Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

Contents lists available at ScienceDirect

Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy journal homepage: www.elsevier.com/locate/saa

Experimental and theoretical approach of photophysical properties of lanthanum(III) and erbium(III) complexes of tris(methoxymethyl)-5-oxine podant Rifat Akbar a, Minati Baral b, B.K. Kanungo a,⇑ a b

Department of Chemistry, Sant Longowal Institute of Engineering & Technology, Longowal, Punjab 148106, India Department of Chemistry, National Institute of Technology Kurukshetra, Haryana 136119, India

h i g h l i g h t s

g r a p h i c a l a b s t r a c t

 The tripodal ligand forms

thermodynamically stable LaIII and ErIII complexes with high log b and pLn values. 3+  [La(LH3)] and [Er(L)] were found at physiological pH, presents interesting photophysical properties.  Nonemitting zwitterionic forms of the ligand quench green fluorescence formed in acidic as well as basic medium.  These compounds render the OFF– ON–OFF type of pH-dependent fluorescent sensors.  Sparkle model clearly attests efficacy for in-silico design of lanthanide– organic frameworks.

a r t i c l e

i n f o

Article history: Received 22 January 2014 Received in revised form 12 March 2014 Accepted 21 March 2014 Available online 1 April 2014 Keywords: Formation constants Lanthanum Erbium Fluorescence OFF–ON–OFF sensors Sparkle

a b s t r a c t With the aim of evaluating the coordination behavior of a novel polydentate tripodal ligand, 5-[[3-[(8hydroxy-5-quinolyl)methoxy]-2-[(8-hydroxy-5-quinolyl)methoxymethyl]-2-methyl propoxy]methyl] quinolin-8-ol (TMOM5OX), towards La(III) and Er(III) metal ions, the detailed investigations of photophysical properties by theoritical and experimental (potentiometric, UV–visible and fluorescence spectrophotometry) methods were carried out. TMOM5OX has been found to form protonated complex [Ln(H4L)]4+ (Ln = La or Er) below pH 3.8, which consecutively deprotonates through one-proton processes with rise of pH. The formation constants (log b) of neutral complexes have been determined to be 36.42 (LaL) and 35.76, 37.62 (for ErL and ErL2, respectively) and the pLn (pLn = log[Ln3+]) values of 24.6 and 27.1 for La(III) and Er(III) ions, respectively, calculated at pH 7.4, indicating TMOM5OX is a good lanthanide synthetic chelator. The absorption spectroscopy of these complexes show marked spectral variations due to characteristic lanthanide transitions, which support the use of TMOM5OX as a sensitive optical pH based sensor to detect Ln(III) metal ions in biological systems. In addition, these complexes have also been shown to exhibit strong green fluorescence allowing simultaneous sensing within the visible region under physiological pH in competitive medium for both La(III) and Er(III) ions. The intense fluorescence from these compounds were revealed to intermittently get quenched under acidic and basic conditions due to the photoinduced intramolecular electron transfer from excited 8-hydroxyquinoline (8-HQ) moiety to metal ion, just an opposite process. This renders these compounds the OFF–ON–OFF type of pH-dependent fluorescent sensors. The complexes coordination geometries were optimized using the

⇑ Corresponding author. Tel.: +91 1672 253166 (O), +91 1672 253167 (R); fax: +91 1672 280057. E-mail address: [email protected] (B.K. Kanungo). http://dx.doi.org/10.1016/j.saa.2014.03.045 1386-1425/Ó 2014 Elsevier B.V. All rights reserved.

366

R. Akbar et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

sparkle/PM6 model and the theoretical spectrophotometric studies were carried out in order to validate the experimental findings, based on ZINDO/S methodology at configuration interaction with single excitations (CIS) level. These results clearly attest for the efficacy of the theoretical models employed in all calculations and create new interesting possibilities for the design in-silico of novel and highly efficient lanthanide–organic frameworks. Ó 2014 Elsevier B.V. All rights reserved.

Introduction In recent years, the chemistry of lanthanides has attracted a great deal of interest owing to their increasing use in chemical, biomedical, industrial, analytical [1,2], and other fields, especially as fluorescent materials and as ideal probes in the studies of biological systems [3,4]. Lanthanide ions with their characteristic magnetic and spectral properties are the potential probes for the study of natural system. Their coordination complexes with organic ligands have long attracted fundamental and applied research interest owing to the possibility of sensitizing rare-earth radiative transitions at visible and near-infrared wavelengths. The promise is to achieve sensitization with high efficiency by resorting advantageous light harvesting properties of the organic part [5], that is, by designing ligands to absorb light with high efficiency in selected spectral regions. The chemistry and spectroscopy of lanthanide ions differ considerably from d-shell transition-metal ions, in that the shielding of the 4f orbitals by the filled 5p66s2 subshells results in special optical features of lanthanide ions with characteristic narrow line-like emissions of optical pure colors that are nearly unaffected by the ligand field [6]. Despite the optical purity and easily recognized emissions of lanthanide ions, the spin- and parity-forbidden nature of the f–f transitions renders direct photoexcitation of lanthanide ions disfavored, which makes their absorption coefficients usually smaller than 10 M1 cm1. One of the most useful strategies that has been employed to circumvent this predicament, is the introduction of organic ligands function as chromophores to absorb light and transfer this energy to the excited states of the central lanthanide ions, which is the so-called ‘‘antenna effect’’, plays a notable role in the improvement of emission efficiency [7]. In order to favor energy transfer from antenna to lanthanide(III) ions effectively, the excited state of the ligand should be higher than the lowest excited state of the lanthanide(III) ion within an appropriate range [8]. This sensitization process consists, initially, of the absorption of light by the ligand in the UV region, followed by a nonradiatively energy generally transferred from the triplet state (T1) of the ligand to an excited level of the Ln(III) ion [9] and finally, the excited Ln ion decays to the ground state via photon emission in the visible region. The excitation in the ligands is much more efficient than that directly in the Ln(III) ion, because absorption coefficients of the ligands are several orders of magnitude larger than the intrinsically low molar absorption coefficients (typically 1–10 M1 cm1) of Ln ions [10]. Therefore, the design of Ln(III) complexes with efficient photonic properties has become an important research goal, working with many different classes of ligands. In fact, organic ligands suitable for sensitizing visible and NIR (near-infrared) emission from lanthanide(III) ions by excitation with visible light have been less explored [11]. As a consequence, a number of chromophoric antenna ligands, especially the b-diketonate [12], carboxylate ligands [13], and 8-hydroxyquinoline [14], which have received the most attention have been developed in an effort to achieve brighter lanthanide luminescence. 8-Hydroxyquinoline and its derivatives are versatile coordination ligands towards a wide range of metal ions, including lanthanide(III) ions. These ligands possess useful photophysical

properties with a triplet state located around 17,100 cm1 (585 nm), making it suitable for the sensitization of Ln(III) luminescence. In fact, lanthanide(III) complexes of 8-hydroxyquinolinates have been considered as one of the most promising materials for the design of electroluminescent devices [15,16]. However, due to the insolubility of the parent lanthanoid-hydroxyquinolines, there is a major disadvantage to obtain the crystalline products for structure determination [17]. It is important to emphasize that emission by Ln(III) complexes is quenched when water or other solvent molecules, which have high-energy vibrational modes, are coordinated to the metal ion [18]. Hence, the design of polydentate ligands to near-saturatively chelate the lanthanide center, thus excluding quenching species such as water, as well as incorporating a sensitizing antenna chromophore within the same molecular framework, has been established as a popular approach to improve the efficiency of lanthanide emission [19]. Most of the studies on quinoline-based chromophores suitable for the sensitisation of lanthanide ions have concentrated on the use of 8-hydroxyquinoline itself as a directlycoordinating bidentate ligand [20,21] or incorporating quinoline moieties into podand structures for multidentate coordination [22]. Quinoline appended macrocycles that bind lanthanides have been developed as sensors for pH [23] and molecular logic gates [24]. Blurring the boundaries between the approaches is an elegant strategy devised by Mazzanti and co-workers, using direct coordination of the lanthanide center by quinoline in league with a macrocyclic binding site to both increase the stability and achieve full coordinative saturation at the lanthanide center [25]. The design and realisation of versatile ligand systems that are able to efficiently sensitize both visible and NIR emitting lanthanide(III) ions and are also able to chelate magnetically active centers whilst exhibiting good stability in polar protic solvents is a challenge. Therefore, it is an essential requirement to hunt for emissive lanthanide complexes that can be sensitized efficiently for application. With trivalent lanthanide ions possessing high coordination flexibility and lack of preferential geometries, the structure–property relationships of lanthanide complexes have gradually caught the researcher’s attentions [26] and prefer high coordination numbers, with 8 or 9 being the most common ones [27]. 8-Hydroxyquinolines being bidentate monoanionic ligands and hence cannot saturate coordination sphere of a lanthanide ion upon formation of a charge-neutral tris-complex. Moreover, in these ligands the hydroxyl group tends to act as a bridge in lanthanide complexes. As a result, the complexation of lanthanide ions with bidentate 8hydroxyquinolines often gives mixtures of mono- and poly-nuclear complexes [28]. To make the outcome of complexation more predictable, it is advantageous to employ polydentate 8-hydroxyquinolines for coordination with lanthanides [29]. Moreover, theoretical tools can be helpful in the investigation of photophysical properties occurring in lanthanide systems. In the field of Ln(III) coordination compounds, the semiempirical sparkle model can be used as an important tool in the design of new light conversion molecular devices (LCMDs) [30], and in the systematic quantum chemical studies on the ordering of the excited electronic states, and a correlation with the structural properties of antenna, could provide a priori screening of efficient chromophores. Since

