Sciencemadness Discussion Board

Preparation of elemental phosphorus

 Pages:  1  2    4  ..  15

BromicAcid - 19-9-2004 at 08:13

Doesn't trisodium phosphate have an impractical temperature of melting? My chem book says that it looses its last water of hydration at 200C but that it decomposes upon melting, I'm assuming into sodium pyrophosphate, which has a melting point of around 800C.

I've read of the method to produce iron phosphide by heating phosphate, iron, and silica and the iron phosphide sinks to the bottom. However converting this to phosphine and reducing seems to be a bit too much work considering the prep for phosphorus from this step is to simply heat iron phosphide with sulfur, or iron sulfide.

From the patents that I found phosphine is usually generated in the lab by the action of water on magnesium phosphide. If magnesium were easy to come by I have no doubt that one could easily make magnesium phosphide by heating a mixture of phosphate, silica, and excess magnesium, however it is not so easy to come by in a reliable form.

In my newest chemistry book it describes the analytical process to determine arsenic content of an unknown sample. Many here are probably familiar with it, ash the sample containing arsenic then dissolve the sample in HCl. The H3AsO4 thus formed is treated with SnCl2 and reduced to H3AsO3, this is in turn reduced with zinc metal according to the reaction:
H3AsO3 + 3Zn +6HCl ---> AsH3 + 3ZnCl2 + 3H2O
So essentially a dissolving metal reaction.

I wonder if dissolving a phosphate in HCl and adding excess metallic tin to the reaction mixture might yield a decent quantity of phosphine? It would preform both reactions simultaneously, however it might be difficult to control.

Just some thoughts.

Edit: Just tried mixing HCl, Sn, and Ca2(PO3)3, upon mixing the HCl with calcium phosphate the solution turned yellow, upon adding the tin the solution decolorized to clear, but shortly thereafter turned black. I caught the faintest whiffs of phosphine but may have been entirely mixtaken due to the biting smell of HCl that overpowered everything. Got the test tube sitting in the sun to see what I end up with, it was odd that it turned black.

Edit #2: If a dissolving metal reduction will reduce nitrate to ammonia or nitro to amine, and the same/similar readuction reduces arsenates to arsine, it would only make sense for a similar reduction for phosphates to phosphine.

[Edited on 9/19/2004 by BromicAcid]

Theoretic - 24-9-2004 at 05:38

"Doesn't trisodium phosphate have an impractical temperature of melting? My chem book says that it looses its last water of hydration at 200C but that it decomposes upon melting, I'm assuming into sodium pyrophosphate, which has a melting point of around 800C."

No, it doesn't. Na3PO4 can only decompose into Na4P2O7 if it loses Na2O, which I'm sure isn't going to happen.

BromicAcid - 24-9-2004 at 08:41

Anhydrous Trisodium Phosphate melting point information:

Melts at 73 - 75 C
Decomposes at 75C
Melts at 1340 C
Decomposes at 75C
Becomes anhydrous at 212F Melts at 75C

So, I'm confused as to what it decomposes to, as if it just lost its water of hydration it would simply state that.

[Edited on 9/24/2004 by BromicAcid]

neutrino - 24-9-2004 at 17:49

According to Lange's 15th, Na<sub>3</sub>PO<sub>4</sub> (anhydrous) melts @1340*C. The 12-hydrate melts @ 73.4*C and looses 11 H<sub>2</sub>O @100*C. No decomposition is mentioned,

BromicAcid - 24-9-2004 at 18:21

That is quite the prohibitively high temperature. I found something interesting though regarding the use of sodium phosphate, particularly in the production of phosphides.

Supposedly according to an older chemistry book I was browsing around today, the only way to make tungsten phosphide is to electrolyze a mixture of tungsten oxide in molten Na3PO4.

But then again electrolyzing calcium phosphate in cryolyte supposedly yields phosphorus directly.

[Edited on 9/25/2004 by BromicAcid]

Theoretic - 1-10-2004 at 08:28

Ca3(PO4)2 dissolves in HCl? Very interesting! Though I would suggest Fe rather than Sn, much more available and dissolves slower than Sn, thus more of the H atoms are utilized due to less chance of recombination at low concentration.
Fe could be used for direct reduction of molten sodium whateverphosphate, more available than zinc and would work the same way, though a weaker reducer.
For carbon reduction, I suggest NH4H2PO4, this would convert to ammonium metaphosphate (AMP), thinly powdered sugar would be thrown in and dehydrated by the metaphosphate and produce a very fine active form of carbon (taken from the CS2 thread), the dihydrogenphosphate (ADHP) thus formed would dehydrate back to metaphosphate, this would be reduced by the active carbon much more actively than by ordinary powdered charcoal. These would be the reactions:

1) NH4H2PO4 => NH4PO3 + H2O

2) 6NH4PO3 + C6H12O6 => 6C + 6NH4H2P2O7

3) NH4H2PO3 => reaction 1)

4) NH4PO3 => reaction 2)

5) 2NH4PO3 + 6C => 6CO + 2NH3 + H2 + 2P

Theoretic - 1-10-2004 at 11:32

A sidenote: what to do with the off-gases after condensing the phosphorus? You could burn them, although that would be a waste (but it would be interesting to see the flame colour, since NH3 burns pale green-yellow, CO blue, H2 colourless :P). You could lead them all to a hot tube with CuO and get copper powder and harmless waste gases (CO2, H2O, N2). Or you could lead them to a very hot ceramic (or with ceramic bits) tube and make HCN from the NH3 and CO, as well as water vapour. After condensing HCN + H2O in ice-cool water, the remaining CO and H2 (in ratio 4CO/H2) could be led into molten NaOH to make sodium formate. The hydrogen left over can be led into a condom and stored. :D Although I'm only half joking, since condoms, when they are first manufactured, are filled with 25 l of air to test for strtength, so one could easily accomodate a mole of H2 (22.4l), the amount that would be left over from a 446 g batch. Then you could take the condom, tie it, attach it with a long gasoline-soaked thread to a brick somewhere in a field, ignite the thread and watch. The thread will ignite, the condom will float into the air and explode in the sky with a big boom when the flame gets to it.
Thus every single reaction product is utilised. :D

elemental phosphorus making

Phosphorus1 - 7-10-2004 at 09:28

Hello everybody!!
this is my first message on this list. I am actually Italian, and I am studying for a PhD in chemistry in the UK, so I apologise if my English is a bit rusty.
Regarding making phosphorus. I have been 'obsessed' with making white phosphorus for a long time. I have actually succeded making about 100 mg of white phosphorus in my landlady's garden a couple of months ago but never tried again since. That night I was so happy I couldn't go to sleep. I obtained a nice yellow 'drop' at the bottom of my condenser which solidified as the water temperature dropped. When I got the pebble out with a spoon, it smoked in air, smelt badly of garlic and yes, it shone green in the dark. It didn't catch fire though, even when pierced with a toothpick. The consistency is that of soft wax. I used a mixture of KH2PO4 and homemade willow charcoal with no sand. The retort was made by 'incapsulating' a 10 ml lab glass vial with 'fire' cement and cooking this in the kitchen oven at 250 C for 1 h. The spout was a piece of copper tubing sealed on the retort with the same fire cement, (this comes as a ready-made putty in my local hardware store). The furnace design is very simple indeed: I have made a refractory kiln with a blowing pipe attached to a 'cold-shot' powerful hair-drier. The furnace was fired with BBQ charcoal. I have successfully melted iron in this, so I guess the temperature at full regime, must have been in excess of 1300 C (it looked so bright it would hurt my eyes to stare at it).
I have tried with thicker glass jars, without success. I guess the glass, which replaces the silica, melts inside the fire cement 'mold' and acts as a flux aiding the melting of the phosphate. The usual reduction reaction then occurs. It took 30 min or so at bright white heat for the first spontaneously flammable bubbles to break the surface of the water in the condenser. They produced white-bright little flames, so I guess some phosphorus got lost in that way. I am now thinking of repeating the experiment with ground glass powder in a slighlty bigger retort.
I still have my pellet of P4 in a small jar or water. It now sits proudly on my desk.

