Fluorocarbons are chemical compounds with carbon-fluorine bonds. Compounds that contain many C-F bonds often have distinctive properties, e.g., enhanced[ clarification needed ] stability, volatility, and hydrophobicity. Several fluorocarbons and their derivatives are commercial polymers, refrigerants, drugs, and anesthetics. [1]
Perfluorocarbons or PFCs, are organofluorine compounds with the formula CxFy, meaning they contain only carbon and fluorine. [2] The terminology is not strictly followed and many fluorine-containing organic compounds are also called fluorocarbons. [1] Compounds with the prefix perfluoro- are hydrocarbons, including those with heteroatoms, wherein all C-H bonds have been replaced by C-F bonds. [3] Fluorocarbons includes perfluoroalkanes, fluoroalkenes, fluoroalkynes, and perfluoroaromatic compounds.
Perfluoroalkanes are very stable because of the strength of the carbon–fluorine bond, one of the strongest in organic chemistry. [4] 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. [1] Furthermore, multiple carbon–fluorine bonds also strengthen the "skeletal" carbon–carbon bonds from the inductive effect. [1] Therefore, saturated fluorocarbons are more chemically and thermally stable than their corresponding hydrocarbon counterparts, and indeed any other organic compound. They are susceptible to attack by very strong reductants, e.g. Birch reduction and very specialized organometallic complexes. [5]
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.
As the high electronegativity of fluorine reduces the polarizability of the atom, [1] 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 and non-polar. Reflecting the weak intermolecular forces these compounds exhibit low viscosities when compared to liquids of similar boiling points, low surface tension and low heats of vaporization. The low attractive forces in fluorocarbon liquids make them compressible (low bulk modulus) and able to dissolve gas relatively well. Smaller fluorocarbons are extremely volatile. [1] There are five perfluoroalkane 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; the most notable exception is perfluorocyclohexane, which sublimes at 51 °C. [6] Fluorocarbons also have low surface energies and high dielectric strengths. [1]
In the 1960s there was a lot of interest in fluorocarbons as anesthetics. The research did not produce any anesthetics, but the research included tests on the issue of flammability, and showed that the tested fluorocarbons were not flammable in air in any proportion, though most of the tests were in pure oxygen or pure nitrous oxide (gases of importance in anesthesiology). [7] [8]
| Compound | Test conditions | Result |
|---|---|---|
| Hexafluoroethane | Lower flammability limit in oxygen | None |
| Perfluoropentane | Flash point in air | None |
| Flash point in oxygen | −6 °C | |
| Flash point nitrous oxide | −32 °C | |
| Perfluoromethylcyclohexane | Lower flammability limit in air | None |
| Lower flammability limit in oxygen | 8.3% | |
| Lower flammability limit in oxygen (50 °C) | 7.4% | |
| Lower flammability limit in nitrous oxide | 7.7% | |
| Perfluoro-1,3-dimethylcyclohexane | Lower flammability limit in oxygen (50 °C) | 5.2% |
| Perfluoromethyldecalin | Spontaneous ignition test in oxygen at 127 bar | No ignition at 500 °C |
| Spontaneous ignition in adiabatic shock wave in oxygen, 0.98 to 186 bar | No ignition | |
| Spontaneous ignition in adiabatic shock wave in oxygen, 0.98 to 196 bar | Ignition |
In 1993, 3M considered fluorocarbons as fire extinguishants to replace CFCs. [9] 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. [10] [11] When combustion does occur, toxic fumes result, including carbonyl fluoride, carbon monoxide, and hydrogen fluoride.
Perfluorocarbons dissolve relatively high volumes of gases. The high solubility of gases is attributed to the weak intermolecular interactions in these fluorocarbon fluids. [12]
The table shows values for the mole fraction, x1, of nitrogen dissolved, calculated from the Blood–gas partition coefficient, at 298.15 K (25 °C), 0.101325 MPa. [13]
| Liquid | 104x1 | Concentration ( mM ) |
|---|---|---|
| Water | 0.118 | 0.65 |
| Ethanol | 3.57 | 6.12 |
| Tetrahydrofuran | 5.21 | 6.42 |
| Acetone | 5.42 | 7.32 |
| Cyclohexane | 7.73 | 7.16 |
| Perfluoro-1,3-dimethylcyclohexane | 31.9 | 14.6 |
| Perfluoromethylcyclohexane | 33.1 | 16.9 |
The development of the fluorocarbon industry coincided with World War II. [14] 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. [15]
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 (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:
Fluoroalkanes are generally inert and non-toxic. [16] [17] [18]
Fluoroalkanes are not ozone depleting, as they contain no chlorine or bromine atoms, and they are sometimes used as replacements for ozone-depleting chemicals. [19] The term fluorocarbon is used rather loosely to include any chemical containing fluorine and carbon, including chlorofluorocarbons, which are ozone depleting. Fluoroalkanes are sometimes confused with fluorosurfactants, which significantly bioaccumulate.[ citation needed ]
Perfluoroalkanes do not bioaccumulate;[ citation needed ] those 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, [20] compared to about a week for perfluorodecalin. [21]
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.[ citation needed ] [22] The global warming potential (compared to that of carbon dioxide) of many gases can be found in the IPCC 5th assessment report, [23] with an extract below for a few perfluoroalkanes.
