Biosorption of phenol using modified barley husk studies on equilibrium isotherm kinetics and thermo

2.1. Materials

PHEN was procured from Sigma-Aldrich Co. De-ionized water was used to prepare all solutions for the experiments. The stock solution of PHEN (1000 mg/L) was prepared in 1000 ml volumetric flask at room temperature by suspending an accurate weight of PHEN in deionized water. The working solutions were prepared by diluting the already prepared stock solution to proper concentrations. The general properties and chemical structure of PHEN are shown in Table 1.

2.2. Adsorbent preparation

The barley husks (BH) were collected from farmlands in Tabriz city, Iran. It was sundried, crushed and sieved to particle sizes in the range of 1–2 mm. The biomass was treated with 0.1 M HCl for a period of 5 h, washed with distilled water and dried in the shade. The resultant biomass was used subsequently in the adsorption experiments.

The specific surface area of the MBH before use was determined by Brunauer-Emmett-Teller nitrogen adsorption (BET-N2) method based on nitrogen adsorption-desorption isotherms at 77 K. Scanning electron microscopy (SEM) of the MBH was carried out using a scanning microscope (Philips, Eindhoven) equipped with an energy-dispersive X-ray (EDX) (SERON). The FTIR spectra of MBH was recorded in the range of 400– 4000 cm-1 by an FTIR machine (Nicolet 5700 instrument, Thermo Corp., USA) to observe the chemical/functional groups which are accountable for the process of PHEN adsorption. To ascertain the point of zero charge (pHZPC) of the MBH and BH, 0.1 g of the MBH and 0.1g of BH were added separately to 50 ml of 0.2 mol/L sodium chloride (NaCl). The initial pH of the solution was adjusted to different pHs from 2 to 12 using 0.1 mol/L sodium hydroxide (NaOH) or 0.1 mol/L hydrogen chloride (HCl) solutions. Then the mixtures were shaken in the shaker at 200 rpm and temperature of 30±2°C for 24 h. After equilibrium, the mixtures were filtered and the zeta potentials were measured using a Zetamaster potentiometer. Table 1. Molecular structure and properties of PHEN. Parameter Molecular structure

CAS number Molecular formula Molecular weight Solubility in water Maximum wavelength (λmax)

2.3. Batch adsorption experiments

The most important effective variables on adsorption from the literature include pH, adsorbent dose, contact time, and adsorbate concentrations. Therefore, the effects of initial PHEN concentration (10-100 mg/L), MBH dosage (0.5-4 g/L), contact time (10-180 min) and pH (3-11) were investigated. The batch experiments were carried out in 200 mL Erlenmeyer flasks. In each adsorption experiment, PHEN solution of a certain concentration was added into the flask. The desired pH was set and then a specific dose of MBH adsorbent was added. Then, the mixture was mixed with a shaker at 120 rpm for a certain time. Then, the mixture was centrifuged at 3000 rpm for 10 min. The amount of PHEN adsorbed was calculated according to the following Equation (1) [24]:

Where qe is the amount of PHEN adsorbed (mg/g), C0 and Ce are the initial and equilibrium concentrations of liquid phase (mg/L), respectively, V is the volume of the solution (L), and m is the mass of the MBH used. All retention measurements by high-performance liquid chromatography (HPLC) were made with a PerkinElmer Series 200 HPLC apparatus (Shelton, CT, USA) equipped with a column C18 and UV detector set at 280 nm. An injector with 40 µL sample loop was used for all sample injections.

