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 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

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

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 log

303.20

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

So

me

sp

eci

es

w/

SO

4-- (

um

ola

l) H2S(aq) SO4--

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

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)