# Using electrolysis of water as a proton generator [closed]

I'm working on a project right now, where the challenge is to try to use electrolysis as a way of controlling the $$\ce{pH}$$ of water.

I've set up an experiment where I have two separate chambers or volumes of water. Both chambers are connected to the same electrolytic unit. The flow of water is from a respective chamber, into the unit and then back into the same chamber again, in an attempt to separate the two half reactions. For this reason I call the chamber that flows over the anode for the anodic chamber, and the chamber that flows over the cathode for the cathodic chamber. The bodies of water are separated in the electrolysis cell by a proton-exchange-membrane (PEM) that should, in theory, only allow positively charged ions to cross.

Apologies for the Danish text on my setup, but the figures should help you understand what I'm trying to do.

The anodic chamber (the water that goes to the anode half-reaction of the electrolysis cell) contains tap water while the cathode chamber contains Milli-Q water with added $$\ce{NaCl}$$ for conductivity purposes.

The $$\ce{pH}$$ of the water is alkaline in both chambers when the experiment starts, but I've also tried to run the experiment where the anodic chamber is acidic.

Anyway, here is what I hoped would happen in alkaline water:

Anode (oxidation):

$$\ce{4OH- (aq) -> 4e- + O2(g) + 2 H2O(l)}$$

Cathode (reduction):

$$\ce{2H2O (l) +2e- -> H2(g) + 2OH-(aq)}$$

In theory, I was expecting the $$\ce{pH}$$ to drop in the anodic chamber, since removing $$\ce{OH-}$$ would necessitate a flow of positively charged ions across the PEM membrane, instead of the $$\ce{OH-}$$ produced in the cathodic chamber in order to satisfy electronegativity. This would also mean that the $$\ce{pH}$$ should rise in the cathodic chamber.

However, what I see is invariably that the $$\ce{pH}$$ rises very rapidly, and continuously, in the cathodic chamber, whereas the $$\ce{pH}$$ in the anodic chamber remains either stagnant or rises very slowly.

Now, part of the reason for the rapid rise in the cathodic chamber is the complete lack of alkalinity, that much is clear to me. However, I don't understand why only the reduction half cell reaction seems to occur.

The only explanation I have been able to come up with is that something else on the anode half-reaction is donating electrons to the reaction. I tried running the experiment at different voltages in the hopes of finding a "sweet spot" that would prevent this "something" from reacting but as of now I haven't been able to do that. I'm worried it may be that my anodic electrode is donating the electrons. My electrodes are nickel based.

A cursory glance at Wikipedia shows the standard electrode potential of nickel as

$$\ce{Ni^2+ + 2 e− <=> Ni(s)} \;\;\; \pu{E^{\circ} = −0.25 V}$$

And for my hydroxides

$$\ce{O2(g) + 2H2O + 4 e− <=> 4 OH−(aq)} \;\;\; \pu{E^{\circ} = +0.401 V}$$

When I run the reaction, I also see a yellowish/greenish precipitate forming in the anodic chamber, which according to some lab technicians at my location may be the nickel reacting.

Anyway, the question(s) are then - what do you think is happening in my system, given this information?

Am I on the right track? should I try to change my electrode?

Have I made some obvious, basic mistakes?

What suggestions do you have, so that I can get this to work?

I hope that you have the time to advice me on this, and I wish you all a pleasant day.