Sciencemadness Discussion Board

AlCl3 by double displacement?

sceptic - 19-11-2022 at 03:42

I've recently been looking for an easy way to get aluminium chloride, and I came up with this idea. I don't know whether it would really work, but I've tried an experiment, and it seemed to work. If anyone's heard of this reaction before, or knows that it definitely won't work, please let me know.


Theory:

A mixture of anhydrous alum (potassium aluminium sulfate) and sodium chloride should undergo a double displacement reaction when heated, forming potassium sulfate, sodium sulfate, and aluminium chloride, which sublimates at 180 Celsius and separates from the mixture.

$$2KAl(SO_4)_2(s) + 6NaCl(s) \longrightarrow K_2SO_4(s) + 3Na_2SO_4(s) + 2AlCl_3(g)$$



Experiment:

I took 1.5 grams of powdered burnt (anhydrous) alum and mixed it thoroughly with 1.0 grams of dried, powdered table salt in a test tube, which I then sealed. The next day, I put a wooden cork with a hole through it in the end of the test tube, and put a balloon over the other end of the cork, in order to allow the air inside the test tube to expand. I heated the bottom end of the test tube in the flame of an alcohol burner. After heating it for 1-2 minutes, I noticed that droplets of water had formed near the unheated end of the test tube. I removed the test tube from the flame, opened it, and wiped the droplets out with a paper towel. I put the end in the flame again. After several minutes, I saw that a few small droplets of water had formed near the unheated end, and that what looked like a thin white mist of particles had formed on the glass inside the test tube about an inch from the mixture at the bottom. This layer was much closer to the heat than the water droplets were. The layer cleared when I moved it toward the flame, and reformed when I removed it. I also saw that the mixture at the hot end turned a yellow, then light orange colour where it was heated most strongly, and seemed to turn white again when the heating was removed. After several minutes, I saw no more changes, and the white layer didn't seem to get any thicker.

I removed the test tube from the heat, opened it, and added a few drops of water. I held a piece of pH paper in the steam that rose, and it turned dark red. I would expect this to mean that aluminium chloride was present, that it hydrolysed to aluminium hydroxide and hydrogen chloride. I assume that the white layer was aluminium chloride. The low yield seems reasonable, since the reaction would only occur where the alum and salt particles were touching.



Questions:

Is this reaction even possible? I've never heard of it before.

If this reaction isn't possible, what would have caused the acidic vapors?

What would have caused the yellow/orange color in the heated part of the test tube?

[Edited on 19-11-2022 by sceptic]

blogfast25 - 19-11-2022 at 08:39

Per se it's not a bad idea because AlCl3 is volatile, so it leaves the reagent mix, thus moving the equilibrium to the right (Le Chatelier Principle).

I had good success with something similar:

$$3\text{ZnCl}_2(anh,s)+2\text{Al}(s)\longrightarrow 3\text{Zn}(s)+2\text{AlCl}3(g)$$

This too works because the AlCl3 leaves the party.

Details are reported somewhere on this site by me. Search and ye shall find.

----------------------

I think at the very least an alcohol burner will not deliver the temperature you need.

And I think your alum/sodium chloride is off.

[Edited on 19-11-2022 by blogfast25]

sceptic - 19-11-2022 at 09:43

I've seen your preparation here, it's what gave me the idea :). However, I don't have access to anhydrous zinc chloride, and it seems no easier to make than aluminium chloride. It seems from what I've found that alum and aluminium sulfate don't significantly hydrolyse when they're heated to dehydration. If this works, it means that aluminium chloride could be made from readily available materials.

clearly_not_atara - 19-11-2022 at 09:59

Every time this discussion comes up I point out that MnCl2 can be dried by heating and should react just fine with Al. So I'll do it again. I don't have the equipment to try it but I do have a keyboard.

blogfast25 - 19-11-2022 at 10:35

Quote: Originally posted by clearly_not_atara  
Every time this discussion comes up I point out that MnCl2 can be dried by heating and should react just fine with Al. So I'll do it again. I don't have the equipment to try it but I do have a keyboard.


There are other chlorides that would work: CuCl2, PbCl2. Maybe even NaCl or CaCl2.

Are you sure MnCl2.2H2O won't hydrolyse or oxidise? I once decarbonated MnCO3 to obtain MnO and had to use an inertial blanket to prevent oxidation to +3/+4.


[Edited on 19-11-2022 by blogfast25]

Rainwater - 19-11-2022 at 16:01

Quote: Originally posted by sceptic  


$$2KAl(SO_4)_2(s) + 6NaCl(s) \longrightarrow K_2SO_4(s) + 3Na_2SO_4(s) + 2AlCl_3(g)$$


Feel free to double check the math but looks like you need more heat.

Screenshot_20221119_190039.jpg - 229kB

Edit:
Quote: Originally posted by blogfast25  


$$3\text{ZnCl}_2(anh,s)+2\text{Al}(s)\longrightarrow 3\text{Zn}(s)+2\text{AlCl}3(g)$$


Screenshot_20221119_190621.jpg - 171kB

[Edited on 20-11-2022 by Rainwater]

DraconicAcid - 19-11-2022 at 17:44

I have no direct experience or citation to back this up, but I suspect you'd be more likely to decompose the sulphate than to get AlCl3.

clearly_not_atara - 19-11-2022 at 20:08

Quote: Originally posted by blogfast25  

There are other chlorides that would work: CuCl2, PbCl2. Maybe even NaCl or CaCl2.

CuCl2 would have an excessively high reaction energy; it's been tried with concerning results. PbCl2 would potentially release lead vapor (10 Pa at 715 C), which is bad. NaCl or CaCl2 will not react, see: https://www.drjez.com/uco/ChemTools/Standard%20Thermodynamic...
Quote: Originally posted by blogfast25  

Are you sure MnCl2.2H2O won't hydrolyse or oxidise? I once decarbonated MnCO3 to obtain MnO and had to use an inertial blanket to prevent oxidation to +3/+4.

Gotta admit I'm going off Wikipedia here, which says the dihydrate dehydrates at 135 C. But I did find a paper showing that manganese (II) malonate dehydrates at 180 C before being oxidized when the malonate anion decomposes around 350 C (attached).

EDIT: Castor and Basolo report that "The stronger metal-chloride bonds deduced for the analogous manganese and nickel [hydrate] complexes favor decomposition to anhydrous chlorides as observed." [brackets mine] I have to admit I'm a little surprised because I thought that NiCl2*xH2O hydrolysed on heating. (attached)

EDIT2: You did this reaction! Apparently NH4Cl is required. Still easier than ZnCl2.
http://www.sciencemadness.org/talk/viewthread.php?tid=10847
Quote:
Feel free to double check the math but looks like you need more heat.

For a gas evolution reaction you don't need to reach equilibrium; if Keq is not extremely low you'll see some gas evolution. This is because the entropy of a gas is measured at atmospheric pressure, but much higher at low partial pressure. With that said, it's not clear that he would have gotten there.
Quote:


Attachment: mohamed1994.pdf (783kB)
This file has been downloaded 236 times

Attachment: castor1953.pdf (537kB)
This file has been downloaded 256 times

[Edited on 20-11-2022 by clearly_not_atara]

sceptic - 19-11-2022 at 21:30
Quote: Originally posted by Rainwater  

Feel free to double check the math but looks like you need more heat.


Thanks for the information! I tried to do those calculations, but I wasn't able to find the entropies of formation for the various sulfates. What calculator are you using?
Quote:

I have no direct experience or citation to back this up, but I suspect you'd be more likely to decompose the sulphate than to get AlCl3.

It did occur to me that if the sulfate was somehow being reduced to sodium sulfide, that would explain the yellow color. However, I can't think of anything that would be acting as a reducing agent.
Quote:

Every time this discussion comes up I point out that MnCl2 can be dried by heating and should react just fine with Al. So I'll do it again. I don't have the equipment to try it but I do have a keyboard.


