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Cou
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Now i'm trying something new: mix hydrochloric acid and bleach in a bottle to make chlorine gas (i already did that) heat a piece of steel wool and
put it in there.
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blogfast25
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Yes, that should work but at a minimum make sure you can lead any excess chlorine through an NaOH scrubber. Chlorine in your immediate environment is
seriously toxic and will also make you choke, even at modest concentrations! Use something like a horizontal test tube with the steel wool at the
bottom, a two holed bung and some tubing to direct the Cl2 to where you want it to be and not where you DON't want it to be. Start heating the bottom
of the test tube when the Cl2 has fully displaced the air in the tube.
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DubaiAmateurRocketry
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Does aluminum chloride decompose at 3100 degree ?
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blogfast25
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The decomposition of AlCl3 is essentially an equilibrium reaction:
AlCl3(g) < === > Al(g) + 3/2 Cl2(g) ... (1), which shift to the right with increasing temperature. The degree of dissociation (decomposition, if
you prefer) is related to temperature T and Gibbs Free energy by Nernst's equation:
ΔG = - RT ln K. with ΔG the change in Gibbs Free energy from left to right and K the equilibrium constant of (1).
At 3100 C I'd imagine the equilibrium to lay very much to the right of (1).
[Edited on 20-5-2013 by blogfast25]
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AJKOER
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The reason AlCl3 turns yellow is, I suspect, it willingness upon exposure to moist air to readily form the hexahydrate AlCl3·6H2O, which also happens
to be white to yellowish in color (see Wikipedia http://en.wikipedia.org/wiki/AlCl3 ).
So anhydrous Aluminium Chloride, by itself in a sealed vessel, doesn't turn yellow, it is transformed upon exposure to water vapor into a hydrate,
which is yellow.
The answer relating to necessity of an Fe impurity is not correct as here is a report for the pure anhydrous AlCl3 (see http://www.grrexports.com/Products/AlCl3.html ) detailing the properties of AlCl3 to quote:
"Aluminium Chloride anhydrous
White powder, easily turns to yellow, green or grey ·Fumes in air ·Strong odor of HCl"
[Edited on 21-5-2013 by AJKOER]
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kmno4
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Quote: Originally posted by AJKOER  | The reason AlCl3 turns yellow is, I suspect, it willingness upon exposure to moist air to readily form the hexahydrate AlCl3·6H2O, which also happens
to be white to yellowish in color (see Wikipedia http://en.wikipedia.org/wiki/AlCl3 ).
So anhydrous Aluminium Chloride, by itself in a sealed vessel, doesn't turn yellow, it is transformed upon exposure to water vapor into a hydrate,
which is yellow.
The answer relating to necessity of an Fe impurity is not correct as here is a report for the pure anhydrous AlCl3 (see http://www.grrexports.com/Products/AlCl3.html ) detailing the properties of AlCl3 to quote:
"Aluminium Chloride anhydrous
White powder, easily turns to yellow, green or grey ·Fumes in air ·Strong odor of HCl"
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It is of course not true.
Just
"Producer, Manufacturer, Supplier of
Aluminium Chloride anhydrous"
wants to sell his coloured product as "pure".
Maybe even it is pure, but 0,1% of Fe (or other matals) is enough to cause coloration.
If it is really Fe - easy to check with thiocyanate.
Pure AlCl3 is perfectly colourless, no matter anhydrous or hydrated.
BTW.
Completely anhydrous AlCl3 is inactive in Friedel–Crafts reactions. Traces of water make it an active catalyst 
Amazing....
[Edited on 22-5-2013 by kmno4]
Слава Україні !
Героям слава !
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woelen
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I have very pure hydrated AlCl3.6H2O and it is perfectly colorless/white. It is a crystalline solid, which is not really interesting. A very tame
reagent, soluble in water, and showing no great reactivity.
I also have anhydrous AlCl3 and that is pale yellow. The yellow color is due to iron contamination. Most commercial anhydrous AlCl3 is yellow. I think
it has to do something with the process, used to make the anhydrous compound. Anhydrous AlCl3 is a completely different beast than AlCl3.6H2O. When I
open the bottle of anhydrous AlCl3, then I am greeted by a big cloud of HCl fumes, which are inside the bottle, under pressure. AlCl3 violently reacts
with water, producing a hissing noise, a lot of heat and fumes of HCl. Handling anhydrous AlCl3 requires great care and storing AlCl3 is hard. It is
very corrosive.
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Bezaleel
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I do have a critical question about the matter of iron contamination.
