Current Environmental Engineering, 2014, 1, 51-58 51
Nano-Adsorbent for Arsenates: Iron Oxyhydroxide Impregnated Microporous Activated Carbon George Z. Kyzas1,2, Eleni A. Deliyanni*,1, Sotiria I. Bele1 and Kostas A. Matis1
1Laboratory of General & Inorganic Chemical Technology, Department of Chemistry, Aristotle University of Thessaloniki, Thessaloniki GR 541 24, Greece 2Department of Petroleum and Natural Gas Technology, Technological Educational Institute of Kavala, Kavala GR 654 04, Greece
Abstract: Impregnation of activated carbon or oxidized activated carbon was carried out using iron(III) nitrate as starting solution and ammonia as precipitating agent, intending the removal from aqueous solution of arsenates. The high affinity of iron oxyhydroxide nanocrystals towards inorganic arsenic species (pollutant) is widely known. Activated carbon can provide high surface area for adsorption. The role of carbon surface chemistry and structural heterogeneity on iron oxyhydroxide, and thus on the adsorption of arsenate, was investigated. The results suggest that a microporous carbon surface can adsorb arsenates. The higher arsenate adsorption of C1Na sample (iron impregnated with Fe(III) nitrate as starting material and ammonia solution as precipitating agent) than the other samples. The study of the microporous materials was carried out by adsorption of nitrogen, FTIR spectroscopy, SEM images, XRD patterns, thermal analysis, and examination of typical laboratory adsorption isotherms.
Keywords: Activated carbon modification, oxidation, iron doping, arsenate removal, metal separation.
1. INTRODUCTION
The occurrence of elevated levels of arsenic (As) in soils and groundwaters can compromise soil and water quality. Arsenic is ubiquitous in the earth’s crust and is highest in marine shale materials, magmatic sulphides and iron oxides, where it occurs as arsenopyrite (FeAsS), realgar (AsS) and orpiment (As2S3). Oxidative weathering and dissolution of As-containing minerals form dissolved inorganic As(III) and As(V), which are transported in surface or groundwater and can be adsorbed on soil and sediment particles [1]. Though the As concentration of uncontaminated soil is normally less than 6 mg/kg, anthropogenic sources of As such as arsenical pesticides, fertilizers, mine drainage, smelter wastes and agricultural drainage water from certain arid regions can elevate the levels of As in soil and water. Arsenic in drinking water poses today a serious threat to the health of people in many countries. The common example constitutes certainly Bangladesh, where around 40 million people consume drinking water with arsenic concentrations exceeding the guideline values of the WHO (10 µg/L) and of the country itself (50 µg/L). The distribution between dissolved As(III) and As(V) is dependent on redox potential. One of the most used adsorbent materials for arsenic removal is activated carbon. Numerous works have been published presenting the adsorption capacity of various activated carbons regarding As(III) or As(V) ions from different type of waters (Table 1). It is nowadays widely known that both As(V) and As(III) are strongly adsorbed to iron oxide and perhaps, this was the
*Address correspondence to this author at the Laboratory of General & Inorganic Chemical Technology, Department of Chemistry, Aristotle University of Thessaloniki, Thessaloniki GR 541 24, Greece; Tel: +30 2310 997808; Fax: +30 2310 997859; E-mail: [email protected]
reason for the respective focus on such sorbents [11]. Iron oxide particles are environmentally benign and exhibit amphoteric adsorption behavior around neutral pH conditions. On the other hand, activated carbon can provide a high surface area for adsorption and strong dispersive forces. The combination of activated carbon and iron loading takes advantage of the strength of these two materials. It was shown that iron preloaded on an activated carbon surface has an increased affinity for arsenic ions. Thus, the effect of iron oxyhydroxide nanocrystal deposition (doping) in the pore system of carbon was investigated [12-15]. A meso/microporous activated carbon (oxidized and ferric iron impregnated) has been recently studied from our research group, and promising results (removals) were exported [16]. In another work, 7 mg As(V) per g adsorbent was reported [17]. It seems that the “secret” for a successful adsorbent, in this case, is the applied method of activated carbon preparation, mainly, that of its following iron doping and perhaps, the experimental conditions (i.e., solution pH etc.). Microporous activated carbons have been elsewhere used for the same scope [18]; 0.03 in the former, and 204.2 mg As/g carbon in the latter were the respective results.
