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Carbon 45 (2007) 3–10

Controlled preparation of Pd/AC catalysts for hydrogenation reactions

Amandine Cabiac a, Gerard Delahay a,*, Robert Durand a, Philippe Trens a,Bernard Coq a, Dominique Plee b

a Laboratoire de Materiaux Catalytiques et Catalyse en Chimie Organique, UMR 5618-CNRS-ENSCM, Institut Charles Gerhardt FR 1878, 8,

rue de l’Ecole Normale, Montpellier Cedex 5, Franceb Groupement de Recherches de Lacq – ARKEMA, rte de Bayonne, 64170 LACQ, France

Received 28 July 2006; accepted 27 August 2006Available online 9 October 2006

Abstract

A commercial activated carbon was submitted to different liquid oxidizing treatments in order to create a high density of carboxylicgroups on the carbonaceous surface. Characterization of the modified activated carbon shows that it is possible to generate a large rangein acidic group concentration without significant change of the textural properties. This ability allows us to introduce, by ion exchange,palladium cations up to 7.3 wt.%. This way of preparation, coupled to soft formaldehyde reduction, leads to highly dispersed metallicpalladium catalysts. The resulting catalysts were tested in hydrogenation of cinnamaldehyde and exhibited a high hydrogenation effi-ciency compared to activities reported in the literature.� 2006 Elsevier Ltd. All rights reserved.

1. Introduction

Noble metal catalysts supported on activated carbon(AC) are widely used in the fine chemical industry. Thechemical inertness, low cost, and the possibility of recyclingcatalyst by recovery of the metal after burn off are the mainadvantages of carbonaceous materials [1]. Oxidizing treat-ments can be used to functionalize the carbon surfacethrough the formation of surface oxygen groups (mainlycarboxylic acids and phenolic groups). These functionsare often introduced to decrease the hydrophobicity ofthe support and enhance the adsorption properties oforganic compounds and heavy metal in purification pro-cesses from water solutions [2,3]. It seems interesting totake advantage of this particular property to the adsorp-tion of noble metal leading to highly dispersed catalysts[4]. For this purpose, both the chemical nature of the sur-face and the porous structure are important. Chemical

0008-6223/$ - see front matter � 2006 Elsevier Ltd. All rights reserved.doi:10.1016/j.carbon.2006.08.016

* Corresponding author. Tel.: +33 467163480; fax: +33 467163470.E-mail address: gerard.delahay@enscm.fr (G. Delahay).

treatments of the support, prior to noble metal loading,shape the properties of the catalytic materials. It shouldbe pointed out that oxidizing treatments carried in liquidphase mainly favor the carboxylic groups formation. Theexact role of oxygen surface groups is not yet established.Recently Yasuda and Nishimura, while studying the prep-aration of Pt/Carbon black, have reported a strong corre-lation between the platinum content deposited and thenumber of surface acid carboxylic groups [5]. Roman-Martinez et al. also observed a stronger interactionbetween the platinum precursors and the surface oxygengroups of an activated carbon [6]. However Mojet et al.,while studying the nature of Pd-carbon fibrils by EXAFemphasized interactions between palladium tetraamineentity and carbon surface. The only role of carboxylicgroups appeared to be a stabilization of the complex [7].

Thus, anionic [8] or cationic [4,9] exchange has beenused for the palladium/activated carbon (Pd/AC) catalystpreparation, depending on the metal precursor and surfaceproperties of the activated carbon support [10].

The first part of this study will deal with the choice ofthe oxidizing treatment through the characterization of

4 A. Cabiac et al. / Carbon 45 (2007) 3–10

the treated supports and therefore will be focused in theoptimization of the palladium cationic exchange capacityof the carbon support. After characterization of the result-ing Pd/AC with the required amount of noble metal, itsactivity in selective hydrogenation of cinnamaldehyde willbe compared to a catalyst prepared by a straightforwardimpregnation.

