Chemical Forums
Specialty Chemistry Forums => Citizen Chemist => Topic started by: woelen on July 25, 2005, 03:36:49 PM
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Just a few days ago, I received a really beautiful sample of metallic sodium (it contains approximately 5 grams). It is stored in crystal clear mineral oil. A picture of the sample follows here:
http://81.207.88.128/science/chem/compounds/sodium.html (http://81.207.88.128/science/chem/compounds/sodium.html)
Now, I am wondering, whether this sample remains so nice, or will it turn dark and look oxidized within a few weeks. If so, can I do something to prevent this sample from deteriorating? Any suggestions are welco
edit by woelen: Actually, the sample is more than 5 grams, it is over 7 grams. It was sold to me as 5 gram-sample, but propbably the seller just wanted to have a full bottle, such that it looks better.
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Jdurg is, of course, the expert here, but from my experience I have found that mineral oil is really not sufficient for keeping sodium metal nice and pretty. I keep a small amount on my lab bench in an identical vial. I have found that the stuff tends to only last a few months before it starts to turn white (NaOH). I would say that if you really want to keep it nice you would need to keep it in a sealed ampoule. You could probably extend the life in that vial by trying to seal it with Parafilm or maybe even wax to prevent oxygen and water from getting in. Be careful though, those caps tend to corrode over time too.
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Sadly woelen, over time it will corrode and within a few weeks you'll see a heavier layer of oxidation on the surface. The cap is not airtight and oxygen dissolves to a certain extent in the oil. As a result, as the oxygen is depleted from the oil, more will dissolve inside of it. This is why the oxidation still happens, but at a reduced rate. The only way to get truly unoxidized sodium metal, permanently, is to take the fresh metal and seal it in an argon filled ampoule. If you look at the periodic table here on this site, you'll see my fresh sodium there which is sealed in a glass ampoule. The stuff, I swear to god, is brighter than fresh silver metal. Quite amazing.
(Also, isn't it remarkable how large a small sample of the stuff is? Now you can see why I faced a dilemma when I foolishly bought two OUNCES of the stuff. I've had fun getting rid of that excess sodium metal, however. With all of the humidity and heat out there today, I decided to take some of my sodium metal and put it on some ceramic bricks outside in the hot sun. Within about an hour, it caught fire and burned with the strong orange/yellow color of sodium).
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Thanks for your reply, both of you. Your answers are clear and I'll see what I can do to keep this sample nice and shiny. I'll store it in another tightly capped bottle, just a little larger than the one, shown on the site. At least, that creates another barrier for oxygen and water vapor to come in. If I want to show it to someone, then I'll take it out, realizing that every time I show it, a little new oxygen ccan enter the larger bottle...
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As I'm sure you already know, sodium is pretty docile in terms of long term storage and stability. Under most normal conditions, all you'll get on it is a light coating of sodium oxide/hydroxide which will give it a dull gray coating. It never really oxidizes and corrodes to the point where it looks like a white mass of crystals. Potassium, if you'll pardon my french, is royal b$*%( to store. >:( When I got my first sample of potassium, it was a small little button that was oxide free when I got it, and over the course of about a year it became this white mass as the oxide coating on it puffed out like cotton. I then upgraded my potassium into a much larger, 1"x1"x0.5" cube of it. This was in September of last year. One side of it was freshly cut when I purchased it, and for a good month or two it remained bright and shiny under the mineral oil. Then the fresh suface took on a violet hue as the oxidation slowly began. Finally, the current state of the metal is a dull white color with a few pockets of bright metal showing through. What's REALLY irking me, however, is the formation of some superoxide on the surface. I see a few areas where a dark orange/brown coloring has begun. That just really pisses me off because now I'll have to carefully excise the peroxide without igniting the rest of the metal. I'll need to find a new airtight container to keep it in, as my current solution just isn't working. :(
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A little bit late for a reply, but I just returned from my holidays... :)
The sodium I have still is as nice and shiny as when I received it. I'm also considering the purchase of a small sample of potassium metal, but the stories about peroxide-forming and the risks of explosion scare me away from purchasing potassium metal. I also would like to have some for experimenting (just as I have for sodium, 1 bottle as collector's item, the other for experiments). If I read about the troubles with the storage of the potassium, then I think that the alkali metals will be represented by a single member, being the sodium metal. Maybe, I buy a compelely sealed sample in argon gas, but unfortunately, these are much more expensive.
