Antimony(III) oxidation and antimony(V) adsorption reactions on synthetic manganite

Chemie der Erde 72 (2012) S4, 41–47

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Antimony(III) oxidation and antimony(V) adsorption reactions on synthetic manganite Xiangqin Wang a , Mengchang He a,∗ , Chunye Lin a , Yuxi Gao b , Lei Zheng b a b

State Key Laboratory of Water Environment Simulation, School of Environment, Beijing Normal University, No. 19 Xinjiekouwai Street, Beijing 100875, China Institute of High Energy Physics, Chinese Academy of Sciences, 19B YuquanLu, Shijingshan District, Beijing 100049, China

a r t i c l e

i n f o

Article history: Received 26 September 2011 Accepted 18 February 2012 Keywords: Antimony Oxidation Adsorption Manganite XANES

a b s t r a c t Oxidation and adsorption processes critically affect the mobility of antimony (Sb) in the environment. Minerals such as manganese oxides appear to be important adsorbents or oxidants for Sb in soils and sediments. In this study, Sb oxidation and adsorption onto synthetic manganite (␥-MnOOH) in aqueous suspensions were comprehensively investigated. The oxidation of Sb(III) by manganite occurred on a time scale of minutes and Sb(V) was released into the suspension as soon as the reaction began. X-ray absorption near edge structure (XANES) analyses showed that Sb(V) was the dominant species adsorbed on manganite, suggesting that manganite acts as a strong oxidant of Sb(III). The effects of temperature on the adsorption rates of Sb(V) were investigated. Thermodynamic parameters indicated that adsorption of Sb(V) by manganite is an exothermic process and is spontaneous at the specific temperatures investigated. A change in ionic strength from 0.001 to 0.1 M NaNO3 had little effect on the adsorption of Sb(V) onto manganite, indicating that Sb(V) forms inner-sphere complexes at the mineral surface. Anions such as phosphate, sulfate, silicate and carbonate, which may be present in natural waters, were ineffective competitors with Sb(V) for sorption sites. © 2012 Elsevier GmbH. All rights reserved.

1. Introduction Antimony (Sb) is consumed in large quantities (>100,000 tons annually worldwide) in a variety of industrial products (Leuz et al., 2006a), e.g., flame retardants, catalysts, ceramics, batteries, and alloys. Elevated concentrations of Sb in soils and water have been detected around mining and smelter areas, along road sides bearing dust from brake pads and tires, and in shooting ranges (Flynn et al., 2003; Wilson et al., 2004; He, 2007; Westerhoff et al., 2008; Wang et al., 2011). Sb lies in the same main group as arsenic (group 15) and, thus, exhibits similar chemical behaviors such as the formation of trivalent and pentavalent species or oxyanions in natural waters due to the similar chemical characteristics (Watkins et al., 2006). In oxic waters, Sb(V) is the predominant species and occurs as Sb(OH)6 − (Cooper et al., 1998; Cai et al., 1998), whereas under anoxic solutions, Sb(III) occurs as Sb(OH)3 and is more stable (Cooper et al., 1998). The Sb species Sb(OH)6 − sorbs less well than the Sb(III) species to Fe and Mn (hydr)oxide surfaces and are more soluble than oxides of Sb(III). Thus, oxidation processes critically affect the mobility of antimony due to the greater solubility of Sb(V) relative to Sb(III)

∗ Corresponding author. Tel.: +86 10 5880 7172; fax: +86 10 5880 7172. E-mail address: [email protected] (M. He). 0009-2819/$ – see front matter © 2012 Elsevier GmbH. All rights reserved. doi:10.1016/j.chemer.2012.02.002