Remember to keep the retort size as SMALL as possible so that it will be easier to achieve and sustain internal high temperatures, and:
do NOT use metal retorts. Too much heat is simply transferred away to the spout and then to the water in the condenser.
Do not use gas torches unless you have an acetylene/oxygen source. Go for a nice charcoal/air furnace which you can make with an old bucket and refractory mix (and a nice hair-drier from your mum/girlfriend/granny...)

Phosphorus-makers of the world ! Get out there NOW and start making phosphorus!!

Phosphorus1

Theoretic - 20-10-2004 at 00:21

If you are ready to rock n’ roll, but lack in a strong heat source, you can do a neat trick or two.
1) Inject a little oxygen in the flame of your good old propane burner.
2) Make a hydrogen flame. The temperature is 1000 C – 2000 C. Plus, it flows around thing very easily, it’s very fluid and relatively long, it will flow around your vessel and heat it evenly.

crushpack - 29-10-2004 at 02:17

In some textbooks I've met with maybe one of the most common method for laboratory phosphorus preparation. I wonder if you know of that.
In substance this method is based on reduction of phosphate, but this phosphate is here sodium ammonium hydrogen phosphate Na(NH4)HPO3. As for reducing agent zinc dust was mentioned but Mg or Al will sure work as well. I've never tried this though I've prepared some sodium ammonium phosphate. If anyone try this with success please tell me.

reaction(I can be wrong):

Na(NH4)HPO3 = heat= NH3 + H2O + NaPO3

4NaPO3 + 5Zn = 2P + 5ZnO + Na4P2O7

totally:

4Na(NH4)HPO3 + 5Zn = 2P + 5ZnO + Na4P2O7 + 4NH3 + 4H2O

FrankRizzo - 29-10-2004 at 18:10

I second that. Just cut a little notch out of the side of a >1982 US penny, hold it with a pliers so that the notch faces downward, and then heat it with a propane torch. When the copper skin of the penny becomes wrinkly, jerk the pliers slightly, and a blob of molten zinc will plop out of the copper shell. I suggest doing this procedure outside, above a non-flammable surface (concrete). Also, the zinc will splash a bit, so wear protection.

[Edited on 30-10-2004 by FrankRizzo]

jimmyboy - 29-12-2004 at 06:49

Has anyone actually tried out Theoretic's method -- maybe melting down some zinc with TSP from the local hardware store into a flask of cold water? sounds neat - toxic as hell but neat :D

S.C. Wack - 29-12-2004 at 10:20

Since this has come up again, this is Rossel's Berichte ref. It is the Al reduction mentioned early on. Unfortunately his Zn ref in Bull. Soc. Chim. isn't so available to me.

Attachment: berichte_P_from_Al_reduction.pdf (210kB)
This file has been downloaded 2103 times


garage chemist - 29-12-2004 at 14:41

S.C. Wack: That's VERY interesting.
The resource says that when sodium phosphate ("NaPO3", obtained by heating NH4NaHPO4) is heated with aluminium powder, elemental phosphorus is directly evolved without the formation of hydrogen phosphide.
Addition of fine SiO2 to the mix improves the yield.
However, the aluminium must not be too fine, or the reaction becomes violent.

Al + NaPO3 + SiO2 is the ideal mixture, it also reacts at the melting point of NaPO3, which makes a strong heat source unnecessary.

The synthesis of phosphorus is not so important for me though, as I once bought 100g of red P from Ebay (was expensive as hell though :mad: ).


[Edited on 29-12-2004 by garage chemist]

neutrino - 29-12-2004 at 17:13

Out of curiosity, how much was it? By the way, NaPO<sub>3</sub> is sodium <i>meta</i>phosphate.

BromicAcid - 29-12-2004 at 20:32

Which is the exact method I use to make phosphorus (the one that works best at least) the name you'll most often find NaPO3 under is sodium hexametaphosphate, it makes a cyclic structure if you see the drawing of it so it is better represented (NaPO3)6 the preparation of (NaPO3)6 is also covered in Gmellin (I think I butchered that name) which I scanned in earlier in this thread.

jimmyboy - 9-4-2005 at 22:06

Read a little on this subject and keep reading over and over that you must use "acid calcium phosphate" Ca(H2PO4)2 made by adding sulfuric to the neutral insoluble calcium phosphate Ca3(PO4)2 or the phosphorus simply won't reduce - maybe everyone knows this already but just checking

jimmyboy - 12-4-2005 at 16:09

has anyone tried melting zinc in sodium metaphosphate yet? sounds so interesting - dangerous as all getout but interesting :)

BromicAcid - 12-4-2005 at 17:14

The 'acid calcium phosphate' is also known as calcium superphosphate, it is made by adding sulfuric acid to calcium phosphate till the precipitate re-dissolves again, calcium phosphate itself does work but the temps are somewhat higher, calcium superphosphate = home production, calcium phosphate = industrial, at least in my opinion.

Zinc does work as a reducing agent, yields are lower then with aluminum reduction from what I've read and just as with aluminum or magnesium reduction yields are also lowered from the reaction of the phosphorus produced with the zinc present to form the phosphide. Although I didn't check into it I have to wonder about the stability of magnesium phosphide, aluminum phosphide and zinc phosphide to heat, if they decompose at elevated temperatures >800 then there would be no need for silica in the reaction because they should just decompose, right?

And while I'm posting to this topic, someone (I will if no one else gets around to it) should attempt the reaction I mentioned earlier on this page of heating iron, a phosphate and silica in a crucible, the iron if in excess reacts with the phosphate and iron phosphide sinks to the bottom of the crucible forming a seperable hunk. Heating this iron phosphide with sulfur (which would imply low temperatures considering there is no mention of pressure) furnishes the distillation of phosphorus.

jimmyboy - 12-4-2005 at 21:48

heck both zinc and tin are better reductants than phosphorus - i might try to use a roll of solder for this

evil_lurker - 13-4-2005 at 15:50

Heres a coupla things I want to try...

According to a patent, can't remember which one right off, good results were obtained when the chemicals to be reduced were ball milled together and compressed into 1 inch or so "briquettes" via hydrolic press at some 15 tons of pressure. A nice side benifit other than good yields was that reduction was completed at 1250C whereas on an industrial scale it is just getting going.

Other thing I want to do is instead of dumping the WP vapor into water, is to have a long pipe as the condenser, with a one way check and ball valve on the end. That way when the RXN is over, the entire vessel can be sealed keeping the atmosphere anoxic, and hopefully keeping the WP from self igniting. When its time to take out the WP, inert gas can be pumped into the end the end with the ball valve while the other unscrewed from the vessel.

Speaking of removing WP... it is soluable in carbon disulfide (which is a nasty substance of its own). So after the condenser is removed and another ball valve screwed on, some carbon disulfide can be introduced into the condenser and sloshed around effectively washing it out.

Well so you got ur WP in the carbon disulfide... what then? Easy, keep it in a glass jug, and sit it out in the sun for a few days. The sun will react with the dissolved WP particles turning them into RP. RP is insoluable in carbon disulfide and it will precipitate and fall out of the solution onto the bottom of the jug.

Easy... lol. :D

Twospoons - 13-4-2005 at 20:55

jimmyboy, solder is tin/lead, not tin/zinc. Unless you've got lead free solder which is ~95% tin, the rest being silver, copper and sometimes indium.

Esplosivo - 13-4-2005 at 21:12

Could any body clear up some of my confusion please? If I were to do this which calcium phosphate should I buy, the monobasic calcium phosphate Ca(H2PO4)2 H2O, dibasic calcium phosphate CaHPO4 or calcium phosphate Ca3(PO4)2? I got a little mixed up, so sorry if this looks stupid, but better be sure. Thank you.

jimmyboy - 13-4-2005 at 23:07

Not so twospoons

http://www.kappalloy.com/tinzinc.html

although yes the more common ones do not contain zinc - i think they also use this alloy for alumaweld rods

[Edited on 14-4-2005 by jimmyboy]

12AX7 - 14-4-2005 at 00:05

Either that, or aluminum-zinc alloys (essentially potmetal). Which would also work quite well as fuel.

