ANALYTICAL CHEMISTRY, VOL. 50, NO. 12, OCTOBER 1978 (13) M. Shinn, Ind. Eng. Chem., Anal. Ed., 13, 33 (1941). (14) G. Macchi and B. Cescon, Anal. Chem., 42, 1809 (1970). (15) M. Nishimura, K. Matsunaga, and K. Matsuda, BunsekiKagaku, 19, 1096 (1970). (17) A. Allen, J . Chem. Soc., 1968 (1954). (18) W. Fischer, 2.Phys. Chem., 65, 61 (1908). (19) T. Turney and G. Wright, Chem. Rev., 59, 497 (1959). (20) F. Helfferich, "Ion Exchange", McGraw-Hill, New York, N.Y., 1962.
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(21) "New Orleans Area Water Supply Study", US. Environmental Protection Agency, Dallas, Texas, 1974.
RECEIVED for review June 16, 1978. Accepted July 25, 1978. Work supported in part by NIH Grant CA-09242 awarded by the National Cancer Institute (PHS, DHEW).
Stability Constants for the Complexation of Copper(I1) Ions with Water and Soil Fulvic Acids Measured by an Ion Selective Electrode William T. Bresnahan,' Clarence L. Grant, and James H. Weber* Department of Chemistry, Parsons Hall, University of New Hampshire, Durham, New Hampshire 03824
To measure metal-humic matter interactions in natural waters, it is necessary to: (1) know and define the concentration of humic matter in the system; (2) isolate and characterize this mixture of compounds; (3) propose model compounds based upon the chemical information obtained in characterization; and (4) evaluate and compare the complexation-chemistry of the natural organic matter and the model compounds. It is the aim of this research to focus upon the fourth point, specifically to measure conditional stability constants between Cu2+and well-characterized water fulvic acid (WFA) and soil fulvic acid (SFA). Reuter and Perdue ( 4 ) recently reviewed the importance of heavy metal-organic matter interactions in natural waters. The review includes a discussion of the advantages and limitations of several analytical techniques that have been used for stability constant determinations. References not in the review include potentiometric titration (5),equilibrium dialysis (6), and ion selective electrode (ISE) techniques (7). Differential pulse polarography (DPP) and differential pulse anodic stripping voltammetry (DPASV) have also been used to determine humic matter-metal ion stability constants (8-10). Several groups (8-12) have emphasized the difficulty of doing D P P and DPASV measurements of free metal ion concentration in the presence of humic matter and other surfactants. For this reason we compared the results of the titration of Cu2+by fulvic acid (FA) using both the DPP and ISE techniques. Because we found the former method to be unreliable in our system, all the useful conditional stability constants originated from the ISE experiments.
We discuss the coordination of Cu2+to welltharacterized soil fulvic acid (SFA) and water fulvic acid (WFA) samples as studied by electron paramagnetic resonance (EPR) spectroscopy, differential pulse polarography (DPP), and a Cu2+ ion selective electrode (ISE). Identical titrations of Cu2+with SFA using DPP and ISE techniques showed that the DPP peak current was not proportional to the hydrated Cu2+ concentration. Cu2+ ISE experiments were done at pH 4.0, 5.0, and 6.0 for both SFA and WFA and also at pH 4.7 for WFA, and conditional stability constants were calculated by two methods. The more useful method-Scatchard plots constructed for each experiment-indicated the presence of two classes of binding sites (confirmed by EPR) with stability constants in all cases of about 1 X 10' and 8 X lo3. From pH 4.0 to 6.0, the total number of binding sites per molecule increased from 0.8 to 4.2 and 0.6 to 2.6 for the soil and water fulvic acids, respectively; most of the increase occurred between pH 5.0 and 6.0. We speculate that the difference in the two classes of sites is not due to the type of donor atoms, but to the geometry of the site.