2. Experimental

2.1. Materials

The starting material was an activated carbon supplied by CECA(namely L3S). Carbon L3S was chemically modified by different oxidizingtreatments in order to create surface oxygen groups. The experimentalconditions are summed up in Table 1.

2.2. Characterization of the supports

The supports were characterized using nitrogen sorption at 77 K, aswell as temperature programmed desorption of the decomposition prod-ucts under helium (TPD-He) and determination of the point of zerocharge (PZC).

Equivalent BET specific surface areas were determined by nitrogenadsorption at 77 K with a Micromeretics ASAP 2000 apparatus on sam-ples (0.05 g) outgassed 7 h at 523 K and 10�4 Pa before analysis. The crosssectional area of nitrogen at 77 K was taken as 0.162 nm2. Evidence ofmicroporosity as well as the external surface area were estimated usingt-plot method with carbon black as Ref. [11].

The oxygen surface groups of the supports were characterized by TPD-He. For each analysis, 0.2 g of the support was heated (5 K/min) underhelium flow (25 mL/min) from room temperature up to 1073 K. The efflu-ent composition was continuously monitored by on line sampling by aquadrupole mass spectrometer (Pfeiffer Omnistar). The amounts of COand CO2 were obtained by using calibration with the decomposition ofa hydrated calcium oxalate standard. The desorption profiles of CO andCO2 are characteristic of specific surface groups (carboxylic acid, lactone,phenol, etc.) [12].

The pH values at the point of zero charge (pHPZC) were measuredusing the so-called ‘‘pH drift’’ method [13]. Activated carbon (0.04 g)was added to 10 mL of a 0.01 M NaCl solution. Initial solution pHs wereadjusted using aqueous HCl or NaOH. The equilibrium pH, reached after48 h, was measured and plotted versus initial pH. The pH at which the ini-tial pH is equal to the final pH is called pHPZC.

2.3. Ion exchange experiments

In order to determine the total sodium exchange capacity, a quantity ofactivated carbon was stirred in NaCl/NaOH solution (m/V = 6.6 g/L�1)at 500 rpm for 48 h at room temperature. The solution was then filtratedand washed several times with distillated water until pH reaches neutral.

Table 1Supports characterization before and after oxidation

Support Oxidizing conditions S‘‘BET’’ m2/g Vmicroa cm3/g

L3S – 901 0.24L3S-NA1 HNO3 1 M, reflux, 1 h 805 0.21L3S-NA13 HNO3 13 M, reflux, 1 h 790 0.21L3S-B NaOCl 30%, 273 K, 5 h 820 0.23L3S-OW H2O2, 298 K, 4 days. 749 0.19

a Calculated with t-plot method.b Estimated from TPD-He experiments.

The Na-exchanged solid was further back exchanged with HCl 0.1 M solu-tion and the desorbed Na quantified by atomic absorption analysis (AAS).The Na quantity was also determined by elementary analysis.

A Pd(NH3)4 (NO3)2 solution (Strem Chemical) was added to Na-exchanged supports at room temperature (pH = 6; mCA/V = 6.5 gCA/L).When equilibrium was reached, the suspension was filtrated, washed severaltimes with deionized water to remove weak adsorbed metal ions on the ACsurface and dried in oven at 353 K. The quantity of palladium adsorbed (qe)was determined by the difference of the metal ion concentration beforeexchange (Co) and after equilibrium and washing (Ce). Palladium analysiswas carried out with a Varian SpectrAA 220 atomic absorption spectropho-tometer equipped with palladium (Varian kPd = 244.8 nm) and sodiumlamp (Varian kNa = 589 nm). The Pd-exchanged solids were stirred with0.1 M HCl solution for 5 h. The amounts of palladium leached in solutionwere analyzed as described above.

2.4. Catalyst reduction and characterization

After introduction of the metallic precursor, the material was reducedeither in liquid media at room temperature with formaldehyde or in H2/Arflow at 523 K for 2 h.

