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Author: Subject: Preparation of elemental phosphorus
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[*] posted on 23-12-2005 at 15:47


Hello everyone, it's been a long time since I've posted here and I can't get my old account to work, but I'm back!

And I have exciting news - I have recently acquired an electric kiln, which should be good for around 800*C. Somewhere I have the manual which tells how long it takes to reach certain temperatures etc. It used to belong to my grandfather, who was a sculptor.

Here are some pics. In the first you can see two holes, just begging to have phosphorus flow out of them! The hole in the side is so you can see the thermometer/thermostat inside, the hole in the top is so you can see your pottery while it's cookin', I think. Both holes are about 1" diameter. The thermometer/thermostat (to be honest, I'm not really sure what it is exactly) is a little block of special material held in a little claw-like thing, which you can see in the second pic. You can get different materials for different temperatures. When the desired temperature is reached, the material softens and bends under the pressure from the claw-like thing. I'm not sure if this then reduces the power to keep it at that temperature, or if it just allows you to see that the desired temperature has been reached.

Anyway, with some plumbing and a tank of inert gas to make sure the pressure in the reaction vessel is always positive, even during the cool-down, combined with some of the lower-temperature methods discussed here, I am quite certain that this thing could produce significant amounts of phosphorus. I'll have to think of a lot of experiments in order to use it all up! And find a good source of phenyl chloride (for triphenylphosphine! :D).

I have absolutely no money at the moment, and so all my projects are in the dolldrums, but as soon as I can afford to and have the time I will start some experiments! I must admit, I am very excited :D.

Also, I will have to think of a safer place to run this thing. It might have to go in the garden when in use...

kiln 001.jpg - 26kB




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[*] posted on 23-12-2005 at 15:52


Second pic. Sorry they're a bit blurry.

I was thinking of phosphorus experiments, and to be honest, I could probably use a lot of it just by investigating all of its many allotropes!

Actually, now that I think about it a bit harder, the number of uses for phosphorus and its simple compounds (halides etc) are staggering.

kiln 002.jpg - 33kB




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[*] posted on 10-2-2006 at 11:42
Bone ash


I've got 4 Kg bones from the local butcher and I'm burning them at the moment. But the bones have some meat around which dehydrates at higher temperatures. The question is if the dehydration residue (carbon) of the meat (proteines) does reduce the Ca-phosphates of the bone at around 700-800°C to any phosphorous-containig compound that leaves the reaction system? The reaction system is in my case a big tin can that is filled with the bones. If its heated beyond the point in which the meat is converted to coal, there are some flames at the top of the tin. This means that here is some circulation with the air that oxydise the carbon or the carbon reduces the phosphate.
After I burned my 4 kg bones I got something around 1kg solid residues (bone ash). Is this normal, I had estimated more residues. And does anyone know if the phosphorous-salt is still in the bone ash or did it left in form of P2O5 or something like that
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[*] posted on 10-2-2006 at 18:07


All the phosphorous is still present, it takes a whole lot more heat to do that.

Don't bother excluding air, unless you want to reduce it with carbon. If you're roasting it, you might as well burn out the carbon.

Bones are porous so you won't seem to get much yield.

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[*] posted on 10-2-2006 at 23:54
Kiln


My kiln will reach Cone 10, 1370 C, and I'm hoping I can find the right chemical combination
to produce phosphorus at that temp. I'll keep looking around and if I have any success,
I'll report back. Interesting discussion, BTW.




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[*] posted on 9-3-2006 at 11:46


Just some other experiment I found: a demonstration experiment to show that bones contain phosphorus.
It is from a german site, the adress is here:
http://cc.upb.de/studienarbeiten/seidel/allgem_chem/versuche...

Translation of the important part:

Cleaned, boiled and dried chicken bones are burned with a bunsen burner on a fireproof surface and directly heated with the flame until they have turned into white ash.

2g of this bone ash are mixed with 0,5g magnesium powder and 0,5g kieselgur.
The mix is heated in a test tube which is plugged with a glasswool plug. After the reaction has finished, it is left to cool and the glasswool plug is removed in a darkened room and observed closely.
A glow is visible on the glasswool.
When the residue in the test tube is mixed with water, gas bubbles are evolved which self-ignite on contact with air. They are phosphine.

