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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+

2

•  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

•  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

Salt bridge

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•  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

Salt bridge

•  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

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•  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:

If we assume [oxidized] = [reduced] = 1 (i.e., at standard state), then:

Oxidized species + e- + H+ ↔ reduced species

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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

or Eh = 2.3 RT pE/F

•  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 H20

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Eh-pH diagrams describe the thermodynamic stability of chemical species under different biogeochemical conditions

Diagram is for 25 degrees C

Example – predicted stable forms of Fe in

aqueous solution:

Example -- Oxidation of H2S released from anoxic sediments into oxic surface water:

Sediment

Water

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•  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)

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

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•  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

DEC

REA

SING

ENER

GY YIELD

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•  Light used directly by phototrophs

•  Hydrothermal energy utilized via heat-catalyzed production of inorganics

Nealson and Rye 2004

Vertical scaling???

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Temporal pattern reflects decreasing energy yield:

1

3 (reactant)

3 (product) 2

4 decreasing redox

•  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

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•  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”

Anoxic sed

Root zone

•  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)

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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

O2

NO3-

Mn2+

SO42-

CH4

Concentration (not to scale)

Dep

th

00

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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

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

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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

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:

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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:

•  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

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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