Do you know at approximately what temperature that would react? I might try that, but it would be difficult to make an apparatus to contain and capture the fumes from a thermite-like reaction.

Rainwater - 20-11-2022 at 01:30

Quote: Originally posted by sceptic  
What calculator are you using?


https://www.sciencemadness.org/whisper/viewthread.php?tid=15...
Quote: Originally posted by sceptic  


Quote:

Every time this discussion comes up I point out that MnCl2 can be dried by heating and should react just fine with Al.


Do you know at approximately what temperature that would react?

I have never performed this reaction, but have been using that excel sheet with great success to do other reactions.

Delta g = zero at 604c
but this is only an estimate so hotter than that. The reagents might need to be a liquid before the reaction will start.
Aluminum melts just above that temperature.
You will have to try it to find out.
Test tube and alcohol lamp will easily reach that point.

[Edited on 20-11-2022 by Rainwater]

sceptic - 20-11-2022 at 03:19

Quote: Originally posted by Rainwater  

Delta g = zero at 604c
but this is only an estimate so hotter than that. The reagents might need to be a liquid before the reaction will start.
Aluminum melts just above that temperature.
You will have to try it to find out.
Test tube and alcohol lamp will easily reach that point.

[Edited on 20-11-2022 by Rainwater]


I probably could reach that temperature, but I don't think my test tube would survive it. According to this source, borosilicate glass begins to deform at 525 Celsius, and shouldn't be heated above 300 Celsius. However, since aluminium chloride is being continuously removed from the products by sublimation, it would drive the the reaction even below that temperature, so it might not be necessary to heat it that much. At least it doesn't look like it would be excessively exothermic, but I would still be hesitant to try it in a test tube.

clearly_not_atara - 20-11-2022 at 08:25

The reaction is endothermic. Oxidation of MnCl2 by atmospheric oxygen might drive the result in some heating.

In the solid phase, the reaction is only very slightly endothermic, deltaH = 37 kJ/mol. I wonder if it could be driven by the solvation enthalpy of AlCl3 in some solvent (like ether) that can be removed by heating.

[Edited on 20-11-2022 by clearly_not_atara]

blogfast25 - 20-11-2022 at 08:52

Quote: Originally posted by clearly_not_atara  



EDIT2: You did this reaction! Apparently NH4Cl is required. Still easier than ZnCl2.
http://www.sciencemadness.org/talk/viewthread.php?tid=10847

Quote:
Feel free to double check the math but looks like you need more heat.




Yes, I remember it now. Using NH4Cl is a trick because its vapour is partly dissociated into HCl + NH3. The former then protects the chloride.

I used this also on NdCl3.

ZnCl2 anh. is cheap and readily available. Very useful for chlorinating alcohols. But buying MnCl2 anh. would be more difficult.

Has anyone tried MnCl2 + Al?


[Edited on 20-11-2022 by blogfast25]

blogfast25 - 20-11-2022 at 09:48

Quote: Originally posted by sceptic  
I put the end in the flame again. After several minutes, I saw that a few small droplets of water had formed near the unheated end, and that what looked like a thin white mist of particles had formed on the glass inside the test tube about an inch from the mixture at the bottom. This layer was much closer to the heat than the water droplets were. The layer cleared when I moved it toward the flame, and reformed when I removed it. I also saw that the mixture at the hot end turned a yellow, then light orange colour where it was heated most strongly, and seemed to turn white again when the heating was removed. After several minutes, I saw no more changes, and the white layer didn't seem to get any thicker.

I removed the test tube from the heat, opened it, and added a few drops of water. I held a piece of pH paper in the steam that rose, and it turned dark red. I would expect this to mean that aluminium chloride was present, that it hydrolysed to aluminium hydroxide and hydrogen chloride. I assume that the white layer was aluminium chloride. The low yield seems reasonable, since the reaction would only occur where the alum and salt particles were touching.


Acc. Wiki, Al sulphate decomposes to alumina and SO3 between 580 and 900 C. So maybe the white layer was SO3?

https://en.wikipedia.org/wiki/Aluminium_sulfate#Chemical_rea...

Rainwater - 20-11-2022 at 11:09

https://www.chemedx.org/JCESoft/jcesoftSubscriber/CCA/CCA3/M...

blogfast25 - 20-11-2022 at 12:38

Quote: Originally posted by Rainwater  
https://www.chemedx.org/JCESoft/jcesoftSubscriber/CCA/CCA3/M...


That's info you can find in Wiki but here we're trying to avoid Cl2 (or HCl).

Rainwater - 20-11-2022 at 17:27

Got ya, so the calculator says silver chloride should work,
$$ 3AgCl + Al \longrightarrow 3Ag + AlCl_3$$
Then sublime the AlCl3 and dissolve excess AgCl in ammonia solution leaving fine silver powder.
Baaa I'm rambling again.
If you got a quarter and some HNO3, AgCl is easy to make and dry.
I'm on holiday and will try before Christmas.


[Edited on 21-11-2022 by Rainwater]

sceptic - 20-11-2022 at 20:25

Quote: Originally posted by blogfast25  

Acc. Wiki, Al sulphate decomposes to alumina and SO3 between 580 and 900 C. So maybe the white layer was SO3?

https://en.wikipedia.org/wiki/Aluminium_sulfate#Chemical_rea...


According to Wikipedia, sulfur trioxide has a boiling point of 45 Celsius, which is below that of water. If the white layer was sulfur trioxide, wouldn't it have formed much farther from the hot end of the test tube, around the area where the water condensed?

Mateo_swe - 21-11-2022 at 01:59

Have you seen Chemplayers video about preparation of anhydrous aluminium chloride from zink chloride and aluminium powder?
He doesnt get very high yields but the preparation is quite simple.
Here is a link to the video on bitchute:

Chemplayers preparation of anhydrous aluminium chloride

[Edited on 2022-11-21 by Mateo_swe]

blogfast25 - 21-11-2022 at 07:04

Quote: Originally posted by Mateo_swe  
Have you seen Chemplayers video about preparation of anhydrous aluminium chloride from zink chloride and aluminium powder?
He doesnt get very high yields but the preparation is quite simple.
Here is a link to the video on bitchute:

Chemplayers preparation of anhydrous aluminium chloride

[Edited on 2022-11-21 by Mateo_swe]


Yep, that's exactly the method I reported here and on the reference (my now defunct site) CP quotes. For some reason he's moved the video from Youtube to Bitchute.

It would be quite easy to improve on the measly yield of AlCl3.

[Edited on 21-11-2022 by blogfast25]

blogfast25 - 21-11-2022 at 07:11

Quote: Originally posted by sceptic  

According to Wikipedia, sulfur trioxide has a boiling point of 45 Celsius, which is below that of water. If the white layer was sulfur trioxide, wouldn't it have formed much farther from the hot end of the test tube, around the area where the water condensed?


Possibly. That would be very easy to test.

DraconicAcid - 21-11-2022 at 08:37

It might have been small droplets of sulphuric acid.

blogfast25 - 21-11-2022 at 10:05

Quote: Originally posted by DraconicAcid  
It might have been small droplets of sulphuric acid.


Good point, squire.

And here's my write up for the ZnCl2/Al method:

http://www.sciencemadness.org/talk/viewthread.php?tid=30150#...

Where have the years gone?


[Edited on 21-11-2022 by blogfast25]

clearly_not_atara - 21-11-2022 at 18:34

If you can buy zinc chloride, why not buy aluminum chloride? I guess you could do the reaction for the interest of it, but usually AlCl3 is a stepping stone to something else.

I guess you could make ZnCl2 from CuCl2 + Zn (less exothermic) and then use that for AlCl3, so the heat is released in two steps instead of all at once.

blogfast25 - 22-11-2022 at 07:31

Quote: Originally posted by clearly_not_atara  
I guess you could do the reaction for the interest of it, but usually AlCl3 is a stepping stone to something else.


Yes, of course.

