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Article

Synthesis of Hydroxyapatite Mulberry Stem Biochar Composites for Efficient Pb(II) Adsorption from Aqueous Solutions

by
Dunqiu Wang
1,2,3,4,5,
Xinyu Zhou
1,2,3,4,5,
Meina Liang
1,2,3,4,5 and
Zimeng Wu
1,2,3,4,5,6,*
1
Guangxi Key Laboratory of Environmental Pollution Control Theory and Technology, Guilin University of Technology, Guilin 541004, China
2
Engineering Research Center of Watershed Protection and Green Development, Guilin University of Technology, Guilin 541006, China
3
Guangxi Engineering Research Center of Comprehensive Treatment for Agricultural Non-Point Source Pollution, Guilin University of Technology, Guilin 541006, China
4
Modern Industry College of Ecology and Environmental Protection, Guilin University of Technology, Guilin 541006, China
5
Collaborative Innovation Center for Water Pollution Control and Water Safety in Karst Area, Guilin University of Technology, Guilin 541006, China
6
Yunnan Xinlian Environmental Technology Co., Ltd., Kunming 650500, China
*
Author to whom correspondence should be addressed.
Water 2025, 17(9), 1389; https://doi.org/10.3390/w17091389
Submission received: 11 April 2025 / Revised: 30 April 2025 / Accepted: 30 April 2025 / Published: 5 May 2025
(This article belongs to the Special Issue Adsorption Technologies in Wastewater Treatment Processes)
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Versions Notes

Abstract

:
In this study, two biochar composites, namely hydroxyapatite/mulberry stem biochar (HMp) and magnesium-doped HMp (Mg0.1-HMp), were prepared using mulberry stem as the major raw material using the sol–gel process. Characterization and batch experiments were carried out on HMp and Mg0.1-HMp to investigate the Pb(II) adsorption mechanism and the factors affecting the adsorption, respectively. The results indicated that carboxylic compounds, phenols, and carbonyl functional groups were formed on the surfaces of HMp and Mg0.1-HMp. At an optimal pH of 5, an adsorption period of 6 h was achieved at an initial Pb(II) concentration of 100 mg/L and adsorbent quantity of 2 g/L. The maximum Pb(II) adsorption capacities of the HMp and Mg0.1-HMp were 303.03 and 312.50 mg/g, respectively, at 25 °C. The maximum Pb(II) adsorption capacity of Mg0.1-HMp was 2.55 times more than that of mulberry stem biochar (MBC). The adsorption of Pb(II) by HMp and Mg0.1-HMp is consistent with the Langmuir isotherm and pseudo-second-order kinetic models, demonstrating a spontaneous, endothermic, and irreversible process dominated by monolayer chemical adsorption. These results show that the mechanisms of Pb(II) by Mg0.1-HMp mainly involved electrostatic interaction, complexation, precipitation, and ion exchange.

Graphical Abstract

1. Introduction

Pb is extensively used for various applications and it is a hazardous, unpredictable, and persistent heavy metal that is present in various environments [1,2,3]. Pb may enter the human body through the food chain and affect the human body [4]. Animal and human health issues such as anemia, kidney failure, mental retardation, and physiological defects are significantly correlated to Pb contamination [5,6]. Many methods such as coagulation, chemical precipitation, solvent extraction, membrane separation, reverse osmosis, ion exchange, and adsorption have been used to remove heavy metals, including Pb(II), from water and wastewater [7,8,9,10]. Among these, owing to its effectiveness, low cost, and simple operation, adsorption is widely employed to remove Pb from dilute aqueous solutions [11].
The International Biochar Initiative defines biochar as a solid material produced by the thermochemical conversion of biomass under oxygen-limited conditions [12]. Waste biomass from various sources [13], such as sawdust, straw, fruit shells, water hyacinth, reed, animal feces, vegetable waste, paper, sludge, and biogas residues, has been used to produce biochar [14,15]. Biomass has a wide range of sources, low price, renewability, and low pollution, and hence, it is suitable for the preparation of heavy-metal adsorbents. For example, under optimal conditions, soybean root biochar can effectively adsorb a quantity of 138.54 mg/g of Cd(II) [16]. The maximum adsorption capacity of straw iron-modified biochar, as a novel adsorbent for As(V), was 28.49 mg/g [17]. The maximum Cr(VI) adsorption capacity of Fe-Mn binary oxide MBC was 37.14 mg/g [18]. The Pb adsorption capacities of biochar derived from different biomass sources significantly vary. The maximum Pb(II) adsorption capacities of wheat straw biochar and buffalo weed biochar were 189.24 and 212.31 mg/g, respectively [19,20,21]. The maximum Pb(II) adsorption capacities of magnetic biochar derived from white-tea trash and sludge biochar were 81.6 and 51.20 mg/g, respectively [22,23]. The results of these studies show that Pb(II) removal efficiencies vary based on the type of biochar. Therefore, absorbents with strong adsorption capabilities for Pb covering a wide concentration range are needed.
Hydroxyapatite (HAp) is a novel and widely researched functional material [24]. Some researchers have discovered that HAp has the capacity to eliminate metal ions such as Pb2+, Cd2+, Cu2+, Zn2+, Co2+, and Ni2+ [25,26,27,28,29,30]. Consequently, several researchers have coupled biochar and HAp to increase the ability to adsorb Pb(II). With a Pb(II) removal efficiency of over 99.5%, the produced HAp rice husk biochar composite outperformed the original rice husk biochar by 24.12%, with a maximum adsorption capacity of 102.2 mg/g [31]. Many limitations in the practical application of HAp biochar for heavy-metal removal exist such as its small specific surface area and fewer active sites compared to those in pure apatite, resulting in a low adsorption capacity and efficiency. However, HAp biochar does not cause secondary pollution of the environment [32].
Consequently, HAp biochar was altered to enhance its capacity to absorb Pb. A HAP-functionalized graded porous/Hami melon biochar composite material (HB-0.5HAp) was prepared using a filtrate rich in Ca2+ ions reusing steel slag as a modifier. Its maximum Pb(II) adsorption capacity was 134.4 mg/g, which was 1.77 times that of unmodified Hami melon biochar [33].
As a major sericulture province, Guangxi produces a large number of discarded mulberry stalks that urgently need to be effectively developed, so they need immediate attention. Pb is a typical contaminant in wastewater, and the current study aims to prepare a novel cost-effective biochar adsorbent to remove Pb. In order to reduce the cost of biochar preparation, mulberry rods are used as raw materials. Mulberry trees have a significant capacity to accumulate Cu, Pb, and Cd. Heavy metals in the atmosphere may also be absorbed by their organs, such as roots, stems, bark, and leaves [34]. Therefore, the adsorption of Pb may be substantially improved by employing mulberry stems as the biomass raw material.
In this study, we prepared MBC, HMp, and Mg0.1-HMp and studied their Pb(II) adsorption characteristics. Factors such as the contact time, pH, and Pb(II) concentration on the Pb(II) adsorption capacities of MBC, HMp, and Mg0.1-HMp were studied in batch experiments. The Pb(II) adsorption characteristics of MBC, HMp, and Mg0.1-HMp were analyzed using the Langmuir adsorption isotherm and pseudo-first-order and -second-order, Bangham, and Elovich kinetics. Mg0.1-HMp was characterized using a scanning electron microscope equipped with an energy-dispersive spectrometer (SEM-EDS), the Brunauer–Emmett–Teller (BET) method, Fourier-transform infrared spectroscopy (FT-IR), X-ray diffraction (XRD), and X-ray photoelectron spectroscopy (XPS) before and after the adsorption of Pb(II) to explore the mechanisms of Pb(II) adsorption onto Mg0.1-HMp.

