*Article* **Adsorption Efficiency and Photocatalytic Activity of Silver Sulfide Nanoparticles Deposited on Carbon Nanotubes**

**Gururaj M. Neelgund 1,\* , Sanjuana Fabiola Aguilar <sup>1</sup> , Erica A. Jimenez <sup>1</sup> and Ram L. Ray <sup>2</sup>**


**\*** Correspondence: gmneelgund@pvamu.edu

**Abstract:** A multimode, dual functional nanomaterial, CNTs-Ag2S, comprised of carbon nanotubes (CNTs) and silver sulfide (Ag2S) nanoparticles, was prepared through the facile hydrothermal process. Before the deposition of Ag2S nanoparticles, hydrophobic CNTs were modified to become hydrophilic through refluxing with a mixture of concentrated nitric and sulfuric acids. The oxidized CNTs were employed to deposit the Ag2S nanoparticles for their efficient immobilization and homogenous distribution. The CNTs-Ag2S could adsorb toxic Cd(II) and completely degrade the hazardous Alizarin yellow R present in water. The adsorption efficiency of CNTs-Ag2S was evaluated by estimating the Cd(II) adsorption at different concentrations and contact times. The CNTs-Ag2S could adsorb Cd(II) entirely within 80 min of the contact time, while CNTs and Ag2S could not pursue it. The Cd(II) adsorption followed the pseudo-second-order, and chemisorption was the rate-determining step in the adsorption process. The Weber−Morris intraparticle pore diffusion model revealed that intraparticle diffusion was not the sole rate-controlling step in the Cd(II) adsorption. Instead, it was contributed by the boundary layer effect. In addition, CNTs-Ag2S could completely degrade alizarin yellow R in water under the illumination of natural sunlight. The Langmuir-Hinshelwood (L-H) model showed that the degradation of alizarin yellow R proceeded with pseudo-first-order kinetics. Overall, CNTs-Ag2S performed as an efficient adsorbent and a competent photocatalyst.

**Keywords:** CNTs; Ag2S; cadmium; adsorption; alizarin yellow R

## **1. Introduction**

Water pollution engrossed by heavy metals is a primary environmental, ecological, and public health concern [1]. Heavy metals are highly soluble in water and are toxic, carcinogenic, and non-degradable [2,3]. Rapid industrialization and technological development have resulted in the discharge of heavy metal-containing effluents into surface and groundwater [4]. The discharged heavy metals are adsorbed by soil and enter the human body through the food chain. Heavy metals accumulate in various organs and body tissues and can cause irreparable damage, including death [5]. Among the heavy metals, cadmium is enormously used and highly toxic [6]. The source for discharging the cadmium into the environment is industrial activities, such as electroplating, smelting, alloy manufacturing, pigments, plastic, battery, fertilizers, pesticides, pigments, dyes, textile operations, and refining [7]. In water, cadmium exists as bivalent, Cd(II), which is responsible for several adverse effects, such as kidney dysfunction, nephritis, hypertension, renal dysfunction, nervous system dysfunction, bone lesions, digestive system dysfunction, cancer, and reproductive organ damage [8–10]. The itai-itai disease, caused by Cd(II)-contaminated water from the Jinzu river in Japan, has instigated severe pain, bone fractures, proteinuria, and osteomalacia [11]. The Cd2+ ions have a high affinity for binding to sulfhydryl (-SH) groups of proteins in biological systems [12]. The presence of Cd(II) in water can lead to severe health and environmental problems. Therefore, its elimination is critically needed. The Cd(II) present in water could be eradicated through chemical precipitation, ion exchange,

**Citation:** Neelgund, G.M.; Aguilar, S.F.; Jimenez, E.A.; Ray, R.L. Adsorption Efficiency and Photocatalytic Activity of Silver Sulfide Nanoparticles Deposited on Carbon Nanotubes. *Catalysts* **2023**, *13*, 476. https://doi.org/10.3390/ catal13030476

Academic Editor: John Vakros

Received: 31 December 2022 Revised: 21 February 2023 Accepted: 21 February 2023 Published: 26 February 2023

**Copyright:** © 2023 by the authors. Licensee MDPI, Basel, Switzerland. This article is an open access article distributed under the terms and conditions of the Creative Commons Attribution (CC BY) license (https:// creativecommons.org/licenses/by/ 4.0/).

solvent extraction, membrane separation, electrochemical removal, coagulation, and adsorption [13–21]. In these techniques, adsorption is superior because of its high efficiency, relative simplicity in design, easy operation, and low operational cost [20,21]. Because of this, many adsorbents have been developed and tested for their efficiency [20–22].

Another class of pollutants that also causes major environmental problems, such as heavy metals, is azo dyes. Azo dyes have attractive colors and are enormously used in industries due to their availability and stability [23]. These dyes contain azo bonds (-N=N-) and substituted aromatic rings [24]. The complex molecular structure of azo dyes has made them recalcitrant and resistant to biodegradation [25]. Azo dyes are mutagenic, teratogenic, and carcinogenic [25]. Moreover, azo dyes can decompose into potentially carcinogenic polychlorinated naphthalenes, benzidine, and amines [26,27]. The toxicity of azo dyes can result in lung cancer, heart diseases, chromosomal aberrations, neurotoxicity, skin disease, and respiratory problems in humans [25]. These dyes can cause disorders of the central nervous system and the inactivation of enzymatic activities [28]. Beyond toxicity, azo dyes are abundantly used in industries because of their cost and advancement in colors that cover the entire spectrum. After their use in industries, about 10-15% of the azo dyes are discharged into the environment through water effluents [29,30]. Releasing azo dyes content water can result in several environmental problems including reduced light penetration in water bodies, which leads to diminished photosynthetic activities, lessened growth and reproduction of aquatic creatures, and aesthetic damage [30,31]. The consequences of the xenobiotic and recalcitrant azo dyes impact the ecosystem's structure and functioning. Considering the adverse effects, it is essential to completely obliterate azo dyes to prevent severe damage and protect health and the environment. Different techniques have been developed to remove the azo dyes in water, viz., coagulation/flocculation, ultrafiltration and membrane processing, chemical precipitation, electrochemical degradation, and ozonation [32–37]. In comparison, photocatalysis is an excellent method for eradicating azo dyes from water in an environmentally friendly approach as the end products of this process are harmless [38,39]. Due to its efficiency in degrading azo dyes, many photocatalysts have been developed to eliminate dyes [37–45]. Among azo dyes, alizarin yellow R (AYR) is a prominently used dye in industries. It is a highly water-soluble anionic dye that contains polycyclic aromatic hydrocarbons [25]. AYR is a derivative of salicylic acid and was prepared in 1887 by Rudolf Nietzki through the reaction of *m*-nitroaniline and salicylic acid [46]. AYR is an industrially important dye and is vastly used in the textile, leather, plastics, paints, and lacquer industries [24]. Furthermore, it is used as an acid-base indicator and employed in histology, stains, and nutrient media preparations [47].

