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ISSN : 1225-0112(Print)
ISSN : 2288-4505(Online)
Applied Chemistry for Engineering Vol.34 No.5 pp.556-565

Glutamic Acid-Grafted Metal-Organic Framework: Preparation, Characterization, and Heavy Metal Ion Removal Studies

Phani Brahma Somayajulu Rallapalli, Jeong Hyub Ha†
Department of Integrated Environmental Systems, Pyeongtaek University, Pyeongtaek 17869, Republic of Korea
Corresponding Author: Pyeongtaek University Department of Integrated Environmental Systems, Pyeongtaek 17869, Republic
of Korea Tel: +82-31-659-8309 e-mail:
August 16, 2023 ; September 12, 2023 ; September 13, 2023


Fast industrial and agricultural expansion result in the production of heavy metal ions (HMIs). These are exceedingly hazardous to both humans and the environment, and the necessity to eliminate them from aqueous systems prompts the development of novel materials. In the present study, a UIO-66 (COOH)2 metal-organic framework (MOF) containing free carboxylic acid groups was post-synthetically modified with L-glutamic acid via the solid-solid reaction route. Pristine and glutamic acid-treated MOF materials were characterized in detail using several physicochemical techniques. Single-ion batch adsorption studies of Pb(II) and Hg(II) ions were carried out using pristine as well as amino acid-modified MOFs. We further examined parameters that influence removal efficiency, such as the initial concentration and contact time. The bare MOF had a higher ion adsorption capacity for Pb(II) (261.87 mg/g) than for Hg(II) ions (10.54 mg/g) at an initial concentration of 150 ppm. In contrast, an increased Hg(II) ion adsorption capacity was observed for the glutamic acid-modified MOF (80.6 mg/g) as compared to the bare MOF. The Hg(II) ion adsorption capacity increased by almost 87% after modification with glutamic acid. Fitting results of isotherm and kinetic data models indicated that the adsorption of Pb(II) on both pristine and glutamic acid-modified MOFs was due to surface complexation of Pb(II) ions with available –COOH groups (pyromellitic acid). Adsorption of Hg(II) on the glutamic acid-modified MOF was attributed to chelation, in which glutamic acid grafted onto the surface of the MOF formed chelates with Hg(II) ions.


    1. Introduction

    The exacerbation of environmental pollution and the degradation of ecosystems, primarily attributed to industrial activities, have posed significant threats to human existence and development from an anthropological perspective[1]. A diverse range of toxic, inorganic and organic chemicals are released into the environment in the form of industrial waste, leading to significant water pollution. Industrial waste typically comprises a diverse array of pollutants, including heavy metal ions, dyes, inorganic anions, and pesticides, which exert toxicity towards various living organisms[2,3].

    Heavy metal pollution of water resources is particularly strongly associated with human survival[4]. Heavy metal pollution arises from both direct and indirect sources. Direct sources include the discharge of effluents from industrial facilities, refineries, and waste treatment plants. Indirect sources encompass the introduction of contaminants into the water supply through soil and groundwater systems, as well as the deposition of pollutants from the atmosphere through rainfall[5,6]. Unfortunately, numerous aquatic ecosystems are confronted with metal concentrations that surpass established safety criteria aimed at safeguarding the environment, as well as the well-being of animals and humans. Heavy metal cations are significant contributors as trace elements in biochemical reactions, and certain heavy metal cations are indispensable for various organisms. However, the presence of heavy metal ions poses significant risks to human beings, animals, and plants due to their propensity for accumulation (at concentration thresholds specific to each metal), inherent toxicity, non-biodegradability, and ability to induce a range of diseases and disorders[7].

    A range of methodologies have been utilized to address heavy metal contamination in diverse wastewater sources. These methodologies commonly involve chemical precipitation, membrane filtration, electrochemical treatment, solvent extraction, ion exchange, and adsorption [8]. Chemical precipitation and electrochemical treatment methods have been found to be ineffective, particularly when the concentration of metal ions in an aqueous solution is within the range of 1 to 100 mg/L. Furthermore, these methods result in the generation of a significant amount of sludge, which poses challenges in terms of treatment and disposal. Membrane filtration, solvent extraction, and ion exchange techniques are cost-ineffective for the treatment of substantial volumes of wastewater that contain low concentrations of heavy metals. Consequently, their application on large scales is unfeasible[9]. Most of the above-mentioned techniques exhibit certain drawbacks, including intricate treatment procedures, elevated expenses, and energy consumption. Adsorption has been regarded as an optimal alternative to these methods due to its operational simplicity and cost-effectiveness[ 10]. The primary characteristics of the adsorbent typically encompass a strong affinity for certain ions and a substantial loading capacity, particularly in relation to the elimination of heavy metals. Furthermore, adsorption is reversible, allowing for the regeneration of the adsorbent through an appropriate desorption process[11]. To date, extensive research has been conducted on various adsorbents, including activated carbon, metal-organic frameworks (MOFs), zeolites, and clay, to explore their efficacy in eliminating heavy metal ions[12-15].

