Unit Five

Day 41: Electrolysis; Commercial Batteries

D41.1 Electrolysis

In an electrolytic cell, supplied electrical energy causes a nonspontaneous redox reaction to occur in a process known as electrolysis. An electrolytic cell is the opposite of a voltaic cell, where a spontaneous redox reaction produces electrical energy.

For example, consider electrolysis of molten sodium chloride. A simplified diagram of the electrolytic cell used for commercial manufacture of sodium metal and chlorine gas is shown in Figure 1. In the molten salt, sodium ions move toward the cathode and chloride ions move toward the anode (note that negative ions move through the circuit in the same overall direction as electrons). A porous screen allows movement of ions but not mixing of the product sodium metal and chlorine gas, which would react spontaneously upon contact.

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 no 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 “Current 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 current source and a negative sign is located to its right. An arrow points downward along the line segment leading to the cathode. This arrow is labeled “e superscript negative.” Below the left side of the diagram is the label “2 C l superscript negative ( a q ) 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 positive ( a q ) plus 2 e superscript negative right pointing arrow 2 N a ( l ).”
Figure 1. Passing an electric current through NaCl(l) decomposes the molten salt into sodium metal and chlorine gas. Care must be taken to keep the products separated to prevent the spontaneous re-formation of sodium chloride.

An external source of electric current forces electrons into the electrode in the cathode compartment, forcing the reduction half-reaction to occur:

Reduction (cathode): 2 × (Na+(aq) + e‾ ⟶ Na(s) )        E°Na+|Na = -2.714 V

(At the temperature of molten NaCl, sodium is a liquid, so E°Na+|Na(s) = -2.714 V is an approximation. E°Na+|Na(l) for liquid sodium is not available in the appendix.)

In the anode compartment, the oxidation half-reaction occurs:

Oxidation (anode): 2 Cl‾(aq) ⟶ Cl2(g) + 2e‾        E°Cl2|Cl‾ = +1.358 V

The electrons formed here are conducted through a wire to the the positive side of the voltage source, completing the electric circuit.

We can calculate E°cell for this electrolysis cell using the same method we used for voltaic cells; that is,

E°cell = E°right half-cell − E°left half-cell = E°cathode − E°anode = -2.714 V − (+1.358 V) = -4.072 V

The overall reaction is:

Overall: 2 Na+(aq) + 2 Cl‾(aq) ⟶ 2 Na(s) + Cl2(g)         E°cell = -4.072 V

The negative E°cell indicates that this reaction is strongly reactant-favored, and under standard-state conditions, the power supply must provide at least 4.1 V to cause the electrolysis reaction to occur. In practice, the applied voltages are higher due to inefficiencies in the process itself and also to help increase the rate of reaction.

With the transition from fossil fuels to renewable energy supplies, an important application of electrolysis is the “splitting” of water into hydrogen gas and oxygen gas (Figure 2). Electric energy from solar panels or wind turbines can be used to synthesize hydrogen for use as a fuel. For current to pass through the solution efficiently, there must be ions present. Hence, acid is typically added to the reaction solution to increase the concentration of ions in solution.

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.
Figure 2. Water decomposes into oxygen gas and hydrogen gas during electrolysis. Sulfuric acid is added to increase the concentration of H+ ions and the total number of ions in solution, but does not take part in the reaction. The volume of hydrogen gas collected is twice the volume of oxygen gas collected, because the reaction produces 2 mol H2 per 1 mol O2.

In 1-M acidic solution:

Oxidation (anode): 2H2O(l) ⟶ O2(g) + 4H+(aq) + 4e‾ E°anode = +1.229 V
Reduction (cathode): 2 × (2H+(aq) + 2e‾ ⟶ H2(g)) E°cathode = 0 V
Overall: 2H2O(l) ⟶ O2(g) + 2H2(g) E°cell = -1.229 V

At least 1.229 V is required to make this reactant-favored process occur in 1-M acidic solution.

Finally, consider what occurs during the electrolysis of 1-M aqueous potassium iodide solution at 25 °C. Present in the solution are H2O(l), K+(aq), and I‾(aq). This example differs from the previous examples because more than one species can be oxidized and more than one species can be reduced.

