FTIR Study of Adsorption and Reactions of Methylamine on Powdered

Jun 5, 2001 - It is found that methylamine is adsorbed molecularly on TiO2 at 35 °C. ... After that, the TiO2 sample was mounted inside the IR cell f...
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J. Phys. Chem. B 2001, 105, 5928-5934

FTIR Study of Adsorption and Reactions of Methylamine on Powdered TiO2 Li-Fen Liao, Wen-Chun Wu, Chih-Chung Chuang, and Jong-Liang Lin* Department of Chemistry, National Cheng Kung UniVersity, Tainan, Taiwan 701, Republic of China ReceiVed: NoVember 22, 2000; In Final Form: April 9, 2001

The adsorption, thermal reactions, and photochemistry of methylamine on TiO2 have been investigated using infrared spectroscopy. It is found that methylamine is adsorbed molecularly on TiO2 at 35 °C. Hydrogen exchange occurs between the NH2 group of adsorbed methylamine and surface OH groups. Thermal reactivity of methylamine is enhanced in the presence of O2, producing NH3(g), CO2(g), and other oxygenated compounds. UV irradiation of TiO2 in the presence of methylamine and oxygen generates NH3(g), CO2(g), H2O(g), NCO(a), and HCOO(a). Possible reaction steps for the photoproducts are discussed.

Introduction Adsorption and thermal reactions of methylamine (CH3NH2) on metal surfaces have been investigated in elucidating the surface bonding and reactivity of competing C-H, N-H, and C-N bond activation. Studies of high-resolution electron energy-loss spectroscopy (HREELS) and infrared reflectionabsorption spectroscopy have suggested that adsorbed methylamine interacts with Cu, Ni, and Cr surfaces through the nitrogen lone pair.1,2 On Ru(001),3,4 Rh(111),5 and Ni(111)6 surfaces, methylamine decomposes with preferential C-H and N-H bond scission. After the bond breakage events, CN groups are suggested to be present on the surfaces by HREELS studies and responsible for the subsequent thermal reaction products formed at higher temperatures. On Mo(100)-c(2 × 2)-N,7 dissociation of adsorbed methylamine is initiated by intramolecular hydrogen transfer to form adsorbed CH2 and NH3 as indicated by D2/CH3NH2 coadsorption study. Reactivity and thermal product distribution of methylamine have been compared on W(100) and its two modified surfaces of W(100)-(5 × 1)-C and W(100)-(2 × 1)-O.8 The bare W(100) surface has the highest reactivity toward methylamine decomposition to form H2, N2, and NH3 among the three surfaces. In contrast, on the oxygen atom preadsorbed tungsten surface methylamine is mainly adsorbed reversibly without showing significant decomposition. Availability of electron density from the surfaces to methylamine’s antibonding orbitals is used to explain the different reactivity of the three surfaces. TiO2 is a chemically stable substrate with a ∼3.2 eV band gap and can induce photoreactions for a wide range of organic molecules.9,10 The photoreactions catalyzed by TiO2 is due to electron-hole pairs generated by absorption of photons with energy higher than its band gap. It has been believed that hole, OH•, and anionic oxygen species (O2-, O3-, and O23-) are the major species responsible for initiation of photoreactions on TiO2. Previously, photoreactions of some nitrogen-containing organics on semiconductors (ZnO, ZnS, or TiO2) have been investigated.11-16 Versatile products can be produced, depending on the molecular structures of the nitrogen-containing reactants. Several examples are listed in Table 1. In the present paper, we report the results of our study on the adsorption, thermal reactions, and photochemistry of methylamine on powdered TiO2 using Fourier transformed infrared spectroscopy.

TABLE 1: Examples of Photoreactions of Nitrogen-Containing Compounds Catalyzed by Semiconductors (1) aniline, toluidine f azo compounds (ref 11) (2) N-acylamines, 5- and 6-membered lactams f imides (ref 12) (3) 4-phenylbutylamine, N-methyl-4-phenylbutylamine f N-formylation of the corresponding reactants (ref 13) (4) primary amines f dialkylamines (refs 14, 15) (5) primary aliphatic amines f N-alkylidene amines (ref 16)

