Sciencemadness Discussion Board

Cobalt Oxalate Precipitation

Vylletra Heart - 7-7-2017 at 14:19

I recently posted a topic about cobalt oxalate complexes on this site and am currently facing new issues that are not exactly related to my previous post.

It is regarding the synthesis of cobalt (II) oxalate, Co(C2O4). According to Wikipedia, the standard procedure is the precipitation of the desired product from mixing CoSO4 and (NH4)2(C2O4).

However, I tried a different procedure: mixing cobalt(II) chloride (made by reacting pottery grade cobalt oxide and double distilled hydrochloric acid toilet cleaner) with oxalic acid (this I got from a chemical supplier, it has the assay written on it so I trust that its pure). The CoCl2 solution is the expected reddish-purple color so I would expect a significant amount of Co2+ ions in solution. However, upon small additions of oxalic acid solution to the CoCl2 batch, no precipitation occurs. May I ask if there is anything wrong with my procedure and what are the specific conditions for cobalt oxalate precipitation? (e.g. type of ions, concentration of ions, pH...) The only probable issue I can think of is that there may have been some volatile chelating agent that distilled together with the HCl (I did make sure to discard everything that came over below 108C in my second distillations though)

The reason I chose not to follow the standard procedure was because I expected that the resultant precipitate would form an ammonium cobalt oxalate soluble complex which I cannot separate from the existing solution.
I'm really irked by my negative results and hope that anyone who has done this before can give me a procedure for a positive result. Thanks!

DraconicAcid - 7-7-2017 at 14:36

The problem is probably that the oxalate is mostly in the form of HC2O4-(aq) instead of C2O4=(aq). You need some base to neutralize the acid.

mayko - 7-7-2017 at 16:14

I have generally added an excess of oxalic acid when preparing transition metal oxalates from soluble salts. In my notes, I precdipitate cobalt oxalate using a ~3x excess of oxalic acid. I also used cobalt nitrate, but I don't think the anion matters too much so long as the cobalt salt is soluble (the chloride might even have the advantage of being driven both by precipitation of the oxalate and evaporation of the HCl formed). I also used heating, which I didn't see you mention.

Vylletra Heart - 7-7-2017 at 17:40

Quote: Originally posted by mayko  
I have generally added an excess of oxalic acid when preparing transition metal oxalates from soluble salts. In my notes, I precdipitate cobalt oxalate using a ~3x excess of oxalic acid. I also used cobalt nitrate, but I don't think the anion matters too much so long as the cobalt salt is soluble (the chloride might even have the advantage of being driven both by precipitation of the oxalate and evaporation of the HCl formed). I also used heating, which I didn't see you mention.


Thanks. It's great to hear that I can continue with using oxalic acid. Why do you use such a huge excess of the acid though?

nezza - 8-7-2017 at 04:02

I have a few observations on Cobalt Oxalate.

Adding stoichiometric amounts of Oxalic acid to Cobalt (II) gives a sand coloured precipitate. Preumably this is Cobalt(II) Oxalate.

The precipitate is almost insoluble in water, but sensitive tests for Co(II) do give a reaction.

It is slightly soluble in Sodium Oxalate solution on heating giving a magenta solution. Presumably a cobalt(II) oxalate complex.

The precipitate is soluble in aqueous solutions of ligands such as ammonia or Ethylene Diamine.

If alkali is added to the solution in sodium oxalate cobalt(II) hydroxide is precipitated. This oxidises on standing.

My conclusion is that any cobalt oxalate complexes are rather weak and unstable and the oxalate can be displaced easily.

I have not been able to oxidise cobalt(II) oxalate to cobalt(III) oxalate complexes as adding H2O2 to any alkaline cobalt oxalate solution/suspension immediately precipitates cobalt(III) hydroxide.

Vylletra Heart - 8-7-2017 at 04:57

Quote: Originally posted by nezza  
I have a few observations on Cobalt Oxalate.

Adding stoichiometric amounts of Oxalic acid to Cobalt (II) gives a sand coloured precipitate. Preumably this is Cobalt(II) Oxalate.

The precipitate is almost insoluble in water, but sensitive tests for Co(II) do give a reaction.

It is slightly soluble in Sodium Oxalate solution on heating giving a magenta solution. Presumably a cobalt(II) oxalate complex.

The precipitate is soluble in aqueous solutions of ligands such as ammonia or Ethylene Diamine.

If alkali is added to the solution in sodium oxalate cobalt(II) hydroxide is precipitated. This oxidises on standing.

My conclusion is that any cobalt oxalate complexes are rather weak and unstable and the oxalate can be displaced easily.

I have not been able to oxidise cobalt(II) oxalate to cobalt(III) oxalate complexes as adding H2O2 to any alkaline cobalt oxalate solution/suspension immediately precipitates cobalt(III) hydroxide.


Yes I have briefly discussed about the oxalate complex in my other recent post: http://www.sciencemadness.org/talk/viewthread.php?tid=75012

I just found it weird that I could not get Co[C2O4] from mixing CoCl2 and oxalic acid, and judging by your sand colored precipitate I would assume you got your reaction to work. I've heard that heat is needed to drive the reaction to completion, and considering that I've only added small amounts of acid to my CoCl2 for fear of wasting it, I think that I may have to throw in the towel and gamble with a larger amount of acid.

nezza - 8-7-2017 at 06:20

The precipitation is slow so from dilute solutions you may just have to wait a while.

mayko - 8-7-2017 at 10:13

Quote: Originally posted by Vylletra Heart  
Why do you use such a huge excess of the acid though?


I'd expect the excess to help the reaction go to completion, since the reaction mixture will still be a strong oxalic solution even when the majority of the cobalt is consumed. I'd rather use a large excess of oxalic acid, which is pretty inexpensive and benign, than an excess of cobalt, which is in comparison more costly and modestly toxic.