Talk:Caesium/Archive 1

Comment
Article changed over to WikiProject Elements format by User:maveric149. Elementbox converted 12:08, 10 July 2005 by Femto (previous revision was that of 12:30, 4 June 2005).

Information Sources
Some of the text in this entry was rewritten from Los Alamos National Laboratory - Cesium. Additional text was taken directly from USGS Cesium Statistics and Information, USGS Periodic Table - Cesium, and from the Elements database 20001107, and Webster's Revised Unabridged Dictionary (1913). Other information was obtained from the sources listed on the main page but was reformatted and converted into SI units.

--

Talk
someone more knowledgable about these things please consider the information at: http://www.cs.rochester.edu/users/faculty/nelson/cesium/cesium_color.html

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color
"Caesium, however, is pale gold in colour" - Shouldn't that be "Caesium, however, is caesium in colour ". The only element I know of that's gold is, well, Gold. ;).


 * It would be nice to know just exactly why cesium is goldish in color. Apart from cesium, copper, and gold, metals are blackish ot whitish silver, no? Some with a bluish hue. But why are these 3 elements so differently colored? Why are there no greenish metals for example? It's probably something about their absorption spectra. But I see no underlying factor why most metals should preferentially absorb in the long-wave part of the visual spectrum, and a few in the short-wave part. Dysmorodrepanis 11:07, 10 February 2007 (UTC)

Bismuth is vaguely pink and I understand some relativistic cause (to do with the speed of electrons in 1s orbitals) underlies this, like the particular (gold!) colour of gold and the liquid state of mercury. No greenish metals appear presumably because the relativistic or overlapping-orbitals mechanisms must, to meet the requirements of the human eye, together see to the absorption of red and blue light to reflect green and maybe that is just too rare given the availability of only 100-odd elements to experimenters. Dajwilkinson 00:13, 17 May 2007 (UTC)
 * Cool! Thanks! Dysmorodrepanis 15:37, 17 May 2007 (UTC)

liquid at room temperature
This article states that ceasium is one of 5 elements that are liquid at room temperature, but the mercury article says mercury is one of 5. Which one is right?

Actually - this says one of 5 metals, mercury lists ceasium, francium, gallim, rubidium and bromine as elements that are liquid. Bromine is a halogen (which is not a metal?), perhaps this is the source of the discrepancy? —Preceding unsigned comment added by 93.209.4.242 (talk) 17:37, 23 April 2009 (UTC)

(Note: both of the above were me - gromgull (talk) 17:38, 23 April 2009 (UTC))

which isotope
I know one isotope is used as a source for cancer therapy (along with Co-60 and a few others). WHich is it? Pakaran. 22:12, 27 Apr 2004 (UTC)

Cesium 137.

It would be nice if articles like this included a link to a page explaining how the pronounciation guide works. ChrisLawson 13:11, 6 June 2006 (UTC)

disputed
Oh come on, there are many stronger bases than hydroxides! LLNL isn't as smart as I'd like. lysdexia 10:00, 19 Oct 2004 (UTC)

Actually, hydroxides are the strongest bases, since they contain the most electropositive elements.

Untrue. Hydroxides can be compounds of a wide range of elements including not particularly electropositive ones (eg copper (II) hydroxide), which are amphoteric rather than strong bases. The most basic hydroxides are the fully dissociated, uncrystallisable ones where the cation is a large complex cation, which need not in fact contain electropositive elements at all. Tetrabutylammonium hydroxide is a stronger base than CsOH and contains C, N and H in the cation! Lysdexia is right that (in nonaqueous environments) there are  H+ acceptors stronger than any hydroxide (eg carbanions), but he/she should have included examples and refs. AGC.

Latin phonetics
Removed from the page:
 * though it goes against Latin phonetics

...apparently in reference to the English pronunciation of the word. First off it doesn't seem relevant (what do Latin phonetics have to do with English pronunciation?) and second off it's inaccurate anyway (The "Caes-" in "Caesar", while not from the same root, was pronounced the same in Latin, /kais/, and the same in English, /si:z/, as this word, so it's not like something inconsistent is happening).

There are exceptions to everything, but the overwhelmingly common pronunciation of "ae" is pronounced like "aye". —The preceding unsigned comment was added by 24.126.5.184 (talk • contribs).

I'm sure there was a reason it was put in though, so if it can be better explained I won't object to putting it back. &#8212;Muke Tever 00:39, 19 Feb 2005 (UTC)

Making Caesium
It can be made by electolysis of CsCl or by reacting CsCl with Ca at a temperature where Cs boils off similar to the Potassium entry.


 * According to http://www.webelements.com/webelements/elements/text/Cs/key.html, Cs is soluble in molten CsCl (like K in KCl). Therefore, the electrolysis method would give low yields.  Instead of Ca, it would probably be better to use a stronger reducing agent such as Na. --Pyrochem 05:51, 20 July 2007 (UTC)

My understanding is the electrolysis method is also done above the Cs boiling point. Also, Ca is about the same reactivity as Na, but has a much higher boiling point, thus using Ca would result in less contamination of Cs by evaporation. In fact the Ca method is what is used commercially. Jokem (talk) 13:45, 14 April 2008 (UTC)

Bromine?
The article currently states that Gallium, Francium, and Mercury are the only three elements that are liquid at or near room temperature. From looking at other wikipedia articles (and high school science classes) it seems to me that we are forgeting Bromine?

Has it been a lie all this time that bromine is liquid at room temperature?

I would go ahead and change it...but I'd like to make sure that I'm right. Would someone please confirm this? —The preceding unsigned comment was added by 69.50.43.220 (talk • contribs).


 * The article specifically narrows the characterization to metals, not elements, so bromine shouldn't be a problem here. Femto 12:08, 16 June 2006 (UTC)


 * I am going to add Bromine, but with the caveat of it being a non metal. Jokem (talk) 13:48, 14 April 2008 (UTC)

Caesium reaction with water
does ceasium react badly with water? i read that it does but im confused!?!?i guess in hot water it would becuse of its low melting point...? —Preceding unsigned comment added by 138.217.219.38 (talk) 02:03, 5 March 2007


 * Alkali metals are highly reactive and react violently in water (very fast and very hot); the further down the periodic table you go, the more violent the reaction.  Squids ' and ' Chips  22:16, 5 March 2007 (UTC)


 * Caesium SHOULD react violently with water and I have seen a video of this happening. However, Theodore Gray (at www.periodictabletable.com) has been trying to react pure caesium with water for a video and the results have so far been disappointing.Fork me (talk) 17:03, 12 February 2008 (UTC)


