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kronos
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what copper compound did I make
Well I just made a little cucl2 using copper wires, hydrogen peroxide (6%) and dilute hydrochloric acid.. All the copper dissolved.. And the turned
a beautiful light blue color.. Then i added excess of sodium carbonate solution.. It bubbled Little which was because of the excess HCl reacting.. A
greenish blue precipitate was formed.. I dried it and after drying i did a few tests... Oje of the tests was adding conc. HCl to it.. If it was
copper carbonate then it should have bubbled.. But it didn't bubble at al. Then what did I prepare? 
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Molecular Manipulations
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How did you dry it? Copper hydroxycarbonate is not very stable, it should be dried at room temperature with a fan on to keep it from decomposing to
copper oxide. Was it still blue when you tried to dissolve it in hydrochloric acid?
-The manipulator
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kronos
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Yes. It was bluish green and i dried it at room temperature. It was drying at normal temperature Outside without Any fan though. It was totally
bluish green no trace of blaack copper oxide in it
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Molecular Manipulations
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Strange. How much acid did you add? Perhaps the acid reacted with the hydroxy group before the carbonate. I can't think of anything else,
Blogfast is the expert of copper hydroxycarbonates.
-The manipulator
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blogfast25
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Just precipitating copper basic carbonate from a neutral solution of CuCl2 will liberate CO2:
2 CuCl2 + 2 Na2CO3 + H2O == > Cu2CO3(OH)2 + 4 NaCl + CO2
'Straight' copper carbonate ('CuCO3') doesn't exist.
Basic copper carbonate (aka Malachite) precipitates as a blue green precipitate, quite different to Cu(OH)2 in texture and colour.
Synthetic Malachite dries to a green powder and resists dehydration/decarbonisation well, up to about 200 C.
I think you should repeat the experiment. If the amount of Malachite used was small then somehow 'missing' the CO2 when treated with HCl is easier
than you might think. Experimental observing is more difficult than you might assume.
When you added the conc. HCl to the Malachite, did it dissolve?
[Edited on 26-2-2015 by blogfast25]
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kronos
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Quote: Originally posted by blogfast25  | Just precipitating copper basic carbonate from a neutral solution of CuCl2 will liberate CO2:
2 CuCl2 + 2 Na2CO3 + H2O == > Cu2CO3(OH)2 + 4 NaCl + CO2
'Straight' copper carbonate ('CuCO3') doesn't exist.
Basic copper carbonate (aka Malachite) precipitates as a blue green precipitate, quite different to Cu(OH)2 in texture and colour.
Synthetic Malachite dries to a green powder and resists dehydration/decarbonisation well, up to about 200 C.
I think you should repeat the experiment. If the amount of Malachite used was small then somehow 'missing' the CO2 when treated with HCl is easier
than you might think. Experimental observing is more difficult than you might assume.
When you added the conc. HCl to the Malachite, did it dissolve?
[Edited on 26-2-2015 by blogfast25] |
yes it did dissolve but it did not bubble at all.. It dissolved slowly
without any bubbles
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Molecular Manipulations
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Try it again, in the right proportions. It was probably some experimental problem.
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blogfast25
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Precisely.
Principle 101 in empiricism: 'repeat' failures.
If failure persists, you may be on to something.
Repeat but better, for example by carefully monitoring all quantities, so a stoichiometrical evaluation can be made.
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Molecular Manipulations
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If you have the reagents, experiment from different angles. Try copper sulfate as well as chloride, use sodium bicarbonate and carbonate, and try it
at different temperatures, concentrations and addition rates. Test one new variable at a time, and use contols to note the differences.
The more variables you individually introduce the more you can learn.
-The manipulator
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Amos
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...what? Basic copper carbonate is plenty stable! Unless you consider temperatures of over 200 celsius to be a normal environment, copper carbonate
isn't going to bother decomposing. Hence why all malachite samples aren't dusty black...