Experimental Section The sample preparation of TiO2 powder supported on a tungsten fine mesh (∼6 cm2) has been described previously.17,18 In brief, TiO2 powder (Degussa P25, ∼50 m2/g, anatase 70%, rutile 30%) was dispersed in water/acetone solution to form a uniform mixture which 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 which was pumped by a 60 L/s 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 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 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. The TiO2 surface after the above treatment still possessed residual isolated hydroxyl groups.19 O2 (99.998%, Matheson), 18O2 (99 at. %, Isotec), and CH3ND2 (98 at. %, Isotec) were used as received in compressed states. CH3NH2 (40% aqueous solution) was purified by several cycles of freeze-pump-thaw and introduced to the cell as the solution was cooled to ∼3 °C to reduce the solubility of CH3NH2 in the water and the evaporation of water molecules. Pressure was monitored with a Baratron capacitance manometer and an ion gauge. In the photochemistry study, both the UV and IR beams were set 45° to the normal of the TiO2 sample. The UV light source used was a combination of a 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 UV absorption of methylamine for the wavelength used in the present study was negligible.20 Infrared spectra were obtained

10.1021/jp004285d CCC: $20.00 © 2001 American Chemical Society Published on Web 06/05/2001

Reactions of Methylamine on Powdered TiO2

J. Phys. Chem. B, Vol. 105, No. 25, 2001 5929

Figure 1. Infrared spectra of a TiO2 surface exposed to 5 Torr of CH3NH2 and then evacuated at the indicated temperatures for 1 min. All the spectra were recorded with 100 scans at 35 °C. The TiO2 sample used was 91 mg.

TABLE 2: Comparison of CH3NH2 IR frequencies (cm-1) mode assigment NH2 stretching

CH3 stretching

vapor (ref 21) solid (ref 22) on TiO2 (this work) 3427 3361

3331

2985 2961

1473 1430

2942 2887 2793 1651 1636 1492 1467 1441

1195 1130 1044

1353 1182 1172 1048

2820 NH2 deformation 1623 CH3 deformation NH2 twist CH3 wagging C-N stretching

3351 3250 3161 2970 2940 2893 2810 1607 1464 1383 1350 1191 1047 1024

with a 4 cm-1 resolution by a Bruker FTIR spectrometer with a MCT detector. The entire optical path was purged with CO2free dry air. The spectra presented here have been ratioed against a clean TiO2 spectrum providing the metal-oxide background and the unit used is absorbance. 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 the full power. Results and Discussion Adsorption of CH3NH2. Figure 1 shows the infrared spectra of a TiO2 surface exposed to 5 Torr of CH3NH2 and then evacuated at the indicated temperatures. In the 35 °C spectrum, absorptions appear at 1024, 1047, 1191, 1350, 1381, 1464, 1607, 2810, 2893, 2940, 2970, 3161, 3250, and 3351 cm-1. As shown in Table 2 comparing the absorption frequencies of CH3NH2 in gas and solid phases21,22 as well as on TiO2 in this study, it is found that the frequencies are similar and are due to vibrational modes of the stretching and bending of CH3, NH2, and C-N functional groups, suggesting that most of CH3NH2 are adsorbed molecularly on TiO2 at 35 °C. Other indirect evidence from

Figure 2. Infrared spectra taken at 35, 125, 260, and 340 °C (a-d) in the course of heating a TiO2 surface in 5 Torr of CH3NH2 in a closed cell. The heating process was controlled at a rate of 2 °C/s from 35 °C to 400 °C and hold at 400 °C for 15 s. Spectrum (e) was taken at 35 °C after the heating process. All the spectra were recorded with five scans. The TiO2 sample used was 50 mg. All the traces in the region 2250-2400 cm-1 have been multiplied by a factor of 3.

our previous studies of CH3OH,23,24 CH3I,25 NH3,26 and N2H426 molecules also supports the intact state of majority CH3NH2 on TiO2. First of all, if the C-N bond in CH3NH2 breaks to generate the moieties of CH3 and NH2 groups, the former is expected to become adsorbed methoxy (CH3O(a)), which exhibits two strong CH3 symmetric and antisymmetric stretching at ∼2830 and ∼2930 cm-1 as observed in the decomposition of CH3I on TiO2 by C-I scission.25 On the other hand, the NH2 moiety is expected to react with surface hydroxyl groups to form adsorbed NH3, which shows strong absorptions at ∼1140 and ∼3395 cm-1 as observed in the decomposition of N2H4 on TiO2 by N-N bond scission.26 However, no such peaks in the 35 °C spectrum in Figure 1 are found. Second, the breakage of the C-H bond in CH3NH2 is unlikely to occur at 35 °C, because this bond dissociation process is not observed in the case of methanol on TiO2 at this temperature23,24 and not suggested by theoretic calculations.27 Third, it is found that NH3 is basically adsorbed molecularly on TiO2 at room temperature through surface Lewis acid sites and hydroxyl groups.28-33 In addition, on the oxygen-preadsorbed W surface, W(100)-(2 × 1)-O, CH3NH2 also shows low reactivity in contrast to the clean W(100) surface.8 Figure 1 also shows the thermal stability of adsorbed CH3NH2 on TiO2. Increasing the temperature under vacuum causes the decrease of peak intensities of CH3NH2 characteristic absorptions and slight shift in peak position, however no new surface absorption peak is found, suggesting that CH3NH2 is desorbed or that if CH3NH2 does decompose with increasing temperature, its reaction products are not stable on the surface. Thermal Reactions of CH3NH2 in the Absence and Presence of O2. Figure 2a-d shows the spectra taken in the course of heating TiO2 initially in 5 Torr of CH3NH2 in a closed cell. For the 35 °C spectrum before heating, in addition to the absorption peaks from adsorbed CH3NH2, those from CH3NH2 in the gas phase appear at 1043, 2819, 2916, and 2962 cm-1. As observed in Figure 1, surface CH3NH2 decreases with the increasing temperature and its peaks are no more observed by