2. EXPERIMENTAL SECTION
2.1. Materials
In this study, the under examination carbon was a microporous spherical activated carbon, Calgon (manufact-ured by the Calgon Company) and designated as C. Part of carbon was oxidized with 70% (v/v) HNO3, by adding 100 mL of nitric acid 70% (v/v) to 10 g of the carbon sample and placed in a glass beaker with a magnetic stirrer. The mixture was stirred for 4 h. In order to remove the excess of acid and the soluble products of surface oxidation, the carbon was
52 Current Environmental Engineering, 2014, Vol. 1, No. 1 Kyzas et al.
extensively washed in a Soxhlet apparatus to constant pH. The oxidized sample was designated as C1. Iron impregnated samples were prepared using a suspension of activated carbon or oxidized activated carbon in aqueous solution of Fe(III) nitrate (0.506 mol/L Fe3+) or chloride in a three-necked round-bottom flask, being placed in a thermostat at 298 K. 0.23 g/dm3 of NH4OH (precipitating agent) or ammonium carbonate were added dropwise, using a dosimetric pump (Metrohm 645 Multi-Dosimat) at a constant flow rate, of 0.15×10-4 dm3/s, until the final pH of the hydrolysis process was adjusted to 8 (measured with a Crison MicropH 2002 instrument). In order to achieve good mixing of the reactants and prevent a possible agglomeration of the gel, vigorous mechanical stirring at 620 rpm was applied. After the addition of the precipitating agent, the stirring was continued for at least 15 min. The product obtained was decanted in a dialysis tubing cellulose membrane (Sigma Co.) and placed in a bath of distilled water, so as the anions of the suspension to be removed by osmosis, through the membrane. The water of the bath was replaced until no more anions were detectable in it. The resulting cake, on the membrane surface, was freeze-dried by a bench-scale instrument (Christ Alpha 1-4), until the temperature of the frozen gel reached the ambient temperature. The samples prepared after the iron impregnation are referred with N for the samples with Fe(III) nitrate as starting material, following an “a” for ammonia solution as the precipitating agent. The samples exposed to arsenate have letter ‘A’ added to their names.
2.2. Methods
The As solutions were prepared from arsenate stock solution (1000 mg/L). The stock solution was prepared by
dissolving reagent-grade 4.1653 g of Na2HAsO4·7H2O (99.8% purity, AnalaR) into 1 L of distilled water. The batch experiments for arsenate removal from dilute aqueous solutions were carried out at ambient temperature, using deionized water and conical flasks (20 mL sample volume), agitated with a reciprocal shaker (160 rpm) for 24 h. This contact time allows equilibrium to be reached, as determined in separate experiments. The pH was measured and found to be about 8 for the iron-modified samples, after precipitation (so, it was kept at this value for all the iron-modified samples). At this pH range, arsenate exists mainly as H2AsO4- and HAsO4
2- [19]. The final pH in all cases was measured and found to change of about ±0.5. It is assumed that these changes do not affect the speciation of arsenate ions in the solutions. The adsorption isotherms, which represent the amount of arsenic adsorbed per gram of carbon, were obtained by varying arsenate concentrations under a fixed dose of adsorbent. For the adsorption experiments, adsorbent concentrations of 1 g/L were applied while the arsenate concentrations were from 10 to 150 mg/L. The residual arsenic (i.e., that remaining in solution after the application of solid/liquid separation of suspended solids by 0.45 µm