Kinetic study of nickel(II) complexation by pteroylglutamic acidAndre s Thomas*, Ezequiel Wolcan*, Mario R. Fe lizà and Alberto L. Capparelli*,§

Instituto de Investigaciones FisicoquõÂmicas TeoÂricas y Aplicadas (INIFTA), Facultad de Ciencias Exactas, UniversidadNacional de La Plata, Casilla de Correo 16, Sucursal 4, (1900) La Plata, RepuÂblica Argentina

Summary

The kinetics of the complexation of NiII by pteroylglu-tamic acid have been studied in the 5±45 �C range, theionic strength �0:6M� being regulated with KNO3, in the5.5±7.0 pH range, using the stopped-¯ow method. Un-der the experimental conditions two processes wereobserved. The faster process was detected in the milli-second range and is associated with the reaction betweenNiII and the ligand. The slower is observed within a fewseconds. Complementary equilibrium studies were madeat 25 �C. The results are consistent with the formation ofa 1:1 complex between the reactants, and a mechanism isproposed to account for the observed behaviour. Equi-librium constants for the NiII plus pteroylglutamic acidsystem, as well as activation parameters, are reported.

Introduction

Pteroylglutamic acid (folic acid) and related substancesexhibit numerous dissociation equilibria in the 1±9 pHrange�1;2�. Folic acid behaves as an acid owing to thepresence of a phenolic group and two carboxylic groups(Figure 1). Several methods (u.v.±vis., 1H- and 13C-n.m.r. studies�2;3�) have been developed to measure pKvalues for folic acid. Equilibrium constants for the re-action between folic acid and some divalent metal ionshave been reported in the literature�1;4;5�.Thermodynamic and spectroscopic data are available

in the literature for this vitamin, but kinetic studies of itsinteraction with divalent transition metal ions arescarce. To the best of our knowledge, the interactionwith cobalt(II) is the only kinetic study reported�6�.Moreover, the equilibrium data available in the litera-ture either exhibit scattering or do not ®t the experi-mental conditions under which the present studies wereperformed.The principal aim of this work was to elucidate the

kinetics and the mechanism of complexation betweennickel(II) and pteroylglutamic acid.

Experimental

Equilibrium studies

The equilibrium constant for complexation between thereactants was determined from absorbance changemeasurements at constant pH and folic acid concen-tration �1� 10ÿ4 M�. The NiII ion concentration rangedbetween 5� 10ÿ4 and 0:3M. Measurements were madeat 390 nm.Typical absorbance changes at 390 nm for dif-

ferent NiII ion concentrations at pH 6.0 are shown inFigure 2.

Kinetics

The kinetic study was carried out on a Durrum D-110stopped-¯ow apparatus at 25 �C. Reagent gradeNiSO4 � 6H2O (Merck), folic acid (Merck), KNO3

(Merck) and bromothymol blue (Merck) were used assupplied. In all experiments the NiII concentrationgreatly exceeded that of folic acid in order to preservepseudo-®rst order conditions. The NiII ion concentra-tion was between 1� 10ÿ3 and 0:3M. The folic acidconcentration was 1� 10ÿ4 M.Owing to a combination of factors, such as the sol-

ubility of folic acid and its complexes in acid solutionand the hydrolysis of the metal ion, kinetic studies wereperformed in the 5.6±7.3 pH range. The pH of eachsolution was adjusted before mixing and this is the pHreported. Absorbance changes were recorded at390 nm. Some experiments were performed in thepresence of 2� 10ÿ5 M of bromothymol blue at 460 and615 nm. No in¯uence of folic acid concentration wasdetected in the present study. In most of the experi-ments the ionic strength was 0:6M (regulated by addingKNO3). Under all conditions, two processes were ob-served. No absorbance changes were observed in blankexperiments.

0340±4285 Ó 1997 Chapman & Hall

Figure 1. Folic acid molecule.

Figure 2. Typical absorbance changes versus [Ni2+] at pH 7.

Transition Met. Chem., 22, 541±544 (1997) NiII complexation by pteroylglutamic acid 541

* Member of CONICET, Argentina.à Member of CICPBA, Argentina.§ Author to whom all correspondence should be directed.

