Sciencemadness Discussion Board - Do calcium hydroxide react with aluminum sulfate?

Do calcium hydroxide react with aluminum sulfate?

bluamine - 29-11-2015 at 12:14

Hi everyone!!
Can this reaction really happen:
3Ca(OH)2+Al2(SO4)3=2Al(OH)3+3CaSO4??!!
I found it on s YouTube video & I know that it is not a trustful science source, especially when it comes to a video existing only on an unknown channel

[Edited on 29-11-2015 by bluamine]

gdflp - 29-11-2015 at 12:25

Under what conditions? If they are both dissolved in water then yes, it is just a simple metathesis reaction.

aga - 29-11-2015 at 12:31

Quote: Originally posted by gdflp

a simple metathesis reaction.

Had to look that up, it being the second time i saw it in relation to something that appeared simple.

Wiki:

"A salt metathesis reaction (from the Greek μετάθεσις, "transposition"), sometimes called a double replacement reaction"

Much better (for me at least) !

bluamine - 29-11-2015 at 12:32

Quote: Originally posted by gdflp

Under what conditions? If they are both dissolved in water then yes, it is just a simple metathesis reaction.

i am comfused about it because calcium hydroxyde is not soluble at least in water

gdflp - 29-11-2015 at 12:34

It is to some extent, about 2g/L at 20°C. Much more soluble than alumina.

[Edited on 11-29-2015 by gdflp]

aga - 29-11-2015 at 12:36

At STP it's not.

As gdflp asked, "under what conditions" ?

Makes a whole lot of difference.

Edit:

Now you mention it (and i look up the other three) none of them are particularly soluble in water apart from the Al Sulph.

[Edited on 29-11-2015 by aga]

bluamine - 29-11-2015 at 12:56

Quote: Originally posted by aga

At STP it's not.

As gdflp asked, "under what conditions" ?

Makes a whole lot of difference.

Edit:

Now you mention it (and i look up the other three) none of them are particularly soluble in water apart from the Al Sulph.

[Edited on 29-11-2015 by aga]

Well i can't increase thé pression so much at home, but if it must be done at a moderately high tempreature, I can do it..

aga - 29-11-2015 at 12:59

It would probably be better blueamine to say what you Want to do, and what reagents you have available.

Please put stuff like this in the Begnings Topic in future.

bluamine - 29-11-2015 at 13:12

Quote: Originally posted by aga

It would probably be better blueamine to say what you Want to do, and what reagents you have available.

I just want to produce some amounts of alumina. I do know many other methods to do that, but if this method work, it would be the easiest one. Unless the sand thermite methods can make me able to formalize alumina

Quote: Originally posted by aga

Please put stuff like this in the Begnings Topic in future.

Sure

aga - 29-11-2015 at 13:23

Hmm.

I didn't know what 'alumina' meant and had to go to wiki again : aluminium(III) oxide looks hard to do.

If you want Aluminium Hydroxide Al(OH)₃ just try making Aluminium Sulphate, screw it up, and it's present in bucketloads.

http://www.sciencemadness.org/talk/viewthread.php?tid=30313

bluamine - 29-11-2015 at 14:27

Quote: Originally posted by aga

I didn't know what 'alumina' meant and had to go to wiki again : aluminium(III) oxide looks hard to do.

Unfortunately, this is what I am looking for

[Edited on 29-11-2015 by bluamine]

aga - 29-11-2015 at 14:33

Buy some ?

Might be easier/cheaper.

gdflp - 29-11-2015 at 15:02

Do you have any basic sodium or potassium salts? Hydroxide, carbonate, and bicarbonate solutions will all precipitate aluminum hydroxide from a solution of an aluminum salt, careful of the release of CO2 if the latter two are used though. A stoichiometric amount should be used as well due to the amphotericity of aluminum hydroxide. From here, calcining the aluminum hydroxide will yield aluminum oxide, in small amounts this should be quite doable in a porcelain crucible with a propane torch or bunsen burner. Aluminum oxide is also quite cheap from many pottery suppliers, only several dollars per pound if you want to go that route.

[Edited on 11-30-2015 by gdflp]

bluamine - 29-11-2015 at 15:34

Quote: Originally posted by aga

Buy some ?

Might be easier/cheaper.

