Sn(IV) Porphyrin-Based Ionic Self-Assembled Nanostructures and Their Application in Visible Light Photo-Degradation of Malachite Green

: A series of porphyrin-based ionic complexes were prepared through the reaction of two porphyrin precursors, 5,10,15,20-tetrakis(4-(2-pyridyl)phenyl)porphyrin H 2 TPhPyP ( 1 ) and trans -dihydroxo [5,10,15,20-tetrakis(4-(2-pyridyl)phenyl)porphyrinato]tin(IV) Sn(OH) 2 TPhPyP ( 2 ), with various acids (HCl, HNO 3 , CF 3 COOH, H 2 SO 4 , H 2 CO 3 , and H 3 PO 4 ). The complexes were characterized via elemental analysis, 1 H nuclear magnetic resonance spectroscopy, electrospray ion-ization mass spectrometry, Fourier transform infrared (FT-IR) spectroscopy, ultraviolet (UV)-visible spectroscopy, ﬂuorescence spectroscopy, and ﬁeld-emission scanning electron microscopy (FE-SEM). Each compound exhibited different results for UV-visible, ﬂuorescence, FT-IR, and FE-SEM studies depending on the counter anions. The complexes possessed different self-assembled nanostructures based on electronic interactions between the cations of compounds 1 and 2 with different counter anions. These aggregated species are stabilized by electrostatic forces and the π - π stacking interactions between the two porphyrin rings, in which the counter anions play an important bridging role. The counter anions also play an important role in controlling the morphology and photocatalytic properties of the as-developed materials. The complexes were then used for the photocatalytic degradation of the malachite green (MG) dye in aqueous media under visible light irradiation for up to 70 min. A morphology-dependent photocatalytic degradation of the MG dye was observed for all the ionic complexes, with efﬁciencies ranging from 50% to 95%. and ﬂuorescence FT-IR spectra were FE-SEM were using The (BET) area using N 2 isotherms at was measured with an The total organic carbon (TOC) values of the degradation were examined with a JEOL to the extent of mineralization.


Introduction
Agricultural and industrial activities generate significant amounts of toxic compounds that are discharged into water ecosystems, causing severe environmental issues. Large quantities of organic dyes, herbicides, phenols, pesticides, plasticizers, biphenyls, amino, and nitro compounds are discharged into water bodies, endangering aquatic life and water potability [1,2]. Several methods to remove hazardous pollutants from wastewater have been reported, including those based on physicochemical techniques, such as adsorption [3], membrane filtration [4], ion-exchange [5], chemical precipitation [6], electrolysis [7], and advanced oxidation processes (AOPs) [8]. Among them, AOPs are the most commonly used methods owing to their simple operation, high degradation rate without the generation of secondary pollutants, and degradation of the entire pollutant to less-toxic compounds (CO 2 and H 2 O). These methods are based on the degradation of pollutant molecules by reactive oxygen species (ROS) generated in situ using a suitable photocatalyst. Photocatalysts absorb light from sunlight and generate ROS, which facilitate the degradation of organic pollutants in water [9][10][11]. TiO 2 has been used as a promising photocatalyst owing to its high stability, low toxicity, and high photocatalytic efficiency for the degradation of pollutants in water. However, a large band gap (~3.2 eV) restricts the use of TiO 2 to only 1 H-NMR spectra of compounds 1 and 2 were recorded in CDCl3 and are shown in Figures S1 and S2. For compound 1, the characteristic signal at 2.73 ppm corresponds to the NH protons of the pyrrole ring, while the protons in the pyridyl ring appear at 8.84, 8.04, 7.90, and 7.35 ppm. The phenyl protons of compound 1 are at 8.35 ppm, and the βpyrrole protons appear at 8.94 ppm. Compound 2 exhibits a peak at 7.43 ppm corresponding to the axial Sn-OH protons. The pyridyl ring protons appear at 8.86, 8.07, 7.92, and 7.38 ppm. The β-pyrrole and phenyl protons of compound 2 appear at 9.23 and 8.45 ppm, respectively. ESI-MS spectra of compounds 1 and 2 are shown in Figures S3 and S4, respectively. Compound 1 has a molecular peak at 927.27 [1 + H] + , while 2 exhibits a peak at 1075.19 [2 + H] + . Furthermore, compound 2 has a low-intensity peak at 538.08, indicating the ion [2 + 2H] 2+ .
The UV-vis absorption spectra of compound 1 and its ionic complexes from various acids are shown in Figure 1. A sharp Soret band (419 nm) and four Q-bands (515, 551, 590, and 645 nm) were observed. After treatment with various acids to form the ionic complexes, the Soret band red-shifted (419→422 nm for HNO3, CF3COOH, and HCl; 419→442 Chart 1. Chemical structures of compounds 1 and 2 used in this study.

Syntheses and Characterization
Scheme 1 illustrates the preparation of compounds 1 and 2. First, H 2 TphPyP (1) was prepared from 4-(2-pyridyl)benzaldehyde and freshly distilled pyrrole in the presence of propionic acid under reflux conditions. Subsequently, compound 1 was treated with tin(II) chloride in the presence of pyridine, followed by hydrolysis, to form Sn(OH) 2 TphPyP (2). Compounds 1 and 2 were fully characterized using various spectroscopic techniques, including elemental analysis, 1 H-NMR spectroscopy, ESI-MS, UV-vis spectroscopy, fluorescence spectroscopy, and Fourier transform infrared (FT-IR) spectroscopy.