 * An episode of the British show "Brainiac: Science Abuse" featured a tiny amount of Caesium being dropped into a bathtub contained within a capsule that dissolved in water. The result was "like setting of a depth charge in a bath", blowing it into two pieces.86.159.20.246 (talk) 19:12, 19 August 2008 (UTC)
 * The Brainiac video is available on both Dailymotion and YouTube, although illegally (so I won't quote the links). This is another convincing video I found (unknown copyright status, but at least not an obvious copyvio!) Physchim62 (talk) 19:31, 19 August 2008 (UTC)
 * The Brainiac video was faked :P--HawkFromHell (talk) 14:08, 9 March 2009 (UTC)

In Alex rider Ark angel the football player is given a caesium medallion that reacts with the water in the shower - Hot water. the article says cold water, seeming to imply that it doesnt react in hot water. can some one pls clarify? —Preceding unsigned comment added by 218.214.18.108 (talk) 02:19, 10 September 2009 (UTC)


 * It is implied that it would react in hot water, albeit more vigorously. --Chemicalinterest (talk) 00:51, 24 May 2010 (UTC)

Death by cesium?
is it possibly that you could die by being exposed to a cesium reaction? in a book i read some soccer player is murdered with it… Milldog 93 22:37, 7 March 2007 (UTC)


 * Possible, yes. Practicable, not so much.  See here for visual demonstration.  Duke Leto (talk) 11:47, 22 November 2007 (UTC)
 * Not so sure about using that as a demonstration. An avid viewer told be the other day that it was faked. Possibly by using larger quantities of the metal above caesium in the periodic table —Preceding unsigned comment added by 82.152.248.87 (talk) 02:07, 19 January 2008 (UTC)

Caesium 137 decay
The article claims caesium 137 is used as a gamma-source, but in the bar on the right indicates that caesium 137 beta decays to a stable isotope. These claims are in contradiction, which is correct? —Preceding unsigned comment added by 72.14.228.89 (talk) 17:30, 12 October 2007 (UTC)

Disparity with Francium article
This article lists the estimated amount of Francium in the Earth's crust at any one time as 550g, while the Francium article lists it as 30g. I'm not gonna take a guess as to which is right but I'm pretty sure one must be wrong. --Duke Leto (talk) 11:36, 22 November 2007 (UTC)

Francium
Aren't the MP and BP for Francium only estimated? I did not think it was possible to get enough together to test it. Jokem (talk) 13:55, 14 April 2008 (UTC)

Cesium(III)
A "slow motion edit war" has been going on for months over the inclusion of Cs(III) in this article, as well as the +3 oxidation state for Cs in other articles about the periodic table and oxidation states. The paragraph in question is as follows:


 * There is an account that caesium, reacting with fluorine, takes up more fluorine than it stoichiometrically should.(Klaus Moock, "Indications of Cesium in a Higher Oxidation State" (1989), Angewandte Chemie International 28,12:1676–8. Correspondence to Konrad Seppelt, Institut für Anorganische und Analytische Chemie der Freien Universität, Berlin.) It is possible that, after the salt Cs+F− has formed, the Cs+ ion, which has the same electronic structure as elemental xenon, can, like xenon, be oxidised further by fluorine and form traces of a higher fluoride such as CsF3, analog of XeF2.

IMO this paragraph should be removed. I looked up the reference given, papers citing it, and any related papers I could find. First, the paper by Moock does not make the claim that it is being used for in the paragraph above. It clearly states that the early reports on the preparation of CsF3 were proven to be false; that later results showed that, under certain matrix isolation conditions, a compound could be isolated but it was Cs+F3− (therefore with Cs(I), not Cs(III)). The paper closes with the statement "the isolation of a Cs3+ species has yet to be reported". Now, the core of the paper is about electrochemical evidence of a possible Cs3+ transient species, although oxidation of the solvent or other species in solution remained a possibility. This is an interesting experiment, but as far as I could find, no one else has confirmed it, and I haven't seen much mention of it by secondary or tertiary sources. Therefore, while mentioning the experiment is verifiable, I think including it may be giving undue weight to an isolated observation that has not been proven conclusively. If it is included at all, it should be presented accurately (i.e., this is an electrochemical observation of an oxidation process that may be attributable to oxidation to Cs3+, but no Cs3+ compounds have ever been isolated) but it does not justify adding Cs3+ to the periodic table, articles about oxidation states, etc., because these focus on oxidation states of species that can be isolated. --Itub (talk) 15:52, 14 April 2008 (UTC)
 * Agree. the sum of the first three ionization energies of Cs are nearly 5600 kJ/mol, far more than could be gained in lattice energy from the formation of Cs(F$–$)$3$. However, the CsF lattice is fairly open, and could easily accomodate F$2$ molecules as an addition compound. I'm even willing to believe in Cs(F$3$$–$), but not in Cs$III$ given the current state of chemical science. Physchim62 (talk) 19:43, 14 April 2008 (UTC)
 * Ionization energy and lattice energy aren't the best thing to use, given that Cs (III) compounds, if they end up existing, are likely to be (partially) covalent rather than ionic. From ionization energy and lattice energy alone, one would come to the conclusion that xenon could never show oxidation states 4, 6, or 8; however, xenon has been found to display all three. The fact that Cs (III) hasn't been isolated yet is quite surprising. I actually think researchers ought to try searching for Cs (V) rather than Cs (III), since Xe (II) compounds are less stable than those of Xe (IV); the first xenon fluoride to be discovered was xenon tetrafluoride, and the difluoride wasn't found until later. So I'd bet that CsF5, or even more likely CsF6-, will be discovered before anyone finds anything with Cs (III). Stonemason89 (talk) 13:33, 31 July 2010 (UTC)
 * I have a chance to look at the paper now: it is an excellent piece of experimental work, but I think the authors have been over-enthusiastic in their conclusions. In particular, their hypothesis of Cs$III$ appears to be (IMHO) unfalsifiable, which may well explain why no one seems to have taken the work any further (there would be no point). The question of possible solvent oxidation seems to be dealt with by ad hoc explanations, despite the fact that the ionization energy of acetonitrile (the solvent used in the experiments) is 1177 kJ/mol (Source: NIST), roughly half the value of the second ionization energy of caesium (2234 kJ/mol). Physchim62 (talk) 16:47, 16 April 2008 (UTC)

Units in isotope halflives
Whats with the wierd units? Megasecond/Terasecond are not common units, and most other element pages list halflives in the normal units of hours, days, and years. Shouldn't this article be the same in that regard? The abbreviation Ms for megasecond is confusing for people who have not encountered it before (ie, the vast majority of the population), because of the very common unit of milliseconds, ms.

I have a suspicion that the megasecond/terasecond halflives were put in by someone trying to push adoption of those units. 38.113.0.254 (talk) 16:16, 4 August 2009 (UTC)

Liquid
Would it be appropriate since this comment is in there...

(Some sodium-potassium alloys are also liquid at room temperature.)

To also mention a Gallium Indium alloy exists which is liquid at room temperature?

