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AndersHoveland
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Reaction of Phosphorous with water?
White phosphorous is a highly reactive element, which is spontaneously oxidized in the presence of air. It can, however, be stored under water without
any observable reaction. But despite this apparent unreactivity, I am going to make an argument that white phosphorous does indeed
react with water.
First, in support of the argument, let us do a calculation to see if such a reaction would be energetically favorable.
2 P + 3 H2O --> PH3 + H3PO3
The heat of formation of solid crystals of phosphorous acid is apparently -964.4 kJ/mol, while a different source gave it as -972. The heat of
formation of crystals of phosphorous acid dissolved in an infinite quantity of water is -6.3 ±0.6 kJ/mol. The heat of formation of phosphine is +5.4
kJ/mol (PH3 prefers to decompose into its separate elements). The heat of formation of water is -285.8 kJ/mol. White phosphorous, as an element, has a
value of 0, while red phosphorous is -17.5 kJ/mol.
-972 + (-6.3) = -285.8
0 + 3(-285.8) = -857.4
+5.4 - 964.4 = -959
Since one mole of PH3 and one mole of aqueous H3PO3 have have a net heat of formation lower than phosphorous and three moles of water, we
would expect that such a reaction would be energetically favorable.
From the Literature
We can also look at some reactions involving phosphorous chemistry to provide support for the idea.
Phophorous can be stored under water, but when finely divided it decomposes water producing hydrogen phosphide. When boiled with water, phosphine and
hypophosphorous acid are produced. (this information came from an online book, but unfortunately I neglected to record the reference when I put
this information in my notes)
One old source says that the reaction of sodium phosphide with water generates small portions of phosphites and hydrogen gas, in addition to the main
products of sodium hydroxide and phosphine, although this could potentially have been caused by impurities.
investigation of sodamide and of its reaction-products with phosphorus" William Phillips Winter p42-43
Aqueous alkali (KOH) dissolves red phopshorous, with the formation of phosphine. Interestingly, when hydrochloric acid was added to the solution, the
phosphorous precipitated back into its elemental form. This could suggest an equilibrium, which shifts depending on pH.
Journal of the Chemical Society, Volume 75 (Great Britain) p.976
So why is there no obvious reaction when bulk white phoshorous is placed in water? There are several possible explanations. There could be a coat of
something forming on the surface that prevents further reaction. For example, trying to dissolve large pieces of phosphorous pentoxide in water can be
difficult if it has not been finely divided before. It has a tendency to form a protective viscous coating that inhibits further hydrolysis. In the
case of elemental phosphorous reacting with water, it might be analogous to trying to hydrate silica. Si(OH)4 is quite stable, either as a solid or in
aqueous solution, but it is essentially impossible to make bulk solid SiO2 react with water. Another possible reason could be activitation energy,
although this seems less likely.
[Edited on 11-5-2012 by AndersHoveland]
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AndersHoveland
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I would wonder whether white phosphorous could be made to react with water in a solvent that could dissolve both reactants. But I am not sure what
solvent could dissolve both that would also not be reduced by the phosphorous. DMSO can dissolve both sulfur and water, but I have a feeling it might
react undesirably with phosphorous. Acetone would probably get reduced by phosphorous also. Perhaps triethylamine would be a good solvent.
N,N-Diisopropylethylamine is another similar amine that is often used as a solvent in organic chemistry.
So dissolve some white phosphorous in pure N,N-Diisopropylethylamine, then add water and see if any bubbles of phosphine form.
[Edited on 12-5-2012 by AndersHoveland]
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arsphenamine
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White
phosphorus, as the tetrahedral P<sub>4</sub> allotrope, has a ΔH<sub>f</sub> of +58.9 kJ/mol.
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AndersHoveland
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That would only make the reaction even more favorable.
Not that it matters for the purpose of this thread, but I just do not see how an element can have non-zero values for all of its allotropes. Are you
sure that value is not for P4 in the vapor phase?
By definition, the heat of formation is the energy released when a compound (or allotrope) is formed from its respective elements, or rather just the
opposite of that value. At least one of the allotropes of phosphorous must have a value of zero.
[Edited on 12-5-2012 by AndersHoveland]
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arsphenamine
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Check the NIST webbook on P4 thermochemical properties.
Red phosphorus is an intermediate state, P2 forms by sundering P4 at T>700C, atomic P requires an even higher temperature.