Considering the health and environmental problems associated with Cd(II) and AYR, we designed the dual applicable nanocomposite, CNTs-Ag2S for efficient Cd(II) adsorption and photodegradation of AYR. The multimode CNTs-Ag2S was produced through the facile deposition process of Ag2S nanoparticles over oxidized carbon nanotubes (CNTs) using the hydrothermal method. The adsorption efficiency of CNTs-Ag2S was estimated by evaluating the adsorption rate of Cd(II) from water. The dynamics and controlling mechanisms of Cd(II) adsorption were assessed by pseudo-first- and second-order kinetic models. The reaction pathways and the rate-controlling step underlying Cd(II) adsorption were evaluated using the Weber−Morris intraparticle pore diffusion model. The adsorption equilibrium was determined by fitting the experimental results with Langmuir, Freundlich, and Temkin isotherm models. Furthermore, the catalytic activity of CNTs-Ag2S was determined through the degradation of AYR under exposure to natural sunlight. The photocatalytic activity of CNTs-Ag2S was quantified using the Langmuir-Hinshelwood (L-H) model.

#### **2. Results and Discussion**

The ATR-FTIR spectrum of CNTs-COOH, shown in Figure 1a, demonstrated a band at 3427 cm−<sup>1</sup> corresponding to the O-H bond. The band for the C=O bond of the -COOH groups appeared at 1699 cm−<sup>1</sup> , and the band for the C=C bonds of CNTs was found at

1554 cm−<sup>1</sup> [48–50]. The spectrum of Ag2S (Figure 1b) displayed the peak of the Ag-S bond at 552 cm−<sup>1</sup> [51]. The spectrum of CNTs-Ag2S (Figure 1c) revealed the characteristic absorption bands related to CNTs and Ag2S. The prominent peaks observed at 1066 and 1096 cm−<sup>1</sup> in Figure 1c were attributed to C-O stretching. The peak at 3574 cm−<sup>1</sup> , was due to the −OH stretching vibrations of adsorbed water molecules. The band due to C=O was observed at 1693 cm−<sup>1</sup> , and the peak, at 1517 cm−<sup>1</sup> , was attributed to C-H. The peak related to Ag-S was situated around 550 cm−<sup>1</sup> . The XRD pattern of CNTs-Ag2S (Figure 2) showed a characteristic (0 0 2) reflection of hexagonal graphite of CNTs at 26.1◦ [52]. It revealed the reflections related to Ag2S at 22.7 (−1 0 1), 25.2 (−1 1 1), 26.5 (0 1 2), 29.2 (1 1 1), 31.7 (−1 1 2), 33.8 (1 2 0), 34.6 (−1 2 1), 34.9 (0 2 2), 36.8 (1 1 1), 37.0 (1 2 1), 37.3 (0 1 3), 37.9 (−1 0 3), 40.9 (0 3 1), 43.6 (0 2 3), 44.4 (−1 3 1), 46.4 (−1 2 3), 48.0 (−2 1 2), 48.9 (0 1 4), 53.0 (0 4 0), 53.5 (−2 1 3), 58.3 (−1 4 1), 58.4 (−2 2 3), 60.2 (−1 0 5), 61.4 (0 1 5), 62.8 (2 3 1), 63.4 (2 1 3), 63.9 (−1, 3 4), and 68.0◦ (2 3 2). These values agree with the values found in the standard pattern of acanthite Ag2S (JCPDS file 14-0072) [52]. The TEM images of CNTs-Ag2S presented in Figure 3 explored the deposition of spherical-resembling Ag2S nanoparticles over the surface of tubular-structured CNTs. The proportion of Ag2S nanoparticles was low, possibly due to the low quantity of precursor AgNO<sup>3</sup> used in the preparation compared to the ratio of CNTs. The presence of Ag2S nanoparticles over the surface of CNTs is perceptible in Figure 3a–f. It is noticeable in Figure 3a that the few Ag2S nanoparticles have combined and formed clusters. Figure 3b,d reveal the ruptured surface of some CNTs. It could have happened through the harsh treatment of CNTs with the mixture of concentrated HNO<sup>3</sup> and H2SO4. This process was required to make the CNTs hydrophilic and provide a suitable environment for the effective deposition of Ag2S nanoparticles. However, the entire surface of the CNTs was not distracted and the smooth surface of the CNT is perceptible in Figure 3e,f. Overall, a few defective sites were formed in CNTs. Figure 3d-f reveals the strong adherence of Ag2S nanoparticles to the smoothsurfaced CNT. The absence of leached or free-standing nanoparticles in the void area shows the strong adherence of Ag2S nanoparticles to the surface of CNTs. The size of the Ag2S nanoparticle present in Figure 3d,e, was around 38 and 15 nm, respectively. Therefore, Ag2S nanoparticles have broad size distribution. The average size of Ag2S nanoparticles, calculated from XRD, was around 29 nm, which agrees with the size determined from the TEM images. The core of the Ag2S nanoparticles (Figure 3d) appears denser than the edge. The CNTs have a width of around 110 nm and a length of several micrometers. Overall, CNTs have effectively exfoliated. Further, the EDS of CNTs-Ag2S (Figure S1, Supplementary Materials) demonstrated the presence of O, Ag, and S in CNTs-Ag2S. The peaks for Cu and Ca were found in Figure S1, those are by the copper grid used in the TEM measurement. *Catalysts* **2023**, *13*, 476 4 of 23

**Figure 1.** ATR‐FTIR spectra of (**a**) CNTs‐COOH, (**b**) Ag2S, and (**c**) CNTs‐Ag2S. **Figure 1.** ATR-FTIR spectra of (**a**) CNTs-COOH, (**b**) Ag2S, and (**c**) CNTs-Ag2S.