R. Akbar et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

sparkle model allows the treatment of a great number of lanthanide complexes efficiently and in a relatively short time than that of HF or DFT methods, the semiempirical approach to the quantum chemical calculation of lanthanide complexes was introduced in 1994 [31]. It replaces the lanthanide ions by a Coulombic charge of +3e, superimposed to a repulsive exponential potential of the form exp(ar), which was introduced to mimic the effect of the size of the ion. Thus, the sparkle model assumes that the angular effects of the f-orbitals are negligible and does not take them into account, being, thus, a spherically symmetric model. This model has been extensively applied to the calculation of ground-state geometry of these molecules, which is necessary to predict spectroscopic properties, singlet and triplet energy positions, and electronic spectra of Ln(III) complexes [32]. With these quantities, it is also possible to build rate equations that involve the energy transfer mechanism to calculate the quantum yields for these complexes [33]. To understand the activities of the lanthanide complexes, the information about the complexation of these metals with a polydentate ligand is of considerable importance and can be predicted on the basis of their formation constants in solution. Therefore, it is necessary to have a detailed knowledge about the formation constants, spectro-chemical behavior, solution equilibria and thermodynamics involved in these reactions. Keeping the above facts in view, with the recently developed encapsulating polydentate tripodal ligand TMOM5OX (Fig. 1) based on 8-hydroxyquinoline which form water-stable Fe+3 and Al+3complexes, systematic study has been undertaken on the interactions of the lanthanide (La3+ and Er3+) ions with TMOM5OX, employing potentiometric, spectrophotometric titration techniques and theoretical studies. We report herein the formation constants, detailed photophysical (complexation and sensitisation) and theoretical (sparkle/PM6 and ZINDO/S) properties along with the other thermodynamic parameters of the podant, TMOM5OX, with the trivalent La and Er ions.

Experimental section Materials and methods All reagents were purchased at the highest commercial quality from Sigma–Aldrich of ultrapure grade and used without further purification unless otherwise stated. For solubility reasons the ligand was converted to its hydrochloride salt, and potentiometric, spectrophotometric and luminescence studies were carried out in water, purified by a Millipore Milli-Q water purification system. Ionic strength was adjusted with 0.1 M KCl. All the stock solutions were prepared by weighing appropriate amount using GR-202 electronic balance (precision 0.01 mg) in Millipore grade deionized

O

O

HO

O N

OH N

HO

367

water. The exact concentration of KOH (0.1 M) and HCl (0.1 M) were determined potentiometrically by acid-base tritration using 0.1 M solutions of succinic acid, potassium hydrogen phthalate (KHP) and oxalic acid as primary standards. Stock solutions of lanthanides were prepared just before use in deionized water from the corresponding nitrate salts, Ln(NO3)3.nH2O (Ln = La and Er; n = 5). All the solutions were prepared with Millipore grade deionized water immediately before use, which was deoxygenated and flushed continuously with Ar (U grade) to exclude CO2 and O2. All measurements were carried out at 25.0 ± 0.2 °C maintained with the help of Julabo F-25 thermostat. Potentiometric measurements Potentiometric titrations were carried out for the determination of the formation constants of the metal complexes. A double wall glass jacketed titration cell connected to a constant temperature circulatory bath was used to maintain temperature at 25 ± 1 °C. The pH measurements were performed using Thermo Scientific ORION STAR-A211 pH meter equipped with combined Ross Ultra pH/ATC glass electrode and the observed pH was measured as log[H+]. The electrode was duly calibrated to read pH according to classical method [34] and titrations were employed with Borosilicate glass burette (1 ml, least count 0.01 ml). A standard KOH (0.095 M) solution was used to titrate a standard hydrochloric acid (0.1045 M) solution and the pH-meter readings were converted to hydrogen ion concentration by calculated hydrogen ion concentrations (pKw = 13.77). The hydrochloride salt of ligand (5 ml, 1  104 M) and mixtures of 1:1 L: Ln+3 (1  104 M) were titrated with standardized 0.095 M KOH for determination of protonation and formation constants, respectively. The solutions were acidified to a pH of 1.98 with standardized HCl (5 ml, 0.1045 M) and the ionic strength was fixed at 0.1 M with KCl. Final concentration of ligand (1  105 M) and metals (1  105 M) were maintained for the different titrations. The titration data {pH range 1.9–12.5, 162 points} were refined by the nonlinear least-squares refinement program HYPERQUAD [35] to determine equilibrium constants (protonation and complexation), while the distribution of species were plotted with the program HYSS 2009 [36]. Spectrophotometric measurements Electronic absorption spectra were recorded on an Agilent-8453 Diode Array Spectrophotometer using 1.0 cm path length Hellma quartz cell and externally connected to a HEWLETT PACKARD Vectra P3327G computer. All titrations were performed in a thermostated (25.0 ± 0.1 °C) glass-jacked vessel. Spectrophotometric titrations for Ln(III) metal complexes were carried out by varying the pH in aqueous medium, in a typical experiment, for monitoring the UV–visible spectra in the titration of stiochiometric 1:1 quantities of TMOM5OX with Ln(NO3)nH2O{5 ml (5  105 M)} (Ln = La and Er; n = 5), acidified with approximately 5 ml (0.1045 M) HCl, ionic strength maintained at 0.1 M with KCl were monitored as a function of pH over the range 1.8–12.3 by titrating with freshly prepared standardized (0.095 M) NaOH solution. After each addition of NaOH and adjustment of pH, it was ensured that complete equilibrium was attained (ca. 5 min) before measurement of the pH and recording of the spectrum. The formation constants from the spectral data were calculated by using a non-linear leastsquare fitting program, HYPSPEC [37]. Fluorescence measurements

N

Fig. 1. 5-[[3-[(8-hydroxy-5-quinolyl)methoxy]-2-[(8-hydroxy-5-quinolyl)methoxymethyl]-2-methyl propoxy]methyl]quinolin-8-ol (TMOM5OX).

Fluorescence measurements were carried out on a Perkin Elmer LS-55 luminescence Spectrophotometer equipped with quartz cuvettes with a path length of 10 mm at 25.0 ± 0.1 °C. Fluorescence

368

R. Akbar et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

spectra were registered with excitation at 332 nm and 393 nm for La and Er complexes, respectively and all excitation and emission slit widths at 5.0 nm unless otherwise indicated. In order to allow correlation of emission intensities, adjustment for instrumental response, inner filter effects and phototube sensitivity were implemented. The linearity of the fluorescence emission vs. concentration was checked in the concentration range used (105–106 M). A correction for wavelength response of the system was performed when necessary. Fluorescence data was also recorded from the same experimental method as used for absorption spectra measurement. For the formation of the luminescent metal complexes of TMOM5OX, titrations were ascertained, as after each addition of base, the pH of the solution was measured using combined Ross Ultra pH/ATC glass electrode and emission spectra was recorded. The formation constants from the spectral data were calculated by using a non-linear least-square fitting program, HYPSPEC [37]. Quantum yields were determined using the relationship:

