ham J Environ. Stadia, 2000, Vol. 57, pp. 597-637 © 2000 OPA (Overseas Publishers Association) N.V.
Reprints available directly from the publisher • Published by license under
Photocopying permitted by license only *<= Gordon and Breach Science
Publishers imprint.
Primed in Malaysia.
CAN FLUORIDATION AFFECT
LEAD(H) IN POTABLE WATER?
HEXAFLUOROSILICATE AND FLUORIDE
EQUILIBRIA IN AQUEOUS SOLUTION*
EDWARD T. URBANSKY* and MICHAEL R. SCHOCK*
United States Environmental Protection Agency (EPA), Office of Research
and Development, National Risk Management Research Laboratory,
Water Supply and Water Resources Division, Treatment Technology
Evaluation Branch, Cincinnati, Ohio 45268 USA
(Received 10 January 1999)
Recent reports have attempted to show that fluoridating potable water is linked to
increased levels of lead(II) in the blood. We examine these claims in light of the estab-
lished science and critically evaluate their significance. The completeness of hexafluoro-
silicate hydrolysis is of paramount importance in ensuring that total water quality^
maintained. The possible impacts of such complexes as Pb —F—SiF5 or PbFJ.
are discussed as are the contributions of fluoridation byproducts to total acid content.
We calculate the fractional distribution of aqueous species based on known chemical
equilibria and show the species concentrations for several different model tap waters.
We discuss and quantitatively show the effects of other complexing anions, such as car-
bonate or hydroxide. Overall, we conclude that no credible evidence exists to show that
water fluoridation has any quantitatable effects on the solubility, bioavailability, bio-
accumulation, or reactivity of lead(O) or lead(II) compounds. The governing factors are
the concentrations of a number of other species, such as (bicarbonate, hydroxide, or
chloride, whose effects far exceed those of fluoride or fluorosilicates under drinking water
conditions. Lastly, we consider some previous epidemiological studies of lead(II) ex-
posure and how recent papers fare methodologically.
Keywords: Fluoridation; fluoride; fluosilicic acid; hexafluorosilicate; hexafluorosilicic
acid; lead(II); potable water; silicofluoride; water treatment
*This paper is the work product of United States government employees engaged
in their official duties. As such, it is in the public domain and not subject to copyright
restrictions.
tCorresponding author. Tel.: 513-569-7655, e-mail: urbansky.edward@epa.gov
*Tel.: 513-569-7412, e-mail: schock.michael@epa.gov
597
-------
598 E. T. URBANSKY AND M. R. SCHOCK
1. INTRODUCTION
Controversy over water fluoridation varies in nature and intensity.
Recent papers have implications for water fluoridation since they sug-
gest that certain adverse health or social conditions may stem from
interactions between lead(II) and inorganic {more-compounds,
specifically, fluorosilicates and fluoride [1-3]. In order to assess the
validity of these assertions, it is necessary to have a firm foundation of
the aqueous chemistry of H2SiF6 and HF. There is a considerable body
of fundamental chemical literature on these species. Nonetheless,
some gaps do remain, and little effort has been expanded in combin-
ing the known chemistry into one comprehensive and authoritative
volume. Accordingly, we believe it to be worthwhile to revisit the con-
cepts involved in water fluoridation at a fundamental level and to
examine some of the relationship suggested by recent papers in light
of well-established science.
The sheer number of people consuming fluoridated potable water
makes fluoridation issues relevant. In 1992, the Centers for Disease
Control Fluoridation Census found that 62.1% of the U.S. popula-
tion served by public suppliers drank fluoridated water [4]. The CDC
also surveyed utilities regarding fluoridating agents (see Tab. I). Most
commonly used are hexafluorosilicic acid (H2SiF6) or its sodium salt
(Na2SiF6), which hydrolyze to produce fluoride ion upon dilution
(l)-(2). However, sodium fluoride (NaF) is sometimes used as a direct
fluoride source (3).
TABLE I Water fluoridation chemicals used by U.S. public water suppliers in 1992*
formula
common
synonyms
population
served
utilities
using
Hexafluorosilicic acid
H2SiF6
fluosilicic acid
fluorosilicic acid
hydrofluosilicic acid
80,019,175
62.6%f
5876
Sodium hexqfluorosilicate
Na2SiF6
sodium silicofluoride
sodium fluorosilicate
36,084,896
28.2%f
1635
Sodium fluoride
NaF
-
11,701.979
9.2%f
2491
•Total US population: 258,544,000. Population and utility data were taken from Ref. [4].
* Percentages are based on total population of 127.8 million persons drinking fluoride-fortified public
water and does not include those drinking water naturally high in fluoride.
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 599
H2SiF6(aq) +4H2O -* 6HF(aq) + Si(OH)4(aq) (1)
Na2SiF6(aq) + 4H2O -» 4HF(aq) + 2NaF(aq) + Si(OH)4(aq) (2)
NaF->Na+(aq) + F~(aq) (3)
Hexafluorosilicic acid is a cheap and readily available source of
fluoride. However, it is difficult to handle and the handling costs can
only be offset by the volume discount in large water treatment plants.
Although more systems rely on sodium fluoride than sodium hexa-
fluorosilicate, these serve only 9.2% of the U.S. population. Because
sodium fluoride is the easiest of the three to handle and dispense,
small systems are the primary users of NaF. Although the EPA regu-
lates drinking water, the US Public Health Service has been involved
in water fluoridation for historical reasons (primarily because the
practice of fluoridation pre-dates EPA). The purpose of fluoridating
water is the prevention of dental caries; therefore, the publication
of water fluoridation how,to manuals falls under the purview of the
CDC. These manuals discuss dosing and other practical matters of con-
cern to the treatment plant operator [5].
2. CAUSE FOR CONCERN?
Potable water from large public water supplies contains a large
number of compounds, including disinfection byproducts (such as
trihalomethanes or haloacetic acids), residual oxidants (such as
chlorine or chloramine), nonspecific dissolved organic matter, trace
metals, minerals (such as sodium chloride or calcium carbonate), and
additives (such as fluoride). Consequently, drinking water science is a
complicated interplay among the chemical constituents as well as the
physical conditions, such as temperature.
This paper evolved in large part as a response to relationships
(posited in the literature [1-3]) between the practice of water fluori-
dation and human blood concentrations of lead(II). It is instructive
to address the issues regarding water fluoridation and lead(II) by
-------
600 E. T. URBANSKY AND M. R. SCHOCK
posing a series of questions regarding the fundamental behaviors of
fluoride, silicon and lead. We will explore the answers to these ques-
tions in the sections that follow.
1. What is the residual concentration of hexafluorosilicate ion
(SiFg~) after the hydrolysis reaction (2) takes place?
2. How fast does reaction (2) occur?
3. Do fluoridation additives affect the pH of the finished water at
the plant or at the tap? If there is a pH change, does it matter?
Can H+ from hydrolysis of residual SiFg~ promote the solubiliza-
tion of lead(II) from the distribution system, thereby increasing
the lead(II) concentration at the tap? '
4. Can F~ or residual SiFg~ complex with lead(II) and make it more
bioavailable? In other words, do fluoro-species complex with
Pb(II), promoting permeation of the gastric mucosa and absorp-
tion into the bloodstream.
5. Can residual SiFg" lower .gastric pH and therefore convert par-
ticulate Pb or lead compounds to bioavailable aqueous lead(II)
ion?
6. How do the lead(II) drinking water regulations and sampling
schemes relate to human health effects?
7. How is human lead exposure measured? What are the weaknesses
in its quantitation?
8. What quality controls exist for drinking water additives? Can
the additives themselves be responsible for contaminants in
the water supply? Specifically, can they be a source of lead
exposure?
9. What are the routes and nature of human lead exposure? What
sorts of factors are linked to human lead exposure?
10. When all of the chemical and physical phenomena are considered
together, what is the magnitude of the effect? That is, what species
represent the greatest fractions of total lead(II) and fluoride con-
centrations under potable water conditions?
Certainly, we cannot fully answer all of these questions in one
paper. However, we can summarize the principal findings and refine
our understanding of the role drinking water plays. Let us begin to
answer these questions at the molecular level by examining what takes
place when sodium hexafluorosilicate is added to the water stream
inside a utility plant.
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 601
3. THE CHEMISTRY OF HEXAFLUOROSILICATE
AND FLUORIDE IN AQUEOUS SOLUTION
3.1. Equilibria and Kinetics of Hexafluorosilicate Hydrolysis
3.1.1. Equilibria
Hexafluorosilicate ion reacts with water to produce fluoride ion and an
assortment of silicon oxyanions [6,7], e.g., SiO3~,SiO4~,Si(OH)O3 .
We represent the oxyanions as SiIV(aq) without further speciation at
this time.
(4)
A principal issue (Ql) at stake is how much residual fluorosilicate is
present, that is, whether Eq. (4) proceeds to completion. Of course,
we must define completion. What fraction of total silicon(IV) is pre-
sent as a fluoro-species? Is it 1%,0.1%, or lower?
The actual speciation of silicon oxyanions is a function of acidity, i.e.
[H+]. Busey.e* al. [8] showed that virtually 100% of the hexafluoro-
silicate is hydrolyzed to silicon oxyanions at pH 6, even when there is a
free fluoride concentration of 0.01 M. Meanwhile, fluoridated drink-
ing water contains only ~1 ppm fluoride, which equates to 5 x 10~ M.
Previous investigations [9, 10] found a non-negligible concentration of
residual SiF41 when this gas was passed through water. Ciavatta et al.
[9] investigated fluorosilicate equilibria with 0.3 < [H+] < 3m[mol F~
(kg water)"1] and ionic strength fixed at 3M, adjusted with LiClO4.
They concluded that the mixed ligand species2 SiF(OH)3 and
SiF(OH)2(H2O)+ are significant contributors to total silicon(IV) in
JWhen dissolved in water, it is probably reasonable to view silicon(IV) compounds
as hexacoordinated so that SiF4(aq) is actually better represented as SiF4(H2O)2(aq).
Possible exceptions to this would involve the formation of pi bonds, such as
Si(*r) <-O(/w) back donation. However, even Si(OH)4 can be thought of as having two
weakly attached axial H2O molecules.
2Choosing a nomenclature system for the mixed ligand species is somewhat compli-
cated. Any name is likely to impart a certain amount of confusion since the naming of
these compounds has historically been lax. For example, SiF(OH)3 might be called
(mono)fluorotrihydroxysilane or (mono)fiuorosilanetriol (organic), silicon (mono)fluor-
ide trihydroxide (inorganic salt), mononuorosilicic acid (nonmetal oxyacid), or fiuoro-
trihydroxosilicon(IV) (coordination complex). A particular name might be more suited
to one context, but there is no truly preferable name. Naming as derivatives of silicic acid
is problematic because there is confusion over whether silicic acid means Si(OH)4
or SiO(OH)2- SiF(OH)2(H2O)+ seems best named as a coordination complex, aquo-
fiuorodihydroxosilicon(IV). Even though that name is cumbersome, the other nomen-
clature systems are incapable of describing such a species.
