| Pages:
1
2 |
maliveline
Harmless
Posts: 7
Registered: 13-11-2014
Member Is Offline
Mood: No Mood
|
|
copper carbonate synth.
ok this is the reaction I'm doing.
2NaHCO3 + CuSO4-5H20 -----> CuCO3 + Na2SO4 +5H20+CO2
I guess My question is do I have to figure in the 5H2O for the molar mass of my copper sulfate. I'm a little rusty on my chemistry. So far the
solution I have seems very acidic and bubbly with everything that is mixed together But I guess that makes sense with the Na2SO4 on the other side of
the reaction. it just seems like I had to use a LOT of copper sulfate pentahydrate or whatever.
|
|
|
Zyklon-A
International Hazard
Posts: 1547
Registered: 26-11-2013
Member Is Offline
Mood: Fluorine radical
|
|
You equation isn't balanced. Look at the the hydrogen in your NaHCO3.
|
|
|
maliveline
Harmless
Posts: 7
Registered: 13-11-2014
Member Is Offline
Mood: No Mood
|
|
2NaHCO3 + CuSO4-5H20 ---> CuCO3 + Na2SO4 + CO2 + 6H2O
My CuCO3 seems to be slightly acidic still maybe from all the Na2SO4 on the right hand side of the reaction? Maybe its neutral I guess because my ph
testing method is pretty poor. I did notice it seemed to still bubble when I stirred everything like it was acidic or something. I'm going to end up
converting my CuCO3 to CuO but I'm concerned that any sulfate contaminate will end up in my CuO final product which I don't want to happen. Maybe it
wont effect the reaction and I can just rinse the final product with some bicarbonate solution or something?
[Edited on 17-12-2014 by maliveline]
[Edited on 17-12-2014 by maliveline]
|
|
|
Zyklon-A
International Hazard
Posts: 1547
Registered: 26-11-2013
Member Is Offline
Mood: Fluorine radical
|
|
Just rinse the precipitate with an excess of water.
Sodium carbonate, sodium sulfate and copper sulfate are all fairly soluble so a good wash will do fine.
However, It's considered best to use sodium carbonate rather then bicarbonate to make copper carbonate.
But if you're just going to decomposes it to oxide then it shaint make much difference.
|
|
|
MrHomeScientist
International Hazard
Posts: 1806
Registered: 24-10-2010
Location: Flerovium
Member Is Offline
Mood: No Mood
|
|
To your original question, yes you must take the crystal water into account when doing stoichiometry. it's important to calculate based on what you
actually have!
Anhydrous copper sulfate is a white powder, and has a molecular weight of 159.5 g/mol.
Hydrated copper sulfate, on the other hand, is a blue crystal, and has a molecular weight of 249.5 g/mol.
Since the latter is what you actually have, that's the number you need to use when doing practical experiments. Similarly, whenever you use a
hydroxide it's a good idea to use a 10% excess to account for water that is inevitably absorbed by the compound. KOH, for example, is said to contain
10-15% water no matter what you do to try and dehydrate it!
[Edited on 12-17-2014 by MrHomeScientist]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
@ maliveline:
Please use the search facility before posting new threads, this is literally the umpteenth thread on copper carbonate, see here for instance:
http://www.sciencemadness.org/talk/viewthread.php?tid=50822
Now trust me on this one: CuCO<sub>3</sub> simply doesn't exist: all (2 known) copper carbonates are copper hydroxycarbonates.
W/o any references this thread belongs in 'beginnings' (see forum rules).
[Edited on 17-12-2014 by blogfast25]
|
|
|
maliveline
Harmless
Posts: 7
Registered: 13-11-2014
Member Is Offline
Mood: No Mood
|
|
you can move the thread if you want sorry i didn't search harder. and to the suggestion that I should rinse the precipitate. I suppose that would work
with vacuum filtration which I dont have. but otherwise filtering CuCO3 is a pain in the a** and takes forever so rinsing it about as much fun as
watching paint dry.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by maliveline  | | you can move the thread if you want sorry i didn't search harder. and to the suggestion that I should rinse the precipitate. I suppose that would work
with vacuum filtration which I dont have. but otherwise filtering CuCO3 is a pain in the a** and takes forever so rinsing it about as much fun as
watching paint dry. |
Like I wrote: CuCO<sub>3</sub> doesn't exist. The two main copper basic carbonates are Malachite,
Cu<sub>2</sub>CO<sub>3</sub>(OH)<sub>2</sub>,and Azurite, Cu<sub>3</sub>(CO<sub>3</sub> <sub>2</sub>(OH)<sub>2</sub>.
