18.3 Galvanic Cells

Learning Objectives

By the end of this section, you will be able to:
  • Use cell notation to describe galvanic cells
  • Describe the basic components of galvanic cells

Galvanic cells, also known as voltaic cells, are electrochemical cells in which spontaneous oxidation-reduction reactions produce electrical energy. In writing the equations, it is often convenient to separate the oxidation-reduction reactions into half-reactions to facilitate balancing the overall equation and to emphasize the actual chemical transformations.

Consider what happens when a clean piece of copper metal is placed in a solution of silver nitrate (Figure 18.3a). As soon as the copper metal is added, silver metal begins to form and copper ions pass into the solution. The blue colour of the solution on the far right indicates the presence of copper ions. The reaction may be split into its two half-reactions. Half-reactions separate the oxidation from the reduction, so each can be considered individually.

[latex]\begin{array}{lr @{{}\longrightarrow{}} l} \text{oxidation:} & \text{Cu}(s) & \longrightarrow \text{Cu}^{2+}(aq)\;+\;2\text{ e}^{-} \\[0.5em] \text{reduction:} & 2\;\times\;[\text{Ag}^{+}(aq)\;+\;\text{e}^{-} & \longrightarrow \text{Ag}(s)]\;\;\;\;\;\;\;\text{or}\;\;\;\;\;\;\;2\text{ Ag}^{+}(aq)\;+\;2\text{ e}^{-}\;{\longrightarrow}\;\text{Ag}(s) \\[0.5em] \hline \\[-0.25em] \text{overall:} & 2\text{ Ag}^{+}(aq)\;+\;\text{Cu}(s) & \longrightarrow 2\text{ Ag}(s)\;+\;\text{Cu}^{2+}(aq) \end{array}[/latex]

The equation for the reduction half-reaction had to be doubled so the number electrons “gained” in the reduction half-reaction equaled the number of electrons “lost” in the oxidation half-reaction.

This figure includes three photographs. In a, a test tube containing a clear, colorless liquid is shown with a loosely coiled copper wire outside the test tube to its right. In b, the wire has been submerged into the clear colorless liquid in the test tube and the surface of the wire is darkened. In c, the liquid in the test tube is a bright blue-green color, the wire in the solution appears dark near the top, and a grey “fuzzy” material is present at the bottom of the test tube on the lower portion of the copper coil, giving a murky appearance to the liquid near the bottom of the test tube.
Figure 18.3a: When a clean piece of copper metal is placed into a clear solution of silver nitrate (a), an oxidation-reduction reaction occurs that results in the exchange of Cu2+ for Ag+ ions in solution. As the reaction proceeds (b), the solution turns blue (c) because of the copper ions present, and silver metal is deposited on the copper strip as the silver ions are removed from solution. (credit: modification of work by Mark Ott in Chemistry (OpenStax), CC BY 4.0).)

Galvanic or voltaic cells involve spontaneous electrochemical reactions in which the half-reactions are separated (Figure 18.3b) so that current can flow through an external wire. The beaker on the left side of the figure is called a half-cell, and contains a 1 M solution of copper(II) nitrate [Cu(NO3)2] with a piece of copper metal partially submerged in the solution. The copper metal is an electrode. The copper is undergoing oxidation; therefore, the copper electrode is the anode. The anode is connected to a voltmeter with a wire and the other terminal of the voltmeter is connected to a silver electrode by a wire. The silver is undergoing reduction; therefore, the silver electrode is the cathode. The half-cell on the right side of the figure consists of the silver electrode in a 1 M solution of silver nitrate (AgNO3). At this point, no current flows—that is, no significant movement of electrons through the wire occurs because the circuit is open. The circuit is closed using a salt bridge, which transmits the current with moving ions. The salt bridge consists of a concentrated, nonreactive, electrolyte solution such as the sodium nitrate (NaNO3) solution used in this example. As electrons flow from left to right through the electrode and wire, nitrate ions (anions) pass through the porous plug on the left into the copper(II) nitrate solution. This keeps the beaker on the left electrically neutral by neutralizing the charge on the copper(II) ions that are produced in the solution as the copper metal is oxidized. At the same time, the nitrate ions are moving to the left, sodium ions (cations) move to the right, through the porous plug, and into the silver nitrate solution on the right. These added cations “replace” the silver ions that are removed from the solution as they were reduced to silver metal, keeping the beaker on the right electrically neutral. Without the salt bridge, the compartments would not remain electrically neutral and no significant current would flow. However, if the two compartments are in direct contact, a salt bridge is not necessary. The instant the circuit is completed, the voltmeter reads +0.46 V, this is called the cell potential. The cell potential is created when the two dissimilar metals are connected, and is a measure of the energy per unit charge available from the oxidation-reduction reaction. The volt is the derived SI unit for electrical potential

