Electrochemistry 1 1

Transcription

1 Electrochemistry 1 1

2 Half-Reactions 1. Balancing Oxidation Reduction Reactions in Acidic and Basic Solutions Voltaic Cells 2. Construction of Voltaic Cells 3. Notation for Voltaic Cells 4. Cell Potential 5. Standard Cell Potentials and Standard Electrode Potentials 2 2

3 6. Equilibrium Constants from Cell Potentials 7. Dependence of Cell Potentials on Concentration 8. Some Commercial Voltaic Cells Electrolytic Cells 9. Electrolysis of Molten Salts 10. Aqueous Electrolysis 11. Stoichiometry of Electrolysis 3 3

4 Learning Objectives Electrochemistry Balancing Oxidation Reduction Reactions in Acidic and Basic Solutions a. Learn the steps for balancing oxidation reduction reactions using the half-reaction method. b. Balance equations by the half-reaction method (acidic solutions). c. Learn the additional steps for balancing oxidation reduction reactions in basic solution using the half-reaction method. d. Balance equations using the half-reaction method (basic solution). 4 4

5 2. Construction of Voltaic Cells Define electrochemical cell, voltaic (galvanic cell), electrolytic cell, and half-cell. Describe the function of the salt bridge in a voltaic cell. State the reaction that occurs at the anode and the cathode in an electrochemical cell. Define cell reaction. Sketch and label a voltaic cell. 5 5

6 3. Notation for Voltaic Cells Write the cell reaction from the cell notation. 4. Cell Potential Define cell potential and volt. Calculate the quantity of work from a given amount of cell reactant. 6 6

7 5. Standard Cell Potentials and Standard Electrode Potentials Explain how the electrode potential of a cell is an intensive property. Define standard cell potential and standard electrode potential. Interpret the table of standard reduction potentials. Determine the relative strengths of oxidizing and reducing agents. Determine the direction of spontaneity from electrode potentials. Calculate cell potentials from standard potentials. 7 7

8 6. Equilibrium Constants from Cell Potentials Calculate the free-energy change from electrode potentials. Calculate the cell potential from free-energy change. Calculate the equilibrium constant from cell potential. 7. Dependence of Cell Potential on Concentration Calculate the cell potential for nonstandard conditions. Describe how ph can be determined using a glass electrode. 8 8

9 8. Some Commercial Voltaic Cells Describe the construction and reactions of a zinc carbon dry cell, a lithium iodine battery, a lead storage cell, and a nickel cadmium cell. Explain the operation of a proton-exchange membrane fuel cell. Explain the electrochemical process of the rusting of iron. Define cathodic protection. 9. Electrolysis of Molten Salts Define electrolysis. 9 9

10 10. Aqueous Electrolysis Learn the half-reactions for water undergoing oxidation and reduction. Predict the half-reactions in an aqueous electrolysis. 11. Stoichiometry of Electrolysis Calculate the amount of charge from the amount of product in an electrolysis. Calculate the amount of product from the amount of charge in an electrolysis

11 Our first step in studying electrochemical cell is to balance its oxidation reduction reaction. We will use the half-reaction method from chapter 5 and extend it to acidic or basic solutions. In this chapter we will focus on electron transfer rather than proton transfer so the hydronium ion, H 3 O + (aq), will be represented by its simpler notation, H + (aq). Only the notation, not the chemistry, is different

12 Oxidation-Reduction Reactions In Chapter 5 we introduced the half-reaction method for balancing simple oxidation-reduction reactions. Oxidation-reduction reactions always involve a transfer of electrons from one species to another. Recall that the species losing electrons is oxidized, while the species gaining electrons is reduced

13 Oxidation-Reduction Reactions Describing Oxidation-Reduction Reactions An oxidizing agent is a species that oxidizes another species; it is itself reduced. A reducing agent is a species that reduces another species; it is itself oxidized. Loss of 2 e -1 oxidation reducing agent Fe + ( s) Cu ( aq) Fe ( aq) + Cu( s) oxidizing agent Gain of 2 e -1 reduction 13 13

14 What reaction is going on and what is the equation? A voltaic cell employs a spontaneous oxidation reduction reaction as a source of energy. It separates the reaction into two halfreactions, physically separating one half-reaction from the other

15 We will be studying more complex situations, so our initial analysis is key. First, we need to identify what is being oxidized and what is being reduced. Then, we determine if the reaction is in acidic or basic conditions

16 Balancing Oxidation-Reduction Equations in Acid Solution 1. Assign oxidation numbers 2. Determine half reactions 3. Complete and balance each half-reaction a. Balance all atoms except O and H b. Balance O atoms by adding H 2 O s to one side c. Balance H atoms by adding H + to one side d. Balance electric charges by adding e - to the more positive side 4. Combine each half reaction to obtain the final balanced oxidation reduction equation (Multiply by appropriate factors and cancel species appearing on both sides of equation 16 16

