Jump to content

Trifluoroacetic acid

From Wikipedia, the free encyclopedia

Trifluoroacetic acid
Names
Preferred IUPAC name
Trifluoroacetic acid
Other names
2,2,2-Trifluoroacetic acid
2,2,2-Trifluoroethanoic acid
Perfluoroacetic acid
Trifluoroethanoic acid
TFA
Identifiers
3D model (JSmol)
742035
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.000.846 Edit this at Wikidata
2729
RTECS number
  • AJ9625000
UNII
  • InChI=1S/C2HF3O2/c3-2(4,5)1(6)7/h(H,6,7) checkY
    Key: DTQVDTLACAAQTR-UHFFFAOYSA-N checkY
  • InChI=1/C2HF3O2/c3-2(4,5)1(6)7/h(H,6,7)
    Key: DTQVDTLACAAQTR-UHFFFAOYAP
  • FC(F)(F)C(=O)O
Properties
C2HF3O2
Molar mass 114.023 g·mol−1
Appearance colorless liquid
Odor Pungent/Vinegar
Density 1.489 g/cm3, 20 °C
Melting point −15.4 °C (4.3 °F; 257.8 K)
Boiling point 72.4 °C (162.3 °F; 345.5 K)
miscible
Vapor pressure 0.0117 bar (1.17 kPa) at 20 °C[1]
Acidity (pKa) 0.52 [2]
Conjugate base trifluoroacetate
-43.3·10−6 cm3/mol
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Highly corrosive
GHS labelling:
GHS05: CorrosiveGHS07: Exclamation mark
Danger
H314, H332, H412
P260, P261, P264, P271, P273, P280, P301+P330+P331, P303+P361+P353, P304+P312, P304+P340, P305+P351+P338, P310, P312, P321, P363, P405, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilInstability 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazards (white): no code
3
1
1
Safety data sheet (SDS) External MSDS
Related compounds
Related perfluorinated acids
 Heptafluorobutyric acid
Perfluorooctanoic acid
Perfluorononanoic acid
Related compounds
 Acetic acid
Trichloroacetic 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 checkY☒N ?)

Trifluoroacetic acid (TFA) is an organofluorine compound with the chemical formula CF3CO2H. It is a haloacetic acid, with all three of the acetyl group's hydrogen atoms replaced by fluorine atoms. It is a colorless liquid with a vinegar-like odor. TFA is a stronger acid than acetic acid, having an acid ionisation constant, Ka, that is approximately 34,000 times higher,[3] as the highly electronegative fluorine atoms and consequent electron-withdrawing nature of the trifluoromethyl group weakens the oxygen-hydrogen bond (allowing for greater acidity) and stabilises the anionic conjugate base. TFA is widely used in organic chemistry for various purposes.

Synthesis

[edit]

TFA is prepared industrially by the electrofluorination of acetyl chloride or acetic anhydride, followed by hydrolysis of the resulting trifluoroacetyl fluoride:[4]

CH
3COCl + 4 HF → CF
3COF + 3 H
2 + HCl
CF
3COF + H
2O → CF
3COOH + HF

Where desired, this compound may be dried by addition of trifluoroacetic anhydride.[5]

An older route to TFA proceeds via the oxidation of 1,1,1-trifluoro-2,3,3-trichloropropene with potassium permanganate. The trifluorotrichloropropene can be prepared by Swarts fluorination of hexachloropropene.[6]

Uses

[edit]
Trifluoroacetic acid in a beaker

TFA is the precursor to many other fluorinated compounds such as trifluoroacetic anhydride, trifluoroperacetic acid, and 2,2,2-trifluoroethanol.[4] It is a reagent used in organic synthesis because of a combination of convenient properties: volatility, solubility in organic solvents, and its strength as an acid.[7] TFA is also less oxidizing than sulfuric acid but more readily available in anhydrous form than many other acids. One complication to its use is that TFA forms an azeotrope with water (b. p. 105 °C).

TFA is popularly used as a strong acid to remove protecting groups such as Boc used in organic chemistry and peptide synthesis.[8][9]

At a low concentration, TFA is used as an ion pairing agent in liquid chromatography (HPLC) of organic compounds, particularly peptides and small proteins. TFA is a versatile solvent for NMR spectroscopy (for materials stable in acid). It is also used as a calibrant in mass spectrometry.[10]

TFA is used to produce trifluoroacetate salts.[11]

Safety

[edit]

Trifluoroacetic acid is a corrosive strong acid[12] but it does not pose the hazards associated with hydrofluoric acid because the carbon-fluorine bond is not labile. TFA is harmful when inhaled, causes severe skin burns and is toxic for aquatic organisms even at low concentrations.