5930 J. Phys. Chem. B, Vol. 105, No. 25, 2001

Figure 3. Infrared spectra taken at 35, 125, 260, and 340 °C (a-d) in the course of heating a TiO2 surface in a mixture of CH3NH2 and O2 in a closed cell. The heating process was controlled at a rate of 2 °C/s from 35 °C to 400 °C and hold at 400 °C for 15 s. Spectrum (e) was taken at 35 °C after the heating process. All the spectra were recorded with 5 scans. The TiO2 sample used was 50 mg. All the traces in the region 2250-2400 cm-1 have been multiplied by a factor of 3.

340 °C in Figure 2. During the heating process up to 400 °C for 15 s, no new peaks appear. To further check the reaction products that are possibly formed due to CH3NH2 decomposition at high temperatures, but not detected, an infrared spectrum was taken at 35 °C after the surface annealing, because the products might be adsorbed on the surface at lower TiO2 temperature. However, as shown in the last spectrum in Figure 2, no peaks other than those of CH3NH2 are observed. This maybe indicate that no reactions occur for CH3NH2 on the surface or that the concentrations of the products, if CH3NH2 does decompose, are too low to be detected by our IR spectrometer. In addition, it is found that the absorption intensities of the characteristic peaks of adsorbed CH3NH2 are only partially recovered in Figure 2e even in the presence of gas-phase CH3NH2. This is in part due to a less adsorption amount of CH3NH2 on the heated surface. Because after annealing TiO2 in a vacuum at higher temperatures, the surface is reduced with less Lewis acid sites (Ti4+) and surface hydroxyl groups to which the CH3NH2 is bonded. To study the O2 effect on CH3NH2 decomposition, 10 Torr of O2 was introduced to the cell at 35 °C after the heating process for Figure 2 and the TiO2 temperature was linearly increased. Figure 3 shows the spectra taken in the course of heating TiO2 in the presence of CH3NH2 and O2 in a closed cell. Unlike the case without O2, substantial CH3NH2 is consumed and new peaks appear at 1560, 1653, 2349, and 3335 cm-1 shown in the 340 °C spectrum. They are assigned to H2O (1560, 1653 cm-1), CO2 (2349 cm-1), and NH3 (3335 cm-1) produced in the gas phase.34 In the heating process up to 400 °C for 15 s, a spectrum taken at this temperature (not shown) shows no other IR bands, except H2O, CO2 and NH3. The last 35 °C spectrum in Figure 3, taken after the surface annealing, shows that CH3NH2 is totally depleted and, in addition to H2O, CO2, and NH3, new peaks appear at 1159, 1270, 1313, 1360, 1394, 1420, 1458, 1576, 1666, 2806, 2857, 2949, and 3257 cm-1. These bands are stable under vacuum. Another comparative experiment shows that if the surface temperature is holding

