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. 2026 Jul 2;16(34):32131–32145. doi: 10.1039/d6ra00235h

Fe3O4 magnetically modified NN and NO donor chelating agents for effective extraction of Cd(ii), Pb(ii) and Cr(vi) ions from water

Siphosethu Maqinana a, Stephen O Ojwach a,✉
PMCID: PMC13325835  PMID: 42395822

Abstract

The NN and ON donor Schiff-base ligands; N-(pyridin-2-ylmethyl)-3-(triethoxysilyl)propan-1-amine (L1) and 2-(((3-(triethoxysilyl)propyl)amino)methyl)phenol (L2) were immobilized on magnetic nanoparticles to afford the functionalized materials L1@Fe3O4 and L2@Fe3O4, respectively. The fabricated materials were characterized using FT-IR for functional group analyses, SEM, TEM, and EDX for surface morphology and elemental composition, nitrogen physisorption for porosity and surface area, VSM for magnetic properties, and PXRD and TGA–DTA/DSC for crystallinity and thermal properties, respectively. The extraction of Cd(ii), Pb(ii), and Cr(vi) metal ions from aqueous solutions using L1@Fe3O4 and L2@Fe3O4 was studied under varied pH, contact time, initial concentration, and adsorbent dosage conditions. Optimal extraction conditions for Cd(ii), Pb(ii), and Cr(vi) ions of pH values of 12 and 3, contact times of 30, 240, and 60 minutes, initial concentrations of 20 and 40 ppm, and adsorbent dosages of 0.5 mg and 20 mg were recorded, respectively, using L1@Fe3O4. Optimal extraction efficiencies of 67%, 82%, and 99% were obtained for Cd(ii), Pb(ii), and Cr(vi) ions, respectively. Adsorbent L1@Fe3O4 showed appreciable adsorption capacities of up to 80% in the naturally heavy metal-polluted river water. The adsorption kinetics for both adsorbents followed pseudo-second-order models. The equilibrium data were best described by the Freundlich isotherm model for both L1@Fe3O4 and L2@Fe3O4 across all the metal ions. The calculated qe values for Cd(ii), Pb(ii), and Cr(vi) ions were 4.2061 mg g−1, 13.1752 mg g−1, and 6.9862 mg g−1 for L1@Fe3O4, and 2.4682 mg g−1, 14.9186 mg g−1, and 14.2116 mg g−1 for L2@Fe3O4, respectively. The chelating agent L2@Fe3O4 exhibited low to moderate regeneration and reusability in the extraction of the Cd(ii), Pb(ii) and Cr(vi) ions.

The NN and ON donor Schiff-base ligands, L1 and L2, were immobilized on magnetic Fe3O4 nanoparticles, yielding the adsorbents L1@Fe3O4 and L2@Fe3O4, respectively.graphic file with name d6ra00235h-ga.jpg

1. Introduction

The presence of pollutants and contaminants, such as organic dyes and heavy metals, has a negative impact on the environment and human health, as they may cause several diseases and side effects to living organisms.1 These toxic heavy metals are discharged into water sources from industrial effluents.2 Amongst the heavy metals, Cd(ii), Pb(ii), and Cr(vi) are considered the most hazardous due to them being persistent, non-essential, and highly toxic even in trace amounts.3,4 The permissible contaminant limits set by the United States Environmental Protection Agency (USEPA) and World Health Organization (WHO) for Cd(ii), Pb(ii), and Cr(vi) are 0.005 mg L−1, 0.01 mg L−1, and 0.05 mg L−1, respectively.5,6 In high concentrations, Cd(ii) causes bone defects, diabetes, hypertension, and chronic kidney disease, Pb(ii) causes irreversible toxic effects on the brain, kidneys, lungs, and liver, and is primarily used in motor vehicles, while Cr(vi) causes lung and gastrointestinal cancer and irreparable damage to the kidneys and central nervous system.6,7

Several techniques, such as adsorption, ion exchange, membrane filtration, reverse osmosis, and chemical precipitation, have been used to remove heavy metals from wastewater. The adsorption method is the most widely used due to its low cost, high efficiency, easy handling, and more available adsorbents for removing heavy metals.8 To date, several materials have been used as adsorbents to remove heavy metal ions from aqueous solutions, including activated carbons, carbon nanotubes, clay, zeolites, graphene, synthetic polymers, ion exchange resins, magnetic nanomaterials, metal oxides, and metal-based nanomaterials, silica-based nanomaterials, and organic functionalized magnetic nanoparticles, amongst others.9,10 Magnetic nanoparticles offer desirable characteristics such as high surface area, affordability, and ease of physical separation, thus making them cost-effective adsorbents.11 However, bare magnetic nanomaterials are readily oxidized under acidic conditions. Additionally, they tend to aggregate due to magnetic forces, which reduces their adsorption capabilities and mechanical strength. Hence, their immobilization on organic agents has been employed to circumvent these drawbacks.11

Currently, considerable attention is being directed towards the use of organic chelating agents to remove heavy metal pollutants from water.12–14 Of these chelating organic agents, Schiff bases have emerged as highly attractive due to the high stability of the metal complexes formed, hence improving extraction efficiencies and limiting leaching.14 For instance, a salicylaldehyde Schiff base ligand was synthesized and immobilized on magnetic nanoparticles for the extraction of Cd(ii) and Pb(ii) ions from aqueous solutions.15 Another hydride Schiff base was prepared by grafting ethylenediamine (EDA), ethyl acrylate (MA), and salicylaldehyde (SA) and immobilized on iron oxide magnetic nanoparticles for the selective extraction of Hg(ii) ions from aqueous solution.16 Furthermore, N2S tridentate organic ligands were synthesized and immobilized on magnetic nanoparticles for the extraction of Cd(ii) and Pb(ii) ions.1 In another contribution, epoxy-triazinetrione (ETT) Schiff-based ligands were immobilized on iron oxide magnetic nanoparticles and utilized for the extraction of Pb(ii) ions.17

A recent study by Nyamato and co-workers reported the use of Fe3O4-functionalized Schiff base tridentate NNN and NNO donor ligands for the removal of Cr(vi), Cd(ii), and Pb(ii) from aqueous solutions under laboratory conditions.18 While that study established the effectiveness of tridentate ligand systems, the influence of donor atom type and coordination mode on adsorption behaviour remains insufficiently explored. In this context, the present study investigated the syntheses and characterization of Fe3O4-supported Schiff base ligands bearing bidentate NN and NO donor atoms, providing a fundamentally different coordination environment compared to the previously reported tridentate systems. This variation in ligand design is expected to influence metal–ligand interactions, selectivity, and extraction efficiency. In addition, the current study goes beyond laboratory-spiked water samples to include the application of these adsorbents to real wastewater samples, validating their practical relevance for wastewater remediation. The synthesized materials were comprehensively characterized for functional groups, morphology, particle size, surface composition, magnetic properties, crystallinity, thermal stability, and elemental composition. The adsorption performance was systematically evaluated through kinetic, isotherm, and mechanistic studies to elucidate the extraction behavior of Cd(ii), Cr(vi), and Pb(ii) ions under various experimental conditions.

