Journal of Colloid and Interface Science 445 (2015) 93–101

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Journal of Colloid and Interface Science www.elsevier.com/locate/jcis

Dispersion–precipitation synthesis of nanosized magnetic iron oxide for efficient removal of arsenite in water Wei Cheng a, Jing Xu a, Yajie Wang a, Feng Wu a, Xiuyan Xu b, Jinjun Li a,⇑ a b

School of Resources and Environmental Sciences, Hubei Key Lab of Bioresources and Environmental Biotechnologies, Wuhan University, Wuhan 430079, China China National Environmental Monitoring Centre, Beijing 100012, China

g r a p h i c a l a b s t r a c t

a r t i c l e

i n f o

Article history: Received 5 November 2014 Accepted 24 December 2014 Available online 6 January 2015 Keywords: Arsenic Adsorption Water purification Iron oxide Nanoparticle

a b s t r a c t Nanosized magnetic iron oxide was facilely synthesized by a dispersion–precipitation method, which involved acetone-promoted precipitation of colloidal hydrous iron oxide nanoparticles and subsequent calcination of the precipitate at 250 °C. Characterization by X-ray diffraction, transmission electron microscopy, Raman spectroscopy, nitrogen sorption, and vibrating-sample magnetometry revealed that the material was a composite of a-Fe2O3 and c-Fe2O3 with primary particle size of 15–25 nm and specific surface area of 121 m2/g, as well as superparamagnetic property. The material was used as adsorbent for the removal of arsenite in water. Batch experiments showed that the adsorption isotherms at pH 3.0–11.0 fit the Langmuir equation and the adsorption obeys pseudo-second-order kinetics. Its maximum sorption capability for arsenite is 46.5 mg/g at pH 7.0. Coexisting nitrate, carbonate, sulfate, chloride, and fluoride have no significant effect on the removal efficiency of arsenite, while phosphate and silicate reduce the removal efficiency to some extent. The As(III) removal mechanism is chemisorption through forming inner-sphere surface complexes. The efficiency of arsenic removal is still maintained after five cycles of regeneration–reuse. Ó 2015 Elsevier Inc. All rights reserved.

1. Introduction Arsenic contamination in natural water poses a significant threat to millions of people because of its high toxicity. Epidemiological studies revealed that long-term exposure to arsenic-contaminated drinking water could lead to serious health problems such as hyperkeratosis, peripheral vascular diseases, diabetes, anemia, cancers, as well as disorders of the immune, nervous, and reproductive systems [1–10]. The inorganic species, including As(III) (arsenite) and As(V) (arsenate), are the main forms ⇑ Corresponding author. E-mail address: [email protected] (J. Li). http://dx.doi.org/10.1016/j.jcis.2014.12.082 0021-9797/Ó 2015 Elsevier Inc. All rights reserved.

of arsenic in natural water. They generally coexist at certain proportions depending on the redox potential of the aqueous system [10]. In many regions, groundwater is the main source of drinking water, however, it is easily contaminated by arsenic due to leaching of As-containing aquifer rocks, and arsenite is often the dominant arsenic species in groundwater because of the reduced condition [11]. Arsenite is well known to be much more toxic than arsenate [4,10]. Therefore, arsenite must be removed before groundwater is delivered into the water-supply network. Among water purification techniques such as adsorption, coagulation, ion exchange, and membrane processes, adsorption stands out as the most common and promising technology for arsenic removal because of its simplicity and high efficiency [2,7,10].

