Question

Depict the galvanic cell in which the reaction \(\mathrm{Zn}(\mathrm{s})+2 \mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s})\) takes place, Further show: (i) which of the electrode is negatively charged, (ii) the carriers of the current in the cell, and (iii) individual reaction at each electrode.

Short answer

Zn electrode is negative. Current carriers are ions (Zn^2+ and Ag^+). Oxidation occurs at Zn, reduction at Ag.

Step-by-step solution

Write the Overall Reaction

The overall reaction given is \( \text{Zn} (\text{s}) + 2 \text{Ag}^+ (\text{aq}) \rightarrow \text{Zn}^{2+} (\text{aq}) + 2 \text{Ag} (\text{s}) \). This reaction involves the oxidation of Zn to Zn^{2+} and the reduction of Ag^+ to Ag.

Identify the Oxidation and Reduction Half-Reactions

To find which electrode is oxidized or reduced: - **Oxidation half-reaction** involves losing electrons: \( \text{Zn} (\text{s}) \rightarrow \text{Zn}^{2+} (\text{aq}) + 2 \text{e}^- \) - **Reduction half-reaction** involves gaining electrons: \( 2 \text{Ag}^+ (\text{aq}) + 2 \text{e}^- \rightarrow 2 \text{Ag} (\text{s}) \).

Determine the Anode and Cathode

The anode is where oxidation occurs, and it typically has a negative charge. Here, zinc is oxidized: \( \text{Zn} \rightarrow \text{Zn}^{2+} + 2 \text{e}^- \). Thus, the zinc electrode is the anode and negatively charged.The cathode is where reduction occurs, and it typically has a positive charge. Here, silver ions are reduced: \( 2 \text{Ag}^+ + 2 \text{e}^- \rightarrow 2 \text{Ag} \). Thus, the silver electrode is the cathode and positively charged.

Identify the Current Carriers

In the solution, the current is carried by ions. \( \text{Ag}^+ \) ions move toward the cathode where they are reduced, and \( \text{Zn}^{2+} \) ions are released into the solution from the anode.

Describe the Individual Reactions at Each Electrode

- At the **anode (Zn electrode)**: Oxidation of zinc occurs, \( \text{Zn} \rightarrow \text{Zn}^{2+} + 2 \text{e}^- \). - At the **cathode (Ag electrode)**: Reduction of silver ions occurs, \( 2 \text{Ag}^+ + 2 \text{e}^- \rightarrow 2 \text{Ag} \).

Key concepts

Electrode Reactions

In a galvanic cell, electrode reactions are fundamental to its operation. The cell converts chemical energy into electrical energy through redox reactions that occur at the two electrodes. In the example of the zinc-silver galvanic cell, the given overall reaction is \( \mathrm{Zn} (\mathrm{s}) + 2 \mathrm{Ag}^{+} (\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+} (\mathrm{aq}) + 2 \mathrm{Ag} (\mathrm{s}) \).
Here, we have two key processes happening simultaneously:

• **Oxidation** at one electrode, where the oxidation of zinc releases electrons.
• **Reduction** at the other electrode, where silver ions gain electrons to form solid silver.

The flow of electrons from the zinc to the silver electrode is what generates the electric current in the external circuit, making electrode reactions pivotal in powering devices using galvanic cells.
Understanding electrode reactions involves balancing the electron transfer during oxidation and reduction, which perfectly coordinates with the principle of charge conservation.

Oxidation and Reduction

Oxidation and reduction are twin processes that together are known as redox reactions, playing a crucial role in galvanic cells. Oxidation refers to the loss of electrons, while reduction is the gain of electrons.
For the zinc-silver reaction, we determine the oxidation and reduction half-reactions from the overall cell reaction:

• **Oxidation half-reaction** involves zinc as it loses electrons: \( \mathrm{Zn} (\mathrm{s}) \rightarrow \mathrm{Zn}^{2+} (\mathrm{aq}) + 2 \mathrm{e}^- \)
• **Reduction half-reaction** involves silver ions gaining electrons: \( 2 \mathrm{Ag}^{+} (\mathrm{aq}) + 2 \mathrm{e}^- \rightarrow 2 \mathrm{Ag} (\mathrm{s}) \)

These half-reactions show the movement and transfer of electrons, with oxidation always occurring at the anode and reduction at the cathode. Together, they ensure the continuity of the current in the cell by allowing electrons to flow through the external circuit. This interconnected process is critical for the generation of electrical energy from chemical reactions.

Anode and Cathode Identification

Anode and cathode identification is essential in galvanic cells as it dictates the direction of electron flow. The anode is where oxidation occurs, and in our zinc-silver cell, the zinc electrode acts as the anode. Here, zinc goes from its elemental form (\( \mathrm{Zn} \)) to its ionized form (\( \mathrm{Zn}^{2+} \)), releasing electrons.
The cathode, on the other hand, is where reduction takes place. Silver ions (\( \mathrm{Ag}^+ \)) journey to the cathode, gaining electrons to form solid silver. In this scenario:

• The **zinc electrode** becomes negatively charged as it gives up electrons and thus is the **anode**.
• The **silver electrode** accepts the electrons thus becoming positively charged and acts as the **cathode**.

This arrangement supports the flow of electrons from anode to cathode through the external wiring, which can be harnessed as electrical power for various applications. Correctly identifying these electrodes helps predict how the cell will behave and which part will undergo each chemical transformation.

Current Carriers in Electrochemical Cells

Current carriers in electrochemical cells facilitate the movement of electrical charges between electrodes. An understanding of current carriers is key to knowing how electrical neutrality is maintained within the cell's compartments.
In a galvanic cell, this is achieved through the migration of ions in the electrolyte solution:

• **Positive ions** (cations) such as \( \mathrm{Ag}^+ \) move towards the cathode, accepting electrons and getting reduced.
• **Negative ions** (anions) and other ions such as \( \mathrm{Zn}^{2+} \) are released into the solution from the anode as zinc oxidizes.

These ions are the true carriers of the current inside the cell. This ionic movement balances the charges as electrons flow externally from the anode to the cathode, thus maintaining the operation of the galvanic cell continuously.
This dual pathway of electrons through wires and ions through solutions ensures that the redox reactions proceed smoothly, providing a practical utility of chemical reactions in producing power.