2. Materials and Methods

2.1. Materials

The mulberry stem for the experiment was collected in Huanjiang County, Hechi City, Guangxi Province, China (24°57′16.30″ N, 108°28′16.57″ E). The peeled and crushed mulberry stem was placed in an oven and dried at 80 °C for 24 h, ground, and passed through a 20-mesh sieve to obtain mulberry stem biomass for later use. Chemicals such as NH3∙H2O, Ca(NO3)2∙4H2O, (NH4)2HPO4, and Mg(NO3)2∙6H2O were obtained from Xilong Chemical Co., Ltd. (Shenzhen, China) Ultrapure water was used in the laboratory and the reagents were analytically pure.

2.2. Preparation of MBC, HMp, and Mg0.1-HMp

Preparation of MBC: the mulberry stem biomass was placed in a muffle furnace, calcined at 700 °C for 2 h, and passed through a 100-mesh sieve to obtain MBC.
Preparation of HMp: A combination of 500 mL of 0.2 M Ca(NO3)2 solution and 300 mL of 0.2 M (NH4)2HPO4 solution with a Ca/P molar ratio of 1.67 was taken and 50 g of MBC was added to the mixture. An 8% NH3∙H2O solution was used to adjust the pH of the mixture to 10.0 and kept for 6 h in a 70 °C water bath after stirring for 30 min at room temperature. The supernatant was centrifuged and washed with ultrapure water after natural cooling to achieve a pH of approximately 7.0. Anhydrous ethanol was used for washing and filtration. The mixture was dried at 80 °C, ground, and sieved through a 100-mesh sieve to obtain HMp.
Preparation of Mg0.1-HMp: A combination of 500 mL of 0.2 M Ca(NO3)2 solution, 50 mL of 0.2 M Mg(NO3)2 solution, and 300 mL of 0.2 M (NH4)2HPO4 solution with a Ca/P molar ratio of 1.67 was taken and 50 g of MBC was added to the mixture. The remaining procedures were identical to those for HMp preparation. Thus, Mg0.1-HMp was obtained.

2.3. Characterization

A Thermo Scientific spectrometer (Thermo Nexus 470FT-IR, Nicorette, Waltham, MA, USA) was used to obtain FT-IR spectra of different materials. The XRD patterns of MBC, HMp, and Mg0.1-HMp adsorbents were recorded using an X’ Pert PRO (Bruker, Germany) with Cu Kα radiation at a scan rate of 0.2°/min (λ = 1.5406 Å, 40 kV, and 40 mA). An SEM-EDS (Hitachi FE-SEMS-4800, Hitachi, Tokyo, Japan) was used to determine surface morphologies. In addition, an XPS (ESCALAB 250Xi, Thermoelectric Corporation, Waltham, MA, USA) was used to study the surface chemical states and elemental compositions of different materials. The specific surface area and porosity were analyzed using BET Surface Area Analyzer (NOVAe1000, Quantachrome, Boynton Beach, FL, USA). The zeta potential values and particle sizes were measured using a zeta potential analyzer (Nano ZS90, Malvern Instruments, Malvern City, UK).

