Sciencemadness Discussion Board - Aluminium ChloroSulphate Crystal

this can never be a pure sample of AlClSO4.6H2O. The latter compound almost certainly must be colorless (or white).

Could you please describe the expected product in terms of possible colours, and textures ?

It would be a great help.

DraconicAcid - 2-6-2014 at 12:26

this can never be a pure sample of AlClSO4.6H2O. The latter compound almost certainly must be colorless (or white).

Could you please describe the expected product in terms of possible colours, and textures ?

It would be a great help.

Halide ions are colourless. Sulphate ions are colourless. Aluminum ions are colourless. Any combination of the above will be either white or colourless, depending on the size of the crystal.

Colours result from the the absorption of visible light. This will only happen in systems that have a) conjugated pi bonds, b) partially-filled d orbitals, or c) the possibility of internal redox reactions (charge transfer). Aluminum doesn't have any of these, and won't form a coloured compound unless its counterion has them.

aga - 2-6-2014 at 12:36

Wow !
Thanks very much for the comprehensive explanation of the ion colours, and also of the Why.
This goes in the Book for sure.

As i will try this experiment tomorrow, could you please give a clue as to the expected texture of the dried crystals, if they do in fact dry at all.

CHRIS25 - 2-6-2014 at 12:50

What purity of chemicals did you use? The solid material is yellow and this can never be a pure sample of AlClSO4.6H2O. The latter compound almost certainly must be colorless (or white).

HCl is Lab quality 37
Aluminium Sulphate I made myself from Lab Quality Sulphuric acid and Aluminium Piping, but the latter was filtered and a clear solution was achieved.

The whole solution started to turn yellow within seconds of heating - what about the colour of a chloride ion, tetra chlorocopper is yellow, and chlorine gas is yellowish/green?

Ishould perhaps also add that I had the boiling tube suspended between a stainless steel wire rack over the heat, because of wind outside I had it slightly covered as well. after 15minutes or so the stainless steel was covered in yellow, it washed off very easily but has stripped the steel of its bright lustre, no problem about this rack, but I thought I would mention it because of the colour of the residue deposited.

[Edited on 2-6-2014 by CHRIS25]

DraconicAcid - 2-6-2014 at 12:59

The whole solution started to turn yellow within seconds of heating - what about the colour of a chloride ion, tetra chlorocopper is yellow, and chlorine gas is yellowish/green?

Chloride ion is colourless (see sodium chloride for an example). Chlorine gas is a different substance, and transition metal complexes with chloride ions have partially-filled d orbitals to explain their colour. I would suspect that the yellow colour is a copper or iron impurity.

The Volatile Chemist - 2-6-2014 at 13:02

What purity of chemicals did you use? The solid material is yellow and this can never be a pure sample of AlClSO4.6H2O. The latter compound almost certainly must be colorless (or white).

HCl is Lab quality 37
Aluminium Sulphate I made myself from Lab Quality Sulphuric acid and Aluminium Piping, but the latter was filtered and a clear solution was achieved.

The whole solution started to turn yellow within seconds of heating - what about the colour of a chloride ion, tetra chlorocopper is yellow, and chlorine gas is yellowish/green?

I should perhaps also add that I had the boiling tube suspended between a stainless steel wire rack over the heat, because of wind outside I had it slightly covered as well. after 15minutes or so the stainless steel was covered in yellow, it washed off very easily but has stripped the steel of its bright lustre, no problem about this rack, but I thought I would mention it because of the colour of the residue deposited.

[Edited on 2-6-2014 by CHRIS25]

I don't know about chlorine gas, but the chlorine-copper complexes get their color from copper. The steel may somehow have influenced the coloration, eg. stirring sodium tetrachlorocuprate(II) with a steel screwdriver seemed to etch the upper layer.

blogfast25 - 2-6-2014 at 13:22

Chris:

This could be as much a success as a failure, right now we don't know that. It's rather encouraging that a solid material did form!

The yellow colour is almost 100 % certain Fe3+. Possibly from the HCl, in which ferric chloride is a common contaminant. You said it was 'lab quality': where did you get it? Test for Fe3+ with thiocyanate.

What's the consistency of this material? Hard/soft/mushy?

Does it dissolve in water and is the solution clear?

The real challenge is now to figure out the composition of the material, i.o.w. does its composition correspond to AlClSO4.6H2O?

If this compound did form then stoichiometrically what happened was:

Al2(SO4

3 + 2 HCl + 6 H2O === > 2 AlClSO4.6H2O + H2SO4

[Edited on 2-6-2014 by blogfast25]

CHRIS25 - 2-6-2014 at 13:24

The only source of iron contamination can be the HCl. My batch contains 0.5 ppm. Whereas a batch from the same company for trace analysis has 0.05 ppm. And yet another batch has 0.1 ppm as we go up in expense.

Blogfast: It dissolves very easily when broken up leaving the water solution clear
I do not have thiocyanate unfortunately.
Consistency is between Hard and soft, but not mushy. Tomorrow I will repeat the experiment but with more heating and try to leave less solution in tube. But will use 2g of sulphate again.
[Edited on 2-6-2014 by CHRIS25]

[Edited on 2-6-2014 by CHRIS25]

[Edited on 2-6-2014 by CHRIS25]

Zyklon-A - 2-6-2014 at 13:29

So none of your solution ever touched the SS? If so it must be an impurity in your reagents.
Chlorine gas is very different, it has covalent bonds, chloride anions are always colorless.
How about the purity of your aluminum sulfate? Was the Al pipe tested? Or you sulfuric acid, Is it also lab grade, or drain cleaner/battery acid.
My guess is iron chloride impurities, very small quantities can discolor a solution.
I don't know how the aluminum sulfate could be colorless, but turn yellow in pure hydrochloric acid, doesn't make sense.
You sure the HCl (aq) is pure?

blogfast25 - 2-6-2014 at 13:29

Are you sure that there wasn't any iron 'pick up', for instance from a steel spatula or such like?

CHRIS25 - 2-6-2014 at 13:34

Ok absolute clear: Not a single bit of the apparatus had any contact at all with anything steel iron, nothing, except glass and I use plastic to transfer solid chemicals, never metal things. The HCl analysis is 0.5ppm iron, the sulphuric acid is 1ppm iron. Both acids are from a reputable lab supplier. Yes my Al sulphate sample was and is pure white. The yellow began almost immediately upon heating.

The Volatile Chemist - 2-6-2014 at 13:38

If you have Tannic acid, you can test for iron complexes with that, or with Potassium Ferro/Ferri-cyanide. The FerroCN makes a blue compound, and the Tannic acid makes a black prec./complex.

Zyklon-A - 2-6-2014 at 13:38

Well then... I'm stumped.
BTW, the yellow coating on your SS wire thing is likly just from acid aerosols or condensed acid on said wire.

blogfast25 - 2-6-2014 at 13:39

I will try and repeat this experiment 'ASAP' [cough!]. It's interesting.

Such an aluminium chlorosuphate could crystallise out from a witches brew of aluminium sulphate and HCl, provided it is less soluble than aluminium sulphate and less soluble than aluminium chloride. It's entirely possible but requires rigorous proof.

