Fluoride

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Contents

Fluoride
F- crop.svg
Fluoride ion.svg
Names
IUPAC name
Fluoride [1]
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
14905
KEGG
MeSH Fluoride
PubChem CID
UNII
  • InChI=1S/FH/h1H/p-1 Yes check.svgY
    Key: KRHYYFGTRYWZRS-UHFFFAOYSA-M Yes check.svgY
  • [F-]
Properties
F−
Molar mass 18.998403163 g·mol−1
Conjugate acid Hydrogen fluoride
Thermochemistry
Std molar
entropy (S⦵298)
145.58 J/mol K (gaseous) [2]
−333 kJ mol−1
Related compounds
Other anions
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Fluoride ( /ˈflʊəraɪd,ˈflɔːr-/ ) [3] is an inorganic, monatomic anion of fluorine, with the chemical formula F−
 (also written [F]−
), whose salts are typically white or colorless. Fluoride salts typically have distinctive bitter tastes, and are odorless. Its salts and minerals are important chemical reagents and industrial chemicals, mainly used in the production of hydrogen fluoride for fluorocarbons. Fluoride is classified as a weak base since it only partially associates in solution, but concentrated fluoride is corrosive and can attack the skin.

Fluoride is the simplest fluorine anion. In terms of charge and size, the fluoride ion resembles the hydroxide ion. Fluoride ions occur on Earth in several minerals, particularly fluorite, but are present only in trace quantities in bodies of water in nature.

Nomenclature

Fluorides include compounds that contain ionic fluoride and those in which fluoride does not dissociate. The nomenclature does not distinguish these situations. For example, sulfur hexafluoride and carbon tetrafluoride are not sources of fluoride ions under ordinary conditions.

The systematic name fluoride, the valid IUPAC name, is determined according to the additive nomenclature. However, the name fluoride is also used in compositional IUPAC nomenclature which does not take the nature of bonding involved into account. Fluoride is also used non-systematically, to describe compounds which release fluoride upon dissolving. Hydrogen fluoride is itself an example of a non-systematic name of this nature. However, it is also a trivial name, and the preferred IUPAC name for fluorane.[ citation needed ]

Occurrence

Fluorite crystals Calcite sur fluorine (USA) 1.JPG
Fluorite crystals

Fluorine is estimated to be the 13th-most abundant element in Earth's crust and is widely dispersed in nature, entirely in the form of fluorides. The vast majority is held in mineral deposits, the most commercially important of which is fluorite (CaF2). [4] Natural weathering of some kinds of rocks, [5] [6] as well as human activities, releases fluorides into the biosphere through what is sometimes called the fluorine cycle.

In water

Fluoride is naturally present in groundwater, fresh and saltwater sources, as well as in rainwater, particularly in urban areas. [7] Seawater fluoride levels are usually in the range of 0.86 to 1.4 mg/L, and average 1.1 mg/L [8] (milligrams per litre). For comparison, chloride concentration in seawater is about 19 g/L. The low concentration of fluoride reflects the insolubility of the alkaline earth fluorides, e.g., CaF2.

Concentrations in fresh water vary more significantly. Surface water such as rivers or lakes generally contains between 0.01 and 0.3 mg/L. [9] Groundwater (well water) concentrations vary even more, depending on the presence of local fluoride-containing minerals. For example, natural levels of under 0.05 mg/L have been detected in parts of Canada but up to 8 mg/L in parts of China; in general levels rarely exceed 10 mg/litre [10]

Fluoride can be present in rain, with its concentration increasing significantly upon exposure to volcanic activity [14] or atmospheric pollution derived from burning fossil fuels or other sorts of industry, [15] [16] particularly aluminium smelters. [17]

In plants

All vegetation contains some fluoride, which is absorbed from soil and water. [10] Some plants concentrate fluoride from their environment more than others. All tea leaves contain fluoride; however, mature leaves contain as much as 10 to 20 times the fluoride levels of young leaves from the same plant. [18] [19] [20]

Chemical properties

Basicity

Fluoride can act as a base. It can combine with a proton (  H+):

F− + H+ → HF

This neutralization reaction forms hydrogen fluoride (HF), the conjugate acid of fluoride.

