The Complexation of AlIII, PbII, and CuII Metal Ions by Esculetin: A

Article pubs.acs.org/JPCA

The Complexation of AlIII, PbII, and CuII Metal Ions by Esculetin: A Spectroscopic and Theoretical Approach Annaïg Le Person,* Aurélien Moncomble, and Jean-Paul Cornard LASIR, CNRS UMR8516, Université Lille 1 Sciences et Technologies, Bât C5 − 59 655 Villeneuve d’Ascq Cedex − France S Supporting Information *

ABSTRACT: UV−visible absorption spectroscopy combined with quantum chemical calculations and, notably, Time-Dependent Density Functional Theory were used to probe the structure of metal complexes with esculetin in dilute aqueous solution, at pH = 5. For the 1:1 complex formation, the studied metal ions can be classified according to their complexing power: aluminum(III) > copper(II) > lead(II). For the three complexes, a chelate is formed with the fully deprotonated catechol moiety and an absorption band is observed at the same wavelength. In all cases, a pronounced ionic character is calculated for metal−ligand bonds. However, the complexes differ in their coordination sphere. Copper and lead are bound to two water molecules leading to a square plane geometry and a hemidirected complex, respectively, whereas aluminum atom has an octahedral environment involving three water molecules and a hydroxide ion. For AlIII only, a 2:1 complex is observed, and the involvement of an aluminum dimer was evidenced.

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INTRODUCTION Dissolved organic matter (DOM) plays an important role in many biogeochemical processes.1−3 The acid soluble part of DOM is called fulvic acids. Smaller than humic acids,4,5 these acids result from the association of relatively small organic molecules.6−8 Among the phenomena that occur in such a medium, metal complexation can influence the transport and the bioavailability to plants of metal ions and organic molecules. Despite a better understanding of this type of interaction,9−11 the complexation of many organic molecules such as esculetin remains barely elucidated. Esculetin (Figure 1) is a hydroxycoumarin (6,7-dihydroxycoumarin) contained in many plants12,13 such as tea14 and

the risks associated with increased exposure to this metal is growing.18,19 In very high doses, aluminum becomes toxic and cause neurotoxicity.20 The contamination of acid soils with aluminum also affects the plant growth.21−23 However, the aluminum toxicity remains much weaker than that of heavy metals such as lead. Lead pollution occurs through a variety of activities, including mining, building construction, metal processing, battery manufacture and disposal, burning of leadcontaining fuels, and application of sludge to agricultural land.24 Lead is a highly poisonous metal, the main target of which is the nervous system. It causes many damages such as brain and blood disorders, it reduces fertility and is responsible of cognitive deficits among children.25−28 Finally, copper is naturally present in the environment and also has industrial (electrical wires, roofing, plumbing and industrial machinery) and agricultural (fungicide) uses. This metal is vital for organisms, but it becomes toxic and responsible of serious health concerns in the case of excessive accumulation.29,30 It also presents long-term effects on crop yields and soil quality31,32 and also risks of accumulation in crop tissues.33 The main objective of this study is to identify the nature of the complexes formed with metal cations and especially (i) the site preferentially involved in the metal fixation, (ii) the protonation state of this site, (iii) the metal cation environment, and (iv) the characteristics of the metal−ligand bonds. Electronic absorption spectroscopy was combined with quantum chemistry calculations in order to characterize the structure of the complexes formed in aqueous solution. The stoichiometries, the formation constants and the pure UV−visible absorption spectra of the

Figure 1. Chemical structure and atomic numbering of esculetin.

medicinal plants.15 Indeed, it presents antioxidative characteristics.16 This molecule potentially present in fulvic acids17 is likely to be complexed by metal ions. This article reports for the first time the study of the complexation with esculetin of three pollutant metal cations: aluminum(III), lead(II), and copper(II) ions. Aluminum and copper are naturally present in the environment, but like most of the metals, they can represent a risk if their concentration becomes too high. Despite aluminum being the most abundant metal in the earth’s crust, it has no known biological function. This metal is transferred to the environment by many anthropic activities such as transportation, packaging, construction, etc. The number of studies pointing out © 2014 American Chemical Society