| Name | Chemical formula | Lifetime (y) | GWP (100 years) |
|---|---|---|---|
| PFC-14 | CF4 | 50000 | 6630 |
| PFC-116 | C2F6 | 10000 | 11100 |
| PFC-c216 | c-C3F6 | 3000 | 9200 |
| PFC-218 | C3F6 | 2600 | 8900 |
| PFC-318 | c-C4F8 | 3200 | 9540 |
The aluminium smelting industry has been a major source of atmospheric perfluorocarbons (tetrafluoromethane and hexafluoroethane especially), produced as by-product of the electrolysis process. [24] However, the industry has been actively involved in reducing emissions in recent years. [25]
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:
As well as several medical uses:
Unsaturated fluorocarbons are far more reactive than fluoroalkanes. Although difluoroacetylene is unstable (as is typical for related alkynes, see dichloroacetylene), [1] hexafluoro-2-butyne and related fluorinated alkynes are well known.
Fluoroalkenes polymerize more exothermically than normal alkenes. [1] 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. [1] The most famous member of this class is tetrafluoroethylene, which is used to manufacture polytetrafluoroethylene (PTFE), better known under the trade name Teflon.
Fluoroalkenes and fluorinated alkynes are reactive and many are toxic for example perfluoroisobutene. [29] [30] To produce polytetrafluoroethylene various fluorinated surfactants are used, in the process known as Emulsion polymerization, and the surfactant included in the polymer can bioaccumulate.
Perfluoroaromatic compounds contain only carbon and fluorine, like other fluorocarbons, but also contain an aromatic ring. The three most important examples are hexafluorobenzene, octafluorotoluene, and octafluoronaphthalene.
Perfluoroaromatic compounds can be manufactured via the Fowler process, like fluoroalkanes, but the conditions must be adjusted to prevent full fluorination. They can also be made by heating the corresponding perchloroaromatic compound with potassium fluoride at high temperature (typically 500 °C), during which the chlorine atoms are replaced by fluorine atoms. A third route is defluorination of the fluoroalkane; for example, octafluorotoluene can be made from perfluoromethylcyclohexane by heating to 500 °C with a nickel or iron catalyst. [31]
Perfluoroaromatic compounds are relatively volatile for their molecular weight, with melting and boiling points similar to the corresponding aromatic compound, as the table below shows. They have high density and are non-flammable. For the most part, they are colorless liquids. Unlike the perfluoralkanes, they tend to be miscible with common solvents.[ citation needed ]
| Compound | Melting point (°C) | Boiling point (°C) |
|---|---|---|
| Hexafluorobenzene | 5.3 | 80.5 |
| Benzene | 5.5 | 80.1 |
| Octafluorotoluene | <−70 | 102–103 |
| Toluene | −95 | 110.6 |
| Perfluoro(ethylbenzene) | 114–115 | |
| Ethylbenzene | −93.9 | 136.2 |
| Octafluoronaphthalene | 86–87 | 209 [32] |
| Naphthalene | 80.2 | 217.9 |
In chemistry, halogenation is a chemical reaction that entails the introduction of one or more halogens into a compound. Halide-containing compounds are pervasive, making this type of transformation important, e.g. in the production of polymers, drugs. This kind of conversion is in fact so common that a comprehensive overview is challenging. This article mainly deals with halogenation using elemental halogens. Halides are also commonly introduced using salts of the halides and halogen acids. Many specialized reagents exist for and introducing halogens into diverse substrates, e.g. thionyl chloride.
Halomethane compounds are derivatives of methane with one or more of the hydrogen atoms replaced with halogen atoms. Halomethanes are both naturally occurring, especially in marine environments, and human-made, most notably as refrigerants, solvents, propellants, and fumigants. Many, including the chlorofluorocarbons, have attracted wide attention because they become active when exposed to ultraviolet light found at high altitudes and destroy the Earth's protective ozone layer.
