Preparation and Analyses of the Multifunctional Properties of 2D and 3D MOFs Constructed from Copper(I) Halides and Hexamethylenetetramine.

In this article, two two-dimensional and three-dimensional metal-organic frameworks are synthesized by the self-assembly of copper(I) halide and the hexamethylenetetramine (hmt) ligand. Compound 1 is a two-dimensional metal-organic framework composed of a pyramidal Cu4I5 cluster and hexamethylenetetramine, in which hmt-bridged Cu clusters form a two-dimensional (4,4)-connected net with a point symbol of (44·62) (44·62). Compound 2 is a homochiral three-dimensional metal-organic framework material generated through an unusual spontaneous crystallization from achiral precursors. The two compounds were characterized by a series of analyses such as infrared spectroscopy, elemental analysis, circular dichroism spectroscopy, and powder X-ray diffraction. Both of them exhibit unexpected stability under a wide range of conditions of acid and base. In addition, the fluorescence intensity changes regularly under acid-base conditions. Stokes shift shows a good linear relationship with -log [H+], which makes them become promising acid-base sensors. Compounds 1 and 2 also display selective adsorption and a significant degradation effect on the organic dye methylene blue. In addition, the fluorescence study indicated that compound 2 could be used as a sensor to detect Cr3+.

Introduction

At present, the assembly of n class="Chemical">metal–organic frameworks is of great significance because of their high specific surface area, porosity, and plasticity, together with structural diversity and interesting topologies. In some aspects of chirality, gas storage, catalysis, zeolite behavior, etc., they show potential applications as functional materials.[1−4] One of the basic strategies for the construction of functional materials is the employment of organic ligands with tetrahedral nodes. Since the hexamethylenetetramine (hmt, by name urotropine) molecule contains four tetrahedral nitrogen atoms after coordination, it becomes a potential multifunctional tetrahedral node. The crystal structures of copper chloride and hmt have been obtained, but the desired crystal structure has not been obtained with copper(I) bromide and iodide as inorganic building centers. This article helps overcome this challenge.[5] From 1923, the molecular structure of hmt began to enter our field of vision.[6,7] Due to the fascinating structure of cage trialkylamine, there has been great interest in the fields of organic synthesis, chemical industry, molecular catalysis, disinfectants, pesticides, curing agents, pharmaceutical physics, and military applications.[8] Hmt exhibits a variety of coordination modes including monodentate(t-), bidentate(μ-), tridentate(μ3-), and tetradentate(μ4-) ligands in combination with metal salts and produces a large number of fascinating functional metal–hmt complexes.[9] It is worth mentioning that there exist coordinatively unsaturated sites (refers to nitrogen) for the first three coordination configurations(t-, μ-, μ3-), which may have laid the foundation for ion recognition. However, it is very difficult to achieve pure chirality from the achiral precursor hmt. The addition of an inducer has paved the way to a new field in the development of pure chiral materials.[10] Due to the relative earth abundance of these metals, the participation of copper species can form functional mononuclear or multinuclear metal–cluster complexes. This has led to an interest in the development and utilization of copper metal salts.[11−18]

In this paper, {[hmt][n class="Chemical">Cu4I5][H2O]5[NH4]+}(1) and {[hmt][Cu3Br3][H2O]}(2) were synthesized. The formed Cu4I5 cluster is connected to four N atoms on hmt. It turns out that multifunctional materials may be produced by clusters instead of metal centers.[19] Since the copper iodide clusters form Cu4I4 (or Cu3I4) as four-connected nodes,[20] the clusters are diverse in form and complicated in structure. The formation of Cu4I5 clusters in this paper may give the material unique functionality. Further, copper(I) bromide forms a three-coordinate structure with hexamethylenetetramine to form a three-dimensional structure. It was also synthesized for the first time.

Our team has long been committed to the construction of organic–inorganic hybrid compounds and the analysis of their physicochemical properties.[21−25] We further studied the properties of the compounds and found that they have high stability under acid–base conditions. The detection of acid–base stability to organometallic compounds was common,[26−29] but there are few examples of this in the field of organic–inorganic hybrid compounds, mainly because of their instability under acidic or basic conditions.[30] Further, compounds 1 and 2 have been found to have good photocatalytic degradation properties for the organic dye n class="Chemical">methylene blue. Since trivalent metal ions play an important role in the human body and in the environment, we chose Cr3+ to evaluate the characteristics of fluorescent probes and found that compound 2 exhibited a high-sensitivity and high-selectivity fluorescence response to Cr3+. It is known that due to complexation with metal ions, charge-transfer interaction is reduced to enhance the fluorescence signal.[31] We conclude that the availability of lone-pair electrons on the uncoordinated nitrogen of compound 2 without coordination will promote the complexation with metal ions.