17 Balancing Oxidation-Reduction Equations in Basic Solution 1. Assign oxidation numbers 2. Determine half reactions 3. Complete and balance each half-reaction a. Balance all atoms except O and H b. Balance O atoms by adding H 2 O s to one side c. Balance H atoms by adding H + to one side d. Balance electric charges by adding e - to the more positive side 4. Note the number of H + ions in the equation. Add this number of OH - ions to both sides of the equation. 5. Simplify the equation by noting that H+ and OH- react to form H 2 O. Cancel waters that occur on both sides of the equation and reduce the equation to simplest terms. 6. Combine each half reaction to obtain the final balanced oxidation reduction equation (Multiply by appropriate factors and cancel species appearing on both sides of equation 17 17

18 Balance in acid solution: Mn +2 + BiO 3 - MnO Bi +3 Balance in Basic solution: S 2- + I 2 SO I

19 Balance in acid solution: Mn +2 + BiO - 3 MnO Bi +3 2(Mn H 2 O MnO H + + 5e - ) 5(6 H + + BiO e- Bi H 2 O) 2Mn H 2 O 2MnO H e - ) 30 H BiO e- 5Bi H 2 O) 2Mn H BiO 3-2 MnO Bi H 2 O 19 19

20 Balance in Basic solution: S 2- + I 2 SO I - 8 OH - +4H 2 O + S 2- SO H OH - + 8e- I e - 2I - 8 OH - +4H 2 O + S 2- SO H OH - + 8e- 4 (I e - 2I - ) 8 OH - +4H 2 O + S 2- SO H 2 O + 8e- 4 I e - 8I - 8 OH - + S I 2 SO I - +4 H 2 O 20 20

21 Dichromate ion in acidic solution is an oxidizing agent. When it reacts with zinc, the metal is oxidized to Zn 2+, and nitrate is reduced. Assume that dichromate ion is reduced to Cr 3+. Write the balanced ionic equation for this reaction using the half-reaction method

22 First we determine the oxidation numbers of N and Zn: Cr 2 O 7 2- (aq) Cr 3+ (aq) and Zn(s) Zn 2+ (aq) Cr was reduced from +6 to +3. Zn was oxidized from 0 to +2. Now we balance the half-reactions

23 Oxidation half-reaction Zn(s) Zn 2+ (aq) + 2e - Reduction half-reaction First we balance Cr and O: Cr 2 O 2-7 (aq) 2Cr 3+ (aq) + 7H 2 O(l) Next we balance H: 14H + (aq) + Cr 2 O 2-7 (aq) 2Cr 3+ (aq) + 7H 2 O(l) Then we balance e - : 6e H + (aq) + Cr 2 O 2-7 (aq) 2Cr 3+ (aq) + 7H 2 O(l) 23 23

24 Now we combine the two half-reactions by multiplying the oxidation half reaction by 3. 3Zn(s) 3Zn 2+ (aq) + 6e - 6e H + (aq) + Cr 2 O 7 2- (aq) 2Cr 3+ (aq) + 7H 2 O(l) 3Zn(s) + 14H + (aq) + Cr 2 O 7 2- (aq) 3Zn 2+ (aq) + 2Cr 3+ (aq) + 7H 2 O(l) Check that atoms and charge are balanced. 3Zn; 14H; 2Cr; 7O; charge

25 To balance a reaction in basic conditions, first follow the same procedure as for acidic solution. Then 1. Add one OH - to both sides for each H When H + and OH - occur on the same side, combine them to form H 2 O. 3. Cancel water molecules that occur on both sides. Balance in Basic solution: S 2- + I 2 SO I - Did This Earlier 25 25

26 Another example Lead(II) ion, Pb 2+, yields the plumbite ion, Pb(OH) 3-, in basic solution. In turn, this ion is oxidized in basic hypochlorite solution, ClO -, to lead(iv) oxide, PbO 2. Balance the equation for this reaction using the halfreaction method. The skeleton equation is Pb(OH) 3- + ClO - PbO 2 + Cl

27 First we determine the oxidation number of Pb and Cl in each species Pb(OH) 3- + ClO - PbO 2 + Cl - Pb is oxidized from +2 to +4. Cl is reduced from +1 to -1. Now we balance the half-reactions

28 Oxidation half-reaction First we balance Pb and O: Pb(OH) 3- (aq) PbO 2 (s) + H 2 O(l) Next we balance H: Pb(OH) 3- (aq) PbO 2 (s) + H 2 O(l) + H + (aq) Then we balance e - : Pb(OH) 3- (aq) PbO 2 (s) + H 2 O(l) + H + (aq) + 2e

