Oxidation-Reduction Dr. Ron Rusay Balancing Oxidation-Reduction Reactions.
Biogeochemical Reduction-Oxidation (Redox) Reactions … · Biogeochemical Reduction-Oxidation...
Transcript of Biogeochemical Reduction-Oxidation (Redox) Reactions … · Biogeochemical Reduction-Oxidation...
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
Export
NPP = net primary production
NEP = net ecosystem production (P-R)