The kinetics and mechanisms of the reactions of aluminium(III) with gallic acid, gallic acid methyl ester and adrenaline

Journal of Inorganic Biochemistry 84 (2001) 1–12 www.elsevier.nl / locate / jinorgbio

The kinetics and mechanisms of the reactions of aluminium(III) with gallic acid, gallic acid methyl ester and adrenaline ´ ´ ´ O’Coinceanainn, Mairtın Michael J. Hynes* Department of Chemistry, National University of Ireland, Galway, Ireland Received 29 August 2000; received in revised form 29 November 2000; accepted 7 December 2000

Abstract The kinetics and mechanisms of the reactions of gallic acid, gallic acid methyl ester and adrenaline with aluminium(III) have been investigated in aqueous solution at 258C and an ionic strength of 0.5 M. A mechanism has been proposed which accounts satisfactorily for the kinetic data. This is consistent with a mechanism in which complex formation takes place almost exclusively by reaction of [Al(H 2 O) 5 OH] 21 with the ligands. [Al(H 2 O) 5 OH] 21 reacts with gallic acid, gallic acid methyl ester and adrenaline with rate constants of 1145, 1330 and 316 M 21 s 21 respectively. These data together with the equilibrium data enable the rate constants for reaction of [Al(H 2 O) 6 ] 31 with both gallic acid and gallic acid methyl ester to be calculated. In view of the dissociative nature of water exchange on [Al(H 2 O) 6 ] 31 and [Al(H 2 O) 5 (OH)] 21 the complex formation rate constants are discussed in terms of the Eigen–Wilkins–Tamm mechanism. The overall mechanisms have been validated using global analysis. The results are compared with previously published data on the complex formation reactions of aluminium(III). In addition, the rate constants and mechanisms for replacement of maltol by gallic acid methyl ester and diethylenetriaminepentaacetic acid (dtpa) have been investigated.  2001 Elsevier Science B.V. All rights reserved. Keywords: Aluminium(III) complexes; Polyphenols; Kinetic studies

1. Introduction Up to twenty years ago aluminium and its salts were described as non-absorbable compounds with few chronic or toxic effects. Because of their assumed safety, aluminium salts were widely adopted for use as food additives, non-prescription drugs and coagulants in water treatment. The discovery of a possible correlation between the aluminium concentration in drinking water and the incidence of Alzheimer’s disease led to a renewed interest in the co-ordination chemistry of aluminium [1]. Green tea (Thea sinensis), a time-honoured drink in Japan for more than 1000 years, is used medicinally and as a refreshment after meals. Recent studies suggest a correlation between the natural anti-oxidants found in green tea and overall good health [2]. The two largest components of green tea are carbohydrates, including cellulosic fiber and protein, both of which are water insoluble. The next largest group comprises the polyphenols which are water soluble *Corresponding author. Fax: 1353-91-525-700. E-mail address: [email protected] (M.J. Hynes).

and may constitute up to 40% the dry weight of green tea. They contain numerous potential metal binding sites and will bind preferentially to hard metals such as aluminium(III) and iron(III). Both epidemiological and laboratory experiments have indicated that polyphenols are useful in the fight against numerous diseases. Tea inhibits the activity of several enzymes related to tumour promotion and cell proliferation including ornithine decarboxylase, protein kinase c, cyclooxygenase and lipoxygenase. It can also inhibit the formation of lung cancer in mice arising from NNK, a common tobacco carcinogen. The antioxidant ability of tea has been partly attributed to its ability to scavenge free copper and iron, thus preventing oxidative damage. Previous studies have indicated that aluminium(III) forms complexes of high stability with ligands containing oxygen donor ligands. Arising from this, polyphenols have the potential to solubilise aluminium and perhaps render it more biologically available. In order to determine whether or not polyphenols act as a means of aluminium uptake into the blood stream and whether or not they nullify the beneficial effects of the ligands and to shed light on the possible mechanisms of these processes we

0162-0134 / 01 / $ – see front matter  2001 Elsevier Science B.V. All rights reserved. PII: S0162-0134( 00 )00232-4