Results

Equilibrium study

The experimental procedure (spectrophotometric titra-tion) for measuring equilibrium constants has beenreported elsewhere(6). The results obtained in the nickel-(II)±folic acid system can be interpreted according to thefollowing complexation scheme:

A nonlinear regression analysis leads to the followingvalues for the equilibrium constants involved in thiscomplexation: K1 � 20:8M

ÿ1, K2 � 2:0� 10ÿ6 M andK3 � 4:2� 10ÿ5.In Figure 2, the solid curve was obtained using these

equilibrium constants. The extinction coe�cient ob-tained for the NiFoOÿ complex is 6300 �300 cmÿ1 Mÿ1at 390 nm.

Kinetic results

Two well time-resolved processes are detected in thepresent system. Both follow a ®rst order law and are pHindependent, as seen in Table 1. They are discussedseparately.

Fast process

This process is observed in the millisecond time range. Itdepends upon the metal ion concentration. The behav-iour of the apparent rate constant k1app with NiII con-centration is shown in Figure 3 and Table 2.The following experimental dependence can be obtainedfrom this experimental behaviour:

k1app � a�M2��=�1� b�M2��� or�k1app�ÿ1 � b=a� �1=a��M2��ÿ1

with b=a � 8:16� 10ÿ3 s Mÿ1 and 1=a � 9:97�10ÿ5 Mÿ1

(see Figure 4).Experiments conducted in the presence of bromo-

thymol blue show that the fast process is accompaniedby a pH decrease of the solutions. The rate constantsmeasured in the presence or absence of this pH indicatorare the same within experimental error.The temperature dependence of the fast process was

measured in the linear region shown in Figure 3 and athigh concentrations of NiII ion. These results are pre-sented in Figure 5. The activation energies E1app andE01app are 45.3 and 49:4 kJmolÿ1, respectively.The fast reaction is associated with the ®rst stage of

the complexation, where the pteridinic moiety of themolecule is involved, i:e::

NiII � FoL(OH)2ÿ ��! ��KNi±FoL(OH)

��!k1 NiFoLOÿ �H�

Table 1. Experimental rate constants at di�erent pH; k1app values wereobtained with [NiII] = 5 ´ 10)3

M whereas k2app values were averagedover di�erent NiII concentrations at the same pH

pH Slow process Fast processk2app �sÿ1� k1app �sÿ1�

5.8 0.156 35.66.0 0.144 36.46.4 0.189 37.06.6 0.152 ±6.8 0.163 ±7.0 0.141 39.07.2 0.146 36.5

Figure 3. Behaviour of k1app versus [NiII].

Table 2. Experimental rate constants at di�erent NiII concentrations;folic acid concentration � 1� 10ÿ4 M

[NiII] (M) Slow process Fast processk2app �sÿ1� k1app �sÿ1�

0.005 0.155 36.50.010 0.161 52.00.020 0.197 69.90.030 0.175 90.10.039 88.00.050 0.1920.100 0.143 106.00.145 122.20.200 0.195 117.7

Figure 4. 1=k1app versus 1/[NiII].

542 Capparelli et al. Transition Met. Chem., 22, 541±544 (1997)

In this mechanism Ni±FoL(OH) is a monodentate in-termediate complex, whereas NiFoLOÿ is a bidentateone. Therefore, the following expression for k1app can beobtained: k1app � k1K�NiII�0=�1� K�NiII�0�. Comparisonwith the experimental expression gives the value ofk1K � a and K � b. At low concentrations of metal ion,k1app � k1K whereas at high concentration k1app � k1.The activation energy, measured in the linear region

shown in Figure 3, is given by E1app � E1 � DH , whereDH is associated with the enthalpy of formation of theintermediate complex. Therefore, the following valuescan be obtained for these magnitudes: E1 � 49:4 kJmolÿ1 and DH � ÿ4:1 kJmolÿ1. The NiFoLOÿ com-plex, formed in this ®rst reaction stage, involves thehydroxy group and the pteridin nitrogen in the mole-cule.Complementary studies were conducted in the system

8-hydroxyquinoline±nickel(II). This ligand resemblesthe pteridinic moiety of folic acid. In this system, only aprocess in the millisecond range is observed, which canbe associated with the interaction of the metal ion withthe 8-hydroxyquinoline. A marked pH decrease is alsodetected during the reaction.Several studies have been reported in the literature

regarding the kinetics of nickel(II) complexation withmono-, bi- and terdentate ligands�7±13�. A very interest-ing review is given in Ref. 14. The e�ect of the chelatesformation can be described through the following re-action scheme.