U ¼ Uref

ðI=AÞ g ðIref =Aref Þ gref

!2

where U is the radiative quantum yield of the sample, Uref is the known quantum yield of quinine sulfate in 1 M aqueous H2SO4 (=0.546) [38], A is the absorbance at the excitation wavelength, I is the integrated emission, and g is the refractive index of the solvent, which is assumed to be the same for the solutions of sample and reference. Theoretical calculations Theoretical basis All computations were carried out on a Pentium IV 3.2 GHz machine with a Linux operating system. Starting structures for optimization were manually drawn using MDL ISIS Draw 2.5 Standalone. The initial geometry optimization of the TMOM5OX and its Ln(III) complexes leading to minimum strain energy were achieved through molecular mechanics calculation using MM force field by Gabedit version 2.4.6 [39]. The periodically search to global minimum energy conformer of the ligand and its metal complex were achieved using molecular dynamics simulation upto 2000 K followed by the geometrical optimization calculations. A bath relaxation time of 0.1 ps and a step size of 0.001 ps were used for dynamic simulation. Then, the minimized structure was further reoptimized through semi-empirical method by applying PM3 self consistent fields (SCF) method, at the Restricted Hartree–Fock (RHF) level using MOPAC2012 [40]. The geometry optimizations were obtained by the application of the Polak–Ribiere algorithm with convergence limit of 0.0001 kcal/mol and RMS gradient of 0.001 kcal/mol. The above optimized molecular structures were then used for sparkle mode calculations. Parametric method number 6, PM6, is one of the latest in a series of semiempirical methods which encompass MNDO, AM1, PM3, and RM1. The accuracy of PM6 in predicting enthalpies of formation, yielding an unsigned mean error of 4.4 kcal mol1 for a representative set of 1373 compounds, exceeds those of Hartree–Fock (7.4 kcal mol1) or B3LYP DFT (5.2 kcal mol1) methods [41]. Recently reported [41] that Sparkle/PM6 is an accurate and statistically valid tool for the prediction of the geometrical features of lanthanide coordination polyhedral and, by design, is expected to perform best with ligands with nitrogen or oxygen as coordinating atoms present in the vast majority of all coordination compounds of the trivalent rare earth metals. Semiempirical (sparkle model) methodologies have been applied [42] to study the spectrophotometric properties of the systems reported in 2007 by DFT and TD-DFT methods [43]. The authors have shown that the semiempirical methods can be used to predict the

spectroscopic properties of europium cryptates with accuracy comparable to that shown by DFT and TD-DFT results and suggested that experimental spectroscopic data are better reproduced in ZINDO calculations when the lanthanide ion is represented by a point charge, also triplet energies calculated by semiempirical methods have errors similar to those obtained by TD-DFT methodology but are hundreds of times faster [43]. It has also been suggested [44–46] that the use of quantum chemical methodology other than semiempirical Sparkle model such as Hartree–Fock (HF) or density functional theory (DFT) using an effective core potential (ECP) to treat the Ln(III) ions is unfeasible owing to the high computational effort needed. Co-ordination scan Molecular mechanics has been successfully utilized to determine the relationship between ligand selectivity and metal ion size. A highly selective ligand for a particular metal would possess a steep curve with a minima close to the ionic radius, conversely a shallow curve would suggest that the ligand is non selective in its metal binding. A related technique is the ‘‘coordination scan’’, wherein similar curves are generated by minimizing complexes with various numbers of water molecules coordinated to the metal ion while changing the M–L bond lengths [47]. The preferred coordination number 9 and 12 were observed for La and Er complexes, respectively, calculated by the coordination scan technique with the stepwise addition of three or more water molecules and calculating strain energy of the molecules using SYBYL. It calculates water as having strain energy of 0.00 kcal/mol; thus the waters added to the complex add no energy other than steric interactions with the ligand. The ground state geometries calculation Full geometry optimizations of all the molecular systems were carried out without symmetry constraint. Frequency calculations and dispersion corrections were performed, and the minima on the potential-energy surfaces of the reported structures were characterized by the absence of negative eigen values in the Hessian matrix. The ground state geometries of [LaTMOM5OX(H2O)3], [ErTMOM5OX(H2O)3] and [Er(TMOM5OX)2] were calculated with the sparkle/PM6 [41] model implemented in the MOPAC2012 software [40] package. The key words used were SPARKLE, PM6, PRECISE, FORCE, BFGS, GNORM = 0.25, SCFCRT = 1.D-10 (to increase the SCF convergence criterion) and XYZ (for Cartesian coordinates). Excited states energies and absorption spectra Optimized geometries from the sparkle/PM6 models have been successfully applied in the prediction of spectroscopic properties such as singlet and triplet energy levels [48] and UV–visible absorption spectra [49]. We predicted the singlet and triplet excited states of all the calculated ground state geometries using the configuration interaction with single excitations (CIS) based on the Zerner’s intermediate neglect of differential overlap/spectroscopic (ZINDO/S) methodology [50], using a point charge of +3e to represent the trivalent lanthanide ion. The CIS space was gradually increased until there were no further meaningful changes in the calculated triplet energies and absorption spectra. A Lorentzian line shape was fitted to the calculated singlet transitions, together with the relative intensities obtained from oscillator strengths and all the simulated spectra had a half-height bandwidth of 25 nm. Results and discussion Complexation: interaction with trivalent lanthanide ions The nature and stability of the species forming between the Ln(III) and Er(III) ions and TMOM5OX have been probed at variable

R. Akbar et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

pH and in equimolar solutions of ligand and metal ions by the combined use of potentiometric (pH range 1.88–12.53), spectrophotometric (pH range 1.79–12.29) and luminescence titrations (pH range 1.84–12.7). These ions were chosen as standard example because La is the very first ion and Er lies in the second half part of the Ln(III) series, and in view of the podand nature of the ligand investigated which behave as predisposed receptor. Not much difference in the stability of the resulting podates was observed along the lanthanide series [51,52]. The potentiometric titration curves of a 1:1 solutions of La(III) and Er(III) ions and ligand TMOM5OX (Fig. 2) show a pH variation indicating the release of protons upon Ln(III) coordination. The curves 2b (for La-TMOM5OX) showed inflection at a = 2, 3, 4 and 2c (for Er-TMOM5OX) a = 1, 3, 4, from the curve for the protonation of ligand alone (a is the number of moles of base per mole of metal). This indicates the formation of species in which different number of protons has been displaced from TMOM5OX by the lanthanide ions. Furthermore, the formation of hydroxo-complexes was also detected at pH above about 10.0 for La-TMOM5OX and about 9.8 for Er-TMOM5OX systems. Several models were tested to refine the potentiometric data by Hyperquad but the best one involves four complexed species for both metal ions in equilibrium and their formation constants are reported in Table 1. The formation of the complex species and their cumulative formation constants, b11n, are defined by Eqs. (1) and (2), respectively. ðn1Þþ þ Ln3þ aq þ L þ nH ½LnðHn LÞ

b11n ¼ ½LnðHn LÞðn1Þþ =½Ln3þ ½L½Hþ 

ð1Þ n

ð2Þ

The spectrophotometric titration were then carried out (Figs. 3 and 4) to correlate the deprotonation of La(III) and Er(III) complexes and the data obtained in two different pH ranges were fitted separately. Between pH 1.79 and 7.2, the best fit of absorbance data corresponds to the presence of only one metal complex in case of La(III) and two metal complexes, in case of Er(III), mainly because several protonated forms of the ligand coexist in this pH range. However, combining these data with the data obtained from the potentiometric titrations allowed us to fit the three protondependent equilibria described by Eqs. (3)–(5). 3þ 4þ Ln3þ þ 2Hþ aq ðH6 LÞ ¡½LnðH4 LÞ

ð3Þ

Fig. 2. Potentiometric titration curves: (a) 5  105 M TMOM5OX, (b) [TMOM5OX]/ [La(III)] = 1/1, 5  105 M, (c) [TMOM5OX]/[Er(III)] = 1/1, 5  105 M. Solvent H2O, I = 0.1 M (KCl), T = 25(2) °C, ‘a’ moles of base added per mole of TMOM5OX-Ln present. Symbols and solid lines represent the experimental and calculated data, respectively.

369

½LnðH4 LÞ4þ ¡½LnðH3 LÞ3þ þ 1Hþ

ð4Þ

½LnðH3 LÞ3þ ¡½LnðH2 LÞ2þ þ 1Hþ

ð5Þ

Above pH 7.2, both potentiometric and spectrophotometric data point to the presence of three other species in case of La(III) and four species, in case of Er(III). Their formation constants defined by Eqs. (6)–(8) were refined by considering two successive deprotonations and the formation of one hydroxo complex at pH higher than 10.0 and 9.8; they are listed in Table 1.

