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Author: Subject: Reaction dynamics of thiosulfate ion
teodor
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[*] posted on 19-7-2019 at 03:17
Reaction dynamics of thiosulfate ion


It was another accidental observation which called me for studying some new part of chemistry.

I just had impure Na2S2O3 mother liqueur and also impure BaCl2 mother liqueur (the rest I had after separation of crystals). I decided to mix them to get BaS2O3 (mostly insoluble in water but soluble in acids). I expected the same speed of the reaction as with formation of BaSO4 precipitation. And nothing. The big flat crystals have appeared the next day.

Schlessinger in his recipe mentions 35 g BaCl2 per 100 ml of water and mix it with solid Na2SO3 (http://www.prepchem.com/synthesis-of-barium-thiosulfate/). It works perfectly and BaS2O3 forms immediately (as a white powder).

Also, the main discoverer of thiosulfates, John Herschell wrote :

"This salt precipitates copiously when muriate of baryta is poured into a solution, _not very dilute_, of hyposulphite of lime... When the solutions from which it is to be precipitated are mixed in a somewhat dilute state .... some minutes elapse before any precipitation or cloudiness commences".

https://archive.org/details/edinburghphilos05edingoog/page/n...

So, my first thought about it that thiosulfate acts as not ionic compound ... but it should be ionic. But why we have no immediate reaction? Just some thermodynamic rules or there are some properties of the ion, so the reaction BaCl2 + Na2S2O3 = BaS2O3 + 2NaCl is not go in a single step?

And an observation of crystal forming process makes me to think that it is not the case of supersaturation, because under some concentrations crystals start forming immediately, just very slow, like one per minute.


[Edited on 19-7-2019 by teodor]
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[*] posted on 20-7-2019 at 00:36


Maybe the slowest step is the aggregation of BaS2O3 miniature (nanometer range?) crystals to bigger, visible ones. Maybe the reaction is immediate between the ions but there's something that prevents the formation of bigger crystals.
Maybe the ionic strength of the solution, maybe something else.

E.g have you tried to heat up a "slow" mixture to see if it has an accelerating effect on crystallisation?
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teodor
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[*] posted on 22-7-2019 at 12:11


Today I made the experiment you suggested.

I did it this way. I took 3 test tubes pouring in each the same amount of Na2S2O3 solution (I would say the concentration was "diluted", I am able to calculate it). I put 2 of these tubes in just boiled water and kept few minutes there. Then I drop the same amount of 2M BaCl2 solution in each of these tubes. So

1. Cold
2. Hot
3. Hot with agitation

The big flat crystals began to form in the first (Cold!) tube. The two hot tubes shown no crystals.

I waited a bit (few minutes, no crystals in hot tubes) and then I put the hold tube in the same water bath as 2 others were rest.

After some time I checked 3 tubes, which now had the same temperature, so:

1. Cold reaction, heating
2. Hot reaction
3. Hot reaction and agitation at the beginning

The same speed ("snow") of crystals forming and the same type of crystals in each tube.
The first tube ("cold start") has the biggest amount of crystals. The second is "hot with agitation at the beginning". The 2nd tube ("hot reaction") has a _very few_ crystals formed at the moment.

So, I feel myself like a stupid guy who has no ability to interpret the data he got.

It seams that the hot tubes reached the same speed of crystal forming just by standing a bit, but the first, cold tube, had this speed at the beginning.




[Edited on 22-7-2019 by teodor]
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[*] posted on 12-9-2019 at 03:22


Some new experiment with BaS2O3.

1) 1 part of 2M BaCl2 , 2 parts of 1M Na2S2O3. No precipitation even after ~0.5 hours
2) 1 part of 2M BaCl2, 2 parts of 1M Na2S2O3, 2 parts H2O. In less than 1 minute the precipitation started to form and in 0.5 hours it was copious amount.

So, I see unusual backward dependability both on temperature and concentration. Of course until some point because in solid form Na2S2O3 precipitate BaS2O3 immediately.
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[*] posted on 12-9-2019 at 22:45


One explanation of the fact that the hot solution does not form crystals is that the solid is more soluble in hot water. So, the complex may be formed faster in hot water, but if it also is more soluble, then it will not precipitate.

Another effect may be that the complex is more labile at high temperature and that it splits up more easily into it constituent ions at higher temperature. This also may prevent precipitation at higher temperature.




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teodor
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[*] posted on 13-9-2019 at 00:37


I found some interesting article which summarize the thiosulfate ion properties - http://digicoll.library.wisc.edu/cgi-bin/JCE/JCE-idx?type=tu...


And there "The precipitation of BaS2O3 from moderately concentrated solution is a slow process that can he accelerated by rubbing the inner wall of the (glass) container with a glass stirring rod - the rubbing procedure makes a useful demonstration".

Can this acceleration by rubbing suggest which version is more presumable?



[Edited on 13-9-2019 by teodor]
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