Environ. Scl. Technol. 1993, 27, 1408- 1414
Response of Copper(11)-Humic Acid Dissociation Kinetics to Factors Influencing Complex Stability and Macromolecular Conformation Andrew W. Rate' and Ronald G. McLaren
Department of Soil Science, P.O. Box 84, Lincoln University, Canterbury, New Zealand Roger S. Swm CSIRO Division of Soils, PMB 2, Glen Osmond, Adelaide, South Australia 5064
Metal-organic interactions, including complexation reactions of metal ions and humic substances, are increasingly becoming recognized as important factors in many natural systems. As a result, the chemical equilibria involved have been studied in great detail, and the complexity of the reactions has led to the formulation of numerous thermodynamic and/or mathematical models and to the methods for resolving physicochemical heterogeneity (1-4). The kinetics of metal-humic acid interactions, however, have received less attention than equilibrium measurements. In composite systems such as soils or natural waters, where many different types of reactive surfaces or compounds exist in two or more phases, individual slow reactions may limit subsequent processes (5). Biological availability of essential or toxic trace metal ions may be limited in many cases by complex dissociation rates rather than thermodynamic stability of complexes (1,6). Chemical and/or diffusion-controlled kinetics may be important in mass-transfer (for example, in solute leaching)or immobilization processes such asincorporation into mineral phases (7). It is considered by some authors (8)that chemical equilibrium is seldom attained in systems such as soils,and thus the usefulness of equilibrium models for predicting speciation may be limited unless models also incorporate the relevant kinetic data (9). Humic acids are known to be polydisperse macromolecules (10, 11), having many carboxylic acid functional
groups which are dissociated and, thus, negatively charged at pH values likely to be encountered in natural systems (12). The most widely accepted model for their tertiary structure is that of a linear macro-ion which coils randomly in time and space (11). The degree of expansion or contraction of solvated macro-ions is known to be sensitive to pH, ionicstrength, and extent of metalion complexation (11,13,14). Briefly, the effect of increasing solution pH is to cause greater acid functional group dissociation of humic acid macromolecules; the resulting increase in charge density expands the macromolecule via electrostatic repulsion. With increasing ionic strength, the macromolecule contracts due to shielding of macro-ionic charge by electrolyte counter-ions and consequent reduction of electrostatic repulsion between charged sites. With increasingdegree of metal ion complexation, macromolecular contraction may occur via charge neutralization due to covalent bonding and/or through intramolecular crosslinking caused by pseudochelation at two or more nonadjacent metal ion binding sites. In addition, cross-linking of separate macromolecules by metal complexation may induce intermolecular aggregation (15). It is hypothesized that if metal ion-humic substance complex dissociation is controlled by intraparticle diffusion, then factors which contract humic acid macromolecules (increasing ionic strength or degree of metal complexation; decreasing pH) will retard complex dissociation rates. Alternatively, it is possible that the dissociation rates of metal ion-humic substance complexes are related to their thermodynamic stability. Factors such as pH, ionic strength, or meta1:ligand ratio also affect ouerall metal ion-humic substance complex stability in the following ways. Increasing pH increases apparent complex stability by decreasing proton competition with metal ions for weakly acidic binding sites (16). Increasing ionic strength decreases overall complex stability by shielding macroionic charge, decreasing complexing metal ion activity in the domain of the macromolecule (14, 17). Increasing metal:humic substance ratio decreases overall complex stability, since metal ions bind to progressively weaker sites on heterogeneous ligands such as humic acids as metal loading increases (18). An alternative hypothesis may thus be stated, that if metal ion-humic substance complex dissociation rate decreases with increasing complex stability, then factors which increase average complex stability (increasing pH, decreasing ionic strength, or metal ion loading) should also lead to an observed decrease in rates of complex dissociation. The objectives of this study were 2-fold. First, to investigate the effects of factors such as pH, ionic strength, or metakhumic substance ratio on the dissociation kinetics
1408 Environ. Scl. Technol., Vol. 27, No. 7, 1993
