Chapter 17: Electrochemistry

Learning Outcomes

  • Describe the function of a galvanic cell and its components
  • Use cell notation to symbolize the composition and construction of galvanic cells

As demonstration of spontaneous chemical change, Figure 17.2.1 shows the result of immersing a coiled wire of copper into an aqueous solution of silver nitrate. A gradual but visually impressive change spontaneously occurs as the initially colorless solution becomes increasingly blue, and the initially smooth copper wire becomes covered with a porous gray solid.

This figure includes three photographs. In the first, a test tube containing a clear, colorless liquid is shown with a loosely coiled copper wire outside the test tube to its right. In the second, the wire has been submerged in the clear colorless liquid in the test tube and the surface of the wire is darkened. In the third, the liquid in the test tube is bright blue-green, the wire in the solution appears dark near the top, and a gray “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 17.2.1. A copper wire and an aqueous solution of silver nitrate (left) are brought into contact (center) and a spontaneous transfer of electrons occurs, creating blue Cu2+(aq) and gray Ag(s) (right).

These observations are consistent with (i) the oxidation of elemental copper to yield copper(II) ions, [latex]\ce{Cu^2+}[/latex](aq), which impart a blue color to the solution, and (ii) the reduction of silver(I) ions to yield elemental silver, which deposits as a fluffy solid on the copper wire surface. And so, the direct transfer of electrons from the copper wire to the aqueous silver ions is spontaneous under the employed conditions. A summary of this redox system is provided by these equations:

[latex]\begin{array}{rl}{}\\ \text{overall reaction:}&2{\ce{Ag}}^{\text{+}}\left(aq\right)+\ce{Cu}\left(s\right)\longrightarrow \ce{2Ag}\left(s\right)+{\ce{Cu}}^{2+}\left(aq\right)\\ \text{oxidation half-reaction:}&\ce{Cu}\left(s\right)\longrightarrow {\ce{Cu}}^{2+}\left(aq\right)+{\text{2e}}^{-}\\ \text{reduction half-reaction:}&2{\ce{Ag}}^{\text{+}}\left(aq\right)+{\text{2e}}^{-}\longrightarrow \ce{2Ag}\left(s\right)\end{array}[/latex]

Consider the construction of a device that contains all the reactants and products of a redox system like the one here, but prevents physical contact between the reactants. Direct transfer of electrons is, therefore, prevented; transfer, instead, takes place indirectly through an external circuit that contacts the separated reactants. Devices of this sort are generally referred to as electrochemical cells, and those in which a spontaneous redox reaction takes place are called galvanic cells (or voltaic cells).

A galvanic cell based on the spontaneous reaction between copper and silver(I) is depicted in Figure 17.2.2. The cell is comprised of two half-cells, each containing the redox conjugate pair (“couple”) of a single reactant. The half-cell shown at the left contains the [latex]\ce{Cu(0)/Cu(II)}[/latex] couple in the form of a solid copper foil and an aqueous solution of copper nitrate. The right half-cell contains the [latex]\ce{Ag(I)/Ag(0)}[/latex] couple as solid silver foil and an aqueous silver nitrate solution. An external circuit is connected to each half-cell at its solid foil, meaning the [latex]\ce{Cu}[/latex] and [latex]\ce{Ag}[/latex] foil each function as an electrode. By definition, the anode of an electrochemical cell is the electrode at which oxidation occurs (in this case, the [latex]\ce{Cu}[/latex] foil) and the cathode is the electrode where reduction occurs (the [latex]\ce{Ag}[/latex] foil). The redox reactions in a galvanic cell occur only at the interface between each half-cell’s reaction mixture and its electrode. To keep the reactants separate while maintaining charge-balance, the two half-cell solutions are connected by a tube filled with inert electrolyte solution called a salt bridge. The spontaneous reaction in this cell produces [latex]\ce{Cu^2+}[/latex] cations in the anode half-cell and consumes [latex]\ce{Ag+}[/latex] ions in the cathode half-cell, resulting in a compensatory flow of inert ions from the salt bridge that maintains charge balance. Increasing concentrations of [latex]\ce{Cu^2+}[/latex] in the anode half-cell are balanced by an influx of [latex]\ce{NO3-}[/latex] from the salt bridge, while a flow of [latex]\ce{Na+}[/latex] into the cathode half-cell compensates for the decreasing [latex]\ce{Ag+}[/latex] concentration.

