Environ. Sci. Technol. 1990, 24, 583-588
Casaletto, G. J.; Gevantman, L. H.; Nash, J. B. USNRDL-TR-565; 1962. Dorfman,L. M.; Hemmer, B. A. J . Chem. Phys. 1954,22, 1555. Eakins, J. D.; Hutchinson, W. P. In Tritium; Moghissi, A. A,, Carter, H. W., Eds.; Messenger Graphics: Phoenix, AZ, 1973; pp 392-399. Sweet, C. W.; Murphy, C. E., Jr. Enuiron. Sci. Technol. 1981, 15, 1485. Easterly, C. E.; Bennett, M. R. Nucl. Technol./Fwion 1983,
(9) Usami, S.; Matsuyama, M.; Watanabe, K.; Takeuchi, T.; Nogami, H.; Asai, Y.; Hasegawa, K. Annu. Rep. Tritium Res. Cent., Toyama Univ. 1987, 7, 81. (10) Usami, S.; Asai, Y.; Hasegawa, K.; Matsuyama, M.; Watanabe, K.; Takeuchi, T. Annu. Rep. Tritium Res. Cent., Toyama Univ. 1988,8,75. (11) Roesch, W. C. P-10 Chemical Equilibria. USAEC Report
HW-17318;Hanford Atomic Products Operation, March 1950.
4 , 116.
Yang, J. Y.; Gevantman, L. H. J . Phys. Chem. 1964,68, 3115. Ichimura, K.; Matsuyama, M.; Watanabe,K.; Takeuchi, T. Fusion Technol. 1985, 8, 2407.
Received for review August 1,1989. Accepted December 11,1989. We acknowledge the subsidy received for this research under the program of Grant-in-Aid for Fusion Research, by the Ministry of Education, Science and Culture of Japan.
and Ionic Strength Effects on Nickel-Fulvic Acid Dissociation Kinetics Stephen E. Cabaniss' Interdisciplinary Programs in Health, Harvard School of Public Health, Harvard University, Boston, Massachusetts
The dissociation kinetics of Ni(I1)-fulvic acid (FA) complexes as a function of FA concentration, pH, and ionic strength (I)were observed by ligand-exchange spectrophotometry. Lowering I and raising pH strongly decreased the dissociation rate; increasing the FA concentration caused a smaller rate decrease. Four data analysis methods estimated component concentrations Ci and rate constants ki for the Ni(I1)-FA mixture. Methods that allowed only either C, or ki to vary in the regression led to a more consistent relationship between these parameters and solution conditions. The most slowly reacting Ni(I1)-FA complexes were least affected by changes in pH and I. The proportion of slowly reacting complexes increased with increasing FA. Consequently, experiments a t environmental Ni:FA ratios are necessary to determine the importance of pH and I in controlling Ni(I1)-FA dissociation in the environment. Introduction Ionic strength, pH, and metal to ligand ratio can strongly influence the ionic equilibria of fulvic acid (FA) ( I ) . This study examines the effects of salt concentration, pH, and Ni:FA ratio on the dissociation rates of Ni(I1)-FA complexes and concludes that these factors are also important variables for modeling metal-FA kinetics. Nickel is an essential micronutrient for some marine algae, and micromolar Ni(I1) is toxic to a variety of algae, invertebrates, and fish (2). Ni(I1) equilibria with FA and dissolved organic matter (DOM) have not been studied frequently ( 2 , 3 ) ,perhaps because of the absence of convenient, sensitive methods for measuring aquo [Ni2+]in the presence of organic complexes ( 4 ) . However, recent voltammetric studies indicate that >95% of Ni(I1) is organically complexed under marine conditions ( I 7). Many metal-organic complexes associate and dissociate rapidly relative to environmental processes, e.g., algal uptake or particle sinking ( 5 ) . These relatively labile systems can be modeled by use of equilibrium considerations alone. However, equilibrium models may not apply to slowly dissociating complexes like Ni(I1)-FA (6) or *Current address: Dept. of Chemistry, Kent State University, Kent, OH 44242. 0013-936X/90/0924-0583$02.50/0
Ca-EDTA (7). These relatively inert systems require combined kinetic and equilibrium models to properly predict metal partitioning. Dissociation of Ni(I1)-FA complexes has been observed via ligand exchange with 4-(2-pyridylazo)resorcinol(PAR). An equilibrated solution of Ni(I1) and FA is mixed with a large excess of PAR, so that the dissociation reaction Ni(I1)-FA
