Junk_Enginerd
Hazard to Others
Posts: 251
Registered: 26-5-2019
Location: Sweden
Member Is Offline
|
|
Thermal decomposition of MgCl2 creates HCl clouds?
I recently found a superbly cheap source of bulk magnesium chloride, and got myself some. Partly as a general purpose magnesium source, and partly
because I'm experimenting with making refractory materials for my furnaces, and there are a few magnesium compounds that seem very fit for the job.
One such compound that I want to explore is magnesium oxide.
I can't find a cheap source of bulk MgO, but I figured there should be a route from MgCl2 to MgO. Indeed it would seem that in air MgCl2 happily
decomposes somewhere around 400-800 °C to form MgO. Great! I couldn't find info on what happens with the chlorine ion, but I thought chlorine gas
would be a reasonable guess.
So I sat myself outside, threw maybe 200 grams of MgCl2 in a stainless steel pot and put it on a burner. It melted and gave off steam as expected. As
temps started rising I sniffed it and started detecting a familiar chlorine smell. Though I didn't see any green hue... I waited and smelled it again.
No this smelled more like HCl... After a short while it was giving off giant clouds. I honestly cannot comprehend how so little material (200 grams)
can create so much mist. Sure the weather was cool and humid, and just on the verge of fog, but this little pot covered a field in mist, something
like 400x200 meters.
I decided to try and figure out if this really was HCl as the smell indicated, and ran my hand through the mist just to see if it would condensate.
Yup. Stings in cuts and tasted like HCl too, albeit somewhat dilute.
I kept it on the burner for something like 30 minutes and it spewed out clouds of HCl mist the entire time and didn't seem like it was going to stop
any time soon. It seemed weirdly endless.
My chemistry knowledge is basic at best, and I don't understand how I'm getting HCl here. Don't get me wrong, it doesn't really matter for what I
tried to do, and if I condense this vapor I have an insanely cheap HCl source which is nice, but it bothers me that I don't understand it. How is this
happening?
|
|
|
DavidJR
National Hazard
Posts: 908
Registered: 1-1-2018
Location: Scotland
Member Is Offline
Mood: Tired
|
|
You have a hydrate of magnesium chloride, not the anhydrous salt. Therefore, there is water molecules that can react as well.
MgCl2 + H2O -> MgO + 2HCl
Many transition metal chlorides have hydrates that decompose like this on heating. So you can't make the anhydrous salts just by heating the hydrates
- unfortunately.
|
|
|
RogueRose
International Hazard
Posts: 1596
Registered: 16-6-2014
Member Is Offline
|
|
I'd check a local farming supply store/feed store and ask for "MagOx" and I'll bet they have it. It's a common feed supplement and it's "cheap"
(compared to tech grade by the lb online). It will probably need to be ground into a powder but that's easier than what you are doing.
Another option is mix it with something like ammonia, or maybe even calcium hydroxide and you'll end up with magnesium hydroxide and either ammonia
chloride or calcium chloride - all readily water soluble while the magnesium hydroxide isn't.
Once you have the mag hydroxide you should filter water through it, or just mix a lot and allow a long time to settle (like 1-4 weeks) - decant and
vacuum out the rest of the water/chloride salt.
Then you can use the hydroxide as is in your refractory as it will just loose it's water in curing/heating. I suspect it would even work better as
the hydroxide than the oxide as it should mix better with the other compounds. If you REALLY want MgO then just heat the hydroxide past it's decomp
point ~665F)
[Edited on 6-26-2019 by RogueRose]
|
|
|
happyfooddance
National Hazard
Posts: 530
Registered: 9-11-2017
Location: Los Angeles, Ca.
Member Is Offline
Mood: No Mood
|
|
As has been said, chloride salts (especially divalent ones) can and do decompose into oxides, HCl, and water, so you haven't discovered anything new
(yet).
