REDOX CHEMISTRY and Eh-pHDiagrams

As we have seen at various previous portions of this course thus far, theredox state of the system is important in understanding speciation,solubility, transport, degradation of organic and inorganic compounds inwater.

Redox reactions are central to understanding the chemistry ofgroundwater, including contaminated surface and groundwater systems.Although O2 is quite abundant and readily supplied in atmospheric andvery near-surface systems, its availability is not assured in deeper surfaceand groundwater systems. Why?

• low solubility of oxygen in water (this sets that maximum aqueouspO2)

• the low diffusivity of oxygen in water (takes time, without wholesalemixing, to bring to depths), and

• the consumption of oxygen by metabolic activities of aerobicorganisms (pO2 is determined by supply - consumption!)

Because of this low solubility (α 1/T) and transport of O2, redox gradientsor "fronts" are a fixture in near surface aqueous systems.

Examples:

Oceans and Lakes (seasonal effects, too)

Groundwater

Contaminant Plumes

OCEAN (high surface NPP, otherwise, shift O2 profile to the right):

GROUNDWATER/CONTAMINANT PLUME:

LAKES (dynamic O2 distribution depends on season and biology):

OTHERS (Flooding as a function of time):

Hence, the redox potential, which is the ability of a system (or a species) toconsume or donate electrons (note parallelism with pH - tendency toconsume or donate protons), is key to understanding metalchemistry/mobility in these systems.

How do we determine the relative tendencies of differentcompounds/metals to be reduced (donate electrons) or oxidized (loseelectrons)?

Recall the galvanic cell: function

• Connect two metal (Zn and Cu) with a conducting wire

• Submerge both in water

• Zn dissolves (LEO – Loss of Electron Oxidation-Anode):

Zn === Zn++ + 2e

• Cu ions reduced to Cu metal and deposited upon Cu electrode(reduction-cathode):

Cu++ + 2e === Cu

• Electrons produced at the Zn electrode from the oxidation of Zn to Znions flow through the external wire to the copper rod. At the Cu endthey are avail to combine with the incoming Cu++ (cupric ions) to formmore copper metal.

• A voltmeter can be introduced within the circuit. The voltage producedby the cell can be measured. Voltage called electromotive force. Tocomplete a porous contact allows flow of IONS back and forth.

• And it is conventional to tabulate voltage corresponding to unitACTIVITY for each ion (one molal).

• Note that for the reaction: Cu++ + 2e- === Cu

2 x Avogadro’s number (6.022 x 1023) of e- will be produced per mole of Cuproduced. The amount of electrical charge associated with one Avogadro’snumber worth of e- is 96,485.3 coulombs or a unit Faraday of charge.

(The electrical work required to move one charge F coulombs through apotential difference ε is w=Fε: joule = coulombs x volts. Hence, theelectrical work required to move n charge is: w=nFε).• Non-IUPAC (Geochemical) Convention ! ∆ Go = -nFEo and ∆ G = -nFE)

Standard Redox Potential - the tendency of a reducing agent to donateelectrons is given by a standard redox potential: Eo

• Definition of Standard Redox Potential - the resultant electromotiveforce (voltage) on a half cell in which oxidants and reductants arepresent in 1.0M concentrations at pH = 7 and 25 oC in equilibrium withelectrode that can reversibly accept electrons from the reductant..

• Reference electrode = hydrogen electrode consists of gaseous H2 at 1atm bubbling over a platinum electrode in an acidic soln with activity of1 for hydrogen ion (1 M HCl).

• Interpretation of redox potentials for REDOX pairs

1. The more negative the redox potential is relative to the Helectrode, the greater reducing potential of the system. The greaterthe tendency to lose electrons (LEO reduces it’s surrounding).

2. If more positive than reference electrode, there are lessertendencies to donate electrons. Rather there is a stronger affinity forelectrons and the element is a good oxidizing agent.

Relationship between Redox potential and Eh

Electrochemical energy is a form of Free Energy for Redox Reactions:

∆ Go = -nFEo and ∆ G = -nFE (non-IUPAC convention, but geochemical)

This expresses fact that EMF (Eo) multiplied by amount of electronic chargemoving through the cell (charge carried by a mole of electron, e, timesnumber of moles) gives maximum electrical work the cell can accomplish(think of galvanic cell).