Syntheses and Characterization
Scheme 1 illustrates the preparation of compounds 1 and 2. First, H2TphPyP (1) was prepared from 4-(2-pyridyl)benzaldehyde and freshly distilled pyrrole in the presence of propionic acid under reflux conditions. Subsequently, compound 1 was treated with tin(II) chloride in the presence of pyridine, followed by hydrolysis, to form Sn(OH)2TphPyP (2). Compounds 1 and 2 were fully characterized using various spectroscopic techniques, including elemental analysis, 1 H-NMR spectroscopy, ESI-MS, UV-vis spectroscopy, fluorescence spectroscopy, and Fourier transform infrared (FT-IR) spectroscopy. 1 H-NMR spectra of compounds 1 and 2 were recorded in CDCl3 and are shown in Figures  The UV-vis absorption spectra of compound 1 and its ionic complexes from various acids are shown in Figure 1. A sharp Soret band (419 nm) and four Q-bands (515, 551, 590, and 645 nm) were observed. After treatment with various acids to form the ionic complexes, the Soret band red-shifted (419→422 nm for HNO3, CF3COOH, and HCl; 419→442 The UV-vis absorption spectra of compound 1 and its ionic complexes from various acids are shown in Figure 1. A sharp Soret band (419 nm) and four Q-bands (515, 551, 590, and 645 nm) were observed. After treatment with various acids to form the ionic complexes, the Soret band red-shifted (419→422 nm for HNO 3 , CF 3 COOH, and HCl; 419→442 nm for H 2 SO 4 ), and an additional band at 443 nm in the case of HCl and CF 3 COOH was observed. In the case of H 2 CO 3 , the Soret band did not shift, but it was slightly broad compared with that of compound 1. Conversely, Q-bands broadened with increasing intensity owing to the formation of ionic complexes. The Soret band of the ionic complexes of compound 1 and H 3 PO 4 showed a broad shift (half-width changed from 396 to 443 nm), and the Q bands red-shifted (515→524, 551→560, 590→596, and 645→654 nm). Similar trends were observed in the UV-vis absorption spectra of compound 2 and its ionic complexes with different acids (Figure 2). Compound 2 exhibited a typical sharp Soret band at 427 nm and Q-bands at 520, 560, and 601 nm. Blue shifts were observed in the Soret bands of all the ionic complexes of the acids: 427→420 nm for HCl, CF 3 COOH, HNO 3 , and H 2 SO 4 ; 427→416 nm for H 2 CO 3 ; 427→418 nm for H 3 PO 4 . The corresponding Q-bands also blue-shifted (551 and 592 nm). However, the Q bands also blue-shifted compared to those of compound 2 (566 and 609 nm).
Catalysts 2022, 12, x FOR PEER REVIEW 4 of 15 nm for H2SO4), and an additional band at 443 nm in the case of HCl and CF3COOH was observed. In the case of H2CO3, the Soret band did not shift, but it was slightly broad compared with that of compound 1. Conversely, Q-bands broadened with increasing intensity owing to the formation of ionic complexes. The Soret band of the ionic complexes of compound 1 and H3PO4 showed a broad shift (half-width changed from 396 to 443 nm), and the Q bands red-shifted (515→524, 551→560, 590→596, and 645→654 nm). Similar trends were observed in the UV-vis absorption spectra of compound 2 and its ionic complexes with different acids (Figure 2). Compound 2 exhibited a typical sharp Soret band at 427 nm and Q-bands at 520, 560, and 601 nm. Blue shifts were observed in the Soret bands of all the ionic complexes of the acids: 427→420 nm for HCl, CF3COOH, HNO3, and H2SO4; 427→416 nm for H2CO3; 427→418 nm for H3PO4. The corresponding Q-bands also blue-shifted (551 and 592 nm). However, the Q bands also blue-shifted compared to those of compound 2 (566 and 609 nm).   Catalysts 2022, 12, x FOR PEER REVIEW 4 of 15 nm for H2SO4), and an additional band at 443 nm in the case of HCl and CF3COOH was observed. In the case of H2CO3, the Soret band did not shift, but it was slightly broad compared with that of compound 1. Conversely, Q-bands broadened with increasing intensity owing to the formation of ionic complexes. The Soret band of the ionic complexes of compound 1 and H3PO4 showed a broad shift (half-width changed from 396 to 443 nm), and the Q bands red-shifted (515→524, 551→560, 590→596, and 645→654 nm). Similar trends were observed in the UV-vis absorption spectra of compound 2 and its ionic complexes with different acids (Figure 2). Compound 2 exhibited a typical sharp Soret band at 427 nm and Q-bands at 520, 560, and 601 nm. Blue shifts were observed in the Soret bands of all the ionic complexes of the acids: 427→420 nm for HCl, CF3COOH, HNO3, and H2SO4; 427→416 nm for H2CO3; 427→418 nm for H3PO4. The corresponding Q-bands also blue-shifted (551 and 592 nm). However, the Q bands also blue-shifted compared to those of compound 2 (566 and 609 nm).   