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DalisAndy
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[*] posted on 27-10-2015 at 18:28
Sodium sulfate hydration


Can Sodium Sulfate have other hydrate states other than 10?



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Metacelsus
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[*] posted on 28-10-2015 at 05:50


The heptahydrate and the anhydrous salt are the other two common ones.



As below, so above.

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annaandherdad
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[*] posted on 28-10-2015 at 06:58


The solubility curve of Na2SO4 has a sharp cusp at the temperature at the transition between two hydrates.



Any other SF Bay chemists?
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Praxichys
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[*] posted on 28-10-2015 at 07:47


Unless specified otherwise, sodium sulfate is usually a mix of air-stable hydrates.

I found a really interesting paper:

"We experimentally show that under real world conditions, both thenardite (Na2SO4) and mirabilite (Na2SO410H2O) precipitate directly from a saturated sodium sulfate solution at room temperature (20°C)."

http://hera.ugr.es/doi/14997496.pdf

Analysis would be the best way to know for sure. Finding water content is easy - just weigh a sample, heat, and then weigh again.




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Upsilon
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[*] posted on 28-10-2015 at 10:10


Quote: Originally posted by Praxichys  
Finding water content is easy - just weigh a sample, heat, and then weigh again.


Thought it appropriate to note that this does not work for all hydrated materials. Chemicals like H2SO4 and P2O5 are almost impossible to dehydrate, and water content would need to be determined by more specialized methods such as titration.
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DistractionGrating
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[*] posted on 28-10-2015 at 14:24


I have a bag of sodium sulfate from this source http://www.dharmatrading.com/chemicals/glaubers-salt.html that claims to be anhydrous, but now I'm dubious. I guess I'll have to do something like creating a solution with a known weight of the sodium sulfate salt, and then adding a slight excess of SrCl2 (I'd use BaCl2, but I don't have any), and then titrate the remaining Sr with EDTA to determine how much Sr precipitated as SrSO4.
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Upsilon
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[*] posted on 28-10-2015 at 15:14


Quote: Originally posted by DistractionGrating  
I have a bag of sodium sulfate from this source http://www.dharmatrading.com/chemicals/glaubers-salt.html that claims to be anhydrous, but now I'm dubious. I guess I'll have to do something like creating a solution with a known weight of the sodium sulfate salt, and then adding a slight excess of SrCl2 (I'd use BaCl2, but I don't have any), and then titrate the remaining Sr with EDTA to determine how much Sr precipitated as SrSO4.


Why go through all of that trouble? As mentioned above sodium sulfate can be largely dehydrated by simple heating.
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DistractionGrating
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[*] posted on 28-10-2015 at 15:58


Quote: Originally posted by Upsilon  

Why go through all of that trouble? As mentioned above sodium sulfate can be largely dehydrated by simple heating.


Because I was paying attention to this:

Quote: Originally posted by Praxichys  
Analysis would be the best way to know for sure.


Plus, I have everything I need to do the analysis I describe readily at hand. That, and the use of the word "largely" make me nervous, especially if I need to have an accurately known concentration of the substance in a stock solution. If nothing else, I could confirm successful conversion of the decahydrate to the anhydrous with this method.

However, I must concede that with a little googling, I found that sodium sulfate is actually used as a primary standard for S, at least by this author: https://books.google.com/books?id=ELQud4ftNX0C&pg=PA446&...
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DalisAndy
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[*] posted on 28-10-2015 at 16:29


So I'm going to assume that H3BO3•Na2SO4 will completely dislocate in water? Making trimethyl borate and need a drying agent and saw that Na2SO4 is one.



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[*] posted on 28-10-2015 at 18:06


You're still wasting your time if you don't have sulfuric acid. It's necessary for a decent yield. In addition, most drying agents are not added to a reflux but rather a water-containing mixture is just dried over it at room temperature.



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DalisAndy
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[*] posted on 28-10-2015 at 18:08


Trimethyl doesn't require sulfuric acid. Or an acid catalyst. I'm just impatient so it should work?



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Amos
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[*] posted on 29-10-2015 at 05:26


Quote: Originally posted by DalisAndy  
Trimethyl doesn't require sulfuric acid. Or an acid catalyst. I'm just impatient so it should work?


Are you trying to make methanol burn green? Or are you trying to produce a significant quantity of trimethyl borate?




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DalisAndy
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[*] posted on 2-11-2015 at 08:26


I'm not trying to make a significant amount, I just wanted a simple experiment that I could test my new lab equipment on



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DistractionGrating
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[*] posted on 6-11-2015 at 14:37


Quote: Originally posted by Upsilon  
Why go through all of that trouble? As mentioned above sodium sulfate can be largely dehydrated by simple heating.


FWIW, I weighed some of my Na2SO4, heated it at 200C for a few hours, cooled, and then weighed again. Both before and after weights were exactly the same.
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