Tim

sparkgap - 14-4-2005 at 04:42

Esplosivo, it's Ca(H<sub>2</sub>PO<sub>4</sub>;)<sub>2</sub> that is needed. :)

Off topic, but solder composition methinks depends on your locale. Lead-tin solder has not yet been phased out in some territories. :D
So you may be both correct. (Personally though, I have yet to see zinc-tin solder.)

evil_lurker's idea sounds good, but just to avoid the odoriferous carbon disulfide, are there substances where white phosphorus is soluble and red phosphorus is not?

sparky (^_^)

garage chemist - 14-4-2005 at 14:30

Conversion of WP into RP via light proceeds incredibly slow, it is useless as a preparative method.
A better method to get out the phosphorus is to fill water into the pipe, warm it to 50°C, let it cool again and take the phosphorus blob out with pliers.

RP is insoluble in all solvents, BTW.

[Edited on 14-4-2005 by garage chemist]

BromicAcid - 14-4-2005 at 15:39

Red phosphorus is soluble in phosphorus tribromide, also the reaction of white phosphorus with light is appreciable, even a solid hunk of white phosphorus will convert to red readily in the presence of light. I have some that I could show you that has turned somewhat rapidly, check my book project under phosphorus at the end and I have a little picture.

evil_lurker - 14-4-2005 at 19:35

The only other half arsed decent WP solvent I can find is benzene. WP is soluable in it up to 1 g/35 ml. WP is very soluable in carbon disulfide... 1g/0.8 ml... 1 liter of it would go a loooong way.

Phosphorus tribromide is not a good candidate for a solvent either. The RP is soluable in it as well as the WP making the final product hard to separate.

As far as sunlight turning the WP in RP, I think it might be worth a shot to expose it to a concentrated UV radiation source. A 400 watt mercury vapor lamp that has had the outer glass bulb broken off will put out a crapload of UV radiation. Coupled with a glass container of proper UV conductivity, it would be interesting to see what would happen.

BromicAcid - 14-4-2005 at 19:58

Yes, it takes a lot of benzene to dissolve appreciable phosphorus. Significantly cheaper though a little worse is olive oil, about one gram per 80 ml, also an option is chloroform, ~1 g / 40 ml. Chemical desctruction of phosphorus in an apparatus is always an option, but as my experience shows it can become coated in... whatever.... which can prevent further oxidation and leave with with a fire hazard later on. Someone should try the reaction between aqueous copper sulfate and white phosphorus for destruction if the situation presents itself, it is one of the methods mentioned in a book I read on the destruction of dangerous reagents in the labratory.

Theoretic - 27-4-2005 at 09:13

Just a useful detail, Ca3(PO4)2 can be converted to Ca(H2PO4)2 by reacting with HCl, and the dihydrogen phosphate should be much easier to reduce (even if it turns to Ca(PO3)2!), while calcium triphosphate is cheaper than straight sodium triphosphate.

JohnWW - 29-4-2005 at 18:11

Are you sure? Sulfuric acid would do that to tricalcium phosphate, but HCl is too volatile and not as strong an acid as phosphoric acid.

neutrino - 29-4-2005 at 18:53

HCl is a strong acid, whereas phosphoric acid is considered a weak acid, especially after loosing two protons and forming HPO<sub>4</sub><sup>2-</sup>. This means that the equilibrium here is shifted strongly to the HPO<sub>4</sub><sup>2-</sup> side.

Marvin - 29-4-2005 at 18:55

Hmm, not as strong as which hydrogen of phosphoric acid? ;)

I suspect this would work. Use of sulphuric produces the problem of separating the sulphate unless you go all the way to phosphoric acid.

Ideally I have the feeling there is a 'magic salt' mixture consisting of a low melting eutectic of metaphosphates that could be regenerated over again with phosphoric acid.

evil_lurker - 5-5-2005 at 13:52

I just read over the lead phosphate reduction patent and it seems though there might be a way to do phosphorus reduction with controlled temps under 500C (sort of).

Let me explain for those that haven't read some of the previous posts or the patent.

According to the patent, lead phosphate or Pb3(PO4)2 is reduced under hydrogen or methane (natural gas comes to mind) with hydrogen resulting in the highest yields and methane about 50% of that.

The reaction consists of three stages:


1. The Pb3(PO4)2 is heated up to 300C to drive off any existing water.

2. Once the temp hits 300C the hydrogen is turned on and the tempurature slowly raised to 500C. The hydrogen reduces the Pb3(PO4)2 by ripping off the oxygen molecules and forming Pb3P2, aka lead phosphide.

3. Upon the cessation of evolution of water, the furnace is again slowly raised up to somewhere between 650-800C. According to the patent, small amounts of PH3 are liberated at around 600C. This makes sense, the Pb3P2 probably starts to break down somewhere around 600C and thus liberates PH3, which subsequently start to be reduced to H2 and elemental P at around 650C, so basically at the beginning of the reduction temp the phosphine being liberated is not hot enough to break down.

What if one were to use small vessels for batch reduction of Pb3(PO4)2. They could be heated in a pot full of molten lead at the proper tempurature more easily that way (everyone knows that without some sort of flux lead doesn't stick to metal worth a damn). The resulting Pb3P2 could then be stored in a container with a dessicant. Once a sufficiant quantity has been made, some sort of reactant could then be slowly dripped thru the product creating phosphine gas, which is routed to a glowing steel tube in fire, where it is effecively reduced to its elements, H and P.

This would eliminate the need for larger (read expensive) reduction vessels, easier cleanup, and less cost.

Alternatively, it would be much easier to react the aluminum with phosphoric acid, and reduce it in the same manner, however no literature that I know of details this.

another trivial experiment

Polverone - 10-5-2005 at 18:49

The last few times I'd attempted some sort of phosphorus production, I started out with mixtures of phosphates and reducing agents. I've been working in borosilicate test tubes, which aren't the most heat-resistant vessels, so I usually ended up melting the glass before I saw any sort of reaction. Today was different.

2 g of lead wire were placed in a borosilicate test tube along with 1 ml of 85% H3PO4. This was heated in a propane torch flame, carefully at first as water was driven off. Heating was increased, and the lead melted under the acid. After a couple minutes a thin stream of whitish smoke started wisping from the test tube. The smoke had the characteristic smell of burning phosphorus. It occurred to me after a bit to turn off the light, and I saw a mysterious and beautiful site: there was a greenish light appearing about halfway down the test tube. The light moved up and down the tube as the heating was increased and decreased in intensity, probably representing the rate of production of flammable vapor vs. its interaction with the atmosphere. After admiring the green glow for a few minutes, I broke off the experiment.

There was something else that appeared in the tube: a reddish coating on the glass above the area with the lead and acid. It looks like read led, so I believe that some of the lead had volatilized and oxidized in the tube.

I found this experiment interesting because the phosphorus production took place entirely at or below a bright red heat, below the melting point of borosilicate glass, while previous experiments rendered the glass unusably soft before showing any signs of success. I realize that test-tube experiments of this sort will never lead to useful production but I find them interesting anyway.

Edit: after I let the tube cool and I broke open the end with the solidified acid and lead residue, I smelled a peculiar phosphorus odor. With the lights out, I could see a faint glow on one particular chunk of residue, but there can't have been much of anything there since warming it in the light didn't reveal any smoke or other visual signs of phosphorus.

