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Research ArticleReceived: 7 September 2010 Revised: 6 October 2010 Accepted: 7 October 2010 Published online in Wiley Online Library: 8 November 2010

(wileyonlinelibrary.com) DOI 10.1002/jctb.2538

Highly stable Fe/γ -Al2O3 catalyst for catalyticwet peroxide oxidationPatricia Bautista, Angel F. Mohedano, Jose A. Casas, Juan A. Zazoand Juan J. Rodriguez∗

Abstract

BACKGROUND: A highly stable Fe/γ -Al2O3 catalyst for catalytic wet peroxide oxidation has been studied using phenol as targetpollutant. The catalyst was prepared by incipient wetness impregnation of γ -Al2O3 with an aqueous solution of Fe(NO3)3· 9H2O.The influence of pH, temperature, catalyst and H2O2 doses, as well as the initial phenol concentration has been analyzed.

RESULTS: The reaction temperature and initial pH significantly affect both phenol conversion and total organic carbon removal.Working at 50 ◦C, an initial pH of 3, 100 mg L−1 of phenol, a dose of H2O2 corresponding to the stoichiometric amount and1250 mg L−1 of catalyst, complete phenol conversion and a total organic carbon removal efficiency close to 80% were achieved.When the initial phenol concentration was increased to 1500 mg L−1, a decreased efficiency in total organic carbon removalwas observed with increased leaching of iron that can be related to a higher concentration of oxalic acid, as by-product fromcatalytic wet peroxide oxidation of phenol.

CONCLUSION: A laboratory synthesized γ -Al2O3 supported Fe has shown potential application in catalytic wet peroxideoxidation of phenolic wastewaters. The catalyst showed remarkable stability in long-term continuous experiments with limitedFe leaching, <3% of the initial loading.c© 2010 Society of Chemical Industry

Keywords: advanced oxidation; industrial effluents; heterogeneous catalysis; catalytic processes; environmental chemistry; wastewater

INTRODUCTIONWater quality regulations have become more stringent inrecent decades due to increasing social concern regarding theenvironment. Industrial effluents frequently contain pollutantsthat are toxic and resistant to conventional treatments. Phenoland phenolic compounds are among the most importantchemical pollutants from industry and are frequently used asmodel compounds for industrial wastewater studies.1 Oxidationtechniques have been applied for the abatement of these typesof pollutants via complete breakdown into carbon dioxide andwater or conversion to easily biodegradable by-products.2,3 Oneof the most effective is the Fenton process, which is based onthe generation of hydroxyl radicals from hydrogen peroxide withiron ions acting as homogeneous catalyst. The hydroxyl radicals,produced upon decomposition of hydrogen peroxide, act asstrong oxidants and are capable of oxidizing organic compoundsunder ambient conditions. One advantage of the Fenton processwith regard to other oxidation techniques is that no energy inputis necessary to activate hydrogen peroxide, making the reactionpossible at atmospheric pressure and at room temperature.Furthermore, this method works with relatively short reactiontimes and uses easy-to-handle reagents.4 Consequently, it hasbeen postulated as one of the most economic approaches5,6 andhas been successfully used with different industrial wastewaters.7

However, this treatment has some drawbacks derived from therelatively high cost of hydrogen peroxide and the fact that thehomogeneous catalyst, added as divalent iron salt, cannot beretained in the process, requiring further separation to prevent

additional water pollution. In order to avoid the continuous lossof catalyst and the need to remove the iron after treatment,which increases the cost, iron can be immobilized in a stableform on a convenient support giving rise to a solid catalystthat can be used in heterogeneous Fenton-like or catalytic wetperoxide oxidation (CWPO). So far, several supports have beenused, such as zeolites,8 – 10 pillared clays,11 – 16 alumina,17,18 silica,19

mesoporous molecular sieves,20,21 mesostructured SBA-15,22,23

ion-exchange resins24 and activated carbon,25,26 mostly in theCWPO of model phenolic compounds. Nowadays, the best resultshave been reported with zeolites, pillared clays and activatedcarbon. The main problem associated with these catalysts arisesfrom the leaching of the iron active phase.