The reducibility of the catalysts was determined by temperature pro-grammed reduction experiments (TPR) in a flow apparatus provided witha quadrupole mass spectrometer (Pfeiffer Omnistar). Prior to analysis, thesample (0.2 g) was flushed one hour with argon (25 mL/min) before heat-ing from 273 K to 523 K (20 K/min) in a H2/Ar mixture (3/97, 15 mL/min). The amounts of H2 consumed or evolved were continuouslymonitored.

XRD powder patterns were collected on a Brucker axs D8 Advancediffractometer using the Ka copper radiation (k = 0.15401 nm) with a stepsize of 0.02�, on a 15�–75� 2h range. Metal crystallite sizes were calculatedfrom the broadening of the Pd (111) peak by using the Scherrer formulawith Warren’s correction for the instrumental contribution.

TEM analyses were performed with a JEOL 1200 EX II instrumentoperating at 80–100 kV. The samples were prepared by suspending andsonicating the reduced catalysts in ethanol. A drop of the suspensionwas deposited on a carbon copper grid. In order to obtain the average pal-ladium particle size in surface, dS ¼ Rnid

3i =Rnid

2i , hundreds of particles

were measured on representatives micrographs.The catalyst activities were compared in benzene hydrogenation reac-

tion. As model structure insensitive reaction, benzene hydrogenation isdirectly correlated to the metallic surface exposed. The reaction was per-formed over catalyst bed (0.1 g) packed within a Pyrex tube of 1.4 cmdiameter. Before reaction, the catalyst was reduced in situ at the desiredtemperature. Benzene hydrogenation was carried out at 393 K using a par-tial pressure of hydrocarbon of 4.7 kPa and a partial pressure of hydrogenof 95.3 kPa. The effluents composition was continuously analyzed with amass spectrometer.

2.5. Catalytic tests

The liquid phase hydrogenation of cinnamaldehyde was carried out ina sealed stainless-steel autoclave (Engineers Autoclave, 100 mL) equippedwith mechanical stirring, sample port, gas inlet, and a vent. The prere-

SEXTa m2/g pHPZC CO2

b mmol/g COb mmol/g

391 8.2 0.07 0.44354 4.2 0.81 1.74350 3.9 1.68 2.08334 4 1.13 1.26337 4.3 0.76 1.95

A. Cabiac et al. / Carbon 45 (2007) 3–10 5

duced Pd/AC catalysts (0.05 g) were reactivated in solvent (cyclohexane)at 333 K or 423 K under 10 bars (1 MPa) H2 pressure for two hours inthe autoclave (1200 rpm). A solution of cinnamaldehyde diluted in cyclo-hexane (0.22 mol/L) was then introduced. The reaction was carried out at333 K under 10 bars (1 MPa) H2 pressure. The H2 pressure was main-tained constant in the reactor by H2 feed regulation. Samples were with-drawn periodically and analyzed on a gas chromatograph (HP-4890 D)using an RTX S-Amine capillary column (30 m · 0.25 mm · 0.52 lm).Undecane was used as internal standard.

3. Results and discussion

3.1. Carbon supports

3.1.1. Textural properties

Nitrogen adsorption at 77 K over the different materialsreveals type I adsorption isotherms according to theIUPAC classification (not shown here). The only differencebetween the different materials upon nitrogen adsorption isthe value of the saturation plateau obtained at high relativepressures, which indicates different micropore volumes. Asshown in Table 1, external surface areas and microporousvolumes of the materials are slightly modified after oxidiz-ing treatments. The decrease in surface area and microporevolume is generally explained by restriction of the pore vol-ume available for N2 uptake due to the formation of oxy-gen groups at the entrance and/or on the walls ofmicropores and by a possible destruction of microporewalls [14,15]. Nevertheless, the H2O2 treatment leads tothe higher change in textural properties of L3S.

