Performance of new metal complexes for anionic and cationic dyes photodegradation: construction, spectroscopic studies, optical properties, and DFT calculations

Abstract

Water pollution is referred to as the introduction of exotic species into an ecosystem, which poses serious difficulties for both the environment and human health. As a result, water treatment has garnered a lot of interest in recent studies, particularly for dyes removal. In an effort to synthesize new photocatalysts with the highest potential for removing pollutants, photocatalytic degradation of two types of dyes; aniline blue, and malachite green by newly synthesized Ag(I), Ni(II), and Co(II) complexes photocatalysts has been investigated here. Ag(I), Ni(II), and Co(II) ions were used to interact with the novel imine ligand, HL, which was produced by condensing 3-acetylcoumarin, and sulfadimidine, in order to make the desired complexes. Different spectroscopic methods were used to properly characterize and confirm the structures of all synthesized compounds. The ligand, HL act as anionic tetradentate through the imine group (C=N), lactone carbonyl group (C=O), sulfonamide group (SO₂NH), and pyrimidine ring nitrogen. The suggested formula of the synthesized complexes were [Ag₂(L)(NO₃)(H₂O)₃], [Co₂(L)Cl₃(H₂O)].2H₂O, and [Ni₂(L)Cl₃(H₂O)].2H₂O with a tetrahedral structure, and nonelectrolyte nature. Through optical studies and XRD measurements, all investigated compounds have the semiconductor nature and nano sizes. TGA-DTG/DSC techniques were utilized to thermal behavior study as well as assess some thermodynamic and kinetic parameters for all degradation stages. The geometry of the proposed structure of all new synthesized compounds was optimized by density functional theory (DFT) calculations. The BET surface area and the porosity data align well with the excellent photocatalytic performance for [Co₂(L)Cl₃(H₂O)].2H₂O complex toward the dyes degradation.

Introduction

Although water is the source of life, and required for different industry fields, industrial development in addition to the human activities in the different fields causes destruction to water supplies. Also, the significant properties of water like the high dissolving capacity in addition to the ease of forming a suspension facilitate the transport of pollutants into water streams, and adversely affect water supplies [1, 2]. Currently, water contamination due to the chemicals either caused by organic or inorganic species is one of the major challenges that affects the whole world. Generally, the various chemicals produced from the different industries are the major contributors to the environmental pollution, and specifically to the aquatic pollution causing serious health issues. Many synthetic dyes are now required in the textile, paint, paper, and leather industries. Industrial effluents from such industries may end up in the water sources causing an observable dyes pollutant that result in the unpleasant odor of the water, and the reduction of the oxygen content, transparence, and sunlight penetration into the water affecting the living organisms, photosynthesis process, and promoting the bioaccumulation of unfavorable species. This is because the dyes are chemical compounds characterize with an attractive color, complex organic structure, highly poisonous, highly stability, low biodegradability, light stability, heat resistant, and can contaminate with water sources for longer times [3, 4]. To reduce or eliminate the impact of pollutants on water supplies, dye-containing wastewater needs to be treated to get rid of any remaining dyes before being released into an environment.

Different strategies have been employed to address this issue, and each one has strengths and weaknesses of its own. Some of these techniques have great problems of high energy, and time consumption, various organic solvents, and complex reagents usage, sludges formation, multistep operation, inefficiency and economical non-satisfactory [5]. Most often, advanced oxidation processes (AOPs) are used, and it is based on the production of extremely reactive free radical species that are highly reactive, and potent oxidizing agents that react with a wide variety of pollutants, and completely mineralize them into carbon dioxide and water. Heterogeneous photocatalysis is one of these methods, in which a semiconductor functioning as the photocatalyst uses radiation to produce enough free radical species with the capacity to pollute to mineralize less harmful compounds [6].

Taking into account all of these factors, our team set out to synthesis a new ligand with multiple donor sites that can effectively chelate with metal ions Ag(I), Ni(II), and Co(II) to form metal complexes that can be utilized as photocatalysts for degradation of two model dyes, aniline blue (an anionic dye), and malachite green (a cationic dye), which serve as models for organic pollutants. Additionally, structure, composition, morphology, optical characteristics, and theoretical calculations for ligand, HL as well as its Ag(I), Ni(II), and Co(II) complexes were considered.

Experimental part

Materials

Sulfadimidine (Sigma); 3-acetylcoumarin (Sigma-Aldrich); silver (I) nitrate (AgNO₃) in addition to hexahydrate chloride slates of cobalt (II), and nickel (II) ions (Merck, Fluka); Aniline blue dye (Sigma-Aldrich), and malachite green dye (Merck). Table 1 shows some of the major properties of the dyes under study.

Table 1 Main characteristics of organic dyes studied; aniline blue, and malachite green

Instrumentation

Quantification of carbon, sulphur, nitrogen, and hydrogen was inquired by an elemental analyzer Perkin-Elmer CHN 2400 type. An electronic melting point device created in Britain by Griffin, and George was employed for melting points measurements. Experimental infrared spectra were carried out on a Perkin-Elmer 437 IR spectrometer that was run as KBr discs covering 4000–400 cm⁻¹ range. Proton nuclear magnetic resonance measurements were carried out using Varian-Mercury spectrometer (300 MHz) in DMSO-d₆ solvent. With a Finnigan MAT SSQ7000 spectrometer operating at 70 eV with an electron ionization technique, mass spectra data was captured. Spectrophotometer Perkin Elmer Lambda 35 with 800–200 nm range was employed for electronic spectra measurements. Using Shimadzu DT-50 thermal analyzer, and Mettler-Toledo DSC 822e module, the thermal study (TG-DTG, DSC) data were collected from 25 to 950 °C with regulated heating rate 10 °C per min in nitrogen environment, and 20 mL/min flow rate. The Jenway 4010 conductometer was employed to measure molar conductance for the created complexes in DMSO, whereas Sherwood magnetic susceptibility balance underwent magnetic measurements. X-ray diffraction data were collected by Philips diffractometer with anode material of Cu. Scanning electron microscopy (SEM) micrographs were captured at a voltage of 10 kV using Quanta FEG-250 microscope. An N₂ adsorption–desorption isotherm analysis was conducted using BELSORP MINI X instrument at 77 K, and a relative pressure (P/P₀) range of 0.00–1.00.

Synthesis of new ligand, HL

Following the described procedures for the synthesis of imine compounds, the ligand, HL, was synthesized using an equimolar ratio of 3-acetylcoumarine, and sulfadimidine as raw reactants. At room temperature, an ethanolic solution of 3-acetylcoumarine (0.01 M) was introduced dropwise over a period to a well-stirred ethanolic solution of sulfadimidine (0.01 M). Drops of glacial acetic acid were added in sequence followed by refluxing for 6 h while being constantly stirred. Throughout this time a complete dissolution, and a change of solution color were seen. The resulting yellowish-red solution was then concentrated on a hot plate and left to cool at ambient conditions. The formed yellow product was separated, followed by washing with ethanol then diethyl ether numerous times, and recrystallized in ethanol. Figure 1 depicts the ligand, HL synthetic route.

