Sciencemadness Discussion Board - Experiments with Cesium chloride?

Experiments with Cesium chloride?

jamit - 30-1-2012 at 00:13

i just received some cesium chloride from a friend... the problem is I don't know what experiment I can do with it?

I don't think I want to make cesium... too dangerous!

There's isn't much info on what to do with it on the web, although there are lots of info on cesium chloride.

Anybody done any useful with cesium chloride? I'd appreciate any suggestion or leads I can research. thanks.

weiming1998 - 30-1-2012 at 01:53

Maybe make some CsOH with it. As it is extremely soluble (300g/100ml), it will actually precipate the NaCl or KCl if you poured a NaOH or KOH solution in it. But be careful with that stuff; it's the strongest non-super-base alkaline available. It would be interesting to perform experiments on it.

[Edited on 30-1-2012 by weiming1998]

strontiumred - 30-1-2012 at 08:38

Hi Jamit,

Try this one - It worked well for me:
http://woelen.homescience.net/science/chem/exps/CsCuCl3/inde...

AirCowPeaCock - 30-1-2012 at 09:39

Aren't Cs compounds important catalysts?

jamit - 30-1-2012 at 13:50

@strontiumred
Thanks for the link. I'll give that a try.

woelen - 31-1-2012 at 23:37

There are more interesting things you can do with it:
- Make Cs2CoCl4 (a blue compound, similar to the experiment as linked above).
- Prepare Cs2Cr2O7 (this compound is almost insoluble in cold water)
- Prepare cesium polyholide compounds, e.g. CsICl4 (if you have conc. HCl and an iodate) or CsBrCl2 from Br2 and CsCl.

Cs also perfectly make nearly insoluble CsClO3. You could make CsClO3 directly from CsCl by means of electrolysis or use NaClO3 and CsCl to make this. This compound makes nice blue flames when mixed with sulphur and ignited (be careful though with chlorate/sulphur mixes!). I also did a similar experiment with CsBrO3 (can easily be made from KBrO3 and CsCl, KBrO3 can easily be made from KBr, see my website).

jamit - 1-2-2012 at 03:07

Thanks woelen. Are the synthesis of these compounds on your website? I didn't see all of them. Again thanks.

alkalimetals - 26-4-2012 at 04:00

Hello! I'm new, I don't speak english very well... Sorry

I've found this interesting discussions in Home Chemistry Forum (italian forum) which talk about Cs+ and his compounds and complexes. I hope these articles will be interesting

Caesium Tetrachlorocobaltate (II) Cs2[CoCl]4
http://homechemistry.gratisforum.tv/viewtopic.php?f=3&t=...

Caesium Tetrachlorocuprate (II) Cs2[CuCl]4
http://homechemistry.gratisforum.tv/viewtopic.php?f=3&t=...

Caesium Dichloroiodide CsICl2
http://homechemistry.gratisforum.tv/viewtopic.php?f=3&t=...

Caesium Tetrachloroiodide CsICl4
http://homechemistry.gratisforum.tv/viewtopic.php?f=3&t=...

Caesium Tribromide CsBr3
http://homechemistry.gratisforum.tv/viewtopic.php?f=3&t=...

Caesium Tetraiodomercurate (II) Cs2[HgI4]
http://homechemistry.gratisforum.tv/viewtopic.php?f=3&t=...

Caesium Dichlorobromide CsBrCl2
http://homechemistry.gratisforum.tv/viewtopic.php?f=3&t=...

Dicaesium Octaiodide Cs2I8
http://homechemistry.gratisforum.tv/viewtopic.php?f=3&t=...

blogfast25 - 26-4-2012 at 06:38

Make your own Cs metal:

http://www.sciencemadness.org/talk/viewthread.php?tid=6981&a...

alkalimetals - 26-4-2012 at 07:14

I've got 100 grams of caesium chloride, I had bought to try to synthesize cesium metal, but after reading the patent, I realized it could not make this because I'll never get a vacuum of at least 0.01 mm Hg...

Adas - 26-4-2012 at 07:41

Quote: Originally posted by alkalimetals

I've got 100 grams of caesium chloride, I had bought to try to synthesize cesium metal, but after reading the patent, I realized it could not make this because I'll never get a vacuum of at least 0.01 mm Hg...

I think there's no need for vacuum, use pure argon. I think it should not be extremely hard to get, neither very expensive.

alkalimetals - 26-4-2012 at 07:56

The patent US3164461 contains a tabel which talk about the conditions of reaction... If you work at 760 mmHg the yield is about 50% and the purity is unknow; and you have to keep the temperature between 675 and 760 °C... this is not very easy!

blogfast25 - 27-4-2012 at 11:56

High vacuum not only eliminates oxygen (to which Cs is very sensitive), it also lowers its boiling point a lot)

AndersHoveland - 27-4-2012 at 12:32

Quote: Originally posted by Adas

... to try to synthesize cesium metal, but after reading the patent, I realized it could not make this because I'll never get a vacuum of at least 0.01 mm Hg
I think there's no need for vacuum, use pure argon.