membrane filtration) was chemically analyzed. The molybdenum blue method was followed, using a double-beam UV-visible spectrophotometer (Hitachi Model U-2000) according to the appropriate standard method, which is based on the development of the molybdate-arsenate complex. As(III) is oxidized with KBrO3 to As(V), which then reacts with ammonium molybdate forming arsenomolybdate salt. This salt (at about 90°C) in the presence of hydrazine sulfate, is reduced to a blue molybdate complex. The absorbance was measured at 800 nm. The samples, which have not been analyzed on the same day of
Table 1. Comparative Table for As(III) or As(V) Removal from Water Using Various Types of Activated Carbons
Nano-Adsorbent for Arsenates Current Environmental Engineering, 2014, Vol. 1, No. 1 53
adsorption experiments, were acidified to pH 1, with concentrated HCl and stored in acid-washed high-density polyethylene containers. All samples were analyzed within three days of collection. The resulted equilibrium data were fitted to the Langmuir (eq. 1), Freundlich (eq. 2) and Langmuir-Freundlich (L-F) (eq. 3) isotherm equations, expressed by the following equations [20-22]:
Qe =
QmKLCe
1+KLCe
(1)
Qe = KFCe1/n (2)
Qe = QmKLF Ce( )1/b
1+KLF Ce( )1/b (3)
where Qe (mg/g) is the equilibrium. As concentration in the solid phase; Qm (mg/g) is the maximum amount of adsorption; KL (L/mg) is the Langmuir adsorption equilibrium constant; KF (mg1–1/n L1/n g-1) is the Freundlich constant representing the adsorption capacity; n (dimensionless) is the constant depicting the adsorption intensity; KLF (L/mg)1/b is the Langmuir-Freundlich constant; b (dimensionless) is the Langmuir-Freundlich heterogeneity constant.
2.3. Characterization
The pH of a carbon sample provides information about the acidity or the basicity of the surface. A sample of 0.4 g of dry carbon powder was added to 20 mL of water, and the suspension was stirred overnight to reach equilibrium. Then the pH of the solution was measured. Nitrogen isotherms were measured using an AS1Win (Quantachrome Instruments) at 77 K. Before experiment, the samples were heated at 393 K and then outgassed at this temperature under a vacuum of 10-4 Torr to constant pressure. The isotherms were used to calculate the specific surface area (SBET), that was calculated from the isotherm data, using the Brunauer, Emmet and Teller model, volume of micropores, Vmic, volume of mesopores, Vmes, and total pore volume, Vt, calculated using the Density Functional Theory (DFT). Thermal analysis was carried out using a TA Instrument thermal analyzer (SDT). The instrument settings were at a heating rate of 10 K/min, and a nitrogen atmosphere with a 100 mL/min flow rate. For each measurement about 25 mg of a ground carbon sample was used. Fourier Transform InfraRed spectroscopy was performed from 4000 to 450 cm-1 with a Perkin-Elmer Spectrum 2000 spectrophotometer. Samples were ground with special grade KBr in a fixed ratio, in an agate mortar. The same amount of mixed powder was also used to prepare the pellet for FT-IR. Scanning electron microscopy (SEM) images were performed at Zeiss Supra 55 VP. The accelerating voltage was 15.00 kV. EDAX analysis was done at magnification 10 K and led to the maps of elements. To test the iron loading on carbon, 0.2 g of carbon was ashed at 600oC and then digested with 25 mL of concentrated hydrochloric acid. The
digestion solutions were analyzed for iron, by a Perkin Elmer A Analyst 400 Atomic Absorption Spectophotometer. Boehm titration technique was used to determine the content of oxygenated surface groups [16].