Ceasius
Doesn't the Latin caesius mean sky blue or heavenly blue, not bluish-gray?
 * No. That's caeruleus. William Avery (talk) 22:26, 28 November 2009 (UTC)


 * Nostris autem veteribus caesia dicts est quae Graecis, ut Nigidus ait, de colore coeli quasi coelia. The colour is blue as the sky and thats what the line in the spectra.--Stone (talk) 08:58, 29 November 2009 (UTC)


 * Interesting he's using it with that classical quote of explanation. Stearne's Botanical Latin, which is good on colours, says "lavender blue", so I can see where the OED is coming from. William Avery (talk) 17:58, 29 November 2009 (UTC)


 * Bunsen visited either a Gymnasium or a Realgymnasium and had to learn more latin than most of the pupils today in a lecture on latin. The amount of classical education was by far greater than today. The education on science was less because less was known at that time.--Stone (talk) 18:34, 29 November 2009 (UTC)


 * The OED has an 1864 quote from Charles Lyell : "He (Professor Bunsen) named the first cæsium, from the bluish-grey lines which it presented in the spectrum." He seems to be supposing what I would call the normal meaning of "caesius". The quote, according to wikisource should be "Nostris autem veteribus "caesia" dicta est, quae a Graecis glaukopis, ut Nigidius ait, "de colore caeli quasi caelia."" For "glaukopis" see this Perseus entry. I still maintain that in general it would be wrong to translate "caesius" as "sky blue". 19:43, 29 November 2009 (UTC)

Caesium oxid
Should we include something about the oxide chemistry of caesium? It is by far the most interesting part, the rest is very similar to the rest of the group.

Caesium and oxygen form a wide variety of chemical compounds Cs7O, Cs4O, Cs11O3, Cs3O, Cs2O, CsO, Cs3O2, CsO2 and CsO3.
 * Cs7O2
 * The orange yellow Cs2O is the only oxide of the anti CdCl2 type it can be produced by heating the Cs7O2 in high vacuum.
 * The dark green Cs3O
 * The superoxide CsO2
 * The ozonide CsO3

Applications rewrite
Obviously, the applications section should be rewritten to use full paragraphs rather than a bulleted list. Any thoughts on how the applications should be grouped? The current organization is logical, but a tad unorthodox. Most applications sections are organized by industry/field. Any thoughts? --Cryptic C62 · Talk 19:22, 12 December 2009 (UTC)
 * First thought is the organization of "other applications" is Ok (in other words, there is no unique way to structure that; other elements are no guide, as application groups are so element specific), but the "Applications" should be merged into "other applications", or these two sections won't make sense.

reactivity with N2
I cannot find any real references for the lack of reactivity with N2 but: a) somebody inserted that CsN3 leads to Cs + N2; therefore it is a definite sink to have metallic Cs in the presence of N2 b) neither Na nor K even remotely touch N2; Li does because Li3N is very stable c) I have not encountered anybody mentioning anything but Li for bringing it into a N2 atmosphere box d) http://www.chemguide.co.uk/inorganic/group1/reacto2.html Anybody got an actual ref about the lack of reactivity? Nergaal (talk) 20:00, 19 December 2009 (UTC)


 * Anybod access to this one ? --Stone (talk) 20:38, 19 December 2009 (UTC)
 * is stable whereas cesium nitride is not, although it takes 1.49 volts more to transfer an electron from lithium to nitrogen than from cesium or Handbook of inorganic chemicals Von Pradyot Patnaik gives also a hint.--Stone (talk) 20:44, 19 December 2009 (UTC)



2.2. Nitrides of the heavier alkali metals

Aside from Li3N, the stabilities of the alkali metal binary nitrides are predicted to be low on the basis of their lattice energies. In fact, only sodium nitride, Na3N, has a calculated heat of formation, ΔH, and extrapolated Gibbs free energy of formation, ΔG, that might suggest a thermodynamically (meta)stable compound (Fig. 4) [21 and 104]. This is also in line with the observed solubilities of nitrogen in the liquid alkali metals. Even in sodium, nitrogen is almost completely insoluble (6×10−11 mol l−1 atm−1) and indicators suggest that nitrogen is dissolved as diatomic N2 rather than as the anion N3− as in liquid lithium [21 and 105]. Correspondingly, the solubility of the respective nominal nitride salt in sodium is also far lower than Li3N in lithium [106 and 107].

The purported synthesis of sodium nitride has been reported under various extreme conditions, yet the compound remains ill-characterised. No binary compound can be prepared from direct reaction of the elements at temperatures up to 800°C, even in the presence of an iron catalyst [104 and 108]. This useful property accounts for the use of nitrogen as a cover gas in liquid metal manipulation and in purifying the alkali metal. Coupled with the appreciable solubilities of lithium and the alkaline earth metals in sodium, this also allows use of liquid Na as a solvent for the preparation of the binary nitrides of these elements (as is highlighted in previous and subsequent sections). The successful synthesis of sodium nitride is only reportedly achieved using electrically activated nitrogen [109 and 110]. At low nitrogen pressures, the products of this sodium–nitrogen reaction are dependent on reaction time; over periods of up to 5 min the only observed product is Na3N whereas over longer discharge durations both nitride and azide, NaN3, are formed [104 and 109]. The identity of the nitride has never been reliably proven, however and no structural information exists. There is some speculation that nitride prepared in this way is in fact a mixture of the azide and sodium metal [12 and 111].

Reports of the synthesis of the other alkali metal nitrides are even more contentious. It is reported that the nitrides of potassium, rubidium and caesium can be synthesised following a similar procedure to the sodium case, reacting the metals with electrically activated nitrogen gas [104 and 110. H. Wattenberg. Ber. D Chem. Gesellschaft 63 (1930), p. 1667. Full Text via CrossRef110]. Here, the species appear to be even more short-lived; further reaction to the respective azides is far more rapid than for ‘Na3N’ with KN3, for example, detectable after only 30 s [110]. The black ‘nitrides’ are reportedly contaminated but, interestingly, reacted with hydrochloric acid to yield ammonia rather than hydrazoic acid supporting the existence of the N3− anion (as opposed to azide) in the binary compounds [109]. Earlier unsubstantiated reports suggest that rubidium and caesium nitride can also be produced by the action of nitrogen gas on the corresponding hydrides [112]. Jansen and Schön recently used global optimisation techniques to calculate possible kinetically stable structures for these hypothetical compounds. A number of viable metastable modifications for Na3N in particular were proposed [113]. Where [112] is: H. Moissan. Compte. Rend. 136 (1903), p. 587.
 * WP:RS for it, so I'm fine now. DMacks (talk) 03:17, 20 December 2009 (UTC)

todo
Nergaal (talk) 20:05, 22 December 2009 (UTC)
 * http://pubs.usgs.gov/of/2004/1432/2004-1432.pdf
 * http://minerals.usgs.gov/minerals/pubs/commodity/cesium/mcs-2009-cesiu.pdf

Health effects
There seems to be a bit of a conflict in this article.

