Sciencemadness Discussion Board - Fluorine gas non-electrical production

This guy Publios (from the video) is a nutcase: to dry this CuF2 on the hob when he doesn't even know whether he's reacted away all the HF or not, is asking for sudden death or at least permanently reduced lung capacity for life.

HF doesn't just cause 'burns on the skin': inhaled it is extremely lethal...

And anyone here contemplating making F2 in a home setting requires sectioning to a very secure madhouse... :mad:

i totally agree with you excess copper hydroxide is absolutely necessary as it is negligible solubility in water to make sure it is fully reacted then adding acetone to help dry it as copper fluoride is insoluble in it.

Jor - 6-5-2011 at 18:01

Sorry but I simply do not agree. I think that HF is quite overrated in it's danger. Yes it is a severe skin contact hazard in high concentrations because of the very high amount of fluoride you can absord from small amounts of liquid (very high molarity, fluoride precitipates calcium in the blood, causing heart failure). The ability to precitipate calcium is the main reason for it's lethality (well yes, I guess the fluoride ion also has some direct toxicity on lung tissue, but just look at the LC50 of HF, it's not that lethal, comparable to bromine and chlorine). To give you an idea: 1mL of conc. HF contains about 20 mmol of HF. 20mmol of HF is (following ideal gas law, in realilty it's probably less) about 500mL of PURE HF gas or 500 L of 1000 ppm. So while liquid, concentrated HF may be very dangerous, gaseous HF is much less so (example: to absorb the same amount of of fluoride wich is present in 10mL of conc. HF you need to inhale 5000 L of 1000ppm HF). Just look at the toxicity data. Is about as toxic as chlorine in gas form. Would you call someone completely mad if he heats chlorine water or bromine? Probably not. That's what I think is a bit weird: everyone put's gaseous HF as EXTREMELY dangerous. Liquid HF (or conc. solutions) is extremely dangerous. And even then it's used massively in analytical labs and in the pharmaceutical industry and there are not that many HF deaths reported.
Another comparison: Gaseous HCN is very hazourdous, but can in no way be compared with liquid HCN. Can you imagine how dangerous that is? Less than a mL through the skin (readily absorbed, although I can imagine a lot evaporated before it can enter the skin) can be fatal, while you still have to inhale quite some gas to die (300 ppm for 10 minutes).

Same for F2. While it may be a very severe reactive hazard (and HF also is hard to handle, it attacks glass) it's not like a nerve gas! Just because it is very obscure and not common and has the reputation of being so reactive and rare doesn't mean it is increadibly toxic. Again it is indeed very toxic, but there are much more toxic gasses like phosgene. So I don't see why you are so mad if you generate mg amounts of fluorine from some CoF3 or similar....
The same with many superacids, while they may sound extremely corrosive and dangerous, many of these acids are not that dangerous at all.

Or can you provide some reference material with lethal (or severe injury) cases or animal studies with very low airbourne concentrations of HF or F2? Then I could be proven wrong ofcourse.

AndersHoveland - 6-5-2011 at 18:48

It is not too dangerous to handle solutions of hydrofluoric acid in water, if the proper saftey precautions are taken (splash-proof goggles, water-proof apron, long rubber gloves, a preprepared solution of calcium acetate in water to apply to the skin in the case of accidental skin-contact). However,be aware that large ammounts of HF solution splashed onto a large enough area of skin can result in death.

Boiling a solution of HF, or gaseous HF stored in a pressurized tank, is far more dangerous. While HF is a very poisonous gas, it is not exceptionally so. There are many other gases or volatile liquids of comparable toxicity which are not uncommon in a professional laboratory. If a vile of concentrated HF solution shatters on the ground, you are certainly not going to die, which is possible with some other extremely deadly reagents. In summary, while HF may not be comparable to many other extremely deadly chemicals that may be found in a laboratory, it should still be treated as if it were more dangerous than it actually is. Handling HF requires caution and is not something to take lightly.

Fluorine gas actually has about the same toxicity as hydrogen fluoride gas, since the main mechanism of toxicity for F2 is the immediate formation of HF since it is so reactive.

[Edited on 7-5-2011 by AndersHoveland]

plante1999 - 7-5-2011 at 02:18

from wiki:

Owing to its extreme reactivity, elemental fluorine was not isolated until many years after the characterization of fluorite. Progress in isolating elemental fluorine was slowed because it could only be prepared electrolytically and even then under stringent conditions, since the gas attacks all but certain exotic materials. In 1886, the isolation of elemental fluorine was reported by French chemist Henri Moissan after almost 74 years of effort by other chemists.[54] The generation of elemental fluorine from hydrofluoric acid proved to be exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These individuals came to be referred to as "fluorine martyrs".

blogfast25 - 7-5-2011 at 06:45

Jor and Anders:

Ok. Keep relativating the dangers of solutions of HF and gaseous HF, probably from the perspective of someone who’s never had to work with the damn things!

Now everything is relative of course but with HF not that much. There’s a story even here on SM of a very experienced experimenter who got serious HF burns from a small amount of inadvertently produced HF (an unexpected hydrolysis), in a professional setting. I think I can find it back if you’re interested.

Wiki linked to a story (no longer available unfortunately) of an American refuse collector who dies from exposure to the gas. How? Some idiot had put an EMPTY plastic bottle of HF solution with the common refuse, and when loaded into the truck, unbeknownst to our poor victim the bottle was crushed and he got spritzed with the ‘content’ of the bottle. He died hours later in hospital, in an agonisingly painful death. Because that’s the other thing: HF burns are very painful, don’t reveal themselves immediately and cause almost irreversible scar tissue.

Sure it’s possible to work safely with HF solutions, soluble fluorides and even anhydrous HF but it takes very strong precautions and clear thinking. Someone I know who works with it on a daily basis professionally refuses to do so w/o full breathing apparatus.

Comparing the toxicity of HF with Cl2 or Br2 is also nonsense. I’ve inadvertently choked on Cl2 several times, having to leave my lab waiting for the concentration to fall. With HF I wouldn’t have lived to tell the tale. Guess what the result would have been with F2…

Bot0nist - 7-5-2011 at 07:09

Wiki linked to a story (no longer available unfortunately) of an American refuse collector who dies from exposure to the gas. How? Some idiot had put an EMPTY plastic bottle of HF solution with the common refuse, and when loaded into the truck, unbeknownst to our poor victim the bottle was crushed and he got spritzed with the ‘content’ of the bottle. He died hours later in hospital, in an agonisingly painful death. Because that’s the other thing: HF burns are very painful, don’t reveal themselves immediately and cause almost irreversible scar tissue.

I immediately thought of that story when I read the last few posts in this thread. I too thought I had read it on wiki, but could no longer find it. Google revealed that the man was Michael Hanley, age 51.

<a href="http://articles.nydailynews.com/1996-11-13/news/18014822_1_hydrofluoric-acid-jug-sanitation-workers">Here's a link to the story.</a>

[Edited on 7-5-2011 by Bot0nist]

BromicAcid - 7-5-2011 at 10:52

And anyone here contemplating making F2 in a home setting requires sectioning to a very secure madhouse... :mad:

I've checked, my local madhouse doesn't have adequate ventilation for the procedure.