Fig. 1

The synthetic route of the ligand, HL, and its metal complexes

Photocatalysts synthesis

To the synthesized ligand, HL dissolved in ethanol under stirring at room temperature, an equimolar ratio of an aqueous solution from metal salts of Ag(I), Ni(II), and Co(II) ions was added dropwise. Following complete addition, the solution was heated gradually to reflux for 4 h while being constantly stirred, at which point the complexes precipitated. The formed precipitated was separated, washed numerous times with ethanol then diethyl ether, and recrystallized by ethanol. Figure 1 depicts the metal complexes synthetic route.

To estimate the metal ions content in their complexes, the investigated complexes were broken down with mixture of strong nitric, and sulfuric acids. Usually, a precisely measured portion of the complexes (0.10 to 0.15 g) was carefully digested in a Kjeldahl flask with a 2:1 (v/v) solution of concentrated nitric, and sulphuric acids. This was done with gradual heating, followed by evaporation until near dryness, at which point the evolution of gases stopped. Following digestion, distilled water was used to dissolve the residue, which had the color of the matching metal salt and made up a volume of 100 mL. The Ni(II) and Co(II) ions content was estimated complexometry, while Ag(I) and Cl(I) ions content was estimated gravimetry [7].

Photocatalytic degradation performance towards organic pollutants

The photocatalytic activity of all synthesized complexes as catalysts was assessed via photodegradation of aniline blue (AB), and malachite green (MG) dyes at room temperature using a self-designed photoreactor. A stock aqueous solution (1000 ppm) of AB, and MG dyes were prepared using distilled water. At first, a portion of an aqueous solution of each dye (50 ppm) was scanned with UV–visible spectrophotometer at room temperature from 200 to 800 nm to obtain their maximum absorption wavelengths. After that a known weight of the catalysts (0.02 g) is transferred into a dry clean conical flask containing the working solution of 25 mL of an aqueous solution of each dye (50 ppm) followed by stirring the resulted suspension for 10 min in darkness to attain a dispersion of catalyst, and a uniform mixture. After that, the suspension was placed 15 cm away from a 300 W tungsten lamp, and the side walls of the flask were covered with an aluminum foil to avoid any effect of other light sources. The experiment was allowed to run while stirring for 80 min. At various time intervals of irradiation, samples of the suspension were withdrawn, and filtered immediately to separate suspended solid. The filtrate was analyzed for the residual AB, and MG dyes concentration by UV–Vis absorption measurements for 2 mL of sample at the characteristic peaks of the AB, and MG dyes at 595, and 617 nm, respectively [8]. Also, an experiment without a catalyst was performed under similar operating conditions.

Density functional theory study

Density functional theory (DFT) for geometry optimization and frontier molecular orbitals were calculated for the complexes by utilizing the Gaussian 09 program package. B3LYP functional was employed with a 6-311G +  + (d, p) basis set for C, H, N, O, and Cl and LANL2DZ for complexes in the gas phase [9].

Results and discussion

Within this work, 3-acetylcoumarin was reacted successfully with sulfadimidine to produce a structurally significance organic ligand, HL (Fig. 1) functionalized with multi donor sites; imine group (C=N), lactone carbonyl group (C=O), sulfonamide group (SO₂NH), and pyrimidine ring nitrogen. The presence of these donor sites in the core structure of ligand, HL facilitates its binding with the metal ions to form complexes by acting as a chelating agent. The analytical data of the created complexes reveal an interaction between metal ions and the ligand, HL occur in 2:1 molar ratio (M:L), which suggests a development of binuclear complexes. An analytical investigation, and physical appearance for the created ligand, HL, and complexes are collected in Table 2. Formularization for synthesized complexes; [Ag₂(L)(NO₃)(H₂O)₃], [Co₂(L)(H₂O)Cl₃]0.2H₂O, and [Ni₂(L)(H₂O)Cl₃]0.2H₂O as given in Fig. 1 were proposed based on correlation of all techniques used in study.

Table 2 Analytical data and some physical properties of the synthesized ligand (HL), and its metal complexes

Infrared spectra investigation

The free ligand, HL, has a complex building structure comprising numerous functional groups, and so, its IR spectrum exhibits unique peaks. Table 3 is a list of the main infrared bands. An important feature to be noted is the lack of the absorption bands characteristic of the ester carbonyl group, and the amino group of sulfadimidine fragment. Instead, the emergence of a novel strong peak at 1620 cm⁻¹ clearly shows the existence of imine linkage confirming the expected structure of the ligand, HL [10]. Further, the peaks at 1736, and 3426 cm⁻¹ are said to be coumarin carbonyl lactone, and NH moiety of sulfadimidine [10, 11], respectively. Also, the IR spectrum displays strong peaks at 1381, and 1142 cm⁻¹ in addition to 872, and 756 cm⁻¹ responsible for symmetric & asymmetric vibration frequencies of the –SO₂– group, vibration mode for S–N, and C–S species, respectively [12,13,14].

Table 3 The major important infrared spectral peaks, (cm^–1) and ¹H NMR spectra data, δ (ppm) and their tentative assignments of the synthesized ligand (HL), and its metal complexes

Upon careful inspection of free ligand, HL spectrum with their metal complexes spectra (Fig. 2), some notices that realize the involvement of coordination sites have been registered as:

• The characteristic vibrations responsible to (C=N) group was red shifted in all metal complexes to 1604–1579 cm⁻¹ with a shift of 41–16 cm⁻¹. Another significant difference is the downward shift of the characteristic vibrations caused by the oxygen atom (C=O) to 1720–1636 cm⁻¹ with an extent of 100–16 cm⁻¹. This obvious shift indicates that the nitrogen, and oxygen atoms of (C=N), and C=O groups in the ligand, HL, has been used to coordinate the metal ions [10]. Observing the non-ligand peaks of medium intensity at a range of 602–548 cm⁻¹ agree with metal–oxygen, and metal-nitrogen vibrations supports all observations mentioned previously [15].

• All metal complexes lack the peak corresponding to NH moiety of sulfadimidine, which suggests that it was deprotonated during the complex formation [16]. Actually, it is difficult to recognize the behavior of NH group in complex formation from IR spectra data as result to presence molecules of water in metal complexes framework.

• However, a deprotonation assumption was postulated based on the conductance measurements in addition to the elemental analysis and confirmed using ¹H-NMR spectra study as explained later. Additionally, the idea that the N-atom of the S–N fragment has been coordinated to the metal ions by its deprotonation is supported by the displacement of the characteristic vibrations caused by (S–N) in all metal complexes from 872 to 880 cm⁻¹ [17].