3.1.2. Oxygen containing functional groups

3.1.2.1. Temperature programmed decomposition. From theheating with a linear ramp of temperature of activated car-bons under inert atmosphere, CO, CO2 and H2O emissionsevolve [12]. Generally, CO2 appears at lower temperaturesthan CO and results from the decomposition of carboxylic,anhydride and/or lactone groups. The CO2 and CO emis-sion profiles of the oxidized samples are reported inFig. 1. L3S raw material exhibited negligible quantities ofevolved CO and CO2 (not shown). The CO2 emission pro-

400 600 800 1000 1200

CO

2 E

mis

sio

ns

(a.u

.)

Temperature (K)

(I) (II) (III)

Fig. 1. CO2 and CO profiles evolving during TPD-He experimen

files of the different samples have the same shape and maycoarsely be divided into three parts. According to the liter-ature, CO2 releases are assigned to decomposition of differ-ent groups

• part I (below 600 K): decomposition of isolated carbox-ylic acids,

• part II: decomposition of lactone groups either presentin the support or generated during heating by condensa-tion of adjacent carboxylic acid and/or phenolic groups[16]. This rearrangement during heating is confirmed byH2O emission peak,

• part III: decomposition of anhydride groups.

The efficiency of surface oxidation treatments to gener-ate carboxylic groups is the following: HNO3 > -NaOCl > H2O2. Further, it can be noted that an increasein nitric acid concentration doubles the final amount of thiskind of oxygenated group.

The CO thermodecomposition profiles of the materialsafter oxidation are similar. The peak centered around1000 K results from the decomposition of a broad familyof compounds: quinone, pyrone, carbonyl, phenolic andanhydride. The amount of CO desorbed is stronglyenhanced after oxidation and always stays higher thanthe quantity of CO2 produced. From the textural and sur-face chemistry characterization it can be concluded that thebest balance towards CO2 is obtained with the bleachtreatment.

3.2. Preparation and characterization of Pd� supported

on oxidized L3S

3.2.1. Ion exchange experiments3.2.1.1. IEP and PZC. The ion exchange experiment condi-tions were determined by measuring the Zeta potential(IEP) and the point of zero charge (PZC). The pHPZC ofthe starting L3S material was 8.2. After treatment, thisvalue decreased in the range of 3–5 depending on the

400 600 800 1000 1200

CO

Em

issi

on

s (a

.u.)

Temperature (K)

ts: (h) L3S-NA13; (s) L3S-NA1; (D) L3S-OW; (q) L3S-B.

6 A. Cabiac et al. / Carbon 45 (2007) 3–10

oxidizing agent. Above the PZC value, the overall surfacewas negatively charged. The cationic exchange could there-fore be achieved in a pH aqueous solution above 5.0, foroxidized samples, and above 9.0, for the initial activatedcarbon. The L3S pHIEP was 4.2, while for the oxidizingsupport, this value dropped down below 1.0 and thus forsome zeta potential measurements, it remained always neg-ative and pHIEP was never reached. Such differencesbetween pHPZC and pHIEP have already been observedand discussed previously [17,18]. The isoelectric point IEPis only representative of the external surface charge ofthe carbon materials whereas the point of zero chargePZC results from the internal and external surface chargecontribution. Generally, experimental exchange conditionswith pre-determined PZC values are in agreement withIEP measures.

3.2.1.2. Preexchange with sodium: Exchange capacity mea-

surement. A preexchange with sodium has been performedin order to determine the cation exchange capacity of thecarbon sample. This H+/Na+ exchange reaction is ratherslow and requires an equilibrium time of at least 48 h(checked before hand by kinetics measurements). Thequantity of sodium introduced increases with the amountof carboxylic functional groups (see Table 2). Fig. 2areveals a good correlation between the sodium uptakesand the number of acidic functions but the ratio Na+/H+

Table 2Cation uptakes of activated carbons before and after oxidation

Support Na uptake (mmol/g) Pd uptake (mmol/g) % Pd (wt.)

