Lecture Summary Physical properties of water exert profound control on nutrient cycling and NPP in lakes Lakes respond dynamically to seasonal climate change The biogeochemical character of lakes is directly linked to the nature of the surrounding landscape and geology Eutrophication is a natural process, which can be accelerated by anthropogenic activities Acid rain has had profound impacts on some lakes; underlying geology is a factor 1
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+ 2
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 Assign O to be (II) & H to be (+I), calculate. Element Nitrogen Sulfur Iron Manganese Oxidation State N (+V) N (+III) N (O) N (III) S (+VI) S (+II) S (O) S(II) Fe (+III) Fe (+II) Mn (+VI) Mn (+IV) Mn (+III) Mn (+II) 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+ 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 3
Voltmeter Agar, KCl e Salt bridge e Pt Pt Cl Cl Fe 2+ Cl Fe 3+ Cl Cl Fe 2+ e = Fe 3+ Fe 3+ + e = Fe 2+ FeCl 2 at different Fe oxidation states in the two sides 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 Voltmeter Agar, KCl e Salt bridge e Pt Pt Cl Cl Fe 2+ Cl Fe 3+ Cl Cl Fe 2+ e = Fe 3+ Fe 3+ + e = Fe 2+ Container on right side is more oxidizing and draws electrons from left side 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 4
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 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 oxidationreduction reaction at equilibrium: K = [ reduced species] [ oxidized species][ e ][ H + ] 5
Oxidized species + e + H + reduced species K = [ oxidized [ reduced species ] species ][ e ][ H ] + log K = log [ red] log [ ox] log [ e log K = p red + p ox + pe + ph ] log [ H + ] If we assume [oxidized] = [reduced] = 1 (i.e., at standard state), then: log K = pe + ph log K = pe + ph The Nernst Equation can be used to relate this equation to measured Ptelectrode voltage (Eh, E h, E H ): F pe pe = Eh or Eh = 2.3 RT pe/f 2.3RT 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 base10 logarithms 6
Used to show equilibrium speciation for reactants as functions of Eh and ph Red lines are practical EhpH limits on Earth F pe = pe = Eh 2.3RT EhpH diagram for H 2 0 EhpH diagrams describe the thermodynamic stability of chemical species under different biogeochemical conditions 1.2 Fe +3 aq FeOH +2 aq Fe(OH) 2 + aq O 2 de/dph = 0.059 Fe(OH) 3 Example predicted stable forms of Fe in aqueous solution: E h (volts) 0.0 Fe +2 aq H 2 O H 2 Fe 3 (OH) 8 Fe(OH) 2 0.6 Diagram is for 25 degrees C 1 7 12 ph 7
Example Oxidation of H 2 S released from anoxic sediments into oxic surface water: Water Sediment 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 halfreaction 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) 8
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 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 highyield electron acceptors to the use of lowyield electron acceptors 9
Why is organic matter an electron donor? Zscheme for photosynthetic electron transport Falkowski and Raven (1997) ADP ATP Ferredoxin Energy Scale ADP ATP to Krebs cycle and carbohydrate formation photooxidation of water energy from sun converted to CC, energy rich, chemical bonds 10
