Sciencemadness Discussion Board - Make Potassium (from versuchschemie.de)

What about catalyst? What comes to mind is that the K forms, but when the Mg runs out the catalyst turns right around and starts consuming the K metal [...].

My thoughts exactly. Then again, I could be very wrong,
[Edited on 9-3-2013 by Mailinmypocket]

You probably are: there just isn't enough t-butanol to consume much K. To be precise, even if your conversion was 100 %, you'd still only get 90 % of the theoretical K because out each 1 mol of actual, dry KOH 0.1 mol gets tied up as K t-BuO at the end of the reaction. It's all explained higher up in much detail. During all of the reaction, including post-reaction, 0.1 mol of each mol of K is 'tied up' as K t-BuO. But 0.1 mol is only 10 % of the total quantity of K (if the KOH/cat. ratio is 1 to 0.1 mol)

No, if you had K forming and then disappearing the most likely explanation is an 'outside' scavenger to reoxidise the K. Oxygen, water or 'something' in the solvent...

[Edited on 10-3-2013 by blogfast25]

blogfast25 - 10-3-2013 at 06:36

If you believe you've failed, at least subject a sample of the gunge to water. ANY K in there will reveal itself quickly.

It could also be useful to subject a sample of the gunge to strong HCl to check solubility and MgO. Then separate off the watery phase and check for magnesium: magnesium carbonate is poorly soluble in water, so after neutralisation you should get some precipitation with Na2CO3.

Mailinmypocket, what temperature control/monitoring do you have in place? Going by your last pictures I see everything is fine but not a hint of thermometer. That's a serious limitation. There's room for a third, small hole in that stopper, for a thermocouple probe. Household digithermometers go up to 200 C, it's what I used.

[Edited on 10-3-2013 by blogfast25]

Mailinmypocket - 10-3-2013 at 07:29

If you believe you've failed, at least subject a sample of the gunge to water. ANY K in there will reveal itself quickly.

It could also be useful to subject a sample of the gunge to strong HCl to check solubility and MgO. Then separate off the watery phase and check for magnesium: magnesium carbonate is poorly soluble in water, so after neutralisation you should get some precipitation with Na2CO3.

Mailinmypocket, what temperature control/monitoring do you have in place? Going by your last pictures I see everything is fine but not a hint of thermometer. That's a serious limitation. There's room for a third, small hole in that stopper, for a thermocouple probe. Household digithermometers go up to 200 C, it's what I used.

[Edited on 10-3-2013 by blogfast25]

I will test the residues later this afternoon following your suggested procedure. The temperature is indeed being monitored by an electronic kitchen thermometer, it can be seen in the background. The probe was buried under the sand, but not directly on the bottom of the pot. The temp fluctuated between 220-240 thanks to my shitty old stove, I would have used the hotplate but the Corning manual says to not heat sandbaths on it so I played it safe. Ill report on the outcome of the gunge test later...

blogfast25 - 10-3-2013 at 08:58

MIMP:

Blindness is a terrible affliction: I didn't see your nice probe there, my bad

. That should really rule out any temperature problems.

elementcollector1 - 10-3-2013 at 20:51

I'm going to see if I can get a stronger heat source. What I have now is a Hamilton electric hot plate, and while it appears to reach 200 C, there's no guarantee that that's true.
A propane burner should be more than good, correct?
Also, per the oxygen and water destroying the K: Is there any way to remove these two without going inside? Again, parents. If I could, I'd flush the container with argon.
After 6 failures in a row, I'm beginning to lose hope for this experiment. I'll keep trying, and posting the results, though.

blogfast25 - 11-3-2013 at 04:16

I'm going to see if I can get a stronger heat source. What I have now is a Hamilton electric hot plate, and while it appears to reach 200 C, there's no guarantee that that's true.
A propane burner should be more than good, correct?
Also, per the oxygen and water destroying the K: Is there any way to remove these two without going inside? Again, parents. If I could, I'd flush the container with argon.
After 6 failures in a row, I'm beginning to lose hope for this experiment. I'll keep trying, and posting the results, though.

Propane should be fine. I used it. Always with sand bath though.

If your plate does 200 C your flask will by definition be slightly cooler. You really should have a temperature probe poking into the reaction chamber...

Argon shouldn't be necessary. If there is something oxidising the K it shouldn't be there in the first place and argon probably wouldn't get rid of it. Maybe it really is time to try Shellsol D?

hyfalcon - 11-3-2013 at 04:51

Maybe it really is time to try Shellsol D?

If you can find a source please u2u me. I've searched and searched for this in the US and can't find it.

blogfast25 - 11-3-2013 at 05:05

I'm not US based so I can't help you there. But higher up in this thread a US based poster did find it without too many problems. The patent is also a US patent, I seem to recall...

elementcollector1 - 11-3-2013 at 07:21

Say, could the lamp oil be dried with NaOH or some such? Sodium metal? (It would be a bit ironic, but still). Anhydrous MgSO4 is good, but these two are a bit easier to prepare.

I'll see if I can pick up a lab thermometer or something from school today (hopefully it can do 200 C). Maybe then I can check the boiling point of this stuff, and see how it matches. If the school is a no-go, then I'll probably get a meat thermometer from Target.

Fantasma4500 - 11-3-2013 at 07:50

i think ive found a person who can approve that its possible to melt potassium nitrate, lead electricity through this molten potassium nitrate and without the formed K metal reacting explosively or at all with the KNO3..
he use KNO3 and KI to get iodine vapours in what i saw, but he said that when a drop of water is added to one of the cathodes (the one with potassium ofcourse) it sparks violently
imagine turning magnesium salts into pure magnesium metal with electricity and abit of heat..
hope some of you guys will be as hopeful towards this as i am.. because i am hoping that this could be possible.. it will open up for so many possibilities and rare metals by electrolysis without +1000*C uncontrolled reactions

elementcollector1 - 11-3-2013 at 07:54

I'm guessing that mixture melts at 500 C? It could work, but any potassium produced is immediately exposed to oxygen. Sodium has been done with NaOH, however, by use of a homemade Castner cell ("Hot electrochemical sodium").
Molten electrolysis isn't exactly a new idea, it's just hard to accomplish effectively. That, and the reactivity of some metals such as K, makes this catalytic synthesis more viable (or does it?).

Fantasma4500 - 11-3-2013 at 08:26

true.. but ive never seen it done more than once with nitrates, and using nitrates would be abit more smart as you dont have highly corrosive liquid (which can etch glass!!)
also its usually easier to get a hold of a metal nitrate than the metal hydroxide (or well atleast easier to make the nitrate, especially when you got 10L HNO3)
usually nitrates have a very low melting point aswell..

elementcollector1 - 11-3-2013 at 08:32

It's viable, but I'd be reluctant to play with a powerful oxidizer and electricity.
Hydroxides have an even lower melting point, but tend to be harder to store or use due to the reaction with glass, as you said.

On a side note, our school thermometers go to only 50 C. Pathetic! I'm going to Target ASAP for a meat thermometer.

Fantasma4500 - 11-3-2013 at 09:06

yes it might be potentially dangerous to play around with molten nitrates.. as it tends to ignite fuels on contact.. (:
the guy i saw do it used graphite electrode, tho.. (strange?)
but the idea is to be able to do it much more controlled than making a thermite and tonnes of tiny pieces of metal..
by electrolysis of lead bromide you get bromine (ofcourse) and lead, the more interesting in this is that it forms on one of the electrodes and then falls to the bottom, which means you can break the solid lump up after its cooled down and you should be able to get a solid piece of metal, i wonder if this works with potassium or sodium aswell, or if its too light to fall to the bottom..

elementcollector1 - 11-3-2013 at 09:08

Again, read "hot electrochemical sodium". It's a stickied thread in Technochemistry, IIRC.

