ar

Ma

a Laboratoire de Chimie Moleculaire et Environnement, Universite de Savoie, Le Bourget du Lac 73376, Franceb Laboratoire de Chimie Analytique et Electroc National Institute of Advanced Industrial Sc

Tsukuba, Ibaraki 305-8569, Japan

A R T I C L E I N F O

Article history:

Received 3 August 2012

Accepted 28 November 2012

Available online 7 December 2012

e adsorption uptake while decreasing adsorption Gibbs energy (nota-

Pharmaceutically active compounds (PhACs) are frequently

Many of these compounds are also suspected to have adverse

impact on humans and wild life [4].

According to literature, ibuprofen (IBP), a non steroidal

anti-inflammatory, is one of the most common drugs found

in water [5]. Indeed, the latter has been detected in the

hazard for human health [7].

electrochemical degradation [10], recognizing extraction

using molecularly imprinted polymers (MIPs) [11,12], adsorp-

tion on mesoporous silica [13] or activated carbons [14,15].

Activated carbon has several applications depending on

its porous structure and adsorption capacity [16]. In the

C A R B O N 5 4 ( 2 0 1 3 ) 4 3 2 4 4 3

Avai lab le at www.sc iencedi rect .com

els* Corresponding author: Fax: +33 479758805.detected in treated wastewater arising from sewage treat-

ment plants, from which they are not completely removed

[1,2]. Moreover, recent studies have reported the presence of

traces or ultra traces (ng L1) of PhACs in drinking water [3].

The removal of IBP was already reported using different

techniques such as ultrasonic irradiation [8], combined ad-

vanced oxidation processes (AOPs) such as sonophotocatalyt-

ic degradation in the presence of Fe3+ and TiO2 [9],bly 7.3 kJ mol1 after sonication in H2O2). The decrease of the adsorption capacity afterbleaching was attributed to the presence of phenolic groups.

2012 Elsevier Ltd. All rights reserved.

1. Introduction effluents from several sewage-treatment plants at concentra-tions up to 24.6 lg L1 [6], possibly presenting a potentialmation of carbonyl gr

helped to enhance th0008-6223/$ - see front matter 2012 Elsevihttp://dx.doi.org/10.1016/j.carbon.2012.11.059

E-mail address: Laurent.duclaux@univ-sachimie, Universite de Tunis, 2092 Manar II, Tunisia

ience and Technology, Energy Technology Research Institute, 16-1 Onogawa,

A B S T R A C T

A commercial microporousmesoporous granular activated carbon was modified by oxida-

tion with either H2O2 in the presence or absence of ultrasonic irradiation, or NaOCl or by a

thermal treatment under nitrogen flow. Raw and modified materials were characterized by

N2 adsorptiondesorption measurements at 77 K, Boehm titrations, pH measurements and

X-ray photoelectron spectroscopy. Ibuprofen adsorption kinetic and isotherm studies were

carried out at pH 3 and 7 on raw and modified materials. The thermodynamic parameters

of adsorption were calculated from the isotherms obtained at 298, 313 and 328 K. The pore

size distribution of carbon loaded with ibuprofen brought out that adsorption occurred

preferentially into the ultramicropores. The adsorption of ibuprofen on pristine activated

carbon was found endothermic, spontaneous (DG = 1.1 kJ mol1), and promoted at acidicpH through dispersive interactions. All explored oxidative treatments led mainly to the for-

oups and in a less extent to lactonic and carboxylic groups. This thenNizar Bellakhal b, Laurent Duclaux a,*The effects of the surface oxidsolution pH and the temperatu

Hanen Guedidi a,b, Laurence Reinert a, Jean-

journal homepage: www.er Ltd. All rights reservedvoie.fr (L. Duclaux).tion of activated carbon, thee on adsorption of ibuprofen

rc Leveque a, Yasushi Soneda c,

evier .com/ locate /carbon.

literature, numerous studies have shown that adsorption by

activated carbon is effective for the removal of various pollu-

tants in water and waste water such as heavy metals [1719],

(2)

C A R B O N 5 4 ( 2 0 1dyes [2025], pesticide [26,27], fluoride [28] and PhACs [29].

Mestre et al. [15] previously reported that IBP adsorption

on activated carbons was enhanced by the presence of supe-

rmicropores and acidic pH values but was not by the variation

of temperature in the range of 298313 K. In addition, they

emphasized the favorable effect of basic carbon surface [14]

on the IBP adsorption process and hypothesized that the pres-

ence of mesopores was crucial to ensure the accessibility to

the inner microporosity which had to be large enough

(0.7 nm) to accommodate IBP moieties. More recently, the

adsorption of IBP on a mesoporous activated carbon prepared

from Artemisia Vulgaris leaves [30] was reported to be endo-

thermic and the maximum uptake was observed at pH 2 sug-

gesting a physisorption mechanism involving the carboxylic

function of the adsorbate molecule.

The modification of the surface chemistry of activated car-

bons by treatments in liquid media with oxidative agents

such as HNO3, (NH4)2S2O8, and H2O2 is known to generate

an increase in the number of acidic surface sites. However,

this also leads to a drastic decrease in the total specific sur-

face area and micropore volume, explained by the presence

of oxygen groups at the entrance and the walls of the pores

[31]. For all cited treatments, except the one with H2O2, a

destruction of the basic sites was observed [31].

Functional oxygen groups play an important role in the

adsorption of aromatic molecules. Haydar et al. [32] reported

that oxidation by HNO3 greatly increased the number of car-

boxylic groups, withdrawing electrons from the graphene lay-

ers and thus reducing the p-nitrophenol uptake. Mattson

et al. [33] suggested that an interaction between carbonyl sur-

face groups (as electron donors) and aromatic rings of phenol

(acceptors) lead to the formation of donoracceptor com-

plexes. Other authors explained the adsorption of phenol

through the formation of hydrogen bonds (of the OH function

of phenol) with carboxylic groups [34].

This present work aims at investigating the IBP adsorption

profile on chemically modified surfaces-bearing activated car-

bons. Thus, a commercial microporousmesoporous acti-

vated carbon was treated in several different ways in order

to obtain adsorbents with different surface properties. The

adsorption properties of the modified and pristine activated

carbons were studied as a function of concentration, pH,

and temperature to investigate the impact of the surface

chemistry on the IBP adsorption.