2. Materials and methods

2.1. Materials and instrumentation

The reagents: (3-aminopropyl)triethoxysilane (3-APTES) (99%), 2-pyridinecarboxaldehyde (99%), sodium borohydride (98%), and salicylaldehyde (98%) were obtained from Sigma-Aldrich and were used as received without further purification. The solvents, methanol, chloroform, and toluene were purchased from Merck, and methanol was used as received without purification. Toluene was distilled and dried in molecular sieves before use. Potassium dichromate (K2Cr2O7), cadmium nitrate Cd(NO3)2, lead nitrate Pb(NO3)2, sodium hydroxide pellets (NaOH), nitric acid (HNO3), iron(iii) chloride hexahydrate (FeCl3·6H2O), and iron(ii) chloride tetrahydrate (FeCl2·4H2O) were bought from Sigma-Aldrich and used without further purification. At room temperature, 1HNMR spectra were recorded on a Bruker Ultrashield 400 spectrometer in CDCl3 solution. Infrared spectra were recorded on a PerkinElmer Spectrum 100 in the range 4000–400 cm−1. The morphology and shapes of the chelating materials were acquired by ZEISS EVO LS 15 scanning electron microscope (SEM), and samples were coated with gold using Quorum and JEOL JEM 1400 transmission electron microscope (TEM). ImageJ software was used to measure the average diameter of the images, and Origin 2021 software was used to plot the histograms and the average distribution of the chelating agents. Powder X-ray diffraction (PXRD) patterns of the chelating agents were captured using an X-ray diffractometer (Bruker D8 Advance). Energy dispersive X-ray (EDX) analysis was carried out using an Oxford-made EDX detector. Data evaluation, using the Brunauer–Emmett–Teller (BET) and Barrett–Joyner–Halenda (BJH) methods, was conducted with BEL master software version 7.3.2.0. Magnetic susceptibility measurements were performed using a Quantum Design MPMS3 Evercool SQUID magnetometer with a 7 Tesla magnet at 300 K. Variable applied DC magnetic fields ranging from −4 T (−40 000 Oe) to 4 T (40 000 Oe) were utilized in DC mode. The thermal stability of the supporting material and chelating agents was recorded using PerkinElmer thermogravimetric analysis, differential thermal analysis, and differential scanning calorimetry (TGA-DTA/DSC) with Pyris software, upon heating from 25 °C to 800 °C at a rate of 20 °C min−1 under a N2 atmosphere. The Metrohm 827 pH meter was used to measure the pH of the working solutions. Standard buffer solutions of pH 4.0 and 7.0 purchased from Merck were used to calibrate the pH meter. The concentration of heavy metal ions in solution after extraction was measured using a Shimadzu Plasma Atomic Emission Spectrometer (ICPE-9820).

3. Results and discussion

3.1. Synthesis and characterization of immobilized chelating agents

The Schiff base compounds (E)-N-(pyridin-2-ylmethylene)-3-(triethoxysilyl)propan-1-amine (S1) and (E)-2-(((3-(triethoxysilyl)propyl)imino)methyl)phenol (S2) were synthesized via condensation reactions of 2-pyridinecarboxaldehyde and salicylaldehyde with (3-aminopropyl)triethoxysilane, following reported procedures.19 Treatments of S1 and S2 with NaBH4 afforded the corresponding reduced analogues N-(pyridin-2-ylmethyl)-3-(triethoxysilyl)propan-1-amine (L1) and 2-(((3-(triethoxysilyl)propyl)amino)methyl)phenol (L2) ligands, respectively (Scheme 1). Immobilization of the chelating agents L1 and L2 on Fe3O4 was carried out as per published protocols20 to give L1@Fe3O4 and L2@Fe3O4 (Scheme 1).

Scheme 1. Synthesis and immobilization of L1 and L2 on Fe3O4 to give L1@Fe3O4 and L2@Fe3O4, respectively.

Scheme 1

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1H NMR was used to interpret the structural descriptions of the synthesized ligands (Fig. S1–S4). The HC Created by potrace 1.16, written by Peter Selinger 2001-2019 N signal at 8.63 ppm (Fig. S2) and 8.36 ppm (Fig. S3) validated the condensation reaction between 2-pyridinecarboxyaldehyde and salicylaldehyde simultaneously with 3-APTES. The H2C-NH signal at 4.7 ppm (Fig. S4) and 3.4 ppm (Fig. S5) suggested the formation of L1 and L2 ligands, respectively. FT-IR was also used to identify functional groups on the prepared ligands and chelating agents (Fig. S6 and S7). For instance, the FT-IR spectra of the imine compounds S1 and S2 showed sharp bands at 1649 cm−1 (Fig. S6) and 1603 cm−1 (Fig. S7), which were attributed to the v(C Created by potrace 1.16, written by Peter Selinger 2001-2019 N) functional group.21 Upon reductions, the v(N–H) bands were observed at 3299 cm−1 and 3246 cm−1 in L1 and L2, respectively (Fig. S6 and S7).22 In addition, the FT-IR spectra of the functionalized chelating agents (L1@Fe2O4 and L2@Fe3O4) exhibited the bands for the v(Fe–O) group at 615 cm−1 and 584 cm−1, respectively, indicating the presence of magnetic nanoparticles on the ligands, which showed complete immobilization.1

The morphology, microstructure, size, shape, and average particle size of the chelating agents were examined using TEM. As depicted in Fig. 1, L1@Fe3O4 and L2@Fe3O4 exhibited agglomerated rod-like particles with dark spots, confirming the incorporation of Fe3O4 magnetic nanoparticles on the organic moieties. ImageJ software (Fig. S8) recorded the average particle size for L1@Fe3O4 and L2@Fe3O4 as 0.926 ± 0.037 nm and 0.949 ± 0.101 nm, respectively. These average particle sizes were relatively smaller compared to those reported by Shaabani et al., and Maleki et al., of 35 nm for NO Schiff bases immobilized on Fe3O4 nanoparticles.21,23 The surface composition of the chelating agents was further studied using the SEM technique (Fig. 1). The surface of L1@Fe3O4 appeared rough, while that of L2@Fe3O4 had small spherical particles that agglomerated to form large spheres.

Fig. 1. TEM images of (A) L1@Fe3O4, (B) L2@Fe3O4, and SEM images of (C) L1@Fe3O4, (D) L2@Fe3O4 chelating agents.

Fig. 1

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The distribution and composition of elements that are present in the chelating agents were determined using EDX analysis (Fig. S9). In both materials L1@Fe3O4 and L2@Fe3O4, it was clear that elemental Fe was present in quantities of 11.45 wt% and 32.43 wt%, respectively, confirming the successful immobilization of the compounds on Fe3O4. The coordinating donor atoms, N and O in L1@Fe3O4 and L2@Fe3O4, were present in quantities of 5.58 wt% and 19.39 wt%, respectively. The elemental mapping (Fig. S9) showed a monotonous and fine distribution of the elements on the surface of the chelating agents.