Results and Discussion

Description of Crystal Structures

Compound 1 indicates that {[hmt][n class="Chemical">Cu4I5]−[H2O]5[NH4]+} (1) crystallizes in a tetragonal space group P4/nmm. The asymmetric unit of the compound contains one hmt ligand and [Cu4I5]− clusters, as shown in Figure S1a. Interestingly, the geometry of the anionic [Cu4I5]− cluster is a square pyramid with four Cu and five I atoms (Figure S1b). The compound is formed by the [Cu4I5]− clusters and the ligand hmt. In compound 1, the hmt ligand is tetradentate (μ4-hmt), linking to four metal ions (Figure a). Two hmt molecules form a 10-membered ring with two [Cu4I5]− clusters. The distance between Cu···Cu in the 10-membered ring is 6.584 Å and between C···C is 3.529 Å (Figure S1c). In view of the topology, however, the minimum closed loops around the nodes involved in symmetry-related ligands and Cu1 ions are (44·62). Thus, the two-dimensional (2D) structure of compound 1 can be described as an interesting four-connected net with the point symbol of (44·62) (44·62) in which Cu(I) centers and ligands act as four-connected nodes (Figure b).

Figure 1

(a) Coordination environment of compound 1 (hydrogen atoms have been omitted for clarity). (b) Topological structure of the 2D network in 1. (c) Stacked view from the a-axis direction. (d) Single-stranded right-handed helical strand twisted along the b-axis.

(a) Coordination environment of compound 1 (hydrogen atoms have been omitted for clarity). (b) Topological structure of the 2D network in 1. (c) Stacked view from the n class="Species">a-axis direction. (d) Single-stranded right-handed helical strand twisted along the b-axis.

The compound {[hmt][n class="Chemical">Cu3Br3][H2O]} (2) crystallizes in the chiral space group P43 with a Flack parameter of 0.09(12). Single-crystal X-ray analysis reveals that 2 represents a rare three-connected three-dimensional chiral network. The structural unit of 2 is shown in Figure S2a. In compound 2, the hmt ligand is tridentate (μ3-hmt), linking to three metal ions (Figure c). There are three different coordination environments for Cu atoms in the structural unit diagram. The Cu1, Cu2, and Cu3 atoms are all coordinated by the N atom from the hmt ligand. Cu1 and Cu2 ions are interconnected by two μ2-Br1 and μ2-Br2, respectively, to form a dinuclear Cu cluster. From the b-axis direction, a distorted quadrangle is formed (Figure S2b). The distorted quadrilateral consists of two hmt ligands and a small quadrilateral formed by two copper atoms and two bromine atoms. The distance between Cu···Cu in the quadrangle is 10.857 Å. From the a-axis direction, units of compound 2 are arranged in parallel to each other in the ABAB manner to produce a chiral structure (Figure c). It forms right-handed helixes along the b axis with large pitches of 27.2948(55) Å. The ligand and metal ion Cu(I) assembly yields a crystal helical chirality (Figure d). Synthesis of a homochiral network that can be established by three methods has been reported: (i) using enantiopure precursors, (ii) through spontaneous resolution occurring during crystallization, and (iii) using achiral precursors in the presence of chirality-inducing agents such as chiral additives, chiral solvents, or chiral catalysts.[41c] The chiral compound 2 in this paper is synthesized by self-assembly and crystallization in solution using CuBr and an organic ligand (hmt), which should be assigned to method (ii).

Acid–Base Resistance

We soaked compounds 1–2 in an aqueous solution of acid and base to explore their stability to acid and base. Compounds 1–2 of the same quality were soaked for 1 day at a range of pH values and then subjected to IR spectroscopy measurements. Compound 1 can maintain its structure at pH = 3–13. In the infrared spectrum, the stretching vibration peak of the C–N group at about 1366 cm–1 gradually disappears, indicating that the acid concentration changes, which causes the skeleton to collapse. There is no peak at pH = 1 and pH = 2 at 3257 cm–1 (Figure S3a). However, for compound 2, at pH = 1, there is no peak at 3257 cm–1. This indicates that compound 2 can remain structurally intact in the range of pH = 2–13 (Figure S3b).

