Sciencemadness Discussion Board - Phosphorus Pentachloride

Phosphorus Pentachloride

cal - 17-11-2012 at 11:15

I recently obtained 1 kilo of calcium phosphide Ca3P2 at a cost to cheap to pass up. I will be attempting the Leonid Lerner Small Scale Synthesis of Phosphorus Pentachloride from calcium phosphide and chlorine at 240 degrees C.
Has anyone here done this reaction? :cool:

UnintentionalChaos - 17-11-2012 at 11:24

Presumably Leonid Lerner has, as he is a member here.

Leonid Lerner

cal - 17-11-2012 at 12:31

I did a search and did not find a member under Leonid Lerner.
Maybe he uses a different name here.

Mailinmypocket - 17-11-2012 at 12:34

He does. It is Len1 I believe

SM2 - 17-11-2012 at 12:42

Careful what you do w/ that PCL5.....or it will Learn you a new lesson. BTW: why on Earth would you be messing with such an infamous and nasty chlorinater? take care, >leaving> now.

kristofvagyok - 18-11-2012 at 06:08

Even working with PCl5 is a really-really bad thing, to prepare it at home could be a nightmare.

I remember when I first worked with PCl5 I used a nearly 10 year old sample what was opened once before me. It was in a large glass bottle, made by Fluka (contained 1 kilo) and when I opened, it just blown down the cap. When it was opened long ago, a little water had been absorbed and it generated a lot HCl, POCl3 during the time, so it became overpressured... Be careful

Magpie - 18-11-2012 at 11:31

Quote: Originally posted by kristofvagyok

Even working with PCl5 is a really-really bad thing....

I'm puzzled by this statement. I have not worked with PCl5 but have used PCl3 - back in my introductory organic class at a university in 1962. It was kept in a reagent bottle with a ground glass stopper in the lab hood. We used it to make acetyl chloride as one of our assigned experiments. I do remember the instructor (a busy PhD grad student) telling us to only open the bottle in the hood.

[Edited on 18-11-2012 by Magpie]

woelen - 18-11-2012 at 11:50

PCl5 is a very nasty chemical and it is hard, really hard, to store this material. I have a small sample of the material and it really eats every cap. Even the beautiful red GL45 caps with teflon liner are eaten slowly and become soft over time.

Even bromine is a toy when compared to PCl5 when it comes to storage!
Als be very careful with the phosphide. With water it releases PH3, which is extremely toxic and this reaction also occurs, even when some water vapor comes into the bottle of the phosphide.

How do you think you are going to make PCl5? Passing chlorine over the phosphide?

[Edited on 18-11-12 by woelen]

stoichiometric_steve - 18-11-2012 at 13:26

you should absolutely buy my mag stirrer before you die.

DJF90 - 18-11-2012 at 13:30

LOL now thats salesmanship

SM2 - 18-11-2012 at 13:51

Quote: Originally posted by woelen

PCl5 is a very nasty chemical and it is hard, really hard, to store this material. I have a small sample of the material and it really eats every cap. Even the beautiful red GL45 caps with teflon liner are eaten slowly and become soft over time.

Even bromine is a toy when compared to PCl5 when it comes to storage!
Als be very careful with the phosphide. With water it releases PH3, which is extremely toxic and this reaction also occurs, even when some water vapor comes into the bottle of the phosphide.

How do you think you are going to make PCl5? Passing chlorine over the phosphide?

[Edited on 18-11-12 by woelen]

woelen, that's amazing that it will attack even polytef. Perhaps the PCL5 can fit between the chain of polytef molecules. My first intuition would tell me no, but if it's just thin (membrane) polytef, maybe that is different. How about a quartz bottle with a ground quartz stopper? I think this chemist would be much safer making his PCL5 from the PCL3 +Cl in situ.

Endimion17 - 18-11-2012 at 15:15

No stoppers, no nothing. Glass seals. End of discussion.
People, learn to make ampoules or you will be taught about the danger the hard way.

Dr.Bob - 19-11-2012 at 09:08

While Teflon is only attacked by a few chemicals, some molecules can diffuse into and through teflon and other fluoropolymers, like other halogens. DCM and TFA both have a slow permeation through teflon as well. That is why Teflon stirbars can turn colored or dark, many material can absorb into the Teflon and get trapped there. It does not mean that they have destroyed the teflon, only been absorbed into it. That can soften it, as Woelen stated. Silicone (and Viton to a lesser extent) cap liners are even worse, many solvents swell them unmercifully, even though the silicone may not dissolve, but the liner is damaged over time by the swelling.

