Experimental and Theoretical Studies on a Simple S-S-Bridged Dimeric Schiff Base: Selective Chromo-Fluorogenic Chemosensor for Nanomolar Detection of Fe2+ & Al3+ Ions and Its Varied Applications.

A simple S-S (disulfide)-bridged dimeric Schiff base probe, L, has been designed, synthesized, and successfully characterized for the specific recognition of Al3+ and Fe2+ ions as fluorometric and colorimetric "turn-on" responses in a dimethylformamide (DMF)-H2O solvent mixture, respectively. The probe L and each metal ion bind through a 1:1 complex stoichiometry, and the plausible sensing mechanism is proposed based on the inhibition of the photoinduced electron transfer process (PET). The reversible chemosensor L showed high sensitivity toward Al3+ and Fe2+ ions, which was analyzed by fluorescence and UV-vis spectroscopy techniques up to nanomolar detection limits, 38.26 × 10-9 and 17.54 × 10-9 M, respectively. These experimental details were advocated by density functional theory (DFT) calculations. The practical utility of the chemosensor L was further demonstrated in electrochemical sensing, in vitro antimicrobial activity, molecular logic gate function, and quantification of the trace amount of Al3+ and Fe2+ ions in real water samples.

Introduction

The rapid advancement of fluorescent chemosensors for the recognition of cations has captivated notable attention of researchers because of their great selectivity, high responsivity, low detection limit, and naked eye recognition. In addition to their fascinating sensing toward any type of anclass="Chemical">alyte species, they are still the pren class="Chemical">ferable techniques utilized in medical, environmental, and biological applications.[1−8] Although various scientific methods are available explicitly for the detection of cations, the colorimetric/fluorescent methods are observed as better ones.[9−12] Besides, the potential of evaluating samples for different targets using a solitary sensor results in quicker analytical performance and probable cost reduction.[13,14] After oxygen and silicon, the third most abundant metallic element (by weight 8.3%) in the geosphere of the earth is aluminum.[15−18] Additionally, it exists in many animals, natural waters, and plants as an Al3+ ion. As such, the Al3+ ion is extensively utilized in different fields, including pharmaceuticals, food packaging, manufacturing industry, and water cleansing.[19−22] As a consequence, it could enter the human body without much of a stretch through food and water because of the wide usage in different platforms. The World Health Organization (WHO) outlined that the average day-by-day intake of aluminum in human body is roughly 3–10 mg and the bearable admission is assessed to be 7 mg/kg of body weight.[23−25] The extreme exposure of human body to Al3+ ions also causes numerous risky diseases such as the progression of bone disease in children, encephalopathy, Alzheimer’s disease, Meknes disorder, and Parkinson’s disease.[26−31] Until now, not very many fluorescent chemosensors have been investigated for selective detection of Al3+ ions.[32] The poor coordination ability, strong hydration capacity, and lack of spectroscopic characteristics of the Al3+ ion have made its screening and detection difficult.[33] Therefore, the advancement of profoundly selective and sensitive fluorescent Al3+ sensors having turn-on sort of fluorescence changes is of critical interest for fast and careful discrimination of Al3+ ions from other potentially competing metal ions at a very low concentration (nanomolar level) in biological and environmental systems.[34−41]

Similarly, class="Chemical">iron is the most imperative bioactive transition n class="Chemical">metal involved in living systems and also the most essential trace element for human sustenance.[42−45] The World Health Organization has reported that every day need of Fe2+ for a human body is around 10–50 mg/day. In addition, it plays a critical role in biochemical processes like oxygen transportation, cellular metabolism, and DNA synthesis and is also involved in electron transfer.[46−49] Accordingly, it is broadly dispersed in environmental and biological materials.[50−53] Be that as it may, Fe2+ ion deficiency or overconsumption causes various diseases such as low blood pressure, heart diseases, kidney damages, hemochromatosis, anemia, cellular damage, atherosclerosis, cancer, and neurological disorders.[54,55] In view of the above reasons, the development of a reliable recognition method for concentration level of Fe2+ ion has consistently attracted a lot of consideration in environmental and scientific fields are important yet additionally opportune.[56−63]

Conventionclass="Chemical">al signn class="Chemical">aling mechanisms such as photoinduced electron transfer (PET), CT, ET, and excimer/exciplex, etc. have been evolved/used for the optical recognition of various analytes based on the principal photophysical properties of fluorescent chemosensors.[64−67] The photoinduced electron transfer (PET) mechanism has been involved most extensively in fluorescent chemosensors among other conventional mechanisms.[68−71] When photoinduced electron transfer (PET) takes place from lone pairs of electronegative atoms (N, O, and S) of a sensor molecule to highest occupied molecular orbital (HOMO) of excited fluorophore, turn-on fluorescence occurs. This PET process causes the quenching of fluorophore and reviving via the inhibition of PET by guest species.[72−74] In continuation of our work toward developing small-molecule-based colorimetric/fluorometric chemosensors, we herein report a simple naphthalene-based Schiff base fluorometric and colorimetric chemosensor for post-transition (Al3+) and transition (Fe2+) metal ions through the photoinduced electron transfer (PET) mechanism. To the best of our knowledge, L is the first naphthalene-based probe having the quality of recognizing both Al3+ and Fe2+ ions via two different detection modes.

Results and Discussion

Synthetic Design of Probe L

Probe L was obtained by a simple single-step class="Chemical">Schiff base reaction between n class="Chemical">cystamine dihydrochloride and 2-hydroxy naphthaldehyde in methanol with good yield (Scheme ). The probe L was characterized by 1H and 13C NMR and liquid chromatography–mass spectrometry (LC-MS) analyses (Figures S2–S4). For an effective chemosensor, signaling (fluorophore) and ion-binding (ionophore) units linked with a suitable spacer are highly essential. In our work, we have chosen naphthalene as a fluorophore because of its excellent photophysical properties, high fluorescence quantum yield, better coordination ability, and competitive stability. Disulfide-containing cystamine dihydrochloride was selected as an ionophore as it is an important constituent in numerous metalloproteins and bioenzymes that plays a vital role in their thermal stability and biocatalytic activities. In this context, our work focus on the development of the typical naphthalene-based probe L for selective detection of Al3+ and Fe2+ by dual mode of detection, fluorogenic and chromogenic, respectively.

Scheme 1

Synthesis of Probe L

Colorimetric Recognition of Fe2+

Response of L toward Fe2+ Ions

The UV–vis absorption spectrum of probe L in class="Chemical">dimethylformamide (n class="Chemical">DMF)-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4) shows two intense bands with λmax at 400 and 422 nm (Figure S5). The selective absorption response of probe L (2 × 10–5 M) to Fe2+ was investigated over various metal ions of environmental and biological significance such as Ag+, Al3+, Ba2+, Bi3+, Ca2+, Cd2+, Ce3+, Co2+, Cr3+, Cu2+, Fe2+, Fe3+, Hg2+, K+, La3+, Li2+, Mg2+, Mn2+, Na+, Ni2+, Pb2+, Sr2+, Zn2+, and Zr2+ in DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4) solution at λex = 340 nm. During the absorption spectral analysis, a distinguishable change in the spectra was observed at 418 nm upon addition of Fe2+. Interestingly, other metal ions produced little or no changes in the absorbance spectra with probe L (Figures and S6). The counter anionic effect of probe L was analyzed with FeSO4, Fe(C2H3O2)2, and Fe(OH)2 metal salts to discriminate Fe2+ ions (Figure S7). Therefore, this result proves that probe L could serve as a potential colorimetric sensor for Fe2+ and the enhancement is due to the inhibition of the photoinduced electron transfer process (PET).

