Environ. Sci. Technol. 1988, 20, 160-165
Horowitz, A.; Elrick, K. U.S. Geol. Surv. Open-File Rep. 1985, 85-78, 1-18. Guy, H. "Techniques of Water Resources Investigations of the U.S. Geological Survey"; 1969, Book 5, Chapter C1, pp 1-58. Carpenter, J.;Bradford, W.; Grant, V. Esturaine Res. 1975,
1, 188.
Received for review April 9,1985. Revised manuscript received August 8,1985. Accepted September 28,1985. The use of brand names is for identification purposes only and does not imply an endorsement by the U S . Geological Survey.
Kinetics of Aluminum Fluoride Complexation in Acidic Waters Brlan J. Plankey and Howard H. Patterson"
Department of Chemistry, University of Maine, Orono, Maine
04469
Chrlstopher S. Cronan
Department of Botany and Plant Pathology, University of Maine, Orono, Maine
04469
where ki = 32.6 M-ls-l, kii = 3.61 X lo3 M-l s-l, kiii = 1.40 M-l s-l, and ki, = 1.1 X lo3 M-l s-l at 25 OC and ionic strength p = 0.10 M. Paths i and iv are independent of the [H+],path ii depends on l/[H+], and path iii depends on [H+]. The pH and temperature dependencies of the overall rate of reaction are discussed along with environmental implications for areas subjected to acidic deposition.
dominant inorganic aluminum species (3). The presence of fluoride can significantly increase total soluble aluminum ( 4 ) or decrease highly toxic monomeric aquo- and hydroxoaluminum (3),depending on the location and rate of mixing and reaction (5). Therefore, aluminum fluoride complexation can potentially have either a mitigating or aggravating effect on acid rain induced aluminum toxicity in natural waters. Unfortunately, there is a significant lack of kinetic data for this process. In this investigation, the kinetics of aluminum fluoride complexation were studied primarily in the pH range of 2.9-4.9. This is roughly the pH range of soil waters and first-order drainage waters in forested watersheds of northeastern North America and northern Europe that are impacted by acidic deposition. The objectives of the investigation were as follows: (i) to quantify the rate of aluminum fluoride complexation as a function of pH, temperature, and concentration; impacted to derive an overall rate expression for the complexation reaction; (iii) to identify a mechanism consistent with the experimental data; and (iv) to determine whether there are ecologically significant rate limitations for the reaction that might affect the biological impact of toxic aqueous aluminum during the mixing of natural waters. The present study builds upon an earlier investigation of the complex formation kinetics of A1F2+ in the pH range 0.9-1.5 (6).
Introduction Studies in North America and Europe have shown that an important consequence of acidic deposition is the transport of aqueous aluminum from terrestrial to aquatic ecosystems. In acidified surface waters, elevated levels of dissolved aluminum may create toxic lethal and sublethal conditions for biological communities, particularly fish (1, 2). However, biotic impacts may be strongly influenced by the chemical speciation of aqueous aluminum. For example, one study in the Adirondack region of New York ( I ) found that fish survival dropped significantly when fish were exposed to 0.5 mg/L aluminum in the pH range 4.4-5.2, but that survival increased when fluoride or citrate was added to the treatment solution (3). These findings imply that our ability to predict aluminum toxicity in natural aquatic environments will require detailed thermodynamic and kinetic understanding of aluminum complexation reactions occurring during the transport and mixing of acidic drainage waters. Thermodynamic calculations indicate that aluminum fluoride complexes can be extremely important in natural waters, and field studies in northeastern North America have confirmed that fluoride complexes are generally the
Experimental Section M sodium fluoride Materials. A stock solution of was prepared by weight from oven-dried reagent grade salt. All subsequent fluoride solutions were prepared from the single stock solution. The fluoride solutions were prepared in glass volumetric flasks but were immediately transferred to polyethylene bottles for storage to minimize the time of contact with glass, Aluminum solutions were prepared fresh each day from reagent grade A1(N03)3.9H20.Initial aluminum concentrations varied between 1.25 X loq5and 4.0 X lo4 M and were always in at least a 2-fold excess over initial fluoride concentrations to ensure only the 1:l complex, A1F2+, formed. Initial fluoride ion concentrations ranged from M. 5.0 X lo4 to 3.0 X All solutions were buffered with either sodium acetate-acetic acid mixtures or sodium dichloroacetate-dichloroacetic acid mixtures, and the ionic strength was adjusted to 0.1 M using sodium chloride. It has been shown that these buffers do not form complexes with aluminum, to any significant extent, in the range of pH and concentration used in this study (7,8). Eight different pHs were studied, ranging from 1.12 to 4.85.
