Sciencemadness Discussion Board

Doable synthesis of F2 in amateur setting?

woelen - 3-2-2018 at 13:15

An interesting read I found is the following:

It states formation of K2MnF6, precipitating from an aqueous solution with 50% HF in it.

The reaction equation is this:
2 KMnO4 + 2 KF + 10 HF + 3 H2O2 --> 2 K2MnF6 + 8 H2O + 3 O2

I have some 48% HF and I have KHF2 (of course I also have KMnO4 and can make quite concentrated H2O2, 30% should be achievable from 12% H2O2). The KHF2 brings in additional HF. So, with the chemicals I have it should be possible to do this.

I am still somewhat reluctant to do this experiment. Although I have 48% HF, the number of experiments I did with this stuff is almost zero. It is so badly toxic and scary. This experiment, however, looks really interesting. I have read papers about making K2MnF6, but they used SbF5 as one of the reactants, totally beyond reach of an amateur-setting, but the chemicals, mentioned in the link above, can be handled in an amateur-setting, albeit that with HF one must be extremely careful.

I have PE-testtubes and beakers. I think I'll give it a try with freezing-cold solutions. If there is anyone here who has done something similar, then I would like to know of that experience. Practical info and theory on paper may make a big difference.

aga - 3-2-2018 at 13:21

Doubtless this does not need saying : be VERY careful !

Our favourite Amateur Chemist needs to preserve himself (while doing the amazing chemistry).

woelen - 3-2-2018 at 13:28

Yes,thanks for your concerns. I will be careful. I know of the dangers of HF. But I think it is doable. I only will do this at ml scale (I only have a little under 100 ml of 48% HF), but I am aware that even ml quantities of HF are dangerous. I especially take utmost care not to get any of this on my skin and I will have a solution of CaCl2 around, ready to grab, next to the experiment to convert the fluoride to insoluble and non-toxic CaF2.

clearly_not_atara - 3-2-2018 at 13:33

Doesn't the synthesis of F2 from fluoromanganates also depend on SbF5?

2 K2MnF6 + 4 SbF5 >> 4 KSbF6 + 2 MnF3 + F2

But maybe you can get away with using PF5 instead of SbF5? PF5 can be made from arsenic trifluoride and phosphorus pentachloride, eliding the need for anhydrous HF:

3 PCl5 + 5 AsF3 >> 3 PF5 + 5 AsCl3

But AsF3 is also generally made with anhydrous HF. But since it is not a Lewis acid you can probably use silver fluoride instead:

3 AgF + AsCl3 >> 3 AgCl + AsF3

AsCl3 can in turn be made by chlorinating arsenic with excellent selectivity. Of course the chemicals here are also toxic. But you knew that.

Vosoryx - 3-2-2018 at 13:39

What would you like written on your headstone? Best to let us know now, otherwise we'll just put "Played with F2"

Don't know enough to really be able to weigh in on this, but looks good. Hope it works out, 'cause i'm sure not going to try it!

Reboot - 3-2-2018 at 14:30

Interesting. Is this reaction is being driven by the acid helping to peel the oxygen atoms off the manganate, with F+ ions then able to associate with it, and the resulting manganate fluoride crashing out of solution (thus driving the reaction forward by removing product from solution....)

Or am I underestimating the oxidizing power of the acid/peroxide mix and it progresses through fluorine (F2 or F• radicals) after all?

SWIM - 3-2-2018 at 15:05

I hope you will consider looking into buying some calcium gluconate.

It seems to be the best emergency treatment to prevent damage during transit to the hospital.

Kays medical has 40 gram tubes for around $20, don't know if they sell to ordinary citizens, but somebody must.

clearly_not_atara - 3-2-2018 at 15:59

I suggest performing the reaction in a xenon atmosphere with UV light. XeF2 is a solid which is much more easily handled and can be stored indefinitely, giving you some record of the experience. It sounds easier than trying to ampoule F2...