[Edited on 2-6-2014 by blogfast25]

CHRIS25 - 2-6-2014 at 13:55

Zyclonb: Yellow coating on stainless steel was condensed acid I know. But would you say that the analysis of both my acids given above provide sufficient reason for the Iron colour?

Blogfast tomorrow I will do the following: Heat up the concentrated 37.5% acid for a minute to see if that remains clear. Then I will add the sulphate, probably more this time, I have to say that upon my first attempt not recorded here, I dissolved the sulphate in cold acid and it remained clear, again, not until heating did the yellow appear. I will see how far I can push the evaporation of Hydrogen chloride before allowing to cool. I think I could get a more hardened solid.

I have tannic acid and ferricyanide, I will get to work.

[Edited on 2-6-2014 by CHRIS25]

I made a quick tannic acid solution, the colour of very weak tea of course. Placed in a sizeable amount of the yellow solid, absolutely no change in colour at all, and no precipitate.

[Edited on 2-6-2014 by CHRIS25]

Zyklon-A - 2-6-2014 at 14:01

Quote:

Zyklonb: Yellow coating on stainless steel was condensed acid I know. But would you say that the analysis of both my acids given above provide sufficient reason for the Iron colour?

Nope. I doubt the impurities could have come from your acid.
Along with testing your acid by heating, test the sulfate by heating, perhaps that's where the impurity come from.

CHRIS25 - 2-6-2014 at 14:08

test the sulfate by heating, perhaps that's where the impurity come from.

The sulphate was created yesterday, and it was at near boiling point before being cooled. As I remember, it turned from clear at dilution stage to a viscous slight off-yellow when concentrated down prior to cooling and getting the alumnium sulphate hydrate. But this was characteristic of the colour you get from sulphuric acid upon heating, and there was excess sulphuric acid in the aluminium sulphate, so maybe something here? But Blogfast did the same thing, I wonder whether the colour of his solution prior to cooling was a slight off yellow, or was it totally clear when slightly viscous?

[Edited on 2-6-2014 by CHRIS25]

The Volatile Chemist - 2-6-2014 at 18:11

CHRIS!!!! I'll P2P you too, but I forgot, DO NOT add Ferricyanide to your HCl, just in case you didn't know. A neutralized AlClSO4 solution will be easy to test with ferricyanide, but Hydrogen Cyanide gas evolves if you mix ferricyanide and a strong acid. BE CAREFUL. If the product itself is acidic in solution, even that could make HCN vapors on mixture w/ HCl. The tannic acid will be fine, though.
Sorry for not saying so earlier, and I apologize if you already knew.

Brain&Force - 2-6-2014 at 19:58

I'm pretty sure the impurity is iron(III) ions. As recommended earlier, I would use thiocyanate to confirm - add some to an HCl sample as well to determine the source of the impurity.

What is the source of the aluminum sulfate? I would also dissolve some aluminum sulfate (if you have any left over) in hydrogen peroxide and allow that to dissolve, then add thiocyanate. If it's the metal ion, it may be hiding as iron(II) which doesn't show up for the thiocyanate test.

aga - 3-6-2014 at 00:03

Here is a repeat CHRIS25's experiment.

Some Aluminium Sulphate (hydrate unknown) was put in a beaker.
Heated until anhydrous, giving 14.41g

Then the equation was to be balanced, in order to determine how much HCl was required.
This failed to work algebraically.

Not knowing the mechanism or quantities, 23g of 25w% HCl and a splash of DIW were added, mixed and heated.

3 mins rapid Boiling, 30mins gentle heat.
Very quickly the liquid became a pale yellow.
By 20 minutes into the gentle heating, the liquid is clearer, and yellower.

Waiting for the 30 mins to elapse, i attempted to work out what was wrong with the equation.

Al₂(SO₄)₃ + HCl -> ALClSO₄ cannot be balanced.

So something else must be going on.

Using CHRIS25's clue regarding ' ...characteristic Sulphuric acid ...', this makes more sense :-

Al₂(SO₄)₃ + 2HCl -> 2ALClSO4 + H₂SO₄

The sulphuric being in equilibrium with the water.

After 30 minutes the liquid was cooled in a water bath.
A this stage the liquid is yellow, and giving off strong fumes of clorine, which are visible.

After 3 mins cooling, the liquid was transfered to a smaller plastic container and further cooled in a water bath.
At 10 minutes the solution has solidified into a hard yellow mass weghing 18.86g

In colour it is more like elemental sulphur than anything else.

So, it is either contaminated by something with yellow ions, or some other reaction is happening, or the quest for The EarWax Recipie is over.

By the weights, that would make it a .16 hydrate. Does that even exist ?

Source of reactants :-
HCl was from a reputable chem supplier.
Al was from aluminium foil.
H₂SO₄ was from OTC drain Openener, colour removed with H₂O₂, boiled
Water is OTC distilled & de-ionised
All implements are glass/plastic in each process.

By the time i finished typing this, the yellow mass has developed white patches, and there is a strong sulphur smell coming from it.

[Edited on 3-6-2014 by aga]

Second Sample

CHRIS25 - 3-6-2014 at 02:18

Ingredients:
4 mLs 37.5% HCl
4 g Aluminium Sulphate (14 hydrate)

Method:
HCl was boiled first, took 10 seconds, but boiled for 1 minute
Sulphate was added to the hot HCl
Boiled continuously for 5 minutes
Cooled in Ice bath for 5 minutes

Observations:
The precipitate is Less yellow and harder than yesterday

Reasonings:
I doubt that there is iron contamination by the distinct lessening of yellow in the sample (Twice as much sulphate). Also by doubling the sulphate and lessening the HCl and boiling continuously I wanted to drive off much more water.

Next step:
Can't think of one just yet.

More Observations:

Point 1. Although not very clear here there are, unlike first sample, far more white areas visible, especially on the underside where the bowl of the boiling tube was.

Point 2. Yesterdays sample has, for all intense and purposes, melted in its own crystalization, Mmm, the colour has gone from this bright yellow to the same colour as the sample I have just made, Loss of yellow in other words.
[Edited on 3-6-2014 by CHRIS25]

[Edited on 3-6-2014 by CHRIS25]

blogfast25 - 3-6-2014 at 04:07

@Volatile chemist:

Calm down. Only in highly acidic conditions with heating do ferro/ferricyanides release any HCN. These complexes are very stable. To do the test safely, just dilute the sample a bit. But ferricyanide won't detect ferric ions, only ferrocyanide does that. Similarly ferricyanide detects ferrous ions. In both cases Prussian Blue is formed.

@Aga:

The stoichiometry is:

Al2(SO4)3 + 2 HCl + 12 H2O === > 2 AlClSO4.6H2O + H2SO4

Without enough HCl in the solution this product will not form.

I doubt very much you observed chlorine: there's nothing in that mixture that can oxidise chloride to chlorine. I think you saw and smelled HCl. Chlorine cannot arise spontaneously: something has to oxidise the chlorine ions. That oxidation requires powerful oxidisers, BTW...