In aqueous solution, fluoride has a pKb value of 10.8. It is therefore a weak base, and tends to remain as the fluoride ion rather than generating a substantial amount of hydrogen fluoride. That is, the following equilibrium favours the left-hand side in water:

F− + H2O ⇌ HF + HO−

However, upon prolonged contact with moisture, soluble fluoride salts will decompose to their respective hydroxides or oxides, as the hydrogen fluoride escapes. Fluoride is distinct in this regard among the halides. The identity of the solvent can have a dramatic effect on the equilibrium shifting it to the right-hand side, greatly increasing the rate of decomposition.

Structure of fluoride salts

Salts containing fluoride are numerous and adopt myriad structures. Typically the fluoride anion is surrounded by four or six cations, as is typical for other halides. Sodium fluoride and sodium chloride adopt the same structure. For compounds containing more than one fluoride per cation, the structures often deviate from those of the chlorides, as illustrated by the main fluoride mineral fluorite (CaF2) where the Ca2+ ions are surrounded by eight F− centers. In CaCl2, each Ca2+ ion is surrounded by six Cl− centers. The difluorides of the transition metals often adopt the rutile structure whereas the dichlorides have cadmium chloride structures.

Inorganic chemistry

Upon treatment with a standard acid, fluoride salts convert to hydrogen fluoride and metal salts. With strong acids, it can be doubly protonated to give H
2F+
. Oxidation of fluoride gives fluorine. Solutions of inorganic fluorides in water contain F− and bifluoride HF−
2. [21] Few inorganic fluorides are soluble in water without undergoing significant hydrolysis. In terms of its reactivity, fluoride differs significantly from chloride and other halides, and is more strongly solvated in protic solvents due to its smaller radius/charge ratio. Its closest chemical relative is hydroxide, since both have similar geometries.

Naked fluoride

Most fluoride salts dissolve to give the bifluoride (HF−
2) anion. Sources of true F− anions are rare because the highly basic fluoride anion abstracts protons from many, even adventitious, sources. Relative unsolvated fluoride, which does exist in aprotic solvents, is called "naked". Naked fluoride is a strong Lewis base, [22] and a powerful nucleophile. Some quaternary ammonium salts of naked fluoride include tetramethylammonium fluoride and tetrabutylammonium fluoride. [23] Cobaltocenium fluoride is another example. [24] However, they all lack structural characterization in aprotic solvents. Because of their high basicity, many so-called naked fluoride sources are in fact bifluoride salts. In late 2016 imidazolium fluoride was synthesized that is the closest approximation of a thermodynamically stable and structurally characterized example of a "naked" fluoride source in an aprotic solvent (acetonitrile). [25] The sterically demanding imidazolium cation stabilizes the discrete anions and protects them from polymerization. [26] [27]

Biochemistry

At physiological pHs, hydrogen fluoride is usually fully ionised to fluoride. In biochemistry, fluoride and hydrogen fluoride are equivalent. Fluorine, in the form of fluoride, is considered to be a micronutrient for human health, necessary to prevent dental cavities, and to promote healthy bone growth. [28] The tea plant ( Camellia sinensis L.) is a known accumulator of fluorine compounds, released upon forming infusions such as the common beverage. The fluorine compounds decompose into products including fluoride ions. Fluoride is the most bioavailable form of fluorine, and as such, tea is potentially a vehicle for fluoride dosing. [29] Approximately, 50% of absorbed fluoride is excreted renally with a twenty-four-hour period. The remainder can be retained in the oral cavity, and lower digestive tract. Fasting dramatically increases the rate of fluoride absorption to near 100%, from a 60% to 80% when taken with food. [29] Per a 2013 study, it was found that consumption of one litre of tea a day, can potentially supply the daily recommended intake of 4 mg per day. Some lower quality brands can supply up to a 120% of this amount. Fasting can increase this to 150%. The study indicates that tea drinking communities are at an increased risk of dental and skeletal fluorosis, in the case where water fluoridation is in effect. [29] Fluoride ion in low doses in the mouth reduces tooth decay. [30] For this reason, it is used in toothpaste and water fluoridation. At much higher doses and frequent exposure, fluoride causes health complications and can be toxic.

Applications

Fluoride salts and hydrofluoric acid are the main fluorides of industrial value.