29 Reduction half-reaction First we balance Cl and O: ClO - (aq) Cl - (aq) + H 2 O(l) Next we balance H: 2H + (aq) + ClO - (aq) Cl - (aq) + H 2 O(l) Finally we balance e - : 2e - + 2H + (aq) + ClO - (aq) Cl - (aq) + H 2 O(l) 29 29

30 Now we combine the half-reactions. Pb(OH) 3- (aq) PbO 2 (s) + H 2 O(l) + H + (aq) + 2e - 2e - + 2H + (aq) + ClO - (aq) Cl - (aq) + H 2 O(l) H + (aq) + Pb(OH) 3- (aq) + ClO - (aq) PbO 2 (s) + Cl - (aq) + 2H 2 O(l) 30 30

31 H + (aq) + Pb(OH) 3- (aq) + ClO - (aq) PbO 2 (s) + Cl - (aq) + 2H 2 O(l) To convert to basic solution, we add OH - to each side, converting H + to H 2 O. H 2 O(l) + Pb(OH) 3- (aq) + ClO - (aq) PbO 2 (s) + Cl - (aq) + 2H 2 O(l) + OH - (aq) Finally, we cancel the H 2 O that is on both sides. Pb(OH) 3- (aq) + ClO - (aq) PbO 2 (s) + Cl - (aq) + H 2 O(l) + OH - (aq) Do exercise 19.1 & 2 and see problems

32 The next several topics describe battery cells or voltaic cells (galvanic cells). An electrochemical cell is a system consisting of electrodes that dip into an electrolyte and in which a chemical reaction either uses or generates an electric current. A voltaic or galvanic cell is an electrochemical cell in which a spontaneous reaction generates an electric current. An electrolytic cell is an electrochemical cell in which an electric current drives an otherwise nonspontaneous reaction

33 Oxidation-Reduction Reactions In this chapter we will show how a cell is constructed to physically separate an oxidation-reduction reaction into two halfreactions. The force with which electrons travel from the oxidation half-reaction to the reduction half-reaction is measured as voltage

34 Electrochemistry An electrochemical cell is a system consisting of electrodes that dip into an electrolyte in which a chemical reaction either uses or generates an electric current. A voltaic, or galvanic, cell is an electrochemical cell in which a spontaneous reaction generates an electric current. An electrolytic cell is an electrochemical cell in which an electric current drives an otherwise nonspontaneous reaction. In this chapter we will discuss the basic principles behind these cells and explore some of their commercial uses

35 Voltaic Cells A voltaic cell consists of two half-cells that are electrically connected. Each half-cell is a portion of the electrochemical cell in which a half-reaction takes place. A simple half-cell can be made from a metal strip dipped into a solution of its metal ion. For example, the zinc-zinc ion half cell consists consists of a zinc strip dipped into a solution of a zinc salt

36 Voltaic Cells A voltaic cell consists of two half-cells that are electrically connected. Another simple half-cell consists of a copper strip dipped into a solution of a copper salt. In a voltaic cell, two half-cells are connected in such a way that electrons flow from one metal electrode to the other through an external circuit. Figure 19.2 illustrates an atomic view of a zinc/copper voltaic cell

37 Figure 19.2: Atomic view of voltaic cell

38 The zinc metal atom loses two electrons, forming Zn 2+ ions. The Cu 2+ ions gain two electrons, forming solid copper. The electrons flow through the external circuit from the zinc electrode to the copper electrode. x Ions flow through the salt bridge to maintain charge balance

39 Voltaic Cells As long as there is an external circuit, electrons can flow through it from one electrode to the other. Because zinc has a greater tendency to lose electrons than copper, zinc atoms in the zinc electrode lose electrons to form zinc ions. The electrons flow through the external circuit to the copper electrode where copper ions gain the electrons to become copper metal

40 Voltaic Cells The two half-cells must also be connected internally to allow ions to flow between them. Without this internal connection, too much positive charge builds up in the zinc half-cell (and too much negative charge in the copper half-cell) causing the reaction to stop. Figure 19.3A and 19.3B show the two half-cells of a voltaic cell connected by salt bridge

41 Figure 19.3: Two electrodes are connected by an external circuit

42 Voltaic Cells A salt bridge is a tube of an electrolyte in a gel that is connected to the two half-cells of a voltaic cell. The salt bridge allows the flow of ions but prevents the mixing of the different solutions that would allow direct reaction of the cell reactants. Figure 19.3C shows an actual setup of the zinc-copper cell

