Synthesis of Two Mononuclear Schiff Base Metal (M = Fe, Cu) Complexes: MOF Structure, Dye Degradation, H2O2 Sensing, and DNA Binding Property.

O-Vanillin and ethylenediamine were used for the synthesis of N,N'-bis(O-vanillinidene)ethylenediamine (O-VEDH2) Schiff base ligand, which was characterized by UV-visible and 1H NMR spectroscopy. The Schiff base ligand O-VEDH2 has been employed to synthesize novel [C18H22FeN2O6]NO3 (1) and C18H16CuN2O5 (2) complexes by the simple slow evaporation method. The single crystals were characterized by X-ray crystallography, as well as Fourier-transform infrared and UV-visible spectroscopy techniques. Complexes 1 and 2 crystallize in monoclinic and orthorhombic space groups with P121/n1 (14) and Pnma (62) point groups, respectively, and both show metal-organic frameworks like structures. Complexes 1 and 2 have optical band gaps of 4.1 and 2.9 eV, respectively, indicating their semiconducting properties. The dye degradation activity and H2O2 sensing of the complexes 1 and 2 were determined in different conditions. The photocatalytic test was performed in the presence of sunlight, methylene blue (MB), and complexes 1 and 2. An interesting result was obtained that complex 2 has degradation ability against MB and the rate constant (K) is 5.46 × 10-5 s-1, whereas complex 1 has H2O2 sensing properties. Both complexes are bound to Calf-thymus DNA by intercalation binding mode, and binding constants (K b) of complexes 1 and 2 are 3.77 × 103 and 1.49 × 104 M-1, respectively.

Introduction

The Schiff base metal complexes are used extensively in the field of materials science, including catalyst, supramolecular chemistry, anticancer, antimicrobial, antioxidants, electrochemistry, photocatalysts, DNA binding, and energy materials.[1−14] In general, the Schiff base can assemble the metal ions to develop metal–organic framework (MOF) compounds and shows different properties like dye adsorption, sensing, semiconductor, and biological properties.[15−23]

In recent years, researchers from around the world have been working to synthesize metal–organic framework complexes to promote the degradation or adsorption of dyes and to eliminate the dyes from chemically contaminated water.[24−27] Because most of the synthetic dyes were used in a wide range of industries, such as paper, plastics, leather tanneries, food technology, and cosmetics,[28,29] waste dyes are discharged into rivers and ponds, thereby contaminating the water. These dye contaminants can cause allergies, skin irritation, liver damage, kidney damage, and also central nervous system disorders in humans and animals.[30] Currently, Yang et al. reported that three-dimensional (3D)-pillared layer copper(II) MOFs have been used for dual purposes, i.e., dye adsorption and separation.[31]

Nowadays, MOFs have often been used as different types of sensing materials, i.e., H2O2 sensing. Reactive oxygen species like H2O2 are nontoxic but produce free radicals such as superoxide anion (O2•–) and hydroxyl radical (•OH) at the time of decomposition, which plays an important role for the initiation of cancer, autoimmunity, neurodegenerative disorders, etc.[32−34] Hydrogen peroxide is widely used in the food, cosmetic, pharmaceutical, paper, and chemical industries.[35,36] Therefore, it is necessary to detect H2O2 for the healthy life of humans and animals. In the last few years, scientists have developed iron- and copper metal-based MOFs and reported variable sensing applications, i.e., fluorescent, luminescent, photocatalytic, and polar aliphatic volatile organic compounds.[37−41]

MOFs, as well as various Schiff base ligands, play an important role in bioapplications. The effective anticancer drugs have an active site primarily targeted to bind the DNA of the cancer cell, reducing the activity of the cancer cell and resulting in cell death.[42−44] Anticancer drugs interact with cancer cell DNA through electrostatic, intercalative, or groove binding, which depends on the binding modes of anticancer drugs.[25,45−47] The Schiff base ligands and the corresponding metal complexes have the active site that interacts with the DNA of the cancer cell.[48,49] Recently, metal–organic framework systems have been used in biomedicine, as reported by Chen et al.[50,51]

