posted by stumped .
Two voltaic cells are to be joined so that one will run the other as an electrolytic cell.
In the first cell, one half-cell has Au foil in 1.00M Au(NO3)3, and the other half-cell has a Cr bar in 1.00M Cr(NO3)3.
In the second cell, one half-cell has a Co bar in 1.00M Co(NO3)2, and the other half-cell has a Zn bar in 1.00M Zn(NO3)2.
A) Calculate the E0cell for each cell.
[2.24V and 0.48V]
B) Calculate the total potential if the two cells are connected as voltaic cells in series.
C) When the electrode wires are switched in one of the cells, which cell will run as the voltaic cell and which as the electrolytic cell.
[voltaic=cell 1, electrolytic=cell 2]
D) Which metal ion is being reduced in each cell?
E) If 2.00 g of metal plates out in the voltaic cell, how much metal ion plates out in the eletrolytic cell.
I undertstand the answers until part C-E. For C, I have the right answer but my logic was based on the spontaneity of Au(s) & Cr reaction, and then with Co(s) with Zn. Is this right or do I have to take the series into account (which I don't know how).
D) this is based on C, but since Au(s) turns into its ionic form, shouldn't it be oxidized?
E) I tried to do a molar ratio (find out the number of moles Au (2.00g/197.0g) then multiply that with Zn molar mass (this is based on the answer for part D) but I still didn't get the answer.
Please explain this since final is approaching.
Chemistry again -