Without washing your product is impure. Try digesting the precipitate by allowing it to stand in a warmish place overnight, before filtering.
|
|
|
maliveline
Harmless
Posts: 7
Registered: 13-11-2014
Member Is Offline
Mood: No Mood
|
|
ya i feel ya thanks.
|
|
|
ChemSwede
Harmless
Posts: 26
Registered: 20-8-2013
Location: Sweden
Member Is Offline
Mood: No Mood
|
|
This seems to be the latest thread about copper carbonate synthesis.
Yesterday I decided to make a batch of 50g of basic copper carbonate starting with CuSO4*5H20 and K2CO3.
I did calculations using the formula:
2 CuSO4 + 2 K2CO3 + H2O → Cu2(OH)2CO3 + 2 K2SO4 +
CO2
(for CuSO4 I took the 5H2O into account).
I dissolved 115g of CuSO4 in 460ml of water and 75,6g of K2CO3 in 550ml of water. The carbonate was in 20% excess.
The two solutions were mixed while stirring and the precipitated product started out as sky bly but quickly turned greener. CO2 was given
off.
I stirred the mix for 20min and then let it sit over night.
The next day the green precipitate had settled at the bottom. I filtered it with a büchner funnel and a handheld vacuum pump (worked very well) and
washed with lots of distilled water. It's now being dried in a can with CaCl2.
After drying I will determine yield.
Questions and reflections:
I noticed that the filtrate was a bit bluish, and that seemed strange since I used an excess of K2CO3. I measured the pH (around
8) and then added some K2CO3, but no new precipitate formed. What could cause the blue colour? I guess that the weakly alkaline
pH suggests that some potassium carbonate was still present.
Any faster way of drying the product? I suppose that basic copper carbonate is not soluble in, or reacts with, acetone or ethanol. Perhaps I could use
any of those as the last step during filtering to get rid of water?
Any advantages or disadvantages using potassium carbonate instead of sodium carbonate? I know that even trace amounts of sodium is bad if you want to
make coloured flames with f ex copper compounds.
Any other improvements I could make?
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Precipitate from a more dilute solution: this avoids occlusion of other ions (like sodium or potassium) into the precipitate's crystal lattice.
Drying in desiccator will be slow. CuCO3(OH)2 resists 100 to 150 C easily without any decomposition, so oven drying is faster. A final rinse with
acetone (on Buchner), then oven drying to constant weight at 100 C is probably fastest and best. It's what I use.
[Edited on 30-12-2014 by blogfast25]
|
|
|
ChemSwede
Harmless
Posts: 26
Registered: 20-8-2013
Location: Sweden
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by blogfast25 |
Precipitate from a more dilute solution: this avoids occlusion of other ions (like sodium or potassium) into the precipitate's crystal lattice.
Drying in desiccator will be slow. CuCO3(OH)2 resists 100 to 150 C easily without any decomposition, so oven drying is faster. A final rinse with
acetone (on Buchner), then oven drying to constant weight at 100 C is probably fastest and best. It's what I use.
[Edited on 30-12-2014 by blogfast25] |
Thanks.
I read in another post on SM that the solutions should not be stronger than 1M so that's what I aimed for. What conc. would you use?
Any idea of why the filtrate was weakly blue? Maybe small particles of product that slipped through the filter paper? Unreacted copper sulfate?
The potassium carbonate should be in excess (75,6g to 115g of copper sulfate), and I stirred the reaction for ca 20min with a magnetic stirrer and
left it to react over night. Strange if any copper sulfate would be left.
Oven drying sounds good. However, I don't have an oven with good temp control, and I don't want to risk decomposing the carbonate to CuO.
|
|
|
Texium
Administrator
Posts: 4508
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
Yeah, I would recommend not even trying to heat copper carbonate. It can decompose very quickly. Just let it dry out in the air, or if you can, set up
a desiccator. That would be the best way to dry it without any decomposition, and it doesn't take all that long to do.
