FTIR Study of Adsorption and Reactions of Ethylene Oxide on

May 8, 2008 - Adsorption, thermal decomposition, and photoreactions of ethylene oxide (EO) on powdered TiO2 have been studied by Fourier-transformed ...
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J. Phys. Chem. C 2008, 112, 8365–8371

8365

FTIR Study of Adsorption and Reactions of Ethylene Oxide on Powdered TiO2 Chen-Fu Lien, Chia-Hsun Ho, Chun-Yi Shieh, Chien-Lin Tseng, and Jong-Liang Lin* Department of Chemistry, National Cheng Kung UniVersity, 1, Ta Hsueh Road, Tainan, Taiwan 701, Republic of China ReceiVed: December 13, 2007; ReVised Manuscript ReceiVed: March 14, 2008

Adsorption, thermal decomposition, and photoreactions of ethylene oxide (EO) on powdered TiO2 have been studied by Fourier-transformed infrared spectroscopy. Most of the adsorbed EO molecules at the saturated coverage on 35 °C TiO2 remain intact. As the temperature is increased (g100 °C), the surface EO molecules dissociate into ethylene glycol-like species (-OCH2CH2OH or -OCH2CH2O-) by ring rupture and crotonaldehyde (CH3CHdCHCHO) possibly via acetaldehyde from EO isomerization. The ring-opening process is enhanced in the presence of coadsorbed H2O molecules. EO with the highly strained three-membered ring is thermally more reactive than 1,4-dioxane on TiO2. EO molecules adsorbed on TiO2 readily undergo photodegradation in O2, forming surface H2O, CO2, HCO3, CO3, and HCOO. The comparative studies of 16 O2 and 18O2 indicate that both O2 and lattice oxygen take part in the EO photodecomposition. The reaction pathways involving the two oxygen species are discussed. Introduction Direct oxidation of ethylene to ethylene oxide is an important industrial process in which silver supported on oxides is often used as catalyst. However, the selectivity for the ethylene oxide formation is significantly diminished due to its further isomerization and oxidation. The effect of bare oxides, such as TiO2, SiO2, and Al2O3, on the reactions of ethylene oxide under conditions relevant to ethylene epoxidation over supported silver has been demonstrated.1 In the case of TiO2, ethylene oxide is converted to acetaldehyde with some minor products of CO, CO2, and dioxane at 239 °C, using feed comprised of ethylene oxide, O2, and He. The isomerization activity is proposed to be related to the surface acidity.2 Vibrational spectroscopy has been employed to investigate the adsorption and thermal decomposition of ethylene oxide on TiO2, but inconsistent assignments for the chemical species on the surfaces are presented from different research groups. Loiko et al. suggest the formation of polymeric molecules and carboxylates, in contrast to ethylene glycol-like species and bicarbonate proposed by Yong et al.1,3 Ethylene oxide is an industrial chemical used to produce glycols, glycol ethers, ethanolamines, polymers, and surfactants and is also used as a disinfectant. It is recognized that ethylene oxide is a hazardous chemical and can cause reproductive and genetic disorder and cancer. Different technologies have been developed to control the emission of ethylene oxide, using wetscrubbers, thermal oxidizers, catalytic oxidizers, and sorbents.4 In the present research, we investigate the adsorption and temperature-dependent decomposition of ethylene oxide on TiO2, focusing on the surface reaction process using FTIR spectroscopy, and make a comparison to the reactions of 1,4dioxane (C4H8O2). The promoting effect of adsorbed water on the decomposition of ethylene oxide is also shown. At a temperature higher than 100 °C, ethylene oxide reacts on the surface, forming crotonaldehyde and ethylene glycol-like species. However, in the presence of coadsorbed water, ethylene oxide readily decomposes into ethylene glycol-like species at * Corresponding author. Phone: 886-6-275-7575 ext. 65326. Fax: 8866-274-0552. E-mail: [email protected].

