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Electrochemical Cells for Every Occasion Perhaps near you right now are a laptop, an MP3 player, and a cell phone—just a few of the devices whose energy comes from an electrochemical cell. In this chapter, you’ll look at the two faces of electrochemistry—reactions in cells, such as those in batteries, that do electrical workApago and reactions in cellsPDF that require Enhancer electrical work to occur. Both are indispensable to our modern way of life.

Electrochemistry: Chemical Change and Electrical Work

21.1 Redox Reactions and 21.3 Cell Potential: Output of a Voltaic Cell 21.6 Corrosion: A Case of Environmental Electrochemical Cells Standard Cell Potentials Electrochemistry Review of Oxidation-Reduction Strengths of Oxidizing and Reducing Corrosion of Iron Concepts Agents Protecting Against Corrosion Half-Reaction Method for Balancing 21.4 Free Energy and Electrical Work 21.7 Electrolytic Cells: Using Electrical Redox Reactions Standard Cell Potential and K Energy to Drive Nonspontaneous Electrochemical Cells Effect of Concentration on Ecell Reactions 21.2 Voltaic Cells: Using Spontaneous Changes in Ecell During Cell Operation Construction and Operation Reactions to Generate Electrical Concentration Cells Predicting Electrolysis Products Energy 21.5 Electrochemical Processes Stoichiometry of Electrolysis Construction and Operation in Batteries Cell Notation Primary (Nonrechargeable) Batteries Why Does the Cell Work? Secondary (Rechargeable) Batteries Fuel Cells siL48593_ch21_922-979 8:11:07 04:52am Page 969 fdfd ve403:MHQY042:siL5ch21:

Chapter Review Guide 969

Section Summary An electrolytic cell uses electrical energy to drive a nonspontaneous reaction. • Oxi- dation occurs at the anode and reduction at the cathode, but the direction of elec- tron flow and the charges of the electrodes are opposite those in voltaic cells. • When two products can form at each electrode, the more easily oxidized substance reacts at the anode and the more easily reduced at the cathode. • The reduction or oxi- dation of water takes place at nonstandard conditions. • Overvoltage causes the actual voltage required to be unexpectedly high and can affect the electrode product that forms. • The amount of product that forms depends on the quantity of charge flow- ing through the cell, which is related to the magnitude of the current and the time it flows. • Biological redox systems combine aspects of voltaic, concentration, and electrolytic cells to convert bond energy in food into electrochemical potential and then into the bond energy of ATP. Chapter Perspective The field of electrochemistry is one of the many areas in which the principles of thermodynamics lead to practical benefits. As you’ve seen, electrochemical cells can use a reaction to generate energy or use energy to drive a reaction. Such processes are central not only to our mobile way of life, but also to our biological existence. In Chapter 22, we examine the electrochemical (and other) methods used by industry to convert raw natural resources into some of the materials modern society finds indispensable.

The following sections provide many aids to help you study this chapter. CHAPTER REVIEW GUIDE (Numbers in parentheses refer to pages, unless noted otherwise.) Learning Objectives These areApago concepts andPDF skills youEnhancer should know after studying this chapter. Relevant section and/or sample problem (SP) numbers 12. How the relative reactivity of a metal is determined by its re- appear in parentheses. ducing power and is related to the negative of its Ehalf-cell (Section 21.3) Understand These Concepts 13. How Ecell (the nonstandard cell potential) is related to G 1. The meanings of oxidation and reduction; why an oxidizing (maximum work) and the charge (moles of electrons times the agent is reduced and a reducing agent is oxidized (Section 21.1; Faraday constant) flowing through the cell (Section 21.4) also Section 4.5) 14. The interrelationship of G , Ecell, and K (Section 21.4) 2. How the half-reaction method is used to balance redox reac- 15. How Ecell changes as the cell operates (Q changes) (Section tions in acidic or basic solution (Section 21.1) 21.4) 3. The distinction between voltaic and electrolytic cells in terms 16. Why a voltaic cell can do work until Q K (Section 21.4) of the sign of G (Section 21.1) 17. How a concentration cell does work until the half-cell con- 4. How voltaic cells use a spontaneous reaction to release electri- centrations are equal (Section 21.4) cal energy (Section 21.2) 18. The distinction between primary (nonrechargeable) and sec- 5. The physical makeup of a voltaic cell: arrangement and com- ondary (rechargeable) batteries and the nature of fuel cells position of half-cells, relative charges of electrodes, and purpose (Section 21.5) of a salt bridge (Section 21.2) 19. How corrosion occurs and is prevented; the similarities be- 6. How the difference in reducing strength of the electrodes de- tween a corroding metal and a voltaic cell (Section 21.6) termines the direction of electron flow (Section 21.2) 20. How electrolytic cells use nonspontaneous redox reactions 7. The correspondence between a positive Ecell and a spontaneous driven by an external source of electricity (Section 21.7) cell reaction (Section 21.3) 21. How atomic properties (ionization energy and electronegativ- 8. The usefulness and significance of standard electrode poten- ity) determine the products of the electrolysis of molten salt mix- tials (Ehalf-cell) (Section 21.3) tures (Section 21.7) 9. How Ehalf-cell values are combined to give Ecell (Section 21.3) 22. How the electrolysis of water influences the products of 10. How the standard reference electrode is used to find an un- aqueous electrolysis; the importance of overvoltage (Section 21.7) known Ehalf-cell (Section 21.3) 23. The relationship between the quantity of charge flowing 11. How an emf series (e.g., Table 21.2 or Appendix D) is used to through the cell and the amount of product formed (Section 21.7) write spontaneous redox reactions (Section 21.3) siL48593_ch21_922-979 8:11:07 04:52am Page 970 fdfd ve403:MHQY042:siL5ch21:

970 Chapter 21 Electrochemistry: Chemical Change and Electrical Work

Learning Objectives (continued) Master These Skills 7. Predicting whether a metal can displace hydrogen or another metal from solution (Section 21.3) 1. Balancing redox reactions by the half-reaction method (Sec- 8. Using the interrelationship of , , and to calculate one tion 21.1 and SP 21.1) G Ecell K of the three given the other two (Section 21.4 and SP 21.5) 2. Diagramming and notating a voltaic cell (Section 21.2 and SP 9. Using the Nernst equation to calculate the nonstandard cell po- 21.2) tential (E ) (SP 21.6) 3. Combining E values to obtain E (Section 21.3) cell half-cell cell 10. Calculating E of a concentration cell (SP 21.7) 4. Using E and a known E to find an unknown E cell cell half-cell half-cell 11. Predicting the products of the electrolysis of a mixture of (SP 21.3) molten salts (SP 21.8) 5. Manipulating half-reactions to write a spontaneous redox reac- 12. Predicting the products of the electrolysis of aqueous salt so- tion and calculate its E (SP 21.4) cell lutions (SP 21.9) 6. Ranking the relative strengths of oxidizing and reducing agents 13. Calculating the current (or time) needed to produce a given in a redox reaction (SP 21.4) amount of product by electrolysis (SP 21.10)

Key Terms These important terms appear in boldface in the chapter and are defined again in the Glossary.

electrochemistry (923) Section 21.2 standard cell potential (Ecell) Section 21.5 electrochemical cell (923) half-cell (930) (935) battery (952) Section 21.1 salt bridge (930) standard electrode (half-cell) fuel cell (955) half-reaction method (924) Section 21.3 potential (Ehalf-cell) (935) Section 21.6 voltaic (galvanic) cell (928) cell potential (Ecell) (934) standard reference half-cell corrosion (956) (standard hydrogen electrolytic cell (928) voltage (934) Section 21.7 electrode) (936) electrode (928) electromotive force (emf) electrolysis (961) electrolyte (928) (934) Section 21.4 overvoltage (963) Faraday constant ( ) (943) anode (929) volt (V) (934)Apago PDF EnhancerF ampere (A) (965) cathode (929) coulomb (C) (934) Nernst equation (946) concentration cell (948)

Key Equations and Relationships Numbered and screened concepts are listed for you to refer to or memorize. 21.1 Relating a spontaneous process to the sign of the cell poten- 21.7 Finding the equilibrium constant from the standard cell po- tial (934): tential (944): E 0 for a spontaneous process RT cell E¡ ln K 21.2 Relating electric potential to energy and charge in SI units cell nF (934): 21.8 Substituting known values of R, F, and T into Equation 21.7 Potential energy/charge or 1 V 1 J/C and converting to common logarithms (945): 21.3 Relating standard cell potential to standard electrode poten- 0.0592 V nE¡ E¡ log K or log K cell (at 298.15 K) tials in a voltaic cell (935): cell n 0.0592 V E E E cell cathode (reduction) anode (oxidation) 21.9 Calculating the nonstandard cell potential (Nernst equation) 21.4 Defining the Faraday constant (944): (946): 4 J RT F 9.6510 (3 sf) Ecell E¡cell ln Q V mol e nF 21.5 Relating the free energy change to the cell potential (944): 21.10 Substituting known values of R, F, and T into the Nernst G nFEcell equation and converting to common logarithms (946): 21.6 Finding the standard free energy change from the standard 0.0592 V cell potential (944): Ecell E¡cell log Q (at 298.15 K) n GnFE cell 21.11 Relating current to charge and time (965): Current charge/time or 1 A 1 C/s siL48593_ch21_922-979 8:11:07 04:52am Page 971 fdfd ve403:MHQY042:siL5ch21:

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Highlighted Figures and Tables These figures (F) and tables (T) provide a visual review of key ideas. F21.1 Summary of redox terminology (924) F21.21 The corrosion of iron (957) F21.3 Voltaic and electrolytic cells (928) F21.25 Tin-copper reaction in voltaic and electrolytic cells (959) F21.5 A voltaic cell based on the zinc-copper reaction (931) T21.4 Comparison of voltaic and electrolytic cells (961) F21.10 The interrelationship of G , Ecell, and K (944) F21.28 Summary of the stoichiometry of electrolysis (965) F21.11 Ecell and log Q for the zinc-copper cell (948)

Brief Solutions to FOLLOW-UP PROBLEMS Compare your solutions to these calculation steps and answers. 2 21.1 6KMnO4(aq) 6KOH(aq) KI(aq) ±£ 21.6 Fe(s)Fe±£ (aq) 2e E0.44 V 2 6K2MnO4(aq) KIO3(aq) 3H2O(l) Cu (aq) 2e ±£ Cu(s) E0.34 V 2 2 2 21.2 Sn(s) ±£ Sn (aq) 2e Fe(s) Cu (aq) ±£ Fe (aq) Cu(s) Ecell 0.78 V [anode; oxidation] So Ecell 0.78 V 0.25 V 1.03 V 6e 14H (aq) Cr O 2 (aq) ±£ 2Cr3 (aq) 7H O(l) 2 2 7 2 0.0592 V [Fe ] [cathode; reduction] 1.03 V 0.78 V log 2 [Cu2 ] 2 2 3Sn(s) Cr2O7 (aq) 14H (aq) ±£ [Fe ] 9 2 3 3.610 3Sn (aq) 2Cr (aq) 7H2O(l) [overall] [Cu2 ] Cell notation: 2 9 9 2 2 3 [Fe ] 3.6 10 0.30 M 1.1 10 M Sn(s) Sn (aq) H (aq), Cr2O7 (aq), Cr (aq) graphite 21.7 Au3 (aq; 2.510 2 M) [B] ±£ 3 4 e– Voltmeter e– Au (aq; 7.010 M) [A] 4 0.0592 V 7.0 10 Anode Cathode Ecell 0 V log 0.0306 V a 3 2.510 2 b Sn (–) NO – K+ (+) C 3 The electrode in A is negative, so it is the anode. Apago PDF Enhancer 21.8 Oxidizing agents: K and Al3 . Reducing agents: F and Br . Al is above and to the right of K in the periodic table, so it has a higher IE: 3 + Al (l) 3e ±£ Al(s) [cathode; reduction] Sn2+ Cr3 , H+, Cr O 2– 2 7 Br is below F in Group 7A(17), so it has a lower EN: 21.3 Br2(aq) 2e ±£ 2Br (aq) Ebromine 1.07 V 2Br (l) ±£ Br2(g) 2e [anode; oxidation] 3 [cathode] 2Al (l) 6Br (l) ±£ 2Al(s) 3Br2(g) [overall] 3 2 2V (aq) 2H2O(l) ±£ 2VO (aq) 4H (aq) 2e 21.9 The reduction with the more positive electrode potential is 3 Evanadium ? [anode] Au (aq) 3e ±£ Au(s); E1.50 V Evanadium Ebromine Ecell 1.07 V 1.39 V 0.32 V [cathode; reduction] 21.4 Fe2 (aq) 2e ±£ Fe(s) E0.44 V Because of overvoltage, O2 will not form at the anode, so Br2 will 2 ±£ 3 form: 2[Fe (aq)Fe(aq) e ] E 0.77 V 2 3 2Br (aq) ±£ Br (l) 2e ; E1.07 V 3Fe (aq) ±£ 2Fe (aq) Fe(s) 2 [cathode; oxidation] E 0.44 V 0.77 V 1.21 V cell 21.10 Cu2 (aq) 2e ±£ Cu(s); therefore, The reaction is nonspontaneous. The spontaneous reaction is 2 mol e /1 mol Cu 2 mol e /63.55 g Cu 3 2 2Fe (aq) Fe(s) ±£ 3Fe (aq) Ecell 1.21 V 2 mol e 2 3 Time (min) 1.50 g Cu Fe Fe Fe 63.55 g Cu 21.5 Cd(s) Cu2 (aq) ±£ Cd2 (aq) Cu(s) 4 9.65 10 C 1 s 1 min G RT ln K 8.314 J/mol rxn K 298 K ln K 1 mol e 4.75 C 60 s 25 143 kJ/mol rxn; K 1.2 10 16.0 min 0.0592 V E¡ log (1.21025) 0.742 V cell 2