Falkowski et al, 2008

The Microbial Engines That DriveEarth’s Biogeochemical Cycles

Redox Chemistry in the Sea – the major driver of biogeochemical cycles

• Chemical reactions that involve transfer of electrons are called RedOx reactions (for reduction-oxidation).

• Redox active chemicals will spontaneously transfer electrons in order to achieve thermodynamic equilibrium (lowest free energy state). Equilibrium chemistry applies.

Fe3+ + e- Fe2+}}{{

}{3

2

eFe

FeKeq

Oxidation = loss of electrons

Reduction = gain of electrons

A chemical that loses elections undergoes oxidation and it oxidation number (valence state) increases. Conversely, a chemical that gains electrons undergoes reduction and its oxidation number (valence state) decreases.

Oxidation#Oxidation

Oxidation#

Reduction

Oxidation states of elements

Redox reactions are chemical reactions which involve the transfer of electrons, or more formally, a change in the oxidation state (or number) of the reactant which occurs as a result of e- transfer.

Fe3+ + e- Fe2+

Fe(III)ox Fe(II)red

Some elements are Redox Active in the environment and some are not. Refer to handout table for key redox-active elements and their important redox states

Examples of elements without appreciable redox chemistry in the environment include:

• Chlorine - nearly all Cl as a reduced form in Cl- in seawater• Major cations in seawater (Na+, K+, Mg2+ and Ca2+) (these elements are already oxidized relative to

their native metallic forms)

Key redox-active elements and their valence states and forms in nature.

Rules for assigning oxidation states

The simplest things to remember are that: 1) Any element in its native state will have an oxidation number of zero. i.e. O2, N2, S, Fe etc. (see the handout table of key elements and their

important oxidation states and forms).

2) In most other cases the element oxygen (O) is assigned the oxidation state of -2 and H = +1. There are exceptions to this, but these are not very common. Hydrogen peroxide, H2O2 is perhaps the most important of the exceptions, in which the O has an oxidation number of -1.

3) The sum of the oxidation numbers in a molecule must equal the charge on the molecule.

Note: the oxidation state of elements is often represented by roman numerals (i.e. -II or +IV). This is to distinguish the oxidation state from the ionic charge on the molecule. For the purposes of calculation, you can use arabic numbers.

Thermodynamic equilibrium

principles apply to the movement

of electrons. When chemicals have

electronic configurations

which are out of equilibrium,

relative to another chemical,

they will spontaneously react

together, transferring

electrons to attain

equilibrium – or the lowest

possible state of free

energy.

Zn(s) + Cu2+ + SO42- <=> Zn2+ + Cu(s) + SO4

2-

What has happened here?

Two redox active chemicals, at non-equilibrium concentrations, will have an electrical potential between them (i.e. a potential to transfer e-).

The electrical potential (E) of the system is called Ecell which is the sum of all half reactions (oxidation and reduction). Ecell is the electrical

potential between chemicals – in Volts

ΔG = -nFEcell

Where n= # moles of electrons transferred and F = 23.062 kcal/(V*mol electrons transferred) (Faraday constant)

The free energy change is related to the cell potential

We can write the reaction for oxidation of zinc as follows:

Zn(s) <=> Zn2+(aq) + 2e- (example of a half-

reaction)

where zinc metal (solid) loses 2 electrons to become the zinc ion. Because electrons cannot exist in a free state, this reaction would not occur if there was nothing to accept the electrons.In order for something to become oxidized, something else must become reduced - hence Red-Ox chemistry

Half cell potentials cannot be measured directly because they must occur in pairs.

A standard half reaction is used is for comparison to all other half reactions. The standard half reaction used is:

2 H+ + 2 e- <=> H2

(the reduction of H+ to hydrogen)

The half cell potential of this hydrogen reduction is set to zero, by convention.

It is said to have an Eho of zero.

The superscript o designates standard conditions (1 molal concentrations, 1 atm pressure etc).

