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Vol. 8. Num. 1.
Pages 1-1592 (January - March 2019)
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Vol. 8. Num. 1.
Pages 1-1592 (January - March 2019)
Original Article
DOI: 10.1016/j.jmrt.2018.05.020
Open Access
Nickel sorption using Bioclastic Granules as a sorbent material: equilibrium, kinetic and characterization studies
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Diego Macedo Veneua,
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diegomveneu@yahoo.com.br

Corresponding author.
, Lídia Yokoyamaa, Osvaldo Galvão Caldas Cunhaa, Claudio Luiz Schneiderb, Marisa Bezerra de Mello Monteb
a Escola de Química, Universidade Federal do Rio de Janeiro, Centro de Tecnologia, Ilha do Fundão, Rio de Janeiro, Brazil
b Laboratório de Química de Superfície, Centro de Tecnologia Mineral, Rio de Janeiro, Brazil
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Tables (6)
Table 1. Effect of the initial concentration of Ni(II) on sorption by bioclastic granules.
Table 2. Isotherm model parameters and their correlation coefficients.
Table 3. RL values for Ni(II) sorption by bioclastic granules.
Table 4. Kinetic parameters obtained from pseudo-first and pseudo-second order kinetic models for BG Ni(II) ions sorption.
Table 5. Fluorescence X-ray analysis of bioclastic granules.
Table 6. Structural properties of bioclastic granules.
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Abstract

The present study evaluated the potential of bioclastic granules (BG) from Lithothamnium calcareum algae as a sorbent material for the removal of Ni(II) from aqueous solutions, through experimental batch tests. Relevant sorption process variables were evaluated, such as pH (2–7), particle size (<38–300μm), initial sorbent concentration (0.5–3.0gL−1), initial metal ion concentration (5–500mgL−1) and contact time (5–240min). The data set regarding the uptake of metal ions at equilibrium was well-fitted to the Langmuir isotherm model (R2 0.957), with a qmax of 54.9mgg−1 and kL 0.014Lmg−1. The kinetic data were better fitted to the pseudo-second order kinetic model (R2 of 0.994), with qeq of 21.74mgg−1 and k2 of 0.028gmg−1min−1. Characterization of BG before and after the sorption tests in the presence of Ni(II) ions was conducted by X-ray fluorescence analysis (XRF), BET surface area, X-ray diffraction (XRD), zeta potential and electron microscopy scanning (SEM-EDS).

Keywords:
Wastewater treatment
Sorption
Nickel
Calcareous algae
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1Introduction

According to Duffus [1], in recent decades the term “heavy metal” has often been used to refer to a group of metals and semimetals (metalloids) associated with contamination and potential toxicity or ecotoxicity. Heavy metal contamination is still an environmental problem in developed and developing countries worldwide [2,3]. With rapid industrialization and economic development, heavy metals continue to be introduced into the environment through point and diffuse sources [4].

While developed countries adopt strict water quality requirements to control water pollution from point and diffuse sources, the situation is different in most developing countries. Wastewater treatment is not prioritized, and, therefore, sanitary and industrial effluents are discharged into water bodies after poor or no treatment at all. According to Rostamian et al. [5], effluents containing low or moderate levels of heavy metals are often found in plating and metal mining industries, in the chemical pharmaceutical industry, in pigments, foundries, battery manufacturing, fertilizers, and many others.

Nickel is a silvery, fairly hard, ductile and malleable metal. It forms soluble inorganic compounds such as hydroxides, sulfates, chlorides and nitrates, as well as insoluble compounds, such as oxides and sulfides. The most common activities that lead to occupational exposure to nickel are mining, milling and smelting of ores obtained from sulfides and oxides, and the use of primary nickel products in both stainless steel and alloys, as is common in foundries. According to Lakshtanov and Stipp [6] nickel contamination from industrial and mining activities may pose a serious risk to groundwater, since this metal is toxic to plants and animals. The International Agency for Research on Cancer (IARC) classifies metallic nickel and alloys (Group 2B) and nickel compounds (Group 1) as carcinogenic to humans [7].

Several treatment technologies have been employed for removing metals from contaminated waters. Conventional techniques usually involve the use of physical or chemical processes, such as ion exchanges, chemical precipitation, oxidation/reduction, adsorption, solvent extraction, biosorption, use of separation membranes (ultrafiltration, nanofiltration and reverse osmosis) and electrochemical methods (electrodialysis) [5,8–12].

It is known that sorption plays an important role in the transport of control and target metal contaminants in the ecosystem. Sorption is, in a broad sense, the mass exchange of a chemical in the solid phase as a result of a mass transfer between fluids and solids, including adsorption, absorption or diffusion in the solid and in some cases, it can advance to surface precipitation to form a cohesive phase which may comprise chemical species derived from both states, i.e. the aqueous solution and the dissolved solid [13,14]. The surface precipitation itself is often induced by surface adsorption (the surface acting as a “template”) and thus called “surface enhanced precipitation” [15].

Various inorganic and organic natural materials, such as activated carbon, resins, clays, silica gel, organic materials and minerals, have been studied regarding sorption of heavy metal ions [16,17]. Calcareous seaweed is a generic term for several species of calcified red algae, in which a photosynthetically active, reddish pink, pigment is present in the living organism [18]. From the deposit of inorganic salts, seaweed die and lose their characteristic red color of the live surface film of the algae crust, developing a rich skeleton composed of calcium carbonate and precipitated magnesium in their cell walls, in the form of calcite crystals and/or aragonite, giving the dead seaweed a white-yellowish coloring [19].

The limestone produced by the extraction of Lithothamnium calcareum algae is often referred to as maerl, biogenic limestone or biodentritic marine and/or bioclastic granules (BG). Dias [19] reports that Brazil has an extensive occurrence of calcareous red algae deposits on the continental shelf in the North-Northeast and Southeast regions, with one of the largest reserves of this renewable mineral source. Estimates of deposits formed by algae in the Brazilian continental shelf have been quoted as up to 2.1011tons of sediment, of which over 75% is CaCO3[20].

