Physiologic Conditions Affect Toxicity of Ingested Industrial Fluoride
Richard Sauerheber, Ph.D.
Department of Chemistry, University of California, San Diego, La Jolla, CA 92037
Palomar Community College, San Marcos, CA 92069 E-mail: firstname.lastname@example.org
Abstract. The effects of calcium ion and pH over broad ranges on the free fluoride ion aqueous concentration were determined. Solubility calculations indicate that blood fluoride concentrations that occur in lethal poisonings would decrease calcium below normal physiological levels. Acute lethal poisoning and also many of the chronic ‘low’ level effects of fluoride are mediated by calcium binding by fluoride ion. At a pH typical of gastric juice, approximately 50% of fluoride is protonated as hydrofluoric acid HF with 50% remaining the free fluoride ion. Industrial fluoride ingested from treated water enters saliva at levels too low to affect caries topically, while blood levels during lifelong consumption can harm heart, bone and brain.
Introduction. Synthetic industrial fluoride compounds lack calcium and are listed toxic substances (Buck, 1, Gleason, 2, Blakiston, 3, Merck, 4). Calcium fluoride is found in nature and is not considered a toxic compound because of its comparatively high lethal oral acute dose in rodents (LD50 = 3,750 mg/kg). The fluoride compounds sodium fluoride and fluosilicic acid, added into municipal water for human ingestion purposes. These are synthesized artificially by industrial reaction and are classed as rodenticides, insecticides, and pediculicides with acute oral lethal doses in experimental animals comparable to arsenic and lead (Merck, 4) (NaF and H2SiF6 LD50 = 125 mg/kg).
Waters in the U.S. can contain natural calcium fluoride along with other calcium salts (U.S. CDC, 5). Although fluoride has been debated to be an agent that alters teeth by ingestion, natural fluoride is accompanied with calcium. Calcium, not fluoride, is the chief ingredient in normal teeth enamel hydroxyapatite. The principal effect of ingested fluoride on developing teeth is to alter the structure of enamel to cause fluorosis, a permanent mottling reported to afflict 5 million teens aged 12-15 in the U.S. as of 2004.
Methods. The concentration of the free fluoride ion was mathematically computed at which the solubility of calcium fluoride would be exceeded with calcium concentrations known to be physiologic. [F-] = (Ksp/[Ca2+])1/2 from the definition of the solubility product constant for insoluble salts where CaF2 Ca2+ + 2 F- and Ksp = [Ca2+][F-]2. Consideration was made for physiologic temperature by using the published value at 37oC.
A 0.9 ppm fluoride solution in distilled water was measured for the free fluoride ion concentration over a wide concentration range of added calcium ion from aliquots of a calcium biphosphate solution. In other experiments a 1.2 ppm fluoride concentration solution was measured for free fluoride ion level as a function of pH. Acidity was adjusted with dilute acetic acid. All readings were made at room temperature with a LaMotte (Maryland) fluoride ion specific electrode calibrated with a 1.00 ppm fluoride standard solution in distilled deionized water. The electrode was rinsed with the solution to be tested for each measurement. The instrument tested as a function of temperature reported accurate readings for known standard solutions to within 0.05 ppm fluoride over the range 15-30oC.
Acute Toxicity. The concentration of fluoride is here calculated that would cause calcium fluoride precipitates to first form from the known solubility product constant for calcium fluoride (Ksp = 8 x 10-11 at 37oC) and the known concentration of calcium ion in normal human blood (2.2 mM) (Davidsohn, 6). The Ksp varies slightly with temperature and may be computed at 37oC (310 Kelvin) from the relation ln(Ksp) = – ∆G/(RT) (Lide, 7) for calcium
fluoride with the free energy for the dissociation of calcium fluoride ∆G = 59 kJ/mol and Ksp = 3.4 x 10-11 at 25oC (298 K).
