Fuel 105 (2013) 695–704

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Fuel

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The role of surface acidity in adsorption of aromatic sulfur heterocycles from fuels

Sachin Nair, A.H.M. Shahadat Hussain, Bruce J. Tatarchuk ⇑Center for Microfibrous Materials Manufacturing (CM3), Department of Chemical Engineering, Auburn University, Auburn, AL 36849, USA

h i g h l i g h t s

" The sulfur adsorption sites of Ag/TiO2 and Ag/Al2O3 were acidic innature.

" The acid sites were generated afteractivation and Ag impregnation.

" The thiophenic species arepreferentially adsorbed on thesurface hydroxyl groups.

0016-2361/$ - see front matter � 2012 Elsevier Ltd. Ahttp://dx.doi.org/10.1016/j.fuel.2012.10.005

⇑ Corresponding author. Tel.: +1 3348442023; fax:E-mail address: tatarbj@auburn.edu (B.J. Tatarchuk

g r a p h i c a l a b s t r a c t

a r t i c l e i n f o

Article history:Received 13 July 2012Received in revised form 25 September2012Accepted 3 October 2012Available online 24 October 2012

Keywords:DesulfurizationAdsorbentTitanium dioxideSurface acidityBronsted sites

a b s t r a c t

The objective of this work is to investigate the relationship between sulfur adsorption capacity and sur-face acidity of supported silver oxide adsorbents from liquid hydrocarbon fuels. Inherently acidic sup-ports such as TiO2 and Al2O3 developed sulfur adsorption capacity following activation in air attemperatures above 350 �C. Addition of Ag to these supports further increased sulfur adsorption capacityby �40%. Efforts were made to understand the thermal activation process and the sulfur adsorptionmechanism. Breakthrough characteristics of various hetero aromatic molecules revealed that thestrengths of adsorption increased with electronegativities of the heteroatoms. Sulfur adsorption capaci-ties of these adsorbents were progressively reduced when poisoned with increasingly basic molecules.Surface acidity was measured using NH3 chemisorption and titration with 2,6-lutidine. Interactions ofprobe molecules on acid centers were also observed from IR spectroscopy where both Bronsted and Lewiscenters were observed. However interactions with 2,6-lutidine and trimethyl chlorosilane indicated thatthe active centers involved in sulfur adsorption were primarily Bronsted in nature. A clear correlationbetween activation temperature, surface acidity, and sulfur adsorption capacity was established for bothblank and Ag loaded samples.

� 2012 Elsevier Ltd. All rights reserved.

1. Introduction

Sulfur is not only a well known environmental pollutant, butalso a poison for catalytic units consisting of precious metal com-ponents such as PEM fuel cells. Use of high sulfur logistic fuelsfor portable power applications involving fuel cells is limited due

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+1 3348442065.).

to their low sulfur tolerance. Adsorption of sulfur at room temper-ature provides a viable solution for fuel clean-up prior to reforma-tion [1]. We have previously reported the substantial adsorptioncapacity of supported Ag adsorbents on TiO2 and Al2O3 for sulfurfrom fuels such as JP5 [2]. We also demonstrated the thermal rege-nerability of the adsorbent using air/hot exhaust gases/steam asstripping medium. This rather unique property of these materialsmakes them adaptable to a multi-bed desulfurizer configured forcontinuous operation. The materials were active in oxidized state

696 S. Nair et al. / Fuel 105 (2013) 695–704

in contrast with the majority of reported adsorbents which areused in reduced state. This accounts for their regenerability in oxi-dative conditions. The absence of activation steps also makes thematerials easier to handle and transport, simplifying regenerationprocedures.

A marked difference in chemistry and morphology from previ-ously reported adsorbents warranted a better analysis of the mech-anism of sulfur adsorption and chemical nature of the activecenters. Our previous work addressed physiochemical character-ization of these materials extensively [3]. XPS and TPR measure-ments showed that part of the Ag phase was present as an oxide.Moreover, TPR indicated that the amount of oxide was inverselyproportional to the amount of Ag dispersed on the support. Major-ity of the species at the adsorption interface was populated byoxide species. Oxygen chemisorption indicated that larger crystalsexisted at higher Ag loadings. Upon further investigation it wasalso observed that Al2O3 and TiO2 contributed to the sulfur capac-ity even in the absence of Ag phase [4]. Sulfur adsorption capacitywas generated by the calcination step during synthesis and was ab-sent in the ‘‘dried’’ adsorbent samples. A clear correlation betweencalcination and surface acidity was observed in our earlier work[5]. All these observations led to the premise that surface aciditymight play a significant role in sulfur adsorption using these mate-rials. Studies have shown that metal impregnation on catalyst sup-ports augments acidity [6,7]. Metal ions such as Cr, Fe, and Co havebeen demonstrated to increase the surface acidity of activated car-bons [8,9]. Modifications of Al2O3, SiO2 and MgO with other oxideshave been reported to influence their acid/base properties as well[9].

