EnvironmentalScienceWater Research & Technology

PAPER

Cite this: Environ. Sci.: Water Res.Technol., 2015, 1, 96

Received 23rd July 2014,Accepted 19th September 2014

DOI: 10.1039/c4ew00020j

rsc.li/es-water

Phosphate removal using modified Bayoxide®E33 adsorption media†‡

Jacob Lalley,ab Changseok Han,ac Gayathri Ram Mohan,c Dionysios D. Dionysiou,b

Thomas F. Speth,a Jay Garlandd and Mallikarjuna N. Nadagouda*a

The adsorption of phosphate onto modified Bayoxide® E33 (E33) and underlying mechanisms were

comparatively investigated by batch kinetics, sorption isotherms, rapid small-scale column tests, and mate-

rial characterization. Synthesis of modified E33 was conducted by the addition of manganese and silver

nanoparticles on the surface of the goethite (α-FeOOH)-based E33. Upon successful adsorbent synthesis,

the materials were characterized using SEM, TEM, XRD, EDX, HR-TEM, and BET surface area analysis. The

adsorption kinetics is described by a fast initial adsorption stage followed by a slow adsorption stage,

complying with pseudo second-order kinetics and Elovich kinetics. Sorption isotherms revealed that the

equilibrium capacity of one of the modified adsorbents (E33/AgII) exceeded that of unmodified E33 likely

due to its increased BET surface area. Column experiments confirmed the results for adsorption kinetics

and equilibrium of phosphate sorption onto E33 and modified E33. This suggests that media modification

has the potential to improve phosphate removal properties.

1. IntroductionPhosphates are one of the common mineral nutrients foundin wastewaters.1 Both natural and anthropogenic sourcessuch as agricultural runoffs2,3 and urban4 and industrialwastes5,6 act as major sources of phosphate in surface waterbodies. In aquatic systems, phosphates occur in variousforms, particulate, organically bound, or dissolved,7 of whichorthophosphates !PO4

3!) are found to be predominant.8

While it is an essential nutrient, higher concentrations ofphosphate in water bodies can be detrimental to aquatic life,human health, and the environment.9 As a limiting nutrientin the environment, increased loads of phosphates can resultin accelerated eutrophication or excessive growth of algae,leading to increased treatment costs and reduced recreationalvalue.10 Therefore, phosphate removal from wastewaters hasbecome an area of emerging concern.

Conventional methods for phosphate removal includephysical, chemical, and biological processes; of which, chem-ical precipitation11 and enhanced biological phosphate

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a The U.S. Environmental Protection Agency, ORD, NRMRL, WSWRD, WQMB,26 W. Martin Luther King Jr. Drive, Cincinnati, OH, 45268, USA.E-mail: Nadagouda.mallikarjuna@epa.gov; Fax: +1 (513) 569 7172;Tel: +1 (513) 569 7232b Environmental Engineering and Science Program, Department of Biomedical,Chemical and Environmental Engineering (DBCEE), University of Cincinnati,OH, 45221, USAc Pegasus Technical Services Inc., 46 E Hollister St, Cincinnati, OH 45219, USAd The U.S. Environmental Protection Agency, ORD, NERL, MCEARD,26 W. Martin Luther King Jr. Drive, Cincinnati, OH, 45268, USA† The U.S. Environmental Protection Agency, through its Office of Research andDevelopment, funded and managed, or partially funded and collaborated in, theresearch described herein. It has been subjected to the Agency's administrativereview and has been approved for external publication. Any opinions expressedin this paper are those of the author(s) and do not necessarily reflect the viewsof the Agency, therefore, no official endorsement should be inferred. Anymention of trade names or commercial products does not constituteendorsement or recommendation for use.‡ Electronic supplementary information (ESI) available. See DOI: 10.1039/c4ew00020j

Water impact

As a limiting nutrient in the environment, increased loads of phosphates can result in accelerated eutrophication or excessive growth of algae, leadingto increased treatment costs and reduced recreational value. Therefore, phosphate removal from wastewater has become an area of emerging concern. Inthis study, the removal of phosphate using newly surface-modified Bayoxide® E33 was thoroughly examined. Underlying mechanisms were comparativelyinvestigated by batch kinetics, sorption isotherms, rapid small-scale column tests, and material characterization. The adsorption kinetics is well describedby a fast initial adsorption stage followed by a slow adsorption stage, complying with pseudo second-order kinetics and Elovich kinetics. The results showedpotential for practical applications of the modified material for phosphate removal.Pu

