Sciencemadness Discussion Board - potassium disposal

potassium disposal

itchyfruit - 19-1-2010 at 06:03

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.

Sedit - 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

Picric-A - 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.

bahamuth - 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.

Picric-A - 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

DJF90 - 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...

bahamuth - 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.

Picric-A - 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

The_Davster - 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]

DJF90 - 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

JohnWW - 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.

12AX7 - 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

Mumbles - 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.

Picric-A - 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.

DJF90 - 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.

a_bab - 20-1-2010 at 04:58

itchyfruit, just how much K metal have you got? (drolling question)

JohnWW - 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]

itchyfruit - 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.

DJF90 - 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!).

a_bab - 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

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 ?

Picric-A - 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...

vulture - 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.

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.

DJF90 - 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.

len1 - 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.

DJF90 - 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!

itchyfruit - 24-1-2010 at 11:19

As the aforementioned pieces of K metal are soon to be the property of yourself i'm sure you'll work it out for good or bad

But i am interested for future reference is it ok to cut K under yellow LIQ paraffin that's used for oil baths ?
The reason i ask is i have quite a lot of this(but no oil bath) and the lady in my local chemist is getting suspicious about my LIQ paraffin purchases.

BTW just PM me a list of the other bits you want, i can't be arsed to do another list!!!!

bahamuth - 24-1-2010 at 12:00

You can cut your K metal under any hydrocarbon solvent that is water and oxygen free, but your liq. paraffin might be a little tedious to wash away when you want your K metal for an experiment.

You want to cut it under some kind of "semi" volatile hydrocarbon, like toluene, xylene or petroleum ether bp. 60-80 or bp. 80-100. Outdoor cooking stove purified gasoline (heptane & its isomers) works great.

len1 - 24-1-2010 at 17:13

Of course, you can do anything, the question only is if the consequences you get are what you want. The potassium after being cut needs to be removed from under the inflammable solvent, which gives you a potassium/air/hydrocarbon surface - if that has not already been achieved during sloshing of the solvent while cutting it. Small pieces of potassium often autoignite in air - that is what pyrophoric means. If its just potassium there will be a light puff, a lightening on the surface of the K due to the protective oxide formed, and that will be the end of it. Under toluene/xylene that will only be the start of it due to the ensuing fireball.

K will also autoignite under kerosene if touched by anything with the slightest moisture content - such as a knife blade.

Pyrophoric metal and inflammable solvent is a bad mixture.

If anyone wants to follow the stupid advice above, you have been warned.

[Edited on 25-1-2010 by len1]

a_bab - 24-1-2010 at 23:47

Potassium metal in the amounts you have there could be disposed off in the snow (providing you have it). I'd say something like up to 5-10 grams at once works great and it's safe.

Just toss in piece by piece outside in the snow; the potassium take fire very fast even if there is parafin oi on it, and burns quietly with no explosions. If you did this in your garden, you'll recover some of the potassium in your vegetables later in the summer

On the other hand just as len1 said, I feel quite nervous just about the thought of having the potassium in something else than parafine oil. Some would even go for ligroin. It's quite pyroforic; it takes fire in the air almost like WP.

bahamuth - 25-1-2010 at 09:11

Tried to provide a means for the recovery of, even if cheap, Potassium, to remove the oxide layer and get fresh potassium.

Just delivered what I have learned and read, in "Hazardous laboratory chemicals disposal guide by Margaret-Ann Armour" in the section concerning Potassium.

If he want to get rid of it, throw it in a lake or something and enjoy the bang, easypeasy.

If recovery was to be made, how do you len1, propose it to be done, and how do they do it in the laboratory when they want fresh exposed K?

Not nearly as reactive, sat for hours cutting about a kilo of Sodium this way under petroleum ether bp. 60-80 C.

Edit:

The more I read this tread the more anxious I get, was not worried by my potassium metal before, had respect, but now, sort of worried....

[Edited on 25-1-2010 by bahamuth]

12AX7 - 25-1-2010 at 11:46

Scroll down to paragraph 6, line 21:
http://204.232.255.211/30/12.html

The Chemical History of a Candle

Michael Faraday

Quote:

I am going, then, to burn this potassium in the carbonic acid, as a proof of the existence of oxygen in the carbonic acid. [In the preliminary process of heating the potassium exploded.] Sometimes we get an awkward piece of potassium that explodes, or something like it, when it burns. I will take another piece, and now that it is heated I introduce it into the jar, and you perceive that it burns in the carbonic acid ...

("Carbonic acid" erroneously referring to the gaseous anhydride, CO2.)

Oh, for the days when lectures were fun.

Tim

[Edited on 1-25-2010 by 12AX7]

len1 - 25-1-2010 at 16:45

Yes I rather thought your experience was more with sodium - but one should never think of the two metals as the same thing.

My method to rid K of its crust is to heat it under paraffin to 70C, the insert a 50ml volumetric pipette with short inlet preliminarily flushed with argon so it punctures the oxide skin by 2-3mm and suction up the desired amount of perfectly clean metal. The floppy gray-blue oxide 'skin' is left behind. Its in none of the books, but thats typical

vulture - 26-1-2010 at 11:27

Quote:

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.

Yes, I cut up to 6g of K into small 2x2mm pieces every two months. No problem at all.

As for destroying K, I did not say you should pour pure iPrOH onto it. Hexane and then adding drops of iPrOH to the hexane works fine.

Quote:

K will also autoignite under kerosene if touched by anything with the slightest moisture content - such as a knife blade.

