Barium perchlorate

Last updated
Barium perchlorate
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.033.359 OOjs UI icon edit-ltr-progressive.svg
EC Number
  • 236-710-4
PubChem CID
RTECS number
  • SC7550000
UNII
  • InChI=1S/Ba.2ClHO4/c;2*2-1(3,4)5/h;2*(H,2,3,4,5)/q+2;;/p-2
    Key: OOULUYZFLXDWDQ-UHFFFAOYSA-L
  • [O-]Cl(=O)(=O)=O.[O-]Cl(=O)(=O)=O.[Ba+2]
Properties
Ba(ClO4)2
Molar mass 336.22 g·mol−1
Appearancewhite powder
Density 3.2 g/cm3
Melting point 505 °C (941 °F; 778 K)
66.48 g/100 mL (25 °C (77 °F; 298 K))
Hazards
NFPA 704 (fire diamond)
NFPA 704.svgHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazard OX: Oxidizer. E.g. potassium perchlorate
3
0
1
OX
Related compounds
Other anions
Other cations
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Barium perchlorate is a powerful oxidizing agent, with the formula Ba(ClO4)2. It is used in the pyrotechnic industry.[ citation needed ]

Contents

Barium perchlorate decomposes at 505 °C (941 °F). [1]

Structure of barium perchlorate trihydrate

Gallucci and Gerkin (1988) analyzed the structure of the hydrate isomer barium perchlorate trihydrate (Ba(ClO4)2·3H2O) by X-ray crystallography. The barium ions are coordinated by six water oxygen atoms at 2.919 Å and six perchlorate oxygens at 3.026 Å in a distorted icosahedral arrangement. The perchlorate fails by a narrow margin to have regular tetrahedral geometry, and has an average Cl-O bond length of 1.433 Å. The space-group assignment of the structure was resolved, with the centrosymmetric assignment of P63/m confirmed. Each axial perchlorate oxygen is hydrogen bonded to three water molecules and each trigonal oxygen is hydrogen bonded to two water molecules. This interaction is the reason that the perchlorate fails to be tetrahedral. Gallucci and Gerkin surmised that the water molecule H atoms lie in the plane at z = ⁠1/4⁠ and ⁠3/4⁠. [2]

Preparation

Barium perchlorate can be prepared using many different reagents and methods. One method involves evaporating a solution containing barium chloride and an excess of perchloric acid. The dihydrate form is produced by recrystallizing and drying to a constant weight. Additional drying over sulfuric acid yields the monohydrate. The anhydrous form is obtained by heating to 140 °C (284 °F) in vacuum. [3] Dehydration of barium perchlorate that does not occur in vacuum will also result in hydrolysis of the perchlorate. [4] Other reactions that produce barium perchlorate are as follows: perchloric acid and barium hydroxide or carbonate; potassium perchlorate and hydrofluosilicic acid followed with barium carbonate; boiling solution of potassium chlorate and zinc fluosilicate. For large-scale manufacturing purposes, barium perchlorate is synthesized by evaporating a solution of sodium perchlorate and barium chloride. [3] Another method of preparation involves the digestion of a saturated solution of ammonium perchlorate with hydrated barium hydroxide in 5-10% excess of the theoretical amount. [5]

Applications

Due to its characteristic as a powerful oxidizing agent, one of barium perchlorate’s primary uses is in the manufacture and preparation of explosive emulsions and other explosive compounds.[ citation needed ] Using an emulsifier makes the process of transporting and handling of the explosive material while still retaining its destructive properties at the end point of use. Perchlorate explosives were mainly used in industrial applications, such as mining during the 1920s. [3] [ better source needed ]

Barium perchlorate is also able to complex with the fluoroquinolone antibacterial agents ciprofloxacin and norfloxacin. [6] FTIR data suggests that CIP and NOR act as bidentate ligands, using the ring carbonyl oxygen and an oxygen of the carboxylic group. This coordination is significant because it increases the solubility of the antibiotics in water and other polar solvents, increasing their uptake efficiency.[ citation needed ]

Because of its high solubility in water, anhydrous barium perchlorate can be used as a dehydrating reagent for other compounds. [3] Due to its high solubility, ease of preparation, low cost,[ citation needed ] stability at high temperatures, and relative ease of regeneration, barium perchlorate is a favored compound as a dehydrating agent.[ citation needed ] The need for dehydrating compounds has increased[ when? ] with the use of chemical reactions employing gases under pressure, as the water must be removed from the air[ which? ] prior to the reaction taking place. [7] [ better source needed ]

Barium perchlorate is also used for the determination of small concentrations (down to 10±1 ppm) of sulfate. [5] In order for the titration to be successful, a high concentration of a nonaqueous solvent, such as ethyl alcohol, 2-propanol, or methanol, must be present. Thorin is typically used as the indicator.[ needs update ]

References

  1. ↑ Haynes, William M. (2011-06-06). CRC Handbook of Chemistry and Physics: A Ready-Reference Book of Chemical and Physical data (92nd ed.). Boca Raton, Florida: CRC Press. ISBN   978-1-4398-5511-9.
  2. ↑ Gallucci, J.C.; Gerkin, R.E. (1988-11-01). "Structure of Barium Perchlorate Trihydrate". Acta Crystallographica Section C. 44 (11): 1873–1876. Bibcode:1988AcCrC..44.1873G. doi:10.1107/s0108270188008200. ISSN   0108-2701. PMID   2855929.
  3. 1 2 3 4 Warren, Francis (1960). Chlorates and perchlorates, their manufacture, properties and uses. Defense Technical Information Center. p. 147.
  4. ↑ Acheson, R.J.; Jacobs, P.W.M. (1969-08-15). "Thermal Decomposition of Barium Perchlorate". Canadian Journal of Chemistry. 47 (16): 3031–3039. doi:10.1139/v69-501. ISSN   0008-4042.
  5. 1 2 Fritz, K.S.; Yamamura, S.S. (September 1955). "Rapid Microtitration of Sulfate". Analytical Chemistry. 27 (9): 1461–1464. doi:10.1021/ac60105a030. ISSN   0003-2700.
  6. ↑ Serafin, A.; Stańczak, A. (2009-02-01). "The complexes of metal ions with fluoroquinolones". Russian Journal of Coordination Chemistry. 35 (2): 81–95. doi:10.1134/S1070328409020018. ISSN   1070-3284. S2CID   95087424.
  7. ↑ Smith, G. Frederick (March 1927). "Anhydrous Barium Perchlorate and Mixed Alkaline-Earth Metal Perchlorates as Dehydrating Reagents". Industrial & Engineering Chemistry. 19 (3): 411–414. doi:10.1021/ie50207a027. ISSN   0019-7866.
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