AndersHoveland
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obscure chemistry of phospine and hypophosphite
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"...it is known that phosphorous acid, reduced with nascent hydrogen*, yields phosphine and hypophosphorous acid
For the formation of phosphine from phosphorous and sodium hydroxide, heated together, the following equation is usually given:
(4)P + (3)NaOH + (3)H2O --> PH3 + (3)NaH2PO2
Without a doubt the reaction is far more complicated.
...what has not been shown, that hydrogen and at least two acids of phosphorous are formed when phosphine is generated in the usual way. ...When
phosphine was obtained, by heating an aqueous solution of potassium hydroxide with phosphorous, only from 10 to 40 percent of the gas obtained was
phosphine [the remainder being hydrogen]. The decomposition of phosphonium iodide, PH4I, is the only method of obtaining pure phosphine.
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"investigation of sodamide and of its reaction-products with phosphorus" William Phillips Winter p42-43
*note that "nascent hydrogen" probably meant zinc in the presence of hydrochloric acid, although it is now known that these reducing regents do not
actually produce any free hydrogen atoms.
The article goes on to say 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.
Hydrogen sulfide is known to be able to reduce sulfur dioxide (actually sulfurous acid in aqueous solution) at room temperature. Perhaps phophine
would similarly be able to reduce phosphorous acid (H3PO3)?
Quote: Originally posted by AndersHoveland  | 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.
Journal of the Chemical Society, Volume 75 (Great Britain) p.976
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As already mentioned in my previous post, it seems apparent that hypophosphorous acid, H3PO2, is reduced by phosphine, PH3, to elemental phosphorous.
The selective oxidation of PH3 by aqueous iodine also produces H3PO2. (in a similar reaction, iodine reduces thiosulfate, S2O3[-2], to tetrathionate,
S4O6[-2] )
The fact that H3PO2 even exists suggests that the chemistry of the P-O bond is more similar to the C-O bond than the S-O bond, which is to say that
H3PO3 is essentially not an oxidizer at ordinary temperatures, like SO2 can be. However, one paper mentions that PH3 can reduce 1-naphthol to
naphthalene: "Phosphine as a Reducing Agent"
SHELDON A. BUCKLER, LOIS DOLL, FRANK K. LIND, MARTIN EPSTEIN. J. Org. Chem., 1962, 27 (3), pp 794–798
So the idea that PH3 could reduce H3PO3 is quite plausible.
Whether PH3 reacts with H3PO3 in aqeous solution is uncertain.
(the paper also describes the reaction of phosphine with nitrobenzene. There was no reaction at neutral conditions, but when sodium hydroxide was
added then azoxybenzene was produced in high yield)
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When heated by itself, hypophosphorous acid is resolved into phosphoric acid and phosphine. Phosphorous acid on heating at 200°C converts to
phosphoric acid and phosphine.
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It should be obvious that PH3 would not reduce H3PO4, just as H2S does not reduce aqueous H2SO4.
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When boiled in alkaline liquids [aqueous KOH] they [hypophosphites] are decomposed into phosphates and hydrogen.
KPH2O2 + (2)KHO --> K3PO4 + H4
The dry salts decompose when heated, giving phosphoretted hydrogen [PH3] (hence they are very flammable) and leaving a residue of pyrophosphate.
(2)BaP2H4O4 --> (2)PH3 + H2O + Ba2P2O7
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"A dictionary of chemistry and the allied branches of other sciences, Volume 4", Henry Watts, p524
Various Other Hydrides of Phosphorous
P2H4, first obtained by P. Thenard (Comptes rendus, 1844, 18, p. 652) by decomposing calcium phosphide with warm water, the products of reaction being
then passed through a U tube surrounded by a freezing mixture (see also L. Gattermann, Ber., 1890, 23, p. 1174). It is a colourless liquid which boils
at 57°-58° C. It is insoluble in water, but soluble in alcohol and ether. It is very unstable, being readily decomposed by heat or light. By passing
the products of the decomposition of calcium phosphide with water over granular calcium chloride, the P2H4 gives a new hydride, P12H6 and phosphine,
the former being an odourless, canary-yellow, amorphous powder. When heated in a vacuum it evolves phosphine, and leaves an orange-red residue of a
second new hydride, P9H2 (A. Stock, W. Bottcher, and W. Lenger, Ber., 9 9, 4 39, 47, 2853).