Received: December 16, 2013 Revised: March 21, 2014 Published: March 21, 2014 2646

dx.doi.org/10.1021/jp412291z | J. Phys. Chem. A 2014, 118, 2646−2655

The Journal of Physical Chemistry A

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(water) was represented by explicit molecules added in the coordination sphere of metal ions (the number of water molecules is determined by examination of the release of molecules during the optimization) and by the addition around this complex of an implicit solvent using the C-PCM model.47 This model is necessary to take into account the fact that solvent molecules around metal cations act both like a solvent and a ligand. An a posteriori study of the solvatation model was carried out on the basis of the actual complexes. The results, reported in the Supporting Information, confirm that the implicit model is not sufficient, especially in the presence of hydroxide ions in the coordination sphere. TD-DFT formalism in the adiabatic approximation was used to compute excitation energies and oscillators strengths.48,49 TDDFT has shown its ability to reproduce absorption and emission properties as illustrated by the increasing number of publications using this method both for small and complex systems.50−53 The 40 lowest excited states (80 in the case of copper complexes) were taken into account to compute vertical excitation energies. Atoms-in-molecules (AIM) computations54 were carried out to give some insight about the bond nature in the studied complexes. To avoid the apparition of non-nuclear attractor critical points, the electron density was computed by a singlepoint calculation on the previously optimized geometry using for Pb the all-electron WTBS basis set55 (no change for other atoms). Bond critical points (BCPs) were localized, and local and integrated properties were computed using the AIMAll software.56

complexes were determined. The molecular geometries of the different species were optimized and the electronic transitions were calculated by methods based on the density functional theory. All the measurements were performed at constant pH (pH = 5) because this pH is commonly observed for fulvic acids and avoids precipitation of metal hydroxides.

1. EXPERIMENTAL AND THEORETICAL METHODS 1.1. Chemicals. Esculetin, sodium chloride (NaCl), sodium hydroxide (NaOH), hexahydrated aluminum chloride (AlCl3(H2O)6), lead chloride (PbCl2), and copper chloride (CuCl2) were obtained from Sigma Aldrich and used as received without further purification. Deionized water was used as the solvent. 1.2. Spectroscopy. UV−visible spectra were recorded on a CARY 100 (Varian) double-beam spectrophotometer, in the region of 200−700 nm using cells of 1 cm path length. 1.3. Experimental Method. A solution containing aqueous esculetin at a concentration of 5.10−5 mol L−1 was mixed with sodium chloride in order to keep the ionic strength constant (0.1 mol L−1). Incremental volumes of a solution of metal salt were added to the solution in order to vary the molar ratio R (nmetal/ nesculetin) and the pH was maintained constant (pH = 5) by small additions of NaOH. A Minipuls II (Gibson) peristaltic pump was used to circulate solution from the titration beaker to the cell (Hellma) for the absorption measurements. UV−visible spectra of the solutions were recorded for each molar ratio. The electronic spectra were corrected due to the dilution of the ligand solution. First, the molar ratio method34 was used from UV− Visible spectra in order to estimate the stoichiometries of the complexes formed. Then, chemometric methods were used for a more extensive treatment of the results. 1.4. Chemometric Methods. The absorption spectra acquired for various metal/ligand molar ratio were analyzed using the Reactlab equilibrium program.35 This program allows the determination of the number of species that contribute to the spectra using a factor analysis procedure (EFA)36−38 and of the electronic spectra of each pure species. The Reactlab equilibrium program has also been used to estimate the apparent formation constants (β) of the different complexed forms. In order to obtain the best fit between the complexation model and the experimental data, several models of complexes were envisaged for the refinement of the formation constants. The apparent formation constant:

