Multifunctional Properties of a 1D Helical Co(II) Coordination Polymer: Toward Single-Ion Magnetic Behavior and Efficient Dye Degradation.

This contribution deals with the synthesis and utilization of a new pyrazole-based unsymmetrical ligand, 3-(3-carboxyphenyl)-1H-pyrazole-5-carboxylic acid (H2CPCA), for generating multifunctional materials. The reaction with the Co(II) salt in the presence of a co-ligand 2,9-dimethyl phenanthroline (dmphen) results in the formation of the helical compound {[Co2(dmphen)2(CPCA)2]DMF} n (1). However, two isostructural monomeric complexes are formed {[M(HCPCA)2(H2O)2], M = Co(II), (2) and Mn(II) (3)} when reactions were carried out in the absence of dmphen. Compound 1 shows some highly encouraging single-ion magnetic (SIM) properties. Detailed magnetic studies unveil slow relaxation of magnetization of compound 1, driven by the higher magnetic anisotropy of the cobalt ion, with the energy barrier of ∼9.2 K and relaxation time of 9.1 × 10-5 s, suggesting a SIM behavior. Moreover, UV-vis and fluorescence studies confirm the selective dye degradation of compound 1 with methylene blue both in the presence and absence of H2O2, with the remarkable degradation efficiency of ∼98 and ∼82%, respectively.

Introduction

Coordination polymers (CPs) featuring superiorities in rational design and synthesis have become a research hotspot in the multifunctional material field.[1] CPs have received much attention from chemists because of their attractive structures and potential applications in various fronts including adsorption, dye degradation, magnetism, luminescence, etc.[2] Significant efforts have been devoted to generate multifunctional CPs for performing various activities with only one system rather than using different single-purpose systems. Interestingly, the functional activity of CPs often holds a one-to-one correspondence with the characteristic features of the ligand or metals used for synthesis.[3] Synergism of the multiple properties in a single system therefore can be originated or induced using an appropriate combination of constituents with prerequisite features. Therefore, careful selection of organic ligands and metal salts is at the heart of CP synthesis for targeted multifunctional activities.[4][5]

Magnetism, inter alia, has remained one of the prime motifs for CP synthesis. However, an unexpected thrust, especially among the scientists of inorganic chemistry background, came from the contemporary discovery of single molecular magnets (SMMs).[5a] Since the discovery of the Mn12 cluster in early 1990s, scientists are actively engaged in generating exchange-coupled polynuclear clusters having SMM property.[5b] However, most of them are traditional discrete inorganic complexes usually based on Fe[6] or Mn.[7] This is because the SMM property is related to an energy barrier for the reversal of the molecular magnetic moment, which is described as Ueff = |D|S2 (for integer spin).[8] The higher the energy barrier, the greater the possibility of SMM behavior.[9−11] On the other hand, the energy barrier can be increased by increasing the ground state spin (S), which is again directly proportional to the spin nuclearity. However, the approach of a multinuclear system also has limitations as quite a large number of systems of high-spin ground states fail to yield higher barriers.[12] This shifts the focus away from the multinuclear systems and leads to an alternate strategy, proposed by Waldmann and later on by Neese and Pantazis,[15] of maximizing the relaxation barriers through increasing the magnetic anisotropy.[13,14] Along this line, Ishikawa introduces mononuclear SMMs in terms of the [LnPc2]− ion (Pc = pthalocyanine).[15c] Subsequently, 3d transition-metal-based mononuclear SMM, known as a single-ion magnet (SIM), was reported by Long for high-spin Fe(II) compounds.[16,17] Recent focus is aiming on Co(II) compounds because the S = 3/2 half-integer spin state should minimize quantum-tunneling effects and lead to favorable relaxation behavior.[18] Furthermore, Co(II) has both orbital and spin contributions toward magnetic moments, which certainly help to have a large number of microstate energy levels in the double-well energy diagram of SMM. Finally, Co(II) compounds have large anisotropy, which is very crucial for enhancing the energy barrier and thus makes them suitable for SMM.[19−21] It is noteworthy here that a large number of Co(II)-based SIMs were reported having different coordination geometries and environments but most of them are with tetrahedral or pseudo-tetrahedral geometry, while a few exhibit higher coordination numbers of 5 or 6.[22]