Sulaiman - 3-2-2018 at 18:46

Form the OP,
extreme oxidiser + PE test tubes and/or beakers does not sound like a good combination.

ninhydric1 - 3-2-2018 at 19:18

He probably meant PTFE. woelen is about to do something I wouldn't do in my life (maybe if someone is paying a million bucks, but unlikely).

unionised - 4-2-2018 at 01:49

Quote: Originally posted by woelen  

I have PE-testtubes and beakers. .

Those are fine for HF, but I suspect that, from the point of the perfluuromanganates, they look like "fuel".
I'd go with PTFE or at least FEP (actually I'd just not go there but...)

j_sum1 - 4-2-2018 at 01:55

I have two probably obvious questions.
At mL scale how will you be certail you have achieved your objective?
Once you have made it... Then what? Surely not storage. But reacting with something (non human)... What will you choose?

This sounds fascinating from a distance. Be sure to set up a camera.

unresearched speculation

Sulaiman - 4-2-2018 at 02:19

If there is such a substance as K2MnCl6 then the synthesis and properties may be similar so maybe you could try that first,
substituting Cl for F avoids the HF fear factor ?

A speculative answer to j_sum1 ... probably something like hopefully :D

woelen - 4-2-2018 at 07:43

I tried, but without success, but I will post what I did and what results I obtained.

I prepared two solutions, both at appr. 5 degrees C:
1) Appr. 1 gram of KHF2 in 3 ml of 48% HF, added appr. 0.5 ml of H2O2 to it (exact concentration unknown, but will certainly be well above 20%, made this by freezing out of a little of 12% H2O2 and removing ice). This solution is colorless and stable. No production of oxygen, no decomposition. Freezing out was not that easy, our fridge is set to -20 C but freezing the H2O2 was very slow. It is remarkable how well KHF2 dissolves in the HF.

2) Appr. 0.5 gram of very finely powdered KMnO4 (crunched to a very fine dark grey powder), added appr. 2.5 ml of 48% HF. The solid quickly dissolves, but almost immediately the liquid started foaming, faster and faster, until it foamed over the rim of the test tube. Also, white fumes and a purple gas were repelled from the test tube. After the reaction, a precipitate settled in the test tube, dark brown. Most likely this is just hydrous MnO2. So, concentrated solutions of KMnO4 in concentrated HF decompose. The gas most likely is oxygen, MnO2 is formed, and KF (or KHF2) goes in solution.

So, making solution (2) failed. I tried again, but now with only appr. 100 mg of KMnO4 in appr. 2.5 ml of cold 48% HF. This time, the solution was stable. It did not foam, there was only very slight bubbling, but the solution could be set aside, and its color was very intense purple, just like an aqueous solution of KMnO4.

At this point I dripped solution (1) very slowly into solution (2). With each drop there was effervescence of a colorless gas (oxygen), but after a few drops the solution was colorless and clear :(
All permanganate simply is reduced to manganese(II) by the H2O2. This is a well-known reaction at low pH. I am quite sure that the solution does not contain any MnF6(2-) ions. According to literature ( ) the ion is light brown in solution and the solid is bright yellow. It can be washed with HF+acetone. They only used 40% HF, I used 48% HF, but they used 35% H2O2, mine most likely is less concentrated. I did not have very good stirring, just swirling of a test tube.

I am wondering why I did not see any formation of the MnF6(2-) complex. Maybe my solutions were too warm, or I added too much H2O2. I need to read more about this. The paper mentions that K2MnF6 is not available commercially due to instability at higher temperatures.

aga - 4-2-2018 at 08:12

Bravo !

Brave of you to even try !

fluorescence - 4-2-2018 at 09:42

@woelen: That's an interesting idea. I had my doubts that would work, too. But if you are already trying, I had the idea to make a video section on Fluorine preparations on my Youtube Channel for perhaps 2019 or so and already collected a few ideas in summer. I don't have that list at the moment but if I find it I give you a copy. Turns out there are quite some chemical (not electrochemical) ways to make F2, at least in trace, as proof of principle method. It would at least be worth trying them out for once but I don't have that much time at the moment.

woelen - 4-2-2018 at 09:57

The reason why I picked this synthesis is that it requires only common chemicals besides the HF, available to many home chemists. The more well known synthesis of K2MnF6 (and from that F2) requires the use of SbF5 and anhydrous HF, both definitely not suitable for home chemistry. Working with 40% HF as is done in the paper, mentioned above, is doable in a home setting, albeit that great care is required.

