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Topic: [SOLVED] Redox Reactions > Primary Cells > Silver Oxide Cell  (Read 10753 times)

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Offline Enit

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[SOLVED] Redox Reactions > Primary Cells > Silver Oxide Cell
« on: April 06, 2008, 12:54:50 AM »
I have a question that I just can't seem to.. understand.

In the question we're dealing with silver oxide cells.

The half reactions are:
Ag2O(s) + H2O(l) + 2e-   -->  2Ag(s) + 2OH-(aq)        Eº = +0.34V

Zn(OH)2(s) + 2e- --> Zn(s) + 2OH-(aq)           Eº = -1.25V

Question: As the cell operates, the

A. [OH-(aq)] increases
B. mass of Zn(s) increases
C. mass of Ag2O(s)
D. mass of Zn(OH)2(s) decreases

The answer key says the answer is C.

This is what I have:

I eliminated A because I identified Zn(s) as the anode and reducing agent in this reaction. If we were to write the net equation for this we would have to flip the second equation (zinc) which would end up being Zn(s) + 2OH- --> Zn(OH)2 + 2e-
Then the 2OH-'s would cancel each other out and so their concentration wouldn't increase.

Then I eliminated B because Zn(s) is our anode and reducing agent. The reducing agent undergoes oxidation and electrons are lost -> mass of anode would decrease. I'm actually not too sure about that.. the whole mass increasing and decreasing thing.. is that related to the loss and gain of electrons at all? Gain of electrons -> mass of cathode increase or is it.. the oxidizing agent? Arghs.. I thought I understood this stuff but it seems like I'm confusing myself more and more!

Oh goodness! Never mind! I have figured it out. I sorted out the concepts and yes.. this question was not as difficult as it seemed.
« Last Edit: April 06, 2008, 02:21:12 AM by Enit »

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