Copper(II) mixed ligands complexes of hydroxamic acids with glycine, histamine and histidine

Article

Copper(II) Mixed Ligands Complexes of Hydroxamic Acids with Glycine, Histamine and Histidine

Maria Celina M.M. Fernandesª, Eucler B. Paniago^b, and Sandra Carvalhoª

^a Depto. de Química-ICEx, UFMG

^bDepto. de Química-ICEB, UFOP

Received: December 19, 1996

Um estudo potenciométrico e espectrofotométrico de complexos mistos de Cu^II, fisiologicamente interessantes, foi realizado envolvendo ácidos hidroxâmicos como ligantes primários (A) e ligantes secundários (B) representados por histamina ou pelos aminoácidos glicina e histidina. Todos são potencialmente capazes de formar quelatos com anéis de cinco ou seis membros. A constante de formação e o espectro de absorção no visível são calculados para cada uma das espécies identificadas, tanto binárias quanto ternárias. As estruturas mais prováveis das espécies mistas são propostas com base nas constantes de formação calculadas, na estabilização em relação às duas espécies binárias correspondentes, bem como nos espectros de absorção no visível. Ficou comprovado que os complexos de Cu-hidroxamato não favorecem significativamente a formação de complexos ternários com aminoácidos que possuem apenas os grupos amino e carboxilato; contudo, a formação de complexos ternários com a histidina é fortemente favorecida, devido à presença do grupo imidazol.

A potentiometric and spectrophotometric study of physiologically interesting Cu^II mixed ligand complexes has been performed involving hydroxamic acids as primary ligands (A) and secondary ligands (B) represented either by histamine or the aminoacids glycine and histidine. All are potentially able to form chelate complexes with either five or six membered rings. The formation constant and the visible absorption spectrum were calculated for each one of the identified species, both binary and mixed ones. The most probable structures of the mixed species are discussed based upon their formation constants, their stabilization with regard to the two binary species as well as their visible absorption. It was evidenced that Cu-hydroxamate complexes do not favour significantly mixed ligand complex formation with aminoacids having only the carboxylate and amino groups; however, these complexes are strongly formed with histidine, due to the presence of the imidazol group.

Keywords: hydroxamic acid, copper(II) complexes, mixed ligand complexes, histidine, glycine

Introduction

Hydroxamic acids having one or more -CONHOH groups have been extensively studied as a consequence of their biological importance which is related with their ability to form metal ion complexes. These acids are important also due to their pharmacological, toxicological and pathological properties¹. Besides, they have a great number of applications in analytical chemistry^2,3.

Monohydroxamic acids (such as acetohydroxamic acid, CH₃CONHOH=Aha) after deprotonation acts as bidentate ligands forming octahedral complexes with a series of metal ions via co-ordination through the two oxygen atoms of the -CONHO^- group. This type of co-ordination has been characterized in previous studies with Fe^III, Cr^III, Ni^II, Co^II, and Zn^II ions, which indicated the formation of octahedral complexes both in the solid state and in solution⁴. With Cu^II ions, the co-ordination again occurs via the two oxygen atoms but the geometry is tetragonal⁴.

Aminohydroxamic acids, due to the protonation of the amino group, have intricate acid base properties as shown by Farkas and Kiss⁵. Anions of a-aminohydroxamic acids (such as glycinehydroxamic acid, NH₂CH₂CONHOH=Gha) are potentially able to form three kinds of a five member chelate complex using either: 1 - the two oxygens of the hydroxamate group, 2 - the two nitrogen atoms or, 3 - the nitrogen atom of the amino group and the oxygen of carbonyl group. As a result, they are powerful ligands for a great number of metal ions.