At one time I very nearly put fourth the time and effort to make a fluorine cell:

http://www.sciencemadness.org/talk/viewthread.php?tid=2358

The procedure and construction was laid out in the Inorganic Synthesis series and claimed to be a safe / reliable procedure if standard precautions were taken against working with the beast. My sticking point at the time wasn't the safety of it (though in retrospect, although I had most everything worked out, it was still quite caviler) but obtaining the necessary quantity of potassium bifluoride and/or anhydrous hydrogen fluoride to keep the cell running.

Was it safe for me to try this? No. Was it sane? Yes, but only because I thought I knew what I was getting into. Would I have survived? I think I would have, but only because of my 10 years of experience as a backyard chemist and my 3 years of college chemistry to that point. Still, my skill with gas handing at the time was abysmal and this is the Tyrannosaurus of the periodic table. Nevertheless, I've always wanted to make perbromic acid, and aside from Xenon Difluoride this was the only way.

[Edited on 5/7/2011 by BromicAcid]

entropy51 - 7-5-2011 at 11:07

Quote: Originally posted by entropy51

In high school we etched glass as Magpie described, in the hood. There were two wax bottles of (aqueous) HF in the stockroom as well. They were slowly emptied by those of us who needed it for home use. IIRC we carried it home in little plastic bottles, most likely in our pocket.

Not that I would do that now, but most likely we didn't wear gloves. No one died. I've since used aqueous HF to make HBF4. It's exciting but no big deal if done carefully with very good ventilation.

Fluorine is another story altogether, but it's unlikely that anyone with enough skill to actually make it couldn't handle it safety.

blogfast25 - 7-5-2011 at 12:37

@BromicAcid:

You also wanted to make anhydrous HF. And then that thread went dead…

Jor - 7-5-2011 at 14:25

Jor and Anders:

Ok. Keep relativating the dangers of solutions of HF and gaseous HF, probably from the perspective of someone who’s never had to work with the damn things!

Now everything is relative of course but with HF not that much. There’s a story even here on SM of a very experienced experimenter who got serious HF burns from a small amount of inadvertently produced HF (an unexpected hydrolysis), in a professional setting. I think I can find it back if you’re interested.

Wiki linked to a story (no longer available unfortunately) of an American refuse collector who dies from exposure to the gas. How? Some idiot had put an EMPTY plastic bottle of HF solution with the common refuse, and when loaded into the truck, unbeknownst to our poor victim the bottle was crushed and he got spritzed with the ‘content’ of the bottle. He died hours later in hospital, in an agonisingly painful death. Because that’s the other thing: HF burns are very painful, don’t reveal themselves immediately and cause almost irreversible scar tissue.

Sure it’s possible to work safely with HF solutions, soluble fluorides and even anhydrous HF but it takes very strong precautions and clear thinking. Someone I know who works with it on a daily basis professionally refuses to do so w/o full breathing apparatus.

Comparing the toxicity of HF with Cl2 or Br2 is also nonsense. I’ve inadvertently choked on Cl2 several times, having to leave my lab waiting for the concentration to fall. With HF I wouldn’t have lived to tell the tale. Guess what the result would have been with F2…

I am not denying the dangers of liquid solutions of HF... I just told you that these are highly lethal even in small amounts because even small amounts of liquid have very high amounts of fluoride and are easily absorbed. I am telling you that gaseous HF is simply not in the same category as liquid HF. HF solutions are so dangerous because it simply is a poison wich is easily absorbed through the skin, most systemic poisons wich are easily absorbed through the skin can be highly dangerous (phenol, dimethyl sulfate, etc.), because in these ways lethal amounts are much easier to absorb quickly than by inhaling it.
Like I told you 1mL of 20 M HF is equivalent to 500mL of pure HF gas! Almost all deaths by HF are by skin contact with concentrated solutions.
It is much harder to absorb lethal amounts of HF by inhalation (unless you use poor ventilation). By inhalation, HF gas is not more hazardous than say bromine.

''With HF I wouldn’t have lived to tell the tale. Guess what the result would have been with F2…''
This is nonsense. How do you know?
F2:

In the rats, the LC50 for exposures of 5, 15, 30, and
60 minutes were 1,088, 606, 420, and 287 mg/m3, respectively

HF:

The 15-minute LC50 values for hydrogen fluoride were
3,362 mg fluoride/m3 for guinea pigs and 2,090 mg fluoride/m3 for Wistar-derived
rats (Rosenholtz. et al. 1963). The 60-minute LC50 for hydrogen
fluoride in ICR-derived mice was 266 mg fluoride/m3 (Wohlslagel et al. 1976),
while in rats it has been reported as 1,084 mg fluoride/m3 for a Sprague-
Dawley-derived strain (Wohlslagel et al. 1976) and 1,016 mg fluoride/m3 for a
Wistar-derived strain (Rosenholtz et al. 1963).

Other investigatorss report even higher concentrations.
http://www.fluoridealert.org/ATSDR-Fluoride.pdf

This means that Br2 is even more toxic than HF by inhalation. Just because skin contact with HF is so damn scary doesn't mean you can call someone a complete madman because he plays with gaseous HF!

I see people distilling 10's of mLs of Br2 and I never see people call them insane...

[Edited on 7-5-2011 by Jor]

blogfast25 - 8-5-2011 at 04:36

Gaseous HF is so poisonous because undissociated HF molecules are readily soluble in fatty tissue. That's precisely why also hydrofluoric acid is a contact poison: HF is a weak acid and in solution it's mostly undissociated HF molecules which are readily absorbed by living tissue and penetrate skin quickly.

And yes, if HF(g) levels in a room would have reached the sort of levels of chlorine that make you choke and leave the room, lung tissue damage would almost certainly kill you.

Comparing LC50 values of different toxic substances is often like comparing apples and oranges.

Bromine has one enormous advantage over HF: it’s very easy to condense, thus easy to contain. Ordinary distillation kits, operated by skilled experimenters, make distilling ml quantities of bromine not at all difficult and not particularly dangerous.

You certainly don’t have to be a madman to manipulate the various forms of soluble fluorides and it all depends on concentrations and quantities. But it remains one of the most dangerous substances you may need to use…

unionised - 8-5-2011 at 04:40

No chemical is safe and every chemical is safe. The difference between "safe" and "unsafe" is what you do with it.
Without any knowledge of that you can't say whether someone's actions are foolish or not.

Jor - 8-5-2011 at 05:52

''Gaseous HF is so poisonous because undissociated HF molecules are readily soluble in fatty tissue.''

This says nothing on the toxicity of the gas. Yes this property makes the liquid so dangerous.

''And yes, if HF(g) levels in a room would have reached the sort of levels of chlorine that make you choke and leave the room, lung tissue damage would almost certainly kill you. ''

No. How do you conclude this? I can also say without any reference if I inhale 1 mg of HCN I die, while it's simply not true. Very bad argumentation. The cause of death by HF inhalation is pulmonary oedema just like other irritant gasses. The simple facts are that (not considering long term chronic exposure the fluoride, only acute) even gases such as Cl2 are more fatal (in ppm levels in the air) than is HF. It's just because liquid HF or it's solutions have gotten such a bad reputation (for a good reason) you just assume gaseous HF is instant death. LC50 is a good indication of toxicity when comparing substances on lethality especially if they have the same mode of action (pulmonary oedema). The above quote is based on nothing (no references, choking concentrations of halogens are about 20ppm, if HF is at this level you ssimply wont die or even get injured unless you're in it for ages), just on your scare of liquid HF (wich IS extremely lethal).
Yes Br2 and Cl2 are a little easier to contain unlike HF wich attacks glass but the former 2 are quite more toxic too. But yes Cl2 and Br2 are so 'normal' as they are used to often to make them pretty harmless. I dare to bet that if 100 people here work with gaseous HF just like they work with Cl2 (assuming they have resistant apparatus) not much worse will happen than when handling thing like Cl2.