Another important aspect of CPs is their application in environmental remediation. Of particular importance is water pollution, which has now become a major global problem with serious consequences on human health and environment. A significant discharge of organic dyes from printing/textile factories is the major source of this problem.[23] Naturally, as a preventive measure, removal of such toxic metals from wastewater warrants remediation prior to their disposal in navigable waters or land. Interestingly, CPs, among the others, have emerged as promising candidates as efficient and low-cost materials for wastewater treatment.[24]

Inspired by the above-mentioned considerations, we endeavored to successfully synthesize three new Co- and Mn-based multifunctional compounds using a new pyrazole-based unsymmetrical ligand, 3-(3-carboxyphenyl)-1H-pyrazole-5-carboxylic acid (H2CPCA) (Scheme ).

Scheme 1

Reaction of Co(II) with H2CPCA and 2,9-dimethyl phenanthroline (dmphen) forms a mononuclear CP {[Co2(dmphen)2(CPCA)2]DMF} (1), whereas two mononuclear complexes [M(HCPCA)2(H2O)2] (M = Co (2) and Mn (3)) were obtained with the reaction of Co(II), Mn(II), and H2CPCA but in the absence of dmphen. Compound 1 shows some interesting magnetic behavior, which qualifies for SIM. The magnetic properties of compounds 2 and 3 have also been studied for comparison. We also report here the chirality and selective dye degradation properties of compound 1.

Experimental Section

Materials and General Method

All reagents and chemicals except ligand were purchased from available commercial sources and were used without further purification. Fourier transform infrared (FT-IR) spectra were obtained on a Nicolet Magna-IR 750 spectrometer with samples prepared as KBr pellets. C, H, and N microanalyses were carried out with a 2400 Series-II CHN analyzer, PerkinElmer. Magnetic measurements were performed using a Quantum Design VSM SQUID magnetometer. The measured values were corrected for the experimentally measured contribution of the sample holder, while the derived susceptibilities were corrected for the diamagnetism of the samples, estimated from Pascal’s tables. All of the compounds were characterized by elemental analysis, FT-IR spectroscopy, and single-crystal X-ray diffraction analysis. Crystallographic data and refinement parameters for all compounds are given in a tabular form (Table S1).

X-ray Crystallography

X-ray diffraction intensities for compounds 1–3 were collected at 100 K on Bruker D8 Venture APEX-3, using Mo Kα radiation with a complementary metal-oxide semiconductor detector. The structures were solved by direct methods in SHELXS and refined by full-matrix least squares on F2 in SHELXL.[25] CIF files for the structures reported in this article have been deposited at the Cambridge Crystallographic Data Centre (CCDC). The CCDC numbers are 1864470–1864472. Crystallographic data and refinement parameters for all compounds are given in a tabular form (Table S1). Bond distances and bond angles of all of the compounds are summarized in Table S2.

Selective Dye Sequestration

In the experiment carried out in the absence of H2O2, freshly prepared compound 1 (20 mg) in a powder form was added to the aqueous solutions of methyl orange (MO), methylene blue (MB), rhodamine B (RhB), and Congo red 2 (BR-2) (5 mL, 10 mg L–1), respectively. After 48 h of impregnation, the compounds were washed with ethanol and dried in air. Similar experiments were performed in the presence of catalytic amount of H2O2, and within 2 h, only methylene blue solution was decolorized.

Synthesis

Caution! Although we do not encounter any problem, appropriate care should be taken in the use of the potentially explosive perchlorate salts.

Synthesis of H2CPCA

Synthesis of Ethyl 4-(3-Cyanophenyl)-2-hydroxy-4-oxobutanoate (A)

Potassium tert-butoxide (4.5 g, 40.10 mmol) was taken in dry diethyl ether (200 mL) at room temperature under a N2 atmosphere under stirring conditions. Diethyl oxalate (4.4 mL, 30.10 mmol) was added to this solution. After 15 min of stirring, 3-acetylbenzonitrile (4.35 g, 29.96 mmol) was added to the resulting mixture, and stirring was further continued for overnight. A light yellow precipitate was obtained, and the whole resulting mixture was taken in a separating funnel for washing with water. Approximately, 20 mL of water was used for this purpose. The entire aqueous layer was then transferred in a beaker and cooled for some time. The resultant solution was then acidified with dilute HCl (2 mL). This results in a yellow precipitate of ethyl 4-(3-cyanophenyl)-2,4-dioxobutanoate, which was filtered and air-dried (3.3 g, 76.74% yield).