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44 Voltaic Cells The two half-cell reactions, as noted earlier, are: Zn(s) Zn 2 Cu 2 + (aq) + 2e + (aq) + 2e The first reaction, in which electrons are lost, is the oxidation half-reaction. The electrode at which oxidation occurs is the anode. Cu(s) (oxidation half-reaction) (reduction half-reaction) The second reaction, in which electrons are gained, is the reduction half-reaction. The electrode at which reduction occurs is the cathode

45 Voltaic Cells Note that the sum of the two half-reactions Zn(s) + Cu (aq) Zn (aq) Cu(s) is the net reaction that occurs in the voltaic cell; it is called the cell reaction Note that electrons are given up at the anode and thus flow from it to the cathode where reduction occurs. The anode in a voltaic cell has a negative sign because electrons flow from it. (See Figure 20.4 and Animation: Anode Reaction) The cathode in a voltaic cell has a positive sign (See 45 Animation: Cathode Reaction) 45

46 Figure 19.4: Voltaic Cell Do exercise 19.3 And look at Problems

47 Notation for Voltaic Cells It is convenient to have a shorthand way of designating particular voltaic cells. The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is written Zn(s) Zn anode (aq) Cu (aq) Cu(s) cathode The anode (oxidation half-cell) is written on the left. The cathode (reduction half-cell) is written on the right

48 The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is written Zn(s) Zn anode (aq) Cu salt bridge (aq) Cu(s) cathode The cell terminals are at the extreme ends in the cell notation. Zn(s) Zn anode (aq) Cu salt bridge (aq) Cu(s) cathode A single vertical bar indicates a phase boundary, such as between a solid terminal and the electrode solution

49 Notation for Voltaic Cells When the half-reaction involves a gas, an inert material such as platinum serves as a terminal and an electrode surface on which the reaction occurs. Figure 19.5 shows a hydrogen electrode; hydrogen bubbles over a platinum plate immersed in an acidic solution. The cathode half-reaction is + 2H (aq) + 2e H2(g) 49 49

50 The notation for the hydrogen electrode, written as a cathode, is + H (aq) H2 (g) Pt To write such an electrode as an anode, you simply reverse the notation. Pt H2(g) H + (aq) In the cell notation, these are written in parentheses. For example, Zn(s) Zn (1.0 M) H (aq) H2(1.0 atm) Pt 50 50

51 Figure 19.5: Hydrogen Electrode 51 51

52 A Problem To Consider Give the overall cell reaction for the voltaic cell Cd(s) Cd (1.0 M) H (aq) H2(1.0 atm) Pt The half-cell reactions are 2+ Cd(s) Cd (aq) + 2e + 2H (aq) + 2e H2(g) Cd(s) + 2H (aq) Cd (aq) + H2(g) Do Exercise 19.4 and see problems

53 Electromotive Force (Cell Potential) The movement of electrons is analogous to the pumping of water from one point to another. Water moves from a point of high pressure to a point of lower pressure. Thus, a pressure difference is required. The work expended in moving the water through a pipe depends on the volume of water and the pressure difference. An electric charge moves from a point of high electrical potential (high electrical pressure) to one of lower electrical potential. The work expended in moving the electrical charge through a conductor depends on the amount of charge and the potential difference

54 Electromotive Force Potential difference is the difference in electric potential (electrical pressure) between two points. You measure this quantity in volts. The volt, V, is the SI unit of potential difference equivalent to 1 joule of energy per coulomb of charge. Electrical work = charge x potential difference Joules = coulombs x volts 1 volt = 1 J C 54 54

55 Electromotive Force The Faraday constant, F, is the magnitude of charge on one mole of electrons; it equals 96,500 coulombs (9.65 x 10 4 C). In moving 1 mol of electrons through a circuit, the numerical value of the work done by a voltaic cell is the product of the Faraday constant (F) times the potential difference between the electrodes. work(j) = F(coulombs) work done by the system volts(j/coulomb) 55 55

56 Electromotive Force In the normal operation of a voltaic cell, the potential difference (voltage) across the electrodes is less than than the maximum possible voltage of the cell. The actual flow of electrons reduces the electrical pressure. Thus, a cell voltage has its maximum value when no current flows

57 Electromotive Force The maximum potential difference between the electrodes of a voltaic cell is referred to as the electromotive force (emf) of the cell, or E cell. It can be measured by an electronic digital voltmeter (Figure 19.6), which draws negligible current

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59 Electromotive Force We can now write an expression for the maximum work attainable by a voltaic cell. Let n be the number of (mol) electrons transferred in the overall cell reaction. The maximum work for molar amounts of reactants is w = max nfe cell Work Exercise 19.6 Look at Problems