35 °C, without forming crotonaldehyde or acetaldehyde. Moreover, the photochemical oxidation of ethylene oxide on TiO2 in the presence of O2 is studied. Oxygen isotopes (16O2 and 18O ) are used to assist in identification of the reaction 2 intermediates and products formed and to understand the roles of O2 and the lattice oxygen in the photoreactions of ethylene oxide on TiO2. TiO2 used as a photocatalyst to oxidize organic compounds has been extensively studied. The photooxidation processes originate from the band gap excitation of TiO2, generating electron-hole pairs. The holes can diffuse to the surface and induce the oxidation of surface molecules. In the present investigation, ethylene oxide is found to be photooxidized to CO2, formate, bicarbonate, and carbonate. The oxygen isotopic study reveals that both TiO2 lattice oxygen and O2 participate in the formation of CO2 and formate. Experimental Section The ethylene oxide used in this study was prepared by the reaction between 2-chloroethanol and potassium hydroxide, following the method reported previously.5 The sample preparation of TiO2 powder supported on a tungsten fine mesh (∼6 cm2) has been described previously.6,7 In brief, TiO2 powder (Degussa P25, ∼50 m2/g, anatase 70%, rutile 30%) was dispersed in water/acetone solution to form a uniform mixture that was then sprayed onto a tungsten mesh. After that, the TiO2 sample was mounted inside the IR cell for simultaneous photochemistry and FTIR spectroscopy. The IR cell with two CaF2 windows for IR transmission down to 1000 cm-1 was connected to a gas manifold that was pumped by a turbomolecular pump with a base pressure of ∼1 × 10-7 Torr. The TiO2 sample in the cell was heated to 450 °C under vacuum for 24 h by resistive heating. The temperature of the TiO2 sample was measured by a K-type thermocouple spotwelded on the tungsten mesh. Before each run of the experiment, the TiO2 sample was heated to 450 °C in a vacuum for 2 h. After the heating, 10 Torr of O2 was introduced into the cell as the sample was cooled to 70 °C. When the TiO2 temperature reached 35 °C, the cell was evacuated for gas dosing and an infrared spectrum was taken as reference background. 16O2 (99.998%) and 18O2 (99

10.1021/jp711700d CCC: $40.75  2008 American Chemical Society Published on Web 05/08/2008

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Figure 1. Infrared spectra of ethylene oxide molecules adsorbed on TiO2 with use of a three-step process. All of the spectra were taken at 35 °C with 50 scans.

atom%) were purchased from Matheson and Isotec, respectively. Pressure was monitored with a Baratron capacitance manometer and an ion gauge. In the photochemistry study, both the UV and IR beams were 45 °C to the normal of the TiO2 sample. The UV light source used was a combination of a 350-W Hg arc lamp (Oriel Corp), a water filter, and a band-pass filter with a bandwidth of ∼100 nm centered at ∼320 nm (Oriel 51650). The photon power at the position of the TiO2 sample was ∼0.24 W/cm2 measured in the air by a power meter (Molectron, PM10V1). Infrared spectra were obtained with a 4 cm-1 resolution by a Bruker FTIR spectrometer with a MCT detector. The entire optical path was purged with CO2-free dry air. The spectra presented here have been ratioed against the TiO2 background spectrum. In the study of photooxidation, the photoirradiation time was started to count as the UV lamp was turned on. It took 40-50 s to reach full power. Results and Discussion In this section, the experimental results of ethylene oxide are presented in the order of adsorption and thermal decomposition, effect of coadsorbed water on the reaction behavior, and photoreactions in the presence of 16O2 or 18O2. Adsorption and decomposition of 1,4-dioxane are also shown in comparison to those of ethylene oxide. Adsorption and Thermal Reactions of Ethylene Oxide on TiO2. The infrared spectra of EO adsorbed on TiO2 at three coverages were obtained after a three-step dosing process and are shown in Figure 1. In the first step, the TiO2 surface at 35 °C was exposed to ∼9.0 × 1018 EO molecules, followed by evacuation and infrared measurement (Figure 1a). After that, the TiO2 was in contact with ∼1.8 × 1019 EO molecules and the infrared cell was evacuated. The spectrum of Figure 1b was then obtained. The TiO2 was further exposed to ∼1.8 × 1019 EO molecules in the final step and Figure 1c was taken in a vacuum. After the three-step dosing process, the TiO2 surface has reached the saturated EO coverage. The three spectra in Figure 1 have a similar infrared absorption feature. The adsorption of EO induces a decrease of isolated OH groups on the TiO2, as evidenced by the presence of the negative peaks between 3600 and 3800 cm-1. The peak frequencies of EO on TiO2 at the saturated coverage (Figure 1c) are listed in Table 1 and compared to those of EO in the solid phase. As shown, most of the observed absorption frequencies in Figure 1 are

