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Electron Transfer Reactions

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.

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Electron Transfer Reactions

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  1. 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.

  2. 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

  3. 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-

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

  5. 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

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

  7. 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.

  8. OXIDATION-REDUCTION REACTIONS Understand a DirectRedox Reaction Oxidizing and reducing agents in direct contact. Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s) PLAY MOVIE

  9. Cu + Ag+Cu2+ + Ag Know how to Balance Equations in a neutral solution

  10. Know how to Balance Equations in a neutral solution Step 1: Divide the reaction into half-reactions, one for oxidation and the other for reduction. Ox Cu Cu2+ Red Ag+ Ag Step 2: Balance each for mass. Already done in this case. Step 3: Balance each half-reaction for charge by adding electrons. Ox Cu Cu2+ + 2e- Red Ag+ + e- Ag

  11. Know how to Balance Equations in a neutral solution 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 Cu2+ + 2e- Oxidizing agent 2 Ag+ + 2 e- 2 Ag Step 5: Add half-reactions to give the overall equation. Cu + 2 Ag+ Cu2+ + 2Ag The equation is now balanced for both charge and mass.

  12. Know how to Balance Equations in an acidic solution Balance the following in acidsolution— 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.

  13. Know how to Balance Equations in an acidic solution 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

  14. Know how to Balance Equations in a basic solution • Two approaches: • Follow the same steps as for an acid solution but add OH- rather than H+ in step 2 • Balance as if it were an acidic solution then add enough OH- to both sides to neutralize the H+

  15. Tips on Balancing 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

  16. Terms Used for Voltaic Cells See Figure 20.6

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

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

  19. 1.10 V 1.0 M 1.0 M CELL POTENTIAL, E Zn and Zn2+, anode Cu and Cu2+, cathode • 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.

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

  21. Using Standard Potentials, EoTable 20.1 • Which is the best oxidizing agent: O2, H2O2, or Cl2? _________________ • Which is the best reducing agent: Hg, Al, or Sn? ____________________

  22. Standard Redox Potentials, Eo • Zn can reduce H+ and Cu2+. • H2 can reduce Cu2+ but not Zn2+ • Cu cannot reduce H+ or Zn2+.

  23. Number of moles of electrons Faraday’s constant

  24. or

  25. Electrolysis Using electrical energy to produce chemical change. Sn2+(aq) + 2 Cl-(aq) Sn(s) + Cl2(g)

  26. Electrolysis of Aqueous CuCl2 Anode (+) 2 Cl-Cl2(g) + 2e- Cathode (-) Cu2+ + 2e- Cu Eo for cell = -1.02 V

  27. Electrolytic Refining of Copper See Figure 22.11 Impure copper is oxidized to Cu2+ at the anode. The aqueous Cu2+ ions are reduced to Cu metal at the cathode.

  28. Producing Aluminum 2 Al2O3 + 3 C f4 Al + 3 CO2 Charles Hall (1863-1914) developed electrolysis process. Founded Alcoa.

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