1. Chapter 21 Electrochemistry

2. Voltaic Cells  Electrochemical cells used to convert chemical energy into electrical energy  Produced by spontaneous redox reactions within the cell

3. Decomposing Water 2 H 2 O  2 H 2 + O 2 The H + is reduced at the (-) cathode, yielding H 2 (g), which is trapped in the tube. The O -2 is oxidized at the (+) anode, yielding O 2 (g), which is trapped in the tube.

4. Electroplating The Ag 0 is oxidized to Ag +1 when the (+) end of the battery strips its electrons off. The Ag +1 migrates through the solution towards the (-) charged cathode (ring), where it picks up an electron from the battery and forms Ag 0, which coats on to the ring.

5.  Half-cell  one part of a voltaic cell in which either oxidation or reduction occurs  Salt bridge  a tube connecting half-cells that contains a strong electrolyte –Allows ions to pass from one half-cell to the other but prevents the solutions from mixing completely –Necessary to complete the circuit and maintain charge neutrality Parts of a voltaic cell

6.  Electrode  a conductor in a circuit that carries electrons to or from a substance –Anode  electrode at which oxidation occurs Electrons are produced; labeled negative electro de –Cathode  electrode at which reduction occurs Electrons are consumed; labeled positive electrode

7. Voltaic Cells How they work  A typical cell looks like this.  The oxidation occurs at the anode.  The reduction occurs at the cathode.

8. Voltaic Cells Once even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would stop.

9. Voltaic Cells  Therefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced. –Cations move toward the cathode. –Anions move toward the anode.

10. Voltaic Cells  In the cell, then, electrons leave the anode and flow through the wire to the cathode.  As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment.

11. Voltaic Cells  As the electrons reach the cathode, cations in the cathode are attracted to the now negative cathode.  The electrons are taken by the cation, and the neutral metal is deposited on the cathode.

12.  The two half-reactions can be summed to show the overall cell reaction  Oxidation: Zn (s)  Zn +2 + 2é  Reduction: Cu +2 + 2é  Cu  Overall: Zn (s) + Cu +2  Zn +2 + Cu (s)

13. Electromotive Force (emf)  Water only spontaneously flows one way in a waterfall.  Likewise, electrons only spontaneously flow one way in a redox reaction—from higher to lower potential energy.

14. Electromotive Force (emf)  The potential difference between the anode and cathode in a cell is called the electromotive force (emf).  It is also called the cell potential, and is designated E cell.  Cell potential is measured in volts

15. Standard Reduction Potentials Reduction potentials for many electrodes have been measured and tabulated.

16. Standard Cell Potentials The cell potential at standard conditions can be found through determining which substance will be oxidized and which will be reduced. Because cell potential is based upon the potential energy per unit of charge, it is an intensive property

17. Cell Potentials For the oxidation in this cell,  For the reduction,  Total voltage  0.76 + 0.34  1.10 Volts E Zn = +0.76 V  E Cu = +0.34 V 

18. Oxidizing and Reducing Agents  The strongest oxidizers have the most positive reduction potentials.  The strongest reducers have the most negative reduction potentials.

19. Oxidizing and Reducing Agents The greater the difference between the two, the greater the voltage of the cell.

21.  Voltage for Al/Pb voltaic cell  Write both reduction potentials then determine which will be forced to be oxidized  Al 3+ + 3é  Al-1.66V  Pb 2+ + 2é  Pb -0.13V  Al  Al 3+ + 3é 1.66V 1.53V  Full reaction  Pb 2+ + Al  Al 3+ + Pb 

23.  Voltage for Sn/Cu voltaic cell  Sn 2+ + 2é  Sn-0.14V  Cu 2+ + 2é  Cu 0.34V  Sn  Sn 2+ + 2é 0.14V 0.48V Full reaction Cu 2+ + Sn  Sn 2+ + Cu

25.  Voltage for Zn/Ag voltaic cell  Zn 2+ + 2é  Zn-0.76V  Ag + + é  Ag 0.80V  Zn  Zn 2+ + 2é 0.76V 1.56V Full reaction Zn + Ag +  Zn 2+ + Ag

26. Applications of Oxidation-Reduction Reactions

27. Hydrogen Fuel Cells

28. Batteries

29. Charging a Battery When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards. This is why the ammeter in your car often goes slightly higher while your battery is charging, and then returns to normal. In your car, the battery charger is called an alternator. If you have a dead battery, it could be the battery needs to be replaced OR the alternator is not charging the battery properly.

30. Dry Cell Battery Anode (-) Zn ---> Zn 2+ + 2e- Cathode (+) 2 NH 4 + + 2e- ---> 2 NH 3 + H 2

31. Alkaline Battery Nearly same reactions as in common dry cell, but under basic conditions. Anode (-): Zn + 2 OH - ---> ZnO + H 2 O + 2e- Cathode (+): 2 MnO 2 + H 2 O + 2e- ---> Mn 2 O 3 + 2 OH -

32. Mercury Battery Anode: Zn is reducing agent under basic conditions Cathode: HgO + H 2 O + 2e- ---> Hg + 2 OH -

33. Lead Storage Battery Anode (-) E o = +0.36 V Pb + HSO 4 - ---> PbSO 4 + H + + 2e- Cathode (+) E o = +1.68 V PbO 2 + HSO 4 - + 3 H + + 2e- ---> PbSO 4 + 2 H 2 O

34. Ni-Cad Battery Anode (-) Cd + 2 OH - ---> Cd(OH) 2 + 2e- Cathode (+) NiO(OH) + H 2 O + e- ---> Ni(OH) 2 + OH -

35. H 2 as a Fuel Cars can use electricity generated by H 2 /O 2 fuel cells. H 2 carried in tanks or generated from hydrocarbons