Selective fluorometric detection of F⁻ and Zn(II) ions by a N, O coordinating sensor and naked eye detection of Cu(II) ions in mixed-aqueous solution

Abstract

Through click chemistry, salicylaldehyde and fluorene groups have been explored to recognize anions through O–H⋯A– hydrogen-bonding complexation followed by an ESIPT mechanism and cation sensing via CHEF and CHEQ mechanisms for Zn²⁺ and Cu²⁺ metal ions respectively. Herein, we demonstrate evidence of the interactions between the O–H bond and F⁻ which were studied by fluorescence spectroscopic, UV-Vis spectroscopic, lifetime spectroscopic, ¹H NMR spectroscopic titrations and theoretical treatment. This sensor could simultaneously detect two biologically important metal ions (Cu²⁺ and Zn²⁺) through colorimetric methods in mixed aqueous solution. The 2 : 1 binding stoichiometry ratio of the ligand and metal in the complexes was established by UV-Vis, fluorescence, ¹H NMR and ESI-MS spectroscopy, and their corresponding association constants, K_assoc observed at 8.13 × 10⁴ and 5.12 × 10⁶ M⁻¹ corresponds to Zn²⁺ and Cu²⁺ metal ions in aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2. In addition, the electronic structures and photo physical properties of the ligand and complexes were calculated by DFT and time-dependent DFT (TDDFT) methods.

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Introduction

The construction of sensory molecules for the recognition and sensing of biologically and chemically important anions and cations is a forefront research topic in chemistry.¹ Among various important anionic analytes, the fluoride ion is one of the most significant due to its role in dental care² and treatment of osteoporosis.³ In particular, the detection of anions by a fluorescence readout has attracted great attention in light of the advantages of fluorescence sensing including easy operation and high sensitivity.⁴

Hydrogen-bond donors such as pyrrole/calixpyrrole, (thio)urea, guanidinium, azophenol, dipyrrolylquinoxalines, indolocarbazoles, (di)amino, amide, and (benzo)imidazole usually act as anion binding sites.⁵ Fluoride is a strong hydrogen bond acceptor and has a high affinity for silicon and boron. There are many reported organoboron compounds which have also been shown to be effective fluorescent sensors for the selective detection of fluoride ions.⁶ The synthesis of such a sensor is not an easy task. Herein we have very easily synthesized a HL sensor containing a phenolic proton with a fluorene moiety which is highlighted with excellent photophysical and photochemical properties such as large extinction coefficients, good photostability with blue emission light in the presence of fluoride anions under a UV lamp. The blue emission is based on the deprotonation of the O–H bond, displaying “off–on” changes in its fluorescence emission upon the addition of fluoride ions followed by the ESIPT⁷ mechanism.

Amongst the various important heavy metal ions in the human body, zinc is the second most abundant metal ion and thus, detection of Zn²⁺ is shown to have considerable interest in chemistry. As we know, the disorders of zinc metabolism are closely associated with many severe neurological diseases such as Alzheimer’s disease (AD) and Parkinson’s disease.⁸ Therefore, the detection of Zn²⁺ in biological samples is of significant interest and importance. Our fluorescent probe can bind selectively to Zn²⁺ in the presence of other important biological metal ions such as Na⁺, K⁺ etc., resulting in a change in the fluorescence intensity with green emission light present under a UV lamp. This change in emission intensity arises due to a decrease in the rate of the PET mechanism acting on the lone pair electrons of the O and N atoms of the sensor as well as increasing the CHEF⁹ process during the complexation. Besides some available Zn²⁺ sensors¹⁰ have difficulty in distinguishing Zn²⁺ from Cd²⁺, since cadmium is in the same group of the periodic table and has similar properties to Zn²⁺. But our designed sensor is a highly selective and sensitive fluorescence sensor for Zn²⁺ detection without any interference from other metal ions, especially Cd²⁺, which is one of the most important objectives.

Copper is third in abundance (after iron and zinc) among the essential transition metal ions in the human body and plays a key role in a multitude of biological processes, mostly in the active center of enzymes where it has a major role in electron transfer reactions.¹¹ Colorimetric sensors have attracted a lot of attention for the “naked-eye” detection properties in a straightforward and inexpensive manner, offering qualitative and quantitative information without using expensive equipment. This sensor changes color from colorless to intense yellow in the presence of a very low concentration of copper ions (10 ppm) in addition to the emission intensity of the sensor being almost completely quenched in the presence of copper ions following the CHEQ mechanism. The World Health Organization (WHO) has recommended the maximum limit of copper in drinking water to be 30 μM.¹²

Some aspects of these unforeseen findings were rationalized by DFT and TDDFT calculations. We also calculated and analyzed the singlet ground state natural transition orbitals (NTOs) derived from the TDDFT results and compared them with the ground state molecular orbitals (MOs) obtained from the DFT calculations. The computational modeling of the NMR parameters is also of abiding interest, and such DFT calculations have emerged as a promising approach for the prediction of nuclear shielding and coupling constants of NMR active nuclei.¹³ Thus, we computed the proton NMR chemical shifts and also the ¹H–¹H spin–spin coupling constants using the gauge-independent atomic orbital (GIAO)-DFT method, which was aimed at providing definitive characterization of the sensor. The solid-state structure of the complex stays in solution, as revealed by the combined experimental and theoretical NMR spectral analyses.

Experimental section

Materials

The transition metal salts used in the present investigation are as follows: Zn(ClO₄)₂·6H₂O, Cr(ClO₄)₃·6H₂O, Cu(ClO₄)₂·6H₂O, Mn(ClO₄)₂·6H₂O, Ni(ClO₄)₂·6H₂O, Co(ClO₄)₂·6H₂O and Cd(ClO₄)₂·6H₂O. The metal salts were procured locally and were used as received. Perchlorate salts were preferred because of the low coordinating ability of their anionic counterparts. Tetrabutylammonium (TBA) salts of the respective anions ([A] = F⁻, Cl⁻, Br⁻, I⁻, PF₆⁻, OAc⁻, Pi⁻, PPi⁻ and CN⁻ “A” stands for anion) were used as received from Sigma-Aldrich, USA. 2-[4-(2-Hydroxyethyl)-1-piperazinyl]ethanesulfonic acid (HEPES), potassium chloride and hydrochloric acid were purchased from Cyno-chem, India. All other solvents and reagents such as dichloromethane (DCM), methanol (MeOH), hexane (HEX), toluene and dimethylformamide (DMF) were of spectroscopic grade (Spectrochem, India) and used after appropriate distillation. The solvents were found to be free from impurities and appeared transparent in the spectral region of interest. The purity was also verified by recording the emission spectra in the studied spectral region. CDCl₃ and CD₃CN for NMR experiments were used as received from Sigma-Aldrich, USA.

