Sciencemadness Discussion Board - Preparation of elemental phosphorus

Preparation of elemental phosphorus

Pages: 1 2 3 4 5 6 .. 15

garage chemist - 15-3-2008 at 11:00

What first paragraph do you mean?
When I click your link, I don't get any german text, just the google page on the book, the book itself seems to be unavailable.
Make sure you posted the correct link.

Magpie - 15-3-2008 at 12:26

GC: The link loads for me just fine. Are you giving it enough time to load?

jimmyboy - 15-3-2008 at 15:09

First paragraph

Nueu Method zur Darstellung von Phosphor

Vf. giebt eine historische Darst. der Entdeckung u. der verschiedenen Gewinnungsarten des Phosphors, welche mit einer eingehenderen Beschereibung des noch jetzt gultigen Verfahrens von NICOLAS-PELLETIER, schliefst, um sich dann der von ROSSEL und dem Vf. studierten Phoshordarst aus Aluminium u. Phosphaten zuzuwenden. Es wurde zunachst durch quantitative Verss. festgestellt, dafs die Einw. von Aluminium auf Natriummetaphosphat im Wasserstoffstrom bei der Temperatur der Geblaseflamme sich nach der Gleichung

6 NaPO3 + 15 Al = 6 NaAlO3 + Al5P3 + 2 Al2O3 + 3 P

vollzieht. Durch Zusatz von Kieselsaure gelingt es, die Gesmatmenge des Phosphors abzudestillieren

6 NaPO3 + 10 Al + 3 SiO2 = 3 Na2SiO3 + 5 Al2O3 + 6 P

garage chemist - 16-3-2008 at 07:02

New method for the preparation of phosphorus

(first sentence is irrelevant)
It was first determined by quantitative experiments that the action of aluminum upon sodium metaphosphate in a stream of hydrogen at the tempertaure of the blowtorch happens according to the equation
6 NaPO3 + 15 Al = 6 NaAlO3 + Al5P3 + 2 Al2O3 + 3 P.

By adding silica, the whole amount of phosphorus contained in the precursors can be distilled off:
6 NaPO3 + 10 Al + 3 SiO2 = 3 Na2SiO3 + 5 Al2O3 + 6 P

Please tell me how you were able to access the text of the book! To me, not even abstracts of the book are available!

[Edited on 16-3-2008 by garage chemist]

jimmyboy - 16-3-2008 at 07:55

Its a pdf link off google books

http://books.google.com/bkshp?hl=en&tab=wp

here is the link to the text again

http://books.google.com/books?id=uTEOAAAAYAAJ&pg=PA1015&...

if you use a few german chemistry terms you get a surprising amount of books on the subject

Do you have windows or maybe you are on a linux platform GC? You also need adobe acrobat as well -- either way if you still can't get it I will place it on rapidshare for you and anyone else who cannot access it

garage chemist - 17-3-2008 at 18:29

The link does not lead to any text.
Is it possible that google books differentiates between different countries so that e.g. users from germany can't access some books that users from other countries can access?

Try to get a different link, one that leads to a certain page in the book.
For example, this is a link that leads to a certain page in a book on google books:
http://books.google.com/books?id=VsdAAAAAIAAJ&printsec=t...
Try to get the link adress from right clicking and copying the link adress from a link that leads to the book text instead of copying it out of your browsers URL window.
The URL window in the browser often does not contain the actual adress of the page you are currently viewing, at least with such applications like google books.

Polverone - 17-3-2008 at 20:16

Quote:

Originally posted by garage chemist
The link does not lead to any text.
Is it possible that google books differentiates between different countries so that e.g. users from germany can't access some books that users from other countries can access?

According to this discussion, different countries do indeed have different access restrictions.

Quote:

At first, I rejoiced and welcomed that addition to the online record, but only until I tried to click on the link in his post. All my joy was quickly dashed, because no matter what I did, and how hard I tried, there was no link to the PDF version of the book, and no preview of any pages within the book. I asked friends in Belgium and the UK to try to click on Dave's link, and they were unsuccessful as well.

I corresponded with the "Internet Help Center" at "Google Books" when I first encountered troubles with viewing some of their books, and as it turns out, certain items are under copyright restrictions outside of the US. That means they are viewable only within the US. It's ironic that the digital version of the book of UK Patents, which was originally published in Britain, is blocked from appearing in the UK and its colonies, let alone the rest of the world.

For this reason it might be worthwhile for US members to download and rehost the PDF copies of interesting material from Google Books.

jimmyboy - 17-3-2008 at 21:42

Well since google books is filtered - I uploaded a copy of the book to rapidshare for anyone that can't access it

http://rapidshare.com/files/100388924/Chemisches_Zentralblat...

manimal - 30-4-2008 at 20:36

I would like to drudge up an idea from earlier in the thread regarding copper phosphide. The Handbook of Modern Chemistry by Tidy says:

"Cupric Phosphide (Cu2P2)

...Properties: A black substance decomposed by heat (3*Cu2P2 -> Cu6P2 + 2*P2)

...Formed as a black precipitate by passing phosphine through a solution of CuSO4."

Other pre-copyright chem books from google corroborate these statements.

That would be a pretty cool experiment if it worked. Maybe the Cu2P2 could be formed by throwing some sodium phosphide or other such thing into a solution of CuSO4., with the phosphine reacting w/ the CuSO4. It doesn't state the decomposition temp for copper phosphide though.

Hopefully it doesn't need arc-furnace tempuratures.

edit: Chemical Abstracts 1892 410 says "a dull, red heat".

[Edited on 30-4-2008 by manimal]

manimal - 21-5-2008 at 23:52

Would the triple calcium superphosphate of commerce be sufficiently free of sulfate to employ the aluminum reduction technique on it? It is formed be decomposing phosphate rock with phosphoric acid, so it shouldnt have much if any calcium sulfate, but you never know.

chloric1 - 22-5-2008 at 18:42

Quote:

Originally posted by manimal
Would the triple calcium superphosphate of commerce be sufficiently free of sulfate to employ the aluminum reduction technique on it? It is formed be decomposing phosphate rock with phosphoric acid, so it shouldnt have much if any calcium sulfate, but you never know.

Noway, it has been stated elsewhere to definately not use superphosphate specifically for this reason. Precipitating dibasic or monbasic calcium phosphate would be easy and inexpensive.

manimal - 22-5-2008 at 19:19

Quote:

Originally posted by chloric1
Noway, it has been stated elsewhere to definately not use superphosphate specifically for this reason. Precipitating dibasic or monbasic calcium phosphate would be easy and inexpensive.

It has been stated elsewhere to not use superphosphate, but I'm asking about triple superphosphate. Superphosphate is loaded with sulfate due to the manufacturing process which decomposes phosphate rock with sulfuric. TSP uses phosphoric acid instead. But the msds labels are reticent about their exact composition.

I'd like to do this reaction sometime this summer, and I have strong ideas on how to set up the necessary apparatus. I think I'll make a mini furnace out of cinderblocks, a heavy iron grate and other materials from my backyard, and borrow bromicacid's modification of the reaction vessel.

[Edited on 22-5-2008 by manimal]

LSD25 - 5-6-2008 at 11:11

I have had this recurring thought and it concerns me 'cos I don't see why it wouldn't work - I thought to myself, shit, I'll just post the idea and wait to get flamed, that'll fix the problem, so here goes....

If phosphoric acid is the water containing version of P2O5/P4O10 thus the oxidised variant of phosphorus (made, suprisingly enough, by oxidising phosphorus) and phosphine PH3 (aka phosphorus hydride), is the reduced/hydride of phosphorus, what happens if you bring the two into contact in the presence/absence of a catalyst? See it seems there is two possible routes to P, either via the reduction of phosphates or the oxidation of phosphines, the target product is in the middle of the two - so why not set out to try and achieve this by reducing one while oxidising the other in water (so the product can be collected)? Phosphine is known as a reducing agent (as are the higher P acids), so for mine this is probably not inherently unfeasible (particularly given that some are suggesting using hydrogen to reduce phosphoric acid and its salts).

Say for instance: 3 P4O10 + 20PH3 => 32P + 30H20, if one used a gas-proof gas line (I'd have to suggest welded/brazed metal/glass here) and then passed the uneacted gasses out of the vessel through either a strong oxidant or a flame, the acute toxicity issue would be dealt with (a flame would also provide a useful source of P4O10). Like I said, I don't know why it wouldn't work.

I do appreciate that surely, if this was going to work, industry would have thought of it - yup, that is likely, but look at what is likely to happen if industry looked at it - they would look at the need for phosphides (and the associated high cost & energy expenditure) and just go via the known route, which is not really a massive problem for industry.

PS I also have noted that industry heats the high P phosphates to drive off the excess P, what would it come off as?

garage chemist - 5-6-2008 at 11:18

Why would you want to use P2O5?
Phosphine is thermodynamically unstable. Simply leading it through a red-hot tube will make it decompose into phosphorus and hydrogen. I've said it multiple times, but people don't seem to listen...

Also, since water and phosphorus will react to phosphine and phosphoric acid at high temperature (Ullmanns says so, it's actually used as a process for phosphine production) your idea will probably only work in the opposite direction.

LSD25 - 5-6-2008 at 13:29

Got a ref or at least some more information than that? Glass, metal or quartz tube? I am more than happy to listen, but you've got to give more than that, although I did find these with a quick search:

http://www.freepatentsonline.com/4175111.html

http://www.jstor.org/pss/95965

http://www.jstor.org/pss/3572795

I found a couple of more articles which suggest that PH3 is adsorbed onto the surface of both Si & Cu, although how the fuck you take advantage of that is going to be interesting to read up on. An idea would be to see if the vapor deposition would grow on the hot P already on the surface of wall.

As to the reaction of phosphorus with hot water, I see no reason to use hot water - this is a simple reduction, heat shouldn't be necessary or if it is, it should be possible to keep the temperature low.

garage chemist - 5-6-2008 at 14:27

Glass or quartz. Anything that doesn't react with phosphorus- many metals form phosphides.

Have a look into a really big comprehensive book on inorganic chemistry- the internet will not be of use here.

Heating PH3 results in the splitting off of hydrogen to form solid, yellow lower phosphines. At higher temperatures, I am sure those will completely decompose into the elements.
As phosphorus has a boiling point of 280°C and you will be working at a much higher temperature, the P will condense as a liquid on the tube walls as the reaction gas exits the hot zone. Just like the unreacted 900°C sulfur vapor from my CS2 synthesis condensed as it left the tube furnace.

