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Chapter 20: Problem 111

Of the following statements concerning electrochemical cells, the correct ones are: (a) The cathode is the negative electrode in both voltaic and electrolytic cells. (b) The function of a salt bridge is to permit the migration of electrons between the half-cell compartments of an electrochemical cell. (c) The anode is the negative electrode in a voltaic cell. (d) Electrons leave the cell from either the cathode or the anode, depending on what electrodes are used. (e) Reduction occurs at the cathode in both voltaic and electrolytic cells. (f) If electric current is drawn from a voltaic cell long enough, the cell becomes an electrolytic cell. (g) The cell reaction is an oxidationreduction reaction.

Short Answer

Expert verified
The correct statements are: (c) The anode is the negative electrode in a voltaic cell, (e) Reduction occurs at the cathode in both voltaic and electrolytic cells, and (g) The cell reaction is an oxidation-reduction reaction.

Step by step solution

01

Evaluate statement (a)

The cathode is the positive electrode in a voltaic cell but becomes the negative one in an electrolytic cell when an external voltage is applied. So, statement (a) is incorrect.
02

Evaluate statement (b)

A salt bridge does not permit the migration of electrons but of ions, maintaining neutrality in the half-cells. Thus, the statement (b) is incorrect.
03

Evaluate statement (c)

The anode is the negative electrode in a voltaic cell, so statement (c) is correct.
04

Evaluate statement (d)

Regardless of the electrodes used, electrons always enter the cell at the cathode and leave the cell at the anode. So, statement (d) is wrong.
05

Evaluate statement (e)

Reduction always occurs at the cathode, whether it is a voltaic or electrolytic cell. Hence, statement (e) is correct.
06

Evaluate statement (f)

Drawing electrical current from a voltaic cell for a long time does not make it an electrolytic cell. Therefore, statement (f) is incorrect.
07

Evaluate statement (g)

The cell reaction in both voltaic and electrolytic cells is always a redox (oxidation-reduction) reaction. Thus, statement (g) is correct.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Voltaic and Electrolytic Cells
Understanding the differences between voltaic (galvanic) and electrolytic cells is crucial for mastering electrochemistry. Voltaic cells convert chemical energy into electrical energy using spontaneous redox reactions. They are the basis for batteries, where the chemical reaction creates a flow of electrons from the anode to the cathode through an external circuit.

Electrolytic cells, on the other hand, require an external source of electricity to induce non-spontaneous chemical reactions. This process is commonly used in electroplating, where a metal is deposited onto an electrode. The key distinction lies in their operating principles: voltaic cells harness spontaneous reactions, while electrolytic cells drive non-spontaneous reactions with external power.

They share similar setups with two electrodes, an electrolyte solution, and sometimes a salt bridge to maintain charge balance by allowing ions, not electrons, to move between compartments. It's important to remember that electrons never physically pass through the salt bridge.
Cathode and Anode in Electrochemistry
In electrochemical cells, electrodes play vital roles and are known as the cathode and anode. The anode is where oxidation occurs; it loses electrons. Conversely, at the cathode, reduction takes place as it gains electrons.

Determining whether an electrode is positive or negative depends on the cell type. In a voltaic cell, the anode is negative since the cell works based on spontaneous reactions, pushing electrons out of the anode through the circuit. The cathode is positive, accepting these electrons to complete the circuit. However, this polarity is reversed in an electrolytic cell because external energy is provided to drive the reactions. Here the cathode is negative, receiving electrons from the external source, while the anode is positive. Keeping these distinctions clear is essential for understanding the flow of electrons and the functioning of electrochemical cells.
Redox (Oxidation-Reduction) Reactions
Redox reactions are at the heart of electrochemical cells. These reactions involve the transfer of electrons between chemical species. When a substance loses electrons, we term it oxidation; when it gains electrons, it's called reduction.

These reactions are coupled鈥攚henever one species is oxidized, another is simultaneously reduced. This electron transfer process is what drives the current in voltaic cells and is forced by an external power source in electrolytic cells. An easy way to remember the distinction is the mnemonic 'OIL RIG': Oxidation Is Loss, Reduction Is Gain.

In voltaic cells, the spontaneous redox reaction generates electric current, while in electrolytic cells, the applied current induces a redox reaction. Both types of cells are fundamental in various technologies, from powering portable electronics to synthesizing valuable chemicals.

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Most popular questions from this chapter

Natural gas transmission pipes are sometimes protected against corrosion by the maintenance of a small potential difference between the pipe and an inert electrode buried in the ground. Describe how the method works.

A common reference electrode consists of a silver wire coated with \(\mathrm{AgCl}(\mathrm{s})\) and immersed in \(1 \mathrm{M} \mathrm{KCl}\) $$\mathrm{AgCl}(\mathrm{s})+\mathrm{e}^{-} \longrightarrow \mathrm{Ag}(\mathrm{s})+\mathrm{Cl}^{-}(1 \mathrm{M}) E^{\circ}=0.2223 \mathrm{V}$$ (a) What is \(E_{\text {cell }}^{\circ}\) when this electrode is a cathode in combination with a standard zinc electrode as an anode? (b) Cite several reasons why this electrode should be easier to use than a standard hydrogen electrode. (c) By comparing the potential of this silver-silver chloride electrode with that of the silver-silver ion electrode, determine \(K_{\mathrm{sp}}\) for \(\mathrm{AgCl}\).

Construct a concept map illustrating the principles of electrolysis and its industrial applications.

Only a tiny fraction of the diffusible ions move across a cell membrane in establishing a Nernst potential (see Focus On 20: Membrane Potentials), so there is no detectable concentration change. Consider a typical cell with a volume of \(10^{-8} \mathrm{cm}^{3},\) a surface area \((A)\) of \(10^{-6} \mathrm{cm}^{2},\) and a membrane thickness \((l)\) of \(10^{-6} \mathrm{cm}\) Suppose that \(\left[\mathrm{K}^{+}\right]=155 \mathrm{mM}\) inside the cell and \(\left[\mathrm{K}^{+}\right]=4 \mathrm{mM}\) outside the cell and that the observed Nernst potential across the cell wall is \(0.085 \mathrm{V}\). The membrane acts as a charge-storing device called a capacitor, with a capacitance, \(C,\) given by $$C=\frac{\varepsilon_{0} \varepsilon A}{l}$$ where \(\varepsilon_{0}\) is the dielectric constant of a vacuum and the product \(\varepsilon_{0} \varepsilon\) is the dielectric constant of the membrane, having a typical value of \(3 \times 8.854 \times 10^{-12}\) \(\mathrm{C}^{2} \mathrm{N}^{-1} \mathrm{m}^{-2}\) for a biological membrane. The SI unit of capacitance is the firad, \(1 \mathrm{F}=1\) coulomb per volt \(=1 \mathrm{CV}^{-1}=1 \times \mathrm{C}^{2} \mathrm{N}^{-1} \mathrm{m}^{-1}\) (a) Determine the capacitance of the membrane for the typical cell described. (b) What is the net charge required to maintain the observed membrane potential? (c) How many \(\mathrm{K}^{+}\) ions must flow through the cell membrane to produce the membrane potential? (d) How many \(\mathrm{K}^{+}\) ions are in the typical cell? (e) Show that the fraction of the intracellular \(K^{+}\) ions transferred through the cell membrane to produce the membrane potential is so small that it does not change \(\left[\mathrm{K}^{+}\right]\) within the cell.

Using the method presented in Appendix \(\mathrm{E}\), construct a concept map showing the relationship between electrochemical cells and thermodynamic properties.
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