Fluoride
A fluoride is the ionic form of fluorine. As a halogen, fluorine forms a monovalent ion (-1 charge). Fluoride forms a binary compound with another element or radical. Examples of common fluoride compounds include hydrofluoric acid (HF), and sodium fluoride (NaF).
Fluoride compounds are used in a wide range of applications.
Where used in very low concentrations (on the order of parts per million), fluorides are used in human health applications; specifically, fluorides such as sodium fluoride (NaF), sodium fluorophosphate (SMFP), tin (II) fluoride (SnFF2), and amine fluoride are common ingredients in toothpaste. Many dentists also give their patients semiannual fluoride treatments.
Similarly, many North American municipalities also fluoridate their water supplies, citing effectiveness in reducing tooth decay, safety of fluoridation, and the low cost to do so. The American Dental Association (ADA), World Health Organization (WHO), and some other health organizations recommend fluoridation of municipal water supplies to a level between 0.7 and 1.2 ppm.
When used in very high concentrations (on the order of 10% by volume or higher) , sodium fluoride may be found in rat poisons, insecticides, and wood preservatives.
Hydrofluoric acid is used in etching glass and industrial applications.
In high concentrations, as with almost all substances, fluoride compounds are toxic. 5 grams of full strength sodium fluoride will kill most adult humans; a lethal dose is approximately 75mg per kilogram body mass. When ingested directly, fluoride compounds are readily absorbed by the intestines; over time, the compound is excreted through the urine, and the half life for concentration of fluorine compounds is on an order of hours. Implied is that fluoride is taken out of circulation by the body and trace amounts bound in bone. Urine tests are a good indication of high exposure to fluoride compounds in the recent past.
In 1973 Jason Burton, a boy in Melbourne, Australia died after swallowing six fluoride tablets; four were recovered after a stomach pumping while the other two had already been absorbed into the body. The hospital staff had assumed it would take well over 100 tablets to be fatal, making the Burton case an object lesson in high concentration fluoride intake.[1]
High concentration salts of fluorine are toxic if swallowed or inhaled. In October 1948 Donora, Pennsylvania had an incident of fluoride poisoning as a result of a temperature inversion in a valley of zinc and steel works; 20 died.[1]
Contact by many fluoride compounds (in high concentrations) with skin or eyes is dangerous. In case of accidental swallowing, give milk, calcium carbonate or milk of magnesia to slow absorption. Eye or skin contact should be treated by removing any contaminated clothing and flushing with water.
Fluoride is best known for its use in small quantities to help reduce Dental caries (cavity) frequency in teeth. A debate continues about whether fluoride ions (F-) are a trace element beneficial to humans for other reasons. Health Canada's current stance is that fluorides are beneficial to teeth, but that other physiological benefits are unproven [1]. The National Academy of Sciences generally agrees with Health Canada's opinion; from a November 1998 letter: "First, let us reassure you with regard to one concern. Nowhere in the report is it stated that fluoride is an essential nutrient. If any speaker or panel member at the September 23rd workshop referred to fluoride as such, they misspoke. As was stated in Recommended Dietary Allowances 10th Edition, which we published in 1989: 'These contradictory results do not justify a classification of fluoride as an essential element, according to accepted standards. Nonetheless, because of its valuable effects on dental health, fluoride is a beneficial element for humans.'"[1]
Fluoride compounds, usually calcium fluoride, are naturally found in low concentration in drinking water and some foods, like tea. Fluoride ions replace hydroxide ions in calcium hydroxyapatite, Ca5[(PO4)3OH], in teeth, forming calcium fluoroapatite, Ca5[(PO4)3F], which is more chemically stable and dissolves at a pH of 4.5, compared to 5.5 pH for calcium hydroxyapatite. This is generally believed to lead to fewer cavities, since stronger acids are needed to attack the tooth enamel.
The only generally accepted adverse effect of low concentration fluoridation at this time is fluorosis. It is a condition caused by 'excessive' intake of fluorine compounds over an extended period of time, and can cause yellowing of teeth, or brittling of bones and teeth. The definition of 'excessive' in the context of fluorosis falls on the order of parts per million and is generally accepted to mean significantly higher than the 0.7 to 1.2 ppm amounts recommended for fluoridated water. One brand of popular bottled water was tested by a critic of the practice of fluoridation to have 3.6 ppm of fluoride; this is a good example.[1]
Fluoride was also tried as a therapy for osteoporosis (on an order of 75mg/day, a relatively high dose). Bone density did increase in the studies. However, some in the study had to reduce their fluoride dosage to avoid side effects. Additionally, overall bone strength was compromised. The treatment created a coarse crystalline matrix with reduced tensile strength, which made the bone (especially the cranked hip joint) susceptible to failure in bending.[1]
Uses
Fluorides and human health
High concentrations
Low concentrations
Research sites