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removal (EBPR)12,13 are commonly used. However, due to anumber of factors including cost, sludge production,required energy, and issues of reliability and stability,14,15

other methods are currently being explored. The quickest,easiest, and most economically known method for theremoval of phosphate ions is adsorption.16 To date, a numberof different adsorbents have been investigated for phosphateremoval including metal oxides,16–22 red mud,23–25 steelslag,26–28 and fly ash.29,30 Table 1 shows different phosphateadsorbents and their reported adsorption capacity at differ-ent parameters. Adsorption also has the benefits of applica-bility at low concentrations, little to no sludge production,and potential for both batch and continuous processes, andthe potential for adsorbent regeneration that could reducewastes and costs.31

One such highly adsorptive material for phosphateremoval is the iron oxide goethite !iron!III) oxide-hydroxide).32–35

The affinity of phosphate adsorption on iron!III) oxide-hydroxide is dependent on phosphate's complexing capacityby ligand exchange as well as electrostatic interactionswith the surface.36 This process involves replacement of aphosphate ion with one or more surface hydroxyl groups,releasing surface structural OH2 and/or OH! into the bulksolution,37 in which stoichiometric changes in the amount ofphosphate adsorbed and OH! released can provide clues onthe type of adsorption mechanism.38 Iron!III) oxide-hydroxideadsorption initially occurs rapidly, followed by a slow stage,39

attributed to micro/mesopore diffusion40 or particle aggrega-tion,41 which is generally dependent on the crystallinity ofsamples.42 Phosphate adsorption onto iron oxides involves aligand exchange mechanism,43 yet the exact mode of phos-phate interaction is not known. IR studies revealed that eachadsorbed phosphate ion replaces two, singly coordinated OHsurface groups, indicating a bridging, binuclear complex.44,45

However, other studies have suggested that monodentate

surface complexes could exist at low adsorbate concentra-tions.46 Using a combination of surface complexes to modeladsorption on iron oxides, Geelhoed et al.47 concluded thatthe bidentate surface complex was most abundant at low tointermediate values of pH.

Goethite (α-FeOOH) is one of the most thermodynamicallystable forms of iron oxide and thus is available in abundantquantities.48 Due to its affinity for ion adsorption, AdEdgeTechnologies Inc. (USA) has produced Bayoxide® E33 media(E33), a commercially available goethite-based adsorbent forenvironmental remediation. Although originally designed forarsenic removal,49,50 E33 has been shown to adsorb a rangeof ions including phosphates. While several studies haveshown that E33 adsorption is an effective technique for phos-phate removal,51–53 surface modifications of E33 can enhanceoverall adsorption. As adsorption is a surface-based phenom-enon dependent on electrostatic interactions and covalentbonding, the addition of di- and tri-valent metals to the sur-face of E33 should correlate with increased phosphateadsorption.

In this study, the surface of E33 media was modified withmanganese and silver. For the removal of phosphate, theadsorptive properties of modified E33 were investigated andthen compared to those of the unmodified E33.

2. Experimental2.1 Materials

The E33 used to study PO43! removal was obtained from

AdEdge Technologies Inc. (USA). Analytical grade sodiumphosphate monohydrate !NaH2PO4·!H2O)) was obtained fromFisher Scientific and used as a stock solution. A stock phos-phate solution at a concentration of "80 mg L!1 was pre-pared by dissolving 1 g of NaH2PO4·!H2O) in 10 L of Milli-Qwater. The pH of the stock phosphate solution was

Table 1 Comparison of phosphate adsorption capacity of different adsorbents

Adsorbent Capacity !PO43! mg g!1) pH Temperature (°C) Reference

Metal oxides synthesized from manganese ore tailing 26.3 6.0 25 16Titanium ion doped ordered mesoporous silica 4.5 7.0 21 17Magnetic iron oxide nanoparticles 3.7 3.0 15 18Zinc–aluminum layered double hydroxides 30.0 9.0 20 19Iron oxide-coated crushed brick 1.8 5.0 20 20Lithium-intercalated gibbsite 0.3 4.5 25 21Fe-loaded biomass char 35.4 3.0 25 22HCl-treated red mud 0.6 5.5 40 23Red mud granular adsorbents 9.8 9.2 37 24Red mud 0.5 4.0 25 25Modified steel slag 11.1 8.0 25 26Steel slag 5.3 5.5 25 27Electric arc furnace (EAF) steel slag 2.2 6.0 23 28Iron oxide/fly ash 27.4 3.0 25 29Fly ash 18.9 3.0 25 29Nanshi-modified fly ash 40.7 4.0 21 30Nanshi-modified zeolite 47.2 4.0 21 30E33 26.8 7.0 21 This studyE33/Mn 17.8 7.0 21 This studyE33/AgI 19.3 7.0 21 This studyE33/AgII 28.0 7.0 21 This study