The speciation of trace metals in natural waters together with their overall concentration has an important effect on the quality of our water resources. The physiochemical state of trace metals in water systems affects their availability to biota, their transport, and their accumulation in sediments. Particularly important is the extent of complexation of trace metals by humic and fulvic acids found in natural waters. The stabilities of metal-organic complexes in natural waters are greater than those of the corresponding inorganic metal complexes. Unfortunately most thermodynamic models have been developed to include the behavior of only inorganic ligands (I),or of inorganic and synthetic organic ligands ( 2 , 3 ) . I t is often assumed that because of the low concentration of the organic matter or the relatively high concentration of major cationic species, the chelation of trace metals is insignificant. In reality, the most important reason humic matter has been omitted from thermodynamic models is that these molecules are chemically very complex and not yet well understood ( 4 ) .
EXPERIMENTAL Apparatus. The differential pulse polarograms were obtained with a Princeton Applied Research (PAR) model 174 Polarographic Analyzer, PAR model 315 Automated Electroanalysis Controller, PAR model 172A Mercury Drop Timer and recorded with a Houston Omnigraph 2000 XY recorder. A platinum wire counter electrode and PAR 9331 saturated calomel reference electrode with salt bridge were used in a PAR 9301 cell fitted with a glass water jacket. Dissolved oxygen was removed from solutions by bubbling nitrogen through the cell for at least 10 min and passing it over the solution during electrolysis. All experiments were performed at 25.0 OC. An Orion model 94-29A Cu2+ISE, the reference electrode of a Corning model 476050 combination electrode, and an Orion model 701 digital pH meter were used for all ISE experiments. The PAR 9301 cell was used to hold the test solutions; again all
'Present address,Department of Chemistry,University of Michigan, Ann Arbor, Mich. 48104. 0003-2700/78/0350-1675$01,00/0
C
1978 American Chemical Society
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ANALYTICAL CHEMISTRY, VOL. 50, NO. 12, OCTOBER 1978
experiments were conducted at 25.0 "C. The cell was covered with aluminum foil to eliminate any slight response of the ISE to ambient light fluctuations. The crystal membrane was polished with a fine alumina paper before each day of use. Electron paramagnetic resonance (EPR) spectra were obtained with a Varian E-4 spectrometer operating at X-band with a 100-kHz magnetic field modulation. The conditions used in obtaining the frozen spectra and for simulation were microwave frequency 9.103 GHz, modulation amplitude 6.3 G, microwave power 5.0 mW, time constant 1.0 s, and scan rate 4000 G/h. Reagents. The isolation and characterization of the SFA and WFA have been described elsewhere (13-15). The FA solids were dried at 100 "C for 2 h before use. FA solutions were adjusted to the pH of the experiment and dispensed with a Gilmont S-3200A Ultra-precision 2.50-mL buret. Stock solutions of Cu2+were prepared by dilution of aqueous 0.1000 f 0.0005 M Cu(NOJ2 and dispensed with 100 and 200 ru.L Oxford Sampler Micropipetting Systems. Procedures. DPP experiments for the determination of conditional stability constants were performed by means of a titration of a Cu2+solution with a FA solution. To the cell was added 25.0 mL of 0.100 M KN03 and an aliquot of a standard Cu2+solution. Two successive replications of the polarogram of this solution were recorded. After an aliquot of FA was added to the cell, the solution was mixed by purging with nitrogen and the polarogram recorded in duplicate. This addition and mixing procedure was repeated for each data point. The pH of the test solution, which was checked before and after the titration, varied less than 0.1 unit. Experiments with the Cu2' ISE were carried out in much the same manner as those with DPP. A 25.0-mL aliquot of 0.100 M KN03 was added to the cell by pipet, and a calibration curve in the 8.0 X to 2.0 X M Cu2+concentration