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teodor
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Quote: Originally posted by Bedlasky  | [
But you are not generating fluorine in the solution. You precipitate CeF4, dry it and than you decompose dry solid. There isn't any solvent involved.
The same goes for drying proces - no elemental fluorine involved during this step. However it is possible that fluoride anions replace chloride in the
SOCl2.
https://en.wikipedia.org/wiki/Thionyl_fluoride |
Yes, I understand that. But drying involves some solvent of crystallization. My point is that it should be a compound in a form XFn where X is any
element, F is fluorine and n is a maximal possible oxidation state of X. Or a mixture of such compounds. Only in this case the fluorine atoms which
are not very tightly bound to Ce has ability not to go with a solvent to which they can have more attraction than to Ce.
I am not talking about getting F2 solution in a liquid, but it would be also interesting experiment with such a solvent.
If one would exclude anhydrouse metal salts which is possible to get with F2 gas only, all other working methods are dependent on some liquid fluoride
- HF or SbF5. You can name it "a solvent" and see this as a necessary condition for any F2 production. This will cover all possible "wet" methods.
For a dry method I can't imagine any gas which can be decomposed to F2 if it is not something you get from F2.
[Edited on 9-2-2025 by teodor]
As an example. 2HF × SiF4 is a liquid, contains 32% HF by weight.
[Edited on 10-2-2025 by teodor]
Probably also I should point out that "the water of crystallization" or, generally "solvent of crystallization" is not just some solvent molecule
catched by a crystall, in case of transitional metals it is unavoidable ligand coming from the reaction medium which could be bound even more tightly
than an acid anion. To make a salt able to give off fluorine we should care about it's coordination sphere, not only about its oxidation state. That's
why we zhould exclude coordination of ghe metal centre with anything which can react with fluorine.
That's why I also mentioned BF3. It is a very strong ligand and fluorine complexes with BF3 are good candidates for thermal decomposition.
[Edited on 10-2-2025 by teodor]
[Edited on 10-2-2025 by teodor]
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clearly_not_atara
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| Quote: Originally posted by woelen |
This is easy, because he uses a chemical, for which fluorine was needed to make it.
An example is decomposition of CoF3 by heating it. This gives CoF2 and F2. But for making CoF3 fluorine is needed, it is not possible to make this
from fluorides only.
Only very recently, a synthetic method was found, in which fluorine can be made without electrolysis of anhydrous HF or molten fluoride salts, and
without using chemicals, which were made with the help of elemental fluorine. This synthesis, however, is not a main route for preparing fluorine, it
is a lab curiousity and is of academic interest, but does not have practical industrial applications. It involves a manganese (IV) intermediate in
concentrated HF and SbF5. Not something for the average home chemist. |
While it's very hard to make fluorine without electrolysis, there was a recent paper about making AgF2 directly through electrolysis without
proceeding via elemental F2:
https://chemistry-europe.onlinelibrary.wiley.com/doi/abs/10....
This would allow some fluorination chemistry -- and it could be slowly decomposed at 700 C to produce F2 in situ by using e.g. a Fresnel lens or a
high-powered laser.
However, my general recommendation for fluorinations is: don't!
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teodor
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SF6 dissociation which release F is starting at 700-900C, and are increasing with the temperature. Because it releases even not F2 but hot F atoms it
reacts immediately with tube and everything inside giving the mix of fluorides and sulfides in a case of metals. I don't know what is the max
oxidation state you can get with that.
The same decomposition occurs with laser and electrical discharge (if you managed to get it).
Of course, some of the gaseous decomposition products could be highly toxic (e.g. S2F10, but who said the fluorine chemistry is tame. You quite can
expect something more potent than phosgene). But considering how many experiments featuring cylinders with SF6 are there on YouTube I suspect you can
buy it in some places. The gas itself is known as a "low voice" gas, so it is not toxic until you break the molecule. Similar to N2 which is quite
inert at low temperatures but even more inert than that.
By the way, thank you for the article, clearly_not_atara. It is very interesting for me as to an armchair fluor chemist.
[Edited on 12-2-2025 by teodor]
[Edited on 12-2-2025 by teodor]
[Edited on 12-2-2025 by teodor]
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j_sum1
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I am not sure I understand the desire to do this without electrolysis.
Or is this purely a theoretical discussion?
