Standard Electrode Potential - Standard Reduction Potential Table

Standard Reduction Potential Table

Since the values are given in their ability to be reduced, the bigger the standard reduction potentials, the easier they are to be reduced, in other words, they are simply better oxidizing agents. For example, F2 has 2.87 V and Li+ has −3.05 V. F2 reduces easily and is therefore a good oxidizing agent. In contrast, Li(s) would rather undergo oxidation (hence a good reducing agent). Thus Zn2+ whose standard reduction potential is −0.76 V can be oxidized by any other electrode whose standard reduction potential is greater than −0.76 V (e.g. H+(0 V), Cu2+(0.16 V), F2(2.87 V)) and can be reduced by any electrode with standard reduction potential less than −0.76 V (e.g. H2(−2.23 V), Na+(−2.71 V), Li+(−3.05 V)).

In a galvanic cell, where a spontaneous redox reaction drives the cell to produce an electric potential, Gibbs free energy ΔG° must be negative, in accordance with the following equation:

ΔG°cell = −nFE°cell

where n is number of moles of electrons per mole of products and F is the Faraday constant, ~96485 C/mol. As such, the following rules apply:

If E°cell > 0, then the process is spontaneous (galvanic cell)
If E°cell < 0, then the process is nonspontaneous (electrolytic cell)

Thus in order to have a spontaneous reaction (ΔG° < 0), E°cell must be positive, where:

E°cell = E°cathode − E°anode

where E°anode is the standard potential at the anode and E°cathode is the standard potential at the cathode as given in the table of standard electrode potential.

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