The P4 allotrope has 3 isoenergetic LUMO's at ~120 kJ/mol, and may indicate the excitation needed to start a reaction.
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AndersHoveland
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| Quote: |
"There are two forms of elemental phosphorous, white phosphorous, which has a standard enthalpy of formation, ∆Hf° = 0 kJ/mol, and red
phosphorous which has a standard enthalpy of formation, ∆Hf° = -17.5 kJ/mol"
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http://chemistry.osu.edu/~woodward/ch121/mid2_f03.pdf
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Lambda-Eyde
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First of all, there are more than two allotropes of phosphorus, and the paper is written in comic sans. I'm having a hard time taking it seriously...
This just in: 95,5 % of the world population lives outside the USA
You should really listen to ABBAPlease drop by our IRC channel: #sciencemadness @ irc.efnet.org
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AndersHoveland
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If phosphorous does have some hydrolysis with water, it could potentially explain the strong toxicity of white phosphorous. Actually, the mechanism of
toxicity of phosphine is not fully understood, but it is generally thought that it binds to the enzyme cytochrome c oxidase. However, I suspect that
as a powerful reducing agent, it could be irreversibly reacting with and inactivating some other important enzyme.
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weiming1998
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Quote: Originally posted by AndersHoveland  | White phosphorous is a highly reactive element, which is spontaneously oxidized in the presence of air. It can, however, be stored under water without
any observable reaction. But despite this apparent unreactivity, I am going to make an argument that white phosphorous does indeed
react with water.
First, in support of the argument, let us do a calculation to see if such a reaction would be energetically favorable.
2 P + 3 H2O --> PH3 + H3PO3
The heat of formation of solid crystals of phosphorous acid is apparently -964.4 kJ/mol, while a different source gave it as -972. The heat of
formation of crystals of phosphorous acid dissolved in an infinite quantity of water is -6.3 ±0.6 kJ/mol. The heat of formation of phosphine is +5.4
kJ/mol (PH3 prefers to decompose into its separate elements). The heat of formation of water is -285.8 kJ/mol. White phosphorous, as an element, has a
value of 0, while red phosphorous is -17.5 kJ/mol.
-972 + (-6.3) = -285.8
0 + 3(-285.8) = -857.4
+5.4 - 964.4 = -959
Since one mole of PH3 and one mole of aqueous H3PO3 have have a net heat of formation lower than phosphorous and three moles of water, we
would expect that such a reaction would be energetically favorable.
From the Literature
We can also look at some reactions involving phosphorous chemistry to provide support for the idea.
Phophorous can be stored under water, but when finely divided it decomposes water producing hydrogen phosphide. When boiled with water, phosphine and
hypophosphorous acid are produced. (this information came from an online book, but unfortunately I neglected to record the reference when I put
this information in my notes)
One old source says that the reaction of sodium phosphide with water generates small portions of phosphites and hydrogen gas, in addition to the main
products of sodium hydroxide and phosphine, although this could potentially have been caused by impurities.
investigation of sodamide and of its reaction-products with phosphorus" William Phillips Winter p42-43
Aqueous alkali (KOH) dissolves red phopshorous, with the formation of phosphine. Interestingly, when hydrochloric acid was added to the solution, the
phosphorous precipitated back into its elemental form. This could suggest an equilibrium, which shifts depending on pH.
Journal of the Chemical Society, Volume 75 (Great Britain) p.976
So why is there no obvious reaction when bulk white phoshorous is placed in water? There are several possible explanations. There could be a coat of
something forming on the surface that prevents further reaction. For example, trying to dissolve large pieces of phosphorous pentoxide in water can be
difficult if it has not been finely divided before. It has a tendency to form a protective viscous coating that inhibits further hydrolysis. In the
case of elemental phosphorous reacting with water, it might be analogous to trying to hydrate silica. Si(OH)4 is quite stable, either as a solid or in
aqueous solution, but it is essentially impossible to make bulk solid SiO2 react with water. Another possible reason could be activitation energy,
although this seems less likely.
[Edited on 11-5-2012 by AndersHoveland] |
Just because a reaction is energetically favourable doesn't mean it happens.