P4H2, first obtained by Le Verrier, is formed by the action of phosphorus trichloride on gaseous phosphine (Besson, Comptes rendus, 111, p. 972); by
the action of water on phosphorus di-iodide and by the decomposition of calcium phosphide with hot concentrated hydrochloric acid. It is a yellow
solid, which is insoluble in water. It burns when heated to about 200° C. When warmed with alcoholic potash it yields gaseous phosphine, hydrogen and
a hypophosphite.
Interesting Non-Reactivity of Phosphorous:
Elemental phosphorous can apparently be dissolved in liquid sulfur dichloride without reaction.
[Edited on 2-12-2011 by AndersHoveland]
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woelen
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I have serious doubt about the reaction between PH3 and phosphorous acid. That would mean that one can make elemental phosphorus in aqueous solution.
I know of the reaction between H2S and SO2 (I actually tried this myself and they react, giving a cloudy liquid, due to finely divided sulphur), but
PH3 and H3PO3 are a different matter. I have done experiments with phosphite and although literature describes this as a reductor, it is a very crappy
one. It's reactions are VERY slow and I hardly managed to get anything reduced by H3PO3. Hypophosphite also is a sluggish reductor, but on heating it
becomes more reactive. But still, also hypophosphite is not really impressive, due to its slow reactions.
I also know of the yellow and reddish stuff, which settles on the glass of tubes in which PH3 (made from white P in warm NaOH solution) is stored. It
takes about one day to form such deposits and the gas volume then has decreased by 10% or so. The remaining gas does not react anymore. I think that
these yellowish and reddish compounds are ill-defined non-stoichiometric oxides or mixed oxides/hydrides of phosphorus, which are very rich in
phosphorus.
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AndersHoveland
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| Quote: Originally posted by woelen | | I also know of the yellow and reddish stuff, which settles on the glass of tubes in which PH3 (made from white P in warm NaOH solution) is stored. It
takes about one day to form such deposits and the gas volume then has decreased by 10% or so. The remaining gas does not react anymore. I think that
these yellowish and reddish compounds are ill-defined non-stoichiometric oxides or mixed oxides/hydrides of phosphorus, which are very rich in
phosphorus. |
Is it possible that this red/yellow substance is elemental phosphorous itself? The literature describes phosphorous forming as a yellow to reddish
flocculent substance in aqueous solution.
Although I also found this:
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"It is extremely doubtful whether phosphorus could be obtained similarly by incomplete oxidation of phosphine..."
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"Chemical periodicity" Robert Thomas Sanderson 1960.
[Edited on 2-12-2011 by AndersHoveland]
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AndersHoveland
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| Quote: |
Precipitated cuprous oxide [Cu2O] is acted on rapidly by phosphine at ordinary temperature, forming a grey white mass and water, the grey substance is
quite insoluble in water, if air be excluded; it is [PCu3]. It melts at a red heat, is rapidly dissolved by nitric acid or bromine water, and is
attacked by hot sulfuric acid, with the formation of sulfur dioxide and phosphine. It does not reduce potassium permanganate.
While phosphine acts differently on the aqueous solutions of different cupric salts, its action of the ammoniacal solutions of all of them (chloride,
sulfate, nitrate, acetate, formate, hydroxide) is the same. Copper phosphide is formed in amount corresponding to two-thirds of the phosphine which
disappears, while the phopshorous of the other third is founf in the liquid, as phosphoruc and hypophosphorous acids in molecular proportions-
(6)PH3 + (12)CuCl2 + (6)H2O + (x)NH3 --> (4)PCu3 + PO4H3 + PO2H3 + (24)HCl + (x)NH3
If, after the absorption of the phosphine, oxygen be admitted, the precipitate redissolves completely, using up a volume of oxygen double that of the
phosphine absorbed in the first place.