Recall:

∆ G = ∆ Go - RT ln IAP

Because ∆ Go = -nFEo and ∆ G = -nFE:

Eh = Eo - (2.303RT/nF ) log ([product]/[reactant])

At STP:

Eh = Eo - (0.0592/n) log ([product]/[reactant])

R = 1.987 x 10-3 kcal/deg/mole, F = 23 kcal/volt /gram equivalent (AtomicWt/valence).

Application:

Consider the half-cell reaction:

Fe3+ + e- = Fe2+

∴ Eh = Eo - (0.059/n) log (aFe2+/aFe3+)

Eo = 0.77 volts

Eh = 0.77 volts - (0.059/1) log (aFe2+/aFe3+)

Eh when aFe3+ = aFe2+ is -0.77 volts.

Another example:

SO42- + 8e- + 10H+ = H2S + 4H2O

Eh = Eo - (0.059/8) log (aH2SaH2O4/ aSO4,2-aH+10)

Eo can be calculated from free energy or electron potential values inCompilations. In the above case: Eo = 0.30 V. At pH = 7, aH2S= aSO4-2

Eh = 0.30 V - (0.059/8) log (1/ 10-7) = 0.216 V

Eh-pH Diagrams (Pourbaix Diagram)

In the example above: Eh = 0.30 V - (0.059/8) log (aH2SaH2O4/ aSO42-aH+10)

IF aH2O= 1, we note:

Eh = 0.30 V - 10(0.059/8) pH + (0.059/8) log (aH2S/ aSO4)

NOW, the line of aSO4= = aH2S is not a specific Eh but a function also of pH

Eh = 0.30 V - 0.074 pH

At pH = 1, Eh = 0.23 V

At pH = 10, Eh = -0.44 V

And so on. Hence, if we plot Eh vs. pH, we have a line.

We could do the same to the H2O System:

4H+ + O2 + 4 e- = 2H2O

Eh = Eo - (0.059/4) log (aH2O2/PO2aH+4)

Or given: aH2O = 1, and maximum PO2 = 1 bar

Eh = Eo + (0.059) log (aH+) = Eo - 0.059 pH

For the above reaction, Eo = 1.23 V, hence:

Eh = 1.23 - 0.059 pH

Again, an equation of a straight line in a Pourbaix Diagram. This definesthe upper Eh-pH stability of water on the Earth’s surface. Because air has

PO2 = 0.3 bars, the upper stability of water in contact with air is (Eh = 1.23 -0.059 pH – 0.523)

The lower stability of water can be defined by the reaction:

2H+ + 2e- = H2

Eh = Eo - (0.059/2) log (PH2/aH+2)

Eh = Eo - (0.059) pH

Eo = 0

because this is the reference electrode reaction:

Eh = - 0.059 pH

Relationship between Redox potential, Eh, and pe

From the reaction:

Fe3+ + e- = Fe2+

Keq = [Fe2+]/([Fe3+][e-]

1/e- = Keq [Fe3+]/([Fe2+]

Or (by analogy to pH):

pe = log Keq + log ([Fe3+]/([Fe2+])

Hence (Eh = (RT/nF) ln Keq + (RT/nF) ln (IAP)): pe = (nF/2.303RT) Eh

pe = 13 + log (([Fe3+]/([Fe2+]) because (log Keq at 25 oC is 13)

pe-pH DIAGRAMS

(Caveat: redox equilibrium is not always attained in natural systems. Thefollowing is intended to show what the equilibrium relations are, giventhat systems tend towards these equilibrium values. In many systems, it isprecisely the availability of energy due to lack of equilibrium that is"catalyzed" by organisms leading to regular patterns of oxidation-reductionreactions).

Fe-O-H2O System:

Consider the species: O2, H2, H2O, Fe2+, Fe3+, Fe3O4, Fe2O3

The high pe limit of water :

0.5 O2 + 2e- + 2 H+ = H2O

log Keq = -0.5log PO2 + 2pe + 2pH = +41.55 (From Thermo Tables)

The limit may be defined as PO2 = 1 atm (hence log PO2 = 0)

Hence: pe = 20.78 - pH (show in a plot; line with slope of -1 and intercept of20.78)

Similarly, the lower pe limit is defined by:

H+ + e = 0.5 H2 (Standard state of formation of H2; Go = log Keq = 0)

The lower limit defined as PH2 = 1 atm (and log PH2 = 0)

