The effective adsorption of arsenic from polluted water using modified Halloysite nanoclay

2.1 Characterization

The morphology of the HS and MHS nanoclay was examined using a JEOL JEM-1230 transmission electron microscope (TEM). An X-pert pro diffractometer from Philips was used to obtain X-ray diffraction (XRD) patterns. The automatic gas sorption system A NOVA 3200e was used to measure the specific surface area (Quantachrome, Boynton Beach, FL, USA).

2.2 Adsorption experiment

The adsorption experiment has carried out in a water samples collected from the polluted water located at Al-Arbaeen lagoon. Al-Arbaeen lagoon is situated on the Red Sea shoreline in Jeddah. With two loops spanning nearly N—S and a network of contacts connecting the interior basins to the open sea, Al Arbaeen lagoon features a more intricate T-shape. El Rayis and Moammar (1998) assert that there is a continuous interchange of water with the open sea at the top two meters of the surface layer. On rare occasions, bottom water will regenerate in rough and stormy conditions (Orif et al, 2017).

To determine the As (III) adsorption capability on Halloysite nanoclay (HS) and 8-hydroxyquinoline-modified HS nanoclay, arsenic adsorption studies were carried out (MHS). In 50 mL centrifuge tubes with 20 mL of an As (III) solution and a predetermined volume of HS or MHS, several experiments were conducted. 0.1 M HCl or 0.1 M NaOH were used to change the pH of the mixture. On a shaker spinning at 350 rpm, all the solutions were mechanically stirred. The initial As (III) solution's pH was not changed during the adsorption kinetics and adsorbent dose studies in order to represent a system with no outside influences, and it was fixed to be at a pH of 3. Each experiment was carried out at room temperature. To ascertain the equilibrium contact duration and maximum adsorption capacity, kinetic studies were carried out at intervals ranging from 0 to 150 min. Through testing several adsorbent dosages between 0.25 and 1.25g, the most effective adsorbent dosage was found. At pH levels ranging from 3 to 9, experiments studying how pH affects adsorption capacity were carried out. After the pH studies, initial As (III) concentration trials were conducted with As (III) initial concentrations ranging from 0.25 mg/L to 10 mg/L at a pH of 8 Then the resultant solution is filtered using the filter paper and the filtrate were tested using ICP-AES.

The adsorption capacity, q_t, at a specific time t and the percent removal of arsenite were calculated based on the following equations:

2.3 Determination of point of zero charge

This study used the salt addition method to estimate the point of zero charge (PZC) of the adsorbent (Bakatula et al. 2018; Rahman and Varshney, 2021). For this, 25 mL of 0.1 M NaNO3 in different pH (3–10) solutions were taken in volumetric flasks, and 0.2 g of prepared adsorbent was then added. By adding 0.1 M HCl or NaOH solutions, the pH of the liquids was regulated. After 24 h, the pH (pHf) of the solutions was assessed. Plotting the initial pH (pHi) readings against the ΔpH (pHi - pHf) values were used to calculate the PZC.

3.1 Characterization

The Fig. 1 shows the TEM characterization of the MHS nanoclay. It is evident that the MHS was made up of a hollow tube-like structure. The TEM image presented in Fig. 1 revealed that the modified HS nanoclay have diameters of 40–60 nm, lumens in the range of 10–15 nm, and lengths within the range of 100–900 nm. The images also show the cross aggregates between the nanotubes brought on by static electric charges. It is evident from the surface morphology that it is a lamellar material made up of various arrays of cleaved sheets (Thommes, 2010; Zhu et al, 2016). An increase in the contact area and active sites for the adsorption of heavy metals from effluents was eventually achieved with the aid of the modification of HS nanoclay, which favors the creation of many cavities and a rather loose surface (Zhang and Wang, 2015). The EDX pattern of MHS nanoclay is shown in the Fig. 2 and the atomic % of each element is listed in the Table 1. It can be seen that the EDX of MHS show the presence of aluminum, silicon, carbon, nitrogen and oxygen. It is very clear from the Fig. 2 that synthesized sample are free from other elemental impurities. The reflection peaks at diffraction angles of 2 = 11.8°, 19.84°, 24.53°, 35.01°, 36.68°, and 38.05°, which correspond to the reflection planes (0 0 1), (1 1 0), (0 0 2), (1 2 2), (2 0 0), and (1 3 1), are explained by the XRD pattern of MHS nanoclay in Fig. 3 (Senoussi et al, 2016). The modified halloysite nanoclay was found to have a basic spacing (0 0 1) of 7.49 (2 = 11.81°).

