Week of May 13th Chemistry Problem: Analyzing a Galvanic Cell

One of the most potentially-intimidating questions you can encounter on the SAT Chemistry Subject Test is one which requires the analysis of a galvanic cell. For example, suppose you were given the following diagram for a galvanic cell and asked to provide analysis:

Where would we even begin?

We would start by noting that in a galvanic cell, a spontaneous reaction is observed (in this case, powering the light bulb). We recall that a galvanic cell always has an anode (where oxidation occurs) and a cathode (where reduction occurs). In this system, we have 4 metallic components: Zn(s), Zn(2+, aq), Cu(s), and Cu(2+, aq). As metal atoms typically are redox active, and since 2-electron oxidized species for each metal are present in solution, we suspect these are the atoms involved in the redox process. On any exam, we would be given the following reduction potentials:

Zn(2+, aq) + 2e-  → Zn(s)          E° = -0.76 V

Cu(2+, aq) + 2e-  → Cu(s)           E° = 0.34 V

and recalling that since a positive potential indicates a spontaneous redox process (which is what a galvanic cell is known for), then it must be the case that Zn should be oxidized at the anode and Cu(2+) should be reduced at the cathode. Therefore, the overall spontaneous process, with a potnetial difference of +1.10 V, should be

Zn(s) + Cu(2+, aq)  →  Zn(2+, aq) + Cu(s)  E° = 0.34 V + -(-0.76 V) = 1.10 V

The following diagram gives us a great visual of what occurs in this galvanic cell:

Keep in mind the salt bridge plays an important role here too, providing the ions necessary to balance excess positive or negative charge as Zn(2+) is formed and Cu(2+) is reduced. For a video on galvanic cells see: http://www.youtube.com/watch?v=0oSqPDD2rMA