Experimental and Computational Studies on the Interaction of a Dansyl-Based Fluorescent Schiff Base Ligand with Cu2+ Ions and CuO NPs

We studied the interaction of Cu2+ ions and CuO nanoparticles with the fluorescent Schiff base ligand H3L, which derives from the condensation of 4-formyl-3-hydroxybenzoic acid with N-(2-aminobenzyl)-5-(dimethylamino)naphthalene-1-sulfonamide (DsA). A detailed assignment of the most significant bands of the electronic and infrared spectra of H3L and DsA was performed using DFT methods, based on both crystal structures. The affinity of H3L to react with Cu2+ ions in solution (KB = 9.01 103 L mol−1) is similar to that found for the Cu2+ ions present on the surface of CuO NPs (KB = 9.84 103 L mol−1). Fluorescence spectroscopic measurements suggest five binding sites for H3L on the surface of the CuO NPs used. The µ-XRF analysis indicates that a polycrystalline sample of CuO-H3L NPs contains 15:1 Cu:S molar ratio (CuO:H3L). ATR-FTIR spectroscopy, supported by DFT calculations, showed that the HL2− (as a phenolate and sulfonamide anion) is coordinated to superficial Cu2+ ions of the CuO NPs through their azomethine, sulphonamide, and phenolic groups. A solution of H3L (126 ppb) shows sensitive responses to CuO NPs, with a limit of detection (LOD) of 330 ppb. The working range for detection of CuO NPs with [H3L] = 126 ppb was 1.1–9.5 ppm. Common metal ions in water, such as Na+, K+, Mg2+, Ca2+, Fe3+, and Al3+ species, do not interfere significantly with the detection of CuO NPs.


Introduction
The unique properties of metal and metal oxide nanomaterials have led to their widespread use in daily life [1][2][3][4][5], which has the clear negative effect of an increasing release of these nanomaterials into the environment [6][7][8][9][10]. Particularly, CuO NPs are manufactured on a large scale for their varied industrial and domestic applications, as they are generally considered to be non-toxic. Thus, one of the common usages of CuO NPs is in water treatment, as an effective adsorbent of several pollutants [4]. However, some concern is arising about the possible adverse health and environmental effects of their release since several studies claim that CuO NPs could have negative impacts on different living organisms [11][12][13].
Therefore, the development and improvement of methods for the detection of CuO NPs is a growing field of study. Thus, Zhao et al. used EDS analysis to quantify the presence of copper on the root surface of an aquatic plant, after CuO NP exposure [14]. ICP-OES Therefore, the development and improvement of methods for the detection of CuO NPs is a growing field of study. Thus, Zhao et al. used EDS analysis to quantify the presence of copper on the root surface of an aquatic plant, after CuO NP exposure [14]. ICP-OES was used by Peng et al. to measure the copper content in root of rice after adding CuO NPs to the soil [15], while Navratilova et al. applied the single particle ICP-MS method for the detection of CuO NPs in colloidal extracts from natural soil samples [16]. In this sense, our research group has also used a pyrrole-based [17], as well as phenolbased Schiff base ligands [18] as fluorescent sensors for the detection of CuO NPs (LOD = 13.83 and 9.80 µgL −1 , respectively). These two N-tosyl imines have a moderate fluorescence emission, which drastically decreases by interaction with copper ions present on the surface of the CuO NPs. In the present study, with the aim of avoiding the disadvantages of using weakly fluorescent ligands, we designed a new sensor with higher fluorescence emission intensity providing potential N and O donor atoms, which have demonstrated to be suitable for coordinating those superficial copper ions present on the surface of the CuO nanoparticles [17,18].
We investigated the ability of the dansyl-based fluorescent Schiff base ligand H3L for detecting CuO NPs in ethanol-water solutions, with short response time and sensing at basic pH. H3L can be easily synthesized (Scheme 1) from the condensation of 4-formyl-3hydroxybenzoic acid with N-(2-aminobenzyl)-5-(dimethylamino)naphthalene-1-sulfonamide (DsA). This dansyl-derived amine had been previously obtained by selective reaction of 2-(aminomethyl)aniline with dansyl chloride [19]. H3L displays an N,N,O-binding domain suitable to interact with Cu 2+ ions, either in solution, or as constituents of a solid matrix as CuO NPs. In this case, additional O atoms belonging to carboxylate and sulfonamide groups can also be involved in the interaction. Scheme 1. Schematic representation of the synthesis of H3L in two steps: (i) selective dansylation of a primary diamine and (ii) aldiminic condensation of DsA with 4-formyl-3-hydroxybenzoic acid.
Since the knowledge of the surface coordination chemistry of NPs [20][21][22][23] is advantageous for the development of chemical sensors for them, we also present here some experimental and computational studies on the interactions of H 3 L with CuO NPs. SEM-EDX, µ-XRF, ATR-IR, and DFT were used to investigate the interaction of H 3 L with the surface of CuO NPs. With the aim of assisting in the interpretation of the results obtained from the study of the interaction between H 3 L and CuO NPs, the copper(II) complex, Cu 2 L 2 (H 2 O) 4 , was synthesized and studied by ATR-IR, UV-Vis, and fluorescence spectroscopies. DFT calculations have been used to predict the effect of the symmetry of DsA, H 3 L, and Cu 2 L 2 (H 2 O) 4 on their vibrational modes and electronic properties. Binding constants have been also determined since these constants characterize the affinity of H 3 L for both NMs and ions. Changes in the fluorescence emission spectra of H 3 L upon increasing addition of CuO NPs have been investigated as well, while taking into consideration possible interferences from some common metal ions in water.

Results and Discussion
With the aim of investigating the ability of a dansyl-based fluorescent Schiff base ligand for detecting copper oxide nanomaterials in ethanol-water samples, we designed the easily synthesized H 3 L (Scheme 1). An ethanol solution of H 3 L at neutral pH emitted a maximal green fluorescence at about 524 nm, when it was exposed to radiation with a wavelength of 400 nm. This fluorescent behavior was assigned to the dansyl moiety, which is well known for a typical charge-transfer band between the donor dimethylamino group and the sulfonyl group [24]. H 3 L can be obtained in a quick two-step synthesis (Scheme 1). In the first step, 2-(aminomethyl)aniline selectively reacts, through the aminomethyl group, with dansyl chloride to yield N-(2-aminobenzyl)-5-(dimethylamino)naphthalene-1-sulfonamide (DsA). Subsequent nucleophilic addition of DsA to 4-formyl-3-hydroxybenzoic acid leads to the desired Schiff base ligand H 3 L [19]. Both reactions were performed at room temperature, in order to prevent the dansylation of both amino groups in the first step, as well as undesired chemical reactions of DsA in the second step. Spectroscopic characterization details of H 3 L and DsA, obtained as powdery solids, are shown in the experimental section. Figures  S1 and S2 show the 1 H NMR spectra of H 3 L and DsA (in dmso-d 6 ), respectively. Single crystals of [H 4 L]Cl·0.92H 2 O were obtained by slow evaporation of an ethanol solution of DsA and 4-formyl-3-hydroxybenzoic acid, to which triethylammonium chloride ([HNEt 3 ] + , pK a = 10.75) had been added.