Tsjerk - 4-2-2018 at 11:06

Cool! Woelen is getting more and more a real mad scientist. I have full trust Woelen is one of the guys around here who could pull this off in a successful way.

About the peroxide, I have at least half a liter what was once bought as 30%. I don't know what percentage it is now, but titration could tell. I would be happy to send you some.

Could you tell/pm the city you live Woelen? I'm traveling quite a bit nowadays, if it is on route I could maybe drop it off if you are interested.

An ampoule of XeF2 would be far beyond cool...

Edit, I have a freezer that goes below -30, I could try to see if I can freeze some water out of my peroxide.

[Edited on 4-2-2018 by Tsjerk]

DraconicAcid - 4-2-2018 at 11:17

Any other chemist would be mad to try it. Woelen, on the other hand....I just hope he's very careful.

[Edited on 4-2-2018 by DraconicAcid]

woelen - 4-2-2018 at 11:30

@DraconicAcid: I understand your concern, I will be very careful.

Just to give an impression of what I do:
- I do not pour the HF.
- I put the bottle of HF in a plastic sink.
- I first put the solids and the H2O2 in the PE-test tubes
- Next, I put the PE-test tubes in a special PE-beaker, with a little water at the bottom
- From the bottle, I transfer 48% HF to both of the test tubes, using a long PE-pipette. I do not have to move around the bottle, nor the test tubes for transferring the HF. If I accidentally spill some HF, then it will be contained in the plastic sink.
- Using a PE-bar, I dissolve the solids and mix the H2O2 with the HF.
- Using another PE-pipette, I carefully drip the H2O2/KHF2 solution in 48% HF into the solution of KMnO4.
The entire procedure is done without a single transfer of test tubes, beakers or bottles out of the plastic sink.
When the experiment is done, I simply run a lot of water in the sink. With a PE-stick I carefully tip over the PE-tubes and beaker and allow a lot of water to dilute all of the stuff and then I let it run away under a running tap, using a chain to open the sink. This is not a really environmental issue, all those people brushing their teeth with fluorinated tooth paste produce MUCH more fluoride waste than I do with my 5 ml of 48% HF.

After flushing all the stuff with water I keep running water around the sink and its walls, and over the beaker, test tubes and pipettes, just to be sure that no drops of HF remain unnoticed and not rinsed away.

Finally, inhalation risk is minute. HF at concentration of 48% does not fume when it is cold (appr. 5 C).

[Edited on 4-2-18 by woelen]

clearly_not_atara - 4-2-2018 at 11:35

I found a paper in which a solution containing HF, KHF2 and KMnO4 forms K2MnF6 in the presence of Si, this being intercalated in the fluorosilicate precipitate:

symboom - 4-2-2018 at 22:25

What is the reaction of hydrofluoric acid and potassium permanganate forms permanganic acid and potassium fluoride

Also manganese pentafluoride decomposes with heat therefore it is a solid storage medium for flourine gas

Now we only need a compound that can store ozone and release it when it is heated to decomposition point

[Edited on 5-2-2018 by symboom]

woelen - 5-2-2018 at 00:01

I have read more about the reaction. Apparently the stirring and the VERY careful addition of H2O2 is most important. You have to add JUST enough H2O2, it is very easy to overshoot and add too much and then the compound is destroyed.

This makes doing this experiment on a test tube scale more difficult. I first will try doing this experiment with dilute solutions (in KMnO4 and H2O2, not dilute in HF), adding single drops with H2O2 and each time swirling. I do not want to handle the test tubes with my hands holding them, I will use some clamps. If this mechanism does not work, then it probably is beyond what I can do. Then I need to work on a bigger scale, in beakers, using a stir-bar and a magnetic stirrer. I, however, do not feel comfortable to use 100 ml or so of a 48% HF solution and stirring that with a magnetic stirrer. It also will use up my entire stock of HF.