Hence: pe = -pH (line with slope of -1)

Recall the Fe3+/Fe2+ reaction earlier (pe = log Keq + log ([Fe3+]/([Fe2+])):

pe = 13 + log (([Fe3+]/([Fe2+])

The line of equal concentration of Fe3+ and Fe2+ is at pe = 13 (horizontalline)

Boundary Between Solids: Fe(OH)2, Fe(OH)3

Fe(OH)3 + e- + H+ = Fe(OH)2 + H2O

pe = log Keq - pH = 3.92 - pH

Boundary Between Solid and Dissolved Ions: Fe2+, Fe(OH)3

Fe(OH)3 + e- + 3H+ = Fe2+ + 3H2O

pe = log Keq - log Fe2+ - 3 pH

Given Fe2+ = 1x10-5 and Keq = 17.2

pe = 22.2 - 3 pH

Other Boundaries are Derived Similarly:

Figure 7-4 of Drever.

Also see Figure 7-5 (Drever) with stability relations using ferrihydrite,instead of hematite, as the solid.

Note toxicological and transport significance!

Thus far, all hydrated ions. The same can be extended with other ligandspresent (e.g. CO2, Sulfur). The same reasoning controls the speciation ofcomplexes that could be present (in terms of the relative dominance ofcomplexes).

NATURAL WATERS: pe-pH Ranges, and Significance

NATURAL WATERS:

Natural water pe-pH ranges are shown in Figure 7.3 (Schlesinger). Thisrange by and large limits the speciation of metals and other aqueousspecies (including complexes) in natural systems (waste water systems andother special systems would be another matter).

" Redox reactions are central to understanding the biogeochemistry oflakes. In lakes and wetlands, oxygen depletion in deeper systems arisesfrom the low solubility of oxygen in water, low diffusivity of oxygen inwater, and consumption of oxygen by metabolic activities of organisms.

" Redox gradients and processes extend beneath thewater/sediment interface.

" In lakes and wetlands, the important species to look out for amongothers are:

(1) O2 (aeration, deaeration)

(2) nutrient N distribution (control on variation with depth and time(NH4

+, NO2-, NO3

-)

(3) sulfur, availability (like N) controls decomposition of organicsand stability of microbes (desulfuvibrio, etc.) H2S, HS-, SO4

(4) CO2, HCO3-, CO3

=, CH4 (also microbial applications)

" Microbial mediation make it possible to accomplish these reactionseven at low temperatures, but obviously, these processes are favored byhigher temperatures.

" The energetically/thermodynamically most favorable reactions usuallypreempts competing reactions because use of common electron denors(organic matter). Because microbial species specializes on specificreactions, they distribute themselves in a manner dictated by thelocation of the specific reactions they want to mediate.

" Redox gradients are (of course) also present in marine sediments, butrelative importance of specific reactions are different (e.g., the sulfate

reducers versus methanogens - compete for organic electron donors).

" 1/4CH2O + 1/4 O2 = 1/4CO2 + 1/4H2O

∆G = -29.9 kcal/mol (aerobic respiration)

" 1/4CH2O + 1/5 NO3- +1/5 H+ = 1/4CO2 +1/10 N2 + 7/10H2O

∆G = -28.4 kcal/mol (denitrifying bacteria)

(-facultative anaerobes zone)

Mn4+ reduction

(¯ obligate anaerobes zone, e.g. Clostridium)

Fe3+ reduction

" 1/4CH2O + 1/8 SO42- +1/8 H+ = 1/4CO2 +1/8 HS- + 1/4H2O

∆G = - 6.1 kcal/mol

(Desulfovibrio, Desulfotomaculum)

" 1/4CH2O = 1/8CO2 +1/8 CH4

∆G = - 5.6 kcal/mol (dissimilators)

" 1/8CO2+1/2 H2 = 1/8CH4 + 1/4H2O

∆G = -4 kcal/mol (reductors)

"Hydrogen Fermentation"

" 1/4CH2O + 1/4 H2O = 1/4CO2 +1/2 H2

∆G = - 1.6 kcal/mol

NITRATE REDUCTION

SULPHATE REDUCTION

" 1/4CH2O + 1/8 SO42- +1/8 H+ = 1/4CO2 +1/8 HS- + 1/4H2O

∆G = - 6.1 kcal/mol

or (at lower pH)

1/4CH2O + 1/8 SO42- +1/4 H+ = 1/4CO2 +1/8 H2S 1/4H2O

(Desulfovibrio, Desulfotomaculum)

" The same microbes responsible for sulphate produce other importantgases such as dimethyl sulfide and dimethyl disulfide (in an analogousreaction SeO4 is reduced to elemental Se rendering Se more toxic towildlife; H2S is also toxic). This sulfur flux was thought to dominatesource of S in the atmosphere during pre-industrial times (with H2S>> methyl sulfides).