2. Experimental

2.1. Ibuprofen

2-[4-(2-Methylpropyl) phenyl] propanoic acid (Fig. 1), also

named ibuprofen, was purchased from the nationalFig. 1 Molecular structure of Ibuprofen.an ultrasonic probe (20 kHz, 43 lm amplitude, 19 mm

diameter, 350 W electrical power, Sonics and Materials

Ultrasonic Processor - VC505) dipped into the carbon

suspension, maintained at 298 K in a cylindrical dou-

bled-jacket cooled reactor.

All the oxidized samples were then recovered by

filtration at 0.45 lm on membrane filters (Durapore-

Millipore), rinsed with distilled water until reaching a

neutral pH (1 L per g) and further dried at 353 K for 24 h.

The bleached AC was named AC[NaOCl]. The samples

oxidized with H2O2, and sonicated in H2O2 were

referred to AC[H2O2] and AC[H2O2]US, respectively.

(3) Sample named AC700N2 was prepared by heating at

973 K (10 K min1, 1 h at 973 K), under a nitrogen flowand kept under this inert atmosphere until cooling to

room temperature.

2.3. N2 adsorptiondesorption at 77 K and CO2 adsorptionat 273 K

N2 adsorptiondesorption isotherms of the activated carbons

were measured using an automatic sorptometer (ASAP 2020,

Micromeritics) at 77 K. The CO2 adsorption isotherms were

measured at 273 K on two samples (AC and AC loaded with

56.7 mg g1 IBP at pH 7). Prior to measurements, samples (ex-cept IBP loaded AC) were degassed for 12 h at 523 K under vac-

uum. IBP loaded AC was outgassed for 4 days at 323 K, below

ibuprofen melting point (348350 K) [39].

The specific surface areas of the activated carbons were

calculated using the BrunauerEmmettTeller (BET) equation,

assuming the area of the nitrogen molecule to be 0.162 nm2.

As negative unrealistic C factors were obtained by applyingthe B

0.3, thFor the bleaching treatment, AC (1 g) was dispersed in

200 mL of a 0.13 mol L1 NaOCl solution and stirred atroom temperature during 24 h.

AC samples (2.5 g) were also oxidized by impregnation

at room temperature for 5 h with 50 mL of 35% H2O2,

or (3) by ultrasound irradiation for 5 h in H2O2 (35%,

50 mL). The sonication treatment was performed usingthe filtrate [38]. After drying at 383 K for 24 h, this sample, re-

ferred to AC, was further chemically modified by four differ-

ent treatments as described here below:

(1)laboratory of the control of drugs (LNCM, Tunisia) and used

as received (99.1% purity). This molecule is very slightly

soluble in water (21 mg L1) [35] but soluble in a majority oforganic solvents [36,37]. The dimensions were estimated

using Chemsketch 3D Viewer assuming a parallelepiped

shape: 1.13 0.50 0.31 nm3.

2.2. Raw and modified activated carbons

A commercial granulated activated carbon (SigmaAldrich,

Ref. 292591, 414 mesh) was milled and washed with a hydro-

chloric acid solution (0.1 mol L1) for 24 h under stirring to re-move metal salt impurities. The resulting suspension was

then filtered off and the recovered carbon was rinsed with

distilled water (200 mL/g) until reaching a constant pH of

3 ) 4 3 2 4 4 3 433ET model in the relative pressure range from 0.05 to

e BET specific surface areas were then preferentially

0 1computed in the relative pressure range from 0.01 to 0.05, as

for microporous materials.

The total pore volume was estimated as the liquid vol-

ume of N2 adsorbed at a relative pressure of 0.995. In addi-

tion, the micropore volume was determined by using the

Dubinin-Radushkevich equation for N2 adsorption (P/P0from 0.01 to 0.05) and for CO2 adsorption [40]. The meso-

pore volume was calculated by difference of the total pore

volume and the micropore volume. The pore size distribu-

tions (PSD) of the pristine AC and IBP loaded AC (pH = 7

and 56.7 mg g1 of IBP) were determined by using NLDFT(non local density functional theory) models applied on

both the adsorption isotherms of N2 at 77 K and CO2 at

273 K. For that, infinite slit pores were assumed for CO2adsorption (pores diameter from 0.35 to 1.1 nm), while finite

slit pores having a diameter-to-width aspect ratio of 12

were used for N2 adsorption simulations (pores diameter

from 3.5 to 250 A) [41]. N2 adsorption data at P/P0 < 0.01

were obtained using incremental fixed doses of 10 cm3 g1(STP), setting the equilibration interval at 300 s [42]. CO2adsorption data were obtained at P/P0 ranging from

4 104 and 3.5 102, using 45 s equilibration interval.

2.4. pHs of carbons suspended in water

The pH of each activated carbon sample (0.5 g) was measured

in a distilled water suspension (12.5 mL) after heating at 90 Cand then cooling to room temperature [43].

The pHPZC (pH value at which the surface has a zero net

charge) was determined by the so-called pH drift method

[34]. The pH of deoxygenized NaCl aqueous solutions (50 mL

at 0.01 mol L1) were adjusted to successive initial values be-tween 2 and 12 by adding either 0.1 mol L1 HCl or NaOH. Sus-pensions were prepared by adding activated carbon (0.15 g) to

each solution. After stirring for 48 h under N2, the final pH

was measured and plotted versus the initial pH. The pHPZCwas determined at the value for which pHfinal = pHinitial.

