217
Ind. Eng. Chem. Res. 1991,30, 217-221
The uptake behavior of several low molecular weight penetrants was also examined to determine the effect of molecular size and the penetrant’s affinity for the resin network on the uptake behavior. While the equilibrium uptake of ethanol, methyl ethyl ketone, and acetone was less than 2.5 g of penetrant/100 g of resin, the resin samples absorbed a significant amount of methanol (approximately 17.0 g/100 g of resin). The dynamic uptake behavior observed for methanol-exposedsamples was anomalous. The inflection observed in these data was significantly more dramatic than observed in the water uptake behavior. The desorption of methanol from methanol-exposed samples indicated that the resin’s volume collapses back to its original unswollen dimensions. Furthermore, the loss of mass signifies that impurities, such as low molecular weight oligomers, have been removed from the original network. The methanol uptake behavior could be correlated with the dynamic changes in the sample’s dimensions. This would indicate whether the rapid change in uptake behavior corresponds with the dramatic inception of volumetric expansion.
Acknowledgment This work was supported in part by a grant from the National Science Foundation (CBT-86-17719). Registry No. (TGDDM)(DDS) (copolymer), 63804-34-2;water, 7732-18-5;acetone, 67-64-1;ethanol, 64-17-5; methanol, 67-56-1; methyl ethyl ketone, 78-93-3.
Literature Cited Antoon, M. K.; Koenig, J. L.; Serafini, T. Fourier-Transform Infrared Study of the Reversible Interaction of Water and A Crosslinked Epoxy Matrix. J . Polym. Sci., Polym. Phys. Ed. 1981, 19, 1567-1575.
Berens, A. R.; Hopfenberg, H. B. Diffusion and Relaxation in Glassy Polymer Powders: 2. Separation of Diffusion and Relaxation Parameters. Polymer 1978,19,489-496. Berens, A. R.; Hopfenberg, H. B. Induction and Measurement of Glassy-State Relaxations by Vapor Sorption Techniques. J . Polym. Sci., Polym. Phys. 1979,17, 1757-1770. Crank, J. The Mathematics of. Diffusion: Oxford University Press: .. London, 1959. Joshi, S.; Astarita, G. Diffusion-Relaxation Coupling in Polymers which Show Two-Stage - Sorption Phenomena. Polymer 1979.20, . . 455-458. Levy, R. L.; Fanter, D. L.; Summers, C. J. Spectroscopic Evidence for Mechanochemical Effects of Moisture in Epoxy Resins. J. Appl. Polym. Sci. 1979,24,1643-1664. May, C. A. Introduction to Epoxy Resins. Epoxy Resins: Chemistry and Technology, 2nd ed.; May, C. A., Ed.; Dekker: New York, 1988. Mijovic, J.; Weinstein, S. A. Moisture Diffusion into a GraphiteEpoxy Composite. Polym. Commun. 1985,26,237-239. Mikols, W. J.; Seferis, J. C.; Apicella, A.; Nicolais, L. Evaluation of Structural Changes in Epoxy Systems by Moisture Sorption-Desorption and Dynamic Mechanical Studies. Polym. Comp. 1982, 3, 118-124. Moy, P.; Karasz, F. E. Epoxy-Water Interactions. Polym. Eng. Sci. 1980,20,315-319. Netravali, A. N.; Fornes, R. E.; Gilbert, R. D.; Memory, J. D. Investigations of Water and High Energy Radiation Interactions in an Epoxy. J . Appl. Polym. Sci. 1984,29,311-318. Netravali, A. N.; Fomes, R. E.; Gilbert, R. D.; Memory, J. D. Effects of Water Sorption at Different Temperatures on Permanent Changes in an Epoxy. J. Appl. Polym. Sci. 1985,30,1573-1578. Ritger, P. L.;Peppas, N. A. Transport of Penetrants in the Macromolecular structure of Coals, 4. Models for Analysis of Dynamic Penetrant Transport. Fuel 1987,66,815-826. Tsou, A.; Peppas, N. A. Transport Properties of Water in Epoxy Resins and Composites. Polym. Mater. Sci. Eng. Proc. 1988,58, 952-956. Wong, T. C.; Broutman, L. J. Water in Epoxy Resins Part 11. Diffusion Mechanism. Polym. Eng. Sci. 1985,25,529-534.