range was determined. Aliquots of a FA solution were added to the solution from which calibration curve data were obtained. The pH was continuously monitored, and slight adjustments were made with a weak solution of KOH in 0.100 M KNOB,dispensed with a Gilson P200 Pipetman variable micropipet. Since the total volume of the adjustments never exceeded 0.1 mL (0.4% of the solution volume), no corrections for dilution were necessary. Typically, 25 aliquots of FA were added during a titration covering a M to 1.5 X M based concentration range of about 1 X on the FA dissociation-corrected number-average molecular weights (14). After the titration, the cell was cleaned and a duplicate standard calibration curve was obtained. To directly study the effect of pH on the chelation of Cu2' to M SFA, 1.96 X M Cu2+was SFA, a solution 5.80 X prepared in 0.100 M KNO, in the PAR cell. The pH of this solution was increased stepwise from about pH 3.5 to 6.0 and then decreased through this same range, while the aquo Cu2' concentration and pH were monitored. Twenty-three 2- t o 4-pL aliquots of a KOH and a HN03 solution were used to adjust the pH. Since the total volume of these adjustments was less than 0.3% of the solution volume, no corrections for dilution were made. A similar experiment was performed with a solution of 2.99 X M SFA and 3.96 X M Cuz+. Calculations. We calculated the conditional stability constants in two ways. In the Scatchard treatment (26),the overall conditional stability constant, KO,can be written for the reaction Cu
(CuFA)
+ FA + CuFA, K O= [Cu](FA
-
CUFA)
(1)
in which charges are omitted for simplicity. Here [CUI is the concentration of hydrated Cu2+,CuFA is the number of moles of complex formed and FA is the total number of moles of fulvic acid in solution. If a particular class of binding sites on the fulvic acid molecule is denoted by i, Ki is the conditional stability constant for that class of sites, and ni is the average number of such sites per molecule of fulvic acid, the expression for Ki has the following form
K. =
'
(CuFA) [Cu]((niFA) - CUFA)
Upon rearrangement, the expression becomes
(2)
where v = (CuFA)/(FA). We calculated moles of FA from tke dissociation-corrected number-average molecular weights of 626 and 644 for WFA and SFA, respectively (14). We determined the [Cu2+]by the ISE, and CuFA by mass balance,
Tc,, = [CU"] t [CUFA]
(3)
where Tcu is the total concentration of Cu". Using values of D obtained from experiments carried out over a wide range of Cu2' concentrations, a plot of o/[Cu2'] vs. P is made. If there is only one class of sites, the plot will yield a straight = line such that when P = 0, 2/[Cuz'] = K,n,. When P/[CU~+] 0, D = n, and the slope of the line is -K,. If there is more than one class of sites, the plot will yield a curve which is concave-up. Then it is necessary to correct the data for the contribution of the first class of sites in order to obtain D values reflecting only the binding of the second class of sites (17). This correction is accomplished by calculating D', where D' = D - nl, nl being the total number of sites of the first type per molecule. If only two classes of sites are present, a plot of D' vs. p'/[Cu] will yield a straight line such that when P'/[CU~+] = 0, P' = n2 and the slope of the line is -K2. The second calculation of conditional stability constants is a method by Cheam (18) and is based on Cu2+complexing with a monoanionic FA species (HL-).
Cu2+ + HL-
CuL
+ H+,K
=
[CULI [H+I
(4)
[Cu2'][HL-]
The K calculation assumes 1:l complexes and a knowledge of the binding site concentrations, which we obtained from functional group analyses (13, 14) by assuming the concentration was one half of the total acidity. The [Cu2+]and [H+]were measured directly, and the [CuL] and the [CuL] and [HL-] were obtained from mass balance relationships.
+ [CUL] TFA = [HL-] + [CUL]
Tcu = [CUZ']
(5)
(6)
Within the pH range studied the [CuOH+],[H,L], and [L2-]were negligible and could be ignored. K was calculated as a function of pH and of [CuL]/TFA.All calculations were done by computer programs written for the DEC system 10.