Oxygen fluorides are compounds of elements oxygen and fluorine with the general formula OnF2, where n = 1 to 6. Many different oxygen fluorides are known:
Tetrafluoromethane, also known as carbon tetrafluoride or R-14, is the simplest perfluorocarbon (CF4). As its IUPAC name indicates, tetrafluoromethane is the perfluorinated counterpart to the hydrocarbon methane. It can also be classified as a haloalkane or halomethane. Tetrafluoromethane is a useful refrigerant but also a potent greenhouse gas. It has a very high bond strength due to the nature of the carbon–fluorine bond.
Octafluoropropane (C3F8) is the perfluorocarbon counterpart to the hydrocarbon propane. This non-flammable synthetic material has applications in semiconductor production and medicine. It is also an extremely potent greenhouse gas.
Tetrafluoroethylene (TFE) is a fluorocarbon with the chemical formula C2F4. It is the simplest perfluorinated alkene. This gaseous species is used primarily in the industrial preparation of fluoropolymers.
Hydrogen fluoride (fluorane) is an inorganic compound with chemical formula HF. It is a very poisonous, colorless gas or liquid that dissolves in water to yield an aqueous solution termed hydrofluoric acid. It is the principal industrial source of fluorine, often in the form of hydrofluoric acid, and is an important feedstock in the preparation of many important compounds including pharmaceuticals and polymers, e.g. polytetrafluoroethylene (PTFE). HF is also widely used in the petrochemical industry as a component of superacids. Due to strong and extensive hydrogen bonding, it boils at near room temperature, much higher than other hydrogen halides.
Xenon difluoride is a powerful fluorinating agent with the chemical formula XeF
2, and one of the most stable xenon compounds. Like most covalent inorganic fluorides it is moisture-sensitive. It decomposes on contact with water vapor, but is otherwise stable in storage. Xenon difluoride is a dense, colourless crystalline solid.
Perfluorohexane, or tetradecafluorohexane, is a fluorocarbon. It is a derivative of hexane in which all of the hydrogen atoms are replaced by fluorine atoms. It is used in one formulation of the electronic cooling liquid/insulator Fluorinert for low-temperature applications due to its low boiling point of 56 °C and freezing point of −90 °C. It is odorless and colorless. Unlike typical hydrocarbons, the structure features a helical carbon backbone.
Organofluorine chemistry describes the chemistry of organofluorine compounds, organic compounds that contain a carbon–fluorine bond. Organofluorine compounds find diverse applications ranging from oil and water repellents to pharmaceuticals, refrigerants, and reagents in catalysis. In addition to these applications, some organofluorine compounds are pollutants because of their contributions to ozone depletion, global warming, bioaccumulation, and toxicity. The area of organofluorine chemistry often requires special techniques associated with the handling of fluorinating agents.
Perfluorodecalin is a fluorocarbon, a derivative of decalin in which all of the hydrogen atoms are replaced by fluorine atoms. It is chemically and biologically inert and stable up to 400 °C. Several applications make use of its ability to dissolve gases.
The carbon–fluorine bond is a polar covalent bond between carbon and fluorine that is a component of all organofluorine compounds. It is one of the strongest single bonds in chemistry, and relatively short, due to its partial ionic character. The bond also strengthens and shortens as more fluorines are added to the same carbon on a chemical compound. As such, fluoroalkanes like tetrafluoromethane are some of the most unreactive organic compounds.
The Fowler process is an industry and laboratory route to fluorocarbons, by fluorinating hydrocarbons or their partially fluorinated derivatives in the vapor phase over cobalt(III) fluoride.
Fluorine is a chemical element; it has symbol F and atomic number 9. It is the lightest halogen and exists at standard conditions as a highly toxic, pale yellow diatomic gas. Fluorine is extremely reactive, as it reacts with all other elements except for the light inert gases.
Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.
Perfluoro-1,3-dimethylcyclohexane is a fluorocarbon liquid—a perfluorinated derivative of the hydrocarbon 1,3-dimethylcyclohexane. It is chemically and biologically inert.
Perfluoromethylcyclohexane is a fluorocarbon liquid—a perfluorinated derivative of the hydrocarbon methylcyclohexane. It is chemically and biologically inert.
Perfluoromethyldecalin is a fluorocarbon liquid—a perfluorinated derivative of the hydrocarbon methyldecalin. It is chemically and biologically inert. It is mainly of interest as a blood substitute, exploiting the high solubility of air in this solvent.
Radical fluorination is a type of fluorination reaction, complementary to nucleophilic and electrophilic approaches. It involves the reaction of an independently generated carbon-centered radical with an atomic fluorine source and yields an organofluorine compound.
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