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Author: Subject: Safety and disposal of potassium fluoride
sludger
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[*] posted on 6-11-2017 at 02:15
Safety and disposal of potassium fluoride


I have about 90 grams of KF that I don't want.

I can neutralise it with an excess of calcium nitrate - anyone know for certain whether the fine CaF2 sludge is safe to dump? - say on the garden with the KNO3 as a bonus fertiliser :)

Also I want to keep some of the KF as a very dilute solution <1%. Are there any hazards with such a dilution (assuming nobody dries it out and adds conc H2S04)? Would a weak organic acid like citric, produce hazardous levels of HF in solution?

I may be being paranoid but I have nightmares about all those fluoride ions reaching my bloodstream...

Any responsible person wanting the KF for free - can come and collect it if they like :)
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[*] posted on 6-11-2017 at 02:42


I would not advise dumping it on your garden....given the circumstances I find this is about the best chance these ions have at reaching your or someboby elses bloodstream as a consequence.
Do you not have a hazardous waste disposal facility in your neighborhood that could take care of that sample in a proper way? I think they might accept it for free or for a small sum of money as the worst case....
If that is not an option then I guess sealing the sample container in a welded plastic vaccuum storage bag and keeping it on storage in a safe place is the second best route you can take. I find there is very little chance of damage happening with this option as compared to trying react it with other substances and trying to spread it across your property in attempts to dilute it.




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[*] posted on 6-11-2017 at 09:16


HF has a pKa of 3.17, the first pKa of citric acid is 3.13 so it isn't suitable. Acetic acid's pKa is 4.76 so it won't react with the potassium fluoride.



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[*] posted on 6-11-2017 at 09:43


[rquote=496649&tid=78063&author=sludger]
Any responsible person wanting the KF for free - can come and collect it if they like :)[/rquote]

Location, even just a country? Might help narrow it down if someone does want it.
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[*] posted on 6-11-2017 at 12:26


[rquote=496680&tid=78063&author=Vosoryx]
Location, even just a country? Might help narrow it down if someone does want it.[/rquote]
Oops! Sorry! Yes it would have helped. I'm in the UK, North Somerset.
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[*] posted on 6-11-2017 at 12:54


[rquote=496652&tid=78063&author=markx]I would not advise dumping it on your garden....given the circumstances I find this is about the best chance these ions have at reaching your or someboby elses bloodstream as a consequence. [/rquote]
Well that was the point of my asking. How exactly would this occur? CaF2 has a NFSA diamond of straight zeros, not dangerous, and only soluble in water to about 16 ppm. And fluorine is present in most soils at around 250 ppm anyway. Calcium precipitation is suggested [url=https://books.google.co.uk/books?id=tzcrAAAAYAAJ&pg=PA86&lpg=PA86#v=onepage&q&f=false]here[/url] but the book is 34 years old and fluorides may have become much more dangerous than they used to be :(

[Edited on 6-11-2017 by sludger]
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[*] posted on 6-11-2017 at 13:13


[rquote=496679&tid=78063&author=LearnedAmateur]HF has a pKa of 3.17, the first pKa of citric acid is 3.13 so it isn't suitable. Acetic acid's pKa is 4.76 so it won't react with the potassium fluoride.[/rquote]

Suitable for what, I wonder? Even with acetic acid there will be some HF in the solution as a minority species - I was wondering about the hazard, not [i][b]trying[/b][/i] to generate HF!

[Edited on 6-11-2017 by sludger]
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[*] posted on 6-11-2017 at 14:25


I think you are over-reacting. Sure Fluoride ions are something to be cautious with but they are not the bogeyman ready to jump up and slay you in your sleep.

Consider this -- normal toothpaste has fluoride present in the order of 0.1% -- often in the soluble form of NaF. That means if you take your KF and dissolve it in 90L of water you have a liquid about as dangerous as 90L of toothpaste. You could probably flush that without too much stress.

What I would do is use the Ca(NO3)2 to precipitate CaF2 as you intended. I would dilute and flush the supernatant liquid. Then to dispose of the calcium fluoride I would mix with sand ad cement to form a nice inert brick. That could go out in the trash. I prefer disposing of solid waste than putting stuff down the drain.

Or... just keep it or give to someone who can use it. In any case, there is no real rush.
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[*] posted on 6-11-2017 at 23:06


[rquote=496697&tid=78063&author=j_sum1]I think you are over-reacting. Sure Fluoride ions are something to be cautious with but they are not the bogeyman ready to jump up and slay you in your sleep.