Standard Hydrogen Electrode

Molal activity

For the reaction: Zn2+

(aq) + H2(g) <=> Zn(s) + 2 H+(aq)

Under standard conditions Ecell = -0.76 Volts

therefore the Eho of the half reaction:

Zn2+ (aq) + 2e- <=> Zn(s) is -0.76 V. ( Eh

o = Ecell relative to the hydrogen half cell potential).

If the reaction is written in reverse, as an oxidation, the Eh

o is +0.76 V. (Eho values are usually

tabulated for the reductions)

Standard Hydrogen Electrode

Consider again the chemical reaction between zinc metal (Zn(s)) and a solution of copper sulfate:

Zn(s) + Cu2+ + SO42- <=> Zn2+ + Cu(s) + SO4

2-

Comparing two half cell reactions we have the

Cu2+ + 2 e- = Cu(s) Eho = +0.34

Zn2+ + 2 e- = Zn(s) Eho = -0.76

The half reaction with the highest reduction potential will undergo reduction and the other half reaction will be an oxidation and will supply the electron(s).

From Table 7.1 in Libes

The reaction will occur spontaneously, if there is a difference in Eh between species

The potential of the oxidation reaction is reversed to +0.76V and the cell potential is therefore:Ecell = 0.34 + 0.76 = 1.10 Volts

Another way of expressing the electrical potential between two chemicals is by treating the electrons as reactants and calculating the “activity” of the electrons. This is often

represented by pe- (-Log{e-}).

Fe3+ + e- Fe2+ }}{{

}{3

2

eFe

FeKeq

eqKFe

Fee

}{

}{}{

3

2

Here you can see that the value of {e-} will be inversely related to Keq

peo = EhoF/2.3RTpe- is also directly related to Eh

}{

}{2

3

Fe

FeKLogpe eqTake (–Log) to get pe-

which is directly related to Log Keq

From Libes Chap 7 (2nd ed)The peo column in Table 7.1 from Libes is messed up. The numbers in the column are “upside down” (they should be in reverse order; numbers at the bottom should be at the top).

For a 1 electron reduction, Log K = peo.

For a 2 electron reduction peo = ½ Log K, etc.

31

13.5

23

13

21.8

20.8

17.1

5.7

-4.3

3.7

-4.75

2.7

2.4

0

8.8

-7.45

-13

-39.7

-46

Correct peo

A high peo means low electron activity (because peo= -Log{e-}), so rxns with high peo have a tendency to accept rather than donate e-. Those same rxns have high Log K indicating the equilibrium favors reduced form.

Reduced form

Relating Red-Ox chemistry to thermodynamic equilibrium

The free energy change in a chemical system governs whether the reaction is spontaneous and how strong the tendency is to proceed.

For any reaction:G = Gproducts - Greactants

The free energy change for the overall reaction is equal to the free energy of the products minus that of the reactants. The free energy change predicts the tendency of the reaction to proceed.

If free energy of products is less than reactants, then rxn is spontaneous

G = Go + RT ln [(C)c(D)d]/(A)a(B)b]

Go = ƩGfoproducts – ƩGf

oreactants

For a redox reaction, the general free energy equation applies:

Which is the same as:

0= Go + RT ln Keq

At equilibrium, ΔG = 0 and [(C)c(D)d]/(A)a(B)b] = Keq

Go = -2.303RT log Keq

Go = - RT ln Keq

ΔGo = -nFEocell

Go = 2.303 nRT (pe-o2 – pe-o

1)Where peo-

1 is the reduction reaction and peo-

2 is the oxidation reaction

Where Eocell = Eh0

1 – Eho

2 and Eho1 is

the reduction reaction

Thermodynamics and redox chemistry

The reaction with the greatest tendency to proceed spontaneously will be the one with the largest equilibrium constant or most negative ΔG value. For redox reactions, this is achieved by pairing the oxidizing agent with the largest Eh (or peo) to the reducing

agent with the smallest Eh (or peo).

In seawater these chemicals are most often O2 and organic matter (reduced carbon)

The superscript –naught (o) means standard conditions (1 molal activities, temperature of 0 oC, 100 kPa of pressure) The subscript – (water) represents special case of standard conditions but with pH at 7.0 and temperature of 25 oC (typical of natural water rxns).

pe and Eh are directly related to one another, and both to ΔG.