Rainwater - 24-11-2022 at 04:28

Smoking a bird board and half lit wishing i had packed some tubes and a retort

Somewhere it was asked what other salts could be used
Heres some thermodynamics values. Took longer to format this table than to do the math
Code:
Name ΔG ΔH ΔS ΔG = 0 when T = AuCl3 -557.97 -466.90 333.38 -1673.66 C BCl -941.25 -1033 -335.85 2802.59 C BCl3 -180.69 -157.32 85.56 -2111.78 C BrCl -317.57 -292.78 90.73 -3500.13 C CdCl2 -100.46 3.72 381.41 -263.39 C CoCl2 -322.77 -231.38 334.59 -964.66 C CrCl2 -63.94 17.15 296.88 -215.37 C CuCl2 -646.28 -551.45 347.18 -1861.5 C FeCl2 -225.57 -143.64 299.95 -752.02 C FeCl3 -231.73 -185.02 171.00 -1355.12 C HCl -286.00 -308.00 -78.00 3647.0 HgCl2 -595.09 -496.22 361.96 -1644.09 C MnCl2 189 275 313 604.24 C NH4Cl 8.71 220.00 775.00 11.2 NiCl2 -353.70 -253.01 368.61 -959.55 C PbCl2 -188.67 -90.79 358.32 -526.54 C S2Cl2 -804.76 -723.41 297.82 -2702.2 C SnCl4 -946.98 -804.16 522.83 -1811.24 C ZnCl2 -25.39 76.15 371.75 -68.31 C

EDIT:
I edited the list of all the compounds that wouldn't provide a negative delta g with a temp less than 1000c
Looks like the reactions that have reported low yields or difficulty in reproduceling are endothermic.

[Edited on 24-11-2022 by Rainwater]

blogfast25 - 24-11-2022 at 07:29

Quote: Originally posted by Rainwater  


Looks like the reactions that have reported low yields or difficulty in reproduceling are endothermic.



Thanks for these values.

Have you taken into account the entropic effect of AlCl3(g)?

[Edited on 24-11-2022 by blogfast25]

Rainwater - 24-11-2022 at 09:51

Quote: Originally posted by blogfast25  

Have you taken into account the entropic effect of AlCl3(g)?

I am not sure what that is so I'm gonna say no.

When I was first reading this thread, not understanding the difficulty in creating the product, it lead to some nice googling, lots of concepts in play.

Those values were generated using the formula ΔG = ΔH - TΔS
Thermodynamic values are from here

The listed values were from reagents, anhydrous, in solid state unless data for solids was not available, at 1 atm, 0c. (Probley got the state wrong for some of those)

"ΔG=0 T=" is formed by reworking the equation to solve for t when delta g is zero

If ΔG =0; T = ΔH ÷ ΔS
No other considerations were used when generating that data.

Ramble:
I have been studying this equation for a few months and im still learning how it is used and properly applied.
Constructive criticism is welcome.

Any reading recommendations are greatly appreciated.
My current textbook isChemistry 2e by OpenStax (I like free)

I assume that when ΔG is below zero,
the chemical equilibrium will favor products.

Then by educated guessing and limited understanding, I compare the delta g of various arrangements, assuming that the probability of product formation is somehow related to the difference in delta g, the more negative delta g, the more likely the arrangement is to form. I have not found an example yet where this is not true.
I suspect this will change when I begin to learn about organic reactions.

blogfast25 - 24-11-2022 at 10:39

Quote: Originally posted by Rainwater  

I am not sure what that is so I'm gonna say no.

When I was first reading this thread, not understanding the difficulty in creating the product, it lead to some nice googling, lots of concepts in play.

Those values were generated using the formula ΔG = ΔH - TΔS
Thermodynamic values are from here

The listed values were from reagents, anhydrous, in solid state unless data for solids was not available, at 1 atm, 0c. (Probley got the state wrong for some of those)

"ΔG=0 T=" is formed by reworking the equation to solve for t when delta g is zero

If ΔG =0; T = ΔH ÷ ΔS
No other considerations were used when generating that data.

Ramble:
I have been studying this equation for a few months and im still learning how it is used and properly applied.
Constructive criticism is welcome.

Any reading recommendations are greatly appreciated.
My current textbook isChemistry 2e by OpenStax (I like free)

I assume that when ΔG is below zero,
the chemical equilibrium will favor products.

Then by educated guessing and limited understanding, I compare the delta g of various arrangements, assuming that the probability of product formation is somehow related to the difference in delta g, the more negative delta g, the more likely the arrangement is to form. I have not found an example yet where this is not true.
I suspect this will change when I begin to learn about organic reactions.


As an 'educator' I can only be glad to see someone studying the subject!

A few points.

ΔGreaction < 0 is not a sufficient condition for the reaction to proceed because themodynamics makes no pronouncements about kinetics.

For the combustion of coal the ΔGreaction is very negative, yet the reaction does not proceed at RT, only after heating. There are tonnes of examples like that.

Entropy: for gases and vapours the ΔS is much higher than lliquids or solids. This helps with the negativity of ΔGreaction, so that effect should be taken into account.

The reaction:

$$\text{CsCl}(s) + \text{Li}(l) \to \text{Cs}(g)+\text{LiCl}(l)$$

carried out at high(ish) temperature and high vacuum proceeds, with gaseous Cs being distilled off. The ΔGreaction the way you calculate it would be positive.

http://www.sciencemadness.org/talk/viewthread.php?tid=6981 (2nd page)

Another exqmple is the silicothermic reduction of magnesia (see Wikipedia)

https://en.wikipedia.org/wiki/Magnesium#Production



[Edited on 24-11-2022 by blogfast25]

Rainwater - 9-12-2022 at 13:16

Got the AgCl drying covered on a watchglass heated to 110c,
cleaned my glass, took off my gloves balled them up and damn!!! I hate silver nitrate

Soon as its dry and the rain lets up im gonna load a sample into a test tube and add some aluminum.
Should be soon

clearly_not_atara - 9-12-2022 at 13:37

On further consideration, CuCl should give a much less exothermic process than CuCl2. And it's insoluble.

blogfast25 - 9-12-2022 at 13:39

Quote: Originally posted by clearly_not_atara  
On further consideration, CuCl should give a much less exothermic process than CuCl2. And it's insoluble.


And it's not so easy to make.

blogfast25 - 9-12-2022 at 13:45

Quote: Originally posted by Rainwater  
Got the AgCl drying covered on a watchglass heated to 110c,

Soon as its dry and the rain lets up im gonna load a sample into a test tube and add some aluminum.
Should be soon


Note that some Cl would not be proof,but it would be evidence...

Try also to sublime at 180 C the suspected AlCl3.

blogfast25 - 9-12-2022 at 13:49

Acc. 'plante' you can make TiCl4 by heating titania with potassium pyrosulphate and NaCl. I wonder if something similar might work for AlCl3?

Rainwater - 9-12-2022 at 16:18

So 3.42g of AgCl. Slightly discolored was my yield.

0.11g Al alloy 6061 was filled and placed into a dry borosilicate tube, closed at 1 end.(dont laugh, my first attempt at glass work. Youtube makes it look too easy)

To this ummm.... math ... . Need 1.75g AgCl
0.41g of AgCl where added.
Reagents were crushed and mixed inside the tube.
Tube heated. Reaction witnessed.
This may produce AlCl3 but your not gonna want to do it in a lab

All contents are contained within the test tube.
Edit #3 i got to read stuff better
Edit #2:
theories,
1) the fine aluminum powder increased the reaction rate too much.
2) the aluminum reacted with oxygen in the tube

Edit #1: added pictures and video
https://youtu.be/Rvq3LjHJvsM

20221209_174644.jpg - 2.3MB 20221209_180517.jpg - 2.2MB 20221209_182513.jpg - 1.2MB 20221209_182947.jpg - 1.3MB

[Edited on 10-12-2022 by Rainwater]

[Edited on 10-12-2022 by Rainwater]

[Edited on 10-12-2022 by Rainwater]

blogfast25 - 9-12-2022 at 17:28

Quote: Originally posted by Rainwater  

This may produce AlCl3 but your not gonna want to do it in a lab



So you're trying 3 AgCl + Al ?