(Leuz et al., 2006b). Evidence of the association of Sb species with hydrous oxides of Mn, Al and Fe has been shown by adsorption studies (Thanabalasingam and Pickering, 1990; Belzile et al., 2001; Leuz et al., 2006a). One of the studies showed that sorption of Sb was greater on the Mn oxide surface than on Al(OH)3 or goethite, indicating that the affinity between Mn oxide surface sites and the Sb species was very high (Thanabalasingam and Pickering, 1990). The oxidizing and adsorption capacities of Mn oxyhydroxides have already been demonstrated for arsenic species (Oscarson et al., 1981; De Vitre et al., 1991). There is a wide range of manganese oxyhydroxide minerals in the environment resulting from the oxidation of soluble, aqueous Mn(II) by both abiotic and biotic means (Huang, 1991; Ehrlich, 1996; Shaughnessy et al., 2003). They are most commonly found as coatings on soil and sediment grains. Manganite has been reported to be present in lakes and rivers in the temperate and subarctic zones of the world (Stumm and Giovanoli, 1976; Hem and Lind, 1983; Murray et al., 1985), and thus, it should play an important role in the geochemical cycling of elements in these areas (Ramstedt et al., 2004). The objective of the present work is to investigate the reaction mechanisms between Sb and manganite in combination with XANES analyses to examine Sb speciation on the mineral. The adsorption of low concentrations of Sb(V) onto manganite as a function of temperature, pH, ionic strength and co-existing anions was investigated.

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X. Wang et al. / Chemie der Erde 72 (2012) S4, 41–47

2. Materials and methods All chemicals were of analytical grade and were purchased from Beijing Chemical Company (Beijing, China). All solutions were prepared from these chemicals without further purification and distilled, deionized water (DDW) with a resistivity of 18 M cm (Millipore Corp., Milford, MA). Polyethylene bottles and glassware were soaked in 10% HNO3 , rinsed with distilled water and then rinsed with DDW water before use. The Sb(III) and Sb(V) stock solutions were prepared from Sb2 O3 dissolved in 2 M HCl and KSb(OH)6 dissolved in water, respectively. 2.1. Adsorbent preparation and characterization Manganite was synthesized following a modification of the method of Giovanoli and Leuenberger (1969). Heating of a 1 L, 0.06 M MnSO4 solution to 60 ◦ C was followed by the addition of 20.4 ml 30% H2 O2 . Then, 300 ml of 0.20 M NH3 was added to the H2 O2 /MnSO4 solution while the solution was vigorously stirred. All three solutions were purged with N2 before and after mixing. The resulting brown suspension was quickly heated to 95 ◦ C and maintained at this temperature for 6 h. While still hot, the dark brown suspension was centrifuged and repeatedly washed with 1000 ml (total) of hot (ca. 80 ◦ C) DDW water. The solid suspension was then dried in a vacuum freeze dryer, crushed, and stored in a freezer. Nitrogen multipoint BET analysis showed that the specific surface area of the synthetic manganite was 64.2 m2 /g. The average oxidation state of manganese was determined by iodometric titration (Clesceri et al., 2000; Wang and Stone, 2006). The sample was reduced by an excess of KI in a 0.01 M HCl solution, yielding a brown I2 solution. Soluble starch indicator was then added and the brown I2 solution was back-titrated using Na2 S2 O3 . The endpoint was reached when the suspension was colorless. The average oxidation state was found to be 2.95. X-ray diffraction (XRD) analysis was performed on an X’Pert PRO MPD diffractometer using Ni-filtered copper K␣ radiation. 2.2. Kinetic experiments Kinetic experiments were performed with a background electrolyte of 0.01 M NaNO3 and 5 mM CH3 COONa at pH 4.0. The desired pH was achieved by addition of 1.0 M HNO3 or NaOH. The pH was kept constant during the experiment with small additions of 0.1 M HNO3 or NaOH. N2 gas was used to purge O2 . For the Sb(III) oxidation experiments, manganite suspensions were allowed to equilibrate in batch reactors in a constant temperature water bath at 25 ◦ C, whereas for the Sb(V) sorption experiment the manganite suspensions were allowed to equilibrate in batch reactors in constant temperature water baths at 0, 25 and 40 ◦ C. During the kinetic experiments, all of the suspensions were stirred with Teflon-coated bars and reaction suspension aliquots were withdrawn at selected intervals. For dissolved Sb(III) and Sb(V), 5 ml aliquots were removed with a plastic syringe and filtered through a 0.22 ␮m cellulose acetate filter immediately and then delivered into polyethylene bottles. The filtered solution was immediately analyzed for separation of Sb(III) and Sb(V). 2.3. Sb LIII -edge analysis Samples of synthetic manganite treated with Sb(III) were prepared for X-ray absorption near edge structure (XANES) analysis by reacting 1.00 g of synthetic manganite with 1.0 L of 8.21 mM antimony (both Sb(III) and Sb(V)) in 0.01 M NaNO3 in a stirred 2 L glass beaker for 24 h at pH 7.0 at 25 ◦ C. An identical preparation was made using Sb(V). After 24 h of equilibration, suspensions of the adsorption experiments were centrifuged and then filtered