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have initiated kinetic and mechanistic studies into the reactions of phenols with aluminium(III). A recently published book on polyphenols [3] contains only two pages dealing with metal-ion complexation. It does refer to the iron-binding properties of enterobactin, a polyphenolbased siderophore produced by Escherichia coli [4]. Our initial studies have concentrated on simple phenols containing only one aromatic ring. In this paper we report the results of our studies of the kinetics and mechanisms of the reactions of aluminium(III) with gallic acid, gallic acid methyl ester and adrenaline. The first two of these are of relevance in themselves as they are the final hydrolysis products of condensed polyphenols found in tea and other foods. The aqueous chemistry of aluminium is complicated by the presence of hydrolytic species. In the case of ligands which exist as protonated and deprotonated species under the reaction conditions, this leads to proton ambiguity when interpreting kinetic data. With an ionic radius of 0.53 ˚ aluminium (III) is a highly charged small cation, which A is easily hydrolysed in aqueous solution in the absence of competing ligands. The pK of [Al(H 2 O) 6 ] 31 to give the monohydroxy species is 5.55 [5], this is considerably higher than the corresponding value of 2.7 for [Fe(H 2 O) 6 ] 31 [6]. Aluminium is sluggishly labile and its intermolecular reactions in H 2 O are relatively slow. The water exchange rate on [Al(H 2 O) 6 ] 31 has been found to be 1.3 s 21 with a volume of activation of 15.7 cm 3 M 21 [7]. This has been interpreted in terms of an Id mechanism, a hypothesis supported by recent molecular dynamic calculations in which it proved impossible to obtain transition states corresponding to either A or Ia mechanisms [8]. Recent work by Nordin et al. [9] has reported a value of 3.1310 4 s 21 for water exchange on [Al(H 2 O) 5 OH] 21 . These authors argue on the basis of the reduced charge and the coordination saturation of the relatively small Al 31 ion that water exchange on [Al(H 2 O) 5 OH] 21 should be more dissociative than on [Al(H 2 O) 6 ] 31 . Apart from stability constant measurements, there have been very few studies of the reactions of [Al(H 2 O) 6 ] 31 with the catecholates and related ligands.

2. Experimental Gallic acid (H 3 ga), 3,4,5-trihydroxybenzoic acid (Aldrich), gallic acid methyl ester (H 2 game), methyl 3,4,5trihydroxybenzoate (Aldrich), adrenaline (H 2 ad) 3,4dihydroxy-a-(methylaminomethyl)benzyl alcohol, (BDH), maltol (Hma) 3-hydroxy-2-methylpyrone (Aldrich) and H 5 dtpa, diethylenetriaminepentacetic acid (Aldrich) were used as supplied. Solutions of Na 2 H 2 edta (BDH) were standardised by titration with standard NaOH. Aluminium ¨ solutions were standardised trichloride (Riedel-de Haen) by addition of excess of standardised EDTA (BDH),

followed by back titration with a standard solution of ¨ Zn 21 (aq) prepared from ZnSO 4 .7H 2 O (Riedel-de Haen), using Solochrome Black indicator (BDH). It has been shown that aluminium(III) does not form any complexes with Cl 2 even at 5 M NaCl [10]. Solutions of the required final pH were made up from deoxygenated stock solutions ¨ of the ligands and using perchloric acid (Riedel-de Haen) as the source of hydrogen ions. All solutions were adjusted to an ionic strength of 0.5 M with sodium perchlorate (Aldrich). In order to avoid the formation of oligomeric species, the aluminium(III) solutions were used immediately after adjustment of the pH. For complex formation reactions at pH ca. 4, both the ligand and metal solutions were buffered either by addition of 0.01 M piperazine-NN9-(bis(4-butanesulphonic acid), pipbs (GFS) or piperazine-N-N9-(bis(3-propanesulphonic acid), pipps (GFS) [11]. At higher pH values pipes (Aldrich) was used. This resulted in a negligible change in pH on formation of the complex. pH measurements were made with an AGB 3000 pH meter equipped with an Amagruss combination electrode. The electrode was calibrated to read hydrogen ion concentration by titrating solutions of perchloric acid (0.001– 0.002 M) with standard sodium hydroxide solutions. Titrations were carried out in jacketed titration vessels through which water at 258C was circulating. Endpoints were determined using the method of Johannson [12]. Most of the kinetic measurements were made by use of a Hi-Tech SF-20 stopped flow apparatus. For slower reactions UV–Vis spectra were obtained as a function of time by use of a Hewlett Packard 8453A diode array spectrophotometer equipped with a Hi-Tech SFA-20 rapid kinetics accessory. The complex formation reactions were studied with the aluminium in pseudo first order excess. For reactions carried out at high acid concentrations, the ligand concentration was ca. 1.0310 23 M while at the higher pH values, it was 0.2310 23 M. The dissociation of the complex was studied by reacting a solution containing metal and ligand at a pH where there was appreciable complex formation with a series of perchloric acid solutions in the stopped-flow apparatus. The reported rate constants are the average of at least three determinations. The pseudo first order rate constants were calculated from the experimental absorbance data using the Olis KINFIT routines [13]. Single value decomposition and global kinetic analysis of spectral-kinetic data obtained using the Hewlett Packard 8453A diode array spectrophotometer was carried out by use of the program SPECFIT [14] which is based on the published works of ¨ Zuberbuhler et al. [15]. Kinetic data were fitted to the proposed mechanisms using a non-linear least squares program [16]. These were then confirmed using SPECFIT. Semi-empirical molecular orbital (MO) calculations were performed by means of PC SPARTAN Pro [17]. Full geometry optimisation was performed with the PM3 MO calculations. Atom charges, the energies of the molecular

M. O´ ’ Coinceanainn, M. J. Hynes / Journal of Inorganic Biochemistry 84 (2001) 1 – 12

orbitals and the contributions of the various orbitals to them were evaluated using the SPARTAN standard procedures.

