# What Happens if You Make a BATTERY Out of CESIUM?

## Метаданные

- **Канал:** Thoisoi2 - Chemical Experiments!
- **YouTube:** https://www.youtube.com/watch?v=eScaAybE380
- **Дата:** 22.11.2025
- **Длительность:** 19:48
- **Просмотры:** 50,067

## Описание

Patreon: https://www.patreon.com/Thoisoi 
Attention! This video shows dangerous experiments! Do not repeat the experiments shown in this video!

Hello everyone! In this video, I experiment with the most reactive alkali metals — from lithium to cesium — to find out which one is truly the most active. I demonstrate how cesium can be obtained in a lab, show its violent reactions with water, and compare them to rubidium, potassium, sodium, and lithium. Along the way, I explain why cesium isn’t used in batteries despite its high reactivity and why lithium remains the leader in energy density. By recreating real chemical reactions, I uncover how atomic size and reaction kinetics determine the explosive power of these metals.

Welcome to my channel! It's dedicated to experiments in inorganic and organic chemistry! Here, you can find a lot of chemical experiments, each of which contains explanations that will be understandable even to people who are not into chemistry. In my video experiments, I also indicate chemical equations that will help you understand the essence of chemical reactions and transformations. If you have problems with the perception of difficult chemical reactions and chemical equations in school, then you can use some of my videos as a self-help guide in chemistry. Also, some experiments from my videos can be repeated at home, of course, in compliance with all safety rules. Many of the experiments that are shown in my videos are shown to children and used as classic demonstration experiments for schoolchildren or students. Each experiment will be explained as clearly as possible. Chemistry is easy for everyone, even for beginners!
#Thoisoi2 #chemistry #metals

## Содержание

### [0:00](https://www.youtube.com/watch?v=eScaAybE380) Segment 1 (00:00 - 05:00)

Hello everyone. In my hands, I have two alkali metals, lightweight lithium and also seesium, which is considered the most reactive metal on Earth. Reaction. Seesium's reaction with water is the most violent of all metals and happens instantly. But if this metal is so chemically active, then why not use it in batteries instead of the rather less reactive lithium? And which metal might actually turn out to be the most reactive? Well, let's figure it out. To start, let's take a closer look at the periodic table of chemical elements and carefully find those very reactive alkalion metals. And here they are, six elements with francium theoretically being the most reactive since it is located at the very bottom of the group. However, this metal is almost never found in nature and is highly radioactive. So even if it could be artificially produced in a particle accelerator in quantities of a few micrograms, it would simply evaporate due to intense self-heating from radioactive decay irradiating everything around it with a significant dose of radiation. Moreover, according to recent calculations, due to its unique electron configuration, seesium will still be more reactive than francium. And in addition, it has stable isotopes like the one in this ampule, for example. So yes, for now, according to reference data, cesium remains the leader in terms of chemical reactivity. I wonder how difficult it would be to obtain it. Actually, it's really not that hard to do. You just need to build something very similar to a moonshine still for distilling metals, which is basically a steel cup made from a piece of steel pipe welded shut at the bottom to make it completely sealed. On top of it, I carefully soldered a copper tube through which the seesium produced in the reaction will be distilled from the reaction of cesium chloride with lithium. And yes, it turns out that a less reactive metal can displace a more reactive one. And later you'll find out why. I decided to carry out the reaction to obtain seesium in a muff furnace at a temperature of about 700° C so that the lithium would start reacting with the molten cesium chloride, which in turn would evaporate and be distilled into this jar filled with argon. In theory, everything looked good. And after heating up the furnace, a light smoke gradually started to enter the flask. Apparently from the reaction of the produced seesium with residual oxygen in the distillation setup. A few minutes later, the first drops of seesium entered the jar. The price per gram of which is currently about €20. But over time, the reaction somehow stalled. So, I decided to additionally heat the copper tube and the upper part of the apparatus with a gas burner to distill the seesium into the colder sections of the tube. By the way, this actually turned out to be a fatal mistake because I accidentally overheated the solder that was attaching the copper tube to the steel lid, which caused some of the cesium to start escaping through the small holes that formed. When I then tried to somehow patch up this burning hole, the copper tube completely detached from the lid. So, I had to stop the reaction. Fortunately, I managed to obtain a couple of grams of precious seesium in an argon atmosphere and see its beautiful golden color. It's a pity that this amount won't be enough for my further planned experiments, but I can show you just how reactive this metal is. For example, if you simply drop it on the table, it will immediately ignite on its own. Cesium droplets react even more vigorously with water. — [snorts] — What's also interesting is that during my experiments, in order to see the beautiful yellowish surface, I poured a small amount of seesium into a jar with dry water. That is a rather heavy and inert liquid, which is actually a florocetone. Later, after filming, I decided to get rid of it by simply throwing it against the walls. Yeah, I haven't heard such a loud bang in a long time. It turned out that highly reactive alkali metals can indeed react with H hallogen derivatives, which is essentially what dry water is. So for safety reasons, never mix these substances or make any sudden movements with them. Still, for further experiments and to determine which metal is the most reactive, I need at least 10 g of cesium. So I'll have to do everything all over again. For this, I decided to assemble a new distillation setup, and some welder friends helped me by making it entirely out of steel and even welding a stainless steel tube to a steel stopper. Unfortunately, my lithium supplies had run out by the time of filming, so I decided to try the classic reaction for obtaining it. Seesium from cesium chloride and metallic calcium. Fortunately, I had plenty of calcium on hand, and it was even finely ground. So without much hesitation, I mixed it with seesium chloride and loaded it into the distillation setup, then tightly screwed on the lid. To help the cesium distill better this time, I thoroughly