Figure 1

UV–vis absorption changes of probe L (2 × 10–5 M) in DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4) in the presence of various metal ions (100 equiv, λex = 340 nm).

UV–vis absorption changes of probe L (2 × 10–5 M) in class="Chemical">DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4) in the presence of various metal ions (100 equiv, λex = 340 nm).

Interference of Other Metal Ions

To understand the reclass="Chemical">alistic utility of probe L, the dun class="Chemical">al metal cross tainting test was executed, as shown in Figure . The competition experiments were carried out by monitoring the changes in the emission intensity before and after adding Fe2+ into the L solution with interferants (100 equiv, DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4)). L treated with 100 equiv of Fe2+ in the presence of other metal ions (100 equiv) (Ag+, Al3+, Ba2+, Bi3+, Ca2+, Cd2+, Ce3+, Co2+, Cr3+, Cu2+, Fe3+, Hg2+, K+, La3+, Li2+, Mg2+, Mn2+, Na+, Ni2+, Pb2+, Sr2+, Zn2+, and Zr2+) did not show any substantial spectral changes. This study reveals that other metal ions did not interfere during the detection of Fe2+ by probe L, as noticed from Figure .

Figure 2

UV–vis absorption changes of probe L with a mixture of dual metal ions (Fe2+ and other stated metal ions) in DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4). Error bars indicate the standard deviation among three samples, and the average is taken from them.

UV–vis absorption changes of probe L with a mixture of duclass="Chemical">al metal ions (Fe2+ and other stated metal ions) in DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4). Error bars indicate the standard deviation among three samples, and the average is taken from them.

Effect of pH and Time Response

To study the practicclass="Chemical">al applicability of probe L as an efn class="Chemical">fective colorimetric sensor, we examined the effect of pH on response of the absorption spectral bands of probe L to Fe2+ at different pH values from 1 to 12 (Figures S8 and S9). For this, the solution was prepared by mixing NaOH and HCl in DMF-H2O (1:1 v/v). The result shows that there is no significant change in probe L at acidic and basic conditions. However, at neutral conditions, L + Fe2+ is quite stable and there is a slight enhancement of absorption above pH 9 and below pH 4. This clearly uncovers the optimal condition of probe L to behave as a better sensor for the analytes under the physiological pH of 7.4. Additionally, a study on the effect of time response was likewise performed to determine the time involved in the binding between probe L and Fe2+ by monitoring the changes in the absorption spectrum (DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4)). The UV–vis absorbance spectra increased, showed the maximum limit until 2 min, and then remained stable for more than 10 min (Figures S10 and S11). Hence, in a short time, probe L favorably recognizes Fe2+ ions and perhaps it can be utilized for sensing of Fe2+ ions in biological and environmental real sample analysis.

Stoichiometry and Binding Mode Studies

To understand the binding mode between probe L and class="Chemical">Fe2+, we carried out UV–vis titrations of probe L by continuous addition of various concentrations of n class="Chemical">Fe2+ ions in DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4). The free probe L showed no significant changes in absorbance at 418 nm, but the absorbance gradually enhanced with an increase in the concentration of Fe2+ ions (0–55 equiv) at 418 nm and got saturated during the addition of 55 equiv of Fe2+ ions (Figure ). This result shows that L/Fe2+ has 1:1 binding stoichiometry. The notable peak at m/z 514.04 for the complex [L + Fe2+] in the mass spectra further advocates the strong binding between probe L and Fe2+ (Figure S20).

Figure 3

UV–vis absorption spectrum of probe L (2 × 10–5 M) in DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4) with different concentrations (0–55 equiv, λex = 340 nm) of Fe2+.

UV–vis absorption spectrum of probe L (2 × 10–5 M) in class="Chemical">DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4) with different concentrations (0–55 equiv, λex = 340 nm) of Fe2+.

To further vclass="Chemical">alidate the stoichiometry of L and n class="Chemical">Fe2+, the absorbance changes were used as a function of mole fraction of Fe2+ in Job’s plot method[82] and the maximum absorbance was observed at 0.5, as shown in Figure , which eventually predicts a 1:1 binding stoichiometry between probe L and Fe2+. The detection limit of probe L was determined as 17.54 × 10–9 M using the formula 3δ/S, where δ denotes the standard deviation of the blank signal and S denotes the slope of the linear calibration plot.[83] Furthermore, the Benesi–Hildebrand nonlinear curve fitting method confirms the above binding stoichiometry (Figure ). The association constant (Ka) of the L + Fe2+ complex is determined by eq (84)where A is the UV–vis absorbance in the presence of Fe2+ and A0 is the UV–vis absorbance in the absence of Fe2+ at 418 nm, respectively. A′ is the maximum absorbance of probe L in the presence of excess amount of Fe2+. Plotting of 1/A – A0 versus 1/[Fe2+] showed a linear relationship (Figure ), which also confirms the 1:1 binding stoichiometry of L + Fe2+. The association constant was calculated as Ka = 2.15 × 102 M–1 for the L + Fe2+ complex.

Figure 4

Job’s plot for the complexation of the [L + Fe2+] system in DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4). Error bars indicate the standard deviation among three samples, and the average is taken from them.

Figure 5

Benesi–Hildebrand nonlinear curve fitting plot (absorbance at 418 nm) of probe L assuming 1:1 binding stoichiometry with Fe2+. Error bars indicate the standard deviation among three samples, and the average is taken from them.

Job’s plot for the complexation of the [class="Chemical">L + n class="Chemical">Fe2+] system in DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4). Error bars indicate the standard deviation among three samples, and the average is taken from them.

Benesi–Hildebrand nonlinear curve fitting plot (absorbance at 418 nm) of probe L assuming 1:1 binding stoichiometry with class="Chemical">Fe2+. Error bars indicate the standard deviation among three sn class="Chemical">amples, and the average is taken from them.

Reversibility of the Probe L

The chemicclass="Chemical">al reversibility of the molecular recognition process of probe L was performed, which is an important requirement for the detection of specific ann class="Chemical">alytes (Figure ). Accordingly, by adding ethylenediaminetetraacetic acid (EDTA; 100 equiv) to a mixture of the L + Fe2+ complex, the original absorption spectra evolved at 418 nm. Hence, the free probe L can again participate in another Fe2+ binding process. This reversible response of probe L with this sort of chelating ability toward Fe2+ meant that the probe L in buffered solution is chemically reversible and can be used for selective recognition of Fe2+ in biological and environmental real samples up to 10 cycles, as shown in Figures b and S12.

Figure 6

UV–vis absorption responses of (a) probe L (2 × 10–5 M) in DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4) upon addition of Fe2+ + EDTA (100 equiv). (b) Number of addition used for the sequential detection of probe L/Fe2+.

UV–vis absorption responses of (a) probe L (2 × 10–5 M) in class="Chemical">DMF-n class="Chemical">H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4) upon addition of Fe2+ + EDTA (100 equiv). (b) Number of addition used for the sequential detection of probe L/Fe2+.