This work reports the results of a combined potentiometric and spectrophotometric study of Cu^II mixed ligand complexes formed by one hydroxamic acid (Aha and Gha) taken as the primary ligand (A) and a secondary ligand (B) represented either by histamine (Ha) or by the aminoacids glycine (Gly) and histidine (His). All these ligands are potentially able to form chelate complexes with either five or six membered rings, displaying different co-ordination sites. Histamine has a N(amino) and a N(imidazol) donor groups, glycine has N(amino) and O(carboxylate) ones, while histidine has these three kinds of donor groups all together: one oxygen of the carboxylate group, one nitrogen of the imidazol group and the amino nitrogen atom. It may, therefore, act as a tridentate ligand while coordinating metal ions with octahedral geometry⁶. With the Cu^II ion it is impossible to have the three donors atoms of histidine simultaneously occupying planar positions⁶. Two of these donor atoms will occupy equatorial positions, while the third one will either remain uncoordinated or occupy an axial position. The importance of the Cu^II histidine co-ordination to explain the transport of copper in the human plasma has led to a series of investigations (involving potentiometry, spectroscopy in the UV, visible, infrared and Raman, calorimetry, X-ray and EPR) attempting to elucidate the structures of several Cu(II)-histidine complexes which are formed in aqueous solutions. However a great deal of controversy still remains^7-13.

The formation constant (log b) and the visible absorption spectrum were calculated for each one of the identified species, both binary and mixed ones. The most probable structures of the mixed species are discussed based upon their formation constants, their stabilization with regard to the two binary species and to their visible absorption spectra.

Experimental

Reagents and Materials

Histamine (dihydrocloride 99%), glycinehydroxamic (90%) and acetohydroxamic (95 to 98%) acids were Sigma products. Glycine and histidine were Merck biochemical grade reagents. Solutions of these ligands were always prepared just before use with bidestilled water and their concentrations checked by direct potentiometric titrations.

A stock solution of copper(II) perchlorate was prepared from copper(II) nitrate and perchloric acid and its concentration determined by EDTA titration, using a copper ion-selective electrode as an indicator¹⁴.

Sodium perchlorate solution, used to maintain constant the ionic strength at 0.1 mol/L, was prepared from sodium carbonate and perchloric acid. The concentration of this solution was determined by titrating the acid eluted after its treatment with a cationic resin (Dowex 50W-X8 Backer), loaded with H⁺ ions, with 0.10 mol/L carbonate free sodium hydroxide solution.

Carbonate free sodium hydroxide solutions were prepared with bidestilled water and standardized with potassium hydrogen phthalate.

All titrations were followed potentiometrically.

General procedure

To determine both the formation constant and the absorption spectrum of a given ternary or mixed species M_pA_qB_rH_s (as define in expression 1) it is necessary to determine first the absorption spectra of the metal ion M and of both ligands A and B and of their protonated species, followed by the determination of the formation constant and the absorption spectrum of each binary species formed by M with both A and B. Therefore, a series of steps are necessary which can be grouped as follows:

1 - Potentiometric titration of an acidified solution (with a strong mineral acid) of each ligand to determine the formation constant of each protonated species, under the same conditions which are being used (temperature, ionic strength, etc.), followed by a simultaneous potentiometric and spectrophotometric titration of the same solution to determine the absorption spectra of the free ligand and of each protonated species;

2 - Potentiometric titrations of solutions containg each ligand and copper(II) ion in several proportions, to determine the formation constants of each binary species M_pA_qH_s and M_pB_rH_s, followed by simultaneous potentiometric and spectrophotometric titration of the same solutions to determine the absorption spectrum of each binary species M_pA_qH_s and M_pB_rH_s;

3 - Potentiometric titrations of solutions containing two ligands and the ion copper(II) in different proportions to determine the formation constants of each ternary species M_pA_qB_rH_s, followed by simultaneous potentiometric and spectrophotometric titrations of the same solutions to determine the absorption spectrum of each ternary species M_pA_qB_rH_s.

Potentiometric measurements

The potentiometric titrations were performed by measuring the electromotive force¹⁵ (emf) of cells of the type (1) using as indicator the Metrohm glass electrode 6.0102.100(PB) and as reference the Metrohm calomel electrode 6.0702.100 (NaCl 0.1 mol/L solution). The Titroprocessor Metrohm 670 was used to measure the electromotive force.