Received for review December 5, 1989 Revised manuscript received July 17,1990 Accepted August 1, 1990
Decomposition of Ozone on a Silver Catalyst Seiichiro Imamura* and Masaaki Ikebata Department of Chemistry, Kyoto Institute of Technology, Matsugasaki, Sakyo-ku, Kyoto 606,Japan
Tomoyasu Ito Department of Chemistry, Faculty of Science, Tokyo Metropolitan University, Setagaya-ku, Tokyo 158, Japan
Takashi Ogita Matsui Kagaku Co. Ltd., Jodori-cho, Fushimi-ku, Kyoto 612,Japan
Decomposition of ozone was carried out on metal oxide catalysts. The activity of the metal oxide catalysts increased roughly in the order of the increase in their surface area and in the amount of surface oxygen on them. Conductance change of these metal oxides on an introduction of ozone suggested that negatively charged oxygen species were formed on their surface. The Ag catalyst showed the highest activity, and the reactivity of the oxygen species produced on the surface of Ag catalyst toward carbon monoxide was much higher than that of the oxygen species on Co, Ni, Fe, and Mn oxides. Formation of superoxide ions and their precursors on the Ag catalyst was suggested by ESR analysis and by an activity measurement of these oxygen species. The latter, probably oxygen ion (0-),seemed to be an active species for low-temperature oxidation of CO. Ozone, a powerful oxidizing agent, is useful for sterilization (Tsuruta, 1988),deodorization (Stevens and Brown, 198% and destruction of pollutants in wastewaters (Gould and Weber, 1976; Teramoto et al., 1981). However, as ozone itself is toxic to organisms, its release into the en-
vironment must be avoided, The most commonly used technique of ozone detoxification is adsorption and decomposition on activated charcoal (Emel’yanova and Atyaksheva, 1979). Although this method is simple and convenient, activated charcoal loses its activity after pro-
0888-5885/91/2630-0217$02.50/0 Q 1991 American Chemical Society
218 Ind. Eng. Chem. Res., Vol. 30, No. 1, 1991
longed use. Ozone is decomposed also on inorganic compounds such as natural sand, which contains SiO,, FezO3, CaO, MgO and other minerals (Suzuki et al., 1979). Various metal oxides have been used for ozone detoxification (Calderbank and Lewis, 1976; Rovero et al., 1983). However, the mechanism of the decomposition of ozone on these metal oxides is scarcely known (Koga et al., 1987). In this work the activity of metal oxide catalysts was examined, and the catalytic action of silver was investigated in detail.
Experimental Section Catalysts. Commercial Moo3, Sn02, and ZnO were calcined at 400 "C in air for 2 h. The V205was obtained by calcining NH4V03a t 400 "C in oxygen. NiO, Co304, CuO, Pb203,Mn203,Fez03,CeO,, and Ag20were prepared from the corresponding nitrates according to the following procedure. Aqueous sodium hydroxide (1 N) was added to an aqueous solution of metal nitrates under magnetic agitation until the pH of the solution was 10-11 (in the case of Pb, 9). The resultant precipitate was washed with deionized water until the pH of the solution was about 8. The precipitate was dried a t 100 "C overnight, followed by calcination a t 400 "C in air for 2 h except for AgzO; Ag,O was calcined at 200 "C in oxygen for 2 h. Biz03was prepared in the same way as above except that bismuthfI1) nitrate was dissolved in a small amount of nitric acid. The form of these oxides was identified as shown above by X-ray analyses except for lead oxide and vanadium oxide. They were deduced as Vz05 and Pb2O3, respectively, judging from the calcination condition. The Ag catalyst exchanged on Na-Y zeolite was prepared as follows. A 50-g sample of Na-Y zeolite provided by Toso Co. Ltd. was dispersed in 500 mL of 2 M silver(1) nitrate solution, and the solution was stirred at 70 "C for 24 h in the dark. This Ag-exchanged zeolite was washed with deionized water several times and was dried at 100 "C overnight. The content of Ag was 30.0 wt %. It was again dispersed in deionized water, and aqueous sodium hydroxide (3 N) was added until the pH of the solution was 10 to produce silver hydroxide inside the cavity of the zeolite. This Ag-containing zeolite was washed with deionized water several times and was dried at 100 "C overnight, followed by calcination in oxygen at 200 "C for 2 h. The form of Ag was identified as Ag,O by an X-ray diffraction analysis. This catalyst was designated as AgZ. Ni was exchanged into Na-Y zeolite from nickel(I1) nitrate according to the same procedure as described above. This Ni-exchanged zeolite (6 w t 70)was treated with aqueous sodium hydroxide (3 N) to produce nickel hydroxide and was heated a t 250 "C in oxygen for 2 h. This catalyst was designated as NiZ, and the form of Ni was identified as NiO by an X-ray diffraction analysis. The catalysts were mixed thoroughly with quartz sand with 1mol 70content as metal ion. The molecular weight of quartz sand was assumed to be 60. The mixture was molded into a disk under pressure and was cut into 814-mesh size before use. Apparatus and Procedure. Reactions were carried out with the use of an ordinary flow reactor under atmospheric pressure. The catalysts were heated at 150 OC under a flow of nitrogen (150 mL/min) for 30 min before the reaction to eliminate adsorbed water. Ozonized oxygen (1L,/min) was produced by a Nippon Ozone 0-3-2 ozone generator from oxygen purified by passing through a molecular sieve column, and a part of it was introduced into the reactor a t a rate of 150 mL/min. Reaction was followed by determining the remaining ozone a t an exit of the reactor. The temperature of the reactor was controlled by circu-
