Applied Catalysis A: General 511 (2016) 87–94
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Core–shell structured, nano-Pd-embedded SiO2 –Al2 O3 catalyst (Pd@SiO2 –Al2 O3 ) for direct hydrogen peroxide synthesis from hydrogen and oxygen Myung-gi Seo a , Seongmin Kim a , Dae-Won Lee b,∗ , Ha Eun Jeong a , Kwan-Young Lee a,c,d,∗∗ a
Department of Chemical and Biological Engineering, Korea University, 145 Anam-ro, Seoul 02841, Republic of Korea Department of Chemical Engineering, Kangwon National University, 1 Kangwondaehak-gil, Chuncheon-si, Gangwon-do 24341, Republic of Korea c Green School, Korea University, 145 Anam-ro, Seoul 02841, Republic of Korea d KU-KIST Graduate School of Converging Science and Technology, Korea University, 145 Anam-ro, Seoul 02841, Republic of Korea b
a r t i c l e
i n f o
Article history: Received 25 August 2015 Received in revised form 18 November 2015 Accepted 21 November 2015 Available online 2 December 2015 Keywords: Palladium catalyst Core–shell structured catalyst Palladium nanoparticle Silica–alumina Direct hydrogen peroxide synthesis
a b s t r a c t In our previous studies, we proved that core/shell-structured Pd/SiO2 catalysts are more active for direct hydrogen peroxide synthesis than the conventional, impregnated Pd/SiO2 catalysts. In this study, the topic of our previous studies was extended to core/shell Pd/SiO2 –Al2 O3 catalysts, through which we examined the influence of acidic shell oxides (SiO2 –Al2 O3 ) on the hydrogen peroxide formation activity. The catalysts were prepared based on the Stöber method, and the reaction tests were performed by adding H3 PO4 (0–0.03 M) and in the presence of KBr (0.9 mM). It was proved that the surface Brønsted acid sites promote hydrogen peroxide formation activity in a manner similar to protons dissolved in a reaction medium (ethanol–water). It was supposed that the influences of heterogeneous and homogeneous acids on catalytic activity are related to how much those acids promote the adsorption of Br− ions on the Pd surface. The highest H2 O2 production rate was approximately 470 mmol H2 O2 /gPd h, which was obtained using core/shell Pd/SiO2 –Al2 O3 catalysts under specific H3 PO4 concentrations. This rate was higher than the highest value (∼ 420 mmol H2 O2 /gPd h) achieved using core/shell Pd/SiO2 catalysts. © 2015 Elsevier B.V. All rights reserved.
1. Introduction Hydrogen peroxide is an environmentally clean oxidant/cleaning reagent, which is widely used in many fields of industry, such as pulp (paper) bleaching, waste water treatment (advanced oxidation processes, AOPs), semiconductor cleaning, and soil remediation [1–4]. In industrial chemistry, the major application of hydrogen peroxide is to synthesize sodium percarbonate and sodium perborate, both of which serve as bleaching reagents in laundry detergent. Hydrogen peroxide is also widely used to synthesize organic peroxides, which serve as an initiator or a promoter in polymerization and are frequently used in
∗ Corresponding author at: Department of Chemical Engineering, Kangwon National University, 1 Kangwondaehak-gil, Chuncheon-si, Gangwon-do 24341, Republic of Korea. Tel.: +82 33 250 6331; fax: +82 33 251 3658. ∗∗ Corresponding author at: Department of Chemical and Biological Engineering, Korea University, 145, Anam-ro, Seoul 02841, Republic of Korea. Tel.: +82 2 3290 3299; fax: +82 2 926 6102. E-mail addresses:
[email protected] (D.-W. Lee),
[email protected] (K.-Y. Lee). http://dx.doi.org/10.1016/j.apcata.2015.11.040 0926-860X/© 2015 Elsevier B.V. All rights reserved.
