Kinetics and Thermodynamics of Uranium (VI) Adsorption onto Humic Acid Derived from Leonardite

Humic acid (HA) is well known as an inexpensive and effective adsorbent for heavy metal ions. However, the thermodynamics of uranium (U) adsorption onto HA is not fully understood. This study aimed to understand the kinetics and isotherms of U(VI) adsorption onto HA under different temperatures from acidic water. A leonardite-derived HA was characterized for its ash content, elemental compositions, and acidic functional groups, and used for the removal of U (VI) from acidic aqueous solutions via batch experiments at initial concentrations of 0–100 mg·L−1 at 298, 308 and 318 K. ICP-MS was used to determine the U(VI) concentrations in solutions before and after reacting with the HA. The rate and capacity of HA adsorbing U(VI) increased with the temperature. Adsorption kinetic data was best fitted to the pseudo second-order model. This, together with FTIR spectra, indicated a chemisorption of U(VI) by HA. Equilibrium adsorption data was best fitted to the Langmuir and Temkin models. Thermodynamic parameters such as equilibrium constant (K0), standard Gibbs free energy (ΔG0), standard enthalpy change (ΔH0), and standard entropy change (ΔS0), indicated that U(VI) adsorption onto HA was endothermic and spontaneous. The co-existence of cations (Cu2+, Co2+, Cd2+ and Pb2+) and anions (HPO42− and SO42−) reduced U(VI) adsorption. The high propensity and capacity of leonardite-derived HA adsorbing U(VI) suggests that it has the potential for cost-effective removal of U(VI) from acidic contaminated waters.


Introduction
Uranium (U) is widely but unevenly distributed in soils with an average concentration of 2.6 mg·kg −1 [1]. Acid mining drainage is a major source of U release into soil and water environments [2,3]. Naturally occurring U consists of three isotopes: U-238 (99.2739-99.2752%), U-235 (0.7198-0.7202%) and U-234 (0.0050-0.0059%). In oxidizing environments U is usually found in hexavalent form. U accumulation moves up the food chain, and eventually, to human organs and tissues, causing severe damage to kidneys, liver and in extreme cases, death [4]. The World Health Organization and US EPA have set the maximum concentration for U in drinking water at 15 and 30 µg·L −1 , respectively [5,6].

Preparation and Characterization of HA
HA was extracted from the leonardite with traditional alkaline-acid protocol [8]. Briefly, 25 g leonardite was placed into a Teflon-container with 250 mL 0.1 M NaOH and sonicated for 30 min. After standing overnight, the supernatant was collected. This process was repeated 2 more times for a total of 3 extractions. The collected supernatants were combined, and small aliquots of 6 M HCl was titrated in, while stirring, until the pH was reduced to 2. The suspensions were then centrifuged at 3000 g for 15 min. The precipitates (HA) were washed three times with distilled water and then freeze-dried for later use.
The physical and chemical properties of the leonardite and derived HA were analyzed as follows: Ash content was determined with ignition in a muffle furnace at 800 • C for 4 h under atmospheric condition. Elemental compositions were determined with an elemental analyzer (Vario micro cube, Elementar, Germany) for dried samples at 80 • C. Functional groups were identified with Fourier transform infrared spectroscopy (Spectrum Two, PerkinElmer, Waltham, MA, USA), and acidic functional groups were quantified with the titration method of the International Humic Substances Society [16].

Adsorption Experiments and Data Processing
All adsorption experiments were conducted in duplicates, including blanks and calibration controls. Briefly, 20 mg of HA was weighed into 50 mL plastic centrifuge tubes (Corning, Corning, NY, USA) with 30 mL U solution, and the pH of the suspension was adjusted to 3.0. The tubes were then shaken for 6 h to achieve equilibrium. Then, the tubes were centrifuged, and the supernatants were filtered through a 0.45 µm membrane (Whatman, Little Chalfont, Buckinghamshire, UK) for analysis of U concentration with an ICP-MS (Varian Inc., Palo Alto, CA, USA). The pH at the beginning and end of adsorption experiment was measured by a pH meter (Oakton, Vernon Hills, IL, USA).
U adsorption on the HA was calculated from the difference in concentrations before and after the adsorption. MS-Excel and OriginPro 8.0 (OriginLab, Wellesley Hills, MA, USA) were used for data processing.

