Effective Removal of Cu(II) Ions from Aqueous Solution by Cross-Linked Chitosan-Based Hydrogels

Abstract

The elimination of metal ions from industrial waste water is one of the most significant environmental needs. For the first time, two chitosan hydrogels that we had previously synthesized, cross-linked with varying concentrations of trimellitic anhydride isothiocyanate (represented by H₁ and H₂), were utilized in this investigation to adsorb Cu(II) ions. We found that pH 6, 25 °C, 200 mg L⁻¹ of Cu(II) ions concentration, and 15 mg of hydrogel dosage were the ideal parameters for Cu(II) ion elimination. The kinetics of their adsorption fitted to the pseudo-second-order model with the highest correlation coefficient (R²) values equal to 0.999 and 1.00 for H₁ and H₂, respectively. The experimental q_e values were found when H₁ was equal to 97.59 mg g⁻¹ (theoretical value is equal to 98.04 mg g⁻¹) and H₂ was equal to 96.20 mg g⁻¹ (theoretical value is equal 99.01 mg g⁻¹). The hydrogels achieved a removal effectiveness of 97.59% and their adsorption isotherms matched the Freundlich model, indicating the multi-layered and homogeneous adsorption nature. The removal of copper ions is significantly driven by the physisorption phenomenon. The hydrogels have a great possibility to be utilized as promising, efficacious, reusable adsorbents for industrial wastewater remediation. Thus, incorporation of a cross-linker, containing binding centers for Cu(II) ions, between chitosan chains is a good way to obtain suitable efficient adsorbents which are good choices for application in the field of metal elimination.

1. Introduction

The issue of heavy metal pollution is very significant for the ecological systems all over the world, particularly aquatic resources, due to the fact that they are not biodegradable. They tend to accumulate in the biological systems of plants and aquatic organisms, resulting in a high risk to living organisms as they are transferred to them through the food chain [1]. Industrial effluents and waste can introduce heavy metals into ecosystems via some industries such as the fertilizer industry, electroplating industry, printing industry, and steel industry, and via some activities such as mining processes, dyes production, and synthesis of disease control agents used for plants, in addition to soil erosion, and natural weathering of the earth’s crust [2]. Several methods have been used to remove or at least minimize the percentage of heavy metal ions in their aqueous solutions, including chemical precipitation [3], ion exchange [4], activated carbon adsorption [5], and membrane filtration [6]. All of these techniques are efficient, but they have certain drawbacks and restrictions, such as expensive costs, high energy consumption, hazardous byproducts from the reagents used, and poor performance at low metal concentrations. An alternative approach is the biosorption process and this approach has received great attention from researchers. Its main characteristics are high efficiency, simplicity, ease of use, short running time, low running costs, no undesired secondary product, environmentally friendliness, abundance of adsorbents, and the simple restoration of adsorbents for future usage. The biosorption process involves an integration of active and passive transport processes. It is the start of a quick and reversible process of accumulation. The second stage, usually called active uptake, is a slower, often irreversible intracellular bioaccumulation linked to metabolic activity [7]. The majority of metals in the periodic table’s fourth period are carcinogenic. The electronic structure of the transition metals is thought to be responsible for this carcinogenicity [8]. Because copper is a metal that is essential in a number of enzymes for all forms of life, problems occur when there is a lack or an excess of it [9]. One of the heavy metals that is most frequently used is copper. Its cupric ion, Cu(II), is the most abundant form in the environment and is toxic to many organisms [10,11,12]. Human beings are being exposed to unusually high levels of toxic copper in their drinking water and food due to certain practices such as cooking in copper-lined or copper-glazed pots and the use of copper pipes for water supply [13]. Moreover, sewage effluent that contains wastes produced from both the fertilizer and the mining industries, as well as wastes resulted from the plating, paints, and pigments baths are other sources of copper [14]. Although a trace amount of copper is essential for living organisms, harmful effects such as neurotoxicity, jaundice, diarrhea, respiratory problems, liver and kidney failure, and even death can result from an excessive amount of copper in water [15,16,17]. The absorption of excessive amounts of copper by the human body leads to severe irritation and corrosion of the mucous membranes of the gastrointestinal tract; widespread capillary, liver and kidneys damage; and irritation of the central nervous system followed by depression [18,19]. A high incidence of cancer among coppersmiths is one of the epidemiological pieces of evidence, suggesting the primary role of copper as a carcinogen [20]. According to the US Environmental Protection Agency (EPA) and the WHO, the acceptable limit of copper in drinking water is 1.3 and 2.0 mg/L, respectively [21]. The biopolymer chitosan has gained great importance in the field of biotechnology in relation to the environment because of its antimicrobial activity [22,23], its metal adsorption ability [24,25], and its capacity to adsorb anionic dyes [26,27]. Chitosan is also abundant, hydrophilic, biocompatible, biodegradable, renewable, and non-toxic [28].