Synthesis of Ethyl 3-(3-Cyanophenyl)-1H-pyrazole-5-carboxylate (B)

Hydrazine monohydrate (0.65 mL) was added dropwise to a ethanolic (40 mL) solution containing ethyl 4-(3-cyanophenyl)-2,4-dioxobutanoate with constant stirring. The resulting solution was stirred and refluxed for 24 h. The reaction mixture was allowed to cool to room temperature, and the solvent was evaporated under reduced pressure. The residue was washed with water and recrystallized from methanol. 4-(3-Cyanophenyl)-2,4-dioxobutanoate was obtained as a light yellow crystalline powder (3.44 g).

Synthesis of 3-(3-Carboxyphenyl)-1H-pyrazole-5-carboxylic Acid (H2CPCA)

Dried 4-(3-cyanophenyl)-2,4-dioxobutanoate (3.44 g, 14.27 mmol) was added portionwise in a water solution of NaOH (880 mg, 22 mmol) under stirring conditions and refluxed for overnight. A clear golden yellow solution was obtained; then, the solution was filtered by a Whatman filter paper. The resultant solution was made just acidic using 6 N HCl. This leads to an off-white precipitate, which was then filtered, air-dried, and recrystallized from methanol (47.5% yield based on 3-acetylbenzonitrile).

1H NMR (dimethyl sulfoxide-d6, 500 MHz): δ 13.7 (3H, br s); 8.09 (3H, d); 8.35 (3H, t); 7.82 (3H, d); 7.53 (3H, t); 6.5 (2H, d).

IR (400–4000 cm–1): 3604 (m); 3197 (br); 2917 (m); 1696 (s); 1515 (s); 1424 (s); 1289 (s).

HRMS: calcd: 232.05, found (M + Na): 255.07.

Elemental analysis: calcd: C, 56.90; N, 3.47; H, 12.06. Observed: C, 52.70; N, 3.70; H, 11.27.

Synthesis of Compound 1

Aqueous solution (1 mL) of Co(ClO4), (37.2 mg, 0.01 mmol) and 2 mL of dimethylformamide (DMF) solution of ligand H2CPCA (23.2 mg, 0.01 mmol) were mixed and allowed to stir for 2 h at room temperature. Methanolic solution (3 mL) of 2,9-dimethyl-1,10-phenanthroline (20.8 mg, 0.01 mmol) was then added and the whole mixture solution was transferred to a Teflon-sealed container and allowed to stand for 72 h at 80 °C. Afterwards, the container was slowly cooled down to room temperature and light-pink, square-shaped crystals were obtained. The crystals were washed with methanol and then kept in dry air (72% yields). FT-IR (400–4000 cm–1): 3423.41 (br), 3134.11 (s), 2974.03 (s), 1745 (m), 1627.81 (s), 1504.37 (s), 1433.01 (s), 1406.01 (s), 1391.86 (s), 1309.58 (s), 1265.22 (s), 1020.27 (s), 871.76 (s). Elemental analysis for compound 1 (C55H47Co2N10O10): calcd: C, 58.62; H, 4.17; N, 12.43. Found: C, 57.92; H, 4.25; N, 11.89.

Synthesis of Compound 2

To a 3 mL methanolic solution of Co(ClO4)2 (37.2 mg, 0.01 mmol), 2 mL DMF solution of ligand H2CPCA (23.2 mg,0.01 mmol) was added. Then, the mixture was refluxed for 48 h at 70 °C and the resulting mixture solution was allowed to cool down to room temperature. Whitish-pink, block-shaped crystals were separated out, which were washed with acetone and kept in air for drying (84% yields). FT-IR (400–4000 cm–1): 3444.63 (s), 3359.77 (br), 3147.61 (s), 2918.10 (s), 2852.52 (s), 2511.15 (m), 1681.81 (s), 1589.23 (s), 1566.09 (s), 1500.52 (s), 1411.80 (s), 1344.29 (s), 1298 (s), 1271 (s), 1180.35 (m), 1207.99 (s), 979.77 (s), 804.26 (s). Elemental analysis for compound 2 (C22H18CoN4O10): calcd: C, 47.36; H, 3.22; N, 10.04. Found: C, 47.16; H, 2.98; N, 11.22.

