Environ. Sci. Technol. 1988, 22, 1454-1459
Effect of Fulvic Acid on the Kinetics of Aluminum Fluoride Complexation in Acidic Waters Brian J. Plankey and Howard H. Patterson"
Department of Chemistry, University of Maine, Orono, Maine Both fluoride ion and fulvic acid are important aluminum binding ligands present in soil and surface waters. As such they play a role in the speciation and toxicity of natural waters that have increased aluminum concentration due to acid precipitation. We report here a kinetic study of aluminum complexation in the presence of both of these naturally occurring ligands. An overall mechanism has been identified and rate constants have been obtained for several of the reactions involved. We find that an a priori model of the two ligands in competition for aluminum is incorrect. In fact, the rate of fluoride ion consumption is increased by the presence of fulvic acid. Evidence is presented that this effect is due to several equilibria, some of which involve mixed-ligand species. The important equilibria in this three-component system are identified and discussed, as are aluminum speciation and toxicity in acidic waters. Introduction
Acid precipitation in northeastern North America has been shown to produce increased aluminum concentrations in both surface waters and soil waters (1-3). Increased concentrations of dissolved aluminum in acidified surface waters have been correlated with toxic conditions for fish and other aquatic animals (1, 4). However, aluminum toxicity to some aquatic animals appears to be strongly dependent on aqueous aluminum speciation (1,4). For example, dissolved inorganic complexes of aluminum, such as hydroxides and fluorides, have been shown to be toxic to young fish, whereas complexes of aluminum with naturally occurring organic ligands, such as fulvic acid, appear to be essentially nontoxic (1,5 ) . In previous investigations, we have reported kinetic study results of both aluminum-fluoride complexation in acidic waters (6) and aluminum-fulvic acid complexation in acidic waters (7). Since both fluoride ion and fulvic acid are present in natural waters, they both play a role in the speciation and toxicity of waters with increased aluminum concentration. In these earlier studies it was found that the reactions of aluminum with both fluoride ion and fulvic acid occur by the same mechanism and at approximately the same rate at 25 "C. However, as the temperature is lowered below 25 "C, the fluoride reaction is slowed considerably while the fulvic acid reaction is only slightly decreased in rate. On the basis of a simple competition model, temperatures below 25 "C ought to favor the complexation of aluminum by fulvic acid over complexation by fluoride and thus reduce aluminum toxicity in natural acidic waters. The present investigation seeks to test this competition model by studying the kinetics of aluminum-fluoride complexation in the presence of fulvic acid. The objectives of the investigation were as follows: (i) to determine the rate of aluminum-fluoride complexation in the presence of fulvic acid as a function of pH and concentration; (ii) to derive an overall rate expression for the complexation reaction; (iii) to identify a mechanism consistent with the experimental data; (iv) to test the simple competition model by comparison of the results of the present study 1454
Envlron. Sci. Technol., Vol. 22, No. 12, 1988
with the results of our two earlier papers (6, 7). Experimental Section
Materials. All materials were prepared exactly as in the aluminum-fluoride and aluminum-fulvic acid studies (6, 7). Initial total aluminum concentrations varied from 4.00 X to 6.00 X M and were always in at least a 2-fold excess over initial fluoride concentrations to ensure only the 1:l aluminum-fluoride complex formed. Initial fluoride ion concentrations ranged from 9.77 X lo4 to 3.19 X M. With 740 for the number-average molecular weight of fulvic acid (7), initial fulvic acid concentrations varied between 6.90 X lo4 and 5.56 X M. All solutions were buffered with sodium acetate-acetic acid mixtures or the pH was adjusted by addition of HC1, and the ionic strength was adjusted to 0.1 M with sodium chloride. It has been shown that neither acetate nor chloride forms complexes with aluminum, to any significant extent, in the range of pH and concentrations used in this study (8-10). Reactions were studied at several different pHs ranging from 2.88 to 4.35. Kinetic Procedure. The reactions were run and monitored the same as were the aluminum-fluoride reactions (6). In a typical run, 101 mL of buffered solution containing both F- and fulvic acid (FA) was thermostated at 25 "C. The [F-] was monitored with a fluoride ion selective electrode and a saturated calomel electrode for reference. The reaction was initiated by injecting a 1-mL sample of the appropriate aluminum nitrate solution and monitored by measuring the potential as a function of time. The response time of the fluoride ion selective electrode has been shown to be much less than 1 s and is capable of monitoring reactions that are faster than the aluminum-fluoride complexation reaction (11, 12). Results and Discussion
