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janger
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[*] posted on 21-8-2004 at 05:34
CuCO3 and CuAc2 for lead dissolution


Hey Polverone, are you around?

You emailed me a method of making lead acetate. Just spent today creating some CuCO3 for the purpose, using the sulphate and NaHCO3. I washed it real well. It settled better with each wash.
However I wanted to make a bit for my experiments. So instead of filtering, which may have taken a while, I decided to decant most of the water off and apply gentle heat with a steam bath to dry it. However, as soon as it began to warm (I mean luke-warm) it started bubbling a little bit.

My good old chem book mentions that the normal copper carbonate doesn't exist, but as the basic carbonate, CuCO3.Cu(OH)2. Is this true, and if so will this cause problems in the acetate production?

Or is it that the basic carbonate is unstable even with low heat?

Another thing. My Cu carbonate reacts slower with vinegar than Ca or Na carbonates. Should it? It's just that the CO2 production with the latter ones is a good indication of how much to add. It's harder to tell with my CuCO3.

Thanks
Dave

[Edited on 26-8-2004 by chemoleo]
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[*] posted on 21-8-2004 at 09:47


A suggestion:
I have been playing around with said chemicals quite a bit.
To my experience, it is easier to start on the copper oxide, and to convert that to acetate, simply due to the ease of filtering.
So - Add to your CuSO4 a stoichiometric amount of NaOH. A turquoise, hard-to-filter preciptiate of Cu(OH)2 forms.
Then you just heat this up, until the colour goes completely black (via some deep olive green colour) - forming CuO
Filter, dry, and (if u are fussy about it) roast it a bit under an open flame, in a soup can or something.
Then dissolve with an excess of acetic acid, and add Pb to it. Some time later you got your Pb(Ac)2 :)

Another method is to react calcium carbonate with HAc, to get calcium acetate. This is mixed with the stoichiometric quantity of CuSO4 - the CuSO4 precipitates and can be filtered, and the Cu(Ac)2 stays in solution.

To answer your question: it doesn't matter, for this particular experiment (lead dissolution). All you want is a slight excess of acetic acid vs copper oxide/carbonate in the end, to make sure your final CuAc2 solution is acidic.



[Edited on 21-8-2004 by chemoleo]




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[*] posted on 21-8-2004 at 11:24


Copper carbonate reacts more slowly - I believe - mostly because of the lower solubility of copper acetate. If you want to speed things up, heating or boiling the vinegar and the carbonate together in a glass or ceramic container should do the trick. Be careful of foaming, though.



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[*] posted on 21-8-2004 at 11:56


Foaming is indeed a problem. When I added acetic acid to a saturated aqueous solution of copper carbonate, it kept flowing over if I didn't stir it once every ten seconds. And that was without any heating! I'd recommend against using a narrow-throated vessel, like a huge jar --- a wide, low plastic box with a lid is what I'll use for my next batch.



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[*] posted on 21-8-2004 at 13:45


"Copper carbonate" precipitated as a light green powder from solution by Na2CO3 solution, which is alkaline due to hydrolysis, is in fact a basic Cu(II) carbonate, containing OH- as well as CO3--. It is in fact the same as the mineral malachite. Azurite, the other naturally occurring Cu carbonate, bright blue in color , has a different structure.

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[*] posted on 21-8-2004 at 18:46


Ok, I think I'll try the CuO route.

But just for interests sake, why does the carbonate sludge bubble when even low heat is applied? Is it something to do with it being the basic carbonate? And what is the gas - CO2?

Dave
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[*] posted on 21-8-2004 at 19:38


Probably CO2 from the surplus of carbonate used.

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[*] posted on 21-8-2004 at 20:19


the green is not from nickel, rather it's from dissolved copper.



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[*] posted on 22-8-2004 at 02:53


This is not going well. I've spent all night and all day searching the net to get this worked out.

How long does dilute acetic (vinegar) take to react with CuO?

First I made CuO by adding excess NaOH, which forms the oxide without the need to heat the hydroxide. Also easier to filter. After drying, added vinegar, and after several hours there is only a slight blue color.

So I made some following chemoleo's method. The oxide doesn't seem to dissolve.

I've tried heating, boiling, adding more acid. Everything I can think of, but no luck.

My books say CuO reacts readily with acids. I expected this would be with acetic as well. Am I wrong?

Basically, I've had more luck with the calcium ac --> cu ac --> pb route,
although I don't think it'd give a very pure product.

If I had solubility tables of the acetates, I could probably recrystallise it. But I can't find any.

Can anyone give me some help with this?
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[*] posted on 22-8-2004 at 10:54


Cu(II) acetate, Cu(CH3COO)2.H20, is supposed to form dark green cystals, and be soluble in cold water (presumably 0ºC) to the extent of 7.2 parts/100 parts water, and be soluble in hot water (presumably 100ºC) to the extent of 20 parts/100 parts water. So the problem should not be due to an insoluble layer of acetate formng on the CuO.

