DrMario
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What is this? (copper inorganic compound)
This is how I obtained the content of that jar: I made a saturated aqueous solution of sodium hydrogencarbonate (NaHCO3) and put some
random pieces of copper into it. These were various copper plates and wires. All reasonably pure, but all, more or less, covered by cupric oxide.
I closed the jar and left it there for several weeks, maybe a month or more. The copper plates became even darker, but the solution became this bright
blue.
What do you think this might be?
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Endo
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I suggest this would be a good post for a beginnings question.
The light blue solution is typical of the copper 2+ ion in solution. The blue color is due to the formation of the complex ion [Cu(H2O)4]2+. Tough
to see what the precipitate looks like. If it is gel like it is probably a copper hydroxide or if not more likely a copper carbonate, or a mix of the
two.
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DrMario
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Quote: Originally posted by Endo  | I suggest this would be a good post for a beginnings question.
The light blue solution is typical of the copper 2+ ion in solution. The blue color is due to the formation of the complex ion [Cu(H2O)4]2+. Tough
to see what the precipitate looks like. If it is gel like it is probably a copper hydroxide or if not more likely a copper carbonate, or a mix of the
two.
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First of all, the precipitate is just undissolved NaHCO3
It's not a beginner's question, since not even you know what exactly is going on. Yes, the cation is clearly Cu2+ - and
then?
[Edited on 10-10-2014 by DrMario]
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hyfalcon
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If you don't have references on the chemistry, then it belongs in the beginners section.
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Endo
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Hmmm,
Your question was
| Quote: |
The copper plates became even darker, but the solution became this bright blue. What do you think this might be? </quote>
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I felt I answered your question clearly. The bright blue color is the complex ion of Cu2+ in solution.
You will need to specify another question if you want more information. -- <b> and then?</b> is a bit of a mystery to me. For now you
have a soup of ions in solution: The copper complex ion [Cu(H2O)4]2+, Na+, and because of equilibrium a range of different carbonate polyatomic ions
in various proportions, and water. It appears you have removed the copper metal.
So, in a closed container, with nothing to drive the reaction that formed the Cu2+ in the first place (no copper metal), this jar will remain sky blue
on your shelf doing very little until you change the conditions and drive an equilibrium in one direction or the other by adding something to the
reaction mix. So your answer to "and then?" is very little.
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AsocialSurvival
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Copper (2+) carbonate. Air oxidizes Copper, and formed Copper Hydroxide reacts with Sodium bicarbonate to form Sodium Carbonate and Copper Carbonate.
You can make your solution even more dark blue and concentrated if you add Carbon Dioxide to it, because Copper bicarbonate will form which is much
more soluble.
For example only 10 mg of Calcium dissolves per liter of water in the form of carbonate, but 95 grams of Ca dissolve in same amount of water if
bicarbonate.
Test it by drying, the remaining powder will look green, not blue, trust me!
[Edited on 10-10-2014 by AsocialSurvival]
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Amos
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Sodium bicarbonate dissociates, creating a low concentration of OH-. Because copper(II) oxide and copper(II) hydroxide are both amphoteric, they will
dissolve in greater amounts than usual with more OH- present. Don't be deceived by the color, though; it's unlikely you have a very significant amount
of copper(II) in solution. It's still overwhelmingly sodium carbonate/bicarbonate.
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DrMario
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| Quote: Originally posted by AsocialSurvival | Copper (2+) carbonate. Air oxidizes Copper, and formed Copper Hydroxide reacts with Sodium bicarbonate to form Sodium Carbonate and Copper Carbonate.
You can make your solution even more dark blue and concentrated if you add Carbon Dioxide to it, because Copper bicarbonate will form which is much
more soluble.
For example only 10 mg of Calcium dissolves per liter of water in the form of carbonate, but 95 grams of Ca dissolve in same amount of water if
bicarbonate.
Test it by drying, the remaining powder will look green, not blue, trust me!
[Edited on 10-10-2014 by AsocialSurvival] |
Thanks a lot.
What is the reason these bicarbonates are so much more soluble than carbonates? Also, I was under the impression that copper carbonate was practically
insoluble in water - so is it possible that there already is copper bicarbonate in the solution?