@Chris:

To avoid the samples redissolving try this. Immediately after preparation put them in the fridge. When cooled sufficiently try washing with small amounts of ice cold water, then pat dry with kitchen towel. It's mainly excess solution that clings to the crystals that causes the latter to re-enter solution, I believe...

To further characterise the material you need to obtain it in a reasonably dry state. Like dried in a CaCl2 desiccator.

So far, the materials obtained look quite different from aluminium sulphate hydrate. That counts for something. My guess remains this is indeed aluminium chlorosulphate but now we need to prove that!

[Edited on 3-6-2014 by blogfast25]

CHRIS25 - 3-6-2014 at 04:42

After cutting, washing in cold deionized water and drying with filter paper they are even less yellow now. Although they are far from crystals, I can break them up between my fingers, and they definitely smell of HCl. They still dissolve into a small amount of water and leave a very clear solution, except now, I can see how impure my tap water is, this stuff is more efficient than aluminium sulphate as a flocculant, you should see the precipitates that this forces out of common tap water.

I will store in a dessicator with calcium chloride.

blogfast25 - 3-6-2014 at 04:50

Well done. The loss of colour is entirely expected: the ferric chloride is present only in SOLUTION, between the crystals. It doesn't crystallise out because ferric chloride is highly soluble. By washing with pure water you've removed part of that solution, thus making the product whiter.

They smell of HCl because of the residual solution clinging to them. And trust me: they are crystalline alright, just not very pretty!

A very small amount of ferric chloride may have ended up trapped in the actual crystal lattice of your product but that's not really a problem.

Oh, and to effectively dry the stuff in a CaCl2 desiccator, crush the lumps to more or less the consistency of white granulated sugar: it will make it easier for the remaining water to evaporate and become absorbed by the CaCl2.

[Edited on 3-6-2014 by blogfast25]

CHRIS25 - 3-6-2014 at 05:10

Ok then. And there I was trying to draw lewis dot diagrams of this to try and see something....(can't manage the lewis dot structure with this one, others I have done but this one is difficult for me. Anyway, so ferric chloride makes sense. No I was expecting a spectacular array of sunshine triangulating its way through crystal Planes of beauty. I still think they are too soft. Will try one more time later on and adjust ratios.

I have been trying for ages to work on this, Normally there are 2 Al ions needed to bond with 3 sulphates. Not having the second Al ion makes it difficult especially since I can not see how on earth to solve this now with the chloride ion. It's now really annoying me since I solved so many double and triple bonds and even a resonance question, I thought I could do this AGHRrrr...

Second question, why has the Al ion in soln a positive 3 charge when it misses 5 electrons from its octet requirement?

The more you learn the more you realize you don't know....and for every question answered three more questions rear their irritating little heads.....

[Edited on 3-6-2014 by CHRIS25]

[Edited on 3-6-2014 by CHRIS25]

blogfast25 - 3-6-2014 at 08:01

Chris:

You're making it more complicated than it is. No Lewis dots needed here. AlClSO4 is slightly unusual (with respect to Al sulphate) in the sense that you have in this type of lattice 1 type of cation (Al) and two types of anion (1 sulphate, 1 chloride). It's uncommon but possible.

Al's electron configuration is [Ne] 3s2 3p1. When it 'loses' these three valence electrons (3s2 3p1

its electronic structure becomes that of Neon. Simply put: Al3+ = [Ne]. And the electron configuration of neon is a full octet.

The only other way Al could acquire an octet structure would be by absorbing 5 electrons, so it would become Al5- (!!!) but that is thermodynamically totally unfavourable because of the mutual electrostatic repulsion of those electrons.

Re. softness, allow to dry, then reassess.

The more you learn the more you realise you know hardly ANYTHING, true for everyone (but few know it).

[Edited on 3-6-2014 by blogfast25]

aga - 3-6-2014 at 08:35

Quote:

I think you saw and smelled HCl

Yes. Very likely.
Bad choice of words on my part due to limited experience.

Nobody is born knowing anything at all.

The point where you think you know everything is precisely the point at which you stop learning.

@CHRIS25
Nice one.
I shall chill some water and see if my earwax also goes Ultrabrite.

[Edited on 3-6-2014 by aga]

CHRIS25 - 3-6-2014 at 08:59

Aga: I don't have the apparatus to do this, but I was reading a patent yesterday and noticed one paragraph. It gave me an idea, maybe you could follow the procedure through to the point where you cool, but don't allow it to solidify, then pump HCl gas into the solution for a few minutes. then allow to cool and solidify? I have no idea but maybe worth trying?

blogfast25 - 3-6-2014 at 09:17

[...] then pump HCl gas into the solution for a few minutes. then allow to cool and solidify? I have no idea but maybe worth trying?

'Gassing' (as we call it) a concentrated solution of Al sulphate with HCl gas is an option I've considered already. But it means making a HCl generator (which really isn't too difficult), it just complicates things a bit. More and a better product may result from it though. Definitely worth trying... The advantage of gassing is that you don't add more water, this should at least give greater yield of product.

@aga:

You suffered from confirmation bias, I think. The smell of HCl, the yellow colour of the solution, possibly visible fumes of HCl... all these fooled your associative brain into 'seeing' chlorine. It's a very common mistake.

[Edited on 3-6-2014 by blogfast25]

aga - 3-6-2014 at 12:18

@blogfast25
Being such a noob, i have never seen actual chlorine gas, so should not have sait it was chlorine.

All i really saw were wisps (3-4 inches) of steam-like stuff twisting out of the beaker, and smelt something clorine-y.

@CHRIS25
'gassing' seems like a Good plan - no added water, no boiling Off excess water.
As the reaction is with HCl, perhaps just Al Sulph (gawd bless 'im) and HCl gas will work much easier.

It sounds like it will be easier to produce a pure Gas than a pure solid or liquid.
Someone jump in and correct me if this is plain stupid.

Busy day tomorrow, so unlikely to get anywhere with it until Thursday.

Now to find out how HCl gas is produced, and how to survive doing so.

blogfast25 - 3-6-2014 at 12:24

I dissolved 20 g of aluminium sulphate hydrate (ASH) in 20 ml of HCl 37 w% by simmering for a bit. I used commercial grade ASH (n = 14.3, MM = 600 g/mol, iron free).

On cooling the whole mass slowly crystallised, so that was no good.

I then added another 20 ml HCl and reheated to complete dissolution (the solid dissolved very quickly). On cooling I obtained a mass of snow white small crystals (0.5 - 1 mm), at a guess about 10 g. They look very different from ASH: much more crystalline.

The supernatant liquid was also yellow because of iron in my HCl.

Now to find out how HCl gas is produced, and how to survive doing so.

1 mol NaCl + 1 mol H2SO4:

NaCl(s) + H2SO4(l) === > HCl(g) + NaHSO4(s)

You don't even need to heat it but you do need 95 - 98 w% H2SO4.

This reaction works because hydrogen chloride is a gas and leaves the reaction. But water is the enemy here because HCl is so damn soluble in it, so water will reduce yield of gaseous hydrogen chloride.

There are various threads on HCl generators on this forum. Search and ye shall find.

[Edited on 3-6-2014 by blogfast25]

aga - 3-6-2014 at 12:41

Only found 3 so far ...