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Fig. 1

TEM characterization of the MHS nanoclay.

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Fig. 2

The EDX pattern of MHS.

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Fig. 3

XRD pattern of MHS nanoclay.

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By using FTIR analysis, the surface functional groups of MHS nanoclay were described. The Fig. 4 shows the FTIR spectra of the pristine and modified HS nanoclay. Both spectra had the characteristic peaks of the HS nanoclay at 3696 cm − 1, 3629 cm − 1, and 1654 cm − 1 and are assigned to the O—H stretching of the inner-surface hydroxyl groups, the O—H stretching of inner hydroxyl groups, and deformation of water, respectively. In addition, the presence of the strong –OH absorption band at 3,433 cm-1 was more pronounce for the modified HS nanoclay due to the presence of 8-HQ at the nanoclay. Also, a band was observed at 1,730 cm-1 for the modified HS nanoclay may be attributed to conjugated C⚌O stretching vibrations due to the presence of 8-HQ at the nanoclay surface. Also, the bands appears at 1763 cm − 1 to 1458 cm − 1, 1384 cm − 1, 814 cm − 1, and 671 cm − 1 correspond to the characteristic quinoline stretching vibrations (Weerasuriya et al, 2017; Shao et al, 2011; Kosa et al, 2012). Also the vibrations in the ranges 1384 cm − 1 to 1465 cm − 1 and 1095 cm − 1 to 1280 cm − 1 correspond to the aromatic amine resonance frequencies for C—N bond and C—O stretching frequencies, respectively (Weerasuriya et al, 2017; Shao et al, 2011; Kosa et al, 2012). According to Fig. 5, which categorizes the N₂ adsorption/desorption isotherms for MHS as type III isotherms, the specific surface (BET) is 72.5 m²/g. The smaller mesopores and higher surface area of HS permitted capillary condensation.

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Fig. 4

FTIR spectrum of the pristine and modified HS nanoclay.

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Fig. 5

BET surface area of Modified Halloysite nanoclay at 77 K.

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3.2 Effect of dosage

Fig. 6 for HS and MHS at various loadings illustrates the impact of the adsorbent dosage on the removal of As (III). Fig. 6 shows that 1 g/L of the modified HS dosage is enough to allow for the quantitative adsorption of As (III). With increasing modified HS dosage, more As (III) is eliminated. There are more active sites available for adsorption due to the rise in the percentage caused by the increase in the amount of existing solid phase (Kavil et al, 2018; Abdel-Fadeel et al, 2022). It is obvious that adsorption onto MHS is caused by the 8-hydroxyquinoline species, which offer complementary adsorption sites for As (III) species, because adsorption on HS is far less significant than that on MHS. The greater adsorbent surface and more available metal ion adsorption sites are the likely causes of the increase in adsorption with adsorbent dosage (Maleki et al, 2015). The overlapping or aggregation of adsorption sites, which lengthens the diffusion path and reduces the overall adsorbent surface area that is available to the metal ions, may be to blame if the adsorption capacity reduced beyond the limit of adsorbent dosage (Crini et al, 2008). In the present study the equilibrium was reached with a dosage of 1 g/L.

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Fig. 6

The effect of HS and MHS dosage on the removal of As (III) at pH 8; Initial concentration 5 ppm; Contact time 90 min.

