Interaction between aspergillic acid and iron(III): A potentiometric, UV–Vis, 1H NMR and quantum chemical study

Polyhedron 28 (2009) 763–768

Contents lists available at ScienceDirect

Polyhedron journal homepage: www.elsevier.com/locate/poly

Interaction between aspergillic acid and iron(III): A potentiometric, UV–Vis, 1H NMR and quantum chemical study Valeria M. Nurchi a,*, Guido Crisponi a, Tiziana Pivetta a, Enzo Tramontano b, Flaminia Cesare Marincola a, Joanna I. Lachowicz a a b

Dipartimento di Scienze Chimiche, Cittadella Universitaria, 09042 Monserrato-Cagliari, Italy Dipartimento di Scienze Applicate ai Biosistemi, Cittadella Universitaria, 09042 Monserrato-Cagliari, Italy

a r t i c l e

i n f o

Article history: Received 21 October 2008 Accepted 19 December 2008 Available online 20 January 2009 Keywords: Aspergillic acid Iron Complex formation constants Potentiometric titration UV–Vis spectrophotometry 1 H NMR

a b s t r a c t The complex formation equilibria of aspergillic acid with iron(III) were studied by potentiometric and spectrophotometric methods. The ligand was prepared by a biosynthetic route and its purity checked by MS and 1H NMR spectroscopies. Some structural features of the ligand are discussed on the basis of NMR results. The iron(III) coordination model is compared with those of two other cyclic hydroxamates, 1-hydroxy-2-pyridinone and 1,4-dihydroxy-2-pyridinone, and its features are discussed on the basis of quantum chemical calculations. Ó 2008 Elsevier Ltd. All rights reserved.

1. Introduction Chelating therapy for iron overload drew our interest during the last 10 years, and our research has developed around this topic [1–7]. In this context, we have studied the complex formation equilibria between Fe(III) and ligands that can be potential iron chelators with an interest in evaluating how structural changes or substituents can affect the coordinating properties. The evaluation of structural and substituent effects on the behaviour of ligands is aimed at designing optimal iron chelators that satisfy both the chemical requirements (high complex stability, reaction rates, selectivity, etc.) and the biological restraints (lipophilicity, bioavailability, etc.). In this work, we present a solution equilibrium study of Fe(III) and aspergillic acid. This molecule can be considered a cyclic hydroxamic acid like 1-hydroxy-2-pyridinone, in respect to which it presents a further nitrogen atom in the ring and two aliphatic substituents that confer high lipophilicity.

* Corresponding author. Tel.: +39 0706754476; fax: +39 0706754478. E-mail address: [email protected] (V.M. Nurchi). 0277-5387/$ - see front matter Ó 2008 Elsevier Ltd. All rights reserved. doi:10.1016/j.poly.2008.12.031

N N OH

Aspergillic acid

O

N

O

OH

1-hydroxy-2-pyridinone

The antibiotic properties of aspergillic acid were recognized in 1943 by White and Hill [8] and subsequently its chemical structure determined by Dutcher [9]. In this last work, the binding of aspergillic acid with Fe(III) was noted because of the intense red colour formation upon treatment of a methanolic solution with ferric chloride. Nevertheless, until now, no studies have been presented on formation equilibria with Fe(III), despite the fact that in a very recent paper Kontoghiorghes and Kolnagou [10] refer to aspergillic acid as a naturally occurring iron chelator with very powerful iron binding properties. The reported value (5.3) for its protonation constant [9] is about 1 unit lower than those of 1-hydroxy-2-pyridinone and 1,4-dihydroxy-2-pyridinone, both of which are cyclic hydroxamic ligands that have been thoroughly examined by the two well-known research groups of Martell and co-workers [11,12] and Raymond and

764

V.M. Nurchi et al. / Polyhedron 28 (2009) 763–768

co-workers [13]. This protonation constant, which is lower because of the structural difference due to the second nitrogen atom in the ring and to the aliphatic substituents, led us to study the aspergillic acid–Fe(III) system. In fact, in a recent paper [14], we pointed out how a decrease in the basicity of coordinating groups, which induces a decrease of both protonation and complex formation constants, can be accompanied by an increase of pFe, i.e. of the real chelating ability of the ligand. The major problem we encountered was the commercial unavailability of aspergillic acid, which lead us to prepare it via a biosynthetic method starting from Aspergillus flavus.

