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Author: Subject: Aluminium Sulphate woes
CHRIS25
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[*] posted on 6-5-2014 at 14:18
Aluminium Sulphate woes


I'll Keep it short. Aluminium dissolved in Sulphuric acid, ended up with a thick mushy white precipitate of Aluminium Sulphate. Left pieces of Aluminium in this mushy precipitate for three days, it thickened the precipitate a bit. I filtered and heated the precipitate to 95c, as expected, it melted into a perfectly clear liquid. Reduced the mass from 95 to 55mLs until no more visible evaporation of water, (tested this by sealing and heating further, no more noticeable condensation appeared). As it evaporated the liquid turned yellow. At this point I removed from heat and allowed to cool fully expecting to get aluminium sulphate precipitated since its very much less soluble at 20 than at 100c. Three hours later and one dip in the freezer - nothing, zilch - hopeless. The liquid has exactly the same viscosity as concentrated sulphuric acid and is of course acidic by litmus test. And No how on earth can this be sulphuric acid unless the aluminium decomposed which it should not have under 300c. So I am baffled by this whole mess.

SOLVED: I was trying to get the anhydrous because the octadecahydrate was such a pain to get. Just took a sample and added a few drops of water and out came the aluminium sulphate. I think now I figured how to get the .18H20 - you have to boil down all the way like I did, until the solution is yellow, then add water drop by drop in order to get a 'dryable' / non mushy precipitate. I wrongly assumed that I could get the anhydrous by simply heating but obviously not.

[Edited on 6-5-2014 by CHRIS25]

[Edited on 6-5-2014 by CHRIS25]
Ok so one extreme amateur learning slowly to make sense of what is happening: Looking at this
2Al(OH)3 + 3H2SO4 = Al2(SO4)3.6H2O
Now obviously I did not have any Al(OH)3 at all, just a heavily saturated Al2(SO4)3 solution. All I did was in reality was to provide (OH)- via water.

[Edited on 7-5-2014 by CHRIS25]




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[*] posted on 7-5-2014 at 04:46


Sounds to me that by concentrating the aluminium sulphate solution you had super saturated it (see also 'Hot Ice': sodium acetate), so it resisted crystallisation. By 'disturbing' it with some water the aluminium sulphate then crystallised out.



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CHRIS25
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[*] posted on 7-5-2014 at 13:16


Quote: Originally posted by blogfast25  
Sounds to me that by concentrating the aluminium sulphate solution you had super saturated it (see also 'Hot Ice': sodium acetate), so it resisted crystallisation. By 'disturbing' it with some water the aluminium sulphate then crystallised out.

Ooo - are you saying that, like sodium acetate, it was 'disturbance' rather than delivering OH- ions? I assumed it was the latter. I am repeating this right now, waiting for another batch to super saturate.




‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)

Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)

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[*] posted on 8-5-2014 at 05:02


It sounds much like a case of super saturation, yes.



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CHRIS25
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[*] posted on 8-5-2014 at 05:22


I just re- did the experiment, this time I really super saturated to the point where is was a very dark golden yellow, thick and very viscous. I let it cool and at the moment there is a creamy yellow tinted precipitate forming from the top down, liquid is underneath.



‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)

Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)

The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by precision and law. (me)
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[*] posted on 8-5-2014 at 05:28


Now allow it to stand: if crystallisation doesn't proceed further you're probably super saturated.

It's quite common: you can do it with water too. Super cooled clean water then freezes over all at once on disturbance.

[Edited on 8-5-2014 by blogfast25]




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[*] posted on 8-5-2014 at 06:51


uuugh i recall that 'periodic videos' on youtube has some video about this guy explaining how they created acetylsalicylic acid, when they had the supersaturated solution they needed to scratch the RBF's with metal spatulas (enough to create an sound most wont like) to get them to attach themselves to a surface

i should really get some useful amount of faily concentrated H2SO4 tho..




~25 drops = 1mL @dH2O viscocity - STP
Truth is ever growing - but without context theres barely any such.

https://en.wikipedia.org/wiki/Solubility_table
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[*] posted on 8-5-2014 at 07:22


Blogfast - you're right, it has crystallized into a thick cake, but looks like I'll need a drill to get it out of the beaker!

Antiswat - yes from aspirin, I was hoping to do this experiment one day, was this from aspirin? but I really don't know what you by "RBF" .




‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)

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[*] posted on 8-5-2014 at 07:30


I've broken several beakers trying to chip out cakes of crystals like that :mad:

edit: RBF = round bottom flask

[Edited on 5-8-2014 by MrHomeScientist]
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[*] posted on 8-5-2014 at 07:30


RBF = Round Bottomed Flask
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[*] posted on 8-5-2014 at 07:44


Quote: Originally posted by aga  
RBF = Round Bottomed Flask

RTVM

(right thankyou very much)




‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)

Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)

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[*] posted on 8-5-2014 at 09:02


Quote: Originally posted by MrHomeScientist  
I've broken several beakers trying to chip out cakes of crystals like that :mad:



Do as I do: use PP (or PE) containers, preferably with a dome shaped bottom. PP is sufficiently flexible (and of course not brittle), so that you can gently squeeze it to liberate the crop of crystals.




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[*] posted on 17-5-2014 at 14:49


Some comments in line with results noted above per Atomistry (link: http://aluminium.atomistry.com/aluminium_sulphate.html ):

"Aluminium sulphate, Al2(SO4)3, is prepared in the anhydrous state by heating the crystalline, hydrated salt. The latter melts in its water of crystallisation, swells up, and eventually leaves a porous, white residue of anhydrous sulphate.

The anhydrous sulphate has a density of 2.713 at 17°, and its specific heat (0° to 100°) is 0.1855. At a red heat it decomposes, leaving a residue of alumina; decomposition becomes appreciable at 770°. It dissolves slowly in water.

A solution of aluminium sulphate is readily prepared by dissolving aluminium hydroxide in dilute sulphuric acid. The solution crystallises with difficulty, the hydrate Al2(SO4)3.18H2O being deposited in thin, six-sided, nacreous plates. This hydrate has also been obtained in the form of tetrahedra. At low temperatures the hydrate Al2(SO4)3.27H2O separates in trigonal crystals (a:c = 1:0.5408). Other hydrates, with 16H2O, 12H2O, 10H2O, 9H2O, 6H2O, 3H2O, and 2H2O have been described. The hydrates with 9H2O and 10H2O are said to be precipitated by alcohol, and to absorb water from a damp atmosphere, forming the hydrate with 18H2O. The hexahydrate results from the action of concentrated sulphuric acid on the hydrate with 18H2O; the trihydrate forms regular tetrahedra. The system aluminium sulphate - water has not yet been systematically investigated.

The hydrate Al2(SO4)3.18H2O is practically insoluble in alcohol. Its density is 1.69, and its specific heat (15° to 52°) is 0.353. As a white, fibrous efflorescence on shale and other rocks, this hydrate occurs as the mineral alunogen."

Not sure if heating the anhydrous salt at 770 C is a practical path to SO3. Also per Wikipedia (link: http://en.m.wikipedia.org/wiki/Aluminum_sulfate ) to quote:

"When dissolved in a large amount of neutral or slightly alkaline water, aluminium sulfate produces a gelatinous precipitate of aluminium hydroxide, Al(OH)3. In dyeing and printing cloth, the gelatinous precipitate helps the dye adhere to the clothing fibers by rendering the pigment insoluble.

Aluminium sulfate is sometimes used to reduce the pH of garden soil, as it hydrolyzes to form the aluminium hydroxide precipitate and a dilute sulfuric acid solution. An example of what changing the pH level of soil can do to plants is visible when looking at the Hydrangea macrophylla. The gardener can add aluminium sulfate to the soil to reduce the pH level which in turn will result in the flowers of the Hydrangea turning a different color."

So slowly dissolving Aluminium sulfate in water and filtering is a possible path to dilute H2SO4.

[Edited on 17-5-2014 by AJKOER]
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[*] posted on 18-5-2014 at 00:23


@AJKoer - a lot of information you have repeated does not really help with the practical issues discussed; firstly I have never come accoss such a stubborn difficult precipitate in all of the precipitate reactions that I have done. After 6 attempts via Al + sodium hydroxide and Al + sulphuric acid, the same undryable glue-like soggy mass of globular annoyance that is unworkable is a s far as I can get, drying it will not. Aluminium Hydroxide in Sulphuric acid is exactly the same as Al in sulphuric acid due to the fact that you get different hydroxides at various PH's between 4 and 7. But the needed OH3 hydroxide begins at PH 5 so keeping the acidity just below this is the way to the Al Sulphate. As for precipitation - trying to get rid of water by evaporation is troublesome since when you heat the gooey globular sticky mass even to below melt temp of 90 it is a no - goer, since it starts to re-dissolve again at 40!!!
Summing up. getting a good Al sulphate solution tested with PH 0.1 step indicators (not the hydroxide) - piece of cake; getting the precipitated Porridge - easy stuff; getting the octadecahydrate - impossible, let alone trying to go below .18H2O




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[*] posted on 18-5-2014 at 05:51


Quote: Originally posted by AJKOER  

So slowly dissolving Aluminium sulfate in water and filtering is a possible path to dilute H2SO4.