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+
Ni2+ FA
-
(A)
is followed by a rapid, irreversible reaction Ni2++ 2PAR
Ni(PAR)2
(B)
where Ni(PAR)2 is a highly absorbing complex (6). Langford and co-workers studied Ni(I1)-FA complexation equilibria using this approach, measuring rates of Ni(11)-FA complex dissociation a t constant reaction pH 7.8 and ionic strength I = 0.125 in solutions equilibrated a t different Ni:FA ratios and a t different pHs. The concentration of Ni(I1)-FA as a function of time was represented by [Ni-FA](t) = c C i exp+lt i= 1
(1)
[Ni-FA](t) is the molar concentration of Ni(I1) complexed by FA as a function of time, determined by [Ni-FA](t) = [NilT - [Ni(PAR),]
(2) where [NilT is the total Ni(I1) concentration. C,is the molar concentration a t time t = 0 of the ith Ni(I1)-FA component. ki is the first-order dissociation constant for reaction A for the ith component. Note that although for a reversible reaction (A + B) the constant It, would be the s u m of the association and dissociation constants,the large excess of PAR makes the reaction essentially irreversible so that It, represents only the dissociation rate constant. The authors modeled the data using four kinetic components, one of which, (i = 1)represented the hexaaquo ion. At constant reaction pH and I, they found that the component rate constants ki varied randomly ca. 25%. In contrast, the component concentrations C,varied systematically up to 3-fold as a function of equilibration pH and Ni:FA ratio, leading the authors to conclude that a fourcomponent model reasonably represented Ni(I1)-FA equilibrium speciation (6).
0 1990 American Chemical Society
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This paper presents PAR-exchange experiments in which reaction conditions were varied to observe changes in the rate constants ki. Ionic strength, pH, and Ni:FA ratio were selected as reaction conditions to vary because they strongly influence metal-FA equilibria (1,8). The primary objective of this paper is to determine which factors may be important for kinetic modeling of Ni(I1) speciation in the environment. Materials and Methods
Suwannee River FA is a standard reference material of the International Humic Substances Society, and its isolation and chemical properties have been previously described (9,10). White Oak River FA was isolated from a pristine coastal watershed in eastern North Carolina, as described earlier (11). Suwannee FA and White Oak FA are 53.8% and 44.5% carbon by weight, respectively. Stock solutions of 50 mg of FA C/L were refrigerated less than 6 weeks before use. Nickel stock solution (1.00 mM) was made from nickel(I1) nitrate hexahydrate. The pH was buffered with 1.0 mM HC0,- and adjusted to pH 7.4 or 8.4 with HClO, and either NaOH or KOH. Ionic strength was adjusted to 0.002, 0.004, 0.010, 0.030, or 0.10 with NaC10, or KNOB. Fresh 1.0 mM PAR (Aldrich)solution was made each week and stored in the dark to prevent decomposition. Dissociation was observed by measuring the color formation accompanying the rapid, irreversible complexation of dissociated Ni2+ by a large excess of PAR. Ligand exchange with PAR has been used to observe dissociation reactions of Cu-DOM (12, 13) and Ni(I1)-FA (6). Absorbance of Ni(PAR)2was measured a t 508 nm with a Perkin-Elmer Model 552 dual-beam spectrophotometer equipped with a strip-chart recorder. Noise and drift over the analysis period were *0.004 absorbance unit (AU). Solutions of PAR and FA used as blanks typically had absorbances of 0.080-0.100 AU. Molar absorptivity of Ni(PAR)2 varied with pH and ionic strength, but was generally near 60000 M-’ cm-’. This value is slightly smaller than that reported at pH 9 (14). The rate constant of Ni(PAR)2 formation from aquo Ni2+varied from 0.25 s-l in 0.1 M NaClO, to 0.40 s-l in 0.002 M NaC104. The estimated error in Ni(PARI2determinations was 0.07 pM, or 1.3% relative error for a 5 pM solution. pH was measured with an Orion Model 601 digital voltmeter and an Orion Model 