I also purchased 75 grams of calcium metal and I received it today. The metal is in a completely sealed metal tin, which, once opened, cannot be closed again. I still left the container closed and before I open up the container, I would like to know what is the best method of storing the calcium metal.
Several years ago, I also had a chunk of calcium metal, which I stored in a glass bottle. A few months later, I only found a white powder of calcium oxide and carbonate, so this should not happen again. Must calcium metal be stored under mineral oil as well or can I store it under ligroin with a low boiling point? Any advice on this would be very welcome.
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Calcium is very similar to sodium in terms of storage. The best thing you can do is keep it stored under mineral oil in a screw top jar. Over time, it will oxidize and corrode at a similar rate to sodium. Like sodium, it will happen very slowly, however. That's the REALLY frustrating thing with metals like Lithium, Sodium, Potassium, Calcium, Lanthanum, etc. They start out just fine and look like they'll stay stable for quite some time, but then one day you'll suddenly see the oxidation starting and BAM, it's fully oxidized. It's really fascinating how the oxidation process takes so long to start, but then once it starts it just races full speed ahead. I swear my potassium was oxide free for a good three or four months, then it turned slightly purple and in about a weeks time was fully oxidized again. The peroxide/superoxide has slowly built up in a few spots, and I really wish I could get a nice dry day so I can excise that. However I'm apt to wait before doing that until I can figure out a much better storage solution for this stuff. There's a lot of 'gunk' on the rind and I have to figure out a way to remove the junk, then seal the potassium so it won't corrode any further. I can always take the corroded portion of it and chuck it in water, but I want to keep the bulk of the metal somewhat fresh looking forever. Ultimately, I'd like to preserve it in the 'just starting to oxidize so there's a purple hue to it' state.
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What I'm wondering about is the colors of the oxidation layers of the different metals. Sodium metal is covered by some white layer, which is perfectly understandable, because Na2O and NaOH are white. If I look at my praseodymium sample, then it turns black, but also with some green specks. The green color is understandable, it is the color of Pr2O3, but what is the black stuff? The same thing for calcium. Calcium oxide, hydroxide and carbonate all are white, so what can explain the dark grey/blue or even grey/black color of the sample of calcium, which I have.
You mention that your K-sample is somewhat purple (I checked your picture and indeed it looks somewhat purple) and now also turns orange. The orange color is understandable, but what it the purple stuff?
Etc. etc. with other metals, I simply do not understand the colors of the oxidation layer on the metal.
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The picture of the potassium was taken in December of last year and that particular side of the K was freshly cut in early September. Now, there are tiny little specs of brown/orange on the top part of the 'fresh cut side'. I have a few friends who work with air sensitive materials, so once I can get a nice container to keep the potassium in I'll head over to their work and excise the super/peroxide.
For the coloring, I believe the purple hue to the oxidation is caused by the excited electrons of the metal. I know that potassium salts impart a violet color to a flame, and that's due to the excitation of the electrons in the potassium ion. I think that when oxidation first begins, and the layer on top of the metal is really thin, some of the electrons get 'excited' by simple light waves and this causes the purple hue. This is just an assumption, however, so I'll have to look into it further.
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I have been able to obtain some potassium metal, but I am a bit worried by jdurg posts concerning the stability of the superoxide...
The situation is, that I bought the potassium with a whole lot of other chems, in an all-in-one sale, this was a couple of month's ago, and now I am going abroad soon for my MSc, and I want to store it responsibly, so no strange things can happen to it...
I've made some fancy pictures (clickable), which are below:
This photo is of the largest chunk there is in the bottle.
