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Chapter 18: Problem 26

How can one construct a galvanic cell from two substances, each having a negative standard reduction potential?

Short Answer

Expert verified
To construct a galvanic cell using two substances A and B with negative standard reduction potentials, first determine the anode (Substance A) and cathode (Substance B) based on their reduction potentials. Then, set up the cell with the notation: Substance A | A^n+ || B^n+ | Substance B. Fill two containers with electrolyte solutions of A and B, place electrodes made of each substance in their respective containers, and connect them with a conducting wire including a voltmeter. Finally, connect the containers using a salt bridge to maintain charge balance.

Step by step solution

01

Identify the substances and their standard reduction potentials

For the sake of this example, let's assume we are given two substances, A and B, with their standard reduction potentials: Substance A has a standard reduction potential E潞_A of -1.2 V, and Substance B has a standard reduction potential E潞_B of -0.8 V.
02

Determine the anode and cathode

To form a spontaneous redox reaction, the substance with the higher reduction potential should be the cathode, and the substance with a lower reduction potential should be the anode. In this case, E潞_B (-0.8 V) is higher than E潞_A (-1.2 V). Therefore, Substance B will act as the cathode, and Substance A will be the anode for this galvanic cell.
03

Describe the cell notation

Now that we've identified the anode and cathode, we can write the complete cell notation. The cell notation for this galvanic cell is given by: Anode | Anode electrolyte || Cathode electrolyte | Cathode For our example: Substance A | A^n+ || B^n+ | Substance B
04

Set up the galvanic cell

To set up the galvanic cell, follow these steps: 1. Place two containers (beakers) side-by-side, one for each half-cell. 2. Fill each container with an electrolyte solution that contains the ions of Substance A and Substance B, respectively. 3. Place an electrode made of Substance A in the container with its electrolyte and an electrode made of Substance B in the container with its electrolyte. 4. Connect the two electrodes with a conducting wire, including a voltmeter in the circuit to measure the cell potential. 5. Connect the two containers via a salt bridge (a U-shaped tube filled with an inert electrolyte). This allows the flow of ions between the half-cells and helps in maintaining charge neutrality of the solutions. Now you have successfully constructed the galvanic cell using two substances with negative standard reduction potentials.

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Most popular questions from this chapter

Look up the reduction potential for \(\mathrm{Fe}^{3+}\) to \(\mathrm{Fe}^{2+} .\) Look up the reduction potential for \(\mathrm{Fe}^{2+}\) to Fe. Finally, look up the reduction potential for \(\mathrm{Fe}^{3+}\) to Fe. You should notice that adding the reduction potentials for the first two does not give the potential for the third. Why not? Show how you can use the first two potentials to calculate the third potential.

An experimental fuel cell has been designed that uses carbon monoxide as fuel. The overall reaction is $$2 \mathrm{CO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)$$ The two half-cell reactions are $$\mathrm{CO}+\mathrm{O}^{2-} \longrightarrow \mathrm{CO}_{2}+2 \mathrm{e}^{-}$$ $$\mathrm{O}_{2}+4 \mathrm{e}^{-} \longrightarrow 2 \mathrm{O}^{2-}$$ The two half-reactions are carried out in separate compartments connected with a solid mixture of \(\mathrm{CeO}_{2}\) and \(\mathrm{Gd}_{2} \mathrm{O}_{3}\) . Oxide ions can move through this solid at high temperatures (about \(800^{\circ} \mathrm{C} ) . \Delta G\) for the overall reaction at \(800^{\circ} \mathrm{C}\) under certain concentration conditions is \(-380 \mathrm{kJ}\) . Calculate the cell potential for this fuel cell at the same temperature and concentration conditions.

Calculate \(\mathscr{E}^{\circ}\) values for the following cells. Which reactions are spontaneous as written (under standard conditions)? Balance the equations that are not already balanced. Standard reduction potentials are found in Table 18.1. a. \(\mathrm{H}_{2}(g) \longrightarrow \mathrm{H}^{+}(a q)+\mathrm{H}^{-}(a q)\) b. \(\mathrm{Au}^{3+}(a q)+\mathrm{Ag}(s) \longrightarrow \mathrm{Ag}^{+}(a q)+\mathrm{Au}(s)\)

A zinc-copper battery is constructed as follows at \(25^{\circ} \mathrm{C} :\) $$\mathrm{Zn}\left|\mathrm{Zn}^{2+}(0.10 M)\right|\left|\mathrm{Cu}^{2+}(2.50 M)\right| \mathrm{Cu}$$ The mass of each electrode is 200. g. a. Calculate the cell potential when this battery is first connected. b. Calculate the cell potential after 10.0 A of current has flowed for 10.0 h. (Assume each half-cell contains 1.00 L of solution.) c. Calculate the mass of each electrode after 10.0 h. d. How long can this battery deliver a current of 10.0 A before it goes dead?

It took 150 . s for a current of 1.25 \(\mathrm{A}\) to plate out 0.109 g of a metal from a solution containing its cations. Show that it is not possible for the cations to have a charge of \(1+.\)
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