Green Synthesis of Silver Oxide Microparticles Using Green Tea Leaves Extract for an Efficient Removal of Malachite Green from Water: Synergistic Effect of Persulfate

Abstract

The removal of water pollutants by photocatalysis is a promising technique, mainly due to its environmentally friendly and sustainable nature. In this study, the degradation of a recalcitrant organic pollutant, malachite green (MG), was investigated in water by a microstructured silver oxide photocatalyst. The silver oxide (Ag₂O) microparticles (MPs) were synthesized by a low-cost, green method, mediated by green tea leaves extract. The surface, morphological and optical properties of the synthesized Ag₂O MPs were determined by scanning electron microscopy (SEM), energy dispersive X-ray (EDX) analysis, X-ray diffraction (XRD) analysis, Fourier transform infrared (FTIR) spectroscopy, and ultraviolet-visible (UV-visible) spectrophotometry. The synthesized Ag₂O MPs showed good photoactivity, represented by 83% degradation of malachite green (MG) ([C]₀ = 0.4 mM, Ag₂O loading = 0.1 g L⁻¹) at neutral pH, in 3 h. Persulfate ions (PS) showed a strong synergistic effect on the efficiency of solar/Ag₂O photocatalysis, represented by complete MG removal in 15 min, in the presence of 1.6 mM PS. The results revealed that solar/Ag₂O, particularly solar/Ag₂O/PS photocatalysis is a promising method for the elimination of toxic organic pollutants, such as malachite green, from the water environment.

1. Introduction

Freshwater resources, especially groundwater, are heavily polluted by the synthetic organic dyes discharged from various industries, particularly textiles, paint, cosmetics, and paper industries [1]. A large amount of synthetic organic dyes is regularly produced, mainly for dying, staining, or coloring purposes, ultimately entering the environment, especially water systems [2]. Many synthetic dyes are toxic organic compounds and cause severe health problems to human beings and animals, including immunological, reproductive, and developmental abnormalities, as well as carcinogenic and mutagenic activity [1,3,4,5]. Thus, water pollution by synthetic dyes is a big risk to the environment.

Malachite green (MG) is an organic dye extensively used in staining cotton, fabrics, silk, and leather [2]. MG is used in the production of paints and inks [6]. MG is also used as an antimicrobial agent in aquaculture [5]. It is a toxic compound of significant environmental concern [7]. Particularly, MG causes risks to human beings, including immunology, reproductive system, and liver issues [3], developmental abnormalities, and carcinogenesis [4]. The removal of toxic organic dyes, especially malachite green (MG), from water is a critical task, and several different methods, including adsorption [8,9], microbial [10], and advanced oxidation processes (AOPs) [11,12,13], have been employed. Among them, advanced oxidation processes (AOPs) were promising techniques because they have low process complexity, produce less harmful by-products, are cost-effective, and have been found to be highly efficient for the removal of toxic organic pollutants [14]. Many pollutants are difficult to remove using only physical, chemical, or biological means in many water and wastewater treatment applications. A more powerful oxidation process, known as advanced oxidation, is required for more efficient removal. AOP produces strong oxidizing agents with the potential to reduce pollutant levels from hundreds of parts per million (ppm) to a few parts per billion (ppb). These radicals attack almost all organic materials because they are non-selective. These contaminants form intermediates after being broken down by the radical once. The intermediates then react with the oxidants to form stable compounds. Some well-known oxidizing agents include hydrogen peroxide, peroxymonosulfate, and persulfate, which degrade organic pollutants relatively efficiently and rapidly when photocatalysts such as metal oxides or Fe(II) in the Fenton or photo Fenton process are used [15,16,17,18].

The advanced oxidation processes (AOPs) are represented by Fenton’s and photo-Fenton’s processes [19,20], ionizing radiation [21], and photocatalysis [22,23]. A common characteristic of the advanced oxidation processes (AOPs) is the generation of highly oxidizing species, mostly hydroxyl radical (HO^•) and/or sulfate radical anion (SO₄^•−) [24]. Both hydroxyl radical (HO^•) and sulfate radical anion (SO₄^•−) have high oxidation potential, i.e., E° (SO₄^•−) = 2.4 V and E° (HO^•) = 2.7 V, and are abundantly used for the oxidation and/or degradation of the most organic pollutants [25]. Advancement and development in the field of advanced oxidation processes (AOPs), mainly for environmental and sustainable purposes, is a rising research field, and solar photocatalysis is one of the most promising options [14].