[latex]\text{volt} = V = \frac{\text{J}}{\text{C}}[/latex]

In this equation C is the charge in coulombs. Note that volts must be multiplied by the charge in coulombs (C) to obtain the energy in joules (J).

This figure contains a diagram of an electrochemical cell. Two beakers are shown. Each is just over half full. The beaker on the left contains a blue solution and is labeled below as “1 M solution of copper (II) nitrate ( C u ( N O subscript 3 ) subscript 2 ).” The beaker on the right contains a colorless solution and is labeled below as “1 M solution of silver nitrate ( A g N O subscript 3 ).” A glass tube in the shape of an inverted U connects the two beakers at the center of the diagram. The tube contents are colorless. The ends of the tubes are beneath the surface of the solutions in the beakers and a small grey plug is present at each end of the tube. The plug in the left beaker is labeled “Porous plug.” At the center of the diagram, the tube is labeled “Salt bridge ( N a N O subscript 3 ). Each beaker shows a metal strip partially submerged in the liquid. The beaker on the left has an orange brown strip that is labeled “C u anode negative” at the top. The beaker on the right has a silver strip that is labeled “A g cathode positive” at the top. A wire extends from the top of each of these strips to a rectangular digital readout indicating a reading of positive 0.46 V that is labeled “Voltmeter.” An arrow points toward the voltmeter from the left which is labeled “Flow of electrons.” Similarly, an arrow points away from the voltmeter to the right which is also labeled “Flow of electrons.” A curved arrow extends from the C u strip into the surrounding solution. The tip of this arrow is labeled “C u superscript 2 plus.” A curved arrow extends from the salt bridge into the beaker on the left into the blue solution. The tip of this arrow is labeled “N O subscript 3 superscript negative.” A curved arrow extends from the solution in the beaker on the right to the A g strip. The base of this arrow is labeled “A g superscript plus.” A curved arrow extends from the colorless solution to salt bridge in the beaker on the right. The base of this arrow is labeled “N O subscript 3 superscript negative.” Just right of the center of the salt bridge on the tube an arrow is placed on the salt bridge that points down and to the right. The base of this arrow is labeled “N a superscript plus.” Just above this region of the tube appears the label “Flow of cations.” Just left of the center of the salt bridge on the tube an arrow is placed on the salt bridge that points down and to the left. The base of this arrow is labeled “N O subscript 3 superscript negative.” Just above this region of the tube appears the label “Flow of anions.”
Figure 18.3b: In this standard galvanic cell, the half-cells are separated; electrons can flow through an external wire and become available to do electrical work (credit: Chemistry (OpenStax), CC BY 4.0).