The aim of this work is to evaluate the potential application of alaboratory synthesized Fe/γ -Al2O3 catalyst for CWPO using phenolas target compound. In earlier works, Al-Hayek et al.17 and Al-Hayekand Dore18 studied the CWPO of phenols and organic acids withFe/γ -Al2O3 catalysts, focusing attention on some features of thecatalyst such as calcination and reduction temperature and thechemical structure determined by Mossbauer spectroscopy. In thepresent work, in order to improve the catalyst performance and

∗ Correspondence to: Juan J. Rodriguez, Ingenier ıa Quımica, UniversidadAutonoma de Madrid, Ctra. de Colmenar Km 15, 28049 Madrid, Spain.E-mail: juanjo.rodriguez@uam.es

Ingenier ıa Quımica, Universidad Autonoma de Madrid, Ctra. de Colmenar Km15, 28049 Madrid, Spain

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establish optimal operating conditions in the CWPO of phenol, theinfluence of the operating conditions, pH, temperature, catalystand H2O2 doses, as well as the initial phenol concentration isanalyzed. The long-term stability of the catalyst is also evaluatedsince it is a main requirement for potential application.

MATERIALS AND METHODSCatalyst preparation and characterizationThe Fe/γ -Al2O3 catalyst was prepared by incipient wetnessimpregnation of γ -Al2O3 supplied by Merck (Germany) with anaqueous solution of Fe(NO3)3· 9H2O. The Fe load was adjusted to anominal 4% (w/w). After impregnation, the catalyst was left for 2 hat room temperature, dried for 18 h at 60 ◦C and calcined at 300 ◦Cfor 4 h. Fe loading and calcination temperature were optimized inprevious work dealing with CWPO of cosmetic wastewaters usingthis catalyst.27

The porous structure of the γ -Al2O3 and the fresh andused catalyst were characterized from the −196 ◦C N2 adsorp-tion–desorption isotherms using a Quantachrome Autosorb-1apparatus (Florida, USA). The samples were outgassed at 250 ◦Cfor 12 h to a residual pressure of 10−3 Torr. The iron content ofthe catalyst was determined by total reflection X-ray fluorescence(TXRF), using a TXRF spectrometer 8030c (FEI Co., Germany). X-raydiffraction (XRD) using a Siemens model D-5000 diffractometer(Germany) with Cu Kα radiation was used to determine thecrystalline phases present in both the catalyst and support.The elemental composition (C, H, N and S) of the fresh andused catalyst was determined with a LECO CHNS-932 analyzer(Michigan, USA). X-ray photoelectron spectroscopy (XPS) using aPhysical Electronics 5700C Multitechnique System (Germany) wasemployed for surface iron analysis.

Characterization of the porous structure of the raw aluminaand the synthesized Fe/γ -Al2O3 catalyst established mesoporouscharacter as the BET and external (or non-microporous) areas arealmost coincident and a very low micropore volume was measured.Impregnation of the support led to a moderate decrease of BETsurface area from 142 to 125 m2 g−1 for alumina and Fe/γ -Al2O3

catalyst, respectively. Analysis by TXRF confirmed the Fe content ofthe catalyst (3.9% w/w). The XRD patterns showed two crystallinephases, corresponding to γ -Al2O3 and hematite (α-Fe2O3). Ahematite mean crystallite dimension of around 20 nm wasdetermined from the Scherrer equation. XPS spectra confirmedthat the main iron species on the surface corresponds to Fe2O3

showing a well defined peak at 710.9 eV and its doublet separatedby 13.6 eV. The ratio of the superficial (XPS) to the total (TXRF) ironcontent of the catalyst was 1.17, indicating a fairly homogeneousdistribution of the active phase on the alumina particles.

CWPO experimentsThe oxidation experiments were carried out in batch modeusing 100 mL stoppered glass bottles placed in a thermostatedshaker, which maintained a constant temperature (±1 ◦C) with astirring velocity equivalent to 200 rpm. The catalyst was usedin powdered form (80 µm < dp < 100 µm). The variablestested and their corresponding ranges were: pHo = 3–9,[phenol]o = 100–1500 mg L−1, [H2O2]o = 250–1000 mg L−1,[catalyst] = 500–2500 mg L−1, T = 25–80 ◦C.

Long-term experiments were performed in a 1 L glass jacketedcontinuous stirred tank reactor at 5 mL min−1 flow rate over 100 husing 1250 mg of catalyst, 50 ◦C temperature and initial pH 3. In

order to avoid the loss of solid catalyst from the reactor a Swagelokfilter (Ø = 15 µm) was placed at the exit. The starting phenol andH2O2 concentrations were 100 and 500 mg L−1, respectively, thelatter corresponding to the stoichiometric amount of H2O2 forcomplete oxidation of phenol up to CO2 and H2O.

