We have studied the pH and the temperature effects on copper ions' adsorption on natural and treated clays from Algeria. The clay was also treated to improve the adsorption capacity. X-ray diffraction identified montmorillonite and kaolinite as major clay minerals. The Langmuir adsorption model was used for the mathematical description of the adsorption equilibrium and the equilibrium data adhered very well to this model. The treated and natural clay had a monolayer adsorption capacity equal to 15.40 and 12.22 mg/g, respectively, at pH value of 6.5 and temperature 20 °C, the adsorption isotherms could be fitted with Langmuir isotherms, and the coefficients indicated favorable adsorption of Cu(II) on the clays. Determination of the thermodynamic parameters, H, S, and G showed that the adsorption process was spontaneous and exothermic accompanied by a decrease in entropy and Gibbs energy. Results of this study will be useful for future scale-up for using this material as a low-cost adsorbent for the removal of Cu(II) from wastewater.

INTRODUCTION

The removal of heavy metals from water and wastewater is important in protecting public health and the environment (Bellir et al. 2005). Many industrial activities such as metal plating, industry fertilizer, mining operations, metallurgy, manufacturing batteries, and dyeing in textile industries introduce heavy metals into the environment via their waste effluents (Benzina 1990). Precipitation, ion-exchange, ultrafiltration, membrane separation, and adsorption are the usual methods for the removal of heavy metal ions from aqueous solutions (Boddu et al. 2003).

Owing to its simplicity and easy operational conditions, adsorption is a widely used process. In this study, we report the adsorption of Cu(II) from aqueous solutions on natural and activated clays.

In continuation of previous works carried out in this field, in the present work, a more efficient method for modification of the clay was studied. One of the objectives of this study was to evaluate the effect of this modification on the capability and mechanism of Cu(II) adsorption from water by the produced adsorbent.

MATERIALS AND METHODS

Preparation of Na-montmorillonite

The natural clays were washed several times with distilled and deionized water and were completely dispersed in water. After 7 h at rest, the dispersion was centrifuged for 1 h at 2,400 rpm. The size of the clay particles obtained was 2 mL.

These clay particles were dispersed in water and heated at 75 °C in the presence of a solution composed of sodium salts of bicarbonate (1 M), citrate (0.3 M), and chloride (2 M) (Hatton & Pickering 1980; Das & Kumar 2006). The purpose of this operation was to eliminate inorganic and organic compounds, aluminum found in the inter-layer spaces and various free captions. Carbonates were removed by treatment with HCl (0.5 M) and chloride was eliminated after several washings. The organic matter was eliminated completely by treatment with H2O2 (30% v/v) at 70 °C. The purified clay was dried at 110 °C, and then saturated with sodium (Na+). To ensure complete transformation into the sodium form, all samples were washed several times with a NaOH solution (1 M). 40 g of Na-montmorillonite was treated with 400 mL of 5 N sulfuric acid (analytical grade) at 90 °C for 3.5 h (Foletto et al. 2003) in a stirred glass reaction vessel with reflux. After the acid treatment, the samples were filtered and washed with distilled water until they were free of . The samples were dried at 60 °C for 12 h and ground to pass through a 0.074 mm sieve. The activated clay acid was characterized using Fourier transform infrared spectroscopy (FTIR), X-ray diffraction (XRD), and differential thermal analysis/thermo gravimetric analysis (DTA/TG) techniques.

RESULTS AND DISCUSSION

X-ray diffraction

The final product of clay was separated by centrifugation, then washed and dried at 60 °C and studied, and XRD patterns of powder samples were recorded at room temperature under air conditions on a Siemens D-5000 instrument, using Cu Kα radiation (λ = 1.5406 Å) and TG/DTA techniques and FTIR spectroscopy.

Results of the X-ray diffraction analysis for natural and activated clay are shown in Figure 1. It clearly shows that the d-spacing of clay increased from 12.93 Å, 2Θ = 6.83 ° to 16.50 Å, 2Θ = 5.35 ° which could be attributed to the natural and activated clay. Quartz (reflection at d = 13.92 Å, 2Θ = 21.07, 26.34 °) and calcite (reflection at d = 4.42 and 2.49 Å, 2Θ = 20.05 °, 35.97 °) are the major impurities. The reflection at d = 7.07 Å, 2Θ = 12.5 ° is characteristic of kaolinite. Montmorillonite and kaolinite as major clay minerals were identified by XRD (Figure 1).
Figure 1

XRD pattern for natural clay and activated clay.

