A chemist assembles the following electrochemical cell. One side of the cell contains a zinc electrode immersed in aqueous zinc nitrate solution. The other side of the cell contains an iron electrode immersed in aqueous iron(II) nitrate solution. In this reaction, zinc is the anode and iron is the cathode. Write equations showing the half-reactions that take place at each electrode. Include phase symbols. Anode half-reaction: Zn(s) -> Zn2+(aq) + 2e- Cathode half-reaction: Fe2+(aq) + 2e- -> Fe(s)
The Correct Answer and Explanation is:
Correct Half-Reactions with Phase Symbols:
Anode half-reaction: Zn(s)→Zn2+(aq)+2e−\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^-
Cathode half-reaction: Fe2+(aq)+2e−→Fe(s)\text{Fe}^{2+}(aq) + 2e^- \rightarrow \text{Fe}(s)
Explanation
In this electrochemical cell, a redox reaction occurs between zinc and iron, converting chemical energy into electrical energy. This setup is known as a galvanic cell or voltaic cell, in which a spontaneous redox reaction drives the flow of electrons through an external circuit.
Each side of the cell is known as a half-cell, consisting of a metal electrode immersed in a solution containing its ions. Here, zinc is in zinc nitrate solution and iron is in iron(II) nitrate solution.
The anode is where oxidation occurs. Zinc serves as the anode, and its atoms lose electrons to form zinc ions: Zn(s)→Zn2+(aq)+2e−\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^-
This means solid zinc is oxidized to aqueous zinc ions, and electrons are released into the external circuit.
The cathode is where reduction occurs. Iron(II) ions in solution accept electrons to form solid iron, which plates onto the iron electrode: Fe2+(aq)+2e−→Fe(s)\text{Fe}^{2+}(aq) + 2e^- \rightarrow \text{Fe}(s)
Thus, iron ions are reduced to metallic iron.
Electrons flow from the zinc electrode (anode) through the wire to the iron electrode (cathode), completing the external circuit. Inside the cell, ions move through a salt bridge to maintain electrical neutrality. Nitrate ions (NO₃⁻) likely move toward the zinc side to balance the Zn²⁺ ions entering solution, while Zn²⁺ ions do not pass through the salt bridge.
This reaction is spontaneous because zinc is higher on the activity series than iron, meaning it more readily loses electrons (is oxidized). The overall cell reaction is: Zn(s)+Fe2+(aq)→Zn2+(aq)+Fe(s)\text{Zn}(s) + \text{Fe}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Fe}(s)
In summary, oxidation of zinc and reduction of iron ions together drive the electrochemical process.