Acidic deposition has an important effect on the transport and speciation of soluble aluminum. Toxicity of aqueous aluminum seems to be strongly dependent on aluminum speciation and the presence of complexing ligands such as fluoride. We report here a study of the complex formation kinetics of AlF2+in the environmentally significant pH range 2.9-4.9. The important paths are
W
k,
AP+ + F-
A1F2+
k-I
A10H2++ F- + H+ z% A1F2++ H 2 0 A13+ + HF
+
__
A10H2+ HF
160
klli
L
k-1,
k-,
A1F2+ + H+ A1F2+ + H 2 0
Environ. Sci. Technoi., Vol. 20, No. 2, 1986
(i) (ii) (iii) (iv)
0013-936X/86/0920-0160$01.50/0
0 1986 American Chemical Society
Kinetic Procedure. The experimental setup was as follows: An Orion single-crystal lanthanum fluoride membrane electrode was immersed in the reaction mixture to be monitored along with a saturated calomel reference electrode, and both were connected to an Orion Model 701A digital pH meter to measure the potential. In each run, 100 mL of a buffered fluoride solution was thermostated at the desired temperature in a plastic beaker and stirred with a magnetic stirring bar coated with Teflon. A 1-mL sample of the appropriate aluminum solution was then injected with an automatic pipet, and the potential was measured as a function of time. Reactions were run at 25, 22.5, 20, and 9 "C. The concentration of H+ was taken as [H'] = 10-PH/yH+, with yH+= 0.78, calculated from the Davies equation (7), where the square brackets denote concentration in mol/L. Solution pH was determined by use of a Fisher universal glass electrode with a Fisher saturated calomel reference electrode connected to an Orion digital pH meter. Standard buffers were used for reference with no corrections made for liquid junction potential differences. The concentrations of all species were determined as molarities rather than activities. However, since the ionic strength was constant throughout, the molarity of each species differed from its activity by only a proportionality constant dependent on the charge of the species.
10.0
9.0 W
0
- 8.0 x
I
LL
7.0
6.0 0
40
20 TIME
60
- SECONDS
Figure 1. [F-] vs. time plot at 25 OC,pH 4.35, and [AI3+] = 2.82 X 10-5 M.
Results and Discussion Determination of Reaction Rate. Frant and Ross (9), developers of the single-crystal lanthanum fluoride electrode, showed that its electromotive force obeys a Nernst-type equation of the form
E = constant - 0.05916 log [F-]
(1)
Consequently, the concentration of free fluoride ion at any time during the course of a reaction at constant ionic strength can be calculated if the initial free fluoride ion concentration and the change in potential are known. The response time of the fluoride ion selective electrode has been shown to be much less than 8 second and is capable of monitoring reactions that are faster than the aluminum fluoride complexation reaction (6, 10). Because H F is a weak acid, the initial free fluoride ion concentration is a function of the hydrogen ion concentration, the relation being
where
(3) Since the value of K m for our experimental conditions was not available in the literature, the constant was determined at an ionic strength of 0.1 M with the fluoride ion selective electrode by use of Bjerrum's method (11). The value obtained, 1196 L/mol, is comparable to those found at 0.1 M ionic strength in other media (12). Values of free fluoride ion concentration as a function of time were plotted in each experiment to obtain the rate of consumption of free fluoride ion. Figure 1shows such a kinetic plot, along with the tangent to the curve at time t = 0. The slope of this tangent gives the initial rate of free fluoride ion consumption, (-d[F-] /dt),,,. Because most of the kinetic runs in the present study gave plots similar to Figure 1, where tangents could be drawn with a slope reproducibility of about f4%, the decision was
6.0 0
5
IO
15
20
IF-I, MxIO6 Flgure 2. Test of rate eq 4 plot of rate/[F-] vs. [F-] at 25 OC and pH 4.35: upper plot, [AI3+]= 2.82 X lo-' M lower plot [AI3'] = 2.35 x 10-5 M.
