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I get that galvanic cells require a salt bridge to maintain neutrality so the cathode doesn't become saturated with electrons, but why is a two-cell setup required? Wouldn't the spontaneous reaction occur and neutral conditions be met if Cu and Zn electrodes were placed in a solution of NaCl?

And vice versa, couldn't an electrolytic cell work just fine with a salt bridge?

This website (and many others) leads me to believe that galvanic cells must have a salt bridge and electrolytic cells must be a single cell. My intuition says that's not right though. Is it because you don't want undesirable compounds reacting at the electrodes? I'd think you'd want the same for both EC cell types if that was the case and electrolytic cells would use salt bridges as well (eg proton exchange membranes). http://chem.libretexts.org/Core/Analytical_Chemistry/Electrochemistry/Electrolytic_Cells

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    $\begingroup$ Sure, galvanic cell would work for a while in a single cell setup, but you'll get a lot of side reactions which you don't want. On the other hand, electrolytic cell would work just fine in a two-cell setup with a salt bridge, but then half of your solution would be effectively excluded from the reaction, which we didn't ask for. Also, welcome to Chem.SE. $\endgroup$ – Ivan Neretin Jan 12 '17 at 6:17
  • $\begingroup$ Thanks. Wouldn't separating the cells in an electrolytic cell be more desirable than galvanic to reduce the back (spontaneous) reaction of by-products? For example, in electrolytic water splitting, the chamber are separated with a proton exchange membrane to reduce back reaction of hydrogen and oxygen to water. Why do introductory texts then claim electrolytic and galvanic cells must be 1 and 2 celled, respectively? Or are they just simplifying it? $\endgroup$ – prof.kvothe Jan 12 '17 at 15:35
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    $\begingroup$ That's it: they are just simplifying. $\endgroup$ – Ivan Neretin Jan 12 '17 at 15:44
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Perhaps this isn't the full answer, but it probably plays a role. It is critical for galvanic cells that there not be any electronic conductivity through the cell between the two electrodes. If there is a path for electrons, they will take it which means the cell will spontaneously discharge and not do any useful work through the external circuit. Pure water is a very poor conductor but becomes conductive once ions are dissolved.

By comparison, the construction of a lithium ion cell is close to what you describe for a single chamber galvanic cell. There is a separator involved but mostly to keep the two electrodes from coming into direct contact. One reason this is possible is that the electrolyte is an organic solvent (usually ethylene carbonate, dimethyl carbonate, etc) and doesn't conduct electrons, even when the appropriate salt (eg. $\ce{LiPF6}$) is dissolved in it. There is no salt bridge in this cell, though it is clearly galvanic.

Regarding the electrolytic cell, I don't see any particular reason that you couldn't build one with a salt bridge, provided the salt-bridge let you balance the charges generated on each side. Maybe they're all presented as one beaker just because it's easier to make and there's not really any advantage to a salt bridge in this case.

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  • $\begingroup$ What would stop electrons moving through an electrically conductive salt bridge then? For example, fuel cells and batteries don't really use a "salt bridge" as we see it drawn often, they use an ionic membrane. Can't electrons pass through that? But even then, why would they want to - aren't they "forced" to go through the electrode via the potential difference? $\endgroup$ – prof.kvothe Jan 12 '17 at 15:32
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    $\begingroup$ Nothing would stop them, that's why salt-bridges aren't electron conductive. I specifically mentioned why Li-ion battery cells don't have a salt bridge, and I think the same applies for fuel cells: the electrolyte is chosen to be non-electronically-conductive. In a galvanic cell, the potential difference is because the electrons can't move between the electrodes internally. In the extreme, imagine a wire connecting the two sides through the electrolyte: the cell would immediately self-discharge and be useless. $\endgroup$ – m3wolf Jan 12 '17 at 16:32

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