Sciencemadness Discussion Board

Preparation of elemental phosphorus

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Polverone - 26-6-2002 at 12:50

Historical and modern preparations of elemental phosphorus from phosphates are straightforward yet inconvenient for an amateur chemist because they require very high temperatures, provided by charcoal/coal fired furnaces in old methods and by electric arcs in newer ones.

In the 1800s and earlier, phosphorus was prepared by a number of processes. The earliest was that of Brandt, who prepared it from human urine and charcoal. Later methods were variations on the theme of "heat bone ash with charcoal at high heat." The ashes of bones contain considerable phosphorus, combined with calcium and oxygen, that can be reduced to the elemental state with enough heat and a suitable reducing agent (carbon).

I will summarize the method given in Muspratt:

Animal bones are strongly heated in air until all organic matter has been destroyed. The bones are powdered and to every 3 parts of powder are added 2 parts concentrated sulfuric acid and 16 to 18 parts water. This causes some of the calcium in the bones to be converted to calcium sulfate, which can then be removed by decantation/filtration.

The liquid thus obtained is evaporated to a thick, sirupy consistency. It is mixed with one-fourth its weight of charcoal powder, and then it is raised to near-red heat to make it perfectly dry. The mass is transferred into a stoneware or iron retort in a furnace. The retort has a copper tube connected to a heated underwater receiver where the phosphorus can condense without oxidation (and without solidifying and blocking the tube); the gases that bubble to the surface are sent back to the chimney of the furnace by a second, smaller copper tube. The furnace temperature is gradually increased to white heat.

First comes off steam, then hydrogen and carbon monoxide and dioxide, and finally, at bright red heat, phosphorus begins to come over, accompanied by phosphine and CO/CO2 (it is difficult to be sure about some of the exact products because of the archaic and sometimes inaccurate terms used in the text).

Muspratt does not say precisely how long all of this takes, but Wagner's Chemical Technology (1872) indicates that heat was maintained for a long time, up to 48 hours.

And here's a summary of the modern method:

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

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?)

I have, a couple of times, tried straightforward phosphate reduction with charcoal and heat. The vessel is a steel pipe with a screwed-on cap at one end and a screwed-on nipple at the other. The nipple has a section of copper pipe inserted in it and sealed with furnace cement. I filled the pipe with a mixture of diammonium phosphate and charcoal on the (admittedly dubious) premise that the ammonium salt would have a lower decomposition temperature and might help the reaction along. Plus it was the only pure phosphate I could find on short notice.

I heated the apparatus with a large gas laboratory burner and had a vessel of warm water to dip the copper pipe into. On my first attempt, I got a lot of strange/unpleasant smelling gases and water condensation at first (I wasn't going to submerge the tube until I was sure something interesting would come out). I also saw some gas leakage around the threads on the pipe. When it looked like the reaction wasn't going anywhere, I removed the nipple/copper tube assembly. I then heated the tube some more just for curiosity's sake. Toward the end I started to see something interesting. The mixture was melting and bubbling out of the tube a bit. I could heat this portion directly with the gas burner, and when I got it red hot I started to see a rather distinctive flame come out of voids in the material. It looked like the flames I'd gotten by igniting red phosphorus (obtained from match box strike strips) and it had the same smell.

After that minor encouragement I figured I'd clean things out and try again, more patiently this time. However, it turns out that whatever hot diammonium phosphate and charcoal turn into, it is hard, insoluble, and tenacious. I had to painstakingly chip/smash slag out of the pipe with a metal rod.

On my second attempt, much later, I kept the copper tube underwater the whole time and tried to be patient with the heating. The gas burner took a while (15 minutes?) to heat the tube up to red heat, and even then could maintain that heat only where the tube directly contacted the flame. It never got hotter than a medium-red. There was a considerable amount of junk deposited in the water - mostly copper salts created by hot/moist exit gases - but no phosphorus that I could see.

I have not yet made a third attempt. I have recently been considering (along with madscientist) alternate ways of preparing phosphorus, ones that might use cleverness to get around the high heat requirement.

Electrochemical methods? I haven't found any, nor do I particularly expect to. The one wacky idea I had was to try electrolyzing phosphides in non-aqueous media (since they react with water) but phosphides are hardly household chemicals and phosphorus is nonconductive, so it would soon block passage of current if it were deposited.

Another wacky idea was to try electrolyzing molten phosphates with carbon electrodes and hope that the intimate contact of phosphate ions and carbon would promote reduction, but this would require temperatures and apparatus similar to the arc furnace.

More active reducing agents? Perhaps aluminum or magnesium powder could be substituted for carbon. I don't know if this would be safe or practical. Would it yield elemental phosphorus or magnesium/aluminum phosphide? Again, I'm not sure.

I AM sure that I want to crack this problem. Phosphorus and its compounds are invaluable reagents for a variety of laboratory procedures, and it's also just plain nifty. Five years ago I could have just ordered some from a chemical supplier, but I didn't have that foresight, and now it cannot be obtained in the United States without a DEA permit.

Such a simple substance, and such simple chemistry, that should pose such challenges...

[Edited on 26-6-2002 by Polverone]

An interesting patent...

Polverone - 28-6-2002 at 11:52

Wow, don't everybody try to answer at once. Anyway, I've just found a very interesting document: US Patent 6207024 (accessible through www.uspto.gov) describes the production of elemental phosphorus from pyrolytic carbon and phosphoric acid, using microwaves to heat the reactants. The reaction takes place at much lower temperatures than the conventional arc furnace process. The problem (or problems), of course, is that the phosphorus still needs to be protected from oxidation, you need a relatively heat-resistant and microwave-transparent reaction vessel, and it's going to be tough to condense and collect the phosphorus if you're trying to come up with something using a domestic microwave oven. Microwave ovens are cheap at thrift stores, but there are still obstacles to overcome for performing this reaction at home...

vulture - 8-8-2002 at 08:05

This is not entirely on the topic but hey.

I found a few mole gassing pyrotechnics.
You light them and put em in the hole where they produce highly toxic PH3.

They contain Ca3(PO4)2 and Al. In order to release PH3, Ca3P2 has to be formed by the following reaction:

3Ca3(PO4)2 + 16Al -> 3Ca3P2 + 8Al2O3

I wonder if one could produce elemental phosphorus from calcium phosphide?
Maybe by substitution with a halogen?

Ca3P2 + 3Cl2 -> 3CaCl2 + 2P

Although somehow this doesn't *feels* right.

Substitution with a non-metal with a lower electronegativity seems more promising, but then you have another problem: How would one get elemental Silicon, Germanium, Antimony or Tellurium?

madscientist - 24-8-2002 at 16:50

Vulture mentioned in another thread that CaC2 would make a good reducing agent in such a reduction as phosphate reduction.

I suspect that sodium polysulfide would work well in a phosphate reduction.

4Na3PO4 + 2Na2S2 ----> 8Na2O + 4SO2 + P4

Calcium phosphide would probably be useful. It probably could be electrolyzed in, say, acetone, with graphite electrodes, which would yield calcium metal (useful) and white phosphorous (holy).

Wait a minute...

Polverone - 24-8-2002 at 21:35

I thought I posted information here on using aluminum and silicon dioxide for low-temp reduction of phosphates. I guess that was on other forums, though. Ahh, yes, here it is, from the E&W Forum:

Oooh, I like that idea Vulture! And here's another gem.

Donald J. Haarmann was kind enough to post this on Usenet in sci.chem as part of a very long post on thermites and Goldschmidt reactions:

Quote:

Aluminium as a Reducing Agent, &c.
L. Franck. Chem. Zeit. 1898, 22, [25], 236-245.
In:- The Journal Of The Society Of Chemical Industry. 17, [6], 612-613.
June 30,1898

Action of Aluminium on Phosphorus Compounds—Phosphorus vapour when led over powdered aluminium, heated to a dull red beat in a current of hydrogen, combines with it with incandescence, forming a dark greyish-black unfused mass, which is decomposed in contact with moist (normal) air, forming PH3, and leaving a greyish-white powder. It is decomposed also by water, aluminium also by water, aluminium hydroxide and a brownish-black residue being left ; and by acids and alkalis, which dissolve it almost completely with evolution of PH3. The compound remains unaltered when heated in air.

At more or less elevated temperatures, all phosphoric, acid compounds (meta-, pyro-, and ortho-salts alike) are decomposed by aluminium. Metaphosphates, however, undergo the most complete change, according to the equation—

6NaP03 + 15AI = 6NaAl02 + 2Al203 + Al5P3 + P3

The addition of silica effects the release of the remaining phosphorus, thus :—

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

Calcium and magnesium salts are as efficacious as those of sodium, but the superphosphates of commerce are not available for the production of phosphorus in this manner. If, however, bone ash be decomposed by hydrochloric acid instead of by sulphuric acid, a material suitable for the purpose is obtained.

Hence phosphorus may be produced, with almost quantitative completeness of yield, at relatively low temperatures...


Now the real question is: what is meant by a "relatively low temperature?" One of these days I may just have to find out!

I hit

Polverone - 24-8-2002 at 21:43

I meant to address this also.

Quote:
Calcium phosphide would probably be useful. It probably could be electrolyzed in, say, acetone, with graphite electrodes, which would yield calcium metal (useful) and white phosphorous (holy).


Mmmm, the holy white phosphorus! Do not approach the glory too closely or you may be hurt! But I think you are jumping a little ahead of yourself. Calcium phosphide may or may not be soluble in acetone, and phosphorus is a non-conductor, so it may be a problem keeping the current flowing even if it is soluble.

Oh, and as long as we're on this topic, in that E&W Forum thread that I've copied part of here, PrimoPyro mentioned that
Quote:
Also, phosphates can be reduced to elemental phosphorus with catalytic hydrogen around 750C (upper limit) and lower down to around 350C.


He didn't provide any more information than that. Care to enlighten us, Primo?

I think the most straightforward method is to build a charcoal furnace and do it the old-fashioned way, but I don't have a furnace nor the room for it. So let's bring on the wacky Rube Goldberg methods!

Damn that post was hard to find!

PrimoPyro - 26-8-2002 at 22:52

I don't have any info to add to that. I referenced that from a post at The Hive. The user "UTFSE" posted it under post no. 210542, in the Stimulants forum, under the thread called, "The Philes on Phos...(so far)".

You should really read that thread, it contains a huge wealth of info and sources. The source for that statement was said to be:

"Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry" Vol 3 Supplement 3; Wiley 1971

Wow! $720 for a used copy! Damn! Check it at:

http://www.amazon.com/exec/obidos/tg/detail/-/0471593095/qid...