advanced oxidation processes as an oxidant to dissociate the organic matters dissolved in waste water. Hydrogen peroxide is almost exclusively manufactured by ‘anthraquinone process, ’ which produces hydrogen peroxide via the redox cycle of ethyl (or amyl) anthraquinone. However, hydrogen peroxide suffers from being the source of environmental hazards and having a relatively high cost for product recovery; both of these issues are related to the heavy use of organic solvents [5]. Accordingly, the direct synthesis of hydrogen peroxide from hydrogen and oxygen (Reaction (1)), which is free of organic solvents and most desirable in terms of atomic efficiency, has been proposed as an alternative route to produce hydrogen peroxide [6,7]. H2 + O2 → H2 O2
(1)
1 H2 + O2 → H2 O 2
(2)
1 O2 2
(3)
H2 O2 + H2 → 2H2 O
(4)
H2 O 2 → H2 O +
However, the process has not reached a level of commercialization because of the low yields of hydrogen peroxide, primarily
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due to the progress of the side reactions producing water (Reactions (2)–(4)). The majority of catalysts developed for direct H2 O2 synthesis are Pd or Pd-alloy supported catalysts. A high level of Pd dispersion (nano-size distribution) is necessarily required for obtaining high H2 and O2 conversions [8], but it is widely known that the morphology (or exposed surface facet of Pd crystal) is a primary factor governing the selectivity of H2 O2 . For example, we proved experimentally that the Pd {1 1 1} facet, which surrounds the exterior of Pd octahedron, is more selective toward H2 O2 than the Pd {1 0 0} facet of a Pd cube [9]. This result is consistent with the DFT simulation results of Tian et al., who claimed that H2 O2 productivity is degraded as the OO dissociation of O2 , OOH (intermediate), and HOOH (H2 O2 , product) are promoted and that Pd {1 1 1} is energetically disadvantageous over Pd {1 0 0} in those OO dissociation reactions [10]. It is also well known that the OO dissociation is better promoted on the edge/corner sites than on the terrace ones of Pd crystal [11,12]. Accordingly, the selective deactivation of edge/corner sites leads to the improvement of H2 O2 selectivity [13,14]. The reaction is usually performed in a water solution containing halide anions (Br− , Cl− ) and protons (H+ ), both of which are known to accelerate H2 O2 formation while suppressing the water-producing side reactions. The major role of halide anions is to enhance the H2 O2 yield (selectivity) via adsorption on the edge/corner Pd sites and deactivating the O–O dissociation reactions. In addition, it is also known that a halide anion provides an additional H2 O2 production route by donating its electron to the non-dissociatively chemisorbed oxygen to form the O2 − intermediate, which then combines with H+ to produce H2 O2 [15]. However, the excess addition of halide anions is rather detrimental to H2 O2 formation because it could decrease the number of beneficial active Pd sites and results in the decrease of H2 conversion [6,16–19]. However, the exact role of H+ is still under debate [6,16–19], with the most widely accepted explanation being that H+ suppresses the decomposition of H2 O2 by accelerating the elimination of H2 O2 adsorbed on the active Pd sites (i.e., by accelerating the decomposition of H–Pd–OOH intermediate) [14]. Another plausible explanation is that H+ lowers the pH of a reaction medium (ethanol–water) below the isoelectric point of the support material, thereby promoting the adsorption of halide anions [15]. It is known that an adsorbed halide anion migrates to the Pd surface with a proton to form Pd·H+ ·X− [14]. However, similar to halide anions, an excessive addition of acid is undesirable for H2 O2 formation because it causes excessive adsorption of halide anions (Br− ), and even the terrace sites of Pd crystal become poisoned, which consequently reduces the H2 conversion [14]. In more serious cases, the leaching of Pd metal may be caused under the highly acidic water condition [20]. Some studies have been reported to use solid acid supports [20–22], such as zeolites, to prevent or reduce these side effects. According to Park et al., the Brønsted acid sites on the catalyst have the effects of stabilizing H2 O2 molecules, similar to the effects of dissolved protons in water [20]. However, to date, the effects of solid acidity on hydrogen peroxide synthesis are not fully known. In our previous studies, we studied the direct H2 O2 synthesis using core/shell Pd/SiO2 catalyst (Pd@SiO2 ) involving nano-sized Pd crystals enclosed by a porous SiO2 shell [23,24]. The catalysts were prepared by forming a silica hydrogel (Stöber method) over a nano-Pd slurry. The pore structure, which is required for gaseous molecules to be transferred, is formed by incinerating organic dispersant molecules (polyvinylpyrrolidone) at a high temperature. The core/shell catalyst structure is advantageous over conventional metal-on-support structure in preventing the nanoPd crystals from being thermally agglomerated during calcination and physically/chemically leached during reaction test. Additionally, the structure provides a solution for catalyst recovery, which
