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We know that very pure water does not conduct electricity, but salt water is a decent conductor. This is commonly explained by saying that "the ions carry the current through the solution", an explanation that does not really make sense because it is not clear what will happen when all of the ions have migrated to the electrodes.

Better explanations of conduction through a salty solution (like this one or this one) explain the conduction of electricity in terms of a reduction reaction that takes place at the anode and an oxidation reaction that takes place at the cathode. In the case of salt water, chlorine gas (Cl2) is formed at the anode and hydrogen gas (H2) is formed at the cathode.

This explanation makes sense to me, but it implies that the conduction of electricity through a solution is fundamentally different than the conduction of electricity through a wire. A copper wire does not change at all after we pass electricity through it. In contrast, in salt water, we are driving a two chemical reactions (one at each electrode), which fundamentally change the composition material.

This implies that it is not possible for a aqueous solution to conduct electricity forever. Since we are driving a chemical reaction, we are either consuming our salt by forming gas (or plating it onto the electrodes) or we are consuming the water by forming H2 or O2 gas.

This is surprising to me (for some reason). So, I'm asking if my thinking is correct: is it possible for a salt solution to conduct electricity forever, or it will it always eventually consume the reactants and stop as I have surmised?

EDIT: See this discussion of the Chemistry SE: https://chemistry.stackexchange.com/questions/7571/can-an-aqueous-solution-conduct-electricity-forever/7610#7610

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closed as off-topic by John Rennie, Waffle's Crazy Peanut, user10851, Brandon Enright, jinawee Dec 21 '13 at 21:46

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  • $\begingroup$ The anode starts to dissolve and puts ions of the anode in the water. I think this will replace any ions that leave the water on the cathode. $\endgroup$ – Brandon Enright Dec 21 '13 at 0:16
  • $\begingroup$ Ah, so there might be a situation where the solution can stay the same because all of the ions are being replaced by ions from the electrodes. But, unless we have infinitely large electrodes, we still cannot go forever. $\endgroup$ – DanHickstein Dec 21 '13 at 0:29
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    $\begingroup$ This question appears to be off-topic because it belongs on chemistry.stackexchange.com $\endgroup$ – John Rennie Dec 21 '13 at 10:05
  • $\begingroup$ Okay, seems like this is more of a chemistry question. Should I delete it and repost to the chemistry SE? $\endgroup$ – DanHickstein Dec 25 '13 at 15:41
  • $\begingroup$ At least part of the answer involves the fact that in addition to making O2 and H2, electrodes produce H+ & OH-. This is not always shown in electrolysis reactions equations, but the ions are there and are detectable with pH indicator. See this question in chemistry stack exchange that addresses the same issues: chemistry.stackexchange.com/questions/44532/… $\endgroup$ – lamplamp May 27 at 20:38
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I believe the answer is no simply because the electric current breaks the $\mathrm{H_2 O}$ bonds and forms $\mathrm{H_2}$ and $\mathrm{O_2}$ gas. Eventually you run out of water due to this process.

Also, in a $\mathrm{H_2 O}$ + $\mathrm{NaCl}$ solution, $\mathrm{Cl_2}$ gas is produced which would eventually affect ion concentrations.

Wikipedia has a pretty detailed description of the electrolysis of water.

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  • $\begingroup$ Yup, this makes sense to me. Good to hear that I'm not totally crazy :). $\endgroup$ – DanHickstein Dec 21 '13 at 0:31
  • $\begingroup$ The Wikipedia article does mention "...the theoretical and real observed threshold of electrolysis to (-)$1.48\: \mathrm{V}$". The reaction doesn't happen below this voltage. I'm pretty sure that means water can't conduct below $1.48\: \mathrm{V}$ but if it can, you could probably have it conduct forever without water loss below this voltage. $\endgroup$ – Brandon Enright Dec 21 '13 at 0:36
  • $\begingroup$ Interesting, but without the chemical reaction, I'm not clear how the electricity would conduct through the solution. This then returns to the overarching question of how electricity conducts through a solution. I'm basically saying that it requires a chemical reaction and wondering if that's accurate. $\endgroup$ – DanHickstein Dec 21 '13 at 0:40
  • $\begingroup$ @DanHickstein yeah I'm pretty sure ion transfer is needed for conduction and that below the threshold voltage it just won't conduct. You might want to ask about this specific point on Chemistry.SE. $\endgroup$ – Brandon Enright Dec 21 '13 at 0:42
  • $\begingroup$ Does this apply to A/C as well? Theoretically the water molecules should be able to reform above these two electrodes and return to the conductor. Also, I'm not sure if chlorine gas would escape from around a sodium-coated electrode rather than reacting with the sodium. (I do understand that the waveform will go haywire because of the threshold voltage) $\endgroup$ – John Dvorak Dec 21 '13 at 8:50

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