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Chapter: 3 Chemistry
    Section: 3.9 Electrochemistry
        SubSection: 3.9.3 Standard potentials at 25 C

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3.9.3  Standard potentials at 25 °C

The standard electrode potential of an electrode reaction is the standard potential difference (or electromotive force, EMF) of a cell whose left hand electrode is a hydrogen gas electrode. This (IUPAC) convention differs from that used in Latimer’s (1938, 1952) classic book on oxidation potentials in that the sign is opposite, i.e. it is a reduction potential.

For example, on the IUPAC convention (Bard et al., 1985)

 

E0(Zn2+ | Zn)= −0.763 V

The recombination of standard potentials is the basis for their utility. If there are n possible electrodes, the total number of cells that can be made from pairs of them is n(n − 1). The standard potential differences of all these can be calculated from a tabulation of n − 1 cells in which n − 1 electrodes are combined in turn with a chosen electrode. This is called the standard reference electrode and, in aqueous solution, and other protic solvents, is taken as the hydrogen gas electrode. Effectively this means E0(H+ | H2) is set equal to zero at all temperatures.

For example, from the standard electrode potential difference of the cell

 

Pt | H2 | H+ || Zn2+ | Zn    E0 = −0.763V

that of the cell

 

Pt | H2 | H+ || Ag+ | Ag     E0 = 0.799 V

may be subtracted to obtain the standard potential difference of the cell

 

Ag | Ag+ || Zn2+ | Zn     E0 = −1.562V

In doing this subtraction, the cell being subtracted is reversed and added so that the electrode in common, the hydrogen gas electrode, is cancelled. In these cell schemes a single vertical bar represents a phase boundary and a double vertical bar represents a liquid–liquid junction, the potential difference across which is assumed to have been minimised.

In 1982 IUPAC recommended adopting a new standard state pressure 105 Pa (or 1 bar) instead of 101 325 Pa (or 1 atm). However, all the tabulated standard potential data refer to the previous standard state condition of 1 atm. The effect of the change on most of the data is less than 0.2 mV, which in most cases is less than the experimental uncertainty. Correction to 1 bar standard state can be made using the equation:

 

(RT/Fv In (101.325) = 0.169Δv mV

where Δv is the increase in number of gas molecules in the cell reaction.


References

W. M. Latimer (1952) Oxidation States of the Elements and their Potentials in Aqueous Solutions, Prentice-Hall,
    New York.
A. J. Bard, R. Parsons and J. Jordan (eds) (1985) Standard Potentials in Aqueous Solution, Dekker, New York.



Standard potentials at 25 °C

Electrode reaction

    Eº /V

Electrode reaction

Eº /V

Li+ + e → Li        .    .    .    .    .

   − 3.045

  AgI + e → Ag + I−     .    .    .    .    .

   − 0.152 2 

K+ + e → K        .    .    .    .    .

   − 2.925

  Sn2+ + 2e → Sn           .    .    .    .    .

− 0.136 

Rb+ + e → Rb    .    .    .    .    .

   − 2.925

  Pb2+ + 2e → Pb           .    .    .    .    .

   − 0.125 1 

Cs+ + e → Cs     .    .    .    .    .

   − 2.923

  2H+ + 2e → H2            .    .    .    .    .

        0 exactly 

Ba2+ + 2e → Ba     .    .    .    .

 − 2.92

  AgBr + e → Ag + Br    .    .    .    .

   + 0.071 1 

Sr2+ 2e → Sr           .    .    .    .

 − 2.89

  Sn4+ + 2e → Sn2+       .    .    .    .    . 

+ 0.15   

Ca2+ + 2e → Ca    .    .    .    .

 − 2.84

  Cu2+ + e → Cu+     .    .    .    .    .     .

 + 0.159  

Na+ + e → Na        .    .    .    .

   − 2.714

  AgCl + e → Ag + Cl     .    .    .    .  

   + 0.222 3 

La3+ + 3e → La     .    .    .    .

 − 2.37

  Hg2Cl2 + 2e → 2Hg + 2Cl   .    .   

     + 0.268 16 

Mg2+ + 2e → Mg       .    .    .

 − 2.56

  Cu2+ + 2e → Cu     .    .    .    .    .     .

 + 0.340  

Sc3+ + 3e → Sc     .    .    .    .

 − 2.03

  Fe(CN)63− + e → Fe(CN)64−     .

   + 0.361 0 

Be2+ + 2e → Be     .    .    .    .

 − 1.97

  Cu+ + e → Cu     .    .    .    .    .    .     .

 + 0.520  

Th4+ + 4e → Th      .    .    .    .

 − 1.85

  I2 + 2e → 2I      .    .    .    .    .     .    .

   + 0.535 5 

Al3+ + 3e → Al       .    .    .    .

 − 1.67

  I3 + 2e → 3I     .    .    .    .    .    .    .

 + 0.536  

Ti2+ + 2e → Ti        .    .    .    .

 − 1.63

  Hg2SO4 + 2e → 2Hg + SO42−   .

  + 0.613   

Mn2+ + 2e → Mn       .    .    .

 − 1.18

  (AuSCN)4 + 3e → Au + 4SCN

  + 0.636   

Zn2+ + 2e → Zn     .    .    .    .

      − 0.762 6

  Fe3+ + e → Fe2+     .    .    .    .    .    .

 + 0.771  

Ga3+ + 3e → Ga    .    .    .    .

   − 0.529

  Hg22+ + 2e → 2Hg      .    .    .    .    .

   + 0.796 0 

Fe2+ + 2e → Fe     .    .    .    .

 − 0.44

  Ag+ + e → Ag     .    .    .    .    .    .    .

   + 0.799 1 

Cr3+ + e → Cr2+    .    .    .    .

   − 0.424

  Hg2+ +2e → Hg22+     .    .    .    .    .

   + 0.911 0 

Cd2+ + 2e → Cd    .    .    .    .

      − 0.042 5

  Pd2+ + 2e → Pd      .    .    .    .    .    .

+ 0.915 

Ti3+ + e → Ti2+        .    .    .    .

 − 0.37

  AuCl4 + 3e → Au + 4Cl      .    .  

+ 1.002 

PbSO4 + 2e → Pb + SO42−

      − 0.350 5

  Pu4+ + e → Pu3+     .    .    .    .    .    .

+ 1.01   

In3+ + 3e → In        .    .    .    .

      − 0.338 2

  Br2(l) + 2e → 2Br      .    .    .    .    .

+ 1.065 

Tl+ + e → Tl        .    .    .    .    .

      − 0.336 3

  O2 + 4H+ + 4e → 2H2O     .    .    . 

+ 1.229 

Co2+ + 2e → Co    .    .    .    .

   − 0.277

  Tl3+ + 2e → Tl+        .    .    .    .    .    .

+ 1.25   

V3+ + e → V2+       .    .    .    .

   − 0.255

  Cl2 + 2e → 2Cl       .    .    .    .    .    .

   + 1.358 3 

Ni2+ + e → Ni   .    .    .    .    .

   − 0.257

  Au3+ + 3e → Au        .    .    .    .    .    .

+ 1.52   

A.K. Covington

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