Two clay raw materials are used in this study, “Aga” clay was taken from the town of Gagnoa located in the Center-West of Côte d’Ivoire, the sampling site coordinates of 6˚07'37.6'N, 5˚56'15.0'W. The second sample noted “AB” comes from the Akouai-Agban deposit, in the suburbs of the town of Bingerville on the lagoon shore in the south of Côte d’Ivoire, the coordinates of the sampling site are 5˚16'12, 883844N, 3˚53'23.325'W. These samples are presented in Figure 1 .

The solution to be treated (from BM) was prepared by introducing 0.0125 g of Methylene Blue powder into a 250 mL volumetric flask containing distilled water which was made up to the mark (i.e., a solution concentration 0.05 g/L). The mixture obtained was homogenized by stirring using a magnetic bar for thirty minutes (30 min).

During this process, the MB solution is subjected to a treatment of hydrogen peroxide (H₂O₂) and ferrous iron (Fe²⁺) in the presence of light in order to find the optimal concentrations of H₂O₂ and ferrous iron Fe²⁺. The whole is placed on a magnetic stirrer at a moderate speed allowing its permanent homogenization via a magnetic bar immersed in it. A manual sample of 2 mL to be taken every 15 minutes for 2 hours using 5 mL syringes. The samples are analyzed using a UV-Visible spectrophotometer. During this approach, to study the influence of clay on the degradation of Methylene Blue by the photo-Fenton process, the optimal concentrations of hydrogen peroxide (H₂O₂) and ferrous iron are added to the BM solution. (Fe²⁺) previously determined by the photo-Fenton process. In the entire mixture, we vary a quantity of clay from 0.05 g to 0.2 g. Everything is placed on a magnetic stirrer at a moderate speed allowing its permanent homogenization via a magnetic bar immersed in it. The sampling and analysis of the homogenized solution follow the approach described above.

The isoelectric point or pH at the point of zero charge is the parameter which corresponds to the pH for which the surface of the solid presents a zero charge [5] . To determine the pH at the point of zero charge of the clays used in this work, the method of Lopez-Ramon et al. taken from the thesis of N’guessan was used [6] [7] . Solutions of 0.1 mol·L⁻¹ of NaCl and pH between 2 and 10 (adjusted by adding an aqueous solution of NaOH or H₂SO₄) were first prepared. The HI 2211 pH meter was used for pH measurement. 0.1 g of the different dried clays is brought into contact with 20 mL of each of the solutions contained in beakers. The solutions are stirred for 3 days at room temperature. Each solution is then filtered using filter paper (Double Rings Filter Paper) and a new pH measurement is carried out. We draw the curve C representing the final pH as a function of the initial pH. The pH_PZC then corresponds to the pH of the solution for which the curve C crosses the first bisector (final pH = initial pH).

The aim of this technique is to determine the chemical composition of the clay sample using the ICP-AES (Industively Coupled Plasma-Atomic Emission Spectromely) Plasma Emission Spectrometer [8] . The sample to be characterized was put into solution and converted into an aerosol, then vaporized, atomized, excited and/or ionized. To do this, 30 mg of each sample is dried at 110˚C for 24 hours in a Teflon tube and the whole is introduced into a microwave device (CEM, MARCH 5).

X-ray diffraction is an analysis method that makes it possible to identify the main crystallized mineral phases present in the material. The X-ray diffraction patterns are obtained using a Bruker D8 Advance type device equipped with a copper anticathode (λ = 1.54 Ǻ).

Infrared spectroscopy is a technique for studying the nature and local environment of chemical bonds. It allows the determination of the functional group present in the material. There are three distinguished areas (near infrared, mid infrared and far infrared). For our study, the IR spectra was carried out in the mid-infrared domain using two spectrometers (Fourier Transform spectrometer (Tensor 27 Golden Gate Single) and Elmer Spectrum 1000 Fourier Transform spectrometer.

Spectrophotometry is a quantitative analytical technique, which consists of measuring the absorbance or optical density of a given chemical substance in solution. The rate of decrease in light intensity as a function of the thickness of the absorbing medium is given by the Beer-Lambert law according to the equation (Equation (1)):

$A=\log \frac{I_0}{I}=\epsilon \cdot l\cdot C$ (1)

With:

I₀: initial intensity of the light passed through; I: intensity of transmitted light; A: absorbance; C: concentration of absorbent species (mol·L⁻¹); l: optical path length in the solution (cm); ε: molecular absorption coefficient (L·mol⁻¹·cm⁻¹).

