Download Wetlands, and biogeochemical redox reactions in aquatic systems

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
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 States
Element
Nitrogen
Sulfur
Iron
Manganese
Oxidation State
Species
NO3NO2N2
NH3, NH4+
SO42S2O32S°
H2S, HS-, S2Fe3+
Fe2+
MnO42MnO2 (s)
MnOOH (s)
Mn2+
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•  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
Salt bridge
•  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
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•  Container on right side is more
oxidizing and draws electrons
from left side
Salt bridge
•  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
•  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:
Oxidized species + e- + H+ ↔ reduced species
If we assume [oxidized] = [reduced] = 1 (i.e., at standard state),
then:
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The Nernst Equation can be used to relate this equation to
measured Pt-electrode voltage (Eh, Eh , EH):
or
Eh = 2.3 RT pE/F
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
•  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
Example – predicted
stable forms of Fe in
aqueous solution:
Diagram is for 25 degrees C
Example -- Oxidation of H2S
released from anoxic
sediments into oxic surface
water:
Water
Sediment
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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 acidbase 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
DECREASING ENERGY YIELD
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•  Light used
directly by
phototrophs
•  Hydrothermal
energy utilized
via heatcatalyzed
production of
inorganics
Nealson and Rye 2004
Vertical scaling???
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Temporal pattern reflects decreasing energy yield:
3 (reactant)
1 decreas
ing r
edox
2
3 (product)
4
•  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
Root zone
Anoxic sed
•  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”
•  Terrestrial
(dry) systems
tend to have
medium NPP,
high + NEP
•  Wetlands
have high
NPP,
+ or - NEP
•  Aquatic
systems have
low NPP,
- NEP
Export
Drained wetlands or aquatic systems are
major sites of “old C” oxidation
NEP = net ecosystem production (P-R)
NPP = net primary production
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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
21.25
19.85
21.0
17.0
16.20
8.2
5.33
3.7
3.06
3.06
Denitrification
1/4CH2O + 1/5NO3 + 1/5H+ = 1/4CO2(g) + 1/10N2(g) +7/20H2O
Manganese Reduction
1/4CH2O + 1/2MnO2(s) + H+ = 1/4CO2(g) + 1/2Mn2+ + 3/4H2O
Iron Reduction
1/4CH2O + Fe(OH)3(s) + 2H+ = 1/4CO2(g) + Fe2+ + 11/4 H2O
Sulfate Reduction
1/4CH2O + 1/8SO42- + 1/8H+ = 1/4CO2(g) + 1/8HS- + 1/4H2O
Methane Fermentation
1/4CH2O = 1/8CO2(g) + 1/8CH4
Free energy available
ΔGr° = - 2.3 RT logK = -5.708 logK
R = 8.314 J deg-1 mol-1
T = °K = 273 + °C
Tracers are circled
0
0
Concentration (not to scale)
O2
NO3-
Depth
Mn2+
SO42-
CH4
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O2
NO3Mn++
Fe++
General guideline for
OATZ progression
Vertical scale changes across
environments
has been traditionally
oversimplified to SO42and H2S
SO42NH4+
S2-
Redox profiling
CH4
Nealson and Stahl 1997
Sulfur redox cycling
general overview
l
Su
f
r
ur
ed
S0
elemental
sulfur
n
tio
c
u
i
ox
ti
da
H2S
hydrogen
sulfide
m
in
era
li
on
SO42sulfate
n
o
ti
c
)
u
ed tor y
r
e la
fat imi
l
Su ass
(
Sulfate reduction
(dissimilatory)
za
ti
on organic S
Megonigal et al. 2004
13
id
ox
on
ati
SO32sulfite
S
S
hydrogen
polysulfide
sulfide HS
H2S
-2
-1
SO32-
tetrathionate
S
S2-
SO42sulfate
SO3
SO32-
thiosulfate
S
S8
S0
Sx
elemental
sulfur
d
re
n
tio
c
u
sulfur oxidation state
0
1
2
3
4
5
6
Sulfur redox cycling
significance
Dissimilatory sulfate reduction:
SO42- + 2CH2O  2HCO3- + HS- + H+
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
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Sulfur redox cycling
important Fe/S chemistry
H2S oxidation (ph >6):
O2 + Fe2+  Fe3+
Fe3+ + H2S  Fe2+ + S(0) as S8 and Sx2FeS formation and dissociation:
Fe2+ + H2S  FeSaq + 2H+
(FeSaq formation is enhanced with increasing temperature)
Pyrite formation:
FeSaq + H2S  FeS2 + H2
or under milder reducing conditions:
FeS(s) + S2O32-  FeS2 + SO32-
•  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
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