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ELECTROCHEMISTRY Chapter 20

ELECTROCHEMISTRY Chapter 20. TRANSFER REACTIONS. Atom/Group transfer HCl + H 2 O  Cl - + H 3 O + Electron transfer Cu(s) + 2 Ag + (aq)  Cu 2+ (aq) + 2 Ag(s). Electron Transfer Reactions. Electron transfer reactions are oxidation-reduction or redox reactions.

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ELECTROCHEMISTRY Chapter 20

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  1. ELECTROCHEMISTRYChapter 20

  2. TRANSFER REACTIONS Atom/Group transfer HCl + H2O  Cl- + H3O+ Electron transfer Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s)

  3. Electron Transfer Reactions • Electron transfer reactions are oxidation-reduction or redox reactions. • Redox reactions can result in the generation of an electric current or be caused by imposing an electric current. • Therefore, this field of chemistry is often called ELECTROCHEMISTRY.

  4. Review of Terminology for Redox Reactions • OXIDATION—loss of electron(s) by a species; increase in oxidation number. • REDUCTION—gain of electron(s); decrease in oxidation number. Oxidation It Loses (electrons) Reduction It Gains (electrons)

  5. OXIDATION-REDUCTION REACTIONS Direct Redox Reaction Oxidizing and reducing agents in direct contact. Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s) PLAY MOVIE

  6. OXIDATION-REDUCTION REACTIONS Indirect Redox Reaction A battery functions by transferring electrons through an external wire from the reducing agent to the oxidizing agent. PLAY MOVIE

  7. Why Study Electrochemistry? • Batteries • Corrosion • Industrial production of chemicalssuch as Cl2, NaOH, F2 and Al • Biological redox reactions The heme group

  8. Electrochemical Cells • An apparatus that allows a redox reaction to occur by transferring electrons through an external connector. • Product favored reactionvoltaic or galvanic cell chemical change produces electric current • Reactant favored reaction electrolytic cell electric current used to cause chemical change. Batteries are voltaic cells

  9. Electrochemistry Alessandro Volta, 1745-1827, Italian scientist and inventor. Luigi Galvani, 1737-1798, Italian scientist and inventor.

  10. Reduction of VO2+ with Zn

  11. Electrons are transferred from Zn to Cu2+, but there is no useful electric current. CHEMICAL CHANGE ELECTRIC CURRENT With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.” Oxidation: Zn(s)  Zn2+(aq) + 2e- Reduction: Cu2+(aq) + 2e-  Cu(s) -------------------------------------------------------- Cu2+(aq) + Zn(s)  Zn2+(aq) + Cu(s)

  12. CHEMICAL CHANGE ELECTRIC CURRENT • To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. This is accomplished in a GALVANIC or VOLTAIC cell. A group of such cells is called a battery.

  13. Zn  Zn2+ + 2e- Cu2+ + 2e-  Cu Oxidation Anode Negative Reduction Cathode Positive • Electrons travel thru external wire. • Salt bridgeallows anions and cations to move between electrode compartments. Anions Cations

  14. The Cu|Cu2+ and Ag|Ag+ Cell

  15. Electrons move from anode to cathode in the wire. Anions & cations move thru the salt bridge. Electrochemical Cell PLAY MOVIE

  16. Terms Used for Voltaic Cells See Figure 20.6

  17. The Voltaic Pile Drawing done by Volta to show the arrangement of silver and zinc disks to generate an electric current. What voltage does a cell generate?