2.4. Batch Experiments

The adsorption procedure was as follows: 0.100 g HMp or Mg0.1-HMp was added to a 100 mL polyethylene plastic centrifuge tube; then, 50.0 mL of Pb(II)-containing solution was added to adjust the pH to a set value by adding 0.01 M HNO3 or 0.01 M NaOH. The tubes were shaken at the desired temperature and duration at 200 rpm in a water bath oscillator. The mixture was separated using centrifugation, and the supernatant was carefully filtered using a 0.45 µm syringe membrane filter. Then, the concentration of Pb(II) was measured using an inductively coupled plasma optical emission spectrometer (ICP-OES).
The effect of pH was studied at pH values of 2.0, 3.0, 4.0, 5.0, 6.0, 7.0, 8.0, 9.0) and the Pb(II)-containing solutions (10, 100, or 500 mg/L) were mixed with HMp or Mg0.1-HMp and shaken for 24 h at 25 °C. To obtain adsorption isotherms, solutions of various Pb(II) concentrations (10, 50, 100, 300, 400, 500, 600, 700, and 800 mg/L) were shaken with HMp or Mg0.1-HMp at pH 5.0 and temperatures of 25, 35, and 45 °C for 24 h. The effect of the contact time of Pb(II) adsorption on HMp or Mg0.1-HMp was evaluated in 50 mL of 10 and 20 mg/L Pb(II) solutions. The concentration of Pb(II) was analyzed at contact times of 0.25, 0.5, 1, 1.5, 3, 6, 10, 18, 24, and 36 h at 25, 35, and 45 °C, respectively.
The adsorption capacity (Qe, mg/g) (Equation (1)) and adsorption rate (R) (Equation (2)) were analyzed using the following formulas:
Q e = ( C 0 − C e ) V / m
R ( % ) = 100 × ( C 0 − C e ) / C 0
where m (g) is the weight of the adsorbent, V (L) is the solution volume, and C0 and Ce are the Pb(II) concentrations (mg/L) before and after the adsorption process, respectively.

2.5. Regeneration and Recycling of Adsorbents

The regeneration and recycling of adsorbents was studied as follows: 0.1 g of HMp or Mg0.1-HMp was taken, filtered, and dried in an oven at 105 °C for 24 h. It was then placed in a 100 mL polyethylene plastic centrifuge tube and 50 mL of HNO3 solution (0.3 mol/L) was added and shaken at 25 °C for 24 h at 200 rpm in a water bath oscillator. Then, the supernatant of the centrifuge tube was drained to regenerate HMp or Mg0.1-HMp for reuse for adsorption. The experimental process for cyclic adsorption was identical to that used for the batch adsorption.

3. Results and Discussion

3.1. Characterization of HMp and Mg0.1-HMp

3.1.1. FT-IR Analysis

Figure 1 shows the FT-IR spectra of HMp, Mg0.1-HMp, and MBC before and after Pb(II) adsorption. Absorption bands were present at wave number (ν) values of 3408–3468, 1619–1652, and 1382–1400 cm−1 in six spectra. The band at ν = 3408–3468 cm−1 indicates the stretching and bending vibration modes of -OH in the hydroxyl group of carboxylic compounds, phenols, or alcohols [35]. The bands at ν = 1619–1652 and 1382–1400 cm−1 correspond to free water absorption [36] and N-O stretching vibration, respectively, which might be attributed to the surface adsorption of NO3- [37]. Meanwhile, the band at ν = 2297 cm−1 in Mg0.1-HMp could be due to the -CH2 asymmetric stretching vibration [38]. Bands at ν = 661 and 563 cm−1 in HMp and Mg0.1-HMp correspond to the tetrahedral symmetric stretching vibration of PO43−. It is demonstrated that HMp and Mg0.1-HMp contain O-H and PO43−, which are characteristic of the HAp structure [39]. Furthermore, compared to MBC, some bands at ν = 1032 and 1037 cm−1 were observed in the FTIR spectra of Mg0.1-HMp and HMp, respectively, which suggests that Mg was successfully loaded on HMp [40].
The band at ν = 3447 cm−1 of the MBC shifted to ν = 3408 cm−1 after Pb(II) adsorption. The band at ν = 3468 cm−1 of the HMp shifted to ν = 3446 cm−1 after Pb(II) adsorption. The wave number of the contraction vibration of the -O-H absorption peak of the Mg0.1-HMp after Pb(II) increased from ν = 3439 to 3444 cm−1. In the FTIR spectra of Mg0.1-HMp after Pb(II) adsorption, the contraction vibration of the -C=O double bond shifted from ν = 1652 to 1633 cm−1, whereas in the FTIR spectra of MBC and HMp, it shifted from ν = 1619 and 1625 cm−1 to 1624 and ν = 1634 cm−1, respectively. This suggests that Pb(II) adsorption occurred on the surfaces of HMp, Mg0.1-HMp, and MBC and was facilitated by –OH and–C=O functional groups [41].

3.1.2. XRD Analysis

The XRD patterns of MBC, HMp, and Mg0.1-HMp are shown in Figure 2. The intense, sharp, and narrow peaks observed in the XRD patterns of HMp and Mg0.1-HMp confirmed their crystallinity [42]. Only one carbon peak was observed in the XRD pattern of MBC. The XRD pattern of Mg0.1-HMp was nearly identical to that of HMp, demonstrating that Mg doping had little impact on the crystal structure of HAp [43]. However, upon comparison with the HAp standard card, the degree of HMp matching was higher than that of Mg0.1-HMp, suggesting that the HMp crystallization process was disturbed following Mg doping along with the production of certain impurities [44].
Following Pb(II) adsorption, the crystal structures of MBC, HMp, and Mg0.1-HMp changed significantly (Figure S1). The characteristic peaks of Pb9C6O24 were observed in the XRD pattern of MBC-Pb. The characteristic peaks of HAp (Pb10(PO4)6(OH)2 were observed in the XRD patterns of both HMp-Pb and Mg0.1-HMp-Pb. The characteristic peaks of (Pb5(PO4)3(OH) and Pb3O2CO3 were observed in the XRD pattern of HMp-Pb. The compounds generated in Mg0.1-HMp-Pb were distinct from those generated in HMp-Pb, including Pb9(PO4)6, Pb24C8O40, and Pb7.38Ca2.62P6O26. The results showed that MBC, HMp, and Mg0.1-HMp mainly removed Pb(II) from aqueous solutions through chemical adsorption; however, Mg0.1-HMp was more effective in removing Pb(II) because the dissolved PO43− directly reacted with Pb2+ to produce Pb9(PO4)6 [45].