Liao et al. at 400 °C for 15 s followed by evacuation at this temperature for few 10 seconds and in the subsequent cooling to 35 °C to remove gaseous species present in the reaction system, no such peaks are observed. Therefore, these peaks are attributed to surface species and are due to reaction intermediates or products that are present and not detected in the gas phase as the TiO2 temperature is high, but are condensed on the TiO2 surface at 35 °C. Since the length of our infrared cell is only ∼10 cm, gaseous species formed in the cell may not be observable if their concentrations or infrared absorption coefficients are not large enough. Although these gaseous reaction intermediates or products are not detected using our infrared spectrometer, they may be observable by flow mode operation in conjunction with mass spectrometry. In Figure 3e, it is difficult to assign the exact compounds responsible for the complexed absorption feature; however, based on the characteristic frequencies of functional groups on TiO2, these condensed species may contain NHx (3257 cm-1), CHx (2806-2949 cm-1), CdO (1666 cm-1), NOx(a) (1270-1576 cm-1),35 HCOO(a) (1360, 1576 cm-1),24 and/or HCONH(a) (1394, 1576 cm-1).36 From these identifiable groups and the gas products of NH3, CO2, and H2O, some of thermal decomposition routes of CH3NH2 on TiO2 in the presence of O2 can be approximately inferred. The NH3 product may result from interaction of surface OH and NH2 generated from C-N bond scission of CH3NH2 or reaction intermediates containing C-NH2 group.26 The carbon in CH3NH2 can be oxidized to carbonyl, carboxylate, and finally to CO2. The nitrogen part of CH3NH2 can be dehydrogenated and oxidized to NOx compounds. The thermal chemistry of a mixture of ∼5 Torr of CH3NH2 and 10 Torr of O2 over a tungsten mesh without TiO2 coating was investigated as well. It is found that the concentration of CH3NH2 is about the same after heating the tungsten mesh to 400 °C for 15 s, although a tiny amount of CO2 is formed. Hydrogen Isotope Interchange between the ND2 Group of Adsorbed CH3ND2 and Surface OH Groups. Figure 4a shows the spectrum of a TiO2 surface exposed to 0.2 Torr of CH3ND2 and then evacuated at 50 °C. From Figure 1, it has been established that methylamine is predominantly adsorbed molecularly on TiO2 at lower temperatures. Therefore, CH3, ND2, and C-N vibrational modes of CH3ND2(a) are expected to be observed. Absorption peaks in Figure 4a indeed exhibit their existence: 1437, 1466, 2812, 2893, 2940, and 2971 cm-1 for the deformation and stretching of CH3; 1201, 2369, 2412, and 2493 cm-1 for those of ND2.21,22 However, additional peaks appear at 1351, 1573, 1654, 2732, 3251, and 3344 cm-1 in Figure 4a. Except for the 2732 cm-1, the other peaks, according to Table 2, can be assigned to the NH2 stretching (3251, 3344 cm-1), NH2 deformation (1573, 1654 cm-1), and NH2 twist (1351 cm-1) of adsorbed CH3NH2. The formation of CH3NH2 after CH3ND2 adsorption on TiO2 must originate from hydrogendeuterium exchange between the surface OH groups and the ND2 group in CH3ND2. The other important evidence for this interchange is the appearance of the 2732 cm-1 peak which is due to surface O-D stretching. The negative peaks in the region 3600-3800 cm-1 indicate that isolated surface OH groups are consumed in the interaction with methylamine by hydrogen bonding and H-D exchange. If the isolated surface OH groups on TiO2 are substituted by OD prior to CH3ND2 adsorption, the formation of CH3NH2 should be reduced. Figure 4b shows the expected result. In Figure 4b before CH3ND2 adsorption, the TiO2 surface was treated with several cycles of D2O adsorption and annealing in a vacuum at 250 °C. After this treatment, most of the OH groups

Reactions of Methylamine on Powdered TiO2

Figure 4. Infrared spectra of a surface exposed to 0.2 Torr of CH3ND2 35 °C and then evacuated at 50 °C. Most of the TiO2 surface OH groups have been replaced by OD before CH3ND2 adsorption for spectrum (b). The TiO2 sample used was 93 mg. All the spectra were recorded with 50 scans.

were substituted by OD groups. Figure 4b shows the effect of OD replacement cutting down the CH3NH2 formation as indicated by the much smaller peak intensities at 1351, 1573, 1654, 3251, and 3344 cm-1 in comparison to Figure 4a. Interactions of CH3ND2 with isolated surface OH and OD groups cause the negative peaks in the regions of 3600-3800 cm-1 and 2600-2800 cm-1, respectively. The absorption frequencies observed in Figure 4b from adsorbed CH3ND2 are similar to those in gas and solid phases and assigned to their corresponding modes as follows: 1143 cm-1 to ND2 twist and/ or CH3 wagging; 1205 cm-1 to ND2 deformation; 2367, 2415, and 2493 cm-1 to ND2 stretching; 2813, 2891, 2933, and 2971 cm-1 to CH3 stretching.21,22 In Figure 4b, the C-N stretching frequency of CH3ND2 on TiO2 seems to be red-shifted, lower than our detection limit of 1000 cm-1. This result is in compliance to the observation that, in solid or vapor phase, the C-N stretching frequency of CH3ND2 is smaller than that of CH3NH2 by 40-50 cm-1.21,22 Photoreactions of Methylamine. Figure 5 shows the infrared spectra taken during UV irradiation of TiO2 initially in contact with a mixture of 7.5 Torr of CH3NH2 and 10 Torr of O2 in a closed cell. Before UV irradiation, only infrared absorptions of CH3NH2 in the gas phase and on TiO2 are observed, but along with the light exposure, new absorptions appear at 2198, 2349, and 3335 cm-1 as well as in the region 1250-1850 cm-1 represented by the narrow peaks at 1560 and 1653 cm-1 located on a broad absorption feature. The peaks of 2349 and 3335 cm-1 indicate the formation of CO2(g) and NH3(g). The peaks of 1560 and 1653 cm-1 represent H2O(g). The 2198 cm-1 peak is likely due to adsorbed isocyanate species, NCO(a), which has been identified on the surfaces of metals, metal/supported oxides, and oxides.37-46 For example, isocyanate can be produced via thermal reactions of NH3/CO on Fe/Al2O343 and NO/CO on Pt/ Al2O340 as well as via photoreactions of formamide and acetonitrile on TiO2.36,46 Absorption frequency of adsorbed NCO