Consider this -- normal toothpaste has fluoride present in the order of 0.1% -- often in the soluble form of NaF. That means if you take your KF and dissolve it in 90L of water you have a liquid about as dangerous as 90L of toothpaste. You could probably flush that without too much stress.

What I would do is use the Ca(NO3)2 to precipitate CaF2 as you intended. I would dilute and flush the supernatant liquid. Then to dispose of the calcium fluoride I would mix with sand ad cement to form a nice inert brick. That could go out in the trash. I prefer disposing of solid waste than putting stuff down the drain.

Or... just keep it or give to someone who can use it. In any case, there is no real rush.[/rquote]

I think you guys are all forgetting something here, CaF2 is the same as the mineral fluorite, which in some areas occurs in rocks the size of boulders. No need for all that mixing with cement and sand, just throw it out the window and you're good to go.
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[*] posted on 6-11-2017 at 23:31


Cement bricks are my standard disposal route for containment and convenience reasons as much as anything.
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[*] posted on 7-11-2017 at 01:30


[rquote=496697&tid=78063&author=j_sum1]I think you are over-reacting. Sure Fluoride ions are something to be cautious with but they are not the bogeyman ready to jump up and slay you in your sleep.

Consider this -- normal toothpaste has fluoride present in the order of 0.1% -- often in the soluble form of NaF. That means if you take your KF and dissolve it in 90L of water you have a liquid about as dangerous as 90L of toothpaste. You could probably flush that without too much stress.

What I would do is use the Ca(NO3)2 to precipitate CaF2 as you intended. I would dilute and flush the supernatant liquid. Then to dispose of the calcium fluoride I would mix with sand ad cement to form a nice inert brick. That could go out in the trash. I prefer disposing of solid waste than putting stuff down the drain.

Or... just keep it or give to someone who can use it. In any case, there is no real rush.[/rquote]
Thanks. That was actually my original idea too - mix the CaF2 in with some concrete. I suppose ordinary landfill is a better place than under a patio.
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[*] posted on 7-11-2017 at 01:47


[rquote=496734&tid=78063&author=Deathunter88]
I think you guys are all forgetting something here, CaF2 is the same as the mineral fluorite, which in some areas occurs in rocks the size of boulders. No need for all that mixing with cement and sand, just throw it out the window and you're good to go. [/rquote]
Thanks.
Ultra-cautious perhaps. But there are a few unknowns in extrapolating from boulders to a fine precipitate. Precipitated CaF2 has an enormous surface/volume ratio. Any water in contact with it will quickly become saturated. 16 ppm is not even remotely hazardous unless you drink it for years but I was wondering whether in naturally acidic (or chelating?) conditions there could be a potential hazard. Or even from fine particles remaining in suspension.

[Edited on 7-11-2017 by sludger]

[Edited on 7-11-2017 by sludger]
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[*] posted on 7-11-2017 at 23:20
Thanks


Thanks everyone for your replies. Been very helpful for keeping things in perspective.
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[*] posted on 8-11-2017 at 08:10


Trying to get rid of more chemicals already, sludger! You only just gave me all that other stuff, I would have gladly taken that potassium fluoride as well as the vanadium(V) oxide although I guess you probably only just found it hiding somewhere. Good luck with your safe disposal!



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[*] posted on 8-11-2017 at 15:12


[rquote=496891&tid=78063&author=18thTimeLucky?]Trying to get rid of more chemicals already, sludger! You only just gave me all that other stuff, I would have gladly taken that potassium fluoride as well as the vanadium(V) oxide although I guess you probably only just found it hiding somewhere. Good luck with your safe disposal![/rquote]

You're still around then? I was afraid you might have had a mishap with the barium nitrate! :) Hope your academic career is going well.
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[*] posted on 17-6-2019 at 05:29