When you look at HF it does actually have pretty high lethal doses ,more than 2 grams, wich is actually quite high for toxic substances. The reason it is so dangerous that it is very easy to quickly absorb this dose via the skin.

Still ofcourse this material is very dangerous and utmost care must be taken when handling it, it's just that I hate it when people make simple assumptions like: OMG HF gas, you must be mad. F2 is instant death, blabla. All these are based on nothing (have you even ever read toxicology reports?) just the image you created yourself for the substance.
The same is true for HCN. Many people make it sound so increadibly lethal because of it's reputation and it's active use as a killing agent while there are many many toxic gasses wich are quite commonly encountered wich are in the same category or are even more toxic.

''You certainly don’t have to be a madman to manipulate the various forms of soluble fluorides and it all depends on concentrations and quantities. But it remains one of the most dangerous substances you may need to use…''
I agree, it is a very dangerous substance, as a gas or liquid.

Sorry for the offtopic discussion on the toxicology of these substances

[Edited on 8-5-2011 by Jor]

The WiZard is In - 8-5-2011 at 06:12

Jor and Anders:
Wiki linked to a story (no longer available unfortunately) of an American refuse collector who dies from exposure to the gas. How? Some idiot had put an EMPTY plastic bottle of HF solution with the common refuse, and when loaded into the truck, unbeknownst to our poor victim the bottle was crushed and he got spritzed with the ‘content’ of the bottle. He died hours later in hospital, in an agonisingly painful death. Because that’s the other thing: HF burns are very painful, don’t reveal themselves immediately and cause almost irreversible scar tissue.

The Analogue Guy again. From my HD.

Trash Collector Dies After Inhaling Discarded Acid
By LAWRENCE VAN GELDER
New York Times 13xi96
[edited /djh/]

A New York City sanitation worker died yesterday after he inhaled the
fumes of a corrosive acid from a discarded container that burst under the
compacting blades of a garbage truck making routine collections in
Brooklyn.

Fire Department hazardous materials experts identified the substance
as hydrofIuoric acid, which is often used to etch glass. The acid
immediately burns the skin like fire and dehydrates it.

A second sanitation worker was injured in the incident, which brought
the Mayor and Sanitation Commissioner rushing to the bum unit of New
York Hospital-Cornell Medical Center, touched off an investigation by the
police and sanitation departments and raised the possibility of homicide
charges against those responsible for leaving the acid to be picked up.

The Brooklyn District Attorney's office said: "The investigation is
continuing. Depending on what evidence is obtained in that investigation,
there could be homicide charges if there is an arrest.

Lucien Chalfen, a spokesman for the Department of Sanitation, said the
incident was avoidable. "The appropriate procedure in this case would
have been to notify the Department of Environmental Protection, and they
will then recommend the proper procedure for disposal," he said.

Normally, the Sanitation Department picks up only residential refuse and
waste from public garbage cans on the street, while private haulers
dispose of industrial garbage. Mr. Chalfen said there was no way of
knowing whether some industrial user had illegally placed the acid where
it could be picked up by sanitation workers.

The dead man was identified as Michael Hanly, 49, of Brooklyn, a
Sanitation Department worker for 22 years. His injured partner, who
suffered bums on the face hands when he came to Mr. Hanly's aid, was
identified as Thomas Giammarino of Staten Island, a member of the
department for years who was listed in stable condition the bum unit,
where he was held for further observation.

"People are unaware of what they dispose of and how they dispose of
things," he said "Frequently, sanitation workers are stuck by needles.
People many times dispose liquid thinners and paints and so forth. When
the cans break open, they spatter. this case, this is a very highly caustic
acid” Mr. Chalfen said yesterday's incident began around 9 A.M. when Mr.
Hanly ,Ad Mr. Giammarino were working their regular collection route.
Their truck was nearly loaded when one of .them picked up a parcel and
threw it in the back.

"One of the sanitation workers Was doing what they call cycling the
blades, to compact, and the other turned away to pick up other material,
and he heard like a pop, and this acid, which is hydrofluoric acid, just
sprayed out," he said.

Gerard Parkin, a professor of chemistry at Columbia University, said
hydrofluoric acid is usually stored in a plastic container. "I presume that's
what the pop was," he said. [Anyone remember when it was stored in
wax containers?]

He said the acid not only burns but ,also dehydrates the skin, adding
that even as little as a drop "generates a Very, very painful experience."

Mr. Hanly inhaled the acid's fumes, leading to his death in the
emergency room of New York Hospital, Mr. Chalfen said. Mr. Giammarno,
who was 10 or 15 feet away collecting other refuse when the acid
container burst, rushed to his aid and pulled his partner away from the
truck.

Later, the truck was impounded. Mr. Chalfen said a gallon container
believed to have held the acid was recovered.

After the two workers were injured, they were taken to the Manhattan
hospital in a Fire Department ambulance, while a department hazardous
materials team set about identifying the volatile material.

Eventually, the Fire Department ordered and took part in the
detoxification of the emergency room, hospital personnel, its own
personnel and ambulance and members of a Bensonhurst ambulance
crew that initially responded.

For a time, the hospital's emergency entrance was closed off for
decontamination. No injuries were reported.

Mr. Chalfen, the Sanitation Department spokesman, said, "Basically this
kind of thoughtless act has created a tremendous amount of pain both for
the families involved as well as for each member of the department."

HF Burns from Industrial Chemical News

The WiZard is In - 8-5-2011 at 06:15

INDUSTRIAL CHEMICAL NEWS
July 1983

by David E. Beesinger, M.D., Director, Hermann Hospital Burn
Unit Assistant Professor of Surgery, University of Texas
Medical School, Houston, and Harold R. Mancusi-Ungaro, Jr.
M.D,, Assistant Professor of Surgery, University of Texas
Medical School.

As the need for industrial solvents ,an cleaning agents grows,
the incidence of chemical injuries escalates. It is estimated that
each year more than 60,000 patients in the United States
require professional care for injuries resulting from exposure to
chemicals. Hospitalization, skin grafting, and prolonged reha-
bilitation time may be needed to deal with the depth and extent
of these injuries. Loss of life or limbs and severe cosmetic
defects may even result.

Hydrofluoric acid is a particularly treacherous example of
such chemicals, causing extensive injury when skin contact
occurs. This acid is frequently used industrially for the etching
of glass and silicon chips, cleaning of machinery, and octane
production. Here's a case history for a typical injury and course
of treatment for exposure to hydrofluoric acid.

Case in point
A 59-year-old male was admitted to Hermann Hospital Burn
Unit, Houston, Texas, after having spilled approximately "one
cup" of hydrofluoric acid onto his left thigh. The worker was an
employee of a petrochemical plant in south Texas where the
acid was used for octane production.

The wound was irrigated at the plant with a solution of sodium
bicarbonate. The patient was then taken to a local hospital.