Fluorescent Recognition of Al3+

Response of L toward Al3+ Ions

Fluorescence emission highly depends upon the nature of the solvent system. In our present work, the maximum fluorescence emission is observed in n class="Chemical">DMF as compared to any other solvents including methanol, ethanol, acetonitrile, dimethyl sulfoxide, and tetrahydrofuran. The outcomes showed that probe L could be used to detect the Al3+ through fluorescent emission in DMF. The λex was fixed at 400 nm for cation binding studies based on the absorption spectrum of L and showed λem around 453 nm in DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4). Afterward, the binding property of probe L with various cations was analyzed by fluorescence spectroscopy upon addition of several cations such as Ag+, Al3+, Ba2+, Bi3+, Ca2+, Cd2+, Ce3+, Co2+, Cr3+, Cu2+, Fe2+, Fe3+, Hg2+, K+, La3+, Li2+, Mg2+, Mn2+, Na+, Ni2+, Pb2+, Sr2+, Zn2+, and Zr2+ in DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4) solution. The free probe L showed a weak fluorescence emission at 453 nm. However, upon addition of various metal ions to the solution of L, there were no prominent changes in the fluorescence intensity except in the case of Al3+ (Figures and S13). The counter anionic effect of probe L was studied with AlCl3 and Al2(SO4)3 metal salts to discriminate Al3+ ions (Figure S14). Therefore, the high fluorescence intensity of probe L during the addition of Al3+ exhibits a selective recognition over the other cations by a fluorescence turn-on mechanism.

Figure 7

Fluorescence spectra of probe L (4 × 10–6 M) in the presence of various cations (100 equiv λex = 400 nm) in DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4).

Fluorescence spectra of probe L (4 × 10–6 M) in the presence of various cations (100 equiv λex = 400 nm) in class="Chemical">DMF-n class="Chemical">H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4).

Interference of Other Cations

To investigate the efclass="Chemical">fect of intern class="Chemical">fering cations such as Ag+, Ba2+, Bi3+, Ca2+, Cd2+, Ce3+, Co2+, Cr3+, Cu2+, Fe2+, Fe3+, Hg2+, K+, La3+, Li2+, Mg2+, Mn2+, Na+, Ni2+, Pb2+, Sr2+, Zn2+, and Zr2+, antijamming tests in the presence of Al3+ in DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4) solution were performed, as shown in Figure . The results show that relatively low interference or no interference was observed for the recognition of Al3+ in the presence of other potentially competing cations. Therefore, the L + Al3+ system shows no changes upon addition of other competitive cations. Thus, probe L can be used for the specific recognition of Al3+ ions in real-sample sensing analysis.

Figure 8

Fluorescence response of the L + Al3+ complex on addition of other metal ions (Ag+, Ba2+, Bi3+, Ca2+, Cd2+, Ce3+, Co2+, Cr3+, Cu2+, Fe2+, Fe3+, Hg2+, K+, La3+, Li2+, Mg2+, Mn2+, Na+, Ni2+, Pb2+, Sr2+, Zn2+, and Zr2+; 100 equiv, λex = 400 nm) in DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4). Error bars indicate the standard deviation among three samples, and the average is taken from them.

Fluorescence response of the class="Chemical">L + Al3+ complex on addition of other n class="Chemical">metal ions (Ag+, Ba2+, Bi3+, Ca2+, Cd2+, Ce3+, Co2+, Cr3+, Cu2+, Fe2+, Fe3+, Hg2+, K+, La3+, Li2+, Mg2+, Mn2+, Na+, Ni2+, Pb2+, Sr2+, Zn2+, and Zr2+; 100 equiv, λex = 400 nm) in DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4). Error bars indicate the standard deviation among three samples, and the average is taken from them.

To comprehend the practicclass="Chemical">al applicability of probe L, a correlation study has been performed between L and n class="Chemical">L + Al3+ at various pH ranges to fix a particular pH to investigate the photophysical properties in DMF-H2O (1:1 (v/v)) (Figures S15 and S16). Experimental outcomes demonstrate that there is a slight quenching of fluorescence intensity in probe L at both high acidic and basic conditions. Conversely, for L + Al3+, at acidic conditions (beneath pH 5), the fluorescence intensity increases with the decreasing pH values, and in the pH range of 6–8 pH, it shows very steady enhancement. Again, it decreases with increasing pH values at basic conditions (above pH 8). These results demonstrate that probe L shows good fluorescence response toward Al3+ under the physiological and neutral pH conditions. Therefore, a pH of 7.4 was fixed as the ideal working condition throughout the spectroscopic analyses. Besides, the fluorescence response of probe L to Al3+ in DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4) versus time (in minutes) was additionally analyzed (Figures S17 and S18). The fluorescence emission intensity increased and achieved the maximum level of saturation in 4 min and kept unfaltering for further 10 min. From this study, plainly the probe L can recognize the Al3+ in a short time frame of 4 min, which could be of potential value for biological and environmental Al3+ detection.

Stoichiometry and Binding Mode Studies

To get further insight into the limit of detection and binding interaction between class="Chemical">Al3+ with L, we performed fluorescence titrations of the probe L in the solution containing difn class="Chemical">ferent concentrations of Al3+ ions (DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4)). The fluorescence intensity increased with the increasing concentration of Al3+ (0–80 equiv) with notable changes. The probe L showed no obvious changes in fluorescence emission intensity at 453 nm when it is excited at 400 nm. Stepwise, progressive additions of Al3+ to probe L showed a strong emission, increasing the fluorescence enhancement at 453 nm (Figure ). Furthermore, the spectral changes arrested when 80 equiv of Al3+ was added. We proposed that this phenomenon could be ascribed to the formation of a ligand–metal complex inhibiting the photoinduced electron transfer (PET) process. The coordination between probe L and Al3+ obstructs the PET mechanism. Nevertheless, it introduces rigidity and consequently reduces the flexibility in probe L. Along these lines, the complexation of Al3+ with probe L recovers emission of naphthol and results in a strong fluorescence emission.

Figure 9

Fluorescence titration spectra of probe L (4 × 10–6 M) upon incremental addition of Al3+ ions (0–80 equiv, λex = 400 nm) in DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4).

Fluorescence titration spectra of probe L (4 × 10–6 M) upon incrementclass="Chemical">al addition of n class="Chemical">Al3+ ions (0–80 equiv, λex = 400 nm) in DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4).

The Job plot technique was utilized to find out the stoichiometry between class="Chemical">Al3+ and probe L. The emission intensity of n class="Chemical">L + Al3+ at 453 nm was plotted as a function of its molar fraction under a constant total concentration. The maximum emission point appeared at a molar fraction of 0.5, which is demonstrative of the 1:1 complexing stoichiometry between probe L and Al3+ (Figure ). The detection limit of probe L was determined as 38.26 × 10–9 M. Moreover, the complexing stoichiometry between probe L with Al3+ ion was also confirmed using the Benesi–Hildebrand nonlinear curve fitting method. The association constant (Ka) of the L + Al3+ complex is determined by eq as follows[84]Here, I and I0 are the fluorescence intensities at 453 nm in the presence and absence of Al3+, respectively; I′ is the saturated intensity of probe L in the presence of excess amount of Al3+; and [Al3+] is the concentration of Al3+ ions added. Plotting of 1/I – I0 versus 1/[Al3+] showed a linear relationship (Figure ), which also strongly supports the 1:1 complexing stoichiometry of L + Al3+, and the association constant Ka was calculated to be 3.81 × 103 M–1.

