11.5 Section Preview/ Specific Expectations In this section, you will ■ explain corrosion as an electrochemical process, and describe some techniques used to prevent corrosion ■ evaluate the environmental and social impact of some common cells, including the hydrogen fuel cell in electric cars ■ explain how electrolytic processes are used in the industrial production of chlorine, and how this element is used in the purification of water ■ research and assess some environmental, health, and safety issues connected to electrochemistry ■ communicate your understanding of the following terms: corrosion, galvanizing, sacrificial anode, cathodic protection, fuel cell, chlor-alkali process Issues Involving Electrochemistry You have probably seen many examples of rusty objects, such as the one in Figure 11.26. However, you may not realize that rusting costs many billions of dollars per year in prevention, maintenance, and replacement costs. You will now learn more about rusting and about other issues involving electrochemistry. Figure 11.26 Rust is a common sight on iron objects. Corrosion Rusting is an example of corrosion, which is a spontaneous redox reaction of materials with substances in their environment. Figure 11.27 shows an example of the hazards that result from corrosion. Many metals are fairly easily oxidized. The atmosphere contains a powerful oxidizing agent: oxygen. Because metals are constantly in contact with oxygen, they are vulnerable to corrosion. In fact, the term “corrosion” is sometimes defined as the oxidation of metals exposed to the environment. In North America, about 20% to 25% of iron and steel production is used to replace objects that have been damaged or destroyed by corrosion. However, not all corrosion is harmful. For example, the green layer formed by the corrosion of a copper roof is considered attractive by many people. Figure 11.27 On April 28, 1988, this Aloha Airlines aircraft was flying at an altitude of 7300 m when a large part of the upper fuselage ripped off. This accident was caused by undetected corrosion damage. The pilot showed tremendous skill in landing the plane. 546 MHR • Unit 5 Electrochemistry Rust is a hydrated iron(III) oxide, Fe2O3 • xH2O. The electrochemical formation of rust occurs in small galvanic cells on the surface of a piece of iron, as shown in Figure 11.28. In each small cell, iron acts as the anode. The cathode is inert, and may be an impurity that exists in the iron or is deposited onto it. For example, the cathode could be a piece of soot that has been deposited onto the iron surface from the air. Water, in the form of rain, is needed for rusting to occur. Carbon dioxide in the air dissolves in water to form carbonic acid, H2CO3(aq) . This weak acid partially dissociates into ions. Thus, the carbonic acid is an electrolyte for the corrosion process. Other electrolytes, such as road salt, may also be involved. The circuit is completed by the iron itself, which conducts electrons from the anode to the cathode. air O2 water rust Fe 2+ iron e− anode cathode Figure 11.28 The rusting of iron involves the reaction of iron, oxygen, and water in a naturally occurring galvanic cell on the exposed surface of the metal. There may be many of these small cells on the surface of the same piece of iron. The rusting process is complex, and the equations may be written in various ways. A simplified description of the half-reactions and the overall cell reaction is given here. Oxidation half-reaction (occurs at the anode): Fe(s) → Fe2+(aq) + 2e− Reduction half-reaction (occurs at the cathode): O2(g) + 2H2O() + 4e− → 4OH−(aq) Multiply the oxidation half-reaction by two and add the half-reactions to obtain the overall cell reaction. 2Fe(s) + O2(g) + 2H2O() → 2Fe2+(aq) + 4OH−(aq) There is no barrier in the cell, so nothing stops the dissolved Fe2+ and OH− ions from mixing. The iron(II) ions produced at the anode and the hydroxide ions produced at the cathode react to form a precipitate of iron(II) hydroxide, Fe(OH)2 . Therefore, the overall cell reaction could be written as follows. 2Fe(s) + O2(g) + 2H2O() → 2Fe(OH)2(s) The iron(II) hydroxide undergoes further oxidation by reaction with the oxygen in the air to form iron(III) hydroxide. 4Fe(OH)2(s) + O2(g) + 2H2O() → 4Fe(OH)3(s) Chapter 11 Cells and Batteries • MHR 547 Iron(III) hydroxide readily breaks down to form hydrated iron(III) oxide, Fe2O3 • xH2O, more commonly known as rust. As noted earlier, the “x” signifies a variable number of water molecules per formula unit. 2Fe(OH)3(s) → Fe2O3 • 3H2O(s) Both Fe(OH)3 and Fe2O3 • xH2O are reddish-brown, or “rust coloured.” A rust deposit may contain a mixture of these compounds. Fortunately, not all metals corrode to the same extent as iron. Many metals do corrode in air to form a surface coating of metal oxide. However, in many cases, the oxide layer adheres, or sticks firmly, to the metal surface. This layer protects the metal from further corrosion. For example, aluminum, chromium, and magnesium are readily oxidized in air to form their oxides, Al2O3 , Cr2O3 , and MgO. Unless the oxide layer is broken by a cut or a scratch, the layer