English

• 0.   Introduction

  The study of polynuclear 3d-metal complexes is of continuous interest to scientists with various disciplines^[1-4]. This interest is orignated from their applications in magnetic materials^[5-11], spintronics^[12-13], magnetoelectronics^[14], biological systems^[15-18], and photocatalytic activity^[19-20]. Successful strategies for building such clusters include employing appropriate bridging ligands, developing simple and efficient synthetic methods, and selecting suitable metal salt sources.

  One favorable method for synthesizing 3d-metal polynuclear complexes is to use oximate ligands. The oximates show a great ability to prepare polynuclear clusters, and the oximates could efficiently convey magnetic interactions^[21-23]. A plethora of clusters based on oximate ligands has been documented^[24]. To our surprise, the clusters involving terminal coordinate- or bridging chloro-ligands are extremely rare^[25]. Chloride ion is a popular ligand in the preparation of molecular magnets since the μ-Cl bridge could transfer ferromagnetic interactions under certain conditions^[26-28]. Hence, the development of efficient approaches that can generate clusters involving both chloro and oximate ligands is in high demand.

  We have recently attempted to synthesize polynuclear complexes in which chloride ions serve as the terminal coordinate or bridging ligand under solvothermal conditions. Our motivation is driven by the following important reasons: (ⅰ) the nuclearities of clusters are partly dependent on the addition of inorganic anion and on the types of the anion^[29-30]. Except that the μ-chloro bridge could transmit ferromagnetic interactions, the involvement of chloride ions in oximate chemistry may result in clusters with an unrevealed nuclearity. (ⅱ) Solvothermal techniques are rarely used in the coordination chemistry of oximate ligands^[31]. The successful preparation of Mn₆^Ⅱ and Ni₆^Ⅱ wheels^[32], which were not accessible via conventional coordination chemistry techniques, suggests that the combination of the solvothermal technique with the employment of chloro ligand may generate the expected complexes. (ⅲ) Studies on the degradation of organic dyes by polynuclear complexes have rarely been reported^[19-20].

  Under these considerations, we chose 2-pyridinealdoxime (HL), which is a favorable ligand in the preparation of single-chain magnets^[33] as the ligand, NiCl₂·6H₂O and Zn(OAc)₂·2H₂O as the metal and chloro ligand sources, and performed the reaction under solvothermal conditions. A complex with the formula [Ni₂Zn₂(L)₄Cl₂(CH₃O)₂] (1) was prepared. Although the utilization of the L^- ligand had led to the formations of a series of Ni₂^{Ⅱ [34]}, Ni₃^{Ⅱ [35]}, Ni^ⅡCr^{Ⅲ [36]}, Ni^ⅡMn^{Ⅱ [37]}, Ni₉^{Ⅱ [38]}, and Ni₁₄^{Ⅱ [39]}, 1 is the only Ni₂Zn₂ cluster based on the L^- ligand.

  At present, many organic dyes are utilized in academic laboratories and industries such as cosmetics, plastics, and pharmaceuticals^[40]. These organic dyes not only seriously pollute the environment but also cause harmful damage to human health^[41-42]. Therefore, the degradation of organic dyes is attracting intensive attention. Recently, many studies on the removal of organic dyes using Zn^Ⅱ and Ni^Ⅱ complexes as photocatalysts have been carried out due to the low cost, low toxicity, and good photocatalytic performance of these complexes^[43-46]. Thus, the photocatalytic activity toward the degradation of dyes and magnetic properties of 1 were reported in this work.

  1.   Experimental

  1.1   Reagents and characterization

  All experiments were carried out under aerobic conditions. The solvents and reagents are purchased from the chemical companies and used without further purification. The C, H, and N analyses were conducted in a Carlo-Erba EA 1110 CHNO-S microanalyzer. IR spectra of the ligand and complex 1 (4 000-400 cm^-1) were measured in a Nicolet MagNa-IR 500 FT-IR spectrometer. Powder X-ray diffraction (PXRD) pattern of 1 was recorded on a Bruker D8 Advance diffractometer (Cu Kα, λ=0.154 06 nm, 2θ=5°-50°, U=40 kV, I=40 mA). Thermoanalysis of complex 1 was performed in a Netzsch STA 449C thermogravimetric analyzer from 30 to 800 ℃ at a 10 ℃·min^-1 rate for heating. Magnetic susceptibility data (2-300 K) were measured using a Quantum Design MPMS-7 SQUID magnetometer. Diamagnetic corrections were performed using Pascal′s constants. The UV-Vis spectra (200-800 nm) were measured using a Shimadzu UV-Vis spectrometer at room temperature.