made to use the initial rate method to analyze the kinetic data, as was done earlier by Srinivasan and Rechnitz (6). It was convenient in this study to define the initial rate of reaction as the initial rate of appearance of A1F2+,(d[A1F2+]/dt),,,, rather than the initial rate of free fluoride ion consumption. If the proton transfer reaction between HF and F- is assumed to be rapid in comparison with the rate of complex formation, the initial rates of formation of A1F2+can be obtained by multiplying the initial rates of consumption of free fluoride ion by the factor -(1 + KHF[H+])(6). Table I gives the values for the calculated initial rates of formation of A1F2+along with initial concentrations of fluoride, aluminum, and hydrogen ions. Rate Equation. As shown in Table I, the initial rate of formation of A1F2+ in all cases was first order with respect to [A13+]. However, this was not the case for the fluoride ion dependence. Instead, rate/ [F-] values increased with increasing [F-],as can be seen in Figure 2. Because the plots were linear and the intercept/[AP+] values as well as the slope/[A13+]values were approximately constant, a rate equation of the following form was indicated: rate = (d[A1F2+]/dt),,, = K1[A13+][F-]+ K2[Al3+][F-I2 (4)
In order to check this, the experimental data at each pH were fit to this rate equation using a nonlinear leastsquares computer program. The calculated rates of 82 runs when compared with the experimentally observed rates gave a percent standard deviation (8)) [ ( l - rate obEnviron. Sci. Technol., Vol. 20, No. 2, 1986
161
served/rate calculated)2/(N - 4)]1/2, of 4.8%. For individual pH values, the highest standard deviation was 7.0% at pH 3.85 and the lowest 2.8% for a limited number of runs at pH 1.62. There was no apparent correlation of percent standard deviation with pH. Mechanism of Complex Formation. A possible mechanism consistent with the experimental data is the so-called Eigen mechanism (13). In this scheme, the complex formation reaction involves the fast formation of an outer-sphere ion pair followed by a first-order process in which the ligand in the outer Coordination sphere replaces a water molecule in the inner coordination sphere of the metal ion. Since Al(II1) undergoes hydrolysis, four paths are possible via the Eigen mechanism: path 1
56
48
-
I
$ 32 -
IO
40
24 16
-
-
8 -
/
/; l
l
1
I
I
I
I
l
l
I
1
-
(H20)6A13++ F- = (H20)5A1(H20),F2+Kosl (5) (H20)5A1(H20),F2+
kl2
(H20)5A1F2++ H2O
(6) (d[A1F2+]/dt),=o= K1[Al3'][F-]
path 2 (H20)5A10H2+ + F- = (H20)5A10H,F+ KO:
(7)
& (H20)4A10HF++ H20
(8)
(H20)5A10H,F+
k84
where Kl =
(H20)5A1F2+
(9)
k43
path 3 (H20)6A13+ + HF = (H20)6A1,HF3+KO: (H20)5A1F2+ + H30+
(H20I6A1,HF3+
k,
and (10) (11)
path 4
+
(H20)&10H2+ HF = (H20)5A10H,HF2+KO:
+
(H20)5A10H,HF2+
(H20)4A10HF+ H30+
ke
-A (H20)5A1F2+
(H20)4A10HF+
(12) (13) (14)
k43
where Kml,KO:, KO:, and KO: are outer-sphere ion-pair association constants. Using the initial rate method, one can neglect back reactions in formulating the rate equation (14), so the initial rate of appearance of A1F2+from the four concurrent reactions is given by (d[A1F2+]/dt),,o = k12Kos1[A13+] [F-] + k23Ko,2[A10H2+][F-]+ k&o,3[HF][A13+] k56KO,4[A1OH2+] [HF] (15)
+
where the coordinated waters have been omitted for simplicity. Applying the steady-state approximation to the rapid protolytic reactions 16 and 17
__
A13+ + F-
k6
ke
A10H2++ HF
A13+ Z A10H2+ + H+ k7 k8
(16) (17)
>> k6 give
[A10H2+],, =
(19)
k32
(H20)4A10HF++ H+
and assuming k,
+ K2[Al3+][F-I2
k,[AP+][F-]
+ k7[A13+]
(18) ~,[H+I Reaction 16 occurs with the anions of weak acids and is 162 Environ. Sci. Technol., Vol. 20, No. 2, 1986
It can be seen that eq 19 is identical with the experimental rate equation, eq 4. Effect of pH on Reaction Rate. The rate of complexation exhibited strong pH dependence. Indeed, if the proposed mechanism is valid, there ought to be a significant dependence of both Kl and K2,and hence the reaction rate, on [H+] as can be seen from eq 20 and 21. Table I1 illustrates the observed Kl and K2values obtained at each pH from the nonlinear least-squares computer program. Equation 21 predicts that a plot of K2 vs l/[H+] ought to be linear. Equation 20 indicates that K1 values should produce a straight line when plotted against l/[H+] at high pH where k7k23Ko,2/kE[H+]predominates over k4&,3Km[H+] and a linear plot when plotted against [H+] at low pH where k4d(,,3Km[H+] is much greater than k7k?3K0,2/kdH'I Figure 3 shows the K2vs l/[H+] plot. As predicted from eq 21, a straight line was obtained between pH 3.36 and 4.35, which appeared to pass through the origin. K2 values were not observed below pH 3.36. Extrapolating the K2 value to pH 2.88 from Figure 3 gave a value of 1.17 X lo5 M-2 s-l. By use of this extrapolated value, the K2 term at pH 2.88 would only contribute a maximum of 5.4% to the total rate. Hence, the K2 term at pH