Fantasma4500 - 11-3-2013 at 09:24

really..?? well thanks.. havent heard of it before tho

blogfast25 - 12-3-2013 at 03:13

Quote: Originally posted by Antiswat

i wonder if this works with potassium or sodium aswell, or if its too light to fall to the bottom..

Downs Cell:

http://en.wikipedia.org/wiki/Downs_cell

elementcollector1 - 13-3-2013 at 08:27

Anyway, back to business. I got a temperature probe with a range of -40 C to 230 C, and while it's not made for long, sustained readings, it seems to work just fine. The first test I made for the BP of the lamp oil oddly resulted in the lamp oil never boiling, and getting stuck at 165 C. Maybe I'll remove some sand from the bath and try that again.

Mailinmypocket - 13-3-2013 at 09:14

Okay, got around to testing the residues and I'm thinking not much is going on as far as the reaction is concerned.

Image 1: Isolated white crust only, in the left tube. Magnesium bits in the right tube. I tested Mg just... Because.

Image 2: After adding approximately 15% HCl to both tubes, the white residue made a frightening cracking sound. I assume this is because it is simply KOH which does not like being thrown into acid, it heated up substantially. The Mg reacted predictably, with violent bubbling and heat to the point of steam production.

Image 3: After neutralization of the tube containing the dissolved residue with sodium carbonate, no precipitate materialized.

The Mg surfaces were dark- almost black, I don't think the size of the particles are small enough...

blogfast25 - 13-3-2013 at 09:48

Quote: Originally posted by Mailinmypocket

[Image 2: After adding approximately 15% HCl to both tubes, the white residue made a frightening cracking sound. I assume this is because it is simply KOH which does not like being thrown into acid, it heated up substantially. The Mg reacted predictably, with violent bubbling and heat to the point of steam production.

Yes, solid KOH + strong HCl means a lot of heat of neutralisation all at once.

You've more or less proved no reaction between the KOH and Mg took place. While particle size could be a factor, I seem to recall someone having success with shavings instead of powder.

In this reduction reaction several steps occur, some of them require solid/liquid reactions, others do not. Without knowing what is the slowest step in the chain it's hard to cast 'the blame' on any given factor.

But by all means try a finer grade of Mg and see what you get....

blogfast25 - 13-3-2013 at 09:51

Anyway, back to business. I got a temperature probe with a range of -40 C to 230 C, and while it's not made for long, sustained readings, it seems to work just fine. The first test I made for the BP of the lamp oil oddly resulted in the lamp oil never boiling, and getting stuck at 165 C. Maybe I'll remove some sand from the bath and try that again.

You said the plate was up to 200 C. What you observe would fit that. With the plate at 200 C the flask can NEVER actually reach that temperature because of insulation provided by the sand and heat losses of the flask to the environment. 165 C is quite different from 200 C. You'll either need a hotter plate or switch to propane.

I think this is real progress for you, EC1...

[Edited on 13-3-2013 by blogfast25]

blogfast25 - 13-3-2013 at 10:05

Just to try and focus the attention a bit, I'll quote myself from here:

http://www.sciencemadness.org/talk/viewthread.php?tid=14970&...

Below is the simplest form of the supposed reaction path (it was later refined a bit).

It’s very clear from Example 1 (potassium from KOH) that the initial hydrogen comes from moisture reacting away (assuming for argument’s sake that neither the patent nor pok are hoaxes). During that step no t-butanol is present yet.

The actual reactions in the second step could possibly be (with R a t-alkyl group):

Initiation:

KOH + ROH < --- > KOR + H2O

Wherever this equilibrium lies, it may be pulled to the right by:

Propagation:

2 KOR + Mg < --- > 2 K + Mg(OR)2

Which may itself by pulled to the right by:

Termination:

Mg(OR)2 + H2O --- > MgO + 2 ROH

Of the last step at least we can be reasonably sure: Mg alkoxides should be quite prone to hydrolysis.

2 KOH + Mg --- > 2 K + MgO + H2O

… would be the overall reaction in that scheme. A calculation using NIST values of HoF and S at 298 K, shows the ΔG = ΔH – TΔS to be about -48.6 kJ/mole (of Mg reacted, @ 298 K) [edit: or about - 100 kJ/mole of K produced], so just barely thermodynamically possible. For higher temperatures this requires a correction that is likely to be small. The low change in Gibbs free energy could explain low reaction rates.
[Edited on 4-12-2010 by blogfast25]

Assuming this is at least conceptually correct, we still don’t know what is the rate limiting step. If the rate limiting step involves solid magnesium, then fineness (total available surface area) would almost certainly affect overall reaction speed, if no then probably not.

[Edited on 13-3-2013 by blogfast25]

elementcollector1 - 13-3-2013 at 20:50

You know, I wonder what it would do to put a piece of sodium in a glass tube or sep funnel at the top of the condenser, to see if it really is atmospheric H2O or O2 that's destroying my K. I have plenty of sodium. This probably won't be my production method, but it would be another interesting test (also it would be interesting to see what the sodium metal does). Now, how can I make some sodium wire? Would a plastic pipette be strong enough, or would the sodium need to be softened first?

violet sin - 13-3-2013 at 22:48

http://youtu.be/btvQdW4fXJE

youtube vid. the guy uses a clay extruder from an arts and crafts supply store to make Na wire. worth a quick look.

condennnsa - 13-3-2013 at 23:20

antiswatt, I think that the guy you mentioned is this guy:
http://www.youtube.com/watch?v=SCiP-COawoU
http://www.youtube.com/watch?v=0pz0qB9yKVk

there is also a patent from 1894 claiming that it's possible to melt kno3 in an aluminum container and electrolyze it to potassium and NO2
here it is

https://docs.google.com/a/google.com/viewer?url=www.google.c...

blogfast25 - 14-3-2013 at 05:51

You know, I wonder what it would do to put a piece of sodium in a glass tube or sep funnel at the top of the condenser, to see if it really is atmospheric H2O or O2 that's destroying my K. I have plenty of sodium. This probably won't be my production method, but it would be another interesting test (also it would be interesting to see what the sodium metal does). Now, how can I make some sodium wire? Would a plastic pipette be strong enough, or would the sodium need to be softened first?

I think you should try and get the temperature right first. When you've eliminated the water as H2, your flask will have been somewhat purged of air (oxygen) because of the H2 flux. And when you have a nice steady simmer you're also replacing air with solvent vapour...

elementcollector1 - 14-3-2013 at 07:38

Going to see if I can borrow a hotplate from the club, then. (I borrow a lot of things, don't I?)

On a side note, 50 pages! Is this the longest running SciMad thread?

EDIT: Trial 7 is running as we speak. Definitely over 200 C, the solvent appears to be boiling (or is that hydrogen?) and things seem to be going smoothly. No visible potassium, though...

[Edited on 15-3-2013 by elementcollector1]

elementcollector1 - 14-3-2013 at 18:10

What in the name of Feynman?
My solvent's turning brown...

EDIT: And now I am out of available tert-butanol. Fantastic.