• The vibrational frequencies due to –SO₂– group are either remain unaffected or observed at slightly lower frequencies. This change in υ(SO₂) modes is linked to changes in the S–N bond upon coordination with metal ions [18].

• All synthesized complexes exhibited water molecules in their structure as suggested from the elemental analysis. The presence of a broad peaks at a range of 3354–3518 cm⁻¹ in all produced complexes indicates that coordinated/lattice water molecules were holed in complexes framework [19]. The thermal analysis study was applied to determine whether there are water molecules in complexes and to further confirm whether they are coordinated or lattice water molecules.

• For Ag(I) complex, the presence of two peaks at 1419, and 1296 cm⁻¹ with difference value 123 cm⁻¹, confirmed an unidentate manner for nitrate group [13].

Fig. 2

IR spectra of the ligand (HL), and its metal complexes

Nuclear magnetic resonance spectroscopic studies

The ¹H NMR investigation for ligand, HL as well as diamagnetic Ag(I) complex are shown in Fig. 3. The free ligand, HL has shown NMR signal at δ ppm of 10.97 comes from the sulfonamide proton (SO₂NH). This signal disappeared with deuterated solvent addition confirming its position [20]. The multiplets in the range δ 6.528–7.659 ppm may be assigned to the phenyl, pyrimidine, and coumarine rings protons [21, 22]. The signals observed at δ 2.499–2.551 ppm range corresponding to methyl protons attached to pyrimidine, and coumarine nucleus [23, 24]. The diamagnetic Ag(I) complex spectrum shows the most significant prominent point is that the signal corresponding to sulfonamide proton (NH) is absent confirming its deprotonation in complex formation [25]. All the other protons were found to be in their expected region as given in Table 3.

Fig. 3

¹H-NMR spectra of the ligand, (HL) and its Ag (I) complex

Mass spectral studies

Mass spectra results for ligand, HL, and its all complexes (Fig. 4) provided additional confirmations for their hypothesized structures. The ion peaks at m/z (intensity) 448.39 (73.96%), 779.13 (36.27%), 726.95 (22.44%), and 725.57 (24.04%) amu are in a good equality with their formula weights 448.54, 779.34, 725.80, and 725.32 amu for ligand HL, and Ag(I), Co(II), and Ni(II) complexes, respectively. These findings support the type of synthesized compounds as C₂₃H₂₀N₄O₄S for ligand HL, and [Ag₂(L)(NO₃)(H₂O)₃], [Co₂(L)(H₂O)Cl₃]0.2H₂O, and [Ni₂(L)(H₂O)Cl₃]0.2H₂O for Ag(I), Co(II), and Ni(II) complexes. Additionally, the mass spectra contain additional multiple peaks certified to the fragmentation of the investigated compounds. A schematic representation of the mass fragmentation for all prepared compounds is shown in Fig. 5a–d.

Fig. 4

Mass spectra of ligand, HL and its metal complexes

Fig. 5

Mass fragmentation pattern of a ligand, HL, and b Ag (I) complex. Mass fragmentation pattern of c Co (II) complex, and d Ni(II)

Molar conductance measurements

Dimethyl sulfoxide solvent was used to generate a solution of 1 × 10⁻³ mol L⁻¹ for each complex to assess their molar conductivities. The estimated values in range 2.42–11.62 Ω⁻¹ mol⁻¹cm² at ambient temperature indicate they were not electrolytic due to lack of counter ions outside coordination sphere [26].

Magnetic behavior of metal complexes

To realize magnetic features for synthesized complexes, and to comprehend their geometry, the magnetic susceptibility of these compounds was investigated. As would be anticipated for a d¹⁰ configuration, the Ag(I) complex is diamagnetic [27]. Complexes of Ni(II) and Co(II) ions were discovered to be paramagnetic, and their observed magnetic moments values at room temperature, Table 4 were well compatible with the expected number of the unpaired electrons. A review of literature revealed that the magnetic moment values for square planar, tetrahedral, and octahedral geometry of Co(II) complexes are, respectively, 2.2–2.9, 4.2–4.8, and 4.7–5.2 B.M [28]. A value 4.72 B.M of Co(II) complex indicates that Co(II) ion had a tetrahedral shape. This result is in accordance with three unpaired electrons, but it is greater than 3.87 B.M for spin-only value because of orbital contribution [29, 30]. A value 2.84 B.M of Ni(II) complex demonstrated existence of two unpaired electrons [29].

Table 4 Magnetic properties, and electronic spectra data (nm) for synthesized ligand (HL), and its metal complexes

Electronic spectral analysis

To learn more about the geometry of the complexes, free ligand, HL, and its Ag(I), Ni(II), and Co(II) complexes were examined for their electronic spectra as presented in Fig. 6a. Two peaks were visible in the ligand, HL spectrum at 325, and 349 nm. The first peak is caused by π − π* transition within phenyl, pyrimidyl, and coumarinyl rings in addition to C=O (coumarin carbonyl lactone), C=N (imine nitrogen), and SO₂NH (sulfonamide) groups in the ligand structure. The other peak is caused by n–π^* transition of the nonbonding electrons of the mentioned groups [31]. Upon ligand HL—metal ions interaction, and metal complexes formation, the distinctive ligand, HL peaks experience notable bathochromic, and hyperchromic shifts in the complexes spectra by sharing the unpaired electrons of C=N, C=O, and SO₂NH species in association with metal ions. Also, other new peaks were seen at a higher wavelength that resulted from charge transfer transitions of the LMCT type [31]. The identified peaks are listed in Table 4. No any d-d transitions are seen in Ag(I) complex spectrum contrary to what is predicted for the d¹⁰ configuration [32]. A spectrum of Ni(II) complex exhibits peak at 690 nm assigned to transition ³T₁(F) → ³T₁(P), whereas, Co(II) complex displays peak at 530 nm assigned to transition ⁴A₂(F) → ⁴T₁(P). A tetrahedral geometry was assigned for both complexes [33, 34].

Fig. 6

For the ligand-HL and its metal complexes; a UV–Vis absorbance spectra. b Dependent of the absorption coefficient (α) on the photon energy (hv). c Tauc’s plot of (ahv)^½svs. photon energy. d Tauc’s plot of (ahv).² vs. photon energy. e Absorption spectra fitting method for n = ½. f Absorption spectra fitting method for n = 2

Optical properties studies

Again, the UV–Vis absorption spectra data can be employed to evaluate the semiconducting properties of synthesized ligand, HL, and its complexes via band gap energy (E_g, eV) estimation [35].

Really, sequence steps are required to determine E_g value of these compounds. The first one is an absorption coefficient (α, cm⁻¹) calculation using absorbance (A), and thickness (t, cm) values [36].

$${\varvec{\upalpha}} \, = \frac{{{\mathbf{2}}{\mathbf{.303}}\left( {\varvec{A}} \right)}}{{\varvec{t}}}$$

For the studied compounds, the dependence of (α) on photon energy (hv) is graphically represented (Fig. 6b). The linear region in this plot is extrapolated to α = zero to get a fundamental absorption edge value (E^e, eV) [37]. The investigated values of E^e for all studied compounds are listed in Table 5.

