Sciencemadness Discussion Board

Unexpected product from Al(OH)3 synthesis...

Junk_Enginerd - 10-8-2019 at 23:36

I want to do some refractory experiments with metal oxides, and one of these is of course alumina. I want alumina that is finely divided from the start since I've no idea how I would grind down something from a thermite reaction or furnace oxidation...

So I decided to dissolve aluminum(junk aluminum sheet metal, certainly alloyed and not pure, may be an important factor) in HCl and neutralize with NaOH to get Al(OH)3 and NaCl which is easily separated, and then calcined to alumina. Before neutralizing, I filtered the solution to get rid of some of the aluminum alloying elements. The problem is I'm confused by the results. Initially it looked exactly as I would expect. On adding NaOH to the HCl/Al mix, a white gel precipitated. I stirred to break it up and kept adding NaOH until no more seemed to be forming.

After this I let it decant once and boiled off the water. This is something I usually do with gelling substances before purifying. It seems to convert the gel to something closer to particles, which stops the gelling and makes filtering and decanting a lot easier. It makes a world of difference for CaOH and MgOH so I figured it'd be handy here as well.

Then I added water again and let it settle. Here's the unexpected part. As far as I can tell I've got two distinctly different products. About a third of the solids settled on the bottom is a white granular powder which is what I expected from AlOH. The other two thirds has settled on top of this and is a more finely divided cream colored mass. The difference in color is quite obvious. What is that? Both seem to be mostly insoluble.

I expected contaminants, sure, but this is such a substantial amount that it has to be aluminum based. I can't imagine any aluminum alloy would have more than 10% impurities, and this is too much to not come from aluminum.

So, which is aluminum hydroxide, and what is the other part?

And while I'm at it I might as well ask what happened when I reacted aluminum with NaOH too. Most of the aluminum I threw in there ended up becoming a medium-grey heavy goop. I thought it would be forming mostly water soluble sodium aluminate? This stuff is not very soluble at all, and as far as I've understood from reading, only a fraction of aluminum reacted with NaOH will go directly to AlOH, and most will form soluble sodium aluminate.

20190810_144140.jpg - 2.6MB

Ubya - 11-8-2019 at 00:55

if you add aluminium to a NaOH solution the pH is strongly basic, aluminium is anphoteric, so it will form the aluminate, it's not strange at all. if you have an acidic aluminium solution (like your solution of AlCl3) when you add NaOH the pH slowly increases, it reaches pH 5.5 and the hydroxide starts to precipitate, the maximum yield of hydroxide is around pH 6-6.5, over 7.5 and you pretty much have only aluminate.
usual aluminium alloys have copper, zinc, magnesium, manganese, most of them are also amphoteric but the concentration is usually subpercent or maximim 5%, it should not be a big deal if used as a refractory. if you want to try something you could add more NaOH to form the aluminate, no nerd to go to pH 14, just 8 or 9 should be enough (hoping there aren't any strong effects that change that), if something remain solid, just remove that with filtration and reprecipitate the hydroxide

unionised - 11-8-2019 at 01:30

Could be something like this

"Polymorphism
Four polymorphs of aluminium hydroxide exist, all based on the common combination of one aluminium atom and three hydroxide molecules into different crystalline arrangements that determine the appearance and properties of the compound. The four combinations are:[5]

Gibbsite
Bayerite
Nordstrandite
Doyleite"

From
https://en.wikipedia.org/wiki/Aluminium_hydroxide

Junk_Enginerd - 12-8-2019 at 00:04

Quote: Originally posted by Ubya  
if you add aluminium to a NaOH solution the pH is strongly basic, aluminium is anphoteric, so it will form the aluminate, it's not strange at all. if you have an acidic aluminium solution (like your solution of AlCl3) when you add NaOH the pH slowly increases, it reaches pH 5.5 and the hydroxide starts to precipitate, the maximum yield of hydroxide is around pH 6-6.5, over 7.5 and you pretty much have only aluminate.
usual aluminium alloys have copper, zinc, magnesium, manganese, most of them are also amphoteric but the concentration is usually subpercent or maximim 5%, it should not be a big deal if used as a refractory. if you want to try something you could add more NaOH to form the aluminate, no nerd to go to pH 14, just 8 or 9 should be enough (hoping there aren't any strong effects that change that), if something remain solid, just remove that with filtration and reprecipitate the hydroxide


Yes, but aluminate is highly soluble in water right? Most of what I got from NaOH+Al wasn't soluble at all. I did some further tests and the grey slop barely even reacts with conc. HCl which is even more confusing.

Oh, so too much NaOH and I might as well dissolve the Al in NaOH directly? Good to know.

I tried neutralizing the AlCl3 with sodium bicarb instead and the results were much better. I guess it limits the maximum PH and makes it more forgiving. Now I just need to figure out a straightforward route to separate it. Good lord the AlOH is the worst slimey jelly-goo I've ever come across. CaOH and MgOH pale in comparison. Boiling off the water and redissolving it in water seemed to help enough to filter it at least, but it's a chore to even just boil it...
Quote: Originally posted by unionised  
Could be something like this

"Polymorphism
Four polymorphs of aluminium hydroxide exist, all based on the common combination of one aluminium atom and three hydroxide molecules into different crystalline arrangements that determine the appearance and properties of the compound. The four combinations are:[5]

Gibbsite
Bayerite
Nordstrandite
Doyleite"

From
https://en.wikipedia.org/wiki/Aluminium_hydroxide


Hmm. I'll look into those and see if they match my results...

S.C. Wack - 12-8-2019 at 13:29

Even with the best procedure for making the hydroxide, filtering it may be a challenge. I'm not sure that the chemistry you think should be happening is so straightforward, so one may want to find a reference with some detail. It would be unsurprising if ammonia was mentioned, or if buying alumina was the better plan.

Doped-Al2O3-fusion - 23-8-2019 at 12:21

If you want pure Al(OH)3, it is a huge pain in the butt to obtain from aluminum foil or soda cans. It's possible to purify it to a high degree with a lot of work and a lot of waste water. Iron will be the most difficult contaminate to contend with unless you use sodium hydroxide in your process, then sodium will be #1 with iron being #2 being the most difficult to remove. EDTA can help remove iron impurities as well as sodium thiosulfate, however this just adds extra steps as I've done this before. You'll just need to play around with adjusting the pH to remove contaminates in stages if using either of those reagents. I find it's easier to remove iron with a strong pH and using hydrogen peroxide to convert as much Fe(II) into Fe(III) which will be easier to filter off as ferric hydroxide. Lower the pH to between approximately 6.5 and 6.8 for best Al(OH)3 formation. From here it's a lot of decanting as the Al(OH)3 gel is too thick for filtering and rinsing. *Working with the acidic solution seems to work better for removing the excess sodium or potassium if using either of those hydroxides.

Potassium hydroxide is a better alternative to sodium hydroxide in my opinion as I find it's easier to remove potassium ions than sodium, but this is purely my opinion. The difference may not be that much either, but as stated, it's just my opinion.