2. RESULTS AND DISCUSSION 2.1. Free Ligand. In a first step, we focused our attention on the spectral properties of the free ligand and notably on its protonation state in our experimental conditions (pH = 5). Indeed, the data found in the literature concerning the pKa values are very inconsistent.57,58 The variation of esculetin UV−visible spectra with pH was studied and shown in Figure 2. For low pH values, the electronic spectrum is mainly characterized by a double band at 295 and 345 nm. By increasing the pH, the band intensity at 345 nm decreased, whereas a new band localized at 385 nm appeared. In addition, an isosbestic point is observed at 355 nm, indicating the presence of an equilibrium between two

βx : y = [MxLy]/([M]x ·[L]y )

results from the equilibrium between the free ligand L, the metal cation M and the complex MxLy (without taking into account the protonation state of the ligand): x M + y L ⇄ M xL y

1.5. Computational Details. All computations were performed using DFT-based methods with the Gaussian 09 software.39 The B3LYP global hybrid functional40−42 was used throughout due to its recognized ability to compute UV−visible spectra43 and the good result obtained for the absorption spectrum of the free ligand. H, C, O, Al, and Cu atoms were represented by the 6-311+G(d,p) basis set44,45 and Pb atom by the Hay-Wadt ECP and the corresponding double-ζ basis set.46 Geometrical optimizations were carried out using standard algorithms without any symmetry constraint and the nature of stationary points was checked by vibrational analysis. The solvent

Figure 2. Evolution with pH of the UV−visible spectra of esculetin (5 × 10−5 mol L−1). 2647



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the lowest energy transition mainly arises from the HOMO → LUMO transition (more than 95%). The KS orbitals involved are depicted on Figure 4 for the fully protonated form, but the main characteristics are the same for the four compounds. The transition consists in an electron transfer from the 6-hydroxyl group to the cyclic moiety; the contribution of the 7-hydroxyl group is nonzero but significantly smaller. This observation allows us to interpret the results reported in Table 1: when the 7hydroxyl group loses a proton, the effect on the computed absorption wavelength is strikingly smaller than when the deprotonation occurs on the 6-hydroxyl group. 2.2. Aluminum(III), Lead(II), and Copper(II) Complexes with Esculetin. 2.2.1. Electronic Spectroscopy Study. Variations of the UV−visible absorption spectra of esculetin upon aluminum(III), lead(II) and copper(II) addition for molar ratio ranging from 0 to 10 are presented on Figure 5. In all cases, the complexation results in a bathochromic shift of the absorption band from 345 to 385 nm. In the case of copper(II) and lead(II) complexation, an isosbestic point is observed at 358 nm for the full molar ratio range, indicating an equilibrium between the complexed and the free forms. The results of the molar ratio method show the formation of a 1:1 stoichiometry complex. In the case of aluminum(III) complexation, the isosbestic point localized at 358 nm disappears for molar ratio higher than 0.7, indicating the formation of a second complex of 2:1 stoichiometry (molar ratio method). The absorption spectra of the pure species (free ligand, 1:1 and 2:1 complexes) and the concentration profiles, both obtained from the experimental spectra using chemometric methods, are presented in Figure 6. Surprisingly, the absorption band in the long wavelength range of the three 1:1 complexes is located at the same wavelength (385 nm). Moreover, the absorption band position of the 2:1 complex of aluminum(III) is found at 377 nm. The concentration profiles show that the maximal concentration of the 1:1 complex of aluminum(III) (noted AlL) is reached for a 1:1 molar ratio. At this stage, 74% of the ligand is complexed: 48% is engaged in the 1:1 complex, and 26% in the 2:1 complex (noted Al2L). From a molar ratio of 3, the ligand is completely consumed, and the 2:1 complex becomes the predominant species in solution. The case of lead(II) and copper(II) complexation is simpler since only one 1:1 complex is formed. At the sight of the concentration profiles, the complexation power is lower for lead(II) ions than for aluminum(III) or copper(II) ions. Indeed, whereas around 74% of the ligand is involved in complexes with aluminum(III) or copper(II) ions for a molar ratio of 1.1, only 20% is consumed in the case of lead(II) complexation. This observation is in agreement with the formation constants of the complex estimated by the chemometric treatments of the spectra and presented in Table 2. The formation constants for aluminum and copper complexation are in the same order of magnitude (log β1:1 = 5.71 and 5.18, respectively), while the formation constant for lead complexation is lower (log β1:1 = 3.76). Finally, a pH decrease was systematically observed after the addition of each metal ion, indicating a deprotonation of the ligand upon complexation. Whereas this decrease is weak in the case of lead(II) and copper(II) complexation, it is higher in the case of aluminum(III) complexation. The pH was readjusted before each spectral measurement. 2.2.2. Theoretical Investigations of the Complexes. Structural Propositions. For the three metal cations, the experiments highlight the formation of a 1:1 complex, while only aluminum(III) permits the formation of a 2:1 complex. For the purpose of a good comparison, we first focused on the study