No, it isn't. The hydrolysis of aluminium sulphate is only very partial. I've dissolved commercial aluminium sulphate many times and the solution comes out near perfectly clear. Your filtrate would contain far more aluminium sulphate than sulphuric acid.

@CHRI$25:

I've become a little unsure about your goal here. Is it to prepare solid aluminium sulphate hydrate?

Certainly dissolving Al scrap in H2SO4 or NaOH has been unproblematic in my experience. You need plenty of acid/NaOH (a good excess) and enough water. If you don't use enough water on cooling the solution will solidify to an unknown hydrate.

Solid aluminium sulphate hydrate (a mixed hydrate) must be possible to obtain by evaporating the water from a solution to the correct degree. Commercial grade aluminium sulphate is obtained by dissolving hydrated alumina in sulphuric acid (with some excess). A solid mass is achieved with a water content of about 43 %, see this product for instance:

http://oxfordchemserve.com/aluminium-sulphate-dye-mordant-ir...

This is a hard white mass, which obviously has been coarsely broken up ('kibbled'). It dissolves easily in hot water to a near perfectly clear solution.

My advice would be to try again, on a smaller scale and using smaller pieces of Al scrap. Try a 30 % H2SO4 solution, using a 30 % excess. After most of the Al has dissolved (getting those last stubborn bits to dissolve can take much time and might not be worth that time), allow to cool. If it completely solidifies you've probably not enough water in there.

[Edited on 18-5-2014 by blogfast25]




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[*] posted on 18-5-2014 at 11:16


I just tested the 'solid white mass' i obtained via the copper sulphate route, and it can't be aluminium sulphate.

Doesn't dissolve in hot water at all. Binned. Start again.




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[*] posted on 18-5-2014 at 12:08


What concentration of CuSO4 did you use?



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[*] posted on 18-5-2014 at 12:10


Therein lieth the answer ...

This was CuSO4 from Spent battery acid, so lead oxide, sulphate, SiC, some random polymers, insects of all natures etc etc etc.




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[*] posted on 18-5-2014 at 12:18


From spent battery acid? You've totally lost me now.



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[*] posted on 18-5-2014 at 14:14


Battery acid from discharged, damaged cells, with some 'desulphator' Magic Battery Rejuvenator thing in there.

Not pure H2SO4 by any stretch of the imagination.

This was then used to make CuSO4 by electrolysis with copper electrodes.

Blue it may have been, Pure Blue, i think not.

To a solution of this was added table salt and aluminium metal (not foil).

The Product may well have been Aluminium Flumpty2Wumbles, given the range of impurities.

It was certainly white, but not Aluminium Sulphate.

I'll try CHRIS25's route, via acid.

[Edited on 18-5-2014 by aga]




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[*] posted on 27-5-2014 at 10:51


Hi, after a lot of experimenting and working like a dog to understand how to interpret and write ionic equations, plus the fact that this is not a single replacement reaction, nor a precipitation reaction despite the insolubility of Aluminium sulphate at 20c, it seems that the sulphate ion is a spectator ion in the reaction between the acid and aluminium metal. In a nutshell, it seems absolutely impossible to get aluminium sulphate from this route. I tried 7 times using the two methods previously discussed. This eighth attempt I have revised as follows: We need to get the hydroxide by pushing the reaction to above PH of 5 or 6 then we avoid all the other hydroxide conglomerations, we only need the Al(OH)3, precipitate the aluminium hydroxide using, well, a hydroxide. Filter this off and then react stoichiometrically with sulphuric acid again but keeping, I imagine, the reaction only very slightly acidic if we follow the stoichiometry.

Why is this the route? If you never look at water again as H2O but as H+(OH)- - as an ionic compound in its own right then:

Al3+ + OH- + H+ + SO2- means that this is a double replacement reaction where the cation Al3+ can now bond with the anion SO2- and the cation H+ can bond with the anion OH-.

