Tripodal amine catechol ligands: A fascinating class of chelators for aluminium(III)

Journal of Inorganic Biochemistry 102 (2008) 1581–1588

Contents lists available at ScienceDirect

Journal of Inorganic Biochemistry journal homepage: www.elsevier.com/locate/jinorgbio

Tripodal amine catechol ligands: A fascinating class of chelators for aluminium(III) Minati Baral a,b,*, Suban K. Sahoo a, B.K. Kanungo a a b

Department of Chemistry, Sant Longowal Institute of Engineering and Technology (Deemed University), Longowal, 148 106 Punjab, India Department of Chemistry, National Institute of Technology, Kurukshetra, 136 119 Haryana, India

a r t i c l e

i n f o

Article history: Received 24 September 2007 Received in revised form 24 January 2008 Accepted 13 February 2008 Available online 4 March 2008

Keywords: Tripodal polycatechol-amine ligands Aluminium Stability constants Molecular mechanics and semi-empirical method

a b s t r a c t Complexation constants based on potentiometric titrations and spectrophotometric measurements in an aqueous medium of 0.1 M KCl at 25 ± 1 °C for the complexes of Al(III) with multidentate tripodal polycatechol-amine ligands, cis,cis-1,3,5-tris[(2,3-dihydroxybenzylamino)aminomethyl]cyclohexane (TMACHCAT, L1) and N1,N3,N5-tris(2-(2,3-dihydroxybenzylamino)ethyl)cyclohexane-1,3,5-tricarboxamide (CYCOENCAT, L2) have been summarized in this paper. Both the ligands released six protons to form various monomeric complexes of the types AlLH3, AlLH2, AlLH and AlL (L = L1 and L2). The first species AlLH3 depicted at low pH for which a monocapped type geometry was suggested, where the ligands were coordinated through three catecholic oxygens at ortho. Other species are formed subsequently from the species AlLH3 in steps upon deprotonation and coordination of the catecholic oxygens at meta to give encapsulated tris(catechol) type complexes. The probable structures of the metal complexes formed in solution were proposed through molecular modeling calculations. The pAl values calculated for AlL1 and AlL2 are appreciably higher than transferrin. The ligand L2 showed higher affinity towards Al(III) than L1 and desferrioxamine (DFO), the only approved drug for the treatment of aluminium intoxication. Ó 2008 Elsevier Inc. All rights reserved.

1. Introduction There is considerable clinical need for the design of new chelators that can be used to reduce metal ions overload diseases and the health effects associated with it. In order to design new chelators for a metal ion, efforts are made to mimic the molecular structure and binding sites of the natural chelators through simple synthetic molecules [1–4]. Such synthetic biomimetic molecules are implemented as metal ion sequestration due to their similar selective and strong binding properties towards the metal ion for which the biomolecules perform [5]. One of the most studied class of biomolecules for designing biomimetic synthetic chelators is siderophores, which are produced by microorganisms in order to extract iron from the external environment and to transport it into the organism [6,7]. Among all the siderophores, enterobactin (Fig. 1) is known to form very strong chelate with iron(III) giving the highest formation constant (log K = 49) [8]. Its efficiency as Fe(III) ion scavenger and carrier has stimulated the synthesis of many analogues containing three catechol units in tripod with respect to their use in iron overload treatment [9]. Previous studies

* Corresponding author. Address: Department of Chemistry, Sant Longowal Institute of Engineering and Technology (Deemed University), Longowal, 148106 Punjab, India. Tel.: +91 1672 284840 (Off.), +91 1672 284841 (Res.); fax: +91 1672 284840. E-mail addresses: [email protected] (M. Baral), [email protected] (B.K. Kanungo). 0162-0134/$ - see front matter Ó 2008 Elsevier Inc. All rights reserved. doi:10.1016/j.jinorgbio.2008.02.006

reveal that most of the analogs of enterobactin are designed by replacing the cyclic L-serine unit with a suitable tri-amine and then treating with catecholamides [1]. Very few attempts were made to design enterobactin analogs by replacing the amide functionals [10]. In light of the widespread biological importance of catecholamines as neurotransmitters [11,12], and pharmacological use for the treatment of Parkinson’s disease [13], hypertension [14], and breast cancer [15]; in our previous communications [16,17] we introduced two novel tripodal polycatechol-amine ligands, cis, cis-1,3,5-tris[(2,3-dihydroxybenzylamino)aminomethyl]cyclohexane (TMACHCAT, L1) and N1,N3,N5-tris(2-(2,3-dihydroxybenzylamino)ethyl)cyclohexane-1,3,5-tricarboxamide (CYCOENCAT, L2), where the cyclic ring and amide linkages of enterobactin have been replaced with cis,cis-1,3,5-tris(aminomethyl)cyclohexane (TMACH) and amine (-NH-) linkage, respectively. The ligands, TMACHCAT and CYCOENCAT provide tris(catechol) compartments for the encapsulation of metal ion such as Fe(III), Cr(III) and Ln(III) (Ln = La, Ga and Lu) [16,17]. Allied to the long established interest in both natural and synthetic chelating agents for Fe(III), currently there is also interest in the implementation of these molecules as chelators for Al(III) [18] because of their common preference towards the hard donor atoms, such as oxygen and similar coordination behaviors. Aluminium, a non-essential element is toxic to the central nervous system. The metal is involved in causing dialysis dementia in patients who are unable to eliminate aluminium because of renal dysfunction [19]. Other possible effects of aluminium on the central nervous