60 A Problem To Consider The emf of the electrochemical cell below is V. Calculate the maximum electrical work of this cell when g H 2 is consumed. 2+ Hg (aq) + 2 H2(g) The half-reactions are 2Hg(l) + 2H + (aq) Hg H 2 (g) (aq) + 2e 2H + 2Hg(l) (aq) + 2e n = 2, and the maximum work for the reaction is written as w w = max nfe cell 4 = 2 max ( C) (0.650 V) 60 60

61 4 w = 2 max ( C) (0.650 V) w = max nfe cell w = max The emf of the electrochemical cell below is V. Calculate the maximum electrical work of this cell when g H 2 is consumed. 2+ Hg (aq) + 2 H2(g) For g H 2, the maximum work is 5 J 2Hg(l) + 2H + (aq) 5 1 mol H J g H 2 = g H 1 mol H J 61 61

62 Standard Cell emf s and Standard Electrode Potentials A cell emf is a measure of the driving force of the cell reaction. The reaction at the anode has a definite oxidation potential, while the reaction at the cathode has a definite reduction potential. Thus, the overall cell emf is a combination of these two potentials. E cell = oxidation potential + reduction potential 62 62

63 A reduction potential is a measure of the tendency to gain electrons in the reduction half-reaction. You can look at the oxidation half-reaction as the reverse of a corresponding reduction reaction. The oxidation potential for an oxidation half-reaction is the negative of the reduction potential for the reverse reaction

64 Standard Cell emf s and Standard Electrode Potentials By convention, the Table of Standard Electrode Potentials (Table 19.1) are tabulated as reduction potentials. Consider the zinc-copper cell described earlier. Zn(s) Zn (aq) Cu The two half-reactions are Zn(s) Zn Cu 2 + (aq) + 2e 2+ (aq) Cu(s) (aq) + 2e Cu(s) 64 64

65 The zinc half-reaction is an oxidation. If you write E Zn for the reduction potential of zinc, then E Zn is the oxidation potential of zinc. Zn 2 + (aq) + Zn(s) Zn 2e 2+ (aq) + Zn(s) 2e (E Zn ) Reduction -(E Zn ) The copper half-reaction is a reduction.. Write E Cu for the electrode potential. Cu 2 + (aq) + 2e Cu(s) (E Cu ) Oxidat. For this cell, the cell emf is the sum of the reduction potential for the copper half-cell and the oxidation potential for the zinc half-cell

66 E E cell cell = = E E Cu Cu + ( E E Note that the cell emf is the difference between the two electrode potentials. In general, E cell is obtained by subtracting the anode potential from the cathode potential. E = cell E cathode Zn Zn ) E anode The electrode potential is an intensive property whose value is independent of the amount of species in the reaction

67 Thus, the electrode potential for the half-reaction 2Cu 2 + (aq) + 4e 2Cu(s) is the same as for Cu 2 + (aq) + 2e Cu(s) 67 67

68 Tabulating Standard Electrode Potentials The standard emf, E cell, is the emf of a cell operating under standard conditions of concentration (1 M), pressure (1 atm), and temperature (25 C). Note that individual electrode potentials require that we choose a reference electrode. You arbitrarily assign this reference electrode a potential of zero and obtain the potentials of the other electrodes by measuring the emf s

69 By convention, the reference chosen for comparing electrode potentials is the standard hydrogen electrode. (see Figure 19.5) Standard electrode potentials (Table 19.1) are measured relative to this hydrogen reference

70 Tabulating Standard Electrode Potentials The standard electrode potential, E, is the electrode potential when concentrations of solutes are 1 M, gas pressures are 1 atm, and the temperature is 25 C. (Table 19.1) For example, when you measure the emf of a cell composed of a zinc electrode connected to a hydrogen electrode, you obtain 0.76 V. Since zinc acts as the anode (oxidation) in this cell, its reduction potential is listed as 0.76 V

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74 Strengths of Oxidizing and Reducing Agents Standard electrode potentials are useful in determining the strengths of oxidizing and reducing agents under standard-state conditions. A reduction half-reaction has the general form oxidized species + ne reduced species The oxidized species acts as an oxidizing agent. Consequently, the strongest oxidizing agents in a table of standard electrode potentials are the oxidized species corresponding to the half-reactions with the largest (most positive) E o values. (For example F 2 (g)) 74 74

75 An oxidation half-reaction has the general form reduced species oxidized species + ne The reduced species acts as a reducing agent. Consequently, the strongest reducing agents in a table of standard electrode potentials are the reduced species corresponding to the half-reactions with the smallest (most negative) E o values. (for example, Li(s)) 75 75