Solutions were added using a Metrohm Dosimat 665 autoburette, except the ligand solutions which were measured with a volumetric pipet. The ionic strength was adjusted to 0.10 ± 0.01 mol/L with a NaClO₄ solution and the temperature was maintained constant inside the cell at 25.0 ± 0.1 °C, by circulating thermostated water. The Metrohm 649 magnetic stirrer was used and all measurements were done under a nitrogen atmosphere. Standardized, carbonate free, sodium hydroxide solutions were used in all titrations.

The electrode system (1) was calibrated for hydrogen ion concentration¹², before and after each series of measurements, by titration of perchloric acid (or sodium hydroxide) with standardized sodium hydroxide solution (or perchloric acid solution) at ionic strength of 0.10 ± 0.01 mol/L. Data were treated with a computer program¹⁶ which besides calculating the standard potential of electrode (E⁰) gives also the calculated water dissociation constant (Kw) under the experimental conditions which were being employed. These values were then used to calculate the hydrogen ion concentration (pH = - log H⁺) from measured potentials.

Treatment of potentiometric data

The equilibrium reactions in the systems which were investigated can be represented by the generic expression (charges omitted for simplification):

The formation constant of the species M_pA_qB_rH_s where M = Cu^II ion; A = ligand A; B = ligand B; H = proton and p, q, r and s represent their stoichiometric coefficients, is being defined as:

Tables 1 and 2 report details of the corresponding experimental conditions, i.e. the initial total concentrations of the reactants, number of titration points included and the pH ranges investigated. The electrode system being calibrated in the concentration scale, pH should be understood in terms of -log H⁺ throughout this study.

Formation constants were calculated using the computer program SUPERQUAD¹⁷ and are show in Tables 3 and ⁵. The standard deviations computed by this program (which refer to random errors only) is being used as an evidence of the presence of one species and of the adequacy of the equilibrium model tested.

Spectrophotometric measurements

The solutions while being potentiometrically titrated were continuously circulated, using a peristaltic pump, through a continuos flow quartz cuvette installed in a model 8451A Hewlett Packard spectrophotometer. This setup allowed simultaneous measurements of the volume of titrant (NaOH) added, of the hydrogen ion concentration and of the absorbance of the solution. Titrations were performed with increments of volume and time intervals both pre-fixed (known as monotonous titration) in order to obtain an absorbance spectrum and the hydrogen ion concentration after each addition of the titrating solution. Spectrophotometric data were treated with the program SQUAD¹⁸.

The measured absorbance (A), in a given wavelength, at each point during the titration must take into account all the absorbing species, which includes the free Cu^II ion, each free ligand and all its protonated derivatives as well as the various binary and mixed ligand Cu^II complexes. If Beers law is valid, then for each solution at any given wavelength, the absorbance (A) is defined by the equation:

where: b_pqrs and e_pqrs are the overall formation constant and the molar absorptivity at the given wavelength of the species M_pA_qB_rH_s, respectively, and l is the pathlength of the cell used.

The values of -log H were obtained from potentiometric measurements and since values of b_pqrs had been already calculated by treating the potentiometric data, the values of e_pqrs could then be calculated. This was done stepwise. First, those of the free Cu^II ion were determined separately. Since the measurements were made in the visible region of the spectra, e of the ligands were made equal to zero. Finaly, e of each one of the binary species were determined by titrating a mixture of Cu^II ion with each of the ligands separately.

Once the binary complexes formation constants were determined, solutions containing two ligands in the presence of copper(II) were titrated first potentiometrically and then spectrophotometrically. There was always an excess of ligand with respect to the metal ion and more than one ratio between the two ligands was used.

Table 4 reports the essentials of the corresponding experimental conditions: the initial total concentrations and the pH ranges investigated. The data treated in each case involved a total of 100 wavelengths (from 400 to 800 nm) and a significant number of spectra, as shown in this table. The spectrophotometric data are show in ^{Table 5}.

Results and Discussion

Binary systems

Ligand protonation constants were potentiometrically redetermined by titrating acidified 2.5 x 10^-2 mol/L solutions of the ligands with 0.1 mol/L NaOH solution. The results, shown in Table 3, are in good agreement with previously reported values^4,19.