Your proposed reaction 2P+3H2O===>PH3+H3PO3 would essentially reverse itself in neutral water, as this:
PH3+H3PO3===>PH4H2PO3 (As PH3 is basic, like NH3), Then it might either get stuck there and the hydrolysis will proceed further to produce a mix of
PH4+ and PO3 (3-) ions, or, if PH4+ is reducing enough, gets oxidized back into P as: PH4H2PO3===>2P+3H2O, essentially driving the reaction back to
the start. So, if I am right, hydrolysis of P is essentially in an equilibrium, with it leaning strongly to the left under neutral conditions. Under
basic conditions, however, the hydrolysis of phosphorus can almost certainly be predicted. In very acidic water, the hydrolysis can probably proceed
as well, as the following happens: 2P+3H2O+H+===>PH4(+)+H3PO3, as the strong acid forms a salt with the PH3, preventing the reverse reaction back
to P from happening, and shifting the equilibrium to the right. But only a basic condition would generate any phosphine at all.
[Edited on 12-5-2012 by weiming1998]
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PHILOU Zrealone
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Just as side notes:
1°)Water dissolves oxygen in some extend...so it is plausible that white P under water will react with the dissolved O2 to produce P oxydes and
finally acids...
This effect must be shown by some phosphorescence or by noticeable pH drop with time.
2°)Red P is said to be stable towards O2 or air at ambiant T°...but if you drop a piece of red P into concentrated H2O2 (50%/200 Volumes) ...you
will get inflamation in a few seconds and evolution of white acidic fumes of H3PO4. I have done the test several times and it works marvelously...but
unfortunatelly I did no test with H2O2 (30%/130 Volumes)...
PH Z (PHILOU Zrealone)
"Physic is all what never works; Chemistry is all what stinks and explodes!"-"Life that deadly disease, sexually transmitted."(W.Allen)
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Zan Divine
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"In water with low oxygen, white phosphorus may react with water to form a compound called phosphine."
-----Agency for Toxic Substances and Disease, National Center for Environmental Health
He who makes a beast of himself gets rid of some of the pain of being a man. --HST
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Magpie
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I recently washed a film of red P off a piece of glassware using aqueous 10% NaOH. After sitting overnight in the solution the small flakes of red P
have disappeared. I saw no bubbles that would indicate generation of PH3.
What has happened to the red P?
The single most important condition for a successful synthesis is good mixing - Nicodem
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Adas
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| Quote: Originally posted by Magpie | I recently washed a film of red P off a piece of glassware using aqueous 10% NaOH. After sitting overnight in the solution the small flakes of red P
have disappeared. I saw no bubbles that would indicate generation of PH3.
What has happened to the red P? |
Probably NaH2PO2 was formed, and the phosphine just dissolved?
Rest In Pieces!
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Eddygp
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Burning e.g. propanol is exothermic, but it doesn't mean it happens spontaneously, not pyrophoric. Compare these two...
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
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Magpie
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| Quote: Originally posted by Adas | | Quote: Originally posted by Magpie | I recently washed a film of red P off a piece of glassware using aqueous 10% NaOH. After sitting overnight in the solution the small flakes of red P
have disappeared. I saw no bubbles that would indicate generation of PH3.
What has happened to the red P? |
Probably NaH2PO2 was formed, and the phosphine just dissolved? |
That seems possible.
Today I placed small bits of red P in a near boiling aqueous solution of 10% NaOH. Quite a lot of very tiny bubbles were emitted.
The disproportionation reaction for the formation of PH3 is:
3NaOH + P4 + 3H20 ---> 3NaH2PO2 + PH3
The single most important condition for a successful synthesis is good mixing - Nicodem
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AndersHoveland
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The reaction of phosphorous with aqueous base produces several mixed reaction products, probably mostly phosphine and hypophosphite, and some
phosphite.
Hypophosphorous acid decomposes when heated to give off phosphine gas. Even phosphorous acid converts to phosphoric acid and phosphine when heated to
200 °C.
One more comment about the structures of hypophosphorous and phosphorous acids. The hydrogen atom bonded to the phosphorous atom is probably
electron-donating to the oxygen atom that is double-bonded to the phosphorous. In other words, the phosphorous could be viewed as actually being in
the +3 state, with the hydrogen ion only weakly coordinated with the phosphorous atom. Indeed, I suspect measurements would show the P—H
bond order in the molecule is less than 1.
Phosphorous acid, O=PH(OH)2 is known to tautomerise into P(OH)3, further supporting the idea that the H can easily ionise off.
[Edited on 25-5-2012 by AndersHoveland]
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AndersHoveland
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I really suspect that, in several ways, phosphorous may behave more like aluminum in its reactivity towards water than like "typical" a non-metalloid.