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Among various adsorbents, iron oxide-based adsorbents [2,5,7–9] have attracted much attention because of their low cost, nontoxicity and great affinity for the arsenic adsorption. In particular, nanostructured iron oxide have significant advantages in sorption applications because of the high specific surface areas associated with their small sizes [2,12]. However, the spent nanostructured adsorbents are often difficult to be separated from aqueous solutions. In recent years, the use of magnetic adsorbents in water purification has stimulated great interest because these materials can be easily separated from aqueous systems by applying external magnetic fields [13]. Maghemite (c-Fe2O3) and magnetite (Fe3O4) are the mostly commonly used magnetic materials [4,8,14,15]. These oxides can be directly used as adsorbents [8], surface-functionalized with organic moieties [14], or incorporated into other porous substrates to provide magnetic properties [15]. Magnetic iron oxides are generally prepared by co-precipitation [4,16–18] or by solvent-thermal method [8]. The ordinary precipitation of ferric and ferrous salts by using alkali [4,16–18] is hard to control the particle size of the product. The solvent-thermal method involves heating solvent to high temperatures, and sometimes expensive and toxic precursors are used. Thus easy and economic methods for the preparation of nanosized magnetic iron oxides are still desirable. Some studies revealed that arsenite is more difficult to be removed from water than arsenate because the former mainly exists as nonionic H3AsO3 at a wide pH range and has low affinity to adsorbent surfaces [4]. In many studies, therefore, pre-oxidation of arsenite to arsenate was performed before adsorption process [4,15,19,20]. For example, when Zhang et al. [20] studied the removal of arsenite by Fe–Mn binary oxide adsorbent, they found that the oxidation of arsenite by manganese oxide promotes arsenic uptake by iron oxide. Recently, a magnetic c-Fe2O3–TiO2 nanocomposite was used to treat arsenic-contaminated water [4]. In that work, treatment involved photocatalytic oxidation of arsenite by TiO2 under UV light and sorption of arsenate on cFe2O3. However, an additional pre-oxidation step might mean more complex process and higher cost. Therefore, nanosized magnetic iron oxides that can be directly used as adsorbent for arsenite without any pre-oxidation step are meaningful. Recently, we reported a dispersion–precipitation synthesis of nanosized manganese oxide [21,22], in which a stable colloidal dispersion containing nanoparticles of hydrous manganese oxide was first prepared, and then the nanoparticles were precipitated by diluting the colloidal dispersion with water to induce a pH increase. The objective of this present work was to further develop a facile dispersion–precipitation method to prepare nanosized iron oxide for removal of arsenite in water. This method involved acetone-promoted precipitation of hydrous iron oxide nanoparticles in colloidal dispersion and subsequent calcination. It is interesting that the nanosized iron oxide was a composite of hematite (aFe2O3) and maghemite (c-Fe2O3), and thus showed superparamagnetic property. Furthermore, results of the batch sorption experiment suggest that it had high adsorption capacity toward arsenite. Since the dispersion–precipitation method is free of toxic chemicals, easily operated and economic, we think it could be an alternative choice to synthesize magnetic iron oxide adsorbents for the adsorption of arsenite.

2. Experimental section 2.1. Materials Iron nitrate nonahydrate, sodium hydroxide, acetone, glacial acetic acid, and hydrochloric acid were purchased from Sinopharm Chemical Reagent Company Limited. Sodium arsenite (NaAsO2)

was obtained from Chengdu Yikeda Chemical Reagent Corporation. All reagents were of analytical grade and were used as received. 2.2. Synthesis of nanosized magnetic iron oxide A colloidal dispersion of hydrous ferric oxide was prepared through a method reported recently with minor modification [23,24]. 16.16 g of iron nitrate nonahydrate was dissolved in 180 mL of ultrapure water, and 6.4 g of sodium hydroxide was dissolved in 120 mL of ultrapure water. Both solutions were then mixed to form a hydroxide slurry. After centrifugation, the hydroxide precipitate was collected, and 6.9 mL of glacial acetic acid was added to it. The resulting slurry was stirred until a colloidal dispersion of hydrous ferric oxide was formed. 400 mL of acetone was added to the above colloidal solution and precipitation occurred immediately. The precipitate was collected by centrifugation, and left in a fume hood overnight to evaporate residual acetone, and then dried at 60 °C. The solid was calcined in a muffle furnace at 100–500 °C for 2 h. The obtained product was labeled as FeMag-T, where T is the calcination temperature (°C). 2.3. Characterizations Fourier transform infrared spectroscopy (FTIR) at room temperature was performed on a Nicolet 5700 (USA). Spectral scans were done from 400 to 4000 cm1 at 4 cm1 resolution. X-ray diffraction (XRD) patterns were collected by using a PANalytical X’Pert Pro X-ray diffractometer with Cu Ka radiation (c = 1.5406 Å) within a 2h range of 10–80°. Raman spectra were collected by confocal Raman microspectroscopy at room temperature on a Reinshaw RM-1000 (Britain) at excitation wavelength of 514.5 nm. An argon laser source was used. Transmission electron microscopy (TEM) images were taken on a JEOL-2100 with an acceleration voltage of 200 kV. Lattice resolution of the instrument was 0.23 nm and it was equipped with a Gatan Orius bottom-mount Slow Scan CCD Camera. Brunauer–Emmett–Teller specific surface areas (SBET) were determined on a Micromeritics ASAP 2000 surface area analyzer. Vibrating-sample magnetometry was conducted on a (PPMS)-9T physical property measurement system to determine the magnetic properties of the iron oxides. X-ray photoelectron spectra (XPS) of the samples were collected on an ESCLAB 250Xi spectrometer with Al Ka radiation. C1s (Eb = 284.6 eV) spectrum was used as an inner standard calibration. 2.4. Batch adsorption experiment A stock solution of NaAsO2 with As(III) concentration of 1 g/L was prepared, and then a series of solutions with arsenic concentrations from 1 to 70 mg/L were prepared by diluting the stock solution with ultrapure water. To obtain the isotherm of arsenite adsorption on iron oxide at pH values from 3.0 to 11.0, the arsenite solutions were adjusted to the desired pH by using hydrochloric acid or NaOH solution. In the adsorption experiment, 10 mL of each arsenite solution was transferred to a 25 mL vial containing 4 mg of adsorbent. The vial was sealed by a cap and then shaken continuously on a platform shaker at 250 rpm and 25 °C for 12 h to ensure that adsorption reached equilibrium. The solid and liquid phases were then separated by centrifugation, and the equilibrium concentrations of arsenite in the supernatant solutions were measured by an 8220 atomic fluorescence spectrophotometry (Beijing Jitian Instrument Company, China) with a total lamp current of 80 mA. High-purity argon was used as carrier and shielding gas (supplied at 300 and 800 mL/min, respectively). The arsenite solutions were diluted with 1.5% hydrochloric acid before measurement, and a working solution was prepared by adding mixture of