through 0.22 ␮m membrane filters prior to analysis. The moist solids of adsorption samples were mounted in a 2-mm thick cell and sealed with adhesive PVC tape for XANES measurements. XANES measurements at Sb LIII edge (4698 eV) were performed by Beijing Synchrotron Radiation Facilities (BSRF), using a Si(1 1 1) double crystal as a monochromator. The XANES data for cells were recorded in total electron yield detection mode. XANES spectra were normalized far from the edge (∼4760 eV) and analyzed in comparison to Sb2 O3 and KSb(OH)6 references to determine the Sb oxidation in the manganite solids. 2.4. Sorption experiments Adsorption tests were performed as triplicates in 50 ml polyethylene bottles to evaluate the adsorption capacity of manganite to Sb. Sorption of Sb(V) on manganite was initiated by the addition of Sb(V) stock solutions to solid suspensions at three ionic strengths (0.001, 0.01, 0.1 M) and different pH values, which were adjusted with NaOH and HCl. The pH of the suspensions was adjusted every 4 h with HCl and NaOH to designated values in the range of 3–12 during the shaking process. After the desired reaction time, suspensions were centrifuged and the supernatant was filtered through a 0.22 ␮m membrane. The Sb(V) sorption isotherm was performed at pH 4.0, 7.0 and 9.0. Initial Sb(V) concentrations varied from 3.99 ␮mol/L to 808 ␮mol/L. To test the effects of co-existing anions (PO4 2− , CO3 2− , SO4 2− and silicate), anion concentrations ranged from 0.2 to 10 mM and ionic strength was adjusted to 0.01 M with NaNO3 . All batch experiments were performed at 25 ± 1 ◦ C with an adsorbent content of 0.4 g/L, and all the suspensions were shaken on an obit shaker at 110 rpm for 24 h. 2.5. Analytical methods Samples for total Sb and Sb(III) determinations were stored in 5% HCl at 4 ◦ C before analysis. Total Sb and Sb(III) were selectively analyzed using a hydride generation atomic fluorescence spectrometer (HG-AFS) (AFS 230, Haiguang Corp., Beijing) as described by Wang et al. (2011). All samples were analyzed within 24 h after collection. 3. Results and discussion 3.1. Sb(III) oxidation by manganite The oxidation of Sb(III) by manganite at pH 4.0 is shown in Fig. 1. Rapid uptake of Sb(III) by manganite was apparent as well as rapid formation of soluble Sb(V). Previous studies have observed similar behavior for the oxidation of arsenite by synthetic birnessite (Scott and Morgan, 1995) and manganite (Chiu and Hering, 2000). The authors suggest that the processes of electron transfer and release of As(V) are fast compared to the adsorption of As(III). Similarly, in this case, the processes of electron transfer and release of Sb(V) into solution were fast compared to the adsorption of Sb(III) by manganite. The decreases in total Sb(III) or dissolved Sb(III) did not exhibit first-order kinetic behavior, which suggests that the Sb(III) uptake reaction was complicated by surface reactions such as manganite surface alterations during the Sb(III) oxidation reaction or competition with Sb(V) product for binding sites. The more manganite particles were added into the suspension, the lower the dissolved Sb(III) detected during the reaction of 38.7 ␮M Sb(III) with synthetic manganite. Approximately 74.0% and 96.4% of the original Sb(III) added in the 200 mg/L and 600 mg/L manganite suspension, respectively, were unrecoverable and presumed to be adsorbed Sb(V).

X. Wang et al. / Chemie der Erde 72 (2012) S4, 41–47

43

a

35

6

25

Normalized absorption

Sb concentrations (μM)

30 Adsorbed Sb

20 15 10 5

γ-MnOOH-Sb(V) 4

γ-MnOOH-Sb(III)

KSb(OH)6

2

Sb2O3 0

0 0.0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

4680

4700

Time (h)

4720

4740

4760

Energy/eV

40 Fig. 2. L-edge XANES spectra for the Sb(III)- and Sb(V)-treated synthetic manganite.

b

Sb(III)

35

Sb(V)

Sb(III+V)

the spectra of reference materials that the absorption edge shifts to higher energy at the higher oxidation state of Sb, indicating that the position of the XANES peak can be used to distinguish between Sb(III) and Sb(V). The absorption edge energy between the KSb(OH)6 and the Sb(III)- and Sb(V)-treated manganite samples is noticeably similar, which provides direct evidence that the Sb(III) oxidized by manganite was adsorbed as Sb(V).