Magpie - 9-8-2015 at 13:22

In theory I have produced 13.4g of PCl5 using the method of garage chemist. I dissolved 2g of P in 100mL of chloroform then sparged it with dry Cl2. The key to this technique is the plunger used to keep the sparge tube end from plugging with PCL5, a white solid.

Originally I tried to do this with a ptfe plunger rod but it was too weak and would break. I then tried to make a heated sparge tube to keep the PCl5 a liquid but I could not make a tube that did not break. This latest version employs a 3/32" ss316 welding wire. This worked well. It is, however, subject to some corrosion giving the solvent a yellow color. Dissolved Cl2 also contributes to this color. Next time I will try an AWG 12 copper wire. The wire is sealed using a gland made of an inverted rubber septum, size 19/22.

Krytox grease was used for all tapered glass joints. This really works well and no hard, sticky white residue was formed. Just a thin layer is needed. I must thank U. Chaos for his suggestion to use this with Cl2.

Tommorrow I will continue with the workup to isolate and weigh the PCl5. I'm thinking storage in an ampule may be best after reading the above posts.

-------------------------------------------------------------
Overnight most all of the dissolved Cl2 outgassed from the solvent as the solvent is almost clear now. Apparently there was very little FeCl3 if any.

PCl5 apparatus

PCl5 in chloroform

[Edited on 9-8-2015 by Magpie]

[Edited on 10-8-2015 by Magpie]

Oscilllator - 9-8-2015 at 17:40

Did the phosphorous you used actually dissolve in the chloroform? I would have thought the polymeric structure of red phosphorous would prevent that. Unless you used white phosphorous

I also noticed that the picture on the wiki page of PCl5 was taken by woelen. It seems he has contributed pictures of just about every chemical under the sun!

Magpie - 9-8-2015 at 18:00

I used white P from some I made several years ago. Yes it dissolves at 1g/40mL. Garage chemist used red P.

The reaction was very smooth but does generate significant heat. Hence the reflux condenser.

chemrox - 9-8-2015 at 21:13

I transferred mine to a Teflon bottle and store that in the fume cupboard .. before I did that I wondered why I was getting acid damage all over the lab.

Magpie - 10-8-2015 at 04:35

chemrox, so you are saying that PCl5 does not destroy Teflon but that the bottle still leaks? I presume then, that your hood fan is on all the time?

PCl5 + H2O ---> POCl3 + 2HCl

So it seems that ingress of H2O is the root of the storage problems with PCl5. This is true also with storing AlCl3.

2AlCl3 + 3H2O ---> 6HCl + Al2O3

With AlCl3 I had a serious problem (bottle destruction) with what was apparently H2O permeating the plastic bottle and reacting to form HCl.

I store muriatic acid (31% HCl) in a 100mL narrow mouth glass bottle (Qorpak) with a Teflon seal in my lab as a day use bottle. I store the bottle in a plastic bag with some NaHCO3. This seems to work OK as my carbon steel tools are not rusting (much).

Magpie - 10-8-2015 at 14:28

...continuation of yesterday's post:

The pot with PCl5 and chloroform was set up for simple distillation to remove the chloroform. An oil bath controlled with a PID was used as a heating source. The chloroform came over at 61°-65°C. There was likely some HCl in the beginning and some CCl4 at the end.

solvent removal from PCl5

PCl5 precipitating during distillation.JPG - 71kB

PCl5 precipitating during solvent removal

As garage chemist indicated the PCl5 still contained some solvent even though I had the bath oil up to 110°C. So, I did as he did, ie, subjected the slush to vacuum distillation at 177mmHg until I saw what likely was sublimation on the RBF walls. 1.2mL of solvent came over.

PCl5 (with stirbar) drying during distillation.JPG - 63kB

PCl5 (w/stirbar) nearly dried

I then reached in the side-neck with a spatula and broke up the crumbly PCl5. I again subjected it to vacuum distillation at 152mmHg but no more solvent came over.

PCl5 vacuum dried

After cooling the powder was transferred to a Qorpak 100mL wide-mouth bottle with ptfe lid seal.

The net weight of PCl5 was 9.1g for an efficiency of 67.7%.