I'll message this user (haven't seen him in awhile, hope he's still around) and ask for a copy of the full reference, word for word.

PrimoPyro

vulture - 28-8-2002 at 09:07

Would that just be a simple high temperature reduction with a hydrogenating catalyst like platinum?

Very likely

PrimoPyro - 28-8-2002 at 18:13

That is very likely. I am very very interested in nickel hydrogenolysis catalysts, especially Urushibara prepared nickels, as opposed to Raney Nickel, which is highly pyrophoric. (Jeeze, it can actually erupt into flames by simple filtration induced friction!)

PrimoPyro

Pyrophoric Nickel...

vulture - 29-8-2002 at 07:58

When I read that in my chemical book I was a bit stunned, you would only expect that with highly reactive metals like Mg.

Did you know Hydrogen autoignites in contact with very fine Pt powder?

I wonder if there is no chemical way to oxidize the P-3 to P from Ca3P2?

Red to White Phosporus

daryl - 21-9-2002 at 04:50

I succeeded in heating red phosporus in a CO2 atmosphere and distilling the phosporus vapour which condensed to white phosporus and looked like candle wax.

Nifty

Polverone - 21-9-2002 at 09:36

I would try the same thing if I had ready access to red phosphorus. However, thanks to the War on Drugs, it is now impossible for somebody who is not licensed by the DEA to buy elemental phosphorus in the United States. Actually, it's illegal to make and own as well. Nevertheless, I am always interested in possible routes to this fascinating element. Phosphorus and its derivatives have so much use in the laboratory that I can't resign myself to saying "oh, that's off-limits."

a_bab - 21-9-2002 at 14:09

It seems that I'm a lucky bastard as I have both red and white P. And it's really nice to play with these. Not talking about Na, K, Li, Br.

daryl - 21-9-2002 at 14:14

bummer, I haven't recently tried to get anything exciting. I suspect it's harder nowadays. In Victoria, they had a crackdown on 'chemical' waste and storage as we've had a few fires over the years. Also there is more caution on chemicals related to explosives such as Ammonium nitrate from garden suppliers etc.

Ramiel - 22-9-2002 at 01:07

I would advise caution in the distillation of phosphorus... especially white phos.

Since it is toxic and u can get a nasty case of phosphorus jaw [jaw slowly rots away]. Not nice eh?

daryl - 22-9-2002 at 17:21

I distilled it in an atmosphere of CO2 and it doesn't burn as long as the CO2 is provided. No real hazard. The P4 just condenses in a test tube in ice water. Because it's always cold it doesn't burn.

pROcon - 10-10-2002 at 05:01

Hey polv!("186" has) missed ya, didn't know you had a forum. Our 19th century muspratt conversation got cut short so I couldnt reply to your note at 'The [bought] Hell' (speculating..uh). I wasn't going to use your email at your college site just in case...

It sounds like you got close with the first attempt, I was under the impression it was a failure altogether from what you had told me, but well done. I'd be very cautious about the Al route, check this out [http://www.uow.edu.au/eng/bmh/explosion.htm] I stumbled on it accidently while trying to source Al powder in australia. The hydrogen gas too.

This diammonium phosphate you played with decomposses at 155C right? You can make it work you already have had success it just needs fine tuning.
I want you to consider the following for your pipe experiment, I think it could get you there.

- Buy a few small pieces of ceramic fibre board. its VERY easily cut. Arrange them like so
here's a cross-section of labtop CF arrangement:
________
| |
| |

use it as an insulating enclosure for your tube, you'll get fantastic temperautures, certaininly in excess of 200C with your torch across the tube. Remember you can easily cut holes in this fibre with a kitchen knife. So you could make a box with it, with a hole at either end to support tube, and a hole for the flame entry.

You know the chemistry of this more than me, so now solve the problem of how to get it under water without oxidizing first. Brainstorm: a bend on the nipple end of tube as it exits the pipe, a longer tube. Also remember the charcoal never reached a temp that would have removed the oxygen would it?

On that note, I'm going to zap the reactants first in microwave to rid of any moisture.
MONOPOTASSIUM PHOSPHATE (Tradename: MKP) KH2PO4)
decomposition point for this one slightly higher at :~250C
This will mean that your C will have reached a higher energy level when its ready, and more of a glow, less oxygen.

Are you sure you want to go the arc route? Already got the tools for it except for refractory cement? What about a cheap electric kiln? they have a vent hole at top and front door. This my intended route,

The refractory vessel shape is basic, I know you can work that out!

Either an upright cone with a flat hem ("lips" sitting on a casserole pot with lips around rim)
|
/\
/ \
= =
|_|
Bad ascii art I know but thats option one.

Option 2 is more ingenius, and I thought it up (heh.. surprising I know)

...cut a kinda tall triangle out of a piece of paper, slice it diagonally from one lateral to the counter-lateral. Now rotate the top piece on the slice axis... the rest is the same as above with lips and casserol [retort].
It should appear like the design above but with the point of the triangle now coming out of the side -- I call it the Muspratt 2002 PDT.
Cone construction in a flash:
http://www.flyingacesclub.net/volare/cone.htm

Construction :
Scrapped the stoneware pot idea. Just make a nice accurate carboard version, pour refractory cement over each of the two pieces in a thoughtful fashion, thats what I'm leaning towards, itll work, ill burn the cardboard off.

Email me if you like polve I have digital images of mine ... and measurements, of everything, and kiln photos if your interested in that option. The tubing could be a winner though!
Have you had time to re-scan that page of muspratt containing the reciever yet? Maybe ill check out your site and see if its updated, you could well have...
The pressure black flow on cooling needs thought.

I have a few questions that may or may not need to be answered (I've forgotten them so I dont know LOL), I must be in a giving mood.

Recent finds I've bookmarked on this topic since we last spoke ...

Fun stuff ...
http://boyles.sdsmt.edu/phosox/oxidation_of_phosphorus.htm

On aluminum as reducing agent.
http://cdl.library.cornell.edu/gifcache/moa/manu/manu0026/00...

More descriptive version including a labtop... in tubing.
http://cdl.library.cornell.edu/gifcache/moa/manu/manu0026/00...


http://cdl.library.cornell.edu is top notch by the way for OLD articles and reference spanning 1815 - 1926, kept it with you in mind two days prior to now.


This post was longer than I intended, but hopefully you'll have a decent read. ttysih.

pRO :)


Hello again, and GOOD INFO!

Polverone - 11-10-2002 at 11:46

It's nice to see you here. When you suggest ceramic fiber board as an insulating material, did you really mean "you'll get fantastic temperautures, certaininly in excess of 200C with your torch across the tube"? Because I can top 200C with a hotplate... I'm hoping you meant to add another zero in there.

Anyway, the problem isn't decomposing the diammonium phosphate so much as it is reducing the decomposition product. I do believe I produced some small amount of phosphorus near the end of my run, but as I've stated in previous posts, I couldn't heat the whole assembly enough to make this method viable. Perhaps with better insulation or a better heat source I could.

But it doesn't look like I even need to do that! That document you turned up on the Cornell site is great. I'll reproduce the text here for folks who don't want to download a gigantic GIF.

Quote:
Aluminum for the Preparation of Phosphorus

The applications of aluminum in the arts multiply with much the same rapidity as do those of electricity. The Berichte describes a new method of preparing phosphorus by its use as a reducing agent. The process is so simple that it can easily be illustrated on the lecture table. Hydrogen ammonium sodium phosphate is fused in a porcelain crucible until it is changed into sodium metaphosphate; aluminum turnings are then dropped into the liquid, and the freed phosphorus bursts into flame. Now, if the experiment is tried with a glass tube, instead of a crucible, a slow current of hydrogen being passed over the mixture of the salt and aluminum, the phosphorus distills into the cooler part of the tube without the formation of any phosphoretted hydrogen. The residue consists of alumina, sodium aluminate, and a phosphide of alumina - Al2P2.

By these steps in the process only 30 per cent of the phosphorus in the mineral used can be obtained; but the phosphide is decomposed entirely by heating it with silica, and this may be added at the beginning of the experiment and the reaction proceeds without difficulty and without loss.

It is advised that for the lecture table a combustion tube a yard long be used; two and a half parts of aluminum, six parts of sodium metaphosphate (obtained from heating previously the hydrogen ammonium sodium phosphate) and two parts of finely pulverized silica are placed in the tube, a slow current of hydrogen is passed through, and heat is applied until the reaction begins. This is shown by sudden incandescence, and phosphorus is seen to condense in globules on the cooler part of the tube, at the end where hydrogen escapes.

Instead of this phosphate, any ordinary phosphate may be used, but experimenters are warned not to use the superphosphates containing calcium sulphate mixed with them, such as are used for fertilizing purposes, because the sulphate is suddenly decomposed by the aluminum with an explosion when a certain temperature is reached.


I doubt that any of us really want to have to use hydrogen gas connected to a burning-hot tube. I would presume, however, that any non-oxidizing gas would serve as a carrier. Nitrogen, argon, helium, and maybe even propane and natural gas might be used in place of hydrogen. Or it may be possible to distill the phosphorus with no carrier gas at all, simply vaporizing it and condensing it in water. However, I would be very careful of suckback with this method, because you certainly don't want a steam explosion in the midst of molten salts and phosphorus.

Really nice news from the above quoted material is that A) the aluminum need not be finely divided; if turnings work, you can use sweepings from a machine shop! and B) silica improves yields greatly but isn't absolutely required. In any case, I think that fine ceramic grade silica powder would be the ticket here.

Be careful when cleaning up your residue if you try this. Aluminum phosphide will, I think, yield phosphine gas on contact with water, and phosphine is very flammable and very poisonous.

Another useful tidbit

Polverone - 12-10-2002 at 15:09

I was searching for phosphates with relatively low melting points. I found that sodium hexametaphosphate melts at 550 C and is used as a water softening agent, so it should not be at all difficult or suspicious to acquire. Diammonium phosphate, the material I previously used for my phosphorus preparation attempts, also has a relatively low melting point, though it partially decomposes as it melts, into what products I'm not sure. The diammonium phosphate was obtained as a nutrient for use in beer/wine brewing.

Hoffmann-LaRoche - 12-10-2002 at 15:52

polverone

Did u already think of using a mix of phosphates?

You could make a mix with another phosphate as this would further decrease the melting point.

