Batteries and Fuel Cells

Galvanic cells designed specifically to function as electrical power supplies are called batteries. A variety of both single-use batteries (primary cells) and rechargeable batteries (secondary cells) are commercially available to serve a variety of applications, with important specifications including voltage, size, and lifetime. Fuel cells, sometimes called flow batteries, are devices that harness the energy of spontaneous redox reactions normally associated with combustion processes. Like batteries, fuel cells enable the reaction’s electron transfer via an external circuit, but they require continuous input of the redox reactants (fuel and oxidant) from an external reservoir. Fuel cells are typically much more efficient in converting the energy released by the reaction to useful work in comparison to internal combustion engines. Spontaneous oxidation of metals by natural electrochemical processes is called corrosion, familiar examples including the rusting of iron and the tarnishing of silver. Corrosion process involve the creation of a galvanic cell in which different sites on the metal object function as anode and cathode, with the corrosion taking place at the anodic site. Approaches to preventing corrosion of metals include use of a protective coating of zinc (galvanization) and the use of sacrificial anodes connected to the metal object (cathodic protection). Nonspontaneous redox processes may be forced to occur in electrochemical cells by the application of an appropriate potential using an external power source—a process known as electrolysis. Electrolysis is the basis for certain ore refining processes, the industrial production of many chemical commodities, and the electroplating of metal coatings on various products. Measurement of the current flow during electrolysis permits stoichiometric calculations.

38.1 Batteries and Fuel Cells

Learning Objectives

By the end of this section, you will be able to:

  • Describe the electrochemistry associated with several common batteries
  • Distinguish the operation of a fuel cell from that of a battery

There are many technological products associated with the past two centuries of electrochemistry research, none more immediately obvious than the battery. A battery is a galvanic cell that has been specially designed and constructed in a way that best suits its intended use a source of electrical power for specific applications. Among the first successful batteries was the Daniell cell, which relied on the spontaneous oxidation of zinc by copper(II) ions (Figure 38.1):

Figure 38.1

Illustration of a Daniell cell taken from a 1904 journal publication (left) along with a simplified illustration depicting the electrochemistry of the cell (right). The 1904 design used a porous clay pot to both contain one of the half-cell’s content and to serve as a salt bridge to the other half-cell.

This figure contains a patent drawing for an electrochemical cell on the left labelled Element Daniell and a diagram of an electrochemical cell on the right. In the diagram, two beakers are shown. Each is just over half full. The beaker on the left contains a blue solution. The beaker on the right 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 grey plug is present at each end of the tube. The plug in the left beaker is labeled “Porous plug.” Each beaker shows a metal strip partially submerged in the liquid. The beaker on the left has a silver strip that is labeled “Z n anode” at the top. The beaker on the right has an orange brown strip that is labeled “C u cathode” at the top. A wire extends up and toward the center from the top of each of these strips before stopping. The end of the left wire points up to a negative sign. The end of the right wire points up to a positive sign. An arrow points toward the left wire which is labeled “Flow of e superscript negative.” A curved arrow extends from the Z n strip into the surrounding solution. The tip of this arrow is labeled “Z n 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 “S O subscript 4 superscript 2 negative.” A curved arrow extends from the solution in the beaker on the right to the C u strip. The base of this arrow is labeled “C u superscript 2 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 “S O subscript 4 superscript 2 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 “Z n superscript 2 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 “S O subscript 4 superscript 2 negative.” Just above this region of the tube appears the label “Flow of anions.”

Modern batteries exist in a multitude of forms to accommodate various applications, from tiny button batteries that provide the modest power needs of a wristwatch to the very large batteries used to supply backup energy to municipal power grids. Some batteries are designed for single-use applications and cannot be recharged (primary cells), while others are based on conveniently reversible cell reactions that allow recharging by an external power source (secondary cells). This section will provide a summary of the basic electrochemical aspects of several batteries familiar to most consumers, and will introduce a related electrochemical device called a fuel cell that can offer improved performance in certain applications.

38.2 Single-Use Batteries

A common primary battery is the dry cell, which uses a zinc can as both container and anode (“–” terminal) and a graphite rod as the cathode (“+” terminal). The Zn can is filled with an electrolyte paste containing manganese(IV) oxide, zinc(II) chloride, ammonium chloride, and water. A graphite rod is immersed in the electrolyte paste to complete the cell. The spontaneous cell reaction involves the oxidation of zinc:

anode reaction:Zn(s)Zn2+(aq)+2eanode reaction:Zn(s)Zn2+(aq)+2e

and the reduction of manganese(IV)

reduction reaction:2MnO2(s)+2NH4Cl(aq)+2eMn2O3(s)+2NH3(aq)+H2O(l)+2Clreduction reaction:2MnO2(s)+2NH4Cl(aq)+2eMn2O3(s)+2NH3(aq)+H2O(l)+2Cl

which together yield the cell reaction:

cell reaction:2MnO2(s)+2NH4Cl(aq)+Zn(s)Zn2+(aq)+Mn2O3(s)+2NH3(aq)+H2O(l)+2ClEcell~1.5Vcell reaction:2MnO2(s)+2NH4Cl(aq)+Zn(s)Zn2+(aq)+Mn2O3(s)+2NH3(aq)+H2O(l)+2ClEcell~1.5V

