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
• 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
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
DG   ni G ( products)   ni G (reactants )
0
R
0
i
i
0
i
i
• Tabulated values of DG0R available…
Equilibrium Constant
 aCc aDd 
RT ln  a b   RT ln Q
 a A aB 
•
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
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
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
Activity Coefficients
• Extended Debye-Huckel approximation
(valid for I up to 0.5 M):
 log g 
Az 2 I
1
2
I  aBI
1
2
 0.2 I
• Where A and B are constants (tabulated),
and a is a measure of the effective
diameter of the ion (tabulated)
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++
Ca(H3SiO4)2
Ca(O-phth)
CaB(OH)4+
CaCH3COO+
CaCO30
CaCl+
CaF+
CaH2SiO4
CaH3SiO4+
CaHCO3+
CaNO3+
CaOH+
CaPO4CaSO4
CaHPO40
• Plus more species  gases and
minerals!!
Mass Action & Mass Balance
c
 n
[CL] [ H ]
i 
c
l
[C ] [ HL ]
mCa   mCa L
2
2 n
x
• 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
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…
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
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 eO2 + 4 e-  2 H2O
4 Fe2+ + O2  4 Fe3+ + 2 H2O
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
ELECTRON ACTIVITY
• Although no free electrons exist in solution, it is useful
to define a quantity called the electron activity:
pe   log ae
• 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.
THE pe OF A HALF REACTION - I
Consider the half reaction
MnO2(s) + 4H+ + 2e-  Mn2+ + 2H2O(l)
The equilibrium constant is
K
aMn2
4
H
a a
2
e
Solving for the electron activity
 aMn2
ae    4
 Ka 
 H




1
2
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:

pe 
Eh
2.303RT
Where  = 96.42 kJ volt-1 eq-1 (Faraday’s constant).
At 25°C, this becomes
pe  16.9 Eh
or
Eh  0.059 pe
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 acidbase pair
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:
 RT
0
EH  EH  
 nF
At 25° C:
a b
  a AaB 
 ln  c d 
  aC aD 
a b

 0.0592V   a A aB 
0
EH  EH  
 log  c d 
n

  aC aD 
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
• 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
-0.32
+0.81
NADPH + H+  NADP+ + 2H+ + 2eO2 + 4H+ + 4e-  2H2O
2 NADPH + O2 + 2H+  2 NADP+ + 2 H2O
+0.32
+0.81
+1.13
ELECTRON TOWER
more negative
more positive
oxidized/reduced forms
potential acceptor/donor
BOM – Figure 5.9
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
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
glucose
• transport of
electrons coupled
to pumping protons
CH2O  CO2 + 4 e- + H+
0.5 O2 + 4e- + 4H+  H2O
e-
Proton Motive Force (PMF)
• Enzymatic reactions pump H+ outside the
cell, there are a number of membranebound 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…
HOW IS THE PMF USED TO
SYNTHESIZE ATP?
• catalyzed by ATP
synthase
BOM – Figure 5.21
ATP generation II
• Alternative methods to form ATP:
• Phosphorylation  coupled to
fermentation, low yield of ATP
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!
Minimum Free Energy for growth
• Minimun free energy for growth = energy
to make ATP?
• What factors go into the energy budget of
an organism??
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.
The Redox ladder
O2
Oxic
Aerobes
H2O
NO3- Denitrifiers
Sub-oxic
anaerobic
N2
MnO2
Mn2+
Manganese reducers
Fe(OH)3
Fe2+
Sulfidic
Iron reducers
SO42H2S
Methanic
Sulfate reducers
CO2
CH4
Methanogens
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)