Efficient Separation of Heavy Metals by Magnetic Nanostructured Beads

inorganics Article Efficient Separation of Heavy Metals by Magnetic Nanostructured Beads Lisandra de Castro Alves , Susana Yáñez-Vilar * , Yolanda Piñeiro-Redondo and JoséRivas Applied Physic Department, NANOMAG Laboratory, Research Technological Institute, Universidade de Santiago de Compostela (USC), 15782 Santiago de Compostela, Spain; [email protected] (L.d.C.A.); [email protected] (Y.P.-R.); [email protected] (J.R.) *Correspondence: [email protected] Received: 9 April 2020; Accepted: 24 June 2020; Published: 26 June 2020   Abstract: This study reports the ability of magnetic alginate activated carbon (MAAC) beads to remove Cd(II), Hg(II), and Ni(II) from water in a mono-metal and ternary system. The adsorption capacity of the MAAC beads was highest in the mono-metal system. The removal efficiency of such metal ions falls in the range of 20–80% and it followed the order Cd(II) >Ni(II) >Hg(II). The model that best fitted in the ternary system was the Freundlich isotherm, while in the mono-system it was the Langmuir isotherm. The maximum Cd(II), Hg(II), and Ni(II) adsorption capacities calculated from the Freundlich isotherm in the mono-metal system were 7.09, 5.08, and 4.82 (mg/g) (mg/L) 1/n , respectively. Lower adsorption capacity was observed in the ternary system due to the competition of metal ions for available adsorption sites. Desorption and reusability experiments demonstrated the MAAC beads could be used for at least five consecutive adsorption/desorption cycles. These findings suggest the practical use of the MAAC beads as efficient adsorbent for the removal of heavy metals from wastewater. Keywords: heavy metals; magnetite nanoparticles; adsorption; nanocomposite; hybrid; multi-metal; water 1. Introduction Water pollution by heavy metals has become a serious problem due to the adverse effects on ecosystems and human health. More specifically, cadmium (Cd), mercury (Hg), lead (Pb), or nickel (Ni) are known to be highly carcinogenic and mutagenic at low concentrations, and may produce acute toxicity or even dead in living organisms, when present slightly above their allowed limits [ 1 – 3 ]. Although diverse technologies have been developed for the removal of heavy metals from water sources [ 4 ], there is still an urgent need for facile cleaning procedures that ensure high efficiency in the low concentration ranges. Chemical precipitation, ion precipitation, ion exchange, and adsorption are some of the most used techniques due to their potential for scaling up. Among them, the use of natural biopolymers, such as alginate, agarose, chitin and pectin [ 5 – 7 ] in metal biosorption from wastewaters has gained much attention in recent years. Alginate is a polysaccharide derived from brown algae and in the majority of the studies it has been used in the form of calcium alginate beads, due to its practical handling [ 8 ]. Nevertheless, the separation of the loaded biomaterial from the medium is often a problem. To overcome this problem, magnetite nanoparticles (Fe 3 O 4 -NPs) are being incorporated onto the biosorbent matrix [ 9 – 11 ] giving the possibility to magnetically manipulate and separate the hybrid materials from the water matrix. Magnetic alginate beads are a very attractive material with multiple properties such as high specific surface area, rapid recovery, cost-effectiveness, and chemical versatility, for which they are amenable to be combined with materials to increase their affinity for pollutants. Humic acid [ 12 ], Cyanex 302 [ 12 ] and microalgae [ 7 ] were incorporated into alginate beads for their Inorganics 2020,8, 40; doi:10.3390/inorganics8060040 www.mdpi.com/journal/inorganics Inorganics 2020,8, 40 2 of 12 affinity to metal ions. Different authors have shown the ability of several alginate beads composition to uptake metal ions from aqueous solutions. It was observed that magnetic alginate beads containing silica coated with iron carbide nanoparticles enhanced more the adsorption of copper ions than alginate beads alone [ 13 ]. In addition, it was studied that magnetic nanoparticles functionalized with citrate ions present an enhanced adsorption of Pb(II) metal ions from solution [ 14 ]. In one of our previous works [ 15 ], magnetic alginate beads tailored with commercial activated carbon revealed a high capacity for cadmium ions uptake. A 35% removal percentage of cadmium ions was achieved over 1 h with less than 15 mg of adsorbent used. These nanostructured beads revealed to be a great assessment for industrial use, for their high adsorption surface area, easy handling, and magnetic separation from any aqueous media. For these reasons, the same nanostructured beads were used in the present study and tested under more realistic conditions by studying their adsorption capacity when exposed to a mixture of heavy metals. With this aim, different adsorption tests were performed on mono-metal and ternary systems, comprising Cd(II), Hg(II) and Ni(II) metals ions, which are commonly found in waste water from industry and mining effluents. To gain insights in the adsorption mechanism other relevant aspects (e.g., metal-adsorbent mechanisms, metal distribution on beads surface and internal structure, metal desorption, reuse, porosity, among others) were studied in detail. 