Download - Redox Geochemistry
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Redox Geochemistry
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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
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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.
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The Redox ladder
H2O
H2
O2
H2O
NO3-
N2 MnO2
Mn2+
Fe(OH)3
Fe2+SO4
2-
H2S CO2
CH4
Oxic
Post - oxic
Sulfidic
Methanic
Aerobes
Dinitrofiers
Maganese 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)
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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 + Ox2LEO says GER
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Half Reactions• Often split redox reactions in two:
– oxidation half rxn • Fe2+ Fe3+ + e-
– Reduction half rxn • O2 + 4 e- + 4 H+ 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
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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-
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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.
e
ape log
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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
24
2
eH
Mn
aa
aK
21
2
4
H
Mne Ka
aa
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THE pe OF A HALF REACTION - II
Taking the logarithm of both sides of the above equation and multiplying by -1 we obtain:
or
Ka
aa
H
Mne
logloglog 21
421
2
Ka
ape
H
Mn loglog 21
421
2
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THE pe OF A HALF REACTION - III
We can calculate K from:
so
65.43)15.298)(10314.8(303.2
))1.453()1.237(21.228(303.2
)2(
303.2log
3
222
RT
GGG
RT
GK
oMnOf
oOHf
o
Mnf
or
83.21log42
12
H
Mn
a
ape
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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.
02
o
ef
oHf
o
HfGGG
12
1
2
H
eH
p
aaK
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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 a H+ = 1 (pH = 0) and p H2 = 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.
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STANDARD HYDROGEN ELECTRODE
Platinumelectrode
a H + = 1
H = 1 atm2
½H2(g) H+ + e-
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ELECTROCHEMICAL CELL
Platinumelectrode
a H+ = 1
H = 1 atm2 VPlatinumelectrode
Salt B ridge
Fe 2+Fe 3+
½H2(g) H+ + e- Fe3+ + e- Fe2+
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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.
ELECTROCHEMICAL CELL
8.12log3
2
Fe
Fe
a
ape
1.148.1205.0log pe
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DEFINITION OF EhEh - 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
EhRT
pe303.2
Ehpe 9.16
peEh 059.0
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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
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FIELD APPARATUS FOR Eh MEASUREMENTS
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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.
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Figure 5-6 from Kehew (2001). Plot of Eh values computed from the Nernst equation vs. field-measured Eh values.
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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.
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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.
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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 Gr ??
Gr = nE or G0r = nE0
– 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
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Nernst EquationConsider 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
Qn
RTEEh log
303.20
100
3
4log8
0592.0
HNO
NH
aa
aEEh
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Let’s assume that the concentrations of NO3- and
NH4+ have been measured to be 10-5 M and
310-7 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
n
GE
or0
volts88.0)42.96)(8(
9.6790
E
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The Nernst equation now becomes
substituting the known concentrations (neglecting activity coefficients)
and
10
3
4log8
0592.088.0
HNO
NH
aa
aEh
volts521.01010
103log
8
0592.088.0 1055
7
Eh
81.8)521.0(9.169.16 Ehpe
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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!
Biology’s view upside down?
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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
20
21
2
1log
0592.0
HO apn
EEh
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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.
20
21
2
1log
0592.0
HO apn
EEh
20
21
2
1log
2
0592.0
HO ap
EEh
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As for the pe-pH diagram, we assume that pO2
= 1 atm. This results in
This yields a line with slope of -0.0592.
221
2log0296.023.1
HO apEh
pHpEh O 0592.0log0148.023.12
volts23.1)42.96)(2(
)1.237(00
n
GE r
pHEh 0592.023.1
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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
pHEh 0592.0
H
H
a
pEEh
21
2log1
0592.00
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O2/H2O
C2HO
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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
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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?
2
30 log
1
0592.0
Fe
H
a
aEEh
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• 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???)
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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)
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Redox titration II
• Let’s modify a bjerrum plot to reflect pe changes
Greg Mon Oct 25 2004
-4 -2 0 2 4 6 8 10 1250
60
70
80
90
100
pe
So
me
sp
eci
es
w/
SO
4-- (
um
ola
l) H2S(aq) SO4--