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June 17 2012 1 17 /06 /June /2012 00:41
I heard from NurdRage's excellent video (http://www.youtube.com/watch?v=BliWUHSOalU) that one can easily extract lithium metal from commonly available Energizer lithium - iron disulfide batteries. So I decided to go and purchase some of these batteries. There are actually two grades of these batteries. The Advanced grade has significantly greater power than an alkaline battery and increased shelf life (10 years). The Ultimate grade has slightly greater power than the Advanced grade and has a 15 year shelf life. Because my goal was extracting the lithium and I wanted to purchase the largest amount of lithium for the money, I purchased 4 Advanced AA batteries for 7.46USD at Wal-mart. The batteries were quite a bit lighter than alkaline batteries (lithium being the lightest metal known).
The next day, I decided to tear open one of the fully charged batteries. These batteries contain lithium metal and iron disulfide (pyrite) as electrodes. Due to my past experience with ripping open batteries, the battery came apart like a charm. It hardly required 5 minutes to open it. (By comparison, the extremely tough sub-C nickel-cadmium batteries in cordless drill battery packs took me about 30 minutes to open.) The lithium almost immediately fell out from between the electrolyte membranes. It was silvery gray but quickly turned golden, then greenish, and gradually began darkening from there. I quickly rolled it up (it was very soft, malleable, and flexible and warm from reactivity with the air) and placed it in some mineral oil I stole from my strontium storage container. It floated to the top, as I expected, but still very little of the lithium was above the liquid due to the container's design.
The iron disulfide stunk like hydrogen sulfide, but did not dissolve to any significant extent in hydrochloric acid. Therefore, it was regarded as useless and discarded.
Soon after I decided to try my first experiment with the lithium. I took a large piece of it and placed it on the bottom of a burning place. I had previously melted selenium on the opposite side in an attempt to form red selenium. Some selenium adhered still to the plate. Despite that, I placed the lithium on the plate and ignited it with a propane torch.
The mineral oil ignited first, which soon ignited the lithium. The lithium shrunk, melting, in the flame. (And this is the highest melting of the alkali metals, the hardest, and the toughest. ) The white glowing spots finally manifested themselves into a real flame. In the video, the bright bottom part of the flame is the burning lithium. The long dim flames are the mineral oil combustion, which ceases about half-way into the video. The lithium burns up, leaving behind glowing formations of white lithium oxide ash. (By the way, this ash absorbs carbon dioxide from the air quite well, to the extent that a hydrated form is used in some spaceships to absorb the astronaut's exhalations.)
However, on the dark underside of the plate, sinister events were occuring. The selenium melted due to the heat of the combustion (selenium has a low melting point of 221 Celsius). Selenium has a high vapor pressure, causing evaporation from the dish and subsequent deposition on the floor. High concentrations of selenium vapor resulted in a black deposit, while low concentrations resulted in a red deposit. One piece of selenium actually dripped from the dish's underside onto the floor. Selenium vapors are also very toxic, making it difficult to work with. Fortunately, this was done in a well-ventilated place, making the fumes less of a concern than the stain. I took hydrogen peroxide and a wash cloth and easily scrubbed all of the red stains off the floor, but the black metallic looking spot which was the real concern remained. I used a torch to try to spread it out, which was a bad idea due to the amount of toxic selenium vapor released. This made it necessary to wait until the selenium fumes were completely ventilated. Again the hydrogen peroxide was applied and the floor scrubbed, without much mitigation of the black spot. A scrub brush and a scraper were both applied, but without effect. Just in time, however, I recalled an earlier experiment where the size of a selenium bead shrunk significantly when immersed in bleach, hinting that selenium is soluble in sodium hypochlorite. Therefore, I decided to apply bleach. It worked wonderfully. The selenium stain disappeared in a minute, and nothing remained except for a slightly whiter spot on the concrete.
