1. Chapter 17 Electrochemistry  Redox review (4.9)  17.1-17.2  17.4-17.5  17.6-17.7

2. Review Oxidation-Reduction  Involves transfer of electrons from reducing agent to oxidizing agent  Oxidation= loss of e - (increase in oxid #)  Reduction= gain of e - (decrease in oxid#)  GER and LEO

3. REVIEW 1. atom in element = 0 2. monatomic ion = charge 3. fluorine = -1 4. oxygen = -2 5. hydrogen = +1 6. sum of oxid. # in compound = 0 7. sum of oxid. # in polyatomic ion = charge on ion

4. Copyright © Houghton Mifflin Company. All rights reserved. 4–44–4 The Half-Reaction Method (Acidic Solution)

5. Review- Balancing Oxidation- Reduction Reactions 1. Separate in ½ reactions 2. Intermediate steps a. balance all elements other than H and O b. balance O with H 2 O c. balance H with H + d. balance charge with (e-) 3. Multiply ½ rxn. so that the number of electrons is same 4. Add ½ rxns.

6. Capture the Energy MnO 4 - + 5Fe 2+  Mn 2+ 5Fe 3+ MnO 4 - and Fe 2+ will react directly in solution. Electrons will be transferred and energy will be released as heat. No useful work will result.

7. Capture the Energy!  Zn + Cu 2+ -  Zn 2+ + Cu  Separate ½ reactions  Connect metals w/ wire (to transfer electrons)  Connect soln w/ bridge (keeps solns separate but allows ions to move)  Converts Chemical Energy to Electrical Energy!!- A Battery!!

8. Galvanic Cell

9. Capture the energy  You have separated the oxidizing agent from the reducing agent  Requires electron transfer through wire  Attach a motor, light bulb, bell etc- the current produced in the wire by e- flow provides work!!

10. Copyright © Houghton Mifflin Company. All rights reserved. 17–10 Figure 17.6 A Galvanic Cell involving the Half-Reactions

11. Cell potential is…..  The pressure of a Galvanic cell to “push” the e- “driving force”  Electromotive Force, emf  Symbol E  Units: Joule/coulomb (=1Volt, V)  Coulomb = unit of charge Specifies # of e-

12. E cell = E anode + E cathode (oxidation)(reduction) pushing e-pulling e- (black wire)(red wire) A spontaneous rxn in a Galvanic cell must be positive. E > 0

13. E 1/2 reactions  P. 796 table  Standard Reduction potentials  1M solutions  1atm gases  25 C  Hydrogen ½ rxn = 0.00V

14. Table 17.1 Standard Reduction Potentials at 25°C (298K) for Many Common Half-Reactions Copyright © Houghton Mifflin Company. All rights reserved. 17–14

15. Helpful Info Need balanced oxidation-reduction rxns from the reduction potentials. One reduction ½ rxn must be reversed. * The ½ rxn with largest positive potential will run as written (reduction). The other ½ rxn will run in reverse (oxidation).

16. Reversing Direction Changes Sign of E Because: E oxidation = - E reduction Then: E cell = E cathode – E anode Examples:

17. Standard Reduction Potential Math Rules # of e - lost must equal # e - gained ½ rxns must be multiplied by integers to balance equations Value of E is not changed when ½ rxn multiplied by an integer. Potential is NOT multiplied by integer. Example….

18. Line Notation Anode listed on left Cathode listed on right Mg (s) l Mg 2+ ll Al 3+ l Al (s) Anode Mg 0 (s) -  Mg 2+ Cathode Al 3+ -  Al 0 (s)

19. Cell Potential & Free Energy A galvanic cell will run in the direction that gives a positive value for E + E corresponds to - G + E and - G indicates a spontaneous reaction. G = -n F E

20. G = nFE  n = # of e- (exchanged in overall rxn)  F = 96,485(c/mol e-) (Faraday’s constant)  Examples:

21. Effects of Concentration on E So far the cells have been under standard conditions…. Le Chatelier’s principle applies if not std. conditions.. Determine if E cell > or < E cell ??

22. To summarize: If E cell not at standard conditions: [Reactant] > 1mol/L E cell > E *cell [Product] E *cell Reverse is also true

23. Concentration cell Same components in cells, but different concentrations. Equilibrium wants these concentrations to be Equal. Examples:

24. Nernst Equation Establishes relationship b/t cell potential and concentration of cell components. For cells not at 1M Concentration: E = E * - RT/nF ln (Q) E * is std cell potential RT/nF ln (Q) is correction factor

25. Common form:  E = E * - RT/nF ln (Q) Commonly written :  E = E * - 0.0591/n log (Q)  Examples:

26. A Battery @ Equilibrium At Equilibrium: E cell = 0 (completely discharged) Q = K and delta G = 0 Using the Nernst Equation: @Equilibrium: 0 = E * - 0.0591/n log(K) Or log K = n E */0.0591

27. Corrosion Process of returning metals to their natural state. Metals oxidize readily resulting in corrosion. Metal ½ rxn is reversed for oxidation. Combined with Oxygen ½ rxn. to give (+) Ecell

28. Electrolysis  Involves forcing current through a cell to produce a chemical change resulting in (-) cell potential.  Example:

29. Copyright © Houghton Mifflin Company. All rights reserved. 17–29 Figure 17.19 a-b (a) A Standard Galvanic Cell Based on the Spontaneous Reaction Zn + Cu 2+ - Zn 2+ + Cu (b) A Standard Electrolytic Cell. A Power Source Forces the Opposite Reaction Cu + Zn 2+ - Cu 2+ + Zn.