Sciencemadness Discussion Board - Why does aluminum chloride turn yellow?

Why does aluminum chloride turn yellow?

Cou - 17-5-2013 at 14:31

So I'm trying to make a collection of metal chlorides, beginning with aluminum chloride. I just put foil in HCl, and it turned cloudy white... but the next day it turns yellow. Is this because it becomes a hydrate? How can I change it back to anhydrous?

[Edited on 17-5-2013 by Cou]

chemcam - 17-5-2013 at 15:29

When you make aluminum chloride from a solution of HCl it will be a hydrate, to make an anhydrous salt I believe you have to run a stream of dry hydrogen chloride over it. There are other ways as well but I don't know them off hand.

Aluminum foil is not pure aluminum, it probably has iron impurities which makes it yellow. Iron chloride is yellow and becomes darker while exposed to air.

Cou - 17-5-2013 at 15:42

Also i tried to make iron chloride for the collection, when i left the piece of iron in there it reacted kind of slowly (not fizzing and bubbling violently all over the place like the aluminum did) but when i came back it looked muddy and orange. is this iron chloride or something else? the iron was also rusted, so it made both FeCl₂ and FeCl₃

[Edited on 17-5-2013 by Cou]

DraconicAcid - 17-5-2013 at 15:44

As the iron reacts with the acid, it forms iron(II) chloride. This will react with the air to form mixed iron(III) chloride/hydroxide.

amazingchemistry - 17-5-2013 at 19:14

Could you purify the aluminum foil and then turn it into aluminum chloride? Here's how I was thinking it would work:
1. Dissolve the Aluminum foil using excess NaOH. Iron would precipitate as iron hydroxides. Which hydroxide forms would depend on the type of iron impurities you have, but all iron hydroxides are insoluble in aqueous solution. Aluminum would first form solid Al(OH)3 and then redissolve as the soluble complex ion [Al(OH4)]-

2. Decant or filter to get rid of precipitate

3. Make the solution slightly acidic. Reference suggests dry ice, maybe because of exothermic concerns. You'll get Al(OH)3 as a precipitate

4. Decant or Filter

Reference: Petrucci et. al's General Chemistry: Principles and Modern Applications.

After this I think you could get the chloride by reacting the hydroxide with hydrochloric acid. I'm not sure about this last step because I think the reaction would be exothermic (and perhaps dangerous). Also, I think the aluminum chloride would react with the water in your HCl solution, giving you back your Al(OH)3.

chemcam - 17-5-2013 at 20:23

Also if you acidify NaOH with HCl won't you then be dealing with sodium chloride impurities?

WIKI: Iron(III) Hydroxide in water 0.115 g/100 mL (20 °C)
That solubility is incorrect according to blogfast25

[Edited on 5-18-2013 by chemcam]

amazingchemistry - 17-5-2013 at 20:38

Perhaps that's why the reference recommended acidifying with dry ice. (step 3). After acidifying, you'd filter your Al(OH)3 precipitate, and the NaOH would be left in the filtrate. The HCl would be used to react with Al(OH)3 which would (presumably) give you aluminum chloride. Then again, the reaction could be dangerously exothermic and the chloride formed could react with the water in your HCl solution giving you back your Al(OH)3 so I'm not so sure about this step.

Prometheus23 - 17-5-2013 at 20:52

Aluminum chloride is such a strong lewis acid that it cannot be prepared from aqueous solutions by any means. The aluminum ion will always bind water molecules and heating will not dehydrate it to the anhydrous form.

Really the only way to prepare anhydrous aluminum chloride is to do it in the complete absence of water. There are several examples of people synthesizing it from aluminum and either dry hydrogen chloride gas or chlorine gas.

amazingchemistry - 17-5-2013 at 21:05

That's what I thought. Could you generate chlorine gas (say, by reacting bleach with acid) and then pass the gas over some high-surface-area aluminum? How would you dry it? Is there a drying agent that does not react with Cl2? Would you have to heat it? I probably wouldn't attempt this synth unless I had done a bunch of research. I have a healthy respect for chlorine gas.

chemcam - 17-5-2013 at 21:58

Chlorine can be dried by calcium chloride. But I would rather use dry hydrogen chloride like I suggested at the top of this thread. HCl gas is easier to work with in my opinion.