Solar photocatalytic advanced oxidation processes (AOPs) are efficient methods for the degradation of recalcitrant organic pollutants in water, such as pesticides, pharmaceuticals, and dyes [22]. Different photocatalysts, such as TiO₂ [26], iron [27], silver [9], zinc [28], palladium [29], and gold [30] nanoparticles, have been used for decomposing toxic organic dyes in water. Metal oxide nanoparticles are considered useful photocatalysts in the dyes’ degradation and microbial inactivation [31].

To synthesize metal oxide nanoparticles or microparticles with photocatalytic applications, green synthesis methods using plant extracts are an emerging field due to their simplicity and high productivity [31]. Unlike physicochemical methods, green synthesis produces no hazardous products or by-products [32]. As a result, green synthesis of nanoparticles using various types of plant extract has been reported [33]. One such example is the synthesis of silver oxide (Ag₂O) nanoparticles [34,35], which is environmentally friendly, cost-effective, biodegradable, and easily accessible for large-scale application [36]. Silver oxide nanoparticles were synthesized using Genwa Leaf [37] and Phoenix dactylifera L. [38]. Similarly, green tea leaf extract was used to synthesize gold nanoparticles (GNP) [39]. The organic coating on the surface of green synthesized nanoparticles, which makes them more stable than chemically produced nanoparticles, is the most important factor to consider [35]. As a result, silver oxide microparticles were greenly synthesized in this study, and their photocatalytic applications were revealed, considering the economic feasibility, environmental friendliness, and simple approach.

This study describes a low-cost green synthesis of silver oxide microparticles (Ag₂O MPs) that employs green tea leaf extract as a reducing and stabilizing agent. SEM, EDX, XRD, FTIR, and UV-visible spectrophotometric techniques were used to identify the structural, morphological, and optical properties of the synthesized Ag₂O MPs. The photocatalytic efficiency of the synthesized Ag₂O MPs in solar radiation was investigated for the degradation of a model organic dye, malachite green (MG), in water. Furthermore, the synergistic effect of persulfate (PS) on the efficiency of solar/Ag₂O photocatalysis for MG degradation was determined. The results showed that solar/Ag₂O photocatalysis, particularly the solar/Ag₂O/PS system, is a promising method for completely removing toxic organic pollutants from the water environment, such as malachite green.

2. Results and Discussion

2.1. Characterization Studies

The structure, morphology, and optical properties of the synthesized Ag₂O MPs were studied by FTIR, XRD, SEM/EDX, and UV/Vis spectrophotometric techniques, and the results are discussed in the succeeding section.

2.1.1. Optical Study

The UV-visible and energy band gap spectra of Ag₂O MPs are shown in Figure 1a,b, respectively. Figure 1a shows a sharp absorbance spectrum at 304 nm, which indicates the presence of silver oxide microparticles. However, Figure 1b shows a 3.95 eV energy band gap for silver oxide microparticles, which was calculated by using Tauc equation αhν = A [ hν − E_g ]ⁿ; where α is the absorption coefficient, hν is the incident photon energy, A is the constant, E_g is the energy band gap, and n is an index with a value 2 [40].

2.1.2. FTIR Study

The FTIR spectrum of the synthesized Ag₂O MPs is shown in Figure 2. The bands observed at 1065 and 516 cm⁻¹ may be assigned to O-Ag-O and Ag-O vibrations [41] existing in Ag₂O MPs, synthesized from green tea leaf extract. The other peaks obtained were probably due to the green tea leaf extract and silver nitrate used in the synthesis process.

2.1.3. X-Ray Diffraction Analysis

Figure 3 represents the X-ray diffraction spectrum of Ag₂O MPs. The spectrum is consistent with the standard spectrum of Ag₂O MPs cited in the literature. The characteristic peaks of diffraction at two theta values confirmed the Miller indices (hkl) of Ag₂O MPs were 27.8° (100), 32.2° (110), 38.2° (111), 44.2° (200), 64.5° (220), and 79.5° (311) analogous with JCPDS card number 00-076-1393 [42]. The crystallite size of Ag₂O MPs was 7.95 nm, as calculated by Deby-Scherrer Equation (1).

τ = κλ/βcosθ

(1)

where ‘τ’ is the size of the particle, ‘κ’ is the Scherer’s constant (κ= 0.94), ‘λ’ is the X-ray wavelength (1.54178Å), ‘β’ is full width at half maximum of the diffraction peak, and ‘θ’ is the Bragg angle.