  1.2   Synthesis of complex 1

  In the 8 mL Pyrex tube, NiCl₂·6H₂O (0.023 7 g, 0.1 mmol), Zn(OAc)₂·2H₂O (0.043 9 g, 0.2 mmol), HL (0.039 6 g, 0.2 mmol), and MeOH/EtOH (2 mL, 1∶1, V/V) were added. The glass tube was sealed and heated in an oven at 120 ℃ for 48 h. The tube was then cooled at a rate of 5 ℃·h^-1 to room temperature. Brownish-red crystals of the product were obtained. The yield of the product was 54% based on HL. Anal. Calcd. for C₂₆H₂₆Cl₂N₈Ni₂Zn₂O₆(%): C, 36.08; H, 3.03; N, 12.95. Found(%): C, 35.66; H, 3.40; N, 12.62.

  1.3   Photocatalysis experiment

  Complex 1 (10 mg) was mixed with methyl orange (MO, 20 mg·L^-1) or rhodamine B (RhB, 20 mg·L^-1), respectively, to obtain suspensions. The suspension solution was stirred in the dark for about 15 min to ensure the establishment of adsorption equilibrium before irradiation. Then, the solution was conducted on an XPA-7-type photochemical reactor equipped with a 500 W Xenon lamp, and the reaction temperature was maintained at about 25 ℃ by circulating cooling water. A certain volume of samples was collected and separated by centrifugation to remove the residual catalyst particles. Then the solution was analyzed by using a Shimadzu UV-Vis spectrometer. The concentration of MO and RhB were estimated by the absorbance at 462 and 552 nm, characteristic of MO and RhB, respectively.

  1.4   X-ray crystallography

  Single crystals of complex 1 were obtained from MeOH/EtOH under solvothermal conditions. Brown bulk crystals (0.20 mm×0.15 mm×0.15 mm) were selected and placed in a single crystal diffractometer. Diffraction data were measured on Bruker Smart Apex Ⅱ CCDC diffractometers equipped with a graphite monochromator utilizing Mo Kα (λ=0.07 073 nm) radiation at 293(2) K. The diffraction intensity data were corrected using the SADABS program by semi-empirical absorption. The structures were solved by direct methods using the SHELXS-97 package. All non-hydrogen atoms and anisotropy parameters were refined using the full-matrix least-squares on F² method by the SHELXL-97 program. The coordinates of all hydrogen atoms were obtained through theoretical hydrogenation. Crystallographic data together with refinement details for 1 are summarised in Table 1. Selected bond lengths and bond angles for the complex are listed in Table 2.

  Table 1

  

  Table 1.  Crystallographic data for complex 1

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  Parameter	1		Parameter	1
  Empirical formula	C26H26Cl2N8Ni2Zn2O6		F(000)	1 744
  Formula weight	865.61		θ range/(°)	1.54-25.00
  Crystal system	Orthorhombic		Limiting indices	-23 ≤ h ≤ 15, -9 ≤ k ≤ 10, -10 ≤ l ≤ 21
  Space group	Pna21		Reflection collected, unique	21 202, 4 153
  a/nm	1.957 26(16)		Data, restraint, number of parameters	4 153, 1, 418
  b/nm	0.912 86(7)		GOF	1.114
  c/nm	1.795 41(17)		R1, wR2 [I > 2σ(I)]	0.035 6, 0.079 4
  V/nm3	3.207(5)		R1, wR2 (all data)	0.045 7, 0.087 7
  Z	4		Largest diff. peak and hole/(e·nm-3)	670, -260
  Dc/(g·cm-3)	1.792

  Table 2

  

