Fluoroboric acid or tetrafluoroboric acid (archaically, fluoboric acid) is an inorganic compound with the simplified chemical formulaH+[BF4]−. Unlike other strong acids like H2SO4 or HClO4, the pure (solvent free) tetrafluoroboric acid (more precisely, pure hydrogen tetrafluoroborate H[BF4]) does not exist. The term "fluoroboric acid" refers to a range of chemical compounds, depending on the solvent. The H+ in the simplified formula of fluoroboric acid represents the solvated proton. The solvent can be any suitable Lewis base. For instance, if the solvent is water, fluoroboric acid can be represented by the formula [H3O]+[BF4]− (oxonium tetrafluoroborate), although more realistically, several water molecules solvate the proton: [H(H2O)n]+[BF4]−. The ethyl ether solvate is also commercially available, where the fluoroboric acid can be represented by the formula [H((CH3CH2)2O)n]+[BF4]−, where n is most likely 2.
It is mainly produced as a precursor to other fluoroborate salts.[3] 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.[2] It is structurally similar to perchloric acid, but lacks the hazards associated with oxidants.
Structure and production
Pure H[BF4] 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 simplified formulas H[PF6] and H[SbF6].[4][5] However, a solution of BF3 in HF is highly acidic, having an approximate speciation of [H2F]+[BF4]− (fluoronium tetrafluoroborate) and a Hammett acidity function of −16.6 at 7 mol % BF3, easily qualifying as a superacid.[6] Although the solvent-free H[BF4] has not been isolated, its solvates are well characterized. These salts consist of protonated solvent as a cation, e.g., H3O+ and H5O+2, and the tetrahedral BF−4 anion. The anion and cations are strongly hydrogen-bonded.[7]
Subunit of crystal structure of [H3O]+[BF4]− highlighting the hydrogen bonding between the cation and the anion
Aqueous solutions of H[BF4] are produced by dissolving boric acid in aqueous hydrofluoric acid.[8][9] 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.[10]
Acidity
The acidity of fluoroboric acid is complicated by the fact that its name refers to a range of different compounds, e.g. [H(CH3CH2)2O]+[BF4]− (dimethyloxonium tetrafluoroborate), [H3O]+[BF4]− (oxonium tetrafluoroborate), and HF·BF3 (hydrogen fluoride-boron trifluoride 1:1 adduct) – each with a different acidity. The aqueous pKa is quoted as −0.44.[3]Titration of [N((CH2)3CH3)4]+[BF4]− (tetrabutylammonium tetrafluoroborate) in acetonitrile solution indicates that H[BF4], i.e., HF·BF3, has a pKa of 1.6 in that solvent. Its acidity is thus comparable to that of fluorosulfonic acid.[2]
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. H[BF4] is also used in aluminum etching and acid pickling.
Solutions of H[BF4] are used in the electroplating of tin and tin alloys. In this application, methanesulfonic acid is displacing the use of H[BF4].[11] Fluoroboric acid is also used for high-speed electroplating of copper in fluoroborate baths.[12]
Safety
H[BF4] is toxic and attacks skin and eyes. It attacks glass.[3] It hydrolyzes, releasing corrosive, volatile hydrogen fluoride.[11]
Other fluoroboric acids
A series of fluoroboric acids is known in aqueous solutions. The series can be presented as follows:[13]
H+[B(OH)4]− (hydrogen tetrahydroxyborate) (not a fluoroboric acid)
^ abcKütt, A., et al., "Equilibrium Acidities of Superacids", J. Org. Chem. 2010, volume 76, pp. 391-395. doi:10.1021/jo101409p
^ abcGregory K. Friestad, Bruce P. Branchaud "Tetrafluoroboric Acid" E-Eros Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.rt035
^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. ISSN1433-7851. PMID15468064.
^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. ISBN9780471596684. OCLC191809598.((cite book)): CS1 maint: location missing publisher (link)
^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
^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
^ abBalaji, 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. S2CID91584456.
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. Vol. 33. pp. 75–83. doi:10.1002/0471224502. ISBN9780471208259.
Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 307. ISBN978-0-13-039913-7.
Meller, A. (1988). "Boron". Gmelin Handbook of Inorganic Chemistry. Vol. 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. ISBN9780849386718.
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.
