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)
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)
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 ?
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 ?
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.
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.
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.
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.
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?
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?
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.
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.
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.
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.
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).
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).
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)
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)
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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
please read forum rules
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
please read forum rules
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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.
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.
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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
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
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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.
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.
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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?
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?
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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).
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).
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VOTE