J. Phys. Chem. B, Vol. 105, No. 25, 2001 5931

Figure 5. Infrared spectra taken after the indicated times during UV irradiation of TiO2 initially in contact with a mixture of 7.5 Torr of CH3NH2 and 10 Torr of O2 in a closed cell (a-f). The spectrum (g) in the region 1250-1850 cm-1 was obtained by subtracting the gaseous peaks of CH3NH2, NH3, and H2O from the (f) spectrum. The UV power used was ∼0.24 W cm-2. The TiO2 sample used was ∼50 mg. All the spectra were recorded with five scans.

falls in the range of 2170-2320 cm-1 for most cases. Especially, the 2198 cm-1 peak in Figure 5 is very close to the previous observation of NCO(a) on TiO2 at 2200 and 2204 cm-1.36,46 The 2198 cm-1 peak is small and disappears for prolonged UV irradiation in Figure 5. Its existence in the methylamine photoreactions on TiO2 will be further supported later by using a lower UV power which can reduce the rate for NCO(a) photodissociation and by using 18O2 to show the NC18O(a) redshift in frequency. In Figure 5, after subtraction of the peaks of H2O(g), NH3(g), and unreacted CH3NH2(g) in the region of 1250-1850 cm-1 from the 180 min spectrum, additional peaks are more clearly observed at 1360, 1385, 1594, 1645, 1684, and 1735 cm-1 in Figure 5g. These absorptions are resemble to those for formamide adsorption on TiO2, which generates adsorbed HCONH2 with peaks at 1310, 1382, 1590, 1675, and 1698 cm-1 as well as η2(N,O)-HCONH with peaks at 1355 and 1568 cm-1.36 A comparative experiment of UV irradiation of TiO2 only in the presence of CH3NH2 has been carried out; however, no photocatalytic reactions were observed. Bare tungsten substrate shows no photocatalytic chemistry for CH3NH2 either. Since the TiO2 temperature was raised to ∼85 °C during the CH3NH2 photoreaction process in Figure 5, a thermal control experiment was carried out by holding the TiO2 temperature at 85 °C in a mixture of 7.5 Torr of CH3NH2 and 10 Torr of O2 for 180 min and taking the infrared measurement during the annealing. It was found that after 180 min annealing at 85 °C, the CH3NH2 amount was about the same, but with slightly enhanced absorptions at 1353 and 1685 cm-1. Therefore, it is concluded that the reaction products of NH3, CO2, H2O, NCO(a) result from photoeffect, instead of heating effect. To further support the formation of NCO(a) species in the photoreaction of CH3NH2 on TiO2, adsorbed CH3NH2 was irradiated at a much lower UV power, only 1/6 of that used in Figure 5, that would reduce NCO(a) dissociation rate once it was formed. Figure 6 shows the infrared spectra taken during UV irradiation CH3NH2-adsorbed TiO2 in 10 Torr of O2. The light exposure induce enhanced absorptions at 1638, 1685, and

5932 J. Phys. Chem. B, Vol. 105, No. 25, 2001

Figure 6. Infrared spectra taken after the indicated times during UV irradiation of adsorbed CH3NH2 in 10 Torr of O2 in a closed cell. The adsorbed CH3NH2 was generated by exposing ∼2 Torr of CH3NH2 to a clean TiO2 surface at 35 °C and heating the surface to 175 °C in a vacuum. The TiO2 sample used was 91 mg. All the spectra were recorded with five scans. The UV power used was ∼0.04 W cm-2.

Figure 7. Infrared spectra taken after the indicated times during UV irradiation of adsorbed CH3ND2 in 10 Torr of O2 in a closed cell. The adsorbed CH3ND2 was generated by exposing ∼0.2 Torr of CH3ND2 to a clean TiO2 surface at 35 °C and heating the surface to 175 °C in a vacuum. Before CH3ND2 adsorption, most of surface OH groups had been replaced by OD. The UV power used was ∼0.04 W cm-2. The TiO2 sample used was 93 mg. All the spectra were recorded with five scans.

