Darkblade48
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Electrolysis of Copper
I've been experimenting with copper lately (just for a little compound collection I started), specifically using it in an electrolytic cell.
Current setup:
Copper anode and cathode
I've tried a magnesium sulfate electrolyte, and the product I got was cupric hydroxide (expected)
However, when I tried it in a salt (NaCl) electrolyte, the precipitate I got was this brown mass, that, upon drying seems to turn a dark brown (cupric
oxide? Unexpected)
Edit: After drying, the solid is clearly a deep brown colour, though there are small red edges to some of the pieces (aha! perhaps it
is cuprous oxide?)
I imagined that electrolysis of copper in a chloride solution would yield cuprous oxide, but instead, I seem to have obtained cupric (I know that the
cuprous is relatively unstable and can convert to the cupric state, but I thought that was only at elevated temperatures).
Just curious as to what other copper compounds can be made, as I started using the copper hydroxide as a starting point to make various things (copper
acetate, sulphate [unfortunately, I can't seem to find it in stores], copper carbonate, etc etc)
Lastly, is it possible to actually crystallize a solution of CuCl2? The last time I tried (from addition of HCl to cupric hydroxide), I
tried to boil away the water, but it left me with a brown/green mess (I figure it's some copper-chloro complex, but maybe someone has a better
idea?)
[Edited on 11-10-2005 by Darkblade48]
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woelen
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Electrolyses of NaCl (or any solution with chloride or bromide in it) with copper electrodes gives an increadibly complex mix. Copper forms all kinds
of complexes, it even forms mixed valence compounds (which are VERY dark brown) and it forms many oxo-chloride species. Copper chemistry, if studied
in more detail, has appeared to me as one of the most complicated transition metal chemistries.
The brown stuff you get most likely is impure copper (I) oxide, contaminated with some copper (II) species. Pure copper (II) oxide is black or very
dark brown and pure copper (I) oxide can have any color between bright yellow and brick red, through all shades of orange.
Pure CuCl2 (anhydrous) is brown, sometimes a little like mustard. It is very hygroscopic though. Pure CuCl2.2H2O has a beautiful nice cyan color,
quite different from CuSO4.5H2O.
Recently, just out of curiousity, I also did the experiment of heating a solution of CuO in HCl and I obtained a brown/yellow mix, with some green
stuff in it. It looked ugly. However, when I added water to it, it dissolved completely and the solution was nice clear and cyan/blue, so apparently
the stuff I made was not that bad.
I think that the ugly color and structure also are due to remnants of HCl in the mix and due to partial loss of water. I have found it really hard to
obtain nice cyan hydrated copper (II) chloride. The brown stuff, however, can be obtained fairly easily by careful heating from HCl.
Here are some pictures I made of some of my (purchased) samples of CuO and CuCl2.2H2O:
http://woelen.scheikunde.net/science/chem/compounds/cupric_c...
http://woelen.scheikunde.net/science/chem/compounds/cupric_o...
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12AX7
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You may've also had copper metal sponge deposit. Sometimes when I electrolyze Cu in NaCl, I get a mushy mess of both that's just ugly... it
has a brownish color. Obviously this sponge hangs off the cathode, until broken up.
The only Cu2O I ever get from the process is solid orange, much brighter than rust (Fe(OH)3 specifically, which tends to be drab orange-brown). From
non-complexing electrolytes I get a low density sky blue precipitate, Cu(OH)2, probably formed by the metal first oxidizing to a microscopic layer of
Cu2O, which being a semiconductor is able to be oxidized further to the +2 state.
The last time I crystallized a CuCl2(aq) solution, I first made sure it was thoroughly green (by adding an oxidizer and if needed, water, but not to
the point of blue color) and only green before drying. Then I let it mostly dehydrate in a warm, dry place. This was green while wet; on further
drying it left a bluish solid which I now have in storage. Despite a varying color (I suspect either Cu(I) impurity or, more likely, oxychlorides) it
appears to dissolve with no insolubles.
Tim
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Darkblade48
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| Quote: | Originally posted by 12AX7
The only Cu2O I ever get from the process is solid orange, much brighter than rust (Fe(OH)3 specifically, which tends to be drab orange-brown).
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How would you go about making Cu2O though? I'd imagine that a solution of NaCl might have worked, but apparently, it
doesn't (good explanation by woelan, complexing with the chloride anion makes sense). However, would the same complex mix of copper products be
formed if I were to use copper electrodes in a CuCl2 electrolyte?
| Quote: | Originally posted by 12AX7
From non-complexing electrolytes I get a low density sky blue precipitate, Cu(OH)2...