Considering the anode first, the possible oxidation reactions are

(i) 2I‾(aq) ⟶ I2(s) + 2e‾ E°anode = +0.535 V
(ii) 2H2O(l) ⟶ O2(g) + 4H+(aq) + 4e‾ E°anode = +1.229 V

(Oxidation of K+ to K2+ is not considered because K+ has a noble-gas electron configuration and a very high ionization energy, making it very difficult to oxidize. Although I2 is generated in aqueous solution, we approximate the E° by using the value for I2(s).)

Because E°cell = E°cathode – E°anode, the more positive the anode half-cell potential is, the more negative the cell potential would be. Therefore, iodide should be oxidized at the anode because it has a less positive half-cell potential. However, the pH of a 1-M KI solution is 7, so [H+] is far from standard-state conditions. Assuming that O2 is produced at 1 bar, applying the Nernst equation to half-reaction (ii) gives:

[latex]\begin{array}{rcl} E &=& {E^ \circ } - \dfrac{RT}{nF}\ln \left( \dfrac{1}{[\text{O}_2][\text{H}^+]^4}\right)\\[0.5em] &=& 1.229\text{ V} - \dfrac{\left(8.314 \frac{\text{J}}{\text{K}\cdot\text{mol}}\right) (298.15\;\text{K})}{4\left(96485\frac{\text{J}}{\text{V}\cdot\text{mol}}\right)} \ln \left(\dfrac{1}{(1)(1 \times 10^{-7})^4}\right)\\[0.5em] &=& 1.229\;{\text{ V}} - 0.414\;\text{V}\; =\; 0.815\;\text{V} \end{array}[/latex]

(Note that the reaction given in the standard half-cell potential table is the reduction reaction “O2(g) + 4H+(aq) + 4e‾ ⟶ 2H2O(l) E° = +1.229 V”, hence reaction quotient Q is expressed in accordance to the given reaction.)

The Nernst equation shows that Eanode = +0.815 V for reaction (ii) at pH = 7, which is still higher than E°anode for reaction (i). Therefore, reaction (i) is the process that occurs at the anode and I2 forms as a product.

Now consider the possible reactions at the cathode (reduction of I− is not considered because I− has a noble-gas electron configuration and it is not energetically favorable to add more electrons):

(iii) 2H2O(l) + 2e‾ ⟶ H2(g) + 2OH‾(aq) E°cathode = -0.8277 V
(ii) K+(aq) + e‾ ⟶ K(s) E°cathode = -2.925 V

For half-reaction (iii), we again need to apply the Nernst equation to calculate E at pH = 7, assuming H2 is produced at 1 bar:

[latex]\begin{array}{rcl} E &=& {E^ \circ } - \dfrac{RT}{nF}\ln \left( \dfrac{[\text{H}_2][\text{OH}^-]^2}{1}\right)\\[0.5em] &=& -0.8277\text{ V} - \dfrac{\left(8.314 \frac{\text{J}}{\text{K}\cdot\text{mol}}\right) (298.15\;\text{K})}{2\left(96485\frac{\text{J}}{\text{V}\cdot\text{mol}}\right)} \ln \left(\dfrac{(1)(1 \times 10^{-7})^2}{1}\right)\\[0.5em] &=& -0.8277\;{\text{ V}} - (-0.4141\;\text{V})\; =\; -0.4136\;\text{V} \end{array}[/latex]

Hence, reduction of water, with Ecathode = -0.4136 V at pH = 7, is much more likely to occur than reduction of K+(aq) with E°cathode = -2.925 V. (This conclusion is supported by the fact that potassium metal reacts vigorously with water to generate K+(aq), hydrogen gas, and hydroxide ions, so if K(s) formed it would immediately react with water.)