The reaction of TFA with bases and metals, especially light metals, is strongly exothermic. For example, reaction of TFA with lithium aluminium hydride (LAH) may result in an explosion.[13]

TFA is a metabolic breakdown product of the volatile anesthetic agent halothane. It is thought to be responsible for halothane-induced hepatitis.[14]

Environment

[edit]

No known natural processes generate trifluoroacetic acid.[15] In the environment, trifluoroacetic acid may be formed by photooxidation of the commonly used refrigerant 1,1,1,2-tetrafluoroethane (R-134a).[citation needed] Moreover, it is formed as an atmospheric degradation product of almost all fourth-generation synthetic refrigerants, also called hydrofluoroolefins (HFO), such as 2,3,3,3-tetrafluoropropene.[citation needed]

Trifluoroacetic acid degrades very slowly in the environment, and has been found in increasing amounts as a contaminant in water, soil, food, and the human body.[16] Median concentrations of a few micrograms per liter have been found in beer and tea.[17] Sea water contains about 200 ng of TFA per liter.[18][19][20] No biodegradation mechanism for the compound is known in water,[21] although biotransformation apparently decarboxylates the acid to fluoroform.[22]

Trifluoroacetic acid is mildly phytotoxic.[23]

See also

[edit]

References

[edit]
  1. ^ Kreglewski, A. (1962). "Trifluoroacetic acid". Welcome to the NIST WebBook. 10 (11–12): 629–633. Retrieved 1 March 2020.
  2. ^ W. M. Haynes.; David R. Lide; Thomas J. Bruno, eds. (2016–2017). CRC Handbook of Chemistry and Physics. CRC Press. pp. 954–963. ISBN 978-1-4987-5429-3.
  3. ^ Note: Calculated from the ratio of the Ka values for TFA (pKa = 0.23) and acetic acid (pKa = 4.76)
  4. ^ a b G. Siegemund; W. Schwertfeger; A. Feiring; B. Smart; F. Behr; H. Vogel; B. McKusick. "Fluorine Compounds, Organic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_349. ISBN 978-3527306732.
  5. ^ Wilfred L.F. Armarego & Christina Li Lin Chai (2009). "Chapter 4 - Purification of Organic Chemicals". Purification of Laboratory Chemicals (6th ed.). pp. 88–444. doi:10.1016/B978-1-85617-567-8.50012-3. ISBN 978-1-85617-567-8.
  6. ^ Gergel, Max G. (March 1977). Excuse me sir, would you like to buy a kilo of isopropyl bromide?. Pierce Chemical. pp. 88–90.
  7. ^ Eidman, K. F.; Nichols, P. J. (2004). "Trifluoroacetic Acid". In L. Paquette (ed.). Encyclopedia of Reagents for Organic Synthesis. New York: J. Wiley & Sons. doi:10.1002/047084289X.rt236.pub2. hdl:10261/236866. ISBN 978-0-471-93623-7.
  8. ^ Lundt, Behrend F.; Johansen, Nils L.; Vølund, Aage; Markussen, Jan (1978). "Removal of t-Butyl and t-Butoxycarbonyl Protecting Groups with Trifluoroacetic acid". International Journal of Peptide and Protein Research. 12 (5): 258–268. doi:10.1111/j.1399-3011.1978.tb02896.x. PMID 744685.
  9. ^ Andrew B. Hughes (2011). "1. Protection Reactions". In Vommina V. Sureshbabu; Narasimhamurthy Narendra (eds.). Amino Acids, Peptides and Proteins in Organic Chemistry: Protection Reactions, Medicinal Chemistry, Combinatorial Synthesis. Vol. 4. pp. 1–97. doi:10.1002/9783527631827.ch1. ISBN 978-3-527-63182-7.
  10. ^ Stout, Steven J.; Dacunha, Adrian R. (1989). "Tuning and calibration in thermospray liquid chromatography/mass spectrometry using trifluoroacetic acid cluster ions". Analytical Chemistry. 61 (18): 2126. doi:10.1021/ac00193a027.
