Nuclear isomer

From Wikipedia, the free encyclopedia

A nuclear isomer is a metastable state of an atomic nucleus, in which one or more nucleons (protons or neutrons) occupy higher energy levels than in the ground state of the same nucleus. "Metastable" describes nuclei whose excited states have half-lives 100 to 1000 times longer than the half-lives of the excited nuclear states that decay with a "prompt" half life (ordinarily on the order of 10⁻¹² seconds). The term "metastable" is usually restricted to isomers with half-lives of 10⁻⁹ seconds or longer. Some references recommend 5 × 10⁻⁹ seconds to distinguish the metastable half life from the normal "prompt" gamma emission half life. Occasionally the half-lives are far longer than this and can last minutes, hours, or years.

 

Sometimes, the gamma decay from a metastable state is referred to as isomeric transition, but this process typically resembles shorter-lived gamma decays in all external aspects with the exception of the long-lived nature of the meta-stable parent nuclear isomer. The longer lives of nuclear isomers' metastable states are often due to the larger degree of nuclear spin change which must be involved in their gamma emission to reach the ground state. This high spin change causes these decays to be forbidden transitions and delayed. Delays in emission are caused by low or high available decay energy.

The first nuclear isomer and decay-daughter system (uranium X₂/uranium Z, now known as ^234m
₉₁Pa
/²³⁴
₉₁Pa
) was discovered by Otto Hahn in 1921.

Nuclei of nuclear isomers

The nucleus of a nuclear isomer occupies a higher energy state than the non-excited nucleus existing at ground state. In an excited state, one or more of the protons or neutrons in a nucleus occupy a nuclear orbital of higher energy than an available nuclear orbital. These states are analogous to excited states of electrons in atoms. 

When excited atomic states decay, energy is released by fluorescence. In electronic transitions, this process usually involves emission of light near the visible range. The amount of energy released is related to bond-dissociation energy or ionization energy and is usually in the range of a few to few tens of eV per bond. 

However, a much stronger type of binding energy, the nuclear binding energy, is involved in nuclear processes. Due to this, most nuclear excited states decay by gamma ray emission. For example, a well-known nuclear isomer used in various medical procedures is ^99m
₄₃Tc
which decays with a half-life of about 6 hours by emitting a gamma ray of 140 keV of energy; this is close to the energy of medical diagnostic X-rays. 

Nuclear isomers have long half lives because their gamma decay is "forbidden" from the large change in nuclear spin needed to emit a gamma ray. For example, ^180m
₇₃Ta
has a spin of 9 and must gamma decay to ¹⁸⁰
₇₃Ta
with a spin of 1. Similarly, ^99m
₄₃Tc
has a spin of 1/2 and must gamma decay to ⁹⁹
₄₃Tc
with a spin of 9/2. 

While most metastable isomers decay through gamma ray emission, they can also decay through internal conversion. During internal conversion, energy of nuclear de-excitation is not emitted as a gamma ray but is instead used to accelerate one of the inner electrons of the atom. These excited electrons then leave at a high speed. This occurs because inner atomic electrons penetrate the nucleus where they are subject to the intense electric fields created when the protons of the nucleus re-arrange in a different way. 

In nuclei which are far from stability in energy even more decay modes are known.

Metastable isomers

Metastable isomers can be produced through nuclear fusion or other nuclear reactions. A nucleus produced this way generally starts its existence in an excited state that relaxes through the emission of one or more gamma rays or conversion electrons. Sometimes the de-excitation does not completely proceed rapidly to the nuclear ground state. This usually occurs when the formation of an intermediate excited state has a spin far different from that of the ground state. Gamma-ray emission is hindered if the spin of the post-emission state varies greatly from that of the emitting state especially if the excitation energy is low. The excited state in this situation is a good candidate to be metastable if there are no other states of intermediate spin with excitation energies less than that of the metastable state. 

Metastable isomers of a particular isotope are usually designated with an "m". This designation is placed after the mass number of the atom; for example, cobalt-58m is abbreviated ^58m
₂₇Co
, where 27 is the atomic number of cobalt. For isotopes with more than one metastable isomer, "indices" are placed after the designation, and the labeling becomes m1, m2, m3, and so on. Increasing indices, m1, m2, etc., correlate with increasing levels of excitation energy stored in each of the isomeric states (e.g., hafnium-178m2 or ^178m2
₇₂Hf
). 

