Biogeochemical Systems -- OCN 40125 September 2012

BiogeochemicalReduction-Oxidation (Redox)

Reactions in Aquatic Systems

Reading: Schlesinger Chapter 7

1. Redox potential• Oxic vs. anoxic environments• 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• Marine sediment profiles• Methanogenesis in wetlands

Outline

Redox Potential: The Basics

• 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 O2because 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

A Simple Electrochemical Cell

• FeCl2 at different redox potentials 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 while 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

Redox Potential in Nature• A mixture of chemicals, not separate electrochemical 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 potentialof an environment

Redox Potential of a Reaction

• General reaction:

Oxidized species + e- + H+ ↔ reduced species

• Redox is expressed in units of “pe,” analogous to pH:

pe = - log [e-]

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

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

log K pe pH= +

pHpeppK

HeoxredK

K

oxred

Hespeciesoxidizedspeciesreduced

+++−=

−−−=

=

+−

+−

log

][log][log][log][loglog

]][][[][

The “Nernst Equation” can be used to relate the above equation to the measured Pt-electrode voltage (Eh, Eh , EH ):

where:Eh = measured voltage

F = Faraday Constant (= 23.1 kcal V-1 equiv-1)

R = the Universal Gas Constant (= 1.99 x 10-3 kcal °K-1 mol-1)

T = temperature (°K)

2.3 = conversion from natural to base-10 logarithms

Note: “pe” is also sometimes written as “pE”

log K pe pH= +

Eh2.3RT

Fpe =

Eh- pH (pe – pH) Diagrams

• Used to show equilibrium speciation for reactants, as functions of Eh (or pe) and pH

•

• Red lines are practical Eh-pH limits on Earth

Eh2.3RT

Fpe =

Eh-pH diagram for H20

2

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

Example – predicted stable forms of Fe in

aqueous solution: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 (vo

lts)

1.2

0.0

-0.6

1 7 12pH

Diagram is for 25 degrees C

pe

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

Sediment

Water

Redox Reactions in Nature

• Example: net reaction for aerobic oxidation of organic matter:

CH2O + O2 → CO2 + H2O

• In this case, oxygen is the electron acceptor – the reduction 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)

• 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

– Cause: high-yield electron acceptors are used before low-yield electron acceptors

Environmentally Important Organic Matter Oxidation Reactions

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

Example: Changing Composition in Flooded Soils

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

Redox Sequence of OM Oxidation in Aquatic Environments

• O2 reduction (aerobic oxidation): first, but [O2] in water is only ~0.2-0.3 mmol/L (mM) -- can run out if organic matter is abundant or circulation is restricted

• NO3 reduction (denitrification): next, but NO3 (typically <0.1 mM) runs out quickly

• Mn reduction and Fe reduction: dependent on soil composition

• SO4 reduction: important in marine environment, but usually minor in fresh water

• CO2 reduction (methanogenesis): very low energy yield, but lots of CO2, so can be very important in freshwater systems

• Only important in organic-rich freshwater environments, or in organic-rich and very restricted marine environments

O2

NO3-

Mn2+

SO42-

CH4

Concentration (not to scale)

Dep

th

Marine Sediment Depth Profiles

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

Methanogenesis in Wetlands

• High OM levels in sediment promote OM oxidation

• CO2 is reduced to CH4during 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

• CH4 oxidation can be predicted from Eh-pH

diagram of C in aqueous solution:

• CO2 and CH4 are released both by direct bubble ebullition (production) and pumping from roots to leaves

• As much as 5-10% of net ecosystem production may be lost as CH4

-0.5

0

0.5

1

2 6 10 14pH

Eh CO2

CH4

HCO3- CO3

2-

Anoxic sed

Root zone

• Terrestrial and wetland methanogenesis is an important source of this “greenhouse gas”

Lecture Summary• 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

– For 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 Next Lecture:

“Lakes, Primary Production, Budgets and Cycling”

Armed with a knowledge of terrestrial biogeochemistry, we’ll look at how lake primary production is closely linked to land-based nutrient supply, and how lakes respond to seasonal climate changes.

Also, we’ll examine how nutrient and carbon budgets provide key means for assessing lake biogeochemistry.

Wetlands Are the Interface Between Terrestrial and Aquatic Systems

• Terrestrial (dry) systems tend to have medium NPP, high pos NEP

• Wetlands have high NPP, pos or neg NEP

• Aquatic systems have low NPP, neg NEP

Drained wetlands or aquatic systems are major sites of “old C” oxidation

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NPP = net primary production

NEP = net ecosystem production (P-R)