RESULT AND DISCUSSION Comparative DPP and ISE Experiments. A comparison of ISE and DPP data from experiments done under identical conditions shows that the DPP peak current is not solely due to diffusion of the hydrated Cu2+ion (Figure 1). If the current were proportional to hydrated Cu2+,the ISE and DPP plots would be identical. The discrepancy could be due to many reasons involving the polarographic behavior of hydrated Cu2+ and its complexes or the interference of organic matter (8, 9, 12). For this discussion, the important point is that the data are not amenable to a modified Lingane analysis (8) for the determination to stability constants, nor can currents be attributed unequivocally to particular species in solution. For these reasons, we used the Cu2' ISE for a determination of conditional stability constants and binding sites of both FA samples. ISE Determination of Conditional Stability Constants and Numbers of Binding Sites. The binding of Cu2+ions of SFA and WFA was studied a t p H 4.0, 5.0, and 6.0 for both acids and also at pH 4.7 for WFA. The results of all titrations were treated by the Scatchard ( 1 6 , 1 7 ) and Cheam (18) methods of computation. After consideration of Cheam's method (1:1 complexes assumed), we will discuss the Scatchard method, which makes no assumptions of stoichiometry. K as defined by Equation 4 assumes 1:l complexes and thus a knowledge of the concentrations of binding sites. We obtained this information from functional group analyses (13,
ANALYTICAL CHEMISTRY, VOL. 50, NO. 12, OCTOBER 1978
24 0 0l 9 o g
‘“i 07
O
20.0
I i
1677
0
0 0
8 A
O
7P
0
0 0
1601
X
0
I
0 0
0 0 0 0 0
I
O0 0 0
40t
a
0 3 t
0
0
040
-v
A 020
A
060
080
Flgure 3. Scatchard plot for the titration of Cu2+ with soil fulvic acid at pH 5.0. Data are from four experiments at different Cu2+concentrations Figure 1. Comparison of differential pulse polarography ( i F / i p o(0)) and ion selective electrode ([Cu2+]/Tc,(A)) for identical titrations of CuZt with soil fulvic acid
Table I. Conditional Stability Constants and Numbers of Binding Sites in the Copper-Soil and -Water Fulvic Acid Systems FA sample
I60
soil soil soil
1401
water water water water
1200
100-
Kl x
lo-’
n,
4.0
4
5.0
10
6.0 4.0 4.7
20 3 10
0.2 0.3 0.7 0.2 0.2
5.0 6.0
9 13
0.2 0.6
pH
K, x 9 12 6 10 7 5 7
n,
n,
+
n,
0.8
0.6 0.7 3.5 0.4 0.6
4.2 0.6 0.8
1.0
1.2
2.0
2.6
1.0
C IY
80-
l
0
o
60-
0
401
0
0
I
L
010
020
[CYLI
030
040
050
TFA
Figure 2. Plot of Rvs. [CuL]/T‘,, for the titration ofCu2+with water fulvic acid at pH 6.0 14) and graphed K against [CuL]/TFA.A typical plot for WFA a t p H 6 is shown in Figure 2. All the graphs have a similar shape including an apparent [CuL]/TF.k break for both SFA and WFA of 0.2 a t pH 6 and ca. 0.1 at pH 4 and 5. After the break, in all cases K is relatively insensitive to [CuL]/TFA. The K values are similar for both acids between pH 4.0 and 6.0. The K values a t pH 6.0 are much greater than the values between p H 4.0 and 5.0 which suggests a significant change in binding a t pH 6. This observation is confirmed in the Scatchard treatment of the data. A major value of the Figure 2 plot is to emphasize strongly that measured “stability constants” are actually conditional stability constants that vary as a function of p H (fraction of FA dissociated) and of metal/ligand ratio. Scatchard plots were also done for all the titrations. Because of the nature of the Scatchard calculations and the interdependence of the variables, no error analysis was
performed. Figure 3, which is for the titration of Cu2+with SFA a t p H 5.0, includes data pooled from four titration experiments with four different Cu2+concentrations and is presented to illustrate the reproducibility. The Scatchard plots for all Cu2+-FA experiments are curved concave-up indicating the presence of more than one class of binding sites. The slope of the steep portion of the curve, -K1 and the intercept of this line a t the D axis, nl was estimated by regression analysis. From Figure 3, for example, K , = 1.0 x lo6 and nl = 0.3. When the contribution of the binding of the stronger (i = 1) class of sites has been removed, D‘ = D ni. A plot of p’/[Cu2+] vs. D’ for the Figure 3 data gives a straight line indicating that no further classes of sites are present. The slope and intercept of the graph show that K 2 = 1.2 X lo4 and n2 = 0.7. Similarly the conditional stability constants and number of sites for all