Perbromate

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Perbromate
Other names Tetraoxidobromate(1−)
Identifiers
InChI InChI=1/BrHO4/c2-1(3,4)5/h(H,2,3,4,5)/p-1
InChIKey LLYCMZGLHLKPPU-REWHXWOFAT
Standard InChI InChI=1S/BrHO4/c2-1(3,4)5/h(H,2,3,4,5)/p-1
Standard InChIKey LLYCMZGLHLKPPU-UHFFFAOYSA-M
CAS number [16474-32-1]
ChemSpider 4574125
Structure[1]
Molecular geometry tetrahedral: r(Br–O) = 161 pm
Thermochemistry[2][3]
Std enthalpy of formation ΔfHo298 +12.1 kJ mol−1 (aq)
Std Gibbs energy of formation ΔfGo298 +120.8 kJ mol−1 (aq)
Standard molar entropy So298 187 J K−1 mol−1 (aq)
Related compounds
Other anions Perchlorate
Periodate
Hypobromite
Bromite
Bromate
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)

A perbromate is a salt or ester of perbromic acid. Perbromates, including perbromic acid,[note 1] are the only well-characterized compounds of bromine in the +7 oxidation state.

Perbromates are thermodynamically more oxidizing than perchlorates or periodates, and eluded preparation for more than 100 years until 1968.[5][6][7] However, they are kinetically quite inert, especially in dilute solution.[6][7][8]

History

Although perchlorates and periodates have been known since the early nineteenth century, perbromates proved far more elusive.[9][10] The first claim of preparation came in 1863,[11] but proved unrepeatable.[12][13] Several later attempts to prepare perbromates were equally unsuccessful.[14][15] Some authors surmised that perbromates were fundamentally unstable.[16][17]

The first clear evidence of the existence of the perbromate anion came from "hot-atom" techniques. A sample of the β−-emittor selenium-83 (t½ = 25 min) was prepared in the form of the selenate anion SeO2−4 and allowed to decay into bromine-83. Approximately 14% of the 83Br activity (β−, t½ = 2.4 h) coprecipitated with rubidium perchlorate, as expected for a perbromate species.[5]

Electrolysis of lithium bromate at a platinum anode gave a yield of about 2% perbromate.[5] Small amounts of perbromate (1–4% yield) were also formed by the xenon difluoride oxidation of bromates, while perxenate, peroxodisulfate and ozone were ineffective as oxidizing agents.[5]

Preparation

The preparation of macroscopic quantities of perbromates relies on the oxidation of sodium bromate[note 2] in alkaline solution (5 M sodium hydroxide) by gaseous fluorine:

BrO−3 + F2 + 2 OH− → BrO−4 + 2 F− + H2O

The yield is about 20%, and the reaction can be operated on a 100-gram scale. Perbromic acid can be prepared from the resulting sodium perbromate by ion exchange, and used to prepare other perbromates.[19]

Structure

  r(Br–O)/pm Ref.
CsBrO4 159.1 [20]
LiBrO4·3H2O 160.6 [21]
Ba(BrO4)2·3H2O 160.8 [22]
EXAFS 161 [1]
[Ni(H2O)6](BrO4)2 162.9 [23]

The perbromate ion has Td symmetry.

Notes and references

Notes

  1. ↑ Perbromic acid is fully dissociated in solution. By analogy with perchloric acid,[4] the solid HBrO4·2H2O is expected to contain [(H2O)2H]+[BrO4]−.
  2. ↑ Sodium bromate is used as the resulting sodium perbromate is very soluble in water; potassium perbromate is much less soluble (as is the case with the corresponding perchlorates).[18]