Synthesis of Compound 3

For the synthesis of compound 3, a similar procedure to that used for compound 2 was followed. DMF solution (2 mL) of ligand H2CPCA (23.2 mg, 0.01 mmol) was added to 3 mL of methanolic solution of Mn(ClO4)2 (36.1 mg, 0.01 mmol). Then, the mixture was refluxed for 48 h at 70 °C and the resulting mixture solution was allowed to cool down to room temperature. White, square-shaped crystals were observed. Then, the crystals were filtered, followed by washing with acetone and keeping in air for drying (86% yields). FT-IR (400–4000 cm–1): 3446.56 (s), 3377.12 (br), 3149.54 (s), 2921.96 (s), 2852.52 (s), 2503.43 (br), 1681.81 (s), 1595.02 (s), 1494.73 (s), 1407.94 (s), 1342.36 (s), 1265.22 (s), 1020.27 (s), 802.33 (s). Elemental analysis for compound 3 (C22H18MnN4O10): calcd: C, 47.71; H, 3.25; N, 10.12. Found: C, 47.03; H, 2.62; N, 11.27.

Results and Discussion

Structural Description

Single-crystal X-ray structure analysis of compound 1 revealed that the compound crystallizes in the P21/n space group with one Co(II) ion, one CPCA2–, dmphen, and one-half of a DMF solvent molecule (Figure S6) in the asymmetric unit. Each cobalt atom acquires a distorted trigonal prismatic geometry, having a N3O3 coordination environment where two N atoms come from the chelating dmphen moiety and one N atom from the CPCA2– ligand, while the O atoms coming from two CPCA2– molecules (Figure b). The structural analysis shows that compound 1 forms an interesting one-dimensional (1D) helical network, facilitated by the V-shaped functional arrangement of the CPCA2– ligand and sustained by π–π stacking between dmphen rings. It is noteworthy here that the conformation of the ligand molecule, especially the angle between coordinating sites linked by a spacer with specific positional orientation, is considered to be the key factor for helical coordination network formation.[26−29] The common trend is that helical network formation becomes favorable as the angle tends to 120°.

Figure 1

(a) Right-hand side single-strand helix of compound 1; (b) packing diagram of compound 1 and coordination geometry around the Co(II) ion (inset). (c) Right-handed isomer (λ-isomer) and left-handed isomer (Δ-isomer); (d) π–π stacked double-stranded helix resulting from interweaving two single-strand helixes.

For compound 1, the helical 1D network propagated through the CPCA2– ligand where one side is coordinated to the carboxylate group and another side to the chelating carboxylate–pyrazole moiety (Figure ). The helical pitch, given by the distance between equivalent atoms generated by one full rotation of the screw axis, can be defined by the alternate Co···Co distance of 10.856 Å. However, a closer inspection of the structure reveals an intriguing aspect of the handedness of helical chains with the coexistence of two kinds of helical networks, a left-handed and a right-handed helical network (Figure b). The alternate helical chain is composed of one particular type of helical chain, either left-handed or right-handed. Such juxtaposition of helical networks brings their dmphens molecules very close to each other to form π–π interactions[30] with 3.72–3.80 Å face-to-face separation (Figure d).

Figure 2

(a) Molecular form of compounds 2 and 3. (b) Coordination geometry around the central metal (Co, Mn) ion. (c) Interplanar spacing diagram of alternative planes.

Crystal Structure of Compounds 2 and 3

Both compounds 2 and 3 crystalize in the Pbca space group and are isostructural in nature. The asymmeric unit of compound 2 (3) contains two HCPCA– ligands, two H2O molecules, and a metal ion (Figure ). The metal ion acquires octahedral geometry having a N2O4 coordination environment (Figure b). The basal plane was formed by two N atoms of two pyrazole moieties of two ligands and two O atoms from the −COOH group of two HCPCA– ligands. The apical positions are occupied by two water molecules. The packing diagram shows that the alternative planes are equidistanced (7.427 Å) and that they are parallel (AA/BB) to each other (Figure S7).