Where can I find it?!

aga - 29-11-2015 at 15:36

Ebay !

http://www.ebay.co.uk/itm/white-fused-alumina-Optical-Glass-...

Cheaper:

http://www.ebay.co.uk/itm/Microdermabrasion-crystals-alumini...

[Edited on 29-11-2015 by aga]

deltaH - 29-11-2015 at 22:03

It also depends on how pure you want your product. If you're using lime, I'd be really worried about also forming calcium aluminate.

If you want a very pure material, then I would precipitate a saturated aluminium nitrate solution containing crushed ice with an ice cold 30% ammonia solution, the cold being employed purely to minimise the escape of noxious ammonia fumes. This would form aluminium hydroxide and ammonium nitrate. I'd then boil the slurry to break the aluminium hydroxide slimes and make powder before filtering (this doesn't always work). Then you will have aluminium hydroxide contaminated with a small amount of ammonium nitrate, however, a mild calcination carried out as a slow heating ramp up to 200°C outside will decompose the ammonium nitrate completely to gases only (nitrous oxide and water vapour), so ultimately you will be left with an extremely pure Al(OH)3 powder.

The result is a material known as gibbsite, Al(OH)3. When you bake gibbsite at higher temperature, it will lose water and convert into boehmite, AlOOH. This is also often called γ-alumina. It has a relatively high specific surface area and is used in the manufacture of catalysts.

If you heat boehmite in a furnace to very high temperature (>1000°C), you dehydrate the boehmite further and form α-alumina, Al2O3. It will no longer be microporous and will have lost its high specific surface area. In this state, it's used as an abrasive and to prepare refractories.

bluamine - 30-11-2015 at 00:52

Quote: Originally posted by aga

Ebay !

http://www.ebay.co.uk/itm/white-fused-alumina-Optical-Glass-...

Cheaper:

http://www.ebay.co.uk/itm/Microdermabrasion-crystals-alumini...

[Edited on 29-11-2015 by aga]

Thanks for the links but.. I can't buy anything online.. At least at the moment.

bluamine - 30-11-2015 at 00:58

Quote: Originally posted by deltaH

It also depends on how pure you want your product. If you're using lime, I'd be really worried about also forming calcium aluminate.

If you want a very pure material, then I would precipitate a saturated aluminium nitrate solution containing crushed ice with an ice cold 30% ammonia solution, the cold being employed purely to minimise the escape of noxious ammonia fumes. This would form aluminium hydroxide and ammonium nitrate. I'd then boil the slurry to break the aluminium hydroxide slimes and make powder before filtering (this doesn't always work). Then you will have aluminium hydroxide contaminated with a small amount of ammonium nitrate, however, a mild calcination carried out as a slow heating ramp up to 200°C outside will decompose the ammonium nitrate completely to gases only (nitrous oxide and water vapour), so ultimately you will be left with an extremely pure Al(OH)3 powder.

The result is a material known as gibbsite, Al(OH)3. When you bake gibbsite at higher temperature, it will lose water and convert into boehmite, AlOOH. This is also often called γ-alumina. It has a relatively high specific surface area and is used in the manufacture of catalysts.

If you heat boehmite in a furnace to very high temperature (>1000°C), you dehydrate the boehmite further and form α-alumina, Al2O3. It will no longer be microporous and will have lost its high specific surface area. In this state, it's used as an abrasive and to prepare refractories.

Thank you for all these important informations. I am looking for both α & γ alumina, & I have no choice, I have to use sodium hydroxide, because bicarbonate is more expensive.
Unfortunately, sand thermite method is the simplest, but it can't help to make something using alumina

deltaH - 30-11-2015 at 01:16

Sodium hydroxide will work fine as well, ammonia is only used when you want it ultrapure. I wouldn't use calcium hydroxide for the reason stated earlier.

Don't forget that boiling sometimes helps to break slimes in precipitations like this. It can make filtering easier if you can convert it to a powdery material first. If you get a powder that filters nicely, then you can easily wash out the sodium chloride.

