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 and their thermodynamic control Redox sequence of OM oxidation in aquatic environments Marine sediment profiles Methanogenesis in wetlands Many elements in the periodic table can exist in more than one oxidation state. Oxidation states are indicated by Roman numerals (e.g. (+I), (-II), etc). The oxidation state represents the "electron content" of an element which can be expressed as the excess or deficiency of electrons relative to the elemental state. Oxidation States Element Nitrogen Sulfur Iron Manganese Oxidation State Species NO 3 - NO 2 - N 2 NH 3, NH 4 + SO 4 S 2 O 3 S H 2 S, HS -, S Fe 3+ Fe 2+ MnO 4 MnO 2 (s) MnOOH (s) Mn 2+ 1
Redox potential expresses the tendency of an environment to receive or supply electrons An oxic environment has high redox potential because O 2 is available as an electron acceptor For example, Fe oxidizes to rust in the presence of O 2 because the iron shares its electrons with the O 2 : 4Fe + 3O 2 2Fe 2 O 3 In contrast, an anoxic environment has low redox potential because of the absence of O 2 the more positive the potential, the greater the species' affinity for electrons and tendency to be reduced FeCl 2 at different Fe oxidation states in the two sides Salt bridge Wire with inert Pt at ends -- voltmeter between electrodes Electrons flow along wire, and Cl - diffuses through salt bridge to balance charge Voltmeter measures electron flow Charge remains neutral 2
Container on right side is more oxidizing and draws electrons from left side Salt bridge Electron flow and Cl - diffusion continue until an equilibrium is established steady voltage measured on voltmeter If container on right also contains O 2, Fe 3+ will precipitate and greater voltage is measured 4Fe 3+ + 3O 2 + 12e - 2Fe 2 O 3 (s) The voltage is characteristic for any set of chemical conditions A mixture of constituents, not really separate cells We insert an inert Pt electrode into an environment and measure the voltage relative to a standard electrode [Std. electrode = H 2 gas above solution of known ph (theoretical, not practical). More practical electrodes are calibrated using this H 2 electrode.] Example: when O 2 is present, electrons migrate to the Pt electrode: O 2 + 4e - + 4H + 2H 2 O The electrons are generated at the H 2 electrode: 2H 2 4H + + 4e - Voltage between electrodes measures the redox potential 3
General reaction: Oxidized species + e - + H + reduced species Redox is expressed in units of pe, analogous to ph: pe = - log [e - ] (or Eh = 2.3 RT pe/f) where [e - ] is the electron concentration or activity pe is derived from the equilibrium constant (K) for an oxidation-reduction reaction at equilibrium: Oxidized species + e - + H + reduced species If we assume [oxidized] = [reduced] = 1 (i.e., at standard state), then: 4
The Nernst Equation can be used to relate this equation to measured Pt-electrode voltage (Eh, E h, E H ): or Eh = 2.3 RT pe/f where: Eh = measured redox potential as voltage R = the Universal Gas Constant (= 8.31 J K -1 mol -1 ) T = temperature in degrees Kelvin F = Faraday Constant (= 23.1 kcal V -1 equiv -1 ) 2.3 = conversion from natural to base-10 logarithms Used to show equilibrium speciation for reactants as functions of Eh and ph Red lines are practical Eh-pH limits on Earth Eh-pH diagram for H 2 0 5
Eh-pH diagrams describe the thermodynamic stability of chemical species under different biogeochemical conditions Example predicted stable forms of Fe in aqueous solution: Diagram is for 25 degrees C Example -- Oxidation of H 2 S released from anoxic sediments into oxic surface water: Water Sediment 6
Example: net reaction for aerobic oxidation of organic matter: CH 2 O + O 2 CO 2 + H 2 O In this case, oxygen is the electron acceptor the half-reaction is: O 2 + 4H + + 4e - 2H 2 O Different organisms use different electron acceptors, depending on availability due to local redox potential The more oxidizing the environment, the higher the energy yield of the OM oxidation (the more negative is ΔG, the Gibbs free energy) 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 ΔG r?? ΔG r = - nie Where n is the # of e- s in the rxn, I is Faraday s constant (23.06 cal V -1 ), and E is electropotential (V) pe for an electron transfer between a redox couple analogous to pk between conjugate acidbase pair 7