1582

M. Baral et al. / Journal of Inorganic Biochemistry 102 (2008) 1581–1588

O

O O O

O O

O HN

NH OH OH

HO HO

O

HN

O

HO HO

ENTEROBACTIN H H H

NH

HN HO

HO

HO

HO

HO

TMACHCAT (L1)

H O O

NH

NH

HN

HO

H H

O

HO HO

HN

HN HO HO

HN

HN HO HO

CYCOENCAT (L2 )

Fig. 1. Molecular structures of enterobactin, TMACHCAT (L1) and CYCOENCAT (L2).

system may be related to Alzheimer’s disease and amyotropic lateral sclerosis [20–22]. The only approved drug currently available for clinical treatment of aluminium intoxication is a natural siderophore desferrioxamine (DFO), an O6 donor [18], whose clinical use suffers from few important drawbacks like high cost, lack of oral efficacy, and major side effects in the long term [23]. Therefore, many attempts have been made to replace DFO by other simple synthetic molecules [24]. Owing to the strong affinity of Al(III) towards O-donor ligands [25] compared to N-donor ones [26], enterobactin analogues are one of the best choice for the development of Al(III) chelators. Keeping in view the above facts, in the present work, we have studied the potentiality of the ligands cis,cis-1,3,5-tris[(2,3dihydroxybenzylamino)aminomethyl]cyclohexane (TMACHCAT, L1) and N1,N3,N5-tris(2-(2,3- dihydroxybenzylamino)ethyl)cyclohexane-1,3,5-tricarboxamide (CYCOENCAT, L2), (Fig. 1) to implement as chelators for Al(III). The formation of Al(III) complexes with the ligands, L1 and L2 have been studied by potentiometric and spectrometric methods in an aqueous medium of 0.1 N KCl ionic strength and 25 ± 1 °C. The formation constants of the various complexes formed in solution, and their selectivity towards Al(III) have been presented. The complementary information on structure of the metal complexes formed in solution was obtained through molecular modeling calculations. 2. Experimental 2.1. Materials and measurements The compounds cis,cis-1,3,5-tris[(2,3-dihydroxybenzylamino)aminomethyl]cyclohexane (TMACHCAT, L1) and N1,N3,N5tris(2-(2,3-dihydroxybenzylamino)ethyl)cyclohexane-1,3,5-tricarboxamide (CYCOENCAT, L2) were synthesized as per our previous reported methods [16,17]. Nitrate salt of aluminium was obtained from Sigma–Aldrich and was used directly. All other chemicals: potassium hydroxide, hydrochloric acid and potassium chloride were obtained from Merck. The formation constants of the metal complexes were determined by potentiometric and spectrophotometric titrations at 25 ± 1 °C maintained from a double wall glass jacketed titration cell connected to a constant temperature circulatory bath. For all

titrations, the observed pH was measured as log [H+] using a ThermoOrion 720 A+ pH meter equipped with a combined glass electrode. The electrode was calibrated to read pH according to the classical method [27]. A standard hydrochloric acid solution was titrated with a standard KOH solution and the calculated hydrogen ion concentrations (pKw = 13.77 ± 0.05) was used to convert the pH-meter reading to hydrogen ion concentration. All solutions were prepared prior to the experiments in double distilled deoxygenated water. KOH solution of 0.1 M was prepared and standardized against potassium hydrogen phthalate. HCl solution (0.1 M) was prepared and standardized against standard KOH. The ionic strength was maintained at 0.1 M by adding appropriate amount of 1 M KCl. Solutions of 0.01 M ligands and 0.01 M metal ions were also prepared in deoxygenated water. Final concentrations of the ligands (1  103 M) and the metal ions (1  103 M and 5  104 M) were maintained for the different titrations. Following titrations with metal-to-ligand molar ratios: CM/CL = 0:1; CM/CL = 1:1, 1:2 were carried out. A non-linear least square computer program Hyperquad 2000 has been used to calculate the formation constants of the metal complexes [28]. The log K values of L1 (11.26, 10.65, 9.80, 8.48, 7.61 and 6.20) and L2 (11.36, 10.67, 9.82, 8.49, 7.62 and 6.27) determined previously [16,17] were used to evaluate the formation constants of the metal complexes. The first three values of log K were assigned to the hydroxyl groups of catechol units at ortho whereas last three values to the secondary amines. Protonation constants for the hydroxyl groups of catechol units at meta could not be evaluated within the adopted experimental conditions (pH 2.5–11.5) [16,17]. In the spectrophotometric studies, a dilute solution of ligand (4.02  105 M) and metal ion (4.02  105 M) was acidified with 0.1 N HCl at an ionic strength of 0.1 M KCl and 25 ± 1 °C, and then titrated with 0.1 N KOH. After each adjustment of pH, an aliquot was taken and spectra were recorded. The formation constants were calculated by global fitting of the whole spectral data using a non-linear least-square fitting program, pHAb [29]. 2.2. Molecular modeling calculations All calculations were carried out on a Pentium IV 3.0 GHz machine on Windows 2000 environment using the computer program CAChe (Computer Aided Chemistry) version 6.1.1 software from

