Sciencemadness Discussion Board - Aluminium Sulphate woes

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.

blogfast25 - 8-5-2014 at 05:02

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

CHRIS25 - 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.

blogfast25 - 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]

Fantasma4500 - 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..

CHRIS25 - 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" .

MrHomeScientist - 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]

aga - 8-5-2014 at 07:30

RBF = Round Bottomed Flask

CHRIS25 - 8-5-2014 at 07:44

RBF = Round Bottomed Flask

RTVM

(right thankyou very much)

blogfast25 - 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.

AJKOER - 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]

CHRIS25 - 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

blogfast25 - 18-5-2014 at 05:51

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]

aga - 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.

blogfast25 - 18-5-2014 at 12:08

What concentration of CuSO4 did you use?

aga - 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.

blogfast25 - 18-5-2014 at 12:18

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

aga - 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]

CHRIS25 - 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⁺ + SO^2- means that this is a double replacement reaction where the cation Al³⁺ can now bond with the anion SO^2- 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]

blogfast25 - 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 SO2. 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 H3O+ 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 SO42-.

[Edited on 27-5-2014 by blogfast25]

CHRIS25 - 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.

aga - 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]

CHRIS25 - 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.

aga - 27-5-2014 at 23:46

http://www.researchgate.net/post/What_is_the_difference_betw...

Interesting discussion of what DIW actually is, and how it is produced.

blogfast25 - 28-5-2014 at 05:05

http://www.researchgate.net/post/What_is_the_difference_betw...

Interesting discussion of what DIW actually is, and how it is produced.

From that source:

"The H+ and OH- combine to form H2O, leaving only the residual H+ and OH- produced by autodissociation (autoprotolysis), H2O = H+ + OH-; the equilibrium constant of this reaction = 1 x 10^-14 at 25 °C"

Even though the H+/OH- is a misrepresentation, once you understand that statement, you'll understand why water is sooooooo NOT an ionic compound. Its dissociation is basically a zero followed by thirteen zeros. NOT ionic at all. The concentration of H3O+ (NOT H+, that's non-sensical) in neutral, distilled water is about 0.0000001 mol/L. Do you get it now?

Please get that filthy thought out of your heads AT ONCE. No kidding.

Will explain more later on.

Researchgate is a very low authority source, BTW. Careful with web pages: no place peddles more misinformation/disinformation than the Tinkerwebs.

[Edited on 28-5-2014 by blogfast25]

CHRIS25 - 28-5-2014 at 07:36

At the risk of fueling a filthy thought

I know that water is a polar covalent dipole molecule where the hydrogen is the weaker and has a slight positive charge and the oxygen has a higher pull over the hydrogen. But it has also the character of an ionic bond in that there is this metal + non-metal bond character. Surely, and withot wanting to appear pedantic, it can not be categorized into either box. I only actually started to investigate this because in all my reactions to date I have un-conciously treated water as if it did not exist, I mean it is water, and very few balanced reactions ever include water in their equations, yet water does play a huge role in affecting the direction of a reactions path. take the ferric chloride solution. Everything seems fine according to this:
Fe + 2HCl = FeCl₂ + H; yet this is misleading due to the fact that you need to add water, and upon doing this you risk then hydrolysis of the ferrous ion to Fe(OH)₃. So now we have to consider water as H⁺ and OH^-

This is why I decided that water is not that neutral solution that one simply ignores in balancing equations, but that water is actually a chemical in its own right that needs to be considered. And then of course we have the same thing happening with aluminium sulphate, this too risks hydrolysis without careful attention to the amount of OH ions balanced against the amount of SO₄ ions available. So my intention is not to argue a point, simply to understand a concept of chemistry that I have ignored because I perceived water, sub-conciously, as a non reactant in chemical equations.

blogfast25 - 28-5-2014 at 09:50

I only actually started to investigate this because in all my reactions to date I have un-conciously treated water as if it did not exist, I mean it is water, and very few balanced reactions ever include water in their equations, yet water does play a huge role in affecting the direction of a reactions path. take the ferric chloride solution. Everything seems fine according to this:
Fe + 2HCl = FeCl₂ + H; yet this is misleading due to the fact that you need to add water, and upon doing this you risk then hydrolysis of the ferrous ion to Fe(OH)₃. So now we have to consider water as H⁺ and OH^-

Certainly water water cannot be ignored in aqueous reactions, yet very often it is nothing but a spectator solvent.

The dissolution of iron in hydrochloric acid to ferrous chloride (1), the oxidation of ferrous chloride to ferric chloride (2) and the precipitation of ferrous hydroxide (3) are three distinctly different phenomena.

(1) HCl is dissociated is water: HCl(aq) + H2O(l) === > H3O+(aq) + Cl-(aq)

Iron is oxidised by H3O+ but only to Fe(II):

Fe(s) + 2 H3O+(aq) === > Fe(2+)(aq) + H2(g) + 2 H2O(l)

(2) Fe(2+)(aq) === > Fe(3+) + e-

But that requires an oxidising agent and it isn't water. Oxygen, peroxide, etc etc can all act as oxidisers here.

(3) Hydrolysis of Fe3+, a multistep (equilibrium) reaction:

[Fe(H2O)6]3+(aq) + H2O(aq) < === > [Fe(OH)(H2O)5]2+(aq) + H3O+(aq)

A second and third step then leads to Fe(OH)3. Here water plays an active part.

Please (PLEASE?) do forget about H+/OH-, H+ cannot exist in watery solution. H+ is a proton, nothing more. In water it immediately bonds to a water molecule: H+ + H2O === > H3O3+. There are NO free protons in watery solution.

[Edited on 28-5-2014 by blogfast25]

[Edited on 28-5-2014 by blogfast25]

aga - 28-5-2014 at 14:48

The most convincing argument would be to show how Aluminium Sulphate is made, and shown to be Aluminium Sulphate as the product.

Define the experiment, and i for one will do it and replicate it, and post the results.

I think CHRIS25 will too, so two verfications.

Once that works as defined, the ionic water which is not ionic argument may be safely culled for now.

[Edited on 28-5-2014 by aga]

aga - 28-5-2014 at 15:15

This is getting creepy now.

Given the sheer Wealth of knowledge and experience out there, and plugged into this board, it is incredible that nobody has jumped in and said 'Nah. This is how it's done. Here's where you went wrong.'

Prime derivations from the known facts are :-

1. Aluminium Suphate is too hard to do.
2. Ignore noobs doing pontless,unexplosive or psychoactive stuff.
3. Aluminium Sulphate is useless, so I as Guru, never tried to make it.
4. It is hard to make, and largely pointless so why bother.
5. I dunno. Never looked.

Thank you blogfast25 for your guidance, as you are our only guidance so far on this topic.

Silence speaks Volumes.

If Aluminium Sulphate synthesis is difficult, pointless and un-useful (regarding the reactions and ionic states) then somebody saying so, and Why, would be very much appreciated.

[Edited on 28-5-2014 by aga]

CHRIS25 - 28-5-2014 at 23:44

Just a quick reply aga. My al solution is progressing with extreme interest. Placing clingfilm and poking 1 hole with a cocktail stick works wonders at preventing any visible loss of water, and so I have complete control now, the hole is simply for the hydrogen. Secondly after 25 hours (over three days) of cut aluminium metal pieces dissolving I have to switch it off every night. In the morning there is a precipitate but still with metal there of course. I have heated up and cooled several times and seen how the quality of the precipitate changes. this morning's precipitate was very different, it is most certainly al sulphate but it had a distinctive crystalline structure that I had not seen in other experiments, pressing lighly with a wooden stick you can feel it crumble, this is very different from the gooey mass of globulated porridge we have had before. There is still fluid in the flask and when you shake the flask vigorously it all but re-dissolves into a milky solution which is expected. So adding a touch more water you get the milkiness disappearing into a very clear solution, the solution now is no longer saturated. Solubility at 20c is about 36g in 100 mLs water so the trick here is to calculate amount of water left over after all is dissolved and do some maths. In view of the fact that this precipitate seems to be quite unique in that I read about the unpredicatability of the number of hydrated variations you can get that includes every single number from .2H₂O to . 28H₂O this explains part of the problem we are having (i am not yet aware of any hydrate compounds or crystals that can have such a variation of water molecules attached to their structure. Also I now realize that you really have to stoichiometrically calculate water content and make sure that the amount of water left in the flask when all is dissolved and you switch off the heat to cool, this water has to be very precise, something I have never needed to do in any other reaction to date. (In the past one just boils down and evaporates subjectively). So when my aluminium is fully dissolved I now feel confident about getting an exact hydrate (the .2H₂O), after some maths. We shall see whether or not I have to humiliate myself by re-tracting the above later on.....