is very inconvenient when applying nano-crystals directly as a catalyst. Crucially, it was revealed in our previous study that the core/shell Pd@SiO2 was more active in direct H2 O2 synthesis than conventional Pd-supported SiO2 catalyst (Pd/SiO2 ): Pd@SiO2 was approximately two-times better than Pd/SiO2 in terms of gram Pdspecific H2 O2 productivity [24]. It was revealed that the prepared core/shell structure preserved the activity of nano-Pd crystals, and its porous shell provides the paths for gas molecules to reach the Pd crystal core. We have continued the study of core/shell structured catalysts for direct H2 O2 synthesis. In this paper, we will present the results from the use of Pd@SiO2 –Al2 O3 catalysts where Pd nano-crystals were encapsulated by acidic shell oxides (silica–alumina). Pd@SiO2 –Al2 O3 of different Si/Al ratios (20, 50 and 100, molar ratio) were prepared and compared to Pd@SiO2 in H2 O2 production activity, by which the influence of solid acidity on H2 O2 synthesis was investigated. To the best of our knowledge, our work is the first study investigating the catalytic activity of ‘metal-in-acidic oxide,’ core/shell type materials. In preparation of the catalysts, we applied the Stöber method for enclosing the nanoPd core with SiO2 –Al2 O3 shell. To induce homogenous mixing of Si and Al over the shell, ‘double alkoxide (di-sec-butoxyaluminoxytriethoxysilane)’ was adopted as a precursor of the SiO2 –Al2 O3 shell and then mixed with tetraethylorthosilicate (TEOS) to adjust the intended Si/Al ratio [25]. The uses of Pd-acidic oxide composite catalysts (Pd/HZSM-5 [20] and palladium-exchanged heteropolyacid [22]) for direct H2 O2 synthesis have been previously reported, but the contribution of surface acid sites to the H2 O2 formation activity was not fully confirmed in those studies. Therefore, in this study, the activity tests of Pd@SiO2 –Al2 O3 were focused on investigating the influences of Si/Al ratios of the catalysts on H2 O2 synthesis activity under different H+ concentrations in the presence of halide anion (Br− ). 2. Experimental 2.1. Chemicals l-Ascorbic acid (ACS reagent, ≥99%), polyvinylpyrrolidone (PVP, K-value 28–34, 55,000 g/mol), potassium bromide (KBr, ACS reagent, ≥99%), sodium tetrachloropalladate (98%, 35.3–37.1% Pd basis), ethanol (ACS reagent, ≥99.5%, absolute), ammonium hydroxide solution (28–33% NH3 basis), and tetraethylorthosilicate (TEOS, ACS reagent, 98%) were purchased from the Sigma–Aldrich Company. Di-sec-butoxyaluminoxy-triethoxysilane ((s-OBu)2 AlOSi(OEt)3 , 95%) was purchased from Gelest Corporation. All of the chemicals were used without further purification. Deionized water (18.2 M) was used in all the procedures involving aqueous solutions and the washing of solid products. 2.2. Synthesis of the catalysts First, Pd nano-crystals were prepared by reducing Na2 PdCl4 with l-ascorbic acid in aqueous solution, to which PVP and KBr were added as stabilizer and capping agents, respectively. The detailed procedure was described in our previous paper [12]. The slurry was washed 5 times with acetone under centrifuge (13,000 rpm, 10 min per each repeat), from which PVP-stabilized Pd slurry was obtained. The prepared Pd nano-crystals were enclosed by a SiO2 or SiO2 –Al2 O3 shell oxide via a modified Stöber method: the PVPstabilized Pd slurry was re-dispersed in 200 mL of ethanol, to which 18 mL of DI water, 10 mL of NH4 OH, and stoichiometric amounts of TEOS and (s-OBu)2 AlOSi(OEt)3 were added. (In the synthesis of Pd@SiO2 , (s-OBu)2 AlOSi(OEt)3 was not added.) The stirring was continued for 24 h at room temperature. The prepared Pd@SiO2 –Al2 O3 particles were collected by centrifugation
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Fig. 1. TEM images of (a) Pd@SA(20), (b) Pd@SA(50), (c) Pd@SA(100), and (d) Pd@S.
at 15,000 rpm for 20 min, and the collected particles were washed with deionized water several times. The washed particles were dried overnight at 333 K and then calcined at 773 K under air for 10 h, followed by reduction under 10 vol.% H2 /N2 gas at 623 K for 2 h. The prepared Pd@SiO2 and Pd@SiO2 –Al2 O3 catalysts were denoted as Pd@S and Pd@SA(X), respectively, in which X represents the Si/Al molar ratio (20, 50, or 100). 2.3. Characterization of the catalysts Transmission electron microscopy (TEM) images were obtained on a Tecnai G2 F30 instrument (FEI Company, OR, USA) operated at 300 kV. The sample was dispersed in ethanol and then dropped over a 300 mesh Cu grid, which was completely dried in air before performing analysis. The analysis was performed at the Seoul Center of the Korean Basic Science Institute (KBSI). A BELSORP-max instrument (BEL Japan Inc., Osaka, Japan) was used to obtain N2 adsorption–desorption isotherms. The specific surface area and total pore volume of each catalyst were calculated using the BET equation and the inner area, and the volume of mesopores for each catalyst was estimated from the t-plot analysis
of the isotherm. The differences between the BET and t-plot data were assigned to the textural properties related to micropores. The crystalline structure of each catalyst was examined using Xray diffraction (XRD: D/MAX-2500/PC, Rigaku). The measurement was performed at the scanning speed of 1◦ /min on the ATX-G X-ray diffractometer (Rigaku) using Cu–K˛ irradiation (˛ = 1.5406 Å). The wt.% of Pd and the Si/Al molar ratio of each catalyst were measured via inductively coupled plasma optical emission spectrometry (ICP-OES) using an iCAP-6300 duo ICP-OES spectrometer (Thermo Scientific, Waltham, MA, USA). Before each measurement, each sample was dissolved in a digestive pump with a liquid mixture (9 mL) of nitric acid, hydrochloric acid, and hydrofluoric acid in 1:3:1 volumetric ratio. The exposed surface areas of Pd on Pd@S and Pd@SA(X) catalysts were estimated using the CO chemisorption method. Measurements were performed at 308 K using an ASAP 2020 chemisorption analyzer (Micrometrics Inc., Norcross, GA, USA). The calculation was based on an assumption that a molecule of CO adsorbs onto an exposed Pd atom. Thus, the number of surface Pd atom was equal to the amounts of CO adsorption. Pd surface area and Pd dispersion were calculated by using the following equations.
