Download Biogeochemical Reduction-Oxidation (Redox) Reactions in Aquatic

Biogeochemical
Reduction-Oxidation (Redox)
Reactions in Aquatic Systems
Biogeochemical Systems -- OCN 401
25 September 2012
Reading: Schlesinger Chapter 7
Outline
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
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 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
A Simple Electrochemical Cell
Voltmeter
• FeCl2 at different redox
potentials in the two sides
Agar, KCl
• Wire with inert Pt at ends -voltmeter between electrodes
Salt bridge
e-
e-
• Electrons flow along wire,
and Cl- diffuses through salt
bridge to balance charge
Pt
Pt
-
Cl
ClFe2+ - e- = Fe3+
Fe2+
-
Cl
Cl- Fe3+
ClFe3+ + e- = Fe2+
• Voltmeter measures electron
flow
• Charge remains neutral
• Container on right side is more
oxidizing and draws electrons
from left side
Voltmeter
Agar, KCl
Salt bridge
• Electron flow and Cl- diffusion
continue while an equilibrium is
established – steady voltage
measured on voltmeter
e-
• If container on right also
contains O2, Fe3+ will
precipitate and greater voltage
is measured
e-
Pt
Pt
ClFe
ClFe2+ - e- = Fe3+
2+
ClCl- Fe3+
ClFe3+ + e- = Fe2+
4Fe3+ + 3O2 + 12e→ 2Fe2O3 (s)
• The voltage is characteristic for
any set of chemical conditions
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 potential
of 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:
K =
[ reduced species ]
[ oxidized species ][ e − ][ H + ]
K=
[ reduced species ]
[oxidized species ][e− ][ H + ]
−
+
logK = log [red] − log [ox] − log [e ] − log [H ]
logK = − pred + pox + pe + pH
If we assume [oxidized] = [reduced] = 1 (i.e., at standard state),
then:
log K = pe + pH
lo g K = p e + p H
The “Nernst Equation” can be used to relate the above
equation to the measured Pt-electrode voltage (Eh, Eh , EH ):
F
Eh
pe =
2.3RT
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”
Eh- pH (pe – pH)
Diagrams
• Used to show equilibrium speciation
for reactants, as functions of Eh (or
pe) and pH
F
• pe =
Eh
2.3RT
• Red lines are practical Eh-pH limits
on Earth
2
Eh-pH diagram for H20
Eh-pH diagrams describe the thermo-dynamic stability of
chemical species under different biogeochemical conditions
1.2
Fe +3 aq
FeOH +2 aq
Fe(OH) 2+ aq
O2
dE/dpH =
-0.059
Fe(OH) 3
Eh (volts)
Example – predicted
stable forms of Fe in
aqueous solution:
Fe +2 aq
0.0
H 2O
H2
Fe 3(OH)8
Fe(OH)2
-0.6
Diagram is for 25 degrees C
7
1
pH
12
Example -- Oxidation of H2S
released from anoxic
sediments into oxic surface
water:
pe
Water
Sediment
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
ΔG
+0.812
-29.9
+0.747
-28.4
+0.526
-23.3
+3H2O
-0.047
-10.1
SO4 + 10H + 8e --> H2S + 4H2O
-0.221
-5.9
-0.244
-5.6
Reduction of O2
+
-
O2 + 4H +4e --> 2H2O
Reduction of NO3-
+
-
2NO3 + 6H + 6e --> N2 + 3H2O
Reduction of Mn (IV)
+
-
2+
MnO2 + 4H + 2e --> Mn +2H2O
Reduction of Fe (III)
+
-
2+
Fe(OH)3 + 3H + e --> Fe
2-
Reduction of SO4
2-
+
-
Reduction of CO2
+
-
CO2 + 8H + 8e --> CH4 + 2H2O
DECREASING ENERGY YIELD
Eh (V)
Reducing Half-reaction
Example: Changing Composition in
Flooded Soils
Relative concentration
Temporal pattern reflects decreasing energy yield:
Easily reducible Mn
3 (reactant)
Eh
1
Fe
O2
2
NO3-
0
1
Exchangeable Mn
3 (product)
2
3
4
5
Days after flooding
6
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
Marine Sediment Depth Profiles
0
0
Reaction
Eh (V)
ΔG
+0.812
-29.9
Concentration (not to scale)
O2
NO3-
Reduction of O2
+
-
O2 + 4H +4e --> 2H2O
Reduction of NO3-
+
Mn2+
-
2NO3 + 6H + 6e --> N2 + 3H2O
+0.747
-28.4
+
-
2+
MnO2 + 4H + 2e --> Mn +2H2O
Reduction of Fe
+0.526
-23.3
-0.047
-10.1
-0.221
-5.9
-0.244
-5.6
3+
+
-
2+
Fe(OH)3 + 3H + e --> Fe +3H2O
Depth
4+
Reduction of Mn
SO42-
2-
Reduction of SO4
2-
+
-
SO4 + 10H + 8e --> H2S + 4H2O
Reduction of CO2
+
-
CO2 + 8H + 8e --> CH4 + 2H2O
CH4
Methanogenesis in Wetlands
•
High OM 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
1
0.5
Eh
• 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
HCO3-
CO2
CO32-
Root zone
0
Anoxic sed
CH4
-0.5
2
6
pH
10
• 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”
14
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
Export
Drained wetlands or aquatic systems are
major sites of “old C” oxidation
NEP = net ecosystem production (P-R)
NPP = net primary production