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teodor
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I know that after Chemical Force shown some decomposition of higer fluorides people started to wonder is it practical to get fluorine by this method.
Just to discover there is no way to get the fluorides he demonstrated without fluorine. That means there were some investigations but no discussion
here. So, here it is.
Another thing that the original question asked by symboom was not answered. As I pointed somewhere already today people are more interested in safety
discussion than doing actual experiments, so almost any topic about fluorine one hundred and one more time goes to discussion of its danger. And
people who had practice did very valuable comments and I read those with great interest, but also people who never were keeping the thing in their
hands are looking not less informed.
So, the original question was is it possible to get F2 from CuF2 made by the method in the video. The answer is "it depends". Because in aqueous
solution you get CuF2 * 2 H2O which decomposes like this:
132C: 2 CuF2 * 2 H2O -> CuF2 * Cu(OH)F + HF + 3 H2O
420C: CuF2 * Cu(OH)F -> CuO + CuF2 + HF
(see Wheeler, Haendler, J.AmSoc 76 [1954] 263/4)
908 - 950C: 2CuF2 -> 2CuF + F2
(see Wartenberg, Z. anorg. Ch 241 [1939] 391/94 and other works I am lazy to cite now)
The problem is that at 420C you get not a pure CuF2 but a mix with CuO, but it is true only when air is excluded. otherwise only CuO (and it's not
known where 2F goes in this case, Wheeler & Haendler didn't figure out that). It is not quite clear how CuO will react with F2 at 950C (but you
can guess, because one of the method of getting anhydrous CuF2 is reaction of CuO with F2) neither how to separate CuO and CuF2 (and to make it more
complex CuF2 is highly hygroscopic, it is converted to CuF2 * 2H2O on air contact). Also, what you can do with F2 (which is highly dissociated to
atoms already) - I see no practical way to isolate the gas. That's why I see there are 2 completely different topics: isolation and production as a
result of reaction. The second is much easier than the first and you can try to do it getting 2CuF from CuO+CuF2 (but be aware that CuF2 is a liquid
at 785C+ and as such is probably much more dangerous than a good old anhydrous HF because the liquid HF never can release a free fluorine except by
electrolysis. Nevertheless preparing eutetic mixture of CuF2 with fluorides with lower m.p. or, on a less extreme side, dissolving it in some
fluorine-resistant solvent (CCl4 is) could be an interesting experiment). But the first (isolating of F2) is also possible if you can bind F2 to some
"solvent" or "adduct" (for example at -40C F2 forms adduct with pyridine. Which unfortunately explodes above -2C. But there are probably another more
oxidation-resistant compounds). And here I miss the knowledge of some theory which motivates me to search the answers.
So I think the topic is quite interesting by several reasons.
[Edited on 13-2-2025 by teodor]
[Edited on 13-2-2025 by teodor]
[Edited on 13-2-2025 by teodor]
[Edited on 13-2-2025 by teodor]
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chornedsnorkack
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How hard is TlF3 to produce?
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teodor
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By the way, the history of tries to make F2 from CuF2 decomposition is a very old one.
I found the article in "Annales de chimie et de physique" published in 1894 by C. Poulenc with review of all known anhydrous fluorides by the date.
You can download it for free: https://gallica.bnf.fr/ark:/12148/bpt6k34902q/f3.item#
Here is the automatic translation of the part regarding CuF2 decomposition:
Action of gaseous hydrofluoric acid on hydrated fluoride
This process, which can be considered as
an application of the previous one, following the decomposition of the fluoride into oxide under the influence of heat,
also gives rise to anhydrous cupric fluoride.
Properties:
White crystalline powder which, exposed
to air, gradually turns green as it hydrates. This
transformation becomes much faster if the anhydrous fluoride is placed directly in contact with water.
With alcohol and ordinary ether, it turns blue in the manner of
anhydrous copper sulfate, which it closely resembles,
and could, like it, serve as an indicator of the presence
of water in these compounds.
Hydrochloric, nitric and hydrofluoric acids dissolve
it very easily. Sulfuric acid decomposes it when hot, with elimination of hydrofluoric acid and
formation of anhydrous sulfate which is deposited in small crystals, when the sulfuric solution is evaporated in a
sand bath.
Heated in the presence of air, it transforms from 300°
into copper oxide; this transformation is complete and
can be used to determine the copper in this compound.