In this study, we have designed and synthesized N,N′-bis(O-vanillinidene)ethylenediamine (O-VEDH2) Schiff base ligand and employed it for the synthesis of new [C18H22FeN2O6]NO3 (1) and C18H16CuN2O5 (2) complexes. The main aim of our research work is to synthesize transition-metal complexes using synthesized Schiff base ligand and study their multifunctional applications, i.e., dye degradation, H2O2 sensing, and DNA binding activity. The Schiff base ligand is characterized by UV–visible and 1H NMR spectroscopies. From optical measurements, complexes 1 and 2 have energy band gaps of 4.1 and 2.9 eV, respectively, and show semiconductor property. Complex 1 can actively sense H2O2, whereas complex 2 has no such sensing ability. On the other hand, complex 2 has photocatalytic activity; it degrades methylene blue (MB) and the rate constant (K) is 5.46 × 10–5 s–1, but complex 1 has no photocatalytic activity. Both the complexes bind to Calf-thymus DNA (CT-DNA) through intercalation binding mode, and their binding constants (Kb) were found to be 3.77 × 103 and 1.49 × 104 M–1, respectively.

Results and Discussion

Structural Description

Complexes 1 and 2 crystallize in monoclinic and orthorhombic space groups with P121/n1 and Pnma, respectively, as shown in Figure . The phase data of the complexes are shown in Table , and the selected interatomic distances and angles are listed in Table . Both complexes have butterfly-like structures. Complex 1 was coordinated by two nitrogen atoms, N(1) and N(2), and two hydroxyl oxygen atoms, O(1) and O(2), of O-VEDH2 moiety and two oxygen atoms, O(5) and O(6), of two water molecules and form an octahedral geometry. Two water molecules are present in the axial position of complex 1. The oxidation state of the metal ion is stabilized by nitrate ion (NO3–) (surrounding the complex) and the corresponding protonated oxygen atom in complex 1, whereas complex 2 was coordinated by two nitrogen atoms, N(1) and N(1)a, and hydroxyl oxygen atoms, O(1) and O(1)a, from one O-VEDH2 moiety and one oxygen (O3) atom and developed square pyramidal structure. The oxidation states of the complexes 1 and 2 are found to be 3+ and 2+, which was also confirmed by the bond valence sum (BVS) calculation shown in Table .[52] The bond lengths between the metal-oxygen and nitrogen of the Schiff base ligand of both complexes are attached through their normal modes of coordination, i.e., for complex 1, Fe1–O1, Fe1–O2, Fe1–O5, Fe1–O6, Fe1–N1, and Fe1–N2 are found to be 1.882(16), 1.866(21), 2.316(13), 2.088(26), 2.094(26), and 2.075(27) Å, respectively, and for complex 2, Cu1–N1/N1′, Cu1–O1/O1′, and Cu1–O3 are found to be 1.942(45), 1.932(28), and 2.330(28) Å, respectively. The intra- and intermolecular H-bonds of complex 1 H5A–O9, H6A–O4, H6A–O9, O1–H6B, H5B–O4, O5–H6B, O5–H6A, O6–H5A, O2–H6A, H6A–N3, H6A–O7, N1–H5A, N2–H6A, N1–H6B, N2–H5A, N2–H5B, O2–H5B, O7–H6B, H6B–N3, H6B–O9, O8–H6B, N3–H6B, N3–H6A, and H6A–O8 are found to be 4.669(48), 4.986(29), 4.382(48), 3.006(17), 4.572(21), 4.554(15), 4.558(13), 4.537(23), 3.045(19), 3.398(29), 3.185(54), 2.793(28), 2.738(23), 2.667(21), 2.771(26), 3.028(30), 2.732(14), 2.851(57), 2.655(29), 3.692(45), 4.753(35), 3.686(26), 2.764(26), and 3.927(39) Å, respectively, as shown in Figure a, whereas complex 2 has no intra- and intermolecular H bonding. Figures b and 3a represent van der Waals interaction between complexes 1 and 2, and their corresponding van der Waals bond lengths are shown in Table , which stabilized the complexes. Complex 2 has an asymmetric unit, as shown in Figure b, whereas complex 1 has no asymmetric unit. The most interesting feature was obtained that both the complexes have built metal–organic frameworks (MOFs) fabricated through Schiff base ligands linked to the entire two-dimensional net and a 3D layer framework for 1 and 2, as shown in Figure a,b, respectively.

Figure 1

(a) Complex 1 and (b) complex 2.