Table 1 Si/Al molar ratio, Pd contents, and the exposed Pd surface area of Pd@SA(X) and Pd@S catalysts. Catalyst
Si/Al molar ratioa
Pd loading (wt.%)a
Exposed Pd area (m2 /g-cat. )b
Exposed Pd atoms (×1019 )b
Pd dispersion (%)b
Pd@S Pd@SA(100) Pd@SA(50) Pd@SA(20)
Si Only 99.7 53.9 19.9
2.89 2.94 3.06 2.98
1.42 1.62 1.68 1.80
1.80 2.09 2.13 2.29
11.0 12.4 12.3 13.7
a b
Si/Al molar ratios and Pd loadings were calculated from the ICP-OES data. Exposed Pd areas, surface Pd atoms and Pd dispersion were calculated based on the CO chemisorption results.
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Pd surface area
=
2.5. Br− ion adsorption experiment
m2 Pd
g-cat
atom number of absorbed CO molecules × 1 COPdmolecule
7.87×10−20 m2 ×
Pd atom
mass of catalyst (g)
Pd dispersion (%) =
the number of surface Pd atoms × 100 total number of Pd atoms in the catalyst
(5)
(6)
NH3 temperature-programmed desorption (NH3 -TPD) analyses were performed using a BELCAT-B device (BEL Japan Inc., Osaka, Japan) to investigate the distribution of acid strength over the catalyst surface. A catalyst (0.06 g) was loaded on sample holder and pretreated at 673 K for 2 h with the supply of a 10 vol.% H2 /N2 gas (50 mL/min). NH3 adsorption was performed with 5% NH3 /He gas (50 mL/min) at 423 K for 1 h. Next, the weakly adsorbed, excess NH3 was removed from the sample with flowing helium (50 mL/min) at 423 K for 1 h. NH3 desorption was recorded while heating the sample from 423 to 773 K at a rate of 5 K/min under a pure helium flow (30 mL/min). Thermal conductivity detector (TCD) was used to detect and quantify the NH3 desorption. Fourier transform infrared (FT-IR) spectroscopy (PerkinElmer, Spectrum GX) was employed to identify the acid sites of a catalyst. Pyridine was used as a probe molecule. Each sample was pressed into a self-supported wafer, which was then evacuated on a sample holder at 573 K for 2 h. The sample was cooled to 423 K, followed by exposure to pyridine vapor for 30 min. The sample was evacuated to remove physisorbed pyridine. The spectra were recorded from 1400 to 1700 cm−1 at a 4 cm−1 resolution using 100 scans per spectrum to increase the S/N ratio. On the Brønsted acid site, pyridine is adsorbed as a pyridinium ion (C5 H5 NH+ ), whereas pyridine is coordinatively bound to Lewis acid sites via sharing of the lonepair electrons of nitrogen. Both of these modes are discernible in the FT-IR spectrum. The Brønsted to Lewis acid site ratio (B/L ratio) was calculated by dividing the integrated absorbance of the IR band at 1545 cm−1 (Brønsted) with that of the band at 1455 cm−1 (Lewis) [26]. 2.4. Reaction test: direct H2 O2 synthesis The reaction tests were performed in a magnetically stirred (1200 rpm) double-jacket glass reactor. A 150 mL of ethanol–water mixture (1:4 in volume ratio) was used as the reaction medium, in which 0.9 mM KBr and phosphoric acid (0, 0.01, 0.02, or 0.03 M) was dissolved. The catalyst loading was 0.05 g and H2 and O2 (1:10 in volume ratio) was fed to the reactor at the total flow rate of 22 mL/min. The reaction tests were performed at 293 K and atmospheric pressure for 3 h. The hydrogen peroxide concentration was estimated using the iodometric titration method [24]. H2 conversion was measured using a gas chromatograph (Younglin, ACME6000) equipped with a TCD and a Carbosieve SII (60-80 mesh)-packed column. H2 conversion, H2 O2 selectivity and gram Pd-specific H2 O2 production rate were calculated using Eqs. (7)–(9), respectively. H2 conversion (%) =
moles of H2 reacted × 100 moles of H2 supplied
H2 O2 selectivity (%) =
moles of H2 O2 formed × 100 moles of H2 reacted
(7) (8)
H2 O2 production rate (mmol/g-Pd h) =
mmoles of H2 O2 formed weight of Pd (g) in catalyst × reaction time (h)
(9)
Each Br− ion adsorption test was performed at 293 K and atmospheric pressure for 3 h. The catalyst loading was 0.2 g, and the volume of ethanol–water mixture was 20 mL. Other conditions were identical to those of the reaction test. The Br− ion concentration was measured using an ion chromatograph (Dionex, thermal conductivity detector) equipped with an IonPac AS14 column. The mobile phase was a mixed aqueous solution of Na2 CO3 (3.5 mM) and NaHCO3 (1.0 mM). The percentage of Br− adsorbed on the catalyst was estimated using the Eq. (10). Br− adsorption (%)