The higher the energy yield, the greater the benefit to organisms that harvest the energy In general: There is a temporal and spatial sequence of energy harvest during organic matter oxidation Sequence is from the use of high-yield electron acceptors to the use of low-yield electron acceptors DECREASING ENERGY YIELD 8
Light used directly by phototrophs Hydrothermal energy utilized via heatcatalyzed production of inorganics Nealson and Rye 2004 Vertical scaling??? 9
Temporal pattern reflects decreasing energy yield: 1 decreasing redox 3 (reactant) 4 2 3 (product) High organic-carbon levels in sediment promote OM oxidation CO 2 is reduced to CH 4 during OM oxidation Release of CH 4 from plant leaves Plants pump air from leaves roots sediment CH 4 is oxidized by O 2 in root zone: CH 4 + 2O 2 CO 2 + 2H 2 O 10
Oxidation can be predicted from Eh-pH diagram of C in aqueous solution CO 2 and CH 4 are released both by direct ebullution and pumping from roots to leaves Root zone Anoxic sed As much as 5-10% of net ecosystem production may be lost as CH 4 Terrestrial and wetland methanogenesis is an important source of this greenhouse gas Terrestrial (dry) systems tend to have medium NPP, high + NEP Wetlands have high NPP, + or - NEP Aquatic systems have low NPP, - NEP NPP = net primary production NEP = net ecosystem production (P-R) Export Drained wetlands or aquatic systems are major sites of old C oxidation 11
Redox sequence of OM decomposition log K log Kw Aerobic Respiration 1/4CH 2 O + 1/4O 2 = 1/4H 2 O + 1/4CO 2 (g) 20.95 20.95 Denitrification 1/4CH 2 O + 1/5NO 3 + 1/5H + = 1/4CO 2 (g) + 1/10N 2 (g) +7/20H 2 O Manganese Reduction 1/4CH 2 O + 1/2MnO 2 (s) + H + = 1/4CO 2 (g) + 1/2Mn 2+ + 3/4H 2 O Iron Reduction 1/4CH 2 O + Fe(OH) 3 (s) + 2H + = 1/4CO 2 (g) + Fe 2+ + 11/4 H 2 O 21.25 19.85 21.0 17.0 16.20 8.2 Sulfate Reduction 1/4CH 2 O + 1/8SO 4 + 1/8H + = 1/4CO 2 (g) + 1/8HS - + 1/4H 2 O 5.33 3.7 Methane Fermentation 1/4CH 2 O = 1/8CO 2 (g) + 1/8CH 4 3.06 3.06 Tracers are circled Free energy available ΔG r = - 2.3 RT logk = -5.708 logk R = 8.314 J deg -1 mol -1 T = K = 273 + C 0 0 Concentration (not to scale) O 2 NO 3 - Mn 2+ Depth SO 4 CH 4 12
O 2 NO 3 - Mn ++ Redox profiling General guideline for OATZ progression Vertical scale changes across environments SO 4 S Fe ++ NH 4 + has been traditionally oversimplified to SO 4 and H 2 S CH 4 Nealson and Stahl 1997 H 2 S hydrogen sulfide Megonigal et al. 2004 Sulfur redox cycling general overview Sulfur reduction oxidation mineralization S 0 elemental sulfur Sulfate reduction (dissimilatory) organic S SO 4 sulfate Sulfate reduction (assimilatory) 13
hydrogen sulfide oxidation S SO 3 sulfite SO 4 sulfate S SO tetrathionate 3 S SO 3 S polysulfide HS - H 2 S S S 8 S 0 S x thiosulfate elemental sulfur sulfur oxidation state SO 3 reduction -2-1 0 1 2 3 4 5 6 Dissimilatory sulfate reduction: Sulfur redox cycling significance SO 4 + 2CH 2 O 2HCO 3 - + HS - + H + Accounts for half or more of the total organic carbon mineralization in many environments Highly reactive HS - is geochemically relevant because of its involvement in precipitation of metal sulfides and potential for reoxidation 14
H 2 S oxidation (ph >6): Sulfur redox cycling important Fe/S chemistry O 2 + Fe 2+ Fe 3+ Fe 3+ + H 2 S Fe 2+ + S(0) as S 8 and S x FeS formation and dissociation: Fe 2+ + H 2 S FeS aq + 2H + (FeS aq formation is enhanced with increasing temperature) Pyrite formation: FeS aq + H 2 S FeS 2 + H 2 or under milder reducing conditions: FeS(s) + S 2 O 3 FeS 2 + SO 3 Redox reactions control organic-matter oxidation and element cycling in aquatic ecosystems Eh ph diagrams can be used to describe the thermo-dynamic stability of chemical species under different biogeochemical conditions Biogeochemical reactions are mediated by the activity of microbes, and follow a sequence of high-to-low energy yield that is thermodynamically controlled Example organic matter oxidation: O 2 reduction (closely followed by NO 3 - reduction) is the highest- yield redox reaction CO 2 reduction to CH 4 is the lowest-yield redox reaction 15
The bottom line Oxic-anoxic transitions There is a principle difference between gradients of compounds used for biomass synthesis and those needed for energy conservation, such as oxygen. Nutrient limitation leads to a decrease of metabolic activity, but absence of an energy substrate causes a shift in the composition of a microbial community, or forces an organism to switch to a different type of metabolism. Brune et al. 2000 16