ΔG small

ΔG large

Eh 1

Eh 2

Eh 2

For reactants with:

Eh 1

Comparison of pe and Eh scales at 25 oC

pεo =

Log

K

pεo =

Log

K

Given a world with 21% O2 in the atmosphere, at thermodynamic equilibrium we would expect nearly all:

C to be found as CO2

N as NO3-

S as SO42-

Fe as FeOOH Mn as MnO2

But in fact we find significant amounts of organic matter (reduced carbon -with -C-C- bonds) and NH4

+, N2, and CH4

Some reduced forms of S, Fe and Mn also exist (i.e. R-SH, Fe2+, Mn2+), especially in biological systems

Thus, the world is in disequilibrium!

The large amount of “unstable” reduced

compounds in nature results mainly from

photosynthesis, which takes advantage of light

energy to drive otherwise thermodynamically

unfavorable reactions. (geochemical energy

does so in certain places)

Photosynthesis

CO2 + H2O <=> CH2O + O2

The Go for this reaction as written is +29.9 kcal/mol, so it is not spontaneous. Energy has to be put in to drive this reaction. The energy can come from the sun, or from chemical oxidation of other matter.

Organic matter does not react spontaneously with O2 under most circumstances (at least not on short time scales) because of kinetic factors related to energy of activation.

If you left a peanut butter sandwich in the air long enough there is the chance that it will explode in a puff of smoke - but not likely (activation energy too great).

But what if you provide a spark?

Respiration with O2 is a perfect balance for photosynthesis:

Photosynthesis

CO2 + H2O <=> CH2O + O2 Respiration

So, Why is there oxygen in the air?

Preservation of organic carbon allows excess O2 to accumulate

But, oxidation of all the organic matter in the current biosphere would lower atmospheric oxygen by only 1%.

Therefore, a large amount of “reducing equivalents” must be buried. Most is as organic carbon in sediments, CH4 hydrates, and peat, but some is in the form of reduced sulfur (i.e. FeS2 -pyrite).

Life does not cease when oxygen disappears!

Anaerobic respiration proceeds in the absence of oxygen and uses alternative electron acceptors.

- The electron acceptor that yields the most energy when coupled with the oxidation of organic matter will be used in preference to all others.

This generates a sequence in electron accepting processes that may be revealed in time or space as e- acceptors are depleted in turn (i.e. vertical profiles in sediments). The sequence in terms of energy yield is:

O2 > NO3- > MnO2 > FeOOH > SO4

2- > CO2 (see handout table)

Electron acceptor

Process name Redox reaction Gow

(Kjoule.mol e-1 )

Relative abundance (conc.) of e- acceptor in coastal sediments

O2 Aerobic respiration (oxygen reduction)

reduction of O2 to

H2O

-119 dissolved gas (150-250 µM

NO3- Denitrification reduction of nitrate to

N2

-113 dissolved ion (1-50 µM)

MnO2 Manganese reduction. reduction of Mn(IV) to Mn(II)

-96.9 solid mineral (1-20 mmol.l-1 sediment)

NO3- Nitrate reduction reduction of NO3

- to

NH3 (DNRA)

-82 dissolved ion (1-50 µM)

FeOOH (amorphous)

Iron reduction reduction of ferric iron [Fe(III)] to Ferrous iron [Fe(II)]

-46.7 solid mineral (10-200 mmol.l-1 sediment)

SO42- sulfate reduction reduction of sulfate

(S(VI)) to sulfide (S(-II))

-20.5 dissolved ion (24-28 mM)

CO2 CO2 reduction

(methanogenesis)

reduction of CO2

(C(IV) to CH4 (C(-

IV))

-17.7 dissolved gas (total CO2 = 2-2.5 mM)

H+ Proton reduction reduction of protons (H+) to H2

-1.1 10-7 - 10-8 M; feasible only when H2 is

removed.

Sequence of e- acceptors for the coupled oxidation of organic matter (OM is e- donor)

Vertical segregation of electron accepting processes in sediments and water columns

O2 - aerobic

NO3- - denitrification

MnO2 - Mn oxide

FeO(OH) - Fe oxide

SO42- - Sulfate reduction

CO2 - Methanogenesis

Dep

th

Interface

The source of electron acceptors is typically from above.

