Question

Sketch a galvanic cell, and explain how it works. Look at Figs. \(18.1\) and \(18.2 .\) Explain what is occurring in each container and why the cell in Fig. \(18.2\) "works" but the one in Fig. \(18.1\) does not.

Short answer

A galvanic cell consists of two half-cells, each containing an electrode and an electrolyte solution. A salt bridge is essential for maintaining electrical neutrality, as it allows the flow of ions between the half-cells. In this cell, oxidation occurs at the anode, and reduction occurs at the cathode, with electrons flowing from the anode to the cathode, generating an electric current. Figure 18.2 "works" because it includes a salt bridge, unlike Figure 18.1, which lacks one. The absence of a salt bridge in Figure 18.1 leads to charge buildup, inhibiting the redox reactions and ceasing the production of electricity.

Step-by-step solution

Sketching a galvanic cell

A galvanic cell comprises two half-cells, each containing an electrode and an electrolyte solution. Start by drawing two containers, one for each half-cell. Label the containers "Reduction Half-cell" and "Oxidation Half-cell." In each container, draw a metal electrode submerged in an electrolyte solution. Label the electrodes and solutions according to their corresponding chemical reactions. You may use any example of a redox pair, such as placing a zinc electrode (Zn) in a solution of zinc sulfate (ZnSO4) for the oxidation half-cell and a copper electrode (Cu) in a solution of copper sulfate (CuSO4) for the reduction half-cell.

Salt bridge

Draw a salt bridge (a U-shaped tube) connecting the two containers—a necessary component to maintain electrical neutrality. The salt bridge contains an inert electrolyte, such as potassium nitrate (KNO3), which allows the flow of ions between the half-cells.

External wire

Draw a wire connecting the two electrodes (one end connected to the zinc electrode and the other end connected to the copper electrode). The wire will facilitate the flow of electrons from the anode (oxidation electrode) to the cathode (reduction electrode).

Voltmeter

Add a voltmeter to the wire circuit between the two electrodes to measure the potential difference (voltage) generated by the galvanic cell. This will help in determining the electric current produced.

Explain the workings of a galvanic cell

In a galvanic cell, redox reactions occur at the electrodes, which leads to the generation of electricity. At the anode (oxidation half-cell), the metal undergoes oxidation, losing electrons that travel through the external wire towards the cathode (reduction half-cell). At the cathode, the electrons are gained by the ions in the solution, resulting in a reduction process. The flow of electrons from the anode to the cathode generates an electric current, which can be measured by the voltmeter. The salt bridge completes the circuit by maintaining electrical neutrality between the half-cells. It allows the flow of ions between the containers to balance the loss of electrons at the anode and the gain of electrons at the cathode. Now, let's analyze Figures 18.1 and 18.2.

Comparison between Figure 18.1 and Figure 18.2

The main difference between the two figures is the presence (or absence) of a salt bridge. In Figure 18.1, there is no salt bridge connecting the two containers, while in Figure 18.2, a salt bridge is present.

Explain why Figure 18.2 "works" and Figure 18.1 does not

In the absence of a salt bridge, as in Figure 18.1, the galvanic cell cannot maintain electrical neutrality, resulting in the buildup of positive and negative charges in respective half-cells. This charge accumulation opposes the flow of electrons, inhibiting the redox reactions and ceasing the production of electricity. On the other hand, Figure 18.2 "works" because it has a salt bridge that maintains electrical neutrality between the two half-cells. The presence of the salt bridge allows ions to flow between the containers, preventing charge buildup and allowing the redox reactions to proceed. This continuous flow of electrons generates electricity. In conclusion, the primary reason why the galvanic cell in Figure 18.2 "works" and the one in Figure 18.1 does not is due to the presence of a salt bridge in the former, which is essential for maintaining electrical neutrality and enabling the production of electricity.

Key concepts

Redox Reactions

Redox reactions are fundamental to the operation of a galvanic cell. **Redox** stands for reduction-oxidation, which are two simultaneous processes that occur in this type of electrochemical cell. One half of the cell undergoes oxidation, where a metal loses electrons. These electrons are released into the external circuit.
On the other side, in the reduction half-cell, electrons are gained by ions in the electrolyte solution, thereby reducing them to form a stable atom or molecule. This continuous exchange of electrons is what powers the galvanic cell and is often represented by a combination of **half-reactions**: an oxidation half-reaction and a reduction half-reaction.

• Oxidation: Loss of electrons (often written with the electrons as products)
• Reduction: Gain of electrons (often written with the electrons as reactants)

In a common galvanic cell setup, the oxidation process typically occurs at the anode, while the reduction process occurs at the cathode.

Electrodes

A galvanic cell contains two electrodes, each of which plays a crucial role in facilitating redox reactions. **Electrodes** are solid conductors that provide a surface for these reactions to occur, as well as a pathway for the electrons to enter or leave the solution.
The two electrodes in a galvanic cell are distinguished as the *anode* and *cathode*:

• **Anode**: Site of oxidation where electrons are released from the metal into the external circuit.
• **Cathode**: Site of reduction where electrons from the circuit are accepted by ions in the solution.

For example in a typical demonstration, a zinc electrode in a solution of zinc sulfate is often used as the anode. In contrast, a copper electrode immersed in a copper sulfate solution is used as the cathode. This setup allows the movement and transfer of electrons, facilitating redox reactions.

Salt Bridge

The salt bridge is a crucial but often overlooked component of a working galvanic cell. Its main function is to maintain electrical neutrality in each half-cell solution, ensuring the continuous flow of electrons through the circuit. Without it, the charges in each half-cell would accumulate and eventually stop the flow of electrons.
A **salt bridge** is usually a U-shaped tube filled with an inert electrolyte, such as potassium nitrate (KNO3). This design allows ions to migrate between half-cells, balancing the charges by moving negative ions towards the anode compartment and positive ions towards the cathode compartment.
There are several functions of a salt bridge in a galvanic cell:

• Prevents charge buildup by allowing ion exchange between half-cells.
• Completes the electrical circuit, allowing continuous redox reactions.
• Maintains chemical neutrality in each cell compartment.

In summary, a salt bridge ensures the longevity and efficiency of a galvanic cell's operation.

Electric Current

Electric current is an essential product of a galvanic cell, generated when electrons flow through the external circuit. The flow of electrons results from the electrochemical potential difference between the two half-cells. This potential difference is measured in volts and can be observed using a voltmeter connected to the external wire circuit.
**Electric current** specifically refers to the movement of electrons from the anode to the cathode through the external conducting path. In the galvanic cell context, this electron flow is driven by the redox reactions powering the cell. A few key points about electric current in this context include:

• Converted from chemical energy stored in the redox active materials.
• Measured as the movement of electrons through the external wire.
• Represents the cell's ability to perform electrical work.

The efficiency and strength of this current are directly influenced by the reaction rates of the involved redox pairs and the conditions provided by the electrodes and salt bridge.