The fluorescence spectra of compound 1 and its ionic complexes with various acids are shown in Figure 3. Compound 1 has a typical two-band emission spectrum (654 and 705 nm). After treatment with various acids to form ionic complexes, the intensities of these two bands decreased without shifting. The fluorescence spectra of compound 2 and its ionic complexes are shown in Figure 4. Compound 2 has two emission bands (601 and 652 nm). Red shifts were observed for HCl and CF 3 COOH ionic complexes (601→606 nm and 651→655 nm, respectively), while the emission intensity decreased considerably. In the case of H 2 SO 4 , HNO 3 , H 2 CO 3 , and H 3 PO 4 ionic complexes, these two bands did not shift, but their intensity decreased compared to that of compound 2. The formation of the ionic complexes was further examined via FT-IR spectroscopy, as shown in Figures S5 and S6. Characteristic peaks were observed for NO 3 − (1351 cm −1 ), SO 4 2− (1110 cm −1 ), CF 3 COO − (1671 cm −1 ), CO 3 2− (1410 cm −1 ), and PO 4 3− (1010 cm −1 ). The fluorescence spectra of compound 1 and its ionic complexes with various acids are shown in Figure 3. Compound 1 has a typical two-band emission spectrum (654 and 705 nm). After treatment with various acids to form ionic complexes, the intensities of these two bands decreased without shifting. The fluorescence spectra of compound 2 and its ionic complexes are shown in Figure 4. Compound 2 has two emission bands (601 and 652 nm). Red shifts were observed for HCl and CF3COOH ionic complexes (601→606 nm and 651→655 nm, respectively), while the emission intensity decreased considerably. In the case of H2SO4, HNO3, H2CO3, and H3PO4 ionic complexes, these two bands did not shift, but their intensity decreased compared to that of compound 2. The formation of the ionic complexes was further examined via FT-IR spectroscopy, as shown in Figures     The fluorescence spectra of compound 1 and its ionic complexes with various acids are shown in Figure 3. Compound 1 has a typical two-band emission spectrum (654 and 705 nm). After treatment with various acids to form ionic complexes, the intensities of these two bands decreased without shifting. The fluorescence spectra of compound 2 and its ionic complexes are shown in Figure 4. Compound 2 has two emission bands (601 and 652 nm). Red shifts were observed for HCl and CF3COOH ionic complexes (601→606 nm and 651→655 nm, respectively), while the emission intensity decreased considerably. In the case of H2SO4, HNO3, H2CO3, and H3PO4 ionic complexes, these two bands did not shift, but their intensity decreased compared to that of compound 2. The formation of the ionic complexes was further examined via FT-IR spectroscopy, as shown in Figures

Supramolecular Ionic Self-Assembly to Nanostructures
The self-assembly of compounds 1 and 2 and their ionic complexes were analyzed via field-emission scanning electron microscopy (FE-SEM). For sample preparation, each complex was suspended in water at a fixed concentration (c = 1 mM) and centrifuged at 13,500 rpm for 10 s. Thereafter, the solution was drop-casted onto the surface of a copper tape for deposition and dried in air, followed by Pt coating prior to the FE-SEM studies. The morphologies of the self-assembled nanostructures are shown in Figures 5 and 6. The average length of the spherical particles of compound 1 was from 110 to 200 nm ( Figure 5a). In contrast, the average particle size of the truncated polyhedron of compound 2 was from 200 to 900 nm with a width of 100 nm (Figure 5b).

Supramolecular Ionic Self-Assembly to Nanostructures
The self-assembly of compounds 1 and 2 and their ionic complexes were analy via field-emission scanning electron microscopy (FE-SEM). For sample preparation, complex was suspended in water at a fixed concentration (c = 1 mM) and centrifuge 13,500 rpm for 10 s. Thereafter, the solution was drop-casted onto the surface of a cop tape for deposition and dried in air, followed by Pt coating prior to the FE-SEM stud The morphologies of the self-assembled nanostructures are shown in Figures 5 and 6. average length of the spherical particles of compound 1 was from 110 to 200 nm (Fi 5a). In contrast, the average particle size of the truncated polyhedron of compound 2 from 200 to 900 nm with a width of 100 nm (Figure 5b). After the formation of ionic complexes of compound 1 with various acids, its m phology changed significantly. In the case of ionic complexes derived from H2CO3 average particle size was from 250 to 350 nm, with a width of 50 nm (Figure 6a). average particle size of the ionic complexes of compound 1 and HNO3 was smaller (7 nm) (Figure 6b). In contrast, the particle size of the ionic complexes of compound 1 H2SO4 was larger than that of compound 1, ranging from 250 to 600 nm (Figure 6c). Fu nanorods were observed in the ionic complexes of compound 1 and HCl, with the ave length ranging from 160 to 40 nm (Figure 6d). The morphology of the ionic complex compound 1 and CF3COOH exhibited rectangular-shaped structures (210 × 170 × 80 n with several small particles (Figure 6e). Flake-shaped particles were observed in i complexes of compound 1 and H3PO4 (Figure 6f), with average particle size ranging f 200 to 260 nm in length and a width of 50 nm. From Figure 6, porphyrin ionic compl of compound 1 and various acids self-assembled into nanostructures, and the morpho of these particles depends on the counter anions.   