One of chitosan’s key drawbacks is thought to be its easy solubility in acidic conditions, particularly because industrial wastewater is often acidic and is employed as an adsorbent material for pollutants removal. Controlling chitosan’s solubility is therefore crucial for these kinds of applications. This is achieved via graft copolymerization [29], polymer blending [30,31] or chemical substitution [32]. Chemical cross-linking of chitosan is used as a promising alternative to improve its mechanical properties, increase its chemical stability by reducing its solubility in acidic solutions, retard its rate of degradation, and extend its product life in different media [33,34,35].

Chitosan-grafted polyacrylonitrile was used for removal of copper. The adsorption of copper obeyed a pseudo-second-order model and the Langmuir model. The optimum conditions were pH = 7.5, the amount of adsorbent = 5 g, and equilibrium time = 5–6 h. The highest adsorption capacity was 239.314 mg g⁻¹ [36]. Chitosan-coated MnFe₂O₄ nanoparticles were utilized to remove a low concentration of toxic Cu(II) ions with a high capacity of adsorption, reaching to 22.60 mg g⁻¹ [37]. The highest amount of Cu (II) ions that could be absorbed by chitosan-coated cotton gauze was 14.40 mg g⁻¹ at pH of 5.

In addition, the removal efficiencies decreased with increasing metal ions concentration, while the temperature did not have any effect on the adsorption process [38]. Xanthate-modified cross-linked magnetic chitosan/poly(vinyl alcohol) showed a maximum adsorption capacity of 139.797 mg g⁻¹ for Cu (II) at 328 K [39]. The maximum adsorption capacity of the cross-linked chitosan bead-grafted poly(methacrylamide) for Cu(II) ions reached 140.90 mg g⁻¹ at pH 4 [40]. Chitosan/poly(vinyl alcohol) beads functionalized with poly(ethylene glycol) showed a percentage of adsorption that reached 99.99% with optimum conditions: 25 mg L⁻¹ Cu(II) ion concentration, 5 h equilibrium time, temperature 45 °C, 1 g L⁻¹ amount of adsorbent and pH 5 [41]. Cross-linked chitosan grafted with polyaniline reached the highest adsorption capacity of Cu(II) ions at 131.58 mg g⁻¹, at a temperature of 40 °C, pH 6, and an initial Cu(II) ion concentration of 100 mg L⁻¹ [42].

Accordingly, for the first time, the present work was directed towards the use of trimellitic anhydride isothiocyanate-cross-linked chitosan hydrogels, that were prepared in our previous study [43], as adsorbents for Cu(II) ions from aqueous solution because of their extreme toxicity. The uptake process of Cu(II) ions by these hydrogels was studied in terms of kinetics and isotherms. The impacts of some parameters on the adsorption process such as concentration of Cu(II) ions, temperature, time, pH, hydrogel dose, and content of cross-linking moieties in the hydrogels were also investigated.

2. Experimental Section

2.1. Materials

Two hydrogels, H₁ and H₂ (Scheme 1 and

Table 1. Synthesis of the cross-linked chitosan hydrogels and their elemental analysis.

Samples	Trimellitic Anhydride Chloride (mmol)	Ammonium Thiocyanate (mmol)	Chitosan (mmol)	Elemental Analysis (%)
				C	H	N	O	S
Chitosan	-	-	-	45.10	6.77	8.43	39.70	-
H1	10.0	10.0	40	46.61	5.48	7.86	36.81	3.24
H2	20.0	20.0	40	47.54	4.68	7.51	35.03	5.24

), used in the present work, were prepared according to our previously described procedure. H₁ and H₂ in the present study have been reported as H₃ and H₄ in our previous work [43]. Copper sulphate pentahydrate (CuSO₄·5H₂O) was purchased from Merck (Darmstadt, Germany).