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 gray 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 rectangle indicating “external circuit” that is labeled “flow of electrons” with an arrow pointing to the right following. 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 salt bridge in the colorless solution is the label “N a superscript plus.” Just above this region of the tube appears the label “Flow of cations.” Just left of the salt bridge in the blue solution is the label “N O subscript 3 superscript negative.” Just above this region of the tube appears the label “Flow of anions.”
Figure 17.2.2. A galvanic cell based on the spontaneous reaction between copper and silver(I) ions.

Cell Notation

Abbreviated symbolism is commonly used to represent a galvanic cell by providing essential information on its composition and structure. These symbolic representations are called cell notations or cell schematics, and they are written following a few guidelines:

  • The relevant components of each half-cell are represented by their chemical formulas or element symbols
  • All interfaces between component phases are represented by vertical parallel lines; if two or more components are present in the same phase, their formulas are separated by commas
  • By convention, the schematic begins with the anode and proceeds left-to-right identifying phases and interfaces encountered within the cell, ending with the cathode

A verbal description of the cell as viewed from anode-to-cathode is often a useful first-step in writing its schematic. For example, the galvanic cell shown in Figure 17.2.2 consists of a solid copper anode immersed in an aqueous solution of copper(II) nitrate that is connected via a salt bridge to an aqueous silver(I) nitrate solution, immersed in which is a solid silver cathode. Converting this statement to symbolism following the above guidelines results in the cell schematic:

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

Consider a different galvanic cell (see Figure 17.2.3) based on the spontaneous reaction between solid magnesium and aqueous iron(III) ions:

[latex]\begin{array}{rl}\\ \text{net cell reaction:}&\ce{Mg}\left(s\right)+2{\ce{Fe}}^{3\text{+}}\left(aq\right)\longrightarrow {\ce{Mg}}^{2+}\left(aq\right)+2{\ce{Fe}}^{2+}\left(g\right)\\ \text{oxidation half-reaction:}&\ce{Mg}\left(s\right)\longrightarrow\ce{Mg}^{2+}\left(aq\right)+\text{2e}^{-}\\ \text{reduction half-reaction:}&2{\ce{Fe}}^{3\text{+}}\left(aq\right)+{\text{2e}}^{-}\longrightarrow 2{\ce{Fe}}^{2+}\left(g\right)\end{array}[/latex]

In this cell, a solid magnesium anode is immersed in an aqueous solution of magnesium chloride that is connected via a salt bridge to an aqueous solution containing a mixture of iron(III) chloride and iron(II) chloride, immersed in which is a platinum cathode. The cell schematic is then written as

[latex]\ce{Mg}(s)\mid {0.1 M} \ce{ MgCl2}(aq)\parallel {0.2 M}\ce{ FeCl3}(aq), {0.3 M}\ce{ FeCl2}(aq)\mid \ce{Pt}(s)[/latex]

Notice the cathode half-cell is different from the others considered thus far in that its electrode is comprised of a substance (Pt) that is neither a reactant nor a product of the cell reaction. This is required when neither member of the half-cell’s redox couple can reasonably function as an electrode, which must be electrically conductive and in a phase separate from the half-cell solution. In this case, both members of the redox couple are solute species, and so Pt is used as an inert electrode that can simply provide or accept electrons to redox species in solution. Electrodes constructed from a member of the redox couple, such as the Mg anode in this cell, are called active electrodes.