Figure 10

Job’s plot for determining the stoichiometry of probe L and Al3+ ions in DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4; λex = 400 nm). Error bars indicate the standard deviation among three samples, and the average is taken from them.

Figure 11

Benesi–Hildebrand plot of the 1:1 complex of probe L and Al3+ ions. Error bars indicate the standard deviation among three samples, and the average is taken from them.

Job’s plot for determining the stoichiometry of probe L and class="Chemical">Al3+ ions in n class="Chemical">DMF-H2O HEPES solution (1:1 v/v, 50 mM, pH = 7.4; λex = 400 nm). Error bars indicate the standard deviation among three samples, and the average is taken from them.

Benesi–Hildebrand plot of the 1:1 complex of probe L and class="Chemical">Al3+ ions. Error bars indicate the standard deviation among three sn class="Chemical">amples, and the average is taken from them.

The coordination modes of probe L with class="Chemical">Al3+ were further studied by the n class="Chemical">1H NMR titration experiment. Figure shows the spectral changes in chemical shifts of probe L with and without Al3+ (0–2 equiv). There is a slight upfield shift in the “–OH” proton and “C=N” proton on addition of 1 equiv of Al3+. This clearly indicates that the O atom of the phenolic −OH group and the N atom of the C=N (amine) group are coordinated strongly to the Al3+ ion. At this point, on increasing the addition of Al3+, there was no reasonable spectral change, which confirms the 1:1 binding of L to Al3+.

Figure 12

1H NMR titration experiment of probe L and its complex with Al3+ ions (0–2 equiv).

class="Chemical">1H NMR titration experiment of probe L and its complex with Al3+ ions (0–2 equiv).

For a class="Chemical">few useful applications, the recognition and reversibility of the probe L toward n class="Chemical">Al3+ is a vital necessity. The reversibility process was inspected during the addition of EDTA (100 equiv) to the complex mixture of L and Al3+ in DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4). The fluorescence intensity at 453 nm disappeared, as shown in Figure a. However, addition of Al3+ again to the mixture of L + Al3+ containing EDTA attained a considerably enhanced fluorescence intensity. As an outcome, the fluorescence intensity changes were recyclable simply through the addition of EDTA. These results showed that EDTA could be a proper chelating reagent, which could completely regenerate the probe L for further repeated usage toward Al3+ as an off–on–off switch chemosensor. The regenerated probe L can be used up to 10 cycles, as shown in Figures b and S19. The comparative examination of probe L with recently reported sensors is outlined in Table .

Figure 13

Fluorescence spectral changes of (a) probe L in DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4) on addition of Al3+ ions and EDTA (100 equiv). (b) Number of reversible cycles of probe L used for the detection of Al3+.

Table 1

Comparison of Probe L with Recently Reported Chemosensors for Fe2+ and Al3+ Ions

mode of detection	detected ions	detection medium	Ka (M–1)	detection limit (M)	interfering ions	refs
fluorometric	Fe2+	THF-H2O (7:3)	1.31 × 105	115.2 × 10–9		(97)
colorimetric and fluorometric	Fe2+	CH3CN-H2O (8:2)	1.0 × 106	0.272 × 10–6		(98)
fluorometric	Fe2+ and Fe3+	CH3CN-H2O (4:1)	−	1.57 × 10–5 and 1.28 × 10–5		(99)
colorimetric and fluorometric	Cu2+ and Al3+	DMSO-H2O (1:1)	6.55 × 104 and 5.16 × 104	0.217 × 10–6 and 49 × 10–9		(100)
colorimetric	Al3+ and Cr3+	CH3CN-H2O (1:1)	2.35 × 105 and 1.26 × 105	0.17 × 10–6 and 0.12 × 10–6		(101)
fluorometric	Al3+	DMSO-H2O (1:2)	9.40 × 104	1.48 × 10–8		(102)
colorimetric and fluorometric	Fe2+ and Al3+	DMF-H2O (1:1)	2.15 × 102 and 3.81 × 103	17.54 × 10–9 and 38.26 × 10–9		probe L

Fluorescence spectrclass="Chemical">al changes of (a) probe L in n class="Chemical">DMF-H2O HEPES (1:1 (v/v), 50 mM, pH = 7.4) on addition of Al3+ ions and EDTA (100 equiv). (b) Number of reversible cycles of probe L used for the detection of Al3+.

Based on the above results, the proposed binding mechanism is illustrated in Scheme . The probe L exhibits weak absorbance and low fluorescence emission as a result of the intramolecular photoinduced electron transclass="Chemical">fer process. Both emission and absorption show strong enhancement upon addition of n class="Chemical">Al3+ and Fe2+ ions to the probe L, respectively. The high enhancement may be due to the suppression of the photoinduced electron transfer process of the probe L and the rigidity of the complex after binding with respective ions. Job’s plot study, Benesi–Hildebrand nonlinear curve fitting analysis, and the mass spectra confirm the binding stoichiometry of [H]/[G] (1:1) of L + Al3+ and L + Fe2+.

Scheme 2

Proposed Binding Mechanism of Probe L with Al3+ and Fe2+ Ions

Microscopic Studies

Furthermore, to acquire a better comprehension of the surface topography changes, scanning electron microscopy (SEM) images of L, class="Chemical">L + Al3+, and n class="Chemical">L + Fe2+ were examined, which are displayed in Figure a–c. Before and after complexation, there are differences in SEM images, showing a crystal-like structure in probe L, which agglomerated upon addition of Al3+ and became plain and giving plate-like structure due to the complexation of L + Fe2+. The chemical compositions of L + Al3+ and L + Fe2+ were estimated by energy dispersive X-ray analysis (EDAX) analysis (Figure d,e), which directly indicates the existence of carbon (C), oxygen (O), nitrogen (N), sulfur (S), aluminum (Al), and Iron (Fe) elements in the synthesized L + Al3+ and L + Fe2+ complexes.

Figure 14

Surface topographic changes in SEM images of (a) probe L, (b) probe L + Al3+, and (c) probe L + Fe2+; EDAX analysis of (d) probe L + Al3+ and (e) probe L + Fe2+.

Surface topographic changes in SEM images of (a) probe L, (b) probe class="Chemical">L + Al3+, and (c) probe n class="Chemical">L + Fe2+; EDAX analysis of (d) probe L + Al3+ and (e) probe L + Fe2+.

FT-IR Analysis

FT-IR anclass="Chemical">alysis was performed to get insight into the structure and binding mode of the complexes. In the FT-IR spectra of probe L, the characteristic absorption bands at 3002, 1570, and 650 cm–1 are due to the OH, C=n class="Chemical">N, and S–S stretching groups, respectively. There is a notable shift in the absorption bands of probe L on addition of Al3+ ions from 3002, 1570 and 650 to 3010, 1576, and 652 cm–1, respectively. In the IR spectra of L + Fe2+, stretching bands are shifted to 3043, 1574, and 652 cm–1 from the bands of free probe L, respectively. These shifts indicate the possible binding of Al3+ ions and Fe2+ ions with the OH and C=N groups of probe L, respectively (Figure ).

Figure 15

FT-IR spectra of (a) probe L, (b) probe L + Al3+, and (c) probe L + Fe2+.

FT-IR spectra of (a) probe L, (b) probe class="Chemical">L + Al3+, and (c) probe n class="Chemical">L + Fe2+.

