Consider the following table of standard electrode potentials for a series of hypothetical reactions in aqueous solution: $$ \begin{array}{lr} \hline \text { Reduction Half-Reaction } & \multicolumn{1}{c} {E^{\circ}(\mathbf{V})} \\ \hline \mathrm{A}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{A}(s) & 1.33 \\\ \mathrm{~B}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{B}(s) & 0.87 \\\ \mathrm{C}^{3+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{C}^{2+}(a q) & -0.12 \\ \mathrm{D}^{3+}(a q)+3 \mathrm{e}^{-} \longrightarrow \mathrm{D}(s) & -1.59 \\\ \hline \end{array} $$ (a) Which substance is the strongest oxidizing agent? Which is weakest? (b) Which substance is the strongest reducing agent? Which is weakest? (c) Which substance(s) can oxidize \(\mathrm{C}^{2+}\) ? [Sections 20.4 and 20.5\(]\)

Understanding the role of oxidizing and reducing agents in electrochemistry is pivotal for analyzing reactions and determining their feasibility. An oxidizing agent, often called an oxidant, gains electrons in a reaction. As it accepts electrons, it causes another species to be oxidized and is therefore reduced itself. Conversely, a reducing agent, or reductant, donates electrons and thereby reduces the other substance and is itself oxidized.

In the given exercise where we have a series of standard electrode potentials for hypothetical reactions, the substances with the higher positive electrode potentials, such as A⁺(aq) with an E⁰ value of 1.33 V, are strong oxidizers because they have a greater tendency to gain electrons. In contrast, substances with negative standards potentials, like D³⁺(aq) with -1.59 V, are strong reducers because they are more inclined to donate electrons. This tendency is crucial for predicting the direction of electron flow and the spontaneity of redox reactions.

At the heart of electrochemical processes are half-reactions. These are the two parts of a whole redox reaction, where one substance loses electrons (oxidation) and another gains them (reduction). Electrochemistry splits these into reduction half-reactions, showing the gain of electrons, and oxidation half-reactions, showing the loss.

In the table provided in the exercise, the reduction half-reactions are listed with their associated standard electrode potentials. To analyze an oxidation half-reaction, one must reverse the reduction half-reaction and change the sign of the standard potential, as seen in the solution step discussing reducing agents. For example, the reduction half-reaction for A is A⁺(aq) + e⁻ → A(s), and the corresponding oxidation half-reaction is A(s) → A⁺(aq) + e⁻ with an E⁰ that is the negative of the original (from +1.33 V to -1.33 V). Half-reactions are integral to understanding the electron transfer process and are essential when balancing redox reactions.

The electrochemical cell potential, represented by E⁰, measures the ability of a reaction to move electrons—that is, its propensity to occur spontaneously as an electrochemical reaction. These potentials are given under standard conditions, typically designated at 1 M concentration, 1 atm pressure, and 298 K (25 °C). A positive E⁰ indicates a reaction that tends to occur spontaneously, while a negative E⁰ suggests non-spontaneity under standard conditions.

To determine whether a specific reaction, like the oxidation of \(\mathrm{C}^{2+}\), can occur, we look for other reaction partners with higher reduction potentials. In our exercise, we compare the reduction potential of \(\mathrm{C}^{2+}\) with the reduction potentials of other species. A⁺(aq) and B²⁺(aq) both have higher standard potentials than the reverse oxidation of \(\mathrm{C}^{2+}\); thus, they can oxidize it, resulting in an overall cell potential which is positive and signifies a spontaneous reaction.