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Author: Subject: Copper Citrate Shenanigans
ManBearSwine
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[*] posted on 27-4-2012 at 18:35
Copper Citrate Shenanigans


Today, I tried to prepare some copper(II) citrate, but the weirdest thing happened....
I made it by neutralizing citric acid with sodium hydroxide, then adding a solution of copper(II) chloride. I expected copper citrate to immediately precipitate, but it instead formed a deep blue-green color. The solution was at about 60 degrees celcius and 0.5 molar with respect to copper citrate. I had not looked up the exact solubility at the time, but had read that it was "sparingly soluble" in water, so I figured that I just had a fairly saturated solution and decided to boil it down until crystals began to form, then chill it to crystallize the product. However, when the solution came to a hard boil, small blue crystals quickly formed and after about fifteen seconds, I had a thick blue slurry. It is definitely not copper hydroxide; the texture is that of small, loose crystals as opposed to a gelatinous suspension. Also, I put a small amount on my hotplate at about 200 degrees, and it did not turn black due to formation of copper oxide. I triple checked my calculations and confirmed that all of my reagents were the anhydrous forms I used in the calculations. After searching the internet for a bit, I found that the solubility of copper citrate is 0.02 g/L, but 60 g/L in a 1:1 mix with sodium citrate. Even this is much lower than my solution, which was about 280 g/L. My theory: I had an incredibly supersaturated solution, which crystallized either from the bumping on the bottom of my beaker or from a small crystal formed from the vaporization of water in contact with the bottom of the beaker. I will repeat this on a smaller scale using copper sulfate to determine if chloride ions play a role, perhaps by forming some type of complex complex ion (containing both citrate and chloride). I do not think the formation of the tetrachlorocuprate ion alone would prevent the precipitation, since there is still a large percent of copper present as the aquo complex even in a concentrated copper chloride solution.
-EDIT-
I just tried it with copper sulfate. The solution was a deep blue color, and did not form a precipitate even when seeded with copper citrate crystals or boiled very hard. Additionally, no crystals appeared when the solution was cooled to 10 degrees. When mixed with an equal volume of alcohol, a dark blue goo formed, which is soluble in water, forming a solution of the same color as the original. When calcium nitrate is added to this solution, no precipitate forms, seeming to indicate that little or so sulfate is present....
I saw a similar "goo" once with copper salicylate, but it doesn't explain the compound's solubility in water. If it was copper citrate alone I had precipitated, it should not be soluble in water.

I apologize for my rambling... does anyone have any idea what I am observing?

[Edited on 4-28-2012 by ManBearSwine]
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Pyridinium
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[*] posted on 27-4-2012 at 20:45


A similar thing can happen with ferric chloride when citrate is present. The hydroxide is prevented from forming a precipitate.

I'd be interested in hearing some more the properties of this precipitate of crystals.
If possible recrystallize them a couple times and see if you get a yellow flame test in with the bluish-green. Sodium forms a compound citrate with copper, I think.
Or it could be something else.
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ManBearSwine
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[*] posted on 27-4-2012 at 21:52


The flame test was inconclusive for both products. They both puffed up into a carbonaceous foam when burnt and gave a bright orange flame without any hint of blue or green. The good news it that this seems to confirm that it was not copper hydroxide that I produced in the first batch. As far as the "goo" goes, it is very sticky and plastic.
What I really don't understand, though, it if I have a 3:2 ratio of copper ions to citrate ions in solution, how could it form a complex? I suppose it would be possible with chloride ions, but I've never heard of sulfate acting as a ligand, so I don't know how to explain the goo. If the goo is a complex, then it must have sodium in it to be neutrally charged. Now, I will burn some more of it, then leech it with water and see if I get a yellow flame.
-EDIT- The ashes do appear to contain sodium. It is difficult to tell whether the compound itself contains sodium or sodium is just an impurity. Perhaps tomorrow I will try again.

[Edited on 4-28-2012 by ManBearSwine]
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Pyridinium
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[*] posted on 28-4-2012 at 07:58


The behavior of that substance to the flame is consistent with a complex compound being decomposed. It might go like this. At first the sodium does not vaporize to color the flame yellow, and neither does the copper participate. The initial stage is burning off of the organics. Once you get them to the ash stage, the flame test would be predominantly yellow, because sodium obscures pretty much everything else.