2.2. pH Value of Zero-Point Charge (pHzpc)

H₁ and H₂ (0.1 g) were individually immersed in NaCl solution (10 mL, 0.1 M) for 24 h. Afterwards, (0.1 N) HCl and (0.1 N) NaOH solutions were used to amend the pH in the range 3–11 and the pH of the solution was always checked by a pH meter (Hanna Model 211). By plotting the initial pH against ΔpH (final pH − initial pH), the pHzpc value was obtained [44,45].

2.3. Adsorption Investigations

A set of experiments was performed to study the adsorption of Cu(II) ions onto H₁ and H₂ hydrogels. In the first experiment, a definite amount of adsorbent (10 mg) was added to Cu(II) ion solution (20–200 mg L⁻¹) with agitation using a water bath shaker (80 rpm) at a constant temperature until an equilibrium was reached. After equilibrium, the solution was passed through filter paper (Whatman, Maidstone, UK, 0.45 µm pore size) and copper ion concentration was determined using an atomic absorption Spectrophotometer (AA-6200 Shimadzu, Kyoto, Japan). Using calibration curves of standard solutions of copper, quantification of the Cu(II) ions was carried out.

To study the effect of pH, the adsorption process was performed over range of pH from 1 to 6 using 0.1 N HCl, while keeping other conditions the same as for the Cu(II) ion solution (10 mL, 100 mg L⁻¹): 10 mg adsorbent, and 25 °C.

To investigate the influence of temperature, the adsorption process was studied at different temperatures of 25, 35, 45 and 55 °C at a pH of the medium equal to 4.

To evaluate the impact of adsorbent dose, the adsorption process was carried out at 25 °C and pH 4 using different doses of the hydrogels (2–15 mg).

All the adsorption capacity experiments were repeated many times and the average of three comparable results were taken.

Equations (1)–(3) can then be applied to calculate the Cu(II) ion quantity adsorbed onto the adsorbent.

$$q_e={\displaystyle \frac{\left(C_o-C_e\right)V}{m}}$$

(1)

$$q_t={\displaystyle \frac{\left(C_o-C_t\right)V}{m}}$$

(2)

$$\% \mathrm{R}\mathrm{e}\mathrm{m}\mathrm{o}\mathrm{v}\mathrm{a}\mathrm{l} \mathrm{e}\mathrm{f}\mathrm{f}\mathrm{i}\mathrm{c}\mathrm{i}\mathrm{e}\mathrm{n}\mathrm{c}\mathrm{y}={\displaystyle \frac{C_o-C_e}{C_o}}\times 100$$

(3)

The capability for adsorption at equilibrium is q_e (mg g⁻¹) and at time q_t (mg g⁻¹). C_o is the concentration of Cu(II) ions before soaking the hydrogel (mg L⁻¹), C_e is the concentration of Cu(II) ions at equilibrium (mg L⁻¹), C_t is the concentration Cu(II) at time t (mg L⁻¹), V is the volume (L) of the solution of Cu(II) ions, and m is the mass used of hydrogel (g).

2.4. Adsorption Kinetic Studies

The study of adsorption kinetics is crucial to comprehending the adsorption rate on the particle surface, because they show the influence of the different conditions on the process rate. This is achieved through the use of models that are able to describe this reaction. It also determines the mechanism by which the metal ions are adsorbed onto the adsorbent material [46].

The kinetic data of Cu(II) ions were modeled using four distinct kinetic models: pseudo-first-order, pseudo-second-order, Elovich, and intraparticle diffusion models.

• Pseudo-first-order model

The correlation between variations in time and adsorption capacity in this kinetic model is of order one. Equation (4) expresses it as follows: k₁ is the rate constant, and q_t and q_e are the adsorption capacities at time t and equilibrium, respectively.

$${\displaystyle \frac{d{q}_t}{dt}}=k_1\left(q_e-q_t\right)$$

(4)

The pseudo-first-order linear Equation (5) is obtained by integrating Equation (4)

$$\mathit{log}\left(q_e-q_t\right)=\mathit{log}{q}_e-{\displaystyle \frac{k_1}{2.303}}t$$

(5)

where t is the time (min) and k₁ (min⁻¹) is the pseudo-first-order rate constant. The values of q_e and k₁ can be found from the slope and intercept of the linear plot of log(q_e − q_t) versus t.