Rainwater - 9-12-2022 at 17:53

Ya
$$Al + 3AgCl \rightarrow AlCl_3 + 3Ag$$

[Edited on 10-12-2022 by Rainwater]

Σldritch - 10-12-2022 at 00:40

Wouldn't PbCl2 be ideal?

Rainwater - 10-12-2022 at 01:02

Quote:
Wouldn't PbCl2 be ideal?

Thermodynamics indicates that lead would be less energetic, another easily sourced metal and easily converted into chloride salt but toxic. It is easy to make anhydrous, melting point is similar. Should work.
Lets find out.

I'm gonna use up the rest of my AgCl, then ill try Pb.
I want a better way to refine silver than lye and sugar.

Rainwater - 10-12-2022 at 02:50

Here are my notes so far very promising
im not sure how to proceed with extraction and further testing.

Attachment: Alcl3 test_221210_054706.pdf (282kB)
This file has been downloaded 200 times

blogfast25 - 10-12-2022 at 08:16

Quote: Originally posted by Rainwater  

I want a better way to refine silver than lye and sugar.


Hmmm... most people would give their right hand to recover a metal with just sucrose and lye and no heat, you know? :D

[Edited on 10-12-2022 by blogfast25]

unionised - 10-12-2022 at 10:22

I wonder if anyone skilled in the use of reaction heat calculators could look at the double decomposition of sodium chloride and aluminium sulphide for me.
I have a hunch that the AlCl3 would distil out.

If it does then there's an interesting possibility of providing the heat for the reaction from a mixture of Al, NaCl and S8.

blogfast25 - 10-12-2022 at 10:54

Quote: Originally posted by unionised  
I wonder if anyone skilled in the use of reaction heat calculators could look at the double decomposition of sodium chloride and aluminium sulphide for me.
I have a hunch that the AlCl3 would distil out.

If it does then there's an interesting possibility of providing the heat for the reaction from a mixture of Al, NaCl and S8.


Al2S3 is one of the stinkiest things I've ever made, it's incredibly sensitive to moisture and water.

"And for that reason, I'm out!"

sceptic - 14-12-2022 at 06:59

Based on the table that Rainwater provided, ammonium chloride looks like the easiest chloride to use for substitution. I don't have any an the moment, but as soon as I can get some, I plan to try that reaction. Another benefit is that all of the side products are gasses, so recovering the aluminium chloride should be easy.

clearly_not_atara - 14-12-2022 at 07:59

I am certain that AlCl3 forms a complex with NH3. Not sure how easy it is to break — maybe just TsOH will do it, but maybe it's intractable.

Rainwater - 14-12-2022 at 08:20

If you could get the salt anhydrous, that stuff is difficult to dry and keep dry. Post some experiments and let us know what you discover.

blogfast25 - 14-12-2022 at 09:50

Quote: Originally posted by Rainwater  
If you could get the salt anhydrous, that stuff is difficult to dry and keep dry. Post some experiments and let us know what you discover.


I don't think moist AlCl3 can be desiccated: once moist the hydrolysis damage has already been done and is irreversible.

But preparing very dry AlCl3 is only a matter of starting from very dry precursors.

[Edited on 14-12-2022 by blogfast25]

Rainwater - 14-12-2022 at 10:59

Sorry, my last post was ambiguous, I meant the ammonia chloride salt. It is reported hygroscopic, heating it above to 300c will dry it out good.
At 337c(wiki) it decomposes but upon cooling will reform.
During my test it was nessacery to have a liquid reagent, in my case that was aluminum metal.
After some quick reading, it appears anhydrous ammonia does not react with aluminum.
Ill look through my reagents, i should have some NH3Cl and see if it works.

After doing some reading, looks like under anhydrous conditions ammonia + aluminum trichloride does form a complex, but google isnt turning up much information

Rainwater - 14-12-2022 at 14:51

Another win. Not proven to be AlCl3 but this time the reaction was easily controlled. Next time i will use an excess of aluminum.

The reaction tube was cleaned and flame dried
0.55g of NH3Cl was added to the tube.
The tube was heated until the steam stopped coming off and the salt began to condense up the tube.
0.07g of aluminum foil was pushed to the bottom of the tube until it was touching the solid ammonia chloride.
The mixture was heated until it began to melt. White smoke formed with the smell of HCl.l and condensed on the cooler parts of the tube
The mixture was kept just above the melting point with a torch until no more smoke appeared.
The apparatus was allowed to cool. All the white precipitate was scrapped off the sides of the tube into the bottom of the tube.

Upon heating, the solid migrated back up the tube with no visible liquid.
I placed a titanium rod, 3mm in diameter inside the tube as I heated it. A white powder formed on the rod.
When I dropped water down the rod, it contacted the white solid.

Dramatic pause...

Click here to see the results

B(a)P - 14-12-2022 at 14:59

I think there is something wrong with the link.

Rainwater - 14-12-2022 at 15:03

Lol. Sorry couldn't resist. It reacted violently and produced a lot of HCl vapor, by the smell of it
Edit: did it again, 1g of ammonia chloride this time.
this reaction was very well-behaved (endothermic) I might scale this up to get enough to properly test.

I can use silver nitrate solution to test the chloride content.
HCl container next to a sample to test for ammonia.
Any other testing suggestions? Whats this used for anyway

[Edited on 14-12-2022 by Rainwater]

clearly_not_atara - 14-12-2022 at 21:28

AlCl3 is a Friedel-Crafts catalyst. Refluxing it in xylene or toluene will cause (slow) disproportionation to benzene and mesitylene. It has significant covalent character and therefore dissolves in many organic solvents, while its hydrates do not. So dissolving it in e.g. chloroform may serve as a test.

sceptic - 17-12-2022 at 01:26

Thanks for testing this, Rainwater, I'm glad to see you were successful!
When you say that you heated the mixture until it began to melt, do you mean that the ammonium chloride began to melt? I'm asking because I had the impression that ammonium chloride didn't really melt, it decomposed into gaseous ammonia and hydrogen chloride. Was the aluminium that you started with used up by the end of the reaction?

Rainwater - 17-12-2022 at 06:21

Ya
So after drying the ammonia by heating it until it decomposes/sublimated. Let everything cool.
I added aluminum foil. What I believe to be a stoichiometric excess, and started heating.
I witnessed a red/orange liquid forming on the bottom of the tube. Likely to be a molton salt, or even molten aluminum. Or the red stain on my test tube left from the silver testing.

Most of the aluminum was consumed in the reaction.
I have some new theories about what is happening and have been thinking of ways to disprove each.
R1)$$ 8NH_4Cl(s) + 2Al \rightarrow 2(NH_4)AlCl_4(s) + 6NH_3(g) + 3H_2(g)$$
R2)$$ 6NH_4Cl(s) + 2Al(s) \rightarrow 2AlCl_3(s) + 6NH_3(g) + 3H_2(g)$$

For reaction #R1, I have been able to find short incomplete references to this compound, no detailed information about its synthesis or properties.

For reaction #R2, i think it is possible do to ammonia gas being lighter than HCl gas. So when the ammonia chloride decomposes, the aluminum is left sitting in a hot bath of HCl gas and the reaction proceeds like so
$$2Al + 6HCl \rightarrow 2AlCl_3 + 3H_2$$

My next experiment will answer these questions.
Q1) At what temperature does the reaction take place
Q2) does the reaction release any gas and if so what
Q3) what is the yield/weight of leftover reagents/products

Proposed appratus
My thinking is to start with a 100mL Rbf, clasien adapter with thermocouple and gas port, tubing and mineral oil bath, under an inverted cup flooded with oil.

