MKSStal - 31-7-2018 at 14:25
I will get about 3 kg of sodium thiosulfate. Any ideas for some cool use of it? Maybe some special metal complex synthesis? Something that is not
iodometry, reaction with acid etc...
CouchHatter - 31-7-2018 at 15:43
Check out the ScienceMadness wiki page for sodium thiosulfate.
clearly_not_atara - 31-7-2018 at 17:28
You can make thiols from sodium thiosulfate via the Bunte salts. Chlorination of the salts leads to the formation of sulfonyl halides.
j_sum1 - 31-7-2018 at 18:01
Sodium thiosulfate is something I always have on hand in my lab. I have more of it than bisulfite. I use it whenever I am working with anything
oxidising -- to neutralise Br2, to clean up iodine spills, to scrub Cl2, to deal with Cr(VI), to neutralise spills of permanganate solution. It is
also useful for iodometric titrations. I really wouldn't be without it.
3kg is quite a bit but IMO not an excessive amount to have on hand.
If you want an interesting investigation on reaction kinetics then the reaction between a thiosulfate solution and an acid is a really easy one to do.
It forms a sulfur precipitate that makes the reaction mixture cloudy. You can measure the time for the turbidity to obscure a black mark on a piece
of paper that the beaker is sitting on. Adjust some concentrations and temperatures and away you go. Classic basic reaction kinetics -- you won't
discover anything new here. But this is a good one for lifting theory out of a textbook and actually seeing it.
And, yes, I believe it has some cool properties in its own right, but I confess that I have never really investigate those. If you can get 3kg for
cheap, then go for it.
Sulaiman - 31-7-2018 at 20:29
I give my hands a rinse with thiosulphate solution immediately after working with silver nitrate,
BEFORE exposing my hands to sunlight.
woelen - 1-8-2018 at 01:50
Another interesting one is investigating complexes of thiosulfate. Try with copper(II) ions and also with ferric ions. Allowing solutions to stand and
working with cold and hot solutions makes a difference. The copper and iron complexes are quite interesting.
You can also do interesting crystallization experiments with the pentahydrate. Gently heat solid Na2S2O3.5H2O in a test tube or a miniature (e.g. 25
ml) beaker. It seems to melt at low temperature (somewhere between 50 and 60 C), but it is not really melting, it just is splitting of water of
hydration and dissolving in this free water. Allow some to cool down again and watch the formation of the solid. Nice to make a movie of that. Adding
a single crystal of solid to the liquid is nice as well.
CobaltChloride - 1-8-2018 at 07:57
Being inspired by woelen's post I did those exact experiments.
I didn't get any color change with ferric chloride, but that might just have been due to chloride interfering so I'm going to prepare some ferric
sulfate in order to try the reaction again.
With copper II (from CuSO4*5H2O) in cold solution I got a light green solution and with copper II in hot solution I got a sort of brown deposit on the
walls of the test tube along with some dark brown precipitate in suspension.
After letting both solutions sit for a couple of hours they look very similar, both having some brown precipitate at the bottom and having a greenish
color. The deposit on the walls of the test tube where I used hot solutions looks almost like cold chocolate in color.
At the moment I'm trying to recover the precipitate in order to characterize it.
[Edited on 1-8-2018 by CobaltChloride]
woelen - 2-8-2018 at 00:21
With copper you indeed get a yellow/green solution, which slowly decomposes on standing, producing a dark brown solid.
With iron(III) you definitely should see a change of color. This gives a very intensely colored purple complex with thiosulfate, which quickly fades
again. With iron(II) there will be no visible color change, the liquid will remain (nearly) colorless. Try cold solutions, then the color will linger
for a somewhat longer time so that you can easily observe it.
CobaltChloride - 2-8-2018 at 00:53
I retried making the Fe (III) complex this time using a solution of ferric sulfate made by adding stoichiometric quantities of dilute H2SO4 and H2O2
(in slight excess) to (NH4)2Fe(SO4)2*6H2O. This resulted in a light yellow solution. Upon adding sodium thiosulfate to this, the solution turned a
slightly purplish caramel color. After swirling the contents of the test tube, the color eventually faded. It seems that the chloride was interfering
in my previous experiment by complexing the Fe 3+ ions to make FeCl4 -.