The surface areas of materials L1@Fe3O4 and L2@Fe3O4 chelating agents were determined using N2 adsorption–desorption isotherms (Fig. S10). The BET results exhibit a type IV isotherm with an H3 hysteresis loop, suggesting the presence of mesoporous structure in the range of 2 nm to 50 nm. The chelating agents, L1@Fe3O4 and L2@Fe3O4, displayed surface areas of 3 m2 g−1 and 6 m2 g−1, respectively (Table S1). The observed low surface areas are attributed to the incorporation of the organic chelating ligands in the Fe3O4 matrix.24 This finding aligns with previous reports, which reported surface areas ranging from 2.56 m2 g−1 to 3.69 m2 g−1 for the magnetite@poly-p-phenylenediamine immobilized salicylaldehyde (SA) composite (Fe3O4@PpPD-SA).25 Due to the high agglomeration of the chelating agents observed in SEM images, the porosity and pore sizes of the samples could not be accurately obtained using the BHJ method.

The saturation magnetization (Ms) values of the chelating agents were measured using a Vibrating Sample Magnetometer (VSM), and the corresponding magnetization curves are presented in Fig. S11. The magnetization saturation of the purely synthesized magnetic was determined to be 75.97 emu per g. A decrease in magnetic moment was observed for the chelating agents compared to the virgin magnetic nanoparticles, with L1@Fe3O4 and L2@Fe3O4 recording values of 0.97 emu per g and 0.60 emu per g, respectively (Table S1). The observed low Ms values can be attributed to the presence of magnetically inactive organic compounds.26,27 Similar results with low Ms values (14.23 emu per g and 0.389 emu per g) were reported for Fe3O4@SiO2-APTMS immobilized on 2-hydroxybenzaldehyde.28 Despite the low Ms values, ferromagnetic interactions and separations were effectively achieved using a bar magnet within 60 s (Fig. S11). The crystallinity of the pure Fe3O4, immobilized materials, L1@Fe3O4, and L2@Fe3O4, was determined using PXRD, as shown in Fig. S12. The hkl (111), (220), (311), (400), (422), (333), (440), (620), (533), and (642) corresponded to the pattern peaks at the 2θ: 18°, 30°, 35°, 43°, 53°, 57°, 62°, 71°, 74° and 87°are assigned to the standard Fe3O4 nanoparticles.1,29,30 The immobilization of the organic ligands on the magnetic nanoparticles did not significantly alter the diffraction peaks of the pure Fe3O4 nanoparticles (Fig. S12). The similarity in the diffraction patterns is expected since the crystallinity of the materials is largely controlled by the Fe3O4 phase, which agrees with the previous reports in the literature.21,30,31 However, the intensities of the peaks of the immobilised materials were weaker compared to the peak intensities in the diffraction patterns of the pure Fe3O4, indicating the presence of the organic groups on the Fe3O4 core.

The thermal stability of materials Fe3O4, L1@Fe3O4, and L2@Fe3O4 was elucidated using TGA–DTA/DSC analyses (Fig. S13 and Table S2). The TGA curves of the materials showed three weight losses. The first weight loss of 10% and 6% between 27 and 200 °C could be assigned to the elimination of water molecules on the materials,32 consistent with the endothermic peak on the DTA curve in the same temperature range. A second weight loss of 19% and 20% for the chelating agents at the temperature range of 219 °C to 448 °C was due to the loss of volatile organic components. Finally, the total weight loss of 29% and 33% for the chelating agents was observed from 449 °C to 800 °C, corresponding to the loss of all organic and inorganic matter.33 The chelating agents exhibited a similar weight loss across all temperature ranges (Table S2). This is expected due to the fabrication of the organic ligands on the same Fe3O4 magnetic nanoparticle. In addition, the two chelating agents, L1@Fe3O4 and L2@Fe3O4, differ only in the pyridine (N) and phenol (O–H) donors and are not expected to significantly influence the thermal properties of the materials, and are in good agreement with previous organic ligand-grafted iron oxide magnetic nanoparticles.15 Furthermore, studies indicate that between 449 °C and 600 °C, the weight loss is attributed to the decomposition of Fe3O4-immobilized Schiff bases and the organic spacer.31 The slightly higher weight loss of 33% for L2@Fe3O4, compared to 29% for L1@Fe3O4, can be explained by the greater magnetic nanoparticle content, as confirmed by EDX analysis.

3.2. Extraction of metal ions from aqueous solutions

3.2.1. Effect of pH on the extraction of metal ions

The pH of the solution is a very important parameter in the extraction of metal ions as it regulates the surface charge of the adsorbent, the nature of the metal cation in the solution, and the hydrolysis of the metal ions.34 For instance, at pH values below the pHPZC, the surface of the chelating agents is protonated and becomes positively charged, thus repelling the positively charged metal ions in the solution. On the other hand, pH values above the pHPZC result in deprotonation of the chelating agents, leading to an overall negative charge.35 The point of zero charge (pHPZC) of L1@Fe3O4 and L2@Fe3O4 were thus determined and obtained as 6.5 and 5.5, respectively (Fig. 2). From the data, the extraction efficiencies of Cd(ii) and Pb(ii) cations increased with the increasing pH. For Cd(ii) cations, a further increase of the pH to 12 resulted in high extraction efficiencies of 97% (Fig. 2C). However, at a high pH of above 10, it is known that Cd(ii) cations precipitate out of solution as Cd(OH)2 rather than undergo extraction via adsorption.36 Thus, a pH of 9 was adopted as the optimum for the extraction of Cd(ii) cations, yielding extraction efficiencies of 74% and 43% for L1@Fe3O4 and L2@Fe3O4, respectively. Similarly, a pH of 9 was optimal for extracting Pb(ii) cations, yielding 90% extraction efficiency with both L1@Fe3O4 and L2@Fe3O4 adsorbents. In contrast, the extraction efficiency of Cr(vi) anions decreased with increasing solution pH. A pH of 3 was optimal for Cr(vi) anions, yielding extraction efficiencies of 76% and 97% for L1@Fe3O4 and L2@Fe3O4, respectively.

Fig. 2. Point of zero charges for (A) L1@Fe3O4, (B) L2@Fe3O4, and extraction efficiency of metal ions at different pH values using 0.02 g of (C) L1@Fe3O4, (D) L2@Fe3O4 chelating agents on 20 ppm solutions for 24 h.