The funny thing is that you mention cheap HCl source, and HCl + H2O can be made from salt water (which is ever present on Earth), for cheaper than
pure water, in my neck of the woods at least.
|
|
|
Junk_Enginerd
Hazard to Others
Posts: 251
Registered: 26-5-2019
Location: Sweden
Member Is Offline
|
|
Quote: Originally posted by DavidJR  | You have a hydrate of magnesium chloride, not the anhydrous salt. Therefore, there is water molecules that can react as well.
MgCl2 + H2O -> MgO + 2HCl
Many transition metal chlorides have hydrates that decompose like this on heating. So you can't make the anhydrous salts just by heating the hydrates
- unfortunately. |
Makes sense. Reminds me of trying to dry ammonium nitrate with heat... Well in my case it doesn't matter much, anhydrous salt isn't the goal. Thanks.
| Quote: Originally posted by RogueRose | I'd check a local farming supply store/feed store and ask for "MagOx" and I'll bet they have it. It's a common feed supplement and it's "cheap"
(compared to tech grade by the lb online). It will probably need to be ground into a powder but that's easier than what you are doing.
Another option is mix it with something like ammonia, or maybe even calcium hydroxide and you'll end up with magnesium hydroxide and either ammonia
chloride or calcium chloride - all readily water soluble while the magnesium hydroxide isn't.
Once you have the mag hydroxide you should filter water through it, or just mix a lot and allow a long time to settle (like 1-4 weeks) - decant and
vacuum out the rest of the water/chloride salt.
Then you can use the hydroxide as is in your refractory as it will just loose it's water in curing/heating. I suspect it would even work better as
the hydroxide than the oxide as it should mix better with the other compounds. If you REALLY want MgO then just heat the hydroxide past it's decomp
point ~665F)
[Edited on 6-26-2019 by RogueRose] |
That's a problem with american forums. The product flora is rarely translatable to the swedish one, especially not brand names/product names...
Regardless if it's a feed supplement some equivalent should be available. Thanks for the tip. Unfortunately googling "Magnesium oxide
buy" is useless since magnesium oxide is a common source of magnesium for people too, so the results are drowned in pills, and 100 g for $10 isn't the
best deal for a furnace liner lol.
I did note that magnesium oxide is soluble in ammonia, and added some to my experiment. This only momentarily stopped the acid fumes though. Maybe
ammonia's low boiling point is a problem here?
Calcium hydroxide is interesting. That would be hydraulic mortar right? As in cement... This seems promising for my purpose.
I mean what I would ideally want to create is a refractory cement, so mortar seems like a natural component. Plus I was intending to use sodium
silicate as a binder for the magnesium oxide... Which has a great reaction with calcium chloride...
So:
MgCl2 + Ca(OH)2 -> Mg(OH)2 + CaCl
CaCl2 + Na2SiO3 -> CaSiO3 + NaCl
Best case I'll have a refractory material consisting of MgO(Melting point 2850°C) and CaSiO3(Melting point 2130°C) with some junk NaCl in there that
may or may not simply melt out or allow itself to be washed out.
Interesting. Thanks!
| Quote: Originally posted by happyfooddance |
The funny thing is that you mention cheap HCl source, and HCl + H2O can be made from salt water (which is ever present on Earth), for cheaper than
pure water, in my neck of the woods at least.
|
Really? How? Only thing that comes to mind is electrodialysis.
|
|
|
Herr Haber
International Hazard
Posts: 1236
Registered: 29-1-2016
Member Is Offline
Mood: No Mood
|
|
You can easily get a mix of MgO and MgOH.
It might not be as cheap as you want but it'll be purer and much finer than the other sources discussed here.
Go to your local sports store or gym, get some of that white powder gymnasts or rock climber use and you're done 
|
|
|
walruslover69
Hazard to Others
Posts: 235
Registered: 21-12-2017
Member Is Offline
Mood: No Mood
|
|
I support Herr Haber and Rogue Rose. You can percipitate out MgCO3 by adding washing soda to the solution, or just purchasing Climbing/gymnastics
chalk. The carbonate decomposes at 350C according to wikipedia.
|
|
|
RogueRose
International Hazard
Posts: 1596
Registered: 16-6-2014
Member Is Offline
|
|
The carbonate will settle much faster but it is SUPER fluffy, so mixing or sodium carbonate/bicarb will give you the carbonate and you can wash out
the NaCl easily with water. Filtering will take a little while but it's doeable with a vacuum pump.