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. The beaker on the right also contains a colorless solution. 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 gray 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, gray, coiled strip that is labeled “M g anode.” The beaker on the right has a black wire that is oriented horizontally and coiled up in a spring-like appearance that is labeled “P t cathode.” Below the coil is the label “F e superscript 3 plus” with a curved right arrowing pointing from that to the label “F e superscript 2 plus.” A wire extends across the top of the diagram that connects the ends of the M g strip and P t cathode just above the opening of each beaker. At the center of the wire above the two beakers is a rectangle labeled “external circuit.” Above the rectangle is the label “flow of electrons” followed by a right pointing arrow. An arrow points down and to the right from the label “N a superscript plus” at the upper right region of the salt bride. An arrow points down and to the left from the label “C l superscript negative” at the upper left region of the salt bride. Below the graylug at the left end of the salt bridge in the surrounding solution in the left beaker is the label “C l superscript negative.” Below the coil on this side is a right arrow and the label “M g superscript 2 plus.” The label “0.1 M M g C l subscript 2” appears beneath the left beaker. The label “0.2 M F e C l subscript 3 and 0.3 M F e C l subscript 2.” appears beneath the right beaker.
Figure 17.2.3. A galvanic cell based on the spontaneous reaction between magnesium and iron(III) ions.

 

Example 17.2.1: Writing Galvanic Cell Schematics

A galvanic cell is fabricated by connecting two half-cells with a salt bridge, one in which a chromium wire is immersed in a 1 M [latex]\ce{CrCl3}[/latex] solution and another in which a copper wire is immersed in 1 M [latex]\ce{CuCl2}[/latex]. Assuming the chromium wire functions as an anode, write the schematic for this cell along with equations for the anode half-reaction, the cathode half-reaction, and the overall cell reaction.

[reveal-answer q=”188028″]Show Solution[/reveal-answer]
[hidden-answer a=”188028″]

Solution

Since the chromium wire is stipulated to be the anode, the schematic begins with it and proceeds left-to-right, symbolizing the other cell components until ending with the copper wire cathode:

[latex]\ce{Cr}(s)\mid{1 M }\ce{CrCl3}(aq)\parallel{1 M } \ce{ CuCl2}(aq)\mid\ce{ Cu}(s)[/latex]

The half-reactions for this cell are

[latex]\begin{array}{rl}{}\text{anode (oxidation):}&\ce{Cr}\left(s\right)\longrightarrow {\ce{Cr}}^{3+}\left(aq\right)+{\text{3e}}^{-}\\ \text{cathode (reduction):}&{\ce{Cu}}^{2+}\left(aq\right)+{\text{2e}}^{-}\longrightarrow \ce{Cu}\left(s\right)\end{array}[/latex]

Multiplying to make the number of electrons lost by [latex]\ce{Cr}[/latex] and gained by [latex]\ce{Cu^2+}[/latex] equal yields

[latex]\begin{array}{rl}{}\text{anode (oxidation):}&2\ce{Cr}\left(s\right)\longrightarrow {2\ce{Cr}}^{3+}\left(aq\right)+{\text{6e}}^{-}\\ \text{cathode (reduction):}&{3\ce{Cu}}^{2+}\left(aq\right)+{\text{6e}}^{-}\longrightarrow {3}\ce{Cu}\left(s\right)\end{array}[/latex]

Adding the half-reaction equations and simplifying yields an equation for the cell reaction:

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

[/hidden-answer]

Check Your Learning

Key Concepts and Summary

Electrochemical cells typically consist of two half-cells. The half-cells separate the oxidation half-reaction from the reduction half-reaction and make it possible for current to flow through an external wire. One half-cell, normally depicted on the left side in a figure, contains the anode. Oxidation occurs at the anode. The anode is connected to the cathode in the other half-cell, often shown on the right side in a figure. Reduction occurs at the cathode. Adding a salt bridge completes the circuit allowing current to flow. Anions in the salt bridge flow toward the anode and cations in the salt bridge flow toward the cathode. The movement of these ions completes the circuit and keeps each half-cell electrically neutral. Electrochemical cells can be described using cell notation. In this notation, information about the reaction at the anode appears on the left and information about the reaction at the cathode on the right. The salt bridge is represented by a double line, ‖. The solid, liquid, or aqueous phases within a half-cell are separated by a single line, |. The phase and concentration of the various species is included after the species name. Electrodes that participate in the oxidation-reduction reaction are called active electrodes. Electrodes that do not participate in the oxidation-reduction reaction but are there to allow current to flow are inert electrodes. Inert electrodes are often made from platinum or gold, which are unchanged by many chemical reactions.