It is, however, already quite interesting to see what kind of reactions there are in concentrated HF-solutions with common salts. It is weird to see that KMnO4 decomposes in conc. HF and that this gives a faint purple vapor. This most likely is MnO3F.

Tsjerk - 5-2-2018 at 01:34

I would like to advice anyone to keep their eyes open forthese kind of pipettes.

New they are very expensive, but second hand you can probably easily find them for a hobbyist friendly price. They are great for precise pipetting of small volumes. The plastic one-time use tips are available in all kinds of material, usually they are made of PP.

I would advice Gilson pipettes, as they are easier to clean/repair compared to Eppendorf pipettes.

Sulaiman - 5-2-2018 at 02:20

Or economise and get 100 disposable PE pipettes for £1 ?
At £0.01 per piece they really are disposable :D
Plus, you can crimp or melt them shut to act as a short term storage vial.
I do not rely on the graduations/callibration however.

fluorescence - 5-2-2018 at 02:49

It would be interesting if you were able to make the Fluoromanganate, but that shouldn't be the only possible compound to release fluorine. By now we are focusing highly fluorinated metals, that would release the Fluorine, which are easy to make if you already have Fluorine, so it's more or less a fluorine storage. The same can be done with highly fluorinated Cerium for example, there is a whole list of chemicals that would do the same kind of chemistry. And it seems good, because the Manganese one can be made without using Fluorine...if it forms like that...

The question is, are you trying to test, if Fluorine can be made at all in a home lab, or are you heading for bigger amounts? Because for the first one, there are a couple of options to test. I'd also move from complex based reactions and possible electrolysis (I don't have a problem with electrolysis, I thought of that as well, but the problem is you would have to pass the F2 through NaF or KF to get rid of the HF in there and this kind of prevents a good circulation and may lead to contact between the Hydrogen and Fluorine).

The question is, either redox-based reactions or thermal ones. The latter ones are already used. There has been a recent paper, where Copper(II)fluoride was used to fluorinate a compound and they assumed elemental Fluorine would form. It seems that with a small stream of Oxygen, Copper(II)Fluoride would form the much more stable CuO and release Fluorine at high temperatures (500-600°C), which I think is a really interesting thing. Same things were tried with Silversubfluoride for example. And the use of certain Phosphates to make Fluorine is really old and can be found in the old Gmelin books. The question still arises how well they detected it back then but I think it would be worth a try to repeat these experiments and really test them.

And the question is also how well you can influence the redox-series here. I mean you can certainly make Sodium using Magnesium, which does not work according to the redox-potentials, but using high temperatures, and the strong bond between Mg and O you can do it. Or Oxygen does not react with Gold, but once you add a strong complexing agent like Thiourea, some say Iodine or Cyanide, it reduces the redox-potential by a lot. And Peroxide is not that far apart from Fluorine, which means, that if we were able to boost the potential of peroxide, we may be able to get at least traces of Fluorine in a high temperature reaction. Peroxodisulfate is even higher than Peroxide, so let's assume we substitute the Peroxide a bit to, let's say a P-O-O-P-type bond. P being extremely oxophilic, O being in the wrong, high energy oxidation state, perhaps we use a cyclic product, with a lot of strain and some high temperatures close to's just an idea but I really believe that if you increase the 'energy' in the system you may be able to overcome this basic though of redox-potentials and one of the oldest works on the synthesis of Fluorine uses Sodium Metaphosphate in combination with NaF at 650-750°C and air to make Sodium Pyrophosphate and Fluorine.

And according to this Copper paper, I guess many other oxophilic metal fluorides may work as well...

This seems like you would require a lot of exotic compounds but that was just the best combination I could think of at the moment, you can certainly reduce a lot here and end up with similar compounds to these easy to acess Phosphates or metal fluorides (although they have to be really dry or HF will form).

Here is a recent example:
Although the idea dates back to 1894 I think.

Harristotle - 5-2-2018 at 03:39

Tsjerk - The Odin sells fixed increment (not "analogue") pipettes for around $25 USD. They are quite accurate - if properly calibrated equivalent to Gilson/Rainin.