Lien et al. close to those of EO molecules in the solid state and can be assigned to the EO fundamental modes, combinations, or overtones, except for the peaks at 1369, 1459, and 2861 cm-1, suggesting that most of the adsorbed-EO molecules are stable on the TiO2 surface at 35 °C. In the previous infrared study of hydrogen-bonded EO-HI complex present in Ar matrix, it has been shown that the ν2, ν5, ν10, and ν12 vibrational modes of EO are red-shifted by 6.5, 9.5, 3.5, and 17.5 cm-1, respectively.9 This result can help to reveal the possible surface bonding of the EO molecules in the present study. The CH2-bending (ν2) frequency of adsorbed EO on TiO2 appears at 1486 cm-1, lower than the 1492 cm-1 for EO in the solid phase8 and the 1496 cm-1 for EO trapped in solid argon.9 This red shift in the infrared absorption can be attributed to the EO molecules interacting with surface OH groups and/or Lewis acid sites through the oxygen atom in EO. The three peaks of 1369, 1459, and 2861 cm-1 not belonging to adsorbed EO on TiO2 are likely due to the glycol-like species of -OCH2CH2OH and/or -OCH2CH2O-. Formation of ethylene glycol-like species from decomposition of EO at elevated temperatures is clearly demonstrated later. Figure 2 shows the infrared spectra of the TiO2 surface exposed to ∼9.0 × 1018 EO molecules at 35 °C, followed by evacuation at the following temperatures, 50, 100, 150, 200, and 250 °C for 1 min. The adsorbed EO molecules almost disappear after heating the surface to 200 °C in a vacuum and new peaks appear at 1070, 1120, 1138, 1360, 1446, 1616, 1635, 2868, 2921, and 2948 cm-1, which are attributed to the formation of crotonaldehyde (CH3CHdCHCHO) and the glycollike species (-OCH2CH2OH and/or -OCH2CH2O-) from thermal decomposition of EO. The CdO stretching band of crotonaldehyde on TiO2 has been shown to appear between ∼1640 and 1660 cm-1, with the CdC band (approximately between 1610 and 1630 cm-1) as a shoulder.10 Ethylene glycollike species have strong absorptions between 1000 and 1200 cm-1 for the C-O stretching and between 2800 and 3000 cm-1 for the CH2 stretching.10 The formation of crotonaldehyde suggests that EO isomerizes to form acetaldehyde, because acetaldehyde molecules recombine readily to generate crotonaldehyde on rutile and anatase TiO2.10,11 Crotonaldehyde on TiO2 may decompose into crotonate, with two broad peaks at ∼1422 and ∼1518 cm-1, at a temperature higher than 200 °C.10,12 Scheme 1 shows the ring-opening reaction pathways of EO in the formation of acetaldehyde involving surface Lewis acid sites13,14 and dioxyethylene involving Ti-O or HO-Ti-O groups. Previously, a mechanism of isomerization of EO to form acetaldehyde via dioxyethylene has been proposed on oxidized silver surfaces only with Lewis acid sites.13,14 However, in our case there is no evidence showing that dioxyethylene can react on TiO2 to form acetaldehyde.10 On silica, CH3CHOH+ · · · -OSi has been suggested to be the surface intermediate in the formation of crotonaldehyde from acetaldehyde reaction.15 The surface CH3CHOH+ can react with the enol form of acetaldehyde (CH2dCH-OH), followed by loss of H+ to form the aldol species CH3-C(OH)H-CH2-C(O)H. Through dehydration the aldol transforms into crotonaldehyde.15 In addition to the chemical reactions of EO on TiO2, the relative integrated peak intensities of EO itself also vary with surface temperature. In Figure 2, the peak area of the 1268 cm-1 (ν3) increases by ∼50% as the temperature is increased from 35 to 50 °C, in contrast to the comparable area of the 3016 cm-1 peak (ν1) at these two temperatures. The area ratios of ν1/ν3 (Iν1/Iν3) in the 35, 50, 100, and 150 °C spectra are 1.46, 1.13, 0.77, and 0.52, respectively. The temperature-dependent peak-area ratios are likely due to