10 20 30 40 50 60 70 80

**2 (degree)**

0

**Figure 2.** XRD of CNTs‐Ag2S.

5

(-101)

(012)

(111)

(-112)

(-121)

(121) (-103)

(031)

(023)

(-123)

(014)

(-213)

(-223)

(015)

(-134)

(120)

10

15

**Intensity**

20

25

4000 3500 3000 2500 2000 1500 1000 500

(a)

(b)

(c)

**Wavenumber (cm-1)**

**Figure 1.** ATR‐FTIR spectra of (**a**) CNTs‐COOH, (**b**) Ag2S, and (**c**) CNTs‐Ag2S.

**Figure 2.** XRD of CNTs‐Ag2S. **Figure 2.** XRD of CNTs-Ag2S.

0

10

20

30

**Transmittance (%)**

40

50

60

**Figure 3.** (**a–f**) TEM images of CNTs‐Ag2S. **Figure 3.** (**a–f**) TEM images of CNTs-Ag2S.

The TGA pattern of CNTs‐Ag2S (Figure 4) displayed four weight loss steps. The ini‐ tial weight loss of 6.5% that occurred before 100 C was attributed to the desorption of physically adsorbed water molecules over the surface of CNTs‐Ag2S. The following weight loss of 15% ensued within the range of 100‐490 C was due to the detachment and decomposition of oxygen‐containing functional groups existing on the surface of the CNTs. The subsequent weight loss of 10.5% transpired within the range of 500‐540 C, The TGA pattern of CNTs-Ag2S (Figure 4) displayed four weight loss steps. Theinitial weight loss of 6.5% that occurred before 100 ◦C was attributed to the desorptionof physically adsorbed water molecules over the surface of CNTs-Ag2S. The following weight loss of 15% ensued within the range of 100–490 ◦C was due to the detachment and decomposition of oxygen-containing functional groups existing on the surface of the CNTs. The subsequent weight loss of 10.5% transpired within the range of 500–540 ◦C, was owing

was owing to the breaking of the bond between silver and sulfide in Ag2S nanoparticles and releasing sulfur. The successive sharp and significant weight loss of 31.5% taken place

metallic silver and silver oxide. The residual weight remained was about 64%. The broad endothermic peak found in the DTA curve around 80 C was because of the dehydration of the sample. The small endothermic peak found at 170 C accounted for the removal of functional groups from the surface of the CNTs. The intense exothermic peak at 545 C was assigned to release sulfur by breaking the bond between silver and sulfide. The UV‐ vis absorption spectrum of CNTs‐Ag2S, shown in Figure 5, exhibited a characteristic ab‐ sorption band of the C=C bonds of CNTs at 232 nm [52]. In addition, the broad absorption tail corresponding to Ag2S nanoparticles was observed in the visible region [52,53]. If sem‐ iconductor nanoparticles are conjugated to CNTs, it shows the charge‐transfer band [52,54]. However, no such band was observed in Figure 5, illustrating that the conjugation

of Ag2S nanoparticles to CNTs has not altered their energy states [52,54].

to the breaking of the bond between silver and sulfide in Ag2S nanoparticles and releasing sulfur. The successive sharp and significant weight loss of 31.5% taken place within the range of 540–550 ◦C was by the decomposition of sulfur and the formation of metallic silver and silver oxide. The residual weight remained was about 64%. The broad endothermic peak found in the DTA curve around 80 ◦C was because of the dehydration of the sample. The small endothermic peak found at 170 ◦C accounted for the removal of functional groups from the surface of the CNTs. The intense exothermic peak at 545 ◦C was assigned to release sulfur by breaking the bond between silver and sulfide. The UV-vis absorption spectrum of CNTs-Ag2S, shown in Figure 5, exhibited a characteristic absorption band of the C=C bonds of CNTs at 232 nm [52]. In addition, the broad absorption tail corresponding to Ag2S nanoparticles was observed in the visible region [52,53]. If semiconductor nanoparticles are conjugated to CNTs, it shows the charge-transfer band [52,54]. However, no such band was observed in Figure 5, illustrating that the conjugation of Ag2S nanoparticles to CNTs has not altered their energy states [52,54]. *Catalysts* **2023**, *13*, 476 6 of 23 *Catalysts* **2023**, *13*, 476 6 of 23

**Figure 4.** TG/DTA of CNTs‐Ag2S. **Figure 4.** TG/DTA of CNTs-Ag2S. **Figure 4.** TG/DTA of CNTs‐Ag2S.

**Figure 5.** UV‐vis absorption spectrum of CNTs‐Ag2S. **Figure 5.** UV-vis absorption spectrum of CNTs-Ag2S.

**Figure 5.** UV‐vis absorption spectrum of CNTs‐Ag2S.

The XPS survey spectrum of CNTs‐Ag2S, shown in Figure 6a, confirmed the presence

The XPS survey spectrum of CNTs‐Ag2S, shown in Figure 6a, confirmed the presence

at 284.5 and 285.5 eV were assigned to the C‐C bonds of the sp2 and sp3 hybridized carbon atoms of CNTs, respectively [55]. The peaks at 286.4 and 287.8 eV were due to the C‐O and C=O bonds of CNTs, respectively [55]. The deconvoluted O 1s peak (Figure 6c) exhib‐ ited a peak at 531.3 eV for the lattice oxygen and a peak at 532.4 eV corresponding to the carbonyl (=C‐O) functional groups of the CNTs [56]. The core‐electron binding energy of Ag 3d3/2 was found at 373.9 eV, and that of Ag 3d5/2 was at 367.9 eV (Figure 6d). The

at 284.5 and 285.5 eV were assigned to the C‐C bonds of the sp2 and sp3 hybridized carbon atoms of CNTs, respectively [55]. The peaks at 286.4 and 287.8 eV were due to the C‐O and C=O bonds of CNTs, respectively [55]. The deconvoluted O 1s peak (Figure 6c) exhib‐ ited a peak at 531.3 eV for the lattice oxygen and a peak at 532.4 eV corresponding to the carbonyl (=C‐O) functional groups of the CNTs [56]. The core‐electron binding energy of Ag 3d3/2 was found at 373.9 eV, and that of Ag 3d5/2 was at 367.9 eV (Figure 6d). The

The XPS survey spectrum of CNTs-Ag2S, shown in Figure 6a, confirmed the presence of C, O, Ag, and S. The high-resolution spectrum of C1s (Figure 6b) divulged four distinct peaks by Gaussian fitting situated at 284.5, 285.5, 286.4, and 287.8 eV. Among these, peaks at 284.5 and 285.5 eV were assigned to the C-C bonds of the sp<sup>2</sup> and sp<sup>3</sup> hybridized carbon atoms of CNTs, respectively [55]. The peaks at 286.4 and 287.8 eV were due to the C-O and C=O bonds of CNTs, respectively [55]. The deconvoluted O 1s peak (Figure 6c) exhibited a peak at 531.3 eV for the lattice oxygen and a peak at 532.4 eV corresponding to the carbonyl (=C-O) functional groups of the CNTs [56]. The core-electron binding energy of Ag 3d3/2 was found at 373.9 eV, and that of Ag 3d5/2 was at 367.9 eV (Figure 6d). The presence of Ag 3d3/2 and Ag 3d5/2 peaks and their positions confirm the Ag<sup>+</sup> state of silver [57]. The difference in the position of Ag 3d5/2 and Ag 3d3/2 peaks was 6 eV, and no shoulder or satellite peaks were observed between them. The high-resolution S 2p spectrum (Figure 6e) for the spin-orbit splitting of S2+ was deconvoluted into the 2p3/2 level peak at 161.2 eV, and the 2p1/2 peak at 162.3 eV. These peaks occurred by the Ag-S bonds of the Ag2S nanoparticles [58]. The additional peak at 163.9 eV is related to the presence of sulfur in CNTs-Ag2S [58]. *Catalysts* **2023**, *13*, 476 7 of 23 presence of Ag 3d3/2 and Ag 3d5/2 peaks and their positions confirm the Ag+ state of silver [57]. The difference in the position of Ag 3d5/2 and Ag 3d3/2 peaks was 6 eV, and no shoulder or satellite peaks were observed between them. The high‐resolution S 2p spec‐ trum (Figure 6e) for the spin‐orbit splitting of S2+ was deconvoluted into the 2p3/2 level peak at 161.2 eV, and the 2p1/2 peak at 162.3 eV. These peaks occurred by the Ag‐S bonds of the Ag2S nanoparticles [58]. The additional peak at 163.9 eV is related to the presence of sulfur in CNTs‐Ag2S [58].