2.1.4. SEM-EDX Analyses

Figure 4a,b show the SEM micrograph of Ag₂O MPs under low and high resolutions. The micrographs indicate irregularly dispersed particles of almost sphere shape with separate limits. The mean particle size, i.e., 0.40 μm (400 nm), was predicted by selecting 100 random features from Figure 4b using Nano Measurer 1.2.5 software. The Ag₂O particle size distribution ranged from 0.06 μm (60 nm) to 1.87 μm (1870 nm), as displayed in Figure 4c.

The EDX spectrum of silver oxide microparticles is shown in Figure 4d, which represents peaks in the range of 0.0–14.0 keV. The four small peaks at 0.2, 0.4, 0.6, and 0.8 keV indicate the presence of carbon, oxygen, copper, and zinc, respectively. A strong band is allocated to silver at 3 keV. The percentage mass of carbon, oxygen, copper, zinc, and silver predicted in the sample were 12.8, 15.4, 2.5, 2.0, and 66.5, respectively. The presence of carbon, copper, and zinc could be due to the green tea leaves extract.

2.2. Degradation of Malachite Green (MG) by Ag₂O MPs

The MG solution (0.4 mM) was examined without the addition of a catalyst (Ag₂O) or oxidant (persulfate) to test the auto/self-degradation in the absence and presence of solar radiation for 3 h at neutral pH. Figure 5a does not demonstrate any significant MG degradation. The degradation of MG by Ag₂O MPs was studied in deionized water at neutral pH and room temperature, i.e., 27 ± 1 °C, without introducing sunlight. In order to avoid the interaction of radiation with the reaction mixture, the glass vial that contained the reaction mixture was wrapped in aluminum foil (Figure 5b). An aqueous solution of MG (0.4 mM) was loaded with 0.1 g L⁻¹ of the synthesized Ag₂O MPs and the time course spectra were recorded (Figure 5b). The MG dye was degraded with the passage of time, and the degradation reaction followed zero-order kinetics with the observed rate constant (k_obs), 2.7 × 10⁻⁷ M min⁻¹ (Figure 5c). The slope of the plot yielded ɛbk_obs according to Equation (2) because we monitored absorbance (A) as a function of time (t). The value of molar absorptivity (ɛ) of MG was determined to be 2411 M^–1 cm^–1 and the path length (b) of the quartz cuvette was 1 cm. The degradation of MG is zero-order, corresponding to Ag₂O, because the reaction was studied under the pseudo-first-order condition with excess MG over Ag₂O. Figure 5d shows that about 32.5% MG was degraded by Ag₂O MPs in the absence of solar radiation in 3 h.

$${\left[\mathrm{A}\right]}_{\mathrm{t}}={\left[\mathrm{A}\right]}_0-{\mathsf{\epsilon }\mathrm{bk}}_{\mathrm{obs}}\mathrm{t}$$

(2)

2.3. Photocatalysis of the Degradation of MG by Ag₂O MPs

Figure 6a shows the spectra of the degradation of MG by solar/Ag₂O MPs in an aqueous medium as a function of time. An aqueous solution of MG (0.4 mM) was loaded with 0.1 g L⁻¹ of the synthesized Ag₂O MPs using a similar procedure as described in Section 2.2. Time course spectra were recorded after exposing the transparent glass vial containing the reaction mixture to direct sunlight, indicating zero-order kinetics corresponding to Ag₂O MPs and 83% MG degradation in 3 h (Figure 6b,c). The observed rate constant, k_obs, of the photocatalytic degradation of MG by solar/Ag₂O was estimated to be 1.2 × 10⁻⁶ M min⁻¹ (Figure 6b). The degradation of malachite green by solar/Ag₂O photocatalysis was due to the generation of hydroxyl radical (HO^•), according to Equations (3)–(10) [43]. Wang et al. reported that Ag₂O is a highly efficient visible light-activated photocatalyst for the degradation of organic contaminants in water [44].

