Sciencemadness Discussion Board - copper carbonate synth.

copper carbonate synth.

maliveline - 16-12-2014 at 20:13

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 - 16-12-2014 at 20:27

You equation isn't balanced. Look at the the hydrogen in your NaHCO₃.

maliveline - 16-12-2014 at 20:48

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 - 17-12-2014 at 08:43

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 - 17-12-2014 at 09:46

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 - 17-12-2014 at 10:52

@ 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: CuCO3 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 - 18-12-2014 at 13:32

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 - 18-12-2014 at 13:43

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: CuCO3 doesn't exist. The two main copper basic carbonates are Malachite, Cu2CO3(OH)2,and Azurite, Cu3(CO3

2(OH)2.

Without washing your product is impure. Try digesting the precipitate by allowing it to stand in a warmish place overnight, before filtering.

maliveline - 18-12-2014 at 16:08

ya i feel ya thanks.

ChemSwede - 30-12-2014 at 10:37

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 - 30-12-2014 at 11:19

Quote: Originally posted by ChemSwede

Any other improvements I could make?

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 - 30-12-2014 at 14:43

Quote: Originally posted 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 - 30-12-2014 at 15:33

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 - 31-12-2014 at 07:38

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 - 1-1-2015 at 10:10

Quote: Originally posted 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 - 2-1-2015 at 10:59

@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 - 3-1-2015 at 06:01

@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 - 3-1-2015 at 06:13

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)4-. 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 Cu2CO3(OH)2(s) + 2 OH-(aq) === > Cu(OH)4-(aq) + 1/2 CO32-(aq)

This adequately explains your blue filtrate.

[Edited on 3-1-2015 by blogfast25]

ChemSwede - 3-1-2015 at 09:56

Quote: Originally posted by blogfast25

The blue filtrate is almost certainly a weak cuprate solution: Cu(OH)4-. 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 Cu2CO3(OH)2(s) + 2 OH-(aq) === > Cu(OH)4-(aq) + 1/2 CO32-(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 - 3-1-2015 at 13:40

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(H2O)6]2+ 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 - 26-1-2015 at 11:53

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 - 26-1-2015 at 17:06

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 - 26-1-2015 at 18:44

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 - 27-1-2015 at 06:49

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 - 27-1-2015 at 17:53

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]

blogfast25 - 28-1-2015 at 06:47

Repeat:

What is "copper ammonium hydroxide"? Formula, how to prepare it and what evidence for its existence, please.

Quote: Originally posted by AJKOER

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...
Still, interestingly, at one stage a mixture containing CuO is referenced.

[Edited on 28-1-2015 by AJKOER]

What 'part of the answer is explained'?

If you calcine malachite (or azurite) at some point you're going to hit CuO, like night follows day.

The other sources you mention are so low level it's unreal. They still crap on about 'mixtures'. Compounds with composition like xCuCO3.yCu(OH)2 (x, y are positive integers) are compounds, probably double salts, NOT mixtures.

AJKOER - 28-1-2015 at 09:19

Currently more popular name (but I have issues with its accuracy, discussed below, so I avoid it) for the hydroxide associated with the copper ammonium complex, for example, is presented in Wikipedia http://en.m.wikipedia.org/wiki/Schweizer%27s_reagent as follows, to quote:

"Schweizer's reagent is the chemical complex tetraamminediaquacopper dihydroxide, [Cu(NH3)4(H2O)2](OH)2. It is prepared by precipitating copper(II) hydroxide from an aqueous solution of copper sulfate using sodium hydroxide or ammonia, then dissolving the precipitate in a solution of ammonia."

My problem is that apparently there is nothing especially special about the tetraamminediaquacopper (II) cation, [Cu(NH3)4(H2O)2]2+, as one can actually create (and I do per my galvanic formation of the complex from elemental copper, dilute ammonia and H2O2) in solution any of [Cu(NH3)4(H2O)2]2+, [Cu(NH3)3(H2O)3]2+, [Cu(NH3)2(H2O)4]2+, or [Cu(NH3)(H2O)5]2+. My search for my prior thread with references, produced the following in the Energetic Materials section of SM:

Quote: Originally posted by AJKOER

I noticed recently that dissolving Cu in aqueous ammonia in the presence of H2O2 and NaCl appears to form the insoluble Copper Ammonium chloride, [Cu(NH3)4(H2O)2]Cl2....
...see "Kinetics and Mechanism of Copper Dissolution In Aqueous Ammonia" available online) .....
Note, as a point of interest, there is actually nothing sacred about the tetraamminediaquacopper cation as detpending on the ammonia concentration in solution, one could have anywhere from [Cu(NH3)(H2O)5]2+ (the greenish-blue thin layer) to [Cu(NH3)5(H2O)]2+ (a royal blue complex), with the latter occurring in very concentrated ammonia solutions. Reference: https://docs.google.com/viewer?a=v&q=cache:IjHK0vuBZhcJ:...
..

Also, some source state that [Cu(NH3)4(H2O)2](OH)2 is formed only with strong ammonia. This statement is also misleading, in my opinion, as I have at times seem layers of layers of color in my galvanic cell of dilute ammonia, including royal blue around the copper at the base of the cell. This anodic zone (the copper source may be several unconnected pieces) is where the oxidation of copper proceeds according to:

Cu + 4 NH3 + 2 H2O --) [Cu(NH3)4(H2O)2]2+ + 2 e-

Reference: "Kinetics and Mechanism of Copper Dissolution In Aqueous Ammonia", at https://www.google.com/url?sa=t&source=web&rct=j&...