3.1.3. SEM-EDS Analysis

The results of SEM/EDS show that MBC, HMp, and Mg0.1-HMp have different pore structures (Figure 3). MBC has a deep pore structure with a reasonably smooth pore wall, and the original biomass structure is almost intact. HMp and Mg0.1-HMp were loaded with HAp and/or Mg oxide particles on the surfaces and pores. MBC has a better pore structure compared to HMp and Mg0.1-HMp, which may be beneficial for the adsorption of heavy metals [46].

Analysis of Specific Surface Area

The specific surface areas of HMp and Mg0.1-HMp determined using the nitrogen adsorption–desorption method (BET-N2) method were 233.78,257.844, and 316.324 m2/g, respectively. The increase in specific surface area indicates that the addition of Mg possibly caused lattice defects. The pore volumes of HMp and Mg0.1-HMp were 0.529 and 0.551 cm3/g, respectively, and the mean particle sizes were 9.729, 8.697, and 7.549 nm, respectively. When the specific surface area of the material increases, more adsorption sites become available, boosting the ability of the material to bind Pb(II). However, a reduction in particle size leads to a reduction in the pore size between particles [43]. From the area and pore size data, it can be seen that when Mg is doped in HMp, the specific surface area will increase to 316.324 m2/g, and the particle size will be reduced by 1.148 nm. It can be seen from Table 1 that HMp and Mg0.1-HMp can significantly improve the pore volume and pore size of MBC, so that its pore volume increases, providing more adsorption check points for the material. The corresponding pore volume increases by 0.319 and 0.341 cm3/g, and the modified pore size was also increased, indicating that there was no pore collapse in HMp and Mg0.1-HMp.

3.2. Effects of Reaction Conditions on Pb(II) Adsorption by HMp and Mg0.1-HMp

3.2.1. Effect of pH

The Pb(II) adsorption capacities of HMp and Mg0.1-HMp at initial Pb(II) concentrations of 10, 100, and 500 mg/L are shown in Figure 2. The Pb(II) adsorption capacities of HMp and Mg0.1-HMp exhibited an increasing trend as the pH increased from 2.0 to 50 mg/g. The Pb(II) adsorption capacities of HMp and Mg0.1-HMp decreased with the increasing pH in the range 5-7.05, and remained stable above pH 7.05. For the initial Pb(II) concentration (C0) of 500 mg/L, the Pb(II) adsorption capacities of HMp and Mg0.1-HMp increased with an increase in pH from 2.0 to 5.0 and decreased with the increasing pH from 5.0 to 8.0. The maximum Pb(II) adsorption capacities (Qm) of HMp and Mg0.1-HMp were 248.1 and 248.5 mg/g, respectively. The Qm of Mg0.1-HMp for Pb(II) is slightly higher than that of HMp, even though they were similar at C0 = 10 and 100 mg/L. The optimum pH for HMp and Mg0.1-HMp adsorption of Pb(II) is from 4.0 to 8.0.
The adsorption behavior of Pb(II) is related to its species in the solution, the surface charge, and dissociation of the functional groups of the biochar [47]. The points of zero charge (pHzpc) of HMp and Mg0.1-HMp were 6.48 and 6.89 (Figure S1). At pH < pHzpc, the HMp and Mg0.1-HMp have a positive surface charge. At pH < 7.0, the predominant Pb(II) species were Pb2+, PbOH+, and Pb2OH3+; an electrostatic discharge action occurs between adsorbent HMp and Mg0.1-HMp to and Pb(II), and the ability of HMp and Mg0.1-HMp to bind Pb(II) is reduced. At pH > 7.0, the precipitation of Pb(OH)2 occurred [42] and Pb(II) hydroxide complexes were formed; these may be the reason why the optimum pH for HMp and Mg0.1-HMp adsorption of Pb(II) is from 4.0 to 8.0.

3.2.2. Adsorption Isotherms

Figure 3 shows the Pb(II) adsorption isotherms for HMp, Mg0.1-HMp, and MBC at 25, 35, and 45 °C. The adsorption capacity increased with the increasing temperature at a specific initial Pb(II) concentration. The Pb(II) adsorption efficiencies of HMp, Mg0.1-HMp, and MBC increased from 305.85 to 318.45, 315.48 to 324.01, and 123.60.0 to 139.55 mg/g at 25, 35, and 45 °C, respectively. The Pb(II) adsorption efficiencies of HMp and Mg0.1-HMp were 2.32 and 2.28 times those of MBC at 45 °C. The maximum Pb(II) adsorption capacity of Mg0.1-HMp was much higher than that of corn straw biochar (74.01 mg/g [48] and composite materials derived from used chicken feathers and Agaricus bisporus mushrooms (70.42 mg/g [49]), but slightly lower than that of nano-HAp [50,51]. The relationship between the temperature and Pb(II) adsorption capacities of HMp and Mg0.1-HMp was positive. An increase in the temperature promoted Pb(II) adsorption by HMp and Mg0.1-HMp, indicating that the adsorption process was endothermic because an increase in the temperature accelerates the chemical adsorption between the adsorbent and adsorbate, which increases the level of adsorbent activation [52].
Isothermal experimental data were fitted using the Langmuir (Equation (3)) and Freundlich (Equation (4)) models to qualitatively determine the nature of solute–adsorbent surface interactions.
C e / Q e = 1 / ( K L Q m ) + C e / Q m
lg Q e = lg K F + 1 / ( n lg C e )
where Ce denotes the solution concentration at adsorption equilibrium (mg/L); Qm and Qe are the maximum and equilibrium adsorption concentrations (mg/g), respectively; KF and KL are the Freundlich and Langmuir constants (L/mg), respectively; and n is the adsorption intensity of the Freundlich isotherm model.
The results of the fitting parameters are displayed in Table 2, which show that the Freundlich equation provides relatively poor fits over the whole concentration range for Pb(II) adsorption onto the MBC, HMp, and Mg0.1-HMp with correlation coefficient (R2) values of 0.8782, 0.8827, and 0.8820, 0.8796, 0.8803, and 0.8785, and 0.9425, 0.9438, and 0.9448 at 25, 35, and 45 °C, respectively. In this study, the 1/n values of the Freundlich equation for Pb(II) adsorption onto the MBC, HMp, and Mg0.1-HMp were 0.2692, 0.2750, and 0.2758, 0.3126, 0.3149, and 0.3143, and 0.4708, 0.4781 and 0.4876, at 25, 35, and 45 °C, respectively, indicating chemical adsorption [53].
Contrastingly, the Langmuir isotherm provides good fits over the entire concentration range for Pb(II) adsorption onto the MBC, HMp, and Mg0.1-HMp, with R2 values of 0.9986, 0.9984, and 0.9987, 0.9999, 0.9999, and 0.9999, 0.9996, 0.9995, and 0.9967 at 25, 35, and 45 °C, respectively. The R2 value of the Langmuir isotherm was superior to that of Freundlich at the same temperature. The fitting of experimental data into the Langmuir isotherm model with high R2 values indicates the homogeneous nature of the adsorbent surface [54]. The maximum Pb(II) adsorption capacities of MBC, HMp, and Mg0.1-HMp were 123.60, 134.05, and 139.55, 305.85, 312.40, and 318.45, and 315.48, 320.76, and 324.01 mg/g at 25, 35, and 45 °C, respectively. The order of the maximum Pb(II) adsorption capacity of the adsorbents was Mg0.1-HMp > HMp > MBC. The introduction of Mg2+ breaks the hexagonal crystal structure of HAp, thereby increasing the specific surface area, number, and utilization efficiency of -OH and PO43−, and facilitating Pb(II) contact with Mg0.1-HMp, resulting in the higher Pb(II) adsorption and immobilization efficiency of Mg0.1-HMp compared to those of MBC and HMp.