Reaction samples were withdrawn at different reaction times (upto 240 or 480 min in the batch runs), and immediately analyzed af-ter filtration through glass microfiber filters (Albet FV-C, Germany).Phenol and aromatic intermediates were identified and quantifiedby means of high performance liquid chromatography (HPLC;Varian Pro-Start 500, USA) with a UV-Vis detector. A Microsorb (TheNetherlands) C18 5 µm column (MV 100, 15 cm long, 4.6 mm diam-eter) was used as stationary phase and 1 mL min−1 of 4 mmol L−1

aqueous sulphuric solution (90%) and acetonitrile (10%) as mobilephase. A UV detector at 210 nm wavelength was used for phenol,catechol and hydroquinone and at 246 nm for p-benzoquinone.Short-chain organic acids were analyzed by ion chromatographywith chemical suppression (Metrohm 790 IC, Switzerland) usinga conductivity detector. A Metrosep (Switzerland) A s upp 5–250column (25 cm long, 4 mm diameter) was used as stationaryphase and 0.7 mL min−1 of an aqueous solution of 3.2 mmol L−1

Na2CO3 and 1 mmol L−1 NaHCO3 as mobile phase. Total organiccarbon (TOC) was determined using an OI Analitic TOC analyzer(model 1010, Texas, USA). The residual H2O2 concentration wasdetermined with a Varian UV/Vis spectrophotometer (model Cary50) at 410 nm as a complex with Ti4+.28 The iron concentration inthe liquid phase was analyzed by colorimetric tritation at 478 nmas a complex with NH4SCN in acidic medium, using a ShimadzuUV/Vis spectrophotometer (model UV-1603, Germany).

RESULTS AND DISCUSSIONFigure 1 shows the time-evolution of phenol and hydrogenperoxide in a typical CWPO experiment carried out at 50 ◦C and thestoichiometric amount of H2O2 with the Fe/γ -Al2O3 catalyst. Ascan be seen, the phenol concentration decreased from the starting100 mg L−1 value to less than 14 mg L−1 after 1.5 h, where ca 29%of H2O2 was decomposed. The primary oxidation intermediates(Table 1) were aromatic compounds (catechol, hydroquinone andp-benzoquinone) resulting from phenol hydroxylation, which thenevolve to short-chain organic acids that partially oxidize to CO2

and H2O. Less than 2 h was required for complete disappearanceof the highly toxic aromatic intermediates, p-benzoquinone andhydroquinone, whose ecotoxicity is, respectively, aproximatelythree and two orders of magnitude higher than that of phenol.29,30

The main organic acids responsible for the residual TOC wereacetic, oxalic and formic, the latter two apparently resistantto oxidation under these operating conditions. Thus, althoughcomplete conversion of phenol was achieved in less than 2 h, aresidual TOC representing 22% of the initial value still remainedafter 8 h of reaction. The measured values of TOC (TOCexp) andthe amount of carbon computed from the residual phenol andthe identified oxidation by-products (�C) were almost balancedfor the first hour of reaction but some differences were observedincreasing as the reaction proceeded. This indicates the formationof unidentified species, which may correspond to condensationproducts, mainly from catechol, for reaction beyond the first hour.

The results obtained from elemental chemical analysis estab-lished that the adsorption of organic species on the catalystsurface was dependent on reaction time (Table 1). The amountof adsorbed carbon increased during 3 h reaction and decreased

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0 50 100 150 200 250

100

200

300

400

500

600

BlankFe/Al2O3 catalyst

H2O

2 (m

g/L

)

Time (min)

0 50 100 150 200 250

25

50

75

100

125

150BlankFe/Al2O3 catalyst

Ph

eno

l (m

g/L

)

Time (min)

Figure 1. Time-evolution of phenol and H2O2 in the presence and absence (blank) of the Fe/γ -Al2O3 catalyst ([phenol]o: 100 mg L−1, [H2O2]o: 500 mg L−1,[catalyst]: 1250 mg L−1, T : 50 ◦C, pHo: 3).

thereafter due to the oxidation process. Thus, taking this adsorp-tion of carbonaceous material into account, the percentage ofmineralization due to CWPO is also shown in Table 1.