Cou - 17-5-2013 at 22:05

Quote: Originally posted by Prometheus23

Aluminum chloride is such a strong lewis acid that it cannot be prepared from aqueous solutions by any means. The aluminum ion will always bind water molecules and heating will not dehydrate it to the anhydrous form.

Really the only way to prepare anhydrous aluminum chloride is to do it in the complete absence of water. There are several examples of people synthesizing it from aluminum and either dry hydrogen chloride gas or chlorine gas.

Can I still prepare hydrous aluminum chloride, or will it always become aluminum hydroxide?

chemcam - 17-5-2013 at 22:30

Yes of course, you would never get a hydroxide by using HCl. You should some reading on acids and bases. It might help you understand a little more. But anyway if you had pure aluminum and used HCl you get hydrate.

Cou - 17-5-2013 at 22:51

I just want to make a collection of chlorides with HCl. Like zinc chloride from pennies, iron chloride, magnesium chloride, nickel chloride, aluminum chloride.

blogfast25 - 18-5-2013 at 05:19

Quote: Originally posted by chemcam

Iron hydroxide has soluabilty of 0.1g in 100ml at 20C so there still could be iron impurities unless its very cold. Also if you acidify NaOH with HCl won't you then be dealing with sodium chloride impurities?

Where do you get this nonsense from? Both Fe (II) and Fe (III) hydroxide are extremely insoluble. Wiki states a solubility of 0.00005255 g/100 ml for Fe(OH)2 and that sounds roughly right. Fe(OH)3 on the other hand is so insoluble that Fe (III) starts dropping out of solution from pH 4 to 5, its solubility product (Ksp) is remarkably small. Careful with spreading disinformation!

Quote: Originally posted by Cou

Can I still prepare hydrous aluminum chloride, or will it always become aluminum hydroxide?

Anhydrous AlCl3 can be prepared from Al powder (or shavings) and dry HCl or dry Cl2 with heat (and a few interesting small scale lab methods). The reaction of Al with dry HCl or Cl2 is NOT for the beginner, except in miniscule quantities!

From aqueous solution, the hexahydrate AlCl3.6H2O can be prepared as follows. Dissolve Al in a large excess of 36 % HCl, then carefully boil in the solution till there’s very little liquor left and ice it. Then gas the cool concentrated AlCl3 solution with HCl gas: AlCl3.6H2O precipitates out because it isn’t soluble in sat. cold HCl solution. I did this once and it works.

Quote: Originally posted by Cou

I just want to make a collection of chlorides with HCl. Like zinc chloride from pennies, iron chloride, magnesium chloride, nickel chloride, aluminum chloride.

You need to read up on ‘hydrolysis’. Several of these chlorides (zinc, iron, aluminium) are difficult to isolate from aqueous liquors because of hydrolysis, causing hydroxychlorides or even hydroxides to form if you don’t get the conditions right. FeCl3 is extremely soluble and thus quite hard to crystallise (as hexahydrate). FeCl2 is much easier to make but oxidises like crazy to Fe(III) when exposed to air.

Hydrolysis of metal chlorides is caused by this equilibrium (example for ferric cations):

[Fe(H2O)_n]³⁺ + H₂O < === > [FeOH(H2O)_n-1]²⁺ + H₃O⁺

Protons in the water mantle of the central cation are repulsed by the central electrostatic field and end up protonating water. These hydrated cations thus act as weak Bronsted acids.

That is what causes these solutions to react acidic and hydroxychlorides to form. Only excess free can acid pushes the equilibrium back to the left. It explains also why ferric solutions are slightly thermochromic: on heating the above equilibrium shifts somewhat to the right and the FeOH²⁺ ion is more darkly coloured than Fe³⁺ which is reportedly almost colourless. Ferric solutions darken visibly when heated.

The preparation of many other D-block chlorides is often even harder because many are predominantly covalent compounds that can't be prepared from aqueous acid: see TiCl4, ZrCl4, HfCl4, WCl6 and many others. AlCl3 is the 'fence sitter' here: predominantly covalent but sufficiently polarised not to fall apart completely in acidified water.