Acidity of both Lewis and Bronsted kind is common on Al2O3

and TiO2 surfaces. Lewis acid sites are coordinatively unsaturatedcenters which are electron acceptors, while Bronsted acid sitesare proton donors. The relationship between strengths and concen-trations of surface acid sites and their catalytic activities has beenestablished by earlier research [10]. For example, a linear relation-ship has been observed between surface acidity and photo-cata-lytic activity of TiO2 [11]. Mixed metal oxides also generatesurface acidity over their individual end members because of an in-crease in polarizability of surface hydroxyl (OH) groups [12,13].Mechanisms wherein terminal OH groups are populated on ana-tase have been presented by previous research [14–16]. Simula-tions of adsorption on TiO2 have shown that only hydroxylatedsurfaces are occupied by organic molecules [17]. Efforts were thusmade to measure the acidity of these Ag based adsorbents, distin-guish between Lewis and Bronsted kinds and then to establish acorrelation with their sulfur adsorption capacities.

Previously it was observed that sulfur selectivity of Ag/TiO2 var-ied significantly with fuel chemistry [18]. Adsorption capacity im-proved with the addition of aromatic rings to the solute moleculethat indicated a substantial contribution of p electron interactionsat the surface. Polarizable molecules like alkenes or aromatichydrocarbons with functional groups can thus overwhelm andcompete with sulfur heterocycles for adsorption sites. These elec-tron-donors may be used to study the acceptor properties of mostreactive surface sites [19]. The selectivity of the adsorbent towardelectron donors should establish the contribution of surface acidityin adsorption. Breakthrough studies of aromatic molecules withfunctional groups varying in their electronegativity values werecarried out to this effect.

NH3 chemisorption may be used to measure acidity of surfaces.Being a strong base, it is readily adsorbed on both strong and weakacid sites resulting from anion vacancies, and used as a probe mol-ecule to test the acidic properties of metal oxide surfaces [20–23].NH3 is chemisorbed on electron deficient metal atoms through theoverlap of free electron pair on NH3 and the vacant orbital of thecation. On TiO2, IR spectroscopy has revealed that NH3 adsorbs

strongly on Ti (IV) sites at low coverage. Lewis acid–base adductsare formed between NH3 and Ti4+ [24]. Strength of interactionweakens at higher coverage due to repulsive lateral interaction be-tween adjacent NH3 molecules [25]. On a hydrated surface, NH3

binds to a proton of a hydroxyl group through hydrogen bonding[24–26]. Dissociation of NH3 on contact with an oxide surfaceresulting in the formation of surface NH2 (or NH) and additionalOH species has also been observed [27]. Complete transfer of pro-tons to the NH3 molecules from Bronsted sites to produce adsorbedNHþ4 ions is also common on anatase [22]. Therefore, NH3 chemi-sorption was carried out to qualify the acid centers present onthe adsorbent surfaces.

A quantitative estimation of Bronsted acid sites was carried outusing standard acid base titration method [28] with 2,6-lutidine asa probe and Hammett indicators [29–31]. Measurement of surfaceacidity using probe molecules is an established and effective meth-od [21,32–35]. 2,6-lutidine is a sterically hindered amine and ispreferentially adsorbed on proton acid sites even in presence ofstrong Lewis acid sites, thus acts as a good adsorbate for theBronsted acid sites [36,37]. The effect of activation temperatureand addition of Ag phase to Al2O3, TiO2, and SiO2 on Bronsted acid-ity was established in this work. In addition, sulfur adsorptioncapacities were determined using JP5 fuel for comparison.

Sulfur removal using silver based adsorbents has been observedthrough three types of mechanisms; the formation of a sulfide onthe surface [38–41], strong chemisorption [42–44], physisorptionor a combination thereof. These mechanisms differ by the strengthsof interactions between the sulfur heterocycles and the active cen-ters. Multi-layer sulfide formation is a possible mechanism as ob-served in the tarnishing of silver objects in air. High selectivity fortrace amounts of sulfur in fuels also points to strong chemisorption.Mechanisms such as p complexation have been previously used toexplain sulfur removal by Cu–Y type zeolites [45–47]. Variation inelectron density of aromatic ring resulting from the sulfur atomhas been attributed to the higher adsorption selectivity of thesematerials. Investigating breakthrough performances, IR signaturesand acidity measurements plausible mechanisms regarding sulfuradsorption by silver adsorbents have been approached in this work.