Reactions of Ethylene Oxide on Powdered TiO2

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TABLE 1: Comparison of the Infrared Frequencies (cm-1) of Ethylene Oxide amorphous solid phasea

crystalline phasea

adsorption on TiO2 (35 °C)

3073 3058 3011 3001 2962 2923

3073 3062 3005 2995 2950 2914 2904

3077

1580 1492 1467 1267 1152 1043 856 821 794 a

1583

3013 2957 2918 2861 1642 1570 1486

1480 1467 1266 1169 1160 1147 1046 859 854 818 796

1459 1369 1268 1156 1124

assignmenta CH2 stretching, ν6 CH2 stretching, ν13 CH2 stretching, ν1 CH2 stretching, ν9 combination, ν2 + ν10 overtone, 2ν2 overtone, 2ν10 overtone, 2ν8 overtone, 2ν15 CH2 bending, ν2 ν13 - 2ν15 CH2 bending, ν10

adsorption on TiO2 (200 °C)

2948 2921 2868 1635 1616

1446 1360 ring breathing, ν3 CH2 wagging, ν11 CH2 twisting, ν14 CH2 wagging, ν4 CH2 twisting, ν7 ring bending, ν12 ring bending, ν5 CH2 rocking, ν8 CH2 rocking, ν15

1138 1120 1070

Reference 8.

Figure 2. Infrared spectra of a TiO2 surface taken after being in contact with ∼9 × 1018 ethylene oxide molecules at 35 °C followed by evacuation at this temperature and then heating at 50, 100, 150, 200, 250, and 300 °C for 1 min in a vacuum. A heating rate of ∼2 deg/s was used and all of the spectra were measured at 35 °C with 50 scans.

the variation in the square of the change in dipole moment with respect to the normal coordinate of the two vibrational modes.16 Previously, the Iν1/Iν3 was also found to be varied with the coverage of EO condensed on a water film at 30 K.8 Effect of Water on the Decomposition of Ethylene Oxide. Figure 3 compares the infrared spectra of TiO2 surfaces covered only with EO, prepared by exposure of ∼6 × 1018 molecules, and covered with both EO and H2O (1621 cm-1) molecules,

prepared by subsequent exposure of ∼3 × 1019 H2O molecules to the EO-covered TiO2. In the latter case, exposure of H2O vapor to the EO-precovered TiO2 surface causes the strong, broad hydrogen-bonding absorption between 2500 and 3700 cm-1 and reduced EO peak intensities. Water-promoted decomposition of EO, forming the glycol-like species, is evidenced by the appearance of its characteristic peaks at 1078, 2862, and 2936 cm-1. The presence of H2O may decrease the activation energy and/or furnish the surface with neighboring hydroxyl groups to accelerate the transformation of EO to the glycollike species. From the thermodynamic point of view, this result is driven by the highly strained three-membered ring of the EO molecule. Ethylene glycol produced by hydration of EO, which is usually operated at 140-230 °C and 1.5-2.5 MPa with a large excess of water, is a well-known process.17 Niobium oxidebased catalysts have been employed to increase the selectivity toward ethylene glycol formation.17 In the present study, it is found that the reaction of EO to form the glycol-like species proceeds readily on TiO2 at 35 °C, so TiO2 can be a potential catalyst for generating ethylene glycol from EO at low temperatures. Adsorption of 1,4-Dioxane on TiO2. Figure 4 shows the infrared spectra of the TiO2 surface after being in contact with ∼3 × 1019 1,4-dioxane molecules at 35 °C, followed by evacuation at the indicated temperatures for 1 min. Note that all of the spectra were measured at 35 °C in a vacuum. In the 35 °C spectrum, the major bands appear at 1050, 1071, 1123, 1254, 1293, 1311, 1371, 1442, 1455, 1465, 2871, 2902, 2932, 2957, and 2967 cm-1. As shown in Table 2, it is found that the frequencies observed after 1,4-dioxane adsorption are similar to those of 1,4-dioxane in the gas and liquid phase. This result indicates that 1,4-dioxane molecules remain intact on TiO2 at 35 °C. The negative bands in the region 3550-3850 cm-1 and the enhanced absorptions between 3000 and 3550 cm-1 reveal that the adsorbed 1,4-dioxane molecules interact with the surface OH groups, leading to the decrease of isolated hydroxyls and formation of hydrogen bonding. After heating to 100 °C, the