Questions, comments, suggestions are welcomed.

[Edited on 11-8-2015 by Magpie]

[Edited on 11-8-2015 by Magpie]

Oscilllator - 10-8-2015 at 20:23

I have a question: Where do you think the remaining 32.4% of the phosphorous went? If the phosphorous remained unreacted, it seems like this could be very dangerous if the PCl5 contaminated with white phosphorous was exposed to the air, as that would end very badly as I'm sure you can imagine.
Is it possible that you lost a significant amount during the time the PCl5 was under vacuum? Wikipedia lists the boiling point as 160.5° so it seems likely that some could come over at the 110° temperature of the oil bath.

Magpie - 10-8-2015 at 21:13

Good question. I know that some of it remains as PCl5 sublimate on the inside of the 250mL RBF. As it was subliming part of it probably went over into the distillate receiver with the chloroform in which it is highly soluble.

Traces of water would convert it to HCl/POCl3 as suggested by garage chemist.

Tomorrow I plan to make a YouTube video of what happens when water is added to the RBF with the sublimate. If it is interesting I will post it in this thread.

Your point about the danger of not accounting for all P is well taken. Sometimes it seems like a wild animal waiting in the bushes to pounce on the unwary chemist. Elemental P is especially dangerous due to its fire hazard.

I have seen no indication that the PCl5 is contaminated with P but I don't see why it couldn't be. Placing some on a watch glass might tell if it starts smoking. But this might be easy to confuse with the HCl/POCl3 being formed when it reacts with the humidity in the air. Right now I have it covered with argon in a tightly closed bottle and hesitate to open it.

I saved the recovered chloroform and plan to distill it one day. I could get some surprises there too.

The highest oil bath temperature was indeed 110°C but it takes about a 20° differential to get sufficient heat transfer.

Oscilllator - 10-8-2015 at 21:37

Quote: Originally posted by Magpie

I have seen no indication that the PCl5 is contaminated with P but I don't see why it couldn't be. Placing some on a watch glass might tell if it starts smoking.

A easy way to tell the difference would be to look at this experiment in the dark to watch for white phosphorous' characteristic glow. I haven't personally experienced this so I don't know how bright it is, but a long exposure photograph should reveal the presence of even small quantities of the stuff.

Marvin - 11-8-2015 at 05:26

Wouldn't the reaction have to pass through the trichloride stage before any pentachloride formed?

Magpie - 11-8-2015 at 06:06

You are probably right Marvin.

Another safety factor is the fact that I used 2X the stoichiometric amount of Cl2. The solvent turned a strong yellow long before I finished gassing. P reacts very strongly with Cl2. It would be extremely strange to find any elemental P anywhere.

I started this experiment back in April but had to suspend it due to plugging of the sparge tube. So, I just placed the 2-neck RBF containing the P dissolved in chloroform in a sand bucket then placed this in my tool shed well away from the house. It sat there without popping a stopper all summer through scorching heat - ambient sometimes over 110°F (43°C)!

When I weighed out the P I wasn't real careful as indicated by the fact I report it with only one significant digit. So I may have used something less than 2.00g (or something more!). Next time I will weigh it more carefully and look for a better accounting of all the P, ie, a mass balance.

Magpie - 11-8-2015 at 11:45

Here's a video of the sublimated PCl5 reacting with water:

https://youtu.be/ZL8VkLEz1vg

Magpie - 12-8-2015 at 08:47

Quote: Originally posted by Marvin

Wouldn't the reaction have to pass through the trichloride stage before any pentachloride formed?

Leonid Lerner in his book shows that there is an equilibrium between PCl5 and PCl3, ie

PCl5 <---> PCl3 + Cl2

K = 0.042 @ 190°C

When I pulled a vacuum on the PCl5 I first noticed the residual chloroform released from the PCl5 condensing/evaporating on the walls of the RBF. Then PCl5 sublimate formed on the walls along with a clear liquid film. I'm thinking now that this liquid was likely PCl3. With the oil bath at 110°C the PCl3 was likely being evaporated and then condensed in the condenser. From there it would mix with the chloroform in the receiving flask. So this is another possible source of loss of PCl5.

Magpie - 19-8-2015 at 09:44

I used 5g of my PCl5 to convert benzophenone oxime to benzanilide using a Beckmann rearrangement. Yield was 64%.