3.1. MBH Characterization Results

The definite surface area is linked to the number of active adsorption sites present on an adsorbent material. The adsorption rate increases with the porosity and specific surface area of an adsorbent. The surface area of MBH was revealed as 176.2 m2/g, which indicates that the MBH has relatively good capability to eliminate PHEN. Scanning electron microscopy (SEM) images were employed to analyze the surface structure of MBH before and after PHEN adsorption (Figures 1a and 1b). The MBH was found to possess a heterogeneous rough surface structure as well as deep pores which is relatively organized with different sizes and shapes. As seen in Figure 1b, the pores of the MBH were filled with PHEN molecules after adsorption. The elemental composition analysis of the adsorbent (MBH) performed using the Energy Dispersive Xray Microanalysis (EDX) is shown in Figure 2. Elemental analysis of MBH shows a high yield of carbon (57.57%), oxygen (35.7%), hydrogen (4.40%) and small amounts of calcium (1.2%), potassium (0.75%), sodium (0.77%), phosphorus (0.017%), chlorine (0.24%), and sulphur (0.11%).

Figure 1. The SEM image of MBH before and after PHEN removal.

In order to understand the interaction between the functional groups existing in MBH and PHEN cations, both MBH before and after use was examined by means of the Fourier transform infrared (FT-IR) spectroscopy. As shown in Figure 3, MBH before and after use showed a similar pattern and the same number of peaks were observed. The peaks at 3410-3430 cm−1 correspond to the O–H stretch due to inter- and intramolecular hydrogen bonding of polymeric compounds for instance alcohols, PHENs, and carboxylic acids, as in pectin, cellulose, and lignin [26]. Thus, showing the presence of free hydroxyl groups on the MPH surface. The peaks at 2928 cm −1 are attributed to the symmetric and asymmetric C–H stretching of aliphatic acids [27]. The peak observed at 1744 cm−1 corresponds to –COOH, –COOCH3 stretching of carboxyl groups and may be assigned to carboxylic acids or their esters. The peaks at 1638 and 1413-1418 cm−1 are due to the asymmetric and symmetric stretching vibrations of C=O in ionic carboxylic groups (– COO−). Aliphatic acid group vibration at 1262 cm−1 may be assigned to C=O and –OH stretch of carboxylic acids and PHEN.

Figure 3. The FT-IR spectra of MBH before and after PHEN removal.

3.2. Effect of pH and determination of point of zero charge

The point of zero charge (pHZPC) is a concept relating to the phenomenon of adsorption. pHZPC can be defined as the pH value at which a solid submerged in an electrolyte exhibits zero net electrical charge on the surface. As observed in Figure 4, the pHZPC values for BH and MBH are approximately 5.9 and 5.1, respectively. This revealed that the activation by HCl increases the acidic properties of MBH as a result of the hydrogen bonding present. The surface charge becomes positive at pH<pHZPC and negative at pH>pHZPC [28].

PHEN removal on MBH decreased significantly by increasing the solution pH from 3 to 11 (Figure 5). One of the most important factors affecting the biosorption of PHEN on MBH was the acidity of the solution since the pH of the solution impacts not only the surface charge of the biosorbent and the dissociation of functional groups of the active sites on the biosorbent surface, but as well as the aqueous chemistry of the PHEN [29]. At this pH solution = pH ZPC, the net surface charge density of MBH is near to zero, and this enhances the ᴫ-ᴫ dispersion interaction. Consequently, when the solution pH>pHZPC, the surface charge density of MBH is negative and PHEN is deprotonated. As a result, the repulsive electrostatic interaction is intensified [30]. This is the main mechanism that contributed to the weak adsorption efficiency onto MBH at pH>pHZPC. These results were consistent with the previous study on the adsorptive removal of PHEN onto Azolla [31]. Kadhim et al. [32] examined the effect of initial pH on the removal of PHEN from aqueous solution onto rice husk. As pH increased from 3 to 11, the adsorption capacity decreased from 15.2 to 5 mg/g at 100 mg/L concentration; maximum uptake was observed at pH 3.0.

Figure 4. Determination of the point of zero charges of BH and MBH.

Figure 5. Influence of pH on PHEN removal (contact time = 90 min, dosage = 3 g/L, PHEN concentration C0: 100 mg/L, temperature = 30 ± 2 ºC).