MrHomeScientist - 30-11-2015 at 07:27

Quote: Originally posted by bluamine

Hi everyone!!
Can this reaction really happen:
3Ca(OH)2+Al2(SO4)3=2Al(OH)3+3CaSO4??!!
I found it on s YouTube video & I know that it is not a trustful science source, especially when it comes to a video existing only on an unknown channel

[Edited on 29-11-2015 by bluamine]

To address the original question, this reaction proceeds because of the high insolubility of calcium sulfate. While 3 of 4 of those compounds are pretty insoluble, the very high insolubility of CaSO4 drives the reaction forward. Because it involves lots of insolubles, I expect the reaction to be very slow and difficult to determine the endpoint.

Aluminum sulfate is reasonably soluble, though, so here's how I would do it: Dissolve Al2SO4 in water, then add enough solid Ca(OH)2 to make a slurry. Vigorously stir and heat for several hours, then filter the solids off. Make use of aluminum's amphoteric nature by adding enough NaOH to dissolve the produced aluminum hydroxide. Filter off the still-insoluble calcium sulfate, then take the solution and re-precipitate the aluminum as hydroxide by lowering the pH with some acid. Convert to alumina by calcining.

Edit: Of course now that I typed all that, I realized you could skip a lot of the intermediate steps by just making a solution of Al2SO4 slightly basic. This precipitates the hydroxide and you can filter and calcine from there. Not too basic, or the aluminum will re-dissolve!

[Edited on 11-30-2015 by MrHomeScientist]

gdflp - 30-11-2015 at 07:38

Calcium sulfate is more soluble in water than calcium hydroxide, (2.05g/L vs. 1.60g/L) and much more so than aluminum hydroxide. The reaction is driven by the low solubility of aluminum hydroxide, which is only about 1mg/L at STP; not that of CaSO4.

Technically this reaction could be cleanly done by fully dissolving the calcium hydroxide in water, mixing with a stoichiometric amount of aluminum sulfate solution, and allowing the precipitate to settle, but the volumes of water involved would be come so large as to be impractical. As long as the precipitate is dense enough, it might be worth a try in a 5 gallon bucket, though mechanical losses involved in filtering the precipitated aluminum hydroxide may kill the yield.

AJKOER - 1-12-2015 at 05:09

OK, Bluamine, so you wish to acquire a large quantity of Al2O3 (although you did not honestly reveal this at first citing a reaction creating Al(OH)3 ), possibly for a thermite. I assume large because in such amounts even NaHCO3 to quote yourself "because bicarbonate is more expensive" and not interested in working with cheap aqueous ammonia, which on a large scale would be problematic.

You have no credit (but claim to be around 23 years of age) or do not wish to leave any paper trail as to your identity/location, as you cannot buy online.

You known little chemistry, are a relatively recent members with 40 posts, and, I would guess, you are not a student, and your goal is to use chemistry for singular not specified reasons.

Excuse me, but I, for one, am suspicious of your intent, and would ask other members to be wary.

[Edited on 1-12-2015 by AJKOER]

MrHomeScientist - 1-12-2015 at 06:33

What possible clandestine use does alumina have?

Also thanks gdflp, I thought calcium sulfate was the nigh-insoluble one.

AJKOER - 1-12-2015 at 07:37

Someone with a background in chemistry could use alumina as a catalyst, as is, or mixed with other materials to act as one.

One may also think (mistakenly) it is employed directly in flash powder or thermites, or can be transformed to aluminum powder.

If I am correct in assuming it is needed/desired, even mistakenly, in large quantities, sorry but it makes me suspicious.

Not doubt, I could be in error and apologize again (perhaps getting paranoid).

[Edited on 1-12-2015 by AJKOER]

bluamine - 1-12-2015 at 11:56

Quote: Originally posted by AJKOER

OK, Bluamine, so you wish to acquire a large quantity of Al2O3 (although you did not honestly reveal this at first citing a reaction creating Al(OH)3 ), possibly for a thermite. I assume large because in such amounts even NaHCO3 to quote yourself "because bicarbonate is more expensive" and not interested in working with cheap aqueous ammonia, which on a large scale would be problematic.

You have no credit (but claim to be around 23 years of age) or do not wish to leave any paper trail as to your identity/location, as you cannot buy online.

You known little chemistry, are a relatively recent members with 40 posts, and, I would guess, you are not a student, and your goal is to use chemistry for singular not specified reasons.

Excuse me, but I, for one, am suspicious of your intent, and would ask other members to be wary.