All our failures are due to the (so I now believe) the fact that the balanced equation 2Al + 3H2SO4 = Al2(SO4)3 + H2 is misleading. For a start off I smell SO2 so that was an indication that something is not right in the theory.

And if anyone reading this spots an error in my thinking, then of course correct me. But I can, after 7 days and much agonizing, no longer see how al + sulphuric acid will ever lead to anything other than a solution of aluminium sulphate mixed in with probable different aluminium hydroxides.


[Edited on 27-5-2014 by CHRIS25]

[Edited on 27-5-2014 by CHRIS25]




‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)

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[*] posted on 27-5-2014 at 12:52


Chris:

There is so much wrong in you reasoning that it's hard to tell where to start refuting it.

But I will tell you this. Whatever you smell isn't SO<sub>2</sub>. Impure metals when dissolved in acids do often emit an odour that is hard to describe but SO2 does not come into it, let me assure you of that. The sulphate ion would have to be reduced to SO2 for that and that is very hard to do and doesn't happen in these mild conditions. Instead it's the H<sub>3</sub>O<sup>+</sup> ion that is reduced to water and hydrogen, as discussed before.

I'm going to make this simple for you. I will prepare a solution of aluminium sulphate by means of sulphuric acid and aluminium, right before your very eyes, so to speak: recipe with photos and results.

You're now resorting to writing nonsense and I can't let that stand either. Sorry.

Water is not an ionic compound, nor is there a 'double replacement' going on. And sulphate ions remain SO<sub>4</sub><sup>2-</sup>.

[Edited on 27-5-2014 by blogfast25]




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[*] posted on 27-5-2014 at 14:10


Whoops! well, I made an attempt to reason, based on lack of experience but yet on what I thought I had learned over the weeks. So I made a fool of myself I really don't mind, I need to learn and need the input. A number of points though require a quick answer probably, hopefully:
1. Hydrogen is in group 1, it is a gas but can be a metal (under extreme conditions that I have read), but it is in group 1 and it is a cation, (as far as I am aware I know of no metals that hydrogen bonds with by itself) Oxygen is an anion, the definition of an ionic compound is in the bonding of a metal and a non-metal, ionic bonding? Then what is it?
2. Aluminium is more prone to hydrolysis than even iron, and I see everywhere that the ideal route to aluminium sulphate was through the Al(OH)3 + Acid. Nowhere did I ever read just the metal; why? copper and sulphuric works, iron and hydrochloric works, no hydroxide routes here. So that alone got me questioning everything since my own attempts could not get passed the gooey sticky very hydrated porridge mixture that would absolutely not dry down enough.




‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)

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[*] posted on 27-5-2014 at 14:31


Chris25 is not posting this stuff by random chance.
It's more a reasoned way to try to understand Why the damnned thing does not want to work, following many failed experiments.

I'll add my 5 to Chris25's 7, making 12 attempts, of which 11 failed dismally.
One 'worked', but is obviously contaminated with copper (faintly Blue crystals) and could well be something else (copper sulphate route).

We've been thru the maths, equations etc, and Chris25 has come up with a possible explanation.

Thinking of Water as as H+ an OH-, which it seems to vaguely behave as, i thought a Good idea, seeing as the experiment has failed to work so many times - i.e. failed to conform to the various Laws (as we understand them).

@blogfast25 - if you can post here a Working synthesis of Aluminium Suphate, please describe How the Product can be tested, what it looks like, basically How you Know it is the required Product.

There is a good probability that Chris25 and i have about 850kg in our garbage.

If 'Water is not an Ionic compound' would it OK for me to cite that in a Mis-Selling Class Action against the De-Ionised Water salesmen ?

[Edited on 27-5-2014 by aga]

[Edited on 27-5-2014 by aga]




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[*] posted on 27-5-2014 at 14:59


Aga, that is a lot of floculant! Ach I would have given up on this sulphate lark 2 weeks ago but for the fact that leaving it un-answered causes me more heart ache than the failure itself. Having done so many precipitant reactions and especially the aluminium potassium sulphate crystals, this one is extremely temperamental and awkward. The aluminium hydroxide route via neutralisation with acid in NaOH is easy, but the precipitated hydroxide is like trying to extract slimy jellyfish pieces with an oily spoon, filtering is impossible since it re-dissolves at the slightest disturbance. It only takes me 5 minutes and pocket money to buy the darn stuff, But if I do that I'm defeated and I have learned nothing. But thanks for the support.



‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)

Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)

The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by precision and law. (me)
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