Fig. 2

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The high optimum pH of 9 observed in the extraction of Cd(ii) and Pd(ii) cations could be accounted for by the deprotonation of N–H and O–H groups, rendering the binding sites negatively charged, hence the ease of coordination of the positively charged Pb(ii) and Cd(ii) cations.37 Similar results (optimum pH of 9) were reported for the extractions of Pb(ii) and Cd(ii) cations using E-2-(((3-(triethoxysilyl)propyl)imino)methyl)phenol, E-2-methoxy-4-(((3-(triethoxysilyl)propyl)imino)-methyl)phenol, and E-4-(((3-(triethoxysilyl)propyl)imino)-methyl)phenol Schiff bases immobilized on magnetic nanoparticles.38,39 On the other hand, hexavalent chromium (Cr(vi)) can exist in several stable negatively charged species such as CrO42−, HCrO4−, Cr2O72−, and HCr2O7−, depending on the concentration and solution pH.40 These species thus require a positively charged adsorbent surface (e.g., –NH+, –OH2+), consistent with the observed higher extractions at a lower acidic pH of 3. Chen et al., reported similar results with an optimum pH of 2 for the extraction of Cr(vi) anions using biomass adsorbent.41

3.2.2. The effect of contact time, initial cation concentration, and adsorbent dosage

The impact of time on the extraction efficiency was assessed by varying the duration of the extraction from 30 min to 480 min using L1@Fe3O4 as the adsorbent (Fig. 3A). From the data in Fig. 3A, contact times of 30 min, 60 min, and 240 min were recorded in the extraction of Cd(ii), Cr(vi), and Pb(ii) ions, with extraction efficiencies of 67%, 99%, and 86%, respectively. The higher extraction efficiencies during the initial minutes (in the case of Cd(ii)) can be attributed to the abundance of available active sites on the surface of the chelating agent, which facilitates the binding of metal ions.42 With prolonged contact time, the number of accessible binding sites decreased, resulting in a reduction in the extraction efficiency of metal ions.42 The influence of the initial metal cation concentration was investigated across a concentration range of 10 ppm to 100 ppm using L1@Fe3O4 (Fig. 3B). Saturation limits of 20 ppm corresponding to extraction efficiencies of 91% and 90% were obtained for Cd(ii) and Pb(ii) cations, respectively. Meanwhile, for Cr(vi) anions, a saturation limit was reached at 40 ppm, with an extraction efficiency of 98%.

Fig. 3. Extraction efficiencies of Cd(ii) and Pb(ii) at pH 9 and pH 3 for Cr(vi) ions at (A) different contact times using 0.02 g of L1@Fe3O4, 20 ppm solutions, (B) different initial concentrations using 0.02 g of L1@Fe3O4 for 4 h, and (C) different chelating dosages (g) on 20 ppm solution for 4 h.

Fig. 3

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Next, the role of adsorbent dosage was investigated across the range of 0.0001 g to 0.02 g using the chelating agent L1@Fe3O4 at optimum pH and contact time (Fig. 3C). From the experimental results, a general increase in extraction efficiencies was noted with an increase in the adsorbent dosage. For example, an increase in the adsorbent dosage from 0.01 g to 0.02 g was marked by an increase in the extraction efficiency of Cr(vi) anions from 20% to 76%, respectively (Fig. 3C). For the Cd(ii) and Pb(ii) cations, optimum adsorbent dosages of 0.0005 g corresponding to extraction efficiencies of 76% and 86%, respectively, were observed. From these data, it is clear that Cr(vi) anions behaved differently and required a high adsorbent dosage, an observation we can loosely attribute to some different interactions between the adsorbent and the respective metal ions.43 Similar results, where high extraction efficiencies (99%) were obtained with lower adsorbent dosages (0.25 mg mL−1) for the extraction of Pb(ii) cations, have been reported in the literature using a NO chelating agent immobilized on SBA-15.44

3.2.3. Investigation of metal ion–ligand interactions

The role of the nature of the chelating ligand/identity of the donor atoms was studied by comparing the relative extraction efficiencies of the NN- and NO-donor chelating agents L1@Fe3O4 and L2@Fe3O4, respectively, towards Cr(vi), Cd(ii), and Pb(ii) ions (Fig. 4). From Fig. 4A, it is evident that the NO-donor chelating agent, L2@Fe3O4 showed a higher extraction efficiency of 96% towards Cr(vi) anions than the corresponding NN-donor adsorbent L1@Fe3O4, which recorded extraction efficiency of 76%. On the other hand, similar extraction efficiencies of 90% for Pb(ii) cations were reported for both NN- and NO-donor adsorbents L1@Fe3O4 and L2@Fe3O. The comparable extraction efficiencies observed for Pb(ii) cations with both L1@Fe3O4 (NN) and L2@Fe3O4 (NO) adsorbents can be attributed to its borderline acid character, which allows favorable interactions with both nitrogen and oxygen donor atoms. Additionally, Pb(ii) exhibits flexible coordination behavior, often adopting variable coordination geometries due to the presence of a stereochemically active 6s2 lone pair. This flexibility enables Pb(ii) to accommodate both N- and O-donor ligands effectively. Furthermore, the steric accessibility and spatial arrangement of donor sites in both adsorbents likely provide similar coordination environments, thereby minimizing selectivity differences, consistent with the comparable extraction efficiencies displayed by both NN- and NO-functionalized materials. For Cd(ii) cations, the NN-donor chelating agent L1@Fe3O4 exhibited a higher extraction efficiency of 73% compared to the respective NO-donor ligand L2@Fe3O4, which achieved an extraction efficiency of 43%. This trend can be explained by the concept of Hard Soft Acid Base (HSAB) theory.45 In this context, N- and O-donor atoms are categorized as intermediate and hard bases, respectively, while Cr(vi) and Pb(ii) ions are classified as hard and intermediate acids, while Cd(ii) cation is a soft acid.45 Thus, ligands L1@Fe3O4 (NN) and L2@Fe3O4 (ON) can be classified as moderate and hard donor ligands, respectively. The higher extraction of Cr(vi) anions, a hard acid by the NO-donor ligand L2@Fe3O4, hard base, agrees with the hard–soft acid–base theory, and so are the results reported for Cd(ii) cations.

Fig. 4. (A) The extraction efficiency of metal ions using 0.02 g of L1@Fe3O4 and L2@Fe3O4 at pH 9 for Cd(ii), Pb(ii), and 3 for Cr(vi) ions, respectively (B) selectivity of L1@Fe3O4 towards Cr(vi), Cd(ii), and Pb(ii) ions on a mixed solution using 0.02 g of adsorbents on 20 ppm solution at pH 3 and 9 for 24 h.

Fig. 4

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Next, investigate how other properties of the chelating agents, such as surface area and particle size, may influence their extraction efficiencies. It is expected that adsorbents with higher surface areas would display higher extraction efficiencies.46 However, these results revealed no significant effect of surface area on the extraction efficiency of the metal ions. For example, L1@Fe3O4, with a surface area of 3 m2 g−1 exhibited extraction efficiency of 76% and 73% for Cr(vi) and Cd(ii) ions, while L2@Fe3O4 with a higher surface area of 6 m2 g−1 achieved extraction efficiencies of 97% and 43% Cr(vi) and Pb(ii) ions, respectively, (Table S1, entries 1 and 2). Furthermore, both chelating agents displayed similar extraction efficiencies of 90% for Pb(ii) cations. These findings suggest that the HSAB concept, specifically the nature of the metal ions and donor atoms of the chelating agents, predominantly governed the observed relative extraction efficiencies but not the surface properties.