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1.5% KBH4 and 0.5% KOH as reducing solution and 1.5% hydrochloric acid as carrier solution. Under this condition, As(III) was converted to AsH3 and then detected by the instrument. The quantity of adsorbed As(III) at equilibrium (qe, mg/g) was calculated as Eq. (1):

ðC 0  C e Þ V m

FeMag-500 FeMag-300

ð1Þ

where C0 (mg/L) and Ce (mg/L) represent the initial and equilibrium concentrations of As(III) in solution, respectively; m is the weight (g) of the adsorbent used and V (L) is the solution volume. The adsorption kinetics of As(III) at pH 7.0 was also studied. The initial concentration of As(III) was fixed at 1 mg/L, 5 mg/L and 10 mg/L, respectively. The dosage of adsorbent was 0.4 g/L, and the total reaction time was 8 h. The solution was sampled after a specified time interval, the solid and liquid phases were separated by centrifugation and the amount of residual As(III) in the solution was measured. To study the influence of naturally coexisting ions on arsenite adsorption, sodium nitrate, sodium carbonate, sodium silicate, sodium sulfate, sodium hydrophosphate, sodium chloride, and sodium fluoride were added to 1 mg/L (0.013 mmol/L) As(III) solutions, respectively. The concentration of each salt was l mmol/L. The pH values were adjusted to 7.0, and then adsorbent was added to each solution at a dosage of 0.4 g/L. The effect of ionic strength on As(III) adsorption at pH 7.0 was also studied by varying NaNO3 concentration from 0.001 to 0.1 mol/L.

Absorbance

qe ¼

1552 14151385

FeMag-250

FeMag-200

FeMag-100

4000

3500

3000

2500

2000

1500

1000

-1

Wavenumbers (cm ) Fig. 1. FTIR spectra of the iron oxides calcined at different temperatures.

3700 cm1 in the spectra of FeMag-100 and FeMag-200 can be attributed to hydrogen-bonded hydroxyl groups on the nanoparticles [27]; the intensity of this band diminished greatly in other samples because of dehydroxylation at elevated temperatures. Fig. 2a shows the nitrogen adsorption–desorption isotherms of FeMag-250 and FeMag-300. Apparently, both samples have type IV

2.5. Regeneration and reuse experiments

3. Results and discussion

(a) Volume adsorbed (cm3 g-1)

Repeated adsorption–desorption experiments were conducted to study the regeneration and reuse of the adsorbents. The adsorption runs were performed at an initial As(III) concentration of 1 mg/L and an adsorbent dosage of 0.4 g/L. After adsorption, the spent adsorbents were soaked in 1 mol/L NaOH solution and then shaken on a platform shaker at 250 rpm for 6 h to achieve regeneration, and then collected by centrifugation, washed with ultrapure water, and finally dried in air at 100 °C. The regenerated adsorbents were directly reused in the next adsorption run without further calcination.

75

60

45

FeMag-250 FeMag-300

30

15

3.1. Synthesis and characterizations of iron oxides

0.0

0.2

0.4

0.6

0.8

1.0

Relative pressure

(b) 0.04

dV/dD (cm3 g-1 nm-1)

Recently, when we used water to dilute a colloidal solution that was prepared by dispersing fresh manganese hydroxide in acetic acid solution, precipitate of nanosized manganese oxide was obtained [21,22]. However, in the present work, the stable colloidal dispersion containing nanoparticles of hydrous ferric oxide did not show any precipitate upon dilution with water. Alternatively, precipitate immediately appeared upon addition of acetone. These observations suggest that the nanoparticles of hydrous ferric oxide were stable in water and unstable in acetone. The acetone-induced precipitate contained organic ligands because acetic acid and acetone can coordinate to iron species, as confirmed by their FTIR spectra (Fig. 1). The absorption band centered at ca. 1552 cm1 in the spectrum of samples calcined below 200 °C could be attributed to surface-coordinated acetate on the hydrous iron oxide nanoparticles [25]; and those at 1415 and 1385 cm1 could be ascribed to asymmetric- and symmetric-bending vibration of methyl groups, respectively [25,26]. These bands almost disappear in the spectra of samples calcined at the temperature above 250 °C, indicating that organic molecules decomposed during calcination of these samples. The broad band at 2700–

0.03

FeMag-250 FeMag-300

0.02

0.01

0.00

0

5

10

15

20

Pore diameter (nm) Fig. 2. (a) Nitrogen adsorption–desorption isotherms and (b) BJH pore size distribution curves of the materials.

