Animated Solution for Chemistry - Electrochemistry: Given the electrochemical cell: $E_{{\text{Cu}}^{2+}|\text{Cu}}^{\circ }=+0.34\text{ V}$ $E_{{\text{Zn}}^{2+}|\text{Zn}}^{\circ }=-0.76\text{ V}$ Identify the incorrect statement from the options below for the above cell.

Galvanic vs Electrolytic Cells

Imagine a tug-of-war between two powerful forces. On one side, you have the natural chemical tendency of a Daniell cell trying to push electrons in one direction. On the other side, you have an external battery trying to force them the opposite way. This problem is a classic exploration of what happens when these two forces clash.

Before we can predict the outcome of this battle, we must first determine the natural strength of our Daniell cell. This is known as the standard cell potential, denoted by $E_{\text{cell}}^{\circ }$. The formula is beautifully simple:

We are given the standard reduction potentials for both electrodes. Copper, with a higher reduction potential of $+0.34\text{ V}$, naturally wants to be reduced, making it the cathode. Zinc, with a lower reduction potential of $-0.76\text{ V}$, naturally wants to be oxidized, making it the anode. Substituting these values into our equation:

So, our Daniell cell naturally generates a voltage of $1.10\text{ V}$. This is the force it brings to the tug-of-war.

Now, let's introduce the external opposing voltage, $E_{\text{ext}}$, and analyze the three possible scenarios.

Scenario 1: The Underdog ($E_{\text{ext}}<1.10\text{ V}$) If the external voltage is weaker than the cell's natural voltage, the cell wins. It operates as a normal Galvanic cell. Electrons flow from the zinc anode to the copper cathode. Zinc dissolves into the solution as ${\text{Zn}}^{2+}$ ions, and copper ions from the solution deposit onto the copper electrode.

Scenario 2: The Stalemate ($E_{\text{ext}}=1.10\text{ V}$) If the external voltage perfectly matches the cell's natural voltage, we have a tie. The opposing forces cancel each other out completely. No electrons flow, no chemical reactions occur, and the system is in a state of perfect equilibrium.

Scenario 3: The Overpowering Force ($E_{\text{ext}}>1.10\text{ V}$) This is where things get interesting. If the external voltage is stronger than the cell's natural voltage, it overpowers the cell. The cell is forced to act as an Electrolytic cell. The external battery dictates the flow of electrons, pushing them in the reverse direction—from copper to zinc.

Because the electron flow is reversed, the chemical reactions reverse as well! Copper is now forced to oxidize and dissolve into the solution, acting as the anode. Meanwhile, zinc ions in the solution are forced to accept electrons and deposit onto the zinc electrode, acting as the cathode.

Let's evaluate the given options based on our findings. Option (a) states: "If $E_{\text{ext}}>1.1\text{ V}$, Zn dissolves at Zn electrode and Cu deposits at Cu electrode."

As we discovered in Scenario 3, when $E_{\text{ext}}>1.10\text{ V}$, the exact opposite happens: Copper dissolves and Zinc deposits. Therefore, statement (a) is fundamentally incorrect, making it the right answer to our question.