At that facility, the wound was further irrigated with water and
Zephiran solution, and dressed with an ointment containing
magnesium sulfate and magnesium hydroxide. Calcium
gluconate was injected into the fatty tissue beneath the wound.
The patient was then transferred to the Hermann Hospital Burn
Unit.

On arrival at the burn center, the wound of the left thigh was
initially estimated to be of second-degree depth. More calcium
gluconate was injected beneath the wound, copious irrigation
was performed, and the wound was covered with Silvadene
(silver sulfadliazine), an antimicrobial cream. to minimize
bacterial contamination. Because of the extreme pain caused
by the injury, narcotics were administered, and the patient was
observed.

Approximately 24 hours after injury, the wound was again
inspected and found to be full-thickness or third-degree in
depth. The patient was taken to the operating room where
upon incision, the injury was seen to extend not only through
all layers of the skin, but also deep into the fat beneath the
skin. Dead tissue was surgically removed all the way down to
the muscle of the thigh. Skin was then taken from the
opposite thigh and placed on the burn wound as a skin graft.

The patient's hospital course was then uneventful. He was
ready for discharge two weeks after his initial injury. The area
of skin grafting had matured to the point that the patient was
able to return to work two months after the injury.

Injury class
Chemical injuries are not burns per se. They do not destroy
tissue by heat contact, but instead cause injury by other
methods of tissue protein destruction. According to their
actions, chemicals capable of skin injury are classified as 1)
oxidizing agents, 2) corrosives, 3) salt formers or cation
binders, 4) desiccants or drying agents and 5) vesicans or
blister producers.

Hydrofluoric acid belongs to category three, acting to bind
calcium in whatever tissues are brought in contact with the
chemical. Since calcium is imperative for cellular life, its binding
or precipitation brings about cell death in a short period of time.
If the wound covers a large percentage of body surface area,
excessive amounts 4 calcium may be inactivated so that
inadequate amounts of the cation are available for normal body
organ functions. Specifically, heart function is diminished and
rhythm of the heart beat becomes abnormal.

Clincally, the action of hydrofluoric acid on skin is manifest in
several significant ways. The destruction of tissue is very rapid,
dictating that the wound be irrigated as soon as possible with
large volumes of water or a dilute solution of sodium
bicarbonate.

In addition, the appearance of the depth of the wound may be
deceiving, as demonstrated in `ihe case just presented.
Because of progressive tissue destruction, wounds that un!ially
appear to be superficial usually are much deeper after 24 to 48
hours. This progressive death of cells again dictates early
removal of the acid by irrigation and provides rationale for the
recommended sub cutaneous injection of calcium gluconate.
The substance binds the fluoride ion as calcium fluoride,
preventing tissue binding of calcium, and thus decreasing
cellular destruction and greatly decreasing pain.

Since hydrogen fluoride injuries are generally full-thickness or
third degree burns, the wound frequently requires surgical
therapy. If damage extends to muscle, tendons, and bone,
amputations of upper or lower extremities may be required. In
other cases, skin and subcutaneous tissue must be removed
and the resultant area covered with skin grafts.

If you are exposed to hydrofluoric acid in your work, you
should be aware of its capacity for injury, and take all
precautions to prevent accidents. You should also be well
versed in appropriate first aid, and be prepared to seek early
professional treatment.

blogfast25 - 8-5-2011 at 07:58

Jor, we’ll just have to aree to disagree. Your argumentation convinces me as much as mine does you. Pot meet Kettle.

Hydrofluoric Acid Burn Treatement

The WiZard is In - 8-5-2011 at 08:35

CRC Handbook of Laboratory Safety
2nd ed. 1971

NB This was originally published 1966 use it
for clinical judgement at your own risk.

djh
----
I shelve :—

Banks, Sharp & Tatlow Editors
Fluorine : The First Hundred Years
Celebration volume to commemorate the
centenary of the isolation of fluorine by
Henri Moissan on 26th June, 1886
Elsevier Sequoia 1996

HF Burns - From MESH

The WiZard is In - 8-5-2011 at 13:32

Taking time out of my medical studies — Gynecology through the
picture study method. too run HF Burns through MESH.

Hits

281 MEDLINE/PubMed - journal citations, abstracts
0 NLM Catalog - books, AVs, serials
1 Bookshelf - full text biomedical books
152 TOXLINE Subset - toxicology citations
0 DART - Developmental and Reproductive Toxicology
0 Meeting Abstracts

Here do be a few of the more current refs. I didn't include
many for which there was no abstract.

Finger burns caused by concentrated hydrofluoric acid, treated with intra-arterial calcium gluconate infusion: case report.
Capitani EM, Hirano ES, Zuim Ide S, Bertanha L, Vieira RJ, Madureira PR, Bucaretchi F.
Sao Paulo Med J. 2009 Nov;127(6):379-81.
Poison Control Center, School of Medicine, University Hospital, Universidade Estadual de Campinas, Campinas, São Paulo, Brazil. capitani@fcm.unicamp.br

CONTEXT: Hydrofluoric acid (HF) is widely used in industry and at home. Severe lesions can occur after contact with highly concentrated solutions, leading to tissue necrosis and bone destruction. Specific treatment is based on neutralization of fluoride ions with calcium or magnesium solutions.

CASE REPORT: A 41-year-old male was seen at the emergency department 35 minutes after skin contact with 70% HF, showing whitened swollen lesions on the middle and fourth fingers of his right hand with severe pain starting immediately after contact. 2.5% calcium gluconate ointment was applied. Twenty-four hours later, the patient was still in severe pain and the lesions had worsened. Considering the high concentration of the solution, early start of severe pain, lesion characteristics and impossibility of administering calcium gluconate subcutaneously because of the lesion location, the radial artery was catheterized and 2% calcium gluconate was administered via infusion pump for 36 hours, until the pain subsided. No adverse effects were seen during the procedure. Ten days later, the lesions were stable, without bone abnormalities on X-rays. Six months later, a complete recovery was seen.

CONCLUSIONS: Intra-arterial calcium gluconate might be considered for finger burns caused by concentrated HF. Complete recovery of wounded fingers can be achieved with this technique even if started 24 hours after the exposure. However, controlled clinical trials are needed to confirm the effectiveness and safety of this intervention.

Survival after hypocalcemia, hypomagnesemia, hypokalemia and cardiac arrest following mild hydrofluoric acid burn.
Wu ML, Deng JF, Fan JS.
Clin Toxicol (Phila). 2010 Nov;48(9):953-5.
Division of Toxicology, Department of Medicine, Taipei Veterans General Hospital, Taipei, Taiwan. mlwu@vghtpe.gov.tw

BACKGROUND: Although hydrofluoric (HF) acid burns may cause extensive tissue damage, severe systemic toxicity is not common after mild dermal exposure.

CASE: A 36-year-old worker suffered a first-degree burn of 3% of his total body surface area as a result of being splashed on the right thigh with 20% HF acid. Immediate irrigation and topical use of calcium gluconate gel prevented local injury. However, the patient developed hypocalcemia and hypomagnesemia, hypokalemia, bradycardia, and eventually had asystole at 16 h post-exposure, which were unusual findings. He was successfully resuscitated by administration of calcium, magnesium, and potassium.

CONCLUSION: This report highlights a late risk of HF acid dermal exposure.

[The patient later died when his ass-hole snapped shut upon receiving the hospital bill.]