Also, as for Zyklon-A's earlier comment about how it's better to use sodium carbonate than sodium bicarbonate, I'd have to disagree. Sodium carbonate
is more basic than sodium bicarbonate in solution, and this promotes more copper hydroxide formation that leaves you with a less pure product. I know
we had a discussion about this somewhere in one of the pre-existing threads. I'm not sure which one though, or if it was ever completely resolved.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by ChemSwede |
I read in another post on SM that the solutions should not be stronger than 1M so that's what I aimed for. What conc. would you use?
Any idea of why the filtrate was weakly blue? Maybe small particles of product that slipped through the filter paper? Unreacted copper sulfate?
The potassium carbonate should be in excess (75,6g to 115g of copper sulfate), and I stirred the reaction for ca 20min with a magnetic stirrer and
left it to react over night. Strange if any copper sulfate would be left.
Oven drying sounds good. However, I don't have an oven with good temp control, and I don't want to risk decomposing the carbonate to CuO.
|
1 M is OK.
The colour is extremely subjective and depends on several factors, like granulometry, moist v dry, oil v water. After washing and drying your product
will look much like anyone else’s.
Oven drying is fine up to at least 150 C without any decomposition.
| Quote: Originally posted by zts16 | Yeah, I would recommend not even trying to heat copper carbonate. It can decompose very quickly. Just let it dry out in the air, or if you can, set up
a desiccator. That would be the best way to dry it without any decomposition, and it doesn't take all that long to do.
Also, as for Zyklon-A's earlier comment about how it's better to use sodium carbonate than sodium bicarbonate, I'd have to disagree. Sodium carbonate
is more basic than sodium bicarbonate in solution, and this promotes more copper hydroxide formation that leaves you with a less pure product. I know
we had a discussion about this somewhere in one of the pre-existing threads. I'm not sure which one though, or if it was ever completely resolved.
|
You go by the many myths that are peddled on this subject.
Malachite is quite stable up to 150 C and even higher. Air drying takes much longer than you might think. Do that to constant weight and you’ll
understand what I mean.
Carbonate v. bicarbonate: another myth. The truth is that it makes not a blinding bit of difference with regards to the final composition of the
product after washing and drying. No copper hydroxide is formed in either case if you respect stoichiometry:
2 CuSO4 + 2 Na2CO3 + H2O → Cu2CO3(OH)2 + 2 Na2SO4 + CO2
2 CuSO4 + 4 NaHCO3 → Cu2CO3(OH)2 + 2 Na2SO4 + 3 CO2 + H2O
As you can see using bicarbonate is actually quite wasteful.
The trouble with copper basic carbonate is that all kinds of assertions are being made without a single actual verification or corroboration. Hence:
much myth and precious little fact. It’s THE subject where every Dick, Tom and Harry feels that they have to ‘contribute’ when most of that is
just waffle.
[Edited on 31-12-2014 by blogfast25]
|
|
|
ChemSwede
Harmless
Posts: 26
Registered: 20-8-2013
Location: Sweden
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by blogfast25 |
1 M is OK.
The colour is extremely subjective and depends on several factors, like granulometry, moist v dry, oil v water. After washing and drying your product
will look much like anyone else’s.
Oven drying is fine up to at least 150 C without any decomposition.
[Edited on 31-12-2014 by blogfast25] |
I followed your advice with oven drying, and it seems to have worked just fine. I don't see any signs of brown or black CuO anywhere. I didn't
monitor the temp, but I kept the oven hatch open.
I weighed the product a few times until the weight was constant.
47,8g water evaporated, so almost half the weight was water. It would have taken forever to get rid of that with a dessicator.
The dry product weighed 51,4g, which gives me a yield of 102,8%.
That's a bit... unusal.
I suppose that the extra weight could come from contamination with potassiumsulfate/carbonate or water.
From what I understand it's also difficult to know how many CO3 -and OH-groups the product contains. A few more of those could cause the
extra weight.
It would be interesting to determine the product's copper content based on 115g of copper sulfate.
I'm still a bit puzzled about the blue colour of the filtrate, not the product's colour.