0013-936X/93/0927-1406$04.00/0
The dissociation reactions of copper(I1)-humic acid complexes were investigated using a competing ligand spectrophotometric technique. The effects of pH, initial Cu: HA ratio, electrolyte concentration, and predissociation equilibration time on the rates of complex dissociation were studied and are discussed in terms of equilibrium speciation and macromolecular conformation. In general, increasing pH or predissociation equilibration time decreased Cu-HA complex dissociation rates; increasing Cu: HA ratio or electrolyte concentration increased Cu-HA dissociationrates. These trends are shown to be consistent with slower dissociation from more thermodynamically stable complexes. The response of the rate of these reactions to the different experimental conditions did not imply an effect on reaction rates due to the degree of macromolecular contraction or expansion. The effect of predissociation equilibration time was attributed to slow, complexation-induced conformational change which promoted formation of more stable complexes and retarded subsequent dissociation. Data analysis was based on proportions of Cu in operationally defined kinetic fractions and on a first-order exponential decay function modified by a log-normal distribution. Introduction
0 1993 American Chemical Society
Table I. Analytical Data for Summit Hill Humic Acid (SHA) and Waimari Peat Humic Acid (WHA) humic acid carbon (% ) hydrogen ( % ) nitrogen ( % ) ash ( % SHA 48.9 5.0 4.6 1.16 WHA 54.2 3.6 3.1 1.30
of copper(I1)-humic acid complexes. The second objective was to evaluate,using the same set of experiments,whether diffusion of dissociating metal ions through the domain defined by a solvated humic substance macro-ion (intraparticle diffusion) might be important in determining complex dissociation rates from Cu-HA complexes. The variable experimental conditions were expected to affect both the equilibrium speciation and the solution conformation of humic substances. Materials and Methods Experimental. Humic acids were extracted from Summit Hill soil, Banks Peninsula, New Zealand (SHA), and Waimari Peat, Canterbury, New Zealand (WHA), using the method adopted by the International Humic Substances Society (IHSS) for soil humic acids (19).The Summit Hill humicacid (SHA) is an IHSS reference humic substance. Elemental compositions and ash contents for both humic acids appear in Table I. 442-Pyridylazo)resorcinol (PAR, Sigma) was obtained as the free acid; stock solutions were prepared by dissolving PAR in an equimolar amount of NaOH solution. All other chemicals were of analytical grade; water was purified to a resistivity of 18.3 M Q cm. Dissociation of Cu-HA complexes was induced by reaction with an excess of the metallochromicligand PAR, and the formation of the CuPAR complex followed spectrophotometrically (510 nm) as a function of time as described by Shuman et al. (1). This method ensures pseudo-first-order dissociation of Cu-HA complexes (11. At copper:humic ratios such as those used in this work, the rate-determining step has been found to be release of Cu2+from the humic acid complex (6). All three solution conditions (pH, electrolyte concentration, initial metal: ligand ratio) were varied individually while holding the other two conditions constant. Solution conditions were constrained to values which maintained both the humic acid and PAR in true solution. In addition, measurements were made of complex dissociation after the Cu2+-HA solutions had equilibrated for different lengths of time (24-168 h). Solutions of Cu2+ and HA (as the sodium humate) at the desired pH, NaN03 concentration, and Cu:HA ratio were prepared and equilibrated in acidwashed plastic containters for the required length of time. This equilibrated Cu-HA solution was mixed with a PAR solution (0.5 mmol L-l) at identical pH and NaN03 concentration in a spectrophotometer cell (1 cm path length) using simultaneously depressed syringes and a T-junction mixing chamber. Kinetic data were acquired for 90 min using a Philips PU8700 spectrophotometer. Numerical Analysis. Simultaneous first-order reaction of a mixture of n Cu-HA complexes CuL1, ..., CuLn with PAR may be represented by: n
Z C u L , + nPAR FInCuPAR 1x1
n
+E L i
(1)
131
A model for describing heterogeneous kinetics, based on a log-normal digtribution of first-order rate constants (20),
has been developed and evaluated (21)for application to Cu2+-humic acid complex dissociation. This model replaces the summation of discrete sites found in eq 1with an integral over a continuum of sites described by a lognormal site density function. The formulationof the model for these experiments is summarized by:
exp[$(?)"]