The present study evaluated the capacity and efficiency of bioclastic granules (BG) from L. calcareum algae limestone to act as a sorbent material for the treatment of aqueous solutions containing Ni(II) ions.

2Material and methods2.1BG and reagents

The L. calcareum (BG) were provided by a company which conducts point extraction in the field located in the state of Espírito Santo, at about 24km from the Itapemirim coastline, at a depth of 13–20m. The extracted material is composed of the mineralized sediment of dead algae. The beneficiation process consists of ore drying, grinding, classification (cyclones), bagging and storage steps. The metal solutions used in the tests were prepared with deionized water from a NiSO4·6H2O salt (99%) supplied by Sigma–Aldrich.

2.2Sorption batch experiments

Factors affecting BG sorption rate and uptake capacity were studied on a bench scale. All assays were performed in 500mL Erlenmeyer flasks, employing 100mL of metal solution, at a rotating speed of 250rpm on a horizontal rotating platform (Cientec CT-712). To elucidate the optimum conditions for the sorption process, the following variables were chosen: (i) pH (2–7); (ii) particle size (<38–300μM); (iii) sorbent concentration (0.5–3.0gL−1); (iv) initial metal concentration (5–500mgL−1); and (v) balancing or contact time (5–240min).

After each assay, BG particles containing Ni(II) ions were concentrated by filtration on a 0.22μm membrane and subsequently removed so that the permeate samples (remaining metal solution) could be collected and acidified with HCl (0.1molL−1) for preservation and subsequent analysis of the residual concentration of the metal ion by elemental ICP-OES analysis on a Perkin Elmer Optima 4300DV equipment. Removal and sorption capacity were calculated through Eqs. (1) and (2), respectively. All assays were performed in triplicate.

Where: R is the removal of Ni(II) ions (%); q is the uptake capacity of Ni(II) ions (mgg−1); Ci is the initial Ni(II) concentration (mgL−1); Ceq is the Ni(II) concentration at equilibrium (mgL−1); V is the volume of the solution containing the Ni(II) ions (L); M is the mass of the sorbent (g).

2.3Sorption isotherm

Due to the diversity of mechanisms that occur, it would be more appropriate to use the term “sorption isotherm” instead of “adsorption isotherm” [21]. According to Porras et al. [22], two methods can be used to study the behavior of sorption isotherms: (i) identification by the type of curvature formed by the isotherm, and (ii) application of formulated mathematical models to describe adsorption behavior. By applying the first method, four basic types of adsorption isotherms have been recognized and used to identify the nature of the solute adsorption from aqueous solutions, namely types C, L, M and S isotherms [23]. The second method generally applies mathematical models proposed in the literature to explain the adsorption equilibrium.

2.3.1Langmuir isotherm

Langmuir's isotherm was initially described for gas adsorption studies on flat surfaces [24]. This model contains a number of assumptions: (i) all binding sites have an equal affinity to the solute, (ii) the sorption is limited to the formation of a monolayer, and (iii) the number of sorbed species does not exceed the total number of surface sites, i.e. there is a 1:1 stoichiometry between the adsorption sites on the surface and the solute [25]. Basically, the Langmuir isotherm equation has a hyperbolic form, given by Eq. (3):

Where: q is the amount of Ni(II) retained in the solid at equilibrium (mgg−1); qmax is the Langmuir parameter relative to sorption capacity (mgg−1); kL is the Langmuir constant relative to sorption energy (Lmg−1); Ceq is the Ni(II) concentration in solution at equilibrium (mgL−1).

According to Foo and Hameed [26], an essential feature of the Langmuir isotherm is expressed in terms of a dimensionless constant, known as the separation factor “RL” expressed by Eq. (4):

Where: RL is the separation factor; kL is the Langmuir constant relative to sorption energy (Lmg−1); and Ci is the initial concentration of Ni(II) in solution (mgL−1).

This factor indicates the nature of the sorption process; for RL>1 (unfavorable) to RL=1 (linear), for 0<RL<1 (favorable) and RL=0 (irreversible).

2.3.2Freundlich isotherm

The Freundlich isotherm model is considered appropriate to describe both multilayer sorption and sorption on heterogeneous surfaces [27]. This model (Eq. (5)) does not predict surface saturation based on the adsorption process, corresponding to an exponential distribution of several adsorption sites with different energies, and can, therefore, be applied to non-ideal systems [28].

Where: q is the amount of Ni(II) retained in the solid at equilibrium (mgg−1); Ceq is the Ni(II) concentration in solution at equilibrium (mgL−1); kF is a constant that indicates sorption capacity (Lmg−1) and n is the constant that indicates sorption strength (Lmg−1).

2.3.3Dubinin–Radushkevich isotherm

The Dubinin–Radushkevich (D-R) model assesses the nature of the adsorption and is more general than the Langmuir isotherm, since it assumes a smooth surface or a constant adsorption potential [29]. Dubinin, in 1960, and Radushkevich, in 1949, reported that the characteristic adsorption curve is related to the porous structure of the sorbent. Eq. (6) represents the Dubinin–Radushkevich isotherm and Eq. (7) represents the Polanyi potential.

Where: q is the amount of Ni(II) retained in the solid at equilibrium (mgg−1); qmax is the maximum sorption capacity of the sorbent (mgg−1); B is the constant related to the adsorption energy (mol2kJ−2); ɛ is the Polanyi potential (kJmol−1); R is the gas constant (0.008314kJmol−1K−1); T is the absolute temperature (K) and Ceq is the Ni(II) concentration in solution at equilibrium (mgL−1).