lating ice water through a jacket outside the reactor (0 "C), by heating the reactor with an electric furnace (>25 "C), or by immersing the reactor into a dry ice-acetone bath (CO "C). Oxidation of carbon monoxide (CO) during the decomposition of ozone was carried out with the same reactor described above. A mixture of CO (10%) and nitrogen was diluted with ozonized oxygen to 1.0%, and this reaction mixture was introduced into the catalyst bed at a rate of 150 mL/min. Temperature-programmed desorption (TPD) of oxygen from the metal oxide catalysts was carried out in a quartz reactor connected to an ordinary glass vacuum line. The metal oxides without silica gel (0.5 g) were heated in oxygen at 400 "C for 1 h except for Ag,O; Ag,O was heated at 200 "C. After cooling to room temperature in oxygen, the gaseous phase was evacuated for 1 h. Helium was introduced into the reactor, and the temperature of the reactor was increased a t a rate of 4 "C/min. The amount of desorbed oxygen was determined by a thermal conductivity detector and a Varian TE-150 quadrupole mass spectrometer. Calibration was made on the basis of the amount of desorbed oxygen from pure Ag,O in the temperature range from 200 to 500 "C. Conductance change of the metal oxides on introducing ozone was measured as follows. Finely powdered metal oxides were mixed with a small amount of water and were spread on a ceramic tube (diameter of 3 mm) as a thin zonal layer of 2-mm width. Ag paste was spread a t both ends of the layer, and an copper wire was connected to this Ag electrode. This specimen was placed in a glass reactor and was connected to an electric circuit equipped with a variable resistance and a voltmeter. A constant voltage was supplied to the circuit, and from the change of the resistance of the specimen on an introduction of ozone, conductance change of the metal oxides was calculated. Analysis. Ozone concentration at the exit of the reactor was determined by a Shimadzu UV-100-02 spectrophotometer. A 10-cm optical cell was used, and the monitoring wavelength used was 285 nm. Carbon monoxide and carbon dioxide (CO,) were determined by a Shimadzu gas chromatograph GC-3BT with a molecular sieve 5A (1.3 m) plus Shimalite (1.7 m) column at 60 "C and a Shimalite (0.5 m) plus activated charcoal (0.5 m) column at 60 "C, respectively. The ESR, ESCA, and X-ray diffraction analyses of the catalysts were carried out by a JEOL JES-PE ESR spectrometer, a Shimadzu ESCA 750 spectrophotometer, and a Rigaku Denki Geigerflex 2012 X-ray analyzer, respectively. The BET surface area of the catalysts was measured by using a conventional gas adsorption apparatus with a glass vacuum line. Nitrogen was used as an adsorbate.
Results and Discussion Decomposition of Ozone on Metal Oxide Catalysts. Table I shows the activity of metal oxide catalysts on the decomposition of ozone as expressed by percentage ozone decomposition a t 0 "C 10 min after an introduction of ozonized oxygen and by the temperatures at which the rate of ozone decomposition reached 50% (Tm) and 95% (Tg5), respectively. In evaluating TW and Tg5,the temperature of the catalyst bed was increased a t a rate of 5 "C/min from 25 "C. As the activity of the catalysts changed during the reaction, evaluation of the catalyst activity was qualitative. The BET surface area and the relative amount of oxygen desorbed from the catalysts are also shown in the table. The activity pattern divided the metal oxides into two groups. One contained Ag20, NiO, Fep03,Co304,
Ind. Eng. Chem. Res., Vol. 30, No. 1, 1991 219 Table I. DecomDosition of Ozone on Metal Oxide Catalysta Sw,b 0,desorbed: 0,decomposed T,," T8s1( catalyst m2/g re1 amt atO°C!% OC OC AgzO 1.87 12.4 79.5 f f NiO 62.90 100 78.1 f f Fez03 35.86 60.3 37.3 f 25 Co304 28.15 62.9 29.4 f 73 CeO, 163.98 80.5 13.0 48 100 MnzOs 15.03 58.4 4.9 62 90 CUO 9.19 49.1 3.0 40 63 PbZO3 1.26 5.7 3.1 104 145 Biz03 2.91 7.5 1.0 120 160 SnO, 19.58 34.0 0 146 205 MOO, 0.81 5.0 0 155 210 V20s 6.65 27.7 0 155 220 SiO# 0 155 205 OCatalyst (1 mol % as metal ion on quartz sand) 1 mL; space velocity (SV)of ozonized oxygen 9000 h-l; [O,]3000 ppm. bBET surface area. '0,desorbed from room temperature to 300 "C; see experimental section. 10 min after introduction of 03.e Tm, temperature of 50% decomposition; TSs,temperature of 95% decomposition. Temperature increase of the catalyst bed 5 OC/min from 25 OC. IT, or TBSwas below 25 "C. #Quartz sand without catalyst.