(1- exp(-e't)) dK (2)
where K = In k, k being the first-order rate constant; [CuPARIt is the concentration of CUPAR at time t; [CuPARIo is the concentration of CuPAR at the experimental time zero; [CuHAlobsis the stoichiometric concentration of Cu-HA complexes; p is the mean of the distribution in In k;and u is the standard deviation of the distribution in In k. This model was fitted to kinetic data by optimizing values for [CuHA],b,, p, and u by nonlinear regression. Using this model, it was found (21) that (i) Cu-HA dissociation kinetics were predicted accurately; (ii) adjustable parameters ([CuHAI,b,, 1.1, and a) were sensitive to changes in experimental conditions, and (iii) changes in the adjustable parameters of the model under different conditions were consistent with real physicochemical processes. The nature of the experimentobserving Cu-HA complex dissociation allows the following method of data analysis. It was always observed that a proportion of Cu-HA complexes dissociated within mixing times; this was evident from nonzero CuPAR absorbance values at the first datum of kinetic runs. In addition, recovery of Cu from Cu-HA complexes was always incomplete, shown by an apparent equilibrium CuPAR absorbance at the end of experiments which was less than that calculated from the total amount of Cu in the system and the molar absorptivity of CuPAR (21). These observations allowed three operational fractions, based on kinetic behavior, to be defined for Cu-HA complex dissociation. The first, rapidly dissociating Cu (CUR),was that proportion of CuHA complexes dissociating within mixing times. This fraction also includes a small proportion of free Cu2+and copper(I1)in hydroxy and nitrate complexes. The second, slowly dissociating Cu (Cus), was the proportion of CuHA complexes whose dissociationreaction occurred during experimental observation times. Cus, being the experimentally observable kinetic fraction, is equivalent to [CuHAl,b, in eq 2. The third, nondissociating Cu (CUN), was comprised both of Cu in Cu-HA complexes which were more thermodynamically stable than CuPAR and Cu-HA complexes with very small dissociation rate constants (Figure 1). Calculations were made using literature formation constants for hydroxy and nitrate complexes of Cu2+(22) and ion-selective electrode measurements of free Cu2+ concentrations in humic acid solutions (23). These indicated that at the copper(I1) and humic acid concentrations used in this work, the only important nonhumic species were Cu(N03)- (ca. 1% 1 and Cu2+(ca. 3.2%), both at pH 5. At higher pH values, 199% of Cu2+was bound to humic acid. Tesh for statistical significance of differences between parameters derived from both methods of data analysis were based on one-way analysis of variance and were performed using MINITAB (24). The term 'significant', Environ. Sci. Technol., Vol. 27, No. 7, 1983 1409
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____________________________________ non-dissocciated (Cu,)
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rapidly-dissociated (Cu,)
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Flgure 1. Simulated CuPAR formation (absorbance versus time) curve R showing definition of operationally defined kinetic fractions. A ~ A is the absorbance of uncomplexed PAR; A. is the absorbance of CuPAR at the expermentaltime zero; A, is the maximum observed CuPAR absorbance;A,, is the maximum possible CuPAR absorbance, if all Cu2+ had reacted with CuPAR.