The computed fluoride level at which an aqueous solution containing physiologic calcium (3 mM) at physiologic temperature (37oC) is precipitated is 0.11 mM fluoride (2 ppm). Here the concentration of fluoride is: [F-] = (Ksp/[Ca2+])1/2 from the definition of the solubility product constant for insoluble salts where CaF2 → Ca2+ + 2F- and Ksp = [Ca2+][F-]2. The concentration of blood fluoride where the blood calcium level would be lowered to the lethal low level of about 1 mM is 0.2 mM fluoride (3.8 ppm).
The calculated calcium levels that would coexist in fluid with a given fluoride level from solubility considerations are compared with actual measurements of blood levels of calcium and fluoride ion in victims of fluoride poisoning (Gessner, 8) in Hooper Bay, Alaska, the largest fluoridation disaster accidental overfeed in the U.S. Calculations show that there is good agreement between theoretic fluoride levels that would lower blood calcium ion to levels below normal and the actual calcium and fluoride ion levels measured in the blood of victims poisoned with fluoridated municipal water. 296 poisoned residents were life-flighted for care. One victim who died of heart failure from fluoride had a measured fluoride level of 0.18 mM, while another victim that survived had blood fluoride at 0.48 mM that caused calcium levels to plummet to a typically-lethal level of 1 mM. These concentrations of fluoride from solubility considerations produce calcium ion lowering to levels reported to decrease beat rates in isolated mammalian heart cells (Wang, 9).
The fact that fluoride lethality occurs at concentrations known to compare with saturation and activity reduction for calcium ion brings forth an aspect of fluoride toxicity that is counter-intuitive. Ionized calcium levels in human plasma can vary in some cases from 1.5 mM in hypocalcemia to 4.5 mM in hypercalcemia (as in hyperparathyroidism or excessive Vitamin D intake) (Davidsohn, 6). The assimilation of ingested fluoride is drastically minimized by calcium ion in the gastrointestinal tract (CDC, 3). This would suggest that fluoride toxicity in the bloodstream would be lowered in individuals with higher blood calcium, but Ksp calculations indicate that higher blood calcium levels require less blood fluoride to reduce calcium activity or mobility. Figure 1A indicates the calcium fluoride precipitation isocurve over a broad concentration range for the two ions. The ions are precipitated when present together at any matched pair of concentrations indicated along the curve. Lower blood calcium levels require higher fluoride blood levels to begin precipitation. The effect is quite substantial, varying from 2.5 to 5 ppm fluoride lethal levels for subjects with physiologic 4.5 and 1.5 mM calcium, respectively. This may help explain the broad variability in reported blood and tissue fluoride levels in lethal acute fluoride toxicity from ingestion in humans.
Wang, Zhang and Wang (9) also found the heart cell beat rate in cultured cells in well-controlled experiments progressively slows with increasing fluoride levels in a concentration-dependent manner at levels comparable to the level calculated here at maximum solubility of calcium fluoride. Unlike skeletal muscle, cardiac muscle requires extracellular calcium ion from the bloodstream to couple electrical excitation of the cell membrane with contraction of cardiac muscle fibers. Each time the heart contracts, calcium fluxes into the heart cells from the extracellular
fluid. When the heart relaxes, the calcium is pumped back out of the cell allowing the fibrils to relax. Lowered extracellular calcium ion levels block contraction of the heart.
These data together suggest that fluoride ingestion is lethal by causing decreased activity of the free calcium ion and blockage of heart contractions. Fluoride acute toxicity is long known to be accompanied with increased blood potassium levels (Burgstahler, 23), where membrane sodium potassium ATPase is also inhibited by fluoride or low calcium. So the sequential mechanism of lethality cannot be precisely stated.
Fluoride tends to associate with and bind calcium anywhere this ion is concentrated throughout the body including teeth, bones, ligaments, aorta, skeletal muscle, and brain. But the most crucial physiologic function requiring calcium that is sensitive to industrial fluoride at acute levels is the beating heart. It is well known that extracellular calcium is an obligatory requirement for heart cells to undergo contraction after electrical excitation. Heart cells do not have well-developed sarcoplasmic reticulum to store calcium for this purpose as does skeletal muscle which does not exhibit this extreme sensitivity to changes in blood calcium. The cellular uptake of calcium occurs during the plateau phase of the cardiac action potential. Extracellular calcium is necessary for the development of contractile force. The strength of contraction (inotropic state) of the heart depends on extracellular calcium with half-maximal contractility at 0.5 mM.
