Sciencemadness Discussion Board

How hazardous is fluorine chemistry?

Sulaiman - 14-2-2024 at 02:13

I'm comfortable with iodine, bromine, chlorine and most of their compounds,
and I'm considering doing a little fluorine chemistry,
which I have until now avoided.

So I want to ask members that have actual physical experience with fluorine and its compounds;
How much more hazardous is fluorine than chlorine?

charley1957 - 14-2-2024 at 04:53

Sulaiman, I haven’t personally done any fluorine chemistry, but I have spent the past several years reading up on it, researching, and building a HF distillation unit. I have a stack of research papers, methods, diagrams, and information that I would be happy to scan and send you if you’d like. I also have 98% acid grade fluorspar, a bunch of it, and would be glad to send you some if you need. U2U me for more information. I know that the discussion of fluorine or even HF in this forum has been a sore spot in the past. Personally I don’t see where working with either of these chemicals with lots of research and preparation is any different than working with dozens of other things some of us routinely do which I feel are more dangerous. There are those here who feel like it is an instant death sentence, however. Anyway, there’s the offer.

Admagistr - 14-2-2024 at 05:20

The fluoride anion binds the calcium cation Ca2+ and CaF2 is insoluble and the calcium bound in this way is definitely lost for the organism! So elemental F and most of its compounds, especially the soluble ones, are highly toxic! It is much worse than the other halogens, although bromine is also quite neurotoxic...Fluorine is also the most reactive element hence its additional risk...



[Edited on 14-2-2024 by Admagistr]

Texium - 14-2-2024 at 08:20

Fluorine chemistry is a "sore spot" both on this forum and among chemists in general for good reason. Spill a few mL of conc. HCl on your hand, you rinse it off and go on with your day. Spill a few mL of conc. HF on your hand, and you could die if it's not immediately treated. If it doesn't kill you, it'll almost certainly give you excruciating pain. That's not an alarmist exaggeration, it's a fact. If enough fluoride ions get into your bloodstream, it can shut down the calcium ion channels that keep your heart beating.

Quote: Originally posted by charley1957  
Personally I don’t see where working with either of these chemicals with lots of research and preparation is any different than working with dozens of other things some of us routinely do which I feel are more dangerous.
While this may be true (though I'd be curious to hear what routine things you think are more dangerous), the way I see it is, why risk it when there's plenty of interesting chemistry that can be done without fluorine?

If you still insist on pursuing fluorine chemistry though, don't even think about doing it without having calcium gluconate gel close at hand. Wear good safety goggles, gloves, and a lab coat. Work on small scales to avoid the possibility of a catastrophic accident that could kill you. Don't work alone.

bnull - 14-2-2024 at 09:04

Skip it. You're a "Hobby Chemist, not Professional or even Amateur". Even the professionals avoid it. The bad reputation F has was very well earned. It's a little hard to see someone who had "actual physical experience with fluorine and its compounds". Or someone who is still healthy with all the fingers still attached. Remember what John D. Clark wrote in Ignition! about ClF3:
Quote:
All this sounds fairly academic and innocuous, but when it is translated into the problem of handling the stuff, the results are horrendous. It is, of course, extremely toxic, but that's the least of the problem. It is hypergolic with every known fuel, and so rapidly hypergolic that no ignition delay has ever been measured. It is also hypergolic with such things as cloth, wood, and test engineers, not to mention asbestos, sand, and water--with which it reacts explosively. It can be kept in some of the ordinary structural metals--steel, copper, aluminum, etc.--because of the formation of a thin film of insoluble metal fluoride which protects the bulk of the metal, just as the invisible coat of oxide on aluminum keeps it from burning up in the atmosphere. If, however, this coat is melted or scrubbed off, and has no chance to reform, the operator is confronted with the problem of coping with a metal-fluorine fire. For dealing with this situation, I have always recommended a good pair of running shoes.


If the tales of mishaps didn't scare you enough, read "Toxic properties of inorganic fluorine compounds" by R. Y. Eager.

And if, after all the gruesome details, you still want to mess with fluorine reactions (and have written your will, don't care about amputations or a slow, painful death, don't intend to have kids and such), do it in a non-DIY glove box. Preferably a glove box distant a couple of miles away from you.

Sulaiman - 14-2-2024 at 09:09

Quote: Originally posted by Texium  
...If you still insist on pursuing fluorine chemistry though, don't even think about doing it without having calcium gluconate gel close at hand. Wear good safety goggles, gloves, and a lab coat. Work on small scales to avoid the possibility of a catastrophic accident that could kill you. Don't work alone.
Although I agree with all of the above,
I work alone
because of the people that I know,
none are interested in chemistry,
and certainly none would be competent enough to rescue me.