Caution! Perchlorate salts are highly explosive, and should be handled with care and in small amounts.

Physical measurements

UV-vis spectra were recorded on a Perkin-Elmer LAMBDA 25 spectrophotometer. IR spectra were obtained with a Perkin-Elmer L-0100 spectrophotometer. Electrospray ionization mass spectrometry (ESI-MS) measurements were carried out on a Micromass Qtof YA 263 mass spectrometer. The molar conductivity was determined using a Systronics Conductivity Meter 304 in acetonitrile solution at room temperature. Elemental analyses (C, H, N) were performed using a Perkin-Elmer 2400 series II analyzer. EPR spectra were recorded for samples in standard quartz EPR tubes using a Varian E-109C spectrometer at the X-band and electrochemical measurements were recorded with a CHI 620A electrochemical analyzer using a platinum electrode under a nitrogen atmosphere. The emission data was collected using a Perkin-Elmer LS 55 fluorescence spectrometer. Quantum yields of the ligand and complex were determined in freeze–pump–thaw-degassed solutions of the complexes with a relative method using quinine sulfate in the same solvent as the standard¹⁴ and calculated using eqn (1),^15a where Φ_r and Φ_std are the quantum yields of the unknown and standard samples [Φ_std = 0.54 (at 298 K) in methanol at λ_ex = 350 nm], A_r and A_std (<0.1) are the solution absorbance values at the excitation wavelength (λ_ex), I_r and I_std are the integrated emission intensities, and η_r and η_std are the refractive indices of the solvents.

The fluorescence decay data were collected with a Hamamatsu MCP photomultiplier (R3809) and were analyzed using IBH DAS6 software. The observed decays of the complex fitted well with a bi-exponential function as in eqn (2) and (3), where τ₁ and τ₂ are the fluorescence lifetimes and α is the pre-exponential factor. For the fits, reduced χ² values are around one, and the distribution of weighted residuals was random among the data channels. τ_f is the mean fluorescence lifetime (the meaning of the symbol is standard).^15b All absorption and emission spectral measurements were performed with appropriate background corrections and with freshly prepared solutions only.

I(t) = [α₁ exp(−t/τ₁) − α₂ exp(−t/τ₂)] (2)

τ_f = α₁τ₁ + α₂τ₂ (3)

Computational details

The geometrical structures of the ligand (HL) in its singlet ground (S₀) and excited (S₁) states were optimized by the DFT¹⁶ and time-dependent DFT (TDDFT)¹⁷ methods with a B3LYP exchange correlation functional¹⁸ approach associated with the conductor-like polarizable continuum model (CPCM).¹⁹ On the basis of the optimized ground (S₀) and excited (S₁) geometries of the ligand, the absorption and the emission properties in solution state were calculated. We computed the lowest 40 singlet–singlet transitions in the ground state (S₀).

To calculate the stability of the keto tautomeric form over the enol form for the ligand (HL) in the excited S₁ state we performed the potential energy scan according to the “distinguished coordinate approach”²⁰ i.e. by specifying a reaction coordinate (in the present case it is the coordinate for the translocation of the proton from O_donor to N_acceptor, i.e. elongation of the O_donor–H bond axis) along which the energy change was observed. For the excited S₁ state all of the other degrees of freedom were relaxed without imposing any symmetry constraints.

The geometrical structures of complexes 1, [ZnL₂], and 2, [CuL₂], in their singlet ground (S₀) state were fully optimized in solution phase without any symmetry constraints. The vibrational frequency calculation was also performed for all the complexes to ensure that the optimized geometries represent the local minima and there are only positive eigen values. Finally to understand the nature of excited states involved in absorption processes, natural transition orbital (NTO) analysis was performed for all complexes.²¹

The effective core potential (ECP) approximation of Hay and Wadt was used for describing the (1s²2s²2p⁶) core electron for zinc and copper whereas the associated “double-ξ” quality basis set LANL2DZ was used for the valence shell.²² For H atoms, we used a 6-311+G basis set and for C, N and O atoms, we employed 6-311+G** as the basis set for the optimization. The calculated electronic density plots for frontier molecular orbitals were prepared using the GaussView 5.0 software. All the calculations were performed with the Gaussian 09W software package.²³ The GaussSum 2.1 program²⁴ was used to calculate the molecular orbital contributions from groups or atoms.

In addition, the ¹H NMR properties of complex 1 were calculated using the magnetic field perturbation method with the GIAO algorithm²⁵ and the NMR = spin–spin keyword incorporated in the Gaussian 09W program. The relative chemical shift of a given nucleus X in the molecule was defined as δ^calc_X [ppm] = σ^ref_X − σ^calc_X where TMS was used as a reference molecule optimized at the same level of theory.²⁶ In order to account for the solvent effect, we used the integral equation-formalism polarizable continuum model (IEFPCM) method.²⁷

Crystallographic studies

The single crystal suitable for X-ray crystallographic analysis of the ligand HL was obtained by slow evaporation of methanol solution. The X-ray intensity data were collected using a Bruker AXS SMART APEX CCD diffractometer (Mo K_α, λ = 0.71073 Å) at 293 K. The detector was placed at a distance of 6.03 cm from the crystal. A total of 606 frames were collected with a scan width of 0.3° in different settings of Φ. The data were reduced with SAINTPLUS²⁸ and empirical absorption correction was applied using the SADABS package.²⁸ The non-hydrogen atoms were emerged from successive Fourier syntheses. The structures were refined with a full matrix least-square procedure on F². All non-hydrogen atoms were refined anisotropically. All calculations were performed using the SHELXTL V 6.14 program package.²⁹ Molecular structure plots were drawn using the (ORTEP)^30a software. Relevant crystal data is given in Table 1.

