The Foundations of Chemistry

STANDARD ELECTRODE POTENTIALS

The potentials of the standard zinc–copper and copper–silver voltaic cells are 1.100 volts

and 0.462 volts, respectively. The magnitude of a cell’s potential measures the spontaneity

of its redox reaction. Higher (more positive) cell potentials indicate greater driving force for the

reaction as written.Under standard conditions, the oxidation of metallic Zn by Cu^2 ions

has a greater tendency to go toward completion than does the oxidation of metallic Cu

by Agions. It is convenient to separate the total cell potential into the individual contri-

butions of the two half-reactions. This lets us determine the relative tendencies of

particular oxidation or reduction half-reactions to occur. Such information gives us a quan-

titative basis for specifying strengths of oxidizing and reducing agents. In the next several

sections we shall see how this is done for standard half-cells.

THE STANDARD HYDROGEN ELECTRODE

Every oxidation must be accompanied by a reduction (i.e., the electrons must have some-

where to go). So it is impossible to determine experimentally the potential of any single

electrode. We therefore establish an arbitrary standard. The conventional reference elec-

trode is the standard hydrogen electrode (SHE).This electrode contains a piece of

metal electrolytically coated with a grainy black surface of inert platinum metal, immersed

in a 1 MHsolution. Hydrogen, H 2 , is bubbled at 1 atm pressure through a glass enve-

lope over the platinized electrode (Figure 21-8).

By international convention, the standard hydrogen electrode is arbitrarily assigned

a potential of exactly0.0000... volt.

SHE Half-Reaction E^0 (standard electrode potential)

H 2 88n2H 2 e exactly 0.0000... V (SHE as anode)

2H 2 e88nH 2 exactly 0.0000... V (SHE as cathode)

We then construct a standard cell consisting of a standard hydrogen electrode and some

other standard electrode (half-cell). Because the defined electrode potential of the SHE

contributes exactly 0 volt to the sum, the voltage of the overall cell then lets us determine

the standard electrode potentialof the other half-cell. This is its potential with respect

to the standard hydrogen electrode, measured at 25°C when the concentration of each

ion in the solution is 1 Mand the pressure of any gas involved is 1 atm.

By agreement, we always present the standard cell potential for each half-cell as a

reductionprocess.

THE ZINC–SHE CELL

This cell consists of an SHE in one beaker and a strip of zinc immersed in 1 Mzinc chlo-

ride solution in another beaker (Figure 21-9). A wire and a salt bridge complete the circuit.

When the circuit is closed, the following observations can be made.

21-12

21-11

864 CHAPTER 21: Electrochemistry

See the Saunders Interactive
General Chemistry CD-ROM,
Screen 21.5, Electrochemical Cells and
Potentials.

See the Saunders Interactive
General Chemistry CD-ROM,
Screen 21.6, Standard Potentials.

H 2 gas

1 atm

1 M HCl

Pt black

Figure 21-8 The standard
hydrogen electrode (SHE). A
molecular-level view of the operation
of the SHE as a cathode is shown in
Figure 21-9 and as an anode in
Figure 21-10.

The superscript in E^0 indicates
thermodynamic standard-state
conditions.