3.3. Effect of adsorbent dosage

The effect of MBH dose on PHEN removal was studied to determine the optimum MBH dosage. As shown in Figure 6, the efficiency of PHEN removal improved with increasing the MBH dosage from 0.5 to 4 g/L; the adsorption capacity decreased with increasing dosage. Lower biosorption capacity at a higher dosage of biosorbent is probably as a result of the decrease of the surface area of the biosorbent by the overlapping or aggregation during the process of sorption [33]. However, a higher dosage of MBH led to the greater availability of active sites for PHEN, resulting in higher PHEN removal [34]. Similarly, Diyanati et al. [31] investigated the influence of dose on the PHEN removal from aqueous solution; the removal efficiency increased from 27.65% to 81.27%, while the qe decreased from 55.3 to 20.31 mg/g by increasing the biosorbent dose from 0.5 to 4 g/L.

Adsorbent dose (g/L) Figure 6. Influence of MBH dose on PHEN removal (contact time = 90 min, pH = 3, PHEN concentration C0 = 100 mg/L, temperature = 30 ± 2 ºC).

3.4. Effect of contact time and initial PHEN concentration

Figure 7 indicates that PHEN removal improved as the contact time was increased from 10 to 180 min. Equilibrium was met at 120 min. Therefore, the optimum contact time for PHEN removal on MBH is considered to be 120 min. At the initial times, PHEN removal was rapid. The gradient of adsorption declined with time which may be due to declining concentration of PHEN and decreasing active adsorption sites on the adsorbent surface area [27]. In the early stages of adsorption, there were lots of empty spaces and they were occupied with time by PHEN molecules [35]. Similarly, Dursun et al. [36] investigated the effect of agitation time on the removal of PHEN from aqueous solution on carbonized beet pulp. The results showed that the adsorption capacity (mg/g) increased with increasing time. The equilibrium was established after 120 min of contact time. PHEN adsorption on MBH decreased with increasing PHEN concentration from 10 to 100 mg/L. The results are shown in Figure 7. As indicated, 93.95% removal was obtained at 10 mg/L concentration; removal of PHEN decreased with increasing initial concentration and reached 72% at the concentration of 200 mg/L. This probably occurred due to the increasing surface charge on the adsorbent; the adsorption sites of the top surfaces of the adsorbent were saturated and the

removal efficiency decreased [37]. Dakhil [38] investigated the effect of concentration on the removal of PHEN from aqueous solution. The removal percent decreased from 90% to 80% with increasing PHEN concentration from 10 to 50 mg/L.

Time (min) Figure 7. Effect of contact time and initial PHEN concentration (pH = 3, dose = 3 g/L, temperature = 30 ± 2 ºC)

3.5. Adsorption isotherms

The equilibrium data is important in the design of adsorption systems [39]. Although several isotherm equations are available, four important isotherms (Freundlich, Langmuir, Temkin, and Dubinin-Radushkevich were selected for this study. The Langmuir isotherm is expressed as Equation (2) [40]: 𝑪𝒆

Where qe (mg/g) and Ce (mg/L) are the equilibrium concentration of PHEN in the solid and liquid phases, respectively; q m (mg/g) is the maximum theoretical biosorption capacity; KL (L/mg) is a measure of biosorption energy, showing the affinity between biosorbent and sorbate. A straight-line plot of Ce/qe versus Ce gives an intercept of 1/qm.KL and slope of 1/qm. The essential feature of the Langmuir isotherm can be expressed in terms of a dimensionless constant called separation factor (RL), also known as equilibrium parameter which is defined by Equation (3) [39, 41]: RL =

Where C0 (mg/L) is the initial PHEN concentration. The value of R L specifies the shape of the isotherms to be either unfavorable (RL>1), linear (RL= 1), favorable (0<RL>1) or irreversible (RL=0). The Freundlich isotherm model can be expressed logarithmically as in Equation (4) [42]: Log qe =

KF [(mg/g)/(mg/l)1/n] is the Freundlich constant related to the sorption capacity of the adsorbent and 1/n is the Freundlich constant associated to the energy heterogeneity of the system and the size of the adsorbed molecule. The values of KF and n can be estimated from the linear plot of Log qe versus Log Ce.