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DrMario
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| Quote: Originally posted by No Tears Only Dreams Now | | Sodium bicarbonate dissociates, creating a low concentration of OH-. Because copper(II) oxide and copper(II) hydroxide are both amphoteric, they will
dissolve in greater amounts than usual with more OH- present. Don't be deceived by the color, though; it's unlikely you have a very significant amount
of copper(II) in solution. It's still overwhelmingly sodium carbonate/bicarbonate. |
OK, thanks, this answers my question quite thoroughly!
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Amos
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| Quote: Originally posted by DrMario | | Quote: Originally posted by No Tears Only Dreams Now | | Sodium bicarbonate dissociates, creating a low concentration of OH-. Because copper(II) oxide and copper(II) hydroxide are both amphoteric, they will
dissolve in greater amounts than usual with more OH- present. Don't be deceived by the color, though; it's unlikely you have a very significant amount
of copper(II) in solution. It's still overwhelmingly sodium carbonate/bicarbonate. |
OK, thanks, this answers my question quite thoroughly! |
Also, I don't know that copper bicarbonate exists by any means, even in solution. I think bicarbonate is restricted to group 1 and 2 elements.
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AsocialSurvival
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| Quote: Originally posted by DrMario | | Quote: Originally posted by AsocialSurvival | Copper (2+) carbonate. Air oxidizes Copper, and formed Copper Hydroxide reacts with Sodium bicarbonate to form Sodium Carbonate and Copper Carbonate.
You can make your solution even more dark blue and concentrated if you add Carbon Dioxide to it, because Copper bicarbonate will form which is much
more soluble.
For example only 10 mg of Calcium dissolves per liter of water in the form of carbonate, but 95 grams of Ca dissolve in same amount of water if
bicarbonate.
Test it by drying, the remaining powder will look green, not blue, trust me!
[Edited on 10-10-2014 by AsocialSurvival] |
Thanks a lot.
What is the reason these bicarbonates are so much more soluble than carbonates? Also, I was under the impression that copper carbonate was practically
insoluble in water - so is it possible that there already is copper bicarbonate in the solution? |
Yes, you can't have any bicarbonate as long as you have one hydroxide (in case of elements with oxidation states +1 and +2), but it can be formed by
absorbing additional CO2 from air. You can imagine those two chemicals seperately, after long period of time, both Sodium hydroxide and Copper
hydroxide will become bicarbonates, because they absorb it from air.
But don't try to get dry Copper bicarbonate, because it only exists in solution, just like Calcium bicarbonate (the alkali metals however can be
isolated dry).
Why?
Well, because Carbon dioxide forms acid with water, and obviously the more CO2 is the more acid is formed, just like adding more NO2 to water to make
more Nitric acid.
The easiest and best way to make bicarbonate is to hold that solution in closed bottle full of CO2 above it, and then periodically add more CO2. That
will raise partial pressure of CO2 which will dissolve more Copper as bicarbonate. However, if you start from Copper metal, you will still need access
to Oxygen so it can oxidize Copper. best to do this method on already made Copper hydroxide or carbonate.
Practicaly, mother nature did that so that you can grow plants 2 times faster than usual, if you increase CO2 concentration in air which is neccessary
for photosynthesis, it will absorb more nutrients from soil in needed proportion.
[Edited on 10-10-2014 by AsocialSurvival]
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blogfast25
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| Quote: Originally posted by No Tears Only Dreams Now | | Sodium bicarbonate dissociates, creating a low concentration of OH-. Because copper(II) oxide and copper(II) hydroxide are both amphoteric, they will
dissolve in greater amounts than usual with more OH- present. Don't be deceived by the color, though; it's unlikely you have a very significant amount
of copper(II) in solution. It's still overwhelmingly sodium carbonate/bicarbonate. |
Is the right answer. Copper is slightly amphoteric and in alkaline conditions forms cuprate anions:
Cu(OH)<sub>4</sub><sup>-</sup>(aq). These are intensely deeply blue coloured and even just a bit of them will colour your
solution light blue. As 'TNODN' claims: there isn't much copper that actually entered the solution.
Copper carbonates or bicarbonates are irrelevant here because the copper is solvated as anions. It's a slow process and as the OP wrote it took weeks.
As a practical way of dissolving significant amounts of copper, forget it.
[Edited on 10-10-2014 by blogfast25]
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DrMario
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Thanks guys/gals.
Before I get rid of the contents of this jar, is there any experiment worthwhile doing?
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blogfast25
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| Quote: Originally posted by DrMario | Thanks guys/gals.