forms. This new band can then be assigned to a monodeprotonated form of the ligand. From the spectral data set, chemometric calculations were performed and allowed the determination of a pKa value of 7.6 ± 0.1. This value is in accordance with the value reported previously by Acero et al.57 Thus, the only species present at pH = 5 appears to be the fully protonated one. Then these results were confronted to those obtained from quantum chemistry calculations. The optimized structure of esculetin presents an intramolecular hydrogen bond in the catechol moiety. The lowest energy transition is computed at 337 nm, in good agreement with the experimental band observed at 345 nm. In the same way, the band observed at 295 nm is well reproduced by the calculation with a value of 289 nm. As depicted on Figure 3, the general shape of the spectrum is well described by the calculated excitation wavelengths.

Figure 3. Computed transitions (blue) and experimental spectrum (black) for esculetin.

From the fully protonated esculetin, two protons were removed successively allowing the proposition of three other structures: the two monodeprotonated (both optimized with an intramolecular hydrogen bond) and the doubly deprotonated. Comparison of computed excitations and experimental data allows us to identify unambiguously the monodeprotonated form, which is present in solution when pH increases. The experimental wavelength observed at 385 nm is very close to the computed value for the 7OH-deprotonated form (378 nm, Table 1). The preferential deprotonation of the 7-hydroxyl group is Table 1. Lowest Energy Transition Wavelength and Oscillator Strength Calculated for the Different Protonation States of Esculetin considered form

wavelength (nm)

transition energy (eV)

oscillator strength

fully protonated 7OH-deprotonated 6OH-deprotonated doubly deprotonated

337 378 445 486

3.68 3.28 2.79 2.55

0.247 0.487 0.128 0.264

consistent with the predictable enhancement of its acidity due to the stabilization of the charge by mesomerism. Thermodynamic data support this point: the Gibbs energy difference between the two monodeprotonated forms is 4.3 kcal mol−1 in favor of the 7OH-deprotonated one. The deprotonation leads to an increase of the charge transfer involved in the transition, explaining a more important measured absorbance and a larger value for the computed oscillator strength. In each case, the contribution to 2648



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Figure 4. Representation of the HOMO (a) and the LUMO (b) of the fully protonated esculetin.

Figure 5. Evolution of the UV−visible absorption spectra of esculetin (5.10−5 mol L−1) upon Al(III), Pb(II) and Cu(II) complexation in aqueous solution (pH = 5) for molar ratio ranging from 0 to 10.

Figure 6. (Top) Electronic absorption spectra of pure species (free ligand L and complexed ligand MxL), and (Bottom) evolution of the different species concentrations versus metal/ligand molar ratio upon Al(III), Pb(II) and Cu(II) complexation by esculetin.

of the former. As in earlier studies,9−11 our approach was to propose hypothetical structures for the different complexes, taking into account the complexation site and mode, but also the

coordination sphere of the cation and the protonation state of the ligand. Ten structures were proposed (Chart 1) for each compound: four for the bidente coordination on the catechol 2649



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Table 2. Logarithm of the Formation Constants β Estimated for the Main Complexes of the Aluminum(III), Lead(II) and Copper(II)−Esculetin Systems at pH = 5 in Aqueous Solution log β

aluminum(III)

lead(II)

copper(II)