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maintained constant at 6, 7 and 9, using 2-!N-morpholino)ethanesulfonic acid (MES), 3-!N-morpholino) propanesulfonicacid (MOPS), and N-tris (hydroxymethyl) methyl-3-amino-propanesulfonic acid (TAPS), respectively, all obtained fromSigma Aldrich. Reagent grade silver nitrate (AgNO3), sodiumborohydride (NaBH4), and manganese acetate tetrahydrate!Mn!CH3COO)2·4H2O) were obtained from Sigma Aldrich andused as received.

2.2 Synthesis of modified E33

For the synthesis of silver-modified E33, an initial mixture of10 g of E33 and 35 mL of Milli-Q water was prepared in a100 mL conical flask. Then, 2 mL of 0.1 N AgNO3 was addedto the flask and shaken by hand for 2 min to ensure thor-ough mixing. To this initial mixture, two different adsorbentswere prepared, one using NaBH4 and the other using teaextract, denoted respectively as (E33/AgI)50 and (E33/AgII)50.For E33/AgI, 10 mL of 0.1 N NaBH4 was added to the initialmixture, hand shaken for 2 min and allowed to react over-night at room temperature. For E33/AgII, 2 g of tea powder(Red Label, India) was boiled in 100 mL of Milli-Q water andthe solution was filtered through a 0.20 μm PTFE filter. Then,20 mL of the tea extract solution was added to the initialmixture, shaken by hand for 2 min, and allowed to react over-night at room temperature. Following the overnight incuba-tion, samples were twice washed with Milli-Q water and air-dried at room temperature prior to characterization andanalysis.50

For the synthesis of manganese-coated E33 (E33/Mn), 50 gof E33 was mixed with 20 mL of Milli-Q water in a chinacrucible. Then, 10 g of manganese acetate tetrahydrate!Mn!CH3COO)2·4H2O, Aldrich product) was added to the cru-cible and the solution was hand-mixed with a spatula for5 min to ensure proper mixing. After mixing was complete,the sample was heated at 310 °C for 1 h in a china crucibleand then stored for later analysis and characterization.

2.3 Sample characterization

X-ray diffraction (XRD) analysis was performed using aPANalytical (X'pert) 2-theta diffractometer (PANalytical, Almelo,the Netherlands) in the 2-theta range of 2–90° under CuKα

radiation and a wavelength of 1.54 μm in order to determinethe adsorbent's crystal structure. An environmental scanningelectron microscope (ESEM, model Philips XL 30 ESEM-FEG)at an accelerating voltage of 30 kV was used to characterizethe surface morphology of the adsorbents. An energy-dispersive X-ray spectrophotometer (EDX) installed in theESEM was used to detect the distribution of elemental phos-phorus as well as the percent of phosphate adsorbed. Fortransmission electron microscopy (TEM), a FEI CM20 TEMoperated at 200 kV was employed. High-resolution TEM(HR-TEM, JEM-2010F (JEOL)) with a field emission gun at200 kV was also employed for the characterization of the pre-pared surface-modified adsorbents. Free software (ImageJ,National Institutes of Health, USA) was used for HR-TEM

image analysis. Adsorbent samples were prepared by ultrason-ication (2510R-DH Bransonic) for 20 min in isopropyl alcohol(99.8%, Pharmco), and then a drop of the supernatant wasfixed on a carbon-coated copper grid (FCF400-Cu, EMS) anddried at room temperature. A Tristar 3000 porosimeter ana-lyzer (Micromeritics) was used to determine the adsorbent'sBET surface area, total pore volume, and porosity. Sampleswere first purged with nitrogen gas at 150 °C for 2 h usingMicromeritics FlowPrep 060.