The voltage (cell potential) of a dry cell is approximately 1.5 V. Dry cells are available in various sizes (e.g., D, C, AA, AAA). All sizes of dry cells comprise the same components, and so they exhibit the same voltage, but larger cells contain greater amounts of the redox reactants and therefore are capable of transferring correspondingly greater amounts of charge. Like other galvanic cells, dry cells may be connected in series to yield batteries with greater voltage outputs, if needed.

Figure 38.2

A schematic diagram shows a typical dry cell.

A diagram of a cross section of a dry cell battery is shown. The overall shape of the cell is cylindrical. The lateral surface of the cylinder, indicated as a thin red line, is labeled “zinc can (electrode).” Just beneath this is a slightly thicker dark grey surface that covers the lateral surface, top, and bottom of the battery, which is labeled “Porous separator.” Inside is a purple region with many evenly spaced small darker purple dots, labeled “Paste of M n O subscript 2, N H subscript 4 C l, Z n C l subscript 2, water (cathode).” A dark grey rod, labeled “Carbon rod (electrode),” extends from the top of the battery, leaving a gap of less than one-fifth the height of the battery below the rod to the bottom of the cylinder. A thin grey line segment at the very bottom of the cylinder is labeled “Metal bottom cover (negative).” The very top of the cylinder has a thin grey surface that curves upward at the center over the top of the carbon electrode at the center of the cylinder. This upper surface is labeled “Metal top cover (positive).” A thin dark grey line just below this surface is labeled “Insulator.” Below this, above the purple region, and outside of the carbon electrode at the center is an orange region that is labeled “Seal.”

Alkaline batteries (Figure 38.3) were developed in the 1950s to improve on the performance of the dry cell, and they were designed around the same redox couples. As their name suggests, these types of batteries use alkaline electrolytes, often potassium hydroxide. The reactions are

anode:Zn(s)+2OH(aq)ZnO(s)+H2O(l)+2ecathode:2MnO2(s)+H2O(l)+2eMn2O3(s)+2OH(aq)¯cell:Zn(s)+2MnO2(s)ZnO(s)+Mn2O3(s)Ecell=+1.43 Vanode:Zn(s)+2OH(aq)ZnO(s)+H2O(l)+2ecathode:2MnO2(s)+H2O(l)+2eMn2O3(s)+2OH(aq)¯cell:Zn(s)+2MnO2(s)ZnO(s)+Mn2O3(s)Ecell=+1.43 V

An alkaline battery can deliver about three to five times the energy of a zinc-carbon dry cell of similar size. Alkaline batteries are prone to leaking potassium hydroxide, so they should be removed from devices for long-term storage. While some alkaline batteries are rechargeable, most are not. Attempts to recharge an alkaline battery that is not rechargeable often leads to rupture of the battery and leakage of the potassium hydroxide electrolyte.

Figure 38.3

Alkaline batteries were designed as improved replacements for zinc-carbon (dry cell) batteries.

A diagram of a cross section of an alkaline battery is shown. The overall shape of the cell is cylindrical. The lateral surface of the cylinder, indicated as a thin red line, is labeled “Outer casing.” Just beneath this is a thin, light grey surface that covers the lateral surface and top of the battery. Inside is a blue region with many evenly spaced small darker dots, labeled “M n O subscript 2 (cathode).” A thin dark grey layer is just inside, which is labeled “Ion conducting separator.” A purple region with many evenly spaced small darker dots fills the center of the battery and is labeled “ zinc (anode).” The very top of the battery has a thin grey curved surface over the central purple region. The curved surface above is labeled “Positive connection (plus).” At the base of the battery, an orange structure, labeled “Protective cap,” is located beneath the purple and blue central regions. This structure holds a grey structure that looks like a nail with its head at the bottom and pointed end extending upward into the center of the battery. This nail-like structure is labeled “Current pick up.” At the very bottom of the battery is a thin grey surface that is held by the protective cap. This surface is labeled “Negative terminal (negative).”

38.3 Rechargeable (Secondary) Batteries

Nickel-cadmium, or NiCd, batteries (Figure 38.4) consist of a nickel-plated cathode, cadmium-plated anode, and a potassium hydroxide electrode. The positive and negative plates, which are prevented from shorting by the separator, are rolled together and put into the case. This is a “jelly-roll” design and allows the NiCd cell to deliver much more current than a similar-sized alkaline battery. The reactions are


When properly treated, a NiCd battery can be recharged about 1000 times. Cadmium is a toxic heavy metal so NiCd batteries should never be ruptured or incinerated, and they should be disposed of in accordance with relevant toxic waste guidelines.