Photoluminescence Properties

The solid-state fluorescence properties of compounds 1–2 at normal temperature were studied, and the fluorescence intensity of compound 1 was weaker than that of compound 2. The corresponding luminescent color of compound 1 is yellow and that of compound 2 is green as shown by CIE (x, y) (Figure a). The emission wavelength of the compounds shifts; we further speculate that the emission of the compound may be caused by a different n class="Chemical">halide–metal charge transfer. In view of the intensity of fluorescence, we monitored the solvent fluorescence changes of compounds at different pH values.[32] The pH is adjusted herein using an aqueous solution of hydrochloric acid and sodium hydroxide to record the corresponding fluorescent emissions. The emission wavelength of compound 1 in an aqueous solution was 383 nm (λex = 290 nm). As the pH gradually reached pH = 3, the maximum emission wavelength was 412 nm. Throughout the experiment, the total Stokes shift reached 20 nm (Figure b). In addition, the Stokes shift showed an overall blue shift, after adding the base. However, as the pH increases (Figure c), the changes in compound 1 are as follows: pH = 8 (314 nm) < pH = 9 (332 nm) < pH = 10 (348 nm) < pH = 12 (353 nm) < pH = 13 (355 nm). In addition, the results show that the Stokes shift has a linear relationship with −log [H+] (Figure d,e). Compound 2 gradually decreases in fluorescence intensity with increasing pH in an acidic aqueous solution (Figure f). As the pH in the alkaline aqueous solution increases, the fluorescence intensity of compound 2 is gradually enhanced (Figure g). The reason for the change in fluorescence intensity under different acid–base conditions is that when a certain group in the compounds is protonated or deprotonated, it is possible to induce a change in the fluorescence intensity of the compounds. The electrons of the electron donor are transferred to the electron acceptor, which produces a fluorescence-quenching phenomenon. When the electron donor is combined with the proton, the electron transfer is inhibited, thereby causing the recovery of fluorescence.[32c]

Figure 2

(a) Emission spectra of compounds 1 and 2 in the solid state at room temperature. (b) Fluorescence analysis of compound 1 in the pH range of 3–6 was processed. (c) Fluorescence analysis of compound 1 in the pH range of 8–13 was processed. (d) Linear relationship between the Stokes shift at pH = 3–6 and the −log [H] value of the solution. (e) Linear relationship between the Stokes shift at pH = 8–13 and the −log [H+] value of the solution. (f) Fluorescence analysis of compound 2 in the pH range of 2–6 was processed. (g) Fluorescence analysis of compound 2 in the pH range of 8–13 was processed.

Dye Adsorption

Organic dyes are used in many fields.[33−35] Therefore, a method for extracting dyes from a solution has been of wide concern.[36−38] The absorption of dye molecules by the use of organon class="Chemical">metallic compounds is very popular. Methylene blue (MB) is used herein as the adsorbed dye. Compared with compound 2, compound 1 has a more obvious adsorption effect at 210 min (Figures S5 and S6). The amount of adsorbed dye can be calculated by the following formula: Qeq = (C0 – Ceq)V/M.[39] It is known from the figure that the adsorption effect of compound 2 is weaker than that of compound 1 (Figure S7), and the adsorption amount of compound 2 is also increased by prolonging the adsorption time. Therefore, both compounds 1 and 2 have a good adsorption capacity for the organic dye MB.

Photocatalytic Activities

The UV–vis diffuse reflectance spectra of a powder sample with a band gap (Eg) of 1 is used, which is reflected using reflectance. The Kubelka–Munk function F is plotted on the energy map E, and F = (1 – R)2/2R is converted from the recorded diffuse reflection data.[40] The gap values are 1.55 and 1.87 ev, as shown in Figure S8. This indicates that compounds 1–2 are semiconducting when exposed to visible light and have possible photocatalytic activity. The photocatalysts (compounds 1–2) are semiconductor materials. Therefore, the organic dye MB was photodegraded by the photoexcitation of compounds 1–2. The degradation efficiency is mainly determined by the free carrier available on the surface of the photocatalysts. Semiconductor materials have been shown to have higher electron transport rates and better material properties. The photocatalytic degradation efficiencies of the organic dye n class="Chemical">MB with respect to compounds 1–2 are shown in Figures S9 and S10, and the absorption spectrum of MB is remarkably lowered as the reaction time increases. Furthermore, the change of the concentration of the MB solution with time was plotted (C0: initial concentration of MB; C: dye concentration at any given time). The results showed that the degradation rate of compound 1 in 40 min was 70%, whereas the degradation rate of compound 2 in 140 min was 16% (Figure a,b). Compound 1 has a high decomposition activity for MB under visible-light irradiation. We studied the reaction kinetics. The catalytic process conforms to the quasi-first-order kinetic equation as shown in Figure c,d. Then, we tested the stability of compounds 1–2 before and after catalysis, and the characteristic peaks before and after the contrast infrared image basically existed (Figure S11). To evaluate the recoverability and reusability of compounds 1 and 2, photocatalytic degradation of MB for three cycles was performed. As can be seen from Figure e,f, similar photocatalytic efficiencies were maintained even after repetition. It can be seen from Figure S11c,d that the powder X-ray diffraction (PXRD) pattern of the compound after photodegradation is almost the same as that of the synthetic sample, which further proves that compounds 1 and 2 do not degrade under photodegradation conditions.

