Sciencemadness Discussion Board - failed attempt for magnesium chloride.

failed attempt for magnesium chloride.

gardul - 27-10-2014 at 14:13

As I am trying to have a decent amount of chlorides in my inventory I have become rather annoyed with this one.

What I thought would be an easy thing has turned into an adventure.

this is how i prepared it.

20g of MgSO4
24g NaHCO3
200ml of warm distilled water.

what my thought was on the reaction

MgSO4 + NaHCO3 + H2O

2 NaHCO3 + MgSO4 = Na2SO4 + Mg(HCO3)2 +H2O

Now watching the reaction, I don't think the solution was a sodium sulfate. And honestly I wasn't sure to start if this was even correct. But I did set it on the hot plate and let it boil. the clear liquid became chloudy and the product started be produced.

After the product is dried, I would assume it would be Magnesium Carbonate.

Now I produced 31g of this. It didn't click before hand until now that I had a flawed product. No way should I have had that much. My thinking now is that it was filled with contamination. but i digress and time to move on.

I then places the 31g in 200ml 14% hydrochloric acid and when boiled down and set out to dry only produced 1.3g of what ever this may be.

While I think I failed at doing it this way, I think perhaps magnesium hydroxide may be better suited for this adventure. and I will post results once I have them.

Now if it is possible to get magnesium chloride this way, ether two things happen.
1) its not really an efficient way
2) I screwed up somewhere and its time to go back to the drawing board.

blogfast25 - 27-10-2014 at 14:38

Try it with sodium carbonate, not sodium bicarbonate (heating the latter to about 200 C gives the former anyway).

Your initial product was probably riddled with sodium carbonate/bicarbonate.

Precipitate the Mg carbonate from a fairly dilute solution, e.g. 1 M MgSO4.

Nowhere did you mention filtration. Filter and wash carefully before drying, you should get reasonably close to Theoretical Yield.

Even for such a simple synthesis it's worth looking up a tried and tested procedure.

UnintentionalChaos - 27-10-2014 at 14:41

Adding carbonate or bicarbonate to a magnesium salt solution precipitates magnesium carbonate and/or basic carbonate. Magnesium bicarbonate itself is moderately soluble and it will offgas CO₂ and precipitate carbonate.

Magnesium chloride crystallizes as the hexahydrate and is extremely hygroscopic. Unless you're living in a desert, it's unlikely that setting a dish of the solution out would ever dry out enough to form crystals.

gardul - 27-10-2014 at 15:08

Actually I do live in a desert. and yes I am fully aware that it is hygroscopic. This is actually one of the reasons why I am temping it. I may be "newish" to chemistry but I'm not a complete fool.

I guess I should have mention the filtration. and I do apologize for not adding that.

I am not really looking for answers here. I was merely stating hey, this doesn't work. I am sure some one new to chemistry may also have interest in chlorides at first, and may stumble a on this issue as well. I will figure this issue out. It is part of the fun. Since I am rather comfortable with my self, I figured I would share a failed attempt in something. We all fail at times. it's just a matter of actually admitting to it or not.

EDIT: the dried substance is actually a bright yellow now. I am very curious in what this may be as well.

[Edited on 27-10-2014 by gardul]

gdflp - 27-10-2014 at 15:16

How strong is the yellow color and how pure is your HCl? I'm guessing the yellow color is some iron contamination from the acid.

gardul - 27-10-2014 at 15:23

Quote: Originally posted by gdflp

How strong is the yellow color and how pure is your HCl? I'm guessing the yellow color is some iron contamination from the acid.

You would be correct on that. it is old 14% hydrochloric acid that has iron impurities. It's actually rather bright Yellow. I am sure it isn't anything fun or interesting. But Il going attempt to put some in a test tube and see if it reacts with anything.

Texium - 27-10-2014 at 15:38

Quote: Originally posted by gardul

Quote: Originally posted by gdflp

How strong is the yellow color and how pure is your HCl? I'm guessing the yellow color is some iron contamination from the acid.

More likely than not it is iron(III) chloride. It can definitely look bright yellow, particularly if it has a backdrop of clear colorless crystals. (on its own it may appear darker)
Unless you're trying to make an iron chloride, that acid that you have will be essentially useless to you if you can't purify it somehow. If you have the proper equipment to distill it, that may be a possible way. You can also just look for some better stuff. The kind that I found at the hardware store is 31.45% (~10 molar) and is actually surprisingly clear. If there's any iron contamination, it's not enough for me to notice.

