REDOX CHEMISTRY and Eh-pH Diagrams

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Terence Ronald Franklin
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1 REDOX CHEMISTRY and Eh-pH Diagrams As we have seen at various previous portions of this course thus far, the redox state of the system is important in understanding speciation, solubility, transport, degradation of organic and inorganic compounds in water. Redox reactions are central to understanding the chemistry of groundwater, including contaminated surface and groundwater systems. Although O 2 is quite abundant and readily supplied in atmospheric and very near-surface systems, its availability is not assured in deeper surface and groundwater systems. Why? low solubility of oxygen in water (this sets that maximum aqueous po 2 ) the low diffusivity of oxygen in water (takes time, without wholesale mixing, to bring to depths), and the consumption of oxygen by metabolic activities of aerobic organisms (po 2 is determined by supply - consumption!) Because of this low solubility (α 1/T) and transport of O 2, redox gradients or "fronts" are a fixture in near surface aqueous systems. Examples: Oceans and Lakes (seasonal effects, too) Groundwater Contaminant Plumes

2 OCEAN (high surface NPP, otherwise, shift O 2 profile to the right): GROUNDWATER/CONTAMINANT PLUME:

3 LAKES (dynamic O2 distribution depends on season and biology): OTHERS (Flooding as a function of time):

4 Hence, the redox potential, which is the ability of a system (or a species) to consume or donate electrons (note parallelism with ph - tendency to consume or donate protons), is key to understanding metal chemistry/mobility in these systems. How do we determine the relative tendencies of different compounds/metals to be reduced (donate electrons) or oxidized (lose electrons)? 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):

5 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 Zn ions flow through the external wire to the copper rod. At the Cu end they are avail to combine with the incoming Cu++ (cupric ions) to form more copper metal. A voltmeter can be introduced within the circuit. The voltage produced by the cell can be measured. Voltage called electromotive force. To complete a porous contact allows flow of IONS back and forth.

6 And it is conventional to tabulate voltage corresponding to unit ACTIVITY for each ion (one molal). Note that for the reaction: Cu e - === Cu 2 x Avogadro s number (6.022 x ) of e - will be produced per mole of Cu produced. The amount of electrical charge associated with one Avogadro s number 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 a potential difference ε is w=fε: joule = coulombs x volts. Hence, the electrical work required to move n charge is: w=nfε). Non-IUPAC (Geochemical) Convention! G o = -nfe o and G = -nfe) Standard Redox Potential - the tendency of a reducing agent to donate electrons is given by a standard redox potential: E o Definition of Standard Redox Potential - the resultant electromotive force (voltage) on a half cell in which oxidants and reductants are present in 1.0M concentrations at ph = 7 and 25 o C in equilibrium with electrode that can reversibly accept electrons from the reductant.. Reference electrode = hydrogen electrode consists of gaseous H 2 at 1 atm bubbling over a platinum electrode in an acidic soln with activity of 1 for hydrogen ion (1 M HCl). Interpretation of redox potentials for REDOX pairs 1. The more negative the redox potential is relative to the H electrode, the greater reducing potential of the system. The greater the tendency to lose electrons (LEO reduces it s surrounding). 2. If more positive than reference electrode, there are lesser tendencies to donate electrons. Rather there is a stronger affinity for electrons and the element is a good oxidizing agent.

7 Relationship between Redox potential and Eh Electrochemical energy is a form of Free Energy for Redox Reactions: G o = -nfe o and G = -nfe (non-iupac convention, but geochemical) This expresses fact that EMF (E o ) multiplied by amount of electronic charge moving through the cell (charge carried by a mole of electron, e, times number of moles) gives maximum electrical work the cell can accomplish (think of galvanic cell).

8 Recall: G = G o - RT ln IAP Because G o = -nfe o and G = -nfe: Eh = E o - (2.303RT/nF ) log ([product]/[reactant]) At STP: Eh = E o - (0.0592/n) log ([product]/[reactant]) R = x 10-3 kcal/deg/mole, F = 23 kcal/volt /gram equivalent (Atomic Wt/valence). Application: Consider the half-cell reaction: Fe 3+ + e - = Fe 2+ Eh = E o - (0.059/n) log (a Fe2 +/a Fe3 +) E o = 0.77 volts Eh = 0.77 volts - (0.059/1) log (a Fe2 +/a Fe3 +) Eh when a Fe3 + = a Fe2 + is volts.