I then dissected another battery which I almost completely depleted by running a 200 mA flashlight lamp on it for 24 hours. The battery had a nominal voltage of 0.12 V and a maximum current of 20 mA. I opened it up. Just about all of the lithium foil was completely reacted. The black iron looked the same (the iron disulfide is reduced to iron). The lithium sulfide was only a smeary residue on the electrolyte papers. I expected a large amount of lithium sulfide to be present, so I placed it in water. I was greeted with the hydrogen release from all of the left-over lithium inside the electrolyte "paper". The solution turned black and has remained black ever since, even after filtering. It looks like my lithium sulfide is irreversibly contaminated and useless. Two dollars almost went down the drain. However, there was an extra piece of lithium in the battery that had escaped the depletion reaction. I decided to react it with water to obtain some lithium hydroxide. To do this, I took a plastic paint mixing container and placed a significantly sized piece of lithium in it. I then took a garden hose sprayer and sprayed a fine stream of water into it. Unexpectedly, the lithium ignited, spewing out a long red flame and melting a hole in the bottom of the container as all of the water was consumed and the flame's heat had its effect on the container. It was fortunate that this experiment was done outdoors on a concrete patio instead of on some flammable object. The lithium ignition occurred in several steps. My hypothesis is that a piece of the lithium in the upper part of the container was temporarily covered in water. The water reacted as it ran off the lithium, producing heat and steam. Despite lithium's high heat capacity, it does not exceed that of water, so the lithium easily heated to a high temperature in the absence of the water. Steam-laden hydrogen passing over the metal from a reaction in the bottom of the container heated the metal so much more. Eventually it reached the ignition point. The lithium ignited, igniting in turn the hydrogen gas. The hydrogen gas was ignited in a steady stream, enabling it to ignite more of the lithium. A runaway reaction occurred, during which the lithium melted through the container.
Excited by the prospect of the lithium's reactivity, I took another piece of lithium, placed it on a brick, and emptied an eyedropper full of water on it. No ignition occurred. This was because of the lack of the specific circumstances stated beforehand which were used to get ignition. No hydrogen laden steam passed over any portion of the lithium, and there was way too little water.
Not to be discouraged, I tried again two more times. Each time was a failure. Finally, I decided to reproduce original conditions. Original conditions produced (nearly) original results. Take a watch. The steam hid much of the reaction from view.
I then hit upon a novel idea. Why not combine my two previous experiments and drop lighted lithium into water? This should produce a more vigorous reaction. I was right. The first experiment involved a small piece of lithium. It was lighted and dropped into the water. The hydrogen immediately caught fire and kept burning throughout most of the video. The lithium reacted completely, leaving only a red-hot sphere of lithium oxide floating on the water by the Leidenfrost effect. When the temperature dropped to a certain extent, it fell down to the bottom with a fizz. This reaction was quite similar to the reaction of a similar piece of sodium with this amount of water.
 However, I was not satisfied with such a small amount of metal. I decided to step it up a bit. I used a larger piece of lithium and got it burning more thoroughly before dropping it into the water. The lithium ignited the hydrogen with a pop, burning brilliantly for a few seconds. However, due to the reaction thinning out the foil, the vigorous bubbling of hydrogen, and the heat of the flame, the lithium foil broke into several pieces and exploded, shooting gorgeous bits of burning lithium up to six feet into the air. This reaction is more worthy of potassium metal than lithium.
I was done with the bangs and flashes of the alkali metals for now but wanted to catch a more close-up video of the combustion of lithium. I did so with a small piece of foil.
The mineral oil first caught fire, melting the lithium. Unlike the alkaline earth metals, the alkali metals have a low melting point and generally melt before burning. The lithium solidified into a blob and began burning. The white flame gradually became brighter as the lithium was completely converted to the oxide. Then it turned red due to the flame spectrum of lithium oxide and faded as the lithium burnt up. The lithium oxide ash appeared in an interesting lump, which I immediately preserved from atmospheric attack in a glass vial, though in a crumbled state. The formation shown here is very delicate.
Lithium oxide
Since I wanted to keep some lithium for further experiments without resorting to breaking open another expensive battery, I ceased experimenting with the metal after this.
If you have any ideas about what do with lithium, drop a note in the comments section below.     

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Published by LanthanumK - in Experiments
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