I believe there are sodium-copper-citrate complexes that are polymeric, because I know tartrate can do that. They would have water in their structure, so they could redissolve. I don't know why you're not getting those with the copper sulfate, but the chloride might be as you say forming a "complex complex ion". Or it might be creating conditions that favor the formation of these polymeric complexes, where the copper sulfate would not.

The presence of sodium, and the polymeric complexes, could explain why it's happening even though your molar ratio of copper doesn't seem "right". Just a theory, of course.

Going back to what I said about the ferric chloride, it would seem not just "any" iron compound behaves quite the same way, although citrate definitely forms complexes with iron. So there would be some interaction with chloride.

Edit:
Quote: Originally posted by ManBearSwine  
If the goo is a complex, then it must have sodium in it to be neutrally charged.


Good observation, and consistent with that theory of sodium / polymeric complexes. These would indeed be electrically neutral and capable of coming down out of solution.




[Edited on 28-4-2012 by Pyridinium]
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rannyfash
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[*] posted on 28-4-2012 at 08:30


what ph is the solution at when the crystals should be precipitating?
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nezza
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[*] posted on 29-4-2012 at 00:07


I think you are basically making a variant of Benedicts solution. With citrate copper II forms a complex which remains in solution in alkaline conditions. This complex is used as a test for reducing sugars similar to Fehlings solution.
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Eddygp
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[*] posted on 29-4-2012 at 00:45


maybe some copper hydroxide made apart from what nezza and Pyridinium said.



there may be bugs in gfind

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AJKOER
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[*] posted on 29-4-2012 at 05:31


If you are still looking to prepare Copper Citrate, may I suggest adding Calcium Citrate (from Ca(OH)2 + Citric acid) to Copper Sulfate (CuSO4).

Per your observations with Sodium Citrate (from NaOH + Citric acid) and CuSO4, if the Copper Citrate continues to be soluble is a positive as the Calcium sulfate probably will not.

If this fails, try dissolving Copper Oxalate (from the action of Copper and H2C2O4) in a Sodium oxalate solution to which you add Calcium Citrate. This will form a mixture of Copper and Sodium Citrate and a precipitate of Calcium oxalate. For more on the solubility characteristics of Calcium and Copper Oxalate see "A dictionary of chemistry and the allied branches of other sciences, Volume 4, by Henry Watts, pages 254 and 257.

Link:
http://books.google.com/books?pg=PA255&lpg=PA250&dq=...
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bquirky
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[*] posted on 29-4-2012 at 11:34


I have found copper citrate usefull for copper electroplating onto steel without ending up with a powdery mess.

My understanding of why this is so is that the copper chelate complex slows the reaction with iron down enough for larger copper crystals to form on the surface. once ive got a thin uniform layer of copper i can switch to a standard copper sulfate bath.

citric acid allso has the advantage of being available at the supermarket which is always a nice property for a reagent to have :)

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ldanielrosa
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[*] posted on 29-4-2012 at 12:47


http://www.lohmann-chemikalien.de/index.php/enhancing-the-so...

So it looks like the sodium citrate route is not good due to the enhanced copper citrate solubility.

@AJKOER, copper citrate should be poorly soluble on it's own so the calcium citrate route may precipitate both salts.

Interesting puzzle. I'll look into potassium and lithium citrate routes.
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AJKOER
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[*] posted on 29-4-2012 at 14:18


Quote: Originally posted by ldanielrosa  
http://www.lohmann-chemikalien.de/index.php/enhancing-the-so...

So it looks like the sodium citrate route is not good due to the enhanced copper citrate solubility.

@AJKOER, copper citrate should be poorly soluble on it's own so the calcium citrate route may precipitate both salts.


May I suggest Calcium citrate and Copper chloride (from CuSO4 + CaCl2) forming Copper citrate and CaCl2.
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99chemicals
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[*] posted on 29-4-2012 at 14:55


I would think that the best way would be some copper carbonate or hydroxide and add it to citric acid until no more dissolves. That way would seem to be the most simple.
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