Adsorption Kinetics Models
Parameters obtained from four adsorption models were used to describe the kinetics of U(VI) adsorption onto HA: pseudo first-order model (Equation (1)) was used to describe the adsorption process in solid-liquid system at the initial phase, which corresponds to a diffusion-controlled process [17,18]; pseudo second-order model (Equation (2)) was used to describe whole adsorption process, involving chemisorption in solid-liquid system [18,19]; the Elovich equation (Equation (3)) was used to describe the chemisorption that occurred on heterogeneous solid surface [20,21]; and the intraparticle diffusion model (Equation (4)) was used to determine the intraparticle diffusion rate constant and the boundary resistance [22]. Detailed descriptions on the models and parameters are available in the literature [17][18][19][20][21][22].

Adsorption Isotherm Models
Four adsorption isotherm models were used to describe U distribution between solution and HA at the equilibrium state: the Freundlich model (Equation (5)) describes both monolayer and multilayer adsorption, which is based on heterogeneous adsorption in solid-liquid system [23,24]; the Langmuir model (Equation (6)) quantifies the adsorption capacity [8,25,26]; the Temkin model (Equation (7)) takes U-HA interaction into account and links adsorption energy to the adsorbent surface [27]; and the Dubinin-Radushkevich (D-R) model (Equation (8)) describes adsorption reaction at low concentration ranges on the homogeneous or heterogeneous surface [28].

Thermodynamic Parameters
The thermodynamic parameters are usually used to illustrate adsorption mechanisms and determine the reaction direction, which can be calculated from the thermodynamic equilibrium constant, K 0 . The standard Gibbs free energy ∆G 0 (kJ·mol −1 ), standard enthalpy change ∆H 0 (kJ·mol −1 ), and standard entropy change ∆S 0 (J·mol −1 ·K −1 ) were determined from the equations as follows: K 0 can be defined as, where R is the gas constant (8.314 J·mol −1 ·K −1 ), T is the temperature in K, C e is the equilibrium concentration (mg·L −1 ), and q e is the amount of adsorption at equilibrium state (mg·g −1 ).

Properties of Adsorbents
The properties of the leonardite and HA are shown in Table 1. HA had lower pH and ash content, but higher C and O contents than leonardite. Both HA and leonardite had abundant acidic functional groups (carboxyl and phenolic-hydroxyl), of which carboxyl groups are considered as the most important for adsorbing metal ions [13]. The FTIR spectra of HA ( Figure 1) confirmed the existence of oxygen-containing functional groups, as shown at wavenumbers of 3201 cm −1 (OH stretch of phenolic-OH), 1704 cm −1 (C=C stretch of COOH groups), 1601 cm −1 (asymmetric -COOstretch), 1426 cm −1 (symmetric -COOstretch), 1368 cm −1 (salts of -COOH), 1204 cm −1 (-C-O stretch and phenolic C-OH) and 1032 cm −1 (O-CH 3 vibrations) [29,30].
K0 can be defined as, where R is the gas constant (8.314 J·mol −1 ·K −1 ), T is the temperature in K, Ce is the equilibrium concentration (mg·L −1 ), and qe is the amount of adsorption at equilibrium state (mg·g −1 ).

Properties of Adsorbents
The properties of the leonardite and HA are shown in Table 1. HA had lower pH and ash content, but higher C and O contents than leonardite. Both HA and leonardite had abundant acidic functional groups (carboxyl and phenolic-hydroxyl), of which carboxyl groups are considered as the most important for adsorbing metal ions [13]. Table 1. The selected properties of leonardite and leonardite-derived humic acid (HA).

Adsorption Kinetics
Figures 2 and 3 show that U(VI) adsorption increased with rising temperature, indicating an endothermic process. This may be due to the increased binding sites of HA at a higher temperature [31]. Similar results were reported in the literature [8,25,31]. The time required for U(VI) adsorption process to reach equilibrium was 1.5 h at 298 K, 2 h at 308 K and 318 K.  Figures 2 and 3 show that U(VI) adsorption increased with rising temperature, indicating an endothermic process. This may be due to the increased binding sites of HA at a higher temperature [31]. Similar results were reported in the literature [8,25,31]. The time required for U(VI) adsorption process to reach equilibrium was 1.5 h at 298 K, 2 h at 308 K and 318 K.  Adsorption kinetics parameters are given in Table 2. The three models fit the adsorption process well (R 2 > 0.95). The Elovich model had the highest R 2 , indicating that U(VI) adsorption onto HA may be chemisorption rather than intraparticle diffusion [18,20]. This was further evidenced by a low R 2 value from the intraparticle diffusion equation (<0.70) in Table 3, which suggests that the adsorption process was not controlled by intraparticle diffusion.  Adsorption kinetics parameters are given in Table 2. The three models fit the adsorption process well (R 2 > 0.95). The Elovich model had the highest R 2 , indicating that U(VI) adsorption onto HA may be chemisorption rather than intraparticle diffusion [18,20]. This was further evidenced by a low R 2 value from the intraparticle diffusion equation (<0.70) in Table 3, which suggests that the adsorption process was not controlled by intraparticle diffusion.