LSD25 - 9-6-2008 at 19:22

Grab this quick before the link dies: http://www.4shared.com/get/29938851/273f5647/Topics_In_Phosp...

LSD25 - 10-6-2008 at 00:41

Quote:

Originally posted by garage chemist
I have not done research into this, but I have heard of patents here that specify the microwave heating of a charcoal/phosphoric acid mix as a method to produce elemental phosphorus.

It would be a simple experiment to see if this is promising: wet some activated charcoal with 85% phosphoric acid, heat in an oven to drive off the water, fill the mix into a loosely stoppered test tube (glass wool or the like) and microwave it for several minutes at full power. See if the mix absorbs the microwaves and heats up to red or even yellow heat. Phosphorus production would be evident if the vapors catch fire upon meeting the air or the green glow is visible, or maybe even yellow drops of P condense on the glass wall.

I do not have my "research microwave" ready, I have dismantled it to make use of the transformer inside. But I still have all the parts needed to drive the magnetron, so I can put it together again when I have time (in spring or summer).

[Edited on 18-1-2007 by garage chemist]

I have already stated that monoammonium phosphate heated in a MW absorbs MW radiation to an incredible degree and from there is converted to polyphosphoric acid (the heat adsorbed is fucking incredible). I have tried to mix wood shavings with the monoammonium phosphate prior to this but every time I used the MW I'd end up by cracking borosillicate (which takes some heat and one hell of a gradient). I don't know if Len1's idea of using cheap-arse porcelain would work better, but it doesn't look like the wood reacts with the acid, then again, perhaps with more heat. Another alternative would of course be to try a reductant, dunno, anyone tried silicon/ferrosilicon?

I also added in a patent I found today on the oxidation of PH3 to hypophosphite using H2O2. What was the procedure used with the hypophosphite to precipitate WP? There is also the proposed P12H6 supposedly formed by treating calcium phosphide with hot, concentrated HCl - which according to Mellor gives P when heated to 70C in a stream of CO2.

http://lateralscience.co.uk/phos/index.html

(Some 4/5 of the way down the page).

Attachment: OxidationPH3toHypophosphite.pdf (463kB)
This file has been downloaded 974 times

Strepta - 15-6-2008 at 11:34

Method for Preparation of Small Amounts of P4.

The attached is a description of a technique I have used to prepare small amounts (< .5g at a time) of white (yellow) phosphorus from sodium hexametaphosphate (calgon). The chemistry for reduction of calgon using aluminum and silica was published by Franck more than a century ago and first mentioned near the beginning of this thread by Polverone 6 years ago.

6NaPO3 + 10AI + 3Si02 = 3Na2Si03 + 5Al203 + 3P2

The reaction, properly implemented, is self -sustaining and initiates at temperatures which can be obtained with a lab (Meker) burner and conducted in a test tube (18mmX150mm, pyrex), which can be re-used. See attached for more details and pictures.

[Edited on by Strepta]

Attachment: Method for Preparation of Small Amounts of P4.doc (286kB)
This file has been downloaded 1914 times

Magpie - 15-6-2008 at 12:44

Very nice, Strepta. And thanks for the nice write-up.

garage chemist - 15-6-2008 at 13:51

Well done!
Incidentally, your choice of raw materials is the same as the one I want to use in my own upcoming synthesis of phosphorus (maybe in about a month, no time for experimentation now).

As the phosphate-based calgon is unavailable here, I will make some sodium dihydrogen phosphate from phosphoric acid and Na2CO3 and calcine it to sodium trimetaphosphate at 550°C.

It's great to see that the reaction already works at bunsen burner temperatures! Though there is one issue: you are using "dark pyro" aluminum, which is difficult to get.
What I want to use is my 75 micron spherical aluminum (filler grade), the only kind that I can buy easily.
Making this react may require the heat of the tube furnace- I'll use a very similar setup as the one I used for the production of SO3 from NaHSO4.

A question: was your test tube attacked on the inside after the reaction? I am worrying about my quartz test tube. Molten sodium silicate on its walls would be very bad, it would cause devitrification, quickly ruining it if used at high temperatures again.
I am thinking about diluting the mix with something inert to "soak up" the molten silicate as it is produced- maybe MgO or Al2O3.
As I am going to use the furnace, it is unimportant for me whether the reaction is self-sustaining or not.

One last thing: I will be using diatomaceous earth (kieselguhr) as the SiO2 source. This really is less than pure, containing maybe 80% SiO2, and also a few percent of Fe2O3. I hope this won't cause any problems- don't want to lose much of the P to ferrosilicon formation!
But it is indeed very finely divided, which is why I purchased it for this purpose (and also as a filtering aid for flocculant and other finely divided suspensions that are hard to filter- it's used to filter yeast from alcoholic beverages, BTW.)

Again, very nice work! Len1 will surely like it!

[Edited on 15-6-2008 by garage chemist]

Strepta - 15-6-2008 at 15:22

Yes, garage chemist, other aluminums will work at higher onset temperatures. I did a number of experiments using a 12 g CO2 cartridge with the neck drilled out to 1/4" (6mm) with a glass tube inserted and sealed with furnace cement. This was led underwater for cooling and recovery. These were charged with other powdered aluminums and the cartridge was heated with a MAPP gas torch, but it was very difficult to keep everything hot enough to force the P out without clogging the outlet tube.

I have used my tube furnace for this as well, and it works just fine. Recovery of the P is more tedious with my furnace design because the quartz reaction tube is also hosting the nichrome heater, which is in turn wrapped with a kaowool blanket. Disassembly of the whole arrangement is necessary each time it is used in order to place the tube underwater, etc, etc.

As to damage to the borosilicate tubes, it's not too bad. I clean the tubes after use with HCl and they recover quite nicely unless too large a charge is used. I would certainly try borosilicate before risking my precious quartz. Light crazing of quartz can be removed by a quick rinse with HF or possibly with NH4HF2 solution.

Looking forward to more of that fine work of yours.

len1 - 16-6-2008 at 03:10

Yes indeed this is fantastc work. I had bought some Calgon a few months ago just for this reaction - good to see it works - I believe this is the first time it has been demonstrated so. A few thoughts

1) The reaction environment is heavily reducing - so it may work in steel - that would be a boon as quartz is precious

2) If the reaction initiates above 550C the Al is molten, in which case its initial state doesnt matter, provided theres good mixing - and so it may be possible to use Al shavings - which are almost free. Al powder is expensive and its purchase is watched due to its use in ammatol, terrorists etc

PS The simplicity of what has just been demonstrated makes a mockery of those who consider the banning of chemical elements consistent with living in a democracy.

[Edited on 16-6-2008 by len1]

LSD25 - 16-6-2008 at 04:04

Nice work Strepta

I converted that to a PDF for you

[Edited on 16-6-2008 by LSD25]

Attachment: Method1.MethodforPreparingSmallAmountsofP4.pdf (319kB)
This file has been downloaded 1259 times

not_important - 16-6-2008 at 05:44

This, and small scale reductions using carbon, are helped by using a few pieces of firebrick to form a small enclosure around the retort, using metal tubing to connect to the retort, and arranging things so some of the exhaust gases from the flame heat the end of that tubing connected to the retort - preheat that end before initiating the reaction.

Soaking the diatomaceous earth in hydrochloric acid for awhile, then rinsing well, can remove some to most of the iron and alkaline earths.

Boric acid can be used in place of silica

6NaPO3 + 10AI + 3B2O3 = 6NaBO2 + 5Al203 + 3P2

the boric acid melts at about the same temperature as the sodium metaphosphate, the sodium metaborate also has a slightly lower melting point than the silicates.

I did this as a recreational exercise decades ago, no pressing need for phosphorus so I didn't go for production data. Using a mix of Ca and Na metaphosphates with B2O3 and SiO2 resulting in some fairly low melting glasses and seemed to work OK with carbon as a reducing agent; I assume because the reaction mix was fairly fluid throughout the reaction giving better contact between all the reactants. On the other hand it is not self-heating.

jimmyboy - 16-6-2008 at 18:06

Nice writeup Strepta - no need for acetylene and steel retorts anymore

[Edited on 17-6-2008 by jimmyboy]

len1 - 16-6-2008 at 20:18

Its worse than I thought, Al powder is not just watched its restricted. (This is ridiculous, I remember a time it was sold in toy shops.)

LSD25 - 16-6-2008 at 20:49

Yeah, obviously some concern that suicide bombers were going to start using thermite

Got any nifty suggestions for powdering Al shavings/filings?

len1 - 16-6-2008 at 20:59

Quote:

Originally posted by LSD25
Yeah, obviously some concern that suicide bombers were going to start using thermite

Got any nifty suggestions for powdering Al shavings/filings?

As I posted, since the Al is actually liquid at reaction temperature, well mixed fine Al shavings should work. ( They wont work in ammatol so theres a relief our govt is protecting us)

PS That should be ammonal, ammatol is the amnitrate with diesel fuel I think

[Edited on 17-6-2008 by len1]

Magpie - 16-6-2008 at 21:29

Quote:

Its worse than I thought, Al powder is not just watched its restricted. (This is ridiculous, I remember a time it was sold in toy shops.)

This is somewhat off-topic but I have to ask: Why are so many chemicals banned/restricted in John Howard land? Was it the bombing in Bora Bora? Or what? Is the meth problem as bad there as in the US?

len1 - 16-6-2008 at 22:40

Quote:

This is somewhat off-topic but I have to ask: Why are so many chemicals banned/restricted in John Howard land? Was it the bombing in Bora Bora? Or what? Is the meth problem as bad there as in the US?

No not really. We are trailing the world in utilisation of chemicals (I wrote before about my attempts to get Ni here) so we are leading the world in the banning of chemicals. Because (honestly) we want to be leaders in something. Thats how the politicians/media have taught the public to think

We dont have much of a chemical industry -> we can ban chemicals willy nilly.

But it will lead to the above statement with the -> reversed. Its the chicken and egg problem

PS Were rid of Howard (he lost his seat in parliament and his party office) but it'll take a long time to reverse the sense of fear he used to govern

[Edited on 17-6-2008 by len1]

Polverone - 16-6-2008 at 22:54

Very nice work, Strepta. I'm a bit envious that you managed to take this to a photo-documented microscale preparation before I did, but I did squander 5 years since I first found mention of aluminum/metaphosphate reductions. Well done!

