RareEarth
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Difficulty forming metal-amino acid chelate, advice needed
In an attempt to form the stable hydrophilic amino acid chelate complex of NiCl2-2(Glycine), I mixed the two together in water, which they are both
soluble in, for 2 days at 25-c. The solution never changed color during this time, still being the green color of nickel. After that, I added sodium
carbonate to the solution and it immediately turned to the blue color of NiCO3, indicating that the nickel did not chelate to the glycine.
What was wrong in my reaction conditions for this complex not to form? The reactants dissolve in the same solvent, and the complex-formation is fairly
irreversible due to the chelate, so reaction conditions should have been prime for formation. Does anyone have available a write up on how to form
this chelates?
edit
Based on some reading I did, I did not think this would require elevated temperatures. I'm going to reflux for an hour to see what happens.
[Edited on 1-5-2015 by RareEarth]
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DraconicAcid
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I didn't think that NiCO3 was blue. If addition of sodium carbonate gave you a blue solution, that is undoubtedly the chelate.
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Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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RareEarth
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Quote: Originally posted by DraconicAcid  | | I didn't think that NiCO3 was blue. If addition of sodium carbonate gave you a blue solution, that is undoubtedly the chelate.
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When I added some pure NiCl2 (hexahydrate) to a solution of Sodium Carbonate, it turned the same blue color. I forgot to mention I did this to verify
whether or not the color was unique to the chelate. Could it be possible that both are coincidentally the same color? Perhaps it is possible that
Sodium Carbonate is coordinating with the chelate in solution? But is it really likely to have that same blue color as nickel carbonate? Chelates are
more often than not different colors than the color of the salt they chelate with.
NiCO3 is also the mineral Hellyerite, which is known to have a light blue color. http://en.wikipedia.org/wiki/Hellyerite
From the time I started this thread, I heated the nickel-glycine solution at 90 degrees for about 2 hours, and then set it out in a tray to evaporate.
When all of the water evaporates, presumably tomorrow, I will be able to compare the color of the remaining solid with the nickel chloride itself.
[Edited on 2-5-2015 by RareEarth]
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DraconicAcid
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But nickel carbonate is also going to be insoluble.
ETA http://en.wikipedia.org/wiki/Nickel%28II%29_carbonate says the carbonates (of various compositions) are light green, which goes along with my
experience. Generally octahedral nickel with six oxygen ligands will be green, with six nitrogen ligands will be purple (e.g., six ammonias or three
ethylenediamines), and with some of each (such as in the expected chelate) will be blue.
[Edited on 2-5-2015 by DraconicAcid]
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mnick12
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Try the dimethoxyethane complex of NiCl2 in something like ether or THF, the aqua may be competing.
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AvBaeyer
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You are not going to form the chelate from NiCl2 unless the chelate is extremely insoluble, which I doubt. The problem is that when glycine reacts
with NiCl2, a molecule of HCl is released each time. What happens is that your reaction reaches an equilibrium which does not favor the chelate. So
you need to start with a nickel compound that does not release acid when it combines with the glycine.
For example, the glycine complex of copper is formed as follows:
Dissolve 1 gram of glycine in 25 ml hot water. Add copper carbonate a little at a time until the reaction ceases. Heat the mixture to boiling , decant
the solution from remaining solids and allow to cool. Blue crystals of the copper complex will form. [From Robertson and Jacobs, Laboratory Practice
of Organic Chemistry]
Upshot, try using nickel carbonate in place of nickel chloride.
AvB
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RareEarth
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I agree with you NiCO3 should be insoluble, but that leaves the question of what is the blue-color that forms when I add NiCl2 to sodium carbonate and
vice versa?
I actually did notice that the PH of the solution of nickel and glycine is considerably lower than the PH of either the NiCl2 or Glycine alone. I
wonder if the complex formed, being acidic, and sodium carbonate reacted with it, deprotonating the free hydrogens and leaving the complex blue. That
could explain the PH I took and the solubility.
| Quote: Originally posted by AvBaeyer | You are not going to form the chelate from NiCl2 unless the chelate is extremely insoluble, which I doubt. The problem is that when glycine reacts
with NiCl2, a molecule of HCl is released each time. What happens is that your reaction reaches an equilibrium which does not favor the chelate. So
you need to start with a nickel compound that does not release acid when it combines with the glycine.
|
I think there is a bit of confusion here. There is a difference between chelated complexes of metals and amino acids, and the salts of amino acids
with those metals. What you are talking about is the formation of one of those salts (which involves HCl elimination). In the chelate complexes I am
referring to, HCl is not eliminated. It's a coordination complex that forms. These chelated coordination complexes bind very tightly together so there
isn't much reversible equilibrium taking place.