Crystal Structures for H 3 L and DsA
The molecular structure of two neighboring [H 4 L] + cations is shown in Figure 1, as an ellipsoid diagram, which includes a labelling scheme equivalent to that used in the crystal structure of DsA (see below). Some significant crystal parameters and refinement data are summarized in Table S1. The asymmetric unit of [H 4 L]Cl·0.92H 2 O consists of two crystallographically independent cationic moieties of [H 4 L] + , showing their amine N atoms protonated (N33 and N73), two chloride counterions, and several water molecules with low occupation sites, which sum up to 0.92, per asymmetric unit. It must be noted that one of the chloride anions, as well as one of the dansyl residues (C63-C75) are disordered on two different occupation sites (0.832 and 0.168).
the easily synthesized H3L (Scheme 1). An ethanol solution of H3L at neutral pH emitt a maximal green fluorescence at about 524 nm, when it was exposed to radiation with wavelength of 400 nm. This fluorescent behavior was assigned to the dansyl moie which is well known for a typical charge-transfer band between the donor dimethylami group and the sulfonyl group [24].
H3L can be obtained in a quick two-step synthesis (Scheme 1). In the first step, (aminomethyl)aniline selectively reacts, through the aminomethyl group, with dans chloride to yield N-(2-aminobenzyl)-5-(dimethylamino)naphthalene- . Subsequent nucleophilic addition of DsA to 4-formyl-3-hydroxybenzoic acid lea to the desired Schiff base ligand H3L [19]. Both reactions were performed at room temp ature, in order to prevent the dansylation of both amino groups in the first step, as well undesired chemical reactions of DsA in the second step. Spectroscopic characterizati details of H3L and DsA, obtained as powdery solids, are shown in the experimental se tion. Figures S1 and S2 show the 1 H NMR spectra of H3L and DsA (in dmso-d6), respe tively. Single crystals of [H4L]Cl·0.92H2O were obtained by slow evaporation of an ethan solution of DsA and 4-formyl-3-hydroxybenzoic acid, to which triethylammonium ch ride ([HNEt3] + , pKa = 10.75) had been added.

Crystal Structures for H3L and DsA
The molecular structure of two neighboring [H4L] + cations is shown in Figure 1, as ellipsoid diagram, which includes a labelling scheme equivalent to that used in the crys structure of DsA (see below). Some significant crystal parameters and refinement data a summarized in Table S1. The asymmetric unit of [H4L]Cl·0.92H2O consists of two cryst lographically independent cationic moieties of [H4L] + , showing their amine N atoms pr tonated (N33 and N73), two chloride counterions, and several water molecules with lo occupation sites, which sum up to 0.92, per asymmetric unit. It must be noted that one the chloride anions, as well as one of the dansyl residues (C63-C75) are disordered on tw different occupation sites (0.832 and 0.168). Geometric parameters (Tables S2 and S3) are within typical ranges observed for re idues related to DsA [25][26][27][28] and 4-formyl-3-hydroxybenzoic acid [29][30][31]. Strong intr molecular H bonds (Table S4)   Geometric parameters (Tables S2 and S3) are within typical ranges observed for residues related to DsA [25][26][27][28] and 4-formyl-3-hydroxybenzoic acid [29][30][31]. Strong intramolecular H bonds (Table S4) (21), are responsible for connecting adjacent molecules, which as a result are forming H-bonded dimers shown in Figure 1. The crystal packing is reinforced through interactions between HMe 2 N + -and HOOC-groups with solvated water molecules, as well as with chloride counterions.
Since the experimental geometry available for the H 3 L ligand corresponds to that of two neighboring cationic units found in the [H 4 L]Cl·0.92 H 2 O crystal, these data were used to obtain the theoretical structure for a H-bonded dimer of H 3 L ( Figure S3). This M062X/6-31G* optimized geometry shows an excellent agreement with the structure of the asymmetric unit of the [H 4 L]Cl·0.92H 2 O crystal ( Figure 1, right), with differences between experimental and theoretical values for bond lengths and bond angles lower than 0.06 Å and 3.5 o , respectively (Table S5).
Moreover, we solved the crystal structure of the precursor DsA, which is shown in Figure 2. Details of the crystal data collection and structure refinement are also summarized in Table S1. Crystals of DsA consist of crystallographically independent molecules, whose main geometric parameters (Tables S2 and S3) are within the typical ranges observed for other dansyl derivatives [25][26][27][28] (Table S6), as typically observed for dansyl [25] and tosyl [32] derivatives. metric unit of the [H4L]Cl·0.92H2O crystal (Figure 1, right), with differences b perimental and theoretical values for bond lengths and bond angles lower than 3.5 o , respectively (Table S5).
Moreover, we solved the crystal structure of the precursor DsA, which Figure 2. Details of the crystal data collection and structure refinement are al rized in Table S1. Crystals of DsA consist of crystallographically independent whose main geometric parameters (Tables S2 and S3) are within the typical served for other dansyl derivatives [25][26][27][28] (Table S6), as typically observed for dansyl [25] and tosyl [32]  Taking this crystallographic information as a starting point, we also o M062X/6-31G* optimized geometry for a single unit of DsA ( Figure S4). The g ment between the experimental structure and the calculated one is shown (right), with differences for bond lengths and bond angles lower than 0.05 respectively (Table S7). The theoretical conformational analysis of DsA suppo preference for the experimentally crystal structure here presented is related to zation facilitated by hydrogen bonds between adjacent molecules in the solid Taking this crystallographic information as a starting point, we also obtained the M062X/6-31G* optimized geometry for a single unit of DsA ( Figure S4). The good agreement between the experimental structure and the calculated one is shown in Figure 2 (right), with differences for bond lengths and bond angles lower than 0.05 Å and 3.5 o , respectively (Table S7). The theoretical conformational analysis of DsA supports that the preference for the experimentally crystal structure here presented is related to the stabilization facilitated by hydrogen bonds between adjacent molecules in the solid phase.