35

Sb concentrations (μM)

25 20 15 10

30 25 20 15 10 5 0 -0.02 0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 0.16 0.18

5

3.3. Sb(V) sorption kinetics

Time (h)

0 0.0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

Time (h) Fig. 1. Concentrations of Sb species in solution during the reaction processes of 38.7 ␮M Sb(III) with synthetic manganite: (a) 200 mg/L manganite and (b) 600 mg/L manganite.

During the initial 7 min of the reaction between Sb(III) with the 600 mg/L manganite suspension, Sb(V) in solution increased rapidly. After 7 min of reaction, a second reaction phase began in which the oxidation rate of Sb(III) slowed down. This second phase was indicated by a lack of Sb(V) increase and a gradual decrease of the oxidation product. Thus, the first 7 min of the reaction between Sb(III) and manganite represented the period of highest reactivity and the most rapid rate of Sb(III) oxidation. This observed decrease in reactivity of manganite after 7 min of reaction with Sb(III) can be attributed to the passivation of the manganite surface (Lafferty et al., 2010), which was likely caused by sorption of Mn(II) and Sb(V). Passivation caused by physically blocking reaction sites likely occurs due to Mn(II) or As(V) sorption at birnessite edge sites, which are the locations of As(III) oxidation (Tournassat et al., 2002; Zhu et al., 2009). At the termination of the reaction (3.0 h), the Sb surface coverage achieved on manganite was 310 ␮mol g−1 (4.84 ␮mol/m2 ) in the 600 mg/L manganite system. 3.2. XANES analysis for solid phase speciation Synthetic manganite was treated with both Sb(III) and Sb(V) for XANES analysis. Fig. 2 shows the normalized L(III)-edge XANES spectra of reference materials (Sb2 O3 and KSb(OH)6 ) and the samples collected from the Sb–manganite system. It is obvious from

Only Sb(V) adsorption kinetics and isotherms were examined because, as described above, Sb(III) can be rapidly oxidized to Sb(V) by manganite. Therefore, the results of Sb(V) adsorption should also be applicable to Sb(III). The sorption of Sb(V) onto manganite at three different temperatures (pH 4.0) is shown in Fig. 3. Similar to Sb(III) uptake by manganite, rapid adsorption of Sb(V) by the mineral was also apparent during the first 7 min. Removal efficiencies of 94.0%, 88.5% and 83.5% for the three temperatures were achieved, respectively, and remained slightly increased until the end of the

65 60 55

Sb(III) concentrations (μΜ)

Sb concentrations (μM)

40

30

50

273 K

298 K

313 K

45 40 35 30 25 20 15 10 5 0 0.00 0.25 0.50 0.75 1.00 1.25 1.50 1.75 2.00 2.25 2.50 2.75 3.00 3.25

Time (h) Fig. 3. Adsorbed Sb(V) on manganite as a function of time at different temperatures (adsorbent concentration: 600 mg/L; Sb concentration: 38.7 ␮M).

X. Wang et al. / Chemie der Erde 72 (2012) S4, 41–47

3.5

700

3.0

600

273K 298K 313K

-1

2.5

Adsorbed Sb (μ mol g )

t/qt (min g/μ mol)

44

2.0 1.5 1.0 .5

pH 3.0 pH 7.0 pH 9.0

500 pH 7

400 300

pH 9

200 100

0.0 0

20

40

60

80

100

120

140

160

180

0

200

0

Time (min) Fig. 4. Pseudo-second-order kinetic plots at different temperatures.