A different kind of metastable nuclear state (isomer) is the fission isomer or shape isomer. Most actinide nuclei in their ground states are not spherical, but rather prolate spheroidal, with an axis of symmetry longer than the other axes similar to an American football or rugby ball. This geometry can result in quantum-mechanical states where the distribution of protons and neutrons is so much further from spherical geometry that de-excitation to the nuclear ground state is strongly hindered. In general, these states either de-excite to the ground state far more slowly than a "usual" excited state, or they undergo spontaneous fission with half-lives of the order of nanoseconds or microseconds—a very short time, but many orders of magnitude longer than the half-life of a more-usual nuclear excited state. Fission isomers are usually denoted with a postscript or superscript "f" rather than "m", so that a fission isomer, e.g. of plutonium-240, is denoted plutonium-240f or ^240f
₉₄Pu
.

Nearly-stable isomers

Most nuclear excited states are very unstable and "immediately" radiate away the extra energy after existing on the order of 10⁻¹² seconds. As a result, the characterization "nuclear isomer" is usually applied only to configurations with half-lives of 10⁻⁹ seconds or longer. Quantum mechanics predicts that certain atomic species will possess isomers with unusually long lifetimes even by this stricter standard and have interesting properties. Some nuclear isomers are so long-lived that they are relatively stable and can be produced and observed in quantity.

The most stable nuclear isomer occurring in nature is ^180m
₇₃Ta
which is present in all tantalum samples at about 1 part in 8,300. Its half-life is at least 10¹⁵ years, markedly longer than the age of the universe. The low excitation energy of the isomeric state causes both gamma de-excitation to the ¹⁸⁰Ta
ground state (which itself is radioactive by beta decay, with a half-life of only 8 hours) and direct beta decay to hafnium or tungsten to be suppressed due to spin mismatches. The origin of this isomer is mysterious, though it is believed to have been formed in supernovae (as are most other heavy elements). Were it to relax to its ground state, it would release a photon with a photon energy of 75 keV. 

It was first reported in 1988 by Collins that ^180mTa
can be forced to release its energy by weaker X-rays. This way of de-excitation had never been observed; however, the de-excitation of ^180mTa
via resonant photo-excitation of intermediate high levels of this nucleus (E~1 MeV) was found in 1999 by Belic and co-workers in the Stuttgart nuclear physics group.

^178m2
₇₂Hf
is another reasonably stable nuclear isomer which possesses a half-life of 31 years and the highest excitation energy of any comparably long-lived isomer. One gram of pure ^178m2Hf
contains approximately 1.33 gigajoules of energy, the equivalent of exploding about 315 kg (694 lb) of TNT. In the natural decay of ^178m2Hf
, the energy is released as gamma rays with a total energy of 2.45 MeV. As with ^180mTa
, there are disputed reports that ^178m2Hf
can be stimulated into releasing its energy. Due to this, the substance is being studied as a possible source for gamma ray lasers. These reports indicate that the energy is released very quickly, so that ^178m2Hf
can produce extremely high powers (on the order of exawatts). Other isomers have also been investigated as possible media for gamma-ray stimulated emission.

Holmium's nuclear isomer, ^166m1
₆₇Ho
has a half-life of 1,200 years, which is nearly the longest half-life of any holmium radionuclide. Only ¹⁶³Ho
, with a half-life of 4,570 years, is longer. 

²²⁹
₉₀Th
has a remarkably low-lying metastable isomer, estimated at only 7.8±0.5 eV above the ground state. After years of failure and one notable false alarm, this decay was directly observed in 2016, producing a gamma ray (defined by its origin, not its wavelength) in the ultraviolet range. The observed energy was between 6.3 and 18.3 eV (200–70 nm). The range is broad because the experiment was optimized for detection rather than precision measurement.

High spin suppression of decay

The most common mechanism for suppression of gamma decay of excited nuclei, and thus the existence of a metastable isomer, is lack of a decay route for the excited state that will change nuclear angular momentum along any given direction by the most common amount of 1 quantum unit ħ in the spin angular momentum. This change is necessary to emit a gamma photon which has a spin of 1 unit in this system. Integral changes of 2, 3, 4, and more units in angular momentum are possible, but the emitted photons carry off the additional angular momentum. Changes of more than 1 unit are known as forbidden transitions. Each additional unit of spin change larger than 1 that the emitted gamma ray must carry inhibits decay rate by about 5 orders of magnitude. The highest known spin change of 8 units occurs in the decay of ^180mTa, which suppresses its decay by a factor of 10³⁵ from that associated with 1 unit. Instead of a natural gamma decay half life of 10⁻¹² seconds, it has a half life of more than 10²³ seconds, or at least 3 × 10¹⁵ years, and thus has yet to be observed to decay. 