Density functional theory (DFT) Studies

To understand the structure of L and its complexation with class="Chemical">Al3+ and n class="Chemical">Fe2+, density functional theory and ab initio calculations at the B3LYP/6-311++G** and HF/6-31G* levels of theory were performed, respectively. The optimized geometries of L, L + Al3+, and L + Fe2+ complexes at the B3LYP/6-311++G** level of theory are shown in Figure a–c. From Figure b, it is observed that L + Al3+ forms a pentadentate complex, wherein Al3+ interacts with nitrogen (Al–N(1), Al–N(2)), oxygen (Al–O(1), Al–O(2)), and sulfur (Al–S(2)) having bond lengths of 1.939, 1.938, 1.942, 1.961, and 2.437 Å, respectively, as shown in Table . Similarly, in the case of L + Fe2+, the structure forms a pentadentate complex, on interacting with two nitrogen atoms, two oxygen atoms, and one sulfur atom (Figure c). On comparing the bond lengths of the L + Fe2+ complex, it was observed that Al–N1, Al–N2, Al–O1, and Al–O2 bond lengths are shorter, indicating that Al3+ could have more profound interaction with the probe L than Fe2+. Overall, the interaction of Fe2+ and Al3+ has a significant effect on the geometry of the probe, as observed from Figure . The optimized energy of the probe L is −2061.91 hartrees and that of its complexes L + Al3+ and L + Fe2+ are −2303.56 and −3325.20 hartrees, respectively. The interaction energies of L + Al3+ and L + Fe2+ complexes are found to be −33.85 and −19.86 eV, respectively. The results clearly indicate that though Fe2+ forms more covalent bonds with the probe L, Al3+ seems to have stronger interaction and hence could be the most suitable metal ion to complex with the probe L.

Figure 16

Optimized geometries of (a) probe L, (b) probe L + Al3+, and (c) probe L + Fe2+ obtained from the B3LYP/6-311++G** level of theory.

Table 2

Structural Parameters of All the Complexes Calculated by the B3LYP/6-311++G** and HF/6-31G* methods

structures		HF/6-31G*	B3LYP/6-311++G**
Bond Length (Å)
L + Al3+	Al–S2	2.410	2.437
	Al–N1	1.936	1.939
	Al–O1	1.923	1.942
	Al–N2	1.940	1.938
	Al–O2	1.945	1.961
L + Fe2+	Fe–S2	2.497	2.292
	Fe–N1	2.034	1.977
	Fe–O1	2.033	2.009
	Fe–N2	2.046	1.986
	Fe–O2	2.060	2.040
Bond Angles (Deg)
L + Al3+	N1–Al–S2	91.63	92.16
	N1–Al–O1	90.26	90.14
	N1–Al–O2	92.02	91.01
	N1–Al–N2	173.5	174.4
L + Fe2+	N1–Fe–S2	91.75	92.83
	N1–Fe–O1	87.08	88.11
	N1–Fe–O2	97.71	93.03
	N1–Fe–N2	174.5	178.5
Dihedral Angles (Deg)
L + Al3+	N1–Al–N2–S2	74.50	84.58
	N1–Al–N2–O1	177.6	–173.5
	N1–Al–N2–O2	–87.62	–79.02
L + Fe2+	N1–Fe–N2–S2	50.10	127.7
	N1–Fe–N2–O1	157.3	–125.2
	N1–Fe–N2–O2	–112.7	–36.61

Optimized geometries of (a) probe L, (b) probe class="Chemical">L + Al3+, and (c) probe n class="Chemical">L + Fe2+ obtained from the B3LYP/6-311++G** level of theory.

Figure shows the FT-IR spectra of L, class="Chemical">L + Al3+, and n class="Chemical">L + Fe2+ structures calculated at the B3LYP/6-311++G** level of theory. The calculated frequencies at the HF/6-31G* and B3LYP/6-311++G** levels of theory are listed in Table . For receptor L, three distinctive absorption bands are observed at 472, 1711, and 3788 cm–1 frequencies, which correspond to S–S, C=N, and O–H vibrations, respectively. Upon interaction of Al3+, the bands undergo significant reduction in the frequency, thereby showing a shift toward the red region. Similarly, in the L + Fe2+ complex, the absorption band is obtained at 498 cm–1 for the S–S vibration, which is a blue shift from the probe, while C–N and O–H vibrations are red-shifted. The frequency calculated at HF/6-31G* for all of the structures matches well with the B3LYP, as seen from Table .

Figure 17

FT-IR spectra of L, L + Al3+, and L + Fe2+ calculated at the B3LYP/6-311++G** level of theory.

Table 3

Selected Unscaled Vibrational Frequencies and Intensity of Probe L and Its Complexes

structures		HF/6-31G*	B3LYP/6-311++G**
L	S–S	511.70 (0.29)	472.59 (5.66)
	C=N	1819.82 (254.28)	1711.78 (132.29)
	O–H	3860.96 (93.03)	3788.45 (84.84)
L + Al3+	S–S	570.62 (8.44)	437.61 (27.90)
	C–N	1920.68 (50.18)	1678.65 (12.59)
	O–H	3541.83 (3113.98)	3701.30 (208.11)
L + Fe2+	S–S	543.62 (8.34)	498.31 (7.27)
	C–N	1852.17 (526.61)	1605.62 (40.20)
	O–H	4062.57 (194.76)	3669.70 (115.68)

FT-IR spectra of L, class="Chemical">L + Al3+, and n class="Chemical">L + Fe2+ calculated at the B3LYP/6-311++G** level of theory.

The highest occupied molecular orbitclass="Chemical">al (HOMO) and lowest unoccupied molecular orbitn class="Chemical">al (LUMO) energies are calculated for the complexes and depicted in Figure . For the receptor L, the HOMO and LUMO energies are −6.042 and −1.703 eV, respectively, and the calculated energy gap value is 4.339 eV. Upon complexation of Al3+ and Fe2+ with the probe L, the energy gap values are found to be 0.101 and 0.100 eV, respectively. Comparatively, L + Al3+ and L + Fe2+ complexes have an almost smaller energy gap, justifying that they are easily polarizable and more reactive.[85,86]

Figure 18

Frontier molecular orbitals of L, L + Al3+, and L + Fe2+ complexes calculated by the B3LYP/6-311++G** method.

Frontier molecular orbitclass="Chemical">als of L, n class="Chemical">L + Al3+, and L + Fe2+ complexes calculated by the B3LYP/6-311++G** method.

The ionization potential (I) and electron affinity (A) n class="Chemical">along with various other chemical reactivity descriptors for the probe L and its complexes are calculated and listed in Table . The chemical potentials of L, L + Al3+, and L + Fe2+ structures are found to be −3.873, −0.479, and −0.370 eV, respectively. A significantly large value for the L + Fe2+ complex clearly indicates that it is more stable and hence does not decompose spontaneously.[87] The global softness value of the L + Fe2+ complex is high compared to the L + Al3+ complex, identifying the former to be more polarized than its counterpart. Other reactivity descriptors further justify the result, making L + Fe2+ as the more suitable complex than L + Al3+. The graphical representation of Mulliken atomic charges for the structures is depicted in Figure . From the Mulliken data, it is observed that all of the hydrogen atoms have negative charges and most of the carbon atoms have positive charges.

