Adsorptive Elimination of Cu(II) Ions from Aqueous Solution onto Chitosan Modified with Uracil

Abstract

The remediation of industrial wastewater to eliminate heavy metal ions represents a pressing environmental requirement. A previously prepared adsorbent, based on chitosan modified with uracil (UCs), was investigated for the first time in this work to eliminate Cu(II) ions. The best conditions for elimination were as follows: adsorbent dose = 0.01 g, Cu(II) ions solution concentration = 0.2 g L⁻¹, pH = 6, and temperature = 25 °C. The adsorption kinetics were favorable for the pseudo second order due to the correlation coefficient’s value being the highest (R² = 1.0). The experimental value of q_e (99.65 mg g⁻¹) was comparable to that of the theoretical one (100 mg g⁻¹). The removal efficiency reached 99.65%, and the adsorption isotherm coincided with the Freundlich model, denoting that the nature of its adsorption was multi-layered. Cu(II) ions removal mainly relies on the physisorption phenomenon. The desorption percentages reached 92.65, 75.29, 49.29, and 19.92% after four successive cycles. So, the insertion of nitrogen-rich uracil along the chitosan chains, as binding sites for Cu(II), is deemed to be an outstanding opportunity to produce an appropriate, efficacious adsorbent that is a good choice to apply in the metal removal domains.

1. Introduction

Global ecological systems, in particular aquatic resources, are greatly impacted by the problem of heavy metals pollution, because these do not degrade naturally. They commonly accumulate in the biological systems of aquatic living organisms, causing a great risk because of their transfer via the food chain [1]. Heavy metals can pollute ecosystems through various activities, including fertilizers, electroplating, printing, and the steel industry. They can also contaminate the environment through some other activities like mining, the production of dyes, and the creation of plant disease control agents, as well as through the natural weathering of the Earth’s crust and soil erosion [2]. The concentration of heavy metal ions in aqueous solutions has been reduced using various methods, including chemical precipitation [3], ion exchange [4], activated carbon adsorption [5], and membrane filtration [6]. However, these techniques have drawbacks, such as a high energy consumption, the generation of hazardous by-products from the reagents used, and a reduced effectiveness at low concentrations of metal. The biosorption process is an alternative approach that has received great attention from researchers. Among its primary features are its short operating time, low operating costs, lack of undesirable secondary products, environmental friendliness, abundance of adsorbents, and ease of restoration for future use. Transport processes both active and passive are involved in the biosorption process. Then, a rapid and reversible process of accumulation begins. The second stage, commonly known as active uptake [7], includes metabolic-activity-related intracellular bioaccumulation that is slower and frequently irreversible. The fourth period of the periodic table has a large number of metals that are thought to cause cancer. The electronic structure of transition metals is believed to contribute to their carcinogenic properties. The presence or absence of copper in a variety of enzymes causes issues for all life forms [8]. Undoubtedly, one of the most widely utilized heavy metals is copper. The most common type of copper ion found in the environment is Cu(II), which is poisonous to a variety of living organisms [9,10]. People’s food and drinking water may contain toxic copper at abnormally high concentrations as a result of certain practices, such as using cookware that has been lined or glazed with copper or using copper pipes for water supply [11]. Additional sources of copper include sewage effluent, which comprises wastes from the mining and fertilizer industries, as well as wastes from baths containing paints, pigments, and plating [12]. Although a trace amount of copper is necessary for life, an excess of copper in water can cause neurotoxicity, jaundice, diarrhea, respiratory issues, liver and kidney failure, and even death [13,14]. The excessive absorption of copper in the body also has the potential to cause severe irritation and deterioration of the gastric mucous membranes, as well as extensive capillary damage and depression [15,16]. One epidemiological finding that points to copper’s primary role as a carcinogen is the high cancer incidence among coppersmiths [17]. The World Health Organization and the US Environmental Protection Agency (EPA) have set acceptable limits for copper in drinking water at 2.0 and 1.3 mg L⁻¹, respectively [18].

Due to the outstanding features of the biopolymer chitosan, such as antimicrobial activity [19], the ability to adsorb metals [20], and an adsorption capacity for anionic dyes [21], it has become increasingly important in environmental biotechnology. Additionally, chitosan is abundant, biocompatible, biodegradable, renewable, and non-toxic [22]. A major drawback of chitosan is its solubility in acidic conditions. Since industrial wastewater is often acidic, this presents a challenge when using chitosan as an adsorbent material for pollutant removal. For these kinds of applications, controlling chitosan’s solubility is crucial. This can be achieved through methods such as graft copolymerization [23], the blending of polymers [24], or chemical substitution processes [25].

The chemical cross-linking of chitosan is employed to enhance its mechanical properties, increase its chemical stability, slow down its degradation rate, and prolong its lifespan in various environments [26]. Copper was removed using chitosan-grafted polyacrylonitrile. The copper adsorption followed the pseudo-second-order model and the Langmuir model. The optimal conditions for adsorption were found to be a pH of 7.5, an adsorbent amount of 5 g, and an equilibrium time from 5 to 6 h, resulting in a maximum adsorption capacity of 239.314 mg g⁻¹ [27]. Additionally, chitosan-coated MnFe₂O₄ nanoparticles were able to eliminate hazardous Cu(II) ions in trace amounts, achieving an adsorption capacity of 22.6 mg g⁻¹ [28]. At pH 5, chitosan-coated cotton gauze could absorb a maximum concentration of 14.4 mg g⁻¹ of Cu(II) ions. Furthermore, as the concentration of metal ions increased, the removal efficiencies decreased as well; the adsorption process was unaffected by temperature [29]. The highest capacity for the adsorption of xanthate-modified cross-linked magnetic chitosan/poly (vinyl alcohol) for Cu(II) at 328 K was reported to be 139.797 mg g⁻¹ [30].

At pH 4, the cross-linked chitosan-bead-grafted poly (methacrylamide) exhibited a maximum capacity of adsorption of 140.9 mg g⁻¹ for Cu(II) ions [31]. The percentage of adsorption of chitosan/poly (vinyl alcohol) beads functionalized with poly (ethylene glycol) reached 99.99% under ideal circumstances, pH 5, 45 °C, 1 g L⁻¹ of adsorbent, 5 h of equilibrium time, and a concentration of 25 mg L⁻¹ Cu(II) ions [32]. The highest adsorption capacity for Cu(II) ions was discovered in polyaniline-grafted cross-linked chitosan, achieving 131.58 mg g⁻¹ at 40 °C, with a pH of 6 and an initial concentration of 100 mg L⁻¹ of Cu(II) ions [33].