### [5:00](https://www.youtube.com/watch?v=eScaAybE380&t=300s) Segment 2 (05:00 - 10:00)

sealed the upper part of the muff furnace with cow wool insulation. This way, the heat inside will be retained more effectively. This time, I decided to increase the distillation temperature since the temperature sensor in the furnace is located at the bottom, and it might not show completely accurate temperature readings for the upper part of the setup. This time I decided to collect the resulting seesium in mineral oil. It seemed to me that this would be easier than dealing with argon. After heating up the setup, the reaction started again and the first golden colored drops began to enter the receiving jar. But wait, how could the weakly reactive metal calcium displace the highly reactive cesium from its chloride? And I think that's a fair question because if you didn't know, unlike the highly reactive seesium, calcium doesn't ignite in air. And it even reacts with water more slowly than lithium does. Usually in chemistry, it's the other way around. More reactive metals displace less reactive ones from their compounds. For example, like in the thermite reaction, which involves aluminum and iron oxide. But here, the less reactive calcium displaces seesium from its compound. Something strange is happening. But in fact, there's nothing strange about it. After all, in this reaction, the metal boils at a high temperature because we gradually remove the seesium from the reaction mixture, leaving calcium chloride inside. In addition, the energy released during the formation of calcium chloride is higher than that of the same chloride of cesium. This makes the reactions thermodynamically more favorable. This phenomenon in chemistry is called the kinetics of chemical reactions, which can also be influenced by pressure and the concentration of substances in the mixture. But let's return to the seesium being produced which eventually stopped being released. For this I had to further increase the temperature in the muff furnace so that the seesium would start distilling faster. According to my calculations I should get about 15 g of cesium. So the reaction is still far from complete. It was truly fascinating to watch the most reactive metal on earth being produced right here in my laboratory and with such a remarkably beautiful golden color which is actually due to some rather unusual relativistic effects. To put it simply, it turns out that cesium like gold for example has a rather large atom and accordingly a large number of electrons at different levels and sub levels. Because of this, the outer electrons start moving very close to the speed of light which leads to an unusual effect called relativistic contraction. As a result, the atom's electron shell is so to speak slightly compressed causing it to absorb not only ultraviolet light but also the blue spectrum. This leaves more reflected red and yellow light and in the end we see the yellow color of the metal. To compare this to everyday life, for example, imagine that if you start spinning a wheel at nearly the speed of light, then to an outside observer, it will actually appear to contract, even though in reality its radius remains exactly the same as before. It sounds rather unusual, but that's exactly how quantum mechanics and the theory of relativity work, which is why gold and cesium have a yellow tint and mercury remains liquid even at room temperature. While I was explaining all that to you, I've already collected enough seesium in my jar. That's enough for my experiments. To determine which metal is actually the most reactive, I decided to mix 10 g of each alkaline metal with water. And to see just how visually intense the reaction would be, setting up not a theoretical, but you could say a field experiment. But still, for my experiment, besides seesium, I need to obtain or somehow acquire the other alkaline metals as well. After seesium comes rubidium which is even more expensive. So I also decided to obtain it by reacting its chloride with metallic calcium. However, for this I first need to clean my apparatus of cesium residues which had tightly clogged the threads making it quite difficult to remove the lid. After that I decided to rinse the steel tube of any remaining residue using ethyl alcohol. There was seesium inside which turned out to be quite a spectacle. Upon contact with the cesium, the alcohol immediately ignited, burning with a beautiful violet flame due to the presence of this metal's atoms. The alcohol also caught fire in the plastic cup. And it's a good thing I had a fire extinguisher nearby just in case. After I cleaned everything from the remains of cesium and other reagents, I loaded the setup again. This time with a mixture of metallic calcium and rubidium chloride, which turned out to be four times more expensive than the same chloride of cesium. In order to prevent the rubidium from oxidizing too much during production, I decided to flush the assembled apparatus with a stream of argon. For distilling rubidium, the temperature of the muffle furnace needs to be set even higher. This time, since rubidium's boiling point is already approximately 40° higher than that of seesium after everything was thoroughly heated up over time, the very first portions of metallic rubidium began to slowly enter the mineral oil. Unlike seesium, rubidium did not remain in a liquid and golden state, but instead immediately solidified into small droplets that quickly became coated with a thin layer of rubidium oxide. The entire process of distilling rubidium

