Chemical properties Perfluoroalkanes are very stable because of the strength of the
carbon–fluorine bond, one of the strongest in organic chemistry. Its strength is a result of the electronegativity of fluorine imparting partial ionic character through
partial charges on the carbon and fluorine atoms, which shorten and strengthen the bond (compared to carbon-hydrogen bonds) through favorable
covalent 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. Fluorocarbons 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. {{Image frame|content=\overset{\delta+}{C}-\overset{\delta-}{F}|caption=The partial charges in the polarized carbon–fluorine bond|width=150}} As the high
electronegativity of fluorine reduces the
polarizability of the atom, Fluorocarbons also have low
surface energies and high
dielectric strengths. In 1993, 3M considered fluorocarbons as fire extinguishants to replace CFCs. This extinguishing effect has been attributed to their high
heat capacity, which takes heat away from the fire. It has been suggested that an atmosphere containing a significant percentage of perfluorocarbons on a space station or similar would prevent fires altogether. When combustion does occur, toxic fumes result, including
carbonyl fluoride,
carbon monoxide, and
hydrogen fluoride.
Gas dissolving properties Perfluorocarbons dissolve relatively high volumes of gases. The high solubility of gases is attributed to the weak intermolecular interactions in these fluorocarbon fluids. The table shows values for the mole fraction, , of nitrogen dissolved, calculated from the
Blood–gas partition coefficient, at 298.15 K (25 °C), 0.101325 MPa.
Manufacture The development of the fluorocarbon industry coincided with
World War II. Prior to that, fluorocarbons were prepared by reaction of fluorine with the hydrocarbon, i.e., direct fluorination. Because C-C bonds are readily cleaved by fluorine, direct fluorination mainly affords smaller perfluorocarbons, such as tetrafluoromethane, hexafluoroethane, and octafluoropropane.
Fowler process A major breakthrough that allowed the large scale manufacture of fluorocarbons was the
Fowler process. In this process,
cobalt trifluoride is used as the source of fluorine. Illustrative is the synthesis of
perfluorohexane: : The resulting cobalt difluoride is then regenerated, sometimes in a separate reactor: : Industrially, both steps are combined, for example in the manufacture of the Flutec range of fluorocarbons by F2 chemicals Ltd, using a vertical stirred bed reactor, with hydrocarbon introduced at the bottom, and fluorine introduced halfway up the reactor. The fluorocarbon vapor is recovered from the top.
Electrochemical fluorination 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, ECF is a rather more direct route to fluorocarbons. The process proceeds 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: : The perfluorinated amine will also be produced: :
Environmental and health concerns Fluoroalkanes are generally inert and non-toxic. Fluoroalkanes are not
ozone depleting, as they contain no chlorine or bromine atoms, and they are sometimes used as replacements for ozone-depleting chemicals. The term fluorocarbon is used rather loosely to include any chemical containing fluorine and carbon, including
chlorofluorocarbons, which are ozone depleting. Perfluoroalkanes used in medical procedures are rapidly excreted from the body, primarily via expiration with the rate of excretion as a function of the vapour pressure; the half-life for
octafluoropropane is less than 2 minutes, compared to about a week for perfluorodecalin. Low-boiling perfluoroalkanes are potent
greenhouse gases, in part due to their very long atmospheric lifetime, and their use is covered by the
Kyoto Protocol. The
global warming potential (compared to that of carbon dioxide) of many gases can be found in the IPCC 5th assessment report, with an extract below for a few perfluoroalkanes. The aluminium smelting industry has been a major source of atmospheric perfluorocarbons (
tetrafluoromethane and
hexafluoroethane especially), produced as by-product of the electrolysis process. However, the industry has been actively involved in reducing emissions in recent years.
Applications As they are inert, perfluoroalkanes have essentially no chemical uses, but their physical properties have led to their use in many diverse applications. These include: •
Perfluorocarbon tracer •
Liquid dielectric •
Chemical vapor deposition •
Organic Rankine cycle • Fluorous biphasic catalysis • Cosmetics • Ski waxes As well as several medical uses: •
Contrast-enhanced ultrasound •
Oxygen Therapeutics •
Blood substitute •
Liquid breathing • Eye surgery • Tattoo removal ==Fluoroalkenes and fluoroalkynes==