(https://www.chemicalforums.com/proxy.php?request=http%3A%2F%2Fi16.photobucket.com%2Falbums%2Fb16%2FTaaie-Neuskoek%2Fth_Largechuck.jpg&hash=ef7eac6447bf215878e48b0582d37d029011fe2e) (http://i16.photobucket.com/albums/b16/Taaie-Neuskoek/Largechuck.jpg)
This is one of the few chunks with an orangish spot on it, it looks like it's taking over the chunk.
(https://www.chemicalforums.com/proxy.php?request=http%3A%2F%2Fi16.photobucket.com%2Falbums%2Fb16%2FTaaie-Neuskoek%2Fth_Smallchunk-orange.jpg&hash=fb419c3ce176d2e3dc4084e3f14d47cb11f91858) (http://i16.photobucket.com/albums/b16/Taaie-Neuskoek/Smallchunk-orange.jpg)
This a shot of hte 100ml beaker I put all the chunks in, I did this because I couldn't see the colour properly through the amber bottle the originals were.
(https://www.chemicalforums.com/proxy.php?request=http%3A%2F%2Fi16.photobucket.com%2Falbums%2Fb16%2FTaaie-Neuskoek%2Fth_Beaker.jpg&hash=da7b30b625d322ac60cad8adb4e1aee53bf83c72) (http://i16.photobucket.com/albums/b16/Taaie-Neuskoek/Beaker.jpg)
After I put the small chunk back into the oil, it started to release bubbles :o I was rather scared, and transferred the chunk as quickly as possible to a separate container with oil. Fortunalty no runaway reaction happened. Here is a photo of the event, just before I transferrred it:
(https://www.chemicalforums.com/proxy.php?request=http%3A%2F%2Fi16.photobucket.com%2Falbums%2Fb16%2FTaaie-Neuskoek%2Fth_Bubbles.jpg&hash=7d94fa4658775f344b15b2651893f8f492a618af) (http://i16.photobucket.com/albums/b16/Taaie-Neuskoek/Bubbles.jpg)
So my question boils down to: What should I do with this stuff? I would love to keep some, as potassium is a rare element, and is not very easy to get.
However, I have my resposibility, and when I leave I will come back only once in the 3 month's, íf even that often.
What are the risks of cutting the oxide layer off under a massive layer of oil, is there ANY chance that it will ignite when the K2O2 gets rubbed into the potassium metal by the action of cutting?
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WOW Taaie. Welcome to the forums and those are some great pictures. It's kind of hard to tell exactly how bad the peroxide/superoxide formation is due to the overall orange hue of the photos (damned incandescent light bulbs), but when I get back home from work/golf today I'll try and clean the pics up a bit in photoshop. That second pictures you've shown is quite disconcerting as even though the orange hue I can see a LOT of super/peroxide on there. You DON'T want to cut that oxide part off. While I do believe that some of the tales of peroxide/superoxide igniting when cut are a bit exaggerated, it's not something I'd mess around with as oil fires and alkali metal fires are next to impossible to put out. The pieces that are heavily oxidized are best disposed of in a VERY large lake. hehe.
So I will say that yes, there is a chance that trying to cut off the super/peroxide layer will result in decomposition and a fire. If you have access to a vacuum device and a vacuum jar, I would suggest putting the potassium in the oil and then vacuum sealing it shut. This will keep it preserved as long as the vacuum remains. For the pieces with a lot of the nasty peroxide and superoxide on there, I would just dispose of them in a very large source of water. I know you want to keep them and getting a hold of potassium is a bit of a pain in the arse, but in reality there's not much you can do with that. :( I think you've got plenty of the stuff that should still be in good shape. ;D
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Jdurg, thanks for you compliments and the warm welcome!
Please don't bother to clean all the pictures up, I've made some in good ol' sunlight.
They are below, the photo was made after work, in a Dutch summer/autumn sun, around 5pm.
It isn't too sharp, but the colours are much better, though the oil has also already a bit of a yellowish shine.
Unfortunatly I do not have a special vacuum bottle, however, I could make some very dry oil/solvent with anhydr. MgSO4, filter, add in a 100ml RBF, add potassium chunks, pull vacuum, release, gently not creating bubbles, flush all the oxgen out of the above air layer (I have an aerosol can therefore), and stopper it. Maybe that is a way forward.