½LnðH2 LÞ2þ ¡½LnðHLÞþ þ 1Hþ

ð6Þ

½LnðHLn Þþ ¡½ðLnLn Þ þ 1Hþ

ð7Þ

½ðLnLÞ¡½LnLðOHÞ þ 1Hþ

ð8Þ

In contemplation to further support the number and nature of species obtained by both potentiometric and spectrophotometric methods and the corresponding formation constants discussed above, luminescence titrations were also carried out for 1:1 solutions of La(III) and Er(III) ions with ligand from 1.8–12.7 and 1.9– 12.6 pH range, respectively. The typical examples of experimental data for the luminescence titrations of the La and Er complexes are depicted in Fig. 5(a) and (b). For La-TMOM5OX system, the best refinement of the luminescence-pH data by the Hypspec program corresponds to the presence of four species, two protonated complex species that is, [La(H3L)]3+ [log b = 43.84(4)] and [La(HL)]+ [log b = 32.19(4)] in concurrence with the potentiometric and spectrophometric data, one neutral species [(LaL)] [log b = 36.49(5)] and one hydroxo-complex [LaL(OH)] [log b = 7.84(3)] were also formed that favourably agree the formation of these species by both potentiometric and spectrophotometric method (the number in parentheses corresponds to the standard deviation in last significant digit). In case of Er-TMOM5OX system in the pH range of 1.7– 6.8, the best fit of the luminescence-pH data returned considering formation of three protonated complexes, Er(H2L)]2+ [log b = 37.81(5)], is in good agreement with the spectrophometric value, while [Er(H3L)]3+ and [Er(HL)]+ species {log b = 45.78(1) and 33.54(3), respectively} which were also in good agreement with both potentiometric and spectrophotometric titrations. Above pH 6.85, the luminescence data impart to the presence of neutral (ErL) [log b = 35.81(5)], (ErL2) [log b = 37.32(3)], and hydroxo [ErL(OH)] [log b = 7.98(3)] species were found in consistent to both potentiometric and spectrophotometric data, which corresponds to the presence of [(ErL2)] [log b = 37.85(4) and 37.39(5), respectively] complex along with (ErL) and [ErL(OH)]. The formation constants {(log b, obtained by all the three techniques, potentiometrically as well as spectrophotometrically (UV–visible and fluorescence techniques)} of these complexes, which were defined above by Eqs. (1)–(8), together with the average log b values are reported in Table 1. On basis of foregoing discussion it is apparent that all species could not be detected by a single method, but combination of potentiometry, UV–visible spectrophotometry and fluorometry (spectro-electrometric), six and seven different species and their formation constants were evaluated for La(III) and Er(III) metal ions, respectively. The corresponding distribution diagrams are shown in Fig. 6(a) and (b). Despite the complexity of the system, LaL is the major species (85.5% present in the range of 7.8–11.5 pH. The other species present at very acidic pH is [La(H4L)]4+ (100% at pH 6 3.4); upon an increase in the pH, successive deprotonation of major species leads to the formation of [La(H3L)]3+, [La(H2L)]2+ and [La(HL)]+ which become the major species at pH 6.0 (98.8%), 8.1(20%) and 8.9(50.2%), respectively. Similarly in case of Er, [Er(H4L)]4+ complex predominates at pH 3.6; an increase in pH leads to diminish this species, and after pH 4.9 new [La(H3L)]3+

370

R. Akbar et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

Table 1 Equilibrium constantsa (log b11n), absorption, and emission characteristics of La(III) and Er(III) complexes of TMOM5OX and corresponding pLn values. Complexes L = TMOM5OX

Species

Log b11nb

Log b11nc

Log b11nd

Average log b11n

kmax (nm)

e (M1 cm1)

kem (nm)

U ePL

pM3+f

(LaL)

[La (H4L)]4+ [La (H3L)]3+ [La (H2L)]2+ [La (HL)]+

54.17(6)

54.21(4)

_

54.19

205, 247, 334

10,000, 45,600, 5150

0.013

24.6

43.94(5)

43.89(5)

43.84(4)

43.91

208, 247, 340

16,000, 39,000, 4000

_

37.41(3)

_

37.41

210, 248, 340

23,000, 30,000, 2000

32.03(3)

32.12(2)

32.19(4)

32.11

215, 250

35,000, 17,000

[(LaL)]

36.41(4)

36.37(3)

36.49(5)

36.42

285, 372

18,000, 1600

7.92(4)

7.89(5)

7.84(3)

07.88

285, 379

29,600, 2900

375, 418 375, 424 375, 515 375, 520 375, 425 375

[Er (H4L)]4+ [Er (H3L)]3+ [Er (H2L)]2+ [Er (HL)]+

_

57.08(4)

_

57.08

210, 248, 330

10,000, 45,000, 5000

440

0.006

45.72(2)

45.81(4)

45.78(1)

45.77

248, 265, 330

38,000, 28,400, 4000

445

0.007

_

37.79(3)

37.81(5)

37.80

215, 248, 265, 392

25,640, 5,140,

515

0.011

33.49(4)

33.52(5)

33.54(3)

33.51

217, 250, 267, 390

36,000, 18,050,

418

0.036

[(Er L)]

35.73(6)

35.76(4)

35.81(5)

35.76

278, 295, 370, 425

29,600, 4000,

520

0.053

[(Er L2)]

37.85(4)

37.39(5)

37.32(3)

37.62

25,600, 4000,

447

0.084

[Er L(OH)]

8.03(5)

8.08(4)

7.98(3)

08.03

275, 290, 320, 373, 420 275, 290, 320

30,000, 3000 25,000, 1600 33,000, 3600 26,000, 2400 22,400,

21,000, 6400

441

0.008

[LaL(OH)] (ErL)

a b c d e f



0.009 0.014 0.035 0.048 0.029 27.1

Numbers in parentheses represent the standard deviation in the last significant digit. In water (I = 0.1 M KCl, T = 25.0 °C). Potentiometric method. UV–visible spectrophotometric method. Luminescence spectrophotometric method. 0.1 M solution of quinine sulfate in 0.5 M H2SO4 as standard (U = 0.546). pLn3+ = log[Ln3+] calculated for [Ln]tot = 106 M, [L]tot = 105 M, and pH = 7.4.

Fig. 3. UV–vis absorption spectra of 1:1 solution of La(III) and TMOM5OX as a function of pH: (a) pH = 1.79–7.2; (b) pH = 8.21–12.29; [TMOM5OX] = [La(III)]tot = 5  105 M, Solvent: H2O, I = 0.1 M(KCl), T = 25.0(2) °C. (c) Inset showing spectral peak around 205 nm–215 nm (at concentration = 5  106 M).

species starts forming with maximum concentration at pH 6.0. Due to further deprotonation, [La(H2L)]2+ was found at pH 5.9 (37%) along with [La(HL)]+ which exists 32% at physiological pH and the neutral complex (ErL) was found during the pH range of 5.8– 9.1. In contrarary to La complex, Er forms, another neutral species (ErL2) with TMOM5OX, which was predominantly present during pH 6.2–12.6 (88%). Further scrutiny of the distribution curves shows a consistent decrease of the free La(III) concentration from pH 2–5.9 followed by slight increase up around pH 8–11 but concentration of Er(III) show steep decrease at pH 3.1 which

completely disappears at pH < 6. This suggests a change in the coordination around the metal ion, as observed for the ferric complex with O-Trensox [53] and also for the Eu(III) complex of Tsox [36]. The complexation efficiency of a metal ion Maq by a given ligand is usually assessed by the pM(=log[Maq]) value calculated at physiological pH for total concentrations of [M]tot = 106 M and [L]tot = 105 M [54]. The corresponding plots of pLa(III) and pEr(III) vs. pH for the TMOM5OX over the pH range of 2.5–14 are presented in Fig. 7. For TMOM5OX, this translates into pLn values of

R. Akbar et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

371

Fig. 4. UV–vis absorption spectra of 1:1 solution of Er(III) and TMOM5OX as a function of pH: (a) pH = 1.71-6.83; (b) pH = 6.85–12.54; [TMOM5OX] = [Tb(III)]tot = 5  105 M, Solvent: H2O, I = 0.1 M(KCl), T = 25.0(2) °C. (c) Inset showing spectral peak around 210 nm –215 nm (at concentration = 5  106 M).

