ElectrolysisAs discussed in Section 10.6, corrosion is the outcome of a spontaneous redox reac-tion. An impressed current (one method that is used to provide cathodic protection) can stop corrosion by reversing the spontaneous corrosion reaction and driving it in the non-spontaneous direction. An impressed current passes an electric current through the structure being protected, such as a pipeline. An external electrical source supplies the current.

Other non-spontaneous redox reactions can be forced to occur in a similar way. One example is recharging the secondary cells in a battery. A battery uses a sponta-neous redox reaction to produce an electric current. Recharging a battery involves applying an external energy source to drive the redox reaction in the reverse, non-spontaneous direction. For example, a lead storage battery in a car can be recharged by applying a current to it (Figure 1). In this section, you will learn how some useful non-spontaneous reactions can be forced to occur. Many of these reactions have important consumer and industrial applications.

10.7

electrolytic cell a cell that uses electrical energy to produce a chemical change that would not occur spontaneously

electrolysis the application of current through a cell to produce a chemical change

Electrolytic CellsA galvanic cell produces a current when an oxidation–reduction reaction happens spontaneously. Another kind of cell, called an electrolytic cell, uses a current from an outside source to make an electrolytic reaction occur. In a process called electrolysis, the current is forced through the cell to produce a chemical change for which the cell potential is negative. That is, electrical energy causes an otherwise non-spontaneous chemical reaction to occur. Electrolysis has great practical importance. For example, electrolysis is used to recharge a battery, produce aluminum metal, and cover objects with a layer of chrome.

Figure 1  A car with a “dead” battery can be restarted by connecting the dead battery to a working battery. It is important to follow the manufacturer’s directions when setting up jumper cables. Otherwise, a spark might cause the hydrogen gas that is produced to explode.

10.7 Electrolysis 663NEL

7924_Chem_CH10.indd 663 5/4/12 3:50 PM

Electrolysis of waterIn Section 10.3, you learned that the reaction in which hydrogen and oxygen com-bine spontaneously to form water can be used to produce an electric current in a fuel cell. Th e reverse reaction is non-spontaneous and is an example of electrolysis. Electric current must be supplied to carry out the electrolysis of water. During electrolysis, water is oxidized at the anode and reduced at the cathode. Th erefore, water is both the oxidizing agent and the reducing agent in this reaction. Only two equations in Table 1, Appendix B7, involve just water as an oxidizing agent or a reducing agent:

O2(g) 1 4 H1(aq) 1 4 e2 S 2 H2O(l) E°r 5 11.23 V2 H2O(l) 1 2 e2 S H2(g) 1 2 OH2(aq) E°r 5 20.83 V

Note that the water in these two equations forms an upward diagonal to the right. This means that the electrolysis reaction is non-spontaneous. However, we can force it to occur by applying a current. The reaction equations can be written as follows:

Figure 2  (a) A galvanic cell produces a current from a spontaneous redox reaction: Zn 1s2 1 Cu21 1aq2 S Zn21 1aq2 1 Cu 1s2  (b) An electrolytic cell requires a power source to drive the opposite, non-spontaneous reaction: Cu 1s2 1 Zn21 1aq2 S Cu21 1aq2 1 Zn 1s2

��

cathode

1.0 mol/L Zn2�

solution1.0 mol/L Cu2�

solution

anode

Zn(s) Cu(s)

e�� �

e�

Cu2�(aq)SO4

2�(aq)

cations

anions

Zn(s) h Zn2�(aq) + 2 e�

(oxidation)Cu2�(aq) + 2 e� h Cu(s)

(reduction)Zn2�(aq) + 2 e� h Zn(s)

(reduction)(a)

Galvanic cell

(b) Cu(s) h Cu2�(aq) + 2 e�

(oxidation)

Zn2�(aq)SO4

2�(aq)

anode

1.0 mol/L Zn2�

solution1.0 mol/L Cu2�

solution

cathode

Zn(s) Cu(s)

e� e�

Cu2�(aq)SO4

2�(aq)

cations

power supplygreater than 1.10 V

anions

Zn2�(aq)SO4

2�(aq)

Electrolytic cell

Figure 2 shows the diff erences between a galvanic cell and an electrolytic cell. Th e galvanic cell (Figure 2(a)) runs spontaneously to produce 1.10 V.