The same is used in industry for melt electrolysis in the production of aluminum and alkali metals.
The fact that impure solid substances have lower melting points, is well known in organic chemistry, where it is used for testing the purity of synthesized substances.
Pure substances have a sharp melting point and usually higher melting points than impure.

HLR

some updated info

vulture - 13-10-2002 at 02:29

When creating Ca3P2 by reducing calciumtriphosphate with Al, the fused mass of Ca3P2 and Al2O3 can't be separated. This leaves only oxidation from PH3 as an option, which is extremely dangerous, because of it's auto-inflammability and it's high toxicity.

Maybe Zn or Mg powder would work?

Something to keep in mind

vulture - 13-10-2002 at 02:34

When directly trying to isolate phosphorus from phosphates or P2O5, you'll have to reduce it, but when isolating form phosphides you'll have to oxidize.

I'm confident that one of the two will show as the most viable method.

Polverone - 13-10-2002 at 10:50

Hoffman: I did think of using a eutectic mixture of phosphates, but I didn't really want to because I'm sure there's some minimum temperature at which the reaction proceeds, and I can easily achieve temperatures up to 800 C or so.

Vulture: Did you see the information from the Cornell site? You'll get a lot of phosphides if you just mix the aluminum and phosphate, but powdered silica will decompose the phosphide to yield phosphorus.

Yesterday, despite needing to do so many other things, I just had to see what I could achieve with a molten phosphate + aluminum. I melted diammonium phosphate in a stainless steel dish with a propane burner until the decomposition bubbling slowed down. Then I added a ball of shredded aluminum foil and mixed it into the viscous liquid in the dish. As the temperature rose, flammable bubbles began popping to the surface and igniting. A steady stream of flammable gas/vapor was produced for a couple of minutes as the foil got hot enough to melt.

I also tried mixing a small quantity of 300 mesh aluminum powder into the melt. After a delay of about half a minute, the whole region where the powder had gone experienced a sudden incandescent ignition. I couldn't tell if this reaction generated further flammable gas because of the intensity of the light. After a few more minutes of heating I turned the heat off and stirred the thick mass with a spoon so that it became crumbly, loose chunks as it cooled (instead of one huge mass stuck in the bottom of the dish).

There were many small bright droplets of unreacted aluminum left in the mix after cooling. This morning, the mixture had adsorbed some water from the air and smelled a bit like garlic (phosphine forming?)

I would like to try this with sodium hexametaphosphate instead because it was hard to tell, with this last attempt, whether I was producing phosphorus vapor or hydrogen. By eliminating all hydrogen sources, a second run could be much more instructive. Also, I think the reaction may go more easily with a sodium compound as the aluminum forms sodium aluminate with it. I would of course *also* like to try adding powdered silica to the mix and seeing how that affects the reaction. Finally, I'd like to try using aluminum turnings/filings instead of powder and foil, since I think they may be more suitable forms (more vigorous than foil, but not dangerously so).

Anyway, I hope that someone else can try this out, as I am extremely busy and I don't know the next time I'll get to make an attempt at this.

This is a possiblity......

Boob Raider - 15-10-2002 at 19:51

Instead of using Na, NH4, etc... PO4's I was thinking of using Pb3(PO4)2, made by using the pH downer in Hydroponics stores and PbO. In the patent I came accross, they used Pb3(PO4)2 prepared from the oxide, and reduced by H2 at about 400*C. The products were P4, Pb and H2O. I guess the same thing can be applied with C as the reducant. Maybe the temp would need to be a little higher, the mix would yeild P4 in a fire place, but I think its worth a shot, I need to make a Cu flask though.

That's a possibility...

Polverone - 16-10-2002 at 00:06

I think the main difficulty, though, is freeing the phosphorus from oxygen, not from whatever cation it's attached to. I could be wrong. I would certainly be interested in hearing of your results in any case. Consider trying aluminum foil/filings/turnings as well, and (if you can) also try adding powdered silica, as it seems to be beneficial in all of the phosphorus-producing methods I have read of...

madscientist - 18-10-2002 at 16:45

It mustn't be too difficult to reduce phosphates with metals - this is what my chemistry book says about it.

The pure acid is a colorless crystalline solid (mp 42.35C). It is very stable and has essentially no oxidizing properties below 350-400C. At elevated temperatures it is fairly reactive toward metals, which reduce it, and it will attack quartz.

It seems that the phosphate anion is reduced by metals at temperatures over 400C.

Marvin - 19-10-2002 at 17:27

Someone suggested CaP might be electrolysed to Ca metal and P. This sounds somewhat unlikley, is acetone even an ionising solvent?

Phosphate salts + silica is only the way to go for much higher temperatures, avoid salts alltogether, and you dont have to melt the silica or even use it at all.

Orthophosphoric acid is easy to reduce, but not the goal. Orthophosphoric acid and ammonium phosphates on strong heating pass through the pyrophosphate stage to the metaphosphoric stage. This is the annoying glassy solid you cant dissolve easily. Empirical formula HPO3 its actually a polymer. The smart money is on reducing this I'm almost sure, less hydrogen means less P lost as phosphine. Theres no cation to form a phosphide, so the only way youd get one is using excess aluminium.

Aluminium sounds rather excessive for this anyway, the reaction might be specitaculaly energetic, maybe a mixture of a small amount of aluminium with a larger amount of carbon would be pratical, the excess thermal energy from the aluminium forcing the mixture to a temperature where the carbon is effective.


Polverone - 19-10-2002 at 21:08

You're right that calcium phosphide can't be electrolyzed in acetone (or any other solvent for that matter). Acetone will dissolve some ionic compounds. I had discussed an old article I found on the production of lithium metal by electrolysis of anhydrous solutions of lithium chloride in various solvents, acetone among them. I believe that's what prompted the idea that one might electrolyze calcium phosphide in the same fashion.

As mentioned a bit upthread, I've already tried using aluminum with diammonium phosphate. Aluminum foil was placid, even sluggish - very manageable. Aluminum dust was much more vigorous; I wouldn't want to try heating large quantities of it with phosphates. The idea of using a mixture of carbon and aluminum sounds good. And, of course, I've never gotten around to trying Vulture's suggestion of using calcium carbide to perform the reduction. One more thing to put on my to-do list...

aluminium and silica

koffee - 30-1-2003 at 02:47

earlier in the thread it says any phosphate can be used for the aluminium reduction, and cilica may be added for a more complete reduction, but heres a hang up i have about that:

If you used a phosphate like mono-calcium phosphate, wouldnt that require much higher temperatures anyway? (defeating the purpose of using the Al for lower reduction temperature)

Same with the silica, wouldnt this require higher temps too?

Would the addition of silica, stop phosphite from being a product? I hope so personally! i know the "carrier gas" would stop any steam from coming in to contact with it, assuming its condensed in to water as polverone suggested. But I'd be more concerned with the carrier gas blowing the phosphite product in to the water, even though it is only a gentle flow of gas.


KABOOOM(pyrojustforfun) - 1-2-2003 at 19:03

can too much heat decompose a phosphide of an unreactive metal?
for example 2Cu3P2 => 6Cu + P4

madscientist - 1-2-2003 at 21:45

Great idea! I think that at the very least the following will occur, considering that cupric chloride decomposes relatively easily into cuprous chloride and chlorine:

4Cu3P2 ----> 4Cu3P + P4

But...

Nick F - 2-2-2003 at 07:06

How would you make the phosphide? I can't think of a good method without elemental phosphorus...

madscientist - 2-2-2003 at 08:54

Igniting a composition of aluminum powder and copper phosphate. If the stoichemitry is done right, the aluminum should be fully oxidized, leaving cupric phosphide, which should decompose as it is formed due to the heat previously liberated by oxidation of aluminum.

Good plan.

Nick F - 2-2-2003 at 09:33

That's the kind of thing I've been thinking about, a method that doesn't require an external heat source once the reaction has started. I was just going to try various Al/C/(HPO3)n mixtures, to see if I could get a self-sustaining reaction that wasn't too vigorous. I can't at the moment see any advantages of using copper phosphate over metaphosphoric acid, but maybe experimentation will prove me wrong.
I couldn't find pure ammonium phosphate in any local shops, only mixtures with sulphates and other crap, so I've bought some by mail order. It should be here within a week, so I'll be able to experiment soon. Also on order is some clay, to make a specially designed cell for making sodium. It's not too special, but would keep air out and help to keep the products seperate. I'll talk about it in the right topic when I actually have something to talk about...

Ramiel - 3-2-2003 at 00:20

I don't know if this will be of any use, bt while I was browsing about the internet, i chanced upon a little article.

caution

KABOOOM(pyrojustforfun) - 4-2-2003 at 19:29

copper phosphide Cu3P2.
Properties: Grayish-black, metallic powder. Insoluble in water; soluble in nitric acid; insoluble in hydrochloric acid. Sp. gr. 6.67.
Derivation: By heating copper and phosphorouse.
Hazard: Dangerous; spontaneously flammable and toxic phosphine is evoloved on reaction with water. May explode when mixed with KNO3.
Use: Manufacturing phosphor-bronze.

trinitrotoluene - 30-3-2003 at 17:24

In a book called "Lange's Handbook of Chemistry" somewhere in the book it says "The chief sources of the element is phosphate rock, generally Ca3(PO4)2, which is treated with sulfuric acid to convert it to phosphoric acid which is then heated with carbon, the crude phosphorus being removed by distillation."

My plans

Flayer - 10-4-2003 at 01:40

On the Sunday I plannig to prepare P from H3PO4 + C = H20 + CO + P. :)
I have real 600-650C in my glove :D

Good luck
Flayer
Sorry for my bad English.:(

trinitrotoluene - 25-4-2003 at 13:41

Maybe I will try this method to make the oxide of phosphorus.
Ca3(PO4)2+3 SiO2= 3 CaSiO3+P2O5.
This does need alot of heat. My idea is CaSiO3 have a melting point of 1540*C and P2O5 sublimbs at 340*C I may be able to distill off the P2O5.

Then the isolateing of P is P2O5+5 C=2P+5 CO2.
The problen is white phosphorus is very reactve and may explosve in contact with air.
And I doubt theres much of a yield of P4. I calculated that out of 310grams of Ca3(PO4)2 you will get around 15 grams of P4 as yield.
I don't really think creating the oxide is too much of a challange. But reducing the oxide and provideing a inert enviorenemt is th biggest challange.