Table 1

Phase Data of Complexes 1 and 2a,b,c,d

parameter	complex 1	complex 2
formula	[C18H22FeN2O6]NO3	C18H16CuN2O5
formula weight (g/mol)	480.2 g/mol	403.8 g/mol
temperature (K)	273 K	99 K
wavelength	0.71073 Å	0.71073 Å
crystal system	monoclinic	orthorhombic
space group	P121/n1 (14)	Pnma (62)
unit cell dimensions a, b, c (Å)	a = 11.8135(2), b = 13.3278(1), c = 13.8168(2)	a = 9.212(2), b = 24.643(5), c = 7.363(2)
α, β, γ (°)	α = 90, β = 111.607(2), γ = 90	α = 90, β = 90, γ = 90
V (Å3)	2022.56(6)	1671.5(7)
Z	4	1
density (g/cm3)	1.577	0.411
radiation type	Mo Kα	Mo Kα
μ (mm–1)	6.504	0.335
absorption correction	multiscan absorption correction	multiscan absorption correction
Tmin, Tmax	0.295, 1.00	0.571, 0.733
Rint	0.029	0.0878
θ (max) (Å–3)	71.360	24.996
refinement R[F2> 2σ(F2)], wR(F2), S	0.0502(3642), 0.1362(3879), 1.043	0.0568(1411), 0.1556(1496), 1.074
goodness of fit on F2	1.041	1.052
no of reflections	3879	1496
no of parameters	284	122
data completeness	0.986	0.995

R1 = ∑(||Fo| – |Fc||)/∑|Fo|.

wR2 = [∑[w(Fo2 – Fc2)2]/∑w[(Fo2)2]]1/2.

S = [∑[w(Fo2 – Fc2)2]/(n – p)]1/2.

w = 1/[∑2(Fo2) + (m*p)2 + n*p], p = [max(Fo2,0) + 2*Fc2]/3, m and n are constants.

Table 2

Selected Bond Distances (Å) and Bond Angles (°) for Complexes 1 and 2a

complex 1		complex 2
atoms 1,2	d 1,2 (Å)	atoms 1,2	d 1,2 (Å)
Fe1–O1	1.882(16)	Cu1–N1/N1′	1.942(45)
Fe1–O2	1.866(21)	Cu1–O1/O1′	1.932(28)
Fe1–O5	2.316(13)	Cu1–O3	2.330(28)
Fe1–O6	2.088(26)
Fe1–N1	2.094(26)
Fe1–N2	2.075(27)

Symmetry codes: primed atoms are related by symmetry −x + 1, y, −z.

Table 3

Bond Valence Sumsb for Complexes 1 and 2a

	complex 1		complex 2
atom	Fe(II)	Fe(III)	Cu(II)	Cu(I)
Fe1	3.29	3.14
Cu1			2.01	1.76

The oxidation state of a particular atom is the nearest integer to the underlined value.

The underlined value is the closest to the charge for which it was calculated.

Figure 2

(a) Intra- and intermolecular H-bonding and (b) van der Waals interaction of complex 1.

Figure 3

(a) van der Waals interaction and (b) asymmetric unit of complex 2.

Table 4

van der Waals Interaction of Complexes 1 and 2

complex 1		complex 2
atoms 1,2	d 1,2 (Å)	atoms 1,2	d 1,2 (Å)
N3–H6B	2.655(29)	O3–C007	3.981(62)
O8–H6B	1.919(39)	O3–C00E	3.657(117)
O8–H16A	2.647(29)	H00E–O2	3.960(36)
O6–O8	2.729(48)	H007–O2	3.395(38)
O8–H14	2.957(43)	H007–O1	3.721(31)
H7–O9	2.929(44)	C007–O2	3.664(70)
O9–H5	2.5834(36)	C00E–O1	3.482(125)
O9–H6A	2.266(48)	H00G–N1	3.477(46)
N3–6A	2.764(26)	H00G–O1	3.160(31)
O7–H6A	2.560(46)	N1–H00E	3.858(50)
O7–H6B	2.851(57)	N1–N1	2.558(70)
H4–O7	2.538(36)	N1–N00C	3.645(70)
		N1–C006	2.930(62)
		O1–N1	2.794(48)
		O1–O1	2.740(35)
		O1–O2	2.589(39)
		O2–H00D	2.649(26)
		O2–C00D	3.692(67)
		C009–O2	3.617(49)
		O1–C00B	3.614(56)
		O1–C00C	3.694(59)
		O1–C009	2.423(54)
		O1–C007	2.952(56)
		O1–H007	3.895(30)

Figure 4

MOF at (010) b axis in central projection for (a) complex 1 and (b) complex 2.