On the topic of chlorides, I actually just posted a picture here of some copper(II) chloride that I made a long time ago. (Which I used the hardware store acid for. As you can see, there's no discoloration)

gardul - 27-10-2014 at 16:00

nice pictures. I have more copper chloride than one would ever need. I love making the stuff and it looks amazing under a microscope.

about the substance... when heated directly in a tube, it turns white. both white and yellow are not very soluble in water. and when both are put to a flame, it gives a brown and yellow flame.

Kitsune - 27-10-2014 at 16:39

Interesting, though for stocking a home laboratory I would stick to something a little more simple, cheaper (at least if you live somewhere HCl is rather hard to come by, like I do) and not to mention more safe; I hope you do not mind me sharing this...

My preferred method for obtaining Magnesium Chloride is the reaction between Magnesium Sulphate (usually Heptahydrate) and Calcium Chloride (sold as a desiccant, not overly pure but it's workable).

Using the closest to stoichiometric amounts as possible, dissolve both separately in minimum water, then pouring them together, the Calcium precipitates as Calcium Sulphate and the Magnesium stays in solution as Magnesium Chloride; be sure to stir it well.

Filtering by gravity is a very very slow and vacuum filtration is just as slow; my preferred way is to leave it to settle (24hr covered) and then decant the Magnesium Chloride bearing solution off from the Calcium Sulphate precipitate.

The Calcium Sulphate can be later washed several times to remove most soluble impurities and then dried; decomposing it to the hemihydrate (0.5 H20) will leave you with Plaster of Paris which is always a useful thing to have around, I have been known to make moulds for electrode casting with mine (I know it's a labour intensive way but it is fun).

The Magnesium Chloride can then be further purified by leaving it to stand (any Calcium Sulphate that was disturbed in decanting will settle), filtering and recrystallising (if required).

Also I will add, the Hydroxide freely precipitates from Magnesium Chloride, leaving Sodium or Potassium chloride in solution (depending on your base), should it be required later.

Best Regards
Kitsune

blogfast25 - 28-10-2014 at 09:25

Quote: Originally posted by UnintentionalChaos

Adding carbonate or bicarbonate to a magnesium salt solution precipitates magnesium carbonate and/or basic carbonate.

I doubt strongly that that is the case with bicarbonate. A solution of NaHCO3 (max about 10 w%) is fairly alkaline (pH ≈ 11) and contains very few carbonate ions.

Furthermore, the Ks of MgCO3 is 10-7.8, the Ks of Mg(OH)2 is 1.5 x 10-11, so the hydroxide is the less soluble of the two.

All this points to Mg(OH)2 as the precipitate.

Acc. A.F.Holleman (‘Inorganic Chemistry’) magnesium salts with sodium carbonate solution precipitate a basic carbonate of the formula

Mg(OH)2.3MgCO3.3H2O, which on treatment with CO2 yields the neutral carbonate.

[Edited on 28-10-2014 by blogfast25]

gardul - 28-10-2014 at 09:43

I never even thought of it actually be Mg Hydroxide. Thank you Blogfast25 for that post. I will also check out the text you mentioned.

I will ave to do another run. study the 1st precipitate a little more see if I can get a 100% in knowing what it is. If it is in deed Mg hydroxide, this method will be a lot easier then the past methods ive been using.

blogfast25 - 28-10-2014 at 10:13

If you're going to check that the obvious way (adding acid will give fizz in the case of a carbonate), make sure you wash the precipitate properly, so as not to get confused by any residual carbonate that might just happen to cling to it.

[Edited on 28-10-2014 by blogfast25]

gardul - 2-11-2014 at 23:33

I washed the new run 3 times. I know that may be a little excessive but I want to be sure it is clean. I added HCl do a sample and it fizzed. To my surprise, i do not think it is Mg(OH)2. I do need to run another test. There were a lot of bubbles and gases being formed in both stages of the experiment.

As of right I will being another run at this. I need to trap the Gases and see what is being produced and well. At that point I think it would settle this for sure. I could be wrong, but I would like to further my knowledge in this since it seems ( not only here) that one person says it is Mg carbonate and another Mg hydroxide.

[Edited on 3-11-2014 by gardul]

blogfast25 - 3-11-2014 at 10:10

Quote: Originally posted by gardul

There were a lot of bubbles and gases being formed in both stages of the experiment.

If CO2 was developed during mixing of the MgSO4 and the NaHCO3 that points to the following.

Bicarbonate is a weak base because of:

HCO3-(aq) + H2O(l) < === > H2CO3(aq) + OH-(aq)

If Mg precipitates as Mg(OH)2, then that causes the first equilibrium to move to the right (Le Chatelier Principle) because of the removal of OH-.

H2CO3 ('carbonic acid') is of course unstable and decomposes into water and CO2, which explains the fizz.

It's possible that your product is a basic carbonate, which would also fizz after washing and treating with HCl.