3.2.3. Adsorption Kinetics

Contact time plays an important role in modeling and designing the adsorption process. During adsorption, the quantity of removal of metal ions increased with time. A 45 h contact time was used to determine its effect on the adsorption of Pb(II) by Mg0.1-HMp. The Pb(II) adsorption capacities of HMp and Mg0.1-HMp steadily increased with the contact time (Figure 4). The adsorption reaction rate was significantly high during the first 3 h, and the Pb(II) adsorption capacities of HMp were 4.89, 49.8, and 244.1 mg/g, respectively, at the end of each hour. The Pb(II) adsorption capacities of Mg0.1-HMp were 4.91, 49.84, and 244.6 mg/g, respectively, at the end of each hour owing to the presence of numerous adsorption sites on its surface at the beginning of the adsorption process. Over time, the adsorption curve becomes smoother as the pace of the adsorption reaction slows and stabilizes as fewer adsorption sites on their surfaces are available. The adsorption equilibrium was reached at 24 h, with a decrease in the adsorption rate of Pb(II).
Experimental adsorption data were fitted linearly using quasi-primary (Equation (5)), quasi-secondary (Equation (6)), Bangham (Equation (7)), and Elovich (Equation (8)) kinetic models.
log ( Q e − Q t ) = log Q e − K 1 t / 2.303
t / Q t = K 2 − 1 Q e − 2 + t / Q e
d Q t / d t = K 3 ( Q e − Q t ) Q e / Q t
Q t = K 4 t 1 / 2 + C
In these equations, t denotes the reaction time, min; Qe and Qt are the adsorption quantities (mg/g) at equilibrium and at time t, respectively; K1, K2, K3, and K4 are rate constants of the quasi-primary, quasi-secondary, Bangham, and Elovich kinetic models (g (mg min)−1), respectively; and C is the thickness of the boundary layer constant.
Pseudo-first-order, pseudo-second-order, Bangham, and Elovich kinetic models were used to match the experimental data obtained from the adsorption process. The closer the R2 is to one, the better the fit of the experimental data to the isotherm. The closer it is to zero, the more unsuitable the isotherm is for these data. Table 3 lists the fitting parameters for the HMp adsorption kinetics. In the pseudo-first-order kinetic adsorption fitting curve, the R2 values for Pb(II) concentrations of 10, 100, and 500 mg/L were 0.88, 0.96, and 0.93, respectively, demonstrating that this equation does not accurately describe the adsorption reaction process. The quasi-second-order kinetic equation described the adsorption process better than the quasi-first-order kinetic fitting equation. The experimentally determined adsorption quantities of 4.89, 49.8, and 244.1 mg/g were close to the estimated theoretical adsorption quantities of 5.00, 50.00, and 250.00 mg/g. The R2 values of Bangham kinetics fitting were 0.88, 0.87, and 0.97. The R2 values of Elovich kinetics fitting were 0.65, 0.64, and 0.87, respectively.
Table 4 lists the fitting parameters for the Mg0.1-HMp adsorption kinetics. According to the pseudo-first-order kinetic adsorption fitting curve, the R2 values of the data for concentrations of Pb solutions of 10, 100, and 500 mg/L were 0.88, 0.97, and 0.90, respectively. This poor correlation suggests that the pseudo-first-order kinetic equation cannot adequately describe the adsorption process. The quasi-second-order kinetic equation describes the adsorption process better than the quasi-first-order kinetic fitting. The linear relationships were strong, with an estimated R2 of 1.000 in all cases. The experimentally determined adsorption quantities of 4.91, 49.84, and 244.6 mg/g are similar to the predicted theoretical adsorption quantities of 5.00, 50.00, and 250.00 mg/g, indicating that the second-order kinetic equation can more accurately explain the adsorption process. The R2 values of Bangham kinetics were 0.88, 0.87, and 0.97. The R2 values of Elovich kinetics fitting were 0.65, 0.64, and 0.87, respectively. Therefore, the quasi-second-order kinetic equation for Pb(II) best describes the adsorption process.