[rquote=496697&tid=78063&author=j_sum1]
What I would do is use the Ca(NO3)2 to precipitate CaF2 as you intended. I would dilute and flush the supernatant liquid. Then to dispose of the calcium fluoride I would mix with sand ad cement to form a nice inert brick. That could go out in the trash. I prefer disposing of solid waste than putting stuff down the drain.
[/rquote]
Well, I do have a use for the KF but I still have about 2l of 5% solution which could be an annoyance to family if I keel over before using it. So I tried precipitating it - I used CaCl2 instead of Ca(NO3)2. The result was a thin, slightly bluish, cloudiness, not at all the dense precipitate I naively expected although it did seem viscous at first. Same result under dilution. It's presumably a colloid. On current showing it seems determined to stay in suspension indefinitely. I wouldn't be happy about sending colloidal CaF2 down the drains and I'm not sure what would happen if it were all mixed directly into a sand-cement mixture. It would be nice to think it would all be taken up into the solid matrix but that seems rather optimistic. So I need to find a way of busting the colloid or else refraining from dying. Any ideas? On the colloid, that is.


[Edited on 17-6-2019 by sludger]
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[*] posted on 17-6-2019 at 09:01


[rquote=615462&tid=78063&author=sludger]
So I tried precipitating it - I used CaCl2 instead of Ca(NO3)2. The result was a thin, slightly bluish, cloudiness, not at all the dense precipitate I naively expected although it did seem viscous at first. Same result under dilution. It's presumably a colloid. On current showing it seems determined to stay in suspension indefinitely. I wouldn't be happy about sending colloidal CaF2 down the drains …
[Edited on 17-6-2019 by sludger][/rquote]

What gives you such qualms about colloidal CaF[sub]2[/sub]? Try evaporating your solution instead, maybe?

I was looking for a way to recrystallise your calcium fluoride, but while the best solvent seems to be acetic acetic, the solubility never exceeds a few tenths of milligrams per litre, even with highly concentrated acetic acid solutions…

EDIT: Even concentrated hydrochloric acid doesn't dissolve CaF[sub]2[/sub] very efficiently. Here you go (attached PDF)

[Edited on 17-6-2019 by Keras]

Attachment: CaF2.pdf (57kB)
This file has been downloaded 5 times

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sludger
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[*] posted on 17-6-2019 at 11:46


[rquote=615477&tid=78063&author=Keras]
What gives you such qualms about colloidal CaF[sub]2[/sub]?[/rquote]
What, flushing it? Just a suspicion that a) it will have significant effect on aquatic life, including that at the sewage works and b) it goes against every regulation, law and bye-law known to man.
[quote]
Try evaporating your solution instead, maybe?
[/quote]
Sure, but he idea is to make things easier if someone [b]else[/b] ever has to dispose of it. Sloshing stuff around in a bucket is one thing, boiling it dry is another - and leaving stuff outside to dry in the sun would result in a dozen or more pet cats dying unpleasant deaths!

I suppose drying it out the CaF2 colloid [b]would[/b] eventually create a porous rock though? There'd be quite a bit of KCl to stop it getting properly compacted. It might be simpler to make several kilos of "fluoridated concrete" with a bit of extra lime using the KF solution directly.
[quote]
I was looking for a way to recrystallise your calcium fluoride, but while the best solvent seems to be acetic acetic, the solubility never exceeds a few tenths of milligrams per litre, even with highly concentrated acetic acid solutions…
EDIT: Even concentrated hydrochloric acid doesn't dissolve CaF[sub]2[/sub] very efficiently. Here you go (attached PDF)
[/quote]
I can't see the document :/ Strange though - all those H+ ions should push [F-] right down and allow a greater [Ca++] for a fixed solubility product. Odd!

Anyway fooling around with conc HCl or glacial AcAc is even less of an option than boiling it down - in this case. Fortunately, the current status is that the fluoride is KF, not CaF2 as I said so I don't have a bucket of dubious CaF2 to worry about.

Oh well, I'm using it up bit by bit. (Like I said, don't ask!) I shall just have to stay alive until it's all gone.

No matter, the main reason I reopened this thread was to inform anyone wanting to dispose of fluoride - and that would include neutralised HF - that just mixing in a calcium salt may not give a decent sludge that you can decant.
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[*] posted on 18-6-2019 at 00:56


Solubility of calcium fluoride is about 20 mg/L (that's a bit conservative, given that Wikipedia reports 16 mg/L).
WHO suggests fluorising drinking water up to 1.5 mg/L to fight off cavities. So, you're about 15 times over that amount (in reality, less than that). It's not going to be a big deal anyway.

