Caesium fluoride: Difference between revisions

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+{{pp-move}}
-{{chembox
+{{Chembox
-| verifiedrevid = 396305984
+| Verifiedfields = changed
-| Name = Caesium fluoride
+| Watchedfields  = changed
-| ImageFile = Caesium fluoride.jpg
+| verifiedrevid  = 444641443
-| ImageName = Caesium fluoride
-| ImageFile1 = Caesium-fluoride-3D-ionic.png
+| Name           = Caesium fluoride
-| ImageName1 = Caesium fluoride
+| ImageFile      = Caesium fluoride.jpg
-| IUPACName = Caesium fluoride
+| ImageName      = Caesium fluoride
-| OtherNames = Cesium fluoride
+| ImageFile1     = Caesium-fluoride-3D-ionic.png
+| ImageName1     = Caesium fluoride
-| Section1 = {{Chembox Identifiers
+| IUPACName      = Caesium fluoride
+| OtherNames     = Cesium fluoride
+| Section1       = {{Chembox Identifiers
 | ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
 | ChemSpiderID = 24179
@@ Line 21: / Line 24: @@
 | CASNo_Ref = {{cascite|correct|CAS}}
 | RTECS = FK9650000
+| PubChem = 25953
+| EC_number = 236-487-3
+| UNII = T76A371HJR
 }}
-| Section2 = {{Chembox Properties
+| Section2       = {{Chembox Properties
 | Formula = CsF
-| MolarMass = 151.90 g/mol
+| MolarMass = 151.903 g/mol<ref name=crc/>
 | Appearance = white crystalline solid
-| Density = 4.115 g/cm<sup>3</sup>
+| Density = 4.64 g/cm<sup>3</sup><ref name=crc/>
-| Solubility = 367 g/100 ml (18 °C)
+| Solubility = 573.0 g/100 mL (25 °C)<ref name=crc>{{RubberBible92nd|page=4.57}}</ref>
+| SolubleOther = Insoluble in [[acetone]], [[diethyl ether]], [[pyridine]] and [[ethanol]]<br>191 g/100 mL in [[methanol]].
-| MeltingPt = 682 °C (955 K)
+| MeltingPtC = 703
-| BoilingPt = 1251 °C (1524 K)
+| MeltingPt_ref = <ref name=crc/>
-| pKb =
+| BoilingPtC = 1251
+| BoilingPt_notes = (2,284 °F; 1,524 K)
+| pKb = −744 kJ/mol
+| RefractIndex = 1.477
+| MagSus = -44.5·10<sup>−6</sup> cm<sup>3</sup>/mol<ref>{{RubberBible92nd|page=4.132}}</ref>
 }}
-| Section3 = {{Chembox Structure
+| Section3       = {{Chembox Structure
 | Coordination = [[Octahedron|Octahedral]]
 | CrystalStruct = [[Cubic crystal system|cubic]],  [[Pearson symbol|cF8]]
+| SpaceGroup = Fm{{overline|3}}m, No. 225<ref name=str>{{cite journal|doi=10.1103/PhysRev.21.143|title=Precision Measurements of Crystals of the Alkali Halides|journal=Physical Review|volume=21|issue=2|pages=143–161|year=1923|last1=Davey|first1=Wheeler P.|bibcode=1923PhRv...21..143D}}</ref>
-|   SpaceGroup = Fm<u style="text-decoration:overline">3</u>m, No. 225
+| LattConst_a =  0.6008 nm<ref name=str/>
+| UnitCellFormulas = 4
+| UnitCellVolume = 0.2169 nm<sup>3</sup><ref name=str/>
 | Dipole = 7.9 [[Debye|D]]
 }}
-| Section7 = {{Chembox Hazards
+| Section4       = {{Chembox Thermochemistry
+| DeltaHf = -553.5 kJ/mol<ref name=b92t>{{RubberBible92nd|page=5.10}}</ref>
-| ExternalMSDS = [https://fscimage.fishersci.com/msds/13612.htm External MSDS]
+| Entropy = 92.8 J/mol·K<ref name=b92t/>
-| EUIndex = Not listed
+| DeltaGf = -525.5 kJ/mol<ref name=b92t/>
-| RPhrases =
+| HeatCapacity = 51.1 J/mol·K<ref name=b92t/>
-| SPhrases =
+}}
+| Section7       = {{Chembox Hazards
+| MainHazards = toxic
+| GHSPictograms = {{GHS05}}{{GHS06}}{{GHS08}}
+| GHSSignalWord = Danger
+| HPhrases = {{H-phrases|301|311|315|318|331|361f}}