referring to differences in parameters derived from different experimental conditions, refers to a probability of less than 0.05 that mean parameter values from different treatments were equal, based on this analysis. Results and Discussion
All of the conditions varied in these experiments (pH, Cu:HA ratio electrolyte concentration, predissociationCuHA equilibration time) resulted in substantial changes in [CuPARl versus time profiles. Some representative data are presented in Figure 2. These changes are reflected in the parameters derived from the kinetic fraction (Table 11) and the log-normal distribution model (Table 111) methods of data analysis. A substantial proportion of CuSHA and Cu-WHA complexes were found to be rapidly dissociating under the various experimental conditions; 3 7 4 4 % of total solution Cu was present in Cu-HA complexes which dissociated within mixing times. Similar proportions of rapidly dissociating Cu, as defined here, have been reported for dissociation of Cu from complexes with humic materials (1,25). The amount of Cu in CuHA complexes with dissociation rates able to be observed directly by these experiments (Cud was in the range 4-45 % , as a proportion of total solution Cu. The nondissociating Cu fraction (CUN)comprised between 5 % and 43 % of total solution Cu. The variation in the proportions of these operationally defined kinetic fractions is considered to offer a simple and effective method of kinetic data analysis. The log-normal model (eq 2) provided excellent agreement between observed and predicted Cu-HA dissociation kinetics for the observable reactions, as found previously (22). Goodness of fit, as defined by R2 values, was in the range 99.35-99.98%. In addition, the adjustable parameters derived from fitting this model to kinetic data often proved to be more sensitive to changes in reaction conditions than the proportions of Cu in the kinetic fractions. It should be noted that the experimental window constrains observable rate constants to lie in the range -8 < In k < 1(k in the range 3.3 x lo4 to 2.7 s-l). This rate constant window may be obtained from the transformation In k = In ( 2 / t ) (25). Where we have presented In k distributions based on application of the log-normal model 1410 Envlron. Scl. Technol., Vol. 27,
No. 7, 1993
(Figure 4), these hypothetical distributions imply In k values outside this window. Exclusion of the portions of these In k distributions outside this range, however, does not affect the information regarding the mean of, and dispersion in, rate constants, or any conclusion that may be made from this method of data analysis. The effects of pH, Cu:HA ratio, electrolyte concentration, and Cu-HA predissociation equilibration time are presented below. pH. Increasing pH for the dissociation reaction, while holding other conditions constant, decreased the rate of dissociationof Cu-HA complexes for both SHA and WHA. Increasing pH caused a consistent increase in the proportion of Cus, indicating a higher proportion of slowly dissociating Cu-HA complexes at higher pH (Figure 3; Table 11). Changes in CURand CUNshowed no consistent trends, however, for either CuSHA or CuWHA dissociation. Analysis of the dissociation data in terms of the log-normal model (eq 2) showed a consistent increase in [CuHAI,b,, confirming the observed increase in Cus. Changes in p and u with changing pH showed less consistent trends; where these changes were statistically significant, an increase in pH resulted in overall decreases in p and increases in u (Figure 4; Table 111). This implies slower dissociation of Cu-HA complexes at higher pH. Similar responses to pH, in that the proportions of more slowly dissociating complexes increased with increasing pH, have been found for nickel(I1)-fulvic acid systems (2627). Copper(I1)-HA Ratio. An increase in initial Cu:HA ratio, by varying HA concentrations with all other conditions held constant, tended to increase the rate of dissociation of both Cu-SHA and Cu-WHA complexes. Analysis of data using kinetic fractions gave conflicting results for Cu-SHA; the amount of Cus showed a consistent and significant decrease (Figure 3; Table 11),but CURand CUNshowed no obvious trends. For Cu-WHA, however, a similar decrease in Cus was accompanied by an increase in CURand a decrease in CUN,which were both consistent and statistically significant (Table 11). The log-normal model, however, showed consistent and significant trends in all adjustable parameters with changing initial Cu:HA ratio (Figure 4; Table 111). Increasing Cu:HA ratio decreased [CuHA],b,, consistent with the observed decrease in CUS. The mean rate constant, p, increased with increasing Cu:HA ratio for both Cu-SHA and Cu-WHA, indicating faster dissociation. Fitted valuesof u decreased; the lower proportions of Cus ([CuHA],b,) contributing to less observed kinetic heterogeneity. Again, these results are in agreement with those found for Ni2+-fulvic acid (26, 27) and Cu2+-humic acid (6) complex dissociation, where dissociation rates increased with increasing metal: humic substance ratio. Electrolyte Concentration. Increasing NaN03 concentration, while holding all other reaction conditions constant, resulted in consistent and significant increases in the proportion of CUSand a decrease in the proportion of CUN. A small decrease was observed for CURwith increasing NaN03 concentration (Figure 3; Table 11; this decrease, however, was significant only for CU-WHA). Application of eq 2 to kinetic data for Cu-SHA and CuWHA dissociation revealed increases in [CuHAI,b,, consistent with the observed increase in Cus. The mean log rate constant, p, decreased with increasing NaNO3 concentration; the associated standard deviation, u, increased with increasing ionic strength (Figure 4; Table 111). These
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Flgure 2. CuPAR formation versus time (mean f standard error) for dissociation of (a) CU-SHA at pH 5.0 (0),pH 6.0 (O), and pH 7.0 (A):(b) CU-WHA at [HA] = 0.025 g L-i (0),0.05 g L-i (0),and 0.10 g L-i (A);(c) CU-SHA at pNa = 1.0 (0),2.0 (O), and 3.0 (A):and (d) CU-WHA after equilibration for 24 h (0)and 96 h (0). ([CUI = 7.87 pmol L-', [SHA] = 0.05 or [WHA] = 0.025 g L-I, pNa = 1.0, pH = 6.0 and Cu-HA equilibration time = 24 h unless otherwise indicated.)