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: γ -Fe2O3

Intensity (a.u.)

: α-Fe2O3

FeMag-500

FeMag-300

FeMag-250 FeMag-200

Fe3O4 (JCPDS No. 19–0629). Because c-Fe2O3 and Fe3O4 have very similar XRD patterns, Raman spectra were collected to further confirm the crystalline structure of the iron oxides, as shown in Fig. 4. Peaks at 225, 295, and 413 cm1 are assigned to a-Fe2O3 [29]. Peaks located at 350, 500 and 700 cm1 are attributed to c-Fe2O3 [30]. Typical peaks for Fe3O4 at 300,530 and 661 cm1 are not observed [30]. These results corroborate the presence of a-Fe2O3 and c-Fe2O3 in the samples. The XRD patterns further reveal that elevated calcination temperatures produced materials with sharper and intensified diffraction peaks, which was a result of particle growth. Moreover, the relative peak intensities of a-Fe2O3 to c-Fe2O3 apparently increase with increasing calcination temperatures, indicating that c-Fe2O3 tends to transform to a-Fe2O3 at higher calcination temperature. Fig. 5 shows the TEM images of FeMag-250. The lowmagnification TEM image (Fig. 5a) shows that the sample is composed of irregular nanoparticles with diameters of 15–25 nm.

FeMag-100

10

20

30

40

50

60

70

80

2-theta (degree) Fig. 3. XRD patterns and of the materials calcined at various temperatures.

700 500

413

350

225

Intensity (a.u)

295

isotherms with hysteresis loops, suggesting the presence of mesopores [28]. The Barrett–Joyner–Halenda (BJH) pore size distribution profiles are present in Fig. 2b, which shows pore sizes in the range of 2–6 nm. The detailed textural properties are summarized in Table S1. With increasing calcination temperature from 250 to 300 °C, the specific surface area (SBET) decreased from 121 to 99 m2/g. FeMag-500 has a small SBET of 16 m2/g, indicating calcination at 500 °C led to significant sintering of the material. Fig. 3 presents the XRD patterns of the materials calcined at different temperatures. No obvious diffraction peaks are present in the patterns of FeMag-100 and FeMag-200, indicating that their crystallites are very small. Actually, owing to their small particle sizes, FeMag-100 and FeMag-200 could be dispersed in water to form transparent colloidal dispersions that showed Tyndall effect (Fig. S1). Fig. S2 shows the TEM image of FeMag-200, which reveals that most of the particles are less than 5 nm. It is possible that at temperatures lower than 200 °C, the existing of organic ligands on the nanoparticles prevented them from agglomerating into large particles. In XRD patterns of samples calcined above 250 °C, characteristic peaks for a-Fe2O3 (JCPDS No. 33–0664) can be found, and other peaks can be assigned to c-Fe2O3 (JCPDS No. 39–1346) or

FeMag-500

FeMag-300

FeMag-250

200

300

400

500

600

700

800

-1

wavenumber (cm ) Fig. 4. Raman spectra of the materials calcined at various temperatures.

Fig. 5. (a) TEM and (b) HRTEM images of FeMag-250.

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The high-resolution TEM (HRTEM) image (Fig. 5b) shows different lattice fringes that represent various crystal faces. The lattice fringe spacing values of 0.27, 0.25, and 0.18 nm correspond to the (1 0 4), (1 1 0), and (0 2 4) crystal planes of a-Fe2O3, respectively. Those of 0.29, 0.25 and 0.14 nm can be assigned to the (2 2 0), (3 1 1), and (5 3 0) crystal planes of c-Fe2O3, respectively. The HRTEM image (Fig. 5b) confirms that FeMag-250 was a composite of a-Fe2O3 and c-Fe2O3, in consistent with the XRD patterns and Raman spectra. Furthermore, the HRTEM image shows that a-Fe2O3 and c-Fe2O3 were both present in the same particle, indicating that they might be chemically bonded together rather than occur as a physical mixture. Fig. 6a illustrates the field-dependent magnetization of FeMag250, in which the magnetization curve and demagnetization curve coincide and no hysteresis occurred. Both coercivity and remnant magnetization were zero, indicating a superparamagnetic characteristic. Under the experimental conditions, the maximum magnetization was about 20 emu/g, which is strong enough to afford magnetic separation from aqueous system (Fig. 6b). Such feature is convenient for application in adsorption because magnetic recovery can be employed.

pH 7 pH 9

As(III) adsorbed (mg/g)

50

pH 5

40

pH 3 30

pH 11 20 10 0 0

10

20

30

40

50

Fig. 7. Isotherms of arsenic adsorption on FeMag-250 at various pH values.