Fig. 5. Dependence of the emission spectra of 1:1 aqueous solution of Ln:TMOM5OX (5.0  105 M) on the change of the pH value from acidic to basic range, T = 25.0(3) °C (excitation and emission slit widths of 5.0 nm) (a) for La[TMOM5OX], with an excitation wavelength of 332 nm (1.81–12.3 pH) and (b) for Er[TMOM5OX], with an excitation wavelength of 393 nm (1.84–12.67 pH).

24.6(La) and 27.1(Er) compared to pEu = 15.6 for Tsox [36] or 19.6 for [Eu(dtpa)]2 (dtpa = diethylenetriaminopentaacetic acid), as computed from known formation constants [55]. The chelate effect of the podant with respect to the 8-hydroxyquinoline building block is impressive and the complexes based on TMOM5OX appear to be sufficiently stable in water for potential in vivo applications. In order to verify the validity and assignment made for the experimental formation constants, aqueous-phase free energy for the entire representative species for both the metal ions were calculated through semi-empirical sparkle/PM6 COSMO quantum mechanical approach. For the metal complex MLH(n1), the log K defined as the negative logarithm of the formation constant of the reaction Lnaq + LHn = [Ln(LH)(n1)]+ + H3O+, is given by the thermodynamics relation DG° = RT(log K/2.303). The change in free energy in aqueous-phase is obtained according to the following equation:

DGaq;½LnðLHÞn  ¼ ðGaq;½LnðLHÞðn1Þ  þ Gaq;H3 Oþ Þ  ðGaq;Ln þ Gaq;LHn þ Gaq;H2 O Þ ð9Þ

Since the metal ions compete with the protons for the coordination sites, the complexation of metal ion with a ligand occurs upon release of hydrogens and greater acidity belongs to the group in which its hydrogen release is easier. A clear increase in DG° was observed for La[TMOM5OX] and Er[TMOM5OX] (Fig. 8a and b) as the deprotonation takes place from fully protonated free ligand, LH3+ 6 , upon complexation with lanthanide metal ions. The validity of this calculated DG° for the different species of La and Er complexes of TMOM5OX were compared with the calculated experimental log K, which resulted an acceptable correlation with R2 = 0.9968 and 0.996, respectively. 3.2. Photophysical properties of La and Er complexes of TMOM5OX 3.2.1. UV–visible spectrophotometric studies The large spectral changes (Figs. 3 and 4) observed in the spectrophotometric titration of the lanthanide TMOM5OX complexes (Ln = La and Er) in acidic medium and at physiological pH is indicative of a change in the coordination of metals with the ligand. At

372

R. Akbar et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

Fig. 6. Species distribution curves of (a) La[TMOM5OX] and (b) Er[TMOM5OX] containing species, computed from the formation constants given in Table 1. Calculated for [TMOM5OX]tot = [Ln(III)]tot = 105 M, Solvent: H2O, I = 0.1 M (KCl), T = 25.0(2) °C.

Fig. 7. Plot of pLn vs. pH for TMOM5OX, pLn = log[Ln(III)], calculated for [Ln(III)] = 106 M and [L] = 105 M. (&) pLa and ( ) pEr.

pH < 2.7, the electronic spectrum of the TMOM5OX with both La and Er ions is symbolized by a strong absorption band at higher energies 247 nm, (e = 45,600 M1 cm1) and 248 nm, (e = 45,000 M1 cm1) assigned for p ? p transitions and a broad band at 334 nm (e = 5,150 M1 cm1) and 330 nm (e = 5000 M1 cm1) assigned for n ? p transitions, respectively. Upon rising pH above 2.7, the behavior of TMOM5OX with Er is inconsistent and more compelling from La metal ion. In case of Er[TMOM5OX], the high energy ligand peak at 248 nm, (e = 38,000 M1 cm1) declines and get shifted towards higher wavelength, 265 nm (e = 28, 400 M1 cm1) with the concomitant ascent in absorbance due to the deprotonation and complexation of Npyr groups, but for La complex very less significant non-structured peak appears at 265 nm. Analogously, the low energy band at 330 nm (e = 4000 M1 cm1) gets shifted towards higher wavelength with the formation of broad band around 392 nm (e = 3000 M1 cm1), but no such shift was observed in case of La complex with the increment of pH. Since the absorption coefficients of these bands

are large for a d–d transition, they are likely due to ligand-to-metal charge transfer (LMCT). However, interpreting the spectra, the appearance of band at higher wavelengths with lower intensity may be attributed to the coordination of pyridine nitrogen atoms to the metal ion upon chelation. It can thus be assumed that the band at 334–340 nm and 392 nm is of the Npyr ? La(III) and Npyr ? Er(III) type, respectively. Similar results have also been reported by ligands Tsox and TsoxMe with above two metal ions [36]. The isosbestic points asserts the formation of four protonated complexes of La and Er with the ligand in the pH range of 1.79–7.2 and 1.71–6.83, subsequently which can be well comprehend from species distribution curves (Fig. 6a and b). Above pH 6.85, the deprotonation and complexation of hydroxyl groups with Er ion, results in the bathochromic and hypochromic shifts at 278 nm, (e = 33,000 M1 cm1) and at 295 nm, (e = 29,600 M1 cm1), respectively with the development of new band at 320 nm (e = 4000 M1 cm1) and the band around 392 nm was elongated upto 425 nm, (e = 3600 M1 cm1) (Fig. 4b). Similarly, the band at higher energies of La[TMOM5OX] was red shifted to 285 nm, (e = 18,000 M1 cm1) and low energy band was also shifted towards near-visible region at 372 nm, (e = 1600 M1 cm1) (Fig. 3b) indicating the formation of neutral complex. As observed in other quinolinate (O-Trensox, Oxinobactin and Sulfoxinobactin) [53,56] and phenolate complexes the bands at 425 nm and 372 nm are assigned to charge transfer by O ? Er+3 and O ? La+3 respectively. The bands due to Npyr ? Ln disappear in this pH range and the change in absorption spectrum can be interpreted by a change in the coordination sphere of the lanthanide ion. In contrarary to acidic ? neutral pH range, the concurrent downslope of LMCT bands of Er complex was observed over the pH range of 6.8511.4 and the behavior of ligand with La ion in the basic medium was somehow similar as observed in physiological pH. No spectral change was observed on further pH variation upto 12.29 for La complex and 12.54 for Er complex, indicating the formation of hydroxo complexes for both metal ions. It can thus be assumed that TMOM5OX is the efficient La(III) and Er(III) chelator in wide range of pH 1.711. In order to validate the above discussion for the formation of lanthanide complexes, further investigation for the electronic

373

R. Akbar et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

Fig. 8. Correlation between the experimental log K and Calculated DG° (a) La[TMOM5OX] (b) Er[TMOM5OX].

transitions and electronic structures of different species formed for both La and Er complexes in solution, were performed from a theoretical stand point through semiempirical sparkle model calculations. The computed optical data of all the PM6/Sparkle optimized species is listed in Table 2 and the simulated electronic absorption spectra of predicted complexes obtained via ZINDOS/CIS method, match with the species obtained from experimentally determined UVvisible spectra (Figs. S1 and S2, Supplementary material). Table 2 presents the calculated singlet ? singlet electronic transitions and corresponding oscillator strengths of different species of metal complexes of TMOM5OX and their assignments to the spectral bands. The ground state optimized structures of La and Er complexes of TMOM5OX are shown in Fig. S3(a–c) (Supplemen-

tary material), respectively. From an electronic point of view, for all the Ln-TMOM5OX acidic forms, the second electronic transition basically corresponds to a HOMO ? LUMO excitation and orbitals have the same character (namely p and p, respectively). On analysing the molecular orbitals, the HOMO is somehow localized on the aromatic ring carrying the OH function while the antibonding LUMO is delocalised over the nitrogen atom of the quinoline system; we expect this transition to be relatively sensitive, in terms of intensity and position, to overall protonation degree of the molecules. In particular, the significant hypochromic shifts are computed going from protonated (acidic) to neutral species for both lanthanide ions and no remarkable blue or red shifts were predicted in going from [Ln(H4L)]4+ to [Ln(HL)]+. While the red shifts

Table 2 Computed optical properties: Absorption-S0 and emission-S1; wavelengths, oscillator strengths (fcalc), electronic transitions and energies calculated at the CIS-ZINDO/S method, using sparkle/PM6 ground state and excited state geometry of the complexes. S0

S1

Complexes

knm (abs)

fcalc

Nature

Transition (excitation energy)

knm(em)

fcalc

Nature

Transition (emission energy)

La(H4L)