=
1−
Br− (mol/L) remained in reactor after adsorption Br− (mol/L) in reactor before adsorption
× 100
(10)
3. Results and discussion 3.1. Characterization of the Pd@S and Pd@SA(X) catalysts The morphologies of the prepared nano-Pd cube crystals were similar to those achieved in our previous study [12]: 7–13 nm in size (9.6 nm in average) and almost uniform in aspect ratio (1.1). The crystals were encapsulated with a SiO2 or SiO2 –Al2 O3 shell, and the TEM images of the core@shell composites are presented in Fig. 1. The images show that each composite contained one or two Pd core crystals and that the shell oxides were almost spherical and uniform in size. This observation indicates that the encapsulation occurred homogeneously over the entire slurry. The Pd core crystals were almost unchanged in size after being encapsulated, which might be due to the shell oxides preventing thermal migration of Pd crystals during calcination. The Si/Al molar ratio, Pd loading, and the exposed Pd area of the prepared catalysts are listed in Table 1. The Si/Al ratios of the Pd@SA(X) catalysts were close to the intended values, and the Pd loading was almost constant (3 wt.%) over the catalysts. The exposed surface areas of metallic Pd were 1.4–1.8 m2 /gPd , indicating the incineration of the PVP-generated pore paths connecting the entrapped Pd core to the outer surface of shell oxides. N2 adsorption–desorption analysis results are presented in Fig. 2(A). All the isotherms could be classified as type-IV isotherms with a hysteresis loop of type H1, which are generally interpreted that the materials have mesoporosity. However, the BJH analyses of the isotherms indicated that the mesopores ranged 20–50 nm (Fig. 2(B)), which were too large to be considered as the pores developed inside the catalyst particle. (The size of catalyst particles ranged 50–70 nm (Fig. 1).) Moreover, if such large-size pores really existed, they should have been identified on the high-resolution TEM images (Fig. 1). Thus, it is reasonable to assign the mesopores to the intra-particle voids. Excluding the hysteresis loops from the N2 isotherms, the isotherms were close to type-I isotherm, which indicated that the samples had microporosity. The t-plot analysis revealed that 75–80% of the total specific surface area and the pore volume were attributed to the micropores (Table 2). The micropores would be generated during incineration of the PVP, which occurred in the course of calcination of the catalysts.The XRD patterns of catalysts are presented in Fig. 3. All the patterns had the characteristic peaks of metallic palladium (JCPDF 46-1043). The peaks at 2 = 40.1◦ , 46.6◦ , and 68.1◦ correspond to the {1 1 1}, {1 0 0}, and {2 2 0} planes of fcc Pd lattice, respectively. Palladium oxides or other Pd-related compounds were not detected. The broad peaks at 2 = 15 30◦ were attributed to the shell oxides (SiO2 and SiO2 –Al2 O3 ). Fig. 4 presents the NH3 -TPD curves of the catalysts. The total acid site densities (mmol-NH3 /g-cat .) were calculated by integrating the
3000
Intensity (Counts)
3
Volume adsorbed (cm /g)
Pd@S
Pd@SA(100)
Pd@SA(50)
Pd(220)
3500
Pd(200)
a)
91 Pd(111)
M.-g. Seo et al. / Applied Catalysis A: General 511 (2016) 87–94
2500
Pd@S 2000
Pd@SA(100)
1500 1000
Pd@SA(50)
500
Pd@SA(20)
Pd@SA(20)
0
10 0.0
0.2
0.4
0.6
0.8
20
30
1.0
40
50
60
70
80
2 theta (degree)
Relative pressure (p/p0) Fig. 3. XRD results of the Pd@SA(X) and Pd@S catalysts.
b) 5
Pd@S
Pd@S Pd@SA(100)
3
Intensity (A.U.)
dV/dlogD
4
Pd@SA(100)
2
Pd@SA(50)
Pd@SA(50) 1
Pd@SA(20)
Pd@SA(20) 0 0
20
40
60
80
450
500
Pore diameter (nm)
Table 2 Specific surface area and pore volume of Pd@SA(X) and Pd@S catalysts.
Pd@S Pd@SA(100) Pd@SA(50) Pd@SA(20)
600
650
700
750
Temperature (K)
Fig. 2. (a) N2 adsorption–desorption isotherms and (b) BJH pore size distribution of the catalysts.
Catalyst
550
Surface area (m2 /g-cat. )
Pore volume (ml/g-cat. )
Totala
Microporeb [%]c
Totala
Microporeb [%]c
335 340 339 306
265 [79] 267 [79] 269 [79] 231 [75]
0.56 0.64 0.61 0.56
0.45 [80] 0.52 [81] 0.49 [80] 0.46 [82]
Fig. 4. NH3 -TPD results of the Pd@SA(X) and Pd@S catalysts.