  11. ^ O. Castano; A. Cavallaro; A. Palau; J. C. Gonzalez; M. Rossell; T. Puig; F. Sandiumenge; N. Mestres; S. Pinol; A. Pomar & X. Obradors (2003). "High quality YBa2Cu3O7 thin films grown by trifluoroacetates metal-organic deposition". Superconductor Science and Technology. 16 (1): 45–53. Bibcode:2003SuScT..16...45C. doi:10.1088/0953-2048/16/1/309. S2CID 250765145.
  12. ^ Henne, Albert L; Fox, Charles J (1951). "Ionization constants of fluorinated acids". Journal of the American Chemical Society. 73 (5): 2323–2325. doi:10.1021/ja01149a122.
  13. ^ Safety data sheet for Trifluoroacetic acid (PDF) from EMD Millipore, revision date 10/27/2014.
  14. ^ "Halothane", LiverTox: Clinical and Research Information on Drug-Induced Liver Injury, Bethesda (MD): National Institute of Diabetes and Digestive and Kidney Diseases, 2012, PMID 31643481, retrieved 15 July 2021
  15. ^ Joudan, Shira; De Silva, Amila O.; Young, Cora J. (2021). "Insufficient evidence for the existence of natural trifluoroacetic acid". Environmental Science: Processes & Impacts. 23 (11): 1641–1649. doi:10.1039/D1EM00306B. hdl:10315/40884. ISSN 2050-7887. PMID 34693963. S2CID 239768006.
  16. ^ Hosea, Leana; Salvidge, Rachel (1 May 2024). "Rapidly rising levels of TFA 'forever chemical' alarm experts". The Guardian. ISSN 0261-3077. Retrieved 29 May 2024.
  17. ^ Marco Scheurer, Karsten Nödler (2021). "Ultrashort-chain perfluoroalkyl substance trifluoroacetate (TFA) in beer and tea – An unintended aqueous extraction". Food Chemistry. 351: 129304. doi:10.1016/j.foodchem.2021.129304. ISSN 0308-8146. PMID 33657499. S2CID 232115008.
  18. ^ Frank, H.; Christoph, E. H.; Holm-Hansen, O.; Bullister, J. L. (January 2002). "Trifluoroacetate in ocean waters". Environ. Sci. Technol. 36 (1): 12–5. Bibcode:2002EnST...36...12P. doi:10.1021/es0221659. PMID 11811478.
  19. ^ Scott, B. F.; MacDonald, R. W.; Kannan, K.; Fisk, A.; Witter, A.; Yamashita, N.; Durham, L.; Spencer, C.; Muir, D. C. G. (September 2005). "Trifluoroacetate profiles in the Arctic, Atlantic, and Pacific Oceans". Environ. Sci. Technol. 39 (17): 6555–60. Bibcode:2005EnST...39.6555S. doi:10.1021/es047975u. PMID 16190212.
  20. ^ Frank, Hartmut; Christoph, Eugen H.; Holm-Hansen, Osmund; Bullister, John L. (2002). "Trifluoroacetate in Ocean Waters". Environmental Science & Technology. 36 (1): 12–15. Bibcode:2002EnST...36...12F. doi:10.1021/es0101532. ISSN 0013-936X. PMID 11811478.
  21. ^ "Refreshingly cool, potentially toxic". Ludwig-Maximilians-Universität (LMU) Munich. 2014. Retrieved 26 July 2018.
  22. ^ Kirschner, E., Chemical and Engineering News 1994, 8.
  23. ^ Boutonnet, Jean Charles; Bingham, Pauline; Calamari, Davide; Rooij, Christ de; Franklin, James; Kawano, Toshihiko; Libre, Jean-Marie; McCul-Loch, Archie; Malinverno, Giuseppe; Odom, J Martin; Rusch, George M; Smythe, Katie; Sobolev, Igor; Thompson, Roy; Tiedje, James M (1999). "Environmental risk assessment of trifluoroacetic acid". International Journal of Human and Ecological Risk Assessment. 5 (1): 59–124. Bibcode:1999HERA....5...59B. doi:10.1080/10807039991289644.
Retrieved from "https://en.wikipedia.org/w/index.php?title=Trifluoroacetic_acid&oldid=1237385523"