Garage chemist, I look forward to your upcoming variation on this preparation. Len1, I would be interested in seeing this tried with Al shavings. I have done it on a test-tube scale with chopped-up Al sheet and sodium hexametaphosphate (just watched the P vapor bubbles burst, no product recovery), and a reaction certainly takes place even at a moderate red heat. The reaction seems pretty slow, though. It may be difficult to get good mixing, even with molten Al, as the melt is pretty viscous. Maybe not_important's suggestion to form a mixture of metaphosphates to lower MP/increase fluidity would help things along. Can you get zinc dust, or is that restricted too? I imagine it would work fine at the temperatures you can reach in your electric furnace, and the finely divided zinc may be more tractable than coarser aluminum even if its intrinsic reducing power is not as great.

LSD25 - 16-6-2008 at 23:08

As to the banning of chemicals here, I'd be fucking suprised if it were possible to source zinc dust,

Although I just checked an auction site which shall remain unnamed (the one you can't buy a condenser on) and it seems that some of the more onerous restrictions are off glassware at last, so anything is possible

That said, I am seriously interested in the idea of using Al shavings - it is either that or build a small ball-mill and make the powder myself (the paint route will get expensive awful quick).

Would you get the hexacalcium metaphosphate from adding calcium carbonate to the monoammonium phosphate melt? It is of academic interest, I can get Calgon, but others cannot.

len1 - 16-6-2008 at 23:16

Thanks for that info Polverone, I think Ill have to think of ways to get good mixing in the molten mix. I have 100gms of Al powder left me for life so theres a great incentive to make the shavings method work, Im absolutely sure its possible.

Zn powder is not restricted (cold-galving, as its application in a paint base is called here is as Australian as BBQ's so I dont think theyll dare touch it) but are you sure it works? ZnO enthalpy of formation is only 350kJ/mol (for Al2O3 its 555kJ/mol oxygen) so it would have to be at a substantially higher T than with Al. But Zn boils at 910C odd, so on the surface it doesnt seem suitable. (I can see our future chemical curriculum: the alkali elements, the rare-earth elements, the banned elements ...)

[Edited on 17-6-2008 by len1]

Polverone - 17-6-2008 at 09:16

The Franck patent that the aluminum reduction method is based upon also mentions the use of zinc as a reducing agent. I have seen a mention much earlier, in the 1855 book Outlines of Chemical Analysis: Prepared for the Chemical Laboratory at Giessen By Heinrich Will, Daniel Breed, Lewis (Google Books) that metaphosphoric acid or metaphosphates will liberate phosphorus when heated before the blow-pipe with a bit of zinc.

I have verified without recovery (like the aluminum sheet experiment) that chopped bits of zinc will cause the liberation of phosphorus from hot fused metaphosphates. Even lead will do it. I think the critical point in going from demonstration to production isn't going to be the reducing power of the metal (within reason), but ensuring that the kinetics and conversion efficiency are optimized for production. Good mixing may strongly influence that (good mixing, or ensuring that oxides are fluxed away to keep exposing fresh metal to the melt).

Strepta - 17-6-2008 at 09:17

Quote:

Originally posted by len1
Its worse than I thought, Al powder is not just watched its restricted. (This is ridiculous, I remember a time it was sold in toy shops.)

len1: By now you've undoubtedly seen the post in the Al Powder Systhesis thread by LSD25. The paper linked provides a wealth of info about producing flake aluminum from Aluminum foil.

I don't suppose you can get magnesium if you can't get aluminum. If you could, you could try to 'cut' the reactivity of the Mg with either Zn or coarse Al. In the experiments I conducted with Mg the reaction mix and phosphorus product exploded out of the tt onto the side wall of my fume hood with a loud BANG but did not break the tt or destroy the setup in any way. This leads me to speculate that this reaction could be toned down fairly easily.

12AX7 - 17-6-2008 at 10:39

Mg and Zn can react in the gas phase at unusually low temperatures. Something to think about.

Tim

chloric1 - 17-6-2008 at 13:49

Quote:

Originally posted by Polverone
Very nice work, Strepta. I'm a bit envious that you managed to take this to a photo-documented microscale preparation before I did, but I did squander 5 years since I first found mention of aluminum/metaphosphate reductions. Well done!

Garage chemist, I look forward to your upcoming variation on this preparation. Len1, I would be interested in seeing this tried with Al shavings. I have done it on a test-tube scale with chopped-up Al sheet and sodium hexametaphosphate (just watched the P vapor bubbles burst, no product recovery), and a reaction certainly takes place even at a moderate red heat. The reaction seems pretty slow, though. It may be difficult to get good mixing, even with molten Al, as the melt is pretty viscous. Maybe not_important's suggestion to form a mixture of metaphosphates to lower MP/increase fluidity would help things along. Can you get zinc dust, or is that restricted too? I imagine it would work fine at the temperatures you can reach in your electric furnace, and the finely divided zinc may be more tractable than coarser aluminum even if its intrinsic reducing power is not as great.

Polverone- you talk about viscosity, what about a flux? Anyone ever melted pyrophosphates? They are less condensed and possibly they can at least be more fluid through the coarse of the reaction before condensing to higher phosphates. Potassium soaps are more liquid that sodium so maybe adding potassium pyrophospate might do the trick. I am shooting in the dark here.

Strepta - 2-7-2008 at 11:29

Tip of the hat to not_important; I tried boron trioxide (from boric acid by heating) in place of silica according to:
6NaPO3 + 10AI + 3B2O3 = 6NaBO2 + 5Al203 + 3P2 and had better yield in terms of less solid residue (ash) after the reaction.

Stoichiometry of the above equation calls for a ratio of calgon/Al/silica of 2.93/1.29/1. I mixed it accordingly and ground it thoroughly in a mortar. The B2O3 is quite hard after it cools and a bit of work is required to pulverize it. The end result is a mix which acts as a fluid, rocking back and forth if swayed and spurting to the top of the tube if the bottom is rapped sharply on a hard surface. Again I took 3 grams of this mix and heated it at the bottom of a pyrex test tube, the other end of which was wrapped with a damp piece of paper towel. CO2 was used a a protective atmosphere and the exhaust was routed through a half full 100 ml cylinder of H2O.

Heating of the tt was via a meker burner. The first noticeable difference in using B2O3 was that the edges of the mix began to shrink and curl before the reaction started, a result of the B2O3 starting to flow.

Once initiated the reaction was, again, self-sustaining, but noticeably slower -12 to 15 seconds- than the 3-4 seconds previously observed with silica.

Yield of P was .265 g from .512 available (identical result from 2 tries). The ash, however, was smaller and almost identical mass in both tries. The residual was .500 g less the initial mass (2.500 vs 3.000). If all the difference is attributed to released P, the efficiency of the reaction is ~98%. the missing 50% P could possibly be lost as P2O5 (quite a bit of gas escapes from the 100 ml cyl. during the reaction.

As not_important suggested, more could be done using a mix of borosilicates and metaphosphates to effect melting at even lower temperatures, if one is so inclined to experiment. I think I will concentrate on a better method of capturing and condensing the vaporous P before it escapes.

Panache - 6-7-2008 at 21:30

Quote:

Originally posted by len1
Its worse than I thought, Al powder is not just watched its restricted. (This is ridiculous, I remember a time it was sold in toy shops.)

Len i read this about a week ago and found it strange, i thought it was available in Melbourne but wanted to establish that definitely before i responded. Anyway, well, yes i have just come from
Barnes&Co, Swan St,Richmond, VIC, online www.barnes.com.au
They have 5kg, 1kg and 500g ultra fine mesh al powder for $59, $20, and $12 respectively and it is ultra ultrafine. They also have other metal powders made from i think a reduction process, let me remember, Sn, Cu, Brass (three kinds??), Iron, Al (ultrafine and course).

The store is a supplier to the resin casting industry and the metal powders make polyester resins and such look metallic.

I shop there for the disposable nitrile gloves and dry IPA. However i noticed today they have four types of organo metallic catalysts avaiable neat, aryl mercuric carboxylate (or slow curing polyurethanes, dibutyl tin dilaurate (for curing silicones), organo tin carboxylate and a platinum one for curing food grade silicones.

I directly asked about their knowledge of the legality of the Al powder, saying i had read in a forum it had become restricted because someone had perhaps made explosives from it. The lady (the daughter of Mr Barnes the stores founder) then said that sounds like crap as they sell it in paint stores, i left it at that with her, however i thought i'd let you know she thinks you're full of shit!

Anyway thought you might like to know.

len1 - 6-7-2008 at 22:38

Yes indeed, this is available! However I can assure you that what I wrote wasnt an off-the-cuff comment, I had actually been told by a reputable chemical company (not a centre stocking trade chemicals) that Al is restricted. Moreover the lady said it was restricted to such an extent that they would not sell it except on an account (after I told her I needed it for my hobby not my work). They sell a lot of chemicals with lower grades of restriction on the production of a drivers license and presentation of EUD - but not Al powder!

It appears now that our authorities are all over the shop with their newly introduced bans on chemicals.

As for paints, I had investigated that as a possible source after the above incident. There is no certainty so called aluminium colourede paints actually contain Al, or that is pure. What is worse, I weighed aluminium coloured paints and could find no noticeable density difference with non-metal paints (density of Al is nearly 3) which suggests there is little Al in them if any.

PS A further thought open for discussion. You have given the shops name, is that good? On the one hand I hate secretive types who dont reveal their 'sources' (and return them the favour by not giving them any help). On the other hand if this forum is watched (as it undoubtedly is at least in the US) these people might now try to bring 'its trade practices into line', as they have done with the chemical suppliers, to the detriment of non-account chemists here

[Edited on 7-7-2008 by len1]

Panache - 7-7-2008 at 18:01

I thought we all readily discussed suppliers and the like here, if because of that the police request a perfectly legitimate company from trading an item then that is fine in my opinion, a mistake had been made by the company. Most large Australian chemical suppliers have at least a permanent part time if not full time police member in their sales office and as such those places readily are updated to changes in the law. This is prudent i believe, as i think it prudent that there is some monitoring of this site, if only to slow morons with a bible or koran making something that goes bang in a public place.
I'm am truly shocked that Aluminum powder has gone past the point of the being restricted as per the Code of Practice and is now banned for all except account holders, that is quite extreme. Is liquid oxygen restricted? Imagine if a truck of liq O2 collided with a coking coal truck. It would be chaos because the drug squad would try and pull jurisdiction over the terrorist squad because coke was involved but really it was just a traffic accident.

[Edited on 7-7-2008 by Panache]

[Edited on 7-7-2008 by Panache]

len1 - 7-7-2008 at 19:11

Quote:

Most large Australian chemical suppliers have at least a permanent part time if not full time police member in their sales office and as such those places readily are updated to changes in the law.