1:1 2:1

5.71 ± 0.02 10.86 ± 0.04

3.76 ± 0.01

5.18 ± 0.01

case of copper(II), the two possibilities mentioned previously were observed depending on the coordination site. For a given structure, two different coordination spheres were sometimes obtained. The hypothetical complex in which the cation is coordinated on the catechol with the deprotonated 6-hydroxyl group (structure C) was not obtained. Indeed, proton transfers were always observed leading to other structures. In the case of lead(II), the coordination number of lead is either 4 or 5, and the geometry around lead is hemidirectional, consistent with our previous studies.62,63 Note, however, that the complex in which the cation is coordinated on the 6-hydroxyl group (structure E) was not obtained: like in the case of copper(II), hydrogen transfer was the cause of this failure. Absorption Properties. Once these structures obtained, the absorption spectra were computed for the 32 hypotheses. The first test to check the validity of a structure is the good reproduction of the lowest energy transition. According to the experimental results, this transition is located at 385 nm for the three complexes. Only transitions calculated between 350 and 425 nm (these values are chosen to correspond to an error of 0.3 eV on the transition energy relative to the experimental value) are reported in table 3. Only transitions that are spin-allowed and with an oscillator strength higher than 0.010 are given. For numerous structures, no transition was computed in this range, and consequently these hypothetical complexes can be directly excluded. By comparison with experimental data, the computed values propose with a good confidence structures for two of the three complexes. Indeed, the values that appear in bold (structure D) in the table for copper(II) and lead(II) are very consistent with the spectral features. The computed vertical transitions and the experimental spectra are depicted on Figure 7. The computed spectrum for the copper(II) complex reproduces all the features of the experimental one; the computed one for the lead(II) complex is quite good with a slight blue shift for the transition observed (experimentally) around 280 nm. Another structure for the copper(II) complex could be proposed with the cation on the catechol with the 7-hydroxyl group deprotonated (structure B); however, the small value of the oscillator strength for the transition occurring at 389 nm let us reject this hypothesis. For

(depending on the protonation of each hydroxyl), two for the monodente coordination with each hydroxyl (depending on the coordinating hydroxyl protonation), one for the monodente coordination on the CO moiety of the lactone and one for the bidente coordination on the lactone. Coordination Sphere. The coordination sphere of aluminum(III) is well-known,59,60 this cation being hexavalent. Thus, water molecules were added around the cation in aluminum(III) complexes to complete the valency of the cation to six. About copper(II), some ambiguities occur due to the Jahn−Teller effect that forbids a perfect hexavalent structure.61 To treat this case, we started geometry optimizations from an hexavalent copper complex, allowing water molecules either to get organized in a distorted octahedral structure or to be released for two of them leading to a square plane geometry. As explained in earlier studies,62,63 lead(II) was more difficult to treat, due to the flexibility of its coordination sphere showing generally a valency of four or five but with large variations. We used a strategy analogous to the one used for copper(II): the valency of lead(II) was initially supposed to be five, a water molecule being possibly released during geometry optimizations. Structure of the 1:1 Complexes. Due to the large number of structures obtained, we will not give detailed results for the whole set but the most characteristic results will be presented. First, the stabilization of a cation in a bidente mode on the lactone (structure J) was impossible independently of the metal cation nature, the structure optimization leading always to the monodente coordination on the CO part. Second, when a hydrogen bond is conceivable on the catechol moiety, it is always present in the computed structure. Except for aluminum(III), for which an octahedral coordination was always observed, several coordination sphere were obtained for other complexes. In the

Chart 1. Hypothetical Structures Used in the Calculations for the Complexes of Esculetin