    MOFs are porous crystalline materials that are synthesized through the coordination of metal ions with organic ligands. One notable distinction between MOFs and other porous inorganic materials, such as zeolites, lies in the composition of MOFs, which consists of organic and inorganic components that are highly adjustable. By altering metal ions and organic ligands, it is possible to customize the framework's topology and chemical functionality[16]. In addition, the organic and inorganic constituents of MOFs facilitate post-synthetic modification through the formation of covalent and coordinate covalent bonds, respectively. MOFs have thus been extensively investigated for a wide range of applications, such as gas storage, gas separation, energy storage and conversion, catalysis, drug delivery, and sensor applications [17-22]. Particularly, the UiO series of zirconium-based MOFs have gained significant recognition due to their exceptional chemical stability. However, it has been observed that various benchmark MOFs, including Cu-BTC, MOF-5, and MOF-177, exhibit instability when exposed to water[26-28].

    Various chemical agents have been employed to date to modify UiO-66 MOFs with the objective of eliminating HMIs from water[29]. One potential concern associated with the utilization of chemical agents is the possibility of inducing significant environmental pollution. Furthermore, it is typically necessary to implement a costly treatment system to manage excess chemicals[30]. It is thus imperative to employ a straightforward approach to functionalize UiO-66 MOFs using reagents that are both non-toxic and biodegradable. Amino acids are such types of reagents that serve as the fundamental constituents of proteins, with certain variants being readily accessible in the market at cost-effective rates[31]. There have been reports on the modification of clays using amino acids[32, 33], while amino acid-modified MOFs have not been reported.

    In this paper, a UiO-66(COOH)2 MOF that contains free carboxylic acid groups was post-synthetically modified with L-glutamic acid. The anticipated reaction involves the interaction between the amino group (-NH2) of the amino acid and the free carboxyl group (-COOH) of the MOF. As a result, the amino acid will be attached to the MOF through the formation of an amide (-CONH-) bond. The O-ligand from carboxylic acid groups and N-ligand of the amino acid function as chelating agents, facilitating the binding of heavy metal ions. Pristine and glutamic acid-treated MOF materials were characterized in detail with several physicochemical techniques. We performed single-ion batch adsorption studies of Pb(II) and Hg(II) ions using pristine as well as amino acid-modified MOFs. Parameters influencing the removal efficiency, such as the initial concentration and contact time, were examined as well. Finally, we tested the applicability of various isotherm and kinetic models.

    2. Materials & methods

    2.1. Materials

    Zr(SO4)2 (extra pure, 98.0%), L-glutamic acid, PbCl2 (extra pure, 98.0%), and HgCl2, (special grade, 99.5%), were purchased from Samchun Chemicals Ltd., South Korea. Pyromellitic acid (ACS reagent, 99.0%) was purchased from Sigma Aldrich.

    2.2. Methods

    2.2.1. Preparation of UIO-66 (COOH)2

    The preparation and purification methods utilized in this study have been adapted from a previous report, with minor adjustments[34]. A total of 56.8 g of Zr(SO4)2.4H2O and 27.2 g of pyromellitic acid were introduced into a 500-mL PTFE bottle, followed by the addition of 320 mL of distilled water. Following a 30-min period of stirring at ambient temperature, an 8-mL volume of sulfuric acid (H2SO4) was gradually introduced into the suspension. The suspension was then agitated for an additional duration of 10 min. The PTFE bottle was closed and subjected to heating by immersion in an oil bath and maintenance of a temperature of 98 °C for 16 h. The reaction mass was cooled to ambient temperature and subjected to multiple water washes until the pH reached a neutral level. The crude product was oven-dried overnight at 80 °C. The crude product was subsequently refluxed in 300 mL of methanol for 18 h to eliminate any pyromellitic acid that had not reacted. Subsequently, the mixture was filtered while still hot and subjected to a series of washes with methanol (100 mL × 3 times). The resulting product was then dried at 80 °C for 3 h, followed by further drying at 150 °C overnight. The mass of the final product was determined to be 31 g. In the following, the bare MOF is referred to as UB.

