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Author: Subject: Copper Carbonate mysteriousness
jgourlay
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[*] posted on 14-3-2014 at 14:00
Copper Carbonate mysteriousness


Why is the copper carbonate that I buy from the pottery shop a malachite green, while the copper carbonate I make (copper sulphate + sodium bicarbonate) is sky blue?
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blogfast25
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[*] posted on 14-3-2014 at 14:37


There are various compounds that are described as 'copper carbonate', like CuCO3 but also Cu2(OH)2CO3 (Malachite), and even Cu3(OH)2(CO3)2 (Azurite). Precipitated 'copper carbonate' is usually of the 'basic' variety (contains OH groups).



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HgDinis25
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[*] posted on 14-3-2014 at 16:08


As stated in it's wiki page (http://en.wikipedia.org/wiki/Copper%28II%29_carbonate) copper carbonate used in pottery is, usually, a mixture of copper carbonate (CuCO3), bacic copper carbonate (Cu2(OH)2CO3) or even copper hydroxide, if made by the following process:
2 Cu (s) + H2O (g) + CO2 + O2 → Cu(OH)2 + CuCO3 (s)
Also to note, the reaction between copper ions and carbonate ions is, acording to wiki:
2 CuSO4 + 2 Na2CO3 + H2O → Cu2(OH)2CO3 + 2 Na2SO4 + CO2
This produces the basic species and not the pure stuff (CuCO3). I presume the same happens when you react Copper sulfate and sodium bicarbonate.
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thesmug
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[*] posted on 14-3-2014 at 16:34


When I made copper carbonate it was the green you describe. The wiki page also lists it as being green. Is it possible you didn't dry your copper carbonate well enough? I heated mine in an evaporating dish for a few hours and then filtered it overnight and it was green. It did, however, start out blue before I did the above procedures to it.
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HgDinis25
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[*] posted on 14-3-2014 at 16:46


Quote: Originally posted by thesmug  
When I made copper carbonate it was the green you describe. The wiki page also lists it as being green. Is it possible you didn't dry your copper carbonate well enough? I heated mine in an evaporating dish for a few hours and then filtered it overnight and it was green. It did, however, start out blue before I did the above procedures to it.


How did you filter it after drying it?
Also, pure CuCO3 should be blue, but inevitable presence of basic carbonate makes it green. Simply heating it won't decompose the basic carbonate back to pure CuCO3.

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jgourlay
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[*] posted on 14-3-2014 at 17:02


Ah...it's still wet. maybe that's it.
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thesmug
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[*] posted on 14-3-2014 at 17:31


Quote: Originally posted by HgDinis25  
Quote: Originally posted by thesmug  
When I made copper carbonate it was the green you describe. The wiki page also lists it as being green. Is it possible you didn't dry your copper carbonate well enough? I heated mine in an evaporating dish for a few hours and then filtered it overnight and it was green. It did, however, start out blue before I did the above procedures to it.


How did you filter it after drying it?
Also, pure CuCO3 should be blue, but inevitable presence of basic carbonate makes it green. Simply heating it won't decompose the basic carbonate back to pure CuCO3.


I just filtered it with regular quantitative slow-flow filter paper in a glass funnel overnight. Nothing reactive ever came into contact with it (except air). And it's probably not copper hydroxide from air because the color is uniform throughout (I powdered it) and started to form while heating. Could it be due to a change of water of crystallization composition?

[Edited on 3/15/14 by thesmug]

[Edited on 3/15/14 by thesmug]
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Texium
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[*] posted on 14-3-2014 at 20:02


Be extremely careful if you choose to heat it! I accidentally heated my first preparation too much and it all became black CuO, which was very disappointing.
Because of this, I decided to instead dry mine in a makeshift desiccator, basically just a large jar with some calcium chloride in the bottom. It's taking a few days but seems to be working very well.

Also, it appears that in the natural world, basic copper carbonate forms the minerals azurite and malachite. Azurite has two hydroxide groups and malachite has one, so based on this it seems that during dehydration, the color will shift from blue like azurite, to green like malachite. That appears to be what is happening with mine.
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DrChemistRabbit
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[*] posted on 15-3-2014 at 05:57


copper carbonate will hydrolize if contact with water or vapour and turn into malachite.
2CuCO3+ H2O == Cu2(OH)2CO3 +CO2
so it's extremely difficult to prepare and store it.




Chem is try.
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blogfast25
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[*] posted on 15-3-2014 at 06:27


Quote: Originally posted by DrChemistRabbit  
copper carbonate will hydrolize if contact with water or vapour and turn into malachite.
2CuCO3+ H2O == Cu2(OH)2CO3 +CO2
so it's extremely difficult to prepare and store it.


I'm fairly sure most 'copper carbonate' sold as CuCO3 is either basic carbonate or at least heavily contaminated with basic carbonate.

Acc. Wiki entry on copper carbonate:

"Pure copper carbonate is obtained from basic copper carbonate in the presence of carbon dioxide at 180 °C and 4.6 MPa (46 atm) pressure."

Assuming this is correct and given the mostly low tech applications of 'copper carbonate' it seems unlikely that most producers would go to that trouble...