[Edited on 5-11-2005 by Polverone]

Esplosivo - 10-5-2005 at 21:24

evil_lurker, why not try heating the Pb3P2 formed with sulfur as was described by BromicAcid in the case of iron phosphide. Might give some results. Besides the reaction temperatures (excluding the last step of phosphine liberation which I am excluding) are surely within reach. If the container was sealed tightly enough a stream of dry hydrogen could be passed quite easily. I'd like to try using metal piping as in the case of Bromic, which seems to be quite leak proof. Reducing with hydrogen does look interesting. Btw, nice results Polverone, the more I hear about this faint glow the more I wish for the summer holidays to start :D

evil_lurker - 11-5-2005 at 10:59

Actually I plan on going down to the steel yard sometime soon and pick up some decent diameter pipe, say in the 4-6 inch neighborhood and some sheet. I can cut the pipe to the desired length, and simply weld on a cap on each end, and a iron pipe fitting on top for a crucible.

I do need to make some modifications to my "hellfire barrel" though. The last experiment in melting down an aluminum engine block was a success, using a 55 gallon drum as a furnace. The bottom of the drum had a 1-2 inch hole burnt in it for the molten aluminum to run out, the side had another hole burnt in it about 3 inches in diameter. In the 3 inch hole, a section of 2 inch copper pipe about 12 feet long was placed and hooked up to a shop vac to provide forced air. The bottom of the barrel was loded with dry oak wood in fairly decent sized chunks. The engine block was set upon this with the barrel top off, and lit with gasoline and a firecracker making a nice mushroom fire/ball cloud go up with a large "WHUMP". It took about an hour, but I melted some odd 70 pounds or so of aluminum leaving all sorts of steel engine parts in the barrel along with some uncombusted charcoal.

I know the inside had to hit at least 660 degrees to melt the aluminum, and probably hit the 800 required for the sodium hexametaphosphate, silicon dioxide, and aluminum reduction.

I did not have a lid on the barrel, so I lost quite a bit of heat coming out the top. I'd say the flames were approx 6 feet high. It would probably work better also if the air stream entering the barrel was at an angle rather than straight in.

One idea for a crucible that was considered was a small portable air tank. They are about $20, but the down side is they are a one use deal.

Another idea was to use several excesses of regular silica sand in a small crucible that has had the powdered sodium hexametaphosphate blended in and using natural gas (mostly CH4, methane) as a reductant to reduce the P205 formed into P4. The excess silica sand would create a P205 coated gas permeable "matrix" thus allowing better yields since the P4 vapor could more easily leave the crucible.

Experiment will show I suppose.

ChemicalBlackArts - 18-8-2005 at 09:18

Pshhh, why make any allotrope of elemental phosphorus? Phosphorus is much better put to use in VX! C11H26NO2PS

sparkgap - 18-8-2005 at 09:28

Because white phosphorus looks so damn cool and toxic. :P

Also makes for a good firestarter.

sparky (~_~)

P.S. I'd only play with VX if I had 10 vials each of pralidoxime and atropine on hand, thank you very much. :P

praseodym - 18-8-2005 at 22:35

Different allotropes of phosphorus could be used for different uses. For example, white phosphorus is used in military applications as incendiary bombs, for smoke-screening as smoke pots and smoke bombs, and in tracer ammunition, while red phosphorus is essential for manufacturing matchbook strikers, flares, and, most notoriously, methamphetamine.

Basically, phosphorus exists in three allotropic forms: white, red, and black. Other allotropic forms may exist. The most common are red and white phosphorus, both of which consist of networks of tetrahedrally arranged groups of four phosphorus atoms. The tetrahedra of white phosphorus form separate groups; the tetrahedra of red phosphorus are linked into chains. White phosphorus burns on contact with air and on exposure to heat or light.

Phosphorus also exists in kinetically and thermodynamically favored forms. They are separated by a transition temperature of -3.8 °C. One is known as the "alpha" form, the other "beta". Red phosphorus is comparatively stable and sublimes at a vapor pressure of 1 atm at 170 °C but burns from impact or frictional heating. A black phosphorus allotrope exists which has a structure similar to graphite – the atoms are arranged in hexagonal sheet layers and will conduct electricity.

tramp - 25-8-2005 at 07:20

white phosphorus dissolves readily in carbon disulfide...

a favorite game of mine in high school (this was several decades ago...) was to dissolve WP in CS2, then fling the liquid on a wall in the dark... it was cool to watch the wall begin glowing green then errupt in flames...

budullewraagh - 2-9-2005 at 05:33

but there are many other allotropes of phosphorus. please see attachment.

i have heard of blue and orange phosphorus as well, although they, like yellow, may not be true allotropes.

a key to the chart:

1] High vapor pressure at room temperatures, [2] heat at 540 C,

[3] heat at 550 C,

[4] heat at 600 C,

[5] heat at 125 C,

[6] heat at 400 C,

[7] heat at 550 C,

[8] heat at 300 C at 8000 atm,

[9] heat at 380 C with Hg or above 250 C at 12 kb,

[10] heat at 400 C with Hg for days,

[11] heat at 200 C at 12000 atm,

[12] heat at 200 C at 15000 atm,

[13] heat at 200 C at 12000 atm,

[14] reversible trasition 50-100 kb,

[15] reversible transition 110 kb,

[16] recrystallize from molten Pb,

[17] heat a PBr3 solution,

[18] reversible transition at 900 C,

[19] reversible transition at 1700 C,

[20] reversible transition at low pressure,

[21] reversible transition at 44.1 C (but can supercool),

[22] reversible transition at -77C or +64 under 1200 atm,

[23] sublime under vacuum,

[24] heat at 220 C at 12 kb,

[25] irradiate with UV at -190 C,

[26] condensation of P2 vapor at -196 C,

[27] heat above -100 C,

[28] heat at low pressure,

[29] boils at 280 C,

[30] heat at 300 C or expose to light or X-rays,

[31] melt about 600 C

P_allotropes.jpg - 45kB

BromicAcid - 2-9-2005 at 07:45

Yeah, but this discussion should be in the thread on phosphorus allotropes, as this table is already there as is the discussion of the different allotropes.

WP and KClO3 detonation

Lambda - 2-9-2005 at 16:31

Quote:
Originally posted by tramp
white phosphorus dissolves readily in carbon disulfide...

a favorite game of mine in high school (this was several decades ago...) was to dissolve WP in CS2, then fling the liquid on a wall in the dark... it was cool to watch the wall begin glowing green then errupt in flames...

If you drop a few drips of this solution (WP in CS2) on a small pile of Potasium chlorate (KClO3), wait a wile until a portion of the CS2 has evaporated, a hugh detonation will take place.

CAUTION: This experiment is very dangerous, for a powerful detonation takes place, with a big bright flame.

How is this for high school fun ?:D

neutrino - 2-9-2005 at 17:13

It spontaneously catches fire? I wonder where the activation energy comes from.

woelen - 3-9-2005 at 10:04

White P is very reactive and on a warm day in summer it can catch fire without heating. I recently received some white phosphorus and I was quite scared when I moved it from one container to another. As soon as I took it from the water a white smoke was released and I put it under water in the other container as fast as I could.
The reason for this reactivity is the strain in the P4 molecules. It is like a spring, which is severely stressed and which is about to break. Only little agitation is needed to break the spring.

BTW, red P and KClO3 also can ignite spontaneously. I put some KClO3 on the ground and added some red P, with the intent to mix them while on the ground, but due to wind, the chems were scattered over a larger area. I took a little stick and wanted to scrape the chems on a little heap. As soon as I touched the chems, a bright white flame was given, some molten KClO3 was sprayed around (also some on my hand :() and all stuff was gone. Fortunately the amounts were small, not more than 100 mg total. I was really surprised, because the chems were not even mixed, there was just a layer of red P on top of a thin layer of KClO3.

[Edited on 3-9-2005 by woelen]

neutrino - 3-9-2005 at 10:53

I must have missed that W, sorry.