Figure 38.4

NiCd batteries use a “jelly-roll” design that significantly increases the amount of current the battery can deliver as compared to a similar-sized alkaline battery.

A diagram is shown of a cross section of a nickel cadmium battery. This battery is in a cylindrical shape. An outer red layer is labeled “case.” Just inside this layer is a thin, dark grey layer which is labeled at the bottom of the cylinder as “Negative electrode collector.” A silver rod extends upward through the center of the battery, which is surrounded by alternating layers, shown as vertical repeating bands, of yellow, purple, yellow, and blue. A slightly darker grey narrow band extends across the top of these alternating bands, which is labeled “Positive electrode collector.” A thin light grey band appears at the very bottom of the cylinder, which is labeled “Metal bottom cover (negative).” A small grey and white striped rectangular structure is present at the top of the central silver cylinder, which is labeled “Safety valve.” Above this is an orange layer that curves upward over the safety valve, which is labeled “Insulation ring.” Above this is a thin light grey layer that projects upward slightly at the center, which is labeled “Metal top cover (plus).” A light grey arrow points to a rectangle to the right that illustrates the layers at the center of the battery under magnification. From the central silver rod, the layers shown repeat the alternating pattern yellow, blue, yellow, and purple three times, with a final yellow layer covering the last purple layer. The outermost purple layer is labeled “Negative electrode.” The yellow layer beneath it is labeled “Separator.” The blue layer just inside is labeled “Positive electrode.”

Lithium ion batteries (Figure 38.5) are among the most popular rechargeable batteries and are used in many portable electronic devices. The reactions are


The variable stoichiometry of the cell reaction leads to variation in cell voltages, but for typical conditions, x is usually no more than 0.5 and the cell voltage is approximately 3.7 V. Lithium batteries are popular because they can provide a large amount current, are lighter than comparable batteries of other types, produce a nearly constant voltage as they discharge, and only slowly lose their charge when stored.

Figure 38.5

In a lithium ion battery, charge flows as the lithium ions are transferred between the anode and cathode.

This figure shows a model of the flow of charge in a lithium ion battery. At the left, an approximately cubic structure formed by alternating red, grey, and purple spheres is labeled below as “Positive electrode.” The purple spheres are identified by the label “lithium.” The grey spheres are identified by the label “Metal.” The red spheres are identified by the label “oxygen.” Above this structure is the label “Charge” followed by a right pointing green arrow. At the right is a figure with layers of black interconnected spheres with purple spheres located in gaps between the layers. The black layers are labeled “Graphite layers.” Below the purple and black structure is the label “Negative electrode.” Above is the label “Discharge,” which is preceded by a blue arrow which points left. At the center of the diagram between the two structures are six purple spheres which are each labeled with a plus symbol. Three curved green arrows extend from the red, purple, and grey structure to each of the three closest purple plus labeled spheres. Green curved arrows extend from the right side of the upper and lower of these three purple plus labeled spheres to the black and purple layered structure. Three blue arrows extend from the purple and black layered structure to the remaining three purple plus labeled spheres at the center of the diagram. The base of each arrow has a circle formed by a dashed curved line in the layered structure. The lowest of the three purple plus marked spheres reached by the blue arrows has a second blue arrow extending from its left side which points to a purple sphere in the purple, green, and grey structure.

The lead acid battery (Figure 38.6) is the type of secondary battery commonly used in automobiles. It is inexpensive and capable of producing the high current required by automobile starter motors. The reactions for a lead acid battery are

anode:Pb(s)+HSO4(aq)PbSO4(s)+H+(aq)+2ecathode: PbO2(s)+HSO4(aq)+3H+(aq)+2ePbSO4(s)+2H2O(l)¯cell:Pb(s)+PbO2(s)+2H2SO4(aq)2PbSO4(s)+2H2O(l)Ecell~2Vanode:Pb(s)+HSO4(aq)PbSO4(s)+H+(aq)+2ecathode: PbO2(s)+HSO4(aq)+3H+(aq)+2ePbSO4(s)+2H2O(l)¯cell:Pb(s)+PbO2(s)+2H2SO4(aq)2PbSO4(s)+2H2O(l)Ecell~2V

Each cell produces 2 V, so six cells are connected in series to produce a 12-V car battery. Lead acid batteries are heavy and contain a caustic liquid electrolyte, H2SO4(aq), but are often still the battery of choice because of their high current density. Since these batteries contain a significant amount of lead, they must always be disposed of properly.

Figure 38.6

The lead acid battery in your automobile consists of six cells connected in series to give 12 V.

A diagram of a lead acid battery is shown. A black outer casing, which is labeled “Protective casing” is in the form of a rectangular prism. Grey cylindrical projections extend upward from the upper surface of the battery in the back left and back right corners. At the back right corner, the projection is labeled “Positive terminal.” At the back right corner, the projection is labeled “Negative terminal.” The bottom layer of the battery diagram is a dark green color, which is labeled “Dilute H subscript 2 S O subscript 4.” A blue outer covering extends upward from this region near the top of the battery. Inside, alternating grey and white vertical “sheets” are packed together in repeating units within the battery. The battery has the sides cut away to show three of these repeating units which are separated by black vertical dividers, which are labeled as “cell dividers.” The grey layers in the repeating units are labeled “Negative electrode (lead).” The white layers are labeled “Postive electrode (lead dioxide).”