Electronic and Vibrational Spectra of H 3 L and DsA
DFT calculations at the M062X/6-31G* level of theory were used to assist in the correct assignment of bands in the electronic and vibrational spectra of H 3 L and DsA.
The theoretical electronic spectrum of DsA ( Figure 3) shows three dominant signals around 180, 220, and 295 nm that appear about 50-70 nm blue-shifted, with respect to the bands experimentally observed (252, 294, and 346 nm). The analysis of their natural transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety.   Table 1.
The theoretical electronic spectrum of H3L ( Figure 4, right) also shows three dominant signals located around 323, 263, and 224 nm, which appear blue-shifted with respect to those experimentally observed (340, 248, and 224 nm) (Figure 4, left). The analysis of the molecular orbitals involved in each electronic transition of H3L, and the corresponding natural transition orbitals ( Table 2), indicates that the bands at 323 and 263 nm correspond to π-π* transitions mainly centered in the quasi-planar and conjugated hydroxybenzoic/aniline region. The occupied orbital for the transition at 263 nm also shows a noticeable contribution from orbitals centered in the C=N group. The low wavelength maximum at 224 nm corresponds to a strongly mixed transition, which can also be mainly characterized as a π-π* transition centered in the orbitals of the π-stacked aromatic rings.  Table 1.
The theoretical electronic spectrum of H 3 L (Figure 4, right) also shows three dominant signals located around 323, 263, and 224 nm, which appear blue-shifted with respect to those experimentally observed (340, 248, and 224 nm) (Figure 4, left). The analysis of the molecular orbitals involved in each electronic transition of H 3 L, and the corresponding natural transition orbitals ( Table 2), indicates that the bands at 323 and 263 nm correspond to π-π* transitions mainly centered in the quasi-planar and conjugated hydroxybenzoic/aniline region. The occupied orbital for the transition at 263 nm also shows a noticeable contribution from orbitals centered in the C=N group. The low wavelength maximum at 224 nm corresponds to a strongly mixed transition, which can also be mainly characterized as a π-π* transition centered in the orbitals of the πstacked aromatic rings.
The assignment of the more significant bands of the infrared spectra of H 3 L (Table 3) and DsA (Table 4) has also been done on the basis of the calculated spectrum using DFT methods. The IR spectrum of H 3 L shows a characteristic new strong sharp band at about 1613 cm −1 , which is attributable to the formation of the imino group (Table 3). This band evidences the condensation through the amino group of DsA with the 4-formyl-3hydroxybenzoic acid. Furthermore, as a result of this condensation, the absence of the two bands at about 3470 and 3383 cm −1 , attributable to ν a NH 2 and ν s NH 2 modes, respectively, is clear in the spectrum of DsA. Figures S5 and S6 show the experimental and calculated IR spectra of H 3 L and DsA, respectively. The characteristic O-H (phenol) and N-H (sulfonamide) stretches are observed as a single signal at 3283 cm −1 in the spectrum of H 3 L. Asymmetric and symmetric O=S=O stretches can also be clearly observed at about 1315 and 1140 cm −1 , respectively, in the IR spectra of H 3 L and DsA. Tables S8 and S9 list a detailed assignment of experimental and theoretical vibrational frequencies of H 3 L and DsA, respectively. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety. transition orbitals (Table 1) indicates that the band at higher wavelengths mainly corresponds to a π-π* transition in the dansyl residue. The band at 220 nm corresponds to a markedly mixed n-π* and π-π* transition in the dansyl unit as well. Finally, among the several electronic transitions that contribute to the intense band at lower wavelengths, the most intense ones are two π-π* transitions located in the dansyl unit, and a π-π* transition located in the aniline moiety.   Table 1.
The theoretical electronic spectrum of H3L (Figure 4, right) also shows three dominant signals located around 323, 263, and 224 nm, which appear blue-shifted with respect to those experimentally observed (340, 248, and 224 nm) (Figure 4, left). The analysis of the molecular orbitals involved in each electronic transition of H3L, and the corresponding natural transition orbitals ( Table 2), indicates that the bands at 323 and 263 nm correspond to π-π* transitions mainly centered in the quasi-planar and conjugated hydroxybenzoic/aniline region. The occupied orbital for the transition at 263 nm also shows a noticeable contribution from orbitals centered in the C=N group. The low wavelength maximum at 224 nm corresponds to a strongly mixed transition, which can also be mainly character-   Table 1.
The theoretical electronic spectrum of H3L (Figure 4, right) also shows three dominant signals located around 323, 263, and 224 nm, which appear blue-shifted with respect to those experimentally observed (340, 248, and 224 nm) (Figure 4, left). The analysis of the molecular orbitals involved in each electronic transition of H3L, and the corresponding natural transition orbitals ( Table 2), indicates that the bands at 323 and 263 nm correspond to π-π* transitions mainly centered in the quasi-planar and conjugated hydroxybenzoic/aniline region. The occupied orbital for the transition at 263 nm also shows a noticeable contribution from orbitals centered in the C=N group. The low wavelength maximum at 224 nm corresponds to a strongly mixed transition, which can also be mainly character-   Table 1. The assignment of the more significant bands of the infrared spectra of H3L (Table 3) and DsA (Table 4) has also been done on the basis of the calculated spectrum using DFT methods. The IR spectrum of H3L shows a characteristic new strong sharp band at about 1613 cm −1 , which is attributable to the formation of the imino group (Table 3). This band evidences the condensation through the amino group of DsA with the 4-formyl-3-hydroxybenzoic acid. Furthermore, as a result of this condensation, the absence of the two bands at about 3470 and 3383 cm −1 , attributable to νa NH2 and νs NH2 modes, respectively, is clear in the spectrum of DsA. Figures S5 and S6 show the experimental and calculated  Table 1.   Table 1. The assignment of the more significant bands of the infrared spectra of H3L (Table 3) and DsA (Table 4) has also been done on the basis of the calculated spectrum using DFT methods. The IR spectrum of H3L shows a characteristic new strong sharp band at about 1613 cm −1 , which is attributable to the formation of the imino group (Table 3). This band evidences the condensation through the amino group of DsA with the 4-formyl-3-hydroxybenzoic acid. Furthermore, as a result of this condensation, the absence of the two bands at about 3470 and 3383 cm −1 , attributable to νa NH2 and νs NH2 modes, respectively, is clear in the spectrum of DsA. Figures S5 and S6 show the experimental and calculated   Table 1. The assignment of the more significant bands of the infrared spectra of H3L (Table 3) and DsA (Table 4) has also been done on the basis of the calculated spectrum using DFT methods. The IR spectrum of H3L shows a characteristic new strong sharp band at about 1613 cm −1 , which is attributable to the formation of the imino group (Table 3). This band evidences the condensation through the amino group of DsA with the 4-formyl-3-hydroxybenzoic acid. Furthermore, as a result of this condensation, the absence of the two bands at about 3470 and 3383 cm −1 , attributable to νa NH2 and νs NH2 modes, respectively, is clear in the spectrum of DsA. Figures S5 and S6 Table 1. The assignment of the more significant bands of the infrared spectra of H3L (Table 3) and DsA (Table 4) has also been done on the basis of the calculated spectrum using DFT methods. The IR spectrum of H3L shows a characteristic new strong sharp band at about 1613 cm −1 , which is attributable to the formation of the imino group (Table 3). This band evidences the condensation through the amino group of DsA with the 4-formyl-3-hydroxybenzoic acid. Furthermore, as a result of this condensation, the absence of the two bands at about 3470 and 3383 cm −1 , attributable to νa NH2 and νs NH2 modes, respectively, is clear in the spectrum of DsA. Figures S5 and S6 show the experimental and calculated   Table 1. The assignment of the more significant bands of the infrared spectra of H3L (Table 3) and DsA (Table 4) has also been done on the basis of the calculated spectrum using DFT methods. The IR spectrum of H3L shows a characteristic new strong sharp band at about 1613 cm −1 , which is attributable to the formation of the imino group (Table 3). This band evidences the condensation through the amino group of DsA with the 4-formyl-3-hydroxybenzoic acid. Furthermore, as a result of this condensation, the absence of the two bands at about 3470 and 3383 cm −1 , attributable to νa NH2 and νs NH2 modes, respectively, is clear in the spectrum of DsA. Figures S5 and S6 Table 1. The assignment of the more significant bands of the infrared spectra of H3L (Table 3) and DsA (Table 4) has also been done on the basis of the calculated spectrum using DFT methods. The IR spectrum of H3L shows a characteristic new strong sharp band at about 1613 cm −1 , which is attributable to the formation of the imino group (Table 3). This band evidences the condensation through the amino group of DsA with the 4-formyl-3-hydroxybenzoic acid. Furthermore, as a result of this condensation, the absence of the two bands at about 3470 and 3383 cm −1 , attributable to νa NH2 and νs NH2 modes, respectively, is clear in the spectrum of DsA. Figures S5 and S6 show the experimental and calculated   Table 1. The assignment of the more significant bands of the infrared spectra of H3L (Table 3) and DsA (Table 4) has also been done on the basis of the calculated spectrum using DFT methods. The IR spectrum of H3L shows a characteristic new strong sharp band at about 1613 cm −1 , which is attributable to the formation of the imino group (Table 3). This band evidences the condensation through the amino group of DsA with the 4-formyl-3-hydroxybenzoic acid. Furthermore, as a result of this condensation, the absence of the two bands at about 3470 and 3383 cm −1 , attributable to νa NH2 and νs NH2 modes, respectively, is clear in the spectrum of DsA. Figures S5 and S6 Table 1. The assignment of the more significant bands of the infrared spectra of H3L (Table 3) and DsA (Table 4) has also been done on the basis of the calculated spectrum using DFT methods. The IR spectrum of H3L shows a characteristic new strong sharp band at about 1613 cm −1 , which is attributable to the formation of the imino group (Table 3). This band evidences the condensation through the amino group of DsA with the 4-formyl-3-hydroxybenzoic acid. Furthermore, as a result of this condensation, the absence of the two bands at about 3470 and 3383 cm −1 , attributable to νa NH2 and νs NH2 modes, respectively, is clear in the spectrum of DsA. Figures S5 and S6 show the experimental and calculated   Table 1. The assignment of the more significant bands of the infrared spectra of H3L (Table 3) and DsA (Table 4) has also been done on the basis of the calculated spectrum using DFT methods. The IR spectrum of H3L shows a characteristic new strong sharp band at about 1613 cm −1 , which is attributable to the formation of the imino group (Table 3). This band evidences the condensation through the amino group of DsA with the 4-formyl-3-hydroxybenzoic acid. Furthermore, as a result of this condensation, the absence of the two bands at about 3470 and 3383 cm −1 , attributable to νa NH2 and νs NH2 modes, respectively, is clear in the spectrum of DsA. Figures S5 and S6 show the experimental and calculated (33%) a The contributions of the represented NTO pairs to the electronic transitions are given between parenthesis.