3. RESULTS AND DISCUSSION
The iron impregnated carbon samples were tested for arsenate adsorption to determine the adsorption capacities. The measured isotherms of arsenate adsorption for the initial and for the oxidized carbon impregnated samples are collected in Fig. (1). The fitting parameters calculated from the Langmuir model, are reported in Table 2 along with the iron content in the samples. Thus oxidation decreases the adsorption capacity of the initial carbon, likely as a result of ash removal [23], while the introduction of iron species to the samples has a positive effect on the adsorption of arsenate. The iron impregnated samples exhibit a good performance, especially for the oxidized sample, where the adsorption capacity of the C1Na sample is about two times better than that of the initial carbon C1. The K values, related to the energy of adsorption [4], are much higher on the non-oxidized samples than those on oxidized ones. This suggests differences in the mechanism of adsorption. Introduction of iron increases the K values for both C and C1. For CNa and C1Na the K values are the highest, which indicates the most heterogeneous surface from the point of view of the energy of the adsorption sites. The analysis of porosity and surface chemistry of the initial and exposed to arsenate adsorbents contributes to better understanding of the factors contributing to an enhancement in the arsenic removal capacity on the modified carbons. The isotherm of the activated spherical carbon Calgon (C) presented in the inset of Fig. (2), which is typical of type I (according to the IUPAC classification) and representative for microporous materials. Moreover, it exhibits a significant increase of nitrogen uptake at P/Po>0.95, suggesting the presence of external surface area and/or textural porosity. The N2 isotherm of the oxidized carbon sample (not shown for brevity) is similar with that of the C sample. The specific surface area (SSA) of the spherical carbon found to be 1168 m2/g, with a micropore volume (0.418 cm3/g) and relatively high macropore volume 0.414 cm3/g, owing to the textural properties of its particles. The parameters of the porous structure for the samples addressed in this paper were collected, calculated from nitrogen adsorption isotherms. The pore size distribution (from 5 to 100 Å) curves for the microporous carbon C and for the oxidized one, based on DFT calculations is presented in Fig. (2). The C sample exhibits a pore size distribution curve with three main maxima (i.e., at ~5, 8 and 12 Å). The estimation (DFT calculations) of the volumes for pores with diameter smaller or even larger of 10 Å, provided some more evidence of the porous structure of the carbons. A significant fraction of the micropore volume of the carbon (i.e., about 32%) is attributed to pores with average diameter larger than 10 Å (up to ca. 20 Å).
54 Current Environmental Engineering, 2014, Vol. 1, No. 1 Kyzas et al.
Modification with iron had no effect in the porosity, as presented by pore size distribution curves (based on DFT
calculations) and shown in Fig. (2). Indeed, there is a low probability that iron introduced to the surface would exist in
Fig. (1). Arsenic adsorption isotherms. Solid lines represent the fitting to Langmuir equation.
Table 2. Iron Content of the Impregnated Samples Determined Using AAS and Fitting Parameters for Langmuir, Freundlich, and L-F Adsorption Isotherms
Fig. (2). Pore size distribution for the samples studied.
0 10 20 30 40 50 60 700
1
2
3
4
5
6
7
8
9
C C1 CNa C1Na
Qe (m
g/g)
Ce (mg/L)
10 1000.00
0.02
0.04
0.06
0.08
0.0 0.2 0.4 0.6 0.8 1.00
250
500
750
N2 v
olum
e ad
sorb
ed
( cm
3 / g
/ STP
)
P / Po
C
Incr
emen
tal p
ore
volu
me
(cm
3 /g)
Pore width (A)
C C1 CNa
Nano-Adsorbent for Arsenates Current Environmental Engineering, 2014, Vol. 1, No. 1 55
pores, suggesting in this way its high dispersion on the surface. That dispersion can be responsible for the activity/high efficiency of this iron phase for reactions with arsenate. The good dispersion of iron on the surface of carbons is observed with EDAX maps as presented in Fig. (3a & b) for the CNa and C1Na impregnated sample, respectively. It is mentioned that the iron content in the modified samples is 2.9% for the initial carbon and 4.4% for the oxidized carbon support (CNa and C1Na, respectively) as was measured by AAS (Table 2). The spherical shape of C is presented in Fig. (3c). The analysis of the surface pH values presented in Table 3 indicates that oxidation increases the acidity of the C sample. Oxidation of Calgon decreases the pH by an increase in the oxygen content. That increase in the oxygen content is related to the formation of acidic groups on the carbon surface. The number of acidic groups (measured using Boehm titration) increased from 0.625 mmol/g for initial C sample to 0.935 mmol/g for oxidized C1 sample. Formation
Table 3. Surface pH Values of the Sorbents Before and After Adsorption of As(V).