Figure 3

(a) Compound 1 and blank photocatalytic degradation of MB dye. (b) Compound 2 and blank photocatalytic degradation of MB dye. (c) Photocatalytic reaction kinetics of compound 1. (d) Photocatalytic reaction kinetics of compound 2. (e) Recycling test with compounds 1 and 2 (f) for MB photodegradation.

(a) Compound 1 and blank photocatalytic degradation of MB dye. (b) Compound 2 and blank photocatalytic degradation of n class="Chemical">MB dye. (c) Photocatalytic reaction kinetics of compound 1. (d) Photocatalytic reaction kinetics of compound 2. (e) Recycling test with compounds 1 and 2 (f) for MB photodegradation.

Circular Dichroism (CD)

Since compound 2 has chiral space groups, solid-state circular dichroism (CD) spectra have been measured for further demonstrating its homochiral nature.[41] The different peaks according to positive and negative Cotton effects appear in the range of 200–260 nm, which are shown in the CD spectra of compound 2 (Figure ) The CD spectrum of 2 exhibits positive Cotton effects centered at 210, 221, and 233 nm and negative Cotton effects centered at 202, 227, and 257 nm. The results well-testify the chirality of compound 2.

Figure 4

Solid-state CD spectra of compound 2.

Sensing of the Metal Cations

Therefore, we use the potential fluorescence properties of chiral compound to detect various metal ions such as K+, n class="Chemical">Cu2+, Zn2+, Cd2+, Co2+, Pb2+, Ni2+, Ba2+, Cr3+, and Fe3+. However, chiral compound recognition is a chiral match between a chiral host and an object, just like a hand and a glove.[42] As shown in Figure a, it is observed that compound 2 has high sensitivity and specificity for the selectivity of Cr3+ ions in aqueous solution. Trivalent chromium (Cr3+) plays an important role in the diet of humans and animals. It is also a stabilizer for nucleic acids, which can prevent cancer. The lack of trivalent chromium can cause diabetes, arteriosclerosis, and high blood pressure. Excessive intake can lead to genotoxic effects. In addition, how to use the convenient and quick method to detect trivalent chromium is an urgent problem to be solved, due to its diversification and participation in industrial and agricultural activities, making trivalent chromium an environmental pollutant.[43] The fluorescence intensity at 392 nm was found to be enhanced. From the fluorescence intensity ratio of different metal ions in the aqueous solution, it is found that the intensity ratio on adding Cr3+ is 22 times that of the blank (Figure b). In contrast, other metal ions would have no obvious influence on the fluorescence intensity of compound 2. The luminescence intensity gradually increases with the increase of Cr3+ ion concentration. The fluorescence intensity of compound 2 at 392 nm is proportional to the concentration of Cr3+ within a certain range. A linear relationship of Cr3+ detection under optimal conditions was obtained at 392 nm with a correlation coefficient of 0.996 (Figure S12). The increase in fluorescence intensity may be due to the formation of an extended electron conjugation system between metal ions, and it is also possible that the chelation of metal ions with compound 2 produces stiffness in the molecule. The regression equation is Y = 282.533 +19.503X. According to the reported method,[44,45] the limit of detection for Cr3+ is up to 0.227 μM. The relationship between the fluorescence intensity and mole fraction is shown in Figure S13. The results show that the emission intensity is close to the maximum when the mole fraction of [Cr3+]/([2]+[Cr3+]) is about 0.5, indicating a 1:1 binding stoichiometry between 2 and Cr3+. All of these results are consistent with the coordination configuration of the ligand μ3-hmt with a single active coordinate N site in compound 2. A possible sensing mechanism of the fluorescence enhancement of compound 2 in the solvent media is proposed, wherein the solvated Cr3+ with a larger ion radius forms complexes with 2 as shown in Figure .[45a] The metal cation Cr3+ was bound to the lone pair on the uncoordinated amino group nitrogen of μ3-hmt. The interaction of the cation Cr3+ with the lone pair on the nitrogen atom in hmt results in an inhibition of the photoelectron transfer process, giving rise to the observed enhancement of fluorescence emission.[45b−45d]