Hydrofluoric acid burns: rational treatment.
Burd A.
J Burn Care Res. 2009 Sep-Oct;30(5):908.

Hydrofluoric acid burns: a 15-year experience.
Stuke LE, Arnoldo BD, Hunt JL, Purdue GF.
J Burn Care Res. 2008 Nov-Dec;29(6):893-6.
Comment in:
• J Burn Care Res. 2009 Sep-Oct;30(5):908. PMID: 19692923.
Department of Surgery, University of Texas Southwestern Medical Center, Parkland Memorial Hospital Burn Center, Dallas, Texas 75390-9158, USA.

Hydrofluoric acid (HF) is a strong inorganic acid commonly used in many domestic and industrial settings. It is one of the most common chemical burns encountered in a burn center and frequently engenders controversy in its management. We report our 15 year experience with management of HF burns. We reviewed our experience from 1990 to 2005 for patients admitted with HF burns. Primary treatment was with calcium gluconate gel. Arterial infusion of calcium and fingernail removal were reserved for unrelenting symptoms. There were 7944 acute burn admissions to our center during this study period, 204 of which were chemical burns. HF burns comprised 17% of these chemical burn admissions (35 patients). All were men, with a mean burn size of 2.1 +/- 1.5% (range, 1-6%) and hospital stay of 1.6 +/- 0.7 days (range, 0-3 days). The most common seasonal time of injury was in the summer. Twelve patients (34%) were admitted to the intensive care unit for a total of 14 intensive care unit days, primarily for arterial infusions. Ventilator support was not required in any patient. No electrolyte abnormalities occurred. All burns were either partial thickness or small full thickness with no operative intervention required and no deaths. The upper extremity was most commonly involved (29 patients, 83%). The most common cause was air conditioner cleaner (8 patients, 23%). HF is a common cause of chemical burns. Although hospital admission is usually required for vigorous treatment and pain control, burn size is usually small and does not cause electrolyte abnormalities, significant morbidity, or death.

Hazardous brick cleaning.
Ferng M, Gupta R, Bryant SM.
J Emerg Med. 2009 Oct;37(3):305-7. Epub 2008 Jun 20.
Department of Emergency Medicine, Rush University Medical Center, Chicago, Illinois, USA

Hypocalcemia, hypomagnesemia, and hypokalemia following hydrofluoric acid chemical injury.
Dalamaga M, Karmaniolas K, Nikolaidou A, Papadavid E.
J Burn Care Res. 2008 May-Jun;29(3):541-3.
Department of Internal Medicine, NIMTS General Hospital, Athens, Greece.

Dermal exposure to hydrofluoric acid could potentially result in severe serum calcium and magnesium depletion induced by binding with fluoride anion. This report describes the case of a 48-year-old man who developed hypocalcemia and hypomagnesemia accompanied by hypokalemia-an interesting finding-following a chemical injury with exposure to 70% hydrofluoric acid. Successful treatment included administration of calcium gluconate and magnesium both intravenously and topically.

Hydrofluoric acid burn to penis.
Schmidt MJ, Bryant SM.
Clin Toxicol (Phila). 2007 Sep;45(6):732.
Department of Emergency Medicine, Northwestern University Feinberg School of Medicine, Chicago, Illinois, USA.

Opical treatment of experimental hydrofluoric acid skin burns by 2.5% calcium gluconate.
Roblin I, Urban M, Flicoteau D, Martin C, Pradeau D.
J Burn Care Res. 2006 Nov-Dec;27(6):889-94.
Laboratoire de Développement Analytique, Agence Générale des Equipements et Produits de Santé, (AP-HP), Faculté des Sciences Pharmaceutiques et Biologiques, Université Paris V, Paris, France.

Topical therapy with 2.5% calcium gluconate gel is considered as the "first-aid" treatment of accidental hydrofluoric acid skin burns. The efficacy of three different gel formulations varying in the amount and/or nature of their gelling and moisturizing agents was experimentally evaluated. Thirty male Wistar-Han rats (250 g) were exposed to 60 mul of 40% hydrofluoric acid for 2 minutes on two spots (4 cm) of skin under pentobarbital anesthesia. One lesion was massaged with 1 g of gel (10 rats/type of gel) at 3 minutes; 30 minutes; 1 hour; 1 hour, 30 minutes; 2 hours; 3 hours; and 4 hours after injury. During the next 3 days, rats received a single daily application of gel. The other lesion for each rat remained untreated (control). From day 1 after injury to the end of the study (day 17), gel therapy reduced the number of extensive (-66%), severe (-44%), and moderate (-34%) lesions (P < .0001). It reduced (P < .001) the median Area Under the Curve day 0-17 of burn injury from 34.0 (25th to 75th percentile: 18.2-44.5; untreated lesions) to 17.7 (7.0-26.7); overall, there was three cases of treatment failure. At day 17, full wound recovery was obtained in 14 cases by gel therapy compared with 6 in the absence of treatment. The efficacy of the three gel formulations was comparable for all evaluated parameters. Repeated applications of a 2.5% calcium gluconate gel is an efficient treatment of experimental 40% hydrofluoric acid skin burn; few differences were observed between evaluated gel formulations.

7 cases of hydrofluoric acid burn in which calcium gluconate was effective for relief of severe pain.
Ohata U, Hara H, Suzuki H.
Contact Dermatitis. 2005 Mar;52(3):133-7.
Department of Dermatology, Nihon University School of Medicine, 30-1 Oyaguchi-kamimachi, Itabashi-ku, Tokyo 173-0032, Japan. utaken@nifty.com

We report 7 cases of chemical burns due to hydrofluoric acid (HF). The patients suffered from severe pain. However, the pain was relieved after treatment with calcium gluconate. 6 out of the 7 cases were men. At the accidental exposures, all the patients had been engaged in washing or cleaning work and received burns on their hands and/or fingers. In one case, the forearm was also involved. During such work, all the patients had used rubber gloves, but the gloves had pinholes. For the treatments, 4% calcium gluconate jelly was applied in 5 cases and 4 of 7 were subcutaneously injected with 8.5% calcium gluconate. The involved nails were removed in 5 cases. It is concluded that physicians should provide calcium gluconate jelly and subcutaneous injections to treat an HF burn and should not hesitate to remove the involved nails. To prevent chemical burn due to HF, education and reeducation of workers regarding the hazard of this chemical are necessary.

Lethal inhalation exposure during maintenance operation of a hydrogen fluoride liquefying tank.
Dote T, Kono K, Usuda K, Shimizu H, Kawasaki T, Dote E.
Toxicol Ind Health. 2003 Jul;19(2-6):51-4.
Department of Hygiene and Public Health, Osaka Medical College, Takatsuki City, Osaka, Japan. rhart99@comcast.net

Calcium sulfate adheres to the inside of liquefying pipes during the production of liquefied hydrogen fluoride. It is regularly washed away with water jets every six months. Two days before the operation, the pipes were experimentally washed down with water and the safety of the operation was confirmed with acidic washing fluid (pH 5). A 65-year-old man was severely sprayed on his face just after the start of the operation. He died half an hour later from acute respiratory failure. High serum concentrations of ionized fluoride indicated massive exposure to hydrofluoric acid (HFA). Pathological findings revealed severe bilateral pulmonary congestion and edema. It was hypothesized that calcium sulfate hardened with the water during the experimental washing and caused some blockages in the pipes. Consequently, choking of the pipes caused the HFA to collect and the washing fluid ran back. Weak HFA is not pungent to skin and mucous membranes. Therefore, it was suggested that a low concentration of HFA was inhaled directly into the peripheral respiratory tracts. No risk management against HFA exposure was in place during the operation. It is especially important to take thorough safety measures against inhalation of HFA. It is also essential that there are no stoppages of the pipes before the operation.