Determination of Reaction Rate. Values of the free fluoride ion concentration, as measured by the fluoride ion selective electrode, were plotted as a function of time. A typical [F-] vs time plot is shown in Figure 1 along with the tangent to the curve at time t = 0. The slope of the tangent gave the initial rate of free fluoride ion consumption. These plots were very similar to the aluminum-fluoride reaction [F]vs time plots in which the initial rate method was used to analyze the kinetic data (6). Therefore, we have also determined the initial rate of free fluoride ion consumption in all of the three-component reactions of aluminum with fluoride ion in the presence of fulvic acid (FA). These initial rates are reported in Table I. We have found that the initial rate of consumption of free fluoride ion is greater when FA is present in the initial reaction mixture than when just aluminum and fluoride are present. This effect is not due to fluoride ion adsorption by FA, because the addition of FA to a fluoride ion solution does not result in a decrease in the fluoride ion concentration. Therefore, fluoride ion complexes aluminum at a greater rate when FA is present. This is an unexpected result if one presumes a simple competition
0013-936X/88/0922-1454$01.50/0
0 1988 Amerlcan Chemlcal Soclety
r
I
I
I
I
I
I
!
'7.60
I
I
I
'4.4
-4.3 log [AI
11
0
20
40
60
TIME, s
Flgure 1. kF-1 vs time at M, [H ] = 1.11 X
[AI3'] = 3.88 X M, and 25
I
'4.2
I
1'
Flgure 2. log initial rate vs log [AI3+] at [F-] = 1.55 X lo-' M, [FA] M, and 25 O C . = 0.69 X M, [H'] = 3.71 X
M, [FA] = 1.04 x
OC.
for aluminum between F- and FA. In this system at least, a simple competition model can be seen to be a naive way of characterizing aluminum complexation in the presence of two coordinating ligands. At the beginning of this investigation, in repeating the measurement of the initial rate of consumption of Fwithout FA present, it was apparent that the absolute initial rate was smaller than in our first study. For ex= 1.80 X lo4 M and ample, in the present study with [H+] M at 25 "C,the initial rate/[AP+] was [ F ] = 1.53 X found to be 1.12 X s-l, whereas in the earlier work under the same conditions initial rate/ [A13+]was deters-l. The reason for this discrepancy mined to be 1.56 X is in the manner in which the initial rate was determined. In the earlier work, [F-] vs time was plotted and tangents were drawn at t = 0 to obtain the initial rate of free fluoride ion consumption. In the present investigation,the plotting and slope determination were done with an Apple IIe computer using a commercial program called CURVE FITTER. The program produced least-squares fittings of data points to exponential equations from which initial rates were obtained directly. In order to make certain that the discrepancies were due to data treatment alone, the data of the earlier run that had given initial rate/ [A13+]= 1.56 X s-l were treated with the CURVE FITTER program. The revised result of initial rate/[A13+] = 1.09 X lod3s-l was in excellent agreement with the 1.12 X s-l value from the present study. Since the results were reproducible in both studies, the present study gave consistently smaller values of initial rate of fluoride ion consumption in the absence of FA. We conclude that the earlier method of obtaining initial rates was perhaps biased by overemphasizing the importance of the first few data points. We attempted to further correct this problem by weighing the individual data points in the least-squares analyses by giving the most weight to the point at t = 0, where no change is occurring, and the least weight to the point at t = 5 s, where mixing may be incomplete. The present computer method is most likely to be reproducible by other investigators, and so the lower values of the present study should be considered better. It should be pointed out that to acknowledge the results of the present study as an improvement over the earlier study changes only the values of the forward and reverse rate constants somewhat (see below). The mechanism and pH dependence are unchanged. Furthermore, the environmental conclusions drawn from half-life arguments are
I
*/
'4.0
'4.7
'4.6
'4.5
log 1F-l
Flgure 3. log initial rate vs log [F-] at [AI3+] = 5.96 X = 0.69 X M, [H+] = 5.31 X M, and 25 'C.
