Sage Journals: Discover world-class research

Abstract

The adsorption of phosphate ion onto natural reed (Arundo donax) was studied in this work. The effect of phosphate initial concentration, adsorbent dose, pH, temperature, and salt addition on adsorption uptake was investigated. The results showed that the adsorption uptake is directly proportional to the phosphate ion initial concentration and inversely proportional to the adsorbent's dose and temperature. A maximum adsorption capacity of 16.2 mg/g was observed at neutral pH. The addition of sodium and potassium chlorides has decreased the adsorption uptake. The adsorption isotherms agree better with the Langmuir model. The negative values of (ΔG) and (ΔH) obtained from the thermodynamic study, indicted that the adsorption process is spontaneous and exothermic. The experimental adsorption data were analyzed using three kinetic models: pseudo-first order, pseudo-second order, and intra-particle diffusion model. The pseudo-second-order model presented the best fit with a determination coefficient (R²) higher than 0.99 and a minimum normalized standard deviation.

Keywords

Phosphate ions adsorption natural reeds kinetics thermodynamics equilibrium

Introduction

The intensive use of fertilizers, pesticides, as well as the discharge of incompletely treated wastewater has become a great environmental concern. Nutrients contained in the aforementioned sources, namely phosphorous (P) and nitrogen loaded into the aquatic environment. This accumulation poses an ecological problem, as phosphorous is a key factor that stimulates significant water eutrophication (Hong-Bing et al., 2012). A variety of phosphate ion ( ${PO}_{4}^{- 3}$ ) removal methods from water are currently available and cited in literature: among these methods are membrane processes (Dietze et al., 2003), chemical separation, biological treatment, coagulation, and filtration (Morse et al., 1998; Rittmann et al., 2011). The previously mentioned methods are characterized of being not cost effective and not satisfactorily efficient in the removal of trace concentrations of phosphate ions. Furthermore, membranes have moderately high maintenance cost and they are susceptible to fouling (Chmielewská et al., 2013). Several studies in literature pointed out that adsorption is a simple and economical method for phosphate removal (Chmielewská et al., 2013; Fazhi et al., 2014). The use of different adsorbents for phosphate removal from water such as granulated ferric hydroxide (GFH) and activated aluminum oxide (AA) was reported (Genz et al., 2004). One drawback for the large-scale use of these commercial adsorbents is their high cost. Therefore, the key for satisfactorily efficient adsorption is looking for novel natural, inexpensive, and locally available adsorbents. Additionally, it is worth to highlight that investigations showed the low efficiency of some natural adsorbents such as sands, limestone, fly ash, or gypsum to remove phosphate ion from water. Not all natural adsorbents are convenient for phosphate removal from water. Some of them show low adsorption capacities such as sands, limestone; fly ash, gypsum, or some have toxic impacts on the aquatic life (Hong-Bing et al., 2012). Elzinga and Sparks (2007) have tested the adsorption of phosphate onto hematite and they reported the existence of phosphate complexes on the hematite surface. In addition, phosphate adsorption on hematite was characterized as a function of pH (3.5–8.9) and phosphate concentration (5–500 µM) by in situ spectroscopy. Khan et al. (2010) suggested the use of low cost phosphate ion adsorbent such as calcareous and alkaline soil. The adsorption data agree with the Freundlich isotherm. Karageorgiou et al. (2007) showed that phosphate species seem to be efficiently removed from solutions using calcite as natural adsorbent. The results of this study showed that the pH plays an important role in the removal of phosphate ion and the removal was more efficient under basic pH. Karaca et al. (2004) showed that the adsorption of phosphate ion from aqueous solution on dolomite involves both physical and chemisorption.

Many adsorbents prepared from agricultural residues have been tested and investigated such as apple pomace (Robinson et al., 2002), coconut husk (Manju et al., 1998), and sawdust (Ajmal et al., 1998). However, the majority of tested adsorbents are not good options for large scale use due to their low phosphate removal capacity and high cost. Nowack and Stone (2006) reported a maximum phosphate adsorption capacity of 5.7 mg ${PO}_{4}^{- 3}$ /g of Geothite. Also, Mallikarjun and Mise (2012) succeeded to reach phosphate adsorption capacity of 4.5 mg/g on black cotton soil as well as red soil.