Ag₂O + hν ⟶ Ag₂O (h⁺_VB + e⁻_CB)

(3)

2Ag₂O + e⁻_CB ⟶ 4Ag⁺ + O₂⁻

(4)

e⁻_CB + O₂ ⟶ O₂⁻

(5)

O₂⁻ + e⁻_CB + 2H⁺ ⟶ H₂O₂

(6)

e⁻_CB + H₂O₂ → HO^• + OH⁻

(7)

h⁺_VB + H₂O → HO^• + H⁺

(8)

Dye + HO^•

(9)

Dye + h⁺_VB

(10)

2.4. Synergistic Effect of Persulfate (PS) on the Degradation of MG by Ag₂O MPs

Figure 7a,b show the degradation efficiency of Ag₂O/persulfate system ([PS]₀ = 1.6 mM, [MG]₀ = 0.4 mM, Ag₂O = 0.1 g L⁻¹, pH = 1), represented by the complete removal of MG in water within 21 min in the absence of solar radiations, with the corresponding observed first-order rate constant (k′_obs), i.e., 0.174 min⁻¹ at room temperature (T = 21 ± 1 °C). The reaction was studied under pseudo-second order condition by keeping PS in excess over MG that was maintained in excess over Ag₂O. The first-order kinetics was observed with respect to MG. The integration method was applied to the absorbance data that was collected as a function of time to plot a linear graph according to Equation (11). The slope of the plot yielded the observed rate constant (k′_obs), i.e., 0.174 min⁻¹. The efficient degradation of MG by the Ag₂O/PS system was due to the generation of SO₄^•−, as expressed by Equation (12). Furthermore, additional HO^• could also be produced due to the reaction of SO₄^•− with H₂O, according to Equation (13), which was scavenged by the addition of protons via studying the degradation at pH = 1.

$$\ln {\left[\mathrm{A}\right]}_{\mathrm{t}}=\ln {\left[\mathrm{A}\right]}_0-{\mathrm{k}}_{\mathrm{obs}}^{\text{’}}\mathrm{t}$$

(11)

The degradation efficiency of the Ag₂O/PS system was further enhanced under the sunlight, i.e., solar/Ag₂O/PS, represented by the complete removal of MG in water at pH = 1. The increased efficiency of the Ag₂O/PS system was demonstrated by a reduction in degradation time of up to 15 min and a high value of the observed first-order rate constant (k″_obs), i.e., 0.229 min⁻¹ (Figure 8a–c). A plausible reason could be the generation of more SO₄^•−, according to Equation (14). However, the increase in acidity could lead to the formation of excess SO₄^•− on scavenging of hydroxyl radicals by excess protons and yielding water. Saleh et al. reported that the methylene blue (MB) degradation efficiency of Ag₂O/TiO₂/nanographene platelet (Ag₂O/TiO₂/NGP) system was enhanced in the presence of H₂O₂ and/or PS [45] in water attributed to the same reasons.

S₂O₈²⁻ + e⁻_CB ⟶ SO₄^•− + SO₄²⁻

(12)

SO₄^•− + H₂O ⟶ HO^• + SO₄²⁻ + H⁺

(13)

S₂O₈²⁻ + hν ⟶ 2SO₄^•−

(14)

2.5. Degradation Kinetics of MG with Respect to the Initial Concentration of MG

The effect of the initial concentration of MG on its degradation kinetics was studied under the pseudo-second-order condition, i.e., PS >> MG >> Ag₂O. The concentration of oxidant and catalyst doses such as [PS]₀ = 1.6 mM, and, Ag₂O = 0.1 g L⁻¹ were constant at neutral pH and room temperature, i.e., 27 ± 1 °C. The initial concentration of MG was varied, such as [MG]₀ = 0.1, 0.2, 0.4, 0.6, and 0.9 mM. The rate of degradation of MG was found to show fractional (1/2 or 0.5) order dependence on the concentration of MG according to Equation (15). Consequently, the reaction followed a fractional (1/2) order in MG. Figure 9a–c display zero ([A]_t vs. time), first (ln[A]_t vs. time), and fractional ([A]_t^0.5 vs. time) order linear plots, respectively, according to the integration method. The best linear regression fit was observed for the fractional (1/2) order kinetics with R² close to 1 (Figure 9c). The fractional order rate constant (k″_obs), however, declined with increasing initial MG concentration and became virtually constant at the highest MG concentration with the corresponding rate constant, (k″_obs)_max, as shown in Figure 9d. The rate constant (k″_obs) was not significantly affected by further increases in the initial MG concentration. The outcome demonstrates that the slowest rate of MG degradation was found at the highest initial concentration of MG (Figure 9e, i.e., plot of percent degradation vs. time). The reciprocal of the initial MG concentration was plotted against the fractional order rate constant (k″_obs) in order to obtain the value of (k″_obs)_max. A plot of k″_obs versus [MG]^–1 yielded a linear fit with R² = 0.97 (Figure 9f). The intercept of the plot yields the ideal value of the rate constant, i.e., 5.73 × 10⁻⁴ M^0.5 min⁻¹ at the infinite concentration of MG that may also refer to the rate constant at the maximum concentration of MG, i.e., (k″_obs)_max, such as the phase when the rate constant becomes independent of the concentration of dye.