Texium - 14-2-2024 at 09:11

Then don't do it. If you can't call out to someone nearby to get you to the emergency room as you're flailing around on the floor with muscle spasms from hypocalcemia, then you shouldn't take the risk.

OK

Sulaiman - 14-2-2024 at 11:00

Postponed indefinitely
Thanks

DraconicAcid - 14-2-2024 at 11:29

Good idea.

(I seem to recall that it's advised not to wear a pair of gloves, but you put on one pair of gloves, cover them with calcium--based gel, and then put on another pair of gloves on top of the first.)

clearly_not_atara - 14-2-2024 at 11:35

It's a pretty broad question, isn't it? As Texium emphasized, HF is extremely dangerous. Many fluorides attack glass, which also creates a barrier to working with them.

But not all sources of fluoride are as dangerous as HF. Hydrogen fluoride, pKa 3.2, crosses the skin as a neutral molecule. However, ionized fluoride does not have this property. Fluoride salts in neutral or alkaline solution are much less dangerous.

Unfortunately, HF is the entry point to most of the interesting fluorine chemistry! So there's not a whole lot that can be done safely. There are a few interesting applications of CsF and a lot of things that you can do with AgF, although caution is advised since the organic halogen exchange with AgF can be very exothermic. The requirement for large quantities of silver limits what you can do, though.

I have considered the preparation of certain 2-fluoroethyl compounds with AgF. Otherwise I think of fluoride as something to be avoided.

woelen - 14-2-2024 at 11:46

I personally have done a few experiments with 48% HF, but I did not feel comfortable at all while working with this.
I know about the extreme risk and the toxic effects. It is much more dangerous than e.g. cyanides or even arsenic. Its effects are extreme. Have a look at internet, when you search for hydrofluoric acid wounds. If you see that, then you understand why I am so reluctant to experiment with this acid.

Working with fluoride salts is less dangerous (e.g. NaF, KHF2, NH4HF2), but still, be very careful with them and do not allow the liquids to touch your skin!
Some fluorine-containing salts can be handled in a somewhat more relaxed way. An example is HBF4 and salts of that. The BF4(-) ion is sufficiently stable to stay around in aqueous solution with only little hydrolysis to HF and boric acid. Salts of this ion are even more stable (e.g. NaBF4, Co(BF4)2, Cu(BF4)2, NH4BF4). The BF4(-) ion is quite inert, and a good counter-ion if you want to experiment with metal complexes. But I prefer ClO4(-) for that purpose, I still feel not entirely comfotable with the fluorine in the BF4(-) ion. It can hydrolyze to HF, especially on long standing in lower concentrations at low pH.

One of the experiments I did with flouride is the following: https://woelen.homescience.net/science/chem/exps/KMnO4+NaF+H...
Another experiment with fluoride: https://woelen.homescience.net/science/chem/riddles/titanium...

As you can see, only small quantities were involved and the starting point is NaF. The fluoride and HF is very dilute, but sufficiently concentrated to cause severe issues if your skin is exposed to this!

Sulaiman - 14-2-2024 at 13:14

Side note... just checked online shopping here,
1M (2%) HF 1litre USD10.45 equivalent...including shipping.

All I'd have to do is distill it :o

Fleaker - 14-2-2024 at 15:45

At work, I once had to deal with anhydrous HF and TaF5 together. It was not a fun experience. Suffice it to say, the combination is extremely aggressive to polymeric materials and chews them up.

HF here in the USA, by the 1000 liter tote--with no valve allowed--about $1.40/lb at 48-49 wt %.

Twospoons - 14-2-2024 at 16:29

Pickling paste for stainless steel welding contains up to 5% HF. I wonder how many welders know just how risky that stuff is? How many would have calcium gluconate gel on hand?

Bedlasky - 14-2-2024 at 17:08

There are some experiments with fluorides which can be done relatively safely. Lanthanides make beautiful insoluble fluorides. You can try making some stable fluoro complexes like [BF4]-, [SiF4]2-, [AlF6]3- etc. Hexafluoroaluminates have pretty low solubility, even Na or K salts. You can decolorize Fe(III)-SCN- complex with fluoride. You can use fluorides to get dissolve some very inert metals like Ti, Zr, Hf, Nb, Ta (use fumehood or work in well ventileted are, because you may need to heat these solutions if needed). Or you can try conductometric titration of fluoride with CaCl2. Just work in small scale and be careful. You don't need HF for these experiments at all, just some soluble fluoride (NaF, KF, NH4F etc.)