3.2.4. Possible Mechanisms Pb(II) Adsorption onto HMp and Mg0.1-HMp

XPS is a useful tool for studying the interactions between adsorbents and adsorbates. XPS analysis was used to understand the surface chemical characteristics of MBC, HMp, and Mg0.1-HMp before and after adsorption to evaluate the adsorption mechanism. High-resolution energy spectra of C1s prior to the adsorption of the three compounds are shown in Figure 5a. The two peaks in MBC, HMp, and Mg0.1-HMp represent C-C and C-O. Figure 5b displays the energy spectrum of C1s following adsorption. MBC has only peaks of C-C and C-O, whereas HMp and Mg0.1-HMp both have three individual component peaks, which represent C-C, C-O, and C=O. The C-C (284.8 eV) peak in MBC was more intense, and the C-O (286.55 eV) changes by +1.75 eV from the C1s energy spectrum before and after Pb(II) adsorption. These observations suggest that both C-C and C-O are involved in the adsorption of Pb(II) by MBC. The binding energies of C-C (283.42 eV) and C-O (284.8 eV) in HMp considerably dropped, and a new C=O (287.84 eV) emerged. Hang et al. [55] suggested that C-C, C-O, and C=O bonds were involved in the adsorption of Pb(II). For Mg0.1-HMp, a new C=O peak (287.82 eV) is observed, although the C-C (283.56 eV) and C-O (284.8 eV) binding energies decrease marginally. This decline was less than that observed for HMp. It is hypothesized that Mg participates in the adsorption process, in addition to C-C, C-O, and C=O.
Similarly, the high-resolution energy spectra of O1s before and after the adsorption of the three compounds are shown in Figure 5c. The two peaks in MBC, HMp, and Mg0.1-HMp represent C-O and O-H. Figure 6d displays the O1s energy spectrum following adsorption, which still contains C-O. In contrast, OH and Pb(II) complexes produce functional groups such as PbOH+ and Pb2OH3+, indicating the significance of surface hydroxyl groups in Pb(II) adsorption [56]. After adsorption, significant changes in the binding energies and peak areas of C-O (531.83 eV) and O-H (533.78 eV) in MBC were not observed compared with those in HMp and Mg0.1-HMp. However, the analysis indicated that the M-OH peak area expanded the most in HMp, leading to peak regions that overlapped with C-O.
This indicates that the apparent hydroxyl groups of HMp are also important. Compared to MBC and HMp, the binding energy of C-O in Mg0.1-HMp changes by −0.25 eV, a significant quantity. The peak area of O-H decreased by 38.65%, which is much less than that of MBC, but not as much as that of HMp, suggesting that doping with Mg ions increases the activity of hydroxyl groups [57]. Therefore, the high-resolution energy spectra of HMp and Mg0.1-HMp were further examined.
Figure 6e shows the high-resolution XPS spectra of Ca2p before and after HMp and Mg0.1-HMp adsorption. Four typical peaks occur at binding energies of 347.9, 351.52, 346.71, and 350.18 eV prior to adsorption, which correspond to the Ca of Ca10(PO4)6(OH)2 in two distinct environments. The first is when Ca joins O at the nine vertices of the tetrahedron, generating the coordinating anion Ca-P6O24, with a coordination number of nine. The entire crystal structure has a sizable channel parallel to the c-axis through this link. A different type of reaction occurs when the channel ion OH combines with the surrounding upper and lower layers of the six Ca atoms to form Ca-OH. In this case, the coordination number between Ca at the top of the cation’s octahedral corner and O and OH at the top of the nearby 4 [PO4] hexagonal corners is seven [58]. The Ca2P3/2 and Ca2P1/2 binding energies of HMp and Mg0.1-HMp moved to +0.34 and +0.19 and −0.23 and −0.08 eV, respectively, after adsorption. This suggests that the synthesis of materials contributes to the adsorption process. Prior to adsorption, the P in the Ca-O-P bond of the corresponding interface P is responsible for the binding energy peaks of HMp and Mg0.1-HMp at 133.66 and 132.14 eV, respectively, whereas the P in PO43− is responsible for the binding energy peaks at 134.55 and 133.71 eV, respectively (Figure 6f). The peak area of PO43− in the HMp sample increased by 21.39% after adsorbing Pb(II), and the binding energy shifted by +0.93 eV, overlapping with the Ca-O-P bond. The PO43− peak intensity of Mg0.1-HMp increased, and the binding energy shifted by +1.57 eV, whereas the Ca-O-P bond’s binding energy shifted by +0.81 eV, indicating that Pb2+ formed PbHPO4, Pb3(PO4)2, and PbHAp with PO43− during this adsorption process [59].
Figure 6g shows the high-resolution XPS spectra of Pb4f following adsorption by MBC, HMp, and Mg0.1-HMp. According to Li et al. [60], MBC has two symmetrical peaks that were centered at 143.78 and 138.85 eV, suggesting no change in the valence state or synthesis of new compounds during the adsorption process. Compared with MBC, HMp and Mg0.1-HMp contained two more peaks. The peaks in the HMp sample with binding energies of 143.95 and 139.55 eV are considered to represent Pb3O2CO3, HAp Pb10(PO4)6(OH)2, and Pb5(PO4)3(OH)2. These findings are supported by the XRD analysis. This suggests that dissolution precipitation is the primary purpose of this adsorption process. The crystals of Ca10−xPbx(PO4)6(OH)2 generated during the procedure are unstable due to the high Ca concentration as the direct exchange of Ca ions is difficult. PbHAp becomes stable when the Ca in the crystal dissolves and is replaced by Pb49. Similarly, Pb10(PO4)6(OH)2 had a binding energy of 144.07 eV in Mg0.1-HMp. A shift and drop in the peak area were observed compared to HMp. Pb9(PO4)6 is represented by a peak with a binding energy of 139.21 eV, suggesting that Pb(II) is adsorbed by Mg0.1-HMp via dissolution and precipitation and that Mg-ion doping causes crystal defects in HAp. These defects increase the specific surface area and improve the efficiency with which the surface functional groups OH and PO43− are used, enabling these groups to interact with Pb2+ more readily. This may be the primary cause for the superior adsorption efficacy of Mg0.1-HMp over HMp [61].
Consequently, Mg0.1-HMp plays a role in the adsorption of Pb(II) via ion exchange, surface complexation, dissolution–precipitation or displacement reactions, and surface and pore adsorption [31]. Van der Waals forces and electrostatic adsorption frequently occur concurrently with ion-exchange activities on the surface of Mg0.1-HMp when Pb(II) approaches it. Ca2+ in Mg0.1-HMp also undergoes ion exchange or dissolution–precipitation when Mg0.1-HMp adsorbed Pb(II) in an aqueous solution by electrostatic contact, fixing the deposited Pb(II). Mg0.1-HMp has better surface characteristics, a larger specific surface area, smaller particles, and a higher degree of Mg solubility than those of HMp. Additionally, it increased the number of surface functional groups such as PO43−, -OH, and -C=O and consumption efficiency. This facilitated Pb2+ production by Pb9(PO4)6 or Pb10(PO4)6(OH)2, which may be the primary cause of the noticeably higher adsorption capacity of Mg0.1-HMp compared to that of HMp. Furthermore, complexation processes involving PO43−, O-H, and Pb2+ would result in the formation of functional groups such as PbOH+ and Pb2OH3+, Pb(OH)2 precipitates, and hydroxide complexes with Pb2+. PbHPO4, Pb3(PO4)2, and PbHAp are formed via complexation reactions between Pb2+ and PO43−. Based on these results, Figure 7 depicts the potential processes for Pb(II) adsorption and immobilization by Mg0.1-HMp. The possible adsorption mechanisms are illustrated in Figure 8.

