Can Sodium Sulfate have other hydrate states other than 10?Metacelsus - 28-10-2015 at 05:50
The heptahydrate and the anhydrous salt are the other two common ones.annaandherdad - 28-10-2015 at 06:58
The solubility curve of Na2SO4 has a sharp cusp at the temperature at the transition between two hydrates. Praxichys - 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)."
Analysis would be the best way to know for sure. Finding water content is easy - just weigh a sample, heat, and then weigh again. Upsilon - 28-10-2015 at 10:10
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.DistractionGrating - 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.Upsilon - 28-10-2015 at 15:14
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.DistractionGrating - 28-10-2015 at 15:58
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&... DalisAndy - 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.Amos - 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.DalisAndy - 28-10-2015 at 18:08
Trimethyl doesn't require sulfuric acid. Or an acid catalyst. I'm just impatient so it should work?Amos - 29-10-2015 at 05:26
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? DalisAndy - 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 onDistractionGrating - 6-11-2015 at 14:37
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.