"most caesium compounds are mildly toxic because of chemical similarity of caesium to potassium. Exposure to large amounts of Cs compounds can cause hyperirritability and spasms, but as such amounts would not ordinarily be encountered in natural sources"

"The median lethal dose (LD50) value for caesium chloride in mice is 2.3 g/kg which is comparable to the LD50 values of potassium chloride and sodium chloride.[50]"

If the LD50 is comparable to that of other alkali chlorides, how can it be "mildly toxic"? Or are the "large amounts" so high that it those quantities of other alkali chlorides would present symptoms of toxicity?

The LD50's for NaCl and KCl are reported as 3-8g/kg and 2.6 g/kg respectively. —Preceding unsigned comment added by 38.113.0.254 (talk) 19:51, 11 January 2010 (UTC)

See also vandalism
There's obvious vandalism under the See also section. Normally, I'd revert it, but it's old enough and there's been enough edits since its creation, that I can't see the original edit and check to see if any other vandalism was done with it.--Prosfilaes (talk) 16:07, 15 January 2010 (UTC)


 * it was a linked template: Nergaal (talk) 16:50, 15 January 2010 (UTC)

Remaining FAC issues
The two big issues from the previous FAC were over-reliance / close paraphrasing / plagiarism of the USGS source and a general need to copyedit the article. Here are the specific issues that were not fully addressed:


 * An article which provides a wikilink for "liquid" but not for "borosilicate glass" is in serious need of its links being reviewed.
 * I think the linking is mostly done, it was already done during FAC mostly --Stone (talk) 21:32, 6 March 2010 (UTC)
 * The section beginning "Acid digestion is the principal commercial method used and usually employs..." again cites the USGS paper but again looks to me like unacceptable close paraphrasing or plagiarism.
 * This looks now a little totally different.--Stone (talk) 21:32, 6 March 2010 (UTC)
 * Next, take the second para under "compounds", which currently begins "Caesium chloride is an important source..." There is only one footnote, at the end of the para. The combination of a lack of context here, and an odd turn of phrase ("The chloride atoms lie upon the lattice points...") makes me suspicious that this has been inappropriately plagiarised or close paraphrased from the source (or from elsewhere). I would be grateful if someone with access to this book would take a look.
 * This phrase under "other uses" - "Concerns about the corrosive action of caesium on spacecraft components, have pushed development in the direction of use of inert gas propellants" - is a verbatim quote from the source (USGS), while other material is a close paraphrase. At the same time, the USGS material has been broken up with this sentence: "It used a method of ionization to strip the outer electron from the propellant by simple contact with tungsten". This fact, however, is in fact not in the USGS study at all, and is therefore completely uncited.
 * Section Atomic clocks; "...control and regulate information flow on the internet". That sounds interesting to me as a non-chemist layman. But no further information is given. It made me curious. Could it be expanded with a brief sentence explaining how it does this?
 * There are some seemingly unnecessary instances of 'even', 'On the other hand', 'In contrast', 'originally' and 'in reality'. In many cases, removing these shouldn't hurt and will tighten up the text.
 * The relevance of much of the sentence that begins "Francium may be more electropositive..." is unclear to me.
 * I fixed this during the FAC: I converted it to a note so that the statement is preserved but also doesn't interrupt the flow of the rest of the text. The statement is relevant because caesium is not necessarily the most electropositive element—it's the most electropositive one for which the electropositivity is known, and that's an important distinction that should be kept in the article. {&#123; Nihiltres &#124;talk&#124;edits&#124;⚡}&#125; 22:28, 6 March 2010 (UTC)
 * Is there any information about why caesium was used for the definition of a second and also why caesium is specifically used for atomic clocks?

I will try to address some of the issues in the coming days.--Stone (talk) 21:54, 3 March 2010 (UTC)

This looks decent: There's an example on p. 256, a table of short-lived isotopes on p. 280, a model graph on p. 291. This ref. includes some discussion of caesium: Here's a ref. on isotopic evidence from the early Solar System: Interestingly, there was also a study of Caesium in the atmosphere of a brown dwarf, where it was used to probe the atmospheric chemistry: But that's a little off topic.&mdash;RJH (talk)
 * The article might mention that caesium is produced by the slow neutron capture process (S-process) in stars; primarily AGB stars.
 * You are right. Do you have any references for that as I am not sure I have any books for it? Nergaal (talk) 06:56, 20 February 2010 (UTC)


 * Would this be OK? Caesium-133 is produced in in stars by the slow neutron capture process (S-process starting from lighter elements.
 * Well it's a true statement, but I think that multiple caesium isotopes are made in that manner, which decay shortly thereafter. I checked the following source and it looks like Caesium-133 is also produced by the r-process, so supernovae could also be added to the list.
 * There's a comparable example in the last paragraph of Xenon.&mdash;RJH (talk)
 * You should fix the URL to http://books.google.com/books?id=Gd_L9binuDsC&pg=PA87, which will both change the interface to English (using google.com instead of google.de) and make the link show the page cited instead of the index page referring to the page cited. Otherwise, it looks great. :) {&#123; Nihiltres &#124;talk&#124;edits&#124;⚡}&#125; 21:20, 21 February 2010 (UTC)
 * Not sure what you're talking about here. The Busso et al. link doesn't have pages blocked as on the google page.&mdash;RJH (talk)
 * You should fix the URL to http://books.google.com/books?id=Gd_L9binuDsC&pg=PA87, which will both change the interface to English (using google.com instead of google.de) and make the link show the page cited instead of the index page referring to the page cited. Otherwise, it looks great. :) {&#123; Nihiltres &#124;talk&#124;edits&#124;⚡}&#125; 21:20, 21 February 2010 (UTC)
 * Not sure what you're talking about here. The Busso et al. link doesn't have pages blocked as on the google page.&mdash;RJH (talk)


 * There is information available about the solar abundance of caesium. I believe it differs from the abundance in the Earth's crust. Abundances of the elements (data page).