The overall reaction is then:

Oxidation (anode): 2I‾(aq) ⟶ I2(s) + 2e‾ E°anode = +0.535 V
Reduction (cathode): 2H2O(l) + 2e‾ ⟶ H2(g) + 2OH‾(aq) Ecathode = -0.4136 V
Overall: 2H2O(l) + 2I‾(aq) ⟶ H2(g) + I2(s) + 2OH‾(aq) Ecell = -0.949 V

Electroplating

An important use for electrolytic cells is electroplating, which results in a thin coating of metal on top of a conducting surface. Metals typically used in electroplating include cadmium, chromium, copper, gold, nickel, silver, and tin. As an example of electroplating, consider how silver-plated tableware is produced (Figure 3).

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.
Figure 3. The metal spoon is connected to the negative terminal of the voltage source and acts as the cathode. The anode is a silver electrode. Both electrodes are immersed in a silver nitrate solution. When electric current is passed through the solution, the net result is that silver metal is removed from the anode and deposited on the spoon.

The anode consists of a silver electrode. The cathode is a spoon made from a less expensive metal. Both electrodes are immersed in a solution of silver nitrate. As the potential from the voltage source is increased, current flows. Silver metal is lost at the anode as it goes into solution:

anode:        Ag(s) ⟶ Ag+(aq) + e−

The mass of the cathode increases as silver ions from the solution are deposited onto the spoon:

cathode:      Ag+(aq) + e− ⟶ Ag(s)

The net result is the transfer of silver metal from the anode to the cathode. The quality of the electroplated object depends on the thickness of the deposited silver and the rate of deposition.

Quantitative Aspects of Electrolysis

The quantity of current that flows in an electrolytic cell is dictated by the amount (mol) of electrons transferred in a redox reaction, which is in turn related to quantities of reactants and products via reaction stoichiometry. Recall that current, I, is related to the total charge, Q:

[latex]I = \dfrac{Q}{t} \;\;\;\;\; \left(\text{SI units:}\;\ A = \dfrac{C}{s}\right)[/latex]

Hence:

Q = (amount e− transferred) × F = I × t

where F is the Faraday constant.

Exercise 1: Electroplating

D41.2 Commercial Batteries

Many of the devices we use every day, such as laptops and smartphones, are powered by batteries. A battery is an electrochemical cell or series of cells that produces an electric current. In principle, any voltaic cell can be used as a battery. An ideal battery would never run down/drain, produce a constant voltage, and be capable of withstanding environmental extremes of temperature and humidity. Real batteries strike a balance between ideal characteristics and practical limitations.

For example, the mass of a car-starter battery is about 18 kg or ~1% of the mass of an average car. This type of battery would supply nearly unlimited energy if used in a smartphone, but would be completely impractical because of its mass and size. Thus, no single battery is “best” and different batteries are selected for particular applications, keeping things like its mass, cost, reliability, and current capacity in mind.

There are two basic types of batteries: primary and secondary. A few batteries of each type are described next.

D41.3 Primary Batteries

Primary batteries are single-use batteries that cannot be recharged.

Zinc-Carbon Battery

A common primary battery is the dry cell (Figure 4), which is a zinc-carbon battery. The zinc serves as both a container and the negative electrode. The positive electrode is a rod made of carbon that is surrounded by a paste of manganese(IV) oxide, zinc chloride, ammonium chloride, carbon powder, and a small quantity of water.

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.”
Figure 4. A cross section of a flashlight battery, a zinc-carbon dry cell.

The reaction at the anode can be represented as the oxidation of zinc:

Zn(s) ⟶ Zn2+(aq) + 2 e‾     E°anode = -0.763 V

The reaction at the cathode is more complicated, in part because more than one reduction reaction is occurring. The series of reactions that occurs at the cathode is approximately:

2MnO2(s) + 2NH4Cl(aq) + 2e‾ ⟶ Mn2O3(s) + 2NH3(aq) + H2O(l) + 2Cl‾(aq)

The overall reaction for the zinc–carbon battery can be represented as:

2MnO2(s) + 2 NH4Cl(aq) + Zn(s) ⟶ Mn2O3(s) + 2NH3(aq) + H2O(l) + 2 Cl‾(aq) + Zn2+(aq)

The cell potential is about 1.5 V initially, and decreases as the battery is used. As the zinc container oxidizes, its contents eventually leak out, so this type of battery should not be left in any electrical device for extended periods.