An important question with regard to the application of thiscatalyst in practice is the possible leaching of iron under theoperating conditions. As can be seen in Table 1, Fe leachingwas always very low, ca 0.9 mg L−1 after 8 h reaction, whichcorresponds to less than 2% of the initial Fe load. This valueis considerably lower than that previously reported by Al-Hayeket al.17 in the CWPO of phenol with a Fe/γ -Al2O3 catalyst. In orderto check some possible contribution of the homogeneous reactionderived from this Fe leaching, a set of experiments was carried outat the same operating conditions but using 1 mg L−1 of dissolvedFe2+ instead of the Fe/γ -Al2O3 catalyst. The TOC reduction after4 h was 21%, which, although much lower than the value obtainedwith the Fe/γ -Al2O3 catalyst (44%) confirms the occurrence ofsome homogeneous reaction contribution. Nevertheless, the useof this Fe/γ -Al2O3 catalyst instead of dissolved Fe2+ salts could bea potential alternative for the oxidation of phenolic wastewaterswith H2O2, avoiding Fe loss and the need for its further removalfrom the final effluent. In order to check this, a comparison betweenthe two oxidation processes, homogeneous and heterogeneous,was carried out (Fig. 2) using the equivalent amount of catalystin both cases. The rate of mineralization is much higher in thehomogeneous process, where the TOC was reduced by 44% in1 h. At that time H2O2 was completely depleted. In the case of theheterogeneous process, 4 h of reaction were needed to achievesimilar results, but it is possible to reach a higher degree ofmineralization (ca 60% after 8 h) without significant leaching ofiron. Furthermore, the specific consumption of H2O2 (amount ofH2O2 consumed per unit TOC removed, w/w) is also depicted inFig. 2. In both processes the specific consumption of H2O2 slightlydecreased with reaction time, but H2O2 was always more efficientlyused in the heterogeneous process. This can be explained becauseof the lower rate of H2O2 decomposition in the heterogeneousprocess, which allows higher availability of hydroxyl radicals duringreaction time, i.e higher degrees of mineralization.

The curves of phenol and H2O2 conversion in Fig. 1 showan initial induction period of more than 30 min. It seems thata slight amount of iron in solution is necessary to promotethe oxidation process. This dissolved iron can initiate thedecomposition of hydrogen peroxide into hydroxyl radicalsfacilitating the hydroxylation of phenol, which gives rise to the

formation of dihydroxybenzenes (catechol and hydroquinone).These compounds seem to have a significant effect on the rate ofoxidation of phenol. In order to check this, a set of experimentswas carried out at the same operating conditions as in Fig. 1, butadding a small amount of catechol (5 mg L−1) (Fig. 3). The rates ofboth H2O2 decomposition and phenol oxidation were significantlyincreased over the first 30 min.

Influence of the operating conditionsFigure 4 shows the evolution of phenol and H2O2 concentrationswith reaction time at different initial pH values. The efficiency ofphenol oxidation was remarkably affected by pH so that completephenol conversion was achieved at pH 3 while it remained verylow after 4 h at pH 6 and 9 (Fig. 4). This trend is similar tothat in the homogeneous Fenton oxidation where a pH around3 has been found optimum. The formation of hydroxyl radicalupon H2O2 decomposition is dependent on pH. As the pH wasincreased from 3, the decomposition of H2O2 was predominatelyto O2, which is not capable of oxidizing phenol under the mildoperating conditions of these experiments. At pH 9 only a 15%H2O2 conversion was observed after 4 h reaction (Fig. 4). The sametrend has been reported for other iron catalysts using differentsupports, such as zeolites,8 pillared clays13 or activated carbon.25

Nevertheless, mesostructured SBA-1522 seemed to be more pHtolerant showing similar profiles of TOC conversion regardless ofthe initial pH within the 2.5–7.0 range.