+| PPhrases = {{P-phrases|201|202|260|261|264|270|271|280|281|301+310|301+330+331|302+352|303+361+353|304+340|305+351+338|308+313|310|311|312|321|322|330|332+313|361|362|363|403+233|405|501}}
+| NFPA-H = 3
+| NFPA-F = 0
+| NFPA-R = 0
+| ExternalSDS = [https://www.fishersci.com/store/msds?partNumber=AC189510250&productDescription=CESIUM+FLUORIDE%2C+99%25+25GR&vendorId=VN00032119&countryCode=US&language=en External MSDS]
 | FlashPt = Non-flammable
 }}
-| Section8 = {{Chembox Related
+| Section8       = {{Chembox Related
-| OtherAnions = [[Caesium chloride]]<br/>[[Caesium bromide]]<br/>[[Caesium iodide]]
+| OtherAnions = [[Caesium chloride]]<br/>[[Caesium bromide]]<br/>[[Caesium iodide]]<br/>Caesium astatide
-| OtherCations = [[Lithium fluoride]]<br/>[[Sodium fluoride]]<br/>[[Potassium fluoride]]<br/>[[Rubidium fluoride]]
+| OtherCations = [[Lithium fluoride]]<br/>[[Sodium fluoride]]<br/>[[Potassium fluoride]]<br/>[[Rubidium fluoride]]<br/>Francium fluoride
 }}
 }}
-'''Caesium fluoride''' ('''cesium fluoride''' in North America), is an [[inorganic compound]] usually encountered as a hygroscopic white solid. It is more soluble and more readily [[Dissociation (chemistry)|dissociated]] than [[sodium fluoride]] or [[potassium fluoride]]. It is available in anhydrous form, and if water has been absorbed it is easy to dry by heating at 100&nbsp;°C for two hours ''[[:wikt:in vacuo|in vacuo]]''.<ref name="Friestad">{{cite book|author=Friestad, G. K.; Branchaud, B. P.|title=Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents|editor=Reich, H. J.; Rigby, J. H.|publisher=Wiley|location=New York|year=1999|pages=99–103}}</ref>  Like all soluble fluorides, it is mildly [[base (chemistry)|basic]]. A notable fact about this compound is that it is the most ionic compound. Caesium has the lowest [[electronegativity]] and fluorine has the highest electronegativity.
+'''Caesium fluoride''' ('''cesium fluoride''' in [[American English]]) is an [[inorganic compound]] with the formula CsF. A [[hygroscopic]] white salt, caesium fluoride is used in the [[organic synthesis|synthesis of organic compounds]] as a source of the fluoride anion.<ref>{{cite book |doi=10.1002/047084289X.rc050.pub2 |chapter=Cesium Fluoride |title=Encyclopedia of Reagents for Organic Synthesis |date=2007 |last1=Friestad |first1=Gregory K. |last2=Branchaud |first2=Bruce P. |last3=Navarrini |first3=Walter |last4=Sansotera |first4=Maurizio |isbn=978-0-471-93623-7 }}</ref>  The compound is noteworthy from the pedagogical perspective as [[caesium]] also has the highest [[electropositivity]] of all commonly available elements and [[fluorine]] has the highest electronegativity.
 ==Synthesis and properties==
+[[File:CsF@DWNT.png|thumb|left|150px|Crystalline CsF chains grown inside double-wall [[carbon nanotube]]s.<ref name=chains>{{cite journal|doi=10.1038/ncomms8943|pmid=26228378|pmc=4532884|title=Single-atom electron energy loss spectroscopy of light elements|journal=Nature Communications|volume=6|pages=7943|year=2015|last1=Senga|first1=Ryosuke|last2=Suenaga|first2=Kazu|bibcode=2015NatCo...6.7943S}} (Supplementary information)</ref>]]
-Caesium fluoride is prepared by the action of [[hydrofluoric acid]] on [[caesium hydroxide]] or [[caesium carbonate]], followed by removal of water.