Table 11. Proportions of Cu (Mean f Standard Error for Experimental Replicates) in Operationally Defined Kinetic Fractions.
kinetic fraction humic acid
PH
[HA1 (g L-9
PNa
PDET (h)
CUR (%)
cw (%)
CUN (%)
SHA SHA SHA SHA SHA SHA SHA SHA WHA WHA WHA WHA WHA WHA WHA WHA WHA WHA WHA Showing the effect of changing pH, Cu:HA ratio, NaNOs concentration, and predissociation equilibration time (PDET). [CUI is always 7.87 pmol L-1 for all experiments. 5.0 6.0 7.0 6.0 6.0 6.0 6.0 6.0 6.0 6.5 7.0 6.0 6.0 6.0 6.0 6.0 6.5 7.0 6.0
0.05 0.05 0.05 0.10 0.025 0.05 0.05 0.05 0.025 0.025 0.025 0.10 0.05 0.025 0.025 0.025 0.025 0.025 0.025
1.0 1.0 1.0 1.0 1.0 2.0 3.0 3.0 1.0 1.0 1.0 1.0 1.0 2.0 3.0 1.0 1.0 1.0 2.0
resulta imply that increasing ionic strength increases the rate of Cu-HA complex dissociation for both Cu-SHA and Cu-WHA. Cabaniss ( 2 3 , investigating Ni2+-FA complexes, also found that increasing the ionic strength increased the overall rate of complex dissociation. It was proposed (27)that only the more rapidly dissociating Ni2+FA complexes responded to changes in electrolyte concentration. In this work, the increase in the proportion
24 24 24 24 24 24 24 96 24 24 24 24 24 24 24 96 168 144 72
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of Cus at the expense Of CUNsuggests an effect on reaction rates for the more slowly dissociating Cu-HA complexes. Predissociation Equilibration Time. Increases in Cu-HA complex equilibration time before reaction with PAR decreased the proportions of CURand increased the proportions of Cus and CUN(Figure 3; Table 11). In a few cases, however, these changes were small and not statistically significant. The parameters generated by fitting Envlron. Sci. Technol., Vol.
27, No. 7, 1993 1411
Table 111. Refined Adjustable Parameters (Mean f Standard Error for Experimental Replicates) Derived from Fitting a Log-Normal Distribution of First-Order Rate Constants (ea 2) to Cu-HA Dissociation Data*
humic acid
pH
[HA] (g L-I)
pNa
PDET (h)
[CUHA],~(pmol L-l)
parameter mean In k ( p )
SHA 5.0 0.05 1.0 24 1.05 f 0.02 -4.43 f 0.05 SHA 6.0 0.05 1.0 24 2.15 f 0.06 -4.6 f 0.2 SHA 7.0 0.05 1.0 24 3.4f 0.1 -5.37 f 0.06 SHA 6.0 0.10 1.0 24 3.00 f 0.08 -5.6 f 0.2 SHA 6.0 0.025 1.0 24 1.52 f 0.05 -3.88 f 0.07 SHA 6.0 0.05 2.0 24 1.07 f 0.01 -4.05 f 0.02 SHA 6.0 0.05 3.0 24 0.64 f 0.02 -3.75 f 0.06 SHA 6.0 0.05 3.0 96 0.90 f 0.03 -4.4 f 0.1 WHA 6.0 0.025 1.0 24 1.24 f 0.02 -3.75 f 0.07 WHA 6.5 0.025 1.0 24 1.64 f 0.03 -4.16 f 0.02 WHA 7.0 0.025 1.0 24 1.73 f 0.07 -3.98 f 0.02 WHA 6.0 0.10 1.0 24 2.88 f 0.03 -5.85 f 0.04 WHA 6.0 0.05 1.0 24 2.15 f 0.09 -5.09 f 0.08 WHA 6.0 0.025 2.0 24 0.74 f 0.04 -3.8 f 0.1 WHA 6.0 0.025 3.0 24 0.29i 0.02 -3.6 f 0.3 WHA 6.0 0.025 1.0 96 1.42f 0.05 -4.17 f 0.01 WHA 6.5 0.025 1.0 168 1.83 f 0.02 -4.86 f 0.07 WHA 7.0 0.025 1.0 144 1.69f 0.05 -4.25 f 0.08 WHA 6.0 0.205 2.0 72 0.68 f 0.09 -4.2 f 0.3 a Showing the effect of changing pH, Cu:HA ratio, NaN03 concentration, and predissociation equilibration time (PDET). 7.87 pmol L-l for all experiments.