M, [FA]
160
"
.'
140
100
t/ 0
20
10
"'1
total 9
30
40
50
M x10'
Flgure 4. Initial ratel[Al3'][F-] vs [FA] at [H'] = 1.11 X lo-' M and 25 O C . (a) [OAc-] = 2.01 X lo-* M; (b) [OAc-] = 0.
also unchanged because initial rates were not used to obtain half-lives. Rate Equation. The dependence of the initial rate on both aluminum concentration and fluoride ion concentration is first order. Figure 2 shows a typical plot of log initial rate vs log [A13+]at constant [F-1, [FA],,, and pH. Figure 3 is a typical plot of log initial rate vs log [F-] at constant [A13+],[FA],,,, and pH. All such plots were linear with slopes ranging from 0.98 to 1.09 with an average Environ. Sci.
Technol., Vol. 22,No. 12, 1988
1455
~~
+
Table I. Initial Rates of Consumption of F-at 26 "C
[H+l, M 5.72 X 5.72 X 5.72 X 5.72 X 5.72 X 5.72 X 1.11 x 1.11 x 1.11x 1.11 x 1.11 x 1-80x 1.80 X 1.80 X 1.80 X 3.71 X 3.71 X 3.71 X 3.71 X 3.71 X 3.71 X 5.31 X 5.31 X 5.31 X 5.31 X 5.31 X 5.31 X 5.31 X 1.62 X 1.62 X 1.62 X
[A13'l, M X lo6
lod 10"' 10-4 10"' 10-4 10-4
10-4 10"'
10"' 10"' 10"' 10"' 10"' 10"'
10"' 10"' 10"' 10"' 10"' 10"'
10"' 10"' 10"' lo4 lo4
M
3.77 4.71 5.65 4.71 4.71 4.71 3.88 4.85 3.88 3.88 3.88 3.92 4.90 5.88 5.88 3.96 4.95 5.94 4.95 4.95 4.95 4.97 5.96 5.96 4.97 3.97 5.96 5.96 4.99 4.99 4.99
VI,
lo6
X
1.54 1.52 1.52 0.989 1.22 2.00 1.52 1.52 1.52 1.55 1.57 1.57 1.58 1.52 2.01 1.55 1.57 1.55 1.56 1.55 1.53 2.16 2.15 1.61 1.58 1.58 2.65 3.15 1.66 1.61 1.62
[FA],
M
X
lo6
A10H2+ FA-
init rate," M s-l X lo8
0.69 0.69 0.69 0.69 0.69 0.69 0.69 0.69 1.38 2.78 4.37 0.69 0.69 0.69 0.69 0.69 0.69 0.69 1.38 2.78 5.56 0.69 0.69 0.69 0.69 0.69 0.69 0.69 1.46 2.91 4.37
20.7 28.4 37.6 22.9 23.3 32.9 6.87 8.72 7.24 8.60 9.86 4.49 5.71 6.70 8.25 2.89 3.67 4.32 3.51 3.99 4.74 3.57 4.27 3.38 2.80 2.17 5.66 6.67 2.05 2.07 2.18
+
where kI and kII are observed rate constants that are functions of the hydrogen ion concentration. In deriving a possible mechanism for the reaction of aluminum with fluoride ion in the presence of FA, the following reactions need to be considered. Coordinated waters have been omitted for simplicity.