38.4 Fuel Cells

A fuel cell is a galvanic cell that uses traditional combustive fuels, most often hydrogen or methane, that are continuously fed into the cell along with an oxidant. (An alternative, but not very popular, name for a fuel cell is a flow battery.) Within the cell, fuel and oxidant undergo the same redox chemistry as when they are combusted, but via a catalyzed electrochemical that is significantly more efficient. For example, a typical hydrogen fuel cell uses graphite electrodes embedded with platinum-based catalysts to accelerate the two half-cell reactions:

Figure 38.7

In this hydrogen fuel cell, oxygen from the air reacts with hydrogen, producing water and electricity.

A diagram is shown of a hydrogen fuel cell. At the center is a vertical rectangle which is shaded dark gray and labeled “Electrolyte.” This region has two labels for H superscript plus in it. To the right and left are narrow vertical rectangles shaded light gray. The one to the right is labeled “Cathode” and the one to the left is labeled “Anode.” To the left of the left-most light gray region is a white region shaped like a closed left bracket. A yellow arrow points in to the white region with the label to show “Fuel In.” In the middle of the white area are two yellow arrows pointing toward the gray shading labeled “H subscript 2.” At the bottom of the white region is a yellow arrow pointing out that is labeled “Excess Fuel.” On the right side is another white region that makes a right closed bracket shape. There are two arrows with the label “Air In” and “H subscript 2 O” in the upper left side of this area pointing in. One arrow is light blue and one is dark blue. In the middle to the white area is a light blue arrow pointing toward the gray shading. The arrow is labeled “O subscript 2.” Below that are two dark blue arrows pointing out from the gray shading to the white area labeled “H subscript 2 O.” At the bottom of the white region are the light blue arrow for O subscript 2 and the dark blue arrow for H subscript 2 O pointing out. This is labeled “Unused Gases Out.” Black line segments extend upward from the light gray shaded regions. These line segments are connected by a horizontal segment that has a curly shape in a circle at the center. This shape is labeled “Electric Current.” In the left light gray shaded region above the yellow arrows is a red arrow pointing up, the label e superscript minus above it, and then another red arrow. The black line segment above this area also has the label e superscript minus. Where the line turns right to connect to the Electric Current shape is a right-facing red arrow. On the other side of the shape where the line turns downward to connect to the other light gray shaded region is a red downward-facing arrow. Below that arrow in the light gray region is the label e superscript minus, followed by a red down arrow, followed by another e superscript minus label that stops before the light blue arrow pointing in to the shaded area.

These types of fuel cells generally produce voltages of approximately 1.2 V. Compared to an internal combustion engine, the energy efficiency of a fuel cell using the same redox reaction is typically more than double (~20%–25% for an engine versus ~50%–75% for a fuel cell). Hydrogen fuel cells are commonly used on extended space missions, and prototypes for personal vehicles have been developed, though the technology remains relatively immature.

38.5 Corrosion

Learning Objectives

By the end of this section, you will be able to:

  • Define corrosion
  • List some of the methods used to prevent or slow corrosion

Corrosion is usually defined as the degradation of metals by a naturally occurring electrochemical process. The formation of rust on iron, tarnish on silver, and the blue-green patina that develops on copper are all examples of corrosion. The total cost of corrosion remediation in the United States is significant, with estimates in excess of half a trillion dollars a year.


Statue of Liberty: Changing Colors

The Statue of Liberty is a landmark every American recognizes. The Statue of Liberty is easily identified by its height, stance, and unique blue-green color (Figure 38.8). When this statue was first delivered from France, its appearance was not green. It was brown, the color of its copper “skin.” So how did the Statue of Liberty change colors? The change in appearance was a direct result of corrosion. The copper that is the primary component of the statue slowly underwent oxidation from the air. The oxidation-reduction reactions of copper metal in the environment occur in several steps. Copper metal is oxidized to copper(I) oxide (Cu2O), which is red, and then to copper(II) oxide, which is black


Coal, which was often high in sulfur, was burned extensively in the early part of the last century. As a result, atmospheric sulfur trioxide, carbon dioxide, and water all reacted with the CuO


These three compounds are responsible for the characteristic blue-green patina seen on the Statue of Liberty (and other outdoor copper structures). Fortunately, formation of patina creates a protective layer on the copper surface, preventing further corrosion of the underlying copper. The formation of the protective layer is called passivation, a phenomenon discussed further in another chapter of this text.

Figure 38.8

(a) The Statue of Liberty is covered with a copper skin, and was originally brown, as shown in this painting. (b) Exposure to the elements has resulted in the formation of the blue-green patina seen today.