I don't know how any splashes/fumes will affect the piston inside the barrel, but no worse than others.

I dunno, this HF business and possibly unknown high oxidation fluorine compounds worries me. It is an interesting thread to watch though...


symboom - 5-2-2018 at 09:31

I find it interesting if hydrochloric acid and potassium permanganate forms chlorine gas but hydrofluoric acid and potassium permanganate just decomposes

If manganese hepoxide releases ozone then may be the reaction with hydrofluoric acid could release fluorine

It seams an oxidizer needs to be more powerful then fluorine
So if the reaction forms ozone it could form fluorine

Perbromates anyone but those can be formed by fluorine gas so even perbromates as a very strong oxidizer is still too weak only thing perbromic acid and manganese hepoxide would be a scary combination but in thory would that combination be a stron enough oxidizer

fluorescence - 5-2-2018 at 10:47

Well there are also Manganese Oxyfluorides....I made some and I think woelen prepared them as well, where you use the Heptoxide and Fluoride and some acid so basically you would also end up fluorinating you Manganese first.

I don't really know what to expect from Perbromate + Mn(VII), they are both strong oxidizers and they are both highly oxidized, it's not like one of them could be oxidized even further besides perhaps the oxygen in there but that's pretty much it.

chornedsnorkack - 5-2-2018 at 11:42

Quote: Originally posted by symboom  
I find it interesting if hydrochloric acid and potassium permanganate forms chlorine gas but hydrofluoric acid and potassium permanganate just decomposes

If manganese hepoxide releases ozone then may be the reaction with hydrofluoric acid could release fluorine

It seams an oxidizer needs to be more powerful then fluorine
So if the reaction forms ozone it could form fluorine

That´s why it cannot.
Fluorine is simply too strong oxidant to form when there is a way to oxidize oxygen instead.

Reboot - 5-2-2018 at 15:13

If you have any interest in working with small volumes of solvent/reagent, you really owe it to yourself to get an adjustable micropipette. :-) Control and reproducibility are miles ahead of what you could hope to get from an eyedropper or traditional pipette.

They come in capacities from a couple microliters up to 10 ml and have gotten quite cheap:

The basic functions are:

1. Rotate the button to change the volume setting.
2. Pressing the button down/letting back up sucks up (or releases) the set amount.
3. Push down past the point of light resistance and it will eject a little air as well to clear the tip.
4. Another button ejects the (disposable) tip.

There are also variants like the tensettes:

These guys only have ten settings (so the 10 ml version will give you 1-10 ml in even 1.0 ml increments.) The advantage is they're very fast to select the volume on (a single twist of the plunger.)

In the bio lab they go through a ton of the disposable tips, but for non-sterile uses there's no reason you can't re-use them.

Cryolite. - 5-2-2018 at 15:27

Woelen, another preparation of potassium fluoromanganate(IV) can be found in preparation 62 from "Inorganic Laboratory Preparations" by Gert Schlessinger in the forum library ( ). The relevant text is reproduced below:

Eighty-nine grams of 48% hydrofluoric acid are mixed with 11ml of water in a 150ml plastic beaker and 8.9g of anhydrous potassium carbonate are added in portions with stirring. When the effervescence has subsided, the solution of potassium hydrogen fluoride (KHF2), is cooled thoroughly in an ice bath and 8g of powdered potassium permanganate are stirred in. Without delay, absolute ether (free from alcohol and peroxides), is added dropwise, with good stirring, from a medicine dropper or a small burette. To avoid any possible reaction with the non-vitreous material, the ether should be run into the center of the solution and not onto the walls of the plastic vessel. About 2ml of ether are required for the complete discharge of the characteristic purple color of the potassium permanaganate. A brown liquid and yellow crystals of product remain at the end of the reaction; the mixture is left for 20 minutes in the ice bath and then the clear liquid is decanted into a second plastic beaker as completely as possible in order to avoid etching the filter flask subsequently. The brown mother liquor is reserved for the isolation of K3[MnF6]. The crystals of product are transferred to the suction filter with about 25-30 ml of glacial acetic acid, pressed dry, washed twice more with 20ml portions of acid, and finally with a similar quantity of acetone. The complex fluoride is dried in vacuo on filter paper. Yield = 7.5-8.0g. The product must be stored out of contact with moist air in plastic or paraffin coated vials because water causes instant hydrolysis to MnO2, and the fluoride, itself, etches glass with the formation of a brown deposit.