Image
The image under the appearance heading in the right box when clicked brings the user to a different image than the one shown. —Preceding unsigned comment added by 174.31.81.2 (talk) 20:04, 23 August 2010 (UTC)
 * I think I have fixed it now. -- Ed (Edgar181) 20:16, 23 August 2010 (UTC)
 * The image on the English WP site is Image:Csmetal.jpg.jpg, which got transfered to commons as commons:Image:Cesium.jpg. But there is already an Image:Cesium.jpg on English that is something different, masking visibility the commons one. All three of these are used in various places in the wiki-world--very confusing! Probably cleanest to rename Image:Cesium.jpg on .en to Image:Cesium2.jpg so that the commons one is visible as expected. Then Csmetal.jpg.jpg (which is a nonstandard naming pattern anyway) can be deleted as dup-of-commons. DMacks (talk) 20:30, 23 August 2010 (UTC)
 * OK, I didn't see your reply until now, but have done basically what you describe. I have transferred the local version of cesium.jpg to Commons:File:Cesium ampule.jpg and deleted the local version.  This allows the Commons version to "show through".  And I have deleted Csmetal.jpg.jpg and converted all the uses to cesium.jpg.  Can you doublecheck that everything worked out right, DMacks?  -- Ed (Edgar181) 20:43, 23 August 2010 (UTC)
 * Looks clean, thanks for the cleanups. File:Csmetal.jpg.jpg is still showing as being used in Caesium, but I don't actually see it there...probably just stuck in some cache somewhere. DMacks (talk) 21:00, 23 August 2010 (UTC)

Some issues, possibly
I will read up more and report later.--Smokefoot (talk) 17:35, 17 October 2010 (UTC)
 * The section on the #1 app, drilling fluids, is at least partly plagiarized from the USGS pamphlet. We might track down the contributing editor and check their other contributions.
 * This has been debated at Wikipedia_talk:WikiProject_Elements/Archive_10. Nergaal (talk) 17:58, 17 October 2010 (UTC)
 * We might need to have a look once more what we can do to improve here. I tried to get out mots of the to close phrases to the USGS paper. --Stone (talk) 20:49, 17 October 2010 (UTC)
 * I question the authority of this pamphlet as well. Ullmann's Encyclopedia of Industrial Chemistry says this about cesium formate "The very low toxicity of the cesium cation ... have led to the suggestion to use these solutions as brines in oil..."  This USGS pamphlet is the most highly cited source in the article, its backbone one could say and it may not measure up, in part because it is US-centric.  Its main virtue may be that it is accessible to editors.
 * Feel free to improve the article. Nergaal (talk) 17:58, 17 October 2010 (UTC)
 * There are other sources for the use as drilling fluid, but they are less credible than the USGS. The web page of the company producing the caesium formate and a oil drilling magazine (I did not like to call it a peer reviewed journal ;-) all give the same impression that this is the largest application. The world wide production is located in one mine and the owner is the company marketing the drilling fluid so it seems locical that the USGS paper is right that the caesium formate is THE caesium application by weight.--Stone (talk) 20:49, 17 October 2010 (UTC)
 * I have not checked carefully, but the initial chem section is somewhat misleading "Isolated caesium is extremely reactive" It is in fact robust, even distillable.
 * If you think caesium metal is not reactive, then feel free to check that with any highschool chemistry book. Nergaal (talk) 17:58, 17 October 2010 (UTC)
 * Okay, good advice...--Smokefoot (talk) 19:32, 17 October 2010 (UTC)
 * We should have a look at potassium and fluorine to compare the standard we use. reactive compared to potassium or reactive compared to iron. This is the question.--Stone (talk) 20:49, 17 October 2010 (UTC)
 * The safety section is hystrionic and semi-hysterical "Caesium metal is one of the most reactive elements and is highly explosive when it comes in contact with water. The hydrogen gas produced by the reaction is heated by the thermal energy released at the same time, causing ignition and a violent explosion." I just dont think that the common person is ever, ever going to encounter metallic Cs, so it can be construed as misleading to the common reader to emphasize such esoteric behavior.
 * I don't think that the common person will come in contact with caesium salts either. This article is supposed to be encyclopedic, not a handguide about what readers should do when they come in contact with whatever form of the metal they could/might encounter in their average day. Nergaal (talk) 17:58, 17 October 2010 (UTC)
 * yes, I got that message, but still ....--Smokefoot (talk) 19:32, 17 October 2010 (UTC)
 * The wording of that sentence could definitely be improved. Caesium is not "highly explosive". There are plenty of videos out there to show that it can "react explosively" with water, although there are some doubts as to how genuine they are (a bit of hydrogen peroxide in the water would certainly add explosive effect). In the absence of oxygen, the reaction certainly isn't explosive. Physchim62 (talk) 20:24, 17 October 2010 (UTC)
 * Right, but like written before if it is compared to iron it is "react explosively" with water compared with potassium it is only a little more reactive, so we should seek a common standard for all the reactive elements.--Stone (talk) 20:49, 17 October 2010 (UTC)
 * Stone, stop saying this. It's just baloney. Alkali metals in water Or Watch to the end. S  B Harris 21:40, 17 October 2010 (UTC)
 * Youtube is NOT a credible source! The introduction of youtube videos as a credible source for the the hazards will jeopardise the FA status. Also self made experiments do not qualify for as a source so lets go to the real sources like MSDS from various sources and there the potassium and  caesium read almost identical. If you have a credible source for the bath tub explosion you are welcome to add it.--Stone (talk) 08:05, 18 October 2010 (UTC)
 * Those are NOT credible sources. See the remark of Physchim62 above; we have no way of knowing it those are real experiments.  And I can dig up a youtube video of an exploding iron bar (IIRC it was a train rail, but I can't find that, but this experiment is similar, though on a wire).  So when the right conditions are applied, iron is just as explosive as Caesium.  --Dirk Beetstra T  C 08:48, 18 October 2010 (UTC)
 * For heavensake, people. I've put all those elements, with the exception of rubidium, into water myself. The effects are just as shown. I did cesium in a very large argon-filled glovebox, (little mercury-like watchglass pool of NaK kept in the corner to show no air leaks!) so there was no flame. However Cs reacts explodively in cold water, no matter how little you use, and whether the H2 ignites or not. I went looking on youtube to show you something I already knew from personal experience, and your reaction is perverse to say the least. You won't believe me, you won't believe videos from several sources, you won't believe a myriad of printed descriptions (google ["cesium metal" water] and just look at the textbooks). So what are you going by? Something user:Stone says? What makes you think you are the *&^%ing expert, here, Stone? S  B Harris 17:23, 18 October 2010 (UTC)
 * Rubidium and cesium also are highly reactive with water, generating hydrogen in the fashion typical of alkali metals.Encyclopaedia Britannica 1981 and Cesium hydroxide (CsOH) is the strongest base (alkali) with the highest pH value of any chemical yet found. It is easy to produce by just placing cesium metal in water (which is very reactive). The history and use of our earth's chemical elements: a reference guide 2006. I am not suggesting I am the expert on caesium, but Encyclopaedia Britannica and others are relaiable sources  for me to check.--Stone (talk) 09:03, 19 October 2010 (UTC)
 * Encyclopaedia Britannica, like all encyclopedias and tertiary sources (including WP itself!) is not a reliable source, by WP's own policies: WP:RS. So if you want to wikilawyer this, try again. Here's a list of textbook sources for the explosive reaction of cesium over potassium, including a note from Beilstein that cesium reacts explosively with dry chloroform. Since I know this topic, having done much element collecting in my chemical synthesis days (my undergrad degree was chemistry, I still do medicinal chem patents, and I still own a 20 year-old cesium sample), I thought I'd help y'all out. I see it's no use-- though you have no personal experience with the heavier alkali metals (this is obvious) you're nevertheless set in your preconceptions. Your mind is like concrete: all mixed up and firmly set. If you don't WANT to believe something, no evidence from others will convince you. So, write what you like! Personally, I'm done with this topic; this kind of argument is why Wikipedia can suck so badly. I'm going to go to other articles and hope its a long time before I have to argue with (say) some biomed grad student about medicine. Yuck. S  B Harris 19:33, 19 October 2010 (UTC)
 * The reactivity of alkali metals toward water is in my view a distraction that non-chemists often highlight because of the drama. But articles are more "interesting" if they can discuss such stuff, but no big deal.--Smokefoot (talk) 13:29, 19 October 2010 (UTC)