2196 cm-1. The last peak demonstrates that CH3NH2 is photocatalyzed to NCO(a) which accumulates on the surface due to the much lower UV power used. The 1638, 1685 cm-1 peaks indicate that the carbon of CH3NH2 is photooxidized to carbonyl. Figure 7 shows the infrared spectra taken during UV irradiation of CH3ND2(a) in 10 Torr of O2. Most of the surface OH groups have been replaced by OD groups prior to CH3ND2 exposure. After 2 min illumination, a peak at 2194 cm-1 appears, indicating NCO(a) production. Along with longer light exposure, this peak splits into 2183 cm-1 and 2208 cm-1 as shown after

Liao et al.

Figure 8. Infrared spectra taken after the indicated times during UV irradiation of adsorbed CH3ND2 in 10 Torr of 18O2 in a closed cell. The adsorbed CH3ND2 was generated by exposing ∼ 0.2 Torr of CH3ND2 to a clean TiO2 surface at 35 °C and heating the surface to 175 °C in a vacuum. Before CH3ND2 adsorption, most of surface OH groups had been replaced by OD. The UV power used was ∼0.04 W cm-2. The TiO2 sample used was 93 mg. All the spectra were recorded with 5 scans.

180 min illumination. Meanwhile, other peaks at 1355, 1382, 1407, 1565, 1641, and 1689 cm-1 are observed. The last two peaks are assigned to species containing carbonyl groups. The positions of the peaks at 1355, 1382, and 1565 cm-1 suggest the formation of HCOO(a). Figure 8 shows the infrared spectra taken during UV irradiation of CH3ND2(a) in 10 Torr of 18O2. Similarly, most of the surface OH groups have been replaced by OD before CH3ND2 adsorption. After 2 min light exposure, a peak appears at 2184 cm-1 which is 10 cm-1 lower than the corresponding peak using 16O2 in Figure 7. This difference is consistent with the previous observation that the frequency of NC16O(a) is lower than that of NC18O(a) by 8-10 cm-1 in the photooxidation of CH3CN on TiO2 using 16O2 and 18O2.46 For longer illumination in Figure 8, the 2184 cm-1 peak splits into 2168, 2184, and 2205 cm-1. Note the 2168 cm-1 peak is not observed in Figure 7 and is assigned to NC18O as well. In addition to the NCO(a) peak, other peaks appear at 1349, 1407, 1557, 1633, and 1687 cm-1 after 180 min illumination. On the basis of the results of Figures 5-8, the infrared spectroscopy evidences the formation of NH3(g), CO2(g), H2O(g), NCO(a), and HCOO(a) in the photocatalyzed reaction of CH3NH2 on the TiO2 in the presence of O2. In the previous studies of amine photoreactions on semiconductors in aqueous or nonaqueous (CH3CN) solutions, it has been proposed that the decomposition of amine is initiated by photohole capture.14,16,47 In the present case of CH3NH2, the CH3NH2+ after hole addition to CH3NH2 may lose H+ to form •CH2NH2. This R-aminoalkyl radical can donate an electron to TiO2 conduction band, producing CH2dNH2+. A similar charge-transfer process has been observed in the electrochemical oxidation of primary aliphatic amines.48 The CH2dNH2+ may react with a CH3NH2 molecule to generate NH3 and CH2dN-CH3 or with a H2O molecule to generate NH3 and CH2dO.16,47 The presence of CH2dN-CH3 and CH2dO in this study is difficult to identify, because they are both photoactive13,16,49 and their infrared absorptions may overlap with or be buried in those from other products. On the other hand, in the photooxidation of CH3NH2

Reactions of Methylamine on Powdered TiO2 SCHEME 1

J. Phys. Chem. B, Vol. 105, No. 25, 2001 5933 Thermal reactivity of TiO2 toward CH3NH2 decomposition is much higher in the presence of O2 than without it. In the photodecomposition of methylamine on TiO2, O2 is essential, leading to the formation of NH3(g), CO2(g), H2O(g), NCO(a), and HCOO(a). Acknowledgment. This research was supported by the National Science Council of the Republic of China under contract NSC-89-2113-M-006-029. References and Notes