Solid-State Circular Dichroism (CD) Analysis of Compound 1

Individual presence of left- and right-handed helical networks prompted us to study the possibility of forming chiral inorganic complexes via spontaneous resolution. Recently, several reports confirmed that chiral inorganic complexes can be generated through spontaneous resolution even in the absence of any optically active ligands. For compound 1, solid-state CD investigations were carried out using a diffused reluctance technique at room temperature over the KBr matrix. The compound shows a positive cotton effect with two broad signals at 251 and 309 nm (Figure ). The experiment was repeated several times from different batches of crystals, and same CD spectra were obtained, confirming the presence of chirality. The most probable reasons for chirality could be the local environment of the octahedral metal center and subsequent transmission of chirality from the metal center to the helical network. Furthermore, this is consistent with earlier reports that when more than one diimine (like 2-2′-bpy, 1,10-phen) chelating ligands prefer to bind to an octahedral metal center characteristic peaks due to the Cotton effect arise in the CD spectrum.[31]

Figure 3

Solid-state CD spectrum of compound 1.

Selective Dye Binding with Compound 1

Organic dyes such as methylene blue (MB), rhodamine B (RhB), and Congo red are the most common organic pollutants of water. In recent years, inorganic polymeric compounds like CPs have been widely used for wastewater treatment.[32] Organic dye sequestration by CPs usually has two different routes: (a) chemical reaction at the surface of the CPs and (b) photocatalytic degradation in the presence of an inorganic catalyst.[33] The later process is particularly suitable for CPs based on metal, which can easily adopt variable oxidation states during the catalytic cycles. Keeping this in mind, we explore the dye degradation capability of compound 1. Considering the insolubility of compound 1 in water, a heterogeneous photocatalytic process using visible light was investigated. Accordingly, freshly prepared compound 1 in a powder form was added into the aqueous solution of methylene blue, rhodamine B, methyl orange, and Congo red, separately. After 48 h, the methylene dye solution was decolorized, while the solutions containing other dyes retained their respective colors (Figure ). This suggests that compound 1 itself selectively degrades methylene blue over other dyes. However, the rate of degradation was relatively slow. This led us to study the degradation with the Fenton oxidation procedure using hydrogen peroxide, which generates hydroxyl radicals and in the presence of Co(II) CPs provides the activation to accelerate the process for effective degradation of organic pollutants. Indeed, upon addition of one drop of catalytic H2O2, dye degradation took place at a considerably faster rate and bleaching of bright blue color of the dye completed within 60 min. The enhancement in the rate could be linked to the attacks of OH• radicals, initiating the oxidation procedure for the degradation, while a plausible mechanism can be envisaged as followsFurthermore, to determine the rate of dye degradation, kinetic studies were performed using UV–vis and photoluminescence (PL) spectroscopy techniques individually both in the presence and absence of H2O2. Accordingly, time-dependent UV–vis spectra of the dye solution containing compound 1 were recorded, which show a distinct trend in the absorption maxima of the dye at 665 nm. Over time, the intensity of the peak shows a steady decrease with a concomitant bleaching of the blue color of the dye after about 1 and 35 h in the presence and absence of H2O2, respectively. Both the dye degradation processes were found to follow pseudo-first-order kinetics for which the rate law can be represented aswhere C is the concentration of dye at time t, A is the absorbance of dye at time t, C0 is the concentration of dye at t = 0, A0 is the absorbance of dye at t = 0, and Kapp is the apparent rate constant.

Figure 4

(a) Photograph of time-dependent noticeable photodegradation of methylene blue dye solution by compound 1 in the presence of H2O2 (within 60 min) and (b) in the absence of H2O2 (within 48 h). (c) Comparative picture showing selective methylene blue dye degradation in comparison to that of Congo red, methyl orange, and rhodamine B.

