Vylletra Heart
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Cobalt Oxalate complexes
Hello, I am currently trying to synthesize a sodium cobalt oxalate complex. However, I am unable to find any extensive research on such a compound. My
current procedure was this:
--Add black cobalt oxide to excess oxalic acid(interestingly enough, I heard distinct crackling during this, like ice melting rapidly but softer)
--Add sodium bicarbonate to the resultant mixture, but still leaving the acid in excess
--The solution was green at first, like that of ferrioxalate(i'm worried that there may be iron oxide impurities)
--Gradually the solution began to turn to a light pink, like that of the hexaaquo cobalt ion. The pink intensified after heating to a boil.
The issues I'm currently having is that I noticed a lot of websites citing the complex with cobalt in the +3 oxidation state. However, this oxidation
state is only achievable with ammine ligands and the anhydrous fluoride compound. I am not so sure if the oxalate ion is bound to the cobalt ion
strongly enough to permit its oxidation from +2.
Has anyone done any experiments similar to this? I would like to know if the final color of my solution(intense pink) is evident of the [Co(C2O4)2] 2-
ion. Does pH of the solution affect the solubility of this complex? Should I have an excess of oxalic acid or sodium bicarbonate during the
crystallization of the cobalt complex?
Additionally, is the Co(III) related complex possible? If it is possible, is it stable in air? What would be its color?
Any additional information on cobalt oxalate complexes is greatly appreciated! Thanks!
[Edited on 25-6-2017 by Vylletra Heart]
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DraconicAcid
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I don't have any experience with the cobalt(III) oxalate complexes, but the tris(carbonato) complex is an intense dark green, and is perfectly stable.
I beleive that carbonate and oxalate have similar field-splitting characterisitics, so I would not be surprised if the tris(oxalato) complex is also
green.
That being said, the tris(oxalato)ferrate(III) complex is light-sensitive because the iron(III) can oxidize the oxalate. Cobalt(III) is a stronger
oxidizing agent, so I'd expect the complex to be even less stable.
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mayko
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In my experience, transition metal oxalates are pretty insoluble, including cobalt oxalate. Merck lists cobalt (II) oxalate as "almost insoluble in
oxalic acid". So, I'm not sure how well a direct reaction from the oxide will work, although it does sound like you've coaxed some cobalt ions into
solution.
Merck also lists cobalt (II) oxalate as "decayed on heating with aqueous KOH or Na2CO3 solution", if that helps.
What I might try, especially if the cobalt source is of lower purity, is extract and refine the cobalt as a salt like the sulfate, and then
precipitate cobalt oxalate powder using oxalic acid. You can recrystallize this from hot HCl if you want. Then, maybe try reacting it with sodium
oxalate somehow? Solubility might still be a barrier.
Cobal (III) is a pretty high oxidation state that I often hear in the context of things like elemental fluorine, and it appears to oxidize mineral
acids when dissolved in them. I'd be pretty surprised if it appeared spontaneously in such mild aqueous conditions.
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Vylletra Heart
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Quote: Originally posted by DraconicAcid  | I don't have any experience with the cobalt(III) oxalate complexes, but the tris(carbonato) complex is an intense dark green, and is perfectly stable.
I beleive that carbonate and oxalate have similar field-splitting characterisitics, so I would not be surprised if the tris(oxalato) complex is also
green.
That being said, the tris(oxalato)ferrate(III) complex is light-sensitive because the iron(III) can oxidize the oxalate. Cobalt(III) is a stronger
oxidizing agent, so I'd expect the complex to be even less stable. |
Thanks for the answer! May I ask for the formula for this carbonato complex? I would assume you mean [Co(CO3)3]3- where cobalt is +3? Correct me if
I'm wrong.
With that being said, I definitely would expect a cobalt carbonate complex to reach +3 oxidation state due to the carbon atom being in the +4
oxidation state, so there isn't much in the immediate vicinity of the cobalt ion to be reduced by.
Would I be right to assume that the carbonato cobalt(II) complex be the same color as that of the hexaaquo cobalt(II) ion?
Because I would assume that ligands bound to the cobalt ion by oxygen atoms from molecules(H2O) or ions([C2O4]2-) would not cause much color change.
So my theory is that the theoretical color for [Co(H2O)6]3+ (which I understand is definitely not possible), or any Co(III) bound completely by oxygen
atom donors would be green. (For example, the ammine cobalt(III) chloride complex has varying colors due to the different instances when NH3 or Cl-
are bound as a ligand, check Wikipedia for a more elaborate quote on this.)
