redox geochemistry
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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 - PowerPoint PPT PresentationTRANSCRIPT
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
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
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)
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
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
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-
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
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
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
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
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
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.
STANDARD HYDROGEN ELECTRODE
Platinumelectrode
a H + = 1
H = 1 atm2
½H2(g) H+ + e-
ELECTROCHEMICAL CELL
Platinumelectrode
a H+ = 1
H = 1 atm2 VPlatinumelectrode
Salt B ridge
Fe 2+Fe 3+
½H2(g) H+ + e- Fe3+ + e- Fe2+
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
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
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
FIELD APPARATUS FOR Eh MEASUREMENTS
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.
Figure 5-6 from Kehew (2001). Plot of Eh values computed from the Nernst equation vs. field-measured Eh values.
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.
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.
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
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
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
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
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?
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
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
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
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
O2/H2O
C2HO
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
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
• 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???)
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)
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--