2. Experimental

2.1. Preparation of adsorbents

TiO2, c-Al2O3, and SiO2 obtained from various vendors (Table 1)were crushed, sieved to size and dried in a convection oven for atleast 6 h at 110 �C prior to use. Wet impregnation was used to de-posit Ag on the support using AgNO3 (99.9%, Alfa Aesar) solution inwater. The concentration of the impregnating solution was ad-justed to obtain the required silver loading on the support. Theresulting particles were then dried at 110 �C for 6 h followed bycalcination in air at 400 �C for 2 h. Silver weight loading on theadsorbent was typically maintained at 4%. The ‘blank’ adsorbentsamples were prepared using identical synthesis procedures asfor Ag supported adsorbent samples, the only difference being thatimpregnation was carried out with dilute HNO3 of identical con-centration. Calcination of the supports in air (Airgas) is referredto as ‘activated’. Typical activation temperature used was 400 �Cunless mentioned otherwise.

2.2. Analysis of sulfur

Sulfur content was analyzed using an Antek 9000VS Total SulfurAnalyzer. The instrument was calibrated using standards of subse-quent dilutions of both real fuels as well as sulfur heterocycles in

Table 1Properties of TiO2, SiO2, and c-Al2O3 used for the preparation of silver sorbents.

Support Vendor Grade BET surface area (m2/g) Pore volume (cc/g)

TiO2 St. gobain nor pro Type 1 153 0.46SiO2 Grace davison Grade 21 319 0.80c-Al2O3 Alfa Aesar Catalyst support 256 1.2

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n-octane. The lower detection limit of the instrument was ob-served to be 200 ppbw.

2.3. Challenge fuels

JP5 with a total sulfur content of 1172 ppmw was obtained fromNAVSEA, Philadelphia. Model fuel consisting of 3500 ± 10 ppmwbenzothiophene (98%, Alfa Aesar) in n-octane (97%, Arcos Organics)was used for breakthrough experiments in various contexts. Modelfuels containing naphthalene, quinoline, and benzofuran (98% AlfaAesar) in n-octane were also prepared for breakthrough studies at3500 ± 10 ppmw.

2.4. Pretreatment with probe molecules for breakthrough experiments

Pretreatment with probe molecules was carried out ex situ byflowing the gas through a bed of adsorbent particles forapproximately 30 min. Volatile probe molecules such as pyridine,2,6-lutidine, and trimethyl chlorosilane (TMCS) (99%, Alfa Aesar)were introduced to the adsorbent by exposure to their vapor. Theadsorbent particles were placed in a vacuum chamber along withthe probe and pressure reduced to ca 100 mPa. Breakthroughexperiments were carried out with the treated adsorbent samplesafter an exposure time of 30 min.

2.5. Adsorbent breakthrough performance

The breakthrough characteristics of the adsorbents were deter-mined in a packed column configuration where the challenge fuelflowed from the bottom of the bed to the top. 10.0 g of adsorbentwas used in all breakthrough studies. The bed was contained in aquartz tube (ID 16 mm and length 62 mm) supported on both endsby quartz wool. None of the adsorbents were not activated in situnor was any step followed to wash the bed with a sulfur-free sol-vent to remove trapped air. The fuel was pumped by a peristalticpump at a flow rate of 0.5 ml/min (LHSV �2 h�1). The bed outputwas sampled at regular intervals for analysis of sulfur content.The concept of t½ was used to estimate the saturation capacityfor cases where sulfur removal was not carried out to bed exhaus-tion. The schematic diagram of the experimental setup can befound elsewhere [2].

2.6. Equilibrium sulfur capacity

For estimation of equilibrium sulfur capacity, Jet fuel JP5 wasagitated mechanically with adsorbent particles in sealed vials for48 h at ambient conditions with a fuel to particle ratio of 20 ml/g. Sulfur content of the remaining fuel was used to determine thesulfur capacity of the adsorbent samples.

2.7. IR spectroscopy

A Thermo Scientific Nicolet IR100 spectroscope was used to ob-tain infrared spectra of adsorbent samples treated with variousprobe molecules. The analysis range was between 400–4000 cm�1

and the spectral resolution was 4 cm�1. A customized IR cell withZnSe window was employed for this study. Each adsorbent sample

was pressed to produce a self-supporting disk before calcination inair for 450 �C for 2 h. The calcined disk was mounted onto the celland pretreated in dry air at 200 �C for 1 h followed by evacuation(at 100 mtorr or 0.013 kPa) to remove residual moisture. After this,the adsorbent sample inside the chamber was treated with probematerials (at room temperatures). Both pretreatment and treat-ment with probe molecules for IR analysis were performed in situ.