2. Experimental 2.1. Reagents NaOH, HCl, HClO4, NaClO4, NaCl, KCl, CHCl3, CDCl3, CH3OH and FeCl3 were purchased from Aldrich. Tryptone was purchased from Microbiol S.n.c. 2.2. Synthesis of aspergillic acid 2.2.1. Growth of Aspergillus flavus The A. flavus strain ATCC 46283 was used in all the experiments (Fig. 1). Spores were grown for 7 days on LB Petri dishes at 25 °C in the dark. An inoculum of 105–106 spores was grown in 1 L of 2% Tryptone, 5% NaCl at 25 °C for 5 days in the dark without shaking. Heavy, white and wrinkled pellicle was formed 48 h after inoculation. After 5 days the pH of the liquid under the pellicle was 7.5–8. The aspergillic acid yield from a litre of culture was 5–10 mg. The Aspergillic acid production did not change significantly with the use of a yeast extract medium and longer growth time. 2.2.2. Isolation of pure aspergillic acid The growth medium after 5 days was filtered to separate the spores from the medium solution. One litre of medium was concentrated by low pressure evaporation to a 100–200 mL final volume. The concentrated medium was then extracted with CHCl3. After chloroform evaporation, the crude aspergillic acid was sublimated at low pressure at 75–80 °C. The isolation of pure product from the culture filtrate performed according to the MacDonald procedure [15] did not allow a significantly higher amount of pure product.

2.3. Purity check The MS experiment was carried out with an Agilent 1200 Series mass spectrometer. Spectra were acquired in the positive ion mode and the solutions were injected into the MS source with a solvent flow rate of 3 lL/min. The temperature of the spray chamber was 200 °C. A capillary 4000 V voltage was applied. Data were obtained in 50–1000 m/z range. Purity was also checked by 1H NMR spectra in CDCl3. 2.4. Potentiometric measurements Potentiometric titrations were performed on CH3OH:H2O 80:20 w/w solutions (1.5 mL total volume) using a MOLSPIN automatic titrator with a Russel CMAW 711 microcombined glass electrode (KCl aqueous solution was replaced with 0.1 M NaCl in CH3OH:H2O 80:20 w/w) at constant temperature (25 °C) and ionic strength (0.1 M NaClO4) under argon flow using NaOH as the titrant. The pH-meter was calibrated daily in terms of hydrogen ion concentration using HClO4 and a pKw 14.42 was established. The purity of the ligand and its exact concentration were determined by the Gran method [16]. NaOH was added using a 0.250 mL micrometer syringe calibrated both by weight titration and by standard material titration. The ligand concentration was 4.7  10 4 M and the metal:ligand ratios were 1:1 and 1:3. Protonation constants and Fe(III) complex formation constants were calculated using SUPERQUAD [17] and HYPERQUAD [18] programs. The standard deviations computed by SUPERQUAD and HYPERQUAD refer to random errors only; they are, however, good indicators of the importance of a particular species at equilibrium [19]. 2.5. UV–Vis absorption spectroscopy Absorption spectra were recorded on a Cary 300 Bio spectrophotometer. The spectra were obtained in the 2–10.5 pH range on solutions containing metal:ligand molar ratios 1:1 and 1:3 at constant ligand concentration 2.4  10 3 M. The molar absorptivities were calculated at the maximum concentration of each species obtained from the potentiometric results. 2.6. NMR measurements 1 H NMR spectra were determined on a Varian Inova-400 spectrometer at 399.94 MHz, with a 5 mm diameter sample tube at 25 °C. 1H NMR spectra were recorded in CDCl3, and chemical shifts were referenced to TMS.

2.7. Quantum chemical calculations Theoretical calculations were performed with the SPARTAN’06 program [20]. The geometries were optimised using semi-empirical [21] methods, with the AM1 [22] and AM1/SM2 [23,24] for ligands and PM3 [25] basis sets for metal complexes.

3. Results and discussion 3.1. Pure ligand

Fig. 1. The Aspergillus flavus ATCC 46283 on the LB Petri dish.

The purity of aspergillic acid obtained with the procedure described in the Section 2 was checked using MS and 1H NMR spectra. The MS spectrum of pure aspergillic acid (C12H20N2O2), showing an ion at 225.1585 m/z and isotope forms at 226.1618 and 227.1562 m/z, is shown in Fig. 2. The aspergillic acid exists also in the minor oxidised form at 223.1439 m/z. The signal due to the isotope form of the oxidised species at 225.1505 m/z is masked

765

V.M. Nurchi et al. / Polyhedron 28 (2009) 763–768

Fig. 2. MS spectrum of aspergillic acid.

