Xenon trioxide

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Xenon trioxide
Structural formula, showing lone pair
Space-filling model
Identifiers
CAS number 13776-58-4 N
ChemSpider 21106493 N
Jmol-3D images Image 1
Properties
Molecular formula XeO3
Molar mass 179.288 g/mol
Appearance colourless crystalline solid
Density 4.55 g/cm3, solid
Melting point 25 °C (77 °F; 298 K) Violent decomposition
Solubility in water Soluble (with reaction)
Structure
Molecular shape trigonal pyramidal (C3v)
Thermochemistry
Std enthalpy of
formation
ΔfHo298
402 kJ·mol−1[1]
Hazards
EU classification not listed
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 4: Very short exposure could cause death or major residual injury. E.g., VX gas Reactivity code 4: Readily capable of detonation or explosive decomposition at normal temperatures and pressures. E.g., nitroglycerin Special hazard OX: Oxidizer. E.g., potassium perchlorateNFPA 704 four-colored diamond
Related compounds
Related compounds Xenon tetroxide
Xenic acid
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
 N (verify) (what is: YesY/N?)
Infobox references

Xenon trioxide is an unstable compound of xenon in its +6 oxidation state. It is a very powerful oxidizing agent, and liberates oxygen from water slowly (and xenon), accelerated by exposure to sunlight. It is dangerously explosive upon contact with organic materials. When it detonates, it releases xenon and oxygen gas.

Chemistry[edit]

Xenon trioxide is a strong oxidising agent and can oxidise most substances that are at all oxidisable. However, it is slow-acting and this reduces its usefulness.[2]

Above 25 °C, xenon trioxide is very prone to violent explosion:

2 XeO3 → 2 Xe + 3 O2

When it dissolves in water, an acidic solution of xenic acid is formed:

XeO3 (aq) + H2O → H2XeO4 is in equilibrium with H+ + HXeO4

This solution is stable at room temperature and lacks the explosive properties of xenon trioxide. It oxidises carboxylic acids quantitatively to carbon dioxide and water.[3]

Alternatively, it dissolves in alkaline solutions to form xenates. The HXeO
4
anion is the predominant species in xenate solutions.[4] These are not stable and begin to disproportionate into perxenates (+8 oxidation state) and xenon and oxygen gas.[5] Solid perxenates containing XeO4−
6
have been isolated by reacting XeO
3
with an aqueous solution of hydroxides. Xenon trioxide reacts with inorganic fluorides such as KF, RbF, or CsF to form stable solids of the form MXeO
3
F
.[6]

Physical properties[edit]

Hydrolysis of xenon hexafluoride or xenon tetrafluoride yields a solution from which colorless XeO3 crystals can be obtained by evaporation.[7] The crystals are stable for days in dry air, but readily absorb water from humid air to form a concentrated solution. The crystal structure is orthorhombic with a = 6.163, b = 8.115, c = 5.234 Å and 4 molecules per unit cell. The density is 4.55 g/cm3.[8]

Xenon-trioxide-xtal-1963-3D-balls.png
Xenon-trioxide-xtal-1963-3D-SF.png
Xenon-trioxide-xtal-1963-Xe-coordination-3D-balls.png
ball-and-stick model of part of
the crystal structure of XeO3
space-filling model
coordination geometry of Xe

Safety[edit]

XeO3 should be handled with great caution. Samples have detonated when undisturbed at room temperature. Dry crystals react explosively with cellulose.[8][9]

References[edit]

  1. ^ Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A23. ISBN 0-618-94690-X. 
  2. ^ Greenwood, N.; Earnshaw, A. (1997). Chemistry of the Elements. Oxford: Butterworth-Heinemann. 
  3. ^ Jaselskis B.; Krueger R. H. (July 1966). "Titrimetric determination of some organic acids by xenon trioxide oxidation". Talanta 13 (7): 945–949. doi:10.1016/0039-9140(66)80192-3. PMID 18959958. 
  4. ^ Peterson, J. L.; Claassen, H. H.; Appelman, E. H. (March 1970). "Vibrational spectra and structures of xenate(VI) and perxenate(VIII) ions in aqueous solution". Inorganic Chemistry 9 (3): 619–621. doi:10.1021/ic50085a037.  edit
  5. ^ W. Henderson (2000). Main group chemistry. Great Britain: Royal Society of Chemistry. pp. 152–153. ISBN 0-85404-617-8. 
  6. ^ Egon Wiberg; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. p. 399. ISBN 0-12-352651-5. 
  7. ^ John H. Holloway; Eric G. Hope (1998). A. G. Sykes, ed. Recent Advances in Noble-gas Chemistry. Advances in Inorganic Chemistry, Volume 46. Academic Press. p. 65. ISBN 0-12-023646-X. 
  8. ^ a b Templeton, D. H.; Zalkin, A.; Forrester, J. D.; Williamson, S. M. (1963). Journal of the American Chemical Society 85 (6): 817. doi:10.1021/ja00889a037.  edit
  9. ^ Bartlett, N.; Rao, P. R. (1963). "Xenon Hydroxide: an Experimental Hazard". Science 139 (3554): 506. Bibcode:1963Sci...139..506B. doi:10.1126/science.139.3554.506. PMID 17843880.  edit

External links[edit]