Try It

  1. Write the following balanced reactions using cell notation. Use platinum as an inert electrode, if needed.
    1. [latex]\ce{Mg}(s)+\ce{Ni^2+}(aq)\longrightarrow \ce{Mg^2+}(aq)+\ce{Ni}(s)[/latex]
    2. [latex]\ce{2Ag+}(aq)+\ce{Cu}(s)\longrightarrow \ce{Cu^2+}(aq)+\ce{2Ag}(s)[/latex]
    3. [latex]\ce{Mn}(s)+\ce{Sn(NO3)2}(aq)\longrightarrow \ce{Mn}\ce{(NO3)2}(aq)+\ce{Au}(s)[/latex]
    4. [latex]\ce{3CuNO3}(aq)+\ce{Au(NO3)3}(aq)\longrightarrow\ce{3Cu(NO3)2}(aq)+\ce{Au}(s)[/latex]
  2. Given the following cell notations, determine the species oxidized, species reduced, and the oxidizing agent and reducing agent, without writing the balanced reactions.
    1. [latex]\ce{Mg}(s)\mid \ce{Mg^2+}(aq)\parallel \ce{Cu^2+}(aq)\mid \ce{Cu}(s)[/latex]
    2. [latex]\ce{Ni}(s)\mid \ce{Ni^2+}(aq)\parallel \ce{Ag+}\ce(aq\ce)\mid \ce{Ag}(s)[/latex]
  3. For the cell notations in the previous problem, write the corresponding balanced reactions.
  4. Balance the following reactions and write the reactions using cell notation. Ignore any inert electrodes, as they are never part of the half-reactions.
    1. [latex]\ce{Al}(s)+\ce{Zr^4+}(aq)\longrightarrow \ce{Al^3+}(aq)+\ce{Zr}(s)[/latex]
    2. [latex]\ce{Ag+}(aq)+\ce{NO}(g)\longrightarrow \ce{Ag}(s)+\ce{NO3-}(aq)\text{(acidic solution)}[/latex]
    3. [latex]\ce{SiO3^2-}(aq)+\ce{Mg}(s)\longrightarrow \ce{Si}(s)+\ce{Mg(OH)2}(s)\text{(basic solution)}[/latex]
    4. [latex]\ce{ClO3-}(aq)+\ce{MnO2}(s)\longrightarrow \ce{Cl-}(aq)+\ce{MnO4-}(aq)\text{(basic solution)}[/latex]
  5. Identify the species oxidized, species reduced, and the oxidizing agent and reducing agent for all the reactions in the previous problem.
  6. From the information provided, use cell notation to describe the following systems:
    1. In one half-cell, a solution of [latex]\ce{Pt(NO3)2}[/latex] forms [latex]\ce{Pt}[/latex] metal, while in the other half-cell, [latex]\ce{Cu}[/latex] metal goes into a [latex]\ce{Cu(NO3)2}[/latex] solution with all solute concentrations 1 M.
    2. The cathode consists of a gold electrode in a 0.55 M [latex]\ce{Au(NO3)3}[/latex] solution and the anode is a magnesium electrode in 0.75 M [latex]\ce{Mg(NO3)2}[/latex] solution.
    3. One half-cell consists of a silver electrode in a 1 M [latex]\ce{AgNO3}[/latex] solution, and in the other half-cell, a copper electrode in 1 M [latex]\ce{Cu(NO3)2}[/latex] is oxidized.
  7. Why is a salt bridge necessary in galvanic cells like the one in Figure 17.2.2?
  8. An active (metal) electrode was found to gain mass as the oxidation-reduction reaction was allowed to proceed. Was the electrode part of the anode or cathode? Explain.
  9. An active (metal) electrode was found to lose mass as the oxidation-reduction reaction was allowed to proceed. Was the electrode part of the anode or cathode? Explain.
  10. The mass of three different metal electrodes, each from a different galvanic cell, were determined before and after the current generated by the oxidation-reduction reaction in each cell was allowed to flow for a few minutes. The first metal electrode, given the label A, was found to have increased in mass; the second metal electrode, given the label B, did not change in mass; and the third metal electrode, given the label C, was found to have lost mass. Make an educated guess as to which electrodes were active and which were inert electrodes, and which were anode(s) and which were the cathode(s).