Table 1 Crystal data and structure refinement parameters for HL

	HL
a R1 = ∑∣∣F0∣ − ∣Fc∣∣/∑∣F0∣.b wR2 = [∑[w(F0^2 − Fc^2)^2]/∑[w(F0^2)^2]]^1/2.
Formula	C20H15NO
Mr	285.33
Crystal system	Monoclinic
Space group	P21/c
a/Å	16.0198(4)
b/Å	8.2129(2)
c/Å	11.1102(3)
α/°	90
β/°	97.047(1)
γ/°	90
V/Å^3	1450.72(6)
Z	4
Dcalcd/mg m^−3	1.306
μ/mm^−1	0.080
θ/°	2.6–27.6
T/K	293(2)
Reflns collected	3366
R1^a, wR2^b [I > 2σ(I)]	0.0547, 0.1671
GOF on F^2	1.03

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Synthesis of ligand (HL)

To a methanolic solution (20 mL) of salicylaldehyde (245 mg, 2 mmol), 2-aminofluorene (365 mg, 2 mmol) was added and the reaction mixture was refluxed in a water bath for 2 hours. After cooling to room temperature, the solvent was removed under reduced pressure. The crude mass was then subjected to column chromatography on a silica gel column (60–120 mesh). A light yellow band was eluted using 5% ethyl acetate in hexane solution. A light yellow colored solid was obtained after the removal of the solvent under reduced pressure to afford the desired ligand (HL). Yield: 419 mg (74%). Elemental anal. calcd for C₂₀H₁₅NO: C, 84.19; H, 5.30; N, 4.91. Found: C, 84.30; H, 5.25; N, 4.82. ¹H NMR {300 MHz, CD₃Cl, δ (ppm), J (Hz)}: 13.41 (H1, bs), 8.71 (H7, s), 7.83–5.93 (11H, ArH), 3.95 (2H, H18 s). ESI-MS (CH₃CN): m/z calcd 286.1232, found: 286.1296 (100%). (HL + H)⁺. IR (KBr, cm⁻¹): ν(C–H) 3443; ν(C=N) 1615; ν(C=C) 1446, 1090 and 620.

Synthesis of the complexes

[Zn(L)₂], 1. Two equivalents of HL (57 mg, 0.2 mmol) was dissolved in 20 mL of MeOH and reacted with a one equivalent amount of zinc acetate dihydrate (22 mg, 0.1 mmol). The mixture was stirred at an ambient temperature for 4 hours followed by the addition of triethylamine (21 mg, 0.2 mmol), the color of the solution changed from light yellow to an intense greenish yellow emission color. The solution was kept for slow evaporation which yielded a dark yellow crude product in a good amount. Yield: 49 mg (76%). Elemental anal. calcd for C₄₀H₂₈N₂O₂Zn: C, 75.77; H, 4.45; N, 4.42. Found: C, 75.84; H, 4.37; N, 4.40. ¹H NMR {300 MHz, CD₃CN, δ (ppm), J (Hz)}: 9.19 (H7, s), 7.90–5.96 (11H, ArH), 3.95 and 3.88 (2H, H18 bs). ESI-MS (CH₃CN): m/z calcd 655.1340, found: 655.1146 (100%) (M + Na)⁺. IR (KBr, cm⁻¹): ν(C–H) 3463, 3029 and 2917; ν(C=N) 1598; ν(C=C) 1436 and 1410. Molar conductance, Λ_M: (CH₃CN) 8 Ω⁻¹ cm² mol⁻¹.

[Cu(L)₂], 2. Two equivalents of HL (57 mg, 0.2 mmol) was dissolved in 20 mL of methanol and reacted with a one equivalent amount of copper acetate monohydrate (20 mg, 0.1 mmol). The mixture was stirred at an ambient temperature for 4 hours followed by the addition of triethylamine (21 mg, 0.2 mmol), the color of the solution changed from light yellow to a yellowish green color. The solution was kept for slow evaporation which yielded a yellow crude product in a good amount. Yield: 52 mg (79%). Elemental anal. calcd for C₄₀H₂₈N₂O₂Cu: C, 75.99; H, 4.46; N, 4.43. Found: C, 76.04; H, 4.47; N, 4.40. ESI-MS (CH₃CN): m/z calcd 632.1525, found: 632.1911 (100%) (M + H)⁺. IR (KBr, cm⁻¹): ν(C–H) 3463, 3035 and 2904; ν(C=N) 1588; ν(C=C) 1443 and 1390. Molar conductance, Λ_M: (CH₃CN solution) 12 Ω⁻¹ cm² mol⁻¹. E_pc (Cu^II/Cu^I): −0.85 V (irr).

Results and discussion

Synthesis

The reaction of Zn(OAc)₂·2H₂O and Cu(OAc)₂·H₂O with HL in a ratio of 1 : 2 in methanol at room temperature in air produces complexes with compositions [Zn(L)₂] 1 and [Cu(L)₂] 2 in excellent yields. In the complexes, the ligand HL binds as a monoanionic N, O coordinating bidentate ligand and forms a neutral complex. The elemental analysis, ¹H NMR, IR spectroscopic and mass spectroscopic measurements confirmed the formation of the synthesized complexes. A schematic representation for the synthesis of the complexes is given in Scheme 1.

Scheme 1

Both the ligand and complexes were diluted with acetonitrile for mass spectrometry. Mass spectral analysis in the positive ion mode showed major peaks at m/z (%) = 286.1296 (100), 655.1146 and 632.1911 (100), which were assigned to the monocationic protonated ligand [HL¹ + H]⁺, [Zn + Na]⁺ and [Cu + H]⁺, the respective spectra are given in ESI (Fig. S1–S3†).

The ligand and complex 1 are diamagnetic and display well resolved ¹H NMR spectra. For the complex, the NMR spectra were measured in CD₃CN solution, whereas, HL was dissolved in CDCl₃ solution. In HL the phenolic proton (H1) was observed as a singlet at a greater deshielded field near 13.41 ppm, whereas the chemical shift of 8.71 ppm was associated with the proton (H7) adjacent to the azomethine group of the coordinated ligand. The aromatic proton span for the ligand was observed in the range 7.83–5.93 ppm. The methylene proton (H18) showed a sharp singlet at 3.95 ppm. As shown in Fig. 1, the peak that appears at a high deshielded field for the phenolic proton (H1) in HL completely disappears in the ¹H NMR spectrum of complex 1. On the other hand, the azo methane proton (H7) undergoes down field shifting to 9.19 ppm and the methylene proton (H18) shows splitting due to the loss of symmetry during complexation and undergoes up field shifting to 3.95 and 3.88 respectively. The correlation between the experimental and calculated ¹H NMR chemical shifts of 1 is shown in ESI Fig. S4† as a representative case.

Fig. 1 The ¹H NMR spectra of HL and 1 with changes in the spectral nature of both the aromatic and aliphatic regions given.