### [10:00](https://www.youtube.com/watch?v=eScaAybE380&t=600s) Segment 3 (10:00 - 15:00)

took about 2 hours, after which I obtained approximately 12 g of this reactive metal. Now I just need to get three more alkaline metals and in principle I can start the experiment. Fortunately, the metals starting from potassium, namely potassium, sodium, and lithium, are several times cheaper than rubidium and seesium. So, you can simply buy them. Nevertheless, I still decided to show you how in principle you can obtain them. To obtain potassium, instead of using its chloride, you can use potassium hydroxide. And instead of expensive calcium, you can use the much cheaper magnesium. Potassium can be produced in the same setup I used to obtain rubidium and seesium. So, visually, everything will look about the same. Fortunately, to avoid all that hassle, I just took a jar of potassium from my own supplies. The metal these days, it's not the cheapest either, and in terms of price, it's comparable to silver, about €1 per gram. To obtain sodium, you can use the same reaction, but due to its smaller atomic size and higher melting point, sodium is actually easier to obtain using electricity. For this, I took some sodium hydroxide and melted it in a steel cup, which was heated from below using a muff furnace. I continued heating the sodium hydroxide for about 20 more minutes to evaporate any possible water impurities which it could have easily absorbed from the air. After that I run a current through the melt. I insert two electrodes into the sodium hydroxide, one nickel electrode as the cathode and a graphite electrode as the anode. And then I start supplying a current of about 2 ampers. In this reaction, under the influence of electricity, metallic sodium should be reduced at the nickel cathode since it accepts electrons. The current flows from the negative electrode. I chose nickel here for a reason because graphite electrodes often fall apart and can contaminate the mixture with carbon particles. However, even after 15 minutes of electrolysis, I did not notice any visible signs of metallic sodium on the cathode. So, I decided to replace the thick nickel plate with a thinner strip of the same nickel. And after some time, small droplets of sodium began to form on it, which can even be collected with a small spatula. Of course, you can't get much sodium with my homemade setup, but I think the principle is clear to you. The rest of the sodium for the experiments I also took from my own supplies. Well, now all that's left is to obtain metallic lithium. What's interesting here is that it can't be displaced by any other metal from its compounds. For example, from lithium chloride or lithium hydroxide. To obtain metallic lithium, only electrolysis is suitable. This time, instead of lithium hydroxide, I use its chloride since electrolysis usually doesn't work with the hydroxide. First, you need to melt lithium chloride, which has a melting point of 614° C. So, this process isn't very fast. When melting, you can notice how the color of the burner flame turns carmine red due to the lithium atoms. After all the lithium chloride has melted, I immerse the same electrodes that I used for the electrolysis of sodium hydroxide. Only instead of a nickel plate, I used a steel knife blade. All that's left is to apply current and wait until a layer of metallic lithium slowly begins to deposit on the nickel cathode. This process turned out to be rather slow, especially considering that the lithium chloride kept solidifying. Apparently, it wasn't getting enough heating from below by the muff. In theory, metallic lithium should be released at the cathode here. And at the anode, gaseous chlorine, which by the way could be detected in the air a little, even under the fume hood. In the end, after 20 minutes of electrolysis, I decided to check whether at least some lithium had formed on my steel cathode. To do this, I took it out of the melt and immersed it in a cup of water. Yes, there is a reaction and it's actually quite vigorous, which means that some metallic lithium did stick to my cathode after all. So, this method of obtaining the metal really does work. But just like with sodium, it's much easier to buy lithium since producing any significant amount of these metals at home by electrolysis is quite a challenge. It's quite an undertaking to say the least. It's a bit more expensive than sodium, but still more affordable than potassium. Well, I finally managed to collect all the alkaline metals. Now let's see how 10 g of each metal will react with water and whether there will be a reaction at all. Seesium reacts with water even more strongly than for example potassium does. I think we can start with seesium which for some reason didn't solidify for me even though the weather outside was quite cold only 10° C while seesium's melting point is as high as 29° C. I wanted to throw a solid piece of cesium into the water. So, I decided to quickly freeze it using dry ice or dry ice snow, which I made from a regular carbon dioxide fire extinguisher. The temperature dropped pretty quickly, and after a few minutes, the cesium seemed to have solidified. I decided to measure out 10 g of cesium using this kind of notched syringe. However, as soon as the seesium warmed up a bit in the air, it melted again, and I still couldn't cut off a solid piece from it. A solid piece. Maybe the