[Edited on 15-3-2013 by elementcollector1]

elementcollector1 - 14-3-2013 at 19:26

Welp. Success?
Changes:
-Used Pok's masses and such instead of NurdRage's (same as Run 6)
-Corning hotplate w/stirring, borrowed from Science Club

I think what happened is this: Potassium metal was formed, but the alcohol left solution (it kept freezing just above, I had to push it back in with a metal rod) or got consumed in some side reaction, and the potassium reacted with oxygen to form a brownish-black potassium oxide. HOWEVER! I have proof that the metal was formed at some point.
When I dumped out the contents into a beaker to inspect them, some was left behind in the flask. I decided (foolishly) to clean this out with water. There was a hiss at first, then a bright flare of lilac and a pop!

I think there are very small K metal pieces dispersed in this opaque solution of potassium oxide. But how to collect them? My best guess would be to heat the beaker to the melting point of potassium (about 65 C) and turn on slow stirring, without a stirbar. The potassium, being paramagnetic, should collect in the center and coalesce, forming a larger and possibly more easily removed sphere.

In the meantime... What do I do with all this potassium oxide? I have the beaker under plastic wrap, in a halfhearted attempt to prevent too much air access (the solvent well do well enough as it is for a short while). Is there any way I could extract it, or turn it back to potassium?

blogfast25 - 15-3-2013 at 04:17

EC1:

Forget about the potassium oxide, it's basically irrecoverable, IMHO and not worth trying.

But clearly you're very nearly there! Too vigorous cooling appears (again) to be your problem: without or with very little t-butanol in the flask, the reaction speed MUST suffer (that is at least one thing we can safely conclude from the proposed reaction scheme). The fact that at least SOME K formed when much of your catalyst wasn't even playing is a good sign. I'm now wondering is some of the other failures we've seen from other experimenters may also be due to freezing t-butanol. You might still be a trailblazer!

MIMP:

I've had solvent turn brown too but it didn't impeded K being formed.

elementcollector1 - 15-3-2013 at 07:14

So, how to solve this? The water in the fume hood is bound to be a little bit warmer... Maybe...
But is there a way to warm the water up so that it won't freeze the t-butanol, but can still recondense it?

Oh, right. Repoured stuff into the flask, and there were many tiny but bright specks of light, likely more potassium. Doesn't look like anything big, but as soon as I get home I'll give coalescence a shot.

[Edited on 15-3-2013 by elementcollector1]

blogfast25 - 15-3-2013 at 08:52

With a BP of only 83 C, the t-butanol will always condense fully, so no worries there. To avoid actual freezing up, the internal wall of your condenser should never be below 25 C. Bear in mind that your boiling solvent will help heat up the condensing channel.

I'd go with air cooling: just a bit of air from an aquarium pump, pumped in from the top inlet. Alternatively, try first with no air at all: just monitor the top of the condensate (aka the 'condensation ring') in the condenser, making sure it doesn't climb too high (that's when loss of vapours could occur).

It's quite interesting that many of us used nothing more than a glass pipe with some moist kitchen towel and didn't suffer these frozen t-butanol problems. Sometimes simples does it well...

[Edited on 15-3-2013 by blogfast25]

elementcollector1 - 15-3-2013 at 10:46

Hmm. I know a friend who has an aquarium pump, I'll ask him.

EDIT: 3 whole milliliters of t-butanol were hiding from me! Bad alcohol!

[Edited on 16-3-2013 by elementcollector1]

blogfast25 - 16-3-2013 at 06:42

Well, good luck. My little finger tells me you're about to make significant balls of potassium.

When I think about it logically, I think too low a temperature (of the plate/flask) goes a long way to explain your fails. Below BP there isn’t much solvent vapour in the head space of the flask but the vapours nonetheless contain much t-butanol, due to its low BP and high vapour pressure at > 160. The too cool (helped by little condensation of solvent, which would otherwise heat up the condenser inner surface quite a bit) condenser surface then condenses the t-butanol, unfortunately freezing it in the process. A downward flow of solvent condensate would also have helped redissolving any solidified t-butanol but there wasn't much of such a flow. Too low reaction temperature combined with subnormal concentration of t-butanol = no K. It all fits rather well…

[Edited on 16-3-2013 by blogfast25]

elementcollector1 - 16-3-2013 at 08:46

Your little finger talks to you?

I'm going to see if I can get adequate cooling with a condenser filled with non-moving water. As the reaction progresses, the water will get steadily more heated, but will never reach 100 C (although it can still reach 85, causing the butanol to boil off). In that case, I'll probably flow a bit more water into there, and shut it off again.

blogfast25 - 16-3-2013 at 10:44

Your little finger talks to you?

Good idea. Make sure to monitor temperature, if only by hand: 4 h is a long time and there's updraft from the sand bath too as an additional heat input to the mante water...

elementcollector1 - 16-3-2013 at 13:37

Well, the sand bath is at least 200 C (probably well over that), the coolant flow is as low as possible (no frozen alcohol visible), and we're 2 hours in. Stirring is at the lowest possible setting. No K visible....

EDIT: Pictures! >

The general setup. Hot plate's on 3, stirring's on 3... if that means anything to anyone.

The mush in the reaction flask. Not much happening, apart from a crust on the top of the oil.

Well, no needle-like crystals visible. I guess the t-butanol is working, then?

[Edited on 16-3-2013 by elementcollector1]

elementcollector1 - 16-3-2013 at 15:48

4 hours in: No visible potassium, and all the magnesium is pretty much still there. And still shiny. Ugh... Maybe #7 was just luck? The only difference is that I didn't add all the catalyst at once, instead doing the 'staggered addition' thing. I guess I'll go back to what I was doing before, but after 8 different tests, this is just getting disappointing. Going to continue reflux for a few hours more, in case a scientific miracle happens.

EDIT: I leave this thing alone for an hour and it turns brown again. Well, it's better than nothing. I did turn up the heat and the stirring at the end, so I guess that's a point for next time. The hot plate surface is oddly discolored slightly yellow, so that's weird. Anyway, I'm expecting there to be small bits of potassium in there, which I'll never get to coalesce, and the majority of the potassium is now the oxide.

So, notes for next experiment:
-bring up the heat more at the beginning
-add the catalyst all at once - again
-More stirring?
-Running out of KOH now. I probably have enough for the next two runs, but not much after that.

You know, I noticed that a significant portion of Mg remained unreacted in both of the last two runs (kind of assuming on this one, since I haven't examined the gunk yet). Some of it obviously reacted, but not all of it. Also, there is still a significant quantity of t-butanol in the mix, as it spewed white gas whenever I put it on a heated surface without cooling (fortunately, it still recondensed).

[Edited on 17-3-2013 by elementcollector1]

elementcollector1 - 16-3-2013 at 19:06

Well, call me a liar!

That represents 0.8g of pure potassium, or a theoretical yield of 18.8%. Pretty bad, but it's something. What's the normal yield for this procedure?
Anyway, the reaction mix had turned brown again, so I assumed no potassium was present. I dumped it out into a strainer, and boy, was I wrong.

Oh, and then stuff lit on fire.

NurdRage told me to clean anything that handled potassium metal with alcohol. I did, and it lit on fire anyway. Melted one of my best plastic containers too. Well, long story short, a towel got burned, the strainer got covered in potassium hydroxide and oxide (thank science I saved most of the potassium in a beaker with fresh solvent first), and nothing else was recoverable afterwards.

