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Author: Subject: Removing Mg 2+ from HCl
Hydronium
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[*] posted on 14-12-2005 at 17:43
Removing Mg 2+ from HCl


Hey guys I am new here, but I will get straight to the point. My AP chemistry teacher has created a little mess and asked us students if we can help him solve it. Apparently he has been dissolving magnesium metal into a solution of hydrochloric acid for demonstrations for the normal chemistry class. But now he wants to remove the Mg 2+ and regain his hydrochloric acid. The information he gave was: It used to be a 6M HCl (but it must be lower now since he was been dissolving magnesium metal which liberates hydrogen gas) but is now unknown and there is Mg 2+ in the solution (the volume seemed to be about 2-3 liters). So what do you think we can do?

So far other classmates are looking into precipitating the Mg 2+ out, yet I feel this is a useless endeavor for most of the compounds for magnesium are soluble and the ones that aren’t are formed by neutralizing he solution (i.e. Mg(OH)2 which would neutralize the HCl). Normal electrolysis would create chlorine gas :(. Using phosphates and sulfates would create phosphoric acid and sulfuric acid which is not what he wants. The closest I have gotten in my mind is using electrolysis using a salt bridge (such as MgCl2). So…what do you guys think?
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[*] posted on 14-12-2005 at 17:58


You can distill the acid and the salt will be left behind.

Or

Maybe ou can remove it by using a centrifuge, which would make all of the salt settle out. then u can filter it.




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[*] posted on 14-12-2005 at 18:04


Ok, but unfortunately my school doesn’t own any distilling equipment (also isn't distilling an acid dangerous?). And how would a centrifuge cause the Mg 2+ come out of solution?
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[*] posted on 14-12-2005 at 18:10


It causes the MgCl2 (salt) which is heavier than the acid to settle out.



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Darkblade48
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[*] posted on 14-12-2005 at 18:17


Why not just buy new HCl? It's a relatively cheap and inexpensive mineral acid
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[*] posted on 14-12-2005 at 18:23


Well, he wanted to see if there was a way. I don't know perhaps he wanted to see how good we are.
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[*] posted on 14-12-2005 at 20:59


Add Mg or MgO or MgCO3 to it until it stops reacting, then dehydrate. You'll get the acid, in the form of its salt :P

Can't do any better without distillation equipment I think.

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[*] posted on 14-12-2005 at 23:56


1. ion exchange:

2R-SO3H + Mg+2 ---> (RSO3)2Mg + 2H+

the column can be then regenerated with conc. acid.

2. ion chromatography

3. the most simple way-buy new acid!
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Marvin
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[*] posted on 15-12-2005 at 05:22


Neutralise with, say magnesium hydroxide. Evaporate to dryness. Add concentrated sulphuric acid, and lead the gas produced into water.

Distillation would get back unused acid, but I assume he wants it all back.

This is not difficult to do with school equipment.
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[*] posted on 15-12-2005 at 07:59


Quote:
Originally posted by Marvin
This is not difficult to do with school equipment.

Just a horrible economical waste, since you're using a more expensive acid to regenerate a cheaper one ;)
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[*] posted on 15-12-2005 at 10:10


Actually, Tim's idea isn't such a bad one. Neutralize with magnesium sulfate or carbonate and evaporate to dryness. This will give you a hydrate of magnesium chloride. Heat at a high heat and lead the HCl and water vapor that comes off into water. The decomposition goes something like:

MgCl<sub>2</sub>.xH<sub>2</sub>O –heat--> MgOCl<sub>2</sub> + 3HCl + yH<sub>2</sub>O

[Edited on 15-12-2005 by neutrino]
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[*] posted on 17-12-2005 at 11:08


"Neutralize with magnesium sulfate "
Pardon?

I thought that most school labs had the equipment for demonstrating the distillation of alcohol from wine. You could use the same equipment (with a thermometer rated for up to about 110C rather than the 100C you need for alcohol/water)
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[*] posted on 17-12-2005 at 19:25


Sorry, I meany hydroxide.
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[*] posted on 22-12-2005 at 13:14


Quote:
Originally posted by neutrino
Actually, Tim's idea isn't such a bad one. Neutralize with magnesium sulfate or carbonate and evaporate to dryness. This will give you a hydrate of magnesium chloride. Heat at a high heat and lead the HCl and water vapor that comes off into water. The decomposition goes something like:

MgCl<sub>2</sub>.xH<sub>2</sub>O –heat--> MgOCl<sub>2</sub> + 3HCl + yH<sub>2</sub>O

[Edited on 15-12-2005 by neutrino]

I don't see the stoichiometry of your reaction. Besides that, the oxidation state of the magnesium seems not OK at the right hand side of the arrow, it would be +4 :cool:.
You mean that the salt hydrolyses and looses HCl, itself becoming a basic chloride? Are you sure that happens for MgCl2.6H2O?




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neutrino
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[*] posted on 22-12-2005 at 14:53


Sorry, I meant Mg<sub>2</sub>OCl<sub>2</sub>

I know from personal experience that heating hydrated magnesium chloride will produce HCl vapor. I don't know how far to completion this reaction goes, though.

I refer to an old post made by vulture on the subject:

Quote:
Hydrated MgCl2 decomposes to magnesiumoxychloride and HCl upon heating:

2MgCl2 + H2O ---> 2HCl + Mg2OCl2
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[*] posted on 22-12-2005 at 19:48


I also believe decomposing the hexahydrate would be the right ticket. But I would not bother leading the fumes into water because water will be mixed with the HCl vapor. Maybe a small amount of water in the recieving flask would absorb the excess HCL vapor. But leading the extremely soluble gasses into water would pose too much of a suckback problem.

Afterwards you would titrate to find what concentration you have and take necessary steps to get 6 M HCl. If the distillate is too weak a small amount of 12 M HCl added surely would be easier than boiling to concentration.




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