Binary complexes formation constants were also redetermined potentiometrically from titration data of solutions containing the ligands in the presence of copper(II), using always an excess of ligand and more than one ligand to metal ratio. The results, also shown in Table 3, are in agreement with literature values^4,19. Treatment of the spectrophotometric data by SQUAD also allowed the calculation of the visible absorption spectrum of all the species which are formed in these systems, as shown in Figs. 1-5.

Bidentate ligand co-ordination in these complex species is either already known or foreseeable⁶.

Cu^II-acetohydroxamic acid: the anion of this acid (pKa = 9.230) co-ordinates through the two oxygen atoms of the hydroxamate group. Hence, spectra 1 and 2 of Fig. 1 are characteristic of an aqueous solution of Cu^II ions bound to one (Cu(Aha)⁺, l_max = 760 nm) and two hydroxamate groups (Cu(Aha)₂, l_max = 660 nm), respectively.

Cu^II-glycinehydroxamic acid: this water insoluble acidcan be dissolved either in acidic or basic solutions, as a result of protonation of the amino group (pKa = 7.615) or deprotonation of the hydroxamic acid group (pKa = 9.159). Even though it is the OH of the CONHOH group that deprotonates, the glicinehydroxamate anion (Gha^-) co-ordinates to Cu^II ions via the two nitrogen atoms, from the amino and the hydroxamate groups^20a,20b. Spectra 1 (Cu(Gha)⁺, l_max = 684 nm) and 2 (Cu(Gha)₂, l_max = 540 nm) of Fig. 2 are therefore characteristic of aqueous Cu^II bound to one and two such groups. Deprotonation of Cu(Gha)₂ (pKa = 9.85) giving the species Cu(Gha)₂OH^- is the result of the hydroxamate group loosing one more proton as has been already shown by X-ray analysis for a Ni(II) analogue complex^20c. Spectrum 3 of this figure shows a resulting shift in l_max from 540 to 504 nm. For the dimeric species Cu₂(Gha)₂OH⁺, (l_max = 644 nm as shown by spectrum 4 of this figure) we have proposed^20a one OH bridging two Cu-Gha units, a hypothesis which is contested by Farkas and col⁵.

Cu^II-glycine: the glycinate anion co-ordinates via the amino (pKa = 9.516) and the carboxylate groups (pKa = 2.302). Therefore, spectra 1 and 2 of Fig. 3 shows the spectra of Cu^II bound to one (Cu(Gly)⁺, l_max = 732 nm) and two glycinate groups (Cu(Gly)₂, l_max = 624 nm), respectively.

Cu^II-histamine: The crystal structure of Cu(Ha)₂(ClO₄)₂ has been determined and no unusual features were found, the nitrogen atoms from two planar imidazol rings occupy trans co-ordination positions²¹. But histamine may also act as a monodentade ligand in which case the imidazol group remains protonated as well as a bidentate one when coordinating via the nitrogen atoms of the amino and imidazol groups. Therefore, spectra 1 and 2 of Fig. 4 shows Cu^II bound to one (Cu(Ha)²⁺; l_max = 680 nm) and two bidentate histamine molecules (Cu(Ha)₂²⁺; l_max = 596 nm), while spectrum 3 of this figure shows Cu^II, as Cu(Ha)₂H³⁺ (l_max = 624 nm), bound simultaneously to a bidentate and a monodentate histamine molecule.

Cu^II-histidine: In histidine, the carboxylate group (pK_a = 1.93), the imidazol nitrogen (pK_a = 6.08) and the amino nitrogen (pK_a = 9.05) become available for co-ordination as the pH increases, as shown by the Scheme 1⁷.

Results of X-ray studies have shown that histidine can use each one of the three possible co-ordination sites for bonding to metal ions¹⁰. Bidentate chelation using either an imidazol-amino or an amino-carboxylate co-ordination pair are feasible. The possibility of imidazol-carboxylate bidentate co-ordination is disfavoured by the seven member ring involved.

