Sciencemadness Discussion Board

MgCl2 hydrate

fitsc - 15-9-2015 at 14:55

If I combine MgO and HCl will the product MgCl2 always be a hydrate? and will it always be 6H2O. If it's mostly 6H2O I can work with that, or any other consistent hydrate. I read that there's no way to burn off the water without changing the product.

gdflp - 15-9-2015 at 15:11

If by HCl, you mean hydrochloric acid, then yes, the magnesium chloride produced will always be hydrated and, if it was prepared at STP, then it will be the hexahydrate salt. At 117°C, the salt will begin to dehydrate so, if you decide to boil the solution, don't allow it to get hotter than that temperature. Attempting to dehydrate the hexahydrate salt with strong heating will leave you with a semisoluble mass of magnesium oxides, oxychlorides, and chlorides, instead of pure anhydrous magnesium chloride.

fitsc - 15-9-2015 at 15:54

Yes, hydrochloric acid. Thanks, that helps. I want my students to determine roughly how much Mg is in an MgO supplement by experiment. So they can calculate the masses of Mg, Cl and 6H2O of the precipitate using percent mass, assuming all the other stuff gets washed off the little buggars. I'll filter the precipiate and dry it off at a low temp.

AJKOER - 16-9-2015 at 12:10

Actually, I recall an article professing the dehydration of the MgCl2 in a steam of hydrogen chloride can produce the anhydrous salt. [Edit] Found it, see https://www.google.com/url?sa=t&source=web&rct=j&... .

Atomistry.com on MgCl2 ( http://magnesium.atomistry.com/magnesium_chloride.html ) states to quote:

"Anhydrous MgCl2 is white, deliquescent, and soluble in water with great evolution of heat. It is formed by the action of chlorine on the metal or on heated magnesium oxide, or, more easily, on a heated mixture of magnesium oxide and carbon.

It has also been prepared by heating NH4Cl.MgCl2.6H2O, by heating the hydroxide in hydrogen chloride, and by heating the hexa-hydrate in vacuo at 175° C.

It is easily distilled in a current of hydrogen, and the cooled product crystallises into shining laminated crystals that have a density of 2.177 and melt at 708° C.

Between 580° C. and 700° C. the reaction

2MgCl2+O2 ⇔ 2MgO+2Cl2

is endothermic.

Between 350° C. and 505° C. the reaction

MgCl2+H2O ⇔ MgCl (HO)+HCl

is exothermic. The oxychloride decomposes from 505°-510° C., and above the latter temperature the reaction

MgCl2+H2O ⇔ MgO+2HCl

is endothermic. Hydrogen chloride has been prepared by this last reaction. "

My recollection appears to agree with the cited thermal decomposition of NH4Cl.MgCl2.6H2O above to MgCl2, as the heating of the ammonium chloride similarly provides a stream of gases including HCl which contributes to limiting O2 exposure.


[Edited on 16-9-2015 by AJKOER]

MrHomeScientist - 17-9-2015 at 08:00

"A st[r]eam of hydrogen chloride" is very different than hydrochloric acid, though.

DistractionGrating - 17-9-2015 at 11:27

Quote: Originally posted by fitsc  
Yes, hydrochloric acid. Thanks, that helps. I want my students to determine roughly how much Mg is in an MgO supplement by experiment. So they can calculate the masses of Mg, Cl and 6H2O of the precipitate using percent mass, assuming all the other stuff gets washed off the little buggars. I'll filter the precipiate and dry it off at a low temp.


What precipitate?

UC235 - 17-9-2015 at 18:22

Hydrated magnesium chloride does not nicely form crystals except from syrupy saturated solution at low temperatures. It is aggressively hygroscopic and deliquescent and you will not get clean data at all regarding magnesium content.


A much better approach would be to use it to teach titration. Mg+2 (assuming no Ca+2 interference) can be smoothly titrated after dissolution in acid using standardized EDTA and Eriochrome Black T as an indicator.

Praxichys - 18-9-2015 at 04:51

On the note of total Mg analysis: dry the pills and dissolve the sample in HCl, then filter the remaining crap after digestion. Then, add strong NaOH solution until pH 14. The magnesium will precipitate as the very insoluble hydroxide (0.00064 g/100 mL @ 20°C according to Wikipedia)

That could turn into a good lesson about quantitative recovery of a precipitate. Ideally vacuum filter with some quick rinsing with cold water, then dry and weigh. It can be dried at up to 300°C without decomposition so heating can expedite the process if time is short. Keep in mind that calcium also has a very insoluble hydroxide and will skew the test if present, and both can react slowly with atmospheric CO2, so keep it in a small airtight container if it must be stored.

You could also try with sodium carbonate instead of NaOH but keep in mind that magnesium carbonate has a basic salt as well as a bunch of hydrates. It may not work as well.

fitsc - 28-9-2015 at 16:29

Ok I thought the MgCl26H2O would precipitate out into nice little crystals