forming NCO(a), it is proposed that formamide is produced as an intermediate based on the following four reasons: (1) similar infrared absorption peaks in Figure 5 to those of formamide adsorption on TiO2 are observed; (2) our previous study has shown that photooxidation of HCONH2 produces NCO(a) and HCOO(a).36 NCO(a) is due to dehydrogenation of formamide, i.e., the O atom in HCONH2 is retained in the NCO(a) photoproduct; (3) from Figures 7 and 8, NCO(a) is formed in the initial stage of CH3NH2 photoreaction and the O atom of O2 is incorporated in NCO(a). This result suggests the O in NCO(a) is directly from O2 instead of TiO2 lattice oxygen. If the O of NCO(a) is from TiO2 oxygen, there should be a 2194 cm-1 peak for NC16O(a) in the 2 or 5 min spectrum of Figure 8 using 18O2. However this is not the case. Formation of formamide is a plausible reaction pathway in CH3NH2 photoreaction for the oxygen atom of O2 to incorporate in NCO(a) product. Note that for prolonged UV irradiation it is also possible to form NC16O(a) in Figure 8, because previous studies of TiO2 UV irradiation in 18O2 has shown that the oxygen atoms in 18O2 and Ti16O2 lattice exchange, producing 16O2 and 16O18O;50 (4) oxidation to carbonyl of the R-carbon in the photocatalytic oxidation of 5- and 6-membered lactams and N-acylamines has been observed in aqueous suspensions of TiO2.12 The photoreaction from methylamine to formamide on TiO2 may be through Russell-like mechanism as shown in Scheme 1. In this mechanism, the hole capture of methylamine forms a R-aminoalkyl radical of NH2CH2• which then incorporates O2 to produce organoperoxy radical, NH2CH2O2•. R-amino radicals recombine with O2 have been proposed in the photocatalytic reactions of 4-phenylbutylamine and N-methyl-4-phenylbutylamine on TiO2.16 This NH2CH2O2• radical can react with •OOH, which is generated by recombination of H+ and O2-, to form NH2CH2OOOOH tetraoxide species. Formamide and water are produced by the decomposition of the tetraoxide. In the radiolysis of aqueous solutions of organic compounds in the presence of O2, tetraoxides have been widely studied and proposed to be reaction intermediates. Recently, organoperoxys and tetraoxides are assumed as well in order to balance the photocatalytic reaction equations of C8 organic films on TiO2 in aqueous solutions.51 Similar mechanisms have been invoked in the studies of photooxidation of alkoxy,23,52 methanol,53 and chlorinated ethanes54 on TiO2. Conclusion Methylamine is molecularly adsorbed on TiO2 at 35 °C and its NH2 hydrogens exchange with surface hydroxyl ones.