9 Another example: SO e H + = H 2 S + 4H 2 O Eh = E o - (0.059/8) log (a H2S a H2O 4/ a SO4,2- a H+ 10) E o can be calculated from free energy or electron potential values in Compilations. In the above case: E o = 0.30 V. At ph = 7, a H2S = a SO4-2 Eh = 0.30 V - (0.059/8) log (1/ 10-7 ) = V Eh-pH Diagrams (Pourbaix Diagram) In the example above: Eh = 0.30 V - (0.059/8) log (a H2S a H2O 4/ a SO42- a H+ 10) IF a H2O = 1, we note: Eh = 0.30 V - 10(0.059/8) ph + (0.059/8) log (a H2S / a SO4 )

10 NOW, the line of a SO4= = a H2S is not a specific Eh but a function also of ph Eh = 0.30 V ph At ph = 1, Eh = 0.23 V At ph = 10, Eh = V And so on. Hence, if we plot Eh vs. ph, we have a line. We could do the same to the H 2 O System: 4H + + O e - = 2H 2 O Eh = E o - (0.059/4) log (a H2O 2/P O2 a H+ 4) Or given: a H2O = 1, and maximum P O2 = 1 bar Eh = E o + (0.059) log (a H+ ) = E o ph For the above reaction, E o = 1.23 V, hence: Eh = ph Again, an equation of a straight line in a Pourbaix Diagram. This defines the upper Eh-pH stability of water on the Earth s surface. Because air has

13 P O2 = 0.3 bars, the upper stability of water in contact with air is (Eh = ph 0.523) The lower stability of water can be defined by the reaction: 2H + + 2e - = H 2 Eh = E o - (0.059/2) log (P H2 /a H+ 2) Eh = E o - (0.059) ph E o = 0 because this is the reference electrode reaction: Eh = ph Relationship between Redox potential, Eh, and pe From the reaction: Fe 3+ + e - = Fe 2+

14 K eq = [Fe 2+ ]/([Fe 3+ ][e - ] 1/e - = K eq [Fe 3+ ]/([Fe 2+ ] Or (by analogy to ph): pe = log K eq + log ([Fe 3+ ]/([Fe 2+ ]) Hence (Eh = (RT/nF) ln K eq + (RT/nF) ln (IAP)): pe = (nf/2.303rt) Eh pe = 13 + log (([Fe 3+ ]/([Fe 2+ ]) because (log K eq at 25 o C is 13) pe-ph DIAGRAMS (Caveat: redox equilibrium is not always attained in natural systems. The following is intended to show what the equilibrium relations are, given that systems tend towards these equilibrium values. In many systems, it is precisely the availability of energy due to lack of equilibrium that is "catalyzed" by organisms leading to regular patterns of oxidation-reduction reactions).

15 Fe-O-H 2 O System: Consider the species: O 2, H 2, H 2 O, Fe 2+, Fe 3+, Fe 3 O 4, Fe 2 O 3 The high pe limit of water : 0.5 O 2 + 2e H + = H 2 O log K eq = -0.5log P O2 + 2pe + 2pH = (From Thermo Tables) The limit may be defined as P O2 = 1 atm (hence log P O2 = 0) Hence: pe = ph (show in a plot; line with slope of -1 and intercept of 20.78) Similarly, the lower pe limit is defined by: H + + e = 0.5 H 2 (Standard state of formation of H 2 ; G o = log K eq = 0) The lower limit defined as P H2 = 1 atm (and log P H2 = 0) Hence: pe = -ph (line with slope of -1)

