1. Redox Reactions (con’t): eg: Al (s) + Cu 2+ (aq)  Cu (s) + Al 3+ (aq) The fact that this reaction takes place tells us that:  Cu 2+ is a stronger oxidizing agent than Al 3+ and/or  Cu 2+ is more easily reduced than Al 3+ “Standard Reduction Potentials” are shown on p. 8 of the data booklet. The strongest oxidizing agent is at the top left (it wants electrons most) The strongest reducing agent is at the bottom right (it wants to give up electrons most)

2. SRP’s of Half-Cells Oxidizing AgentsReducing Agents STRENGTH OF OXIDIZING AGENT STRENGTH OF REDUCING AGENT WEAK STRONGWEAK STRONG Cu 2+ + 2e - Cu (s) Al 3+ + 3e - Al (s)

3. A substance on the left will take electrons away from (react with) any substance on the right that is lower down. eg: If you combine F 2 and Cu, F 2 will take electrons from Cu because F 2 is a stronger oxidizing agent than Cu 2+ : F 2 + Cu  Cu 2+ + 2F - eg: Write the reaction that takes place between I 2 and Zn. I 2 + 2e -  2I - Zn 2+ + 2e -  Zn from table “flip” I 2 + Zn  2I - + Zn 2+

4. eg: Write the reaction that takes place between Cu 2+ and Hg. Hg 2+ + 2e -  Hg from table Cu 2+ + 2e -  Cu Here we see that Hg 2+ wants to hold on to electrons more than Cu 2+ wants to take them away and no reaction takes place.

5. eg: Why is it not safe to store a solution of CuCl 2 in an iron container? What kind of metal container could be used instead? Cu 2+ will react with Fe (s). Fe (s) + Cu 2+ (aq)  Cu (s) + Fe 2+ (aq) Better metals to use would be silver or gold. Au and Ag will not be oxidized by Cu 2+.