Uracil is a naturally originating component, one of the pyrimidine class members, and represents one of the four bases of biopolymer nucleic acid RNA. Two H-bonds are formed between uracil and adenine in RNA. Uracil is replaced by thymine as the nucleobase in DNA. It can be deemed as a version of demethylated thymine [34,35]. Uracil-based compounds have attracted much attention from chemists and medicinal researchers for discovering new promising drugs [36]. Uracil assists enzymes to carry out different processes, so it is essential in the medication industry and drug delivery systems. In addition, the capacity of uracil compounds to create H-bonds is the main reason for their use beyond their distinguished biological features [37]. During the last two decades, several researchers have developed efficacious methods to synthesize and manipulate some functionalized uracil moieties. The latter have attracted great interest in the domain of drug evolution and other relevant applications, due to their assorted scope of biological characteristics, in addition to their make-up availability [38]. The diversified functionality in uracil-based compounds is ascribed to the multiple substituent groups at the N1, N3, C5, and C6 sites. C5 in uracil-based derivatives is known to be a nucleophilic position. Their reactivities are remarkably governed via their substituent’s nature at their backbone. It is found that the C5 place becomes highly nucleophilic if there is a substituent -NH₂ group on C6 in uracil relative to the un-substituted uracil. Thus, one of the substantial methods utilized for the diversity-directed preparation of uracil-based derivatives is the modification of the substituent groups at the C6 place on the uracil ring [39]. Furthermore, 2-thiouracil can be adsorbed chemically on the surfaces of copper, gold, and silver electrodes via its deprotonated sulfur atom [40]. The macro- and microscopic characteristics of solid humic acid (HA) extracted from a German peat (GHA) and a New Hampshire soil (NHA) were determined using adsorption, as analytical probes, for copper(II) strongly bound with uracil, the nucleic acid constituent. Cu(II) decreased the amount of uracil adsorbed by GHA [41].

In view of the above considerations, we investigated uracil-modified chitosan (UCs), which was prepared in our previous study [42], as an adsorbent for Cu(II) ions from aqueous solutions, for the first time in this work. The kinetics and isotherms of the Cu(II) ions’ uptake process by the UCs were studied. The current work was also extended to evaluate the effect of different parameters on the adsorption process, including the adsorbent dose, temperature, time, pH, and concentration of Cu(II) ions.

2. Experimental Section

2.1. Materials

Chitosan, having a molecular weight of 1.0–3.0 × 10⁵ g mol⁻¹ and a degree of deacetylation of 98%, was provided by Acros Organics (Fair Lawn, NJ, USA). Epichlorohydrin and benzaldehyde were obtained from Pan-Reac. AppliChem-ITW Reagent ((Darmstadt, Germany). 6-Amino-1,3-dimethyl uracil was purchased from Sigma-Aldrich (Munich, Germany). The other solvents and chemicals were supplied by Aldrich (Darmstadt, Germany).

2.2. Preparation of Chitosan Modified with Uracil (UCs)

The adsorbent under investigation (UCs, Scheme 1) was obtained following the four-step method reported in our earlier work [42]. In brief, the 1^ry -NH₂ groups were initially shielded via the production of a chitosan Schiff’s base through a reaction with benzaldehyde. Secondly, the epichlorohydrin reaction was intended for the 1^ry -OH groups to obtain an epoxy chitosan Schiff’s base. In the third step, 6-amino-1,3-dimethyluracil opened the epoxide rings to obtain the uracil chitosan Schiff’s base. The latter, in the fourth step, was subjected to acid hydrolysis to remove its benzaldehyde moieties and recover the 1^ry -NH₂ groups, generating a uracil-modified chitosan (UCs) adsorbent. The incorporation of basic functional groups of 6-amino-1,3-dimethyluracil together with both the amino and hydroxyl groups in one structure can effectively improve the removal of metals such as Cu(II) ions from aqueous solutions.

2.3. Characterization of UCs

The chemical structures, internal morphologies, and topographical features of the modified chitosan derivatives were characterized using some convenient analytical techniques, including FTIR spectroscopy (a Thermo Scientific Nicolet 6700 FTIR spectrometer (Yokohama, Japan) from 4000 to 500 cm⁻¹ wave number range), X-ray diffractometry (a Rigaku Ultima-IV wide-angle X-ray diffractometer (Tokyo, Japan) at range of diffraction angles (2θ) from 5 to 80° with a 5° min⁻¹ velocity), and scanning electron microscopy (a field emission scanning electron microscope JSM-7610F (Jeol, Freising, Germany) at a 15 kV accelerating voltage and 8000× magnification). The results, which are available in Supplementary Figure S1A, were consistent with that of our previous investigation [42].

In comparison to the FTIR spectrum of virgin chitosan (Supplementary Figure S1A), the chitosan Schiff’s base showed some additional new absorption peaks at 3052 and 3027 cm⁻¹ corresponding to C-H (Ar), at 1691 cm⁻¹ related to C=N groups, at 1600, 1579, 1493, and 1454 cm⁻¹ assigned to C=C (Ar), and at 757 and 692 cm⁻¹ attributed to a mono-substituted aromatic moiety. The epoxy chitosan Schiff’s base showed another new peak at 1250 cm⁻¹ that was indicative of the presence of epoxide nuclei. The latter peak disappeared, as shown in the uracil chitosan Schiff’s base spectrum, together with the manifestation of a new peak at 1000 cm⁻¹ ascribed to the C-H of pyrimidine moieties. At 1657 cm⁻¹, the peak associated with the uracil moieties’ carbonyl groups overlapped with that of the amide I of chitosan. Finally, the demise of the peaks at 757 and 692 cm⁻¹, corresponding to the mono-substituted aromatic nuclei, confirmed the elimination of the benzaldehyde nuclei to produce the UCs [42].

In comparison to the XRD pattern of virgin chitosan (Supplementary Figure S1B), the modified chitosan derivatives showed a lower crystallinity. This was demonstrated by the vanishing of the peak near 2θ = 10° and a reduction in the peak intensity around 2θ = 20°. This was due to the chitosan functionality being significantly altered, with a great lowering in the hydrogen bonds owing to the exhaustion of its polar groups (−NH₂ and/or −OH) and the insertion of moieties of the modifiers. The latter separated the chitosan chains from one another, leading to an increase in the amorphous region and a decrease in the crystalline region [42].