Quick thought : why does the 'gassing' need to be done in a beaker ?
It doesn't does it ?

The surface area will be an important aspect - more = better/faster.

Al Sulph Hydrate (OldAl) when heated goes liquid.

So, if i make a box that is gas-tight, pump far too much HCL into it (purge the Air) and then Spray the OldAl solution in an aerosol Up into that gas, then keep heating/pumping OldAl until it pumpeth no more, then the reaction should be complete.

I guess the product would be spattered all over the container, but it should be pretty pure.

Does that sound like madness ?

[Edited on 3-6-2014 by aga]

blogfast25 - 3-6-2014 at 13:21

Does that sound like madness ?

[Edited on 3-6-2014 by aga]

Erm... affirmative. Repeat: AFFIRMATIVE, roger.

Why go to that trouble if you can simply lead HCl gas into a beaker with saturated aluminium sulphate solution? The gas will dissolve in the solution and on cooling (presumed) aluminium chloride sulphate should crystallise out. Or so says the plan.

aga - 3-6-2014 at 13:26

OK. affirmative. We got a Green Light then.

Insane method versus Beaker & inverted glass funnel gas dispenser/anti-suck-back device.

Experimentation will tell us which works best.

No bets though.

Hang on.
Even if OldAl sticks to the sides Un-reacted, the presence of the gas will cause a reaction, so adding the liquid to the Gas in this way should work out better, as in faster, seeing as the available surface area for the reactants to encounter each other will be many times higher than a bubble of HCl.

Clearly the Gas concentration must remain high at all times, so just need to keep pumping it in.

Equally clearly, the hydrate required needs water.
Calculation required.

[Edited on 3-6-2014 by aga]

Zyklon-A - 3-6-2014 at 13:30

Be careful to avoid suck-back, the high solubility of HCl (g) will cause the solution to flow back into the gas generator.
A simple anti suck-back devise will contain any and all problems, should they occur/when they occur.

[EDIT], Ah, I see you already know about such things.

BTW, would chlorine work to make AlClSO₄, quite a bit cheaper.

[Edited on 3-6-2014 by Zyklonb]

aga - 3-6-2014 at 13:35

Ah, I see you already know about such things.

Such things yes, but not That thing !

What is it called, and where can i get one ?

BTW, would chlorine work to make AlClSO4, quite a bit cheaper.

I have no idea.
Will bubbling Chlorine thru OldAl make the same as bubbling HCl through it ?

[Edited on 3-6-2014 by aga]

Zyklon-A - 3-6-2014 at 13:47

After looking for more than 5 minutes I couldn't find one anywhere except Alibaba:http://heqi.en.alibaba.com/product/256344917-200694327/anti_suck_back_bubbler.htmlhttp://heqi.en.alibaba.com/product/256344917-200694327/anti_suck_bac k_bubbler.html :mad:

You have to buy only 100 pieces, you'll need that many right?

I don't have one, although it would be nice, a vacuum flask or similar will work just as well.

Will bubbling Chlorine thru OldAl make the same as bubbling HCl through it ?

[Edited on 3-6-2014 by aga]

Thats what I was wondering, I work with chlorine quite a bit, as long as you have a gas mask, it's a very pleasant substance to work with.

[Edited on 3-6-2014 by Zyklonb]

CHRIS25 - 3-6-2014 at 14:08

Blogfast: used 0.033 mol sulphate and an eventual 0.472 mol HCl. This is a 14:1 ratio; I used a 7:1 ratio. It seems more acid produces smaller crystals then. Tomorrow, if not raining still, I will try midway between these two just for fun, and then call it a day.

ZyKlonb: That piece of apparatus does not look suitable anyway, major problem that I see is no way to get any solid out that may form inside. It seems only suitable for gases and liquids.

aga - 3-6-2014 at 14:18

It's an experiment, so i guess the inverted funnel will be OK.

After all, this isn't Industrial Science Best Practice is it ?

Anyhoo, y'all jump in and do a simple experiment and show us noobs how it is done.

My bet is on my idea - squirt the OldAl soln into the gas.

aga - 3-6-2014 at 14:22

Quote:

and then call it a day.

Not a Chance.

Got to hammer it flat or it will forever be wobbly.

Quote:

Thats what I was wondering, I work with chlorine quite a bit

If you work with Chlorine, then you are in a much better position to answer that question than most.

Assuming that the chlorine Ion is involved in the reaction, and already knowing that water is present, so suspect that Cl dives into the water and so Cl ions happen.

Sequester a litre or 22.4 and try it, and tell us if it works.
Simples.

[Edited on 3-6-2014 by aga]

blogfast25 - 4-6-2014 at 01:49

@aga:

Good luck!

@Zb:

Boy, is that piece overkill or is it overkill?

All you need is an empty container with an unobstructed inlet (connected to the generator) and an unobstructed outlet (connected to the container you're gassing). In case of suckback the liquid is sucked into the container. Been there, done that. The piece you show would be suitable for gassing something but even for that purpose it's overkill. And what Chris said...

BTW, it's not the solubility of HCl in water that can cause suckback. And with the inverted funnel method there can be no suckback anyway. But I still prefer leading the gas directly into the solution to be gassed...

Chlorine won't work here: there's no way to reduce it to chloride without additional chemicals.

@Chris:

Careful drawing conclusions like that: how do you know it wasn't crystallisation temperature for instance that caused this more compact, microcrystalline product?

But by all means try my ratio...

@All:

FORGET chlorine. There's no way it can react here to chloride and chloride is what we need here.

[Edited on 4-6-2014 by blogfast25]

aga - 4-6-2014 at 03:00

I just had a quick go at the HCl gas idea.

First off i knocked up and tested the gas generator (used coffee jar+irrigation valves+lost of hot melt glue), and the reaction of NaCl + H₂SO₄ seemed quite slow, so i add quite a lot of salt and controlled the addition of the acid.

Interestingly the HCl gas feels 'hot' on skin, wheras the actual reaction solution doesn't feel even warm through the glass.

Then 5g of Alumnium Sulphate (unknown hydrate) as added to a vacuum flask with no additional water and the tube attached.
After running the generator for 30 seconds, a Pressure Sensor was fitted to contain the gas in the flask, and verify gas production.

Almost immediately the white Aluminium Sulphate crystals turned yellow.

Afer 30 minutes the crystals are showing white spots, mostly at the edges, and appear 'wetter' than before, in that they leave traces on the glass as the flask is swirled.

Zyklon-A - 4-6-2014 at 05:44

FORGET chlorine. There's no way it can react here to chloride and chloride is what we need here.

Ah right, I don't know how I missed that.
Al₂(SO₄)₃ (aq) + 2 HCl (g) → 2 AlClSO₄ + H₂SO₄, that works, but
Al₂(SO₄)₃ (aq) + Cl₂ (g) → 2 AlClSO₄ + SO₄^-, there's no cation to combine with the sulfate anion. [Face-palm]

Almost immediately the white Aluminium Sulphate crystals turned yellow.

After 30 minutes the crystals are showing white spots, mostly at the edges, and appear 'wetter' than before, in that they leave traces on the glass as the flask is swirled.