IR spectra of the ligand and complexes were recorded as KBr disks. HL showed a broad band near at 3500 cm⁻¹ which arises due to asymmetric vibration of both the O–H and C_sp2–H groups. In both complexes this band appears as a very small peak due to absence of the OH group during complexation (ESI Fig. S5†). The band at 1615 cm⁻¹ due to the azomethine group of the Schiff base underwent a shift to a lower frequency (1580–1590 cm⁻¹) after complexation, indicating the coordination of the azomethine nitrogen to the metal atom and this can be explained by the donation of electrons from nitrogen to the metal atom. The bands in the region of 1500–1350 cm⁻¹ are consistent with the skeletal vibration of the aromatic system. The nature of metal–ligand bonding is confirmed by the newly formed band at ∼430 cm⁻¹ in the spectrum of the complexes which is tentatively assigned to M–N vibration.

Cyclic voltammetry was performed for complex 2 in methanol solution at room temperature under a nitrogen atmosphere, with tetraethylammonium perchlorate (TEAP) as the supporting electrolyte using a Pt electrode as the working electrode. The potentials were referenced to the saturated calomel electrode (SCE) without junction correction. The complex also showed a quasi-reversible peak in the negative region, characteristic of the Cu(II) to Cu(I) couple at E_pc = −0.85 V, with the associated irreversible anodic peak at E_pa = 0.82 V which is diagnostic of a quasireversible one-electron transfer controlled by diffusion following the equation Cu(III)L + e → Cu(II)L.^30b

Complex 1 is EPR silent both in solution and in the solid state whereas complex 2 is EPR active. The EPR spectrum of complex 2 in the solid state is given in Fig. 2. The spectrum is characterized by an axial pattern with the features g_| > 2.1 > g_⊥ > 2 (g_| = 2.282, g_⊥ = 2.05). Fig. 2c depicts the spin density plot of complex 2. The electron density of the coordinating atom (N, O) of the sensor is distributed over the empty d-orbital of the copper metal ion. This type of distribution of electron density of a metal center is the origin of an EPR active complex.

Fig. 2 (a) Cyclic voltammograms of a ∼10⁻³ M solution of 2 in methanol solution; (b) X-band EPR spectra for 2 in the solid state at room temperature. Instrument settings: microwave frequency, 9.1 GHz; microwave power, 30 dB; modulation frequency, 100 kHz; modulation amplitude, 12.5 G; sweep centre, 3200 G; (c) spin density plot of complex 2 in the ground state.

Crystal structures

The molecular structure of HL was determined by single-crystal X-ray diffractometry. The ligand crystallizes in the P2₁/c space group (Fig. 3). The imine C=N distance is 1.268 Å. The internuclear distance between the imine nitrogen atom and phenolic oxygen atom is 2.587 Å. There is a possibility of intramolecular hydrogen bonding between the above stated atoms. In the crystal packing, several non-covalent interactions such as hydrogen bonding, C–H/π, π⋯π were found. The C–H/π interactions are generated by interactions of H2, H16 and H26 with the centroid of the rings Cg(1), Cg(3) and Cg(4) respectively. The π–π stacking interaction is found to be involved in the Cg(1)–Cg(4) rings with a centroid-to-centroid distance of 3.273(18) Å. An extended 2D network perspective view is given in ESI Fig. S6.†

Fig. 3 ORTEP plot and atom labeling scheme of HL. Hydrogen atoms have been omitted for clarity.

Ligand photophysical study

The UV-vis absorption spectrum of HL shows a well resolved peak around 350 nm in different solvents at room temperature (ESI Fig. S7†). The band is attributed mainly to the n–π* electronic transition as the polarity of the solvent increases, the band gradually undergoes hypsochromic (blue) shift. The blue shift of the ligand in the UV-vis spectrum with increasing solvent polarity was verified by theoretical calculation. We calculated TDDFT and GaussSum 2.0 analyses on the basis of the optimized geometry of HL using different polarity solvents. In every case it was observed that the lowest lying distinguishable singlet → singlet absorption band near 350 nm can be attributed to ILCT character. It was further observed that the energy gap between the HOMO and LUMO in the S₀ → S₁ transition gradually increased with increasing solvent polarity (Table 2). The fluorescence behavior of the ligand discussed in this present work was studied at room temperature in CH₃OH solution. The ligand displayed red shifted emission spectra at 512 nm with large Stokes shift. In solution, at room temperature, the ligand was a weak emitter with low quantum yield (Φ_F = 0.011) and a lifetime (τ) of 7.2 ns. The luminescence parameters for the ligand are listed in Table 3.

Table 2 Main calculated optical transition for HL with composition in terms of the molecular orbital contribution of the transition, vertical excitation energies and oscillator strength in solvents of different polarity

Solvent	Hexane	Toluene	DCM	CH3OH	DMF
Transition (S0 → S1)	^1ILCT	^1ILCT	^1ILCT	^1ILCT	^1ILCT
Oscillator strength (f)	0.8325	0.5856	0.5877	0.9725	0.6019
λtheo (nm)	362.19	357.78	352.01	349.40	342.89
λexp (nm)	358	354	350	347	344
Excitation energy (eV)	3.6987	3.7231	3.7420	3.7679	3.7756

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Table 3 Selected parameters for the vertical excitation (UV-vis absorptions) and the emission of HL; electronic excitation energies (eV) and oscillator strengths (f), configurations of the low-lying excited states of HL; calculation of the S₀–S₁ energy gaps based on optimized ground-state geometries (UV-vis absorption) and the optimized excited-state geometries (fluorescence). (CH₃OH used as solvent)

	Process	Electronic transitions	Composition	Excitation energy	Oscillator strength (f)	CI	λexp (nm)
HL	Absorption	S0 → S1	HOMO → LUMO	3.7679 eV (349.40 nm)	0.9725	0.6990	347
	Emission	S1 → S0	LUMO → HUMO	2.7569 eV (504.63 nm)	0.5051	0.7057	511
FHL	Absorption	S0 → S1	HOMO → LUMO	3.8933 eV (325.07 nm)	0.2730	0.4879	322
	Emission	S1 → S0	LUMO → HUMO	3.1005 eV (439.70 nm)	0.5263	0.7057	450

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To get a better insight on the experimental photophysical values, TDDFT calculations were performed for the ligand on the basis of the optimized geometry both in the singlet ground state (S₀) and excited state (S₁) in solution. The absorption energies associated with their oscillator strengths, the main configurations and their assignments calculated using the TDDFT method using the S₀ geometry for HL are given in Table 3. In the case of the ground state (S₀) the electron density at the HOMO is delocalized over the entire system, while in the case of the LUMO it originates from the contribution of both the salicylaldehyde (54%) and azomethine (35%) moieties. The excitation energy is 3.7377 eV and this is due to HOMO → LUMO transitions with ILCT character. Frontier molecular orbitals involved in the UV-vis absorption of HL are given in Fig. 4a. In summary, the TDDFT calculations revealed that the ILCT process occurred from the aminofluorene moiety (donor) to the salicylaldehyde moiety (acceptor).