    2.2.2. Post-synthetic modification of UIO-66 (COOH)2 with L-glutamic acid

    A total of 3 g of UB and 1.5 g of L-glutamic acid were combined in a mortar and thoroughly grounded until a homogeneous mixture was obtained. The powder sample was placed into a porcelain crucible and heated at 180 °C for 5 h. The reaction mass was cooled to ambient temperature and introduced into 500 mL of water and agitated for 6 h to eliminate unreacted L-glutamic acid. The suspension was filtered and washed with water (100 mL × 2 times). It was then dried at 80 °C for 3 h, followed by further heating at 150 °C overnight. The weight of the final product was measured to be 3.21 g. In the following, the L-glutamic acid-modified MOF is referred to as UG.

    2.3. Batch adsorption experiments

    The procedure in this study have been adapted from our previous report[ 35]. To prepare stock solutions with a concentration of 1000 ppm, HgCl2 and PbCl2 were dissolved in distilled water using a 1-L volumetric flask. The resultant solution was subsequently diluted using deionized water until the desired volume was attained. Following this, a sequence of 100 mL solutions comprising heavy metal ions with different initial concentrations ranging from 1 to 150 ppm were prepared by progressively diluting the original stock solutions. Adsorption experiments were performed utilizing 50-mL conical flasks equipped with a rubber septum. The experimental procedure entailed the introduction of an adsorbent into a 10-mL solution containing heavy metal ions. Subsequently, the conical flask was sealed tightly using a rubber septum and placed onto an open-air shaker (OS-4000, Jeio Tech, Korea) where it undergo continuous agitation at a rate of 150 rpm. This measure was implemented in order ensure sufficient adsorption under the specific conditions of an ambient temperature of 25 ± 2 °C. The suspension was filtered using a disposable syringe filter with a pore size of 0.45 μm. Subsequently, the filtrate was subjected to analysis to ascertain the concentration of residual heavy metal ions.

    The heavy metal ion (%) removal was calculated using the following equation:

    ( % ) R e m o υ a l = ( C o C e ) C o × 100

    The heavy metal ion removal capacity at equilibrium was calculated using the following equation:

    R e m o υ a l C a p a c i t y ( m g g ) = ( C o C e ) × V m

    Where Co and Ce are the initial and equilibrium concentrations in mg/L of heavy metal ions, respectively, V is the volume of the solution in liters, and m is the mass of adsorbent in grams.

    2.4. Instrumentation

    Fourier-transform infrared (FTIR) spectra were recorded using a IRSpirit spectrometer (Shimadzu, Japan) with a resolution of 1.0 cm-1. The spectra were obtained in the wavelength range of 400~4000 cm-1 at ambient temperature. Powder X-ray diffraction (PXRD) patterns were collected at room temperature using a X-ray diffractometer (Bruker AXS GmbH, D8 advance Eco, Germany) at 40 kV and 25 mA with Cu Kα (λ = 0.154 nm) wavelength. The angle range for data collection was set from 2θ = 5° to 50°, and the scan speed was set at 0.1 °/s. The surface morphology of the samples was analyzed using a field emission scanning electron microscope (FE-SEM), and elemental mapping images were obtained simultaneously by coupling energy dispersive X-ray spectroscopy (EDS) with the FE-SEM instrument (JSM-7001F, JEOL, USA). Before FTIR and SEM-EDS analysis, the samples were oven-dried at 150 °C overnight. Elemental (CHN) analysis was conducted using an automatic elemental analyzer (FLASH 2000, Thermo Fisher Scientific, Italy). Inductively coupled plasma optical emission spectrometry (ICP-OES; Avio 200, Perkin Elmer Inc., USA) equipped with auto sampler was used to quantify the concentration of residual HMIs in the aqueous solution.