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Journal of the Society of Chemical Industry, Volume 18, p716
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AJKOER
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Here is an interesting find (link: http://chemistry.mdma.ch/hiveboard/acquisition/000396783.htm... ), but more on the electrochemistry side:
"electrolysis of phosphoric acid Bookmark
This is something I found in a 1930´s chemistry book a couple of days ago:
"Dill obtained phosphorus by electrolisys of hot 65% phosphoric acid mixed with 20-25% of powdered carbon using 80-150Amp and 120V(...)"
Add the comment by AndersHoveland:
"As already mentioned in my previous post, it seems apparent that hypophosphorous acid, H3PO2, is reduced by phosphine, PH3, to elemental
phosphorous."
So, it may indeed be possible to form elemental phosphorous via the electrolysis of H3PO4 and C, being that the acid is reduced by the nascent H2 to
say, phosphorous acid, H3PO3, then, hypophosphorous acid, H3PO2,...There is also, with heating, the formation of PH3 and reactions therewith.
My take on some possible reactions involving nascent hydrogen (H2*) per electrolysis and heat:
H3PO4 + H2* --> H3PO3 + H2O
Then, per Wikipedia: "Phosphorous acid on heating at 200 °C converts to phosphoric acid and phosphine:
4 H3PO3 → 3 H3PO4 + PH3 "
Then, a reaction cited by AndersHoveland:
PH3 + H3PO3 -->PH4H2PO3 --> 2P + 3 H2O
......
Now, the role of the carbon may be electrochemical, or, it could be there to capture PH3 gas, and thereby facilitate the reduction process.
I will let others extrapolate on other possible reactions as the electolysis proceeds (I do believe this is a good exercise as it reveals any
necessary conditions, as for example above, the temperature at which PH3 forms). Comments as to how this all could be yet another outbreak of
sciencemadness are welcomed as well.
Now, per Woelen remarks, in the best case, I would not expect good yields, but one could be surprised nevertheless! The bad news is that if PH3 is
really an evolving gas in the electrolysis due to heating, one really has to use ventilation and protective breathing equipment (nothing is actually
without cost).
[Edited on 17-4-2013 by AJKOER]
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blogfast25
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AJ: what happened to:
"Also, I am not a subscriber to a unified nascent hydogen theory, but, at the same time, I do want to dismiss all the observed results as without
basis either."
For someone who's 'not a subscriber' you don't half invoke it a lot of lately.
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AJKOER
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Blogfast:
I recall reading that the electrolysis of hypophosphorous acid, H3PO2, liberates P, or I would write:
H2PO2 + H2* --> 2 H2O + P
If you want to explain (or depict) the reaction otherwise, please proceed. I stay with a convenient & understandable short hand.
---------------------------------------------------------------------------
Also,on the reaction:
PH3 + H3PO3 -->PH4H2PO3 --> 2P + 3 H2O
I should have mentioned, even in the presence of Carbon to hopefully hold some of the PH3, stirring may improve the yield.
[Edited on 18-4-2013 by AJKOER]
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elementcollector1
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Does this really mean I could electrolyze phosphoric acid to get phosphorus? It seems that if it were that easy, someone would have done it
ages ago...
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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AndersHoveland
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| Quote: Originally posted by AJKOER |
This is something I found in a 1930´s chemistry book a couple of days ago:
"Dill obtained phosphorus by electrolisys of hot 65% phosphoric acid mixed with 20-25% of powdered carbon using 80-150Amp and 120V(...)"
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I do not think that has anything to do with the chemistry I was referring to in the first post. Seems to me that the reaction you are describing
required a substantial ammount of heat.
| Quote: Originally posted by elementcollector1 | | Does this really mean I could electrolyze phosphoric acid to get phosphorus? It seems that if it were that easy, someone would have done it
ages ago... |
The source I cited described the reduction of phosphorous acid. Phosphoric acid is much more difficult to reduce.
