[PDF] [PDF] Environmental Standard Reduction Potentials

[PDF] APPENDIX H Standard Reduction Potentials* - CSUN

Standard Potentials in Aqueous Solution (New York: Marcel Dekker, 1985) Reduction potentials for 1 200 free radical reactions are given by P Wardman,

[PDF] Standard Reduction Potentials at 25°C 287 H2O2 + 2 H + 2 e - → 2

Standard Reduction Potentials at 25°C E°r ed (V) F2 (g) + 2e- → 2F- 2 87 H2O2 + 2 H + + 2 e - → 2 Η2O 1 763 PbO2 (s) + 4H + + SO4 2- + 2e

[PDF] Standard Reduction Potentials - Chem21Labs

Standard Reduction Potentials at 25°C Half-Reaction E° (V) Ag + (aq) + e - → Ag (s) +0 799 AgBr (s) + e - → Ag (s) + Br - (aq) +0 095 AgCl (s) + e

[PDF] Standard Reduction Potentials - Ars- Chemia

Standard Reduction Potentials at 25 oC Half-Cell Reactions Eo Half-Cell Reactions Eo F2 + 2 e- 2 F- +2 866 Cu2+ + e- Cu+ +0 153 OF2 + 

[PDF] CRC Handbook of Chemistry and Physics, 89th Edition - Chm Ulaval

Each table lists standard reduction potentials, E° values, at 298 15 K (25°C), in the order of decreasing potential and range from 0 0000 V to –4 10 V The 

[PDF] Environmental Standard Reduction Potentials

Environmental Standard Reduction Potentials We will consider a simple reversible redox reaction for which we are able to measure directly the free energy 

[PDF] Standard Reduction Potential

The standard reduction potential is in a category known as the standard cell potentials or standard electrode potentials The standard cell potential is the 

[PDF] Standard Reduction Potentials:

Standard Reduction Potentials: Half Cell E°(V) Na + (aq) + e - → Na(s) - 2 7144 Y 3+ (aq) + 3e - → Y(s) -2 370 Mg 2+ (aq)+ 2e - → Mg(s) -2 3568 U 3+

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Environmental Standard Reduction Potentials

We will consider a simple reversible redox reaction for which we are able to measure directly the free energy change, G, with a galvanic cell. The redox reaction is the reversible interconversion of 1,4-benzoquinone (BQ) and hydroquinnone (HQ) with a platinum electrode immersed in an aqueous solution buffered at pH 7. This half-cell is coupled to a standard hydrogen electrode (SHE), which is an aqueous solution at pH 0 ([H ] = 1 M) with H 2 (g) bubbling thru at 1 atm and a platinum electrode. The reactions occurring at each electrode are then: H2 (g) ==== 2 H + 2e

OHHOOO+2e2H+ BQ HQ

where we omit the denotion (aq) for dissolved species.

The overall reaction is;

BQ + H

2 (g, 1 atm) + 2H (aq, 10 -7

M) ==== HQ + 2H

(aq, 1 M) and thus 2 272
101
H

P)]M(H][BQ[)]M(H][HQ[lnRTGQlnRTGG

and since [H (1M)] = 1, and = 1, we obtain 2 H P 14

10]BQ[]HQ[lnRTGG

and we can find the cell voltage from:

G = nFE

H where E H is the cell voltage referenced to the SHE.

We can also write:

2]H[]BQ[]HQ[

ln nFRT EE o HH or 2

059160

]H[]BQ[]HQ[logn.EE o HH

In our example,, if [BQ] = [HQ] = [H+

] = and for this reaction, is +0.70 v at 25 C. However we are dealing with redox reactions occurring in natural waters and are more interested in redox potential values that are more representative of typical natural conditions and not solutions at pH = 0 ([H ] = 1 M). We therefore define a value o HH EE o H E (E oH )W (the W indicating conditions typical for natural waters) by setting the pH = 7, [Cl ] = 10 -3

M for a declorination reaction, [Br

] = 10 -5

M for a debromination reaction, [HCO

3 10 -3

M and so on, but by leaving the oxida

nt and reductant at unit activity.

Hence for our reaction;

v./..log..]H[]BQ[]HQ[logn.E)W(E oHoH

280205916014700

1011

2059160700059160

272
In this example, we considered a reversible redox reaction with an overall transfer of two electrons. Since in most abiotic multi-electron redox processes, particularly if organic compounds are involved, the actual electron transfer occurs by a sequence of one- electron transfer steps. There are intermediates formed which are very reactive, that is, they are not stable under environmental conditions. In the example, BQ is first reduced to the corresponding semiquinone (SQ) which is then reduced to HQ:

OHHOOOOHO+e+H+e+H

BQ SQ HQ

Each of these subsequent one-electron steps has its own value, which we can denote and : )W(E oH )W(E H1 )W(E H2