Fig. 4 Frontier molecular orbitals involved in the UV-vis absorption and emission of HL (a) and [FHL]⁻ (b). CT stands for conformation transformation. Excitation and radiative decay process are marked as solid lines and the non-radiative processes are marked by dotted lines.

In order to study the emission properties, a potential energy scan of HL was performed. As shown in ESI Fig. S8,† the keto (dihedral angle between the two planes made by the salicylaldehyde moiety and the aminofluorene moiety is 0.09°) tautomeric form is more stable by 1.882 kcal mol⁻¹ than the corresponding less planar enol form (dihedral angle between the two planes is 6.1°). In the excited state proton transfer takes place from the less planar enol tautomeric from to the more stable coplanar keto tautomeric form. Hence the emission of the ligand is attributed to the ESIPT mechanism. The energy gap between the S₀ and S₁ states, calculated with the optimized S₁ state geometry, is the fluorescence emission wavelength. This geometry relaxation upon photo excitation imparts a remarkable effect on the energy level of the molecular orbitals. In the case of HL the LUMO is stabilized by 0.4829 eV at the S₁ state geometry compared to that at the S₀ state geometry while the HOMO is destabilized by 0.5794 eV for the S₁ state geometry compared to that at the S₀ state geometry. As a result, the energy difference between the HOMO and LUMO is greatly decreased at the S₁ state compared to that at the S₀ state geometry and this geometry relaxation is the main origin of the large Stokes shift. The fluorescence wavelength was calculated as 504.63 nm (in CH₃OH) which is in very good agreement with the experimental value of 511 nm.

Anion sensing study

UV-vis titration of HL with F⁻

With the ligand in hand, we evaluated its response towards anions by absorption, emission, NMR and lifetime spectroscopy. The free sensor HL displayed a major absorption band near 350 nm in HEPES aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2. However, the addition of F⁻ induced a large blue shift in the absorption peaks at 322 nm passing through an isosbestic point at 329 nm, indicating the formation of a new species upon treatment of HL with F⁻. Interactions with other anions (Cl⁻, Br⁻, I⁻, OAc⁻, SO₄²⁻, Pi⁻, PF₆⁻, CN⁻ and PPi⁻) were investigated using spectrophotometric titrations. The strong absorption intensity at 322 nm was observed only for the fluoride anion while the other anions remained silent. This may be due to electronic excitation through charge transfer from the aminofluorene and the fluoride moiety to the salicylaldehyde moiety of the ligand (Fig. 4b). The inset of Fig. 5 gives [A⁻]/([A] + [HL]) = 0.5 which confirms 1 : 1 binding with HL and the fluoride anion to form the [FHL]⁻ compound. To get better agreement with the experimental photophysical values, TDDFT calculations were performed for the [FHL]⁻ on the basis of the optimized geometry both in the singlet ground state (S₀) and excited state (S₁) in methanolic solvent.

Fig. 5 (a) Spectrophotometric titrations of HL (10 μM) with various numbers of equivalents of TBAF ([F⁻] = (0–12 × 10⁻⁶ M)) in HEPES aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2. Inset: the corresponding titration profiles for Job’s method to confirm the 1 : 1 (HL : F⁻) binding stoichiometry; (b) UV-vis spectra of HL with different anions in HEPES aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2.

The absorption energies associated with their oscillator strengths, the main configurations and their assignments calculated using the TDDFT method and the S₀ geometry for [FHL]⁻ is discussed here (vide infra). The energy difference between the HOMO and the LUMO is 3.9 eV. The calculated absorption bands are located at 322.66 nm, which is in good agreement with the experimental result. These absorption bands can be assigned to ICT with some rearrangement of electron density from fluoride to the ligand moiety. Here, the expectation is that the deprotonation facilitated by the fluoride anion will generate strong intramolecular charge transfer (ICT), which would lead to the enhancement of the HOMO–LUMO energy of the FHL compound which is 0.1254 eV higher than the free ligand form, finally resulting in hypsochromic shift around 30 nm.

Emission titration of HL with F⁻

Upon gradual addition of F⁻ to the receptor solution, the emission intensity increased by 12 fold at 450 nm, the emission maxima showed a blue shift of 61 nm and the new maxima was centered around 450 nm in HEPES aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2. This clearly indicates a ‘turn-on’ response of the synthesized chemo sensor of HL. The inset of Fig. 6 shows a plot of emission intensity at 450 nm against the titration of F⁻ from 0 to 1.1 equivalents. It is clear from the plot that the fluorescence intensity reaches a plateau after the addition of exactly 1.0 equivalent of F⁻ ions and there is no significant enhancement of the fluorescence intensity on further addition of F⁻. This result strongly corroborates with the formation of the 1 : 1 [FHL]⁻ compound.

Fig. 6 (a) Spectrofluorimetric titrations of HL (10 μM) with various numbers of equivalents of TBAF ([F⁻] = (0–12 × 10⁻⁶ M) in HEPES aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2. Inset: the corresponding titration profiles confirm the 1 : 1 (HL : F⁻) binding stoichiometry; (b) emission spectra of HL with different anions at room temperature; (c) photographs of sensor HL in HEPES aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2 in the presence of different anions under UV light.

Theoretical treatment of FHL

To scrutinize the emission properties a potential energy scan of [FHL]⁻ was performed. As shown in Fig. 7 during the addition of the fluoride anion to the ligand (HL), initially some hydrogen bond interactions take place between the fluoride anion and the phenolic proton [form A] which was confirmed by NMR titration (vide infra). Finally the hydrogen atom almost detaches from the ligand moiety and binds with the fluoride anion [form B]. Fig. 7 shows [B] is energetically 0.081 eV more stable than [A]. The selective C–O, O–H and F–H bond distances of the two forms clearly indicate that the [A] form consists of the ground state enol form and almost converted into the keto tautomeric [B] form (Table 4), as a result of which the [B] form (dihedral angle 1.09°) becomes more planar leading to a strong emission intensity (dihedral angle of [A] is 41.7°).