The aqueous BM solution was prepared by dissolving powdered BM in distilled water. The maximum absorption wavelength (λ_max) was obtained by scanning the λ between 200 and 800 nm. The peak of our dye is obtained at 663.1 nm. In order to obtain the MB calibration curve we prepared a stock solution with a concentration of 50 mg·L⁻¹. From the stock solution we determined the concentrations (1, 2, 3, 4, 5 mg·L⁻¹). These will subsequently be analyzed by UV-visible spectrophotometry. Figure 2 shows the BM calibration curve. The curve is linear over the interval of chosen concentrations, therefore the Beer-Lambert law is verified for this concentration range.

The percentage of dye degradation was calculated using the equation (Equation (2)):

$q_e=\frac{V}{m}\times \left(C_0-C_e\right)$ (2)

$E\left(\%\right)=\frac{C_0-C_e}{C_0}\times 100$ (3)

q_e: the quantity adsorbed per unit of clay (mg·g⁻¹); V: volume of MB solution (mL); m: the mass of the clay (g); C₀: the initial concentration of BM (mg·L⁻¹); C_e: the equilibrium concentration of BM (mg·L⁻¹); E: MB elimination rate (%).

Several kinetic models have been used to interpret the experimental data [9] . The two commonly used kinetics models are: the pseudo-first-order model (PPO), the pseudo-second-order model (PSO).

1) Pseudo-first order model

Lagergren proposed a pseudo-first order kinetic model expressed by the following relation [9] (Equation (4)):

$\frac{\text{d}{q}_t}{\text{d}t}=k_1\left(q_e-q_t\right)$ (4)

with k₁ the rate constant for pseudo first order kinetics (min⁻¹), t the contact time (min), q_t and q_e the adsorption capacities at time t (mg of adsorbat·g⁻¹ of adsorbent) and at equilibrium (mg of adsorbat·g⁻¹ of adsorbent), respectively. L’intégration de l’équation (Equation (4)) donne:

$\ln \left(q_e-q_t\right)=\ln q_e{k}_{1}t$ (5)

The value of k₁ is obtained from the slope of the linear plot of $\ln \left(q_e-q_t\right)=f\left(t\right)$.

2) Pseudo-second order model

An expression also very often used is that of pseudo-second order. It is represented by the formula (Equation (6)) [10] :

$\frac{\text{d}{q}_t}{\text{d}t}=k_2{\left(q_e-q_t\right)}^2$ (6)

With k₂ the rate constant for second order kinetics (g·mg⁻¹·min⁻¹), q_t and q_e the adsorption capacities at time t (mg of adsorbate·g⁻¹ of adsorbent) and at equilibrium (mg of adsorbate·g⁻¹ of adsorbent), respectively. The integration of equation (Equation (6)) leads to (Equation (7)):

$\frac{t}{q_t}=\frac{t}{q_e}+\frac{1}{k_2\cdot q_e^2}$ (7)

The pseudo-second order model is verified when the plot against t gives a linear relationship, with a slope equal to ( $\frac{1}{k_2\cdot q_e^2}$).

1) pH at point of zero charge (pH_PZC)

The pH_PZC curves were obtained according to our experimental procedure described previously. Figure 3 shows that the pHs of zero charge (pH_PZC) are 5.5 for AB clay and 6.08 for Aga clay. Both clays have a slightly acidic pH_PZC.

The curves made it possible to determine the pH_PZC of the different clays. This value is 5.5 for AB and 6.08 for Aga. At pH < pH_PZC, the surface of the clay acquires positive charges. There will, therefore, be an excess of H⁺ ions which will compete with the cations of the dye to access the clay adsorption sites [11] . Electrostatic repulsion causes a decrease in the amount of dye absorption. At pH > pH_PZC values, the surface of the clay will be negatively charged facilitating an association between the cationic dye [12] . The pH of the BM solution is 7 higher than the pH at the zero charge point of the clays, so the pH of the solution will be used for the rest of our work.

2) Chemical analysis

The chemical analysis of AB and Aga clays was carried out by ICP-AES Plasma Emission spectrometry after chemical solution using a microwave. The chemical composition expressed as oxide is given in Table 1 . The results show a predominance of silica (SiO₂) and alumina (Al₂O₃) in the two raw materials. Potassium, sodium and titanium oxides are also present in relatively small quantities. Calcium and manganese are found in trace amounts in clays. These high contents of SiO₂ and Al₂O₃ show that they are aluminosilicates [13] . The silica/alumina ratio (SiO₂/Al₂O₃) is 10.44 for Aga and 2.18 for AB. This ratio is high in both samples instead of 1.18 for pure kaolins [14] . These high values suggest the presence of a large amount of free silica and type 2/1 clay.