To further understand the relative selectivity of the chelating agents towards the three metal ions, L1@Fe3O4 was used in the extraction of mixed metal ions (real wastewater environment). This was accomplished by adding 0.02 g of the chelating agent to a solution containing the three metal ions with an initial concentration of 20 ppm at different pH values of 3 and 9 (Fig. 4B). As depicted in Fig. 4B, L1@Fe3O4 displayed high extraction efficiency of 98% for Cr(vi) anions at pH 3. However, at a high pH of 9, a low extraction efficiency of 40% was recorded towards Cr(vi) anions. Meanwhile, the Cd(ii) and Pb(ii) cations exhibited moderate to high extraction of 74% and 88% at pH 3, while maximum extraction efficiencies of 99% and 98% were recorded at pH values of 9, respectively. This finding points to the significant role of pH in the relative extraction of these metal ions and indicates that, by controlling solution pH, selective extraction of the respective metals can be fine-tuned. Interestingly, Fig. 4B shows that the extraction efficiency of all metal ions improved in the mixed-metal solution compared to the individual solutions. For instance, at pH 3, L1@Fe3O4 displayed extraction efficiencies of 76% and 98% for Cr(vi) anions in single- and mixed-metal solutions, respectively. The enhanced extraction efficiencies observed in the mixed-metal system may be attributed to changes in solution chemistry, such as ionic strength and metal speciation.47,48 These factors can influence the availability of active sites, the distribution of surface charge, and electrostatic interactions, thereby enhancing the extraction capacity of specific ions under certain conditions. Another plausible explanation, the observed synergy in mixed-metal systems can be rationalized by cooperative mechanisms at the adsorbent surface.49 The presence of multiple metal ions can induce local modifications of the surface charge, enhancing electrostatic attraction for other cations. Moreover, cooperative chelation at adjacent donor sites is possible, where the binding of one metal ion facilitates the coordination of another at nearby sites.50

The selectivity of L1@Fe3O4 toward different metal ions in mixed-metal systems was further assessed by calculating the separation factor (α) during competitive extraction experiments using eqn (1), where α12 represents the separation factor of metal ion 1 relative to metal ion 2, q1 and q2 are the equilibrium extraction capacities of metal ions 1 and 2, respectively, and C1 and C2 are their corresponding equilibrium concentrations in solution. A separation factor value greater than unity (α12 > 1) indicates preferential extraction of metal ion 1, whereas a value less than unity (α12 < 1) indicates preferential extraction of metal ion 2.51,52

3.2.3. 1

The separation factor results presented in Table 1 demonstrate that the extraction selectivity of L1@Fe3O4 is strongly dependent on the solution pH. For instance, under acidic conditions (pH 3), the adsorbent exhibited high selectivity towards Cr(vi) and Pb(ii) ions, with Cr(vi) showing the highest preference among the investigated ions. This is evidenced by the high αCr(VI) Cd(II) (16.421) and αCr(VI) Pb(II) (6.683) values. In addition, the αPb(II) Cd(II) value of 2.457 confirmed that Pb(ii) cation was favoured over Cd(ii) cation under acidic conditions. Consequently, the overall selectivity sequence of L1@Fe3O4 at pH 3 followed the order: Cr(vi) > Pb(ii) > Cd(ii). The enhanced selectivity toward Cr(vi) anions at low pH may be attributed to the predominance of protonated functional groups on the adsorbent surface, which promote strong electrostatic interactions with anionic chromate species present in acidic media.40 Similarly, the higher selectivity towards Pb(ii) cations than the Cd(ii) cations could be attributed to the preference of the Pb(ii) cations for the moderate and hard-donor N- and O-donor atoms, since Cd(ii) is a soft metal as discussed previously.45

Table 1. Separation factor of heavy metal ions under different pH conditions in mixed solution using L1@Fe3O4.
Metal ions solution α Cd(II) Pb(II) α Cd(II) Cr(VI) α Pb(II) Cd(II) α Pb(II) Cr(VI) α Cr(VI) Cd(II) α Cr(VI) Pb(II)
pH 3 0.407 0.061 2.457 0.149 16.421 6.683
pH 9 1.908 147.506 0.524 77.319 0.007 0.013
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In contrast, under basic conditions (pH 9), the selectivity behaviour changed significantly, with L1@Fe3O4 displaying very low selectivity towards Cr(vi) anions, as reflected by the very low values of αCr(VI) Cd(II) (0.007) and αCr(VI)/Pb(II) (0.013). Interestingly, the adsorbent showed high preferential extraction for Cd(ii) and Pb(ii) cations, with Cd(ii) being the most favoured metal. The high αCd(II) Pb(II) (1.908) and exceptionally large αCd(II) Cr(VI) (147.506) values confirmed the strong preference for Cd(ii) extraction under alkaline conditions. The observed high selectivity of Cd(ii) could be attributed to its faster complexation kinetics and greater complex stability, enabling it to occupy available adsorption sites more efficiently than Pb(ii) ions in competitive studies at pH 9. The selectivity trend at pH 9 followed the order: Cd(ii) > Pb(ii) > Cr(vi).

3.2.4. Regeneration and reuse of the adsorbent

The main objective of immobilizing the organic chelating agent on Fe3O4 was to achieve adequate separation and reuse of the adsorbents. The regeneration of the used L2@Fe3O4 chelating agent was carried out using 0.5 M HNO3 for Cd(ii) and Pb(ii) cations and 0.5 M NaOH for Cr(vi) anions. The regeneration studies were carried out using single-metal systems, where each metal ion was adsorbed independently. This regeneration process was repeated for four cycles (Fig. 5). As depicted in Fig. 5, the percentage extraction for Cd(ii) cations showed a significant decrease from 98% to 34% in the first and fourth regeneration cycles, indicating poor stability and recyclability of the adsorbent. Similarly, the extraction efficiency of Pb(ii) cations decreased from 98% to 55% from the first to the fourth cycles, respectively (Fig. 5), while the extraction efficiency for Cr(vi) anions decreased from 73% to 53% over the same period. To gain further insight into the changes that occur to the adsorbent surface during extraction and account for the observed decrease in extraction efficiencies in the recycling experiments, the morphology and surface composition of the chelating agent post-extraction were analyzed using SEM and TEM (Fig. S14). The results revealed that the spent chelating agents were less agglomerated and exhibited dark spots (Fig. S14), compared to the agglomerated morphology of the fresh adsorbents (Fig. 1B and D). This observation suggests that the organic ligands remained immobilized on the nanoparticles after the desorption cycles. The average particle size of the adsorbents also increased from 0.949 ± 0.101 nm to 2.088 ± 0.223 nm after extraction, which can be attributed to the incorporation of the metal ions in the adsorbent matrix. It is therefore reasonable to argue that clogging of the active sites could be responsible for the drastic declines in extraction efficiencies in the recycling studies.

Fig. 5. Regeneration cycles of L2@Fe3O4 chelating agent after extraction of Cd(ii), Pb(ii), and Cr(vi) metal ions.