A long time agon, I dissolved a few pieces of household aluminium foil in some 50 ml of HCl solution. This would not form crystals after letting stand
in a dry atmosphere. When I evaporated the solution with an alcohol burner, an intensely yellow solution formed, and ultimately a light yellow
compound that looked similar to floury potatoes that cooked too long. This compound was extremely hygroscopic, much more than CaCl2.
On another occasion, I took a piece from the foil and dissolved it in NaOH solution. In the end I had obtained a clear, colourless solution, WITHOUT
any precipitate on the bottom of the beaker (even the greyish oxides from the foil had dissolved after a week or so). When it dried in, off-white
crusts formed in the beaker.
If the yellow colour is indeed due to iron from the foil, then why is de solution of the foil in NaOH solution colourless and does it have no
precipitate of ironhydroxides?
If I may suggest an alternative explanation for the yellow colour (I'm only guessing here) -- could it be due to the combined effect of aluminium
hydrolysis and presence of excess chloride to form chlorine oxides of any kind? Similar to what happens on evaporating a solution of a chloride with
some HCl and H2O2.
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blogfast25
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There are no coloured compounds of Al(III), unless it's combined with a coloured anion and perhaps in some organometallic compounds.
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Bezaleel
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ClO2, crystallised as adduct, causes a yellow colour.
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AJKOER
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OK, I think it is interesting that the described colors of AlCl3 upon exposure to air and moisture (reputedly forming the hydrate AlCl3.6H2O) include
yellow and green.
Perhaps the hydrate has traces of chlorine water? But what could be the source of the chlorine? Perhaps trace oxidation of the reported gaseous HCl
with O2 in the presence of light or/and Al2O3 or otherwise?
Now, here is are simple potential tests for the chlorine contamination hypothesis. Expose a small amount of yellow AlCl3.6H2O, spread out in a thin
layer, to sunlight or gentle heat. If the sun (heating) does not remove the color, it could be an Iron (or other metal) impurity afterall.
Sunlight reactions, for example:
2 Cl2 + 2 H2O <---> 2 HOCl + 2 HCl
2 HOCl --uv--> 2 HCl + O2 (g)
------------------------------------------
Net
2 Cl2 + 2 H2O ---> 4 HCl + O2 (g)
One can also apply a direct test for Fe, as has been suggested.
[Edited on 22-5-2013 by AJKOER]
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woelen
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| Quote: Originally posted by Bezaleel | | [...]If I may suggest an alternative explanation for the yellow colour (I'm only guessing here) -- could it be due to the combined effect of aluminium
hydrolysis and presence of excess chloride to form chlorine oxides of any kind? Similar to what happens on evaporating a solution of a chloride with
some HCl and H2O2. |
I hardly can believe that this occurs. You need a really strongly oxidizing environment
to get oxides of chlorine and with aluminium around in HCl you have a strongly reducing environment. I see no mechanism at all, which leads to
formation of chlorine oxides.
It is an interesting observation though that in alkaline environments you don't get the yellow color.
Iron(III) in combination with chloride ion has a very strong yellow color. Even tiny amounts of this can give a yellow color to a white or colorless
compound. But this also is not a really strong explanation. In alkaline environment I would expect at least a tiny amount of flocculent brown stuff
from Fe2O3 or flocculent greyish stuff from Fe3O4.
Maybe the yellow color is due to some other metal in the aluminium. Copper might be a candidate (this gives yellow complexes at very high chloride
concentration and it also is noticeably yellow/mustard colored when present as contamination in certain solid chlorides which also have acid as
impurity). In strongly alkaline solution, the copper may be present as blue complex, but this color is weaker and you may have missed it. Again, this
is just some guestimating from my side.
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ScienceSquirrel
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| Quote: Originally posted by Bezaleel | I do have a critical question about the matter of iron contamination.
A long time agon, I dissolved a few pieces of household aluminium foil in some 50 ml of HCl solution. This would not form crystals after letting stand
in a dry atmosphere. When I evaporated the solution with an alcohol burner, an intensely yellow solution formed, and ultimately a light yellow
compound that looked similar to floury potatoes that cooked too long. This compound was extremely hygroscopic, much more than CaCl2.
On another occasion, I took a piece from the foil and dissolved it in NaOH solution. In the end I had obtained a clear, colourless solution, WITHOUT
any precipitate on the bottom of the beaker (even the greyish oxides from the foil had dissolved after a week or so). When it dried in, off-white
crusts formed in the beaker.
If the yellow colour is indeed due to iron from the foil, then why is de solution of the foil in NaOH solution colourless and does it have no
precipitate of ironhydroxides?