91-05 combination pH electrode, calibrated daily. Experiments were conducted at room temperature, 25 f 2 “ C . Procedure. The pH of a Ni(I1)-FA solution was adjusted with acid or base, and the solution was then equilibrated for 12-24 h in a covered, opaque Teflon beaker. If the pH drifted more than 0.1 pH unit during this period, it was readjusted to the desired value 1h before beginning the kinetic measurements. Kinetic measurements were performed in a 1-cm quartz cell. An Eppendorf pipet delivered 2.50 mL of Ni(I1)-FA solution into the cell. Data acquisition was initiated when a second pipet delivered 0.250 mL of 1.0 mM PAR to the cell. The initial time t o was known to f0.3 s, but mixing required 1 s, so data acquisition began at 2 s. Seventy data points were collected for each run a t unequal time spacings, beginning with 1-s intervals up to t = 9 s and increasing to 300-s intervals at t > 1200 s. Maximum absorbance of a kinetic run with total nickel concentration [NilT = 5 pM was typically 0.370-0.400 AU, and recovery of nickel was good (97-102%). Kinetic data discussed in this paper represent the average of two to four separate kinetic runs, which agreed within 2 % for all data
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Environ. Sci. Technol., Vol. 24, No. 4, 1990
points. Experiments with each set of reaction conditions were repeated with new solutions on separate days to compensate for day-to-day variations in instrument performance and solution quality. Data Analysis. All experimental data analysis was based on eq 1. A Laplace transform technique and four statistical approaches were used. The transform method assumes a very large n and produces a continuous concentration distribution. The statistical approaches produce small ( n < 10) sets of components, which are more convenient for calculations. For all sets of components, C1and k, will refer to the fastest dissociating component, usually aquo Ni2+and kinetically indistinguishable species. C, and h, will refer to the most slowly dissociating component. A. The kinetic spectrum technique expresses the data as a concentration density function H ( k ) (13). H ( k ) is equal to the difference between the first and second derivatives of [Ni-FA] with respect to In t and plotted versus log k. Peaks of a defined half-width appear a t log k corresponding to the kiof the solution components, and the peak areas correspond to the Ci.Kinetic spectra were calculated with a spreadsheet using digital filters and derivatives rather than the smoothing spline used by Olson and Shuman (13). B. Nonlinear regression. The kinetic spectrum was used to generate initial guesses for Ciand kifor one-, two-, and three-component systems. Nonlinear regression was used to refine these guesses and match the data. The “best” number of components was chosen based on the ratio of model error to estimated analytical error, which should be close to 1, and the magnitude of the relative standard deviations of the parameter estimates, which should be small. C. Nonlinear regression with component stripping. In this method, suggested by Lavigne et al. (6), the k, and C, of the slowest reacting component were estimated from the slope and intercept of a straight line fitted to a plot of log [Ni-FA] versus t at long times. The number of data points considered to be “at long times” was somewhat arbitrary and was usually 12-14 points (determined by requiring the linear r2 to be >0.99). The rate constant of the aquo nickel ion component (k,)was known from blank runs. These three parameters were used to reduce the number of variable parameters in the nonlinear regression. The “best”number of components was selected by using the criteria in B above. D. Discrete kinetic spectra. Discrete kinetic spectra are analogous to the Laplacian transform technique discussed in A, except that only a finite number of log k values are permitted instead of a continuous distribution. Discrete spectra were calculated by postulating a small (
