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itchyfruit
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[*] posted on 19-1-2010 at 06:03
potassium disposal


I recently read that old stocks of potassium should be disposed of by dissolving it in Propan-2-ol.
Firstly i would like to know if this dissolves the metal or just the oxide coating, in either case what would the product be.
Secondly would Butan-2-ol also work.
Most importantly how safe is this procedure.

I don't actually intend to try this (unless it produces a interesting product) i do have quite a lot of potassium and some of it is rather old, but i can't help thinking the best(and most fun)way to dispose of it is to launch it into a decent size expanse of water.
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[*] posted on 19-1-2010 at 07:52


It may not be enviromentaly correct but you can bet I would sling shot it into the middle of a lake with a camcorder ready so you can show me the results:D



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Picric-A
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[*] posted on 19-1-2010 at 07:52


With propan-2-ol it would produce potassium isopropoxide, a fairly useful organic alkali used in condensations ect.
It will work with any alcohol, producing potassium isobutoxide with butan-2-ol. It dissolved the metal and the oxide.

It would be much more usefull dissolving the potassium in an alcohol than water, firstly its safer and secondly it yields a usefull product.
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bahamuth
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[*] posted on 19-1-2010 at 08:10


I would just cut of the oxidized layers under regular gasoline (if you got cubes or large rods) and put it in fresh anhydrous hydrocarbon (heptane etc), and only get rid of the oxide layers.

Butyl rubber gloves would be good for protection.




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Picric-A
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[*] posted on 19-1-2010 at 08:17


The reason you hear old stock of potassium should be disposed of is due to the formation of explosive superoxides and peroxides, these should appear as a purple/blue layer on the metal.

For this reason it is suggested you dissolve the metal to dispose of it instead of scraping off the oxide layer which could cause an explosion, deadly if you are doing it under volatile, highly flammable gasoline :o
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DJF90
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[*] posted on 19-1-2010 at 08:55


Not quite sure where you got blue/purple from, but last time I checked the colour to watch out for was yellow/orange...
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bahamuth
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[*] posted on 19-1-2010 at 09:26


Read somewhere yellow/red color for superoxides/peroxides.

Myself, I just, when I need small amounts, carve off the oxides in an nitrogen atm.

My Potassium (200ish gram) is black/bluish in color and is at least 30 years old and stored in kerosene.







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Picric-A
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[*] posted on 19-1-2010 at 11:25


I once saw a sample of potassium which had a purple layer, might have been another compound but you right superoxides are orange.

[Edited on 19-1-2010 by Picric-A]

imagesCAJQ9IH6.jpg - 2kB
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The_Davster
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[*] posted on 19-1-2010 at 12:12


Most/all of my Na and K is blue. So is the Li to some degree. It seems to form when only a slight amount of O2 can reach the metal, such as when poorly sealed(with air trapped inside, yet a good glass seal) in an ampule, or kept in a glass stoppered vial. I think the color comes from solvated electrons from the alkali dissolving in this oxide. Needless to say, a solvated alkali metal is very very reactive, so it is no surprise to me that this effect is only seen when a finite amount of air is allowed to oxidize the metal, and then no more oxidizer is present.

A member here(len1? Garage Chemist? I forget...) did tests on heavily superoxidized K, and found it to be overrated. I had some K with chunky yellow oxide on it, and I had no issues cutting it off. Mind you, I also did not cut it under a flammable solvent, just damp with oil in a plate.

[Edited on 19-1-10 by The_Davster]




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DJF90
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[*] posted on 19-1-2010 at 13:45


The_Davster: Yes I remember seeing a picture posted here with significant amounts of a yellow oxide "growth" on it. Quite a cool looking picture if I remember correctly too :D
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[*] posted on 19-1-2010 at 14:18


The reaction of metallic K with isopropanol to form the isopropoxide (a reagent which can be used to introduce isopropoxyl groups in organic synthesis, e.g. by nucleophilic substitution on alkyl and aryl halides) would be quite highly exothermic, and indeed possibly highly explosive, with evolution of H2 which could form an explosive mixture with air. You had better take appropriate precautions.
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[*] posted on 19-1-2010 at 19:19


Quote: Originally posted by The_Davster  
I think the color comes from solvated electrons from the alkali dissolving in this oxide.


Hmm, non-stoichiometric oxide?

Tim




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[*] posted on 19-1-2010 at 22:58


It was a sodium amalgam, but I do believe I witnessed someone disposing of excess alkali by covering it with a stirred hydrocarbon (toluene or xylene), and dripping in alcohol to have a slow steady reduction. Probably helps the metal was liquid though.
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Picric-A
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[*] posted on 19-1-2010 at 23:59


Quote: Originally posted by JohnWW  
The reaction of metallic K with isopropanol to form the isopropoxide (a reagent which can be used to introduce isopropoxyl groups in organic synthesis, e.g. by nucleophilic substitution on alkyl and aryl halides) would be quite highly exothermic, and indeed possibly highly explosive, with evolution of H2 which could form an explosive mixture with air. You had better take appropriate precautions.


The reaction is exothermic but not enough to be classed 'highly explosive'.

As long as you dissolve it in small chunks, put the beaker in cold water and use an exess of isopropanol you should be fine.
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DJF90
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[*] posted on 20-1-2010 at 04:24


And if for some reason the beaker were to break??! I wouldn't like to be in that shit storm...

Fieser & Fieser say this in their book "Reagents for organic synthesis" under "Potassium":

Quote:

Procedures for safe handling. For the preparation of a potassium t-alkoxide it is necessary to remove the outer oxide-coated layer from the 20-g. lumps of commercial potassium, weigh the clean metal, and transfer it under nitrogen to a flask containing the alcohol. Two preparative procedures are described below.