The objective of the current work is to study the adsorption of phosphate ion onto Arundo donax reeds as natural, available, and inexpensive adsorbent. Arundo donax reeds grow and spread particularly in the vicinity of water. In Jordan, it grows in a natural ecosystem in Wadi Shueib region located in Al-Balqa' Directorate (Ammari, 2014). Arundo donax reeds suggest a broad potential application as they are abundant in nature (eco-biosorbents), they need low cost operation, and they have high capacity of phosphate removal. Therefore, it has merits over other biosorbents. The removal of phosphate ion by Arundo donax reeds was investigated at different operating conditions namely the pH, the temperature, the contact time, salt addition, and the initial adsorbent concentration. Equilibrium, kinetic and thermodynamic studies were also performed. The results of the current paper would help in developing or designing cost effective and environmentally compatible technique to remove phosphate from water bodies.

Experimental analysis

Preparation of Arundo donax reeds

The Arundo donax reed material used in this study as an adsorbent was obtained from Jordan. Prior to their use as adsorbents, they were prepared according to the following procedure:

The reeds were collected, dried naturally under the sunshine, and then crushed into small pieces. They were ground in order to get smaller sizes.

Afterward, the reeds were sieved in order to obtain particles with diameters in the range of 0.15 to 0.85 mm.

The particles with the appropriate size were washed with distilled water, and then they were dried at 108℃ for 8 h.

The phosphate aqueous concentration (20, 40, 60, and 100 mg/L) was determined by the Vanadomolybdo-phosphoric acid colorimetric method, using a Hach DR/4000 spectrophotometer (APHA, 1989). The phosphate solution containing reeds was shaken with a shaking water bath. The pH of all samples was measured using a pH meter.

Phosphate removal

Effect of adsorbent and adsorbate dosage on phosphate removal

A stock solution of 1000 mg/L of phosphate was prepared by dissolving potassium dihydrogen phosphate (KH₂PO₄) in distilled water. The stock solution of the adsorbate was diluted to an initial concentration of 100 mg/L. The pH of the standard solution was adjusted to 6.5. The following masses: 0.5, 1, 1.5, and 2 g of adsorbent were added to 20 mL of the standard in 100 mL glass bottles, and the mixtures were shaken using a rotary shaker at 100 rpm and allowed to equilibrate over 24 h at 25℃. However, the minimum time needed to reach equilibrium is about 1 h. The bottles were then removed from the shaker and sent for analysis.

The previous procedure was repeated using different initial phosphate concentrations (20, 40, and 60 mg/L), respectively. Each run was made in duplicate.

Effect of pH on phosphate removal

The effect of pH on phosphate removal was investigated in a range from 2 to 10. The pH was varied using hydrochloric acid (HCl) and sodium hydroxide (NaOH). The experiment was run at 25℃ using an initial phosphate concentration of 100 mg/L with 5 mg/mL of reeds. Mixtures were shaken at 100 rpm and allowed to equilibrate over a period of 24 h. Each run was made in triplicate.

Effect of temperature on phosphate removal

Batch adsorption experiments were carried out to study the effect of temperature on phosphate ion removal. The uptake of phosphate ion was measured using 5 mg/mL of reeds and at different initial phosphate concentrations (20, 40, 60, and 100 mg/L). These experiments were performed at various temperatures namely 25℃, 35℃, and 45℃.

Effect of salt addition on phosphate removal

This part was devoted to examine the effect of the presence of different salts on phosphate removal. Specifically, the effect of sodium chloride and potassium chloride was tested using salt concentrations ranging from 0 to 0.5 M. Each run was made in triplicate.

Kinetic study

Twenty milliliters of the stock solution (diluted to initial concentration of 100 mg/L) was placed into glass bottles containing 0.1 g of the adsorbent (reeds). The bottles were placed and agitated using isothermal bath shaker adjusted to the required value. Samples were taken out at different time intervals (15, 30, 45, 60, 180, and 300 min). Each run was made in triplicate.