-------
602 E. T. URBANSKY AND M. R. SCHOCK
addition to SiF4, SiF^~ and HSiFJ under these conditions. None-
theless, their results showed that fluoro-complexes comprised less
than 5mol% of the total silicon(IV) in 0.01 mH+ and 10~4mF~.
Korobitsyn et al. [10] examined the hydrolysis of sodium hexafluor-
osilicate in sodium carbonate solution. Their work was geared towards
an industrial process for producing sodium fluoride and is not directly
applicable here.
The use of chemical shift information derived from 19FNMR spec-
trometry in understanding the formation of fluoro-ligated species
is well-established [11-16]. Fluoride ligand exchange occurs rapidly be-
tween HF and SiFg~ at temperatures above -10°C [13], and the identi-
fication of aqueous fluorosilicate species and the measurement of the
concommitant equilibrium constants has been done almost entirely
by 19F NMR spectroscopy and spectrophotometry [14-16]. The
Gmelin Handbook of Inorganic Chemistry tabulates values for the equi-
librium constants expression (6) of the hydrolysis reaction (5) at tem-
peratures from 0 to 60°C [17].
j:_[Si(OH)4][H+]4[F-]6 (6)
The smallest value at ambient temperature reported for K is 10~31'6.
Using this value at [H+] = 10~6M and [F~] = 5 x lO^M, the ratio
[Si(OH)4]/[SiF^~] = 1.6 x 1018. Note that less than 1% of fluoride
exists as HF at drinking water acid levels (i.e., pH > 5.2) since
ptffF = 3.17 [18]. Even if the hydrolysis constant were off by a factor
of 1000, it would not matter. There would still be essentially no hexa-
fluorosilicate ion. A fractional distribution plot in Gmelin [17] shows
that other fluorosilicates (i.e., SiF4 and SiFJ) also drop oft" dramati-
cally as free fluoride concentration, and not [F~]T, decreases towards
10~4M, even in silica-saturated 4M perchloric acid. For this solution,
total fluoride concentration is expressible as (7), neglecting any mixed
fluorohydroxo-ligated species:
[F-]T = [HF] + [F-] + 4 [SiF4] + 5 [SiFJ] + 6 [SiF^] (7)
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 603
Crosby studied the dissociation of sodium hexafluorosilicate and
hexafluorosilicic acid in deionized water [19]. He found that about
99mol% of the hexafluorosilicate had hydrolyzed when added to
water to produce a 1 ppm fluoride solution; however, the pH of this
solution was 4.20, considerably below a potable water pH. An im-
portant factor must be considered in potable water fluorid'ation as
Crosby explains:
It should be remembered that the actual ionic population of
most public'drinking-water supplies is somewhat different from
the experimental conditions used in the present and previous
studies. Thus, the pH is normally adjusted to about 7 to 8, and
the presence of additional salts may further influence the equi-
librium owing to the formation of complexes with calcium and
other metals.
If the pH of a treated drinking water is too low, it is adjusted to
comply with regulations and minimize corrosion. We do not dispute
Crosby's results. His water was demineralized and completely devoid
of buffering agents. Consequently, the dissociation of hexafluorosili-
cate was hindered by the drop in pH. Crosby's fractional dissociation
data cannot be applied directly to a potable water supply without cor-
recting them for pH. Of course, that correction is the effect we have
computed above, namely, the complete hydrolysis of fluorosilicates.
This is precisely what Crosby was emphasizing. This observation
hints at the answer to Q3 (effect on pH), which we shall come back
to shortly.
Interestingly enough, a number of species actually promote the
dissociation of hexafluorosilicate, including ferric ion [20]. While the
compound PbSiF6-2H2O can be synthesized, it decomposes quickly
in moist air and slowly when dry [21]. Perhaps then lead(II) itself
promotes hexafluorosilicate decomposition, such as through the for-
mation of plumbous fluoride. Because moist air promotes this
compound's destruction, we can infer that it would not be stable in
aqueous solution at all. Based on the above computations and observa-
tions, we can dispense with the issue of incomplete hydrolysis entirely.
There is essentially no hexafluorosilicate remaining in drinking water
at equilibrium.
-------
604 E. T. URBANSKY AND M. R. SCHOCK
3.1.2. Kinetics and Mechanisms
Now that we have answered Ql by concluding that the hydrolysis
reaction (4) does in fact proceed to completion, we must consider how
fast it reaches that state (Q2). The first major kinetics studies were
performed in the early 20th century by Hudleston and Bassett [22],
who studied the reaction between hexafluorosilicic acid and sodium
hydroxide by adding an excess of the acid to a known volume of base
and measuring the time for the phenolphthalein color to fade, and
then titrating to the end point with more base. The problem with such
a design is that excess acid stabilizes the hexafluorosilicate and the
silicon tetrafluoride from hydrolyzing. Hudleston and Bassett did
conclude that SiF4 hydrolysis was fast; however, it was not favored
under their conditions. Rees and Hudleston added various volumes
of standard NaOH(aq) solution to a set volume of hexafluorosilicic
acid solution [23]. There was no difference in the time required for the
phenolphthalein color to fade. Consequently, they concluded that re-
action was zerqth order in [OH~] and first order in [SiF§~]. At the
time, the study of chemical kinetics was in its infancy, and the con-
cepts of rapid pre-equilibrium and steady-state intermediates were not
well-established. Suppose the reaction mechanism proceeds along
these lines:
(8)
(9)
(10)
(11)
This mechanism immediately suggests a steady state in the concentra-
tion of tetrafluorosilane. The steady state approximation requires an
unstable intermediate of limited lifetime, capable of remaining pre-
sent at very low levels as it is simultaneously produced in (9) and con-
sumed by (-9) and (10). Applying the steady state approximation, we
set the rate of change of this species to zero, i.e., d[SiF4]M/d* = 0.
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 605
The net reaction rate is expressible as (12):
= fcio[SiF4UOH-]' '(12)
Making use of the steady state approximation, we obtain the following
differential rate expression (13):
rate = kwk9K&[Si¥26-}(OU-}/{(¥-](k^[F-} + *io[OH-])} (13)
At sufficiently high hydroxide concentrations where fc]0 [OH~]»fc_9
[F~], the dependence of the rate on hydroxide concentration would
disappear. Although an inverse variation with fluoride concentration
would exist, Rees and Hudleston did not vary fluoride concentration.
Note that we do not advocate this mechanism per se, but point out
that it is consistent with Rees and Hudleston's data. Since they did
not vary fluoride concentration, it is impossible to say whether their
data support the inverse dependence suggested by Eq. (13). To our
knowledge, the kinetics and mechanisms of hexafiuorosilicate hydro-
lysis have not been adequately investigated at this time other than to
suggest that the rate is fast enough for equilibrium to be achieved in a
period of less than 30 minutes. Given the limits of laboratory apparati
and computing/data analysis technology in the 1920s and 1930s, it is
unsurprising that the authors did not carry the work further.
In the 1970s, Plakhotnik conducted studies into the effects of lithium
and calcium cations on the rate of hexafiuorosilicate (and tetrafluoro-
borate) hydrolysis [24,25]. Plakhotnik concluded that a rapid pre-
equilibrium existed between the hexafiuorosilicate and the metal cation
with a fluoride ion acting as an inner sphere bridging ligand to both
the Si™ and the Li+ or Ca2+. Plakhotnik found that the reaction was
catalyzed by both metal cations. He allowed for contributions from
hydroxide assistance, but did not quantitate these. Although solutions
were reported to have been made alkaline, neither a pH nor a hydrox-
ide concentration is given. Presumably, all reactions were carried
out at the same base concentration. Calcium accelerated the reaction
by a factor of more than 10, while lithium accelerated it by a factor
of perhaps 1.2. Based on Plakhotnik's results, we calculate that the
hydrolysis would be 99mol% complete in 12 minutes if carried
entirely by the uncatalyzed pathway. That notwithstanding, natural
water supplies do contain calcium and other divalent metals as well as
-------
606 E. T. URBANSKY AND M. R. SCHOCK
trivalent metal cations (e.g., A13+, Fe3+) either native or introduced
during other treatment processes, such as agglutination (coagulation).
Therefore, we can state that the actual hydrolysis rate would be even
faster. And so we have answered Q2: the hydrolysis of hexafluoro-
silicate proceeds to completion before water reaches the consumer's
tap.
Based on the above information on both the thermodynamics of the
hydrolysis reaction and its kinetics, we can safely conclude that there
is essentially no hexafluorosilicate remaining in drinking water at equi-
librium and that equilibrium is rapidly reached from the combined
uncatalyzed and metal-catalyzed reactions.
3.7.5. Significance of Common Practices in Analytical
Chemistry
Having spent so much time demonstrating the completeness of hexa-
fluorosilicate hydrolysis, we fear that someone who is familiar with
two common practices in fluoride measurement, but not the chemical
logic behind them, may now be puzzled. Accordingly, we take a mo-
ment to explain.
3.1 J.I. On the Determination of Fluoride in Potable Water One
might wonder about the sample pre-treatment steps used in the po-
icmiometric measurement of fluoride in potable water as one basis
for concluding that hydrolysis is incomplete. In the procedure, sever-
al reagents are added to a potable water sample in the form of a total
ionic strength adjustment buffer (TISAB). The TISAB contains so-
dium chloride, acetic acid, and cyclohexanediaminotetraacetic acid
(CyDTA) which forms more stable and less labile complexes than
the more common EDTA. As a chelating agent, CyDTA binds up
metal cations (e.g., A13+) that might otherwise complex with fluoride
and thereby interfere in the titration. Complexation of fluoride with
silicon(IV) is not a concern.
3.1.3.2. On the Determination of Total Fluoride in Sodium Hexafluoro-
silicate In this procedure, the analyte is total fluoride. Since most
of the fluoride is in the form of SiF|~, [F-]T « 6[SiF^-]0 (the initial
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 607
concentration of hexafluorosilicate). In the sample of reagent material,
[SiFg~]0 is presumed to be at a concentration of 0.01 M or higher,
and is titrated with 0.2 M NaOH(aq) solution while boiling. Be-
cause these conditions are far more forcing than what occurs in water
treatment, do these standard practices constitute a basis for
concluding that hexafluorosilicate hydrolysis is incomplete? Yes. Of
course the hydrolysis will be incomplete in the raw material. The to-
tal fluoride and total acid are much higher than drinking water
conditions. In fact, these conditions approach (or possibly exceed)
Crosby's (vide supra). As with any titrimetric method, one strives for
a very rapid reaction that occurs immediately upon addition of a drop
of titrant. Boiling accomplishes this and it minimizes carbon dioxide
introduction. With a phenolphthalein end point, it is necessary that
the reaction reach completion on a millisecond time scale; otherwise,
a falsely low end point could result.