Chronic Toxicity. The mechanism by which fluoride from blood at desired ‘low’ levels irreversibly accumulates in bone (National Research Council, 10). It does not involve precipitation of ionized calcium because fluoride is below the Ksp for direct precipitation. Instead, an ion exchange mechanism occurs at extremely minute fluoride levels where the fluoride ion merely by diffusion exchanges with hydroxide on bone hydroxyapatite. A fluoride ion solution made in soft or distilled water has a very high chemical activity, or chemical potential compared to the activity of the ion at the same concentration when accompanied by calcium or magnesium in solution. Although much less sensitive and exquisite than an actual biological cell membrane, a fluoride specific electrode senses such a difference.
Fluoride electrode measurements of a solution of sodium fluoride fixed at 0.8 mg/L (ppm) (0.042 mM) in de-ionized water were examined at various calcium levels over a wide range. 30 mM calcium causes substantial inter-ionic interactions with fluoride that significantly lower diffusion or Brownian motion of the fluoride ion because of the relatively massive divalent positive charge on the compact calcium ion. Further addition of calcium to 650 mM causes progressive decreases in the free fluoride ion measured level due to precipitation of calcium fluoride particles that the electrode cannot detect. The calcium level calculated to first begin fluoride precipitation for 0.9 ppm fluoride is approximately 30 mM which is consistent with Figure 1B data.
The phenomenon of fluoride-induced decreased chemical activity (i.e. in the absence of precipitation) applies to Group II cations including magnesium ion prevalent in all foods and natural hard waters. In contrast, fluoride accompanied in solution with Group I metal cations, such as sodium or potassium, exhibit no decline in activity over a broad range of cation concentration because these ions are only monovalent in charge (not shown).
Activity coefficients for the fluoride ion are substantially reduced in the presence of calcium and magnesium divalent cations (Moore, 11). This effect may be compared to the phenomenon of attraction between fluoride ion
and hydrogen atoms in water known as hydrogen bonding which decreases the Brownian motion and diffusion of the ion. These factors determine the overall biologic effect of fluoride ion in living organisms where calcium decreases assimilation through the gastrointestinal tract but in the bloodstream lowers the fluoride required for calcium sequestration. Further, membranes exhibit complex structural and functional features that are calcium-determined (Sauerheber and Gordon 12). Fluoride diffusion from a solution containing calcium ion may be impaired even though far below the level required for binding as calcium fluoride precipitate. The higher the calcium concentration of a region, the less fluoride is able to diffuse away from it. This electrical attractive force is also responsible for the fact that fluoride at levels below the Ksp is trapped in bone by ion exchange.
As shown in Figure 2, changes in pH affect the percent of fluoride that converts to HF. As HF, fluoride gains entry into the bloodstream because HF is a neutral small molecule comparable in size to the water molecule and is freely permeable through the biologic membrane (Whitford, 13). The Ka for HF indicates it is a weak acid but as a small molecule HF is the most penetrating corrosive. Its assimilation is most efficient at stomach pH and in the absence of calcium. Industrial fluoride in drinking water can cause GI distress in human subgroups because fluoride converts to HF in the stomach (NRC, 10, p. 268). Even at low concentrations HF can aggravate and prevent healing of ulcerated tissue. Industrial fluoride in drinking water at gastric pH is mostly protonated to HF. At pH 2, 96% of fluoride is HF (NRC, 10), in agreement with calculated levels here (Figure 2B). Structural damage to mucosa is detected with 20 ppm HF in 15 minutes, and lower concentrations cause pain without visible damage (NRC, 10). Human case studies proved that abdominal pain occurs from drinking 1 ppm artificially fluoridated water (Waldbott, 27; Petraborg, 28). In a careful placebo controlled double blind clinical study, 1 ppm artificial fluoridated water caused GI discomfort in 7% of subjects (Grimbergen, 29). The longer the exposure, and the more elderly the person with slower gastric mucosa turnover, the more overt symptoms could potentially become. In the presence of ulcers, industrial fluoridated water is contraindicated.