Proposed procedure

A1) By monitoring the temperature vs power input, I should be able to see that at some point the apparatus requires more heat/power to maintain/increases the apparatus temperature. Being an endothermic reaction this will signal that the reaction has started and ended. This will require another software rewrite to add more data collection so may not happen this time. But deffenetly sticking a thermocouple in there

A2) By using a bubbler, i hope to collect any gases that are generated by this reaction. Maybe even use the gas as a method to measure reaction rate and endpoint

A3) by taking accurate weight measurements of all components of the appratus both before and after the reaction, then using the formulas R1 and R2. One formula should predict the observed weights

Edit:
I always miss something obvious, any suggestions?

[Edited on 17-12-2022 by Rainwater]

blogfast25 - 17-12-2022 at 09:20

Quote: Originally posted by Rainwater  

I always miss something obvious, any suggestions?



I remain skeptical re. your proposed reaction equations: it's easy to out write 'pretty balanced equations', but much harder to prove their veracity...

But I'll wait for your next results, before critiquing (or not?)

clearly_not_atara - 17-12-2022 at 12:53

Ammonium tetrachloroaluminate is described as melting cleanly 303 C:
https://pubs.acs.org/doi/pdf/10.1021/ic50148a012

This seems basically consistent with the other observations. The ammine AlCl3*NH3 is also described, and it is suggested to form upon boiling at 402 C. I don't know if it is feasible to obtain AlCl3 from the ammine. It does work with ZnCl2*2NH3, which suggests that the NH4Cl + Zn reaction could be worth investigating (beware of Zn vapor).

blogfast25 - 18-12-2022 at 08:22

Quote: Originally posted by clearly_not_atara  
Ammonium tetrachloroaluminate is described as melting cleanly 303 C:
https://pubs.acs.org/doi/pdf/10.1021/ic50148a012

This seems basically consistent with the other observations. .


I'm very skeptical that NH4AlCl4 can be prepared as described above. The alkali metal tetrachloroaluminates are prepared very differently, as is ammonium tetrachloroaluminate, by your own reference (which starts from AlCl3!)

As regards the 'other observations', there really isn't all that much.

Also if it was possible to oxidise Al 0 to Al+3 we would have heard about it by now.

[Edited on 18-12-2022 by blogfast25]

unionised - 18-12-2022 at 08:58

Quote: Originally posted by blogfast25  


Al2S3 is one of the stinkiest things I've ever made, it's incredibly sensitive to moisture and water.

"And for that reason, I'm out!"

In principle, I'm giving you a way to get rid of it.
:-)

Yes it's a PITA to work with, but the interesting plan it to create it and then react it, without ever needing to isolate it.

Maybe if I get bored some time...

blogfast25 - 18-12-2022 at 09:01

Quote: Originally posted by unionised  

Maybe if I get bored some time...


Please do!

Rainwater - 18-12-2022 at 16:02

Quick question, what liquid boils around -33.6c and turns blue when sodium metal is added :)
https://youtu.be/im3LMio-dyE

20221218_160332.jpg - 2.5MB20221218_155738.jpg - 2.3MB
20221218_161041.jpg - 2.4MB20221218_161407.jpg - 3.1MB

If i get to name something im calling it "ashley's oxidation"

Lab notes
Code:
2022-12-18 Test tube placed in sand bath, 900watt heater controled by a triac in insulated pot. Duckseal will be used to secure gas collection tube in reaction vessel. Tub leads to a test tube(15ml) in a dry ice acetone bath. Thermocouple is tapped to bottom of the reaction vessel. Reaction vessel weights 23.35g 0.48g of aluminum foil was weighed. When added to the reaction vessel the total weight is 24.01g That is a difference of 0.66g I will trust the weight of the aluminum packed into the test tube vs the weight of the foil in air. Math section NH4Cl mol weight 53.49g Al mol weight 26.982g NH3 mol weight 17.031g Ammonia density 0.6819 g/cm3 at −33.3 °C Predicted products are H2 gas and NH3 gas 0.0244607516mol of aluminum are in the reaction vessel Gas collection test tube is 15ml By reaction R1 8mol of NH4Cl = 6mol NH3 By reaction R2 6mol of NH4Cl = 6mol NH3 5ml of liquid ammonia seems like a safe limit. By reaction R2, which will produce the most NH3 per mol of HN4Cl, 5ml * .6819g/ml ÷ 53.49g/mol NH3 = 0.0637408861 mol of NH3 to collect, which requires 6mols or 20.47g of NH4C and 1.71g of al. So aluminum was added until 1.71g was the weight And i dont know if i can fot that much salt into the test tube Switching to rbf 100ml weight 63.54g Rbf w/ al 65.29g = 1.75g difference 1630 heat applied. 20% power Sand bath 16c Cold trap -67.8c 1640 no observation Sand bath 112.5c Cold trap -67.2c Heating played with. Full power 30seconds then minimum power, repeat when temp quits rising 1646 sand bath 155.0c no observations 1648 sand bath 165.0c no observations 1650 sand bath 178.0c no observations 1656 sand bath 193.8c no observations 1700 sand bath 202.5c no observations 1704 sand bath 226.8c no observations 1710 sand bath 271.5c no observations 1712 sand bath 282.0c no observations 1714 sand bath 292.3c no observations 1716 sand bath 309.4c no observations 1718 sand bath 321.5c no observations 1720 sand bath 330.1c no observations Heating adjusted to maintain slow temp increase 1725 sand bath 332.1c power ~50% 1726 sand bath 335.6c reaction vessel appears to have white dust on the inside where their was none before 1730 sand bath temp 334.6c amd falling, nonadjustment to power has been made since 1725 1734 sand bath 332.3c heating power increased 1738 sand bath 352.7c reaction vessel has much more dust on the inside 1742 sand bath 358.8c heating power increased 1746 sand bath 389.5c 1753 sand bath 414.0c heating stopped. Chear Liquid has condenced in my test tube. Verified by smell, ventilation was inadequate, need bigger fans Added sodium metal. Condensate turned blue. Appratus cooled and capped gently


Edit:
As it has been kindly pointed out to me, this proves nothing but
1) you can make silver metal explosively from AgCl + Al
2) you can collect HCl Solution from the solid product of (NH4Cl + Al) plus H2O
3) you can create NH3 from NH4Cl + Al
4) 2 wide open garage doors and 2 box fans are not enough to deal with liquid ammonia

[Edited on 19-12-2022 by Rainwater]
Looks like im not the first to propose this reaction but can find nothing on it being tested
https://www.sciencemadness.org/whisper/viewthread.php?tid=53...

[Edited on 19-12-2022 by Rainwater]

[Edited on 19-12-2022 by Rainwater]
Code:
2022-12-19 After reviewing videos, liquid ammonia was at first clear, but befor adding sodium, volume collected had doubled ( around 1ml total) and had a yellow tent. Reaction vessel was checked for damage, none found. A white powder condenced in the tubing leading to the cold trap Lots of white solids where scraped off the reaction vessel walls. Nothing was removed from reaction vessel. When setting up the water trap, some of the solids fell from the tubing into the universal indicator solution changing it from a light green(ph 6.9~7) To a yellow green(ph 6.4~6.7) Appratus was reconfigured to add a stright column to the reaction vessel to catch sublimated compounds A water trap with bromothymol blue solution was connected to the gas outlet in such a manner that suck-back can not occurr 0905 heating started 25% power


[Edited on 19-12-2022 by Rainwater]

blogfast25 - 19-12-2022 at 08:50

Quote: Originally posted by Rainwater  
Quick question, what liquid boils around -33.6c and turns blue when sodium metal is added :)
https://youtu.be/im3LMio-dyE

Edit:
As it has been kindly pointed out to me, this proves nothing but
1) you can make silver metal explosively from AgCl + Al
2) you can collect HCl Solution from the solid product of (NH4Cl + Al) plus H2O
3) you can create NH3 from NH4Cl + Al
4) 2 wide open garage doors and 2 box fans are not enough to deal with liquid ammonia



So did you measure the temperature of your presumed NH3?

That video is pure crap and doesn't seem to show anything relevant anyway.

