A voltaic cell is constructed that uses the following...

A voltaic cell is constructed that uses the following half-cell reactions: $$ \begin{aligned} \mathrm{Cu}^{+}(a q)+\mathrm{e}^{-} & \longrightarrow \mathrm{Cu}(s) \\ \mathrm{I}_{2}(s)+2 \mathrm{e}^{-} \longrightarrow & 2 \mathrm{I}^{-}(a q) \end{aligned} $$ The cell is operated at $298 \mathrm{~K}$ with $\left[\mathrm{Cu}^{+}\right]=0.25 \mathrm{M}$ and $\left[\mathrm{I}^{-}\right]=3.5 \mathrm{M} .$ (a) Determine $E$ for the cell at these concentrations. (b) Which electrode is the anode of the cell? (c) Is the answer to part (b) the same as it would be if the cell were operated under standard conditions? (d) If $\left[\mathrm{Cu}^{+}\right]$ was equal to $0.15 \mathrm{M},$ at what concentration of $\mathrm{I}^{-}$ would the cell have zero potential?

A voltaic cell is constructed that uses the following half-cell reactions:
$$
\begin{aligned}
\mathrm{Cu}^{+}(a q)+\mathrm{e}^{-} & \longrightarrow \mathrm{Cu}(s) \\
\mathrm{I}_{2}(s)+2 \mathrm{e}^{-} \longrightarrow & 2 \mathrm{I}^{-}(a q)
\end{aligned}
$$
The cell is operated at $298 \mathrm{~K}$ with $\left[\mathrm{Cu}^{+}\right]=0.25 \mathrm{M}$ and $\left[\mathrm{I}^{-}\right]=3.5 \mathrm{M} .$ (a) Determine $E$ for the cell at these concentrations. (b) Which electrode is the anode of the cell? (c) Is the answer to part (b) the same as it would be if the cell were operated under standard conditions? (d) If $\left[\mathrm{Cu}^{+}\right]$ was equal to $0.15 \mathrm{M},$ at what concentration of $\mathrm{I}^{-}$ would the cell have zero potential?

Key Concepts

Voltaic Cell

A voltaic (or galvanic) cell is an electrochemical cell that converts the energy released from spontaneous redox reactions into electrical energy. It is constructed from two half-cells connected by a salt bridge or porous membrane, allowing ion transfer while electrons flow through an external circuit.

Half-cell Reactions

These are the individual oxidation and reduction reactions that occur at the two electrodes of an electrochemical cell. Each half-cell reaction reveals the transfer of electrons and the species involved, which is critical for determining the overall cell potential.

Anode and Cathode Identification

In any electrochemical cell, the anode is the electrode where oxidation occurs, and the cathode is where reduction takes place. Identifying the correct electrode for each half-reaction is essential for understanding the direction of electron flow and the cell’s operation under various conditions.

Nernst Equation

The Nernst equation is a fundamental tool in electrochemistry that relates the electrode potential to the standard electrode potential and the activities or concentrations of the chemical species involved in the redox reaction. It accounts for non?standard conditions by incorporating temperature, number of electrons transferred, and reaction quotient into the calculation of cell potential.

Concentration Effects on Electrode Potential

The cell potential is influenced by the concentrations of the reactants and products, as demonstrated in the Nernst equation. Changes in ion concentrations affect the reaction quotient, thereby altering the potentials of the half-cells and, consequently, the overall cell potential.

Electrochemical Equilibrium and Zero Cell Potential

This concept refers to the state where there is no net electron transfer between the electrodes because the driving force for the reaction is balanced. When the cell potential is zero, the system is at equilibrium, and adjusting the concentrations of reactive species can bring the cell to this state.

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