How Do Electrolytes Conduct Charge Without Redox Reactions?

[mic*]

First up, what is the prefix thing about? I picked "B" but I have no idea whether this is correct.

My question is pretty simple. How does an electrolyte conduct charge below the potential threshold of any redox reactions (ie the threshold for electron transfer)?

Is it purely due to EDLE? If so, why is the conductivity sustainable for extended periods and does not diminish as charged ions accumulate around the respective electrode?

 

[mjc123]

Question makes no sense. Have you omitted to post something? (You could always type the question out.)

 

[mic*]

If I use two platinum rods as electrodes and put them in say 1M NaCl(aq), then I apply a potential of 0.5V across the electrodes, a current will flow.

This is not enough potential for any possible pairs of redox reactions where actual electrons are exchanged at the electrode-electrolyte interfaces.

If the current flow is due to ion mobility and the work done by the emf at the electrode to move the ions in solution (such as with the Electric Double Layer Effect - EDLE), then why does current continue to flow after periods of hours, days, weeks etc?

Wouldn't it be reasonable to expect an equilibrium of sorts where the electrodes are "charged"?

Wouldn't it be reasonable to expect the conductivity of the system to vary with time, as it does with typical capacitive systems when they are charged and discharged?

 

[BvU]

 

[mic*]

Thanks BvU. Maybe an Android thing. It doesn't display like that on my device.

 

[mic*]

I would say it should be I or A. Not B.

 

[DrDu]

If I use two platinum rods as electrodes and put them in say 1M NaCl(aq), then I apply a potential of 0.5V across the electrodes, a current will flow.

This is not enough potential for any possible pairs of redox reactions where actual electrons are exchanged at the electrode-electrolyte interfaces.

Are you sure that the effect really happens with pure solution, free of gases like oxygen?
Are you sure, that there is really not enough potential for any possible pair of redox reactions?
How about ##\mathrm{Cl^-+H_2O \rightarrow ClO^{-} +2H^+ +2e^-}## and ##2 \mathrm{H^++2e^-\rightarrow H_2}##?

 

[vanhees71]

I don't understand the question. What have chemical reaction to do with the conductivity of a solution of salt? It's conducting electric currents, because it contains the ions of the salt in the solution (here ##\mathrm{Na}^+## and ##\mathrm{Cl}^-##), free to move in the applied electric field. For a quite detailed theory, see

https://en.wikipedia.org/wiki/Debye–Hückel_theory

 

[DrDu]

I don't understand the question. What have chemical reaction to do with the conductivity of a solution of salt? It's conducting electric currents, because it contains the ions of the salt in the solution (here ##\mathrm{Na}^+## and ##\mathrm{Cl}^-##), free to move in the applied electric field. For a quite detailed theory, see

https://en.wikipedia.org/wiki/Debye–Hückel_theory

The point is that the ions which travel to the electrodes have to be oxidized/reduced. If not, charge would accumulate until the external potential is compensated, i.e. exactly what happens in an electrolyte condensator.

 

[mic*]

DrDu, I haven't performed the experiment checking the concentration of dissolved oxygen or in a controlled gas environment. Seems a bit extreme, but maybe that's what would need to be done.

I'd considered the base view of that redox pair and the potential difference is 0.9V... (it's kind of why I picked a low potential like 0.5V)

I take from your answer that you feel confident it is just a chemistry problem and it is a case of investigating what reactions are occurring? Ie I am not missing something obvious in my theoretical approach.

 

[DrDu]

I think it is a chemical problem, yes. Especially, I suspect that dissolved oxygen could react directly at the cathode, like ## \mathrm{4H^++O_2 +4e^-\rightarrow 2 H_2O}##.

 

[mic*]

Vanhees71, thank you for your post. I definitely think Debye-Huckel theory is related.

I am not an expert, so please correct me if I am wrong, but the part that i believe nullifies its specific relevance here is "When conductivity is measured the system is subject to an oscillating external field due to the application of an AC voltage to electrodes immersed in the solution." (From wiki page).

To be fair, I did not explicitly state that I was using a DC potential.

I feel like ionic motion will nessecitate *some* continued current to "hold the electrodes charged". I'm just not sure where it is predicted in models like the Gouy-Chapman approximation of Helmholtz EDLE, or of there is superceeding model or theory I am missing.

 

[vanhees71]

Sure, if you use DC, you operate the whole thing as a capacitor ;-).

 

[mic*]

DrDu, the reduction of O2 that you just mentioned, paired with the oxidation of chlorine to hypochlorite that you previously mentioned, progresses naturally with a potential difference of 0.33V.

I think that might be the one I'm looking for! Or at least one of them Thank you for your provocation haha.

I hadn't really considered the scenario as follows being significant - the above pair progresses naturally and quickly towards dissolved O2 and Cl- being consumed and turned into hypochlorite when NaCl is added to H2O. Then, when a potential above 0.33V is applied, the reactions are reversed at the respective electrode.

Definitely need to check DO.

 

[mic*]

Yes vanhees71, exactly, and then you arrive at my question when you find an anomalously high current draw at a stable voltage.

 

[vanhees71]

Sure, I obviously didn't understand what's behind your question :-(.

 

[DrDu]

It should also work with chloride being oxidized to chlorine instead of hypochlorite, but I don't have the numbers here at the moment.

 

[mic*]

Nope. Reduction potential of chlorine is greater than that of O2. And Cl2 gas is not available by any other route to oxidise (the oxygen in) H2O.

 

[DrDu]

Nope. Reduction potential of chlorine is greater than that of O2. And Cl2 gas is not available by any other route to oxidise (the oxygen in) H2O.

I meant:
##\mathrm{Cl_2+2e^- \rightarrow 2Cl^-}##, U=1.36 V at the anode and
and ##\mathrm{4H^++O_2 +4e^-\rightarrow 2 H_2O}## U= 1.229 V, at the cathode,
so that ##\Delta U=0.13 V## (which still depends on pH, but rather little).

 

[mic*]

Okay, that reaction is plausible by application of an external potential >0.13V, but there is no coulombic attraction between O2 and the cathode. Wouldn't that make it... um... slow I guess.

 

[mic*]

And possibly almost unoticable...? To try and support why I see no gas formation at the anode... Kinda need that bit.

If you got two Cl2 molecules at the anode for every O2 that happened to bump into the cathode, maybe these Cl2's just stay in solution and then proceed to spontaneously oxidise H2O back into the original reactants.

No visible gas formation at anode. Plausible? I am not a chemist... :(

 

[DrDu]

Chlorine is quite well soluble in water. I also don't expect large currents, but you where asking for a possible mechanism why there are any currents after all. You could add some drops of phenolphthalein or some other indicator to see whether the cathode gets basic. If chlorine or hypochlorite are formed at the anode, they will probably bleach out some colorants.

 

[mic*]

That is a great, sensible idea. With the right pH indicator, would it be reasonable to also expect to see indications of acidity at the anode as H+ accumulates?

 

[DrDu]

Yes