A determining factor of the economy of Fenton and Fenton-likeoxidation processes is the H2O2 consumption. Thus, the H2O2

dose must be conveniently adjusted in order to optimize thatreagent consumption, taking also into account that the remainingunconverted amount can not be recovered and moreover has tobe removed before discharging the effluent since the ecotoxicitymust be controlled. Figure 5 shows the time-evolution of TOCand H2O2 conversion at different H2O2 doses corresponding to0.5, 1 and 2 times the stoichiometric amount. The rate of TOCconversion substantially increased with the H2O2 dose up tothe stoichiometric ratio but no significant improvement occuredbeyond that level. The rate of formation of hydroxyl radicalsis limited by the availability of active iron sites on the catalystsurface and on the other hand a higher concentration of thoseradicals favours scavenging reactions. As can be seen, there wasa residual TOC value which represents 25–30% of the initial evenat the highest H2O2 dose and 4 h reaction time. That residual TOC

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Tab

le1

.Re

sult

so

bta

ined

fro

mC

WPO

ofp

hen

olw

ith

the

Fe/γ

-Al 2

O3

cata

lyst

([p

hen

ol] o

:100

mg

L−1,[

H2

O2

] o:5

00m

gL−1

,[ca

taly

st]:

1250

mg

L−1,T

:50

◦ C,p

Ho

:3)

Tim

e(m

in)

Phen

ol

(mg

L−1)

Cat

ech

ol

(mg

L−1)

HQ

1

(mg

L−1)

p-B

Q2

(mg

L−1)

Mal

eic

(mg

L−1)

Ace

tic

(mg

L−1)

Oxa

lic(m

gL−1

)Fo

rmic

(mg

L−1)

TOC

exp

(mg

L−1)

�C

(mg

L−1)

XH

2O

2(%

)C

ads

(mg

L−1)

Min

eral

izat

ion

(%)

Fele

ach

ed(m

gL−1

)

010

00.

00.

00.

00.

00.

00.

00.

076

.676

.60.

0–

0.0

–

1510

00.

10.

00.

00.

00.

00.

00.

079

.377

.24.

2–

0.0

0.1

3095

.70.

20.

00.

00.

00.

00.

00.

075

.874

.27.

8–

1.0

0.2

6075

.78.

20.

35.

70.

02.

10.

02.

471

.668

.812

.54.

11.

10.

4

9013

.219

.31.

26.

40.

72.

40.

07.

050

.130

.728

.515

.814

.10.

4

120

0.8

3.5

0.0

0.5

1.3

2.8

0.0

9.6

33.2

7.4

46.1

22.8

26.9

0.4

180

0.0

0.1

0.0

0.0

1.2

3.0

0.9

10.2

24.2

4.6

59.5

24.0

37.0

0.5

240

0.0

0.0

0.0

0.0

1.4

3.3

3.5

8.9

22.0

5.1

66.8

20.9

44.0

0.7

360

0.0

0.0

0.0

0.0

1.3

3.5

7.2

6.5

19.2

5.6

75.6

16.9

53.0

0.9

480

0.0

0.0

0.0

0.0

0.4

3.7

10.1

3.9

17.1

5.4

81.2

13.6

59.8

0.9

1H

idro

qu

ino

ne

2p-

ben

zoq

uin

on

e

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0 50 100 150 200 250 300 350 400 450 500

20

40

60

80

100

Time (min)

x TO

C (

%)

Heterogeneous reactionHomogeneous reaction

2

4

6

8

10

12

14

16

18

Sp

ecif

ic H

2O2

con

sum

pti

on

(gH

2O2/

gT

OC

)

Figure 2. TOC conversion (solid symbols) and H2O2 specific consumption(open symbols) in the homogeneous ([Fe2+]: 50 mg L−1) and heteroge-neous ([Fe/γ -Al2O3 catalyst]: 1250 mg L−1) process ([phenol]o: 100 mg L−1,[H2O2]o: 500 mg L−1, T : 50 ◦C, pHo: 3).

0 50 100 150 200 250

20

40

60

80

100

120

Time (min)

Ph

eno

l (m

g/L

)

PhenolPhenol+catechol

0

100

200

300

400

500

H2O

2 (m

g/L

)

Figure 3. Time-evolution of phenol (solid symbols) and H2O2 concen-tration (open symbols) in the CWPO of phenol and catechol with theFe/γ -Al2O3 catalyst ([phenol]o: 100 mg L−1, [catechol]o: 5 mg L−1, [H2O2]o:500 mg L−1, T : 50 ◦C, pHo: 3).

corresponds to the aforementioned organic acids resistant toCWPO oxidation. Thus, using a H2O2 dose above the stoichiometricdoes not imply any advantage but actually results in a less efficientuse of the reagent. The amount of TOC removed per unit H2O2

(w/w) in the conditions of Fig. 5 was 0.11 at the stoichiometricratio and was reduced to half that value when doubling the H2O2

dose. Furthermore, the H2O2 remaining after treatment must beremoved prior to final discharge or biological treatment due tothe associated toxicity.