I found Mg(OH)2 much more difficult to settle out of a solution as it seems to make a paste where the water just seems to bind.
If wiki is correct it says the hydroxide and carbonate decompose at the same temp which seems questionable. IDK which would be better to work with
for refractory but I suspect you might be able to use a slurry of hydroxide as the liquid and not worry about dehydrating it completely before using
it. Just add water to dilute any salt if you worry about that, could do it a few times and allow it to settle, decant, repeat. If you had a
centrifuge that would work best to seperate the hydroxide and water.
|
|
|
RogueRose
International Hazard
Posts: 1596
Registered: 16-6-2014
Member Is Offline
|
|
Magnesium oxide is produced by the calcination of magnesium carbonate or magnesium hydroxide. The latter is obtained by the treatment of magnesium
chloride solutions, typically seawater, with lime.[10]
Mg2+ + Ca(OH)2 → Mg(OH)2 + Ca2+
Calcining at different temperatures produces magnesium oxide of different reactivity. High temperatures 1500 – 2000 °C diminish the
available surface area and produces dead-burned (often called dead burnt) magnesia, an unreactive form used as a refractory. Calcining
temperatures 1000 – 1500 °C produce hard-burned magnesia, which has limited reactivity and calcining at lower temperature, (700–1000 °C)
produces light-burned magnesia, a reactive form, also known as caustic calcined magnesia. Although some decomposition of the carbonate to oxide occurs
at temperatures below 700 °C, the resulting materials appears to reabsorb carbon dioxide from the air.[11]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by DavidJR | You have a hydrate of magnesium chloride, not the anhydrous salt. Therefore, there is water molecules that can react as well.
MgCl2 + H2O -> MgO + 2HCl
Many transition metal chlorides have hydrates that decompose like this on heating. So you can't make the anhydrous salts just by heating the hydrates
- unfortunately. |
A 2011 source paper on "Thermal decomposition mechanisms of MgCl2·6H2O and MgCl2·H2O" at https://www.sciencedirect.com/science/article/pii/S016523701... noting a complex process, to quote:
"there were six steps in the thermal decomposition of MgCl2·6H2O: producing MgCl2·4H2O at 69 °C, MgCl2·2H2O at 129 °C, MgCl2·nH2O (1 ≤ n ≤
2) and MgOHCl at 167 °C, the conversion of MgCl2·nH2O (1 ≤ n ≤ 2) to Mg(OH)Cl·0.3H2O by simultaneous dehydration and hydrolysis at 203 °C, the
dehydration of Mg(OH)Cl·0.3H2O to MgOHCl at 235 °C, and finally the direct conversion of MgOHCl to the cylindrical particles of MgO at 415 °C. "
[Edited on 29-6-2019 by AJKOER]
|
|
|
happyfooddance
National Hazard
Posts: 530
Registered: 9-11-2017
Location: Los Angeles, Ca.
Member Is Offline
Mood: No Mood
|
|
Yes, exactly. It costs less money to make a given value of HCl than an equal value of pure water, starting from salt water (which is the natural
resource of choice for producing chlorine). Some ops burn it with the hydrogen from the same (or different) processes for energy and use the HCl
produced, others do different things, but I can't think of a single chemical producer or industry that doesn't produce enough HCl from its' waste
processes that they need to buy HCl. I think dealing with the excess is a more common problem. I am sure some places buy HCl, but I can't think of one
off-hand.
I also can't think of many industrial processes that produce pure water as a cheap (or costless) byproduct. The only ones I can think of have H20
coming off as a top product, which is pretty demanding as far as energy requirements go
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