  Table 2.  Selected bond lengths (nm) and bond angles (°) of complex 1

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  Ni1—N1	0.211(7)	Ni2—N4	0.206(7)	Zn1—O5	0.201(6)
  Ni1—N2	0.205(7)	Ni2—N5	0.207(7)	Zn1—Cl1	0.219(4)
  Ni1—N6	0.208(9)	Ni2—N7	0.211(9)	Zn2—O2	0.196(6)
  Ni1—N8	0.205(7)	Ni2—O5	0.210(6)	Zn2—O4	0.199(7)
  Ni1—O5	0.213(6)	Ni2—O6	0.212(6)	Zn2—O6	0.202(6)
  Ni1—O6	0.212(6)	Zn1—O1	0.197(6)	Zn2—Cl2	0.219(4)
  Ni2—N3	0.212(7)	Zn1—O3	0.197(7)
  N1—Ni1—O5	167.4(2)	O6—Ni1—O5	79.0(2)	O5—Ni2—O6	79.6(2)
  N1—Ni1—O6	93.7(3)	N3—Ni2—O6	165.7(2)	O5—Ni2—N7	164.1(3)
  N2—Ni1—N1	78.8(3)	N4—Ni2—N5	162.8(3)	O1—Zn1—O3	95.8(3)
  N2—Ni1—O5	92.4(3)	N4—Ni2—N7	91.3(3)	O1—Zn1—O5	103.4(2)
  N2—Ni1—O6	100.8(3)	N4—Ni2—N3	78.3(3)	O1—Zn1—Cl1	121.2(2)
  N2—Ni1—N6	89.3(3)	N4—Ni2—O5	101.1(3)	O3—Zn1—O5	104.8(3)
  N2—Ni1—N8	163.5(3)	N4—Ni2—O6	91.1(3)	O3—Zn1—Cl1	115.6(2)
  N6—Ni1—N1	96.1(3)	N5—Ni2—N3	89.8(3)	O5—Zn1—Cl1	113.3(2)
  N6—Ni1—O5	92.7(3)	N5—Ni2—N7	78.3(3)	O2—Zn2—Cl2	117.3(2)
  N6—Ni1—O6	167.1(3)	N5—Ni2—O5	91.8(3)	O2—Zn2—O4	98.8(3)
  N8—Ni1—N1	91.1(3)	N5—Ni2—O6	102.5(2)	O2—Zn2—O6	102.4(2)
  N8—Ni1—O5	99.5(2)	N7—Ni2—O6	90.2(3)	O4—Zn2—Cl2	116.8(2)
  N8—Ni1—N6	78.7(3)	N7—Ni2—N3	99.5(3)	O4—Zn2—O6	104.2(3)
  N8—Ni1—O6	92.8(3)	O5—Ni2—N3	92.9(3)	O6—Zn2—Cl2	114.9(2)

  2.   Results and discussion

  2.1   Synthesis of complex 1

  Complex 1 was synthesized in EtOH/MeOH (Eq.1) under solvothermal conditions. The 1∶1∶2 reaction among NiCl₂·6H₂O, Zn(OAc)₂·2H₂O and HL gave tetranuclear 1. We have also attempted the synthesis of 1 by conducting the experiments at room temperature at 101.325 kPa. The reaction of NiCl₂·6H₂O, HL, and Zn(OAc)₂·2H₂O in EtOH under 101.325 kPa and room temperature resulted in a known complex [Ni(L)Cl₂]^[47].

  $ \text{HL}+\text{NiC}{{\text{l}}_{2}}\cdot 6{{\text{H}}_{2}}\text{O}+\text{Zn}(\text{OAc})\cdot 2{{\text{H}}_{2}}\text{O}\xrightarrow{\text{MeOH}/\text{EtOH}, T, p}\left[ \text{N}{{\text{i}}_{2}}\text{Z}{{\text{n}}_{2}}{{\left( \text{L} \right)}_{4}}\text{C}{{\text{l}}_{2}}{{\left( \text{C}{{\text{H}}_{3}}\text{O} \right)}_{2}} \right] $

  (1)

  2.2   Description of the structure

  The coordination environment of complex 1 is shown in Fig. 1. 1 crystallizes in the orthorhombic crystal system of the Pna2₁ space group. In the structure of 1, there are two Ni^Ⅱ ions, two Zn^Ⅱ ions, four L^- ligands, two Cl^- ions, and two CH₃O^- anions. The two Ni^Ⅱ and two Zn^Ⅱ ions are linked together by four L^- ligands and two CH₃O^- anions to form a butterfly-like structure. The Ni1(Ni2) ion is six-coordinated by four nitrogen atoms from two L^- ligands, and two oxygen atoms (μ₃-O) from two deprotonated methanol molecules, forming a distorted octahedron geometry. The Zn1(Zn2) ion is tetra-coordinated by two oxygen atoms from two L^- ligands, one terminal chloride ion, and one oxygen atom (μ₃-O) from one CH₃O^- anion, exhibiting a distorted tetrahedron configuration. The four L^- ligands all display η¹∶η¹∶η²∶μ₂ coordination mode. The ring skeleton of the two Ni^Ⅱ and two Zn^Ⅱ ions of 1 is shown in Fig. 2.