Figure 1

XRD pattern for natural clay and activated clay.

Infrared spectroscopy study

The IR spectra of the natural and activated samples were recorded over the spectral range 400–4,000 cm−1. In fact, IR techniques have been used by many researchers for identification of natural clay minerals (Hajjaji et al. 2001).

The characteristic vibrations of hydroxyl groups, the silicates anions, and the octahedral captions are present in the IR spectra of the studied samples, as shown in Figure 2. The peak positions agree with the values given in the literature (Madejova 2003). In the region of 3,700–3,200 cm−1, a number of sharp peaks at 3,625 and 3,460 cm−1 are observed; the region above 3,000 cm−1 wavenumber contains information about the silanols.
Figure 2

FTIR spectra of natural and activated clay.

Figure 2

FTIR spectra of natural and activated clay.

The characteristic vibration peaks of montmorillonite are at 3,625 cm−1 (O–H stretching), 1,040 cm−1 (Si–O stretching), 628 cm−1 (Al–O–Si stretching), 525 cm−1 (Si–O–Al bending), and 467 cm−1 (Si–O–Si bending). The small peaks in the region of 1,650–1,500 cm−1 are due to O–H stretching vibrations from H–O–H, which implies the presence of little amount of adsorbed water in the clay samples. The peak at 914 cm−1 and weak band at 440 cm−1 were assigned to O–Si–O asymmetric stretching; the doublet at 850 and 790 cm−1 indicates the presence of quartz. The characteristic bands of montmorillonite-Na were observed at 1,039 and 467 cm−1 (Madejova 2003). The high intensity of the peak appearing at 1,039 cm−1 is an indication of the large amount of this mineral in the sample.

Thermo-gravimetric analysis

The thermal analysis diagrams for natural and activated clays are shown in Figures 3 and 4.
Figure 3

TGA curve of natural clay and treated clay.

Figure 3

TGA curve of natural clay and treated clay.

Figure 4

DTA curve of natural clay and modified clay: (a), dehydration; (b), dehydroxylation; (c), calcination.

Figure 4

DTA curve of natural clay and modified clay: (a), dehydration; (b), dehydroxylation; (c), calcination.

The thermogravimetric results for activated clay (Figure 3) revealed a weight loss corresponding to free and absorbed water on the outer surface of montmorillonite and organic materials (Heller-Kallai et al. 2006) due to acid treatment in the range of 34 to 133 °C. The second peak occurring between 397 and 536 °C refers to the loss of hydroxyl in the structure (Heller-Kallai et al. 2006). The corresponding weight loss for natural and activated clays was 8.69% and 2.04%, respectively.

Results showed an endothermic peak in the DTA curve of natural clay in the range of 25 to 130 °C due to the dehydration of clay minerals as shown in Figure 4. The second endothermic phenomenon took place between 350 and 600 °C; this peak refers to the loss of hydroxyl in the structure (Ghosh & Bhattacharyya 2002). The corresponding loss of bicarbonates and carbonates are related to the endothermic peak around 283 and 702 °C, respectively.

Adsorption of copper ions

Experimental procedure

Solutions of Cu(II) concentrations in the range of 10–100 mg/L were prepared from a stock solution of CuSO4, 5H2O. The pH was adjusted with 0.1 M NaOH or HNO3.

Amounts of 0.2 g clay were dispersed in the different copper salt solutions and shaken for 2 h (Yavuz et al. 2003). The dispersions were filtered, and the copper concentration was determined by spectrophotometry at a wave-length of 805 nm (visible spectrophotometer using KBr disc method). The amounts of Cu(II) adsorbed were calculated from the concentration differences.

Effect of pretreatment on adsorption

The increase of the adsorption capacity of Cu(II) ions added is shown in Figure 5. In addition, it is clearly seen that the adsorption capacity, for adsorption of Cu(II) on activated clay was significantly higher than adsorption on natural clay from the aqueous solution.
Figure 5

Pretreatment on the adsorption capacity of Cu(II) on the natural clays and activated clay adsorbents at 20 °C, pH 6.5.