L3S 0.12 0.09 0.95L3S-NA1 0.43 0.29 3.08L3S-NA13 0.97 0.69 7.34L3S-B 0.80 0.44 4.68L3S-OW 0.39 0.27 2.87

0

0.2

0.4

0.6

0.8

1

1.2

1.4

0 0.2 0.4 0.6 0.8 1 1.2

Na

upta

ke (m

mol

. g-1

)

Acidic groupssurface (mmol. g-1)

L3S

L3S-NA13

L3S-B

L3S-OWL3S-NA1

a b

Fig. 2. Relationship between: (a) Na uptakes vs. s

(mol/mol) of 0.75 deduced from the linear curve fit, is byfar lower than 1.00. This lower value may be explainedby: (i) an overestimation of the exchangeable H+ fromthe TPD-He analysis and (ii) a certain non accessibilityof the acidic functions to Na+ cations since some acidicfunctions can be incorporated in the microporous network.

3.2.1.3. Sodium exchange by palladium cations. After preex-change with sodium, the Na-activated carbon samples werefurther exchanged with palladium cations from tetraminepalladium nitrate aqueous solution. The preliminarykinetic studies showed that equilibrium was reached inone hour but five hours of stirring were preferred so thatequilibrium is not questionable. The adsorption isothermsof the different supports are shown in Fig. 3. After wash-ing, filtration and drying, the Pd/AC samples were backexchanged in HCl solution by H+ and more than 95% ofpalladium was released in solution. Therefore, very fewpalladium species remained trapped in the carbonaceousnetwork. It can also be assumed that an almost reversibleadsorption of palladium through an ion exchange mecha-nism involving the acidic functional groups occurred. Nareleased in solution after total palladium adsorption pre-cisely matched the initial sodium capacity. The relationbetween initial Na content and Pd maximum uptake isdescribed in Fig. 2b. The Na+/Pd2+ ratio, close to 1.5, isvery inferior to 2.0. A Na+/Pd2+ ratio equal to 2.0 corre-sponds to the following cation exchange reaction:

½PdðNH3Þ4ðH2OÞ2�2þ þ 2Naþ � � � ½�OOCR�2L3S

$ ½PdðNH3Þ4ðH2OÞ2�2þ . . . ½�ðOOCÞR�2L3Sþ 2Naþ

Nevertheless, hydrolysis of [Pd(NH3)4(H2O)2]2+ mayoccur with the formation, for example, of [Pd(OH)(NH3)4(H2O)]+ and may explain the lower Na+/Pd2+ ratio.Similar observations have been noticed with other metals[19–21]. Jia and Thomas while studying the Na+/Cd2+

0.2 0.4 0.6 0.8 1 1.2 1.40

0.2

0.4

0.6

0.8

1

1.2

1.4

Pd uptake (m

mol. g

-1)

Nasurface (mmol. g-1)

L3S

L3S-NA13

L3S-B

L3S-OWL3S-NA1

urface acidic groups; (b) Pd2+/Na+ exchange.

0.0 0.5 1.0 1.5 2.00.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

qe

(mm

ol/g

)

Ce (mmol/g)

Fig. 3. Adsorption isotherms of palladium onto different oxidisedsupports: (h) L3S-NA13; (X) L3S-B; (s) L3S-NA1; (D) L3S-OW.

523523473

CO

2 E

mis

sio

n (

a.u

.)

H2 C

on

sum

pti

on

(a.

u.)

Temperature (K)373

Fig. 4. Temperature Programmed Reduction profiles of: Pd/L3S-B and(j) Pd/L3S.

A. Cabiac et al. / Carbon 45 (2007) 3–10 7

exchange found a ratio equal to 2 at low equilibrium con-centration but noticed a slight decrease at higher concen-trations [19]. The adsorption mechanism is not wellknown; nevertheless it is pretty sure that carboxylic acidgroups are involved in the palladium adsorption. Withoutthese functions, the L3S activated carbon only adsorbs0.09 mmol/g, eight times less than after HNO3 treatment.

3.2.2. Catalyst temperature programmed reduction

The ionic exchange experiments allowed finding theoptimal preparation method of a 4 wt.% Pd/AC catalyst.The NaOCl oxidation was preferred for practical, safetyand economical reasons. Moreover, this procedure doesnot modified textural properties as some other oxidantsdo. For the sake of comparison, a 4 wt.% Pd/L3S, whosethe carbonaceous support was not oxidized before, wasprepared by impregnation in the aqueous phase with thesame palladium precursor.

