Sciencemadness Discussion Board

Aluminum phosphate decomposition

Rattata2 - 18-8-2009 at 17:28

So I read on wiki that aluminum phosphate can be thermally decomposed to form aluminum oxide and phosphorus pentoxide (aka the anhydride of phosphoric acid.) I'd not heard of that before and it seems rather useful for creating other anhydrides and the like - however the wiki article states no references and I can't seem to find anything else anywhere on this.

"[Aluminum phosphate] is also used industrially as a high-temperature dehydrating agent. When strongly heated, aluminium phosphate decomposes into aluminium oxide and phosphorus pentoxide, the latter of which is very effective at absorbing water. Thus, aluminium phosphate can be used to dehydrate acetic acid through the following series of reactions:

4AlPO4 → 2Al2O3 + P4O10[citation needed]
6CH3COOH + P4O10 → 4H3PO4 + 6CH2CO
CH2CO + CH3COOH → (CH3CO)2O"
http://en.wikipedia.org/wiki/Aluminum_phosphate

Would this reaction actually take place and in what kind of yield? What temperatures and conditions are required to thermally decompose aluminum phosphate?

I can't seem to find any references. Would try it myself but I dnt have the means to atm :p I want to know if this is actually viable or if someone is just spreading mis-information on wiki.

not_important - 18-8-2009 at 18:29

It's used to make dimethyl ether by dehydrating methanol, the target application uses syngas and forms the methanol in situ with the dehydration catalyst.
Quote:
The preferred operating conditions of the process are a pressure range from about 200 psig to 2000 psig, more preferably from about 400 psig to about 1500 psig; a temperature range from about 200° C. to about 350° C.; and a space velocity in excess of 50 standard liters of synthesis gas per kilogram of catalyst per hour, more preferably in the range from about 1,000 to about 15,000 standard liters of synthesis gas per kilogram of catalyst per hour.


U.S. patent 5753716
and
http://www.osti.gov/bridge/product.biblio.jsp?osti_id=801223

For dehydrating acetic acid, the reaction is more like dehydrating it to ketene, which then reaxts with further acid to form the anhydride. Vapour phase at about 700 C as I recall.





kclo4 - 18-8-2009 at 18:32

I think it would happen, but I think you'd need some serious high temperatures, and have difficulty reaching them...
The phosphorous pentoxide would likley distill off (unless it forms some weird compound that quacks like a mixture of aluminum oxide, and phosphorous pentoxide but is not.) So you'd have to have a high temp still set up...

Polyphosphoric acid, as has been shown here is a decent dehydrator and I believe there has been success by heating phosphoric acid in a copper crucible aka a copper pipe cap thing. It has the ability to make SO3 from H2SO4 IIRC.


Edit: posted before I did, wow that isn't nearly as hot as I thought it would be! This might be useful..

[Edited on 19-8-2009 by kclo4]

JohnWW - 18-8-2009 at 21:38

Finely-divided amorphous Al2O3 has been used for the same purpose - high-temperature dehydration catalysis - e.g. in production of ketene from acetic anhydride. Also, AlPO4 is isoelectronic with SiO2, so I wonder how successfully finely-divided amorphous SiO2 has been tried for the same purpose; results on Google indicate that it can be done. However, Al2O3 and SiO2 (as well as AlPO4) are much less soluble and hygroscopic for water than is P2O5, alleged above to be the reason for the successful use of AlPO4 (with its decomposition) for the purpose; and it appears that they work by adsorption of the organic vapor on their surfaces.

Rattata2 - 19-8-2009 at 00:06

So...the operating conditions for it working as a dehydration catalyst in dimethyl ether synthesis are between 250-350C at increased pressure...but that wouldn't mean the compound decomposes at that temperature, otherwise it would be more of a reactant than a catalyst.

I'm more interested in whether this compound can be used to collect phosphorus pentoxide (or well, tetraphosphorus decoxide for the technical :p) and at what temperatures/pressures, and reaction conditions.

garage chemist - 19-8-2009 at 00:41

Just as a reminder: P2O5 and acetic acid don't give acetic anhydride. I have tried this several times, and other members of the german forum have, too. It has been discussed here as well.
It becomes black and gooey, some of the acetic acid can be distilled off, and the residue fails to give off anything upon stronger heating. Upon hydrolyzing the residue, the aqueous solution strongly foams when shaken.
Do not waste your time on P2O5 production in the futile hope of getting Ac2O.

P2O5 and H2SO4 do make SO3, but I think it's a waste of P2O5.

Polyphosphoric acid from heating H3PO4 in a crucible also make some SO3 when heated with H2SO4, but the yield is minuscule and absolutely not worth the work.

S.C. Wack - 19-8-2009 at 03:04

Use Ullmann's and Kirk-Othmer for industrial processes, and not unreferenced material...I assume there weren't references?

The melting point of aluminum (aluminium...it was written by someone outside of the US) phosphate as told by the wikipedia may be right though.

(edit: copy/paste from K-O: Aluminum phosphate [7784-30-7], AlPO4, is a highly insoluble, hard, and unreactive material with a high melting point (>1800C) which is used as a refractory material.)

There was some article published in French a while back and not worth saving where some was heated white hot in strong vacuum and it lost some weight.

[Edited on 19-8-2009 by S.C. Wack]

Rattata2 - 19-8-2009 at 13:30

Yeah I looked in Ullman's (2002, most recent I have) and didn't find anything about aluminum phosphate. Neither is included as a reference on that page.