3) X-ray diffraction

The X-ray diffractograms of the two samples AB and Aga are shown in Figure 4 . All peaks have been assigned. The X-ray spectrum of Aga clay highlights the characteristic lines of kaolinite (12˚; 20˚; 25˚; 35˚), illite (37˚), quartz (20˚; 26˚; 45˚; 50˚) and goethite (41˚). The diffractogram of AB presents the lines of kaolinite (12˚; 25˚; 35˚), quartz (26˚; 36˚; 50˚) and illite (9˚; 19.70˚; 29.36˚) [15] .

The results obtained in XRD are consistent with those of the chemical analysis.

4) Mineralogical composition

The mineralogical composition of AB and Aga clays was estimated from the chemical analysis and the chemical composition of the mineralogical phases detected by X-ray diffraction. The results obtained are recorded in Table 2 .

5) Infrared spectroscopy

The infrared spectra of Aga clay and AB clay are shown in Figure 5 . In the 3000 - 4000 cm⁻¹ range, the following are observed:

For the Aga and AB samples, bands at (3700 cm⁻¹ and 3624 cm⁻¹) and (3695 cm⁻¹ and 3619 cm⁻¹) which are characteristic of kaolinite due to stretching vibrations of the hydroxyl groups [16] [17] . The band at 3700 cm⁻¹ in particular corresponds to the vibrations of the external hydroxyls of the kaolinite and that at 3624 cm⁻¹ to the vibrations of the internal hydroxyls located between the layers. This same band at 3624 cm⁻¹ is also characteristic of the stretching vibration of the OH bond of the mainly Al-OH-Al network of illite. These bands have been the subject of several studies. The band at 3695 cm⁻¹ is attributed to the hydroxyls at the edges of the sheet. The one located at 3619 cm⁻¹ is linked to the internal hydroxyls.

In the range 400 - 1800 cm⁻¹:

For Aga clay:

For AB clay:

1) Effect of hydrogen peroxide

All experiments were carried out under daylight irradiation. To study the action of hydrogen peroxide on the degradation of Methylene Blue, it is important to determine the optimal quantity of H₂O₂ necessary for maximum degradation of the dye. To do this, 20 mg of ferrous iron was fixed, then the quantity of H₂O₂ (30%) was varied in steps of 0.1 mL in an interval of 0.1 mL to 0.5 mL, for a solution of BM concentration 50 mg·L⁻¹. Figure 6 shows the different yields for each quantity of H₂O₂ after 2 hours.

The degradation rate of MB reached more than 90% after 120 minutes of reaction. The optimal quantity of H₂O₂ retained for the degradation of MB is 0.1 mL because it makes it possible to obtain a satisfactory rate for a small quantity of reagent.

2) Kinetics of MB degradation by hydrogen peroxide

The pseudo-first order and pseudo-second order kinetic laws were investigated in this study.

The plot of ln(q_e − q_t) as a function of time t for the degradation of MB by hydrogen peroxide gives linear shapes ( Figure 7 ).

The values of k₁ were calculated from the slopes of these lines, and q_e,_cal from the ordinate at the origin. The kinetic constant of the pseudo-first-order degradation reaction k₁, the quantity of MB adsorbed at equilibrium q_e, and the correlation coefficient $R_1^2$ are shown in Table 3 . The results in Table 3 show that the correlation coefficients obtained are 0.4308 for 0.2 mL and 0.1227 for the quantity of 0.5 mL. Despite the high correlation coefficient values, the experimental q_e do not agree with the q_e calculated from linearized forms of pseudo-first order kinetics. Therefore, the pseudo-first order kinetic model is not well suited to describe the degradation of MB by hydrogen peroxide.

The plot of versus time t for the degradation of MB by hydrogen peroxide gives the linear forms ( Figure 8 ). The values of k₂ and q_e can be determined from the slopes and y-intercepts of these lines.