[Edited on 18-5-2013 by blogfast25]

Cou - 18-5-2013 at 08:09

Oh, so it looks like I have to get a tank of HCl gas instead of acid

blogfast25 - 18-5-2013 at 09:22

Quote: Originally posted by Cou

Oh, so it looks like I have to get a tank of HCl gas instead of acid

No, you can make small amounts of fairly pure HCl gas safely by using HCl generators: NaCl + conc. H2SO4. There are various threads on it on this forum, just use the search facility. For anhydrous AlCl3, your HCl has to be very dry though, so you need a scrubber to get the last bits of water out.

To precipitate AlCl3.6H2O 'wet' HCl gas is obviously not a problem.

chemcam - 18-5-2013 at 09:30

Quote: Originally posted by blogfast25

Quote: Originally posted by chemcam

Straight from wiki Iron(III) Hydroxide
Solubility in water 0.115 g/100 mL (20 °C)

Oops just noticed it redirects to Iron(III) oxide-hydroxide.

[Edited on 5-18-2013 by chemcam]

Cou - 18-5-2013 at 09:39

so last question: can i make zinc chloride, magnesium chloride and lead chloride by dissolving the metals in HCl acid, and boiling the acid to leave behind solids? or will it change to hydroxides and everything just like aluminum does?

chemcam - 18-5-2013 at 09:46

You can boil away most of the excess water/HCl but don't go to dryness, let the rest evaporate on its own or you risk decomp to OXIDE, depending on each metal though.

-EDIT
I like to use steam to heat the container holding the metal salt, keeps temperatures down.

[Edited on 5-18-2013 by chemcam]

blogfast25 - 18-5-2013 at 11:12

Quote: Originally posted by chemcam

Straight from wiki Iron(III) Hydroxide

[Edited on 5-18-2013 by chemcam]

For hydroxides, simply use the following rule of thumb: all hydroxides are insoluble (to varying degrees, see Ksp), except for those of Group I which are highly soluble and those of Group II, some of which are sparingly soluble.

blogfast25 - 18-5-2013 at 11:17

Quote: Originally posted by Cou

so last question: can i make zinc chloride, magnesium chloride and lead chloride by dissolving the metals in HCl acid, and boiling the acid to leave behind solids? or will it change to hydroxides and everything just like aluminum does?

Zinc and magnesium chloride can be crystallised from acid solution but do this GENTLY.

Lead chloride is poorly soluble in water and lead is poorly soluble in most acids, including HCl, except in HNO3. Lead chloride is prepared by precipitating it from a solution of lead nitrate or lead acetate with a minimum amount of a soluble chloride (sodium, potassium or ammonium). Filter off, wash and dry.

Pb(NO3)2(aq) + 2 NH4Cl(aq) === > PbCl2(s) + 2 NH4NO3(aq)

[Edited on 18-5-2013 by blogfast25]

Fantasma4500 - 18-5-2013 at 13:33

sorry if i missed out on major points, but aluminium chloride isnt easy to make really..
it tends to go into hydroxides
you should try NaCl + 98% H2SO4 and then cap the beaker off with many layers of aluminium foil.. this works for aluminium nitrate i know, or at least with NO2

very interesting idea to collect metal chlorides..

for FeCl3, i read in this thread it turns into hydroxides, that sounds reasonable.. to isolate that i think you should just put it in a airproof container, perhaps include a little bag of SiO2 as they put in shoes etc.

i noticed FeCl3 decomposes somewhat at about past 100*C, as i was once looking into it as a method for making Fe2O3
unfortunately this creates anhydrous HCl gas we have concluded (yes.. this might be of use to you if you dont want to use Cl2!)

or no wait.. thats incorrect, it contains water molecules so it will not be 100% anhydrous, concentrated but not concentrated enough to make other anhydrous salts, dont know if you can somehow make it 100% anhydrous

blogfast25 - 18-5-2013 at 13:39

Quote: Originally posted by Antiswat

sorry if i missed out on major points, but aluminium chloride isnt easy to make really..
it tends to go into hydroxides
you should try NaCl + 98% H2SO4 and then cap the beaker off with many layers of aluminium foil.. this works for aluminium nitrate i know, or at least with NO2

very interesting idea to collect metal chlorides..