50 mg/g. The arsenic adsorption could be affected by the surface property of the adsorbent and the existing species of the adsorbate, which are both pH-dependent. H3AsO3 has a pKa of 9.2, and at pH below 9.2, As(III) is mainly present as neutral H3AsO3, while at pH above 9.2, H2AsO 3 dominates [31]. Iron oxides possess surface hydroxyl groups, which can be protonated or deprotonated in solution depending on the pH [32]. Measurements of the zeta potentials at various pH values showed that the point of zero charge (PZC) of FeMag-250 was about pH 7.5, indicating that the adsorbent is likely to be protonated and therefore positively charged below pH 7.5, while deprotonated and negatively charged at above pH 7.5. Interaction/reaction between hydroxyl groups on adsorbent and those on arsenic species might play important roles in adsorption via forming either inner- or outer-sphere complexes [33]. The surface of FeMag-250 could retain a large proportion of unaltered hydroxyl groups at pH 7.0 since this pH is near the PZC of the adsorbent. This condition might favor adsorption of arsenic in the form of H3AsO3. Lowering of the sorption capacity with decreasing pH value may be attributed to the protonation of the surface hydroxyl groups on the adsorbent, which leads to fewer unaltered hydroxyl groups on the adsorbent surface. At higher pH, deprotonation of the hydroxyl groups lead to a negatively charged surface while arsenite tends to exist as an anion, and therefore electrostatic repulsion between arsenite anion and the negatively charged adsorbent surface also impedes arsenic adsorption. Two popular adsorption models, Langmuir and Freundlich models, were used to fit the isotherm data. The Langmuir model is based on the hypothesis that the adsorbent surface is homogeneous and adsorbs only monolayer adsorbates without involving intermolecular forces. Adsorption finally reaches a dynamic equilibrium, which can be represented as follows:

3.2. Effect of calcination temperature on As(III) removal Batch adsorption of As(III) on the adsorbents calcined at various temperatures were studied at an initial arsenite concentration of 5 mg/L and an adsorbent dosage of 0.4 g/L. The amount of adsorbed As(III) on FeMag-250, FeMag-300 and FeMag-500 were 11.4, 10.9 and 3.3 mg/g, respectively. This reflects that on the whole a higher specific surface area favors the arsenic uptake. However, if the amounts of As(III) adsorbed on the iron oxides were divided by their respective specific surface areas, we would found that that the As(III) uptake on the surface of FeMag-250, FeMag-300 and FeMag-500 were 0.09, 0.11 and 0.20 mg/m2, respectively. This indicates that increasing calcination temperature actually enhances the surface affinity to arsenic; however, it induces sintering of particles, which limits the total amount of adsorbed arsenic. As reflected by XRD, higher temperature promotes transformation of c-Fe2O3 to a-Fe2O3. It is possible that a-Fe2O3 is more favorable for arsenic uptake than c-Fe2O3. 3.3. Effect of solution pH on As(III) removal and sorption isotherms Fig. 7 shows the isotherms for As(III) adsorption on FeMag-250 at various pH values. Evidently, the pH had a profound effect on arsenite adsorption. Overall, FeMag-250 had high sorption capacities (above 20 mg/g) within a broad pH range (from 3.0 to 11.0). At pH 7.0, it manifested the highest sorption capacity, approaching

10

10

8

8

6

6 0

4

4

2

2

(a)

Magnetization (emu/g)

20

10 024681

0

-10

-20 0

-20000

-10000

0

0

10000

60

As(III) equilibrium concentrations (mg/L)

20000

Applied magnetic field (Oe) Fig. 6. (a) Magnetization curves of FeMag-250 and (b) illustration of the magnetic separation of FeMag-250 from water by a magnet.

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Qe ¼ Qm

bC e 1 þ bC e

ð2Þ

whose linearized form can be written as

1 1 1 ¼ þ Q e Q m bQ m C e

ð3Þ

where Ce (mg/L) is the equilibrium concentration of As(III) in solution, Qe (mg/g) is the amount of adsorbed As(III) at equilibrium, Qm (mg/g) is the maximum uptake capacity, and b (L/mg) is the Langmuir adsorption constant. The Freundlich model describes adsorption on heterogeneous surfaces. Its empirical formula is