I thought I could tell your jokes - but here Im not sure - please help me - youre not serious???

Quote:

This is prudent i believe, as i think it prudent that there is some monitoring of this site, if only to slow morons with a bible or koran making something that goes bang in a public place.

Well this has two implications. First as youd know the police's and jurisdiction marks in chemistry are not high enough to place them in the best position to differentiate between the former and you. On slow days, you might just do. Second Al powder can indeed be used to make bang (and it doesnt care about the place). So if thats the criterion, its the end of chemistry right there.

Or maybe I got you wrong and you wrote all the above to scare the undesirables away. To be honest I dont think there'll be a rush to the place you mentioned to buy Al powder

[Edited on 8-7-2008 by len1]

Panache - 8-7-2008 at 18:03

Quote:

Originally posted by len1

Quote:

Most large Australian chemical suppliers have at least a permanent part time if not full time police member in their sales office and as such those places readily are updated to changes in the law.

I thought I could tell your jokes - but here Im not sure - please help me - youre not serious???

Unfortunately this is absolutely true. Actually let me qualify, i know each largish company has a liason type officer who may have at most three companies they deal with so their presence is constant and frequent. However depending on the circumstance, a large company say like Sigma in Australia will have both the police person all the sales staff know about, they do the updating etc etc. They will also have a second undercover sales staff, who monitories staff. Only a few of the senior exec's know who this person is. That's all i know, i have heard more frightening things, but find them unbelievable.

Quote:

This is prudent i believe, as i think it prudent that there is some monitoring of this site, if only to slow morons with a bible or koran making something that goes bang in a public place.

Well this has two implications. First as youd know the police's and jurisdiction marks in chemistry are not high enough to place them in the best position to differentiate between the former and you. On slow days, you might just do. Second Al powder can indeed be used to make bang (and it doesnt care about the place). So if thats the criterion, its the end of chemistry right there.

Or maybe I got you wrong and you wrote all the above to scare the undesirables away. To be honest I dont think there'll be a rush to the place you mentioned to buy Al powder

[Edited on 8-7-2008 by len1]

I guess i was over-summarising my position on the issue above as this was kind of an aside in an already very long thread and as such what i said sounded a little oversimplified. I will however defer my actual position on this matter until a more appropriate opportunity arises.
This thread is about P, which i have nothing to add to given the only time i needed 2g's of red phosporous i spent four hours scratching off the 'matchheads' only to be told (luckily i asked) by Sauron that the red P was in the strikers!! Ha. Whats in the matchheads anyway, they are red, its so confussing.

497 - 8-7-2008 at 19:02

The red is just dye

IIRC they're mostly KClO3, glue, and ground glass for sensitizing. I'm probably forgetting something else.

Edit: I'm confused as to why it would be good to use a mix of Ca and Na metaphosphates. In US patents 4725368 it states that sodium-calcium metaphosphate (1:1 ratio) melts around 750*C. That's much higher than 550*C stated for straight sodium metaphosphate. Calcium metaphosphate melts above 1000*C IIRC. Maybe when mixed at a different ratio they for a eutectic?

I did find another patent US 5529960 that has a bunch of information on low melting metaphosphate glasses, some with transition temperatures as low as 230*C. I don't know how that compares to "melting point" though since they are a glass after all.

The simplest and easiest one I have seen so far is (in molar %):

46% P2O5
4% Al2O3
25% K2O
25% Na2O

It has a transition temperature of 260*C, I don't know how much of an effect this would have on the reaction, but I can't see how it could hurt. It might require an initial melting of 1000*C though...

Now if this was combined with a lower melting reducing agent it might react at a much lower temperature. Tin comes to mind, being amazingly low melting and more reducing than lead. An added bonus is it's very easy to get as solder.

Also are there any other options besides B2O3 and SiO2?

Second edit:

I don't know if this would work or not, but there are some very low melting oxide eutectics. US patent 4743302 and 5145803 contain some good data. There are some with a softening point of down to 225*C.

The simplest and most OTC one I have found yet consists of about 87% PbO and 13% B2O3 and has a softening temperature of around 330*C and melting at 450*C. If I'm thinking right it should form a sodium plumbite/borate mix after the reaction. PbO melts at 888*C.

For the absolute lowest temperature a V2O5/PbO eutectic can be used, but V2O5 is not too easy to get, and I don't think lower softening temps would help that much anyway.

So if all goes according to planed (unlikely) this should be the reaction for the majority of components:

4Na(K)PO3 + 6Sn + 2PbO ---> 2Na(K)2PbO2 + 6SnO2 + P4

Should proceed at 300-400*C. Maybe it could work?

[Edited on 9-7-2008 by 497]

[Edited on 9-7-2008 by 497]

DJF90 - 9-7-2008 at 15:06

Using solder would actually be a great idea... Not only would it melt at a sufficiently low temperature (lower than lead or tim on its own IIRC?), there are two reducing agents present (it was noted earlier that the reaction would run with lead). Solder is fairly cheap too, and can be bought "fluxed" or "flux-free", depending on whether a flux would be needed and/or the purity of the end product?

Harvesting Aluminum from Paint

bigbigbeaker - 9-7-2008 at 15:59

len1 - See Richard Nakka's website for details. I used RustOleum aluminum enamel and some cheaper brands when I was a kid. The AL payoff is good especially if you don't have another source. Nakka says " Fortunately, in the true spirit of experimental rocketry, a source of flake aluminum is as close as your nearest Home Depot or hardware store. Aluminum paint is largely a mixture of linseed oil and metallic aluminum flake. According to various MSDS's, such paint contains 20%-30% aluminum, by mass." The link is http://www.nakka-rocketry.net/igniter.html#Aluminum. Here in the US we can buy all the AL powder we want off Ebay or from a pyro supplier. The only restriction is that most suppliers won't sell AL in the same order with potassium perchlorate because flash powder is the intended product. And there is a limit on suppliers selling more than two pounds per year to any one customer. Since I'm not a flash powder maker, I don't care except that restrictions always get tighter as the years pass. They are never relaxed. The sodium and phosphorus guides are great. Fine work, great pics, totally enjoyable reading. Thank you.

JohnWW - 10-7-2008 at 19:36

Illegal to sell aluminium powder in Oz, and the Pigs actually try to enforce that?? (BTW The Australian Pigs are as corrupt as hell, especially the New South Wales, Queensland, Victoria, and Federal Pigs, while the West Australian Pigs are more noted for their brutality, so you could probably bribe them). That is utterly pointless, because all you have to do to make your own is to get some offcuts of aluminium joinery extrusions (which would otherwise be sent for recycling by melting-down at a foundry) from any maker of aluminium windows and doors (of whom there are hundreds around), and grind them into powder by holding them onto the face of an electric grinder or grinding wheel. Or, you could buy aluminium ingots from a recycler, or from the aluminium smelter at Tiwai Point, near Invercargill, New Zealand, and grind those down. (I think there may also be an aluminium smelter in Tasmania).

[Edited on 11-7-08 by JohnWW]

497 - 19-7-2008 at 01:14

So today I went to the store in search of some NaPO3. The only Calgon they had was non phosphate crap, not surprising. They did have another brand "White King" water softener that stated Na2CO3 and Sodium tripolyphosphate as ingredients. I bought it, for $5 I figured I couldn't go wrong. I haven't been able to find an MSDS for it so I don't really know what all is in it. It consists of a mixture of clear crystals a bit larger than table salt and opaque somewhat amorphous granules of variable size. The crystals are much more soluble and I determined that they are the at least substantially Na2CO3. I purified about a kilo of it by simply adding minimum amount of water to dissolve most of the crystals and then filtering out the undissolved suspected tripolyphosphate. It seems to work well.

My question is will the tripolyphosphate work? If it won't, how can I convert it to NaPO3? It melts at around 600*C and decomposes at above 1000*C according to MSDSs. Not sure what it decomposes to. It is said to eventually hydrolyze to orthophosphate, but I don't have any specifics on that. If possible I would also like to make half of it potassium rather than sodium, it forms a eutectic. I have yet to work out a way to do the efficiently.

I built a phosphorus "retort" out of a piece of 4130 thin wall steel tube welded closed on one end. It has a pipe thread attachment welded to the other end. I took a pipe cap and welded on a fitting that attaches to automotive brake line. I bent a piece of said brake line to about the right shape to bubble into water, it is ready to run AFAIK. All steel, it should hold up hot phosphorus vapor and such shouldn't it? It has a pretty large capacity, about 350 cc. I hope the propane torch will handle that. Always have acetylene as backup though.

Also I have a box of borax that needs to be made into B2O3, not sure the best way to do that. Maybe add nitric acid since sodium nitrate is so soluble and boric acid is not?

One final problem is the reducing agent. I don't really have access to fine aluminum. Carbon works, and since my setup should handle the temperature, maybe that would be best. I like the idea of using solder because it melts so low, but I'm not sure if it will work. And I'm not sure how to easily grind it into a fine powder. Cooling it and trying to make it form alpha tin would be nice, but the 5% antimony makes that hard to do. Maybe with dry ice...

[Edited on 19-7-2008 by 497]

Sauron - 6-11-2008 at 06:15

In the course of my investigations into preparing PCl3 I was taking a little detour to P2O3 (P4O6) and ran into an interesting paper by Krafft and coworkers in Ber. 34 565 (1901).

According to Krafft, the trioxide, trisulphide and trichloride of phosphorus (P2)3, P2S3 and PCl3) can all be reduced by metallic As or Sb to red P and the corresponding compounds of As or Sb.

P2O3 is far from readily available, but P2S3 is not. It is used in the match industry. PCl3 is hard to come by and hard to prepare but PBr3 is commercially available. Antimony metal is not hard to come by. So the reaction of P2S3 or PCl3 with Sb per Krafft may be of preparative value.

A while ago I posted the Brauer prep of vermillion P from PBr3 and Hg. That one produces an impure red P containing varying amounts of H6 bromides difficult to completely remove.

And recently in another thread I posted Michaelis' prep of red P from H3PO3 and PCl3 which I believe can be extended to PBr3. Ann., 325, 365 (1902). The product is 93% red P balance Cl.

I just wanted to call these reactions to your attention.

garage chemist - 3-1-2009 at 11:52

I recently prepared a little phosphorus in a simple proof-of-concept experiment.