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1,4-dioxane absorption feature, such as peak frequency and intensity, is similar to the 35 °C one. At 150 °C, all of the bands reduce in intensity without new infrared bands being detected. This result is attributed to 1,4-dioxane desorption. Further heating to 200 °C causes continuous decrease of the 1,4-dioxane bands. In the 250 °C spectrum, the adsorbed 1,4-dioxane is almost removed from the surface, as evidenced by the very small band at 1254 cm-1, but the relatively strong bands located at 1071, 2874, and 2932 cm-1 indicate the formation of the glycollike species. These bands almost disappear after heating the surface to 300 °C (not shown). No characteristic peaks showing crotonaldehyde formation are observed. Furthermore, in contrast to the effect of H2O on the decomposition of ethylene oxide at 35 °C, 1,4-dioxane molecules remain intact on TiO2 with coadsorption of H2O even when the temperature is raised to 75 °C (Supporting Information, Figure 1). Photodecomposition of EO on TiO2. Figure 5 shows the infrared spectra taken before (a) and after 2, 5, 30, 90, and 180 min (b-f) of photoirradiation during UV exposure of TiO2 covered with EO molecules initially in 10 Torr of O2 in a closed cell. The EO-covered surface was prepared by exposing a clean TiO2 surface to ∼9.0 × 1018 EO molecules at 35 °C, followed by evacuation. The temperature of the TiO2 surface subjected to the photoillumination was increased to 68 °C. Spectra g and

Figure 3. Infrared spectra of a TiO2 surface covered with ethylene oxide only and with both ethylene oxide and water at 35 °C, showing the effect of H2O on the enhancement of ethylene oxide ring rupture.

h of Figure 5 were taken after the irradiation and subsequent evacuation of the reaction cell. The amount of adsorbed EO decreases gradually with an increase of photoirradiation time, as evidenced by the reduction of both its characteristic peaks at 1270 and 3018 cm-1. In the spectrum of Figure 5g, new peaks appear at 1119, 1218, 1327, 1357, 1438, 1561, 1575, 1620, 2361, 2872, 2919, 2952, and 3606 cm-1. In addition, the infrared absorptions from surface OH groups between 3600 and 3800 cm-1 are enhanced. The 2361 cm-1 peak is due to adsorbed CO2.21 On ZnO, it has been reported that bidentate bicarbonate (HCO3) has four bands at 1224, 1425, 1636, and 3616 cm-1

Figure 4. Infrared spectra of a TiO2 surface taken after being in contact with ∼3 × 1019 1,4-dioxane molecules at 35 °C, followed by evacuation at the following temperatures, 50, 100, 150, 200, and 250 °C. A rate of ∼2 deg/s was used to raise the surface temperature. All of the spectra were obtained at 35 °C with 50 scans.

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TABLE 2: Comparison of the Infrared Frequencies (cm-1) of 1,4-Dioxane vapora 1052 1086 1136 1256 1291 1369 1378 1449 1457

liquidb 1046 1081 1120 1251 1286 1318 1368 1446

2863

2860 2899 2921

2970

2967

a

1049 1083 1125 1256 1289 1322 1366 1374 1449 1454 1470 2852 2889 2914 2960

adsorption on TiO2, 35 °C

modec

1050 1071 1123 1254 1293 1311 1371

CH2 rocking CH2 rocking C-O stretching CH2 twisting CH2 twisting combination CH2 wagging CH2 wagging CH2 bending CH2 bending combination CH2 stretching combination combination CH2 stretching CH2 stretching

1442 1455 1465 2871 2902 2932 2957 2967

Reference 18. b References 19 and 20. c Reference 20.

Figure 6. Infrared spectra taken before (a) and after 2, 5, 30, 90, and 180 min (b-f) of photoirradiation during UV exposure of TiO2 covered with ethylene oxide molecules initially in 10 Torr of 18O2 in a closed cell. The EO-covered TiO2 surface was prepared by doing ∼9.0 × 1018 molecules at 35 °C, followed by evacuation. Both g and h are postirradiation spectra. The ranges between 2250 and 2500 cm-1 and between 2650 and 3150 cm-1 have been multiplied by a factor of 5 and 3, respectively.