1 lnQ e ¼ lnK f þ lnC e n

ð4Þ

where Qe (mg/g) represents the amount of adsorbed As(III), and Ce (mg/L) is the equilibrium As(III) concentration of the solution. The Freundlich constant Kf (mg/g) denotes the adsorption capacity of the adsorbent, and the constant n indicates the effect of concentration on the capacity and represents the adsorption intensity [34]. By fitting the experimental data with the above models, the constants and correlation coefficients (R2) are derived (Table 1). Comparison of the R2 values revealed that the results fit better into the Langmuir equation than the Freundlich equation. The adsorption capacities at various pH values calculated from Langmuir equation are shown in Fig. S3. The calculated maximum arsenite uptake at pH 7.0 is 46.5 mg/g. To further understand the adsorption process, a dimensionless constant separation factor, RL, was calculated to determine the feasibility of the isotherm criterion [1], which is defined as

RL ¼

1 1 þ bC 0

ð5Þ

where b is the Langmuir constant (L/mg) and C0 (mg/L) is the initial arsenic concentration. The value of RL indicates whether the adsorption is favorable. For favorable adsorption, 0 < RL < 1, while for unfavorable adsorption, RL > 1 [1]. In this study, RL values for various initial arsenic concentration was ranging from 0.02 to 0.72, indicating favorable adsorption of arsenic on FeMag-250.

3 5 7 9 11

3.4. Kinetics of As(III) adsorption To reveal the adsorption mechanism and to evaluate the rate of arsenite uptake, an investigation of AS(III) adsorption kinetics was conducted, the adsorption kinetics of As(III) on FeMag-250 at various concentrations are shown in Fig. 8a. Both pseudo-first-order and pseudo-second-order kinetic models were fitted with the experimental data, and their linearized formulas are given in Eqs. (6) and (7), respectively [38].

lnðQ e  Q t Þ ¼ ln Q e  k1 t

ð6Þ

t 1 1 ¼ þ t Q t k2 Q 2e Q e

ð7Þ

where Qe (mg/g) is the amount of As(III) adsorbed at equilibrium and Qt (mg/g) is the amount of As(III) adsorbed at sampling time t (min); and k1 and k2 are rate constants of sorption. The initial sorption rate (h0) can be calculated using Eq. (8):

h0 ¼ k2 Q 2e

Table 1 Parameters of sorption isotherms. pH

To assess the As(III) removal performance of FeMag-250, the Qm values of FeMag-250 are compared with maximum sorption capacities of other adsorbents described in literatures (Table 2). These data indicate that FeMag-250 has relatively high sorption capacity. We noticed that some previously reported adsorbents, including Fe2O3@C, c-Fe2O3–TiO2 and ascorbic acid-coated Fe3O4, had higher specific surface areas than FeMag-250, whereas their sorption capacities were not higher than that of the latter [2,4,8]. These reported adsorbents were composites of iron oxide and other components, and the other components might play important roles in maintaining high specific surface areas of these materials. It is known that iron oxide could readily uptake arsenic via forming surface complexes. However, the other components might not have so strong affinity to arsenic. Therefore, the higher specific surface areas of a composite might not always mean more active sorption sites for arsenic on the surface. In another work, Fe–Mn binary oxide had a high specific surface area of 265 m2/g and showed a maximal As(III) adsorption capacity of up to 132.8 mg/g [36]. Manganese oxide was capable of inducing the oxidation of As(III), which could promote arsenic adsorption on the solid surface [37].

Langmuir isotherm

Freundlich isotherm

Qm (mg/g)

b (L/mg)

R2

RL range

Kf (mg/g)

n

R2

32.1 35.2 46.5 40.9 24.1

1.21 1.07 0.54 0.82 0.19

0.993 0.993 0.992 0.995 0.993

0.01–0.29 0.01–0.32 0.03–0.48 0.02–0.38 0.07–0.72

12.9 14.1 13.0 11.1 4.3

3.02 2.58 2.29 2.36 1.92

0.961 0.968 0.971 0.991 0.969

ð8Þ

The pseudo-first-order and pseudo-second-order plots for the As(III) adsorption on FeMag-250 are shown in Fig. 8b and c, respectively. By fitting the experimental data with the above models, the kinetic parameters and correlation coefficients (R2) are derived (Table 3). Fig. 8a reveals that the adsorbent rapidly takes up As(III) and the adsorption almost reaches equilibrium within 200 min, suggesting that FeMag-250 had strong affinity toward As(III). The adsorption of As(III) fits better into a pseudo-second-order rate

Table 2 Maximal As(III) adsorption capacity of FeMag-250 and those of related adsorbents in literatures. Adsorbents

Qm (mg/g)

Initial As(III) concentration (mg/L)

pH

Specific surface area (m2 g1)

Ref.