Before I get rid of the contents of this jar, is there any experiment worthwhile doing? |
Take a small amount and slowly neutralise it with strong HCl (lots of bubbles). At the end of that you should have a faintly green solution of cupper
(II) chloride.
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xfusion44
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What would I get if I were to connect power to the copper leads in DrMario's case? (NaHCO3, water, copper leads) I assume that I would get copper
carbonate, but faster, right?
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unionised
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A solution of bicarbonate left lying around won't have a lot of hydroxide in it, but it will contain carbonate ions.
Coper forms a complex ion with carbonate and My guess is that's the most likely cause of the blue colour.
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AsocialSurvival
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| Quote: Originally posted by xfusion44 | | What would I get if I were to connect power to the copper leads in DrMario's case? (NaHCO3, water, copper leads) I assume that I would get copper
carbonate, but faster, right? |
Right! Electric current speeds up oxidation!
[Edited on 12-10-2014 by AsocialSurvival]
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unionised
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After a while the copper would dissolve from one electrode and be plated out at the other. The net concentration of copper in the solution would
remain unchanged.
Not really very useful unless you are seeking to electrolytically refine copper.
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DrMario
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| Quote: Originally posted by blogfast25 | | Quote: Originally posted by DrMario | Thanks guys/gals.
Before I get rid of the contents of this jar, is there any experiment worthwhile doing? |
Take a small amount and slowly neutralise it with strong HCl (lots of bubbles). At the end of that you should have a faintly green solution of cupper
(II) chloride. |
I appreciate the idea, but the outcome seems a bit too obvious. I'd rather do an experiment that determines the anions present in the solution, or
perhaps something that would determine the exact complex. I do have some EDTA-Na 2 available. And of course ammonium hydroxide, the Big
Three acids etc.
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AsocialSurvival
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| Quote: Originally posted by unionised | After a while the copper would dissolve from one electrode and be plated out at the other. The net concentration of copper in the solution would
remain unchanged.
Not really very useful unless you are seeking to electrolytically refine copper. |
Not really! That is only correct for more soluble Copper salts. In this case, Hydrogen would reduce and only a fraction of Copper at a time (which is
present in microgram to miligram quantities in solution). And Copper would rather get oxidized than Hydroxide aka Oxygen from water.
I tried this many times, and collected lots of Copper carbonate (blueish green or greenish blue in solution, green when dry).
[Edited on 12-10-2014 by AsocialSurvival]
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Amos
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| Quote: Originally posted by DrMario | | Quote: Originally posted by blogfast25 | | Quote: Originally posted by DrMario | Thanks guys/gals.
Before I get rid of the contents of this jar, is there any experiment worthwhile doing? |
Take a small amount and slowly neutralise it with strong HCl (lots of bubbles). At the end of that you should have a faintly green solution of cupper
(II) chloride. |
I appreciate the idea, but the outcome seems a bit too obvious. I'd rather do an experiment that determines the anions present in the solution, or
perhaps something that would determine the exact complex. I do have some EDTA-Na 2 available. And of course ammonium hydroxide, the Big
Three acids etc.
|
I mean, what you have there is pretty straightforward. Just because it's changed in color doesn't mean the composition has changed much at all, and I
doubt you have significant enough amounts of copper in there to do much experimentation on in the realm of solid compounds, but complexing in solution
can be explored. Adding HCL or even just salt will make it turn green, adding ammonium hydroxide will produce a precipitate of Cu(OH)2 in
small amounts and in larger ammounts form the tetraammine copper(II) complex, which is ultramarine blue. If both of these happen with the same
solution, you can pretty much bet that what's in solution is the [Cu(H2O)6]2+ complex, as the H2O ligands
very readily detach themselves to make room for others. This is the same thing you usually get with solutions of copper(II) salts. If you had some new
and different complex, it might not react so readily.
[Edited on 10-12-2014 by No Tears Only Dreams Now]
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xfusion44
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| Quote: Originally posted by AsocialSurvival | | Quote: Originally posted by xfusion44 | | What would I get if I were to connect power to the copper leads in DrMario's case? (NaHCO3, water, copper leads) I assume that I would get copper
carbonate, but faster, right? |
Right! Electric current speeds up oxidation!
[Edited on 12-10-2014 by AsocialSurvival] |
I'm asking that, because I already done this and I could copper plate key
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