Sodium trimetaphosphate (made by very slow heating of monosodium phosphate to 500°C in an oven, yielding a porous, almost unsintered mass and losing the theoretical amount of weight) was ground together in a mortar with a stochiometric amount of aluminum powder (75 micron, spherical, cheap epoxy filler grade that is of little use for pyrotechnics) and a stochiometric amount (corresponding to an assumed SiO2 content of 70%) of kieselgur (diatomite).
The thoroughly homogenized mixture was stored in a tightly capped bottle.

A 7cm long, 12mm OD piece of quartz tubing was melted shut at one end (using an oxy-propane welding torch), a small amount (ca. 1cm high) of the mixture added and the opening plugged tightly with glass wool.

The assembly was clamped at a slight upwards angle and the mixture heated strongly from all sides with the welding torch. At a bright yellow heat (1200-1300°C from the looks of it), the mixture sintered together and evolved large amounts of phosphorus vapor. An orange-red sublimate settled to the tube walls.

After cooling, the glass wool plug was removed. A greenish chemoluminescence could be observed in the dark.
By scraping the walls, some pieces of phosphorus could be removed, which promptly caught fire upon removal from the tube and burned with the characteristic yellow flame of phosphorus.

The porous, very hard glassy plug of slag could only be removed with much difficulty and also melted to the tube walls.
So the quartz tube is essentially a one-use item (it did cost less than 50 cent though), which is why I will not do this reaction on a larger scale in my quartz test tube.

So I can reasonably assume that the aluminum reduction of sodium trimetaphosphate with added kieselgur runs absolutely fine if the heat is strong enough, and the only thing I need to do now is developing suitable apparatus technology.
I am certain that quartz glass is a bad choice of material- once it has been contaminated by fluxes/alkalies, it slowly crystallizes all the way through like the supercooled melt it is, destroying it by the volume expansion during crystallization.

My next task is to manufacture a ceramic retort that fits into my tube furnace and can be connected to a glass tube that leads into a receiver.
Mullite-alumina ceramic, comprised of 3 parts alumina and 2 parts kaolin, made into a slip and cast into plaster molds, shall be the material I will work with over the next months.
Anon from the ABYMC metalcasting forums has developed and popularized this ceramic material for use as a crucible, due to its excellent reistance against thermal shock. The added alumina content also gives it a high flux resistance.
It normally has to be fired to 1500-1600°C due to the abscence of feldspar, but prolonged holding time at 1300°C could work as well, so I could fire it in the same tube furnace it will later be used in.

[Edited on 3-1-2009 by garage chemist]

Magpie - 3-1-2009 at 12:44

Very nice progress garage chemist. As usual your careful research and preparation are yielding fruitful results. I'm sure you will soon have developed a practical lab scale synthesis.

Do you anticipate the need for an inert carrier gas, or do you think that a simple retort design will suffice?

Readman-Parker process

Gamal - 8-3-2009 at 16:25

I was looking around for an elextrolysis process for making phosphorus, and found the Readman-Parker process.
It's described at

http://chestofbooks.com/crafts/scientific-american/sup7/The-...

It's a two state Process:
"...by dispensing altogether with the use of sulphuric acid for decomposing the phosphate of lime which forms the raw material of the phosphorus manufacturer, and also with the employment of fire clay retorts for distilling the desiccated mixture of phosphoric acid and carbon which usually forms the second stage of the operation."
I think it could be used in the lab as it won't need the high temperatures as the usual method, and the anode material is graphite.
I was trying to find more info about the process, without luck. If anyone know more than mee...

I work at an aluminum smelter, in Sweden, so I know a little about production by electrolysis in the industry.

Magpie - 8-3-2009 at 17:15

Gamal I read the SA paper. This is not an electrolytic process but an electric furnace process. Ie, the electricity is just used to produce heat through ohmic resistance. As far as I know this is the standard industrial process that has been used worldwide for some time. Nice try though.

If I come to Sweden will you give me a tour of your aluminum smelter? I love industrial processes.

[Edited on 8-3-2009 by Magpie]

Gamal - 8-3-2009 at 17:36

You are welcome Magpie! Just tell me when.

I think you are right about your comment. I now understand why I didn't find anything on electrolysis when I was searching for manufacture of phosphorus. Sad but true!

Globey - 9-3-2009 at 07:03

This is awesome work! What would be nice, is if someone made (and marketed) miniaturized versions of the Downs cell, a minerature ammonia factory (using only air, water, and initial activation energy), and a minerture liquid air factory. I'm envisioning all these made out of the same type of rugged/resistant types of steel the real plants are made out of, except instead of weighing in at tons, each of these plants would be perhaps a foot tall. You could plug in your Downs cell, lead a tube down the drain for your chlorine (or use it for other purposes), add some salt, plug in, and viola.

I am contemplating buying a lathe and trying this myself, but it's not easy, and will be very time consuming. Now, I realize this is wrong thread, and that people have made sodium (Castner), but my interest is limited in scope to OTC availability. NaCLO3 is another which can be made 100% OTC. That is my interest, small versions of plants using only 100% OTC. Awesome work on the P BTW, truly awesome!!! =) It was a crime to ever make P a controlled "chemical"...it's elemental, my dear Watson!

kilowatt - 28-4-2009 at 16:20

I need about 50-100g of white phosphorus (unless there is another way to prepare sodium hypophosphite that I'm not aware of). I am yet unsure what to do with the phosphine that is liberated by this reaction but I suppose I will need to find a way to neutralize or destroy it. This will need to be prepared from bone ash as it is the main economical phosphate source in my case and which I have plenty of. However this thread has become so lengthy and bloated that I have found it nearly impossible to summarize the useful information and figure out what procedures people here have actually tried successfully.

So, my understanding is that the metaphosphate is required for the direct production of phosphorus by distillative reduction. The violently exothermic aluminum reductions of metaphosphate are largely useless to me; these are uncontrollable, cannot be contained, and easily produce lethal amounts of phosphine. A carbon reduction is much more preferable, for example:

2Ca(PO3)2 + SiO2 + 10C --> Ca2SiO4 + 10CO + 4P

My understanding is that this can only be carried out in a practically one-use retort made of high temperature ceramics or fused silica. I see that BromicAcid constructed a steel retort, but does not mention what mix was used in that or how severely the walls were phosphided. Was it used with a formula such as this?

I am also unclear on whether the reduction to Ca3P2 can actually be performed with orthophosphate or if this too requires the metaphosphate. I have seen the (again practically useless) aluminum analog of the reaction stated here a few times with phosphate. Any ideas what reaction temperatures would be required for the carbon reduction to phosphide?

Ca3(PO4)2 + 8C --> Ca3P2 + 8CO

This reaction may be of interest because it would not require the lengthy conversion of ortho- to meta-phosphate, and would not emit the highly reactive gaseous phosphorus. I have seen some references to possible electrolytic separation of Ca3P2. Are there any "cold" electrolytes which might work for such a method?

Also someone mentioned in passing a fused salt/bone ash electrolysis route. Such fused salt electrolytic routes may be of particular interest to me since I am in the process of constructing a multi-purpose fused salt cell capable of extracting a variety of metals and gases via airtight traps from about 650°C to 1000°C. The cell can operate under inert gas and the traps are designed so that air cooling can be used to maintain a frozen layer of electrolyte on them, keeping their exterior surfaces electrically insulated. Of particular interest are CaCl2/CaF2 eutectic electrolytes. I am unsure how calcium phosphide would behave in such an system, but I have previously derived phosphoric acid from the electrolysis of bone ash in a small experimental version of such a cell. Since that was an open cell with no electrode traps, that is not very indicative of whether P2O5 itself was formed at the anode or whether phosphorus/phosphine was released as this would have immediately oxidized to P2O5. Even if phosphorus could be derived electrolytically from a fused salt cell, the reactivity of phosphorus with virtually all metals (stainless steel in this case) would be of concern.

Basically what I have available to me are any number of wet methods for the salt conversions, steel vessels for distillative reduction, and my fused salt cell.

[Edited on 29-4-2009 by kilowatt]

Formatik - 29-4-2009 at 15:09

You might be able to find more inormation on the phosphide and phosphate formations or reactions in Mellor or Gmelin.

garage chemist - 29-4-2009 at 22:52

Why would the aluminum reduction of metaphosphate produce phosphine? The reactants are (NaPO3)n, Al and SiO2, there's no hydrogen in there. Only if water is added or aerial moisture admitted can phosphine be formed.
The phosphide is just an intermediate product. Strong heating with SiO2 converts it to phosphorus.

Don't try to use carbon as the reductant, the temperatures required are in the region of 1500°C and aren't attainable in homebuilt furnaces. Stick with the aluminum method.

And IF you get phosphine, you can destroy it by leading it through a heated glass pipe. It's thermodynamically unstable and will decompose into the elements on heating.
You could put the phosphide slag from the Al/NaPO3 reaction on a flask, lead a stream of nitrogen (to displace oxygen) through it followed by a heated glass tube and slowly drip water onto the phosphide. P sublimate will deposit in the glass pipe.

kilowatt - 30-4-2009 at 15:47

Do you think I can get decent results using Ca(PO3)2 or would it be worth the trouble to use NaPO3 made from Na3PO4 with phosphoric acid which I could obtain by running bone ash in my hot cell? If Ca(PO3)2 can work nearly as well then that would definitely be the way to go here.

Can aluminum directly reduce Ca3(PO4)2 to Ca3P2 which could then be acidified in a controlled fashion and led through a quartz tube at yellow heat and then a cooler and into an argon filled collection flask? This method would seem cleaner than trying to run phosphorus in a metal apparatus. If I run such a setup as that I would like to have a low pressure relief valve in case anything should plug with phosphorus or other material, in which case a rapid chemical means of destroying phosphine would be desirable.

ferrophosphorus

ammonium isocyanate - 14-7-2009 at 00:03

Im in the USA and here we have ALOT of trouble aquiring reagents (or any lab stuff for that matter- in some states erlenmeyer flasks are illegal!).

One chemical that is impossible to buy as an individual in any reasonable quantity is phosphorus. And making it generally entails reducing phosphate rock at extreem temperatures and condensing the phosphorus vapours. The equipment required for this is certainly beyond my reach (and I tried with a basic distillation set-up, one that as a result, I no longer have).

I have been playing with the idea of using phosphides to make phosphorus, but they are pretty nasty and I havent found a workable procedure for this yet.

But I haved found a substitute, of sorts. It isn't perfect, but ferrophosphorus, an alloy of iron and phosphorus, works to produce many phosphorus compounds (including the very interesting and obscure dessicant/reagent, phosphorus pentasulfide). And, best of all, it can be produced from inorganic phosphates using a thermite mixture.