The pseudo-second order degradation kinetic constant k₂, the quantity of MB adsorbed at equilibrium q_e, and the correlation coefficient $R_2^2$ are indicated in Table 4 . From these results, we note that the correlation coefficients $R_2^2$ are greater than the $R_1^2$ coefficients for all quantities of hydrogen peroxide. Furthermore, the experimental q_e values are much closer to the q_e values calculated from the linear forms of the pseudo-second order kinetics for all quantities of hydrogen peroxide. From these results, it appears that the pseudo-second order kinetic model gives a better description of the kinetics of the MB degradation reaction with hydrogen peroxide, unlike the pseudo-first order kinetic model.

3) Effect of iron (Fe²⁺)

Ferrous iron plays an important role in the efficiency of the photo-Fenton process. For the determination of its optimal value for the degradation of methylene blue is essential. For this, the optimal quantity of H₂O₂ is set at 0.1 ml and the mass of Iron is varied from 5 to 25 mg in steps of 5 mg. The results are presented in Figure 9 .

The degradation rate of Methylene Blue changes rapidly when the quantity of iron varies from 0 to 15 minutes. After 15 minutes, the yield is constant with an elimination rate of more than 90% for the different quantities. For the quantity of 10 mg, the elimination rate is BM is 92% after 120 minutes. This yield being the best, the quantity of iron of 10 mg will be retained as the optimal value for the rest of the work.

4) Kinetics of MB degradation by ferrous iron

The pseudo-first order and pseudo-second order kinetic laws were used in this study.

The plot of ln(q_e − q_t) versus time t for the degradation of MB by iron gives linear shapes ( Figure 10 ).

The values of k₁ were calculated from the slopes of these lines, and q_e,_cal from the ordinate at the origin. The pseudo-first order degradation kinetic constant k₁, the quantity of MB adsorbed at equilibrium q_e, and the correlation coefficient $R_1^2$ are shown in Table 5 .

Table 5. Pseudo-first order kinetic constants for the degradation of MB iron.

The results in Table 5 show that, the correlation coefficients obtained are 0.9402 for 10 mg; 0.9651 for 15 mg and 0.9837 for 25 mg. Despite the high correlation coefficient values, the experimental q_e do not agree with the q_e calculated from linearized forms of pseudo-first order kinetics. Therefore, the pseudo-first-order kinetic model is not well suited to describe the degradation of MB by iron.

The plot of versus time t for the degradation of MB by ferrous iron gives the linear shapes ( Figure 11 ). The values of k₂ and q_e can be determined from the slopes and y-intercepts of these lines. The kinetic constant of the pseudo-second order adsorption reaction k₂, the quantity of phenol adsorbed at equilibrium q_e, and the correlation coefficient $R_2^2$ are shown in Table 6 .

From these results, we notice that the $R_2^2$ correlation coefficients are greater than the $R_1^2$ coefficients for all pseudo-first order iron quantities. Furthermore, the experimental q_e values are much closer to the q_e values calculated from the linear forms of the pseudo-second order kinetics for all quantities of iron. From these results, it appears that the pseudo-second order kinetic model gives a better description of the kinetics of MB degradation by iron, unlike the pseudo-first order kinetic model.

5) Influence of clay

For the study of the influence of clay on the removal of BM the optimal values of peroxide and iron were used in the photo-Fenton process. Then the masses of the AB and Aga clays were varied from 0.05 g to 0.2 g in steps of 0.05 g over a period of 2 hours. The results obtained are presented in Figure 12 . For all quantities of clay used, the removal rate is high (ƞ > 90%). Among the clay masses used, for masses of 0.05 to 0.15 g, a rapid increase in the elimination rate is observed, from 15 to 40 minutes the yield is constant at 98%. Beyond 40 minutes a decrease in the elimination rate of methylene blue is observed. Thus, in the case of our study, the optimal mass of clay necessary for the effective elimination of methylene blue (50 mg/L) is 0.1 g.

6) Kinetics of MB degradation by the influence of clay

During this study, pseudo-first order and pseudo-second kinetic laws were used. The plot of ln(q_e − q_t) as a function of time t for the degradation of MB with the influence of clay gives linear shapes ( Figure 13 ). The values of k₁ were calculated from the slopes of these lines, and q_e,_cal from the ordinate at the origin.

The pseudo-first order degradation kinetic constant k₁, the quantity of MB adsorbed at equilibrium q_e, and the correlation coefficient $R_1^2$ are shown in Table 7 .