4. Conclusions

Currently, there is a great need to develop fit-for-purpose metal-containing wastewater treatment methods using low-cost materials. In this study, we highlight that mulberry stems are an abundant biomass that can serve as a potential low-cost source. The adsorption and fixation efficiency of Pb(II) by HMp and Mg0.1-HMp were slightly inhibited by ionic strength. The higher temperature is conducive to the adsorption of Pb(II) on HMp and Mg0.1-HMp, which is a monolayer adsorption of Pb(II) by HMp and Mg0.1-HMp. The adsorption of Pb(II) by HMp and Mg0.1-HMp is a spontaneous endothermic physicochemical process. The main mechanisms of Pb(II) adsorption on HMp and Mg0.1-HMp involve electrostatic interactions, surface adsorption, and ion exchange. In addition, 0.1 mol/L of HNO3 had a strong desorption effect on the adsorption of Pb(II) by Mg0.1-HMp. After four cycles, Mg0.1-HMp has good regeneration properties and is a promising composite material that can be used as a fixative for the extraction of Pb(II) from contaminated water, and these results are essential for the control of toxic and hazardous metal-contaminated water bodies and the reduction in hazards associated with them. In addition, this process improves the safe recovery rate of agricultural biowaste and reaffirms the applicability of functionalized biochar for the removal of harmful heavy metals from the water environment.

Supplementary Materials

The following supporting information can be downloaded at: https://www.mdpi.com/article/10.3390/w17091389/s1, Figure S1: Zeta potentials of HMp, Mg0.1-HMp, and MBC, Figure S2: Freundlich equation of HMp, Mg0.1-HMp, and MBC, Figure S3: Langmuir equation of HMp, Mg0.1-HMp, and MBC, Figure S4: Pseudo-first-order-kinetics of HMp and Mg0.1-HMp, Figure S5: Pseudo- second-order kinetics of HMp and Mg0.1-HMp, Figure S6: Bangham kinetics model of HMp and Mg0.1-HMp, Figure S7: Elovich kinetics model of HMp and Mg0.1-HMp.

Author Contributions

Conceptualization, X.Z.; methodology, D.W.; formal analysis, X.Z. and; investigation, X.Z., Z.W., and M.L.; resources, Z.W.; data curation, X.Z.; writing—original draft, X.Z. and M.L.; writing—review and editing, M.L. and X.Z.; visualization, X.Z.; supervision, Z.W.; project administration, D.W.; funding acquisition, M.L. All authors have read and agreed to the published version of the manuscript.

Funding

This research was funded by the National Natural Science Foundation of China (52260023) and the Guangxi Natural Science Foundation (2025GXNSFDA069043).

Data Availability Statement

The data that support the findings of this study are available from the corresponding author upon reasonable request.

Conflicts of Interest

Author Zimeng Wu was employed by the company Yunnan Xinlian Environmental Technology Co, Ltd. The remaining authors declare that the research was conducted in the absence of any commercial or financial relationships that could be construed as a potential conflict of interest.