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Are there any other non-complexing electrolytes that you might be aware of? I was thinking perhaps Na2CO3 or NaHCO3 might not form complexes.
| Quote: | Originally posted by 12AX7
The last time I crystallized a CuCl2(aq) solution, I first made sure it was thoroughly green (by adding an oxidizer and if needed, water, but not to
the point of blue color) and only green before drying. Then I let it mostly dehydrate in a warm, dry place.
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Is this to say that attempting to crystallize a CuCl2 solution via hot plate is a big no-no? Apparently, the mass of liquid I had is a dark brown
goop, perhaps indicating some kind of complex (It can't be the anhydrous, since it's in aqueous solution)
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12AX7
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| Quote: | Originally posted by Darkblade48
How would you go about making Cu2O though? I'd imagine that a solution of NaCl might have worked, but apparently, it
doesn't (good explanation by woelan, complexing with the chloride anion makes sense). |
Well, it does for me, at least when it doesn't plate sponge metal out on the cathode.
You've searched, right?
| Quote: | | However, would the same complex mix of copper products be formed if I were to use copper electrodes in a CuCl2 electrolyte? |
Heh...
Well first of all you have to consider the electrode reactions. Since Cu(0) cannot be oxidized by water, it will deposit from solution in proportion
to the charge passed. Likewise, copper anode will be oxidized to copper ions in the same quantity. Works great for purifying copper, but rather
useless for making oxides.
An alkali metal salt, such as NaCl or Na2SO4, works because the sodium ions, when reduced by the cathode, then reduce the water. Or just as likely,
the sodium ions are never deposited at all and instead protons are deposited (being reduced from 2H+ + 2e- > H2 by the electron exchange), since
water has a lower reduction potential than sodium.
Whatever the case be, removing acid ions or concentrating alkali ions, the cathode becomes basic and only removes hydrogen from solution, rather than
the cation of the salt in solution, when the cation's potential is above that of water.
| Quote: | | Are there any other non-complexing electrolytes that you might be aware of? I was thinking perhaps Na2CO3 or NaHCO3 might not form complexes.
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Read the thread linked above. Apparently, I got cuprate from alkali salts. Very little copper actually did anything, but some did plate out.
I'm guessing ammonium salts would be a waste, since copper oxides are actually soluble due to the even stronger complexation.
Acetate might make Cu2O. Sodium or potassium nitrate and sodium chlorate and perchlorate would work for Cu(II) hydroxide. (KClO3 and -ClO4 would
work if they were soluble worth a damn.)
Nitrate, acetate, chlorate and perchlorate would also work for lead, if you are interested in oxidizing that.
| Quote: | | Is this to say that attempting to crystallize a CuCl2 solution via hot plate is a big no-no? Apparently, the mass of liquid I had is a dark brown
goop, perhaps indicating some kind of complex (It can't be the anhydrous, since it's in aqueous solution) |
Yeah, that's a complex between Cu(I) and Cu(II). If you dilute it with water until it turns blue, you'll get a white precipitate of CuCl,
which is no longer soluble in the blue hydrate complex solution.
Cu(II) solutions can be reduced in part to Cu(I) by copper metal and acidic solution (comparable to FeCl3 used to etch metals), but it isn't
quantitative, at least under a breathable atmosphere. Woelen would know more about this and other complexation behavior than I do.
Evaporating your solution will probably yield a mixture of CuCl and CuCl2, with a variable amount of Cu(OH,Cl) type products depending on conditions.
In my experience, brown means a lower oxidation state is present - if you add an oxidizer (H2O2, bleach (mind the sodium or calcium ions you're
adding!), TCCA, etc.) it should turn green. Now you can evaporate to form crystals; remember that this salt forms a mat of acicular (needle-like)
crystals that tend to be impossible to decant or filter without vacuum.
Tim (I like NiCl2 better, who cares if it's carcinogenic..)
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Darkblade48
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| Quote: | | Quote: | Originally posted by 12AX7
Well, it does for me, at least when it doesn't plate sponge metal out on the cathode.
You've searched, right?
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Yes, I have searched, however, that thread you linked me to was for the cupric oxide and not the cuprous (though I did have fun reading that thread as
well)
| Quote: | Originally posted by 12AX7
Cu(II) solutions can be reduced in part to Cu(I) by copper metal and acidic solution (comparable to FeCl3 used to etch metals), but it isn't
quantitative, at least under a breathable atmosphere.
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So in this case, the Cu(II) in the CuCl2 solution I had was reduced down to Cu(I) because of the excess HCl (and probably copper metal contaminants)?
| Quote: | Originally posted by 12AX7
Evaporating your solution will probably yield a mixture of CuCl and CuCl2, with a variable amount of Cu(OH,Cl) type products depending on conditions.