pH pH-As
C 5.57 6.65
CNONH 5.85 6.86
C1 4.16 6.43
C1NONH 6.10 6.85
of acidic groups on the carbon surface suggests their involvement either in reaction of these groups with ammonium hydroxide or in iron deposition. The modification applied makes the surface more basic on the average for both carbon samples. The impregnation with iron increases the surface pH of the oxidized carbon (C1Na) of about 2 units. After exposure to arsenate, an increase in the average pH for the all samples was noticed. The oxidation of the surface by As(V) did not occur, which is opposite to the
(a) (b)
(c) (d)
Fig. (3). EDAX maps for (a) CNa and (b) C1Na; (c) SEM image of CNa; (d) EDAX of C1Na
56 Current Environmental Engineering, 2014, Vol. 1, No. 1 Kyzas et al.
case of our previous work. The samples with iron exposed to arsenate show a visible decrease in the acidity, which indicates reaction of iron oxide/hydroxides with arsenate anions and formation of salts. This was indicated by Lorenzen and co-workers [9] who suggested formation of Fe(AsO4)2, FeAsO4 x Fe(OH)3 and FeAsO4 x H2O. Mohan and Pittman proposed the following reaction [24]:
FeOOH + 3H2AsO4
− + 3H+ → Fe H2AsO4( )3+ 2H2O
(4)
The effect in surface chemistry on the final chemistry of the carbons modified with iron species and then on arsenate removal was also evaluated using thermal analysis. The differential thermogravimetric (DTG) curves were obtained in nitrogen where the weight loss is associated with the removal/decomposition of species present on the carbons surface (Fig. 4). The DTG curve of the initial C sample reveals three peaks at 80°C, ~200°C and ~700-800°C. The first peak (at about 80°C) is common for all carbon samples and is related to the removal of water [25]. Peaks at about 700 and 800°C could correspond to the evolution of CO attributed to the decomposition of phenol, ether, carbonyl and quinone functional groups. An appearance of the new peaks in a 200-500°C temperature range for the C1 carbon corresponds to the decomposition of carboxylic groups, carboxylic anhydrides and/or lactone groups [25] formed by oxidation of carbon surface with HNO3. For the iron-impregnated samples, the broad peak between 50 and 120°C representing removal of water is more intense than that for the initial C sample and expands to 150°C with maximum at 100°C. Since these samples were prepared with NH4OH as the precipitating agent, this is likely the result of removal of water and ammonia from the decomposition of ammonium ions or complexes. For these impregnated with iron samples, a significant part of the weight loss at temperatures above 750°C must be related to
the gradual dehydratation/dehydroxylation of iron oxyhydroxides. These oxyhydroxides are finally reduced in two steps to metallic iron between 750 and 900°C [12]. The effects of oxidation, iron impregnation and arsenate adsorption on the C carbon are also seen on the FTIR spectra collected in Figs. (5a, b). For all carbon samples, the band at about 1580-1590 cm−1 can be assigned to aromatic ring stretching vibrations highly conjugated with CO and/or to CO vibrations of carboxylates. The very broad band at 1100-1150 cm−1 is assigned to C–O stretching and O–H bending modes of alcoholic, phenolic and carboxylic groups and the weak band at about 1700-1725 cm−1 is assigned to CO vibrations of the carboxylic acid group. Finally, the weak band at about 1260 cm−1 can be assigned to S=O2 and C=S stretching due to the sulfur content of this carbon. For the oxidized samples not exposed to the adsorbate, an increase in the intensities of the bands at 1720 and 1100 cm−1, representing carboxylic acids and phenols respectively, is seen, while a band at about 1380 cm−1 can be attributed to HNO3 oxidation. For the iron impregnated carbons it is important to mention that the carbon surface was further oxidized during the impregnation process. That oxidation of carbon with Fe(NO3)3 results in the increase of the intensity of the band respresenting lactone, as seen in the FTIR spectrum for the CNa sample. The peaks (shown in Fig. 4b, for the CNa sample) at about 680 cm-1 is characteristic of the iron oxyhydroxide phase (akaganeite). To that iron phase, which can react with arsenates, could be attributed the better performance, since iron oxy-hydroxides can react with arsenates as illustrated in Fig. (6). The peaks for C1Na sample at about 871 cm-1 and 716 cm-1are characteristics of the iron oxyhydroxide phase (goethite) (Fig. 5b). The peak at 1400 cm-1, which represents the vibrations of nitrogen in NH4
+ inorganic ion on the
Fig. (4). Comparison of the DTG curves in nitrogen for the initial and exhausted samples.