12AX7 - 3-9-2005 at 12:39

P + KClO3 is well known to be rather dangerously unstable, after all. :P

Tim (has put a crystal of KClO3 between two match strike pads and struck with hammer)

neutrino - 3-9-2005 at 17:15

Yes, I'm well versed in the processes of natural selection. :P

BromicAcid - 3-9-2005 at 17:45

Every post since ChemicalBlackArts' sarcastic post, posts have been straying from the topic at hand, preparation of elemental phosphorous. Indeed, conversion of the white allotrope to the others and the properties of the white allotrope are indeed things that could be covered under this heading, but it seems to me the discussion of these things would be more sutible for some of the other threads hanging around here, let's save this thread for actual experimentation and ideas for the preparation of elemental phosphrous and keep these other discussions in other threads more suited for this such as:

what's the usual way to make phosphorus? (white, then red from white, then black?)
Interesting allotropes, glassifications, and polymorphs.
Experiments with the luminosity of white P
Unusual phosphorus allotropes from red P
Favourite Pnictogen

And there are of course threads on potassium chlorate/phosphorus mixtures, something most of us here know as a more dangerous mixture then we want to mess with, and there are still other threads on accidents, there are better places for this little off shoot of the topic that has budded here.

Edit: Didn't realize this would be on a new page, let's try to make this page more reflective of pages 1-5 rather then page 6, okay :)

[Edited on 9/4/2005 by BromicAcid]

phossyburn3.jpg - 94kB

garage chemist - 5-9-2005 at 00:26

Here's a thread where someone actually managed to make a considerable amount of white phosphorus:
http://www.scienceforums.net/forums/showthread.php?t=3083&am...

The trick is to use metaphosphoric acid (oh how useful this substance turns out to be, it will become one of my most used chems very soon! You can make oleum from it, and white phosphorus!) with charcoal.
The reaction happens at a temperature easily reachable with a bunsen burner and can be conducted in a test tube.
Stuff the mouth of the test tube with glass wool, otherwise the P will burn away as it forms.

A method to actually collect the phosphorus in solid form is to bend the test tube in the middle. The upper part of the test tube iss arranged horizontally for the white P to collect as a puddle and solidify upon cooling.
The reaction mix is contained in the lower, bent part of the test tube.

The necessary metaphosphoric acid can be made by strongly heating ammonium phosphate (available as fertilizer).

neutrino - 5-9-2005 at 03:17

I got a different impression from that thread. It seemed like what was being described was a quick-and-dirty synthesis with ordinary phosphoric acid (note the part about dissolving the decomposition product of (NH<sub>4</sub>;)<sub>2</sub>HPO<sub>4</sub> in water), absurdly high temperatures (the burner melted the test tube), and a small yield (.1g).

BromicAcid - 5-9-2005 at 08:12

I also got a different impression then that thread was giving a method for making useable amounts of phosphrous. Just like earlier in this thread with Flayer who also attempted this, some of us were eventually concerned for him due to this method producing phosphine which is quite toxic, in that thread as well it is mentioned that phosphine is produced.

It is a decent method, maybe dehydrating the phosphoric acid to metaphosphate first then grinding and mixing with carbon might help rather then just mixing with phosphoric acid initally and dehydrating them both together. I can personally attest to the sodium hexametaphosphate and aluminum working well for this. Maybe all that individual needed was a better apparatus rather then a test tube, the only problem being that the procedure uses the free acid and it wouldn't be good to do this in anything short of a somewhat inert metal or glass, the iron works that I normally utilize would be out of the question.

Polverone - 6-9-2005 at 01:25

Inspired by YT2095's posts, I tried test-tube phosphorus production again with charcoal. I added a couple mL of 75% H3PO4 to a test tube. Initial heating produced some strong bumping so I added a bit of silica powder, and this maintained constant placid bubbling. In order to avoid heating too quickly, I started out with an alcohol lamp. After some minutes of heating the liquid began growing thicker and translucent, and over a time of less than a minute slackened and then ceased to bubble altogether. It thickened to the point of solidification in the tube, and application of the alcohol flame could not make it flow again. Is this metaphosphoric acid? I have been unable to find a melting point reference for metaphosphoric acid. I understand that it may be hard to isolate as a pure substance, but I wish I could find at least a ballpark figure to guess whether I'm seeing metaphosphoric acid or some strange glass/acid reaction product.

I added a small amount of charcoal powder on top of the slug of presumed metaphosphoric acid, wetted it with a couple drops of H3PO4, and fired up the propane torch and began heating the tube again.

The liquid acid seemed to permit the charcoal and plug of presumed metaphosphoric acid to mix. There was bubbling and the glass rapidly softened. Bubbles of gas started forcing their way through the glass. I maintained heating for several minutes, until the bottom of the test tube was thoroughly deformed and pierced by bubbles. I was working in semi-darkness, but never saw the greenish glow of phosphorus appearing in the tube as I did with my reduction experiments using lead. Neither did I ever smell the garlic-like odor I associate with phosphorus production and combustion.

After the glass had melted through, I was able to direct the bright blue cone of the flame against the presumed metaphosphoric acid solidified at the bottom of the tube. This produced a reflected flame tinged light yellow/white and green at the edges, such as I associate with phosphorus, but of course this is no step toward useful production.

Given that I encountered the same difficulties with phosphoric acid/charcoal reaction as I have seen before, that the pyrex tube was destroyed well before there was any noticeable phosphorus production, and that every reference I've ever seen on the topic indicates that carbon reduction of phosphoric acid or phosphates is significant only at temperatures well above the softening point of pyrex, I am very curious how YT2095 managed to produce 100 mg of elemental P from a test tube and propane burner setup.

Thinking perhaps there was some magic to his exact procedure, I combined charcoal and phosphoric acid as he suggested, heating the beaker until I had a just-slightly-damp mixture like moist corn meal. This was loaded into another test tube and again heated ferociously with a propane torch, until the mass was solid, the tube was horribly deformed and glowing bright orange, and I was tired of waiting for results (several minutes). I caught only the barest whiff of what might be the telltale odor, and that only for a short time, and did not see the characteristic glow in the tube that I saw with my experiments where I used metallic lead.

I fear YT2095 is omitting some key details or is talking about things that he hasn't actually done, though I hesitate to challenge the integrity of a man with a postcount of over 8000.

Some details that make me doubtful:

He's asking the forum at large how he can make phosphorus in his first post, but less than 2 weeks later he's claiming success in producing phosphorus. Beginner's luck? An unusually fast and skilled experimenter despite his original unfamiliarity with the topic?

His first instructions are
Quote:
add 25% of (the phosphoric acid you just made) mass with carbon (lumpwood chacoal powdered will do). mix well and allow to got as dry as possible.

step 4: put a little of this mix in a heatproof test tube and heat it over a bunsen burner (in a dark room is best). you will see Phosphorus begin to distill up the tube and maybe catch fire with a green glow when it hits the air

Now anybody who's actually tried this or related attempts at home phosphorus knows that a standard Bunsen burner is insufficient for carbon-reduction phosphorus production. His rather vague description puts me in mind of someone trying to fake his way through a procedure he's only imagined executing.

Later he said he managed to save 100 mg of red phosphorus, produced from a larger stock of WP that mostly went up in flames. He never described how he performed the allotropic transformation, while it's actually worthy of some discussion.

Yet later he said:
Quote:
did you leave the carbon and the conc phos acid to stand a while? I left mine for just over 24 hours ontop of a room heater to make sure it was completely dry and intimately mixed as absorbed into the pores of the carbon, then gently heated it over a flame after to drive off any excess moisture, as soon as there was the garlic type smell I capped the tube with a bleeder pipe, THEN you turn up the heat, the carbon will start to Glow dull red at 1`st then bright orange, leave it like for a good five minutes

It would be quite a feat to begin producing phosphine (the "garlic type smell";) at only a few hundred degrees, below a red heat.

But the #1 fakery tipoff: he says "Sayo I have the pics all leading up to this event, just not on my website yet as I wanted the complete set, you`ll forgive me for not taking pics during the fire I hope!" on 3/13/2004, but he doesn't post after that date until someone bumps the thread in August 2005. Wouldn't everyone here post pictures of their work if they had actually made usable quantities of phosphorus, started a thread crowing about it, and already taken the photographs?

