H+ + 2e- NADPH + H+ -0.32 O2 + 4H+ + 4e- 2H2O +0.81 NADPH + H+  NADP+ + 2H+ + 2e- +0.32 O2 + 4H+ + 4e- 2H2O +0.81 2 NADPH + O2 + 2H+ 2 NADP+ + 2 H2O +1.13

21. ELECTRON TOWER more negative oxidized/reduced forms potential acceptor/donor more positive BOM – Figure 5.9

23. Microbes, e- flow • Catabolism – breakdown of any compound for energy • Anabolism – consumption of that energy for biosynthesis • Transfer of e- facilitated by e- carriers, some bound to the membrane, some freely diffusible

24. NAD+/NADH and NADP+/NADPH • Oxidation-reduction reactions use NAD+ or FADH (nicotinamide adenine dinucleotide, flavin adenine dinucleotide). • When a metabolite is oxidized, NAD+ accepts two electrons plus a hydrogen ion (H+) and NADH results. NADH then carries energy to cell for other uses

25. glucose e- • transport of • electrons coupled • to pumping protons CH2O  CO2 + 4 e- + H+ 0.5 O2 + 4e- + 4H+ H2O

26. Proton Motive Force (PMF) • Enzymatic reactions pump H+ outside the cell, there are a number of membrane-bound enzymes which transfer e-s and pump H+ out of the cell • Develop a strong gradient of H+ across the membrane (remember this is 8 nm thick) • This gradient is CRITICAL to cell function because of how ATP is generated…

27. HOW IS THE PMF USED TO SYNTHESIZE ATP? • catalyzed by ATP synthase BOM – Figure 5.21

28. ATP generation II • Alternative methods to form ATP: • Phosphorylation  coupled to fermentation, low yield of ATP

29. ATP • Your book says ATP: “Drives thermodynamically unfavorable reactions”  BULLSHIT, this is impossible • The de-phosphorylation of ATP into ADP provides free energy to drive reactions!

30. Minimum Free Energy for growth • Minimun free energy for growth = energy to make ATP? • What factors go into the energy budget of an organism??

31. 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.

32. O2 Aerobes Oxic H2O Denitrifiers NO3- N2 Manganese reducers Sub-oxic anaerobic 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)