for FeCl3, i read in this thread it turns into hydroxides, that sounds reasonable.. to isolate that i think you should just put it in a airproof container, perhaps include a little bag of SiO2 as they put in shoes etc.

i noticed FeCl3 decomposes somewhat at about past 100*C, as i was once looking into it as a method for making Fe2O3
unfortunately this creates anhydrous HCl gas we have concluded (yes.. this might be of use to you if you dont want to use Cl2!)

or no wait.. thats incorrect, it contains water molecules so it will not be 100% anhydrous, concentrated but not concentrated enough to make other anhydrous salts, dont know if you can somehow make it 100% anhydrous

No, no: FeCl3.6H2O is possible to prepare as a solid, it's just not easy because it only crystallises quite slowly and only from very concentrated, very acidic solutions. There's a part of a thread by me on it, with pics of the yellow solid. Search.

Anhydrous FeCl3 requires dry iron and very dry Cl2 and some heat, much like anh. AlCl3.

amazingchemistry - 19-5-2013 at 09:14

Am I correct in assuming that AlCl3 is a very strong Lewis acid because of both the chlorine electron withdrawing groups and the fact that Al has an incomplete octet?

blogfast25 - 19-5-2013 at 10:16

Quote: Originally posted by amazingchemistry

Am I correct in assuming that AlCl3 is a very strong Lewis acid because of both the chlorine electron withdrawing groups and the fact that Al has an incomplete octet?

Yes. Anhydrous AlCl3, anhydrous FeCl3 (also anh. WCl6) are strong Lewis acids. AlCl3 is the main catalyst used in Friedel-Crafts alkylation and acylation reactions. But small amounts of water are disastrous for them due to intense hydrolysis.

[Edited on 19-5-2013 by blogfast25]

Cou - 19-5-2013 at 17:02

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.

blogfast25 - 20-5-2013 at 04:04

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.

DubaiAmateurRocketry - 20-5-2013 at 08:47

Does aluminum chloride decompose at 3100 degree ?

blogfast25 - 20-5-2013 at 10:00

Quote: Originally posted by DubaiAmateurRocketry

Does aluminum chloride decompose at 3100 degree ?

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]

AJKOER - 21-5-2013 at 13:41

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]

kmno4 - 21-5-2013 at 23:56

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"

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]

woelen - 22-5-2013 at 00:40

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.

Bezaleel - 22-5-2013 at 04:52

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.

blogfast25 - 22-5-2013 at 04:55

There are no coloured compounds of Al(III), unless it's combined with a coloured anion and perhaps in some organometallic compounds.

Bezaleel - 22-5-2013 at 06:28

ClO2, crystallised as adduct, causes a yellow colour.

AJKOER - 22-5-2013 at 06:44

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]

woelen - 22-5-2013 at 06:59

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.

ScienceSquirrel - 22-5-2013 at 08:50

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!

blogfast25 - 22-5-2013 at 08:58

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?

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]

AJKOER - 23-5-2013 at 19:25

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.

Mildronate - 24-5-2013 at 07:30

how about using hcl solution in organic?

blogfast25 - 24-5-2013 at 09:01

Quote: Originally posted by Mildronate

how about using hcl solution in organic?

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]

Mildronate - 25-5-2013 at 08:32

no hcl in toluene + alumiiniem =?

blogfast25 - 25-5-2013 at 12:02

Quote: Originally posted by Mildronate

no hcl in toluene + alumiiniem =?

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.

Mildronate - 26-5-2013 at 20:07

bulshit thermoddinamicaly its posible look at gibs energy. there need al without oxide layer

woelen - 26-5-2013 at 22:38

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.

blogfast25 - 27-5-2013 at 04:44

Quote: Originally posted by Mildronate

bulshit thermoddinamicaly its posible look at gibs energy. there need al without oxide layer

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]

Bezaleel - 28-5-2013 at 04:53

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.