Theoretic - 24-7-2003 at 08:00

Hmmm... I think that whatever phosphate you have, it's a good idea to convert it to H3PO4 by boiling in conc. H2SO4, then convert it to trisodium phosphate by reacting it (a melt is better than solution - - H3PO4 melts at 44C) with CALCINATED soda - water is highly undesirable.
It's melting point is 73 - 77 C.:)

Theoretic - 24-7-2003 at 09:13

And the equation for the reaction:
3Na3PO4+5Al+2SiO2=3P+5NaAlO2+ +2Na2SiO3
The sodium salts left over could be used as either alkalis (the corresponding insoluble oxide precipitates - sometimes as a gel:() or as sources of the corresponing chemically active oxide.

maybe I have to revise my statement

Organikum - 24-7-2003 at 21:39

that the Castner-Tiegel is the best way for electrolytic sodium production.

The following is quoted from a post a certain "Polverone" made at the HIVE whereby he quotes somebody he correspondended with in usenet. A quote of quotes so to say. ;)
Quote:
As for the sodium, I regret to inform you that I did not make it in my own castner cell I used the most ingenious and safe way that I could find, melt sodium nitrate, obtain a closed container made of soda lime glass and seal it with a wire touching the inside, use that as the electrode where the sodium forms and the ions will migrate though the glass and sodium will be deposited inside, it actually works reasonably well, although more recently I just buy my sodium from eBay.

I always had another opinion on the conductivity of glass, but there maybe a trick. No temperatures nor voltages or currents given. Somebody ever heard of this?

[Edited on 25-7-2003 by Organikum]

Polverone - 27-7-2003 at 21:41

I've heard of the method before on Usenet... it's been mentioned as appearing in kids' science fair projects. The person who I corresponded with on these topics is now a member of this forum... care to speak up, You Know Who You Are?

vulture - 28-7-2003 at 08:45

Why wouldn't it work? You don't necessarily need two wires, although I know it's hardwired into our brains that you do. The electrode on the inside is one path for the current, while the migration of ions through the glass is another path for the current.

Your body cells, for example, use the potential difference created by ion pumps to perform biochemical work.

There are many electrodes for quantitative analysis that use the migration of ions through glass. Especially sodium ions can be exchanged on glass surfaces.

[Edited on 28-7-2003 by vulture]

BromicAcid - 28-7-2003 at 16:54

Okay, I'll speak up, I didn't want to post in response because this is the Phosphorus topic, not the Sodium but I will. The basis for the procedure came from an off hand mention to it's production mentioned in an eBook on glassware manipulation, it can be found at Rhodium, the text relevant is as follows:
[ Diffusion through Glass
The mobility of sodium atoms in a soda-lime-silica glass at elevated temperatures is fairly high; if an evacuated bulb of such a glass is dipped into molten sodium nitrate and electrolysis is brought about by bombarding the inside of the bulb with electrons, the circuit being completed with an electrode in the sodium nitrate, then metallic sodium appears in the bulb. By immersing the bulb in other molten salts the sodium ions can be replaced by ions of silver, copper, thallium, and vanadium. ]
The name of the book, etc., were not copied so I cannot give credit where it is deserved.
I am fairly new to electrolysis so there was a lot of trial and error, I have no volt meter so Amps were not measured, I carried out experiments with a big 9V battery for a lantern and some wall adapters, greater then the 9V and some other side reactions seemed to dominate like NO2 formation otherwise it appeared to be N2 and O2 being formed. The glass MUST be thin, I got soda-lime glass from eBay and my first try was an antique bottle, way too thick. Also the vessel must be evacuated, it helps to pull the sodium in I guess keeps the equilibrium favored, and the wire must touch the inside of the bulb below the liquid sodium nitrate line but not at the bottom, if sodium pools around the wire the reaction will go to something else. It doesn't make a whole lot of sodium at a time 5 - 20 g but I'm sure someone could tinker with it and get more approximated values and increase yields, etc. The sodium is liquid at this temperature.

On the topic of phosphorus, just out of curiosity I looked it up in a book of electrode potentials to see if there was an electrolysis procedure for elemental phosphorus manufacture. As I just stated my electrochemistry needs some tweaking, here is what I found.

Electrically conducting phase P(red)
Solvent Water
Temperature 25C
Pressure 760 Torr
Electrode reaction HPO4(2-) + 7H(+) + 5e(-) <---> P + 4H2O
Standard Value -0.288V

There are numerous listed but all the electrically conducting phases are either red phosphorus or white phosphorus and all the solvents are water, how would such a reaction be run?

BromicAcid - 26-12-2003 at 17:50

Wow, I was the last one to post here, and off topic. Anyway, I don't know much about alloys but I know that they use a copper-phosphourus alloy in welding.

Often copper is refined by using a copper cathode and an impure copper anode, the anode dissolves in solution and replates on the cathaode and the sludge of impurites ends up below it in a puddle.

I know that when making red phosphourus by mixing white with molten lead and heating for several days you take the left over lead and electrolyze it similarly to revcover the bulk of the red phosphourus formed.

This copper-phosphorus alloy is readily availible, so would it be possible to isolate the phosphorus in this way?

unionised - 3-1-2004 at 06:59

If you can get hold of a phosphide and acidify it you get PH3. This will decompose on heating to give P4.
Reduction of phosphate by a reactive metal should give the phosphide.
Sulphur reacts with phosphorus so I doubt polysulphide would work as a reductant

[Edited on 3-1-2004 by unionised]

KABOOOM(pyrojustforfun) - 5-1-2004 at 20:11

veeeeery interesting unionised. can you give more details? specially the dec temperature. does it completely decompose to P<sub>4</sub> and hydrogen or some higher phophines form too?

Marvin - 5-1-2004 at 23:00

BromicAcid,

Interestesing post about sodium, it matches what Ive read with a few alterations though....

One book it is mentioned in is Shakashiri's Chemical demonstrations. A bulb is partially immersed in molten sodium nitrate, its powered from an isolating transformer and 2kv is placed between the filement and the container which can be AC or DC the right way round. The hot filment turns the bulb into a vacuum tube and the current flowing through the glass (sodium ions) produces in a short time a mirror coating to the inside of the bulb.

Seeing as it needed 2kv I dismissed the idea, I'm very curios how you got it to work with only 9v and with such large amounts of sodium produced......

Aparently a lot of other ions will replace sodium, but the glass tends to get very brittle and shatters.

The phosphorous by electrolysis looks very strange. I was under the impression Red P or White P didnt conduct electricity. I dont follow either why the acid would be reduced and not the water....
I think perhaps the reaction was measured in reverse to provide a value and that the reaction cannot be run preperativly.

[Edited on 6-1-2004 by Marvin]

BromicAcid - 14-1-2004 at 15:25

For some reason I'm always thinking of phosphorus.... I was going through my notes today (The book is now over 150 pages long with tons of photocopies and handwritten things on 3 x 5 cards jammed in it) and came across:
Quote:

....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 sulphur, hydrogen sulphide, carbon disulphide, etc., furnished phosphorus.

Inorganic and Theoretical Chemistry: pg 740 (That's all I have on where it came from, neglected to write down the authors name on the photocopy:( )

All it has for the R. A. Brooman reference in the index was i.b., 2294, 1864 I wish I could track it down because it seems interesting in that it could be carried out by heating with sulfur, I assumed that any kind of sulfur reduction would result in sulfur-phosphorus compounds, I guess their low decomposition temperature got them and in addition such a reaction seems like it would run at a low temperature... well hopefully.

Also I dug up my references from Gmelin for "Labratory Preparation (of phosphorus)" I've looked it over, it's in German of course but I believe there is nothing new there, if someone wants to look at it I can scan it in, it covers Reduction of CaHPO4 with carbon, reduction under HCl or Cl2, Reduction of Hg6(PO4)2, and reduction of sodium hexametaphosphate by aluminum powder, by the way, sodium hexametaphosphate is made by decomposing NH4NaHPO4.

Finally, I talked with one of my teachers about electrolysis of aqueous solutions to yield phosphorus, he told me it can be done with degassed water at temperatures just above the melting point of white phosphorus. He said it melts off and sinks to the bottom, but when I asked which salts could be used he gave one of those "Well it was back when I was in school..." speeches and expected me to forgive him for not remembering, I felt like ringing his neck! :mad:

Marvin - 16-1-2004 at 09:37

I would like to see the Gmelin scanned, particulaly to do with the temperature of carbon reductions.

Your reference is from 'A comprehensive treatise on inorganic and theoretical chemistry' by JW Mellor.

The Brooman reference itself I think is a british patent number, so no luck there.

I will see if I can find info on electrolytic production of P, but its unlikley I have anything.

BromicAcid - 16-1-2004 at 11:20

Ask and ye shall receive... Scanned in Gmelin for the production of phosphorus.

(Sorry that it's a little big...)

[Edited on 1/16/2004 by BromicAcid]

BromicAcid - 9-2-2004 at 14:44

Quote:

If you can get hold of a phosphide and acidify it you get PH3. This will decompose on heating to give P4.

Funny thing, there was at one time an industrial process relying on this method. I ran across it today in the chemistry abstracts from the turn of the last century. Basically it was the reduction of calcium phosphate in the presence of excess iron and silica to yield iron phosphide which was then reacted with acid to yield the phosphine which was decomposed at high temperature. I will get the exact reference on Wednesday, although I really wouldn't want to work with phosphine, it is interesting.

Also when I was looking though the early chemistry abstracts today I repeatedly saw the regular reduction of calcium phosphate, usually with just silica and occasionally with silica and carbon and with silicon carbide. But... I also saw an iron phosphate catalyst being used, saying that its use lowered the reaction temperature, very interesting.... Maybe current literature neglects this because for the longest time now they have been using an arc furnace to make phosphorus and adding the catalyst only lowered the activation energy a slight amount.... to them! Heck, I'll use any advantage I can get.

Finally, my last find of the day. I found a reference for the production of phosphorus from electrolysis!!!

But... it's the electrolysis of bone ash in molten cryolite :( ... yeah... in case you don't know what's wrong with that I should point out the melting point of cryolite is about 1000 C, so you wouldn't save much in terms of how many BTU's you would need to make your phosphorus, and I really don't want to think about what's being evolved at the anode. Maybe there's an eutectic mixture that could be used, using the eutectic finder http://ras.material.tohoku.ac.jp/~molten/molten_eut_query1.p...
I found some that melted around 600 C, somewhat lower, but would they even work? The stuff is soluble in concentrated sulfuric, maybe dissolve it in 100% acid and as long as the ions are there it might facilitate electrolysis... I know, it's a stretch, but I'm still trying to figure out new ways to get to the phossy prize.