The adsorbent was separated from the samples by filtration using a 0.1 µm filter paper.

The amount of phosphate ion adsorbed at equilibrium (q_e, mg/g) was calculated using the following equation:

q_{e} (mg / g) = \frac{(C_{0} - C_{e}) \times V_{s}}{m}

(1) where C₀ and C_e (mg/L) are the initial and the equilibrium concentrations of phosphate ion in solution, respectively; Vs (L) is the volume of phosphate solution, and m (g) is the weight of sorbent.

Three models (pseudo-first order, pseudo-second order, and the intra-particle diffusion model) were investigated in this part in an attempt to identify the adsorption rate and mechanism.

The relevant equations are (Ho, 2004; Hong-Bing et al., 2012):

Pseudo-first order model:

\log (q_{e} - q_{t}) = \log q e - k_{1} t

(2) where q_e is the amount of phosphate ion adsorbed onto the reeds at equilibrium in (mg phosphate/g of reeds), q_t is the amount of phosphate adsorbed onto the reeds at any time in (mg phosphate/g of reeds), and k₁: kinetics rate constant of the pseudo-first-order model (min⁻¹).

Pseudo-second-order model (Ho, 2004; Hong-Bing et al., 2012):

\frac{t}{q_{t}} = \frac{1}{k_{2} {q_{e}}^{2}} + \frac{t}{q_{e}}

(3) where k₂ is the kinetics rate constant of the pseudo-second-order model (g.mg⁻¹.min⁻¹).

Intra-particle diffusion model

q_{t} = k_{p} t^{0.5} + C

(4) where k_p is the intra-particle rate constant (g/mg.min^0.5) and C is a constant.

Validity and significance of each model can be verified using a normalized standard deviation Δq%, which can be calculated as follows:

Δ q_{t} (%) = 100 \sqrt{\frac{Σ {[(q_{t, \exp} - q_{t, cal}) / / q_{t, \exp}]}^{2}}{N - 1}}

(5) where N is the number of data points, q_t,exp and q_t,cal are the experimental and calculated adsorption uptake, respectively in (mg/g).

Adsorption isotherms

For the adsorption isotherm study, 0.1 g of reeds was added into glass bottles containing 20 ml of different phosphate initial concentrations (20, 40, 60, and 100 mg/L) at three temperatures (25℃, 35℃, and 45℃). The bottle-point technique was used to conduct the adsorption equilibrium test. The mixture was allowed to equilibrate over 5 h in a shaking water bath.

Results and discussion

Effect of adsorbent and adsorbate concentration

The effect of adsorbent concentration on phosphate adsorption onto reeds was investigated. Batch experiments were carried out with four different adsorbent amounts (5, 10, 15, and 20 mg/mL) at a fixed temperature, pH, contact time, and initial adsorbate concentration. This procedure was repeated with different initial phosphate concentrations (20, 40, 60, and 100 mg/L).

Figure 1 depicts the variation of phosphate ion uptake with the concentration of adsorbent in addition to the initial concentration of phosphate ion. It can be clearly seen that the phosphate ion removal process is monotonically increasing with phosphate ion initial concentration and monotonically decreasing with the amount of adsorbent. A maximum phosphate ion uptake of 16.2 mg/g was obtained at 100 mg/L and 5 mg/mL as phosphate initial concentration and reeds concentration, respectively. The decrease in the amount of phosphate adsorbed with increasing the adsorbent mass is due to the concentration gradient between the phosphate concentration in the bulk solution and the liquid phosphate concentration on the surface of the adsorbent. Therefore, with increasing the mass of adsorbent, the mass of phosphate adsorbed per unit weight of adsorbent get reduced (Ammari, 2014). The same concept is valid when increasing the initial concentration of phosphate in solution as this means altering the equilibrium and increasing the driving force of adsorption, which is the concentration gradient between the solution and the surface of adsorbent.

Figure 1.

Effect of reeds concentration (ms) and initial phosphate concentration on phosphate (PO4⁻³) uptake (q) (rpm: 100; pH: 6.5; T: 25℃).