In addition, addition of potassium bifluoride to the brown supernatant liquid precipitates potassium fluoromanganate(III) on cooling.

clearly_not_atara - 5-2-2018 at 23:10

Sounds like the wrong salt though. We want K2MnF6 not K3MnF6.

woelen - 5-2-2018 at 23:43

I understand from Cryolite's post that you get both salts. The yellow precipitate is K2MnF6 and the brown material is K3MnF6. The latter one is obtained as a bonus ;)

I certainly will try this synthesis, it sounds more doable than the one with H2O2. I will do this experiment with KHF2 instead of K2CO3 and adjust the concentration of HF accordingly. This is a nice thing to try next weekend.

Pok - 7-2-2018 at 04:11

Quote: Originally posted by woelen  
The more well known synthesis of K2MnF6 (and from that F2) requires the use of SbF5 and anhydrous HF, both definitely not suitable for home chemistry.

But to make F2 from K2MnF6 you will need SbF5 anyway, as clearly_not_atara pointet out earlier.

In this tread I already mentioned some detailed literature procedures for the H2O2 method. The descriptions there are different from your procedure. They first mixed HF, KHF2 and KMnO4 and added H2O2 in drops afterwards. Also, large volumes are required and the precipitated amount of K2MnF6 will be very small in comparison the the reaction volume.

I'm sceptical about the description of Schlessinger. This procedure is based on experiments from 1899 (here). The yield seems to be much smaller than with the H2O2 method (reference), however the Schlessinger procedure claims a yield of about 62 % which would be good. Also, the ratio of product weight and reaction volume is higher, so this might be worth a try especially when working with small volumes.

[Edited on 7-2-2018 by Pok]

woelen - 7-2-2018 at 06:14

I now remember I did this before, using the procedure in the other thread. First at room temperature, later with chilled solutions. The result was unsatisfactory. All manganese was reduced to mangense(II), at that time I also noticed the decomposition of KMnO4 in HF. I simply have forgotten, the results were not that interesting.

clearly_not_atara - 7-2-2018 at 14:14

Found a ball-milling method. This one is also dependent on intercalation:

Wt% of K2MnF6 in the product tops out at 1.1% but longer ball milling times were not tested in the paper. The product is also a phosphor, which is kind of cool.

DubaiAmateurRocketry - 4-8-2018 at 18:35

Any updates on this?

I am interested in synthesizing the HOF-CH3CN complex by passing F2 into a CH3CN water solution, but of course, I dont have the F2 :)

woelen - 6-8-2018 at 01:40

I have done a little more experimenting in this direction, but without success. I do not have SbF5 (and I think I also do not want that stuff around in my house) and my experiments with KMnO4, H2O2 and HF all have been without result. Sooner or later, the purple color of the permanganate disappears and I get an off-white solid or clear solution. The hydrogen peroxide simply reduces the manganese(VII) to manganese(II) in the HF-solution.

SbF5 (or other covalent fluorides) are really nasty stuff. they are strongly fuming compounds, which react with water, giving dense fumes of HF and the corresponding oxide of the other elements or some intermediate oxyfluoride. This is stuff, which hardly can be handled in a home-setting. It is beyond what I can do safely, and hence I do not wish to have SbF5, PF5 or other easily hydrolyzed fluorine compound.

clearly_not_atara - 6-8-2018 at 14:53

I really don't understand why H2O2 is used in this prep? It's known to catalyze the decomposition of permanganate. So why would you use it for a reaction with permanganate?