FeMag-250 Fe2O3@C Ultrafine Fe2O3 c-Fe2O3–TiO2 Iron oxide-coated biomass M-FeHT Ascorbic acid-coated Fe3O4 R/S/FeMn Fe–Mn binary oxide Fe3O4-RGO Fe–Mn binary oxide

46.5 29.4 10.5 33.0 0.9 0.1 46.1 13.5 1.7 13.1 132.8

1–70 0.2–30 0.1–0.5 0.1–50 0.1 0.1–1 0–70 5–150 0.05–5 3–7 0.5–50

7.0 – 6.5 7.0 6.0 9.0 5.0 3.3 7.0 7.0 5.0

121 877 – 154 2 – 179 48 15 117 265

This study [2] [3] [4] [5] [7] [8] [10] [19] [35] [36]

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(a)

16 1 mg/L 5 mg/L 10 mg/L

Q t (mg/g)

12

8

4

0 0

100

200

300

400

500

t (min) 250

150

1 mg/L 5 mg/L 10 mg/L

2

1mg/L 5mg/L 10mg/L

200

t / Qt (min.g.mg -1)

3

(c)

1

In (Q e -Q t)

(b)

100

0 -1 -2

50

-3

0 0

100

200

300

400

500

0

100

t (min)

200

300

400

500

t (min)

Fig. 8. (a) Kinetics of As(III) adsorption at different concentrations, (b) pseudo-second-order plots and (c) pseudo-first-order plots.

Table 3 Kinetic parameters for As(III) sorption.

1 5 10

Pseudo-second order

Pseudo-first order

Qe (mg/g)

k2 (g mg1 mim1)

h0 (mg g1 min1)

R2

Qe (mg/g)

k1 (mim1)

R2

2.3 9.0 14.5

2.01 ⁄ 101 2.84 ⁄ 103 1.96 ⁄ 103

0.105 0.231 0.412

0.998 0.997 0.999

1.6 6.1 13.7

1.70 ⁄ 102 7.63 ⁄ 103 1.13 ⁄ 102

0.982 0.975 0.977

kinetic model, which indicates that arsenite removal mechanism is chemisorptions [38]. Moreover, the calculated value of Qe is in good agreement with that obtained from experiment. The rate of As(III) adsorption was enhanced with the increase of the initial As(III) concentration. To further understand the sorption process, the Weber and Morris intraparticle diffusion model is applied to predict the rate-limiting step, which can be expressed as follows [6,10]:

16

12

Qt (mg/g)

Initial As(III) concentration (mg/g)

10 mg/L 5 mg/L 1 mg/L

8

4

Q t ¼ kp t0:5 þ c

ð9Þ

where Qt is the amount adsorbed at time t, kp is the intraparticle diffusion rate constant and c is the intercept. If the plot of Qt vs. t0.5 gives straight lines and passes through the origin, the adsorption process is controlled by the intra-particle diffusion; or the intraparticle diffusion is not the only rate-limiting step [6,10]. In this study the plots for intra-particle diffusion were multilinear and did not pass through the origin (Fig. 9), indicating that the mechanism of As(III) removal was multistep and both the film diffusion and intra-particle diffusion contributed to the rate-determining step [6,10]. The first sharper portion of the plot was attributed to the diffusion of arsenic through the solution to the readily available sites of FeMag-250 [10]. The second line described the gradual adsorption stage of arsenic to the macro- and mesopore of FeMag-250, where intra-particle diffusion was rate-limiting step [10]. When the initial As(III) concentration was 1 mg/L, the slope of the second line was very small, which indicates that almost all As(III) was adsorbed on the external surface of the

0 0

5

10

t

0.5

15

20

0.5

(min )

Fig. 9. Plots of the intraparticle diffusion model.

FeMag-250 during the first stage. With the increase of the initial As(III) concentration, the slope of the second line increased distinctly, indicating the enhanced contribution of intra-particle diffusion in the rate-limiting step. 3.5. Effect of coexisting Ions and Ionic strength on As(III) removal Arsenic removal by iron oxides is generally achieved by forming surface complexes at active sites on the surface of iron oxides [36]. In ground water, some other anions always co-exist and can also

100

W. Cheng et al. / Journal of Colloid and Interface Science 445 (2015) 93–101

80

80

Arsenic removal (%)

100

Arsenic removal (%)

(a) 100

60

40

20

0

60

40

20

blank CO32- SiO 2- SO42-HPO 2- Cl3 4

-

F

-

0 1

NO3

2

3

4

5

Cycle number

Co-existing ions

Fig. 11. Arsenic removal efficiency on regenerated FeMag-250.

Arsenic removal (%)

(b) 100

efficiency of As(III) removal hardly changed, implying an innersphere adsorption mechanism [36].