HF at high concentrations frosts glass and permeates porcelain. This is not due to the hydrogen or the fluoride ions since strong acids and soluble fluoride salts do not have this ability. Corrosiveness of HF is due to its extremely tiny uncharged structure, intermediate between fluorine F (0.8 Å3, 128 pm diameter) and the 10-fold larger volume fluoride ion F- (7.9 Å 3, 272 pm diameter).
Natural and Industrial Fluoride in Water. Assimilation of ingested fluoride is largely inhibited by calcium in the GI tract (CDC, 5). The conversion of fluoride to HF measured with an electrode is reduced in the presence of calcium at levels far below saturation (data not shown). Thus, the toxic effects of fluoride found naturally in some waters from calcium fluoride do not compare with water treated with an equal level of industrial fluoride. All artificial industrial fluoride compounds are toxic calcium chelators. Salmon are unaffected by natural 1 ppm fluoride in ocean water where calcium is extremely high, but are narcotized in fresh soft water at only 0.3 ppm industrial fluoride (Damkaer, Dey, 36) . The ratio of calcium ion molarity (around 0.12 mM) to added fluoride molarity (0.05 mM) in soft water states, particularly in the Pacific Northwest, in an artificially fluoridated city is very low. In moderate hardness water States, the ratio is typically about 80-100:1 or more so hard water regions are more protected from fluoride assimilation.
Blood levels appear to average 0.21 ppm from ingesting 1 ppm fluoride water (NRC, 10) (p. 70). The fluoride levels currently allowed in drinking water in the U.S. were found to not be protective of human health (NRC, 10), having adverse effects on brain and bone, which is the final resting site for 95% of all ingested fluoride. In bone, fluoride accumulates from consumption in a 1 ppm fluoride water region to 2,500 mg/kg in two years and to 3-4,000 mg/kg lifetime. Bone weakening occurs above 3,000 mg/kg. In elder years with fluoride-loaded bone, continuous consumption of fluoride, with reduced binding sites available in bone, should cause accumulation in soft tissues, including brain and, as discussed in Goodman and Gilman (21) in ligaments and tendons. Fluoride, at blood levels found from consumption of 1 ppm fluoridated water, can cross the blood-brain barrier in man (Yiamouyiannis, 34). This is perhaps by passage as trace neutral HF comparable in size to a water molecule. In animals (Reddy, 35), this blood level causes direct histopathology observed by electron microscopy. Long term exposure appears to decrease IQ in children (Connett, 2010). Variations in biologic outcome of course occur because of differences in water hardness and calcium content.
Studies at the Veterans Administration Health Care Center, Los Angeles report that systemic fluoride at sub-acute levels incorporates into atherosclerotic plaque in coronary vessels in cardiovascular disease patients directly revealed in PET scans (Yuxin, 14). Chronic ingestion of sub-acute concentrations of fluoride from drinking water weakens heart muscle in animal studies (U.S. CDC, 5) and causes alterations in heart function in humans (Varol, 15,16). Per capita cardiovascular deaths increased after Grand Rapids, Michigan and Newburgh, New York began industrial fluoridation (USPHS, 32). Increased incidence of EKG abnormalities occurred in patients with tooth fluorosis (Xu, 33). Our data suggest these effects may result from alterations in calcium homeostasis. Bone fluoride is retained for an estimated 20 years (NRC, 10) that affects calcium mobilization into blood.
