Fluoridation: Aspects of toxicity

[Malcolm Harris, Ph.D. (Wales), B.Pharm (Wales), FPS, Fss, FRSH]

Fluoridation: Errors and Omissions in Experimental Trials

[Chapters 19, 20 and 21. Philip Sutton. Originally published in 1960]

The Kinetics of Acetylcholinest'se Inhibition and the Influence of Fluoride and Fluoride Complexes on the Permeability of Erythrocyte Membranes

[Westendorf, 1975]

Introduction | Contents | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9

The Kinetics of Acetylcholinesterase Inhibition and the Influence of Fluoride and Fluoride Complexes on the Permeability of Erythrocyte Membranes - Page 5.

c. Dissociation Behavior of Several Fluoride Complexes at pH 7.4 

Fluoride is, in our opinion, not only found as F^- in both the living and non-living realm of nature, but often exists in complex bound form as well. It is still largely unknown if complex fluorides are of biological importance. If, or to be precise in which form, they enter the organism can be studied with the help of the radioactive isotopes ¹⁸F and ³¹Si (in the case of fluorosilicates). Since these isotopes were until now rarely available to us, we could only use them to carry out a few orienting preliminary tests.   

We therefore now occupied ourselves with the following questions. Which of the named complexes are stable at pH 7.4 (pH of the blood)? Is an influence on biochemical processes possible? And in particular, can the AChE inhibition be increased through the use of complex fluorides without thereby further raising the fluoride concentration? We studied the complexes BF₄^-, AlF₆^3-, SiF₆^2-  GeF₆^2-, SnF₆^2-, and PF₆^-. Of these, only the Al, Si, and P complexes are of natural importance, the latter however not in the form of PF₆^-, but instead as PO₃F^2-. The phosphoric acid residue may also be bound to organic residues (carbohydrates or adenosine). The remaining complexes we only studied for the sake of completeness, in order to possibly determine a relationship between the radius of the ion and the charge of the complex, and their effectiveness as inhibitors of enzymes. We studied the dissociation behavior of the complex fluorides by dissolving the complex salt in a buffer system and determining the free F^_ concentration with the help of an ion selective electrode. 

Properties of the Fluoride Electrode  

The fluoride electrode is a solid membrane electrode. The active electrode phase forms a single LaF₃ crystal, which is doped with Eu²⁺ to diminish the electrical resistance. The crystal can conduct fluoride. The external side contacts the test solution while the internal side contacts a fixed ion solution, which closes the measurement chain by way of a Ag/AgCl half element. The EMK of the measurement chain tracks the fluoride ion activity in the test solution. The NERNST equation yields the mathematical expression for the EMK trace. 

(equation 25)                      

If one chooses  pF value as an expression for the F^- activity (analogous to the pH value; ie the negative decadic logarithm of the F^- concentration) equation 25 can be rewritten in the form:

In our case the pF value was displayed directly by way of a digital voltmeter.  

The fluoride electrode possesses an unusually large selectivity, so that even a 1000 times excess of foreign ions does not bother it. Its functional region lies between 1-10^-5M F^-. Since the display of the instrument is influenced by a number of controlled variables (kind of buffer, pH value, temperature, stirring speed) it is necessary to record a calibration curve for each set of measurements and to maintain the controlled variables as exactly as possible. We carried out each of our measurements in 200ml Veronal/HCl buffer with a pH of 7.4 at 37^o C. Figure 15 indicates the course of the calibration curve recorded under these conditions. 

Figure 15 – pF Value as a Function of the F^- Concentration in Veronal/HCl-Buffer of PH 7.4 at 37^o C. 

Next we determined the level of hydrolysis of the individual complexes as a quotient of the concentration of free F^- (which can be derived from the calibration curve) and the total concentration of the fluoride atoms bound to the complex before the hydrolysis.

(equation 27)

In order to determine this value we first submerged the electrode in 200ml of buffer and waited until a constant p_F value was displayed, which was caused by the F^- that had gone into solution from the electrode. Then we added the complex salt as a solid and tracked the change in p_F value as a function of time until saturation. We determined the fluoride concentrations corresponding to the measured p_F values from the calibration curve and lastly calculated the level of hydrolysis a using equation 27.  

(Translator’s note: There is no text for nor any equation numbered “26 ”)

Hexafluorosilicate (as MgSiF₆)  

We tracked the speed of hydrolysis for two MgSiF­₆ concentrations.  

Figure 16 - Dependence of the Level of Hydrolysis of SiF₆^2- on Time. 

1) c = 5.7 x 10^-4M

2) c = 1.01 x 10^-4M  

The hydrolysis initially occurs very quickly. No more change occurred in the level of hydrolysis after only 15 minutes. We observed the process for several hours. Since the smaller concentration yielded a larger value for a we examined two further concentrations, for which we however only recorded the saturation value and plotted the level of hydrolysis as a function of MgSiF₆ concentration. 