Procedure:
The following are ground and mixed in a suitable container: 16.28g Na3PO4, 5.89g powdered magnesium, 45.53g hydrated iron (III) oxide (rust), 10.85g powdered aluminium, 3.09g silicon dioxide (crushed glass). This mixture is poured onto a flat rock, several grams of powdered magnesium are added to the top, and a magnesium ribbon is stuck in as a fuse. Light the ribbon, and stand back to enjoy the show.

Notes/Findings:
After cooling, a hard outer shell of metal oxides is formed, with mostly porous ferrophosphorus on the inside. Be very careful around the product, as a side reaction generates magnesium phosphide, which gives off deadly phosphine and diphosphine fumes, resulting in a foul odor.
I used an excess of of thermite to ensure enough heat to drive the reaction and trap the phosphorus produced.
I used less SiO2 than required to convert all the Na2O to Na2SiO3, partly because I didn't want to contaminate the melt too much, partly because it was all the SiO2 I had on hand.

I haven't had a chance yet to do quantitative analysis on the product, but qualitative tests confirm the presence of phosphorus alloy.

I hope someone finds this useful.

By the way, does anyone know where to purchase sulfuryl chloride and ammonium fluoride? I have seen them listed among people's recent chemical orders, but haven't been able to find a supplier, and they're a real hassle to make.

*Any mods please feel free to move this to the phosphorus thread in general chemistry, I hadn't seen the thread until after I posted this.

[Edited on 14-7-2009 by ammonium isocyanate]

blogfast25 - 13-10-2009 at 08:53

Well, ammonium isocyanate, I haven't checked the stoichiometry of your mixture but I can assure you that SiO2 (and thus also Na2SiO3) is reduced by both Al and Mg to Si. So, BTW, is Na2O (see the NaOH/Na2CO3 'thermites' for 'chemical sodium' elsewhere on this forum) to Na.

Why do you feel certain that the ferrophosporus is trapped "inside"?

And what's the function of the Mg? It would appear to me that Al is capable of reducing P2O5 - I remember having made some calcs to that effect.

In my opinion, a simpler stoichiometric mix of Na3PO4, ferric oxide, Al powder and perhaps some Fluorite (CaF2) as a fluxing agent, should lead to ferrophosporus and alumina. Possibly even neatly separated.

What's the P-content you're aiming for in the ferrophosporus?

[Edited on 13-10-2009 by blogfast25]

12AX7 - 13-10-2009 at 11:03

FYI, FeP is the last phosphide compound / "intermetallic" formed in the Fe-P system. It has a high melting point, 1723K. Fe2P has a somewhat lower melting point, and Fe3P decomposes at 1431K (peritectic). The Fe-Fe3P eutectic consists of 16.18%at phosphorus, melting at 1325K.

Likely you'll need a mild oxidizer to displace phosphorus. Perhaps Fe(x)P can be displaced with sulfur, yielding P(g) and FeS or FeS2 (possibly liquid, both having lower melting points than any FeP). This may also be possible in aqueous solution, grinding it up and bubbling Cl2 through the mush.

Sodium probably has low solubility in iron, but most of it will go off as gas anyway. The same is true of hydrogen, if you use pyrophosphoric acid or ammonium pyrophosphate instead. Gasses are best avoided in a thermite reaction.

Calcium phosphate may be a better substrate, since calcium won't be reduced and it contains no gasses. The resulting slag will contain calcium aluminate, which is slaggy enough as it is without silica. I don't see any need for silica: with metallic iron around, I think iron phosphide will form instead of calcium phosphide.

I guess the question is, how much energy is there in forming (CaO)x.Al2O3 versus Ca3P2 versus Fe(x)P? Is it better to have three immiscible materials, Ca3P2, Fe(l) and Al2O3(l), or two, (CaO)x.Al2O3 and Fe(x)P? This is complicated by (CaO)x.Al2O3 being a glass, while effectively trying to reduce calcium with iron (even though the reaction doesn't produce Ca, Ca is always oxidized). What if Ca3P2 is soluble in Fe altogether?

An example reaction might be:
Fe2O3 + 2 Al ==> Al2O3 + 2 Fe
3 Ca3(PO4)2 + 10 Al ==> 9 CaO + 5 Al2O3 + 6 P
which might be mixed in proportions of CaO.Al2O3 giving:
3 Ca3(PO4)2 + 18 Al + 4 Fe2O3 ==> 9 CaO.Al2O3 + 4 FeP + 2 Fe2P
The resulting metallic(?) mixture has solidus 1534K and liquidus 1660K.

Tim

blogfast25 - 13-10-2009 at 11:30

Tim:

You wrote:

3 Ca3(PO4)2 + 18 Al + 4 Fe2O3 ==> 9 CaO.Al2O3 + 4 FeP + 2 Fe2P
The resulting metallic(?) mixture has solidus 1534K and liquidus 1660K.

... is thermodynamically by far the most likely outcome, IMHO. As you state, there is nothing in the mix that can reduce CaO, thus there is no reason to suspect calcium phosphide can form.

Calcium monoaluminate (CaO.Al2O3) is supposed to have a low melting point (about 1,100 C is a figure I've seen bandied around), so that would make the use of a slag fluidiser like CaF2 not necessary at first glance. But my own (very limited) experience with using CaO as a slagformer in thermite reactions (to promote formation of the fluid CaO.Al2O3) wasn't positive though...

I'd suggest a formulation comprised of Ca3(PO4)2, Fe2O3 (or Fe3O4) and Al in controlled ratios. Test with and without CaF2.

I'd also suggest the typical 'backyard thermite' set-up consisting of a smaller flowerpot with the thermite mixture in it, embedded in a larger flowerpot filled with dry sand. This will allow the molten mixture of reaction products to collect at the bottom of the inner flowerpot and allow the Fe-P and slag to separate from each other.

According to this source http://answers.yahoo.com/question/index?qid=20081020091007AA... the HoF for P2O5 ≈ -1530 kJ/mole, so the reaction:

P2O5 + 10/3 Al ---> 2 P + 5/3 Al2O3

... has a HoR of about a whopping - 1260 kJ/mole (of P2O5)! Assuming the estimated HoF of P2O5 is correct, together with the F2O3 thermite reaction this mixture should run very, very hot! Assuming the CaO behaves indeed inertly, it should act as a heat sink...

[Edited on 13-10-2009 by blogfast25]

12AX7 - 13-10-2009 at 19:59

Plenty hot indeed; should be plenty of room there to add silica or fluorite for whatever purpose.

As sand containers go, you can pour a bed layer of sand, place a paper tube on it, and fill inside with thermite and outside with sand. Go a layer at a time, packing lightly then withdrawing the tube, so it doesn't get stuck. Or leave it in (just a rolled up sheet of paper isn't much), it'll outgas easily into the surrounding sand.

Tim

blogfast25 - 14-10-2009 at 05:59

In fact, the system could be seen as a Classic Thermite (TM) but heat boosted with phosphate: 1.5 mole of P2O5 per mole of Fe2O3! It should be hoped the 9/4 mole of CaO (per mole of Fe2O3) slurps up enough reaction enthalpy, otherwise just about everything might just be blown off and one might just end up with an empty crucible... Isocyanate's mention of "porous" ferrophosphorus may already be an indication of things running too hot...

Myfanwy - 19-11-2009 at 12:58

Sorry i don`t have the time to read the whole 15 pages. Did anyone tried the Method with Calcium dihydrogenphosphate?

Ca(H2PO4)2 is available in Germany as "Superphosphat".
You can make it by the reaction of Calcium phosphate with sulfuric acid.

Ca3(PO4)2 + 2H2SO4 -> Ca(H2PO4)2 + 2CaSO4

The Calcium dihydrogenphosphate decomposes at 200°C to Water and Calcium metaphosphate.

Ca(H2PO4)2 -> 2H2O + Ca(PO3)2

The Calcium metaphosphate can reduced with Carbon releasing carbon monoxide and P4 Fumes.The P4 is led in Water.
2Ca(PO3)2 + 5C -> 2P + 5CO + Ca2P2O7

I dont now at which temperature this reaction is done, but i think at least 800°C

Taoiseach - 20-11-2009 at 00:38

The chemistry is all figured out already. Dozens of books out there (some available 4free @archive.org) which cover the ancient P production methods. Besides it is pretty much irrelevant what sort of phosphorous matter you ignite with C (and possibly SiO2) - they all yield elemental P. Even urine does.

If you want to contribute something useful to the discussion, go find a suitable reaction vessel. Not only does it need to withstand the extremely reactive P but also it has to withstand the extreme temperature gradient of heating one end to bright redness while dipping the other end into cold water. There's not a whole lot of materials to fulfill these requirements. Ancient chemists used clay retorts, but these are not easily improvised. Definetly not from OTC ingredients.

len1 - 20-11-2009 at 03:42

Well he has obviously got home something dozens of books couldn't. Its not true that any phosphorous containing salt will do. Carbon will not reduce ordinary phosphate. One needs stoichiometry richer in P2O5 than that - which the metaphosphate provides as can be seen by the equivalence

3Ca(PO3)2 -> Ca3(PO4)2 + P4O10

Magpie - 20-11-2009 at 09:11

If you can find sodium hexametaphosphate, (NaPO3)6, you will have your metaphosphate specie directly. This used to be sold OTC as Calgon. It may still be available in some countries.

Len1, why do you say that orthophosphates can't be reduced with carbon? Isn't this the common industrial method, ie,:

Ca3(PO4)2 + SiO2 +5C = 3CaSiO3 + 2P + 5CO

blogfast25 - 20-11-2009 at 10:47

Why not use dried and calcined chicken carcasses (or other bones?) They contain phospate, the exact type of phosphate matters little. Reduce with C or Al (see above)...

len1 - 20-11-2009 at 19:37

Magpie, carbon will not reduce ordinary phosphate. That is the reason for the manipulations of the Scheele process which is run at bright red heat (1000-1300C), and it is indeed what the german poster described. The process you have written uses sand which converts the phosphate to silicate with the phosphorus being reduced by carbon. But it is a much higher temp process conducted at white heat and neads an electric arc furnace.

S.C. Wack - 21-11-2009 at 08:51

Quote: Originally posted by len1

carbon will not reduce ordinary phosphate

Ca3(PO4)2 and coke:
US2860037

Mercury and lead phosphates are also known to give phosphorus directly.