Table 1 Experimental and calculated MS data for aspergillic acid. Molecular formula

Experimental m/z

Calculated m/z

C12H20N2O2

225.1585 226.1618 227.1562

225.1602 226.1633 227.1661

223.1439 224.1564

223.1446 224.1477 225.1505

C12H19N2O2

H

14

CH3

6

10

6 7 8 9 10 11 12 13,14

7.284 3.085 1.315 1.747 0.967 2.691 2.226 0.867

data

da (ppm)literature

data

[26]

7.26 3.06 1.29 1.67 0.87 2.66 2.16 0.86

H

11

H

H3C H3C

da (ppm)experimental

11

N

H 8 7

9H

Proton number

Proton numbers according to Fig. 3. a The 1H NMR shifts are referred to TMS.

12 13

H3C

Table 2 Experimental and literature-based chemical shifts of 1H NMR spectra of aspergillic acid.

H H

N O

by the signal at 225.1585 m/z of aspergillic acid. All the obtained results are in good agreement with the freeware ISOPRO 2.1 computer calculations (Table 1) and confirm that the synthesized product was aspergillic acid. The 1H NMR spectrum of aspergillic acid (Fig. 3) also confirms the purity of product. The d (ppm) values of the signals agree with those reported by Brotzu et al. [26], even though a different appearance of the 400 MHz spectrum with respect to that at 100 MHz presented in the cited paper. In Table 2 1 H NMR data with their relative assignments are reported. As can be observed in Fig. 3, the signals of the different protons are split in two distinct parts that can be attributed to two different isomers of aspergillic acid. Impurities of the same magnitude order of the main molecule can in fact be ruled out on the basis of the mass spectrum in Fig. 2. We can then hypothesize, in agreement with Brotzu et al. [26], that there exist two isomers in which the two

O 4

H

9

Fig. 3. 1H NMR spectrum of aspergillic acid at 400 MHz.

A 6000

B

4000

Normalized peak heights

Molar absorptivity

pH = 9.85 pH = 3.07

2000

0

1.0

0.8

0.6

0.4

0.2

0.0

280

320 360 Wavelength (nm)

400

2

4

6 pH

8

10

Fig. 4. (A) Absorptivity spectra of aspergillic acid in the acidic and basic forms. (B) Distribution plots of aspergillic acid calculated with log K = 6.07 potentiometrically determined (continuous lines), and normalized heights (symbols) of the peaks at 308 nm (acidic form HL) and 344 nm (basic form L ).

766

V.M. Nurchi et al. / Polyhedron 28 (2009) 763–768

alkylic groups are interchanged. This is supported by the findings of MacDonald [27] who, working with radioactive 14C leucine and isoleucine, demonstrated that these compounds were incorporated in the formed aspergillic acid. It would be thus equivalent to obtain one or the other of the two isomers, since the aspergillic acid biosynthesis occurs from these two amino acids. Aspergillic acid is a monoprotic acid (HL) characterized by a hydroxamic acid proton. The potentiometric titration in methanol:water 80:20 w/w solution shows one protonation with log K 6.07, which seemingly differs from the log K 5.3 reported by Dutcher [9], or from the log K 5.5 reported in the Merck Index [28]. It should be noted that the medium is reported neither by Dutcher nor by the Merck Index. The log K value of 6.07 obtained by potentiometric titrations is supported by the results of spectrophotometric measurements. Twelve spectra were collected in the pH interval 2–10, of which those at pH 3.07 and 9.85 are reported in Fig. 4A as an example. The 12 spectra were decomposed into the component Gaussian peaks with the SPECPEAK program [29]. The normalized heights of the band centred at 308 nm (attributable to the acidic form HL) and of the band at 344 nm (attributable to the basic form L ) are reported as symbols in Fig. 4B. On the same figure, the distribution plots are also reported as continuous lines, calculated with the log K = 6.07 obtained from potentiometric measurements; a perfect agreement can be observed. 3.2. Iron complexes

Table 3 Complex formation constants estimated from potentiometric and spectrophotometric measurements for the Fe(III)–aspergillic acid system*. Wavelengths of maxima and absorptivities for ligand and all complex species. Species

Log b

k (nm)

e (cm

HL

6.07(4)

FeL

7.92(2)

324 (HL) 339 (L ) 310 464 310 311

5400 5900 6000 772 6255 6872

FeH FeH

1L

3.69(5) 0.13(5)

2L

* Hydroxide formation constants were FeH 1 2.56, FeH FeH 4 21.88, Fe2H 2 2.84 and Fe3H 4 6.05 [34].

