posted by Katherine on .
Consider two cells, the first with Al and Ag electrodes, and the second with Zn and Ni electrodes, each in appropriate 1.00M solutions of their ions.
If 2.10g of metal is plated in the voltaic cell, how much metal is plated in the electrolytic cell?
Another plated question... I'm sorry, normally I include an explanation of what I'm confused on but I just do not know how to approach this problem at all. I know that when Ag and Ni get plated they will have a 2.98 V and when Ag and Zn get played they have a 1.98 V
Reduction Half Reaction & Standard Potential Ered°
Al3+(aq) + 3e– → Al(s) –1.66 V
Ag+(aq) + e– → Ag(s) +0.80
Zn2+(aq) + 2e– → Zn(s) –0.76
Ni2+(aq) + 2e– → Ni(s) –0.26
Voltaic cell is
Al + 3Ag^+ ==> Al^3+ + 3Ag(s)
2.1g is how many mols of Ag?
2.1/107.86 = 0.0195 and that will be deposited by 0.0195*96,485 coulombs = about 1878 but you can be more accurate than that.
When in the electrolytic cell, the Al is being plated and Ag is going into soln. The Al plated will be 26.98/3 = about 9 g with one(1) coulomb; therefore,
9g x (1878/95,485) = about 0.175g Al plated. You can arrive at the same answer by taking a shortcut.
The equivalent weight of Al is about 9, that of Ag is about 108; therefore,
2.1 x (9/108) = about 0.175)
Using this method i did not arrive at the right answer for the problem
the steps are right but its not al its zn.
a)Plated Ni and Ag at 2.98 V
b)Plated Ag and Zn at 1.98 V
this is proven
STILL DOESNT WORK
b) Ag & Zn, 0.80 - (-0.76) = 1.56V
Can someone explain how to calculate part a and b?