The sorption energy value (kJmol−1) are related to the constant B, as shown in Eq. (8). It is defined as the change in free energy when one mol of Ni(II) is transferred from infinity, inside the solution, to the solid surface, and is related to the sorption phenomenon that happens in the sorbent/solute system [30]. If Es<8kJ mol−1, the adsorption process is of a physical nature; If 8kJmol−1<Es<16kJmol−1, adsorption occurs by ion exchange; if 16kJmol−1<Es, adsorption is of a chemical nature.

Where: Es is the mean sorption free energy (kJmol−1).

2.4Sorption kinetics

Sorption kinetics is expressed as the solute removal rate that controls solute residence time in the solid–liquid interface. The sorption rate for a given system is among the most important factors in the design of a sorption system, because the system kinetics determines the solute residence time and reactor dimensions [31]. Various kinetic models have been established to understand sorption kinetics and determining the rate of the limiting step. Pseudo-first order and pseudo-second order kinetic models are the most used to study metal sorption kinetics to solids.

2.4.1Pseudo-first order model

The Lagergren model was the first to be developed for the study of sorption processes of solid–liquid systems [32]. This is the most widely used model to determine the solute sorption rate in a liquid solution, and can be represented by Eq. (9).

Eq. (10) is obtained by integrating Eq. (9) from t=0 to t=t and qt=0 to qt=qt:
Where: qeq is the amount of Ni(II) retained in the solid at equilibrium (mgg−1); q is the amount of Ni(II) retained in the solid at time “t” (mgg−1) and k1 is the pseudo-first order reaction velocity constant (gmg−1min−1).

2.4.2Pseudo-second order model

The pseudo-second-order model is also based on the sorption capacity of the sorbent [31]. However, unlike the Lagergren model, this model predicts the kinetic behavior over the entire sorption time range [33]. The pseudo-second order model is represented by Eq. (11):

Eq. (12) is obtained by integrating Eq. (11) from t=0 to t=t and qt=0 to qt=qt:
Where: qeq is the amount of Ni(II) retained in the solid in equilibrium (mgg−1); q is the amount of Ni(II) retained in the solid at time “t” (mgg−1) and k2 is the pseudo-second order reaction velocity constant (gmg−1min−1).

2.5Characterization2.5.1X-ray fluorescence

The elemental chemical quantitative determination of the BG material was performed by X-ray fluorescence analysis (XRF) on a PanAnalytical Axios Max model in the standardless method (semiquantitative analysis). The samples were prepared by melting at a 1:10 ratio using a borate mixture (Li2B4O7–LiBO2) from Maxxifluxi as flux.

2.5.2Surface area measurements of specific and porosity

Adsorption–desorption isotherms in N2 were performed on a Tristar II equipment (Surface Area and Porosity – Micromeritics) at liquid nitrogen temperature of 77K. The samples were degassed at 180°C 24h before the measurements. The surface area was obtained by the Brunauer–Emmett–Teller (BET) method and the pore size distribution was calculated from the adsorption isotherm derivation of the Barrett–Joyner–Halenda (BJH) model.

2.5.3X-ray diffraction

X-ray diffraction was conducted by the analysis of approximately 3.0g of BG powder, previously homogenized on a Bruker D4 Endeavor equipment, under the following operating conditions: Co Kα radiation (35kV/40mA); goniometer speed 0.02° 2θ per step with 1s count time per step and collected from 5 to 80° 2θ. The qualitative interpretation spectra were made by comparison with standards contained in the Bruker DiffracPlus software PDF02 database.

2.5.4Zeta potential measurements

The zeta potential measurements of the BG colloidal particles were performed on a Zetasizer Nano (Nano-ZS) apparatus coupled to a multipurpose titrator (MPT-2) (Malver). Measurements were performed for the particle charges at different pH values. A 0.01mM NaCl solution prepared with Milli-Q water was employed as the indifferent electrolyte for determining the zeta potential of the particles. The colloidal suspension was prepared in water at 0.5% mV−1. Separate 0.1M HCl and 0.01M NaOH solutions used for the pH adjustments by the automatic titrator were also prepared.

2.5.5Scanning electron microscopy

To evaluate BG morphology, micrographs were obtained by X-ray spectrometry (EDS) on a Scanning Electron FEI Quanta 400 Microscope, equipped with a Bruker XFLASH 4030 Sparse Energy System. The data acquisition and processing were obtained by use of the Esprit version 9.1 Brucker software. The samples were dried in an oven for 24h at 60°C, placed in metal discs tapes supported on carbon and underwent a micrometric carbon film coating from evaporations in a metallizing chamber. The elemental analysis by EDS was performed at different surface points, in order to minimize any possible damage resulting from the heterogeneity of the analyzed surface.

3Results and discussion3.1Initial pH effect

Fig. 1 displays the removal and uptake percentages of Ni(II) ions in the pH range of 2–7. The most significant removal was obtained at pH 5, corresponding to 37.4%, and a residual concentration of 32.3mgL−1. At pH 6 and 7 a decrease in removal was observed, corresponding to 31.9 and 33.7%, respectively. The least significant removal values were obtained at pH 2–3, corresponding to 17.9 and 13.3%, respectively. The maximum uptake value was of 19.3mgg−1.

Fig. 1.

Effect of pH on Ni(II) sorption by bioclastic granules (initial metal concentration: 50mgL−1; sorbent concentration: 1gL−1; particle size: 150–106μm; temperature: 25°C, and contact time: 60min).

(0.06MB).

The initial pH of the metal solutions were measured and adjusted to different values, after the addition of BG. At the end of the tests, new pH measurements were carried out. The final pH values ranged from 7.7 to 8.3 in all assays, regardless of the initial pH.

Similar behavior was found by Sdiri et al. [34], who observed that systems composed of different types of limestone and various metal ions (Pb2+, Zn2+, Cd2+ and Cu2+) always showed a final pH of 7.5, and suggested that the cause of the relatively constant final pH is the buffering capacity of the carbonate system.