    3. Results & discussion

    3.1. Characterization

    The CHNS elemental analysis technique is an economically efficient approach used to ascertain the levels of carbon (C), hydrogen (H), nitrogen (N), and sulfur (S) in both organic and inorganic samples. The technique relies on the combustion of a sample in the presence of oxygen at a temperature of 1000 °C, enabling the detection of minute amounts of elemental components. We performed CHNS analysis of the as-synthesized UB and UG, and the results were as follows: the percentages of C, H, S, and N in UB were 27.70, 0.30, 0.20, and 0.00%, respectively, and those in UG were 28.49, 1.43, 0.12, and 0.90%, respectively. Sulfur was present in the samples due to the usage of H2SO4 as a mineralizing agent in the preparation of UB. The increment of C% and H% as well as the presence of N in UG indicates the presence of glutamic acid. UG contained 0.90% of N, which is equivalent to 0.095 g of glutamic acid per 1 g of UG sample. This value was slightly higher than the experimental value (0.07 g glutamic acid per 1 g of UG) obtained from the weight of the product (section 2.2.2), which may be due to handling errors. The results indicated that, in this solid-solid reaction approach, 9.5% of glutamic acid was grafted on the surface of UB. The reduction in the S% value in UG might be due to excessive washing with water during sample purification or displacement of anionic sulphate ions with the carboxylic acid groups of glutamic acid.

    The X-ray diffraction (PXRD) patterns of the as-synthesized UB and UG samples are illustrated in Figure 1a. The observed peak positions of UB at 7.4° and 8.0° align with the results reported in a previous study[34]. A similar peak pattern was observed for the compound UG at 8.0° and 8.5°. The peaks exhibited a rightward shift and broadening because of the grafting of glutamic acid molecules, which compressed the UB lattice. Comparable peak shifts have been reported in a previous study[36].

    The FTIR spectra of pyromellitic acid, glutamic acid, UB, and UG are shown in Figure 1b. The spectrum of pyromellitic acid showed a broad peak in the region between 2500 and 3155 cm-1, which corresponds to the bending vibrations of -OH in free -COOH groups. The sharp peak at 1693 cm-1 corresponds to the stretching vibrations of -C=O in free -COOH groups. The spectrum of UB clearly showed a peak at 1708 cm-1 (-C=O), which confirmed the presence of free -COOH groups in the bare MOF and was in good agreement with a previous report[34]. However, the intensity of the broad peak (-OH) may have diminished because of its thermal activation. According to previous reports, heating of UB (> 100 °C) induces formation of anhydride bonds, which may account for the decrease in intensity of the related broad peak[37]. The spectrum of glutamic acid also showed a broad peak at 3000 cm-1 and sharp peaks at 1640 and 1290 cm-1, which correspond to -OH, -C=O, and -C-N groups. Similar to the spectrum for UB, the intensity of the peak related to -OH groups in the spectrum of UG was reduced. Apart from -C=O and -C-N stretching bands, a new peak at 1568 cm-1 was observed in the spectrum for UG, which was not present in the spectra for UB and glutamic acid. We assume that the free -COOH groups of UB and the -NH2 groups of glutamic acid reacted with each other to form new amide (-CONH-) bond when the bare MOF (UB) and glutamic acid were heated together. It was also likely that the grafting of glutamic acid onto the surface of the MOF took place via formation of amide (-CONH-) bonds. The new peak observed at 1568 cm-1 corresponds to the antisymmetric stretching bands of -CON functional groups[38].

    The SEM and elemental mapping images and the map sum spectra of as-synthesized UG and UB are depicted in Figure 2(a-f). The SEM images of UB suggest that the surface was porous and amorphous in nature. After modification with glutamic acid, the surface roughness was increased. The presence of N in the elemental mapping image and map sum spectrum of UG (Figure 2e and 2f) clearly indicate the presence of glutamic acid on its surface. To record SEM and EDS images for samples after adsorption, UB and UG were treated with 100 ppm of a Pb(II)/Hg(II) binary mixture. The SEM and elemental mapping images and map sum spectra of heavy metals adsorbed to UG and UB are depicted in Figure 3(a-f). The elemental mapping images and map sum spectra of UB and UG after adsorption clearly indicate the presence of Pb and Hg on their surfaces.

