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Lead(II) perchlorate

Lead(II) perchlorate
Trihydrate
Names
Other names
  • Plumbous perchlorate
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.033.736 Edit this at Wikidata
EC Number
  • 237-125-7
UN number 1470
  • InChI=1S/2ClHO4.Pb/c2*2-1(3,4)5;/h2*(H,2,3,4,5);/q;;+2/p-2
  • [O-]Cl(=O)(=O)=O.[O-]Cl(=O)(=O)=O.[Pb+2]
Properties
Pb(ClO4)2
Molar mass 406.10 g/mol
Appearance White solid
Density 2.6 g/cm3
Boiling point 250 °C (482 °F; 523 K) (decomposes)
256.2 g/100 ml (25 °C)
Vapor pressure 0.36 Torr (trihydrate)
Hazards
GHS labelling:
GHS03: OxidizingGHS07: Exclamation markGHS06: ToxicGHS09: Environmental hazard
H272, H302, H332, H360Df, H373, H410
P210, P260, P273, P301+P312, P304+P340, P308+P313
Related compounds
Other cations
 Mercury(II) perchlorate; Tin(II) perchlorate; Cadmium perchlorate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Lead(II) perchlorate is a chemical compound with the formula Pb(ClO4)2·xH2O, where is x is 0,1, or 3. It is an extremely hygroscopic white solid that is very soluble in water.[1]

Preparation

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Lead perchlorate trihydrate is produced by the reaction of lead(II) oxide, lead carbonate, or lead nitrate by perchloric acid:

Pb(NO3)2 + HClO4 → Pb(ClO4)2 + HNO3

The excess perchloric acid was removed by first heating the solution to 125 °C, then heating it under moist air at 160 °C to remove the perchloric acid by converting the acid to the dihydrate. The anhydrous salt, Pb(ClO4)2, is produced by heating the trihydrate to 120 °C under water-free conditions over phosphorus pentoxide. The trihydrate melts at 83 °C.[1] The anhydrous salt decomposes into lead(II) chloride and a mixture of lead oxides at 250 °C.[1][2] The monohydrate is produced by only partially dehydrating the trihydrate, and this salt undergoes hydrolysis at 103 °C.[3]

The solution of anhydrous lead(II) perchlorate in methanol is explosive.[1]

Applications

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Lead perchlorate has a high nucleon density, making it a viable detector for hypothetical proton decay.[4]

References

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  1. ^ a b c d H. H. Willard; J. L. Kassner (1930). "PREPARATION AND PROPERTIES OF LEAD PERCHLORATE". Journal of the American Chemical Society. 52 (6). ACS Publications: 2391–2396. doi:10.1021/ja01369a027.
  2. ^ Zinov'ev, A. A.; and Kritsov, N. V. (1960). Zhur. Neorg. Khim. issue 5: p. 1418, as cited in Giridharan, A. S.; Udupa, M. R.; Aravamudan, G. (February 1975). "Thermal behaviour of thallous perchlorate". Journal of Thermal Analysis. 7 (1): 65–71. doi:10.1007/BF01911626. ISSN 0022-5215.
  3. ^ A. V. Dudin (1993). "Water-vapor pressure and thermodynamics of the dehydration of manganese, nickel, cadmium, and lead perchlorate hydrates". Russian Chemical Bulletin. 42: 417–421. doi:10.1007/BF00698419.
  4. ^ Boyd, R. N.; Rauscher, T.; Reitzner, S. D.; Vogel, P. (2003-10-31). "Observing nucleon decay in lead perchlorate". Physical Review D. 68 (7). arXiv:hep-ph/0307280. doi:10.1103/PhysRevD.68.074014. ISSN 0556-2821.