Fluoroboric acid

Fluoroboric acid or tetrafluoroboric acid (archaically, fluoboric acid) is an inorganic compound with the chemical formula [H+][BF4−], where H+ represents the solvated proton. The solvent can be any suitably Lewis basic entity. For instance, in water, it can be represented by H
3OBF
4 (oxonium tetrafluoroborate), although more realistically, several water molecules solvate the proton: [H(H2O)n+][BF4−]. The ethyl ether solvate is also commercially available: [H(Et2O)n+][BF4−], where n is most likely 2. Unlike strong acids like H2SO4 or HClO4, the pure unsolvated substance does not exist (see below).

Fluoroboric acid
Names
Preferred IUPAC name
Tetrafluoroboric acid
Other names
tetrafluoroboric acid, oxonium tetrafluoroboranuide, oxonium tetrafluoridoborate(1-)
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.037.165
EC Number
  • 240-898-3
21702
MeSH Fluoroboric+acid
RTECS number
  • ED2685000
UNII
UN number 1775
Properties
BF4H
Molar mass 87.81 g·mol−1
Appearance Colourless liquid
Melting point −90 °C (−130 °F; 183 K)
Boiling point 130 °C (266 °F; 403 K)
Acidity (pKa) ~1.8 (MeCN solution)[1]
Hazards
Safety data sheet External MSDS
C
R-phrases (outdated) R34
S-phrases (outdated) (S1/2), S26, S27, S45
NFPA 704 (fire diamond)
Flammability code 0: Will not burn. E.g. waterHealth code 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
0
3
0
Related compounds
Related compounds
 Hydrogen fluoride

Triflic acid

Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YN ?)
Infobox references

It is mainly produced as a precursor to other fluoroborate salts.[2] It is a strong acid. Fluoroboric acid is corrosive and attacks the skin. It is available commercially as a solution in water and other solvents such as diethyl ether. It is a strong acid with a weakly coordinating, non-oxidizing conjugate base.[1] It is structurally similar to perchloric acid, but lacks the hazards associated with oxidants.

Structure and production

Pure HBF4 has been described as a "nonexistent compound", as a sufficiently 'naked' proton is expected to abstract a fluoride from the tetrafluoroborate ion to give hydrogen fluoride and boron trifluoride: [H+][BF4–] → HF + BF3. (The same holds true for the superacids that are known by the simple formulas HPF6 and HSbF6.)[3][4] However, a solution of BF3 in HF is highly acidic, having an approximate speciation of [H2F+][BF4–] and a Hammett acidity function of –16.6 at 7 mol % BF3, easily qualifying as a superacid.[5] Although the solvent-free HBF4 has not been isolated, its solvates are well characterized. These salts consist of protonated solvent as a cation, e.g., H3O+ and H
5O+
2, and the tetrahedral BF−
4 anion. The anion and cations are strongly hydrogen-bonded.[6]

Subunit of crystal structure of H3OBF4 highlighting the hydrogen bonding between the cation and the anion

Aqueous solutions of HBF4 are produced by dissolving boric acid in aqueous hydrofluoric acid.[7][8] Three equivalents of HF react to give the intermediate boron trifluoride and the fourth gives fluoroboric acid:

B(OH)3 + 4 HF → H3O+ + BF−
4 + 2 H2O

Anhydrous solutions can be prepared by treatment of aqueous fluoroboric acid with acetic anhydride.[9]

Acidity

The acidity of fluoroboric acid is complicated by the fact that the name refers to several different species H(OEt2)+BF−
4, H3O+BF−
4, and HF.BF3 – each with a different acidity. The aqueous pKa is quoted as −0.44.[2] Titration of NBu+
4BF−
4 in acetonitrile solution indicates that HBF4, i.e., HF.BF3, has a pKa of 1.6 in that solvent. Its acidity is thus comparable to that of fluorosulfonic acid.[1]

Applications

Fluoroboric acid is the principal precursor to fluoroborate salts, which are typically prepared by treating the metal oxides with fluoroboric acid. The inorganic salts are intermediates in the manufacture of flame-retardant materials and glazing frits, and in electrolytic generation of boron. HBF4 is also used in aluminum etching and acid pickling.

Organic chemistry

HBF4 is used as a catalyst for alkylations and polymerizations. In carbohydrate protection reactions, ethereal fluoroboric acid is an efficient and cost-effective catalyst for transacetalation and isopropylidenation reactions. Acetonitrile solutions cleave acetals and some ethers. Many reactive cations have been obtained using fluoroboric acid, e.g. tropylium tetrafluoroborate (C
7H+
7BF−
4), triphenylmethyl tetrafluoroborate (Ph
3C+
BF−
4), triethyloxonium tetrafluoroborate (Et
3O+
BF−
4), and benzenediazonium tetrafluoroborate (PHN+
2BF−
4).

