1. 18.4 Standard Electrode Potential

2. Standard Electrode Potential

3. Standard Electrode Potential (continued)

4. Let’s Try a Practice Problem! Reduction Half ReactionE o (V) Pb 2+ (aq) + 2e -  Pb(s)-0.13 Cr 3+ (aq) + 3e -  Cr(s)-0.73

5. Let’s Try Another! An electrode has a negative electrode potential. Which statement is correct regarding the potential energy of an electron at this electrode? (a)An electron at this electrode has a lower potential energy than it has at a standard hydrogen electrode. (b)An electron at this electrode has a higher potential energy than it has at a standard hydrogen electrode. (c)An electron at this electrode has the same potential energy as it has at a standard hydrogen electrode. (b) An electron at this electrode has a higher potential energy than it has at a standard hydrogen electrode.

6. Predicting the Spontaneous Direction of a Redox Reaction The more negative the E o cell is, the more likely the electrode is to lose electrons. Meaning according to the table 18.1 (on page 873 of your text) which orders standard electrode potentials. The more positive values at the top of the table, the more likely that to be reduced. So as you go further down the table, oxidation occurs more easily. The oxidation reactions proceed in the reverse direction. Any reduction reaction in table 18.1 is spontaneous when paired with the reverse of any of the reactions listed below it on the table.

7. Let’s Try a Couple of Practice Problems! Without calculating the standard cell potentials, predict whether the following redox reactions are spontaneous under standard conditions? (a)Zn(s) + Ni 2+ (aq)  Zn 2+ (aq) + Ni(s) (b)Zn(s) + Ca 2+ (aq)  Zn 2+ (aq) + Ca(s) (a)Since zinc is being oxidized and has a more negative standard electrode potential, the reaction will occur spontaneously. (b)Since zinc is being oxidized but has a larger (more positive) standard cell potential, this reaction will not occur spontaneously. Reduction Half ReactionE o (V) Ni 2+ (aq) + 2e -  Ni(s)-0.23 Zn 2+ (aq) + 2e -  Zn(s)-0.76 Ca 2+ (aq) + 2e -  Ca(s)-2.76

8. Let’s Try Another!!! A solution contains both NaI and NaBr. According to table 18.1 (on pg. 873 of your text), which oxidizing agent could you add to the solution to selectively oxidize I - (aq) but not Br - (aq)? (a)Cl 2 (b)H 2 O 2 (c)CuCl 2 (d)HNO 3

9. Predicting Whether a Metal Will Dissolve in an Acid Metals whose reduction half-reactions are listed below the reduction reaction of H + in table 18.1 dissolve in acids, while those above do not. (an exception in nitric acid) Basically if the metal is more easily oxidized than the acid, it will dissolve in the acid.

10. Let’s Try a Practice Problem! Use table 18.1 to determine which metal will dissolve in HNO 3 but not HCl? (a)Fe (b)Au (c)Ag

11. 18.4 pg. 905 #’s 48, 52,58, and 62(a) Read 18.5-18.6 pgs. 877-885