278.42 535.28 291.46 572.78 299.23 551.23 289.75 547.14 312.51 546.46 300.60 505.51 265.30 385.63 262.57 376.76 270.18 399.14 273.07 391.07 280.37 445.32 299.85 355.57 459.52 300.78 365.57 482.47

0.3581 0.0396 0.1996 0.0394 0.1766 0.0336 0.1068 0.0922 0.0831 0.0241 0.0697 0.0467 0.1302 0.0573 0.1496 0.0590 0.1381 0.0740 0.1474 0.0293 0.0901 0.0134 0.0873 0.0661 0.0095 0.1002 0.0579 0.0289

HOMO3 ? LUMO+2 HOMO ? LUMO HOMO2 ? LUMO+1 HOMO ? LUMO HOMO3 ? LUMO+1 HOMO ? LUMO HOMO2 ? LUMO+1 HOMO ? LUMO HOMO1 ? LUMO+2 HOMO1 ? LUMO HOMO1 ? LUMO+1 HOMO2 ? LUMO HOMO3 ? LUMO+2 HOMO ? LUMO HOMO2 ? LUMO+2 HOMO ? LUMO HOMO2 ? LUMO+1 HOMO ? LUMO HOMO2 ? LUMO+1 HOMO ? LUMO HOMO1 ? LUMO+2 HOMO ? LUMO+1 HOMO1 ? LUMO+2 HOMO2 ? LUMO+1 HOMO ? LUMO+1 HOMO1 ? LUMO+1 HOMO2 ? LUMO+1 HOMO ? LUMO+1

116 ? 122 (3.829 eV) 119 ? 120 (0.526 eV) 117 ? 121 (3.543 eV) 119 ? 120 (0.683 eV) 116 ? 121 (2.916 eV) 119 ? 120 (0.652 eV) 117 ? 121 (2.670 eV) 119 ? 120 (0.911 eV) 118 ? 122 (1.877 eV) 118 ? 120 (0.809 eV) 118 ? 121 (1.234 eV) 117 ? 120 (0.768 eV) 116 ? 122 (1.293 eV) 119 ? 120(0.674 eV) 117 ? 122 (1.202 eV) 119 ? 120 (0.661 eV) 117 ? 121 (1.115 eV) 119 ? 120 (0.592 eV) 117 ? 121 (1.109 eV) 119 ? 120 (0.521 eV) 118 ? 122 (0.899 eV) 119 ? 121 (0.211 eV) 118 ? 122 (0.892 eV) 117 ? 121 (0.929 eV) 119 ? 121 (0.247 eV) 118 ? 121 (0.887 eV) 117 ? 121 (0.893 eV) 119 ? 121 (0.173 eV)

376.83

0.0791

HOMO1 ? LUMO+1

118 ? 121 (0.463 eV)

375.72 524.87 525.84 638.03 525.84

0.0681 0.1134 0.1274 0.0712 0.1629

HOMO ? LUMO+1 HOMO ? LUMO HOMO2 ? LUMO HOMO ? LUMO HOMO1 ? LUMO

119 ? 121 119 ? 120 117 ? 120 119 ? 120 118 ? 120

410.46

0.1632

HOMO1 ? LUMO

118 ? 120 (1.139 eV)

378.17

0.0903

HOMO ? LUMO+1

119 ? 121 (0.694 eV)

425.28

0.0513

HOMO1 ? LUMO+1

118 ? 121 (0.312 eV)

502.58

0.0774

HOMO ? LUMO

119 ? 120 (0.737 eV)

515.03

0.0897

HOMO1 ? LUMO

118 ? 120 (0.798 eV)

520.15 624.06 518.23 620.83 426.28

0.0965 0.0587 0.2860 0.0672 0.3260

HOMO2 ? LUMO HOMO2 ? LUMO+1 HOMO1 ? LUMO+1 HOMO ? LUMO HOMO1 ? LUMO+1

117 ? 120 117 ? 121 118 ? 121 119 ? 120 118 ? 121

418.14 547.31

0.9670 0.0641

HOMO1 ? LUMO+1 HOMO ? LUMO

118 ? 121 (0.847 eV) 119 ? 120 (0.401 eV)

La(H3L) La(H2L) La(HL) LaL LaL(OH) Er(H4L) Er(H3L) Er(H2L) Er(HL) ErL ErL2

ErL(OH)

(0.431 eV) (0.897 eV) (0.924 eV) (0.437 eV) (1.126 eV)

(0.839 eV) (0.334 eV) (1.997 eV) (0.483 eV) (2.054 eV)

374

R. Akbar et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

were observed of significant magnitude from [ErL] to [ErL(OH)], substantial blue shifts (of about 45 nm) were observed in going from neutral to basic forms (LaL to [LaL(OH)]). A quantitative explanation of these variations in absorption transitions and energies can be given by the inspection of the HOMO and LUMO orbitals computed for all species shown in Fig. S4 (Supplementary material). Indeed protonation of the N atom increases the acceptor character of the quinoline ring. Thus, it is expected to raise the LUMO energy and increases the vertical S0 ? S1transition energy, in agreement with the above mentioned computed hypochromic shift on deprotonation of NH+ upon complexation. On the other hand, the deprotonation of the hydroxyl function will destabilize the HOMO, thus closing the gap and giving rise to smaller transition. Notably both the above-mentioned effects (i.e. stabilization of the LUMO and destabilization of the HOMO) are concomitant when going from Ln(H4L)]4+ to [LnL(OH)]. Not surprisingly these species are computed to have the HOMO ? LUMO gap (3.061 eV) and a significant red shift of the second transition (98 nm) is consistently computed particularly going from Er(H4L)]4+ to [ErL(OH)]. The first excited state basically corresponds to a HOMO3 ? LUMO+1 transition, and it is the only intense transition in this spectral region but Er(L2) and [ErL(OH)] that show a second transition of approximately the same intensity at 365 nm and 370 nm, respectively. This later corresponds, in both cases, HOMO to LUMO+1 excitation, the LUMO+1 being a p orbital delocalized over all quinoline rings with negligible contribution from the acidic functions. All of these results precisely concede the hypothesis that the influence of the lanthanide ion on the absorption and electronic properties (and consequently on the ground state geometries) of the complexes is a small effect. A further important point is that taking into account, at least partially; the core interaction between the metal ion (+3 point charge) and its neighbor atoms appreciably improves the quality of the theoretical results. Luminescence studies Before the pH dependent luminescence properties of the lanthanide complexes of TMOM5OX can be expounded, it is necessary to give a very brief description of the fluorescence properties of pH-dependent fluorescent sensors in which 8HQ moiety is utilized as fluorescent receptor unit. The fluorescent intensity of reported pH sensors was found to change depending only on the change of pH value in a monotonous way, which gets increased or decreased along with the change (either increase or decrease) in the pH value in a consistent way due to the only one photoinduced electron transfer mechanism between pyridyl –N and OH– group of quinoline unit, resulting in the OFF–ON or ON–OFF type of fluorescent pH sensors [57]. There are still very few reports on the pH sensors whose fluorescent intensity changes in acidic and basic systems along with decrease and increase in the pH value, respectively, based on the opposite photoinduced electron transfer processes between a receptor and signaling unit [58]. To scrutinize the pH-dependent optical properties of La and Er complexes of TMOM5OX with unique pyridyl –N and –OH structural characteristics, the pH vs. fluorescence titration experiments were carried out in aqueous medium with an appropriate volume of 0.1 M HCl and 0.1 M NaOH used for protonation of pyridyl –N group and deprotonation of the –OH group, respectively (discussed in complexation section). As shown in Figs. 3 and 4, the electronic absorption spectrum of complexes of TMOM5OX with both lanthanide ions exhibits the exemplary behavior (red and hypochromic shifts) along with increasing the pH value of the system. In good contrast, the changes in the blue-green emission maxima located at 370 nm and 435 nm, which were excited at 332 nm and 393 nm, were monitored at pH of aqueous solution of La and Er with TMOM5OX in 1:1 ratio, ranging from 1.81–12.3 and 1.84–