  18. BATTERIESPrimary, Secondary

  19. Dry Cell Battery Primary battery — uses redox reactions that cannot be restored by recharge. Anode (-) Zn Zn2+ + 2e- Cathode (+) 2 NH4+ + 2e-  2 NH3 + H2

  20. Alkaline Battery Nearly same reactions as in common dry cell, but under basic conditions. PLAY MOVIE Anode (-):Zn + 2 OH-ZnO + H2O + 2e- Cathode (+): 2 MnO2 + H2O + 2e-  Mn2O3 + 2 OH-

  21. Lead Storage Battery • Secondary battery • Uses redox reactions that can be reversed. • Can be restored by recharging

  22. Lead Storage Battery Anode (-)Eo = +0.36 V Pb + HSO4-PbSO4 + H+ + 2e- Cathode (+) Eo = +1.68 V PbO2 + HSO4- + 3 H+ + 2e- PbSO4 + 2 H2O PLAY MOVIE

  23. Ni-Cad Battery Anode (-) Cd + 2 OH-Cd(OH)2 + 2e- Cathode (+) NiO(OH) + H2O + e- Ni(OH)2 + OH- PLAY MOVIE

  24. 1.10 V 1.0 M 1.0 M CELL POTENTIAL, E • Electrons are “driven” from anode to cathode by an electromotive force or emf. • For Zn/Cu cell, this is indicated by a voltage of 1.10 V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M. Zn and Zn2+, anode Cu and Cu2+, cathode

  25. CELL POTENTIAL, E • For Zn/Cu cell, potential is +1.10 V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M. • This is the STANDARD CELL POTENTIAL, Eo • —a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 ˚C.

  26. Calculating Cell Voltage • Balanced half-reactions can be added together to get overall, balanced equation. Zn(s)  Zn2+(aq) + 2e- Cu2+(aq) + 2e-  Cu(s) -------------------------------------------- Cu2+(aq) + Zn(s)  Zn2+(aq) + Cu(s) If we know Eo for each half-reaction, we could get Eo for net reaction.

  27. CELL POTENTIALS, Eo Can’t measure 1/2 reaction Eo directly. Therefore, measure it relative to a STANDARD HYDROGEN CELL, SHE. 2 H+(aq, 1 M) + 2e- e H2(g, 1 atm) Eo = 0.0 V

  28. Negative electrode Positive electrode Zn/Zn2+ half-cell hooked to a SHE. Eo for the cell = +0.76 V Supplier of electrons Acceptor of electrons Zn  Zn2+ + 2e- Oxidation Anode 2 H+ + 2e-  H2 Reduction Cathode

  29. Reduction of H+ by Zn See Active Figure 20.13

  30. Overall reaction is reduction of H+ by Zn metal. Zn(s) + 2 H+ (aq)  Zn2+ + H2(g) Eo = +0.76 V Therefore, Eo for Zn  Zn2+ (aq) + 2e- is +0.76 V Zn is a (better) (poorer) reducing agent than H2.

  31. Cu/Cu2+ and H2/H+ Cell Eo = +0.34 V Positive Negative Acceptor of electrons Supplier of electrons Cu2+ + 2e-  Cu Reduction Cathode H2 2 H+ + 2e- Oxidation Anode

  32. Cu/Cu2+ and H2/H+ Cell Overall reaction is reduction of Cu2+ by H2 gas. Cu2+ (aq) + H2(g) Cu(s) + 2 H+(aq) Measured Eo = +0.34 V Therefore, Eo for Cu2+ + 2e- Cu is +0.34 V

  33. + Zn/Cu Electrochemical Cell Zn(s)  Zn2+(aq) + 2e- Eo = +0.76 V Cu2+(aq) + 2e-  Cu(s) Eo = +0.34 V --------------------------------------------------------------- Cu2+(aq) + Zn(s)  Zn2+(aq) + Cu(s) Eo (calc’d) = +1.10 V Anode, negative, source of electrons Cathode, positive, sink for electrons

  34. Uses of Eo Values Organize half-reactions by relative ability to act as oxidizing agents Cu2+(aq) + 2e-  Cu(s) Eo = +0.34 V Zn2+(aq) + 2e-  Zn(s) Eo = –0.76 V Note that when a reaction is reversed the sign of E˚ is reversed!

  35. Uses of Eo Values • Organize half-reactions by relative ability to act as oxidizing agents • Table 20.1 • Use this to predict direction of redox reactions and cell potentials.