Table 5 Optical band-gap energy, and some optical parameters of the synthesized ligand (HL), and its metal complexes

Continuously, Tauc’s plot is constructed using the calculated α–values to determine the band gap energy (E_g) for the studied compounds [36].

$${\varvec{\upalpha}}\,{\varvec{hv}} = {\varvec{C}}\left( {{\varvec{hv}} - {\mathbf{E}}_{{\mathbf{g}}} } \right)^{{\varvec{n}}}$$

h (Planck constant), ʋ (incident light frequency), C (proportionality constant), and n (optical frequency expressing electronic transition nature). The exponent ‘‘s’’ exhibits the values 1/2, 2, 3/2, and 3 for a transition mode; allowed direct, allowed indirect, forbidden direct, and forbidden indirect, respectively.

A matching Tauc’s plot for the different possible values of n = 1/2, and 2 is given in Fig. 6c, d that depicts graphing of (αhν)ⁿ along the Y-axis with incident photon energy (hs) along x-axis. At the intersection of the extrapolated linear portion of the figure and the photon energy axis, both direct (E^d_g), and indirect (Eⁱ_g) energy band gaps values were calculated and listed in Table 5 [38].

According to a literature review, the absorption spectra fitting (ASF) method is the one that accurately predicts the energy band gap. In this approach, the parameter (Aλ⁻¹)^1/n along the y-axis for the different possible values of n = 1/2, and 2 is plotted against the parameter (λ⁻¹) along the x-axis (Fig. 6e, f). Following this, a linear section of the graph was extrapolated onto x-axis, then λg values were found at the cut-off point. Hence, direct, and also, indirect energy band gap values (E_g^ASF) for the studied compounds are calculated and inferred in Table 5 [39].

$${\varvec{E}}_{{\varvec{g}}}^{{{\varvec{ASF}}}} = \frac{{{\mathbf{1239}}{\mathbf{.83}}}}{{{{\varvec{\uplambda}}}_{{\varvec{g}}} }}$$

Comparison of all obtained energy values as shown in Table 5, led to the suggestion that all synthesized compounds might undergo direct permitted transitions. Furthermore, using the previously determined values of the fundamental absorption edge (E^e), the parameter Ln (α hν) along the y-axis is plotted against the parameter Ln (hν-Eg) along the x-axis (Fig. 7a) to confirm the optical band transition mode of all synthesized compounds. A straight line is produced by this relationship, and its slope yields the power factor (n) [40]. The values of (n) discovered for the investigated compounds were found to be between 0.5, and 1.00, supporting the direct transition mode.

Fig. 7

For the ligand-HL and its metal complexes; a Plots of Ln(αhν) vs Ln (hν − E_g). b Variation of Ln(α) vs photon energy to find Urbach energy. c Variation of the extinction coefficient vs the wavelength. d Variation of the skin depth vs the wavelength

In accordance with the reported values (2 to 4 eV) of the variable semiconductors [41], the synthesized compounds can function as semiconductor materials with band gaps in the range 3.20–3.32 eV. Additionally, the metallization criterion (M) based on the band gap energy values can be estimated confirming their semiconductor nature [42, 43].

$${\varvec{M}} = \sqrt {\frac{{{\varvec{Eg}}}}{{{\mathbf{20}}}}}$$

Additionally, a graph drawn for ln (α) along y-axis with hv along x-axis (Fig. 7b) generates straight line allowing to estimate the Urbach energy (another sort of energy). Urbach energy (E_U, eV) is calculated using the inverse slope value for a straight line and listed in Table 5 [44].

After considering these factors, it is possible to analyze various optical properties for all synthesized compounds under study, such as refractive index (n; 2.31–2.33), steepness coefficient (σ_s; 6.05–9.35 × 10⁻²¹), and optical electronegativity (X*; 0.860 − 0.892). Additionally, an extinction coefficient (k), and skin depth (δ) were assessed [45,46,47,48].

$${\varvec{n}}^{4} = \frac{{{\mathbf{95}}{\varvec{eV}}}}{{{\varvec{E}}_{{\varvec{g}}}^{{\left( {\varvec{d}} \right)}} }};\quad {\varvec{\sigma}}_{{\mathbf{s}}} = \frac{{{\varvec{k}}_{{\varvec{B}}} {\varvec{T}}}}{{{\varvec{E}}_{{\varvec{U}}} }};\quad X* \, = \, 0.{\text{2688E}}_{{\text{g}}} ;\quad {\varvec{k}} = \frac{{{{\varvec{\upalpha}}}{\varvec{\lambda}}}}{{4{\varvec{\pi}}}};\quad {\varvec{\delta}} = \frac{{\mathbf{1}}}{{{\varvec{\upalpha}}}}$$

Moreover, the fluctuation for an extinction coefficient (k), and skin depth (δ) with wavelength is graphically depicted in Fig. 7c, d.

Thermal properties study (TG/DTG, and DSC)

Through using TGA-DTG, and DSC techniques, a thermal study for synthesized ligand, HL and, all complexes were done to explore their thermal response as well as to investigate the water molecules nature, and the metal content via weight loss estimation as a function of temperature [49]. TGA-DTG, and DSC curves of all studied compounds are given in Fig. 8a, b. The decomposition steps, temperature range, mass loss (%), mass loss assignments, leaving species, and the final residue are listed in Table 6.

Fig. 8

a TG-thermograme of the ligand, HL and its metal complexes. b DSC-grame of the ligand, HL and its metal complexes. c Coats-redfern plots of the ligand, HL and its metal complexes log y = log[log{W_∞(W_∞ − W)⁻¹}T⁻²

Table 6 Thermogravimetric data of the synthesized ligand (HL), and its metal complexes

The ligand, HL was found to have a good thermal stability, and its thermogram shows no weight loss was observed below 166 °C indicating its higher thermal stability. The ligand, HL, underwent two thermal breakdown stages at temperatures between 166–290, and 511–635 °C, respectively. The first stage of decomposition accompanied with degradation of the organic moiety, and evaporation of N₂, SO₂, 2CO, 5HC≡CH gases with an observed weight loss of 62.24% (Calcd: 62.05%). This stage of decomposition is characterized by the presence of two endothermic peaks from 197 to 293 °C in the DSC curve of the ligand, HL. The peak temperature, extrapolated onset temperature, and decomposition enthalpy are 197 °C, 195 °C, and − 218 mJmg⁻¹ for the first process, and 292 °C, 287 °C, and − 138 mJmg⁻¹ for the second process, respectively. The second step of decomposition with a weight loss 37.76% (Calcd: 37.59%) corresponding to a complete degradation of all organic moiety, and evaporation of ½N₂, CN, 5HC≡CH gases. At the end there is no residue left where a total weight loss is 100.00% (calcd 100.00%).