1583

M. Baral et al. / Journal of Inorganic Biochemistry 102 (2008) 1581–1588

the Oxford Molecular Group [30]. The probable structures of the metal complexes formed in solution were drawn using CAChe workspace and then the geometry was optimized through molecular mechanics calculation using MM3 force field and adopting the eigen-vector following (EF) method. Molecular mechanics calculations using MM3 force field is known to predict the 3D-model structure of various metal complexes adequately [31,32]. Since, the ligands L1 and L2 have the probability to undergo ring flipping from their cis,cis-equatorial to axial conformer upon complexation, optimization process were undertaken for the same complex where the coordinating arms are present either at equatorial or axial position with respect to the cyclohexane ring. For all the possible metal chelates, six coordinated structures were drawn by adding appropriate number of water molecules with the metal ion. The semi-empirical calculations were carried out using MOPAC 2000 program implemented in CAChe. The MM3 minimized structures are re-optimized using semi-empirical PM3 self-consistent fields (SCF) method, at the Restricted Hartree–Fock (RHF) level with convergence limit of 0.0001 kcal/mol and RMS gradient of 0.001 kcal/mol. 3. Results and discussion 3.1. Metal complex formation Potentiometric titration curves of TMACHCAT (L1) and CYCOENCAT (L2) in the absence and presence of Al(III) ion at l = 0.1 M KCl and 25 ± 1 °C in aqueous medium are shown in Fig. 2, where the solid symbols represent equilibrium points collected when no solid phase was present in solution while dotted line represent points collected when turbidity or precipitation appeared in the solution. The deviation in the metal-ligand titration curves from the free ligand curve implies the formation of metal complexes. Also, the shape of titration curves qualitatively indicates that both the ligands have considerable affinity for the metal ion. The equilibrium points collected before a = 0, where ‘a’ is moles of base added per mole of ligand present, is due to the neutralization of excess acid present in the solution. The deviation in the metal-ligand titration curves from the free ligand curves started after a = 0 and the first break in the metal-ligand titration curves obtained at a = 3, where the curve for L2 lying at low pH region was expected to form stronger complexes than L1. When the pH increases further, the equilibrium points collected nearly up to a = 6 indicates that both the

1

8

pH

+

L :H 1 +3 L :Al 2 + L :H 2 +3 L :Al

9

M þ L þ 3H MLH3 ;

b113 ¼

M þ L þ 2H MLH2 ;

b112 ¼

½MLH3 

ð1Þ

½M½L½H3 ½MLH2 

ð2Þ

½M½L½H2 ½MLH M þ L þ H MLH; b111 ¼ ½M½L½H ½ML M þ L ML; b110 ¼ ½M½L

ð3Þ ð4Þ

The pK (pK = log K) values of the metal complexes are also given in Table 1, which can be represented by Eqs. (5)–(7); and derived by considering MLH3 as the first species, which is formed by the reaction of M and LH3 followed by dissociation of protons in steps from MLH3 to form other species MLH2, MLH and ML. The gradual increase in pK values (Table 1) indicates that the successive release of protons from MLH3 is easier in the formation of MLH2 as compared to MLH and ML. Also, considering the similar protonation constants of L1 and L2, shape of the titration curves and the calculated formation constants for the different species, it can be suggested that both the ligands forming complexes with similar structures using the same donor atom sets. Orvig and coworkers [33,34] also suggested that the ligands with same protonation state formed similar kinds of complexes in solution using the same donor atom sets. ½MLH2 ½H ½MLH3  ½MLH½H MLH MLH2 MLH þ H; K MLH2 ¼ ½MLH2  ½ML½H ML MLH ML þ H; K MLH ¼ ½MLH 2 K MLH MLH3 ¼

MLH3 MLH2 þH;

ð5Þ ð6Þ ð7Þ

Although Al(III) gave no specific band in the electronic spectrum to ascertain the different chelate(s) formed in solution but the modes of coordination can be explained from the shift in the intra-ligand transition with the supporting help of the potentiometric results. Similar to potentiometric method, spectrophotometric titrations of L1 and L2 were carried out in 1:1 metal-ligand ratio by keeping the ligand concentration [L] = 4.02  105 M and metal ion concentration [M(III)] = 4.02  105 M with varying pH between 3.5 and

11 10

ligands are able to release six protons prior to the precipitation occurred. Keeping in view these preliminary observations many sets of possible complexes were tested in the minimization programme, and the best-fit models were obtained when formation of species of the types MLH3, MLH2, MLH and ML were considered. No additional species were detected from the data obtained between 1:1 and 1:2 metal-ligand molar ratios. Inclusions of dinuclear species in the model for complexes of L2 worsen the fit. The overall formation constants (log b) of the species were calculated using Hyperquad 2000 program are summarized in Table 1. The equilibrium reactions for the overall formations of the complexes are given by the Eqs. (1)–(4) (charges are omitted for clarity):

7

Table 1 Overall formation constants (log b) of the metal complexes at 25 ± 1 °C and l = 0.1 M KCl, (A = potentiometry and B = spectrophotometry)

6 5

Equilibrium

4 3 -6

-4

-2

0

2

4

6

a Fig. 2. Potentiometric titration curves of L1 and L2 in the absence and presence of Al(III) in 1:1 ligand-metal molar ratio, where ‘a’ is moles of base added per mole of ligand present.

