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PART IV Electrochemical Methods electrochemical PART IV Methods Chapter 18 Introduction to Electrochemistry Chapter 19 Applications of Standard Electrode Potentials Chapter 20 Applications of Oxidation/Reduction Titrations Chapter 21 Potentiometry Chapter 22 Bulk Electrolysis: Electrogravimetry and Coulometry Chapter 23 Voltammetry Copyright 2013 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. Introduction to chapteR 18 electrochemistry From the earliest days of experimental science, workers such as Galvani, Volta, and cavendish realized that electricity interacts in interesting and important ways with animal tissues. ­Electrical charge causes muscles to contract, for example. Perhaps more surprising is that a few animals such as the torpedo (shown in the photo) produce charge by physiological means. More than 50 billion nerve terminals in the torpedo’s flat “wings” on its left and right sides rapidly emit ace- tylcholine on the bottom side of membranes housed in the wings. The acetylcholine causes sodium ions to surge through the membranes, producing a rapid separation of charge and a correspond- ing potential difference, or voltage, across the membrane.1 The potential difference then gener- ates an electric current of several amperes in the surrounding seawater that may be used to stun or kill prey, detect and ward off enemies, or navigate. Natural devices for separating charge and creating electrical potential difference are relatively rare, but humans have learned to separate charge mechanically, metallurgically, and chemically to create cells, batteries, and other useful charge storage devices. e now turn our attention to several analytical methods that are based on W oxidation/reduction reactions. These methods, which are described in chapters 18 through 23, include oxidation/ reduction titrimetry, potentiometry, coulometry, electrogra- vimetry, and voltammetry. In this chapter, we present the fundamentals of electrochemistry © Norbert Wu/Minden Pictures/Corbis that are necessary for understanding the principles of these procedures. CharaCterIzIng OxIdatIOn/reduCtIOn 18a reaCtIOns Oxidation/reduction reactions are In an oxidation/reduction reaction electrons are transferred from one reactant to sometimes called redox reactions. another. An example is the oxidation of iron(II) ions by cerium(IV) ions. The reac- tion is described by the equation 41 21 31 31 Ce 1 Fe 8 Ce 1 Fe (18-1) 21 41 31 31 A reducing agent is an electron In this reaction, an electron is transferred from Fe to Ce to form Ce and Fe 41 donor. An oxidizing agent is an ions. A substance that has a strong affinity for electrons, such as Ce , is called an electron acceptor. oxidizing agent, or an oxidant. A reducing agent, or reductant, is a species, such 1Y. Dunant and M. Israel, Sci. Am. 1985, 252, 58, DOI: 10.1038/scientificamerican0485-58. Copyright 2013 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 18A Characterizing Oxidation/reduction Reactions 443 as Fe21, that donates electrons to another species. To describe the chemical behavior 21 41 41 represented by Equation 18-1, we say that Fe is oxidized by Ce ; similarly, Ce is reduced by Fe21. We can split any oxidation/reduction equation into two half-reactions that show which species gains electrons and which loses them. For example, Equation 18-1 is the sum of the two half-reactions Ce41 1 e2 8 Ce31 (reduction of Ce41) Fe21 8 Fe31 1 e2 (oxidation of Fe21) It is important to understand ❮that while we can write an equation for a half-reaction in The rules for balancing half-reactions (see Feature 18-1) are the same as those for which electrons are consumed other reaction types, that is, the number of atoms of each element as well as the net or generated, we cannot observe charge on each side of the equation must be the same. Thus, for the oxidation of an isolated half-reaction 21 2 experimentally because there Fe by MnO4 , the half-reactions are must always be a second half- MnO 21 5e2 1 8H1 8 Mn21 1 4H O reaction that serves as a source 4 2 of electrons or a recipient of 5Fe21 8 5Fe31 1 5e2 electrons. In other words, an individual half-reaction is a In the first half-reaction, the net charge on the left side is (21 25 1 8) 5 12, theoretical concept. which is the same as the charge on the right. Note also that we have multiplied the second half-reaction by 5 so that the number of electrons lost by Fe21 equals the 2 number gained by MnO4 . We can then write a balanced net ionic equation for the overall reaction by adding the two half-reactions 2 21 1 21 31 MnO4 1 5Fe 1 8H 8 Mn 1 5Fe 1 4H2O 18A-1 Comparing Redox Reactions to Acid/Base Reactions Recall that in the Brønsted/ ❮Lowry concept an acid/base Oxidation/reduction reactions can be viewed in a way that is analogous to the reaction is described by the Brønsted-Lowry concept of acid/base reactions (see Section 9A-2). In both, one or equation more charged particles are transferred from a donor to an acceptor—the particles acid1 1 base2 8 base1 1 acid2 Copyright 1993 by permission of Johnny Hart and Creator's Syndicate, Inc. Copyright 2013 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 444 ChApteR 18 Introduction to electrochemistry Feature 18-1 Balancing redox equations Knowing how to balance oxidation/reduction reactions is essential to understanding all the concepts covered in this chapter. Although you probably remember this tech- nique from your general chemistry course, we present a quick review to remind you of how the process works. For practice, we will complete and balance the following 1 2 equation after adding H , OH , or H2O as needed. 2 2 21 2 MnO4 1 NO2 8 Mn 1 NO3 2 First, we write and balance the two half-reactions. For MnO4 , we write 2 21 MnO4 8 Mn To account for the 4 oxygen atoms on the left-hand side of the equation, we add 4H2O on the right-hand side. Then, to balance the hydrogen atoms, we must provide 8H1 on the left: 2 1 21 MnO4 1 8H 8 Mn 1 4H2O To balance the charge, we need to add 5 electrons to the left side of the equation. Thus, 2 1 2 21 MnO4 1 8H 1 5e 8 Mn 1 4H2O For the other half-reaction, 2 2 NO2 8 NO3 we add one H2O to the left side of the equation to supply the needed oxygen and 2H1 on the right to balance hydrogen: 2 2 1 NO2 1 H2O 8 NO3 1 2H Then, we add two electrons to the right-hand side to balance the charge: 2 2 1 2 NO2 1 H2O 8 NO3 1 2H 1 2e Before combining the two equations, we must multiply the first by 2 and the second by 5 so that the number of electrons lost will be equal to the number of electrons gained. We then add the two half reactions to obtain 2 1 2 2 2MnO4 1 16H 1 10e 1 5NO2 1 5H2O 8 21 2 1 2 2Mn 1 8H2O 1 5NO3 1 10H 1 10e This equation rearranges to the balanced equation 2 1 2 21 2 2MnO4 1 6H 1 5NO2 8 2Mn 1 5NO3 1 3H2O Copyright 2013 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 18A Characterizing Oxidation/reduction Reactions 445 being electrons in oxidation/reduction and protons in neutralization. When an acid donates a proton, it becomes a conjugate base that is capable of accepting a pro- ton. By analogy, when a reducing agent donates an electron, it becomes an oxidizing agent that can then accept an electron. This product could be called a conjugate oxidant, but that terminology is seldom, if ever, used. With this idea in mind, we can write a generalized equation for a redox reaction as Ared 1 Box 8 Aox 1 Bred (18-2) In this equation, Box, the oxidized form of species B, accepts electrons from Ared to form the new reductant, Bred. At the same time, reductant Ared, having given up elec- trons, becomes an oxidizing agent, Aox. If we know from chemical evidence that the equilibrium in Equation 18-2 lies to the right, we can state that Box is a better elec- tron acceptor (stronger oxidant) than Aox.
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