2.5. Boehm titrations

Boehm titrations quantify the basic and oxygenated acid sur-

face groups on activated carbons [44]. Surface functional

groups such as carboxyl (RCOOH), lactone (ROCO), phenol

(ArOH), carbonyl or quinone (RR 0C@O) and basic groups weredetermined using different reactants, assuming that: NaOC2-H5 reacted with all groups; NaOH did not react with RR 0C@Ogroups; Na2CO3 did not react with RR 0C@O nor ROH groupsand that NaHCO3 only reacted with RCOOH groups. The pro-

cedure was as follows: about 0.15 g of each carbon sample

was mixed in a closed erlenmeyer with 50 mL of a 0.1 mol L1

aqueous reactant solution (NaOH, Na2CO3 or NaHCO3). In the

case of NaOC2H5, only 0.1 g of carbon sample was added to

50 mL of 0.01 mol L1 solutions. The mixtures were stirredfor 24 h at constant speed (650 rpm) and room temperature

and then filtered off. Back-titrations of the filtrate (10 mL)

were then achieved with standard HCl (0.01 mol L1) to deter-mine the oxygenated groups contents. Basic groups contents

were also determined by back titration of the filtrate with

434 C A R B O N 5 4 ( 2NaOH (0.01 mol L1) after stirring of the activated carbon(0.15 g) in HCl (50 mL, 0.01 mol L1) for 24 h.2.6. X-ray photoelectron spectroscopy

X-ray photoelectron spectroscopy (XPS) measurements were

performed using an ESCALAB 250 spectrometer (Thermo Fish-

er Scientific) at monochromated Al-Ka anode X-ray radiation,

on a 150 800 lm2 analysis region, under 2 109 mbar vac-uum. The high-resolution scans (0.1 eV) were obtained over

the 280.1299.9 eV (C1s) and 523.1539.9 eV (O1s) energy

ranges with a pass energy of 20 eV. After baseline subtraction,

the curve fitting was performed assuming a mixed Gaussian

Lorentzian peak shape (the ratio of Gaussian to Lorentzian

form was equal to 0.3). The carbon 1s electron binding energy

corresponding to graphitic carbon was referenced at 284.6 eV

for the calibration [45].

2.7. Adsorption experiments: kinetics and isotherms

The IBP solutions were prepared from UHQ water (Ultra High

Quality, 18.2 MX purity) containing 10 vol.% of methanol

(99.9%, SigmaAldrich) in order to increase the solubility.

The kinetics were studied at room temperature at pH 3 and

7. Suspensions of 10 mg of the differently treated ACs in

30 mL of 100 ppm IBP solution were put to stirring at

650 rpm and then filtered at different intervals between

15 min and 7 days. The IBP concentrations were subsequently

determined by UVvisible spectrometry (Varian, Cary50). The

maximum absorbance was observed at 220 nm at pH = 3, and

222 nm at pH = 7.

The isotherms of adsorption of IBP on the various acti-

vated carbons were studied at three constant temperatures:

298, 313 and 328 K, at pH 3 and 7. Activated carbons (10 mg)

were introduced in IBP solutions (30 mL) of varying concentra-

tions (5100 mg L1) and stirred for 5 days.The equilibrium IBP uptake Qads (mg g

1) was calculatedfrom equation:

Qads Ce Ci V

m

where V is the solution volume (mL), Ci is the initial IBP con-

centration (mg L1), Ce is the equilibrium IBP concentration(mg L1) and m is the weight of the dry activated carbon (mg).

3. Results and discussion

3.1. Characterization of activated carbons

3.1.1. Surface chemistryThe Boehm titrations of AC revealed the overall presence of

more basic than acidic functional groups (Table 1), confirmed

by the pH value (9.05). When AC was oxidized respectively by

NaOCl or by H2O2 with and without ultrasonic irradiation, the

pH values decreased due to the increase of the oxygenated

surface groups concentrations from 0.37 to 2.1, 3.9 and3.5 meq g1. The value of pHPZC slightly decreased for thebleached activated carbon (from 7.76 to 7.60) but increased

for both AC[H2O2] and AC[H2O2]US (Table 1). This slight in-

crease in pHPZC could be attributed to the higher amount of

carbonyl groups (2.10 and 2.25 meq g1 for AC[H2O2] and

3 ) 4 3 2 4 4 3AC[H2O2]US, respectively) and lactonic groups (0.77 and

1.32 meq g1 for AC[H2O2] and AC[H2O2]US, respectively)

deconvolution and brought out the highest content in ad-

sorbed O detected by XPS (Table 2).

Whatever the oxidized activated carbons, their contents in

graphitic carbon (at 284.6 eV) were lower than in pristine AC,

and the percentages of oxygen bonded to carbon increased

ti

AC

Fig. 2 O1s XPS spectra of raw and modified activated

carbons.

C A R B O N 5 4 ( 2 0 1 3 ) 4 3 2 4 4 3 435which possess high pKa values (pKa-carbonyl 1620, pKa-lactonic 8.2).

The oxidation by ultrasonic irradiation in H2O2 (AC[H2O2]-

US sample) compared to the silent analog (AC[H2O2] sample)

yielded to a slight decrease of the carbonyl groups. By con-

trast, additional phenolic groups (0.27 meq g1) were foundunder ultrasound irradiation, whose origin can be attributed

to the reaction between atoms of carbon and free OH radicalsgenerated by the sonolytic degradation of H2O2 [46]. Ultra-

sonic and silent treatments with hydrogen peroxide gener-

ated the highest contents in carbonyl and lactonic groups

(Table 1) while the bleaching treatment yielded to the highest

amount of phenolic groups (0.5 meq g1). The carbonyl andlactonic contents were lower in bleached samples

(1 meq g1) than in samples treated with H2O2 (>2 meq g1).The ultrasonic treatment also induced a significant particle

size reduction due to the cavitation phenomenon [47]. Laser

granulometry characterization (not shown) confirmed that

AC[H2O2]US was mainly formed of submicrometric particles

(median size d50 = 1 lm) while oxidation in H2O2 without

ultrasound preserved the micrometric particles

(d50 = 100 lm), such as for the raw AC (d50 = 700 lm).

Boehm titrations (Table 1) revealed quite different surface

groups contents for a fresh sample characterized just after

annealing under N2 atmosphere, in comparison to a one-year

air atmosphere aged sample. The fresh AC700N2 sample

was found to be the most basic sample, which is in agreement

Table 1 Surface groups (in meq g1) obtained from Boehmcarbons.