So, bottom line, the more technical name possibly implying to some, that it is "the" copper ammine complex is misleading, especially, in my case, if I am recommending an electrochemical approach to the complex employing dilute ammonia. Now, the older naming convention (which I employ) is not without issues as it can create confusion when a double salt of ammonia is also possible, but this is not the case for the hydroxide.
-------------

To answer your other comment, I repeat a direct quote, provide above with link, from an obvious authority, in my opinion, on minerals where the words "and varying portions" is actually applied to natural occurring basic copper carbonate crystals, that are apparently quite beautiful:

"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."

[Edited on 28-1-2015 by AJKOER]

blogfast25 - 28-1-2015 at 11:03

Quote: Originally posted by AJKOER

To answer your other comment, I repeat a direct quote, provide above with link, from an obvious authority, in my opinion, on minerals where the words "and varying portions" is actually applied to natural occurring basic copper carbonate crystals, that are apparently quite beautiful:

"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."

All of this is well known and discussed ad nauseam in the relevant copper basic hydroxide thread in 'beginners':

http://www.sciencemadness.org/talk/viewthread.php?tid=50822&...

Any significant contributions are best parked there.

Wonderful photos of malachite/azurite combinations here:

http://www.bing.com/images/search?q=azurite+malachite&FO...

How does the copper tetrammine complex relate to all this? What diabolical plan do you have in mind?

[Edited on 28-1-2015 by blogfast25]

AJKOER - 28-1-2015 at 11:20

Actually, I am a few years too late, as what I am suggesting has been recently patented. To quote from the patent, "Direct synthesis of copper carbonate", US 7411080 B2, Link: http://www.google.com/patents/US7411080 :

"...What is needed is a method of directly forming a variety of basic copper salts from copper metal.

SUMMARY OF THE INVENTION
In a broad aspect the invention encompasses a novel method of synthesizing selected basic copper salts including basic copper carbonate, basic copper sulfate, and basic copper acetate directly from copper metal using water, an amine (preferably ammonia), an oxidizing source (preferably oxygen), and a source of anions. Advantageously the reaction takes place in a single reactor. Advantageously the oxidant is a gas comprising oxygen, such as oxygen, an oxygen/inert gas mixture such as oxygen and nitrogen, air, or any combination thereof. .."

although, I can up with the idea to aid home chemists in their quest for copper salts and such.

[Edited on 28-1-2015 by AJKOER]

blogfast25 - 28-1-2015 at 12:01

From that patent:

Quote:

EXAMPLE 1
A 500-ml beaker with magnetic stir bar was charged with about 300 ml water and 3 g of ammonia (0.176 moles of NH4OH, added in the form of 10 grams of concentrated aqueous ammonium hydroxide). CO2 gas was bubbled into the solution with agitation until a pH of 8.5 was obtained. The solution was heated on a hot plate and maintained at temperature of about 50° C. Approximately 20 g (0.315 moles) of copper powder (˜325 mesh) was added to the beaker under good agitation. The molar ratio of copper to ammonia was 1.8, and the copper loading was 66.6 grams per liter of water. Air was introduced at the same time. During the course of the reaction the pH increased gradually, so CO2 gas was occasionally introduced to maintaining a pH of around 8.5. After about 10 hrs no metallic copper was visible, and a slurry of green BCC precipitate was obtained, which held a stable pH of 8.50 without need for further CO2 addition. No metallic copper was observed. The reaction mixture was filtered to yield a green solid cake and a near colorless filtrate. The cake, after drying in a 50° C. oven overnight, was found to contain 56.2% copper. The theoretical copper content of basic copper carbonate is 57.48% but commonly 56.0% is as high as conventionally practical to manufacture. The weight mean particle size (“WMPS”, the diameter at which half of the weight of material is present as particles having an effective diameter which is less than the weight mean particle size) of the basic copper carbonate precipitate was 15.2 microns and the packed bulk density was 0.70 lbs/ft3. The particle size was determined using a MicroTrac™ S3500/S3000 laser scattering device. The packed bulk density “PABD” was determined using a JEL Stampfvolumeter™ STAV 2003.

I'm more than willing to believe that. Ammonia acts as a catalyst here because it speeds up the oxidation/solubilisation of Cu(0) to Cu(II).

It's rather funny though how much effort they put into a physical characterisation of their product and how little into a chemical characterisation! At a very minimum a CO2 and O content determination should be called for.

And of course like all patents it's riddled with *rse covering nonsense by keeping x and y deliberately vague and wide.

Copper metal is also less economical than using copper salts and Azurite this method will not make.

[Edited on 28-1-2015 by blogfast25]

zed - 1-2-2015 at 14:46

Skinflint that I am, I recently used this reaction to treat some wood. The treated wood at Home Depot is just too damned expensive.

I Just rolled some concentrated Copper Sulfate Solution onto some boards with a paint roller, and let 'em soak it in for a bit. The boards showed little apparent color change.

Then, I dumped a box of Bicarbonate into a bucket of water, dissolved it well, and rolled the boards again. The boards took on a few green highlights, but no big deal. Thereafter, I set them out, for a few days to let them dry a bit.

Coated them with a deck sealer, let 'em dry again for a few days, and observed they were continuing to "Green".

By the time, I put a final coat of paint on them, they were looking very green indeed. In the grain of the wood, it was taking a little while for the copper ions and carbonate ions to co-operate into their insoluble magic.

So, there it is....my new front porch. Even done to code, which requires that such exposed woodwork be constructed of rot resistant, or treated wood, but does not stipulate method of treatment.

I suppose, I'll know how well the treatment actually worked, in twenty-five years or so. If the porch is still present, it was a success. If not, well...there is a decent chance I will be too dead to care.

[Edited on 1-2-2015 by zed]