Organofluorine chemistry

Organofluorine compounds are pervasive. Many drugs, many polymers, refrigerants, and many inorganic compounds are made from fluoride-containing reagents. Often fluorides are converted to hydrogen fluoride, which is a major reagent and precursor to reagents. Hydrofluoric acid and its anhydrous form, hydrogen fluoride, are particularly important. [4]

Production of metals and their compounds

The main uses of fluoride, in terms of volume, are in the production of cryolite, Na3AlF6. It is used in aluminium smelting. Formerly, it was mined, but now it is derived from hydrogen fluoride. Fluorite is used on a large scale to separate slag in steel-making. Mined fluorite (CaF2) is a commodity chemical used in steel-making. Uranium hexafluoride is employed in the purification of uranium isotopes.

Cavity prevention

Fluoride is sold in tablets for cavity prevention. Sodium fluoride tablets.jpg
Fluoride is sold in tablets for cavity prevention.

Fluoride-containing compounds, such as sodium fluoride or sodium monofluorophosphate are used in topical and systemic fluoride therapy for preventing tooth decay. They are used for water fluoridation and in many products associated with oral hygiene. [31] Originally, sodium fluoride was used to fluoridate water; hexafluorosilicic acid (H2SiF6) and its salt sodium hexafluorosilicate (Na2SiF6) are more commonly used additives, especially in the United States. The fluoridation of water is known to prevent tooth decay [32] [33] and is considered by the U.S. Centers for Disease Control and Prevention to be "one of 10 great public health achievements of the 20th century". [34] [35] In some countries where large, centralized water systems are uncommon, fluoride is delivered to the populace by fluoridating table salt. For the method of action for cavity prevention, see Fluoride therapy. Fluoridation of water has its critics (see Water fluoridation controversy). [36] Fluoridated toothpaste is in common use. Meta-analysis show the efficacy of 500 ppm fluoride in toothpastes. [37] [38] However, no beneficial effect can be detected when more than one fluoride source is used for daily oral care. [39] [ need quotation to verify ]

Laboratory reagent

Fluoride salts are commonly used in biological assay processing to inhibit the activity of phosphatases, such as serine/threonine phosphatases. [40] Fluoride mimics the nucleophilic hydroxide ion in these enzymes' active sites. [41] Beryllium fluoride and aluminium fluoride are also used as phosphatase inhibitors, since these compounds are structural mimics of the phosphate group and can act as analogues of the transition state of the reaction. [42] [43]

Dietary recommendations

The U.S. Institute of Medicine (IOM) updated Estimated Average Requirements (EARs) and Recommended Dietary Allowances (RDAs) for some minerals in 1997. Where there was not sufficient information to establish EARs and RDAs, an estimate designated Adequate Intake (AI) was used instead. AIs are typically matched to actual average consumption, with the assumption that there appears to be a need, and that need is met by what people consume. The current AI for women 19 years and older is 3.0 mg/day (includes pregnancy and lactation). The AI for men is 4.0 mg/day. The AI for children ages 1–18 increases from 0.7 to 3.0 mg/day. The major known risk of fluoride deficiency appears to be an increased risk of bacteria-caused tooth cavities. As for safety, the IOM sets tolerable upper intake levels (ULs) for vitamins and minerals when evidence is sufficient. In the case of fluoride the UL is 10 mg/day. Collectively the EARs, RDAs, AIs and ULs are referred to as Dietary Reference Intakes (DRIs). [44]

The European Food Safety Authority (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR. AI and UL are defined the same as in the United States. For women ages 18 and older the AI is set at 2.9 mg/day (including pregnancy and lactation). For men, the value is 3.4 mg/day. For children ages 1–17 years, the AIs increase with age from 0.6 to 3.2 mg/day. These AIs are comparable to the U.S. AIs. [45] The EFSA reviewed safety evidence and set an adult UL at 7.0 mg/day (lower for children). [46]

For U.S. food and dietary supplement labeling purposes, the amount of a vitamin or mineral in a serving is expressed as a percent of Daily Value (%DV). Although there is information to set Adequate Intake, fluoride does not have a Daily Value and is not required to be shown on food labels. [47]