The thermogram of [Ag₂(L)(NO₃)(H₂O)₃] complex demonstrates three stages of breakdown between 131 and 598 °C supporting the hypothesis that no solvent or water molecules (hydrated) outside the coordination sphere [50]. A weight loss of 30.47% (calculated: 30.56%) was recorded during the first stage of decomposition, which began at 131 °C, and continued up to 310 °C. This coincides with the release of three water molecules (coordinated) in addition to SO₂, 2N₂, and 2O₂ gases. The DSC curve of Ag(I) complex offers two exothermic peaks from 283 to 392 °C at this stage of breakdown. The peak temperature, extrapolated onset temperature, and decomposition enthalpy for the first process, are 283 °C, 271 °C, and 228 mJmg⁻¹, whereas for the second process are 378 °C, 338 °C, and 33 mJmg⁻¹ respectively. The second stage of degradation begins at 401 °C and ends at 445 °C corresponding to a removal of CO, HCN, CH₄, and 2HC≡CH gases, resulting in loss of weight 16.17% (calcd 15.80%). The third decomposition stage manifested at a temperature of 572–598 °C, and mass loss of 15.80% (calcd 16.71%) was noted, which accounted for the loss of five molecules of the HC≡CH gas. The total weight loss of the decomposition steps was 62.44% (calcd 63.07%) leaving two silver, and six carbon atoms (found 37.56; calcd 36.93%).

The decomposition of [Ni₂(L)Cl₃(H₂O)]0.2H₂O complex began at 35 °C, and completed in three stages at 410 °C. The first one at 35–95 °C and associated with weight loss 3.75% (calcd 3.73%) corresponding to loss of one, and half molecule of water (hydrated). Within 102–200 °C temperature range, a second decomposition occurs with weight loss of 3.61% (calcd 3.73%) attributed to loss half molecule of water (hydrated), in addition to other one water molecule (coordinated). This stage is characterized by the presence of an endothermic peak in DSC curve of the Ni(II) complex takes place from 134 to 161 °C. The peak temperature, extrapolated onset temperature, and decomposition enthalpy are 161 °C, 109 °C, and − 352 mJmg⁻¹, respectively. H₂S, 2CO₂, 4CN, 3HCl, and 7HC≡CH gases were removed in the last stage resulting in weight loss 71.12% (calcd 71.40%). This stage is characterized by the presence of an endothermic peak in DSC curve takes place from 405 to 417 °C. The peak temperature, extrapolated onset temperature, and decomposition enthalpy are 417 °C, 388 °C, and − 506 mJmg⁻¹, respectively. Two nickel and three carbon atoms (found 21.52%; calcd 21.14%) were left behind after the disintegration stages, with an overall weight loss 78.48% (calcd 78.86%).

Decomposition of [Co₂(L)Cl₃(H₂O)]0.2H₂O complex took place in four steps at 35–105, 235–310, 350–460, and 499–529 °C, respectively. Two molecules of hydrated water should have been eliminated to account for the first mass loss of 5.40% (calcd 4.96%) at about 35–105 °C. This stage of decomposition is characterized by the presence of an endothermic peak from 165 to 194 °C as shown in DSC curve. The peak temperature, extrapolated onset temperature, and decomposition enthalpy are165°C, 145 °C, and 301 mJmg⁻¹, respectively. In the 235–310 °C range, weight loss 4.79% (calcd 4.41%) was due to removal of one water molecule (coordinated), and ½N₂ gas. The DSC curve exhibits an endothermic peak during this stage of breakdown, which occurs between 296 and 323 °C. The peak temperature, extrapolated onset temperature, and decomposition enthalpy are 296 °C, 281 °C, and 24 mJmg⁻¹, respectively. The third decomposition step begins from 350 to 460 °C, with weight loss 58.56% (calcd 59.65%), and production of various gases, including H₂S, HCN, N₂, CO₂, 2CO, ½Cl₂, and 8HC≡CH. This stage of decomposition is accompanied by the appearance of an endothermic peak in DSC curve which occurs between 372 and 392 °C. The peak temperature, extrapolated onset temperature, and decomposition enthalpy are 372 °C, 339 °C, and − 335 mJmg⁻¹, respectively. The last step occurs at 499–529 °C and is caused by the release of one Cl₂ gas molecule along with a weight loss of 9.90% (calcd 9.77%). Two cobalt and three carbon atoms (found 21.35%; calcd 21.21%) were left behind after the disintegration stages, with an overall weight loss of 78.65% (calcd 78.79%).

Kinetic and thermodynamic study

The Coats-Redfern method was employed to evaluate an essential kinetic and thermodynamic parameter for all various thermal events for synthesized ligand, HL and all complexes under investigation over the detected temperature range [51].

$${\mathbf{log}}\left[ {\frac{{{\mathbf{log}}\left\{ {{\mathbf{w}}_{{\mathbf{f}}} / \left( {{\mathbf{w}}_{{{\mathbf{f}} - }} {\mathbf{w}}} \right)} \right\}}}{{{\mathbf{T}}^{2} }}} \right] = \, {\mathbf{log}}\left[ {\frac{{{\mathbf{AR}}}}{{{\mathbf{\theta E}}^{*} }} \left( {1 - \frac{{2{\mathbf{RT}}}}{{{\mathbf{E}}^{*} }}} \right)} \right] - \frac{{{\mathbf{E}}^{*} }}{{{\mathbf{2}}{\mathbf{.303}}{\mathbf{RT}}}}$$

W_f (final weight for sample), W (actual weight for sample), T (temperature), A (Arrhenius pre-exponential factor), R (universal gas constant), θ (heating rate), and E* (activation energy). A left-hand side of an above equation can be plotted against 1/T as shown in Fig. 8c to produce straight line whose intercept yields an A value, while its slope yields an E* value [52].

$${\mathbf{Slope}} \, = \frac{{\mathbf{E}}}{{\mathbf{R}}}\varvec{ };\quad {\mathbf{intercept}} = {\mathbf{log}}\left[ {\frac{{{\mathbf{log}}\left\{ {{\mathbf{w}}_{{\mathbf{f}}} / \left( {{\mathbf{w}}_{{{\mathbf{f}} - }} {\mathbf{w}}} \right)} \right\}}}{{{\mathbf{T}}^{2} }}} \right]$$

Using the Eyring equation, an activation enthalpy, entropy, and free energy change (ΔH*, ΔS*, and ΔG*, respectively) were estimated [52]. All results are listed in Table 7.

$${\mathbf{\Delta H}}^{*} = {\mathbf{E}}* - {\mathbf{RT}};\quad {\mathbf{\Delta S}}^{*} = {\mathbf{2}}.{\mathbf{303}}\left( {{\mathbf{log}}\frac{{{\mathbf{Ah}}}}{{{\mathbf{KT}}}}} \right){\mathbf{R}};\quad {\mathbf{\Delta G}}^{*} = {\mathbf{\Delta H}}^{*} - {\mathbf{T}}\,{\mathbf{\Delta S}}^{*}$$

k (Boltzman constant), and h (Planck constant).