2

6.20, FeH

1

3

mol

1

L)

11.41,

the formation constant 7.92 for FeL2+ complex. The potentiometric data relative to the titrations at different molar ratios were then processed with the HYPERQUAD program [18] keeping fixed the value 7.92 for FeL2+ formation. The complex formation model reported in Table 3 with the relative stability constants was obtained, and the distribution curves are illustrated in Fig. 6. The chelating behaviour of aspergillic acid can be compared with those of 1-hydroxy-2pyridinone and 1,4-dihydroxy-2-pyridinone (Table 4 and Fig. 7), both of which are characterized by a hydroxamic group inserted into an aromatic ring. The distribution curves together with the numerical values enable a direct comparison of the formation constants, taking into account the proton competition and the different stoichiometries. As far as the 1-hydroxy-2-pyridinone is

The complex formation equilibria between aspergillic acid and Fe(III) were studied both potentiometrically and spectrophotometrically. In fact, since iron was almost completely complexed at the starting of the potentiometric titration, the first complex formation constant of ML species had to be determined only spectrophotometrically. A set of 12 solutions in the 1–2.5 pH range was prepared by additions of HCl, then spectra were collected. In all the work pH represents –log[H+] at constant ionic strength 0.1 M; in strongly acidic solutions the ionic strength 0.1 M was reached by proper amount of HCl and NaCl. The spectra show a band centred at 470 nm of increasing intensity, whose appearance changes when the pH further increases (Fig. 5). The visible bands in the molar ratio 1:3 solutions progressively disappear after pH 6 with an increase of the baseline, indicating the breaking of iron complexes with the formation of iron hydroxides and free ligand. In solutions with a 1:1 molar ratio, the precipitation of iron takes place at pH 4. The analysis of the spectrophotometric data in the pH range 1– 2.5 with the SPECFIT program [30–32] allowed the evaluation of

Abundance % relative to Fe(III)

100

Fe(III)

FeH-3 FeL

80

FeLH-2

60 40 FeH-4

FeLH-1

20

FeH-1 FeH -2 0 0

2

4

6

8

10

pH Fig. 6. Distribution plot of the Fe(III)–aspergillic acid system at 1:3 metal/ligand ratio, with a total iron concentration of 1.5  10 4 M, and a total ligand concentration of 4.7  10 4.

0.8

pH 9.7 Table 4 Literature protonation and iron(III) complex formation constants for 1-hydroxy-2pyridinone and 1,4-dihydroxy-2-pyridinone.

Absorbance

0.6

pH 5.8

Species*

0.4

[12]

pH 2.5 0.2

pH 0.8 0 400

440

1-Hydroxy-2-pyridinone

480 520 Wavelength (nm)

560

600

Fig. 5. Visible spectra of the system Fe(III)–aspergillic acid with a total iron concentration of 8  10 4 M and a total ligand concentration of 2.4  10 3, optical path 0.5 cm.

LH LH2 LH3 MLH ML ML2H2 ML2H ML2 ML3H2 ML3H ML3 *

5.86 7.06

10.61

[13] 5.78 4.88

8.28 14.42 14.54 18.5

10.3

20.11

19.3

27.21

26.9

Charges were omitted for clarity.

1,4-Dihydroxy-2pyridinone [11]

34.26 25.62 21.31 41.35 35.02 28.18

767

V.M. Nurchi et al. / Polyhedron 28 (2009) 763–768

concerned, we remark the significative difference between the second protonation constants of Li and Martell [12] and Scarrow et al. [13], log K2 1.2 and 0.9, respectively, that at acidic pH values greatly affects the distribution plots for the pure ligand and for iron complexes. The found differences are to be ascribed to the fact that potentiometric and/or spectrophotometric measurements were used to determine the protonation constant bH2L and the complex formation constant bFeL. In the actual cases in which the equilibria lie outside the range of potentiometric measurements, the use of spectrophotometric measurements allows to obtain more reliable results. At any rate, the formation of the FeL2+, FeL2+ and FeL3 complexes is unquestionable, as for most hydroxypiridinones, with the last complex completely formed at pH 4. More interesting is the case of 1,4-dihydroxy-2-pyridinone that results in the formation of FeLH2+, FeL2H2+ and FeL3H2 (similar to the previous FeL2+, FeL2+ and FeL3 complexes with the phenolic group in the protonated form). These occur at pH values higher than those observed in the case of 1-hydroxy-2-pyridinone complexes. In our opinion, this is to be ascribed to the fact that the two protonation constants are in reality macroconstants, and that the protonation of the