Effect of contact time

Batch assays were conducted to examine the phosphate uptake onto reeds during a contact time of 300 min using different initial phosphate ion concentrations (20, 40, 60, and 100 mg/L). The time course of the phosphate uptake in these batch tests is presented in Figure 2. As it is illustrated in Figure 2, for initial phosphate concentrations of 20, 40, and 60 mg/L, most of the phosphate ion is adsorbed during the first 15 min. After that, the rate of phosphate uptake started to stabilize with no further uptake noticed after 45 min. At this point, it was presumed that equilibrium was reached. At the initial phosphate ion concentration of 100 mg/L, the phosphate ion removal was rapid during the first 30 min, then proceeded at a slower rate until it reached equilibrium after 60 min. Hence, it was noticed that as the initial concentration increased, longer time may be required to attain equilibrium. This result is consistent with what was previously reported in literature (Kumar et al., 2010).

Figure 2.

The kinetic curves of the amount of phosphate adsorbed onto reeds at different initial phosphate concentrations (rpm: 100; pH: 6.5; T: 25℃, ms: 5 mg/mL).

The initial rapid adsorption could be attributed to the high affinity sites and adsorption. Then, as the reaction proceeds, the adsorption of phosphate ion increases the overall surface negative charge, causing the adsorption rate to slow down. This is also consistent with what was previously reported in literature (Fazhi et al., 2014; Luengo et al., 2006; Xiaoming et al., 2013; Yue et al., 2010). In these studies, the time required to reach equilibrium for phosphate ion adsorption on modified giant reed, Geothite and Hematite was about 25 min.

Effect of initial pH

The significant impact of the aqueous solution pH on the sorption process of phosphate ion was addressed in several studies (Biswas et al., 2008; Krishnan and Haridas, 2008). pH has an important effect in determining the phosphate ionic form. Hence, the effect of initial pH on phosphate ion adsorption was studied in this work. It was evident from Figure 3 that the pH had a significant influence on the amount of phosphates adsorbed onto reeds. In the pH range from 2 to 7, the removal efficiency of phosphate from the solution increased until it reached a peak (16.2 mg/g) at the neutral pH value and then descended as the solution had become more alkaline. These findings are similar to the results reported by Yue et al. (2010) and Xu et al. (2011). In these studies, it was shown that the maximum effective removal of phosphate ions by cotton stalk/wheat stalk and modified giant reed occurred in the pH range of 4–9. These results were explained by the presence of dominant dihydrogen and hydrogen phosphate $H_{2} {PO}_{4}^{-}$ and ${HPO}_{4}^{- 2}$ , ions that had a strong affinity for the adsorption sites when pH ranges from 2 to 7. It was reported by Razaq (1989) that the pK₁ and pK₂ of H₃PO₄ occur at pH values of 2 and 7, respectively. Consequently, the observed peak at pH of 7 in our case indicates that ${HPO}_{4}^{- 2}$ has a far greater affinity to be adsorbed on the surface than $H_{2} {PO}_{4}^{-}$ . However, as the pH increases, the concentration of ${HPO}_{4}^{- 2}$ ion increases making the surface more negative resulting in higher electrostatic repulsion. This repulsion becomes more significant at pH higher than 7, which makes adsorption potential of ${HPO}_{4}^{- 2}$ much weaker (Razaq, 1989).

Figure 3.

Adsorption of phosphate onto reeds as a function of pH (rpm: 100; T: 25℃, Co: 100 mg/L; ms: 5 mg/mL).

Effect of NaCl and KCl addition

This part is devoted to investigate the effect of interfering ions on the adsorption of phosphate ion onto reeds. Different concentrations of NaCl and KCl salts were added to the phosphate-containing solution. It can be noticed from Figure 4 that both salts adversely affected the adsorption of phosphate. The latter observation could be ascribed to a competitive adsorption with the co-existing chloride ions especially at the high salt concentration used. Fazhi et al. (2014) demonstrated that Cl⁻ ions had negative effect on PO4⁻³ adsorption on magnesium oxide nanoflake-modified diatomite adsorbent. The phosphate uptake from a solution containing 5 mg/L PO4⁻³ was decreased when chloride-containing salt concentration was increased from 1 to 10 mM. However, the effect of chloride ions is different when adding it as NaCl or KCl salts. This could be supported by the work of Razaq (1989) in which he mentioned that phosphate adsorption varies with the nature of the ionic composition and the ionic strength of the supporting matrix solution. In other words, it depends on the cations added to phosphate containing solutions. It was reported also that the effect of different salts on adsorption depends strongly on pH value. In the later work, phosphate adsorption capacity was different in the presence of potassium and sodium cations at pH of 5.5, which is in accordance with findings in our work. However, this effect of cations on PO4⁻³ adsorption cannot be generalized as it changes with pH (Razaq, 1989).