The thermodynamically most favorable electron acceptors become depleted at depth as they are used to oxidize organic matter. After depletion of one acceptor the next most favorable one is used, generating the vertical sequence at left.

Organic mattere- acceptors

The concentration or supply rate of a given e- acceptor determines its importance for overall carbon oxidation in a given system

• The concentration of O2 is relatively low, but it is diffusible

• The concentration of nitrate is generally very low, but it can be produced in sediments provided oxygen is sufficient to sustain nitrification

• The concentration of sulfate is very high and it can diffuse

• The concentration of FeOOH is high, but it is a solid and so it cannot diffuse. The same can be said for MnO2. In the case of both Fe and Mn, the oxidized forms are insoluble whereas the reduced forms (Fe2+ and Mn2+) are highly soluble and diffusible.

-25

-20

-15

-10

-5

0

0 50 100

De

pth

(z

)

O2

NO3-

NH4+

Mn2+

Fe2+

CH4

SO42-

HS-

SWI

Relative concentration of electron acceptors or reduced end-products in sediment pore waters

Iron and manganese oxides are insoluble, thus they will be in the solid phase. The reduced Fe(II) and Mn(II) are soluble, hence they appear in the pore water. Both Mn(II) and Fe(II) form insoluble sulfides so they decrease with depth in the sulfide zone.

Ammonium (NH4+) comes

mainly from organic matter degradation – not from NO3

- reduction!

The use of various electron acceptors in sediments results in sharp gradients of e- acceptors and reduced end-products

Libes Chap 12

Example of hypothetical pore water distributions of selected redox species in a coastal sediment

Another view:

Note arbitrary depth scale – the depths over which these gradients occur will vary greatly, depending on several factors, including supply of labile organic matter, temperature (affects rates of respiration), sediment porosity, and availability of the different electron acceptors.

O2 would occur only in the very surface layer

Oxidation of organic carbon (CH2O) with different electron acceptors (stoichiometry for 1 electron transfer)

Complete oxidation of “Redfield” ideal organic matter (C:N:P = 106:16:1) with different electron acceptors.

Note that oxidation of 16 mole of ammonia level N to 16 moles of nitrate requires 32 moles of O2

Different mineral forms of MnO2

Different mineral forms of Fe2O3

Aerobes release NH4

+, which is then oxidized by nitrifying prokaryotes

Modified from Canfield, Thamdrup & Kristensen. Aquatic Geomicrobiology, 2005

The Eoh’s of these

redox couples is directly related to the Go values for coupling organic matter oxidation to the electron acceptors on the left.

e- acceptor

Reduced product

Fe(II)

Aerobic respiration

Denitrification

Manganese reduction

Nitrate reduction to ammonia

Iron reduction

Sulfate reduction CO2 reduction (Methanogenesis)

Proton reductionAcetogenesis

Electron accepting Process

All these reactions are for oxidation of organic matter coupled to reduction of listed electron acceptor

All these reactions are for oxidation of the listed substrate with OXYGEN as the electron acceptor

Photo drivenShould be H2O oxidation which produces O2 in

photosynthesis

Does not occur

Chemo-autotrophic processes

Microbially-mediated Red-Ox reactions

Depth distributions of redox active chemical species in a portion of the water column of the Black Sea near the oxic-anoxic interface.

The σt (density) values represent depth since density increases downward.

From Konovalov et al, 2005

Aerobic respiration

Use of molecular oxygen as electron acceptor in oxidation of organic matter

For sediments at these water column depths

The steeper gradients in shallower sediments is due to more labile organic matter present in those compared to deep sea sedimentsFrom Canfield, Thamdrup & Kristensen. Aquatic Geomicrobiology, 2005

Nitrate reduction (Denitrification)

Next most energetically favorable e- acceptor after O2.