As illustrated in SEM images (Supplementary Figure S1C), the smooth surface of chitosan was changed after its modification into raucous surfaces, having lumps of different sizes owing to the different sizes of the inserted modifiers. The distribution of these lumps was homogeneous in all derivatives, confirming that the all stages of the chitosan modification process were successfully accomplished. The inclusion of modifiers separated the polymer chains, reduced the formation of interchain hydrogen bonds, and created porous matrices with high surface areas [42].

Thus, the FTIR spectra, XRD patterns, and SEM images confirmed that all the stages of the chitosan modification process were successfully accomplished. However, elemental analyses of both the chitosan and UCs adsorbent were measured using the EDS technique to give additional proof of the successful synthesis of the UCs. The results of the elemental analysis illustrated appreciable increases in both the carbon and nitrogen elements at the expense of the oxygen elements of the UCs adsorbent (56.4, 18.4, and 25.2%, respectively) relative to that of the virgin chitosan (51.6, 7.0, and 41.4%, respectively), due to the incorporation of the uracil moiety, as represented in Figure 1. This supports the results of the above-mentioned techniques used for proving the successful preparation of UCs adsorbent.

It is worth pointing out that the prepared UCs had a substitution degree of 47.85, which was calculated on the basis of its C/N value, obtained from the elemental analysis (EDS, Figure 1), following the formula reported previously [43].

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

Several samples of UCs adsorbent (100 mg) were separately sunk in 10 mL of a 0.1 M aqueous sodium chloride solution for 24 h. Thereafter, aqueous solutions of hydrochloric acid and sodium hydroxide at a 0.1 N concentration were employed for adjusting the pH values to be between 3 and 11, which were measured using a pH meter (Hanna Model 211, Padova, Italy). The value of the pHzpc was procured from the plot of the initial pH versus the ΔpH (Final pH − Initial pH) [44,45].

2.5. Adsorption Measurements

A series of experiments was carried out for studying the Cu(II) ions’ adsorption onto the UCs adsorbent. A known weight of UCs adsorbent (10 mg) was introduced into conical flasks containing a Cu(II) ions aqueous solution, with concentrations ranging from 20 to 200 mg L⁻¹, and agitated (80 rpm), employing a water bath equipped with a shaker at a fixed temperature until it reached equilibrium. Afterwards, the solutions were left along a Whatman filter paper with a 0.45 µm pore size, and an atomic absorption spectrophotometer (AA-6200 Shimadzu, Kyoto, Japan) was used to measure the concentrations of Cu(II) ions. The latter were quantified employing the calibration curves of copper standard solutions. The absorption capacity estimation was repeated several times, and the average of three results was obtained.

For studying the pH impact, the adsorption procedures proceeded at a pH range from 1 to 6 via dropping HCl (0.1 N), while holding other parameters the same, such as a 10 mL Cu(II) ions solution with a 100 mg L⁻¹ concentration, 10 mg of UCs, and 25 °C.

For investigating the temperature effect, the adsorption procedure was carried out at diverse temperatures (25, 35, 45, and 55 °C), a 10 mL Cu(II) ions solution with a 100 mg L⁻¹ concentration, 10 mg of UCs, and at a solution pH of 4.

For evaluating the adsorbent dose influence, adsorption experiments were performed utilizing different UCs doses that ranged from 2 to 10 mg, a 10 mL Cu(II) ions solution with a 100 mg L⁻¹ concentration, a pH of 4, and a temperature 25 °C.

The adsorbed amount of Cu(II) ions by the UCs adsorbent can then be calculated applying the Equations from (1) to (3).

$${\mathrm{q}}_{\mathrm{e}}={\displaystyle \frac{\left({\mathrm{C}}_{\mathrm{o}}-{\mathrm{C}}_{\mathrm{e}}\right)\mathrm{V}}{\mathrm{m}}}$$

(1)

$${\mathrm{q}}_{\mathrm{t}}={\displaystyle \frac{\left({\mathrm{C}}_{\mathrm{o}}-{\mathrm{C}}_{\mathrm{t}}\right)\mathrm{V}}{\mathrm{m}}}$$

(2)

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

(3)

These equations denote the following: the adsorption capacities at equilibrium (q_e, mg g⁻¹) and at time t (q_t, mg g⁻¹). C_o (the Cu(II) ion concentration prior immersing the UCs, mg L⁻¹), C_e (the Cu(II) ion concentration at equilibrium, mg L⁻¹), C_t (the Cu(II) ion concentration at time t, mg L⁻¹), V (the Cu(II) ion solution volume, L), and m (the UCs mass, g).

2.6. Adsorption Kinetics

The kinetics of adsorption should be studied in order to comprehend the rate of adsorption onto the surface of the adsorbents, since they manifest the different conditions’ impacts on the rate of the adsorption procedure. This can be achieved using some models that have a capability for describing this process. They also can determine the metal ions’ adsorption mechanism by the adsorbent matrix [46].

Four distinct kinetic models—the pseudo-first-order, pseudo-second-order, Elovich, and intraparticle diffusion models—were used to model the kinetic data of the Cu(II) ions’ adsorption by the UCs adsorbent.

• Pseudo-First-Order Model

In this model, the liaison between the adsorption capacity and time variations is of the order one. It is expressed by Equation (4), where k₁ is the rate constant.

$${\displaystyle \frac{{\mathrm{dq}}_{\mathrm{t}}}{\mathrm{dt}}}={\mathrm{k}}_1\left({\mathrm{q}}_{\mathrm{e}}-{\mathrm{q}}_{\mathrm{t}}\right)$$

(4)

The linear equation of the pseudo-first order model (Equation (5)) is given via the integration of Equation (4).

$$\log \left({\mathrm{q}}_{\mathrm{e}}-{\mathrm{q}}_{\mathrm{t}}\right)=\log {\mathrm{q}}_{\mathrm{e}}-{\displaystyle \frac{{\mathrm{k}}_1}{2.303}}\mathrm{t}$$

(5)

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

• Pseudo-Second-Order Model

The kinetic model of the pseudo second order, which supplies the second-order relation between the concentrations and the adsorption capacities, is expressed by Equation (6). It proposes that Cu(II) ions’ adsorption from the aqueous solution is proceeded by a chemisorption reaction via either a cationic exchange procedure or a chemical exchange between the dissolved Cu(II) ions and the surface of the adsorbent.