The liquid that went through the filter was clear but weakly blue. That seemed strange since all the copper sulfate should have reacted. Addition of
some K2CO3 caused no more precipitate, and the filtrate's pH was slightly alkaline, around 8.
|
|
|
Texium
Administrator
Posts: 4508
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
@blogfast: I go by some educated guesses, yes, but also by direct experience, in that the first time I made copper carbonate almost a year ago, using
sodium carbonate, gentle heating on a hotplate was enough to decompose it. I have also seen it withstand much more heat. I am not sure why that is.
Since then, the other couple of times that I've made it, I filter it and then use a desiccator to dry it, which takes a day or two.
As for the bicarbonate vs carbonate, when precipitated with sodium carbonate, it is much more blue, whereas with bicarbonate it is green like
malachite. And I always use stoichiometric amounts.
If I had a better, more analytical way of determining what sort of products I'm getting, I would use it, but unfortunately I do not.
I think it would be bet if we were to settle these debates about copper carbonates once and for all. I'd happily participate in contributing to that.
[Edited on 1-2-2015 by zts16]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
@zts16:
If your copper basic carbonate decomposes on a hot plate, either the setting is too high or Cu(OH)2 was present. The latter may be caused by poor
washing or excess CuSO4 still being present. Cu2CO3(OH)2 can be dried on a medium hot plate without problems.
Re the colour, it's what gets many into trouble because they think the product is different, if the 'perceived' colour is different. My copper basic
carbonate from sodium carbonate and CuSO4 is almost always blue immediately on precipitation, then takes on its true Malachite green on washing and
drying.
Blue and green are very subjective and the colour of such a precipitate depends on many factors. Even my Cu2Cl(OH)3 looked blue when freshly
precipitated, only to become a deep Malachite green on working up.
Re. bicarbonate v. carbonate, use experiments and the stoichiometric equations above to determine weight loss and how much CO2 (mol) is released per
mol of carbonate or bicarbonate used. You'll find these ratios are fixed. The mass balances don't lie: there is only Cu2CO3(OH)2.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by ChemSwede | The dry product weighed 51,4g, which gives me a yield of 102,8%.
That's a bit... unusal.
I suppose that the extra weight could come from contamination with potassiumsulfate/carbonate or water.
From what I understand it's also difficult to know how many CO3 -and OH-groups the product contains. A few more of those could cause the
extra weight.
It would be interesting to determine the product's copper content based on 115g of copper sulfate.
I'm still a bit puzzled about the blue colour of the filtrate, not the product's colour.
The liquid that went through the filter was clear but weakly blue. That seemed strange since all the copper sulfate should have reacted. Addition of
some K2CO3 caused no more precipitate, and the filtrate's pH was slightly alkaline, around 8. |
The blue filtrate is almost certainly a weak cuprate solution: Cu(OH)<sub>4</sub><sup>-</sup>. I think you may have used a bit
too much K2CO3. Adding more carbonate to cuprate solutions will not cause any precipitation. See also Edit.
'102.8 %' is nothing more than measuring error. Small amounts of occlusion can cause that.
The CO3/OH ratio is fixed, if you prepared the product properly it will be 1:2. That this ratio 'can vary' is an old and obsolete belief.
Better than determining copper content is to either determine CO2 content or Molar Mass of the product.
Edit:
Higher up you wrote:
"The potassium carbonate should be in excess (75,6g to 115g of copper sulfate), and I stirred the reaction for ca 20min with a magnetic stirrer
and left it to react over night. Strange if any copper sulfate would be left."
The Molar Mass of K2CO3 is 138.2 g/mol
The Molar Mass of CuSO4.5H2O is 249.5 g/mol
The reaction is K2SO4 + CuSO4 + 1/2 H2O == > 1/2 Cu2CO3(OH)2 + K2SO4 + 1/2 CO2
75.6 g K2CO3 thus requires 136.5 g CuSO4.5H2O, NOT 115 g CuSO4.5H2O. 75.6 g K2CO3 is almost twice the stoichiometric amount. At such
high alkalinity it's possible that some of the copper basic carbonate re-enters solution as cuprate, especially on standing
overnight:
1/2 Cu<sub>2</sub>CO<sub>3</sub>(OH)<sub>2</sub>(s) + 2 OH<sup>-</sup>(aq) === >
Cu(OH)<sub>4</sub><sup>-</sup>(aq) + 1/2 CO<sub>3</sub><sup>2-</sup>(aq)
This adequately explains your blue filtrate.