Gamma emission is impossible when the nucleus begins in a zero-spin state, as such an emission would not conserve angular momentum.

Applications

Hafnium and tantalum isomers have been considered in some quarters as weapons that could be used to circumvent the Nuclear Non-Proliferation Treaty, since it is claimed they can be induced to emit very strong gamma radiation. This claim is generally discounted. DARPA has (or had) a program to investigate this use of both nuclear isomers. The potential to trigger an abrupt release of energy from nuclear isotopes, a prerequisite to their use in such weapons, is disputed. Nonetheless a 12-member Hafnium Isomer Production Panel (HIPP) was created to assess means of mass-producing the isotope.

Technetium isomers ^99m
₄₃Tc
(with a half-life of 6.01 hours) and ^95m
₄₃Tc
(with a half-life of 61 days) are used in medical and industrial applications.

Nuclear batteries

Nuclear decay pathways for the conversion of lutetium-177^m to hafnium-177

 

Nuclear batteries use small amounts (milligrams and microcuries) of radioisotopes with high energy densities. In one design, radioactive material sits atop a device with adjacent layers of P-type and N-type silicon. Ionizing radiation directly penetrates the junction and creates electron-hole pairs. Nuclear isomers could replace other isotopes, and with further development, it may be possible to turn them on and off by triggering decay as needed. Current candidates for such use include ¹⁰⁸Ag, ¹⁶⁶Ho, ¹⁷⁷Lu, and ²⁴¹Am. As of 2004 the only isomer which had been successfully triggered was ^180mTa which required more photon energy to trigger than was released.

An isotope such as ¹⁷⁷Lu releases gamma rays by decay through a series of internal energy levels within the nucleus, and it is thought that by learning the triggering cross sections with sufficient accuracy, it may be possible to create energy stores that are 10⁶ times more concentrated than high explosive or other traditional chemical energy storage.

Decay processes

An isomeric transition (IT) is the decay of a nuclear isomer to a lower-energy nuclear state. This radioactive decay process involves emission of a gamma ray. The actual process can have two effects:

• γ (gamma) emission (emission of a high-energy photon)
• internal conversion (the energy is used to excite the atom's electrons)

Isomers may decay into other elements, though the rate of decay may differ between isomers. For example, ^177mLu can beta decay to ¹⁷⁷Hf with a half-life of 160.4 d, or it can undergo isomeric transition to ¹⁷⁷Lu with a half-life of 160.4 d which then beta decays to ¹⁷⁷Hf with a half-life of 6.68 d.

The emission of a gamma ray from an excited nuclear state allows the nucleus to lose energy and reach a lower energy state, sometimes its ground state. In certain cases, the excited nuclear state following a nuclear reaction or other type of radioactive decay can become a metastable nuclear excited state. Some nuclei are able to stay in this metastable excited state for minutes, hours, days, or occasionally far longer, before undergoing gamma decay, in which they emit a gamma ray. 

The process of isomeric transition is similar to any gamma emission from any excited nuclear state, but differs by involving excited metastable states of nuclei with longer half lives. These states are created, as all nuclei that undergo gamma radioactive decay, following the emission of an alpha particle, beta particle, or occasionally other types of particles that leave the nucleus in an excited state. 

The gamma ray may transfer its energy directly to one of the most tightly bound electrons causing that electron to be ejected from the atom, a process termed the photoelectric effect. This should not be confused with the internal conversion process, in which no gamma ray photon is produced as an intermediate particle.

Organofluorine chemistry

From Wikipedia, the free encyclopedia

Some important organofluorine compounds.

 

A: fluoromethane
B: isoflurane
C: a CFC
D: an HFC
E: triflic acid
F: Teflon
G: PFOS
H: fluorouracil
I: fluoxetine

 

Organofluorine chemistry describes the chemistry of the organofluorines, organic compounds that contain the 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.

The carbon–fluorine bond

Fluorine has several distinctive differences from all other substituents encountered in organic molecules. As a result, the physical and chemical properties of organofluorines can be distinctive in comparison to other organohalogens.