Fig. 5

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3.2.5. Extraction of metal ions from the northern effluent wastewater treatment works (WWTW) water sample

One of the primary objectives of this study was to assess the practical application of these immobilized chelating agents in the removal of heavy metal ions from naturally contaminated waters. To this end, L2@Fe3O4 was employed to extract the metal ions from real wastewater samples collected from the Northern Effluent Wastewater Treatment Works (WWTW) in KwaZulu-Natal province, South Africa. Following digestion with concentrated nitric acid, the wastewater sample had an acidic pH of 1.50. The adsorption efficiencies (%) and adsorption capacities (mg g−1) of Cd(ii), Pb(ii), Cr(vi), and co-existing metal ions (Cu(ii) and Zn(ii)) using L2@Fe3O4 are summarized in Table 2. The material demonstrated variable extraction efficiencies of 80%, 62%, and 35% for Cd(ii), Pb(ii), and Cr(vi) ions, respectively (Table 2, entries 1–3). Notably, the removal efficiencies of Cd(ii) and Pb(ii) cations from the real wastewater sample were significantly higher than those obtained under single-metal systems (stock solutions) at comparable acidic conditions (pH 3), where lower efficiencies of up to 23% were observed (Fig. 2D). The higher extraction efficiencies in the real water samples thus suggests that L2@Fe3O4 exhibits significant adsorption capacity in multicomponent matrices, highlighting its high potential for practical wastewater treatment. In contrast, the extraction efficiency of Cr(vi) anions decreased markedly in the real wastewater sample (35%) compared to 97% efficiency achieved in a single-metal system (Fig. 2D). This reduction can be attributed to competitive adsorption arising from the presence of multiple co-existing metal ions, which likely limits the availability of active binding sites for Cr(vi) anions. Additionally, the removal of the Zn(ii) ions displayed moderate efficiency of 68% (Table 2, entry 5), further demonstrating the versatility of L2@Fe3O4 toward the adsorption of different metal species.

Table 2. Adsorption efficiencies (%) and capacities (mg g−1) of different metal ions from water sample collected from the northern effluent wastewater treatment works (WWTW) in KwaZulu-Natal, South Africa, using 0.02 g of L2@Fe3O4, at room temperature for 2 h.
Entry Heavy metal ions Initial concentration (mg L−1) Final concentration (mg L−1) Adsorption efficiency (%) Adsorption capacity (mg g−1)
1 Cd 1.38 0.28 80 0.82
2 Pb 12.11 4.61 62 5.63
3 Cr 16.54 10.84 35 4.28
4 Cu 6.21 3.30 47 2.18
5 Zn 0.67 0.21 68 0.34
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3.3. Adsorption kinetics

The adsorption kinetics of Cd(ii), Pb(ii), and Cr(vi) ions onto L1@Fe3O4 and L2@Fe3O4 were evaluated using the pseudo-first-order and pseudo-second-order kinetic models in order to investigate the nature of the interactions between the chelating agents and metal ions. The kinetic model equations and corresponding parameters are presented in Table S3, while the linear kinetic plots are shown in Fig. 6 and S15. The calculated adsorption parameters and correlation coefficients (R2) are summarized in Table S4. In general, relatively high R2 values ranging from 0.9466 to 0.9999 were obtained for the pseudo-second-order model, indicating that the adsorption process was more suitably described by the pseudo-second-order kinetic model than by the pseudo-first-order model. The calculated equilibrium adsorption capacities (qe,calc) values for L2@Fe3O4 were determined as 14.2116 mg g−1 for Cr(vi), 2.4682 mg g−1 for Cd(ii), and 14.9186 mg g−1 for Pb(ii) ions (Table S4). Similarly, the corresponding qe,calc values for L1@Fe3O4 were obtained as 6.9862 mg g−1, 4.2061 mg g−1, and 13.1752 mg g−1 for Cr(vi), Cd(ii), and Pb(ii), respectively (Table S4). Comparison between the experimental adsorption capacities (qe,exp) and the calculated equilibrium adsorption capacities (qe,calc) values revealed that the pseudo-second-order model provided good agreement for Pb(ii) and Cr(vi) adsorption for the L2@Fe3O4 adsorbent respectively. However, a noticeable discrepancy was observed for Cd(ii) adsorption onto L1@Fe3O4, where the calculated adsorption capacity (qe,calc) of 4.2061 mg g−1 was significantly lower than the experimental (qe,exp) of 13.5675 mg g−1, despite the relatively high R2 value (0.9894). The deviations observed for Cd(ii) adsorption onto L1@Fe3O4, may be attributed to surface heterogeneity, diffusion limitations, or the coexistence of multiple adsorption mechanisms not fully accounted for by the simplified pseudo-second-order model.53 Nevertheless, the predominance of the pseudo-second-order model and the high correlation coefficients obtained suggest that chemisorption played a significant role in the adsorption process of Cd(ii), Cr(vi), and Pb(ii) ions onto both L1@Fe3O4 and L2@Fe3O4,53 consistent with previous reports by Mohan et al. and Yuan et al. on Cr(vi) adsorption systems.40,54

Fig. 6. (a) Pseudo-first-order kinetics and (b) pseudo-second-order kinetics for the extraction of metal ions using L2@Fe3O4.

Fig. 6

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3.4. Adsorption isotherms

The equilibrium data for the extraction of Pb(ii), Cr(vi), and Cd(ii) ions at different concentration ranges were analysed using the Langmuir, Freundlich, and Temkin isotherm models (Fig. 7 and S16), and the calculated parameters are summarised in Table S6. Overall, the adsorption data for both L1@Fe3O4 and L2@Fe3O4 showed the best correlation with the Freundlich model, followed by the Temkin model, whereas the Langmuir model generally provided poorer fits (Table S6). The Freundlich model produced high correlation coefficients (R2 = 0.96998–0.99998) (Table S6), indicating that the adsorption processes were better described by heterogeneous surface interactions and multilayer adsorption behaviour rather than ideal monolayer adsorption.

Fig. 7. (a) Langmuir model, (b) Freundlich model, and (c) Temkin model for the extraction of metal ions using L2@Fe3O4.

Fig. 7

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In contrast, the Langmuir model displayed comparatively low R2 values, particularly for Pb(ii) adsorption onto L2@Fe3O4 (R2 = 0.38376) and Cr(vi) adsorption onto L2@Fe3O4 (R2 = 0.45343) (Table S6). The Langmuir model assumes homogeneous adsorption sites, identical adsorption energies, and monolayer coverage, with no adsorbate–adsorbate interactions.55,56 Furthermore, the co-existence of chelation, electrostatic interactions, and surface complexation may contribute to non-uniform adsorption behaviour, particularly for Pb(ii) and Cr(vi) ions. Similar observations of Pb(ii) and Cr(vi) adsorption on different adsorbents through surface complexation (inner-sphere coordination with functional groups such as carboxyl, hydroxyl, or amine) alongside electrostatic interactions have been reported, which can cause non-uniform energy distribution.57,58 Thus, the high qm (maximum adsorption capacity) value predicted by the Langmuir model for Cd(ii) adsorption onto L2@Fe3O4 (1.4982 × 103 mg g−1) contradicts the very low KL (binding strength) value (4.7585 × 10−4 L mg−1) (Table S6), which suggests weak adsorbate–adsorbent interactions and a greater tendency towards desorption. These inconsistencies may further account for the poor fits observed for the Langmuir model.