[reveal-answer q=”437960″]Show Selected Solutions[/reveal-answer]
[hidden-answer a=”437960″]

  1. The balanced reactions are as follows:
    1. [latex]\ce{Mg}(s)\mid \ce{Mg^2+}(aq)\parallel \ce{Ni+}(aq)\mid \ce{Ni}(s)[/latex]
    2. Stoichiometric coefficients do not appear in cell notation [latex]\ce{Cu}(s)\mid \ce{Cu^2+}(aq)\parallel \ce{Ag+}(aq)\mid \ce{Ag}(s)[/latex]
    3. Spectator ions do not appear in cell notation [latex]\ce{Mn}(s)\mid \ce{Mn^2+}(aq)\parallel \ce{Sn^2+}(aq)\mid \ce{Sn}(s)[/latex]
    4. Neither stoichiometric coefficients nor spectator ions appear in cell notation. Platinum electrode needed [latex]\ce{Pt}(s)\mid\ce{Cu^+}(aq),\ce{Cu^2+}(aq)\parallel\ce{Au^3+}(aq)\mid \ce{Au}(s)[/latex]
  2. The balanced reactions are as follows:
    1. [latex]\ce{Mg}(s)+\ce{Cu^2+}(aq)\longrightarrow \ce{Mg^2+}(aq)+\ce{Cu}(s)[/latex]
    2. [latex]\ce{2Ag+}(aq)+\ce{Ni}(s)\longrightarrow \ce{Ni^2+}(aq)+\ce{2Ag}(s)[/latex]
  3. Species oxidized = reducing agent:
    1. [latex]\ce{Al}[/latex](s)
    2. [latex]\ce{NO}[/latex](g)
    3. [latex]\ce{Mg}[/latex](s)
    4. [latex]\ce{MnO2}[/latex](s)
  4. Species reduced = oxidizing agent:
    1. [latex]\ce{Zr^4+}[/latex](aq)
    2. [latex]\ce{Ag+}[/latex](aq)
    3. [latex]\ce{SiO3^2-}(aq)[/latex]
    4. [latex]\ce{ClO3-}(aq)[/latex]
  5. Without the salt bridge, the circuit would be open (or broken) and no current could flow. With a salt bridge, each half-cell remains electrically neutral and current can flow through the circuit.
  6. Active electrodes participate in the oxidation-reduction reaction. Since metals form cations, the electrode would lose mass if metal atoms in the electrode were to oxidize and go into solution. Oxidation occurs at the anode.

[/hidden-answer]

Glossary

active electrode: electrode that participates in the oxidation-reduction reaction of an electrochemical cell; the mass of an active electrode changes during the oxidation-reduction reaction

anode: electrode in an electrochemical cell at which oxidation occurs; information about the anode is recorded on the left side of the salt bridge in cell notation

cathode: electrode in an electrochemical cell at which reduction occurs; information about the cathode is recorded on the right side of the salt bridge in cell notation

cell notation: shorthand way to represent the reactions in an electrochemical cell

cell potential: difference in electrical potential that arises when dissimilar metals are connected; the driving force for the flow of charge (current) in oxidation-reduction reactions

galvanic cell: electrochemical cell that involves a spontaneous oxidation-reduction reaction; electrochemical cells with positive cell potentials; also called a voltaic cell

inert electrode: electrode that allows current to flow, but that does not otherwise participate in the oxidation-reduction reaction in an electrochemical cell; the mass of an inert electrode does not change during the oxidation-reduction reaction; inert electrodes are often made of platinum or gold because these metals are chemically unreactive.

voltaic cell: another name for a galvanic cell

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