Sciencemadness Discussion Board - Aluminium chloride hexahydrate: strange behaviour?

Aluminium chloride hexahydrate: strange behaviour?

blogfast25 - 30-6-2011 at 07:44

I needed a fairly concentrated aqueous solution of AlCl3, so I thought I might as well make the solid hydrate AlCl3.H2O.

6.2 g (0.23 mol) of Al chips were added to 85 ml (100 g) of 37 % HCl (low Fe), gradually to avoid boiling over (the reaction is very exothermic, presumably due to great solvation energy of Al3+

.

The solution was then hot filtered and boiled in slowly. It’s borderline bumping but then after about ½ h on a rolling boil, boiling behaviour changed a lot: it became smoother. Shortly after that I stopped boiling and allowed to cool. Sure enough, a white material dropped out, with some more crystallites forming also at the top and a small ‘lake’ of solution squeezed between these two layers. I let it cool further to RT.

Hours later the whole mass had turned into a clear, thick liquid (no crystals observable) and after leaving the stuff overnight it had congealed into a moist butterlike mass, slightly off-white. Half a teaspoon of it was dissolved into about 10 ml of DIW: it forms a perfectly clear solution; no sign of hydrolysis whatsoever.

I’m now wondering whether I left enough water for the formation of the hexahydrate and if not, what the nature of the butterlike mass is: hexahydrate with (an)other hydrate(s)?

About 45 g of product was obtained (theor. = 56.3 g) and 20 ml of DIW has now been added to it and reheated to BP. It’s now cooling. But AlCl3 has a flat solubility temperature curve…

m1tanker78 - 30-6-2011 at 18:47

You probably lost a great deal of water (exothermic + 'misting') when you added the Al to the strong HCl. Barring that, it sounds like your buttery stuff is a polymerized variety of the chloride.

Tank

woelen - 30-6-2011 at 23:44

I do not think it has anything to do with polymerics.

I have had similar experiences with other chemicals. I once tried to make cobaltous nitrate. I did this by dissolving cobaltous carbonate in 30% nitric acid. I added so much cobaltous nitrate that almost all of it dissolved but a slight amount remained unreacted, even when the liquid was heated.

After that, I vacuum filtered the solution and I obtained a clear solution of a nice rose/red color, looking somwhat like red wine (slightly lighter). This liquid I boiled down and to my surprise I could keep on boiling until I had a kind of syrup. At that point I stopped boiling. On cooling down, the material became really sticky, it became a thick clear deep red syrup, almost like a paste. It remained like that for many hours. One day later though, it had solidified to a glass-like material, which only could be taken out of its container with great difficulty.

This glass-like material dissolves in water, giving a clear solution and to my surprise, this solution is very acidic.

I have the impression that when metal salts are made of acids and metals (or oxides/carbonates of metals) and not all acid is used up, that it becomes very hard to obtain true crystals and that instead a kind of paste/syrup is formed which still contains quite some acid besides the metal salt.

In this particular case, I took some of the glassy material and heated that again. It melted and became a syrup again and it started bubbling/foaming on further heating. At a certain point I could see formation of dense white fumes and the air above the material also became pale brown. At that point I knew that the heating was gobe too far and acid was driven off and decomposition occurred. What remained was a basic material, which still dissolved in water, but the solution of it is not entirely clear.

Probably the only way to get a nice crystalline salt is to boil down, but not too much and then let the solution stand in a warm and dry place, such that water can evaporate slowly from this.

blogfast25 - 1-7-2011 at 04:43

Tank:

It’s not very likely: a material is either polymeric or it isn’t.

Woelen:

Yeah, this isn’t the first time for me neither: I’ve had several similar experiences, notably with FeCl3, which took weeks to complete crystallisation. The glassiness of your cobalt salt points of course to amorphousness: a supercooled liquid ‘solidified’ but not crystallised.

The solution obtained yesterday by dissolving the ‘butter’ back in 20 ml water hasn’t budged and I doubt if it will (I think it’s now sub-saturated, going by Wiki’s solubility data). So it looks like I’ll have to reduce it again a bit.

It’d be interesting to hear from anyone here who has obtained the crystals of AlCl3.6H2O...

Arthur Dent - 1-7-2011 at 06:03

@ blogfast25: What were the sources of your aluminium and HCl?

I have been tempted for a while to synthetize this chemical, but was hesitant to use aluminium foil because of possible impurities. But your experiment prompted me to try it out... So I have rummaged through my parts bin and selected a nice little heat sink that looks very much like solid aluminium. I'll do a volume to weight ratio calculation to make sure it is.