1. The carbon–fluorine bond is one of the strongest in organic chemistry (an average bond energy around 480 kJ/mol). This is significantly stronger than the bonds of carbon with other halogens (an average bond energy of e.g. C-Cl bond is around 320 kJ/mol) and is one of the reasons why fluoroorganic compounds have high thermal and chemical stability.
2. The carbon–fluorine bond is relatively short (around 1.4 Å).
3. The Van der Waals radius of the fluorine substituent is only 1.47 Å, which is shorter than in any other substituent and is close to that of hydrogen (1.2 Å). This, together with the short bond length, is the reason why there is no steric strain in polyfluorinated compounds. This is another reason for their high thermal stability. In addition, the fluorine substituents in polyfluorinated compounds efficiently shield the carbon skeleton from possible attacking reagents. This is another reason for the high chemical stability of polyfluorinated compounds.
4. Fluorine has the highest electronegativity of all elements: 3.98. This causes the high dipole moment of C-F bond (1.41 D).
5. Fluorine has the lowest polarizability of all atoms: 0.56 ${\displaystyle \times }$ 10⁻²⁴ cm³. This causes very weak dispersion forces between polyfluorinated molecules and is the reason for the often-observed boiling point reduction on fluorination as well as for the simultaneous hydrophobicity and lipophobicity of polyfluorinated compounds whereas other perhalogenated compounds are more lipophilic.

In comparison to aryl chlorides and bromides, aryl fluorides form Grignard reagents only reluctantly. On the other hand, aryl fluorides, e.g. fluoroanilines and fluorophenols, often undergo nucleophilic substitution efficiently.

Types of organofluorine compounds

Fluorocarbons

Formally, fluorocarbons only contain carbon and fluorine. Sometimes they are called perfluorocarbons. They can be gases, liquids, waxes, or solids, depending upon their molecular weight. The simplest fluorocarbon is the gas tetrafluoromethane (CF₄). Liquids include perfluorooctane and perfluorodecalin. While fluorocarbons with single bonds are stable, unsaturated fluorocarbons are more reactive, especially those with triple bonds. Fluorocarbons are more chemically and thermally stable than hydrocarbons, reflecting the relative inertness of the C-F bond. They are also relatively lipophobic. Because of the reduced intermolecular van der Waals interactions, fluorocarbon-based compounds are sometimes used as lubricants or are highly volatile. Fluorocarbon liquids have medical applications as oxygen carriers. 

The structure of organofluorine compounds can be distinctive. As shown below, perfluorinated aliphatic compounds tend to segregate from hydrocarbons. This "like dissolves like effect" is related to the usefulness of fluorous phases and the use of PFOA in processing of fluoropolymers. In contrast to the aliphatic derivatives, perfluoroaromatic derivatives tend to form mixed phases with nonfluorinated aromatic compounds, resulting from donor-acceptor interactions between the pi-systems. 

Segregation of alkyl and perfluoroalkyl substituents.

 

Packing in a crystal pentafluorotolan (C₆F₅CCC₆H₅), illustrating the donor-acceptor interactions between the fluorinated and nonfluorinated rings.

Fluoropolymers

Polymeric organofluorine compounds are numerous and commercially significant. They range from fully fluorinated species, e.g. PTFE to partially fluorinated, e.g. polyvinylidene fluoride ([CH₂CF₂]_n) and polychlorotrifluoroethylene ([CFClCF₂]_n). The fluoropolymer polytetrafluoroethylene (PTFE/Teflon) is a solid.

Hydrofluorocarbons

Hydrofluorocarbons (HFCs), organic compounds that contain fluorine and hydrogen atoms, are the most common type of organofluorine compounds. They are commonly used in air conditioning and as refrigerants in place of the older chlorofluorocarbons such as R-12 and hydrochlorofluorocarbons such as R-21. They do not harm the ozone layer as much as the compounds they replace; however, they do contribute to global warming. Their atmospheric concentrations and contribution to anthropogenic greenhouse gas emissions are rapidly increasing, causing international concern about their radiative forcing. 

Fluorocarbons with few C-F bonds behave similarly to the parent hydrocarbons, but their reactivity can be altered significantly. For example, both uracil and 5-fluorouracil are colourless, high-melting crystalline solids, but the latter is a potent anti-cancer drug. The use of the C-F bond in pharmaceuticals is predicated on this altered reactivity. Several drugs and agrochemicals contain only one fluorine center or one trifluoromethyl group. 

Unlike other greenhouse gases in the Paris Agreement, hydrofluorocarbons have other international negotiations.

In September 2016, the so-called New York Declaration urged a global reduction in the use of HFCs. On 15 October 2016, due to these chemicals contribution to climate change, negotiators from 197 nations meeting at the summit of the United Nations Environment Programme in Kigali, Rwanda reached a legally-binding accord to phase out hydrofluorocarbons (HFCs) in an amendment to the Montreal Protocol.