2.8. Ammonia chemisorption

Ammonia uptakes were measured using a Quantachrome Auto-sorb AS1 instrument. Acidity analysis included sample pretreat-ment (in situ) consisting of raising sample temperature at 10 �C/min in flowing air (zero grade, Airgas) and then holding at that tem-perature for 2 h. The dried support samples (SiO2, Al2O3, and TiO2)were calcined at 200, 400, and 600 �C, while the Ag supported adsor-bent samples were calcined at 400 �C. Following air treatment, sam-ple temperature was lowered to 175 �C and then evacuated for30 min. Pretreatment and analysis conditions were taken from liter-ature study [48]. Analysis was performed at 175 �C using a 10 com-bined point and 10 weak point titration method. The first isothermcollected after the sample preparation step represents the com-bined contribution of chemisorption (on strong sites) and physicaladsorption on the adsorbent sample. After obtaining the combinedisotherm, the system was evacuated to remove all physically ad-sorbed NH3. Chemisorbed gas is only removed from the surfacewhen evacuation is carried out at a much higher temperature thanthe adsorption temperature. The weak isotherm was obtained byadmitting NH3 following the evacuation step. The isotherm for thechemisorbed gas was thus obtained by subtracting the weak fromthe combined isotherm.

2.9. Potentiometric titration

0.2 g of the adsorbent sample was added to a series of vials(around eight to ten vials) and 10 ml of dry Benzene (99%, Alfa Ae-sar) was added to each. 0.01 M solution of 2,6-lutidine (98 + %, AlfaAesar) was added to each of the adsorbent samples and agitatedmechanically for 4 h. The equilibrated adsorbent samples weresubsequently tested with 0.1 wt% solution of methyl yellow (98%,Alfa Aesar) indicator (pKa = +3.3). After agitating for 4 h the colorof the solid adsorbent samples were observed. In presence of theindicator the adsorbent samples turned red when acidic(pKa < +3.3) and yellow when basic (pKa > +3.3). For every adsor-bent sample, the volume of 2,6-lutidine solution was increasedsequentially in each vial so as to determine the amount of base re-quired to neutralize the acid surface sites through titration. Thepreceding steps were repeated using smaller stepwise increasesin volume of base between the ranges found in the previous trialand the amount of base was bracketed to even smaller intervals.Following this procedure, the titer values were approximated towithin ±0.5 lmoles/g for TiO2 samples and ±10 lmoles/g forAl2O3 samples. The interval for silver supported adsorbent sampleswas ±25 lmoles/g. Thus the end points were determined fromchange in color of adsorbed Hammett indicators following a seriesof successive approximations. Details of the titration procedure arepresented elsewhere [28].

698 S. Nair et al. / Fuel 105 (2013) 695–704

3. Results

3.1. Influence of the hetero atom

From a comparison of breakthrough characteristics of severalsulfur species in model fuel compositions as well as JP5, it was ob-served that sulfur capacity of Ag/TiO2 was influenced by the struc-tures of the sulfur heterocycles [4]. Breakthrough studies usingsulfur free solute molecules containing other heteroatoms was car-ried out to observe selectivity of the adsorbent. The data may bethen correlated to the chemical nature of the active centers.

Breakthrough curves of benzothiophene, quinoline, naphtha-lene, and benzofuran over Ag/TiO2 adsorbent are shown in Fig. 1.These solutes were selected based on their structural similarityto benzothiophene; in quinoline, the sulfur atom replaced by nitro-gen, in benzofuran by oxygen and naphthalene had no heteroatom.The initial concentration of each component was 3500 ppmw in n-octane. The highest saturation capacity was observed for quinoline(44.60 mg/g) followed by benzothiophene (22.30 mg/g) and naph-thalene (20.44 mg/g). The lowest capacity was exhibited towardbenzofuran (14.62 mg/g). Thus the affinity of the active center to-ward the heteroatom was in the order N > S > O which corre-sponded to the order of electronegativity values of these atoms.Therefore electron rich or basic heteroatoms imparted higher affin-ity for the active center which indicated that the active center wasacidic or electron deficient in nature.

3.2. The influence of acidic centers on sulfur adsorption

To ascertain the relationship between surface acidity and sulfurcapacity, Ag/TiO2 adsorbent was treated with basic probe mole-cules NH3, pyridine, 2,6-lutidine, CO, and TMCS. The adsorbent par-ticles were saturated with probe molecules as explained earlierand sulfur capacity established using JP5 fuel as challenge.

Two sets of breakthrough studies were carried out on the adsor-bent samples; the first on blank TiO2 and the second on adsorbentsamples supported with silver. In order to distinguish the effect ofprobe molecules on the support devoid of silver phase, activatedAl2O3 and TiO2 were treated with NH3 and breakthrough behaviorobserved. Sulfur capacities of ‘dried’ adsorbent samples were alsomeasured for comparison purposes. The various stages of activa-tion and pretreatment with probe molecules are schematicallyshown in Fig. 2.

Pretreatment of the blank TiO2 samples with NH3 deactivatedthe active centers completely (Fig. 3a). Similar deactivation was

Fig. 1. The effect of electron density on the heteroatom on adsorption efficiency ofquinoline, benzothiophene, and benzofuran.

also observed on the Al2O3 support (Fig. 3b). Therefore the activesites generated on the ‘blank’ supports resulting from activationwere acidic.