Right now it looks strongly like your R1/R2 is indeed proceeding.


Rainwater - 19-12-2022 at 09:49

Ya i had a ktype thermocouple submerged in the ammonia.
It started at -67c and warmed up until it was boiling off at -33.6c
close enough for me, that's within the error of my equipment and tolerance of my nose

I agree the video is crap,i was freaking out that this worked
And i mean i was $<"(>,% thrilled. Hands shaking, knees wobbly,
I may have wet myself a little bit.
If anyone wants better video then please petition a youtuber, nerdrage is my favorite
Im a watcher not a maker.

Updated Pdf including this mornings experiments
https://www.dropbox.com/s/ahcf7pm7pd4wpsr/Alcl3%20test%20202...
Got a bunch of pictures in it 30meg file

Im trying to narrow down tempature the reaction takes place
Looks to be just north of 350c

Lab notes
Code:
2022-12-19 After reviewing videos, liquid ammonia was at first clear, but before adding sodium, the volume collected had doubled ( around 1ml total) and had a yellow tent. The reaction vessel was checked for damage, but none was found. A white powder condensed in the tubing leading to the cold trap Lots of white solids were scraped off the reaction vessel walls. Nothing was removed from the reaction vessel. When setting up the water trap, some of the solids fell from the tubing into the universal indicator solution changing it from a light green(ph 6.9~7) To a yellow green(ph 6.4~6.7) Apparatus was reconfigured to add a straight column to the reaction vessel to catch sublimated compounds A water trap with bromothymol blue solution was connected to the gas outlet in such a manner that suck-back can not occur from the larger gas trap colored with the same indicator T1 = top of the column T2 = bottom of the column T3 = sand bath (top of sand) T4 = sand bath (bottom of reaction vessel) 0905 heating started 25% power T1 18.1c T2 18.0c T4 18.3c 0915 T1 18.1c T2 18.1c T4 44.5c 0925 T1 18.5c T2 19.0c T4 124.6c the liquid is starting to rise in the gas trap bubbler. The water trap solution has lost all yellow color. Gas trap and water trap show ph of 6.9~7 0935 T1 18.4c T2 19.1c T4 182.2c 0940 T1 16.0c T2 15.7c T4 202.3c Doors were opened, outside temp -1c 0945 T1 14.5c T2 14.3c T4 214.8c 0952 T1 13.2c T2 15.3c T4 232.9c Doors shut heating power 40% 1002 T1 17.2c T2 20.4c T4 277.0c no observations 1007 heating power lowered to slow rate in increase 1027 T1 20.4c T2 26.3c T4 343.3.1c 1038 T1 20.6c T2 28.8c T3 232.7c T4 355.0c T3 was added to the top of the sand bath 1040 T1 20.4c T2 27.6c T3 241.8c T4 359.1c Color change in water trap, darker shade of green when compared to previous picture New image on left, old on right 1050 T1 21.9c T2 33.8c T3 261.9c T4 371.4c 1055 T1 22.1c T2 34.6c T3 274.3c T4 379.5c gradient of 105.2c 1100 T1 22.0c T2 36.1c T3 282.4C T4 381.4c time laps video started, 1104 T1 21.8c T2 42.3c T3 295.1c T4 387.8c found sd card, video started The strong smell of ammonia tell me my appratus is leaking Bottom of column has began showing signs of sublimation 1113 T1 23.7c T2 40.1c T3 309.6c T4 390.2c Bubbbles in water trap and gas trap with significant color change in water trap 1120 T1 23.1c T2 45.0c T3 326.5c T4 346.2c heating lowered, bubbling was aggressive 1127 T1 21.6c T2 35.5c T3 333.3c T4 317.6c Doors opened outside temp 0.5c Heating increased, reaction has begain suck-back 1130 T1 20.0c T2 34.4c T3 336.9c T4 323.3c bubbles started again, gas trap has started turning blue Heating lowered smallest ammount possible 1140 T1 18.2c T2 42.9c T3 348.7c T4 320.0c reaction completly stopped suck-back max Heating increased 1143 T1 17.8c T2 69.8c T3 354.5c T4 324.2c Reaction started again, suckback has decreased Cracking was heard. Heating off 1147 T1 20.2c T2 77.5c T3 366.7c T4 327.2c Powder made it to the water trap Reacted violently upon contact generating lots of gas Solution color went fron dark blue to light blue Shortly after the powder did not have the same effect when toughing the water 1150 T1 19.5c T2 68.2c T3 378.1c T4 329.1c heating turned back on 1205 T1 20.1c T2 48.8c T3 413.3 T4 349.9c no observations 1210 Experment stopped. Lunch Appratus cooling.


So while i wait for everything to cool down i found limited solubility data regarding AlCl3 and no data for the ammonia complexs mentioned by others.
http://chemister.ru/Database/properties-en.php?dbid=1&id...
http://chemister.ru/Database/properties-en.php?dbid=1&id...
States that NH4Cl is "sparingly soluble" in acetone while
AlCl3 is 57.5 (18°C) per 100g of solvent.

Are there any known reactions between acetone and AlCl3 which would prohibit it being used as an extraction solvent at room temp?

Can anyone provide relevant data regarding the solubility of the other perdicted products?

[Edited on 19-12-2022 by Rainwater]

Glass wool fibers that where inserted into the column as a filter, where placed into a test tube containing 1 drop of the generated ammonia solution, 5ml of water and 1 drop of universal indicator. When the fibers contacted the solution, the ph went from 8.1~12.5 to 6.4.

[Edited on 19-12-2022 by Rainwater]

blogfast25 - 19-12-2022 at 11:33

Quote: Originally posted by Rainwater  

States that NH4Cl is "sparingly soluble" in acetone while
AlCl3 is 57.5 (18°C) per 100g of solvent.

Are there any known reactions between acetone and AlCl3 which would prohibit it being used as an extraction solvent at room temp? You would then sublime the AlCl3 off, to separate it from the excess Al.



It doesn't look like it, no.

But having to carry out a post-reaction separation (work-up) would be a drawback of this method. Maybe it would be wise to try again with a strong stoichiometric excess of Al with respect to the NH4Cl?

[Edited on 19-12-2022 by blogfast25]

[Edited on 19-12-2022 by blogfast25]

clearly_not_atara - 19-12-2022 at 11:38

Quote: Originally posted by blogfast25  

Also if it was possible to oxidise Al 0 to Al+3 we would have heard about it by now.

What did you mean by this? Aluminum oxidizes very readily as long as its oxide layer is somehow compromised (many methods have been described).

As for the rest, I would agree that there are many other possible products comprising various condensations of AlCl3 and NH3. But, certainly, aluminium can be oxidised.

In fact even the reaction of aluminum with vapor-phase ammonium chloride has been reported (producing aluminum nitride):

https://www.sciencedirect.com/science/article/pii/S004060901...

and further AlN by microwaving Al+NH4Cl+urea:

https://link.springer.com/content/pdf/10.3103/S1061386210030...

Of course, AlN (mp > 1000 C) certainly can't be the liquid product. But all references I found on this reaction focus on the production of AlN.

blogfast25 - 19-12-2022 at 12:41

Quote: Originally posted by clearly_not_atara  
Quote: Originally posted by blogfast25  

Also if it was possible to oxidise Al 0 to Al+3 we would have heard about it by now.

What did you mean by this?


I meant in @Rainwater's context, of course: oxidation by NH4Cl to AlCl3. Sorry for not being clearer!

Rainwater - 19-12-2022 at 14:10

I just want to say thanks to everyone who has responded to this thread.
Currently, I'm overly excited that we have been able to apply the scientific method and developed a theory, that has merit and is worthy of further investigation and experimentation.

Joyful rant over:
Quote: Originally posted by blogfast25  

But having to carry out a post-reaction separation (work-up) would be a drawback of this method.