[Edited on 29-5-2014 by CHRIS25]

So I figured this out, it looks good, but only the ongoing experiment will confirm:
2Al + 3H₂SO₄ = Al₂(SO₄)₃ +3H₂

The net ionic equation: 2Al⁰ + 3H₂⁺ + 3SO₄^2- = Al₂³⁺ + (SO₄)₃^2- +3H₂⁰

But since we know that the sulphate comes in many many different hydrates (as we have experienced) we need to calculate water at the very end just before cooling and evaporate or add accordingly in thus manner:

Water = 18g/mol so if we want to insure aluminium sulphate as Al₂(SO₄)₃.2H₂O then calculating the theoretical yield of our product (will show you if you need help here) or just assuming the mols Al we already used (either way), we can do the following:

0.5 mol Al₂(SO₄)₃ + 2 attached water molecules to each ionic molecule needs 0.5 x 2
thus: 0.5 x 2 = 1. And so this must be 18g/mol, in other words 18 mLs water must be left in the flask BEFORE any cooling off is allowed, but certainly no more than this. This is theoretical and I would leave a little less say 15 to allow for always a lesser yield than predicted. And this whole idea only came about since reading that this sulphate could precipitate out in dozens of hydrates ranging from 2 to 28. Allowing for the fact that an average room is 14c I would put it in the fridge, at 18 mLs at 0 degrees we have solubility of 31g per 100 mLs . So the fridge methinks would ensure absolute precipitation with just 18 mLs of water, Mmm...

Now at this juncture I am expecting a : "Chris all your reasoning is wrong, in fact it is absurd - with a slap on the wrist" But secretly I am hoping that there is sense to be found hidden somewhere in this brainstorming.

[Edited on 29-5-2014 by CHRIS25]

[Edited on 29-5-2014 by CHRIS25]

[Edited on 29-5-2014 by CHRIS25]

blogfast25 - 29-5-2014 at 04:52

Aga:

All in good time. I'm a bit stretched here but hope to do it tonight, if not Sunday.

You're reading far too much into too little.

So I figured this out, it looks good, but only the ongoing experiment will confirm:
2Al + 3H₂SO₄ = Al₂(SO₄)₃ +3H₂

The net ionic equation: 2Al⁰ + 3H₂⁺ + 3SO₄^2- = Al₂³⁺ + (SO₄)₃^2- +3H₂⁰

We would not write the equations the way you do. 3H₂⁺ and (SO₄)₃^2- are definitely nonsense.

Will elaborate tonight. To write reaction equations (and not just stoichiometry equations) you need to break down what is happening and then write the corresponding equations, you're not doing that...

[Edited on 29-5-2014 by blogfast25]

CHRIS25 - 29-5-2014 at 05:03

This is exactly how it is demonstrated on so many video tutorials, admittedly maybe I should have called it simply the ionic equation, or showing the charges, because I have not really done the "net" ionic I suppose, but the charges are correct and maybe on the right side I should have typed Al2(SO4)3 since this is the precipitate and all that actually changes is that the H reduces and the Al oxidizes. I really do not see what I am doing wrong, yet again?

blogfast25 - 29-5-2014 at 05:15

Many video 'tutorials' are made by people who couldn't tell beryllium from carbon. 'A little knowledge is a dangerous thing'.

No serious chemist would write these equations the way you do. They are plain wrong. 3H2+ or even H2+ e.g. doesn't exist although it would be a highly interesting quantum system if it did

. Ditto your 'trisulphate' thingy.

Laters...

[Edited on 29-5-2014 by blogfast25]

CHRIS25 - 29-5-2014 at 05:17

ah well then, I'll get back to my embroidery.

http://www.occc.edu/kmbailey/chem1115tutorials/Net_Ionic_Eqn...
just one of many tutorials, and most of the videos are from chem univerisities or schools, but they all follow this type of notation, The only thing I missed was to put aq, i, or s afterwards.

[Edited on 29-5-2014 by CHRIS25]

AJKOER - 29-5-2014 at 05:58

Here is an experiment that is on my to do list. Not sure if it will work, so if anyone else attempts it, please limit the size of the experiment.

1. Prepare a clear jelly like form of Al(OH)3 from say the action of dilute ammonia on Aluminum foil (slow, days).

2. Remove the solution over the precipiate, add distilled water and stir to dilute out the NH3 (I suspect this key step will have to be repeated several times).

3. Add the product, Al(OH)3 as a clear pure gel, to a concentrated MgSO4 solution.

If a white precipitate of Mg(OH)2 forms, a positive indication of the following reaction:

2 Al(OH)3 + 3 MgSO4 ----) Al2(SO4)3 + 3 Mg(OH)2 (s)

One could also more rapidly form Al(OH)3 from, say, boiling Al in boiling Na2CO3 to produce Al(OH)3 as a white precipitate. Filter and thoroughly wash as leftover Na2CO3 will react with the MgSO4. If it dissolves in aqueous MgSO4 with the formation of a visibly different (?) white precipitate, then possibly a more rapid path.

[Edited on 29-5-2014 by AJKOER]

blogfast25 - 29-5-2014 at 06:42

[...] but they all follow this type of notation, The only thing I missed was to put aq, i, or s afterwards.

It's not the type of notation that I object to (I use it too) but what you do with it: there is no such thing as (SO4

32- and you have never seen this on a reputable site. N-E-V-E-R or in any half decent book. Please don't tell me otherwise.

Please ignore the village idiot AJKOER. Here he comes again peddling his Al + NH3 imbecility. He'll never learn.

[Edited on 29-5-2014 by blogfast25]

blogfast25 - 29-5-2014 at 08:01

The displacement reaction:

2 Al(OH)3 + 3 MgSO4 ----) Al2(SO4)3 + 3 Mg(OH)2 (s)

... cannot proceed: poorly soluble as Mg(OH)2 is, it is far, far more soluble than Al(OH)3 (in neutral conditions).

Replacing NH3 with sodium carbonate changes little: both are weak bases.

AJKOER - 29-5-2014 at 09:18

The displacement reaction:

2 Al(OH)3 + 3 MgSO4 ----) Al2(SO4)3 + 3 Mg(OH)2 (s)

... cannot proceed: poorly soluble as Mg(OH)2 is, it is far, far more soluble than Al(OH)3 (in neutral conditions).

Replacing NH3 with sodium carbonate changes little: both are weak bases.

Yes, I agree with your comment with the qualifier "neutral solution". However, Blogfast, please explain how a concentrated MgSO4 solution, the product of a strong mineral acid and a very weak base, Mg(OH)2, could produce a neutral solution?

My experience (in agreement with your unqualified statement) is that the white precipitated Al(OH)3 is still hard to dissolve in even non-neutral solutions, but the clear geltaneous form of Al(OH)3 (formula Al(OH)3(H2O)3 or commonly still expressed as Al(OH)3 ignoring water, see http://www.science.uwaterloo.ca/~cchieh/cact/applychem/alumi... ) in mildly acidic solutions as compared to solid Mg(OH)2 in the same medium?

I agree the odds are still not likely, but excuse me if I want to verify as it may provide a novel path to other Aluminum salts as well.

[Edited on 29-5-2014 by AJKOER]

DraconicAcid - 29-5-2014 at 09:31

Magnesium hydroxide is a very weak base only because of its low solubility. The magnesium ion does not hydrolyze appreciably in aqueous solution. Merck Index says "Its aqueous soln is neutral. pH = 6-7"

blogfast25 - 29-5-2014 at 09:35

It looks like I won’t be doing this today: it’s raining and part of it has to be done outside. So I’ll go back to theory a bit.