Optical Properties

To explain the conductivity of the isolated complexes, the optical band gap energy of the complexes was calculated by the Tauc formula according to eq where α denotes the absorption coefficient, which is calculated as α = (2.303 × absorption)/t, t = cubet thickness (1 cm); A is a constant; Eg indicates the band gap energy; exponent n depends on the type of transition; and h denotes Planck’s constant.

Equation was used to plot (αhυ)2 vs hυ (Figure ) from which band gap energy (Eg) of the complexes was calculated by extrapolating the linear portion of the curve to (αhυ)2 = 0. The band gap energy of complexes 1 and 2 were 4.1 and 2.9 eV, respectively. In the solar cell, optoelectronic, and electronic applications, this result was significant because the small band gap energy expanded the electronic progress between highest occupied molecular orbital and lowest unoccupied molecular orbital energy levels and increased the electrical conductivity of the complexes.[53−56] The band gap value suggests that complexes are semiconductor materials and complex 2 has shown photocatalytic activity due to the lower band gap energy value (2.9 eV).

Figure 5

Band gap energies of (a) complex 1 and (b) complex 2.

Photocatalytic Activity

We continued our experiment to investigate the photocatalytic activity of the complexes against different types of organic dyes because the band gap of the complexes suggests that complexes are semiconductor materials. The photocatalytic activity of the complexes was investigated against MB, Congo red (CR), methyl orange (MO), rhodamine B (RB), and naphthalol orange (NO) in the presence of sunlight (Figure a). Among the two complexes, complex 2 has photocatalytic activity against MB (Figure b) due to lower band gap energy. The degradation rate of MB by complex 2 in 210 min was 50%. Photocatalytic degradation of MB by complex 2 follows the pseudo-first-order kinetics, and rate constant was determined by the following relation: , where C0 and C denote the MB dye concentration at times t = 0 and t = t, respectively. The inset in Figure b shows a good linear plot of vs t, in which the rate constant was 5.46 × 10–5 s–1.

Figure 6

(a) Kinetics of the dye degradation and (b) MB degradation by complex 2 and rate of reaction.

Complex 1 has no photocatalytic activity due to higher band gap energy value, as a result of which the electron cannot make an effective transition from valence band to conduction band in the presence of sunlight. The proposed mechanistic pathway for the degradation of MB in the presence of complex 2 is shown in Figure .

Figure 7

Possible mechanism and reason for dye degradation by complex 2 in the presence of sunlight.

H2O2 Sensing by Complexes

To detect the sensing capacity of the complexes in the presence of H2O2 by adding 1 mL of 20 mM H2O2 to 3 mL solution of the complexes, the electronic spectrum was recorded at regular intervals. Figure shows that absorbance of complex 1 decreases with increasing time in the presence of H2O2. After 10 min, we have observed that the characteristic peak of complex 1 at 274 nm disappears and the color intensity of complex 1 turns from brown to colorless. It indicates that complex 1 has H2O2 decomposing ability, whereas complex 2 has no sensing ability. This function of complex 1 can be used as the measure to rapidly detect H2O2 in unknown samples. The addition of H2O2 to the Fe(III) complex results in the formation of free radicals, which initiated the degradation of iron complex, and Fe(III) gets reduced to Fe(II) and decreases the absorbance of complex 1.

Figure 8

H2O2 sensing by Complex 1.