Table 7 Thermodynamic data of the thermal decomposition of the synthesized ligand (HL), and its metal complexes

The data obtained, shown in Table 7, show an increase in activation energy and free energy change of activation values from one step to another. This increase can be attributed to the complexes higher stability and the rigidity of the fragment produced by the decomposition process in comparison to the original state. Additionally, the positive ΔH, and negative ΔS values show the endothermic degradation process as well as the absence of any spontaneous process at all [53].

Powder X-ray diffraction analysis

There have been several attempts to get a precise structure for investigated ligand, HL, and all complexes using single crystals, but none of them have been successful. Otherwise, X-ray powder diffraction measurements were performed to obtain further evidence about the studied compounds, and to test the crystallinity degree in addition to other geometric parameters investigation. X-ray diffraction (XRDP) for the examined compounds, were performed at range of 2θ from 5 to 80° and are graphically displayed in Fig. 9. All information related to the diffractograms of the examined compounds that were collected, including inter-planar spacing (d, Å), diffraction angle (2θ, degrees), and relative intensity (I/I_o, %) for the peak are tabulated in Table 8. The observed and estimated d-values utilizing Bragg’s equation [54] have quite agreement, as shown in Table 8.

$$n\lambda = { 2}d\,{\text{sin}}\theta$$

where n is a diffraction order, and λ (nm) is the wavelength.

Fig. 9

X-ray diffraction patterns of ligand, HL and its metal complexes

Table 8 X-ray diffraction data of ligand (HL), and its metal complexes

The recorded XRDP of the free ligand, HL exhibits four diffraction peaks centered at 2θ values of 9.30, 18.44, 24.64, and 19.70° with d-spacing values of 9.51, 4.81, 3.61, and 4.50 Ǻ having an intensity (%) 100, 24.72, 11.92, and 22.12, respectively. The observed two characteristic peaks at 18.44, and 19.70° confirm the ligand, HL formation due to imine group [55].

After complexation process with metal ions, the recorded XRDP for metal complexes, Fig. 9 shows that the position of this peaks shift to higher 2θ diffraction angles with different intensity in addition to an appearance of other new peaks at different values of 2θ, and d-spacing that cannot be detected for free ligand, HL. Additionally, the obtained diffractograms of the metal complexes differ from each other in their diffraction patterns. All this confirms the metal complexes formation. The crystallite size (D) of the examined samples, Table 9 was determined using Debye–Scherrer method [55].

$$D = \frac{K\lambda }{{\beta \,{\text{cos}}\theta }}$$

The ligand HL had the largest crystallite size, measuring 37.16 nm, while Ag(I), Ni(II), and Co(II) complexes had crystallite sizes 30.66, 16.65, and 29.29 nm, respectively, as given in Table 9 showing that the examined compounds are nanocrystalline. The dislocation density (δ) can be approximated based on the D value. Furthermore, other important structural parameters are also investigated to strengthen the XRDP analysis, namely the microstrain (ε), distortion (g), stacking fault (SF), and the average crystallite separation (R) listing in Table 9 [55,56,57].

$$\delta = \frac{1}{{D^{2} }};\quad \varepsilon \, = \frac{\beta }{{4 {\text{tan}} \theta }} ;\quad g = \frac{\beta }{{ {\text{tan}} \theta }};\quad {\text{SF}} = \left[ {\frac{{2\pi^{2} }}{{45\left( {3 {\text{tan}}\theta } \right)^{1/2} }}} \right] \beta ;\quad R \, = \frac{5 \lambda }{{8 {\text{sin}} \theta }}$$

Table 9 The calculated geometric parameters data of ligand (HL), and its metal complexes

Surface morphology study

The surface morphology of the synthesized ligand, HL, was subjected to the scanning electron microscopy (SEM) analysis before, and after complexation. SEM images of all investigated compounds are depicted in Fig. 10. It is evident that, metal ions coordination with donor sites in ligand, HL, and formation of complexes causes that SEM micrographs to differ greatly [58]. Additionally, each metal complex shows a different typical surface morphology in its micrograph as needles (Ag(I) complex), uneven (Ni(II) complex), and flowers (Co(II) complex), whereas ligand, HL, has a block form in the micrograph. This demonstrates that metal ions affect how a surface morphology of metal complexes changes.

Fig. 10

SEM study of ligand, HL and its metal complexes

Theoretical investigations

Theoretical investigations utilizing density functional theory were employed to clarify and assign a greater degree of specificity to the structure of all synthesized compounds. Ligand, HL, and all complexes all had optimized geometries. Geometrical optimization is performed to obtain the lowest energy configuration from the initial guess of the geometry. The optimized structure and frontier molecular orbitals for the compounds under investigation are given in Figs. 11, 12.

Fig. 11

Geometrical optimized structure for the ligand HL, and its metal complexes

Fig. 12

DFT computed energy level diagram for the frontier molecular orbitals of the ligand, HL and its metal complexes in the gas phase

For all new synthesized compounds, the total electronic energy (E_total), molecular orbital energies that highest occupied (E_HOMO), and lowest unoccupied (E_LUMO) as well as dipole moment (D) were calculated then collected within Table 10. Additional parameters include HOMO–LUMO gap energy (ΔE), absolute hardness (η), chemical potential ($\pi$), absolute electronegativity (χ), global electrophilicity (ω), absolute softness (σ), ionization potential (IP), global softness (S), electron affinity (EA), and additional electronic charge (ΔNs_max) were estimated which are also shown in Table 10 [59, 60].

$${\mathbf{\Delta }}2E = 2E_{{{\mathbf{LUMO}}}} - 1muE_{{{\mathbf{HOMO}}}} ;\quad {\mathbf{X}}1mu = \frac{{ - \left( {2E_{{{\mathbf{HOMO}}}} + {\text{ }}2E_{{{\mathbf{LUMO}}}} } \right)}}{{\mathbf{2}}};\quad {\text{ }}2\eta 1mu = \frac{{2E_{{{\mathbf{LUMO}}}} - {\text{ }}2E_{{{\mathbf{HOMO}}}} }}{{\mathbf{2}}};\quad {\text{ }}2\pi = 1mu - {\text{ }}{\mathbf{X}};$$

$${\varvec{\sigma}} \, = \frac{1}{{\varvec{\eta}}};\quad {\varvec{S}} = \frac{1}{{2{\varvec{\eta}}}};\quad {\varvec{\omega}} \, = \frac{{{\varvec{\pi}}^{2} }}{{2{\varvec{\eta}}}};\quad {{\varvec{\Delta}}}{\varvec{N}}_{{{\mathbf{max}}}} = \frac{{\varvec{ \pi } }}{{\varvec{\eta}}};\quad {\mathbf{IP}} = - {\varvec{E}}_{{{\mathbf{HOMO}}}} ; \quad {\mathbf{EA}} = - {\varvec{E}}_{{{\mathbf{LUMO}}}}$$

Table 10 The calculated quantum chemical parameters data of ligand (HL), and its metal complexes

There are several significant notes based on the fundamental idea for molecular orbital theory and an information in Table 10 as follows:

• The estimated E_total was found to be − 1796.643 (ligand HL), and − 2204.315 to − 1816.061 (metal complexes). The more negative values of metal complexes suggest their better stability than its parent ligand, HL [61].