Sample AC AC[NaOCl]

Carboxylic groups 0 0.15Lactonic groups 0.03 0.35Phenolic groups 0.10 0.50Carbonyl groups 0.24 1.08Total oxygenated groups 0.37 2.08Total basic groups 0.99 0.99pH 9.05 7.18pHPZC 7.76 7.60a Fresh AC700N2.b Aged AC700N2.with its measured pH and pHPZC values. Although the anneal-

ing under N2 atmosphere removed many oxygen surface

groups, the sample underwent again oxidation with time

when exposed to air atmosphere (Table 1), leading to a similar

carbonyl content (1 meq g1) than the bleached sample. Infact, a N2 thermal treatment is known to be only temporarily

efficient for removing the oxygen containing surface groups,

as it also generates very reactive surface sites that further re-

act with oxygen from air, enabling the formation of new oxy-

genated surface groups [48].

Surface chemical groups of the different activated carbons

were also characterized by XPS. Fig. 2 presents the O1s spec-

trum of the raw and treated activated carbons. Deconvolution

of the C1s and O1s spectra gave respectively five and four

peaks (Table 2).

XPS analysis of the aged AC700N2 sample confirmed a

slight oxidation compared to raw AC, through C1s spectrumtrations, pH and pHPZC values of raw and modified activated

[H2O2] AC[H2O2]US AC700N2a AC700N2

b

0.32 0.37 0 01.32 0.77 0 0.230 0.27 0 0.032.25 2.10 0.10 1.023.89 3.51 0.10 1.280.65 0.62 1.14 0.778.38 8.03 10.14 8.747.80 8.03 9.08 8.44after oxidation (Table 2), in agreement with the Boehm

titrations.

The C1s spectra (Table 2) displayed higher carbonyl (or qui-

none) and carboxylic (or ester or lactone) groups contents in

AC[H2O2] (8.05% of C@O group; 5.82% of OC@O groups) andAC[H2O2]US (9.82% of C@O group; 7.69% of OC@O groups)compared to the bleached AC[NaOCl] sample (4.45% of O

C@O, and 7.02% of C@O), for which the levels were quite sim-ilar to the AC sample. By contrast, Cagnon et al. [49] reported

an increase in carboxylic and carbonyl groups XPS contents

after bleaching treatment of an activated carbon but without

mentioning the used concentration of NaOCl. Perrad et al. [50]

also showed by XPS an increase in the formation of carboxylic

groups after oxidation of an activated carbon cloth by NaOCl

(0.53 M) for 90 min.

XPS O1s (Fig. 2) analyses confirmed that oxidation was

more efficient under ultrasonic irradiation in hydrogen

This reduction after oxidation treatment was the highest for

the AC[NaOCl] sample. A blockage of the micropores due to

the formation of surface functional groups might be the cause

of this slight reduction in the surface area as already stated

before [53]. By contrast, the very small increase in the BET

specific area (SBET = 809 m2 g1) and in the micropore volume

(0.32 cm3 g1) for AC700N2 compared to raw AC, might beattributed to the removal of the functional groups from the

inner pores through the thermal treatment (Table 1).

The ultramicropore volumes obtained from the DFT PSD

s)

42042

5977

8

P/P0

Fig. 3 N2 adsorptiondesorption isotherms at 77 K of the

pristine AC (h: adsorption, j: desorption) and AC[NaOCl] (4:adsorption, m: desorption).

0 1peroxide (14.37 at.0% in sample AC[H2O2]US) compared to the

silent analog reaction (11.97 at.0% in sample AC[H2O2]).

XPS analyses (Table 1), indicated higher carbonyl and lac-

tonic groups amount in AC[H2O2]US compared to AC[H2O2],

disagreeing with the contents determined by Boehm titra-

tions that are known to be only semi-quantitative.

After NaOCl treatment, the increase of the number of C-O

bonds from phenols and lactones (16.77% of CO for AC[NaO-

Cl] compared to 10.22% for AC) was consistent with the evolu-

tion observed by Boehm titrations.

Whatever the activated carbon, the O1s analyses (Table 2)

showed OC at.% lower than the CO content obtained from

the C1s spectra, which could be the signature of the presence

of ether bridges in all of the activated carbons.

The O1s analyses after deconvolution (Table 2) indicated

also O@C at.% lower than the C@O content obtained fromthe C1s spectra, particularly for raw AC and samples treated

with H2O2 as already reported for other oxidized activated car-

bons [51,52]. This revealed the difficulties for XPS analysis to

take into account all the surface oxygen groups species in

particular ethers and anhydrides, etc.

The peak at the lowest energies (530.42531.01 eV) related

to the presence of metal oxides, increased after oxidation by

H2O2. Indeed, the X-ray micro-analysis showed the presence

of metals, such as Al and Fe on the surface of AC which might

have formed metal oxides after their oxidation by hydrogen

Table 2 Atomic percentages of carbon (C1s) and oxygen (O1

Signal attribution E(eV) AC

C1s CC (284.6) 62.0CO (285.13285.55) 10.2C@O (285.99286.75) 6.5OC@O (286.83289.14) 4.4Carbonate, pp transition (288.57291.48) 9.9

O1s Metal oxide (530.42531.01) 0.8O@C (531.8532.29) 1.8OC (532.96533.77) 2.8Chemisorbed O (534.7535.22) 1.2

Total atomic O% 6.8

436 C A R B O N 5 4 ( 2peroxide.

3.1.2. Porosity characterizationThe nitrogen adsorptiondesorption isotherms for the raw

(Fig. 3) and modified carbons are quite similar (not shown),

except for AC[NaOCl] sample (Fig. 3). They were typical of

microporousmesoporous adsorbents (type I and IV in the IU-

PAC classification) with a hysteresis loop (H4 types) in the

desorption branch at relative pressures above 0.5. The adsorp-

tiondesorption hysteresis on AC isotherm showed clearly

that liquid nitrogen was condensed in slit-shaped mesopores.

The N2 isotherms at 77 K showed a slight decrease of the

BET surface area (Table 3, SBET = 709 m2 g1 for AC[NaOCl],

781 m2 g1 for AC[H2O2] and 762 m2 g1 for AC[H2O2]US) of

the oxidized carbons compared to the pristine AC (SBET = 800 -

m2 g1). The decrease in the specific surface areas after oxida-tion was related only to the decrease in the micropore

volumes observed either by N2 or CO2 adsorption (Table 3).obtained by XPS analysis of raw and modified AC.