[Edited on 1-12-2015 by AJKOER]

1.I supposed that the title is clear, so I don't have to mention that. Sorry for the inconvenient.
2.I know my questions seem pretty weird, & I was accused of trolling in another forum, but I have serious problems which makes many chemicals unavailable for me. Unfortunately, ammonia is one of them.. I am not sure if i can buy calcium cyanamide, if so I would use it to produce ammonia, because I need it to produce magnetite (I would discuss that in another topic).
3.I prefer using hydroxide because I wish to make it at home (1kg of sodium chloride costs here about .3 $ very cheap). On the other side, I can buy 15g of bicarbonate for about 0.05 $.
4.If I was looking for thermite, I would use hematite & aluminum or any other method (I mentioned sand thermite which I know how to make it) without using aluminum oxide.
5. I did not lie when I wrote that I am a 23 yo student, & I am using chemistry for both personal reasons & others reasons related of my study (though unfortunately my major is not chemistry).
6.Some members post here many much more dangerous topics (like RDX's one for example), & I did not read anything like this reply there.
7.I am just a guest here, if moderators believe that I must not post anything here, they have completely the right to ban me

[Edited on 1-12-2015 by bluamine]

[Edited on 1-12-2015 by bluamine]

bluamine - 1-12-2015 at 12:10

Quote: Originally posted by MrHomeScientist

What possible clandestine use does alumina have?

Also thanks gdflp, I thought calcium sulfate was the nigh-insoluble one.

I always loved your videos on YouTube, & this is the first time I see you here on this blessed forum. Mrhomescienrist allow me to talk to you a little in french: je vous tire Mon chapeau. Well I see all of you are asking about my purpose.. Here is my answer: I will use it mainly to make some acid resistant ceramic, because I will use some acids such as HF in some of my applications. I know some of you will say to me you can just buy it, believe me I can't find such a thing in stores.

MrHomeScientist - 1-12-2015 at 13:50

Thanks! Always glad to hear someone's seen my videos, and I appreciate the compliments. I actually haven't been in the lab in quite a while.

I'm not sure why AJ is so suspicious. I've never heard of alumina being used for anything nefarious.

Hopefully you're aware, but I feel obligated to give the warning anyway: HF is extremely dangerous stuff. Very small amounts spilled on the skin can and will kill you. Any noticeable effects may be delayed up to a few days, by which time it is too late to save your life. I highly, highly suggest avoiding this chemical in a home lab setting. If you feel compelled to use it, you must buy some calcium gluconate gel. In case of spills, immediately apply to the affected area then immediately go to the hospital anyway.
An acid-resistant ceramic might not be necessary; most uses of HF work perfectly well in plastic HDPE beakers.

bluamine - 1-12-2015 at 14:35

Quote: Originally posted by MrHomeScientist

Thanks! Always glad to hear someone's seen my videos, and I appreciate the compliments. I actually haven't been in the lab in quite a while.

I'm not sure why AJ is so suspicious. I've never heard of alumina being used for anything nefarious.

Hopefully you're aware, but I feel obligated to give the warning anyway: HF is extremely dangerous stuff. Very small amounts spilled on the skin can and will kill you. Any noticeable effects may be delayed up to a few days, by which time it is too late to save your life. I highly, highly suggest avoiding this chemical in a home lab setting. If you feel compelled to use it, you must buy some calcium gluconate gel. In case of spills, immediately apply to the affected area then immediately go to the hospital anyway.
An acid-resistant ceramic might not be necessary; most uses of HF work perfectly well in plastic HDPE beakers.

Thanks for advices. I do already know that HF is very dangerous, fluorine itsself is poisonous, but I thought before that a diluted sodium bicarbonate solution is enough to prevent any side effects..
HF is not the only acid I will use (HF is what I need to get calcium fluoride & magnesium fluoride to be clear to prevent any suspicious ideas can enter someone's brain) I have also nitric acid which can't even be stored in somethibg plastic made (though it doesn't react with silicon), in addition some reactions using such acids are endothermic or exothermic to a point that can make you fear about it (for example reaction between nitric acid & sugar to make oxalic acid). This is why I want to have an acid-resistant (& heat redistant) ceramic

[Edited on 1-12-2015 by bluamine]

[Edited on 1-12-2015 by bluamine]