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3.3 Effect of contact time

Equilibrium between the adsorbent and the solution's solute must be attained for adsorption to be fully effective. Therefore, a specific amount of time is needed for the equilibrium's contacts in order to guarantee that adsorption is attained. The contact time is the length of time needed to reach equilibrium. Numerous writers have investigated how contact time affects adsorption. Their findings demonstrated that the rate of metal ion adsorption rises over time, reaching an optimal value at which point no further metal ion removal occurs (Azouaou et al, 2010; Martinez et al, 2006; Montazer-Rahmati et al, 2011). The Fig. 7 shows the effect of contact time on the adsorption of As (III) using HS and MHS nanoclay. The present study achieved an adsorption efficiency of 83.6 % with the MHS nanoclay over the contact time of 150 min. During this experiment the parameters like pH, the dosage of adsorbent and the temperature were kept constant. The adsorption up to 90 min were quite fast and afterwards it was showing a slight elevation in the adsorption for both HS and MHS nanoclays. The maximal adsorption capacity attained by the adsorbent under the operating conditions is reflected in the quantity of metal ions adsorbed at the time of contact. Adsorption is observed to be nearly constant beyond the equilibrium time, which is explained by the availability of empty surface sites during the initial stages of adsorption and the difficulty of filling the empty sites after equilibrium has been reached due to the repellent forces between the heavy metal molecules on the HS/MHS and bulk solution (Alley, 2007). The similar observations were noted in elsewhere (Selomulya et al, 1999; Attia et al, 2010; Brahmaiah et al, 2015).

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Fig. 7

The effect of contact time on the adsorption of As (III) using HS and MHS nanoclay at pH 8;; Initial concentration 5 ppm; Dosage 1 g/L.

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3.4 Effect of pH

In the process of heavy metal adsorption, the pH level of the metal solution is crucial. It makes a significant adsorption capacity contribution. The aqueous solution's behavior and the surface attachment locations are both impacted by pH (Ushakumari and Madhu, 2013). The adsorbent surface charge has an impact on pH. The pH affects the surface charge, ionization potential, and distribution of metal ions in the adsorbent (Park et al, 2010). The 8-hydroxyquinoline (8-HQ), also known as 8-quinolinol or oxine, is a well-known monoprotic, bidentate chelating agent. It is a compound that contains an oxygen donor atom and a nitrogen donor atom that can both bind to metal atoms. When the pH of an adsorbent surface is reduced, the H + ions that are present in high numbers neutralize the negatively charged surface, reducing the barrier to diffusion and increasing the rate of adsorption. The effect of solution pH on this study is shown in the Fig. 8. The study was done in a range of pH between 3 and 9. In comparison to acidic medium, overall heavy metal removal values in alkali solutions were significantly greater. Fig. 8 illustrates this phenomenon, showing that metal ion adsorption is minimal in the acidic zone (pH = 3), increases with rising pH, and reaches its highest value at pH 8. These findings suggest that the ability of the material to adsorb metal ions was lowered by protonation of more basic pyridine-type nitrogen in the quinoline rings. The high concentration of hydroxyl ions (HO) in the solution may be the cause of the decrease in metal ion absorption at pH levels higher than 8 (Ghaemy et al, 2015). Lowering the positive charges of amino groups reduces the pH values, which enhances adsorption behavior (Maleki et al, 2015). In line with prior arsenite adsorption experiments (Yuh-Shan, 2004; Ho and Mckay, 1998; Weber and Morris, 1963), the current investigation found that As (III) adsorption increased with increasing pH up to pH 9. At a pH less than 8, attraction and repulsion are ineffective against neutral H₃AsO₃ because the adsorption process may be regulated by H₃AsO₃ surface complexation that causes deprotonation or dissociation (Bhowmick et al, 2014). Adsorption occurs most frequently in the neighborhood of pH 8. Moreover, the point of zero charge of the adsorbent was found to be 8 (Fig. 9). The balancing effect of the As (III) solution at high pH caused the reaction mixture's pH to rise above 8, which likely caused a decrease in adsorption and the onset of precipitation/coagulation (Tandon et al, 2013).

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Fig. 8

The effect of solution pH on the adsorption of As (III) using HS and MHS nanoclay at Initial concentration 5 ppm; Dosage 1 g/L; Contact time 90 min.

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Fig. 9

Determination of point of zero charge of the adsorbent.

