Problem 605
A voltaic cell is made up of a silver electrode in a \(1 \mathrm{M}\) silver nitrate solution and an aluminum electrode in a \(1 \mathrm{M}\) aluminum nitrate solution. The half reactions are 1) \(\mathrm{Ag}^{+}+\mathrm{e}^{-} \rightarrow \mathrm{Ag}\) (s) \(\mathrm{E}^{\circ}=.80\) volt, and (2) \(\mathrm{Al}^{3+}+3 \mathrm{e}^{-} \rightarrow \mathrm{Al}(\mathrm{s})\) \(E^{\circ}=-1.66\) volt, calculate the standard free energy change, \(\Delta \mathrm{G}^{\circ}\), in \(\mathrm{Kcal} /\) mole. \(1 \mathrm{cal}=4.184\) joule \(1 \mathrm{~F}=96,500\) coul, \(1 \mathrm{Kcal}=1000 \mathrm{cal}\)
Problem 5601
\(.0324\) Faradays (F) liberated \(.651 \mathrm{~g}\) of Calcium. What is the atomic weight of Calcium?
Problem 5602
Two electrolytic cells were placed in series. One was composed of \(\mathrm{AgNO}_{3}\) and the other of \(\mathrm{CuSO}_{4}\). Electricity was passed through the cells until \(1.273 \mathrm{~g}\) of \(\mathrm{Ag}\) had been deposited. How much copper was deposited at the same time?
Problem 5603
The same quantity of electricity was passed through two separate electrolytic cells. The first of these contained a solution of copper sulfate \(\left(\mathrm{CuSO}_{4}\right)\) and exhibited the following cathode reaction (reduction): \(\mathrm{Cu}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Cu}(\mathrm{s})\) The second of these contained a solution of silver nitrate \(\left(\mathrm{AgNO}_{3}\right)\) and exhibited the following cathode reaction: \(\mathrm{Ag}^{+}+\mathrm{e}^{-} \rightarrow \mathrm{Ag}(\mathrm{s})\) If \(3.18 \mathrm{~g}\) of Cu were deposited in the first cell, how much (Ag) was deposited in the second cell?