$${\displaystyle \frac{{\mathrm{dq}}_{\mathrm{t}}}{\mathrm{dt}}}={\mathrm{k}}_2{\left({\mathrm{q}}_{\mathrm{e}}-{\mathrm{q}}_{\mathrm{t}}\right)}^2$$

(6)

Equation (7) displays the linearized version of the pseudo-second-order model, which is obtained by integrating Equation (6).

$${\displaystyle \frac{\mathrm{t}}{{\mathrm{q}}_{\mathrm{t}}}}={\displaystyle \frac{1}{{\mathrm{k}}_2{\mathrm{q}}_{\mathrm{e}}^2}}+{\displaystyle \frac{\mathrm{t}}{{\mathrm{q}}_{\mathrm{e}}}}$$

(7)

where k₂ is the pseudo-second-order constant (g mg⁻¹ min⁻¹). From the intercept and the slope of the linear plot of t/q_t vs. t, the k₂ and q_e values can be obtained, respectively.

The rate of the initial adsorption h (mg g⁻¹ min⁻¹) can be determined from Equation (8).

$${\mathrm{h}=\mathrm{k}}_2{\mathrm{q}}_{\mathrm{e}}^2$$

(8)

• Elovich Model

This model is of great importance in interpreting the active chemical sorption mechanisms. It is used for generally studying chemical sorption kinetics, in a considerable domain of slow adsorption reactions, and considering the adsorbent having heterogeneous surfaces. It can be denoted by Equation (9)

$${\displaystyle \frac{\mathrm{d}\mathrm{q}}{\mathrm{d}\mathrm{t}}}={\alpha \mathrm{e}}^{-\beta \mathrm{q}}$$

(9)

Integrating the rate constant results in the linearized form is expressed in Equation (10).

$${\mathrm{q}}_{\mathrm{t}}\mathrm{ }=\mathrm{ }{\displaystyle \frac{1}{\beta }}\mathrm{ }\mathrm{l}\mathrm{n}\left(\alpha \beta \right)+{\displaystyle \frac{1}{\beta }}\mathrm{l}\mathrm{n}\mathrm{t}$$

(10)

where β is the constant of the rate (likewise recognized as the constant of the desorption), which is related to the energy of activation of chemisorption and surface coverage (g mg⁻¹), and α is the initial adsorption process rate constant (mg g⁻¹ min⁻¹). From the linear relation obtained from plotting q_t versus ln t, the α and β values are determined.

• Intraparticle Diffusion Model (the Weber–Morrison model)

This model was utilized for calculating the rate control of the adsorption procedure. It can be expressed by Equations (11) and (12).

q_t = (k_int t^1/2) + C

(11)

q_t = (k_int t^1/2)

(12)

where C (mg g⁻¹) is a constant and is directly proportionated to the intensity of the boundary stratum, and k_int (mg g⁻¹ min^−1/2) is the constant of the intraparticle diffusion. The intercept and the slope, obtained from the linear relation of plotting q_t vs. t^1/2, give the values of C and k_int, respectively.

2.7. Adsorption Isotherms

The adsorption isotherm can describe the following: the distribution of the molecules of the adsorbate at equilibrium between the solid and liquid phases, the style by which the interaction between the adsorbate and the adsorbent takes place, and the nature and characteristics of adsorption, which are the three prime aspects to consider. To recognize the adsorption manner, the experimental data were examined for fitting to some models, including the Langmuir, Freundlich, Temkin, and Dubinin–Radushkevich (D-R) models.

• The Langmuir isotherm model can be expressed by Equations (13)–(15).

$${\mathrm{q}}_{\mathrm{e}}={\displaystyle \frac{{\mathrm{q}}_{\mathrm{m}\mathrm{a}\mathrm{x}}{\mathrm{K}}_{\mathrm{L}}{\mathrm{C}}_{\mathrm{e}}}{1+{\mathrm{K}}_{\mathrm{L}}{\mathrm{C}}_{\mathrm{e}}}}$$

(13)

$${\displaystyle \frac{{\mathrm{C}}_{\mathrm{e}}}{{\mathrm{q}}_{\mathrm{e}}}}={\displaystyle \frac{1}{{\mathrm{q}}_{\mathrm{m}\mathrm{a}\mathrm{x}}{\mathrm{K}}_{\mathrm{L}}}}+{\displaystyle \frac{{\mathrm{C}}_{\mathrm{e}}}{{\mathrm{q}}_{\mathrm{m}\mathrm{a}\mathrm{x}}}}$$

(14)

$${\mathrm{R}}_{\mathrm{L}}={\displaystyle \frac{1}{(1+{\mathrm{K}}_{\mathrm{L}}{\mathrm{C}}_{\mathrm{o}})}}$$

(15)

where K_L (L mg⁻¹) is a constant that is correlated with the energy of the adsorption, R_L is the Langmuir separation factor (essential isotherm character and dimensionless), and $q_{max}$ (mg g⁻¹) is the capacity for the sorption of a monolayer. The R_L elucidates the isotherm linear nature (R_L = 1), unfavorable (R_L > 1), favorable (0 < R_L < 1), and irreversible (R_L = 0).

• The Freundlich isotherm model can be expressed by Equations (16) and (17).

$${\mathrm{q}}_{\mathrm{e}}={\mathrm{K}}_{\mathrm{F}}{\mathrm{C}}_{\mathrm{e}}^{{\displaystyle \frac{1}{\mathrm{n}}}}$$

(16)

$$\ln {\mathrm{q}}_{\mathrm{e}}=\ln {\mathrm{K}}_{\mathrm{F}}+{\displaystyle \frac{1}{\mathrm{n}}}{\mathrm{l}\mathrm{n}\mathrm{C}}_{\mathrm{e}}$$

(17)

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

• Temkin isotherm Model

This is a paramount model which deems the variables which interpret the interaction of the adsorbate with the adsorbent. In accordance with this model, the distribution of the binding energy is uniform and the heat of adsorption linearly drops when coverage is attained. Equations (18) and (19) are used to express this model.

$${\mathrm{q}}_{\mathrm{e}}={\displaystyle \frac{\mathrm{R}\mathrm{T}}{{\mathrm{B}}_{\mathrm{T}}}}\ln {\mathrm{K}}_{\mathrm{T}}{\mathrm{C}}_{\mathrm{e}}$$

(18)

qe = B_TInK_T + B_TIn C_e

(19)

where B_T (Temkin constant, J mol⁻¹) is governed by temperature, and K_T (L g⁻¹) is the Temkin isotherm binding constant. The plot of q_e versus ln C_e is used to obtain the intercept and the slope from which the constants K_T and B_T values are derived.