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3.5 Effect of initial concentration

The initial concentration of the solution affects the rate of adsorption. Adsorption experiments from the past have demonstrated that at lower concentrations, the rate of adsorption increases (Shaban et al, 2019; Alfarawati et al, 2020; Kavil et al, 2020). The deposit of metal at the interface between the solid and liquid phases is the broad definition of the mass transfer process known as metal adsorption. In order to investigate this, the initial metal concentration was varied between 0.25, 1, 3, 5, 7, and 10 ppm of As (III). The effect of initial concentration for the adsorption of As (III) in the present study is shown in the Fig. 10. It is clearly observed that the lower concentration was adsorbed quickly compared to the higher level of As (III). This is due to the fact that most metal ions remained unadsorbed when the concentration was high due to saturation of the adsorption sites (Azouaou et al, 2010). Because there was more adsorbent surface area available at lower concentrations, the rate of adsorption increased (Kumar et al, 2010).

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Fig. 10

The effect of initial concentration for the adsorption of As (III) at pH 8; Dosage 1 g/L; Contact time 90 min.

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3.6 Adsorption kinetics

Adsorption rate predictions offer crucial details about adsorption mechanisms. The experimental data at different adsorption durations corresponding to the changes in adsorption capacity were fitted using the pseudo-first-order model to assess the adsorption mechanism onto MHS. The following equations was used to determine the pseudo-first-order model, $$\frac{{\mathit{dq}}_t}{\mathit{dt}}=k_2{\left(q_e-q_t\right)}^2$$ $$\frac{t}{q_t}=\frac{1}{k_2{q}_t^2}+\frac{t}{q_e}$$

Where q_e (mg/g) is the quantity of As (III) adsorbed at equilibrium, q_t (mg/g) is the amount of As (III) adsorbed at a specific time t (min), and k₂ (g/mg min-1) is the pseudo-second-order reaction rate. The terms k₂ and q_e can be calculated from the intercept and slope of the (t/q_t) vs t plot, respectively.

In order to create the proper adsorbent materials for environmental remediation, it is helpful to investigate the process kinetically in order to develop an adequate mathematical model for a better description of the interactions between pollutants and solid adsorbents. The experimental removal capabilities, or the quantity of metal ions adsorbed from the desalination outfall sample by MHS nanoclay (qt), are shown to vary with interaction time in Fig. 11. The adsorption/removal of As (III) by MHS nanoclay was discovered to reach equilibrium within 90 min, and further extending the contact period had little effect on the removal capabilities. The Lagergren pseudo-first-order (PFO) kinetic model was used to kinetically analyze the adsorption experiment results displayed in Fig. 12. The current data demonstrate the superior values of the PSO kinetic models' regression coefficients when compared to the PFO kinetic models, demonstrating the suitability of the PSO kinetic model for explaining the removal of As (III). The pattern is shown in the Fig. 13. The adsorption process was also partially modeled by the Elovich model, as shown in Fig. 13, which also shows that the adsorption sites were heterogeneous and that a variety of activation energies occurred during the adsorption process. Given that this model had the highest R², a comparison of the R² values clearly demonstrates that the adsorption kinetics followed this model (Rahman and Haseen, 2014; Rahman et al, 2021). The MHS was found to fit well for the entire adsorption period, as evidenced by the very high correlation coefficient, as shown in the pseudo-second-order model fit in Fig. 13. Fig. 13′s calculated equilibrium capacity was not showing much variations to that determined by experimental investigation (Table 2).

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Fig. 11

The quantity of metal ions adsorbed from the desalination outfall sample by MHS nanoclay (qt).

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Fig. 12

The Lagergren pseudo-first-order (PFO) kinetic model.

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Fig. 13

The Pseudo Second Order (PSO) kinetic model (Elovich model) List of Tables.

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3.7 Isotherm studies

The isotherm study shows the efficiency of the interaction between the biosorbent and adsorbate as well as how metal ions can be dispersed across the liquid and solid phases at different equilibrium concentrations (Maleki et al, 2015). The experimental data have been described using a variety of isotherm models, including Langmuir and Freundlich, to improve the design of the sorption system for the removal of metal ions from solutions.