[Edited on 3-1-2015 by blogfast25]
|
|
|
ChemSwede
Harmless
Posts: 26
Registered: 20-8-2013
Location: Sweden
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by blogfast25 |
The blue filtrate is almost certainly a weak cuprate solution: Cu(OH)<sub>4</sub><sup>-</sup>. I think you may have used a bit
too much K2CO3. Adding more carbonate to cuprate solutions will not cause any precipitation. See also Edit.
'102.8 %' is nothing more than measuring error. Small amounts of occlusion can cause that.
The CO3/OH ratio is fixed, if you prepared the product properly it will be 1:2. That this ratio 'can vary' is an old and obsolete belief.
Better than determining copper content is to either determine CO2 content or Molar Mass of the product.
Edit:
Higher up you wrote:
"The potassium carbonate should be in excess (75,6g to 115g of copper sulfate), and I stirred the reaction for ca 20min with a magnetic stirrer
and left it to react over night. Strange if any copper sulfate would be left."
The Molar Mass of K2CO3 is 138.2 g/mol
The Molar Mass of CuSO4.5H2O is 249.5 g/mol
The reaction is K2SO4 + CuSO4 + 1/2 H2O == > 1/2 Cu2CO3(OH)2 + K2SO4 + 1/2 CO2
75.6 g K2CO3 thus requires 136.5 g CuSO4.5H2O, NOT 75.6 g CuSO4.5H2O. 75.6 g K2CO3 is almost twice the stoichiometric amount. At such
high alkalinity it's possible that some of the copper basic carbonate re-enters solution as cuprate, especially on standing
overnight:
1/2 Cu<sub>2</sub>CO<sub>3</sub>(OH)<sub>2</sub>(s) + 2 OH<sup>-</sup>(aq) === >
Cu(OH)<sub>4</sub><sup>-</sup>(aq) + 1/2 CO<sub>3</sub><sup>2-</sup>(aq)
This adequately explains your blue filtrate.
[Edited on 3-1-2015 by blogfast25] |
So formation of cuprate-ion could explain the blue colour. Well, I suppose it wasn't much, since the colour was weak and my yield was high.
I used 115g of copper sulfate to a 20% excess of potassium carbonate.
I aimed for 50g of Cu2(OH)2CO3 with a molar mass of 221 = 0,23 moles
I used the formula 2 CuSO4 + 2 K2CO3 + H2O → Cu2(OH)2CO3 + 2 K2SO4 + CO2
For CuSO4*5H2O I calculated (0,23*2)*250 = 115g
For K2CO3 I calculated (0,23*2)*138 = 63,5 + 20% excess = 75,6g
So I used 115g of CuSO4 (not 75,6g) to 75,6g of K2CO3.
Maybe it would be sufficient with 10% excess of K2CO3 to minimize formation of cuprate.
The pH of the filtrate after one day was around 8, so not that high though.
I checked most of the threads on copper carbonate before I performed the synth, and there I read that different basic copper carbonates ( f ex
malachite or azurite) are formed depending on pH, conc, temp etc. This is not true then? With this method the only product should be Cu2(OH)2CO3?
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by ChemSwede |
I checked most of the threads on copper carbonate before I performed the synth, and there I read that different basic copper carbonates ( f ex
malachite or azurite) are formed depending on pH, conc, temp etc. This is not true then? With this method the only product should be Cu2(OH)2CO3?
|
The thing about the 'different basic basic carbonates' is what you could call a half truth and from there originates much nonsense.
Azurite, Cu3(CO3)2(OH)2 (or 2CuCO3.Cu(OH)2, if you prefer), does indeed exist but every authorative source will tell you that this cannot be prepared
by simple precipitation of a cupric salt with a soluble carbonate or bicarbonate.