The pieces with orange are indeed applying for a single trip to a large amount of water. Is a nice project, and íf I do it, I will try to catch it on tape and post it if anybody is interested...
Ok, the pic, here it is:
(https://www.chemicalforums.com/proxy.php?request=http%3A%2F%2Fi16.photobucket.com%2Falbums%2Fb16%2FTaaie-Neuskoek%2Fth_IMG_0983.jpg&hash=cab166b4fce20e8000cc3ed0dea5e9eb1db652df) (http://i16.photobucket.com/albums/b16/Taaie-Neuskoek/IMG_0983.jpg)
I know I can't do anything with the potassium, but it just doesn't feel good to chuck it all away, but maybe it is the best option.
BUT, if you think of the peroxidised parts of it, if it reacts under a few cm of oil with the K, and there are signs of fire, like a spark or something, to which extend would that be dangerous? The oil is plain parrafin, and cannot burn without oxygen, so there is simply no other oxygen there than is present in the K2O2 form... so under a large layer of oil, (I say large, I mean large, something like 10-15cm in a 600ml beaker) it should be able to abuse the chunks severely without any harmfull consequense for the experiment executioner, shouldn't it?
what do you think? It is handy nor easy to cut with a long knife under a large layer of oil, with a breaking index which is probably going to make you seesick, but I have done worse things...
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The problem is that as the peroxides and superoxides decompose, they generate oxygen gas which can make things ignite. The heat from the decomposition can make the oil hot enough to the point that it can ignite due to dissolved oxygen in the oil. It really doesn't take much at all. You are also right about not wanting to throw away all that potassium, but after taking a picture and adjusting the color balance in photoshop, there is a LOT of peroxide and superoxide covering the surface of all of your potassium. I would be very careful when moving the stuff around and cutting chunks of it.
What I would first do is take a small piece of it, preferably one that has a lot of the orange discoloration, and SLOWLY start to excise the orange colored parts of it with a VERY sharp exacto knife. If you do this, you could try doing this under a little bit of non-flammable oil and have a giant pile of sand nearby to quench any fires that might start. Sadly, looking at all your pieces I see that this must have come from a very old batch of potassium as much of it has corroded. For storage, I think your method of purging the oxygen and getting all of the oxygen/moisture absorbed with a dessicant attached to the cap is a very good thing to do. All I'll say is try not to handle the potassium too much as right now it's already very heavily corroded and oxidized. It may be safe cutting it one time, but then at any random time it could ignite while you're cutting it.
(Oh yeah, heaving chunks of it into water is fun. I did the same with some ten gram lumps of sodium metal a few weeks ago and the explosions were great!. ;D )
(https://www.chemicalforums.com/~jdurg/ShakingNa.jpg)
Force of the explosion shook the camera.
(https://www.chemicalforums.com/~jdurg/Sodium.jpg)
Tail end of a sodium fireball at night.
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What methods do university laboratories use to maintain a long shelf life of the alkali metals such as potassium, sodium, and lithium? I can not believe that they maintain them in oil in a vacuum chamber.
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Universities have enough money where they can just throw away the unused stuff and buy new stuff. They typically have the alkali metals stored in oil in a tightly sealed container which is placed in a fireproof cabinet filled with dessicant. Upon opening the container, they are required to put the date it was opened on the label. Depending on the regulations for the university, after a certain period of time has passed they are forced to send the container to a disposal company which handles the stuff. They typically purchase the potassium in quantities that they'll go through, and if any's left they have enough money to deal with the loss.
Also, sodium and lithium do not form unstable/dangerous peroxides unless subjected to a VERY rich oxygen atmosphere. (I don't even think lithium CAN form peroxides). It's potassium that has the tendency to form those nasty oxygen compounds which is why it has a shelf life.
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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.
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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.
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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
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Pink... If I recall correctly pink layer of (per)oxide is created on potassium - but I can be completely wrong.
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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.
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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!
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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).
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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.