three temperature runs (180 min). Sb(V) sorption onto manganite can be formulated as follows: Sb(aq) + Mn–OH(s) → Mn–Sb(s) , where Sb(aq) is the Sb(V) concentration in the aqueous phase, Mn–OH(s) is the available reactive surface of the media for Sb(V) adsorption (i.e., adsorption sites), and Mn–Sb(s) is the concentration of antimony in the solid phase. Sorption kinetics provide valuable insights into the reaction pathways and into the mechanism of sorption reactions. Because both the Sb(V) concentration and active sorption sites on manganite play important roles in sorption processes, we used a pseudo-second-order kinetic model to clarify the adsorption kinetics of Sb(V) onto manganite at three temperature conditions. The form of the pseudo-second-order model is as follows: t 1 1 = + t, qt qeq kq2eq where k (g ␮mol−1 min−1 ) is the rate constant of the second-order equation, qt (␮mol g−1 ) is the amount of adsorption at time t (min), and qe is the amount of adsorption at equilibrium (␮mol g−1 ). Our results indicate that the kinetics of adsorption conform to the pseudo-second-order reaction very well (Fig. 4) with a correlation coefficient of R2 = 1 (Table 1). 3.4. Sb(V) adsorption thermodynamics Thermodynamic parameters such as enthalpy change [H◦ ] and free energy change [G◦ ] for the sorption reaction are estimated using the following equations: G◦ = −RT Ln Kd , −H ◦ Ln Kd = + constant RT where G◦ is the change in free energy, kJ/mol; H◦ is the change in enthalpy, kJ/mol; R is the gas constant (8.3145 J/mol K); and Kd is the equilibrium constant (adsorption or distribution coefficient) which at infinite dilution can be written as: Kd =

Sb(V)

pH 3

[Mn–Sb(s) ] [Sb(aq) ]

.

A plot of Ln Kd versus 1/T yields a straight line, and H◦ can be estimated from the slope of the line. The Kd values and thermodynamic parameters for Sb(V) adsorption summarized in Table 1 reveal that increasing temperature decreases the values of Kd . The

100

200

300

400

500

Dissolved Sb (μM) Fig. 5. Adsorption isotherms of Sb(V) at different pHs (solid line, Langmuir model; dotted line, Freundlich model).

negative sign of H◦ (−7.0 kJ/mol) signifies that the adsorption of Sb(V) onto manganite is an exothermic processes; i.e., the adsorption capacities decrease with increasing temperature. The Gibbs free energy changes (G◦ ) were calculated to be −22.4, −21.5 and −20.6 kJ/mol for 0, 25 and 40 ◦ C, respectively. The negative G◦ indicates that the adsorption is spontaneous, and the increase in G◦ with temperature implies less efficient adsorption at higher temperatures. 3.5. Sb(V) sorption equilibrium Sorption isotherms on manganite were generated to estimate maximum sorption density. The experimental adsorption data fitted both Langmuir and Freundlich isotherms very well (R2 > 0.93) (Fig. 5), and the corresponding sorption parameters were derived by plotting the linear forms of the isotherms (Table 2). The total site density of manganite was constant, and adsorption varied with the types of adsorbates and the composition of the water (Bang, 2003). The predominant crystal faces of manganite consist of >Mn–OH and Mn > OH2 (without >Mn–O) over a wide pH range and are therefore responsible for oxyanion adsorption (Foster et al., 2003). The predominant crystal plane (0 1 0) hosts 7.9 sites/nm2 (Ramstedt et al., 2004; Zhu et al., 2009). The surface site density (sites Sb/nm2 ) was calculated from the maximum adsorption density using the following mass balance: D=

Qmax × NA , S

where Qmax is the maximum adsorption density of Sb species on manganite, NA is Avogadro’s constant, and S is the surface area of the manganite mineral particles. When equilibrium pH values were 3.0, 7.0 and 9.0, the maximum sorption densities for Sb(V) were 784, 711 and 635 ␮mol g−1 manganite, which corresponded to 93.2%, 84.5% and 75.5% of the Mn–OH site density, respectively. Table 3 compares the adsorption capacity of manganite for Sb(V) with that of various sorbents reported in the literature (Ambe, 1987; Thanabalasingam and Pickering, 1990; Hasany and Chaudhary, 1996; Deorkar and Tavlarides, 1997; Madrid et al., 1998; Watkins et al., 2006; Biswas et al., 2009; Sari et al., 2010). As shown in Table 2, manganite has important potential for the adsorption of Sb(V) from aqueous solutions at wide pH ranges (pH 3.0, 7.0 and 9.0). The adsorption capacity of Sb(V) decreased with increasing pH. It is well known that anions tend to adsorb more strongly to a positively charged surface of metal oxides and