AC[NaOCl] AC[H2O2] AC[H2O2]US AC700N2

53.02 51.77 46.74 54.7416.77 12.07 18.84 11.577.02 8.05 9.82 7.214.45 5.82 7.69 4.986.49 10.38 2.53 10.51

0.83 2.31 2.48 1.026.28 3.51 5.40 4.914.63 3.93 5.71 2.400.51 2.16 0.78 2.67

12.25 11.91 14.37 11

0

50

100

150

200

250

300

350

0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1

Vads

(cm

3 g-1

STP

)

3 ) 4 3 2 4 4 3analysis of CO2 isotherms (Table 3) were found quite similar

(about 0.14 cm3 g1) for all the modified activated carbonsand slightly lower than to the one of AC (0.18 cm3 g1).

The N2 DFT pore size distribution of the pristine AC (not

shown) confirmed that the AC material was both of micropo-

rous and mesoporous nature. The micropore and mesopore

volumes computed from the N2 DFT model were 0.31 and

0.21 cm3 g1, respectively.The N2 adsorptiondesorption isotherm of raw AC and IBP-

loaded AC (Fig. 4 left) had a very similar profile but the iso-

therm of the loaded sample was shifted towards lower

adsorption volumes. This difference is due to a decrease of

the microporous volume concomitant to IBP adsorption. The

volume occupied by IBP calculated from loading content:

0.05 cm3 g1 was similar than to the volume calculated fromthe difference of microporous volumes before after loading

(0.05 cm3 g1 from N2 adsorption) (Table 3). An activated car-

bon (AC or treated AC) saturated by IBP at pH 3, where the up-

rption and CO2 adsorption studies.

l] AC[H2O2] AC[H2O2]US AC700N2 IBP loaded AC

781 762 809 6550.31 0.30 0.32 0.260.25 0.25 0.23 0.220.56 0.55 0.55 0.480.26 0.25 0.27 0.180.14 0.14 0.15 0.12

0 1 3 ) 4 3 2 4 4 3 437Table 3 Textural properties obtained by N2 adsorption/deso

Sample AC AC[NaOC

BET surface area SBET (m2 g1) 800 709

Micropore volumea (N2) (cm3 g1) 0.31 0.28

Mesoporous volume (cm3 g1) 0.21 0.21Total pore volumeb (cm3 g1) 0.52 0.49Micropore volumea (CO2) (cm

3 g1) 0.27 0.23Ultramicropore volumec (CO2) (cm

3 g1) 0.18 0.14a From Dubinin-Radushkevich.b Single point adsorption at P/P0 = 0.995.c From NLDFT slit pore model.

Isotherms

0

50

100

150

200

250

300

350

400

0.000001 0.00001 0.0001 0.001 0.01 0.1P/P0

Vads

(cm

3 g-1

STP

)

C A R B O N 5 4 ( 2take is maximum, contains 130165 mg g1 of adsorbate,which means a volume of 0.11 cm3 g1 for the uptake volumeof IBP (taking into account the volume of IBP : 0.23 nm3 men-

tioned in Section 2.1). As the IBP uptake volume is lower than

the microporous volumes of the studied activated carbons,

this molecule can be accommodated in the micropores (Vol-

ume of micropores 0.31 cm3 g1 for AC). The ultramicroporevolume values, investigated by CO2 adsorption before and

after loading (Table 3), have confirmed that these smaller

micropores were occupied by IBP (IBP uptake volume deduced

from ultramicropore volumes: 0.06 cm3 g1). The PSD plots of

pristine AC and IBP-loaded AC (Fig. 4 right) clearly displayed

that the adsorption occurred in the ultramicropores of diam-

eter lower than 0.6 nm. Similarly we have studied the micro-porosity of a IBP-loaded AC700N2 by CO2 adsorption (Volume

of ultramicropores 0.09 cm3 g1 for IBP-loaded AC700N2),and we have confirmed that IBP was also adsorbed into ultra-

micropores. Mestre et al. [14] studied a series of activated car-

bons with various surface chemistry and showed that a large

fraction of ultramicropores were not appropriate to effectively

remove ibuprofen. In our activated carbon, IBP is preferen-

tially adsorbed in the slit shape ultramicropores which are

accessible to the molecule in terms of dimensions.

3.2. Adsorption kinetics of raw activated carbon

Fig. 5 shows the effect of contact time on the adsorption of IBP

onto AC. The equilibrium adsorption time was found at about

4000 min. The kinetics of adsorption of IBP by AC was

Fig. 4 Adsorption isotherms of N2 at 77 K (s: AC, h: 56.7 mg g

loaded AC), and NLDFT pore size distributions of AC (4) and IBP1

PSD

0

0.01

0.02

0.03

0.04

0.05

0.06

3 5 7 9Pore width ()

Incr

emen

tal v

olum

e (c

m3 g

-1)investigated by four models, namely, the Lagergren pseudo-

first-order model (1) [54], the pseudo-second-order model (2)

[55], the Elovich model (3) [56] and the intraparticle diffusion

model (4) [57].

The first-order equation is:

dqtdt

k1qe qt 1

where k1 is the rate constant of pseudo-first-order adsorption

(min1), qe and qt (mg g1) are the amounts of adsorbed IBP at

equilibrium and at time t (min), respectively.

The second-order equation is:

dqtdt

k2qe qt2 2

where k2 is the rate constant of pseudo-second-order adsorp-

tion (g mg1 min1), qe and qt (mg g1) are the amounts of ad-

sorbed IBP at equilibrium and at time t (min), respectively.

1 IBP loaded AC at pH 7), and of CO2 at 273 K (4: AC, : IBPloaded AC (), from CO2 isotherms.

00.10.20.30.40.50.60.70.80.9

1

0 1000 2000 3000 4000 5000 6000Time (min)

C/C0

Fig. 5 Kinetic result of Ibuprofen adsorption onto AC at

25 C and pH 3 ([IBP] = 100 mg L1, madsorbent = 10 mg,VIBP = 30 mL, stirring speed = 650 rpm).

The Elovich equation is:

qt 1b

lna b 1b

lnt t0 3

where a is the adsorption initial rate (mg g1 min1, t0 = 1/(a b)) and b is a constant (g mg1) related to the external sur-face area and activation energy of adsorption

(chemisorption).