The results in Table 7 show that the correlation coefficients obtained are 0.6305 for 0.1 g of AB and greater than 0.95 for the quantity of 0.1 g of Aga clay. Despite the high correlation coefficient values, the experimental q_e does not agree with the q_e calculated from linearized forms of pseudo-first order kinetics. Therefore, the pseudo-first order kinetic model is not well suited to describe the degradation of MB with the influence of clay.

The plot of versus time t for the degradation of MB with the influence of clay gives linear shapes ( Figure 14 ).

The kinetic constant of the pseudo-second order adsorption reaction k₂, the quantity of phenol adsorbed at equilibrium q_e, and the correlation coefficient $R_2^2$ are shown in Table 8 .

The best model established for the study of degradation kinetics is chosen according to the correlation factor. The higher this factor is and closer to 1, the more favorable the model is for studying the degradation process [23] . In Table 7 , we notice that the $R_2^2$ correlation coefficients are greater than the $R_1^2$ coefficients for all quantities of clays. In addition, the experimental q_e values are much closer to the q_e values calculated from the linear forms of the pseudo-second order kinetics for all quantities of hydrogen peroxide. From these results, it appears that the pseudo-second order kinetic model gives a better description of the reaction kinetics of MB degradation with the influence of clay, unlike the pseudo-first order kinetic model.

The degradation of Methylene Blue by the photo-Fenton process was carried out in the presence of natural light. First, the volume of peroxide required for maximum dye removal was determined by setting the iron at 20 mg and varying the peroxide from 0.1 mL to 0.5 mL. Then, the optimal amount of iron was determined by setting the optimal volume of peroxide determined previously and varying the amounts of iron from 5 mg to 25 mg. Indeed, due to its low oxidizing power (1.8 V/ENH compared to 2.8 V/ENH for the hydroxyl radical), hydrogen peroxide is not capable of degrading MB alone [24] . On the other hand, when associated with ferric ions, it produces hydroxyl radicals which eliminate the dyes. Nicol showed that the use of 1 to 2 M hydrogen peroxide would improve the efficiency of dye processing, but only when H₂O₂ was coupled to ferric ions at 0.2 M [25] . The amount of 0.1 mL of hydrogen peroxide allowed the elimination of 90% of the dye, so it was chosen as the optimal value. Wei et al. showed that an excess amount of H₂O₂ could decrease the degradation rate due to competitive reactions between H₂O₂ and free radicals, which could reduce the efficiency of the process [26] . The optimal amount of ferrous iron for dye degradation is 10 mg for a yield of 92%. Beyond this value, the yield decreases. This observation could be explained by the formation of aggregates which would reduce the number of active sites [27] [28] . The decolorization rate of Methylene Blue when adding 0.1 g of AB and Aga clays made it possible to obtain a removal of 98.43% for AB and 98.13% for Aga. This could be explained by two phenomena that occur simultaneously during processing: the adsorption of the dye by the clay and the photo-Fenton process. Indeed, the mineralogical composition of the clays showed that these clays contain a significant proportion of clay minerals. These minerals have a strong adsorbent power due to the small size of their particles which gives them a large specific surface area. Indeed, Karim et al. in their study, showed that the adsorption process on raw MB clay is very rapid during the first 10 minutes of treatment and evolves slowly then stabilizes after the 40th minute [29] . Their results show that 95% of the adsorbed quantity was reached during the first ten minutes. Beyond this quantity of clay, we observe a reduction in the rate of elimination of the pollutant. This may be due to the difficulty of MB molecules in reaching the absorbing sites [30] . When the quantity of clay increases, we witness the phenomenon of agglomeration of clay particles. The solution becomes cloudier, which leads to a decrease in the effectiveness of the catalyst observed by the drop in the amount of light in the liquid, consequently the number of hydroxyl radicals responsible for degradation is reduced. Furthermore Badis and Manaa showed it for the elimination of two pollutants by photo-induction and the elimination of pharmaceutical compounds by photo-Fenton [2] [31] . This same observation was made by Saoudi and Hamouma [32] . According to them, this behavior can be explained:

It is also noted that the amount of adsorption is slightly higher for AB clay than Aga clay. The comparison of the two clays allows us to conclude that AB clay has greater adsorbent power than Aga clay. This is reflected in the characteristics of the clays contained in Table 2 . Indeed, AB clay is richer in clay minerals with a rate of 70.06% unlike Aga clay which has 21.44% clay minerals. In addition, goethite, which represents 3.12% of Aga clay and which is absent in AB clay, is a mineral which also has adsorption properties. It contributes to the high adsorption capacity of Aga clay.