Lots of threads seem to mention small bits about qualitative metal ion analysis in a home chemistry lab, but none seem to go into detail about it. I
took out one of my old general chem lab manuals which has a procedure for separating Group I-IV cations. The H2S source seems to be a 1M solution of
thioacetamide, which looks to be rather expensive and hard to come by. This thread mentions substituting thiourea with thioacetamide as a H2S source, has anyone actually tried this with success? It looks plausible, the
only difference between the two compounds is a methyl group being substituted for a primary amine group. I'm guessing, if thiourea is plausible, due
to the lower solubility of thiourea, more has to be used to achieve similar sulfide ion concentration?Boffis - 30-4-2014 at 16:59
@gdflp Thiourea does not react in the same way, it doesn't act as an H2S donor as easily as thioacetamide and so metals like silver don't precipitate
black silver sulphide with thiourea, indeed thiourea dissolve silver sulphide. (hence its use in silver cleaning products like Godhards silver dip
etc). At one time BDH (long lost chemical supplier) sold a hydrogen sulphide substitute for just this type of analysis under the name "Emdite" which I
believe was Diethylamine Diethylthiocarbamate. If this is correct you could make it by dissolving 1 mol equivalent of CS2 in two of diethylamine. The
sodium salt is a well known reagent for copper so probably reacts similarly and it regularly crops up on ebay.
Alternatively you can make thioacetamide by reacting H2S with acetonitrile (see Organic Chemistry by Richter v1 p282) though not stated in Richter the
reaction is catalysed by traces of water. There is another preparation given in Org Syn under "2,5 Dimethylthiazole" Col vol 3 p332 but this uses the
rather inaccessible P2S5.
Also white precipitated zinc sulphide can be used as a sulphide donor, if precipitated into the pores of filter paper it reacts quickly the powder
more slowly.Dan Vizine - 2-5-2014 at 07:55
I always go with what is simplest. You can buy sodium sulfide on eBay from several sources. Just acidify this with sulfuric acid.
Remember that the toxicity exceeds that of HCN.Mildronate - 2-5-2014 at 08:36
Yes we used sodium sulphide in in university. I hated qualitative course, this old way is pain in ass, especially when you analyze destillated H2O and
think that there is ions gdflp - 2-5-2014 at 09:30
Using H2S has limitations for the following reasons. As mentioned before it is extremely toxic, and difficult to lead into solution if it is only a
small amount in a test tube. In addition, it is much more difficult to specifically control the sulfide ion concentration to allow separation of
Group II, Group III and Group IV ions.blogfast25 - 2-5-2014 at 09:57
So [S<sup>2-</sup>] can be manipulated to some extent via pH (buffering). A paper I read years ago showed that the total solubility of
H<sub>2</sub>S is remarkably independent of pH over quite a wide range. That makes selecting the right pH to exceed (or not, depending on
what element you want to precipitate) the solubility product of a selected sulphide much easier.
Separating elements from others by means of pH selective precipitation as sulphides has certainly been used in the past in aqueous analytical
chemistry.
[Edited on 2-5-2014 by blogfast25]Metacelsus - 2-5-2014 at 16:50
Having any appreciable amount of sulfide ions in aqueous solution requires an extremely high pH (the pKa of hydrosulfide is 12.9). In most cases the
dominant species is hydrosulfide.Boffis - 23-5-2014 at 05:53
While rummaging through some old papers yesterday I found this paper about the generation of H2S from acetamide and I thought it might be useful to
some. Of particular interest though is the inadvertent claim that in the presence of ammonia thioacetamide hydrolyses to ammonium sulphide and
acetamidine. The later compound is a difficult to obtain compound used in the synthesis of heterocyclics. The use of hydrazine hydrate may therefore
yield the even more obscure acetamidazine, the starting material for the preparation of some tetrazole compounds.
Attachment: 30_38.pdf (361kB) This file has been downloaded 976 times
Having any appreciable amount of sulfide ions in aqueous solution requires an extremely high pH (the pKa of hydrosulfide is 12.9). In most cases the
dominant species is hydrosulfide.
True as it is it doesn't contribute anything to this discussion.
To exceed the solubility product of for instance CuS (8 x 10<sup>-37</sup> you don't need 'appreciable' amounts of sulphide ions, indeed very small amounts suffice.
For other sulphides the [S<sup>2-</sup>] won't be enough to reach solubility limit.
[S<sup>2-</sup>] can be manipulated to some degree with pH.
you don't need 'appreciable' amounts of sulphide ions, indeed very small amounts suffice.