In the prep of K2MnCl6, potassium chloromanganate, no peroxide is used, but the hydrochloric acid must be "fuming", that is, concentrated hydrochloric acid saturated with HCl gas:

Presumably it should be possible to convert K2MnCl6 to K2MnF6 by treatment with silver fluoride, which is not likely to have been tested in the literature, but many, many other chlorine compounds are converted to the corresponding fluorides by AgF. This may thus be a workable two-step route, particularly considering that fuming HCl, no matter how concentrated, is not nearly as frightening as anhydrous HF -- and it won't eat glass!

However, dissolving the K2MnCl6 in some solvent so as to react it with AgF is a bit of a puzzle. Water is clearly out of the question. Hydroxyl-bearing solvents in general are not acceptable. DMSO will be oxidized. DMF will probably be oxidized. Propylene carbonate and acetonitrile stand a chance, I think. The tetramethylammonium salt is probably more stable, according to my bad understanding of which cations stabilize reactive anions.

Another possibility is to convert MnF2 to xenon fluoromanganate by reaction with XeF2:

The fluorination of xenon is kind of hard, though...

[Edited on 6-8-2018 by clearly_not_atara]

DubaiAmateurRocketry - 6-8-2018 at 23:02

and if we do obtain XeF2, we can obtain the F2 gas relatively easily, right?

symboom - 6-8-2018 at 23:18

This would only bedoable in a florinated hydrocarbon sich as freon or other florinated hydrocarbons

Silver fluoride soubility
83g/100 g (11.9 °C) in hydrogen fluoride
1.5g/100 mL in methanol(25 °C)

Maybe just find what hydrogen fluoride desolves in a good site with preparations

Everything would have to be made in situ
The HF generated by an acid boric acid and silver fluoride
Generates insouble silver borate

Some sort of modification of the electrofluorination process
To work with a safe form of HF

[Edited on 7-8-2018 by symboom]

Dan Vizine - 7-8-2018 at 15:20

I looked at a number of routes to F2. I settled on the attached as this method relies on easy-to-obtain chemicals, a simple apparatus, and was actually a laboratory-adapted method.

I didn't pursue it as I didn't see a way to produce samples. The amount of preparation needed to end-cap hydroxyl groups and all of the rest that needs to be done to preserve samples was simply excessive.

Albert L. Henne
J. Am. Chem. Soc., 1938, 60 (1), pp 96–97
Publication Date: January 1938

Good luck in your endeavor. The scientists of SM, professional and amateurs, have made just about everything else so why not F2 or BrF3?
Attachment: 1 Page from Henne.pdf (224kB)
This file has been downloaded 209 times

[Edited on 8/8/2018 by Dan Vizine]

symboom - 7-8-2018 at 18:30

Acually your right
There is alot isolated who would have thought
Ceasium. potassium and phosphorous
Here is a periodic table of whats been isolated

The blue is ones that have been isolated
The purple is ones that can be isolated
The orange is radioactive
The Red is dangerous

ChangeColor_7-8-2018-7-0-4.jpg - 88kB

Fluorine could be proven isolated by making a fluorocarbon by electrolysis producing HF and F in situ
[Edited on 8-8-2018 by symboom]

[Edited on 8-8-2018 by symboom]

MrHomeScientist - 8-8-2018 at 06:01

Fluorinated hydrocarbon, you say? I have about 2 gallons of Fluorinert FC-40 that I haven't found a use for. If anyone can show me a solid plan and prove their determination to actually try an experiment using it, I may be willing to part with some. It's very expensive and rare, so I'm very selective with it.

DubaiAmateurRocketry - 27-8-2018 at 15:41

What process do you rip of fluorines off these flourocarbons?

If we use electrodes, would the F2 react with the electrode itself ?

[Edited on 27-8-2018 by DubaiAmateurRocketry]

symboom - 27-8-2018 at 16:41

It is a carbon anode
Here a diagram

images (1).jpeg - 9kB

[Edited on 28-8-2018 by symboom]

metalresearcher - 28-8-2018 at 11:26

Quote: Originally posted by symboom  
It is a carbon anode
Here a diagram

[Edited on 28-8-2018 by symboom]

Watch out when making such a cell ! Keep the H2 and F2 STRICTLY SEPARATED !
These gases do react explosively when they come in contact with each other, they are hypergolic !