[MLH3]/[M][L][H]3 [MLH2]/[M][L][H]2 [MLH]/[M][L][H] [ML]/[M][L] [MLH2][H]/[MLH3]a [MLH][H]/[MLH2]a [ML][H]/[MLH]a a

TMACHCAT

CYCOENCAT

A

B

A

B

45.86 ± 0.03 41.19 ± 0.02 35.67 ± 0.03 28.53 ± 0.09 3.67 5.52 7.14

45.36 ± 0.03 41.09 ± 0.02 35.57 ± 0.03 28.44 ± 0.09 4.27 5.52 7.13

45.51 ± 0.03 40.74 ± 0.02 35.35 ± 0.03 29.78 ± 0.09 4.77 5.39 5.57

45.48 ± 0.01 40.72 ± 0.04 35.31 ± 0.08 29.74 ± 0.05 4.76 5.41 5.57

Represents pK values (pK = log K).

1584

M. Baral et al. / Journal of Inorganic Biochemistry 102 (2008) 1581–1588

7.5. Above the pH 7.5, the solution became turbid. The experimental electronic spectra of L1- and L2-Al(III) systems are given in Fig. 3. The ligand peak at 280 nm was shifted towards higher wavelength with concomitant rise in the absorbance upon successive rise in pH led to the formation of two isosbestic points, which indicate the formation of metal complexes. The variations in electronic spectra for both the ligands at different pH are almost similar with respect to the metal ion suggesting similar complexation reactions in solution. The experimental electronic spectral data were refined using a nonlinear least square program, pHAb and the formation constants calculated for the best-fit models were summarized in Table 1 and their predicted electronic spectra are given in Fig. 3c and d. In order to explain the possible modes of coordination in the various species depicted in solution, it is necessary to explain about the deprotonation groups of the free ligands (L1 and L2). Both the ligands within the adopted experimental conditions (pH 2.5– 11.5) gave six protonation constants assigned to three secondary amino groups and three more acidic catecholic oxygens at ortho. No values were determined for the catecholic oxygens at meta. Thus, the neutral and fully protonated form of ligands L1 and L2 can be represented by LH3 and LHþ3 6 , respectively. From the pH dependent species distribution curves (Fig. 4), it was found that complexation begins from pH 3.0 with the formation of partially

a

protonated species MLH3 (L = L1 or L2). This species formed with the interaction of M with LH3. The fully protonated form, LH3þ 6 (Fig. 5a) releases three protons to give the coordinated ligand LH3 in MLH3 and obviously these protons must be released from the three more acidic catecholic oxygens at ortho. The complex MLH3 can be structurally represented by the monocapped geometry as shown in Fig. 5b, where three secondary amine nitrogen atoms remain in protonated form. Similar monocapped type coordination mode in tripodal ligands with different trivalent metal ions has been reported by Orvig and coworkers in both the solid and solution states [35–37]. Again, since it has been well established that Al(III) exists as [M(H2O)6]3+ in aqueous solution [38], three water molecules may be coordinated to the metal ion in MLH3 to fulfill the hexa-coordinated configuration. Moreover, the formation of MLH3 complex can be supportively suggested by comparing its kmax with that of LH3. The species LH3 shows maximum absorbance at 294 nm [16,17], whereas in 1:1 M/L solution the said peak appeared at 281 nm assignable to the coordination of the three catecholic groups to the metal ion forming the species, MLH3. As the pH increases subsequently, successive deprotonation from MLH3 led to the formation of three different complexes MLH2, MLH and ML prior to the solid phase appeared. The extrusion of protons from the complex MLH3 may take place either from

b

0.6

1.4 0.5 1.2

Absorbance

Absorbance

0.4

pH=7.02 0.3

0.2

pH = 7.47

1.0 0.8 0.6 0.4

0.1

pH=3.57

pH = 3.67

0.2 0.0 240

260

280

300

320

340

225

250

Wavelength, nm

c

300

325

350

d 30 30

2

ML

Molar absorptivity X 103

Molar absorptivity X 103

275

Wavelength, nm

2

ML H

20

2

2

ML H3

ML H2

10

20

ML

3

3

ML H 3

ML H2

10

3

ML H3 240

260

280

300

Wavelength, nm

320

340

225

250

275

300

325

350

Wavelength, nm

Fig. 3. Experimental electronic spectra at 1:1 metal to ligand ratio for (a) Al(III)-L1and (b) Al(III)-L2 systems, and the predicated electronic spectra of the four species for (c) Al(III)-L1 and (d) Al(III)-L2 systems.