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copper(II) complexes, the difference between distorted octahedral and square plane geometries is very low: the absorption properties are not so different for the high-intensity transition (less than 0.1 eV) that shows that the implicit model associated with a valency of 4 is sufficient in the case of study. In both cases, the cation binds on the catechol that releases two protons (structure D) and is surrounded by two water molecules. The structures are depicted on Figure 8. The geometry around copper is square plane, whereas lead presents an hemidirectional geometry. It is noteworthy that an AIM study of the lead complex structure (see below for the AIM results) shows the presence of a minimum of the laplacian of the electronic density around the cation that supports the presence of a lone-pair implying a hemidirected geometry. However, in the case of aluminum(III), we were not able to draw a conclusion from these data: no good correlation is observed between experimental and computed wavelengths. The proximity of the recorded absorption wavelengths of the three complexes leads us to propose that the three electronic structures must be very close. So the aluminum(III) should be coordinated to the doubly deprotonated catechol (structure D), and this cation should have a global +2 charge. This latter could be achieved by replacing a water molecule by a hydroxide ion (in axial or equatorial position). These propositions are consistent with previous works about the coordination of aluminum(III).9 So these structures were optimized, and their electronic spectra were computed (Table 4). The computed and experimental wavelengths are this time in very good agreement (Figure 9). It can be noticed that removing a proton from an equatorial or axial water molecule did not significantly modify the spectra (the difference in the computed wavelength for the lowest energy transition is 2 nm (0.01 eV in this wavelength range), considerably below the precision of the theoretical method used here). Bond Analysis. Results reported above let us assume that the main influence of metal cations is due to their charge. To challenge this hypothesis, we proceeded to a thorough examination of geometrical and electronic structures of the three complexes and to AIM analyses of the MOcatechol bonds. First, the comparison of the structure of the ligand between the three complexes was performed and some geometrical parameters are given in Table 5. These values show that esculetin structure is modified when it interacts with a cation, but that the dispersion between the values for the three complexes involving a cation with a global +2 charge is very low: less than 1 pm for bond lengths and about 2° for angles. The only general trend is the increase of angles and of metal−oxygen bonds with

Table 3. Wavelength and Energy of Transition Calculated between 350 and 425 nm for Each Hypothetical Structure, Discrepancy with the Experimental Values, and Oscillator Strengths transitions between 350 and 425 nm water molecules

cation

structure

AlIII

A B C D E F G H I experiment A A B

4 4 4 4 5 5 5 5 5

B D

4 2

E E F G G H H

3 5 3 3 5 3 5

I I experiment A A B C D F G H I experiment

3 5

CuII

PbII

2 4 2

2 3 3 2 2 3 3 3 4

λabs (nm, eV) 363 (3.41) 346 (3.58) 366 (3.39) 385 (3.22) 389 (3.18) 359 (3.45) 369 (3.36) 385 (3.22) 353 (3.51) 360 (3.45) 423 (2.93) 368 (3.37) 361 (3.44) 354 (3.51) 385 (3.22) 357 (3.48) 383 (3.24) 353 (3.51) 354 (3.50) 354 (3.50) 385 (3.22)

Δλ (nm, eV)

f

22 (0.19) 39 (0.36) 19 (0.17)

0.326 0.182 0.189

4 (0.04) 26 (0.23) 16 (0.14) 0 (0.00) 32 (0.29) 25 (0.23) 38 (0.29) 17 (0.15) 24 (0.22) 31 (0.29)

0.018 0.264 0.364 0.317 0.027 0.035 0.046 0.027 0.261 0.274

28 (0.26) 2 (0.02) 32 (0.29) 31 (0.28) 31 (0.28)

0.208 0.447 0.288 0.424 0.300

Figure 7. Computed transitions (blue) and experimental spectra (black) for the complexes of Cu(II) (a) and Pb(II) (b). 2651



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Figure 8. Structure of the copper(II)−esculetin (left) and lead(II)−esculetin (right) complexes.