2. Results and Discussion 2.1. Morphology of Magnetic Beads The morphology of the MAAC beads and commercial activated carbon was examined using scanning electron microscopy (SEM) and transmission electron microscopy (TEM). In Figure 1a, SEM micrograph shows that commercial activated carbon presents a disordered layer-like structure, which was further studied with TEM, reveals a wide variation in the layer size range between 40 µ m to 320 µ m (Figure 1b). SEM micrograph (Figure 1c) of a representative MAAC bead shows a spherical morphology with a porous and layer-like structure on the surface inherited from the precursor commercial activated carbon. The internal structure, as can be observed in Figure 1d, is a combination of activated carbon layers within the interconnected porous network of alginate in MAAC beads. Inorganics 2020, 8, x 2 of 13 were incorporated into alginate beads for their affinity to metal ions. Different authors have shown the ability of several alginate beads composition to uptake metal ions from aqueous solutions. It was observed that magnetic alginate beads containing silica coated with iron carbide nanoparticles enhanced more the adsorption of copper ions than alginate beads alone [13]. In addition, it was studied that magnetic nanoparticles functionalized with citrate ions present an enhanced adsorption of Pb (II) metal ions from solution [14]. In one of our previous works [15], magnetic alginate beads tailored with commercial activated carbon revealed a high capacity for cadmium ions uptake. A 35% removal percentage of cadmium ions was achieved over 1 h with less than 15 mg of adsorbent used. These nanostructured beads revealed to be a great assessment for industrial use, for their high adsorption surface area, easy handling, and magnetic separation from any aqueous media. For these reasons, the same nanostructured beads were used in the present study and tested under more realistic conditions by studying their adsorption capacity when exposed to a mixture of heavy metals. With this aim, different adsorption tests were performed on mono-metal and ternary systems, comprising Cd (II), Hg (II) and Ni (II) metals ions, which are commonly found in waste water from industry and mining effluents. To gain insights in the adsorption mechanism other relevant aspects (e.g., metal-adsorbent mechanisms, metal distribution on beads surface and internal structure, metal desorption, reuse, porosity, among others) were studied in detail. 2. Results and Discussion 2.1. Morphology of Magnetic Beads The morphology of the MAAC beads and commercial activated carbon was examined using scanning electron microscopy (SEM) and transmission electron microscopy (TEM). In Figure 1a, SEM micrograph shows that commercial activated carbon presents a disordered layer-like structure, which was further studied with TEM, reveals a wide variation in the layer size range between 40 µm to 320 µm (Figure 1b). SEM micrograph (Figure 1c) of a representative MAAC bead shows a spherical morphology with a porous and layer-like structure on the surface inherited from the precursor commercial activated carbon. The internal structure, as can be observed in Figure 1d, is a combination of activated carbon layers within the interconnected porous network of alginate in MAAC beads. Figure 1. Commercial activated carbon SEM (a) and TEM (b) micrograph. SEM image of the MAAC bead surface (c) and internal structure (d). 2.2. Structural and Textural Characterization The X-ray diffraction pattern of the MAAC beads in Figure 2a shows the presence of sharp diffraction peaks located at 2θ = 30.1, 35.5, 43.2, 53.5, 57.1, 62.7°. This is the characteristic diffraction pattern arising from the reflection of planes (022), (113), (004), (224), (115) and (044) corresponding to crystalline Fe3O4-NPs embedded in the porous beads. The MAAC bead’s surface was also analyzed by Fourier transform infrared (FT-IR)spectroscopy. As shown in Figure 2b, the peaks at 3228 cm−1 and 1076 cm−1 are related to the -OH and -C-O stretching vibration bands of alginate, respectively, Figure 1. Commercial activated carbon SEM ( a ) and TEM ( b ) micrograph. SEM image of the MAAC bead surface (c) and internal structure (d). 2.2. Structural and Textural Characterization The X-ray diffraction pattern of the MAAC beads in Figure 2a shows the presence of sharp diffraction peaks located at 2 θ =30.1, 35.5, 43.2, 53.5, 57.1, 62.7 ◦ . This is the characteristic diffraction pattern arising from the reflection of planes (022), (113), (004), (224), (115) and (044) corresponding to crystalline Fe 3 O 4 -NPs embedded in the porous beads. The MAAC bead’s surface was also analyzed by Fourier transform infrared (FT-IR)spectroscopy. As shown in Figure 2b, the peaks at 3228 cm −1 and Inorganics 2020,8, 40 3 of 12 1076 cm −1 are related to the –OH and –C–O stretching vibration bands of alginate, respectively, while peaks at 1585 and 1286 cm−1are attributed to the asymmetric and symmetric stretching vibrations of the carboxyl groups of alginate, respectively [ 16 ]. Finally, the band at 558 cm −1 is due to collective vibrations of the magnetite lattice [ 17 ], which confirms the successful incorporation of Fe 3 O 4 -NPs into the alginate matrix. Inorganics 2020, 8, x 3 of 13 while peaks at 1585 and 1286 cm −1 are attributed to the asymmetric and symmetric stretching vibrations of the carboxyl groups of alginate, respectively [16]. Finally, the band at 558 cm −1 is due to collective vibrations of the magnetite lattice [17], which confirms