experiments were calculated with the aid of regression analysis (Table I). Both the nl and n2 sites represent a variety of chemically similar metal complexes. Gamble (1g21) in thorough studies of FA polyelectrolytes reported that (a) most of the coordinating ability is due to carboxyl, phenol, or both functional groups and (b) because of the irregular nature of the polymer probably, no two carboxyl (or phenol) groups are chemically identical. Therefore, unequivocal interpretation of the Scatchard plots is difficult. The manner in which slopes were assigned to the curves also introduced error which would be reflected in both Ki and ni, since the slope and intercept are interdependent. The data for the first site of Cu-SFA a t pH 6.0 will serve as an example. When a line is fit to the first seven points, K1 = 1.3 X lo6 and nl = 0.86. When only the first five points are used (a better fit), K1 = 2.0 X lo6 and nl = 0.73. In view of the complexity of the system and the nature
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ANALYTICAL CHEMISTRY, VOL. 50, NO. 12, OCTOBER 1978
0
n
40 -
SIMULATED
* 0
-X
30-
8 0
EXPERIMENTAL
A 0
A
A0 0
35
40
45
50
55
60
PH
Flgure 4. Plot of B/[Cu2+] vs. pH for a solution of 5.80 X M soil fulvic acid and 1.96 X lo-' M Cu2+, showing data for increasing pH (O), and decreasing pH (A)
of the calculations, the accuracy of the complexation parameters in Table I is acceptable. Over the 4 to 6 p H range studied, K , values for both SFA and WFA average about 8 x lo3 and do not show any systematic trend with change in pH. Kl values increase systematically with increasing pH. The number of binding sites per molecule, both nl and n2,also increase with increasing pH. If each class of chelating sites were truly one type of acidic group, a 10-fold increase of stability constant would be expected for each increase of 1.0 pH unit. The results presented here clearly do not follow that prediction. A plot of ij/[Cu2+]vs. p H for a solution of 5.80 x M SFA and 1.96 X M Cu2+ (Figure 4) illustrates the significant increase in binding above pH 5.0, in agreement with the Scatchard results (Table I). The chelation is reversible with respect to pH for these concentrations as shown by the coincidence of points obtained when the pH was first increased from 3.5 to 6.0 and then decreased to 3.5 again. A similar experiment was performed with a solution of 2.99 X M SFA and 3.96 X M Cu2+. Here the chelation was not immediately reversible with respect to pH, that is, values of ij/[Cu2+] were greater while the pH was decreased. Intermolecular FA interactions, which would be more important a t the higher FA concentration, may be responsible for this behavior. Since both nl and n2 increased markedly between p H 5.0 and 6.0, the increased binding is common to some structures of both classes of chelating sites (see below). Nature of the Binding Sites. A room temperature EPR spectrum of a solution of 1.0 X M Cu2+and 7.0 X M SFA in 0.100 M KNO, a t pH 5.0 consisted of one broad, asymmetric band. Because freely tumbling Cu2+complexes typically give four-line room temperature spectra, the EPR results suggest that the Cu2+complex has a molecular weight of a t least 1000. Since the dissociation-corrected numberaverage molecular weight of SFA is 644, a C U ( F A )complex ~ under the conditions of excess SFA would have an appropriate molecular weight. EPR spectra were recorded at 77 K for the Cu-SFA system a t Cu/SFA = 0.03, 0.14, 0.45, and 0.71. An experimental spectrum of the 0.03 Cu/SFA ratio and a computer simulation are given in Figure 5. The spectra of the 0.03 and 0.14 ratio solutions were similar, but the 0.45 and 0.71 ratios showed only one broad asymmetric peak due to the presence of significant amounts of hydrated Cu2+(measured by the ISE). The physical parameters of the simulation of the EPR spectrum of the Cu/SFA = 0.03 solution are given in Table 11. Two binding sites in a 813 ratio are necessary to approximate the experimental spectrum. In addition the A
Y 2600
3000
2800
3200
GAUSS
Figure 5. Ex erimental and simulated EPR spectra of solution of 1.O X M CuR and 3.4 X IO', soil fulvic acid in 0.100 M KNO, at pH 5.0 and 77 K
Table 11. Parameters for the Simulated First-Derivative X-Band Spectrum of a Copper-Soil Fulvic Acid Complex at pH 5.0 and 77 K site weight 8
A II
A53 cm gL
IO4
I
cm" x Wl"
gll
lo4
WII"