The intraparticle diffusion model is formulated by:

qt kpt0:5 4

where kp (mg g1 min1/2) is the intraparticle diffusion rate

constant.

adsorption uptake at pH = 7, the latter increased with the oxi-

dative treatment by H2O2 with or without ultrasound (Fig. 7).

In addition, IBP adsorption is claimed to operate mainly

through interactions with the p-electron system of the

relation coefficient > 0.91) for AC sample (pseudo-first order,

pH = 3 pH = 7

132.90 67.140.087 0.1210.949 0.912138.10 69.75

) 0.000436 0.0011470.985 0.9522148 14920.092 0.183

0

20

40

60

80

100

120

140

160

180

0 1000 2000 3000 4000 5000 6000

Time (min)

Ibup

rofe

n up

take

(mg

g-1

)

AC at pH=3

pseudo 1storderpseudo 2ndorderIntraparticlediffusionElovich

438 C A R B O N 5 4 ( 2 0 1 3 ) 4 3 2 4 4 3The fitted parameters of adsorption kinetics at pH 7 and 3

(Table 4) were calculated from the non linear regressions of

the integrated Eqs. (1)(3). The profiles of each fitted curve dis-

played in Fig. 6, in comparison with the experimental kinetics

at pH 3, were very similar to the ones of fitted kinetic plots at

pH 7 (not shown). Whatever the pH value, the intraparticle

diffusion model (Fig. 6) was not in good agreement with the

experimental kinetics. The best fit of adsorption kinetics at

pH 3 (R2 = 0.985) was obtained for the pseudo-second-order

model. The IBP uptake decreased while the pH increased from

3 to 7. The pseudo-second-order constant k2 obtained at

pH = 7 was higher than at pH = 3 meaning that the IBP adsorp-

tion at pH = 7 required higher amount of AC than at pH = 3. As

a consequence, activated carbon AC has more affinity for IBP

at acidic pH. On the other hand, the Elovich model showed a

satisfactory correlation coefficient (R2 = 0.982 at pH = 3). Thus,

both pseudo-second-order and Elovich models are applicable

to reproduce the adsorption kinetics of IBP on AC at pH = 3.

Besides, the Elovich model was found to be the best model

at pH = 7 (R2 = 0.997). The adsorption initial rate of ibuprofen

(a) decreased as the pH of the ibuprofen solution became ba-

sic. The value of 1/b, indicative of the number of sites avail-

able for adsorption, was higher at pH = 3 (1/b 11) than atpH = 7 (1/b 5).

3.3. Adsorption isotherms

3.3.1. Effect of pHpH remains one of the most important factors controlling the

adsorption capacity of activated carbons, by influencing both

the chemistry of the surface groups and the IBP charge (pKa of

IBP equals to 4.91 [58]). Its impact was subsequently studied at

Table 4 Results of the best IBP kinetic fittings (R-square corpseudo-second order and Elovich models).

Kinetic model Parameters

Pseudo-first-order qe (mg g1)

k1 (min1)

R2

Pseudo-second-order qe (mg g1)

k2 (mg g1 min1

R2

Elovich a (mg g1 min1)b (g mg1)

R2pH pKa-2 and pH pKa+2 (i.e. pH = 3 and 7). For pH valueshigher than pKa(IBP), the molecule was present under its an-

ionic form. The pHPZC of AC was equal to 7.76 (Table 1). At

pH < pHPZC, the activated carbon was positively charged.

Thus, at pH = 7, the negatively charged IBP species were at-

tracted by the positive AC surface (electrostatic interactions),

which was not the case at pH = 3. At such a value, the surface

of AC remained positively charged (pH < pHPZC) whereas IBP

moieties existed under their neutral form disabling the favor-

able electrostatic interactions between both. However, the

dominant dispersive interactions explained the higher

adsorption capacity at this low pH. Indeed, the uptake of

IBP by AC was found higher at pH = 3 (Fig. 7, qmax = 130 -

mg g1) than at pH = 7 (Fig. 7, qmax = 75 mg g1).

Whatever the explored pH, adsorption capacities of the

raw and the modified carbons were different. As the studied

carbons probed by N2 adsorption at 77 K and CO2 adsorption

at 273 K were very similar in term of microporous and ultra-

microporous accessible volumes (Section 3.1.2, Table 3), nota-

bly lower than the uptake volume of IBP (0.11 cm3 g1,Section 3.1.2); the difference in uptakes of the various carbons

can only be attributed to the variety of their surface

chemistry.

While the bleaching treatment led to a decrease of the

Fig. 6 Fitted profiles of ibuprofen adsorption kinetics of AC

sample at pH 3 () by pseudo-first-order, pseudo-secondorder, Elovich and intraparticle diffusion models.0.982 0.997

pH=7

0

20

40

60

80

100

120

140

160

180

0 20 40 60 80

Ibup

rofe

n up

take

(mg/

g)

an

P]

0 1 3 ) 4 3 2 4 4 3 439graphitic rings. One thus would expect that additional acidic

surface oxygen groups might have withdrawn the electrons

from the p-system. However, this explanation is not fully sat-

isfactory here as the oxidation by H2O2 treatment led to an in-

crease in the adsorption uptake.

Indeed, the knee of the adsorption isotherm of AC[H2O2]US

was more pronounced compared to any other samples indi-

cating that strong interactions with this sample occurred

even at low concentration. As both XPS characterization and

Boehm titrations (Section 3.1.1) showed that the main differ-

ence in terms of resulting surface chemistry, between the

H2O2 and NaOCl oxidation treatments, laid in the amount of

carbonyl and lactonic groups; the higher IBP uptake could

be thus related to the higher amount of carbonyl (also from

quinones) and lactonic groups, created after H2O2 oxidation.

These carbonyl groups could have interacted with the aro-

pH=3

0

20

40

60

80

100

120

140

160

180

0 20 40 60 80Equilibrium solution concentration (mg/L)

Ibup

rofe

n up

take

(mg/

g)

Fig. 7 Adsorption isotherms of ibuprofen at 25 C onto rawAC[H2O2]US, *: AC[NaOCl] and m: AC700N2) at pH 3 and 7 ([IBtime = 5 days, stirring speed = 650 rpm).

