Perfluorohexane, a stable fluoroalkane liquid
Fluorocarbons, sometimes referred to as perfluorocarbons or PFCs, are organofluorine compounds that contain only carbon and fluorine bonded together in strong carbon–fluorine bonds. Fluoroalkanes that contain only single bonds are more chemically and thermally stable than alkanes. However, fluorocarbons with double bonds (fluoroalkenes) and especially triple bonds (fluoroalkynes) are more reactive than their corresponding hydrocarbons. Fluoroalkanes can serve as oil-repellent/water-repellent fluoropolymers, solvents, liquid breathing research agents, and powerful greenhouse gases. Unsaturated fluorocarbons tend to be used as reactants.
Perfluoroisobutene, a reactive and toxic fluoroalkene gas
Many chemical compounds are labeled as fluorocarbons, perfluorinated, or with the prefix perfluoro- despite containing atoms other than carbon or fluorine, such as chlorofluorocarbons and perfluorinated compounds; however, these molecules are fluorocarbon derivatives, and not true fluorocarbons. Fluorocarbon derivatives share many of the properties of fluorocarbons, while also possessing new properties due to the inclusion of new atoms. For example, fluorocarbon derivatives can function as fluoropolymers, refrigerants, solvents, anesthetics, fluorosurfactants, and ozone depletors.
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Contents
- 1 Usage of term
- 2 Properties
- 2.1 Physical properties
- 2.1.1 London dispersion force reduction
- 2.2 Fluoroalkane stability
- 2.3 Fluoroalkene and fluoroalkyne reactivity
- 3 Manufacture
- 3.1 Fowler process
- 3.2 Electrochemical fluorination
- 4 Derivatives
- 4.1 Fluorosurfactants
- 4.2 Anesthetics
- 4.3 Halogenated derivatives
- 4.4 Hydrofluorocarbons
- 5 Environmental and health concerns
- 6 See also
- 7 References
- 8 External links
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Usage of term
The formal IUPAC definition of a fluorocarbon is a molecule consisting wholly of fluorine and carbon.[1] However, other fluorocarbon based molecules that are not technically fluorocarbons are commonly referred to as fluorocarbons,[2] because of similar structures and identical properties. Compounds with atoms other than carbon and fluorine are not true fluorocarbons and they are considered as fluorocarbon derivatives in a separate section below.
Properties
Perfluorodecalin, a dense clear liquid
Physical properties
Fluorocarbon liquids are colorless. They have high density, up to over twice that of water, due to their high molecular weight. Low intermolecular forces give the liquids low viscosities when compared to liquids of similar boiling points. Also, low surface tension, heats of vaporization, and refractive indices are notable. 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). The number of carbon atoms in a fluorocarbon molecule largely determines most physical properties. The greater the number of carbon atoms, the higher the boiling point, density, viscosity, surface tension, critical properties, vapor pressure and refractive index. Gas solubility decreases as carbon atoms increase, while melting point is determined by other factors as well and is thus not readily predicted.
Tetrafluoromethane, a fluorocarbon gas
London dispersion force reduction
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. 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.[3] 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.[4] Fluorocarbons also have low surface energies and high dielectric strengths.[2]
The partial charges in the polarized carbon–fluorine bond
Fluoroalkane stability
Fluorocarbons with only single bonds are very stable because of the strength and nature of the carbon–fluorine bond. It is called the strongest bond in organic chemistry.[5] Its strength is a result of the electronegativity of fluorine imparting partial ionic character through partial charges on the carbon and fluorine atoms.[5] 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. However, fluoroalkanes are not inert. They are susceptible to reduction through the Birch reduction.
Tetrafluoroethylene, an important reactant
Fluoroalkene and fluoroalkyne reactivity
When fluorocarbons are unsaturated, they are less stable and more reactive than fluoroalkanes, or comparable hydrocarbons, due to the electronegativity of fluorine. The reactivity of the simplest fluoroalkyne, difluoroacetylene, is an example of this instability; difluoroacetylene easily polymerizes.[2] Another example is fluorofullerene, which has weaker and longer carbon–fluorine bonds than saturated fluorocarbons.[6] It is reactive towards nucleophiles and hydrolyzes in solution.[6] Additionally, the polymerization of the fluoroalkene tetrafluoroethylene (which results in PTFE) is more energetically favorable than that of ethylene.[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]
One notable exception to this trend is fluorobenzene, which is stabilized by its aromaticity.
Manufacture
Prior to World War II, the only known route to fluorocarbons was by direct reaction of fluorine with the hydrocarbon. This highly exothermic process was capable only of synthesising tetrafluoromethane, hexafluoroethane and octafluoropropane; larger hydrocarbons decomposed in the extreme conditions. The Manhattan project saw the need for some very robust chemicals, including a wider range of fluorocarbons, requiring new manufacturing methods. The so-called "catalytic" method involved reacting fluorine and hydrocarbon on a bed of gold-plated copper turnings, the metal removing the heat of the reaction (so not really acting as a catalyst at all), allowing larger hydrocarbons to survive the process. However, it was the Fowler process that allowed the large scale manufacture of fluorocarbons required for the Manhattan project.