After drying at 353 K, these two catalysts were reducedeither in formaldehyde solution at room temperature or ina H2/Ar (3/97) flow at 523 K and 773 K. Temperature pro-grammed reduction profiles in 3% H2/Ar flow are shown inFig. 4. In addition to the hydrogen consumption, the for-mation of CO2, H2O and CH4 were monitored. Reductionof palladium occurred at a lower temperature in the sam-ple, which does not contain significant amount of carbox-ylic functions. Therefore, palladium cations in the samplePd/L3S-B, are more difficult to reduce since they are betterstabilized. This stabilization may be obviously attributed tothe (�OOC) surface groups. In the course of this reduction,carboxylic functions decompose which is followed by thereduction of the carbonaceous surface into CH4. However,palladium cation begins to be reduced before the appear-ance of CO2. After reduction, the specific surface areaincreases with reduction temperature. The textural proper-ties of Pd/L3S-B-773 looked like those obtained with thestarting material (Table 3). Indeed, the surface functions

decomposed gradually leading to the recovery of the micro-porous network.

3.2.3. Palladium particle size

Fig. 5 shows XRD patterns of the Pd/L3S-B reduced bydifferent methods. After reduction with formaldehyde atroom temperature, no Pd0 diffraction peak was detected,which is an evidence of particle size inferior to 2 nm. Thereduction of sample in H2/Ar flow (3/97) at 523 K resultsin the presence of a diffraction peak indicating a particlesize broadening. A higher reduction temperature did noteven slightly modify the diffraction peak intensity. Pd/L3S (4 wt.%), reduced in formaldehyde, exhibited an aver-age palladium particle of 11.4 nm (see Fig. 5). In this case,soft reduction conditions are not sufficient to generate highpalladium dispersion. The absence of acidic groups on thecarbonaceous surface seems to be the determinant param-eter. This ascertains the efficiency of the ion exchange tech-nique to prepare well dispersed 4wt.% Pd catalyst.

Transmission electron micrographs of Pd/L3S-B pro-vided more obvious evidence of the influence of the reduc-tion procedure over the particle sintering (Table 3). Theaverage metal particle size shifted from 1.5 nm up to3.8 nm after reduction at 523 K. This difference could beexplained by carboxylic acid groups decomposition duringreduction at high temperature as shown above. The H2/Pdratio during reduction was superior to 1, which is the stoe-chiometric ratio to reduce Pd2+ in Pdo. Dantas Ramoset al. observed the same behavior with carbon supportedpalladium catalysts [22]. They attributed the hydrogenoverconsumption either to (i) spillover phenomena, cata-lyzed by the surface oxygen groups: hydrogen that spillsover the palladium could stay adsorbed to the surface orto (ii) the reduction of surface groups and hydrogasifica-tion reactions. This last point confirms the role of carbox-ylic anchoring groups to stabilize palladium cations. It is

Fig. 5. XRD patterns of Pd/AC: (a) L3S-B-F; (b) L3S-B-523; (c) L3S-IMP-F.

Table 3Some characteristics of Pd/L3S-B catalyst reduced at different temperatures

Support Reduction temperature (K) S‘‘BET’’a (m2/g) Vmicro

a,b (cm3/g) dPd (TEM) (nm) dPd (XRD) (nm) rBc (lmol/gPd/s�1)

L3S-B-F 298 813 0.21 1.5. nd. 52L3S-B-523 523 880 0.24 3.8 4 8.42L3S-B-773 773 922 0.25 4 3.8 6.38L3S-IMP-F 298 729 0.17 20 11.4 1.38

a Data estimated with corrected masses (mcorrected = mcatalyst · (1-%Pd)).b Calculated with t-plot method.c Rate of benzene hydrogenation.