W.S. Johnson's procedure (W.S. Johnson and G.H.Daub, Org. Reactions, 6, 42, (1951), and W.S. Johnson and W.P. Schneider, Org. Syn., Coll. Vol., 4, 132 (1963)).

(The following is a brief outline; the original should be consulted for details and precautions).
The oxidised surface of a lump of metal is cut off with a knife under xylene in a mortar and each scrap is transferred immediately with tweezers to a second mortar containing xylene. ...

... All metal scraps and residue are decomposed immediately by placing the mortar at the rear of the hood, making ready with a square of asbestos with which to cover the vessel of the liquid catches fire, and adding t-Butanol in small portions from a medicine dropper at a rate that the reaction does not become too vigorous.
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[*] posted on 20-1-2010 at 04:58


itchyfruit, just how much K metal have you got? (drolling question)
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JohnWW
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[*] posted on 20-1-2010 at 07:35


Fieser & Fieser, quoted above, refer to handling K metal under nitrogen. That is clearly an error, because K burns vigorously in N2 to form the ionic nitride, K3N. Also formed may be the pernitride, K2N2, and supernitride, KN2, analogous to the peroxide and superoxide. This would also form on the surface even if there was no spontaneous combustion. So you should substitute "argon" for "nitrogen".

[Edited on 20-1-10 by JohnWW]
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[*] posted on 20-1-2010 at 13:55


Wow.. Thanks for all the info allthough i'm not sure if i'm any the wiser. I have 3 separate containers one contains just over 100g which is fairly new and looks ok ones about 10years old and contains a solid cube of about 50g which again looks ok but the one i'm concerned about is at least 20 years old (sigma) it contains 3 sticks about 50g and their all covered in a white oxide (they actually look like white phosphorous) their all in mineral oil.
My concern is if i try to remove the oxide layer they'll go pop so i'm looking at slinging one stick into water just to see if there is still any k metal their or if it's jusy a stick of koh.
Djf90 and panziandy you've seen them so you must have some idea on their condition.

A_bab I'm in England and i'm fairly sure your not, if you we're you'd be welcome to a bit, however i do know of a supplier in Europe which you may allready know about if not PM me and i'll give you their details.

I'm sure i've heard abourt K reacting with N2 before but can't think where.
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[*] posted on 20-1-2010 at 16:09


Yes, they looked fine Itchy, although I didnt see the sticks. Just keep an eye out for yellow/orange oxides; white ones are fine (I've seen a jar filled with what must have been 1kg ish of Sodium metal before and it was literally covered in oxide, maybe to a depth of 2mm or so!).
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[*] posted on 21-1-2010 at 10:55


Thanks for your offer itchyfruit; I know about that cheap supplier in Europe and I don't need K. But I always like chems stories, like some (men) like to share women stories. Just realising what a sick person I must be :D

About the K sticks - although they look bad, I'm sure there's a fair amount of K left in them; the KOH layer could be just 1-2 mm.

My K developed black - blueish color; I kept it in some HDPE jar under mineral oil some 10 years in a basement (transfered from the original glass containers). It was much older than that when I got it (some 20 years more). I eventually realized that PE is penetrated by water and other gas molecules; a closer inpection revealed that some oxidation took place, but not what I would have expected. Now it's sitting back in a glass jar. I didn't attempt to remove the oxide layer, but I did feel a bit nervous when I did the transfer.

Just wondering, is there any documented incident involving an explosion caused by heavily oxidized potassium metal apart from the old book stories when potaasium was isolated by cooking it's salts with charcoal and these people used to blow up from time to time ?
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[*] posted on 21-1-2010 at 11:03


Alkali metals should be stored in PET bottles ( if you must use plastic- metal is usually preferred) due to PET's complete umpermeability however its reactivity to alkali hydroxides may be a problem...
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[*] posted on 21-1-2010 at 14:43


I've disposed of heavily oxidized and contaminated Na and K by putting it in hexane and slowly adding isopropanol while stirring.

This way your isopropanol is diluted and the hexane acts as a buffer for the heat.

The same method works for destroying n-BuLi, NaH and KH. Could also work for aluminiumhydrides using ethylacetate instead of isopropanol.

Methanol is not recommended as I've witnessed a few bursts of flames when this method was used. :D

Cutting of black residue from potassium in air is perfectly safe as long as you don't mess around for half an hour. I do this on a regular basis to make Na/K alloy for drying THF and ether.




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DJF90
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[*] posted on 21-1-2010 at 22:25


Personally I would store alkali metals in a glass jar, to avoid the aforementioned issues with plastic. Metal is good, but has limitations. A friend of mine had a tin of potassium metal which had eaten through the can, despite being in its original unopened state.
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[*] posted on 24-1-2010 at 04:49


Sodium destroyed via isopropanol and cut under xylene/gasoline is generally OK.

I wonder however if those referring to treating potassium that way have actually much experience with potassium? Its a completely different beast. What you get destroying K under isopropanol is an alcohol flame in the beaker. If by any chance you happen to be mad enough to pour it straight out of the bottle into the beaker then the flame will most likely take up the bottle as well.

Cutting K under xylene is like playing russian roulette, small pieces of K are pyrophoric and you get a mixed K/xylene fire. Have a big bowl of sand or salt plus strong nerves at the ready.

I recommend cutting K under paraffin, with this I never had any problems.
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[*] posted on 24-1-2010 at 07:50


len1 - this is why I would use t-butanol, as Fieser also advocates. The reaction is much more controlled, although adding the alcohol to potassium in a solvent (to act as a dilutant) would likely be a better method than they suggest. Also, I have never heard of potassium forming nitrides as JohnWW says, bar atom collision experiments. My impression was that only lithium could form a stable nitride, and thats what I recall my thermodynamic calculations showing last year!
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