Fluorocarbenes

As indicated throughout this article, fluorine-substituents lead to reactivity that differs strongly from classical organic chemistry. The premier example is difluorocarbene, CF₂, which is a singlet whereas carbene (CH₂) has a triplet ground state. This difference is significant because difluorocarbene is a precursor to tetrafluoroethylene.

Perfluorinated compounds

Perfluorinated compounds are fluorocarbon derivatives, as they are closely structurally related to fluorocarbons. However, they also possess new atoms such as nitrogen, iodine, or ionic groups, such as perfluorinated carboxylic acids.

Methods for preparation of C–F bonds

Organofluorine compounds are prepared by numerous routes, depending on the degree and regiochemistry of fluorination sought and the nature of the precursors. The direct fluorination of hydrocarbons with F₂, often diluted with N₂, is useful for highly fluorinated compounds:

R
₃CH + F
₂ → R
₃CF + HF

Such reactions however are often unselective and require care because hydrocarbons can uncontrollably "burn" in F
₂, analogous to the combustion of hydrocarbon in O
₂. For this reason, alternative fluorination methodologies have been developed. Generally, such methods are classified into two classes.

Electrophilic fluorination

Electrophilic fluorination rely on sources of "F⁺". Often such reagents feature N-F bonds, for example F-TEDA-BF₄. Asymmetric fluorination, whereby only one of two possible enantiomeric products are generated from a prochiral substrate, rely on electrophilic fluorination reagents. Illustrative of this approach is the preparation of a precursor to anti-inflammatory agents:

Electrosynthetic methods

A specialized but important method of electrophilic fluorination involves electrosynthesis. The method is mainly used to perfluorinate, i.e. replace all C–H bonds by C–F bonds. The hydrocarbon is dissolved or suspended in liquid HF, and the mixture is electrolyzed at 5–6 V using Ni anodes. The method was first demonstrated with the preparation of perfluoropyridine (C
₅F
₅N) from pyridine (C
₅H
₅N). Several variations of this technique have been described, including the use of molten potassium bifluoride or organic solvents.

Nucleophilic fluorination

The major alternative to electrophilic fluorination is, naturally, nucleophilic fluorination using reagents that are sources of "F⁻," for Nucleophilic displacement typically of chloride and bromide. Metathesis reactions employing alkali metal fluorides are the simplest.

R
₃CCl + MF → R
₃CF + MCl (M = Na, K, Cs)

Alkyl monofluorides can be obtained from alcohols and Olah reagent (pyridinium fluoride) or another fluoridating agents. 

The decomposition of aryldiazonium tetrafluoroborates in the Sandmeyer or Schiemann reactions exploit fluoroborates as F⁻ sources.

ArN
₂BF
₄ → ArF + N
₂ + BF
₃

Although hydrogen fluoride may appear to be an unlikely nucleophile, it is the most common source of fluoride in the synthesis of organofluorine compounds. Such reactions are often catalysed by metal fluorides such as chromium trifluoride. 1,1,1,2-Tetrafluoroethane, a replacement for CFC's, is prepared industrially using this approach:

Cl₂C=CClH + 4 HF → F₃CCFH₂ + 3 HCl

Notice that this transformation entails two reaction types, metathesis (replacement of Cl⁻ by F⁻) and hydrofluorination of an alkene.

Deoxofluorination

Deoxofluorination agents effect the replacement hydroxyl and carbonyl groups with one and two fluorides, respectively. One such reagent, useful for fluoride for oxide exchange in carbonyl compounds, is sulfur tetrafluoride:

RCO₂H + SF₄ → RCF₃ + SO₂ + HF

Alternates to SF₄ include the diethylaminosulfur trifluoride (DAST, NEt₂SF₃) and bis(2-methoxyethyl)aminosulfur trifluoride (deoxo-fluor). These organic reagents are easier to handle and more selective:

From fluorinated building blocks

Many organofluorine compounds are generated from reagents that deliver perfluoroalkyl and perfluoroaryl groups. (Trifluoromethyl)trimethylsilane, CF₃Si(CH₃)₃, is used as a source of the trifluoromethyl group, for example. Among the available fluorinated building blocks are CF₃X (X = Br, I), C₆F₅Br, and C₃F₇I. These species form Grignard reagents that then can be treated with a variety of electrophiles. The development of fluorous technologies (see below, under solvents) is leading to the development of reagents for the introduction of "fluorous tails".