Fowler process
The Fowler process uses cobalt fluoride to moderate the reaction. In the laboratory, this is typically done in two stages, the first stage being fluorination of cobalt difluoride to cobalt trifluoride.
- 2 CoF2 + F2 → 2 CoF3
During the second stage, in this instance to make perfluorohexane, the hydrocarbon feed is introduced and is fluorinated by the cobalt trifluoride, which is converted back to cobalt difluoride. Both stages are performed at high temperature.
- C6H14 + 28 CoF3 → C6F14 + 14 HF + 28 CoF2
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
An alternative technique, 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 + 45 H2
The perfluorinated amine will also be produced:
- N(C6H13)3 + 42 HF → 2 N(C6F13)3 + 21H2
Both of these products, and others, are manufactured by 3M as part of the Fluorinert range.
Derivatives
See also: Fluoropolymer, Fluorotelomer, Perfluorinated compound, and Organofluorine compound
Fluorocarbon derivatives are highly fluorinated molecules that can be commonly referred to as fluorocarbons. They are economically useful because they share part or nearly all of the properties of fluorocarbons. Some fluorocarbon derivatives have markedly different properties than fluorocarbons. For example, fluorosurfactants powerfully reduce surface tension by concentrating at the liquid-air interface due to the lipophobicity of fluorocarbons,[7] 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.
Fluorosurfactants
- Perfluorooctanesulfonic acid (PFOS)
- Perfluorooctanoic acid (PFOA)
- Perfluorononanoic acid
Anesthetics
- Methoxyflurane (contains chlorine)
- Enflurane (contains chlorine)
- Isoflurane (contains chlorine)
- Sevoflurane
- Desflurane
Halogenated derivatives
- Polychlorotrifluoroethylene ([CFClCF2]n)
- Perfluorooctyl bromide (perflubron)
- Dichlorodifluoromethane
- Chlorodifluoromethane
Hydrofluorocarbons
- Polyvinylidene fluoride ([CH2CF2]n)
- Tetrafluoroethane
Environmental and health concerns
Despite the presence of some natural fluorocarbons such as tetrafluoromethane, which has been reported in rocks,[8] man-made fluorocarbons are potent greenhouse gases.
Another important aspect in terms of environmental concerns, is certain fluorocarbons' bioaccumulative properties. Fluorocarbons are extremely stable and can be stored in the bodies of both humans and animals. Examples of harmful fluorocarbons include PFOA (perfluorooctanoic acid) and PFOS (perfluorooctane sulfonate), frequently present in water resistant textiles and sprays conferring water resistant properties to textiles.[9] Data from animal studies of PFOA indicate that it can cause several types of tumors and neonatal death and may have toxic effects on the immune, liver, and endocrine systems. Data on the human health effects of PFOA are however sparse.[10]
The fluorocarbon, PFOA and PFOS have both been subject for numerous investigations by the EU and the United States Environmental Protection Agency (EPA) regarding them being harmful to the environment.[9]
See also
- Category:Fluorocarbons
- Trifluoromethyl
- Fluorographene
- Liquid breathing
References
- ^ Nic, M.; Jirat, J.; Kosata, B., eds. (2006–). "fluorocarbons". IUPAC Compendium of Chemical Terminology (Online ed.). doi:10.1351/goldbook.F02459. ISBN 0-9678550-9-8. http://goldbook.iupac.org/F02459.html.
- ^ 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.
- ^ http://www.ornl.gov/~webworks/cpr/v823/rpt/108771.pdf
- ^ 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.
- ^ 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.
- ^ a b 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. http://pubs.acs.org/doi/abs/10.1021/cr00026a005.
- ^ Mason Chemical Company: "Fluorosurfactant - Structure / Function" Accessed November 1, 2008.
- ^ Murphy CD, Schaffrath C, O'Hagan D.: "Fluorinated natural products: the biosynthesis of fluoroacetate and 4-fluorothreonine in Streptomyces cattleya" Chemosphere. 2003 Jul;52(2):455-61.
- ^ a b US Environmental Protection Agency. "FAQ". Perfluorooctanoic Acid (PFOA) and Fluorinated Telomers. http://www.epa.gov/oppt/pfoa/pubs/faq.html#risks. Retrieved 11 May 2011.
- ^ 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. http://ehp03.niehs.nih.gov/article/fetchArticle.action?articleURI=info%3Adoi%2F10.1289%2Fehp.0901827. Retrieved 2011-05-11.
External links
- Fluorocarbons and Sulphur Hexafluoride, proposed by the European Fluorocarbons Technical Committee
- CFCs and Ozone Depletion Freeview video provided by the Vega Science Trust.
- Introduction to fluoropolymers
- Organofluorine chemistry by Graham Sandford