Azurite is a fairly rare mineral that seems to form in very specific conditions of copper(II) concentration, calcium bicarbonate concentration and
CO2 pressure. It is metastable and tends to revert back to Malachite in a process known as 'pseudomorphing':
2 Cu3(CO3)2(OH)2 + H2O === > 3 Cu2CO3(OH)2 + CO2
The required conditions to form Azurite cannot be achieved in simple precipitation conditions.
Here's a thread where we're working hard to create these conditions, so far without luck.
http://www.sciencemadness.org/talk/viewthread.php?tid=50822&...
Re. cuprate and using an excess carbonate. The filtrate was light blue because there wasn't much cuprate in it and because cuprate is very intensely
blue. pH 8 only confirms that: at that pH the cupric ion [Cu(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup> cannot
exist. Copper is slightly amphoteric though.
Using an excess of one reagent is usually good practice but here it is not necessary: the precipitation reaction is very swift and runs to completion
even in the absence of an excess. Use exact stoichiometric amounts.
[Edited on 3-1-2015 by blogfast25]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is a partial extract from Atomistry on Copper carbonates (link: http://copper.atomistry.com/cupric_carbonates.html ):
"The normal salt has not been prepared. Malachite, CuCO3,Cu(OH)2, occurs in monoclinic crystals, density 3.7 to 4. It has been produced artificially.
Azurite, 2CuCO3,Cu(OH)2, forms monoclinic crystals, density 3.5 to 3.88. It has been obtained by a laboratory method.
The formation of basic carbonates of copper by the interaction of solutions of cupric sulphate and of the carbonates of sodium has been investigated
by Pickering. Sodium carbonate precipitates a blue, basic carbonate, 5CuO,2CO2,nH2O, which is converted by drying over sulphuric acid at 100° C. into
another hydrate of green colour, 5CuO,2CO2,3H2O. In moist air the green hydrate becomes reconverted into the blue form:
5CuSO4 + 8Na2CO3 + 3H2O = 5CuO,2CO2 + 5Na2SO4 + 6NaHCO3.
The blue carbonate is transformed by concentrated aqueous sodium carbonate into cupric hydroxide, and by aqueous sodium hydrogen carbonate into
malachite, 2CuO,CO2,H2O. Pickering considered ordinary commercial copper carbonate to be similar in constitution to malachite, a view questioned by
Dunnicliff and Lai.
Sodium hydrogen carbonate and cupric sulphate react to precipitate a blue, basic carbonate, 5CuO,3CO2,nH2O, converted by drying at 100° C. into
another blue hydrate, 5CuO,3CO2,7H2O. Another basic carbonate is also produced in the same reaction. It has the formula 8CuO,3CO2,6H2O, is dark blue
in colour, and becomes green at 100° C. No other basic carbonate was isolated by Pickering. All the products are insoluble in water and
sodium-carbonate solution, but dissolve slightly in solutions of carbon dioxide and of sodium hydrogen carbonate, with production of the normal
carbonate or a double carbonate.
Feist has prepared a basic carbonate, 7CuO,4CO2,H2O, by powdering together crystallized cupric sulphate and sodium carbonate, and then adding water.
It is difficult to separate the substance from a basic cupric sulphate simultaneously formed. Auger has described an amorphous basic carbonate of the
formula 8CuO,5CO2,7H2O. Another basic carbonate, 7CuO,2CO2,5H2O, has been prepared by the interaction of a mixture of sodium carbonate and sodium
hydrogen carbonate with cupric sulphate in aqueous solution. Complex carbonates of copper with sodium and potassium have also been obtained. An
example of this type of double salt of the formula Na2Cu(CO3)2,3H2O crystallizes on addition of a solution of cupric acetate to one of sodium
carbonate and sodium hydrogen carbonate at 50° C. It forms needles or rosettelike agglomerations, and above 100° C. "
As background on Atomistry.com, it is essentially provides an extract from various Chemical journals (not specified) by authors (not identified) over
time (a significant omission to assess accuracy).
Note, per Atomistry, Blogfast is correct in his assertions that it is unlikely that one actually will prepare pure CuCO3, and not one of many basic
copper carbonates.
Assuming the historical literature as presented above is accurate, one forms varying basic carbonates of copper depending on such factors as combing
dry or a concentrated aqueous Na2CO3 (or, even dilute, for example), employing NaHCO3 in place of Na2CO3 (so pH differences are material), and even
"powdering together crystallized cupric sulphate and sodium carbonate", along with other variables including, for example, the creation of double
salts with the presence of aqueous Na or K compounds.