+Caesium fluoride can be prepared by the reaction of [[caesium hydroxide]] (CsOH) with [[hydrofluoric acid]] (HF) and the resulting salt can then be purified by recrystallization. The reaction is shown below:
+:CsOH + HF → CsF + H<sub>2</sub>O
-Caesium fluoride reacts usually as a source of fluoride ion, F<sup>-</sup>. It therefore undergoes all of the usual reactions associated with soluble fluorides, for example:<ref name="greenwood">Greenwood, N.N.; Earnshaw, A. ''Chemistry of the Elements'', Pergamon Press, Oxford, UK, 1984.</ref>
+Using the same reaction, another way to create caesium fluoride is to treat [[caesium carbonate]] (Cs<sub>2</sub>CO<sub>3</sub>) with hydrofluoric acid and again, the resulting salt can then be purified by recrystallization. The reaction is shown below:
-:2 CsF + CaCl<sub>2</sub>  → 2 CsCl  + CaF<sub>2</sub>
+:Cs<sub>2</sub>CO<sub>3</sub> + 2 HF → 2 CsF + H<sub>2</sub>O + CO<sub>2</sub>
+CsF is more soluble than [[sodium fluoride]] or [[potassium fluoride]] in organic solvents. It is available in its anhydrous form, and if water has been absorbed, it is easy to dry by heating at 100&nbsp;°C for two hours ''[[in vacuo]]''.<ref name="Friestad">{{cite book|author1=Friestad, G. K. |author2=Branchaud, B. P. |title=Handbook of Reagents for Organic Synthesis: Acidic and Basic Reagents|editor1=Reich, H. J. |editor2=Rigby, J. H. |publisher=Wiley|location=New York|year=1999|pages=99–103|isbn=978-0-471-97925-8}}</ref> CsF reaches a [[vapor pressure]] of 1 [[kilopascal]] at 825&nbsp;°C, 10 kPa at 999&nbsp;°C, and 100 kPa at 1249&nbsp;°C.<ref>{{cite book | editor = Lide, D. R. | title = CRC Handbook of Chemistry and Physics | edition = 86th | location = Boca Raton (FL) | publisher = CRC Press | year = 2005 | isbn = 0-8493-0486-5 |page=6.63|chapter=Vapor Pressure|chapter-url=https://www.physics.nyu.edu/kentlab/How_to/ChemicalInfo/VaporPressure/morepressure.pdf}}</ref>
-==Crystal structure==
-Caesium fluoride has the halite structure, which means that the Cs<sup>+</sup> and F<sup>&minus;</sup> pack in a [[cubic closest packed]] array as do Na<sup>+</sup> and Cl<sup>&minus;</sup> in sodium chloride. Caesium cations are larger than fluoride anions, whereas in the lithium, sodium, potassium, and rubidium halides, the cations are smaller than the anion. <ref name="greenwood"/><ref name="CRC">''Handbook of Chemistry and Physics'', 71st edition, CRC Press, Ann Arbor, Michigan, 1990.</ref>
-==Applications==
+==Structure==
+Caesium fluoride has the [[halite]] structure, which means that the Cs<sup>+</sup> and F<sup>&minus;</sup> pack in a [[cubic closest packed]] array as do Na<sup>+</sup> and Cl<sup>&minus;</sup> in sodium chloride.<ref name=str/> Unlike sodium chloride, caesium fluoride's anion is smaller than its cation, so it is the anion size that sterically inhibits larger coordination numbers than six under normally encountered conditions. A larger halide ion would allow for the eight-coordination seen in other caesium halide crystals.
-===In organic synthesis===
-Being highly dissociated it is a more reactive source of fluoride than related salts.  CsF is less [[hygroscopic]] alternative to [[tetra-n-butylammonium fluoride]] (TBAF) and [[TASF reagent|TAS-fluoride]] (TASF) when anhydrous "naked" [[fluoride]] [[ion]] is needed.
+==Applications in organic synthesis==
-====As a base====
+Being highly dissociated, CsF is a more reactive source of fluoride than related alkali metal salts. CsF is an alternative to [[tetra-n-butylammonium fluoride]] (TBAF) and [[TASF reagent|TAS-fluoride]] (TASF).