I have no idea where you get this outrageous claim. I always take blocks of K out of the parrafin using a knife which has been exposed to air. After cutting, the K & Na go into THF which comes from a fresh bottle, but obviously isn't moisture or oxygen free. Neither does this autoignite. In wet diethylether you get vigorous bubbling, sure, but no ignition.

[Edited on 26-1-2010 by vulture]

[Edited on 26-1-2010 by vulture]

len1 - 26-1-2010 at 13:35

Quote:

I have no idea where you get this outrageous claim

Experience.

And the same shows it is to be trusted more than somebodys post.

vulture - 26-1-2010 at 14:41

You do realize that your experience is also nothing more than something described in a post? Which gives it the same value as mine?

Are you insinuating I make this all up?

You think cutting solid K is unsafe but pipetting the molten metal is safe?

I really do not understand the aim of your scaremongering towards handling potassium.

Quote:

the insert a 50ml volumetric pipette with short inlet preliminarily flushed with argon

You do also realize that you are contradicting your claim here of traces of moisture igniting K? Merely flushing with argon does not remove traces of moisture.

[Edited on 26-1-2010 by vulture]

[Edited on 26-1-2010 by vulture]

len1 - 26-1-2010 at 21:49

Yes I was going to the trouble of excluding oxygen by purging my pipettes with argon, but not baking them at 150C to remove the moisture, because I was so stupid.

Thanks for clearing this up. Now I will not have exploding pipettes anymore.

JohnWW - 27-1-2010 at 00:18

On the subject of K metal, along with Rb and Cs, together with inert gases such as argon, that reminds me: I wonder if anyone has yet succeeded in isolating an ionic compound, containing K+ or Rb+ or Cs+, in which the alkali metal has given up its lone valence electron to either He or Ne, to form an anion He- or Ne- which would be isoelectronic with metallic Li or Na? Such a compound would be extremely reactive, much more so than Li or Na. In this connection, alkali-metal anions have been isolated, containing Li- or Na- or K-, but they would be more easily formed than He- or Ne- because of the greater atomic charge and the stabilization afforded by spin-pairing in the s orbital.

The possibility has been thought of, e.g., regarding hypothetical cesium neonide, see http://www.eric.ed.gov/ERICWebPortal/recordDetail?accno=EJ26... or http://pubs.acs.org/doi/abs/10.1021/ed059p637 , which was originally in Journal of Chemical Education, v59 n8 p637-39 Aug 1982. Also http://pubs.acs.org/doi/abs/10.1021/ed061p566.4 . If it could be produced, it would mean that K, Rb, and Cs cannot be safely handled and stored under He or Ne, but can still be under Ar or Kr or Xe.

vulture - 27-1-2010 at 11:50

Quote:

Yes I was going to the trouble of excluding oxygen by purging my pipettes with argon, but not baking them at 150C to remove the moisture, because I was so stupid. Thanks for clearing this up. Now I will not have exploding pipettes anymore.

There is no need for sarcasm. It's not about what you know. It's about what you say in your post. You lecture others about recklessness and bad information when it comes to handling potassium, yet you omit an essential safety procedure from your own post!

You also glossed over other important points I raised without even addressing them.

Do I really have to make a movie next time when I cut potassium to prove to you I'm not selling BS? Or will you also dismiss that with sarcasm without addressing the real issue?

turd - 27-1-2010 at 12:47

I can also not relate to the inherent danger of potassium. We used it to dry THF for many years and treated it just like sodium. We even used the same rusty butter knife for both metals.

Maybe we were just lucky, but nobody ever caught fire. The big advantage of K over Na is of course that there is no need to make small pieces. Now NaK I hear is an entirely different matter, but unfortunately I've never seen it in action - maybe I should make some one day (in an ampoule).

vulture - 27-1-2010 at 13:41

Quote:

I can also not relate to the inherent danger of potassium. We used it to dry THF for many years and treated it just like sodium. We even used the same rusty butter knife for both metals.

We use Na-K alloy to dry THF and ether. It is prepared by cutting the Na and K and putting them in the respective solvent. The alloy doesn't form until the solvent is refluxed, after which it forms silver globules which keep a fresh layer on the outside and stay liquid at RT. Obviously this should be kept under argon at all times.

entropy51 - 27-1-2010 at 17:37

What's the big deal about Na-K? It has been used to cool computers.

turd - 28-1-2010 at 02:31

Quote: Originally posted by vulture

What's the advantage of keeping it liquid at RT? That it keeps the THF dry over night?

JohnWW - 28-1-2010 at 04:47

The eutectic alloy Na-K composition is liquid at room temperature. It consists of 78% potassium and 22% sodium, is liquid from −12.6 to 785 °C, and has a density of 866 kg/m³ at 21°C and 855 kg/m³ at 100°C, lighter than water. The entire composition range of 40% to 90% K is in fact liquid at room temperature. See http://en.wikipedia.org/wiki/NaK .

Uses for it, e.g. as a liquid coolant in heat exchangers, with a thermal conductivity much greater than that of water (but cheaper and lighter and more conducting than Hg or Ga) have been based on this property, particularly in experimental fast-neutron nuclear reactors. It could also be used in thermometers. However, of course, it suffers from the disadvantage of being dangerous if exposed to moisture or alcohols (or several other types of highly polar organic compounds) or any gas other than the elemental inert gases. It also has only about 1/4 of the thermal heat capacity of water, and a high surface tension.

turd - 28-1-2010 at 07:21

Huh?
That didn't answer my question at all. I wanted to know what the advantage of using NaK/THF over K/THF is, given that K is also liquid in refluxing THF.