A special but significant application of the fluorinated building block approach is the synthesis of tetrafluoroethylene, which is produced on a large-scale industrially via the intermediacy of difluorocarbene. The process begins with the thermal (600-800 °C) dehydrochlorination of chlorodifluoromethane:

CHClF₂ → CF₂ + HCl

2 CF₂ → C₂F₄

Sodium fluorodichloroacetate (CAS# 2837-90-3) is used to generate chlorofluorocarbene, for cyclopropanations.

¹⁸F-Delivery methods

The usefulness of fluorine-containing radiopharmaceuticals in ¹⁸F-positron emission tomography has motivated the development of new methods for forming C–F bonds. Because of the short half-life of ¹⁸F, these syntheses must be highly efficient, rapid, and easy. Illustrative of the methods is the preparation of fluoride-modified glucose by displacement of a triflate by a labeled fluoride nucleophile:

Biological role

Biologically synthesized organofluorines have been found in microorganisms and plants, but not animals. The most common example is fluoroacetate, which occurs as a plant defence against herbivores in at least 40 plants in Australia, Brazil and Africa. Other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate which are all believed to be biosynthesized in biochemical pathways from the intermediate fluoroacetaldehyde. Adenosyl-fluoride synthase is an enzyme capable of biologically synthesizing the carbon–fluorine bond. Man made carbon–fluorine bonds are commonly found in pharmaceuticals and agrichemicals because it adds stability to the carbon framework; also, the relatively small size of fluorine is convenient as fluorine acts as an approximate bioisostere of the hydroxyl group. Introducing the carbon–fluorine bond to organic compounds is the major challenge for medicinal chemists using organofluorine chemistry, as the carbon–fluorine bond increases the probability of having a successful drug by about a factor of ten. An estimated 20% of pharmaceuticals, and 30–40% of agrichemicals are organofluorines, including several of the top drugs. Examples include 5-fluorouracil, fluoxetine (Prozac), paroxetine (Paxil), ciprofloxacin (Cipro), mefloquine, and fluconazole.

Applications

Organofluorine chemistry impacts many areas of everyday life and technology. The C-F bond is found in pharmaceuticals, agrichemicals, fluoropolymers, refrigerants, surfactants, anesthetics, oil-repellents, catalysis, and water-repellents, among others.

Pharmaceuticals and agrochemicals

The carbon-fluorine bond is commonly found in pharmaceuticals and agrochemicals because it is generally metabolically stable and fluorine acts as a bioisostere of the hydrogen atom. An estimated one fifth of pharmaceuticals contain fluorine, including several of the top drugs. Examples include 5-fluorouracil, flunitrazepam (Rohypnol), fluoxetine (Prozac), paroxetine (Paxil), ciprofloxacin (Cipro), mefloquine, and fluconazole. Fluorine-substituted ethers are volatile anesthetics, including the commercial products methoxyflurane, enflurane, isoflurane, sevoflurane and desflurane. Fluorocarbon anesthetics reduce the hazard of flammability with diethyl ether and cyclopropane. Perfluorinated alkanes are used as blood substitutes.

Inhaler propellant

Fluorocarbons are also used as a propellant for metered-dose inhalers used to administer some asthma medications. The current generation of propellant consists of hydrofluoroalkanes (HFA), which has replaced CFC-propellant-based inhalers. CFC inhalers were banned as of 2008 as part of the Montreal Protocol because of environmental concerns with the ozone layer. HFA propellant inhalers like FloVent and ProAir ( Salbutamol ) have no generic versions available as of October 2014.

Fluorosurfactants

Fluorosurfactants, which have a polyfluorinated "tail" and a hydrophilic "head", serve as surfactants because they concentrate at the liquid-air interface due to their lipophobicity. Fluorosurfactants have low surface energies and dramatically lower surface tension. The fluorosurfactants perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (PFOA) are two of the most studied because of their ubiquity, toxicity, and long residence times in humans and wildlife.

Solvents

Fluorinated compounds often display distinct solubility properties. Dichlorodifluoromethane and chlorodifluoromethane were widely used refrigerants. CFCs have potent ozone depletion potential due to the homolytic cleavage of the carbon-chlorine bonds; their use is largely prohibited by the Montreal Protocol. Hydrofluorocarbons (HFCs), such as tetrafluoroethane, serve as CFC replacements because they do not catalyze ozone depletion. Oxygen exhibits a high solubility in perfluorocarbon compounds, reflecting again on their lipophilicity. Perfluorodecalin has been demonstrated as a blood substitutes, transporting oxygen to the lungs. 

