mycotheologist
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The electrolysis experiment thread
I decided to start satisfying my fascination with electrochemistry by performing electrolysis on all sorts of different compounds so I'll post all my
findings here. I just started 5 minutes ago so give it time to grow. First though, heres the electrolysis setup I'm using: A canning jar with two
thick graphite electrodes blue attached to the inner walls of the jar with blue tack (I was previously putting the electrodes through two holes in the
jars lid but I later concluded that the metal lid was short circuiting the device). Each end of the wire from the power source is stripped down and
the copper is coiled around the graphite rods and blue tacked in place. Alright so heres the experiment:
Sodium bicarbonate
As soon as I plugged in the power source, big bubbles started forming on one of the electrodes. These bubbles are way bigger than the Cl2
bubbles I've seen in your average NaCl cell. These bubbles can't really be anything other than CO2. Nothing astonishing there, it was
pretty obvious that the bicarbonate would turn into CO2 at the cathode. Not sure if this would be a practical method of generating CO2 or
not, it was like what you see when you dissolve an Alka-Seltzer in water but it was confined to the cathode.
Citric acid
Small bubbles (far smaller than the CO2 bubbles from the previous test) began to form at the same electrode that the CO2 bubbles
formed during the previous test. I'm confused. If these bubbles are H2 then they should be forming at the opposite electrode. I tried it
on sodium bicarbonate again and noticed that there were bubbles forming on the other electrode too but it was very mild in comparison. You could
literally hear the sizzling from the bubbling in the bicarbonate cell but the citric acid cell was silent the whole time.
Potassium bromide
This ones far more interesting. There are a lot of bubbles forming on one of the electrodes but its the other electrode that is turning brown. I see
now that its the hydroxide formation that was causing the bubbles in the previous experiments. I wonder why the bicarbonate cell bubbled so much more
vigorously than the others. Maybe it was more saturated.
Potassium permanganate
I couldn't see the electrodes because the solution was too purple but pink foam eventually started forming on top of the solution. When I washed out
the jar, there was some manganese dioxide at the bottom of the jar. I'm guessing the foam at the top of the solution was oxygen expelled during the
reduction of the permanganate ion to manganese dioxide.
Boric acid
Nothing happened at all as far as I could see. I had to clean the electrodes and all that crap after the permanganate experiment so maybe I
accidentally disconnected a wire. The boric acid wasn't very soluble in the water at all though so that could also have had something to do with it. I
read somewhere that electrolysis of boric acid yields elemental boron, I'll have to look into that further.
Sodium chloride
I just did this because I want to make some hypochlorite. I made an interesting observation though. The anode (the one where the Na+)
deposites was fizzing/bubbling like mad. I've never seen it react that vigorously before. In another thread, I said that I hooked my power supply up
to the graphite electrodes but barely any current flowed. Only thing I changed now is that I'm no longer using a metal lid for the jar to hold the
electrodes in place. The metal lid must have short circuited the device so instead of the electricity flowing through the NaCl solution, it flowed
through the metal jar lid lol. I've had this cell running for a few hours now, I just checked on it and noticed one of the electrodes is still
bubbling vigorously. Can't remember if its the anode of the cathode though. I think its the anode cuz the sodium reacts violently with water but at
the cathode, the chlorine formation should be a pretty mild reaction. Thats just my opinion though, correct me if I'm wrong.
I'm wondering what to do next. The ones I've done so far aren't very interesting though, I wanna make this thread more interesting. I wish I had a
camera so I could add pictures of the results to the thread. The KBr one was pretty cool, the cathode started turning brown/orange. I wanna do
something that will produce a coloured gas or change the color of the water or something.
[Edited on 9-4-2012 by mycotheologist]
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elementcollector1
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When I ran the bicarbonate single-cell for a while, it made a brownish-red solution and carbon particles were observed decaying off one of the
electrodes (I don't know which). What is this dark, soluted product?
Is there a way to get NaOH (or KOH) from this route?
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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ManBearSwine
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@mycotheologist
Electrolysis of sodium bromide produces hydrogen at the cathode and elemental bromine at the anode, which partially dissolves in water to form a
yellow solution. It reacts with the hydroxide produced at the cathode to form colorless bromate ions.
Citric acid is a weak acid, so there are very few ions to conduct the current. This is why you saw only slight bubbling.
