[Taaie-Neuskoek]

Ok, I played around, and let you know my findings...

I've cut off the oxide layer of my potassium chunks, after I tried to lit a piece of potassiu, to see how it would burn. It didn't. I did heat it with a propane torch, but all happend was a controlled oxidation to KOH.

So witha new scalpel I went to cut, it went very smoothly, there was a small spark at one one point which freaked me out, but for the rest it went very ok. I stored the clean chunks in oil, and hope they stay ok there. I had a lot of fun with trowing the pieces into water, I did a fair amount during the day, but kept a bit for the night... Actually a pity I did trow so much in during the day, at night it was absolutely beautifull!! I wrapped the peices in a tissue (not more than ~2grams I think) and they were trown into the water, first is a big puple flame, than a pooof, and then there are 20 purple small flames swimming over the water, sometime a crack here or there... absolutely cool! A friend of mine catched in on tape, I hope to post the video here when it's ready.
I also did a small chunk into water for my family, a piece of 1,5x0,5x0,3cm was thrown into a plastic bucket with some water, but it reacted the same as sodium, a bit fizzing and a puple flame.

From my experiences are the stories about potassium with big explosions are a bit overexaggerated, if the pieces are a bit cut to small chunks they can be safely thrown into water.

[jdurg]

From my experiences are the stories about potassium with big explosions are a bit overexaggerated, if the pieces are a bit cut to small chunks they can be safely thrown into water.

That is the key sentence right there.  If cut into small chunks, yes they can be safely chucked into water.  The problem is, most 'kids' see pieces of sodium metal being thrown into water and after they see something the size of your thumb thrown into water as sodium, they do the same thing with potassium.  The problem is, potassium reacts much faster than Na so they don't get away in time and the caustic reaction mixture flies onto them.  With the Na I threw into the water, there was a good time delay before it started to go off.  Plenty of time to step back and make sure you were far enough away.  Then it suddenly went 'BOOM!' and any unreacted sodium went flying up into the air.  With potassium, you don't get nearly the same amount of time before it goes 'BOOM!'

A good control for the explosion size is the water itself.  If you have shallow water, the explosion will be MUCH greater as you won't have as much of a cooling effect from the massive amount of water which can slow down the reaction.  In shallow water, the forming gas can quickly find some oxygen and ignite.  When I went up to the lake, I took a good ten grams or so of sodium and wrapped it in some kleenex and attached it to a rock.  I then threw the rock into the center of the lake.  I got maybe one or two bubbles to show up at the surface, but that's it.  There was so much water on top of the sodium that it could not ignite the evolving hydrogen gas and instead just fizzed away at the bottom of the lake.  Kind of dissapointing really.

Probably the best way to show how the water itself controls the explosion is with cesium and rubidium.  Both Cs and Rb are denser than water, so they will sink to the bottom of the reaction vessel.  If you have a deep source of water, it's not going to be able to react as well.  If it's a shallow pan of water, then it will explode violently as the gas and metal are exposed to atmospheric air.

[woelen]

Jdurg, I just received the following two samples. A very nice sample of lithium (for just $6 from a german guy, who sells lots of chems to chemistry hobbyists):

http://woelen.scheikunde.net/science/chem/compounds/lithium.html

From that same guy I also received some sodium metal (appr. one ounce), but this metal looks very strange. It is covered by a pink layer and it is immersed in oil, which contains a lot of turbid yellow/brown stuff. I'm wondering what that pink stuff is. Any idea?

http://woelen.scheikunde.net/science/chem/pics/natrium.jpg

[Borek]

Pink... If I recall correctly pink layer of (per)oxide is created on potassium - but I can be completely wrong.

[Borek]

When I went up to the lake, I took a good ten grams or so of sodium and wrapped it in some kleenex and attached it to a rock.  I then threw the rock into the center of the lake.  I got maybe one or two bubbles to show up at the surface, but that's it.  There was so much water on top of the sodium that it could not ignite the evolving hydrogen gas and instead just fizzed away at the bottom of the lake.  Kind of dissapointing really.

Perhaps the problem was that evolving hydrogen bubbles out from the vicinity of sodium, and there is no oxygen to mix with thus no matter how hot sodium gets - there is nothing flammable to ignite in vicinity.