• Dubinin–Radushkevich (D-R) isotherm Model

The D-R isotherm is employed to differentiate between physisorption and chemisorption. It can be used to interpret the construction of multi-layers on the adsorbents of micro-porosity. Thus, it is a more generic model in comparison to the Langmuir isotherm because it does not imply the homogeneity of the surface. On the other hand, poor coverage may refer to the heterogeneous surfaces of adsorbents. This model can be expressed by Equations (20) and (21).

$${\mathrm{q}}_{\mathrm{e}}={\mathrm{q}}_{\mathrm{m}}{\mathrm{e}\mathrm{x}\mathrm{p}}^{{–\beta \epsilon }^2}$$

(20)

Ln q_e = ln (q_m) − βε²

(21)

q_m (mg g⁻¹) is the capacity of monolayer saturation. Equation (22) is used to obtain the Polanyi potential (ε).

$$\epsilon \mathrm{ }=\mathrm{ }\mathrm{R}\mathrm{T}\mathrm{ }\mathrm{l}\mathrm{n}\mathrm{ }(1+{\displaystyle \frac{1}{{\mathrm{C}}_{\mathrm{e}}}})$$

(22)

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

From the plot of ln q_e vs. ε², the ln q_m and β values, Equation (21), can be derived from the slope and the intercept, respectively. The mean free energy is potentially calculated by applying Equation (23).

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

(23)

if 8 < E < 16 kJ mol⁻¹, the procedure is a chemisorption, whereas if the E value is below 8 kJ mol⁻¹, the procedure is a physisorption.

2.8. Desorption Behavior

The removal of Cu(II) ions from the adsorbent was conducted via rinsing it with distilled water for eliminating any unabsorbed Cu(II) ions. Then, 10 mg of the adsorbent was submerged in HNO₃ acid (10 mL, 0.1 M) at 25 °C overnight. Equation (24) was employed to quantify the quantity of Cu(II) ions that were desorbed from the adsorbent.

% Cu(II) ions desorption = q_d/q_a × 100

(24)

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

2.9. Theoretical Calculation Details

All calculations were conducted using density functional theory (DFT) with the Gaussian 16 software package, employing the B3LYP functional and the LANL2DZ basis set. The total density of states (DOS) was analyzed using GaussSum software. The adsorption energy $\left({\mathrm{E}}_{\mathrm{a}\mathrm{d}\mathrm{s}}\right)$ was determined using Equation (25).

$${\mathrm{E}}_{\mathrm{a}\mathrm{d}\mathrm{s}}={\mathrm{E}}_{\mathrm{U}\mathrm{C}\mathrm{s}-\mathrm{C}\mathrm{u}}-{\mathrm{E}}_{\mathrm{U}\mathrm{C}\mathrm{s}}{-\mathrm{E}}_{\mathrm{C}\mathrm{u}}$$

(25)

where $\left({\mathrm{E}}_{\mathrm{U}\mathrm{C}\mathrm{s}-\mathrm{C}\mathrm{u}}\right)$ is the total energy of the UC adsorbents after the adsorption of Cu cations, $\left({\mathrm{E}}_{\mathrm{U}\mathrm{C}\mathrm{s}}\right)$ is the energy of the UC adsorbents, and $\left({\mathrm{E}}_{\mathrm{C}\mathrm{u}}\right)$ is the energy of the single Cu atom.

3. Results and Discussion

3.1. pH of Zero-Point Charge (pHzpc)

With specified conditions of aqueous solution composition and temperature, pHzpc is known as the pH at which the overall charges along the surface equal zero. This does not mean that no charges are located on the surface at pHpzc, but actually that an equalized number of negatively and positively charged ions exist on the surface. The quantity of the charges on the surface relies on the kinds and numbers of functional groups, as well as the pH of the media. The pHzpc plays a significant role in the characterization of the surface because it determines to which extent the adsorbent can potentially bind dangerous ions. This is owed to the fact that the adsorbent surface affords a total negative charge at pH > pHzpc, which makes cationic species adsorption more appropriate. On the contrary, if pH < pHzpc value, the surface of the adsorbent possesses a total positive charge (${-\mathrm{N}\mathrm{H}}_2^{+}$), which has a repulsion effect on cations. The results of the salt addition procedure revealed that the pHzpc value of the UC adsorbents was 5.4 (Figure 2). With increments in the pH, both the oxygen and nitrogen atoms on UC keep their lone pairs of electrons and consequently can effectively bind metal cations [44,45].

3.2. Adsorption Optimization

3.2.1. Influence of the Initial Cu(II) Ions Solution Concentrations

From Figure 3A, it can be noted that the capacity of adsorption of the UCs adsorbent ranges from 16.9 to 183.96 mg g⁻¹ when the initial concentrations of the Cu(II) ion aqueous solutions ranges between 20 and 200 mg L⁻¹. This may be ascribed to the fact that, at low Cu(II) ion concentrations, the quantity of Cu(II) ions and the total number of available adsorption sites differ slightly, resulting in a partial adsorption which does not depend on the initial concentration. At high Cu(II) ion concentrations, few sites of adsorption on the UCs adsorbent are attainable, so the adsorption relies on the initial concentrations. Thus, as the initial Cu(II) ion concentrations rise, the adsorption capacities increase [46,47]. Furthermore, the Cu(II) ions’ mass transfer resistance from the aqueous phase to the solid one could prevail over utilizing a higher concentration gradient as a drive power [48].

3.2.2. Impact of the Temperature

The results in Figure 3B indicate that the removal efficiency of Cu(II) ions using the UCs adsorbent at 25 °C reaches 81.6%, followed by a decrease to become 75.25% at 35 °C. This may be due to the physisorption process, in which the metal ions are removed by exothermic reaction. But when the temperature increases from 35 to 55 °C, the removal efficiency slightly rises from 75.25 to 75.81%, indicating that the removal endothermic process happened by chemisorption. The mobility of the Cu(II) ions usually increases as the temperature increases, thus, these ions possess an adequate energy for interaction with the active binding sites on the adsorbent. Additionally, the temperature increase results in a decrease in the viscosity of the solution and an increase in the rate at which the Cu(II) ions diffuse across the adsorbent’s outer boundary layer. This allows many Cu(II) ions to permeate into the porous structure of the adsorbent, enhancing its removal efficiency [49].

