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Author: Subject: Preparation of elemental phosphorus
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[*] posted on 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.




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[*] posted on 15-3-2008 at 12:26


GC: The link loads for me just fine. Are you giving it enough time to load?
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[*] posted on 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
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[*] posted on 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]




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[*] posted on 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
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[*] posted on 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.




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[*] posted on 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.




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[*] posted on 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...
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[*] posted on 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]
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[*] posted on 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.
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[*] posted on 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.




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[*] posted on 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]
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[*] posted on 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?




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[*] posted on 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.




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[*] posted on 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.




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[*] posted on 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.




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[*] posted on 9-6-2008 at 19:22


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



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[*] posted on 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 786 times





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[*] posted on 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)
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[*] posted on 15-6-2008 at 12:44


Very nice, Strepta. And thanks for the nice write-up.
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[*] posted on 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]




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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 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)
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[*] posted on 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.
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