Reducing Halfreaction E h (V) Δ G Reduction of O 2 O 2 + 4H + +4e > 2H 2 O +0.812 29.9 Reduction of NO 3 2NO 3 + 6H + + 6e > N 2 + 3H 2 O +0.747 28.4 Reduction of Mn (IV) MnO 2 + 4H + + 2e > Mn 2+ +2H 2 O +0.526 23.3 Reduction of Fe (III) Fe(OH) 3 + 3H + + e > Fe 2+ +3H 2 O 0.047 10.1 Reduction of SO 4 SO 4 + 10H + + 8e > H 2 S + 4H 2 O 0.221 5.9 Reduction of CO 2 DECREASING ENERGY YIELD CO 2 + 8H + + 8e > CH 4 + 2H 2 O 0.244 5.6 Light used directly by phototrophs Hydrothermal energy utilized via heatcatalyzed production of inorganics Nealson and Rye 2004 11
Vertical scaling??? Temporal pattern reflects decreasing energy yield: Relative concentration 1 O 2 NO 3 E h decreasing redox 2 Easily reducible Mn 3 (reactant) Exchangeable Mn 3 (product) 0 1 2 3 4 5 Days after flooding Fe 2+ 6 4 12
High organiccarbon 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 Oxidation can be predicted from EhpH diagram of C in aqueous solution 1 0.5 CO 2 and CH 4 are released both by direct ebullution and pumping from roots to leaves Eh 0 CO 2 Root zone HCO 3 CO 3 As much as 510% of net ecosystem production may be lost as CH 4 0.5 Anoxic sed CH 4 2 6 10 14 ph Terrestrial and wetland methanogenesis is an important source of this greenhouse gas 13
Terrestrial (dry) systems tend to have medium NPP, high + NEP Wetlands have high NPP, + or NEP Aquatic systems have low NPP, NEP Export Drained wetlands or aquatic systems are major sites of old C oxidation NEP = net ecosystem production (PR) NPP = net primary production Redox sequence of OM decomposition log K log Kw Aerobic Respiration 1/4CH2O + 1/4O2 = 1/4H2O + 1/4CO2(g) 20.95 20.95 21.25 19.85 21.0 17.0 16.20 8.2 5.33 3.7 3.06 3.06 Denitrification 1/4CH2O + 1/5NO3 + 1/5H+ = 1/4CO2(g) + 1/10N2(g) +7/20H2O Manganese Reduction 1/4CH2O + 1/2MnO2(s) + H+ = 1/4CO2(g) + 1/2Mn2+ + 3/4H2O Iron Reduction 1/4CH2O + Fe(OH)3(s) + 2H+ = 1/4CO2(g) + Fe2+ + 11/4 H2O Sulfate Reduction 1/4CH2O + 1/8SO4 + 1/8H+ = 1/4CO2(g) + 1/8HS + 1/4H2O Methane Fermentation 1/4CH2O = 1/8CO2(g) + 1/8CH4 Tracers are circled Free energy available ΔGr = 2.3 RT logk = 5.708 logk R = 8.314 J deg1 mol1 T = K = 273 + C 14
0 0 Concentration (not to scale) O 2 Reaction E h (V) Δ G NO 3 Reduction of O 2 O 2 + 4H + +4e > 2H 2 O +0.812 29.9 Reduction of NO 3 2NO 3 + 6H + + 6e > N 2 + 3H 2 O +0.747 28.4 Reduction of Mn 4+ MnO 2 + 4H + + 2e > Mn 2+ +2H 2 O +0.526 23.3 Reduction of Fe 3+ Fe(OH) 3 + 3H + + e > Fe 2+ +3H 2 O 0.047 10.1 Reduction of SO 4 SO 4 + 10H + + 8e > H 2 S + 4H 2 O 0.221 5.9 Depth Mn 2+ SO 4 Reduction of CO 2 CO 2 + 8H + + 8e > CH 4 + 2H 2 O 0.244 5.6 CH 4 Microbial Mats steep gradients 15
Wisconsin & Australia Hamelin Bay, Australia 16
Laboratory Feoxidizing microbial cultures Sobolev et al. 2001, Roden et al. 2004 Banded iron formation 17
Hydrothermal vents O2 NO3Mn++ Fe++ SO42S NH4+ Redox profiling General guideline for OATZ progression Vertical scale changes across environments has been traditionally oversimplified to SO42and H2S CH4 Nealson and Stahl 1997 18
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) 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 19
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 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 20
Redox reactions control organicmatter oxidation and element cycling in aquatic ecosystems Eh ph diagrams can be used to describe the thermodynamic stability of chemical species under different biogeochemical conditions Biogeochemical reactions are mediated by the activity of microbes, and follow a sequence of hightolow 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 lowestyield redox reaction The bottom line Oxicanoxic 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 21