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I was not aware that CuCl would even exist at SATP.
I would imagine it to be hygroscopic enough to pull water from the air?
| Quote: | Originally posted by 12AX7
In my experience, brown means a lower oxidation state is present - if you add an oxidizer (H2O2, bleach (mind the sodium or calcium ions you're
adding!), TCCA, etc.) it should turn green. Now you can evaporate to form crystals. |
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So addition of an oxidizer will oxidize up the remaining Cu(I) into Cu(II) so that I can evaporate and have a (decently) pure sample of CuCl2?
Also, on another note, I was reading the thread on reducing copper sulfate using ascorbic acid to yield copper powder. However, when I tried using the
method outlined in the thread, I could never get it to work. The solution I had merely went from a blue colour to a light green shade, and did not go
any further. Perhaps chemoleo could shed some light on this?
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12AX7
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| Quote: | Originally posted by Darkblade48
So in this case, the Cu(II) in the CuCl2 solution I had was reduced down to Cu(I) because of the excess HCl (and probably copper metal contaminants)?
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Yeah, more or less. In this case, you get the reproportionation:
CuCl2 + Cu = 2CuCl (all aq)
CuCl stays in solution as complexed Cu+ ions (CuCl3(2-) or CuCl4(3-)?).
| Quote: | I was not aware that CuCl would even exist at SATP.
I would imagine it to be hygroscopic enough to pull water from the air? |
Nope! Deposits from solution as white anhydrous powder, well I should say no water of hydration anyway, of course there's plenty of water
adsorbed to, and inside, the crystals. However, in a moist oxidizing atmosphere, the slight solubility in water allows it to be oxidized to Cu(2+),
which of course causes a shortage of anions, so one is taken from the water. You get something like
2H2O + O2 = 4OH- (not what happens, but sums up the catalytic action of water)
4CuCl + 4OH- = 4Cu(OH,Cl)2
If oxidized and hydrated stoichiometrically, it'll end up 1:1 Cl- to OH- in the basic chloride precipitate. In reality you'll get a variety
of non-stoichiometric compounds, most likely with one or two crystal structures (out of the 19 or so oxychlorides known) dominating the composition,
plus some CuCl residue that escaped reaction.
| Quote: | | So addition of an oxidizer will oxidize up the remaining Cu(I) into Cu(II) so that I can evaporate and have a (decently) pure sample of CuCl2?
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Ya. Don't forget to remove the cation; I would recommend TCCA (which hydrolyzes to CO2 and HCl + HOCl, IIRC) or calcium hypochlorite followed by
sulfuric acid to precipitate the calcium. Common bleach is sodium hypochlorite so you'll be left with a salt impurity.
Back to the subject at hand, you want Cu2O eh? The easiest way would be electrolysis in NaCl solution as I said. Maybe you have too much or too
little voltage on your cell? Try like 5V maybe, and a concentrated brine solution.
You can add CuCl2 and Cu metal (use a lot of surface area like fine wire, MAKE SURE IT ISN'T ENAMELED  ) to an acid solution, that should reduce it at least partially to Cu(I), which you can then dilute to break the
chloro-complex and precipitate CuCl, which can be neutralized with NaHCO3 or whatever to get orange to brick red Cu2O. The only problem is, reducing
Cu(II) to Cu(I) has a really low yield, so you'd have to evaporate a whole lot of the mother liquor after dilution to repeat the process
like...five times...
Tim
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chloric1
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This weekend I embarked on an electrolytic cupric hydroxide quest. Friday I used dilute sodium sulfate solution between copper anode and cathode.
Copper hydroxide started to form but the current jumped up to 4 amps and the whole works(approx. 20 fl. oz.) heated to 40°C and promptly decopmposed
most of the hydroxide to oxide. Some copper plated back on the cathode so I got some nasty mixes. Today I am trying a very dilute bicarbonate
electrolyte. It is much less conducting and the hydroxide forms slowly and with grace. I shall patiently wait a week to see if my prudence rewards
me.
Fellow molecular manipulator
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triggernum5
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What kind of current density are you dealing with in terms of amps/electrode_surface_area.. I find that when electrolizing Cu solutions imparticular
that sponge deposits like 12AX7 mentioned can be eliminated by droping the current.. I'm pretty sure thats fairly common knowledge though.. I
apologize if I'm pointing out the obvious, I'm out of my league in most threads..
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tentacles
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Darkblade: This is a bit OT, but where in Canada are you that you can't find CuSO4? My Home Depot has it, though it's not with the drain cleaners.
edit: It's by the solder and drain snakes, with the septic tank boosters and crap.
[Edited on 24-3-2008 by tentacles]
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