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Al Koholic
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[*] posted on 4-12-2003 at 22:38
Copper Sulfate


I'm attempting to produce a good amount of copper sulfate for some copper plating experiments I want to try in the very near future. What I have been doing lately is running the 5V, 20A DC line from a computer power supply (AT) right into a pyrex beaker containing 800ml of 30% H2SO4 and 2 copper electrodes.

It was working great for a while and I was getting erosion of the copper anode and the solution was turning blue. The way I figure it, each molecule of H2 that leaves at the cathode is another Cu++ ion for my solution.

Anyway, after a few hours of operation I started to get a nice reddish buildup on the cathode as well. I think this may be copper (II) oxide although I am not sure. It did not have good adhesion to the surface of the electrode. However, I reasoned that since hydrogen was still coming off, I MUST be putting copper cations into solution still to balance the charge.

Now I can barely sustain a reduction of hydrogen for more than 2 minutes because the cathode gets coated with this red coating and just quits on me. What I think may be happening is that because I am at such a high voltage I am possibly plating copper back out onto the cathode as well now which might be responsible in some way for the coating that builds up. Since no more hydrogen is being evolved, am I simply moving copper from one electrode to the other? I think I may be...

This brings up my real question. After doing some reading on galvanic cells and electrolytic cells and consulting the handbook of chem and phys, I noted that Cu++ + 2e- => Cu is a 0.34 volt potential...ie: spontaneous. Sooooo I need to drive the reaction the other way with a voltage source greater than 0.34V. Now I am beginning to think that the buildup is a result of having both electrodes in the same bath...should I separate them with a salt bridge?

Also, I am now also going to modify a power supply (an old one with the big ass transformer and cap) by placing an AC dimmer switch between the power connection and the wall socket, then remove the voltage control board from the circuit of the supply so that power goes right from the transformer through a rectifier and into the cap providing me with essentially variable DC, controlled by the dimmer. This would allow me to run at say...0.5V and hopefully avoid providing force for other side reactions.
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Saerynide
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[*] posted on 5-12-2003 at 02:47


I was also planning on making copper sulfate. Could I make copper sulfate by using a copper anode and a carbon or stainless steel cathode and using a solution of magnesium sulfate? I was hoping to get Mg(OH)2 and CuSO4.
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[*] posted on 5-12-2003 at 05:50


Hmm, a few things I dont understand.
If you were indeed just transferring copper from one electrode to the other, then I wouldnt understand why the electrode would 'quit on you', as freshly plated copper also conducts electricity... so the reaction (even if you just electrolyse water) should continue until the electrode is corroded completely.
You want the copper sulphate for plating (sorry to disappoint you it doesnt work very well)... so I guess you could use CuCl2, too? In that case I seem to remember that Cu dissolves in HCL, with H2O2 - but dont take my word for it. Someone else in the forum will remember details for sure.




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[*] posted on 5-12-2003 at 08:56
Copper plating.


Chemoleo is right. Smooth, shiny copper plating using copper sulfate is not easy. I tried every possible combination. Only made it right once, using low voltage, very low current and graphite substrate.

I believe you have to use cyanide (poison!)salt compositions to have a beautifull finish.
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[*] posted on 5-12-2003 at 10:18
over engineering


If you want to eat up the copper with the sulfuric acid, why don't you just use alternating current? You don't need any fancy transformers, just put a safe load in series with the electrolysis cell. Start off with a 60 watt light bulb, which will give you about .5 amp of current through the solution. You will have to be careful that you don't get across the two electrodes, and get electrocuted when it is out of the solution, as the voltage will rise to line voltage when the resistance of the cell goes up. You will have to replace the acid as it is used up, and you willl have to replace the copper as it goes into solution. The solution (electrolyte) will also get warm if you use too much current. You can regulate the current with different size loads. Another safer method would be to use a stepdown transformer with AC output and it would we less likely to shock you. Have fun.
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Al Koholic
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[*] posted on 5-12-2003 at 10:43


This is all what I though would happen when I first started too....I assumed total corrosion of electrodes which I weighed so I would know when the reaction was done by the weight of the remaining copper.

The red coating on the cathode still confuses me and I can only think it means I am using too high a voltage. Something I also noticed was that yes, before I implemented fan cooling of the solution it would get quite warm after some time. Once it had warmed up sufficiently it began to produce gas at the anode as well!!! Must have been oxygen due to splitting water. At this point the reddish copper oxide coating began to deposit even more rapidly. Perhaps I need to keep the temp down.

Also I thought I would let you all know about my experiments with just this acid solution and plating. I managed to nicely plate a butterfly knife just by doing an acid dip and then a CuSO4/acid dip. The coating was not removed by vigorous cleaning with steel wool. It has begun to corrode because of all the use I give it but that was expected....adhesion still seems to be good.
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[*] posted on 5-12-2003 at 13:55


I know the main idea of this is making copper sulfate, but at least in America, it is common in the plumbing isle. It is used as a root killer for septic tanks. "Drain Line Root Killer" or the like. It is, however, a pentahydrate, or hexahydrate, I cannot remember. It is available though. Resume conversation.