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Water 17 01389 g001
Figure 1. FTIR spectra of before and after Pb(II) adsorption onto HMp, Mg0.1-HMp, and MBC.
Figure 1. FTIR spectra of before and after Pb(II) adsorption onto HMp, Mg0.1-HMp, and MBC.
Water 17 01389 g001
Water 17 01389 g002
Figure 2. XRD spectra of before and after Pb(II) adsorption onto HMp, Mg0.1-HMp, and MBC.
Figure 2. XRD spectra of before and after Pb(II) adsorption onto HMp, Mg0.1-HMp, and MBC.
Water 17 01389 g002
Water 17 01389 g003
Figure 3. SEM spectra of MBC, HMp, and Mg0.1-HMp.
Figure 3. SEM spectra of MBC, HMp, and Mg0.1-HMp.
Water 17 01389 g003
Water 17 01389 g004
Figure 4. Effect of pH on Pb(II) adsorption by HMp and Mg0.1-HMp.
Figure 4. Effect of pH on Pb(II) adsorption by HMp and Mg0.1-HMp.
Water 17 01389 g004
Water 17 01389 g005
Figure 5. Adsorption isotherms of Pb(II) adsorption onto (a) HMp, (b) Mg0.1-HMp, and (c) MBC.
Figure 5. Adsorption isotherms of Pb(II) adsorption onto (a) HMp, (b) Mg0.1-HMp, and (c) MBC.
Water 17 01389 g005
Water 17 01389 g006aWater 17 01389 g006b
Figure 6. High-resolution energy spectra of MBC, HMp, and Mg0.1-HMp before and after Pb(II) adsorption. C1s (a) before and (b) after Pb(II) adsorption. O1s (c) before and (d) after Pb(II) adsorption. (e) Ca2p before and after Pb(II) adsorption. (f): P2p before and after Pb(II) adsorption. (g) Pb4f after Pb(II) adsorption.
Figure 6. High-resolution energy spectra of MBC, HMp, and Mg0.1-HMp before and after Pb(II) adsorption. C1s (a) before and (b) after Pb(II) adsorption. O1s (c) before and (d) after Pb(II) adsorption. (e) Ca2p before and after Pb(II) adsorption. (f): P2p before and after Pb(II) adsorption. (g) Pb4f after Pb(II) adsorption.
Water 17 01389 g006aWater 17 01389 g006b
Water 17 01389 g007
Figure 7. Effect of contact time on Pb(II) adsorption onto (a) HMp and (b) Mg0.1-HMp.
Figure 7. Effect of contact time on Pb(II) adsorption onto (a) HMp and (b) Mg0.1-HMp.
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Water 17 01389 g008
Figure 8. Possible mechanisms of Pb(II) adsorption and immobilization onto HMp and Mg0.1-HMp.
Figure 8. Possible mechanisms of Pb(II) adsorption and immobilization onto HMp and Mg0.1-HMp.
Water 17 01389 g008
Table 1. The analysis of surface area and pore size of HMp, Mg0.1-HMp, and MBC.
Table 1. The analysis of surface area and pore size of HMp, Mg0.1-HMp, and MBC.
SampleSBET/(m2/g)Vtotal/(cm3/g)DBET/nm
MBC233.780.2109.729
HMp257.880.5298.697
Mg0.1-HMp316.3240.5517.549
Table 2. Isotherm parameters of Pb(II) adsorption onto MBC, HMp, and Mg0.1-HMp.
Table 2. Isotherm parameters of Pb(II) adsorption onto MBC, HMp, and Mg0.1-HMp.
SampleTemperature (°C)Langmuir EquationFreundlich Equation
Qm(mg/g)KL(L/mg)R2KF(L/mg)1/nR2
MBC25123.460.13210.998626.9290.26920.8782
35135.140.12760.998428.5800.27500.8827
45138.890.14780.998730.0090.27580.8820
HMp25303.031.37500.9999104.9100.31260.8796
35312.501.33330.9999108.5810.31490.8803
45322.581.40910.9999113.4660.31430.8785
Mg0.1-HMp25312.500.94120.999634.5220.47080.9425
35322.580.91180.999537.6040.47810.9438
45322.580.79490.996741.0220.48760.9448
Table 3. Kinetic parameters of Pb(II) adsorption onto HMp.
Table 3. Kinetic parameters of Pb(II) adsorption onto HMp.
Initial Pb Concentration (mg/L)Pseudo-First-Order-KineticsPseudo-Second-Order Kinetics
R2 K1 (1/min)Qmax (mg/g)R2K2 (g∙mg/min)Qmax (mg/g)
100.87620.00214.99451.00000.06885
1000.96930.002249.99551.00000.080050
5000.93620.0023248.4831.00000.0013250
Initial  Pb Concentration (mg/L)Bangham kinetics modelElovich kinetics model
R2K3  (g∙mg/min)R2K4  (g∙mg/min)
100.88860.090.65680.0086
1000.87960.050.64200.0048
5000.97261.500.87060.1538
Table 4. Kinetic parameters of Pb(II) adsorption onto Mg0.1-HMp.
Table 4. Kinetic parameters of Pb(II) adsorption onto Mg0.1-HMp.
Initial Pb Concentration (mg/L)Pseudo-First-Order KineticsPseudo-Second-Order Kinetics
R2K1 (1/min)Qmax (mg/g)R2K2 (g∙mg/min)Qmax (mg/g)
100.88840.00154.9961.00000.08335
1000.97980.001849.9971.00000.043450
5000.90920.0031248.621.00000.0011250
Initial  Pb Concentration (mg/L)Bangham kinetics modelElovich kinetics model
R2K3  (g∙mg/min)R2K4  (g∙mg/min)
100.90460.0100.62300.0060
1000.97840.0010.88730.0052
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Wang, D.; Zhou, X.; Liang, M.; Wu, Z. Synthesis of Hydroxyapatite Mulberry Stem Biochar Composites for Efficient Pb(II) Adsorption from Aqueous Solutions. Water 2025, 17, 1389. https://doi.org/10.3390/w17091389

AMA Style

Wang D, Zhou X, Liang M, Wu Z. Synthesis of Hydroxyapatite Mulberry Stem Biochar Composites for Efficient Pb(II) Adsorption from Aqueous Solutions. Water. 2025; 17(9):1389. https://doi.org/10.3390/w17091389

Chicago/Turabian Style

Wang, Dunqiu, Xinyu Zhou, Meina Liang, and Zimeng Wu. 2025. "Synthesis of Hydroxyapatite Mulberry Stem Biochar Composites for Efficient Pb(II) Adsorption from Aqueous Solutions" Water 17, no. 9: 1389. https://doi.org/10.3390/w17091389

APA Style

Wang, D., Zhou, X., Liang, M., & Wu, Z. (2025). Synthesis of Hydroxyapatite Mulberry Stem Biochar Composites for Efficient Pb(II) Adsorption from Aqueous Solutions. Water, 17(9), 1389. https://doi.org/10.3390/w17091389

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