Figure 5. Infrared spectra taken before (a) and after 2, 5, 30, 90, and 180 min (b-f) of photoirradiation during UV exposure of TiO2 covered with ethylene oxide molecules initially in 10 Torr of 16O2 in a closed cell. The EO-covered TiO2 surface was prepared by doing ∼9.0 × 1018 molecules at 35 °C, followed by evacuation. Both g and h are postirradiation spectra. The ranges between 2250 and 2500 and between 2650 and 3150 cm-1 have been multiplied by a factor of 5 and 3, respectively.

and bidentate carbonate (CO3) has bands at 1030, 1330, and 1520 cm-1.22 On TiO2, the carbonate-like bands have been observed at 1220, 1337, 1440, 1520, and 1577 cm-1.22 Accordingly, in our case the peaks observed at 1218, 1438, and 3606 cm-1 are attributed to the formation of bicarbonate and are assigned to the OH bending, OCO stretching, and OH stretching modes, respectively. The 1327 and 1575 cm-1 peaks are assigned to the OCO and CdO stretching vibrations of bidentate carbonate.21,22 The postirradiation evacuation (Figure 5h) leads to the disappearance of adsorbed CO2 and significant reduction of both carbonate species. In addition to bicarbonate and

carbonate from the EO photodecomposition, formate (HCOO) is generated as well, with the characteristic bands located at 1357, 1561, 2872, 2919, and 2952 cm-1. HCOO also has been observed previously in the dissociative adsorption of formaldehyde and formic acid on TiO2.23,24 The 1357 and 1561 cm-1 bands can be assigned to the OCO symmetric and antisymmetric stretching: 2872 cm-1 to CH stretching; 2919 and 2952 cm-1 to the combination of the OCO stretching and CH bending modes.23 The 1620 cm-1 band in Figure 5h is due to the absorption of H2O generated in the reaction. The broad absorption at ∼1119 cm-1 can be ascribed to -OCH2CH2Oor --OCH2CH2OH formation, but due to the thermal effect upon photoirradiation of the TiO2 surface. This effect is revealed by the thermal control experiment carried out by using the same experimental conditions as those in Figure 5, but annealing the surface at 68 °C to replace UV exposure (Supporting Information, Figure 2). As a brief conclusion, photooxidation of EO on TiO2 in the presence of O2 generates CO2, HCO3, CO3, HCOO, and H2O on the surface. This EO photodecomposition process catalyzed by TiO2 is facile, as indicated by the fast appearance of the product peaks in the 2-min spectrum in Figure 5. To investigate the participation of dioxygen and TiO2 lattice oxygen in the photoproduct formation, 18O2 was employed in the photoirradiation of TiO2 covered with EO molecules. The infrared spectroscopic results are shown in Figure 6. UV irradiation of TiO2 covered with EO in 18O2 should have the same photochemistry as in 16O2; however, the study of isotopic oxygen incorporated in the photooxidation products measured by peak shift in frequency can shed light on the reaction mechanism. Previously, oxygen isotopic exchange has been shown in the photoirradiation of C16O2, C16O3, or HC16O16O

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on TiO2 in the presence of 18O2 for prolonged times.21 Therefore, in the present study identification of the photooxidation products containing 18O is particularly focused on the initial stage of the photoirradiation of the TiO2 surface. In the 2- and 5-min spectra of Figure 6, the absorption band for the adsorbed CO2 in the 2250-2450 cm-1 region becomes broader, as compared to the 16O case. The absorption maxima are located at 2342 cm-1, 2 indicating that 18O can be incorporated into the carbon dioxide from EO photodecomposition. However, not all of the carbon dioxide molecules contain 18O atoms, because substantial absorption of adsorbed C16O2 is found to be the left-hand shoulder at 2361 cm-1. The CO3 peak is also red-shifted from 1573 to 1565 cm-1 as 16O2 is replaced by 18O2. Note that the 1573 cm-1 has been assigned to the CdO stretching mode of bidentate carbonate. In the range of 1290-1410 cm-1, the absorption at ∼1330 cm-1 relative to that at ∼1360 cm-1 is enhanced in 18O2, as compared to the case of 16O2. This enhancement is due to the formation of adsorbed formate groups with 18O atoms. It has been reported that the symmetric