prevents further corrosion. In contrast, rust easily flakes off from the surface of an iron object and provides little protection against further corrosion. Corrosion Prevention Corrosion, and especially the corrosion of iron, can be very destructive. For this reason, a great deal of effort goes into corrosion prevention. The simplest method of preventing corrosion is to paint an iron object. The protective coating of paint prevents air and water from reaching the metal surface. Other effective protective layers include grease, oil, plastic, or a metal that is more resistant to corrosion than iron. For example, a layer of chromium protects bumpers and metal trim on cars. An enamel coating is often used to protect metal plates, pots, and pans. Enamel is a shiny, hard, and very unreactive type of glass that can be melted onto a metal surface. A protective layer is effective as long as it completely covers the iron object. If a hole or scratch breaks the layer, the metal underneath can corrode. It is also possible to protect iron against corrosion by forming an alloy with a different metal. Stainless steel is an alloy of iron that contains at least 10% chromium, by mass, in addition to small quantities of carbon and occasionally metals such as nickel. Stainless steel is much more resistant to corrosion than pure iron. Therefore, stainless steel is often used for cutlery, taps, and various other applications where rust-resistance is important. However, chromium is much more expensive than iron. As a result, stainless steel is too expensive for use in large-scale applications, such as building bridges. Galvanizing is a process in which iron is covered with a protective layer of zinc. Galvanized iron is often used to make metal buckets and chain-link fences. Galvanizing protects iron in two ways. First, the zinc acts as a protective layer. If this layer is broken, the iron is exposed to air and water. When this happens, however, the iron is still protected. Zinc is more easily oxidized than iron. Therefore, zinc, not iron, becomes the anode in the galvanic cell. The zinc metal is oxidized to zinc ions. In this situation, zinc is known as a sacrificial anode, because it is destroyed (sacrificed) to protect the iron. Iron acts as the cathode when zinc is present. Thus, iron does not undergo oxidation until all the zinc has reacted. 548 MHR • Unit 5 Electrochemistry Cathodic protection is another method of preventing rusting, as shown in Figure 11.29. As in galvanizing, a more reactive metal is attached to the iron object. This reactive metal acts as a sacrificial anode, and the iron becomes the cathode of a galvanic cell. Unlike galvanizing, the metal used in cathodic protection does not completely cover the iron. Because the sacrificial anode is slowly destroyed by oxidation, it must be replaced periodically. If iron is covered with a protective layer of a metal that is less reactive than iron, there can be unfortunate results. A “tin” can is actually a steel can coated with a thin layer of tin. While the tin layer remains intact, it provides effective protection against rusting. If the tin layer is broken or scratched, however, the iron in the steel corrodes faster in contact with the tin than the iron would on its own. Since tin is less reactive than iron, tin acts as a cathode in each galvanic cell on the surface of the can. Therefore, the tin provides a large area of available cathodes for the small galvanic cells involved in the rusting process. Iron acts as the anode of each cell, which is its normal role when rusting. Sometimes, the rusting of iron is promoted accidentally. For example, by connecting an iron pipe to a copper pipe in a plumbing system, an inexperienced plumber could accidentally speed up the corrosion of the iron pipe. Copper is less reactive than iron. Therefore, copper acts as the cathode and iron as the anode in numerous small galvanic cells at the intersection of the two pipes. Build on your understanding of corrosion by completing the following practice problems. Figure 11.29 Magnesium, zinc, or aluminum blocks are attached to ships’ hulls, oil and gas pipelines, underground iron pipes, and gasoline storage tanks. These reactive metals provide cathodic protection by acting as a sacrificial anode. Practice Problems 25. (a) Use the two half-reactions for the rusting process, and a table of standard reduction potentials. Determine the standard cell potential for this reaction. (b) Do you think that your calculated value is the actual cell potential for each of the small galvanic cells on the surface of a rusting iron object? Explain. 26. Explain why aluminum provides cathodic protection to an iron object. 27. In the year 2000, Transport Canada reported that thousands of cars sold in the Atlantic Provinces between 1989 and 1999 had corroded engine cradle mounts. Failure of these mounts can cause the steering shaft to separate from the car. The manufacturer recalled the cars so that repairs could be made, where necessary. The same cars were sold across the country. Why do you think that the corrosion problems showed up in the Atlantic Provinces? 