Electroplating

Solutions of HBF4 are used in the electroplating of tin and tin alloys. In this application, methanesulfonic acid is displacing the use of HBF4.[10]

Safety

HBF4 is toxic and attacks skin and eyes. It attacks glass.[2] It hydrolyzes, releasing corrosive, volatile hydrogen fluoride.[10]

Other fluoroboric acids

A series of fluoroboric acids is known in aqueous solutions. The series can be presented as follows:[11]

  • H[B(OH)4]
  • H[BF(OH)3]
  • H[BF2(OH)2]
  • H[BF3(OH)]
  • H[BF4]

See also

References
  1. Kütt, A., et al., "Equilibrium Acidities of Superacids", J. Org. Chem. 2010, volume 76, pp. 391-395. doi:10.1021/jo101409p
  2. Gregory K. Friestad, Bruce P. Branchaud "Tetrafluoroboric Acid" E-Eros Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.rt035
  3. Juhasz, Mark; Hoffmann, Stephan; Stoyanov, Evgenii; Kim, Kee-Chan; Reed, Christopher A. (2004-10-11). "The Strongest Isolable Acid". Angewandte Chemie International Edition. 43 (40): 5352–5355. doi:10.1002/anie.200460005. ISSN 1433-7851. PMID 15468064.
  4. Reed, Christopher A. (2005). "Carborane acids. New "strong yet gentle" acids for organic and inorganic chemistry" (PDF). Chem. Commun. 0 (13): 1669–1677. doi:10.1039/B415425H. ISSN 1359-7345. PMID 15791295.
  5. Olah, George A.; Surya Prakash, G. K.; Sommer, Jean; Molnar, Arpad (2009-02-03). Superacid chemistry. Olah, George A. (George Andrew), 1927-2017,, Olah, George A. (George Andrew), 1927-2017. (2nd ed.). Hoboken, N.J. ISBN 9780471596684. OCLC 191809598.
  6. Mootz, D.; Steffen, M. "Crystal structures of acid hydrates and oxonium salts. XX. Oxonium tetrafluoroborates H3OBF4, [H5O2]BF4, and [H(CH3OH)2]BF4", Zeitschrift für Anorganische und Allgemeine Chemie 1981, vol. 482, pp. 193-200. doi:10.1002/zaac.19814821124
  7. Brotherton, R. J.; Weber, C. J.; Guibert, C. R.; Little, J. L. "Boron Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a04_309.
  8. Flood, D. T. (1933). "Fluorobenzene" (PDF). Organic Syntheses. 13: 46.; Collective Volume, 2, p. 295
  9. Wudl, F.; Kaplan, M. L., "2,2′-Bi-L,3-Dithiolylidene (Tetrathiafulvalene, TTF) and its Radical Cation Salts" Inorg. Synth. 1979, vol. 19, 27. doi:10.1002/9780470132500.ch7
  10. Balaji, R.; Pushpavanam, Malathy (2003). "Methanesulphonic acid in electroplating related metal finishing industries". Transactions of the Imf. 81 (5): 154–158. doi:10.1080/00202967.2003.11871526. S2CID 91584456.
  11. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.

Further reading
  • Albert, R.; Dax, K.; Pleschko, R.; Stütz, A. E. (1985). "Tetrafluoroboric acid, an efficient catalyst in carbohydrate protection and deprotection reactions". Carbohydrate Research. 137: 282–290. doi:10.1016/0008-6215(85)85171-5.
  • Bandgar, B. P.; Patil, A. V.; Chavan, O. S. (2006). "Silica supported fluoroboric acid as a novel, efficient and reusable catalyst for the synthesis of 1,5-benzodiazepines under solvent-free conditions". Journal of Molecular Catalysis A: Chemical. 256 (1–2): 99–105. doi:10.1016/j.molcata.2006.04.024.
  • Heintz, R. A.; Smith, J. A.; Szalay, P. S.; Weisgerber, A.; Dunbar, K. R. (2002). Homoleptic Transition Metal Acetonitrile Cations with Tetrafluoroborate or Trifluoromethanesulfonate Anions. Inorganic Syntheses. 33. pp. 75–83. doi:10.1002/0471224502. ISBN 9780471208259.
  • Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 307. ISBN 978-0-13-039913-7.
  • Meller, A. (1988). "Boron". Gmelin Handbook of Inorganic Chemistry. 3. New York: Springer-Verlag. pp. 301–310.
  • Perry, D. L.; Phillips, S. L. (1995). Handbook of Inorganic Compounds (1st ed.). Boca Raton: CRC Press. p. 1203. ISBN 9780849386718.
  • Wamser, C. A. (1948). "Hydrolysis of Fluoboric Acid in Aqueous Solution". Journal of the American Chemical Society. 70 (3): 1209–1215. doi:10.1021/ja01183a101.
  • Wilke-Dörfurt, E.; Balz, G. (1927). "Zur Kenntnis der Borfluorwasserstoffsäure und ihrer Salze". Zeitschrift für Anorganische und Allgemeine Chemie. 159 (1): 197–225. doi:10.1002/zaac.19271590118.
This article is issued from Wikipedia. The text is licensed under Creative Commons - Attribution - Sharealike. Additional terms may apply for the media files.