The fact that aluminum foil can be stored in water without any reaction is not proof that aluminum cannot react with water at room temperature. Just
like with phosphorous, if a little NaOH is added the aluminum will begin to rapidly react.
Elemental phosphorous vapor will react with moist silver nitrate paper, causing it to blacken. This suggests to me that the P4 may at least have some
equilibrium hydrolysis with water.
Another example, silicon nano-powder reacts with water at a temperature of only 70-90 °C, releasing hydrogen, even though bulk silicon is essentially
unreactive towards water.
(John Foord at Oxford University)
I just think there may be more to the interraction between water and phosphorous than is commonly realised. I think everyone has just erroneously
assumed for a long time that there is no interraction just because a piece of phosphorous placed in water does not seem to react.
[Edited on 4-6-2012 by AndersHoveland]
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TheAMchemistry87
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i also think this because if phosphorus react with air that has mainly oxygen water is also made with oxygen and will be displaced with hydrgen so it
should react to form an hydrogen phosphide but in a very long time a good idea to test it is boil the water and put molten phosporus if the Ph
decreases it reacted with the water to form hydrogen phosphide so maybe a safer way without any errors of reacting with water it should be kept under
mineral oil
[Edited on 13-6-2012 by TheAMchemistry87]
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AndersHoveland
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| Quote: Originally posted by TheAMchemistry87 | | a good idea to test it is boil the water and put molten phosporus if the Ph decreases it reacted with the water to form hydrogen phosphide so maybe a
safer way |
I am not so sure that this would work. There is obviously some coating that spontaneously forms on the surface preventing further reaction.
Phosphorous pentoxide, for example, does not melt until 340 °C.
I would think the best way to test this would be to first dissolve the phosphorous into N,N-Diisopropylethylamine solvent before seeing if
there is a reaction with water. This would prevent any protective coating from forming. Analogous to dissolving aluminum into liquid mercury amalgam
before adding water.
[Edited on 13-6-2012 by AndersHoveland]
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eastern24
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Elemental phosphorus reacts indeed with water at high temperature 280-300C.
In fact this has been an industrial process of phosphine for more 30years US3371994.
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AJKOER
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Here is a reaction idea that I think would be interesting to perform.
1. Add Ca3P2 powder to the bottom of a tall tower like vessel and cover it with a mesh. Note, Calcium phosphide is cheaply and commonly available due
to its use as a fumigant and phosphine source (see page 163 at http://www.scribd.com/doc/63860142/Anorganic-Reagents ).
2. Add a 1% (or even more dilute) NaOCl solution to the tower which can be sealed for safety as the net reaction (given below) should not result in a
significant gas formation. The dilute hypochlorite solution has been treated with a small amount of acid to bring the solution closer to a neutral
pH.
Possible reactions:
3 Ca3P2 + 18 H2O --> 9 Ca(OH)2 + 6 PH3 (g)
2 PH3 + 3 NaOCl --> 3 H2O + 2 P (s) + 3 NaCl
and the solution of the hypochlorite actually gets more dilute.
Some side reactions:
2 P + 5 NaOCl + 3 H2O --> P2O5 + 3 H2O + 5 NaCl ---> 2 H3PO4 + 5 NaCl
where some water is re-consumed. Now, as the solution is very dilute, no Cl2 is formed, just HOCl from the created acid:
6 NaOCl + 2 H3PO4 --> 2 Na3PO4 + 6 HOCl
6 HOCl + 4 PH3 --> 6 HCl + 6 H2O + 4 P (s)
Net reaction (so far):
3 Ca3P2 + 12 H2O + 14 NaOCl ---> 9 Ca(OH)2 (s) + 6 HCl + 8 NaCl + 4 P (s) + 2 Na3PO4
and the pH of the solution should be monitored to be close to neutral.
Well, at least, this is the theory, and somewhat unbelievable, as it may form free phosphorous at room temperature by a relatively safe procedure. I
must be wrong somewhere.
Comments welcomed.
[Edited on 9-5-2013 by AJKOER]
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Adas
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AJOKER, the NaOCl would probably directly oxidize the Ca3P2 into calcium phosphate (or phosphite).
Rest In Pieces!
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I doubt calcium phosphide is readily available to anyone without a permit and a pest control qualification.
It reacts with water to form highly poisonous phosphine.
Aluminium phosphide, which reacts similarly is highly controlled. It is used for kiiling moles by licensed operators only.
