Competition of Ca(II) and Mg(II)with Ni(II) for Binding by aWell-Characterized Fulvic Acid inModel SolutionsR U P A S R I M A N D A L ,M O H A M E D S . A . S A L A M ,J O H N M U R I M B O H , N O U R I M . H A S S A N ,C H U N I L . C H A K R A B A R T I , * A N DM A R G A R E T H . B A C K

Ottawa-Carleton Chemistry Institute, Department ofChemistry, Carleton University, 1125, Colonel By Drive,Ottawa, Ontario K1S 5B6, Canada

D E N I S C . G R EÄ G O I R E

Geological Survey of Canada, 601 Booth Street,Ottawa, Ontario K1S 0E8, Canada

Competition of Ca(II) and Mg(II) with Ni(II) ions forbinding sites of a well-characterized fulvic acid (FA) inmodel solutions at constant pH and ionic strength wasinvestigated. The Competing Ligand Exchange Method withChelex-100 and dimethyl glyoxime as the competingligands was employed to measure the rate of free Ni2+

ion release using graphite furnace atomic absorptionspectrometry and adsorptive cathodic stripping voltammetry,respectively. The Windermere Humic Aqueous Modelwas used to predict the effect of competition of Ca(II) andMg(II) on the binding of Ni(II) by the FA. The resultsshow that the presence of high concentrations of Ca(II)and Mg(II) in model solutions has considerable effects onthe binding of Ni(II) by the FA. Since the concentrationsof Ca(II) and Mg(II) used are 4 orders of magnitude higherthan those of Ni(II), Ca(II) and Mg(II) can outcompeteNi(II) for sites where electrostatic interactions dominate,resulting in Ni(II) forming weak Ni(II)-FA complexes that arelabile. The significance is that in freshwaters containinghumic substances and trace quantities of nickel and majorcations, Ca2+ and Mg2+, the competition of Ca2+ andMg2+ with Ni2+ for binding sites of humic substancesproduces weak Ni(II)-humate complexes that are labile,releasing free Ni2+ ions.

IntroductionMetal complexation by naturally occurring, organic com-plexants in freshwaters controls the speciation of transitionmetals and, hence, the bioavailability and toxicity of metals.One such category of complexants is humic substances (e.g.,humic and fulvic acids), which arise from chemical andbiological degradation of plant remainders and other debriseither in aquatic systems or in soils. Trace metal complexationby humic substances plays an important role in solubilizingmetals and in regulating transport, bioavailability, and toxicityof metals in freshwaters.

In our previous publications (1, 2), we did not explicitlytake into consideration the contribution of Coulombicinteractions to the total bonding between the metals and theorganic complexants on the grounds that, at the fixed pHand the fixed ionic strength used in those works, thecontribution of the Coulombic interactions to the totalbonding remained constant and was therefore not respon-sible for any of the observed changes. In the present study,we have explicitly taken into consideration the contributionof Coulombic interactions between the negatively chargedsurfaces of humic substances and the major cations, Ca2+

and Mg2+, in the total bonding of the target trace metal ion,Ni2+, by humic substances.

In unpolluted freshwaters, alkali and alkaline earth metalions may influence the complexation of trace metals by humicsubstances since Ca2+ and Mg2+ ions are generally presentat many orders of magnitude greater concentrations (∼10-3

M) (than those of trace metal ions∼<10-8 M), which outweightheir weaker complexing power. The literature containsseveral reports of experimental studies on the competitionbetween Ca2+ or Mg2+ and copper species for binding byhumic substances. Morel and Hering (3) have reported thatCa2+ does not have any competitive effect on Cu(II)-humatebinding in the experimental conditions used by them andhave concluded that different ligand sites must be involvedin calcium and copper binding. Recently, Cao et al. (4) havereported that Ca2+ does have a measurable effect on thebinding of trace metals [Cd(II), Pb(II), and Cu(II)] by humicacid. The effect however was pH dependent, and at pH > 6,the effect on the Cu(II)-humate binding was negligible. SinceCa2+ mainly forms outer-sphere complexes with ligands, theeffect of adding Ca2+ might have been due to an overall changein electrostatic attraction. Cabaniss and Shuman (5) andSunda and Hanson (6) also have reported small but measur-able competition between calcium and copper in fulvic acidtitrations. Buffle et al. (7) have measured copper binding byorganic matter as a function of added Ca2+ ions in freeze-concentrated marsh waters using an ion-selective electrodeover the pH range of 5-6.5. They have reported a decreasein bound copper with increasing calcium concentration (from1 × 10-4 up to 4 × 10-3 M) and that for pH > 6 and [Ca2+]< 5 × 10-3 M (typical of freshwaters) the influence of [Ca2+]on the complexation of Cu2+ by fulvic acid is negligible. vanden Hoop et al. (8) have observed a decreased associationof Cd2+ and Zn2+ with the humate polyanion at pH 5 and atconstant ionic strength, suggesting that Ca2+ ions competewith the heavy metal ions for binding by humic material.Ephraim et al. (9) have suggested that Ca2+ will have minimalinterference in the interactions of strongly complexing metalssuch as Cu2+ with FA but could reduce the binding of weakcomplexing metals such as Cd2+ and Zn2+. The differentexperimental conditions used by the above workers, assummarized in Table 1, may account for the difference intheir observations. Most of the above workers have reportedthat Ca(II) and Mg(II) have a measurable effect on the bindingof trace metals by humic substances, with the exception ofMorel and Hering (3), who found no effect of the competitionof Ca(II) on the Cu(II)-humate binding. This may be due tohigh pH 8.2 and ionic strength 0.5 M used by them, resultingin very strong Cu(II)-humate bindingsstrong intrinsic(chemical) binding further enhanced by strong Coulombicattraction (10). Apparently, the effect of Ca(II), competingfor specific functional groups and decreasing (shielding) theCoulombic attraction of the neighboring ionized groups, istoo small to be measurable.