This figure contains two photos of the Statue of Liberty. Photo a appears to be an antique photo which shows the original brown color of the copper covered statue. Photo b shows the blue-green appearance of the statue today. In both photos, the statue is shown atop a building, with a body of water in the background.

Perhaps the most familiar example of corrosion is the formation of rust on iron. Iron will rust when it is exposed to oxygen and water. Rust formation involves the creation of a galvanic cell at an iron surface, as illustrated in Figure 38.8. The relevant redox reactions are described by the following equations:

anode:Fe(s)Fe2+(aq)+2eEFe2+/Fe°=−0.44 Vcathode:O2(g)+4H+(aq)+4e2H2O(l)EO2/O2°=+1.23 Voverall:2Fe(s)+O2(g)+4H+(aq)2Fe2+(aq)+2H2O(l)Ecell°=+1.67 Vanode:Fe(s)Fe2+(aq)+2eEFe2+/Fe°=−0.44 Vcathode:O2(g)+4H+(aq)+4e2H2O(l)EO2/O2°=+1.23 Voverall:2Fe(s)+O2(g)+4H+(aq)2Fe2+(aq)+2H2O(l)Ecell°=+1.67 V

Further reaction of the iron(II) product in humid air results in the production of an iron(III) oxide hydrate known as rust:


The stoichiometry of the hydrate varies, as indicated by the use of x in the compound formula. Unlike the patina on copper, the formation of rust does not create a protective layer and so corrosion of the iron continues as the rust flakes off and exposes fresh iron to the atmosphere.

Figure 38.9

Corrosion can occur when a painted iron or steel surface is exposed to the environment by a scratch through the paint. A galvanic cell results that may be approximated by the simplified cell schematic Fe(s) | Fe2+(aq) ||O2(aq), H2O(l) | Fe(s).

A grey rectangle, labeled “iron,” is shown with thin purple layers, labeled “Paint layer,” at its upper and lower surfaces. A gap in the upper purple layer at the upper left of the diagram is labeled “Cathodic site.” A blue droplet labeled “water” is positioned on top of the gap. A curved arrow extends from a space above the droplet to the surface of the grey region and into the water droplet. The base of the arrow is labeled “O subscript 2” and the tip of the arrow is labeled “H subscript 2 O.” A gap to the right and on the bottom side of the grey region shows that some of the grey region is gone from the region beneath the purple layer. A water droplet covers this gap and extends into the open space in the grey rectangle. The label “F e superscript 2 positive” is at the center of the droplet. A curved arrow points from the edge of the grey area below to the label. A second curved arrow extends from the F e superscript 2 positive arrow to a rust brown chunk on the lower surface of the purple layer at the edge of the water droplet. A curved arrow extends from O subscript 2 outside the droplet into the droplet to the rust brown chunk. The grey region at the lower right portion of the diagram is labeled “Anodic site.” An arrow extends from the anodic site toward the cathodic site, which is labeled “e superscript negative.”

One way to keep iron from corroding is to keep it painted. The layer of paint prevents the water and oxygen necessary for rust formation from coming into contact with the iron. As long as the paint remains intact, the iron is protected from corrosion.

Other strategies include alloying the iron with other metals. For example, stainless steel is an alloy of iron containing a small amount of chromium. The chromium tends to collect near the surface, where it corrodes and forms a passivating an oxide layer that protects the iron.

Iron and other metals may also be protected from corrosion by galvanization, a process in which the metal to be protected is coated with a layer of a more readily oxidized metal, usually zinc. When the zinc layer is intact, it prevents air from contacting the underlying iron and thus prevents corrosion. If the zinc layer is breached by either corrosion or mechanical abrasion, the iron may still be protected from corrosion by a cathodic protection process, which is described in the next paragraph.

Another important way to protect metal is to make it the cathode in a galvanic cell. This is cathodic protection and can be used for metals other than just iron. For example, the rusting of underground iron storage tanks and pipes can be prevented or greatly reduced by connecting them to a more active metal such as zinc or magnesium (Figure 38.10). This is also used to protect the metal parts in water heaters. The more active metals (lower reduction potential) are called sacrificial anodes because as they get used up as they corrode (oxidize) at the anode. The metal being protected serves as the cathode for the reduction of oxygen in air, and so it simply serves to conduct (not react with) the electrons being transferred. When the anodes are properly monitored and periodically replaced, the useful lifetime of the iron storage tank can be greatly extended.

Figure 38.10

Cathodic protection is a useful approach to electrochemically preventing corrosion of underground storage tanks.

A diagram is shown of an underground storage tank system. Underground is a metal tank-like structure, labeled “Sacrificial anode” which is vertically oriented. M g is on the tank, followed by a right arrow, followed by M g superscript 2 plus. A black line extends upward from the center of the tank, but stays underground. A horizontal black line segment continues right underground. 2 e superscript minus is followed by an arrow that points right below the line segment. A vertical black line segment leads downward to a horizontal grey tank which is labeled “Object to be protected (cathode).” 2 e subscript minus is on the tank with an arrow pointing from it to the ground below the tank. Below that arrow is “2 H superscript plus plus O subscript 2 arrow 2 H subscript 2 O.”