the successful incorporation of Fe 3 O 4 -NPs into the alginate matrix. (a) (b) Figure 2. (a) X-ray diffraction and infrared spectra and (b) IR spectra of MAAC beads. The mean particle size of the Fe 3 O 4 nanoparticles employed for the synthesis of the MAAC beads were calculated from the X-ray diffraction (XRD) pattern (Figure S1) according to the linewidth of the (113) plane refraction peak using Scherrer equation: L = Kλ/βcosθ (1) where L is the mean size of the ordered (crystalline) domains, λ is the X-ray wavelength, β is the width of the XRD peak at half height, K is a shape factor, about 0.9 for magnetite and θ is the Bragg angle The particle diameter calculated from the X-ray diffractogram, using the Scherrer equation was 9.58 nm. By TEM the image (Figure S1) we can observe nanoparticles with a size of 20 nm. The difference between TEM and XRD can be attributed to the fact that “crystallite size” is not synonymous with “particle size”, while XRD is sensitive to the crystallite size inside the particles. 2.3. Specific Surface Are and Pore-Distribution The textural properties of commercial activated carbon (commercial AC) and MAAC beads were analyzed with Brunauer–Emmett–Teller (BET) porosimetry and shown in Table 1. The surface area of commercial AC (849.32 m 2 /g) is within the theoric range for activated carbons (500 to 3.000 m 2 /g) [8]. The surface area of the MAAC beads is 107.13 m 2 /g which is two orders of magnitude higher than the reported 6.25 m 2 /g for calcium alginate beads [18], and can be ascribed to the presence of activated carbon. The pore volume of the MAAC beads (0.075 cm 3 /g) was smaller than the one of commercial AC (0.28 cm 3 /g). In addition, the pore diameter size of the MAAC beads (1.24 nm) is slightly smaller than the commercial AC (1.26 nm). This could be attributed to the deposition of the commercial AC on the surface of the MAAC beads (both surface and internally), leading to a complete filling of the smaller pores. Table 1. Textural parameters of Commercial AC and MAAC beads. Samples BET Surface Area (S BET ) 1 (m 2 / g) Pore Volume (cm 3 / g) Pore diamete r 2 (nm) Commercial AC 849.32 0.28 1.26 MAAC 107.13 0.075 1.24 1 S BET is the BET surface area evaluated at a relative pressure (p/p0) of 0.99. 2 Pore diameter calculated using the Barrett-Joyner-Halenda (BJH) method. Figure 2. (a) X-ray diffraction and infrared spectra and (b) IR spectra of MAAC beads. The mean particle size of the Fe 3 O 4 nanoparticles employed for the synthesis of the MAAC beads were calculated from the X-ray diffraction (XRD) pattern (Figure S1) according to the linewidth of the (113) plane refraction peak using Scherrer equation: L=Kλ/βcosθ(1) where L is the mean size of the ordered (crystalline) domains, λ is the X-ray wavelength, β is the width of the XRD peak at half height, K is a shape factor, about 0.9 for magnetite and θ is the Bragg angle The particle diameter calculated from the X-ray diffractogram, using the Scherrer equation was 9.58 nm. By TEM the image (Figure S1) we can observe nanoparticles with a size of 20 nm. The difference between TEM and XRD can be attributed to the fact that “crystallite size” is not synonymous with “particle size”, while XRD is sensitive to the crystallite size inside the particles. 2.3. Specific Surface Are and Pore-Distribution The textural properties of commercial activated carbon (commercial AC) and MAAC beads were analyzed with Brunauer–Emmett–Teller (BET) porosimetry and shown in Table 1. The surface area of commercial AC (849.32 m 2 /g) is within the theoric range for activated carbons (500 to 3.000 m 2 /g) [ 8 ]. The surface area of the MAAC beads is 107.13 m 2 /g which is two orders of magnitude higher than the reported 6.25 m 2 /g for calcium alginate beads [ 18 ], and can be ascribed to the presence of activated carbon. The pore volume of the MAAC beads (0.075 cm 3 /g) was smaller than the one of commercial AC (0.28 cm 3 /g). In addition, the pore diameter size of the MAAC beads (1.24 nm) is slightly smaller than the commercial AC (1.26 nm). This could be attributed to the deposition of the commercial AC on the surface of the MAAC beads (both surface and internally), leading to a complete filling of the smaller pores. Inorganics 2020,8, 40 4 of 12 Table 1. Textural parameters of Commercial AC and MAAC beads. Samples BET Surface Area (SBET)1(m2/g) Pore Volume (cm3/g) Pore Diameter 2(nm) Commercial AC 849.32 0.28 1.26 MAAC 107.13 0.075 1.24 1 S BET is the BET surface area evaluated at a relative pressure (p/p0) of 0.99. 2 Pore diameter calculated using the Barrett–Joyner–Halenda (BJH) method. 