These points define the standard electrochemical reduction potentials, rep-
resented by E°. A list of standard reduction potentials is given in Table 8.2. You
should know and be able to apply these conventions in order to successfully
work with electrochemistry.
As an aside, it should be pointed out that conventions do change occasion-
ally. It used to be the convention to list half-reactions as oxidationreactions,
not reduction reactions. You may occasionally find an old book or table that
lists half-reactions in that manner, and you should be cautious. Also, the SHE
is not the only possible standard electrode against which other half-reactions
can be measured. Another common one is the saturated calomel electrode,
which is based on the half-reaction
Hg 2 Cl 2 2e^ →2Hg () 2Cl^ (8.24)
E° 0.2682 V versus SHE
(The common name for mercury(I) chloride is calomel.) This half-reaction is
sometimes preferable because it doesn’t use hydrogen gas, which is a potential
explosion hazard. If it is used, then all of the standard reduction potentials are
shifted by 0.2682 V from the standard reduction potentials listed with respect
to SHE.
In order to use the standard potentials for an electrochemical reaction of
interest, simply separate the reaction into its half-reactions, find the standard
potential from a table, reverse one (or more) of the reactions to make it an
oxidation reaction, and negate (that is, change the sign of ) its E° value. A prop-
erly balanced redox reaction has no leftover electrons, so one or more of the
reactions must be multiplied by some integral constant so that the electrons
cancel. However, the E° values are not multiplied by that same constant. E’s are
electric potentials and are intensivevariables, which are defined as independent
of the amount (as opposed to extensivevariables, which are dependent on the
amount).
Finally, standard potentials are strictly additive only for overall electro-
chemical reactions in which there are no unbalanced electrons. If there are un-
balanced electrons in the overall reactions, the E° values are not strictly addi-
tive. Consider as an example the following:Fe^3 3e^ →rxn 1 Fe (s) E°
0.037 VFe (s) →rxn 2 Fe^2 2e^ E° 0.447 VOverall rxn: Fe^3 e^ →overall rxn Fe^2 E° ? 0.410 V
A look at Table 8.2 shows that the reduction reaction Fe^3 e^ →Fe^2 has
an E° of 0.771 V, nowhere close to the predicted 0.410 V.E° values are not ad-
ditive if electrons do not cancel.
However, by Hess’s law,energiesare additive. What must be done for the
above example is to convert each E° into an equivalent G°, add the G°val-
ues together for the overall reaction as allowed by Hess’s law, and then convert
the final G° into a final E° for the new half-reaction. For the above example,
we getRxn 1: G°
(3 mol e^ )96,485
moC
le^( 0.037 V) 10,700 J
Rxn 2: G°
(2 mol e^ )96,485
moC
le^( 0.409 V)
86,300 J
216 CHAPTER 8 Electrochemistry and Ionic Solutions
Table 8.2 Standard reduction potentials
Reaction E°(V)
F 2 2e^ →2F^ 2.866
H 2 O 2 2H^ 2e^ →2H 2 O 1.776
N 2 O 2H+ 2e^ →N 2 H 2 O 1.766
Au^ e^ →Au 1.692
MnO 4
4H^ 3e^ → 1.679
MnO 2 2H 2 O
HClO 2 H^ 3e^ → 1.63
^12 Cl 2 2H 2 O
Mn^3 e^ →Mn^2 1.5415
MnO 4
8H^ 5e^ → 1.507
Mn^2 4H 2 O
Au^3 3e^ →Au 1.498
Cl 2 2e^ →2 Cl^ 1.358
O 2 4H^ 4e^ →2H 2 O 1.229
Br 2 2e^ →2Br^ 1.087
2Hg2+ 2e^ →Hg 22 0.920
Hg^2 2e^ →Hg 0.851
Ag^ e^ →Ag 0.7996
Hg 22 2e^ →2Hg 0.7973
Fe^3 e^ →Fe^2 0.771
MnO 4
e^ →MnO 42
0.558
I 3
2e^ →3I^ 0.536
I 2 2e^ →2I^ 0.5355
Cu^ e^ →Cu 0.521
O 2 2H 2 O 4e^ →4OH^ 0.401
Cu^2 2e^ →Cu 0.3419
Hg 2 Cl 2 2e^ →2Hg 2Cl^ 0.26828
AgCl e^ →Ag Cl^ 0.22233
Cu^2 e^ →Cu^ 0.153
Sn^4 2e^ →Sn^2 0.151
AgBr e^ →Ag Br^ 0.07133
2H^ 2e^ →H 2 0.0000
Fe^3 3e^ →Fe 0.037
2D^ 2e^ →D 2 0.044
Pb^2 2e^ →Pb 0.1262
Sn^2 2e^ →Sn 0.1375
Ni^2 2e^ →Ni 0.257
Co^2 2e^ →Co 0.28
PbSO 4 2e^ →Pb SO 42
0.3588
Cr^3 e^ →Cr^2 0.407
Fe^2 2e^ →Fe 0.447
Cr^3 3e^ →Cr 0.744
Zn^2 2e^ →Zn 0.7618
2H 2 O 2e^ →H 2 2OH^ 0.8277
Cr^2 2e^ →Cr 0.913
Al^3 3e^ →Al 1.662
Be^2 2e^ →Be 1.847
H 2 2e^ →2H^ 2.23
Mg^2 2e^ →Mg 2.372
Na^ e^ →Na 2.71
Ca^2 2e^ →Ca 2.868
Li^ e^ →Li 3.04