Figure 17 - Level of Hydrolysis of SiF₆^2- as a Function of the Concentration 

The change in level of hydrolysis as a function of SiF₆^2- concentration is relatively small. Extrapolating the curve to even smaller concentrations should yield a level of hydrolysis for physiological concentrations of not more than 0.67, which corresponds to the splitting of four fluoride atoms from the complex.   

If one assumes that a uniform product forms as a result of hydrolysis, complex ions of the type [SiF₂(OH)₄]^2- should be present under these conditions, which by way of the pH value and temperature approximated physiological conditions. A coordination number other than 6 is not to be expected for the Si in aqueous solution. The small concentration inhibits chain formation, as it is often observed in silicon chemistry. Of course this possibility can nonetheless not be ruled out.   

Hexafluorogermanate (as K₂GeF₆)  

We produced this compound for ourselves in the following way. We dissolved germanium dioxide (GeO₂) in a platinum dish while heating in an excess of 30% hydrofluoric acid (H₂F₂). By adding the calculated amount of potassium carbonate (K₂CO₃) we precipitated the highly insoluble (0.542g/100ml at 18^o C) salt. We filtered out the precipitate, flushed it out with 3% hydrofluoric acid, and dried it in the exsiccator over phosphorus pentoxide (P₂O₅).  

Reactions:

We then determined the hydrolytic behavior of the complex using the technique used for MgSiF₆. 

Figure 18 - Hydrolysis of 1.1 x 10^–3 M GeF₆^2- as a Function of Time 

The initial slow climb of hydrolysis is noteworthy. Since hydrolysis represents an ionic reaction, one night expect equilibrium to be established quickly. But the course of this curve may reflect dissolving speed of the salt, (a diffusion-dependent processon proportional to t ^1/2 by. Fick's Rule). A plot of a vs. t^1/2 is linear from t = 0 to t = 50, which speaks for this suspicion. 

Figure 19 - Level of Hydrolysis of K₂GeF₆ as a Function of t^1/2 

The dissociation level of the saturation, at 0.83, corresponds exactly to the splitting of 5 F^- out of the complex. If a complex of the form [GeF(OH)₅]^2- exists, or if higher molecular aggregates form through condensation, can not be determined using the available materials.   

Hexafluorstoannate (as K₂SnF₆·H₂O)  

The salt was produced using the procedure applied for K₂GeF₆. However, we used SnCl₄ as the initial substance. Chlorine was expelled from this substance as HCl by way of repeated steaming with 30% HF. The potassium salt crystallizes with one mole of crystal water. The hydrolysis experiment yielded a complete breakdown of the substance after only five minutes. Since nothing else special occurred a further representation of the experiment will be omitted.  

Hexafluoroaluminate (as Na₃AlF₆)  

This compound is of greater biological importance since it is widespread in nature in the form of cryolite and can therefore be taken up by the human body. This compound appears at elevated concentrations in the exhaust and wastewater near aluminum factories, which use this substance as a fluxing material in melt-electrolysis, so that a burden for humans and animals beyond the physiologically justifiable region can arise. The solubility of this compound in water is minimal (0.042g/100ml). We studied a concentration of 0.03725 g/l = 1.78 x 10^-4M.

Figure 20 - Hydrolysis of 1.78 x 10^-4M AlF₆^3- as a Function of Time 

The rate of hydrolysis is slower than in the case of SiF₆^2-, but faster than in the case of GeF₆^2-. A constant value of a = 0.695, which lies only slightly above the value for a separation of 4 fluoride atoms (a = 0.67), is reached after 40 min. The hydrolytic behavior of the cryolites is thereby similar, at this pH value of 7.4, to that of the hexafluorosilicates.  

Hexafluorophosphate and Tetrafluoroborate (as KPF₆ and KBF₄) 

We included these two substances in the study as representatives of monovalent complexes. We used concentrations of: KPF₆ = 1.67 x 10^-4M and KBF₄ = 2.34 x 10^-4M 

Figure 21 - Hydrolysis of PF₆^- and BF₄^- as a Function of Time 

These two substances are remarkably stable in comparison to those dealt with up to now. The level of hydrolysis at saturation in both cases lies below the value for the separation of one mole F^- per mole of complex: 

KPF₆: a_s = 0.0209 ; a_1/6 = 0.17

KBF₄: a_s = 0.068 ; a_1/4 = 0.25 

To these considerations we also add an overview of the hydrolytic behaviors of the studied complexes in the form of the following table.

Table 2. Degree of Complex Dissociation at Physiological Conditions, pH 7.4, T = 37^o C  

The experiments showed that several fluoride complexes, of which the hexafluorosilicate and the cryolites are found in nature, do not fully hydrolyze under "quasi-physiological" conditions. When these compounds are ingested as part of the nutrition, one must expect the appearance of such partially hydrolyzed "intermediate complexes" in the body, (assuming re-absorption).  These complexes are most likely to appear when resorption occurs in the acidic medium of the stomach, in which case hydrolysis only begins in the blood. If the complexes first reach lower sections of the intestines they will be more extensively dissociated because of the alkaline medium that prevails there. It will be possible to follow the resorption of these compounds with the help of the isotopes ¹⁸F and ³¹Si.

Introduction | Contents | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9