laplacian of the electronic density at the BCP clearly characterize a closed-shell interaction between the metal cation and the catechol. By comparison of the charges and Hb values, the complexes can be ordered by their ionicity binding character (AlIII > PbII > CuII) in accordance to their hardness.64 In the case of lead(II) and copper(II), the negative value for the total electronic energy density shows a slight dative character of the metal-catechol bonds while the alumium(III)-catechol bond appears to be strictly ionic (Hb ≈ 0). Structure of the 2:1 Complex for Aluminum(III). After the study of the 1:1 complex for aluminum ion, we tried and shed some light on the nature of the 2:1 complex observed. Two major hypotheses are conceivable: (i) the coordination of the metal cation to the lactone moiety and (ii) the formation of aluminum dimers in solution that bind on the catechol moiety (Chart 2). These hypothetical structures were optimized as in the previous section, and their UV−visible spectra were simulated (see Supporting Information for the detail of the computed wavelengths). For case (i), due to the results obtained for the 1:1 complex study that show that the bidente binding on the lactone was not stable, the sole monodente binding was studied varying the protonation state of water molecules around the two aluminum ions. For these structures, the lowest-energy transition was computed between 396 and 423 nm (2.93 and 3.13 eV) (increasing more the number of hydroxides was not chemically reasonable). Even if these values are not so far from the experimental one (377 nm, 3.29 eV), we did not consider them as representative because of the variation compared to 1:1 complex (a small redshift is computed while a small blueshift is measured). For case (ii), we computed some structures corresponding to aluminum dimers linked by either two or three μ2 hydroxides. Depending on the protonation state of water molecules around the two aluminum ions, the lowest-energy transition was computed between 366 and 397 nm (3.12 and 3.39 eV). This values are in very good agreement with the measured one.

Table 4. Wavelength and Energy of Transition Calculated between 350 and 425 Nm for Hypothetical Structure D Varying the Coordination Sphere of Aluminum(III), Discrepancy with the Experimental Values, and Oscillator Strengths transitions between 350 and 425 nm cation AlIII

water molecules

coordinating hydroxide

4 3 3 experiment

0 1 (axial) 1 (equatorial)

λabs (nm, eV) 363 (3.41) 382 (3.25) 380 (3.26) 385 (3.22)

Δλ (nm, eV)

f

22 (0.19) 3 (0.03) 5 (0.04)

0.326 0.311 0.309

the ionic radius of the metal. On the contrary, a difference is observed when compared to the complex involving a cation with a +3 charge: C−O bond lengths are increased, certainly due do the lower possibility for oxygen electrons to be involved in those bonds. Second, the molecular orbitals involved in the transition located at 385 nm are mainly the frontier orbitals (HOMO → LUMO transition). These orbitals are depicted for the three actual complexes on figure 10. The differences between the three orbitals pairs are quite low. These orbitals are mainly localized on the ligand (this effect is particularly important in the case of aluminum, see AIM study below) and the main features of the HOMO and LUMO are the same in the three cases. The transition, similar to the one computed for the free ligand, implies mainly a charge transfer from the 6-hydroxyl group (with a small contribution from the 7-hydroxyl group) to a π* orbital localized on the rings. Last, some characteristic values obtained by the AIM study are reported in Table 6 for each of the 1:1 complexes. The difference between the MO6 and MO7 bond characteristics reflects the difference between bond lengths reported in Table 5. The quite low value of the density and highly positive value of the

Figure 9. Structure and electronic spectrum of the aluminum(III)−esculetin complex. 2652



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Table 5. Bond Lengths (Å), Angles (°) and Dihedral Angles (°) Relative to the Catechol Group of Esculetin, Doubly-Deprotonated Ligand, and Complexes (in Structure D) esculetin ligand (dep) AlIII complex AlIII(OH) complex CuII complex PbII complex

O6−C6

C6−C7

C7−O7

O6−M

O7−M

O6−C6−C7

C6−C7−O7

O6−C6−C7−O7

1.365 1.293 1.352 1.340 1.344 1.342

1.414 1.507 1.436 1.442 1.437 1.444

1.361 1.273 1.334 1.326 1.328 1.322

1.934 1.855 1.897 2.176

1.966 1.872 1.909 2.217

120.2 119.6 114.8 115.0 116.4 117.1

114.8 120.7 115.2 115.2 116.8 117.9

0.0 0.0 0.0 −0.4 0.2 −0.1

Figure 10. HOMO and LUMO for the three 1:1 complexes.