+
k
AlOHF+ + HzO
k-1
k-a
+
k
A10H2+ HF &A1F2++
A13+ + FA1456
k4
H20
A1(FA)2+
tnviron. Sci. Technol., Vol. 22, No. 12, 1988
k
(10)
k 10
Al(FA)OHF z-
(11)
A13+
k 11
T T
+
A10H2+ H+
(12)
A P + F- F A10H2++ HF A13+ + FA-
k13
(13)
+
A10H2+ HFA
%-
(14)
Reactions 12-14 are all protolytic equilibria and are assumed to be rapid in comparison to complex formation (8, 9). Since the initial rate method is being used to analyze the kinetic data, back-reactions can be neglected in formulating the rate equation (15). The initial rate of consumption of fluoride ion is then given by
+
+
-(d[F-] /dt),=, 121[Al3+][F-] k2[A10H2+][F-] k9[A1(FA)2+][F-]+ klo[Al(FA)OH+][F-] (15)
(2)
But since both [A1(FA)2+]and [Al(FA)OH+]are zero at t = 0, the last two terms of eq 15 are equal to zero. Applying the steady-state approximation using reactions 12-14 to obtain the initial concentration of A10H2+gives
(3)
[A10H2+]i,iti,l = kll [A13+] kI2[Al3+][F-]
+
.Iz-ll[H+]+ k-l,[HF]
(4) (5)
k k-6
(9)
In this reaction scheme, HFA represents a fully protonated aluminum binding site on FA, and FA- represents a deprotonated site (13). This notation is used only to indicate that aluminum binding sites on FA are also proton binding sites (13). It is not intended to identify the actual charge, the type of aluminum binding sites, or the proton binding behavior of FA. Also, for simplicity we have assumed FA to be a monodentate ligand. This, however, does not affect the kinetic analysis if any subsequent ring closure is very rapid and thus not rate determining. Rapid ring closure has been observed for the reaction of Fe3+with FA (14). Reactions 2-5 have been shown to be the important reactions in the complexation of aluminum by fluoride alone (6). Similarly, reactions 6-9 were found to be the reactions by which aluminum is complexed by fulvic acid alone (7). Reactions 10 and 11,however, involve mixedligand species in which both fluoride and fulvic acid are bound to aluminum. In addition to the above reaction scheme, the following equilibria must be considered.
k
A13+ + HF 2A1F2++ H+
k 2 Al(FA)OH+ + H+ k-a
Al(FA)OH+ F-
= initial rate = k1[Al3+][F-l+ k11[Al3+1[F-I [FAlbbi (1)
A10H2+ F-
(8)
+ $Al(FA)F+
i
k + F- & A1F2+ k-1
+
k-7
A1(FA)2+ F-
slope of 1.03. This is good evidence of first-order dependence of the rate on both [F-] and [A13+]. Log-log plots of log rate vs log [FA],, were not linear and indicated a fractional order with respect to FA],^ Linear plots with nonzero intercepts were obtained when initial rate/ [Al3+][F-] was plotted against [FA],, at constant pH. Such a plot is shown in Figure 4a. On the basis of the results summarized in Figure 2-4a, the experimental rate equation for the aluminum complexation of fluoride ion in the presence of FA can be expressed as
A13+
k
+
(7)
k-6
A13+ HFA .& A1(FA)2+ H+ A10H2+ + HFA
The average standard deviation of initial rate values was 3.2% at pH 4.06 and below.
(-d[F-]/dt),,,
& Al(FA)OH+
(6)
+ k19[A13+][FA-] + 12_13[HFA]
(16)
Both reactions represented by rate constants 12-12 and involve bond breaking in addition to bond making. Thus, it is likely that the rate of reaction represented by rate constant k-ll [which is limited only by the rate of diffusion (16)]is much greater than the reverse of reactions 13 and 14 (11). We can therefore assume that 12-13
k-ll[H+] >> k-l,[HF]
+ k-,,[HFA]
(17)
The expression for [A10H2+]initidthen becomes [A10H2+]initid= k11[Al3+] k12[Al3+][F-]
+
[H+I,
+ k13[A13+][FA-]
k-ll[H+l
(18)
+
(-d[F-]/dt),=o = k1[A13'][F-] kzk1l[Al3+][F-]/k-ll[H+] k2k1z[Al3+][F-I2/k-11[H+] + kzk13[Al3+1[F-] [FA-]/k-ll[H+l (19) If, however, h i
kidFA1
>> k12[F-1
(20)
