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pKa of perfluorooctanoic acid

Posted by srayne on 04 Jan 2011 at 05:03 GMT

In this manuscript, the authors state that "A small quantity of TFA (1%) was included in all the solutions to protonate the carboxyl group of PFOA (the pKa value of PFOA is 3.8 at infinite dilution [18]), which is expected to disrupt the ionic interactions between PFOA and peptides, therefore helps evaporating PFOA.", where ref. [18] is the following article: Burns DC, Ellis DA, Li H, McMurdo CJ, Webster E (2008) "Experimental pKa determination for perfluorooctanoic acid (PFOA) and the potential impact of pKa concentration dependence on laboratory-measured partitioning phenomena and environmental modeling." Environ Sci Technol 42: 9283-9288.

The pKa of PFOA has been the subject of much debate in the literature over the past couple years, and the following experimental and theoretical studies appear to have shown that the reported infinite dilution pKa for PFOA at +3.8 by Burns et al. is in error, and that the actual pKa of PFOA is about 0 (i.e., effectively equivalent to trifluoroacetic acid):

Goss KU, The pKa values of PFOA and other highly fluorinated carboxylic acids. Environ Sci Technol. 2008 Jul 1;42(13):5032. (also see Comments and Replies-to-Comments in Environ Sci Technol on this article and the work by Burns et al.)

Cheng J, Psillakis E, Hoffmann MR, Colussi AJ (2009). "Acid Dissociation versus Molecular Association of Perfluoroalkyl Oxoacids: Environmental Implications". J. Phys. Chem. A 113 (29): 8152–6. doi:10.1021/jp9051352

Rayne S, Forest K. Theoretical studies on the pKa values of perfluoroalkyl carboxylic acids. Journal of Molecular Structure: THEOCHEM, Volume 949, Issues 1-3, 15 June 2010, Pages 60-69. doi:10.1016/j.theochem.2010.03.003

In light of these other works collectively agreeing that the pKa of PFOA is ~0 (and not +3.8), the authors may wish to reconsider some experimental interpretations in their study.

No competing interests declared.

RE: pKa of perfluorooctanoic acid

mmiyagi replied to srayne on 06 Jan 2011 at 18:37 GMT

Dear Dr. Rayne,
We appreciate your comment on the pKa of PFOA and agree that the acid dissociation constant (Ka) of PFOA is an important factor to determine the rate of its evaporation from the peptide samples. In our solvent system (ethanol:ethylacetate:water:TFA, 0.33:0.33:0.33:0.01), PFOA exist as a neutral or anionic species. We do not know the equilibrium between the two species; however, we found that adding acids (e.g., TFA) to the peptide samples significantly facilitates the evaporation of PFOA from the peptide samples. As discussed in the paper, we think this is because acids disrupt the ionic interaction between the anionic PFOA and peptides by protonating the anionic PFOA. In our experiment, protonated (neutral) PFOA is quickly removed from the system by evaporation, thus more protonated PFOA is constantly generated to return to the equilibrium and eventually the concentration of PFOA in the samples becomes extremely low. We think this mechanism can happen even when the pKa of PFOA is assumed to be much lower (pKa = 0) than the value (pKa = 3.8) cited in the paper in our system.

No competing interests declared.

RE: RE: pKa of perfluorooctanoic acid

srayne replied to mmiyagi on 09 Jan 2011 at 19:00 GMT

The results of the experiments do not appear to be in doubt, but the mechanism for PFOA removal is of interest. As we discuss in our paper (doi:10.1016/j.theochem.2010.03.003), and as also noted by Cheng et al. (doi:10.1021/jp9051352), we would expect PFOA to have a strongly concentration dependent pKa due to its propensity to begin aggregating at very low concentrations (perhaps below detection limits of most instruments; Cheng et al. discuss this issue). Aggregation of PFOA reduces the exposure of the carboxyl group to a polar solvent/co-solvent (such as water), thereby raising the effective pKa (analogous to moving a general acid HA from a polar protic solvent such as water to a polar aprotic solvent such as acetonitrile or DMSO). The PFOA concentrations in the extracts appear to be relatively high, and we may expect PFOA to be aggregated in the solutions. TFA has a much lower tendency to aggregate compared to TFA (due to the lack of a long hydrophobic perfluoroalkyl chain). Consequently, even though TFA and PFOA are believed to have the same infinite dilution pKa in water (about 0 to +0.5; there is evidence that PFOA may be a slightly stronger acid than TFA at infinite dilution in water), the effective pKa of PFOA would be expected to be much higher than that of TFA in concentrated solutions (where PFOA is highly aggregated, but TFA is not). Thus, TFA can protonate PFOA under such conditions because it is the stronger effective acid for the particular solution being investigated (but TFA does not appear to be a stronger acid than PFOA at infinite dilution in water). The effects of a partially nonaqueous solvent system that the authors used will also likely hinder ionization of both TFA and PFOA, but the nonaqueous pKa values of TFA and PFOA at infinite dilution under the experimental conditions (or a close approximation) are not known. Using the SPARC nonaqueous pKa calculator (http://ibmlc2.chem.uga.ed...), and assuming a solution composition of ethanol:ethyl acetate:water of 0.33:0.33:0.34, the nonaqueous pKa values at infinite dilution for TFA and PFOA are estimated at 3.44 and 2.12, respectively (i.e., TFA is a weaker acid in such a solution, perhaps semiquantitatively similar to the situation at infinite dilution in water). Assuming these results are at least qualitatively accurate, this suggests the mechanism for PFOA removal may be that its aggregation sufficiently raises its effective pKa to be substantially higher than that of TFA under the experimental conditions, allowing its protonation by TFA and subsequent removal from solution.

No competing interests declared.