Several studies report that the metal ion sorption capacity curve relative to pH decreases significantly near the pH range between 6.0 and 8.5 in the presence of CO2, due to the formation of carbonate-metal aqueous species [35]. Furthermore, the surface chemistry of carbonate minerals in contact with water is described in terms of the coordination reactions that these surface sites suffer with species in solution, such as Ca2+, CO32−, HCO3 and CaHCO3+[36].

To further elucidate the possible species involved in the removal of Ni(II) ions in the present study, a logC graph in function of pH speciation diagram was constructed (Fig. 2), with the aid of the computational software program MEDUSA – Make Equilibrium Diagrams Using Sophisticated Algorithms. As can be seen in Fig. 2, as the pH of the system rises, the concentration of H+ ions decreases significantly, increasing the formation and concentration of various soluble species such as the following:

Fig. 2.

Speciation diagram for the CaCO3 and Ni2+ system ([CO32−]=8.0mM; [Ca2+]=8.0mM, and [Ni2+]=0.85mM).

(0.1MB).

According to Kwon et al. [35], heavy metal ions in aqueous solutions can undergo solvation, hydrolysis and polymerization, and may form various hydrolysis products, which exist under different conditions. As seen in Fig. 2, at pH around 5, the Ni(II) ion concentration decreases gradually, resulting in the formation of more stable compounds. The reactions that describe the formation of these compounds are:

According to Pokrovsky and Schott [37], high amounts of H+ ions are present in solution at lower pHs, which can compete with metal ions for CO32−, inhibiting metal precipitation due to the carbonate present on the surface of the sorbent. Fig. 2 demonstrates that NiCO3 and NiOH+ begin forming between pH 3 and 3.6, respectively, with gradual increases in concentration until pH 7.3. From this pH onwards, Ni(OH)2 emerges and is maintained practically stable. Despite the fact that NiCO3 and NiOH+ coexist at about the same pH range, NiCO3 presents a more favorable equilibrium constant value (k=6.9×106), than NiOH+ (k=3.16×10−10).

According to several authors, solids with chemical composition based on CaCO3 can be considered good sorbents for heavy metals. Due to high percentages of calcium carbonate, calcareous algae have low metabolic rates and low proportion of metabolic tissue, which favors the accumulation of metals in relation to non-calcified organisms [38,39]. Moreover, the presence of soluble carbonate species causes precipitation in the form of metal oxides and probably carbonates [40]. This carbonate system tends to neutralize any acidity resulting from the presence of H+ ions formed in the reactions described above. This behavior has been previously described by Stumm and Morgan [41], in a system buffered by calcite.

3.2Effect of particle size

Fig. 3 displays the removal and uptake percentages of Ni(II) ions in the particle size range of <38–300μm. Particle size effects on sorption followed a distinct trend, in which increases in particle size increased Ni(II) removal and decreased Ni(II) uptake. The highest removal and uptake values were observed in the particle size range of <38μm corresponding to 46.4% and 26.8mgg−1, respectively.

Fig. 3.

Effect of particle size on Ni(II) sorption by bioclastic granules (initial pH: 5.0; initial metal concentration: 50mgL−1; sorbent concentration: 1gL−1; temperature: 25°C, and contact time: 60min).

(0.07MB).

Several authors describe a similar behavior in mineral sorption processes, such as calcite and aragonite, for different metal ions. Du et al. [42] evaluated the effect of particle size on the sorption of Pb, Cd and Zn in aragonite and calcite, and found that particle size for both carbonates and metal ions, with the exception of Cd, removed by calcite, show the same behavior, of higher uptake occurring in the presence of smaller particles, with higher sorption capacities in particle sizes ranging between 38 and 75μm.

Sánchez and Ayuso [43] observed that the maximum uptake values (qmax) of Cr, Zn and Cd in calcite drastically decreased with increasing particle size range of <100μm to 200–1000μm. According to the authors, this decline is related to the decreased surface area available to react with the metal solution.

3.3Effect of sorbent concentration

The effect of BG concentration on the removal and uptake of Ni(II) ions was evaluated in the range 0.5–3.0gL−1, as displayed in Fig. 4. The best removal and uptake results were obtained at 3.0 and 0.5gL−1, corresponding to 34.1% and 22.7mgg−1, respectively. Removal gains obtained by increasing concentrations from 1.0 to 3.0gL−1 were negligible, from 30 to 34.1%, respectively, while Ni(II) uptake decreased markedly from 22.7mgg−1 (0.5gL−1 concentration) to 5.9mgg−1 (3.0gL−1 concentration).

Fig. 4.

Effect of sorbent concentration on Ni(II) sorption by bioclastic granules (initial pH: 5.0; initial metal concentration: 50mgL−1; particle size: <38μm; temperature: 25°C, and contact time: 60min).

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The increase in removal efficiency by increasing the initial sorbent concentration is due to a greater number of active sorption sites and CO23− ions released in the solution, increasing the probability of forming precipitates [44,45]. In actual processes, higher sorbent concentrations are applied to ensure complete removal of metal ions in effluents, since they should be in accordance to the limits established by the responsible agencies. Nevertheless, the results of the present study suggest that sorbent concentration must be taken into consideration when CaCO3 based materials are used, in order to control the cost/benefit ratio. Both uptake capacity and removal efficiency of Ni(II) ions are also important in the sorption process.

3.4Effect of Ni(II) concentration

Sorption tests were conducted using BG and different metal solutions with initial Ni(II) concentrations in the range of 5–500mgL−1. Table 1 shows the removal and uptake values of Ni(II) ions at different initial concentrations at pH 5.0; initial sorbent concentration of 1gL−1; particle size: <38μm; temperature: 25°C; and contact time: 60min.