LH2+

100

L-

LH

FeL2 % formation relative to Fe(III)

% formation relative to ligand

100

hydroxamic group takes place at a pH higher than in the case of 1-hydroxy-2-pyridinone. The proton competition with iron for the hydroxamic group is consequently higher and the complex formation is delayed. The binding modes of aspergillic acid are completely different from those of the other cyclic hydroxamates; in fact, it does not form complexes of 1:2 and 1:3 stoichiometries, but the FeL2+ complex stabilizes with the loss of two protons from the coordinated water molecules, with pK values of 3.8 and 4.2. This is in strong contrast with the behaviour of the two hydroxypyridinones, especially taking into account the medium used. In order to find an explanation for this different behaviour of aspergillic acid, we performed a quantum mechanical study of the most stable isomer of the complexes with different stoichiometries, i.e. FeL2+ and FeL3 on one side, and FeLOH+ and FeL(OH)2 on the other. Apart from the energetic results, which are not directly comparable, the analysis of the structures presented in Fig. 8 allows us to obtain some significant insight: the FeL2+ and above all the FeL3 structures show that hydrophobic cages around the central metal ion are formed when aspergillic acid coordinates FeL2+, at the same time neutralizing the charge of the complex. This fact, also observed

80 60 40 20

FeL3

80 60 40

FeL

FeH-3

20 FeH-4

Fe(III)

0

0

0

2

4

6

8

10

0

2

4

pH

10

FeL3

80 60 40 20

80

FeL2

60 40 FeH-3

FeL 20

FeH-4

0

0

0

2

4

6

8

10

Fe(III) 0

2

4

pH 100

LH-

LH2

100

L2-

80 60 40

+

LH3

20 0 2

4

6 pH

8

10

FeLH FeL2H2 FeL3

80 FeL3H2

60

FeL3H 40 FeH-3

20

8

10

0

FeH-4

FeL2

Fe(III)

0

0

6 pH

% formation relative to Fe(III)

% formation relative to ligand

8

100

L-

LH

% formation relative to Fe(III)

% formation relative to ligand

100

6 pH

2

4

6

8

10

pH

Fig. 7. Protonation and Fe(III) complex distribution plots for system at a total metal concentration of 1  10 3 M and a total ligand concentration of 3  10 3 M; in the upper plots the ligand is 1-hydroxy-2-pyridinone, stability constants from Ref. [12]; in the medium plots the ligand is 1-hydroxy-2-pyridinone, stability constants from Ref. [13] and in the lower plots the ligand is 1,4-dihydroxy-2-pyridinone, stability constants from Ref. [11].

768

Fig. 8.

V.M. Nurchi et al. / Polyhedron 28 (2009) 763–768

SPARTAN’06

calculated structures of the most stable isomer of the FeL2+ (upper left), FeL3 (upper right), FeLOH+ (lower left), and FeL(OH)2 (lower right) complexes.

in the low water medium used in our study, prevents a positive interaction of the FeL2+ and FeL3 complexes with the solvent, and therefore, creates unfavourable conditions for the formation of such complexes. On the other hand, the formation of the hydroxo complexes FeLOH+ and FeL(OH)2 stabilizes the already hydrophobic 1:1 complex FeL2+ by creating positive interactions of the coordinated OH groups via hydrogen bonds with the solvent. Summing up, even though aspergillic acid cannot be classified as an iron chelator and taken into account for the treatment of iron overload, its characteristic coordination model is useful to clearly demonstrate the effect of a hydrophobic side chain in determining the coordination modes of the ligand. Acknowledgements V.M.N. thanks Fondazione Banco di Sardegna for the financial support. J.I.L. is in-debt with Prof. H. Kozlowski for the availability of the Molspin titrator and for the helpful discussions and with N. Zinnarosu for NMR spectra collection. All the authors express gratitude to Prof. S. Cosentino and Dr. B. Pisano for their useful advice and assistance in biological procedures. References [1] G. Faa, G. Crisponi, Coord. Chem. Rev. 184 (1999) 291. [2] G. Crisponi, V.M. Nurchi, R. Silvagni, G. Faa, Polyhedron 18 (1999) 3219. [3] E. Gumienna-Kontecka, R. Silvagni, R. Lipinski, M. Lecouvey, F. Cesare Marincola, G. Crisponi, V.M. Nurchi, Y. Leroux, H. Kozlowski, Inorg. Chim. Acta 339 (2002) 111. [4] G. Crisponi, V.M. Nurchi, T. Pivetta, J. Inorg. Biochem. 102 (2) (2008) 209. [5] V.M. Nurchi, G. Crisponi, T. Pivetta, M. Donatoni, M. Remelli, J. Inorg. Biochem. 102 (4) (2008) 684.