stretching frequencies of -C16O16O-, -C16O18O-, and -C18O18O-, measured for H13C16O16ONa, H13C16O18ONa, and H13C18O18ONa, are 1340, 1315, and 1297 cm-1, respectively.25 The antisymmetric stretching frequencies of -C16O16O- for H12C16O16ONa and -C18O18O- for H12C18O18ONa are 1607 and 1587 cm-1.26 Although 18O-incorported formate is observed in Figure 6, the relatively strong band at 1360 cm-1 clearly reveals the formation of HC16O16O in the presence of 18O2. The oxygen isotope effect displayed in Figures 5 and 6 shows that O2 and TiO2 lattice oxygen involve in the product formation of EO photoreaction on TiO2. As TiO2 particles absorb photons with energy enough to induce band-to-band transition, electron-hole pairs can be generated and diffuse to the surfaces. The photogenerated holes may recombine with adsorbed molecules, trapped at surface oxygen sites to form Ti-O-• species, and react with surface OH or H2O to form OH•. The photoreaction of EO adsorbed on TiO2 in the presence of O2 may proceed with diverse pathways. The photoreaction routes that possibly occur for EO are proposed

Reactions of Ethylene Oxide on Powdered TiO2 and shown in Scheme 2. Pathway A in Scheme 2, involving the dioxy four-membered ring (A2), is similar to the mechanism that has been proposed for 1,4-dioxane photoreaction catalyzed by TiO2 dispersed in aerated water to form H(O)COCH2CH2OC(O)H (ethylene glycol diformate).27 The A3 species derived from A2 rearrangement decomposes on TiO2 to generate 18O-containing products. In this pathway, no lattice oxygen is involved in the product formation. However, it is believed that EO must be bonded to the lattice oxygen in the photochemical process to generate the products of HC16O16O and C16O2, as observed in Figure 6. The B1 shown in pathway B, possibly formed through recombination of A1 and Ti-O-• or through the reaction between the lattice oxygen and EO+ produced right after the hole capture of adsorbed EO molecules, is proposed to be the chemically bonded EO. In pathway C, the dioxy four-membered ring (C1) is generated from B1 after hole capture and 18O2 addition and may decompose into 18O-containing products on TiO2. Pathway D shows the formation of the tetraoxide species (D1). Tetraoxides have been proposed to be the reaction intermediates in the radiolysis of aqueous solutions of organic compounds in the presence of O2 and in the photooxidation of C8 organics with TiO2-coated glass microbubbles.28 Moreover, in the study of photooxidative degradation of CH3OH vapor in contact with Pt/TiO2 the mechanism involving organoperoxy and tetraoxide has also been reported: CH3OH + h+ f •CH2OH + H+, •CH2OH + O2 f HOCH2OO•, HOCH2OO• + •OOH f HOCHOOOOH f HCOOH + O2 + H2O.29 D1 can decompose into water and the three-membered lactone ring (D2), which is likely to be the short-lived intermediate for the formation of HC16O16O and C16O as it dissociates in the subsequent photochemical steps on the TiO2 surface. Conclusion Most of the adsorbed EO molecules decompose on TiO2 at a temperature g100 °C to generate ethylene glycol-like species and crotonaldehyde which originate from EO ring rupture and isomerization, respectively. EO is more reactive than its homologue, 1,4-dioxane, which decomposes at a TiO2 temperature g200 °C. The EO ring rupture process is found to be promoted in the presence of water molecules on the surface. The photodegradative reaction of EO on TiO2 in the presence of O2 is facile. Both dioxygen and TiO2 lattice oxygen play an important role in the EO photoreactions. Acknowledgment. This research was supported by the National Science Council of the Republic of China (NSC-952113-M-006-017).