[Edited on 2-27-2015 by Amos]
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blogfast25
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| Quote: Originally posted by Amos | ...what? Basic copper carbonate is plenty stable! Unless you consider temperatures of over 200 celsius to be a normal environment, copper carbonate
isn't going to bother decomposing. Hence why all malachite samples aren't dusty black...
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This is in response to what or whom?
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gdflp
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| Quote: Originally posted by blogfast25 | | Quote: Originally posted by Amos | ...what? Basic copper carbonate is plenty stable! Unless you consider temperatures of over 200 celsius to be a normal environment, copper carbonate
isn't going to bother decomposing. Hence why all malachite samples aren't dusty black...
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This is in response to what or whom? |
I believe that it was in response to this. | Quote: Originally posted by Molecular Manipulations | | How did you dry it? Copper hydroxycarbonate is not very stable, it should be dried at room temperature with a fan on to keep it from decomposing to
copper oxide. Was it still blue when you tried to dissolve it in hydrochloric acid? |
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Molecular Manipulations
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Probably me:
I said that because some people don't know how unstable some transition metal carbonates are. Something that decomposes at 200 deg.C isn't what I'd
call very stable. Also, that 200 deg. C isn't when it starts decomposing, it's just when the rate is high enough to be practical. I have experience
with that as I once dried it overnight in an oven at 110 deg. C and after 13 hours, I measured over 30% had decomposed.
-The manipulator
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Amos
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| Quote: Originally posted by Molecular Manipulations | Probably me:
I said that because some people don't know how unstable some transition metal carbonates are. Something that decomposes at 200 deg.C isn't what I'd
call very stable. Also, that 200 deg. C isn't when it starts decomposing, it's just when the rate is high enough to be practical. I have experience
with that as I once dried it overnight in an oven at 110 deg. C and after 13 hours, I measured over 30% had decomposed. |
When someone describes a compound as unstable, I'm not sure decomposing at 200C is what they mean. Drying overnight in the oven gave you a product of
what color? And how can you be so sure that it decomposed if you were putting it in the oven to remove an unknown amount of water?
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diddi
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I think a vis/UV would be needed to demonstrate whether decomposition has occurred. and to me 200C is a bit harsh to be expecting stability.
Beginning construction of periodic table display
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blogfast25
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| Quote: Originally posted by diddi | | I think a vis/UV would be needed to demonstrate whether decomposition has occurred. and to me 200C is a bit harsh to be expecting stability.
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DSC or thermogravimetry would be used to measure any actual transformations of the dry material. There are some papers on copper basic carbonate with
that method.
At a home lab level determination of CO2 content of partially decomposed material could be used as a fairly crude measure of the degree of
decomposition.
[Edited on 27-2-2015 by blogfast25]
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Molecular Manipulations
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| Quote: Originally posted by Amos |
When someone describes a compound as unstable, I'm not sure decomposing at 200C is what they mean. |
Like I said, at less than 200 deg. C it begins to decompose, it's just quite slow.
| Quote: Originally posted by Amos | | how can you be so sure that it decomposed if you were putting it in the oven to remove an unknown amount of water? |
Simple my friend. The water content wasn't strictly unknown, the total copper ion was known, the solubility product of copper basic-carbonate was
known and the carbonate ion was in excess, thus the quantity of copper basic-carbonate was known. The only major source of error was effectiveness of
my filtering and not losing any product in transporting, which was easy, as I used only two vessels for the whole process.
Furthermore the color was quite dark and I tested the carbonate content by measuring the carbon dioxide produced from addition of acid. Taking into
account solubility of carbon dioxide, the equilibrium of of the reaction and the water vapor from collecting over water.
I'm pretty sure my accuracy was within 5%.
[Edited on 27-2-2015 by Molecular Manipulations]
-The manipulator
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clearly_not_atara
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My God, if it were a snake, you'd all be wriggling bulges in a snake-stomach because -- Jesus, take the wheel. Carlos, take the stereo.
You used too much HCl. The Na2CO3 you added was all decomposed on first addition and when you saw a precipitate it was because you oversaturated the
solution and precipitated CuCl2. That's why it didn't bubble in HCl, and it's also why you saw a...
| Quote: | | greenish blue precipitate was formed.. |
instead of, you know, this:
https://en.wikipedia.org/wiki/Patina#mediaviewer/File:Minnea...