• Pseudo-second-order model

The pseudo-second-order kinetic model, which provides the second-order relationship between adsorption capacity and concentration, is represented by Equation (6). This model suggests that a chemical sorption is involved in the adsorption of dissolved Cu(II) ions by either a cation exchange process or a chemical exchange between these dissolved ions and the adsorbent surface.

$${\displaystyle \frac{d{q}_t}{dt}}=k_2{\left(q_e-q_t\right)}^2$$

(6)

The linear version of pseudo-second-order Equation (7) is obtained by integrating Equation (6):

$${\displaystyle \frac{t}{q_t}}={\displaystyle \frac{1}{k_2{q}_e^2}}+{\displaystyle \frac{t}{q_e}}$$

(7)

where the pseudo-second-order constant is denoted by k₂ (g mg⁻¹ min⁻¹). The values of q_e and k₂ are derived from the slope and intercept, respectively, of the linear plot of t/q_t vs. t.

• Elovich model

This model holds significance in explaining the mechanism of activated chemisorption. It can be used to study the chemisorption kinetics in general, to an extensive range of slow processes of adsorption, and to heterogeneous adsorbent surfaces. Equation (8) is used to express it.

$${\displaystyle \frac{dq}{dt}}={\alpha e}^{-\beta q}$$

(8)

Integration of the rate constant leads to the linear form as shown in Equation (9):

$$q_t={\displaystyle \frac{1}{\beta }} \mathit{ln}(\alpha \beta )+{\displaystyle \frac{1}{\beta }}lnt$$

(9)

where β is the rate constant, also known as the desorption constant, which is correlated with the activation energy of chemisorption and surface coverage (g mg⁻¹), and α is the initial constant rate of the adsorption process (mg g⁻¹ min⁻¹). Plotting q_t against lnt yields a linear connection from which values of α and β can be derived.

• Intraparticle diffusion model

This model is used to calculate the adsorption process’s rate control. Equations (10) and (11) have its expression:

q_t = (k_int t^1/2) + C

(10)

q_t = (k_int t^1/2)

(11)

where C (mg g⁻¹) is a constant that is directly proportional to the boundary layer thickness, and k_int (mg g⁻¹ min^−1/2) is the intraparticle diffusion constant. Both the values of k_int and C can be computed from the slope and intercept of the linear curve (obtained by plotting q_t vs. t^1/2).

In order to quantitatively compare the applicability of each model, a normalized standard deviation Δq is calculated.

$$\mathsf{\Delta }\mathrm{q}(\%)=100\times \sqrt{{\displaystyle \frac{{\left[\right(q_{t,exp}-q_{t,cal})/q_{t,exp}]}^2}{n-1}}}$$

2.5. Investigation of Adsorption Isotherms

The adsorption isotherm is one of the most crucial instruments for comprehending the characteristics of the adsorbent surface. By choosing the proper adsorption equation for various concentration ranges, a clear picture of the surface can be created. The following can be described with the use of the adsorption isotherm: The distribution of adsorbate molecules in equilibrium between the liquid and solid phases, the way the adsorbate interacts with the adsorbent, and the kind and properties of adsorption are the three main aspects to consider. To identify the adsorption system, the data were fitted to a number of models, including the Dubinin–Radushkevich (D-R), Freundlich, Langmuir, and Temkin isotherm models [47].

• The Langmuir isotherm model is expressed by Equations (12)–(14):

$$q_e={\displaystyle \frac{q_{max}{K}_L{C}_e}{1+K_L{C}_e}}$$

(12)

$${\displaystyle \frac{C_e}{q_e}}={\displaystyle \frac{1}{q_{max}{K}_L}}+{\displaystyle \frac{C_e}{q_{max}}}$$

(13)

$$R_L={\displaystyle \frac{1}{(1+K_L{C}_o)}}$$

(14)

where K_L is a constant associated with adsorption energy (L mg⁻¹), R_L is the Langmuir separation factor (essential isotherm character and dimensionless), and $q_{max}$ is the monolayer’s sorption capacity (mg g⁻¹). The isotherm’s linear (R_L = 1), unfavorable (R_L > 1), favorable (0 < R_L < 1), and irreversible (R_L = 0) characteristics are explained by the R_L.

• The Freundlich isotherm model is expressed by Equations (15) and (16):

$$q_e=K_F{C}_e^{{\displaystyle \frac{1}{n}}}$$

(15)

$$\mathit{ln}{q}_e=\mathit{ln}{K}_F+{\displaystyle \frac{1}{n}}{lnC}_e$$

(16)

where K_F and n are empirical constants associated with the adsorption’s capacity and intensity, respectively.