0 200 400 600 800 10000.0
0.5
1.75
1.80
1.85
1.90
1.95
2.00
0 200 400 600 800 10000,0
0,1
0,2
0,3
0,4
Wei
ght l
oss
deri
vativ
e [% /
o C ]
Temperature [ oC ]
C C1
Wei
ght l
oss
deriv
ativ
e (%
/o C)
Temperature (oC)
CNa C1Na
Nano-Adsorbent for Arsenates Current Environmental Engineering, 2014, Vol. 1, No. 1 57
surface of very acidic carbon is not revealed. NH4+ for this
carbon. NH4+ is a Brönsted acid, can react with deprotonated
carboxylic type acidic sites to ammonium salts, get adsorbed as NH4
+ on the metal acidic sites while the presence of iron, favors Lewis acid-base interactions, offering in this way binding sites for NH4
+ ions [26]. On the other side, the ammonium sites present on the adsorbent surface of this sample have some affinity to the arsenate anions. At the pH range the adsorption experiments were performed (pH 8), arsenate species appear as monovalent H2AsO4
- and divalent HAsO4
2- ions [27] so the ammonium moieties are capable to
attract arsenate anionic species by electrostatic forces facilitating the arsenate adsorption via the following reaction:
≡ Fe −OH + R − NH4+ + HAsO4
2− →
≡ Fe −O − AsO3H....+ NH4 − R + H2O (5)
Opposite to the case of our previous work [16], for this carbon it seems like this reaction did not occur having as result the less arsenate removal. Furthermore, iron hydroxide nanoparticles are anchored to oxygen functional groups (-
(a) (b)
Fig. (5). FTIR spectra for the initial and exhausted samples.
58 Current Environmental Engineering, 2014, Vol. 1, No. 1 Kyzas et al.
COOH, -OH, -C=O) via chemical bond (Fig. 6), which inhibits their leaching to treated water. Another comment which has to be done is the possible competitiveness between As(V) and other ions. In groundwaters, pollutants other than As(V) (as carbonates, sulfates, phosphates, silicates, magnesium, manganese, etc) are present and can interfere in the removal of anionic arsenic species. Competitive adsorption of As(V) in the presence of co-existing ions has been determined in the literature and was found that PO3
4- and SO24- can affect the
As(III) removal because they have very similar chemical structure. On the contrary, silicate, manganese, magnesium, calcium or carbonate slightly reduce As(V) adsorption due to the different molecular structure of these species.
4. CONCLUSION
The results presented in this paper suggest that a microporous carbon surface can adsorb arsenates. However, a reduction was noticed this time in the capacity, to around 8 from the ~32 mg arsenic per g carbon that we obtained previously for the meso/microporous sample, with exactly similar experimental conditions. The higher arsenate adsorption of C1Na sample than the other samples can be attributed to the higher percentage of iron for this sample in the form of iron oxide/hydroxides that can react with arsenate anions. CONFLICT OF INTEREST
The authors confirm that this article content has no conflict of interest.
ACKNOWLEDGEMENTS
Declared none.
REFERENCES [1] Deliyanni EA, Peleka EN, Gallios GP, Matis KA. A critical review
of the separation of arsenic oxyanions from dilute aqueous solution (the contribution of LGICT). Int J Environ Waste Managem 2011; 8: 286-304.
[2] Huang CP, Fu P. Treatment of As(V) containing water by activated carbon. J Water Pollut Contr Fed 1984; 56: 232.
[3] Pattanayak J, Mondal K, Mathew S, Lalvani SB. Parametric evaluation of the removal of As(V) and As(III) by carbon-based adsorbents, Carbon 2000; 38: 589-96.
[4] Chuang CL, Fan M, Xu M, et al. Adsorption of arsenic(V) by activated carbon prepared from oat hulls, Chemosphere 2005; 61: 478-83.
[5] Gu Z, Fang J, Deng B. Preparation and evaluation of GAC-based iron-containing adsorbents for arsenic removal. Environ Sci Technol 2005; 39: 3833-43.
[6] Navarro P, Alguacil FJ. Adsorption of antimony and arsenic from a copper electrorefining solution onto activated carbon. Hydrometallurgy 2002; 66: 101-5.