Experimental and Theoretical Studies on the Copper(II) Complex
Cu 2 L 2 (H 2 O) 4 was obtained at room temperature from the reaction between H 3 L and Cu(OAc) 2 , and alternatively, by the electrochemical oxidation of a copper anode immersed in a solution of H 3 L. It must be noted that the electrochemical procedure saves time and leads to higher yields (see experimental section). The copper(II) complex obtained from H 3 L has been investigated in solution by a combination of UV-Vis and fluorescence spectrometries. In the solid state, this complex has been investigated by elemental analysis and ATR-FTIR spectroscopy. Characterization details are shown in the experimental section. The results of the CHNS analyses showed that the copper(II) complex was obtained with a 1:1 metal:ligand stoichiometry. This finding is coherent with the formation of a sulfonamidobridged dimeric complex, as those reported by us for complexes of related Schiff base ligands [17,33,34].
The IR spectrum of this copper(II) complex showed a broad band centered at about 3360 cm −1 attributable to ν(OH), which evidenced the formation of the hydrated complex ( Figure S7). The slight shifts of ν(NC), ν as (SO 2 ), and ν s (SO 2 ) to lower wavenumbers (1-5 cm −1 ), as well as the absence of the band related to the phenolic COH bending (1412 cm −1 ), is a clear sign of the coordination of the Schiff base ligand through its O phenol , N imine , and N sulfonamide donor atoms. The observation of two bands attributable to ν as (COO − ) and ν s (COO − ) at about 1574 and 1393 cm −1 , respectively, instead of a single band at 1686 cm −1 indicates the deprotonation of the carboxylic group of the ligand in this complex. The net charge zero of the complex can be easily achieved, as the ligand units can adopt the zwitterionic form, in which the carboxylic groups are deprotonated, while the amine N atoms are protonated.
The electronic spectrum of the dimeric copper(II) complex ( Figure S8) shows three dominant signals around 219, 244, and 336 nm appearing ca. 5 nm blue-shifted with respect to the free ligand (224, 248, and 340 nm). The metal complexation of H 3 L with Cu 2+ is related to a new absorption band at about 410 nm, and which has been assigned to a ligand-to-metal charge transfer (LMTC).
The observation of a new absorption band at about 680 nm attributable to a d-d transition ( 2 E g ← 2 T 2g ) in the UV-Vis spectrum of the dinuclear copper(II) complex (Figure 5), is a sign of six-coordination [35]. Therefore, two water molecules must complete the coordination sphere around each metal ion, occupying the vacant positions left by the dianionic Schiff base ligand.
The geometrical structure of the copper(II) complex has been analyzed by means of theoretical methods. Based on experimental results (CHNS analyses, infrared, and electronic spectra) and previous studies [17,33,34], we were able to obtain the optimized geometry of a sulfonamido-bridged dimeric complex represented in Figure 6. The structure is symmetric, with an inversion center in the middle of the distorted square formed by the two sulfonamido bridges. A detailed analysis of the metal ion environment shows a six-coordination pattern comprising the expected N,N,O binding site of a ligand unit, a sulfonamido N-bridge of the neighboring ligand unit, and two monodentate water molecules completing its coordination sphere. A comparison of this optimized structure with a related dimeric copper(II) complex, which is also bridged by sulfonamide N atoms, showed a great similarity. The reported complex shows Cu-N distances in the range 1.969(4)-2.004(4) Å, and bond angles in the range 85.12-98.4 [34]. sition ( 2 Eg ← 2 T2g) in the UV-Vis spectrum of the dinuclear copper(II) is a sign of six-coordination [35]. Therefore, two water molecules mus dination sphere around each metal ion, occupying the vacant positio ionic Schiff base ligand. The geometrical structure of the copper(II) complex has been an theoretical methods. Based on experimental results (CHNS analyses tronic spectra) and previous studies [17,33,34], we were able to obtai ometry of a sulfonamido-bridged dimeric complex represented in Fig  is symmetric, with an inversion center in the middle of the distorted sq two sulfonamido bridges. A detailed analysis of the metal ion enviro coordination pattern comprising the expected N,N,O binding site of fonamido N-bridge of the neighboring ligand unit, and two monodent completing its coordination sphere. A comparison of this optimized lated dimeric copper(II) complex, which is also bridged by sulfonamid a great similarity. The reported complex shows Cu-N distances in 2.004(4) Å, and bond angles in the range 85.12-98.4 [34].  The geometrical structure of the copper(II) complex has been analyzed by means of theoretical methods. Based on experimental results (CHNS analyses, infrared, and electronic spectra) and previous studies [17,33,34], we were able to obtain the optimized geometry of a sulfonamido-bridged dimeric complex represented in Figure 6. The structure is symmetric, with an inversion center in the middle of the distorted square formed by the two sulfonamido bridges. A detailed analysis of the metal ion environment shows a sixcoordination pattern comprising the expected N,N,O binding site of a ligand unit, a sulfonamido N-bridge of the neighboring ligand unit, and two monodentate water molecules completing its coordination sphere. A comparison of this optimized structure with a related dimeric copper(II) complex, which is also bridged by sulfonamide N atoms, showed a great similarity. The reported complex shows Cu-N distances in the range 1.969(4)-2.004(4) Å, and bond angles in the range 85.12-98.4 [34]. Finally, with the aim of knowing the affinity of H3L to Cu 2+ ions at room temperature, the value of the binding constant (KB) has been determined ( Figure S9, KB = intercept/slope). For this purpose, we used the Benesi-Hildebrand equation [36], using data from the UV-Vis absorption (KB = 8.459 10 3 M −1 ), as well as from the fluorescence emission Finally, with the aim of knowing the affinity of H 3 L to Cu 2+ ions at room temperature, the value of the binding constant (K B ) has been determined ( Figure S9, K B = intercept/slope). For this purpose, we used the Benesi-Hildebrand equation [36], using data from the UV-Vis absorption (K B = 8.459 10 3 M −1 ), as well as from the fluorescence emission intensity (K B = 7.190 10 3 M −1 ).
Since the reaction of H 3 L with Cu 2+ ions resulted in a decrease of its fluorescence emission, we investigated the quenching mechanism with Stern-Volmer plots [37] (F 0 /F = 1 + K SV [Cu 2+ ]). Figure S10 shows that the value of the quenching constant decreases with increasing temperature, which is a sign of static quenching. It must be noted that for static quenching (τ 0 /τ = 1), the K SV value matches that of the K B for the complex formation (K SV = 9.010 10 3 M −1 ). Hence, K SV show similar values to K B for the reaction of H 3 L with Cu 2+ ions in a 1:1 molar ratio, supporting the accuracy of these measurements.