### [15:00](https://www.youtube.com/watch?v=eScaAybE380&t=900s) Segment 4 (15:00 - 19:00)

seesium was somehow self-heating here due to a reaction with impurities in the oil. It's hard to say. In the end, I just decided to use another syringe to measure out 10 g of liquid seesium and placed it in a different jar. This time under a layer of kerosene. For safety, I will be mixing the metals with water. I'll be doing this on this stand. And for all the reactions, I'll be using warm water at a temperature of about 30° C. Well, let's see how 10 g of the most reactive metal will react with water. Yes, the reaction is pretty fast, but the explosion wasn't particularly loud, even though it seemed to happen twice. That's strange. I thought that the reaction with cesium would just blow this metal dish apart. Now, let's compare it with rubidium. Rubidium itself, when weighing it out, I washed it in cold kerosene and later stored it under a layer of butane to prevent it from oxidizing. Now, let's take a look at the reaction. 10 g of rubidium with warm water here. The bang was noticeably stronger, which you can also see in the camera footage. It's strange as if the less reactive rubidium reacted more violently with water than the more reactive cesium. For better clarity, let's compare it with potassium, which was much easier to measure out. Because of its rather low density, 10 g of potassium looked quite impressive. Let's take a look at its reaction with water. Very strangely, the reaction here looks even more intense. What's going on here? Maybe 10 g of sodium won't react as violently. Let's check. Oh, I think I got a bit stunned after that reaction. It turns out that 10 g of sodium reacted more vigorously than 10 g of cesium with water. But how is that possible? And this is where the laws of chemistry come into play. The thing is the size of a seesium atom is about six times larger than that of a sodium atom. A sodium atom that is. And accordingly, if you take 10 g of sodium, that piece will contain six times more sodium atoms than there are atoms in 10 g of cesium. Because of this, the energy released when 10 g of sodium react with water will be about six times higher than from the reaction of 10 g of cesium with water. If we take lithium, the lightest metal on Earth, then 10 g of it would look roughly like this since its density is twice as low as that of water. However, it will react with water much more calmly than sodium does. It's all because when lithium reacts with water, it forms slightly soluble lithium hydroxide, which creates a protective film on the surface of the metal that prevents the reaction from progressing further. Even if you throw burning lithium into water, it will just catch fire, but there won't be an explosion. But still, sometimes if on the contrary, you pour water onto burning lithium, the reaction can actually be even more vigorous than with sodium. And yes, here there were only 2 g of lithium, but judging by the sound, the reaction was stronger than with 10 g of sodium. As a result, the smaller the size of the metal atom, the more energy is released per gram during its chemical reaction. That is when its electron is transferred to another atom. Of course, here so-called redux potentials also play a major role. But if we really simplify things, then basically yes, the smaller the atom, the more energy per gram. That's exactly why lithium is currently the leading component in lithium ion batteries since such batteries have the highest capacity per kilogram of mass. If cesium were used in them, the capacity would in fact decrease by about 15 times. Additionally, there would be excessive heating during discharge due to the large and sluggish cesium atoms, seesium. So to sum up, in terms of chemical reactivity, that is the speed at which the chemical reaction occurs. Seesium is the clear leader among the alkalion metals. But when it comes to energy density, lithium comes out on top simply because of its exceptionally small atomic size. Well, I think that after watching this video, you now have a much better understanding of how cesium differs from lithium and which metal is considered the most energy dense. But if you enjoyed this video, as always, please don't forget to like and subscribe to the channel in order to discover many more new and interesting things in the future. Beautiful.

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*Источник: https://ekstraktznaniy.ru/video/20441*