Weird. Why is it always turning yellow, some kind of side reaction? Doesn't make much sense...

[Edited on 4-6-2014 by Zyklonb]

MrHomeScientist - 4-6-2014 at 06:06

So, if i make a box that is gas-tight, pump far too much HCL into it (purge the Air) and then Spray the OldAl solution in an aerosol Up into that gas, then keep heating/pumping OldAl until it pumpeth no more, then the reaction should be complete.

I guess the product would be spattered all over the container, but it should be pretty pure.

Does that sound like madness ?

This sounds terrifying and overcomplicated, to say the least. The inverted funnel trick should work perfectly here, considering HCl's high solubility. It's also a lot less crazy, without pumps, boxes full of pressurized corrosive gas, and tedious scraping of contaminated product off the walls ala an exploded microwave burrito.

In case it isn't entirely clear, here's a picture I dug up from the internet on how the inverted funnel trick works.

If suckback starts to occur, the liquid level outside the funnel falls until it breaks the seal and equalizes the pressure. In the case of very soluble gases like HCl, you may not even need to immerse the funnel in the solution at all - just suspending it a few mm above (while stirring) may work.

Edit: The rest of that page is actually pretty interesting. It's the classic ammonia fountain demonstration, except with HCl instead:
http://www.nuffieldfoundation.org/practical-chemistry/proper...

There's also a nice graphic of an HCl(g) generator:

[Edited on 6-4-2014 by MrHomeScientist]

Yellow after all?

CHRIS25 - 4-6-2014 at 07:55

I am beginning to question the iron theory. there are hardly any images of this chemical to be found, here is the only one I could find:
http://www.feralco.com/ES/UK/page858-phal-10.php

Not perhaps best but this also:
http://www.lehighminerals.com/Z4auction339oct22.htm just scroll down to KLEINITE now I know there is mercury here but mercury yellow? No.

It does appear so far that just maybe our chlorosulphate should be yellow after all, and that yellow is provided by ???

I am not discounting the ferro idea, but where is Iron coming from in HCl gas, nowhere....

[Edited on 4-6-2014 by CHRIS25]

blogfast25 - 4-6-2014 at 08:30

I am beginning to question the iron theory. there are hardly any images of this chemical to be found, here is the only one I could find:
http://www.feralco.com/ES/UK/page858-phal-10.php

Not perhaps best but this also:
http://www.lehighminerals.com/Z4auction339oct22.htm just scroll down to KLEINITE now I know there is mercury here but mercury yellow? No.

It does appear so far that just maybe our chlorosulphate should be yellow after all, and that yellow is provided by ???

I am not discounting the ferro idea, but where is Iron coming from in HCl gas, nowhere....

[Edited on 4-6-2014 by CHRIS25]

I don't think the feralco product is the same as ours. The 'poly' I a bit of a giveaway. The feralco product is for water treatment: a bit of ferric chloride in there would do no harm. It's also a technical grade: they probably started from technical HCl solution.

It also doesn't explain why my crystals are snow white and only the supernatant liquid is yellow. I will test it later on with ammonium thiocyanate.

Also, Al cations, sulphate ions and chloride ions are all perfectly colourless: hard to see how something made from these 'building blocks' would be coloured...

What is very puzzling is that search terms like 'AlClSO4', 'aluminium chlorosulphate', 'aluminium chloride sulphate' yield almost nothing (except really for the low authority 'atomistry' and NEVER a CAS number or similar known chemical identifier (EINECS and such like).

CHRIS25 - 4-6-2014 at 08:43

be yellow after all, and that yellow is provided by ???
Also, Al cations, sulphate ions and chloride ions are all perfectly colourless: hard to see how something made from these 'building blocks' would be coloured...

What is very puzzling is that search terms like 'AlClSO4', 'aluminium chlorosulphate', 'aluminium chloride sulphate' yield almost nothing (except really for the low authority 'atomistry' and NEVER a CAS number or similar known chemical identifier (EINECS and such like).

Ok I do understand, but what about aluminium chloride most of the products are yellow?

You are right about the search, I have in between other things spent two days on and off trying to find anything, nothing at all, except a few patents.

Zyklon-A - 4-6-2014 at 09:07

Per Wikipedia:

Quote:

Aluminium chloride (AlCl3) is the main compound of aluminium and chlorine. It is white, but samples are often contaminated with iron trichloride, giving it a yellow colour.

So, it's yellow color is due to common impurities, has nothing to due with aluminum chloride, which is always white.
That's not to say that aluminum chlorosulfate isn't off-white, but I can't see how that's possible.

blogfast25 - 4-6-2014 at 09:57

Re, AlCl3, the AlCl3.6H2O I once synthesized was also completely white.

aga - 4-6-2014 at 10:31

Elemental sulphur.
Thats the kind of colour it has, and it smells sulphurous.
Also chlorine-y.
However, CHRIS25 washed his in cold water, and the colour was washed out, so it's probably a Yellow Ion.

Does anybody have an ion/colour list readily to hand in order to set up the Candidates ?

Here's a list of Stuff that can be Yellow that i have found so far.

Elements/compounds of:-

Chlorine
Sulphur
Sodium
Barium
Uranium

Ions :-

Iodide
Tetrachloro-copper complex
Cobalt-ammonium complex
Iron(III)
Chromate
Pervanadyl

Sources:-
http://en.wikipedia.org/wiki/Color_of_chemicals
http://www.sanjuan.edu/webpages/dkrenecki/files/ColorsofSubs...

If i were using Aluminium Sulphate derived from the CuSO4 route, i'd suspect the Tetrachloro-copper complex, however this Aluminium sulphate was made via the Sulphuric acid + Aluminium route.

The prime candidates for the Yellow colour are Chlorine and Sulphur.
Iron is less likely, in that it is hard to see where it came from.
If it were copper or Chromate ions, that could come from cheapo chrome plated copper spatulas, however they were not used in any part of the processes.

If it were Uranium, i would definitely remember adding it.

[Edited on 4-6-2014 by aga]

The Volatile Chemist - 4-6-2014 at 12:07

Elemental sulphur.
Thats the kind of colour it has, and it smells sulphurous.
Also chlorine-y.
However, CHRIS25 washed his in cold water, and the colour was washed out, so it's probably a Yellow Ion.

Does anybody have an ion/colour list readily to hand in order to set up the Candidates ?

Here's a list of Stuff that can be Yellow that i have found so far.

Elements/compounds of:-

Chlorine
Sulphur
Sodium
Barium
Uranium
[...]

Barium? yellow? Looking at your first source, it says a flame test gives green/yellow. If that's where you got the idea barium compounds could be yellow...

A flame test has little to do with actual coloration of compounds. If you do actually know of a yellow barium compound, tell me! I want to bring some color to one of my mostly white chemicals

blogfast25 - 4-6-2014 at 12:16

Volatile Chemist:

Barium chromate.

Jokes aside though, allow me to test for the most LIKELY culprit first (no, aga, it's NOT uranium

) before we we set off on a wild goose chase, i.e. ferric ions.

The Volatile Chemist - 4-6-2014 at 12:18

Volatile Chemist:

Barium chromate.