There has been controversy regarding the solution structure of Cu(His)H²⁺. Coordination of both the amino and imidazol groups is popular⁷, but a glycine-like coordination has also been proposed²². This structure seems to us the most probable since there is close similarity between the spectrum of this species, shown in Fig. 5 (l_max = 716 nm), and the one of Cu(Gly)⁺ (l_max = 732 nm). Therefore there is co-ordination in this complex of the amino and of the carboxylate groups to the Cu^II ion, the imidazol group remaining protonated, as shown in Structure 1. This is the same type of co-ordination observed in crystals of Cu(HisH)₂²⁺, as shown by X ray analysis²³.

The absorption spectrum of Cu(His)⁺ (l_max = 664 nm) resembles that of Cu(Ha)²⁺ (l_max = 680 nm) and it is generally agreed that in this species, histidinate bonding to the Cu^II ion occurs via two nitrogen donor atoms in the co-ordination plane, with an axial carboxylate weakly co-ordinated or uncoordinated, as shown in Structure 2.

It has been suggested⁷ that in Cu(His)₂H⁺ (l_max = 608 nm) co-ordination occurs through a combination of the ones occurring in both CuHis⁺ and CuHisH²⁺. However Structure 3, which has already been suggested²⁴, seems the most probable, considering the close similarity between its spectrum and the one of Cu(Ha)₂²⁺.

Several propositions have been made for the structure of the species Cu(His)₂ (l_max = 640 nm). Attempts to isolate Cu(l-His)₂ in neutral pH were unsuccessful. Camerman¹⁷ has succeeded in crystallizing Cu(l-His)(d-His) and has shown that, in the solid state, co-ordination occurs via the nitrogen atoms of the amino and imidazol groups. Two water molecule are co-ordinated to the Cu(II) axial positions and they are H-bonded to carboxylate groups of each histidine molecule. Considering the shift in the spectrum of 608 nm of the protonated species to 640 in this neutral one, it does not seem reasonable to propose that this arrangement remains in solution. Actually, arguments have already been presented by Sigel and McCormick, considering this species to consist of two histidine anions with trans amino groups, the tetragonal plane consisting of a substituted glicinate from one histidine and two nitrogens from the other, as shown in Structure 4. This structure would be dynamic with Jahn-Teller distortions carrying first one set of trans imidazol-carboxylate groups into an axial position and then the other²⁵.

Finally, for the species Cu(His)₂OH^- (l_max = 636 nm), the most probable structure results simply from a loss of a proton from one of the imidazol groups shown in Structure 4.

Mixed ligand systems

A parameter, known as log X, is frequently used to characterise the stability of ternary or mixed complexes. It measures the tendency of one mol each of the binary complexes MA₂ and MB₂ to disproportionate forming two moles of MAB i.e.

It is therefore calculated by: log X = 2 log b_MAB - (log b _MA2 + log b_MB2). The value of the constant X expected on statistical grounds is 4. Whenever it deviates from this value, it must be the result of intraligand electronic and/or steric interactions²⁶.

Whenever possible, this parameter was calculated and the results are shown in ^{Table 5}, where it can be seen that it has always a value greater than the statistically expected one.

In our case, it seemed also relevant to calculate the equilibrium constant corresponding to the addition of ligand B to the 1:1 complex MA

A constant corresponding to this reaction was calculated and is also shown in ^{Table 5}.

Cu^II-acetohydroxamic acid mixed ligand complexes

The value of log X which is equal to 4.65 for Cu(Aha)(Ha)⁺ is almost as large as the one found for the Cu^II mixed complex with histamine and pyrocatecholate, which is 4.86 under the same conditions²⁷. The values of log X for Cu(Aha)(Gly) and Cu(Aha)(His) are smaller, but still much larger than the statistically expected value, demonstrating a significant stabilization of these mixed ligand complexes.

