Chemistry again

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.
[2.72V]
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?
[Au, Zn]
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.
Thanks.

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