X. Wang et al. / Chemie der Erde 72 (2012) S4, 41–47

45

Table 1 Thermodynamic parameters of the pseudo-second-order kinetics of Sb(V) adsorption on manganite at different temperatures. Temp. ◦ C

K (g ␮mol−1 min−1 )

qe (␮mol g−1 )

R2

Kd (g/ml)

0±2 25 ± 2 40 ± 2

0.0516 0.0236 0.0176

62.9 62.1 62.1

1 1 1

54,183 44,986 43,614

G◦ (kJ/mol)

H◦ (kJ/mol)

−22.4 −21.5 −20.6

−7.0

Table 2 Langmuir and Freundlich isotherm parameters for Sb(V) sorption on manganite. Langmuir constants −1

Qmax (␮mol g 3 7 9

)

784.53 711.24 635.48

Freundlich constants b (1 ␮mol

−1

)

0.019 0.003 0.002

hydroxides, i.e., a surface at pH below the pH of point of zero charge (pHpzc ) and that the adsorption capacity typically decreases with increasing pH (Stumm, 1992; Pena et al., 2005). Increasing pH results in a large number of Mn–OH sites, a higher density of negative surface sites, and a higher extent of hydration (Hingston, 1981). All of these factors disfavor Sb(OH)6 − adsorption onto manganite as pH increases.

3.6. Effects of ionic strength and other co-existing anions on Sb(V) adsorption by manganite Between pH 3 and 8, the sorption of Sb(V) was above 88% at all three ionic strengths (0.001, 0.01 and 0.1 M, Fig. 6). As discussed in the previous section, a decrease in sorption occurred when the pH increased. From a pH of 9 to 12, the adsorption decreased sharply. It is clear that a change in ionic strength from 0.001 to 0.1 M NaNO3 had little effect on the adsorption of Sb(V), which indicates that Sb(V) is bound to the surface of manganite as an inner-sphere surface complex. This finding agrees with the results of Scheinost et al. (2005), who has confirmed that the dominant surface interaction between Mn and Sb(V) is inner-sphere in soil by extended X-ray adsorption fine structure (EXAFS). Aqueous anions can compete with Sb anions for adsorption sites and influence the surface charge of the oxide surface. The latter effect will be significant only if a charged surface complex is formed on the manganite surface (Chiu and Hering, 2000). Adsorption experiments were performed with Sb(V) (at an initial concentration of 160 ␮M) in the presence of silicate, H2 PO4 − , CO3 2− and SO4 2− .

2

R

Kf

n

R2

0.964 0.996 0.981

49.51 7.98 5.00

2.01 1.52 1.45

0.985 0.989 0.992

22 20 18

Adsorbed Sb (μmol/g)

pH

16 14 12 10 8

I=0.001M I=0.01M I=0.1M

6 4 2 0 2

3

4

5

6

7

8

9

10

11

12

13

pH Fig. 6. Sorption edges of Sb(V) on manganite at different ionic strengths.

The effects of the four oxyanions at three concentration levels (0.2, 1.0 and 10 mM) and a fixed pH of 6.3 ± 0.1 are illustrated in Fig. 7. The addition of CO3 2− and SO4 2− had no discernible effect on the overall amount of Sb(V) adsorption on manganite at pH 6.3. Silicate species are common oxyanions in natural water, with concentrations ranging from 0.45 to 14 mg/L Si (Clesceri et al., 2000). An understanding of silicate competition will benefit the development

400 Table 3 Comparison of Sb sorption capacity of manganite with that of different adsorbents (␮mol g−1 ). Sb(III)

Diatomite MnOOH

289 160

FeOOH

45.0

AlOOH

33.0

Goethite Zr(IV)-loaded SOW Fe(III)-loaded SOW Chemically bonded adsorbent Spiruline platensis Haro River sand ␣-Fe2 O3 Humic acid Manganite

500 ± 65 940 1120 180

Sb(V)

Reference Sari et al. (2010) Thanabalasingam and Pickering (1990) Thanabalasingam and Pickering (1990) Thanabalasingam and Pickering (1990) Watkins et al. (2006) Biswas et al. (2009) Biswas et al. (2009) Deorkar and Tavlarides (1997)

Adsorbed Sb (μmol/g)

Sorbents

399

0 mM

0.90 64.5 57.5 7.77 784.5

Madrid et al. (1998) Hasany and Chaudhary (1996) Ambe (1987) Pilarski et al. (1995) This study

1mM

10mM

398

397

396

395 1.23

0.2mM

Silicate

3-

2-

PO4

CO3

2-

SO4

Anions Fig. 7. Effects of co-existing anions on Sb(V) adsorption at a fixed initial Sb(V) concentration (160 ␮M) (pH 6.3 ± 0.1, 400 mg/L suspension).