Temkin isotherm proposes an equal distribution of binding energies over the number of the exchanging sites on the surface. The linear form of Temkin isotherm can be written as [43]: qe = B Ln A + B Ln Ce

Where B=RT/b, T is the absolute solution temperature and R is the universal gas constant (8.314 J/mol.K). A is the equilibrium binding constant and B is related to the heat of sorption. The values of B and A can be determined from the plot of q e versus Ln Ce. The Dubinin-Radushkevich (D-R) model can be expressed independently of solution temperature using the adsorption potential (ε) (Equation 6) [44]: 𝟏

The D-R isotherm assumes a Gaussian-type distribution for the characteristic curve and the model can be described by Equation (7) [45, 46]: Ln qe = Ln qm – βε2

Where qm is the D-R constant (mol/g) and B gives the mean sorption free energy, E (kJ/mol) per molecule of sorbate at the moment of its transfer to the solid surface from the solution and can be computed using Equation (8) [45]: E = (−2K)−1/2

Plotting Ln qe versus ε2, using Eq. (8), results in a straight-line of slope, β, and intercept, Ln (qm). Figures 8a-d show the Freundlich, Langmuir, Temkin, and Dubinin-Radushkevich plots, with the calculated parameters are summarized in Table 2. As seen in Table 2, D-R isotherm favors the PHEN uptake by MBH (R2=0.971) than the other models. The values of R2 reveal that the experimental data fitted into the D-R isotherm model. It is reported that when the value of E is below 8 kJ/mol, the process of adsorption can be considered as physisorption [47]. In contrast, if the value of E is located within the range of 8-16 kJ/mol, it is a chemisorption process. From Table 2, it can be seen that the obtained value of mean free energy, E, is 0.267 kJ/mol. Based on these, it can thus be concluded that the effect of physical adsorption will play a dominating role in the process of PHEN adsorption on the MBH. Furthermore, Freundlich isotherm constant can be used to explore the favorability of the process of adsorption. The process is said to be favorable when the value of n satisfies the condition 1<n<10, otherwise it is unfavorable. The 1/𝑛 value of 0.667 (Table 2) for the adsorption of PHEN on MBH is found to be within 0 and 1 indicating desirable adsorption process and heterogeneity of the process [48]. Table 2. Adsorption isotherm constants for the adsorption of PHEN onto MBH. qm 111.1

Figure 8. Linear fitting plots of (a) Langmuir, (b) Freundlich, (c) Temkin and (d) D-R isotherm models.

3.6. Adsorption kinetics

In order to examine the mechanism in control of the adsorption process, pseudo-second-order, pseudo-first-order, and intraparticle diffusion kinetic equations were used to examine the kinetic data. The pseudo-first-order rate equation is expressed as Equation (9) [49, 50]: Log (qe − qt) = Log qe −

Where qe and qt are the amounts of PHEN adsorbed (mg/g) at equilibrium and at time t (min), respectively; K1 is the rate constant of adsorption (min−1). K1 values were estimated from Log (qe - qt) versus t plots at various concentrations. The pseudo-second-order model equation is expressed as Equation (10) [51, 52]: 𝒕 𝒒

Where K2 is the rate constant of pseudo-second-order (g mg −1min ); q and qe are the amounts of solute adsorbed at any time and at equilibrium (mg/g), respectively. The two models were able to define the adsorption data. The agreement between the calculated data and experimental values is expressed by the correlation coefficient (R2). The results are presented in Table 3 and Figures 9-10. Kinetic data of PHEN onto MBH fit the pseudo-second-order model (R2>0.992) than the pseudo-first-order model (R2>0.971) proposing a chemisorption process [47]. The intraparticle diffusion rate equation is expressed as Equation (11) [53]:

Where C is the intercept and Ki is the intraparticle diffusion rate constant (mg/gmin ), which can be evaluated from the slope of the linear plot of qt versus t1/2. Intra-particle diffusion is the rate-limiting step when the plot passes through the origin (C=0). If the value of C is higher than zero, the difference in the rate of mass transfer during the early and final stages occurred. It is observed from Figure 11 that the plot did not pass through the origin (the value of C>0), hence the PHEN adsorption mechanism is probably a combination of boundary layer and pore diffusion. Both contributed to the rate-determining step. Therefore, the adsorption of PHEN ions on the MBH surface is complex, involving more than one mechanism as shown in Figure 11. Table 3. Calculated parameters of kinetics for PHEN adsorption on MBH. Co (mg/L)

K1 qe (mg/g) R2 (min−1) 0.046 3.11 0.971 0.043 7.66 0.988 0.039 17.45 0.983 0.042 25.75 0.976 0.049 32.19 0.992

K2 (g qe (mg/g) mg−1min−1) 0.025 3.401 0.009 8.403 0.004 16.39 0.002 23.80 0.002 31.25

Ki C (mg/gmin1/2) 0.155 1.406 0.403 30.14 0.824 5.556 1.256 7.178 1.735 7.774

t (min) Figure 9. Pseudo-first-order kinetic model for PHEN adsorption on MBH.

t (min) Figure 10. Pseudo-second-order kinetic model for PHEN adsorption on MBH. 50 mg/L

t1/2 Figure 11. Intraparticle diffusion plot of PHEN adsorption at various concentrations.

3.7. Thermodynamic analysis

To evaluate the effect of temperature on the PHEN adsorption on MBH, the free energy change (ΔGo), enthalpy change (ΔHo), and entropy change (ΔSo) were determined using the following equations [54-56]: Ln KL = -

Where KL is the Langmuir equilibrium constant (L/mol); R is the gas constant (8.314 J/mol K); T is the solution temperature (K). Considering the connection between ΔGo and KL, ΔHo and ΔSo were evaluated from the slope and intercept of the Van’t Hoff plots of Ln (KL) versus 1/T (Figure 12). Table 4 presents the parameters of thermodynamic at different temperatures (273, 283, 293, 303, 313 and 323 K). The

negative values of ΔGo confirm the feasibility and spontaneity of the process [57]. ΔGo values were found to decrease from −0.056 to −3.73 kJ/mol. The decrease in the ΔGo negativity value with increasing temperature indicates that the PHEN removal on MBH becomes more favorable at greater temperatures. The ΔHo and ΔSo values calculated from the plot of Ln (K L) versus 1/T are given as 23.88 kJ/mol and 0.087 kJ/mol.K, respectively. Zheng et al. [58] suggested that the ΔHo of physisorption is smaller than 40 kJ/mol. The value of ΔHo (23.88 kJ/mol) suggests that the process is a physisorption process. The ΔHo value was positive, showing that the reaction was endothermic. The positive value of ΔSo reflects the affinity of the MBH for PHEN; this also suggests some structural changes in PHEN and MBH [57]. Table 4. Calculated parameters of thermodynamics for PHEN adsorption on MBH.

Temperature (°𝐊) 273 283 293 303 313 323 1,8 1,6 1,4 1,2 1 0,8 0,6 0,4 0,2 0 -0,20,003

3.8. Comparison of MBH with other adsorbents for PHEN removal

See Table 3 by Igwegbe et al. [57] which shows the comparison of gypsum with other materials for PHEN reduction. For this study, maximum PHEN removal of 93.95% was obtained using MBH at pH 3, MBH dose: 3 g/L, PHEN concentration: 10 mg/L, and contact time: 180 min at a constant temperature of 30 ± 2 ºC. This indicates that MBH can also be considered as a good biosorbent for the removal of phenol from water.