The possible reaction mechanism of H2O2 decomposition by complex 1 is given below.In the possible proposed reaction mechanism, eq was an intermediate of the decomposition of H2O2 by complex 1.[57]

DNA Binding

Before adding CT-DNA to complexes, the stabilities of complexes 1 and 2 in Tris–HCl buffer solution were measured by a UV–visible spectrophotometer at room temperature for 1 h (each 15 min interval measured by UV) and the absorption peak was not changed. Figure indicates that the UV–vis absorption spectrum of complexes 1 and 2 have absorption peak at 274 and 374 nm, respectively, which can be assigned d–d transition. To demonstrate the binding characteristics of the complexes with CT-DNA, we used electronic absorption spectroscopy through changes in absorbance and shift in wavelength. The wavelengths of complexes 1 and 2 exhibit blue shift (for complex 1, the wavelength shifts from 274 to 274, 273, 272, 271, and 271 nm with increasing concentration of CT-DNA of 20, 40, 80, 160, and 320 μL, respectively, and similarly in complex 2, also wavelength shifts from 374 to 371, 371, 370, 370, and 369 nm with increasing concentration of CT-DNA of 20, 40, 80, 160, and 320 μL, respectively), when the concentration of CT-DNA was increased.

Figure 9

Absorption spectra of complexes in the absence and presence of CT-DNA: (a) complex 1 and (b) complex 2.

The intercalability of a metal complex depended on the conjugation, planarity, ligand donor atoms, and the formation of the geometry of the complex.[25] The absorption spectra of the complex binding to DNA through intercalation reveal that blue and red shifts occur due to strong staking interaction between ligand of complexes and base pairs of DNA. The binding constant Kb of the complexes was determined by eq (58)where εa, εf, and εb denote extinction coefficients of the complex, CT-DNA, and bound complex, respectively. The binding constants of complexes 1 and 2 were 3.770 × 103 and 1.488 × 104 M–1, respectively, as shown in Table . Therefore, binding constant values lie within the range of a characteristic of the CT-DNA binding by complexes through intercalative mode.[59,60]

Table 5

Absorption Titration Data of Complexes 1 and 2

Conclusions

The Fe(III) and Cu(II) complexes of the designed Schiff base ligand N,N′-bis(O-vanillinidene)ethylenediamine (O-VEDH2) were successfully synthesized and characterized by single X-ray crystallography and Fourier-transform infrared (FT-IR) spectroscopy. Several applications of the metal complexes were investigated, such as MOFs, conductivity, photocatalytic activity, H2O2 sensing, and DNA binding activity. Complexes 1 and 2 have MOF structures and semiconductor properties, and their optical band gap energy values were 4.1 and 2.9 eV, respectively. Complex 1 has sensing ability to decompose H2O2, but it has no photocatalytic activity, whereas complex 2 has photocatalytic activity against MB degradation but no sensing ability to decompose H2O2. Both the complexes were bound to CT-DNA via intercalative mode of interaction, and binding constants were 3.770 × 103 and 1.488 × 104 M–1, respectively.

Experimental Section

Materials and Methods

All chemical reagents used were commercially available and used without further purification. The solvents were used as received.

Synthesis of Schiff Base Ligand

A methanolic solution of O-Vanillin (0.608 g) was mixed with the methanolic solution of ethylenediamine (0.133 g) under constant stirring for 15 min at room temperature, resulting in the formation of a yellow precipitate. The precipitate was filtered, washed with methanol, and finally dried in vacuum. The yield of Schiff base ligand is 80% (0.486 g). 1H NMR (400 MHz, CDCl3) δ: 8.34 (s, 1H), 6.92 (d, J = 4 Hz, 1H), 6.87 (m, 1H), 6.82 (d, J = 8.3 Hz, 1H), 4.73 (s, J = 7.1 Hz, 1H), 3.97 (t, J = 12.1 Hz, 2H), and 3.50 (s, 3H), as shown in Figure S1 (Supporting Information). Solubility: acetonitrile, dimethyl sulfoxide, dimethylformamide (DMF). Scheme presents the synthesis of Schiff base ligand N,N′-bis(O-vanillinidene)ethylenediamine (O-VEDH2).