* Corresponding author telephone: (613)520-2600-3839; fax: (613)-520-3749/3830; e-mail: chuni•chakrabarti@carleton.ca.

Environ. Sci. Technol. 2000, 34, 2201-2208

10.1021/es991074h CCC: $19.00 2000 American Chemical Society VOL. 34, NO. 11, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 9 2201Published on Web 04/22/2000

Tipping (11) has used his Humic Ion-Binding Model V topredict competition effects in the binding of trace metalspecies and alkaline earth cations (Ca2+, Mg2+) by fulvic-type humic substances. Recently, Tipping (12) has proposedHumic Ion-Binding Model VI, which gives an improveddescription of the interactions of protons and metal ionswith humic substances and accounts for competition andionic strength effects and for proton-metal exchange. Thispaper forms a part of the research program of this laboratoryto develop comprehensive schemes of chemical speciationfor metals and metalloids in freshwaters, rainwater, and snowbased on kinetic approaches (13-21).

Objective. The objective of this research is to study thecompetitive binding of Ca(II), Mg(II), and Ni(II) by a well-characterized FA in model solutions containing low Ni2+ ionconcentration (10-7 M) and high concentrations of Ca2+ andMg2+ ions (∼10-3 M), typical of freshwaters.

TheoryCompeting Ligand Exchange Method. In freshwaters, metalsmay be present as free ions and also as complexes of variouscomplexants, organic and inorganic. Humic substances arenaturally occurring organic complexants that are macro-molecular, polyfunctional, polydisperse, oligoelectrolytical,and ubiquitous in freshwaters. We have used a well-characterized fulvic acid as an organic complexant for metalsand a kinetic method for characterization of the metalcomplexes. We have adapted the kinetic model proposed byOlson et al. (22) and have developed it further (18) to studythe dissociation kinetics of complex ML, where M is a metalion and L is a macromolecular, polyfunctional, multisite,heterogeneous, organic complexant, such as a fulvic acid.Consider an aqueous mixture of n components in whicheach component, designated MLi, undergoes a first-order orpseudo-first-order reaction simultaneously with all othercomponents:

where k1 and k-1 are the rate coefficients for the forward andthe reverse reaction, respectively. The charges on M and Lhave been neglected in this general statement of the reaction.In the Competing Ligand Exchange Method (CLEM) with acompeting ligand P, the latter reacts with M as follows:

where P is solid Chelex-100 cation-exchange resin or asolution of dimethyl glyoxime (DMG). The model assumesthat (a) reaction 2 is much faster than reaction 1 and (b) [P]. [M].

Because P is added in large excess, [P] . [M], and reaction2 is pseudo-first-order. Since k2 is large, as has beendetermined by inductively coupled plasma mass spectrom-etry for Chelex (18, 19) and by AdCSV for DMG (23), k2[Chelex]. k-1[L].

Hence, the overall reaction

is irreversible. The rate of formation of the MP complex ishence determined by the rate of dissociation of ML in reaction1, and the rate expression is simply:

In the Chelex technique, if each complex, MLi, undergoesindependently and simultaneously a first-order or pseudo-first-order dissociation reaction, the sum of the concentra-tions of all components remaining in the solution at time t,CML(t), decreases exponentially as

where Ci0 is the initial concentration of MLi, the ith com-

ponent. In the AdCSV technique, the concentration of freemetal ion at any time t, CM(t), is described by an exponentialfunction that increases to a limiting value:

Experimental SectionReagents and Materials. Stock solution (1000 µg/mL) of Ni2+

was prepared by dissolving an appropriate quantity of nickelmetal powder (SPEX 99.999%) in ultrapure nitric acid (UltrexII, J. T. Baker Inc., Phillipsburg, NJ) with heating and dilutingto the appropriate volume with ultrapure water; the finalsolution contained 1% (v/v) ultrapure nitric acid. Standardsolutions (1000 µg/mL) of Ca2+ (from carbonate) and Mg2+

(from nitrate) were purchased from Fisher Scientific Com-pany, NJ. Ultrapure water of resistivity 18.2 MΩ‚cm wasobtained direct from a Milli-Q-Plus water purification system(Millipore Corporation).

A well-characterized FA (24) extracted from a Bh horizonsoil and collected from Armadale, Prince Edward Island, wassupplied by Dr. D. S. Gamble, Agriculture Canada, Ottawa,ON, Canada. However, the “well-characterized FA” termrequires some elaboration. The FA was characterized bypotentiometric titration (24-27). Gamble (26, 27) has dealtwith the errors involved in the potentiometric titration ofthe FA. The total number of phenolic OH and carboxyl groupsdetermined by potentiometric titration are 3.0 (24, 25) and7.71 mmol/g fulvic acid (24, 26, 27), respectively. Thebidentate complexing capacity was therefore approximately5.4 mmol/g fulvic acid (28). It should be noted that for anaturally occurring, heterogeneous complexant, such as theFA the molar mass is not directly accessible. However, a molarconcentration has been calculated, based on the bidentatecomplexing capacity, 5.4 mmol/g, of the Armadale fulvic acid,as determined by Gamble et al. (28). A stock solution (1.0219g/L) of the Armadale FA was prepared by dissolving ap-propriate amount of the freeze-dried FA in ultrapure water.The stock solution was stored in the dark at 4 °C until use.The stock solution of the Armadale fulvic acid 1.0219 g/Lcorresponded to 5.5 × 10-3 mol/L solution of the FA,calculated on the basis of the bidentate complexing capacityof the FA. Chelex-100 resin (Bio-Rad, 100-200 mesh, sodiumform) was soaked in ultrapure water and stored as such. TheChelex resin was taken out of the water solution and filteredjust before use. The stock solution of 0.1 M DMG was preparedby dissolving an appropriate amount of the solid DMG (FisherScientific, certified) in methanol (Spectro grade). For AdCSV,