Table 1.

Effect of the initial concentration of Ni(II) on sorption by bioclastic granules.

Initial concentration (mgL−1Removal (%)  Uptake (mgg−1
63.74  3.63 
20  52.48  12.70 
50  29.49  14.83 
70  28.59  17.90 
100  25.81  26.07 
150  25.32  39.00 
200  18.52  40.00 
250  17.12  44.33 
300  14.66  45.00 
400  10.84  43.67 
500  8.86  45.00 

As displayed in Table 1, the removal of Ni(II) ions followed a trend of dramatically decreased removal with increasing concentrations, of 63.7% at a concentration of 5mgL−1 and 8.8% at 500mgL−1. Reddy et al. [46] observed similar behavior in assays applying calcite as a filter medium, in which Cd removal drastically decreased with the increase of the initial concentration in the range of 15–300mgL−1. Uptake behavior seems to follow an opposite trend, of gradual increases with initially high concentrations. For Ni(II) ions, the lowest uptake was obtained at 5mgL−1 corresponding to 3.6mgg−1, while more significant uptake values seem to remain in a concentration range of 250–500mgL−1, corresponding to 43.7–45.0mgg−1.

3.4.1Sorption isotherms

Assays were performed at initial concentrations from 5 to 500mgL−1, to calculate maximum BG saturation capacity. Fig. 5 displays the amount of sorbed Ni(II) ions (q) versus the concentration at equilibrium (Ceq), as well as the employment of Langmuir, Freundlich and Dubinin–Radushkevich isotherms. Fig. 5 indicates that the BG Ni(II) ion sorption isotherm profile resembles the “L” (Langmuir) type. This isotherm is given to solids with relatively small external surfaces with adsorption forming only a monolayer. The relationship between the concentration of solute remaining in solution and that sorbed on the solid decreases as the solute concentration increases, resulting in a concave curve, suggesting a progressive saturation of the solid. Table 2 shows the parameters obtained from the isotherm models, as well as their respective correlation coefficients.

Fig. 5.

Langmuir, Freundlich and Dubinin–Radushkevich isotherms for sorption of Ni(II) ions by bioclastic granules (pH: 5.0; concentration of sorbent: 1gL−1, particle size: <38μm; temperature: 25°C and contact time: 60min).

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Table 2.

Isotherm model parameters and their correlation coefficients.

LangmuirFreundlichDubinin–Radushkevich
qmax (mgg−1kL (Lmg−1R2  kF (Lmg−1n  R2  qmax (mgg−1B (mg2kJ−2R2 
54.9  0.014  0.957  5.5  2.74  0.908  44.2  29.57  0.946 

In general, as seen in Table 2, Ni(II) sorption by BG are well-fitted to the Langmuir (R2=0.957) and Dubinin–Radushkevich (R2=0.946) models. Aziz et al. [40], in their studies on Cu2+, Zn2+, Pb2+, Cd2+, Ni2+ and Cr3+ sorption in limestone also observed that the data were well-fitted to the Langmuir model (R2>0.90), suggesting that metal ion sorption was predominantly carried out as a monolayer on the limestone surface.

The qmax and kL values obtained by the Langmuir isotherm model for Ni(II) ions were 54.9mgg−1 and 0.014Lmg−1, respectively. The qmax value represents the maximum adsorption capacity when the surface of the sorbent is completely covered by metal ions and the high kL value indicates a high affinity between BG and Ni(II) ions. The separation factor or equilibrium parameter, RL, calculated from BG Ni(II) adsorption data is displayed in Table 3. All RL values obtained for the various concentrations from 5 to 500mgL−1 are between 0 and 1, suggesting that Ni(II) ions sorption is favorable.

Table 3.

RL values for Ni(II) sorption by bioclastic granules.

Initial concentration (mgL−1)
20  50  70  100  150  200  250  300  400  500 
0.935  0.781  0.588  0.505  0.417  0.323  0.263  0.222  0.192  0.152  0.125 

Ma et al. [47], in his studies on the sorption of Ni(II) ions in a CaCO3–altose hybrid material demonstrated that the Langmuir model presented R2=0.981. The maximum adsorption capacity (qmax) and the constant kL for Ni(II) ions obtained by the Langmuir isotherm were 769.23mgg−1 and 0.003Lmg−1, respectively. The RL values were all in the range from 0 to 1, confirming the positive uptake of these ions in the hybrid material. For the Freundlich model, the kF and n constant values were 5.84Lg−1 and 1.45, respectively. In the present study, the kF constant (5.5Lmg−1) represents a relative measure of the adsorption capacity and n (2.74) is related to the adsorption intensity. As observed in Table 2, the BG sorption data for Ni(II) ions was well-fitted to the Dubinin–Radushkevich model, with an R2 of 0.946. The qmax and B values obtained by the D-R model were 44.2mgg−1 and 297.57mg2kJ−2 respectively. According to Wang et al. [30], the mean free sorption energy, Es, can be related to the type of adsorption. For Es values <8kJmol−1, the adsorption is predominantly of a physical nature, i.e. BG Ni(II) sorption occurs in this condition, with a value of 0.04kJmol−1.

3.5Effect of contact time

According to Davis et al. [48] the sorption rates of metallic ions by calcium carbonate minerals have generally been characterized by both fast and slow processes, with reaction half-lives of the order of a few minutes up to hours and days. The effects of contact time on BG Ni(II) sorption are displayed in Fig. 6.

Fig. 6.

Effect of contact time on the sorption of Ni(II) ions by Bioclastic Granules (pH: 5.0; sorbent concentration: 1gL−1; initial metal concentration: 50mgL−1; particle size: <38μm, and temperature: 25°C).