**Figure 6.** *Cont*.

**Figure 6.** (**a**) XPS survey spectrum of CNTs‐Ag2S. (**b**) High‐resolution spectrum of C1s. (**c**) High‐ resolution spectrum of O1s. (**d**) High‐resolution spectrum of Ag3d. (**e**) High‐resolution spectrum **Figure 6.** (**a**) XPS survey spectrum of CNTs-Ag2S. (**b**) High-resolution spectrum of C1s. (**c**) High-resolution spectrum of O1s. (**d**) High-resolution spectrum of Ag3d. (**e**) High-resolution spectrum of S2p.

The adsorption ability of CNTs‐Ag2S was determined by estimating the Cd(II) ad‐ sorption present in water. The contact time between adsorbent and adsorbate is critical

for contact with CNTs‐Ag2S at different times and compared with that of CNTs and Ag2S (Figure 7). The plot of the adsorption capacity, qt versus t, is presented in Figure S2. Cd(II) adsorption is time dependent and occurs as a gradient function of time. Hence, contact time is crucial in controlling the Cd(II) adsorption. The tendency in Cd(II) adsorption by CNTs, Ag2S, and CNTs‐Ag2S was identical like the adsorption was rapid, then gradually reduced, and finally attained equilibrium. The initial rapid Cd(II) adsorption could have occurred due to the difference in the concentration of Cd2+ ions in the solution and over the surface of the CNTs‐Ag2S, which facilitated the movement of Cd2+ ions from the solu‐ tion to the surface of the CNTs‐Ag2S [8]. Also, in the beginning, many active sites were available on the surface of CNTs‐Ag2S to occupy by Cd2+ ions [8]. With the elapsed contact time, the active sites of the CNTs‐Ag2S were predominantly occupied by Cd2+ ions. Apart, there is a possibility of repulsion between Cd2+ ions in the solution and Cd2+ ions exist over the surface of CNTs‐Ag2S. Owing to these reasons, the adsorption rate could gradually decrease and eventually attain equilibrium [8]. Within 80 min of the contact time, CNTs‐ Ag2S could adsorb the Cd(II) completely; however, underidentical conditions, the adsorp‐ tion rates of CNTs and Ag2S were 79 and 53%,respectively. Thus, the adsorption efficiency of CNTs and Ag2S was significantly improved by their conjugation in CNTs‐Ag2S. Fur‐ ther, to identify the adsorption kinetics of CNTs‐Ag2S, the data were simulated with the

pseudo‐first‐order kinetic model using the expression presented in Equation (1).

(min), respectively, and *k*<sup>1</sup> (min−1) is the pseudo‐first‐order rate constant.

where *qe* and qt are the quantity of adsorbate (mg/g) at equilibrium and particular time, *t*

ሺ െ ௧ሻ *=* െ ଵ*t* (1)

of S2p.

The adsorption ability of CNTs-Ag2S was determined by estimating the Cd(II) adsorption present in water. The contact time between adsorbent and adsorbate is critical and it controls adsorption. Therefore, the Cd(II) adsorption was determined by allowing for contact with CNTs-Ag2S at different times and compared with that of CNTs and Ag2S (Figure 7). The plot of the adsorption capacity, q<sup>t</sup> versus t, is presented in Figure S2. Cd(II) adsorption is time dependent and occurs as a gradient function of time. Hence, contact time is crucial in controlling the Cd(II) adsorption. The tendency in Cd(II) adsorption by CNTs, Ag2S, and CNTs-Ag2S was identical like the adsorption was rapid, then gradually reduced, and finally attained equilibrium. The initial rapid Cd(II) adsorption could have occurred due to the difference in the concentration of Cd2+ ions in the solution and over the surface of the CNTs-Ag2S, which facilitated the movement of Cd2+ ions from the solution to the surface of the CNTs-Ag2S [8]. Also, in the beginning, many active sites were available on the surface of CNTs-Ag2S to occupy by Cd2+ ions [8]. With the elapsed contact time, the active sites of the CNTs-Ag2S were predominantly occupied by Cd2+ ions. Apart, there is a possibility of repulsion between Cd2+ ions in the solution and Cd2+ ions exist over the surface of CNTs-Ag2S. Owing to these reasons, the adsorption rate could gradually decrease and eventually attain equilibrium [8]. Within 80 min of the contact time, CNTs-Ag2S could adsorb the Cd(II) completely; however, under identical conditions, the adsorption rates of CNTs and Ag2S were 79 and 53%, respectively. Thus, the adsorption efficiency of CNTs and Ag2S was significantly improved by their conjugation in CNTs-Ag2S. Further, to identify the adsorption kinetics of CNTs-Ag2S, the data were simulated with the pseudo-first-order kinetic model using the expression presented in Equation (1).

$$
\ln(q\_\varepsilon - q\_t) = \ln q\_\varepsilon - k\_1 t \tag{1}
$$

where *q<sup>e</sup>* and *q<sup>t</sup>* are the quantity of adsorbate (mg/g) at equilibrium and particular time, *t* (min), respectively, and *k*<sup>1</sup> (min−<sup>1</sup> ) is the pseudo-first-order rate constant. *Catalysts* **2023**, *13*, 476 10 of 23

**Figure 7.** Adsorption kinetics of Cd(II) over CNTs, Ag2S, and CNTs‐Ag2S. **Figure 7.** Adsorption kinetics of Cd(II) over CNTs, Ag2S, and CNTs-Ag2S.