Fig. 7 Potential energy curves of FHL calculated at the DFT/B3LYP level.

Table 4 Optimized structural parameters of two different forms of the [FHL]⁻ compound

	C–O (Å)	O–H (Å)	F–H (Å)	Energy	(C6, C7, N1, C8) (°)
A[F⋯HL]	1.3941	1.0206	1.5849	−1001.07617	41.7
B[FH⋯L]	1.3001	1.4132	1.0421	−1001.07915	1.09

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This type of geometrical relaxation is the origin of the emission enhancing properties as well as the proton transfer from the oxygen atom to the fluoride atom which is the mechanistic cause of the emission properties (ESIPT). The potential energy scans of both systems (HL and FHL) indicate that they pass through a minimum and the optimized configurations at these minima appeared due to proton transfer which takes place from the oxygen atom to the nitrogen atom for HL and the fluoride ion for FHL. As shown in Fig. 4 the HOMO–LUMO energy gap of the S1 state of the polar FHL system is more than for HL itself, as a result the former gives comparatively shorter wavelength emission maxima (λ_em 450) than the latter (λ_em 515) and continuously emits blue emission light in the presence of a UV lamp (Fig. 6c). The quantum yield determination of FHL gave Φ = 0.127, while the free ligand was found to be very weakly fluorescent Φ = 0.011. The increase in emission intensity was approximately 10 times in the presence of F⁻ which accounted for the transfer of the phenolic proton and could be used as a fluoride ion selective luminescent probe. All these findings indicate that HL behaves as a highly sensitive fluorescent F⁻ sensor.

¹H NMR titration of HL with F⁻

To gain insight into the interaction between HL and fluoride, ¹H NMR titration was carried out by using CD₃CN as a solvent at room temperature (Fig. S9†). The proton NMR spectra were very much affected during the enhancement of fluoride ion concentration. The phenolic peak which appeared at 13.41 ppm was shifted towards the downfield region and the peak intensity decreased with broadening, hence we can conclude that fluoride and hydrogen interactions took place. When an excess amount of anion was added complete deprotonation occurred. The methylene proton (H7) and the salicylaldehyde ring protons (H2 and H4) were affected a little by the anion, the methylene proton undergoes a downfield shift while the salicylaldehyde protons appeared in the shielding region. In both cases the peak intensity remains unaltered. The protons of the fluorene moiety remained almost silent during the whole titration.

Lifetime titration of HL with F⁻

The changes in the emission decay profile of HL upon the incremental addition of F⁻ show that the decay exhibited complex kinetics that can be fitted with a sum of two exponentials compared with the mono exponential decay behavior of the free receptor (Fig. 8). Moreover, the lifetime of the second component gradually increases whereas the lifetime of the first component gradually decreases. Lifetime data indicate that the deprotonated receptor lives longer than its protonated counterpart.

Fig. 8 Changes in time resolved fluorescence spectrum of HL in HEPES aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2 upon the addition of F⁻ ions and lifetime values are also given in the inset of the figure.

Cation sensing study

UV-vis titration of HL with cations

The UV-vis spectrum of sensor HL was recorded in HEPES aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2 which displayed well-defined bands at 272 nm and 350 nm. The cation binding affinities of HL toward Na⁺, K⁺, Ca²⁺, Mn²⁺, Cr³⁺, Co²⁺, Ni²⁺, Cu²⁺, Cd²⁺ and Zn²⁺ were investigated using UV-vis spectroscopy. However, no such significant change in the UV-vis spectrum was observed for intracellular metal ions such as K⁺, Na⁺ and Ca²⁺ whereas other transition metal ions including Mn²⁺, Cr³⁺, Co²⁺, Ni²⁺ and Cd²⁺ induced a small bathochromic shift of ∼10 nm. Surprisingly a larger amount of red shift (∼60 nm) was observed when the experiment was carried out with Zn²⁺ and Cu²⁺ ions. Fig. 9 shows a representative UV-vis titration curve of HL with various concentrations of Zn²⁺ ions. After the addition of Zn²⁺ ions a change in color of the solution from colorless to yellow occurs, which can be seen by the naked eye. It was observed that with increasing the concentration of Zn²⁺ ions the absorption intensity of the free ligand HL at 273 nm and 352 nm slightly decreased, while a new peak appeared at a longer wavelength (405 nm) with three well-defined isosbestic points at 249 nm, 285 nm and 376 nm. It should be noted that there were no changes in the absorption intensity both at 273 nm, and 405 nm after the addition of in excess of 1.0 equivalent of Zn²⁺ ions with respect to 2.0 equivalents of HL. The UV-vis titration of HL with the Cu²⁺ ion also maintained a similar trend. It passes through two distinguished isosbestic points at 299 nm and 381 nm. However, a significant change in the spectrum was observed during the addition of Cu²⁺ ions, the absorption at 273 nm increased while for Zn²⁺ the reverse phenomenon happened (ESI Fig. S10†). A Job’s plot of absorption intensity shows maxima which correspond to a ∼0.5 mole fraction indicating the 2 : 1 complex formation of HL with Zn²⁺ (Fig. 9) as well as with Cu²⁺ (Fig. S10†). Compared with other detection methods, colorimetric detection offers the advantages of simplicity, rapidity, and cost effectiveness.³¹ Therefore, the colorimetric detection method is extremely attractive in chemical and biological analyses. The N, O bidentate Schiff base type receptor HL showed color changes from colorless to intense yellow in the presence of Cu²⁺ and Zn²⁺ ions, while other metal ions caused no change in color (Fig. 9). Thus, it is very difficult to characterize these two specific metal ions with the naked eye, but we can easily distinguish them by their emission spectral behavior. Fortunately, complex 1 exhibited fluorescence with strong green emission under a UV lamp but 2 did not.

Fig. 9 (a) UV-vis spectra of HL with different cations in HEPES aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2; (b) photographs of the chemo sensor HL in HEPES aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2 in the presence of different cations under visible light; (c) spectrophotometric titrations of HL (10 μM) with various numbers of equivalents of Zn²⁺ at room temperature ([Zn²⁺] = (0–7 × 10⁻⁶ M)); (d) the corresponding titration profiles confirm the 2 : 1 (HL : Zn²⁺) binding stoichiometry.