38.6 Electrolysis

Learning Objectives

By the end of this section, you will be able to:

  • Describe the process of electrolysis
  • Compare the operation of electrolytic cells with that of galvanic cells
  • Perform stoichiometric calculations for electrolytic processes

Electrochemical cells in which spontaneous redox reactions take place (galvanic cells) have been the topic of discussion so far in this chapter. In these cells, electrical work is done by a redox system on its surroundings as electrons produced by the redox reaction are transferred through an external circuit. This final section of the chapter will address an alternative scenario in which an external circuit does work on a redox system by imposing a voltage sufficient to drive an otherwise nonspontaneous reaction, a process known as electrolysis. A familiar example of electrolysis is recharging a battery, which involves use of an external power source to drive the spontaneous (discharge) cell reaction in the reverse direction, restoring to some extent the composition of the half-cells and the voltage of the battery. Perhaps less familiar is the use of electrolysis in the refinement of metallic ores, the manufacture of commodity chemicals, and the electroplating of metallic coatings on various products (e.g., jewelry, utensils, auto parts). To illustrate the essential concepts of electrolysis, a few specific processes will be considered.

38.7 The Electrolysis of Molten Sodium Chloride

Metallic sodium, Na, and chlorine gas, Cl2, are used in numerous applications, and their industrial production relies on the large-scale electrolysis of molten sodium chloride, NaCl(l). The industrial process typically uses a Downs cell similar to the simplified illustration shown in Figure 38.11. The reactions associated with this process are:


The cell potential for the above process is negative, indicating the reaction as written (decomposition of liquid NaCl) is not spontaneous. To force this reaction, a positive potential of magnitude greater than the negative cell potential must be applied to the cell.

Figure 38.11

Cells of this sort (a cell for the electrolysis of molten sodium chloride) are used in the Downs process for production of sodium and chlorine, and they typically use iron cathodes and carbon anodes.

This diagram shows a tank containing a light blue liquid, labeled “Molten N a C l.” A vertical dark grey divider with small, evenly distributed dark dots, labeled “Porous screen” is located at the center of the tank dividing it into two halves. Dark grey bars are positioned at the center of each of the halves of the tank. The bar on the left, which is labeled “Anode” has green bubbles originating from it. The bar on the right which is labeled “Cathode” has light grey bubbles originating from it. An arrow points left from the center of the tank toward the anode, which is labeled “C l superscript negative.” An arrow points right from the center of the tank toward the cathode, which is labeled “N a superscript plus.” A line extends from the tops of the anode and cathode to a rectangle centrally placed above the tank which is labeled “Voltage source.” An arrow extends upward above the anode to the left of the line which is labeled “e superscript negative.” A plus symbol is located to the left of the voltage source and a negative sign it located to its right. An arrow points downward along the line segment leading to the cathode. This arrow is labeled “e superscript negative.” The left side of below the diagram is the label “2 C l superscript negative right pointing arrow C l subscript 2 ( g ) plus 2 e superscript negative.” At the right, below the diagram is the label “2 N a superscript plus plus 2 e superscript negative right pointing arrow 2 N a ( l ).”

38.8 The Electrolysis of Water

Water may be electrolytically decomposed in a cell similar to the one illustrated in Figure 38.12. To improve electrical conductivity without introducing a different redox species, the hydrogen ion concentration of the water is typically increased by addition of a strong acid. The redox processes associated with this cell are

anode:2H2O(l)O2(g)+4H+(aq)+4eEanode°=+1.229 Vcathode:2H+(aq)+2eH2(g)Ecathode°=0 V¯cell:2H2O(l)2H2(g)+O2(g)Ecell°=−1.229 Vanode:2H2O(l)O2(g)+4H+(aq)+4eEanode°=+1.229 Vcathode:2H+(aq)+2eH2(g)Ecathode°=0 V¯cell:2H2O(l)2H2(g)+O2(g)Ecell°=−1.229 V

Again, the cell potential as written is negative, indicating a nonspontaneous cell reaction that must be driven by imposing a cell voltage greater than +1.229 V. Keep in mind that standard electrode potentials are used to inform thermodynamic predictions here, though the cell is not operating under standard state conditions. Therefore, at best, calculated cell potentials should be considered ballpark estimates.

Figure 38.12

The electrolysis of water produces stoichiometric amounts of oxygen gas at the anode and hydrogen at the anode.

