Radium chloride

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Radium chloride
Radium Chloride
Identifiers
10025-66-8 N
ChemSpider 20138060 YesY
Jmol 3D model Interactive image
  • InChI=1S/2ClH.Ra/h2*1H;/q;;+2/p-2 YesY
    Key: RWRDJVNMSZYMDV-UHFFFAOYSA-L YesY
  • InChI=1/2ClH.Ra/h2*1H;/q;;+2/p-2
    Key: RWRDJVNMSZYMDV-NUQVWONBAG
  • [Ra+2].[Cl-].[Cl-]
Properties
RaCl2
Molar mass 296.094 g/mol
Appearance Colorless solid[1]
Density 4.9 g/cm3[1]
Melting point 900 °C (1,650 °F; 1,170 K)[1]
245 g/L (20 °C)[2]
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YesYN ?)
Infobox references

Radium chloride (RaCl2) is a chemical compound of radium and chlorine, and the first radium compound isolated in a pure state. Marie Curie and André-Louis Debierne used it in their original separation of radium from barium.[3] The first preparation of radium metal was by the electrolysis of a solution of this salt using a mercury cathode.[4]

Preparation

Radium chloride crystallises from solution as the dihydrate. It may be dehydrated by heating to 100 °C in air for one hour followed by ​5 1⁄2 hours at 520 °C under argon.[5] If the presence of other anions is suspected, the dehydration may be effectuated by fusion under hydrogen chloride.[6]

Radium chloride can also be prepared by heating radium bromide in a flow of dry hydrogen chloride gas, or by dehydrating radium sulfate with dry air and then heating the sulfate in a stream of hydrogen chloride.[1]

Properties

Radium chloride is a colorless-white salt with a blue-green luminescence, especially when heated. Its color gradually changes to yellow with aging, whereas contamination by barium may impart a rose tint.[1] It is less soluble in water than other alkaline earth metal chlorides – at 25 °C its solubility is 245 g/L whereas that of barium chloride is 307 g/L, and the difference is even larger in hydrochloric acid solutions. This property is used in the first stages of the separation of radium from barium by fractional crystallization.[2] Radium chloride is only sparingly soluble in azeotropic hydrochloric acid and virtually insoluble in concentrated hydrochloric acid.[7]

Gaseous radium chloride exists as RaCl2 molecules, as with other alkaline earth metal halides. The gas shows strong absorptions in the visible spectrum at 676.3 nm and 649.8 nm (red): the dissociation energy of the radium–chlorine bond is estimated as 2.9 eV,[8] and its length as 292 pm.[9]

Contrary to diamagnetic barium chloride, radium chloride is weakly paramagnetic with a magnetic susceptibility of 1.05×106. It also differs from barium chloride by the flame color, which is red as opposed to green for barium chloride.[1]

Uses

Radium chloride is still used for the initial stages of the separation of radium from barium during the extraction of radium from pitchblende. The large quantities of material involved (tonnes of ore for milligrams of radium) favour this less costly (but less efficient) method over those based on radium bromide or radium chromate (used for the later stages of the separation).

It is also used in medicine to produce radon gas which in turn is used as a cancer treatment.[citation needed]

Radium-223 dichloride (USP, radium chloride Ra 223), tradename Xofigo (formerly Alpharadin), is an alpha-emitting radiopharmaceutical. Bayer received FDA approval for this drug to treat prostate cancer osteoblastic bone metastases in May 2013. Radium chloride Ra 223 is one of the most potent drugs known.[citation needed] One dose (50 kBq/kg in an adult is about 60 nanograms; this amount is 1/1000 the weight of an eyelash (75 micrograms).

References

  1. ↑ 1.0 1.1 1.2 1.3 1.4 1.5 Kirby, p. 5
  2. ↑ 2.0 2.1 Kirby, p. 6
  3. ↑ Curie, M.; Debierne, A. (1910). C. R. Hebd. Acad. Sci. Paris 151:523–25.
  4. ↑ Kirby, p. 3
  5. ↑ Lua error in package.lua at line 80: module 'strict' not found.
  6. ↑ Lua error in package.lua at line 80: module 'strict' not found.
  7. ↑ Lua error in package.lua at line 80: module 'strict' not found.
  8. ↑ Lagerqvist, A. (1953). Arkiv Fisik 6:141–42.
  9. ↑ Karapet'yants, M. Kh.; Ch'ing, Ling-T'ing (1960). Zh. Strukt. Khim. 1:277–85; J. Struct. Chem. (USSR) 1:255–63.

Bibliography

Sources

  • Gmelins Handbuch der anorganischen Chemie (8. Aufl.), Berlin:Verlag Chemie, 1928, pp. 60–61.
  • Gmelin Handbuch der anorganischen Chemie (8. Aufl. 2. Erg.-Bd.), Berlin:Springer, 1977, pp. 362–64.
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