Same thing with sulfuric acid. Sulfurous acid (SO2 dissolved in water) is easy to reduce, but dilute sulfuric acid is nearly impossible to reduce. In
the reaction of zinc with dilute sulfuric acid, none of the sulfate ions get reduced. A similar situation exists with dilute perchloric acid.
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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woelen
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Indeed, the highest oxidation state salts and acids usually are much less reactive than the salts (and acids) of lower oxidation states:
- nitrite is more reactive than nitrate
- sulfite is more reactive than sulfate
- chlorate is much more reactive than perchlorate
- bromate is more reactive than perbromate
- phosphite is more reactive than phosphate, although the difference is less than with the other examples above.
The lower oxidation state material can both act as oxidizer or as reductor, although oxidation of chlorate and especially bromate is not easy at all.
Nitrate and sulfite can easily be oxidized, phosphite can fairly easily be oxidized, albeit somewhat slowly.
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AJKOER
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If one is generating some PH3 as an problematic side product (see MSDS highlighting toxicity at https://www.google.com/url?sa=t&source=web&rct=j&... ), or otherwise, may I suggest scrubbing it in a large vessel filled with chlorine
(or safer, have the PH3 generator sitting in a large container of chlorine). Typical cited reaction:
PH3 + 3 Cl2 = PCl3 + 3 HCl (See, for example, https://books.google.com/books?id=5Hf-uK58OZIC&pg=SA5-PA... )
But as historical texts describe this reaction as occurring with a flash of light (see, for example, discussion at https://books.google.com/books?id=S_gJAAAAIAAJ&pg=PA412&... ), I would believe that this reaction could possibly proceeds in two steps with
the liberation of phosphorus and its subsequent combustion in chlorine:
2 PH3 + 3 Cl2 = 2 P + 6 HCl
2 P + 3 Cl2 = 2 PCl3
As such, even with the correct proportions of gases, the product is not likely, in my opinion, pure PCl3 due to uneven combustion and also:
PCl3 + Cl2 = PCl5
could occur due to randomness in gas mixing.
Note: PCl5 reacts with dry NH3 gas forming a stable nitride, which on heating to 850 C, proceeds to liberate P and N2 (see https://en.m.wikipedia.org/wiki/Triphosphorus_pentanitride ).
[Edit] Just found this from Atomistry on PCl3, to quote:
" It [PCl3] chlorinates arsenic (in the presence of a little AsCl3) at 200° to 300° C., also antimony, phosphine and arsine, giving phosphorus in
each case. "
Which implies the reaction:
PCl3 + PH3 = 2 P + 3 HCl
Link: http://phosphorus.atomistry.com/phosphorus_trichloride.html
Also, interesting is the reaction of dry NH3 on PCl3 forming the unstabe P(NH2)3 at very low temperatures, the implied reaction being:
PCl3 + 3 NH3 = P(NH2)3 + 3 HCl
To quote from Atomistry on this compound:
"Phosphorus Triamide, P(NH2)3, is said to be formed by the action of dry ammonia on phosphorus tribromide at -70° C. It was a yellow solid which
decomposed at 0° C. giving a brown substance, diphosphorus triimide, P2(NH)3, and was further decomposed on further heating into phosphorus, nitrogen
and ammonia."
Link: http://phosphorus.atomistry.com/phosphorus_triamide.html
So, upon warming to room temperature, under an inert gas (like CO2), P(NH2)3 may produce P, N2 and NH3. Or, the visible products of the direct
reaction of dry NH3 and very cold PCl3 could include phosphorus (which may further react) and clouds of NH4Cl, hopefully not in the form of an
explosion. Best would be with NH3 dissolved in cold CCl4 acting on PCl3 as to quote from Atomistry on PCl3 again:
"By the action of ammonia on phosphorus trichloride in carbon tetrachloride ammoniates such as PCl3.6NH3 and PCl3.8NH3 have been obtained. On heating,
these ammines are decomposed with the formation of a phosphamide and ammonium chloride. "
[Edited on 7-4-2017 by AJKOER]
[Edited on 7-4-2017 by AJKOER]
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