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blogfast25
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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]
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CHRIS25
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Quote: Originally posted by blogfast25  | | Quote: Originally posted by CHRIS25 | | 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. |
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]
‘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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aga
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| Quote: Originally posted by blogfast25 | 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!
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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.
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blogfast25
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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.
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DraconicAcid
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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.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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MrHomeScientist
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| Quote: Originally posted by AJKOER | | 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]
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blogfast25
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| Quote: Originally posted by aga | [
Would it not make more sense to drive off all the water and then rehydrate to get the actual hydrate you want ?
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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.
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blogfast25
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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]
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blogfast25
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| 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]
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CHRIS25
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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]
‘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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aga
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| 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
| Quote: Originally posted by blogfast25 | | 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]
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AJKOER
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| Quote: Originally posted by blogfast25 | .......
The Ksp of Al(OH)3 = 3 x 10<sup>-34</sup>. 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<sup>-12</sup>.
For your idea to work you'd have to find a form of Al(OH)3 that's roughly (AT LEAST) 10<sup>-12</sup>/10<sup>-34</sup> =
10<sup>22</sup> more soluble than what the tables tell you. A 1 followed by 21 zeros: some Al(OH)3 is that!
............
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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]
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CHRIS25
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@ 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.
‘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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AJKOER
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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.
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CHRIS25
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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
‘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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blogfast25
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So you don't even understand what neutral actually means?
Do you know what the concentration of H<sub>3</sub>O<sup>+</sup> is at pH 5.5? Here, let me help you:
10<sup>-5.5</sup> = 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.
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AJKOER
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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
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blogfast25
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You've already got your answer: whatever type of hydrated alumina you want to use it will have to be about 10<sup>22</sup> more soluble
than listed values. Good luck with that.
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aga
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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]
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blogfast25
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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 H<sub>2</sub> 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]
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blogfast25
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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.
"Al<sub>2</sub>" 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]
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CHRIS25
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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]
‘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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aga
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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]
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aga
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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.
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blogfast25
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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) === > Al<sup>3+</sup>(aq) + 3 e<sup>-</sup> (2)
And:
3 H3O+ (aq) + 3 e<sup>-</sup> === > 1.5 H2(g) + 3 H2O(l) (3)
(2) + (3) = (1)
Clearer now?
[Edited on 30-5-2014 by blogfast25]
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