D41.2 Using Standard Half-Cell Potentials

Based on the methods described in the preceding section, standard half-cell potentials have been determined for many half-cells. The table below gives half-cell potentials for selected half-cells. Click on the table to open a version that can be enlarged or printed.

There are several important aspects to note:

• The oxidizing agent is on the reactant side of a reduction half-reaction equation.
• The strongest oxidizing agents (the substances most easily reduced) have the largest positive E° values and are at the top of the table. (For example, F₂(g) is a very strong oxidizing agent.)
• The reducing agent is on the product side of a reduction half-reaction equation.
• The strongest reducing agents (the substances most easily oxidized) have the most negative E° values and are at the bottom of the table. (For example, Li(s) and K(s) are very strong reducing agents.)
• A redox reaction is product-favored when a stronger oxidizing agent reacts with a stronger reducing agent. This results in a positive value for E°_cell when the oxidation half-reaction is combined with the reduction half-reaction.
• Half-cell reactions are reversible and the direction a half-cell reaction goes depends on the potential of the other half-cell to which it is connected in a cell.

The table of standard half-cell potentials can be used to determine the E°_cell for any voltaic cell and predict whether a specific redox reaction is product-favored. For example, for the cell:

Because the value of E°_cell is positive, the redox reaction corresponding to this cell notation is product-favored and the voltaic cell can produce electrical energy. The overall reaction can be obtained from the cell notation by writing the oxidation and the reduction half-reactions, multiplying each half-reaction by an appropriate number to balance electrons, and summing the two half-reactions:

Note that:

• Even though the reduction half-reaction is multiplied by 2, the E°_cathode is not multiplied by 2 when E°_cell is calculated. This is because electrical potential is a ratio of energy per coulomb of charge transferred (V = J/C). If the half-reaction is doubled, both energy and charge transferred are doubled, leaving the ratio constant.
• The oxidation half-reaction is the reverse of the reaction as shown in the half-cell potential table, but E°_anode is not multiplied by -1 . This “-1” is accounted for by the fact that E°_anode is subtracted from E°_cathode. (In this way, the cell potential does not depend on the direction of a half-reaction; it depends only on the composition of the half-cell.)

E° for many half-cells are included in the appendix. Tables like these make it possible to determine the E°_cell for many redox reactions. Moreover, by comparing the standard potentials, we can discern which species is easier to reduce/harder to oxidize/a stronger oxidizing agent (higher/more positive E°) and which species is easier to oxidize/harder to reduce/a stronger reducing agent (lower/more negative E°).