2.4 Adsorption studies

To determine various adsorption parameters of thesurface-modified adsorbent, both batch and flow-throughcolumn experiments were employed. Batch adsorption kineticstudies were carried out to determine the amount of phos-phate adsorbed with respect to time. Experiments wereconducted in 250 mL Nalgene polypropylene bottles giventhe following conditions: (i) temperature: 20–25 °C, (ii) adsor-bent amount: 5 g L!1, (iii) pH: 6, 7, and 9, (iv) influent phos-phate concentration: "80 mg L!1. A constant pH in the batchreactor was maintained using 5 mM MES, MOPS, and TAPSbuffer (for pH 6, 7, and 9, respectively) and was continuallymixed at a constant speed of 150 rpm (G10 Gyrotary shaker,New Brunswick Scientific Co. Inc., USA). At higher mixingspeed, E33 samples started breaking down. Grab sampleswere collected at various predetermined intervals (e.g. after 5,10, 15, 30, 45 min and 1, 2, 3, 4, 5 h) and filtered using aWhatman 0.45 μm nylon syringe filter. Experiments wereconducted in duplicate while each grab sample was a singlemeasurement.

Rapid small scale column tests were performed in Harvelplastic columns 80 cm high and 1.9 cm in diameter. Onegram of adsorbent media with a fixed height of 1.27 cm, withsand and gravel above and below, was placed in the columnswith a stainless steel sieve at the bottom end of the columnto prevent washout. Using a Thermo Scientific™ FH100MSeries Peristaltic pump, the phosphate solution was passedthrough the column at a rate of 2 mL min!1 at room tempera-ture. Similar to the kinetic study, solution pH was adjustedinitially and buffered to remain constant using 5 mM MES,MOPS, and TAPS. The column effluent samples were col-lected and analyzed for phosphate concentration at varioustime periods and run until complete breakthrough (i.e. satu-ration). Flow through experiments were conducted in dupli-cate while each grab sample was a single measurement.

Variable-dose isotherm adsorption studies were conducted.Different amounts of adsorbent (ranging from 100 to 600 mg)were placed in 250 mL Nalgene polypropylene bottles. Onehundred milliliters of the same 80 mg L!1 phosphate solutionwas then added to the bottles. Buffers (5 mM MES, MOPS,and TAPS) were used to maintain constant pH conditions toevaluate adsorption characteristics at varying pH values. Thesealed bottles were placed on a tabletop rotating shaker andshaken at a constant 150 rpm at 21 ± 1 °C for 2 weeks, whichpreliminary test results revealed was sufficient time for

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equilibrium to be reached. The phosphate concentration inall experiments was analyzed by a colorimetric measurementtechnique in which ammonium molybdate and potassiumantimonyl tartrate react in an acidic solution with orthophos-phate to form phosphomopydbic acid which can be reducedby ascorbic acid to form an intense blue color.54 The absor-bance due to the blue complex was monitored at 880 nmusing UV-Vis spectrophotometry (Hach model number DR2700). Experiments were conducted once and samples weremeasured in triplicate and averaged.

3. Results3.1 Adsorbent characterization

Table 2 summarizes the physical characteristics of E33 andmodified E33 samples. The BET surface area of E33 exceededthat of the modified samples except for E33/AgII, which hada slightly higher BET surface area (142.0 m2 g!1 compared to140.4 m2 g!1). The manganese-coated E33 had a significantlyreduced BET surface area (102.8 m2 g!1) compared to E33.This suggests that at least a portion of the manganese oxidenanoparticles occupied the pore structure of E33. This tablealso summarizes structural characteristics of the adsorbentprior to and after phosphate adsorption. Fig. S1 (a) and (b)‡show N2 adsorption and desorption isotherms for all sam-ples before and after using for phosphate removal. They pres-ent typical Type IV adsorption isotherms, indicating theformation of mesoporous materials.55

3.1.1 SEM characterization. Fig. 1 illustrates SEM imagesof E33 and modified E33. Fig. 1(a) and (b) representunmodified E33 at 12k and 50k magnification, respectively.Based on Fig. 1(b), it could be determined that the particles arespherical and, on average, 75 nm in diameter. Fig. 1(c) and (d)illustrate the surface morphology of E33/Mn at 12k and 50kmagnification, respectively. As can be seen in Fig. 1(d), thespherical particles range in size from 50 to 200 nm indiameter. Also, several particle aggregates 500 nm indiameter were present. Fig. 1(e) and (f) illustrate SEM imagesof E33/AgI. Fig. 1(f) (at 50k magnification) reveals that theaverage particle size was 75 nm which was similar to theunmodified E33. Fig. 1(e) (at 12k magnification) reveals asmoother surface when compared to the other adsorbents.The surface morphology for E33/AgII is illustrated inFig. 1(g) and (h), at 12k and 50k magnification, respectively.The average particle size was determined to be 100 nm, and allparticles were spherical in shape. The images for EDXmapping show adsorbed phosphate (yellow) on the surface of

E33 and E33/AgII. More phosphate was detected on thesurface of E33/AgII compared to that of unmodified E33(please see Fig. S2 in the ESI‡).