BromicAcid - 14-2-2024 at 20:52

Fluorine chemistry can be dangerous. The old standby to have on hand is calcium gluconate. But I posted in a different thread before that became common it was just magnesium sulfate paste in ethanol (you'd have to find the thread to be sure) or something that was pretty accessible.

I've ended up working with quite a bit of fluorine chemistry over the years. There are precautions to take for sure and at one point I even entertained building a fluorine cell. But there are other things we deal with that are quite dangerous as well. I would hardly say the danger is in a league of its own but it does pose some unique hazards you need to consider whenever you experiment.

chornedsnorkack - 15-2-2024 at 10:10

Quote: Originally posted by woelen  
I personally have done a few experiments with 48% HF, but I did not feel comfortable at all while working with this.
I know about the extreme risk and the toxic effects. It is much more dangerous than e.g. cyanides or even arsenic.

What is so bad about arsenic?
HF and HCN seem to have some of the similar hazards - both are low boiling (HF at 19,5 Celsius, HCN at 25,5 Celsius), weak acids, can pass through intact skin. Both are weak acids so you can suppress volatilization (and skin absorption?) by keeping solutions neutral/basic. HF is stronger acid than HCN, so easier to keep in neutral or mildly basic solutions.
In contrast, much of the As compounds is not so volatile. OK, arsine and halides are, but it seems that HAsO2 does not tend to volatilize as much?

Quote: Originally posted by woelen  

One of the experiments I did with flouride is the following: https://woelen.homescience.net/science/chem/exps/KMnO4+NaF+H...
Another experiment with fluoride: https://woelen.homescience.net/science/chem/riddles/titanium...

As you can see, only small quantities were involved and the starting point is NaF. The fluoride and HF is very dilute, but sufficiently concentrated to cause severe issues if your skin is exposed to this!

Does this HF also etch your glassware? You don´t seem to mention it.
Also, are there any nice colour reactions for low concentrations of HF?

[Edited on 15-2-2024 by chornedsnorkack]

[Edited on 15-2-2024 by chornedsnorkack]

unionised - 16-2-2024 at 07:17

Quote: Originally posted by clearly_not_atara  



I have considered the preparation of certain 2-fluoroethyl compounds with AgF.

There's a plant which makes 2 fluoroethyl compounds.
https://en.wikipedia.org/wiki/Dichapetalum_cymosum

"commonly known as gifblaar from Afrikaans, "
"gifblaar " means "poison leaf.

So this stuff- which is toxic enough for a plant to be named after it- is one of the less dangerous fluorine compounds...

Sulaiman - 17-2-2024 at 03:32

Quote: Originally posted by unionised  
Quote: Originally posted by clearly_not_atara  



I have considered the preparation of certain 2-fluoroethyl compounds with AgF.

There's a plant which makes 2 fluoroethyl compounds.
https://en.wikipedia.org/wiki/Dichapetalum_cymosum

"commonly known as gifblaar from Afrikaans, "
"gifblaar " means "poison leaf.

So this stuff- which is toxic enough for a plant to be named after it- is one of the less dangerous fluorine compounds...
Thanks, good to know,
Where can I get some?

EF2000 - 17-2-2024 at 07:38

Quote: Originally posted by unionised  
Quote: Originally posted by clearly_not_atara  



I have considered the preparation of certain 2-fluoroethyl compounds with AgF.

There's a plant which makes 2 fluoroethyl compounds.
https://en.wikipedia.org/wiki/Dichapetalum_cymosum

"commonly known as gifblaar from Afrikaans, "
"gifblaar " means "poison leaf.

So this stuff- which is toxic enough for a plant to be named after it- is one of the less dangerous fluorine compounds...

Gifblaar contains sodium fluoroacetate, the infamous "compound 1080".
There's also a lot of fluoroacetate-containing plants in Australia, mainly in genus Gastrolobium, also called "poison pea" or "poison brush", or just "poison".

Though, if you get to Australia, much easier is to extract 1080 from poison baits used for control of invasive species.

clearly_not_atara - 17-2-2024 at 16:32

Quote: Originally posted by unionised  
Quote: Originally posted by clearly_not_atara  



I have considered the preparation of certain 2-fluoroethyl compounds with AgF.

There's a plant which makes 2 fluoroethyl compounds.
https://en.wikipedia.org/wiki/Dichapetalum_cymosum

"commonly known as gifblaar from Afrikaans, "
"gifblaar " means "poison leaf.

So this stuff- which is toxic enough for a plant to be named after it- is one of the less dangerous fluorine compounds...

