This is “Electrochemistry”, chapter 19 from the book Principles of General Chemistry (index.html) (v. 1.0M). This book is licensed under a Creative Commons by-nc-sa 3.0 (http://creativecommons.org/licenses/by-nc-sa/ 3.0/) license. See the license for more details, but that basically means you can share this book as long as you credit the author (but see below), don't make money from it, and do make it available to everyone else under the same terms. This content was accessible as of December 29, 2012, and it was downloaded then by Andy Schmitz (http://lardbucket.org) in an effort to preserve the availability of this book. Normally, the author and publisher would be credited here. However, the publisher has asked for the customary Creative Commons attribution to the original publisher, authors, title, and book URI to be removed. Additionally, per the publisher's request, their name has been removed in some passages. More information is available on this project's attribution page (http://2012books.lardbucket.org/attribution.html?utm_source=header). For more information on the source of this book, or why it is available for free, please see the project's home page (http://2012books.lardbucket.org/). You can browse or download additional books there. i Chapter 19 Electrochemistry In oxidation–reduction (redox) reactions, electrons are transferred from one species (the reductant) to another (the oxidant). This transfer of electrons provides a means for converting chemical energy to electrical energy or vice versa. The study of the relationship between electricity and chemical reactions is called electrochemistry1, an area of chemistry we introduced in Chapter 4 "Reactions in Aqueous Solution" and Chapter 5 "Energy Changes in Chemical Reactions". In this chapter, we describe electrochemical reactions in more depth and explore some of their applications. In the first three sections, we review redox reactions; describe how they can be used to generate an electrical potential, or voltage; and discuss factors that affect the magnitude of the potential. We then explore the relationships among the electrical potential, the change in free energy, and the equilibrium constant for a redox reaction, which are all measures of the thermodynamic driving force for a reaction. Finally, we examine two kinds of applications of electrochemical principles: (1) those in which a spontaneous reaction is used to provide electricity and (2) those in which electrical energy is used to drive a thermodynamically nonspontaneous reaction. By the end of this chapter, you will understand why different kinds of batteries are used in cars, flashlights, cameras, and portable computers; how rechargeable batteries operate; and why corrosion occurs and how to slow—if not prevent—it. You will also discover how metal objects can be plated with silver or chromium for protection; how silver polish removes tarnish; and how to calculate the amount of electricity needed to produce aluminum, chlorine, copper, and sodium on an industrial scale. 1. The study of the relationship between electricity and chemical reactions. 2272 Chapter 19 Electrochemistry A view from the top of the Statue of Liberty, showing the green patina coating the statue. The patina is formed by corrosion of the copper skin of the statue, which forms a thin layer of an insoluble compound that contains copper(II), sulfate, and hydroxide ions. 2273 Chapter 19 Electrochemistry 19.1 Describing Electrochemical Cells LEARNING OBJECTIVE 1. To distinguish between galvanic and electrolytic cells. In any electrochemical process, electrons flow from one chemical substance to another, driven by an oxidation–reduction (redox) reaction. As we described in Chapter 3 "Chemical Reactions", a redox reaction occurs when electrons are transferred from a substance that is oxidized to one that is being reduced. The reductant2 is the substance that loses electrons and is oxidized in the process; the oxidant3 is the species that gains electrons and is reduced in the process. The associated potential energy is determined by the potential difference between the valence electrons in atoms of different elements. (For more information on valence electrons, see Chapter 7 "The Periodic Table and Periodic Trends", Section 7.3 "Energetics of Ion Formation".) Because it is impossible to have a reduction without an oxidation and vice versa, a redox reaction can be described as two half-reactions4, one representing the oxidation process and one the reduction process. For the reaction of zinc with bromine, the overall chemical reaction is as follows: Equation 19.1 2+ − Zn(s) + Br2(aq) → Zn (aq) + 2Br (aq) The half-reactions are as follows: Equation 19.2 2. A substance that is capable of donating electrons and in the − − process is oxidized. reduction half-reaction: Br2(aq) + 2e → 2Br (aq) 3. A substance that is capable of accepting electrons and in the Equation 19.3 process is reduced. 