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The removal and uptake of Ni(II) ions reached 32% and 17.2mgg−1 at 5min. They both then rise slowly, reaching 38.3% and 20.5mgg−1 at 20min, and from 20min onwards a plateau is present until 240min. Non-representative increases in removal and uptake were observed from 30 to 240min, accounting for removals in the range from 38.6 to 41% and uptakes from 20.6 to 22mgg−1, respectively.

Several authors describe that the rate of metal ion sorption in calcite are related to the isotopic exchange of Ca2+ between solutions (aqueous and solid), where the first adsorption phase, faster, was due to changes on the calcite surface, while the slower exchange process was the result of a structural diffusion and recrystallization on the calcite surface [48–52]. In fact, in the present study, the concentration of Ca(II) ions in solution during the sorption process increased as the metal ions were removed from the aqueous solution.

According to Prieto et al. [14] in their studies with two calcium carbonate polymorphs (calcite and aragonite) used for removal of metal ions, found that at baseline with aragonite Ca2+ ions were virtually absent, and that concentrations increased rapidly during the first 3h. This rapid dissolution was accompanied by an increase in the concentrations of carbonate species in solution, which led to a quick increase in pH during the first 1h, followed by a slow increase toward a limit pH value of 8.2. During the calcite assays, however, Ca2+ release was significantly slower and of Ca2+ concentrations after long reaction times were lower compared to assays conducted with aragonite.

As the initial pH of the solutions containing Ni(II) were conducted in a slightly acid medium (5.0), when adding the GB (that consists of CaCO3), species like Ca(HCO3)+ and Ni(HCO3)+ (Fig. 1) predominated in the solid interface/solution. Since the solubility product of these species corresponds to 3.67×10−12 and 3.39×10−13, respectively, a greater abundance of the Ni(HCO3)+ species, which is a precursor of NiCO3 (Kps=1.3×10−7), formed on the GB surface until the time of 20min. From this contact time, the surface area available for dissolution is reduced, leading to lower dissolution rates of the species described above. As the aqueous carbonate concentrations decrease, the same occurs with the NiCO3 supersaturation and, therefore, with its precipitation rate up to 240min.

3.5.1Sorption kinetics

During the sorption process, the reaction rate is given by measuring the amount of Ni(II) ions that disappear from the solution in a given time interval. On the other hand, the order of the reaction, or more precisely the form of expression of its velocity, is considered a magnitude that is determined experimentally, which results from the mechanism which regroups the atoms and bonds between the Ni(II) ions and the BG. Fig. 7 shows the pseudo-first order and pseudo-second order kinetic models for BG sorption of Ni(II) ions.

Fig. 7.

Pseudo-first and pseudo-second order models for BG Ni(II) sorption (pH: 5.0; sorbent concentration: 1gL−1; initial metal concentration of: 50mgL−1, particle size: <38μm, and temperature: 25°C).

(0.06MB).

In many cases, the sorption data applied to the Lagergreen equation of the pseudo-first order kinetic model fit well only to the initial stage of the adsorption process, and it is assumed that the rate limiting step may be the chemical adsorption [53,54]. Several researchers have reported that the sorption of divalent metal species in heterogeneous sorbents follows a pseudo-second order kinetic mechanism [55,56]. Table 4 displays the kinetic parameters obtained using both a pseudo-first and pseudo-second order model for Ni(II) ions.

Table 4.

Kinetic parameters obtained from pseudo-first and pseudo-second order kinetic models for BG Ni(II) ions sorption.

Pseudo-first orderPseudo-second orderqeqexp (mgg−1
qeq (mgg−1k1 (min−1R2  qeq (mgg−1k2 (gmg−1min−1R2   
21.02  0.288  0.979  21.74  0.028  0.994  21.97 

The obtained kinetic data were well-fitted to the kinetic pseudo-first order model, with R2 0.979; qeq of 21.02mgg−1 and k1 of 0.288min−1. However, the sorption data for Ni(II) ions exhibited a better fit to the pseudo second-order kinetic model, with R2 0.994; calculated qeq of 21.74mgg−1, very close to the qeq obtained experimentally (21.97mgg−1) and k2 0.028gmg−1min−1.

Du et al. [42] observed that the data sorption of Pb, Cd and Zn in calcite and aragonite were well-fitted to the pseudo-second order kinetic model. All sorption data showed an initial rapid phase, followed by a slow phase and partial equilibrium. After this initial phase, a layer is formed by the adsorption and precipitation of metals on the surface of both aragonite and calcite, significantly reducing the available sorption sites and the CO32− released from these adsorbents, making the sorption rate slower. This slower sorption rate in the second phase can also be explained by steric hindrance from the precipitates of the sorbed metal ions and metal diffusion within the porous structure and/or the crystal lattice of biogenic sorbents.

3.6Characterization3.6.1X-ray fluorescence

Table 5 shows the semi-quantitative analyses of oxides of BG constituent elements obtained by X-ray fluorescence (XRF). Calcium oxide (CaO) comprises about 45.5% of the material, followed, in much smaller proportions, by magnesium oxide (MgO – 3.5%), silica (SiO2 – 1.5%) and sodium (Na2O –1.2%). The other oxides account for just over 3.5% of the material. Approximately 44.7% of the material was lost on ignition.

Table 5.

Fluorescence X-ray analysis of bioclastic granules.

Constituent (%)
CaO  L.O.I.a  MgO  SiO2  Na2Al2O3  Cl  SO3  Fe2O3  SrO  P2O5  K2TiO2 
45.50  44.70  3.54  1.50  1.21  0.93  0.73  0.72  0.50  0.39  0.18  0.06  0.04 
a

Loss on ignition.

Different authors have obtained similar compositions to those found in this study for carbonate minerals. Yavuz et al. [57], characterizing a natural calcite from Turkey, noted that the chemical composition corresponded to 56% CaO, 17.0% MgO and 0.7% Fe2O3. Aziz et al. [40], when analyzing the chemical composition of a limestone applied in the removal of Ni(II), produced as a byproduct of the marble industry, had the following percentages: 53.52% CaO, 1.44% of SiO3, 1.14% MgO, 0.27% Fe2O3, 0.21% Al2O3, 0.06% K2O, 0.02% P2O5, 0.01% MnO, 0.01% of SrO and 43.3% CO2.