3.2. Potential for Solubilizing and Concentrating Plumbous
Ion in the Distribution System
In Q3, we inquired whether additives could affect pH. Since Eq. (2)
produces hydrogen ions, a pH drop is the only possibility. Conse-
quently, we would like to ascertain whether this pH drop is capable
of either dissolving otherwise insoluble lead(II) compounds that coat
pipe walls (14) -(16) or dissolving metallic lead (as from solder or
brass), specifically particulate lead(0)?
PbX(s) + 2 H+(aq) ^ Pb2+(aq) + H2X(aq) ( 14)
(15)
Pb3(C03)2(OH)2(s) + 4H+(aq) ^ 3Pb2+(aq)
+ 2H20(1) (16)
where X2~ is some anion that forms an insoluble lead(II) compound
which exists at least partly as particulate matter (i.e., in the solid
state and not a soluble complex), for example, CO2", O2~, or SO2;".
Hydrogen ion can react with metallic lead in a redox reaction
(17), but this direct reaction is not observable under drinking water
-------
608 . E. T. URBANSKY AND M. R. SCHOCK
conditions. Other species that affect lead(II) solubility, such as carbo-
nate and hydroxide, are affected by [H+], but this is not via a redox
phenomenon.
Pb + 2HF(aq)^Pb2+(aq)+2F-+H2 (17)
The solubilizing effects of hydroxide, bicarbonate, carbonate and
even chloride (though to a lesser extent) far outweigh the effect of this
small amount of acid. In addition, we must not ignore the buffering
capacity of a natural water resulting predominantly from bicarbonate
and carbonate and to a lesser degree from organic matter, silicates,
and hydrolytic (Lewis acidic) metal cations (e.g., Fe3+,Mg2+,Ca2+).
Moreover, the oxidizing capacity of the residual chlorine, which may
be as high as 1 mg LT1 (= 14uM = 28uN) is about half that of the
HF, which is present at 53 uM ( = 53 uN).
A later section will describe the speciation of lead(II) in drinking
water using a model constructed from multiple simultaneous equi-
libria. The concentrations of PbF+(aq) and PbF2(aq) will be shown
to be negligible compared to carbonate-, hydroxo- and other ligand
complexes of the plumbous cation.
A more quantitative treatment of buffer capacity will come later.
To answer Q3 for now, we can say that pH is controlled by many fac-
tors of greater magnitude than the hydrogen ion produced by (2).
Furthermore, the small change in pH induced by (2) cannot solubilize
lead(II) to the extent that the basic ligands, such as carbonate can.
3.3. Potential Effects on Lead(II) Unavailability
Now we will consider chemistry under some physiological conditions.
Even though we have demonstrated that there is no hexafluorosilicate
remaining by the time water reaches the consumer's tap, let us assume
momentarily that this is not the case. The following beyond-worst-case
scenarios nicely illustrate the magnitude of the effects on lead(II).
•2-
5.3.7. Possible fi-fluoro Complexes ofPb" and SiF\
In Q4, we ask whether fluoride or hexafluorosilicate could complex
with lead(II). Let us first consider SiF2". The only coordination option
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 609
for Pb" and SiF2" is a bridging fluoride ion (symbolized by /z-F)
attached to both the lead(II) and the silicon(IV).
To produce 1.0 ppm fluoride requires an initial hexafluorosilicate
concentration of 8.8 uM. The hydrolysis reaction (2) is a reversible
equilibrium, and in the most acidic gastric conditions, the pH could
be as low as 1.5 so that [H+] = KT1'5 M.3 Using this hydrogen ion con-
centration, we calculate the ratio [Si(OH)4]/[SiF^~] = 4.5 x 10s. This
means that only 0.00022% of the total silicon(IV) is present as the
hexafluorosilicate ion so that [SiF2" ] — 1.9 x 10~nM = 19 picomolar
(pM).
Haque and Cyr showed that hexafluorosilicate anion forms com-
plexes with several metal cations: Cu", Ni", Co11, and Fem [26]. The
largest stability constant they obtained was for the reaction with
ferrous ion, with K= 1.2 = 10°'08. Let us assume that the lead(II) ion
forms a stabler complex and set the stability constant for Eq. (18) to
an arbitrarily high value of 100. In addition, we shall pretend that
the hydrolysis computed above has not occurred, that all 8.8 uM of
the silicon(IV) remains in the form of hexafluorosilicate ion.
Pb2+ + SiF2.- ^ PbSiF6(aq), K = 100 (hypothetical) (18)
In this worst case, only 0.088 mol% of the total lead(II) would be in
the form of a hexafluorosilicate complex. The /z-fluoro ligand would
serve as a link between the silicon(IV) and lead(II). How significant is
0.088 mol% of the lead(II)?
Now consider the actual case where [SiF2,"] = 19 pM. If the
equilibrium constant for (18) were larger, say 106, so that the reaction
could be treated as going to near completion, there would still be less
than 19pM Pb—p-F—SiFs. Because of the magnitude of the equilib-
rium constant for (6), the equilibrium constant for (18) would have
to exceed 1025 in order to have a quantitatable effect by prevent-
ing hexafluorosilicate hydrolysis. There is no basis in fact for such
3We assume activity coefficients of unity so that activities and molar concentrations
are approximately equal. Although one could attempt to account for the differences by
an extended Debye-Huckel treatment, there is little point since there is rarely more than
0.4 logarithm unit variation. Given the uncertainties in the equilibrium constants and the
magnitude of the concentrations, a variation of even a factor of 100 is not especially
significant, let alone what would be expected from the activity coefficient corrections.
The same applies to differences in temperature.
-------
610 E. T. URBANSKY AND M. R. SCHOCK
an assertion. As a final point, we note that the national primary
drinking water standards are intentionally predicated on the assump-
tion that all lead is bioavailable, and the water utilities should be
complying with these standards.
Before we deal with the other part of Q4, fluoro-complexes of
lead(II), let us consider how residual hexafluorosilicate might affect
gastric pH (Q5).
5.5.2. Post-ingestion Gastrointestinal Chemistry
Imagine that the hexafluorosilicate ion has escaped hydrolysis up to
the moment of consumption of a glass of water. Imagine further that
it is suddenly hydrolyzed upon entering the stomach, where it now
forms hydrofluoric acid.
How does the pH drop from the hexafluorosilicate hydrolysis
compare with that from other sources of acidity? In other words, how
many protons (hydrogen ions) are produced? Drinking water treated
to contain 1 ppm fluoride would contain 53 (J.M HF(aq). If one were to
drink this solution of 53 p.M HF(aq), which is 93mol% dissociated
to hydrogen and fluoride ions, it would contain 49 (iM H+ and its
solution would have a pH of 4.3. Meanwhile, the high extreme for
stomach pH (lowest acidity) is about 3 (1000 nM H+); the lowest
stomach pH is about 1.5 (for optimal pepsin enzymatic activity in the
digestion of protein). At pH 3, roughly half of the HF will not ionize
since it is a weak acid. Meanwhile, some foods are equally or more
acidic, for example, apple (pH 2.9) or tomato juice (pH 4.4). Using
this logic, we would expect those persons who consume large amounts
of acidic foods or drinks to suffer from lead toxicity.
One might be tempted to counterargue that a person consuming
fruit juices has excluded particulate lead (from the distribution system)
except that many fruit drinks are reconstituted using tap water, and
soda fountain drinks are prepared likewise. Thus, consumers of soft
drinks should be at especially higher risk, given the high concentra-
tions of complexing organic acids (e.g., citric and tartaric acids in
powdered fruit drink mixes) or inorganic acids (e.g., phosphoric and
carbonic acids in colas). In fact, Coleman et al, showed that chelating
organic bases (e.g., citrate, ascorbate, EDTA) promote the transport
of lead(II) in the small intestine [27]. The acidic components of these
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 611
beverages completely overwhelm the contribution from HF in the
water used to prepare them. Whether any of these other species is pre-
sent in sufficient concentration to influence bioavailability is unknown.
Regardless, acid from hexafluorosilicate-based fluoridation is negli-
gible compared to other sources of acid in the American diet. Conse-
quently, one cannot demonstrate that an increase in blood lead(II)
ion levels can be linked to acidity from SiFg" hydrolysis as opposed
to consuming fruit drinks, carbonated beverages, or frozen orange
juice on a city-wide scale. Since we expect consumption of these pro-
ducts to occur without respect to water utility, suggestions of effects
due to SiFg~ within the digestive tract cannot play an important role
in lead(II) gastroenterochemistry.
A pH effect-if it actually existed-would only pertain to the
stomach. The small intestine does not absorb divalent ions well. In
addition, bile (from the gall bladder) and bicarbonate (secreted by
the pancreas) raise the pH and effectively buffer against pH change.
Partly digested food in the chyme will also act as a buffer. Moreover,
normal gastric biophysiology resists changes in acidity by a mecha-
nism involving gastrin secretion and activity for which a detailed
description is beyond the scope of this work. In conclusion, the produc-
tion of acid from fluoridation of water is insignificant when compared
to other acids and bases supplied by a normal diet or physiological
mechanisms.
3.3.3. Fluoro-complexes of the Plumbous Ion
What about the effect of the fluoride itself (Q5)? Can it promote
lead(II) bioabsorption? Is there an association between lead(II) and
fluoride?
HF is a weak acid with a pKa of 3.17 [18]. Therefore, NaF does
affect pH via (19). However, its effect can generally be neglected since
the pH of drinking water is controlled by many different buffering
species.
F- + H20 ^ HF + OH-, Kb w 1(T10-5 (19)
While it may be calculated, the very small association [28,29] between
Na+ and F~ (20) would be insignificant, and can be safely ignored.
-------
612 E. T. URBANSKY AND M. R. SCHOCK
The magnitude of this stability constant is so small as to be negligible;
however it can still be calculated. On the other hand, other cations
present in reasonably high concentrations, most notably aluminum,
bind to fluoride much more strongly (21)-(32). Table II summarizes
these chemical equilibria and their stability constants.
There are many metal cations competing for the fluoride; therefore,
the free fluoride available to complex with the lead(II) ion is very
small. In addition, most, if not all, of the competing metal cations
are in greater abundance than lead(II) by orders of magnitude. Fur-
ther reducing the lead(II) are such ligands as hydroxide, chloride,
carbonate, bicarbonate, and sulfate, all of which compete with fluo-
ride for the lead(II) and are present in far greater concentrations.
Table III summarizes these equilibria and their stability constants.
For pH > 6, the free lead(II) concentration drops off dramatically
from hydroxo- and (bi)carbonato-complexation. That drinking water
contains a substantial fraction of fluoro-aluminum complexes (21)-
(26) rather than free fluoride was highlighted by Fitter as a concern
because free fluoride is more effective in protecting against tooth decay
[30]. We shall take these and other factors into account in speciating
the lead(II).