If, on the contrary, this action is carried out away from the air,
such as, for example, in a platinum tube closed at one
end, it is observed that the cupric fluoride melts and dissociates
in part. But this dissociation is slow
even at 600° - 700°C and the temperature to which the platinum tube must be brought to make it complete makes it impossible
to use this process as a preparation of fluorine.
Hydrogen reduces it with great ease, which
allows it to be used as a method of determination for copper.
Water vapor decomposes it at low temperature into
copper oxide and hydrofluoric acid.
Under the same conditions, with hydrogen sulfide, copper sulfide will be obtained and, with gaseous hydrochloric acid, cupric chloride.
The molten alkali carbonates transform cupric fluoride into copper oxide and atkaline fluoride.
So, as Richard Wagner said "Respect the old masters and your head will be in order".
The magic part is a transformation of CuF2 (even in a form of CuF2 * Cu(OH)F which is relatively easy to get) to CuO on heating. Nobody explained the
process as far as I aware of, so we can assume at this temperature the compound already is a good florinating agent (this is my point about difference
of fluorine isolation vs getting)
TlF3: I am scary even to hear this formula.
For somebody interested in a discussion of a suitable solvent systems, which as I said I think is the most important condition of any chemical F2
generation I'd like to say that I have some new information, but because it looks like this topic has a little interest in relation to the present
discussion I would not share it here.
[Edited on 14-2-2025 by teodor]
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teodor
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I am keeping in my hand the book "Inorganic solid fluorides: Chemistry and Physics" edited by Paul Hagenmuller, Academic Press 1985. This is the last
systematic review of the topic I was able to find. It obviously doesn't contain yet electrochemical method of AgF2 generation, but as I said, it is
the last systematic review I found. Please comment if you know a more recent one.
According to this review the wet methods for known anhydrous fluoride syntheses use those solvents:
- water solution of HF. It is possible to crystallize some double salts in anhydrous state
- BrF3
- BrF5
- IF5
- SeF4
- SbF5
- VF5
- anhydrous HF
- SO2
Another group of methods include heating (fusing) with solids which work also as solvents just having m.p. above the room temperature:
- NH4HF2 and KHF2 which I can consider as a solvent belonging to an anhydrous HF system of solvents
- Hydrazinium Fluoride (what?), N2H6F2 (@210C)
Dehydration of some hydrates just by heating them is also mentioned, e.g.
Fe2F5 * 2H2O -(170C)-> Fe2F5 * H2O -(230C)-> FeF3 * H2O + FeF2 -> (250C) -> FeF3 + FeF2
also CsMnF4 from a dyhidrate at 100C and Rb2MnF5 from a hydrate.
I would omit reactions between 2 solids in a vapor phase here and other reactions type limiting only to "wet" method.
SO2 and molten difluorides are quite accessible methods but the number of explored reactions in SO2 was not very high, just good to know it appears an
inert solvent e.g. for SbF5 and ReF6 when you need to react them with something. For difluorides worth to mention making (NH4)2MgF4 straight from
MgCO3 and (NH4)3VF6 straight from V2O3. May be some limited interest could be paid for the reaction:
RuI3 + 3KHF2 -> K3RuF6 + HI
but I doubt many labs have Ru on their shelves.
Update:
H.J. Emeleus (Simons, "Fluorine Chemistry" vol. I, 1950) provides the list of hydrated binary fluorides which could be dehidrated by heat only:
- ZnF2, CdF2, NiF2, CoF2, FeF3, CrF3
[Edited on 21-2-2025 by teodor]
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teodor
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I continue to check various preparative methods for anhydrous fluorides. The general remark from the book "Fluorine chemistry" by Simons:
"Reaction with fused salts: this type of reactions has not been extensively applied, possibly because of the abailability of simpler methods".
Well, if you have a fluorine lab with a special vacuum line probably it is easier to use F2. For amateurs, as I already said, NH4F or NH4HF2 are quite
convenient compounds. So, probably we can investigate what is possible to get with them, because this method had never bean "extensively applied".
As an example, the compound in question, CuF2 could be generated this way:
"Hydrated fluoride is fused with NH4F and the product is heated to 260C in a steam of CO2 to remove the excess ammonium salt".
This is interesting, because it uses volatility of ammonium cation and stream of CO2 to remove it and not HF as in classical preparation.