• Temkin isotherm model

An important model that considers the variables that explain the interaction between the adsorbent and adsorbate is the Temkin model. According to the Temkin model, binding energies are dispersed uniformly and the adsorption heat drops linearly when coverage is achieved. Equations (17) and (18), which represent this model, can be used:

$$q_e={\displaystyle \frac{RT}{B_T}}\mathit{ln}{K}_T{C}_e$$

(17)

q_e = B_TlnK_T + B_TlnC_e

(18)

where B_T (J mol⁻¹) is Temkin controlled by temperature constant, and K_T (L g⁻¹) is the constant of Temkin isotherm binding. Plotting q_e against lnC_e yielded the slope and intercept, from which the values of the constants B_T and K_T were derived.

• Dubinin–Radushkevich (D-R) isotherm model

The D-R isotherm model can be utilized to distinguish between chemisorption and physisorption. The formation of multilayers on the microporous adsorbents can be explained by this isotherm. It is consequently more general than the Langmuir isotherm since it does not imply a homogeneous surface. On the other hand, low coverage may indicate the adsorbent’s surface heterogeneity. Equations (19) and (20) provide the D-R isotherm model.

$$q_e=q_m{exp}^{– \beta \epsilon 2}$$

(19)

lnq_e = ln(q_m) − βε²

(20)

Equation (21) can be used to obtain the Polanyi potential, ε, which is related to the adsorption mean free energy per mole of metal ions when transferred from infinity in the solution to the adsorbent surface, and q_m (mg g⁻¹), the monolayer saturation capacity:

$$\mathsf{\epsilon }=\mathit{RT} \mathit{ln}(1+{\displaystyle \frac{1}{C_e}})$$

(21)

where T (K) is the absolute temperature and R (8.314 J mol⁻¹ K⁻¹) is the gas constant.

The values of β and lnq_m in Equation (20) can be determined from the intercept and slope, respectively, of a plot of lnq_e vs. ε². It is possible to calculate the mean free energy using Equation (22).

$$E=({\displaystyle \frac{1}{{\left(2\beta \right)}^{0.5}}})$$

(22)

The process is chemisorption if the mean free energy value is 8 < E < 16 kJ mol⁻¹, while if it is less than 8 kJ mol⁻¹, the process is known as physisorption [48].

2.6. Desorption Investigation

The Cu(II) ions were desorbed from the adsorbent, firstly by washing it utilizing distilled water to eliminate any non-adsorbed Cu(II) ions. Secondly, the adsorbent (10 mg) was soaked in 10 mL of HNO₃ acid (0.1 M) at 25 °C for 24 h. The amount of the desorbed Cu(II) ions from the adsorbent was determined using Equation (23):

$$\mathrm{\%} \mathrm{C}\mathrm{u}\left(\mathrm{I}\mathrm{I}\right) \mathrm{i}\mathrm{o}\mathrm{n}\mathrm{s} \mathrm{d}\mathrm{e}\mathrm{s}\mathrm{o}\mathrm{r}\mathrm{p}\mathrm{t}\mathrm{i}\mathrm{o}\mathrm{n}=q_d/q_a\times 100$$

(23)

where $q_d$ and $q_a$ are the amounts of Cu(II) ions desorbed from and adsorbed onto the adsorbent (mg g⁻¹), respectively [24].

3. Results and Discussion

3.1. pH of Zero-Point Charge (pHzpc)

In specific temperature and aqueous solution composition scenarios, the pH value at which the total charge on the surface equals zero is known as the pH of the zero-point charge (pHzpc). This is not to say that there is zero charge; rather, there is an equal quantity of positive and negative charges on the surface at pHzpc. The amount of the surface charge is determined by the pH of the solution, the kind and quantity of functional groups, or both. Because it measures how easily an adsorbent may bind potentially hazardous ions, pHzpc is significant in surface characterization. This is because there is a net negative charge on the adsorbent surface for pH values greater than pHzpc, which promotes the adsorption of cationic species. In contrast, at pH values below pHzpc, the adsorbent’s surface has a net positive charge of -NH₂⁺, which causes it to reject cations. The value of the pHzpc for H₁ and H₂, obtained by the method of salt addition, was 6.7 and 7.5, respectively (Figure 1). As the pH increases, the surface of H₁ and H₂ becomes more anionic due to the deprotonation of the carboxyl groups of the cross-linking bonds to carboxylate anions (-COO⁻) [44,45].