Investigations on the Interaction of H 3 L with CuO NPs
Spectrofluorimetry, Attenuated Total Reflection-Infrared spectroscopy (ATR-IR), Micro-X-ray Fluorescence (µ-XRF), Scanning Electron Microscopy-Energy Dispersive X-ray spectroscopy (SEM-EDX), and Density Functional Theory (DFT) were used to investigate the reaction product of H 3 L with CuO NPs.

Spectrofluorimetry
The interaction of H 3 L with CuO NPs leads to the emission of a yellow fluorescence when a solid sample of the reaction product is placed under a light source UV of 365 nm (Figure 7). Finally, with the aim of knowing the affinity of H3L to Cu 2+ ions at room temperature, the value of the binding constant (KB) has been determined ( Figure S9, KB = intercept/slope). For this purpose, we used the Benesi-Hildebrand equation [36], using data from the UV-Vis absorption (KB = 8.459 10 3 M −1 ), as well as from the fluorescence emission intensity (KB = 7.190 10 3 M −1 ).
Since the reaction of H3L with Cu 2+ ions resulted in a decrease of its fluorescence emission, we investigated the quenching mechanism with Stern-Volmer plots [37] (F0/F = 1 + KSV[Cu 2+ ]). Figure S10 shows that the value of the quenching constant decreases with increasing temperature, which is a sign of static quenching. It must be noted that for static quenching (τ0/τ = 1), the KSV value matches that of the KB for the complex formation (KSV = 9.010 10 3 M −1 ). Hence, KSV show similar values to KB for the reaction of H3L with Cu 2+ ions in a 1:1 molar ratio, supporting the accuracy of these measurements.

Investigations on the Interaction of H3L with CuO NPs
Spectrofluorimetry, Attenuated Total Reflection-Infrared spectroscopy (ATR-IR), Micro-X-ray Fluorescence (µ-XRF), Scanning Electron Microscopy-Energy Dispersive X-ray spectroscopy (SEM-EDX), and Density Functional Theory (DFT) were used to investigate the reaction product of H3L with CuO NPs.

