blueberry58 - 18-11-2014 at 16:29
For some weeks I had trying to prepare Magnesium thiosulfate (MgS2O3) in small quantites from the following reaction:
MgCl2*6H2O+NaS2O3*5H2O= MgS2O3*6H2O+2NaCl
Separation from NaCl ocurrs by crystallization of the product at low temperature (-9,5 Celsius degrees).
On mixing the reactants at normal temperature nothing happen. By lowering the temperature to - 10 degrees, a tenuous white precipitate appears but no
crystallization progresses because the driving force for this reaction is the removal the precipitate from the solution.
Apart of filtering, exist some method to remove this product?
Thank you.
CHRIS25 - 19-11-2014 at 01:49
At the risk of appearing ignorant, why don't you want to filter? I would simply wash the magnesium several times in the filter paper to flush out any
remaining salt and then allow to dry? The Thiosulphate won't lose water until about 160c, therefore you could heat away a lot of water since you know
how much salt there may well be in there? Just curious. The crystalization to the hexahydrate won't happen in solution until you evaporate, I presume.
Chemosynthesis - 19-11-2014 at 01:52
I agree with asking the above question; other than filtering, you're talking about something like ion exchange chromatography, which is far more
laborious and material intensive.
blueberry58 - 19-11-2014 at 08:05
It`s too tenuous to filter it.
I think this is a difficult issue facing many resources available to the chemical. I wonder if there will be another technology that promotes the
precipitation of a salt. Similar to EM fields of specific frequencies that dissociate water molecules. (Lower the activation energy to break as does
an enzyme)
Chemosynthesis - 19-11-2014 at 09:18
Perhaps if you leave the solution undisturbed in a lab refrigerator overnight, you will find your product has formed a good nucleation site and
crystallized cleanly. Sometimes a seed crystal is preferred. You may also try blowing inert gas over your flask to help speed up evaporation before
leaving it undisturbed, which should assist you. A large part of crystallization is not to introduce additional energy into your solvent via agitation
or heating. It may even be beneficial to add more sodium chloride to remove water from hydrating/solvating your product.
This really isn't a big problem for the chemical industry. There are other techniques such as sonocrystallization, but they do not lower activation
energies as enzymes do, but only adjust the kinetics of a reaction rather than the thermodynamics, intermediates, or transition states. Any technique
such as coulometry to reduce out metal cations (not applicable here due to sodium and the water solvent), ion exchange chromatography, etc. is usually
planned out in advance.
If you find yourself unable to purify a product, it's often because you didn't plan the experiment to completion, and there may be better methods for
getting product. For example, it may be possible to use a technique analogous to that on Wikipedia's sodium thiosulfate entry and just substitute
magnesium sulfite instead. There are patents out there making the sulfite in situ (US2412607 entails the ammonium compound, and US 6921523 B2 appears
somewhat similar for magnesium).
[Edited on 19-11-2014 by Chemosynthesis]
blueberry58 - 19-11-2014 at 09:49
Thank you Chemosynthesis.
Adding some ammonia to the solution, perhaps promote MgS2O3 cristallization. Solubilities of both salts are different to ammonia.
I am looking for a straightforward way to synthezise MgS2O3 in a simple way. These patents look good but I worry about the use of SO2.
This is unortodox medical project to propose to some people and I am working to simplify it and do it less expensive.
www.google.com/patents/US6921523
Perhaps this patent may be implemented in a simple way.
Thank you for your help.