Treatment of hydrofluoric acid burn to the face by carotid artery infusion of calcium gluconate.
Nguyen LT, Mohr WJ 3rd, Ahrenholz DH, Solem LD.
J Burn Care Rehabil. 2004 Sep-Oct;25(5):421-4.
Comment in:
• J Burn Care Rehabil. 2005 May-Jun;26(3):291. PMID: 15879755.
Regions Hospital, St. Paul, Minnesota 55101, USA.

Hydrofluoric acid (HF) is highly corrosive substance often used in industrial processes. HF burns to the skin cause local tissue injury. Systemic hypocalcemia may ensue, with the potential to produce life-threatening arrhythmias. Medical treatment consists of local application of topical calcium gels, subcutaneous injection of calcium gluconate, and intravenous or intra-arterial infusion of calcium gluconate. Calcium gluconate infusions have been used for HF burns on distal extremities and digits. We report a case of HF burn to the face that was treated by the use of calcium gluconate infusion via the external carotid artery.

[Hydrofluoric acid burns. A rare chemical emergency situation].
Richter H, Hollenberg S, Sachs HJ, Oeltjenbruns J, Weimann J.
Anaesthesist. 2005 Feb;54(2):123-6.
[Article in German]
Klinik für Anaesthesiologie und operative Intensivmedizin, Charité-Universitätsmedizin Berlin, Campus Benjamin Franklin, Berlin. hrichter@zedat.fu-berlin.de

Burns caused by hydrofluoric acid can be life-threatening. Of special significance is the often underestimated local and sometimes delayed deep action of the highly diffusible free fluoride ions and the accompanying systemic toxicity. The specific antidote calcium gluconate can be topically applied, injected into tissue or infused intra-arterially. Because of the extreme danger of systemic toxicity even after seemingly trivial injuries, monitoring in the intensive care station, especially by measuring the calcium concentration in blood and electrocardiography, and therapy is recommended.

. Hydrofluoric acid-induced burns and life-threatening systemic poisoning--favorable outcome after hemodialysis.
Björnhagen V, Höjer J, Karlson-Stiber C, Seldén AI, Sundbom M.
J Toxicol Clin Toxicol. 2003;41(6):855-60.
Department of Reconstructive Plastic Surgery, Karolinska Hospital, Stockholm, Sweden

BACKGROUND: Skin contact with hydrofluoric acid (HF) may cause serious burns and life-threatening systemic poisoning. The use of hemodialysis in fluoride intoxication after severe dermal exposure to HF has been recommended but not reported.

CASE REPORT: A 46-year-old previously healthy man had 7% of his body surface exposed to 71% HE Despite prompt management, with subsequent normalization of the serum electrolytes, recurrent ventricular fibrillation occurred. On clinical suspicion of fluoride-induced cardiotoxicity, acute hemodialysis was performed. The circulatory status stabilized and the patient fully recovered. High fluoride levels in the urine and serum were confirmed by the laboratory.

DISCUSSION: There is no ultimate proof that the favorable outcome in this case was significantly attributable to the dialysis. However, most reported exposures of this magnitude have resulted in fatal poisoning. As our patient had normal serum electrolytes and no hypoxia or acidosis at the time of his arrhythmias, it was decided that all efforts should be focused on removing fluoride from his blood. The rationale for performing hemodialysis for this purpose is clear, even though such intervention is more obviously indicated in patients with renal failure.

CONCLUSION: Hemodialysis may be an effective and potentially lifesaving additional treatment for severe exposure to HF when standard management has proven insufficient.

Treatment of hydrofluoric acid burns.
Schiettecatte D, Mullie G, Depoorter M.
Acta Chir Belg. 2003 Aug;103(4):375-8.
Department of Plastic Surgery, AZ Sint-Jan Brugge, Belgium.

Hydrofluoric acid injuries have a potential for both systemic as well as severe local tissue destruction. In this article the different treatment modalities will be presented. Hydrofluoric acid is frequently found in the semiconductor industry, in rust removers and façade cleansers. The negligence or carelessness of workers and ignorance of the risks of hydrofluoric acid promote the incidence of these severe burns. To prevent these burns, adequate information for the workers is necessary. Splash goggles and neoprene gloves as well as laboratory coats should be worn at all times to prevent eye and skin contact. In cases of exposure, therapy should be accurate and immediate.

----
I disposed of my bottle of HF acid some years back (by dilution)
the plastic bottle was looking Ifffff-yyy! I had visions of me picking
it and having it splinter.

Say - anyone remember the days before plastic — when HF acid
came in wax bottles? You cut a small piece of the tube projecting
out of the top off with a hot knife, and then heated it to seal it off
again.

I remember chloroacetic acid came in a glass bottle with a
tube projecting from the top, you cut it open, later flamed
it shut again.

CA Acid is some corrosive ----.

fluorine, mangansese trifluoride

AndersHoveland - 9-5-2011 at 13:59

So to summarize, one should take every precaution possible to prevent any HF solution from coming in contact with skin, which results in severe burns deep into the tissue. If some of the solution nevertheless gets onto your skin, IMMEDIATELY throw water onto the area as fast as you can, even if you do not feel any pain. A delay of only a few seconds can make a big difference in the severity of the burn.
(preferably you should have on big plastic container of calcium acetate solution (without a lid, and within easy reach next to the expiriment) to dump on the area. I would advise doing the expiriment outside, to avoid fumes. Sometimes the fan in the fumehood can be too noisy and distracting. You will be safer if you are using all your concentration to avoid an accident. If doing the expiriment outside, be sure not to abandon the HF solution at any time. Someone else could find your expiriment and not realize the danger.

MnF3 decomposes to manganese(II) fluoride above 600°C.
In situ time-resolved X-ray diffraction study of manganese trifluoride thermal decomposition , J.V. Raua, V. Rossi Albertinib, N.S. Chilingarova, S. Colonnab, U. Anselmi Tamburini, Journal of Fluorine Chemistry 4506 (2001) 1–4

(2)KMnO4 + (2)KF + (10)HF + (3)H2O2 --> (2)K2MnF6 + (8)H2O + (3)O2
(solution is 50% aqueous HF, the potassium fluoromanganate precipitates out)
http://www.thieme-chemistry.com/fileadmin/Thieme/HW-100/pdf/...

MnF3 itself hydrolyzes with water.

Potassium hexafluoronickelate likely coud be prepared in a similar reaction. At 400°C, a mixture of solid K2NiF6 and KF disproportionates/decomposes to form K3NiF6 and F2. (the reaction is reversible at 250°C).

also see wikipedia: http://en.wikipedia.org/wiki/Manganese(IV)_fluoride

[Edited on 9-5-2011 by AndersHoveland]

Bubafat - 20-10-2011 at 09:28

To preface: I am a fluorine chemist with ~8 years of handling HF, F2, Metal fluorides, XeF2, etc etc etc etc.