  Figure 1

  

  Figure 1.  Coordination environment of complex 1

  Ellipsoidal probability: 25%.

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  Figure 2

  

  Figure 2.  Ring skeleton of two Ni^Ⅱ and two Zn^Ⅱ ions of complex 1

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  The stacking diagram of 1 is shown in Fig. 3. The layers are interconnected by C3—H⋯O3 (0.254 nm) and C21—H⋯O4 (0.241 nm) bonds, while the layers are interconnected by C11—H⋯O1 (0.248 nm) bonds to form a 2D planar structure, and the faces are interconnected by C—H⋯π bonds, expanding into a 3D structure.

  Figure 3

  

  Figure 3.  Structure stacking diagram of complex 1

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  2.3   IR spectra

  The IR spectra of HL and complex 1 were determined in a range of 4 000-400 cm^-1. As shown in Fig. 4a, the stretching vibration peak of O—H in oxime (HL) appeared at 3 188 cm^-1, and the peak moved to 3 429 cm^-1 in 1. A C=N bending vibration absorption peak in HL appeared at 1 593 cm^-1, and the peak moved to 1 610 cm^-1 in 1. These signals show that the nickel ion is coordinated with the hydroxyl group in the HL ligand.

  Figure 4

  

  Figure 4.  IR spectra of HL (a) and complex 1 (b)

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  In the IR spectrum of 1 (Fig. 4b), the C—H absorption peaks of pyridine were found at 3 077 and 3 005 cm^-1. The signal at 1 540 cm^-1 is assigned to the stretching vibration absorption peak of the aromatic ring. These results revealed that the skeleton structure of the HL ligand remains intact during the coordination process.

  2.4   Thermal properties and PXRD analysis

  To identify the thermal stabilities of complex 1, thermogravimetry (TG) measurements were performed. The TG curve of 1 is shown in Fig. 5a. The weight loss of 3.81% around 100 ℃ may be the loss of solvent water in 1. Then, 1 lost weight dramatically at 294 ℃, and the weight loss was about 34.5%, due to the removal of the pyridine ring in the complex (Calcd. 36.5%). The third weight losses occurred in the interval from 330 to 796 ℃, 8.2% and 19.5% respectively, which may be due to the removal of the coordinated alkoxy group (OCH₃) and the remaining structure (C=N, Cl^-) (Calcd. 7.1%, 20.2%). Finally, the remaining 33.7% may be a mixture of ZnO and NiO (Calcd. 36.0%). The plateau appearing at 294-330 ℃ may be due to the absorption of energy by alkoxy groups during the bond-breaking process.

  Figure 5

  

  Figure 5.  TG curve (a) and simulated and experimental PXRD patterns (b) of 1

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  The PXRD pattern was measured at room temperature to confirm the purity of the phase of 1. The experimental and simulated PXRD patterns are shown in Fig. 5b. The simulated PXRD pattern was in good agreement with that of the experimental ones, proving the phase purity of 1.

  2.5   Magnetic properties

  Magnetic property investigations on transition-metal complexes continue to attract interest from scientists. These studies can provide fundamental factors controlling their magnetic properties. The dc magnetic properties for complex 1 were measured at an applied magnetic field of 1 kOe in a temperature range of 2-300 K (Fig. 6).

  Figure 6

  

  Figure 6.  Magnetic susceptibility for 1 in the form of χ_MT vs T at 1 kOe in a temperature range of 2-300 K, where the solid line corresponds to the best fit from 300 to 2 K

  Inset: magnetic susceptibility data for 1 in the form of χ_M^-1 vs T at 1 kOe in a temperature range of 45-300 K, where the solid line corresponds to the best fit from 300 to 45 K.