When the electrochemical cell is constructed in this fashion, a positive cell potential indicates a spontaneous reaction and that the electrons are flowing from the left to the right. There is a lot going on in Figure 18.3b, so it is useful to summarize things for this system:

  • Electrons flow from the anode to the cathode: left to right in the standard galvanic cell in the figure.
  • The electrode in the left half-cell is the anode because oxidation occurs here. The name refers to the flow of anions in the salt bridge toward it.
  • The electrode in the right half-cell is the cathode because reduction occurs here. The name refers to the flow of cations in the salt bridge toward it.
  • Oxidation occurs at the anode (the left half-cell in the figure).
  • Reduction occurs at the cathode (the right half-cell in the figure).
  • The cell potential, +0.46 V, in this case, results from the inherent differences in the nature of the materials used to make the two half-cells.
  • The salt bridge must be present to close (complete) the circuit and both an oxidation and reduction must occur for current to flow.

There are many possible galvanic cells, so a shorthand notation is usually used to describe them. The cell notation (sometimes called a cell diagram) provides information about the various species involved in the reaction. This notation also works for other types of cells. A vertical line, │, denotes a phase boundary and a double line, ‖, the salt bridge. Information about the anode is written to the left, followed by the anode solution, then the salt bridge (when present), then the cathode solution, and, finally, information about the cathode to the right. The cell notation for the galvanic cell in Figure 18.3b is then:

[latex]\text{Cu}(s){\mid}\text{Cu}^{2+}(aq\text{,}\;1\;M){\parallel}\text{Ag}^{+}(aq\text{,}\;1\;M){\mid}\text{Ag}(s)[/latex]

Note that spectator ions are not included and that the simplest form of each half-reaction was used. When known, the initial concentrations of the various ions are usually included.

One of the simplest cells is the Daniell cell. It is possible to construct this battery by placing a copper electrode at the bottom of a jar and covering the metal with a copper sulfate solution. A zinc sulfate solution is floated on top of the copper sulfate solution; then a zinc electrode is placed in the zinc sulfate solution. Connecting the copper electrode to the zinc electrode allows an electric current to flow. This is an example of a cell without a salt bridge, and ions may flow across the interface between the two solutions.

Some oxidation-reduction reactions involve species that are poor conductors of electricity, and so an electrode is used that does not participate in the reactions. Frequently, the electrode is platinum, gold, or graphite, all of which are inert to many chemical reactions. One such system is shown in Figure 18.3c. Magnesium undergoes oxidation at the anode on the left in the figure and hydrogen ions undergo reduction at the cathode on the right. The reaction may be summarized as

[latex]\begin{array}{lr @{{}\longrightarrow{}} l} \text{oxidation:} & \text{Mg}(s) & \longrightarrow \text{Mg}^{2+}(aq)\;+\;2\text{ e}^{-} \\[0.5em] \text{reduction:} & 2\text{ H}^{+}(aq)\;+\;2\text{ e}^{-} & \longrightarrow \text{H}_2(g) \\[0.5em] \hline \\[-0.25em] \text{overall:} & \text{Mg}(s)\;+\;2\text{ H}^{+}(aq) & \longrightarrow \text{Mg}^{2+}(aq)\;+\;\text{H}_2(g) \end{array}[/latex]

The cell used an inert platinum wire for the cathode, so the cell notation is

[latex]\text{Mg}(s){\mid}\text{Mg}^{2+}(aq){\parallel}\text{H}^{+}(aq){\mid}\text{H}_2(g){\mid}\text{Pt}(s)[/latex]

The magnesium electrode is an active electrode because it participates in the oxidation-reduction reaction. Inert electrodes, like the platinum electrode in Figure 18.3c, do not participate in the oxidation-reduction reaction and are present so that current can flow through the cell. Platinum or gold generally make good inert electrodes because they are chemically unreactive.