The pseudo‐first‐order plot obtained for Cd(II) adsorption over CNTs, Ag2S, and CNTs‐Ag2S is shown in Figure 8. The correlation coefficients (R2) perceived for CNTs, The pseudo-first-order plot obtained for Cd(II) adsorption over CNTs, Ag2S, and CNTs-Ag2S is shown in Figure 8. The correlation coefficients (R<sup>2</sup> ) perceived for CNTs, Ag2S, and CNTs-Ag2S were 0.9906, 0.9540, and 0.9930, respectively. The k<sup>1</sup> and q<sup>e</sup> were computed

Ag2S, and CNTs‐Ag2S were 0.9906, 0.9540, and 0.9930, respectively. The k1 and qe were computed using the slope and intercept values of the straight lines acquired in Figure 8.

for the Cd(II) adsorption could not be projected by the pseudo‐first‐order kinetic model. Alternatively, it was analyzed using the pseudo‐second‐order kinetic model using Equa‐

<sup>మ</sup> +

௧ 

(2)

௧ 

where k2 [g/(mg.min)] is the pseudo‐second‐order rate constant.

ൌ <sup>ଵ</sup> మ

tion (2).

using the slope and intercept values of the straight lines acquired in Figure 8. The estimated values of k<sup>1</sup> and q<sup>e</sup> is summarized in Table 1. The calculated qe(cal) could not agree with the experimental value qe(exp) for all samples. Thus, the experimental data for the Cd(II) adsorption could not be projected by the pseudo-first-order kinetic model. Alternatively, it was analyzed using the pseudo-second-order kinetic model using Equation (2).

$$\frac{t}{q\_t} = \frac{1}{k\_2 \ q\_\varepsilon^2} + \frac{t}{q\_\varepsilon} \tag{2}$$

where *k<sup>2</sup>* [g/(mg.min)] is the pseudo-second-order rate constant.

**Figure 8.** Pseudo‐first‐order kinetics for Cd(II) adsorption over CNTs, Ag2S, and CNTs‐Ag2S. **Figure 8.** Pseudo-first-order kinetics for Cd(II) adsorption over CNTs, Ag2S, and CNTs-Ag2S.



CNTs 98.75 4.7785 0.0608 0.9906 109.8 0.8360 0.9963 Ag2S 67.50 4.0970 0.0519 0.9540 72.62 1.7750 0.9985 The pseudo‐second‐order model plot is illustrated in Figure 9, and the determined parameters are presented in Table 1. The R2 values derived from the pseudo‐second‐order plots for CNTs, Ag2S, and CNTs‐Ag2S were 0.9963, 0.9985, and 0.9980, respectively. The R2 value of the second‐order plot was relatively higher than that found for the first‐order plot, which was close to 1. The qe(cal) of the second‐order model matched the qe(exp) for all samples. The linearity of the pseudo‐second‐order plot and the agreeing values of qe(exp) and qe(cal) show that the Cd(II) adsorption could be analyzed using the pseudo‐ second‐order kinetics rather than the pseudo‐first‐order kinetics. The agreement of the The pseudo-second-order model plot is illustrated in Figure 9, and the determined parameters are presented in Table 1. The R<sup>2</sup> values derived from the pseudo-secondorder plots for CNTs, Ag2S, and CNTs-Ag2S were 0.9963, 0.9985, and 0.9980, respectively. The R<sup>2</sup> value of the second-order plot was relatively higher than that found for the firstorder plot, which was close to 1. The qe(cal) of the second-order model matched the qe(exp) for all samples. The linearity of the pseudo-second-order plot and the agreeing values of qe(exp) and qe(cal) show that the Cd(II) adsorption could be analyzed using the pseudo-second-order kinetics rather than the pseudo-first-order kinetics. The agreement of the second-order kinetics demonstrates that chemisorption was the rate-determining step in the Cd(II) adsorption [9]. Moreover, it represents the rapid transfer of Cd(II) from a solution with a lower initial concentration to the surface of the CNTs-Ag2S due to

second‐order kinetics demonstrates that chemisorption was the rate‐determining step in

tration gradient [9]. Further, the experimental data were explored using the Weber−Morris intraparticle pore diffusion model to evaluate the diffusion mechanism using Equation

where kid (mg/g.min) is the intraparticle diffusion rate constant, which can be calculated from the linear plot of qt versus t0.5. c (mg/g) is the intraparticle diffusion constant, esti‐ mated from the intercept and directly proportional to the boundary layer thickness. It is assumed that the higher the intercept value, the more significant the contribution of the

௧ *=* ௗ .<sup>ହ</sup> *+ c* (3)

(3).

the concentration gradient [9]. Further, the experimental data were explored using the Weber−Morris intraparticle pore diffusion model to evaluate the diffusion mechanism using Equation (3).

$$q\_l = k\_{il}t^{0.5} + \mathcal{C} \tag{3}$$

where *kid* (mg/g.min) is the intraparticle diffusion rate constant, which can be calculated from the linear plot of *q<sup>t</sup>* versus *t* 0.5. c (mg/g) is the intraparticle diffusion constant, estimated from the intercept and directly proportional to the boundary layer thickness. It is assumed that the higher the intercept value, the more significant the contribution of the surface adsorption in the rate-controlling step. If the regression of the *q<sup>t</sup>* versus *t* 0.5 plot is linear and passes through the origin, intraparticle diffusion plays a significant role in controlling the kinetics of the adsorption process. If it does not pass through the origin, intraparticle diffusion is not the only rate-limiting step. Instead, it is contributed by the boundary layer effect [59]. The intraparticle diffusion plot of *q<sup>t</sup>* versus *t* 0.5 acquired for the Cd(II) adsorption on CNTs-Ag2S is demonstrated in Figure 10. The plot in Figure 10 does not pass through the origin of the coordinates (Figure S3). So, intraparticle diffusion is not the sole rate-limiting step in Cd(II) adsorption over CNTs-Ag2S. The association of the two straight lines in the intraparticle diffusion plot (Figure S3) and the intercept values (Table S1) support that the intraparticle diffusion was not only a rate-controlling step in the Cd(II) adsorption; it also contributed through the boundary layer effect [59]. The multilinearity of the plot in Figure 10 (Figure S3) shows that the Cd(II) adsorption transpires through multiple phases [60]. Among these, the initial stage occurred through the rapid adsorption of Cd2+ ions, the second phase was owing to the diffusion of Cd2+ ions into the pores of CNTs-Ag2S, and the third phase was due to equilibrium of the adsorption that caused a chemical reaction/bonding [60]. surface adsorption in the rate‐controlling step. If the regression of the qt versus t0.5 plot is linear and passes through the origin, intraparticle diffusion plays a significant role in con‐ trolling the kinetics of the adsorption process. If it does not pass through the origin, intra‐ particle diffusion is not the only rate‐limiting step. Instead, it is contributed by the bound‐ ary layer effect [59]. The intraparticle diffusion plot of qt versus t0.5 acquired for the Cd(II) adsorption on CNTs‐Ag2S is demonstrated in Figure 10. The plot in Figure 10 does not pass through the origin of the coordinates (Figure S3). So, intraparticle diffusion is not the sole rate‐limiting step in Cd(II) adsorption over CNTs‐Ag2S. The association of the two straight lines in the intraparticle diffusion plot (Figure S3) and the intercept values (Table S1) support that the intraparticle diffusion was not only a rate‐controlling step in the Cd(II) adsorption; it also contributed through the boundary layer effect [59]. The multi‐ linearity of the plot in Figure 10 (Figure S3) shows that the Cd(II) adsorption transpires through multiple phases [60]. Among these, the initial stage occurred through the rapid adsorption of Cd2+ ions, the second phase was owing to the diffusion of Cd2+ ions into the pores of CNTs‐Ag2S, and the third phase was due to equilibrium of the adsorption that caused a chemical reaction/bonding [60].