3. Results and discussion While a relatively large number of species may be formed on hydrolysis of aqueous solutions of alumin¨ ium(III), Ohman et al. suggested that up to pH 4.2, only the following species need be considered [5]. Baes and Mesmer [18] conclude that: Al 31 1 H 2 OáAl(OH)21 1 H 1 , 1 3Al 31 1 4H 2 OáAl 3 (OH) 51 4 1 4H ,

b 211 5 10 25.52

(1)

b 243 5 10 213.57 (2)

1 13Al 31 1 32H 2 OáAl 13 O 4 (OH) 71 24 1 32H ,

b 23213 5 10 2109.2

3

Fig. 1 shows the various ligand species present in the gallic acid–aluminium(III) system [19]. When the metal is present in at least a five-fold excess, three complexes have to be considered, [Al(HL)] 1 , [Al(L)] and 41 [Al 3 (OH) 4 (H 2 L)] . Fig. 1 shows that up to a pH of 3.5 only the [Al(HL)] 1 species is present while at higher pH values, some of the [Al(L)] and [Al 3 (OH) 4 (H 2 L)] 41 species may be present. While equilibrium measurements do indicate that the [Al 3 (OH) 4 (H 2 L)] 41 is formed, it is highly unlikely that it will be formed as an initial product when aluminium(III) is reacted with gallic acid. Moreover, this species has been formulated as Al 3 OH 4 (H 2 L)41 [19] which suggests that it is derived by reaction of Al 3 (OH) 451 with a ligand species. Under the experimental conditions of the present work, Al 3 (OH) 51 is present in extremely 4 low concentration and furthermore would be expected to be relatively inert compared to other aluminium(III) species present such as Al(OH)21 . A relatively recent study carried out over the pH range 0.97–3.65 which used

(3)

(a) mononuclear hydrolytic species are formed rapidly and reversibly, (b) small polynuclear species such as Al 2 (OH) 41 and Al 3 (OH) 51 are formed less rapidly, (c) 2 4 71 species such as Al 13 O 4 (OH) 24 are formed more slowly still and (d) in general dimeric and other oligomeric species are only of minor importance at 258C and are based largely on measurements at higher temperatures. In this work, the experimental conditions were such that only the Al(OH)21 species had to be considered in the interpretation of the kinetic data. In the present investigation, the slowest k obs measured (0.045 s 21 ) represents a half-life of 15 s such that the measurement time was always less than 1 min.

Fig. 1. Species distribution of aluminium(III) with gallic acid [19].

Fig. 2. Time dependent spectra at 258C on reaction of aluminium(III) with (a) gallic acid methyl ester, pH 4, DT50.2 s and (b) gallic acid, pH 4.3, DT50.2 s. [Aluminium] total 54.0310 23 M and [ligand] total 50.23 10 23 M.

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M. O´ ’ Coinceanainn, M. J. Hynes / Journal of Inorganic Biochemistry 84 (2001) 1 – 12

Fig. 3. The first three eigenvectors of the factor analysis (both time-dependent and wavelength-dependent) for reaction of aluminium(III) with (a) the gallic acid and (b) the gallic acid–methyl ester systems at pH 4.

total aluminium(III) concentrations significantly higher than those used in the present work did not indicate any contributions from oligomeric species such as Al 3 (OH) 51 4

to the solvent exchange process [9]. Considerations such as the above are not relevant to aluminium(III)–gallic acid methyl ester complexes due to the lack of a displacable

M. O´ ’ Coinceanainn, M. J. Hynes / Journal of Inorganic Biochemistry 84 (2001) 1 – 12

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Fig. 3. (continued)

carboxylic acid proton on the ligand. The pKa values of the phenol groups on gallic acid have been reported to be 8.44 and 11.08 [20] so that under the experimental conditions of

the present work, the phenolic groups are fully protonated in the free ligand. The same is true for gallic acid methyl ester and adrenaline.