3.2.3. Influence of pH of the Medium

Since the charges on the adsorbent surface (either negative or positive) depend on the pH value of the medium, the latter is an important factor in the metal ions’ process. To evaluate the influence of the pH of the medium on the effectiveness of the UCs adsorbent in removing the Cu(II) ions from their solutions, the examined values of pH were confined to 1–6, as shown in Figure 3C, since the Cu(II) ions were precipitated as Cu(OH)₂ at pH ≥ 7 [50,51,52]. The results manifested that the removal effectiveness dropped from 96.93 to 76.75% when the pH values decreased from 6 to 1. This may be ascribed to the fact that increments in the acidity of the adsorption solution resulted in a contest between the Cu²⁺ and H⁺ to bind with the active sites on the UCs adsorbent. Owing to the smaller size of the H⁺ in comparison to that of Cu²⁺, the adsorption of H⁺ ions preponderated.

3.2.4. Effect of UCs Dose

To investigate the impact of changing the UCs dose (2, 5, and 10 mg), the removal efficiency of Cu(II) ions was studied when keeping other variables constant (the Cu(II) ion solution volume and concentration were 10 mL and 100 mg L⁻¹, respectively, with a temperature of 25 °C, pH of 4, stirring speed of 80 rpm, and contact time of 24 h), as illustrated in Figure 3D. It is worth mentioning that the Cu(II) ion removal efficiency increased from 75.5 to 81.6% as the dose of the UCs adsorbent increased from 2 to 10 mg. This may be ascribed to the increment in the adsorbent’s surface area with an increase in its dose, and consequently to the increment in the available commutable binding sites on its surface that were willing to uptake Cu(II) ions [53,54,55,56,57].

3.3. Adsorption Kinetics

The results of the adsorption kinetics of the Cu(II) ions by the UCs uptake using 50 mg of UCs, 50 mL of Cu(II) ion solution at a 100 mg L⁻¹ concentration, a pH of 6, 25 °C, and at a time of 5–180 min via applying the Elovich, intraparticle diffusion (the Weber–Morrison model), pseudo-second-order, and pseudo-first-order models are presented in Figure 4. Furthermore, the kinetic constants and the correlation coefficients (R²) are summarized in

Table 1. Correlating coefficients and kinetic model constants for Cu(II) ions adsorption by UCs adsorbent.

Kinetic Models	Parameters	UCs
qe·exp (mg g−1)		99.65
Pseudo-first-order model $\log \left({\mathrm{q}}_{\mathrm{e}}-{\mathrm{q}}_{\mathrm{t}}\right)=\log {\mathrm{q}}_{\mathrm{e}}-{\displaystyle \frac{{\mathrm{k}}_1}{2.303}}\mathrm{t}$	R2	0.752
	qe·cal (mg g−1)	1.948
	k1 (min−1)	0.009
Pseudo-second-order model ${\displaystyle \frac{\mathrm{t}}{{\mathrm{q}}_{\mathrm{t}}}}={\displaystyle \frac{1}{{\mathrm{k}}_2{\mathrm{q}}_{\mathrm{e}}^2}}+{\displaystyle \frac{\mathrm{t}}{{\mathrm{q}}_{\mathrm{e}}}}$ ${\mathrm{h}=\mathrm{k}}_2{\mathrm{q}}_{\mathrm{e}}^2$	R2	1
	qe·cal (mg g−1)	100
	k2 (10−5) (g mg−1 min−1)	0.011
	h (mg g−1 min−1)	96.15
Elovich ${\mathrm{q}}_{\mathrm{t}}={\displaystyle \frac{1}{\beta }}$ $\ln (\mathsf{\alpha }\beta )+{\displaystyle \frac{1}{\beta }}$ln t	R2	0.924
	β (g mg−1)	0.593
	α (1023) (mg g−1 min−1)	4.108
Intraparticle diffusion (the Weber–Morrison model) qt = (kint t1/2) + C	R2	0.840
	k (mg g−1 min−1/2)	0.483

. The results indicate a lack of compliance with the pseudo-first-order model, as evidenced by the low correlation coefficient, R², in the linear plot. This weak R² value indicates a weak correlation between the parameters and highlights that the adsorption mechanism does not adhere to the pseudo-first-order kinetics model. The highest R² value was acquired via applying the pseudo-second-order model (R² = 1.000) in comparison to the other used models (Table 1). This refers to the adsorption process precisely fitting the pseudo-second-order model. The determined rate constant, k₂, was 0.011 (g mg⁻¹ min⁻¹), and the initial adsorption rate, h, was 96.15 mg g⁻¹ min⁻¹. Moreover, the value of the experimentally obtained q_e (99.65 mg g⁻¹) was very close to the theoretically calculated one (100 mg g⁻¹), as given in Table 1, suggesting a robust consistency with the pseudo-second-order model. Thus, this model reflects and is perfectly valid to describe the process of the adsorption of Cu(II) ions onto the UCs adsorbent. The Elovich kinetic model is widely used for its ability to describe adsorption processes and heterogeneous adsorbent surfaces, making it a valuable tool for modeling a range of adsorption systems. The model’s higher R² value (0.924) supports its suitability to describe the process of the adsorption of Cu(II) ions onto UCs. Although the pseudo-second-order-model successfully describes the Cu(II) adsorption, it does not provide clear insights into the rate-limiting step. To address this, the Weber–Morrison model, which represents intraparticle diffusion, was utilized to determine whether film diffusion or intraparticle diffusion controlled the process. As shown in Figure 4 and Table 1, the intraparticle diffusion model for Cu(II) ions’ adsorption onto the UCs adsorbent had the lowest agreement with the experimental results (R² = 0.840).

3.4. Adsorption Isotherms

The outcomes of the isotherms of the adsorption of Cu(II) ions onto the UCs adsorbent, which was carried out using 10 mg of UCs adsorbent, pH 4, a 298 k temperature, an 80 rpm agitation velocity, and for 24 h, are represented in Figure 5.