There is so much inaccurate information in this thread that I'm not even going to start to address, instead I'll focus on the posters question.

CuF2 --950C--> Cu + F2

While on paper this may be possible, feasibly it is not. You will not be able to find a reactor that is stable to fluorine at these temperatures. Graphite will react. Teflon will degrade. Even corrosive resistant metals and alloys such nickle 201, inconel, monel, etc will react and flake away.

HF/F2 is dangerous...but it can be handled safely. Oddly enough, in my corresponences with docters who treat HF burns, the group of people who most frequently are admitted and treated for HF burns are people who have fancy expensive rims on their cars. The cleaners for these aluminum rims has HF on it and they don't read the warning label and don't wear gloves.

Bubafat - 20-10-2011 at 09:32

BTW...if you're going to be working with HF or any metal fluorides which can hydrolyze to HF (i.e. with the moisture on your skin), order some calcium gluconate gel (Brand name Calgonate). Use butyl rubber gloves and garments (NOT latex/nitrile!). Keep a spray bottle of dilute ammonium hydroxide around and spray everything you think might have been contaminated. Basic fluoride is MUCH less dangerous than acidic fluoride.

watson.fawkes - 20-10-2011 at 10:34

Quote: Originally posted by Bubafat

You will not be able to find a reactor that is stable to fluorine at these temperatures. Graphite will react. Teflon will degrade. Even corrosive resistant metals and alloys such nickle 201, inconel, monel, etc will react and flake away.

I take it in these conditions that fluorine will displace oxygen in ceramics and create fluxing compounds. Correct?

AndersHoveland - 20-10-2011 at 12:01

Precautions for Fluorine Gas
Fluorine ignites on contact and easily burns through concrete. Stainless steel also burns easily in fluorine gas, but generally does not spontaneously react without ignition. Stainless steel valves have been used to handle fluorine gas, despite the danger. The inside of the valve must be free from any contaminants that could cause ignition of the steel. The valve is then gradually exposed to progressively increased concentrations of fluorine mixtures to oxidize any trace contaminants. Despite the precautions, for reasons not entirely understood, the valves have often spontaneously ignited when, after several minutes of operation without incident, the flow rate was suddenly increased.

More Ideas
The below are ideas for extreme reactions, which have not yet been investigated.

H2O2 + 6HF + FeO4(-2) --> FeF6(-2) + 4H2O + O2
The H2O2 and HF must be premixed, then added to ferrate solution. A similar reaction is known to occurr for permanganate.

After completion of reaction, a solution of calcium nitrate is added, which would cause CaFeF6 to precipitate out. The solution is filtered, and the precipitate is dried.

The fluoroferrate ion FeF6(-2) is probably only stable in alkaline solution. Addition of acid likely would cause hydrolysis, with the loss of oxygen:
4 FeF6(-2) + 2 H2O + 8 H+(aq) --> 4 FeF3 + 12 HF + O2

Calcium fluoroferrate could possibly be used as a powerful dehydrating agent. While acting as a base towards nitric acid, it would simultaneously dehydrate a small portion of the acid to dinitrogen pentoxide.
(2)CaFeF6 + (24)HNO3 --> (4)Ca[NO3]2 + (4)Fe[NO3]3 + (24)HF + (2)N2O5 + O2
Ideally the concentration of the nitric acid used in the reaction should be greater than 98%.

MnF3 decomposes to manganese(II) fluoride above 600°C.
(2)KMnO4 + (2)KF + (10)HF + (3)H2O2 --> (2)K2MnF6 + (8)H2O + (3)O2
(solution is 50% aqueous HF, the potassium fluoromanganate precipitates out)

MnF3 itself hydrolyzes with water.

Potassium hexafluoronickelate likely coud be prepared in a similar reaction. At 400°C, a mixture of solid K2NiF6 and KF disproportionates/decomposes to form K3NiF6 and F2. (the reaction is reversible at 250°C). One would think this would also work with K2FeF6.

blogfast25 - 20-10-2011 at 12:09

Quote: Originally posted by watson.fawkes

I take it in these conditions that fluorine will displace oxygen in ceramics and create fluxing compounds. Correct?

Not sure about that at RT but certainly fluorination of D and F block elements by treating their (highly stable) oxides with F2 is an industrial practice. I would think ceramic would be attacked somewhat slowly.

Unlike this chicken that got fried with F2 by some German scientists:

http://www.youtube.com/watch?v=M5_9z1TxUfg

'Vorsprung durch nonsense', something like that...

Neil - 23-10-2011 at 07:55

oh, you beat me to it.

http://www.youtube.com/watch?v=wqLnSkLalOE

fluorine + brick

blogfast25 - 23-10-2011 at 08:18

Quote: Originally posted by Neil

fluorine + brick

Yeah, it's quite remarkable and at RT too! Makes me think that ceramics won't stand much chance either, although they're compacter and that might help a little...

Panache - 23-10-2011 at 17:51

Quote: Originally posted by Bubafat

To preface: I am a fluorine chemist with ~8 years of handling HF, F2, Metal fluorides, XeF2, etc etc etc etc.

So does it have an odor, f2 that is? Taste? is it soluble in anything at cryotemps? Is it reactive with liquid nitrogen at those temperatures. Can i burn off a wart with it? Can i turn a disused copper water line into a fluorine manifold at home?

Some of the later questions are silly, oh and also the taste one, but the others are of interest. Do F2 researchers buy the fluorine in or make it in the lab. If its bought in what delivery protocol does it have? Is it a cylinder handcuffed to a guy with a machine gun or does USPS delivery it. These are all fascinating questions i would love to know the answer to if you can be bothered.

blogfast25 - 24-10-2011 at 05:07

Quote: Originally posted by Panache

Do F2 researchers buy the fluorine in or make it in the lab. If its bought in what delivery protocol does it have?

It's bottled, of course. Sold only to 'authorised personnel' I would imagine, using specialist couriers for transport.

Wiki states that 'about 17,000 tonnes of fluorine are produced per year by 11 companies in G7 countries'. That's quite a bit of F2... I used to live relatively near one ('F2 Chemicals' in Preston).

Bubafat - 24-10-2011 at 21:05

Quote:

I take it in these conditions that fluorine will displace oxygen in ceramics and create fluxing compounds. Correct?

Yes.

Quote:

Stainless steel valves have been used to handle fluorine gas, despite the danger.

Same goes with teflon and/or FEP tubing. Even at room temperature, if you flow fluorine through them too fast they WILL catch on fire (and scare the crap out of you in the process). Thankfully the "fire" can be extinguished by simply turning off or slowing down the fluorine.

Quote:

So does it have an odor, f2 that is?

As you might imagine, it smells very similar to a combination of chlorine and ozone. More chlorine than ozone.

Quote:

Taste?

Dunno...not dumb enough to try that one...since it would immediately start my tongue on fire.

Quote:

Is it soluble in anything at cryotemps?

It's a liquid with a significant vapor pressure at -196C. Meaning that while you can condense it, if you pull a vacuum on it, it will bubble up. Soluble in anything at that temp...dunno, never tried condensing too much of it.

Quote:

Is it reactive with liquid nitrogen at those temperatures?

No. Fluorine requires either high temperatures, a catalyst or a spark to react with nitrogen (producing NF3, N2F4, etc).

Quote:

Can i burn off a wart with it?