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When solutions containing aluminium(III) and solutions containing either gallic acid or gallic acid methyl ester were reacted in the stopped-flow apparatus, absorbance increases resulting from formation of the complex were observed in the vicinity of 240 and 330 nm while a decrease in the absorbance of the ligand peak was observed in the vicinity of 270 nm. Analysis of the kinetic data at the three wavelengths gave identical pseudo-firstorder rate constants. Fig. 2 shows time-dependent spectra obtained for reaction of aluminium(III) with both gallic acid and gallic acid methyl ester. For both systems, the isosbestic points are sharp and suggest a smooth conversion of reactants to product. Furthermore, for both ligands there is no shift in the isosbestic points of time dependent spectra on going from pH 3.5 to 4. This suggests that only a single absorbing complex species is present in both systems. This hypothesis is supported by single value decomposition analysis of the diode array spectra of the type shown in Fig. 2 using SPECFIT [14]. Fig. 3 shows the first three time dependent and wavelength dependent eigenvectors of the factor analysis for both gallic acid and gallic acid methyl ester. The results are unambiguous and show that only two vectors are contributing to the absorbance changes. The third eigenvector shows no time or wavelength dependence. In the case of gallic acid methyl ester, the only complex formed is the [Al(L)] 1 species where H 2 L represents gallic acid methyl ester. Using the ¨ potentiometric results reported by Ohman et al. [19,20] it can be shown that at equilibrium, with a gallic acid 24 concentration of 2.0310 M, a five-fold excess of aluminium(III), and a pH of 3.5, 91% of complexed aluminium is present as [Al(HL)] 1 with 9% present as [Al(L)], a species in which the carboxylic acid group is dissociated, while at pH 4, the percentages are 71 and 23 respectively with approximately 5% of the Al 3 (OH) 4 (H 2 L)41 . However, analysis of the spectral changes observed in the aluminium(III)–gallic acid system, either by visual inspection or by global analysis using SPECFIT [14] clearly demonstrates that the spectral characteristics of [Al(HL)] 1 and [Al(L)] are indistinguishable. This is hardly surprising in light of the work of Kipton et al. [21]. They monitored the spectra of gallic acid as a function of pH and found that dissociation of the carboxylic acid had little or no effect on the spectral parameters. Point charge calculations carried out using SPARTAN suggest relatively small changes in the charges on either of the phenolic oxygens or the carboxylic acid oxygens on removal of the carboxylic acid proton. For the complex dissociation studies in which solutions containing appreciable concentration of the complex were reacted with an excess of acid, an absorbance decrease was observed. Fig. 4(a), (b) and (c) show plots of the kinetic data obtained during both the complex formation and dissociation reactions of aluminium(III) with the methyl ester of gallic acid. The kinetic data for the reactions with gallic acid and gallic acid methyl ester are shown in Tables

1 and 2 respectively. The linearity of the plots in Fig. 4 is consistent with a reaction that is first-order in both ligand and metal ion. It is clear that k obs shows little dependence on the concentration of aluminium(III) at the lower pH values. The kinetic data for dissociation of the 1:1 complex of aluminium(III) with methyl gallate are shown in Fig. 4(c). In this instance, the plots of k obs against [H 1 ] curve upwards, thus indicating that the order of the reaction with respect to hydrogen ion is greater than one. The kinetic data are rationalised by the following equations. k1

Al 31 1 H 3 L á Al(HL)1 1 2H 1

(4)

k 21

kh

Al 31 á Al(OH)21 1 H 1

(5)

k2

Al(OH)21 1 H 3 L á Al(HL) 1 H 1

(6)

k 22

According to this scheme and assuming that Eq. (5) represents a rapid equilibrium, k obs is described by Eq. (7) where [M o ] is the total concentration of aluminium(III) and Kh is the hydrolysis constant of Al 31 to give Al(OH)21 , 2.82310 26 [5].

H

J

k1 k2 k obs 5 ]]]] 1 ]]]] [M o ] 1 1 Kh / [H 1 ] 1 1 [H 1 ] /Kh 1 k 21 [H 1 ] 2 1 k 22 [H 1 ]

(7)

Under the present experimental conditions, Kh / [H 1 ]<1 and [H 1 ] /Kh 41 so that Eq. (7) may be written as Eq. (8).

H

J

k 2 Kh 1 2 1 k obs 5 k 1 1 ]] [M o ] 1 k 21 [H ] 1 k 22 [H ] [H 1 ]

(8)

The complex dissociation data were used to obtain values for k 21 and k 22 and these were then used to fit the formation data Eq. (8). This yielded values 1.1(61.0) M 21 s 21 for k 1 and 781(622) M 22 s 21 for k 21 and 1145(6213) and 139(62) M 21 s 21 for k 2 and k 22 respectively for reaction of gallic acid with aluminium(III). The kinetic data in Table 2 yielded values 3.0(60.6) M 21 s 21 for k 1 and 810(640) M 22 s 21 for k 21 and 1328(688) and 135(66) M 21 s 21 for k 2 and k 22 respectively for reaction of gallic acid methyl ester with aluminium(III). Analysis of kinetic data showed that the k 1 pathway made only a small contribution to complex formation. At the highest pH used in the present investigation, some of the carboxylic acid on the gallic acid would be dissociated (pKa 54.3). If solvent exchange were the rate determining step in the complex formation reaction, then it might be expected that this species would have a higher

M. O´ ’ Coinceanainn, M. J. Hynes / Journal of Inorganic Biochemistry 84 (2001) 1 – 12

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Table 1 Kinetic data for reaction of aluminium(III) with gallic acid in aqueous solution at 258C and I50.5 M NaClO 4 and l 5330 nm 10 3 [H 1 ] / M 25.0 50.0 75.0 100 125 150 175 225 10.0 a 10.0 a 10.0 a 10.0 a 10.0 a 1.00 a 1.00 a 1.00 a 1.00 a 1.00 a 0.316 b 0.316 b 0.316 b 0.316 b 0.316 b 0.100 b 0.100 b 0.100 b 0.100 b 0.100 b