Anode half-reaction equation: Zn 1s2 S Zn21 1aq2 1 2 e2

Cathode half-reaction equation: Cu21 1aq2 1 2 e2 S Cu 1s2Figure 2(b) shows an external energy source forcing electrons through the cell in

the direction opposite to that shown in Figure 2(a). Th is requires an external potential greater than 1.10 V, because the external potential must be applied in opposition to the natural cell potential. Th e device in Figure 2(b) is an electrolytic cell. Notice that the anode and the cathode are reversed in Figures 2(a) and 2(b). In a galvanic cell, zinc is the anode. In an electrolytic cell, copper is the anode. Also, ions fl ow through the salt bridge in opposite directions in the two cells.

664 Chapter 10 • Electrochemical Cells NEL

7924_Chem_CH10.indd 664 5/4/12 3:50 PM

Anode half-reaction equation: 2 H2O(l) S O2(g) 1 4 H1(aq) 1 4 e2

Cathode half-reaction equation: 4 H2O(l) 1 4 e2 S 2 H2(g) 1 4 OH2(aq) [multiply by 2 to balance electrons]Net ionic equation: 6 H2O 1l2 S 2 H2 1g2 1 O2 1g2 1 4 H1 1aq2 1 4 OH2 1aq2

4 H2O 1l2or 2 H2O 1l2 S 2 H2 1g2 1 O2 1g2

The standard cell potential for the electrolysis of water can be determined using the same equation that is used for galvanic cells:

∆E°r (cell) 5 E°r (cathode) 2 E°r (anode)

5 20.83 V2 (11.23 V) ∆E°r (cell) 5 22.06 VIf you tried to electrolyze pure water by using platinum electrodes connected to

a 6 V battery, you would observe no reaction. This is because pure water contains so few ions that almost no current can flow. However, dissolving even a small amount of an electrolyte, such as a soluble salt or a strong acid, in the water causes the immediate production of bubbles of hydrogen and oxygen (Figure 3). The electrolysis of water supplies oxygen to submarine crews and to astronauts in the International Space Station.

Comparing Galvanic and Electrolytic CellsA galvanic cell and an electrolytic cell are both types of electrochemical cell. They have some characteristics in common and some differences. The main characteristics of galvanic cells and electrolytic cells are compared in Table 1.

A galvanic cell converts chemical energy into electrical energy through a sponta-neous chemical reaction that takes place at the electrodes. Oxidation occurs at the anode, giving the anode a negative charge. Reduction occurs at the cathode, giving the cathode a positive charge. A wire connects the anode and the cathode, allowing electrons to move from the anode to the cathode. The electrodes are also connected by a salt bridge or a similar ionic conductor, which permits ions to flow between them. The flow of ions prevents a charge from building up in one of the electrodes. Since the redox reaction that takes place in a galvanic cell is spontaneous, a galvanic cell can be used as an energy source for an electrical device.

In contrast, an electrolytic cell requires an external energy source to operate. The basis for an electrolytic cell is a non-spontaneous redox reaction. Therefore, energy must be provided to drive the reaction. Like a galvanic cell, an electrolytic cell has an anode (the site of oxidation) and a cathode (the site of reduction). Electrons leave the anode along the wire, pass through the source of electrical energy, and travel to the cathode.

Table 1  Comparison of Galvanic and Electrolytic Cells 

Galvanic cell Electrolytic cell

involves a spontaneous redox reaction involves a non-spontaneous redox reaction

produces electrical energy requires electrical energy to drive the reaction

uses a salt bridge to prevent a charge buildup may or may not use a salt bridge

results in oxidation at the anode and reduction at the cathode 

results in oxidation at the anode and reduction at the cathode

The electrode that would be the anode of a galvanic cell is the cathode of an electrolytic cell.