    3.2. Effect of initial concentration on the Pb(II) and Hg(II) ions adsorption

    To ascertain the nature of the interaction between adsorbate molecules and adsorbent, as well as calculate the maximum adsorption capacity, it is imperative to establish an adsorption isotherm by examining adsorbate ions throughout a range of initial concentrations. Single ion batch adsorption experiments using UB and UG were conducted with initial concentrations between 1 and 150 ppm of Pb(II) and Hg(II) ions. The effect of the initial concentration on Pb(II) removal efficiency is shown in Figure 4a. The Pb(II) removal efficiency (%) of UB increased from 87.27 (1 ppm) to 97.58% (10 ppm) and decreased gradually to 58.02% (150 ppm). The Pb(II) removal efficiency (%) of UG increased from 87.95 (1 ppm) to 96.87% (10 ppm) and decreased gradually to 24.93% (150 ppm). The effect of the initial concentration on Hg(II) removal efficiency is shown in Figure 4b. The removal efficiency (%) of Hg(II) ions increased from 3.9 to 80% between 1~5 ppm and decreased gradually to 14.80% for UG, whereas much lower Hg(II) removal efficiency (%) values (0~2.7%) were obtained for UB. The initial concentration of ions plays a crucial role in facilitating the transportation of adsorbent ions from the bulk solution to active adsorption sites on the adsorbent material[39]. The lower removal percentage observed at an initial concentration of 1 ppm can be attributed to a reduced concentration gradient in the vicinity of the adsorbent interface. Consequently, the diffusion of Pb(II)/Hg(II) ions was hindered at lower initial concentrations. Previous studies reported similar findings[40]. The further decrease in the removal percentage of Pb(II)/Hg(II) ions with increasing initial concentrations was because, at a constant adsorbent dosage, the total number of available adsorption sites of the adsorbent are fixed and adsorb nearly the same amount of adsorbate ions, resulting in a decrease in the removal percentage of heavy metal ions, which corresponds to an increase in the initial concentration.

    3.2.1. Adsorption isotherms

    The adsorption isotherms of Pb(II) and Hg(II) ions, obtained at a temperature of 25 °C, are illustrated in Figure 4c. Except for the Pb(II) adsorption isotherm for UB, all other isotherms reached saturation within the range of concentrations measured. The adsorption capacitiesof UB and UG for Pb(II) were found to be 261.87 and 102.66 mg/g, respectively. The potential adsorption sites for Pb(II) ions on UB are unoccupied carboxyl (-COOH) groups originating from pyromellitic acid. Despite the relatively low grafting amount of glutamic acid onto the surface of UB, the adsorption capacity for Pb(II) was significantly diminished after modification with glutamic acid. The larger glutamic acid molecules may impede the reactivity of active -COOH groups on UB. Therefore, it is possible that hindrance contributed to the decrease in the adsorption capacity for Pb(II) in UG. The adsorption capacities of UB and UG for Hg(II) were 10.54 and 80.6 mg/g, respectively. Notably, the adsorption of Hg(II) on UB was significantly lower than that of Pb(II). This observation suggests that the active -COOH sites present did not exhibit a significant affinity for Hg(II) ions. Following the introduction of glutamic acid, the adsorption of Hg(II) was substantially enhanced by approximately 87%. The observed behavior can be explained by referring to the Hard and Soft Acids and Bases (HSAB) theory. According to this theory, Pb(II) ions are classified as borderline acids, while Hg(II) ions are classified as soft acids[41]. The carboxyl (-COOH) groups present in pyromellitic acid are classified as hard bases, which explains the adsorption of Pb(II) ions on UB. Glutamic acid possesses an amino group (-NH2), and upon grafting onto UB, a novel amide bond is generated, thus characterizing it as a soft base. Therefore, these soft-base sites can adsorb soft Hg(II) ions. Table 1 compares the Hg(II) ions maximum adsorption capacity for various adsorbents reported in the literature.

    3.2.2. Isotherm models

    The experimental data were fitted using the non-linear Langmuir and Freundlich isotherm models given in equations (3) and (4), respectively, and the results are shown in Figure 4c.