Nitrogen would probably work. Despite being considered so "reactive", elemental caesium is actually completely unreactive towards nitrogen gas. Lithium is the only alkali metal that can burn in nitrogen. (caesium nitride, Cs3N, is actually unstable, and it has probably never been prepared, although some of the older literature refers to caesium azide, CsN3, as "caesium nitride")

[Edited on 27-4-2012 by AndersHoveland]

Eddygp - 27-4-2012 at 12:48

@AndersHoveland However, a vacuum is needed or something to keep the temperature at more than 1000K...

AndersHoveland - 27-4-2012 at 13:12

A vacuum is not absolutely needed, but it would certainly help. Without a vacuum, the distillation will have to reach caesium's standard boiling point, at 671°C. This temperature is going to be fairly difficult to reach, and one would need an iron retort, as glass would melt and fuse into the alkali metals.

One method that might be considered is igniting a "thermite" mixture of CsOH and Mg in a closed metal pot. (the formation of CsH would not be problematic, as this compound decomposes at around 170 °C )

2 CsOH + Mg --> MgO + 2 Cs + H2

(but remember that caesium is still considered the most "reactive" element on the periodic table, it can reduce the chloride salts of all of the other metallic elements)
MgCl2 + 2 Cs --> 2 CsCl + Mg

[Edited on 27-4-2012 by AndersHoveland]

Eddygp - 27-4-2012 at 13:41

Not an easy isolation, that one of elemental Cs

Maybe some electrolysis, in a solution of a cesium salt?

blogfast25 - 27-4-2012 at 14:20

Quote: Originally posted by AndersHoveland

A vacuum is not absolutely needed, [blahdiblahdiblahdiblah di blah diblah blah blah]
[Edited on 27-4-2012 by AndersHoveland]

Is there anything you actually DO KNOW??? :mad:

Any misinformation you're NOT willing to cackle away about???

elementcollector1 - 27-4-2012 at 20:47

Whoa there, blogfast, even if I don't know about the vacuum part, the rest of his post holds up to scrutiny.
As for Eddygp, that would only result in cesium hydroxide, as cesium reacts with water. This is the common industrial method of producing alkali hydroxides from other alkali compounds.

Eddygp - 28-4-2012 at 01:09

And calcinating an hydroxide would only bring up an oxide, I believe. Moreover, caesium isn't considered the most reactive -element-. Maybe it is the most reactive metal or alkali metal, but fluorine and other halogens are much more reactive.

┼┼┼Eddygp┼┼┼

[Edited on 28-4-2012 by Eddygp]

weiming1998 - 28-4-2012 at 03:13

What, now there are copycat bots now?

Spam reported.

Eddygp - 28-4-2012 at 04:17

The bond between Cs and oxygen might be quite complicated to break. After all, it will turn rapidly to cesium hydroxide in presence of water. The problem is that many ions will attach to Cs instead of other metal, because of its properties. Therefore, the vacuum would be the best way, or the only possible one with acceptable yield...

blogfast25 - 28-4-2012 at 05:20

Quote: Originally posted by elementcollector1

Whoa there, blogfast, even if I don't know about the vacuum part, the rest of his post holds up to scrutiny.

No, it bloody well doesn't!

'Thermiting' alkali hydroxides with Mg at best leaves a dangerous mess (especially in the case of Cs) from which any metal formed cannot be recovered. And in closed chamber, the hydride will form when the assembly cools down, at least assuming its formation is exothermic.

Cs and Rb are best prepared by reducing with Li or Ca. Vacuum is needed to:

a. eliminate oxygen
b. remove the Cs as vapour, thereby driving the equilibrium MCl(s) + Li(l) < === > M(g) + LiCl(s) to the right (Le Chatelier).

AH loves talking through his *rse. He's got previouses the length of the Golden Gate Bridge.

blogfast25 - 28-4-2012 at 05:47

Quote: Originally posted by Eddygp

The bond between Cs and oxygen might be quite complicated to break. After all, it will turn rapidly to cesium hydroxide in presence of water. The problem is that many ions will attach to Cs instead of other metal, because of its properties. Therefore, the vacuum would be the best way, or the only possible one with acceptable yield...

Oxides and hydroxides of alkali metals are ionic compounds par excellence. There are no bonds other than M+ (M a generic alkali metal) ions forming an ionic lattice M+ A- (A- a generic anion) with the anions. Much (but not all) of the Standard Heat of Formation of such compounds in the solid form is released when the lattice is formed, when Coulombic attraction between the cation-anion ‘pairs’ is converted to lattice energy (heat).

In the case of the reduction of a Ceasium or Rubidium salt (usually the chloride) the reduction of the metal can be effectuated by means of the equilibrium CsCl (s) + Li (l) < === > Cs(g) + LiCl (l) at quite high temperature. But even in those conditions the equilibrium point is still too far to the left to yield much metal. To remediate that, high vacuum is pulled so the Cs vapour distils off. This, following Le Chatelier, ’pulls’ the equilibrium to the right and the yield of Cs metal is near quantitative.