then the term in eq 19 with [F-I2 is comparatively small and the [PI2 dependence of the initial rate will not be seen. This is quite reasonable for two reasons. First, as will be shown below, k13 is quite large in comparison to kll. Second, in the reaction of aluminum with fluoride alone, kll alone was not very much greater than klz[F-], and the [F-I2 dependence of the rate was observed (6). However, this was true only at relatively high pH. Below pH 3.36 the [ P I 2dependence was not observed, indicating that klz was relatively small. In fact, in repeating runs with no FA present and with the computer to determine initial rates, k12 was small enough so that the [PI2dependence was not observed at any pH for the F- concentrations used in this study. Assuming, therefore, that eq 20 is valid gives (-d[F-]/dt),,o
1.68 x 1.65 x 1.44 x 1.35 x 3.74 X 3.74 x 3.74 X 3.71 X 3.71 X 3.71 x 3.02 X 2.19 X 2.19 X 2.19 X
(22)
~ I =I
(23)
lo4
At this point it appeared that FA was increasing the initial rate of fluoride complexation by increasing the initial [A10H2+]/[A13+]ratio according to eq 18. If this were so, acetate, which was used as buffer, could do likewise according to eq 24. To test this, several runs were made with
[FA], M x 105
init rate: M ~1 x 108
0.987 1.59 1.70 1.74 1.59 2.14 2.67 1.62 1.63 1.63 1.88 1.43 1.44 1.44 1.47 1.50 1.50 1.52
0 4.37 1.46 2.91 0 0 0 1.46 2.91 4.37 0 1.46 2.91 4.37 0 1.46 2.91 4.37
0.329 0.790 0.798 0.893 2.12 2.93 3.73 2.53 3.51 3.87 2.76 4.08 5.40 6.59 6.95 8.21 9.53 11.9
Table 111. Comparison of Initial Rates with and without Acetate Buffer
[FAI~M, M
init rate/ [A13+][F-1, M-1 s-l
[H+] = 3.71 X 10"' 0 4.37 x 10-5 0 4.37 x 10-5
26.9 48.0 40.2 57.9
[OAK], M 0 0 2.04 X 2.04 X
[H+] = 1.11 X 10"' 0 0 2.01 x 10-2 2.01 x 10-2
and kzki3/k-ii[H+]
[PI, M x 105
"The average standard deviation of initial rate values was 3.5%.
Equation 21 can be seen to be identical with the experimental rate equation, eq 1,with
= kl + k2kll/k-ll[H+I
106
4.99 4.99 4.99 4.99 4.96 4.96 4.96 4.95 4.95 4.95 4.95 4.92 4.92 4.92 4.85 4.85 4.85 4.85
10-3 10-3 10-3 10-3 10"' 10-4 lo4 10"' lo4 10-4 lo4
10"' lo4 1.11 X 10"' 1.11 X 10"' 1.11X 10"' 1.11 X 10"'
= k1[A13+][F-] + k2h[Al3+1[F-l/k-dH+l + kzk13[A13+1[F-I [FA-]/k-ll[H+I (21)
kI
[A13+],
Mx
M
Substitution of eq 18 into the rate expression, eq 15, yields
+
Table 11. Initial Rates of Consumption of F- at 25 OC in the Absence of Acetate Ion
?*
100
-
00
-
97.5 161 107 162
0 4.37 x 10-5 0 4.37 x 10-5 I
I
2
4
7
E
+
A13+ + OAc- + A10H2+ HOAc (24) no acetate present (the pH being adjusted by addition of HC1) both with and without FA present. The results are shown in Table 11. As expected, the presence of acetate did indeed increase the initial rate of aluminum-fluoride complexation. A comparison of initial rates obtained with and without acetate present are summarized in Table 111. Figure 4b shows a plot of initial rate/[AF+][P] vs [FA],a with no acetate present. Comparison with Figure 4a, in which acetate was used as buffer, again shows the increase in initial rate caused by the presence of acetate. Effect of pH on Reaction Rate. Whether the solutions were buffered with acetate or not, in the pH range 2.88-4.06 there was a steady increase in the initial rate of F- consumption with increase in pH. Plots of initial rate/ [Al3+1[F-1vs 1/ [H+] at any total [FA] were linear in this pH range, as shown for example in Figure 5a with M and no acetate buffer present. [FA],,, = 1.46 X Above pH 4.06 there was an increase in rate such that
0
6
0
Figure 5. Initlal rate/[A13+][F-] vs l/[Ht] at [FA] = 0.69 X O C . (a) [FA],, = 1.46 X IO-' M; (b) [FAltOhl= 0.
