Trifluoroacetic acid

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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).
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Trifluoroacetic acid (TFA) is an

conjugate base. TFA is widely used in organic chemistry
for various purposes.

Synthesis

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

Trifluoroacetic acid in a beaker

TFA is the precursor to many other fluorinated compounds such as

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

Boc used in organic chemistry and peptide synthesis.[8][9]

At a low concentration, TFA is used as an ion pairing agent in

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

Trifluoroacetic acid is a corrosive strong acid

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.

TFA's reaction with bases and metals, especially light metals, is strongly exothermic. The reaction with lithium aluminium hydride (LAH) results in an explosion.[13]

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

Environment

No known natural processes generate trifluoroacetic acid.

2,3,3,3-tetrafluoropropene.[citation needed
]

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

Trifluoroacetic acid is mildly

phytotoxic.[22]

See also

References

  1. ^ Kreglewski, A. (1962). "Trifluoroacetic acid". Welcome to the NIST WebBook. 10 (11–12): 629–633. Retrieved 1 March 2020.
  2. 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. ^
    ISBN 978-3527306732
    .
  5. 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? (PDF). Pierce Chemical. pp. 88–90.
  7. ISBN 978-0-471-93623-7
    .
  8. PMID 744685
    .
  9. ISBN 978-3-527-63182-7
    .
  10. doi:10.1021/ac00193a027
    .
  11. S2CID 250765145
    .
  12. doi:10.1021/ja01149a122
    .
  13. ^ Safety data sheet for Trifluoroacetic acid (PDF) from EMD Millipore, revision date 10/27/2014.
  14. PMID 31643481
    , retrieved 15 July 2021
  15. S2CID 239768006
    .
  16. S2CID 232115008
    .
  17. PMID 11811478
    .
  18. PMID 16190212
    .
  19. PMID 11811478
    .
  20. ^ "Refreshingly cool, potentially toxic". Ludwig-Maximilians-Universität (LMU) Munich. 2014. Retrieved 26 July 2018.
  21. ^ Kirschner, E., Chemical and Engineering News 1994, 8.
  22. doi:10.1080/10807039991289644
    .
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