10 3 [Al(III)] / M c

0.00 0.00 c 0.00 c 0.00 c 0.00 c 0.00 c 0.00 c 0.00 c 5.00 10.0 15.0 20.0 25.0 5.00 10.0 15.0 20.0 25.0 1.00 2.00 3.00 4.00 5.00 1.00 2.00 3.00 4.00 5.00

k obs / s 21

k calc / s 21

3.40 8.31 14.0 22.0 31.0 39.8 49.2 71.0 1.36 1.41 1.45 1.50 1.55 0.136 0.155 0.204 0.233 0.293 0.051 0.063 0.071 0.087 0.090 0.039 0.075 0.095 0.145 0.180

4.05 9.09 15.0 22.1 30.0 39.0 48.9 71.7 1.41 1.43 1.45 1.47 1.49 0.170 0.201 0.232 0.263 0.293 0.056 0.067 0.079 0.091 0.102 0.043 0.071 0.097 0.125 0.153

[Gallic acid]51.0310 23 M. [Gallic acid]50.2310 23 M. c Dissociation of the complex. a

b

Fig. 4. (a) and (b). Plots of k obs against [aluminium(III)] for reaction of aluminium(III) with gallic acid methyl ester. (c) Plot of k obs against [H 1 ] for acid catalysed dissociation of the 1:1 complex of aluminium(III) with gallic acid methyl ester. All data were obtained at 258C and I50.5 M NaClO 4 at 330 nm.

reactivity than the parent gallic acid corresponding to its higher outer-sphere association constant with Al(OH)21 , 1.18 compared to 0.3 for the fully protonated ligand. However, solvent exchange is not the rate determining step in these reactions, vide infra, and as a result it appears that there is little difference in the reactivity of either species towards aluminium(III). This conclusion is supported by forward evolving factor analysis of spectra-kinetic data using SPECFIT32. This did not indicate the appearance of two species over different time scales as would be expected if one form of the ligand was more reactive. Furthermore, attempts to fit the kinetic data to a mechanism involving a separate pathway for the deprotonated species were unsuccessful. Reaction of aqueous solutions of aluminium(III) with adrenaline results in a single exponential. This was confirmed by global analysis of the kinetic data. The kinetic data are shown in Table 3. Analysis of stability constants data for adrenaline [22] showed that the stability constants of aluminium(III) complexes with this ligand are some twenty-fold less than with gallic acid. As a result, there is little complex formation below pH 4. Consequently the rate of complex formation was ascertained from the

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Table 2 Kinetic data for reaction of aluminium(III) with gallic acid methyl ester in aqueous solution at 258C and I50.5 M NaClO 4 and l 5330 nm 10 3 [H 1 ] / M 25.0 50.0 75.0 100 125 150 175 10.0 a 10.0 a 10.0 a 10.0 a 10.0 a 1.00 a 1.00 a 1.00 a 1.00 a 1.00 a 0.316 b 0.316 b 0.316 b 0.316 b 0.316 b 0.0794 b 0.0794 b 0.0794 b 0.0794 b 0.0794 b

10 3 [Al(III)] / M c

0.00 0.00 c 0.00 c 0.00 c 0.00 c 0.00 c 0.00 c 5.00 10.0 15.0 20.0 25.0 5.00 10.0 15.0 20.0 25.0 1.00 2.00 3.00 4.00 5.00 1.00 2.00 3.00 4.00 5.00

k obs / s 21

Kcalc / s 21

3.31 8.15 15.3 20.5 31.0 38.9 47.8 1.38 1.45 1.43 1.48 1.52 0.187 0.208 0.245 0.260 0.280 0.060 0.081 0.092 0.115 0.135 0.045 0.090 0.128 0.196 0.265

3.88 8.78 14.7 21.6 29.6 38.5 48.5 1.42 1.43 1.45 1.47 1.49 0.174 0.208 0.242 0.275 0.310 0.059 0.074 0.088 0.103 0.118 0.058 0.105 0.151 0.195 0.240

[Gallic acid]51.0310 23 M. [Gallic acid]50.2310 23 M. c Dissociation of the complex. a

b

complex dissociation data and the published literature values of the equilibrium constant values. This gave an estimated rate constant of 316 M 21 s 21 for k 2 with k 21 and k 22 equal to 2.06(60.22)310 3 M 22 s 21 and 388(627) M 21 s 21 respectively. Table 4 summarises the kinetic data for all three ligands. The validity of the proposed mechanism may be tested by using the rate constants derived to calculate stability

Table 3 Kinetic data for the reaction of aluminium(III) with adrenaline in aqueous solution at 258C and I50.5 M and l 5310 nm 10 3 [H 1 ] / M