I have never heard of a carbonato complex with cobalt before though, I've only seen a vague quote on the copper carbonate complex [Cu(CO3)2]2- on
Wikipedia but even then it leaves me with a lot of questions about how to make it. I would love to hear your experimental results on this if you have
any! 
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Vylletra Heart
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| Quote: Originally posted by mayko | In my experience, transition metal oxalates are pretty insoluble, including cobalt oxalate. Merck lists cobalt (II) oxalate as "almost insoluble in
oxalic acid". So, I'm not sure how well a direct reaction from the oxide will work, although it does sound like you've coaxed some cobalt ions into
solution.
Merck also lists cobalt (II) oxalate as "decayed on heating with aqueous KOH or Na2CO3 solution", if that helps.
What I might try, especially if the cobalt source is of lower purity, is extract and refine the cobalt as a salt like the sulfate, and then
precipitate cobalt oxalate powder using oxalic acid. You can recrystallize this from hot HCl if you want. Then, maybe try reacting it with sodium
oxalate somehow? Solubility might still be a barrier.
Cobal (III) is a pretty high oxidation state that I often hear in the context of things like elemental fluorine, and it appears to oxidize mineral
acids when dissolved in them. I'd be pretty surprised if it appeared spontaneously in such mild aqueous conditions. |
Yea I kinda realised my mistake with my initial process and have moved on to your proposed solution. Thanks for that!
You mentioned the oxalate compound being insoluble in oxalic acid. Would I be right in saying that the free, protonated acid is not a
chelating/complexing agent? Does the presence of Na+ affect the solubility of cobalt oxalate regardless of pH?
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DraconicAcid
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| Quote: Originally posted by Vylletra Heart |
Thanks for the answer! May I ask for the formula for this carbonato complex? I would assume you mean [Co(CO3)3]3- where cobalt is +3? Correct me if
I'm wrong. |
That is correct.
| Quote: | Would I be right to assume that the carbonato cobalt(II) complex be the same color as that of the hexaaquo cobalt(II) ion?
Because I would assume that ligands bound to the cobalt ion by oxygen atoms from molecules(H2O) or ions([C2O4]2-) would not cause much color change.
So my theory is that the theoretical color for [Co(H2O)6]3+ (which I understand is definitely not possible), or any Co(III) bound completely by oxygen
atom donors would be green. (For example, the ammine cobalt(III) chloride complex has varying colors due to the different instances when NH3 or Cl-
are bound as a ligand, check Wikipedia for a more elaborate quote on this.)
I have never heard of a carbonato complex with cobalt before though, I've only seen a vague quote on the copper carbonate complex [Cu(CO3)2]2- on
Wikipedia but even then it leaves me with a lot of questions about how to make it. I would love to hear your experimental results on this if you have
any! |
I've made the tris(carbonate)cobaltate(III) complex several times, but I don't think it forms a complex with cobalt(II). Carbonate isn't much of a
ligand.
If you're on ResearchGate, you can read this article: https://www.researchgate.net/publication/231503890_A_general...
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Vylletra Heart
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| Quote: Originally posted by DraconicAcid | | Quote: Originally posted by Vylletra Heart |
Thanks for the answer! May I ask for the formula for this carbonato complex? I would assume you mean [Co(CO3)3]3- where cobalt is +3? Correct me if
I'm wrong. |
That is correct.
| Quote: | Would I be right to assume that the carbonato cobalt(II) complex be the same color as that of the hexaaquo cobalt(II) ion?
Because I would assume that ligands bound to the cobalt ion by oxygen atoms from molecules(H2O) or ions([C2O4]2-) would not cause much color change.
So my theory is that the theoretical color for [Co(H2O)6]3+ (which I understand is definitely not possible), or any Co(III) bound completely by oxygen
atom donors would be green. (For example, the ammine cobalt(III) chloride complex has varying colors due to the different instances when NH3 or Cl-
are bound as a ligand, check Wikipedia for a more elaborate quote on this.)
I have never heard of a carbonato complex with cobalt before though, I've only seen a vague quote on the copper carbonate complex [Cu(CO3)2]2- on
Wikipedia but even then it leaves me with a lot of questions about how to make it. I would love to hear your experimental results on this if you have
any! |
I've made the tris(carbonate)cobaltate(III) complex several times, but I don't think it forms a complex with cobalt(II). Carbonate isn't much of a
ligand.
If you're on ResearchGate, you can read this article: https://www.researchgate.net/publication/231503890_A_general... |
Unfortunately I am not a user of the site. I tried making an account but it seems like you need to be affiliated with an organization to be able to
request for the full text(and I don't believe fictitious inputs would help). Sorry to trouble you but maybe you could just give me the basic steps and
conditions to make the complex? I would really love to see the green for myself!