Fluoroacetate is special, you know that. It specifically inactivates aconitase. Much more toxic than 3-fluoropropionate for example, and also more toxic than difluoroacetate or trifluoroacetate. It's also technically not a fluoroethyl compound if we're being pedantic.

unionised - 18-2-2024 at 03:40

Quote: Originally posted by clearly_not_atara  
It's also technically not a fluoroethyl compound if we're being pedantic.

No, but it's the metabolic endpoint of some fluoroethyl compounds...

Keras - 18-2-2024 at 04:15

I have a 1L bottle of 60% HF on my shelves. Never opened it. I have a tube of calcium gluconate gel at hand, but up to know, I didn't dare opening the bottle.

chornedsnorkack - 19-2-2024 at 13:34

Some instructive text from professionals:
https://ehs.wisc.edu/wp-content/uploads/sites/1408/2020/08/S...

Precipitates - 19-2-2024 at 20:08

Quote: Originally posted by Sulaiman  
Side note... just checked online shopping here,
1M (2%) HF 1litre USD10.45 equivalent...including shipping.

All I'd have to do is distill it :o


70% HF is available here for around 6 USD per litre, which is kind of terrifying


dicyanin - 20-2-2024 at 07:52

Quote: Originally posted by Sulaiman  
Quote: Originally posted by unionised  
Quote: Originally posted by clearly_not_atara  



I have considered the preparation of certain 2-fluoroethyl compounds with AgF.

There's a plant which makes 2 fluoroethyl compounds.
https://en.wikipedia.org/wiki/Dichapetalum_cymosum

"commonly known as gifblaar from Afrikaans, "
"gifblaar " means "poison leaf.

So this stuff- which is toxic enough for a plant to be named after it- is one of the less dangerous fluorine compounds...
Thanks, good to know,
Where can I get some?


Africa Seeds carries it, but it's really expensive for what is essentially a native weed, $35 for 8 seeds. Monofluoroacetate is toxic to humans, but other mammals are much more sensitive to it. Attached paper discusses its properties and the preparation of some derivatives.
They also prepare monofluoroacetic acid by refluxing methyl monoiodoacetate with silver fluoride at 170°C in a platinum apparatus:D, followed by saponification of the ester. Apparently when a platinum refluxing apparatus is not available, the authors note a lead one can be used.

gifblaar.png - 509kB

Attachment: MONOFLUOROACETIC ACID the toxic principle of Gifblaar (Dichapetalum cymosum) _marais1944_20_1.pdf (664kB)
This file has been downloaded 96 times

sceptic - 20-2-2024 at 11:52

After reading this thread, I'm in no hurry to do so, but if I wanted to extract fluorine compounds from gifblaar, does anyone know of a reasonably safe way to do it? I ask because here (southern Africa) I have an almost unlimited supply of that plant available.

Texium - 20-2-2024 at 13:18

I found the previous paper published by the author of the one above. I wasn't able to access the full text, but fortunately the relevant portion was available:
Quote:
Extract 10 kg. dried and finely ground plant material with 96% EtOH for 36-48 hrs., distil the alc. under reduced pressure, dissolve the sirupy mass in 4 l. water, add 500 cc. 10% H2SO4 and let stand for 1-2 days to allow precipitated material to settle. Filter the solution by suction, extract the filtrate repeatedly with Et2O, neutralize the combined extracts to phenolphthalein with N KOH and sep. the aqueous solution which contains the K cymonate. Evaporate the solution to about 300 cc. on the water bath, liberate the free acid by adding the equivalent amount of dilute H2SO4, and repeatedly extract the solution with Et2O until practically all the acid is removed. Decolorize the Et2O extract with activated charcoal, dry overnight with anhydrous Na2SO4, filter and distil the Et2O at low temperature Transfer the concentrate to a small distillation flask, collect the distillate in fractions of 10° up to 160°, carefully neutralize (N KOH and phenolphthalein) the individual fractions, as well as the initial Et2O distillate, and evaporate the K salt solutions separately to dryness on a water bath. Wash the residues with acetone, dry at 100° and crystallize from 96% EtOH. The best yields of the K salt are obtained from the 110-160°, fractions. After repeated recrystallization from 96% alc. the salt melts at 213° with decomposition; insoluble in anhydrous organic solvents except MeOH.
J. S. C. Marais, Onderstepoort Journal of Veterinary Research, Volume 18, Pages 203-206, 1943

Apparently before they figured out it was monofluoroacetate, they called the mystery anion "cymonate."

teodor - 21-2-2024 at 02:11

This is molten NH4F inside my DIY home fume hood.
And I am still alive. So, the answer: it depends ...