Let's recap: vague descriptions of the processes, no photographic evidence posted though he claimed to have it more than a year ago, and so-far irreproducible results that seemingly contradict the experiences of many others (he wasn't the first to think of charcoal and H3PO4 in a test tube over a burner, but he claims to be the first to get useful results this way).

Now the opposing view, the "maybe he's not bullshitting" view. He said that he actually used a homemade forced air gas burner rather than a Bunsen burner, which could deliver quite a bit more heat. I would expact the test tube glass to lose structural integrity before any usable quantity of phosphorus would be produced, but perhaps a core of carbon and condensed phosphoric acids supported it. He mentioned making his phosphoric acid by double decomposition of fertilizers, followed by heating the produced ammonium hydrogen phosphate until it ceases to smell of ammonia. Is it possible that some impurities were retained that allow his phosphoric acid to be reduced by charcoal with unusual ease? I don't think so, but that's the best I can come up with on short notice.

Given that he started another thread to say he believed himself to be close to a complete theory of the Big Bang and the universe, I'm inclined to believe that he is given to flights of fancy, has never made any macroscopic quantity of phosphorus, and certainly doesn't have 100 mg of self-made red phosphorus in his collection.

I will contritely retract these words if he or anyone else can post instructions for <b>reproducibly</b> duplicating his claimed phosphorus production. Charcoal and acid in a test tube at orange heat do virtually nothing. Where's the key innovation that makes it do something useful?

Edit: garage chemist, where did you get the details about bending the test tube and stuffing its neck with glass wool? I didn't see those details in the thread you linked to. Are those your ideas about successfully implementing the procedure?

[Edited on 9-6-2005 by Polverone]

garage chemist - 6-9-2005 at 02:34

Bending the test tube is my process for collecting lager quantities of white phosphorus (3 grams), which I make by sublimation of commercially bought red phosphorus (a non-controlled substance in germany, still difficult to get though).
The white P collects as a liquid puddle and solidifies to a slab after complete cooling. It can then be extracted with pliers.

I did not think of the possibility that YT2095 may have faked his success and told of things that he wasn't actually making.
But now that you outlined the flaws in his report, this becomes a possibility.

I have the impression that we need a fundamentally different experimental procedure than "mix a few substances, heat them as high as you can and hope for phosphorus being produced".

Especially the reactor design (glass or metal pipes heated from the outside) doesn't seem that covenient.
We should abandon external heating and try to switch to internal heating, like it's done in industry.
Using an electric arc would be my approach if I would try to make white P without using red P.

A metal reactor, maybe made from a pipe and two endcaps and with a narrow pipe for the P vapor to escape, has a carbon rod inserted from the top, electrically insulated of course. The reaction mix (carbon, phosphate and silica- not too finely powdered, so that not too much dust is produced) is placed in there and the reactor closed.
The pipe is heated from the outside to 300°C in order for the P to stay gaseous.
Then the current is switched on, I'm thinking of a MOT (2000V, 0,5A) as the power source. The gases are filtered from dust by leading them throug a heated glass wool filter. Then the P is condensed by leading the gases through warm water (for the P to stay liquid and not condense to fine flaky material which is hard to melt together- I know what I'm talking about).

Maybe this simulation of the industrial process works better than the ancient process from the times where electricity wasn't available as a heat source.

Lambda - 6-9-2005 at 08:05

Garage chemist, I strongly advise against using a MOT transformer. The reason is twofold:

1 - These MOT transformers are not designed to run continuously, even at the given specifications. These transformers (the modern consumer ones) are designed at borderline specifications to save on core dimensions (using less Silicium Iron), and especially expensive Copper wire. They soon run hot, and overheat.

2 - You want to create an Arc with the output power, and the resistance of an Arc is very low. The voltage drop is also very low, and thus will result in excessive power soak requirements in series with the Arc current. Assuming there will be a current of 0.5 ampere running through the Arc at a very low resistance, then the voltage drop (maybe < 20 Volts) will also be very low. This means that if you apply Ohm's law of U = I * R then you can calculate the power (P = I * U) to be P = 0.5 * ~20 = ~ 10 Watts heat input into your reaction mix meanwhile, the rest has to be soaked up (P = 0.5 * <2000 = <1000 Watts). Much less than 1000 watts, because it is a current transformer, and probably more than 10 Watts (the short circuit current may be higher than 0.5 Ampere).

It would be better to use an Arc-welding machine. These are also current transformers, only with the difference that the current is much and much higher (<160 Amps). You then pump P = <160 * ~20 = <~3200 Watts of energy into your Arc. But here again, these welding transformers are not designed to run continuously. The cheap Aluminum wire wound hobby stuff, have a very low turn-on ratio, and even the (semi)professional and Copper wire wound one I have (TICO 160), has a turn-on ratio of only 25 %. But in this case, even if the Arc output power would be far less than 3200 Watts, it will still be a lot more efficient than a MOT transformer out of a Magnetron.

[Edited on 6-9-2005 by Lambda]

garage chemist - 6-9-2005 at 12:29

Lambda, you're right about a MOT having not enough amperage and not being designed to run continously.

However, an arc welder which might seem ideal at first has an important disadvantage: the arc needs to be started manually because of the low voltage.
Maybe a bridge wire between carbon electrode and metal bottom which burns through when the power is switched on can start the arc.
Or maybe the charcoal powder is conductive enough to allow enough current to flow through the mixture for heating it high enough. If not, an admixture of graphite powder might be the solution.

If the problem of starting the arc inside the chamber can be solved, an arc welder would be perfect for production of phosphorus in a semi-industrial reactor configuration.

Lambda - 6-9-2005 at 13:20

Garage chemist, you can simply start the arc with a high voltage burst. You may need to place a third electrode though, and just use a simple Geiser or Heater Igniter (Coil induction or Piezoelectric). Maybe just shortly tipping the two Carbon rods together, and then pulling them apart may work out easier. Wants the arc has started, you can pull them quit a distance apart.

The biggest problem may still be, the good conductivity of fused salts. Excessive heating, vaporization and maybe even explosions may be the result of this. When these salts have molted, they become conductive. If one electrode is placed in the molten salt, then the other electrode may be used to create an arc towards the molten salt surface, not touching it though. This may be done in a ceramic (heat resistant) vessel or maybe even a Steal pipe (with one Ceramic Electrode isolator).

Please check these threads out for castable Ceramics:

Kiln progress:
https://sciencemadness.org/talk/viewthread.php?tid=4381

Reaching higher temperatures:
https://sciencemadness.org/talk/viewthread.php?tid=4376

[Edited on 6-9-2005 by Lambda]

A few more odds and ends

Polverone - 7-9-2005 at 00:01

I had a lot of acid-impregnated charcoal left after my last experiments. So I maintained heating all night. Today it was slightly less damp than yesterday, but not much. Anyway, I tried again with the charcoal, this time grinding in some silica powder. There was a little bit of phosphorus production. It wasn't much, but it was steady as long as I kept the base of the tube at a bright orange. I can't imagine how many hours it would take to get 100 mg out of a test tube like this, even if I could collect the phosphorus instead of burning it.

I tried mixing zinc powder and then zinc powder plus silica powder with the acid charcoal. Both of these reactions went very poorly. I didn't notice any increased production of phosphorus; in fact I couldn't see any production at all since the zinc volatilized and left opaque oxide coatings on the inside of the tube, but neither could I see any white smoke in the light, so I don't think much P was being produced. Some zinc phosphide was formed, evidenced by the scent observed upon adding hot water to the cooled tubes.

Aluminum worked much better. In the first attempt, I placed cut-up pieces of a soda can's pull tab in the bottom of a test tube and poured the acid charcoal over it. This showed the most rapid and easy production of phosphorus, giving a healthy green combustion front racing up the tube as soon as the bottom reached red heat. The rate of reaction slackened considerably after that first burst, but it was still considerable compared to my earlier efforts. All of the successful reactions leave a white ring (presumably of phosphorus oxides) at the point in the tube where the combustion front spends most of its time; this one's white ring had some visible thickness by the time I was done. I scraped it with a bamboo skewer and the residue seemed to absorb water from the air. This showed an acid reaction with litmus (unsurprising).

