1. Reduction- Oxidation Reactions (1) 213 PHC 9 th lecture Dr. mona alshehri (1) Gary D. Christian, Analytical Chemistry, 6 th edition. 1

2. By the end of the lecture the student should be able to: Understand the principals of electrochemical cells. Understand the principals of electrode potential. Calculate the cell potential. 2

3. The reducing or oxidizing tendency of a substances will depend on its reduction potential 3

4. Electrochemical Cells Oxidation-reduction reactions take place in electrochemical cells. 4 electrochemical cells Galvanic cells Electrolytic cells

5. ◦ Galvanic cells  A spontaneous reaction occurs and produce electrical energy. ◦ Electrolytic cells  Electrical energy is used to force a non- spontaneous reaction to occur. 5

6. Both cells contain electrodes where the oxidation and reduction reactions occur:electrodes Oxidation occurs at the electrode called the anode. Reduction occurs at the electrode called the cathode. Electrons flow from the anode to the cathode. 6

7. The anode of a galvanic cell is negatively charged, since the spontaneous oxidation at the anode is the source of the cell's electrons or negative charge. The cathode of a galvanic cell is its positive terminal. 7

8. Galvanic cell 8

9. Cu +2 (aq) + Zn (s) ---> Cu(s) + Zn +2 (aq) At the anode, the oxidation of zinc occurs: Zn ---> Zn +2 + 2e - At the cathode, the reduction of copper occurs: Cu +2 + 2e - ---> Cu 9

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11. The two cells are connected via a salt bridge. A salt bridge is a porous barrier which prevents the spontaneous mixing of the aqueous solutions in each cell, but allows the migration of ions in both directions to maintain electrical neutrality. 11

12. The two half-cells are also connected externally. Electrons provided by the oxidation reaction are forced to travel via an external circuit to the site of the reduction reaction. Each electrode will adopt an electrical energy called the electrode potential. Once these half cells are connected, the difference in electrode potential can be measured. This difference in potential energy between the 2 electrodes is measured in terms of volts. 12

13. Electrode Potential (E o ) Each half-reaction will generate a potential. Individual electrode potential can’t be measured. The difference between 2 electrode potentials can be measured. The standard hydrogen electrode is used to measure the potential of any half reaction because it’s potential is zero. 13

14. The more +ve E o = (oxidation). The more -ve E o = (reduction). E cell = E + - E - 14

15. Fe 3+ + e -  Fe 2+ E o = 0.771 V Sn 4+ + 2e -  Sn 2+ E o = 0.154 V 2Fe 3+ + Sn 2+  2Fe 2+ + Sn 4+ 0.617 E cell = 0.771- 0.154 = 15

16. Questions? 16

17. Homework: Write the 2 halves reaction: 17

18. Homework: What is the overall cell reaction and the cell potential for the two half-reactions? A) Cu 2+ + 2e = Cu E o = 0.34 V Zn 2+ + 2e = Zn E o = -0.76 V B) Fe 3+ + e = Fe 2+ E o = 0.77 V Ti 4+ + e = Ti 3+ E o = 0.15 V 18

19. Summary: Electrochemical cells. Electrode potential. Calculation of cell potential. 19

20. THANK YOU 20