Denitrification reduces nitrate (NO3-) to nitrogen gas

(N2), with net production of small amounts of N2O

Denitrification removes biologically-available nitrogen from the ecosystem

Dentrification occurs in the marine water column in strong oxygen minimum zones, and possibly microzones

Denitrification in estuarine sediments can remove 50% or more of N inputs to estuaries

Oceanic (global) denitrification may control ocean primary production over long time scales i.e. glacial/interglacial

Oxidized metals Fe(III) and Mn(IV) are highly insoluble at pH of seawater and in the presence of O2 they form insoluble oxides (e.g. FeOOH and MnO2)

Metal oxides can be used as e- acceptors by bacteria, but these oxides also are chemically labile

Some of the most primitive of life forms among the Bacteria and Archaea are metal reducers, suggesting a role for metal reduction in early evolution.

Reduced end-products (Fe2+ and Mn2+) are highly soluble under anoxic conditions and therefore diffusible. They are subject to oxidation either chemically or biologically – especially when they reach zones where O2 is around.

Reduction/oxidation cycles of the most abundant metals (typically Fe and Mn) greatly influence the chemistry of other trace metals

Metal oxide reduction (FeOOH and MnO2)

Iron oxide

Manganese oxide

Dissimilatory Sulfate Reduction

Redfield organic matter(CH2O)106 (NH3)16 (H3PO4)1 + 53 SO4

2- --> 106 CO2 + 16 NH3 + 53 S2- + H3PO4 + 106 H2O

note that the sulfide and ammonia stay in reduced forms under anoxic conditions.

2 moles of carbon are oxidized per mole of sulfate reduced.

No free intermediates of oxidation state between +6 and -2 are known to be released during sulfate reduction.

In contrast, many intermediates (So, SO3-, S2O3

- etc) are released during oxidation of sulfide to sulfate.

Sulfate reduction is one of the most important biogeochemical processes responsible for oxidation of organic matter (due to high [SO4

2-) concentration in seawater; ~ 28 mM.

Responsible for ~50% of carbon oxidation in coastal marine sediments

Generates highly reactive sulfide (HS-) and contributes to alkalinity

Sulfide reacts with important metals especially Fe, forming insoluble metal sulfides, thereby greatly affecting metal chemistry

Dissimilatory sulfate reduction dominates the natural sulfur cycle in terms of mass flux. However, most of this is within aquatic systems - exchange of sulfur with atmosphere is primarily via organic sulfur (i.e. DMS, COS etc.)

Biogenesis of methane occurs by two main pathways:

CO2 + 4 H2 CH4 + 2 H2OAutotrophic methanogenesis

H2 is produced during fermentation and other anaerobic processes

CH3COOH CH4 + CO2

Acetate fermentation (acetoclastic methanogenesis)

But can also have Methylotrophic Methanogesis where methylated compounds such as methanol, methylamines and dimethylsulfide (DMS) are converted to CH4 and CO2.

Methanogenesis

Chapter 13, Bianchi.

CO2 is the electron acceptor for CH4 formation but it is generally not limiting in most anoxic waters)

Segregation of methane accumulation from zone of sulfate reduction. When sulfate is depleted, CH4 is produced (if enough organic matter is present to allow methanogenesis)

Various abiological chemical reactions involving electron acceptors can take place

such as

2FeOOH + 3 H2S 2FeS + So + 4 H2O

where iron oxides are reduced by H2S chemically. Similarly, sulfide can react with MnO2 or O2 and become oxidized. In fact, most of the sulfide generated in anoxic environments is reoxidized, either abiotically or biologically.

The electron Bully

e-

e-e-

e-

e-e-e-

e-

My Eh is higher than yours, so give me those

electrons!

Oxygen (O2)

Hydrogen sulfide (H2S)

My pe- is so low that I guess I have no

choice

Oxidation of H2S with O2 occurs spontaneously (abiotically) but is relatively slow, which allows microbes to enzymatically do the same reaction and harness the free energy released

Vertical segregation of electron accepting processes in sediments and water columns

O2 - aerobic

NO3- - denitrification

MnO2 - Mn oxide

FeO(OH) - Fe oxide

SO42- - Sulfate reduction

CO2 - Methanogenesis

Dep

th

Interface

Less energy is gained from the organic matter as less favorable electron acceptors are used

Organic mattere- acceptors

Where did the energy go?