10
M
and 25
initial rate/ [A13'] [F-] values fell above the straight line plots. This has been observed before and is thought to be caused by more highly hydrolyzed species of Al(II1) providing additional paths for aluminum-fluoride complexation (6-9, 17). The intercept of Figure 5a is near zero while the slope equals 1.26 X lo-, s-l. According to eq 21, the intercept equals kl and the slope equals k2kll/k-11 + k2k13[FA-]/k-,l. As the total FA concentration was increased, plots of initial Environ. Scl. Technol., Vol. 22, No. 12, 1988 1457
c
Table IV. Comparison of Rate Constant Values Obtained from Present and Previous Studies
kl,
M-'s-'
kz, M-'s - ~ kp, M-'S-' k-1, S-' k-2, S-' k-d, M-' s-'
present study
previous study
>2 (3.15 f 0.11) x 103 (9.61 f 0.34) X lo2 >8 X lo-' (5.43 0.19) x 10-3 2.10 f 0.07
37 3.61 x 103 1.10 x 103 1.42 X 6.22 x 10-3 2.27
rate/[A13+][F-] vs l/[H+] at constant [FA],, yielded increasing slopes and the intercepts remained nearly zero, as expected. For example, at [FA],oM= 2.91 X and 4.37 X M, the intercepts were 2.87 and 1.66 M-l s-l, respectively, and the slopes were 1.46 X lo-, and 1.81 x lo-, s-'. While these results can be used to obtain values for the rate constants kl and k,, the complex mixture nature of FA renders the values somewhat doubtful. These values have therefore been determined in the absence of FA, as described in the next section. Fulvic acids contain a distribution of different proton binding sites with varying acid dissociation constant values, which may be somewhat different depending on the source of FA (13). As a result there is no single k,, value for any FA. Furthermore, the acid dissociation constant values of a FA are not really constants at all. They depend on the degree of ionization, and hence the pH (17). At higher pH, the stronger proton acceptors of FA will be deprotonated and therefore an apparent k13 will increase with pH. This could affect both the slopes and intercepts of the initial rate/[Al3+J[F-] vs 1/[H+] plots when FA is present. Rate Constants. Since the complexity of FA prevents accurate determination of kl and k, values when FA is present, the runs made without FA present, and unbuffered as well, were used to obtain values for these rate constants. A plot of initial rate/[A13+][F-] vs l/[H+] for these runs is shown in Figure 5b. The essentially zero intercept indicated kl N 0, and from the slope of k2k11./k-ll = 1.12 X s-l, k2 = 3.15 X lo3 M-l s-l was obtained, using 3.55 X lo4 M for k l l / k l i (8). These values can be compared to our earlier study done in the presence of acetate that gave k, = 37 M-l s-l and kz = 3.61 X lo3 M-l s-l. The increase in rate due to the presence of acetate used as buffer exaggerated the kl value in our earlier study while not affecting the k, value a great deal. In the absence of acetate, the value of k4 can be recalculated by again assuming that the mechanism for ligand penetration for reactions 3 and 5 are both dissociative (6, 18,19). In that case, the rate constants for ligand penetration for these two reactions are equal, and we can obtain k4 using the value of k, and the Fuoss equation (6, 20). Doing so yielded k, = 9.61 X 10, M-l s-'. As was done earlier (6,9),we have calculated values for the rate constants for the reverse reactions from the forward rate constants and equilibrium quotient values at = 0.1 M derived from the literature (8, 21). Both the forward and reverse rate constant values are reported in Table IV along with a comparison of values of these same rate constants obtained in our earlier paper. The lower values obtained in the present study should be considered better, as discussed above. Overall Rate and Environmental Considerations. It appears that the k2k13[A13+] [F-][FA-]/k-ll[H+l term in eq 19 was responsible for the increased initial rate of Fconsumption in the presence of FA. This term arises from eq 14 and leads to an increased initial rate of F- consumption, apparently because reaction 14 yielded a sig1458
Environ. Sci. Technol., Vol. 22, No. 12, 1988
160
7
E
120 140
1
A
I
40 20
-
0
2.9
3.1
3.3
3.5
3.7
3.9
4.1
PH Figure 6. Initial rate/[AI3+][F-] vs pH at several different values of [FA]: (0)[FA] = 0; (A)[FA] = 1.46 X lo-' M; (0) [FA] = 2.91 X lo-' M; (W) [FA] = 4.97 X M.