2.4. Magnetic Properties of MAAC Beads Figure 3a shows the variation of magnetization, M, as a function of temperature of MAAC beads in the range 5 to 350 K in an external magnetic field of 100 Oe recorded in zero-field cooling (ZFC) and field cooling (FC). From the curves it is clearly observed the superimposition of the ZFC and FC curves take place at 275 K. The superimposition of ZFC and FC curves is one of the characteristic features of a superparamagnetic system. The magnetic content on the MAAC beads was calculated by thermogravimetric analyses (TGA), which was equal to 23%. Figure 3b illustrates the magnetization curves of bare Fe 3 O 4 -NPs and of the MAAC nanocomposite beads. The saturation of magnetization (Ms) for the synthesized Fe 3 O 4 -NPs (69.23 emu/g) was higher than the observed for the MAAC beads (48.62 emu/g). This may be attributed to the coating effect of alginate trapping the Fe 3 O 4 -NPs in the gel matrix. However, the MAAC beads have superparamagnetic behavior and are easily separated from solution with the help of an external magnetic force. Inorganics 2020, 8, x 4 of 13 2.4. Magnetic Properties of MAAC Beads Figure 3a shows the variation of magnetization, M, as a function of temperature of MAAC beads in the range 5 to 350 K in an external magnetic field of 100 Oe recorded in zero-field cooling (ZFC) and field cooling (FC). From the curves it is clearly observed the superimposition of the ZFC and FC curves take place at 275 K. The superimposition of ZFC and FC curves is one of the characteristic features of a superparamagnetic system. The magnetic content on the MAAC beads was calculated by thermogravimetric analyses (TGA), which was equal to 23%. Figure 3b illustrates the magnetization curves of bare Fe 3 O 4 -NPs and of the MAAC nanocomposite beads. The saturation of magnetization (Ms) for the synthesized Fe 3 O 4 -NPs (69.23 emu/g) was higher than the observed for the MAAC beads (48.62 emu/g). This may be attributed to the coating effect of alginate trapping the Fe 3 O 4 -NPs in the gel matrix. However, the MAAC beads have superparamagnetic behavior and are easily separated from solution with the help of an external magnetic force. (a) (b) Figure 3. ZFC and FC curve recorded at 100 Oe (a) and magnetization curve of MAAC bead and Fe 3 O 4 -NPs at 25 °C (b). 3. Adsorption Study 3.1. Desorption and Reusability The reusability of the MAAC bead was study by repetitive adsorption and desorption cycles of Cd (II) metal ions using 0.01 M HCl solution, as the desorption solution. The q (mg/g) desorption of cadmium was calculated directly from the amount of cadmium adsorbed and amount of cadmium desorbed using Equation (2): 𝑞(𝑑𝑒𝑠𝑜𝑟𝑝𝑡𝑖𝑜𝑛)=(𝑀 −𝑀 )𝑉 𝑀 (2) where M adsorb and M desorb are the adsorbed and desorbed amount of metal ions (mg/g), respectively; V (L) is the volume of desorption solution and M (g) the mass of the MAAC beads used. The desorption percentage was calculated using the following Equation (3): % 𝐷𝑒𝑠𝑜𝑟𝑝𝑡𝑖𝑜𝑛 = 𝑀 𝑀 × 100 (3) The desorption of cadmium from MAAC bead showed a fair desorption percentage over the five cycles, as can be observed in Figure 4. However, the adsorption capacity decreased with increasing regeneration cycle number, except for on cycle number two. The cadmium metal ions have entered inside of the beads structure after the first cycle (Figure S2) resulting in the maximum adsorption capacity on cycle two. Continuous adsorption resulted on internal pore saturation and consequently, on the decreased of adsorption capacity and increase amount of desorbed cadmium metal ions found Figure 3. ZFC and FC curve recorded at 100 Oe ( a ) and magnetization curve of MAAC bead and Fe3O4-NPs at 25 ◦C (b). 3. Adsorption Study 3.1. Desorption and Reusability The reusability of the MAAC bead was study by repetitive adsorption and desorption cycles of Cd(II) metal ions using 0.01 M HCl solution, as the desorption solution. The q (mg/g) desorption of cadmium was calculated directly from the amount of cadmium adsorbed and amount of cadmium desorbed using Equation (2): q(desorption)=(Madsorb −Mdesorb)V M(2) Inorganics 2020,8, 40 5 of 12 where M adsorb and M desorb are the adsorbed and desorbed amount of metal ions (mg/g), respectively; V(L) is the volume of desorption solution and M(g) the mass of the MAAC beads used. The desorption percentage was calculated using the following Equation (3): %Desorption =Mdesorb Madsorb ×100 (3) The desorption of cadmium from MAAC bead showed a fair desorption percentage over the five cycles, as can be observed in Figure 4. However, the adsorption capacity decreased with increasing regeneration cycle number, except for on cycle number two. The cadmium metal ions have entered inside of the beads structure after the first cycle (Figure S2) resulting in the maximum adsorption capacity on cycle two. Continuous adsorption resulted on internal pore saturation and consequently, on the decreased of adsorption capacity and increase amount of desorbed cadmium metal ions found in solution. Furthermore, the results suggest a reduction of the metal ions from the matrix in the first fourth cycles and more than 50% of metal was recovered. After the fifth cycle, the desorption percentage decreased to 20%, since cadmium metal ions remained inside of the beads structure after the third desorption cycle. Thus, it can be said that these beads have the potential to be reuse up to four cycles under the chosen conditions. For further studies, factors, such as time and desorbent concentration must be considered for maximum desorption capacity for more than 5 cycles. In previously studies [ 19 ], hydrochloric acid proved to be a good desorbent within 2 h of repetitive adsorption and desorption cycles. Inorganics 2020, 8, x 5 of 13 in solution. Furthermore, the results suggest a reduction of the metal ions from the matrix in the first fourth cycles and more than 50% of metal was recovered. After the fifth cycle, the desorption percentage decreased to 20%, since cadmium metal ions remained inside of the beads structure after the third desorption cycle. Thus, it can be said that these beads have the potential to be reuse up to four cycles under the chosen conditions. For further studies, factors, such as time and desorbent concentration must be considered for maximum desorption capacity for more than 5 cycles. In previously studies [19], hydrochloric acid proved to be a good desorbent within 2 h of repetitive adsorption and desorption cycles. Figure 4. Adsorption and Desorption cycles of cadmium from MAAC beads (mean ± SE of 15 replicates). 