Redox Geochemistry
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Transcript of Redox Geochemistry
Redox Geochemistry
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
Fundamental electromagnetic relations:• Electric charge (q) is measured in coulombs (C).
– The magnitude of the charge of a single electron is 1.602 x 10-19 C. 1 mole of electrons has a charge of 9.649 x 104 C which is called the Faraday constant (F)
– q=n*F• The quantity of charge flowing each second through a circuit is called the
current (i). The unit of current is the ampere (A) 1 A = 1 C/sec • The difference in electric potential (E) between two points is a measure
of the work that is needed when an electric charge moves from one point to another. Potential difference is measured in volts (V) 1 V = 1 J/C – The greater the potential difference between two points, the stronger
will be the "push" on a charged particle traveling between those points. A 12 V battery will push electrons through a circuit 8 times harder than a 1.5 V battery.
• Ohm’s Law: V = I R potential is equal to current * resistance
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
Examples
Balance these and write the half reactions:
• Mn(IV) + H2S Mn2+ + S0 + H+
• CH2O + O2 CO2 + H2O
• H2S + O2 S8 + 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-
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:
• 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.
e
ape log
THE pe OF A HALF REACTION - I
Consider the half reactionMnO2(s) + 4H+ + 2e- Mn2+ + 2H2O(l)
The equilibrium constant is
Solving for the electron activity
24
2
eH
Mn
aaa
K
21
2
4
H
Mne Ka
aa
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
efo
Hfo
HfGGG
12
1
2
H
eH
paa
K
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
aa
pe
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
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– 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 Equation
Consider 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 log303.20
100
3
4log8
0592.0
HNO
NH
aa
aEEh
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.
CONVERTING ELECTRODE READING TO Eh
Once a stable potential has been obtained, the reading can be converted to Eh using the equation
Ehsys = Eobs + EhZobell - EhZobell-observed
Ehsys = the Eh of the water sample.
Eobs = the measured potential of the water sample relative to the reference electrode.
EhZobell = the theoretical Eh of the Zobell solution
EhZobell = 0.428 - 0.0022 (t - 25)
EhZobell-observed = the measured potential of the Zobell solution relative to the reference electrode.
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.
Figure 5-6 from Kehew (2001). Plot of Eh values computed from the Nernst equation vs. field-measured Eh values.
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.
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.
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
Som
e sp
ecie
s w
/ S
O4-- (
umol
al) H2S(aq) SO4
--
The Redox ladder
H2O
H2
O2
H2ONO3
-
N2 MnO2
Mn2+
Fe(OH)3
Fe2+SO4
2-
H2S CO2
CH4
Oxic
Post - oxic
Sulfidic
Methanic
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)