Spectrofluorimetry
The interaction of H3L with CuO NPs leads to the emission of a yellow fluorescence when a solid sample of the reaction product is placed under a light source UV of 365 nm (Figure 7).  The fluorescence spectrum of the product obtained from the reaction between H 3 L and CuO NPs revealed a quenching of the fluorescence emission at 524 nm, in comparison with the fluorescence intensity of H 3 L measured under similar conditions. Stern-Volmer plots evidenced a static quenching mechanism, and therefore K SV provides a measure of the bonding affinity between H 3 L and CuO NPs (Figure 8, top). As expected, K SV for the reactions of H 3 L with CuO NPs and Cu 2+ ions have similar values (Table 5). Values of Gibbs free energy, which have been determined from K SV , indicate that the dimeric copper(II) complex (−22.19 kJ mol −1 ) and H 3 L-CuO NPs (−22.41 kJ mol −1 ) also show similar stabilities. the bonding affinity between H3L and CuO NPs (Figure 8, top). As expected, KS reactions of H3L with CuO NPs and Cu 2+ ions have similar values (Table 5). V Gibbs free energy, which have been determined from KSV, indicate that the dime per(II) complex (−22.19 kJ mol −1 ) and H3L-CuO NPs (−22.41 kJ mol −1 ) also show stabilities.   The number of interaction binding sites (n) were determined by the Scatchard Equation (1), where F 0 and F are the relative fluorescence of H 3 L in absence and presence of NMs, respectively, K B is the binding constant, and [Q] is the quencher concentration [38]. The Scatchard equation linear graph of H 3 L and NMs is shown in Figure 8 (bottom). The number of binding sites (n) is around 5 at 293 K, suggesting five binding sites for H 3 L on the CuO NPs surface. The interaction ratio increased slightly with rising temperature, probably because temperature favors the reactivity of superficial Cu 2+ ions.

Attenuated Total Reflection-Infrared Spectroscopy
The IR spectrum of H 3 L-CuO NPs clearly shows the absence of the band at about 3283 cm −1 attributable to ν OH/ν NH modes, which evidences the bideprotonation of the Schiff base ligand. Furthermore, a slight shift to lower wavenumbers of the bands corresponding to ν(NC), ν as (SO 2 ), and ν s (SO 2 ) can be observed The similarity between the ATR-IR spectra of CuO-H 3 L NPs and Cu 2 (HL) 2   The number of interaction binding sites (n) were determined by the Scatcha tion (1), where F0 and F are the relative fluorescence of H3L in absence and pre NMs, respectively, KB is the binding constant, and [Q] is the quencher concentra The Scatchard equation linear graph of H3L and NMs is shown in Figure 8 (bott number of binding sites (n) is around 5 at 293 K, suggesting five binding sites fo the CuO NPs surface. The interaction ratio increased slightly with rising temp probably because temperature favors the reactivity of superficial Cu 2+ ions. log − = log + log

Attenuated Total Reflection-Infrared Spectroscopy
The IR spectrum of H3L-CuO NPs clearly shows the absence of the band 3283 cm −1 attributable to ν OH/ν NH modes, which evidences the bideprotonati Schiff base ligand. Furthermore, a slight shift to lower wavenumbers of the ban sponding to ν(NC), νas(SO2), and νs(SO2) can be observed The similarity between IR spectra of CuO-H3L NPs and Cu2(HL)2(H2O)4·10H2O (Figure 9) indicates that t action between H3L and NMs surface occurs via metal-ligand coordination thr Nsulfonamide, Nimine, and Ophenol atoms. The observation of two bands attributable to ν and νs(COO − ) at about 1574 and 1391 cm −1 , respectively, is a clear sign of the depro of the carboxyl group present in the ligand, which is coherent with the formati carboxylate sodium salt.  The µ-XRF analysis indicates that the polycrystalline sample of CuO-H 3 L NPs contains copper and sulfur in a 30:1 ratio by weight, which is coherent with a 15:1 molar ratio (CuO:H 3 L). Moreover, µ-XRF mapping was used to explore the distribution of copper and sulfur within a sample of CuO-H 3 L NPs (Figure 10, inset). The homogeneity observed supports the effective interaction between the CuO NPs and H 3 L.  Figure 10 shows the peaks corresponding to the characteristic X-ray fluorescence emission of the chemical elements present in the reaction product of CuO NPs with H3L. The µ-XRF analysis indicates that the polycrystalline sample of CuO-H3L NPs contains copper and sulfur in a 30:1 ratio by weight, which is coherent with a 15:1 molar ratio (CuO:H3L). Moreover, µ-XRF mapping was used to explore the distribution of copper and sulfur within a sample of CuO-H3L NPs (Figure 10, inset). The homogeneity observed supports the effective interaction between the CuO NPs and H3L. Figure 10. XRF spectrum of the reaction product of CuO NPs with H3L. X axis: energy (keV), Y axis: intensity. Inset: XRF mapping of the reaction product between CuO NPs and H3L with elemental distribution of sulfur (red) and copper (green) on the sample surface.

Scanning Electron Microscopy-Energy Dispersive X-ray Spectroscopy
The SEM micrograph shown in Figure 11 show that the NPs are agglomerated to form clusters, which are made up of nanocrystals of sizes in the range of 70 to 40 nm, with a rather irregular shape.

Scanning Electron Microscopy-Energy Dispersive X-ray Spectroscopy
The SEM micrograph shown in Figure 11 show that the NPs are agglomerated to form clusters, which are made up of nanocrystals of sizes in the range of 70 to 40 nm, with a rather irregular shape.
A combination of SEM and EDX has been used to generate a "map" of element distributions in the raw sample. The EDX analysis shows the uniform presence of Cu, O, C, N, S, and Na elements with similar spatial distribution patterns, which supports the interaction of H 3 L (as carboxylate sodium salt) with CuO NPs.

DFT Studies
Several models have been considered to mimic the interaction between H 3 L and CuO NPs using DFT calculations [39] ( Figure S11). The largest models studied comprise a single unit of the ligand interacting with a sheet of 32 units of CuO. The geometry of these surfaces has been taken from one of its common crystallographic structures (available from CSD) and choosing that facet containing the largest number of copper atoms suitable to interact with the ligand, in this case, the (1,0,0) facet. The size of the largest models has been chosen to allow a complete interaction between H 3 L and the NPs surface. Particular attention has been devoted to establishing the preference of a dianionic (as phenolate and sulfonamide anion) or a monoanionic (only as phenolate) form of the ligand to interact with the superficial Cu 2+ ions on the NPs. Figure 12 shows the DFT models for the interaction of CuO NPs with these dianionic and monoanionic forms of the ligand.
A combination of SEM and EDX has been used to generate a "map" of element distributions in the raw sample. The EDX analysis shows the uniform presence of Cu, O, C, N, S, and Na elements with similar spatial distribution patterns, which supports the interaction of H3L (as carboxylate sodium salt) with CuO NPs.