3.2. Effect of pH The influence of pH on Cd (II), Hg (II) and Ni (II) adsorption capacity by the MAAC beads was studied at pH range of (2.0 to 9.0). In Figure 5, the adsorbed metal ions per adsorbent mass q (mg/g) are presented for all tested pH values. As can be seen, the adsorption process for Cd (II) and Ni (II), was constant between the pH values of 4.5 to 7.0, followed by an increase at higher pH values. On contrary, Hg (II) adsorption depends in a non-predictable way on the pH, attaining a maximum adsorption capacity at pH 4.5, followed by a steep decrease at pH 6.5 and a large increment up to pH 9.0. The higher concentration of hydrolyzed ions such as H + allow an enhanced binding of Hg (II) metal ions to the sodium alginate surface [19]. Besides this binding mechanism being more enhanced at pH 2.0 for Hg (II) ions, at pH 4.5 all metal ions revealed to have a constant and high adsorption capacity at this pH, being the selected for the experiments at the mono and ternary system. The increasing adsorption at higher pH values for all metal ions may be attributed to the formation of hydroxyl ions [20]. Figure 5. Effect of initial pH on Cd (II), Hg (II) and Ni (II) adsorption onto MAAC beads (mean standard deviation ± 0.3). Figure 4. AdsorptionandDesorptioncyclesofcadmiumfromMAACbeads(mean ± SEof15replicates). 3.2. Effect of pH The influence of pH on Cd(II), Hg(II) and Ni(II) adsorption capacity by the MAAC beads was studied at pH range of (2.0 to 9.0). In Figure 5, the adsorbed metal ions per adsorbent mass q (mg/g) are presented for all tested pH values. As can be seen, the adsorption process for Cd(II) and Ni(II), was constant between the pH values of 4.5 to 7.0, followed by an increase at higher pH values. On contrary, Hg(II) adsorption depends in a non-predictable way on the pH, attaining a maximum adsorption capacity at pH 4.5, followed by a steep decrease at pH 6.5 and a large increment up to pH 9.0. The higher concentration of hydrolyzed ions such as H + allow an enhanced binding of Hg(II) metal ions to the sodium alginate surface [ 19 ]. Besides this binding mechanism being more enhanced at pH 2.0 for Hg(II) ions, at pH 4.5 all metal ions revealed to have a constant and high adsorption capacity at this pH, being the selected for the experiments at the mono and ternary system. The increasing adsorption at higher pH values for all metal ions may be attributed to the formation of hydroxyl ions [20]. At higher pH values, the decreased adsorption capacity by Cd(II) coincides with the decreasing concentration of Cd(II) and the precipitation of Cd(OH) 2 into solution. The Cd(OH) 2 ionic species are adsorbed by the MAAC beads occupying the available sites and preventing the further adsorption of Cd(II) ions [ 21 ]. For nickel an increasing trend at higher pH values was observed. This characteristic trend is attributed to the specific adsorption of cationic hydroxo-complexes, which is the pH range Inorganics 2020,8, 40 6 of 12 more favorable for the formation of these species [ 22 ]. The hydrolysis of mercury, on the other hand, begins at very low pH values (pH <4.5) with the formation of Hg(OH) + and at pH values (pH >4.5) the adsorption capacity suddenly increases because of the formation of mercury neutral species Hg(OH)2[23]. Inorganics 2020, 8, x 5 of 13 in solution. Furthermore, the results suggest a reduction of the metal ions from the matrix in the first fourth cycles and more than 50% of metal was recovered. After the fifth cycle, the desorption percentage decreased to 20%, since cadmium metal ions remained inside of the beads structure after the third desorption cycle. Thus, it can be said that these beads have the potential to be reuse up to four cycles under the chosen conditions. For further studies, factors, such as time and desorbent concentration must be considered for maximum desorption capacity for more than 5 cycles. In previously studies [19], hydrochloric acid proved to be a good desorbent within 2 h of repetitive adsorption and desorption cycles. Figure 4. Adsorption and Desorption cycles of cadmium from MAAC beads (mean ± SE of 15 replicates). 3.2. Effect of pH The influence of pH on Cd (II), Hg (II) and Ni (II) adsorption capacity by the MAAC beads was studied at pH range of (2.0 to 9.0). In Figure 5, the adsorbed metal ions per adsorbent mass q (mg/g) are presented for all tested pH values. As can be seen, the adsorption process for Cd (II) and Ni (II), was constant between the pH values of 4.5 to 7.0, followed by an increase at higher pH values. On contrary, Hg (II) adsorption depends in a non-predictable way on the pH, attaining a maximum adsorption capacity at pH 4.5, followed by a steep decrease at pH 6.5 and a large increment up to pH 9.0. The higher concentration of hydrolyzed ions such as H + allow an enhanced binding of Hg (II) metal ions to the sodium alginate surface [19]. Besides this binding mechanism being more enhanced at pH 2.0 for Hg (II) ions, at pH 4.5 all metal ions revealed to have a constant and high adsorption capacity at this pH, being the selected for the experiments at the mono and ternary system. The increasing adsorption at higher pH values for all metal ions may be attributed to the formation of hydroxyl ions [20]. Figure 5. Effect of initial pH on Cd (II), Hg (II) and Ni (II) adsorption onto MAAC beads (mean standard deviation ± 0.3). Figure 5. Effect of initial pH on Cd(II), Hg(II) and Ni(II) adsorption onto MAAC beads (mean standard deviation ±0.3). 