Estimated daily intake

Daily intakes of fluoride can vary significantly according to the various sources of exposure. Values ranging from 0.46 to 3.6–5.4 mg/day have been reported in several studies (IPCS, 1984). [28] In areas where water is fluoridated this can be expected to be a significant source of fluoride, however fluoride is also naturally present in virtually all foods and beverages at a wide range of concentrations. [48] The maximum safe daily consumption of fluoride is 10 mg/day for an adult (U.S.) or 7 mg/day (European Union). [44] [46]

The upper limit of fluoride intake from all sources (fluoridated water, food, beverages, fluoride dental products and dietary fluoride supplements) is set at 0.10 mg/kg/day for infants, toddlers, and children through to 8 years old. For older children and adults, who are no longer at risk for dental fluorosis, the upper limit of fluoride is set at 10 mg/day regardless of weight. [49]

Examples of fluoride content
Food/DrinkFluoride
(mg per 1000g/ppm)		PortionFluoride
(mg per portion)
Black tea (brewed)3.731 cup, 240 g (8 fl oz)0.884
Raisins, seedless2.34small box, 43 g (1.5 oz)0.101
Table wine1.53Bottle, 750 mL (26 imp fl oz)1.150
Municipal tap-water,
(Fluoridated)		0.81Recommended daily intake,
3 litres (0.79 US gal)		2.433
Baked potatoes, Russet0.45Medium potato, 140 g (0.31 lb)0.078
Lamb0.32Chop, 170 g (6.0 oz)0.054
Carrots0.031 large carrot, 72 g (2.5 oz)0.002
Source: Data taken from United States Department of Agriculture, National Nutrient Database Archived 2014-03-01 at the Wayback Machine [50]

Safety

Ingestion

According to the U.S. Department of Agriculture, the Dietary Reference Intakes, which is the "highest level of daily nutrient intake that is likely to pose no risk of adverse health effects" specify 10 mg/day for most people, corresponding to 10 L of fluoridated water with no risk. For young children the values are smaller, ranging from 0.7 mg/d to 2.2 mg/d for infants. [51] Water and food sources of fluoride include community water fluoridation, seafood, tea, and gelatin. [52]

Soluble fluoride salts, of which sodium fluoride is the most common, are toxic, and have resulted in both accidental and self-inflicted deaths from acute poisoning. [4] The lethal dose for most adult humans is estimated at 5 to 10 g (which is equivalent to 32 to 64 mg elemental fluoride per kg body weight). [53] [54] [55] A case of a fatal poisoning of an adult with 4 grams of sodium fluoride is documented, [56] and a dose of 120 g sodium fluoride has been survived. [57] For sodium fluorosilicate (Na2SiF6), the median lethal dose (LD50) orally in rats is 125 mg/kg, corresponding to 12.5 g for a 100 kg adult. [58]

Treatment may involve oral administration of dilute calcium hydroxide or calcium chloride to prevent further absorption, and injection of calcium gluconate to increase the calcium levels in the blood. [56] Hydrogen fluoride is more dangerous than salts such as NaF because it is corrosive and volatile, and can result in fatal exposure through inhalation or upon contact with the skin; calcium gluconate gel is the usual antidote. [59]

In the higher doses used to treat osteoporosis, sodium fluoride can cause pain in the legs and incomplete stress fractures when the doses are too high; it also irritates the stomach, sometimes so severely as to cause ulcers. Slow-release and enteric-coated versions of sodium fluoride do not have gastric side effects in any significant way, and have milder and less frequent complications in the bones. [60] In the lower doses used for water fluoridation, the only clear adverse effect is dental fluorosis, which can alter the appearance of children's teeth during tooth development; this is mostly mild and is unlikely to represent any real effect on aesthetic appearance or on public health. [61] Fluoride was known to enhance bone mineral density at the lumbar spine, but it was not effective for vertebral fractures and provoked more nonvertebral fractures. [62] In areas that have naturally occurring high levels of fluoride in groundwater which is used for drinking water, both dental and skeletal fluorosis can be prevalent and severe. [63]

In 2024, a 300-page report by the National Institutes of Health linked flouridation of drinking water in the United States, in areas where levels are slightly more than twice the recommended limit, to lower IQ in developing children. [64] [65] [66] [67]