This figure shows an apparatus used for electrolysis. A central chamber with an open top has a vertical column extending below that is nearly full of a clear, colorless liquid, which is labeled “H subscript 2 O plus H subscript 2 S O subscript 4.” A horizontal tube in the apparatus connects the central region to vertical columns to the left and right, each of which has a valve or stopcock at the top and a stoppered bottom. On the left, the stopper at the bottom has a small brown square connected just above it in the liquid. The square is labeled “Anode plus.” A black wire extends from the stopper at the left to a rectangle which is labeled “Voltage source” on to the stopper at the right. The left side of the rectangle is labeled with a plus symbol and the right side is labeled with a negative sign. The stopper on the right also has a brown square connected to it which is in the liquid in the apparatus. This square is labeled “Cathode negative.” The level of the solution on the left arm or tube of the apparatus is significantly higher than the level of the right arm. Bubbles are present near the surface of the liquid on each side of the apparatus, with the bubbles labeled as “O subscript 2 ( g )” on the left and “H subscript 2 ( g )” on the right.

38.9 The Electrolysis of Aqueous Sodium Chloride

When aqueous solutions of ionic compounds are electrolyzed, the anode and cathode half-reactions may involve the electrolysis of either water species (H2O, H+, OH-) or solute species (the cations and anions of the compound). As an example, the electrolysis of aqueous sodium chloride could involve either of these two anode reactions:

(i)2Cl(aq)Cl2(g)+2eEanode°=+1.35827 V(ii)2H2O(l)O2(g)+4H+(aq)+4eEanode°=+1.229 V(i)2Cl(aq)Cl2(g)+2eEanode°=+1.35827 V(ii)2H2O(l)O2(g)+4H+(aq)+4eEanode°=+1.229 V

The standard electrode (reduction) potentials of these two half-reactions indicate water may be oxidized at a less negative/more positive potential (–1.229 V) than chloride ion (–1.358 V). Thermodynamics thus predicts that water would be more readily oxidized, though in practice it is observed that both water and chloride ion are oxidized under typical conditions, producing a mixture of oxygen and chlorine gas.

Turning attention to the cathode, the possibilities for reduction are:

(iii)2H+(aq)+2eH2(g)Ecathode°=0 V(iv)2H2O(l)+2eH2(g)+2OH(aq)Ecathode°=−0.8277 V(v)Na+(aq)+eNa(s)Ecathode°=−2.71 V(iii)2H+(aq)+2eH2(g)Ecathode°=0 V(iv)2H2O(l)+2eH2(g)+2OH(aq)Ecathode°=−0.8277 V(v)Na+(aq)+eNa(s)Ecathode°=−2.71 V

Comparison of these standard half-reaction potentials suggests the reduction of hydrogen ion is thermodynamically favored. However, in a neutral aqueous sodium chloride solution, the concentration of hydrogen ion is far below the standard state value of 1 M (approximately 10-7 M), and so the observed cathode reaction is actually reduction of water. The net cell reaction in this case is then

cell: 2H2O(l)+2Cl(aq)H2(g)+Cl2(g)+2OH(aq)Ecell°=−2.186 Vcell: 2H2O(l)+2Cl(aq)H2(g)+Cl2(g)+2OH(aq)Ecell°=−2.186 V

This electrolysis reaction is part of the chlor-alkali process used by industry to produce chlorine and sodium hydroxide (lye).



An important use for electrolytic cells is in electroplating. Electroplating results in a thin coating of one metal on top of a conducting surface. Reasons for electroplating include making the object more corrosion resistant, strengthening the surface, producing a more attractive finish, or for purifying metal. The metals commonly used in electroplating include cadmium, chromium, copper, gold, nickel, silver, and tin. Common consumer products include silver-plated or gold-plated tableware, chrome-plated automobile parts, and jewelry. The silver plating of eating utensils is used here to illustrate the process. (Figure 38.13).

Figure 38.13

This schematic shows an electrolytic cell for silver plating eating utensils.

This figure contains a diagram of an electrochemical cell. One beakers is shown that is just over half full. The beaker contains a clear, colorless solution that is labeled “A g N O subscript 3 ( a q ).” A silver strip is mostly submerged in the liquid on the left. This strip is labeled “Silver (anode).” The top of the strip is labeled with a red plus symbol. An arrow points right from the surface of the metal strip into the solution to the label “A g superscript plus” to the right. A spoon is similarly suspended in the solution and is labeled “Spoon (cathode).” It is labeled with a black negative sign on the tip of the spoon’s handle above the surface of the liquid. An arrow extends from the label “A g superscript plus” to the spoon on the right. A wire extends from the top of the spoon and the strip to a rectangle labeled “Voltage source.” An arrow points upward from silver strip which is labeled “e superscript negative.” Similarly, an arrow points down at the right to the surface of the spoon which is also labeled “e superscript negative.” A plus sign is shown just outside the voltage source to the left and a negative is shown to its right.