Figure 4.

Effect of salt addition on the uptake of phosphate by reeds (rpm: 100; pH: 6.5; T: 25℃; Co: 100 mg/L; ms: 5 mg/mL).

Effect of temperature

The effect of temperature on phosphate uptake by reeds was investigated using various phosphate initial concentrations. It is apparent from Figure 5 that the increase in temperature has a negative effect on the adsorption of phosphate, which leads to the conclusion that the sorption process of phosphate is of exothermic nature. This result was supported by a study conducted by Yue et al. (2010), using modified giant reed as a bio sorbent. They reported a similar temperature pattern as the phosphate adsorption capacity has decreased as the temperature increased from 20 to 60℃. This result was explained by the exothermic nature of the adsorption process.

Figure 5.

Effect of solution temperature on the uptake of phosphate by reeds (rpm: 100; pH: 6.5; T: 25℃; ms: 5 mg/mL).

Thermodynamic studies

The thermodynamic parameters, namely the Gibbs free energy (ΔG), the enthalpy (ΔH) of phosphate ion adsorption and the adsorption entropy (ΔS) onto reeds were determined from the Langmuir isotherms at different temperature using Van't Hoff equation:

G = - RT \ln K_{D}

(6)

\ln K_{D} = \frac{- Δ H}{RT} + \frac{Δ S}{R}

(7) where R is the universal gas constant (8.314 kJ/kmol⋅k), T is temperature in kelvin (K), and K_D is the adsorption equilibrium constant. The values of the enthalpy (ΔH) and entropy (ΔS) were obtained graphically from the slope and intercept of van't Hoff plots of ln K_D versus 1/T (Min et al., 2001). The (ΔG), (ΔH), and (ΔS) values are listed in Table 1. The negative value of the enthalpy and the free energy indicates the exothermic and spontaneous nature of the process, respectively, which coincides with the previous findings in this paper. In view of the fact that the value of ΔH was found to be smaller than an absolute value of 62.7 kJ/mol, it is suggested that the removal process was due to the physical adsorption mechanism (Fogler, 1992; Yue et al., 2010). In addition, the decrease in ΔG with increasing temperature suggested that the adsorption process is less favorable at higher temperatures. The negative value of ΔS indicates a decrease in the degree of freedom at the solid liquid interface throughout the adsorption process (Kennedy et al., 2007).

Table 1.

Thermodynamic parameters of phosphate onto reeds at different temperatures.

T (K)	b (L/g)	ΔG (kJ/mol)	ΔH (kJ/mol)	ΔS (kJ/mol)
298.15	0.014	−10.5348	−23.1	−0.113
308.15	0.010	−11.7114
318.15	0.008	−12.7867

Kinetic studies

A kinetic study was performed to determine the rate at which the adsorption reaction is taking place as well as the adsorption mechanism. Three kinetic models were proposed in the analysis of the sorption data of phosphate onto reeds: pseudo-first order model, pseudo-second-order model, and the intra-particle diffusion model. The linearized form of the three models are depicted in Figures 6 –8.

Figure 6.

First-order adsorption plot for various phosphate concentrations onto activated reeds 25℃ and pH 6.5.

Figure 7.

Second-order adsorption plot for various phosphate concentrations onto activated reeds 25℃ and pH 6.5.

Figure 8.

Intra-particle diffusion model for adsorption of various phosphate concentrations onto activated reeds 25℃ and pH 6.5.

The kinetic parameters, the determination coefficients, and the standard deviation of the three models are shown in Table 2.