In the context of an amateur chemist forming this compound,
with the intent to use it for a greater purpose,
I agree 100%

I think that the observations I have witnessed were occurring in the gas phase.
Which is incredibly dangerous and hard on equipment.
I let that flask cool for 4 hours. It is currently 175c and dropping nice and slow.
I fully expect the rbf and column to be unfit for further use after this mornings experiments.
All downsides to this procedure. But a proof of concept for further research

I'm hoping that using acetone or some other solvent, this reaction can take place at lower temperatures in solution.

Thermodynamics says 11c.
Kinetics says chemicals need to be able to touch each other.

In the gas phase, ammonia chloride produces HCl which strips the passiveation layer off the aluminum letting it be attacked and freeing ammonia. I'm hoping something similar will happen in non aqueous solution.

I am also focusing efforts to easily obtained reagents.

sceptic - 19-12-2022 at 20:45

Nice experiment!
Quote: Originally posted by Rainwater  

Are there any known reactions between acetone and AlCl3 which would prohibit it being used as an extraction solvent at room temp?

Can anyone provide relevant data regarding the solubility of the other perdicted products?


I haven't been able to find any good references for a reaction, but there is this site which mentions that aluminium chloride forms complexes with ketones, which is what I would expect given how strong a Lewis acid it is. According to this site it should be soluble in a variety of haloalkanes.

Rainwater - 20-12-2022 at 07:07

So here's what I got so far
20221220_072617.jpg - 2.1MB20221220_075502.jpg - 2.1MB20221220_072634.jpg - 2.5MB
The reaction vessel (100ml RBF) has survived. And currently contains 8.02g of material.
The total reaction weight was 22.24g meaning 14.22g has moved

white solids recovered by scraping was weighed 3.37g and labled "crude"
20221220_094055.jpg - 2.2MB

At least twice that was rock hard in the column and would not scrape off.
Column washed with room temp(16c 250ml) acetone and contents scraped off with a metal rod.
Currently removing acetone by distillation.
Acetone turned to milk upon contact with products.

Compounds will be processed separately.

Noticed that at first, all the easy-to-remove powder, was free-flowing and clump-free. After moments in the air it begain to get clumpy and lumpy. Powder stored and capped as fast as possible.

Crude powder turns damp ph paper red upon contact.
Acetone solution turns damn ph paper red upon contact
Di water being used does not change ph paper color
Acetone being used does not change ph paper color

Next test will be sublimation tempature of both samples.
Edit:
Lab notes
Code:
2022-12-20 All appratus has cooled over night, capped off. Column and rbf needed to have the joint tourched to free then. Nothing appears damaged. Glass fiber that was in column was used to clean column of powder and column scrapped throughly. Powder extracted from appratus. 3.37g white/gray and placed in a test tube labled "crude" RBF weight - 71.65g Solid black mass, like concrete is stuck to the bottom Column was coated in white concrete like powder. 250ml of 16c acetone effectively cleaned column. Washing solvent was setup for distillation. Sandbath at 90c was used Solids from rbf where scraped out. All but 0.14g where recovered. Total recovery from rbf is 7.84g and placed in a test tube labled "rbf" Material left in the funnel from rbf reacts with air changing to a green color, material in test tube is a mix of black rock and gray powder Samples from di water where placed on ph paper, no color change Sample from acetone source was placed on damp ph paper, no color change. Sample from "crude" the size of a grain of table salt was placed on damp filter paper, color changed to dark red(ph 1~2) Sample from acetone roughly 0.10ml was placed on damp filter paper, color changed to dark red(ph 1~2) Sample from funnel( green compound) was placed on damp filter paper, color changed to dark red(ph 1~2) A black sample from "rbf" the size of a grain of table salt was placed on damp filter paper, color changed to a darker orange (ph 6). After time it was observed the ph paper turned dark red(ph 1~2) A gray sample from "rbf" the size of a grain of table salt was placed on damp filter paper, color changed to dark red(ph 1~2) Waiting on distillation to complete. Funnel is being left out on air to see what happens Beaker used in acetone washing is coated in white power. This powder has not changed colors upon sitting in air. Powder has a pasty cream like texture and turns ph paper bright red before paper is damp. Upon wetting paper, ph shows 3


[Edited on 20-12-2022 by Rainwater]
More notes and rambling numbers
Code:
Distillation complete, Stir bar removed. Placed on ph paper and wetted. The compound did not effect the ph paper Material was white/pink and hot. Boiling flask shows color change It should be notes that A vinyl glove was damaged and entered into the acetone wash. 7.39g of powder was recovered from the acetone wash and placed into a test tube labled "acetone" Total amount of recovered compounts 18.60g estimated 1g not recovered From appratus and wash beaker. No accurate information available for weight of gas products. Experment summary 1.71g al metal = 0.06337 mol/ al 20.47g NH4Cl = 0.38025 mol/ NH4Cl Total reagents 22.18g 6.476g of NH3 theoretical 100% yield Completed reaction 18.60g RBF 7.84g White powder collected 3.37g Acetone extracted 7.39g 83.86% mass retained 16.14% mass lost Reaction was first ran until gas products were released then stopped Next day reaction was ran untill no more gas products where released. Then held above reaction temp (350c-400c) for at least 30 minutes. Reaction vessel was slowly cooled to 100c over 6 hours. Reagents sublimated from the reaction vessel to the condenser column and no longer reacted. 8NH4Cl(s)+2Al→2(NH4)AlCl4(s)+6NH3(g)+3H2(g) 481.89g → 102.186g NH3 or 21.205% ammonia by mass 6NH4Cl(s)+2Al(s)→2AlCl3(s)+6NH3(g)+3H2(g) 176.784g → 102.186g NH3 or 57.80% ammonia by mass Acetone test tube placed in vacuum chamber

Pdf updated
https://www.dropbox.com/s/ovr1as4g39g468x/Alcl3%20test.pdf?d...

[Edited on 20-12-2022 by Rainwater]

Rainwater - 21-12-2022 at 11:50

Got something after holding 190c for 30 minutes.
Currently going to 230c.
20221221_135922.jpg - 1.2MB 20221221_140035.jpg - 2.7MB

Edit: at 230c a liquid was observed in the test tube labeled "RBF"
It quickly solidifies when removed from heat

[Edited on 21-12-2022 by Rainwater]

Σldritch - 21-12-2022 at 12:51

Why would anyone want 'AlCl3' so contaminated with Ammonium Chloride? Seems pretty useless.

If you want to moderate the reaction I would use Lead/Aluminium-alloy instead of Aluminium with PbCl2. With the lead melting and the possibility of sublimation separation should be a breeze.

Rainwater - 21-12-2022 at 13:15

I agree, but thats not my goal. It would be much easier to generate cl2 gas and burn aluminum, than to explore a previously undocumented chemical reaction
Edit: sorry thats not ment to sound as blunt as its read

This method for producing a pure reagent
...sucks...
All I have been able to do is, show without a dought that the production of anhydrous AlCl3 from the experiments described above is maybe possible.

All n all ive enjoyed myself. And learned a few things.
Best i can conclude from the data available to me is both reactions R1 and R2 are taking place.
From the product subliming out at 190c to the observed liquid at 230c. The ph of each sample shows results expected with these products. And the violent reaction with water, hygroscopic property's.

If my lab technique was more developed and matured i belive i would have been able to obtain a pure sample of the product.
I made many mistakes that were easy to avoid.

Shuting down my reaction at the first sign of gas production was a big one, insufficient surface area to condense the products, another.


[Edited on 21-12-2022 by Rainwater]

blogfast25 - 21-12-2022 at 13:57

Quote: Originally posted by Rainwater  
I agree, but thats not my goal.


So what is your goal then?

Rainwater - 21-12-2022 at 14:30

To prove whether or not this reaction takes place. Once that's done, which I believe it is, then optimization of recovery of the product
Lab notes
https://www.dropbox.com/s/ovr1as4g39g468x/Alcl3%20test.pdf?d...