10 3 [Al(III)] / M

k obs / s 21

k calc / s 21

25.0 50.0 75.0 100 125 150 0.0794 0.0794 0.0794 0.0794 0.0794

0.00 0.00 0.00 0.00 0.00 0.00 1.00 2.00 3.00 4.00 5.00

8.50 22.2 40.1 63.1 81.7 102 0.068 0.099 0.152 0.210 0.254

11.0 24.5 40.7 59.4 80.6 104 0.052 0.103 0.154 0.206 0.257

constants and then comparing them with independently determined values. Using the potentiometric data of ¨ ¨ Ohman and Sjoberg for reaction of aluminium(III) with gallic acid [20] a value of 1.3310 25 may be calculated for Q where Q5k 1 /k 21 5k 2 Kh /k 22 . The kinetic data for k 2 and k 22 yield a value 2.3310 25 for Q in fair agreement with the directly determined. In view of the small contribution made by the k 1 pathway to complex formation and the resultant uncertainty, it is not considered sufficiently reliable for calculation of Q. The limited equilibrium data available for reaction of aluminium(III) with gallic acid methyl ester [23] suggest that the stability of the 1:1 complex is similar to that of the corresponding gallic acid complex. This is confirmed by the kinetic measurements which give a value of 2.65310 25 for Q. Tables 1–3 give the calculated rate constants obtained using Eq. (8) and the fitted values of the rate constants. For both H 3 ga and H 2 game the agreement is excellent over a range of almost four orders of magnitude of hydrogen ion concentration and is good evidence in support of the proposed mechanism. In view of the fact that solvent exchange on both [Al(H 2 O) 6 ] 31 and [Al(H 2 O) 5 OH] 21 proceeds by an Id mechanism, the complex formation reactions may be interpreted in terms of the Eigen–Wilkins–Tamm dissociative mechanism [24] which involves the formation of an outer-sphere complex in rapid equilibrium with the reactants. Using the values of the water exchange rate constants together with the outer-sphere association constant of 0.3 calculated using the modified form of the Eigen–Fuoss equation [25], values of approximately 0.3 and 6750 M 21 s 21 may be calculated for reaction of [Al(H 2 O) 6 ] 31 and [Al(H 2 O) 5 OH] 21 respectively with uncharged ligands using Eq. (9) [26] where k f is the rate constant for complex formation, k s is the rate of solvent exchange and Ko is the outer sphere association constant. Kf 5 0.75 k s Ko

(9)

It is clear that the rate constants obtained for reaction of [Al(H 2 O) 5 OH] 21 with both gallic acid and its methyl ester are retarded by a factor of between 5 and 8. Secco and Venturini [27] investigated the reactions of aluminium(III) with salicylate ions. At the time of their investigations, no water exchange data were available for [Al(H 2 O) 6 ] 31 or [Al(H 2 O) 5 OH] 21 . Using the data now available the rate constants for reactions of [Al(H 2 O) 6 ] 31 and [Al(H 2 O) 5 OH] 21 with salicylate ion may be calculated using Eq. (9). These together with a selection of rate constants for other ligands are presented in Table 4. The results are quite interesting and demonstrate a remarkable inconsistency in the complex formation behaviour of [Al(H 2 O) 5 OH] 21 in particular. Previous work has shown that complex formation reactions involving fully protonated ligands tend to be retarded compared to the corresponding reactions with the deprotonated ligands [28–30]. Sutin el al. [29] have argued

M. O´ ’ Coinceanainn, M. J. Hynes / Journal of Inorganic Biochemistry 84 (2001) 1 – 12

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Table 4 Comparison of calculated and experiment rate constants for complex formation reactions of [Al(H 2 O) 6 ] 31 and [Al(H 2 O) 5 OH] 21 in aqueous solution a Reaction

I/M

Ko

k f calc / M 21 s 21

k f expt / M 21 s 21

k(calc) / k(expt)

Ref.

Al 31 –H 3 ga Al 31 –H 2 game Al(OH)21 –H 3 ga AlOH 21 –H 2 game Al 31 –H 3 ad 1 AlOH 21 –H 3 ad 1 Al 31 –SC 21 Al(OH)21 –SC 12 Al(OH)21 –SSC 22 Al(OH)21 1NSC 12 Al 31 1SXO 22 Al 31 1SMTB 22 Al(OH)21 1SXO 22 Al(OH)21 1SMTB 22 Al 31 1H 4 dfb 1 Al(OH)21 1H 4 dfb 1 Al 31 1H 4 DTPA2 Al 31 1H 3 DTPA22 Al(OH)21 1H 3 DTPA22

0.5 0.5 0.5 0.5 0.5 0.5 0.1 0.1 0.1 0.1 0.6 0.6 0.6 0.6 2.0 2.0 0.5 0.5 0.5

0.3 0.3 0.3 0.3 0.043 0.043 5.3 2.1 13.6 2.1 13.7 13.7 3.9 3.9 0.087 0.133 2.29 16.7 4.44

0.30 0.30 6975 6975 0.043 1000 5.2 4.9310 4 3.2310 5 4.9310 4 13.4 13.4 9.1310 4 9.1310 4 0.085 3092 2.23 16.3 1.0310 5

1.1 3.0 1145 1330 4.7310 23 316 0.91 1.95310 3 1.22310 3 0.77310 3 24.3 6.83 1.51310 2 63.3 0.0211 189 2.1 19.3 1.34310 3