2+ − 4. Reactions that represent either oxidation half-reaction: Zn(s) → Zn (aq) + 2e the oxidation half or the reduction half of an Each half-reaction is written to show what is actually occurring in the system; Zn is oxidation–reduction (redox) reaction. the reductant in this reaction (it loses electrons), and Br2 is the oxidant (it gains 2274 Chapter 19 Electrochemistry electrons). Adding the two half-reactions gives the overall chemical reaction (Equation 19.1). A redox reaction is balanced when the number of electrons lost by the reductant equals the number of electrons gained by the oxidant. Like any balanced chemical equation, the overall process is electrically neutral; that is, the net charge is the same on both sides of the equation. Note the Pattern In any redox reaction, the number of electrons lost by the reductant equals the number of electrons gained by the oxidant. 5. An apparatus that generates electricity from a spontaneous oxidation–reduction (redox) In most of our discussions of chemical reactions, we have assumed that the reaction or, conversely, uses reactants are in intimate physical contact with one another. Acid–base reactions, electricity to drive a for example, are usually carried out with the acid and the base dispersed in a single nonspontaneous redox phase, such as a liquid solution. With redox reactions, however, it is possible to reaction. physically separate the oxidation and reduction half-reactions in space, as long as 6. An electrochemical cell that there is a complete circuit, including an external electrical connection, such as a uses the energy released wire, between the two half-reactions. As the reaction progresses, the electrons flow during a spontaneous oxidation–reduction (redox) from the reductant to the oxidant over this electrical connection, producing an reaction (ΔG < 0) to electric current that can be used to do work. An apparatus that is used to generate generate electricity. electricity from a spontaneous redox reaction or, conversely, that uses electricity to 5 7. An electrochemical cell that drive a nonspontaneous redox reaction is called an electrochemical cell . consumes electrical energy from an external source to drive a nonspontaneous There are two types of electrochemical cells: galvanic cells and electrolytic cells. A (ΔG > 0) galvanic (voltaic) cell6Galvanic cells are named for the Italian physicist and oxidation–reduction (redox) physician Luigi Galvani (1737–1798), who observed that dissected frog leg muscles reaction. twitched when a small electric shock was applied, demonstrating the electrical 8. A solid metal connected by an nature of nerve impulses. uses the energy released during a spontaneous redox electrolyte and an external reaction (ΔG < 0) to generate electricity. This type of electrochemical cell is often circuit that provides an electrical connection between called a voltaic cell after its inventor, the Italian physicist Alessandro Volta 7 systems in an electrochemical (1745–1827). In contrast, an electrolytic cell consumes electrical energy from an cell (galvanic or electrolytic). external source, using it to cause a nonspontaneous redox reaction to occur (ΔG > 8 9. One of two electrodes in an 0). Both types contain two electrodes , which are solid metals connected to an electrochemical cell, it is the external circuit that provides an electrical connection between the two parts of the site of the oxidation half- system (Figure 19.1 "Electrochemical Cells"). The oxidation half-reaction occurs at reaction. one electrode (the anode9), and the reduction half-reaction occurs at the other (the 10 10. One of two electrodes in an cathode ). When the circuit is closed, electrons flow from the anode to the electrochemical cell, it is the cathode. The electrodes are also connected by an electrolyte, an ionic substance or site of the reduction half- solution that allows ions to transfer between the electrode compartments, thereby reaction. 19.1 Describing Electrochemical Cells 2275 Chapter 19 Electrochemistry maintaining the system’s electrical neutrality. In this section, we focus on reactions that occur in galvanic cells. We discuss electrolytic cells in Section 19.7 "Electrolysis". Figure 19.1 Electrochemical Cells A galvanic cell (left) transforms the energy released by a spontaneous redox reaction into electrical energy that can be used to perform work. The oxidative and reductive half-reactions usually occur in separate compartments that are connected by an external electrical circuit; in addition, a second connection that allows ions to flow between the compartments (shown here as