USGS citation thing
BTW, my comments are intended to improve an already excellent article. My main concern is the heavy reliance (25 citations) on the USGS document: --Smokefoot (talk) 13:29, 19 October 2010 (UTC) I wouldn't be so damning about the USGS reviews as you are: for me they serve a very useful purpose. They serve as secondary sources for uses of the mineral commodities they cover: the USGS scientists have already done a selection of which commercial claims to believe and which not to. If we didn't use them (and obviously we don't have them for all commodity chemicals), we are limited to repeating the claims of commercial suppliers or restricting ourselves to possibly outdated textbooks. Physchim62 (talk) 03:00, 20 October 2010 (UTC)
 * T he current article (10-19), both in the body and the lede, indicates that the major application for Cs is in drilling fluids, whereas Ullmann's Encyclopedia of Industrial Chemistry implies that this application is minor/exploratory compared to other applications. The USGS document is highly geo-oriented, understandably so, being written by a geology office.
 * I do not question the good intentions of the USGS document, but most of its references would not be allowed in Wikipedia, so the source itself is therefore of questionable authority. It seems nonideal to cite a "geology document" to support statements on basic physical and chemical properties.
 * Documents by the US government - or any government - are probably less objective than non-government sources, such as textbooks and scholarly works.
 * Perhaps we should restrict citations to the USGS document to the sections on occurrence and ore extraction, i.e. areas where the expertise of geologists apply.
 * Addendum: I check Chemical Abstracts tonight on caesium formate - 200+ hits, many of which are within the past few years and almost all were patents for drilling fluids.--Smokefoot (talk) 02:46, 20 October 2010 (UTC)
 * USGS covers the information gap between science and companies, and I know no good alternative for that. I'm sure you realize (@smokefoot) that the mining/production/use numbers from "reliable scientific sources" are silently copied from either USGS, or, more often from much less reliable sources. Materialscientist (talk) 04:13, 20 October 2010 (UTC)

If you're going to assign Caeseum formate as an important ingredient of Drilling fluid you should probably say in the drilling fluid article as to what its function is, so that people who are interested in the function of drilling might learn something.WFPM (talk) 10:56, 20 October 2010 (UTC)

"one of only five metals that are liquid at or near room temperature"
The article Metal states that the word can also mean an alloy of metallic elements. Then there would be not only 5 "metals" but many more as in Low-melting alloy. I propose changing the sentence in the headline to "one of only five elemental metals that are liquid at or near room temperature". Michbich (talk) 06:57, 29 November 2010 (UTC)
 * Agree and done. mgiganteus1 (talk) 10:57, 29 November 2010 (UTC)

Reaction with water
If it explodes when exposed to water, how did they discover it in water, not exploding?


 * What the discoverers found in mineral water was cesium's ionic compounds, not the pure element. S B Harris 22:35, 29 November 2010 (UTC)

Spelling
Moved to Talk:Caesium/archive/Spelling Lanthanum-138 (talk) 06:31, 3 February 2011 (UTC)

Nuclear Reactor waste?
Where's the evidence that Caesium comes from radioactive waste at nuclear reactors? The fact doesn't seem to be referenced anywhere (No original research, remember?). Besides, Caesium wouldn't make good heavy water anyway... —Preceding unsigned comment added by 99.238.195.225 (talk) 00:47, 27 April 2011 (UTC)

Cesium v. Caesium
There should really be something on the name of it, as there is discrepancy on the spelling, like which is more used, which are accepted by IUPAC, NIST, etc. 141.161.179.25 (talk) 13:16, 30 June 2011 (UTC)
 * It's discussed in Note 1.--agr (talk) 17:15, 30 June 2011 (UTC)

Higher oxidation states
A recent paper in Nature Chemistry, authored by Mao-sheng Miao, found that under conditions of extreme pressure (greater than 30 gigapascals), the inner electrons (i.e those in the d and f shells), could possibly be used to form chemical bonds. This discovery indicates that compounds such as caesium triflouride or caesium pentaflouride could exist under such conditions.

Do we really need this?

--Stone (talk) 19:06, 22 November 2013 (UTC)


 * I would say yes. You would include it if they were synthesized, right? So I think they deserve to be mentioned, as it's especially interesting for an alkali metal (which "should" by all rights have +1 as its maximum oxidation state, not +5!). I think this is kind of like HgF4 before its synthesis; a predicted breach of the inner subshells of a main group element by fluorine, though not confirmed experimentally as yet. (Note, it's not the d- and f-electrons that are bonding; it's the p-electrons. In these higher fluorides Cs behaves as though it were the seventh 5p element, after Xe.) Double sharp (talk) 13:45, 23 November 2013 (UTC)
 * P.S. A nice guy on Reddit has posted a link to the fulltext of the paper: . Check it out! Double sharp (talk) 13:46, 23 November 2013 (UTC)


 * It is a primary source of a fact which is not in reach for quite some time, why not wait for a review picking up this? --Stone (talk) 22:44, 23 November 2013 (UTC)
 * I think it should still be in there, but only as a note. It's not really important as yet, being unsynthesized, but once it is it certainly would be classified as important (along with HgF4). It is still interesting. Double sharp (talk) 04:27, 24 November 2013 (UTC)
 * The paper also states that Ba could probably act as the eighth 5p element under similarly extreme conditions, and Rb the seventh 4p element. Calculations have not yet been done for Rb and Ba, though, so I think only the Cs results could or should be mentioned. Double sharp (talk) 09:28, 15 May 2014 (UTC)

ENGVAR: en-GB
Just discovered that this article is in British English. So I have set the infobox en-GB accordingly (to show vapour, vaporise, colour). Phosphorus is the other one in en-GB. -DePiep (talk) 17:58, 16 December 2014 (UTC) To be clear: the element name is & stays 'caesium' by IUPAC definition. -DePiep (talk) 17:59, 16 December 2014 (UTC)

test
test archive routine -DePiep (talk) 21:34, 15 April 2017 (UTC)