0.08 0.10 6.1 5.2 0.11 3.2 6 25 262 63 0.55 2 600 1438 4 16.4 1.1 1.02 77

This work This work This work This work This work This work [27] [27] [45] [46] [47] [47] [47] [47] [48] [48] [49] [49] [49]

a

GA, gallic acid; GAME, gallic acid methyl ester; SC, salicylate; SSC, 5-sulfosalicylate; NSC, 5-nitrosalicylate; SXO, semi-xylenol orange; SMTB, semi-methylthymol blue; H4dfb, fully protonated desferrioxamine.

that at least three factors must be considered when considering the complex formation reactions of multidentate protonated ligands, (a) relatively strong intramolecular hydrogen bonds which may convert the protonated ligand into a poor entering group, (b) the energetics of proton release which may slow down the rate of ring closure and (c) ring strain associated with the formation of six-membered chelates. The latter is not a consideration in the present work as the chelates formed contain five-membered rings. It is usually extremely difficult to distinguish between factors (a) and (b). The reaction of [Al(H 2 O) 6 ] 31 with H 2 ad is faster than its reaction with either H 3 ga or H 2 game while the reaction of [Al(H 2 O) 5 (OH)] 21 with H 2 ad is somewhat slower than its reactions with the other two ligands. The rate constants for the reverse reactions, k 21 and k 22 are larger than for the other two ligands. It is clear that for all of the reactions listed in Table 4, water release is not the rate determining factor. However, the degree of retardation is not uniform and indeed some of the results are rather puzzling. While the degree of retardation in the present work is consistent with that observed for other ligands, the greatly increased retardation on going from salicylate to 5-sulfosalicylate is rather surprising.

3.1. Reactions at physiological pH The preparation of aluminium solutions from aluminium salts at neutral pHs is difficult due to problems of precipitation and as a result the concentration of

[Al(H 2 O) 6 ] 21 will be extremely low. 3-Hydroxy-2methyl-4-pyrone (maltol, Hma) has been used to form a soluble aluminium(III) complex that does not undergo hydrolysis [31,32]. Distribution calculations using published values of the stability constants [33] show that at a ligand to metal ratio of four, virtually all of the aluminium is present as [Al(ma) 3 ] at a pH of 6.7. When solutions containing 4310 24 M aluminium(III) and 1.6310 23 M maltol were reacted with a pseudo first order excess of acid, three separate reactions were observed. The first reaction was extremely rapid with a rate constant of ca. 500 s 21 . For the second reaction, k obs had the form of Eq. (10) while the third step was independent k obs 5 5 1 160[H 1 ]

(10)

of hydrogen ion concentration and had a rate constant of 5310 23 s 21 . Variable pH 27 Al experiments of solutions containing maltol and aluminium(III) show a single resonance at pH 7 while at lower pH values, three resonances were observed which were assigned to [Al(H 2 O) 6 ] 31 [Al(ma)(H 2 O) 4 ] 21 and [Al(ma) 2 (H 2 O) 2 ] 1 at 0, 13 and 26 ppm respectively [34]. Using these values and the frequency at which the spectra were recorded (78.1 MHz) it is readily shown that the half-lives of the various species are greater than 0.16 ms. This is not inconsistent with the complex dissociation data in view of the fact that a rate constant of 500 s 21 corresponds to a half-life of 1.4 ms. When solutions containing 4310 24 M aluminium(III) and 1.6310 23 M maltol were reacted with a pseudo first order excess of gallic acid methyl ester, a single reaction was observed with a rate constant of 5.04310 22 s 21

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which was independent of the concentration of gallic acid methyl ester. Diethylenetriaminepentaacetic acid (DTPA) has been used as an agent for chelating ingested plutonium [35] and has been investigated as a replacement for desferrioxamine, the most widely used chelating agent for treatment of ingested aluminium in order to reduce the body burden [36]. As a result, its reactions with aluminium complexes at physiological pH are of interest. Global analysis of the spectral data obtained on reaction of aluminium maltolate with a pseudo first order excess of DTPA at pH 6.7 showed that the overall reaction consisted of two steps, the first of which had a half-life of ca. 12 s while the second which was associated with a relatively small absorbance change had a half life of ca. 350 s. Only the faster of the two steps was investigated in detail as it is unlikely that the slow step was associated with ligand substitution. The faster step had a rate constant of 4.66(60.15)310 22 s 21 and was independent of the DTPA concentration. This rate constant is in good agreement with that found for reaction of aluminium(III) maltolate with gallic acid methyl ester. It is clear that the reactions of [Al(ma) 3 ] with both gallic acid methyl ester and DTPA proceed by a limiting dissociative mechanism with the three maltol ligands being removed in stepwise fashion. The initial step involves cleavage of a metal–oxygen bond, with the now vacant co-ordination site being replaced by a donor atom on the incoming ligand. The high rate constant observed for the acid catalysed dissociation of [Al(ma) 3 ] is at first rather surprising. In view of the fact that in [Al(ma) 3 ], the Al 31 cation is co-ordinatively saturated, it might be expected that the rate constant for acid catalysed dissociation would be similar to that obtained for replacement of maltol by gallic acid methyl ester and by DTPA. The magnitude of the observed rate constant may be rationalised by a mechanism which involves proton assisted dissociation. In this mechanism, one of the co-ordinated oxygen atoms becomes protonated in a rapid reaction following which the metal–oxygen atom is cleaved. There are two possible sites for the proton attack. However, from an examination of previous examples of reactions where this mechanisms has been shown to be operative, it is probable that the attack occurs at the .C5O oxygen. In an investigation of the formation and dissociation of lanthanide(III) complexes of 1,4,7,10-tetraazacyclododecane-N,N9,N0,N--tetraacetic acid (H 4 DOTA), the linear dependence of k obs on [H 1 ] was explained in terms of a proton-catalysed dissociation of [Gd(DOTA)] 2 in which a proton attached to a carboxylate oxygen prior to dissociation [37]. Similar observations have been made by Sherry et al. [38] arising from their kinetic investigations of the dissociation of gadolinium(III) complexes of triazacyclodecanetriacetate ligands and their investigations of formation and dissociation kinetics of the magnesium(II)complex of 1,4,7-triazacyclononane-1,4,7-tris(methylenemethylphosphinic acid) [39].