80

60

3.6. Mechanism of arsenite removal by FeMag-250 40

20

0 0.001M NaNO3

0.01M NaNO3

0.1M NaNO3

Ionic strength Fig. 10. Effect of (a) coexisting ions and (b) ionic strength on As (III) removal.

form surface complexes with iron oxides. Thus, it is essential to evaluate the competition between arsenic and common coexisting ions. As illustrated in Fig. 10a, nitrate, carbonate, sulfate, chloride, and fluoride ions hardly affected the efficiency of arsenic removal under the experimental conditions. In contrast, silicate and hydrophosphate decreased the efficiency of arsenic removal from 93.9% to 78.5% and 60.9%, respectively. This indicates that silicate and hydrophosphate slightly competed for adsorption sites of the adsorbent, which could be associated with their specific species in aqueous solution. The pKa1 and pKa2 values of silicic acid are 9.77 and 11.80, respectively. Concentrations of various species of silicate calculated on the basis of these acidity coefficients revealed that the dominant species of silicate at pH 7.0 was H2SiO3 (Fig. S4). Since the pKa1, pKa2, and pKa3 of phosphoric acid are 2.15, 7.20, and 12.38, the major species of phosphate at pH 7.0 is H2PO 4 (Fig. S5). Both H2SiO3 and H2PO 4 bear hydroxyl groups, which could compete with arsenite for adsorption sites. In particular, phosphorus and arsenic are located in the same main group; their anions are structurally similar and may thus strongly compete with each other. Nevertheless, even when the phosphate concentration (1 mmol/L) was much greater than that of arsenic (0.013 mmol/ L), the removal efficiency of the latter was still above 60%, indicating that FeMag-250 still holds promise as an arsenic adsorbent for water purification. Either inner-sphere or outer-sphere surface complexation is considered as mechanism for ion adsorption. The effect of ionic strength on adsorption can be used as an indication for identifying the mechanism [33]. Co-existing electrolyte can also form outersphere complexes through electrostatic forces, which could suppress the outer-sphere adsorption of arsenite due to competition [36]. In contrast, the inner-sphere surface complexation shows little sensitivity to ionic strength [33]. Fig. 10b indicates that ionic strength had no significant effect on the adsorption of As(III). When the concentration of NaNO3 increased from 0.001 to 0.1 mol/L, the

The studies of adsorption kinetics and ionic strength effect suggested that the As(III) removal process on FeMag-250 might be chemisorption via inner-sphere surface complexation. To further reveal the adsorption mechanism, XPS and FTIR spectra of the used adsorbent were recorded. The As(3d) binding energy of the used adsorbent is 44.5 eV, while that of sodium arsenite is 44.0 eV (Fig. S6), indicating that the arsenic anions or molecules are adsorbed onto iron oxide [39]. In the FTIR spectrum of the spent adsorbent (Fig. S7), a peak at 780 cm1 could be assigned to uncomplexed As(III)–O stretching vibrations. According to Pena et al. [40], the uncomplexed As(III)–O bond of dissolved arsenite species is at 795 cm1, and if inner-sphere complexes are formed, red shift of this band would occur because the formation of Fe– O–As bond would decrease the strength of the As(III)–O bond [40]; in contrast, outer-sphere adsorption that results in protonation of the solid surface would not cause such a red shift [40]. Therefore, the FTIR spectra also reveal that the mechanism of arsenite removal is inner-sphere surface complexation. 3.7. Efficiency of regeneration and reusability In practical sorption applications, the high efficiency of adsorbent regeneration and reusability is of great significance. When FeMag-250 with adsorbed arsenite is stirred with NaOH solution, arsenite anions could desorb from the surface of FeMag-250 nanoparticles through hydroxyl exchange. Electrostatic repulsion might also play an important role in this regeneration process [7]. Fig. 11 shows the efficiency of As(III) removal by FeMag-250 nanoparticles during five adsorption–regeneration cycles. The removal efficiency after five cycles could still reach 95%, and no significant loss in adsorption capacity occurred, indicating that it could afford cyclic use in a cost-effective manner. 4. Conclusion We adopted a facile dispersion–precipitation method to produce nanosized magnetic iron oxide. This method involves dispersion of ferric hydroxide in acetic acid solution to obtain a colloid, precipitation of colloidal particles of hydrated ferric oxide by acetone, and calcination at proper temperature. The product had high specific surface area, high uptake capability for As(III), and excellent regenerability. The maximum sorption capacity for As(III) is 46.5 mg/g at pH 7.0. Adsorption isotherms fit well into the

W. Cheng et al. / Journal of Colloid and Interface Science 445 (2015) 93–101

Langmuir equation, and sorption obeys pseudo-second-order kinetics. Coexisting phosphate and silicate anions hindered removal of As(III), whereas nitrate, carbonate, sulfate, chloride, and fluoride had no significant adverse effect. The effect of ionic strength on arsenic adsorption and FTIR spectra confirmed that As(III) removal mechanism is inner-sphere surface complexation. The magnetic iron oxide nanoparticles have potential as adsorbent for the removal of arsenite from water. Acknowledgments

[11] [12] [13] [14] [15] [16] [17] [18] [19] [20]

This work was supported by the National Natural Science Foundation of China (21477090, 21477092) and the Natural Science Foundation of Hubei Province (2014CFB182).

[21] [22] [23] [24]

Appendix A. Supplementary material Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.jcis.2014.12.082.

[25] [26] [27] [28]

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