Theoretical interpretation of the UV-vis study

The geometry optimization of complexes 1 and 2 was performed in solution phase (methanol) in their singlet spin states. The optimized structures of 1 and 2 are shown in ESI Fig. S11.† The partial frontier molecular orbital compositions and energy levels of the complexes in their singlet ground states (S₀) are listed in ESI Table S1.† To gain better insight on the experimental absorption values, TDDFT calculations were performed for complexes 1 and 2 on the basis of the optimized geometry. In order to analyze the nature of the absorption, we performed an NTO analysis based on the calculated transition density matrices. Based on our TDDFT NTO analysis the bands in the region 420–390 nm for complex 1 can be characterized as ILCT states with some rearrangement of electron density over the ligand moiety, while complex 2 showed MLCT as well as ILCT transition character. As illustrated in ESI Fig. S12,† optical excitations occur from the occupied (hole) transition orbitals to the unoccupied (electron) transition orbitals. Hole NTOs contributing to the bands are localized on the π orbitals of the ligands while the electron NTOs are mainly delocalized over the π* orbitals of the ligand moiety. The lower energy absorption band (∼400 nm) can be assigned to the S₀ → S₁ and S₀ → S₂ transitions for the zinc complex whereas S₀ → S₂ and S₀ → S₃ for the copper complex. In both cases the S₀ → S₂ transition had higher oscillating strengths. Fig. 10 shows the accompanying electron density redistribution of complexes 1 and 2.

Fig. 10 Differences in the electron density upon excitation from the ground state (S₀) to allowed singlet states (40 singlet to singlet excitations) for complexes 1 and 2 determined with TDDFT (B3LYP/CPCM-methanol) calculations. Turquoise and purple colors show regions of decreasing and increasing electron density respectively.

Metal d-orbital involvement was detected in complex 2 while complex 1 remained silent. The calculated absorption bands are located at 400 and 270 nm for both complexes which is in good agreement with the experimental results respectively (Table 5).

Table 5 Main calculated optical transitions for complexes 1 and 2 with composition in terms of the molecular orbital contribution of the transition, vertical excitation energies, and oscillator strength in methanol

Complex	Electronic transitions	Composition	Excitation energy	Oscillator strength (f)	CI	Assign	λexp (nm)
1	S0 → S1	H → L	3.0984 eV (400.15 nm)	0.1378	0.69830	^1ILCT	405
	S0 → S2	H − 1 → L	3.1205 eV (397.32 nm)	0.6449	0.30222	^1ILCT
		H → L + 1			0.63466
	S0 → S21	H − 11 → L	4.5541 eV (272.25 nm)	0.6297	0.24933	^1ILCT	273
		H − 11 → L + 1			0.11578	^1ILCT
		H − 10 → L + 1			0.12071	^1ILCT
		H − 9 → L			0.14564	^1ILCT
		H − 8 → L + 1			0.28112	^1ILCT
		H − 7 → L			0.17342	^1ILCT
		H − 7 → L + 1			0.24887	^1ILCT
		H − 6 → L			0.31692	^1ILCT
		H − 6 → L + 1			0.10197	^1ILCT
		H − 5 → L + 1			0.15031	^1ILCT
		H − 4 → L			0.13388	^1ILCT
		H − 4 → L + 1			0.10667	^1ILCT
2	S0 → S2	H → L	3.0107 eV (411.81 nm)	0.7891	0.68245	^1MLCT/^1ILCT	413
		H → L + 1			0.16712	^1MLCT/^1ILCT
	S0 → S3	H → L + 1	3.0232 eV (410.10 nm)	0.1090	0.67032	^1ILCT
		H → L			0.16889	^1MLCT/^1ILCT
		H − 2 → L			0.12722	^1ILCT
	S0 → S18	H − 4 → L	4.4161 eV (280.75 nm)	0.1778	0.59422	^1ILCT	275
		H − 4 → L + 1			0.10913	^1ILCT
		H − 3 → L			0.31065	^1ILCT
		H − 2 → L			0.11227	^1MLCT/^1ILCT
	S0 → S19	H − 8 → L	4.4558 eV (278.26 nm)	0.1234	0.36746	^1ILCT
		H − 4 → L + 3			0.18716	^1MLCT/^1ILCT
		H − 4 → L + 4			0.47200	^1MLCT/^1ILCT
		H − 3 → L + 2			0.19413	^1MLCT/^1ILCT
		H − 3 → L + 4			0.16535	^1ILCT
		H − 2 → L + 3			0.19163	^1MLCT/^1ILCT
		H − 2 → L + 4			0.23925	^1MLCT/^1ILCT
		H → L + 3			0.12363	^1MLCT/^1ILCT
		H → L + 4			0.17214	^1MLCT
		H → L + 5			0.18848	^1MLCT

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Fluorogenic Zn(II) sensing

To determine the practical applications, the fluorescence response behavior of the ligands were examined upon treatment with various metal ions in HEPES aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2. Fig. 11 shows the fluorescence intensity of HL in the presence of different metal ions. Only Zn²⁺ resulted in a pronounced fluorescence enhancement, whereas other transition metal ions Ni²⁺, Co²⁺, Mn²⁺, Cd²⁺ and Cr³⁺ did not induce fluorescence even at concentrations 10-fold higher than the corresponding Zn²⁺ ion concentration, as there is a higher probability of an electron and energy transfer between the metal ion and ligands with Zn²⁺.

Fig. 11 (a) Metal-ion sensitivity towards HL via a spectrofluorimetric study; (b) blue bars represent the fluorescence sensitivity of HL (2 × 10⁻⁵ M) toward various metal ions. Purple bars represent the fluorescence response measured after the addition of Zn²⁺ (1 × 10⁻⁵ M) to the indicated metal ion-complex of 1 (10 : 1 for transition metal ions and 100 : 1 for alkali and alkaline earth metal ions) in aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2 following excitation at 390 nm (slit width 5 nm).