Hazard maps for fluoride in groundwater

Around one-third of the human population drinks water from groundwater resources. Of this, about 10%, approximately 300 million people, obtain water from groundwater resources that are heavily contaminated with arsenic or fluoride. [68] These trace elements derive mainly from minerals. [69] Maps locating potential problematic wells are available. [70]

Topical

Concentrated fluoride solutions are corrosive. [71] Gloves made of nitrile rubber are worn when handling fluoride compounds. The hazards of solutions of fluoride salts depend on the concentration. In the presence of strong acids, fluoride salts release hydrogen fluoride, which is corrosive, especially toward glass. [4]

Other derivatives

Organic and inorganic anions are produced from fluoride, including:

See also

Related Research Articles

<span class="mw-page-title-main">Halogen</span> Group of chemical elements

The halogens are a group in the periodic table consisting of six chemically related elements: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and the radioactive elements astatine (At) and tennessine (Ts), though some authors would exclude tennessine as its chemistry is unknown and is theoretically expected to be more like that of gallium. In the modern IUPAC nomenclature, this group is known as group 17.

<span class="mw-page-title-main">Potassium</span> Chemical element with atomic number 19 (K)

Potassium is a chemical element; it has symbol K and atomic number 19. It is a silvery white metal that is soft enough to easily cut with a knife. Potassium metal reacts rapidly with atmospheric oxygen to form flaky white potassium peroxide in only seconds of exposure. It was first isolated from potash, the ashes of plants, from which its name derives. In the periodic table, potassium is one of the alkali metals, all of which have a single valence electron in the outer electron shell, which is easily removed to create an ion with a positive charge. In nature, potassium occurs only in ionic salts. Elemental potassium reacts vigorously with water, generating sufficient heat to ignite hydrogen emitted in the reaction, and burning with a lilac-colored flame. It is found dissolved in seawater, and occurs in many minerals such as orthoclase, a common constituent of granites and other igneous rocks.

<span class="mw-page-title-main">Distilled water</span> Water that has had many of its impurities removed through distillation

Distilled water is water that has been boiled into vapor and condensed back into liquid in a separate container. Impurities in the original water that do not boil below or near the boiling point of water remain in the original container. Thus, distilled water is a type of purified water.

<span class="mw-page-title-main">Skeletal fluorosis</span> Medical condition

Skeletal fluorosis is a bone disease caused by excessive accumulation of fluoride leading to weakened bones. In advanced cases, skeletal fluorosis causes painful damage to bones and joints.

<span class="mw-page-title-main">Water fluoridation</span> Addition of fluoride to a water supply to reduce tooth decay

Water fluoridation is the addition of fluoride to a public water supply to reduce tooth decay. Fluoridated water contains fluoride at a level that is effective for preventing cavities; this can occur naturally or by adding fluoride. Fluoridated water operates on tooth surfaces: in the mouth, it creates low levels of fluoride in saliva, which reduces the rate at which tooth enamel demineralizes and increases the rate at which it remineralizes in the early stages of cavities. Typically a fluoridated compound is added to drinking water, a process that in the U.S. costs an average of about $1.32 per person-year. Defluoridation is needed when the naturally occurring fluoride level exceeds recommended limits. In 2011, the World Health Organization suggested a level of fluoride from 0.5 to 1.5 mg/L, depending on climate, local environment, and other sources of fluoride. In 2024, the Department of Health and Human Services' National Toxicology Program found that water fluoridation levels above 1.5 mg/L are associated with lower IQ in children. In 2024, U.S. court rulings have raised concerns about the potential health risks of water fluoridation, including findings by the EPA and new risk assessments that suggest the benefits may be waning. Bottled water typically has unknown fluoride levels.

Fluoride toxicity is a condition in which there are elevated levels of the fluoride ion in the body. Although fluoride is safe for dental health at low concentrations, sustained consumption of large amounts of soluble fluoride salts is dangerous. Referring to a common salt of fluoride, sodium fluoride (NaF), the lethal dose for most adult humans is estimated at 5 to 10 g. Ingestion of fluoride can produce gastrointestinal discomfort at doses at least 15 to 20 times lower than lethal doses. Although it is helpful topically for dental health in low dosage, chronic ingestion of fluoride in large amounts interferes with bone formation. In this way, the most widespread examples of fluoride poisoning arise from consumption of ground water that is abnormally fluoride-rich.