[7] Eguez HE, Cho EH. Adsorption of asrenic on activated charcoal. J Metals 1987; 39: 38-41.
[8] Manju GN, Raji C, Anirudhan TS. Evaluation of coconut husk carbon for the removal of arsenic from water. Water Res 1998; 32: 3062-70.
[9] Lorenzen L, van Deventer JSJ, Landi WM. Factors affecting the mechanism of the adsorption of arsenic species on activated carbon. Miner Eng 1995; 8: 557-69.
[10] Budinova T, Petrov N, Razvigorova M, Parra J, Galiatsatou P. Removal of arsenic(III) from aqueous solution by activated carbons prepared from solvent extracted olive pulp and olive stones. Ind Eng Chem Res 2006; 45: 1896-901.
[11] Deliyanni EA, Bakoyannakis DN, Zouboulis AI, Matis KA. Sorption of As(V) ions by akaganéite-type nanocrystals. Chemosphere 2003; 50: 155-63.
[12] Deliyanni E, Bandosz TJ. Importance of carbon surface chemistry in development of iron-carbon composite adsorbents for arsenate removal. J Hazard Mater 2011; 186: 667-74.
[13] Torres-Perez J, Gerente C, Andres Y. Conversion of agricultural residues into activated carbons for water purification: Application to arsenate removal. J Environ Sci Health Part A Toxic Hazard Subst Environ Engin 2012; 47: 1173-85.
[14] Tuna AOA, Özdemir E, Şimşek EB, Beker U. Removal of As(V) from aqueous solution by activated carbon-based hybrid adsorbents: Impact of experimental conditions. Chem Eng J 2013; 223: 116-28.
[15] Zhu H, Jia Y, Wu X, Wang H. Removal of arsenic from water by supported nano zero-valent iron on activated carbon. J Hazard Mater 2009; 172: 1591-6.
[16] Deliyanni E, Bandosz TJ, Matis KA. Impregnation of activated carbon by iron oxyhydroxide and its effect on arsenate removal. J Chem Technol Biotechnol 2013; 88: 1058-66.
[17] Gu Z, Deng B, Yang J. Synthesis and evaluation of iron-containing ordered mesoporous carbon (FeOMC) for arsenic adsorption. Microporous Mesoporous Mater 2007; 102: 265-73.
[18] Fierro V, Muñiz G, Gonzalez-Sánchez G, Ballinas ML, Celzard A. Arsenic removal by iron-doped activated carbons prepared by ferric chloride forced hydrolysis. J Hazard Mater 2009; 168: 430-7.
[19] Lehmann M, Zouboulis AI, Matis KA, Grohmann A. Sorption of arsenic oxyanions from aqueous solution on goethite: A study of process modelling. Microchim Acta 2005; 151: 269-75.
[20] Langmuir I. The adsorption of gases on plane surfaces of glass, mica and platinum. J Am Chem Soc 1918; 40: 1361-403.
[21] Freundlich H. Over the adsorption in solution. Z Phys Chem 1906; 57: 385-470.
[22] Tien C. Adsorption Calculations and Modeling. Butterworth-Heinemann; Boston, USA 1994.
[23] Kolthoff IM, Elving PJ. Treatise in Analytical Chemistry. Interscience; New York, USA 1962.
[24] Mohan D, Pittman CU Jr. Arsenic removal from water/wastewater using adsorbents—A critical review. J Hazard Mater 2007; 142: 1-53.
[25] Bashkova S, Bandosz TJ. The effects of urea modification and heat treatment on the process of NO2: removal by wood-based activated carbon. J Colloid Interface Sci 2009; 333: 97-103.
[26] Huang CC, Li HS, Chen CH. Effect of surface acidic oxides of activated carbon on adsorption of ammonia. J Hazard Mater 2008; 159: 523-7.
[27] Przepiórski J, Skrodzewicz M, Morawski AW. High temperature ammonia treatment of activated carbon for enhancement of CO2: adsorption. Appl Surf Sci 2004; 225: 235-42.
Received: September 12, 2013 Revised: October 16, 2013 Accepted: October 17, 2013