16 Recall the Fe 3+ /Fe 2+ reaction earlier (pe = log K eq + log ([Fe 3+ ]/([Fe 2+ ])): pe = 13 + log (([Fe 3+ ]/([Fe 2+ ]) The line of equal concentration of Fe 3+ and Fe 2+ is at pe = 13 (horizontal line) Boundary Between Solids: Fe(OH)2, Fe(OH)3 Fe(OH) 3 + e - + H + = Fe(OH) 2 + H 2 O pe = log K eq - ph = ph Boundary Between Solid and Dissolved Ions: Fe 2+, Fe(OH)3 Fe(OH) 3 + e - + 3H + = Fe H 2 O pe = log K eq - log Fe ph Given Fe 2+ = 1x10-5 and K eq = 17.2 pe = ph Other Boundaries are Derived Similarly:

17 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 ligands present (e.g. CO 2, Sulfur). The same reasoning controls the speciation of complexes that could be present (in terms of the relative dominance of complexes).

18 NATURAL WATERS: pe-ph Ranges, and Significance

19 NATURAL WATERS: Natural water pe-ph ranges are shown in Figure 7.3 (Schlesinger). This range by and large limits the speciation of metals and other aqueous species (including complexes) in natural systems (waste water systems and other special systems would be another matter).

23 " Redox reactions are central to understanding the biogeochemistry of lakes. In lakes and wetlands, oxygen depletion in deeper systems arises from the low solubility of oxygen in water, low diffusivity of oxygen in water, and consumption of oxygen by metabolic activities of organisms. " Redox gradients and processes extend beneath the water/sediment interface. " In lakes and wetlands, the important species to look out for among others are: (1) O 2 (aeration, deaeration) (2) nutrient N distribution (control on variation with depth and time (NH 4 +, NO 2 -, NO 3 - ) (3) sulfur, availability (like N) controls decomposition of organics and stability of microbes (desulfuvibrio, etc.) H 2 S, HS -, SO 4 (4) CO 2, HCO 3 -, CO 3 =, CH 4 (also microbial applications) " Microbial mediation make it possible to accomplish these reactions even at low temperatures, but obviously, these processes are favored by higher temperatures. " The energetically/thermodynamically most favorable reactions usually preempts competing reactions because use of common electron denors (organic matter). Because microbial species specializes on specific reactions, they distribute themselves in a manner dictated by the location of the specific reactions they want to mediate. " Redox gradients are (of course) also present in marine sediments, but relative importance of specific reactions are different (e.g., the sulfate

24 reducers versus methanogens - compete for organic electron donors). " 1/4CH 2 O + 1/4 O 2 = 1/4CO 2 + 1/4H 2 O G = kcal/mol (aerobic respiration) " 1/4CH 2 O + 1/5 NO /5 H + = 1/4CO 2 +1/10 N 2 + 7/10H 2 O G = kcal/mol (denitrifying bacteria) (-facultative anaerobes zone) Mn 4+ reduction ( obligate anaerobes zone, e.g. Clostridium) Fe 3+ reduction " 1/4CH 2 O + 1/8 SO /8 H + = 1/4CO 2 +1/8 HS - + 1/4H 2 O G = kcal/mol (Desulfovibrio, Desulfotomaculum) " 1/4CH 2 O = 1/8CO 2 +1/8 CH 4 G = kcal/mol (dissimilators) " 1/8CO 2 +1/2 H 2 = 1/8CH 4 + 1/4H 2 O G = -4 kcal/mol (reductors)

25 "Hydrogen Fermentation" " 1/4CH 2 O + 1/4 H 2 O = 1/4CO 2 +1/2 H 2 G = kcal/mol

26 NITRATE REDUCTION SULPHATE REDUCTION " 1/4CH 2 O + 1/8 SO /8 H + = 1/4CO 2 +1/8 HS - + 1/4H 2 O G = kcal/mol or (at lower ph) 1/4CH 2 O + 1/8 SO /4 H + = 1/4CO 2 +1/8 H 2 S 1/4H 2 O (Desulfovibrio, Desulfotomaculum) " The same microbes responsible for sulphate produce other important gases such as dimethyl sulfide and dimethyl disulfide (in an analogous reaction SeO 4 is reduced to elemental Se rendering Se more toxic to wildlife; H 2 S is also toxic). This sulfur flux was thought to dominate source of S in the atmosphere during pre-industrial times (with H 2 S >> methyl sulfides). " H 2 S generated may be re-oxidized by sulfur bacteria as it diffuse to overlying water column, react with organic matter to form organically-bound S (vulcanization; to 80 % of S in peat), or -- " H 2 S may react with Fe 2+ to form sulfide minerals (this cycle is linked with the P cycle because FeS precipitated means less Fe to precipitate or sorb phosphorus (see lecture on NPP and Nutrients). Steel structures through wetlands/lakes are likewise "attacked" by H 2 S to form FeS.