The Freundlich constant n, which is associated with adsorption intensity and surface heterogeneity, was determined to be 0.9799, 0.6216, and 0.9425 for Cd(ii), Cr(vi), and Pb(ii) adsorption onto L2@Fe3O4, respectively. For L1@Fe3O4, the n values were 1.2107, 0.9273, and 1.1462 for Cd(ii), Cr(vi), and Pb(ii), respectively (Table S6). Generally, n values between 0 and 1 indicate favourable adsorption, while values slightly above 1 suggest strong adsorption interactions and near-linear adsorption behaviour.59 Accordingly, the adsorption of Pb(ii) and Cd(ii) onto L1@Fe3O4, with n values greater than 1, may indicate stronger affinity and more uniform adsorption interactions compared to the other systems.59

The Temkin model showed a moderate fitting behaviour (R2 = 0.70117–0.95399) (Table S6), suggesting that adsorbate–adsorbent interactions and a decrease in adsorption heat may have contributed to the extraction process.60 The BT and KT parameters further confirmed the presence of appreciable interactions between the metal ions and the adsorbent surfaces. The isotherm analysis demonstrates that the adsorption of Pb(ii), Cr(vi), and Cd(ii) onto L1@Fe3O4 and L2@Fe3O4 is predominantly governed by heterogeneous multilayer adsorption mechanisms rather than by ideal monolayer adsorption described by the Langmuir model.61

3.5. Mechanism of extraction and coordination chemistry

A plausible mechanism underlying the extraction of Cd(ii), Cr(vi), and Pb(ii) ions by the L2@Fe3O4 and L1@Fe3O4 chelating agents is depicted in Fig. S17. The chelating agents incorporate N, O, and Fe3O4 groups that undergo protonation when exposed to acidic conditions, leading to a positive charge.62 This positive charge facilitates the attraction of negatively charged chromium species through electrostatic interactions.63 Conversely, in basic conditions, these groups become negatively charged due to deprotonation, causing them to attract positively charged metals such as Cd(ii) and Pb(ii) cations.63 Changes in the FT-IR signals of the functional groups of the fresh and used adsorbents were assessed as shown in Fig. S17 and S18. For instance, the v(C Created by potrace 1.16, written by Peter Selinger 2001-2019 N) band of the NN-donor L1@Fe3O4 shifted from 1591 cm−1 in the fresh adsorbent to 1602 cm−1 in the used adsorbent. Similarly, the v(Fe–O) and the v(Si–O) bands shifted 544 cm−1 to 561 cm−1 and 1113 cm−1 to 1046 cm−1 on the fresh and used adsorbent, respectively. A new band was observed at 806 cm−1 on the used adsorbent for L1@Fe3O4. These changes hint at the possible coordination of the metal ions to the N-donor atoms. Similarly, the disappearance of the v(O–H) band at 1324 cm−1 in the spectrum of the NO-donor chelating agent L2@Fe3O4 on the used adsorbent is consistent with deprotonation of the O–H group prior to coordination and eventual extraction of the metal ions.

To further elucidate the mechanism of metal–ligand/adsorbent interactions, we isolated the Cd(ii) complexes of the pre-immobilized homogeneous ligands (L1 and L2) and characterized the complexes using FT-IR (Fig. S19), 1H NMR (Fig. S20) and mass spectrometry (Fig. S21). The FT-IR spectra of the isolated Cd(ii) complexes were consistent with those of the used adsorbents (Fig. S17 and S18). The 1H NMR spectra of the isolated complexes were also helpful in understanding the nature of the metal-adsorbent interactions. For example, in ligand L2, a signal at 8.29 ppm assigned to the O–H proton was not observed in the respective complex [Cd(L2)], pointing to the deprotonation of the phenolic proton prior to coordination to the Cd atom (Fig. S20). On the other hand, the presence of the N–H protons in both ligand L1 and its corresponding complex [Cd(L1)Cl2], at 2.68 ppm and 3.17 ppm, respectively, attested to the absence of deprotonation prior to coordination (Fig. S20). The significant shift of this N–H signal in the Cd complex relative to the ligand further established the coordination of the Cd atom to the N–H motif.

Further information on the nature of the complexes formed was derived from the mass spectral data of the Cd(ii) complex. Both Cd(ii) and Pb(ii) cations are expected to adopt similar coordination behaviour, while a different coordination chemistry exists for the Cr(vi) anions. For instance, the Cd(ii) complex derived from the NO-donor ligand (L2) exhibited a m/z signal at 787 amu [M + Na] and agrees with the formation of a bis(chelated) complex [Cd(L2)], (Fig. S21a). This data is consistent with both the FT-IR and 1H NMR spectral data, where deprotonation of the phenolic proton was evident. On the other hand, the mass spectrum of the Cd(ii) complex of the NN chelating ligand L1 showed an m/z peak at about 786 amu [M+ − Na] and corresponds to the formula of [Cd(L1)2Cl2] (Fig. S21b). From these spectral data, the proposed nature of the coordination compounds formed during the extraction experiments is given in Scheme S2. This is supported by the increased extraction efficiencies of all the metal ions with increased adsorbent dosages of L1@Fe3O4 materials (Fig. 3C), as two equivalents of the adsorbents are required to coordinate one metal atom.

3.6. Comparison of L1@Fe3O4 and L2@Fe3O4 chelating agents with previously reported adsorbents

The current chelating agents (L1@Fe3O4 and L2@Fe3O4) used for the extraction of Cd(ii), Pb(ii), and Cr(vi) ions were compared with similar ligands immobilized on magnetic nanoparticles at optimum concentrations, adsorbent dosages, and contact time reported in Table 3. As an illustration, the extraction performance of adsorbent L2@Fe3O4 exhibited adsorption capacities for Cd(ii) and Pb(ii) of 74.3 mg g−1 74.5 mg g−1 using adsorbent dosage and initial metal concentrations of 1.3 g L−1 and 80 mg L−1, respectively, at a contact time of 120 min (Table 3, entry 7). These values are slightly lower than the adsorption capacities for Cd(ii) and Pb(ii) of 78 mg g−1 and 120 mg g−1 recorded for the previously reported adsorbent VSB-Mag at lower adsorbent dosages (1.0 g L−1) and shorter contact times of 60 min (Table 3, entries 2 vs. 7). On the other hand, adsorbent L2@Fe3O4 exhibited higher extraction capacities at significantly lower adsorbent dosage (1.3 g L−1) for Cd(ii) of 74.3 mg g−1 and Pb(ii) of 74.5 mg g−1, compared to the previously reported adsorbent Fe3O4@SiO2-NH-HCGs, which required a much higher dosage (20 g L−1) to achieved extraction capacities of 45 mg g−1 and 61 mg g−1 for Cd(ii) and Pb(ii) cations respectively (Table 3, entries 3 vs. 7). Both L1@Fe3O4 and L2@Fe3O4 showed much lower adsorption capacities for Cr(vi) of 14.7 mg g−1 and 7.9 mg g−1, respectively, compared to the adsorption capacity of 30.2 mg g−1 displayed by the previously A-M-S-MNPs adsorbent (Table 3, entries 6 vs. 7–8), consistent with the moderate donor abilities of our adsorbents.