Table 4

Calculated Chemical Reactivity Descriptors at the B3LYP/6-311++G** Levela

structures	I	A	Eg	χ	η	ζ	μ	Ψ
L	6.042	1.703	4.339	3.873	2.170	0.230	–3.873	3.456
L + Al3+	0.529	0.428	0.101	0.479	0.051	9.804	–0.479	2.249
L + Fe2+	0.420	0.320	0.100	0.370	0.050	10.0	–0.370	1.369

I, ionization potential; A, electron affinity; Eg, HOMO–LUMO energy gap; χ, electronegativity; η, chemical hardness; ζ, global softness; μ, chemical potential; and Ψ, electrophilicity index. All of the values are given in units of electronvolts.

Figure 19

Mulliken atomic charges of all of the structures obtained at the B3LYP/6-311++G** level of theory.

Mulliken atomic charges of class="Chemical">all of the structures obtained at the B3n class="Gene">LYP/6-311++G** level of theory.

I, ionization potential; A, electron affinity; Eg, HOMO–LUMO energy gap; χ, electronegativity; η, chemicn class="Chemical">al hardness; ζ, global softness; μ, chemical potential; and Ψ, electrophilicity index. All of the values are given in units of electronvolts.

The experimentclass="Chemical">al and theoreticn class="Chemical">al 1H and 13C NMR spectra of the probe L are shown in Figure . The corresponding chemical shift values are presented in Table . The NMR spectra of the probe L are theoretically calculated by gauge-independent atomic orbitals (GIAOs) with respect to tetramethylsilane (TMS) at the B3LYP/6-311++G** level of theory in the gas phase. The aromatic proton peaks in the benzene rings are usually observed between 7 and 8 ppm.[88] In this study, the chemical shifts for the benzene ring of protons are experimentally observed to be in the range of 7.21–8.08 ppm. The theoretically calculated chemical shift for the same is found to be between 7.49 and 7.93 ppm and matches well with our experimental result. Experimentally, the chemical shift values of H21 and H19 atoms that belong to the methylene group are 3.98 and 3.17 ppm, respectively. The calculated theoretical values of H21 (3.84) and H19 (2.68) also match with the experimental result. The chemical shift values of aromatic carbons are usually found to be overlapped in the range of 150–100 ppm.[89] In our case, the chemical shift values of the aromatic carbon in the benzene ring are found to be in the range of 106.45–137.55 ppm experimentally and 117.74–136.40 ppm theoretically. Both theoretical and experimental values correlate with the earlier study.[89] The chemical shifts of methylene carbon (C25, C23) are experimentally observed at 38.46 and 50.38 ppm, respectively. The calculated chemical shift values of C25 and C23 are found to be 54.14 and 49.16 ppm, respectively, and they match well with our experimental results.

Figure 20

Theoretical and experimental 1H and 13C NMR isotropic chemical shifts of probe L.

Table 5

Experimental and Theoretical 1H and 13C NMR Isotropic Chemical Shifts (with Respect to TMS) of Probe L

atoms	theoretical	experimental
H14	9.22	9.15
H4	8.29	7.64
H6	7.93	8.08
H1	7.85	7.75
H3	7.76	7.43
H2	7.49	7.21
H21	3.84	3.98
H19	2.68	3.17
C21	163.24	160.10
C7	136.40	137.55
C6	134.05	134.64
C4	129.72	129.35
C2	128.71	128.37
C13	128.11	125.82
C11	126.03	122.77
C10	120.51	119.07
C8	117.74	106.45
C23	49.16	50.38
C25	54.14	38.46
S1	314.02	−
N1	392.89	−

Theoreticclass="Chemical">al and experimentn class="Chemical">al 1H and 13C NMR isotropic chemical shifts of probe L.

The correlation graph between the experimentclass="Chemical">al and theoreticn class="Chemical">al 1H and 13C NMR chemical shifts of the probe L is depicted in Figure . The correlation coefficients of 1H and 13C NMR were calculated to be 0.993 and 0.991, and they are in good agreement.

Figure 21

Correlation graph between the experimental and theoretical (B3LYP) 1H and 13C NMR chemical shifts of probe L.

Correlation graph between the experimentclass="Chemical">al and theoreticn class="Chemical">al (B3LYP) 1H and 13C NMR chemical shifts of probe L.

Further time-dependent density functionclass="Chemical">al theory (TD-DFT) cn class="Chemical">alculations using the B3LYP/6-311++G** level were performed for L and L + Fe2+ structures to study their UV absorption characteristics. The simulated absorption spectra of L and L + Fe2+ structures are shown in Figure , and the calculated results are listed in Table . The UV–visible spectra for the probe L exhibit the maximum absorption peak at 320.06 nm in the UV region (E = 3.874 eV and f = 0.158 au) along with two absorption peaks located at 316.20 and 310.71 nm (E = 3.921 eV, f = 0.088 au and E = 3.990 eV, f = 0.001 au, respectively). In the case of the L + Fe2+ complex, the strongest absorption peak is featured at 1639.50 nm (E = 0.756 eV and f = 0.001 au), which is red-shifted from the corresponding strongest absorption peak of the probe L. Two additional small peaks are identified at 1274.76 and 863.28 nm (E = 0.973 eV, f = 0.001 au and E = 1.436 eV, f = 0.001 au, respectively) for the complex.

Figure 22

UV–vis absorption spectra of the probe L and complex L + Fe2+ in the gas phase.

Table 6

Values of Excitation Energy (E), Oscillator Strength (f), and Wavelength (λmax) of L and the L + Fe2+ Complex

structures	excitation energy E (eV)	oscillator strength f (au)	λmax (nm)
L	3.8737	0.1575	320.06
	3.9211	0.0875	316.20
	3.9903	0.0014	310.71
L + Fe2+	0.7562	0.0007	1639.50
	0.9726	0.0014	1274.76
	1.4362	0.0005	863.28

UV–vis absorption spectra of the probe L and complex class="Chemical">L + Fe2+ in the gas phase.

Application Studies

Determination of Al3+ & Fe2+ Ions in Real Water Samples

To demonstrate the practicclass="Chemical">al utility of probe L, we have estimated the most abundant n class="Chemical">Al3+ and Fe2+ ions in different water samples to find the feasibility of probe L via fluorescence and UV–vis absorption techniques. The samples were carefully filtered using a 0.5 μL membrane and utilized for the sensing studies. Through atomic absorption spectroscopy (AAS analysis), the reading of 0.5–6 ppm was registered for Al3+ and Fe2+ ion concentrations. Two different commercially available water samples were analyzed by this method, all within the Port Blair, Andaman, region (Table ). The spiked samples were analyzed and confirmed with known standard Al3+ and Fe2+ ion solutions added to probe L (4 μM). The result indicates that the probe L could be potentially utilized for the recognition of the selected metal ions in real water samples without any interferences of other coexisting metal ions.

Table 7

Detection of Al3+ and Fe2+ Ions in Water Samplesa

test sample	concentration of Al3+ present in blank (ppm) (AAS)	Al3+ ion spiked (ppm)	Al3+ ion found (ppm) (fluorescence) (mean ± SD)	Concentration of Fe2+ present in blank (ppm) (AAS)	Fe2+ ion spiked (ppm)	Fe2+ ion found (ppm) (colorimetry) (mean ± SD)
Delanipur ground water	0.934	2	1.98 ± 0.21	1.058	2	2.06 ± 0.10
		4	4.01 ± 0.12		4	4.14 ± 0.14
		6	6.10 ± 0.15		6	5.97 ± 0.14
Chakargaon ground water	0.861	2	2.03 ± 0.11	0.536	2	2.02 ± 0.12
		4	3.88 ± 0.17		4	4.0 ± 0.10
		6	5.97 ± 0.14		6	6.03 ± 0.15

The results are the mean ± SD (n = 3).