$${\left[\mathrm{A}\right]}_{\mathrm{t}}{}^{0.5}={\left[\mathrm{A}\right]}_0{}^{0.5}-0.5{\left(\mathsf{\epsilon }\mathrm{b}\right)}^{0.5}{\mathrm{k}″}_{\mathrm{obs}}\mathrm{t}$$

(15)

2.6. Degradation Kinetics of MG with Respect to PS

The effect of variation in PS concentration on the kinetics of MG degradation was examined under the pseudo-second-order condition of PS >> MG >> Ag₂O. The MG concentration and catalyst dose, such as [MG]₀ = 0.4 mM and Ag₂O = 0.1 g L⁻¹, were constant at neutral pH and room temperature (27 ± 1 °C). The initial concentration of PS was varied, such as [PS]₀ = 0.8, 1.2, 1.6, 2.0, and 2.4 mM. The results revealed that the rate of MG degradation has a fractional (0.5) order dependence on the concentration of MG. As a result, the reaction occurred in MG in a fractional order. According to the integration method, Figure 10a–c show zero ([A]_t vs. time), first (ln[A]_t vs. time), and fractional ([A]_t^0.5 vs. time) order linear plots, respectively. For the fractional order kinetics, a linear regression fit with an R² close to 1 was found to be the best (Figure 10c). However, the fractional order rate constant rose linearly with increasing initial PS concentration, revealing first order in PS, and the negative value of the intercept indicates that MG is degrading independently of PS by the catalyst Ag₂O as was discussed in Section 2.2 and Section 2.3 (Figure 10d). According to Equations (16–18), the slope of the plot results in the overall fractional (3/2) order degradation rate constant (k = 0.794 M^−0.5 min⁻¹), as seen in Figure 10d.

$$\mathrm{Rate}=\mathrm{k}{\left[\mathrm{MG}\right]}^{0.5}\left[\mathrm{PS}\right]{\left[{\mathrm{Ag}}_2\mathrm{O}\right]}^0$$

(16)

We can say that the pseudo-second-order condition was kept by keeping PS in excess over MG, which was kept in excess over Ag₂O. Equation (16) is therefore simplified to Equation (17).

$$\mathrm{Rate}={\mathrm{k}″}_{\mathrm{obs}}{\left[\mathrm{MG}\right]}^{0.5}$$

(17)

Since,

$${\mathrm{k}″}_{\mathrm{obs}}=\mathrm{k}\left[\mathrm{PS}\right]$$

(18)

However, Figure 10d demonstrates that an outlayered value of the observed fractional order rate constant (k″_obs) is attained when the concentration of PS is 0.8 mM, which is only two-fold the concentration of MG. As a result, it has been declared that the just-doubled PS concentration to MG concentration does not satisfy the pseudo-second-order condition. In order to maintain the pseudo-order condition, the concentration of PS should be in excess of the concentration of MG. However, Figure 10e represents the degradation kinetics of MG as a function of time at a varying concentration of PS.

Figure 9. Kinetic study of the degradation of MG corresponding to the initial dye concentration at neutral pH, 1.6 mM PS, 0.1 g L⁻¹ Ag₂O, and 27 ± 1 °C. (a) Zero-order plot. (b) First-order plot. (c) Fractional (0.5) order plot. (d) Plot of fractional-order rate constant versus [MG]. (e) Plot of percent degradation of MG as a function of time at a varying concentration of MG. (f) Plot of the fractional-order rate constant versus reciprocal of [MG].

Figure 9. Kinetic study of the degradation of MG corresponding to the initial dye concentration at neutral pH, 1.6 mM PS, 0.1 g L⁻¹ Ag₂O, and 27 ± 1 °C. (a) Zero-order plot. (b) First-order plot. (c) Fractional (0.5) order plot. (d) Plot of fractional-order rate constant versus [MG]. (e) Plot of percent degradation of MG as a function of time at a varying concentration of MG. (f) Plot of the fractional-order rate constant versus reciprocal of [MG].

Figure 10. Kinetic study of the degradation of MG corresponding to PS concentration at neutral pH, 0.4 mM MG, 0.1 g L⁻¹Ag₂O, and 27 ± 1 °C. (a) Zero-order plot. (b) First-order plot. (c) Fractional (0.5) order plot. (d) Plot of fractional-order rate constant versus [PS]. (e) Plot of percent degradation of MG as a function of time at a varying concentration of PS.