As expected, increasing the catalyst concentration increased therate of phenol conversion (Fig. 6) since there were more active sitesavailable for H2O2 decomposition into HO· radicals, as consistentlyshown in the time-evolution curves of H2O2 conversion in thesame figure. It is more important in that figure to point out therate increase in the early stages of reaction (the aforementionedinduction period). Increasing the catalyst concentration increasedsomewhat Fe leaching from 0.1 to 0.3 mg L−1 for 500 and2500 mg L−1 of catalyst, respectively, after 5 min reaction time.As indicated before, this can promote the generation of HO·radicals through homogeneous-phase decomposition of H2O2

thus allowing phenol hydroxylation.Figure 7 shows the results obtained at different temperatures in

the range 25 to 80 ◦C. The rate of TOC conversion increased with

temperature consistently with the rate of H2O2 decomposition. Amoderate increase of 15 ◦C, from 35 to 50 ◦C, had an enormouseffect on TOC removal. At 50 ◦C phenol and the aromaticintermediates were completely removed and only organic acids,mainly acetic, formic and oxalic, remained in the liquid after4 h. Thus, the ecotoxicity was drastically reduced. A residualpercentage (15–20%) of TOC remained even at the highesttemperature tested after 4 h of reaction, confirming the resistanceof these organic acids to oxidation. The induction period becameless noticeable as the temperature was increased and wasnot apparent at 80 ◦C. This can be explained as the result ofincreased Fe leaching, which enhanced early decomposition ofH2O2 into HO· radicals. Decomposition of H2O2 on the Fe activesites must be accelerated as well at increasing temperature.The relative distribution of the short-chain organic acids wastemperature dependent. After 4 h reaction time, oxalic acid,which was more resistant to oxidation, represented ca 97%of the total amount of organic acids at 80 ◦C, while at 50 ◦Cit was only ca 20%. The loss of Fe from the catalyst byleaching was somewhat increased with temperature, but didnot exceed 3 mg L−1 after 4 h, corresponding to around 6% ofthe initial Fe load. From the results obtained, a temperaturearound 50 ◦C can be considered as the optimum for thisprocess.

Table 2 summarizes the results obtained at different initialphenol concentrations. As can be seen, although completephenol conversion can be achieved independently of theinitial concentration, the percentage of mineralization decreasedsubstantially as that initial concentration was increased. Althoughthe remaining TOC corresponded mainly to organic acids oflow ecotoxicity, significant amounts of aromatic intermediatesremained after 4 h reaction time. These intermediates can becompletely converted by increasing the reaction time as indicatedby the shape of the concentration–time curves. Among the organicacids, oxalic deserves particular attention due to its quantitativeimportance and its detrimental effect on Fe leaching. However, thisFe leaching was always lower than that observed for Fe/activatedcarbon catalyst tested in CWPO of phenol.25

Stability of the catalystTo address catalyst stability, a long-term experiment was carriedout in a continuous stirred tank reactor at 50 ◦C and a space-timeof 3.92 kgcat h mol−1 phenol. The results are shown in Fig. 8. Thecatalyst showed a high stability since phenol and TOC reductionsclose to 100 and 50%, respectively, were maintained over the100 h on stream. Analysis by TXRF confirmed that the loss of Fefrom the catalyst by leaching was low, since the Fe content ofthe catalyst after that time on stream (3.7%) was similar to thatof the fresh one (3.9%). The same trend was observed regardingthe BET surface area, which was hardly modified after reaction for100 h. To gain more conclusive information the catalyst from thelong-term experiment was characterized for elemental chemicalanalysis. A small deposition of residual carbon-containing species(1.4% w/w) on the surface of the catalyst used was measured,although this did not seem to involve any significant blockage ofactive sites that could affect the activity of the catalyst. In previouswork, the XRD of the fresh and used Fe/γ -Al2O3 catalyst in CWPOof cosmetic wastewaters showed similar profiles except for a newpeak appearing in the used catalyst at 2θ around 18◦, which maycorrespond to carbonaceous deposits since it was removed uponcalcination at 550 ◦C.31

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0 50 100 150 200 250

20

40

60

80

100

Time (min)

Ph

eno

l (m

g/L

)

pH = 3pH = 6pH = 9

0 50 100 150 200 250

100

200

300

400

500

Time (min)

H2O

2 (m

g/L

)

pH = 3pH = 6pH = 9

Figure 4. Influence of pH ([phenol]o: 100 mg L−1, [H2O2]o: 500 mg L−1, [catalyst]: 1250 mg L−1, T : 50 ◦C).