Figure 5

Pretreatment on the adsorption capacity of Cu(II) on the natural clays and activated clay adsorbents at 20 °C, pH 6.5.

Up to 50 mg/g and 27.5 mg/g of Cu(II) was successfully adsorbed by activated and natural clay. This again reveals that a part of Cu(II) ions was bonded by cation exchange on the negative surface sites. The involved mechanism was the exchange of proton hydrogen ions of the activated clay by copper ions. This proves the feasibility of activated clay as an effective adsorbent.

Adsorption isotherms of Cu(II) on the natural clays and activated clay is shown in Figure 5.

Effect of pH on adsorption

Adsorption of heavy metal ions is often pH dependent (Hatton & Pickering 1980; Stumm 1992). The influence of pH is shown in Figure 6. Adsorption of copper ions was maximal at pH 6.5. Similar results were reported by Bellir et al. (2005). After activation, the same effect was observed (Figure 7). At lower pH, the adsorption of protons of water competed with the adsorption of copper ions, and at higher pH copper hydroxide was precipitated.
Figure 6

Adsorption of Cu(II) on natural clay at different pH values.

Figure 6

Adsorption of Cu(II) on natural clay at different pH values.

Figure 7

Adsorption of Cu(II) on activated clay at different pH values.

Figure 7

Adsorption of Cu(II) on activated clay at different pH values.

Effect of temperature on adsorption

The Langmuir isotherm showed a fit with the experiment data (Figures 8 and 9). With increased temperature, the adsorption of Cu(II) increased (Figure 8) (Benzina 1990) confirming that the process was exothermic. For the activated clay, the same effect was observed (Figure 9).
Figure 8

Effect of initial concentration on adsorption under different temperatures at pH 6.5.

Figure 8

Effect of initial concentration on adsorption under different temperatures at pH 6.5.

Figure 9

Effect of initial concentration on adsorption under different temperatures at pH 6.5.

Figure 9

Effect of initial concentration on adsorption under different temperatures at pH 6.5.

Chemical composition of the natural clay and activated clay is shown in Table 1.

Table 1

Chemical composition of the natural clay and activated clay (in mass %)

SampleActivated clayNatural clay
SiO2 67.27 54.782 
Al2O3 21.25 30.794 
SO3 0.314 0.1 
K22.10 1.199 
MgO 2.204 1.761 
CaO 0.037 0.015 
Fe2O3 3.069 2.636 
Na22.57 4.33 
TiO2 0.265 0.1 
P2O5 0.015 0.03 
SampleActivated clayNatural clay
SiO2 67.27 54.782 
Al2O3 21.25 30.794 
SO3 0.314 0.1 
K22.10 1.199 
MgO 2.204 1.761 
CaO 0.037 0.015 
Fe2O3 3.069 2.636 
Na22.57 4.33 
TiO2 0.265 0.1 
P2O5 0.015 0.03 

The acid activation of clay is a two-step procedure in which the splitting of particles within the octahedral sheet takes place. In the first step the exchangeable cations are replaced by protons (H+). The second step involves the leaching of octahedral cations such as AL3+, Mg2+, and Fe3+ from the octahedral and the tetrahedral sheets (Steudel et al. 2009). The octahedral Al3+ cation could be more easily leached by acid attack.

Adsorption isotherms

The Langmuir model is the simplest and the most commonly used model to represent the adsorption from a liquid phase by a solid phase (Boddu et al. 2003). This model assumes a monolayer adsorption.

The obtained adsorption data were fitted by the linearized Langmuir equation: 
formula
1
where Ceq is the equilibrium adsorptive concentration in solution, Xm the monolayer capacity, (x/m) is the specific amount, and b is related to the adsorption energy.

The parameters derived from the least-squares fitting of the isotherms by the linearized Langmuir equation are given in Table 2.