The electrode that would be the cathode of a galvanic cell is the anode of an electrolytic cell.

electrochemical cell a general term that is used to refer to both a galvanic cell and an electrolytic cell

Figure 3  The electrolysis of water produces hydrogen gas at the cathode (on the right) and oxygen at the anode (on the left). Note that twice as much hydrogen as oxygen is produced.

g

10.7 Electrolysis 665NEL

7924_Chem_CH10.indd 665 5/4/12 3:50 PM

Commercial Electrolytic ProcessesAn important characteristic of metals is their ability to donate electrons (become oxidized) to form ions. Since metals are usually good reducing agents, most are found in nature in ores. Ores are mixtures of ionic compounds, which oft en contain oxide, sulfi de, and silicate anions. Th e noble metals, such as gold, silver, and platinum, are more diffi cult to oxidize and are more oft en found as pure metals.

Electrolysis of Aqueous SolutionsMany industries use electrolytic cells. One important application of electrolysis involves plating one metal onto another metal. For example, in Section 10.6 you learned that an eff ective way to prevent iron from corroding is to coat it with zinc. In this process, called galvanizing, the metal object to be plated is made the cathode of the electrolytic cell by attaching it to the positive terminal of the electrical energy supply (Figure 4). An inert electrode, such as graphite, is oft en used as the anode. Th e electrodes are then placed in an aqueous solution that contains cations of the metal to be plated onto the cathode.

For galvanizing, a zinc sulfate solution is used. As the cell operates, a shiny layer of zinc precipitates onto the iron cathode. Bubbles of oxygen rise from the anode, sug-gesting that the water is oxidized. Th e reaction equations are given below:

Anode half-reaction equation: 2 H2O(l) S O2(g) 1 4 H1(aq) 1 4 e2

Cathode half-reaction equation: 2 Zn21(aq) 1 4 e2 S 2 Zn(s) [multiply by 2 to balance electrons]Net ionic equation: 2 H2O(l) 1 2 Zn21(aq) S

2 Zn(s) 1 O2(g) 1 4 H1(aq)

4 H2O 1l2Th e standard cell potential for this cell is∆E°r (cell) 5 20.76 V2 (11.23 V)

∆E°r (cell) 5 21.99 VIn practice, a potential diff erence of more than 1.99 V must be applied for zinc to

plate onto iron.

mini Investigation

In this investigation, you will use electrical energy to make a non-spontaneous reaction occur. The graphite in pencil “lead” serves as the electrode for both anode and cathode half-reactions.

Equipment and Materials:  chemical safety goggles; lab apron; Petri dish; 2 small pencils sharpened at both ends; 2 connecting wires with alligator clips; 9 V battery; 0.5 mol/L potassium iodide, KI(aq); dropper bottle of phenolphthalein indicator

  1.  Put on your chemical safety goggles and lab apron.

  2.  Half-fi ll the Petri dish with the potassium iodide solution.  

  3.  Add about 10 drops of phenolphthalein to the solution.  

  4.  Place one end of each pencil in the dish. Keep the pencils far apart from each other. 

  5.  Use the connecting wires to connect the dry end of each pencil to a terminal on the battery. 

  6.  After about 30 s, disconnect the pencils from the battery. 

  7.  Look for evidence of chemical change. Remove each pencil, and check for an odour.   

  8.  Dispose of the contents of the dish according to your teacher’s instructions.

  A.   Identify the anode and the cathode of this cell.  T/I

  B.   What evidence of chemical change did you observe while the pencils were connected?  T/I

  C.  Based on your evidence, predict the identity of one substance that was produced at the anode and one substance that was produced at the cathode. Justify your predictions.  T/I

Pencil Electrolysis

Skills: Performing, Observing, Analyzing

mini Investigation

Figure 4  A model of the electrolytic cell that is used industrially to electroplate zinc onto iron

ee

power supply

solution containing Zn2 (aq) ions

cations

anions cathode(iron)

anode(graphite)

eSKILLS

HANDBOOK A1

666 Chapter 10 • Electrochemical Cells NEL

7924_Chem_CH10.indd 666 5/4/12 3:50 PM

Aluminum Production Since aluminum is a very active metal, it occurs naturally as an oxide in an ore called bauxite. The production of aluminum metal from its ore proved to be more difficult than the production of most other metals. In 1854, a process for producing metallic aluminum using sodium was discovered. However, aluminum remained extremely expensive.