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  At 300 K, the χ_MT product of 1 was determined to be 1.99 cm³·mol^-1·K. This value was very close to the calculated value for the two Ni(Ⅱ) ions (2.00 cm³·mol^-1·K). Cooling down the sample of 1 over the whole temperature range, the χ_MT value decreased and reached 0.02 cm³·mol^-1·K at 2 K, revealing the antiferromagnetic interactions between two Ni(Ⅱ) ions. The magnetic properties were analyzed by the reported equation^[48] for two S=1 metal centers. Two parameters, g=2.08 and J=-4.22 cm^-1, were given. The magnetic data of χ_M^-1 vs T from 45-300 K (linear part) were also fitted by the Curie-Weiss law, which yields C=2.005 cm³·mol^-1·K and θ= -11.366 K, confirming the presence of antiferromagnetic exchange coupling interactions among the metal centers.

  2.6   Photocatalytic properties

  In recent years, an increased number of organic dyes have been produced in chemical industries^[49-53]. These organic dyes have played an important role in industrial production^[28-32]. Although these organic dyes provide great convenience in life, they are difficult to degrade by biological processes in the short term due to the stability of their structures. At present, the detection of the dye is mainly through a variety of chromatography and electrochemical methods, but these methods have the disadvantages of pretreatment and high price. It is the focus of scientists′ attention to detect dyes quickly, accurately, and economically.

  Due to its heterometallic nature, complex 1 may be an effective catalyst for the photocatalytic degradation of dyes in wastewater under UV irradiation. Here, we tested the catalytic capability of 1 for MO and RhB under UV irradiation.

  The photocatalytic activity of the catalysts is affected by their band gaps (E_g). Therefore, the E_g of 1 was calculated by the energy axis, and the line extrapolated from the linear portion of the adsorption edge in a plot derived from the Kubelka-Munk function: (αhν)² vs hν. As shown in Fig. 7, the band-gap energy of 1 was 2.86 eV. This reveals that 1 can be activated by radiation from ultraviolet and visible lights, indicating that it may be used as a photocatalyst.

  Figure 7

  

  Figure 7.  Plot of (αhν)² vs hν of complex 1

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  Then, control experiments were performed to prove the properties of 1 for the photocatalytic degradation of dyes. Fig. 8 and 9 show the time-dependent absorption spectra and concentration change of the MO and RhB solutions degraded by H₂O₂ (Fig. 8a and 9a)/1(Fig. 8b and 9b)/H₂O₂ and 1 (Fig. 8c and 9c). The degradation efficiencies of MO and RhB solutions were less than 0.58% and 5.12%, respectively, in the absence of photocatalyst and H₂O₂. The degradation efficiency of MO (RhB) was 30.2% (17.6%) in the presence of H₂O₂. The degradation efficiency of MO (RhB) was 48.9% (46.9%) in the presence of 1. In contrast, the degradation efficiency of MO (RhB) was greatly improved to 83.9% (71.1%) in the presence of H₂O₂ and 1.

  Figure 8

  

  Figure 8.  Time-dependent absorption spectra of MO upon the addition of H₂O₂ (a), complex 1 (b), and H₂O₂+1 (c)

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  Figure 9

  

  Figure 9.  Time-dependent absorption spectra of RhB upon the addition of H₂O₂ (a), complex 1 (b), and H₂O₂+1 (c)

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  To better understand the photocatalytic activity of 1 toward the degradation of MO and RhB, we constructed a plot of c/c₀ versus time (t) for analysis (Fig. 10a and 11a), where c and c₀ are the instantaneous and before irradiation concentrations of MO and RhB, respectively. It was shown that the degradation of MO or RhB was almost negligible in the absence of 1. On the other hand, when H₂O₂ was introduced, 1 displayed good catalytic activity, which implies that the degradations of MO and RhB are due to the photocatalytic activity of 1. To further explore the kinetics of degradation of MO and RhB, we performed kinetic experiments on the degradations of MO and RhB. As a result, the pseudo-first-order kinetic model was verified by the equation -ln(c/c₀)=kt (t: irradiation time; k: total reaction rate constant). The degradation reactions were also analyzed kinetically. As shown in Fig. 10b and 11b, the highest reaction rate constants were 0.009 06 min^-1 for MO and 0.007 14 min^-1 for RhB. The reaction rate constant of the system of 1+H₂O₂ was 2.7 times that of complex 1, and 4.7 times that of H₂O₂ for MO (Fig. 10c). The rate constant of the mixture of 1 and H₂O₂ was 2.5 times that of complex 1, and 7.2 times that of H₂O₂ for RhB (Fig. 11c). These results prove that 1 is photocatalytically active toward the degradations of MO and RhB.