Table 3 Acidic properties of Pd@SA(X) and Pd@S catalysts. Catalyst
Pd@S Pd@SA(100) Pd@SA(50) Pd@SA(20)
Brønsted/Lewis acid site ratioa
Acid site density (mmol-NH3 /g-cat )
N/A 0.7 1.1 1.3
Totalb
Brønstedc
Lewisc
0 0.142 0.202 0.526
0 0.058 0.106 0.297
0 0.084 0.096 0.229
a
The total specific surface area and pore volume were estimated by BET method. Microporous surface area and micropore volume were estimated by subtracting t-plot data from BET data. c The percentage of the total surface area or total pore volume that micropores account for. b
NH3 -TPD curves, and the results are listed in Table 3. As shown in Fig. 4, Pd@S had no detectable acid sites, while Pd@SA(X) catalysts showed a single and distinctively sharp NH3 peak at ∼523 K, which increases in peak intensity with the increase of Al content (i.e., decrease of Si/Al ratio, X). The peaks were sharp and narrow (FWHM = ∼10 K) compared to the typical NH3 -TPD peaks of commercial silica–alumina materials, which were generally larger than 100 K in FWHM and composed of multiple peaks [27]. This result indicates that the prepared SiO2 –Al2 O3 shell was more uniform in Si/Al distribution than the conventional silica–alumina materials; this result could be attributed to the use of double alkoxide (di-sec-butoxyaluminoxy-triethoxysilane) as a precursor [25]. The total acid density increased from 0.142 to 0.526 mmol NH3 /gcat as
a
The acid site densities were obtained by integrating NH3 -TPD curves in Fig. 4. B/L ratios were obtained from the pyridine adsorption-FTIR spectroscopy results in Fig. 5. c Brønsted and Lewis acid density were calculated using total acid density and B/L ratio. b
the Si/Al ratio (X) decreased from 100 to 20 (Table 3). In the pyridine adsorbed FT-IR results (Fig. 5), the intensity of pyridinium ion adsorbed on the Brønsted acid sites increased with the increase of Al content in Pd@SA(X) catalysts. Assuming the adsorption sites of pyridine are identical to those of ammonia, the densities of the Brønsted and Lewis acid sites were calculated by multiplying total acid site density by
B/L B/L+1
and
1 B/L+1
, respectively. Table 3
indicates that the Brønsted sites became richer (B/L ratio > 1) than the Lewis sites as the Si/Al ratio (X) decreased below 50. Pd@SA(20) was 1.8 times higher than Pd@SA(100) in terms of the Brønsted acid site density.
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0.6
a)
H2 conversion (%)
0.3
Pd@SA(50) 0.2
70
15
60 50 10 40 30 5
20
H2O2 selectivity (%)
Absorbance (%)
80
H2O2 selectivity (Pd@S)
Pd@SA(100)
0.4
H2 conversion (Pd@S)
(1455 cm )
(1545 cm )
0.5
90
20
-1
-1
10
0.1 0
Pd@SA(20)
0 0M
1650
1600
1550
1500
0.02 M
0.03 M
H3PO4 concentration (M)
0.0
1700
0.01 M
1450
-1
Wave number (cm )
Br- ion adsorption (%)
40
Pd@S without H2
Pd@S with H2
Pd@SA(100) without H2
Pd@SA(100) with H2
Pd@SA(50) without H2
Pd@SA(50) with H2
Pd@SA(20) without H2
Pd@SA(20) with H2
30
20
10
0 0.00
S, SA(20), without H2
b)
500 H2O2 production rate-Pd@S
H2O2 production rate (mmol/gPdh)
Fig. 5. FT-IR spectra of pyridine adsorbed onto Pd@SA(X) catalysts.
400
300
200
100
0
0M
0.01 M
0.02 M
0.03 M
H3PO4 concentration (M) 0.01
0.02
0.03
H3PO4 concentration (M) Fig. 6. Adsorption of Br− ion on Pd@S and Pd@SA(X) catalysts. Test conditions: 293 K, 1 atm, 0.2 g of catalyst, 20 mL of ethanol–water (4:1) mixture (containing 0.9 mM KBr and 0, 0.01, 0.02 or 0.03 M H3 PO4 ), stirring rate = 1200 rpm, total gas flow rate = 0 or 22 mL/min, H2 /O2 = 1/10 (v/v), and adsorption time = 3 h.