1) Sulphuric acid solution:

When H2SO4 is mixed with water, the acid deprotonates as follows:

H2SO4(aq) + 2 H2O(l) < === > 2 H3O+ (aq) + SO42-(aq)

Although this is strictly an equilibrium, sulphuric acid is a very strong acid and the deprotonation is more or less complete (this is a slight simplification but OK for this discussion).

2) Oxidation and solvation of the aluminium:

The oxonium ions oxidise the Al to Al cations:

Al(s) + 3 H3O+ (aq) === > Al3+(aq) + 3/2 H2(g) + 3 H2O(l)

Solvation of the Al3+ ions:

Al3+(aq) + 6 H2O(l) === > [Al(H2O)6]3+(aq)

This hexaaqua aluminium cation is how Al exists in strongly acidic watery solutions.

3) Hydrolysis of [Al(H2O)6]3+(aq):

This cation is prone to hydrolysis acc.:

[Al(H2O)6]3+(aq) + H2O(l) < === > [AlOH(H2O)5]2+(aq) + H3O+ (aq) (I)

A second and third step of deprotonation can give rise to solvated [Al(OH)2]+ and even Al(OH)3 (hydrated alumina, aka aluminium hydroxide).

But here’s an important note: when dissolving aluminium metal in acids, we use a strong excess of acid (I will be using 30 % in excess of the stoichiometric requirement) and the excess acid, which provides extra H3O+, pushes the equilibrium reaction (I) (as well as the two subsequent deprotonation reactions) strongly to the left and explains why no Al(OH)3 forms. In short, in strongly acidic solutions there is no significant hydrolysis of the hexaaqua aluminium cation.

4) Crystallisation of aluminium sulphate hydrate:

As long as the solubility limit is not exceeded (for Al2(SO4)3 that is 36.4 g / 100 ml of water, at 20 C (Wiki) but it depends very strongly on temperature) the solution consist of [Al(H2O)6]3+, oxonium ions, sulphate ions, the excess of acid and of course, water.

When we exceed the solubility limit, solid aluminium sulphate hydrate must crystallise out. Assuming it is mostly octadeca (18) hydrate:

2 [Al(H2O)6]3+(aq) + 3 SO42-(aq) + 6 H2O(l) === > Al2(SO4

.18 H2O(s)

As you can see, only if crystallisation occurs, does the sulphate ion reappear in the equation, that’s why we call it a spectator ion. In essence it is there to ensure electrical neutrality of the solution/crystals.

[Edited on 29-5-2014 by blogfast25]

blogfast25 - 29-5-2014 at 09:38

[...] if I want to verify as it may provide a novel path to other Aluminum salts as well.

[Edited on 29-5-2014 by AJKOER]

Twits like you should try SIMPLE things, not 'novel' paths that are likely not to work or prove highly impractical.

Thanks, DA. Faced with such stubborn stupidity help is appreciated.

[Edited on 29-5-2014 by blogfast25]

AJKOER - 29-5-2014 at 10:10

Excuse my simplicity, but the compound we should be comparing for relative solubility to Mg(OH)2 at near neutral pH is (see my edited thread for clarity and reference link) is Al(OH)3(H2O)3 and not any published value for Al(OH)3. Note, higher aqua cations like the hexaaqua aluminium cation discussed above, are very soluble.

As such, I still argue there is some possibility of success.

If I do succeed, unless there is some modification of the dialogue on what l have clearly outlined as a speculative preparation, sorry, but I may not even consider sharing on this forum. But the debate has fueled my desire to perform this experiment for certain!

[Edited on 29-5-2014 by AJKOER]

blogfast25 - 29-5-2014 at 10:16

If I do succeed, unless there is some modification of the dialogue on what l have clearly outlined as a speculative preparation, sorry, but I may not even consider sharing on this forum.

Could you make that: 'I won't share on this forum', please? With a firm signature?

Thank you.

[Edited on 29-5-2014 by blogfast25]

AJKOER - 29-5-2014 at 10:30

Interestingly Blogfast, no chemistry based retort!

Please don't tell me I am starting to make sense, I am creating doubt, or worst you are starting to worry that my speculation is even possibly correct?!!

[Edited on 29-5-2014 by AJKOER]

CHRIS25 - 29-5-2014 at 11:00

It looks like I won’t be doing this today: it’s raining and part of it has to be done outside. So I’ll go back to theory a bit.

1) Sulphuric acid solution:

When H2SO4 is mixed with water, the acid deprotonates as follows:

H2SO4(aq) + 2 H2O(l) < === > 2 H3O+ (aq) + SO42-(aq)

Although this is strictly an equilibrium, sulphuric acid is a very strong acid and the deprotonation is more or less complete (this is a slight simplification but OK for this discussion).

2) Oxidation and solvation of the aluminium:

The oxonium ions oxidise the Al to Al cations:

Al(s) + 3 H3O+ (aq) === > Al3+(aq) + 3/2 H2(g) + 3 H2O(l)

Solvation of the Al3+ ions:

Al3+(aq) + 6 H2O(l) === > [Al(H2O)6]3+(aq)

This hexaaqua aluminium cation is how Al exists in strongly acidic watery solutions.

3) Hydrolysis of [Al(H2O)6]3+(aq):

This cation is prone to hydrolysis acc.:

[Al(H2O)6]3+(aq) + H2O(l) < === > [AlOH(H2O)5]2+(aq) + H3O+ (aq) (I)

A second and third step of deprotonation can give rise to solvated [Al(OH)2]+ and even Al(OH)3 (hydrated alumina, aka aluminium hydroxide).

But here’s an important note: when dissolving aluminium metal in acids, we use a strong excess of acid (I will be using 30 % in excess of the stoichiometric requirement) and the excess acid, which provides extra H3O+, pushes the equilibrium reaction (I) (as well as the two subsequent deprotonation reactions) strongly to the left and explains why no Al(OH)3 forms. In short, in strongly acidic solutions there is no significant hydrolysis of the hexaaqua aluminium cation.

4) Crystallisation of aluminium sulphate hydrate:

As long as the solubility limit is not exceeded (for Al2(SO4)3 that is 36.4 g / 100 ml of water, at 20 C (Wiki) but it depends very strongly on temperature) the solution consist of [Al(H2O)6]3+, oxonium ions, sulphate ions, the excess of acid and of course, water.

When we exceed the solubility limit, solid aluminium sulphate hydrate must crystallise out. Assuming it is mostly octadeca (18) hydrate:

2 [Al(H2O)6]3+(aq) + 3 SO42-(aq) + 6 H2O(l) === > Al2(SO4

This raises so many more questions than it gives answers, i do understand mostly, and certainly get it. But I could spend a whole afternoon on asking questions. So I think the best thing to do is simply comment and say that I am astounded and yet suspected this as so, that WATER plays much more of part and poses more of a threat to this whole experiment. It would be absolutely impossible to work all this out (the water part) without experience. None of this could be predicted unless you are an experienced chemist I suppose. There is so much here that annoys me to be honest. Simply because I have watched literally dozens and dozens of ionic equations, reaction theories, goodness knows what else, and all from colleges and universities, not bumpkins on the side street, but yet not even one single paper or video has ever covered this topic of how water interferes with reactions not even mentioning reactions involving the H and O ions in water. This is why I am mad. I could never have worked this out. I simply slowly realised that water was a problem and I had to solve it in this reaction that I am still doing.

I guess to sum up this is an extremely complex reaction and too advanced for beginners really. Th reaction can be performed but the water involvement is a bit too difficult to understand and how you arrive at all those figures and formulas for the involvement of water. Yes I am still mad, mostly because nobody covers this topic in any thing I have read or watched.

[Edited on 29-5-2014 by CHRIS25]

blogfast25 - 29-5-2014 at 11:23

Interestingly Blogfast, no chemistry based retort!