" H2S generated may be re-oxidized by sulfur bacteria as itdiffuse to overlying water column, react with organic matter toform organically-bound S (vulcanization; to 80 % of S in peat),or --

" H2S may react with Fe2+ to form sulfide minerals (this cycle is linkedwith the P cycle because FeS precipitated means less Fe to precipitateor sorb phosphorus (see lecture on NPP and Nutrients). Steelstructures through wetlands/lakes are likewise "attacked" by H2S toform FeS.

" δ34S is useful in examining which S has been through thedissimilatory sulphate pathway (depleted in 34S). Even pyrite mayutilize S that has been through this pathway instead of normalprecipitation routes (can be examined in ores, too).

l These pathways can be traced all the way back to the atmosphere (32Sdepleted SO4 may be derived from oxidation of wetland-reduced Sgases.

METHANOGENESIS

" 1/4CH2O = 1/8CO2 +1/8 CH4 (CH3COOH is preferred "CH2O"substrate)

∆G = - 5.6 kcal/mol (acetate dissimilation)

1/8CO2+1/2 H2 = 1/8CH4 + 1/4H2O

∆G = -4 kcal/mol (CO2 reduction)

Hydrogen Fermentation provides the H2 for CO2 reducers:

1/4CH2O + 1/4 H2O = 1/4CO2 +1/2 H2

∆G = - 1.6 kcal/mol

" In marine environments, methanogenesis is limited by the abundanceof sulphate, which competes formidably for acetate and H2.Methanogenesis is thus observed only in regimes where SO4 is low tobegin with or depleted to a significant degree by sulphate reduction(not much overlap between sulphate reducers and methanogens insediments).

" Furthermore, even when methanogenesis takes place in marinesystems, they tend to utilize CO2 reduction pathways because acetateis complete utilized by sulphate reducing bacteria.

" In terrestrial systems, especially wetlands, methane production byeither mechanisms is not inhibited (except in very saline marshes).

" What controls CH4 flux, however, is not production rate but oxidationduring ascent. Observed dependence of flux on moisture content ofbogs and tundra is related to efficient oxidation (10 to 100%) ofasccending methane during dry seasons. The physics of bubbleformation and ascent, by-pass provided by hollow plant stems are alsosignificant factors.

" Rice paddies account for about half of wetland emission of methane.

" In lakes, methane oxidation is also key. Much of the oxidation cantake place in overlying sediments (to 50% in Lake Washington), butin reduced (stratified systems) much of the oxidation occur in thewater column.

" On a global scale, as much as 2% of NPP in lakes eventually end up

as CH4 in the atmosphere (that is not a trivial amount, especiallyconsidering isotopically peculiar properties of CH4 derived bymethanogenesis (below).

" Acetate dissimilation (∆13C = -50 to -65 per mil) and CO2 reduction(∆13C = -60 to -100 per mil) are distinguishable pathways of methaneformation (recall Whiticar's talk).

" (bio)Methylation is another process that operates in very reducedenvironments of wetlands and lakes. Bacteria, again, utilize "freeenergy" from reactions such as:

Hg2+ + RCH3 = HgCH3+ + R+ or Hg2+ + 2R-CH3 = Hg(CH3)2 + 2R+

The concern with biomethylation is that for Hg and other metals, themethyl form is either more easily assimilated or dramatically moretoxic than the non-methylated species. James Bay, Palawan, goldpanning.

" GLOBAL IMPORTANCE OF REDOX PROCESSES

When one realize that 1/4 of total organic storage of the world are inwetland and tundra soils, it is easier to appreciate that the fate of thisstored carbon has a lot of say in global-scale balance.

GEOCHEMICAL WORKBENCH:

Please visit ACT2 beginning with p. 37 of the User'sGuide. Show Eh-pH speciation for URANIUM-H2Osystem.