3.1.2 XRD analysis. Fig. 2 shows the XRD patterns of allsamples used for phosphate removal. The peaks for allsamples were defined using reference codes in XRD analysissoftware corresponding to different materials such as00-03-0249, 00-001-1053, 00-014-0557, and 00-036-0045 forgoethite, hematite, iron manganese oxide hydroxide, and ironphosphate hydroxide, respectively. The XRD patterns of allsamples are shown in Fig. S3!a)–!d).‡ The well-defined peaksdemonstrated typical XRD patterns of goethite of iron oxidesin E33 and Ag-modified E33 in Fig. 2(a), indicating that Agmodification did not affect the phase transformation of goe-thite. Moreover, the pH of solution for phosphate removaldid not affect the goethite phase of iron oxide (see Fig. S3(a),(c), and (d)‡). However, the peaks, which correspond toAg, were not detected due to low concentration of Ag (seeFig. 2(a)). For Mn-modified E33, the peaks corresponding tohematite of iron oxide were observed, indicating phase trans-formation from goethite to hematite during the synthesisprocess. Well-defined peaks corresponding to iron manga-nese hydroxide were observed, indicating that the surface ofE33 was successfully modified with manganese (see Fig. 2(a)).Fig. 2(b) shows XRD patterns of E33 and surface-modifiedE33 after completion of the phosphate removal studies. Thepeaks corresponding to iron phosphate hydroxide in E33 andall surface-modified E33 were also observed, which indicatesthat the media complexed with the dissolved phosphate in theinfluent water.

3.1.3 TEM and HR-TEM characterization. Fig. 3 representsTEM images at 100k magnification of E33 and the modifiedadsorbents. The silver-modified adsorbents (E33/AgI and E33/AgII), in Fig. 3(c) and (d), had a rod-like shape, similar to theunmodified E33. These particles were determined to be, onaverage, 150 nm long and 20 nm across. It seems that theeffect of silver modification on the morphological structureof E33 is negligible. On the other hand, the manganese-modified E33 had spherically shaped particles 50 nm indiameter on average. In order to investigate the structure ofthe crystal in detail, HR-TEM was employed. Fig. 4 shows theimages of HR-TEM analysis for unmodified and modifiedE33 samples. Fig. 4(a) shows that the measured latticespacing of the sample was 0.42, 0.25, and 0.17 nm, whichcorrespond to the (110), (111), and (221) planes of goethite ofiron oxide, respectively.56 In Fig. 4(b), the lattice fringe with ad-spacing of 0.27 nm for the (222) plane of Mn2O3 was

Table 2 Structural characteristics of E33 and modified E33 before and after PO43! adsorption

Sample

Before adsorption After adsorption

SBET (m2 g!1) Total pore volume (cm3 g!1) Porosity (%) SBET (m2 g!1) Total pore volume (cm3 g!1) Porosity (%)

E33 140.4 0.583 16.2 145.0 0.637 15.0E33/AgI 124.6 0.456 12.6 136.8 0.449 11.1E33/AgII 142.0 0.553 15.4 143.1 0.618 14.7E33/Mn 102.8 0.441 12.3 99.5 0.623 14.8

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measured, which indicated the formation of Mn2O3 duringthe surface modification of E33 with manganese.57–59 Afterthe surface modification of E33 with Ag, the presence of Agwas detected with a lattice spacing of 0.24 nm, correspondingto the (111) plane of Ag. The plane of (200) with a measuredlattice spacing of 0.20 nm was also observed, indicating thatthe surface of E33 was successfully modified with Ag,56,60

although the peaks corresponding to Ag were not detected byXRD analysis due to the low concentration of Ag on the sur-face of E33.

3.2 Adsorption studies

3.2.1 Effect of contact time. Fig. 5 shows the removal ofphosphate as a function of reaction time from 0 up to 220 h.

Fig. 1 SEM images of E33 at 12k (a) and at 50k (b), E33/Mn at 12k (c) and at 50k (d), E33/AgI at 12k (e) and at 50k (f), and E33/AgII at 12k (g) and at50k (h) magnification.