nificant increase in the initial concentration of A10H2+, which is known to complex ligands at a greater rate than AP+ (6,8). Therefore, the protolytic equilibrium represented by eq 14 seems to be quite important in determining the initial rate at which aluminum was complexed by both fluoride ion and fulvic acid in this three-component system. In order for the kzk13[A13+] [F-][FA-]/kl1[H+]term in eq 21 to make a significant difference in the overall initial rate, hi3 and/or [FA-] must be fairly large. Despite the fact that ki3 is neither a single rate constant nor even constant over any significant pH range, we can obtain an approximate average value for k13/kil by plotting the slopes of the initial rate/[AP+][F] vs l/[H+] plots against [FA],w From the intercept of this plot k2k11/k-ll = 1.07 X s-l and the slope kZk13/k-li = 156 M-l s-l, we obtained k13/k11 = 1.45 X lo4 M-l. This is only an approximate average ratio over the pH range of about 3-4. It does, however, indicate why even small concentrations of FA caused observable increases in the initial rate of aluminum-fluoride complexation. In addition, the concentration of deprotonated FA- sites will increase with increase in pH at any given [FA],,. The combined effect was that even very low concentrations of FA increased the initial rate of aluminum-fluoride complexation and the increase in initial rate due to FA increased as the pH was raised. Figure 6 shows the pH dependence of the initial rate as a function of [FA],,,d. In addition to the increased initial rate of fluoride complexation in the presence of FA, eq 15 indicates that when t # 0, the rate of F- consumption may be further increased because of the presence of Al(FA)2+and Al(FA)OH+. This effect was observed by mixing the A13+and FA solutions together and allowing them to react -1 h prior to injecting into a fluoride ion solution. With [H+] = 5.31 x 10-4 M, [PI = 1.60 x 10-5M, [AP+I = 5.96 x 10-5 M, and [FA] = 6.90 X lo4 M, the initial rate of F- conM s-l. Under the same condisumption was 4.54 X tions without prior mixing of A P and FA solutions, the initial rate was 3.38 x M s-l. This clearly shows the contribution of at least one of the last two terms of eq 15 to the rate when [A1(FA)2+]and [Al(FA)OH+] are not
initially zero. These last two terms arise from reactions 10 and 11, and one or both of these equilibria were apparently important in determining the overall rate of reaction and reaction half-lives. Therefore, the overall rate of consumption of F in the presence of FA was increased beyond that caused by increased [A1OH2+Iinitid due to reaction 14. By analogy with the relative forward rate constants of reactions 2 and 3, reaction 11 is likely to be the most important of reactions 10 and 11, both in affecting the overall rate and in determining speciation in mixed-ligand complexes. As a consequence of increased overall rates due to reaction 14 and probably reaction 11, half-lives for aluminum-fluoride complexation were shorter in the presence of FA and decreased with increase in FA conM, [F-] = centration. For example, at [H+] = 2.19 X 1.44 X 10" M, and [A13+] = 4.92 X M, the first half-life period decreased from 243 to 180 to 145 s in going from M. At [FA],, = 1.46 X to 2.91 X 10" to 4.37 X M, and [A13+] = [H+] = 3.71 X 104M, [F-] = 1.60 X 4.96 X M, the first half-life period decreased from 632 s with no FA present to 253 s at [FA],, = 4.37 X lo4 M. The presence of FA therefore reduced the time to equilibrium of aluminum-fluoride complexation. However, aluminum speciation was also changed with such mixedligand complexes as Al(FA)F+,Al(FA)OH+,and