28. (a) Use a table of standard reduction potentials to determine whether elemental oxygen, O2(g), is a better oxidizing agent under acidic conditions or basic conditions. (b) From your answer to part (a), do you think that acid rain promotes or helps prevent the rusting of iron? Chapter 11 Cells and Batteries • MHR 549 Automobile Engines The internal combustion engine found in most automobiles uses gasoline as a fuel. Unfortunately, this type of engine produces pollutants, such as carbon dioxide (CO2 ), nitrogen oxides (NOx ), and volatile organic compounds (VOCs). These pollutants contribute to health and environmental problems, such as smog and the greenhouse effect. In addition, the internal combustion engine is very inefficient. It converts only about 25% of the chemical energy of the fuel into the kinetic energy of the car. Electric cars, such as those shown in Figure 11.30, may provide a more efficient and less harmful alternative. Figure 11.30 Many electric cars are currently under development. In fact, the idea of electric cars is not new. In the early days of the automobile, electric cars were more common than gasoline-powered cars. The production of electric cars peaked in 1912, but then completely stopped in the 1930s. Manufacturers and researchers have attempted to power electric cars with rechargeable batteries, such as modified lead-acid and nickel-cadmium batteries. However, rechargeable batteries run down fairly quickly. The distance driven before recharging a battery may be 250 km or less. The battery must then be recharged from an external electrical source. Recharging the lead-acid battery of an electric car takes several hours. Cars based on a version of the nickel-cadmium battery can be recharged in only fifteen minutes. However, recharging the batteries of an electric car is still inconvenient. A new type of power supply for electric cars eliminates the need for recharging. A fuel cell is a battery that produces electricity while reactants are supplied continuously from an external source. Because reactants continuously flow into the cell, a fuel cell is also known as a flow battery. Unlike the fuel supply of a more conventional battery, the fuel supply in a fuel cell is unlimited. As in the combustion of gasoline in a conventional engine, the overall reaction in a fuel cell is the oxidation of a fuel by oxygen. The space shuttle uses a fuel cell as a source of energy. This cell depends on the oxidation of hydrogen by oxygen to form water. The fuel cell operates under basic conditions, so it is sometimes referred to as an alkaline fuel cell. Figure 11.31, on the next page, shows the design of the cell. The half-reactions and the overall reaction are as follows. Reduction (occurs at the cathode): O2(g) + 2H2O() + 4e− → 4OH−(aq) Oxidation (occurs at the anode): H2(g) + 2OH−(aq) → 2H2O() + 2e− Multiply the oxidation half-reaction by 2, add the two half-reactions, and simplify to obtain the overall cell reaction. 2H2(g) + O2(g) → 2H2O() Notice that the overall equation is the same as the equation for the combustion of hydrogen. The combustion of hydrogen is an exothermic reaction. In the fuel cell, this reaction produces energy in the form of electricity, rather than heat. 550 MHR • Unit 5 Electrochemistry anode e− e− cathode H2 O2 inert carbon electrode containing Ni inert carbon electrode containing Ni and NiO hot KOH solution Figure 11.31 In an alkaline fuel cell, the half-reactions do not include solid conductors of electrons. Therefore, the cell has two inert electrodes. The hydrogen fuel cell produces water vapour, which does not contribute to smog formation or to the greenhouse effect. This product makes the hydrogen fuel cell an attractive energy source for cars. Also, the hydrogen fuel cell is much more efficient than the internal combustion engine. A hydrogen fuel cell converts about 80% of the chemical energy of the fuel into the kinetic energy of the car. There is a possible problem with the hydrogen fuel cell. The cell requires hydrogen fuel. Unfortunately, uncombined hydrogen is not found naturally on Earth. Most hydrogen is produced from hydrocarbon fuels, such as petroleum or methane. These manufacturing processes may contribute significantly to pollution problems. However, hydrogen can also be produced by the electrolysis of water. If a source such as solar energy or hydroelectricity is used to power the electrolysis, the overall quantity of pollution is low. The following practice problems will allow you to test your understanding of fuel cells. Practice Problems 29. Calculate E˚cell for a hydrogen fuel cell. 30. In one type of fuel cell, methane is oxidized by oxygen to form carbon dioxide and water. (a) Write the equation for the overall cell reaction. (b) Write the two half-reactions, assuming acidic conditions. 31. Reactions that occur in fuel cells can be thought of as being “flameless combustion reactions.” Explain why. 32. If a hydrogen fuel cell produces an electric current of 0.600 A for 120 min, what mass of hydrogen is consumed by the cell? Chapter 11 Cells and Batteries • MHR 551
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