Table 2. Parameters of different isotherm models for adsorption of Cu(II) ions by UCs adsorbent.

Models	Parameter	UCs
Langmuir ${\displaystyle \frac{{\mathrm{C}}_{\mathrm{e}}}{{\mathrm{q}}_{\mathrm{e}}}}={\displaystyle \frac{1}{{\mathrm{q}}_{\mathrm{m}\mathrm{a}\mathrm{x}}{\mathrm{K}}_{\mathrm{L}}}}+{\displaystyle \frac{{\mathrm{C}}_{\mathrm{e}}}{{\mathrm{q}}_{\mathrm{m}\mathrm{a}\mathrm{x}}}}$ ${\mathrm{R}}_{\mathrm{L}}={\displaystyle \frac{1}{(1+{\mathrm{K}}_{\mathrm{L}}{\mathrm{C}}_{\mathrm{o}})}}$	qmax (mg g−1)	133.33
	RL	0.126–0.590
	KL (L mg−1)	0.035
	R2	0.833
Freundlich $\ln {\mathrm{q}}_{\mathrm{e}}=\ln {\mathrm{K}}_{\mathrm{F}}+{\displaystyle \frac{1}{\mathrm{n}}}{\mathrm{l}\mathrm{n}\mathrm{C}}_{\mathrm{e}}$	1/n	1.427
	Kf (mg g−1)	3.016
	R2	0.960
Temkin qe = BTInKT + BTIn Ce	BT (kJ mol−1)	95.79
	KT (L g−1)	3.205
	R2	0.799
D-R Ln qe = ln (qm) − βε2 $\mathrm{E}=({\displaystyle \frac{1}{{(2\beta )}^{0.5}}})$	qm (mg g−1)	133.49
	ε	22.452
	β × 10−6	5
	kJ mol−1	0.316
	R2	0.796

clearly demonstrates the Langmuir isotherm’s inability to provide a satisfactory fit for the experimental data across all concentration ranges, as evidenced by the low R² value (0.833). Despite the Langmuir model’s successful application in describing the sorption of cadmium, chromium, and copper onto activated carbon [58] and zeolite [59], it fails to adequately explain the current experimental data. The model’s poor performance in representing the experimental data could be a result of its inability to account for the interactions between the adsorbate–adsorbate and the heterogeneous surface of the solid.

The Freundlich model results reveal a 1/n value above one, suggesting cooperative adsorption. Its high R² value (0.960) further supports the suitability of the Freundlich isotherm, which is applicable to multi-layered adsorption on heterogeneous solid surfaces. Consequently, this isotherm is valid for both weak van der Waals adsorption and strong chemisorption processes [60,61].

The Temkin model provides insights into the changing heat of adsorption and the interactions between adsorbate–adsorbate molecules as the adsorbed layer progresses. The correlation coefficient R² of 0.799 suggests a poor fit for the adsorption of Cu(II) ions with the UCs adsorbent.

The D–R isotherm does not fit well with the experimental data of Cu(II) ions’ adsorption (R² = 0.796). This indicates that the D–R model was not suitable for describing the sorption equilibrium of Cu(II) ions on UCs. The calculated E value (the mean free energy of Cu(II) ions’ adsorption by UCs adsorbent) was found to be 0.316. It is evident from the results that the E value is less than 8 kJ mol⁻¹, showing the essential features of physical adsorption [62].

Consequently, one can conclude that the adsorption is multi-layered and the Cu(II) ions are not distributed uniformly. Similar studies have documented the use of chitosan-coated superparamagnetic iron oxide nanoparticles for the removal of chromium [45], as well as the removal of Cr(VI) ions from water by using Γ-Fe₂O₃ nanoparticles enclosed in millimeter-sized magnetic chitosan beads [63].

3.5. Desorption Assessment

To boost the economic feasibility of UCs adsorbent, consecutive adsorption/desorption runs were performed. Subsequent to desorption, UCs adsorbent can be reutilized to minimize the likelihood to incur high costs to synthesize other new adsorbents. The results of the desorption of Cu(II) ions from UCs using nitric acid are represented in Figure 6. The percentages of Cu(II) ion desorption were 92.65, 75.29, 49.29, and 19.92% at four successive cycles. These results show the possibility to recycle UCs as an adsorbent for the removal of Cu(II) ions from aquatic media.

3.6. Adsorption Mechanism

To further unveil the adsorption mechanism between uracil-modified chitosan (UCs) and Cu²⁺ ions, a series of adsorption properties was explored using density functional theory (DFT) calculations [64,65,66,67,68]. These studies provided deeper insights into the interaction sites and energetic stability of the adsorption process. A uracil-modified chitosan structure was constructed (Figure 7a), representing a promising functional material for capturing metal ions. Two primary adsorption sites were identified, each contributing uniquely to the binding mechanism and stability.

The first adsorption site, denoted as UCs–Cu(I) (Figure 7b), involves the interaction of Cu(II) ions with the oxygen atoms in the glucosamine ring and the neighboring hydroxyl groups. This interaction highlights the intrinsic affinity of the chitosan backbone for metal coordination, emphasizing the role of its hydroxyl and amino functionalities. Such interactions are often associated with hydrogen bonding and electrostatic attractions, which collectively enhance the stability of the adsorbed state.

The second adsorption site, denoted as UCs–Cu(II) (Figure 7c), involves the coordination of Cu(II) with the nitrogen atom of the uracil ring and a neighboring oxygen atom. This site reflects the contribution of the uracil modification, which introduces additional active binding sites. The nitrogen in the uracil ring, being electron-rich, readily forms coordination bonds with Cu(II) ions, thereby enhancing the adsorption capacity of the modified chitosan. The neighboring oxygen atom further stabilizes the adsorption through complementary interactions, likely involving orbital overlap and charge transfer.

These findings underscore the significance of uracil modification in improving the adsorption properties of chitosan. The modification not only increases the number of active sites, but also enhances the strength of the interactions with Cu²⁺ ions. By analyzing the electronic and geometric structures of UCs–Cu(I) and UCs–Cu(II), it becomes evident that the interplay between the chitosan backbone and uracil moieties provides a synergistic effect, resulting in an improved adsorption performance. These observations highlight the potential of uracil-modified chitosan for applications in metal ions’ removal, catalysis, and environmental remediation.