For the final reaction I ground 400 mesh aluminum (the only sort of particulate aluminum I have) with the acid charcoal and loaded it into a test tube. There was some exotherm and funny smells even before I applied heat. I ran a very small batch, less than 1 gram of mixture, because I was wary of what might happen in the event of a violent reaction or accidental tube break. The reaction actually seemed harder to initiate than the one using chopped-up soda can bits. It never got as vigorous either, but it did all right.

I don't think any of these reactions showed marked improvements over the ones I did a few months ago with metallic lead and phosphoric acid.

I will probably continue to try things with my acid charcoal, but this path (and all those using free acid) looks unpromising compared to sodium hexametaphosphate plus aluminum. Now I just need to get some sodium hexametaphosphate.

Of course, even if I can make tongues of flaming phosphorus shoot from a test tube, I can't recover it very easily. But one step at a time, eh?

[Edited on 9-7-2005 by Polverone]

BromicAcid - 9-9-2005 at 20:19

A while back I posted an idea to obtain phosphorus by the basic method of electrorefining copper metal. I.e. putting two electrodes of phosphor-bronze (avalible for welding) into a solution with some H<sub>2</sub>SO<sub>4</sub> and preform electrolysis, the copper going into solution from the anode and plating out on the cathode, leaving the phosphorus particles in solution. Today however when looking up metaphosphoric acid I came across this somewhat relevent entry:
Quote:
Anodic oxidation of copper phosphide yields up to 60% of the dissolved phosphorus as hypophosphate (A. Rosenheim and J. Pinsker, Ber. 1910, 43, 2003).
So maybe this wouldn't be the best method in the world as I didn't anticipate anodic oxidation to be an issue here. Although phosphor-bronze doesn't contain much phosphorus it may be useful, personally I think more hope would lie in the nickel platings that do not require electrolysis to produce and other high-phosphorus nickel alloys (which are fast becoming popular, 8-15% phosphorus in some!) these may prove useful for making phosphorus and again they may not and only be useful for preparing hypophosphoric acid and the like.

jimmyboy - 10-9-2005 at 10:13

pyrex test tubes simply wont do for this i think - all the texts i read use steel or a clay retort of some sort and say 1400 - 1600 degrees C - so i guess an acetylene torch and a small fireclay container might work - maybe even try mixes of thermite - carbon and acid phosphate or maybe use the thermite to heat it in some fashion - set it off with long strip of magnesium - i remember someone making bits of sodium this way somewhere

S.C. Wack - 17-9-2005 at 17:49

Bromic scanned a page of Gmelin's earlier, but I decided to make a pdf of the same page anyways, and also the page before (red P) and after (purification). Because I already had the photocopies, so I might as well make them available, FWIW.

Attachment: P.pdf (583kB)
This file has been downloaded 1320 times


garage chemist - 18-9-2005 at 02:41

Interesting passages from the text:

"Hydrogen, when led over P2O5, HPO3 or alkali- or earthalkaliphosphates at 800 to 900°C, contains phosphorus vapor which can be condensed to white or red P."

"When a melt of NaPO3 is electrolyzed with a nickel cathode, phosphorus is depositing there."

"When an evaporated leachate of bone ash with H2SO4 is heated with charcoal in a porcelain tube, P evolution begins at 740°C, the largest part of P goes over at 960°C and at 1170°C a 92% yield is obtained."

"NaPO3 produced by melting NH4NaHPO4 is mixed with Al powder and heated. Already at red heat, the mass begins to glow and emit P vapors. Other phosphate salts can also be used, even the Ca and Mg salts."

jimmyboy - 27-9-2005 at 22:06

I found something else pretty interesting - it may not be too much use to any of us - unless someone knows jack about ultrasonics

METHOD OF SONOCHEMICAL REDUCTION OF PHOSPHATE SALTS
Valentin V. Oshchapovsky, Serikjan D. Kulbekov
1
, Victoria V. Shylo
2
Lviv State Institute of Fire Safety, P.O.Box 10676, Lvov, 79000, Ukraine
E-mail: asu@franko.lviv.ua and oshchapovsky@yahoo.com
1
Chimkent Corporation “Phosphorus”, Chimkent, 486025, Kazakhstan
2
Lviv National University, 1, Universitetska Str., Lvov, 79000, Ukraine
Reduction of phosphate salts is the high-temperature process. To lower the temperature of process
we have proposed to reduce the phosphates ores in ultrasonic field [1]. The reduction of Vyatsko-
Kamsky phosphorite was carried out in solid-phase with the application of natural gas as reducer.
We have studied the influence of temperature, process duration, frequency of ultrasound oscillations
and other factors on the degree of phosphorus extraction. The experiments were conducted in the
electric furnace with the ultrasound generator. The experiments showed that in the ultrasonic field
the process of phosphate ion reduction and extraction of phosphorus out of phosphates accelerated
considerably. The temperature of the beginning of the process dropped to 300-350°. A marked
extraction of phosphorus (12.5%) was fixed already at 700° C. Depending on the frequency of US-
oscillations at 900° C the degree of phosphorus extraction reached 90 % and more. Thus, the
application of US-oscillations intensifies the process of phosphate salts reduction considerably and
lowers the temperature of the process.
[1] Kulbekov S.D.; Oshchapovsky V.V. et al., Author Svid. of USSR No 1369176, 1987.
Submitted by:
Associate Professor V.V. Oshchapovsky, Department of Fundamental Disciplines, Lviv State
Institute of Fire Safety, P.O.Box 10676, Lvov, 79000, Ukraine
Phone: 38-0322-33-02-02; Fax: 38-0322-33-00-88.
E-mail: asu@franko.lviv.ua and oshchapovsky@yahoo.com and ndr@lipb.lviv.ua
Preferred Scientific Session / Presentation:
7D. Novel phosphorus chemistry / Poster
Signed:
Date: 02 December 2003
Page 2

Kman100 - 7-10-2005 at 18:09

Hey Bromic,

I've obtained some Iron Phosphide and I am considering performing the reaction that you mentioned earlier on in this topic.

Do you have any further information on this reaction? I'm looking for specifics...

Also, as for the reaction vessel, what do you recommend? What is the latest status on your P production? What is the current vessel you are using?

If I decide to perform this, I'll post pictures and progress.

Thanks all!

BromicAcid - 7-10-2005 at 18:37

I really wish I knew more, I don't know if I posted what I comprehensively know word for word so here it is, from Comprehensive Tretise on Inorganic and Theoretical Chemistry
Quote:
E. Minary and R. Soudry prepared phosphorus from a mixture of iron phosphate and coke; and R. A. Brooman heated a mixture of silica, iron, coal, and calcium phosphate so as to form a fusible slag and iron phosphide. The latter when heated with sulfur, hydrogen sulfide, carbon disulfide, etc., furinshed phosphorus.
Somewhat straight foreward but still, what about the reference?
Quote:
A. Nicolle, Brit. Pat. No. 1693, 1888 ; R. Lammy, ib., 311, 1857 ; J. H. Player, ib., 660, 1866 ; 1064, 1867 ; R. A. Brooman, ib., 2294, 1864 ; J. Townsend, ib., 1862, 1872
And what does that refer to? I don't know, I never learned much about references and when I took the book to my librarian and pointed at the reference they didn't know either (Mind you I was in the 'Reference Help Office' of the library!) I would really like some feedback on where I might find this reference, I was told by a librarian there may be a table that it refers back to but they couldn't find anything, I have a feeling ib means something simple and once that is discerned I will be able to find the reference (though I looked up the meaning online and it was no help.)

