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




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[*] posted on 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?
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[*] posted on 8-3-2009 at 16:25
Readman-Parker process


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.
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[*] posted on 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]
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[*] posted on 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!:(
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[*] posted on 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!
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[*] posted on 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]




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





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[*] posted on 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.




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[*] posted on 14-7-2009 at 00:03
ferrophosphorus


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




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

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

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[*] posted on 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
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[*] posted on 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





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




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