In the Langmuir model, a single layer of adsorbed solute at a constant temperature is assumed together with homogeneous sorption energy. The most often used isotherm in the Langmuir model, the monolayer sorption isotherm, is expressed as follows (McRae et al, 2008): $$q_{e=\left(\frac{C_e\times Q_0\times K_L}{1+K_L{C}_e}\right)}$$

Where q_e, C_e, Q₀, and K_L stand for, respectively, the quantity of solute adsorbed at equilibrium (mg/g), the concentration of adsorbate at equilibrium (mg/L), the maximal adsorption capacity (mg/g), and the Langmuir constant (L/mg). The Langmuir equation's linear form is $$\frac{C_e}{q_e}=\frac{1}{K_L{Q}_e}+\frac{C_e}{Q_0}$$

A linear plot of C_e/q_e against C_e was created to examine the Langmuir isotherm model's applicability for the metals adsorption onto MHS. The higher R² values demonstrated the Langmuir isotherm's suitability for arsenic adsorption upon MHS (Turk et al, 2020). The values of Q₀, K_L, and R² (correlation coefficient) are displayed in Table3.

The Freundlich isotherm is employed in the adsorption process on heterogeneous surfaces (Akin et al.2012). These sources provide the Freundlich isotherm equation: $$\log Q_e=log{K}_F+\frac{1}{n}log{C}_e$$

Where n is the adsorption intensity and KF is the Freundlich constant (Turk, 2017). The slope and intercept of the logQ_e against logC_e linear plot are 1/n and logK_F, respectively. From linear graphs, the Freundlich isotherm parameters in Table 3 were derived. In these models, the correlation coefficient (R²) was 0.98. According to Chetia et al. (2012), the KF and n values for As (III) adsorption on MHS were determined to be 94.75 and 1.33, respectively (if 1/n is less than 1, adsorption is appropriate). Arsenic adsorption on MHS is better suited by the Freundlich model, which has a higher correlation coefficient (R² [0.98]), however the difference between the R² Langmuir adsorption is very low.

3.8 Thermodynamic studies

The actual indicators for the practical application of an adsorption process are thermodynamic parameters like the Gibbs energy (ΔG), enthalpy (ΔH), and entropy (ΔS). Using the following equations [30, 37], the thermodynamic parameters were calculated: $$\mathrm{\Delta }\mathrm{G}=\mathrm{\Delta }\mathrm{H}-\mathrm{T}\mathrm{\Delta }\mathrm{S}$$ $$K_c=\frac{q_e}{C_e}$$ $$\mathit{ln}{K}_c=\frac{\mathrm{\Delta }\mathrm{S}}{R}-\frac{\mathrm{\Delta }\mathrm{H}}{\mathit{RT}}$$

Where Kc, qe, Ce, R, and T are the equilibrium constants, the equilibrium quantity of metal adsorbed on the adsorbent (mol/g), the equilibrium concentration of metal in the solution (mol/L), the gas constant (8.314 J/mol K), and the absolute temperature (K), respectively. The values of ΔH and ΔS can be inferred from the slopes and intercepts of the graph of lnKc vs 1/T. Table 4 contains the calculated thermodynamic parameters. The endothermic reaction and enhanced randomization at the solid/solution interface during the adsorption of arsenic onto MHS are suggested by the positive ΔH and ΔS values. The adsorption process is implied to be spontaneous by the negative ΔG values. Additionally, the fact that ΔG values drop as temperature rises suggests that adsorption is more spontaneous at higher temperatures (Maleki et al, 2015). Additionally, MHS adsorbents' ΔG° values above −20 kJ/mol, indicating that chemisorption was the mechanism (Horsfall Jr et al. 2006).

3.9 Comparison with other adsorbents

Due to variations in experimental circumstances and methods, it is challenging to make a direct and precise comparison between the prepared adsorbent and other studies in the literature. Table 5 compares the adsorption capacity of the developed MHS with that of other modified clay adsorbents and adsorbents with related chemical characteristics. As can be shown, compared to certain other less expensive adsorbents including granular ferric hydroxide, fungal biomass, and iron oxide coated sand, the MHS adsorbent material exhibits superior removal efficiency. Some of the described clays, such iron- and manganese-pillared MMT, had much lower Langmuir adsorption capabilities. The current material exhibits a comparatively higher adsorption rate and performs best at a pH that is somewhat alkaline, which is characteristic of waters that have been contaminated with arsenic (Almasri et al, 2018).