Ironically, the level of fluoride in saliva that filters from the bloodstream after swallowing water with 1 ppm fluoride is a miniscule 0.02 ppm average (NRC, 10). This is unable to influence teeth cavities at 75,000 times lower concentration than in toothpaste at 1,500 ppm. Consistent with this, as found in the largest U.S. epidemiologic studies we have, industrial fluoridated cities had comparable caries incidence as non-fluoridated cities (Hileman, 17). The largest international study we have indicated lowest caries incidence in cities with lowest levels of water fluoride with calcium-sufficient diets (Teotia and Teotia, 22). The U.S. CDC (18) published findings that although high levels of topical fluoride might have a caries effect, systemic blood-borne fluoride from swallowing does not. In fact, systemic fluoride plays the most major role in causing the current U.S. high incidence of tooth fluorosis in children that prompted the U.S Health and Human Services to request in 2011 that water fluoride be lowered.
Normal teeth enamel is a calcium phosphate matrix. The original correlation in the U.S. between NATURAL calcium fluoride in water and caries incidence attributed the effect to fluoride. But calcium and magnesium levels totaled 302 ppm in water at the time. Ziegelbecker found the original correlation was mere scatter and that no relationship exists between caries incidence and water fluoride level when all data are considered (Connett, 19).
Water districts most commonly now inject artificial UNNATURAL industrial synthetic compounds into water to increase blood fluoride levels in consumers (personal communication CA Dept. of Public Health official).
Fluosilicic acid H2SiF6 lacking calcium is the most widely used substance (Connett, 19). Controlled human clinical trials for safety and effectiveness have never been completed with water treated with either sodium fluoride or fluosilicic acid. The U.S. Food and Drug Administration has never approved fluoride compounds for ingestion in the U.S. The FDA has written that fluoride is not a mineral nutrient and has labeled fluoride in water an uncontrolled use of an unapproved drug. In 1966 the FDA banned the sale of fluorides intended for ingestion by pregnant women because of lack of effectiveness in teeth of offspring.
Early interpretations of data collected from experiments in Newburgh, N.Y. and Grand Rapids, MI where synthetic industrial sodium fluoride was first infused into public water supplies in 1945 have also been widely discredited. The assimilated fluoride in the treated cities was associated with delayed teeth eruption where missing teeth were scored as absence of cavities (Sutton 24).
Other Substantial Fluorosilicic Acid Products in Water. Fluosilicic acid is not a source for fluoride in any natural water supply. Its infusion adds three ingredients that are not components of pristine fresh drinking waters (such as in the U.S. Pacific Northwest): fluoride, silicic acid, and sodium. Neither fluoride nor silicic acid are listed in the Merck Medical Manual or in Clinical Chemistry texts as constituents of normal human blood because neither are mineral requirements. Silicic acid H4SiO4 from fluorosilicic is typically 0.6 ppm when fluoride is adjusted to 0.8 ppm. Silicic with a low dissociation constant remains the intact acid even at alkaline pH and is the suspected ingredient that leeches lead salts in lead-based plumbing. The U.S. EPA has no MCL for silicic acid because it is not found as a contaminant in water supplies. NMR studies confirm that fluorosilicic acid fully dissociates into fluoride ion and silicic acid at community water pH but forms a silicofluoride complex at pH 3 (identified as SiF5-) (Finney, 30) as in stomach acid. A positive view has been presented for possible benefits of silicic acid consumption to cause soft fingernails and changes in skin structure (Barel, 31) due to stimulation of collagen formation by fibroblasts, but this effect may be non-physiologic.
In Figure 3 notice Southern California public drinking water sodium levels increased from industrial emissions along the source Colorado River to 85 ppm in 2006 prior to artificial fluoridation. After fluosilicic acid injections began in 2007 with sodium hydroxide required to neutralize acidity, the sodium level reached 93 ppm. Many plant species that have thrived in this region including the widely grown avocado crop are known to be saline intolerant (Musyimi, Netondo and Ouma, 20). Avocado leaf number, chlorophyll content, chloride content, root weight, and transpiration water loss rate are all altered by sodium in irrigation water. At 345 ppm sodium, chlorophyll content is reduced in leaves by 40%, chloride content increased 42% and transpiration rate of water loss decreased 21% after only 7 days treatment. The U.S. Health and Human Services January, 2011 took a commendable step and requested that fluoride infusions not exceed 0.7 ppm in public water supplies. This request was followed and the current level of sodium became reduced again to 85 ppm.