M. Baral et al. / Journal of Inorganic Biochemistry 102 (2008) 1581–1588

a

Al

% of formation

change from equatorial to axial form upon complexation. Previously, we have found that the relative strain energies for the cis,cis-equatorial conformation of the ligands L1 and L2 were stable by 50 and 102 kcal/mol, respectively than cis,cis-axial conformers [17]. Since, ring-flipping upon complexation was reported for many tripodal aminophenolate ligands derived from cyclohexane based tripodal triamine, cis,cis-1,3,5-trisaminocyclohexane (TACH) [40–42] it can be expected that an additional strain energy required to be compensated in order to flip from one conformation to another. Again, the intimate relationship of complex stability to the size of the metal ion and size of the chelate ring is well documented [43,44] and generally stability increases as the chelate ring size changes from six to five [44]. On the basis of these observations, one can rationalize the preference of tris(catecholate) bonding for Al(III) than the tris(aminocatecholate). It has already been reported that tripodal catecholate ligands yielded strong tris(catechol) type complexes with group 13 metal ions [45–47]. Thus, oxygen–oxygen encapsulated structures as shown in Fig. 5c–e can be suggested for the species MLH2, MLH and ML, and consequently, the amino groups remain in the protonated forms. Similar type of coordination with aluminium, where the amino groups are in protonated form was also reported earlier for some hydroxamate containing ligands derived from 1,4,8,11tetraazacyclotetradecane and ferrioxamine B [48,49].

100

80

1

AlL H3

1

AlL H 1

AlL H2

60

1

AlL 40

20

0 3

4

5

6

7

pH

b

100

Al

% of formation

80

1585

2

AlL H3 2

AlL

60 2

AlL H2

3.2. Selectivity and binding abilities of the ligands 2

40

AlL H

20

0 3

4

5

6

7

8

pH Fig. 4. Speciation diagram (% of formation vs. pH) calculated for 1:1 Al(III)L, where L = (a) L1 and (b) L2.

the NHþ 2 groups for which the protonation constant values have been evaluated (can be represented by Fig. 5f–h) or from catechol units, for three of which the protonation constant values are unknown (can be represented by Fig. 5c–e). These two probabilities lead to two distinct geometries, either nitrogen–oxygen encapsulated or oxygen–oxygen encapsulated. In the first case, upon coordination if the metal ion facilitates to enter into the nitrogen–oxygen cavity with release of proton from the protonated secondary amine nitrogen atom, it may encounter with a repulsive force due to its higher charge-to-ionic radii in the presence of other two protonated NHþ 2 groups. However, in the second case of oxygen–oxygen encapsulation, the ligand coordinates through its catecholic oxygens as anionic donors. Al(III) is the hard metal ion, which shows strong preference towards hard donor atoms such as negatively charged oxygen atoms. Its affinity towards nitrogen atoms is very poor [26]. Farkas and Csoka [39] demonstrated the aqueous coordination properties of Al(III) by using a ligand 2,3dihydroxy-phenylalanine-hydroxamic acid (DOPAHA) which has three different bonding sites: catecholate, hydroxamate and amine types, where the formation of both catecholate and hydroxamate chelates were reported but not the amine-type coordination mode with Al(III). Also, from Fig. 3c and d, it was observed that the deprotonation and subsequent coordination in the species MLH3 resulted a bathochromic shift in absorbance maximum. Since, the catechol units are the only chromophores in the ligands, shift in the kmax can be assigned to the further deprotonation and/or coordination of catecholic oxygens. In nitrogen–oxygen encapsulated case, there is more probability for the ligands to face the conformational

Table 1 reveals that both the ligands, L1 and L2 form strong complexes with aluminium(III). The ligand L2showed higher affinity with approximately one log b unit higher than the ligand L1. The higher selectivity of ligand L2 over L1 may be due to its more flexible and pre-organized structure for metal ion encapsulation. As reported in our previous communication [17], the ligand L2 is pre-organized due to the presence of intra- and intra-strand Hbonding, which brings the three coordinating sites closer to each other and the cavity made between the three catechol units is suited for metal ion encapsulation. On comparing the log bFeL(L1 = 27.14 and L2 = 31.02) [17], the Fe(III)L1 complex showed lower stability by approximately one log unit from Al(III)L1. The higher stability of Al(III)L1 over Fe(III)L1 complex may be due to the differences in the preferred orientations of the donor groups in the coordination sphere of the metal ion and the forced distortion from the regular octahedron [50] as in case of Ga(NOTA) (log b = 30.98) and Fe(NOTA) (log b = 28.30) [51]. The ligands (L1 and L2) are weak acids, the proton competition occurs depending on their pKa and the pH. Thus, the pM (pM = log [Mn+]) is a better value of the relative complexation efficiency of the ligands under given conditions of pH, Mn+ and L concentrations. The pM (M = Al3+) values (Table 2) concerning L1 and L2 calculated at pH 7.4, [L]total = 5  104 M (L = L1 or L2) and [Al3+]total = 5  105 M are 18.62 (L1) and 19.59 (L2), respectively. This result provides evidence that, at physiological pH, ligand L2 binds Al(III) more selectively than the ligand L1. The pH dependence of the selectivity was examined owing to the fact that in particular biological compartments or circumstances, pH values far from 7 can be observed. Hence, it is important to see whether the selectivity of ligand still remains in a biological relevant pH range. The plots of pM as a function of the pH are given in Fig. 6 for both L1 and L2. They indicate that the selectivity of ligands L1 and L2 for Al(III) is maintained over the pH range 4–8. 3.3. Comparisons with other multidentate ligands The ligands L1 and L2 have two compartments: (i) nitrogen–oxygen [tris(aminocatecholate)] and (ii) oxygen–oxygen [tris(catecholate)], which allow to make comparisons with some tripodal imino/