Anyway, heat the two reactants together, I don't know if excess S is added to drive the reaction foreward, because excess S would react with phosphorus. Also, I don't know if you need to distill the phosphorus from the mixture or if it pools on top. Finally since H<sub>2</sub>S is listed as a sulfur source it means that it doesn't have to occur in the liquid phase, it may be gaseous sulfur running over iron phosphide. None of these details are given. Given enough time one could calculate the thermodynamic values for the system to see at what temperatures it becomes spontaneous (though we would of course know nothing of the rate). One good idea if distillation seemed probable then you could run it between 280 C (boiling point of phosphorus) and under 444 C (the boiling point of the sulfur) with a stoichiometric or slightly less quantity of sulfur maybe?

Sulfur is corrosive to many metals in the liquid state so thick metal would be preferable, or its still within the relm of glass heating so it could be done in that medium. Hopefully more imput from other members will come and you will be sucessful in your efforts.

IrC - 7-10-2005 at 19:59

It looks like British patent number and year, as in patent number 1693 in the year 1888.

Speaking of: are any of these helpful?

4,204,925 Recovery of phosphorus from sludge
4,689,121 Recovery of phosphorus from sludge
3,113,839 Recovery of phosphorus from sludge
3,442,621 Phosphorus Production
2,050,796 Recovery of Phosphorus
4,462,973 Phosphorus purification
4,608,241 Production of phosphorus and phosphoric acid

ibidem

Magpie - 7-10-2005 at 20:28

Bromic Acid says:


Quote:

I have a feeling ib means something simple


I'm guessing this means ibidem, which means "in the same place." In the USA this is usually abreviated "ibid." It is, or was, commonly used to refer to a reference already cited. In your case it is just the reference for the 1st author.

S.C. Wack - 7-10-2005 at 20:44

Yep.

GB patents up to the 191?'s are found by using the year followed by a 5 digit number. The first would be gb188801693, if only the coverage on espacenet was not so spotty for the older GB patents. Some are there, others like this one aren't. IIRC the earliest GB patent that I've found there was from 1886.

I hate it when there are US patents for the corresponding GB or DE ones, but the US number isn't mentioned - supposedly all US patents are online.

IrC - 12-10-2005 at 13:20

BromicAcid, I recently nabbed two check valves on ebay and I wonder if they could be used. They are threaded for standard half inch plumbing pipe at both ends (inside the ID). They would screw perfectly into a section of pipe, are spring loaded and open at a couple PSI. The problem is the big ball inside of each is some type of plastic and I am not sure if they would melt or if there would be some reaction with the material. I do not know what the material is, I am half tempted to take one apart and see if I can find a perfect ball bearing to replace the plastic ball with, around 11 mm or so in diameter.

Going back to Polverone's beginning post on page one:

"Sadly, this reaction requires 1000-1500 C to operate. It's done in an arc furnace. Building a suitably airtight, nonconducting, refractory vessel for an arc furnace is something well beyond my current engineering skills/resources (any brilliant suggestions?)"

Why does the vessel have to be an insulator? If you had say a 9 inch diameter 11 inch tall graphite crucible with one inch thick walls (and bottom) as the grounded electrode (the whole crucible itself), could you not use castable refractory to create a gastight lid which also holds the second termimal (say a one inch diameter graphite rod going down into the crucible stopping an inch or so before the bottom)? The refractory would be the insulation to keep the center electrode from contacting the outer electrode (i.e., the crucible itself). A suitable hole could be cast in the refractory top where the exit pipe came out.

Sitting here looking at this exact crucible I had this idea. If you do not ground the crucible, another cast refractory piece could be made which holds the crucible keeping it insulated from the stand, assuming you are using a floating source of power such as a two bushing pole transformer (also sitting here looking at one of those).

[Edited on 13-10-2005 by IrC]

leu - 13-10-2005 at 04:48

Quote:
I really wish I knew more, I don't know if I posted what I comprehensively know word for word so here it is, from Comprehensive Tretise on Inorganic and Theoretical Chemistry

quote:E. Minary and R. Soudry prepared phosphorus from a mixture of iron phosphate and coke; and R. A. Brooman heated a mixture of silica, iron, coal, and calcium phosphate so as to form a fusible slag and iron phosphide. The latter when heated with sulfur, hydrogen sulfide, carbon disulfide, etc., furinshed phosphorus.

Somewhat straight foreward but still, what about the reference?

quote:A. Nicolle, Brit. Pat. No. 1693, 1888 ; R. Lammy, ib., 311, 1857 ; J. H. Player, ib., 660, 1866 ; 1064, 1867 ; R. A. Brooman, ib., 2294, 1864 ; J. Townsend, ib., 1862, 1872

And what does that refer to? I don't know, I never learned much about references and when I took the book to my librarian and pointed at the reference they didn't know either (Mind you I was in the 'Reference Help Office' of the library!) I would really like some feedback on where I might find this reference, I was told by a librarian there may be a table that it refers back to but they couldn't find anything, I have a feeling ib means something simple and once that is discerned I will be able to find the reference (though I looked up the meaning online and it was no help.)


Mellor's citations were written in the ancient style; clearly the reference librarian didn't have enough grey hairs; the citations are on pp 743-4; and are:

Minary & Soudry

Rev Univ Mines, (1), 16. 360, 1864

Brooman

Ann Chim Phys 2294, 1864

The latter article may be available from Gallica; but two citations previously searched for from that publication given in Mellor's were erroneous, whether this is due to errors made by the author, typesetter or proofreader is unknown :(

jimwig - 13-10-2005 at 08:19

wp using microwaves.
hey sounds cute
will post patent number as soon as I find it - its somewhere around here I just know it.

WOWWIE - I R A HAZARD TO OTHERS. WHOOPPEEEEE- I was just getting used to being a hazard to my own damn self.

[Edited on 13-10-2005 by jimwig]

well my new title requires that I actually have to think about it some more - seems like an ion exchange resin was involved along with phosphoric acid - kinda vague but will not let my fellow officers and crew down - NOSIREEE!!!

(Jeez give a guy a little fucking title?)

[Edited on 13-10-2005 by jimwig]

Looking For Phoshorus?

HomeAlchemy - 26-10-2005 at 12:08

I am amazed that it is so hard to find Phosphorus in the US!

It is very frustrating trying to do neat experiments at home -- it sounds pretty harsh in the US:(

It's not that easy to find in Canada, but at least you are allowed to own some.

If you are looking for a little to play with just e-mail me @ darien88@hotmail.com

I also have many other neat, hard to get exotic chemical elements, metals, and non-metals if you are interested

I have shipped this stuff to the USA without problems. Just a few friends who needed some.:cool:

Dr. Beaker - 3-12-2005 at 15:43

some general comments about high tem. decomposition of phosphates (also applicable for sulfates, CaSO4 reduction thread), and please correct me if I'm wrong:
I believe that the actual reduction occurs with the oxide (i.e P4O10 or SO3 and C, Al, ...) the very high temp. is needed for the decomposition of the very stable ions: 4[PO4(-3)] -> 4O(-2) + P4O10.
now, as the polarizing power of the cation increase, the thermal stability of the salt decrease and it decomposes at lower temp. i.e AlPO4 should be a better choice then Ca3(PO4)2 and Fe III better then Fe II

NJF - 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

NJF - 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

Bone ash

hinz - 10-2-2006 at 11:42

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

12AX7 - 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.

Tim

Kiln

MadHatter - 10-2-2006 at 23:54

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.

garage chemist - 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.

Magpie - 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]

garage chemist - 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.).

skippy - 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.

common phosphorous sources

skullandfeather - 14-3-2006 at 11:20

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

densest - 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.

12AX7 - 17-3-2006 at 01:20

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

Tim

densest - 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.

Cyrus - 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.

garage chemist - 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]

hinz - 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]

densest - 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?

12AX7 - 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.

Tim

Agamemnon82 - 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.

jimmyboy - 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?

12AX7 - 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.

Tim

12AX7 - 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

Tim

jimmyboy - 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]

12AX7 - 5-4-2006 at 13:44

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

Tim

Agamemnon82 - 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

 Pages:  1  2    4  ..  15