Table 3. Comparison of L1@Fe3O4 and L2@Fe3O4 chelating agents on extraction efficiencies and extraction capacities (mg g−1) of Cd(ii), Pb(ii), and Cr(vi) ions with other similar adsorbents.

Entry Adsorbenta pH Time (min) Volume (mL) Adsorbent dosage (g L−1) Initial concentration (mg L−1) Extraction efficiencies (%) Extraction capacity (mg g−1) Ref.
1 Mag-ligand (EDTA) 7 120 50 2.0 10 Cd2+: 95 Cd2+: 79.4 22
15 Pb2+: 97 Pb2+: 100.2
2 VSB-Mag >4.6 60 25 1.0 70 Cd2+: >60 Cd2+: 78 39
7 30 Pb2+: >80 Pb2+: 120
3 Fe3O4@SiO2-NH-HCGs 5–8 20 50 20 10 Cd2+: 96 Cd2+: 45 64
Pb2+: 96 Pb2+: 61
4 Fe3O4@APS@AA-co-CA 5.5 45 50 1.0 100 Cd2+:95 Cd2+: 29.6 65
Pb2+: 95 Pb2+: 166.1
5 MAMNPs 6 100 20 1.5 20 Cd2+: 98 Cd2+: 91.5 66
Pb2+: 98 Pb2+: 118.5
6 A-M-S-MNPs 2 30 25 1.4 50 Cr6+: >80 Cr6+: 30.2 67
7 L2@Fe3O4 12 120 15 1.3 80 Cd2+: 99 Cd2+: 74.3 This work
9 120 80 Pb2+: 99 Pb2+: 74.5
3 60 100 Cr6+: 97 Cr6+: 7.9
8 L1@Fe3O4 9 240 15 1.3 20 Cd2+: 74 Cd2+: 49.5 This work
9 240 20 Pb2+: 90 Pb2+: 48.2
3 60 40 Cr6+: 99 Cr6+: 14.7
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a

Mag-EDTA: magnetic coated with ethylenediaminetetraacetic acid, VSB-Mag: magnetic modified with valine Schiff base, Fe3O4@SiO2-NH-HCGs: magnetic modified with [HCG = py (2-pyridinyl); pyd (3-pyridazinyl); pya (2-pyrazinyl); pym (4-pyrimidinyl)], Fe3O4@APS@AA-co-CA: magnetic nanoparticles modified with 3-aminopropyltriethoxysilane (APS), acrylic acid (AA) and crotonic acid (CA), MAMNPs: mercaptoethylamino monomer modified with magnetic nanoparticles, A-M-S-MNPs: silica magnetite nanoparticles (S-MNPs) embedded in cetyltrimethylammonium bromide (CTAB) resulting in (M-S-MNPs) and modified with 3-aminopropyltriethoxysilane (APTES) to form (A-M-S-MNPs).

4. Conclusions

In summary, NN (L1) and ON (L2) donor ligands were synthesized and immobilized onto magnetic Fe3O4 nanoparticles to give L1@Fe3O4 and L2@Fe3O4 chelating agents. The immobilized materials were characterized by FT-IR, SEM, TEM, EDX, PXRD, BET, VSM, and TGD-DTA/DSC techniques. The chelating agents displayed effective extraction efficiencies (of up to 97%) for Cd(ii), Pb(ii), and Cr(vi) metal ions at pH 9 and 3, contact time of 120 min, initial concentration of 20 ppm, and dosage of 20 mg. The pH of the solutions controlled the extraction efficiencies of each metal cation. Both chelating agents, L1@Fe3O4 and L2@Fe3O4, displayed higher extraction efficiencies in mixed metal ions and naturally heavy metal-polluted waters. Adsorbent L2@Fe3O4 demonstrated appreciable loss of adsorption capacities in the recycling experiments due to possible saturation of the adsorption sites and material clogging. Kinetics isotherms for L2@Fe3O4 and L1@Fe3O4 fitted pseudo-second order isotherms, while the equilibrium data were best described using the Freundlich model.

5. Limitations of the study

While organic ligands offer superior stability and selectivity for heavy metal extraction compared to readily available biomass-based alternatives, their relatively high synthesis cost and the use of organic solvents may limit large-scale application, particularly for domestic water treatment. In addition, material loss during repeated adsorption–desorption cycles presents a challenge, as it may reduce process efficiency and introduce the risk of secondary contamination.

6. Recommendations for future work

Future studies should investigate the effect of temperature on the adsorption process to determine key thermodynamic parameters, including enthalpy, entropy, and Gibbs free energy changes, thereby providing a deeper understanding of the adsorption mechanism. To further evaluate the practical applicability of the adsorbents, vibrating sample magnetometry (VSM) analysis should be conducted on regenerated materials to assess any changes in saturation magnetization (Ms) and confirm the stability of magnetic separation performance over repeated use. In addition, X-ray photoelectron spectroscopy (XPS) analysis before and after metal ion adsorption is recommended to elucidate the binding interactions, oxidation states/speciation, and surface chemical changes associated with the adsorption process. Furthermore, toxicity assessments of the treated water should be performed to evaluate potential environmental risks, including the possibility of secondary pollution arising from ligand leaching or material degradation during reuse. Lastly, future studies should explore more sustainable and environmentally friendly synthetic strategies for the chelating ligands, including the use of green solvents, solvent-free reactions, or bio-based precursors.

Author contributions

S. M. conceptualization, methodology, software, validation and editing, writing – original draft preparation, data acquisition, and analyses. S. O. O.: conceptualization, writing reviews and editing, supervision, and funding. All authors have read and agreed to the published version of the manuscript.

Conflicts of interest

There are no conflicts to declare.

Supplementary Material

RA-016-D6RA00235H-s001

Acknowledgments

This work was supported by the Department of Science and Innovation (South Africa) through the National Research Foundation (NRF), research grant number: SRUG200408511611 (S. O. O). The authors would like to thank NRF for the financial support (bursary to SM). The authors would also like to thank Markus Schmidt (RPTU) for the VSM analyses and Rene Bachmann and Wolfgang Kleist of RTPU-Germany for the BET analyses.

Data availability

The data associated with this manuscript is available as supplementary information (SI). Supplementary information: detailed synthetic procedures, spectral (NMR, FT-IR, MS) and the materials' analytical and physical characterization data. Additional extraction data, graphs, kinetics, and adsorption isotherms. See DOI: https://doi.org/10.1039/d6ra00235h.

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Associated Data

This section collects any data citations, data availability statements, or supplementary materials included in this article.

Supplementary Materials

RA-016-D6RA00235H-s001
RA-016-D6RA00235H-s001.pdf (2.7MB, pdf)

Data Availability Statement

The data associated with this manuscript is available as supplementary information (SI). Supplementary information: detailed synthetic procedures, spectral (NMR, FT-IR, MS) and the materials' analytical and physical characterization data. Additional extraction data, graphs, kinetics, and adsorption isotherms. See DOI: https://doi.org/10.1039/d6ra00235h.

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