Final answer:
The anode tends to have a smaller (more negative) standard electrode potential (°E) compared to the cathode. In a spontaneous redox reaction, the cathode's potential is greater than the anode's potential.
Step-by-step explanation:
When determining whether the anode or the cathode has a smaller standard electrode potential (°E), it is essential to recall that the electrode potential is a measure of the tendency of a half-cell to be reduced. By convention, the electrode potentials listed in tables are reduction potentials, even for anodes where oxidation occurs. Therefore, the reaction with the more negative standard reduction potential tends to act as the anode.
In a redox reaction under standard conditions, a spontaneous reaction occurs when °E cell is positive, which corresponds to the condition °Ecathode > °Eanode. This implies that the cathode, which houses the reduction reaction, will have a greater (less negative or more positive) standard reduction potential than the anode. To illustrate this with an example, zinc (°E Zn/Zn²+ = -0.76 V) often serves as an anode to iron (°E Fe/Fe²+ = -0.447 V) because its standard reduction potential is more negative.
A better understanding comes from the use of the Nernst equation, which accounts for reaction conditions like temperature and concentration. The electrode potential will fluctuate during the operation of a cell as the reagent concentrations change, but the initial conditions dictate that a smaller standard electrode potential (°E) is associated with the anode.