TABLE 1. Summary of Experimental Conditions Used byDifferent Research Groups

research groups pH

ionicstrength

(M)[Cu(II)]

(M)

[organicmatter](mg/L)

[Ca2+]or [Mg 2+]

(M)

Morel & Hering (3) 8.2 0.50 10-5 0.3 10-2

Cao et al. (4) 6, 7, 8 0.05 10-5 5.0 10-6-10-4

Cabaniss et al. (5) 4-9 0.10 10-6 5.0 10-4-10-2

Buffle et al. (7) 5-6.5 0.10 10-5 35 10-4-10-3

MLi \k1

k-1M + Li (slow) (1)

P + M \k2

k-2MP (fast) (2)

P + MLi f MP + Li (3)

d[MP]dt

) -d[MLi]

dt) k1[MLi] (4)

CML(t) ) ∑i)1

n

Ci0 exp(-kit) (5)

CM(t) ) ∑i)1

n

Ci0[1 - exp(-kit)] (6)

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the pH buffer solution was 2 M Hepes (BDH 99%) plus 2 Maqueous NaOH (Anachemia).

Cleaning Procedure. All containers used in this work wereprecleaned by soaking in 10% HNO3 (AR grade) for 1 weekat room temperature, followed by rinsing 5 times withultrapure water. Finally, they were soaked in ultrapure wateruntil they were used.

Samples. Three types of model aqueous solutions wereused: (1) solutions containing fixed concentrations of Ni(II)[3.4 × 10-7 M] and FA [2.5 × 10-5 M] and varying concentra-tions of Ca(II); (2) solutions containing fixed concentrationsof Ni(II) [3.4 × 10-7 M] and FA [2.5 × 10-5 M] and varyingconcentrations of Mg(II); (3) solutions containing fixedconcentrations of Ni(II) [3.4 × 10-7 M] and FA [2.5 × 10-5 M]and varying concentrations of Ca(II) and Mg(II). For thiswork the [Ni]/[FA] mole ratio of 0.01 was used. If the FA has∼1% strong binding sites, then the mole ratio of metal to theFA for the occupation only of the strong binding sites of theFA was ∼0.01. In our previous paper (2), we have shown thatthe dissolved organic carbon (DOC) in the Rideau Riversurface water samples was 6.5 mg of C/L. Assuming that∼75% of the DOC was present as a humic substance, theconcentration of humic substances in the Rideau Riversurface waters was estimated to be roughly 3 × 10-5 M,assuming that its bidentate complexing capacity for Ni(II) isof the order of 5.4 × 10-3 mol/g (28). Therefore, for this work,2.5 × 10-5 M FA was used to simulate freshwaters in thelaboratory. Model solutions were prepared by spiking ul-trapure water with appropriate volumes of standard solutionsof the metals and of the FA solution (1.0219 g/L). The spikedsamples were equilibrated for 24 h, and then the pH of themodel solutions was measured using a Fisher ScientificAccumet 925 pH/ion meter. The pH meter was calibratedwith two standard pH buffer solutions (pH 7 and pH 10). Forgraphite furnace atomic absorption spectrometry (GFAAS),the pH of all the test solutions was adjusted to pH 8.0 ( 0.5(to simulate freshwaters) using dilute HNO3 (Ultrex II) ordilute NaOH (electrolytically purified) solutions. For AdCSV,the pH of the test solutions was adjusted to 8.0 ( 0.1 using2 M Hepes/NaOH buffer solution. For GFAAS, the changesin the pH between the beginning and the end of theexperiment were found to be < (0.5 pH unit. The variationsof the pH units were not systematic. Sodium perchlorate(NaClO4) was used to fix the ionic strength of all the testsolutions at 6 × 10-3 M.

Apparatus. Screw-capped Teflon cylindrical bottles (500mL capacity) were machined to fabricate the Reactors forkinetic studies of the test solutions using GFAAS (21). Teflonbottles of 1000 mL capacity were used as sample storagecontainers. The concentration of nickel in the model solutionswas determined by GFAAS using the graphite platformtechnique for metal atomization. For the GFAAS, Perkin-Elmer Zeeman atomic absorption spectrophotometer, model4100ZL, equipped with an AS-70 autosampler, pyrolyticallycoated graphite tubes (Perkin-Elmer) and laboratory-madepyrolytic graphite platforms were used. The instrumentalparameters used in the GFAAS determination of nickel areas follows: analysis line, 232.2 nm; ashing temperature, 1000°C; atomization temperature, 2300 °C. For the determinationof nickel concentration, 10 µL of the test solution was injectedinto the graphite furnace where it was dried, ashed, andatomized. The signal was measured in the peak area mode.Each completed determination was followed by a 4-s cleanupcycle of the graphite furnace at 2600 °C. During the drying,ashing, and cleanup cycles, the internal argon gas was passedthrough the graphite furnace at 300 mL/min. The internalargon gas flow was interrupted during the atomization cyclebut was restored for the cleanup cycle. Voltammetric meas-urements were made using a Autolab-Ecochmie-PGSTAT 30with a Metrohm 663 VA stand. A large static mercury drop

electrode (SMDE) with a 0.52-mm2 surface area was employedas the working electrode. The reference electrode was Ag/AgCl, filled with saturated KCl, and the counter electrodewas a Pt rod (Metrohm).