**Figure 9.** Pseudo‐second‐order kinetics for Cd(II) adsorption over CNTs, Ag2S, and CNTs‐Ag2S. **Figure 9.** Pseudo-second-order kinetics for Cd(II) adsorption over CNTs, Ag2S, and CNTs-Ag2S.

**Figure 10.** Weber‐Morris intraparticle diffusion plot for the Cd(II) adsorption over CNTs‐Ag2S. **Figure 10.** Weber-Morris intraparticle diffusion plot for the Cd(II) adsorption over CNTs-Ag2S.

The experimental equilibrium parameters for Cd(II) adsorption were determined by applying three isotherm models: Langmuir, Freundlich, and Temkin. The Langmuir model predicts that the adsorbed molecules form a monolayer and adsorption emerges at a static number of adsorption sites, each of which is equivalent in efficiency. Moreover, each molecule has a constant enthalpy and adsorption activation energy, it means that all molecules have an affinity equal to entire adsorption sites [61]. With the help of Langmuir isotherm model, it could be find the value of the maximum adsorption capacity of the adsorbent using Equation (4) [62]: ൌ 1 The experimental equilibrium parameters for Cd(II) adsorption were determined by applying three isotherm models: Langmuir, Freundlich, and Temkin. The Langmuir model predicts that the adsorbed molecules form a monolayer and adsorption emerges at a static number of adsorption sites, each of which is equivalent in efficiency. Moreover, each molecule has a constant enthalpy and adsorption activation energy, it means that all molecules have an affinity equal to entire adsorption sites [61]. With the help of Langmuir isotherm model, it could be find the value of the maximum adsorption capacity of the adsorbent using Equation (4) [62]:

$$\frac{\mathbf{C}\_{\varepsilon}}{q\_{\varepsilon}} = \frac{\mathbf{C}\_{\varepsilon}}{q\_{m}} + \frac{1}{K\_{L}}q\_{m} \tag{4}$$

is the concentration of Cd(II) at equilibrium; qm is the maximum amount of the Cd(II) adsorbed per unit mass of CNTs‐Ag2S to form a complete monolayer on the surface‐bound at high Ce. KL is the Langmuir adsorption constant related to the free energy of adsorption. The Langmuir plot for Cd(II) adsorption over CNTs‐Ag2S is presented in Figure 11. The estimated parameters are shown in Table 2. The maximum adsorption capacity (qm) cal‐ culated for Cd(II) adsorption over CNTs‐Ag2S was 256.4 mg/g. where *q<sup>e</sup>* (mg/g) is the amount of adsorbed Cd(II) per unit mass of CNTs-Ag2S; *C<sup>e</sup>* (mg/L) is the concentration of Cd(II) at equilibrium; *q<sup>m</sup>* is the maximum amount of the Cd(II) adsorbed per unit mass of CNTs-Ag2S to form a complete monolayer on the surface-bound at high *C<sup>e</sup>* . *K<sup>L</sup>* is the Langmuir adsorption constant related to the free energy of adsorption. The Langmuir plot for Cd(II) adsorption over CNTs-Ag2S is presented in Figure 11. The estimated parameters are shown in Table 2. The maximum adsorption capacity (qm) calculated for Cd(II) adsorption over CNTs-Ag2S was 256.4 mg/g. *Catalysts* **2023**, *13*, 476 14 of 23

**Figure 11.** Langmuir isotherm plot for Cd(II) adsorption over CNTs‐Ag2S. **Figure 11.** Langmuir isotherm plot for Cd(II) adsorption over CNTs-Ag2S.

**Table 2.** Parameters calculated for Cd(II) adsorption over CNTs‐Ag2S using Langmuir, Freundlich,

256.4 −0.6830 0.9972 305.0 −31.21 0.3562 1.1 × 10256 0.4565 0.0019

The Freundlich isotherm is developed based on the assumption that the adsorption sites are distributed exponentially concerning the heat of adsorption [62,63]. It provides the relationship between the equilibrium of liquid and solid phase capacities with multi‐ layer adsorption properties consisting of the heterogeneous surface of the adsorbent. The Freundlich isotherm that supports multilayer adsorption agrees with the Langmuir model over moderate ranges of concentrations but differs at low and high concentrations. The

**<sup>g</sup>−1) <sup>n</sup> R2Fre <sup>A</sup> <sup>B</sup> R2Tem**

 

ൌ (6)

(5)

linear form of the Freundlich isotherm can be represented b Equation (5):

ൌ ி

where qe (mg/g) is the amount of Cd(II) adsorbed at equilibrium. KF and n are Freundlich constants. KF symbolizes the affinity of the adsorbent, and n signifies the adsorption in‐ tensity. Ce (mg/L) is the Cd(II) concentration at equilibrium. The Freundlich isotherm plot attained for Cd(II) adsorption is shown in Figure S4, and the calculated parameters are

The Temkin isotherm model is based on the fact that, during adsorption, the heat of all molecules decreases linearly with an increase in the coverage of the adsorbent surface and that adsorption is characterized by the uniform distribution of binding energies up to the maximum binding energy [62,64]. The linear form of the Temkin isotherm model is

and Temkin adsorption isotherm models.

**KL (L mg−1) R2Lan**

**qm (mg g−1)**

included in Table 2.

shown in Equation (6):


**Table 2.** Parameters calculated for Cd(II) adsorption over CNTs-Ag2S using Langmuir, Freundlich, and Temkin adsorption isotherm models.

The Freundlich isotherm is developed based on the assumption that the adsorption sites are distributed exponentially concerning the heat of adsorption [62,63]. It provides the relationship between the equilibrium of liquid and solid phase capacities with multilayer adsorption properties consisting of the heterogeneous surface of the adsorbent. The Freundlich isotherm that supports multilayer adsorption agrees with the Langmuir model over moderate ranges of concentrations but differs at low and high concentrations. The linear form of the Freundlich isotherm can be represented b Equation (5):

$$\ln q\_{\varepsilon} = \ln \mathcal{K}\_{\mathcal{F}} + \frac{\ln \mathcal{C}\_{\varepsilon}}{n} \tag{5}$$

where *q<sup>e</sup>* (mg/g) is the amount of Cd(II) adsorbed at equilibrium. *K<sup>F</sup>* and n are Freundlich constants. *K<sup>F</sup>* symbolizes the affinity of the adsorbent, and n signifies the adsorption intensity. *C<sup>e</sup>* (mg/L) is the Cd(II) concentration at equilibrium. The Freundlich isotherm plot attained for Cd(II) adsorption is shown in Figure S4, and the calculated parameters are included in Table 2.