76 Which is the stronger reducing agent under standard conditions: Sn 2+ (to Sn 4+ ) or Fe (to Fe 2+ )? Which is the stronger oxidizing agent under standard conditions: Cl 2 or MnO 4-? The stronger reducing agent will be oxidized and has the more negative electrode potential. The standard (reduction) potentials are Sn 2+ to Sn 4+ E = 0.15 V Fe to Fe 2+ E = 0.41 V The stronger reducing agent is Fe (to Fe 2+ )

77 The stronger oxidizing agent will be reduced. The standard (reduction) potentials are Cl 2 to Cl - E = 1.36 V MnO 4- to Mn 2+ E = 1.49 V The stronger oxidizing agent is MnO

78 Cd Calculating Cell emf s from Standard Potentials The emf of a voltaic cell constructed from standard electrodes is easily calculated using a table of electrode potentials. Consider a cell constructed of the following two halfreactions. Ag Cd 2 + o = + (aq) + 2e (aq) + 1e Cd(s); Ag(s); E E o = 0.40 V 0.80 V You will need to reverse one of these reactions to obtain the oxidation part of the cell reaction. 2 + o = (aq) + 2e Cd(s); E 0.40 V Ag + (aq) + 1e Ag(s); E o = 0.80 V 78 78

79 Cd This will be Cd, because has the more negative electrode potential. Ag 2 + o = (aq) + + 2e (aq) + 1e Cd(s); Ag(s); E E o = 0.40 V 0.80 V Therefore, you reverse the half-reaction and change the sign of the half-cell potential. Cd(s) Ag + Cd (aq) + 1e (aq) o = 2e Ag(s); ; E E o = 0.40 V 0.80 V We must double the silver half-reaction so that when the reactions are added, the electrons 79 cancel. 79

80 Cd(s) Ag + Cd (aq) + 1e (aq) o = 2e ; Ag(s); E E o = 0.40 V 0.80 V This does not affect the half-cell potentials, which do not depend on the amount of substance. Cd(s) Cd + 2Ag(aq) + 2e (aq) o = 2e 2Ag ; (s); E E o = 0.40 V 0.80 V Now we can add the two half-reactions to obtain the overall cell reaction and cell emf. Cd(s) 2Ag + Cd (aq) + (aq) o = 2e 2e ; E 2Ag(s); E o = 0.40 V 0.80 V 80 80

81 2Ag Now we can add the two half-reactions to obtain the overall cell reaction and cell emf. + (aq) + Cd(s) + 2Ag Cd(s) + 2e Cd (aq) Cd (aq) o = 2Ag(s); 2e ; E E o = 2 + o (aq) + 2Ag(s); E = cell The corresponding cell notation would be 0.40 V 0.80 V 1.20 V Cd(s) Cd (1M) Ag (1M) Ag(s) Do Exercise 19.7 See Problems

82 Note that the emf of the cell equals the standard electrode potential of the cathode minus the standard electrode potential of the anode. E o cell = E o cathode E o anode Calculate the standard emf for the following voltaic cell at 25 C using standard electrode potentials. What is the overall reaction? Al(s) Al (aq) Fe (aq) Fe(s) 82 82

83 The reduction half-reactions and standard potentials are Al Fe 3 + o = (aq) + 3e (aq) + 2e Al(s); Fe(s); E 2 + o = E 1.66 V 0.41V Al(s) Al (aq) Fe (aq) Fe(s) You reverse the first half-reaction and its half-cell potential to obtain Al(s) Fe Al (aq) + 2e (aq) + 3e 3 + o = ; Fe(s); E 2 + o = E 1.66 V 0.41V 83 83

84 To obtain the overall reaction we must balance the electrons. 2Al(s) 2Al 3Fe Al(s) Al (aq) + 6e (aq) Fe (aq) + 6e 3 + o = ; 3Fe(s); (aq) Fe(s) E 2 + o = E 1.66 V 0.41V Now we add the reactions to get the overall cell reaction and cell emf. 2Al(s) 2Al 3Fe 2Al(s) + 3Fe (aq) + 6e (aq) + 6e 3 + o = ; 3Fe(s); E 2 + o = (aq) 2Al E (aq) + 3Fe(s); 1.66 V 0.41V o = E 1.25 V 84 84

85 How do you determine the direction of spontaneity? G Look at example 19.7 and example 19.8 See problems and

86 Look at Concept check 19.2 Page

87 Equilibrium Constants from emf s Some of the most important results from electrochemistry are the relationships among E cell, free energy, and equilibrium constant. In Chapter 18 we saw that DG equals the maximum useful work of a reaction. For a voltaic cell, work = -nfe o, so when reactants are in their standard states o G = nfe The measurement of cell emf s gives you yet another way of calculating equilibrium constants. o 87 87

