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Environmental fate & pathways

Hydrolysis

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Description of key information

The hydrolysis of BF4− occurs stepwise to BF3OH−, BF2(OH)2− , and BF(OH)3− and ultimately B(OH)4− accompanied with the formation of hydrofluoric acid.

Due to the strong affinity of boron for fluoride, complete hydrolysis of BF4- occurs at low concentrations, yet with a slow initial step, especially at normal temperature and at a neutral or alkaline pH.

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Additional information

Hydrolysis : weight of evidence approach

 

The fluoroboric anion is stable in concentrated solutions, and hydrolyses slowly in aqueous solution to hydroxyfluoroborates. Wamser described the hydrolysis reaction of fluoroboric acid (HBF4 + H2O ↔ HBF3OH + HF) and determined the equilibrium constants for the hydrolysis at 25°C. The hydrolysis of aqueous fluoroboric acid solutions (at equilibrium) was measured over the concentration range 0.001M to 5.41M, the % hydrolysis at equilibrium ranged from 77.7% to 5.47% respectively and the equilibrium constant between 2.75E-3 to 17.6E-3 respectively. At 0.01M HBF4 the hydrolysis was about 35% and the equilibrium constant 1.98E-3 (Wamser, 1948).

 

The equilibrium quotient Q in 1 molal NaCl at 25°C shows the strong affinitiy of boron for fluoride (Papcun, 2000).

B(OH)3 + 4 F+ 3 H+ ↔ BF4+ 3 H2O logQ = 19.0 ± 0.1 (Mesmer, Palen, & Baes Jr., 1973)

 

The hydrolysis of BF4occurs stepwise to BF3OH, BF2(OH)2, and BF(OH)3 and ultimately B(OH)4accompanied with the formation of hydrofluoric acid. Studies demonstrated that BF4is very stable to hydrolysis, but is slow to form from BF3OHand HF. Kinetic results and 19F NMR experiments illustrated that the hydroxyfluoroborates are in rapid equilibrium and easily exchange fluoride (Mesmer et al., 1973; Wamser, 1951). By boiling solutions containing NaBF4 in the presence of calcium carbonate a small amount of hydrolysis was effected and the 19F and 11B spectra were observed for what was believed to be BF3OH- (Mesmer and Rutenberg, 1973).

The rate constant obtained from the earlier part of the hydrolysis reaction of 0.1 M HBF4 was found to be 0.00090 min-1 (Wamser, 1948). Wamser also stated that a 0.001M solution of fluoroboric acid requires about two months to come to equilibrium and that the salts of fluoboric acid hydrolyze very slowly in the presence of alkali (rapidly only at the boiling point and in the presence of excess calcium ions) (Wamser, 1948).

 

Anbar and Guttmann demonstrated that the rate of hydrolysis of fluoroborate ions was first order for BF4and first order for the hydrogen ion concentration in the acid region; hydrolyis rate = k [H+] [BF4]. There was found no catalytic effect of HF initially present on the rate of hydrolysis. This result is confirmed by the fact that no autocatalytic deviation from first rate law was observed. In another series of experiments bisulfate ions were added up to 0.9 M keeping the pH constant, no appreciable change of the rate of hydrolysis could be detected. The activation energy ΔE = 25.1 kcal./mole-1, was derived from data at 25°, 37°, (1.23 ± 0.05 X 10-3 mol-1L. sec.-1), 60° (k = 1.33 ± 0.05 X 10-2 X 10-3 mol-1L. sec.-1) and at 100° (k = 1.10 X 100mol-1L. sec.-1). The specific rate constant of fluoroborate hydrolysis in acid solution is therefore k = 2.3 x 1014e-25100/RTmol-1L. sec.-1. In neutral and basic solutions the rate of hydrolysis is much lower and no effect of alkalinity on the rate of hydrolysis could be detected from measurements at 60°C and at 100°C. The following rate constant for non-acid hydrolysis of fluoroborate ion was found: k = 8 x 105e-15500/RTsec.-1 (Anbar & Guttmann, 1960)

 

The specific rate constants of the monomolecular dissociation of BF4at 25° is slower by a factor of about 1010 from that of HBF4, although its energy of activation is lower by about 9.6 kcal/mole. This points to a spectacular increase in the entropy of activation of the dissociation process on addition of a proton to the BF4ion (Anbar & Guttmann, 1960).

 

The result for the recently conducted OECD guideline study 111 (Tier 1) is well in line with literature (TNO Triskelion BV, 2012). The degree of hydrolysis as observed in this study, i.e. between 42 -56 % after 5 days at 50°C is in line with the result observed by Wamser, i.e. 35% (Wamser, 1948).

Also the recent F-NMR investigations (Hildebrand, 2012) confirmed the existing literature, i.e. at a higher concentration tested ( 0,5%, i.e. 0.4 M) the hydrolysis of the BF4- was less pronounced, i.e. ± 5% after 5 days at room temperature and ± 15% after 5 days at 50°C.

 

Conclusion

The hydrolysis of BF4occurs stepwise to BF3OH, BF2(OH)2, and BF(OH)3 and ultimately B(OH)4accompanied with the formation of hydrofluoric acid. Due to the strong affinity of boron for fluoride, complete hydrolysis of BF4- occurs at low concentrations, yet with a slow initial step, especially at normal temperature and at a neutral or alkaline pH. 

  

References

- Anbar, M., & Guttmann, S. (1960). The isotopic exchange of fluoroboric acid with hydrofluoric acid. J. Phys. Chem.,64, 1896–1899.

- TNO Triskelion BV (2012). Hydrolysis of the BF4- ion of KBF4 as a function of pH according to OECD guideline 111, a preliminary test (Tier 1) and identification of hydrolysis products. Unpublished report.

- Mesmer, R. E., Palen, K. M., & Baes Jr., C. F. (1973). Fluroborate Equilibria in Aqueous Solutions. Inorganic Chemistry,12(1), 89–95.

- Mesmer, R.E. & Rutenberg, A.C. (1973). Fluorine-19 Nuclear Magnetic Resonance studies on fluoroborate Species in Aqueous Solutions. Inorganic Chemistry, 12(3), 699 -702.

- Papcun, J. R. (2000). Fluorine Compounds, Inorganic, fluoroboric Acid and Fluoroborates. Kirk-Othmer Encyclopedia of Chemical Technology(pp. 1–13). John Wiley & Sons, Inc. doi:10.1002/0471238961.0612211516011603.a01

- Wamser, C. A. (1948). Hydrolysis of fluoboric Acid in Aqueous Solution. J. Am Chem. Soc.,70, 1209–1215.

- Wamser, C. A. (1951). Equilibria in the System Boron Trifluoride-Water at 25 °C. J. Am Chem. Soc.,73, 409–416.

- Hildebrand, M (2012). KBF4 hydrolytical stability in water. Unpublished report