On the adsorbent samples supported with silver, all the probemolecules acted as poisons. The poisons had three effects on theperformance of the adsorbent; the first group reduced the sulfurcapacity, the second group completely deactivated the adsorbentand the third group also deactivated the adsorbent, but reversibly.This behavior was explained based on the strength of interactionbetween the acidic center and the basicity of the probe molecule.The probe molecules with more basicity irreversibly deactivatedthe acid centers while the less basic had a milder interaction andwas replaced by the sulfur aromatics during desulfurization(Fig. 4).

Interaction with CO molecules reduced the sulfur capacity ofthe adsorbent. However the effect was the least pronouncedamong the probe molecules. This indicated a gradation in strengthof surface sites. CO being a much milder base compared to NH3,interacted with stronger of acid sites. NH3 and TMCS completelyand irreversibly deactivated the adsorbent. This interaction withTMCS indicated that the sulfur removal centers were predomi-nantly single hydroxyl groups.

Pyridine and 2,6-lutidine are common probe molecules used inthe study of surface acidity [36]. They demonstrated a reversibleinteraction, pyridine being the stronger among the two. 2,6-lutidineis almost selectively adsorbed on Bronsted sites while pyridinereacts with both Bronsted and Lewis sites [49]. Methyl groups onthe 2,6-lutidine molecule sterically hinder the nitrogen atom so thatthe molecule fails to interact with Lewis sites compared to pyridine.2,6-lutidine only interacts with Bronsted sites because the hydroxylgroups extend away from the substrate surface and make the inter-action possible. Both pyridine and 2,6-lutidine treated adsorbentsamples showed no breakthrough capacity for sulfur and was inac-tive during the initial phase of breakthrough. However, with timethe sulfur molecules were able to displace the active centers popu-lated by both pyridine and 2,6-lutidine, more easily in the case of2,6-lutidine. From the above breakthrough behavior, not only wasthe acidic nature of the active center reiterated but also that theywere Bronsted sites consisting mainly of single hydroxyl groups.

3.3. Infrared spectroscopy

Identification of the surface functional groups resulting frominteractions with probe molecules was performed via IR for provid-ing further information on the nature of active centers. Infrared

Ag (4 wt%)/TiO2 established from breakthrough of 3500 ± 10 ppmw naphthalene,

Fig. 2. The sequence of activation, deposition of silver phase (silver oxide), andpoisoning of active sites.

S. Nair et al. / Fuel 105 (2013) 695–704 699

spectra of calcined Ag/TiO2 were compared following treatmentwith NH3 and TMCS. NH3 was exposed to the adsorbent sampleat atmospheric pressure for 30 min followed by evacuation for an-other 30 min at 100 mtorr. Exposure to TMCS vapor was performedat vacuum (100 mtorr) for 30 min. Ample amount of adsorbate wasused in each test for complete saturation. Both the treatmentswere performed at room temperature.

NH3: Spectra of calcined Ag/TiO2 before and after treatment withNH3 were compared as shown in Fig. 5. Table 2 illustrates the peaksand their corresponding assignments. For neat Ag/TiO2, The OHstretching vibrations can be seen in the range of 3600–3750 cm�1

Fig. 3. The effect of poisoning of acidic adsorption centers using NH3 on activated (a) TiOJP5 fuel (1172 ppmw sulfur) as challenge.

Fig. 4. The effect of poisoned adsorption centers on Ag (4%)/TiO2 with probe molecules o

(Fig. 5) [50,51]. These bands disappeared when treated with ammo-nia, indicating adsorption of ammonia on these sites. Both theadsorbent samples showed an absorption band at 1620 cm�1,although a small shift to lower wavenumber (ca. 5 cm�1) was ob-served for the ammonia adsorbed sample. This band has been asso-ciated with the OH bending (t2) vibration of residual adsorbedwater. Studies have shown that monomeric water has its t2 vibra-tion at 1600, dimeric water at 1620 and polymeric water at1633 cm�1 [52]. Therefore there was an abundance of dimeric wateron the surface of Ag/TiO2. The band at 1620 cm�1 has also been ob-served in the case of dried (at 200 �C) anatase-TiO2 by otherresearchers [55]. A broad band centered at 1170 cm�1 was observedthat denotes t2 vibrations of coordinatively bound ammonia on Le-wis acid sites [54,56]. A weak shoulder was also observed at1225 cm�1 showing the presence of dissociated ammonia.

Apart from surface hydroxyl groups, the presence of ammoniumions on the surface upon NH3 absorption also indicates the pres-ence of Bronsted centers. A broad band observed at 1430–1470 cm�1 has been reported to be the t4(das) asymmetric defor-mation vibration of the NHþ4 species on surface Bronsted sites[53]. This indicates the protonation of part of the adsorbed NH3

and the presence of Bronsted acid sites on the sample surface[54]. Thus IR spectroscopy on NH3 adsorbed adsorbent sampleswas able to identify both Bronsted and Lewis acid sites on theadsorbent surface.