MethodologyMethod for the Kinetic Measurement. For the GFAAStechnique, the rate of Ni2+ ion release from Ni(II)-FAcomplexes was measured using GFAAS. For the measurementof the rate by GFAAS, the pretreated Chelex-100 resin wasadded to a 300-mL test solution placed in the reactor, andthe mixture was stirred with a Teflon-coated magnetic stirringbar. Test solutions for analysis were withdrawn from thereactor through a nylon membrane filter placed at the endof the reactor tube, which filtered out the Chelex-100 resinand injected a clear filtrate into the graphite furnace. Thezero time for the kinetic measurement was taken as the timeat which the Chelex-100 resin first came in contact with thetest solution in the reactor. The sum of the concentrationsof all the kinetically distinguishable components of the nickelcomplexes remaining in the test solution, CML(t), (eq 5) wasmeasured as a function of time.

For the AdCSV technique, 50 mL of the test solution wasbuffered at pH 8.0 ( 0.1 with 2 M Hepes/NaOH buffer solutionin a Teflon cell (Metrohm). It was deoxygenated for 10 minwith prepurified nitrogen (extra-dry BOC Gases). Then 500µL of 0.1 M DMG was added. The preconcentration potentialwas applied for 30 s using a new mercury drop while thesolution was stirred at a rotational speed 2000 rpm. The testsolution was then allowed to become quiescent for 5 s, andthe adsorbed Ni(DMG)2 complex was stripped using a squarewave waveform. The dissociation rate coefficients of theNi(II)-FA complexes were obtained from the ip/i0 ) f(t)curves, where ip and i0 are the current for Ni2+ ion generatedby dissociation of the nickel-FA complexes and the referencecurrent under conditions of no complexation of Ni2+,respectively. The stripping conditions were as follows: purgetime, 5 s; deposition potential, -0.7 V; deposition time, 30s; equilibrium time, 5 s; frequency, 160 Hz; initial potential,-0.7 V; final potential, -1.3 V; step potential, 5 mV/s;amplitude, 20 mV.

Data Analysis for the Kinetic Measurement. The ex-perimental data were analyzed for discrete values of thedissociation rate coefficients by a method based on nonlinearregression analysis using the Marquardt-Levenberg algo-rithm. This method analyses the data assuming the decreasein the metal concentration, which represents dissociation ofthe metal complexes as a sum of exponential terms. All theexperimental data were used in the fitting.

Method for the Measurement of Stability Constants. Thepurpose of determining stability constants of Ni(II) complexesin the model solutions is to study the effects of thecompetition of major cations, Ca2+ and Mg2+, in modelsolutions on the Ni(II)-humate binding. The above com-parison required the use of competiting ligandsspoly-carboxylate ligands similar to humate ligands. However, forNi, which is insoluble in mercury, the determination of thestability constant of Ni(II)-humate complex precluded theuse of anodic stripping voltammetry (ASV). Instead, we usedthe adsorptive cathodic stripping voltammetry, which wasused by Qian et al. (29) for Co(II) complexes in freshwaters.We followed the method of competing ligand exchange (CLE)/adsorptive cathodic stripping voltammetry (AdCSV) of vanden Berg (30). Because the CLE/AdCSV method required theuse of a competing ligand that would form a much strongercomplex than the Ni(II)-humate complex, we were com-pelled to use ethylenediaminetetraacetic acid (EDTA) as apolycarboxylic acid even though EDTA (a chelate-former parexcellence) is not strictly comparable to the simple poly-carboxylic acids of humate ligands. The strength of AdCSV

VOL. 34, NO. 11, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 9 2203

compared to ASV is that the former has much better limitsof detection; its weakness is that the addition of the competingligand may perturb the chemical equilibria; and also, thelimited range of Ni(II) concentration in the test solutionsrequired for relating the signal to the Ni(II) concentration isa constraint on its wide application.

Results and DiscussionEffect of the Various Concentrations of Ca(II) and Mg(II)on the Lability of the Ni(II)-FA Complexes. For the sake ofcompleteness of experimental data, we studied the effect ofvarying concentrations of Ca(II) on the release of Ni(II) andCa(II), using the same model solutions and the sameexperimental conditions. Figure 1a shows the results of theformer, and Figure 1b shows the results of the latter. Figure1a is interpreted as follows. In the absence of Ca(II), thestrong binding sites of the FA are occupied by Ni(II), formingstrong complexes that are inert (top curve). However, thepresence of high and sharply increasing concentrations ofCa(II) (1 × 10-4-2 × 10-3 M) (each curve from top to bottom)

has a dramatic effect on the lability of the Ni(II)-FAcomplexes, where high concentrations of Ca(II) compensatesfor its relatively weak fulvic-binding affinities and results information of weak Ni(II)-FA complexes that are labile. Theexperimental data presented in Figure 1a were analyzed usinga nonlinear regression method. The values for kineticparameters of the Ni(II)-FA complexes in the presence ofvarious concentrations of Ca(II) and constant concentrationsof Ni(II) and FA are presented in Table 2, which shows that,in the presence of the relatively high concentration of Ca(II)(2 × 10-3 M), the Ni(II)-FA complex has a faster dissociationrate coefficient (k1 ) 7.0 × 10-2 s-1) than in the presence ofthe relatively lower concentration of Ca(II) (1 × 10-3 M) (k1

) 6.5 × 10-3 s-1), in the presence of a still lower concentrationof Ca(II) (1 × 10-4 M) (k1 ) 2.6 × 10-3 s-1), and in the absenceof Ca(II) (k1 ) 2.0 × 10-4 s-1). These results are consistentwith the interpretation of the lability/inertness of the Ni(II)-FA complexes given above. Figure 1b shows that the Ca(II)-FA complexes are ∼100% labile [almost all the Ca(II) arereleased within ∼120s] even at the relatively low concentra-