88 Combining the previous equation, G o = -nfe o cell, with the equation G o = -RTlnK, we get nfe o cell = RT lnk Or, rearranging, we get E o cell = 2.303RT nf log K Substituting values for the constants R and F at 25 o C gives the equation E o cell = n log K 88 (values in volts at 25 o C) 88

89 Do Exercise 19.9 and see problems and Look at Examples 19.9 and Do Exercises and Look at problems & 74, 81 &

90 Figure 19.7 summarizes the various relationships among K, G o, and E o cell

91 A Problem To Consider The standard emf for the following cell is 1.10 V. E o cell = Zn(s) Zn (aq) Cu (aq) Cu(s) Calculate the equilibrium constant K c for the reaction Zn(s) + Cu (aq) Zn (aq) Cu(s) Note that n=2. Substituting into the equation relating E o cell and K gives nf RT log K 1.10 V = log K 91 91

92 Solving for log K c, you find log K = 37.2 Now take the antilog of both sides: K = c anti log(37.2) = The number of significant figures in the answer equals the number of decimal places in 37.2 (one). Thus K = c

93 Dependence of emf on Concentration Recall that the free energy change, G, is related to the standard free energy change, G, by the following equation. G = G o + RT lnq Here Q is the thermodynamic reaction quotient. Combining the previous equation, G o = -nfe o cell, with the equation G o = -RTlnK, we get o nfe = cell nfe + cell RT lnq 93 93

94 Dependence of emf on Concentration The result rearranges to give the Nernst equation, an equation relating the cell emf to its standard emf and the reaction quotient. E cell = E o cell 2.303RT nf logq Substituting values for R and F at 25 o C, we get E cell = E o cell n logq (values in volts at 25 o C) 94 94

95 Dependence of emf on Concentration The result rearranges to give the Nernst equation, an equation relating the cell emf to its standard emf and the reaction quotient. The Nernst equation illustrates why cell emf decreases as the cell reaction proceeds. As reactant concentrations decrease and product concentrations increase, Q increases, thus increasing log Q which in turn decreases the cell emf

96 A Problem To Consider What is the emf of the following voltaic cell at 25 C? Zn(s) Zn ( M) Cu The standard emf of the cell is 1.10 V. The cell reaction is Zn(s) + Cu (aq) Zn (0.100M) Cu(s) (aq) Cu(s) The number of electrons transferred is 2; hence n = 2. The reaction quotient is [Zn Q = [Cu ] 1.00 = 10 ] =

97 E The standard emf is 1.10 V, so the Nernst equation becomes cell = E cell = 1.10 V E o cell n logq log( E cell = 1.10 V ( 0.12) = 1.22 V 4 ) The cell emf is 1.22 V. Do Exercise Look at Problems &86 97

98 Concept Check 19.3 Page

99 Some Commercial Voltaic Cells The Leclanché dry cell, or zinc-carbon dry cell, is a voltaic cell with a zinc can as the anode and a graphite rod in the center surrounded by a paste of manganese dioxide, ammonium and zinc chlorides, and carbon black, as the cathode. The electrode reactions are Zn(s) Zn 2+ (aq) + 2e anode + 2NH (aq) + 2MnO (s) e ) Mn O (s) + H O(l) NH3(aq cathode (see Figure 19.9) 99 99

100 Figure 19.9: Leclanché Dry Cell

101 Some Commercial Voltaic Cells The alkaline dry cell, is similar to the Leclanché cell, but it has potassium hydroxide in place of ammonium chloride. The electrode reactions are Zn(s) Zn MnO (s) + H O(l) + 2 2e 2+ (aq) + 2e 2 Mn O (s) OH (aq) anode cathode (see Figure 19.10)

102 Figure 19.10: A Small Alkaline Dry Cell

103 Some Commercial Voltaic Cells The lithium-iodine battery is a solid state battery in which the anode is lithium metal and the cathode is an I 2 complex. The solid state electrodes are separated by a thin crystalline layer of lithium iodide. Although it produces a low current, it is very reliable and is used to power pacemakers. (see Figure 19.11)

104 Figure 19.11: Solid-State Lithium-Iodine Battery

105 Some Commercial Voltaic Cells The lead storage cell (a rechargeable cell) consists of electrodes of lead alloy grids; one electrode is packed with a spongy lead to form the anode, and the other electrode is packed with lead dioxide to form the cathode. The electrode reactions are + + HSO (aq) PbSO (s) H (aq) 2e Pb(s) 4 + PbO (s) + 3H (aq) + 2 HSO4 (aq) + 2e PbSO (s) 2H2O(l) 4 + anode cathode (see Figure 20.12)