The Temkin isotherm model is based on the fact that, during adsorption, the heat of all molecules decreases linearly with an increase in the coverage of the adsorbent surface and that adsorption is characterized by the uniform distribution of binding energies up to the maximum binding energy [62,64]. The linear form of the Temkin isotherm model is shown in Equation (6):

$$q\_{\varepsilon} = B \ln A + B \ln \mathbb{C}\_{\varepsilon} \tag{6}$$

where B = RT/KT, K<sup>T</sup> is the Temkin constant related to the heat of adsorption (J/mol); A is the Temkin isotherm constant (L/g), R is the gas constant (8.314 J/mol K), and T is the absolute temperature (K). The Temkin isotherm fitting plot for the Cd(II) adsorption is shown in Figure S5. The values estimated from Figure S5 are illustrated in Table 2.

The R<sup>2</sup> Lan value found for the Langmuir isotherm plot was 0.9972; for Freundlich and Temkin isotherm plots, the R<sup>2</sup> Fre and R<sup>2</sup> Tem values were 0.3562 and 0.0019, respectively. The Freundlich and Temkin isotherm models provided scattered points and failed to accomplish linearity. Therefore, the Freundlich and Temkin isotherm models were unsuitable for explaining Cd(II) adsorption over CNTs-Ag2S. However, the Langmuir isotherm model produced linearity. Therefore, the Langmuir model is appropriate for elucidating the Cd(II) adsorption. The validation of the Langmuir model indicates that the Cd(II) adsorption over CNTs-Ag2S ensued with monolayer molecular covering and chemisorption, together [65].

Further, the photocatalytic activity of CNTs-Ag2S was investigated through the degradation of AYR under exposure to sunlight. The degradation of AYR by CNTs-Ag2S is presented in Figure 12. The activity of CNTs-Ag2S was compared with that of CNTs and Ag2S. The CNTs-Ag2S could degrade the AYR completely within 120 min of illumination. However, CNTs and Ag2S were capable to degrade 77.3 and 41% of AYR, respectively, in 120 min. Hence, the conjugation of CNTs and Ag2S was substantially improved the photocatalytic activity in CNTs-Ag2S. The degradation of AYR was quantified by the reduction in the electronic absorption band located at 373 nm (Figure 13). The hybrid nano-architecture of CNTs-Ag2S was proficient in adsorbing higher number of AYR molecules and capable in absorption of sunlight. In addition, the effectual hindrance of recombining of electrons and holes during photocatalysis facilitated the degradation process. The photocatalytic activity of CNTs-Ag2S was further explored by finding the apparent rate constants using

the Langmuir-Hinshelwood (L-H) model and with the help of Equation (7) that could be applicable for low concentrations of dyes [38,39,44,45]. stants using the Langmuir‐Hinshelwood (L‐H) model and with the help of Equation (7) that could be applicable for low concentrations of dyes [38,39,44,45].

where B = RT/KT, KT is the Temkin constant related to the heat of adsorption (J/mol); A is the Temkin isotherm constant (L/g), R is the gas constant (8.314 J/mol K), and T is the absolute temperature (K). The Temkin isotherm fitting plot for the Cd(II) adsorption is shown in Figure S5. The values estimated from Figure S5 are illustrated in Table 2.

The R2Lan value found for the Langmuir isotherm plot was 0.9972; for Freundlich and Temkin isotherm plots, the R2Fre and R2Tem values were 0.3562 and 0.0019, respectively. The Freundlich and Temkin isotherm models provided scattered points and failed to accom‐ plish linearity. Therefore, the Freundlich and Temkin isotherm models were unsuitable for explaining Cd(II) adsorption over CNTs‐Ag2S. However, the Langmuir isotherm model produced linearity. Therefore, the Langmuir model is appropriate for elucidating the Cd(II) adsorption. The validation of the Langmuir model indicates that the Cd(II) ad‐ sorption over CNTs‐Ag2S ensued with monolayer molecular covering and chemisorption,

Further, the photocatalytic activity of CNTs‐Ag2S was investigated through the deg‐ radation of AYR under exposure to sunlight. The degradation of AYR by CNTs‐Ag2S is presented in Figure 12. The activity of CNTs‐Ag2S was compared with that of CNTs and Ag2S. The CNTs‐Ag2S could degrade the AYR completely within 120 min of illumination. However, CNTs and Ag2S were capable to degrade 77.3 and 41% of AYR, respectively, in 120 min. Hence, the conjugation of CNTs and Ag2S was substantially improved the pho‐ tocatalytic activity in CNTs‐Ag2S. The degradation of AYR was quantified by the reduc‐ tion in the electronic absorption band located at 373 nm (Figure 13). The hybrid nano‐ architecture of CNTs‐Ag2S was proficient in adsorbing higher number of AYR molecules and capable in absorption of sunlight. In addition, the effectual hindrance of recombining of electrons and holes during photocatalysis facilitated the degradation process. The pho‐ tocatalytic activity of CNTs‐Ag2S was further explored by finding the apparent rate con‐

$$\ln \frac{\mathcal{C}\_0}{\mathcal{C}} = k\_{app} \, t \,\tag{7}$$

where *C<sup>0</sup>* is the initial concentration of AYR and *C* is the concentration at a particular time of irradiation. *kapp* is the apparent rate constant of the reaction, and t is the irradiation time. where C0 is the initial concentration of AYR and C is the concentration at a particular time of irradiation. kapp is the apparent rate constant of the reaction, and t is the irradiation time.

*Catalysts* **2023**, *13*, 476 15 of 23

together [65].

**Figure 12.** The degradation of alizarin yellow R in the presence of CNTs, Ag2S, and CNTs‐Ag2S under the illumination to sunlight. **Figure 12.** The degradation of alizarin yellow R in the presence of CNTs, Ag2S, and CNTs-Ag2S under the illumination to sunlight. *Catalysts* **2023**, *13*, 476 16 of 23

**Figure 13.** Reduction in the electronic absorption of alizarin yellow R in the presence of CNTs‐Ag2S under the irradiation to sunlight. **Figure 13.** Reduction in the electronic absorption of alizarin yellow R in the presence of CNTs-Ag2S under the irradiation to sunlight.