J. Phys. Chem. C, Vol. 112, No. 22, 2008 8371 Supporting Information Available: Figure 1 shows the infrared spectra of 1,4-dioxane on TiO2 with and without coadsorbed H2O at 35 and 75 °C, and Figure 2 shows the infrared spectra obtained for the thermal experiment in contrast to the photoirradiation of ethylene oxide on TiO2 after holding the surface at 68 °C for 180 min in 10 Torr O2. This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes (1) Yong, Y.-S.; Kennedy, E. M.; Cant, N. W. Appl. Catal. 1991, 76, 31–48. (2) Lee, J. K.; Verykios, X. E.; Pitchai, R. Appl. Catal. 1988, 44, 223– 237. (3) Loiko, V. E.; Davydov, A. A. Zh. Prikl. Spektrosk. 1987, 47, 231– 237. (4) Arsenijevic´, Z, L.; Grbic´, B. V.; Radic´, N. D.; Grbavcˇic´, Zˇ. B. Chem. Eng. J. 2006, 116, 173–178. (5) Cox, J. D.; Warne, R. J. J. Chem. Soc. 1951, 1893. (6) Basu, P.; Ballinger, T. H.; Yates, J. T., Jr ReV. Sci. Instrum. 1988, 59, 1321. (7) Wong, J. C. S.; Linsebigler, A.; Lu, G.; Fan, J.; Yates, J. T., Jr J. Phys. Chem. 1995, 99, 335. (8) Schriver, A.; Coanga, J. M.; Schriver-Mazzuoli, L.; Ehrenfreund, P. Chem. Phys. 2004, 303, 13–25. (9) Bernadet, P.; Schriver, L.; Schriver, A.; Perchard, J.-P. J. Phys. Chem. 1988, 92, 7204–7210. (10) Wu, W.-C.; Yang, S.-J.; Ho, C.-H.; Lin, Y.-S.; Liao, L.-F.; Lin, J.-L. J. Phys. Chem. B 2006, 110, 9627–9631. (11) Rekoske, J. E.; Barteau, M. A. Langmuir 1999, 15, 2061–2070. (12) Coloma, F.; Baeza-Bachiller, B.; Rochester, C. H.; Anderson, J. A. Phys. Chem. Chem. Phys. 2001, 3, 4817–4825. (13) Grant, R. B.; Lambert, R. M. J. Catal. 1985, 93, 92–99. (14) Tan, S. A.; Grant, R. B.; Lambert, R. M. J. Catal. 1987, 106, 54– 64. (15) Young, R. P.; Sheppard, N. J. Catal. 1967, 7, 223–233. (16) Colthup, N. B.; Daly, L. H. Introduction to Infrared and Raman spectroscopy; Academic Press, Inc.: New York, 1990. (17) Li, Y.; Yan, S.; Qian, L.; Yang, W.; Xie, Z.; Chen, Q.; Yue, B.; He, H. J. Catal. 2006, 241, 173–179. (18) Shimanouchi, T. Tables of Molecular Vibrational Frequencies Consolidated, Vol. II. J. Phys. Chem. Ref. Data 1972, 6, 993. (19) Malherbe, F. E.; Bernstein, H. J. J. Am. Chem. Soc. 1952, 74, 4408. (20) Ellestad, O. H.; Klaboe, P.; Hagen, G. Spectrochim. Acta 1971, 27, 1025. (21) Liao, L.-F.; Lin, C.-F.; Shieh, D.-L.; Chen, M.-T.; Chen, M.-T.; Lin, J.-L J. Phys.Chem. B. 2002, 106, 11240–11245. (22) Boccuzzi, F,; Chiorino, A J. Phys. Chem. 1996, 100, 3617–3624. (23) Busca, G.; Lamotte, J.; Lavalley, J.-C.; Lorenzelli, V. J. Am. Chem. Soc. 1987, 109, 5197–5202. (24) Chuang, C.-C.; Wu, W.-C.; Huang, M.-C.; Huang, I.-C.; Lin, J.-L. J. Catal. 1999, 185, 423–434. (25) Spinner, E. Spectrochim. Acta 1987, 43A, 301. (26) Spinner, E.; Rowe, J. E. J. Chem. 1979, 32, 481. (27) Hill, R. R.; Jeffs, G. E.; Roberts, D. R. J. Photochem. Photobiol., A 1997, 108, 55. (28) Schwitzgebel, J.; Ekerdt, J.; Gerischer, H.; Heller, A. J. Phys. Chem. 1995, 99, 5633. (29) Sadeghi, M.; Liu, M.; Zhang, T.-G.; Stavropoulos, P.; Levy, B. J. Phys. Chem. 1996, 100, 19466.

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