Polverone - 9-2-2004 at 15:54

I still think you had a great idea when you conceived of dissolving those phosphorus-copper alloys (used for making phosphor bronze) electrolytically. I know it isn't really a preparation, but it's the next best thing. I can't recall if cyanides complex copper strongly enough to dissolve the metal in aqueous solution... if they do, that could be an easier/faster way of dissolving considerable quantities of the alloy without simultaneously oxidizing the phosphorus.

BromicAcid - 9-2-2004 at 16:07

Here's another one about dissolving metals. The electo-less nickel plating is high in phosphorus, much higher then the phosphor bronze, like anywhere from 3 - 12% by weight. Not only can you buy it because the phosphorus adds desired properties but you could make it in an electro-less plating bath. Then you perform electrolysis between two pieces of nickel to purify them like you would to refine copper, the phosphorus would sink to the bottom of the anode, or if you did it industrially you would wrap the anode in a bag to prevent the phosphorus from getting out into the solution. I thought I already posted about that but I guess I just had it in my notes.... That's a good idea now that I reflect on it.... and as a by product you could nickel plate things, the more I think about it the better it sounds... Weird...

BromicAcid - 13-2-2004 at 11:20

Quote:

....particulaly to do with the temperature of carbon reductions.

I found a reference to reaction temperatures the other day for carbon. Straight calcium phosphate with carbon in the form of coke begins at 1000C and is complete by 1600C.

If someone were to go the method of production involving the high temperature decomposition of phosphine might it be possible to do a dissolving metal reduction of a phosphate to yield phosphine? I know that using antimonates and arsenates you get stibine and such so wouldn't phosphate produce phosphine?

Oh, and I hit the library again, went though the chemistry abstracts from vol 1 to vol 21, they were all covered in dust and looking through them gave me a sore throat. :D Anyway, I was getting really really sick of seeing almost every abstract for a new modification of reduction of calcium phosphate with either carbon, silicon dioxide, or aluminum oxide to yield some form of slag that could be used for concrete but I came across some interesting stuff.

Patent 1988387 has some interesting designs for phosphorus producing apparatuses. Reduction of phosphorite containing 1-2% moisture with a 10-15% excess carbon charge yields phosphorus in a nearly theoretical amount a 500 - 600C which is really low.

Reduction of bone ash with peat moss yields CaCN2 and Phosphorus or just the regular reduction with an excess of Carbon under a nitrogen atmosphere yields cyanimide.

[Edited on 2/13/2004 by BromicAcid]

It's the season of phosphorus!

BromicAcid - 24-4-2004 at 12:24

I decided that since I might be doing more experiments soon that I would post pictures of the evolution of my phosphorus production apparatuses so here they are in chronological order. (I remade all the apparatuses of parts I had handy so they are not the originals except for a few parts however they convey the concepts used in the designing of each incarnation)

My Oldest design

For this one I thought that clay would be the way to go. So I got one of those pots with the high magnesium content and I got one of the clay bottoms to sit it on. I drilled out the bottom to make it into a lid and put a ran a 3/4 inch pipe through it and put a flange on it and held that in place with screws. Then I drilled 4 holes evenly spaced along the edges. I bought 4 hooks with threads and put them in the holes. They hung down with the hook part at the bottom and the threads at the top. I put the pot in place and put the hooks under the bottom rim of the pot and tightened the bolts on the ends of the hooks to hold the pot in place. Within minutes of heating though the pot started to crack. Further heating produced more cracks and I tried to add fire place mortar for a repair but it was falling apart too quickly so I took it of the heat and never tried using pots again.

My First Design That Worked

I got results with this one. This is similar to the design that I used in my somewhat infamous description of phosphorus production gone wrong. With the exception that the exit was exactly at the same height as the bottom of the original container. Phosphorus was produced in the bottom of what is pictured here to be a propane can and from there the gaseous phosphorus traveled upward and into the downward pipe where it went down and into water to condense. This design was subject to burn through and corky phosphorus which I rectified on later designs by adding a piece of screen around the exit gas area on later designs. (That bolt in the back ground along the condenser pipe is just being used to prop up the pipe to take the picture.)

This One is More Stout

The whole of the reaction area is somewhat thicker metal then I was used to so the problem of burn through was more distant then before. Still the rest of the apparatus stayed the same. Except the vessel that it lead to full of water had a stopper on it and the only way for gasses to exit or enter was through the one-way gas valve that I had made (not pictured) this also prevented suckback which although not a point of previous experiments (the large diameter pipe prevented it somewhat) was still a concern from time to time. In addition I started to burn my exit gasses by placing a torch so that the flame was positioned over the exit point.

My Radical New Design

The charge that makes the phosphorus will be placed in the bottom of the pipe on the far left. That will be the section that is heated. The elbow on the bottom to the right will be filled with water a little past the bend to force the phosphorus rich gasses to bubble though it and condense out the phosphorus. Finally there is a one-way gas valve on the end to prevent air from reentering the system. Less can go wrong in this one in that it is all self contained. At the end open the end with the one-way valve and dump out the water probably full of colloidal phosphorus.

Now I just need to design my apparatus for the reduction of lead phosphate with hydrogen.

vulture - 24-4-2004 at 13:49

Quote:

If you can get hold of a phosphide and acidify it you get PH3. This will decompose on heating to give P4.


PH3 made this way is usually spontaneously flammable in air because of P2H6 content. Not to mention the toxicity of PH3.

In Belgium DIY store sell a rodenticide that uses Al (german dark it seems!) to reduce calciumtriphosphate to calciumphosphide, which then slowly reacts with moisture to form PH3.

I dissected one of these things, but didn't further investigate because the toxicity of PH3 is/was worrying me.

I might just go for P2O5 + Al.

a thought for your safety

Magpie - 24-4-2004 at 21:14

Don't get me wrong - I love pipefitting, but looking at your apparatus scares me -as you are depending on a check valve for a vent. If I was going to heat an enclosed pipe system I would feel a lot better about it if it included a rupture disk and/or a pressure indicator.

But it has to be one way....

Hermes_Trismegistus - 24-4-2004 at 23:33

All things considered Bromics' design is superb.

The gravity ball valve is contemporary technology and is the valve of choice for high pressure/hard vacuum.

And If I'm not mistaken the rupture disk you speak of could act like a burst diaphram. In this instance, the consequences of the vessel exploding really isn't that much more severe than having the contents splatter all over the inside of the phossy kiln.

The wisdom of the "Normal" safety procedures for pressure vessels get a little fuzzy when working with anthing more dangerous than your average pressurized vapours or gases.

In short, flying bits of metal puncturing major organs might actually be a plus when a man is running around in circles in his yard with his facial features running down his shirt.

Did you see the look on the Nazi's faces when they opened the Ark of the Covenant in Indiana Jones and the Raiders of the Lost Ark?

I'm certain they would have agreed.

In all seriousness, (not that gouts of White Phosphorus vapour condensing on your delicate fleshy parts is a joke) but Bromic has explained to me, his device in no small detail and I feel very confident in his ability to acheive success safely.
:cool::cool::cool:

Good Work

thecheman - 29-4-2004 at 06:42

Bromic Acid: Well done on your new apparatus! I have been contemplating making my own for a while now, but couldn't decide what I should make it from... What material did you use for the pipes? I'm assuming that it is standard plumbing (?) - the photos look somewhat like stainless steel... And what reaction did you end up using? Some phosphate + Al for reduction and sand? I have read that this reaction can be done without the sand, but it only yields about a third as much - what sort of yeild have you produced?
It would be interesting to see what happens with the hydrogen / lead phosphate reaction you mentioned. I have read somewhere about using 'catalytic hydrogen', this is a bit more advanced to the sort of reactions I'm used to... (I'll keep reading!)

Cheers

Not sand!

Polverone - 29-4-2004 at 12:27

You want very fine silica. Either go to a ceramics supply place and pick up some 300 mesh silica there, or make sodium silicate from your sand with NaOH, acidify the sodium silicate solution, and wash/dry the precipitate to yield clean SiO2 dust.

BromicAcid - 29-4-2004 at 13:13

Yes (playsand will not work unless brought to ungodly temperatures 1500C+), the finer the SiO2 you use the faster/better the reaction. The finest silica you can find is fumed silica. I bought some and 100g takes up nearly 500 cubic centimeters. Very very light powerdery stuff. The best reaciton that I used to produce phosphorus I mentioned further up thread:

6(NaPO3) + 10Al + 3SiO2 ---> 3Na2SiO3 + 5Al2O3 + 1 1/2P4

Zinc can also be used and I bet magnesium could be used but the reaction might get violent.

One other interesting reaction that I found that I may have mentioned before is R.A. Brooman's procedure, he heated a mixture of silica, iron, coal, and calcium phosphate, this formed a fuseable slag and iron phospide. He took the phosphide and heated with sulfur and distilled the phosphorus.

[Edited on 4/29/2004 by BromicAcid]

Excellent

thecheman - 1-5-2004 at 03:32

Thanks for the info, I'll give it a try soon...
Cheers

Arc furnace

thecheman - 3-5-2004 at 07:05

Just in case anyone is interested in building an Arc Furnace (ie temp 3000C+)

I found an article posted on the Popular Science site:

http://www.popsci.com/popsci/science/article/0,12543,611070,...

They give a brief description of using an arc welder connected onto graphite rods to yield extreme temps (if you want to melt tungsten, etc).

Perhaps it could be scaled down? Maybe put the electrodes into a mixture of powdered carbon + [the material you want to heat] ?

We already have easier ways to make phosphorus (thanks guys!!!), but maybe this will be helpful to someone...

calcium carbide

Magpie - 3-5-2004 at 18:11

Maybe this is the way to make some calcium carbide. I understand you need a furnace reaching 3000 deg C.

Theoretic - 4-5-2004 at 03:36

Hmmm... I think that the phosphide ion is an intermidiate in the phosphorus preparation. The aluminium/carbon/else strips the phosphate ion of oxide ions, and the electrons from the reductant transfer to the ion, this happens step by step untill a phosphide ion is obtained. This then comes into an equilibrium:

3PO4--- + 5P--- <=> 12O-- + 8P

The oxide ions are scavenged by silicon dioxide, which drives the equilibrium to the right. If no silica is present, the phosphorus produced immediately disproportionates, and eventually all the phosphate is reduced to phosphide so no equilibrium can establish.
The phosphide is the general product if a reductant is present, but no acidic oxide like SiO2. Production of calcium phosphide by carbon reduction works at 1200 C, while production of phosphorus requires 1400 C.