But isn't the chromium the cause of the coloration, not the barium???

blogfast25 - 4-6-2014 at 12:19

VC:

Of course but you asked for a yellow barium compound? Your will is our command!

The Volatile Chemist - 4-6-2014 at 12:26

VC:

Of course but you asked for a yellow barium compound? Your will is our command!

I hardly use the word, but LOL. I bring it up because barium ions just don't cause yellow color, which is what we're talking about. But chromium compounds? I love em

aga - 4-6-2014 at 13:03

Apparently some Barium ions/compounds/complexes Can give a Yellow/Green colour.

I included Everything i could find a reference to that *could* give a Yellow colour.
It is highly unlikely to be Barium, seeing as i have never used barium or a barium compound in these pieces of glass.

I'd put Barium contamination on a par with Uranium.

Sulphur or Chlorine.
My bet is on Sulphur.
Each way bet. Might be Chlorine.

[Edited on 4-6-2014 by aga]

The Volatile Chemist - 4-6-2014 at 13:12

Apparently some Barium ions/compounds/complexes Can give a Yellow/Green colour.

I included Everything i could find a reference to that *could* give a Yellow colour.
It is highly unlikely to be Barium, seeing as i have never used barium or a barium compound in these pieces of glass.

I'd put Barium contamination on a par with Uranium.

Sulphur or Chlorine.
My bet is on Sulphur.

Sure, I'll bet it's chlorine dioxide...

I wasn't saying it would be barium, though it's more likely than uranium (minimal still) because glass can be made of barium for spectroscopic reasons, but still ridiculously (to the point of being absurd) unlikely.

aga - 4-6-2014 at 13:17

Quote:

glass can be made of barium for spectroscopic reasons

Worth mentioning.
I didn't know that, and Glass has been used extensively.
Borosilicate glass actually.
Seems Silicon & Boron aren't Yellow though.

blogfast25 - 4-6-2014 at 13:20

My bet is on Sulphur.
Each way bet. Might be Chlorine.

How much are you willing to wager? I'm not a betting man but I do like a dead cert (Fe3+

aga - 4-6-2014 at 13:24

A Small Hadron Collider.

Spit. Deal Done !

I doubt that i have the reagents to test for Fe ions.
I'd best buy some to settle the matter.

[Edited on 4-6-2014 by aga]

blogfast25 - 4-6-2014 at 13:26

Not fair.

You're already sending me one of those.

If you want to chicken out, just say so!

aga - 4-6-2014 at 14:25

Oh dear.
You said the C word.
You have no idea.
None, due to the unseen Dimensions at work behind that word.

P'kaaark !

[Edited on 4-6-2014 by aga]

aga - 4-6-2014 at 15:32

I do like a dead cert (Fe3+

Ok.
How did Fe enter the chain ?

Candidates are :-
1. OTC Aluminium Foil
2. OTC Drain Opener (brown) Sulphuric acid
3. OTC 3% H2O2
4. 25 w% HCl from a reputable dealer
5. poor physical hygiene : dirty glassware, iron fingers etc.
6. poor mental hygiene (imagining it was all simple)

blogfast25 - 5-6-2014 at 05:10

In order of likelihood (my opinion): 4, 1, 2. Unlikely: 3, 5, 6

CHRIS25 - 5-6-2014 at 06:29

The following Not possible: 1,2,3,5,6, that leaves 4. I did not use foil I used aluminium metal, cleaned and washed, I am not aware that aluminium metal pipe can contain Iron? Mine came from years in the sea, close to 50year old dead boats. No signs of pitting or discolouration. Maybe trace amounts of iron in some manufacturing process? Possible? No idea.

[Edited on 5-6-2014 by CHRIS25]

blogfast25 - 5-6-2014 at 07:18

Chris:

Trace amounts Fe are possible in Al commercial metal, possibly more in some alloys. But your ASH was snow white, indicating very low levels of Fe.

CHRIS25 - 5-6-2014 at 07:42

Chris:

Trace amounts Fe are possible in Al commercial metal, possibly more in some alloys. But your ASH was snow white, indicating very low levels of Fe.

Ok then. Just found a whole lot of people saying No not possible, but then came across one comment from someone who works in an aluminium sheet factory who said that he has to test the final product for a trace analysis and there is always Iron to be found usually less than 5% by weight. Yes my Al sulphate is snowy white, no visible yellowing at all.

DraconicAcid - 5-6-2014 at 08:24

Chris:

Trace amounts Fe are possible in Al commercial metal, possibly more in some alloys. But your ASH was snow white, indicating very low levels of Fe.

The yellow colour of iron(III) is much more intense when it's got chloride ions as ligands. I'm not surprised that the sulphate isn't yellow, but only turns yellow when you add HCl.

blogfast25 - 5-6-2014 at 08:42

Left: supernatant liquid from that run, light yellow.

Right: same but with a bit of 1 M KSCN added.

The colour is due to the complex FeSCN2+ cation, which is deep wine red in high concentration.

In the mean time I’ve also broken up my crystalline mass of product and washed it with small aliquots of iced water. Normally I would suck them dry on my Buchner but it’s not available right now. So I patted them dry, put them on a piece of filter paper (and some kitchen towel underneath) and will allow them to dry in the fridge. Then some more drying in a silicagel desiccator.

A small amount of material was tested for solubility with RT water: it dissolved quickly to a clear solution.

All being well I will determine the Al content of this sample, as a first step to composition analysis.

[Edited on 5-6-2014 by blogfast25]

blogfast25 - 5-6-2014 at 08:47

The ASH I used has a reported Fe content of max 0.006 w%. In my case much of this iron must come from the HCl, which I know contains some.

CHRIS25 - 5-6-2014 at 09:17

Very interesting and conclusive Gert.

As a side not, I have had my broken up chlorosulphate in a dessicator bag with calcium chloride as dessicant, outside in the unusually bright sun for 8 hours. I have just opened the bag and got an obnoxious pong of HCl fumes. The formally yellow crystalline solid is for the best part now utterly white, though the larger bits still a bit yellow. They are solid But with one proviso, they can still be squeezed like icing sugar texture between your fingers, no change in weight though, so no evaporation of water occurred, and the calcium chloride smelled of HCl. So there was some absorption caused by the heat, but no change at all in the "wetness" of the product.

This I did not expect: After an hour in doors and in open air the chlorosulphate has gone completely and totally and utterly snow white, the whole lot, every grain, and still smells strongly of the Chloride, so it has not decomposed to aluminium sulphate.

[Edited on 5-6-2014 by CHRIS25]

blogfast25 - 5-6-2014 at 11:56

Chris:

It cannot be excluded at this point that the product is decomposing while one dries it. It would then emit HCl gas, almost certainly...

Once my product is sufficiently dry I will test for decomposition. If it decomposes that easily, we may never know its exact, initial composition.

It's possible (as a hypothesis) that as a solid it can only exist in very strongly acidic solutions.

Describe what you understand by 'wetness', please?

[Edited on 5-6-2014 by blogfast25]

[Edited on 5-6-2014 by blogfast25]

CHRIS25 - 5-6-2014 at 14:17

Chris:

Describe what you understand by 'wetness', please?