Then i'll drop the aluminium in a RBF with the acid and when it's all dissolved, i'll heat the stuff (gently) in a sand bath.

I have read that Aluminium trichloride is notoriously tough to dehydrate, so i'll be happy just to obtain the hexahydrate in its liquid form. On that note, will it keep well in a glass bottle? Or will I need to seal it tight to avoid hydrolysis?

AlCl3•6H2O is said to be corrosive, does it have a high vapor pressure and does it emit toxic fumes at room temperature? I wonder because if I'm to keep some of it in my lab, it has to be relatively stable (all unstable and overly "gassy" or oxidizing chems are kept outside in my shed).

Robert

m1tanker78 - 1-7-2011 at 06:12

Hmmm, Aluminum chlorohydrate is a polymer.

Quote:

no sign of hydrolysis whatsoever.

How so?

Tank

woelen - 1-7-2011 at 06:51

@Arthur Dent: AlCl3.6H2O is not very corrosive. It even is used as one of the active ingredients in deodorant. On the other hand, AlCl3 is very corrosive. When I open the bottle of AlCl3, then a lot of hissing noise is produced, like opening a bottle of Coca Cola, but instead of CO2, lots of intensely corrosive fumes of HCl are produced. Immediately after opening the bottle my hands and forearms are immersed in a nice cloud of HCl-fumes!

@m1tanker78: We are not talking about aluminium chlorohydrate, but about hydrated aluminium chloride. It might be, however, that on too much heating and on loss of hydrochloric acid a solution of AlCl3.6H2O may be converted to aluminium chlorohydrate. But formation of aluminium chlorohydrate only occurs if the solution of AlCl3.6H2O becomes deficient in HCl, due to hydrolysis.

[Edited on 1-7-11 by woelen]

Arthur Dent - 1-7-2011 at 07:52

@ Woelen: Gasp! I'll stick to the hexahydrate then!!! LOL

You mentioned that the solution may hydrolize if it becomes deficient in HCl, so does that means that it is by nature quite acidic and will emit a certain quantity of HCL fumes?

Is there a "perfect balance" where the solution will have just enough HCl to be stable and not emit too many HCl fumes? I have a 100ml pyrex orange-capped reagent bottle, will that be adequate to store that stuff safely?

I'm asking because anything I store in my lab must be stable, or in an airtight container so I can avoid all of my electronic eqipment and tools becoming "HCL hazy"... Anything that too reactive or overly oxidizing is relegatd to the shed, and I keep an absolute mimimum of stuff in there because it's icy cold in the winter and boiling hot in the summer...

EDIT: Sorry, blogfast25, I didn't mean to hijack your thread. I owe you a beer for that!

Robert

[Edited on 1-7-2011 by Arthur Dent]

blogfast25 - 1-7-2011 at 08:08

Well, at least it’s consistent: after reducing the solution (a bit) I got some crystals at the top, some at the bottom and a small amount of liquid. But it all already seems to be turning back to a liquid again… So I’m a bit at a loss how to proceed to powder now. I’m thinking recrystallisation from alcohol…

@ Robert: no problemo. AlCl3.6H2O is one of the safest synthesis you can undertake. It’s not a particularly toxic or corrosive substance (as hydrate - the anhydrous form can only be produced by direct action of dry HCl or dry Cl2 on the heated metal). I use the tabs from beer cans: nice size and easy to weight up.

Take care of one thing though: if using 20 % or more HCl, beware of just how exothermic the dissolution really is. It’s best to add your metal in three equal portions, allowing each portion to react away before adding the next, otherwise it’s a recipe for boil-overs (this is of course more true the finer your metal is - chunks are slower). It’s also best carried out outside: you get quite a lot of hydrogen (laced with HCl and some fairly smelly stuff)… Cover your beaker with something made of glass as a primitive refluxer.

[Edited on 1-7-2011 by blogfast25]

m1tanker78 - 1-7-2011 at 08:58

Quote: Originally posted by woelen

@m1tanker78: We are not talking about aluminium chlorohydrate, but about hydrated aluminium chloride. It might be, however, that on too much heating and on loss of hydrochloric acid a solution of AlCl3.6H2O may be converted to aluminium chlorohydrate. But formation of aluminium chlorohydrate only occurs if the solution of AlCl3.6H2O becomes deficient in HCl, due to hydrolysis

Hello Woelen,

Can the reaction and subsequent heating be performed with a reflux condenser to prevent hydrolysis. I've never had much luck otherwise (shooting for 6-hydrate crystals).