27 " δ 34 S is useful in examining which S has been through the dissimilatory sulphate pathway (depleted in 34 S). Even pyrite may utilize S that has been through this pathway instead of normal precipitation routes (can be examined in ores, too). l These pathways can be traced all the way back to the atmosphere ( 32 S depleted SO 4 may be derived from oxidation of wetland-reduced S gases. METHANOGENESIS " 1/4CH 2 O = 1/8CO 2 +1/8 CH 4 (CH 3 COOH is preferred "CH 2 O" substrate) G = kcal/mol (acetate dissimilation) 1/8CO 2 +1/2 H 2 = 1/8CH 4 + 1/4H 2 O G = -4 kcal/mol (CO 2 reduction) Hydrogen Fermentation provides the H 2 for CO 2 reducers: 1/4CH 2 O + 1/4 H 2 O = 1/4CO 2 +1/2 H 2 G = kcal/mol

28 " In marine environments, methanogenesis is limited by the abundance of sulphate, which competes formidably for acetate and H 2. Methanogenesis is thus observed only in regimes where SO 4 is low to begin with or depleted to a significant degree by sulphate reduction (not much overlap between sulphate reducers and methanogens in sediments). " Furthermore, even when methanogenesis takes place in marine systems, they tend to utilize CO 2 reduction pathways because acetate is complete utilized by sulphate reducing bacteria. " In terrestrial systems, especially wetlands, methane production by either mechanisms is not inhibited (except in very saline marshes). " What controls CH 4 flux, however, is not production rate but oxidation during ascent. Observed dependence of flux on moisture content of bogs and tundra is related to efficient oxidation (10 to 100%) of asccending methane during dry seasons. The physics of bubble formation and ascent, by-pass provided by hollow plant stems are also significant factors. " Rice paddies account for about half of wetland emission of methane. " In lakes, methane oxidation is also key. Much of the oxidation can take place in overlying sediments (to 50% in Lake Washington), but in reduced (stratified systems) much of the oxidation occur in the water column. " On a global scale, as much as 2% of NPP in lakes eventually end up

29 as CH 4 in the atmosphere (that is not a trivial amount, especially considering isotopically peculiar properties of CH 4 derived by methanogenesis (below). " Acetate dissimilation ( 13 C = -50 to -65 per mil) and CO 2 reduction ( 13 C = -60 to -100 per mil) are distinguishable pathways of methane formation (recall Whiticar's talk). " (bio)methylation is another process that operates in very reduced environments of wetlands and lakes. Bacteria, again, utilize "free energy" from reactions such as: Hg 2+ + RCH 3 = HgCH R + or Hg R-CH 3 = Hg(CH 3 ) 2 + 2R + The concern with biomethylation is that for Hg and other metals, the methyl form is either more easily assimilated or dramatically more toxic than the non-methylated species. James Bay, Palawan, gold panning. " GLOBAL IMPORTANCE OF REDOX PROCESSES When one realize that 1/4 of total organic storage of the world are in wetland and tundra soils, it is easier to appreciate that the fate of this stored carbon has a lot of say in global-scale balance. GEOCHEMICAL WORKBENCH:

30 Please visit ACT2 beginning with p. 37 of the User's Guide. Show Eh-pH speciation for URANIUM-H 2 O system.

Oxidation States. 1. Redox potential Oxic vs. anoxic Simple electrochemical cell Redox potential in nature

Oxidation States. 1. Redox potential Oxic vs. anoxic Simple electrochemical cell Redox potential in nature 1. Redox potential Oxic vs. anoxic Simple electrochemical cell Redox potential in nature 2. Redox reactions Redox potential of a reaction Eh ph diagrams Redox reactions in nature 3. Biogeochemical reactions