The solvent 1,1,1,2-tetrafluoroethane has been used for extraction of natural products such as taxol, evening primrose oil, and vanillin. 2,2,2-trifluoroethanol is an oxidation-resistant polar solvent.

Organofluorine reagents

The development of organofluorine chemistry has contributed many reagents of value beyond organofluorine chemistry. Triflic acid (CF₃SO₃H) and trifluoroacetic acid (CF₃CO₂H) are useful throughout organic synthesis. Their strong acidity is attributed to the electronegativity of the trifluoromethyl group that stabilizes the negative charge. The triflate-group (the conjugate base of the triflic acid) is a good leaving group in substitution reactions.

Fluorous phases

Of topical interest in the area of "Green Chemistry," highly fluorinated substituents, e.g. perfluorohexyl (C₆F₁₃) confer distinctive solubility properties to molecules, which facilitates purification of products in organic synthesis. This area, described as "fluorous chemistry," exploits the concept of like-dissolves-like in the sense that fluorine-rich compounds dissolve preferentially in fluorine-rich solvents. Because of the relative inertness of the C-F bond, such fluorous phases are compatible with even harsh reagents. This theme has spawned techniques of "fluorous tagging and fluorous protection. Illustrative of fluorous technology is the use of fluoroalkyl-substituted tin hydrides for reductions, the products being easily separated from the spent tin reagent by extraction using fluorinated solvents.

 

Hydrophobic fluorinated ionic liquids, such as organic salts of bistriflimide or hexafluorophosphate, can form phases that are insoluble in both water and organic solvents, producing multiphasic liquids.

Organofluorine ligands in transition metal chemistry

Organofluorine ligands have long been featured in organometallic and coordination chemistry. One advantage to F-containing ligands is the convenience of ¹⁹F NMR spectroscopy for monitoring reactions. The organofluorine compounds can serve as a "sigma-donor ligand," as illustrated by the titanium(III) derivative [(C₅Me₅)₂Ti(FC₆H₅)]BPh₄. Most often, however, fluorocarbon substituents are used to enhance the Lewis acidity of metal centers. A premier example is "Eufod," a coordination complex of europium(III) that features a perfluoroheptyl modified acetylacetonate ligand. This and related species are useful in organic synthesis and as "shift reagents" in NMR spectroscopy. 

Structure of [(C₅Me₅)₂Ti(FC₆H₅)]⁺, a coordination complex of an organofluorine ligand.

 

In an area where coordination chemistry and materials science overlap, the fluorination of organic ligands is used to tune the properties of component molecules. For example, the degree and regiochemistry of fluorination of metalated 2-phenylpyridine ligands in platinum(II) complexes significantly modifies the emission properties of the complexes.

The coordination chemistry of organofluorine ligands also embraces fluorous technologies. For example, triphenylphosphine has been modified by attachment of perfluoroalkyl substituents that confer solubility in perfluorohexane as well as supercritical carbon dioxide. As a specific example, [(C₈F₁₇C₃H₆-4-C₆H₄)₃P.

C-F bond activation

An active area of organometallic chemistry encompasses the scission of C-F bonds by transition metal-based reagents. Both stoichiometric and catalytic reactions have been developed and are of interest from the perspectives of organic synthesis and remediation of xenochemicals. C-F bond activation has been classified as follows "(i) oxidative addition of fluorocarbon, (ii) M–C bond formation with HF elimination, (iii) M–C bond formation with fluorosilane elimination, (iv) hydrodefluorination of fluorocarbon with M–F bond formation, (v) nucleophilic attack on fluorocarbon, and (vi) defluorination of fluorocarbon". An illustrative metal-mediated C-F activation reaction is the defluorination of fluorohexane by a zirconium dihydride, an analogue of Schwartz's reagent:

(C₅Me₅)₂ZrH₂ + 1-FC₆H₁₃ → (C₅Me₅)₂ZrH(F) + C₆H₁₄

Fluorocarbon anions in Ziegler-Natta catalysis

Fluorine-containing compounds are often featured in noncoordinating or weakly coordinating anions. Both tetrakis(pentafluorophenyl)borate, B(C₆F₅)₄⁻, and the related tetrakis[3,5-bis(trifluoromethyl)phenyl]borate, are useful in Ziegler-Natta catalysis and related alkene polymerization methodologies. The fluorinated substituents render the anions weakly basic and enhance the solubility in weakly basic solvents, which are compatible with strong Lewis acids.