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Saerynide
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[*] posted on 6-12-2003 at 00:20


So what about the magnesium sulfate idea? Does that work, since magnesium hydroxide isnt very soluble?
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[*] posted on 6-12-2003 at 03:30


I'd have to look up the solubillities, but I don't think copper hydroxide is much more soluble than Mg(OH)2. On heating it converts to CuO which is very insoluble. At best, you will have to keep the solution cold.

Why go to all this trouble? Cu disolves in hot H2SO4 anyway. (Do this outside)
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[*] posted on 6-12-2003 at 05:27


I dont got expensive-fancy-online-ordered chemicals. I have to make do with whats easily obtainable at home :P
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[*] posted on 6-12-2003 at 12:00


But if you're in the US, you don't need to order H2SO4 online. Go to a hardware store and look at the drain openers until you find one that has the container sealed inside a plastic bag. Look at the warning labels. It's probably concentrated H2SO4.

Edit: Ah, if your signup IP correctly indicates your location, then you're certainly not going to be able to get chemicals like you can in the US! I don't think your electrochemical method will work. However, copper + an acid + H2O2 will cause the metal to dissolve pretty easily and yield the corresponding metal salt. You could use vinegar and H2O2 to make copper (II) acetate, then precipitate it as carbonate (with addition of sodium carbonate). Now you can react the basic copper carbonate with even weak or dilute acids to get the corresponding copper salt. Can you find dilute H2SO4? It's used for adjusting pool pH, or (somewhat more concentrated) replacing the electrolyte in lead/acid storage batteries.

Actually, if you can get one of these forms of H2SO4, you might be able to use it directly with H2O2 and copper metal.

[Edited on 12-6-2003 by Polverone]
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[*] posted on 6-12-2003 at 15:23


Most places in the world have car batteries that contain sulphuric acid. Diulte sulphuric acid can be concentrated quite effectively by boiling it to drive off the water (You need an inert container to do this). If you do that with copper present you will get copper sulphate (and SO2, unfortunately).
You could add hydrogen peroxide to avoid the waste of acid but copper (and quite a lot of other things) catalyse the decomposition of peroxide so the yield based on peroxide will be low, and that assumes you can get it.
I don't know where you live but if it happens to be warm and dry you can even concentrate acid by just leaving it for the water to evaporate at room temperature. At 50% RH you get up to about 45% W/W. Better if it's hotter or drier.
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[*] posted on 6-12-2003 at 20:54


I wish I was still in Canada where you could get pretty much anything from Home Depot... you cant get anything here :(

And you know whats stupid? Half the drain cleaners here dont even got warning labels/ingredients lists at all! They dont even say the're caustic and that therefore you should not stick you hand in them :o No wonder why they never work when Im trying to unclog my drain :mad:

[Edited on 4-1-2004 by Saerynide]
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Al Koholic
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[*] posted on 6-12-2003 at 22:25


As far as I'm concerned...conc. sulfuric is just too valuble to me to be messing around with in the production of CuSO4. I'm sure if you have unending supplies you would disagree but I have just lost my source of good concentrated acid. Now I can only get battery electrolyte OTC and thats not good for reacting with Cu.

Anyway, electrolysis is just the way to go for making something like this IMHO. It would be nice if I could just make my own H2SO4 from something cheap and simple like sani-flush (NaHSO4 + heat --> Na2S2O7 + heat --> Na2SO4 + SO3 --> SO3 + H2O --> H2SO4). This is off topic and another project entirely...
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[*] posted on 7-12-2003 at 04:20


NaHSO4 mixed with alcohol gives a solution of H2SO4 in alcohol and Na2SO4 as a precipitate.
Boil, or better, distil off the alcohol and get the acid.
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[*] posted on 7-12-2003 at 07:13


What kind of alcohol? Or does any kind work?
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[*] posted on 7-12-2003 at 11:25
Electrolysis


Saerynide-- I have experimented with MgSO4 electrolysis and I think you could make some CuSO4. The technique is to get a small(4 oz) flower pot and plug the hole with plumbers putty. This will be your diaphragm. Use a plastic container with a lid. Cut a hole above the flower pot and insert a copper anode. Then insert a lead cathode adajacent to the pot. The lid should fit snuggly over the top of the flower pot. Use a high voltage 12 volts or more because you will have a voltage drop across the walls of the pot. You are using the pores of the ceramic to separate compartments. the flower pot should contain copper sulfate in solution and the cathode compartment will have a pasty Mg(OH)2 deposit. thsi method will also work with NaCl and Fe anodes to make FeCl2!! USe Nickel or stainless cathode for NaCl electrolyte!!