References

  1. ↑ 1.0 1.1 Levason, William; Ogden, J. Steven; Spicer, Mark D.; Young, Nigel A. Characterisation of the oxo-anions of bromine BrOx− (x = 1–4) by infrared, Raman, nuclear magnetic resonance, and bromine K-edge extended X-ray absorption fine structure techniques. J. Chem. Soc., Dalton Trans. 1990 (1), 349–53. DOI: 10.1039/DT9900000349.
  2. ↑ Johnson, Gerald K.; Smith, Peter N.; Appelman, Evan H.; Hubbard, Ward N. Thermodynamic properties of the perbromate and bromate ions. Inorg. Chem. 1970, 9 (1), 119–25. DOI: 10.1021/ic50083a025.
  3. ↑ Lee, C. L.; Lister, M. W. The Decomposition of Aqueous Sodium Bromite. Can. J. Chem. 1971, 49 (17), 2822–26. DOI: 10.1139/v71-470.
  4. ↑ Olovsson, Ivar Hydrogen-Bond Studies. XXIX. Crystal Structure of Perchloric Acid Dihydrate, H5O2+ClO4−. J. Chem. Phys. 1968, 49 (3), 1063–67. DOI: 10.1063/1.1670193.
  5. ↑ 5.0 5.1 5.2 5.3 Appelman, Evan H. The Synthesis of Perbromates. J. Am. Chem. Soc. 1968, 90 (7), 1900–1. DOI: 10.1021/ja01009a040. Appelman, Evan H. Perbromic acid and perbromates: synthesis and some properties. Inorg. Chem. 1969, 8 (2), 223–27. DOI: 10.1021/ic50072a008.
  6. ↑ 6.0 6.1 Greenwood, Norman N.; Earnshaw, A. Chemistry of the Elements; Pergamon: Oxford, 1984; pp 1020–22. ISBN 0-08-022057-6.
  7. ↑ 7.0 7.1 Cotton, F. Albert; Wilkinson, Geoffrey Advanced Inorganic Chemistry, 5th ed.; Wiley-Interscience: New York, 1988; pp 568–69. ISBN 0-471-84997-9.
  8. ↑ Kjaer, A. M.; Ulstrup, J. Electron-transfer reactions between the perbromate ion and iron(II) complexes of 2,2'-bipyridine and substituted 1,10-phenanthrolines. Inorg. Chem. 1982, 21 (9), 3490–94. DOI: 10.1021/ic00139a044.
  9. ↑ Herrell, A. Y.; Gayer, K. H. The elusive perbromates. J. Chem. Educ. 1972, 49 (9), 583. DOI: 10.1021/ed049p583.
  10. ↑ Appelman, Evan H. Nonexistent compounds. Two case histories. Acc. Chem. Res. 1973, 6 (4), 113–17. DOI: 10.1021/ar50064a001.
  11. ↑ Kämmerer, Hermann Notizen. J. Prakt. Chem. 1863, 90, 190. DOI: 10.1002/prac.18630900126. Kaemmerer, Hermann Studien über die Sauerstoffverbindungen der Halogene. Ann. Phys. (Berlin) 1869, 214 (11), 390–417. DOI: 10.1002/andp.18692141104.
  12. ↑ Pattison Muir, M. M. Note on the perbromates. J. Chem. Soc. 1876, 30, 469. DOI: 10.1039/JS8763000469.
  13. ↑ Wolfram, Guido Ueber die Darstellung der Perbromsäure. Justus Liebigs Ann. Chem. 1879, 198 (1–2), 95–98. DOI: 10.1002/jlac.18791980107.
  14. ↑ Cook, Ernest H. Effect of heat on iodates and bromates. I. Potassic iodate and bromate. J. Chem. Soc., Trans. 1894, 65, 802–14. DOI: 10.1039/CT8946500802.
  15. ↑ Bancroft, G. M.; Gesser, H. D. The search for perbromate—I: The thermal decomposition of bromates. J. Inorg. Nucl. Chem. 1965, 27 (7), 1545–56. DOI: 10.1016/0022-1902(65)80016-1.
  16. ↑ Hugus, Z. Z., Jr. The Possible Use of 4f Orbitals in Bonding: the Enhanced Stability of the Higher Oxidation States of Iodine, Tellurium and Antimony; the Non-existence of Perbromic Acid. J. Am. Chem. Soc. 1952, 74 (4), 1076–77. DOI: 10.1021/ja01124a502.
  17. ↑ Urch, D. S. The perbromate problem. J. Inorg. Nucl. Chem. 1963, 25 (7), 771–78. DOI: 10.1016/0022-1902(63)80360-7.
  18. ↑ Ukeles, S. D.; Freiberg, M. Bromine, Inorganic Compounds. In Kirk-Othmer Encyclopedia of Chemical Technology; John Wiley: New York, 2002. DOI: 10.1002/0471238961.021815131001031.
  19. ↑ Appelman, E. H. Perbromic acid and potassium perbromate. Inorg. Synth. 1972, 13, 1–9. DOI: 10.1002/9780470132449.ch1.
  20. ↑ Gebert, Elizabeth; Peterson, Selmer W.; Reis, Arthur H., Jr.; Appelman, Evan H. The crystal structure of cesium perbromate. J. Inorg. Nucl. Chem. 1981, 43 (12), 3085–89. DOI: 10.1016/0022-1902(81)80067-X.
  21. ↑ Blackburn, A. C.; Gallucci, J. C.; Gerkin, R. E.; Reppart, W. J. Structure of lithium perbromate trihydrate. Acta Crystallogr., Sect. C: Cryst. Struct. Commun. 1993, 49 (8), 1437–39. DOI: 10.1107/S0108270193000538.
  22. ↑ Gerkin, R. E.; Reppart, W. J.; Appelman, E. H. The structure of barium perbromate trihydrate Ba(BrO4)2.3H2O. Acta Crystallogr., Sect. C: Cryst. Struct. Commun. 1988, 44 (6), 960–62. DOI: 10.1107/S0108270188002215.
  23. ↑ Gallucci, J. C.; Gerkin, R. E.; Reppart, W. J. Structure of nickel(II) perbromate hexahydrate at 296 K. Acta Crystallogr., Sect. C: Cryst. Struct. Commun. 1990, 46 (9), 1580–84. DOI: 10.1107/S0108270189013533.

Further reading

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