Completing the biogeochemical cycles - oxidation of reduced end-products

The reduced end products of respiration reactions (H2O, N2/NH4, Mn(II), Fe(II), S2-, CH4 and H2) contain “free energy” in amounts inverse to that of the yield from the electron accepting (organic carbon respiration) processes that produced them.

Free Energy Content (when coupled to O2 reduction)

Low High

H2O N2/NH4 Mn(II) Fe(II) S2- CH4 H2

Just as respiration generates oxidized carbon and reduced inorganic chemicals as end-products (H2S, Fe2+, Mn2+, NH4

+ etc), Chemoautotrophy completes the biogeochemical cycles and utilizes the energy in reduced chemicals for the fixation of inorganic carbon (and hence production of reduced carbon biomass for growth of microorganisms).

Examples include:

Sulfide oxidation

H2S + 2O2 SO42- + 2H+

Ammonia Oxidation

NH4+ + O2 NO2

-

NO2- + O2 NO3

-

Methane oxidation

CH4 + 4O2 CO2 + 2 H2O

Iron and manganese oxidation (etc)

Chemoautotrophic processes

Can be coupled with CO2 fixation into biomass

Reactions written for 1 electron transfers. The free energy change would be multiplied by # of electrons per mole of substrate oxidized

ΔGow = -29.95

Oxidation of reduced chemicals by molecular oxygen

Organic matter

Anammox – a recently discovered reaction in the nitrogen cycle. (it is a form of denitrification)

16NH4+ + 16 NO2

- 16 N2 + 32 H2O

• The nitrite comes from the denitrification pathway

• Discovered only in the mid 1990’s!

• Carried out by a unique group of bacteria within the Planctomyces

• Major role in the ocean N cycle identified only in 2003! May account for 15-30% of N2 production

• Occurs in sediments and anoxic water columns (e.g. Black Sea)

Anaerobic oxidation of ammonia

Gradients in electron acceptors can be very sharp – as revealed by microelectrode measurments

Red-ox coupling in sediments – Sulfide oxidation coupled to O2 reduction – via Mn- or Fe-oxide intermediates.

O2 and H2S never meet directly, but H2S oxidation in linked to O2

From Jorgensen

Black smoker vents (photo by Emory Kristof - National Geographic Society

FeSFe2+, Mn2+

HS-

Vent community - sustained by chemoautotrophic sulfide oxidation (photo by Emory Kristof - National Geographic Society

Hemoglobin of tube worms carries both H2S and O2 to bacterial symbionts that oxidize the sulfide with O2.

The Brine Pool is a crater-like depression on the seafloor filled with very concentrated brines coming from the Luann Salt Layer. The brine contains a high concentration of methane gas that supports a surrounding dense mussel bed. (Image based on a mosaic created by Dr. Ian McDonald, Texas A&M University). http://oceanexplorer.noaa.gov/explorations/02mexico/background/brinepool/media/brine_pool.html

Methane oxidation supports chemoautotrophic communities are brine pools in the Gulf of Mexico

Giant bacteria of the genus Thioploca stretch from a coastal sediment to accumulate nitrate from the overlying water column. Nitrate then is transported by the bacteria to several centimeters sediment depth, where it is used to oxidize sulfide. Markus Heuttel – Florida State University

The “Troph” metabolic mode guide

There are prokaryotes that fit each of these modes and some prokaryotes can carry out mixed mode metabolisms.

Eukaryotes are generally chemo-organo-heterotrophs.

Sulfide oxidizers, ammonium oxidizers (nitrifiers) and methane oxidizers are all chemo-litho-autotrophs

Many “heterotrophic” marine bacteria contain proteorhodopsin and can use light energy (but don’t fix CO2) – thus they are photo-organo-heterotrophs

Energy Source Electron Donor Carbon Source

Chemo

Chemo

Photo

Litho

Organo

Autotroph

Heterotroph

(inorganic)

organic)

(light)

(fix CO2)

(incorporate C from organic matter)

Thiomargarita namibiensis

A Chemoautotrophic bacterium that Oxidizes sulfide with NO3

- (or O2) as electron acceptor

Photo by Heide Shultz

Elemental sulfur (S8) granules

Nitrate filled vacuole

This is one of the largest bacteria known – mm in dia!