This figure contains a diagram of an electrochemical cell. Two beakers are shown. Each is just over half full. The beaker on the left contains a colorless solution and is labeled “Solution of M g C l subscript 2.” The beaker on the right contains a colorless solution and is labeled “Solution of H C l.” A glass tube in the shape of an inverted U connects the two beakers at the center of the diagram. The tube contents are colorless. The ends of the tubes are beneath the surface of the solutions in the beakers and a small grey plug is present at each end of the tube. At the center of the diagram, the tube is labeled “Salt bridge. Each beaker shows a metal coils submerged in the liquid. The beaker on the left has a thin grey coiled strip that is labeled “M g coil.” The beaker on the right has a black wire that is oriented horizontally and coiled up in a spring-like appearance that is labeled “Non reactive platinum wire.” A wire extends across the top of the diagram that connects the ends of the M g strip and platinum wire just above the opening of each beaker. This wire is labeled “Conducting wire.” At the center of this wire above the two beakers near the center of the diagram is a right pointing arrow with the label “e superscript negative” at its base. Bubbles appear to be rising from the coiled platinum wire in the beaker. These bubbles are labeled “Bubbles of H subscript 2.” An arrow points down and to the right from a plus sign at the upper right region of the salt bride. An arrow points down and to the left from a negative sign at the upper left region of the salt bride. A curved arrow extends from the grey plug at the left end of the salt bridge into the surrounding solution in the left beaker. The label “Oxidation M g right pointing arrow M g superscript 2 plus plus 2 e superscript negative” appears beneath the left beaker. The label “Reduction 2 H superscript plus plus 2 e superscript negative right pointing arrow H subscript 2” appears beneath the right beaker.
Figure 18.3c: The oxidation of magnesium to magnesium ion occurs in the beaker on the left side in this apparatus; the reduction of hydrogen ions to hydrogen occurs in the beaker on the right. A nonreactive, or inert, platinum wire allows electrons from the left beaker to move into the right beaker. The overall reaction is: Mg + 2 H+ → Mg2+ + H2, which is represented in cell notation as: Mg(s) │ Mg2+(aq) ║ H+(aq) │ H2(g) │ Pt(s) (credit: Chemistry (OpenStax), CC BY 4.0).

Example 18.3a

Using Cell Notation

Consider a galvanic cell consisting of

[latex]2\text{ Cr}(s)\;+\;3\text{ Cu}^{2+}(aq)\;{\longrightarrow}\;2\text{ Cr}^{3+}(aq)\;+\;3\text{ Cu}(s)[/latex]

Write the oxidation and reduction half-reactions and write the reaction using cell notation. Which reaction occurs at the anode? The cathode?

Solution

By inspection, Cr is oxidized when three electrons are lost to form Cr3+, and Cu2+ is reduced as it gains two electrons to form Cu. Balancing the charge gives

[latex]\begin{array}{lr @{{}\longrightarrow{}} l} \text{oxidation:} & 2\text{ Cr}(s) & \longrightarrow 2\text{ Cr}^{3+}(aq)\;+\;6\text{ e}^{-} \\[0.5em] \text{reduction:} & 3\text{ Cu}^{2+}(aq)\;+\;6\text{ e}^{-} & \longrightarrow 3\text{ Cu}(s) \\[0.5em] \hline \\[-0.25em] \text{overall:} & 2\text{ Cr}(s)\;+\;3\text{ Cu}^{2+}(aq) & \longrightarrow 2\text{ Cr}^{3+}(aq)\;+\;3\text{ Cu}(s) \end{array}[/latex]

Cell notation uses the simplest form of each of the equations, and starts with the reaction at the anode. No concentrations were specified so:

[latex]\text{Cr}(s){\mid}\text{Cr}^{3+}(aq){\parallel}\text{Cu}^{2+}(aq){\mid}\text{Cu}(s)[/latex]. Oxidation occurs at the anode and reduction at the cathode.

 

Example 18.3b

Using Cell Notation

Consider a galvanic cell consisting of

[latex]5\text{ Fe}^{2+}(aq)\;+\;\text{MnO}_4^{\;\;-}(aq)\;+\;8\text{ H}^{+}(aq)\;{\longrightarrow}\;5\text{ Fe}^{3+}(aq)\;+\;\text{Mn}^{2+}(aq)\;+\;4\text{ H}_2\text{O}(l)[/latex]

Write the oxidation and reduction half-reactions and write the reaction using cell notation. Which reaction occurs at the anode? The cathode?