[Edited on 5-6-2014 by blogfast25]

[Edited on 5-6-2014 by blogfast25]

In this context speaking subjectively, when you squash a piece between your fingers it feels like pinching icing sugar except that there is less friction due to a slippery surface. Wow I have never had to describe wetness like this before.

aga - 5-6-2014 at 15:06

i just tried the Aluminium Suphate Hydrate crystals with Chlorine Gas.

The Cl₂ generator was OTC Sodium Hypochlorite and HCl.

Having now Seen, Smelt and Suffered from Chlorine gas exposure, i retract unreservedly my earlier comment about something smelling 'chloriney'.

It didn't.
That was HCl.
This is Cl₂

Remarkable difference. Truly remarkable.
I used to think i knew what chlorine smelt like.
I did not know at all.

Anyway, as before, the gas was led into a vacuum flask into which was placed some aluminium sulphate (unknown) hydrate crystals.

Most of the crystals remained totally white.

Some are taking on a very small portion of Yellow, with most of the mass remaining white.

So, HCl makes it go Ape, Cl₂ does not.

Now, if you will excuse me, i will go and seek out a Beer, an ENT specialist, a pneumothorax specialist, and posssibly an entire replacement body.

The Volatile Chemist - 5-6-2014 at 18:56

Excused, aga. This yellow color is indeed interesting!

blogfast25 - 6-6-2014 at 03:46

Chris:

It would be simpler to just call it: 'moisture content'

Aga:

Not sure what you wanted to achieve there. For Cl2 to do anything, it would need something to oxidise but there's nothing there. Everything in that flask is already at its maximum oxidation state, except oxygen [in the sulphate ions] which is at -2 but O (II-) cannot be oxidised by Cl2 in any way, shape or form. Mixing things for the sake of mixing things isn't very useful...

Trying to react chlorine with a solid at RT isn't easy anyway: most things would need some heat to get started.

Far more interesting remains to gas a saturated solution of ASH with HCl: there we can expect something to precipitate/crystallise but it's by no means certain.

As regards the 'remarkable' difference? What did you expect? They are two very different substances!

[Edited on 6-6-2014 by blogfast25]

CHRIS25 - 6-6-2014 at 05:22

Chris:

It would be simpler to just call it: 'moisture content'

[Edited on 6-6-2014 by blogfast25]

You said "Describe" I would have failed 'O' level English if I had simply said "moisture content"

Zyklon-A - 6-6-2014 at 09:47

_{aga, it was already established that chlorine wouldn't work for this reaction.

It is impossible to balance the equation: Al₂(SO₄)₃ (aq) + Cl₂ (g) → 2 AlClSO₄ +}SO₄^-, there's no cation to combine with the sulfate anion.
With HCl (g) the hydrogen atoms combine with the resulting sulfate^- ions to form sulfuric acid.

[Edited on 6-6-2014 by Zyklonb]

blogfast25 - 6-6-2014 at 09:55

You said "Describe" I would have failed 'O' level English if I had simply said "moisture content"

Yes but now I know what you meant!

The Volatile Chemist - 6-6-2014 at 09:58

aga, it was already established that chlorine wouldn't work for this reaction. It is impossible to balance the equation: Al₂(SO₄)₃ (aq) + Cl₂ (g) → 2 AlClSO₄ + SO₄^-, there's no cation to combine with the sulfate anion.
With HCl (g) the hydrogen atoms combine with the resulting sulfate^- ions to form sulfuric acid.

Although that reaction works, it seems like it'd be a bit unstable, sulfuric acid being stronger than HCl.

Zyklon-A - 6-6-2014 at 10:51

Quote: Originally posted by The Volatile Chemist

Although that reaction works, it seems like it'd be a bit unstable, sulfuric acid being stronger than HCl.

That reaction doesn't work, that's my point, you can't just have a sulfate^- ion floating around in solution.
BTW, HCl (aq) is a stronger acid than H₂SO₄.
Besides, that's hardly relevant. I think it forms AlClSO₄ due to it's insolubility, which drives the equilibrium to the right. Has little to do with acid strength .

[Edited on 6-6-2014 by Zyklonb]

aga - 6-6-2014 at 12:08

@blogfast25 : it was suggested, seemed easy to do, so i did it just to see for myself is all. Good practice, and experience.

Never actually smelt chlorine gas before yesterday. Ever.
Assumed that it was the same as what you get off an overdosed pool.
Boy was i wrong !
Well, now i know.

@ZyklonB

Will bubbling Chlorine thru OldAl make the same as bubbling HCl through it ?

Thats what I was wondering, I work with chlorine quite a bit

I have never encountered chlorine gas before yesterday.
Just did the experiment and there's the results.
No point telling me that it Obviously Cannot work after suggesting that it might, expecially citing experience with chlorine gas.

Looking at the Aluminium Sulphate crystals today, they are in fact slightly yellow in parts.
Nothing at all like the dramatic change with HCl.

I suspect that some Cl₂ reacts to make HCl 'on-site' so to speak, either from atmospheric water or from water of crystallisation, and then reacts.

Overall the crystals are still mainly white.

blogfast25 - 6-6-2014 at 12:45

Quote: Originally posted by The Volatile Chemist

[Although that reaction works, it seems like it'd be a bit unstable, sulfuric acid being stronger than HCl.

That reaction DOESN'T work. On the left you have Cl2, oxidation state 0. On the right the Cl shows up as chloride, Cl-, oxidation state -1. So a reduction took place: Cl(0) + e ===> Cl(-1). For that to occur something else would have to be oxidised, to provide that electron. But nothing else has changed oxidation state, so this reaction CANNOT proceed. As is also evidenced by the fact the reaction equation cannot be balanced.

Zyklon-A - 6-6-2014 at 13:03

Quote:

No point telling me that it Obviously Cannot work after suggesting that it might, especially citing experience with chlorine gas.

I know, I made a mistake, Blogfast quickly debunked the idea, soon after I suggested it:

But by all means try my ratio...
FORGET chlorine. There's no way it can react here to chloride and chloride is what we need here.

I quickly realized my mistake:

My experiences with chlorine had little to do with my assumption as to whether it would work, I just wasn't thinking clearly.
The good thing is, you've dealt with the green gas now. You always remember your first

Oh the smell... I remember my first time with the bitch. Coughing uncontrollably on the floor. My mom came in the lab to see what happened, then ran back in the house slamming the door in my face, "Don't get that 'death gas' in my kitchen!" Then I vomited 20 minutes later, and had the poison control on the phone... Good times.

aga - 6-6-2014 at 13:14

Quote:

"Don't get that 'death gas' in my kitchen!" Then I vomited 20 minutes later,

I said the same, although my vomiting was due to over-cavorting rather than the Cl₂

To be fair, i was prepared for it, doing the whole thing in a fume hood.
I thought it worthwhile to grab a gloved-handful and cautiously have a sniff.

Experience it is, although not a good one !

blogfast25 - 6-6-2014 at 13:16

"Don't get that 'death gas' in my kitchen!" Then I vomited 20 minutes later, and had the poison control on the phone... Good times.