8 A. Cabiac et al. / Carbon 45 (2007) 3–10

well known from the literature that the benzene hydrogena-tion is considered as structure insensitive in the case of sup-ported palladium catalysts. The relative rate (rB) is thendirectly linked to the number of accessible active sites[23,24]. This experiment is another way to provide informa-tion about the dispersion and the metal particle size. Theresults are summed up in Table 3. The reaction ratedecreases by a factor seven after reduction at high temper-ature (from ca. 50 to 7 lmol/gPd/s�1). The impregnationmethod performed with a non-oxidised activated carbonyielded poorly dispersed catalysts as the benzene hydroge-nation shows (1.4 lmol/gPd/s�1). These values are in agree-

O

CAL

HCAL

COL

H2

H2

Fig. 6. Reaction pathway of the hyd

ment with Krishnanttucky and Vannice [25]. The resultsfrom TEM, benzene hydrogenation and XRD analysisconverge towards the fact that Pd/L3S-B-F catalyst is themost dispersed.

3.3. Catalytic tests

Prior to the reaction, the pre-reduced catalyst was acti-vated in situ at 333 K or 423 K under the reaction hydro-gen pressure. Typically, cinnamaldehyde (CAL)hydrogenation leads to a mixture of cinnamyl alcohol(COL), saturated hydrocinnamaldehyde (HCAL) and phe-nylpropanol (PP) (Fig. 6). With palladium catalysts, theadsorption of cinnamaldehyde through C@C bond isfavored [26]. This feature could also be observed since nocinnamyl alcohol was detected. The final products are thusmade of hydrocinnamaldehyde and the full hydrogenatedproduct.

The reduction temperature decreased the initial rate ofreaction by a factor ten between reduction at room temper-ature with formaldehyde and reduction in hydrogen flow at773 K (Table 4). The rates are lower than those found byGiroir-Fendler et al. while working with Pd ion-exchangedcatalysts (4 wt.% Pd/AC, dTEM = 1.7 nm) [27]. Howeverthe use of isopropanol/water as reacting media comparedto apolar solvent could explain this difference [26]. TheHCAL selectivity of 76–87% at 40% conversion is generallyhigher compared to those found in literature [28].

The temperature of the in situ reactivation processbefore the catalytic test also influences the activity. Thereis a two-fold decrease of the rate when the activation

O

OH

OH

PPH2

H2

rogenation of cinnamaldehyde.

Table 4Some properties of the catalysts in the hydrogenation of cinnamaldehyde

Catalyst Activationtemperature(K)a

V0 (mol/s/gPd) · 104

SelectivityHCAL (%mol)b

L3S-B-F 333 78 87423 39 86

L3S-B-523 333 43 85423 15 87

L3S-B-773 333 8 76L3S-IMP-F 333 36 80

a Temperature of in situ activation before the catalytic test.b Determined at 40% conversion.

Fig. 7. XRD patterns of L3S-B-F: (a) prior to activation; (b) after 333 Kactivation; (c) after 423 K activation.

A. Cabiac et al. / Carbon 45 (2007) 3–10 9

temperature goes up from 333 K to 423 K. The decrease ofrate can be assigned to an increase of Pd particle size. TheXRD analysis of L3S-B-F shows a shift from 2 nm to12 nm of the mean Pd particle size on the used catalystafter the catalytic run (Fig. 7).

4. Conclusion

The carbonaceous surface can be functionalized withcarboxylic groups in a broad range of content by modify-ing the conditions of the oxidizing liquid treatment. Thesecarboxylic groups are necessary functions, in order to dis-perse at an atomic level, palladium cations throughexchange reactions: H+/Pd2+ or H+/Pd(OH)+. Reductioninto mild conditions, for example in formaldehyde solutionat room temperature of a 4 wt.% Pd/AC leads to the bestdispersion of palladium metal particles by preventing sin-tering. The resulting catalyst contains palladium particlessmaller than 2 nm which are much more smaller than par-ticle size of a 4 wt.% Pd/AC whose the carbon surface wasnot oxidized.

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