It is my understanding that the very old methods that don't have Si present mostly used phosphoric acid instead of a salt of it, and that the methods with silica and apatites are using the SiO2 as a flux. Ullmann's: "Reaction Mechanism. Phosphate and quartzite form a melt in the furnace at the reaction temperature of 1400–1500C, and the carbon reacts with the ionic phosphate–silicate–fluoride melt according to
2 PO4(-3) + Si6O15(-6) +5 C → P2 + 5 CO + 2 Si3O9(-6)
The phosphate ion is reduced to gaseous P2, transferring its charge to the polymeric silicate ions, whose degree of polymerization is thereby reduced."

Skip the lame references and go right to the better literature, such as Mellor's P chapter:
http://www.sciencemadness.org/library/index.html

and J.B. Readman's detailed entry for phosphorus in Thorpe:
http://books.google.com/books?id=5nnPAAAAMAAJ&pg=PA195

[Edited on 21-11-2009 by S.C. Wack]

len1 - 21-11-2009 at 13:12

Im sure silica actually participates in the reaction, its not just a flux, as per the formula, which is ubiquitous. As for the patent Im afraid the patent world is full of bullshit - but thats nothing new if one regularly visits the world of practical chemistry. In this instance you can also deduce it by thinking about it, why would people go through the complicated procedure of making the metaphosphate if the ordinary phosphate works? To disprove me Im afraid you'll have to do much more than quote a patent - go into the lab and make the phosphrus.

blogfast25 - 21-11-2009 at 14:12

The description of the discovery of Phosphorus from phosphates in concentrated urine (see Wiki, down right now) doesn't feature silica either.

The high temperature reduction of phosphates with carbon must be possible due to the volatility of both reaction products, P4 and CO.

S.C. Wack - 21-11-2009 at 14:38

Quote: Originally posted by len1

Im sure silica actually participates in the reaction

Unfortunately it appears that no one else is as sure as you are, according to Kirk-Othmer, Gmelin's, and Ullmann's, so please excuse me for sticking with the literature.

Why should I have to build an arc furnace to prove Hermann Hilbert, Albert Frank, Jonas Kamlet, and Nikodem Caro, PhD's all, full of shit, because you don't believe what they say, even though no one has any reason to doubt them?

BTW the phosphates of iron and aluminum are also known to be reduced by carbon, in addition to the mercury and lead, long ago.

len1 - 21-11-2009 at 16:19

No its your interpretation of what you have read that I dont take for gospel- why should I, as you havent actually done anything.

If all you do is read stuff without ever trying to do any of it, not only are you a receptacle for a lot of rubbish as well as all the good stuff, you have no ability to be able to differentiate amongst it. Thats something I commented to Sauron about as well. You at least do not write your posts as if youre about to actually do something practical.

By the way I should add the results are in now - several thousand posts of what he was going to do - he's gone, none of it done. It was all bullshit.

[Edited on 22-11-2009 by len1]

DJF90 - 21-11-2009 at 16:45

Len1 - I think you'll find a chemists eyesight is a necessary peripheral; without such, many things cannot be done. I have not heard from Sauron from a long time. Short of kicking the bucket, this is the only factor I can think of as to why he has not replied to emails/posted here etc, especially as his eyesight was failing when he was last active here. Even other friends of his have not heard from him. If anyone has been in contact recently, please U2U me.

entropy51 - 21-11-2009 at 18:16

Quote: Originally posted by len1

No its your interpretation of what you have read that I dont take for gospel- why should I, as you havent actually done anything.

If all you do is read stuff without ever trying to do any of it, not only are you a receptacle for a lot of rubbish as well as all the good stuff, you have no ability to be able to differentiate amongst it. Thats something I commented to Sauron about as well. You at least do not write your posts as if youre about to actually do something practical.

By the way I should add the results are in now - several thousand posts of what he was going to do - he's gone, none of it done. It was all bullshit.

[Edited on 22-11-2009 by len1]

It must be such a burden to have to deal with those of us who believe that most of the chemical literature is not bullshit. Sounds like the voice of someone who's not published much himself, else he would know that peer review, though imperfect, does tend to filter out the rubbish.

Quote: Originally posted by len1

As for the patent Im afraid the patent world is full of bullshit - but thats nothing new if one regularly visits the world of practical chemistry.

The literature is bullshit. Patents are bullshit. What are we to do? Whom are we to believe? Luckily we have len1 to show us the truth!

[Edited on 22-11-2009 by entropy51]

[Edited on 22-11-2009 by entropy51]

len1 - 21-11-2009 at 18:28

Anyonw who says the above has just not tried to do much in the real world. Lost in the clouds.

DJF90 - well in that case he could have written Im sorry I can not do this or that, can someone check it out for me. That would have earned much more respect. When you start lying, youre finished as a scientist.

PS I thought Ill add to this, because the relevance of the second sentence of the above post at first escaped me. What the poster is saying is that maybe in my experimental work here I was free to flag some literature work as bullshit, but if I tried to publish it I would be put straight. Yes that is correct, part of the human condition is promulgating undeserved respect to others at the expense of honesty. I have several published papers to the effect that this or that publication is 'bullshit'. Each time I had to fight to get them accepted - sometimes in a different journal, one in particular took 3 years. What the poster is saying is that we soon learn not to write such things, in essence helping each other lie. But I dont want to give the impression that sources such as 'kirk-othmer' deliberately lie. It contains errors because its a compilation, its writers mostly have no direct knowledge of what they are reporting. Errors creap in because of misunderstandings with what the original work was trying to say, and because the science in some of the many original works it references is faulty. In the given instance Im not saying the phosphrus entry lies, the poster does.

[Edited on 22-11-2009 by len1]

blogfast25 - 22-11-2009 at 06:48

Quote: Originally posted by len1

Anyonw who says the above has just not tried to do much in the real world. Lost in the clouds.

DJF90 - well in that case he could have written Im sorry I can not do this or that, can someone check it out for me. That would have earned much more respect. When you start lying, youre finished as a scientist.

PS I thought Ill add to this, because the relevance of the second sentence of the above post at first escaped me. What the poster is saying is that maybe in my experimental work here I was free to flag some literature work as bullshit, but if I tried to publish it I would be put straight. Yes that is correct, part of the human condition is promulgating undeserved respect to others at the expense of honesty. I have several published papers to the effect that this or that publication is 'bullshit'.

Those are serious allegations, len1.

You've been able to flag some literature as bullshit and I can easily believe that. But it appears to be on those grounds that you believe also to be right in this case?

That clearly seems to be the train of thought:

"I showed some literature to be BS and I was right.
I believe this literature to be BS too, and since I was right about these other cases, ergo I must be right on this one too."

Not very empirical either.

I take it then that you have compared, back to back (silica/no silica), in controlled conditions, the alleged effect of silica?

Where you write "Yes that is correct, part of the human condition is promulgating undeserved respect to others at the expense of honesty.", I completely agree. But your statement doesn't mean you're right on this particular case.

From a thermochemical point of view, I don't think it matters much what form the phosphorus is in, as long as the species derives from P2O5. High temperature reaction of the phosphate with carbon pushes the equilibrium:

phosphate + C ---> n P4 (g) + m CO (g)

to the right due to the volatility and distilling off of the P4 and the CO - Law of Mass Action.

A Gibbs Free Energy change calculation ΔG = ΔH - TΔS could probably show a negative number (or close to 0) due to a strong entropic term.

It's possible the reaction is somehow facilitated by the presence of silica but the evidence doesn't seem to point strongly to that.

This reaction is similar to the reduction of chlorates (+V to -I), perchlorates (+VII to -I), nitrates (+V to o) or sulphates (+VI to -II) in thermite reactions.

len1 - 22-11-2009 at 09:18

Hey? No this was just my reply to that one who believes the literature is all good.

The lack of participation of silica in the reaction, was claimed by that wack in the tirade. Still the formula including silica is in the sources he mentioned, so it seemed little point discussing it. Instead I pointed out the interpretations of someone who only reads literature and imagines he can understand everything from that are likely to be misguided. I found the same to be the case with Sauron, he posted interesting literature, but his interpretation was of someone who had little first hand experience.

In the high temperature phosphate reaction the first stage is generally written as

Ca3(PO4)2 + 3SiO2 -> 3CaSiO3 + P2O5

with subsequent reduction by carbon. To see how silica drives the equilibrium, one only has to calculate the free energy change in the formation of CaSiO3 from the two oxides.

At very high temperatures you can get the reduction to proceed without silica as you say but the whole point of the metaphosphate process is that it goes at much easier attained temperatures

blogfast25 - 22-11-2009 at 10:28

If you suppose a significant difference in DH for these reactions:

Ca3(PO4)2 ---> 3 CaO + P2O5

3 CaO + 3 SiO2 ---> 3 CaSiO3

Then

Ca3(PO4)2 + 3 SiO2 --> 3 CaSiO3 + P2O5

could be significantly exothermic. Do you have some data?

And roughly what temperature for the metaphosphate process?

Here's a source that does confirm use of silica:

http://mysite.du.edu/~jcalvert/phys/phosphor.htm#Mine

"Elemental phosphorus is produced in electric [arc?] furnaces by heating phosphate rock, sand and coke. The silica combines with the phosphate rock to give calcium silicate, a slag, and phosphorus pentoxide. The pentoxide is then reduced by the coke to give phosphorus vapor and carbon monoxide."

Personally, I like garage Chemist's idea of thermiting a phosphate with Al powder (previous page of this thread), I'm just not sure why he wanted silica in there. Same reason as here?

I don't think it needs the silica: sulphates for instance are very successfully (and very hotly) reduced by Al. Al is of course much, much 'keener' on oxygen than carbon...

[Edited on 22-11-2009 by blogfast25]

S.C. Wack - 22-11-2009 at 12:37

No, you want to obscure with your assclown trolling the fact that you've got nothing to back you up, while I can find supporting references for my position only.

So I have to build an arc furnace to be right about carbon reducing phosphate without silica? Nonsense. Why can't someone else use theirs? Where in the references given, and in any of my posts, have I got something wrong? Feel free to look through all 5 1/2 years, to point out these misguided interpretations which you accuse me of. Why is it your word that must be taken as gospel when everyone disagrees with you, and how is your value higher than that of those who specialized in this field for years?

I merely present relevant work of certain published scientists at face value for informational purposes, because of the dearth, especially early on here, of same; sometimes for comparison with internet ramblings of bumbling anonymous amateurs, and invite debate. Sadly len wants none of this when it does not support him, not unlike Muslims in backwards countries banning schools. If you don't believe real scientists (especially when they patent their discovery) who tend to bother with details and analyses, then what can anyone here say about anything?