Errrmm... you didn't read up on chlorine a bit first? See, this is why some parents get the jitters when their precious brood decides to 'do some chemistry'!

Zyklon-A - 6-6-2014 at 13:26

Not as much as I should have!
My parents are very understanding, letting me learn the hard way is something they do a lot it seems.

aga - 6-6-2014 at 13:32

Strange how attitude changes over time.
When i was young and with parents, i made gunpowder, rockets, Exciting stuff !

Now i'm old, have grown-up children and Life Assurance, those things are of little interest.

Basically i have fulfilled my primary purpose for edisting, and am now expendable, however all the Fun chemistry holds little attaction.

Odd.

The Volatile Chemist - 6-6-2014 at 14:46

Sheesh, y'all attacked me behind my back (aka not on the computer

)! I meant the second reaction wouldn't be stable,

Quote:

With HCl (g) the hydrogen atoms combine with the resulting sulfate- ions to form sulfuric acid.

but zyklonb was right, I guess HCl is stronger, and that it wouldn't matter anyways.
I knew the Cl2 wouldn't work.

AJKOER - 6-6-2014 at 22:02

Here is preparation I plan on trying that avoids H2SO4 and any Iron issues. Consider the action [Edit: possible action, no experimental evidence as of yet] Aluminium chlorohydrate on pure aqueous MgSO4:

2 AlCl2(OH) (aq) + MgSO4 (aq) ---?---) Mg(OH)2 (s) + Al2(SO4)Cl4

forming Aluminum chlorosulfate. Now, with respect to the preparation of AlCl2(OH), to quote from Wikipedia (http://en.m.wikipedia.org/wiki/Aluminum_chlorohydrate ):

"Aluminium chlorohydrate can be commercially manufactured by reacting aluminium with hydrochloric acid. A number of aluminium-containing raw materials can be used, including aluminium metal, alumina trihydrate, aluminium chloride, aluminium sulfate and combinations of these. The products can contain by-product salts, such as sodium/calcium/magnesium chloride or sulfate.[3]"

where one may be able to form aqueous AlCl3 from the action of CaCl2 on Al2(SO4)3:

3 CaCl2 + Al2(SO4)3 -----) 3 CaSO4 (s) + 2 AlCl3

See also comments at http://www.google.com/patents/US5124139

[Edited on 7-6-2014 by AJKOER]

CHRIS25 - 7-6-2014 at 01:21

Ajkoer@ Just to get back on track, It was the Aluminium Chlorosulphate "Crystal" I was trying to obtain and this can only be done via HCl, now if anyone wants I can dig up the one and only reference in the whole world where this is said, This is what caught my attention. Otherwise it is common knowledge that chlorosulphate is easy to make and is used by the water industry to floculate impurities.

Everyone Else@ I thought it was common knowledge from early 1st year school age that chlorine, hydrogen chloride and hydrogen sulphide gases were poisonous based on our history lessons?

Or has history been diluted and polluted so much with the history of waste products and European union indoctrination?

blogfast25 - 7-6-2014 at 04:39

Ajkoer@ Just to get back on track, It was the Aluminium Chlorosulphate "Crystal" I was trying to obtain and this can only be done via HCl, now if anyone wants I can dig up the one and only reference in the whole world where this is said, This is what caught my attention. Otherwise it is common knowledge that chlorosulphate is easy to make and is used by the water industry to floculate impurities.

Hmmm... it's far from clear whether the products used by the water industry correspond to the empirical formula AlClSO4. Mostly what are being described in various patents are basic polyaluminium chlorosulphates. These contain OH ions. Their formulas would be something like Al2Cl2SO4(OH)2 or similar...

'The only reference in the whole world' you are referring to is:

http://aluminium.atomistry.com/aluminium_sulphate.html

But interesting as the atomistry site generally is, atomistry isn't a very high authority source and doesn't provide any references to its assertion re. AlClSO4 whatsoever. That reduces its credibility considerably. A peer reviewed source would not allow an entry on AlClSO4 without some credible references to its existence or preparation.

Note also that we haven't proved the composition of our product yet either.

[Edited on 7-6-2014 by blogfast25]

blogfast25 - 7-6-2014 at 05:12

Quote: Originally posted by AJKOER

2 AlCl2(OH) (aq) + MgSO4 (aq) ------) Mg(OH)2 (s) + Al2(SO4)Cl4

forming Aluminum chlorosulfate. Now, with respect to the preparation of AlCl2(OH), to quote from Wikipedia (http://en.m.wikipedia.org/wiki/Aluminum_chlorohydrate ):

"Aluminium chlorohydrate can be commercially manufactured by reacting aluminium with hydrochloric acid. A number of aluminium-containing raw materials can be used, including aluminium metal, alumina trihydrate, aluminium chloride, aluminium sulfate and combinations of these. The products can contain by-product salts, such as sodium/calcium/magnesium chloride or sulfate.[3]"

where one may be able to form aqueous AlCl3 from the action of CaCl2 on Al2(SO4)3:

3 CaCl2 + Al2(SO4)3 -----) 3 CaSO4 (s) + 2 AlCl3

See also comments at http://www.google.com/patents/US5124139

Do you have any references for that equation is it just another one of your 'pretty equations'?

"Aluminium chlorohydrate can be commercially manufactured by reacting aluminium with hydrochloric acid. A number of aluminium-containing raw materials can be used, including aluminium metal, [...]" is plain nonsense. With HCl, Al goes all the way down to AlCl3. Unless they mean Al metal as a raw material in the broadest sense of the word...

As regards AlCl3 from Al sulphate and CaCl2, fine if you want a product with residual CaSO4 in it (it has some solubility). BaCl2 would work much better because BaSO4 is truly insoluble.

Still, why one would resort to displacement reactions when aqueous AlCl3 is so easy to prepare remains a mystery. Fear of HCl perhaps? If you can't stand the heat, get out of the kitchen.

[Edited on 7-6-2014 by blogfast25]

AJKOER - 7-6-2014 at 08:28

Blogfast:

Based on limited test so far, I have qualified the reaction to a maybe, perhaps requiring an elevation in pH. The original statement has been appropriately edited in my thread.

I disagree [Edited typo] with your qualification of the Wikipedia statement as "nonsense". I actually found it appropriate with my limited experimenting with aqueous Aluminum salts so far. For example, Aluminum carbonate hardly forms at all by the action of Na2CO3 on aqueous AlCl3, just mostly Al2O3 with a possible touch of the Aluminum carbonate. Similarly, one can prepared aqueous Aluminum hypochlorite, but it readily decomposes into Al2O3,..... Also, Aluminum chloride readily forms the basic salt. The chemistry of Aluminum is a little more unique, in my opinion.

[Edited on 8-6-2014 by AJKOER]

blogfast25 - 7-6-2014 at 09:16

'dismiss agree'? That's a new one...

aga - 8-6-2014 at 10:11

Why argue if there's an experiment to be done ?

AJKOER : post the parameters for an experiment, do it, i will also do it and we can post the results.

Much more interesting than randon theorising.

If you have an idea, let's give it a whirl and Test it by experimentation.

blogfast25 - 8-6-2014 at 12:23