3.6.2Surface area measurements of specific and porosity

The nitrogen adsorption–esorption isotherm for <38μm BG samples is displayed in Fig. 8. The N2 isotherm shows a profile corresponding to the type III isotherm with hysteresis loop type H3, with a sharp rise from the relative pressure of 0.8, suggesting the presence of macropores in the BG. According to Sing et al. [58] this type of hysteresis isotherm indicates the presence of a slot, cone and/or pyramid-shaped macroporous structure.

Fig. 8.

N2 adsorption–desorption isotherm by <38μm bioclastic granules.

(0.05MB).

According to Webb and Orr [59], some structural properties should be considered when working with fine particles, such as size, surface area, volume, and pore size. Table 6 shows some of the structural properties of bioclastic granules. As seen in Table 6, both the micropore and the external surface areas exhibit values close to 1.5112 and 1.2982m2g−1, respectively. Therefore, the BET surface area corresponds to 2.8094m2g−1. Aziz et al. [40], measuring the surface area of a limestone obtained as by-product of the marble industry by the N2 BET adsorption method found a value of 0.04m2g−1, while Gómez del Río et al. [60] reported that the calcite used in their sorption studies for removal of Cd, Zn and Co had a surface area of 0.46m2g−1.

Table 6.

Structural properties of bioclastic granules.

Particle size  Area (m2g−1)Volume (cm3g−1Pore diameter (Å) 
  Micropores  External surface  BET surface     
<38μm  1.5112  1.2982  2.8094  0.0066  298.8 

According to Sing et al. [58] pores with diameters <20Å are classified as micropores, from 20 to 500Å as mesopores and >500Å as macropores. As BG have pore diameters of 298.8Å, they can be classified as mesopores. The total pore volume, which corresponds to all breaks, cracks, holes and channels volumes within the particle bodies [61,62] corresponds to the value of 0.00660cm3g−1.

Prieto et al. [14] comparing three types of calcium carbonate in relation to surface area and total pore volume, observed that calcite showed values of 0.066m2g−1 and 0.00021cm3g−1, non-biogenic aragonite showed values of 0.29m2g−1 and 0.00059cm3g−1, and biogenic aragonite (Cerastoderma edule) showed values 2.66m2g−1 and 0.0102cm3g−1, respectively. According to the authors, the higher uptake of Cd ions by biogenic aragonite is due to its higher specific surface than other types of carbonates, because each particle is a highly structured crystalline aggregate, with a more reactive surface than a fragment of a single similar crystal.

3.6.3X-ray diffraction

The crystal properties depend on the arrangement of their atoms in the unit cell and their periodicity. Calcium carbonate is a polymorphic material that naturally presents three crystal structures, calcite, aragonite and vaterite. According to Fisler et al. [63], CaCO3 crystals may present some punctuate impurities caused by isomorphic substitution of Ca by some metallic elements in the crystal lattice. Fig. 9 displays the BG X-ray diffraction patterns before and after the Ni(II) ions sorption process.

Fig. 9.

X-ray diffractograms of bioclastic granules before and after the sorption process with Ni(II) ions.

(0.08MB).

Localized peaks at angles 27.18; 34.81; 46.8; 51.26; 56.83; 57.83; 68.84; 73.01 and 77.96°, corresponding to the interplanar spacings of 3.27; 2.57; 1.93; 1.78; 1.61; 1.59; 1.36; 1.29 and 1.22Å, belonging to the calcite phase, respectively, can be observed (Fig. 9). The angles of 30.52; 31.70; 38.62; 44.24; 50.24; 53.76; 59.09 and 61.68° correspond to the interplanar spacing's of 2.92; 2.81; 2.32; 2.04; 1.81; 1.70; 1.56 and 1.50Å, characteristic of the aragonite phase, respectively. As seen by these results BG have two predominant phases, calcite and aragonite, as expected, since calcite is a major constituent of lime sandstones from thick marine deposits of calcium carbonate, which may acquire the structure of aragonite if intensively ground [64]. Lim and Arizona [65] using dead calcareous skeletons (Scleractinians) for the removal of Cd in aqueous solutions noted that the CaCO3 phases obtained by XRD were aragonite and calcite.

As suggested by some authors, the incompatibility of a network is a key factor in determining whether a film grows on a given substrate [14,66,67]. When observing Fig. 9, it can be seen that a single peak relating to the gaspeite phase, was identified at a 38.64° angle, corresponding to the interplanar spacing of 2.32Å. This large lattice incompatibility between BG and the solid nickel carbonate solution may not favor heteroepitaxial growth, i.e. the sorption process is not based solely on metal chemistry, but also on their crystallographic relations with the sorbent.

Zachara et al. [68] suggest that Ni ions form hydrated complexes on the calcite surface and are then incorporated into the mineral structure by recrystallization. Several authors, using various surface analysis techniques, reached a similar conclusion, that Ni was incorporated in calcite by rearrangement of the solid, but that uptake was much slower than for other metals [69–71].

3.6.4Zeta potential measurements

According to Fenter et al. [72], interactions between minerals and metals may be susceptible to variations in pH-dependent surface species and charge. Several authors have studied the electrical properties of carbonate mineral surfaces, mainly calcium carbonate, by electrokinetic studies [36,73–76]. Fig. 10 shows the zeta potential curves as a function of pH for BG before and after sorption of Ni(II) ions in the pH range of 2–10.

Fig. 10.

Bioclastic granules zeta potential before and after sorption of Ni(II) ions (initial metal concentration: 50mgL−1; sorbent concentration: 1gL−1; particle size: <38μm; temperature: 25°C, and contact time: 240min).