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The effluent presented a decreasing phosphate concentrationwith increasing reaction times. While E33/Mn and E33/AgIhad lower reductions in phosphate concentration, E33/AgIIperformed similarly to the unmodified E33. The phosphateadsorption capacity of the different adsorbents correlates toBET surface area (Table 2). The solution pH in the batchreactor noticeably impacted the efficiency of adsorbents forphosphate removal. For example, E33/AgII at pH 7 removedroughly 65% of phosphate after a 3 h reaction time, yet onlyabout 35% at pH 9. Influent pH also impacted overall adsorp-tion. E33/AgII at pH 7 reduced phosphate concentration(initially at 82.66 mg L!1) to 1.71 mg L!1, while with aninfluent of pH 9, phosphate concentration was reduced to5.71 mg L!1. These graphs also show that E33/AgII removed

50% of the phosphate in the first hour of adsorption. It thentook nearly five more days to reach equilibrium (Fig. 5(b)).

3.2.2 Phosphate adsorption kinetics. The results of kineticexperiments for phosphate adsorption are illustrated inFig. 6. It can be seen that the adsorption rates of E33 andE33/AgII were slightly higher than those of E33/AgI andE33/Mn. Yet all adsorbents portrayed the typical, two-stageadsorption process, with a rapid initial phase and a slowerphase, which likely corresponds to interparticle diffusion. Forexample, after 2 h of adsorption, the phosphate removal forE33, E33/Mn, E33/AgI, and E33/AgII was 66, 54, 38, and 63%,respectively.

The kinetic data for phosphate adsorption seen in Fig. 6were fitted to several kinetic models including the pseudo-

Fig. 2 XRD patterns of E33 and surface-modified E33: (a) before and (b) after using for phosphate removal.

Fig. 3 TEM images at 100k magnification of adsorbent media (a) E33, (b) E33/Mn, (c) E33/AgI, and (d) E33/AgII. (All scale bars: 50 nm.)

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second-order equation and the simple Elovich equation bynon-linear regression using SigmaPlot software. The kineticequations used and their estimated parameters, with thecorrelation coefficient R2, are illustrated in Table 3. Based onthe values of R2, the kinetics of phosphate adsorption on E33and modified E33 can be satisfactorily described by boththe pseudo-second-order equation and the simple Elovichequation; thus, Fig. 6 has both equations plotted also. Theapparent applicability of the Elovich equation, which wasoriginally developed to describe chemisorption of gases onheterogeneous solid surfaces,61 suggests that the sorptionsites involved in phosphate sorption are also of a heteroge-neous nature. Obtained results were generally in agreementwith other results from the literature and can properly explainthe kinetics of phosphate adsorption onto commonly foundsoil minerals such as goethite.62–64

3.2.3 Column study. Column experiments were conductedto quantify the mechanism of phosphate adsorption aswould be seen in an industrial-scale fixed bed adsorber. Thebreakthrough curves were constructed by plotting the ratio ofPO4

3! concentration at time t to the initial influent concen-tration (C/C0) versus time (t). Fig. 7 shows the typical "S"shape of the breakthrough curves indicating the effects ofmass transfer parameters as well as internal resistance within

the column. Phosphate adsorption was initially high, decreas-ing with time until fully saturated. The impact of influent pHon phosphate adsorption was also evident where at higherpH (pH 9), breakthrough time (1.5 hr) was longer, while thetime to reach saturation was shorter (15 hr), as opposed tothe pH 7 where the breakthrough occurred after 1 hr, buttook 18 hr to reach full saturation. The cumulative adsorp-tion capacity of the columns was determined. At pH 6, thecumulative column adsorption capacity for E33, E33/Mn,E33/AgI, and E33/AgII was 31.5, 28.5, 29.9, and 40.1 mg g!1,respectively. These results show that the phosphateadsorption capacity of the adsorbents in columns were muchhigher when compared to batch experiments (e.g. E33/AgIIat pH 6 had a capacity of 17.0 mg g!1 for batch studies and40.1 mg g!1 for column experiments). However, this islikely due to the volume of solution treated. Batch experi-ments were conducted using 0.1 L of phosphate solutionwhile the continuous column experiments passed nearly3 L of phosphate solution through the sorbents. Theseresults indicate that E33/AgII provided a greater capacityfor phosphate adsorption when compared to E33 at high(pH 9) and neutral (pH 6) pH values. The adsorptioncapacities for both batch and column tests are summa-rized in Table S1.‡

Fig. 4 HR-TEM analysis for unmodified and modified E33 used for phosphate removal: (a) E33, (b) E33/Mn, (c) E33/AgI, and (d) E33/AgII. (All scalebars: 2 nm.)

