teodor
National Hazard
Posts: 876
Registered: 28-6-2019
Location: Heerenveen
Member Is Offline
|
|
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]
|
|
|
Pumukli
National Hazard
Posts: 705
Registered: 2-3-2014
Location: EU
Member Is Offline
Mood: No Mood
|
|
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?
|
|
|
teodor
National Hazard
Posts: 876
Registered: 28-6-2019
Location: Heerenveen
Member Is Offline
|
|
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]
|
|
|
teodor
National Hazard
Posts: 876
Registered: 28-6-2019
Location: Heerenveen
Member Is Offline
|
|
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.
|
|
|
woelen
Super Administrator
Posts: 8012
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
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.
|
|
|
teodor
National Hazard
Posts: 876
Registered: 28-6-2019
Location: Heerenveen
Member Is Offline
|
|
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]
|
|
|
teodor
National Hazard
Posts: 876
Registered: 28-6-2019
Location: Heerenveen
Member Is Offline
|
|
I tried to look into the article I posted here but the link is not working anymore. But eventually I found it, it is attached.
Attachment: in_praise_of_thiosulfate.pdf (3.4MB)
This file has been downloaded 475 times
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