• The computed ΔE for the ligand, HL, is found to be 6.893 eV as opposed to 6.819 (Ag(I) complex), 7.128 (Ni(II) complex), and 3.720 (Co(II) complex) eV. Co(II) complex with lowest ΔE value is discovered to be the more chemically active, and the observed reactivity pattern as Co(II) complex > Ag(I) complex > Ni(II) complex. The low ΔE reflected the system's increased reactivity and reduced kinetic stability [59].

• It is clear that Ni(II) complex, the chemically hardest complex among all metal complexes, has the highest value (3.564 eV) in contrast to other complexes. Ni(II) complex > Ag(I) complex > Co(II) complex is the observed tendency in η values of metal complexes. Additionally, the metal complex σ values reveal that the soft Co(II) complex (0.538 eV) is the one. This pattern resembles the order of ΔE, in which the large ΔE value for hard molecule whereas, small ΔE value for soft molecule [61].

• Co(II) complex is shown to have the best value of IP among the metal complexes (5.374 eV). The lowest IP value that demonstrated more activity for the Co(II) complex is the optimal one [62].

• Meal complexes display higher dipole moment values than the ligand HL, and the Ag(I) complex has the maximum polarity.

Textural properties of metal complexes as catalysts

To further investigate the microstructure of the synthesized metal complexes, their surface textural properties including the specific surface area (m²g⁻¹), the pore diameter (nm), and the pore volume (cm³g⁻¹) were determined from N₂ adsorption–desorption isotherms. All obtained results are listed in Table 11, and graphically presented in Fig. 13a, b. The specific surface area, and the pore size distributions were calculated using the Brunauer- Emmett–Teller (BET), and Barrett-Joyner-Halenda (BJH) methods, respectively. From the BET, and BJH methods, an average specific surface area (average pore diameter) of 4.77, 16.41, and 3.65 m²g⁻¹ (29.02, 20.71, and 30.47 nm) was obtained for Ag(I), Co(II), and Ni(II)-complexes, respectively with a narrow pore size distribution centered at 0.04–0.14 nm. All metal complexes were characterized as mesopores as they had a porosity in range of 20.71–30.47 nm (The porosity of materials is usually classified as mesopores in a range of 2–50 nm) [63]. As can be seen from Table 11, Co(II) complex possesses the largest specific surface area, and smaller mesopores values of 16.41 m²g⁻¹, and 20.71 nm, respectively. This significant property provides more contact area for pollutants which could be more efficiently trapped and would be expected to enhance the photocatalytic performance [64].

Table 11 Textural properties of the synthesized metal complexes

Fig. 13

a The N₂ adsorption–desorption isotherm curves of metal complexes. b The pore size distribution by the BJH method of metal complexes

Catalytic photodegradation of aniline blue, and malachite green dyes

Generally, series of considerable factors associated with catalytic potentiality of the materials acting as catalysts such as reaction temperature, medium pH, ionic strength, solvent using, contact time between the catalyst, and pollutant species, pollutant nature, catalyst properties, pollutant concentration, catalyst dosage, interactions nature between catalyst, and pollutant species, and others as they control the rate of pollutant degradation [65]. In a current study, and to discover the catalytic potentiality for new synthesized complexes as catalysts, we chose two dyes: aniline blue (AB-dye), and malachite green (MG-dye) as exemplary organic pollutants. The experiments were run at room temperature, dyes concentration was 50 ppm, catalysts dosage was 0.02 g, and pH was that corresponding to dyes aqueous solution over 80 min. We focus on the type of selected dyes where AB-dye is anionic dye, and MG-dye is cationic dye. The photodegradation progress was spectrophotometrically monitored, and the observed outcomes are represented graphically in Fig. 14. The aqueous solution of AB, and MG dyes characterize with an intense absorption peak at 595, and 617 nm, respectively. The gradual reduction in such peaks accompanied with a color decolorization during the reaction progress illustrates the successful photodegradation process of the studied dyes. Moreover, all synthesized complexes exhibit an ability to degrade the studied dyes with different impacts resulting in different percentage of degradation efficiency (DE) [66].

$${\text{DE}}\left( \% \right) = \left( {1 - \frac{{A_{t} }}{{A_{o} }}} \right) \times 100 = \left( {1 - \frac{{C_{t} }}{{C_{o} }}} \right) \times 100$$

where A_t, and C_t, refer to absorbance reading, and concentration for the reaction time t, respectively, whereas A_o, and C_o, are absorbance reading, and the concentration at the initial time.

Fig. 14

Photodegradation process of aniline blue, and malachite green dyes up to 80 min using metal complexes as catalyst

The values of C_t were determined spectrophotometry. Degradation efficiency (%) with function of time was estimated for all metal complexes, given in Table 12, and depicted in Fig. 15a, b. The highest degradation percentage was observed for Co(II) complex up to 87.20, and 94.55% for AB-dye, and MG-dye, respectively after 80 min as compared to other complexes. The overall photo degradation power follow Co(II) complex > Ag(I) complex > Ni(II) complex towards both dyes. Furthermore, in the absence of the studied catalysts, and under the same conditions the photodegradation process indicates the removal of only 14.00, and 20.00% for AB and MG dyes, respectively after 80 min.