DO NOT USE ISOPROPANOL TO CLEAN POTASSIUM METAL. I used 99% by volume from the local pharmacy, and it caught fire anyways. Then the lamp oil got mixed in, and that didn't help much. Well, nothing important caught on fire, so nothing harmed too badly.

Well, anyway, after this little fiasco I cleaned and coalesced the rest of the potassium, but quite a number of 1mm spheres of the stuff are left in the beaker. I can't seem to melt or coalesce these, and I'm frankly too scared to get rid of them with isopropanol. Anything I can do to safely destroy them?

I put the potassium metal spheres, once they had cooled, into a test tube with fresh mineral oil - good enough for a while. Any tips on ampouling for a longer storage?

By the way, this stuff is SOFT - even when it was solid and cold, the tweezers used to scoop it out into the test tube severely dented the surface. I mean, that's to be expected - sodium can be cut with a knife. But still, it's such an interesting property for a metal to be soft.

Elementcollector1, signing off - it's been a long run after 8 tries, but it was worth it! And then I have to do it again in front of my peers.

...Monologue over.

mr.crow - 16-3-2013 at 19:13

Congrats!!!

Try using an inert solvent and adding the isopropanol very slowly. You could also use your tert-butanol

Yes, its been a long thread

condennnsa - 17-3-2013 at 01:10

congratulations, elementcollector, your success is encouraging to me I plan to try this again soon

Mailinmypocket - 17-3-2013 at 03:41

Well you son of a....!

Lol

very nice though, good job! Ill have to read over your last couple posts in detail when I have time, but this is very promising!

K-ongratulations!

blogfast25 - 17-3-2013 at 06:05

EC1: there you go; welcome to the weird world of K! (My little finger didn't lie after all!)

A few points.

Stirring: if your solvent was slowly simmering then a very slow speed of stirring should be sufficient. I never used magnetic stirring at all. Beware that molten K can react with Teflon (stirrer bars coating) according to:

2n K+ (CF2)n [Teflon] === > 2n KF + n C (potentially very energetically!) Check your stirrer bar for black/grey spots/streaks!

Yield: 50 % or more is easily achieved. A couple more runs and you’re probably there.

Disposal of K: reaction with alcohols like IPA or methylated spirits is much less vigorous than with water. I’m very surprised in your case the K caught fire with IPA, was the K still warm by any chance?

So be careful with admonitions like: "DO NOT USE ISOPROPANOL TO CLEAN POTASSIUM METAL": many of use use IPA for refreshing the surface of potassium balls, PRECISELY because it reacts with K only quite slowly.

Ampouling K: we’ll keep that for another day. For now, storing under clean oil is enough.

Heat: apply full heat right from t = 0.

Beware: from your 3rd photo on 16/03 I can see that your bottom water outlet tube was attached to the condenser in quite a dangerous way. If that hose flips off you've got cold water hitting a hot flask: thermal shock, possibly cracking of the flask and hot K to boot!!!!

[Edited on 17-3-2013 by blogfast25]

blogfast25 - 17-3-2013 at 06:24

And a quick word on yield.

Theoretically achievable yield is less than what most here calculate. Firstly, your KOH contains about 10 % water, so you need to subtract that from the charge to get the actual amount of pure KOH available for reaction.

Secondly, even if your reaction proceeded ‘100 %’, an amount of K that’s stoichiometrically equivalent to the amount of t-butanol used is ‘locked’ away as K t-butoxide at the end of the reaction. That’s a simple consequence of how this thing works. So if you used 0.1 mol t-butanol, 0.1 mol of K is ‘lost’ because it remains tied up as K t-BuO.

Any calculation of so called Actual Yield must take these realities into account.

m1tanker78 - 17-3-2013 at 07:34

EC1: That's the prize of stubbornness. Good job!

Note that you can use brake fluid to destroy remaining K bits in the reaction flask without so much drama. Add it then loosely cap the flask in case any of the bits float up. Swirl occasionally if needed.

Tank

blogfast25 - 17-3-2013 at 08:23

Tank:

I remember trying to coalesce K with DOT and that the K reacted with the brake fluid, albeit very slowly. Presumably K attacks the oxygen in the polyethers of the DOT. But these atoms are much more sluggish than the -OH groups in alcohols. Ethanol, propanol etc all react well with K/Na without much drama.

m1tanker78 - 17-3-2013 at 08:37

Blog: I do recall your experiment. EC1 claimed that his mixture caught fire when he added IPA. I suggested brake fluid as a low or no drama solution but you're right, alcohols will react much more vigorously. Methanol would probably be the best since absolute IPA and ethanol aren't always OTC.

To paraphrase EC, "little fiasco" = he wasn't prepared for the fire but luckily no harm done.

Tank

elementcollector1 - 17-3-2013 at 10:16

Quote: Originally posted by m1tanker78

To paraphrase EC, "little fiasco" = he wasn't prepared for the fire but luckily no harm done.

I think the K was still quite warm, as I remember having trouble getting the solution down to room temperature.

Oddly enough, the stirbar did react, but only in the middle, forming a yellowish-gray coating. Wiped off easily, so no worries there.

Anyway, I guess I'll try alcohol to destroy the bunch of little bits of K, but I'm going to do it outside this time.

What I had done was I got a metal strainer and a plastic container, and strained the mix through (here I first saw the odd, lumpy bits of K). I scooped these out into a beaker of fresh lamp oil, and some smaller bits of K had fallen through the strainer (without my knowledge). I poured isopropanol over the mix in a futile attempt to clean the strainer (too big to submerse in alcohol, but couldn't hold any of the liquid either), and several small bits of K rose to the surface, quickly reacting. And then they caught fire. I suppose it was a combination of the alcohol, kerosene, and warm K metal that did it, but still, the fire was for the most part contained. I took a metal spoon, and scooped the flaming plastic container into a larger metal one, setting that on the floor and placing the steel pot with the sand bath over it (this cut off oxygen, and extinguished the flames handily). Now, some bits of flaming material had also made it onto a plastic bag and pink towel on the floor, and these caught fire too. I stamped those fires out, and proceeded to throw the towel, strainer and bucket with the now-charred mix outside, where it was heavily raining. I thought I saw a tiny droplet of K leap out of the metal container, hit the ground, and produce a tiny lilac flame, but it was quickly extinguished. After this, I kept the garage door open to reduce the smell of burnt plastic, and focused on coalescing the K.

We'll be back with 'Misadventures in Chemistry' every Sunday morning, stay tuned!

I guess I'll use alcohol, but I'm probably going to pour off the lamp oil first, to avoid too much kerosene in the mix.

That brings me to an important note - when I was coalescing the K in fresh lamp oil, this yellowish fluff started to appear, turning the beaker cloudy. What could this be?

m1tanker78 - 17-3-2013 at 12:05

when I was coalescing the K in fresh lamp oil, this yellowish fluff started to appear, turning the beaker cloudy. What could this be?

Impurities in the lamp oil, most likely. I always get some sort of grunge when I coalesce sodium or NaK. What brand is your lamp oil? If you're in the U.S., the best lamp oil I've used so far is from Academy. Here is a link to the product.