TABLE II Cumulative stability constants for formation of fluoro-complexes*
Fluoro-complexation Equations log (3
Na++F~^NaF(aq) (20) -0.24*
(21) 7.0*
(22) 12.7
) (23) 16.8
(24) 19.4*
(25) 20.6*
(26) 20.6*
(27) 5.2*
(28) 9.1*
) (29) H.9*
(30) 0.94*
(31) 1.82*
(32) 1.2*
(33) 3.18*
(34) 3.76*
•These stability constants are used for the construction of Figures 3,4 with the exception of Eqs.
(27)-(29) and (32).
'Values taken from Ref. [29].
'Values taken from Ref. [18].
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 613
TABLE III Lead(II) equilibria and constants*
Equilibrium
Pb2++H2O^PbOH +
Pb2 + + 2 H2O ^ Pb(OH)2(aq) + 2 H +
Pb2+ + 3 H2O ^ Pb(OH)J + 3 H+
Pb2+ + 4 H20 ^ Pb(OH)2- + 4 H+
2 pb2+ +H2O?=±Pb2OH3+ +H +
3 pb2+ + 4 H20 ^ Pb3 (OH)2* + 4 H+
Pb2++ CO3~ ^PbCO3(aq)
Pb2++2C02-^Pb(C03)^
Pb2i"+ H+ + CO2' ^ PbHCOj
Pb2"1" + SO2" ^ PbSO4(aq)
Pb2++2SOj-^Pb(S04)2-
Pb2 + + Cl" ^ PbCl +
Pb2++2Cr^PbCl2(aq)
Pb2++ 3 Cl" ^ PbClJ
Pb2+ + 4 Cl" ^ PbCl2"
Pb2+ +p~?=;pbF'1"
Pb2++2F-^PbF2(aq)
Pb2 + + H4SiO4(aq) + 4H+ + 6F~^
Pb-F-SiF5(aq) + 4H20
Equations
(35)
(36)
(37)
(38)
(39)
(40)
(41)
(42)
(43)
(44)
(45)
(46)
(47)
(48)
(49)
(50)
(51)
(52)
logp
-7.22
-16.91
-28.08
-39.72
-6.36
-23.86
7.10
10.33
12.59
2.73
3.50
1.6
1.8
1.7
1.4
2.06
3.42
32.18f
•Values derived from Tables IV-XVI in Ref. [31] at 25°C and zero ionic strength. These equilibria are
used in the construction of Figures 1 -4.
* Computed from combining the dissociation constant for the reaction Si(OH),, + 4 H+ + 6 F ^
SiF^- +4H2O,logK = 30.18 (from Ref. [29]) with Eq. (18). We believe this value to be an inten-
tional overestimate by a factor of at least 10-20 over the likely value of the true stability constant,
which has not been measured.
One might logically inquire whether PbF2 can precipitate under
drinking water or physiological conditions. Combining Eqs. (51) and
(53) gives the solubility of the aqueous uncharged difluorolead(II) co-
ordination complex: tPbF2(aq)]max=9.5 x 10~5M.
PbF2(s) ^ Pb2+ + 2F-, Ksp - KT7-44 (Ref. [18]) (53)
Recall that a 1.0 ppm fluoride solution contains only 5.3 x 10~5M
fluoride ion. Even if the equilibrium constants have some error in them,
the competing equilibria ensure that plumbous fluoride does not pre-
cipitate. Aluminum, iron(III), calcium, magnesium, and copper(II)
compete with lead(II) for fluoride. Meanwhile hydroxide, carbonate,
phosphate, and sulfate compete with fluoride for lead(II). The net
result of these simultaneous competitions is that PbF2 cannot
precipitate as a solid. Even with 90th percentile lead(II) levels of
~210ngL~1(~l|iM), no plumbous fluoride would precipitate. The
formation of soluble fluoro-complexes of Pb(II) is governed solely by
-------
614 E. T. URBANSKY AND M. R. SCHOCK
the stability equilibria of (50)-(51). Therefore, no simple stoichio-
metric ratio exists among the concentrations of lead(II), fluoride, and
the fluoro-complexes.
Pretending that there are no competing metal cations and no com-
peting coordinating ligands, the total Pb(II) concentration is given
by (54):
[Pbn]T = [Pb2+] + [PbF+] + [PbF2] = [Pb2+](l + ft [F-] + /?2[F-]2)
(54)
where /?i and /32 come from (50) and (51), respectively. In l.Opprafree
fluoride (5.3 x 10~5M) solution, the fractional speciation is as follows:
/pb2+ = 99.904%,/PbF+ = 0.096% and/PbF2(aq) = 0.000099%. We draw
attention to the fact that, in fluoridated tap water, the number
5.3 x 10~5M really refers to the total fluoride, which is expressible as
(55):
[F~]T- = [F-] + [PbF+] + 2 [PbF2] = 5.3 x 10~5 M (55)
Nevertheless, because [F~]T« [F~] (less than 0.1% difference), there is
no point in distinguishing between these two concentrations. However,
in a real water, the competition of other metal cations for fluoride
would substantially reduce free fluoride. We will shortly demonstrate
this fact and that [Pbn]T»[Pb2+] in any real potable water.
We emphasize that - for practical purposes - there is no endless
supply of lead(0/II) awaiting complexation by fluoride (or anything
else for that matter). In other words, one cannot argue that the reac-
tion between lead(II) and complexing anions is driven forwards by a
readily available reservoir of lead that provides as much as the
ligands can react with. The dissolution of lead is confounded by three
possible barriers: (1) the kinetics of the oxidation of lead(O) are in-
sufficiently facile for equilibrium to be approached, (2) the reaction is
limited by the total oxidizing capacity of the water, and (3) the pre-
sence of passivating films such as Pb3(CO3)2(OH)2(s), Pb3(PO4)2(s),
PbCO3(s), or Pb5(PO4)3(OH)(s) controls solubility [29].
If ligand availability alone were the determining factor, chloride
itself would be far more important than fluoride. A chloride con-
centration of 50ppm (= 1.4mM = 1400uM) is about 26 times the
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 615
fluoride concentration. Between the chloride concentration and the
chloro-complex formation constants in (46)-(49), concentrations of
lead(II) complexes with chloride exceed those of fluoride in drinking
water.
We feel strongly that there is no scientific justification for asserting
that water fluoridation chemicals can have any quantitable impact
on human health via lead(II) exposure. There is no sound chemical
explanation for invoking any interactions between plumbous species
and hexafluorosilicic acid, hexafluorosilicate, or fluoride.
4. METHODOLOGICAL CONSIDERATIONS
IN STUDYING THE IMPACT OF FLUORIDATION
ONLEAD(H)
4.1. Measurement and Significance of Lead(II)
Concentrations in Tap Water
The total lead in a first draw sample tells more about the quality and
construction of a building plumbing system than it does about the
quality of the post-treatment water. Comprehensive water sampling
for epidemiological and other health effects studies for lead(II) is
logistically complicated and expensive; therefore, it is very tempting
to try to use available regulatory tap water monitoring data for this
purpose. The temptation must be resisted, however as the monitoring
program specified in the United States drinking water regulations is
both statistically and physically invalid for this purpose. The vast
preponderance of the lead(II) in nearly all tap waters originates from
the plumbing materials located between the water distribution mains
and the end of the faucet used by the consumer. Individuals consume
water under innumerable combinations of volumes of water, interior
plumbing system configurations and ages, and lengths of stagnation
of the water in the plumbing between uses. Data reported from many
tap water sampling experiences throughout the US and Europe indi-
cate tap water lead levels tended to follow a log-normal distribution,
and both within-site and between-site variability tended to be large
relative to the lead(II) concentrations. Keeping this in mind, the
American standard for lead in drinking water was crafted to focus
-------
616 E. T. URBANSKY AND M. R, SCHOCK
on the lowering of lead(II) levels by central water treatment for the
plumbing configurations most likely to represent nearly the worst
cases for the most vulnerable humans, i.e., infants, children, pregnant
women (Q6). After a cursory examination of the requirements for a
statistically valid sampling program accounting for needed levels of
predictive confidence across all sources of variability observed, one
realizes that it would take literally hundreds or thousands of samples
at great frequency from cities of all sizes to try to adequately charac-
terize tap water lead levels for even a single uniformly applied national
sampling protocol.
Obviously, the water chemistry at the point the distributed finished
water enters the domestic or commercial building plumbing system
plays a very significant role in affecting lead release into the water,
but many other physical factors also operate [31-40]. The water at
this point may have undergone chemical changes during its passage
through the distribution system from the treatment plant or well, and
changes in treatment or changes in water sources may also cause the
chemical characteristics of the water to change periodically, especial-
ly in such important aspects as pH, and concentrations of alkalinity,
natural organic matter, oxidant levels, and a variety of potentially
aggressive anions. Even the season may influence lead levels in compli-
cated ways, by changes in ground temperature, or temperatures in
buildings where pipes run through basements, unheated crawl spaces,
concrete slabs, or nearby heating or air conditioning ducts. A single
snapshot sampling event cannot capture this.
The drinking water literature is full of papers that show how diffi-
cult it is to correlate lead levels with any one or even a mix of several
water quality parameters, and a complete discussion of the matter is
beyond the scope of this article. There may be countless other physi-
cal or chemical quantities that may be statistically correlated with
lcad(II) levels but nonetheless be totally unrelated mechanistically.
Clearly, aggregate measures such as a small number of first-draw or
fully-flushed water samples taken infrequently from an intentionally
biased relatively small pool of sampling sites throughout a water sys-
tem cannot quantitatively and precisely predict the exposure of any
individuals to lead from drinking water. To accurately determine
lead(II) intake, sampling schemes using diverters or proportional sam-
pling devices that capture a representative fraction of the water
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 617
actually drawn at the faucet by the consumer seem to be the only fea-
sible approach [31]. Informed speculation suggests that other water
and plumbing characteristics, which were not measured in the studies
of Ref. [2] could correlate with lead(II) levels with equal or greater
statistical significance than those relationships that were put forth.
4.2. Measurement and Significance of Lead(II)
Concentrations in Blood
How is lead exposure assayed? How should it be measured? These
are complicated questions, and whole volumes can be dedicated just to
Q7. The discussion here will be confined to certain techniques and
their weaknesses.
We suggest that the apparent link between hexafluorosilicate usage
and children's blood lead(II) concentrations is the result of an unfor-
tunate coincidence. The use of blood lead(II) levels is problematic
for a number of reasons. Masters and Coplan [2] report these results
(reproduced here in Tab. IV); however, they do not report the uncer-
tainties associated with measurement or the variation of the sample
space.
A recent paper by Skerfving et al., reports that blood lead levels
are not linked in a linear fashion with exposure or risk and that bio-
logical monitoring is complicated by these factors [41]. Consequently,
it is unclear whether a difference of 7.1 ugLT1 is truly significant.