<span class="mw-page-title-main">Sodium fluoride</span> Ionic compound (NaF)

Sodium fluoride (NaF) is an inorganic compound with the formula NaF. It is a colorless or white solid that is readily soluble in water. It is used in trace amounts in the fluoridation of drinking water to prevent tooth decay, and in toothpastes and topical pharmaceuticals for the same purpose. In 2021, it was the 291st most commonly prescribed medication in the United States, with more than 600,000 prescriptions. It is also used in metallurgy and in medical imaging.

<span class="mw-page-title-main">Fluorine deficiency</span> Medical condition

Fluoride or fluorine deficiency is a disorder which may cause increased dental caries and possibly osteoporosis, due to a lack of fluoride in diet. Common dietary sources of fluoride include tea, grape juice, wine, raisins, some seafood, coffee, and tap water that has been fluoridated. The extent to which the condition truly exists, and its relationship to fluoride poisoning has given rise to some controversy. Fluorine is not considered to be an essential nutrient, but the importance of fluorides for preventing tooth decay is well-recognized, despite the effect is predominantly topical. Prior to 1981, the effect of fluorides was thought to be largely systemic and preeruptive, requiring ingestion. Fluoride is considered essential in the development and maintenance of teeth by the American Dental Hygienists' Association. Fluoride incorporates into the teeth to form and harden teeth enamels. This makes the teeth more acid resistant, as well as more resistant to cavity-forming bacteria. Caries-inhibiting effects of fluoride were first noticed 1902, when fluoride in high concentrations was found to stain teeth and prevent tooth decay.

The water fluoridation controversy arises from political, ethical, economic, and health considerations regarding the fluoridation of public water supplies. For deprived groups in both maturing and matured countries, international and national agencies and dental associations across the world support the safety and effectiveness of water fluoridation. Proponents of water fluoridation see it as a question of public health policy and equate the issue to vaccination and food fortification, citing significant benefits to dental health and minimal risks. In contrast, opponents of water fluoridation view it as an infringement of individual rights, if not an outright violation of medical ethics, on the basis that individuals have no choice in the water that they drink, unless they drink more expensive bottled water. A small minority of scientists have challenged the medical consensus, variously claiming that water fluoridation has no or little cariostatic benefits, may cause serious health problems, is not effective enough to justify the costs, and is pharmacologically obsolete.

<span class="mw-page-title-main">Fluoride therapy</span> Medical use of fluoride

Fluoride therapy is the use of fluoride for medical purposes. Fluoride supplements are recommended to prevent tooth decay in children older than six months in areas where the drinking water is low in fluoride. It is typically used as a liquid, pill, or paste by mouth. Fluoride has also been used to treat a number of bone diseases.

<span class="mw-page-title-main">Dental fluorosis</span> Tooth enamel discoloration due to excessive fluoride ingestion

Dental fluorosis is a common disorder, characterized by hypomineralization of tooth enamel caused by ingestion of excessive fluoride during enamel formation.

<span class="mw-page-title-main">Sodium monofluorophosphate</span> Chemical compound

Sodium monofluorophosphate, commonly abbreviated SMFP, is an inorganic compound with the chemical formula Na2PO3F. Typical for a salt, SMFP is odourless, colourless, and water-soluble. This salt is an ingredient in some toothpastes.

<span class="mw-page-title-main">Hexafluorosilicic acid</span> Octahedric silicon compound

Hexafluorosilicic acid is an inorganic compound with the chemical formula H
2SiF
6. Aqueous solutions of hexafluorosilicic acid consist of salts of the cation and hexafluorosilicate anion. These salts and their aqueous solutions are colorless.

<span class="mw-page-title-main">Ammonium bifluoride</span> Chemical compound

Ammonium bifluoride is an inorganic compound with the formula [NH4][HF2] or [NH4]F·HF. It is produced from ammonia and hydrogen fluoride. This colourless salt is a glass-etchant and an intermediate in a once-contemplated route to hydrofluoric acid.