Reactions are on the site that I posted.

The important feature here is the use of magnesium instead of the often- used aluminium. Mg reacts at a much lower temperature than Al.
The SiO2 must be finely dispersed in the mix, hence the use of kieselgur. Quartz sand is not fine enough, even after good grinding.
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[*] posted on 9-3-2006 at 14:24


Garage Chemist this is potentially very good news. I must give this a try.

Why do you think this works without the application of high external heat? Does the reaction itself generate a very high temperature?

[Edited on 9-3-2006 by Magpie]




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[*] posted on 9-3-2006 at 14:33


Yes, the reaction of calcium phosphate with Mg is very exothermic.
Good heat is still necessary though.
The following reaction, the reaction of the phosphide with SiO2, is the actual phosphorus producing reaction, and the important bottleneck in DIY P manufacture.
The use of kieselgur promotes the reaction by increasing contact area due to its fine microporous structure.
The ingredients should also be grinded together very well, due to this reason (dont use too fine Mg powder though, otherwise it might be more of a pyrotechnic mix. But the SiO2 and bone ash should be intimately mixed.).
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[*] posted on 11-3-2006 at 12:24


Maybe fine magnesium would be ok if the powdered composition was pressed into a large pellet or into something like a rocket grain (maybe with a small amount of non water based binder) I kind of like the idea of a self sustaining reaction. Light fuse, retire, and see what happens.
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[*] posted on 14-3-2006 at 11:20
common phosphorous sources


how musch phosphoric acid is in coca cola?
would there be enough carbon in the corn syrup and stuff to
reduce it if you poured it on sand and distilled it?
it'd be a neat parlour trick:)

i was going to drink it and eat alot of bananas and distill my personal by-product, but i like this idea better
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[*] posted on 16-3-2006 at 17:53


Quote:
Originally posted by garage chemist
Yes, the reaction of calcium phosphate with Mg is very exothermic.
Good heat is still necessary though.
The following reaction, the reaction of the phosphide with SiO2, is the actual phosphorus producing reaction, and the important bottleneck in DIY P manufacture.
The use of kieselgur promotes the reaction by increasing contact area due to its fine microporous structure.
The ingredients should also be grinded together very well, due to this reason (dont use too fine Mg powder though, otherwise it might be more of a pyrotechnic mix. But the SiO2 and bone ash should be intimately mixed.).


Would other metal ions interfere with the reaction? A flux could be added so that the SiO2 melts. Perhaps some very finely powdered soda lime glass? Easily available. Naive reading of the electrochemical series suggests that Na+ and Ca++ wouldn't react much.
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[*] posted on 17-3-2006 at 01:20


Dunno, alkalis like to be reduced by Mg and Al remember. Ca should stay put, though.

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[*] posted on 19-3-2006 at 20:59


If Na glass would get reduced, maybe K2CO3 would make a good flux. MP circa 1000C is a lot better than silica at approx 1700C.
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[*] posted on 19-3-2006 at 22:24


(Cyrus is catching the phosphorus bug again, but this time I'm much more prepared for high temperature reactions...)

K is an alkali metal (I'm sure you know that, but it was mentioned above that Mg reduces alkali metals) , and so I doubt it would work.

Hmm. The idea of heavily fluxed silica is appealing, as it would enable the silica to react better IMO.... I was thinking of borosilicate glass, as everyone has plenty of that, and it melts very easily, but B2O will react with Mg to produce B. Mg is just too darned reactive. IIRC there isn't anything on the CaO - SiO2 phase diagram that melts at a low enough temperature for us to like it. :P

I'm interested in that reaction involving phosphates and carbon that was mentioned earlier in this thread. If the temperatures are low enough that steel can be used as a reaction vessel, then it should be pretty doable.

Side note- It seems that lots of phosphine is produced by adding water to Al slag from pop cans according to 12AX7. (2AlP + 3H2O -> Al2O3 + 2PH3 ) I wonder if the amount produced is useful in making phosphorus... there can't have been tremendous amounts of phosphine around, as I'm still alive.