Blood lead(II) levels were reported in a previous study conducted
by the Dartmouth-Hitchcock Medical Center [42,43]. We have some
concerns regarding the secondary use of these data. Samples were col-
lected by a wide variety of personnel with undocumented training and
experience in the matter of collecting samples for trace metal analysis
and collection conditions were unspecified. Without a protocol that
TABLE IV Blood lead(II) concentrations and fluoridating agents*
Treatment [pl>U]nood' V-8 *•"'
H2SiF6 ~ 2?!
Na2SiF6 26-6
NaF 20.7
20.2
none
•Taken from Ref. [2].
-------
618 E. T. URBANSKY AND M. R. SCHOCK
includes the collection of spiked samples or recovery check stand-
ards, it is impossible to know if an individual collection site may be
responsible for a determinate error in all of the samples in a parti-
cular region. Trace metals are extremely susceptible to surface ad-
sorption, and lead(II) is notorious for plating out (reduction to Pb°)
under nonacidic conditions. Because the procedure was unspecified,
sample preservation and integrity are unknown.
The subsequent determination of lead(II) in blood was carried out
by the semi-quantitative method of Piomelli [44]. In this method, a mod-
erately linear relationship is found to exist between free erythro-
cyte porphyrins and total lead(II) concentration (as determined by
atomic absorption photometry). A plot of FEP concentration against
blood total lead(II) concentration in the original reference shows
so much scatter that the uncertainty in any blood lead concentration
must be at least ±10 ugL"1, and Piomelli gives an estimate of
±15ugL~1 at a concentration of SO^gL"1. Accordingly, we feel that
there is no statistical difference in the blood total lead values for
any set of conditions. Piomelli does not list any interferences other
than iron-deficiency anemia and erythropoietic protoporphyria;
however, we suspect that other elevated levels of other metals may
also bring about elevated levels of free porphyrins in the blood.
In defense of the Sargent et al., studies, we want to emphasize that
the original authors never intended for their lead(II) screening data
to be used in the way that Masters and Coplan have used them. The
original studies did not attempt to construct the sort of hard and fast
links of the Masters and Coplan papers, and they made no mention of
tap water whatsoever. In fact, senior authors Sargent and Bailey have
specifically stated that they do not agree with the Masters and Coplan
thesis (personal communication). To be sufficiently rigorous for a
study of drinking water and blood, more accurate, precise, and rugged
techniques than the FEP test are required, for instance, inductively-
coupled plasma or atomic absorption spectrophotometry.
4.3. Difficulties in Correlating Lead(II) Concentrations
in Water and Blood
The authors did not give the total lead concentrations in the first draw
water samples, so we cannot directly compare blood lead levels with
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 619
water lead levels [2]. They did give blood lead levels divided up by
those water systems where first draw samples were divided by a cut-off
of 15 ugL-1 of lead(II); see Table V.4
Because the sodium hexafluorosilicate data are based on one sys-
tem with [PbII]water> 15\igL~1, we discount it as an outlier without
further consideration. Without some estimation of the uncertainties
of [Pbn]blood, we cannot be assured that 23ngL~1 is distinct from
33 ugL"1. Also, we do not believe that it is possible to report 3 signi-
ficant digits in the blood lead level. We expect that the numbers
are probably good to about 10-15%. Masters and Coplan also failed
to include the possibility of naturally occurring fluoride and silicates
in the unfiuoridated water systems, which would be necessary to
substantiate their thesis, as the effects should be the same whether
the fluoride is an additive or native constituent. The authors them-
selves acknowledge the poor correlation between blood and water
lead(II) levels [2]:
Whereas a community's average uptake of lead by children is
weakly associated with the so-called "90th percentile first draw"
lead levels of lead in public water supplies (adjusted r2 = 0.02),
the fluoridation agents used in water treatment have a major
effect on lead levels in children's blood.
There appears to have been no effort to correlate an individual
building or house with blood lead levels. From the details in the paper,
there is no indication that there is any connection between sampled
taps and sampled persons. In other words, nothing indicates that a
person living in the sampled house had his blood drawn. Instead, the
TABLE V [Pb"]b,ood (ugL~') for fluoridation processes*
{Pb ]waKr
< ISugL-'
n —
> ISugL-1
n =
none
19.7
86
21.8
29
NaF
21.1
31
19
8
Na2SiF6
23.7
6
43.8
1
H2SiF6
23.1
26
32.7
25
•Taken from Ref. [2].
4 It is unclear from where these numbers originated. Reference [2] mentions averaged
90th percentile values. We take this to mean that 90th percentile values from two or
more rounds of regulatory testing were averaged.
-------
620 E. T. URBANSKY AND M. R. SCHOCK
authors rely on quartile divisions of both water lead levels and blood
lead levels. For there to be a correlation between the lead(II) levels
in blood and water, there must be a link between the samples. It is
possible that the highest blood levels of lead(II) are closely linked
to other exposures, such as paint, soil, or mine run-off. It would ap-
pear that the authors did not believe their own correlation coefficient
of 0.02. When these methodological problems are coupled with the
failure to account for fundamental chemical interactions, the rela-
tionships posed between lead(II) and water fluoridation become
unsubstantiatable.
5. PURITY AND INTEGRITY OF WATER
FLUORIDATING AGENTS
Water treatment chemicals are subject to National Sanitation
Foundation specifications, which require that additives contain a maxi-
mum allowable level (MAL) less than or equal to 10% of the maximum
contaminant level (MCL) for any regulated contaminant in the na-
tional primary drinking water standards [45]. This third-party stand-
ard sets the limits for quality control (Q8). Nevertheless, one might
ask if there is a possibility that the materials purchased by utilities
could be contaminated in some way. Let us consider how these com-
pounds are produced and handled.
Most hexafluorosilicate and fluorosilicic acid are derived from the
processing of phosphate rock by the fertilizer industry [46]. In this
process, apatite and fiuorapatite (which can be thought of as a blend
of fluorite and apatite for this purpose) are decomposed with sul-
furic acid:
CaF2-Ca3(P04)2 + 10H2SO4 + 20H2O -» 2HF + 6H3PO4
-t-10CaSO4-2H2O (56)
Ca5F(P04)3 + 5 H2S04 + 10 H2O -» HF + 3 H3PO4
+ 5CaSO4-2H2O (57)
After the mined rock is ground, it is mixed with dilute phosphoric
acid. The wet crushed rock is moved to a reactor, treated with sulfuric
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 621
acid, and heated to 75-80°C for several hours. The phosphate and
fluoride minerals make up only a portion of the rock. Part of the re-
mainder is silica, which reacts with the hydrofluoric acid:
Si02(s) +4HF(aq) -* SiF4(g) + 2H2O(1) (58)
Both the silicon tetrafluoride (tetrafluorosilane) and hydrofluoric acid
are swept away by an air stream into an absorption tower where hexa-
fluorosilicic acid is generated.
SiF4 + 2HF-»H2SiF6 ' (59)
The resulting material is stored in containers of high density poly-
ethylene or other unreactive materials. Because the HF and SiF4 are
removed as gases, there is little chance of lead contamination from
the crushed rock. However, 23% w/w hexafluorosilicic acid is a strong
acid and quite corrosive; therefore, any lead would have to be the
result of improper storage or handling. While it is possible for this to
be a source of lead in drinking water, there is no evidence to suggest
contamination. A service draw of water should contain the same con-
centration of lead(II) as a first draw if such contamination occurs,
and the authors themselves admit that service draws are character-
istically much lower in lead. Testing either the water at the plant or
the stock fluoridating agent itself would also be sufficient to rule out
this possible route of exposure.
6. EPIDEMIOLOGY OF LEAD EXPOSURE
AND ABSORPTION
Exposure to lead can occur in many ways (Q9). The varying contri-
butions of these exposure routes have been the subject of multiple
studies. The problem is further complicated by incomplete under-
standing of subacute toxicity and dose-response [47]. Much of the
exposure to lead occurs through dust, air-borne particulates, soil,
paint, ceramic glazes, and sundry other sources, including drinking
water [48-51]. One of the special concerns for drinking water is that
the lead(II) appears to be far more bioavailable [48]. This is probably
because aqueous lead(II) is far more likely to pass through mucous
membranes than insoluble plumbous minerals. However, there is
-------
622 E. T. URBANSKY AND M. R. SCHOCK
some evidence to suggest that even insoluble minerals can release
lead(II) when ingested under the right conditions [52]. A number of
studies5 have concentrated on other factors affecting bioavailability
and bioabsorption, including other nutrients, alcohol, cigarettes, water
hardness, plumbing, and lifestyle [53-59]. The main conclusion that
"can be drawn from these studies is that the biological availability,
absorption, and accumulation of lead and its compounds depend on
a wide variety of factors, making this a very complicated, puzzle to
solve [60, 61].
7. FRACTIONAL SPECIATION MODELING
While many of the conclusions reached thus far have been based on
calculations, the complicated interplay of multiple species, which was
alluded to in the introduction, has heretofore been avoided. It is
now time to rigorously account for all of these interactions in a quanti-
tative fashion (Q10).
We have taken into account equilibria of lead(II), aluminum,
calcium, and other metals for such ligands as carbonate, chloride,
hydroxide, sulfate, and, of course, fluoride. .These were given earlier in
Tables II and III. Other necessary equilibria and their constants that
we have used for this modeling exercise are shown in Table VI. From
the following graphs, it is immediately clear that hexafluorosilicate
and fluoride complexes play no role in the chemistry of lead(II) 'in a
drinking water matrix. They are of inconsequential concentration
under the conditions encountered in treated potable water.
To test different hypotheses about the impacts of fluoride ligands
on lead solubility, several solutions were modeled using the computer
program MINEQL + [62]. The effect of various background ions
such as CX>3~, HCO^ and PO^" and water quality parameters such as
pH have been extensively investigated and reported in the water treat-
ment literature [29,31,35,59-68]. Free lead(II) ion, Pb2+, is a very
minor fraction of the soluble lead in most drinking water systems;
^he references cited here include a representative sampling over the last two decades
of the kinds of work that have been done. These references are not intended to comprise
a complete listing or review of the studies in this area.