3.2. Adsorption of Br− ions on the Pd@S and Pd@SA(X) catalysts As explained in the introduction, the use of the proper amount of halide ion adsorption is very helpful to enhance the selectivity and the yield of hydrogen peroxide, and the amount of halide ion adsorption is influenced by proton concentration of aqueous reaction medium. We examined how Pd@S and Pd@SA(X) behave regarding the adsorption of halide anion (Br− ) with the change of proton concentration in the reaction medium. Fig. 6 shows that Br− ions adsorbed onto the Pd@S and Pd@SA(X) catalysts under acidic condition. Because Br− adsorption was negligible on the pure silica (S) and silica–alumina (SA(20)), it was supposed that a Br− ion adsorbed and settled on Pd, forming a relatively stable intermediate, Pd·H+ ·Br− [14]. For all the catalysts, the amount of Br− adsorption increased with the increase of H3 PO4 concentration. The Br− adsorptions were suppressed when H2 gas was flowed into the reaction medium, which implies that a competitive adsorption occurred between H+ Br− and H2 on the Pd sites [14]. A noticeable result is that the Br− adsorption amount was higher on acidic Pd@SA(X) than on Pd@S and increased with the increase of the surface acidity of the catalyst (i.e., decrease of the Si/Al ratio, X). The trend was maintained, even in the absence of an external proton
Fig. 7. Direct synthesis of hydrogen peroxide using Pd@S catalyst under different H3 PO4 concentrations and in the presence of Br− ion: (a) H2 conversion and H2 O2 selectivity; (b) H2 O2 productivity (Pd-weight-specific H2 O2 production rate). Test conditions: 293 K, 1 atm, 0.05 g of catalyst, 150 mL of ethanol/water (4/1, v/v) mixture (containing 0.9 mM KBr and 0, 0.01, 0.02 or 0.03 M H3 PO4 ), stirring rate = 1,200 rpm, total gas flow rate = 22 mL/min, and H2 /O2 = 1/10 (v/v).
source (H3 PO4 concentration = 0 M). These observations suggest that the Brønsted acid sites on Pd@SA(X) affect the adsorption of halide anions in a manner similar to protons in the reaction medium. 3.3. Reaction results: influence of surface acidity on the H2 O2 formation activity Fig. 7 shows the activity test results of Pd@S under different H+ concentrations and in the presence of Br− ion. The figure shows that the H2 conversion, H2 O2 selectivity, and H2 O2 production rate increased with the increase of H3 PO4 concentration from 0 to 0.02 M, but turned to decrease from 0.03 M. The results were typical of Pd-catalyzed H2 O2 synthesis in the presence of protons and halide anions [15,16,28]: The added protons increase the reaction rate by facilitating the reductive elimination of the H–Pd–OOH intermediate (a rate-determining step) and improve H2 O2 selectivity by protecting H2 O2 from decomposition [14]. At the same time, the increase of H+ concentration would increase Br− adsorption, thereby amplifying the beneficial effects of halide adsorption, such as deactivation of edge/corner Pd sites and activation of the H2 O2 formation route via combination of O2 − and H+ [15]. At 0.03 M H3 PO4 , the Br− adsorption reaches the highest value among the test conditions considered (Fig. 6), but the reversal of the H2 O2 activ-
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a) 25
100
60
40
b) 500
H3PO4: 0M
H3PO4: 0.01M H3PO4: 0.02M
20 Pd@S
Pd@SA(100)
Pd@SA(20)
Pd@SA(50)
Pd@S
Pd@SA(100)
Pd@SA(20)
Pd@SA(50)
Pd@S
Pd@SA(50)
Pd@SA(20)
0
Pd@SA(100)
Pd@S
10
Pd@SA(100)
15 Pd@SA(20)
80
Pd@SA(50)
20
5
0
H3PO4: 0.03M
of terrace Pd sites were poisoned and the activity was reduced as a result. Thus, it could also be supposed that such unfavorable events did not occur for ‘Pd@SA(100) or Pd@S under 0.02 M H3 PO4 ,’ which would be low in the actual acidity compared to the previous two catalysts. Under the 0.02 M H3 PO4 condition, the best activity (470 mmol H2 O2 /gPd h) was achieved using Pd@SA(100), which was intermediate in acidity among the tested catalysts. With the increase of H3 PO4 addition to 0.03 M, the activity was depressed, regardless of catalyst; this result is due to the high amount of Br− adsorption. Thus, the activity order of catalysts was reverse to the order regarding the surface acidity, and the highest H2 O2 production rate (377 mmol H2 O2 /gPd h) was observed with Pd@S, which had almost no acidity. The results in Figs. 6 and 7 support the idea that the dissolved acids in the reaction medium could be replaced by the solid acid sites (Brønsted site) on a Pd@SA(X) catalyst. Further study is required to establish a quantitative model to determine the optimal combinations of surface acidity, dissolved acid and halide anion concentrations to produce the best activity in a Pd@SA(X) catalyst.