After the formation of ionic complexes of compound 1 with various acids, its morphology changed significantly. In the case of ionic complexes derived from H2CO3, the average particle size was from 250 to 350 nm, with a width of 50 nm (Figure 6a). The average particle size of the ionic complexes of compound 1 and HNO3 was smaller (70-20 nm) (Figure 6b). In contrast, the particle size of the ionic complexes of compound 1 and H2SO4 was larger than that of compound 1, ranging from 250 to 600 nm (Figure 6c). Fused nanorods were observed in the ionic complexes of compound 1 and HCl, with the average length ranging from 160 to 40 nm (Figure 6d). The morphology of the ionic complexes of compound 1 and CF3COOH exhibited rectangular-shaped structures (210 × 170 × 80 nm 3 ) with several small particles (Figure 6e). Flake-shaped particles were observed in ionic complexes of compound 1 and H3PO4 (Figure 6f), with average particle size ranging from 200 to 260 nm in length and a width of 50 nm. From Figure 6, porphyrin ionic complexes of compound 1 and various acids self-assembled into nanostructures, and the morphology of these particles depends on the counter anions.  After the formation of ionic complexes of compound 1 with various acids, its morphology changed significantly. In the case of ionic complexes derived from H 2 CO 3 , the average particle size was from 250 to 350 nm, with a width of 50 nm (Figure 6a). The average particle size of the ionic complexes of compound 1 and HNO 3 was smaller (70-20 nm) (Figure 6b). In contrast, the particle size of the ionic complexes of compound 1 and H 2 SO 4 was larger than that of compound 1, ranging from 250 to 600 nm (Figure 6c). Fused nanorods were observed in the ionic complexes of compound 1 and HCl, with the average length ranging from 160 to 40 nm (Figure 6d). The morphology of the ionic complexes of compound 1 and CF 3 COOH exhibited rectangular-shaped structures (210 × 170 × 80 nm 3 ) with several small particles (Figure 6e). Flake-shaped particles were observed in ionic complexes of compound 1 and H 3 PO 4 (Figure 6f), with average particle size ranging from 200 to 260 nm in length and a width of 50 nm. From Figure 6, porphyrin ionic complexes of compound 1 and various acids self-assembled into nanostructures, and the morphology of these particles depends on the counter anions. The morphologies of the self-assembled nanostructures obtained upon the formation of ionic complexes derived from compound 2 with various acids are shown in Figure 7. In the case of ionic complexes of H 2 CO 3 , the average particle size of the sphere ranged from 70 to 140 nm (Figure 7a). Fused nanoflakes were observed in the ionic complex of compound 2 with HNO 3 (Figure 7b). A square-shaped morphology was observed in the complex of compound 2 with H 2 SO 4 , with the average length of the fused square particles being 290 nm (Figure 7c). Fused rectangular-shaped structures (240 × 160 × 55 nm 3 ) were observed for the ionic complexes of compound 2 and HCl (Figure 7d). The morphology of the ionic complexes of compound 2 and CF 3 COOH was considerably different from that of compound 2 (Figure 7e) and demonstrated rectangular-shaped structures (220 × 90 × 40 nm 3 ) along with fused spheres (average diameter of~120 nm). The morphologies of the ionic complexes of compound 2 and H 3 PO 4 consisted of stone-like particles with lengths ranging from 90 to 300 nm and a width of 20 nm (Figure 7f). From Figure 7, it can be concluded that the counter anions affect the morphology of the ionic complexes derived from Sn(IV) porphyrin 2 with different acids. The morphologies of the self-assembled nanostructures obtained upon the formation of ionic complexes derived from compound 2 with various acids are shown in Figure 7. In the case of ionic complexes of H2CO3, the average particle size of the sphere ranged from 70 to 140 nm (Figure 7a). Fused nanoflakes were observed in the ionic complex of compound 2 with HNO3 (Figure 7b). A square-shaped morphology was observed in the complex of compound 2 with H2SO4, with the average length of the fused square particles being 290 nm (Figure 7c). Fused rectangular-shaped structures (240 × 160 × 55 nm 3 ) were observed for the ionic complexes of compound 2 and HCl (Figure 7d). The morphology of the ionic complexes of compound 2 and CF3COOH was considerably different from that of compound 2 (Figure 7e) and demonstrated rectangular-shaped structures (220 × 90 × 40 nm 3 ) along with fused spheres (average diameter of ~120 nm). The morphologies of the ionic complexes of compound 2 and H3PO4 consisted of stone-like particles with lengths ranging from 90 to 300 nm and a width of 20 nm (Figure 7f). From Figure 7, it can be concluded that the counter anions affect the morphology of the ionic complexes derived from Sn(IV) porphyrin 2 with different acids.

