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Chapter 18: Problem 2

In making a specific galvanic cell, explain how one decides on the electrodes and the solutions to use in the cell.

Short Answer

Expert verified
To decide on the electrodes and solutions for a specific galvanic cell, follow these steps: 1. Choose electrodes made of metals that participate in the desired redox reaction, are stable, and have a positive cell potential (\(E_{cell}\)). 2. Select solutions containing ions involved in the redox reaction, while considering concentration, safety, and availability. 3. Set up the cell by connecting electrodes to a voltmeter, placing them in the solutions, and adding a salt bridge. 4. Observe voltmeter readings and adjust ion concentrations if necessary to ensure the desired redox reaction occurs and provides the desired output.

Step by step solution

01

Understanding Galvanic Cells

A galvanic cell is an electrochemical cell that converts chemical energy into electrical energy through spontaneous redox reactions. The main components of a galvanic cell are two electrodes (an anode and a cathode), an electrolyte solution in which the electrodes are placed, and a salt bridge that allows the flow of ions between the two compartments without mixing the solutions.
02

Choosing the Electrodes

The electrodes are chosen based on the redox reaction we want to occur in the galvanic cell. One of the electrodes will be the anode, where oxidation occurs, and the other will be the cathode, where reduction occurs. When selecting electrodes, consider the following factors: 1. The type of metal: The electrodes should be made of a metal that can participate in the redox reaction. The most common choice is the metal that is a constituent of the species being oxidized and reduced. 2. The standard reduction potential: Using the standard reduction potential values in a table, predict whether the redox reaction will be spontaneous or not. If the cell potential (\(E_{cell}\)) is positive, the reaction is spontaneous. 3. The Stability: The electrodes should be stable and not corrode or dissolve in the solution.
03

Choosing the Solutions

The solutions to be used along with the electrodes in the galvanic cell should contain the ions that are involved in the overall redox reaction. This can be achieved by choosing solutions containing different ions and observing the changes in cell potential until the desired redox reaction occurs. When selecting solutions, consider the following factors: 1. Concentration: Choose the appropriate concentration of the ions to favor the desired redox reaction. 2. Safety: The solutions should be safe to handle and not dangerous, as some chemicals can be highly reactive or toxic. 3. Availability: The chosen solutions should be readily available and cost-effective.
04

Finalizing the Galvanic Cell

Once the electrodes and the solutions are selected, set up the galvanic cell by connecting the electrodes to a voltmeter and placing them in their respective solutions. Add a salt bridge to facilitate ion transfer between the compartments. Finally, observe the voltmeter readings to ensure that the desired redox reaction is occurring, and the cell provides the desired output. Adjust the concentrations of the ions in the solution if necessary.

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Most popular questions from this chapter

Gold is produced electrochemically from an aqueous solution of \(\mathrm{Au}(\mathrm{CN})_{2}^{-}\) containing an excess of \(\mathrm{CN}^{-}\). Gold metal and oxygen gas are produced at the electrodes. What amount (moles) of \(\mathrm{O}_{2}\) will be produced during the production of \(1.00 \mathrm{~mol}\) gold?

The overall reaction and equilibrium constant value for a hydrogen-oxygen fuel cell at \(298 \mathrm{~K}\) is $$2 \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(I) \quad K=1.28 \times 10^{83}$$ a. Calculate \(\mathscr{C}^{\circ}\) and \(\Delta G^{\circ}\) at \(298 \mathrm{~K}\) for the fuel cell reaction. b. Predict the signs of \(\Delta H^{\circ}\) and \(\Delta S^{\circ}\) for the fuel cell reaction. c. As temperature increases, does the maximum amount of work obtained from the fuel cell reaction increase, decrease, or remain the same? Explain.

The overall reaction and standard cell potential at \(25^{\circ} \mathrm{C}\) for the rechargeable nickel-cadmium alkaline battery is \(\mathrm{Cd}(s)+\mathrm{NiO}_{2}(s)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) \(\mathrm{Ni}(\mathrm{OH})_{2}(s)+\mathrm{Cd}(\mathrm{OH})_{2}(s) \quad \mathscr{E}^{\circ}=1.10 \mathrm{~V}\) For every mole of Cd consumed in the cell, what is the maximum useful work that can be obtained at standard conditions?

. When jump-starting a car with a dead battery, the ground jumper should be attached to a remote part of the engine block. Why?

Consider the standard galvanic cell based on the following halfreactions: $$\begin{array}{r}\mathrm{Cu}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Cu} \\ \mathrm{Ag}^{+}+\mathrm{e}^{-} \longrightarrow \mathrm{Ag}\end{array}$$ The electrodes in this cell are \(\mathrm{Ag}(s)\) and \(\mathrm{Cu}(s)\). Does the cell potential increase, decrease, or remain the same when the following changes occur to the standard cell? a. \(\operatorname{CuSO}_{4}(s)\) is added to the copper half-cell compartment (assume no volume change). b. \(\mathrm{NH}_{3}(a q)\) is added to the copper half-cell compartment. [Hint: \(\mathrm{Cu}^{2+}\) reacts with \(\mathrm{NH}_{3}\) to form \(\left.\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}^{2+}(a q) .\right]\) c. \(\mathrm{NaCl}(s)\) is added to the silver half-cell compartment. [Hint: \(\mathrm{Ag}^{+}\) reacts with \(\mathrm{Cl}^{-}\) to form \(\left.\mathrm{AgCl}(s) .\right]\) d. Water is added to both half-cell compartments until the volume of solution is doubled. e. The silver electrode is replaced with a platinum electrode. $$\mathrm{Pt}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Pt} \quad \mathscr{E}^{\circ}=1.19 \mathrm{~V}$$
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