1586

M. Baral et al. / Journal of Inorganic Biochemistry 102 (2008) 1581–1588

H

H H H + H N + H N H HO HO HO HO

+

NH H OH OH

(a) LH 6 -3H +

H

Coordination modes: Tris(catecholate) type

-H +

NH H O

H + N

H O M

O

OH

Coordination modes: Tris(aminocatecholate) type

-H +

H H

= CH 2 (L 1 ); CO N H CH 2 CH 2 (L 2 )

H + N

+

H + H NH N + H H N H O O O OH HO M HO (b) M LH 3 +

H

H H

H

H H

H O

O

HO

OH

H O

H + N H O

HO

HO

H + N

NH M

(f) M LH 2

(c) M LH 2

-H +

-H + H H

H

NH H O

H + N

H + N

+

H O

H O M

M

O HO

OH

(d) M LH

(g) M LH

-H +

-H +

NH H O

+

HN H O M O

+

H + N H O

O

O

OH

NH M

HO

NH

NH

O

O

HO

(e) M L

H + N H O

H H

H

H H

H

O

NH

NH O

HO

O

O

H H

H

HO

(h) M L 1

Fig. 5. Possible coordination modes for the various species formed upon complexation of ligands L and L2 with Al(III).

aminophenolate ligands, EDTA, DTPA, transferrin and desferrioxamine B. The pAl values (Table 3) indicate that the ligands L1 and L2 show higher preference than the tripodal N3O3-donors imino/ aminophenolate ligands and other ligands (EDTA and DTPA) containing nitrogen donors. The enhanced stability of the ligands L1 and L2 is due to the formation of tris(catecholate) type complexes

and this is axiomatic because of the high affinity of Al(III) for the anionic oxygen donor. Both the ligands showed appreciably high affinity towards Al(III) than the transferrin. It is known that most of the aluminium circulating in blood plasma is bound to transferrin through the sites left vacant by iron [56]. Because of the labile nature of the aluminium transferrin complex [57–59], the alumin-

M. Baral et al. / Journal of Inorganic Biochemistry 102 (2008) 1581–1588 Table 2 Comparative pM (pM = log [Mn+], where Mn+ = Al3+) values calculated at pH 7.4 (L:M = 10, [L] = 5  104 M ) Ligands

pM

References

Transferrin TAMS6 DTPA TAPS6 EDTA Saltriens Saltames L1 Desferrioxamine B L2

14.5 15.2 15.55 15.8 15.89 15.9 16.9 18.62 19.4 19.59

[50] [32] [51] [32] [52,54] [53,55] [53,55] Present work [47] Present work

22

1587

ium overload is cured by sequestering agents and the only approved drug available is DFO [18]. The calculated pAl values for L1 and L2 are quite close to DFO (Table 2) where the ligand L2 showed slightly higher affinity than DFO. Concerning the behaviour of the ligands (L1 and L2) and their aluminium chelates i.e., high affinity for Al(III), comparable affinity with DFO, aqueous solubility over a wide range of pH and no hydroxo species with Al(III), make these two ligands useful for aluminium sequestration. Moreover, Santos and coworkers [48] suggested that sequestering ligands with free amines as in desferrioxamine B may improve their usefulness as a drug, where the amine groups are used as a point of attachment to a polymeric solid matrix, which enhanced the properties of removing ‘‘hard” ions from water solutions. Such free amines are also present in ligands L1 and L2 and also in their complexes, depicted in solution. 3.4. Molecular modeling calculations

20 18

AlL

1

AlL

2

16 14

pM 12 10 8 6 4 4

5

6

7

8

pH Fig. 6. Plot of pM (pM = log [Mn+]) vs. pH. pM was calculated for [L] = 5  104 M (L = L1 or L2), [Mn+] = 5  105 M (Mn+ = Al3+) using the deprotonation constants of ligands L1 and L2, and the complexation constants b11n.

Table 3 Calculated least strain energy (ETin kcal/mol) for MLH3 and ML type complexes

Molecular modeling studies provide complementary information on the coordination modes. The experimentally proposed coordination modes for the different species MLH3, MLH2, MLH and ML [L = L1 and L2, M = Al(III)] are shown in Fig. 5. Since, the ligands are isolated in cis,cis-equatorial forms, it may be expected that the encapsulation of metal ion may take place with or without any change in conformation to cis,cis-axial forms. Molecular mechanics calculations using MM3 force field were done for the initial (AlLH3) and final (AlL) species using both axial as well as equatorial forms of the ligands. The calculated strain energies (Table 3) reveal that both the ligands retain their cis,cis-equatorial conformations in their respective metal chelates. The least strain structures of AlL1and AlL2 complexes were reoptimized by applying semi-empirical PM3 method. The optimized structures of AlL1and AlL2chelates (Fig. 7) predicted a distorted octahedral type geometry and showed the presence of intramolecular hydrogen bonds between amine protons and catecholic oxygens. Such intramolecular H-bonds are known to provide an extra stability for the formation of tripodal tris(catechol) type encapsulated complex [1,47]. 4. Conclusion

L

MLH3

ML

3-axial

3-equitorial

3-axial

3-equitorial

L1 L2

489.11 582.54

532.96 593.11

359.35 370.92

365.61 412.91

The tripodal amine catechol ligands L1 and L2 formed stable complexes with aluminium(III) in solution, where only the catechol units were taken part in the complex formation to give

Fig. 7. Optimized structures of the (a) AlL1 and (b) AlL2 species through semi-empirical PM3 method.