3.2. Optimizing Adsorption

As illustrated in Figure 2, when Cu(II) ions are initially concentrated from 20 to 180 mg L⁻¹, their adsorption capacity increases from 17.21 to 154.93 mg g⁻¹ for H₁ and from 18.79 to 161.51 mg g⁻¹ for H₂. This suggests that at lower concentrations of Cu(II) ions, there is a tiny ratio between the quantity of Cu(II) ions and the total number of adsorption sites accessible, which leads to fractional adsorption that is independent of the initial concentration. At greater concentrations, fewer adsorption sites are accessible, hence the adsorption of Cu(II) ions depends on the initial concentration. As a result, as the initial concentration of Cu(II) ions increases, the capacity keeps rising [49,50]. Moreover, the resistance to Cu(II) ion mass transfer between the aqueous and solid phases can be overcome by using the greater concentration gradient as a driving force [51].

3.2.1. Effect of Temperature

The findings showed that, from 25 to 35 °C, the percent removal efficiency of Cu(II) ions by H₁ and H₂ hydrogels dramatically dropped (Figure 3). The percentage removal efficiency by H₁ and H₂ at 25 °C was 81.91 and 86.08%, respectively, whereas at 35 °C it was 76.34 and 77.04%, respectively. This suggests that the removal process is exothermic at this point and happens by physisorption. However, as the temperature rose from 35 to 55 °C, the removal efficiency as a percentage gently increased from 76.336 to 77.506%, and from 77.037 to 83.364% by H₁ and H₂, respectively (Figure 3). This indicates that in this stage the removal process has an endothermic nature and occurred via chemisorption. Cu(II) ions became more mobile as a result of the temperature increase, and they were able to obtain sufficient energy to interact with the hydrogels’ sites of activity. Further, the increase in temperature lowered the solution viscosity, increased the diffusion rate of the Cu(II) ions through the external boundary layer of the hydrogels, and enhanced the swelling of the internal structure of the hydrogels. The latter is in accordance with our previous study [43], in which these hydrogels have shown a high degree of swell ability that improved with increasing temperature from 25 to 45 °C. This permitted more Cu(II) ions to infiltrate inside the pores of the hydrogels and improved the percent removal efficiency [52].

3.2.2. Effect of the pH

Because the positive and negative surface charges of the adsorbent are correlated with the pH value, the pH of the solution is crucial for the adsorption of metal ions. To identify the impact of pH of the solution of Cu(II) ions on their removal performance by H₁ and H₂ hydrogels, pH values that ranged from 1 to 6 were examined (Figure 4), because at pH ≥ 7 the Cu(II) ions precipitate as Cu(OH)₂ [53,54,55]. The findings revealed that the removal efficiency decreased (from 96.202 to 76.026 for H₁, and from 97.587 to 76.026 for H₂) with the rising the acidity of the Cu(II) ion solution from pH 6 to pH 1. The rivalry between H⁺ and Cu²⁺ for the active sites on hydrogels could be the cause of this; the adsorption of H⁺ ions predominates due to their smaller size relative to that Cu²⁺ ions.

3.2.3. Hydrogel Dosage Effect

While holding the other variables constant (pH = 4, temperature = 25 °C, stirring speed = 80 rpm, and contact time = 24 h), different doses of the hydrogels (2, 5, 10, and 15 mg) in 100 mL of Cu(II) ion solution (100 mg L⁻¹) were used to study the dosage effect on the percentage removal of Cu(II) ions, as shown in Figure 5. The findings demonstrate that the percentage removal of Cu(II) ions by H₁ and H₂ rose from 74.1 to 81.34% and from 75.08% to 82.28%, respectively, as the adsorbent dose increased from 2 to 15 mg. This is due to the increase in the adsorbent surface area with the increase in adsorbent dosage, and the availability of more exchangeable binding sites on the surface, ready for the uptake of Cu(II) ions [56,57,58,59,60].

3.2.4. Cross-Linking Content Effect

Although the cross-linking moieties content of H₂ is greater by two times than that of the H₁ (Table 1), there is no big difference between them in Cu(II) ion removal (Figure 2, Figure 3, Figure 4 and Figure 5). This could be explained by the existence of residual primary amino groups in H₁ that did not participate in the cross-linking process but which efficiently shared in the removal process of Cu(II) ions, indicating that there is no need for a high degree of cross-linking to achieve the targeted Cu(II) ion removal.