The rate constants were subsequently determined by plotting ln(C/C0) versus time, which fairly follows the pseudo-first-order kinetics with a linear fitting[34] (Figure ). Apparently, compound 1 shows higher activity as a catalyst to degrade the dye in a Fenton-like process with Kapp = 8 × 10–2 s–1 compared to that in the nonradical process, showing Kapp = 9.9 × 10–3 s–1 (Table ). Furthermore, the dye degradation efficiency was also calculated from the data, and it was found that degradation efficiencies are ∼82 and ∼98% in the absence and presence of H2O2, respectively.

Figure 5

Time-dependent UV–vis spectral study of dye degradation by compound 1 in the (A) presence and (B) absence of H2O2. Photodegradation efficiency in the (C) presence and (D) absence of H2O2. Plot of ln(C/C0) vs T in the (E) presence and (F) absence of H2O2.

Table 1

Rate Constant and Degradation Efficiency of Compound 1 in Different Conditions

reaction condition	Kapp (s–1)	degradation efficiency (%)
in the presence of H2O2	8.00 × 10–2	98
in the absence of H2O2	9.9 × 10–3	82

Moreover, photoluminescence property was further investigated for compound 1, both with and without H2O2. Interestingly, when the experiments were carried out in the absence of H2O2 and a decrease in intensity was a recurrent theme, a red shift of ∼53 nm (λem = 734 nm) (Figure ) was also observed. This unusual luminescent red shift can be attributed to the interaction of dye with compound 1.[35]

Figure 6

(a) PL spectra of compound 1 in the presence and (b) absence of H2O2; notice a clear red shift in the emission spectrum.

The powder X-ray diffraction (PXRD) pattern recorded for compound 1 after the degradation of dye (Figure ) shows an interesting feature. Although the PXRD patterns satisfactorily match with the simulated pattern, some additional new peaks also appeared. When these samples were further analyzed by FT-IR spectra, some interesting observations were noted. Although the fingerprint region before and after interaction with the dye remains the same, shifting of some characteristic peaks was observed (e.g., 3423–3275 nm; 2974–2751 nm; 1627–1447 nm; 1391–1206 nm), which can be attributed to the interaction of the dye molecule and compound 1. To inspect whether compound 1 adsorbs the dye molecule or has any accessible void space, Brunauer–Emmett–Teller (BET) adsorption was studied and surface area was calculated (Figure S9). However, it shows very poor adsorption capability with BET surface area of 2.6058 m2 g–1. Therefore, it can be concluded that the dye molecule interacts with compound 1 through a helical grove but does not adsorb at the surface.

Figure 7

(a) Comparison of the simulated PXRD pattern and the PXRD patterns before and after dye interaction with compound 1; comparison of PXRD patterns of compound 1 with and without H2O2.

Magnetic Study

Direct current (dc) magnetic susceptibility data of all of the compounds were collected on polycrystalline samples in the temperature range of 2–300 K at a field of 0.1 T (Figure ). The room temperature χMT values are 6.27, 2.76, and 3.94 cm3 mol–1 K for compounds 1, 2, and 3, respectively. For compounds 1 and 2, experimental χMT values are higher than spin-only values of 3.75 and 1.875 cm3 mol–1 K for two and one Co(II) centers, respectively. Higher experimental values can be explained on the basis of orbital contribution of the Co(II) ion (Table ). The experimental value of compound 3 is slightly lower than the theoretical values of 4.375 cm3 mol–1 K for one Mn(II) spin center having S = 5/2, g = 2.0.

Figure 8

Temperature-dependent dc magnetic susceptibility of compounds 1–3. The red lines are the best fit.

Table 2

Comparison of Experimental and Theoretical χMT (cm3 mol–1 K) Values of Compounds 1 and 2

compound	theoretical spin only, χMT	theoretical (spin + orbital), χMT	experimental
1	3.75	6.705	6.27
2	1.875	3.352	2.76

For all three compounds 1–3, χMT values gradually decrease with the temperature probably due to weak antiferromagnetic interactions with the nearest metal centers or due to the crystal field effect.