(0.06MB).

BG surface charge is positive in the pH range from 2 to 4 (Fig. 10), pH 4.2 corresponds to the isoelectric point, and above this pH BG charge becomes positive. Somasundaran and Agar [77], based on thermodynamic calculations of a calcite/aqueous solution/atmosphere system, suggest that Ca2+, HCO3, CO32−, H+ and OH are the potential-determining ions. Sondi et al. [36], on the other hand, describe that the surface charge and, consequently, the electrokinetic potential of calcite, depends on the redistribution of the potential-determining ions, Ca2+, HCO3, CO32−, H+ and OH, between the surface and the solution, so H+ and OH would be of secondary importance, and would have direct influence on the electrical properties of the surface. Their function would be only speciation regulation of the carbonate ions in the solution. Thus, the positive zeta potential values from 18 to 1.8mV (pH 2–4) are probably due to Ca2+ ions. In the range of pH 5–10, this charge becomes negative (−8, 5 to −18.3mV) due to the increased concentrations of HCO3 and CO32− from the carbonate system. Similar results were observed by Gómez del Río et al. [60], in which gradual increases in pH led to more negatively charged calcite particles, until approximately pH 9.0, where the negative charge was then maintained almost constant, until pH 10.5.

In the present study, all sorption tests, even those with initial pH in the acidic range, showed a natural tendency to rise to values in the range of 7.7–8.3. According to Reddy et al. [46], in their assays using calcite as filter media for the removal of Ni(II) ions, pH ranged from 6.3 to 9.6, a range in which functional groups on the calcite surface are neutral or slightly negative, thus favoring metal adsorption. In fact, as can be seen in Fig. 10, after the sorption process with Ni(II) ions, the BG charges remained essentially constant throughout the studied pH range, with zeta potential values varying from −2.6 to 3.3mV.

3.6.5Scanning electron microscopy

Analyses were performed with a scanning electron microscope (SEM) in order to evaluate BG morphology before and after the Ni(II) ion sorption process. Fig. 11 displays BG before and after Ni(II) sorption. Fig. 11A and B show BG before the sorption process, with very irregular surfaces and the prevalence of different-sized cavities in the range of 7.1–9.5μm and smaller fragments within and around these cavities.

Fig. 11.

SEM images of bioclastic granules before (A and B) and after (C and D) sorption of Ni(II) ions, with magnification of (A) 1200×; (B) 5000×; (C) 5000×; and (D) 5000×.

(0.58MB).

Fig. 11C and D shows bioclastic granules after Ni(II) sorption. Fig. 11C shows BG cavities mostly taken up by NiCO3 precipitates. Lee et al. [78], in studies on the characterization of a calcium carbonate granule employed to remove Ni, found that most of the ions removed from the aqueous phase were incorporated as oxides, hydroxides or calcium salts and calcium carbonate on the surface of the CaCO3 particles. The SEM micrographs suggest that the particles formed by the reaction are mainly composed of precipitated colloidal calcium carbonate.

Fig. 11D shows a layer formed by NiCO3 precipitates on BG surfaces, with a thickness of approximately 1.75μm. Xu et al. [79], by Atomic Force Microscopy, observed that, after an initial growth phase both laterally and vertically, otavite crystals reached a height of 2.2–2.7nm, remaining constant over the time interval from 2 to 6h, suggesting that the end times are independent of the initial concentration of Cd2+ in the solution, and that crystal growth is limited by the diffusion in monolayers formed on the surface.

Fig. 12 displays BG surface spectra by energy dispersive X-ray (EDX) before and after adsorption of Ni(II). The BG spectrum before sorption shows three well-pronounced peaks related to Ca, C and O, characteristic of CaCO3 minerals, in addition to peaks related to elements Mg, Al and Si. After sorption, Ni(II) peak appear, and decreases the three significant peaks Ca, C and O, especially the Ca peak.

Fig. 12.

Bioclastic granules EDX spectrum before and after Ni(II) sorption.

(0.08MB).
4Conclusions

Based on the results obtained in the present study, it can be concluded that at an initial pH of 5 removal of 37.4% and uptake of 19.3mgg−1 Ni(II) was observed. At BG particle size of <38μm, removal and uptake were 46.4% and 26.8mgg−1, respectively. The initial sorbent concentrations of 3.0 and 0.5gL−1 showed the best removal and uptake results, corresponding to 34.1% and 22.7mgg−1, respectively. The highest Ni(II) removal (63.7%) was observed at the initial Ni(II) concentration corresponding to 5mgL−1, while the highest uptake was observed at 500mgL−1 (45mgg−1). The sorption data was well-fitted to the Langmuir isotherm model (R2=0.957) presenting qmax=54.9mgg−1 and kL=0.014Lmg−1. The sorption process proved very fast in the first 20min of contact time, reaching Ni(II) removal and uptake of 38.3% and 20.5mgg−1. The uptake data were well-fitted to the pseudo-second order kinetic model (R2=0.994), with qmax=21.74mgg−1 and k2=0.028Lmg−1. Fluorescence X-rays analysis indicated that CaO, MgO, Na2O and SiO2 are the major BG constituents. The BET surface area of the particles BG was 2.8094m2g−1. X-ray diffraction showed the presence of both a calcite and aragonite phase. Zeta potential measurements demonstrated that BG surface charge is positive in the pH range from 2 to 4, with the isoelectric point present at pH 4.2. Above this pH, the charge becomes positive. Scanning electron microscopy indicated that BG have a very uneven surface, with many different sized cavities, which are subsequently filled by NiCO3 precipitates.

Conflicts of interest

The authors declare no conflicts of interest.

Acknowledgements

The authors would like to acknowledge the support given by Center of Mineral Technology (CETEM/MCTI).

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