  36. Potential Ladder for Reduction Half-Reactions See Figure 20.14 Best oxidizing agents Best reducing agents

  37. oxidizing o ability of ion E (V) 2+ Cu + 2e- Cu +0.34 + 2 H + 2e- H 0.00 2+ Zn + 2e- Zn -0.76 reducing ability of element TABLE OF STANDARD REDUCTION POTENTIALS 2

  38. Standard Redox Potentials, Eo Any substance on the right will reduce any substance higher than it on the left. • Zn can reduce H+ and Cu2+. • H2 can reduce Cu2+ but not Zn2+ • Cu cannot reduce H+ or Zn2+.

  39. Cu(s) | Cu2+(aq) || H+(aq) | H2(g) Cathode Positive Anode Negative Electrons r Cu2+ + 2e-  Cu Or Cu  Cu2+ + 2 e- H2 2 H+ + 2 e- or 2 H+ + 2e-  H2

  40. Cu(s) | Cu2+(aq) || H+(aq) | H2(g) Cathode Positive Anode Negative Electrons r Cu2+ + 2e-  Cu H2 2 H+ + 2 e- The sign of the electrode in Table 20.1 is the polarity when hooked to the H+/H2 half-cell.

  41. 2+ Cu + 2e- Cu +0.34 + 2 H + 2e- H2 0.00 2+ Zn + 2e- Zn -0.76 Standard Redox Potentials, Eo Ox. agent Red. agent Any substance on the right will reduce any substance higher than it on the left. • Northwest-southeast rule: product-favored reactions occur between • reducing agent at southeast corner • oxidizing agent at northwest corner

  42. Using Standard Potentials, EoTable 20.1 • In which direction do the following reactions go? • Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s) • Goes right as written • 2 Fe2+(aq) + Sn2+(aq) 2 Fe3+(aq) + Sn(s) • Goes LEFT opposite to direction written • What is Eonet for the overall reaction?

  43. 2+ Cu + 2e- Cu +0.34 + 2 H + 2e- H2 0.00 2+ Zn + 2e- Zn -0.76 Standard Redox Potentials, Eo CATHODE ANODE • Northwest-southeast rule: • reducing agent at southeast corner = ANODE • oxidizing agent at northwest corner = CATHODE

  44. Standard Redox Potentials, Eo E˚net = “distance” from “top” half-reaction (cathode) to “bottom” half-reaction (anode) E˚net = E˚cathode - E˚anode Eonet for Cu/Ag+ reaction = +0.46 V

  45. Eo for a Voltaic Cell Cd Cd2+ + 2e- or Cd2+ + 2e- Cd Fe Fe2+ + 2e- or Fe2+ + 2e- Fe All ingredients are present. Which way does reaction proceed?

  46. Eo for a Voltaic Cell From the table, you see • Fe is a better reducing agent than Cd • Cd2+ is a better oxidizing agent than Fe2+ Overall reaction Fe + Cd2+Cd + Fe2+ Eo = E˚cathode - E˚anode = (-0.40 V) - (-0.44 V) = +0.04 V

  47. More About Calculating Cell Voltage 2 H2O + 2e- H2 + 2 OH- Cathode 2 I-I2 + 2e- Anode ------------------------------------------------- 2 I- + 2 H2O I2 + 2 OH- + H2 Assume I- ion can reduce water. Assuming reaction occurs as written, E˚net = E˚cathode - E˚anode = (-0.828 V) - (+0.535 V) = -1.363 V Minus E˚ means rxn. occurs in opposite direction

  48. Michael Faraday1791-1867 Originated the terms anode, cathode, anion, cation, electrode. Discoverer of • electrolysis • magnetic props. of matter • electromagnetic induction • benzene and other organic chemicals Was a popular lecturer.

  49. Eo and Thermodynamics • Eo is related to ∆Go, the free energy change for the reaction. • ∆G˚ proportional to –nE˚ ∆Go = -nFEo where F = Faraday constant = 9.6485 x 104 J/V•mol of e- (or 9.6485 x 104 coulombs/mol) and n is the number of moles of electrons transferred

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