Someone in an IRC channel mentioned that aluminum phosphate doesn't actually have an exact decomposition point, but rather it decomposes according to said formula above about 800C, and the more energy supplied the more easily it will decompose but that there's no definite point when it will start decomposing.

[Edited on 19-8-2009 by Rattata2]

Rattata2 - 13-9-2009 at 17:32

Update: The guy who wrote that bit included a reference but it says NOTHING. It's starting to bug the hell outta me.

And yeah I have more than one use planned for P2O5 if I am able to obtain it...not necessarily the production of A.A.

Edit: Aha, found this MSDS http://www.spipharma.com/downloads/Products/AntacidActives/S... which states that aluminum oxide as well as oxides of phosphorus may form.

[Edited on 14-9-2009 by Rattata2]

Magpie - 15-3-2010 at 20:45

I've made some AlPO4 precipitate from trisodium phosphate and aluminum sulfate. First I gave the ppt a good washing, then dried it in a 93C oven, then calcined it in a muffle furnace. The first batch I calcined at 1000C and got an 18% wt reduction. The 2nd batch, calcined at 600C, yielded a 11% wt reduction. Loss of 2H2O/molecule would be a 23% wt reduction.

The calcining at 1000C also resulted in a big volume loss and a hardening of the salt.

I can't find much on AlPO4 in my handbooks or the internet. Has anyone had similar experiences? What's going on here?

watson.fawkes - 16-3-2010 at 07:28

Quote: Originally posted by Magpie  
The calcining at 1000C also resulted in a big volume loss and a hardening of the salt.
The hardening also seems similar to what happens with dead-burnt lime. In that case, it's impurities that flux the CaO and make a glassy exterior coating, rendering it fairly unreactive. My guess (it's only that) is that you've got enough sodium in your precipitate to act as a flux. If a glass layer is forming, it would also explain why you're getting less-than-complete drying. The sodium could be in there by some double crystallization; see the Wikipedia article on Tarankite for a candidate.

It's also possible that you're simply sintering the AlPO4, getting densification instead of porosity. If this is happening on the surface before the interior, you could get the same sort of sealing effect.

I note with some amusement that the Wikipedia article on aluminum phosphate references this very thread.

Strepta - 16-3-2010 at 09:05

@Magpie: Apparently AlPO4 can exist as AlPO4.xH2O where x is less than three:

http://pubs.acs.org/doi/abs/10.1021/ja01652a016

This may be what you have (and I had as well) as I completed my second attempt at reduction of AlPO4 with C over the weekend. Some P4 was released but quite a lot of H2O as well at high temperature. The flames emerging from the exit gas beaker were more than likely PH3 as smoke was evolved. I retract my earlier statement about CO being the culprit.

Magpie - 16-3-2010 at 10:23

Thanks watson & Strepta for the information on AlPO4.
Quote: Originally posted by Strepta  

The flames emerging from the exit gas beaker were more than likely PH3 as smoke was evolved.

I agree that the burning gas was most likely phosphine. I had later come to that conclusion also. What else could burn so readily after emerging from a water bath.

blogfast25 - 16-3-2010 at 12:42

There's some indirect evidence that the dissociation of orthophosphates into their basic oxides and P2O5 (or P4O10) is considerably harder to achieve (higher ΔG to overcome) than for analogous dissociations of stable carbonates and sulphates I believe that's why silica seems to increase the yield of phosphate reductions by means of C or Al.

For instance, a stoichiometric mixture of CaSO4 and Al can be lit fairly easily and 'burns' at very high temperature (due to the formation of alumina). But an equivalent stoichiometric mix of Ca3(PO4)2 and Al could not be lit (at least not by me), despite the fact that the reduction of P2O5 with Al is expected to be highly exothermic.


[Edited on 16-3-2010 by blogfast25]

S.C. Wack - 16-3-2010 at 15:47

Sur les phosphates d'aluminium.

Attachment: compt_rend_234_1777_1952.pdf (199kB)
This file has been downloaded 1034 times

Magpie - 17-3-2010 at 11:22

"Study of kinetics and thermodynamics of the dehydration reaction of AlPO4 · H2O", has been kindly retrieved by jokull. This paper indicates that for AlPO4*H2O dehydration is complete at 573K (300C) and a transition from amorphous to crystalline occurs at 864K (591C). See http://www.sciencemadness.org/talk/viewthread.php?tid=4388&a...

My lab work indicated some shrinkage after 1 hour at 600C but the AlPO4 was still soft and easily powdered. But after 1 hour at 1000C it had increased in bulk density from 0.4g/mL to 1.3g/mL. And it was so hard I was concerned that it might damage my porcelain mortar & pestle.

I wonder which form (amorphous or crystalline) can be expected to perform better as a substrate for the formation of P via high temperature reaction with carbon?


watson.fawkes - 17-3-2010 at 16:34

Quote: Originally posted by Magpie  
I wonder which form (amorphous or crystalline) can be expected to perform better as a substrate for the formation of P via high temperature reaction with carbon?
I would imagine that the amorphous form would react better, because there's free surface energy remaining that hasn't gone to driving crystallization, a process which moves to minimize this energy term. Also, all other things being equal (such as granule size and shape), the specific surface area is inversely proportional to density, so reaction rate should be higher for the less compacted material.

Panache - 23-3-2010 at 21:59

Quote: Originally posted by watson.fawkes  
Also, all other things being equal (such as granule size and shape), the specific surface area is inversely proportional to density, so reaction rate should be higher for the less compacted material.


Unless your co-reactants are co-compacted, then the reverse is true. Intimacy of your reagants becomes important.

[Edited on 24-3-2010 by Panache]