Table 2.

Kinetic parameters for adsorption rate of phosphate on activated reeds at different initial concentrations.

	First-order-model			Second-order-model				Intra-particle diffusion model
C₀ (mg/L)	k₁ (min⁻¹)	R ²	Δq_t (%)	k₂ (mg/g.min)	h	R ²	Δq_t (%)	k_p (mg/g.min^1/2)	R ²	Δq_t (%)
20	0.049	0.9926	13.8	0.023	0.353	0.9975	2.75	0.190	0.7906	5.76
40	0.048	0.9859	23.2	0.020	1.047	0.9985	3.04	0.266	0.8857	2.91
60	0.046	0.9305	33.61	0.017	1.590	0.9921	6.46	0.222	0.9350	1.11
100	0.048	0.9949	24.5	0.002	1.769	0.9665	9.05	0.814	0.8310	4.20

Based on the determination coefficients and normalized standard deviation, it was found that the adsorption data can be best fit by the second-order models. This model has shown (R²) values higher than 0.99 at all different initial phosphate concentrations. This indicates that the external adsorption is more appearing than micropores adsorption (Mezenner and Bensmaili, 2009). It can be noticed that the second-order model constant decreased with higher initial phosphate concentration. Additionally, the initial adsorption rate h, was directly proportional to the increase in initial phosphate concentration, which could be attributed to the increase in concentration gradient and consequently in the mass transfer driving force (Liu et al., 2008). The intra-particle diffusion model, which suggests that the rate of adsorption depends on the speed of adsorbate diffusion toward adsorbent surface, was not applicable (R² = 0.79–0.93).

Equilibrium isotherms

Various adsorption models have been used in literature to describe the equilibrium relationship between the liquid phase concentration and the surface concentration of the adsorbate. Two different models, Langmuir and Freundlich are used in this section to simulate the adsorption isotherms of phosphate ion onto reeds. The linear equation for Langmuir isotherm is as follows:

\frac{1}{q_{e}} = \frac{1}{Q_{m}} + (\frac{1}{C_{e}} \times \frac{1}{b \times Q_{m}})

(8) where C_e is the phosphate ion concentration in solution at equilibrium (mg/L). Q_m and b are the Langmuir coefficients associated with the maximum theoretical adsorption capacity and energy of adsorption, respectively. These coefficients can be calculated from the slope and intercept from (6). The Langmuir isotherm for phosphate adsorption on reeds at different temperatures is shown in Figure 9.

Figure 9.

Langmuir isotherm for the adsorption of phosphate on reeds (rpm: 100; pH: 6.5; ms: 5 mg/mL).

The Freudlich isotherm theory states that the ratio of q_e to C_e is not constant at different concentrations (Aksu, 2002). The isotherm is described by the following equation:

q_{e} = K_{F} \times {C_{e}}^{\frac{1}{n}}

(9)

The linearized form of equation (9) is:

\ln q_{e} = \ln K_{F} + \frac{1}{n} \ln C_{e}

(10)

K_F is the Freundlich constant related to the capacity of adsorption, 1/n is the Freundlich constant related to the intensity of adsorption. The Freundlich linear relations was used to fit the equilibrium data as shown in Figure 10.

Figure 10.

Freundlich isotherm for the adsorption of phosphate on reeds (rpm: 100; pH: 6.5; ms: 5 mg/mL).

The calculated coefficients Q_m, b, K_F, and n for both Langmuir and Freundlich's isotherms as well as the determination coefficient (R²) were determined at different temperatures and listed in Table 3.

Table 3.

Langmuir and Freundlich's constants related to adsorption of phosphate from aqueous solution by onto reeds.

Adsorption temp. (℃)	Langmuir isotherm			Freundlich isotherm
Adsorption temp. (℃)	Q_m (mg/g)	b (L/g)	R ²	K_F	n	R ²
25	50	0.014	0.96	1.67	1	0.90
35	47	0.010	0.97	2.33	1.2	0.94
45	39	0.008	0.96	0.737	1	0.94

Based on the determination coefficients, it can be concluded that the data were best represented by Langmuir isotherm. Consequently, it can be suggested that the adsorption took place in the form of monolayer on a homogeneous surface in terms of energies of sorption. It assumes also that no transmigration of sorbate occurs on the plane of the surface (Wu, 2007). This result is in good agreement with the results reported by Yue et al. (2010). As it is also indicated in Table 3, the adsorption capacity Q_m decreases with temperature, which is in agreement with the exothermic nature of the sorption process previously mentioned.