When I did planning to start playing with F I've selected boron trifluoride and BF4 anion as an entry level but still very interesting F compound.

qqq.jpg - 200kB

ItalianChemist - 21-2-2024 at 03:58

BF4 anion is very interesting in coordination chemistry since it forms many compounds that can be easily crystallized from water.
Also I don't think it is very dangerous unless it is heated. (It will release some amounts of HF I suppose)

teodor - 21-2-2024 at 06:01

NH4BF4 is quite volatile. I kept it in an open container for drying and it covered inside walls of my cabinet.
I plan to use it for generate BF3 which have many interesting applications in inorganic as well as in organic chemistry. Also it's very potent drying agent.

[Edited on 21-2-2024 by teodor]

EF2000 - 21-2-2024 at 06:44

Quote: Originally posted by Texium  
I found the previous paper published by the author of the one above. I wasn't able to access the full text, but fortunately the relevant portion was available:
J. S. C. Marais, Onderstepoort Journal of Veterinary Research, Volume 18, Pages 203-206, 1943

Apparently before they figured out it was monofluoroacetate, they called the mystery anion "cymonate."

Here's the paper, found it in open access on the University of Pretoria website: https://repository.up.ac.za/handle/2263/59331

Attachment: 19marais1943_18.pdf (279kB)
This file has been downloaded 87 times

mr_bovinejony - 25-2-2024 at 06:12

Quote: Originally posted by teodor  
NH4BF4 is quite volatile. I kept it in an open container for drying and it covered inside walls of my cabinet.
I plan to use it for generate BF3 which have many interesting applications in inorganic as well as in organic chemistry. Also it's very potent drying agent.

[Edited on 21-2-2024 by teodor]


Sodium tetrafluoroborate can be used as well for this, does NH4BF4 work better? I'd like to make trifluoroacetic acid but I have found no other way than to use antimony fluoride compounds.

EF2000 - 26-2-2024 at 07:50

Quote: Originally posted by dicyanin  


Africa Seeds carries it, but it's really expensive for what is essentially a native weed, $35 for 8 seeds.

Gastrolobium calycinum is ~5 AU dollars for packet from Australian Seed: https://www.australianseed.com/shop/item/gastrolobium-calyci...
Gastrolobium bilobum is 6.5 USD for 10 seeds from rarepalmseeds.com, even though it's not rare and not a palm. And $14 for 1,000 seeds, meaning that for the price of 8 giftblaar seeds one can get ~2,500 poison brush seeds and plant an entire garden with it. Very toxic garden.

I also found a little instruction how to pre-treat seeds and sow them.

Attachment: Gastrolobium species.pdf (339kB)
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Rainwater - 27-2-2024 at 08:20

Side note, he is some dude spraying cold fluorine gas onto a fire brick (time index 11:12)
https://youtu.be/KX-0Xw6kkrc?t=672

Dr.Bob - 26-3-2024 at 11:14

I did work in a fluorochemical synthesis lab for a few years with my old friend Ed, who was much more jaded about working with fluorine than most people are today, having worked with it for 50+ years. But only long after working there did I realize how many things he had there were incredibly toxic, like cylinders of F2, HF, SF4, and many others, plus some other things like monofluoroacetic acid and related compounds. But he was somehow lucky, and I think the F2 just made him inert to all other chemicals, as he lived to be about 98.

But even Ed had Calcium Gluconate kits and other safety gear, plus some very large scrubbers which most of the gas streams went out through. Mostly I worked with Freons, which are quite safe to handle, but the process of fluorinating hydrocarbons is quite exciting, and involves some very specialized equipment.

I have also used anhydrous liquid HF for cleaving and deprotecting peptides, which is mostly not done any more now (mostly used for tBoc type peptide chemisty, which is rarer now). That was almost as crazy, but done in special hood with special equipment, and 2 thick gloves, special aprons, face mask, etc. That was just as nuts, and done in big pharma lab. I would not do it for fun or much else now. But I have done other fluorine type chemistry that is safer and many fluoro groups are fine, like many, but not all, -CF3 type reagents. So the safety is quite dependent on the exact type of chemistry, the equipment, and the real need to do it. Like making some fluorophosphates, minor changes can make drastic differences in toxicity.

So I would say that some fluorine chemistry is quite safe, some is dangerous, and some is nuts. Even Ed tried to use ClF3 a few times and decided that it was not able to be used safely for anything, depsite being a very powerful fluorinating agent.

[Edited on 26-3-2024 by Dr.Bob]

BromicAcid - 26-3-2024 at 14:23

Had a very nasty experience with SF4 at one point Dr.Bob, always interesting to see people mention it since it tends to fly under the radar of most chemistry.