Fluorocarbon

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Fluorocarbons, sometimes referred to as perfluorocarbons or PFCs, are, strictly speaking organofluorine compounds with the formula CxFy, i.e. they contain only carbon and fluorine.[1] The terminology is not strictly followed.[2] Compounds with the prefix perfluoro- are hydrocarbons, including those with heteroatoms, wherein all C-H bonds have been replaced by C-F bonds. Fluorocarbons and their derivatives are useful fluoropolymers, refrigerants, solvents, anesthetics. They are also the basis of some controversies, e.g., the ozone depletors and problems with bioaccumulation of fluorosurfactants.

Perfluoroalkanes[edit]

Chemical properties[edit]

Perflluoroalkanes are very stable because of the strength of the carbon–fluorine bond, one of the strongest in organic chemistry.[3] Its strength is a result of the electronegativity of fluorine imparting partial ionic character through partial charges on the carbon and fluorine atoms.[3] The partial charges shorten and strengthen the bond through favorable coulombic interactions. Additionally, multiple carbon–fluorine bonds increase the strength and stability of other nearby carbon–fluorine bonds on the same geminal carbon, as the carbon has a higher positive partial charge.[2] Furthermore, multiple carbon–fluorine bonds also strengthen the "skeletal" carbon–carbon bonds from the inductive effect.[2] Therefore, saturated fluorocarbons are more chemically and thermally stable than their corresponding hydrocarbon counterparts. In terms of chemical reactivity, C-F bonds are very inert. They are susceptible to attack by very strong reductants, e.g., Birch reduction and very specialized organometallic complexes.,[4] Because of the resiliency of the C-F bond, many pharmaceuticals contain fluorine, including several top drugs.[5]

Physical properties[edit]

Fluorocarbon are colorless and have high density, up to over twice that of water. They are not miscible with most organic solvents (e.g., ethanol, acetone, ethyl acetate and chloroform), but are miscible with some hydrocarbons (e.g., hexane in some cases). They have very low solubility in water, and water has a very low solubility in them (on the order of 10 ppm). . They have low refractive indices.

As the high electronegativity of fluorine reduces the polarizability of the atom,[2] fluorocarbons are only weakly susceptible to the fleeting dipoles that form the basis of the London dispersion force. As a result, fluorocarbons have low intermolecular attractive forces and are lipophobic in addition to being hydrophobic/non-polar. Reflecting the weak intermolecular forces these compounds exhibit low viscosities when compared to liquids of similar boiling points, low surface tension, heats of vaporization.Thus fluorocarbons find applications as oil-, water-, and stain-repellents in products such as Gore-Tex and fluoropolymer carpet coatings. The reduced participation in the London dispersion force makes the solid polytetrafluoroethylene (PTFE) slippery as it has a very low coefficient of friction. Also, the low attractive forces in fluorocarbon liquids make them compressible and gas soluble while smaller fluorocarbons are extremely volatile.[2] There are five fluoroalkane gases; tetrafluoromethane (bp −128 °C), hexafluoroethane (bp −78.2 °C), octafluoropropane (bp −36.5 °C), perfluoro-n-butane (bp −2.2 °C) and perfluoro-iso-butane (bp −1 °C). Nearly all other fluoroalkanes are liquids with the exception of perfluorocyclohexane, which sublimes at 51 °C.[6] As a result of the high gas solubility of fluorocarbon liquids, they have been the subject of medical research as blood carriers because of their oxygen solubility.[7] Fluorocarbons also have low surface energies and high dielectric strengths.[2]

The partial charges in the polarized carbon–fluorine bond

Fluoropolymers[edit]

Fluoropolymers are widely used commercially. There are two main classes, perfluorinated polyolefins, which are perfluoroalkanes. A well known example is teflon. The second class of perfluorinated polymers contain ether groups. Such polymers are produced by free-radical polymerization from the perfluoroalkene and perfluorovinyl alkyl ethers.

Fluoroalkenes and fluoroalkynes[edit]

unsaturated fluorocarbons are far more reactive than fluoroalkanes.

Fluoroalkenes polymerize more exothermically than normal alkenes.[2] Unsaturated fluorocarbons have a driving force towards sp3 hybridization due to the electronegative fluorine atoms seeking a greater share of bonding electrons with reduced s character in orbitals.[2] The most famous member of this class is tetrafluoroethylene, which converts to Teflon. Hexafluorobenzene is robust.

Although difluoroacetylene is unstable (as is typical for related alkynes, see dichloroacetylene),[2] hexafluoro-2-butyne and related fluorinated alkynes are well known.

