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Chapter 20 Principles of Chemical Reactivity: Electron Transfer Reactions. Electron Transfer Reactions. Electron transfer reactions are oxidation-reduction or redox reactions.

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Chapter 20 Principles of Chemical Reactivity: Electron Transfer Reactions


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    1. Chapter 20Principles of Chemical Reactivity: Electron Transfer Reactions

    2. 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.

    3. 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. • OXIDIZING AGENT:electron acceptor; species is reduced. • REDUCING AGENT: electron donor; species is oxidized.

    4. Electrochemistry Applications: • Batteries • Corrosion • Industrial production of chemicals such as Cl2, NaOH, F2 and Al • Biological electron transfer reactions The heme group

    5. Electrochemical Cells • An apparatus that allows an electron transfer reaction to occur via an external connector. • Voltaic or galvanic cells: Reactions that produce a chemical current that are product favored • Electrolytic cell: Reactions which require an electric current to cause chemical change. (reactant favored) Batteries are voltaic cells

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

    7. Cu(s) + 2Ag+ (aq)  Cu2+ (aq) + 2Ag(s) Oxidation/Reduction Reactions

    8. Oxidation-Reduction Reactions: “Redox” Direct Redox Reaction Oxidizing and reducing agents in direct contact. Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s)

    9. Oxidation-Reduction Reactions: “Redox” Indirect Redox Reaction A battery functions by transferring electrons through an external wire from the reducing agent to the oxidizing agent.

    10. Oxidation-Reduction Reactions: “Redox” In any Oxidation Reduction reaction: • One species is Oxidized and one species is Reduced. • Neither can occur alone. • That which is reduced is the Oxidizing Agent. • That which is Oxidized is the Reducing Agent. • The total number of electrons lost in oxidation must equal the total number of electrons gained in reduction. • Redox reactions must therefore be balanced for mass and charge.

    11. Balancing Redox Equations Step 1: Divide the reaction into half-reactions, one for oxidation and the other for reduction. Ox Cu(s)  Cu2+(aq) Red Ag+(aq)  Ag(s) Step 2:Balance each for mass. Already done in this case. Step 3:Balance each half-reaction for charge by adding electrons. Ox Cu(s)  Cu2+(aq) + 2e Red Ag+(aq) + e Ag(s)

    12. Balancing Redox Equations Step 4: Multiply each half-reaction by a factor so that the reducing agent supplies as many electrons as the oxidizing agent requires. Reducing agent: Cu(s)  Cu2+ (aq) + 2e Oxidizing agent: 2Ag+(aq) + 2e 2Ag(s) Step 5: Add half-reactions to give the overall equation. Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2Ag(s) The equation is now balanced for both charge and mass.

    13. Balancing Equations for Redox Reactions Some redox reactions have equations that must be balanced by special techniques MnO4- + 5 Fe2+ + 8 H+ Mn2+ + 5 Fe3+ + 4 H2O Mn = +7 Fe = +2 Mn = +2 Fe = +3

    14. Reduction of VO2+ by Zn

    15. Balancing Redox Equations Balance the following in acid solution: VO2+ + Zn  VO2+ + Zn2+ Step 1: Write the half-reactions Ox Zn  Zn2+ Red VO2+ VO2+ Step 2: Balance each half-reaction for mass. Ox Zn  Zn2+ Red 2 H++ VO2+ VO2+ + H2O Add H2O on O-deficient side and add H+ on other side for H-balance.

    16. Balancing Redox Equations Step 3: Balance half-reactions for charge. Ox Zn  Zn2+ + 2e- Red e- + 2 H+ + VO2+ VO2+ + H2O Step 4: Multiply by an appropriate factor. Ox Zn  Zn2+ + 2e- Red 2e- + 4 H+ + 2 VO2+  2 VO2+ + 2 H2O Step 5: Add balanced half-reactions Zn + 4 H+ + 2 VO2+ Zn2+ + 2 VO2+ + 2 H2O

    17. Tips on Balancing Redox Reaction Equations • Never add O2, O atoms, or O2- to balance oxygen. • Never add H2 or H atoms to balance hydrogen. • Be sure to write the correct charges on all the ions. • Check your work at the end to make sure mass and charge are balanced. • PRACTICE!

    18. Chemical Change & Electric Current Electrons are transferred from Zn to Cu2+, but there is no useful electric current. Oxidation: Zn(s)  Zn2+(aq) + 2e- Reduction: Cu2+(aq) + 2e- Cu(s) -------------------------------------------------------- Cu2+(aq) + Zn(s)  Zn2+(aq) + Cu(s) With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.”

    19. Chemical Change & Electric Current • To obtain a current that can do work, the oxidizing and reducing agents are separated so that electron transfer occurs via an external wire. This is accomplished in a GALVANIC or VOLTAIC cell. A group of such cells is called a battery.