Tank

Arthur Dent - 1-7-2011 at 09:03

Quote: Originally posted by m1tanker78

(shooting for 6-hydrate crystals).
Tank

I think this might be a hard thing to accomplish, MSDS says the melting point of the hexahydrate is 0 deg. C

EDIT: Then again, another MSDS states 100 deg. C melting point... never mind the comment above...

Robert

[Edited on 1-7-2011 by Arthur Dent]

m1tanker78 - 1-7-2011 at 09:18

Quote: Originally posted by Arthur Dent

Supposedly, the melting point of the hexahydrate is 0*C. Though I didn't have any luck at ~ -10*C on a couple of prior attempts.

S.C. Wack - 1-7-2011 at 10:55

Does the method in Brauer not work?

blogfast25 - 1-7-2011 at 12:30

Quote: Originally posted by S.C. Wack

Does the method in Brauer not work?

Aha, S.C.Wack and Brauer to the rescue:

I quote from Brauer:

Aluminum Chloride Hydrate

A1C13 • 6 H2O

At 0°C, A1C13-6H3O is slightly soluble (21 mg./lOO ml.) in
saturated aqueous HC1 and is therefore easily isolated from such
a solution. The aluminum is dissolved in concentrated HC1 and the
solution is transferred into a three-neck flask fitted with a stirrer,
an inlet tube for HC1 gas and an outlet tube. The flask is cooled to
0°C and HC1 gas is introduced into the continuously stirred and
cooled solution until it is saturated. The inlet tube should not dip
into the solution, since it might become clogged with salt, but
sufficient absorption of the HC1 is ensured by vigorous stirring.
A wash bottle with concentrated H2SO4 connected to the outlet of
the flask serves to indicate the progress of the saturation. The
precipitated hydrate is rapidly suction-filtered and, while still
cold, washed with some ether and dried on a clay plate.

I might just try adding conc. HCl to my AlCl3 syrup and then ice it to see what happens…

As regards the melting point of 0 C, that must be in error. Wiki also states that number, alongside with a nice photo of a solid powder, presumably shot at RT...

[Edited on 1-7-2011 by blogfast25]

not_important - 1-7-2011 at 12:49

That's similar to the industrial method, using HCl(g) to saturate a cold solution of AlCl3. Drying with ether is preferred, the salt is soluble in the lower alcohols and ketones SFAIR.

Heating without an excess of HCl is likely to lead to chlorohydrates, adding HCl converts those back to the chloride by the conversion process may be slow as some of the chorohydrates form fairly stable clusters.

blogfast25 - 2-7-2011 at 09:39

The mass of AlCl3.6H2O this time had more or less held its own, with a harder and clearly crystalline mass at the bottom. To it I added 50 ml of fuming HCl and reheated to BP. The solid didn’t dissolve much, if at all. It’s now being iced to see if any more crystals form.

And I’m getting somewhat similar results with Al2(SO4)3. About 0.232 mol of this sulphate was synthesised by dissolving 12.5 g of Al chips into about 50 % H2SO4 (with a good excess to ensure speedy dissolution, so that’s theor. 79.5 g of (anh.) Al2(SO4)3. On cooling the whole mass crystallised into soft crystals, about 243 g of it.

100 ml of water was added and the mass redissolved and hot filtered. On cooling to RT no crystals formed at all. On fridging (about 5 C) overnight a small amount of well-formed crystals, 53 g, formed.

The whole amount of crystals, water and Al sulphate solution was then brought to the boil and 50 g of water carefully distilled off. On cooling again almost the entire mass crystallised, leaving only a few ml of solution.

Al sulphate has multiple hydrates and probably a complicated phase diagram with water. It seems to get crystals + solution the amount of water used is quite critical…

[Edited on 2-7-2011 by blogfast25]

Neil - 7-7-2011 at 17:20

I was recycling the aluminum chloride left from dropping a batch of tin. I tried filtering the last of the dropped tin from a moderately concentrated solution and it seems to me that the concentrated solution is rather non-Newtonian.

It has a weird viscosity, it's thick-ish like a syrup while not being “sticky”. Stirring it initially causes it to thicken before it suddenly thins out.

A stir bar spins with clear resistance before the fluid around it suddenly thins and then it turns with ease but only is able to affect a very limited area. As the speed is increased the rest of the fluid visibly thins and moves with sudden ease.