Conclusion. The data in this article indicate the decision to infuse industrial fluoride compounds into public water supplies to treat blood and organs of consumers with fluoride ion was in error, particularly for regions with calcium-deficient water. This conclusion was made in recent reports on the very high incidence of neurologic, cardiovascular, muscuskeletal and other pathology in silicofluoridated soft water Southern Ireland (Waugh, 25) compared to nonfluoridated Northern Ireland. Such sentiment was recently Nationally publicized in the U.S. (RT News, 26), and it is hoped that this policy will be discontinued.
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Acknowledgements. Much appreciation is expressed for the late, world-renown calcium physiologist Dr. C. Baird Hastings and colleague Dr. Andrew A. Benson, discoverer of the carbon dioxide fixation product in plant photosynthesis, in his laboratory today at the Scripps Institution of Oceanography, University of California, San Diego, for discussing the chemistry of industrial synthetic fluoride.
Figure 1. Calcium concentrations (mM) that precipitate various fluoride concentrations.
A. From the known solubility product constant Ksp for calcium fluoride, the calculated isocurve for CaF2 precipitation is shown as a function of varying concentrations for the two ions. Calcium levels in normal blood (1 – 3 mM) reach precipitation maximum solubility with 0.1 – 0.18 mM (1.9 – 3.4 ppm) fluoride. For 3 mM calcium, precipitation begins to occur when fluoride is 3% of calcium and for 1 mM at 18%. The Hooper Bay poisoning incident produced a lethal 5 ppm blood fluoride level in one victim. Whether it is a low Ca/F ratio or rather calcium fluoride saturation that is required for acute toxicity is not known.
B. A 0.9 ppm fluoride solution in distilled water was measured for free fluoride ion level with a LaMotte fluoride specific electrode calibrated with 1.00 ppm sodium fluoride in distilled de-ionized water at room temperature. Calcium ion was adjusted over a wide range by addition of aliquots of calcium biphosphate. Fluoride readings progressively decrease with increasing calcium concentration as expected, over the range 30 – 650 mM.
Figure 2. Fluoride protonation depends on prevailing acidity, following the equilibrium reaction F- + H+ ↔ HF. A. All measurements were made with a LaMotte Instruments fluoride ion specific electrode (calibrated with a 1.00 ppm fluoride standard solution in distilled deionized water at room temperature). Readings for the 1.2 ppm true concentration solution progressively decrease as pH decreases. Acidity was adjusted with dilute acetic acid. At stomach acid pH, fluoride is about 50% protonated as hydrofluoric acid HF, and 50% is free fluoride. B. The observed values agree with F- levels calculated theoretically from the Henderson Hasselbach equation [F-]/[HF] = 10(pH – pKa) over the pH range here, utilizing the known Ka for HF. HF decreases (diamonds) while F- increases (squares) from pH 1 to 8. At pH 3.14 (= pKa = -log 7.2 x 10-4), HF is half dissociated.
Figure 3. Data are from public published water quality reports from the Metropolitan Water District, Los Angeles for sodium as a function of year. The curves increase progressively after 2007 when industrial fluosilicic acid with caustic soda injections began. Every 24 tons of industrial fluosilicic acid requires 14 tons of sodium hydroxide to maintain pH at 8.4 (two H+ ions from H2SiF6 requires two sodium ions). Sodium at 116 ppm has been found to decrease yields and affect vegetable and fruit quality. Sodium is released into the Colorado River by scores of industries lining the river. The EPA Salt Abatement Program limits releases to one ton daily per site, but with so many sites has led to this level. The EPA secondary standard for TDS (500 ppm) is exceeded but is not enforced–plants can tolerate natural TDS from 800-1000 ppm. No MCL standards have been developed by EPA for sodium, since fresh water has historically been low in sodium. Sodium in blood is 3,000 ppm but is 0-10 ppm in pristine fresh drinking water with a national average at 15 ppm.