Ni�H2O�2�6 �A±BH±C �! �KIPNi�H2O�2�6 A±BH±C

Ni�H2O�2�6 A±BH±C �! �k0

kÿ0Ni�H2O�2�5 ±A±BH±C�H2O

In our case A, BH and C can be identi®ed as the pter-idinic N, pteridinic hydroxy group and p-aminobenzoicamine group, respectively.This scheme can be employed to explain the ®rst stage

in the complexation, involving an ionic pair and amonodentate complex. If complexation is determined(in this stage) by the release of a water molecule from the

coordination sphere of the metal ion, an apparent rateconstant, ca: k0KIP, should be expected. Here k0 meansthe water exchange rate in the metal ion�k0 � 3� 104 sÿ1� and KIP is the ionic pair equilibriumconstant, which can be calculated using the well-knownFuoss±Bjerrum equation�15�. Assuming that in the co-ordination site the net charge is almost zero (the car-boxylate groups are far from this site) then KIP � ca: 0:3and the rate should be 1� 103 Mÿ1 sÿ1 or higher. Valuesof this order have been observed in the kinetics ofcomplexation with pyridine, bipyridine and terpyridine�4� 103, 1:5� 105 and 1:4� 103 Mÿ1 sÿ1, respective-ly)�9;10�.In our case, the experimental rate constant can not be

represented by this simple consideration. Therefore, therate-determining step can not be associated with thewater exchange rate in the coordination sphere ofthe nickel(II) ion. The complex dependence of k1 withmetal ion concentration and the value for K and k1 (81and 123, respectively) implies that, prior to chelation, amonodentate intermediate is formed and that the che-lation is controlled by deprotonation of the OH group inthe pterine moiety. The values of K measured in ourexperiments are quite similar to those reported formonodentate ligands�16� (i:e: in pyridine logK � 1:87).

Slow process

This process is detected within a few seconds. Its be-haviour is similar to that observed for the cobalt(II)±folic acid system�6�, i:e: the apparent rate constant, k2app,is independent of the metal ion concentration (Table 2).Experiments in the presence of bromothymol blue showa rapid pH decrease associated with the earlier stage ofthe reaction, followed by a pH increase in the timewindow where the slow process is detected. In a similarway, as mentioned previously for the fast process, therate constants measured in the presence or absence ofindicator are the same.The temperature dependence of the experimental rate

of the slow process leads to an activation energy of52.2 kJ mol)1.Previous studies on the cobalt(II)±folic acid system

reveal only one process within the same timescale andsimilar apparent rate constants as measured for the slowprocess detected in the system reported in this paper.The following comments about the mechanism can be

inferred from the previous experimental results.The slow process shows similar characteristics to those

observed for the cobalt(II)±folic acid system. This processis independent of pH and metal ion concentration, andthe pH behaviour during complexation is the same inboth systems. The global rate constants and the activa-tion energy are similar. The principal di�erences betweenthe systems are associated with the detection of the fastprocess in presence of nickel(II). The rapid processobserved in this system, but not detected under the sameexperimental conditions in the cobalt(II)±folic acidsystem, is probably associatedwith the di�erence betweenthe water exchange rate in the coordination spheresof nickel (II) �3:0� 104 sÿ1� and cobalt(II) �2:0 �106 sÿ1��17±19�.The rate constant for the slower process is very low in

comparison with other complexation rate constants.However, the reactions of nickel(II) with 2-2(pyridazo)-1-naphthol dyes (a-PAN)(13) are characterized by

Figure 5. Arrhenius plot for k1app at di�erent nickel concentrations.[NiII] = (d) 0.01; (s) 0.1 M.