(Libes)

Eh

[O2]

Anaerobic methane oxidation

Oxic

Anoxic

There are other chemoautotrophic reactions not shown here (e.g. Fe2+ and Mn2+ oxidations)

End

O2 disappears

below 125 m

Saanich Inlet – a fjord with anoxic bottom water

Pore water profiles of O2 and NO3- in shelf sediments

from Antarctic Peninsula (600 m water depth)

O2 goes to zero at 2 to 3 cm depth in sediment – even in these cold waters!

Solid Phase

Solid Phase

Pore water

Pore water

Pore water

Pore water

Sediment profiles from high Arctic shelf sediments

Sta 17 (208 m depth)

Sta 18 (340 m depth)

Fe(II)

Fe(II)

20

10

5

0

-5

-10

15

pe-

0

-0.5

0.5

1.0

Eh (V)

Comparison of pe and Eh scales at 25 oC

pe and Eh are directly related to one another, and both to ΔG.

ΔG small

ΔG large

Eh 1

Eh 2

Eh 1

Eh 2

For reactants with:

The cell potential is directly related to the free energy change for a reaction (though opposite in

sign):

Go = -nF(Eocell)

where Go is the standard free energy change, n = # moles of e- transferred, F is the Faraday constant (23.066 kcal . volts-1 . electron transferred-1).

RTGeq

o

eK /)(

RTnFEeq

cello

eK /)(

Likewise, the free energy change for a reaction is related to the equilibrium constant, and the equilibrium constant is therefore related to the cell potential.

peo = EhoF/2.3RT

“ …it is a convenience of chemistry to express the oxidizing power of an environment in terms of its reduction-oxidation (redox) potential, measured electrically and expressed in volts. It is, in fact, no more than the voltage of a hypothetical battery with one electrode in the oxygen and the other in the food”.

- James E. Lovelock – Gaia: A new look at life on Earth

Marvin-DiPasquale & Capone, 1998

Low sulfate High sulfate

Sulfate depletion @ depth

Sulfide accumulation @ depth

Freshwater Seawater

pe = peo – 1/n log [(D)d(E)e]/(B)b(C)c]

1 NH4+ + 1.32 NO2

- + 0.066 HCO3- + 0.13 H+ = 1.02 N2 + 0.26 NO3

- + 0.066 CH2O0.5N0.15 + 2.03 H2O

Annamox stoichimetry

Respiratory electron chain.

Glucose (or other “food”)

Used to reduce NAD+ to NADH

So, organic matter (food) and O2 don’t react directly together, but electrons are transferred from one to the other via intermediate e- carriers

NADH

This is analogous to O2 serving as the ultimate electron acceptor for most respiration deep in sediments – reducing equivalents work their way back to O2 by diffusion

Element Redox state

OxidationNumber

+6 +5 +4 +3 +2 +1 0 -1 -2 -3 -4

Valence VI V IV III II I 0 -I -II -III -IV

Hydrogen H+ H2

Oxygen O2H2O2 O-R

H2O

Carbon CO2 HCOO (formate)

Co

HCHOCH3OH CH4

Nitrogen NO3- NO2

- N2O N2 NH3

Sulfur SO42- SO3

2- S2O3-2 S8

o H2S

Phosphorous PO43- Phosphonates

(e.g. CH3PO32-)

P PH3

Iron Fe3+, FeO(OH)FeCl3

Fe2+

FeS, FeCO3

Feo

Manganese MnO2 Mn2+

MnS, MnCO3

Mno

Iodine IO3- I2 I-

Selenium SeO42- SeO3

2- Seo H2Se

Copper Cu2+ Cu+ Cuo

Cobalt Co2+ Co+ Coo

Zinc Zn2+ Zno

Cadmium Cd2+ Cdo

Mercury Hg2+ Hgo

Update this table with one in Word file