2.067 8.00 15.0 2.310 179.0 20.0 2.058 11.00 15.0 2.297 185.0 20.0 a Anisotropic linewidth parameter, half-width at halfheight for the absorption spectrum in gauss. 3
values for the Cu-SFA complexes are unusually large, but a similar value of 188 f 10 X lo4 cm-' was found for an alkaline Cu2+meso-tartrate complex (22). We attempted to model the Cu-SFA system with l j 2 Cu/ligand solutions of salicylic acid and phthalic acid and 1/1 mixtures of these ligands. The EPR spectra of these model systems were not similar to the Cu-SFA spectra. Although these small compounds have the donor groups which are thought to be responsible for the chelation of metal ions to FA (19),the effects of additional substituents on the benzene ring and differences in the geometry of the binding sites of the large molecules probably account for the difference in spectra. Considering the complexity of fulvic acids, it is not surprising that these simple compounds alone do not provide a model for the chelating sites of Cu2+. Functional group analyses (13) strongly suggest that the donor groups of FA are predominantly carboxylate, phenolate, and carbonyl. However, the behavior of the model systems makes it necessary to look beyond simple phthalic acid and/or salicylic acid as bidentate chelating groups in FA systems. We speculate that in the stronger class of sites (low Cu/FA), four (or three) donor atoms chelate the Cu2+. These groups are not likely to originate from a molecule having the usual geometry of a quadridentate ligand. The donor atoms could be from different molecules, in agreement with the observation that di- and trivalent metal ions increase the average molecular weight of humic acids (23). In either case, these complexes would be expected to be found in solutions of low metal to FA ratio (i.e., a t the end of the titration of Cu2+with FA).
ANALYTICAL CHEMISTRY, VOL. 50, NO. 12, OCTOBER 1978
When the Cu/FA ratio is large (the beginning of the titration), the more numerous weaker sites predominate. Under these conditions, water is more competitive as a ligand and possibly only two FA donor atoms are bonded to each Cu2+. The excess Cu2+would tend to cause ligand dissociation of the second type of quadridentate chelation (a) into weaker, bidentate sites (b). We arbitrarily considered two salicylic
n2d
.OH,
HO ,
OH,
ibl
acid groups, but any combination of phthalic and salicylic acid groups would be equally justified. This model is also in agreement with the larger number of weak than strong sites. The 1/1 species of (b) might be included in the weaker class of sites observed in the Scatchard analysis. The functional group analyses of SFA and WFA (13,14)gave the following values in equiv/mol for SFA and WFA, respectively: total acidity (8.6, 6.61, carboxyl (5.3, 3.91, phenol OH (3.3, 2.7), and carbonyl (2.3, 2.7). With the exception of the carbonyl analysis, the SFA functional group concentrations are 1.2 to 1.4 times greater than the corresponding WFA values. The Scatchard data (Table I) indicate that the number of binding sites are also greater for SFA than WFA in approximately the same ratio. The 8.6 equiv/mol total acidity of SFA is about twice the 4.4 binding sites/molecule for SFA at pH 6.0. The corresponding values for WFA are 2.6 sites/molecule and 6.6 equiv/mol; for the same ratio, 3.3 sites/molecule would be necessary. This discrepancy is not surprising because of the complexity of the system and the error in the Scatchard treatment. We chose, therefore, one-half the WFA and SFA total acidity values as a measure of the total number of binding sites/molecule for the above Cheam (18)K data treatment. For example, SFA has 4.3 equiv chelating sites/mol or 1 equiv chelating site/l50 g. An additional observation is significant when discussing the nature of FA binding sites. Their number shows a dramatic increase between pH 5 and 6 (Table I and Figure 4). The rapid increase in binding sites cannot be due to increased availability
1679
of anionic donor atoms due to proton dissociation. The most reasonable explanation is that a structural change in the FA macromolecules brings more functional groups to the surface where they can act as ligands. Such a change might be the dissociation of dimers or aggregates, or some more subtle conformational change triggered by the pH change since gel filtration studies have shown no change in the size of humic acids over this pH range, whereas significant changes do occur below pH 4 and above pH 7 (24).