Please don't tell me I am starting to make sense, I am creating doubt, or worst you are starting to worry that my speculation is even possibly correct?!!

[Edited on 29-5-2014 by AJKOER]

I'll gladly put my hand in the fire that you are wrong.

The Ksp of Al(OH)3 = 3 x 10-34. That value may vary a bit from source to source and from product to product etc.

For Mg(OH)2 we have 5.6 x 10-12.

For your idea to work you'd have to find a form of Al(OH)3 that's roughly (AT LEAST) 10-12/10-34 = 1022 more soluble than what the tables tell you. A 1 followed by 21 zeros: some Al(OH)3 is that!

Hand in fire, thousands of dollars betted: betting against you succeeding is the safest bet I'd ever have placed.

http://www.ktf-split.hr/periodni/en/abc/kpt.html

[Edited on 29-5-2014 by blogfast25]

blogfast25 - 29-5-2014 at 11:29

So I think the best thing to do is simply comment and say that I am astounded and yet suspected this as so, that WATER plays much more of part and poses more of a threat to this whole experiment.

The water doesn't pose a 'threat'. You just need enough of it. Most of it gets recycled: turned into oxonium ions, which the in turn react away to hydrogen and water, for instance.

Hopefully the demo will shed some more light.

aga - 29-5-2014 at 11:49

Wow.

No wonder us noobs are confused.

In my ignorance, i made up 4 pots of some aluminium + sulphuric acid solutions today.
The quantities were carefully worked out, then adjusted - the acid volume increased by 30% and some more water added.
They are all reducing on the hotplate at the mo.

I'll post the results later, if there are any worth reporting.

AJKOER's suggestion almost had me throw some aluminium in ammonia, but it occurred to me that i have no idea what aluminium does in an organix mix, so decided against it.

As a great help, could someone say what the products look like ?
i.e.
is aluminium hydroxide Grey in solution (aka greysludge)
is aluminium sulphate .18H2O sort-of wet, a bit like toothpaste ?

blogfast25 - 29-5-2014 at 12:00

Aga:

To make matters more complicated: if you are going to evaporate water from an Al sulphate solution until solid Al sulphate hydrate starts appearing, then chances are you'll be crystallising a lower hydrate, not 18. The highest hydrates tend to crystallise at the lowest temperatures. Unfortunately I don't have a two phase system for Al sulphate/water.

Throw some Al in even strong NH3, by all means. Almost nothing will happen: NH3 is a very weak alkali, too weak to dissolve Al at any appreciable rate.

"Al hydroxide in solution" is an oxymoron: if it dissolved it's not hydroxide anymore. Freshly precipitated Al(OH)3 is a white (but it can appear greyish), voluminous precipitate.

Freshly prepared Al sulphate 18 hydrate is kind of white with the texture of toothpaste, On drying it hardens up. The stuff I sell you could break windows with!

[Edited on 29-5-2014 by blogfast25]

CHRIS25 - 29-5-2014 at 12:03

So I think the best thing to do is simply comment and say that I am astounded and yet suspected this as so, that WATER plays much more of part and poses more of a threat to this whole experiment.

Don't you mean I need as little as is possible (nearing the end of all metal dissolving) so that I end up with a saturated solution of Al sulphate Before I then cool it. To keep the solubility at the lowest possible place around 30 or so grams per 100 mLs at 0 degrees?
And I also suppose that this is why you never read Al + sulphuric acid in any document, it is always aluminium hydroxide that people use to react with sulphuric acid. This prevents hydrolysis and all the complications, is this more controllable?

[Edited on 29-5-2014 by CHRIS25]

[Edited on 29-5-2014 by CHRIS25]

aga - 29-5-2014 at 12:13

Freshly prepared Al sulphate 18 hydrate is kind of white with the texture of toothpaste, On drying it hardens up. The stuff I sell you could break windows with!

Brilliant. That helps enormously.
The Window test is new to me

So far we've seen quite a lot of toothpaste.
In my case, it failed to dry - ever.
It stayed toothpasty.

Would it not make more sense to drive off all the water and then rehydrate to get the actual hydrate you want ?

As regards AL + NH₃, i think i'll leave that to it's inventor.

blogfast25 - 29-5-2014 at 12:23

Chris:

If you are going to aim for as little water as possible near the end of the dissolution, you will run into other kind of problems. Very high concentrations of Al sulphate (low water content) tend to reduce the activity of the remaining acid and the dissolution will slow down, even stall. Better to have more water, then remove it later on by evaporation, if it is solid Al sulphate hydrate you want.

Al hydroxide + sulphuric acid is more discussed because industrially speaking starting from the metal doesn't make much sense. Al metal production is a very energy consuming process because it requires electrolysis at high temperature. Scrap Al metal is recycled mostly as Al metal, there's not much point in dissolving it in acid.

But for a chemist dissolving Al metal in various acids is a good way to obtain the various Al salts: sulphate, chloride, nitrate etc.

Aluminium hydroxide that is easy to dissolve in acids is not so easy to obtain either. Industrial producers of Al sulphate (Feralco, for instance) probably use taylor made grades for that purpose: dried to just the right degree to allow easy re-dissolution.

DraconicAcid - 29-5-2014 at 12:26

If your product is going to have the texture of toothpaste, why go for that product? It's much easier to make the alum (if you have access to KOH), which crystallizes nicely.

MrHomeScientist - 29-5-2014 at 12:27

Please don't tell me I am starting to make sense, I am creating doubt, or worst you are starting to worry that my speculation is even possibly correct?!!

If I had any doubts about your lack of credibility before, this has firmly erased them. The attitude you show here is the defining attribute of the crackpot. Next you'll be claiming there's some sort of "conspiracy" against your "research", and that all us "mainstream" scientists are too caught up in our "dogma" to see your correctness. We simply tire of refuting your nebulous, overly wordy posts of what you admit are completely untested and "in [your] opinion"-type reaction schemes.

If you actually sat down and did any of these reactions you talk about, and posted about your experiments on the board, then one of us would be proven wrong and that would be the end of the discussion.

Honestly I hope your ideas work, for then you will have added valuable knowledge and experience to the community. But armchair theorizing without ever doing any experimentation adds little.

As I saw elsewhere on the forum,
"There is no need to argue if an experiment can be made."
- H. St. C. Deville

=================
=================

I suppose I should add something on-topic here. I don't know how applicable this is to the discussion, but I've made potassium alum very easily before, which is a double sulfate salt: KAl(SO₄)₂ * 12H₂O. This crystallizes perfectly with no trouble. I start by dissolving Al in KOH solution. This produces a clear solution of soluble potassium aluminate. I then neutralize with H₂SO₄. This initially precipitates Al(OH)₃ as a white slime, but keep on adding acid and it all redissolves to sulfates. Finally, cool and evaporate to yield crystals. I don't see how making aluminum sulfate on its own would be difficult, but I could be wrong. I'm looking forward to reading about blogfast's experiment.

Edit: typo.

[Edited on 5-29-2014 by MrHomeScientist]

blogfast25 - 29-5-2014 at 12:28

[
Would it not make more sense to drive off all the water and then rehydrate to get the actual hydrate you want ?

Try it. Anhydrous Al sulphate exists, so it's only a matter of time and heat, in principle. Could take a while though. Record weight until no further weight loss is observed. That should be your anhydrous Al sulphate.

blogfast25 - 29-5-2014 at 12:33

Mr HS:

Yes, alum that way works well. Been there, done that.

As regards AJKOER, I so wish he'd take up crop circles 'research' as a hobby...

[Edited on 29-5-2014 by blogfast25]

blogfast25 - 29-5-2014 at 12:35

Quote: Originally posted by DraconicAcid

If your product is going to have the texture of toothpaste, why go for that product? It's much easier to make the alum (if you have access to KOH), which crystallizes nicely.

Because alum does not equate to aluminium sulphate? They are related, yet different substances. Alum is more 'interesting', I'll easily grant you that. But the single sulphate is also cheaper, as a dye mordant for instance.