  Figure 10

  

  Figure 10.  (a) Relationships between c/c₀ of MO and irradiation time under different conditions; (b) Plots of -ln(c/c₀) vs t and the corresponding linear fitting for the pseudo-first-order kinetic equation; (c) Reaction rate constants under different conditions

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  Figure 11

  

  Figure 11.  (a) Relationships between c/c₀ of RhB and irradiation time under different conditions; (b) Plots of -ln(c/c₀) vs t and the corresponding linear fitting for the pseudo-first-order kinetic equation; (c) Reaction rate constants under different conditions

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  As reported in the literature^[54-57], when semiconductor particles were excited by UV light, an electron was transferred from HOMO to LUMO. For the complex, the process of degradation was considered to be the transfer of electrons from the HOMO of the 2p bonding orbital of oxygen or nitrogen to the LUMO of the empty metal d orbital. Consequently, the HOMO had the same number of positive slot vacancies (h⁺) and strongly pushes the electron back to a stable state. Next, the electrons from the water were captured, creating the same amount of ·OH radicals. Meanwhile, LUMO electrons react with oxygen atoms adsorbed on the surface of 1 to produce an ·O^2- radical. As a result, dyes can be effectively degraded by free radical species, and the photocatalytic degradation process was completed as shown in Fig. 12. To confirm the degradation process, free radical capture experiments were performed.

  Figure 12

  

  Figure 12.  Mechanism of dye degradation catalyzed by complex 1 under UV light

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  BQ (benzoquinone) (10.8 mg) and TBA (tert-butylalcohol) (1.55 g) were added to the reaction system and used to investigate the main oxidative species ·O^2- and ·OH. As can be seen in Fig. 13, in the presence of TBA, the degradation efficiency (η) was significantly reduced to 11.4% for MO and 9.2% for RhB, respectively. In the presence of BQ, the degradation efficiencies were reduced to 31.4% for MO and 47.2% for RhB, respectively. In conclusion, the presence of BQ and TBA inhibited the degradation efficiencies of dyes. It shows that the active substances in photocatalytic degradation are mainly ·O^2- and ·OH.

  Figure 13

  

  Figure 13.  Degradation efficiencies of MO and RhB under different conditions

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  3.   Conclusions

  In summary, a new Ni^ⅡZn^Ⅱ cluster complex [Ni₂Zn₂(L)₄Cl₂(CH₃O)₂] was prepared by using the solvothermal method. The complex was fully characterized by IR spectroscopy, elemental analysis, and single-crystal X-ray diffraction. It is a new addition to the group of complexes based on 2-pyridinealdoxime ligands. The magnetic properties and photocatalytic activity toward the degradation of dyes were studied. It is found that there exist antiferromagnetic Ni^Ⅱ⋯Ni^Ⅱ exchange coupling interactions in the complex. Moreover, the complex was a good candidate for the photocatalytic degradation of MO and RhB, and the degradation mechanism shows that ·OH and ·O^2- are active species in the degradation.

  
  
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  Figure 1  Coordination environment of complex 1

  Ellipsoidal probability: 25%.

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  Figure 2  Ring skeleton of two Ni^Ⅱ and two Zn^Ⅱ ions of complex 1

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  Figure 3  Structure stacking diagram of complex 1

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  Figure 4  IR spectra of HL (a) and complex 1 (b)

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  Figure 5  TG curve (a) and simulated and experimental PXRD patterns (b) of 1

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  Figure 6  Magnetic susceptibility for 1 in the form of χ_MT vs T at 1 kOe in a temperature range of 2-300 K, where the solid line corresponds to the best fit from 300 to 2 K

  Inset: magnetic susceptibility data for 1 in the form of χ_M^-1 vs T at 1 kOe in a temperature range of 45-300 K, where the solid line corresponds to the best fit from 300 to 45 K.