Transition Met. Chem., 22, 541±544 (1997) NiII complexation by pteroylglutamic acid 543

abnormally low rate constants (ca. 20 s)1 at 298 K).These values have been ascribed to the formation of aterdentate chelate. In these systems the coordinationpattern resembles that proposed for the nickel(II)±folicacid or cobalt(II)±folic acid complexes (see Scheme 1and Ref. 6).Therefore, this second process allows for a similar

interpretation to that advanced in the cobalt(II)±folicacid system, i.e. the bidentate complex formed in the®rst step of the complexation (NiFoLO)) undergoes amajor conformational change leading to a terdendatespecies(6, 20). This complex is partially reprotonated in afast step as depicted in Scheme 1.

The second process, observed in nickel(II)±folic acidand cobalt(II)±folic acid systems, requires the transfor-mation of a bidentate into a terdentate complex,involving the nitrogen in the pterydin and benzyl rings.This terdentate complex contains two ®ve-memberedrings positioned around the central ion. The intra-molecular distances involving the di�erent groups aresimilar (dOAN � 2:77 ÊA in the pteridyl moietyand dNAN � 2:73 ÊA between N in the pteridyl and N inthe p-aminobenzoic acid moieties), allowing the coor-dination pattern seen around the central ion in this

complex(6,20). The activation energy obtained for theslow process is of the same order of magnitude as thatobserved for the cobalt(II)±folic acid system. Therefore,conformational changes during chelation involving theligand may constitute the rate-limiting step associatedwith the interaction of divalent metal ions and folic acid.

Acknowledgements

This work was supported by the Consejo Nacional deInvestigaciones Cientõ cas y Te cnicas de la Repu blicaArgentina (CONICET), the Comisio n de Investig-aciones de la Provincia de Buenos Aires (CICPBA) andUniversidad Nacional de La Plata. One author (A.T.)thanks CONICET for a graduateship.

References(1) R. Nayan and A. K. Dey, Z. Naturforsch., Teil B, 52, 1453 (1970).(2) M. Poe, J. Biol. Chem., 252, 3724 (1977).(3) M. Poe, J. Biol. Chem., 248, 7025 (1973).(4) A. Albert, Biochem J., 54, 646 (1953).(5) J. T. H. Roos and D. R. Williams, J. Inorg. Nucl. Chem., 39, 367

(1977).(6) A. H. Thomas, M. R. Feliz and A. L. Capparelli, Transition Met.

Chem., 21, 317 (1996).(7) P. K. Chattopadhyay and B. Kratochvil, Can. J. Chem., 54, 2540

(1976).(8) P. K. Chattopadhyay and B. Kratochvil, Inorg. Chem., 18, 2953

(1979).(9) R. G. Wilkins, Comments Inorg. Chem., 2, 187 (1983).(10) R. G. Wilkins, Acc. Chem. Res., 3, 408 (1970).(11) D. M. W. Buck and P. Moore, J. Chem. Soc., Dalton Trans., 2082

(1974).(12) R. B. Jordan and B. E. Erno, Inorg. Chem., 18, 2895 (1979).(13) R. L. Reeves, G. S. Calabrese and S. A. Harkaway, Inorg. Chem.,

22, 3076 (1983).(14) R. G. Wilkins, Kinetics and Mechanism of Reactions of Transition

Metal Complexes, 2nd Edit, VCH, New York, 1991 p. 219.(15) R. M. Fuoss, J. Am. Chem. Soc., 80, 5059 (1958).(16) A. E. Mantell and R. M. Smith, Critical Stability Constants,

Plenum Press, New York, 1975, Vol. 2, p. 165.(17) Y. Ducommun, K. E. Newman and A. E. Merbach, Inorg. Chem.,

19, 3696 (1980).(18) Y. Ducommun, D. Zbinden and A. E. Merbach, Helv. Chim. Acta,

65, 1385 (1982).(19) P. Bernhard, L. Helm, I. Rapaport, A. Ludi and A. E. Merbach,

Inorg. Chem., 27, 873 (1988).(20) E. A. Miromov and V. S. Nabokov, Khim.-Farm. Zh., 10, 136

(1976).

(Received 28 October 1996Accepted 29 April 1997) TMC 3839

Scheme 1.

544 Capparelli et al. Transition Met. Chem., 22, 541±544 (1997)