[jdurg]

Very cool woelen.  I too recently updated my lithium sample to a much bigger brick of the stuff.  I'll need to upload the photos eventually when I get some time.  Amazing how the Li floats on the oil, isn't it?

For the pink on your sodium, it makes me think that it was stored in something similar to a kerosene.  After storing Na in kerosene, it will start to take on a reddish hue when it's removed.  That hue just won't go away either.  Sodium peroxide is VERY difficult to form under atmospheric conditions.  You typically need an incredibly high pressure, or a very rich oxygen atmosphere to get peroxide to form over the oxide.  Even then, the sodium peroxides aren't reddish in color.  They're actually pretty colorless.

For the Na in the lake Borek, that's the assumption that I made as well.  I was hoping that the water logged paper towels would break and the H2 gas and molten sodium metal would rise to the top and go KABOOM!

[woelen]

Na stored in kerosene ?? Hmm, that's a strange story. I'll inform at the eBay seller from whom I bought the sodium metal. Do you think it does harm? I want to use this sodium, but I'm not feeling like scrubbing the surface of these sodium samples to get rid of that pink crap.

I removed the dirty oil, rinsed the sodium a few times with low-boiling ligroin. The pink stuff sadly enough did not dissolve in the ligroin. Now I put the sodium in clear and colorless mineral oil. It looks better now, but still I have the pink crap around it.

The Li-sample indeed is very cool. The sample I have has a weight of approximately 3 grams and that is sufficiently large for me. The picture shows nice details and it nicely shows how it floats in the oil.

This afternoon, I also made my own almost waterfree bromine sample from KBrO3, NaBr and conc. HCl (the KBrO3 I have made with electrolysis, see thread on SFN). Have a look at that:

http://woelen.scheikunde.net/science/chem/compounds/index2.html

It is really nice. Now it is still in the little vial, but I quickly have to find a way to find a more permanent storage. The cap of the vial will be eaten away within a few weeks. It, however, perfectly closes the vial, no smell of bromine at all. Part of it I used in an experiment with Al-foil. Really impressive to see that the metal catches fire in the bromine and continues burning for a while!

It is fun to make your own elements from cheap chemicals or from chems you have no application for (such as the KBrO3, which simply is too extreme for pyro-experiments).

 

[jdurg]

Your last paragraph says it best.  Of all my elements, the chlorine sample I have is my absolute favorite because I made it myself.  I didn't have to buy it, only the VERY cheap chemicals needed to produce it.  There's just a great feeling of accomplishment when I look at it and say 'Wow.  I made that myself!"  

For bromine, it's very easy to make it relatively pure as the stuff is so incredibly volatile.  You can slowly heat the bromine water up and the bromine will escape from the water very quickly.  Just direct it into a vial that is sufficiently chilled and the bromine will collect in there.  For storage, you are correct on your website.  There is no way to store the stuff WITHOUT it destroying the container.  In labs, they have it in amber glass bottles with tightly sealed caps, but that bottle is also stored in a metal can which is stored in a vented cabinet.  Over time, it will eat through almost anything.  There are two ways to permanently store the bromine.

The first way is to seal it in a glass ampoule like I have done with mine.  It will never escape the ampoule unless you break the glass, and you can also store it easily and see the red-brown vapor.  (I took some white cardboard and created a container for my fluorine, chlorine, bromine, and iodine ampoules.  It's REALLY neat to see the colorless fluorine tube, the pale green chlorine tube, the intense red-orange bromine tube, and the faint violet iodine tube.  The fluorine tube is colorless because pure fluorine is barely visible anyway, and the sample I have is only 33% fluorine.  The rest is helium which dilutes the fluorine and prevents it from eating the glass).  

The other method of storage is to store the bromine in a freezer.  Bromine's melting point is 266 Kelvin, so a cold freezer will cause the bromine to solidify.  When solid, it's much easier to keep confined.  Just keep it in a tightly stoppered glass vial, and then put the glass vial in yet another glass vial for safety and store it in a freezer.

The aluminum/bromine reaction is incredible.  Especially because of the delay before it really gets going.  It takes some time for the aluminum oxide to go bye-bye, but once it does the reaction proceed vigorously and fumes are spread all over the place.