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 623
TABLE VI Other equilibria used to calculate the fractional distribution of aqueous
species*
Equilibrium
C02(g)-C02(aq)
CO2(aq) + H2O ^ HCOJ + H^
HCOJ ^ CO2- + H-
Na" + HCO3- ** NaHC03(aq)
Na^ + COf- ^ NaCOj
Ca2+ +HCOJ r=i CaHCO^
Ca2" 4- CO2- ^ CaC03(aq)
Mg2- + HCOJ ^ MgHCOj
Mg2- + CO2- ^ MgC03(aq)
Si(OH)4(aq) ^ SiO^H)^ 4- H"
Si(OH)4(aq) ^ SiO2(OH)J- + 2H+
HSOJ ^ SOj- + H+
Ca2^ + SO2- ^ CaS04(aq)
Mg2- + SOj- ^ MgSO4 (aq)
A13+ + SOl~ ** A1SO4
A13" + 2SOJ-^A1(S04)2
A13+ + HSO4 ^ AlHSO^"
H2O^H- + OH-
Na^ + H2O ^ NaOH(aq) + IT
Ca2" + ftO ^ CaOHT + H"
Mg2^ + H2O ^ MgOH" + H-
AP* + H2O ^ AlOH2^ + H-
A13+ + 2 H2O ^ A1(OH)2 + 2 H*
Al3" + 3 H20 ^ Al(OH)3(aq) + 3 H~
Al3^ + 4 H2O ^ A1(OH)4 + 4 H-
Equations
(60)
(61)
(62)
(63)
(64)
(65)
(66)
(67)
(68)
(69)
(70)
(71)
(72)
(73)
(74)
(75)
(76)
(77)
(78)
(79)
(80)
(81)
(82)
(83)
(84)
Iog0
-1.468
-6.352
-10.329
-0.25
1.27
1.106
3.224
1.07
2.98
-9.83
-23.0
-1.988
2.30
2.37
3.02
4.92
0.46
-14.00
-14.18
-12.78
-11.44
-5.00
-10.1
-16.9
-22.7
* Values taken from Ref. [29]. These equilibria are used in the construction of Figures 3,4. Equations
(59), (60), (67), (68) and (75) are used for Figure 5.
therefore, sophisticated methods must be used to determine the overall
distribution of lead species in the water.
Calculations were performed for the following hypothetical water
solutions, as a means to test some plausible limits on when fluoride
or fluorosilicate complexes might be of consequence with respect to
solubility. Conditions are summarized in Tables VII and VIII. The
lead(II) concentration used for Figures 1 and 2 is 1000 times the 90th
percentile action level for public water supplies. For all modeling, tem-
perature was set to 25°C, and an ionic strength of 0.005 M was
assumed. This is characteristic of many New England waters and
close enough that small differences will not significantly affect lead
solubility and speciation projections. Test calculations showed this was
very close to the actual ionic strength of the hypothetical waters for
which other major constituents were included.
-------
624 E. T. URBANSKY AND M. R. SCHOCK
TABLE VII Water quality parameters for speciation modeling*
Species
[SiOJ-r1
rp— j 1
[COJT°
Concentration, mgL '
Figures 1, 2
Figures 3, 4
Figures 1, 3, 4
Figure 2
5.0
15
0.015
1.0
5.0
50
Concentration, molL '
8.3 x 10~5
7.2 x 10~s
7.2 x 10~8
5.3 x 10 5
4.2 x 10~4
4.2 x 10~3
* These concentrations used for Figures 1-4 except as noted.
'[SiOJr = total si!icon(IV) concentration, expressed as silicon dioxide.
J[Pb2'*"]T=total lead(H) concentration, all species.
1[F-]T=total fluoride concentrationp"] + [HF] + Zn[MF^~">}.
°[CO2]T = [CO2(aq)] + [H2CO3] + [HCOJ] + [CO^] (dissolved inorganic carbon). Mass-based con-
centration is expressed as C not CO2).
TABLE VIII Concentrations of background ions used in speciation modeling*
Ion Concentration, mgZ."' Concentration, molL~*
Ca2+ Io 1.2 x 10-"
Mg2+ 2.0 8.2x10-'
Na+ 10.0 4.4 x 1Q-"
A13+ 0.20 7.4 xlO"6
Cr . 10.0 2.8 x 1Q-"
SO2.- 5.0 5.2 x 10~5
•These values used for Figures 3, 4.
Figures 1, 2 show the fractional distribution of species in water that
contains only dissolved lead(II), inorganic silicates, fluoride, and carbo-
nates, where we find that the principal lead(II) species are the aquo
ion, hydroxo-complexes, and carbonate-complexes. We note that
there is less than one molecule of PbSiF6 per liter of water even
though there is no competition for fluoride by aluminum or other
metal cations and competition for lead(II) by only the predominant
complexing species (but not minor ones). Figure 1 shows that
Pb2+(aq) is the dominant species for pH < 6.8, at which point the
carbonatolead(II) complex begins to dominate. By pH«7.2, the
hydroxolead(II) ion also exceeds the free lead(II). As pH increases to
~ 8.4, only ~ 1 % of the total lead is the free aquated ion. In Figures 1
and 2, we see that the mono- and difluorolead(II) complexes always
account for less than 1% of the total lead(II). Note that the species
PbSiF6 is present at such low concentrations that we would expect
to find only one. molecule of this complex in 1000 liters of tap water
at pH 6. Note the broken ordinate.
-------
FLUORIDATED POTABLE WATER AND LEAD(II)
625
-4.0
1 %
Pbz*
Pb(OH)2AQ
PbOH*
Pb2OH3*
PblOH)3-
o Pb(OH),.2'
-a _ Pb3(OH)42+
_e_ PbHC03 +
_«._ PbC03 AQ
_*.. Pb(C03)22-
—«— PbF*
__ _ PbF, AQ
FIGURE 1 Species distribution for lead in system containing only carbonate (5mgC/
L), silicate (5mg SiO2/L), and fluoride (1 mg/L) at 25°C, 7=0.005. PbSiF°6 complex was
included in the model, with assumption of log /3-2.
In Figure 3, we show the minor species, including the sulfato-,
fluoro-, and chloro-complexes under the influences of other common
background ions (Tab. VIII). The carbonate-complexes of lead(II) are
much stronger than the halo-complexes - as reflected by their stability
constants, which are 5-8 orders of magnitude higher than those of
the comparable halide complexes. As a consequence, we did not con-
sider higher [CO2]T concentrations where the Sucre-complexes become
even less significant.
The concentration of aluminum in Table VIII represents a mod-
erate to somewhat high residual carried over from agglutination
-------
626
E. T. URBANSKY AND M. R. SCHOCK
1 %
Pb" '
Pb(OH)2AQ
PbOH*
Pb2OH3*
Pb(OHJ3-
_e_ Pb(OH)t2-
_a - Pb3(OH)42H-
_c_ PbHCO/
_e,_ PbC03 AQ
__«,.. Pb(CO,)22-
-A- PbF*
_, _ PbF2 AQ
10
FIGURE 2 Species distribution for lead in system containing only carbonate (50 mg C/
L), silicate (5mg SiO2/L), and fluoride (1 mg/L) at 25°C, 7=0.005. PbSiFjj complex was
included in the model, with assumption of log 0 = 2.
(coagulation) with potassium aluminum sulfate (alum), as commonly
occurs with surface water treatment plants adjusted for corrosion
control and the Lead and Copper Rule. Pertinent results with respect
to lead speciation and the impact of fluoride and SiFg~ complex
formation on it, are shown in Figures 1-3.
The insignificance of any SiFg~ can be logically determined another
way. Even if the formation constant for a hypothetical PbSiF6(aq)
complex were ten times higher than the strongest complex found
by Haque and Cyr [26], it would have a similar stability to PbF+.
Assuming all of the fluoride present in drinking water were SiFg~,
-------
FLUORIDATED POTABLE WATER AND LEAD(II)
627
-4.0
1 %
10
FIGURE 3 Minor lead species distribution for lead in hypothetical "New England"
water described in Table VII. Computations were done for 25°C, 7=0.005. PbSiFg
complex was included in the model, assuming of log 13=2.
Figures 1-3 show that it would still be approximately 3 (pH6) to
more than 6 (pH 10) orders of magnitude lower than the soluble lead
level, which is governed by the concentrations of other Lewis bases.
Because complexation with carbonate and bicarbonate dominates
aqueous lead speciation at drinking water pH, the increased [CC>2]T
level of 50 mg L~' (Fig. 2) makes contribution of the fluoro-complexes
to [Pbn]T even less significant. The bar graph in Figure 4 clearly illus-
trates how free lead(II), hydroxo-, and (bi)carbonato-complexes domi-
nate the speciation of lead(II) at all drinking water pH values while
fluoro-complexes are always in the minority.
-------
628
E. T. URBANSKY AND M. R. SCHOCK
100
10
o
u_
0.1
8
PH
10
FIGURE 4 Illustration of fractions of soluble lead bound to different ligand groups for
hypothetical representative "New England" water, assuming 15ng/LPb and back-
ground ion concentrations given in Table VII of the text.
7.1. Buffer Intensity (Buffer Capacity)
We have previously stated that naturally occurring buffers have a
significant impact on drinking water chemistry. At this point, we will
quantitatively illustrate the magnitude of this impact. Figure 5 shows
the buffer intensity (capacity) 38 as a function of pH.6 The buffer
6Buffer intensity is usually represented by the symbol 0, which we find to be an un-
fortunate coincidence as it leads to confusion between this quantity and cumulative
stability constants, for which ft is often used simultaneously. As a result, we have broken
with convention and used the symbol 38 to stand for buffer intensity.
-------
FLUORIDATED POTABLE WATER AND LEAD(II)
10'3
629
4.00 5.00 6.00 7.00 8.00 9.00
' pH
FIGURE 5 Buffer intensity as a function of pH for an aqueous solution containing
[CO;>]T=42nM and [SiO2h- = 83uM with 7=0.005 M, T=25°C.
intensity is a quantitative description of a solution's resistance to
changes in pH upon addition of acid or base and is defined as (85):
B = -dCa/d(pH) - dCb/d(pH)
(85)
where Ca is the formal concentration (formality) of added acid and Cb
is the formality of added base. Because we are discussing infinitesimal
(differential) quantities of added acid or base, the effects are equal and
opposite for added acid versus added base. Note that the derivative
described in (83) is the slope of a curve of Cb versus pH, which is the
inverse of a "titratipn curve" where base molecules are added directly
to a solution of acid so that the liter is zero and there is no change in
volume, only changes in concentration. In terms of a real system, one
can imagine dropping NaOH pellets or bubbling HC1 into a pond.
Alternately, a large body of water into which small volumes of acidic
or alkaline solution are added experiences essentially no change in
volume and can be represented in this fashion.
Virtually all potable waters contain some dissolved inorganic
carbon, represented here as [CO2]-r; therefore, the buffer intensity will
-------
630 E. T. URBANSKY AND M. R. SCHOCK ;
be controlled by the simultaneous conjugate acid-base equilibria
of the carbon dioxide-carbonic acid-bicarbonate-carbonate system.
Although not conceptually difficult to understand, the derivation of a
quantitative definition of the buffer intensity @ for a given system can
be cumbersome due to the lengthy algebraic expressions and differ-
ential calculus. For "this reason, we include the derivation for this
system. Let us consider the situation where the system resists a change
in pH upon addition of an infinitessimal (differential) amount of
alkali. The same thought processes would apply for the introduction
of a small amount of acid.