Table 2. Results obtained at different initial phenol concentrations after 8 h of reaction ([H2O2]o: stoichiometric to [phenol]o, [catalyst]: 1250 mg L−1,T: 50 ◦C, pHo: 3)

Phenolo(mg L−1)

Xphenol(%)

XTOCexp(%)

Cads(mg L−1)

Mineralization(%)

XH2O2(%)

�aromatics(mg L−1)

�acids(mg L−1)

Oxalic(mg L−1)

Feleached(mg L−1)

100 100 77.6 13.6 59.8 81.2 0.0 18.1 10.1 0.9

500 100 60.0 21.8 54.3 67.6 1.2 130.0 40.7 3.9

1000 99.7 48.9 25.5 45.6 58.6 22.2 208.8 44.6 5.3

1500 99.4 45.1 27.6 42.7 54.2 74.9 312.7 67.4 7.4

0 50 100 150 200 250

25

50

75

100250 mg/L500 mg/L1000 mg/L

XT

OC ,

XH

2O2 (

%)

Time (min)

Figure 5. Time-evolution of TOC (solid symbols) and H2O2 (open symbols)conversion at different H2O2 doses ([phenol]o: 100 mg L−1, [catalyst]:1250 mg L−1, T : 50 ◦C, pHo: 3).

CONCLUSIONSA laboratory synthesized catalyst based on Fe supported onγ -Al2O3 has shown potential application for catalytic wet peroxideoxidation of phenolic wastewaters. Working at 50 ◦C, pH 3 andwith the stoichiometric amount of H2O2, complete removal ofphenol and aromatic intermediates was achieved with a highreduction of TOC (∼80%). The ecotoxicity was drastically reducedsince the remaining TOC corresponds to low weight organic acids,mainly formic, acetic and oxalic, of low toxicity. The reactiontemperature and the initial pH significantly affect both phenolconversion and TOC removal efficiency. A temperature of 50 ◦Cand pH of 3 can be considered as optimal. Increasing the initialphenol concentration up to 1500 mg L−1 did not affect phenol

0 50 100 150 200 250

20

40

60

80

100

Time (min)

Ph

eno

l (m

g/L

)

500 mg/L1250 mg/L2500 mg/L

20

40

60

80

100

XH

2O2 (

%)

Figure 6. Time-evolution of phenol concentration (solid symbols) andH2O2 conversion (open symbols) in the CWPO of phenol with theFe/γ -Al2O3 catalyst ([phenol]o: 100 mg L−1, [H2O2]o: 500 mg L−1, T : 50 ◦C,pHo: 3).

conversion but substantially reduced the percentage of TOCremoval and increased the Fe leaching as a consequence of higherconcentration of oxalic acid. The catalyst showed remarkablestability, since in long-term continuous experiments, phenol andTOC reductions were maintained constant over 100 h time onstream. Fe leaching from the catalyst over that time on streamwas lower than 3% of the initial load. Small amounts of carbon-containing species were deposited on the surface of the catalystbut did not affect its activity.

ACKNOWLEDGEMENTSThe authors greatly appreciate financial support from the SpanishMICINN through the projects CTQ2008-03988/PPQ and CTQ2007-

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0 50 100 150 200 250

100

200

300

400

50050 °C65 °C80 ºC

25 °C35 °C

H2O

2 (m

g/L

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Time (min)

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20

30

40

50

60

70

80

90

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%)

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Figure 7. Effect of temperature ([phenol]o: 100 mg L−1, [H2O2]o: 500 mg L−1, [catalyst]: 1250 mg L−1, pHo: 3).

0 20 40 60 80 100

20

40

60

80

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Phenol

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Figure 8. Long-term performance of the catalyst ([phenol]o: 100 mg L−1,[H2O2]o: 500 mg L−1, T : 50 ◦C, pHo: 3, τ : 3.92 kgcat h mol−1).

61748/PPQ. We also express our gratitude to the Consejerıa deEducacion y Ciencia of the CM for financial support to the project(S-0505/AMB/000395). The valuable contribution of Miss L.Gonzalez is also acknowledged.

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