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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
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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
Oxidation States
Element Nitrogen
Sulfur
Iron
Manganese
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 State Species NO3
-
NO2-
N2 NH3, NH4
+
SO42-
S2O32-
S° H2S, HS-, S2-
Fe3+
Fe2+
MnO42-
MnO2 (s) MnOOH (s) Mn2+
3
Oxidation States
Element Nitrogen
Sulfur
Iron
Manganese
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 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 NO3
-
NO2-
N2 NH3, NH4
+
SO42-
S2O32-
S° H2S, HS-, S2-
Fe3+
Fe2+
MnO42-
MnO2 (s) MnOOH (s) Mn2+
Assign O to be (-II) & H to be (+I), calculate.
• Redox potential expresses the tendency of an environment to receive or supply electrons
– An oxic environment has high redox potential because O2 is available as an electron acceptor
For example, Fe oxidizes to rust in the presence of O2 because the iron shares its electrons with the O2:
4Fe + 3O2 → 2Fe2O3
– In contrast, an anoxic environment has low redox potential because of the absence of O2
the more positive the potential, the greater the species' affinity for electrons and tendency to be reduced
4
• FeCl2 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
e-
e-
Voltmeter
Pt
Agar, KCl
Pt
Fe2+ - e- = Fe3+ Fe3+ + e- = Fe2+
Fe2+
Cl-Cl-Cl-
Cl-Cl-
Fe3+
Salt bridge
• 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 O2, Fe3+ will precipitate and greater voltage is measured
4Fe3+ + 3O2 + 12e-
→ 2Fe2O3 (s)
• The voltage is characteristic for any set of chemical conditions
e-
e-
Voltmeter
Pt
Agar, KCl
Pt
Fe2+ - e- = Fe3+ Fe3+ + e- = Fe2+
Fe2+
Cl-Cl-Cl-
Cl-Cl-
Fe3+
Salt bridge
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• 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 = H2 gas above solution of known pH (theoretical, not practical). More practical electrodes are calibrated using this H2 electrode.]
– Example: when O2 is present, electrons migrate to the Pt electrode:
O2 + 4e- + 4H+ → 2H2O
– The electrons are generated at the H2 electrode:
2H2 → 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 oxidation-reduction reaction at equilibrium:
]][][[][
+−=Hespeciesoxidized
speciesreducedK
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If we assume [oxidized] = [reduced] = 1 (i.e., at standard state), then:
logK pe pH= +
pHpeppKHeoxredK
K
oxred
Hespeciesoxidizedspeciesreduced
+++−=
−−−=
=
+−
+−
log][log][log][log][loglog
]][][[][
Oxidized species + e- + H+ ↔ reduced species
The Nernst Equation can be used to relate this equation to measured Pt-electrode voltage (Eh, Eh , EH):
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
logK pe pH= +
Eh2.3RTFpEpe =≡ or Eh = 2.3 RT pE/F
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• Used to show equilibrium speciation for reactants as functions of Eh and pH
• Red lines are practical Eh-pH limits on Earth
• Eh2.3RTFpepE ==
Eh-pH diagram for H20
Eh-pH diagrams describe the thermodynamic stability of chemical species under different biogeochemical conditions
Fe+3 aqFeOH+2 aqFe(OH)2
+ aqO2
H2O
Fe(OH)3
Fe+2 aq
H2
Fe3(OH)8
Fe(OH)2
dE/dpH = -0.059
E h (v
olts
)
1.2
0.0
-0.6
1 7 12pH
Diagram is for 25 degrees C
Example – predicted stable forms of Fe in
aqueous solution:
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Example -- Oxidation of H2S released from anoxic sediments into oxic surface water:
Sediment
Water
• Example: net reaction for aerobic oxidation of organic matter:
CH2O + O2 → CO2 + H2O
• In this case, oxygen is the electron acceptor – the half-reaction is:
O2 + 4H+ + 4e- → 2H2O
• 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)
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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 ΔGr ??
ΔGr = - 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 acid-base 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 high-yield electron acceptors to the use of low-yield electron acceptors
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Why is organic matter an electron donor?
Z-scheme for photosynthetic electron transport
to Krebs cycle and carbohydrate formation
Ferredoxin
energy from sun converted to C-C, energy rich, chemical bonds
photooxidation of water
ADP→ATP
ADP→ATP
Falkowski and Raven (1997)
Ener
gy S
cale
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Reducing Half-reaction Eh (V) ΔGReduction of O2
O2 + 4H + +4e- --> 2H2O +0.812 -29.9Reduction of NO3
-
2NO3- + 6H+ + 6e- --> N2 + 3H2O +0.747 -28.4
Reduction of Mn (IV) MnO2 + 4H+ + 2e- --> Mn2+ +2H2O +0.526 -23.3Reduction of Fe (III) Fe(OH)3 + 3H+ + e- --> Fe2+ +3H2O -0.047 -10.1
Reduction of SO42-
SO42- + 10H+ + 8e- --> H2S + 4H2O -0.221 -5.9
Reduction of CO2
CO2 + 8H+ + 8e- --> CH4 + 2H2O -0.244 -5.6
DEC
REA
SING
ENER
GY YIELD
• Light used directly by phototrophs
• Hydrothermal energy utilized via heat-catalyzed production of inorganics
Nealson and Rye 2004
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Vertical scaling???