CuO is supposed to be soluble in acids (except where insoluble salts result), cyanide solution, and NH4Cl solution. It should also be soluble in concentrated alkali metal hydroxide solutions to form cuprates.

John W.
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[*] posted on 22-8-2004 at 15:07


Quote:
Originally posted by JohnWW
Cu(II) acetate, Cu(CH3COO)2.H20, is supposed to form dark green cystals, and be soluble in cold water (presumably 0ºC) to the extent of 7.2 parts/100 parts water, and be soluble in hot water (presumably 100ºC) to the extent of 20 parts/100 parts water. So the problem should not be due to an insoluble layer of acetate formng on the CuO.

If these properties are correct, it should be easy to purify the CuAc via crystalization. There's actually some beautiful green crystals in a beaker I left from one of my tests.

I took some of the filtered Cu carbonate (from previous tried method), dissolved it in vinegar and added lead. The blue color has disappeared somewhat this morning, with a bit of copper on the lead pieces, but a white powder has also precipitated. Any idea what this is? My lead is fairly clean AFAIK.

Dave
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[*] posted on 22-8-2004 at 15:34


Pb(II) acetate is very soluble in water, 19.7 parts/100 parts at 0ºC, and 221 parts/100 parts at 50ºC, for the anhydrous salt. So the white precipitate should not be Pb acetate. But it partly hydrolyses in solution in the absence of excess acetic acid, so the white stuff is probably PbO.

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[*] posted on 22-8-2004 at 16:41


Thanks John.

Wish I hadn't got rid of my best chem books.

I thought it may have been lead hydroxide, since the basic copper carbonate contains Cu(OH)2.

I think I'll try and isolate pure CuAc crystals first then continue.

Thanks again
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[*] posted on 22-8-2004 at 16:57


I confess that I gave advice without first trying my own advice; though I had previously made copper acetate using the carbonate and pure acetic acid, I had not attempted it with vinegar. Today I attempted to dissolve copper carbonate in hot vinegar. I used a considerable excess of vinegar and raised the temperature to slow boiling, but after an hour a considerable amount of carbonate-colored powder remained on the bottom of the flask.

However, when I added lead shavings to this mixture, there was a vigorous frothing and additional gas generation. The shavings were rapidly covered in loosely-adhering copper. I continued to add more shavings to make sure lead was in excess relative to copper. Within 15 minutes all of the copper carbonate had dissolved and the solution had turned clear (somewhat clouded darkly by fine metallic particles). I filtered the mixture through a plug of cotton. The filtrate is colorless and clear. It gave a white precipitate with sodium hydroxide solution, which was oxidized to a golden yellow by addition of calcium hypochlorite.

I am now evaporating the remaining solution down. I would suggest that the best method of working may therefore be to heat the lead, copper carbonate, and vinegar all together. I have not experienced hydrolysis problems so far. I think that isolating pure copper acetate first may be an unnecessary and time-consuming step.




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[*] posted on 22-8-2004 at 21:03


Quote:
Originally posted by Polverone
I think that isolating pure copper acetate first may be an unnecessary and time-consuming step.
But the crystals look so damn cool:)

The CuO method seems to take too long with dilute acetic. I think your carbonate method is the way to go.

One problem was my lead pieces were too large. What is a good way to make shavings - just use an old knife?

Could you tell me if you made the carbonate, or purchased it? And is it ok to dry the CuCO3 on a steam bath?

I'm getting pretty close to a result now.

Dave
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[*] posted on 22-8-2004 at 22:08


I already had the carbonate, purchased some time back from a ceramic supply place. Yes, I just used a knife to make shavings from the lead. You shouldn't need to dry the CuCO3. Once it's been formed, filtered, and washed, I would directly add the wet precipitate to your vinegar and lead.

My lead acetate solution seems to have formed a small amount of white precipitate as it has evaporated down. I will later see if it goes clear again with additional vinegar. You're probably going to apply the lead acetate as a solution anyway, right? You may not need to forcibly evaporate it to concentrate it.




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[*] posted on 23-8-2004 at 03:36


I'm glad someone is repeating this. Was getting very tiresome:(

Since you're getting the white precipitate as well, it looks like it's not caused by my synth'd carbonate. A few hours ago I did another test batch. Had no problem dissolving my CuCO3 without heat. On adding shiny lead shavings, they were instantly coated in copper. But once again, a white ppt has formed. What seems to be happening is the copper drops out, then reacts somehow to form the white stuff. I say this because there doesn't seem to be much copper there again. This has happened with every batch.

This has occured both from the CO3 route and the CuO method. Heating doesn't dissolve it, and adding more vinegar doesn't seem to do anything.

I might just filter the solution and see if lead acetate crystallizes on evaporation.

Quote:
You're probably going to apply the lead acetate as a solution anyway, right?
Well actually, for making PbO2 anodes, the concentration is supposed to be fairly low, about 0.5 mol/litre. That's about 30g/l. Which is why I thought vinegar would be good enough for this. However, I'd like to try and isolate relatively pure PbAc so I don't need to go through this every time. The whole point of this is to find a method of creating PbO2 electrodes, using readily available chemicals. Hopefully this carbonate route will work, so the most expensive chemical would be the copper sulfate.