In 1886, Charles M. Hall in the United States and Paul Héroult in France almost simul-taneously discovered a practical electrolytic process for producing aluminum. The key factor in the Hall–Héroult process is the use of molten cryolite, Na3AlF6(l), as the solvent for the aluminum oxide (also called alumina). Today, aluminum is relatively inexpensive and has a wide variety of applications, such as building materials (Figure 5).

Electrolysis is possible only if ions can move to the electrodes. A common method for making ions mobile is to dissolve the substance to be electrolyzed in water. However, water cannot be used for the electrolysis of the aluminum salt because water is more easily reduced than aluminum ions:

2 H2O 1l2 1 2 e2 S H2 1g2 1 OH2 1aq2 E°r 5 20.83 V

Al31 1aq2 1 3 e2 S Al 1s2 E°r 5 21.66 V

Thus, aluminum metal cannot be plated out of an aqueous solution of aluminum ions.Generally, ions can be made mobile by melting the salt. However, the melting

point of aluminum oxide, Al2O3, is much too high (2050 °C) for electrolysis of the molten oxide to be practical. A mixture of aluminum oxide and cryolite, however, has a melting point of 1000 °C, and the resulting molten mixture can be electrolyzed to obtain aluminum metal. As a result of the Hall–Héroult discovery, the price of alu-minum plunged and its widespread use became economically feasible. The industrial production of aluminum still uses the Hall–Héroult process.

Bauxite is not pure aluminum oxide. It also contains oxides of iron, silicon, and titanium, along with various silicate materials. To obtain the pure hydrated alu-minum oxide, Al2O3

# nH2O, the bauxite is treated with aqueous sodium hydroxide. Aluminum oxide is amphoteric and thus reacts in a basic solution:

Al2O3 1s2 1 2 OH2 1aq2 S 2 AlO22 1aq2 1 H2O 1l2

The other metal oxides, which are basic, remain as solids. (Recall Section 8.6.) The solution that contains the aluminate ion, AlO2

2(aq), is separated from the sludge of the other oxides and is acidified with carbon dioxide gas, causing the hydrated alumina to re-form:

2 CO2 1g2 1 2 AlO22 1aq2 1 1n 1 12H2O 1l2 S 2 HCO3

2 1aq2 1 Al2O3# nH2O(s)

Then, the purified aluminum oxide is mixed with cryolite, Na3AlF6, and melted at about 1000 °C. The dissociated aluminum ions are reduced to aluminum metal in an electrolytic cell (Figure 6).

Figure 5  The low mass, the lustre, and the corrosion resistance of aluminum make it a useful material for cladding the outsides of buildings.

Figure 6  This electrolytic cell is used to produce aluminum by the Hall–Héroult process. Since molten aluminum is denser than the molten mixture of cryolite and alumina, the metal settles to the bottom of the cell and is drawn off periodically. 

carbon dioxideformed at anodes

carbon-linediron tank

molten Al2O3/Na3AlF6(l) mixture

to externalpower supply

anodes(graphite rods)

molten aluminum

plug

10.7 Electrolysis 667NEL

436900_Chem_CH10.indd 667 5/9/12 9:56 AM

The electrolyte solution in the Hall-Héroult cell includes many different ions that contain aluminum, so the chemistry is not completely clear. However, aluminum oxide probably reacts with the cryolite anion:

Al2O3 1s2 1 4 AlF632 1l2 S 3 Al2OF6

22 1l2 1 6 F2 1l2 The following half-cell reactions are also thought to occur:

Cathode half-reaction equation: AlF632 1l2 1 3 e2 S Al 1s2 1 6 F2 1l2

Anode half-reaction equation: 2 Al2OF6

22 1l2 1 12 F2 1l2 1 C 1s2 S4 AlF6

32 1l2 1 CO2 1g2 1 4 e2

The overall cell reaction equation can be written as

2 Al2O3 1s2 1 3 C 1s2 S 4 Al 1s2 1 3 CO2 1g2The aluminum produced in this electrolytic process is 99.5 % pure. Most applications of aluminum use alloys that include other metals. For example,

to be useful as a structural material, aluminum is alloyed with metals such as zinc (for trailer and aircraft construction) and manganese (for cooking utensils, storage tanks, and highway signs).

Unfortunately, the production of aluminum has negative effects on the environ-ment. Since a great deal of energy is consumed during the process of aluminum refining, aluminum production contributes to greenhouse gas emissions, such as perfluorocarbons (PFCs), polycyclic aromatic hydrocarbon (PAH), fluoride, F–(aq), sulfur dioxide, SO2(g), and carbon dioxide, CO2(g). Placing aluminum smelting plants near renewable sources of electrical energy, such as waterfalls, reduces the emissions from electricity production. However, the consumption of the carbon elec-trodes during the electrolytic process also produces PAH and carbon dioxide.

Recycling used aluminum requires only about 5 % of the energy that is used to make new aluminum. Effective recycling programs not only reduce manufacturing costs, but also reduce the harmful emissions associated with producing aluminum.

Electrorefining MetalsThe purification of metals is another important application of electrolysis. Impure copper from the chemical reduction of copper ore is cast into large slabs that serve as the anodes for electrolytic cells. Aqueous copper sulfate is the electrolyte. Thin sheets of ultrapure copper function as the cathodes (Figure 7). When a potential difference is applied across the electrodes, only copper is deposited on these cathodes. All the

Figure 7  Ultrapure copper sheets, which serve as the cathodes of an electrolytic cell, are lowered between slabs of impure copper, which serve as the anodes. It takes about four weeks for the anodes to dissolve and for the pure copper to be deposited on the cathodes.

668 Chapter 10 • Electrochemical Cells NEL

7924_Chem_CH10.indd 668 5/4/12 3:50 PM

Figure 9  Silver plating is often used to beautify and protect cutlery and tableware, such as this elegant tea set. 

cathodeanode

e

e

e e

power supply

Ag Ag

AgAg

Figure 8  In an electroplating cell, the item to be plated (the spoon) is the cathode and a silver bar is the anode. Silver is plated out at the cathode: Ag1 1aq2 1 e2 S  Ag 1s2Note that a salt bridge is not needed because silver ions are involved at both electrodes.

other impurities are left as sludge in the cell. Ultrapure copper is important for the production of semiconductors and other electronic components.

Th e main anode half-reaction equation can be written as follows:

Cu 1s2 S Cu21 1aq2 1 2e2

Other metals, such as zinc and iron, are also oxidized from the impure anode:

Zn 1s2 S Zn21 1aq2 1 2 e2 and Fe 1s2 S Fe21 1 2e2

Noble metal impurities in the anode, such as silver, gold, and platinum, are not oxidized at the voltage used. Th ey fall to the bottom of the cell to form a sludge, which is then processed to remove them. Th e copper(II) ions from the solution are reduced and deposited onto the cathode:

Anode half-reaction equation: Cu21 1aq2 1 2 e2 S Cu 1s2Th is process produces copper metal that is 99.95 % pure. High-purity gold is also produced by electrorefi ning.

Electroplating MetalsMetals that readily corrode can oft en be protected by applying a thin coating, or “plating,” of a metal that resists corrosion. Examples are “tin” cans (which are actually steel cans with a thin coating of tin), chrome-plated steel car bumpers, and silver-plated jewellery and decorative items.

