Chemistry Reference
In-Depth Information
Table 17.2
Standard Reduction Potentials (25
C)
ε
(V)
ε
(V)
F
2
2 e
£ 2 F
Sn
4
2 e
£ Sn
2
2.87
1.82
0.13
0.000
0.13
0.13
0.25
0.44
0.74
0.76
Co
3
e
£ Co
2
2 H
2 e
£ H
2
MnO
4
8 H
5 e
£
Sn
2
2 e
£ Sn
Mn
2
Pb
2
2 e
£ Pb
4 H
2
O
1.51
1.36
Cl
2
2 e
£ 2 Cl
Ni
2
2 e
£ Ni
Cr
2
O
7
2
14 H
6 e
£
Fe
2
2 e
£ Fe
2 Cr
3
Cr
3
e
£ Cr
2
7 H
2
O
1.33
0.80
0.77
0.53
0.34
0.15
Ag
e
£ Ag
Zn
2
2 e
£ Zn
Fe
3
e
£ Fe
2
2 H
2
O
2 e
£
Cu
e
£ Cu
H
2
2 OH
0.828
1.66
2.71
Cu
2
2 e
£ Cu
Al
3
3 e
£ Al
Cu
2
e
£ Cu
Na
e
£ Na
reduction without the other. We can measure the standard cell potential of any
cell. How can we determine the half-cell potential of each half-reaction? We
define
the half-cell potential of the hydrogen/hydrogen ion standard half-cell as
0.000 V, then measure the potential of another half-cell, such as the copper half-
cell, with that half-cell. The entire potential of this cell can then be assigned to
the other (copper) half-cell, because the potential of the hydrogen half-cell is
zero. Of course, we can do the same thing for every other half-cell, but we need
not do so, since the hydrogen/hydrogen ion half-cell is hard to work with
because it involves a gas. We can get an unknown half-cell potential from its
cell potential with any half-cell of known potential. For example, once we get
the copper half-cell potential, we can use it to calculate the unknown zinc half-
cell potential from the Daniell cell potential. A collection of half-cell potentials,
all written as reductions, is presented as Table 17.2.
The following principles allow calculation of cell potentials from reduction
potentials:
1. The two half-cell potentials (one for the oxidation and the other for the
reduction) add up to the cell potential for a complete cell.
2. Writing the chemical equation in the reverse direction requires changing
the sign of the cell or half-cell potential. Note that all the equations in
Table 17.2 refer to reduction half-reactions, but each complete cell requires
one oxidation and one reduction. Thus one of the half-cell equations must
be reversed (and the sign of its potential changed) to add to the other to
make a complete cell equation.
3. Multiplying the coefficients of either type of equation (half-cell or cell)
does
not change
the associated potential.
4. If the sign of a cell potential is positive, the reaction will go as written; if
it is negative, the opposite reaction will proceed.
When an equation is reversed,
the sign of the potential is
changed.
Multiplying the coefficients of
either type of equation (half-
cell or cell) does not change the
associated potential.
EXAMPLE 17.2
Given the half-cell potentials in Table 17.2, calculate the cell potential of
(a) the Daniell cell. (b) the copper/silver cell.