Scheme 1

Preparation of Schiff Base Ligand N,N′-Bis(O-vanillinidene)ethylenediamine (O-VEDH2)

Synthesis of [C18H22FeN2O6]NO3 (1)

FeCl3 (0.487g, 3 mmol) was dissolved in 10 mL of acetonitrile, followed by dropwise addition of O-VEDH2 (0.305 g, 1 mmol) (separately soluble in MeCN solution) and stirred for another 2 h until a clear solution is obtained. Finally, the resulting solution was filtered and one part of the brown filtrate was kept for slow evaporation and another part is layered in diethyl ether. Within a week, high-quality black single crystals resulted under both conditions in high yield, which were washed with Et2O (two to three times) and dried under vacuum. Yield: 35% (0.17 g). Anal. calcd for complex 1 C18H22FeN3O9: C, 45.01; H, 4.62; N, 8.75%. Found: C, 45.11; H, 4.82; N, 8.56%. Selected IR data in cm–1: 633 (s), 739 (s), 1080 (m), 1218 (m), 1384 (b), 1442 (s), 1550 (s), 1618 (b) [s = strong, m = medium, b = broad] (for a detailed discussion of IR spectrum of 1, see Figure S2a).

Synthesis of C18H16CuN2O5 (2)

A similar method was followed for the synthesis of complex 2. Copper sulfate (0.250 g, 1 mmol) was dissolved in 10 mL of methanol followed by dropwise addition of acetonitrile solution (0.61 g, 2 mmol) of O-VEDH2 with the other conditions remaining constant. Within a week, high-quality green single crystals appeared, which were washed with Et2O (two to three times) and dried under vacuum. Yield: 30% (0.075 g). Anal. calcd for complex 2 C18H16CuN2O5: C, 53.54; H, 3.99; N, 6.94%. Found: C, 53.23; H, 3.85; N, 6.85%. Selected IR data in cm–1: 499 (b), 601 (s), 772 (s), 1080 (m), 1218 (s), 1394 (m), 1437 (s), 1544 (s), 1636 (b) [s = strong, m = medium, b = broad] (for a detailed discussion of IR spectrum of 2, see Figure S2b).

Physical Measurements

1H NMR spectroscopy was performed on a JEOL-400 MHz spectrometer. Infrared spectra were recorded in the solid-state (KBr pellets) using a PerkinElmer FT-IR spectrometer in the 400–4000 cm–1 range. Bruker SMART APEX performed elemental analysis (C, H, N) and X-ray crystallography. UV–visible spectrophotometer (Shimadzu UV-1800) was used to analyze the electronic spectrum of complexes.

Crystallographic Data Collection and Refinement

Single precious stone X-beam information from complexes 1 and 2 was collected at 293 and 99 K on the Bruker SMART APEX CCD diffractometer using SMART/SANTI programming.[61] Intensity data were collected with Mo Kα graphite monochromatized radiation (0.71073 Å) at 100 K. The structures were understood by a coordinate strategy using the SHELXS-2013[62] program, which was incorporated in WinGX.[63,64] Exact corrections to the absorption were linked to SADABS.[65] All nonhydrogen particles were refined with an anisotropic displacement coefficient. The hydrogen carbon-bonded particles were incorporated in geometric positions, and the given warm parameters are one to two times those of the molecule to which they were attached. An initial search for reciprocal space exposed that complexes 1 and 2 crystallize in monoclinic and orthorhombic space groups with P121/n1 and Pnma, respectively. A and B alerts are observed in the checkCIF file of the complexes due to the presence of high disorder in the structure and solvent molecules.

The results of the precious crystal data and structure-refined information of complexes 1 and 2 are available in Table .

Dye Degradation

Rhodamine B (RB), Congo Red (CR), methyl orange (MO), naphthalol orange (NO), and methylene blue (MB) were separately dissolved in water. The complex (10 mg) was soaked in 10 mL of dye solution and kept under sunlight. Furthermore, we examined the complex’s specific photocatalytic activity, and readings with UV spectroscopy were taken regularly.

Hydrogen Peroxide-Sensing Capacity of the Complexes

The detection of H2O2 by both complexes is investigated by the protocol developed by Bera et al. (2013).[66] Initially, complexes 1 and 2 were dissolved in CH3CN and DMF, respectively. A suitable dilute solution of the complexes (3 mL) was noted for the initial of UV–visible spectrum. In the complex solution, 1 mL of 20 mM H2O2 was added. The readings were taken at regular intervals with UV–visible spectroscopy.

DNA Binding Study

Calf-thymus (CT) DNA was dissolved in Tris–HCl buffer (pH = 7.25) solution, and its purity was checked from the absorbance ratio A260/A280, which is in the range 1.8–1.9. Metal complexes 1 and 2 were dissolved, and absorption titration experiments were performed using increased concentration of CT-DNA with the constant amount with the metal complex.