If I may suggest an alternative explanation for the yellow colour (I'm only guessing here) -- could it be due to the combined effect of aluminium
hydrolysis and presence of excess chloride to form chlorine oxides of any kind? Similar to what happens on evaporating a solution of a chloride with
some HCl and H2O2. |
If you used hardware store or even technical grade hydrochloric acid I suspect that the iron came from the acid and not the aluminium.
This would explain why you did not see any in the second base reaction.
I favour the presence of trace amounts of iron and even a trace would be enough to colour the solid. A little gunk goes a long way!
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blogfast25
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| Quote: Originally posted by Bezaleel |
If I may suggest an alternative explanation for the yellow colour (I'm only guessing here) -- could it be due to the combined effect of aluminium
hydrolysis and presence of excess chloride to form chlorine oxides of any kind? Similar to what happens on evaporating a solution of a chloride with
some HCl and H2O2. |
No, I can’t see that happening either. I’ve evaporated AlCl3 solutions to near dryness and didn’t see anything like that happening.
| Quote: Originally posted by AJKOER | Perhaps the hydrate has traces of chlorine water? But what could be the source of the chlorine? Perhaps trace oxidation of the reported gaseous HCl
with O2 in the presence of light or/and Al2O3 or otherwise?
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Any colour is almost 99.9 % due to contaminating coloured cations. Certainly ferric ions explain a lot, perhaps there are others too. I’ve seen
yellow zirconyl chloride too, also due to iron (III). You and Bezaleel are looking to far ashore.
As Squirrel states, hardware store HCl is almost always iron contaminated. The only Fe free HCl I have is pro analysis.
[Edited on 22-5-2013 by blogfast25]
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AJKOER
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Per one of my favorite sources (http://aluminium.atomistry.com/aluminium_trichloride.html ) to quote:
"A simple method of preparation [referring to AlCl3] is said to consist in heating crude alumina or clay to redness in a current of hydrogen chloride
and carbon disulphide vapour, and purifying the aluminium chloride so obtained by sublimation over iron filings.
Auminium chloride, purified by sublimation over aluminium, forms white, lustrous, six-sided plates which are said by Seubert and Pollard to possess
rhombic symmetry. The slightly impure chloride is usually yellow owing to the presence of a little ferric chloride."
So yes, the AlCl3 that is yellow is prepared via the crude method with sublimation over iron filings, hence the Fe contamination. The good stuff is
white.
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Mildronate
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how about using hcl solution in organic?
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blogfast25
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You mean an acid solution of AlCl3 as catalyst? Once hydrolysed/solvated AlCl3 isn't a Lewis acid anymore.
Anhydrous AlCl3 can accept to share an electron pair from a donor to complete it's shell, that's what a Lewis acid does and that is the basis of it's
catalytic power. But hydrated (solvated) Al3+ cations or hydrolysed versions of them can't do that because they already have full shells.
Look up 'Lewis acid'.
[Edited on 24-5-2013 by blogfast25]
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Mildronate
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no hcl in toluene + alumiiniem =?
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blogfast25
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There's thread on that by 'peach', using dichloromethane as a 'solvent'. The results were inconclusive at best. In my view room temperature reaction
of Al with dry HCl in an inert solvent cannot proceed without some catalysis.
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Mildronate
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bulshit thermoddinamicaly its posible look at gibs energy. there need al without oxide layer
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woelen
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Theory and practice are very different things. Even if a reaction is thermodynamically favorable, then it may be difficult for the reaction to
proceed. It might be that the activation energy is too high, or simply no mechanistic pathway exists for the reaction.
For the reaction between Al and HCl it is necessary to break the H-Cl bond and some energy is needed for that. If that energy is not available (e.g.
low temperature), then the reaction does not occur. A nice analogue is a ball sitting on the top of a large hill, but on the top there is a little
wall around the ball. The ball just lies there, sitting against the wall and some energy is needed to push it over the wall, before it can roll down
the hill and release its potential energy.
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blogfast25
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By your reasoning the whole world would be on fire, constantly.
Take for instance C + O2 === > CO2. It's highly thermodynamically favourable but do you see any coal or cokes spontaneously
bursting into flames? The reason that doesn't happen is that at RT the average speed of the atoms/molecules is far too low for reactive (as
opposed to elastic) collisions to take place. To overcome this so called kinetic obstacle the reagents have to be heated up, which increases the
average speed (and thus their average kinetic energy) of their atoms/molecules. Once a threshold is exceeded a sufficient proportion of collisions
become reactive collisions and the reaction starts proceeding on a significant, macroscopic scale.
Search for 'collision theory chemistry'.
[Edited on 27-5-2013 by blogfast25]
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Bezaleel
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Moreover, Al is normally coated by a thin (but protective) layer of oxide. That's why if you put a piece of Al foil in muriatic acid, the reaction
takes some time to get started.
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