Trifluoroacetic acid

From Wikipedia, the free encyclopedia
Jump to: navigation, search
Trifluoroacetic acid
Trifluoroacetic acid.svg
Trifluoroacetic-acid-3D-vdW.png
Trifluoroacetic-acid-elpot.png
Trifluoro acetic acid 1ml.jpg
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 (Jmol)
ChEBI
ChemSpider
ECHA InfoCard 100.000.846
RTECS number AJ9625000
UNII
Properties
C2HF3O2
Molar mass 114.02 g/mol
Appearance colorless liquid
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
Acidity (pKa) 0.23 [1]
-43.3·10−6 cm3/mol
Hazards
Main hazards Highly corrosive
Safety data sheet External MSDS
R-phrases (outdated) R20 R35 R52/53
S-phrases (outdated) S9 S26 S27 S28 S45 S61
NFPA 704
Flammability code 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g., canola oil Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium Special hazards (white): no codeNFPA 704 four-colored diamond
Related compounds
Related perfluorinated acids
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 YesYN ?)
Infobox references

Trifluoroacetic acid (TFA) is an organofluorine compound with the chemical formula CF3CO2H. It is a structural analogue of acetic acid with the three hydrogen atoms replaced by three fluorine atoms and is a colorless liquid with a sharp odor similar to vinegar. TFA is a stronger acid than acetic acid, having an acid ionisation constant that is approximately 34,000 times higher,[2] as the highly electronegative fluorine atoms and consequent electron-withdrawing nature of the trifluoromethyl group weakens the oxygen-hydrogen bond 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 and acetic anhydride, followed by hydrolysis of the resulting trifluoroacetyl fluoride:[3]

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

Where desired, this compound may be dried by addition of trifluoroacetic anhydride.[4]

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.

TFA occurs naturally in sea water, but only in small concentrations (<200 ng/L).[5][6]

Uses[edit]

TFA is the precursor to many other fluorinated compounds such as trifluoroacetic anhydride, trifluoroperacetic acid, and 2,2,2-trifluoroethanol.[3] 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 in peptide synthesis and other organic synthesis to remove the t-butoxycarbonyl protecting group.[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 acid but it does not pose the hazards associated with hydrofluoric acid because the carbon-fluorine bond is not labile. Only if heated or treated with ultrasonic waves will it decompose into hydrofluoric acid. TFA is harmful when inhaled, causes severe skin burns and is toxic for water organisms even at low concentrations.

TFA's reaction with bases and metals, especially light metals, is strongly exothermic. The reaction with lithium aluminium hydride (LAH) results in an explosion.[12]

Environment[edit]

TFA is neither readily nor inherently biodegradable [13] although it does not readily bioaccumulate. In water, it is toxic to algae.

See also[edit]

References[edit]

  1. ^ Ref 1 in Milne, J. B.; Parker, T. J. (1981). "Dissociation constant of aqueous trifluoroacetic acid by cryoscopy and conductivity". Journal of Solution Chemistry. 10 (7): 479. doi:10.1007/BF00652082. 
  2. ^ Note: Calculated from the ratio of the Ka values for TFA (pKa = 0.23) and acetic acid (pKa = 4.76)
  3. ^ a b G. Siegemund, W. Schwertfeger, A. Feiring, B. Smart, F. Behr, H. Vogel, B. McKusick (2005), "Fluorine Compounds, Organic", Ullmann's Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, doi:10.1002/14356007.a11_349 
  4. ^ Wilfred L.F. Armarego & Christina Li Lin Chai. "Chapter 4 - Purification of Organic Chemicals". Purification of Laboratory Chemicals (6th ed.). doi:10.1016/B978-1-85617-567-8.50012-3. 
  5. ^ "Trifluoroacetate in ocean waters". Environ. Sci. Technol. 36 (1): 12–5. January 2002. Bibcode:2002EnST...36...12P. doi:10.1021/es0221659. PMID 11811478. 
  6. ^ "Trifluoroacetate profiles in the Arctic, Atlantic, and Pacific Oceans". Environ. Sci. Technol. 39 (17): 6555–60. September 2005. Bibcode:2005EnST...39.6555S. doi:10.1021/es047975u. PMID 16190212. 
  7. ^ Eidman, K. F.; Nichols, P. J. (2004). "Trifluoroacetic Acid". In L. Paquette. Encyclopedia of Reagents for Organic Synthesis. New York: J. Wiley & Sons. doi:10.1002/047084289. 
  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. "1. Protection Reactions". In Vommina V. Sureshbabu; Narasimhamurthy Narendra. Amino Acids, Peptides and Proteins in Organic Chemistry: Protection Reactions, Medicinal Chemistry, Combinatorial Synthesis. 4. doi:10.1002/9783527631827.ch1. 
  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. 
  12. ^ Safety data sheet for Trifluoroacetic acid (PDF) from EMD Millipore, revision date 10/27/2014.
  13. ^ "GPS Safety Summary: Trifluoroacetic Acid" (PDF). Retrieved October 18, 2016.