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Chapter 17: 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 make a specific galvanic cell, follow these steps: 1. Determine the redox reaction you want to study, which involves both oxidation and reduction half-reactions. 2. Choose suitable electrode materials for each half reaction that can facilitate the redox process and are good conductors of electricity. 3. Select solutions containing ions necessary for each half reaction with suitable concentrations. 4. Connect the compartments using a salt bridge, allowing the flow of ions and maintaining electrical neutrality. 5. Measure the cell potential by connecting the electrodes with a voltmeter or another device to calculate the difference in electrode potentials using the Nernst equation. For example, for the redox reaction \(Zn(s)+Cu^{2+}(aq) \rightarrow Zn^{2+}(aq)+Cu(s)\), use a zinc strip as the anode, a copper strip as the cathode, \(ZnSO_4\) solution in the anode compartment, and \(CuSO_4\) solution in the cathode compartment.

Step by step solution

01

1. Find the Redox Reaction to Study

Determine the specific redox reaction that you are interested in studying or for which you want to create an electrochemical cell. It will involve both oxidation and reduction half-reactions, which will help in determining the appropriate electrodes and solutions to use.
02

2. Determine the Electrode and Half-Cells

For each half-reaction, select an electrode that will facilitate the redox process. The electrode should be a good conductor of electricity and be either involved in the reaction or chemically inert to the redox process. For example, let us consider the following redox reaction: \(Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)\) Here, the half-reactions are: - Oxidation: \(Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-\) - Reduction: \(Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)\) We will use zinc strip as the anode (oxidation electrode) and copper strip as the cathode (reduction electrode).
03

3. Choose the Solutions

For each half-reaction, use a solution that contains the necessary ions to facilitate the reaction. The solution should have a concentration suitable for the desired experiment. In our example, - In the anode compartment, we will use a solution containing zinc ions, such as \(ZnSO_4\). - In the cathode compartment, we will use a solution containing copper ions, such as \(CuSO_4\).
04

4. Salt Bridge

Connect the anode and cathode compartments with a salt bridge containing a neutral and inert electrolyte, which allows the flow of ions between the compartments without physically mixing the solutions. This maintains the electrical neutrality of the cell.
05

5. Measure the Cell Potential

Close the circuit by connecting the electrodes with a voltmeter or another device that measures the cell potential. The cell potential, also known as the cell voltage, is the difference in the electrode potentials of the anode and the cathode. It can be calculated using the Nernst equation: \(E_{cell} = E_{cathode} - E_{anode}\) By following these steps, you can create a specific galvanic cell and decide on the electrodes and solutions to use in the cell.

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Most popular questions from this chapter

Consider a galvanic cell based on the following theoretical half-reactions: $$\begin{array}{lr} & \mathscr{E}^{\circ}(\mathrm{V}) \\ \mathrm{Au}^{3+}+3 \mathrm{e}^{-} \longrightarrow \mathrm{Au} & 1.50 \\ \mathrm{Mg}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Mg} & -2.37 \end{array}$$ a. What is the standard potential for this cell? b. A nonstandard cell is set up at \(25^{\circ} \mathrm{C}\) with \(\left[\mathrm{Mg}^{2+}\right]=\) \(1.00 \times 10^{-5} M .\) The cell potential is observed to be 4.01 V. Calculate \(\left[\mathrm{Au}^{3+}\right]\) in this cell. \mathscr{E}^{\circ}

The ultimate electron acceptor in the respiration process is molecular oxygen. Electron transfer through the respiratory chain takes place through a complex series of oxidationreduction reactions. Some of the electron transport steps use iron-containing proteins called cytochromes. All cytochromes transport electrons by converting the iron in the cytochromes from the +3 to the +2 oxidation state. Consider the following reduction potentials for three different cytochromes used in the transfer process of electrons to oxygen (the potentials have been corrected for \(\mathrm{pH}\) and for temperature): $$\begin{aligned} &\text { cytochrome } \mathrm{a}\left(\mathrm{Fe}^{3+}\right)+\mathrm{e}^{-} \longrightarrow \text { cytochrome } \mathrm{a}\left(\mathrm{Fe}^{2+}\right)\ &\mathscr{E}^{\circ}=0.385 \mathrm{V}\\\ &\text { cytochrome } \mathbf{b}\left(\mathrm{Fe}^{3+}\right)+\mathrm{e}^{-} \longrightarrow \text { cytochrome } \mathrm{b}\left(\mathrm{Fe}^{2+}\right)\ &\mathscr{E}^{\circ}=0.030 \mathrm{V}\\\ &\text { cytochrome } c\left(\mathrm{Fe}^{3+}\right)+\mathrm{e}^{-} \longrightarrow \text { cytochrome } \mathrm{c}\left(\mathrm{Fe}^{2+}\right)\ &\mathscr{E}^{\circ}=0.254 \mathrm{V} \end{aligned}$$ In the electron transfer series, electrons are transferred from one cytochrome to another. Using this information, determine the cytochrome order necessary for spontaneous transport of electrons from one cytochrome to another, which eventually will lead to electron transfer to \(\mathrm{O}_{2}\)

Sketch the galvanic cells based on the following overall reactions. Show the direction of electron flow, and identify the cathode and anode. Give the overall balanced equation. Assume that all concentrations are \(1.0 M\) and that all partial pressures are 1.0 atm. a. \(\mathrm{Cr}^{3+}(a q)+\mathrm{Cl}_{2}(g) \rightleftharpoons \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{Cl}^{-}(a q)\) b. \(\mathrm{Cu}^{2+}(a q)+\mathrm{Mg}(s) \rightleftharpoons \mathrm{Mg}^{2+}(a q)+\mathrm{Cu}(s)\)

The solubility product for \(\operatorname{CuI}(s)\) is \(1.1 \times 10^{-12} .\) Calculate the value of \(\mathscr{E}^{\circ}\) for the half-reaction $$ \operatorname{CuI}(s)+\mathrm{e}^{-} \longrightarrow \mathrm{Cu}(s)+\mathrm{I}^{-}(a q) $$

Given the following two standard reduction potentials, $$\begin{array}{ll} \mathrm{M}^{3+}+3 \mathrm{e}^{-} \longrightarrow \mathrm{M} & \mathscr{E}^{\circ}=-0.10 \mathrm{V} \\ \mathrm{M}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{M} & \mathscr{E}^{\circ}=-0.50 \mathrm{V} \end{array}$$ solve for the standard reduction potential of the half-reaction $$ \mathbf{M}^{3+}+\mathbf{e}^{-} \longrightarrow \mathbf{M}^{2+} $$ (Hint: You must use the extensive property \(\Delta G^{\circ}\) to determine the standard reduction potential.)
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