the log-normal model to these kinetic data were much more sensitive to the changes in Cu-HA complex dissociation rate. Examination of these derived parameters showed consistent increases in [CuHA],b,, decreases in p , and increases in a (Figure 4; Table 111). These changes were statistically significant more often than those derived from kinetic fraction analysis but were still not significant in a few cases even though the direction of change in the adjustable parameters ([CuHA],b,, p, and a) was the same for each experiment. This is considered to result from the small size of the effects observed. These observations are indicative of an overall decrease in the rate of Cu-HA complex dissociation with increasing reaction time for the initial Cu-HA system, which is supported by inspection of the raw data (e.g., Figure Id). The rate of dissociation of Cu-HA complexes thus appears to be retarded by increasing pH, decreasing Cu: HA ratio, decreasing ionic strength, and longer predissociation equilibration time for the Cu-HA complex. The first three factors (increasing pH, decreasing Cu: HA ratio, and decreasing ionic strength) were expected to induce a more expanded conformation for the solvated humate macro-ion and to facilitate faster dissociation, if the rate of metal ion diffusion through the macromolecular domain was rate-determining. Increasing pH, decreasing Cu:HA ratio, and decreasing ionic strength would all tend to increase overall Cu-HA complex stability, as discussed above. The response of Cu-HA dissociation rates to these factors therefore suggests that thermodynamic complex stability is a more important determinant of Cu-HA dissociationkinetics than any potential diffusiveprocesses. Since factors which increase overall Cu-HA complex stability also decrease Cu-HA complex dissociation rates, then it appears that more thermodynamically stable CUHA complexes also dissociate more slowly. Average molecular weights of approximately 50 000have been presented (28) for soil and peat humic acids. Molecular weights have not been definitively determined for SHA or WHA. The observation, however, that a substantial proportion of SHA elutes near the exclusion limit of Sephadex G-150 chromatography columns (29) supports the premise that these materials have similarly 1412 Envlron. Scl. Technol., Vol. 27, No. 7, 1993
SD In k (a) 2.15f 0.08 2.05 f 0.09 2.41f 0.02 2.54 f 0.06 1.82 f 0.03 1.89 f 0.02 1.68 f 0.06 2.14 f 0.06 1.72 f 0.04 2.02 f 0.02 1.79f 0.05 2.64 f 0.03 2.39 f 0.09 1.8 f 0.1 1.7 f 0.3 1.91 f 0.02 2.41 f 0.06 2.0 f 0.1 2.1 f 0.4
[CUIis always
high molecular weights. This would make intraparticle diffusion a definite possibility. The observation that intraparticle diffusion appears to be less important than complex stability in determining rates of aqueous Cu-HA complex dissociation has several possible explanations. It may be that the time scale of conformational change for randomly coiling humic acid macromolecules is short compared with the rate of metal ion release. In this case, it is envisaged that individual Cu2+ binding sites spend a proportion of time on the ‘exterior’ of the random coil, making diffusionunnecessary. It is also possible that none of the conditions which prevailed in this set of experiments induced a sufficiently contracted conformation for the Cu-HA complexes for diffusiveprocesses to become rate-limiting for dissociation. We consider that the retardation in Cu-HA complex dissociation rates caused by longer predissociation equilibration times may be explained by the proposal that the Cu-HA complex formation reaction involves some very slow processes. Lehmann and Harter (30) found