(a) Complete the calculation of the standard reduction potentials for each half- cell in Section C. Then, arrange the standard reduction potentials and corresponding half-reactions in order, starting with the most positive at the top of the table and ending with the most negative at the bottom. By doing this you are making a table of reduction potentials to compare with your textbook. NOTE: All solutions are 1.0 M, except for Ag, which is 0.1 M. You should use the Nernst equation, -log1oQ(at 25 C) 0.0592V mol = E n "cell cell to determine Ecell for the cell involving Ag+. Include literature values (your textbook would be useful) for the standard reduction potentials in your table. Reference your source! (b) Which is the strongest oxidizing agent in the table? Which is the strongest reducing agent in the table? (c) Why is it important that the precious metals Au, Ag, and Pt are poor reducing agents? What would happen to them if they weren't? don't get oxidized easily - world rest (d) Write a balanced redox equation for the spontaneous reaction between the Cu/ Ag half-cells.