Table 6. Charge of Metal Ion and Catechol Oxygen Atoms, and Electronic Density, Laplacian of the Electronic Density, Kinetic Energy Density, and Total Electronic Energy Density at the Bond Critical Point (in Atomic Units) for the Three Actual Complexes

a

cation

AIM cation charge

bond

AIM O charge

ρ(BCP)

Δρ(BCP)

Gb

Hb

AlIII(OH)

2.55a

CuII

1.18

PbII

1.49

Al−O6 Al−O7 Cu−O6 Cu−O7 Pb−O6 Pb−O7

−1.32 −1.32 −1.12 −1.12 −1.22 −1.22

0.073 0.070 0.101 0.098 0.077 0.071

0.491 0.460 0.520 0.509 0.390 0.347

0.1222 0.1147 0.1432 0.1389 0.1041 0.0918

0.0005 0.0002 −0.0101 −0.0117 −0.0065 −0.0051

This value does not take into account the presence of a −1 charge in the coordination sphere of aluminum.

present in the medium. The nature of the 2:1 complex is then very different from the one we gave evidence for chrysazin complexation59 and quite surprising because the presence of dimers of aluminum ions in water is not their predominant form.65

Chart 2. Proposed Structure for the 2:1 Complex of Aluminum(III) (n = 2 or 3)

■

However, we were not able to discriminate between these hypotheses. Nevertheless, the important width of the absorption band compared to that of the 1:1 complex (see Figure 6) allows us to propose that several of these complexes are simultaneously

CONCLUSION

A detailed study of the complexation of esculetin with three different metal cations is presented in this work. Lead(II) and copper(II) lead to the formation of a 1:1 complex, while aluminum(III) allows the formation of two complexes of 2653



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stoichiometry 1:1 and 2:1 simultaneously. The association of UV−visible experiments with quantum chemistry computations showed that, in each case, the metal cation binds on the doubly deprotonated catechol to form the 1:1 complex. Moreover, the coordination sphere was determined that allows us to propose chemical formula for the complexes: two water molecules bind the metal cation in the complex with copper(II) (Cu(C9H4O4)(H2O)2) and with lead(II) (Pb(C9H4O4)(H2O)2), while three water molecules and one hydroxide anion bind aluminum(III) in its complex with esculetin (Al(C9H4O4)(H2O)3(HO)). A thorough investigation of the bonds involved in the formation of the complexes emphasizes the predominant role of the global charge of the metal cation (together with its coordination sphere) in its coordination with esculetin. In the case of aluminum, the nature of the 2:1 complex was not fully characterized. However, it is shown that, even in the presence of another potential binding site, the 2:1 complex involves an aluminum dimer on the catechol moiety. The values of the formation constants obtained for the three complexes show that the ligand presents a significant chelating power with regard to the three studied metals. Our previous work66 has shown that esculetin was a product of the photodegradation of caffeic acid (ubiquitous in the plant kingdom) but also of its metal complexes in solution. Thus it is quite conceivable that the formed esculetin and the metal ions released from caffeic acid complexes during photodegradation lead to the formation of new complexes.

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ASSOCIATED CONTENT

S Supporting Information *

The full reference 39, the study of the solvatation, the wavelengths computed for the 2:1 complexes involving aluminum(III), and the Cartesian coordinates for esculetin and the three actual complexes are given in Supporting Information. This information is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Tel.: +33 3 20 43 49 14. E-mail address: annaig.le-person@ univ-lille1.fr. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors are very grateful to Jérémy Gaillard for spectra measurements. Furthermore, this work was granted access to the HPC resources of CINES under the allocations 2013086933 and 2014086933 made by GENCI (Grand Equipement National de Calcul Intensif). We also thank the CRI (Centre de Ressources Informatiques) of the University Lille 1 for providing computing time for part of the theoretical calculations.

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REFERENCES

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The Complexation of AlIII, PbII, and CuII Metal Ions by Esculetin: A

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