FIGURE 1. (a) Nickel remaining bound to the FA as a function of time in model solutions containing varying concentrations of Ca(II),measured by 4100ZL GFAAS. Chelex-100 was the competing ligand. Concentrations of Ni(II) and FA are the same for all curves. Bidentatecomplexing capacity of the FA, 5.4 mmol/g; pH, 8.0 ( 0.5; ionic strength, 6 × 10-3 M; temperature 296 K: (O) [Ni(II) 3.4 × 10-7 M] + [FA2.5 × 10-5 M]; (4) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 1 × 10-4 M] + [FA 2.5 × 10-5 M]; (0) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 1 × 10-3 M] + [FA2.5 × 10-5 M]; ()) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 2 × 10-3 M] + [FA 2.5 × 10-5 M]. (b) Calcium remaining bound to the FA as a functionof time in model solutions containing varying concentrations of Ca(II), measured by 4100ZL GFAAS. Chelex-100 was the competing ligand.Concentrations of Ni(II) and FA are the same for all curves. Bidentate complexing capacity of the FA, 5.4 mmol/g; pH, 8.0 ( 0.5; ionic strength,6 × 10-3 M; temperature, 296 K. (4) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 1 × 10-4 M] + [FA 2.5 × 10-5 M]; (0) [Ni(II) 3.4 × 10-7 M] + [Ca(II)1 × 10-3 M] + [FA 2.5 × 10-5 M]; (]) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 2 × 10-3 M] + [FA 2.5 × 10-5 M].

TABLE 2. Kinetic Parameters of Ni(II)-FA Complexes in Model Solutions in the Presence of Varying Concentrations of Ca(II)and/or Mg(II) and Constant Concentrations of Ni(II) and the FAa

relative distribution of components of Ni(II) complexes dissociation rate coeff of Ni(II) complexesb

[Ca(II)] (M) [Mg(II)] (M) C1 (%) C2 (%) inertc k1 × 103 (s-1) C1 k2 × 105 (s-1) C2

nil nil 6.1 ( 0.1 94.0 ( 0.2 0.2 ( 0.1 0.1 ( 0.01 × 10-4 nil 12.2 ( 0.3 87.7 ( 0.5 2.6 ( 0.1 0.3 ( 0.01 × 10-3 nil 20.0 ( 0.6 79.8 ( 0.4 6.5 ( 0.1 3.2 ( 0.12 × 10-3 nil 39.0 ( 2.0 48.9 ( 1.0 12.0 ( 0.4 70.0 ( 0.1 85.0 ( 0.0nil 1 × 10-4 10.1 ( 0.2 90.0 ( 1.0 2.6 ( 0.2 0.4 ( 0.1nil 1 × 10-3 24.5 ( 1.0 75.0 ( 1.0 6.9 ( 0.3 3.6 ( 0.1nil 2 × 10-3 41.5 ( 1.0 44.5 ( 1.0 14.0 ( 0.3 75.0 ( 0.3 80.0 ( 0.21 × 10-4 1 × 10-4 13.3 ( 0.3 87.2 ( 0.5 3.5 ( 0.1 0.4 ( 0.05 × 104 5 × 10-4 32.9 ( 0.6 67.1 ( 0.4 8.9 ( 0.1 4.8 ( 1.01 × 10-3 1 × 10-3 48.9 ( 2.0 39.1 ( 1.0 10.8 ( 0.4 85.0 ( 0.2 83.0 ( 4.0

a Measured by 4100ZL GFAAS; pH, 8.0 ( 0.5; ionic strength, 6 × 10-3 M; temperature, 298 K. [Ni(II) 3.4 × 10-7 M] + [FA 2.5 × 10-5 M] + [Ca(II)]and/or [Mg(II)]. b k1 and k2 are the dissociation rate coefficients of the faster and the slower component, respectively. c Inert component is definedas having k < 10-6 s-1.

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tion of 1 × 10-4 M Ca(II) as compared to ∼90% of Ni(II)-FAcomplexes remaining inert.

Similar studies were done on the effect of varyingconcentrations of Mg(II) on the release of Ni(II) and Mg(II),using the same model solutions and the same experimentalconditions. Figure 2a shows the results of the former, andFigure 2b shows the results of the latter. Figure 2a,b showsa behavior similar to that in Figure 1a,b. In the absence ofMg(II), nickel occupies the small number of strong bindingsites of the FA, forming strong Ni(II)-FA complexes that areinert (top curve). However, the presence of high and sharplyincreasing concentrations of Mg(II) (1 × 10-4-2 × 10-3 M)(each curve from top to bottom) has a dramatic effect on thelability of the Ni(II)-FA complexes, where high concentra-tions of Mg(II) compensate for relatively weak fulvate-bindingaffinities, resulting in the formation of weak Ni(II)-FAcomplexes that are labile. Figure 2b shows that the Mg(II)-FA complexes are ∼100% labile [almost all the Mg(II) arereleased within ∼120 s] even at the relatively low concentra-tion of 1 × 10-4 M Mg(II) as compared to ∼90% of Ni(II)-FAcomplexes remaining inert.