3.3. Adsorption of Mono and Ternary Systems Batch adsorption experiments were performed using a defined amount of MAAC beads in mono and ternary systems containing Cd(II), Ni(II) and Hg(II) metal ions. The effect of the initial metal ion concentration on adsorption was studied for solutions prepared with a set of concentrations 10 to 250 mg/L, and applying previously optimized conditions (magnetic agitation at 300 rpm, using 14 mg of adsorbent during 6 h) at fixed pH =4.5. The equilibrium adsorption capacity ( qe ) was calculated according to the following equation [24]: qe=(C0−Ce)V M(4) The removal efficiency (R%) of Cd(II), Hg(II) and Ni(II) in the mono-component system was calculated using the following equation: R%=(C0−Ce) C0 ×100 (5) where qe is the equilibrium adsorption capacity (mg/g); C0 and Ce are the initial and equilibrium concentration (mg/L) of metal ions, respectively; Vis the volume of working solution (L) and Mis the weight (g) of adsorbent used. Alternatively, q, can be expressed in terms of molarity, q(mol/g), to gain insights into the number of moles (atoms) that adsorb on the cleaning beads (Figure S3). In Figure 6, the adsorption capacity, q(mg/g), of Cd(II), Ni(II) and Hg(II) by the MAAC beads, for the mono and ternary metal adsorption batch tests are presented versus the initial concentration in the batch solution of each metal. From all tested metals under the current experimental conditions, it is evident that Cd(II) ions are preferentially adsorbed in both tests, being in the mono-metal case (Figure 6a) the adsorption capacity of Cd(II) ions twice the adsorption observed for Ni(II) and Hg(II). In addition, in the mono-metal test, the adsorption of cadmium and nickel ions increases uniformly indicating ongoing adsorption process onto the MAAC beads, on contrary to mercury, which shows a decreased adsorption at the initial concentration of 150 mg/L. In the ternary metal system (Figure 6b), the adsorption capacity is generally smaller than in the mono-metal case, indicating a competition between the metal ions in the ternary system for available binding sites on the MAAC beads [ 25 ], with a striking exception at initial metal concentration of C=150 mg/L, where all metal ions are adsorbed Inorganics 2020,8, 40 7 of 12 with more efficacy. Moreover, the similarities of the adsorption capacity curve of all ions in the ternary system with the mercury adsorption feature in the mono-metal system suggests that mercury has an important role modulating the adsorption mechanism, which will deserve future studies. The removal efficiency decreases when the initial metal ions concentration is increased (Figure 7), a trend that was already observed in a previous study [ 15 ]. This behavior is mainly ascribed to the saturation of the available binding sites during the adsorption process, leading to a reduction of the adsorption capacity. Inorganics 2020, 8, x 6 of 13 At higher pH values, the decreased adsorption capacity by Cd (II) coincides with the decreasing concentration of Cd (II) and the precipitation of Cd(OH)2 into solution. The Cd(OH)2 ionic species are adsorbed by the MAAC beads occupying the available sites and preventing the further adsorption of Cd (II) ions [21]. For nickel an increasing trend at higher pH values was observed. This characteristic trend is attributed to the specific adsorption of cationic hydroxo-complexes, which is the pH range more favorable for the formation of these species [22]. The hydrolysis of mercury, on the other hand, begins at very low pH values (pH < 4.5) with the formation of Hg(OH)+ and at pH values (pH > 4.5) the adsorption capacity suddenly increases because of the formation of mercury neutral species Hg(OH)2 [23]. 3.3. Adsorption of Mono and Ternary Systems Batch adsorption experiments were performed using a defined amount of MAAC beads in mono and ternary systems containing Cd (II), Ni (II) and Hg (II) metal ions. The effect of the initial metal ion concentration on adsorption was studied for solutions prepared with a set of concentrations 10 to 250 mg/L, and applying previously optimized conditions (magnetic agitation at 300 rpm, using 14 mg of adsorbent during 6 h) at fixed pH = 4.5. The equilibrium adsorption capacity (𝑞) was calculated according to the following equation [24]: 𝑞=(𝐶−𝐶 )𝑉 M (4) The removal efficiency (R%) of Cd (II), Hg (II) and Ni (II) in the mono-component system was calculated using the following equation: 𝑅% = (𝐶−𝐶 ) 𝐶 × 100 (5) where 𝑞 is the equilibrium adsorption capacity (mg/g); 𝐶 and 𝐶 are the initial and equilibrium concentration (mg/L) of metal ions, respectively; V is the volume of working solution (L) and M is the weight (g) of adsorbent used. Alternatively, q, can be expressed in terms of molarity, q (mol/g), to gain insights into the number of moles (atoms) that adsorb on the cleaning beads (Figure S3). In Figure 6, the adsorption capacity, q (mg/g), of Cd (II), Ni (II) and Hg (II) by the MAAC beads, for the mono and ternary metal adsorption batch tests are presented versus the initial concentration