    20. Chemical Change & Electric Current Anode • The electrode where oxidation occurs • The electrode where mass is lost • The electrode that attracts anions Cathode • The electrode where reduction occurs • The electrode where mass is gained • The electrode that attracts cations

    21. Zn  Zn2+ + 2e- Cu2+ + 2e-  Cu •Electrons travel thru external wire. • Salt bridge allows anions and cations to move between electrode compartments. Oxidation Anode Negative Reduction Cathode Positive Anions Cations

    22. The Cu|Cu2+ & Ag|Ag+ Cell

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

    24. Terminology for Voltaic Cells

    25. The Voltaic Pile An arrangement of silver and zinc disks used to generate an electric current is show in this drawing by Alessandro Volta.

    26. Cell Potential, E 1.10 V 1.0 M 1.0 M • 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

    27. 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, E° • E°is a quantitative measure of the tendency of reactants to proceed to products at standard state conditions: 1 atm, 1M solutions @ 25 °C. • Cell potential in a voltaic cell is the electrical potential energy difference between the cathode and anode. • Cell potential depends on the ease of reduction at the cathode verses that of the anode.

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

    29. Cell Potentials, E° • Individual oxidation/reduction half reaction Eo values cannot be measured directly. • They are measured relative to a standard reference cell. The Standard Hydrogen Electrode (SHE) 2 H+(aq, 1 M) + 2e H2(g, 1 atm) The potential of the cell is set to an E° = 0.0 V

    30. Negative electrode Positive electrode Zn/Zn2+ half-cell hooked to a SHE. E° for the cell = +0.76 V Electron donor Electron acceptor Zn  Zn2+ + 2e- Oxidation Anode 2 H+ + 2e-  H2 Reduction Cathode

    31. Reduction of H+ by Zn

    32. Reduction of H+ by Zn Overall reaction is reduction of H+ by Zn metal. Zn(s) + 2 H+ (aq)  Zn2+ + H2(g) E° = +0.76 V Therefore, E° for Zn  Zn2+ (aq) + 2e- is +0.76 V Zn is a strongerreducing agent than H2.

    33. Cu/Cu2+ and H2/H+ Cell E° = +0.34 V Positive electrode Negative electrode Electron donor Electron acceptor Cu2+ + 2e- f Cu Reduction Cathode H2f 2 H+ + 2e- Oxidation Anode

    34. 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 E° = +0.34 V Therefore, E° for Cu2+(aq)+ 2e-  Cu(s) is +0.34V

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

    36. E° Values Half-reactions are organized by relative ability to act as oxidizing agents (Reduction Potential) Cu2+(aq) + 2e-  Cu(s) E° = +0.34 V Zn2+(aq) + 2e-  Zn(s) E° = –0.76 V Note that when a reaction is reversed the sign of E˚ is reversed!

    37. E° Values Table 20.1 tabulates the standard reduction potentials for half reactions from highest (most positive) to lowest (most negative) The greater the reduction potential (more positive) the greater the tendency to undergo reduction. One can use these values to predict the direction and cell potentials of redox reactions.

    38. Potential Ladder for Reduction Half-Reactions Stronger oxidizing agents Stronger reducing agents

    39. Using Standard Potentials, E° • Which species is the strongest oxidizing agent: O2, H2O2, or Cl2? _________________ • Which is the strongest reducing agent: Hg, Al, or Sn? ____________________

    40. Using Standard Potentials, E° • Which species is the strongest oxidizing agent: O2, H2O2, or Cl2? H2O2 > Cl2 > O2 • Which is the strongest reducing agent: Hg, Al, or Sn? ____________________

    41. Using Standard Potentials, E° • Which species is the strongest oxidizing agent: O2, H2O2, or Cl2? H2O2 > Cl2 > O2 • Which is the strongest reducing agent: Hg, Al, or Sn? Al > Sn > Hg

    42. Standard Redox Potentials, E° oxidizing 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 2

    43. Standard Redox Potentials, E° A 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+.

    44. Cu(s) | Cu2+(aq) || H+(aq) | H2(g) Cathode Positive Anode Negative Electrons r H2(g) Cu(s) Salt Bridge KNO3(aq) 1 M C u ( N O ) + 1 M H O 3 2 3 Cu2+ + 2e-  Cu Or Cu  Cu2+ + 2 e- H2 2 H+ + 2 e- or 2 H+ + 2e-  H2

    45. Galvanic Cells Zn (s) + Cu2+(aq) Cu (s) + Zn2+(aq) • The difference in electrical potential between the anode and cathode is called: • cell voltage • electromotive force (emf) • cell potential [Cu2+] = 1 M and [Zn2+] = 1 M Cell Diagram phase boundary Zn (s) | Zn2+ (1 M)|| Cu2+ (1 M) | Cu (s) anode cathode salt bridge

    46. Standard Redox Potentials, E° 2+ Cu + 2e- Cu +0.34 + 2 H + 2e- H2 0.00 2+ Zn + 2e- Zn -0.76 Ox. agent Red. agent Any substance on the right will reduce any substance higher than it on the left. • Northwest-southeast rule: product-favored when… • The reducing agent is the species at southeast corner • The oxidizing agent is the species at northwest corner

    47. Using Standard Potentials, E° In which direction do the following reactions go?

    48. Using Standard Potentials, E° In which direction do the following reactions go? Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s)

    49. Using Standard Potentials, E° • In which direction do the following reactions go? • Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s) • Goes right as written