12.67, respectively (Fig. 5a–b) at room temperature. The emission intensity at 370 nm and 435 nm (with meager quantum yield) as well as maximum wavelength remains almost unchanged under acidic conditions (within the pH range of 1.8–4.2 and 1.8–4.0 in case of La and Er complexes, respectively). Most interestingly, with an increase in the pH value from 5.0 and 4.2, the evolution of emission peak at about 508 nm and 520 nm (for La and Er complexes, respectively) results along with no change in blue-green emission at 370 nm and 435 nm while the fluorescence intensity at 508 nm and 520 nm get enhanced upto 2–3 fold, along with an increase in the pH value upto 6.7. The intense fluorescence of both complexes promptly get blue shifted at 428 nm and 440 nm after pH 7.5 and 7.9 with little growth of emission in case of La[TMOM5OX] whereas significant enhancement of emission of Er[TMOM5OX] upto 3-fold times more intense than at 6.8 pH, which get 4-fold times quenched under basic conditions (pH 8.45), in a similar manner as observed under acidic conditions. Similarly the peak at 428 nm gets blue shifted upto 370 nm for La[TMOM5OX] which shows insignificant change under basic conditions (9.5–12.5). The greater enhancement of fluorescence intensity of Er complex at pH 7.94 around 440 nm may be attributed to the formation of ErL2 species at this particular pH (cf. species distribution curves, Fig. 6b). The significant fluorescence of this complex is due to the effective excitation of erbium ion by two hexadentate chelators which increases the antenna effect attached to metal ion by providing 12 coordination number. The strong luminescence is due to the chemical constraint of chromophoric ligands in that the coordination cavity provides enough donor atoms to saturate the coordination number which encapsulate, absorb energy, transfer it efficiently to the central metal and protect from the solvent molecules (H2O) which usually quench the fluorescence efficiency. The organized molecular architecture [Er(TMOM5OX)2], shown in Fig. S3(c) (Supplementary material) formed above physiological pH (7.94) proved TMOM5OX can be used in efficient light conversion devices. Thus, the main significance of this spectral behavior is the use of this (ErL2) complex as luminescent structural probe for large biological molecules. It is worth noting that more significant changes under physiological pH and less significant changes under acidic and basic medium confirm that protonation of N atom lowers the LUMO energy and decreases the vertical S1 ? S0 transition energy and the deprotonation of the hydroxyl function will also destabilize the HOMO, thus closing the gap and giving rise to smaller transition, while in neutral form the HOMO–LUMO gap increases which corresponds significant red shifts along with intense fluorescence. Thus among all the species formed during the pH range of 1.8–12.5, Ln(HL)]+, [LnL] and [Ln(L)2] were found to be more fluorescent which were formed under physiological pH in case of both La and Er complexes. While the more protonated and hydroxo complexes were found to be less fluorescent. These results reveal the typical characteristics of a pH-fluorescent probe through the photoinduced proton transfer (PPT) enhancement process and photoinduced electron transfer (PET) quenching processes [58] between receptor and fluorescent signaling group for this tripodal sensor. In other words, these two processes between the TMOM5OX receptor and metal ions are responsible for the particular fluorescent properties along with the pH range of 1.8–12.5. The unique pH-dependent fluorescent property of these complexes indicates the potential application, as pH indicators under both acidic and basic conditions and will act as novel OFF–ON–OFF type of fluorescent pH sensors. For the purpose of understanding the interesting pH-dependant fluorescence characteristics of these La and Er complexes of TMOM5OX, the theoretical calculations for excited state sparkle/PM6 optimized structures based on CIS method using ZINDO/S for all corresponding species were carried out. The calculated emission spectra of these species are shown in Fig. S5(a–b) (Supplementary

R. Akbar et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

material). In comparison, the computed emission wavelengths (kem) agree well with experimental data, particularly with the experimental values of 370 nm, 420 nm and 525 nm for protonated, hydroxo and neutral complexes, respectively, whose excellent agreement was not only obtained for the kem numerical values but also for the shape of the experimental emission spectra. Computed vertical S1 ? S0 emission energies (that is, fluorescence reported in Table 2) are, as expected, all red-shifted with respect to the corresponding absorption values. This observation is in line with the very small structural changes computed for each form when going from the ground to the excited state. Indeed, the largest structural evolutions are related to a modification of the hydrogen bond network at the excited state due to the photoacidity or photobasicity of the different functional groups. In particular, the NH+/N pair seems to act as a photobase, the nitrogen atoms interacting stronger with water’s hydrogens at the excited state than at the ground state. On the contrary, the OH function acts as a photoacid, its hydrogen interacting stronger with water’s oxygen at the excited state than at the ground state. These results can justify the fluorescence behavior observed for these complexes in solution. So the present examples appear to represent the novel pH fluoresent sensors, with TMOM5OX as receptor, in particular with OFF–ON–OFF type nature, which should be helpful for designing and preparing novel versatile fluorescent sensors with potential applications in chemical and biological fields.

Conclusion The solution studies described in this paper shows that tripod TMOM5OX leads to soluble and thermodynamically stable Ln(III) complexes in water, featuring resistance toward hydrolysis in philological pH, with pLn values in the range of 24–27. The single major species present at physiological pH, [La(LH3)]3+ and [Er(L)] with log b values of 43.91 and 35.76, respectively, shows interesting photophysical properties in ultraviolet and visible range. The photophysical properties of the complexes formed studied by means of interaction with single excitations (CIS) based on ZINDO/S methodology supplemented the observations. In particular, computed absorption and emission spectra of all the pH-dependent species were found in good agreement with experimentally claimed data. Some insights are shared on the combined effects of 8-hydroxyquinoline binding units of TMOM5OX and Ln ions on the geometry, electronic structure, and optical properties of complexes. Our findings also suggest that experimental spectroscopic data are better reproduced in ZINDO calculations when the lanthanide is represented by a point charge and these methods also enabled the calculation of luminescent properties. The results of these methods not only reproduced well, the experimental values but also helped in explaining the low values of quantum efficiency observed for these complexes. Our expectation is that this set of semiempirical methodologies can be used in the theoretical design of new luminescent Lanthanum and Erbium complexes. From a mechanistic point of view, the absence of green fluorescence of acidic forms of Ln complexes experimentally observed claimed to be related to the formation of the nonemitting zwitterionic forms of quinoline was here proven and validated. In addition, the possibility of excitation in the visible range is a definitive asset for probes intended for bioanalyses. Altogether, this molecular engineering opens astounding perspectives for the development of metal chelators and luminescent bioprobes since the synthesis of TMOM5OX can be undertaken on large scale effortlessly and derivatization of this chelating agent is also easy at hand, in particular with respect to the grafting of central unit and functional groups for the coupling with large biological molecules.

375

Acknowledgement The financial support from the DST, New Delhi is gratefully acknowledged.