in the batch solution of each metal. From all tested metals under the current experimental conditions, it is evident that Cd (II) ions are preferentially adsorbed in both tests, being in the mono-metal case (Figure 6a) the adsorption capacity of Cd (II) ions twice the adsorption observed for Ni (II) and Hg (II). (a) (b) Figure 6. Effect of initial metal concentration on the adsorption capacity of MAAC beads for (a) monometal and (mean SD ± 1.45) (b) ternary experiment at pH 4.5 (mean SD ± 1.27). Figure 6. Effect of initial metal concentration on the adsorption capacity of MAAC beads for (a) mono-metal and (mean SD ±1.45) (b) ternary experiment at pH 4.5 (mean SD ±1.27). Inorganics 2020, 8, x 7 of 13 In addition, in the mono-metal test, the adsorption of cadmium and nickel ions increases uniformly indicating ongoing adsorption process onto the MAAC beads, on contrary to mercury, which shows a decreased adsorption at the initial concentration of 150 mg/L. In the ternary metal system (Figure 6b), the adsorption capacity is generally smaller than in the mono-metal case, indicating a competition between the metal ions in the ternary system for available binding sites on the MAAC beads [25], with a striking exception at initial metal concentration of C = 150 mg/L, where all metal ions are adsorbed with more efficacy. Moreover, the similarities of the adsorption capacity curve of all ions in the ternary system with the mercury adsorption feature in the mono-metal system suggests that mercury has an important role modulating the adsorption mechanism, which will deserve future studies. The removal efficiency decreases when the initial metal ions concentration is increased (Figure 7), a trend that was already observed in a previous study [15]. This behavior is mainly ascribed to the saturation of the available binding sites during the adsorption process, leading to a reduction of the adsorption capacity. Figure 7. Removal efficiency or mono-metal and ternary experiments at pH 4.5. 3.4. Competitive Adsorption Evaluation The interactive effect of Cd (II), Hg (II) and Ni (II) was investigated in the ternary system. For that, an evaluation ratio was introduced to assess the type of adsorption competition between each metal ion in the system and combined [26]. The evaluation ratio is expressed by the following Equation (6): 𝐸=𝑄 󰆒 𝑄 (6) where 𝐸, is the evaluation ratio; 𝑄 󰆒 (mg/g) is the amount of metal ions adsorbed in a ternary system and 𝑄 (mg/g) is the amount of metal ions adsorbed in the mono-metal system. If the evaluation ratio, E > 1, the presence of other metal ions have enhanced the adsorption of other metal ions in solution (synergism effect); when E = 1, this means that the presence of other metal ions would not influence the adsorption of another metal ion; and when E < 1, the presence of another metal ion would suppress the adsorption of one another (antagonism effect). The evaluation ratios of individual Cd (II), Hg (II) and Ni (II) in the ternary system and the whole mixture are listed in Table 2. Figure 7. Removal efficiency or mono-metal and ternary experiments at pH 4.5. 3.4. Competitive Adsorption Evaluation The interactive effect of Cd(II), Hg(II) and Ni(II) was investigated in the ternary system. For that, an evaluation ratio was introduced to assess the type of adsorption competition between each metal ion in the system and combined [ 26 ]. The evaluation ratio is expressed by the following Equation (6): E=Q0 e Qe(6) where E , is the evaluation ratio; Q0 e (mg/g) is the amount of metal ions adsorbed in a ternary system and Qe (mg/g) is the amount of metal ions adsorbed in the mono-metal system. If the evaluation ratio, E>1, the presence of other metal ions have enhanced the adsorption of other metal ions in solution (synergism effect); when E=1, this means that the presence of other metal ions would not influence the Inorganics 2020,8, 40 8 of 12 adsorption of another metal ion; and when E<1, the presence of another metal ion would suppress the adsorption of one another (antagonism effect). The evaluation ratios of individual Cd(II), Hg(II) and Ni(II) in the ternary system and the whole mixture are listed in Table 2. Table 2. Individual and sum of the evaluation ratios of Cd(II), Hg(II) and Ni(II) in the ternary system for MAAC beads. Initial Concentration (mg/L) 10 70 150 200 250 Adsorbent Metal Ions Evaluation Ratios MAAC Cd(II) 1.08 0.58 1.10 0.35 0.56 Hg(II) 0.14 0.66 1.46 0.44 1.36 Ni(II) 1.05 0.63 1.38 0.20 0.26 Cd(II) +Hg(II) +Ni(II) 0.68 0.71 1.48 0.32 0.48 Cadmium and nickel in both mono and ternary systems were most efficiently adsorbed than mercury (Figure 6a,b). On contrary, the study of the evaluation ratios in terms of competitive scenario reveals that mercury ions are the least suppressed of all three metal ions. The results compiled in Table 2, reveal an antagonistic effect (E<1) almost in all concentrations in the individual metals and on the whole mixture of the ternary system, except at C=150 (mg/L), where E=1.48 (E>1), indicating an average synergistic behavior between the metal ions. For high values of initial metal concentrations of C=200 (mg/L) and C=250 (mg/L), the degree of suppression is large for nickel and cadmium ions, while mercury adsorption is favored by a synergistic effect. Moreover, for C ≥ 70 (mg/L) mercury is the least suppressed metal and benefits from an enhancement effect, in part due to the increase electrostatic repulsion among the cations that would limit the adsorption of the metal ion [27]. 