Materials science

Organofluorine compounds enjoy many niche applications in materials science. With a low coefficient of friction, fluid fluoropolymers are used as specialty lubricants. Fluorocarbon-based greases are used in demanding applications. Representative products include Fomblin and Krytox, manufactured by Solvay Solexis and DuPont, respectively. Certain firearm lubricants such as "Tetra Gun" contain fluorocarbons. Capitalizing on their nonflammability, fluorocarbons are used in fire fighting foam. Organofluorine compounds are components of liquid crystal displays. The polymeric analogue of triflic acid, nafion is a solid acid that is used as the membrane in most low temperature fuel cells. The bifunctional monomer 4,4'-difluorobenzophenone is a precursor to PEEK-class polymers.

Biosynthesis of organofluorine compounds

In contrast to the many naturally-occurring organic compounds containing the heavier halides, chloride, bromide, and iodide, only a handful of biologically synthesized carbon-fluorine bonds are known. The most common natural organofluorine species is fluoroacetate, a toxin found in a few species of plants. Others include fluorooleic acid, fluoroacetone, nucleocidin (4'-fuoro-5'-O-sulfamoyladenosine), fluorothreonine, and 2-fluorocitrate. Several of these species are probably biosynthesized from fluoroacetaldehyde. The enzyme fluorinase catalyzed the synthesis of 5'-fluoro-5-deoxyadenosine (see scheme to right).

History

Organofluorine chemistry began in the 1800s with the development of organic chemistry. The first organofluorine compounds were prepared using antimony trifluoride as the F⁻ source. The nonflammability and nontoxicity of the chlorofluorocarbons CCl₃F and CCl₂F₂ attracted industrial attention in the 1920s. In the 1930s, scientists at duPont discovered polytetrafluoroethylene. Subsequent major developments, especially in the US, benefited from expertise gained in the production of uranium hexafluoride. Starting in the late 1940s, a series of electrophilic fluorinating methodologies were introduced, beginning with CoF₃. About this time, electrochemical fluorination ("electrofluorination") was announced, having been developed in the 1930s with the goal of generating highly stable perfluorinated materials compatible with uranium hexafluoride. These new methodologies allowed the synthesis of C-F bonds without using elemental fluorine and without relying on metathetical methods. In 1957, the anticancer activity of 5-fluorouracil was described. This report provided one of the first examples of rational design of drugs. This discovery sparked a surge of interest in fluorinated pharmaceuticals and agrichemicals. The discovery of the noble gas compounds, e.g. XeF₄, provided a host of new reagents starting in the early 1960s. In the 1970s, fluorodeoxyglucose was established as a useful reagent in ¹⁸F positron emission tomography. In Nobel Prize-winning work, CFC's were shown to contribute to the depletion of atmospheric ozone. This discovery alerted the world to the negative consequences of organofluorine compounds and motivated the development of new routes to organofluorine compounds. In 2002, the first C-F bond-forming enzyme, fluorinase, was reported.

Environmental and health concerns

Only a few organofluorine compounds are acutely bioactive and highly toxic, such as fluoroacetate and perfluoroisobutene. 

Some organofluorine compounds pose significant risks and dangers to health and the environment. CFCs and HCFCs (hydrochlorofluorocarbon) deplete the ozone layer and are potent greenhouse gases. HFCs are potent greenhouse gases and are facing calls for stricter international regulation and phase out schedules as a fast-acting greenhouse emission abatement measure, as are perfluorocarbons (PFCs), and sulphur hexafluoride (SF₆). 

Because of the compound's effect on climate, the G-20 major economies agreed in 2013 to support initiatives to phase out use of HCFCs. They affirmed the roles of the Montreal Protocol and the United Nations Framework Convention on Climate Change in global HCFC accounting and reduction. The U.S. and China at the same time announced a bilateral agreement to similar effect.

Persistence and bioaccumulation

Because of the strength of the carbon–fluorine bond, many synthetic fluorocarbons and fluorcarbon-based compounds are persistent in the environment. Fluorosurfactants, such as PFOS and PFOA, are persistent global contaminants. Fluorocarbon based CFCs and tetrafluoromethane have been reported in igneous and metamorphic rock. PFOS is a persistent organic pollutant and may be harming the health of wildlife; the potential health effects of PFOA to humans are under investigation by the C8 Science Panel.