2 and (b) Al2O3 on sulfur adsorption capacity indicated by breakthrough tests using

f varying acidity on desulfurization using JP5 fuel (1172 ppmw sulfur) as challenge.

Fig. 5. In situ IR Spectra (in transmission mode) of calcined Ag (4 wt%)TiO2 (a)before treatment and (b) after treatment with NH3.

Fig. 6. In situ IR Spectra (in transmission mode) of calcined Ag (4 wt%)TiO2 (a)before treatment and (b) after treatment with TMCS.

700 S. Nair et al. / Fuel 105 (2013) 695–704

TMCS: Fig. 6 shows the IR signatures of Ag/TiO2 samples beforeand after treatment with TMCS. The peaks of the spectra have beendescribed in Table 2. Generally, TMCS reacts selectively with thesurface hydroxyl groups. There are three types of surface OHgroups; vicinal, germinal, and hydrogen bonded. Only vicinal andgerminal OH groups react with TMCS. Single hydroxyl groups withoxygen atoms facing away from the plane of the surface latticepreferentially react 1:1 with methyl chlorosilane [57]. Hydrophilicnetworks formed among OH groups deter the hydrogen bondedcenters from reacting with TMCS [57]. From the figure it can beseen that the peaks at 3650–3800 cm�1 were diminished in inten-sity in the TMCS treated adsorbent sample. Data presented in Fig. 6therefore indicates the adsorbent to contain a mixture of hydroxylgroups (Bronsted acid sites). While vicinal and germinal OH groupsreacted with TMCS, the remaining broad peak indicates the pres-ence of hydrogen bonded groups. The band at 1620 cm�1 has beendescribed earlier. The bands at 1255 cm�1 (t4 CH3 symmetricdeformation), 1449 cm�1 (t16 CH3 nonsymmetric deformation),and 2963 cm�1 (surface attached TMCS groups) were also observedin the spectra of TMCS treated Ag/TiO2 as shown in Fig. 6 [58].

3.4. Ammonia chemisorption

NH3 chemisorption measurements were carried out to gaugeoverall surface acidity of the adsorbent samples as NH3 interactswith both Lewis and Bronsted sites. The adsorbent samples werecalcined in situ, cooled to the analysis temperature (175 �C) priorto obtaining the isotherms. NH3 uptake of SiO2, Al2O3, and TiO2

Table 2IR peaks and their respected assignments for calcined Ag/TiO2 before and after treatment

Sample Peak wavenumber (cm�1) Assignment

Ag (4 wt%)/TiO2 1322 Vibration re1620 –OH bendin3600–3750 –OH stretch

Ag (4 wt%)/TiO2 + NH3 1170 t2 Vibration1225 Dissociated1430–1470 t4(das) Asym1600–1630 –OH bendin

Ag (4 wt%)/TiO2 + TMCS 1255 t4 Symmetri1449 t16 Nonsymm2963 Surface attac

was obtained. The amount of chemisorbed NH3 was obtained fromisotherms representing combined and weak adsorption as ex-plained earlier (Fig. 7).

NH3 uptake was in the order TiO2 > Al2O3� SiO2. This confirmedthe earlier observation that the surface of SiO2 was non-acidic.Increasing calcination temperature to 600 �C did not significantlyimpact NH3 uptake on SiO2. However, on TiO2 and Al2O3, increasein calcination temperature from 200 to 400 �C increased NH3

uptake. The increase was higher on TiO2 than Al2O3. NH3 uptake ofTiO2 decreased at activation temperature of 600 �C. This was attrib-uted to collapse of TiO2 pore structure. Previous work on the TiO2

support also indicated a significant loss in surface area at calcinationtemperatures above 500 �C [59].

The addition of silver to SiO2 and Al2O3 resulted in an increase inNH3 uptake as shown in Fig. 7. However for TiO2 the NH3 uptake de-creased with the addition of silver. The silver phase occupies acidsites during impregnation and itself adds more acid sites. Whenthe number of acid sites added by silver phase is greater than thesites removed due to its deposition, the total acidity increases. Inthe case of TiO2 the net acidity was lower. We hypothesize that sul-fur adsorption occurs on Bronsted acid sites; both on TiO2 supportand AgxO phase. Since NH3 chemisorption measured overall surfaceacidity, titration with mild organic bases were employed to esti-mate Bronsted acidity as described in the following section.

3.5. Potentiometric titration

Introduction of probe molecules to the surface prior to adsorp-tion studies indicated that Bronsted centers play a substantial role

with NH3 and TMCS.