So, just to give you impression that there could be always some different method.
Update:
Another reaction which gives F2 is this (R. Salih Hisar, Bl. Soc. chim. 1952 308):
2NaF + (NaPO3)2 + 1/2 O2 -> Na4P2O7 + F2
It happens at 650 - 750C. 650C is the max working temperature of Nickel in regard to fluorine.
Probably the main question for this route, how to get (NaPO3)2.
[Edited on 21-2-2025 by teodor]
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metalresearcher
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| Quote: Originally posted by teodor | I continue to check various preparative methods for anhydrous fluorides. The general remark from the book "Fluorine chemistry" by Simons:
"Reaction with fused salts: this type of reactions has not been extensively applied, possibly because of the abailability of simpler methods".
Well, if you have a fluorine lab with a special vacuum line probably it is easier to use F2. For amateurs, as I already said, NH4F or NH4HF2 are quite
convenient compounds. So, probably we can investigate what is possible to get with them, because this method had never bean "extensively applied".
As an example, the compound in question, CuF2 could be generated this way:
"Hydrated fluoride is fused with NH4F and the product is heated to 260C in a steam of CO2 to remove the excess ammonium salt".
This is interesting, because it uses volatility of ammonium cation and stream of CO2 to remove it and not HF as in classical preparation.
So, just to give you impression that there could be always some different method.
Update:
Another reaction which gives F2 is this (R. Salih Hisar, Bl. Soc. chim. 1952 308):
2NaF + (NaPO3)2 + 1/2 O2 -> Na4P2O7 + F2
It happens at 650 - 750C. 650C is the max working temperature of Nickel in regard to fluorine.
Probably the main question for this route, how to get (NaPO3)2.
[Edited on 21-2-2025 by teodor] |
I have some (NaPO3)6 ordered from ebay, so that is rather easy to obtain. It is the six-mere instead of the dimer, but I don't think that is an issue,
it is the same salt.
I got that for isolating another dangerous element: white P4.
[Edited on 2025-2-21 by metalresearcher]
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teodor
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I have no access to the original publication HISAR, R. (1952). * NOTE SUR LA DECOMPOSITION DES HALOGENURES ALCALINS SOUS LACTION DU METAPHOSPHATE DE
SODIUM. BULLETIN DE LA SOCIETE CHIMIQUE DE FRANCE, 19(3-4), 308-308. So, I don't know what is the point to make equation with the dimer. I found the
citation in another book.
Let's suppose (NaPO3)6 can work.
12 NaF + 2 (NaPO3)6 + 3 O2 -> 6 Na4P2O7 + 6 F2
It was mentioned to heat the mixture in open air.
So, I suppose the experiment could be performed in a nickel crucible and above the mixture could be fixed some compound to detect fluorine. Quite
straightforward except the detection part which hardly can proof we get fluorine itself and not some compound of fluorine. But at least we can try.
Above 650C I assume nickel can react with F2. Covering the bottom with NaF and putting NaF/(NaPO3)6 mixture on top should make some barrier. Also
there are another ideas how to make F2 resistant high temperature crucible I can share. But I have no (NaPO3)6 and my fume hood is in the process of
moving from my old lab to my new lab, so I am unable to perform the experiment myself.
I am very interested in getting the original french article to better understand the method.
[Edited on 21-2-2025 by teodor]
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clearly_not_atara
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Anhydrous FeF3 may bee had by rxn of anhydrous FeCl3 with anhydrous HF. This compound thermally decomposes to the difluoride:
https://academic.oup.com/bcsj/article-abstract/76/6/1165/734...
However, pretty much every step in this process is extremely dangerous and should not be attempted.
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teodor
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So, one more production method.
In some review I encountered the remark "there are million of chemical methods to produce fluorine but none of them has a reliable proof it can be
used to produce fluorine gas".
The problem of isolation and keeping is the killer.
(As for electrolitical cells, they can explode, it depends on construction & knowledge, so don't tell it is an accessible method).
Fluorine has very poor solubility in any solvents except when it forms adducts with some additional dissolved substance. The most useful is the adduct
with pyridine. It is possible to get it using e.g. CsCoF4 or KCoF4, so no decomposition is required.
Another option is to prepare an eutetic (liquid) mixture of fluorides with high and normal oxidation state. This way you can get mobility of fluorine
atoms in a liquid which is quite similar to solution.