Figure 3 presents the effect of the varying concentrationsof Ca(II) and Mg(II) on the release of free Ni2+ ion from theNi(II)-FA complexes in the model solutions and shows abehavior similar to that in Figures 1 and 2. In the absenceof Ca(II) and Mg(II), nickel occupies the small number of thestrong binding sites of the FA, forming strong Ni(II)-FAcomplexes that are inert (top curve). However, the presenceof high and sharply increasing concentrations of Ca(II) andMg(II) (1 × 10-4-1 × 10-3 M) (each curve from top to bottom)has a dramatic effect on the lability of the Ni(II)-FAcomplexes, where high concentrations of Ca(II) and Mg(II)compensate for their relatively weak fulvate-binding affinities,resulting in the formation of weak Ni(II)-FA complexes thatare labile. The values for the kinetic parameters of Ni(II)-FAcomplexes in the presence of varying concentrations of Ca(II)and Mg(II) and constant concentrations of Ni(II) and FA arepresented in Table 2. Comparison of Figures 1 and 2 withFigure 3 reveals that Ni(II)-FA complexes are more labile in

the presence of Ca(II) and Mg(II) when they are together ascompared with the lability of the Ni(II)-FA complexes wheneither Ca(II) or Mg(II) is present alone. We have investigatedwhether this is simply a mass effect or a synergistic effect.Figures 1 and 3 and Table 2 show that in the presence of 1× 10-3 M Ca2+ alone (Figure 1, the second curve from bottom),the percentage of the labile fraction of the nickel-fulvic acidcomplexes (C1 ) 20.0%) is lower than that in the presenceof 5 × 10-4 M Ca2+ and 5 × 10-4 M Mg2+ when they arepresent together (Figure 3, the second curve from bottom,C1 ) 32.9%). Hence, the enhanced lability of the Ni(II)-FAcomplexes when Ca(II) and Mg(II) are present together ascompared with the lability of the Ni(II)-FA complexes when

FIGURE 2. (a) Nickel remaining bound to the FA as a function of time in model solutions containing varying concentrations of Mg(II),measured by 4100ZL GFAAS. Chelex-100 was the competing ligand. Concentrations of Ni(II) and FA are the same for all curves. Bidentatecomplexing capacity of the FA, 5.4 mmol/g; pH, 8.0 ( 0.5; ionic strength, 6 × 10-3 M; temperature, 296 K. (O) [Ni(II) 3.4 × 10-7 M] + [FA2.5 × 10-5 M]; (4) [Ni(II) 3.4 × 10-7 M] + [Mg(II) 1 × 10-4 M] + [FA 2.5 × 10-5 M]; (0) [Ni(II) 3.4 × 10-7 M] + [Mg(II) 1 × 10-3 M] + [FA2.5 × 10-5 M]; ()) [Ni(II) 3.4 × 10-7 M] + [Mg(II) 2 × 10-3 M] + [FA 2.5 × 10-5 M]. (b) Magnesium remaining bound to the FA as a functionof time in model solutions containing varying concentrations of Mg(II), measured by 4100ZL GFAAS. Chelex-100 was the competing ligand.Concentrations of Ni(II) and FA are the same for all curves. Bidentate complexing capacity of the FA, 5.4 mmol/g; pH, 8.0 ( 0.5; ionic strength,6 × 10-3 M; temperature, 296 K. (4) [Ni(II) 3.4 × 10-7 M] + [Mg(II) 1 × 10-4 M] + [FA 2.5 × 10-5 M]; (0) [Ni(II) 3.4 × 10-7 M] + [Mg(II)1 × 10-3 M] + [FA 2.5 × 10-5 M]; ()) [Ni(II) 3.4 × 10-7 M] + [Mg(II) 2 × 10-3 M] + [FA 2.5 × 10-5 M].

FIGURE 3. Nickel remaining bound to the FA as a function of timein model solutions containing varying concentrations of Ca(II) andMg(II), measured by 4100ZL GFAAS. Chelex-100 was the competingligand. Concentrations of Ni(II) and FA are the same for all curves.Bidentate complexing capacity of the FA, 5.4 mmol/g; pH, 8.0 ( 0.5;ionic strength, 6 × 10-3 M; temperature, 298 K. (O) [Ni(II) 3.4 × 10-7

M] + [FA 2.5 × 10-5 M]; (4) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 1 × 10-4

M] + [Mg(II) 1 × 10-4 M] + [FA 2.5 × 10-5 M]; (0) [Ni(II) 3.4 × 10-7

M] + [Ca(II) 5 × 10-4 M] + [Mg(II) 5 × 10-4 M] + [FA 2.5 × 10-5

M]; (]) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 1 × 10-3 M] + [Mg(II) 1 ×10-3 M] + [FA 2.5 × 10-5 M].

VOL. 34, NO. 11, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 9 2205

either Ca(II) or Mg (II) is present alone is due to synergy andis not a mass effect.

Figures 1-3 present the effects of the varying concentra-tions of Ca(II) and Mg(II) on the release of Ni(II) from theNi(II)-FA complexes, measured by the GFAAS technique.The CLEM studies were also done using the AdCSV technique,and the results are presented in Figures 4 and 5 and Table3. The results of both GFAAS and AdCSV show the sametrend of increasingly greater fractions of the Ni(II) becomingmore labile in the presence of increasing concentrations ofCa(II) and Mg(II). The difference between dissociation ratecoefficients and the percentage of the labile fraction of theNi(II)-FA complexes obtained by the two techniques isattributable to their respective analytical windowsstheAdCSV technique allows measurement of faster kinetics andthe associated kinetically distinguishable components,whereas the GFAAS technique allows measurement of only

slower kinetics and the associated kinetically distinguishablecomponents.

Effect of the Various Concentrations of Ca(II) and Mg(II)on the Stability Constants of the Ni(II)-EDTA Complexes.Table 4 shows the effect of increasing concentrations of Ca(II)and Mg(II) on the stability constants of the Ni(II)-EDTAcomplexes measured by AdCSV. The results show that thepresence of high concentrations of Ca(II) and/or Mg(II) hasconsiderable effects on the stability constants of the Ni(II)-EDTA complexes, which decrease in the presence of in-creasing concentrations of Ca(II) and/or Mg(II). The aboveresults confirm the trend of increasing lability of the Ni(II)-FA complexes in the presence of increasing concentrationsof Ca(II) and/or Mg(II) reported above.