We begin with the charge balance (86) and mass balance (87)
expressions:
[H+] + [M+] = [OH-] + [HCO^] + 2 [COM (86)
[C02]T = [COf-] + [HCOJ] + [C02] (87)
where M+ represents the cation of some monohydroxide base (MOH).
Because H2CO3 can be thought of as CO2-H2O and water has activity
equal to unity, there is no need to include carbonic acid separately in
this process, K&\ includes both CO2 and CO2-H2O. As we are adding
portions of base directly, there is assumed to be no change in volume.
For this exercise, it is convenient to express the carbon dioxide acid
equilibria in terms of the cumulative protonation constants rather
than the stepwise acid dissociation constants given in Table VI.7
= [HC03-]/([H+][COM) (88)
= [C02]/([H-f [CO2-]) (89)
Making use of (88) and (89), we obtain the following expression for
the total carbon dioxide concentration (dissolved inorganic carbon):
[C02]T = [COf-] (1 + A [H+] + /?2[H+]2) (90)
'Because our equilibrium constants are adjusted for ionic strength and temperature,
we have not included correction factors or activity coefficients; however, we emphasize
that one cannot simply substitute any values into these equations without appropriate
consideration of relevant conditions.
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 63 1
Since [M+] is equal to the formal concentration of the added base,
we replace [M+] with Cb. We express the bicarbonate and carbonate
concentrations in (86) in terms of [CO2]T, using (88)-(90). Lastly,
we substitute £w/[H+] for [OH"] and rearrange, solving for Cb:
Cb = KW/[H+] - [H+] + (A [H+] + 2)[C02]T/(1 + A[H+j + /32[H+]2)
(91)
Differentiating with respect to the hydrogen ion concentration and
combining like terms yields (92):
dCb -#w
-
d[H+] [H+]2
Since d[H+]/d(pH) = ~p+](ln 10), we make use of the chain rule for
composite functions to obtain the buffer intensity SS, which is given
by (93) after simplification:
= dCb/d(pH) - dCb/d[H+] • d[H+]/d(PH)
\1
)\
(93)
where SS has units of M (pH unit)"1 when all concentrations are ex-
pressed in molarities. Upon inspection of (93), it can be seen that the
buffer intensity can readily be divided into contributions from [OH~],
[H+] and [COJ-r. As Butler [68] notes, the cubic terms are negligible
under most conditions and are often dropped from these expressions.
Contributions from additional buffers, £%h may be accounted for by
additional terms. For example, ^s;o2 is given by (94):
= (In 10)[Si02]T(£1[H+]3 + 4*T2[H+]2 + K1K2[H+])
(94)
where [SiO2]T - [Si(OH)4] + [SiO(OH)^] + [SiO2(OH)n and K,
and K2 take their forms from Eqs. (69) and (70), respectively. Note
that the multiplier of [buffer]T is unitless as it represents the active
-------
632 E. T. URBANSKY AND M. R. SCHOCK
fraction of the total buffer. The inclusion of other components, in-
cluding other weak Br0nsted-Lowry bases (silicates, phosphates, etc.),
or Lewis acidic (hydrolytic) metal cations (Mg2+,Al3+, etc.), serves
to add terms and complicates the mathematics. However, the same
fundamental principles and logic used here apply equally to such
derivations.
Figure 5 shows that much of the buffer intensity is derived from the
carbon dioxide-bicarbonate-carbonate system.8 For the example New
England water used in the calculations, the minimum buffer inten-
sity contributed by the [COJy-and water is ^ = 0.25mM (pH unit)"1.
The contribution of acid from undissociated SiFg~ can again be
proved negligible from the following extreme example. Even if 10%
of the total fluoride input were as SiFg~ at pH 7, the acid input would
be ACa = 5xl(T6M. /«0.005 M, T=25°C. Thus, we can
compute the change in pH directly using the buffer intensity calculat-
ed above: A(pH) w -ACJ® = -0.0050 mM/[0.25 mM (pH unit)"1^
-0.020 pH unit. This value is within the limits of a linear approxima-
tion of buffer capacity. Such a small effect on pH is undetectable and
inconsequential with respect to other sources of variability in factors
affecting lead release from plumbing materials. The concepts of chemi-
cal equilibria are well-established and measured equilibrium constants
are sufficiently accurate and precise to show that fluoride and fluoro-
silicate essentially do not affect the distribution of lead(II) species
under potable water conditions.
8. CONCLUSION
Recent reports on the possible effects of water fluoridating agents,
such as hexafluorosilicic acid, sodium hexafluorosilicate, and sodium
fluoride are inconsistent with accepted scientific knowledge, and the
authors fail to identify or account for these inconsistencies. Many of
the chemical assumptions are scientifically unjustified, and alternate
explanations (such as multiple routes of Pb11 exposure) have not been
satisfactorily addressed. At present, there is no evidence to suggest
8The atomic mass of carbon is 12.011 gmol ', so concentrations of [CO2]x expressed
as mgCL"1 can be converted to units of molL"1 by dividing by 12,011.
-------
FLUORIDATED POTABLE WATER AND LEAD(H) 633
that the common practice of fluoridating drinking water has any unto-
ward health impacts via effects on lead(II) when done properly under
established guidelines so as to maintain total water quality. Our
conclusion supports both EPA and PHS/CDC policies on water
fiuoridation.
LIST OF SYMBOLS
!% buffer capacity
buffer capacity of carbon dioxide
buffer capacity of silicon dioxide
J3 stability constant
/j. micro, bridging ligand
m molality
[X}r total concentration of X
/pbz- fraction of lead(II) ion
/PbF- fraction of lead(II) ion
/pbF2 (aq) fraction of lead(II) fluoride
A'sp solubility product
References
|lj R. D. Masters and M. Coplan, "Brain Biochemistry and the Violence Epidemic:
Toward a 'Win-Win' Strategy for Reducing Crime", Global Legal Policy, S. Nagel,
Ed., St. Martin's Press: New York (in press).
[2] R. D. Masters and M. Coplan, "Water Treatment with Silicofluorides and Lead
Toxicity", Intern. J. Environmental Studies 56, 435-449 (1999).
[3] Crime Times: Research Reviews and Information on Biological Causes of Criminal,
Violent, and Psychopathic Behavior 4(4), 1-2 (1998).
[4] US Department of Health and Human Services, US Public Health Service, Centers
for Disease Control and Prevention. Water Fiuoridation Census 1992. (July, 1993).
[5] US Department of Health and Human Services, Public Health Service, Centers for
Disease Control and Prevention, (a) Water Fiuoridation: A Manual for Engineers
and Technicians. (September, 1986); (b) Water Fiuoridation: A Manual for Water
Plant Operators. (August, 1993).
[6] C. F. Baes, Jr. and R. E. Mesmer, The Hydrolysis of Cations. Krieger Publishing
Co.: Malabar, Fla., 1976; Ch. 15.
[7] R. H. Busey and R. E. Mesmer, "lonization Equilibria of Silicic Acid and
Polysilicate Formation in Aqueous Sodium Chloride Solutions", Inorg. Chem.
16(10), 2444-2450 (1977).
[8] R. H. Busey, E. Schwartz and R. E. Mesmer, "Fluorosilicate Equilibria in Sodium
Chloride Solutions from 0 to 60°C", Inorg. Chem. 19(3), 758-761 (1980).
-------
634 E. T. URBANSKY AND M. R. SCHOCK
[9] L Ciavatta, M. lulianno and R. Porto, "Fluorosilicate Equilibria in Acid
Solutions", Polyhedron 7(18), 1773-1779 (1988).
[10] A. S. Korobitsyn, G. N. Bogachov and D. F. Rize, "Influence of Temperature on
the Rate of Decomposition of Crystalline Sodium Fluosilicate by Sodium Carbo-
nate Solutions", J. Appl. Chem. (USSR) 45(4), 736-740 (1972). Translated
from Zh. Prikl. Khim. 45(4), 722-726 (1972).
[11] R. E. Connick and R. E. Poulson, "19F Nuclear Magnetic Resonance of Various
Metal-Fluoride Complexes in Aqueous Solution", J. Phys. Chem. 63, 568-569
(1959).
[12] E. L. Mutterties and W. D. Phillips, "Structure and Exchange Processes in Some
Inorganic Fluorides by Nuclear Magnetic Resonance", J. Am. Chem. Soc. 81,
1084-1088(1959).
[13] A. G. Kucheryaev, V. A. Lebedev and I. M. Ovchinnikov, "Stability of the Silicon
Hexafluoroanion in Hydrofluoric Acid, from 19F NMR Data", /. Struct. Chem.
(USSR) 11(5), 858-859 (1970). Translated from Zh. Strukt. Khim. 11(5), 925-926
(1970).
[14] P. M. Borodin and N. K. Zao, "Equilibria in the Li2SiF6-HClO4-H2O System by
the 19F Nuclear Magnetic Resonance Method", Russ. J. Inorg. Chem. 16(12),
1720-1722(1971).
[15] P. M. Borodin and N. K. Zao, "Study of Mobile Equilibria in the HF-SiO2-H2O
System by the 19F NMR Method", Russ. J. Inorg. Chem. 17(7), 959-962 (1972).
[16] P. M. Borodin, N. K. Zao and N. S. Petrov, "Stability of the Hexafluorosilicate Ion
in Aqueous Solutions of Fluorosilicic Acid System, from I9F NMR Data",
J. Struct. Chem. (USSR) 14(4), 564-568 (1973). Translated from Zh. Strukt.
• Khim. 14(4), 605-610 (1973).
[17] Gmelin Handbook of Inorganic Chemistry, 8th edn. Silicon. Springer-Verlag:
New York, 1995; "H2SiF6", Supplement Vol. B8.
[18] R. M. Smith and A. E. Martell, Critical Stability Constants: Inorganic Complexes.
Plenum Press: New York, 1976; Vol. 4, pp. 96-103.
[19] N. T. Crosby, "Equilibria of Fluorosilicate Solutions with Special Reference to
the Fluoridation of Public Water Supplies", J. Appl. Chem. 19, 100-103 (1969).
See also references therein.
[20] V. N. Krylov, E. V. Komarov and M. F. Pushlenkov, "Stability of the SiF^ Ion in
Aqueous Solutions in the Presence of Ions of Various Valence", Sov. Radiochem.
13(3), 445-449 (1971). Translated from Radiokhim. 13(3), 430-434 (1971).
[21] K. K. Worthington and M. M. Haring, "Methods for Determining the Solubilties
of Some Fluosilicates", Ind. and Eng. Chem. 3(1), 7-9 (1931).
[22] L. J. Hudleston and H. Bassett, "Equilibria of Hydrofluosilicic Acid", J. Chem.
Soc. (London) 119, 403-416 (1921).
[23] A. G. Rees and L. J. Hudleston, "The Decomposition of the Fluorosilicate Ion in
Aqueous and in Aqueous Salt Solutions", J. Chem. Soc. (London) 288, 1334-1338
(1936).