For form I (Figure 7b), the adsorption energy between UCs and Cu²⁺ was calculated to be −0.53 eV, while form II (Figure 7c) exhibited a lower adsorption energy of −0.61 eV. The more negative adsorption energy for form II indicates a more stable configuration, making it thermodynamically favored and easier to form compared to form I. These results highlight the strong affinity of uracil-modified chitosan for Cu²⁺ ions, as reflected in the negative adsorption energies, which signify spontaneous adsorption processes. The difference in adsorption energies between the two configurations further emphasizes the role of specific interaction sites in determining the stability of the adsorbed complexes.

The incorporation of uracil into chitosan significantly enhances the material’s adsorption capacity by introducing additional active sites. Uracil modification provides nitrogen and oxygen atoms with high electron densities, which contribute to a stronger coordination with Cu²⁺ ions. This modification not only stabilizes the adsorption process, but also increases the binding versatility of the chitosan, allowing it to capture Cu²⁺ ions more effectively. These findings suggest that the uracil-modified chitosan structure could be an excellent candidate for applications requiring a high metal ion adsorption efficiency, such as water purification and heavy metals recovery.

To further understand the adsorption mechanism at the electronic level, the density of states (DOS) of UCs, UCs–Cu(I), and UCs–Cu(II) were analyzed (Figure 7d, Figure 7e, and Figure 7f, respectively). The total DOS revealed significant electron orbital hybridization between the adsorbed Cu ions and the oxygen atoms of the UCs structure. This hybridization indicates strong covalent and electrostatic interactions, which are essential for the stability of adsorbed complexes. In particular, the DOS analysis of UCs–Cu(II) (Figure 7f) showed more pronounced hybridization peaks compared to UCs–Cu(I) (Figure 7e), further corroborating the higher stability of form II.

Moreover, the electronic structure changes upon Cu²⁺ adsorption illustrate the redistribution of electron density, particularly around the adsorption sites. This redistribution enhances the material’s electronic and chemical properties, enabling it to act as an efficient adsorbent for Cu²⁺ ions. These insights into the electronic interactions provide a deeper understanding of the adsorption mechanism, highlighting the synergistic effects of the uracil incorporation and the chitosan backbone.

3.7. Comparison of Different Adsorbents’ Capacities for Adsorbing Cu(II) Ions

To remove Cu(II) ions, a wide variety of adsorbents were reported in some earlier works [33,64,65,66,67,68,69,70,71], as shown in

Table 3. The maximum capacities of adsorption (q_max) of various reported Cu(II) ion adsorbents.

Adsorbent	qmax (mg g−1)	Temperature (°C)	Cu(II) Ions 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	[64]
Horn core calcined at 400 °C	99.98	25	100–500	0.02	5	[65]
Magnetic chitosan@bismuth tungstate coated by silver	181.8	20–40	20–120	0.02	6	[66]
Xanthate-modified magnetic chitosan	34.5	25	100	-	5	[67]
Epichlorohydrin cross-linked xanthate chitosan	43.47	50	100	-	5	[68]
Medicinal plant (Salvadora persica)	74.30	25	100	-	4	[69]
Pecan nutshell (Carya illinoinensis)	23.37	30	-	-	5	[70]
H1	96.20	25	100	0.01	6	[71]
H2	97.59	25	100	0.01	6	[71]
UCs	99.65	25	100	0.01	6	Present study

. It is obvious that the capacity for the adsorption of Cu(II) ions by UCs is higher or lower than that of the other reported options.

In addition to the data listed in Table 3, the adsorption data of Cu(II) ions onto chitosan-grafted-polyaniline fitted to the pseudo-second-order model. The regeneration efficiency and removal efficiency of the regenerated adsorbent were 97.7 and 90.4%, respectively [33]. The adsorption mechanism of a quaternized chitosan cationic polyelectrolyte microsphere unveiled that the cation exchange, chelating effect, proton exchange, and complex formation mainly contributed to the ultrafast and ultra-efficient removal of heavy metal ions, in addition to its superb recyclability (within 10% loss after five cycles), antibacterial viability, and biodegradability (165 days) [64]. The adsorption process of calcined horn cores was identified as a cationic exchange between Ca(II) from bones and Cu(II) in solution and followed the pseudo-second-order and Langmuir models. This was spontaneous and endothermic [65]. Magnetic chitosan@bismuth tungstate coated by silver demonstrated the excellent adsorption/photocatalytic removal of Cu(II) in aqueous solution. It was found that illumination could enhance the adsorption of Cu(II) by this adsorbent. Meanwhile, the composite exhibited a desirable adsorption ability of Cu(II) after five cycles. The adsorption fitted well with the pseudo-second-order kinetic model and its isotherm followed the Freundlich model [66]. Xanthate-modified magnetic chitosan showed that the adsorption of Cu(II) followed the Langmuir mode. The thiol and amino groups participated in the adsorption of Cu(II) [67]. The adsorption of Cu(II) by epichlorohydrin cross-linked xanthate chitosan strongly depends on pH and temperature. Its adsorption capacity increased with an increasing temperature, indicating the endothermic nature of the adsorption process. The adsorption process was found to follow the pseudo-second-order and Langmuir models [68]. The Langmuir model fits better to the experimental data of copper sorption onto activated carbon prepared from Salvadora persica [69].

4. Conclusions

In this study, Cu(II) ions were adsorbed utilizing UCs adsorbent for industrial wastewater remediation. The UCs adsorbent were characterized by a distinguished removal performance of Cu(II) ions, reaching a maximum (99.65%) at a Cu(II) ions concentration of 200 mg L⁻¹, UCs dose of 10 mg, pH 6, and a temperature of 25 °C. The adsorption of Cu(II) ions by the UCs adsorbent obeyed both the pseudo-second-order kinetic and Freundlich isotherm models, indicating the multi-layer nature of the adsorption process. There are several advantages of these UCs adsorbent, included forbidding chitosan from solubility in acidic adsorption media, incorporating additional functional groups such as -OH, -NH, C=O, and -N-CH₃, which can act as active sites to bind Cu(II) ions, and banning the predictable decrease in the adsorption of capacity of Cu(II) ions, which usually results from the consumption of the amino groups of chitosan during its modification process. In addition, UCs adsorbent possesses the prospect to be employed as a promising efficient, re-utilizable adsorbent for removing Cu(II) ions. The synthesis of UCs adsorbent represents a credible approach to create novel adsorbents suitable for removing heavy metal ions to remedy industrial wastewater.