C A R B O N 5 4 ( 2matic ring of ibuprofen in a donoracceptor bond as previ-

ously reported [40]. In fact, the oxygen lone pair electrons

are known to be responsible of the formation of donoraccep-

tor complexes via a dipole moment interaction; the carbonyl

groups acting as the electron donor and the aromatic ring act-

ing as the acceptor. As the carbonyl oxygen groups have more

intense dipole moments than carboxylic acid oxygens, they

have been reported as stronger donors [33]. In their molecular

form, the carboxylic acid could also interact in a different

manner by Van der Waals, hydrogen bonds with the carbox-

ylic groups of IBP, enabling the formations of dimers [34].

The functionalization of activated carbons by oxidation with

NaOCl yielded to specific oxygen surface groups (mainly phe-

nolic). Some of these groups might be inappropriate for ben-

eficial interaction with IPB, leading to a decrease of the

dispersive interactions, through decreasing the p-electron

density of the graphene layers [59]. Thus the uptake of

AC[NaOCl] remained rather similar to AC at pH 3 and lower

at pH 7.

A significant increase of the adsorption capacity compared

to AC was observed for AC700N2 aged sample: the uptake,

equal to 92 mg g1 at pH = 7, reached a maximum of

160 mg g1 at pH = 3, showing that the adsorption was alsopromoted on slightly oxidized activated carbons. Moreover,

at pH 7, the electrostatic attraction contribution might in-

crease slightly in the AC700N2 sample due to the shift of

pHPZC towards basic values (pHPZC of AC700N2 equals to

9.08). At pH 3, the increase in adsorption uptake for

AC700N2, was attributed to beneficial interactions of IBP with

the oxygenated surface groups (mainly carbonyl groups).

3.3.2. Effect of temperatureThe effect of temperature was studied at pH 3 in order to

determine the thermodynamic parameters of adsorption of

IBP on three different activated carbons (Figs. 810: samples

AC, AC[H2O2]US and AC700N2, respectively) and to compare

the interactions energies.

As the temperature increased from 298 to 328 K, the

adsorption capacities increased from 130 to 180 mg g1 for

Equilibrium solution concentration (mg/L)

d modified activated carbons (: AC, d: AC[H2O2], j:= 5100 mg L1, madsorbent = 10 mg, VIBP = 30 mL, contactAC (Fig. 8), from 140 to 160 mg g1 for AC[H2O2]US (Fig. 9)and from 163 to 230 mg g1 for AC700N2 (Fig. 10).

To investigate the temperature dependency of the adsorp-

tion capacity, four equilibrium models were analyzed, including

Langmuir, Freundlich, LangmuirFreundlich and Redlich

Peterson. All the isotherms were simulated using an iterative

0

50

100

150

200

250

0 20 40 60 80 100Ce (mg L-1)

Qad

s(m

g g-

1 )

Fig. 8 Experimental adsorption isotherms of ibuprofen on

AC at pH 3, and at 298 K (), 313 K (m) and 328 K (j) fitted bythe LangmuirFreundlich model (lines).

0 10

50

100

150

200

250

0 10 20 30 40 50 60 70 80Ce (mg L-1)

Qad

s(m

g g-

1 )

Fig. 9 Experimental adsorption isotherms of ibuprofen on

AC[H2O2]US at pH 3, and at 298 K (), 313 K (m) and 328 K (j)fitted by the LangmuirFreundlich model (lines).

150

200

250

s(m

g g-

1 )

440 C A R B O N 5 4 ( 2procedure based on a non-linear least-squares algorithm. The

Langmuir adsorption isotherm equation, expressed as follows

requires for its applicability a mono-layered coverage on the

surface of adsorbent [60]:

qe qmaxKLCe1 KLCe

where qe, KL, qmax and Ce are the ibuprofen uptake at equilib-

rium (mg g1), the Langmuir constant (L mg1), the monolayeradsorption capacity (mg g1) and the solution concentrationat equilibrium (mg L1), respectively.

The Freundlich equation is applicable for multi component

adsorption. The Freundlich isotherm is expressed by:

qe KF C1=ne

where KF is the Freundlich constant (mg1-1/n L1/n g1) and n is

the Freundlich exponent.

The Freundlich equation was successfully used to fit the

experimental data for many fluidsolid systems, but one of

the main drawbacks of this equation is that it doesnt tend

to a limiting value while the equilibrium concentration in-

creases. The LangmuirFreundlich equation is a way to over-

come this problem [61]. It is expressed as follows:

qe qmaxKLC

ne

1 KLCne

0

50

100

0 10 20 30 40 50 60 70 80Ce (mg L-1)

Qad

Fig. 10 Experimental adsorption isotherms of ibuprofen on

AC700N2 at pH 3, and at 298 K (), 313 K (m) and 328 K (j)fitted by the LangmuirFreundlich model (lines).The RedlichPeterson isotherm, proposed by considering

the limitations of Freundlich and Langmuir isotherms [62],

is given by:

qe KRCe

1 aRCbewhere KR (L g

1) and aR (L mg1) are the RedlichPeterson iso-

therm constants and b is the exponent, ranging between 0

and 1.

The experimental data of IBP adsorption on AC and modi-

fied carbons at pH 3 were well fitted by the LangmuirFreund-

lich and RedlichPeterson models. In terms of R2, the

LangmuirFreundlich model provided better fitting

(R2 > 0.96). Figs. 810 present the fittings by the Langmuir

Freundlich equation of isotherms at pH 3.

The estimated parameters for the LangmuirFreundlich

model are reported in Table 5.

The adsorption capacities, qmax increased as temperature

increased for each carbon indicating an endothermic adsorp-

tion. The LangmuirFreundlich constants KL were the highest

for AC[H2O2]US suggesting that this activated carbon has a

higher adsorption affinity to IBP than AC and AC700N2. Com-

petitive adsorption of water molecules and IBP might have

also occurred as the carbon became more hydrophilic upon

oxidation by H2O2 with ultrasonic irradiation [34]. Moreover,

as n value is near to unity for AC[H2O2]US, the adsorption pro-

cess of IBP rather follows the Langmuir model. The n increase

for AC and AC700N2 together with temperature suggested

that adsorption affinity of IBP for the substrates increased

as heating.