Refs.

sulting from SO4 impurities [51,55]g (t2) vibration of dimeric water [52]ing bands [50,51]

s of coordinatively bound ammonia [54,56]ammonia [56]metric deformation vibration of NHþ4 [53,54]g (t2) vibration of Monomeric, dimeric, and polymeric waters [52]

c deformation of –CH3 [57]etric deformation of –CH3 [57]

hed TMCS groups [58]

Fig. 7. Ammonia adsorption isotherms of SiO2, Al2O3, and TiO2 samples calcined in air at 200, 400 and 600 �C compared to Ag (4 wt%)/SiO2, Ag (4 wt%)/Al2O3, and Ag (4 wt%)/TiO2 samples calcined at 400 �C; Isotherms obtained at 175 �C.

S. Nair et al. / Fuel 105 (2013) 695–704 701

in sulfur adsorption. For quantitative measurements, Bronstedsites were selectively titrated using 2,6-lutidine solution. The effectof activation on Bronsted acidity of Al2O3, TiO2, and SiO2 were stud-ied. Pre-dried adsorbent samples were calcined at various temper-atures and titrated with the base. Equilibrium sulfur capacities ofthe adsorbent samples were obtained using JP5 fuel. Calcinationtemperature of TiO2 was limited to 550 �C as the pore structurecollapsed at higher temperatures [59].

The pKa of the methyl yellow indicator is +3.3. The pKa of sam-ples that required no titer (indicated as zero lmoles/g) was greaterthan +3.3. It was observed that with the increase in activation tem-perature, the concentration of Bronsted sites increased (Fig. 8).Activation temperature of Al2O3 was higher (250 �C) compared toTiO2 (100 �C). The concentration of Bronsted sites on Al2O3 surfacewas higher compared to TiO2. Previous research has indicated that

Al2O3 mostly comprises of Bronsted sites while Lewis sites domi-nates TiO2 [60]. SiO2 was basic after calcination at 450 �C andhad very little impact on its sulfur adsorption capacity. An identicaltrend was observed from NH3 chemisorption experiments pre-sented earlier. The order of support samples in terms of the con-centration of proton donating surface acid sites was:Al2O3� TiO2 > SiO2. Sulfur adsorption capacities of the adsorbentsamples using JP5 fuel were measured to correlate with the mea-sured Bronsted acidity shown in Fig. 8. The relationship betweencalcination temperature and proton donating surface acidity aswell as sulfur removal capacity was evident. Although both Al2O3

and TiO2 have almost the same sulfur adsorption capacity, thebronsted acidity of Al2O3 was much higher than TiO2. The probableexplanation might be the difference in the types of bronsted acidsites (hydroxyl groups). It was predicted earlier that sulfur is

Fig. 8. The relationship between bronsted surface acidity measured by potentiometric titration and sulfur adsorption capacity using JP5 fuel (1172 ppmw sulfur) as challengefor activated (a) Al2O3 and (b) TiO2 at various activation temperatures.

702 S. Nair et al. / Fuel 105 (2013) 695–704

captured only on the single surface (isolated) hydroxyl groups,while other hydroxyl groups (e.g. vicinal and hydrogen bonded) re-mains inert to them. Although Al2O3 has a lot of hydroxyl groups,most of them might be the vicinal and hydrogen bonded hydroxylgroups.

The effect of addition of the silver phase to TiO2 on the Bronstedacidity was also investigated. The amounts of base required to neu-tralize samples of Ag (4 wt%)/SiO2, Ag (4 wt%)/Al2O3, and Ag(4 wt%)/TiO2 were determined using titration listed in Table 3.There was a considerable increase in the amount of base requiredfor neutralization which demonstrates that metal impregnationenhanced Bronsted surface acidity.

4. Discussion

Activated Al2O3 and TiO2 demonstrated a sulfur capacity of ca.0.15 mmol/g using a challenge of 3500 ppmw benzothiophene inn-octane. Addition of silver phase increased the sulfur capacityfrom 0.15 to 0.3 mmol/g as shown in Fig. 9a. Breakthrough studyof different heteroatoms pointed out the basicity of the organicmolecules as a function of adsorption on silver. NH3 chemisorptionindicated that activated Al2O3 and TiO2 were considerably moreacidic than SiO2. A comparison of NH3 uptake (760 mm) betweenactivated supports and silver supported adsorbent samples isshown in Fig. 9b. The acidity was acquired during activation pro-cess and increased with increasing activation temperature. How-ever this was not observed when high temperatures resulted inthe collapse of pore structure as in TiO2 [59]. Addition of silverphase resulted in a modest increase in the acidity of Al2O3 andSiO2 while witnessed a decrease in the case of TiO2. Similar trendswere observed on Bronsted acidity measurements by titration.Bronsted acidity increased with activation temperature. Additionof silver increased the Bronsted acidity of TiO2 and SiO2 but thatof Al2O3 did not vary. Generation of Bronsted acidity during calci-nation is a result of defect mediated dissociation of water mole-cules on oxygen vacancies or Lewis centers. Water moleculesexist on support surfaces in varying configurations such as hydro-nium ions and dimers. High temperature treatments can drive off

Table 3Bronsted acidity of activated (450 �C/2 h in air) and Ag loaded supports obtained fromtitration measurements.