    q e = b q m C e 1 + b C e

    q e = k C e 1 / n

    Where Ce (mg/L) represents the equilibrium concentration of Hg2+ ions; qe (mg/g) represents the amount of Hg2+ ions adsorbed per unit mass of adsorbent at equilibrium; qm (mg/g) is the Langmuir constant that represents the maximum adsorption capacity assuming a monolayer coverage of adsorbate over a homogenous adsorbent surface; b (L/mg) is a kinetic parameter representing the adsorption energy of the adsorbent for the adsorbate; k (mg/g) is the Freundlich constant related to the adsorption capacity; and 1/n is an empirical parameter related to adsorption intensity or surface heterogeneity[35]. Table 2 presents the parameters corresponding to the two isotherm models described above. Based on R2 values, Pb(II) and Hg(II) ion adsorption to UB followed the Langmuir model, with R2 values of 0.9872 and 0.9784, respectively. The Langmuir model can be attributed to the formation of a monolayer of heavy metal ions on the surface of the UB material. Upon completion of monolayer formation, an observable plateau was detected in the Langmuir fit, which signified the saturation of all the active sites[46]. The Pb(II) qmax values obtained from Langmuir model fits were 300.42 and 102.66 mg/g for UB and UG, respectively. Heavy metal ion adsorption on UG following the Freundlich and Langmuir model fits did not converge. The applicability of the Freundlich model indicated a heterogeneous nature of the adsorbent, i.e., more than one adsorption site was present on the surface.

    3.3. Effect of contact time on the Pb(II) and Hg(II) ions adsorption

    The study of adsorption kinetics offers insights into both the mechanism and rate at which the adsorption process occurs. Using UB and UG, single ion batch adsorption experiments were conducted between 15 and 150 min, with an initial concentration of 10.0 ppm for Pb(II) and Hg(II) ions. Figure 5a depicts the effect of contact time on Pb(II) removal efficiency. The Pb(II) ion removal efficiency for UB at 15, 30, and 150 min was 91.86, 94.64, and 97.19%, respectively, and that for UG at 15, 30, and 150 min was 85.57, 91.17, and 95.9%, respectively. A minimal increase in the RE% was detected in both materials after a contact time of 30 min. Figure 5b depicts the effect of contact time on Hg(II) removal efficiency. The Hg(II) ion removal efficiency for UG at 15 and 150 min was 19.09 and 30.2%, respectively, whereas that for UB at 15 and 150 min was 0.03 and 0.73%, respectively. The removal percentage for UG constantly increased in the time interval between 15 and 150 min, which indicates that it did not saturate during that time.

    3.3.1. Kinetic models

    Nonlinear pseudo-first-order (PFO) (equation 5), pseudo-second-order (PSO) (equation 6), and intraparticle diffusion kinetic models (equation 7) were used to fit the experimental adsorption kinetics data. The kinetic model fittings are shown in Figure 5c.

    q t = q e ( 1 e k 1 t )

    q t = q e 2 k 2 t q e k 2 t + 1

    q t = k i t 1 / 2 + C q t

    Where qe (mg/g) and qt (mg/g) represent the amount of Hg2+ ions adsorbed per gram of adsorbent at equilibrium and any time (min), respectively. k1 (min−1) and k2 ( are the rate constants of the PFO and PSO kinetic models, respectively. ki (mg/g.min1/2) is the rate constant, and C (mg/g) is the thickness of the boundary layer of the intraparticle diffusion kinetic model[35]. Table 3 presents the parameters corresponding to the three kinetic models described above. The experimental data for Pb(II) ion adsorption using both UB and UG were best fit by the PSO model. The R2 values obtained for UB and UG were 0.9962 and 0.9972, respectively. Notably, surface adsorption, including chemisorption, is the rate-limiting factor in this model[47]. In the case of Hg(II) ion adsorption on UB, the difference obtained in the R2 values of both the PFO (0.8451) and PSO models (0.8296) was very small. Hence, both physisorption and chemisorption were involved. The experimental data for Hg(II) ion adsorption on UG were better fit by the intraparticle diffusion model (R2 = 0.9543) than by the PSO model. The fit did not converge in the PFO model. The kinetics of adsorption encompass three distinct mass transfer phenomena: external diffusion, internal diffusion, and adsorption onto active sites. External diffusion refers to the movement of adsorbate within the liquid film surrounding the adsorbent. Internal diffusion pertains to the transfer of adsorbate within the pores of the adsorbent material (IPD model). Lastly, adsorption onto active sites denotes binding of the adsorbate molecules to reactive sites on the adsorbent surface. Generally, it is assumed that if the function represented by qt vs t1/2 passes through the origin (i.e., C = 0), adsorption solely follows intraparticle diffusion; if that function does not pass through the origin, adsorption is controlled by multiple processes[48,49]. The plot of the function qt vs t1/2 for UG/Hg(II) adsorption is shown in Figure 5d. The plot clearly shows that the regression does not pass through the origin (C = 5.007), which indicates that the rate-controlling step is not interparticle diffusion alone but that additional processes, such as adsorption onto active sites, may influence the adsorption rate as well. Based on the fitting results of the isotherm and kinetic data models, the adsorption of Pb(II) on both UB and UG can be attributed to surface complexation of Pb(II) ions with available –COOH groups (pyromellitic acid) on the respective surfaces. The adsorption of Hg(II) on UG was attributed to chelation, in which the glutamic acid grafted onto UG formed chelates with Hg(II) ions.