Figure 10. Kinetic study of the degradation of MG corresponding to PS concentration at neutral pH, 0.4 mM MG, 0.1 g L⁻¹Ag₂O, and 27 ± 1 °C. (a) Zero-order plot. (b) First-order plot. (c) Fractional (0.5) order plot. (d) Plot of fractional-order rate constant versus [PS]. (e) Plot of percent degradation of MG as a function of time at a varying concentration of PS.

2.7. Effect of pH on the Degradation Kinetics of MG

The degradation of MG was studied at various pH levels, including acidic, neutral, and basic. The reaction was studied under the pseudo-second-order condition, i.e., PS >> MG >> Ag₂O. Except for pH, the experimental conditions such as [MG]₀ = 0.4 mM, [PS]₀ = 1.6 mM, and Ag₂O = 0.1 g L⁻¹ were kept constant at 27 ± 1 °C. To study the effect of pH on the degradation of MG, pH levels of 1, 3, 5, 7, and 8 were varied. Figure 11a–f reveal the effect of pHs on the degradation kinetics of MG. In acidic pHs of 1, 3, and 5, the MG degradation reaction followed a first-order kinetics corresponding to MG (Figure 11c). However, at neutral and basic pHs of 7 and 8, respectively, MG degradation occurred in the usual fractional order (Figure 11b). The results also show a linear relationship between the observed first-order rate constant (k′_obs) and the reciprocal of pHs in the acidic range, with the best linear regression R² value near 1 (Figure 11f). The results helped to conclude that the formation of hydroxyl radical (Equation (13)) is reduced during the reaction in an acidic medium by yielding water, and thus the degradation is accelerated by the synergistic effect of PS ion. As a result, the reaction order with respect to MG concentration is raised to first from fractional (1/2) order. The rate of percent degradation of MG shows the following decreasing order with reference to pH of the solution. Because of the scavenging of hydroxyl radical by free protons and hydroxyl ion at pHs 1 and 8, degradation was rapid. It is worth noting that a minimum 1 M nitric acid solution acts as an oxidizing agent; we used a 0.1 M solution to maintain pH [46].

1 > 8 > 3 > 5 > 7

2.8. Reusability of Ag₂O MPs

The reusability of Ag₂O MPs was tested over four cycles by keeping reaction parameters constant such as [PS]₀ = 1.6 mM, [MG]₀ = 0.4 mM, Ag₂O = 0.1 g L⁻¹, and 27 ± 1 °C at neutral pH. Following filtration of the first and subsequent used solution, the Ag₂O MPs were collected on filter paper and washed with deionized water, dried, and reused. The catalyst performs admirably, beginning with 100% degradation in the first use and progressing to 97.3% degradation in the third and fourth uses. This minor variation in catalyst efficiency may be encountered by using the catalyst again after collecting from filter paper, which may result in a minor weight loss of the catalyst at the filter paper. Meanwhile, in the first use, the MG degradation kinetics followed fractional order as expected, but from the second to fourth use, it went through first order, possibly due to surface changes in the catalyst (Figure 12a–c). From the first to the fourth use, the degradation rate of MG was gradually reduced may because of the weight loss, as mentioned earlier. The results are depicted in Figure 12d,e.

2.9. Comparative Study of the PS Activation by Iron(II) Sulfate and Its Effect on the Degradation of MG

At 0.2 g L⁻¹ of the catalyst, i.e., iron(II) sulfate in 0.4 mM aqueous solution of MG and 1.6 mM PS, a comparative examination of the PS activation by iron(II) sulfate rather than silver oxide was conducted in room temperature (26 ± 1 °C) at neutral pH. Figure 13a,b show that the rate of MG degradation in the FeSO₄/PS system was lower than that in the Ag₂O/PS system. Ag₂O/PS system completely degraded MG in 36 min, which is about >2 times faster than the FeSO₄/PS system, which demonstrated 99.3% degradation of MG in 79 min. A comparison of the degradation of MG under various experimental circumstances is shown in

Table 1. Comparative analysis of the degradation of MG under different experimental conditions.