Isothermal magnetization (M/Nβ vs H) for compound 1 was plotted for the temperature range of 2–10 K from 0 to 5 T, which shows that the plot does not saturate even at the highest field of 5 T (Figure ). The maximum M/Nβ value of 5.5Nβ was obtained at 5 and 2 K. The maximum experimental value is lower than the maximum theoretical value of 6Nβ, suggesting the presence of antiferromagnetic interactions or anisotropy. To examine the anisotropy of the system, M/Nβ was plotted versus H/T and this shows that all of the isotherms do not fall on the same master plot, which clearly reveals the existence of anisotropy of compound 1. Therefore, we were interested in studying alternating current (ac) measurements. Compounds 2 and 3 did not show any significant peak for SMM. If we closely look at the geometry around Co(II) centers of compounds 1 and 2, compound 1 has distorted trigonal prismatic geometry, which may originate magnetic anisotropy suitable for slow relaxation of magnetization.[12] On the other hand, compound 2 has almost perfect octahedral geometry, which may be responsible for the absence of any slow relaxation of magnetization.[36] The anisotropy parameters of Co(II) centers in compounds 1 and 2 have been calculated by fitting the susceptibility data and magnetization data using PHI software;[37] during the fitting procedure, the g tensor was considered to be isotropic. The best fits of the experimental data gave D = 38.2(4) cm–1, E = 1.2(7) × 10–4 cm–1, and g = 2.28 for 1; D = 31.5(1) cm–1, E = 4.6(4) × 10–2 cm–1, and g = 2.21 for 2 (where D and E represent the single-ion axial and rhombic zero-field splitting parameters, respectively). The values of the anisotropy parameters for 1 are found to be good enough for six-coordinated Co(II) complexes, showing single-ion magnetic behavior.[38] The positive sign of the D parameter arises from the interaction between the ground and excited electronic states coupled through spin–orbit coupling. However, for compound 2, the D value may not be good enough in that temperature region due to the geometry and crystal field effect for not showing any SMM behavior.

Figure 9

(a) M/Nβ vs H plot and (b) M/Nβ vs H/T plot for compound 1. The red lines are the best fit.

To check the type of exchange interaction between nearby Mn(II) centers in compound 3, we have attempted to fit the susceptibility data using PHI software. The best fit resulted in JMn-Mn = −0.11 ± 0.05 cm–1 and g = 2.0 ± 0.04 (Figure ), suggesting very weak antiferromagnetic interactions between the nearby Mn(II) centers in compound 3. J indicates the exchange interaction between the nearby neighbors.

To investigate the SMM behavior of compound 1, ac susceptibility measurements were carried out in the temperature range of 2–10 K. At 0 dc field, both in-phase and out-of-phase susceptibility measurements did not show any significant peak due to quantum tunneling of magnetization (QTM). To suppress the QTM, ac measurements were performed at different dc fields. At an optimized dc field of 0.5 T, both in-phase and out-of-phase maxima were observed (Figure ). Temperature-dependent out-of-phase ac susceptibility plots reveal the existence of slow relaxation of magnetization and SMM behavior.

Figure 10

Temperature dependence of the in-phase (χ′) (a) and out-of-phase (χ″) (b) ac susceptibility for compound 1 under a field of 0.5 T.

To estimate the relaxation time and energy barrier of compound 1, ln(ω) was plotted versus 1/T and fitted using the Arrhenius equation: ln(ω) = ln(1/τ0) – (Eeff/kT) (Figure ), where ω is the frequency, τ0 is the relaxation time, and Eeff is the energy barrier. From the fitting, energy barrier (E/k) = 9.2 K and relaxation time (τ0) = 9.1 × 10–5 s were obtained. The relaxation time is in the range of SMM.

Figure 11

Plot of ln(ω) vs 1/T for complex 1. The solid line indicates fitting of the plot using the Arrhenius equation.

Conclusions

The major impetus for this study was to generate multifunctional coordination complexes using a novel pyrazole-based ligand. Accordingly, we have successfully synthesized three coordination compounds, among which a coordination polymer exhibits multifunctional properties. The helical CP shows an interesting chirality originated from the local coordination environment of the metal ion with corroborating CD spectrum. The Co(II)-based CP shows an interesting magnetic feature of single-ion magnetism, while the corresponding relaxation time falls well within the normal range of SMM. The magnetic results presented herein clearly indicate our strategy of using Co-based CP to maximize the relaxation barrier through increasing magnetic anisotropy for achieving SMM behavior. The Co-based CP also shows excellent photocatalytic degradation of toxic dye molecules with a potential application in wastewater purification.