As far as I know, sodium bicarbonate is a relatively inert electrolyte. Carbon dioxide will not be produced by electrolysis alone; the only mechanism
I can think of would be a reaction between the hydronium ions at the anode and the bicarbonate ions in solution to form carbonic acid, which
decomposes to carbon dioxide and water. This seems unlikely, as hydroxide ions are also drawn to the anode, and would prevent the formation of carbon
dioxide. Did you test any of the gases being evolved, or did you just assume that it was carbon dioxide?
@elementcollecteor1
I haven't had any problems with my graphite electrodes corroding, leading me to believe that yours are either impure or have things deposited on them
(metals or metal oxides) from other experiments.
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nora_summers
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i'm still new here, so i'm not sure, but i think electrolysis and related electrochemistry is something meant for the "technochemistry" section of
this forum rather than general chemistry.
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CrEaTiVePyroScience
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Well I have made an (I think) interesting electrolysis I have posted a video about it on my youtube channel.
But you do need a gold / platinum electrode for that.
https://www.youtube.com/watch?v=XSXRG0JbWTI (making sulfuric acid from copper sulfate)
In this video you can see a very nice deposit of copper on the electrode
I also got just a more basic electrolysis of iron oxide which is very basic and used in pyrotechnics its also on my channel.
Over few days I will post another video of another electrolysis of potassium chlorate, made from potassium chloride which is availble as salt that
contains no sodium and is used often for diets almost all grocery stores have them!
[Edited on 14-4-2012 by CrEaTiVePyroScience]
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bbartlog
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I think you might benefit from a little more theory.
| Quote: | | The anode (the one where the Na+) deposites was fizzing/bubbling like mad. |
Metals plate out at the cathode, generally (things get reduced at the cathode and oxidized at the anode).
| Quote: | | Nothing astonishing there, it was pretty obvious that the bicarbonate would turn into CO2 at the cathode. |
Not obvious at all (actually it would be astounding). You should get hydrogen at the cathode and oxygen at the anode in such a cell, *possibly* with
small amounts of CO2 if the current heated your electrolyte enough to drive it out of solution (but that would not be an electrolytic effect). CO2
evolution would stop fairly quickly in any case since the electrolyte would become more basic over time.
| Quote: | | The boric acid wasn't very soluble in the water at all though so that could also have had something to do with it. |
Boric acid should be a lousy electrolyte because it has a pKa of 9.2.
The less you bet, the more you lose when you win.
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woelen
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I also have done many electrolysis experiments and some of the most remarkable I put on my website. Some of the experiments require special chemicals,
but most of them can be done with chemicals which are not that difficult to obtain (e.g. pottery stored, ebay). Instead of using a platinum wire as
anode, one can use graphite as well.
http://woelen.homescience.net/science/chem/exps/electrolysis...
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Pyridinium
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Quote: Originally posted by bbartlog  |
Not obvious at all (actually it would be astounding). You should get hydrogen at the cathode and oxygen at the anode in such a cell, *possibly* with
small amounts of CO2 if the current heated your electrolyte enough to drive it out of solution (but that would not be an electrolytic effect). CO2
evolution would stop fairly quickly in any case since the electrolyte would become more basic over time.
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I would add that as oxygen is formed at the anode, there also has to form H+, according to the reaction
2H2O ------> O2 + 4H+ + 4e-
(Edit: meant to say 4e- not 2e-. Charge balance.)
This, again, wouldn't directly be an electrolytic effect, but definitely still a result of the electrolysis anyway. It should, however, cause CO2
evolution to last for essentially the duration of the experiment.
See, I wasn't planning to do any experiments this afternoon, and now look what y'all did to me. 
I am going to have to route some of the putative CO2 into some limewater and see how much precipitate we get. Too bad my Mettler is broken. (I
forgot to ask if Woelen or one of you guys already did this experiment, but now you got me started).
[Edited on 14-4-2012 by Pyridinium]
[Edited on 15-4-2012 by Pyridinium]
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woelen
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If an anionic species is not oxidized during electrolysis at the anode, then the main reaction is oxidation of water to oxygen and free acid
(Pyridinium already gave the correct net equation). This free acid in turn may react further. A nice example is the electrolysis of bicarbonate salts.
The acid reacts with the HCO3(-) ion to water and CO2. So, what is produced is a mix of O2 and CO2 at the anode, when a bicarbonate salt is
electrolysed.
I have done this experiment a long time ago already and even tried to determine the ratio of O2 and CO2 experimentally. I always had too high a value
for oxygen, but later, when I understood things better I could explain this. Oxygen hardly dissolves in cold water, while CO2 does dissolve to some
extent.
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