In the figure, the anode consists of a silver electrode, shown on the left. The cathode is located on the right and is the spoon, which is made from inexpensive metal. Both electrodes are immersed in a solution of silver nitrate. Applying a sufficient potential results in the oxidation of the silver anode

anode: Ag(s)Ag+(aq)+eanode: Ag(s)Ag+(aq)+e

and reduction of silver ion at the (spoon) cathode:

cathode: Ag+(aq)+eAg(s)cathode: Ag+(aq)+eAg(s)

The net result is the transfer of silver metal from the anode to the cathode. Several experimental factors must be carefully controlled to obtain high-quality silver coatings, including the exact composition of the electrolyte solution, the cell voltage applied, and the rate of the electrolysis reaction (electrical current).

38.10 Quantitative Aspects of Electrolysis

Electrical current is defined as the rate of flow for any charged species. Most relevant to this discussion is the flow of electrons. Current is measured in a composite unit called an ampere, defined as one coulomb per second (A = 1 C/s). The charge transferred, Q, by passage of a constant current, I, over a specified time interval, t, is then given by the simple mathematical product


When electrons are transferred during a redox process, the stoichiometry of the reaction may be used to derive the total amount of (electronic) charge involved. For example, the generic reduction process


involves the transfer of n mole of electrons. The charge transferred is, therefore,


where F is Faraday’s constant, the charge in coulombs for one mole of electrons. If the reaction takes place in an electrochemical cell, the current flow is conveniently measured, and it may be used to assist in stoichiometric calculations related to the cell reaction.

EXAMPLE 38.10.1

Converting Current to Moles of Electrons

In one process used for electroplating silver, a current of 10.23 A was passed through an electrolytic cell for exactly 1 hour. How many moles of electrons passed through the cell? What mass of silver was deposited at the cathode from the silver nitrate solution?


Faraday’s constant can be used to convert the charge (Q) into moles of electrons (n). The charge is the current (I) multiplied by the time
n=QF=10.23 Cs×1 hr×60 minhr×60 smin96,485 C/mol e=36,830 C96,485 C/mole=0.3817 mol en=QF=10.23 Cs×1 hr×60 minhr×60 smin96,485 C/mol e=36,830 C96,485 C/mole=0.3817 mol e

From the problem, the solution contains AgNO3, so the reaction at the cathode involves 1 mole of electrons for each mole of silver

cathode: Ag+(aq)+eAg(s)cathode: Ag+(aq)+eAg(s)

The atomic mass of silver is 107.9 g/mol, so

mass Ag=0.3817 mol e×1 mol Ag1 mol e×107.9 g Ag1 mol Ag=41.19 g Agmass Ag=0.3817 mol e×1 mol Ag1 mol e×107.9 g Ag1 mol Ag=41.19 g Ag

Check Your Learning

Aluminum metal can be made from aluminum(III) ions by electrolysis. What is the half-reaction at the cathode? What mass of aluminum metal would be recovered if a current of 25.0 A passed through the solution for 15.0 minutes?


Al3+(aq)+3eAl(s);Al3+(aq)+3eAl(s); 0.0777 mol Al = 2.10 g Al.

EXAMPLE 38.10.2

Time Required for Deposition

In one application, a 0.010-mm layer of chromium must be deposited on a part with a total surface area of 3.3 m2 from a solution of containing chromium(III) ions. How long would it take to deposit the layer of chromium if the current was 33.46 A? The density of chromium (metal) is 7.19 g/cm3.


First, compute the volume of chromium that must be produced (equal to the product of surface area and thickness):
volume=(0.010 mm×1 cm10 mm)×(3.3m2×(10,000cm21m2))=33 cm3volume=(0.010 mm×1 cm10 mm)×(3.3m2×(10,000cm21m2))=33 cm3

Use the computed volume and the provided density to calculate the molar amount of chromium required:

mass=volume×density=33cm3×7.19 gcm3=237 g Crmass=volume×density=33cm3×7.19 gcm3=237 g Cr
mol Cr=237 g Cr×1 mol Cr52.00 g Cr=4.56 mol Crmol Cr=237 g Cr×1 mol Cr52.00 g Cr=4.56 mol Cr

The stoichiometry of the chromium(III) reduction process requires three moles of electrons for each mole of chromium(0) produced, and so the total charge required is:

Q=4.56 mol Cr×3mol e1 mol Cr×96485 Cmol e=1.32×106CQ=4.56 mol Cr×3mol e1 mol Cr×96485 Cmol e=1.32×106C

Finally, if this charge is passed at a rate of 33.46 C/s, the required time is:

t=QI=1.32×106C33.46 C/s=3.95×104s=11.0 hrt=QI=1.32×106C33.46 C/s=3.95×104s=11.0 hr

Check Your Learning

What mass of zinc is required to galvanize the top of a 3.00 m ×× 5.50 m sheet of iron to a thickness of 0.100 mm of zinc? If the zinc comes from a solution of Zn(NO3)2 and the current is 25.5 A, how long will it take to galvanize the top of the iron? The density of zinc is 7.140 g/cm3.


11.8 kg Zn requires 382 hours.


Previous Citation(s)
Flowers, P., et al. (2019). Chemistry: Atoms First 2e. https://openstax.org/details/books/chemistry-atoms-first-2e (16.5-16.7)

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