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[*] posted on 7-12-2003 at 12:34


When I was experimenting with making Cu(OH)<sub>2</sub> from electrolysis, I found that an NaCl electrolyte would cause the red covering but Na<sub>2</sub>CO<sub>3</sub> wouldn't. I may try chloric1's method on some salts when I'm bored. ;)



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[*] posted on 17-12-2003 at 18:43


Quote:
Originally posted by chloric1
Saerynide-- I have experimented with MgSO4 electrolysis and I think you could make some CuSO4. The technique is to get a small(4 oz) flower pot and plug the hole with plumbers putty. This will be your diaphragm. Use a plastic container with a lid. Cut a hole above the flower pot and insert a copper anode. Then insert a lead cathode adajacent to the pot. The lid should fit snuggly over the top of the flower pot. Use a high voltage 12 volts or more because you will have a voltage drop across the walls of the pot. You are using the pores of the ceramic to separate compartments. the flower pot should contain copper sulfate in solution and the cathode compartment will have a pasty Mg(OH)2 deposit. thsi method will also work with NaCl and Fe anodes to make FeCl2!!


I've done this even simpler. I had two bowls of MgSO4 solution, with a paper towel bridging them. I had copper electrodes on both sides. Anode side the solution gradually turned blue-green. Cathode side I did get some flaky magnesium hydroxide, and also some thick white powder which i had to scrape off the electrode from time to time (MgO?). About the time I calculated that enough amp-hours had passed through the solution to replace all the magnesium ions with copper ions, I noticed the solutions started trying to mix on their own (i.e. the green color started to flow through the paper towel towards the cathode, and some of the white precipitate started to flow into the anode side. I stopped the current flow at this point. Not sure of the ultimate purity but the anode side ended up with a beautiful blue-green color and when the water evaporated left blue-green crystals.

Later I took some of the CuSO4 I had made and electrolyzed using a carbon rod at the anode. I ended up with copper at the cathode. It was spongy and did not stick very well, but I verified that it conducted electricity. Once the blue color was gone I assumed I had a weak H2SO4 solution. I verified this by using reacting it with NaHCO3 and noting the fizzing.

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[*] posted on 17-12-2003 at 22:17


Sweet!!! I always wondered why no one talks about making H2SO4 by electrolysing CuSO4, so I thought it wasnt possible, but I guess I was wrong :D

Thanks so much for posting that.
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[*] posted on 18-12-2003 at 12:28


You can make H2SO4 that way, I used to electrolyse copper sulphate till most of the blue colour had gone. Shake it up with lead shavings and filter it (Pb +CuSO4--> PbSO4 +Cu) to remove the copper then boil it to remove the water. Guess what! I never made a lot of acid.
BTW the trick with NaHSO4 uses dry ethyl alcohol. I think the azeotrope would do.
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wink.gif posted on 19-12-2003 at 04:24


The reaction with sodium bicarbonate isn't the method here, it fizzes with CuSO4 as well. There's a substitution reaction, and the "copper bicarbonate" decomposes to copper carbonate, carbon dioxide and water (only alkali metal and ammonium bicarbonates are stable in a solid state). I tried it. It worked. :D:D:D
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[*] posted on 20-12-2003 at 17:35


Quote:
Originally posted by Theoretic
The reaction with sodium bicarbonate isn't the method here, it fizzes with CuSO4 as well. There's a substitution reaction, and the "copper bicarbonate" decomposes to copper carbonate, carbon dioxide and water (only alkali metal and ammonium bicarbonates are stable in a solid state). I tried it. It worked. :D:D:D


Okay, it looks like I did not have a good test then. I know when I mix MgSO4 with NaHCO3 there is no visible reaction. But aparently there is a reaction with CuSO4. I still have the solution - when I get my pH meter that I ordered from E-Bay I will measure the pH and compare it to the normal pH for a CuSO4 solution. This should tell whether I actually have a weak H2SO4 solution or just an even weaker (since no color visible) CuSO4 solution.

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[*] posted on 21-12-2003 at 05:27


Sodium carbonate will fizz with acid but not with copper sulphate (unless you heat it rather strongly).
If you don't happen to have sodium carbonate you can make it by boiling sodum bicarbonate solution.
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[*] posted on 21-12-2003 at 09:44


Are you positive that boiling a solution of Sodium Bicarbonate will yield Sodium Carbonate? I have seen the decomposition temperature referenced at 270<sup>o</sup>C numerous times. Last time I checked boiling water gets no where near this, unless you have some sort of super autoclave.

http://www.hummelcroton.com/nahco3_m.html

http://www.tabex.com/Tabex%20MSDS/Total%20Alkalinity%20up.pd...

[Edit] Ok, on further searching I found that 270 was the complete decomposition temperature. It starts to decompose around 93. So yes, it would would form some Carbonate, but no where near complete conversion.

[Edited on 12-21-2003 by Mumbles]
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