Tank

Lambda-Eyde - 17-3-2013 at 12:39

Congratulations, ec1. I have run this reaction once and I didn't get any potassium either. I used paraffin oil as the solvent, Merck t-BuOH, old KOH (which was >85,5% a few decades ago) and a freshly opened bottle of Mg turnings. I performed the reaction in a 250 erlenmeyer and used a 300 mm vigreux as a condenser. A bubbler was attached to the vigreux. I can't say what the temperature was for sure, my thermocouple gave very different reading depending on the placement of the probe. It peaked at around 370 C (!) in some positions, but the sand bath seemed to be stable at 220-250 C. Whether or not my probe is faulty or if this is a typical reading I don't know. There was a very visible reaction between the KOH and the Mg when the reaction reached a certain temperature. The vigreux was lukewarm through the whole procedure. I can't remember how I added the catalyst. I will come back with a more detailed report when I get home to my notebook. I will try the experiment again as well.

blogfast25 - 17-3-2013 at 13:45

I poured isopropanol over the mix in a futile attempt to clean the strainer (too big to submerse in alcohol, but couldn't hold any of the liquid either), and several small bits of K rose to the surface, quickly reacting. And then they caught fire.

Part of your problem was that there wasn't much liquid about. When the K starts reacting with the alcohol, heat is generated, say ΔH. This heats up the liquid acc. ΔH = m c ΔT, with m mass of liquid, c heat capacity of liquid and ΔT the temperature rise. When m is small, ΔT is larger. Presumably you exceeded the flashpoint of the alcohol/kero mix. And higher ΔT also accelerates the reaction: runaway!

To avoid this in the future, slowly dump all your suspected K into a larger volume of alcohol, stirring continuously. The greater alcohol mass will then not heat up so much.

Quote: Originally posted by Lambda-Eyde

A bubbler was attached to the vigreux.

What do you mean by a 'bubbler' in this context?

[Edited on 17-3-2013 by blogfast25]

Lambda-Eyde - 17-3-2013 at 14:20

Quote: Originally posted by Lambda-Eyde

A bubbler was attached to the vigreux.

What do you mean by a 'bubbler' in this context?

It was an improvised bubbler, made by attaching a tubing adapter to the top of the vigreux, and leading a hose into a beaker of water, in order to keep additional air out. It was probably overkill, but I thought it couldn't hurt.

Sublimatus - 17-3-2013 at 15:53

One could always use a balloon to achieve the same effect, assuming the increase of gas volume inside the apparatus won't be too great.

elementcollector1 - 17-3-2013 at 16:05

I think we decided against balloons a while back. I used Al foil wrapped fairly tightly around the top of the condenser.

blogfast25 - 18-3-2013 at 06:26

The whole 'perforated balloon' idea rests on the assumption that some material (low boilers from the solvent?) escapes, thus keeping the balloon inflated and any air out. But if there really is such an updraft then it purges the apparatus of air anyway. And if there's not, this primitive septum is unlikely to keep anything out. It's a colourful waste of effort, in short.

The only way to work fully anaerobically is to fully purge the apparatus with O2 free argon or nitrogen and keep a low flux of either throughout the reaction. This has been shown conclusively not to be necessary.

[Edited on 18-3-2013 by blogfast25]

elementcollector1 - 18-3-2013 at 18:36

Just cut through a ball of potassium today to make it fit in a test tube. The metal was very silvery... for about 10 seconds. Now it's whitish and bluish. Also switched solvents from mineral oil to lamp oil, so now all my potassium spheres are more whitish than silverish in color. I'll probably just clean them when I'm coalescing stuff from the final demo...

Anyway, the extremely good news is, my coach ordered 100mL of tert-butanol from Sargent-Welch as an "official Science Club chemical", so coupled with a good source of magnesium and potassium hydroxide, if I wanted, I could either scale this synthesis way up or do it a lot. Or examine the properties of the tert-butanol itself.

I should get him to order chemicals through the club more often...

blogfast25 - 19-3-2013 at 06:28

Perhaps you could try and make some k t-butoxide from K and your t-butanol, and see how it reacts in 'pok's conditions' with Mg? That would be a novel contribution to this thread.

elementcollector1 - 19-3-2013 at 07:14

Perhaps you could try and make some k t-butoxide from K and your t-butanol, and see how it reacts in 'pok's conditions' with Mg? That would be a novel contribution to this thread.

Hmm. That might be something to do. Wouldn't I have to re-calculate the mass of K-t-butoxide to use? Should be a 1:1 mole ratio.
But wait - isn't there water needed to regenerate the catalyst from the Mg-t-butoxide? Or does the catalyst not need to be regenerated?

blogfast25 - 19-3-2013 at 09:52

There is no ‘regeneration of the catalyst’ in that case, because there's no catalyst. The reaction simply becomes:

2 K t-BuO(dissolved) + Mg(s) === > 2 K(l) + Mg(t-BuO)2(?)

With magnesium t-butoxide as one of the end products.

One of the embodiments of the patent deals with reducing an alkali metal t-butoxide with magnesium but I forget what alkali metal they used.

Edit: in fact it was potassium t-amylate they used:

http://www.freepatentsonline.com/4725311.html

Example 3.

[Edited on 19-3-2013 by blogfast25]

Lambda-Eyde - 19-3-2013 at 09:57

Perhaps you could try and make some k t-butoxide from K and your t-butanol, and see how it reacts in 'pok's conditions' with Mg? That would be a novel contribution to this thread.

This has been a plan of mine for a long time. Doing this reaction successfully would definitely prove that the intermediate is the K butoxide. Only problem is that I haven't had success with the reaction yet.

blogfast25 - 19-3-2013 at 10:20

Quote: Originally posted by Lambda-Eyde

What do you have in mind to synth the K alkoxide?

Lambda-Eyde - 19-3-2013 at 10:28

What do you have in mind to synth the K alkoxide?

The plan is to first succeed in the potassium synthesis, then quench the potassium with t-BuOH in parafin (or any other suitable solvent), and then run the reaction with the solution (or dispersion) in order to minimize the degradation of the butoxide.

MrHomeScientist - 26-3-2013 at 10:30

Over a year since I first posted on this topic expressing my interest, I am finally nearly ready to attempt this experiment myself! It took quite a while to get the right equipment and chemicals (particularly a good hot plate and the t-butanol).

I have spent the last few days reading the entire thread (yes, the WHOLE thing!) to review the progress of others. I copied the bits and pieces I found useful and put them in a Word document, which I've attached here in case anyone else might find it helpful. It's quite long itself, but at least shorter than 51 pages! Don't be offended if your posts didn't make the cut, or if I misattributed some peoples' ideas - I was simply looking for information relevant to my own plans for this experiment and this was just a quick notepad for myself. I just figured others might benefit from my tired eyes

There's lots more great information here, and the document is really no substitute for reading the thread in its entirety.

Attachment: Potassium Synthesis Notes.docx (325kB)
This file has been downloaded 1216 times

==========================

Anyways what I really wanted to report is that for my solvent, I purchased some Lamplight Ultra-Pure Lamp Oil (MSDS) from a local arts and crafts store. They were unfortunately out of the colorless variety, so instead I bought one dyed blue. I figured the dye shouldn't interfere, and there wouldn't be much in there anyway.