(1) Baca, A. G.; Schulz, M. A.; Shirley, D. A. J. Chem. Phys. 1985, 83, 6001. (2) Maseri, F.; Peremans, A.; Darville, J.; Gilles, J. M. J. Electron. Spectrosc. Relat. Phenom. 1990, 54/55, 1059. (3) Weinberg, W. H.; Johnson, D. F.; Wang, Y.-Q. Surf. Sci. Lett. 1990, 235, L299. (4) Sasaki, T.; Aruga, T.; Kuroda, H.; Iwasawa, Y. Surf. Sci. Lett. 1991, 249, L347. (5) Hwang, S. Y.; Kong, A. C. F.; Schmidt, L. D. J. Phys. Chem. 1989, 93, 8327. (6) Chorkendorff, I.; Russell, J. N., Jr.; Yates, J. T., Jr. J. Chem. Phys. 1987, 86, 4692. (7) Bafrali, R.; Bell, A. T. Surf. Sci. 1994, 316, 267. (8) Pearlstine, K. A.; Friend, C. M. J. Am. Chem. Soc. 1986, 108, 5842. (9) Fox, M. A.; Dulay, M. T. Chem. ReV. 1993, 93, 341. (10) Hoffmann, M. R.; Martin, S. T.; Choi, W.; Bahnemann, D. W. Chem. ReV. 1995, 95, 69. (11) Kasturirangan, H.; Ramakrishnan, V.; Kuriacose, J. C. J. Catal. 1981, 69, 216. (12) Pavlik, J. W.; Tantayanon, S. J. Am. Chem. Soc. 1981, 103, 6755. (13) Fox, M. A.; Chen, M.-J. J. Am. Chem. Soc. 1983, 105, 4497. (14) Nishimoto, S.; Ohtani, B.; Yoshikawa, T.; Kagiya, T. J. Am. Chem. Soc. 1983, 105, 7180. (15) Yanagida, S.; Kizumoto, H.; Ishimaru, Y.; Pac, C.; Sakurai, H. Chem. Lett. 1985, 141. (16) Fox, M. A.; Younathan, J. N. Tetrahedron 1986, 42, 6285. (17) Basu, P.; Ballinger, T. H.; Yates, J. T., Jr. ReV. Sci. Instrum. 1988, 59, 1321. (18) Wong, J. C. S.; Linsebigler, A.; Lu, G.; Fan, J.; Yates, J. T., Jr. J. Phys. Chem. 1995, 99, 335. (19) Suda, Y.; Morimoto, T.; Nagao, M. Langmuir 1987, 3, 99. (20) Calvert, J. G.; Pitts, J. N., Jr. Photochemistry; John Willey & Sons: New York, 1966. (21) Gray, A. P.; Lord, R. C. J. Chem. Phys. 1957, 26, 690. (22) Durig, J. R.; Bush, S. F.; Baglin, F. G. J. Chem. Phys. 1968, 49, 2106. (23) Chuang, C.-C.; Chen, C.-C.; Lin, J.-L. J. Phys. Chem. B 1999, 103, 2439. (24) Chuang, C.-C.; Wu, W.-C.; Huang, M.-C.; Huang, I.-C.; Lin, J.-L. J. Catal. 1999, 185, 423. (25) Su, C.; Yeh, J.-C.; Chen, C.-C.; Lin, J.-C.; Lin, J.-L. J. Catal. 2000, 194, 45. (26) Chuang, C, -C.; Shiu, J.-S.; Lin, J.-L. Phys. Chem. Chem. Phys. 2000, 2, 2629. (27) Bates, S. P.; Gillan, M. J. J. Phys. Chem. B 1998, 102, 2017. (28) Parfitt, G. D.; Ramsbotham, J.; Rochester, C. H. Trans. Faraday Soc. 1971, 67, 841. (29) Ramis, G.; Forzatti, P. Appl. Catal. 1990, 64, 243. (30) Went, G. T.; Bell, A. T. Catal. Lett. 1991, 11, 111. (31) Went, G. T.; Leu, L.-J.; Lombardo, S. J.; Bell, A. T. J. Phys. Chem. 1992, 96, 2235. (32) Pittman, R. M.; Bell, A. T. Catal. Lett. 1994, 24, 1. (33) Markovits, A.; Ahdjoudj, J.; Minot, C. Surf. Sci. 1996, 365, 649. (34) The Aldrich Library of FT-IR Spectra, 1st ed.; Aldrich Chemical Co., Inc.: Milwaukee, 1985. (35) Hadjiivanov, K.; Bushev, V.; Kantcheva, M.; Klissurski, D. Langmuir. 1994, 10, 464. (36) Wu, W.-C.; Liao, L. -F,; Chuang, C.-C.; Lin, J.-L. J. Catal. 2000, 195, 416. (37) Celio, H.; Mudaligek, K.; Millis, P.; Trenary, M. Surf. Sci. 1997, 394, L168. (38) Solymosi, F.; Bansagl, T. J. Am. Chem. Soc. 1979, 83, 552. (39) Morrow, B. A.; Sont, W. N.; Onge, A. S. J. Catal. 1980, 62, 304. (40) Solymosi, F.; Vo¨lygesi, L.; Sa´rka´ny, J. J. Catal. 1978, 54, 336. (41) Kameoka, S.; Chafik, T.; Ukisu, Y.; Miyadera, T. Catal. Lett. 1998, 51, 11.

5934 J. Phys. Chem. B, Vol. 105, No. 25, 2001 (42) Kameoka, S.; Chafik, T.; Ukisu, Y.; Miyadera, T. Catal. Lett. 1998, 55, 211. (43) Johnston, C.; Jorgensen, N.; Rochester, C. H. J. Chem. Soc., Faraday Trans. 1 1989, 85, 1111. (44) Paul, D. K.; Worley, S. D. J. Phys. Chem. 1990, 94, 8956. (45) Solymosi, F.; Sa´rka´ny, J.; Schauer, A. J. Catal. 1977, 46, 297. (46) Zhuang, J.; Rusu, C. N.; Yates, J. T., Jr. J. Phys. Chem. B 1999, 103, 6957. (47) Low, G.; Mcevoy, S. R.; Mathews, R. W. EnViron. Sci. Technol. 1991, 25, 460. (48) Barnes, K. K.; Mann, C. K. J. Org. Chem. 1967, 32, 1474.

Liao et al. (49) Muggli, D. S.; Lowery, K. H.; Falconer, J. L. J. Catal. 1998, 180, 111. (50) Sato, S. J. Phys. Chem. 1987, 91, 2895. (51) Schwitzgebel, J.; Ekerdt, J.; Gerischer, H.; Heller, A. J. Am. Chem. Soc. 1995, 99, 5633. (52) Wu, W.-C.; Chuang, C.-C.; Lin, J.-L. J. Phys. Chem. B 2000, 104, 8719. (53) Sadeghi, M.; Liu, W.; Zhang, T.-G.; Stavropoulos, P.; Levy, B. J. Phys. Chem. 1996, 100, 19466. (54) Mao, Y.; Scho¨neich, C.; Asmus, K.-D. J. Phys. Chem. 1991, 95, 10080.