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nezza
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I have made some Cobalt(III) complexes before including the carbonate which is bright green. I have never tried to make the oxalate, but I'll have a
go at it now.
Most of them are fairly easy to make from cobalt(II), the complexing reagent and hydrogen peroxide.
If you're not part of the solution, you're part of the precipitate.
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DraconicAcid
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| Quote: Originally posted by Vylletra Heart | | Quote: Originally posted by DraconicAcid | | Quote: Originally posted by Vylletra Heart |
Thanks for the answer! May I ask for the formula for this carbonato complex? I would assume you mean [Co(CO3)3]3- where cobalt is +3? Correct me if
I'm wrong. |
That is correct.
| Quote: | Would I be right to assume that the carbonato cobalt(II) complex be the same color as that of the hexaaquo cobalt(II) ion?
Because I would assume that ligands bound to the cobalt ion by oxygen atoms from molecules(H2O) or ions([C2O4]2-) would not cause much color change.
So my theory is that the theoretical color for [Co(H2O)6]3+ (which I understand is definitely not possible), or any Co(III) bound completely by oxygen
atom donors would be green. (For example, the ammine cobalt(III) chloride complex has varying colors due to the different instances when NH3 or Cl-
are bound as a ligand, check Wikipedia for a more elaborate quote on this.)
I have never heard of a carbonato complex with cobalt before though, I've only seen a vague quote on the copper carbonate complex [Cu(CO3)2]2- on
Wikipedia but even then it leaves me with a lot of questions about how to make it. I would love to hear your experimental results on this if you have
any! |
I've made the tris(carbonate)cobaltate(III) complex several times, but I don't think it forms a complex with cobalt(II). Carbonate isn't much of a
ligand.
If you're on ResearchGate, you can read this article: https://www.researchgate.net/publication/231503890_A_general... |
Unfortunately I am not a user of the site. I tried making an account but it seems like you need to be affiliated with an organization to be able to
request for the full text(and I don't believe fictitious inputs would help). Sorry to trouble you but maybe you could just give me the basic steps and
conditions to make the complex? I would really love to see the green for myself! |
I'll post them tomorrow from work.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Vylletra Heart
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| Quote: Originally posted by nezza | I have made some Cobalt(III) complexes before including the carbonate which is bright green. I have never tried to make the oxalate, but I'll have a
go at it now.
Most of them are fairly easy to make from cobalt(II), the complexing reagent and hydrogen peroxide. |
Interesting! From your photo I can tell you've been able to get your hands on some really hardly accessible chemicals. Congratulations on that aspect!
However I would like to know if you have made any complexes that are more achievable in terms of reactants to the average Joe(any ventures into the
chemical complexes of other transition metals would be welcome as well ).
I would especially love to know about your work on the carbonate complex seeing your proof of product! Nevertheless, I look forward to hearing your
results on the oxalate complex as well! I'm too scared to scale up my experiments because I'm still unaware if it works and I really don't want to
waste some expensive cobalt...
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AJKOER
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As I have noted on prior occasions (like http://www.sciencemadness.org/talk/viewthread.php?tid=69355#... ), the combination of a transition metal in a lower valent state, an acid source
and O2 (or H2O2 producing a fenton reaction in the case of iron, or HOCl creating a so called fenton-type reaction, but not with H2C2O4 as the acid
source, as it also forms CO2 and Cl2 vigorously, see my comments at http://www.sciencemadness.org/talk/viewthread.php?tid=63645&... or employ KMnO4, but again not with oxalate also liberating CO2, see https://www.google.com/url?sa=t&source=web&rct=j&... ) undergoes an electrochemical reaction with the metal moving to a higher valent
state (think of the process of iron rusting) consuming oxygen and acid. This may be occurring here also.
It can be accelerated in the presence of a good electrolyte (especially sea salt) and boiling the mix in air.
The comments in parentheses are provided for guidance if experimenting with different paths.
[Edited on 26-6-2017 by AJKOER]
[Edited on 26-6-2017 by AJKOER]
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DraconicAcid
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https://www.google.ca/url?sa=t&rct=j&q=&esrc=s&a...
This is the paper you want.
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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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Vylletra Heart
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Thanks for the help!
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nezza
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The cobalt(III) carbonate complex is easy enough.
Mix cold sodium bicarbonate solution with dilute H2O2.
Add a cobalt(II) solution. There will be a muddy green precipitate and some effervescence.
After a few minutes the green precipitate will go into solution.
The solution will still release oxygen slowly and it is not indefinitely stable.
It is not cobalt(II) carbonate as that is pink or blueish and insoluble in water.
If you're not part of the solution, you're part of the precipitate.
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