First thing's first: I assembled my setup to test the hot plate. I don't have a thermometer that goes high enough, so I wanted to test if I was able to boil the solvent (which would give me an idea if my hotplate is suitable). A thermocouple is in the mail, but we'll see when it actually gets here. Apparatus: a sand bath covering a 24/40 100mL RBF containing 50mL of the oil, and connected to a liebig condenser (with nothing connected to it, just cooled by ambient air). I slowly brought the temperature up to 9/10 on the knob, and observed that mirage-like effect of convection in the liquid. It never boiled, but handily enough the blue color slowly faded! By the time I took it off heat, it was completely colorless. This was about 3 hours later - unless I'm doing something wrong, sand baths take forever to heat up!
I also did another trial with the exact same setup, except adding carbon boiling chips. This time a very slow boil was observed (about 1.5 bubbles / sec). Condensation was observed at the very bottom of the condenser (not even up to the jacket). This means that my plate should be able to reach the necessary temperature. This batch had a slight yellow tinge to it, probably from carbon particles. Hopefully those will have settled next time I check on it. Carbon shouldn't interfere, but I'll at least run it through a filter.

A question on the sand bath, though, as I've never used one before. Should temperature be slowly turned up, or is it alright to go straight to max setting? I've seen some hot plates warn against using sand baths on them - what's the reasoning on that, and is there anything I should do to mitigate the risks?

Hopefully I'll be set to try this out this weekend. I'll use the decolorized lamp oil in my first trials.

[Edited on 3-26-2013 by MrHomeScientist]

elementcollector1 - 26-3-2013 at 11:23

Not sure, as my hot plate could take both ways. Turning it all the way up seems like the best idea, just to save time. Then again, it was a Corning hotplate with stirring.
What's your sand bath? I used a stainless pot full of beach sand, and it worked pretty well.
Your dye was probably destroyed by heat - interesting...
What's your hot plate, anyway?

MrHomeScientist - 26-3-2013 at 12:09

My hot plate was a recently-bought eBay auction, a "Thermolyne Nuova II Hot Plate Stirrer Model SP18425." It was some type of estate sale, you could tell they didn't really know what it was

I use the same setup for my sand bath, except my sand was taken from the excess bricklaying material from a nearby house under construction. Interestingly, after taking my flask out of it I saw a lot of glittery flakes adhering to the glass that reminded me of mica. Might be a component of mortar.

By the way I forgot to mention: Nurdrage has apparently taken down all his potassium synthesis videos (both the synthesis and coalescence experiments). Why would he do this? I was hoping to use them to help me prepare.

Mailinmypocket - 26-3-2013 at 12:51

I think sandbaths can confuse the hotplate's thermostat (if it has one) and cause it to overheat and burn out. The manual for my hotplate says to not heat metal vessels directly on the "pyroceram" surface, nor to heat sandbaths. I'm not the expert on hotplate's though so maybe other members could confirm if this is why sandbaths are not recommended. A cheap kitchen hotplate with coil element might be a safer option if in doubt...

blogfast25 - 26-3-2013 at 13:04

Nice idea to summarise this long thread, Mr HS, long overdue IMHO...

And get that thermocouple in there: w/o it you're driving blindfolded, so to speak. Temperature is too important a factor here to leave it to chance.

[Edited on 26-3-2013 by blogfast25]

elementcollector1 - 26-3-2013 at 15:00

My hot plate was a recently-bought eBay auction, a "Thermolyne Nuova II Hot Plate Stirrer Model SP18425." It was some type of estate sale, you could tell they didn't really know what it was

I really don't know why he took down the video. K3wls? Noobs? Apparently his flash powder is down too, so that's what I'd expect.
Mica, huh? Interesting. I used to be a mineral collector.
The coiled hotplate is a no-no, from my experience. It contributed to 6/7 of my failures. Use something stronger.

blogfast25 - 27-3-2013 at 05:01

Nurdrage is moving jobs and will not have the use of that laboratory his current employer kindly allowed him to use for the Nurdrage channel. He's now bidding for donor money to try and establish something independent. All this might have something to do with it.

MrHomeScientist - 27-3-2013 at 05:31

But then why only take down some videos, and not all? Also his Instructibles page on potassium remains up, but the video link in it is broken. Very strange to me.

To come somewhat back on topic, I tried filtering my second batch of oil but the light yellow color remains. Probably miniscule carbon particles my paper can't catch. Also, interestingly I seem to have "reodorized" my deodorized lamp oil by heating it...

blogfast25 - 27-3-2013 at 05:38

To come somewhat back on topic, I tried filtering my second batch of oil but the light yellow color remains. Probably miniscule carbon particles my paper can't catch. Also, interestingly I seem to have "reodorized" my deodorized lamp oil by heating it...

Just try it: if anything is to be learned from the whole thread it's that the reduction is fairly robust with respect to the solvent used.

elementcollector1 - 27-3-2013 at 08:49

Blogfast succeeded with lemon-flavored oil; I think a few particles of carbon won't do much to your reaction.
(Seriously, that's going to be an eternal joke for me. Lemon-flavored potassium.)

MrHomeScientist - 27-3-2013 at 08:53

I know, I'm really not concerned about the stuff in the oil. I just like to eliminate potential uncertainties whenever I can, especially because there's been so many failures on this experiment. Really, using the blue oil straight out of the bottle probably won't hurt anything. My next concern is that my oil has a listed boiling point of 250+ C, so there may not be much agitation going on at the reaction temperature. I'll occasionally swirl the flask if this is the case. Thermocouple is in the mail, but it probably won't make it here by the weekend.

blogfast25 - 27-3-2013 at 09:25

I'll occasionally swirl the flask if this is the case. Thermocouple is in the mail, but it probably won't make it here by the weekend.

Occasionally swirling should be just fine. Remember that if the reaction proceeds, hydrogen will evolve, providing also a modicum of turbulence.

First Attempt: Failure!

MrHomeScientist - 30-3-2013 at 17:01

My thermocouple arrived much sooner than expected, so I was able to try this out today. I somewhat followed the procedure on Woelen's page, with the exception of adding the alcohol all in at the beginning (the so-called "one pot").

In a 100mL RBF, the following were mixed:
- 7.0g KOH flakes
- 3.5g Mg granular powder (~80 mesh)
- 0.8mL tert-butanol
- 50mL deodorized and decolorized lamp oil (see my previous post about the oil pre-treatment)

You can see the Mg is a fairly coarse powder. This was connected to a condenser and placed in a sand bath. I put a layer of aluminum foil on top of the condenser with two small holes poked in it, to allow for hydrogen escape and to attempt to exclude air. The thermocouple was stuck in the sand near to the flask, so the temperature recorded is probably not the reaction temperature.

Shortly after the thermocouple read 100 C, about 20 minutes into heating, copious amounts of white smoke bubbled out of the liquid. Much of this blew out the holes in the top of the condenser, and you can see that in the picture. This completely vanished 2 minutes later.

After the smoke dissipated, the solution gently bubbled throughout the remainder of the run. The solid contents initially appeared to "fluff up" quite a bit, likely to be dehydrated KOH. About 45 minutes in, the temperature read around 200 C, and I fiddled with the dial to keep it there for the rest of the run. (The max temperature I saw at any point was about 225 C.) I swirled the flask once, gently, around this time.

An hour and a half in, a white crust started appearing around the sides of the flask and as a little island in the middle. I swirled again to break this up, and it never reappeared.

After a total of 4 hours since reaching 200 C, I stopped the experiment and removed the flask from heat. The final solution has turned yellow, with no sign of potassium metal. Throughout the whole run I never once saw any hint of metal formation.

Possible Sources of Error
- Mg fineness. Next run, I'll try a mix of 20% powder and 80% turnings as suggested to work by len1. The turnings will be freshly drilled from a firestarter block.

- Too low temperature. My thermocouple probe is bare wire, so I can't put it straight in the reaction mix. It may be that the temperature I see is greater than that in the flask, so ramping it up so it reads, say, 250 C might be a solution.