Anywho, these crystals look familiar to you? I'm guessing you have a few sitting around.
https://en.wikipedia.org/wiki/Copper%28II%29_chloride#mediav...
[Edited on 3-3-2015 by clearly_not_atara]
[Edited on 3-3-2015 by clearly_not_atara]
[Edited on 3-3-2015 by clearly_not_atara]
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Amos
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I can't believe I didn't mention this earlier, you may have produced dicopper chloride trihydroxide(commonly called copper oxychloride), which can
result when an already acidic solution of copper chloride is neutralized with base. My method for producing it in fact involves adding dilute sodium
carbonate to an acidified solution of copper chloride and sodium chloride.
clearly_not_atara's theory doesn't seem likely, as it was clearly mentioned that only very dilute reagents were used, and copper
chloride is very soluble in water. The condescension isn't really necessary, either.
[Edited on 3-3-2015 by Amos]
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clearly_not_atara
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Copper hydroxy-chloride is a convenient possibility but it has the unfortunate property of being bright green. I am more willing to believe that he
got the concentration of a reagent wrong than that he can't see the difference between blue and green -- but I guess we can never really know.
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Zyklon-A
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I think Amos is right. I had forgotten that copper oxychloride is insoluble. If however kronos had added more acid
it would have formed copper (II) chloride and water again, but apparently he didn't.
What's that supposed to mean? kronos said blue-green:
Also copper oxychloride is blue-green, depending on the lighting, purity and water present it can even look more blue than green:
[Edited on 3-3-2015 by Zyklon-A]
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blogfast25
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Well done, 'Amos'.
Copper hydroxychloride (Cu<sub>2</sub>(OH)<sub>3</sub>Cl) is indeed what I should have thought of too! Doh!
Especially because I've prepared it myself, stoichiometrically:
4 CuCl2 + 3 Na2CO3 + 3 H2O === > 2 Cu2(OH)3Cl + 6 NaCl + 3 CO2
All basic Cu(II) salts (carbonate, chloride and sulphate) precipitate as sandy green blue precipitates which on drying turn fairly green.
It also fits observing no bubbles when dissolving the hydroxychloride in HCl.
'clearly_not_atara's' first link is to copper basic carbonate, the second to CuCl2, which is too soluble for 'kronos'' scenario.
In the case of these compounds the difference is subtle and quite subjective. Dry/wet, granulometry and lighting all influence that subjective
experience. And I think we now DO know: the explanation is simple and straightforward and fits the facts well.
[Edited on 3-3-2015 by blogfast25]
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Amos
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| Quote: Originally posted by clearly_not_atara | | Copper hydroxy-chloride is a convenient possibility but it has the unfortunate property of being bright green. I am more willing to believe that he
got the concentration of a reagent wrong than that he can't see the difference between blue and green -- but I guess we can never really know.
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Copper oxychloride exists as several polymorphs that have different structures, densities, and colors. The kind I've produced in a way similar to the
method being discussed is a blue-green when wet and later dries to a very pale blue-green, almost white.
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blogfast25
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| Quote: Originally posted by Amos |
Copper oxychloride exists as several polymorphs that have different structures, densities, and colors. |
That's a very interesting statement. Do you have any references backing that up?
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Amos
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http://en.wikipedia.org/wiki/Dicopper_chloride_trihydroxide#...
The four polymorphs listed there are atacamite, paratacamite, clinoatacamite, and botallackite. You can do a search of the minerals to see different
colors, they're just varying shades of green and blue, that's copper for you.
I was incorrect to use the word density; I was more referring to the volume and consistency of the chemical product formed through various routes; use
of sodium carbonate or bicarbonate has produced an extremely fine and chalky precipitate in my experience, while that formed using ammonia as the base
gave a more gelatinous and pale green-colored precipitate. But this difference could be due to the impurity of the result in each case; online
products labelled copper oxychloride often state a purity well below what one would consider lab grade.
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