What I suspect is that the aquo complex is competing with formation of the chelate, and preventing it. One of the preps I recently read they performed
the reaction in acetic acid and concentrated ethanol to form a palladium amino acid complex, and the chelate preciptated out, but was highly water
soluble in nature.
Tomorrow I am going to try to run the reaction in acetic acid and ethanol and see what happens. I am worried that I may have to dehydrate the nickel
chloride before use.
[Edited on 2-5-2015 by RareEarth]
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Nicodem
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| Quote: Originally posted by RareEarth | | I think there is a bit of confusion here. There is a difference between chelated complexes of metals and amino acids, and the salts of amino acids
with those metals. What you are talking about is the formation of one of those salts (which involves HCl elimination). In the chelate complexes I am
referring to, HCl is not eliminated. It's a coordination complex that forms. These chelated coordination complexes bind very tightly together so there
isn't much reversible equilibrium taking place. |
You might want to check the charges in your equation. The ways for having no HCl elimination is for hexacoordinate complexes like
H2[Ni(Gly)2Cl2] to form (or H[Ni(Gly)Cl2(H2O)2], or the tetracoordinated
H[Ni(Gly)Cl2]. This is unlikely for some reasons, among it being the acidity decomposing the chelating ligand (as already mentioned).
Another reason is that the two chloride ligands would likely be substituted with the aqua ligands forming the likely more stable
[Ni(Gly)2(H2O)2] species. Furthermore, the prevalent complex species in solution depends also on the stoichiometry.
If you use a 1 : 3 ratio, then H[Ni(Gly)3 might prevail.
But in any case, you cannot talk about a preparation of a complex, unless you crystallize it. In solution you can only have equilibriums of ligand
exchanging species. For this reason, it appears unusual to me that you are trying to invent a new synthesis for a specific complex, rather than first
repeat a procedure from literature. Only then should you attempt new syntheses. If nothing else, because you would better have some reference material
to verify the product identity.
And obviously, most if not all transition metal salts of alpha-amino acids are chelates. It is highly unlikely they could crystallize without forming
the coordination bond with both, the oxygen and the nitrogen of the alpha-aminocarboxylate anion.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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Boffis
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Try using the sodium salt of glycine ie dissolve the glycine in a molar amount of NaOH, or try using nickel acetate instead of the chloride.
By the way NiCO3 is vivid yellow Green and called Gaspeite. Hellyerite is a hellish unstable mineral that dehydrate both in air and even in slightly
warm solution.
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AvBaeyer
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Nicodem is precisely correct. You need to be well aware of the total charge balance of what you are proposing. If you have both glycine and chloride
combined with the nickel, you will have some sort of "nickelate" species (do the charge balance keeping mind the charges on glycine as a function of
pH).
Here is a report for the preparation of the nickel-alanine complex:
N i C l 2 ⋅ 6 H 2 O (20 mL, 0.1 M), KOH (20 mL, 1.0 M), and L-alanine (20 mL, 0.2 M) were mixed. The mixture was made basic with pH = 8
and turned from green to pale blue. The flask solution was left at room temperature. After standing for two weeks, pale-blue tablet-shaped crystals
were obtained, removed, and dried under vacuum. The isolated crystals were subjected to X-ray studies.
This came from: http://www.hindawi.com/journals/ijic/2009/168416/
Perhaps this should work with glycine also.
The following may also be informative: www.scirp.org/journal/PaperDownload.aspx?paperID=21812
There appears to be a fair amount of Ni-Glycine information available for the looking.
AvB
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DraconicAcid
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Collect. Czech. Chem. Commun. 1988, 53, 563-570
doi:10.1135/cccc19880563
Gives a synthesis for k[Ni(gly)3]*3H2O, according to http://cccc.uochb.cas.cz/53/3/0563/
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