[Edited on 21-12-2022 by Rainwater]

Experiment ended with the addition of water to the 3 samples

Rbf was the most reactive. Boiling the water upon contact. Strong smell of hcl, addition of water had to be keep slow to prevent boil over. ph 2~3

Crude was second most reacting with the tube heating up amd a slight smell of hcl. ph 4

Acetone was the least reactive with the tube only warming up. ph 6

All 3 tubes now have a merky black solution with a red tent

[Edited on 22-12-2022 by Rainwater]

sceptic - 22-12-2022 at 03:30

Quote: Originally posted by Σldritch  
Why would anyone want 'AlCl3' so contaminated with Ammonium Chloride? Seems pretty useless.


It would be quite difficult for me to make aluminium chloride from any reaction involving gasses as reactants, because I don't have the equipment. With this, I should be able to make small amounts in a test tube. There are some purposes for which it doesn't have to be pure. I'm planning to use it to make sodium tetrachloroaluminate.

Rainwater - 22-12-2022 at 06:02

My new theory is that during the first attempt, AlCl3 was formed. Then during the second test. when i had the powder/smoke flood past and into the water trap, i lost a large amount of product in the water trap
.
Additional losses occured with the mixture of AlCl3 and NH3 gas in the appratus

l've gots one more thing I want to try before I move on from ammonia chloride.

A mineral oil bath, at 189c.
it's a long shot but I think if it can evolve ammonia gas the oil will act as a barrier to protect the compounds formed. Ammonia is insoluble in mineral oil, so is ammonia chloride. But if the smallest amount of NH3Cl will dissolve, the reaction should take place.
Even if it is impracticaly slow

As for the liquid salt observed. Very unexpected and literature is either non-existant or behind a paywall.

The only thing I have found that fits the observations is AlH3

I've got 3 samples im going seal in glass, willing to send them to someone if they would agree to perform a proper analysis and publish all results.

Moving on from there, i have some stannous chloride(dihydrate)
Not sure if it can be made anhydrous by heating above 250c
But if it can that is a much lower reaction temperature than 330c
The second problem is thermodynamics indicates it will be more exothermic than silver chloride.
Back to the test tube.
After that I will attempt zinc and lead, ive found documents showing the successful synthesis using these reagents, so thats a little boring.
Edit:


[Edited on 22-12-2022 by Rainwater]

20221222_142451.jpg - 2.5MB

blogfast25 - 22-12-2022 at 06:35

Quote: Originally posted by Rainwater  

Moving on from there, i have some stannous chloride(dihydrate)
Not sure if it can be made anhydrous by heating above 250c
But if it can that is a much lower reaction temperature than 330c

The second problem is thermodynamics indicates it will be more exothermic than silver chloride.


It will not be possible to dehydrate SnCl2.2H2O by mere heating: the hydrate will hydrolyse badly without some kind of protection, like from NH4Cl or thionyl chloride.

But I can't see greater exothermicity being much of a problem.

Rainwater - 23-12-2022 at 17:57

Out of that list, it looks like the only salts that can easily be made anhydrous are
Code:
Name ΔG ΔH ΔS ΔG = 0 when T = AuCl3 -557.97 -466.90 333.38 -1673.66 C CoCl2 -322.77 -231.38 334.59 -964.66 C HgCl2 -595.09 -496.22 361.96 -1644.09 C MnCl2 189 275 313 604.24 C NH4Cl 8.71 220.00 775.00 11.2 NiCl2 -353.70 -253.01 368.61 -959.55 C PbCl2 -188.67 -90.79 358.32 -526.54 C

I'm going to try to optimize the NH4Cl process because that's the same stuff I put on my roses to make them red and I have plenty in stock. It's relatively nontoxic and the reaction is easy to control. And mostly because I don't have any of the others.

Using mineral oil, nothing. No gas production of any kind, 50ml of baby oil, 2g of aluminum foil, 6g of NH4Cl, 24hours at 190c with stirring. Not one bubble collected
Oil turned dark.

At 330c i started distilling the oil, and gas production began. Solids where highy reactive to water.
Next test will be directed at making a cheap disposable apparatus designed to cool the gasses below 190c and collect crude product for sublimation.

Rainwater - 26-12-2022 at 06:58

After several experments this is the cheapest appratus that worked

It goes without saying the appratus and all components must be dry before assembly, this should be performed with adequate ventilation and gas capture equipment.

This reaction is very well controlled but must be executed with extreme caution. Products are highly reactive with water. Side products are dam near explosive with water, boiling it instantly.

The entire reaction in the described appratus took less than 5 minutes once started. Multiple products are formed including 1 or more unknown compounds which react violently with water

Products:
ammonia gas
Hydrogen gas
Anhydrous aluminum trichloride
Unknown solid, liquid at 220c for a limited time. I Guess it decomposes

BoM
(1) steel food can
(1) soda can
Aluminum foil/shavings
high temp sealant, im using hvac rated "duckseal"
Metal pipe, atleast 1/4in diameter
Acetone based cleaner
Ammonia chloride 30g

Step 1) cut the bottom off the soda can and discard the rest, ensure the soda can fits into the food can
Step 2) Clean the plastic liner out of the steel can and soda can with acetone.
Step 3) drill hole into center of food can lid and soda can bottom, the diameter of the pipe used
Step 4) place pipe into center of can and loosly pack foil/shavings around it
Step 5) insert soda can over pipe, your making a bowl
Step 6) fill bowl with salt.
Step 7) place food can top over pipe and seal appratus very well.

The apparatus is burried in a sand bath with the salt at the top.

Heat is applied from the bottom. Ensuring the aluminum will be
At sufficient tempature (350c)before the salt begins to decomposes(330c).
aluminum should be in the appratus lose enough to create large surface area and maximum exposure

The outlet pipe shoud enter a cooled flask to collect the solid product.
Be prepared for large quantities of hazardous gas.

Weight of the aluminum should 3 times or more exceed the weight of ammonia chloride used. If not the output product will be contaminated with reagents. This is due to the endothermic nature of the primary reaction cooling the aluminum below the requires tempature.
When this happens reagents pass through unreacted and contaminate the product.

The additional thermal mass of the aluminium as well as the reactor arrangement in a sand bath ensures the aluminium temperature exceeds 350c for the duration of the reaction. By bottom heating quickly a temperature gradient will form, ensuring the proper tempature is maintained to prevent contamination of the final product.

After the appratus has cool it can be neturalized by slightly opening the top can and dunking the entire appratus in an extream excess of water.

My bests yield was 15.71g with minor contamination.
After sublimation 11.42g was recovered.

30g NH4Cl ÷ 53.491 g/mol NH4Cl * 2mol AlCl3 ÷ 6mol NH4Cl * 133.34g/mol AlCl3 = 24.93g maximum yield

11.42g ÷ 24.93g = 45% yield
Lots of room for improvement, great source of anhydrous ammonia.
Edit: fixed typo

[Edited on 26-12-2022 by Rainwater]

[Edited on 26-12-2022 by Rainwater]

blogfast25 - 26-12-2022 at 14:05

Quote: Originally posted by Rainwater  

Lots of room for improvement, great source of anhydrous ammonia.
Edit: fixed typo



Note that the first letter of a chemical element symbol is always a capital letter: H, N, A, O, Fe etc etc.

Nice write up.

sceptic - 29-12-2022 at 08:39

Nice documentation!

Anhydrous copper (II) chloride dissolved in acetone should be reduced by aluminium to copper and aluminium chloride. Based on this data it looks like copper chloride might have retrograde solubility in acetone, so cooler temperatures should increase solubility, as well as reducing acetone losses to evaporation. I've tried a similar experiment with copper chloride in propylene carbonate, and it turned from a dark brown solution to clear, which leads me to think that it did form aluminium chloride.

unionised - 29-12-2022 at 12:16

I'm pretty sure that acetone will react with AlCl3.

sceptic - 29-12-2022 at 20:26

I would have thought so too, but Rainwater apparently used it successfully for extraction.

Rainwater - 30-12-2022 at 02:09

It did react but after distillation