Photocatalytic Performance for the Degradation of MG Dye
We have examined the photocatalytic activity of the photocatalysts developed in this study for the degradation of MG dye in aqueous solutions under visible-light irradiation. A negligible decay of the MG dye was observed in the absence of either a photocatalyst or visible light ( Figure S7). ESI-MS spectra were recorded after 30 min in each photodegradation experiment. New small molecules from the degradation of the MG dye were formed, which was confirmed by the appearance of new peaks in the mass spectra. The time-dependent absorption spectra of MG in the presence of photocatalysts generated from compound 2 with H3PO4 under visible light irradiation are shown in Figure S8. From Figures 8 and 9, all the photocatalysts (generated from compounds 1 and 2 with various acids) were effective for the photodegradation of MG, and their photocatalytic performance depended on the morphology of the nanoaggregates. The degradation rates of MG dye in the presence of a photocatalyst can be determined from its degradation efficiency, (C0 − C)/C0, where C0 is the initial concentration of MG, and C is the concentration at time t. The observed degradation rates of MG were 50% for compound 1 and 53, 60, 66, 73, 86, and 90% for the photocatalysts derived from compound 1 with H2CO3, HNO3, H2SO4, HCl, CF3COOH, and H3PO4, respectively (Figure 8).

Photocatalytic Performance for the Degradation of MG Dye
We have examined the photocatalytic activity of the photocatalysts developed in this study for the degradation of MG dye in aqueous solutions under visible-light irradiation. A negligible decay of the MG dye was observed in the absence of either a photocatalyst or visible light ( Figure S7). ESI-MS spectra were recorded after 30 min in each photodegradation experiment. New small molecules from the degradation of the MG dye were formed, which was confirmed by the appearance of new peaks in the mass spectra. The time-dependent absorption spectra of MG in the presence of photocatalysts generated from compound 2 with H 3 PO 4 under visible light irradiation are shown in Figure S8. From Figures 8 and 9, all the photocatalysts (generated from compounds 1 and 2 with various acids) were effective for the photodegradation of MG, and their photocatalytic performance depended on the morphology of the nanoaggregates. The degradation rates of MG dye in the presence of a photocatalyst can be determined from its degradation efficiency, (C 0 − C)/C 0 , where C 0 is the initial concentration of MG, and C is the concentration at time t. The observed degradation rates of MG were 50% for compound 1 and 53, 60, 66, 73, 86, and 90% for the photocatalysts derived from compound 1 with H 2 CO 3 , HNO 3 , H 2 SO 4 , HCl, CF 3 COOH, and H 3 PO 4 , respectively (Figure 8). In contrast, the observed degradation rate of MG was 65% for comp 70, 75, 80, 91, and 95% for the photocatalysts derived from compound H2SO4, HCl, H2CO3, CF3COOH, and H3PO4, respectively ( Figure 9). From t vation, it was clear that the photocatalysts derived from 1 and 2 with H3P optimum performance towards the decay of MG dye. To further elucida kinetics for the degradation of the MG dye, the pseudo-first-order rate equ for the photocatalytic degradation experiments. The equation was express kt, where C0 is the initial concentration of dye, C represents dye concentra tion time t, and k denotes the pseudo-first-order rate constant. Using the in Figures 8 and 9, the reaction kinetics of the MG dye degradation are p ures S9 and S10. The first-order rate constant for the degradation of MG b was 0.009 min −1 , and the other rate constants were 0.011, 0.013, 0.015, 0 0.032 min −1 for the photocatalysts derived from compound 1 with H2CO3 HCl, CF3COOH, and H3PO4, respectively ( Figure S9).  In contrast, the observed degradation rate of MG was 65% for comp 70, 75, 80, 91, and 95% for the photocatalysts derived from compound H2SO4, HCl, H2CO3, CF3COOH, and H3PO4, respectively (Figure 9). From t vation, it was clear that the photocatalysts derived from 1 and 2 with H3P optimum performance towards the decay of MG dye. To further elucid kinetics for the degradation of the MG dye, the pseudo-first-order rate equ for the photocatalytic degradation experiments. The equation was express kt, where C0 is the initial concentration of dye, C represents dye concentr tion time t, and k denotes the pseudo-first-order rate constant. Using the in Figures 8 and 9, the reaction kinetics of the MG dye degradation are p ures S9 and S10. The first-order rate constant for the degradation of MG b was 0.009 min −1 , and the other rate constants were 0.011, 0.013, 0.015, 0 0.032 min −1 for the photocatalysts derived from compound 1 with H2CO3 HCl, CF3COOH, and H3PO4, respectively ( Figure S9).  In contrast, the observed degradation rate of MG was 65% for compound 2 and 67, 70, 75, 80, 91, and 95% for the photocatalysts derived from compound 2 with HNO 3 , H 2 SO 4 , HCl, H 2 CO 3 , CF 3 COOH, and H 3 PO 4 , respectively (Figure 9). From the above observation, it was clear that the photocatalysts derived from 1 and 2 with H 3 PO 4 showed the optimum performance towards the decay of MG dye. To further elucidate the reaction kinetics for the degradation of the MG dye, the pseudo-first-order rate equation was used for the photocatalytic degradation experiments. The equation was expressed as ln(C 0 /C) = kt, where C 0 is the initial concentration of dye, C represents dye concentration at irradiation time t, and k denotes the pseudo-first-order rate constant. Using the data displayed in Figures 8 and 9, the reaction kinetics of the MG dye degradation are presented in Figures S9 and S10. The first-order rate constant for the degradation of MG by compound 1 was 0.009 min −1 , and the other rate constants were 0.011, 0.013, 0.015, 0.017, 0.028, and 0.032 min −1 for the photocatalysts derived from compound 1 with H 2 CO 3 , HNO 3 , H 2 SO 4 , HCl, CF 3 COOH, and H 3 PO 4 , respectively ( Figure S9).