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[*] posted on 22-3-2014 at 20:39


Ok, so this seems rather strange. For some reason, the copper carbonate which I had prepared a couple weeks ago has slowly changed from a nice, pretty, bluish green powder like it's supposed to look like, to a dark gray, clumpier powder. It has remained in a small closed vial, completely dry, free from outside contaminants, and at room temperature.
Does anyone have any idea what might have happened to it?
The only compound that I would guess it to be is CuO, but it would have had to thermally decompose to form CuO, right?
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blogfast25
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[*] posted on 23-3-2014 at 05:47


Quote: Originally posted by zts16  
Ok, so this seems rather strange. For some reason, the copper carbonate which I had prepared a couple weeks ago has slowly changed from a nice, pretty, bluish green powder like it's supposed to look like, to a dark gray, clumpier powder.


When you prepared the product, did you wash it with copious amounts of cold water? Did you check the pH of the wash water? I'm kind of wild guessing here but maybe small amounts of the alkali you used left in the product continued slowly to react, even after drying?

Another factor is the concentration of Cu salt from which you precipitated the 'carbonate'? Higher concentrations promote occlusions which can be hard to wash out.

The colour does point to CuO.

[Edited on 23-3-2014 by blogfast25]




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[*] posted on 23-3-2014 at 09:03


Ok, well, this is the procedure that I used:
I prepared the copper carbonate from copper sulfate and sodium carbonate, with an excess of sodium carbonate. The product was very cloudy, so I centrifuged it. It separated into easily separable layers of clear sodium sulfate/carbonate solution and a thick blue layer of copper carbonate paste.
I pipetted off the clear layer and the very top of the blue layer to ensure that I was getting as much sodium out as possible. After this, I left the copper carbonate paste in an evaporating dish inside a desiccator with calcium chloride for a week.
At that point, it was completely dried. There was also a tiny bit of white powder around the edge, which I thought was probably sodium sulfate, which I discarded. The copper carbonate has been in a clean vial ever since. The change in color started happening about two days after I put it in the vial.

Looking back… I realize that it could very well be the lack of washing, as there was most definitely sodium sulfate left over. There are even a couple tiny blue crystals that are visible now in my end product that look like copper sulfate. Also, I don't have any idea of what exactly my concentrations were because I don't have a balance yet (I'm getting one in tomorrow).
I'm still curious as to why it would form CuO though...
I guess I'll chuck this stuff out and try again next weekend!
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blogfast25
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[*] posted on 23-3-2014 at 10:33


It's the excess carbonate that is probably causing the problem here.

Next time, at least do as follows.

1. Make sure the solutions of CuSO4 and Na2CO3 aren't too concentrated. About 1 M I'd say. Mix with intensive stirring.

2. Only a small excess of Na2CO3 is needed: about 20 %

3. Filter the product (a coffee filter should work here) and wash it with copious amounts of COLD tap water. To do this efficiently always allow the previous wash water to filter down, only then add the next batch of water. Finally rinse once with deionised water.

4. Allow to drip 'dry' overnight, then spread out on filter paper to allow to dry. NEVER heat.

[Edited on 23-3-2014 by blogfast25]




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[*] posted on 23-3-2014 at 15:50


Yeah, I figured out the never heating it part the hard way, the first time I attempted it. I used a desiccator the second time because I had read somewhere else on here (I can't remember which thread) that it doesn't filter well. Regardless, I'll definitely be more careful next time. Having a balance should definitely help a lot, and washing it thoroughly.
Thanks for all of the advice!
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[*] posted on 30-3-2014 at 10:21


Quote: Originally posted by DrChemistRabbit  
copper carbonate will hydrolize if contact with water or vapour and turn into malachite.
2CuCO3+ H2O == Cu2(OH)2CO3 +CO2
so it's extremely difficult to prepare and store it.


My take on what is occurring, consider two steps in the above reaction:

CuCO3 + H2O = Cu(OH)2 + CO2

Cu(OH)2 + CuCO3 = CuCO3Cu(OH)2

the last step proceeding due apparently to the greater thermodynamic stability of the combined (or double?) salt. Reference, see http://books.google.com/books?id=0o4DnYptWdgC&pg=PA152&a...
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[*] posted on 30-3-2014 at 13:52


Yes, that is what happens, but that is no reason that it should turn black. Basic copper carbonate is green, and also what I was originally intending to prepare. The black color would have had to come from copper oxide due to an error, or combination of errors that I made, as blogfast explained.
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[*] posted on 31-3-2014 at 10:17


Quote: Originally posted by AJKOER  

My take on what is occurring, consider two steps in the above reaction:

CuCO3 + H2O = Cu(OH)2 + CO2

Cu(OH)2 + CuCO3 = CuCO3Cu(OH)2

the last step proceeding due apparently to the greater thermodynamic stability of the combined (or double?) salt. Reference, see http://books.google.com/books?id=0o4DnYptWdgC&pg=PA152&a...

For this reason, nickelcarbonate is prepared from a solution of nickelchloride and a solution of sodiumbicarbonate. If the pH of a carbonate solution becomes too high, OH- will start to bind to the nickel. The lower pH of a bicarbonate solution prevents this, and prevent a basic carbonate from being formed. I can imagine that the same trick may work with copper(II).
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[*] posted on 31-3-2014 at 10:47


Copper very easily forms basic carbonates and precipitation with sodium carbonate will always form some basic carbonate. As Bezaleel has said sodium bicarbonate is a better reagent for the preparation of heavy and transition metal carbonates as the pH of the solution is kept relatively low.
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