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Metallic cesium is NOT yellow
The article wrongly stated that metallic cesium is yellow-colored. It's not. Pure metallic cesium is colorless like most metals, but takes on a yellow hue when exposed to oxygen. See here http://www.chemistryexplained.com/Ce-Co/Cesium.html I have corrected the error. HandsomeMrToad (talk) 12:08, 15 January 2017 (UTC)
 * I should have thought that Greenwood and Earnshaw was a more reliable source than chemistryexplained.com. The oxidation explanation has been given by many others, but was already refuted in 1984. Addison's The Chemistry of the Liquid Alkali Metals explains: "When free from surface contamination, the metals Li, Na, K and Rb are silver-bright and lustrous. Caesium, however, is pale gold in colour in both solid and liquid states; coupled with the low density, and the low viscosity of the liquid, it surely must be the most attractive element in the periodic table. Earlier workers attributed the colour to the presence of trace amounts of oxide impurity, and certainly the gold colour does darken on the addition of oxygen. However, intensive purification of the metal by all available methods in the writer's laboratory over several years has not succeeded in removing the colour, and we have no doubt that caesium metal is indeed pale golden in colour." An explanation in terms of sd overlaps follows. I've reinstated the statement of golden-yellow colour. (Similarly, the heavy alkaline earth metals Ca, Sr, Ba are pale yellow, as are the divalent lanthanides Eu and Yb.) Double sharp (talk) 09:11, 16 January 2017 (UTC)


 * 1. Look at the image in the article.  Does that look "pale yellow" to you?


 * 2. One chemist's failure to remove the yellow color experimentally is not "refutation" of the assertion that the pure metal is colorless.  It shows only that total purification is a formidable challenge.


 * 3. And, if, as you say, the heavy alkaline earth metals Ca, Sr, Ba are pale yellow, as are the divalent lanthanides Eu and Yb, then the wiki article on copper should say so, should not say "copper is one of only four elements with a color" as it does.


 * Cheers. HandsomeMrToad (talk) 20:40, 16 January 2017 (UTC)
 * I agree that chemistryexplained.com can't be taken as a definitive source here. Nor does looking at images online really help; exact colour reproduction is notoriously unreliable.  But for what its worth, I think the image does indeed look slightly yellowish. SpinningSpark 00:04, 17 January 2017 (UTC)
 * Exact colour reproduction is horribly unreliable for this sort of thing: at least we're not dealing with spectra, which would be completely impossible. And yet, the picture does look yellowish enough to tell, and the fact is already long since standard: everyone who is an actual chemistry professor is quoting it, and it is in such standard texts as Greenwood and Earnshaw's Chemistry of the Elements. I will change the copper article. Double sharp (talk) 06:28, 17 January 2017 (UTC)


 * Well this may not be exactly relevant and I'm not trying to be rude here, only entertaining, BUT, having said that.... LOL LOL the actual chemistry professor in your link (Richard Kaner of UCLA) says cesium is "the only other metal besides gold and copper that is not silvery in color." BUT osmium is blue!!  (Wrawwwwwng, Prof!)  HandsomeMrToad (talk) 09:54, 17 January 2017 (UTC)


 * Possible resolution/consensus: how about describing the question of the color of cesium as CONTROVERSIAL (both here and in the wiki article on copper)? Some sources say pale yellow, others say colorless, the truth is very difficult to measure because it's hard to verify that a sample is really pure of oxides.  HandsomeMrToad (talk) 09:47, 17 January 2017 (UTC)
 * Osmium is a bit blue rather than pure metallic-silver according to our article text, and its photo, and all over the internet. DMacks (talk) 11:02, 17 January 2017 (UTC)
 * Osmium is silvery-bluish, and the blue tinge is not anywhere near as intense as it is for the yellow of caesium and gold. Double sharp (talk) 12:45, 17 January 2017 (UTC)
 * chemistryexplained.com cites three refs:
 * Emsley, John (2001). Nature's Building Blocks: An A-Z Guide to the Elements. New York: Oxford University Press.
 * Greenwood, Norman N., and Earnshaw, A. (1997). Chemistry of the Elements, 2nd edition. Boston: Butterworth-Heinemann.
 * Lide, David R., ed. (2000). CRC Handbook of Chemistry & Physics, 81st edition. New York: CRC Press.
 * Double sharp says that this Greenwood and Shaw ref actually says it's actually gold. I just checked the Emsley ref's entry for ""Caesium or Cesium (USA)", which states that it is "gold-coloured" without further qualification. Anyone have a CRC handy? DMacks (talk) 11:13, 17 January 2017 (UTC)
 * Here is the quote from Greenwood and Earnshaw (p. 74): "All [group 1 metals] are silvery-white except caesium which is golden yellow; in fact, caesium is one of only three metallic elements which are intensely coloured, the other two being copper and gold (see also pp. 112, 1177, 1232 [these pages detail some of the less intensely coloured metals, like Ca, Co, Zn, Sr, Cd, Ba, Eu, Yb])." Double sharp (talk) 12:44, 17 January 2017 (UTC)

I don't think that it is reasonable to describe the issue as controversial, unless there are some recent RS saying, for instance, that Greenwood & Earnshaw (or whoever) are talking out of the wrong orifice. Or we have an RS that out and out declares that the issue is controversial. Using out of date or not so reliable sources is not good enough to demonstrate a controversy. Saying in the article that there is a controversy without something reliable to back it up is WP:SYNTH. SpinningSpark 15:44, 17 January 2017 (UTC)

an attempt to be helpful Comment
as proven in other articles (thankfully few) i have been proven to know very little about how to 'set things out' in wikipedia... however this comment is intended to help. last thing i want ever is to start a war. so please dont get the wrong idea. im no scientist - i really only know stuff from looking it up. this article says caesium is one of five elements... i googled 'what elements are liquid at room temperature' and got the following answer. Liquids around room temperature. The only liquid elements at standard temperature and pressure are bromine (Br) and mercury (Hg). Although, elements caesium (Cs), rubidium (Rb), Francium (Fr) and Gallium (Ga) become liquid at or just above room temperature. my knowledge isnt near good enough to debate. i wanted only to comment and hope someone found it useful, then go away. have a great day. — Preceding unsigned comment added by Sanasazi (talk • contribs) 03:18, 28 March 2017 (UTC)
 * It's five elemental metals in the article, so bromine doesn't count. Double sharp (talk) 23:54, 20 April 2017 (UTC)