similar behavior for a Cu2+-soil system. They proposed that the observed retardation of Cu2+desorption rate from soil, with increasing Cu2+-soilreaction time, could be explained by slow movement of Cu2+into less labile or less sterically accessible sites, formation of a separate copper(I1)containing solid phase, or chemical alteration of binding sites after prolonged contact with Cu2+. This work shows that complex stability is rate-determining for dissociation of Cu-HA complexes; therefore, it is most likely that prolonged reaction of Cu2+with humic acid allows Cu2+ to be complexed by more thermodynamically stable sites than is possible for a reaction with a shorter time scale. In addition, long-term reaction of Cu2+ with HA may induce conformational changes in the Cu-HA complex, such that steric accessibility does become important in determining dissociation reaction rates. The evidence presented here for a very slow component in aqueous-phase Cu-HA complex formation reactions has serious implications for experiments designed to observe metal-humic substance complex equilibria. The effect of slow complexation processes will be more critical for reactions forming very stable complexes. For example, if
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24 96 Pre-reaction equllibratlon time / h Flgure 3. Proportions of Cu in kinetic fractions, showlng effect of (a) pH for CU-SHA dissociation; (b) Cu:HA ratio for CU-WHA dissoclation; (c) [NaN03]for Cu-SHA dissocition, and(d) predissodationequlllbration time for CU-WHA dissociations. CURis rapidly dissociating Cu; Cus Is slowly dlssoclatlng Cu; CUNIs nondissociatingCu. ([Cu] = 7.87 pmol L-l, [SHA] = 0.05 or [WHA] = 0.025 g L-', pNa = 1.0, pH = 6.0, and Cu-HA equilibration time = 24 h unless otherwise indicated.)
-12
-10
-8
-6
-4
In k Flgure 4. Hypothetical log-normal dlstrlbutions of firstorder rate constants for slowly dissociated CU-SHA complexes showing effect of (a)pHforCuSHA; (b)Cu:HAratloforCuWHA;(c)NaNO3concentratlon for CuSHA. and (d) predissociatlonequiilbratlontlme for CuWHA. ([Cu] = 7.87 pmoi L-I, [SHA] = 0.05 or [WHA] = 0.025 g L-I, pNa = 1.0, pH = 6.0, and CU-HA equliibratlon time = 24 h unless otherwise Indicated.)
Environ. Sci. Technol., Vol. 27, No. 7, 1993 1413
99.9% of a metal ion is bound to a humic acid (not an unreasonable figure for Cu2+),and further slow reaction allows the proportion of bound Cu2+to increase by 0.05 % to 99.95% ,then the concentration of unbound Cu2+would decrease by 50%. In contrast, for a more weakly bound metal ion where the proportion bound is 90% (e.g., Cd2+), a similar 0.05% increase in bound metal would decrease the unbound metal ion concentration by only 4.5 % . This work implies that aqueous-phase Cu-HA complex formation reactions are still proceeding after at least 24 h. A typical experiment, involving continuous titration of a humic acid with metal ion a t constant pH, would be unlikely to allow 24 h between incremental metal ion additions. These experiments have shown that relatively small changes in pH, copper:humic acid ratio, ionic strength, and length of copper-humic acid reaction time can have a marked effect on the dissociation kinetics of copper(11)-humic acid complexes. The fact that experiments were performed in the aqueous phase, at relatively high humic acid and copper concentrations, means that this work is most applicable to processes occurring in the soil solution, especially in peats and soil horizons with higher organic content. The effects of changing experimental conditions are particularly relevant for cases of soil contamination or for soils undergoing acidification or salinity changes.