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j_sum1
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Again, I am not sure I understand the desire to avoid electrolysis.
Labcoatz recently performed electrolysis on potassium bifluoride in a specially constructed copper cell with graphite electrodes. https://www.youtube.com/watch?v=IcC8_CX9ud0
That is about as accessible as you can get for such a dangerous chemical.
One of the advantages of electrolysis is the ability to switch off and halt production immediately -- something that is significantly less difficult
than quenching a reaction.
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metalresearcher
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I stumbled upon an even better F2 video: https://youtu.be/UzIH6raTxyE . One of the best chemistry porn I have ever seen !
Here three chemical Youtubers (NileRed, Fire&Explosions and Advanced Tinkering together in a professional fluorine lab). I miss Cody's Lab here in
...
When is the first ClF3 video ? ClF3 is even more aggressive than F2 because it is a liquid at (almost) ambient temperature (bp = 12 C).
Even the Nazis were defeated by its aggressiveness, they built Falkenhagen bunker in 1945 to produce several tons a month for chemical warfare, but it
was too aggressive.. Shortly after, the war was over.
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teodor
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| Quote: Originally posted by j_sum1 | Again, I am not sure I understand the desire to avoid electrolysis.
Labcoatz recently performed electrolysis on potassium bifluoride in a specially constructed copper cell with graphite electrodes. https://www.youtube.com/watch?v=IcC8_CX9ud0
That is about as accessible as you can get for such a dangerous chemical.
One of the advantages of electrolysis is the ability to switch off and halt production immediately -- something that is significantly less difficult
than quenching a reaction. |
Thank you for the video, I was not aware of it.
It is not easy to "switch off" and halt fluorine production. The main problem with electrolisys it always accumulates dangerous amount of fluorine in
the electrolitic cell and tubes. If fluorine will come in contact with hydrogen in the other part of the apparatus it will cause explosion. It can
happen also if the output line is blocked and the fluorine goes to the second part of the cell by pressure. Carbon particles detached from electrodes
also can cause little explosions. To shut down the cell all parts which could contain fluorine should be flushed with nitrogen, untill that it is just
a little bomb and if the electrolite is still melted a very dangerous one. This type of problems make the process quite complex to handle out of
professional lab. For people who want safe experimenting with microquantities of wild fluorine it would be nice to have a different method.
Searching for different method demands a research and understanding properties of the gas and its compounds and this is the most interesting part of
it.
Fluorine reacts with water vapors forming mixture of HF, OF2, O3 and O2. Generally 1 mol of water destroys 2 mols of fluorine. So, all experiments
with microquantities require very dry environment.
It is not quite clear at which temperature it reacts with nitrogen and wether different transition metal compounds can activate nitrogen to react with
fluorine at lower temperatures. They potentially can, so decomposition of transition metal fluoride not necessarily can give accessible F2, it could
be NF3 if the experiment is performed not in vacuum.
[Edited on 24-2-2025 by teodor]
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teodor
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| Quote: Originally posted by Bedlasky | | CeF4 thermally decompose to CeF3 and F2. You can precipitate CeF4.H2O from aqueous solution, problem is that monohydrate produce during decomposition
little to no fluorine, water is oxidized instead. If you come with way how to dehydrate CeF4.H2O first (mayble SOCl2?), you could obtain CeF4 and
thermally decompose it. |
I checked this, the information probably is not correct.
Indeed, there was an old article in Russian Journal of Inorganic Chemistry, Batsanova et al, 18 [1973] 476/8. I don't have it (you can try to find it
probably in the library genesis). Gmelin only mentiones that there was detected an "endotermic effect" on the thermogram @ 270C which "was interpreted
as reaction 2CeF4 -> 2CeF3 + F2". But this was only some speculative hypothesis which was not confirmed by direct experiment in argon or vacuum.
CeF4 doesn't decompose even at much hier temperatures. See the attached articles.
There is one more reason why I have attached those. It contains also proof of applicability of "fusing with NH4F" dehidration method mentioned by me
few time already to partial (non stoichiometric) dehydration of CeF4 * H2O. It should be probably mentioned that non-stoichiometric compounds are
ordinarily things among fluorides due to the fact they often had very stressed/irregular/damaged crystalls (and hexafluorides have no crystalls at all
being just molecular solids).