Manufacture[edit]

The development of fluorocarbon industry coincided with World War II. Prior to that, fluorocarbons were prepared by reaction of fluorine with the hydrocarbon, i.e., direct fluorination. This highly exothermic process mainly affords smaller perfluorocarbons, such as tetrafluoromethane, hexafluoroethane, and octafluoropropane. The problem is that C-C bonds are cleaved by fluorine.[8]

A major breakthrough was the Fowler process, which allowed the large scale manufacture of fluorocarbons. This process uses cobalt fluoride in place of fluorine gas. Illustrative is the synthesis of perfluorohexane:

C6H14 + 28 CoF3 → C6F14 + 14 HF + 28 CoF2

The resulting cobalt trifluoride is then regenerated, often in a separate reactor:

2 CoF2 + F2 → 2 CoF3

Industrially, both steps are combined, for example in the manufacture of the Flutec range of fluorocarbons, using a vertical stirred bed reactor, with hydrocarbon introduced at the bottom, and fluorine introduced half way up the reactor. The fluorocarbon vapor is recovered from the top.

Electrochemical fluorination[edit]

Electrochemical fluorination (ECF) (also known as the Simons' process) involves electrolysis of a substrate dissolved in hydrogen fluoride. As fluorine is itself manufactured by the electrolysis of hydrogen fluoride, this is a rather more direct route to fluorocarbons. The process is run at low voltage (5 - 6 V) so that free fluorine is not liberated. The choice of substrate is restricted as ideally it should be soluble in hydrogen fluoride. Ethers and tertiary amines are typically employed. To make perfluorohexane, trihexylamine is used, for example:

2 N(C6H13)3 + 90 HF → 6 C6F14 + 2 NF3 + 81 H2

The perfluorinated amine will also be produced:

N(C6H13)3 + 39 HF → N(C6F13)3 + 39H2

Applications[edit]

Many fluorocarbons are useful. For example, fluorosurfactants powerfully reduce surface tension by concentrating at the liquid-air interface due to the lipophobicity of fluorocarbons,[9] due to the polar functional group added to the fluorocarbon chain. Other groups or atoms for fluorocarbon based compounds the oxygen atom incorporated into an ether group for anesthetics, and the chlorine atom for chlorofluorocarbons (CFCs). In a sharp contrast to true fluorocarbons, the chlorine atom produces a chlorine radical which degrades ozone.

Environmental and health concerns[edit]

Fluoroalkanes are generally very inert. Fluoroalkenes and fluorinated alkynes are highly reactive and toxic.

Although teflon and related fluoropolymers are nontoxic, their production often involved the use PFOS (perfluorooctane sulfonate), which bioaccumulates. Data on the human health effects of PFOA are however sparse.[10]

See also[edit]

References[edit]

  1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "fluorocarbons".
  2. ^ a b c d e f g h i Lemal DM (January 2004). "Perspective on fluorocarbon chemistry". J. Org. Chem. 69 (1): 1–11. doi:10.1021/jo0302556. PMID 14703372. 
  3. ^ a b O'Hagan D (February 2008). "Understanding organofluorine chemistry. An introduction to the C–F bond". Chem. Soc. Rev. 37 (2): 308–19. doi:10.1039/b711844a. PMID 18197347. 
  4. ^ Kiplinger JL, Richmond TG, Osterberg CE (1994). "Activation of Carbon-Fluorine Bonds by Metal Complexes". Chem. Rev. 94 (2): 373–431. doi:10.1021/cr00026a005. 
  5. ^ Ann M. Thayer "Fabulous Fluorine" Chemical and Engineering News, June 5, 2006, Volume 84, pp. 15-24. http://pubs.acs.org/cen/coverstory/84/8423cover1.html
  6. ^ http://www.ornl.gov/~webworks/cpr/v823/rpt/108771.pdf
  7. ^ Lewandowski G, Meissner E, Milchert E (August 2006). "Special applications of fluorinated organic compounds". J. Hazard. Mater. 136 (3): 385–91. doi:10.1016/j.jhazmat.2006.04.017. PMID 16759798. 
  8. ^ G. Siegemund, W. Schwertfeger, A. Feiring, B. Smart, F. Behr, H. Vogel, B. McKusick "Fluorine Compounds, Organic" in "Ullmann's Encyclopedia of Industrial Chemistry" 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a11_349
  9. ^ Mason Chemical Company: "Fluorosurfactant - Structure / Function" Accessed November 1, 2008.
  10. ^ Steenland, Kyle; Fletcher, Tony; Savitz, David A. (2010). "Epidemiologic Evidence on the Health Effects of Perfluorooctanoic Acid (PFOA)". Environmental Health Perspectives 118 (8): 1100–8. doi:10.1289/ehp.0901827. PMC 2920088. PMID 20423814. Retrieved 2011-05-11. 

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