The L‐H plots perceived for the degradation of AYR were linear, as depicted in Fig‐ ure 14. The linearity of the L‐H plots reveals that the degradation of AYR ensue with pseudo‐first‐order kinetics. The apparent rate constants determined by the L‐H plots for CNTs, Ag2S, and CNTs‐Ag2S were 0.0108, 0.0035, and 0.0378 min‐1, respectively. The value determined for CNTs‐Ag2S was about four‐fold higher than CNTs and eleven‐fold greater The L-H plots perceived for the degradation of AYR were linear, as depicted in Figure 14. The linearity of the L-H plots reveals that the degradation of AYR ensue with pseudofirst-order kinetics. The apparent rate constants determined by the L-H plots for CNTs, Ag2S, and CNTs-Ag2S were 0.0108, 0.0035, and 0.0378 min−<sup>1</sup> , respectively. The value determined for CNTs-Ag2S was about four-fold higher than CNTs and eleven-fold greater

than Ag2S. Therefore, the conjugation of CNTs and Ag2S has magnificently improved the photocatalytic activity in CNTs‐Ag2S. Due to the conjugated structure, containing aro‐

Accordingly, AYR degradation could happen by breaking the conjugated system in the presence of CNTs‐Ag2S under sunlight irradiation. With this assumption, the possible mechanism for the rapid degradation of AYR by CNTs‐Ag2S could be explained as follows

> CNTs‐Ag2S + h e<sup>−</sup> + h+ h+ + H2O HO<sup>−</sup> + H+ e<sup>−</sup> + O2 O2

> > <sup>−</sup> + H+ HO2

O2

−

<break>AYR + HO H2O + CO2 + nontoxic products

(Figure 15).

HO2

+ H2O H2O2 + HO

than Ag2S. Therefore, the conjugation of CNTs and Ag2S has magnificently improved the photocatalytic activity in CNTs-Ag2S. Due to the conjugated structure, containing aromatic, carbonyl, and azo groups, the AYR solution possessed an intense color in water. Accordingly, AYR degradation could happen by breaking the conjugated system in the presence of CNTs-Ag2S under sunlight irradiation. With this assumption, the possible mechanism for the rapid degradation of AYR by CNTs-Ag2S could be explained as follows (Figure 15).

CNTs-Ag2S + hν → e <sup>−</sup> + h<sup>+</sup> h <sup>+</sup> + H2<sup>O</sup> <sup>→</sup> HO<sup>−</sup> + H<sup>+</sup> e <sup>−</sup> + O<sup>2</sup> → O<sup>2</sup> •− O<sup>2</sup> •− + H<sup>+</sup> <sup>→</sup>HO<sup>2</sup> • HO<sup>2</sup> • + H2O → H2O<sup>2</sup> + HO• AYR + HO• → H2O + CO<sup>2</sup> + nontoxic products *Catalysts* **2023**, *13*, 476 17 of 23 *Catalysts* **2023**, *13*, 476 17 of 23 5 Ag2S

**Figure 14.** Langmuir‐Hinshelwood plot for the degradation of alizarin yellow R in the presence of CNTs‐Ag2S under illumination to sunlight. **Figure 14.** Langmuir-Hinshelwood plot for the degradation of alizarin yellow R in the presence of CNTs-Ag2S under illumination to sunlight. **Figure 14.** Langmuir‐Hinshelwood plot for the degradation of alizarin yellow R in the presence of CNTs‐Ag2S under illumination to sunlight.

Further, to find the stability of CNTs‐Ag2S, it was recovered by centrifugation after photocatalysis. It was washed with DI water, dried and used in further two cycles of the **Figure 15.** A possible mechanism for the degradation of alizarin yellow R in the presence of CNTs‐ Ag2S under exposure to sunlight. **Figure 15.** A possible mechanism for the degradation of alizarin yellow R in the presence of CNTs-Ag2S under exposure to sunlight.

photocatalysis. The reused CNTs‐Ag2S was able to degrade AYR completely in further two cycles without any significant reduction in activity (Figure S6). The XRD of CNTs‐ Ag2S recorded before and after using in the three degradation cycles of AYR could not show any structural modification or rapture (Figure S7). Consequently, CNTs‐Ag2S is ta‐

Further, to find the stability of CNTs‐Ag2S, it was recovered by centrifugation after

The CNT/Ag2S nanocomposite prepared by Di et al. and applied it to the degradation of rhodamine B revealed its higher activity than bare Ag2S nanoparticles under illumina‐ tion to visible and near‐infrared (NIR) light [57]. In addition, the recycled CNT/Ag2S nano‐ composite could not lost the activity [57]. The Ag2S‐CNT nanocomposite, reported by

photocatalysis. The reused CNTs‐Ag2S was able to degrade AYR completely in further two cycles without any significant reduction in activity (Figure S6). The XRD of CNTs‐

show any structural modification or rapture (Figure S7). Consequently, CNTs‐Ag2S is ta‐

The CNT/Ag2S nanocomposite prepared by Di et al. and applied it to the degradation of rhodamine B revealed its higher activity than bare Ag2S nanoparticles under illumina‐ tion to visible and near‐infrared (NIR) light [57]. In addition, the recycled CNT/Ag2S nano‐ composite could not lost the activity [57]. The Ag2S‐CNT nanocomposite, reported by

Ag2S under exposure to sunlight.

ble and suitable for reuse.

ble and suitable for reuse.

Further, to find the stability of CNTs-Ag2S, it was recovered by centrifugation after photocatalysis. It was washed with DI water, dried and used in further two cycles of the photocatalysis. The reused CNTs-Ag2S was able to degrade AYR completely in further two cycles without any significant reduction in activity (Figure S6). The XRD of CNTs-Ag2S recorded before and after using in the three degradation cycles of AYR could not show any structural modification or rapture (Figure S7). Consequently, CNTs-Ag2S is table and suitable for reuse.

The CNT/Ag2S nanocomposite prepared by Di et al. and applied it to the degradation of rhodamine B revealed its higher activity than bare Ag2S nanoparticles under illumination to visible and near-infrared (NIR) light [57]. In addition, the recycled CNT/Ag2S nanocomposite could not lost the activity [57]. The Ag2S-CNT nanocomposite, reported by Meng et al. [66], efficiently degraded texbrite BA-L in presence of visible light. In this study, the Ag2S-CNT nanocomposite was for four degradation cycles of texbrite BA-L [66]. Further, the photocatalytic activity of CNTs-Ag2S in the degradation of AYR was compared with different photocatalysts and presented in Table 3 [24,25,46,67–69].


**Table 3.** Comparison of the degradation rate of AYR by CNTs-Ag2S with reported photocatalysts.

#### **3. Experimental**

*3.1. Materials*

The chemicals and CNTs were purchased from Millipore Sigma and used as received. The aqueous solutions were prepared using ultrapure water obtained by the Milli-Q Plus system (Millipore; Burlington, MA, USA).