Adsorption Isotherms
As shown in Figure 4, adsorption capacity increased with U concentrations. The parameters from fitting adsorption data into four isotherm models are given in Table 4.  The n values of Freundlich equation were higher than unity, indicating adsorption may be chemical rather than physical in nature with a high affinity of HA for U(VI), thus a high adsorption capacity [25,32]. Constant, kF, was related to adsorption capacity. Its increase with temperature also confirmed that U(VI) adsorption on HA was endothermic. Adsorption data fit the Langmuir model The n values of Freundlich equation were higher than unity, indicating adsorption may be chemical rather than physical in nature with a high affinity of HA for U(VI), thus a high adsorption capacity [25,32]. Constant, k F , was related to adsorption capacity. Its increase with temperature also confirmed that U(VI) adsorption on HA was endothermic. Adsorption data fit the Langmuir model well (R 2 > 0.95). The maximum adsorption capacity (q L ) at the concentration range of 0-100 mg/L increased with temperature. Even at acidic condition (pH 3), q L of 68.6 mg·g −1 was higher than the adsorption capacities of common adsorbents (kaolin, biochar, activated carbon, hematite, and bentonite) at near-neutral pH that would not be observed in acidic effluents ( Table 5). The large adsorption capacity of HA for U is in agreement with its abundant carboxyl group [13]. The good fit of experimental data with Temkin equation (R 2 > 0.97) implied that U(VI) adsorption onto HA involved chemisorption [33]. This was further supported by the results of pseudo second-order and Elovich equations. The q D values of D-R model were not consistent with the q L calculated from the Langmuir isotherm as show in Figure 4 and Table 4. Fitting of adsorption data into the D-R model produced the lowest R 2 in Table 4, further suggesting U(VI) adsorption onto HA was not a physical process [25,28,34,35].

Adsorption Thermodynamics
The values of lnK 0 at different temperatures were determined by linear plotting ln(q e /C e ) versus q e , assuming q e as zero as described in Figure 5a [8,14]. ∆G 0 values were calculated from Equation (9) as displayed in Table 6. ∆H 0 and ∆S 0 were determined on the bases of Equation (10) by plotting lnK 0 versus 1/T, included in Figure 5b. The negative ∆G 0 indicated that the adsorption reaction was spontaneous, and its extent of spontaneity increased with rising temperature. A positive ∆S 0 = 114.3 J·mol −1 ·K −1 suggested that U(VI) adsorption onto HA was endothermic, which was supported by the higher adsorption capacity at higher temperature. A positive ∆H 0 = 23.13 kJ·mol −1 revealed that the HA had a high affinity for U(VI). Further, ∆H 0 was a useful value to distinguish physisorption from chemisorption. In general, ∆H 0 for physisorption is small, 2.1-20.9 kJ·mol -1 , whereas ∆H 0 for chemisorption is large, 20.9-418.4 kJ·mol −1 [38,39]. The value of ∆H 0 in the range of 20.9-418.4 kJ·mol −1 indicated that the adsorption of U(VI) onto HA involved chemisorption [39].
C: the U 6+ concentration range; The qm was calculated from the Langmuir equation.

Adsorption Thermodynamics
The values of lnK0 at different temperatures were determined by linear plotting ln(qe/Ce) versus qe, assuming qe as zero as described in Figure 5a [8,14]. ΔG 0 values were calculated from Equation 9 as displayed in Table 6. ΔH 0 and ΔS 0 were determined on the bases of Equation 10 by plotting lnK0 versus 1/T, included in Figure 5b. The negative ΔG 0 indicated that the adsorption reaction was spontaneous, and its extent of spontaneity increased with rising temperature. A positive ΔS 0 = 114.3 J·mol −1 ·K −1 suggested that U(VI) adsorption onto HA was endothermic, which was supported by the higher adsorption capacity at higher temperature. A positive ΔH 0 = 23.13 kJ·mol −1 revealed that the HA had a high affinity for U(VI). Further, ΔH 0 was a useful value to distinguish physisorption from chemisorption. In general, ΔH 0 for physisorption is small, 2.1-20.9 kJ·mol -1 , whereas ΔH 0 for chemisorption is large, 20.9-418.4 kJ·mol −1 [38,39]. The value of ΔH 0 in the range of 20.9-418.4 kJ·mol −1 indicated that the adsorption of U(VI) onto HA involved chemisorption [39].