Similarly, the first-order rate constant for the degradation of MG by compound 2 was 0.014 min −1 and the other rate constants were 0.015, 0.016, 0.018, 0.022, 0.032, and 0.039 min −1 for the photocatalysts derived from compound 2 with HNO 3 , H 2 SO 4 , HCl, H 2 CO 3 , CF 3 COOH, and H 3 PO 4 , respectively ( Figure S10). Previously, various investigations concerning the MG dye degradation have been reported in the literature. Our observed rate constants were comparable to or even better than those reported for other photocatalysts such as TiO 2 (0.0064 to 0.140 min −1 ) [47], Zn(II)porphyrin with H 2 O 2 (0.0042 to 0.0054 min −1 ) [48], Au-Ag/TiO 2 (0.00987 to 0.2595 min −1 ) [49], and CuFe 2 O 4 (0.0142 min −1 ) [50]. In addition, the recovery of the photocatalysts from the reaction vessel was easily handled by a continuous filtration-wash-dry procedure. The efficiency of the photocatalyst remained unchanged after recovery, and it showed excellent recyclability for the degradation of MG dye ( Figure S11). It should be noted that the efficiency of photocatalyst (derived from 2 with H 3 PO 4 ) remained almost intact even after 10 consecutive cycles. As depicted in Figure S11, 95% MG was degraded within 70 min for the first run, and the photocatalytic performance slightly decreased, marked by 91% MG degradation, after 10 consecutive cycles. The morphology of the catalysts remained intact after their use in the degradation experiment, as confirmed by the FE-SEM image of the photocatalyst derived from compound 2 with H 3 PO 4 ( Figure S12).
To examine the effect of dye/catalyst ratio on the performance of the photocatalysts, different concentrations of the MG dye aqueous solution (5,10,15,20,30, and 40 mg L −1 ) with a fixed amount of catalyst (5 mg each time) were used for each time. From Figure S13, it was revealed that the degradation rate gradually decreased as the concentration of MG dye increased. The effect of the temperature and the pH on the degradation of MG dye has been also examined. From the Figure S14, it was observed that degradation efficiency increased with increasing temperature. As shown in Figure S15, the pH of the MG dye solution affected the degradation rate. The rate of degradation increased from pH 2 to pH 7, and then decreased up to pH 11. Several possible mechanisms for the photodegradation of pollutant dyes in aqueous solutions using porphyrin-based photocatalysts have been reported [11,[15][16][17]. The mechanism consists of five steps for porphyrin-based nanomaterials (P). When the photocatalyst absorbs visible-light irradiation, electrons in the valence band (VB) are excited to the conduction band (CB). The electron-hole pairs (e − /h + ) are formed through this process at the surface of the photocatalyst (step 1). Intermolecular π-π interactions stabilized the electronic delocalization over the nanomaterials and minimize the recombination energy of the excited electrons. Next, the photogenerated holes (h + ) react with H 2 O to produce reactive species of hydroxyl radicals ( • OH). On the other hand, the reaction of the electrons with the dissolved O 2 generate superoxide radical anions (O 2 −• ) (steps 2 and 3). Then these reactive species (superoxide radical anions and hydroxyl radicals) react with the MG dye and mineralized it into low molecular compounds (steps 4 and 5).
• OH + MG→Degraded products (4) Reactive species involved in the photocatalytic degradation of the MG dye by the photocatalyst generated from compound 2 and H 3 PO 4 were identified via radical trapping analysis [51]. To investigate the effect of different reactive species on the photodegradation of the MG dye, tert-butyl alcohol ( t BuOH) and para-benzoquinone (BQ) were used to capture • OH and O 2 −• , respectively ( Figure S16). From Figure S16, the degradation rate was affected by the presence of t BuOH as well as BQ, suggesting that hydroxyl radicals and superoxide radicals are the major reactive species that contribute to the catalytic oxidation of the MG dye. For further details on the detection of all intermediate species involved in the dye degradation process (all oxygen species, such as 1 O 2 and 3 O 2 , • OH, or O 2 −• ), readers can refer to our previous reports [11,[15][16][17].
In order to examine the adsorption properties of the photocatalysts towards the MG dye, the Brunauer-Emmett-Teller (BET) surface area of the photocatalysts (derived from 1 and 2 with H 3 PO 4 ) was measured. Figure S17 presents the N 2 adsorption-desorption isotherms of the photocatalysts (derived from 1 and 2 with H 3 PO 4 ). Each photocatalyst exhibited a low surface area, for example the BET surface area of the best performing photocatalysts (derived from 1 and 2 with H 3 PO 4 ) were estimated to be 4.5 m 2 /g and 12.7 m 2 /g, respectively. As shown in Figure S18, it took about 30 min to reach the absorption-desorption equilibrium, 2% and 5% of MG dye were adsorbed by the photocatalysts derived from 1 and 2 with H 3 PO 4 , respectively. Lower adsorption of dye on the surface of photocatalysts have some advantages as the bare active sites are readily available for photodegradation. However, the absorption of dye on the surface of photocatalyst not only depends on the surface area of the catalyst but also the interaction of dye with active sites of the photocatalysts and various interaction at the junction of solid-liquid interface.