Dave
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[*] posted on 23-8-2004 at 04:31


Wait a minute. Due to the low concentration of vinegar, could the white ppt be lead carbonate?

If that's the case, I may have to go back to square one:
CaAc --> CuAc --> PbAc.
I'm pretty sure that one will work. Although it will be essential to crystallize out the pure PbAc, I think.

Dave
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[*] posted on 23-8-2004 at 11:31


I doubt the white ppt is lead (II) carbonate. In my case I tried collecting a small sample of such ppt and carried out a few tests on it, basically mixing it with nitric acid. No effervescence or reaction whatsoever occured and the white ppt remained even after heating. Most probably it is some lead (II) sulphate, due to impurities present in the cheap vinegar solution.



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[*] posted on 23-8-2004 at 11:52


Acetic acid being a stronger and less volatile acid, any PbCO3 would dissolve as acetate, liberating CO2.

BTW If the white precipitate is PbSO4, and not PbO, as you suspect from its unreactivity with HNO3, it would indicate adulteration of the (cheap) vinegar with an appreciable amount of H2SO4. You should conduct further qualitative tests for sulfate, e.g. by testing it with BaCl2 or BaNO3 solution - it would give a white precipitate of highly insoluble BaSO4, which can be distinguished fom BaCO3 through the sulfate not effervescing when tested with HCl. If this test is positive, take it to your Public Health Dept. to have it officially analysed by them.

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[*] posted on 23-8-2004 at 14:20


Yeah, I realised after posting it couldn't be lead carbonate. However I don't think it's all vinegar impurities. Even a small sample ends up with a fair amount of this ppt. Then again, I am using a fair amount of vinegar to make sure it's in excess. Just found out my cleaning vinegar is naturally brewed, so would have some impurities. But I did try one test using a cheap synthetic vinegar. Same result. Only other thing I could think of is impurites from the lead. So I'm now going to form some pure lead particles via electrolysis. Fairly simple to do.

BTW JohnWW, You've posted some solubility data previously. As I don't have access to such info, do you happen to have a little more detail on the sol. of the copper, calcium and lead acetates, from 0 to 100°C. If I begin with calcium acetate, I would be able to recrystallize it and remove most impurities from the vinegar, if any exist.

Thanks for all your help
Dave
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[*] posted on 23-8-2004 at 17:50


The solubility data for Cu and Pb(II) acetates I gave are all that Perrys Chemical Engineers Handbook had for them. Only the International Critical Tables, which I do not have and which have not, as far as I know, been released in electronic form, would have significantly fuller solubility data.

Another possible product of reaction of PbO with acetic acid is basic acetates, Pb(CH3COO)2.2Pb(OH)2.H20, or (more basic still) Pb(CH3COO)2.2Pb(OH)2. The first of these is "very soluble" in water; the second has solubilities of 5.55 parts/100 parts at ºC, and 18.2 parts/100 parts at 100ºC

Lead also formers a Pb(IV) acetate, Pb(Ch3COO)4, presumably by reaction of PbO2 with acetic acid at temperatures too low for Pb(IV) oxidation of acetate to be significant. However, I have no data on it.

Calcium acetate has a solubility with varies from 52 parts/100 parts water at 0ºC, to 45.5 parts/100 parts at 80ºC, a slight decrease with temperature. This may be due to hydrolysis.

John W.
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[*] posted on 24-8-2004 at 14:41


Here's a crappy pic of my acetate solution made from CuO. The oxide had not dissolved very well even after a few days, so I added lead to it and things started happening. The brown at the bottom is the copper particles.


Here's a quickly thrown together plating cell for testing. Gutter guard is handy stuff:)

I first electrolysed a 10% solution of NaOH to clean the carbon rod. The rod was then sat in vinegar for a while to neutralize the alkali.

And another with about 90mA flowing through the lead/copper acetate soln.

After an hour of electrolysis, solution is clear as the copper has dropped out. The current has dropped to about 50mA due to less conductivity. Seem to have a PbO2 coating. I'll post more on results when I can.

Dave

[Edited on 24-8-2004 by janger]
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[*] posted on 24-8-2004 at 14:48


I'd love to see the pictures that you've obviously forgotten to include.... :)

Edit: Viewing thye page source you seem to have linked to files on your local drive --- you have to attach them or put them on a webserver and make links to their (world-reachable) location.


[Edited on 2004-8-24 by axehandle]




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[*] posted on 24-8-2004 at 15:09


Quote:
Originally posted by axehandle

Edit: Viewing thye page source you seem to have linked to files on your local drive --- you have to attach them or put them on a webserver and make links to their (world-reachable) location.


[Edited on 2004-8-24 by axehandle]
Sorry 'bout that. Was trying to find out why they weren't loading from the web. Damn capitals:)
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