Furthermore, it is important to note that the adsorption of Cu(II) ions is dependent on the chitosan’s major amino groups. However, one of chitosan’s biggest drawbacks is that it dissolves easily in acidic solutions, which is particularly problematic when using it as an adsorbent for Cu(II) ions in acidic medium. Thus, the chemical cross-linking process of chitosan using trimellitic anhydride isothiocyanate is one of the few promising alternatives for insolubility of produced cross-linked chitosan-based hydrogels in acidic solutions as reported in our previous work [43]. Additionally, this leads to the incorporation of functional groups that can serve as sites for binding Cu(II) ions, such as thiourea, amides, and carboxylic groups. This prevents the anticipated drop in Cu(II) ions adsorption capacity that is caused by the consumption of chitosan’s primary amino groups during the cross-linking process. As would be predicted, the elimination efficiency percentages of Cu(II) ions (adsorbent weight = 10 mg, Cu(II) ion solution (10 mL, 100 mg L⁻¹), pH = 6, and temperature = 25 °C) by H₁ and H₂ were 96.20%, and 97.59%, respectively, which seem to be comparable to that of the parent chitosan (98.43%).

3.3. Adsorption Kinetics

Figure 6 and Figure 7 and Online Supplemental Figures S1 and S2 illustrate the outcomes of the adsorption kinetics of Cu(II) ions onto H₁ and H₂: adsorbent dose = 50 mg, Cu(II) ion solution concentration (50 mL, 100 mg L⁻¹), 25 °C, and solution pH 6.

Table 2. Kinetic model constants and correlating coefficients of adsorption of Cu(II) ions onto H₁ and H₂.

Kinetic Models	Parameters	H1	H2
qe.exp (mg g−1)		96.2	97.59
Pseudo- first-order	R2	0.645	0.944
	qe.cal (mg g−1)	3.043	1.98
	k1 (min−1)	0.0151	0.0166
	Δq	47.14	52.23
Pseudo- second-order	R2	0.999	1
	qe.cal (mg g−1)	98.04	99.01
	k2 (10−5)	0.005	0.0146
	(g mg−1 min−1)
	Δq	3.2	2.93
Elovich	R2	0.722	0.979
	β (g mg−1)	0.361	0.698
	α (1013)	2.714	1.105 × 1015
	(mg g−1 min−1)
	Δq	85.44	42.64
Intraparticle diffusion	R2	0.778	0.967
	k (mg g−1 min−1/2)	0.853	0.446
	Δq	89.79	93.47

provides an overview of the various kinetic characteristics of the Elovich, pseudo-second-order, pseudo-first-order, and intraparticle diffusion models for the adsorption of Cu(II) ions onto H₁ and H₂.

The pseudo-second-order model had the highest correlation coefficient (R²) values and the low values of the normalized standard deviation Δq_e (%) when compared to the other models (Table 2). For the adsorption of Cu(II) ions onto H₁ and H₂, they are 0.999 and 1.00, respectively. This suggests that the pseudo-second-order model accurately describes the Cu(II) ions’ adsorption by H₁ and H₂. Furthermore, the experimental q_e values (97.59 and 96.20 mg g⁻¹ for H₁ and H₂, respectively) are similar to the calculated ones (98.04 and 99.01 mg g⁻¹ for H₁ and H₂, respectively) as presented in Table 2, indicating a strong agreement with pseudo-second-order kinetics. This is a reflection of and validation of the pseudo-second-order kinetic model’s perfect match to the Cu(II) ions adsorption process onto H₁ and H₂.

3.4. Isotherms of Adsorption

The results of adsorption isotherms of Cu(II) ions onto H₁ and H₂ (T = 298 K, pH = 4, shaking speed = 80 rpm, hydrogel dose = 10 mg, and time = 24 h) are shown in Figure 8 and Figure 9 and in Online Supplemental Figure S3 and Online Supplemental Figure S4, respectively. The correlation coefficient (R²) value of the Freundlich isotherm linear plot (0.965 and 0.952 for H₁ and H₂, respectively) was satisfactorily higher than that obtained by the Langmuir (0.014 and 0.744 for H₁ and H₂, respectively), Temkin (0.802 and 0.843 for H₁ and H₂, respectively), and D-R models (0.785 and 0.765 for H₁ and H₂, respectively), as shown in

Table 3. Parameters of the adsorption isotherms for the removal of Cu(II) ions by H₁ and H₂ hydrogel.