1588

M. Baral et al. / Journal of Inorganic Biochemistry 102 (2008) 1581–1588

tris(catecholate) type of complex. Molecular modeling calculations predicted a distorted octahedral geometry for the metal complexes, where the ligands retained their cis,cis-equatorial conformations. The ligand L2 showed higher affinity for Al(III) than L1. The pAl values calculated for AlL1 and AlL2 were appreciably higher than transferrin. The ligand L2 showed slightly higher affinity towards Al(III) than desferrioxamine (DFO), the only approved drug for the treatment of aluminium intoxication. The high affinity towards Al(III) comparable pM values with DFO and the water solubility over a wide range of pH without formation of hydroxo species, hopefully make these ligands L1 and L2, useful for Aluminuium sequesteration.

[14] [15] [16] [17] [18]

5. Abbreviations

[24] [25] [26]

M L L1

Metal ion Ligand cis,cis-1,3,5-Tris[(2,3-dihydroxybenzylamine)aminomethyl]cyclohexane (TMACHCAT) L2 N1,N3,N5-Tris(2-(2,3-dihydroxybenzylamino)ethyl)cyclohexane-1,3,5-tricarboxamide (CYCOENCAT) DOPAHA 2,3-Dihydroxy-phenylalanine-hydroxamic acid NOTA 1,4,7-Triazacyclonane-N,N0 ,N000 -triacetic acid DFO Desferrioxamine B TAMS6 1,1,1-Tris(2-hydroxo-5-sulfobenzylaminomethyl) ethane DTPA Diethylenetriamine-pentaacetic acid TAPS6 1,2,3-Tris(2-hydroxo-5-sulfobenzylamino)propane EDTA Ethylenediamine-tetraacetic acid Saltriens 1,10-Bis(2-hydroxy-5-sulfobenzylidene)-1,4,7,10-tetraazadecane Saltames 1,1,1-Tris(((2-hydroxy-5-sulfonatobenzylidene)amino)methyl)ethane TACH cis,cis-1,3,5-Trisaminocyclohexane

References [1] K.N. Raymond, Coord. Chem. Rev. 105 (1990) 135–153. [2] J.R. Budge, P.E. Ellis, R.D. Jones, J.E. Linard, F. Basolo, J.E. Baldwin, R.L. Dyer, J. Am. Chem. Soc. 101 (1979) 4760–4762. [3] R.D. Jones, D.A. Summerville, F. Basolo, Chem. Rev. 79 (1979) 139–147. [4] A.R. Battersby, A.D. Hamilton, J. Chem. Soc. Chem. Commun. (1980) 117–119. [5] W.R. Harris, C.J. Carrano, K.N. Raymond, J. Am. Chem. Soc. 101 (1979) 2213– 2214. [6] K.N. Raymond, G. Muller, B.F. Matzanke, in: F.L. Boscheke (Ed.), Topics in Current Chemistry, vol. 123, Springer, New York, 1984, pp. 49–102. [7] A.L. Crumbliss, in: G. Winkelmann (Ed.), Handbook of Microbial Iron Chelates, CRC Press, New York, 1991, pp. 177–233. [8] L.D. Loomis, K.N. Raymond, Inorg. Chem. 30 (1991) 906–911. [9] R.C. Hider, G.S. Tilbrook, in: A. Sigel, H. Sigel (Eds.), Iron Transport and Storage in Microorganisms, Plants and Animals, Marcel Dekker, New York, 1998, pp. 691–730. [10] M. Hayashi, K. Hiratani, S.I. Kina, M. Ishii, K. Saigo, Tetrahedron Lett. 39 (1998) 6211–6220. [11] L. Stryer, Biochemistry, WH Freeman, San Francisco, CA, 1975. pp. 796–797. [12] R. Fried, J. Chem. Educ. 45 (1968) 322–335. [13] G.C. Cotzias, M.H. Van Woert, L.M. Schiffer, New Engl. J. Med. 276 (1967) 374– 379.

[19] [20] [21] [22]

[23]

[27] [28] [29] [30] [31] [32] [33] [34] [35] [36] [37] [38] [39] [40] [41] [42] [43] [44] [45] [46] [47] [48] [49] [50] [51] [52] [53] [54] [55] [56] [57] [58] [59]