Just my theory so...

Edit: Hold on:

"Reduction of bone ash with peat moss yields CaCN2 and Phosphorus or just the regular reduction with an excess of Carbon under a nitrogen atmosphere yields cyanimide."

Ca3(PO4)2 + x => 3CaCN2 + 2P + 8H2O/4CO2 ?

Ca3(PO4)2 + 7C + 3N2 => 2P + 3CaCN2 + 4CO2 ? Some temperature you need for that...

Ca3(PO4)2 + Ca(CN)2 => 2CaCO3 + 2P + 2CaO + N2?

Ca3(PO4)2 + Ca(CN)2 + 2SiO2 => 2CaCO3 + 2P + 2CaO + N2?

[Edited on 4-5-2004 by Theoretic]

1600C+ furnace

thecheman - 5-5-2004 at 07:24

hmmmm.... just found a site with a home-made furnace capable of 3000F+ (ie 1600C+) from gas!

http://www.reil1.net/Furnace.shtml

May be useful for making phos and melting certain metals, too...

Lestat - 18-5-2004 at 18:49

This is in response to chemoleo's advice on the thread "nerve damage" suggesting as i had made WP to post here.

I have made about 5 grammes of WP a few days ago, by painstakingly scraping off the striking strips from safety match boxes. I filtered off the paper from the RP using nail varnish remover, then filtered the red P off.

I then placed a large amount of RP in a test tube, with a tube for distillation into cold water and condensed the WP.

This works fairly well, but a warning, Don't even let the smallest trace of this touch your skin, this happened to me, I accidentally got a small streak of white P on my right hand which I immediately scrubbed off. 2-3 days later and I cannot write at all and can hardly move my arm, due to great weakness, this also affects the grip of my hand too. Just be careful not to touch WP, even in the smallest quantities.

Proteios - 18-5-2004 at 21:48

Quote:
Originally posted by Lestat

I have made about 5 grammes of WP a few days ago, by painstakingly scraping off the striking strips from safety match boxes. I filtered off the paper from the RP using nail varnish remover, then filtered the red P off.


Daym that a clever synthesis, the more so given that the strips on safty match boxes contain a phosphorus sulfide AND NOT RED PHOSPHORUS.

whoopsie-daisy

Polverone - 18-5-2004 at 22:50

No, the strikers really do contain red phosphorus. If you burn one in, say, the inverted bottom of an aluminum soft drink can, and then rub your fingers in the sticky residue (ahh, healthy!), your fingers will smoke a bit and glow in the dark from the WP in the residue. The glow/smoke will be more intense if you rub your fingers together, at least until the phosphorus is consumed.

Proteios - 19-5-2004 at 13:49

i gotta admit i was talkin from memory of the uses of Phosphorus - charcogenide glasses in greenwood and earnshaw. Having had a scratch around the web Ive found stuff that says the striker is P4S3, red P and EVEN white P:- none of these are really concisely written, and i wouldnt trust any of these as much as G and E. However I may well have to eat humble pie on this one.

thunderfvck - 19-5-2004 at 13:59

Meth heads make their speed using a red P/I2 reduction of pseudoephedrine. The red phosphorous is obtained mainly from match strikers. Yields are poor. About 2 grams per 250 strikers, or something like that anyway. I've never done it, but I recall it being pretty sad.

Lestat - 19-5-2004 at 13:59

If you want to see it in very small amounts for yourself, just scrape the powder off about 5 boxes of safety match strips, put it into a neat little pile on top of a lid from a tin can and ignite, when the smoke has died down, knock away the ash and scrape the stain on the metal with a toothpick or sumthing, you will see sort of expanding rings of burning P as it gets consumed wrim the center.

Thats how i first did it, out of curiosity from the way the above said meth heads obtained their red P :P

[Edited on 19-5-2004 by Lestat]

Theoretic - 31-5-2004 at 09:26

"My plans

On the Sunday I plannig to prepare P from H3PO4 + C = H20 + CO + P.
I have real 600-650C in my glove

Good luck
Flayer
Sorry for my bad English."
by Flayer.
This was his first and last post at MSDB.
I'll tell you why.
When carbon reduces an oxygen atom off the H3PO4 molecule, H3PO3 results. That disproportionates quantitatively back into H3PO4 and PH3, which he wasn't expecting... so the guy's liver probably gave way in a dramatic fashion, terminating him on the spot.

Half-joking :) but...
This could be used for phosphorus production after all.
Phosphine would react with CuCl2 or FeCl3 and reduce them to the metals, while turning into white phosphorus vapour, a byproduct is HCl gas.

BromicAcid - 31-5-2004 at 09:33

Quote:
"My plans

On the Sunday I plannig to prepare P from H3PO4 + C = H20 + CO + P.
I have real 600-650C in my glove

Good luck
Flayer
Sorry for my bad English."
by Flayer.
This was his first and last post at MSDB

Kind of chilling (although according to his profile he was active 5 days ago). Phosphine was a major problem for most of my attempts, but luckily (kind of) lots of it was generated with appreciable diphosphine to render it spontaneously flammable. Great idea there about oxidizing the phosphine using CuCl2 or other salts. Might have to give that a try, usually this method is done by taking very high strength phosphoric acid, 95%+ and kneading in saw dust then heating the resulting doughy mass.

[Edited on 5/31/2004 by BromicAcid]

S.C. Wack - 31-5-2004 at 11:15

What would stop the P from giving the phosphides, if this works? Also, H3PO4 will not react as mentioned. Obviously it will not be in that form at the temperatures C reduces at. Even so, the reaction 4H3PO4 + 16C = 6H2 + 16CO + P4, is all over the literature, so it ought to work, somehow. But I haven't done it, so I wouldn't know for sure.

Marvin - 31-5-2004 at 21:11

Orthophosphoric acid is less than ideal for a lot of reasons.

Dehydration produces pyrophoric acid, and then metaphosphoric acid, but this tends to volatalise before the temperature is reached where reduction by carbon starts. I speak both from lit and experience.

Aside from the lead phosphate process, the most promising 'low temperature' carbon method I can find starts with calcium dihydrogen othophosphate. This is used as a fertiliser as 'triple superphosphate', strong heat converts this to calcium metaphosphate, this when reduced by carbon produces phosphorous and neutral calcium orthophophate. Its still a white heat process though.

S.C. Wack - 10-6-2004 at 12:32

Was looking at espacenet last night for P pats. that I could use. There are a lot of patents for sure and really only looked at a few. Someone with a lot of time on their hands could find better than these, and if they speak German, even more. Click on the "requested patent" #'s.

The microwave patent Polverone mentioned:
http://l2.espacenet.com/espacenet/viewer?PN=US6207024&CY...

The or a Pb phosphate/H patent:
http://l2.espacenet.com/espacenet/viewer?PN=US4287165&CY...

An iron or Al/H3PO4/silicic acid/700C pat:
http://l2.espacenet.com/espacenet/viewer?PN=GB320598&CY=...

Different, but similar, silicon and ferrophosphorus, to illustrate somewhat of a variation of the above patent:
http://l2.espacenet.com/espacenet/viewer?PN=US1836618&CY...

PClx directly from phosphate at 700C with Cl:
http://l2.espacenet.com/espacenet/viewer?PN=US1926072&CY...

This one is at 1250C, sort of in the usual thermal way, with variations and great experimental lab detail:
http://l2.espacenet.com/espacenet/viewer?PN=US2897057&CY...

Cyrus - 26-8-2004 at 14:04

Today I tried to make some elemental phosphorus, using trisodium phosphate as a fine powder (all ground by hand, my hand hurt for a while, I must be holding the pestle wrong or something), fine silica, and aluminum, in the form of snipped up wires. There was an excess of Al because I figured it was the reactant that would get mixed and used the most inefficiently.

I heated about 50 g total reactants in the distilling apparatus described in my furnace thread for about 2 hours on "hellfire" :). (the part of the apparatus in the furnace was glowing reddish orange. The only difference from the apparatus I used than the one shown in that thread was that instead of bubbling the exit gasses through a tin can soldered on, which I tried but wouldn't hold water, I put another 90 deg elbow on the end of the pipe and a short section pointing upwards, this part was filled with water.

As the thing was heated, phosphine (so I think) started coming out of the end as a white mist, so I burned it off with my propane torch, it made popping sounds and the mist disappeared.

After this, the water started getting milky, so I figured there was some phosphorus in there, but at the very end of the run, I heated the water up until it boiled, and then dumped what I supposed would be a water/phosphorus mix into a tin can filled with water. All that came out was water. Since the furnace ate a handful of wood or two every few minutes, I had to stoke the fire a LOT, and the only way to add more fuel was to take off the lid, set it down on some bricks, add more fuel, and then put the lid back on. Every time I did this some of the water spilled out. I don't think phosphorus is a good grass fertilizer. :) The furnace is still cooling down (I also fired some pottery) which takes about 5-10 HOURS! Thusly, I cannot check for more details. :(

The most "scary" part of the whole procedure had nothing to do with the reaction...

So I was sitting there watching the 30 cm high flames scream out of the furnace lid as glowing ash flew around me like snow (I had protection on) mostly in view of about 3 neighbor houses, and all of a sudden, a FIRE TRUCK comes barreling down my street! You can guess what I was thinking. The truck came closer to my house, I saw firemen suited up to do battle with their enemy, and they saw me suited up to do battle with my enemy, the truck slowly screeched to a halt in front of my house. I started thinking, WHICH LOUSY NEIGHBOR TURNED ME IN?!,

Then the firetruck left and kept going!

It turns out they come once a year to adjust fire hydrants. There is a fire hydrant near my house :) WHEW!

Three things I have learned

1 do not stand on top of a thin piece of wood on top of a bucket with a saw in one hand and stomp on the wood to break it in half.

2 take altoids before putting on gas mask. :P

3 making phosphorus might not be as easy as it sounds. :(

BromicAcid - 26-8-2004 at 14:52

Quote:
As the thing was heated, phosphine (so I think) started coming out of the end as a white mist, so I burned it off with my propane torch, it made popping sounds and the mist disappeared.


I got the white mist too at parts, I believe this is phosphorus pentoxide, it might have also been phosphorus vapor (too much furnace heat and too short of a condenser may have made most of the phosphorus resist turning to a liquid resulting in it staying in the gas phase and oxidizing on exit.) Plus there may have been some phosphine/diphosphine in there, I knew I had some form of phosphine because the bubbles were almost exploding when the exited the water.