------------------------------------------------------------------
Something interesting I have been thinking about is forming copper ammonium hydroxide and adding Mg(HCO3)2. Separating out the Mg(OH)2, and letting
the copper ammonium bicarbonate stand in the open and slowly evaporate.
Note, if I formed copper ammonium sulfate instead, and let so evaporate, I would expect CuSO4. Also, the similar evaporation of copper ammonium
hydroxide forms Cu(OH)2 nanotubes. So, I would expect a basic copper carbonate also, but the different in pH, particle size,..., not sure if the
product would agree with any of the basic carbonates cited above.
[Edited on 26-1-2015 by AJKOER]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
AJ:
atomistry is for the most part unreliable Tinkerweb filler, specialising in quote mining without actual references. Take it with a bag of salt and
prefer much more primary and authorative sources.
Mentioning "Dunnicliff and Lai" (WHO???) makes the text look more intelligent without actually informing anyone about anything.
These pseudo references are in any case very old: no one uses notations like "7CuO,2CO2,5H2O" anymore.
It's now generally agreed that there are two basic carbonates of copper: Malachite and the much more elusive Azurite, PERIOD. And no 'straight' CuCO3,
for the likely and simple reason that it is too soluble with respect to Cu(OH)2.
Please refrain from bringing up old crap in new clothes, something atomistry excels at. It only muddies the waters.
Perhaps we should also go back to Bohr's atomic model, just because there are still some references as to its validity?
What you're doing here is promoting old myths: that's very antithetical to good science.
As regards:
| Quote: Originally posted by AJKOER | Something interesting I have been thinking about is forming copper ammonium hydroxide and adding Mg(HCO3)2. Separating out the Mg(OH)2, and letting
the copper ammonium bicarbonate stand in the open and slowly evaporate.
|
What makes you think that would work and what is it supposed to yield?
[Edited on 27-1-2015 by blogfast25]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Actually, whatever happens I probably find interesting. The expected is some basic copper carbonate or a mixture of such and Cu(OH)2, and the very
less likely alternatives are pure Cu(OH)2 or CuCO3.
The action on pure Copper of dilute ammonia in dilute H2O2 (and in a hurry, add a tiny amount of NaCl electrolyte), readily forms the copper ammonium
hydroxide. My experience is upon adding MgSO4 (or Mg(NO3)2), does form a white suspension of Mg(OH)2 and the corresponding aqueous royal blue copper
ammonium salt. However, to obtain the solid copper ammonium salt, one could evaporate in a stream of NH3, else one gets just the copper salt without
ammonia.
Working with the unstable Magnesium bicarbonate, should correspondingly form Mg(OH)2 and Copper ammonium bicarbonate. The later should form the
highly unstable Copper bicarbonate, which upon evaporation in open air removing the ammonia, readily creates basic Copper carbonate..
Now with respect to as whether there are just two basic copper salts, to quote Wikipedia on Copper carbonate:
""Copper carbonate" was the first compound to be broken down into several, separate elements (copper, carbon, and oxygen). It was broken down in 1794
by the French chemist Joseph Louis Proust (1754–1826). When heated, it thermally decomposes to form CO2 and CuO, cupric oxide, a black solid. The
basic copper carbonates, malachite and azurite, both decompose forming CO2 and CuO, cupric oxide.[5]"
Your point of two salts may not refute the similarly derived ratio of elements per the Wikipedia quote above as computed by the chemists cited in
Atomistry. The actual product produced per the varying paths may just contain some free Cu(OH)2 or a weighted average of the underlying two salts,
which means you may still have to work with the indicated salt mixture as if it where a true compound.
[Edited on 27-1-2015 by AJKOER]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Copper ammonium hydroxide: what evidence do you have for its existence?
| Quote: Originally posted by AJKOER |
Your point of two salts may not refute the similarly derived ratio of elements per the Wikipedia quote above as computed by the chemists cited in
Atomistry. The actual product produced per the varying paths may just contain some free Cu(OH)2 or a weighted average of the underlying two salts,
which means you may still have to work with the indicated salt mixture as if it where a true compound. |
What are the chances of such mixtures corresponding to very specific ratios like "5CuO,2CO2,3H2O", "5CuO,3CO2,7H2O" or "8CuO,3CO2,6H2O" or
"7CuO,4CO2,H2O" et al? These are ratios specific to compounds and double salts, not mixtures.