Appendix A. Supplementary material Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.saa.2014.03.045. References [1] F. Bonnet, M. Visseeaux, A. Pereira, D. Barbier-Baudry, Macromolecules 38 (2005) 3162–3170. [2] X. Liu, S. Wang, C. Peng, J. Rare Earths 28 (2010) 965–968. [3] H. Shargi, M.A. Nasser, Bull. Chem. Soc. (in Jpn.) 76 (2003) 137–142. [4] J.C.G. Bunzli, Lanthanide Probes in Life, Chemical and Earth Sciences, Elsevier, Amsterdam, 1989. [5] J.C.G. Bunzli, C. Piguet, Chem. Soc. Rev. 34 (2005) 1048–1077. [6] (a) M.D. Allendorf, C.A. Bauer, R.K. Bhakta, R.J.T. Houk, Chem. Soc. Rev. 38 (2009) 1330–1352; (b) K. Binnemans, Chem. Rev. 109 (2009) 4283–4374. [7] L.D. Carlos, R.A.S. Ferreira, V.D. Bermudez, S.J.L. Ribeiro, Adv. Mater. 21 (2009) 509–534. [8] S.V. Eliseeva, J.C.G. Bunzli, Chem. Soc. Rev. 39 (2010) 189–227. [9] N. Sabbatini, M. Guardigli, J.M. Lehn, Coord. Chem. Rev. 123 (1993) 201–228. [10] Y.H. Kim, N.S. Baek, H.K. Kim, ChemPhysChem 7 (2006) 213–221. [11] (a) P.A. Vigato, V. Peruzzo, S. Tamburini, Coord. Chem. Rev. 253 (2009) 1099– 1201; (b) T.S. Kang, B.S. Harrison, M. Bouguettaya, T.J. Foley, J.M. Boncella, K.S. Schanze, J.R. Reynold, Adv. Funct. Mater. 13 (2003) 205–210. [12] C. Freund, W. Porzio, U. Giovanella, F. Vignali, M. Pasini, S. Destri, A. Mech, S. Di Pietro, L. Di Bari, P. Mineo, Inorg. Chem. 50 (2011) 5417–5429. [13] C. Marchal, Y. Filinchuk, X.Y. Chen, D. Imbert, M. Mazzanti, Chem. Eur. J. 15 (2009) 5273–5288. [14] C.W. Tang, S.A. Vanslyke, Appl. Phys. Lett. 51 (1987) 913–915. [15] (a) M. Albrecht, R. Frohlich, J.C.G. Bunzli, A. Aebischer, F. Gumy, J. Hamacek, J. Am. Chem. Soc. 129 (2007) 14178–14179; (b) A. Meyers, A. Kimyonok, M. Weck, Macromolecules 38 (2005) 86718678. [16] (a) L. Ning, M.I. Trioni, G.P. Brivio, J. Mater. Chem. 17 (2007) 44644470; (b) R. Kikkeri, L.H. Hossain, P.H. Seeberger, Chem. Commun. (2008) 2127–2129. [17] R.G.W. Hollingshead, Oxine and Its Derivatives, vol. I, Butterworths, London, 1954. [18] E.E.S. Teotonio, J.G.P. Espínola, S.F. Oliveira, D.L. Farias, O.L. Malta, H.F. Brito, Polyhedron 21 (2002) 1837–1844. [19] E.G. Moore, A.P.S. Samuel, K.N. Raymond, Acc. Chem. Res. 42 (2009) 542–552. [20] (a) R. Van Deun, P. Fias, K. Driesen, K. Binnemans, C. Gorller-Walrand, Phys. Chem. Chem. Phys. 5 (2003) 2754–2757; (b) N.M. Shavaleev, R. Scopelliti, F. Gumy, J.C.G. Bünzli, Inorg. Chem. 47 (2008) 9055–9068. [21] G. Bozoklu, C. Marchal, J. Pecaut, D. Imbert, M. Mazzanti, Dalton Trans. 39 (2010) 9112–9122. [22] M.E. Masaki, D. Paul, R. Nakamura, Y. Kataoka, S. Shinoda, H. Tsukube, Tetrahedron 65 (2009) 2525–2530. [23] C.P. McCoy, F. Stomeo, S.E. Plush, T. Gunnlaugsson, Chem. Mater. 18 (2006) 4336–4343. [24] T. Gunnlaugsson, D.A. MacDónaill, D. Parker, J. Am. Chem. Soc. 123 (2001) 12866–12876. [25] A. Nonat, D. Imbert, J. Pecaut, M. Giraud, M. Mazzanti, Inorg. Chem. 48 (2009) 4207–4218. [26] X.H. Yan, Z.H. Cai, C.L. Yi, W.S. Liu, M.Y. Tan, Y. Tang, Inorg. Chem. 50 (2011) 4857–4867. [27] J.C.G. Bunzli, C. Piguet, Chem. Soc. Rev. 34 (2005) 1048–1077. [28] F. Artizzu, L. Marchio, M.L. Mercuri, L. Pilia, A. Serpe, F. Quochi, R. Orru, F. Cordella, M. Saba, A. Mura, G. Bongiovanni, P. Deplano, Adv. Funct. Mater. 17 (2007) 2365–2376. [29] (a) M. Albrecht, O. Osetska, J. Klankermayer, R. Frohlich, F. Gumy, J.C.G. Bunzli, Chem. Commun. 18 (2007) 1834–1836; (b) M. Albrecht, S. Mirtschin, O. Osetska, S. Dehn, D. Enders, R. Frohlich, T. Pape, E.F. Hahn, Eur. J. Inorg. Chem. 20 (2007) 3276–3278. [30] R.O. Freire, A.M. Simas, J. Chem. Theory Comput. 6 (2010) 2019–2023. [31] A.V.M. de Andrade, N.B.Jr. Da Costa, A.M. Simas, G.F. de Sa, Chem. Phys. Lett. 227 (1994) 349–353. [32] M.E. Mesquita, F.R.G. Silva, R.Q. Albuquerque, R.O. Freire, E.C. da Conceição, N.B.C. Júnior, J.E.C. da Silva, G.F. de Sá, J. Alloys Compd. 366 (2004) 124–131. [33] R.O. Freire, M.O. Rodrigues, F.R.G. Silva, N.B.C. Júnior, M.E. Mesquita, J. Mol. Model. 12 (2005) 16–23. [34] A.E. Martell, R.J. Motekaitis, Determination and Use of Stability Constants, Second ed., VCH, New York, 1992. [35] L. Alderighi, P. Gans, A. Ienco, D.C. Peters, A. Sabatini, A. Vacca, Coord. Chem. Rev. 184 (1999) 311–318.

376

R. Akbar et al. / Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 129 (2014) 365–376

[36] C. Steve, I. Daniel, C. Anne-Sophie, G.B. Jean- Claude, Inorg. Chem. 45 (2006) 732–742. [37] F.L. Javier, N. Cristina, N.F. Olalla, L.C. Jose, L. Carlos, J. Sérgio Seixas de Melo, S.L. Carlos, Inor. Chim. Act. 381 (2012) 218–228. [38] M. Albrechet, K. Witt, R. Frohlich, O. Kataeva, Tetrahedron (2002) 561–567. [39] A.R. Allouche, J. Comput. Chem. 32 (2011) 174–182. [40] J.J.P. Stewart, MOPAC2012, Version 12.239W, in: S.C. Chemistry (Eds.), Colorado Springs, USA, 2012. [41] J.G. Santos, J.D.L. Dutra, A. Severino.Jr, R.O. Freire, N.B. da Costa Jr., J. Phys. Chem. A 116 (2012) 4318–4322. [42] D. Guillaumont, H. Bazin, J.M. Benech, M. Boyer, G. Mathis, ChemPhysChem 8 (2007) 480–488. [43] M.O. Rodrigues, F.A. Almeida, et al., J. Phys. Chem. B 113 (2009) 12181–12188. [44] R.O. Freire, G.B. Rocha, A.M. Simas, J. Mol. Model. 12 (2006) 373–389. [45] M. Dolg, H. Stoll, A. Savin, H. Preuss, Theor. Chim. Acta 75 (1989) 173–194. [46] K. Hegetschweiler, R.D. Hancock, M. Ghisletta, T. Kradolfer, V. Gramlic, H.W. Schmalle, Inorg. Chem. 32 (1993) 5273–5284. [47] P.P. Lima, S.S. Nobre, R.O. Freire, S.A. Junior, R.A. Sa, et al., J. Phys. Chem. C 111 (2007) 17627–17634. [48] P.P. Lima, R.A. Sa, et al., ChemPhysChem 7 (2006) 735–746.

[49] (a) J.E. Ridley, M.C. Zerner, Theor. Chim. Acta 42 (1976) 223–236; (b) M.C. Zerner, G.H. Loew, R.F. Kirchner, U.T. Mueller-Westerhoff, J. Am. Chem. Soc. 102 (1980) 589–599. [50] F. Bravard, C. Rosset, P. Delangle, Dalton Trans. 13 (2004) 2012–2018. [51] D. Imbert, N. Fatin-Rouge, J.C.G. Bunzli, Eur. J. Inorg. Chem. 7 (2003) 1332–1339. [52] G. Serratrice, H. Boukhalfa, B. Claude, P. Baret, C. Caris, J.L. Pierre, Inorg. Chem. 36 (1997) 3898–3910. [53] K.N. Raymond, G. Muller, F. Matzanke, Top. Curr. Chem. 123 (1984). 49-102. [54] A.E. Martell, R.M. Smith, Critical Stability Constants, Plenum Press, New York, 1974. [55] D.H. Amaury, G. Gisèle, P. Christian, G. Serratrice, Inorg. Chem. 51 (2012) 12142–12151. [56] (a) C.N. Baki, E.U. Akkaya, J. Org. Chem. 66 (2001) 1512–1513; (b) T. Chen, W. Zhu, Y.F. Xu, S.Y. Zhang, X.J. Zhang, X.H. Qian, Dalton Trans. 39 (2010) 1316–1320. [57] C. Yuting, W. Hailong, W. Liang, B. Yongzhong, J. Jianzhuang, J. Org. Chem. 76 (2011) 3774–3781. [58] (a) Y. Urano, M. Kamiya, T. Kanda, T. Ueno, K. Hirose, T. Nagano, J. Am. Chem. Soc. 127 (2005) 4888–4894; (b) K. Tanaka, T. Miura, N. Umezawa, Y. Urano, K. Kikuchi, T. Higuchi, T. Nagano, J. Am. Chem. Soc. 123 (2001) 2530–2536.