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3.2.4 Adsorption isotherm study. The following equationwas used to determine the amount of phosphate anionsadsorbed per unit weight of adsorbent (qe):

q c c Vme �( *)0

where c0 is the initial concentration (mg L!1), c* is the equi-librium concentration (mg L!1), V is the solution volume (L),and m is the adsorbent mass in grams. Table 1 illustrates thephosphate adsorption capacity of a wide range of adsorbentsincluding metal oxides, red mud, steel slag, and fly ash, as

well as E33 and modified E33. The most efficient adsorbentswere the Nanshi-modified zeolite, the Nanshi-modifiedfly ash, and Fe-loaded biomass char. However, the stocksolution influent pH for these adsorbents was maintainedatypically low (pH 3 and 4) to promote enhanced adsorption.At neutral pH, E33/AgII had a phosphate adsorption capacityof 28.0 mg PO4

3! g!1. This was slightly higher than that ofthe unmodified E33 (26.3 mg PO4

3! g!1), likely due to theincreased BET surface area of the E33/AgII sample. Thecapacities from this study compare favorably with the capaci-ties of other adsorbents as reported in literature.

Adsorption isotherms are used to relate the bulk adsorbateconcentration and the amount adsorbed at the surface inter-face. The isotherm results were analyzed using the Langmuirand Freundlich isotherms. The Langmuir isotherm model,which assumes monolayer saturation with no lateral inter-action between adsorbed particles, has previously been

Fig. 5 Rate of E33 and modified E33 adsorption of phosphate atvarious pH values: (a) pH 6, (b) pH 7, and (c) pH 9.

Fig. 6 Phosphate adsorption kinetic data for E33 and modified E33 at20 °C, pH 7, and an influent concentration of 83 mg P L!1.

Table 3 Estimated parameters for the kinetic model of phosphateadsorption on E33 and modified E33

Power function equation (q = atb)a b R2

E33 8.259 0.1715 0.8718E33/Mn 6.396 0.1793 0.9464E33/AgI 5.728 0.2288 0.9343E33/AgII 8.139 0.1742 0.8817

Simple Elovich equation (q = a + b ln t)a b R2

E33 8.228 2.071 0.9600E33/Mn 6.513 1.619 0.9894E33/AgI 5.895 2.045 0.9645E33/AgII 8.130 2.070 0.9645

Pseudo second-order equation (qt = k2qe2t/1 + k2qet)

qe,cal k2 (g mg!1 min!1) R2

E33 16.34 0.0641 0.9725E33/Mn 13.74 0.0645 0.9021E33/AgI 15.46 0.0272 0.9547E33/AgII 16.31 0.0616 0.9639

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employed in similar adsorption experiments.62–64 The Langmuiradsorption model is mathematically expressed by:

q Q K CK Ce

L e

L e

�max

1

where qe (mg g!1) is the amount of adsorbed phosphate perunit mass of adsorbent, Ce (mg L!1) is the unadsorbed phos-phate concentration in solution at equilibrium, Qmax is themaximum amount of phosphate per unit mass of sorbent to

form a complete monolayer on the surface, and KL (L mg!1)is a constant related to the affinity of the binding sites.The Langmuir adsorption equation can be described by itslinearized form:

1 1 1 1q K Q C Qe L e

§

©¨

·

¹¸ �

max max

The linear plot of the above equation (Fig. 8) yields essen-tial specifications of the system. The slope and intercept ofthe plot were used to determine the Langmuir constants Qmax

and KL, which are presented in Table 4.

Fig. 7 Column experiments for E33 and modified E33 phosphateadsorption at (a) pH 6, (b) pH 7, and (c) pH 9.

Fig. 8 Langmuir isotherm graphs for (a) pH 7 and (b) pH 9.