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blogfast25
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| Quote: Originally posted by AJKOER | Here is a reaction idea that I think would be interesting to perform.
3 Ca3P2 + 12 H2O + 14 NaOCl ---> 9 Ca(OH)2 (s) + 6 HCl + 8 NaCl + 4 P (s) + 2 Na3PO4
and the pH of the solution should be monitored to be close to neutral.
[Edited on 9-5-2013 by AJKOER] |
As long as the ‘performing’ isn’t done by AJ, naturellement…
As regards that last pretty equation, how are solid Ca(OH)2, HCl and a lake of water supposed to co-exist, huh?
I see AJ’s temporary exile hasn’t decreased his propensity to try and sell b*llcrap. A bit like his friend ‘Mr H’. The Flat Earth Society
beckons both of you.
[Edited on 9-5-2013 by blogfast25]
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AJKOER
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Blogfast, thanks for your comment as I did want to emphasized, as I noted, that the net equation caption "so far" as, in my opinion, there are several
reactions that could continue. For example, some of the HCl could react with the Ca(OH)2 (as you suggested), or react with the NaOCl (forming more
HOCl as I have also indicated), or directly with Ca3P2 as follows forming more PH3 for processing:
Ca3P2 + 6 HCl --> 3 CaCl2 + 2 PH3 (g)
perhaps more readily than does water.
Also, Adas noted "the NaOCl would probably directly oxidize the Ca3P2 into calcium phosphate (or phosphite". I would bring Adas attention to the fact
that I already have this reaction which I would argue actually proceeds in stages over time similar to the action of HOCl on Sulfur ending in H2SO4. I
would also point out the parallel use of NaOCl as a scrubber for H2S resulting largely in a fine Sulfur suspension. But, I should state, it is my
intention to quickly harvest the Phosphorous to achieve a maximum yield. The equation I referred to is as follows:
2 P + 5 NaOCl + 3 H2O --> P2O5 + 3 H2O + 5 NaCl ---> 2 H3PO4 + 5 NaCl
Note, this reaction requires more NaOCl, but employing a dilute hypochlorite may obviate this path. However, there is one troubling documentation
problem (yes, I search for source references for my suggested preparation) for the assumed reaction:
3 HOCl + 2 PH3 --?--> 3 H2O + 2 P + 3 HCl
as the only source (see "Handbook of Chemistry", Volume 2, by Leopold Gmelin, page 143 http://books.google.com/books?pg=PA143&dq=hypochlorous+a... ), I have found states "and aqueous hypochlorous acid yields phosphoric and
hydrochloric acid. (Balard.) Solution of hypochlorite of lime (chloride of lime) acts in the same way". [EDIT] I did find the following interesting
reference which is possibly supportive (or, upon the rapid addition of some H2O2 to cold dilute NaOCl, as PH3 is also being formed, can be made to
closely replicate the conditions specified), to quote from page 145:
"It [referring to PH3] likewise deposits phosphorus when standing over water (containing air?) in the dark, especially on cooling, and subsequently
does not take fire in the air at ordinary temperatures, excepting when brought in contact with the air in large quantity at once. (Vauquelin.)"
So a possible mechanism here is that oxygenated water in the dark slowly attacks PH3, releasing P into the water, where the Phosphorous is apparently
not as readily attacked by the oxygen containing cold water.
In addition, with respect to alternate preparations, the same source on page 142 states "Bromine precipitates phosphorus from the spontaneously
inflammable gas [PH3], and forms hydrobromic acid, with evolution of heat. (Balard.)—", which suggests a different preparation path involving a
bromide salt in the event that the dilute hypochlorous route prove ineffective.
There is also a major issue that is raised by ScienceSquirrel. While I do have a reference (link provided above) to the claim of available of Ca3P2,
well at least at a point in time perhaps many years ago, this may no longer, in fact, be the case. As I am interested in this preparation route, if
for no other reasons that theoretical feasibility, I will discreetly investigate. Non-US members can do likewise, to see if this restriction
geographically varies.
---------------------------------------
[EDIT] ScienceSquirrel point appears justified. For example (see http://pmep.cce.cornell.edu/profiles/extoxnet/pyrethrins-zir... ) to quote:
"Zinc Phosphide is a Restricted Use Pesticide (RUP). RUPs may be purchased and used only by certified applicators."
[Edited on 10-5-2013 by AJKOER]
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