-As with other soluble fluorides, CsF is moderately basic, because [[hydrofluoric acid|HF]] is a weak acid. The low [[nucleophile|nucleophilicity]] of fluoride means it can be a useful base in organic chemistry.<ref name="greenwood"/>Caesium fluoride is a useful base in [[organic chemistry]], due the fact that fluoride [[ion]] is a relatively poor [[nucleophile]]. CsF gives higher yields in [[Knoevenagel condensation]] reactions than [[potassium fluoride|KF]] or [[sodium fluoride|NaF]].<ref name="Fiorenza">{{cite journal|year=1985|journal=Tetrahedron Letters|volume=26|page=787|doi=10.1016/S0040-4039(00)89137-6|title=Fluoride ion induced reactions of organosilanes: the preparation of mono and dicarbonyl compounds from β-ketosilanes|author1=Fiorenza, M|author2=Mordini, A|author3=Papaleo, S|author4=Pastorelli, S|author5=Ricci, A}}</ref>
-====Formation of C-F bonds====
+===As a base===
+As with other soluble fluorides, CsF is moderately basic, because [[hydrofluoric acid|HF]] is a [[weak acid]]. The low [[nucleophile|nucleophilicity]] of [[fluoride]] means it can be a useful base in [[organic chemistry]].<ref name="greenwood">{{Greenwood&Earnshaw1st|pages=82–83}}</ref> CsF gives higher yields in [[Knoevenagel condensation]] reactions than [[potassium fluoride|KF]] or [[sodium fluoride|NaF]].<ref name="Fiorenza">{{cite journal|year=1985|journal=Tetrahedron Letters|volume=26|pages=787–788|doi=10.1016/S0040-4039(00)89137-6|title=Fluoride ion induced reactions of organosilanes: the preparation of mono and dicarbonyl compounds from β-ketosilanes|author1=Fiorenza, M|author2=Mordini, A|author3=Papaleo, S|author4=Pastorelli, S|author5=Ricci, A|issue=6}}</ref>
-Caesium fluoride is also a popular source of fluoride in [[organofluorine chemistry]]. For example, CsF reacts with [[hexafluoroacetone]] to form a caesium perfluoroalkoxide salt, which is stable up to 60 °C, unlike the corresponding [[sodium]] or [[potassium]] salt. It will convert [[electron-deficient]] [[aryl chloride]]s to [[aromatic|aryl]] fluorides ([[halex reaction]]).<ref>{{cite journal|author=F. W. Evans, M. H. Litt, A. M. Weidler-Kubanek, F. P. Avonda|title=Reactions Catalyzed by Potassium Fluoride. 111. The Knoevenagel Reaction|year=1968|journal=Journal of Organic Chemistry|volume=33|pages=1837–1839|doi=10.1021/jo01269a028}}</ref>
-====Deprotection agent====
+===Formation of Cs-F bonds===
+Caesium fluoride serves as a source of fluoride in [[organofluorine chemistry]]. Similarly to [[potassium]] fluoride, CsF reacts with [[hexafluoroacetone]] to form a stable perfluoroalkoxide salt.<ref>{{cite journal|author1=Evans, F. W. |author2=Litt, M. H. |author3=Weidler-Kubanek, A. M. |author4=Avonda, F. P. |title=Formation of adducts between fluorinated ketones and metal fluorides|year=1968|journal=Journal of Organic Chemistry|volume=33|pages=1837–1839|doi=10.1021/jo01269a028|issue=5}}</ref> It will convert [[electron-deficient]] [[aryl chloride]]s to [[aromatic|aryl]] fluorides ([[Halex process]]), although potassium fluoride is more commonly used.