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[*] posted on 20-3-2006 at 07:17


How about first reacting the phosphate with magnesium (without SiO2), and THEN adding the SiO2 with the flux of your choice?

This should eliminate the problems with reduction of the flux, since there will be no magnesium to react with it.

Remember: It is very easy to reduce phosphates to phosphides with magnesium, the reaction is exothermic and self- sustaining.
The subsequent production of phosphorus from the phosphide is the tricky part.

[Edited on 20-3-2006 by garage chemist]
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[*] posted on 27-3-2006 at 09:23


Today, I've heated some matchbox-striking-surfaces in a test tube. Bevore I did that, I eliminated all oxygen by venting the tube with gas from my burner and lighting the mixture afterwards. Then I plugged the tube with glass wool and heated the red P of the matchbox till some brown-yellow vapour apeared. Now I switched out the light and saw no caracteristic glow of the WP (P4).
The room was not completly dark, so I went into my cellar where it's really dark and watched the test tube aggain, nothing. I got a knife and pushed the glass wool plug down and took it out aggain, then I put the striker residues from the test tube together with the plug on my hands and crushed it gently. Now I saw the WP glowing very weak in a greenish-white on my hands in the dark.When I rubbes the plug aggain it ignited with a fizzling sound and I threw it away imidiately. (It wasn't surely very healty to have some WP on my hand but I expected nothing of it to be there.)

What I like to see with this mayby a bit trivial experiment is that the phosphorous glows so weak that someone working in a not really dark room doesn't see the glow of it for which everyone around here is looking so much. Maybe several people here prepared some WP but didn't knew it in the lack of seeing the glow of it.

In an earlier experiment I heated some crude boneash with quarz sand and charcoal powder in a quarz glass tube plugged at the bottom and the top with some glass wool.
The source of the tube was a bit tricky and I'm not absolutly sure if it's really quarz. I got it from an old microwave ofen with integrated grill, the heating coils were in those tubes. ( they aren't clear there are some bubbles inside)

Those tubes are also found inside there:
http://www.dicker-bauchladen.de/catalog/images/reer-heizstra...

When I heated the tube, PxHy and PH3 came out and I burned it with a green phosphor-like flame. Other people had the same thing, but I ask myself if those gasses (PH3 etc.) are also containing phosphorous vapours, which everyone wanted to recognize by their glowing, but which is to weak to be seen in an alluminated area.


[Edited on 27-3-2006 by hinz]
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[*] posted on 31-3-2006 at 18:34


I'm going to live up to my name - I don't understand how Mg can reduce Na or K. The enthalpy of formation of (say) MgCl2 vs. 2 NaCl favors Na reducing Mg, if I read the tables correctly. Furthermore, this came off of the Purdue chemistry review:
Code:
n the basis of many such experiments, the common oxidation-reduction half-reactions have been organized into a table in which the strongest reducing agents are at one end and the strongest oxidizing agents are at the other, as shown in the table below. By convention, all of the half-reactions are written in the direction of reduction. Furthermore, by convention, the strongest reducing agents are usually found at the top of the table. The Relative Strengths of Common Oxidizing Agents and Reducing Agents K+ + e- <----> K Best Ba2+ + 2 e- <----> Ba reducing Ca2+ + 2 e- <----> Ca agents Na+ + e- <----> Na Mg2+ + 2 e- <----> Mg H2 + 2 e- <----> 2 H- Al3+ + 3 e- <----> Al

I understand that removing the correct reaction products can drive the reaction "backwards" at an energy cost.
A clue, please?
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[*] posted on 31-3-2006 at 19:56


It's probably in the binding energy of MgO and Al2O3, which have very high melting points and thus are very stable. Chlorides tend to melt under orange heat (800C) so when you heat one up with another metal, it'll go basically by ionic rules.

It's perfectly safe to melt aluminum or magnesium and use a sodium, potassium, calcium (and others) chloride (and fluoride) melt to flux it.

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[*] posted on 1-4-2006 at 23:03


I'd like to try making WP from off the shelf stuff, here's how i propose to do it.