Attachment: asker1965_1.pdf (490kB)
This file has been downloaded 26 times
Attachment: asker1965_2.pdf (337kB)
This file has been downloaded 25 times
Comparing the false hypothesis of CeF4 decomposition at @270C with that strange behaviour of CuF2 @300C in the open air (100% convertion to CuO2) I
now have my own hypothesis on reaction of higher transitional metal fluorides with air in 250-300C range.
Update:
I hope some inorganic chemists will enjoy by being aware of the reaction:
Ce2(SO4)3 + Na2SiF6 + 4H2O -> 2CeF3 + Si(OH)4 + Na2SO4 + 2H2SO4
(Tsubaki, Namikoshi, Bunseki Kagaku 20 [1971] 781/3)
It's interesting not only because it produces anhydrous CeF3 and uses SiF6(2-) as a source of anhydrous fluoride, but because 1 mol of dry mixture can
react with 4 mols of water producing ... a dehydration agent.
[Edited on 27-2-2025 by teodor]
[Edited on 27-2-2025 by teodor]
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teodor
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Some easy reading for today:
Attachment: ruff_eng.pdf (3.3MB)
This file has been downloaded 33 times
This is an automatic translation. The German original is here:
https://onlinelibrary.wiley.com/doi/10.1002/zaac.19160980103
This is only a part of the full story.
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teodor
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Some comments regarding my yesterday post about PbF4 decomposition.
The first fact that PbF2 in a solid form is a conductor of electricity, that fact was discovered by Faraday. The reason is that F- ions have mobility
in it. Some fluorides are "fast ionic conductors" or "solid electrolites". They could be seen as a solid solution of fluorine ions. This could
explain, e.g., the fast reaction with various materials and air (water vapours) on the surface of crystalls.
As Otto Ruff mentioned, molten PbF4 * X salt easily fluorinates platinum.
There was another story. Any attempts to do experiments with heating pure PbF4 were failed because in the process of heating some PbF4 is converted to
PbF2 and the mixture of those two is an eutetic one. It's already a liquid at not so high temperature (didn't find the data of the melting point but
there is a data of required PbF2/PbF4 relation). The very high mobility of F- ions in PbF2 in this liquid is combined with their excess in PbF4. As a
result the experiments with that liquid just ended with destroyed apparatus and probably nobody even measured its melting point as well as chemical
properties. I think they are close to that what a concentrated fluorine solution could have at that temperature.
There are few more interesting topics regarding getting fluorine from fluorides. There were many attempts to decompose CeF4 and all those attempts are
just explained with 2 notes published in 1881 by Oscar Löw. He provided the first proof that the mineral Antozonite contains a free florine. He
mistaken with the explanation of the phenomenon but as a real genius this error didn't lead to a wrong conclusion.
Almost 150 years chemists doubted his results. Only in 2012 the reliable proof was gotten:
Attachment: kraus2012.pdf (658kB)
This file has been downloaded 29 times
The modern explanation is that it is because of uranium impurities. Beta radiation splits CaF2 to elemental Ca and F2 and due to fluorite crystal
properties they both could be accumulated in the crystall without reacting with each other. When somebody breaks or just rub the mineral it "stinks"
of fluorine.
But Löw had hypothesis that it is no because of U (the beta-decay was discovered about 10 years later) but because of CeF4 which is prone to
decomposition to free fluorine. And he found that exactly those fluorites which contain Ce have fluorine smell (and they had uranium as well). This
error is exactly why many chemists did experiments with CeF4 decomposition.
It's interesting that according to the experiments CeF4 * H2O and not anhydrous CeF4 gives traces of F2. So, decomposition is started by water but
looks like can give small detectable amount F2 as a byproduct.
Probably Löw experiments worth to attach to this thread to show why CeF4 is still mentioned as a chemical method of generating F2.
The German original:
Attachment: loew1881_1.pdf (157kB)
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Attachment: loew1881_2.pdf (95kB)
This file has been downloaded 23 times
The automatic English translation:
Attachment: loew1881_1eng.pdf (451kB)
This file has been downloaded 25 times
Attachment: loew1881_2eng.pdf (254kB)
This file has been downloaded 24 times
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teodor
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| Quote: Originally posted by metalresearcher | | Quote: Originally posted by teodor |
Another reaction which gives F2 is this (R. Salih Hisar, Bl. Soc. chim. 1952 308):
2NaF + (NaPO3)2 + 1/2 O2 -> Na4P2O7 + F2
It happens at 650 - 750C. 650C is the max working temperature of Nickel in regard to fluorine.