DFT Studies
Several models have been considered to mimic the interaction between H3L and CuO NPs using DFT calculations [39] ( Figure S11). The largest models studied comprise a single unit of the ligand interacting with a sheet of 32 units of CuO. The geometry of these surfaces has been taken from one of its common crystallographic structures (available from CSD) and choosing that facet containing the largest number of copper atoms suitable to interact with the ligand, in this case, the (1,0,0) facet. The size of the largest models has been chosen to allow a complete interaction between H3L and the NPs surface. Particular attention has been devoted to establishing the preference of a dianionic (as phenolate and sulfonamide anion) or a monoanionic (only as phenolate) form of the ligand to interact with the superficial Cu 2+ ions on the NPs. Figure 12 shows the DFT models for the interaction of CuO NPs with these dianionic and monoanionic forms of the ligand. As a result of these CuO NP models, we can deduce that the interactions of the dianionic form of the ligand show higher stabilization energies (−273.4 kcal/mol) than those corresponding to the monoanionic ones (−170.0 kcal/mol). A similar trend is observed for all the diverse models considered, even with the smallest models employed here, which only consider the interaction between the ligand and a single CuO unit ( Figure S11). These As a result of these CuO NP models, we can deduce that the interactions of the dianionic form of the ligand show higher stabilization energies (−273.4 kcal/mol) than those corresponding to the monoanionic ones (−170.0 kcal/mol). A similar trend is observed for all the diverse models considered, even with the smallest models employed here, which only consider the interaction between the ligand and a single CuO unit ( Figure S11). These theoretical results confirm that the interaction between H 3 L and the NMs surface takes place through the O phenol , N imine , and N sulfonamide atoms of the ligand, as occurring for the dimeric copper(II) complex. In addition, the interactions of the superficial Cu 2+ ions with the oxygen atoms of the carboxylate and the sulfonamide groups appear to contribute to the stabilization of the systems studied.

Fluorescence Emission Studies on H 3 L upon Addition of CuO NPs
In order to check the usefulness of H 3 L as a chemosensor for the detection of CuO NPs, we selected a concentration of H 3 L of 126 ppb. The fluorescence emission varied linearly in the range 0-10.5 ppm with the gradual addition of CuO NPs. Concentration values of CuO NPs higher than those indicated would oversaturate the system. The decrease by over 90% in the fluorescence emission of ethanol-water solutions (in 80:20 v/v) of H 3 L (λ em = 540 nm) upon addition of CuO NPs (λ em = 538 nm) is shown in Figure 13. dimeric copper(II) complex. In addition, the interactions of the superficial the oxygen atoms of the carboxylate and the sulfonamide groups appear to the stabilization of the systems studied.

Fluorescence Emission Studies on H3L upon Addition of CuO NPs.
In order to check the usefulness of H3L as a chemosensor for the det NPs, we selected a concentration of H3L of 126 ppb. The fluorescence emiss early in the range 0-10.5 ppm with the gradual addition of CuO NPs. Conc ues of CuO NPs higher than those indicated would oversaturate the system by over 90% in the fluorescence emission of ethanol-water solutions (in 80 (λem = 540 nm) upon addition of CuO NPs (λem = 538 nm) is shown in Figur   Figure 13. Fluorescence emission spectra calibration curve of H3L in 80:20 ethanol:w pH 11, with gradual addition of CuO NPs (until the fluorescence intensity stops de = 126 ppb. NaOH 0.1 M was used to deprotonate H3L. Final volume of the solutio solution was sonicated for 1 min before being measured. λexc = 400 nm.
The limit of detection (LOD) and the limit of quantification (LOQ) of H expressed as LOD = 3SD/M and LOQ = 10SD/M, where SD is the standard d response and M is the slope of the calibration curve. LOD and LOQ were 33 ppm, respectively. It must be noted that these values are higher than tho reported by us for a pyrrole-based Schiff base ligand (13.83 µgL −1 and 46.05 tively) [17] and a phenol-based Schiff ligand (9.8 µgL −1 and 32.6 µgL −1 , resp The working range of H3L has been obtained using LOQ, as the minimum be measured, and the highest value of [NMs] at which linearity is still main maximum value. Therefore, the working range for detection of CuO NPs w ppb was 1.1-9.5 ppm. In view of the affinity of H3L to CuO NPs, we studied the selectivity of for the cited NPs. As a criterion for interference, a ±10% variation of the av cence intensity was used. The selectivity of H3L as a probe for CuO NPs wa presence of common metal ions in water, such as Na + , K + , Mg 2+ , Ca 2+ , Fe 3+ , ions, in the same concentration as CuO NPs (Figure 14). These results sh cited species do not significantly interfere with the detection of CuO NPs. A tolerated in concentrations of at least 4 ppm without interference. Each solution was sonicated for 1 min before being measured. λ exc = 400 nm.
The limit of detection (LOD) and the limit of quantification (LOQ) of H 3 L have been expressed as LOD = 3SD/M and LOQ = 10SD/M, where SD is the standard deviation of the response and M is the slope of the calibration curve. LOD and LOQ were 330 ppb and 1.09 ppm, respectively. It must be noted that these values are higher than those previously reported by us for a pyrrole-based Schiff base ligand (13.83 µgL −1 and 46.05 µgL −1 , respectively) [17] and a phenol-based Schiff ligand (9.8 µgL −1 and 32.6 µgL −1 , respectively) [18]. The working range of H 3 L has been obtained using LOQ, as the minimum value that can be measured, and the highest value of [NMs] at which linearity is still maintained, as the maximum value. Therefore, the working range for detection of CuO NPs with [H 3 L] = 126 ppb was 1.1-9.5 ppm.
In view of the affinity of H 3 L to CuO NPs, we studied the selectivity of H 3 L as a probe for the cited NPs. As a criterion for interference, a ±10% variation of the average fluorescence intensity was used. The selectivity of H 3 L as a probe for CuO NPs was tested in the presence of common metal ions in water, such as Na + , K + , Mg 2+ , Ca 2+ , Fe 3+ , and Al 3+ metal ions, in the same concentration as CuO NPs (Figure 14). These results showed that the cited species do not significantly interfere with the detection of CuO NPs. Al 3+ ions can be tolerated in concentrations of at least 4 ppm without interference.

Materials and Methods
All starting materials and reagents were commercially available, and they were used without further purification. 1 H NMR spectra (400 MHz) were measured in deuterated solvents using a Varian Inova 400 Spectrometer. J values are given in Hertz. The IR spectra were measured in the range 4000 to 400 cm −1 wavenumber using a FTIR spectrometer PerkinElmer Spectrum Two coupled with Platinum Diamond ATR, which consists of a diamond disc as an internal reflection element. Elemental analyses were performed on an elemental analyzer FISONS model EA 1108. Fluorescence emission studies were performed on a Shimadzu RF-600 Spectro Fluorophotometer. X-ray fluorescence was measured under vacuum (19.6 mbar) using an M4 TORNADO system (BRUKER), with the Rh tube, operating at 50kV and 200 µA. The fluorescence maps were collected for a total time of 2 ms per pixel, with a pixel size of 20 µm. Scanning Electron Microscopy (SEM) was applied to investigate the size and morphology of the NPs, and it was carried out under a ZEISS FESEM ULTRA Plus with EDX microanalysis.