[Edited on 29-5-2014 by blogfast25]

CHRIS25 - 29-5-2014 at 12:52

At this moment in time my solution is still dissolving after three days. The acid route is so slow. Anyway, I added a little more metal because there is more than enough excess acid. At this moment I have 0.14 mols aluminium and 0.43 mols of sulphuric in 165 mLs of that hated awkward substance called water

My intention is to dissolve everything - naturally - and chase away the water at high temperature to about, well, that will have to be calculated based on theoretical yield and solubility predictions, sould I say INsolubility prediction at 0 degrees, yes it's going into the fridge, (that is what I meant Gert sorry, probably my ineffective communication),

I have also done really big alum crystals, such a pleasure! But this sulphate purpose is about my pride, integrity and determination

Besides if i have enough over I might send some to aga since it kills spanish slugs.

For anyone that might be interested, though I presume everyone knows more than me anyway: Pages 23 - 25 are interesting here: http://pub.epsilon.slu.se/8299/1/torapava_n_110826.pdf

and this: http://lanthanumkchemistry.over-blog.com/article-hydrolysis-...

I only found these because I was able to ask the right question after what blogfast had said earlier on. If anyone has similar material that reasonably straightforward to swallow I would be happy to read.
[Edited on 29-5-2014 by CHRIS25]

[Edited on 29-5-2014 by CHRIS25]

aga - 29-5-2014 at 13:19

Quote: Originally posted by DraconicAcid

If your product is going to have the texture of toothpaste, why go for that product? It's much easier to make the alum (if you have access to KOH), which crystallizes nicely.

This is simply because we have not been sucessful at making Aluminium Sulphate.

To drop that and switch to an easier product would leave the Reasons Why unanswered, and nobody would learn anything.

The objective is to Learn, not to make Al₂(SO₄)₃ of any specific hydrate

Try it. Anhydrous Al sulphate exists, so it's only a matter of time and heat, in principle. Could take a while though.

Ok. Will do !

[Edited on 29-5-2014 by aga]

AJKOER - 29-5-2014 at 16:40

.......

The Ksp of Al(OH)3 = 3 x 10-34. That value may vary a bit from source to source and from product to product etc.

For Mg(OH)2 we have 5.6 x 10-12.

For your idea to work you'd have to find a form of Al(OH)3 that's roughly (AT LEAST) 10-12/10-34 = 1022 more soluble than what the tables tell you. A 1 followed by 21 zeros: some Al(OH)3 is that!
............

Why is the discussion focused on Al(OH)3 when the actual compound to be tested with concentrated aqueous MgSO4 is Al(OH)3(H2O)3, also called Aluminium triaquatrihydroxy complex.

[Edit] Found some interesting comments on this compound at http://www.chemthes.com/entity_datapage.php?id=4011 to quote:

"Real, long lived, electronically neutral reagent chemical.
Gram formula weight (molecular mass) = 132.04
Solubility in organics = insoluble
Solubility dependent upon pH. The trisaquatrishydroxyaluminium neutral complex is white, (even colourless), gelatinous material. Solubility is dependent upon pH."
----------------------------------------

Also, per this source http://www.drugs.com/pro/magnesium-sulfate.htm a 5% Epsom soultion can have a pH between 5.5 to 7, so much for the neutral solution assumption.

The comment I do agree with completely is why argue when one can do the experiment. I have already place some Al foil flakes in a vinegar solution, boiled, and let stand for a few hours. After rinsing off the flakes, which have been now stripped of protective coating, I did notice an immediate and obvious reaction with dilute household ammonia. The experiment proceeds!

[Edited on 30-5-2014 by AJKOER]

CHRIS25 - 29-5-2014 at 23:47

@ AJKOER: even if I did not understand it I would never take properties of chemicals from a drug store only from an authorised chemistry source. It's a drug store! Secondly I think you should read this more carefully: Each mL contains: Magnesium Sulfate heptahydrate 500 mg; Water for Injection q.s. Sulfuric acid and/or sodium hydroxide may have been added for pH adjustment THe Ph can be adjusted by the addition of .....it is not the inherent Ph of mag sulphate.

AJKOER - 30-5-2014 at 03:49

Chris:

A couple of points. First, the reason that the pH of the MgSO4 solution exposed to air can be as low 5.5 could be due simlpy because of CO2 (please see discussion at http://www.researchgate.net/post/What_is_the_difference_betw... ).

Second, drug/foodgrade compounds are generally low in heavy metals (a large problem if there are not). This makes them generally high purity and afforable chemicals for the home chemist.

CHRIS25 - 30-5-2014 at 04:05

I have anhydrous mag sulphate and some that has absorbed an unknown amount of water and is exposed for weeks. I have just measured its Ph using both a universal indicator strip and a comparitor strip with 0.3 increments. First one read between 7 and 8, the comparitor strip read 7.4

blogfast25 - 30-5-2014 at 04:33

Also, per this source http://www.drugs.com/pro/magnesium-sulfate.htm a 5% Epsom soultion can have a pH between 5.5 to 7, so much for the neutral solution assumption.

So you don't even understand what neutral actually means?

Do you know what the concentration of H3O+ is at pH 5.5? Here, let me help you: 10-5.5 = 0.0000003 mol/L.

Do you think this is going to dissolve your 'Al(OH)3.3H2O'?

You really, really haven't got the foggiest clue, do you?

I feel that if you keep on trolling here you are risking a repeat of your three week posting rights suspension of last year.

AJKOER - 30-5-2014 at 05:41

With respect to Al(OH)3(H2O)3, also called Aluminium triaquatrihydroxy complex, I repeat the interesting comments on this compound at http://www.chemthes.com/entity_datapage.php?id=4011 to quote:

"Real, long lived, electronically neutral reagent chemical.
Gram formula weight (molecular mass) = 132.04
Solubility in organics = insoluble
Solubility dependent upon pH. The trisaquatrishydroxyaluminium neutral complex is white, (even colourless), gelatinous material. Solubility is dependent upon pH."

The central, to be determined, question relates not to absolute solubility at any given pH, but relative solubility of this complex versus Mg(OH)2 at, say, pH 5.5 and the ability of the proposed reaction below to move to the right:

2 Al(OH)3(H2O)3 (aq) + 3 MgSO4 (aq) --?--) Al2(SO4)3 (aq) + 3 Mg(OH)2 (s) + 6 H2O

blogfast25 - 30-5-2014 at 06:26

You've already got your answer: whatever type of hydrated alumina you want to use it will have to be about 1022 more soluble than listed values. Good luck with that.

aga - 30-5-2014 at 08:56

seeing as the Product required is Aluminium Sulphate (some number)Hydrate,
surely the equation is this :

2Al + 3H₂SO₄ + nH₂O -> AL₂(SO₄)₃ . nH₂O + 3H₂

where n is the hydrate number required.

Also, seeing as this is an Ionic equation, the reaction needs some water as a solvent in whch to actually take place.

Via sheer dint of experimentation we now have a repeateable method to produce an unknown hydrate.

Pushing the theory forwards will require another set of experiments, which are on-going as i type.

This should hopefully vield a formula into which the required mass and hydrate required are entered, and dictate accurately the reactant volumes/masses required.

Thinking about it, having the Solvent end up in the Product is probably why this is a tricky balancing act.

[Edited on 30-5-2014 by aga]

[Edited on 30-5-2014 by aga]

blogfast25 - 30-5-2014 at 09:44

Here’s my aluminium sulphate hydrate experiment.

6.9 g of Al chips (these little levers used to open Al pop and beer cans with, each about 0.2 g), 191 g of water and 26 ml of 98 w% H2SO4 were loaded into a 500 ml beaker and heated on a hotplate, medium heat, maintaining heat all through the dissolution process.