Experimental Conditions at Room Temperature	Percent Degradation (Time in Minute)
0.4 mM MG/0.1 g L−1 Ag2O/pH(neutral)/dark	32.5 (180)
0.4 mM MG/0.1 g L−1 Ag2O/pH(neutral)/sunlight	83 (180)
0.4 mM MG/0.1 g L−1 Ag2O/1.6 mM PS/pH(1)/without sunlight	100 (21)
0.4 mM MG/0.1 g L−1 Ag2O/1.6 mM PS/pH(1)/sunlight	100 (15)
0.4 mM MG/0.1 g L−1 Ag2O/1.6 mM PS/pH(neutral)/without sunlight	100 (36)
0.4 mM MG/0.1 g L−1 Ag2O/1.6 mM PS/pH(8)/without sunlight	99.9 (27)
0.4 mM MG/0.2 g L−1 FeSO4/1.6 mM PS/pH(neutral)/withouht sunligh	99.3 (79)

.

2.10. Proposed Degradation Mechanism of MG

According to Equations (3)–(10), the hydroxyl radical produced by the photo-excitation of the electron from the valence band to the conduction band of silver oxide plays a leading role in the degradation of MG dye. This electron reacts with oxygen to form the hydroxyl radical, which degrades MG in an aqueous solution (Equations (5)–(9)). In the absence of sunlight, however, the minor change in the concentration of MG could be due to MG adsorption on the surface of Ag₂O MPs. Another mode of degradation in the presence of sunlight could be the role of MG as a photosensitizer after adsorption on the surface of Ag₂O MPs. The surface-adsorbed MG is excited and donates an electron to the conduction band of Ag₂O MPs by absorbing solar light. These electrons were scavenged by the surface-adsorbed O₂ and reacted with surface-adsorbed H₂O to form H₂O₂, which then produces hydroxyl radicals and degrades the surface-adsorbed MG dye. When a dye molecule degrades, another dye molecule is adsorbed on the surface of Ag₂O MPs, restarting the photocatalytic cycle [47]. Meanwhile, the synergistic effect of persulfate may follow the mechanism of photo-excitation of Ag₂O and generation of SO₄^•− radical according to Equations ((3)–(6), (12)–(14)), which helps to completely degrade the MG dye as seen in the spectrum (Figure 8a). We see no significant peak in the UV-visible region after 15 min. Figure 6a and Figure 8a show a difference in the intensity of the π–π* transition in the MG absorption spectrum because of different degradation mechanisms and scavenging of hydroxyl radicals by the high concentration of protons at pH = 1.

3. Materials and Methods

3.1. Materials

Dried green tea leaves (Camellia sinensis) were purchased from a local market in Peshawar, Pakistan. Analytical grade sodium dodecyl sulfate (SDS) (CH₃(CH₂)₁₁OSO₃Na), sodium persulfate (PS) (Na₂S₂O₈), iron(II) sulfate heptahydrate (FeSO₄∙7H₂O), malachite green oxalate salt (C₂₃H₂₅N₂∙C₂HO₄∙0.5C₂H₂O₄), and silver nitrate (AgNO₃) were purchased from Sigma-Aldrich (St. Louis, MO, USA). The AgNO₃ was used as a pioneer salt for the synthesis of silver oxide (Ag₂O) microparticles, while green tea was used as a reducing and stabilizing agent. The pH of the solution was adjusted by using 0.1 M HNO₃ (Merck, 67% extra pure) and KOH (Sigma-Aldrich). Deionized water was used for the preparation of all solutions. All the chemicals were used as received.

3.2. Green Tea Leaves Processing

About 7.4 g of green tea (GT) leaves powder was introduced into a beaker containing 100 mL deionized water and stirred gently for 20 min at 70 °C. The solution was left for 2 h (at room temperature, i.e., 25 ± 1 °C), and the undissolved matter was separated out. The green tea leaves extract was filtered using Whatman filter paper, and the solution was stored at 5 °C for further use (Figure 14).

3.3. Synthesis of Silver Oxide Microparticles (Ag₂O MPs)

The synthesis of silver oxide microparticles was carried out according to the cited procedure [39]. Briefly, 5 mL of the green tea leaves extract was added to 2 mM (100 mL) AgNO₃ solution after the addition of 0.0015 g SDS. To stabilize the Ag₂O MPs, SDS was added to the reaction mixture. The resulting mixture was stirred for 2.5 h at 40 °C in the dark. The subsequent precipitate was separated by centrifugation at 8000 rpm for 30 min, followed by washing with deionized water. The resulting dark brown pellet of Ag₂O MPs was dried in a petri dish at 90 °C in oven and stored in an air-tight glass vial (Figure 14).

Figure 14. Preparation of green tea leaves extract and synthesis of silver oxide microparticles.

Figure 14. Preparation of green tea leaves extract and synthesis of silver oxide microparticles.