Sample Preparation
Samples for investigations on the interaction of H3L with CuO NPs were obtained by stirring an ethanol solution of H3L at pH 9, and a suspension of NPs in a 2:1 molar ratio, at room temperature for about 15 min. Then, the sample was air-dried for about 48 h.

Crystal Structure Analysis Data
Diffraction data for prismatic crystals corresponding to DsA and H3L were collected at 100(2) K, using graphite-monochromatized Mo-Kα radiation (λ = 0.71073 Å) from a fine focus sealed tube, using a Bruker D8 Venture Photon III-14 diffractometer. Data were routinely processed and corrected for Lorentz and polarization effects. Multi-scan absorption corrections were performed using the SADABS routine [40]. The structure was solved by standard direct methods using SHELXT [41], and then refined by full matrix least squares on F 2 by using the SHELXL program [42]. All non-hydrogen atoms were anisotropically refined. Hydrogen atoms were mostly included in the structure factor calculation in geometrically idealized positions, with thermal parameters depending on the parent atom,

Materials and Methods
All starting materials and reagents were commercially available, and they were used without further purification. 1 H NMR spectra (400 MHz) were measured in deuterated solvents using a Varian Inova 400 Spectrometer. J values are given in Hertz. The IR spectra were measured in the range 4000 to 400 cm −1 wavenumber using a FTIR spectrometer PerkinElmer Spectrum Two coupled with Platinum Diamond ATR, which consists of a diamond disc as an internal reflection element. Elemental analyses were performed on an elemental analyzer FISONS model EA 1108. Fluorescence emission studies were performed on a Shimadzu RF-600 Spectro Fluorophotometer. X-ray fluorescence was measured under vacuum (19.6 mbar) using an M4 TORNADO system (BRUKER), with the Rh tube, operating at 50 kV and 200 µA. The fluorescence maps were collected for a total time of 2 ms per pixel, with a pixel size of 20 µm. Scanning Electron Microscopy (SEM) was applied to investigate the size and morphology of the NPs, and it was carried out under a ZEISS FESEM ULTRA Plus with EDX microanalysis.

Sample Preparation
Samples for investigations on the interaction of H 3 L with CuO NPs were obtained by stirring an ethanol solution of H 3 L at pH 9, and a suspension of NPs in a 2:1 molar ratio, at room temperature for about 15 min. Then, the sample was air-dried for about 48 h.

Crystal Structure Analysis Data
Diffraction data for prismatic crystals corresponding to DsA and H 3 L were collected at 100(2) K, using graphite-monochromatized Mo-Kα radiation (λ = 0.71073 Å) from a fine focus sealed tube, using a Bruker D8 Venture Photon III-14 diffractometer. Data were routinely processed and corrected for Lorentz and polarization effects. Multi-scan absorption corrections were performed using the SADABS routine [40]. The structure was solved by standard direct methods using SHELXT [41], and then refined by full matrix least squares on F 2 by using the SHELXL program [42]. All non-hydrogen atoms were anisotropically refined. Hydrogen atoms were mostly included in the structure factor calculation in geometrically idealized positions, with thermal parameters depending on the parent atom, by using a riding model. H atoms potentially involved in H-bonding were localized in Fourier maps and then refined with a thermal parameter depending on the parent N atom. More details of the refinement, as well as crystal data are collected in Table  S1 of the ESI. CCDC 2,121,699 and 2,121,700 contain the supplementary crystallographic data for this paper. These data can be obtained free of charge from the Cambridge Crystallographic Data Centre via www.ccdc.cam.ac.uk/data_request/cif (accessed on 20 September 2022)

Computational Methods
Theoretical calculations on DsA, H 3 L, copper(II) complex, and ligand-NPs models have been performed employing density functional theory (DFT) methods, as implemented in Gaussian 09 [39]. Thus, M062X/6-31G* geometrical optimization of different starting structures of DsA, H 3 L, and its copper(II) complex have been performed. The most stable conformers found have been characterized by calculation of their vibrational frequencies.
For DsA and H 3 L, we tested the idea that the M062X/6-31G* level provided results of comparable quality to those obtained with a more extended basis set as 6-311++G** (not included for brevity), but with a significantly smaller computational cost. For comparison with the experiment results, the theoretical IR spectra of DsA and H 3 L have been obtained using the harmonic vibrational frequencies scaled by 0.947 and a FWHM of 4 cm −1 .
The electronic spectra of DsA and H 3 L have been obtained employing the M062X/6-31G* level, and the influence of the solvent has been included using the PCM method and ε = 24.852 for ethanol. The main electronic transitions have been analyzed by computing their natural transition orbitals, which provide a picture of the orbitals involved in any electronic transition, which is easier to interpret [43].
Models of different size have been used to evaluate the interaction of the ligand with CuO NPs. The geometry for the smallest model, comprising only one H 3 L unit and a single CuO unit, has been optimized without restrictions. In the remaining cases, the geometries of the ligand over the surface have been freely optimized, keeping a frozen structure for the metal obtained from crystallographic data.
Interaction energies, E int , for these systems have been evaluated as E int = E complex − (E H3L + E metal model ), where the E complex is the total energy of the interacting ligand-NPs, E H3L is the total energy of the monoanionic or dianionic form of the H 3 L ligand, and E metal is the energy of the different metal models indicated above. These interaction energies have been evaluated by means of the NEDA method included in the NBO7.0 program [44], so that the basis set superposition error is corrected by means of the counterpoise method.

DsA
The compound N-(2-aminobenzyl)-5-(dimethylamino)naphthalene-1-sulfonamide was prepared following a synthetic procedure previously used with imidazolopyridine compounds [45], but with some modifications. In a 250 mL round flask, 2.35 mmol (0.6339 g) of dansyl chloride was dissolved in 60 mL of CH 2 Cl 2 . In a second flask, 2.59 mmol (0.3162 g) of the 2-aminobenzylamine was dissolved in 40mL of CH 2 Cl 2 , at which 2.59 mmol (0.2631 g) of triethylamine was added. The content of the second flask was added to the dansyl chloride solution, giving a dark yellow solution which was stirred at room temperature for 18 h. During this time period, the solution changed its color from dark to light yellow. Evaporation under reduced pression of the solution to a third of its volume, and the addition of diethyl ether until the observation of turbidity, led to the formation of a light brown precipitate, which is a small unreacted 2-aminobenzylamine fraction. Filtration of the mixture led to the separation of 2-aminobenzylamine as a powdery solid. Evaporation to dryness under vacuum of the filtrate led to the precipitation of a light yellow powdery solid, which was washed with deionized water to separate triethylamine hydrochloride. The yellow solid was filtrated and air-dried to be characterized as DsA.