The quantities correspond to about 0.25 mol Al, 20 w% H2SO4 solution and a stoichiometric excess of about 30 %. Here it is after about half an hour, vigorous H2 production can be seen (the white layer is hydrogen foam, nothing else):

After about 1 h the reaction had slowed down a lot and there’s still some unreacted metal. Here’s my little trick: the solution was ran through a sieve and the recovered unreacted metal mixed with some water and some fresh H2SO4, this dissolved those last bits quite easily and that solution was added to the mother liquor and mixed in. 25 ml of the final solution, now approx 0.5 M in aluminium sulphate (0.125 mol Al sulphate/0.25 L solution) was set aside in a smaller beaker and to the rest was added about 40 g of solid K2SO4. Both were then simmered; left: alum solution, right: 25 ml of about 0.5 M aluminium sulphate:

After about 5 minutes of simmering both solutions were perfectly clear and the alum solution transferred into a PP beaker for cooling, chilling overnight and crystallisation of alum. A seed crystal of alum was added.

The right hand side 25 ml aluminium sulphate solution I continued to boil until it started to get quite syrupy but still perfectly clear. I then stopped heating and observed. After a bit of cooling crystallisation started. Here it is seen about half way through (view from top of beaker):

Almost completely solidified, a few minutes later:

This product a kind of a waxy solid, harder than I expected and much stiffer than toothpaste. That suggests tentatively a lower hydrate. Commercial 14.3 hydrate is hard, like hard candle wax.

Watching this process confirms it’s a crystalline product, even though the crystals don’t appear well formed at all. Which hydrate (or mixture of hydrates) formed here would have to be determined separately.

On Sunday (I’m off to Liverpool tomorrow): yield determination of the alum and some tests on the aluminium sulphate hydrate.

[Edited on 30-5-2014 by blogfast25]

blogfast25 - 30-5-2014 at 09:49

Aga:

Strictly speaking your equation is called a stoichiometrical equation: it tells you what reacts with what, in what ratios and what you get and how much. Ionic equations are the ones I wrote on the previous page, post of 9:35.

"Al2" remains a no-no, though.

That notation suggests diatomic molecules of aluminium.

Yes, the fact that aluminium sulphate hydrate crystallises with so much lattice water has to be factored in, otherwise it's a recipe for trouble.

[Edited on 30-5-2014 by blogfast25]

CHRIS25 - 30-5-2014 at 10:58

Hi I am not quite following this bit: "....this dissolved those last bits quite easily and that solution was added to the mother liquor and mixed in. 25 ml of the final solution...."
Do you mean that you took the filtered solution and added it to just 25 mLs of the original solution? You split the one original solution into exactly two halves?

Why has mine taken three days and still not finished? Do aluminium can opener levers dissolve super fast then? I even added aluminium wire and this also seems to take almost two days - it's only 3.7g aluminium in total in 150 mLs water and plenty of acid.

My theoretical yields and limiting product:
Al is the limiting product
giving a yield of 23.58g of sulphate
Sulphuric acid theoretical yield is 48.88g of sulphate so you can see that I have loads of excess acid.
It's just that you did this in a day?

I know - I am being a pain now; but I would just like to know how you arrived at this, the second half of this equation
Al(s) + 3 H3O+ (aq) === > Al3+(aq) + 3/2 H2(g) + 3 H2O(l)
Everything hinges on understanding how you get to 3/2 H2 + 3 H2O. I take it by 3/2 you mean 1.5? I can count the Hydrogens and see the balance on both sides, it's the reasoning I am trying to work on. I studied today quite a lot about hydrolysis of different categories of metals and the acid and base hydrolysis, is it like this - that each hydronium donates a proton to the solution, the aluminium loses 3 electrons and one each go to each of the Hydrogen which then becomes the gas and this releases 3 extra H2O molecules?
[Edited on 30-5-2014 by CHRIS25]

[Edited on 30-5-2014 by CHRIS25]

[Edited on 30-5-2014 by CHRIS25]

[Edited on 30-5-2014 by CHRIS25]

aga - 30-5-2014 at 11:41

Yay !

Many many thanks blogfast !

I'll post our recipie when it has been verified by yet aother experiment, or two.

It was only in the last experiment when the idea came to make some sort of reflux mechanism (RBF on top of beaker) so there could be some idea of the water remaining ... and there you have it in shining technicolour !

[Edited on 30-5-2014 by aga]

Sorry about the 2Al₂. I got carried away with the 2s.

[Edited on 30-5-2014 by aga]

aga - 30-5-2014 at 12:29

Hooray ! Ours works too, at last (and verified by more than 1 experiment).

2.5g Al (whizzed up al foil was used)
63ml Water
9.5ml 12[M] suphuric acid

Mix in a beaker, cover with an RBF with a bit of water in it, and either wait overnight, or put on a hotplate for about 1 hour 30 mins.
The solution starts off grey, and becomes clearer towards the end of the reaction.
On the hotplate it goes frothy on top, and the bits of Al dance up and down a lot.

Filter, giving a Clear liquid.
Measure the liquid volume.
Assuming that this is all Water with stuff dissolved in it, calculate the amount of water required to make the Hydrate you want, and work out the Difference.

In this last experiment, 40mls was left, and a .16H₂ hydrate was the target, so there was 13mls too much water.
Boil off that quantity of water (the difference) then leave to cool (or put in fridge).

I put this last one in a bucket of water to cool, and it forms a skin within a minute.
A few seconds later, the whole mass begins to solidify.

For Anhydrous (possibly .2H₂O) just keep boiling until dry.
Be careful though, as the last pockets of water boiling off in the hardended mass cause explosions, and the rock-hard sulphate escapes from the beaker.

Many Thanks again blogfast25.

I'm glad you beat us to it.

All Fame, Glory and Endless Riches are rightfully yours.

blogfast25 - 30-5-2014 at 14:17

Chris:

I simply took out the unreacted aluminium by sieving (fitration, if you prefer) and set the filtrate aside. The metal that was caught on the sieve was added to a bit of water and fresh acid and then it reacted away quickly. That second solution was added to the part that was set aside. This way all the dissolved Al was in the same solution.

You aren’t heating, right? Remember that higher temperature speeds up chemical reactions?

Al(s) + 3 H3O+ (aq) === > Al3+(aq) + 3/2 H2(g) + 3 H2O(l) (1)

3/2 (“three over two”) = 1.5

This is simply the balanced redox reaction for:

Al(s) === > Al3+(aq) + 3 e- (2)

And:

3 H3O+ (aq) + 3 e- === > 1.5 H2(g) + 3 H2O(l) (3)

(2) + (3) = (1)

Clearer now?

[Edited on 30-5-2014 by blogfast25]

aga - 30-5-2014 at 14:38

Quote:

.16H2 hydrate was the target, so there was 13mls too much water.

CHRIS25 just pointed out that the 13mls was what was required, and not what needed boiling off ...
Bugger.

[Edited on 30-5-2014 by aga]

CHRIS25 - 31-5-2014 at 10:13

@Blogfast: """""You aren’t heating, right? Remember that higher temperature speeds up chemical reactions?""""
have I really projected such an ignoramous persona on this forum

(I hear some say yes...) that's like asking why my potatoes aren't cooking - were they being boiled? (actually I forgot to switch the electric on....)

But thankyou for your explanation anyway. So acid hydrolysis occurs when the sulphate and Hydrogen ion bond is broken by water resulting in the taking of a proton by water resulting in the hydronium ion. The H ion reduces the Al ion which is oxidised by the loss of 3 electrons, each of these electrons is then snapped up by Hydrogen which then results in a neutral hydrogen atom escaping now as gas, and hydronium ions returning as water molecules, but naturally all this is self repeating and ongoing for as long as there is excess H ions from the acid.

Is this ok as far as trying to break it down? I am trying to visualize it as a picture, simplified I am sure.

[Edited on 31-5-2014 by CHRIS25]

[Edited on 31-5-2014 by CHRIS25]

blogfast25 - 31-5-2014 at 11:49

Chris:

A few points:

(I) The speed at which Al (as well as other metals) dissolves in acids (or alkali, in the case of amphoteric elements like Al) depends on several factors, mainly:

1) Specific surface area: grades with more m2 surface area per kg will dissolve much faster than coarser materials
2) Alloying elements: since as we’re using scrap metal and not 99.9999 % aluminium, one type of scrap aluminium may dissolve more easily than another
3) Temperature
4) Acid/alkali concentration

So different experimenters will find different dissolution rates.