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  Figure 7  Plot of (αhν)² vs hν of complex 1

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  Figure 8  Time-dependent absorption spectra of MO upon the addition of H₂O₂ (a), complex 1 (b), and H₂O₂+1 (c)

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  Figure 9  Time-dependent absorption spectra of RhB upon the addition of H₂O₂ (a), complex 1 (b), and H₂O₂+1 (c)

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  Figure 10  (a) Relationships between c/c₀ of MO and irradiation time under different conditions; (b) Plots of -ln(c/c₀) vs t and the corresponding linear fitting for the pseudo-first-order kinetic equation; (c) Reaction rate constants under different conditions

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  Figure 11  (a) Relationships between c/c₀ of RhB and irradiation time under different conditions; (b) Plots of -ln(c/c₀) vs t and the corresponding linear fitting for the pseudo-first-order kinetic equation; (c) Reaction rate constants under different conditions

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  Figure 12  Mechanism of dye degradation catalyzed by complex 1 under UV light

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  Figure 13  Degradation efficiencies of MO and RhB under different conditions

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  Table 1.  Crystallographic data for complex 1

  Parameter	1		Parameter	1
  Empirical formula	C26H26Cl2N8Ni2Zn2O6		F(000)	1 744
  Formula weight	865.61		θ range/(°)	1.54-25.00
  Crystal system	Orthorhombic		Limiting indices	-23 ≤ h ≤ 15, -9 ≤ k ≤ 10, -10 ≤ l ≤ 21
  Space group	Pna21		Reflection collected, unique	21 202, 4 153
  a/nm	1.957 26(16)		Data, restraint, number of parameters	4 153, 1, 418
  b/nm	0.912 86(7)		GOF	1.114
  c/nm	1.795 41(17)		R1, wR2 [I > 2σ(I)]	0.035 6, 0.079 4
  V/nm3	3.207(5)		R1, wR2 (all data)	0.045 7, 0.087 7
  Z	4		Largest diff. peak and hole/(e·nm-3)	670, -260
  Dc/(g·cm-3)	1.792

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  Table 2.  Selected bond lengths (nm) and bond angles (°) of complex 1

  Ni1—N1	0.211(7)	Ni2—N4	0.206(7)	Zn1—O5	0.201(6)
  Ni1—N2	0.205(7)	Ni2—N5	0.207(7)	Zn1—Cl1	0.219(4)
  Ni1—N6	0.208(9)	Ni2—N7	0.211(9)	Zn2—O2	0.196(6)
  Ni1—N8	0.205(7)	Ni2—O5	0.210(6)	Zn2—O4	0.199(7)
  Ni1—O5	0.213(6)	Ni2—O6	0.212(6)	Zn2—O6	0.202(6)
  Ni1—O6	0.212(6)	Zn1—O1	0.197(6)	Zn2—Cl2	0.219(4)
  Ni2—N3	0.212(7)	Zn1—O3	0.197(7)
  N1—Ni1—O5	167.4(2)	O6—Ni1—O5	79.0(2)	O5—Ni2—O6	79.6(2)
  N1—Ni1—O6	93.7(3)	N3—Ni2—O6	165.7(2)	O5—Ni2—N7	164.1(3)
  N2—Ni1—N1	78.8(3)	N4—Ni2—N5	162.8(3)	O1—Zn1—O3	95.8(3)
  N2—Ni1—O5	92.4(3)	N4—Ni2—N7	91.3(3)	O1—Zn1—O5	103.4(2)
  N2—Ni1—O6	100.8(3)	N4—Ni2—N3	78.3(3)	O1—Zn1—Cl1	121.2(2)
  N2—Ni1—N6	89.3(3)	N4—Ni2—O5	101.1(3)	O3—Zn1—O5	104.8(3)
  N2—Ni1—N8	163.5(3)	N4—Ni2—O6	91.1(3)	O3—Zn1—Cl1	115.6(2)
  N6—Ni1—N1	96.1(3)	N5—Ni2—N3	89.8(3)	O5—Zn1—Cl1	113.3(2)
  N6—Ni1—O5	92.7(3)	N5—Ni2—N7	78.3(3)	O2—Zn2—Cl2	117.3(2)
  N6—Ni1—O6	167.1(3)	N5—Ni2—O5	91.8(3)	O2—Zn2—O4	98.8(3)
  N8—Ni1—N1	91.1(3)	N5—Ni2—O6	102.5(2)	O2—Zn2—O6	102.4(2)
  N8—Ni1—O5	99.5(2)	N7—Ni2—O6	90.2(3)	O4—Zn2—Cl2	116.8(2)
  N8—Ni1—N6	78.7(3)	N7—Ni2—N3	99.5(3)	O4—Zn2—O6	104.2(3)
  N8—Ni1—O6	92.8(3)	O5—Ni2—N3	92.9(3)	O6—Zn2—Cl2	114.9(2)

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