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"Einstein's proof that the speed of light is the most constant dimension in the universe"
Should "Einstein's proof that the speed of light is the most constant dimension in the universe" be revised?
 * 1) Since the speed of light changes with medium (Jean Foucault 1850) is it a constant?
 * 2) Is the speed of light properly called a "dimension" like height, length, width, & time?
 * 3) "in the universe?" Did Einstein explore every corner of the universe to "prove" this?
 * 4) Proof? Does this statement confuse "the best known explanation based on evidence" with proof? (PeacePeace (talk) 15:27, 21 September 2017 (UTC))

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Colour
I think that the colour of caesium is due to relativistic quantum chemistry, but this is not mentioned in the article. Axl ¤ [Talk] 09:32, 6 April 2018 (UTC)
 * Relativistic effects are fairly small at Z = 55 (only a little earlier Ag is silver as would have been expected without including relativity at all). A better explanation is that the plasmonic wavelength of Cs creeps into the violet end of the visible spectrum (it is in the ultraviolet for most metals) and hence some violet light is absorbed, giving a yellowish colour. A weaker effect may be present for Ca, Sr, Ba, Eu, and Yb, which according to Greenwood and Earnshaw are pale yellow. Double sharp (talk) 09:56, 6 April 2018 (UTC)
 * Would you add this information to the article, please? Axl ¤ [Talk] 08:39, 7 April 2018 (UTC)
 * I've added it. This incidentally makes me wonder if anyone has tried to calculate what the colour of francium might be, because then relativistic effects should become important. If the transition is to the lowest excited state (which is p1 instead of s1), then relativistic effects widen the 7s/7p gap enough that the energy required to excite the electron goes back up, so that the value of Fr is close to the value for Rb. This bit of OR suggests that francium should be silvery, but unfortunately I can't find an actual source for this (and according to the ref I added the exact electron transition that creates the golden colour of Cs was not known at the time, so it is probably not this simple). Double sharp (talk) 16:59, 7 April 2018 (UTC)
 * Thank you for adding the statement. Could you perhaps also include the reason why violet (and not ultraviolet) light is absorbed? I shall leave the investigation of francium to you. :-) <b style="color:#808000">Axl</b> ¤ <small style="color:#808000">[Talk] 20:58, 7 April 2018 (UTC)
 * Ultraviolet is absorbed too. Ultraviolet and violet are below the plasmonic wavelength of Cs and so for those wavelengths Cs preferentially transmits light rather than reflects it. The transmission is not perfect so you will not see transparency unless you prepare a very thin film of the metal, with a thickness of a few atoms. ^_^ Double sharp (talk) 10:12, 8 April 2018 (UTC)
 * I have expanded the explanation, in the process correcting my awful mistake in typing it up (the lower-frequency colours are preferentially reflected and not absorbed). T_T Double sharp (talk) 15:24, 8 April 2018 (UTC)
 * Thank you. <b style="color:#808000">Axl</b> ¤ <small style="color:#808000">[Talk] 17:10, 8 April 2018 (UTC)

Why could it have -1?
It is the least electronegative... So why??? Alfa-ketosav (talk) 15:59, 7 July 2018 (UTC)
 * Just because an atom or element itself might be the least electronegative among the elements doesn't mean there could not be other less-electronegative combinations or ways of forcing electrons onto it. The electronegativity is not "zero". See the cited ref to learn about the chemicals that give this result. DMacks (talk) 16:09, 7 July 2018 (UTC)
 * Says nothing about which compound... I know the EN is always positive, and I see the cited ref about ... and even Fr is more electronegative. Alfa-ketosav (talk) 18:05, 7 July 2018 (UTC)
 * The ref (10.1002/anie.197905871) seems quite likely to tell you examples of such chemicals. And refs to more articles about them. Somewhere along that chain will definitely be exact procedures for how thay are made. The "why" of your question doesn't make sense..."because in the product that someone made, there are 56 formal electrons assigned to the caesium atom." If you mean "how", then you'll need to follow the cited ref, and refs it cites, and so on... A chemical need not be stable in order to exist. It could be "metastable"...stuck in a high-energy way with no direct way to resolve the problems. Electron configuration can be changed by a variety of factors. Here (10.1038/ncomms5861) is using pressure to raise the energy of lithium's valence electron, giving it the ability to donate to caesium. Possibly up to –2(!). Here (10.1002/rcm.4913) is another approach, using decomposition of carboxylate salts. There is (as I suggested earlier) a more unstable electron donor than a simple "less electronegative" atom. DMacks (talk) 19:58, 7 July 2018 (UTC)
 * Actually wait, that first ref might actually answer your question as asked: "because under high pressure, Li becomes a better electron donor and Cs does not, so a Li–Cs bond becomes polarized as Liδ+–Csδ-". DMacks (talk) 20:33, 7 July 2018 (UTC)
 * Alkalides are known, such as [Cs(cryptand-222)]+•Cs-. Burzuchius (talk) 20:28, 23 July 2018 (UTC)

English or Oxford English
The talk page states this article is written in English with Oxford/IUPAC Spellings, but the edit notice (when editing) states it is written in English with -ise etc., anyone know why this is happening? Shadowssettle(talk) 08:40, 1 April 2020 (UTC)

Crystal structure
As I understand, this metal is amorphous at ambient temperature. The crystal structure shown in the info box applies to extreme low temperatures (T ≤ 78 K). This is not made clear in the article. Simon de Danser (talk) 22:02, 3 October 2020 (UTC)

Semi-protected edit request on 7 October 2020
A-Venturi Sebastiano (talk) 15:49, 7 October 2020 (UTC)
 * Red question icon with gradient background.svg Not done: it's not clear what changes you want to be made. Please mention the specific changes in a "change X to Y" format and provide a reliable source if appropriate. William Avery (talk) 15:58, 7 October 2020 (UTC)

caesium

Semi-protected edit request on 7 October 2020 (2)
A-Venturi Sebastiano (talk) 16:05, 7 October 2020 (UTC)
 * Red question icon with gradient background.svg Not done: it's not clear what changes you want to be made. Please mention the specific changes in a "change X to Y" format and provide a reliable source if appropriate. Eggishorn (talk) (contrib) 22:03, 8 October 2020 (UTC)

Biology
Cesium is absorbed by animal and plant cells competitively with potassium, but cesium has no any known beneficial function; however, at high concentrations, can cause toxicity in plants, seen as growth inhibition. In fact, the mammalian organisms have, during evolution, began to distinguish useless non-radioactive cesium from potassium, which is essential in the Na + / K + pump of animal cell membranes. This is clearly visible in the poor uptake and selectivity for cesium in liver and fetuses in the Nelson's autoradiographies .. The human organism expels the useless Cesium through two emunctories: the salivary glands and the exocrine pancreas, which filter, and eliminate cesium with their secretions (saliva and pancreatic  juice) in the intestine. In fact, the "Prussian Blue" (ferric ferrocyanide) in the intestine is able to chelate cesium, thus preventing its reabsorption, and to eliminate it in the faeces.