Sciencemadness Discussion Board

Cobalt Oxalate complexes

Vylletra Heart - 25-6-2017 at 00:56

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]

DraconicAcid - 25-6-2017 at 08:45

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.

mayko - 25-6-2017 at 10:22

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.

Vylletra Heart - 25-6-2017 at 18:59

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! :)

Vylletra Heart - 25-6-2017 at 19:08

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?

DraconicAcid - 25-6-2017 at 19:53

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...

Vylletra Heart - 25-6-2017 at 22:38

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!

nezza - 25-6-2017 at 22:53

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.

Cobalt (III) complexes.jpg - 231kB

Most of them are fairly easy to make from cobalt(II), the complexing reagent and hydrogen peroxide.

DraconicAcid - 26-6-2017 at 00:22

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.

Vylletra Heart - 26-6-2017 at 02:51

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...

AJKOER - 26-6-2017 at 03:29

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]

DraconicAcid - 26-6-2017 at 09:50

https://www.google.ca/url?sa=t&rct=j&q=&esrc=s&a...

This is the paper you want.

Vylletra Heart - 27-6-2017 at 04:23

Quote: Originally posted by DraconicAcid  
https://www.google.ca/url?sa=t&rct=j&q=&esrc=s&a...

This is the paper you want.


Thanks for the help!

nezza - 28-6-2017 at 02:36

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.