An object can be plated by making it the cathode in a tank that contains ions of the plating metal. Silver plating is an example of this electrolytic process (Figure 8). Th e silver plating solution contains ions that form complexes with the silver ion. By lowering the concentration of available silver ions, a smooth, even coating of silver forms over the surface of the object (Figure 9).  CAREER LINK

Metal plating processes have brought many benefi ts but also pose a range of envi-ronmental concerns. Heavy metals, such as chromium, lead, and cadmium, are used in these processes. Cyanide is a poisonous compound that is used in plating baths. Th e release of any of these toxic metals or their ions into the environment can cause serious health problems for organisms and persistent damage to ecosystems.

10.7 Electrolysis 669NEL

7924_Chem_CH10.indd 669 5/4/12 3:51 PM

Questions

1. Explain, using a graphic organizer of your choice, the differences between a galvanic cell and an electrolytic cell. Refer to type of reaction and energy flow. K/U

2. Describe a consumer product that contains a galvanic cell or battery and is sometimes used as part of an electrolytic cell. A

3. Predict the half-reaction that will take place at the cathode and anode when each salt is melted and electrolyzed. T/I

(a) NiBr2(l)(b) AlF3(l)(c) MnI2(l)

4. Explain why aluminum recycling can be good for both business and the environment. T/I

5. A steel ring can be electroplated with gold in an aqueous solution that contains gold(III) ions. Sketch and label a diagram of a cell that could be used to produce gold-plated steel. Identify the anode and the cathode in the cell. Indicate the direction of electron flow. T/I

6. Draw a diagram of a process that can be used to electroplate copper onto a steel object, such as a coin. Include in your diagram suggested electrodes, the ions in solution, power supply and connectors, the sign of the electrodes, and the direction of electron and ion flow. T/I C

7. During the electrolysis of a tin(II) chloride solution, SnCl2(aq), a silver-coloured metal, is deposited onto the cathode. Bubbles of a gas are observed at the anode. A bleach-like odour is detected. Use a redox table to predict the anode, the cathode, and the net ionic equation. T/I C

8. An electrolytic cell is set up using inert electrodes in an aqueous copper(II) sulfate solution. After an hour of electrolysis, the solution is tested. The concentration of copper(II) ions has decreased, while the concentration of hydrogen ions has increased. The mass of one electrode has also increased. Explain these observations using the half-reaction equation for each electrode. T/I

9. Consider the equations for the anode and cathode half-reactions for the Hall–Héroult process, given in this section. T/I A

(a) Combine these two equations to determine the overall cell reaction equation.

(b) Combine your answer in (a) with the equation for the reaction of aluminum oxide with the cryolite ion, AlF6

32(l), to give the overall equation for the production of aluminum:

2 Al2O3(s) 1 3 C(s) S 4 Al(s) 1 3 CO2(g) 10. Research and explain the steps involved in

removing tarnish from a silver object using aluminum foil and baking soda. Include the equation for the chemical reaction. T/I

11. Electroplating companies use a variety of hazardous substances. Research answers to the following questions, and summarize your findings in a brief report: T/I A C

• What environmental hazards are associated with the electroplating industry?

• What health and safety hazards are associated with the waste that is generated by electroplating companies?

• What steps do electroplating companies take to minimize any negative health or environmental effects of their industry?

Summary

• Electrolysis involves forcing a current through a cell to cause a non-spontaneous redox reaction to occur.

• Electrolytic and galvanic cells are both electrochemical cells. They contain an anode where oxidation occurs and a cathode where reduction occurs.

• An electrolytic cell, unlike a galvanic cell, requires an external energy source and involves a non-spontaneous redox reaction rather than a spontaneous redox reaction.

• Industrial applications of electrolysis include the electrolysis of water and other aqueous solutions, as well as recharging batteries, producing aluminum, electrorefining metals, and electroplating.

Review10.7

WEB LINK

670 Chapter 10 • Electrochemical Cells NEL

7924_Chem_CH10.indd 670 5/4/12 3:51 PM