It is interesting to note that the fluorescence intensity of HL in the presence of Cu²⁺ is significantly quenched. When the experiment was carried out with ubiquitous intracellular metal ions such as K⁺, Na⁺ and Ca²⁺, which exist at very high concentrations inside cells, no significant fluorescence was observed, even at concentrations that were 100-fold higher than the Zn²⁺ ion concentration. Metal-ion selectivity was also examined to probe if HL could be used as a selective sensor for Zn²⁺ in the presence of other competitive cations found in biological systems. Emission spectra were measured for a 2 : 1 mixture of HL and Zn²⁺ in the presence of other metal ions. The prominent fluorescence enhancement observed upon mixing HL and Zn²⁺ remained unchanged, even in the presence of a 100-fold excess of metal ions such as K⁺, Na⁺, and Ca²⁺ (Fig. 11, purple bars). This confirms the excellent selectivity of HL for Zn²⁺ over other abundant cations. Notably, the fluorescence intensity of the zinc complex was partially quenched in the presence of metal ions such as Mn²⁺, Cr³⁺, Ni²⁺, Cd²⁺ and Co²⁺ when used in 10-fold excess.

Titration of HL with quencher

An exceptional case appeared in that the fluorescence intensity of HL in the presence of Zn²⁺ was significantly quenched by Cu²⁺ metal ions probably due to their strong binding affinity towards the ligand. Fig. 12 shows the fluorescence spectra of HL in the presence of different concentrations of Cu²⁺ excited at 415 nm in aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2. It can be seen that the ligand showed one intense peak at 512 nm during excitation at 405 nm with a larger slit width (ex/em: 10/10) and the fluorescence intensity gradually decreases in presence of Cu²⁺. It is observed that there is no change in the intensity at 512 nm after the addition of an excess of 1.0 equivalent of the Cu²⁺ ion with respect to 2.0 equivalents of HL.

log(F₀ − F)/F = log K + n log[Mⁿ⁺] (4)

Fig. 12 (a) Fluorescence titration of HL (20 μM) with gradual addition of Cu²⁺ and Zn²⁺ respectively, 1–12 μM in aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2; (b) emission intensity at 512 nm vs. [M²⁺]; (c) linear plot for F₀ − F/F vs. [M²⁺]. Slit width 10/10. λ_ex = 405 nm (M = Cu and Zn).

This result corroborated with the formation of a 1 : 2 (M : L) complex in solution. This change in fluorescence intensity at 512 nm is used to estimate K for the binding of Cu²⁺ to HL using eqn (4).³² Here F₀ and F are the fluorescence intensity of the fluorophore, at 512 nm in the absence and the presence of different concentrations of Cu²⁺ respectively. Fig. 12(c) shows a linear plot passing through the origin for (F₀ − F)/F vs. [Cu²⁺] (n = 1). From this, according to eqn (4), the value of K was estimated to be 5.12 × 10⁶ M⁻¹, for Cu²⁺ towards HL. The reaction of Cu²⁺ with the chelating agent HL induced rigidity in the resulting molecule and produced a large CHEQ effect which further induced a decrease in the fluorescence intensity.

Titration of HL with fluorophore

In the fluorescence titration experiment, receptor HL was subjected to excitation at 405 nm and was monitored after each stepwise addition of Zn²⁺ ions to the solution in aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2. The sensor showed weak emission probably because of quenching by the occurrence of a photo induced electron transfer (PET) process due to the presence of a lone pair of electrons on the donor atoms in the ligands (N, O donor).³³ A gradual enhancement (∼12 fold) of the fluorescence intensity was observed at 512 nm upon increasing the concentration of Zn²⁺ ions. The reaction of Zn²⁺ with the chelating agent HL induced rigidity in the resulting molecule and produced a large CHEF effect which further induced the large enhancement of the fluorescence.³⁴ Fig. 12(b) shows a plot of emission intensity at 512 nm against the titration of Zn²⁺ from 0 to 1.2 equivalents. It is clear from the plot that the fluorescence intensity reaches a plateau after the addition of exactly 1.0 equivalent of Zn²⁺ ions and there is no significant enhancement of the fluorescence intensity on further addition of Zn²⁺. This result strongly corroborates with the formation of a 1 : 2 (M : L) complex. Assuming a 1 : 2 association between Zn²⁺ and HL, the binding constant is determined with the help of the Benesi–Hildebrand equation, given as:³⁵

(F_max − F₀)/(F − F₀) = 1 + 1/K[Zn²⁺]

where, F₀ is the fluorescence of HL in the absence of externally added Zn²⁺, F is the fluorescence obtained at different [Zn²⁺] and F_max is the fluorescence of HL with [Zn²⁺] in large excess. As shown in Fig. 12, the plot of {(F_max − F₀)/(F − F₀)} vs. 1/[Zn²⁺] yields a straight line with slope = (12.3 ± 0.06) × 10⁻⁶, indicating that HL indeed associates with Zn²⁺ in a 2 : 1 stoichiometry. The intercept value 1.05 ± 0.5, close to 1.0, also manifests the self-consistency of the experimental data. Therefore, the ligand association constant K is the reciprocal of slope, 8.13 × 10⁴ M⁻¹. The 1 : 2 complex formation in solution was further confirmed by ESI-MS⁺-(m/z) analysis (see Experimental section).

Conclusion

In summary, we have rationally designed and synthesized fluorene based on the new phenolic proton platform as a novel ratiometric fluorescent sensor for both fluoride anions and zinc cations in aqueous buffer–CH₃OH (2 : 1, v/v) at pH 7.2. The sensor exhibited a large blue shift (32 nm) in absorption and a drastic ratiometric fluorescent response (I₄₅₀/I₅₁₁ = 7.4) to fluoride anions with intense blue light emission, whereas it exhibited a large red shift (61 nm) in absorption and a drastic ratiometric fluorescent response (I₅₂₀/I₅₁₁ = 10.3) to zinc cations with intense green light emission. Density functional theory and time-dependent density functional theory calculations were conducted to rationalize the optical response of the sensor. We expect that the unique ESIPT character of the phenolic proton to the azomethine platform for anion sensing and the CHEF character for cation sensing will be widely employed to construct a wide variety of ratiometric fluorescent sensors based on the ICT signaling mechanism. In addition this sensor could simultaneously detect two Cu²⁺ and Zn²⁺ metal ions through colorimetric methods in mixed aqueous solution.

Acknowledgements

Amar Hens acknowledges UGC, New Delhi for the research fellowship. We are also thankful to the Department of Science and Technology (DST), New Delhi, India for the data collection on the CCD facility setup (Jadavpur University) under the DST-FIST program. We also acknowledge CAS, Department of Chemistry, Jadavpur University and the DST-PURSE program for other facilities.

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Footnote

† Electronic supplementary information (ESI) available: X-ray crystallographic file in CIF format for HL; Fig. S1–S12 and Table S1. CCDC 1054382. For ESI and crystallographic data in CIF or other electronic format see DOI: 10.1039/c5ra05145b

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