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 Plus more species  gases and minerals!!
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) CaF CaOH+
Ca(O-phth) CaH2SiO CaPO4-
CaB(OH) CaH3SiO CaSO4
CaCH3COO CaHCO 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 Fe e-
4 Fe2+ + O2  4 Fe 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.
A defined above, the pe is a measure of the oxidation-reduction capacity of a natural water. It is analogous to pH = -log aH+. A natural water with a high pe (low electron activity) would be considered to be oxidizing, and a water with a low pe (high electron activity) would be considered to be reducing. Just as pH is important to mineral solubilities and the speciation of acid-base pairs, pe is important to the solubilities of minerals containing elements with variable oxidation states, and the speciation of redox pairs.

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
Calculation of the pe of a half reaction provides us with another way to determine which way a reaction will go. Consider the above half cell reaction in which pyrolusite is reduced to Mn2+. We start by writing the mass-action expression for this half reaction. We then rearrange the mass-action expression so as to get the electron activity on the left side and everything else on the right side.

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  = kJ volt-1 eq-1 (Faraday’s constant).
At 25°C, this becomes or
As mentioned before, Eh is the potential of a solution relative to the SHE. This slide shows that there is a fairly simple relationship between pe and Eh. The symbol “Eh” comes from the fact that “E” is the normal symbol for a potential (or electromotive force), and the lower case “h” reminds us that the reference electrode is the SHE.
Eh and pe measure the same thing. High values of Eh or pe correspond to oxidizing conditions, and low values of Eh or pe correspond to reducing conditions.

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
O2 + 4H+ + 4e-  2H2O
NADPH + H+  NADP+ + 2H+ + 2e
O2 + 4H+ + 4e-  2H2O
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
Oxic waters - 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.
Barcelona and Holm (1991) presented the classification above for natural waters with respect to Eh.

32. The Redox ladder
H2O H2 O2
NO3- N2
MnO2
Mn2+
Fe(OH)3
Fe2+
SO42-
H2S
CO2
CH4
Oxic
Sub-oxic
anaerobic
Sulfidic
Methanic
Aerobes
Denitrifiers
Manganese reducers
Sulfate reducers
Methanogens
Iron reducers
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)