Adsorption Mechanism
FTIR is a useful tool to probe adsorption behavior of cations onto adsorbents [8,23,40]. The vibration frequency changes in characteristic peaks of HA before and after adsorption (Figure 1) include the shifts of the symmetric -COO − stretch frequency from 1601 to 1590 cm −1 (red shift), symmetric -COO − stretch frequency from 1426 to 1416 cm −1 (red shift), salts of -COOH stretch frequency from 1368 to 1360 cm −1 (red shift), and phenolic C-OH stretch frequency from 1204 to 1219 cm −1 (blue shift). Thus, U(VI) reacted with HA through functional groups [8,41]. The FTIR analysis further elaborated that U(VI) adsorption onto HA was via chemisorption. The adsorption process could be controlled by surface or intraparticle diffusion, and the intraparticle diffusion model is often used to make the judgment [22,42]. The parameters and R 2 of data fitting into the intraparticle diffusion model are given in Table 3. The low R 2 (< 0.7) suggested that the adsorption process was not controlled by intraparticle diffusion. In other words, surface diffusion was the dominant process for U(VI) adsorption onto HA via chemisorption, such as ion-exchange, complexation and chelation [25,30].

The Effects of Cations and Anions on U(VI) Adsorption
Anions and cations are common in acidic U contaminated water and in soil environment [43]. They may affect U(VI) adsorption onto HA. Figure 6 shows the effect of common cations and anions on U(VI) adsorption onto HA. The presence of Cu 2+ , Co 2+ , Cd 2+ and Pb 2+ cations reduced U(VI) adsorption capacity, which could be explained by the competitive adsorption of the cations for U(VI) [44,45]. However, they are not good competitors for U(VI) and Pb 2+ was the least competitive.
The U(VI) adsorption decreased as co-existing cation concentrations increased, which is consistent with previous studies [46,47]. The presence of anions HPO 4 2− and SO 4 2− greatly reduced the adsorption capacity of HA for U(VI) as shown in Figure 6b. For SO 4

The Effects of Cations and Anions on U(VI) Adsorption
Anions and cations are common in acidic U contaminated water and in soil environment [43]. They may affect U(VI) adsorption onto HA. Figure 6 shows the effect of common cations and anions on U(VI) adsorption onto HA. The presence of Cu 2+ , Co 2+ , Cd 2+ and Pb 2+ cations reduced U(VI) adsorption capacity, which could be explained by the competitive adsorption of the cations for U(VI) [44,45]. However, they are not good competitors for U(VI) and Pb 2+ was the least competitive. The U(VI) adsorption decreased as co-existing cation concentrations increased, which is consistent with previous studies [46,47]. The presence of anions HPO4 2− and SO4 2− greatly reduced the adsorption capacity of HA for U(VI) as shown in Figure 6b. For SO4 2− , the reduced adsorption may be caused by the competition between SO4 2− and HA for UO2 2+ , or the formation of negatively charged complexes with UO2 2+ [43,48,49]. At acidic condition, the HPO4 2− can react with H + to form H2PO4 − and H3PO4 [50]. HPO4 2− had stronger effects than SO4 2− . This may be caused by the formation of precipitation between UO2 2+ and HPO4 2− , H2PO4 − and H3PO4, which could prevent UO2 2+ being adsorbed onto HA surface [48,[50][51][52].

Conclusions
HA derived from leonardite was an effective adsorbent for removing uranium from aqueous acid solutions. The adsorption increased as temperature increased. Data fitting into kinetic models and large ∆H 0 suggested that the adsorption involved chemisorption. The thermodynamic parameters indicated that the adsorption process was endothermic and spontaneous. Co-existing cations and anions had negative effects on U(VI) adsorption onto HA. Because of its wide availability and low-cost HA has a potential for use in the treatment of acidic mining effluents.