Table 2

Langmuir parameters of adsorption isotherms at 20 °C

Langmuir equation
AdsorbentsXm (mg/g)b (L.mg−1)R2
Natural clay 12.22 0.62 0.992 
Activated clay 15.40 0.79 0.992 
Langmuir equation
AdsorbentsXm (mg/g)b (L.mg−1)R2
Natural clay 12.22 0.62 0.992 
Activated clay 15.40 0.79 0.992 

The monolayer capacity (Xm) for activated and natural clay was 15.40 and 12.22 mg/g, respectively. The higher b value of natural clay compared with that of activated clay showed that the adsorption of copper ions on the raw clay required more energy.

Thermodynamic parameters

The thermodynamic parameters for the adsorption of Cu(II) by natural and activated clay, such as the enthalpy change (ΔH°), the Gibbs free energy change (ΔG°), and the entropy change (ΔS°), can be calculated from the variation of maximum adsorption with temperature (T) using the following basic thermodynamic relations (Seki & Yurdakoc 2006): 
formula
2
 
formula
3
 
formula
4
where R is the gas constant, R = 8.314 × 10−3 kJ/mol.K; Kads is the equilibrium constant, T is the absolute temperature; ΔG° is the change in free energy, kJ/mol; ΔH° is the change in enthalpy, kJ/mol; ΔS° is the change in entropy, kJ/mol. According to Equation (2), the mean value of the enthalpy change due do the adsorption of Cu(II) by natural and activated clay over the temperature range studied can be determined graphically by the linear plotting of ln Kads against 1/T using the least squares analysis shown in Figure 10.
Figure 10

ln Kads as a function of 1/T at PH 6.5.

Figure 10

ln Kads as a function of 1/T at PH 6.5.

The mean enthalpy change can be determined from the slope of the straight line. The variation of Gibbs free energy and entropy change with temperature can be calculated using Equations (3) and (4), respectively; the results are shown in Table 3. An important result obtained from Table 3 is that the Gibbs free energy (ΔG°) is small and negative with its value decreasing with increasing temperature. This indicates that the adsorption processes of Cu(II) by natural and activated clay can be enhanced by decreasing temperature. The values of entropy change (ΔS°) are positive and remain almost constant with temperature. This provides evidence that structural changes in Cu(II) and natural and activated clay occur during the adsorption process (Figure 10). The negative values of enthalpy change (ΔH°) for adsorption are lower than 80 kJ/mol, suggesting the physical nature of the sorption, i.e., physisorption conducted with van der Waals forces.

Table 3

Thermodynamic parameters

 − ΔG° (kJ/mol)
− ΔH° ΔS°
293 °K303 °K313 °K323 °K(kJ/mol)(J/mol.K)
Natural clay 3.140 4.873 6.375 8.534 48.572 176.306 
Activated clay 4.039 5.649 6.981 9.155 33.391 127.769 
 − ΔG° (kJ/mol)
− ΔH° ΔS°
293 °K303 °K313 °K323 °K(kJ/mol)(J/mol.K)
Natural clay 3.140 4.873 6.375 8.534 48.572 176.306 
Activated clay 4.039 5.649 6.981 9.155 33.391 127.769 

The enthalpy changes (ΔH°) for adsorption are negative −33.3915 to −48.5720 kJ/mol, indicating physical adsorption of Cu(II) on to natural and activated clay (Seki & Yurdakoc 2006). The values of ΔG° are negative (Table 3), which means that the reaction of Cu(II) adsorption is spontaneous and exothermic (Suhas & Ribeiro 2007).

CONCLUSION

Adsorption isotherms were both well described by the Langmuir model. A maximum capacity of Cu(II) adsorbed on natural and activated clays at equilibrium was 12.22 and 15.40 mg/g, respectively. The negative value of ΔH° indicated that the adsorption process was exothermic in natural and activated clays, and the negative values of ΔG° at different temperatures (293–323 °K) indicated the spontaneous nature of Cu(II) adsorption. The adsorption reached a maximum at pH 6.5 and increased with temperature. The Cu(II) adsorbed by natural clay was lower compared with activated clay.

ACKNOWLEDGEMENTS

The authors would like to thank the University Of Technology Faculty of Chemistry, Oran, Algeria and Mascara's Laboratory University. The authors also thank Prof. A. Djafri for her assistance in XRD and IR measurements and educational experience, and not forgetting her helpful discussions. Finally, the authors offer a special thanks to Prof. M. Bouchekara for his guidance and support throughout this research.

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