Sorbent 2,6-Lutidine titer/sorbent wt. (lmoles/g)

Activated Ag loaded

TiO2 7.5 100Al2O3 700 700SiO2 0.0 225

surface water and OH groups resulting in Lewis sites. Increasingsurface temperature may also convert vicinal OH groups to eithergerminal or isolated OH groups. Such changes to surface watershown schematically in Fig. 10 significantly increases surface acid-ity and are reflected in NH3 chemisorption and titration measure-ments. In situ IR spectroscopy to monitor these interactions wascarried out that indicated the presence of both Bronsted and Lewisacid centers. Bronsted centers consisting of hydroxyl groups wereidentified for both NH3 and TMCS treated adsorbent samples. Pre-viously TPR measurements had indicated the removal of stronglybound acidic groups on activated TiO2 by the addition of silverphase [61]. Similar observation was also made from IR spectra ofNH3 treated adsorbent samples. Therefore silver occupies acidicsites on the surface. It is also clear from potentiometric titrationdata that silver phase consists of additional Bronsted centers.Therefore there would be a net increase in surface acidity whenBronsted acidity added by the silver phase is higher than the acidsites occupied by the dispersed silver. A comparison of Bronstedacidity between activated and silver supported adsorbent samplesis listed in Table 3. Bronsted acidity of the TiO2 sample increasedfrom 7.5 to 100 lmol/g and that of SiO2 from 0 to 225 lmol/g forSiO2. The Bronsted acidity of Al2O3 remained unchanged at700 lmol/g.

To investigate whether the surface acid sites are responsible forsulfur adsorption, the adsorbent samples were purposefully poi-soned with different probe molecules which were specific to cer-tain surface acid sites. The breakthrough tests reveal that thesulfur adsorption capacity was decreased substantially when theBronsted acid sites were poisoned. Comparison of data from poten-tiometric titration and equilibrium saturation experiments furtherillustrated that Bronsted acid sites were responsible for sulfuradsorption.

Sulfur adsorption mechanism on silver adsorbents at room tem-perature are postulated through an interaction between the p elec-trons on sulfur molecules and the acidic OH groups on the surfaceas represented in Fig. 11. The hydroxyl groups present in titaniaparticipates in the adsorption of sulfur compounds. Calcinationaugments the hydroxyl group concentration thereby increasingsulfur adsorption capacity. Silver impregnation on titania reducesoverall surface acidity but increases Bronsted acidity, eventuallyenhancing sulfur adsorption. There are also possibilities of sulfuradsorption on this complex adsorbent formulation through otherminor pathways, which has been described previously [4]. How-ever, for this type of oxide adsorbents, there is a clear relationshipbetween sulfur adsorption capacity and surface acidity. Futurework will focus on more extensive researches regarding this rela-tionship and the application of more acidic materials in sulfuradsorption.

Fig. 9. A comparison of (a) sulfur adsorption capacity and (b) NH3 uptake between activated TiO2, Al2O3 and SiO2 supports and samples supported with 4 wt% Ag.

Fig. 10. Conversion of vicinal OH, hydronium ions, and water in dimeric form togeminal and isolated OH’s.

Fig. 11. Adsorption of benzothiophene through interaction between p-electronsand proton donating acid sites.

S. Nair et al. / Fuel 105 (2013) 695–704 703

5. Conclusions

Active sulfur adsorption centers on supported silver adsorbentswere acidic in nature. This was confirmed using several techniques.The adsorption selectivity among aromatic species was observed tobe quinoline > naphthalene � benzothiophene > benzofuran, whichis similar to the order of basicity of the molecules. The poisoningeffect of probe molecules on adsorption of benzothiophene wasNH3 > TMCS > Pyridine > 2,6 Lutidine > CO also in the order of theirbasicity. IR signatures of surface complexes with Lewis andBronsted acid sites were observed on adsorbent samples treatedwith NH3 and TMCS. Surface acidity was measured by NH3 chemi-sorption and Bronsted acidity using titration with 2,6-lutidine.Activation in air increased the Bronsted surface acidity from 0 to10 lmol/g between calcination temperatures 20 and 550 �C andthe addition of silver to the TiO2 surface increased acidity from 5to 100 lmol/g. Therefore Bronsted acid centers account for the sul-fur adsorption capacity of these materials.

Acknowledgements

Supports from the Office of Naval Research and US Army aregratefully acknowledged; also the contributions of Donald Cahela,Alexander Samokhvalov, and Wenhua Zhu of the Center for Micro-fibrous Materials at Auburn University and Hongyun Yang of Intra-micron Inc. for their inputs on various aspects of this work.

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