400
4. Conclusion
Pd@S
Pd@SA(100)
Pd@SA(20)
Pd@SA(50)
Pd@S
Pd@SA(100)
Pd@SA(50)
Pd@SA(20)
Pd@S
Pd@SA(100)
Pd@SA(20)
Pd@SA(50)
0
Pd@SA(100)
100
Pd@S
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Pd@SA(20)
300
Pd@SA(50)
H2O2 production rate (mmol/gPdh)
H3PO4: 0.01M H3PO4: 0.02M H3PO4: 0.03M
H2O2 selectivity (%)
H2 conversion (%)
H3PO4: 0M
93
Fig. 8. Direct synthesis of hydrogen peroxide using Pd@SA(X) catalysts (X = 20, 50 and 100) under different H3 PO4 concentrations and in the presence of Br− ion: (a) H2 conversion and H2 O2 selectivity; (b) H2 O2 productivity (Pd-weight-specific H2 O2 production rate). Test conditions: 293 K, 1 atm, 0.05 g of catalyst, 150 mL of ethanol/water (4/1, v/v) mixture (containing 0.9 mM KBr and 0, 0.01, 0.02 or 0.03 M H3 PO4 ), stirring rate = 1,200 rpm, total gas flow rate = 22 mL/min, H2 /O2 = 1/10 (v/v).
ity (Fig. 7) implies that the adsorption exceeded a proper level, thus deactivating the terrace Pd sites and suppressing the H2 O2 formation rate. Fig. 8 presents the activity test results of Pd@SA(X), which were compared with those of Pd@S. When H3 PO4 was not added (H3 PO4 : 0 M), all of the activity indices (H2 conversion, H2 O2 selectivity and H2 O2 production rate) increased in the order of Pd@SA(20) > Pd@SA(50) > Pd@SA(100) > Pd@S, which was exactly consistent with the order of the surface (Brønsted) acidity. This result implies that the surface acid sites of Pd@SA(X) act similar to a dissolved acid (H3 PO4 ) regarding promotion of the H2 O2 synthesis activity. With the addition of 0.01 M H3 PO4 , the activities were largely improved, but the order of the catalysts in terms of activity was identical to the case of 0 M H3 PO4 . The most acidic catalyst, Pd@SA(20), showed the highest H2 O2 production rate (468 mmol H2 O2 /gPd h). However, when the H3 PO4 addition increased to 0.02 M, the change of activity was different according to the surface acidity of a catalyst: Pd@SA(20) and Pd@SA(50), which were the highest two of the tested catalysts regarding surface (Brønsted) acidity, were depressed regarding H2 conversion and H2 O2 selectivity. In contrast, Pd@SA(100) and Pd@S, which were the lowest two regarding surface (Brønsted) acidity, were largely improved in terms of their activities. Considering that the surface Brønsted acid sites are involved in the reaction and catalysis in a manner similar to that of the dissolved acid, it could be supposed that the actual acidity level of ‘Pd@SA(20) or Pd@SA(50) under 0.02 M H3 PO4 ’ was high enough to make Br− adsorption excessive, by which a part
The core/shell structured, Pd@SiO2 –Al2 O3 catalysts (Pd@SA(X), X = (Si/Al)mol = 20, 50, 100) were successfully prepared by applying Stöber method. One or two Pd crystals (3 wt.%) were encapsulated by a SiO2 –Al2 O3 shell (50–70 nm), and the size of Pd crystals (7–13 nm) was almost unchanged after calcination. The micropores were well developed in the catalysts, and the exposed surface area of Pd was in the range of 1.4–1.8 m2 /gPd . The SiO2 –Al2 O3 shell exhibited surface (Brønsted) acidity that increased along with the Al content. In the Br− adsorption experiments, Br− adsorption increased with the increase of the surface acidity on the Pd@SA(X) catalysts. It was supposed that Br− ions adsorbed on Pd, thereby forming Pd·H+ ·Br− complexes and that the Brønsted acid sites on Pd@SA(X) affect the Br− ion adsorption in a manner similar to a dissolved acid. The H2 O2 synthesis activity of the Pd@SA(X) catalysts was largely influenced by the surface (Brønsted) acidity of the catalyst and the concentration of dissolved acid (H3 PO4 ). The more acidic Pd@SA(X) catalyst exhibited a higher H2 O2 formation activity when the concentration of dissolved acid was lower than 0.01 M. As the concentration of H3 PO4 increased, the activity was suppressed in the order of surface acidity. The least acidic catalyst (Pd@S) exhibited the highest activity when the concentration of H3 PO4 was higher than 0.03 M. It was supposed that the influence of the acids (either heterogeneous or homogenous acids) on H2 O2 formation activity is related to how much the adsorption of halide anion is promoted. The use of Pd@SA(X) catalysts makes it possible to reduce the addition of an external proton source, which is typically required to enhance the H2 O2 synthesis activity of Pd catalysts. It is anticipated that the use of Pd@SA(X) catalysts could not only help to reduce the process cost but also relieve the side effects caused by the use of acidic liquid medium. Acknowledgments This work was supported by the National Research Foundation of Korea (NRF) and grant funded by the Korea Government (MSIP) (NRF-2012M1A2A2671682). References [1] X.S. Chai, Q.X. Hou, Q. Luo, J.Y. Zhu, Anal. Chim. Acta 507 (2004) 281–284. [2] M. Ksibi, Chem. Eng. J. 119 (2006) 161–165. [3] W. Kern, J. Electrochem. Soc. 137 (1990) 1887–1892.
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