(II) For hydrolysis I prefer this explanation.

The hexaaqua aluminium cation (see above) tends to expel protons from its 6 ligand water molecules, these protons are then absorbed by solvent water molecules, to form oxonium ions. In watery solutions with high H3O+ concentration this tendency is suppressed. But in alkaline conditions (small H3O+ concentration) the tendency is favoured and hydrolysis can proceed all the way down to Al(OH)3.

(III) With regards to the oxidation of Al metal by the oxonium (aka hydronium) ion H3O+, when the latter ‘receives’ an electron (thus being reduced), a neutral hydrogen atom, here represented by H* (* indicates hydrogen's lone 1s1 electron), is split off:

H3O+(aq) + e- === > H2O(aq) + H*(g)

Two H* then combine to form diatomic hydrogen gas:

2 H*(g) === > H2(g)

The notation that is most commonly used to represent the oxidation of a metal, e.g. here Al:

Al(s) === > Al3+(aq) + 3 e-

... which I've used much above is actually a serious simplification that doesn't correspond well to reality. Unfortunately that reality will have to wait a bit before I try and describe it here, because I don't want to create an 'information overload' on your system. Maybe later...

Does all that make any sense at all to you?

[Edited on 31-5-2014 by blogfast25]

aga - 31-5-2014 at 13:06

I would argue that Physical Reality doesn't fit well with Current Models at all.

However, current models do seem to work, mostly.

Certainly i cannot currently offer a Better model, so go with the flow.

Success with a decent hydrate

CHRIS25 - 1-6-2014 at 03:01

It has the texture and hardness of a chocolate bar out the fridge when scraped, is actually whiter than the photo, and is 2.5M @ 0.125mol

Weight of this product is 76g
equivalent anhydrous would be 42.75
Logically therefore 33.25g is water

This is more like a .16H₂O:
33.24g water = 2 mol approx
0.125/2 = 16

[Edited on 1-6-2014 by CHRIS25]

aga - 1-6-2014 at 03:17

Congratulations !

Woe-lessness all round.

blogfast25 - 1-6-2014 at 04:01

@Aga:

The best models are the ones that work best, although they don't necessarily work very well. Further research and understanding then refines the theory, making it fit 'reality' better. I use quote marks because in the quantum world 'reality' is a far cry from what we normally understand by it.

@Chris:

I calculate yours more like 14.8 but that's due to your rounding off which is fine. Well done.

Bear in mind though that the number is kind of an average: not only does it probably contain multiple hydrates, it almost certainly contains small pockets of saturated solution (inbetween the small hydrate crystals). Assuming you refrigerated this material to obtain it, this small amount of solution will contain still about 25 w% of aluminium sulphate (expressed as anhydrate). Will try and elaborate a bit on this later on.

Those last bits of solution could probably be removed (if it was needed to do so) by vacuum drying, freeze drying or even prolonged drying in air.

Despite that, your product would definitely be usable. It would be interesting to determine Al content by, dare I say... titration?

[Edited on 1-6-2014 by blogfast25]

CHRIS25 - 1-6-2014 at 04:30

Interesting Blogfast, ""So small pockets of saturated solution and multiple hydrates."" I would not have considered this, but plenty to learn still I am interested, what I really need though is to see a visual of what you have described; I will titrate for Al content though. Obviously vacuum and freeze drying I can not do, unless I invest in a CO2 fire extinguisher, but air drying - I am sure that I read that this compound is hygroscopic.
I don't have EDTA and have no idea what else could be used.

[Edited on 1-6-2014 by CHRIS25]

blogfast25 - 1-6-2014 at 05:08

My Al sulphate, after refrigeration overnight is now also a quite hard, white mass. Although I didn’t do a mass balance I suspect it too is in the n = 14 – 16 range. For the alum I only got about 60 % actual yield but that is largely because I used a weaker solution than I normally would.

Even though I cannot find a two phase system for Al sulphate + water, I can construct something that is kind of a ‘next best thing’. I used the Al sulphate solubility data from Wiki’s table and converted them to weight % aluminium sulphate and plotted them against temperature:

The data point line is called the ‘saturation line’. To the left all combinations of w% and temperature form completely liquid, non-saturated solutions. To the right, hydrates begin to form, in equilibrium with a saturated solution (noted as 'liquid' on the diagram).

Take the point A for instance (40 w% at 100 C), which is a completely liquid, non-saturated solution of Al sulphate in water. Now cool this down to about 70 C and it hits the saturation line, which means solid aluminium sulphate hydrate now starts forming, in equilibrium with saturated solution at 40 w% and 70 C.

Further cooling forces more aluminium sulphate hydrate out of the solution (unfortunately my diagram cannot tell which hydrates form) and the liquid phase composition now ‘slides down’ along the saturation line, until at 0 C it is about 25 w%. Because aluminium sulphate forms such high hydrates, the solid takes with it much of the solvent. That is why the solution solidifies but some solution remains nonetheless.

[Edited on 1-6-2014 by blogfast25]

blogfast25 - 1-6-2014 at 05:10

I am sure that I read that this compound is hygroscopic.
I don't have EDTA and have no idea what else could be used.

[Edited on 1-6-2014 by CHRIS25]

Actually, I don't think it is. Certainly my commercial grade isn't hygroscopic at all.

Yes, for Al, EDTA titration is great.

[Edited on 1-6-2014 by blogfast25]

CHRIS25 - 1-6-2014 at 05:18

I used the Al sulphate solubility data from Wiki’s table and converted them to weight % aluminium sulphate and plotted them against temperature:

Sorry, but normally I understand weight % of a chemical, but I am lost at what you mean here by converting the al sulphate from wiki into weight percent? and then on your graph 40%, what is this percent?

blogfast25 - 1-6-2014 at 05:31

[Sorry, but normally I understand weight % of a chemical, but I am lost at what you mean here by converting the al sulphate from wiki into weight percent? and then on your graph 40%, what is this percent?

For instance, Wiki says at 70 C, an aluminium sulphate solution in water can contain a MAXIMUM of 66.2 g Al2(SO4)3 per 100 g of water. Convert that to w% Al2(SO4)3:

w% Al2(SO4)3 = 66.2 g Al2(SO4)3 / (66.2 g Al2(SO4)3 + 100 g water) x 100 % = 39.8 w% Al2(SO4)3 (and 60.2 w% water).

Clear?

CHRIS25 - 1-6-2014 at 06:08

Stupid me yes absolutely. Clear thankyou.

Aluminium chlorosulphate crystal - next side line project.

AJKOER - 1-6-2014 at 06:12

With respect to Al(OH)3(H2O)3, also called Aluminium triaquatrihydroxy complex, I repeat the interesting comments on this compound at http://www.chemthes.com/entity_datapage.php?id=4011 to quote:

"Real, long lived, electronically neutral reagent chemical.
Gram formula weight (molecular mass) = 132.04
Solubility in organics = insoluble
Solubility dependent upon pH. The trisaquatrishydroxyaluminium neutral complex is white, (even colourless), gelatinous material. Solubility is dependent upon pH."

The central, to be determined, question relates not to absolute solubility at any given pH, but relative solubility of this complex versus Mg(OH)2 at, say, pH 5.5 and the ability of the proposed reaction below to move to the right:

2 Al(OH)3(H2O)3 (aq) + 3 MgSO4 (aq) --?--) Al2(SO4)3 (aq) + 3 Mg(OH)2 (s) + 6 H2O

No reason to argue any more, I have good indications of a complete success!

First, I developed a fast route to Al(OH)3(H2O)3 by using my previous work on a galvanic cell of NaClO in the presence of Al and Cu with an electrolyte of NaCl (as one my references I have given in previous threads, please see http://www.exo.net/~pauld/saltwater/ ). After a few hours, just pour out some of the galvanic cell content leaving any undissolved Al, Cu or other solid residues behind.

What is actual created is a basic solution containing unreacted NaClO, NaCl and partially undissolved hydrated alumina, so the true magnitude of alumina in solution is not revealed until one adds a small amount of a weak acid (I used seltzer water which has a pH between 3 and 4). Immediately a thick suspension of hydrated alumina is revealed. Results of the creation of this compound can actually be obtained in minutes!