The Windermere Humic Aqueous Model (WHAM), achemical equilibrium model and a computer code for waters,sediments, and soils incorporating a discrete site/electrostaticmodel of ion binding by humic substances (31), was used topredict the competition effects of Ca(II) and Mg(II) on thebinding of Ni(II) by the FA and to compare its predictionswith the above experimental results. The WHAM takes intoaccount the interactions of metals with (i) inorganic ligands(OH-, HCO3

-, CO32-, SO4

2-, Cl-), using conventationalequilibrium formulations and equilibrium constants fromthe literature, and (ii) humic substances using Humic Ion-Binding Model V (32). The WHAM predictions presented inTable 5 agree reasonably well with the experimental results,

FIGURE 4. Effect of varying concentrations of Ca(II) on the kineticsof Ni(II)-FA complexes in model solutions by CLE/AdCSV at a SMDE.Concentrations of Ni(II) and FA are the same for all curves. Bidentatecomplexing capacity of the FA, 5.4 mmol/g; pH, 8.0 ( 0.1; ionicstrength, 6 × 10-3 M; temperature, 297 K. (b) reference solution[Ni(II) 3.4 × 10-7 M]; (O) [Ni(II) 3.4 × 10-7 M] + [FA 2.5 × 10-5 M];(0) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 1 × 10-4 M] + [FA 2.5 × 10-5 M];(4) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 1 × 10-3 M] + [FA 2.5 × 10-5 M];(]) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 2 × 10-3 M] + [FA 2.5 × 10-5 M].

FIGURE 5. Effect of varying concentrations of Ca(II) and Mg(II) onthe kinetics of Ni(II)-FA complexes in model solutions by CLE/AdCSV at a SMDE. Concentrations of Ni(II) and FA are the same forall curves. Bidentate complexing capacity of the FA, 5.4 mmol/g;pH, 8.0 ( 0.1; ionic strength, 6 × 10-3 M; temperature, 297 K. (b)reference solution [Ni(II) 3.4 × 10-7 M]; (O) [Ni(II) 3.4 × 10-7 M] +[FA 2.5 × 10-5 M]; (0) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 1 × 10-4 M]+ [Mg(II) 1 × 10-4 M] + [FA 2.5 × 10-5 M]; (4) [Ni(II) 3.4 × 10-7

M] +[Ca(II) 5 × 10-4 M] + [Mg(II) 5 × 10-4 M] + [FA 2.5 × 10-5

M]; ()) [Ni(II) 3.4 × 10-7 M] + [Ca(II) 1 × 10-3 M] + [Mg(II) 1 × 10-3

M] + [FA 2.5 × 10-5 M].

TABLE 3. Kinetic Parameters of Ni(II)-FA Complexes in ModelSolutions in the Presence of Varying Concentrations of Ca(II)and/or Mg(II) and Constant Concentrations of Ni(II) and theFAa

relative distribution ofcomponents of Ni(II)

complexesdissociation rate coeff

of Ni(II) complexes

[Ca(II)](M)

[Mg(II)](M)

C1

(%)C2

(%)k1 × 103

(s-1) C1

k2 × 105

(s-1) C2

nil nil 82.0 ( 1.0 18.0 ( 0.0 2.9 ( 0.1 0.2 ( 0.01 × 10-4 nil 85.0 ( 1.0 15.0 ( 0.0 4.9 ( 0.9 3.9 ( 0.11 × 10-3 nil 85.0 ( 1.0 15.0 ( 1.0 3.2 ( 0.2 11.3 ( 1.02 × 10-3 nil 89.0 ( 2.0 11.0 ( 1.0 5.4 ( 0.5 15.2 ( 2.0nil 1 × 10-4 83.0 ( 2.0 16.0 ( 1.0 3.3 ( 0.2 1.1 ( 0.1nil 1 × 10-3 85.0 ( 1.0 15.0 ( 1.0 4.3 ( 0.3 1.5 ( 0.1nil 2 × 10-3 88.0 ( 1.0 12.0 ( 1.0 4.9 ( 0.3 3.8 ( 0.21 × 10-4 1 × 10-4 83.0 ( 1.0 17.0 ( 1.0 3.3 ( 0.1 3.6 ( 0.05 × 10-4 5 × 10-4 85.0 ( 1.0 15.0 ( 1.0 4.7 ( 0.3 12.3 ( 1.01 × 10-3 1 × 10-3 89.0 ( 2.0 11.0 ( 1.0 3.7 ( 0.2 27.1 ( 4.0

aMeasured by AdCSV. pH, 8.0 ( 0.1; ionic strength, 6 × 10-3 M;temperature, 298 K. [Ni(II) 3.4 × 10-7 M] + [FA 2.5 × 10-5 M] + [Ca(II)]and/or [Mg(II)]. b kl and k2 are the dissociation rate coefficients of thefaster and the slower component, respectively.

TABLE 4. Stability Constants for Various Ni(II)-EDTA Systemsas a Function of Concentrations of Ca(II) and Mg(II) UsingCLE/AdCSV at a SMDE and DMG as the Competing Liganda

[Ca(II)] (M) [Mg(II)] (M) log KNiEDTA

18.65 ( 0.011 × 10-4 18.01 ( 0.025 × 10-4 17.97 ( 0.031 × 10-3 17.77 ( 0.03

1 × 10-4 18.46 ( 0.075 × 10-4 18.23 ( 0.041 × 10-3 18.03 ( 0.02

1 × 10-4 1 × 10-4 17.88 ( 0.055 × 10-4 5 × 10-4 17.78 ( 0.021 × 10-3 1 × 10-3 17.52 ( 0.06

a [EDTA], 5 × 10-8 M; pH, 8.0 ( 0.1; ionic strength, 6 × 10-3 M;temperature, 297 K.