3.5. Adsorption Isotherms To determine the adsorption capacity of the MAAC beads in a mono-metal and ternary system, adsorption studies with initial concentrations ranging from C=10 (mg/L), to C=250 (mg/L), were carried out. The adsorption equilibrium was studied fitting the experimental data to the linear equations of Langmuir and Freundlich isotherm models. Langmuir isotherm model [28]: Ce qe =1 KLqm +Ce qm (7) Freundlich isotherm model [29]: logqe=logKF+1 nlogCe(8) where qe (mg/g) is the amount of metal ions adsorbed; Ce (mg/L) is the adsorbate concentration in solution, both at equilibrium; K L (L/mg) is the Langmuir adsorption constant; and qm (mg/g) is the maximum adsorption capacity for monolayer formation on the adsorbent. The value KF can be defined as the adsorption or distribution coefficient and represents the quantity of metal ions adsorbed onto the beads. The value of 1 /n indicates surface heterogeneity, which becomes more heterogeneous as its value gets closer to zero. A fundamental characteristic of the Langmuir isotherm is to predict the affinity between sorbate and sorbent using a dimensionless constant, known as separation factor RL , which can be represented as: RL=1 1+KLC0(9) where C0 (mg/L) is the adsorbate initial concentration. The value of RL stands between 0 and 1 for favorable adsorption, while RL >1 represents unfavorable adsorption, RL =1 represents linear adsorption and RL=0 for irreversible adsorption processes [28]. Inorganics 2020,8, 40 9 of 12 Table 3lists the parameters of Langmuir and Freundlich isotherms models computed from the experimental tests using MAAC beads for mono-metal and ternary systems adsorption. Table 3. Langmuir and Freundlich isotherm parameters. Adsorption System Metal Ions Langmuir Parameters Freundlich Parameters qm (mg/g) KL (L/mg) RLR2n1/nKF (mg1−(1/n)L1/ng−1)R2 Cd(II) Cd(II) 59.17 0.048 0.103 0.981 2.52 0.397 7.09 0.983 Hg(II) Hg(II) 25.00 0.040 0.118 0.922 3.94 0.254 5.08 0.865 Ni(II) Ni(II) 37.04 0.031 0.144 0.945 2.79 0.358 4.82 0.995 Cd(II) +Hg(II) + Ni(II) Cd(II) 56.18 0.029 0.339 0.758 2.49 0.401 6.11 0.889 Hg(II) 181.82 0.007 0.164 0.681 1.50 0.668 3.35 0.901 Ni(II) 172.71 0.001 0.456 0.561 1.14 0.878 0.35 0.953 The models of Langmuir and Freundlich equations in general described the data well, although the mono-metal system seems to be better described by Langmuir while the ternary system by the Freundlich isotherm, as suggested by the correlation coefficient values R 2 . The Langmuir constant, K L values were higher for the mono-metal system, revealing a higher adsorption capacity. Contrary to the ternary system, where the presence of other metal ions in the system decreased the adsorption capacity due to the competition of available adsorption sites onto the MAAC beads. The R L parameter values stand below 1, indicating favorable and weakly reversible adsorption of studied metal ions onto MAAC beads. The values of ndetermined with the Freundlich equation were generally higher than 1.0, indicating heterogeneous adsorption process for all metal ions onto the beads. Among all metals, Cd(II) was the most highly adsorbed with a Freundlich KF constant of 7.09 and 6.11 for both mono-metal and ternary system, respectively. In our study, the adsorption cannot be simply related to the physicochemical properties of metal cations, since cadmium with higher atomic weight and ionic radius than nickel was more intensely adsorbed [ 30 ]. Thus, adsorption in the ternary system was better described by the Freundlich adsorption isotherm, revealing heterogeneous surface with different affinity sites on MAAC beads. 4. Materials and Methods 4.1. Synthesis of MAAC Beads TheprocedureoftheMAACbeadssynthesisisdescribedinourpreviousstudy[ 15 ]. Briefly,Fe 3 O 4 -NPs were synthesize by reverse coprecipitation method [ 16 ]. In the synthesis, 15 mL of 1.0 M FeCl 3· 6H 2 O (Alfa Aesar, Madrid, Spain) and 0.5 M FeSO 4· 7H 2 O (Sigma, St. Louis, MO, USA) were mixed and added dropwise into a 3.5 M NH 4 OH solution of 20 mL at 60 ◦ C. The reaction proceeded for 30 min under mechanical agitation. The magnetic nanoparticles were then washed and re-dispersed in distilled water. Beads were prepared in cross-linking solution using calcium chloride solutions as the cross-linking agent. Next, 2.0 g of sodium alginate (Sigma, St. Louis, MO, USA) was subsequently added to the previously prepared Fe 3 O 4 -NPs solution (35 mL). After obtaining a homogeneous solution, 3.0 g of commercial activated carbon was added and the solution was mechanically agitated for 4 h. The obtained suspension was added dropwise into a previously prepared bath of 0.13 M CaCl 2 (Sigma, St. Louis, MO, USA) and 450 µ L Tween 20 (Fluka, Steinheim, Germany) under continuous magnetic speed of 450 rpm using a New Era NE-300 syringe pump (Biogen, Madrid, Spain). Beads were instantly formed and were left in the bath around 30 min for hardening. Afterwards, the beads were collected with a magnet and cleaned with distilled water. Finally, the magnetic beads were dried at 60 ◦C. 4.2. Effect of pH The effect of pH on adsorption capacity of Cd(II), Hg(II) and Ni(II) metal ions was conducted individually by mixing (14 mg) of adsorbent with 20 mL of 10 mg/L −1 metal ions concentration. The adsorption capacity was studied at pH values of (2.0, 3.0, 4.5, 5.0, 6.5, 7.0, 9.0) under magnetic