106 Figure 19.12: Lead Storage Battery

107 Some Commercial Voltaic Cells The nickel-cadmium cell (nicad cell) consists of an anode of cadmium and a cathode of hydrated nickel oxide on nickel; the electrolyte is potassium hydroxide. The electrode reactions are + 2OH (aq) Cd(OH) (s) + 2e Cd(s) 2 NiOOH(s) + H O(l) + e 2 Ni(OH) (s) + 2 OH (aq) anode cathode (see Figure 19.14)

108 Figure 19.14: Nicad Storage Batteries

109 Some Commercial Voltaic Cells A fuel cell is essentially a battery, but differs by operating with a continuous supply of energetic reactants, or fuel. For a hydrogen-oxygen fuel cell, the electrode reactions are 2H (g) + 4OH (aq) H O(l) e (g) + 2H O(l) + 4e O2 2 4OH (aq) anode cathode (see Figure 19.15)

110 Figure 19.15: Hydrogen-Oxygen Fuel Cell To do: Research how a fuel cell works and what different types exist

111 Figure 19.16: The Electrochemical Process Involved in the Rusting of Iron

112 Figure 19.17: Cathodic Protection of Buried Steel Pipe

113 Figure 19.18: A Demonstration of Cathodic Protection Unprotected nail Fe 2+ + Ferricyanide

114

115 Electrolytic Cells An electrolytic cell is an electrochemical cell in which an electric current drives an otherwise nonspontaneous reaction. (See Video: Electrolysis of Water) The process of producing a chemical change in an electrolytic cell is called electrolysis. Many important substances, such as aluminum metal and chlorine gas are produced commercially by electrolysis

116 Figure 19.19: Electrolysis of Molten Sodium Chloride

117 Electrolysis of Molten Salts A Downs cell is a commercial electrochemical cell used to obtain sodium metal by electrolysis of molten NaCl. (see Figure 19.20) A number of other reactive metals are obtained by the electrolysis of a molten salt. Lithium, magnesium, and calcium metals are all obtained by the electrolysis of the chlorides. Do Exercise See Problems

118 Figure 19.20: Downs Cell for Preparation of Sodium Metal

119 Figure 19.22: Chlor-Alkali Membrane Cell

120 Figure 19.23: Chlor-Alkali Mercury Cell

121 Figure 19.24: Purification of Copper by Electrolysis

122 Figure 19.24: Purification of Copper by Electrolysis (cont d)

123

124 Stoichiometry of Electrolysis What is new in this type of stoichiometric problem is the measurement of numbers of electrons. You do not weigh them as you do substances. Rather, you measure the quantity of electric charge that has passed through a circuit. To determine this we must know the current and the length of time it has been flowing

125 Electric current is measured in amperes. An ampere (A) is the base SI unit of current equivalent to 1 coulomb/second. The quantity of electric charge passing through a circuit in a given amount of time is given by Electric charge(coul) = electric current (coul/sec) time lapse(sec)

126 A Problem To Consider When an aqueous solution of potassium iodide is electrolyzed using platinum electrodes, the halfreactions are 2I (aq) I2 (aq) + How many grams of iodine are produced when a current of 8.52 ma flows through the cell for 10.0 min? 2e 2H2O(l) + 2e H2(g) + 2OH (aq) When the current flows for 6.00 x 10 2 s (10.0 min), the amount of charge is ( = A) ( s) 5.11 C

127 Figure 19.21: Electrolysis of Aqueous Potassium Iodide

128 A Problem To Consider When an aqueous solution of potassium iodide is electrolyzed using platinum electrodes, the halfreactions are 2I (aq) I2 (aq) + How many grams of iodine are produced when a current of 8.52 ma flows through the cell for 10.0 min? 2e 2H O(l) + 2e 2 H2(g) + 2OH (aq) Note that two moles of electrons are equivalent to one mole of I 2. Hence, 1 mol e 1 mol I g I C = g I C 2 mol e 1 mol I

129 Do Exercise See Problems Do Exercise See Problems Do Exercise See problems

130 Operational Skills Balancing oxidation-reduction reactions Sketching and labeling a voltaic cell Writing the cell reaction from the cell notation Calculating the quantity of work from a given amount of cell reactant Determining the relative strengths of oxidizing and reducing agents Determining the direction of spontaneity from electrode potentials Calculating the emf from standard potentials

131 Operational Skills Calculating the free-energy change from electrode potentials Calculating the cell emf from free-energy change Calculating the equilibrium constant from cell emf Calculating the cell emf for nonstandard conditions Predicting the half-reactions in an aqueous electrolysis Relating the amounts of product and charge in an electrolysis

132 End of Chapter