[Edited on 28-4-2012 by blogfast25]

AndersHoveland - 28-4-2012 at 11:39

Quote: Originally posted by blogfast25

'Thermiting' alkali hydroxides with Mg at best leaves a dangerous mess (especially in the case of Cs) from which any metal formed cannot be recovered.

If I remember correctly, one of the members on this forum isolated potassium from this method, but was unable to isolate sodium. It was speculated that this was either because of potassium's higher volatility, or because the sodium had a higher affinity for hydrogen, and that NaH was forming instead of sodium.

Quote: Originally posted by blogfast25

And in closed chamber, the hydride will form when the assembly cools down,

This is not likely to happen, as the decomposition temperature of CsH (160°C) is much lower than the boiling point of caesium (671°C). The hydrogen may not have enough surface area to react with the caesium rapidly enough before it cools.

Quote: Originally posted by blogfast25

Vacuum is needed to:
a. eliminate oxygen
b. remove the Cs as vapour, thereby driving the equilibrium to the right (Le Chatelier).

As I previously mentioned, a vacuum would certainly help drive the equilibrium, but may not absolutely be required. There are other inert gases that could potentially be used instead of a vacuum. Nitrogen would be suitable if metallic sodium was being used to distill caesium vapor out from a caesium salt. Li, Mg, and Ca can all burn in N2 however, so if using these elements, it would necessitate using argon or a vacuum instead.

[Edited on 28-4-2012 by AndersHoveland]

Eddygp - 28-4-2012 at 11:40

@blogfast, So we finish where we had started: a vacuum is needed, along with a very high temperature.

blogfast25 - 28-4-2012 at 12:57

In a nutshell.

alkalimetals - 29-4-2012 at 00:00

If I use a fridge compressor, reversed, can I obtain a vacuum below 10 mmHg? It's the maximum value described in the patent...

blogfast25 - 29-4-2012 at 04:38

Quote: Originally posted by alkalimetals

If I use a fridge compressor, reversed, can I obtain a vacuum below 10 mmHg?

Personally I doubt that. But I'm no expert on fridge compressors...

alkalimetals - 29-4-2012 at 08:28

Thanks! But maybe it's sufficient a vacuum below 100 mmHg (I'm sure, in this case, fridge compressor works quite well...)

blogfast25, what do you think about?

blogfast25 - 29-4-2012 at 11:32

Try it. But fridge compressors aren't designed to be run as vac pumps, remain aware of that. Only use this idea when you've conclusively proved the fridge 'vac pump' can deliver and over prolonged periods of time too...

BackyardScience2000 - 2-5-2020 at 22:14

I just want to throw out there that I was able to distill cesium metal with no inert atmosphere, at normal pressures without many problems other than a not so great yield. On my best run I was able to turn 20g of CsCl into 10g of Cesium metal. It's definitely not the most efficient way to do it and nowhere near to being the best way. But it can be done. An inert atmosphere and reduced pressure are not necessary. Only highly preferred to increase yields.

j_sum1 - 2-5-2020 at 22:16

Quote: Originally posted by BackyardScience2000

I just want to throw out there that I was able to distill cesium metal with no inert atmosphere, at normal pressures without many problems other than a not so great yield. On my best run I was able to turn 20g of CsCl into 10g of Cesium metal. It's definitely not the most efficient way to do it and nowhere near to being the best way. But it can be done. An inert atmosphere and reduced pressure are not necessary. Only highly preferred to increase yields.

Welcome to SM.
Care to post photos if what you accomplished?

nezza - 4-5-2020 at 07:38

If you have any chlorate or perchlorate Caesium chlorate and perchlorate are pretty insoluble and easy to precipitate. It can then be used for pyrotechnics and gives an interesting purplish coloured flame.

Bedlasky - 4-5-2020 at 08:35

Quote: Originally posted by nezza

If you have any chlorate or perchlorate Caesium chlorate and perchlorate are pretty insoluble and easy to precipitate. It can then be used for pyrotechnics and gives an interesting purplish coloured flame.

Caesium perchlorate is sparingly soluble (1,974g/100ml at 25°C; 0,8g/100ml at 0°C). I don't know why K, Rb and Cs perchlorates are described as insoluble, while CaSO4 is described as sparingly soluble and have lower solubility at room temperature then these perchlorates. But yes, they have really low solubility which isn't common among perchlorates.

But caesium chlorate is soluble. I never read about any insoluble nitrate or chlorate.

Fery - 4-5-2020 at 09:49

I've bought some CsCl just only for demonstration of flame coloring, thanks for everyone for a lot of colorful experiments which are possible with Cs+

Quote: Originally posted by Bedlasky

I never read about any insoluble nitrate or chlorate.

Although not truly insoluble, just only poorly soluble - basic bismuth nitrate should be the less soluble anorganic nitrate in water (less than 1g / 100 ml).