The fuck I'm a brainless receptacle incapable of critical thinking, I reject your rubbish because I do know it when I see it. Naturally anything published before 1930, especially in IEC, is bullshit because chemists were inept, unaware liars then. Other spouters of nonsense even in 1986 cite them anyways, such as in http://dx.doi.org/10.1007/BF02657149

Attachment: tirade.zip (1.8MB)
This file has been downloaded 724 times

entropy51 - 22-11-2009 at 13:33

Quote: Originally posted by len1

When you start lying, youre finished as a scientist.

I have several published papers to the effect that this or that publication is 'bullshit'. Each time I had to fight to get them accepted - sometimes in a different journal, one in particular took 3 years.

How very odd. I have published four journal articles reporting that I was unable to replicate another scientist's results. I had no problem whatsoever getting those papers into print. No fight. No delay. In fact they were all published in better journals than the original work which I could not reproduce.

Did your manuscripts perhaps contain the words "liar", "dishonest" or "bullshit"? This might explain your problems.

[Edited on 23-11-2009 by entropy51]

[Edited on 23-11-2009 by entropy51]

blogfast25 - 22-11-2009 at 14:03

Well, the paper attached by S.C.Wack (tirade.zip or Reduction of Tricalcium Phosphate by Carbon, Effect of Silica and Alumina on the Reaction, by K. D. Jacob, D. S. Reynolds, and W. L. Hill in IEC, Nov. 1929) actually strongly supports the role of silica (but also alumina and for similar reasons) as a rate-determining factor in the high temperature reduction of Ca triphosphate with carbon. And a very through investigation it was indeed.

blogfast25 - 25-11-2009 at 12:51

I wanted to test the hypothesis that it's possible to react tricalciumphosphate (TCP) with Al powder in an aluminothermic reaction:

Ca3(PO4)2 + 10/3 Al ---> 4/3 CaO + 5/3 CaAl2O4 + 1/2 P4

But I've no TCP or bone ash lying around. I did however have quite a surplus of ammonium dihydrogenphosphate (NH4H2PO4 or ADHP) from a crystal growing set (it grows lovely crystals) and decided to convert some to TCP by neutralising the ADHP to tertiary phosphate with NaOH and adding CaCl2 to that solution.

I set out to prepare 50 g of TCP (about 1/6 mol) from 37.2 g ADHP, 25.8 g NaOH and 53.7 g CaCl2 according:

NH4H2PO4 + 2 NaOH + 3/2 CaCl2 ---> 1/2 Ca3(PO4)2 + NH4Cl + 2 NaCl + 2 H2O

The solutions were quite concentrated: without reactions the solution strength would have been:

about 2/3 M ADHP
about 4/3 M NaOH
about 1 M CaCl2
about 500 ml.

After the first neutralising step (ADHP + NaOH, a whiff of ammonia can then be observed) I added the concentrated solution of CaCl2 to the tertiary phosphate solution. To my surprise, at first nothing seemed to happen. But the solution turned quite viscous, as a perfectly transparent gel of CTP started to form. Stirring caused the gel to start turning opaque/white, as did adding small amounts of water. I didn't expect such behaviour from such a thoroughly insoluble substance... The changes came about over several minutes.

I then added about another 500 ml of water to the 500 ml of 'slurry', stirred it and brought the thick, gelatinous mass slowly to the boil. The slurry gradually thinned out but remained viscous enough for the precipitate to remain fully suspended. I'm now filtering, washing and drying it.

I may be talking through my backside here but this type of precipitating may be useful in bone development, with newly excreted TCP gelling onto existing bone and gradually grafting itself onto it and maturing over time.

The dry and clean technical TCP should ready for the weekend...

Oxydro - 25-11-2009 at 16:27

I'm contemplating a phosphorus retort, although with me there are a great deal of projects planned that never happen. Money is the limiting factor, with time running a close second - I'm sure many of you are familiar with this!

The idea is to build a small electric furnace, fired with MoSi2 elements. Inside the furnace, I will have a "flask" made of either clay or iron. The problem with iron is that I don't know how long it will last at these temperatures. The problem with clay is that I don't know how to make airtight connections between sections (that have to be removable)

From the flask, a pipe runs into a water trap, in the same style as BromicAcid's experiments. However, I intent to use an inert gas flood over my watertrap to prevent problems with the P floating to the surface (As someone experienced, I think Bromic again). Is there a gain to be had from running inert gas through the reaction space as well? I suspect it might help the P distill over, but it will also carry valuable heat away from the reaction.

Nitrogen, CO2 or argon are all possibilities for an inert gas over the collection vessel. However, if I was to be running it through the hot section, I don't think CO2 would work ( carbon will likely reduce it to monoxide ). I have access to cylinders of gas, I work at a garage that is also a Linde distributor.

Problems to be solved include deciding what refractory from which to build the furnace, locating suitable moly disilicide elements, figuring out a power supply, and picking the best material for the reaction vessel.

Sedit - 25-11-2009 at 17:24

Quote:

The problem with clay is that I don't know how to make airtight connections between sections (that have to be removable)

I have a kiln and am pretty well versed in working with ceramics making me contemplate making a phosphorus retort on a number of occasions. Next time I fire up the kiln I no doubt will have a couple of retorts and ceramic tubes for other high temperature works. Any way my point is I thought about the same problem about sealing up the two sections until my common sence kicked in and I realized all I had to do is buy some very well named retort cement. Its sodium silicate and as it heats up it turns to SiO2 so it basicly just melts then solidifys as glass yet will still break along the seams if you want to. We use it quite often when working with boilers.

entropy51 - 28-11-2009 at 13:45

Quote:

Any way my point is I thought about the same problem about sealing up the two sections until my common sence kicked in and I realized all I had to do is buy some very well named retort cement

For what it's worth, some of the 200 year old books on www.archive.org give recipes for "luting" or something like that. They used all kinds of strange ingredients, as you will see if you investigate. Is wonder if sodium silicate by itself will stand up to making phosphorus. What about some of the stuff made for repairing fireplace brick?

12AX7 - 29-11-2009 at 00:35

"Furnace Cement" is bonded with sodium silicate, as far as I know. Sticky stuff, and very thick (like clay) fresh from the bucket. It tends to skin over, so you'd probably end up with steam bubbles (or worse, cracked glass) if you tried using it directly on glass joints.

Tim

blogfast25 - 29-11-2009 at 06:43

Well, my TCP is ready to go but it's raining cats and dogs here.

From the standard heats of formation of TCP (-4,132 kJ/mol, Wiki), CaO (- 635 kJ/mol, Wiki) and Al2O3 (-1,676 kJ/mol, NIST) can be deduced that the standard heat of reaction of:

Ca3(PO4)2 + 10/3 Al ---> 3 CaO + 5/3 Al2O3 + 2 P

must be about -566 kJ/mol (of TCP). That's pretty exothermic, almost as high as the HoF of CaO!

Hopefully the rain clears up later today or tomorrow...

blogfast25 - 1-12-2009 at 11:21

Well, well, well, somewhat unexpectedly I couldn't get the mixture of Ca3(PO4)2 powder = 1 mol, Al powder = 10/3 mol to light using a bit of KClO3/Al mix and a piece of Mg ribbon.

It was a small test: total batch weight 4.75 g.

I will now add some CaSO4/Al booster mix, the formulation will then be:

Ca3(PO4)2.............1 mol
CaSO4....................0.4 mol
Al...............................4.4 mol

The CaSO4 booster reacts like this:

CaSO4 +8/3 Al ---> CaS + 4/3 Al2O3

and will hopefully provide the missing heat.

garage chemist - 1-12-2009 at 12:42

You need either SiO2, or calcium metaphosphate, Ca(PO3)2. CaO will not form in this reaction.
All you will get from Al and orthophosphate is phosphide.

blogfast25 - 1-12-2009 at 14:50

Quote: Originally posted by garage chemist

You need either SiO2, or calcium metaphosphate, Ca(PO3)2. CaO will not form in this reaction.
All you will get from Al and orthophosphate is phosphide.

What makes you so sure of that?

garage chemist - 1-12-2009 at 19:22

There are cartridges for rodent eradication that consist of a simple phosphate-aluminum mixture. When ignited, they burn like thermite, but slower, producing a slag containing phosphide which then reacts with aerial and soil moisture generating phosphine.
They do not release any phosphorus during burning.
A patent for such a device is buried in this or another thread.

Additionally, a german article has been posted here which also concerned Al reduction of phosphates and clearly demonstrated the essential role of silica- unless reactive silica (diatomite, kieselguhr) was added to the mixture, the phosphorus yield was small.
Sodium metaphosphate (NaPO3 in any of its modifications) heated with Al or better Mg will produce some phosphorus of its own, but only a fraction of that which is released when silica is present.

See my last experiment in this thread, and try to repeat what is known to work. My recommendation.

len1 - 1-12-2009 at 19:41

Stephan my hat of to you for having the patience to explain when asked in such a tone. I have long since stopped bothering

kclo4 - 1-12-2009 at 21:12

Quote: Originally posted by Sedit

Have you ever thought about selling ceramic retorts that you make?
If you made them well, I bet a lot of people would be interested in this sort of thing, and might aid a few people in producing phosphorous, and who knows what other cool things.

Sedit - 2-12-2009 at 06:11

Haven't made one yet although it should be a breeze and I never really thought about selling them. I normaly sell the sculptures and potts not the morter/pestles and weird things like that I make. I suppose if I get good at it then Ebay couldn't hurt. Its been quite a while since I fired that kiln up because everytime I make something my kids destroy it before it makes it to the kiln. Lord knows I could use a little extra cash. Perhaps trade people for some real glassware

[Edited on 2-12-2009 by Sedit]

unome - 3-12-2009 at 00:20

Often wondered.... As silicon is fairly easily arrived at via a mild variation of the standard thermite process (such as here), could we utilize the reducing power of Si itself in this reaction? I mean, I don't pretend to be a scientist - not even a little bit (and try hard not to tell porkies, even of the ever-present social variety, makes for some tough social encounters let me tell you), but isn't Si above P in the table (kind of a shock really, given that it doesn't burst into flame on contact with air)?

Just asking a question, but flame, flame away if it so pleases you... Just seemed to me that a solid crystal mass (like that Si block) should be a sight easier to crush/break up into small enough pieces than Al...

[Edited on 3-12-2009 by unome]

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