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Author: Subject: Why does aluminum chloride turn yellow?
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[*] posted on 17-5-2013 at 14:31
Why does aluminum chloride turn yellow?


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]
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[*] posted on 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.




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[*] posted on 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]
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[*] posted on 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.



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[*] posted on 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.
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[*] posted on 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]




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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 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.



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[*] posted on 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?
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[*] posted on 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.



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[*] posted on 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.
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[*] posted on 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]




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[*] posted on 18-5-2013 at 08:09


Oh, so it looks like I have to get a tank of HCl gas instead of acid
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[*] posted on 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.




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[*] posted on 18-5-2013 at 09:30


Quote: Originally posted by blogfast25  
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!


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]




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[*] posted on 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?
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[*] posted on 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]




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[*] posted on 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.




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[*] posted on 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]




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[*] posted on 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




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[*] posted on 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.




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[*] posted on 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?
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[*] posted on 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]




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