Vacuum filtering the solution at ~30°C caused the solution to massively thicken at the point where it ran through the filter and forms very syrupy, almost molten glass like, drips off the funnels tip. The syrup lands and gels, before slowly re-liquifying.

The vacuum was ~26”Hg. I ended up watering the 500ml solution down with 200ml water to get it to flow through the filter faster, it is now drying.

I found references to aluminum chloride solutions in combinations with other salts being non-Newtonian and for other aluminum salts, but not only aluminum chloride in water.

Some neat mechanisms at work here.

blogfast25 - 8-7-2011 at 06:15

That's interesting. Could very much point to polymeric structure: most high polymer melts behave non-Newtonianly.

My AlCl3.6H2O (under 37 % HCl, at RT) still hasn't dissolved, despite repeated stirring. I think it might be possible to get reasonable yields without HCl gassing, just using concentrated HCl added to a concentrated solution of AlCl3.

[Edited on 8-7-2011 by blogfast25]

Waffles SS - 11-8-2016 at 21:19

After adding AcOH to solution of aluminium chloride(made from Al and HCl) white precipitate similar to AlCl3.6HCl produce.I did it many times

AJKOER - 18-8-2016 at 11:44

Blogfast:

I suspect exposure of the Al/HCl mix to O2 is creating some of the inorganic polymer Aluminium chlorohydrate.

Logic: To cite one of my recent comments (link: http://www.sciencemadness.org/talk/viewthread.php?tid=22414#... ):

Quote: Originally posted by AJKOER

Per this source: Alok D. Bokare and Wonyong Choi, "Review of iron-free Fenton-like systems for activating H2O2 in advanced oxidation processes", published in Journal of Hazardous Materials, May2014. Link: https://www.google.com/url?sa=t&source=web&rct=j&...

To quote from the article, page 126:

"The use of bare ZVAl (without surface modification or pre-treatment) as Fenton-type catalyst to generate HO• was first demonstrated by Bokare and Choi [86]. In the presence of O2, in situ generation of H2O2 and the subsequent decomposition into HO• was achieved by electron transfer from commercial ZVAl samples under acidic condition (Fig. 2b). After the dissolution of the native surface oxide (Al2O3) layer on ZVAl at acidic pH(pH ≤ 4) to expose the bare Al metal surface, the sequential generation of H2O2 and HO• was utilized for the oxidative mineralization of organic pollutants (4-chlorophenol, phenol, nitrobenzene and sodium dichloroacetate) [86]."

where ZVAl stands for zero valent aluminum.

Apparently, in addition to Al(OH)3, there are created some reactive oxygen species (ROS), including the superoxide radical anion, which in the presence of an acid forms H2O2, in situ. This can then drive a Fenton-like reaction with Aluminum going from a valent state of zero to +3, along with the creation of hydroxyl radicals, HO•, capable of mineralizing organic pollutants. My take on a possible radical pathway:

2 Al + 6 O2 = 2 Al(3+) + 6 O2•-

6 O2•- + 6 H+ → 3 H2O2 + 3 O2 (See, for example, https://books.google.com/books?id=PQ3OBQAAQBAJ&pg=PA491&... )

Al + 3 H2O2 → Al(3+) + 3 HO• + 3 OH-

Al(3+) + 3 OH- → Al(OH)3 (s)

....

It would not be surprising to me, in the presence of OH- and HCl, if also some Aluminium chlorohydrate, which has been described as a group of aluminium salts with the general formula AlnCl(3n-m)(OH)m, is formed. Interestingly per Wikipedia ( https://en.m.wikipedia.org/wiki/Aluminium_chlorohydrate ), to quote:

"Aluminium chlorohydrate can be commercially manufactured by reacting aluminium with hydrochloric acid."

which is along the lines of what you did. Note, the cited Brauer preparation above likely limits air exposure.

If I happen to be correct, managing exposure to oxygen would be one possible fix.
----------------------------------------

Interestingly also, in a system with generated hydrogen (from HCl + Al ) in the presence of any formed hydroxyl radicals, the following radical creation could take place as well:

•OH + H2 = •H + H2O

which is the atomic hydrogen radical. Source: See Buxton work,Table 2 on page 16, available at: https://www.google.com/url?q=http://scholar.google.com/schol...

The reactive properties of this radical have long been observed and attributed here to "nascent hydrogen". The formation of the latter is, if I recall properly, said to increase in very effervescent reactions, which, in my opinion, should increase the level of dissolved oxygen in solution, thereby feeding the Fenton-like reaction with Aluminum.

[Edited on 18-8-2016 by AJKOER]