Conclusion

This work investigated the possibility of using Arundo donax reeds as effective, natural, and inexpensive adsorbent to remove phosphate ions from aqueous experimental solutions. The adsorption uptake was dependent on the pH range and favored at neutral pH (6.5–7) showing a maximum q_e of 16.2 mg/g, which is relatively high when compared with capacities reported in literature. It was thought the predominant phosphate form at this pH is HPO4⁻².The high initial phosphate concentration has also a positive effect on phosphate adsorption since the sorption process is mass transfer driven. The phosphate uptake was found to be inversely proportional to the adsorbent dose and temperature. The highest adsorption capacities were observed at 25℃, which indicates the exothermic nature of the process. The equilibrium isotherms indicated that the experimental data were well represented by the Langmuir model. The kinetic study conducted in this work showed that the data follow the pseudo-second-order model (R² = 0.96 to 0.99) suggesting that the external adsorption is the key adsorption process. Finally, the (ΔH) and (ΔG) values were determined from a thermodynamic analysis and the negative values of both confirm the spontaneity and exothermic nature of the process.

Footnotes

Declaration of Conflicting Interests

The author(s) declared no potential conflicts of interest with respect to the research, authorship, and/or publication of this article.

Funding

The author(s) received no financial support for the research, authorship, and/or publication of this article.

References

Ajmal

Khan

Ahmad

(1998) Role of sawdust in the removal of copper(II) from industrial wastes. Water Research 32: 3085–3091.

Aksu

(2002) Determination of the equilibrium, kinetic and thermodynamic parameters of the batch biosorption of nickel(II) ions onto Chlorella vulgaris. Process Biochemistry 38: 89–99.

Ammari

(2014) Utilization of a natural ecosystem bio-waste; leaves of Arundo donaxreed, as a raw material of low-cost eco-biosorbent for cadmiumremoval from aqueous phase. Ecological Engineering 71: 466–473.

American Public Health Association; American Water Works Association; Water Environment Federation (1989) Standards Methods for the Examination of Water and Wastewater. 17th ed. Washington, DC: American Public Health Association.

Biswas

Inoue

Ghimire

et al. (2008) Removal and recovery of phosphorus from water by means of adsorption onto orange waste gel loaded with zirconium. Bioresource Technology 99: 8685–8690.

Chmielewská

Hodossyová

Bujdoš

(2013) Kinetic and thermodynamic studies for phosphate removal using natural adsorption materials. Polish Journal of Environmental Studies 22(5): 1307–1316.

Dietze

Wiesmann

Gnirss

(2003) removal with membrane filtration for surface water treatment. Water Science and Technology: Water Supply 3: 23–30.

Elzinga

Sparks

(2007) Phosphate adsorption onto hematite: An in situ ATR-FTIR investigation of the effects of pH and loading level on the mode of phosphate surface complexation. Journal of Colloid and Interface Science 308: 53–70.

Fazhi

Fengchang

Guijian

et al. (2014) Removal of phosphate from eutrophic lakes through adsorption by in situ formation of magnesium hydroxide from diatomite. Environmental Science & Technology 48: 582–590.

10.

Fogler

(1992) Elements of Chemical Reaction Engineering, 2nd ed. New Jersey, NJ: Prentice Hall.

11.

Genz

Kornmüller

Jekel

(2004) Advanced phosphorus removal from membrane filtrates by adsorption on activated aluminium oxide and granulated ferric hydroxide. Water Research 38: 3523–3530.

12.

(2004) Citation review of Lagergren kinetic rate equation on adsorption reactions. Scientometrics 59: 171–177.

13.