The voltage delivered by a battery is the same regardless of the size of a battery. For this reason, D, C, A, AA, and AAA batteries all have the same voltage. However, larger batteries can deliver more moles of electrons and will therefore last longer if powering the same device.

Alkaline Battery

Alkaline batteries (Figure 5) were developed in the 1950s partly to address some of the performance issues with zinc–carbon dry cells, and are manufactured to be their exact replacements. As their name suggests, these types of batteries use alkaline electrolytes, often potassium hydroxide.

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).”
Figure 5. Alkaline batteries were designed as direct replacements for zinc-carbon batteries.

The reactions are:

Oxidation (anode): Zn(s) + 2OH‾(aq) ⟶ ZnO(s) + H2O(l) + 2e‾ E°anode = -1.28 V
Reduction (cathode): 2MnO2(s) + H2O(l) + 2e‾ ⟶ Mn2O3(s) + 2OH‾(aq) E°cathode = +0.15 V
overall: Zn(s) + 2MnO2(s) ⟶ ZnO(s) + Mn2O3(s) E°cell = +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 sometimes leak potassium hydroxide, so these should also 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.

D41.4 Secondary Batteries

Secondary batteries are rechargeable; that is, the reaction that powers the battery can be reversed so that the original reactants can be regenerated. Secondary batteries are found in smartphones, electronic tablets, automobiles, and many other devices.

Lead-Acid Battery

The lead-acid battery (Figure 6) is the type of secondary battery used to start gasoline-powered automobiles. It is inexpensive and capable of producing the high current required by the starter motors when starting a car. They are heavy because of lead’s high density, they contain highly corrosive concentrated sulfuric acid, and must be disposed of properly to avoid lead-poisoning hazards. But they can produce a lot of current in a short time so for certain applications they are the best choice.

The reactions for a lead acid battery are:

Oxidation (anode): Pb(s) + HSO4‾(aq) ⟶ PbSO4(s) + H+(aq) + 2e‾ E°anode = -0.359 V
Reduction (cathode): PbO2(s) + HSO4‾(aq) + 3H+(aq) + 2e‾ ⟶ PbSO4(s) + 2H2O(l) E°cathode = +1.690 V
overall: Pb(s) + PbO2(s) + 2H2SO4(aq) ⟶ 2PbSO4(s) + 2H2O(l) E°cell = +2.049 V

Each cell produces 2.05 V, so six cells can be connected in series to produce a 12-V car battery.

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).”
Figure 6. The lead acid battery in an automobile consists of six cells connected in series to give 12 V. The low cost and high current output makes the battery suitable for providing power for a car’s starter motor.

In each cell, the lead electrodes are immersed in sulfuric acid. The anodes are spongy lead metal and the cathodes are lead impregnated with lead oxide. As the battery is discharged, a powder of PbSO4 forms on the electrodes. When a lead-acid battery is recharged by a car’s alternator, electrons are forced to flow in the opposite direction which reverses the reactions at anode and cathode, in other words, the cell undergoes electrolysis reactions to replenish the substances that have reacted away.

Practically, the concentrated sulfuric acid becomes quite viscous when the temperature is low, inhibiting the flow of ions between the plates and reducing the current that can be delivered. This effect is well-known to anyone who has had difficulty starting a car in cold weather. These batteries also tend to slowly self-discharge, so a car left idle for several weeks might be unable to start. And after thousands of discharge-charge cycles, PbSO4 that does not get converted to PbO2 gradually changes to an inert form which limits the battery capacity. Also, “fast” charging causes rapid evolution of potentially explosive H2 gas from the water in the electrolyte (electrolysis of water); the gas bubbles form on the lead surface and can tear PbO2 off the electrodes. Eventually enough solid material accumulates at the bottom of the electrolyte to short-circuit the battery, leading to its permanent demise.

