[blueberry58]

I `m no chemist but need to produce in low quantity this product.

I had problems writing the reaction, so I am going to write the reactants

Magnesium chloride (aq) more Sodium thiosulfate gives Magnesium thiosulfate and 2 mols of sodium chloride.

I need to perform this to produce some Magnesium thiosulfate.  How find the temperature, time and salt out temperature? and other important considerations.

Where find I information on this kind of reactions?

Thank you

[Arkcon]

I don't think you can make it that way, without excessive purification steps.  Maybe you're better off just purchasing it from a chemical company, if you only need small amounts.

[blueberry58]

This reaction is taken the Encyclopedia de alkaline earth from Richard C Ropp-

The author states this reaction is used and need a salt out temperature from 15 °C or 9,4°C to crystallize the compound.

MY problem is to know the temperature at the begining and the time needed for the reaction.

Thank you

 Edit: Email address removed per forum rules
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[blueberry58]

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?

[Arkcon]

blueberry58:, I hope you don't mind me combining your two posts.  Given the topics are similar enough, its worthwhile to have them and their answers all together.

So, your plan is to mix two soluble salts, and get two soluble products.  You have an ancient reference, that says by cooling, you'll get one product to precipitate out.  I warned you, this wasn't going to be easy.

The precip is too tenuous, so filtering won't work well.  And you probably don't have a centrifuge.  At least not yet.  But maybe, if you keep it cold long enough, it will settle, and you can decant the solution.  Hey, I never said it would be particularly productive.

[Borek]

a tenuous white precipitate appears but no crystallization progresses because the driving force for this reaction is the removal the precipitate from the solution.

That would be the first time I hear that presence of precipitate stops precipitation. Doesn't sound correct to me. Solid precipitates till the concentration in the liquid phase is that defined by the solubility product. I have a feeling someone tried to use LeChatelier's principle not understanding why it doesn't apply in this case (or rather: how to apply it correctly in this case).

[blueberry58]

You are right Borek. I realized not is needed to remove physically the precipitate because its ions are out of the solution.

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)

[blueberry58]

Looking prepare some MgS2O3 I find this other way:

Na2SO3(aq)+MgCl2*6H2O(aq)=MgSO3()+2NaCL(aq)
8MgSO3+S8=8MgS2O3

I could prepare MgSO3 easily.

Now
MgSO3 + S = MgS2O3

It required several hours to boiling temperature and constant agitation. I wonder if any catalyst or change in pH can make it happen faster.

[Archy12345]

If you're still looking for a way to help increase the amount of precipitate that forms, use an excess of MgCl₂ by increasing it's concentration. Common ion effect.

[blueberry58]

Thank you Archy12345.

With a little effort I could prepare magnesium thiosulfate. But I have some doubts because I got a very hygroscopic salt. The test with potassium iodide exhibits a clear aspect, silver nitrate produces yellow opacity and then black. Both indicate that it is a thiosulfate. By mixing with hydrochloric acid diluted a haze white-yellow appears and odor of sulfur dioxide. Also indicates that thiosulphate is not magnesium sulfate.

Has anyone had experience with this salt?

[blueberry58]

The reaction MgSO3 + S = MgS2O3

it is inefficient because their reactants are scarcely soluble in water.
Altough some researcher magnesium sulphate improves the solubility of MgSO3  http://pubs.acs.org/doi/abs/10.1021/je00013a045
but perhaps may be difficult to cristalize the final product.

I thought to change the solvent adding  DMSO 20% ?   
Is there any way to increase the solubility of reactants to raise the yield ?

Thank you.

[blueberry58]

Regarding the preparation of MgS2O3 I have found three feasible reactions

MgCl2*6H2O + Na2S2O3*5H2O = MgS2O3*6H2O + 2NaCl+ 6H2O
MgSO4*7H2O + Na2S2O3*5H2O = MgS2O3*6H2O + Na2SO4*5H2O + H2O
MgSO3*6H2O + Na2S2O3*5H2O = MgS2O3*6H2O + Na2SO3*5H2O

MgO and MgCO3 were avoided because their low solubility

Now, my problem is to separate both salts in each process. Supposedly MgS2O3 precipitates at -9.5 C (salt out temp) but in several trials I couldnt get this.

Commercially and for many years Na2S2O3 is prepared using

Na2SO3+S=Na2S2O3.

I get to prepare MgS2O3  in the same way

MgSO3+S=MgS2O3

But reaction last many hours and need active stirring because solubility of S is poor or none.

What salt would be preferable? or there is a method to precipitate some of these salts by changing the pH or creating another insoluble compound?

Thank you