J. Fermaglich, T.N. Chase, Lancet 1 (1973) 1261–1262. B.A. Stoll, Lancet 1 (1972) 431–432. S.K. Sahoo, M. Baral, B.K. Kanungo, Polyhedron 25 (2006) 722–736. S.K. Sahoo, PhD Thesis, Punjab Technical University, Jalandhar, India, 2007. R.C. Hider, A.D. Hall, in: R.W. Hay, J.R. Dilworth, K.B. Nolan (Eds.), Perspectives in Bioinorganic Chemistry, vol. 1, JAI Press, London, 1991, pp. 209–253. A.C. Alfrey, G.R. LeGendre, W.D. Kaehny, New Engl. J. Med. 294 (1976) 184– 188. I. Klatzo, H. Wisniewski, E. Streicher, J. Neuropathol. Exp. Neur. 24 (1965) 187– 199. D.P. Perl, D.C. Gajdusek, R.M. Garruto, R.T. Yanagihara, C.J. Gibbs, Science 217 (1982) 1053–1055. N.C. Bowdler, D.S. Beasley, E.C. Fritze, A.M. Goulette, J.D. Hatton, J. Hession, D.L. Ostman, D.L. Rugg, C.J. Schmittdiel, Pharmacol. Biochem. Behav. 10 (1979) 505–512. T.P.A. Kruck, E.A. Fisher, D.R.C. McLachlan, Clin. Pharmacol. Ther. 48 (1990) 439–446. S. Descroches, F. Biron, G. Berthon, J. Inorg. Biochem. 75 (1999) 27–35. A.E. Martell, R.J. Motekaitis, R.M. Smith, Polyhedron 9 (1990) 171–187. F. Mulla, F. Marsicano, B.S. Nakani, R.D. Hancock, Inorg. Chem. 24 (1985) 3076– 3080. A.E. Martell, R.J. Motekaitis, The Determination and Use of Stability Constants, VCH Publishers, New York, 1992. P. Gans, A. Sabatini, A. Vacca, Talanta 43 (1996) 1739–1753. P. Gans, A. Sabatini, A. Vacca, Ann. di Chim. 89 (1999) 45–49. User guide manual for CAChe version 6.01, Fujitsu limited, 2003. B.P. Hay, D.A. Dixon, R. Vargas, J. Garza, K.N. Raymond, Inorg. Chem. 40 (2001) 3922–3935. M. Seitz, H.G. Alt, J. Mol. Cat. A: Chem. 257 (2006) 73–77. A.K.W. Stephens, C. Orvig, J. Chem. Soc., Dalton Trans. (1998) 3049–3056. P. Caravan, C. Orvig, Inorg. Chem. 36 (1997) 236–248. M.P. Lowe, P. Caravan, S.J. Rettig, C. Orvig, Inorg. Chem. 37 (1998) 1637–1647. P. Caravan, T. Hedlund, S. Liu, S. Sjoberg, C. Orvig, J. Am. Chem. Soc. 117 (1995) 11230–11238. S. Liu, L. Gelmini, S.J. Rettig, R.C. Thompson, C. Orvig, J. Am. Chem. Soc. 114 (1992) 6081–6087. F.A. Cotton, G. Wilkinson, C.A. Murillo, M. Bochmann, Advanced Inorganic Chemistry, sixth ed., John Wiley & Sons Inc., Singapore, 1999. E. Farkas, H. Csoka, J. Inorg. Biochem. 89 (2002) 219–226. J.E. Bollinger, J.T. Mague, W.A. Banks, A.J. Kastin, D.M. Roundhill, Inorg. Chem. 34 (1995) 2143–2152. J.E. Bollinger, J.T. Mague, D.M. Roundhill, Inorg. Chem. 33 (1994) 1241–1242. J.E. Bollinger, J.T. Mague, C.J. O’Connor, W.A. Banks, D.M. Roundhill, J. Chem. Soc., Dalton Trans. (1995) 1677–1688. R.D. Hancock, A.E. Martell, Chem. Rev. 89 (1989) 1875–1914. R.D. Hancock, J. Chem. Educ. 69 (1992) 615–621. S.M. Moerlein, M.J. Welch, K.N. Raymond, F.L. Weitl, J. Nucl. Med. 22 (1981) 710–719. S.M. Moerlein, M.J. Welch, K.N. Raymond, J. Nucl. Med. 23 (1982) 501–506. S.M. Cohen, K.N. Raymond, Inorg. Chem. 39 (2000) 3624–3631. M. Gaspar, R. Grazina, A. Bodor, E. Farkas, M.A. Santos, J. Chem. Soc., Dalton Trans. (1999) 799–806. A. Evers, R.D. Hancock, A.E. Martell, R.J. Motekaitis, Inorg. Chem. 28 (1989) 2189–2195. K. Wieghardt, U. Bosseck, P. Chaudhuri, W. Herrman, B.C. Menke, J. Weiss, J. Inorg. Chem. 21 (1982) 4308–4314. E.T. Clarke, A.E. Martell, Inorg. Chim. Acta 186 (1991) 103–111. W.R. Harris, J. Sheldon, Inorg. Chem. 29 (1990) 119–124. D.A. Aikens, F.J. Bahbah, Anal. Chem. 39 (1967) 646–649. G. Schwarzenbach, R. Gutt, G. Anderegg, Helv. Chim. Acta 37 (1954) 937– 957. D.F. Evans, D.A. Jakubovic, Polyhedron 7 (1988) 1881–1889. S.J.A. Fatemi, F.H.A. Kadir, G.R. Moore, Biochem. J. 280 (1991) 527–532. R.B. Martin, J. Savory, S. Brown, R.L. Bertholf, M.R. Wills, Clin. Chem. 33 (1987) 405–407. G. Berthon, Coord. Chem. Rev. 149 (1996) 241–280. H.M. Marques, J. Inorg. Biochem. 41 (1991) 187–193.