Hong-Bing

Fei

Hao

et al. (2012) Adsorption Control 319 performance of phosphorus removal from agricultural non-point source pollution by nano-aperture Lanthanum-modified active alumina. Advance Journal of Food Science and Technology 4: 337–343.

14.

Karaca

Gürses

Ejder

et al. (2004) Kinetic modeling of liquid-phase adsorption of phosphate on dolomite. Journal of Colloid and Interface Science 277: 257–263.

15.

Karageorgiou

Paschalis

Anastassakis

(2007) Removal of phosphate species from solution by adsorption onto calcite used as natural adsorbent. Journal of Hazardous Materials 139: 447–452.

16.

Kennedy

Vijaya

Sekaran

et al. (2007) Equilibrium, kinetic and thermodynamic studies on the adsorption of m-cresol onto micro- and mesoporous carbon. Journal of Hazardous Materials 149: 134–143.

17.

Khan

Rehman

et al. (2010) Comparison of different models for phosphate adsorption in salt inherent soil series of Dera Ismail Khan. Soil & Environment 29: 11–14.

18.

Krishnan

Haridas

(2008) Removal of phosphate from aqueous solutions and sewage using natural and surface modified coir pith. Journal of Hazardous Materials 152: 527–535.

19.

Kumar

Sudha

Chand

et al. (2010) Phosphate removal from aqueous solution using coir-pith activated carbon. Separation Science and Technology 45: 1463–1470.

20.

Liu

Sun

Yin

et al. (2008) Removal of phosphate by mesoporous ZrO2. Journal of Hazardous Materials 151: 616–622.

21.

Luengo

Brigante

Antelo

et al. (2006) Kinetics of phosphate adsorption on goethite: Comparing batch adsorption and ATR-IR measurements. Journal of Colloid and Interface Science 300: 511–518.

22.

Mallikarjun

Mise

(2012) A study of phosphate adsorption characteristics on different soils. IOSR Journal of Engineering 2: 13–23.

23.

Manju

Raji

Anirudhan

(1998) Evaluation of coconut husk carbon for the removal of arsenic from water. Water Research 32: 3070–3602.

24.

Mezenner

Bensmaili

(2009) Kinetics and thermodynamic study of phosphate adsorption on iron hydroxide-eggshell waste. Chemical Engineering Journal 147: 87–96.

25.

Min

Ahn

Lee

et al. (2001) Adsorption characteristics of phenol and chlorophenols on granular activated carbon. Microchemical Journal 70: 123–131.

26.

Morse

Brett

Guy

et al. (1998) Review: Phosphorus removal and recovery technologies. Science of the Total Environment 212: 69–81.

27.

Nowack

Stone

(2006) Competitive adsorption of phosphate and phosphonates onto goethite. Water Research 40: 2201–2209.

28.

Razaq IBA (1989) Effect of pH and exchangeable metals on phosphate adsorption by soils. Retrospective Theses and Dissertations. Paper 9170.

29.

Rittmann

Mayer

Westerhoff

et al. (2011) Capturing the lost phosphorus. Chemosphere 84: 846–853.

30.

Robinson

Chandran

Nigam

(2002) Removal of dyes from a synthetic textile dye effluent by biosorption on apple pomace and wheat straw. Water Research 36: 2824–2830.

31.

(2007) Studies of the equilibrium and thermodynamics of the adsorption of Cu(2+) onto as-produced and modified carbon nanotubes. Journal of Colloid and Interface Science, 311: 338–346.

32.

Xiaoming

Fan

Wenfeng

et al. (2013) Sparks characteristics of phosphate adsorption-desorption onto ferrihydrite: Comparison with well-crystalline Fe (Hydr) oxides. Soil Science 178: 1–11.

33.

Gao

et al. (2011) Characteristics of diethylenetriamine-crosslinked cotton stalk/wheat stalk and their biosorption capacities for phosphate. Journal of Hazardous Materials 192: 1690–1696.

34.

Yue

Wang

Gao

et al. (2010) Phosphate removal from aqueous solution by adsorption on modified giant reed. Water Environment Research 82: 374–381.

Kinetic and thermodynamic study of phosphate removal from water by adsorption onto ( Arundo donax ) reeds