3.4. Thermodynamic parameters

The isosteric Gibbs free energy DG of adsorption of IBP ontothree activated carbons (AC, AC[H2O2]US and AC700N2) was

calculated using the results obtained from the simulations

(Table 5) and the following equation:

LnKd LnCa=Ce DG=RTwhere Kd is the equilibrium constant, Ca is the amount of IBP

adsorbed at equilibrium (mg L1), Ce is the concentration ofIBP remaining in the solution at equilibrium (mg L1), T isthe solution temperature (K) and R is the perfect gas constant.

The isosteric enthalpy DH and entropy DS of adsorptionwere also calculated from the slope and the intercept of the

plot of Ln(Kd) versus 1/T, respectively, using the equation:

DG = DH TDS. Thermodynamic parameters were calcu-lated for the adsorption of IBP onto raw and modified acti-

vated carbons at constant adsorption uptake value equal to

100 mg g1 and at pH = 3 (Table 6).At pH = 3, the adsorption capacity of IBP onto all studied

activated carbons increased as the temperature increased

from 298 K to 328 K, that was consistent with the positive val-

ues obtained for DH. This suggested that the adsorption ofIBP by the three activated carbons was an endothermic pro-

cess at this pH in agreement with the work of Dubey et al.

[30] for the adsorption of IBP at natural pH on an activated

carbon prepared from Artemisia vulgaris. At pH 3, the esti-

3 ) 4 3 2 4 4 3mated values for DG were 1.1, 7.3 and 4.5 kJ moL1,respectively for AC, AC[H2O2]US and AC700N2, indicating a

pti

A

f ad

0 1spontaneous physisorption process [63]. The DH values wereless endothermic as increasing the amount of oxygenated

groups, suggesting an exothermic interaction of some of

them with IBP.

The positive values of DS showed an increase of the disor-der at the solid-solution interface, probably due to the desorp-

tion of solvent molecules (water and methanol) prior to IBP

adsorption. The positive values of DS could indicate the com-petitive adsorption between IBP and solvent molecules.

The highest absolute value of Gibbs energy (7.3 kJ mol1)for the AC[H2O2]US sample suggested the presence of high en-

ergy adsorption sites related to particular oxygenated groups

formed by sonication in hydrogen peroxide. The characteriza-

tion of the surface chemistry suggested that these sites might

be carbonyl (or lactonic) groups mainly formed through

hydrogen peroxide oxidation. The increase in the adsorption

Table 5 Parameters of the LangmuirFreundlich fits of adsorat temperature in the range 298328 K (qmax in mg g

1).

T (K) AC

qmax KL n qmax

298 160.0 0.61 0.25 146.6313 179.1 0.45 0.83 152.0328 185.2 0.75 0.86 159.9

Table 6 Isosteric Gibbs free energy, enthalpy and entropy oQads = 100 mg g

1) at pH 3.

Sample AC

DG (kJ mol1) 1.1DH (kJ mol1) 74.6DS (kJ K1 mol1) 254.1

C A R B O N 5 4 ( 2energy compared to pristine AC (4.5 kJ mol1 compared to1.1 kJ mol1) was also observed for aged AC700N2, maybethrough the interaction with quite similar beneficial oxygen-

ated groups, but in a lower amount.

4. Conclusion

We showed that the adsorption of ibuprofen occurred prefer-

entially into the ultramicropores of the raw and N2 heat trea-

ted activated carbons. The adsorption of IBP is strongly

dependent on pH and is promoted by dominant dispersive

interactions which are increasing when the molecule is pres-

ent on its neutral form. For this reason at pH 7, as the electro-

static interaction predominates, the adsorption uptake

subsequently decreases compared to pH 3 where the adsorp-

tion was found endothermic and entropically promoted as

the competition with solvent molecules occurs. We brought

out that some particular surface chemistry nature, such as

carbonyl or lactonic or quinone groups can influence posi-

tively the adsorption uptake of IBP.

As an example, oxidation treatment in hydrogen perox-

ide combined or not with ultrasonic irradiation enables

increasing slightly the adsorption capacity, and the Gibbsenergy of adsorption at acidic pH. This might be attributed

to a donoracceptor mechanism between the created car-

bonyl groups (also in lactonic and carboxylic groups) and

the p aromatic ring of ibuprofen. Both appropriate oxidation

or reduction treatments can increase the adsorption uptake,

but the reduction of the surface functionality (by annealing

at 973 K under N2) also gives a higher adsorption capacity

whatever the pH value as the carbon becomes slightly oxi-

dized after aging, favoring dispersive interactions. However

treatment with NaOCl does not benefit to ibuprofen uptake

because it preferentially leads to the formation of phenolic

groups and a lack of carbonyl adsorption sites [32]. The

LangmuirFreundlich model allows describing well the ibu-

profen adsorption at pH 3. DH values indicate that theadsorption of ibuprofen on the studied activated carbons

is an endothermic process, in relation with solvent

AC[H2O2]US AC700N2

7.3 4.58.1 31.9

51.6 122.2sorption of ibuprofen (at constant adsorption uptake valueon isotherms of ibuprofen on the different activated carbons

C[H2O2]US AC700N2

KL n qmax KL n

1.23 0.95 190.7 0.30 0.700.98 1.52 212.5 0.46 0.871.22 1.08 256.2 0.35 0.87

3 ) 4 3 2 4 4 3 441interaction.

Acknowledgments

Generous MIRA grant of the Region Rhone-Alpes (France) is

thankfully acknowledged.

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C A R B O N 5 4 ( 2 0 1 3 ) 4 3 2 4 4 3 443

The effects of the surface oxidation of activated carbon, the solution pH and the temperature on adsorption of ibuprofen1 Introduction2 Experimental2.1 Ibuprofen2.2 Raw and modified activated carbons2.3 N2 adsorptiondesorption at 77K and CO2 adsorption at 273K2.4 pHs of carbons suspended in water2.5 Boehm titrations2.6 X-ray photoelectron spectroscopy2.7 Adsorption experiments: kinetics and isotherms

3 Results and discussion3.1 Characterization of activated carbons3.1.1 Surface chemistry3.1.2 Porosity characterization

3.2 Adsorption kinetics of raw activated carbon3.3 Adsorption isotherms3.3.1 Effect of pH3.3.2 Effect of temperature

3.4 Thermodynamic parameters

4 ConclusionAcknowledgmentsReferences