*Warning imprecision present due to a scale unable to measure small enough amounts*
In the lab earlier today I mixed a solution of cobalt nitrate with around 2mL of 3% H2O2 and then with tetrasodium EDTA. It immediately turns a much
darker purple/red color with the release of a gas. The order in which the reagents are added does not appear to matter. Could this be a cobalt peroxo
EDTA complex? I am bad with organics and complex chemistry so I'm unsure of what happened. Either way it is a very cool reaction.DraconicAcid - 17-6-2014 at 15:30
It probably just oxidizes the cobalt(II) EDTA complex to the cobalt(III) complex.
Oh yes I forgot to mention that that shouldn't be because I could not replace the H2O2 with sodium persulfate (although the persulfate does not
interfere).
Can H2O2 oxidize EDTA (I am clueless about organics)? Perhaps cobalt catalyzes the reaction. Again I know nothing of organics.
[Edited on 17-6-2014 by bismuthate]nezza - 18-6-2014 at 11:04
I have made quite a few cobalt complexes. Cobalt II is stable in acid solution but tends to oxidise to cobalt III in alkali. Mixing a cobalt II salt
with EDTA gives no apparent reaction, but there is no precipitate on making the solution alkaline. Adding H2O2 gives a bluish solution of an alkaline
Co(III)EDTA complex. If this solution is acidified you get a deep magenta solution of a cobalt(III)EDTA complex which is stable in acid. I have
attached a picture of some Co(III) complexes.
DraconicAcid - 18-6-2014 at 11:36
Nice pictures. Any idea why the EDTA complex changes colour with pH?blogfast25 - 18-6-2014 at 11:56
Nice pictures. Any idea why the EDTA complex changes colour with pH?
Not sure about colour but buffering to a certain pH is often needed for the complex to form in the first place. This is true at least in part because
the complexation releases H<sup>+</sup>. EDTA titrations are almost always buffered to a prescribed pH, depending on the target cation.DraconicAcid - 18-6-2014 at 12:20
Not sure about colour but buffering to a certain pH is often needed for the complex to form in the first place. This is true at least in part because
the complexation releases H<sup>+</sup>. EDTA titrations are almost always buffered to a prescribed pH, depending on the target cation.
Yeah- I'm just wondering if the acidic version is protonated (as in Co(HEDTA), with the hydrogen attached to one of the acetate oxygens, or
Co(HEDTA)(H2O), with a pentadentate EDTA) and the basic version the simple Co(EDTA) anion, or if the acidic version is the simple anion,
and the basic version a hydroxy complex.
Inorganic Chemistry 1995, 34, 6409 gives the synthesis of the seven-coordinate [Co(HEDTA)(H2O)] complex (but doesn't state its colour), and
references JACS 1959, 81, 549 for the synthesis of NH4[Co(EDTA)]. I don't have access to either full paper.
[Edited on 18-6-2014 by DraconicAcid]bismuthate - 19-6-2014 at 06:37
Well as it turns out the complex was indeed the Co (III) EDTA complex except in a neutral state with a VERY dark color. Perhaps my sodium persulfate
is fake. It should have oxidized the Co(II) like the H2O2. DraconicAcid - 19-6-2014 at 11:02
Well as it turns out the complex was indeed the Co (III) EDTA complex except in a neutral state with a VERY dark color. Perhaps my sodium persulfate
is fake. It should have oxidized the Co(II) like the H2O2.
When I was looking for preps for the Co(EDTA) anion, I found one that suggested that silver ions were required to catalyze the oxidation of Co(II) by
persulphate.unionised - 20-6-2014 at 05:38
Hmmm...I tried the prep for K[Co(EDTA)], but it seems to be *very* water-soluble, and reluctant to precipitate even upon addition of copious amount of
alcohol. I could have sworn I had some tetraethylammonium chloride to try a larger cation, but it turns out that I don't.