Solution

By inspection, Fe2+ undergoes oxidation when one electron is lost to form Fe3+, and MnO4− is reduced as it gains five electrons to form Mn2+. Balancing the charge gives
[latex]\begin{array}{lr @{{}\longrightarrow{}} l} \text{oxidation:} & 5\;\times\;[\text{Fe}^{2+}(aq) & \longrightarrow \text{Fe}^{3+}(aq)\;+\;\text{e}^{-}] \\[0.5em] \text{reduction:} & \text{MnO}_4^{\;\;-}(aq)\;+\;8\text{ H}^{+}(aq)\;+\;5\text{ e}^{-} & \longrightarrow \text{Mn}^{2+}(aq)\;+\;4\text{ H}_2\text{O}(l) \\[0.5em] \hline \\[-0.25em] \text{overall:} & 5\text{ Fe}^{2+}(aq)\;+\;\text{MnO}_4^{\;\;-}(aq)\;+\;8\text{ H}^{+}(aq) & \longrightarrow 5\text{ Fe}^{3+}(aq)\;+\;\text{Mn}^{2+}(aq)\;+\;4\text{ H}_2\text{O}(l) \end{array}[/latex]

Cell notation uses the simplest form of each of the equations, and starts with the reaction at the anode. It is necessary to use an inert electrode, such as platinum, because there is no metal present to conduct the electrons from the anode to the cathode. No concentrations were specified so: [latex]\text{Pt}(s){\mid}\text{Fe}^{2+}(aq)\text{,}\;\text{Fe}^{3+}(aq){\parallel}\text{MnO}_4^{\;\;-}(aq)\text{,}\;\text{H}^{+}(aq)\text{,}\;\text{Mn}^{2+}(aq){\mid}\text{Pt}(s)[/latex].

Oxidation occurs at the anode and reduction at the cathode.

Exercise 18.3a

Use cell notation to describe the galvanic cell where copper(II) ions are reduced to copper metal and zinc metal is oxidized to zinc ions.

Check Your Answer[1]

Links to Interactive Learning Tools

Explore Galvanic Cells from the Physics Classroom.

Key Equations

[latex]\text{volt} = V = \frac{\text{J}}{\text{C}}[/latex]

Attribution & References

Except where otherwise noted, this page is adapted by David Wegman from “17.2 Galvanic Cells” In General Chemistry 1 & 2 by Rice University, a derivative of Chemistry (Open Stax) by Paul Flowers, Klaus Theopold, Richard Langley & William R. Robinson and is licensed under CC BY 4.0. ​Access for free at Chemistry (OpenStax)​.

  1. From the information given in the problem:

    [latex]\begin{array}{lr @{{}\longrightarrow{}} l} \text{anode (oxidation):} & \text{Zn}(s) & \longrightarrow \text{Zn}^{2+}(aq)\;+\;2\text{ e}^{-} \\[0.5em] \text{cathode (reduction):} & \text{Cu}^{2+}(aq)\;+\;2\text{ e}^{-} & \longrightarrow \text{Cu}(s) \\[0.5em] \hline \\[-0.25em] \text{overall:} & \text{Zn}(s)\;+\;\text{Cu}^{2+}(aq) & \longrightarrow \text{Zn}^{2+}(aq)\;+\;\text{Cu}(s) \end{array}[/latex]
    Using cell notation: [latex]\text{Zn}(s){\mid}\text{Zn}^{2+}(aq){\parallel}\text{Cu}^{2+}(aq){\mid}\text{Cu}(s)[/latex].
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Enhanced Introductory College Chemistry Copyright © 2023 by Gregory Anderson; Caryn Fahey; Jackie MacDonald; Adrienne Richards; Samantha Sullivan Sauer; J.R. van Haarlem; and David Wegman is licensed under a Creative Commons Attribution 4.0 International License, except where otherwise noted.

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