    4. Conclusions

    We successfully modified UIO-66 (COOH)2 with L-glutamic acid via the solid-solid reaction route. Nitrogen was found in the SEM-EDS elemental mapping images of the L-glutamic acid-modified MOF, which confirmed the presence of glutamic acid on its surface. We further observed a decrease in the Pb(II) adsorption capacity of the bare MOF after modification. This MOF showed a higher ion adsorption capacity for Pb(II) (261.87 mg/g) than for Hg(II) ions (10.54 mg/g) at an initial concentration of 150 ppm. In contrast, we observed an increment in the Hg(II) ion adsorption capacity for the L–glutamic acid-modified MOF (80.6 mg/g) as compared to the bare MOF. The ion adsorption capacity for Hg(II) increased by almost 87% after modification of UIO-66 (COOH)2 with L-glutamic acid. The Langmuir model well fit the Pb(II) adsorption isotherms, and the Pb(II) qmax values obtained for the fits of this model were 300.42 and 102.66 mg/g for UB and UG, respectively. The Freundlich model well fit the Hg(II) adsorption isotherms. Adsorption of Pb(II) ions on the bare and L-glutamic acid-modified MOFs were best fit by the PSO model, which indicates that the adsorption process was chemisorption. Adsorption of Hg(II) ions on UG followed intraparticle diffusion and subsequent adsorption onto active sites. The preparation of L-glutamic acid-modified MOFs is a cost-effective process.


    This work was supported by the Korea Environment Industry and Technology Institute through the Prospective Green Technology Innovation Project, funded by the Korea Ministry of Environment (2021003160013).


    (a) PXRD patterns of UB and UG (b) FTIR patterns of pyromellitic acid, glutamic acid, UB and UG.
    SEM-EDS analysis of as-synthesized samples; (a) SEM image of UB, (b) EDS elemental mapping image of UB, (c) map sum spectrum of UB, (d) SEM image of UG, (e) EDS elemental mapping image of UG, (f) map sum spectrum of UG.
    SEM-EDS analysis of samples after adsorption; (a) SEM image of UB, (b) EDS elemental mapping image of UB, (c) map sum spectrum of UB, (d) SEM image of UG, (e) EDS elemental mapping image of UG, (f) map sum spectrum of UG.
    (a) Effect of the initial concentration on Pb(II) removal efficiency [volume: 10 mL; amount of adsorbent: 30 mg; contact time: 150 min], (b) effect of the initial concentration on Hg(II) removal efficiency [volume: 10 mL; amount of adsorbent: 30 mg; contact time: 150 min], (c) isotherm model fittings (solid lines indicate Langmuir model fits and dashed lines indicate Freundlich model fits).
    (a) Effect of contact time on Pb(II) removal efficiency [initial concentration: 10.0 ppm; volume: 10 mL; amount of adsorbent: 30 mg], (b) effect of contact time on Hg(II) removal efficiency [initial concentration: 10.0 ppm; volume: 10 mL; amount of adsorbent: 30 mg], (c) kinetic model fittings (dotted lines indicate PFO model fits, solid lines indicate PSO model fits, and dashed lines indicate intraparticle diffusion model fits), (d) qt vs t1/2 plot of Hg(II) adsorption on UG.


    A Comparison of the Maximum Adsorption Capacities of Various Adsorbents for Hg(II) Ions
    *since the Langmuir fit was not converged, experimental adsorption capacity was given.
    Isotherm Parameters of Pb(II) and Hg(II) Ion Adsorption on UB and UG
    - Fit did not converge.
    Kinetic Parameters of Pb(II) and Hg(II) Ion Adsorption on UB and UG
    - Fit did not converge.


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