The reaction mixture was analyzed after 30 min irradiation of visible light using ESI-MS ( Figure S19) to study the details of the degradation pathway of the MG dye. New peaks appeared in the mass spectra imply that the MG dye was degraded into new small molecules [52,53]. By analyzing mass spectral data in Figure S19, possible intermediates for the degradation of the MG dye are given in Figure 10. Hydroquinone is further fragmented into benzoquinone and then to maleic acid. Finally, all intermediate molecules were further fragmented and mineralized into CO 2 and H 2 O. Moreover, the total organic carbon (TOC) value was measured to evaluate the removal of the MG dye by photocatalysts [54]. The TOC removal percentage achieved using the photocatalyst generated from compound 2 and H 3 PO 4 was only 75%.

Materials and Methods
All chemicals were purchased from Tokyo Chemical Industry Co. LTD. (Tokyo, Japan) and used without further purification. Elemental analysis was conducted using an EA 1110 Fisons analyzer (Used Lab Machines Limited, London, UK). 1 H-NMR spectra were recorded on a Bruker BIOSPIN/AVANCE III 400 spectrometer at 293 K (Bruker Bio-Spin GmbH, Silberstreifen, Rheinstetten, Germany). ESI-MS spectra were recorded using a Thermo Finnigan linear ion trap quadrupole mass spectrometer (Thermo Fisher Scientific, Waltham, MA, USA). Steady-state UV-vis and fluorescence spectra were recorded on a Shimadzu UV-3600 spectrophotometer and a Shimadzu RF-5301PC fluorescence spectrophotometer, respectively. FT-IR spectra were measured using a Shimadzu FTIR-8400S spectrophotometer (Shimadzu, Tokyo, Japan). FE-SEM images were obtained using a MAIA III system (TESCAN, Brno, Czech Republic). The Brunauer-Emmett-Teller (BET) surface area using N2 adsorption isotherms at 77 K was measured with an analyzer (BEL-SORP-mini volumetric adsorption equipment). The total organic carbon (TOC) values of the degradation of MG dye were examined with a JEOL JEM-3010 total organic carbon analyzer to calculate the extent of mineralization.

Materials and Methods
All chemicals were purchased from Tokyo Chemical Industry Co. LTD. (Tokyo, Japan) and used without further purification. Elemental analysis was conducted using an EA 1110 Fisons analyzer (Used Lab Machines Limited, London, UK). 1 H-NMR spectra were recorded on a Bruker BIOSPIN/AVANCE III 400 spectrometer at 293 K (Bruker BioSpin GmbH, Silberstreifen, Rheinstetten, Germany). ESI-MS spectra were recorded using a Thermo Finnigan linear ion trap quadrupole mass spectrometer (Thermo Fisher Scientific, Waltham, MA, USA). Steady-state UV-vis and fluorescence spectra were recorded on a Shimadzu UV-3600 spectrophotometer and a Shimadzu RF-5301PC fluorescence spectrophotometer, respectively. FT-IR spectra were measured using a Shimadzu FTIR-8400S spectrophotometer (Shimadzu, Tokyo, Japan). FE-SEM images were obtained using a MAIA III system (TESCAN, Brno, Czech Republic). The Brunauer-Emmett-Teller (BET) surface area using N 2 adsorption isotherms at 77 K was measured with an analyzer (BELSORP-mini volumetric adsorption equipment). The total organic carbon (TOC) values of the degradation of MG dye were examined with a JEOL JEM-3010 total organic carbon analyzer to calculate the extent of mineralization.  (1) Freshly distilled pyrrole (0.38 mL, 5.5 mmol) was added dropwise to a solution of 4-(2pyridyl)benzaldehyde (1.00 g, 5.5 mmol) in propionic acid (200 mL) under reflux. After 3 h, the solvent was evaporated to dryness under vacuum to obtain an oily residue. Thereafter, the oily residue was neutralized with trimethylamine and dissolved in a minimum amount of N,N-dimethylformamide. The reaction mixture was stored in a refrigerator overnight. The resulting solid was filtered, washed with hot water, and dried. The crude product was purified using column chromatography (SiO 2 , eluent: CH 2 Cl 2 /MeOH = 90:10) to afford 5,10,15,20-tetrakis [4-(2-pyridyl)phenyl]porphyrin 1. The crude product was recrystallized from CH 2 Cl 2 /n-hexane to obtain a violet-purple crystalline powder. Yield: 0.127 g (10%).  (2) H 2 TPhPyP (0.092 g, 0.1 mmol) was dissolved in pyridine (50 mL), and SnCl 2 ·2H 2 O (0.22 g, 0.85 mmol) was added to the solution and refluxed for 15 h. Subsequently, pyridine was removed under vacuum, and the residue was dissolved in CHCl 3 and filtered through a celite pad to remove excess SnCl 2 . The filtrate was evaporated, and the crude was redissolved in a solvent mixture (50 mL) consisting of tetrahydrofuran (THF; 40 mL) and H 2 O (10 mL). K 2 CO 3 (0.33 g, 1.45 mmol) was added to the reaction mixture, and it was refluxed for 10 h. THF was removed under reduced pressure conditions, and the mixture was cooled to 5 • C to precipitate the product. The solid compound was then filtered and dried in an oven. Recrystallization (CHCl 3 /CH 3 (1) and Sn(OH) 2 TPhPyP (2) Using Various Acids

Synthesis of the Ionic Complexes of H 2 TPhPyP
Ten milligrams of each of the porphyrin compounds (1 and 2) were dissolved in a 1% aqueous solution (5 mL) of each of the following acids-HCl, HNO 3 , CF 3 COOH, H 2 SO 4 , H 2 CO 3 , and H 3 PO 4 , and acetone (20 mL) was layered over for slow diffusion for 7 d. Subsequently, the crystalline materials were filtered, washed with water, and air-dried.