Ingredients:

Phosphoric acid - H3PO4 (85% solution, Ph down, from Pet store)
Caustic soda - NaOH (98% from Supermarket)
Ammonia - NH3 (8% soluction, from Hardware)
Silica - SiO2 (100% little packets with shoes and stuff)
Powdered aluminium - Al (100% Printing additive)


1. H3PO4 + NaOH ==> NaH2PO4 + H2O

2. NaH2PO4 + NH3 ==> (NH4)NaHPO4

3. Heat (NH4)NaHPO4 ==> NH3 + H2O + NaPO3

4. Heat 12NaPO3 + 20Al +6SiO2 ==> 6NaSiO3 + 10Al2O3+ 3P4

I'm planning to collect the Phosphorus under water same way as has already been described in this thread.

I don't think anyone has mentioned making the Microcosmic salt from Sodium dihydrogen sulphate and ammonium. Any reason why this wouldn't work?

Then on heating it should reduce to Sodium Phosphite?

I'm very interested to hear anyones thoughts... this is a great thread.
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[*] posted on 5-4-2006 at 03:39


has anyone tried the electrolysis of molten trisodium phosphate? any opinions? i have been searching for any examples of this and nothing - it might be something to try with some co2 to flush out the air in a retort - i dont think phosphine would be liberated since no hydrogen is there - would all the oxygen be removed at the anode?
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[*] posted on 5-4-2006 at 06:14


Hmm, sodium phosphide..??

Can't find a melting point for TSP, even an MSDS for "anhydrous" claims the same melting point as the dodecahydrate. :mad:

Anyways, I would imagine electrolysis would reduce the phosphate first. You could then add some, say, sulfur, and distill off the phosphorous, leaving sodium sulfide.

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[*] posted on 5-4-2006 at 06:19


Hmm, sodium phosphide..??

Can't find a melting point for TSP, even an MSDS for "anhydrous" claims the same melting point as the dodecahydrate. :mad:

Anyways, I would imagine electrolysis would reduce the phosphate first. You could then add some, say, sulfur, and distill off the phosphorous, leaving sodium sulfide.

Hmm, efficiency could be low.. sodium isn't going to want to form with phosphorous available so I think it would form a molten phosphide, which would then oxidize from the O2 bubbles produced by the anode...

Heh, there's another thought: y'think carbon could displace phosphorous, in say, molten lithium phosphate? Ooo, and phosphorous' electronegativity *is* lower than carbon's. I'm thinking Li3PO4(l) + 5C(s) = 3/4 Li4C(l/s) + P(g/l) + 4CO(g) (aww crap, the hell with balancing a 4/3rds ratio!). Lithium forms a carbide right? :P

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[*] posted on 5-4-2006 at 12:38


nevermind i looked a few posts back - anhydrous sodium phosphate is hell to melt - so i guess the only way through this is a carbon arc in a vacuum chamber of some sort or maybe just one filled with co2 - hot phosphorus doesnt reduce co2 does it? - pretty fancy setup - but im sure it would work

According to Lange's 15th, Na3PO4 (anhydrous) melts @1340*C. The 12-hydrate melts @ 73.4*C and looses 11 H2O @100*C. No decomposition is mentioned,

or CRC sodium pyrophosphate is 988*C either way far too hot

[Edited on 5-4-2006 by jimmyboy]

[Edited on 6-4-2006 by jimmyboy]
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[*] posted on 5-4-2006 at 13:44


Wow, that's gotta be the most refractory sodium compound there is!

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[*] posted on 6-4-2006 at 13:35


I've had some success so far.

Phosphoric acid combines well with Sodium Hydroxide. It gets pretty hot obviously, so its important to combine the solutions slowly :)

Next i mixed the Sodium Dihydrogen Phospate solution with and excess of 30% ammonia solution.

I'm pretty sure this reacted in the solution to form (NH4)NaHPO4. Anyway, on heating it gave off copious amount of foul smelling gas which i presume to be NH3 + H2O. Leaving me with a white powder which should be NaPO3. Probably Sodium Hexametaphosphate right?

Does anyone know a good test to check i've produced (NaPO3)6 ?

Next is the fun part:

4. Heat 12NaPO3 + 20Al +6SiO2 ==> 6NaSiO3 + 10Al2O3+ 3P4
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