Any thoughts from anyone? I'll try one or both of these ideas tomorrow morning, and see how it goes.

m1tanker78 - 30-3-2013 at 18:31

Forgive me if this has been mentioned or done before:

Why not add KOH + solvent to the flask and heat for a period of time, cool, add Mg + catalyst, ramp up to reaction temp?? The idea is to drive off excess moisture from the KOH before the reaction is run. Maybe this is why Mg turnings seem to produce better results compared to powder. Less surface area to initially react with moisture...

One other thing, could there be a product from a side reaction that MUST be vented off in order to keep it from reacting with K? That could explain why Blogfast and others(?) had success with a glass pipe and wet paper towel while others(?) who use fancy schmancy condensers fail sometimes in spite of careful prep, etc. Perhaps something is fouling the reaction or newly formed K because it condenses and returns to the pot??

Tank

MrHomeScientist - 30-3-2013 at 19:59

I believe the consensus was that powder is beneficial to the dehydration step, and turnings are better for the potassium formation. Dehydration isn't just driving the water off by heating - I don't think you can do that to KOH without very high temperature. Rather, its a reaction with magnesium: Mg + H2O --> MgO + H2

On Woelen's page he mentions airing out the apparatus after the dehydration step, to eliminate any remaining water vapor. But, he uses the "staggered addition" approach and I think in my case that would lead to loss of much of the alcohol as vapor. Plus, I believe the initial reaction and reflux serves to saturate the air above the flask with other vapors, which mostly keeps out air.

blogfast25 - 31-3-2013 at 05:33

Tank:

There are no solid reasons to believe that one total reflux system is to be preferred over another. Water cooled (mantle, like Liebig or Allihn) could potentially lead to freezing the catalyst onto the cooling surface, as observed by elementcollector1.

Mr HS:

I wouldn’t go so far as to call it a ‘consensus’: there was len1’s spectacular failure with shavings, as reported on a pre-pok thread, for instance.

Again, I don’t really see the point in ‘airing’ the set up. Water is supposed to be scavenged chemically by that excess Mg we add. But as I’ve postulated, very small amounts of water may be needed to regenerate the catalyst.

elementcollector1 - 31-3-2013 at 07:31

Only outside hose water led to freezing, the water form the fume hood was fine.

HOLD THE PHONE: It did work!

MrHomeScientist - 31-3-2013 at 12:01

Like elementcollector1 did some time ago, it looks like I spoke too soon! My first run on this experiment did indeed produce elemental potassium! I just didn't notice because the oil had gone opaque, and as it turns out the metal barely did any coalescing. Take a look at the sand leftover from the reaction:

Zoom in: all those bright pieces are thousands of tiny potassium spheres! I am astounded by how many there are. I managed to salvage most of the larger balls, the largest of which is only about 4 mm across.

So it appears my problem is coalescence. What has been people's experience getting these things to come together? I read the coalescence thread, but nothing really jumped out at me as being particularly good. If I remember right, the method Nurdrage used was to heat the salvaged balls to melting, then using the magnetic stirrer to push them together (taking advantage of paramagnetism). That video is gone, of course, so I can't review it. Even so, that method is more for easily recoverable balls, not my potassium sand. Any ideas?

Possible Sources of Problems
- Flask geometry. Since I am using a round-bottom flask, there's a lot of space in the solvent for potassium to float around without touching each other. The narrow neck of an erlenmeyer might be beneficial here, constricting the balls to run into each other more often.
- Not enough agitation. Since my oil boils at ~250C, there was only a gentle boil. Perhaps a lower boiling solvent or higher temperatures (see below) will cause more agitation and allow things to come together neater.

==============================

In the meantime, I'm running Trial #2 to test if higher temperature will help. This time I'm keeping things at 250 C instead of 200. Most of the way through and it's looking much like Trial #1, but with a more steady boil.

Interestingly, this time the white smoke actually solidified in the condenser:

After a second burst of smoke, this condensate melted and now looks like a small amount of liquid on the condenser walls. This is pretty convincing to me that this smoke is mostly t-butanol, and it looks like I might be losing a good amount of my catalyst as it escapes.

elementcollector1 - 31-3-2013 at 12:07

That smoke would be your alcohol. Too powerful cooling, is my guess.
I had heavy stirring; this led to a large, very irregular lump of K in the center (not at all spherical).
Be sure to store this potassium well, as it annoyed me to no end trying to store this stuff under oil and kerosene. I'll have to try ampouling when I next get a chance.
To coalesce these later, simply heat to the melting point of potassium, and swirl. Repeat as necessary.
An additional note: NurdRage's video on the subject was misleading. The potassium does not stir itself, even when molten (no instance of paramagnetism). I wonder if it's my hotplate or my potassium, though...

blogfast25 - 31-3-2013 at 13:53

Nice work, Mr HS. I had a feeling yours would work.

Coalescence is the thing that takes up the most time in this experiment. Flask geometry may have a little effect but from my experiments its influence is much smaller than you might expect. Most of us used flat bottom reactors too.

Coalescence is essentially a game of chance, as len1 once put it. Most collisions between droplets of molten potassium do not result in union because there’s a layer of oil separating them. Once in a while two drops merge. What you see in your first photo is of course in itself the result of even smaller droplets having united into larger ones. Keep going and going and with swirling of the flask gradually even larger balls of K will result.

It can help a lot to separate the larger ones from the sand by straining it (COLD, of course!), then continuing the coalescence more or less in the absence of MgO sand, at about 100 C. But you’ll lose many fines (K) in the process.

So the crucial question is: what was your total run time?

MrHomeScientist - 31-3-2013 at 19:09

blogfast,

Thanks a lot for the vote of confidence! I had thought straining it might be my best bet, too. Many of those spheres are just miniscule, though, so I would definitely lose a lot. Personally I think a somewhat lower boiling solvent will help, as it provides some more agitation. I'll investigate the results of Trial 2 tomorrow afternoon and see if the higher running temperature (and thus more vigorous bubbling) helped at all.

Trial #1 was run for 3h 45min after reaching the running temperature of 200C (5 hours if you count the initial warm-up), and Trial #2 ended 3 hours after the running temperature of 250C was reached (4 hours counting warm up). #2 ended a bit early because I had to leave for Easter dinner.

So both were a bit shorter than recommended, but I'd have thought I'd get bigger spheres than I did - Woelen had 3 huge balls 3.5 hours in to the one on his page. It's also possible my actual temperature is a good bit lower than my measured temperature, since I can't measure the mix directly.

blogfast25 - 1-4-2013 at 06:11

Mr HS:

Yes, with about 4 h run time, most of us get better coalescence and bigger globules than you achieved. Like you wrote, perhaps a slightly lower boiling solvent (250 C really is borderline) could help with the coalescence because of slightly improved agitation (boiling).

Several of us worked on the problem of coalescence but nothing very effective was unearthed. Perhaps Nurdrage’s use of Tetralin, in which molten potassium floats, is the one thing that seemed to cut down run time considerably: he reported one experiment that was over and done with in about 1 h, with good yield and coalescence. But Tetralin is expensive and hard to get for non-institutional chemists.

It could also be interesting to play with the clean K balls in clean solvent, to try and merge them into larger units: you'll experience just what a slow process it is!

Question: when and how did you add the catalyst?