Fig. 5. Charge distribution on oxygen atoms of gallic acid.

3.2. Molecular orbital calculations Calculations have been made for the MO structure of both gallic acid and gallic acid methyl ester. The geometry was optimised and the electron densities were calculated in order to predict both the spatial arrangement of the hydroxyl groups and the initial point of attachment to the metal ion. Fig. 5 shows the optimised structure produced by PC SPARTAN Pro together with the net charges calculated for oxygens 1, 2 and 3. The arrangement of the phenolic groups is identical to that suggested by Pauling [40,41] on the basis of infra-red measurements in pyrogallol. Two bands were observed, one at 7050 cm 21 and another with doubled area at 6960 cm 21 . The arrangement is facilitated by the fact that the C–O bond in phenols has some double-bond character which causes the hydrogen atoms to lie in the plane of the benzene ring. Arising from this, it is probable that the initial attack by aluminium occurs at O(1).

4. Conclusions The aqueous hydrolytic behaviour of aluminium(III) is quite complicated with the possibility of producing a range of both mononuclear and polynuclear species such as Al(OH)21 , Al 3 (OH) 51 and Al 13 O 4 (OH) 71 4 24 . However species distribution plots demonstrate that by working at pHs less than about 4 and relatively low concentration of metal ion, the influence of the polymeric species may be minimised. In addition, under the conditions of the present work, the Al 13 O 4 (OH) 71 24 species is unlikely to be formed [42] and furthermore the half lives of its acid catalysed breakdown have been reported to be between 350 and 4300 s. The corresponding rate constants are much less than those reported in Tables 1–3. The current work shows that [Al(H 2 O) 5 (OH)] 21 is the major species involved in complex formation. At first sight, this is rather surprising in view of the relatively small fraction of the metal present at this form, even at pH 4. However based on the water exchange rates (k s 51.3 s 21

M. O´ ’ Coinceanainn, M. J. Hynes / Journal of Inorganic Biochemistry 84 (2001) 1 – 12

11

Fig. 6. Single wavelength comparisons of the data (+), fit (–), and residuals from the global fit to the kinetic model given above for the formation of the aluminium(III) gallic acid complex in aqueous solution at T5258C.

for [Al(H 2 O) 6 ] 31 ; k s 53.1310 4 s 21 for 21 [Al(H 2 O) 5 (OH)] , this species is some 2.38 X 10 4 more 31 reactive than Al . Hence, its dominant role in complex formation at the pHs used in the current work. Such behaviour is not unprecedented as a relatively similar situation obtains in the reactions of iron(III) [43] and chromium(III) [44] where the main pathways for complex formation involve reaction of the M(OH)21 species. The kinetic data were obtained in the present investigation cover over almost four orders of magnitude of hydrogen ion concentration. Both the complex formation and dissociation data fit the reaction scheme proposed. It is most unlikely that a reaction scheme in which the hydrogen ion stoichiometry was incorrect would give such a good fit over such as large hydrogen ion concentration. Fig. 6 shows the single wavelength comparisons of the data (+), fit (–), and the residuals trace from the global fit to the

kinetic model given above, for the spectral data in Fig. 2(b) for the formation of the aluminium(III) gallic acid complex in aqueous solution at 258C. This returned a globally optimised value of 1115 (65) M 21 s 21 for reaction of Al(OH)21 compared to 1145 M 21 s 21 for single wavelength data. The choice of the complexed species selected for gallic acid is supported by the fact that the data for gallic acid methyl ester yield very similar rate constants as would be expected.

Acknowledgements The work is supported by Enterprise Ireland Basic Research Grant SC / 1998 / 487 /. One of us (MOC) thanks Enterprise Ireland for a Basic Research Award.

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