Electrochemical Behavior of Probe L with Al3+ and Fe2+ Ions

Electrochemicclass="Chemical">al sensing execution of probe L toward n class="Chemical">Al3+ and Fe2+ ions has been investigated. The cyclic voltammogram (CV) of probe L was estimated in the presence and absence of various cations such as Al3+, Fe2+, Ni2+, and Hg2+, respectively. Here, probe L exhibited some incredible changes upon addition of Al3+ metal ions in DMF solution. Also, electrochemical oxidation of the first peak of probe L was observed, which corresponds to the reversible behavior.[90−93] It might be due to the oxidation of the imine group in which the first oxidation potential observed is −0.96 V. As seen in Figure a, probe L alone demonstrated the reduction potential (Erdx) of −0.96 V. Conversely, on addition of Al3+, the oxidation potential of L + Al3+ shifted to −0.62 V. Interestingly, on simultaneous addition of Al3+ and Fe2+, the peak potential again shifted to −0.57 V. This notable potential shift (Eox) reveals the bonding attraction between the probe L and Al3+ ions. Figure b shows that the addition of Ni2+ and Hg2+ could not cause any shift in the potential of probe L. Meanwhile, under comparative conditions, probe L on investigation with the previously described cations such as Fe2+, Ni2+, and Hg2+ did not show any significant potential shift or changes in the oxidation peak (Figure a,b). All of the above outcomes showed that the probe L is highly sensitive and selective to Al3+ ions than other aggressive cations. From these investigations, the overall order of selectivity is Al3+ > Fe2+ > Ni2+ > Hg2+ (Table S1). Finally, the results provide useful information about the electrochemical sensing nature of probe L toward Al3+.

Figure 23

Cyclic voltammograms obtained on (a) probe L through Al3+, Fe2+, and both of Al3+ and Fe2+ and (b) various metal ions Ni2+ and Hg2+ (100 equiv) with the 0.1 M TBAP supporting electrolyte in DMF solution.

Cyclic voltammograms obtained on (a) probe L through class="Chemical">Al3+, n class="Chemical">Fe2+, and both of Al3+ and Fe2+ and (b) various metal ions Ni2+ and Hg2+ (100 equiv) with the 0.1 M TBAP supporting electrolyte in DMF solution.

Antimicrobial Activity of Probe L with Al3+ and Fe2+ Ions

The in vitro biologicclass="Chemical">al screening efn class="Chemical">fects of the probe L and its complexes L + Al3+ and L + Fe2+ were examined against potential pathogens including both bacteria and fungi such as Staphylococcus aureus, Escherichia coli, and Aspergillus flavus by the disc diffusion method. Empty sterile discs added with the probe L and complexes L + Al3+ and L + Fe2+ were incubated carefully. The test solution was spread out, and the growth of the inoculated pathogen was found to be affected during this period. The inhibition zone developed in the plates was measured (Figure S21). The results show that both the complexes have moderate activity against bacterial and fungal species (Table ). L + Al3+ was found to be more active than probe L and probe L + Fe2+ in the bacterial species S. aureus and E. coli. The results on antifungal activity of probe L + Al3+ show moderate activity in A. flavus, whereas probe L and L + Fe2+ show less activity comparatively.

Table 8

Antimicrobial Activity of Probe L alone and with Al3+ and Fe2+ Ions by the Disc Diffusion Methoda

	zone of inhibition (mm)
sample	S. aureus (mean ± SD)	E. coli (mean ± SD)	A. flavus (mean ± SD)
control	25 ± 0.2	22 ± 0.1	−
L	5 ± 0.4	5 ± 0.3	25 ± 0.3
L + Al3+	15 ± 0.1	10 ± 0.3	36 ± 0.4
L + Fe2+	2 ± 0.1	4 ± 0.2	32 ± 0.1

The results are the mean ± SD (n = 3).

Molecular Logic Gate Function

A molecular Boolean logic function[94−96] was designed based on the reversible fluorescent behavior of the probe L in the presence of class="Chemical">Al3+/n class="Chemical">Fe2+ and EDTA (Figure ). The emission intensity (λmax = 453 nm) and absorbance (λmax = 418 nm) of the probe L were taken as the two outputs of the logic gate function, which were obtained by the addition of Al3+/Fe2+ and EDTA as the two inputs. In this system, strong fluorescence or absorption was considered as the ON mode (output = 1), while weak fluorescence or absorption was considered as the OFF mode (output = 0). In stage 1, the free probe L was taken as the base constituent and it showed no obvious change in the emission intensity and absorbance with the addition of EDTA (input 2) alone as well as with the addition of equal proportion of Al3+/Fe2+ and EDTA (inputs 1 and 2). However, the fluorescent emission and absorbance were switched “ON” by the treatment of Al3+/Fe2+ (input 1) alone. Thus, the presence of Al3+/Fe2+ was represented through a NOT gate to realize the logic function. Since the simultaneous presence of Al3+/Fe2+ and EDTA did not produce any significant change in the output of probe L, this combination (L + Al3+/Fe2+ + EDTA) was taken as the base constituents in stage 2. The results obtained in this stage were found to be identical to those of stage 1. Figure b shows the truth table formulated on the basis of the obtained emission intensity and absorbance changes in both the above stages. The input–output relationships clearly show that the molecular interactions of each stage mimic a NOR gate function (with inverted input 1) outlined in Figure a.

Figure 24

Fluorescence emission and absorption spectra of probe L and L + Al3+/Fe2+ (stage 1) and fluorescence emission and absorption spectra of L + Al3+/Fe2+ with the simultaneous addition of EDTA and Al3+/Fe2+ (stage 2).

Figure 25

Molecular logic gate (a) NOR gate circuit of probe L; (b) truth table of probe L.

Fluorescence emission and absorption spectra of probe L and class="Chemical">L + Al3+/n class="Chemical">Fe2+ (stage 1) and fluorescence emission and absorption spectra of L + Al3+/Fe2+ with the simultaneous addition of EDTA and Al3+/Fe2+ (stage 2).

Conclusions

In conclusion, we have developed an elegant, efclass="Chemical">fective, and economic n class="Chemical">naphthalene-based probe L as a chemosensor for the detection of Al3+ and Fe2+ ions by two different spectroscopic systems. Probe L exhibits selective detection of Al3+ ions by fluorimetry with a detection limit of 38.26 × 10–9 M and of Fe2+ ions by colorimetry with a detection limit of 17.54 × 10–9 M. The recognition of Al3+ and Fe2+ by probe L is free from the interference of Fe2+ and Al3+ ions, respectively, and also from other potentially competing cations. The probe L can be regenerated from its complexes with a suitable chelating agent such as EDTA and thus showing its reversible nature. The potential applications of probe L are demonstrated in real-water sample analysis, electrochemical sensing, in vitro antimicrobial studies, and molecular logic function. Therefore, probe L might serve toward the development of cation-targeting chemosensors in biological, environmental, and medical monitoring systems. Further tuning of the spacer and fluorophore toward the novel construction of chemosensors is currently underway in our laboratory.

