17.2 Galvanic Cells

Write cell schematics for the following cell reactions, using platinum as an inert electrode as needed.

(a) $\text{Mg} \left(s\right) + \text{Ni}^{2+} \left(a q\right) \rightarrow \text{Mg}^{2+} \left(a q\right) + \text{Ni} \left(s\right)$

(b) $2 \text{Ag}^{+} \left(a q\right) + \text{Cu} \left(s\right) \rightarrow \text{Cu}^{2+} \left(a q\right) + \text{2Ag} \left(s\right)$

(c) $\text{Mn} \left(s\right) + \text{Sn}(\text{NO}_{3} \left.\right)_{2} \left(a q\right) \rightarrow \text{Mn}(\text{NO}_{3} \left.\right)_{2} \left(a q\right) + \text{Sn} \left(s\right)$

(d) $3 \text{CuNO}_{3} \left(a q\right) + \text{Au}(\text{NO}_{3} \left.\right)_{3} \left(a q\right) \rightarrow \text{3Cu}(\text{NO}_{3} \left.\right)_{2} \left(a q\right) + \text{Au} \left(s\right)$

Write a balanced equation for the cell reaction of each cell in the previous exercise.

Identify the oxidant and reductant in each reaction of the previous exercise.

Why is a salt bridge necessary in galvanic cells like the one in the figure?

An active (metal) electrode was found to lose mass as the oxidation-reduction reaction was allowed to proceed. Was the electrode an anode or a cathode? Explain.