Dan Vizine - 28-8-2018 at 13:54

Quote: Originally posted by metalresearcher  
Quote: Originally posted by symboom  
It is a carbon anode
Here a diagram

[Edited on 28-8-2018 by symboom]

Watch out when making such a cell ! Keep the H2 and F2 STRICTLY SEPARATED !
These gases do react explosively when they come in contact with each other, they are hypergolic !

I think that the V-shaped reactor is inherently safer by design.

Tdep - 28-8-2018 at 20:46

Hey, here is a legit easy, HF free way for you to get F2 in a home setting, enough to smell at least. Radiation damage inside Calcium fluoride causes defects and F2 to build up inside the crystal. When the crystal is crushed, the fluorine is released. A lot of fluorite in the environment is already like this, due to naturally occurring uranium.

diddi - 28-8-2018 at 21:21

CaF2 crystals in nature occur with such free F2. it is known as antozonite, or "stink stone" because when crushed the F2 reacts with O2 to form ozone

[Edited on 29-8-2018 by diddi]

DubaiAmateurRocketry - 29-8-2018 at 10:27

although its do-able, sounds like an extremely low yield procedure

phlogiston - 19-10-2018 at 06:31

unless you have access to a very intense source of ionizing radiation.

Toth et al produced fluorine directly by gamma radiation of a mixture of fluorides.
But they used a dose rate of 10E7-10E8 R/h
To put that into perspective: At 10E8 R/h it takes only 0.02 seconds to receive a lethal dose for humans.

But maybe if you mix coursely ground CaF2 with high grade uranium ore and let the mixture sit for a few years it may yield enough to be noticeable/measurable. At least then you can claim having purposefully made F2.

Dan Vizine - 12-12-2018 at 13:01

A bit off-topic but I received an interesting email re. F2....

Hi Dan, FYI, I forgot if I mentioned that I found and purchased a cylinder of F2 gas diluted in Helium for an eximer laser to make fluorine element ampules. Total net F2 gas content is 1.5Lbs or 15 Ft3 @ STP. I'm planning on cryogenically transferring the F2 gas to lecture bottles,minus helium, which hold approx. 2 Ft3 at 2200 psig. so I need minimum of 8 bottles to completely empty the cylinder and I'd prefer to not fill the lecture bottle to max. pressure for safety reasons, maybe half. So I have enough fluorine gas for a lifetime.

Absolutely terrifying and yet so attractive.

Dan Vizine - 13-12-2018 at 14:14

And a further email states that...

I only need a few bottles for my ampuling project so I'll have about a dozen of these [half-filled] lecture bottles left so probably will offer some for sale to qualified buyers.

He expects that next summer is a likely time for this to occur. So, anybody feel qualified?

[Edited on 12/13/2018 by Dan Vizine]

DoctorOfPhilosophy - 13-12-2018 at 14:20

Qualified but not certified

btw, was this fluorine leftovers from SupaVillan's excimer laser?

[Edited on 13-12-2018 by DoctorOfPhilosophy]

Tsjerk - 13-12-2018 at 14:27

Nice, wouldn't want to have it under my responsibility though.

Dan Vizine - 13-12-2018 at 15:17

Quote: Originally posted by DoctorOfPhilosophy  
Qualified but not certified

btw, was this fluorine leftovers from SupaVillan's excimer laser?

[Edited on 13-12-2018 by DoctorOfPhilosophy]

This individual is a rocket technician, doesn't visit here. He's one of those guys handling (at one time) UDMH, RFNA, N2O2, MMH, and various military high explosives. His testicles, I'm told, are Inconel. Former Edwards AFB guy, a real gem. One of the most amazing collectors I've ever met.

Qualified but not certified describes most of us with many of our projects, doesn't it?

[Edited on 12/13/2018 by Dan Vizine]

clearly_not_atara - 13-12-2018 at 16:09

Actually the reaction in that thread from K2NiF6 sounds interesting. Could this salt be produced by e.g. electrolysis of an HF solution of NiF2/KF?