Table 12 Photodegradation efficiency evaluation (%) for the aniline blue, and malachite green using the metal complexes as catalysts

Fig. 15

(a-h): a, b Variation of photodegradation efficiency (%) with time (min); c, d Photodegradation profiles; and e, f Kinetic plot

Degradation kinetic, and suggested mechanism

The kinetics study for photodegradation process was examined to learn more about the general photodegradation process that was aided by the synthesized complexes. Assuming that adsorbed pollutant molecules amount is independent on concentration of pollutant in the medium, photocatalytic degradation kinetic of organic pollutants may typically be simplified as pseudo-first order [66].

$${\varvec{Ln}}\left( {\varvec{ }\frac{{{\varvec{A}}_{{\varvec{o}}} }}{{{\varvec{A}}_{{\varvec{t}}} }}} \right) = {\varvec{Ln}}\left( {\varvec{ }\frac{{{\varvec{C}}_{{\varvec{o}}} }}{{{\varvec{C}}_{{\varvec{t}}} }}} \right) = {\varvec{k}}_{{{\mathbf{app}}}} \,{\varvec{t}};\quad {\varvec{t}}_{\raise.5ex\hbox{$\scriptstyle 1$}\kern-.1em/ \kern-.15em\lower.25ex\hbox{$\scriptstyle 2$} } = \frac{{{\mathbf{0}}{\mathbf{.693}}}}{{{\varvec{k}}_{{{\varvec{app}}}} }}$$

where k_app (min⁻¹) refers to apparent photocatalytic rate constant, t (min) refers to reaction time, and t_1/2 (min) is the half-life time. Plotting of $\left(\frac{{C}_{t}}{{C}_{o}}\right)$ with time (Fig. 15c, d) shows fast AB, and MG dyes degradation was observed after about 20, and 15 min of illumination, respectively. Additionally, straight line diagram with slope equal to k_app, and extinction coefficient correlation (R²) values of 0.9516–0.9911, and 0.9699–0.9255, respectively, was obtained by plotting Ln $\left(\frac{{C}_{o}}{{C}_{t}}\right)$ with time (Fig. 15e, f), supporting the evidence that degradation process follows first order type [67].

An evaluated rate constant for Co(II) complex was 0.041, and 0.027 min⁻¹ with a corresponding half-life time 16.72, and 25.80 min towards AB, and MG dyes photodegradation, respectively. This results proved the superiority of Co(II) complex reaffirming its excellent photocatalytic traits confirmed by its lowest band gap energy value. Moreover, the enhanced specific surface area for Co(II) complex as clear from BET analysis could bring more active centers, leading to good adsorption performance, and enhanced photocatalytic performance.

It can be seen, there are differences in the decolorization efficiency of the studied dyes solutions by the metal complexes, where the removal efficiency for MG-dye was higher than AB-dye. In our opinion, this was attributed to the difference in the adsorption capacities of both studied dyes arising from the two studied dyes have different nature, functional groups, molecular weight, and molecular size as located in Table 1. Furthermore, an easy penetration for MG-dye into the catalyst molecules on comparison to AB-dye due to lack of steric hindrance, and less molecular size causes greater adsorption of MG-dye into the catalyst molecules [68].

A mechanism for photocatalytic degradation process was proposed as [69, 70];

• Photocatalytic degradation takes place at catalyst’s surface, and the extent of photocatalytic activity is largely reliant on how well the photocatalyst can bind dye molecules. The removal process’ efficiency enhanced with an increase for dye adsorption on the catalyst.

• When catalyst’s surface is exposed to light with the proper energy (hy), electron from catalyst’s valence band (VB) is stimulated to conduction band (CB) and leaves positive charge hole (h +) there. A most important factor in improving photocatalytic activity is the successful suppression of photogenerated electron–hole pair recombination.

• When the photocatalytic process runs in an aqueous solution, a generated h + in VB oxidizes water molecules into hydroxyl free radicals (OH) while, electrons in CB reduces oxygen to superoxide ions (O₂⁻).

• Both O₂⁻, and OH are responsible to degrade the adsorbed dye molecules into CO₂ and H₂O.

• A schematic presentation of the dyes degradation process by the created complexes is illustrated in

Fig. 16, and possible steps are given below.

$${\mathbf{M}} - {\mathbf{complex}} \, + \varvec{h\nu } \, \to {\mathbf{M}} - {\mathbf{complex}} \, \left\{ {{\varvec{h}}^{ + } \left( {{\mathbf{VB}}} \right) + {\mathbf{e}}^{ - } \left( {{\mathbf{CB}}} \right)} \right\}$$

$${\mathbf{M}} - {\mathbf{complex}} \, \left\{ {{\varvec{h}}^{ + } \left( {{\mathbf{VB}}} \right)} \right\} + {\mathbf{H}}_{{\mathbf{2}}} {\mathbf{O}} \, \to {\mathbf{OH}} \, + \, {\mathbf{H}}^{ + }$$

$${\mathbf{M}} - {\mathbf{complex}} \, \left\{ {{\mathbf{e}}^{ - } \left( {{\mathbf{CB}}} \right)} \right\} + \, {\mathbf{O}}_{{\mathbf{2}}} \to \, {\mathbf{O}}_{{\mathbf{2}}}^{ - }$$

$${\mathbf{OH}}/{\mathbf{O}}_{{\mathbf{2}}} + {\mathbf{dyes}}_{{{\mathbf{ads}}}} \to \, {\mathbf{Intermediate}} \, + {\mathbf{H}}_{{\mathbf{2}}} {\mathbf{O}} \, + \, {\mathbf{CO}}_{{\mathbf{2}}}$$

Where; M-complex represents Ag(I), Ni(II) or Co(II) complex.

Fig. 16

Suggested illustration of aniline blue, and malachite green dyes photodegradation

Comparative analysis of catalytic performance of new created complexes with other reported catalysts

Huge number of photocatalysts have been prepared, studied, and reported in literature for.

AB, and MG dyes degradation. We were concerned to compare how well the newly synthesized complexes worked as catalysts with reported results for earlier investigations on degradation of AB and MG dyes [69, 71,72,73,74,75,76,77,78,79]. The results are presented in Table 13. The as-produced Ag(I), Ni(II), and Co(II) complexes shown adequate performance for AB, and MG- dyes degradation, and can be employed for successful dyes removal from water when compared to other photocatalysts.

Table 13 Comparison of aniline blue, and malachite green dyes degradation by Ag(I), Co(II), and Ni(II) complexes with other reported catalysts

Conclusion and future perspectives

In conclusion, by condensing sulfadimidinem with 3-acetylcoumarin, we were able to design and synthesize the ligand, HL a novel multi donor functionalized ligand. [Ag₂(L)(NO₃)(H₂O)₃], [Co₂(L)Cl₃(H₂O)]0.2H₂O, and [Ni₂(L)Cl₃(H₂O)]0.2H₂O complexes for ligand, HL were synthesized. All new synthesized compounds were fully described with the help of spectroscopic techniques, and elemental analysis. Ligand, HL interact with metal ions forming tetrahedral binuclear complexes. The semiconductor nature was suggested for all studied compounds. Density functional theory was developed to verify their optimized structures. Some quantum chemical parameters were also performed to give more insight on their reactivity, and kinetic stability. The photocatalytic activity of all synthesized complexes was also assessed in mineralization for AB, and MG dyes. The findings of this test indicate that a set of physical and chemical characteristics plays a crucial role in the governing of catalytic activity. All synthesized metal complexes are promising advanced photocatalytic material, which can be used for various environmental applications. The BET surface area, and the porosity data align well with the relative photocatalytic performances of the synthesized metal complexes. An excellent photocatalytic performance was obtained for Co(II)-complex toward the AB, and MG dyes degradation. Additionally, these complexes have practical significance due to how easily these catalysts can be separated from the cleaned water.