Exercise 2: Lead-acid Batteries

Lithium Ion Battery

Lithium ion batteries (Figure 7) are among the most popular rechargeable batteries and are used in many portable electronic devices because their advantages outweigh the disadvantage of higher cost. In a typical Li-ion battery the reactions are:

Oxidation (anode): LiCoO2 ⟶ Li1-xCoO2 + xLi+ + xe‾
Reduction (cathode): xLi+ + xC6 + xe‾ ⟶ xLiC6
overall: LiCoO2 + xC6 ⟶ Li1-xCoO2 + xLiC6

(x is no more than about 0.5.) The battery voltage is about 3.7 V.

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.
Figure 7. In a lithium ion battery, charge flows between the electrodes as the lithium ions move between the anode and cathode.

Lithium batteries are popular because they can provide a large amount of 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.

Exercise 3: Lithium-Ion Batteries

D41.5 Fuel Cells

Suppose a voltaic cell is constructed such that the substance that is oxidized at the anode and the substance that is reduced at the cathode (the reactants in the overall redox reaction) are both supplied continuously. Such a battery would never run down because reactant concentrations or partial pressures would never decrease. Such a device is a fuel cell, which produces electricity as long as fuel is available. Hydrogen fuel cells have been used to supply power for satellites, space capsules, automobiles, boats, and submarines (Figure 8).

A diagram is shown of a hydrogen fuel cell. At the center is a narrow vertical rectangle which is shaded tan and labeled “Electrolyte.” To the right is a slightly wider and shorter green rectangle which is labeled “Cathode.” To the left is a pale blue rectangle of the same size which is labeled “Anode.” White rectangles, each with an inlet at the top and an outlet at the bottom are at the right and left sides, attached to the green and blue rectangles. On the right side O subscript 2 enters at the top, moves inward and along the interface with the green region, and exits to the lower right. H subscript 2 O also exits at the lower right. Pale gray diatomic O subscript 2 molecules move through this region and at the bottom are converted to a single pale gray atom with two smaller bright red atoms, H 2 O molecules. A similar pathway is on the left, allowing entry of H subscript 2 from the upper left along the interface with the blue rectangle, allowing for the exit of H subscript 2 out to the lower left of the diagram. Small red diatomic H subscript 2 molecules ar shown in this region. Black line segments extend upward from the blue anode and purple cathode regions. These line segments are connected by a horizontal segment that has a yellow light-bulb shape at the center. H subscript 2 molecules travel into the anode where they lose electrons and become single, small, red circles labeled H superscript plus. H superscript plus ions are shown traveling though the electrolyte to the green cathode where they interact with O subscript 2 molecules to form water molecules. The electrolyte is a special membrane permeable to H superscript plus but not to electrons.
Figure 8. In this schematic of a hydrogen-oxygen proton-exchange fuel cell, oxygen from the air reacts with hydrogen, producing water and electricity.

In a hydrogen-oxygen proton-exchange fuel cell, the cell potential is about 1 V, and the reactions involved are:

Oxidation (anode): 2 × (H2(g) ⟶ 2H+(aq) + 2e‾)
Reduction (cathode): O2(g) + 4H+(aq) + 4e‾ ⟶ 2H2O(l)
overall: O2(g) + 2H2(g) ⟶ 2H2O(l)

The efficiency of fuel cells is typically about 40-60%, which is higher than the typical internal combustion engine (25-35%). Moreover, in the case of the hydrogen fuel cell, nearly pure water is produced as exhaust. Currently, fuel cells are comparably more expensive and contain features that may cause a higher failure rate.

Podia Question

Consider this graph, which shows cell potential on the vertical axis and fraction of reactants remaining (battery life remaining) on the horizontal axis for a commercial battery.

  • Write an explanation in appropriate scientific language for why the graph has the shape it has.
  • Describe why the shape of this graph is a positive feature when batteries power devices such as smartphones and laptop computers.

Two days before the next whole-class session, this Podia question will become live on Podia, where you can submit your answer.

Comments. If you found any inconsistencies, errors, or other things you would like to report about this module, please use this link to report them. A similar link will be included in each day’s material. We appreciate your comments.

License

Icon for the Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License

Chemistry 109, Fall 2020 Copyright © by John Moore; Jia Zhou; and Etienne Garand is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.