[Edit] Next, add dilute H2O2 to convert any unreacted NaClO to NaCl releasing O2 in the sample solution of hydrated alumina. On some runs, I did not perform this step, which can alter the final products.

Then, I recommend creating two working solutions of the hydrated alumina for comparison purposes. In the first solution, add MgSO4 dissolved in seltzer water. The second solution, add no MgSO4, just seltzer water. On standing in both test solutions, a similar precipitate is formed (I think the 1st solution precipitate is a little more fluffy).

Now to test the formation of Mg(OH)2, add vinegar as Alumina will not dissolve in weak vinegar, but Mg(OH)2 will. Results: the control solution remains cloudy, but the MgSO4 solution becomes completely clear as water. To test further, add some solid MgSO4 to the control solution. It eventualy becomes clear as well!

[Edit] First, I should mention that I used an excess of NaCl, which in the role of an electrolyte is essentially not significant, but for immediate side reactions and subsequent reactions involving H2CO3, amphoteric Al2O3 and MgSO4, this actually may have a significant impact on the 'activity coefficient' (which can significantly increase the functional acid strength of even weak acids as I have previously discussed on SM).

Also, caution, as while the presence of heavy metals in solution is much lower than direct routes employing especially impure (especially Iron) H2SO4, the longer the galvanic cell runs, potentially more toxic Copper is placed into solution. Allowing the final aqueous solution to sit in a piece of Aluminum (previously boiled and allowed to stand for a couple hours in vinegar) could mitigate this contaminant.

[Edited on 1-6-2014 by AJKOER]

[Edited on 1-6-2014 by AJKOER]

blogfast25 - 1-6-2014 at 06:31

JOKER:

You've proved nothing. You're just cackling on. Where's your proof of aluminium sulphate? There is none.

Chris:

I can even go further with this diagram, by making some simple (but probably inaccurate to some degree) assumptions. I’ll assume the main hydrate forming is n = 15 in my conditions. That has a molecular weight of 342 + 15 x 18 = 612 g/mol and a w% aluminium sulphate (AS) of 342/612 x 100 = 55.9 w%.

I’ll also assume I’m starting from a hot, 50 w% AS solution and allow it to cool to 0 C.

Say I start from 100 g hot solution and that after cooling it contains H g of AS hydrate and (100 – H) g of saturated solution. The diagram tells me that solution contains 26 w% of AS.

My original solution contained (100 x 50) / 100 = 50 g AS. Since as I didn’t add or remove any AS during the cooling I must find the same amount of AS in the cooled stuff:

Mass balance: 50 = (100 – H) x 26/100 + H x (55.9/100) or:
H = (50 – 26) / (0.559 – 0.26) = 80.3 g

So the 100 g of cooled solution would contain 80.3 g of solid hydrate and 19.7 g of saturated solution.

CHRIS25 - 1-6-2014 at 06:54

Sorry Blogfast - it is hard for me to follow this because I just about 'cope' with working on stoichiometry and the maths involved in so many things, that to start trying to theorize with mathematical ideas stretches my capacity to nervous breakdown. I understand your chart, the principles and so on, and functional maths, but I simply can not think this way, (ex trucker and visual artist/photographer = wrong side of the brain).

aga - 1-6-2014 at 07:00

What the hell.

AJKOER : post a Full experiment, including volumes, masses, concentrations etc and i will give it a go and post the results.

Why not ?

Start a new thread for it though.

CHRIS25 no longer has any Alumnium Sulphate Woes.

[Edited on 1-6-2014 by aga]

[Edited on 1-6-2014 by aga]

blogfast25 - 1-6-2014 at 07:13

What the hell.

AJKOER : post a Full experiment, including volumes, masses, concentrations etc and i will give it a go and post the results.

Why not ?

Most of all, I'd like to see a tangible product, with proof that it is (or is mainly) aluminium sulphate. Any hydrate will do.

Joker's latest belief system is crutched up by one paragraph on Al(OH)3.3H2O:

"Real, long lived, electronically neutral reagent chemical.
Gram formula weight (molecular mass) = 132.04
Solubility in organics = insoluble
Solubility dependent upon pH. The trisaquatrishydroxyaluminium neutral complex is white, (even colourless), gelatinous material. Solubility is dependent upon pH."

He forgets there's nothing novel about that: Al(OH)3 dissolves in sufficiently acidic conditions to hexaaqua aluminium cations, in sufficiently alkaline conditions to aluminate, Al(OH)4-, In neutral condition it is totally insoluble. Nothing new, nothing to see here, the wheels on the bus are turning.

[Edited on 1-6-2014 by blogfast25]

[Edited on 1-6-2014 by blogfast25]

Various Hydroxides

CHRIS25 - 1-6-2014 at 10:31

blogfast25 - 1-6-2014 at 11:41

Note that nearly all of these species are solvated, so for instance:

[Al(OH)2(H2O)4]+

[Al(OH)(H2O)5]2+

etc.

Note also extremely low concentration: 5 x 10-6 mol/L

[Edited on 1-6-2014 by blogfast25]

CHRIS25 - 1-6-2014 at 12:29

Note that nearly all of these species are solvated, so for instance:

[Al(OH)2(H2O)4]+

[Al(OH)(H2O)5]2+

etc.

Note also extremely low concentration: 5 x 10-6 mol/L

[Edited on 1-6-2014 by blogfast25]

Yes I know and that is the problem with Aluminium along with Iron, but aluminium is twice as bad due to its higher positive charge and the large size of its atom (though I confess I do not really understand the implications of the latter); anyway what is curious, is that I have read that Al + OH in soln becomes Al(OH)₄^- in alkali conditions and Al(H₂O)₆⁺³ in acidic. Question - Al(OH)₃ is formed here in slightly acidic conditions plus the concentration is low, does this mean that the concentration of Al ions must be kept low by addition of plenty of free OH ions, since I am now dissolving more Al in sulphuric to make only the Hydroxide so that I don't have to use Al metal in other experiments + it is more efficient, I am curious.

[Edited on 1-6-2014 by CHRIS25]

blogfast25 - 2-6-2014 at 04:36

The tendency for hydrolysis of Al3+ and Fe3+ is actually fairly comparable. The Al3+ ion is smaller and has the same charge as Fe3+. For the former, being smaller the electrical field surrounding it is greater and it has a greater potential to expel protons from these 6 water ligands.

I'm not sure I understand the question. Can you rephrase?

Aluminate ions can form from hydroxide by so-called "ligand substitution":

Al(OH)3(H2O)3 + OH- === > [Al(OH)4(H2O)2]- + H2O

One water molecule ligand substituted by a hydroxide ion ligand.

CHRIS25 - 2-6-2014 at 04:55

I'm not sure I understand the question. Can you rephrase?

I want to make Al hydroxide, I will precipitate it with 35% ammonia solution. Ph has to be carefully monitored and I have to arrive at 6 - 6.5. However you mentioned the percent of Al ions in solution that were given on this graph was extremely low. Is this really worth being too concerned about? (I have to confess in not being able to understand yet the maths nomenclature of what this concentration actually means).

[Edited on 2-6-2014 by CHRIS25]

blogfast25 - 2-6-2014 at 08:22

5 x 10-6 = 5 times 10 to the power minus six = 0.000005. It's a convenient format to describe small numbers with.

With ammonia you don't need to worry about pH and aluminate formation: ammonia is far too weak an alkali to form aluminates with.

Just respect the following stoichiometry (simplified notation):

Al3+ + 3 NH3 + 3 H2O === > Al(OH)3 + 3 NH4+

If you start from Al sulphate, you'll need 6 mol of ammonia per mol of Al sulphate.

After the addition of the ammonia, check for pH, it should be pH > 7. If it isn't yet add more ammonia in small aliquots until pH > 7.

After filtering you can evaporate the filtrate to recover ammonium sulphate, if you want to.

DraconicAcid - 2-6-2014 at 10:32