Literature Cited (1) Shuman, M. S.;Collins, B. J.;Fitzgerald, P. J.; Olson, D. L. In Aquatic and Terrestrial Humic Materials, 1st ed.; Christman, R. F., Gjessing, E. T., Eds.; Ann Arbor Science: Ann Arbor, MI, 1983;pp 349-370. (2)Turner, D. R.; Varney, M. S.; Whitefield, M.; Mantoura, R. F. C.; Riley, J. P. Geochim. Cosmochim.Acta 1986,50,289297. (3) Sposito, G. Crit. Rev. Environ. Control 1986,16,193-228. (4) Nederlof, M. M.; van Riemsdijk, W. H.; Koopal, L. K. J. Colloid Interface Sci. 1990,135,410-426. (5) Morgan, J.J.;Stone, A. T. In Chemical Processes in Lakes, 1st ed.; Stumm, W., Ed.; Wiley-Interscience: New York, 1985; pp 389-426. (6) Hering, J. G.; Morel, F. M. M. Enuiron. Sci. Technol. 1990, 24, 242-252. (7) van der Zee, S. E. A. T. M.; van Riemsdijk, W. H.; van Grimsven, J. J. M. Neth. J . Agric. Sci. 1989,37,47-60. (8) Sparks, D. L. In Advances in Agronomy; Academic Press: New York, 1985;Vol. 38,pp 231-266.
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(9) Wu, S.-C.; Gschwend, P. M. Enuiron. Sci. Technol. 1986, 20,717-725. (10) Hayes, M.H. B.; Swift, R. S. In The Chemistry of Soil Constituents, 1st ed.; Greenland, D. J., Hayes, M. H. B., Eds.; Wiley: Chichester, 1978,pp 179-320. (11) Swift, R. S. In Humic Substances II: The Search for Structure, 1sted.; Hayes, M. H. B., MacCarthy,P.,Malcolm, R. L., Swift, R. S., Eds.; Wiley: Chichester, 1989;pp 449465. (12)Perdue, E. M. In Humic Substances in Soil, Sediment and Water, 1st ed.; Aiken, G. R., McKnight, D. M., Wershaw, R. L., MacCarthy, P., Eds.; Wiley-Interscience: New York, 1985;pp 493-526. (13) Marinsky, J. A.; Ephraim, J. Environ. Sci. Technol. 1986, 20,349-354. (14) Morawetz, H. Macromolecules in Solution, 2nd Ed.; Wiley-Interscience: New York, 1975;pp 344-396. (15)Teasdale, R. D. J. Soil Sci. 1987,38,433-442. (16)Kerndorff, H.; Schnitzer, M. Geochim. Cosmochim. Acta 1980,44,1701-1708. (17) Ephraim, J.; Marinsky, J. A. Environ. Sci. Technol. 1986, 20,367-376. (18) Gamble, D. S.;Underdown, A. W.; Langford, C. H. Anal. Chem. 1980,52,1901-1908. (19) Swift, R. S.;Delisle, G.; Leonard, R. L. Sci. Total Environ. 1987,62,423-430. (20) Albery, W. J.; Bartlett, P. N.; Wilde, C. P.; Darwent, J. R. J. Am. Chem. SOC.1985,107,1854-1858. (21) Rate, A. W.; McLaren, R. G.; Swift, R. S. Enuiron. Sci. Technol. 1992,26,2477-2483. (22) Lindsay, W. L. Chemical Equilibria in Soils, 1st ed.; Wiley: New York, 1979;pp 224-225. (23) Rate, A. W. Ph.D. Thesis, Lincoln University,New Zealand, 1990. (24)Minitab Reference Manual Release 8; Minitab Inc.: State College, PA, 1991;Chapter 8. (25) Olson, D. L.; Shuman, M. S. Anal. Chem. 1983,55,11031107. (26) Lavigne,J. A,; Langford, C. H.; Mak, M. K. S. Anal. Chem. 1987,59,2616-2620. (27) Cabaniss, S.E. Environ. Sci. Technol. 1990,24,583-588. (28) Stevenson, F. J. Humus Chemistry 1st ed.; Wiley: New York, 1982;pp 285-308. (29) Leonard, R. L. Lincoln University, New Zealand, personal communication, 1989. (30) Lehmann, R. G.; Harter, R. D. Soil Sci. SOC.Am. J . 1984, 48,769-772. Received for review November 10, 1992.Revised manuscript received March 9,1993.Accepted March 18,1993.