AJKOER - 1-12-2015 at 15:36

Here is something that may be helpful, "The thermal decomposition of aluminium sulfate", by T.J. TruexR.H., Hammerle and R.A. Armstrong. To quote the abstract:

"The mechanism of thermal decomposition of aluminum sulfate has been investigated in the 500–700°C temperature range using a flow reactor system with the emitted gaseous sulfur oxides collected in a Goksøyr—Ross coil and a hydrogen peroxide impinger. Sulfur trioxide (SO3) was found to be the primary sulfur oxide released during thermal decomposition (1). Less than 3% of the released sulfur oxides were sulfur dioxide (SO2), indicating that the SO3 dissociation reaction (2) is slow relative to the residence time of the SO3 in the reactor (∼ 1 sec). The experimental technique should be readily adaptable to the study of the thermal decomposition of other metal sulfates."

Link: http://www.sciencedirect.com/science/article/pii/00406031778...

Also, of possible interest, "Thermal Decomposition of Aluminum Sulfate and Hafnium Sulfate", by H.A. Papazian, P.J. Pizzolato and R.R. Orrell, to quote from abstract:

"The thermal decomposition of aluminum sulfate and hafnium sulfate was studied in air by thermogravimetric analysis (TG) and in vacuum by simultaneous TG and evolved gas analysis (EGA). No SO3 was detected by the mass spectrometer. The primary products of decomposition appear to be SO and O2 for both sulfates. For aluminum I sulfate, the Arrhenius relationship shows two activation energies, whereas for hafnium sulfate there is only 'one activation energy in vacuum and two activation energies in air. X-ray data show the solid products of the reaction to be η-Al2O3 and HfO2. "
Link: http://www.sciencedirect.com/science/article/pii/S0040603172...

bluamine - 2-12-2015 at 03:32

Quote: Originally posted by AJKOER

Here is something that may be helpful, "The thermal decomposition of aluminium sulfate", by T.J. TruexR.H., Hammerle and R.A. Armstrong. To quote the abstract:

"The mechanism of thermal decomposition of aluminum sulfate has been investigated in the 500–700°C temperature range using a flow reactor system with the emitted gaseous sulfur oxides collected in a Goksøyr—Ross coil and a hydrogen peroxide impinger. Sulfur trioxide (SO3) was found to be the primary sulfur oxide released during thermal decomposition (1). Less than 3% of the released sulfur oxides were sulfur dioxide (SO2), indicating that the SO3 dissociation reaction (2) is slow relative to the residence time of the SO3 in the reactor (∼ 1 sec). The experimental technique should be readily adaptable to the study of the thermal decomposition of other metal sulfates."

Link: http://www.sciencedirect.com/science/article/pii/00406031778...

Also, of possible interest, "Thermal Decomposition of Aluminum Sulfate and Hafnium Sulfate", by H.A. Papazian, P.J. Pizzolato and R.R. Orrell, to quote from abstract:

"The thermal decomposition of aluminum sulfate and hafnium sulfate was studied in air by thermogravimetric analysis (TG) and in vacuum by simultaneous TG and evolved gas analysis (EGA). No SO3 was detected by the mass spectrometer. The primary products of decomposition appear to be SO and O2 for both sulfates. For aluminum I sulfate, the Arrhenius relationship shows two activation energies, whereas for hafnium sulfate there is only 'one activation energy in vacuum and two activation energies in air. X-ray data show the solid products of the reaction to be η-Al2O3 and HfO2. "
Link: http://www.sciencedirect.com/science/article/pii/S0040603172...

Should I heat it more if I need γ alumina?

AJKOER - 2-12-2015 at 05:56

Here is a link discussing forms of Al2O3:

http://www.researchgate.net/post/What_is_the_difference_betw...

[Edited on 2-12-2015 by AJKOER]

bluamine - 2-12-2015 at 09:33

Quote: Originally posted by AJKOER

Here is a link discussing forms of Al2O3:

http://www.researchgate.net/post/What_is_the_difference_betw...

[Edited on 2-12-2015 by AJKOER]

Unfortunately I guess all informations on the link above are about physical properties in the microscopic level. I read on Wikipedia that γ alumina is acid-resistant & α alumina is not. This is why I prefer γ alumina, but because I am lazy, I would choose α alumina for some other applications