What are the chances of Cu(OH)2 and CuCO3 combining into such a plethora of double salst, all thermodynamically stable enough to exist? Nil, nada,
zilch.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Perhaps part of the answer is explained by this study of the thermal decomposition stages and intermediate compounds: "Thermal decomposition of the
basic copper carbonates malachite and azurite", by I.W.M. Brown, K.J.D. Mackenzie, G.J. Gainsford...
"Abstract
Thermogravimetry (TG) and evolved gas analysis (EGA) studies of malachite, CuCO3 · Cu(OH)2, and azurite, 2 CuCO3 · Cu(OH)2, heated in helium carrier
gas at 10° min−1 show that malachite decomposes in a single step at 380°C, in which water and CO2 are lost simultaneously. By contrast, the two
azurites investigated both decompose under these conditions in two approximately equal steps, losing one-half of their CO2 and water content in each
step. The product formed in the first stage of the decomposition is a mixture of tenorite (CuO) and material with X-ray characteristics similar to
azurite, ruling out reaction sequences involving malachite or CuCO3. From structural considerations, a decomposition mechanism is proposed which is
consistent with the observed intensity changes in the X-ray pattern of the azurite-like intermediate phase."
Note, the study was conducted in helium gas, and I would guess that more dated research was not so carefully controlled. Still, interestingly, at one
stage a mixture containing CuO is referenced.
Link: http://www.sciencedirect.com/science/article/pii/00406031848...
Another source states:
"Copper Carbonate has a fairly complex decomposition. The accompanying curve shows the history of weight loss as this material is fired (courtesty of
Bob Hickerson, World Metal, LLC). It is interesting to compare this chart with the one for Copper Hydroxide to see the difference in the amount of
weight lost, and when and how fast it occurs."
Link: http://digitalfire.com/4sight/material/copper_carbonate_basi...
An educational sight which I suspect to be less accurate, nevertheless makes this interesting observation:
"The colour can vary from bright blue to green, because there may be a mixture of both copper carbonate and basic copper carbonate in various stages
of hydration. It was formerly much used as a pigment, and is still in use for artists colours. "
Link: http://www.rsc.org/learn-chemistry/resource/rws00013799/copp...
Also, here is an interesting patent that claims, to quote:
"The basic copper carbonate produced has the formula: (CuCO3)x(Cu(OH2)y, where y is 1 and x is between 0.1 to less than 1; or where y is 1 and x is 1,
or where y is 1 and x is between 0.5 to less than about 0.95, or where y is 1 and x greater than 1."
Reference: Direct synthesis of copper carbonate, US 7411080 B2
Link: http://www.google.com/patents/US7411080
The following comment concerns the natural transformation of azurite into malachite (so a weighted average of the two compounds can arise even
naturally):
"A pseudomorph of malachite after azurite retains the same shape as the original azurite crystal but is composed of malachite rather than azurite. The
pseudomorph is therefore malachite green in color rather than azurite blue.
The chemical formula describing the inversion of azurite to malachite is:
2 [Cu(OH)2 • 2(CuCO3)] + H2O ----------> 3 [Cu(OH)2 • (CuCO3)] + CO2
2 azurite + water ------------------> 3 malachite + carbon dioxide
Mineral specimens containing only azurite, only malachite, and varying portions of each substance exist. The contrast between azurite's intense blue
and malachite's bright green is very pleasing to the eye. Samples in which the transformation process has begun but remains incomplete can therefore
be quite beautiful. "
Link: http://dave.ucsc.edu/myrtreia/specimens.html
Finally, here is a pdf that alludes to naturally occurring Zn/Cu carbonates, so a Zinc presence may also present an issue for dated research.
Link: https://www.google.com/url?sa=t&source=web&rct=j&...
[Edited on 28-1-2015 by AJKOER]
|
|
|
| Pages:
1
2 |
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