Experimental Section

Materials and Instruments

class="Chemical">All chemicn class="Chemical">als and solvents (of analytical reagent grade and spectroscopic grade) used for synthesis were purchased from a commercial source (Sigma-Aldrich) and used without any purification. The 1H NMR and 13C NMR spectra of the ligand and complex were recorded on Bruker 400 and 100 MHz spectrometers (DMSO-d6), respectively, with tetramethylsilane (SiMe4) as an internal standard. The chemical shifts (δ) were expressed in ppm. LC-MS was performed on a LC/MS TOF mass spectrometer. Fourier transform infrared spectrum was recorded on a Shimadzu IRPrestige-21 spectrophotometer. A Jasco V-730 spectrophotometer was used for recording UV–vis absorption spectra at 24 ± 1 °C. JEOL model JSM-6390 was used for examining scanning electron microscope (SEM) studies. Fluorescence emission spectral measurements were carried out using a Jasco FP-8200 spectrophotometer at 24 ± 1 °C. A stock solution of probe L (2 × 10–3 M) was prepared freshly in the system of DMF–H2O (1:1 (v/v), 50 mM, pH = 7.4, HEPES buffer) before the experiments for all spectroscopic analyses. The stock solutions of cations for binding studies were prepared using high-purity chloride and nitrate salts of different metals of Ag+, Al3+, Ba2+, Bi3+, Ca2+, Cd2+, Ce3+, Co2+, Cr3+, Cu2+, Fe2+, Fe3+, Hg2+, K+, La3+, Li2+, Mg2+, Mn2+, Na+, Ni2+, Pb2+, Sr2+, Zn2+, and Zr2+ by dissolving in DMF–H2O HEPES buffer solution (1:1 (v/v), 50 mM, pH = 7.4). The fluorescence titration of probe L (4 × 10–6 M) was performed with a series of solutions containing various equivalents of Al3+ ions. The colorimetric titration of probe L (2 × 10–5 M) was carried out with a series of solutions containing various equivalents of Fe2+ ions. The electrochemical measurements were performed with a glassy carbon (Alfa Aesar with a purity of 99.99% and with an exposed surface area of 0.07 cm2) electrode used as a working electrode and Pt wire as a counter electrode, in which Ag/AgCl with 3.0 M KCl assisted as the reference electrode.

Synthesis of the Chemosensor (L)

Probe L was obtained by stirring a mixture of class="Chemical">cystamine dihydrochloride (0.5 g, 4.40 mmol) and n class="Chemical">2-hydroxy naphthaldehyde (1.52 g, 8.80 mmol) in methanol (30 mL). A catalytic amount (5 drops) of TEA was added dropwise into the above mixture and stirred at 70 °C for 2 h. The precipitate formed was filtered, washed with methanol thoroughly, dried at room temperature, and recrystallized in ethanol to give 1-((E)-(2-(2-(2-((E)-(2-hydroxynaphthalen-1-yl) methylene amino)ethyl)disulfanyl)ethylimino)methyl)naphthalen-2-ol L as a brown microcrystalline solid. Yield: 85% mp: 89 °C. 1H NMR (400 MHz, DMSO, ppm): δ 14.11–14.09 (d, J = 8.8 Hz, 1H), δ 9.15–9.13 (d, J = 9.6 Hz, 1H), δ 8.08–8.06 (d, J = 8.4 Hz, 1H), δ 7.75–7.72 (d, J = 9.6 Hz, 1H), δ 7.64–7.63 (d, J = 7.6 Hz, 1H), δ 7.43–7.39 (t, J = 15.2 Hz, 1H), δ 7.21–7.17 (t, J = 14.8 Hz, 1H), δ 6.76–6.74 (d, J = 9.2 Hz, 1H), δ 3.98–3.97 (d, J = 4.8 Hz, 2H), δ 3.17–3.14 (t, J = 13.2 Hz, 2H). 13C NMR (100 MHz, DMSO, ppm): δ177.03, 160.10, 137.55, 134.64, 129.35, 128.37, 125.82, 122.77, 119.07, 106.45, 50.38, 38.46. Elemental analysis: C26H24N2O2S2; calcd.: C, 67.80; H, 5.25; N, 6.08; found: C, 67.76; H, 5.24; N, 6.02. LC-MS calcd. for C26H24N2O2S2: [M+] 460; found: [M+ + H]+ 461.

Computational Details

The geometries of L, class="Chemical">L + Al3+, and n class="Chemical">L + Fe2+ structures were fully optimized by density functional theory (DFT) and ab initio using Becke’s three-level parameter Lee–Yang–Parr (B3LYP)[75,76] and Hartree Fock (HF)[77] levels along with 6-311++G** and 6-31G* basis sets, respectively. Vibrational frequency analysis was performed for all of the structures, confirming them to occupy the local minima. The interaction energies of all of the complexes were calculated as the difference between the total energy of the complex and the sum of isolated fragments, and their results were corrected for the basis set superposition error (BSSE) by the counterpoise correction method of Boys and Bernardi, in eq , as follows.[78]where EAB is the total energy of the complex and EA and EB are the energies of their constituent monomers or fragments. HOMO–LUMO energies and molecular electrostatic potential (MEP) analysis were performed at the B3LYP/6-311++G** level of theory to find the reactive sites of electrophilic and nucleophilic reactions.[79] All of the theoretical calculations and the molecular representations are done using the Gaussian 09 program package[80] and chemcraft software,[81] respectively.

Electrochemical Studies

BioLogic class="Chemical">SP-150 potentiostats, (France) were employed for the electrochemicn class="Chemical">al investigation. All examinations were completed in an electrochemical cell (10 mL) with the conventional three-electrode setup at room temperature. In the electrochemical measurements, the GC electrode used as a working electrode and Pt wire as a counter electrode, in which Ag/AgCl with 3.0 M KCl assisted as a reference electrode. Likewise, all of the electrochemical experiments were done in 2 × 10–5 M probe L with a 0.1 M TBAP supporting electrolyte. In addition, all of the electrochemical studies were carried out in a N2 atmosphere. The cyclic voltammogram was obtained in the potential range from −1.5 to 0.5 V at a sweep rate of 0.05 V s–1. Before each estimation, the electrode was polished with emery paper.

In Vitro Antimicrobial Activity

The antibacterial activities of probe L and its complexes were performed against the bacterin class="Chemical">al species S. aureus and E. coli by the disc diffusion method. The standard antibacterial agents used in this study were kanamycin and chloramphenicol. Nutrient agar medium was used for the growth of test organisms in Petri plates. The compounds were dissolved in DMSO and soaked in Whatman’s no. 3 filter paper disc of 5 mm diameter and 1 mm thickness. Each sterile disc was incorporated individually with prepared compounds. Precautions were taken to prevent the flow of the compounds from the discs to the outer surface. The prepared compounds were applied in small quantities on discs, and they were allowed to dry in air. Finally, they were incubated at 37 °C for 24 h. After incubation, the diameter of the inhibition zones was measured in mm and recorded. The antifungal activities of probe L and its complexes were also checked against A. flavus cultured on potato dextrose agar (PDA) medium and also carried out by the disc diffusion method at 30 °C for 4–15 days. The agar plates were plunged equally with a swab dipped in the standard inoculum suspension. To allow any excess surface moisture to be absorbed into the agar before compound impregnation in discs, the lids were kept in a laminar flow for 3 min. The test agents were applied to the surfaces of inoculated plates, and the plates were inverted and incubated at 30 °C for 4–7 days for fungal growth. Subsequently, the zone of inhibition was recorded in mm.