Quote:
After this, the water started getting milky, so I figured there was some phosphorus in there, but at the very end of the run, I heated the water up until it boiled, and then dumped what I supposed would be a water/phosphorus mix into a tin can filled with water. All that came out was water.


What color was your final water, for some reason my pipe corroded at the water interface and resulted in a nasty orange color. What was the purpose of heating the water till it boiled? To get the phosphorus to condense into single blobs?

Quote:
I don't think phosphorus is a good grass fertilizer.


Actually, when I did a similar experiment and clouds of P2O5 rode across the land, a few weeks later there were mysterious dark green spots that coincidentally ran across the same areas phosphorus clouds had roamed a few weeks before.

Personally I think the most reveling thing you could do is look at the contents of your reaction vessel. Being that the aluminum was not finely divided it may have reated on the surface forming a thick Al2O3 coating which held in the molten aluminum and prevented further reaction. That is the best I can come up with as to why this does not seem to have worked well. Should be easy to tell if you look at the reaction cake.

Marvin - 27-8-2004 at 05:26

Some of this stuff has been covered allready. Trisodium orthophosphate is not a good candidate. Sodium dihydrogenorthophosphate, or a mixture of trisodium and phosphoric acid or ammonium hydrogen phosphates that will get you sodium dihydrogen phosphate should work much better, as the sodium metaphosphate this will form in the slow heating phase (you need a slow heating phase to dehydrate and avoid phosphine formation) reduces much more easily.

I would not add silica to this. Under best conditions with carbon this would only leave the trisodium othophosphate as the product, with aluminium expect some phosphides but you want the mixture to stay put while it reduces rather than form oxides and acids of phosphorous that distill out. I think like bromic this is the mist you saw. Metaphosphoric acid itself is not a good candidate because its too volatile for carbon. Maybe aluminium would would work but if you stick to a salt of metaphosphoric acid you stay in more documented chemistry. The only problem I can see if that if the aluminium is active at the temperature of dehydration this will reduce the yeild.

Phosphorous is an oxidising agent too, excess aluminium favours phosphide formation. The majority of the flammable gas should be CO. If you have phosphine it means you have hydrogen and something isnt being done right.

I'm not convinced wood will get the same sorts of temperatures as charcoal. I might be inclined to try a reduction with carbon using a charcoal or coke run. Aluminium wire cant be cost effective. I also think you need to scale up. 50 grams total reactants is nothing, this isnt a reaction you can do well small scale.

Edit, something else occured to me. What stops the aluminium reducing the silica instead?

[Edited on 27-8-2004 by Marvin]

PHOSPHORUS!!!

Cyrus - 27-8-2004 at 17:54

Ok, I did make some phosphorus. :) But last night I confused it with campbell's mushroom soup concentrate. :o

Right after the fire truck came, I decided to stop the reaction, and so poured the water/phosphorus mix into a soup can filled with water. I thought the soup can was clean, but later when I saw white stuff all along the edges, it looked like soup conc. Honestly!. However, the next morning, all of the white stuff above the waterline was gone. Soup doesn't do that. I then scraped the remaining white stuff (about 90% of it was above the water line:() into a beaker filled with water, it sank to the bottom, good.
It doesn't melt at 55+deg. C, so it can't be pure phosphorus. :(

When I dissasembled the apparatus, there was a white coating on most of the walls which did not react with water, and the tube inside the furnace had some water in it from suckback. The slag was black and hard, and I'll investigate that some more tonight. There was a nice sized chunk of phophorus (>1cc) at the spot where I had expected phosphorus to be (but could not tell w/o disassembing the apparatus), but as I was d. the a., that chunk disappeared. :mad: I don't know where it went, and I wanted to rescue and save that little guy, to keep him forever. :( Now I have <1 cc of impure white phosphorus. Hey, all I wanted it for was the challenge and thrill of making some phosphorus, and keeping it, and it looks like I've succeeded, barely. :)

I want more, of course, but I think I'll take a break from that and use the same apparatus to make some sodium, working with high temps, electricity, fire, toxic, unfamiliar, illegal, and explosive chemicals, snoopy neighbors, and firetrucks can be mentally exhausting. I feel much safer w/ sodium. ;)


Quote:
Originally posted by Marvin
Trisodium orthophosphate is not a good candidate. Sodium dihydrogenorthophosphate, or a mixture of trisodium and phosphoric acid or ammonium hydrogen phosphates that will get you sodium dihydrogen phosphate should work much better, as the sodium metaphosphate this will form in the slow heating phase (you need a slow heating phase to dehydrate and avoid phosphine formation) reduces much more easily.

---I tried Na3PO4 to eliminate all hydrogen from the equation, to prevent formation of phosphine, but next time I'll try sodium metaphosphate.

I would not add silica to this.

---That's what BromicAcid used, and it seemed to work there. :)

Phosphorous is an oxidising agent too, excess aluminium favours phosphide formation.

---I know the Al will form phosphides. :P I also know that until I use powdered Al, the Al will not be mixed very well. Would it have helped to have a stoichiometric amount of Al? Maybe. Frankly, I don't think anyone can be sure without a lot of tests. Besides, there was only about 1.1 times as much Al as there "should" have been.

The majority of the flammable gas should be CO.

--- From the reaction chamber? I don't see how we are getting carbon to appear from sodium phosphate, silica, and aluminum. From the carbon in the iron container???


I'm not convinced wood will get the same sorts of temperatures as charcoal.

--- Charcoal is made from wood. Charcoal is cleaner, but it's also not free. I doubt it's appreciably hotter.

I might be inclined to try a reduction with carbon using a charcoal or coke run.

---Maybe I'll try that next.

Aluminium wire cant be cost effective.

---Hmm, my total cost for reactants was $0.40. (for SiO2) I have several pounds of pure Al in the form of thick wires, picked up for free from a construction site. Since I only needed about 12 g IIRC, this seemed fine to me.

I also think you need to scale up. 50 grams total reactants is nothing, this isnt a reaction you can do well small scale.

---If you remember what happened when I tried the sodium hydroxide/Mg thermite, you can understand why I tried this on a small scale!

Edit, something else occured to me. What stops the aluminium reducing the silica instead?

---I thought of this too. I have no answer.

[Edited on 27-8-2004 by Marvin]


Edit- the color of the water was light yellow/orange, which I also assumed was from iron corrosion. And the reason I heated the water to boiling, it was mostly an accident. When I added another load of wood, it didn't fit all the way down the furnace, so the lid didn't fit on tightly, causing flames to shoot a foot long out all sides, this tends to boil water. Also, I wanted to be sure that the water was above 44 deg. C, so I heated it a little. It heated more quickly than was anticipated.



[Edited on 28-8-2004 by Cyrus]

JohnWW - 28-8-2004 at 05:44

There is one important matter that has been overlooked on this thread. P exists as two different allotropes: white P, which is very poisonous and destructive to the bones (it caused "phossy jaw" in workers in old-time match factories which used the stuff, arsenic having a similar effect), being still used in some liquid vermin poisons; and red phosphorus, which is much safer to handle and which is used in organic and inorganic syntheses including of methamphetamine from ephedrine. However, both are slowly inflammable (hence "phosphorescent";) in air, and have to be kept under water or other non-oxidizing liquid.

Which of these does your reduction processes produce?

John W.

vulture - 28-8-2004 at 07:58

Phosphorus preparation will always yield white phosphorus. This is because all forms of phosphorus convert to white when heated sufficiently.

Marvin - 28-8-2004 at 14:50

Cyrus,

"It doesn't melt at 55+deg. C, so it can't be pure phosphorus"

We dont care about the melting point, the only thing we're interested in is how well does it glow :)

One matter I'm not clear on, the phosphate you used, was it trisodium orthophosphate as I assumed, or trisodium polyphosphate?

"I would not add silica to this.

---That's what BromicAcid used, and it seemed to work there"

The point of silica is to liberate the phosphate and leave a silicate. Otherwise in a carbon reduction all you can get is the phosphide (prioduction of carbide is not going to happen). You shouldnt need this with aluminium (particually with a metaphsphate, but it should leave an aluminate), nor is it useful in low temperature reductions.

"The majority of the flammable gas should be CO.

--- From the reaction chamber? I don't see how we are getting carbon to appear from sodium phosphate, silica, and aluminum. From the carbon in the iron container??? "

Thats because I temperarily mislaid my brain. You are right, no carbon, but then there shouldnt be any hydrogen either to make phosphine. Did you dehydrate the phosphate first?

"--- Charcoal is made from wood. Charcoal is cleaner, but it's also not free. I doubt it's appreciably hotter. "

Charcoal is a much hotter fuel than wood. I *think* this is because overall with wood you are producing a lot more gas (mainly water vapour) so you are heating a larger volume of gas. This would be akin to oxygen producing a much hotter temperature than air, the nitrogen doesnt interfere, but you have more gas to heat up with the same energy.

"---If you remember what happened when I tried the sodium hydroxide/Mg thermite, you can understand why I tried this on a small scale! "

Point taken, now I'm being unsafe. Carbon runs are almost certainly poor at that level and I'm stuck in that mindset.

JohnWW,

Phosphorous exists in a lot more than 2 forms. White, Red, Violet, Black ... I think someone said there were about 7 known forms at least. vulture is quite correct, the vapour is in the form of P2 molecules, it will always condense to white initially under normal conditions, high pressure and long cooking usually with a catalyst is used to make the other forms.

Red phosphorous does not glow and is virtually unreactive to air. It is not kept under water.

The term phosphorescent in english includes phosphorous, but in chemistry the more accurate term would be chemiluminescent. Phosphorescence is used soley to describe emission of light between states of different spin multiplicities, ie it glows after you remove the source of the energy.

BromicAcid - 28-8-2004 at 16:47

Quote:
You shouldnt need this with aluminium (particually with a metaphsphate, but it should leave an aluminate), nor is it useful in low temperature reductions.


I used SiO2 with my procedure because the references for producing phosphorus using aluminothermic reduction of phosphates stated that without silica a large portion of the phosphorus becomes irrecoverable due to conversion to aluminum phosphide, the SiO2 reduces this and can improve yields from about 20% to begin with to about 75%. I belive it also helps to decrease the violence of this reaction, (NaPO3)6 powdered reacting with aluminum powder without SiO2 is almost like flash in decent amounts rendering almost impractically fast.

Quote:
Thats because I temperarily mislaid my brain. You are right, no carbon, but then there shouldnt be any hydrogen either to make phosphine. Did you dehydrate the phosphate first?