-Due to the strength of the [[silicon|Si]]–[[fluorine|F]] bond, fluoride ion is useful for [[desilylation]] reactions (removal of Si groups) in [[organic chemistry]]; caesium fluoride is an excellent source of anhydrous fluoride for such reactions.  Removal of silicon groups ([[desilylation]]) is a major application for CsF in the laboratory, as its [[anhydrous]] nature allows clean formation of [[water (molecule)|water]]-sensitive intermediates. Solutions of caesium fluoride in [[THF]] or [[Dimethylformamide|DMF]] attack a wide variety of [[organosilicon]] compounds to produce an organosilicon fluoride and a [[carbanion]], which can then react with [[electrophile]]s,<ref name="CRC"/> for example:<ref name="Fiorenza"/>
+===Deprotection agent===
+Due to the strength of the [[silicon|Si]]–[[fluorine|F]] bond, fluoride is useful for [[desilylation]] reactions, i.e., cleavage of Si-O bonds in [[organic synthesis]].<ref>{{OrgSynth |author=Smith, Adam P. |author2=Lamba, Jaydeep J. S. |author3=Fraser, Cassandra L. |name-list-style=amp |title=Efficient Synthesis of Halomethyl-2,2'-bipyridines: 4,4'-Bis(chloromethyl)-2,2'-bipyridine|volume=78|page=82 |year=2002|collvol= 10|collvolpages=107 |prep=v78p0082}}</ref> CsF is commonly used for such reactions. Solutions of caesium fluoride in [[THF]] or [[Dimethylformamide|DMF]] attack a wide variety of [[organosilicon]] compounds to produce an organosilicon fluoride and a [[carbanion]], which can then react with [[electrophile]]s, for example:<ref name="Fiorenza"/>
 :[[Image:CsF desilylation.png|500px]]
-Desilylation is also useful for the removal of [[silyl]] [[protecting group]]s.<ref>Adam P. Smith, Jaydeep J. S. Lamba, and Cassandra L. Fraser “Efficient Synthesis of Halomethyl-2,2'-bipyridines: 4,4'-Bis(chloromethyl)-2,2'-bipyridine” Organic Syntheses, Vol. 78, p. 82 (2002); Collected Volume 10, p.107 (2004).</ref>
-===Other uses===
-Single crystals of the salt are transparent into the deep [[infrared]]. For this reason it is sometimes used as the windows of cells used for [[infrared spectroscopy]].
 ==Precautions==
-Like other soluble fluorides, CsF is moderately toxic.<ref name="msds-csf">[https://www.hazard.com/msds/f2/bms/bmsqc.html MSDS Listing for cesium fluoride]. ''[https://www.hazard.com/ www.hazard.com].'' MSDS Date: April 27, 1993. Retrieved on September 7, 2007.</ref> Contact with [[acid]] should be avoided, as this forms highly toxic/corrosive [[hydrofluoric acid]]. Caesium [[ion]] (Cs<sup>+</sup>), or caesium chloride, is generally not considered toxic.<ref name="msds-cscl">"[https://hazard.com/msds/mf/baker/baker/files/c1903.htm MSDS Listing for cesium chloride]." ''[https://www.jtbaker.com/ www.jtbaker.com].'' MSDS Date: January 16, 2006. Retrieved on September 7, 2007.</ref>
+Like other soluble fluorides, CsF is moderately toxic.<ref name="msds-csf">{{cite web |url=https://www.hazard.com/msds/f2/bms/bmsqc.html |title=MSDS Listing for cesium fluoride |date=April 27, 1993 |website=hazard.com |access-date=September 7, 2007 |url-status=usurped |archive-url=https://web.archive.org/web/20120209103840/https://www.hazard.com/msds/f2/bms/bmsqc.html |archive-date=2012-02-09 }}</ref> Contact with [[acid]] should be avoided, as this forms highly toxic/corrosive [[hydrofluoric acid]]. The caesium [[ion]] (Cs<sup>+</sup>) and [[caesium chloride]] are generally not considered toxic.<ref name="msds-cscl">{{cite web |url=https://hazard.com/msds/mf/baker/baker/files/c1903.htm |title=MSDS Listing for cesium chloride |date=January 16, 2006 |website=hazard.com |publisher=JT Baker |access-date=September 7, 2007 |url-status=usurped |archive-url=https://web.archive.org/web/20120313200630/https://hazard.com/msds/mf/baker/baker/files/c1903.htm |archive-date=2012-03-13 }}</ref>
 ==References==
 {{reflist}}
-==External links==
+{{Authority control}}
-*[https://www.npi.gov.au/database/substance-info/profiles/44.html National Pollutant Inventory-Fluoride and compounds fact sheet]
+{{Caesium compounds}}
+{{Fluorides}}
 [[Category:Fluorides]]
 [[Category:Caesium compounds]]
 [[Category:Metal halides]]
+[[Category:Alkali metal fluorides]]
+[[Category:Rock salt crystal structure]]
-{{good article}}
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