Probably the main question for this route, how to get (NaPO3)2.
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I have some (NaPO3)6 ordered from ebay, so that is rather easy to obtain. It is the six-mere instead of the dimer, but I don't think that is an issue,
it is the same salt.
I got that for isolating another dangerous element: white P4.
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I didn't find any of 1952 journals but I've got 1951.
There are 2 articles by the turkish chemist here.
It looks like the dimer is essential. In the first article of the November issue Hisar provides the method to synthesize it.
Reaction which evolves fluorine is a logical continuation of his experiments published in 1951. Thus, he was studying similar reaction
2NaNO3 + Na2P2O6 -> Na4P2O7 + 2NO2 + 1/2 O2
and found it is reversible, so N2O5 could be used to get the dimer. Then he discovered that other acid anhydrides work similar way.
I attach both original French paper and the automatic translation to English. Google Translate omits some indexes in chemical formulas, so look into
both.
Attachment: hisar.pdf (7.5MB)
This file has been downloaded 27 times
Attachment: hisar1951_1eng.pdf (1.5MB)
This file has been downloaded 25 times
Update: from the 1951 article is still not evident is the dimer is essential, but because his discovery about fluorine was published the next year I
assume it is.
But if somebody has hexametaphosphate why wouldn't try its reaction with NaNO3 to get a clue.
(General remark. Sometimes fluorine could be replaced with oxigen and oxigen couldn't be replaced with fluorine. It's because oxigen use additional
bonds (pi type) from its p orbital and fluorine doesn't. That's why there are solid electrolites with fluorine, it is because its bonding is a pure
ionic one).
[Edited on 5-3-2025 by teodor]
[Edited on 5-3-2025 by teodor]
[Edited on 5-3-2025 by teodor]
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teodor
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Posts: 1001
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Unfortunately the interest to metaphosphates faded during 1950x, so the Hisar's article never got serious review. And there is really a lot of
questions here. The structure of the acid was given incorrect, the existence of the dimer in the polyphosphate melt at suggested temperatures was
never checked with modern techniques but nobody said it is impossible either. Some author statements looks quite speculative (e.g. the reaction with
NaNO3), at least without checking other his reports, but they were published only locally.
In few words: still a big field for amateur experiments.
I wouldn't make this comment if accidentally I found some indirect proof of the possibility to get fluorine this way.
In parallel to metaphosphate chemistry the chemistry of fluorophosphates was developed, mainly by Willy Lange and his group (at the same time Sanders
discovered the toxic action of their esters, but it is another story).
In vol.2 of Inorganic Synthesys Willy Lange published procedures for preparation of ammonium monofluophosphate and difluophosphate (by heating to 135C
the mix of P4O10 with NH4F). And in vol.3 Audrieth (who also tried to systematize the chemistry of phosphates, see his excelent article https://pubs.acs.org/doi/pdf/10.1021/ed025p80 ) gave another procedure:
Na3P3O9 + 3NaF -> 3Na2PO3F (@800C).
It is very similar to Hisar's claim except the temperature.
This is the first part of the story.
The second part, that
"Na2PO3F melts at 625C and on further heating is decomposed to pyrophosphate". I found this information in different sources but not sure about the
experimental details/first reference, need to search. There is some additional detail: "also in wet air with evolution of HF". But nobody informs what
happens in dry air.
So, I suppose there is some field of experiment involving heating Na2PO3F melt with different metals/salts to see how much oxidation effect in dry
conditions could be.
Later comments:
It was actually Lange and his assistance who first discovered the toxic effect of diesters. Schrader was aware of this work. The method of getting
esters from nonvolatile acids was used to determine their molecular weigh (by vapours density). It was extensively used for metaphosphate
characterisation.
It was a very serious doubt about safety of usage Na2PO3F and some extensive studies were performed. After getting the result that the toxity of the
salt of this acid is tame some company patented the usage of Na2PO3F in toothpaste. So different story for so close chemicals.
[Edited on 7-3-2025 by teodor]
[Edited on 7-3-2025 by teodor]
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