Al(FA)OHF being formed. The situation with regard to toxicity is uncertain. Since aluminum-organic complexes appear to be essentially nontoxic to young fish and aluminumfluoride complexes have been shown to be toxic (I, 5), it is possible that mixed organic-fluoride complexes of aluminum are more toxic than aluminum-organic complexes and less toxic than aluminum-fluoride complexes. This needs to be investigated. On the basis of this study, however, it is clear that a simple competition model will not suffice in predicting aluminum speciation or toxicity. Thus, having FA present in soils and waters subject to increased aluminum concentration due to acid precipitation will not guarantee a reduction in aluminum toxicity simply because FA outcompetes inorganic ligands such as fluoride. Conclusion. A critical factor in determining the rate at which aluminum is complexed by naturally occurring ligands is the [A10H2+]/[A13+]ratio. A10H2+has been shown to react with complexing ligands at a much greater rate than AP'. Any substance that increases [A10H2+] relative to [A13+]will thus increase the rate at which aluminum is complexed by all complexing ligands present. Therefore, soluble sulfates, carbonates, fluorides, fulvic acids, and hydroxides in natural systems will increase the rate at which aluminum is complexed by all naturally occurring aluminum binding ligands present. In contrast,
soluble chlorides and nitrates cannot increase A10H2+ concentrations in natural waters. For example, the equilibrium
+
(H20)6A13+ C1-
+
(H20)5A10H2+ HCl
(25)
lies entirely to the left and therefore the [A10H2+]/[A13*] ratio is unaffected by the presence of chloride. Thus, the rate of aluminum complexation in natural systems will be unaffected by the presence of soluble chlorides and nitrates. Registry No. Al, 7429-90-5. Literature Cited Driscoll, C. T., Jr.; Baker, J. P.; Bisogni, J. J., Jr.; Schofield, C. L. Nature (London)1980, 284, 161-164. Cronan, C. S.;Schofield, C. L. Science (Washington,D.C.) 1979,204, 304-306. Johnson, N. M.; Driscoll, C. T.; Eaton, J. S.; Likens, G. E.; McDowell, W. H. Geochim. Cosomochim.Acta 1981,45, 1421-1437. Burrows, W. D. CRC Crit. Reu. Environ. Control 1977, 7, 167-216. Johnson, N. M.; Likens, G. E.; Feller, M. C.; Driscoll, C. T. Science (Washington, D.C.) 1984,225, 1424-1425. Plankey, B. J.;Patterson, H. H.; Cronan, C. S. Environ. Sei. Technol. 1986,20, 160-165. Plankey, B. J.; Patterson, H. H. Environ. Sci. Technol. 1987, 21, 595-601. Perlmutter-Hayman, B.; Tapuhi, E. Inorg. Chem. 1977,16, 2742-2745. Secco, F.; Venturini, M. Inorg. Chem. 1975,14,1978-1981. Holmes, L. P.; Cole, D. L.; Eyring, E. M. J. Phys. Chem. 1968, 72, 301-304. Srinivasan, K.; Rechnitz, G. A. Anal. Chem. 1968, 40, 1818-1825. Thompson, H. I.; Rechnitz, G. A. Anal. Chem. 1972,44, 300-305. Saar, R. A.; Weber, J. H. Environ. Sci. Technol. 1982,16, 510A-517A. Langford, C. H.; Kahn, T. R. Can. J. Chem. 1975, 53, 2979-2984. Espensen, J. H. Chemical Kinetics and Reaction Mechanisms; McGraw-Hill: New York, 1981; p 81. Laidler, K. T. Chemical Kinetics;McGraw-Hilk New York, 1965; pp 535-537. Gamble, D. S. Can. J. Chem. 1970,48, 2662-2669. Perlmutter-Hayman, B.; Tapuhi, E. Inorg. Chem. 1979,18, 875-877. Hasinoff, B. B. Can. J. Chem. 1976,54, 1820-1826. FUOSS, R. M. J. Am. Chem. SOC.1958,80,5059-5061. Baumann, E. W. J. Inorg. Nucl. Chem. 1969,31,3155-3162.
Received for review May 11,1987. Revised manuscript received February 24,1988. Accepted June 20, 1988. This is a contribution from the ALBIOS project, EPRI Contract RP-2365-01.
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