Photocatalytic Investigation
The photocatalytic efficiency of the self-assembled ionic nanostructures was studied by monitoring the degradation of MG dye in an aqueous solution by the catalysts, using malachite oxalate green salt as the source of MG. In a typical procedure, 5 mg of the photocatalyst was added to 150 mL of aqueous MG solution (20 mg L −1 , distilled water at pH 7.0) with stirring. The mixture was then placed in the dark for 30 min to attain an adsorption-desorption equilibrium. The irradiation process was conducted using a 150 W xenon arc lamp with a UV cut-off filter (ABET technologies, Old gate lane Milford, Milford, CT, USA) at room temperature. At specified intervals, 3 mL of the suspension was centrifuged to remove the photocatalyst. Subsequently, the solution was analyzed by recording absorption variations in the UV-Vis spectra at 617 nm to determine the concentration of MG.

Conclusions
In this study, various porphyrin-based ionic complexes were prepared through the reaction of two porphyrin precursors, compounds 1 and 2, with various acids (HCl, HNO 3 , CF 3 COOH, H 2 SO 4 , H 2 CO 3 , and H 3 PO 4 ). Depending on the counter anion [55,56], these complexes readily assembled into different nanostructures, based on the electronic interaction of cations of compounds 1 or 2 with different counter anions. These nanostructures can be further stabilized through π-π stacking interactions between the phenyl-extended porphyrin rings. Furthermore, the morphology-dependent photocatalytic degradation of MG dye in aqueous solution under visible light irradiation for 70 min was observed for all the ionic complexes, with efficiencies ranging from 50% to 95%. Therefore, counter anions can affect the morphology as well as the optical properties of porphyrin-based species. In addition, the aggregates are good photocatalysts for the treatment of wastewater.
Supplementary Materials: The following supporting information can be downloaded at https: //www.mdpi.com/article/10.3390/catal12070799/s1, 1 H-NMR spectra, mass spectra, and results of the photocatalytic degradation of MG are shown in Figures S1-S19. Figure S1: 1 H-NMR spectrum of compound 1; Figure S2: 1 H-NMR spectrum of compound 2; Figure S3: ESI-MS spectrum of compound 1; Figure S4: ESI-MS spectrum of compound 2; Figure S5: FT-IR spectra of compound 1 and its ionic complexes; Figure S6: FT-IR spectra of compound 2 and its ionic complexes; Figure S7: UV-vis spectroscopic observation of the photocatalytic degradation of MG dye in aqueous solution under visible light irradiation by photocatalyst 1 and 2; Figure S8: Absorption spectral change of MG dye for the photocatalytic degradation by photocatalyst generated from 2 and H 3 PO 4 in aqueous solution under visible light irradiation; Figure S9: Kinetics for the photocatalytic degradation of MG under visible light irradiation by photocatalyst generated from 1 and various acids; Figure S10: Kinetics for the photocatalytic degradation of MG under visible light irradiation by photocatalyst generated from 2 and various acids; Figure S11: Recyclability of the photocatalyst generated from 2 and H 3 PO 4 towards the degradation of MG dye; Figure S12: FE-SEM images of photocatalyst generated from 2 and H 3 PO 4 (after and before the degradation of MG dye); Figure S13: Effect of initial concentration of MG on the degradation by the photocatalyst (5 mg) generated from 2 and H 3 PO 4 ; Figure S14: Effect of temperature on the MG degradation by the photocatalyst generated from 2 and H 3 PO 4 ; Figure S15: Effect of pH of the solution on the MG degradation by the photocatalyst generated from 2 and H 3 PO 4 ; Figure S16: Photocatalytic degradation of MG dye in aqueous solution by the photocatalyst generated from 2 and H 3 PO 4 with the addition of different scavengers under visible light irradiation; Figure S17: N 2 adsorption-desorption isotherms at 77 K for the photocatalysts derived from 1 and 2 with H 3 PO 4 ; Figure S18: Adsorption of MG dye for the photocatalysts derived from 1 and 2 with H 3 PO 4 ; Figure S19: ESI-MS spectrum (positive ion mode) of the reaction mixture of MG with the photocatalyst generated from 2 and H 3 PO 4 after 30 min of visible light irradiation.
Author Contributions: Investigation, methodology, data curation, visualization, formal analysis, validation, software, and writing, N.K.S.; conceptualization, review, editing, supervision, project administration, and funding acquisition, H.-J.K. All authors read and agreed to the published version of the manuscript.