[6] G. Crisponi, V.M. Nurchi, T. Pivetta, J. Gałe˛zowska, E. Gumienna-Kontecka, T. Bailly, R. Burgada, H. Kozłowski, J. Inorg. Biochem. 102 (7) (2008) 1486. [7] G. Crisponi, M. Remelli, Coord. Chem. Rev. 252 (2008) 1225. [8] E.C. White, J.H. Hill, J. Bacteriol. 45 (1943) 433. [9] J.B. Dutcher, J. Biol. Chem. 171 (1947) 321. [10] G.J. Kontoghiorghes, A. Kolnagou, Curr. Med. Chem. 12 (2005) 2695. [11] E.T. Clarke, A.E. Martell, Inorg. Chim. Acta 196 (1992) 185. [12] Y.J. Li, A.E. Martell, Inorg. Chim. Acta 214 (1993) 103. [13] R.C. Scarrow, P.E. Riley, K. Abu-Dari, D.L. White, K.N. Raymond, Inorg. Chem. 24 (1985) 954. [14] V.M. Nurchi, G. Crisponi, T. Pivetta, J.I. Lachowicz, J. Inorg. Biochem., doi:10.1016/j.jinorgbio.2008.10.011. [15] J.C. MacDonald, Can. J. Biochem. 31 (1973) 1311. [16] G. Gran, Analyst 77 (1952) 661. [17] P. Gans, A. Sabatini, A. Vacca, J. Chem. Soc., Dalton Trans. (1985) 1195. [18] P. Gans, A. Sabatini, A. Vacca, Talanta 43 (1996) 1739. [19] D. Valensin, M. Luczkowski, F.M. Mancini, A. Legowska, E. Gaggelli, G. Valensin, K. Rolka, H. Kozlowski, J. Chem. Soc., Dalton Trans. (2004) 1284. [20] SPARTAN’06, Wavefunction, Inc. [21] J.A. Pople, D.A. Beveridge, Approximate Molecular Orbital Theory, McGrawHill, New York, 1970. [22] M.J.S. Dewar, E.G. Zoebisch, E.F. Healy, J.J.P. Stewart, J. Am. Chem. Soc. 107 (1985) 3902. [23] C.C. Chambers, G.D. Hawkins, C.J. Cramer, D.G. Truhlar, J. Chem. Phys. 100 (1996) 16385. [24] C.J. Cramer, D.G. Truhlar, Structure and Reactivity in Aqueous Solution, in: ACS Symposium Series, 568, American Chemical Society, Washington D.C., 1994. [25] J.J.P. Stewart, J. Comput. Chem. 10 (1989) 209. [26] G. Brotzu, G. Crisponi, R. Garzia, A. Lai, G. Paschina, Z.L. Rossetti, G. Saba, Rendiconti Seminario Facoltà Scienze, Cagliari University, XLV (1975) 225. [27] J.C. MacDonald, J. Biol. Chem. 236 (1961) 512. [28] The Merck Index, 11th ed., Rahway, 1990. [29] M.C. Aragoni, M. Arca, G. Crisponi, V.M. Nurchi, Anal. Chim. Acta 316 (1995) 195. [30] H. Gampp, M. Maeder, Ch.J. Meyer, A.D. Zuberbuhler, Talanta 32 (1985) 1133. [31] H. Gampp, M. Maeder, Ch.J. Meyer, A.D. Zuberbuhler, Talanta 33 (1986) 943. [32] H. Gampp, M. Maeder, Ch.J. Meyer, A.D. Zuberbuhler, Anal. Chim. Acta 193 (1987) 287. [34] C.F. Baes, R.E. Mesmer, The Hydrolysis of Cations, John Wiley and Sons, Inc., New York, 1976.