Easily reducible Mn
O2
Eh
Exchangeable MnNO3-
Days after flooding
Rel
ativ
e co
ncen
trat
ion
60 1 2 3 4 5
Fe2+
Temporal pattern reflects decreasing energy yield:
1
3 (reactant)
3 (product) 2
4 decreasing redox
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• High organic-carbon levels in sediment promote OM oxidation
• CO2 is reduced to CH4 during OM oxidation
• Release of CH4 from plant leaves
• Plants pump air from leaves → roots → sediment
• CH4 is oxidized by O2 in root zone:
CH4 + 2O2 → CO2 + 2H2O
• Oxidation can be predicted from Eh-pH diagram of C in aqueous solution
• CO2 and CH4 are released both by direct “ebullution” and pumping from roots to leaves
• As much as 5-10% of net ecosystem production may be lost as CH4
• Terrestrial and wetland methanogenesis is an important source of this “greenhouse gas”
-0.5
0
0.5
1
2 6 10 14pH
Eh CO2
CH4
HCO3- CO32-
Anoxic sed
Root zone
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• Terrestrial (dry) systems tend to have medium NPP, high + NEP
• Wetlands have high NPP, + or - NEP
• Aquatic systems have low NPP, - NEP
Drained wetlands or aquatic systems are major sites of “old C” oxidation
Export
NPP = net primary production
NEP = net ecosystem production (P-R)
Redox sequence of OM decomposition log K log Kw Aerobic Respiration 1/4CH2O + 1/4O2 = 1/4H2O + 1/4CO2(g) 20.95 20.95
Denitrification 1/4CH2O + 1/5NO3 + 1/5H+ = 1/4CO2(g) + 1/10N2(g) +7/20H2O
21.25 19.85 Manganese Reduction 1/4CH2O + 1/2MnO2(s) + H+ = 1/4CO2(g) + 1/2Mn2+ + 3/4H2O 21.0 17.0 Iron Reduction 1/4CH2O + Fe(OH)3(s) + 2H+ = 1/4CO2(g) + Fe2+ + 11/4 H2O
16.20 8.2 Sulfate Reduction 1/4CH2O + 1/8SO4
2- + 1/8H+ = 1/4CO2(g) + 1/8HS- + 1/4H2O 5.33 3.7
Methane Fermentation 1/4CH2O = 1/8CO2(g) + 1/8CH4 3.06 3.06
Tracers are circled
Free energy available ΔGr° = - 2.3 RT logK = -5.708 logK R = 8.314 J deg-1 mol-1
T = °K = 273 + °C
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O2
NO3-
Mn2+
SO42-
CH4
Concentration (not to scale)
Dep
th
Reaction Eh (V) ΔGReduction of O2
O2 + 4H + +4e- --> 2H2O +0.812 -29.9Reduction of NO3
-
2NO3- + 6H+ + 6e- --> N2 + 3H2O +0.747 -28.4
Reduction of Mn4+
MnO2 + 4H+ + 2e- --> Mn2+ +2H2O +0.526 -23.3
Reduction of Fe3+
Fe(OH)3 + 3H+ + e- --> Fe2+ +3H2O -0.047 -10.1
Reduction of SO42-
SO42- + 10H+ + 8e- --> H2S + 4H2O -0.221 -5.9
Reduction of CO2
CO2 + 8H+ + 8e- --> CH4 + 2H2O -0.244 -5.6
00
Microbial Mats steep gradients
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Laboratory Fe-oxidizing microbial cultures
Sobolev et al. 2001, Roden et al. 2004
Banded iron formation
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Hydrothermal vents
Redox profiling
General guideline for OATZ progression
O2
NO3-
Mn++
Fe++
SO42-
S2- NH4+
CH4
has been traditionally oversimplified to SO4
2- and H2S
Nealson and Stahl 1997
Vertical scale changes across environments
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Sulfur redox cycling general overview
SO42-
sulfate H2S
hydrogen sulfide
Sulfate reduction (dissimilatory)
Sulfur r
educti
on S0
elemental sulfur
oxidation
organic S
Sulfate
reducti
on
(assimila
tory)
mineralization Megonigal et al. 2004
SO42-
sulfate
SO32-
sulfite
SO32- S
S SO3
tetrathionate
SO32- S
thiosulfate
S S8 S0
elemental sulfur
S2- Sx
polysulfide
H2S
HS- hydrogen
sulfide
0 -2 -1 1 2 3 4 5 6
reducti
on
oxidation
sulfur oxidation state
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Sulfur redox cycling significance
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
SO42- + 2CH2O 2HCO3
- + HS- + H+
Dissimilatory sulfate reduction:
Sulfur redox cycling important Fe/S chemistry
O2 + Fe2+ Fe3+
Fe3+ + H2S Fe2+ + S(0) as S8 and Sx2-
H2S oxidation (ph >6):
Fe2+ + H2S FeSaq + 2H+
(FeSaq formation is enhanced with increasing temperature)
FeS formation and dissociation:
FeSaq + H2S FeS2 + H2
or under milder reducing conditions: FeS(s) + S2O3
2- FeS2 + SO32-
Pyrite formation:
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• 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:
• O2 reduction (closely followed by NO3- reduction) is the
highest- yield redox reaction
• CO2 reduction to CH4 is the lowest-yield redox reaction
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