BQ + H

+ e ==== SQ = +0.10 v )W(E H 1

SQ + H

+ e ==== HQ = +0.46 v )W(E H 2 From these values, we see that that the transfer of the first electron to BQ is much less spontaneous (smaller positive voltage), as compared to the transfer of the second electron to SQ. In general, we can assume that the formation of an organic radical is much less favorable from an energetic point of view, as compared to the formation of an organic species exhibiting an even number of electrons. From this we may conclude that the first step of a two-electron transfer between an organic molecule and an electron donor or acceptor is frequently the rate-limiting step. Thus, when we are interested in relating thermodynamic and kinetic data (e.g., thru LFER's), we need to consider primarily the E H values of this rate limiting step, that is, the E H value of the first one-electron transfer. The following table summarizes standard reduction potentials of some environmentally important redox couples. Standard reduction potentials at 25C of some redox couples that are important in natural redox processes

Half-Reaction

oH (w) (v) oH (v) G° H (w) / n (kJ.mol -1 1 O 2 (g) + 4 H + 4e 2 H 2

O +0.81 1.22 78.3

2 2 NO

3 + 12 H + 10e N 2 (g) + 6 H 2

O +0.74 1.24 71.4

3 MnO 2 (s) + HCO 3 (10 -3

M) + 3 H

+ 2e MnCO 3 (s) + 2 H 2

O +0.52

50.2
4 NO 3 + 2 H + 2e NO 2 + H 2

O +0.42 0.83 40.5

5 NO 3 + 10 H + 8e NH 4 + 3 H 2

O +0.36 0.88 34.7

6 FeOOH(s) + HCO

3 (10 -3

M) + 2 H

+ e FeCO 3 (s) + 2 H 2

O 0.05 4.6

7 Pyruvate + 2 H

+ 2e lactate 0.19 18.3 8 SO 42
+ 9 H + 8e HS + 4 H 2

O 0.22 0.25 21.3

9 S(s) + 2 H

+ 2e H 2

S(g) 0.24 0.17 23.5

10 S(s) + H

+ 2e HS 0.27 11 CO 2 (g) + 2 H + 8e CH 4 (g) + 2 H 2

O 0.25 0.17 23.5

12 2 H

+ 2e H 2 (g) 0.41 0 39.6

13 6 CO

2 (g) + 24 H + 24e C 6 H 12 O 6 + 6 H 2

O 0.43 0.01 41.0

Environmental Standard Conditions are taken as; [H ] = 10 -7

M, [Cl

] = [HCO 3 ] = 10 -3

M, [Br

= 10 -5 M It should be pointed out that many of the half-reactions that we consider do not occur reversibly at an electrode surface, so that we would not be able to measure the corresponding E H values using a galvanic cell.

Nevertheless, it is very convenient to

express the free energy change of a half-reaction by assigning the appropriate reduction potentials, that is, E H = G/nF. One possibility is to calculate such reduction potentials from thermodynamic data, for example, from standard free energies of formation, G f of the various species involved in the half-reaction. To illustrate, consider the half-reaction: 2 NO 3 + 12 H + 10e N 2 (g) + 6 H 2 O which is catalyzed by microorganisms and is commonly referred to as denitrification.

From compilations of G

f values, we calculate G rxn G H

1200.6 kJ.mol

-1 (using G f (H G f (e) = 0). oH

E = G

H /nF = +1.24 v.

The Nernst equation then gives us:

62122
3 2

10059160

]OH[P]H[]NO[log.EE No HH

With all species except H

(10 -7

M) at standard conditions, we obtain

v...)(log.E)W(E o Ho H

7405002411010059160

127
Redox Processes That Determine Redox Conditions In The Environment From the data in the table above (Standard Reduction Potentials At 25C Of Some Redox Couples That Are Important In Natural Redox Processes), we can get a general idea about the maximum free energy that microorganisms may gain from catalyzing redox reactions. On earth, the maintena nce of life resulting directly or indirectly from a steady input of solar energy is the main cause for nonequilibrium redox conditions. In the process of photosynthesis, organic compounds exhibiting reduced states of carbon, nitrogen and sulfur are synthesized, and at the same time oxidized species including molecular oxygen, O 2 (oxic photosynthesis) or oxidized sulfur species (anoxic photosynthesis) are produced. Using glucose as a model organic compound, we can express oxic photosynthesis by combining equati ons (1) and (13) from the above table.

Since we are looking at the overall process, it

is convenient to write the reaction with a stoichiometry corresponding to the transfer of one electron (remembering that and (w) are independent of the number of electrons transferred). (w) G° H (w)/n (v) (kJ.mol -1

From equation (13) above;

41
CO 2 + H + e 241
C 6 H 12 O 6 (glucose) + 41
H 2

O 0.43 +41.0

We take the reversed form of equation (1) - changing the sign: 21
H 2 O 41
O 2 (g) + H + e 0.81 +78.3

Overall;

41
CO 2 41
H 2 O 241
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