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of effective treatment processes for Sb removal (Bang, 2003). However, the addition of silicate species caused only a slight decrease in the manganite adsorption efficiency of Sb. A small decrease was observed when the silicate concentration was as high as 10 mM, but the total adsorption was still over 99%. Adsorbed Sb(V) concentrations were slightly decreased by the presence of phosphate, which is consistent with a competitive adsorption process. Phosphate always competes with other anions such as arsenate anions in natural solids, biological and water treatment systems (Smedley and Kinniburgh, 2002; Roberts et al., 2003). The reason that phosphate slightly impacted the adsorption of Sb could be that phosphate formed inner-sphere complexes on manganite and impacted the adsorption of Sb(V). However, even in the presence of a competing adsorbate, manganite is an effective adsorbent for Sb(V). 4. Conclusions Kinetic studies of Sb(III) reacted on the manganite surface showed that equilibrium was achieved within 15 min. Rapid oxidation of Sb(III) by manganite was apparent as well as formation of soluble Sb(V) in a few minutes. The oxidation of Sb(III) by manganite occurs on the time scale of minutes. XANES analyses showed that Sb(V) was the dominant species adsorbed on manganite, suggesting that manganite acts as a strong oxidant toward Sb(III). The effects of temperature on the adsorption rates of Sb(V) were investigated. The adsorption kinetics conform to the pseudo-second-order reaction very well. Thermodynamic parameters showed that the adsorption of Sb by manganite is an exothermic process that is spontaneous at the specific temperatures investigated. The four tested anions (silicate, PO4 3− , CO3 2− and SO4 2− ) were not effective competitors with Sb(V) for sorption sites and the Sb(V) adsorption was not significantly hampered even when these anions occurred in relatively high concentrations. A change in ionic strength from 0.001 to 0.1 M had little effect on the adsorption of Sb(V). Acknowledgements This work was supported by the National Natural Science Foundation of China (40873077, 21177011), Nonprofit Environment Protection Specific Project (201009037-06, 201209013) and the Program for Changjiang Scholars and Innovative Research Team in University (No. IRT0809). References Ambe, S., 1987. Antimony(V) sorption and mobility in calcareous soils. Langmuir 3, 489–493. Bang, S., 2003. Effects of anions on arsenic adsorption by iron hydroxides. Ph.D., Stevens Institute of Technology New Jersey, pp. 94. Belzile, N., Chen, Y.-W., Wang, Z., 2001. Oxidation of antimony(III) by amorphous iron and manganese oxyhydroxides. Chemical Geology 174, 379–387. Biswas, R.K., Inoue, J., Kawakita, H., Ohto, K., Inoue, K., 2009. Effective removal and recovery of antimony using metal-loaded saponified orange waste. Journal of Hazardous Materials 172, 721–728. Cai, J., Salmon, K., DuBow, M.S., 1998. A chromosomal ars operon homologue of Pseudomonas aeruginosa confers increased resistance to arsenic and antimony in Escherichia coli. Microbiology 144, 2705–2729. Chiu, V.Q., Hering, J.G., 2000. Arsenic adsorption and oxidation at manganite surfaces. 1. Method for simultaneous determination of adsorbed and dissolved arsenic species. Environmental Science & Technology 34, 2029– 2034. Clesceri, L.S., Eaton, A.D., Greenberg, A.E., 2000. Standard Methods for the Examination of Water and Wastewater. American Public Health Association, American Water Works Association, Water Pollution Control Federation, Washington, DC. Cooper, W.J., Zika, R.G., Petasne, G., Fischer, A.M., 1998. Sunlight-induced photochemistry of humic substances in natural waters: major reactive species. In: Aquatic Humic Substances, vol. 219. American Chemical Society, pp. 333–362. Deorkar, N.V., Tavlarides, L.L., 1997. A chemically bonded adsorbent for separation of antimony, copper and lead. Hydrometallurgy 46, 121–135.

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