2206 9 ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 34, NO. 11, 2000

considering that the WHAM calculations have not taken intoaccount the particular characteristics of the Armadale FAand that the WHAM is based on equilibrium calculationsand does not take into account nonequilibrium speciation.Also, the Model V [unlike the Model VI (12)] is not optimizedfor the competition of major ions, Ca(II) and/or Mg(II), andthe ionic strength effects.

Figures 1-5 and Tables 2, 3, and 5 show that when highlevels of Ca(II) and Mg(II) (∼10-3 M) are present with Ni(II)in model solutions, increasing the concentrations of Ca(II)and/or Mg(II) results in releasing greater fractions of thebound nickel from Ni(II)-FA complexes as free Ni2+ ion. Theresults described above are probably the composite effect ofthe following factors. When Ca(II) and Mg(II) are present inhigh molar concentrations [several orders of magnitudegreater than that of Ni(II)], the high concentrations of Ca(II)and Mg(II) compensate for their relatively weak fulvic acid-binding affinities. The high concentrations of Ca2+ and Mg2+

have important effects on the total binding energy involvedin the binding of Ni(II) by the FA, i.e., on the sum total ofthe energy involved in the covalent bonding and theCoulombic attraction. It appears that the type of the FAdiscrete sites that dominate the binding under the experi-mental conditions used here are the sites where electrostaticinteractions dominate; for these sites, there is also directcompetition between Ca(II), Mg(II), and Ni(II). Also, there iscompetition for counterion accumulation in the diffuse partof the electrical double layer of the negatively charged fulvateanions and reduction in the net electric charge on the fulvicacid molecule by Ca2+ and Mg2+ ions binding to certain sites,which diminishes the electrostatic contribution to the Ni2+

binding at other sites, resulting in formation of weak Ni(II)-FA complexes that are labile. The predictions of the WHAMare that, for the target trace metal Ni2+ ions, binding occursalmost exclusively at bidentate sites, at both low and highionic strengths. However, the concentrations of Ca2+ andMg2+, and its pKMHA and pKMHB values [where A and B aretwo types of proton-binding groups, one (type A) beingrelatively acidic (mainly carboxyl groups) and the other (typeB) less acidic (phenolic groups)] are such that Ca2+ and Mg2+

ions bind mostly at the monodentate and diffuse doublelayer sites (33). This is significant for the competition, becauseit means that the target trace metal Ni2+ ions and majorcations, Ca2+ and Mg2+, are substantially separated in theirbinding so that to a considerable degree they are notcompeting for the same sites (11). The significance of theseresults is that in model solutions simulating freshwaterscontaining humic substances and the target trace metal Ni(II)and major cations, Ca(II) and Mg(II), the competitive binding

of Ni(II), Ca(II), and Mg(II) by humic substances makesNi(II)-humate complexes labile, releasing free Ni2+ ion,which is reported to be toxic. This knowledge will allowdecision-makers to make better pollution risk assessmentand to adopt environmental policies and regulations con-straining discharge of effluents from mining and metallurgicaloperations into the natural environment, taking into con-sideration the presence of calcium and magnesium carbonatebedrocks and acid rains to dissolve them in the aquaticenvironment.

AcknowledgmentsThe following financial support is gratefully acknowledged:Nickel Producers Environmental Research Association, USA;Inco Ltd., Canada; and Falconbridge Ltd., Canada, for threeresearch contracts; Natural Sciences and Engineering Re-search Council of Canada for a research grant and for theMetals in the Environment-Research Network grant; OntarioPower Generating Company; and the Mining Association ofCanada. We are grateful to Dr. E. Dabek-Zlotorzynska,Environment Canada, for providing some equipment andchemicals; to Dr. D. S. Gamble for supplying the ArmadaleFA; and to Dr. E. Tipping for providing guidance with theWHAM calculations. M.S.A.S. is grateful to the EgyptianMinistry of High Education, J.M. is grateful to the NaturalSciences and Engineering Research Council of Canada, andN.M.H. is grateful to the Government of Libya for the awardof graduate scholarships.

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TABLE 5. WHAM Computer Calculations for Speciation ofNickel in Model Solutions in the Presence of VaryingConcentrations of Ca(II) and Mg(II) and ConstantConcentrations of Ni(II) and the FAa

WHAM predictions experimental

[Ca(II)](M)

[Mg(II)](M)

free Ni2+

ions plusother smallNi species

(%)

Nielectrostat

bound(%)

Nicovalbound

(%)

free Ni2+

ions plusother smallNi species

(%)

1 × 10-4 nil 13.4 1.0 85.6 12.21 × 10-3 nil 36.2 0.2 63.6 20.02 × 10-3 nil 45.6 0.1 54.2 39.0

1 × 10-4 1 × 10-4 18.9 0.6 80.4 13.35 × 10-4 5 × 10-4 36.2 0.2 63.6 32.91 × 10-3 1 × 10-3 45.6 0.1 54.2 48.9

a Measured by 4100ZL GFAAS; pH, 8.0 ( 0.5; ionic strength, 6 × 10-3

M; temperature, 296 K.

VOL. 34, NO. 11, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 9 2207

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Received for review September 20, 1999. Revised manuscriptreceived February 21, 2000. Accepted March 3, 2000.

ES991074H

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