The two mixed complexes with histamine and histidine have almost the same formation constant, while the binary Cu(His)₂complex (log b = 18.220) is much more stable than the Cu(Ha)₂²⁺ (log b = 15.97). But while log X of Cu(Aha)(His) is smaller than the one of Cu(Aha)(Ha)⁺, the value of K_ad of Cu(Aha)(His) is greater than the one of Cu(Aha)(Ha)⁺. Therefore, the greater stabilization of Cu(Aha)(Ha)⁺, when compared to Cu(Aha)(His), can only be attributed to the greater stability of binary histidine complex, which must be destroyed in the disproportionation process that supposedly takes place to form the mixed ligand complex.

Both Cu(Aha)(Gly) (l_max = 652 nm) and Cu(Aha) (Ha)⁺ (l_max = 624 nm) have foreseeable coordinations, shown in Structures 5 and 6. The same does not happen with the histidine complexes Cu(Aha)(His) (l_max = 624 nm) and Cu(Aha)(His)OH^- (l_max = 584 nm). However, considering the coincidence of l_max of Cu(Aha)(Ha)⁺ and Cu(Aha)(His), shown in Fig. 6, it seems reasonable to propose that both have the same co-ordination (shown in Structure 6), the carboxylate group of histidine remaining uncoordinated. As for the structure of Cu(Aha)(His)OH^- (l_max = 564 nm) it is probably similar to the one of Cu(Aha)(His), a proton being lost by the imidazol group of histidine, since both have similar spectra, as shown in Fig. 6.

Cu^II-glycinehydroxamic acid mixed ligand complexes

In this case none of the structures is foreseeable, since while in the mixed complexes with Aha its co-ordination must remain O,O, in the ones with Gha it could be co-ordinated either as a O, O donor (like Aha) or as a N, N, donor like in its binary complexes with the Cu^II ion^20b. Therefore, starting with Cu(Gha)(Gly) (l_max = 596 nm) there is always more than one possibility of co-ordination for the glycinehydroxamate anion. By comparison, Cu(Gha)(Gly) is much more stable than Cu(Aha)(Gly) and they evidently do not have the same structure, since their l_max are quite different. Structure 7, in which Gha is N,N co-ordinated, seems therefore the most probable one.

The Gha complexes with both histamine and histidine, with the exception of Cu(Gha)(His)H⁺, all have quite similar absorption spectra. They have formation constants larger than the corresponding complexes with Aha, but their values of log X, as well as of K_ad, are smaller. Evidently this is a result of a different co-ordination, since there is also a difference in the absorption spectra of Cu(Gha)(Ha)⁺ (l_max = 564) when compared to Cu(Aha)(Ha)⁺ and of Cu(Gha)(His) (l_max = 572) when compared to Cu(Aha)(His). The smaller values of log X can be explained as the result of a larger stability of the Gha-Cu^II binary complexes. Consequently, for these complexes the co-ordination shown in Structure 8 for Cu(Gha)(Ha)⁺ is the most probable, in which the carboxylate group would be co-ordinated axially, to account for the small difference in l_max in these two complexes.

The structure of Cu(Gha)(His)H⁺ (l_max = 612 nm) probably has histidine co-ordinated glycine-like and Gha as in its binary complexes.

A final conclusion can be formulated as follows: Cu-hydroxamate complexes do not significantly favour mixed ligand complex formation with aminoacids having only the carboxylate and amino groups, however these complexes are strongly formed with histidine, due to the presence of the imidazol group.

Acknowledgment

The authors are grateful for financial support from Fundação de Amparo à Pesquisa do Estado de Minas Gerais (FAPEMIG), Conselho Nacional de Desenvolvimento Científico e Tecnológico (CNPq) and Financiadora de Estudos e Projetos (FINEP). During this work one of us (MCF) received a maintenance grant from CNPq.