Table 4 Isotherm parameters at pH 7

Sample

Langmuir isotherm constantsFreundlichisotherm constants

Qmax (mg g!1) KL (dm3 mg!1) R2 KF 1/n R2

E33 25.71 0.82 0.9999 14.72 0.15 0.7389E33/Mn 18.73 0.24 0.9836 9.59 0.15 0.8133E33/AgI 20.58 0.36 0.9642 12.87 0.10 0.6718E33/AgII 28.01 0.69 0.9917 15.75 0.15 0.7893

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Essential characteristics of the Langmuir isotherm modelcan be expressed in terms of a constant dimensionless sepa-ration factor, RL, given by the equation:

RK CLL

�� �1

1 0

where C0 (mg L!1) is the initial adsorbate concentration insolution and KL (L mg!1) is the Langmuir constant. The valueof RL can indicate if adsorption is unfavorable (RL > 1), linear(RL = 1), irreversible (RL = 0), or favorable (0 < RL < 1). RL

values for E33, E33/Mn, E33/AgI, and E33/AgII were 0.015,0.050, 0.034, and 0.018, respectively. As all values computedfor RL were less than 1 and greater than 0, adsorption wasdetermined to be favorable.

The Freundlich isotherm is an empirical equation used toexplain heterogeneous systems. The Freundlich equation isgiven by:

qe = KF(Ce)1/n

where KF ((mg g!1) (L mg!1)1/n) is the adsorption capacity ofthe adsorbent and n gives an indication of adsorption favor-ability, for example, values of n > 1 indicate favorable adsorp-tion. Freundlich isotherm constants KF and n can be deter-mined by the linearization of the Freundlich equation andgraphing ln!qe) vs. ln!Ce):

ln lnq Kn

Ce� � � � � � �� �F e1

Values of KF and n can be determined from the interceptand slope of the plot for the linearized Freundlich equation.The linearized Freundlich isotherm plots for phosphateadsorption onto E33 and modified E33 are presented inFig. 9. The Freundlich isotherm constants, KF and 1/n, weredetermined by linear regression (Table 4).

Fig. 10 shows phosphate adsorption isotherms for E33and modified E33 under identical conditions, i.e. DI water atpH 7. The results show that E33/AgII had the highest capacityof phosphate adsorption for the modified adsorbents as wellas unmodified E33. This is in accordance with BET surfacearea analysis, which revealed that E33/AgII had the highestspecific surface area. Likewise, E33/Mn, which had the lowestphosphate adsorption capacity, also had the lowest BET sur-face area. Isotherm results were evaluated using both

Fig. 9 Freundlich isotherm graphs for influent pH values of (a) 7 and(b) 9.

Fig. 10 Isotherm graph for phosphate adsorption at pH 7.

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Langmuir and Freundlich isotherm equations. The deter-mined isotherm constants for both models are summarizedin Table 4. As can be seen from the coefficients of determina-tion, the Langmuir isotherm appears better suited to explainthe adsorption equilibrium data.

4. ConclusionThree modified adsorbents, based on the commercially avail-able E33 adsorptive medium, were successfully synthesizedand characterized using SEM, TEM, XRD, porosimetry analy-sis and EDX. The adsorption capacity of the synthesized sam-ples was investigated for the removal of phosphate. Theresults of the characterization of synthesized samples indi-cated that the surface of E33 was successfully modified withnanostructured Ag and manganese. In particular, for Ag-modified E33, the phase of E33 was not affected by Ag modi-fication while the transformation of goethite to hematite wasobserved in Mn-modified E33.

Batch and column experiments were conducted to deter-mine the adsorbents' capacity for phosphate adsorption. Thephosphate removal was influenced by the pH of solutions. Asthe solution pH decreased, the amount of phosphateadsorbed increased, presumably because the surface chargeof E33 is pH dependent, thus lower pH values correlate to amore positive surface charge on the E33, thus promotingincreased PO4

3! adsorption. Of the modified adsorbents syn-thesized, Ag-modified E33 using tea extract (E33/AgII) had aslightly higher phosphate adsorption capacity compared tounmodified E33. This increase in adsorptive capacity is likelydue to an increased BET surface area for E33/AgII. AlthoughMn-modified E33 (E33/Mn) was believed to promoteincreased adsorption, a reduction in BET surface area (likelycaused by the aggregation of particles resulting in a sphericalshape for the E33/Mn particles as revealed by TEM) resultedin a pore blockage and lower phosphate adsorption capacitywhen compared to unmodified E33. The increased adsorp-tion for E33/AgII suggests the potential for improving the per-formance of commercially available adsorbents for phosphateremoval.

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