[Edited on 1-4-2013 by blogfast25]

MrHomeScientist - 1-4-2013 at 10:27

The catalyst was added in the very beginning along with all the rest of the chemicals, before the apparatus was assembled, ala your "one-pot" method. I'm fairly certain the plumes of white smoke I observed around 100C were mostly the alcohol boiling off and recondensing. I'm not sure if this is detrimental in any way.

One thing I thought of, beyond straining my mix of potassium and MgO/Mg sand, is to separate via density. Magnesium oxide (and to a lesser extent, unreacted magnesium) is much more dense than potassium, so there must be an appropriate solvent in which K floats and MgO doesn't. The picture I have in my head is having this solvent with a layer of mineral oil on top (which my K sinks in), and pouring the sand mix into that. The K should sit at the interface, remaining protected from air, while the rest sinks to the bottom. Dioxane seemed a good choice, but synthesizing it totally dry would be difficult and I don't really have a good source for purchasing solvents.

blogfast25 - 1-4-2013 at 13:07

So further evidence that one pot works. I remember seeing such a mist too, it could be H2 dragging some stuff with it. It's nothing remarkable.

Re. floatation, unfortunately molten K has a density that's very close to most hydrocarbon solvents. To make them heavier you need to introduce 'exo atoms' in the chains or branches but most exo atoms render the solvent reactive towards K! A conundrum... Dioxane does offer a way out (tested and tried) but like you say, it's not easy to get, and hard to home make completely anhydrous. I bought some glycol for that purpose but never got round to trying to synth it.

MrHomeScientist - 1-4-2013 at 19:04

Hm. I guess if I can find a fine strainer I'll try that, or just put the whole mix back in the sand bath for more time. I'll look around for a lower boiling oil as well.

I took a look at the products of Trial #2, and they are much the same as #1 - lots of tiny K balls mixed in with fine Mg/MgO sand. I'd hoped this one would yield larger spheres because of increased agitation by heating closer to the boiling point, but maybe the somewhat shortened run time counteracted that. I also took one of the dirtier balls and dropped it in some water, resulting in the beautiful lilac flame of K. First time I'd seen that with my own eyes, so it was very satisfying!

blogfast25 - 2-4-2013 at 10:09

Careful with the disposal of fines! I once had a bit of 'bang!' when putting spent MgO down the drain. I had rinsed the MgO with alcohol first to react away the K but a bit must have been left behind. There was no damage to anything but it was quite startling!

MrHomeScientist - 2-4-2013 at 10:35

I actually had a similar experience. I had been rinsing my larger balls in some mineral oil before transfering them to my storage bottle, and this rinse oil had a lot of K-fines in it. To destroy them, I added a few drops of isopropyl alcohol (rubbing alcohol) and swirled. It made a crackling sound, and I had a bit of a "flare-up" when some of the alcohol on the surface ignited. It was just a flash and disappeared quickly, but was definitely startling I agree! It worries me greatly how I going to dispose of what I can't coalesce. Woelen uses t-butanol in his disposal, but I do not want to use up my hard-to-obtain catalyst.

[Edited on 4-2-2013 by MrHomeScientist]

blogfast25 - 2-4-2013 at 12:20

I actually had a similar experience. I had been rinsing my larger balls in some mineral oil before transfering them to my storage bottle, and this rinse oil had a lot of K-fines in it. To destroy them, I added a few drops of isopropyl alcohol (rubbing alcohol) and swirled. It made a crackling sound, and I had a bit of a "flare-up" when some of the alcohol on the surface ignited. It was just a flash and disappeared quickly, but was definitely startling I agree! It worries me greatly how I going to dispose of what I can't coalesce. Woelen uses t-butanol in his disposal, but I do not want to use up my hard-to-obtain catalyst.

[Edited on 4-2-2013 by MrHomeScientist]

Use lots of IPA or methylated spirits: lots of it so the alcohol won't heat up too much and won't ignite. Stir well to keep temperature homogeneous. Allow time for reaction to complete.

[Edited on 2-4-2013 by blogfast25]

elementcollector1 - 12-4-2013 at 13:22

Running a one-pot reaction as we speak, we'll see if any K results.
Any tips on removing solvent before ampouling? It's been a nasty problem for my sodium...

elementcollector1 - 12-4-2013 at 16:07

Reaction: Success
Extraction: Partial success
I waited until the stuff was slightly warm, then strained out the potassium. Surprise! Even at that temperature, it caught fire. I managed to salvage a few tiny bits of potassium, which will coalesce into a sphere about a quarter-inch big by my estimate, but nowhere near my actual yield.

Fortunately, I was expecting it this time around, so nothing was hurt.

Reaction time seems to be significantly faster with the one-pot method - the reaction was done in just 2 hours!

blogfast25 - 13-4-2013 at 04:54

Reaction: Success
Extraction: Partial success
I waited until the stuff was slightly warm, then strained out the potassium. Surprise! Even at that temperature, it caught fire. I managed to salvage a few tiny bits of potassium, which will coalesce into a sphere about a quarter-inch big by my estimate, but nowhere near my actual yield.

Fortunately, I was expecting it this time around, so nothing was hurt.

Reaction time seems to be significantly faster with the one-pot method - the reaction was done in just 2 hours!

Fire again? Strange. Still glad to see you got K!

If one-pot did seem shorter (I didn't conclusively observe that), it may simply be that adding the catalyst from the start cuts down time by avoiding that silly 'timed, step by step' addition.

elementcollector1 - 13-4-2013 at 15:35

You know, I always see my reaction mix turn chocolate brown. Although it's a good indicator that the reaction took place, I have to wonder whether it's K2O and why it's there: Is my potassium getting consumed again somehow? I don't see anyone else having this problem.

elementcollector1 - 13-4-2013 at 18:33

Well, this one worked! I got a sphere about 1/2" in diameter, and nothing else (everything coalesced, apparently). Unfortunately, the sphere was so large that it wouldn't fit through the neck of the RBF, and I was forced to mutilate it...

Oh well, I'll just coalesce them again tomorrow.
Pictures may follow, once I get the camera up and running.

blogfast25 - 14-4-2013 at 04:33

You know, I always see my reaction mix turn chocolate brown. Although it's a good indicator that the reaction took place, I have to wonder whether it's K2O and why it's there: Is my potassium getting consumed again somehow? I don't see anyone else having this problem.

I’ve had browning too, up to the point where the potassium is hard to see. I think it may be due to a bit of Diels-Alder on the solvent but that is merely an idea of mine right now.

Sounds like you’re in business now. Pix would be nice.

elementcollector1 - 14-4-2013 at 21:18

It doesn't look nearly as pretty as the last bunch, but I digress: I'm probably just going to ampoule it anyway.

MrHomeScientist - 15-4-2013 at 11:19

@elementcollector1: I've also seen browning of the solvent, which you can see in one of my posts a few pages back. It turns out all the color is just a suspension of extremely fine particles - upon standing overnight, it all settles out and leaves perfectly clear oil. Looks like you're having great results now! So yours finished reacting and coalesced in only 2 hours? What's the boiling point of the oil you use?

@AJKOER: This whole nascent hydrogen business is unnecessarily complex. You're introducing new things that aren't required for this reaction to run (as far as we understand it). What you also aren't understanding is that len1 and blogfast have essentially proven that MgO is the final product by measuring gas outflow. If you don't want to scroll all the way back to their testing, I believe I included it in my word document summary on page 51. Again, there's no need to postulate extra things. Why don't you do some experimenting yourself?

elementcollector1 - 15-4-2013 at 11:55