The bifluoride ion is an inorganic anion with the chemical formula [HF2]−. The anion is colorless. Salts of bifluoride are commonly encountered in the reactions of fluoride salts with hydrofluoric acid. The commercial production of fluorine involves electrolysis of bifluoride salts.

<span class="mw-page-title-main">Remineralisation of teeth</span>

Tooth remineralization is the natural repair process for non-cavitated tooth lesions, in which calcium, phosphate and sometimes fluoride ions are deposited into crystal voids in demineralised enamel. Remineralization can contribute towards restoring strength and function within tooth structure.

<span class="mw-page-title-main">Water fluoridation in the United States</span>

Water fluoridation in the United States is common amongst most states. As of May 2000, 42 of the 50 largest U.S. cities had water fluoridation. On January 25, 1945, Grand Rapids, Michigan, became the first community in the United States to fluoridate its drinking water for the intended purpose of helping to prevent tooth decay.

<span class="mw-page-title-main">Biological aspects of fluorine</span>

Fluorine may interact with biological systems in the form of fluorine-containing compounds. Though elemental fluorine (F2) is very rare in everyday life, fluorine-containing compounds such as fluorite occur naturally as minerals. Naturally occurring organofluorine compounds are extremely rare. Man-made fluoride compounds are common and are used in medicines, pesticides, and materials. Twenty percent of all commercialized pharmaceuticals contain fluorine, including Lipitor and Prozac. In many contexts, fluorine-containing compounds are harmless or even beneficial to living organisms; in others, they are toxic.

Defluoridation is the downward adjustment of the level of fluoride in drinking water. Worldwide, fluoride is one of the most abundant anions present in groundwater. Fluoride is more present in groundwater than surface water mainly due to the leaching of minerals. Groundwater accounts for 98 percent of the earth's potable water. An excess of fluoride in drinking water causes dental fluorosis and skeletal fluorosis. The World Health Organization has recommended a guideline value of 1.5 mg/L as the concentration above which dental fluorosis is likely. Fluorosis is endemic in more than 20 developed and developing nations.

Topical fluorides are fluoride-containing drugs indicated in prevention and treatment of dental caries, particularly in children's primary dentitions. The dental-protecting property of topical fluoride can be attributed to multiple mechanisms of action, including the promotion of remineralization of decalcified enamel, the inhibition of the cariogenic microbial metabolism in dental plaque and the increase of tooth resistance to acid dissolution. Topical fluoride is available in a variety of dose forms, for example, toothpaste, mouth rinses, varnish and silver diamine solution. These dosage forms possess different absorption mechanisms and consist of different active ingredients. Common active ingredients include sodium fluoride, stannous fluoride, silver diamine fluoride. These ingredients account for different pharmacokinetic profiles, thereby having varied dosing regimes and therapeutic effects. A minority of individuals may experience certain adverse effects, including dermatological irritation, hypersensitivity reactions, neurotoxicity and dental fluorosis. In severe cases, fluoride overdose may lead to acute toxicity. While topical fluoride is effective in preventing dental caries, it should be used with caution in specific situations to avoid undesired side effects.

References

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  2. ↑ Chase, M. W. (1998). "Fluorine anion". NIST. pp. 1–1951. Retrieved 4 July 2012.
  3. ↑ Wells, J.C. (2008). Longman pronunciation dictionary (3rd ed.). Harlow, England: Pearson Education Limited/Longman. p. 313. ISBN   9781405881180.. According to this source, /ˈfluːəraɪd/ is a possible pronunciation in British English.
  4. 1 2 3 4 Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, René; Cuer, Jean Pierre (2000). "Fluorine Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a11_307. ISBN   978-3527306732.
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  8. ↑ "Ambient Water Quality Criteria for Fluoride". Government of British Columbia. Archived from the original on 24 September 2015. Retrieved 8 October 2014.
  9. ↑ Liteplo, Dr R.; Gomes, R.; Howe, P.; Malcolm, Heath (2002). FLUORIDES - Environmental Health Criteria 227 : 1st draft. Geneva: World Health Organization. ISBN   978-9241572279.
  10. 1 2 Fawell, J.K.; et al. "Fluoride in Drinking-water Background document for development of WHO Guidelines for Drinking-water Quality" (PDF). World Health Organization. Retrieved 6 May 2016.
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