Figure 5

(a) Fluorescence intensity of compound 2 in aqueous solution with 1.0 × 10–2 mol/L metal ions (K+, Cu2+, Zn2+, Cd2+, Co2+, Pb2+, Ni2+, Ba2+, Cr3+, and Fe3+); λex = 345 nm. (b) Fluorescence intensity ratio of different metal ions in aqueous solution (I and I0 denote the fluorescence intensities of compound 2 with and without the metal ions of interest, respectively).

Figure 6

Possible sensing mechanism of compound 2 for the recognition of Cr3+.

(a) Fluorescence intensity of compound 2 in aqueous solution with 1.0 × 10–2 mol/L metal ions (K+, n class="Chemical">Cu2+, Zn2+, Cd2+, Co2+, Pb2+, Ni2+, Ba2+, Cr3+, and Fe3+); λex = 345 nm. (b) Fluorescence intensity ratio of different metal ions in aqueous solution (I and I0 denote the fluorescence intensities of compound 2 with and without the metal ions of interest, respectively).

Thermogravimetric (TG) Studies

The TG curves of compounds 1–2 are provided in Figure S14. The frameworks of compounds 1 and 2 were stable up to ca. 300 and 240 °C, and so the overall thermal stability is high. The TG trace of compound 1 exhibits two main steps of n class="Disease">weight loss, which correspond to the release of coordinate inorganic groups and decomposition of organic groups. The organic ligand of compounds 1–2 decompose, which causes a decrease of 250–300 °C. The decomposition of inorganic substances mainly occurs at 600–800 °C. Due to the difference in the coordination mode of the ligands, the stable temperatures of compounds 1 and 2 are quite different. It appears that compound 2 has only one mass drop and continuously loses weight from 300–800 °C, indicating that the decomposition of the ligand is completed in one step.

Conclusions

In summary, we produced two compounds by the self-assembly of n class="Chemical">copper(I) halides and a tetradentate hmt as a ligand. A two-dimensional metal–organic framework compound with Cu4I5 clusters was synthesized by reaction with CuI, and it was reacted with CuBr to form a three-coordinate three-dimensional chiral structure. The study found that the compounds have good acid–base stability and the Stokes shift has a good linear relationship with −log [H+]. In addition, we found that the compounds have the potential to adsorb and cause photocatalytic degradation of methylene blue by visible light. The highly selective and sensitive detection method of compound 2 for Cr3+ has been successfully applied to aqueous solutions.

Experimental Section

Synthesis of 1

Hexamethylenetetramine (n class="Chemical">hmt) (0.0014 g) was added to a stirred solution of colorless CuI (0.0019 g) in 1.5 mL of dimethylformamide (DMF) and 1 mL of NH3OH in the presence of excess KI (0.0066 g). The resulting mixture was stirred for 5 min and filtered. The filtrate was allowed to stand and volatilized at room temperature to form pale-yellow crystals for about 3 months, and the collection rate was 20.05%. IR (KBr cm–1): 3331 (m), 2980 (w), 1599 (w), 1448(w), 1623 (w), 1242 (s), 1215 (m), 1009 (s), 822(s), 701(s). Elemental analysis for C6H26Cu4I5N5O5 (%): Calcd: C, 6.34; H, 2.30; N, 6.16. Found: C, 6.38; H, 2.35; N, 6.21.

Synthesis of 2

Hexamethylenetetramine (n class="Chemical">hmt) (0.0014 g) was added to a stirred solution of colorless CuBr (0.0014 g) in 1.5 mL of DMF and 1 mL of NH3OH in the presence of excess KBr (0.0048 g). The resulting mixture was stirred for 5 min and filtered. The filtrate was allowed to stand and volatilized at room temperature to form pale-yellow crystals for about 3 months, and the collection rate was 35%. IR (KBr cm–1): 3441 (w), 3296 (s), 2888 (w), 1595 (s), 1457 (s), 1250 (s), 1028 (s), 991(s), 819 (s), 673(s). Elemental analysis for C6H14Br3Cu3N4O (%): Calcd: C, 12.24; H, 2.39; N, 9.52. Found: C, 12.20; H, 2.42; N, 9.56.