ACKNOWLEDGMENT We thank N. Dennis Chasteen for helpful discussions of the EPR spectra. LITERATURE CITED (1) F. Morel, R. E. McDuff, and J. J. Morgan, Mar. Chem., 4, 1-26 (1976). (2) A. Lerman and C. W. Childs, in "Trace Metals and Metal-Organic Interactions In Natural Waters", P. C. Singer, Ed., Ann Arbor Science Publishers, Ann Arbor, Mich., 1973, p 201-235. (3) F. Morel, R. E. McDuff, and J. J. Morgan, in Ref. 2, pp 157-200. (4) J. H. Reuter and E. M. Perdue, Geochim. Cosmhim. Acta, 41,325-334 (1977). (5) F . J. Stevenson, Soil Sci., 123, 10-17 (1977). (6) R. D. Guyand C. L. Chakrabarti, Can. J. Chem., 54, 2600-2611 (1976). (7) J. Buffle, F-L. eeter, and W. Haerdi, Anal. Chem., 49, 216-222 (1977). (8) R. Ernst, H. E. Allen, and K. H. Mancy, WaterRes., 9, 969-979 (1975). (9) P. L. Brezonlk, P. A. Brauner, and W. Stumm, Water Res., I O , 605-612 (1976). (IO) 1.A. O'Shea and K. H. Mancy, Anal. Chem., 48, 1603-1607 (1976). (11) A. M. Bond and G. Hefter, J. Ekctroanal. Chem., 31, 477-485 (1971). (12) E. Jacobsen and H. Lindseth, Anal. Chim. Acta, 86, 123-127 (1976). (13) J. H. Weber and S. A. Wilson, Water Res., 9, 1079-1084 (1975). (14) S. A. Wilson and J. H. Weber, Chem. Geol., 19, 285-293 (1977). (15) S. A. Wilson and J. H. Weber, Anal. Lett, 10, 75-84 (1977). (16) G. Scatchard, Ann. N.Y.Acad. Sci., 51, 660-672 (1949). (17) G. Scatchard, J. S. Coleman and A. L. Shen, J . A m . Chem. Soc., 79, 12-20 (1957). (18) V. Cheam. Can. J . Soil Sci., 5 3 , 377-382 (1973). (19) D. S. Gamble, M. Schnitzer, and I. Hoffman, Can. J . Chem., 48, 3197-3204 (1970). (20) D. S. Gamble, Can. J . Chem., 48, 2662-2669 (1970). (21) D. S. Gamble, Can. J . Chem., 51, 3217-3222 (1973). (22) A. J. Fatiadl. J . Res. Nat. Bur. Stand., Sect. A , 74, 723-731 (1970). (23) 6. Kribek, J. Kaigl, and V. Oruzinsky, Chem. Geol., 19, 73-81 (1977). (24) E. T. Gjessing, Schweiz. 2. Hydro/., 3 3 , 592-600 (1971).
RECEIVED for review March 13,1978. Accepted July 24, 1978. This work was partially supported by National Science Foundation Grant OCE77-08390 and by the Central University Research Fund of the University of New Hampshire.