References

1. Paniago, E.B.; Carvalho, S. Ciência e Cultura 1988, 40, 629.

2. Davidson, D.J. J. Chem. Ed. 1940, 17, 81.

3. Agrawal, Y.K. Rev .Anal. Chem. 1980, 5, 3.

4. Kurzak, B; Kozlowski, H.; Farkas, E. Coord. Chem. Rev 1992, 114, 169.

5. Farkas, E.; Szoke, J.; Kiss, T.; Kozlowzki, H.; Bal, W. J. Chem. Soc. DaltonTrans. 1989, 2247.

6. Sigel, H. In Metal ions in biological systems; Sigel, H., Ed.; Marcel Dekker; New York, v.2, 1973, p.63.

7. Sundberg, R.J.; Martin, R.B. Chemical Reviews 1974, 74, 471.

8. Wilson, E.W.; Kasperian, M. H.; Martin, R.B. J. Am. Chem. Soc. 1970, 92, 5365.

9. Pettit, L.D.; Swash, L.M. J. Chem. Soc., Dalton Trans. 1976, 588.

10. Camerman, N.; Fawcett, J.K.; Kruck, T.P.A.; Sarkar, B.; Camerman, A. J. Am. Chem. Soc. 1978, 100, 2690.

11. Sovago, I.; Kiss, T.; Gergely, A. J .Chem. Soc., Dalton Trans. 1978, 964.

12. Sarkar, B. In Metal ions in biological systems; Sigel, H., Ed.; Marcel Dekker; New York, v. 12, 1981, p. 233.

13. Patel, P.; Bhattacharya, P. K. J. Inorg. Biochem. 1994, 54, 187.

14. Flasckka, H.A. In EDTA Titrations - An Introduction to Theory and Practice, Pergamon Press, London, 1959, p. 78.

15. Molina, M; Melios, C.; Tognolli, J.O.; Luchiari, L.C.; Jafellici, J.R.M. J. Electroanal. Chem. 1979, 105, 237.

16. Duarte, H.A.; Paniago, E.B.; Carvalho, S.; Fernandes, M.C.; Yamaki, R.T. de Resumos - SBQ 1994, QA-119

17. Gans, P.; Sabatini, A.; Vacca, A. J. Chem. Soc., Dalton Trans. 1985, 1195.

18. (a) Leggett, D.J. Anal. Chem. 1977, 49, 276. (b) Leggett, D.J.; Kelly, S.L.; Shine, L.R.; Wu, Y.T.; Chang, D.; Kadish, M. Talanta 1983, 30, 579.

19. (a) Sillen, G. L.; Martell, A.E. Stability constants of metal complexes, The Chemical Society, Special Publication No. 25, Supplement No 1 to Special Publication No. 17; The Chemical Society: London, 1971. (b) Martell, A.E.; Smith, R. M. Critical Stability Constantsv. 1. Amino Acids. Plenum Press: New York,. 1974.

20. (a)Paniago, E.B.; Carvalho, S. Inorg. Chim. Acta 1984, 92, 253. (b) Miranda-Pinto, C.O.; Paniago, E.B.; Tabak, M.; Carvalho, S.; Mascarenhas, Y.P. Inorg.Chim. Acta 1987, 137, 145. (c) Julien-Pouzol, M.; Jaulmes,. S.; Laruelle, P.; Carvalho, S.; Paniago, E.B. Acta Crystallog. 1985, C41, 712.

21. Bonnett, J.J.; Jeannin, Y. Acta Crystallogr., Sect. B 1970, 26, 318.

22. Wilson, E.W.; Kasperian, M.H.; Martin, R.B. J. Am. Chem. Soc. 1970, 92, 5365.

23. Evertsson, B. Acta Crystallogr.Sect. B. 1969, 25, 30.

24. Meyer, J.L.; Bauman, J.E. J. Am. Chem. Soc. 1970, 92, 4210.

25. Sigel, H.; McCormick, D.B. J. Am. Chem. Soc. 1971, 93, 2041.

26. Laurie, S.H.; James, C. Inorg. Chim. Acta, 1983, 78, 225.

27. Huber, P.R.; Griesser, R.; Prijs, B.; Sigel, H. Eur. J. Biochem. 1969, 10, 238.

Publication Dates

• Publication in this collection
  10 Sept 2010
• Date of issue
  1997

History

• Received
  19 Dec 1996