Strontium - 8-7-2010 at 09:30
Hi,
This week I have been looking at Cysteine complexes of various transition metals and I was wondering if anybody else had investigated these? I
couldn't find much on Google about this (other than in pay per view articles).
I was looking to see if the thiol group would act as a ligand, but there is also the possibility of N, O or even chelation from this molecule:
HOOC-CH(NH2)-CH2-SH
The reaction that I couldn't fathom at all was with Fe 3+
I added 50mg of Ammonium Iron (III) Sulfate to 1 ml of distilled water and dissolved (Usual brown partial hydrolysis occured). I then added 1ml of 5%
L-Cysteine solution. At first a deep blue colouration of the liquid followed by the evolution of a colourless gas and within 1 second the liquid had
completely cleared again. Then the liquid turned hazy and after 10 minutes a snow white precipitate sat beneath the colourless liquid.
I'm guessing that the Iron III was probably reduced to Iron II but wondered what the blue colour was? Could it be a mixed valency multi-centred iron
intermediate (Similar to Prussian Blue), or a transitory blue complex of Iron III that is destroyed by redox? The final preciptate was white, and I
didn't detect any sulfurous smell.
Out of interest my observations of the other metals was:
VO 2+: No Visible reaction.
Cr 3+: Dark purple colour (Initially pale blue-violet).
Mn 2+: No Visible Raction.
Fe 2+: No Visible reaction.
Co 2+: Very strong orange-brown inky colour, especially after heating.
Ni 2+: No Visible reaction.
Cu 2+: Grey-blue precipitate, became pale yellow on heating.
Ag+: Heavy yellow crust
Au 3+: White haze.
Regards,
Strontium.
woelen - 8-7-2010 at 12:06
This sounds interesting what you are doing. I also am interested in metal complexes, but I've never experimented with L-cysteine (I do not have that
compound). If I look at the structure then I think that it coordinates at the S-atom. Iron(III) gives remarkable complexes with many
sulphur-containing compounds in which there is a terminal S-atom or an S-atom with an H atom attached (not with S=O bonded compounds, such as
sulfate).
What you describe in your experiment sounds very much like the experiment I've done with thiosulfate. When thiosulfate is added to a solution with
iron(III) then you get a deep purple complex which fairly quickly fades (takes a few minutes at most) and a colorless liquid remains. Internal redox
inside the complex is responsible for the fading and so I expect that the same thing happened in your experiment.
Could you post some pictures of youir results?
Strontium - 9-7-2010 at 02:43
Hi Woelen,
I tried replicating your thiosulfate experiment and got the purple complex that fades after a few minutes so it may be a similar effect. When I get
some time I'll try varying the concentrations of my solutions and see if I can make my blue complex last long enough to photograph.
Love the photography on your website by the way, a work of art.
Ozone - 9-7-2010 at 06:47
Interesting. I am always interested in Fe complexes and redox cycling (especially with phenolics).
Note that ferric ammonium sulfate is the dodecahydrate at 482.25 g/mol. I am assuming 12H2O to give a mole fraction of 55.845/482.25 = 0.116 of
Fe(III). 50mg = 5.79 mg Fe3+/1000 = 0.0058 g/55.845 g/mol = 0.000104 moles. Assuming % w/w, 1 mL of 5% = 0.05g L-CYS/121.16 g/mol = 0.000413 moles.
0.000413/0.000104 = 1:4 stoichiometry (or simple excess of L-CYS).
A guess:
The Fe(III) formed a transient complex wherein it was reduced to Fe(II) and the cysteine R-SH was oxidized, probably to yield the disulfide-dimer
R-S-S-R, cystine, which precipitated.
The outgassing is interesting, perhaps some concerted decarboxylation (e.g. -CO2)? This was at rt? If that is the case, the product would be
+H3N-(CH2)2-S-S-(CH2)2-NH3+ (the solution should be acidic, so I figured that the amino groups would be quats).
I'll duplicate this today if I have time.
Cheers,
O3
Strontium - 9-7-2010 at 07:49
Thanks Ozone,
I tried with more dilute solutions and this time no gas was evolved. The blue colouration also lasted longer but (unsurprisingly) was less intense.
Perhaps the first time the iron solution had been standing for a few minutes and was more acidic?