Models	Parameter	H1	H2
Langmuir	qmax (mg g−1)	2500	227.27
	RL	0.333–0.833	0.357–0.833
	KL (L mg−1)	0.002	0.068
	R2	0.014	0.744
Freundlich	1/n	0.969	0.681
	Kf (mg g−1)	6.33	18.05
	R2	0.965	0.952
Temkin	B (kJ mol−1)	62.7	44.02
	KT (L g−1)	2.921	1.066
	R2	0.802	0.843
D-R	qm (mg g−1)	113.52	101.19
	E kJ mol−1	0.408	0.845
	B × 10−6	3	0.7
	R2	0.785	0.765

. This means that the Freundlich isotherm model fits the experimental data quite well and is very suitable to describe the adsorption equilibrium of Cu(II) ions on H₁ and H₂. From this, it can be concluded that the nature of the adsorption is multi-layered; the distribution is not uniform. Similar experiments were performed by Samrot et al. [45] using chitosan-coated SPIONs for chromium removal, and by Jiang et al. [61] using Γ-Fe₂O₃ nanoparticles encapsulated in millimeter-sized magnetic chitosan beads for removal of Cr(VI) from water.

3.5. Desorption Evaluation

The capability for reusing adsorbents is considered as one of the worthiest features that lowers the adsorption technique costs through the restoration of the adsorbents [24]. The desorption of the Cu(II) ions was calculated using Equation (23). Nitric acid, as a desorption medium, manifested that the percentage of desorption attained was 70.11% and 61.57% from H₁ and H₂, respectively. These findings confirmed that the hydrogels can be effectively reutilized for adsorbing Cu(II) ions from the aqueous solutions.

3.6. Comparison of Adsorption Capacity of Different Adsorbents for Cu(II) Ions

For the removal of Cu(II) ions, different adsorbents were previously reported [42,62,63,64,65,66,67,68].

Table 4. The maximum adsorption capacities (q_max) of different reported adsorbents for Cu(II) ions.

Adsorbent	Adsorption Capacity (mg g−1)	Temperature (°C)	Metal Conc. (mg L−1)	Adsorbent Dose (g)	pH	Ref.
Cross-linked chitosan grafted with polyaniline	131.58	20–40	100	0.05	6	[42]
Quaternized chitosan@chitosan cationic polyelectrolyte microsphere	687.6	25	0–2000	0.075	5	[62]
Horn Core calcined at 400 °C (P400)	99.98	25	100–500	0.02	5	[63]
Magnetic chitosan@bismuth tungstate coated by silver (MCTS-Ag/ Bi2WO6)	181.8	20–40	20 -120	0.02	6	[64]
Xanthate-modified magnetic chitosan	34.5	25	100	-	5	[65]
Epichlorohydrin cross-linked xanthate chitosan	43.47	50	100	-	5	[66]
Medicinal plant (Salvadora persica)	74.30	25	100	-	4	[67]
Pecan nutshell (Carya illinoinensis)	23.37	30	-	-	5	[68]
H1	96.20	25	100	0.01	6	Present study
H2	97.59	25	100	0.01	6	Present study

lists the adsorption capabilities of various different kinds of sorbents. It is clear that the adsorption capacity of the H₁ and H₂ hydrogels with respect to Cu(II) ions is better or even comparable to that of the reported adsorbents [42,62,63,64,65,66,67,68].

4. Conclusions

The adsorption of Cu(II) ions from aqueous solution for the treatment of industrial wastewater was studied using two chitosan-based hydrogels cross-linked with varying concentrations of trimellitic anhydride isothiocyanate. The hydrogels’ strong adsorption capabilities for Cu(II) ions were demonstrated by the results. The Freundlich isotherm and pseudo-second-order models might more accurately explain the equilibrium results, suggesting that the adsorption is multilayered in nature. It was discovered that pH 6, 25 °C, 200 mg L⁻¹ of Cu(II) ion concentration, and 15 mg of the hydrogels were the ideal values for the elimination of copper ions. For H₁ and H₂, the removal efficiencies were 97.59% and 96.20%, respectively. Therefore, chitosan cross-linking with trimellitic anhydride isothiocyanate demonstrated several benefits: (1) preventing chitosan from becoming soluble in acidic solutions (adsorption media); (2) adding functional groups such as thiourea, amide, and carboxylic groups that can serve as sites for binding Cu(II) ions; and (3) preventing the anticipated decrease in Cu(II) ion adsorption capacity as a result of chitosan’s primary amino groups being consumed during the cross-linking process. This is a sensible strategy that could encourage the creation of novel adsorbents for the extraction of heavy metal ions from industrial waste water. Further, the findings proved that these hydrogels can be efficiently reused for the adsorption of Cu(II) ions from the aqueous solution.