Yes. My graduate adviser told me stories of how his boss used to detect pinhole leaks in his F2 line by licking his finger and then passing it over the vacuum line...if it got hot = leak.

Quote:

Can i turn a disused copper water line into a fluorine manifold at home?

I wouldn't. Stainless steel would be better, but without proper training you'd be putting yourself and others at significant risk. Fluorine is nothing to laugh at.

Quote:

Do F2 researchers buy the fluorine in or make it in the lab?

Most buy it, however due to some pesky laws, researchers in Japan have to make it (via electrolysis) and consume it immediately.

Quote:

If its bought in what delivery protocol does it have?

Can't say. :cool:

strontiumred - 25-10-2011 at 04:27

Thanks for all the warnings and advice guys. I've finally managed to get a water soluble tantalum compound (K2TaF7) and will treat this with great care.

I made a 2% solution of this and boiled it for just a minute - Test tube turned white!

2K2TaF7 + 3H2O ---> K2Ta2O3F6 + 2KF + 6HF

I was careful but still didn't expect such a corrosive mix.

vertexrocketry - 5-2-2025 at 01:57

chemical force made fluorine witha test tube and a fume hood so making it is not very hard

teodor - 7-2-2025 at 09:52

By some reason, at least according to wikipedia, electrolisys of anhydrous HF at low temperatures is still a main industrial method despite the fact it constantly corrodes platinum electrodes, so it could not be very economical. I am very sceptical about demonstrations were no proof they get F2 was given.
But in most cases "in situ" generation of fluorine could be more than enough. Getting F2 as a stream in a pipe is a different task than just make some particular reaction with F2 happen.

BromicAcid - 7-2-2025 at 18:57

By some reason, at least according to wikipedia, electrolisys of anhydrous HF at low temperatures is still a main industrial method despite the fact it constantly corrodes platinum electrodes, so it could not be very economical. I am very sceptical about demonstrations were no proof they get F2 was given.
But in most cases "in situ" generation of fluorine could be more than enough. Getting F2 as a stream in a pipe is a different task than just make some particular reaction with F2 happen.

At one of my company sites fluorine is generated on site via electrolysis, electrodes are carbon and constantly generate CF4. Fluorine is piped directly to several of the hoods and to some of the equipment. A little outside my comfort zone but beyond cool to see.

teodor - 8-2-2025 at 01:47

Thank you for the proof that the electrolisys is still the best practical method if somebody needs F2 in a pipe, BromicAcid. I was also thinking about the possibility of carbon usage by the cost of inert CF4 but didn't know what is the real proportion (CF4 to F2). Interesting.
The reactions of disproportionation mentioned above requires carefull investigation. Dozens of patents and ideas are less than a single experiment.
It is relatively safe to get SiF4 or BF3 and use it as a proof you had F2 on some intermediate step. But it doesn't mean you can isolate it as a pure compound, I can imagine this requires some tricky techniques.

[Edited on 8-2-2025 by teodor]

[Edited on 8-2-2025 by teodor]

metalresearcher - 8-2-2025 at 03:57

So F2 cannot be isolated chemically only ?
I know that in 1986 at the centennial of Henri Moissan's first isolation of F2, a team succeeded to isolate it chemically.
https://en.wikipedia.org/wiki/Fluorine#Laboratory_routes
But this is well beyond reach of laboratory chemist (unless you are NileRed or ChemicalForce).

Bedlasky - 8-2-2025 at 11:39

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.

teodor - 9-2-2025 at 04:19

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.

1. Use not aqueous solution. But I doub't it can solve the problem, fluorine can oxidise almost everything.
2 . Use BF3 as dehydrating agent, it can be generated by heating the mixture of CaF2 (pulverised, this could be a problem), H2SO4 and a boric acid/anhydride. It doesn't attack glass if dry.
3. The similar about mixture of CaF2, H2SO4 and glass powder for SiF4 which is also doesn't attack glass.

But in 2-3 we need a method to separate CeF4 from B/Si ions.

As for SOCl2 I am not sure. Exotermic reaction, fluorine, SOF2.

[Edited on 9-2-2025 by teodor]

teodor - 9-2-2025 at 04:34

Quote: Originally posted by metalresearcher

So F2 cannot be isolated chemically only ?
I know that in 1986 at the centennial of Henri Moissan's first isolation of F2, a team succeeded to isolate it chemically.
https://en.wikipedia.org/wiki/Fluorine#Laboratory_routes
But this is well beyond reach of laboratory chemist (unless you are NileRed or ChemicalForce).

SbF5. To prepare you need anhydrous HF. Not to mention SbCl5 which is also a devil but not in comparison. If it is your routine compound already why bother with other steps. Add KF and use electricity.

[Edited on 9-2-2025 by teodor]

[Edited on 9-2-2025 by teodor]

teodor - 9-2-2025 at 06:06

To understand why HF is almost unavoidable for F2 synthesys (excluding anhydrous transitional metal compounds which could be prepared only from F2 itself) it is better to look at it as to a solvent system. To perform a reaction we need something in a form X-Fn as a solvent where X is already in highest oxidation state. For other galogens it is not so difficult, there are a lot of liquids. But binary fluorine compounds have peculiar phisical properties - when X is a metal, the melting point is TOO high, when X is not metal - too low. Only HF and SbF5 fill the range. Most close compounds are SF6 wigh m.p. -50C and BiF5 with m.p. 150C, but it seams no way to synthesyse without F2. TlF2 - 327C. HgF2 - 650 approx. , Pb, Sn - 850, the rest is zbove 1000C or below -100C.
I think some mixtures of fluorides can work depressing the melting point.
Having such a solvent it is possible not only get weak HF solutions but also synthesise anhydrous metal fluorides.

Bedlasky - 9-2-2025 at 10:25

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

woelen - 9-2-2025 at 10:58

Quote: Originally posted by vertexrocketry

chemical force made fluorine witha test tube and a fume hood so making it is not very hard

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.

teodor - 9-2-2025 at 11:58

[

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]

clearly_not_atara - 11-2-2025 at 14:46

Quote: Originally posted by woelen

Quote: Originally posted by vertexrocketry

chemical force made fluorine witha test tube and a fume hood so making it is not very hard

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!

teodor - 12-2-2025 at 03:32

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]

j_sum1 - 12-2-2025 at 17:22

I am not sure I understand the desire to do this without electrolysis.
Or is this purely a theoretical discussion?

teodor - 13-2-2025 at 07:38

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]

chornedsnorkack - 13-2-2025 at 12:45

How hard is TlF3 to produce?

teodor - 14-2-2025 at 02:34

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]

teodor - 20-2-2025 at 11:40

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]

teodor - 21-2-2025 at 01:44

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]

metalresearcher - 21-2-2025 at 13:06

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]

teodor - 21-2-2025 at 14:48

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]

clearly_not_atara - 22-2-2025 at 09:45

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.

teodor - 22-2-2025 at 12:45

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.

j_sum1 - 22-2-2025 at 18:05

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.

metalresearcher - 23-2-2025 at 07:50

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.

teodor - 24-2-2025 at 00:57

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]

teodor - 27-2-2025 at 07:58

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]

teodor - 3-3-2025 at 10:07

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.

teodor - 4-3-2025 at 14:20

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)
This file has been downloaded 24 times

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

teodor - 5-3-2025 at 03:43

Quote: Originally posted by metalresearcher