I actually think my water for phosphine formation came from the sand that I used. SiO2 does absorb water, does it have to be specially prepared to absorb water or does it just collect on its own?

[Edited on 8/29/2004 by BromicAcid]

Marvin - 28-8-2004 at 17:24

I stand corrected on the addition of silica. I'm still confused as to the deeper reason though.

I understood it only absorbed water on the surface normally so fumed silica might take up a reasonable amount but sand shouldnt. Something like silica gel is a hydrated sillic acid.

vulture - 29-8-2004 at 04:34

Quote:

Something like silica gel is a hydrated sillic acid.


Silica gel is made from H4SiO4 in solution, but this acid is unstable and polymerizes to SiO2. Yes, SiO2 is actually a polymer. So it seems likely that the clusters on the outside still have hydroxy groups attached. The smaller the gel the more hydroxylgroups you will have.

Cyrus - 29-8-2004 at 14:06

I think the water came, as BromicAcid suggested, from any water absorbed by the very fine SiO2 or from any residual water in the sodium phosphate I used.

The sodium phosphate was made from NH4H2PO4 and NaOH, so I think that the sodium phosphate was Na3PO4.


Edit- about charcoal vs plain wood, When I added the fresh, slightly damp, cold wood, within 0.5 to 1 min. the part of the apparatus within the furnace would be glowing reddish orange, but when the wood had turned to charcoal, the apparatus cooled down, iirc, to a very dull red glow. :o

Heat, fires, summer, hot and thick clothes, all outside. Maybe I should go for ultra LOW temps instead. :)

I'll probably decide to do that about, oh, January. :P

[Edited on 31-8-2004 by Cyrus]

BromicAcid - 1-9-2004 at 14:04

Tried phosphorus today using bone ash (Calcium Phosphate). I ran three attempts, a stoichiometric mix of SiO2, Al, and Ca3(PO4)2 this attempt showed no results with my heating method, a MAPP gas torch, all reactants were very fine powders, the SiO2 was fumed silica, aluminum was 300 mesh, and the calcium sulfate was a very fine powder. The second attempt was half way between the reduction with SiO2 and the reduction with just Al, this also was uninitiateable with a MAPP gas torch. Finally I ran the reaction using entirely Al as the reluctant. This formed a blob at the bottom that was at least partially composed of the phosphide as evidenced by the phosphine smell, nothing came over in the retort though.

Also I ran the reaction using sodium hexametaphosphate, silicon dioxide, and aluminum powder. All my previous attempts with (NaPO3)6 used regular play sand as I was unable to obtain the fine SiO2 that I currently have. Therefor this one was different. It formed a solid cake at the bottom, similar to the calcium reduction with straight aluminum. This cake was full of bubbly holes and when broken apart flared up in spots from phosphorus. But nothing came over in the retort. I believe on my previous attempts that the SiO2 in its corse form helped to aerate the mixture and allow it to react more thoroughly without forming that cake.

Regardless, going to try again next week possibly. I need to get a picture up of my setup still, it is pretty neat looking.

Edit: My retort



1) That is one of those hand held extension burners for soldering and such, it is hooked to MAPP gas and is held with one of those grippy claw things that go on the ring stand. 2)That is a glass pipette that is attached to 3) A thick vacuum hose running to 4) A ball check valve to prevent suckback of both water and air into the system. This provides the burning method of the exit gasses as the pipette runs into the flame, or at least close by it, when it ran right into the flame the predictable result happened, the tip melted and blocked the exit. 5) A spaghetti sauce jar with two holes put into the lid. These holes are the exit and entry points and are sealed with an epoxy around each pipe. The entry pipe pictured on the left has a piece of screen surrounding it to prevent massive bubbles and prevent corky phosphorus. This jar is filled with water. 6) The body of the retort, covered in fiberglass insulation to prevent phosphorus from solidifying on the inside. It is made from 1/4 inch piping except right at the reaction vessel 7) Which is made from a 1 1/2 inch endcap connected to a 1 1/2 - 1/2 inch brushing, the brushing runs to a 1/2 inch 90 deg elbow to a short 1/2 inch pipe to adapt it to a 45 deg 1/2 - 1/4 adapter.

Simple with only a few joints and prevents suck back and exit gasses are held in check. I like it
;)

[Edited on 9/1/2004 by BromicAcid]

Cyrus - 1-9-2004 at 18:46

Are you sure that the "jam jar" will not implode under vacuum? That would be unfortunate. :o Might a metal container be better?

In order to prevent suckback, how about this?

Well the picture isn't working, my idea was to have a tee connection in between the heated part and the water with a ball valve connected to that, on the other side of the ball valve, a balloon filled with an inert gas. When heating is done, slowly open the ball valve.

[Edited on 2-9-2004 by Cyrus]

balloon.bmp - 390kB

BromicAcid - 1-9-2004 at 19:30

Your design looks nice but with the scale of my design (my reaction vessel is only 36 cm 3) There really isn't enough area to generate a significant vacuum. Besides, my ball valve check is not totally air tight, it is home-aid, and the spaghetti jar is probably not air tight either, I am not afraid of it imploding, although worse things have happened :(

BromicAcid - 9-9-2004 at 09:54

Alternative phosphorus production using an interesting mechanism I saw the other day.

Zn3P2(s) + HCl(aq) ---> ZnCl2(aq) + PH3(g)
PH3(g) + Ni(CH3COO)2(aq) ---> P(x)Ni(x)(s) + xCH3COOH(aq)

Commercially available zinc phosphide in the form of mole killing pellets is reacted in an airtight vessel with concentrated hydrochloric acid at a controlled rate. The exit gasses consisting mostly of phosphine are lead though a concentrated solution of a soluble nickel salt. Phosphine reduces the nickel forming nickel-phosphorus 'alloys' of variable composition. These alloys are filtered from the solution and subject to treatment with dillute potassium dichromate/sulfuric acid solutions to yeild free phosphorus which is melted under water to consolidate it into one blob.

I was thinking of this due to a passage that I read the other day that bubbling phosphine though soluble nickel salts results in the formation of insoluble nickel alloys of high phosphorus content, under suitable conditions true alloys can be formed. The dichromate sulfuric acid is the normal procedure to wash phosphorus and hopefully it would dissolve out the nickel given time. According to my calculations one container of mole pellets (250 g) should yield 1 g of phosphorus, but it's an alternative method. However ... the extreme toxicity of phosphine (and its long term effects) make this method an adventure that I would not be willing to take.
:o

chemoleo - 9-9-2004 at 10:56

Actually that is a method I'd happy to undertake as long as it's done outside, and upwind...
Sounds like an easy method otherwise!
Apart from this pellet - are there any other sources, or ways, to get Zn3P3, or other alternativve phosphides?
I mean, think about it, do you think sublimating phosphorous is much more pleasant, particulalry at the temperatures normally required?

Also - what is the mechanism of it? i.e. the reduction of PH3 to P, forming this 'alloy'?

JohnWW - 9-9-2004 at 11:04

Unless you use a gas mask, or better still breathing apparatus like what firefighters use. The toxic mechanism of PH3 is similar to that of CO, H2S, HCN, (CN)2, azides, etc, due to its strongly complexing Fe in hemoglobin. In fact, it may be more toxic than even H2S.

I think Zn3P2 is also used for killing rats and especially rabbits, by putting the pellets into their burrows.

John W.

BromicAcid - 9-9-2004 at 11:35

I chose Zn3P2 because it is the most widely avalible, as John WW said it is for killing rabits and other wildlife, moles like I said. Around here it comes in 250 g containers and is between 2 and 4% Zn3P2.

The mechanism is just phosphine reducing the nickel and in the process being oxidized to elemental phosphorus.

2PH3(g) + 3Ni2+(aq) ----> Ni3P2(s) + 6H+(aq)

That's just an approximate equation, I will re-look up the reference tommorow since there is intrest but it's not a phosphide that is formed, it is a mixture, so that formula would not be accurate. The book stated that the mixture formed is of a highly variable composition depending on condition. Will post more tommorow.

BromicAcid - 10-9-2004 at 15:49

Taken form "Chemical Elements and their Compounds" Vol I 1962 Sidgwick pg 730

Quote:
Phosphine will reduce nickel salts in aqueous solution, forming alloy-like P-Ni compounds with 0.4 - 0.1 P to 1 Ni : under special conditions, definite phosphides Ni3P, Ni2P, and NiP are obtained. 457


457 R. Scholder, A. Apel, and H.L. Haken, Z. anorg. Chem. 1937, 232, 1.

Zinc shall save you

Theoretic - 18-9-2004 at 12:39

I've had an idea. Zinc could be the perfect reducing agent. Look what happens if you have molten Na3PO4 and you throw in a gob of zinc (in stoichiometric proportion, of course).
First zinc reduces PO4--- to P---:
PO4--- + 4Zn => P--- + 4ZnO.
The thus formed oxide film on zinc is dissolved away by the phosphate ions:
ZnO + 2PO4--- => ZnO2-- + P2O7----,
then when ortophosphate ions run out its:
ZnO + P2O7---- => ZnO2-- + 2PO3-.
Of course all of the PO4--- turns into PO3- and P--- before the zinc runs out, then phosphide and metaphosphate start conproportionating:
5P--- + 12PO3- => 8P + 9PO4---,
PO4--- reacts with ZnO as was shown above, the exposed zinc surface reduces some more PO4---, and so you again have phosphide and ortophosphate, they react and the cycle repeats untill zinc runs out and you have a solution of ZnO in molten Na2ZnO2 in your vessel, and dough* (sorry, white phosphorus) in the condenser. The advantages are that the zinc can be a gob instead of a powder, it's a much less vigorous reducer than Al and so the reaction can't get out of hand, there are two components instead of three (phosphate and zinc as opposed to phosphate, Al and SiO2), the reaction isn't stopped by a tough oxide layer, is faster and goes much nearer to completion. The disadvantage is that aluminium is very easy, getting sizeable amounts of zinc is harder.

*dough - taken from Stupid White Men by Michael Moore, where he says: More dough for YOU! meaning money, I meant phosphorus :P.

[Edited on 18-9-2004 by Theoretic]

To boot

Al Koholic - 18-9-2004 at 13:12

Acquiring Zn has been discussed in another thread but I'd just like to point out that by melting a few bucks worth of US pennies one can obtain quite a bit of Zn. Just scrape off the slag formed after the melt to get reasonably pure material. I also doubt impurities will really matter.
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