[24] V. N. Plakhotnik, "Kinetics and Mechanism of the Hydrolysis of Some
Acidofluoro-Complexes in the Presence of Polarizing Cations", Kinet. Catal.
16(6), 1211-1214 (1975). Translated from Kinet. Katal. 16(6), 1395-1399 (1975).
[25] V. N. Plakhotnik, "The Catalytic Influence of Li"1" on the Alkaline Hydrolysis of
the Hexafluorosilicate Anion", Kinet. Catal. 14, 1417-1418 (1973).
[26] R. Haque and N. Cyr, "Fluorine Magnetic Resonance Studies of Metal-
Hexafluorosilicate Ion-Pairs in Solution", Faraday Soc. Trans. 66, 1848-1853
(1970).
[27] I. P. L. Coleman, J. A. Blair and M. E. Hilburn, "Effects of Dietary and Synthetic
Chelating Agents on the Intestinal Absorption of Lead", Intern. J. Environmental
Studies IS, 187-191 (1982).
[28] A. E. Martell and Robert M. Smith, Critical Stability Constants, First Supplement.
Plenum Press: New York, 1982; Vol. 5, pp. 415-417.
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 635
[29] D. K. Nordstrom, L. N. Plummer, D. Langmuir, E. Busenberg, H. M. May, B. F.
Joiws and D. L. Parkhurst, "Revised Chemical Equilibrium Data for Major Waterr
V,ifi«ral Reactions and their Limitations". In: Chemical Modeling of Aqueous
Svhtems II. ACS Symposium Series 416, R. L. Bassett and D. C. Melchior, Eds.,
American Chemical Society, Washington, DC, 1990; Ch. 31.
[30] P. Fitter, "Forms of Occurrence of Fluorine in Drinking Water", Water Res. 19(3),
281-284(1985).
[31] M. R. Schock, I. Wagner and R. J. Oliphant, "Corrosion and Solubility of Lead in
Drinking Water", In: Internal Corrosion of Water Distribution Systems, 2nd edn.
American Water Works Association Research Foundation: Denver, Colo., 1996;
Ch. 4.
[32] M. R. Schock, "Understanding Corrosion Control Strategies for Lead", /. Am.
Water Works Assn. 81, 88-100 (1989).
[33] N.-H. Lin, A. Torrents, A. P. Davis, M. Zeinali and F. A. Taylor, "Lead Corrosion
Control from Lead, Copper - Lead Solder, and Brass Coupons in Drinking Water
Employing Free and Combined Chlorine", J. Environ. Sci. Health A32(4), 865 - 884
(1997).
[34] D. A. Lytle and M. R. Schock, "Impact of Stagnation Time on the Dissolution of
Metal from Plumbing Materials", Proc. AWWA Annual Conf. (Atlanta, Ga., June
15-19, 1997).
[35] M. R. Schock and M. C. Gardels, "Plumbosolvency Reduction by High pH
and Low Carbonate-Solubility Relationships", J. Am. Water Works Assn. 75(2),
87-91 (1983).
[36] J. H. Colling, P. A. E. Whincup and C. R. Hayes, "The Measurement of
Plumbosolvency Propensity to Guide the Control of Lead in Tapwaters", J. Inst.
Water and Environ. Management (J. IWEM) 1(3), 263-269 (1987).
[37] A. D. Hulsman, "Paniculate Lead in Water Supplies", J. IWEM 4(1), 19-25
(1990).
[38] J. H. Colling, B. T. Croll and P. A. E. Whincup, "Plumbosolvency Effects and
Control in Hard Waters", J. IWEM 6(3), 259-268 (1992).
[39] R. Levin, M. R. Schock and A. Marcus, "Exposure to Lead in U.S. Drinking
Water", Twenty-third Annual Conference on Trace Substances in Environmental
Health (Cincinnati, Ohio, May 30-June 1,1989).
[40] M. R. Schock, "Causes of Temporal Variability of Lead in Domestic Plumbing
Systems", Environ. Monitoring Assessment 15, 59-82 (1990).
[41] S. Skerfving, L. Gerhardsson, A. Schiitz and U. Strombert, "Lead-Biological
Monitoring of Exposure and Effects", J. Trace Elems. Exptl. Med. 11(2-3),
289-301 (1998).
[42] J. D. Sargent, M. J. Brown, J. L. Freeman, A. Bailey, D. Goodman and D. H.
Freeman Jr., "Childhood Lead Poisoning in Massachusetts Communities: Its
Association with Sociodemographic and Housing Characteristics", Am. J. Public
Health 85(4), 528-534 (1995).
[43] A. J. Bailey, J. D. Sargent, D. C. Goodman, J. Freeman and M. J.
Brown, "Poisoned Landscapes: The Epidemiology of Environmental Lead Ex-
posure in Massachusetts Children 1990-1991", Soc. Sci. Med. 39(6), 757-766
(1994).
[44] S. Piomelli, "A Micromethod for Free Erythrocyte Porphyrins: The FEP Test",
J. Lab. Clm. Med. 81(6), 932-940 (1973).
[45] N. I. McClelland, D. A. Gregorka and B. D. Carlton, "The Drinking-Water
Additives Program", Environ. Sci. Technol. 23(1), 14-18 (1989).
[46] P. J. Chenier, Survey of Industrial Chemistry, 2nd rev. edn. Wiley-VCH: New York,
1992; pp. 50, 51.
[47] A. C. Todd, J. G. Wetmur, J. M. Moline, J. H. Godbold, S. M. Levin and P. J.
Landrigan, "Unraveling the Chronic Toxicity of Lead: An Essential Priority for
Environmental Health", Environ. Health Perspect. 104(1), 141-146 (1996).
-------
636 E. T. URBANSK.Y AND M. R. SCHOCK
[48] G. M. Raab, D. P. H. Laxen and M. Fulton, "Lead from Dust and Water as
Exposure Sources for Children", Environ. Geochem. & Health 9(3-4), 80-85
(1987).
[49] C. D. Strehlow and D. Barltrop, "Temporal Trends in Urban and Rural Blood
Lead Concentrations", Environ. Geochem. & Health 9(3-4), 74-79 (1987).
[50] J. L. Pirkle, R. B. Kaufmann, D. J. Brody, T. Hickman, E. W. Gunter and
D. C. Paschal, "Exposure of the U.S. Population to Lead, 1991-1994", Environ.
Health. Perspect. 106(11) (1998).
[51] M. B. Rabinowitz and D. C. Bellinger, "Soil Lead-Blood Lead Relationship
Among Boston Children", Bull. Environ. Contam. Toxicol. 41,791-797(1988).
[52] U. G. Gasser, W. J. Walker, R. A. Dahlgren, R. S. Borch and R. G. Burau, "Lead
Release from Smelter and Mine Waste Impacted Materials Under Simulated
Gastric Conditions and Relation to Speciation", Environ. Sci. Technol. 30(3),
761 -769 (1996).
[53] C. R. Angle and K. L. 'Stelmak, "Lead and Iron Deficiency", Trace Substances in
Environmental Health-IX. Proceedings of the University of Missouri's 9th Annual
Conference on Trace Substances in Environmental Health. D. D. Hemphill, Ed.
(Columbia, Mo., June 10-12, 1975); pp. 377-386.
[54] D. Barltrop and H. E. Khoo, "Nutritional Determinants of Lead Absorption",
Trace Substances in Environmental Health-IX. Proceedings of the University of
Missouri's 9th Annual Conference on Trace Substances in Environmental Health.
D. D. Hemphill, Ed. (Columbia, Mo., June 10-12, 1975); pp. 369-376.
[55] M. Rabinowitz, G. Wetherill and J. Kopple, "Absorption, Storage, and Excretion
of Lead by Normal Humans", Trace Substances in Environmental Health-IX.
Proceedings of the University of Missouri's 9th Annual Conference on Trace Sub-
stances in Environmental Health. D. D. Hemphill, Ed. (Columbia, Mo., June
10-12, 1975); pp. 361-368.
[56] S. J. Pocock, A. G. Shaper, M. Walker, C. J. Wale, B. Clayton, T. Delves, R. F.
Lacey, R. F. Packham and P. Powell, "Effects of Tap Water Lead, Water
Hardness, Alcohol, and Cigarettes on Blood Lead Concentrations", /. Epidemiol.
& Community Health 37, 1 -7 (1983).
[57] H. M. James, M. E. Hilburn and J. A. Blair, "Effects of Meals and Meal Times on
Uptake of Lead from the Gastrointestinal Tract in Humans", Human Toxicol. 4,
401-407(1985).
[58] M. Berode, V. Wietlisbach, M. Rickenbach and M. P. Guillemin, "Lifestyle
and Environmental Factors as Determinants of Blood Lead Levels in a Swiss
Population", Environ. Res. 55, 1 -17 (1991).
[59] J. Sherlock, G. Smart, G. I. Forbes, M. R. Moore, W. J. Patterson, W. N. Richards
and T. S. Wilson, "Assessment of Lead Intakes and Dose-Response for a Popu-
lation in Ayr Exposed to a Plumbosolvent Water Supply", Human Toxicol. 1,
115-122(1982).
[60] H. F. Thomas, P. C. Elwood, C. Toothill and M. Morton, "Blood and Water Lead
in a Hard Water Area", Lancet pp. 1047, 1048 (May 9, 1981).
[61] W. L. Marcus, "Lead Health Effects in Drinking Water", Toxicol. Ind. Health 2(4),
363-407(1986).
[62] W. D. Schecher and D. C. McAvoy, MINEQL + : A Chemical Equilibrium
Modeling System, Version 4.0 for Windows. Environmental Research Software:
Hallowell, Maine, 1998.
[63] M. R. Schock, "Response of Lead Solubility to Dissolved Carbonate in Drinking
Water", J. Am. Water Works Assn. 72(12), 695-704 (1980).
[64] M. R. Schock, "Response of Lead Solubility to Dissolved Carbonate in Drinking
Water", Erratum. J. Am. Water Works Assn. 73(3), 36 (news) (1981).
[65] M. R. Schock, "Internal Corrosion and Deposition Control", In: Water Quality
and Treatment: A Handbook of Community Water Supplies, 4th edn. American
Water Works Association, McGraw-Hill, Inc.: New York, 1990; Ch. 17.
-------
FLUORIDATED POTABLE WATER AND LEAD(II) 637
[66] M. R. Schock and J. A. Clement, "Control of Lead and Copper with Non-Zinc
Orthophosphate", /. New Engl. Water Works Assn. 112(1), 20-242 (1998).
[67] R. Gregory and P. J. Jackson, "Central Water Treatment to Reduce Lead
Solubility", Proc. AWWA Annual Conference (Dallas, Tex., 1984).
[68] J. N. Butler, Ionic Equilibrium. Wiley Interscience: New York, 1998; p. 392.
-------
------- |