Toothpaste holds something like 1450 ppm of NaF, or 0.15%, w/w. So roughly 12 g of toothpaste contains 20 mg of soluble fluorine anions. I'm sure a family with two adults and two kids use more than 12 g of toothpaste a day, and they use probably less than one litre of water to rince their mouths. That doesn't seem to pose a serious threat to sewering systems.

Besides, fluorine is mainly toxic because of its action on blood calcium, nerves, heart and bones. Its effect on bacteria must be totally different.

If I had to neutralise HF, I would simply throw limestone gravel in it until it stopped bubbling!
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biggrin.gif posted on 18-6-2019 at 07:20


Quote: Originally posted by Keras
||
|| Solubility of calcium fluoride is about 20 mg/L (that's a bit conservative, given that Wikipedia reports 16 mg/L).
|| WHO suggests fluorising drinking water up to 1.5 mg/L to fight off cavities. So, you're about 15 times over that amount (in reality, less than that). It's not going to be a big deal anyway.
||
|| Toothpaste holds something like 1450 ppm of NaF, or 0.15%, w/w. So roughly 12 g of toothpaste contains 20 mg of soluble fluorine anions. I'm sure a family with two adults and two kids use more than 12 g of toothpaste a day, and they use probably less than one litre of water to rince their mouths. That doesn't seem to pose a serious threat to sewering systems.
||
|| Besides, fluorine is mainly toxic because of its action on blood calcium, nerves, heart and bones. Its effect on bacteria must be totally different.
||
|| If I had to neutralise HF, I would simply throw limestone gravel in it until it stopped bubbling!

Taking the last point first. This thread is about KF; it has nothing to do with HF. Furthermore, if you attempt to neutralise HF with a solid calcium compound, the CaF2 will form an insoluble layer around it, greatly reducing the reaction rate, possibly stopping it altogether. Also, any "bubbling" is just the H+ ions reacting with (solid) CO3-- to produce CO2 and H2O. There's be no reaction with KF. And if the HF had other cations present, you would be left with XF solution, which may very well defeat the object. I am not even sure a reasonable concentration of HF would be acidic enough to self-destruct with lime anyway. It's a weak acid - even citric acid will liberate it from a fluoride. There are protocols for neutralising HF and other fluoride solutions with calcium, but they still leave you with the CaF2. It used to be acceptable to flush it, but I'm not sure that would go down well (if you'll pardon the expression) today.

As regards solubility, yes, a saturated solution of fluorite is about 15 ppm or 10 times the WHO recommendation for drinking water. But what of it? If I had a way of precipitating lumps of CaF2 from my KF solution we would not be having this discussion.

Toothpaste - 12 g of .15% NaF toothpaste contains about 18 mg (your 20?) of soluble fluoride. It does not contain 20 mg of soluble fluoride ions. More like 8 mg. But no matter, your family will undoubtedly use 100 or so litres of water so the average reaching the sewage works will be less than .1 ppm F-. Most natural waters have a lot more. If the family use a 200 ppm mouthwash - say 100 ml/day - as well that's about twice as much again. It is still easily soluble even in very hard water. KF is very soluble and I have 100 g in solution. Your suggested family would be using the equivalent of about 30 mg so my stock represents 3000 household-days.

It is certainly true that many micro-organisms are tolerant of fluoride. Insects are much more sensitive and I presume that would go for other invertebrates, some of which play a vital role at the sewage treatment plant. But the bottom line is that if someone else has to dispose of my left-over fluoride, I want them to be able to do so
a) at negligible cost and
b) with a clear conscience.

[Edited on 18-6-2019 by sludger]

[Edited on 18-6-2019 by sludger]
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[*] posted on 18-6-2019 at 07:52


UPDATE
The gelatinous precipitate seems to dissolve in excess CaCl2. Cl- does form soluble complexes with many metals whose simple chlorides are insoluble, so I was wondering whether something similar is happening here. So I repeated the experiment with Ca(NO3)2 and, sure enough, got a more persistent precipitate even when an excess was added. Unfortunately, when I then added a large excess, the precipitate dissolved. I don't know if this is because the nitrate has a lot of chloride impurity or what. But in any case, it makes the idea of precipitating the fluoride out pretty impractical without a lot of tedious titration and, even then, several litres of CaF2 jelly (jello to you Americans) is more of a nuisance than the original KF. Time to do precisely nothing with it I think!

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[*] posted on 20-6-2019 at 03:46


Thanks Keras for the pdf, which I can now see.
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