3.4. Characterization Techniques

The optical properties of the synthesized Ag₂O MPs were studied by a UV-visible spectrophotometer (UV-1800, Schimadzu, Kyoto, Japan) in the wavelength range of 400–800 nm. The crystallinity of Ag₂O MPs was determined by X-ray diffraction (XRD) (Malvern Panalytical, Malven, UK). The functional groups present in the Ag₂O MPs were investigated by Fourier transform infrared (FTIR) spectrophotometer (Schimadzu IR Prestige-21, Duisburg, Germany), using molecular and atomic vibration in the 400 to 4000 cm⁻¹ region. The morphology and composition of Ag₂O microparticles were studied by using a scanning electron microscope (SEM) (TESCAN Vega 3, Brno, Czech Republic) and energy dispersive X-ray (EDX) spectroscope (Vario EL CHNS, Elementar Analysensysteme GmbH, Hanau, Germany), respectively.

3.5. Degradation Experiments

The degradation of malachite green was studied in deionized water on a lab scale using a transparent glass vial as a photoreactor and sunlight as a photo/light source. A typical experimental setup includes 10 mL of the dye solution in the transparent glass vial that was loaded with 1 mg of Ag₂O MPs to reveal their catalytic and photocatalytic effect on MG degradation. To investigate the effect of solar light on the degradation of MG without/after introducing Ag₂O MPs, the vial was left in the sun for three hours without being stirred in each case. In the case of the MG degradation study in the dark, the vial was wrapped in aluminum foil without/after introducing Ag₂O MPs and left for 3 h without stirring in both cases. A 3 mL solution was taken at specific time intervals and analyzed by UV/Visible spectrophotometer to record the absorption spectra of MG. Following the spectral analysis, the solution was transferred back to the vial from the cuvette. The typical formulation was used to calculate percent degradation in such a way that the concentration of MG at any specific time point t (C_t) was subtracted from the initial concentration of MG at zero time point (C₀) before introducing solar radiation/catalyst. The result was divided by the initial concentration of MG at zero time point (C₀), and the result was multiplied by 100 to yield percent MG degradation as follows: [(C₀ − C_t)/C₀] × 100. However, the concentration of MG was calculated using the Beer–Lambert law at the selected wavelength of MG, i.e., the maximum wavelength of MG in the visible region (617 nm). The concentration of MG was calculated by dividing the absorbance of MG by the molar absorptivity of MG at 617 nm (2411 M⁻¹ cm⁻¹) in a quartz cuvette with a 1 cm path length. Meanwhile, for kinetic analysis, absorbance was plotted directly without conversion to concentration using the integration method and its formulations for different kinetic orders, as detailed in Section 2. The typical experimental step for the synergistic effect of persulfate ion contained persulfate concentration within the overall 10 mL solution of MG and 1 mg of Ag₂O MPs. The kinetic study was conducted on the spectrophotometer’s spectrum mode with overlay or kinetic mode with a time interval of at least 1 min or more to observe the rate of degradation by measuring changes in absorbance as a function of time.

4. Conclusions

Silver oxide microparticles were successfully synthesized by using green tea leaf extract, which is an environmentally friendly and cost-effective method as compared to conventional chemical and physical methods. The FTIR and UV-visible spectroscopy helped to identify the characteristic peaks of Ag₂O MPs, while the morphology, size, and percent composition of Ag₂O MPs were explained by SEM, XRD, and EDX analysis, respectively. The mean particle size, i.e., 0.40 μm (400 nm), and the crystallite size, i.e., 7.95 nm of the synthesized Ag₂O MPs were determined by SEM and XRD, respectively, indicating approximately 50 times smaller size for the crystallite than the particle. The solar photocatalytic degradation of MG using Ag₂O MPs was 83% during 180 min. Interestingly, Ag₂O/PS system showed excellent efficiency in removing MG from water, represented by the complete removal of MG in 36 min at neutral pH and 21 min at pH = 1 ([PS]₀ = 1.6 mM), with the corresponding rate constant of 0.174 min⁻¹ without introducing sunlight. The degradation efficiency of the Ag₂O/PS system was further increased in the presence of sunlight, represented by complete MG degradation in 15 min with the rate constant of 0.229 min⁻¹ when the pH of the solution was 1 (i.e., 1.3 times rise). The results revealed that solar/Ag₂O, particularly solar/Ag₂O/PS photocatalysis is a promising method for the elimination of toxic organic pollutants, such as malachite green, from the water environment.