Table 111. Oxidation of CO during Decomposition of Ona reactionb 03 catalyst temp, OC decomposed, % CO conv, % 58.9 AgzO 51 100 NiO 40 100 16.2 16.4 Co304 40 97.2 Mn203 38 40.3 0 Fez01 38 98.0 0 SlO, 25 0 0 aCatalyst (1 mol % as metal ion on quartz sand) 1 mL; [O,] 5000 ppm; [CO] 10000 ppm; [O,]89.5%; [N,] 9%; SV 12000 h-l; measured a t 20 min. Temperature increased due to the exothermic reaction. cConversion of CO without O3a t 51 OC w a 6.1%. dQuartz sand without catalyst.
Table IV. DecomDoeition of O1 at 0 O c a rate of O3 decompn, mol of O,/(h.mol of Ag or Ni) catalyst 1 min 20 min 60 min AgZ 1.61 2.97 4.44 NiZ 0.56 0.84 1.62 DCatalyst 0.069 mmol of metal in quartz sand plus zeolite Y;
[O,]3000 ppm; SV 9000 h-l. Table 11. Conductance Change on an Introduction of 02 catalyst semiconductor typeb conductance change' Sn02 n ca. 1 Bi203 ca. 1 Pb203 ca. 1 CeOz n ca. 1 MOO, n ca. 1 n ca. 1 v2°5 n 0.05d Mn203 P 5e c0301 P looe Ag,O P 5800O CUO i (n) 1200e NiO P 6OOd ~~~~
~
[O,]3000 ppm. bn, n-type; p, p-type; i, intrinsic. 'Relative to original conductance. Minimum conductance. e Maximum conductance. fMonotonic increase to 6000 after 30 min.
CeO,, Mnz03,and CuO, which had high or moderate activity. The other contained Pb203,BizO3, SnO,, Moo3, and Vz05,which were inactive or had, if any, only low activity. The former group except for Ag,O had large surface area and a large amount of surface oxygen which desorbed easily on heating to 300 "C. The surface area and the amount of surface oxygen were small for the latter group. Generally, large surface area and a large amount of easily desorbed oxygen are characteristic of combustion catalysts, while low surface area and a small amount of surface oxygen are seen for partial oxidation catalysts. Therefore, the nature of the catalytic ozone decomposition resembles that of the catalytic combustion reaction in which mobile surface oxygen is involved. Although Ag is not commonly used as a combustion catalyst due to its instability to aggregate as metallic Ag, it is remarkably active in the oxidation of CO a t low temepratures when it retains its oxidized state (Imamura et al., 1988). Table I1 shows the conductance change of the catalysts on introduction of ozone. The mode of conductance change was complex. For example, NiO showed a monotonic conductance increase which reached 6000 times the original value after 30 min, while the conductance of Ag20 reached a maximum (5800 times the original value) 3 min after an introduction of ozone and, then, decreased to 500 times the original value after prolonged contact with ozone (18 min). Although no detailed inspection of the above phenomena was carried out, the following qualitative deduction was derived. If oxygen species produced in the
decomposition of ozone abstract an electron from the catalyst and are present as an anionic state, the population of the acceptor level of p-type semiconductors will increase and, therefore, their conductance will increase. On the other hand, the conductance of n-type semiconductors will decrease as a result of the decrease in the population of the donor level (Hattori, 1966). Introduction of ozone leads to the increase in the conductance of Mnz03, Co304,AgzO, CuO, and NiO, which are p-type semiconductors except for CuO (Hattori, 1966). Koga et al. reported the same phenomenon (Koga et al., 1987). The only n-type metal oxide that exhibited a conductance decrease was Vz05. Although these results did not present conclusive evidence, we deduced that negatively charged oxygen species were formed in the decomposition of ozone on these metal oxide catalysts, judging mainly from the behavior of p-type metal oxides. In order to see the reactivity of the oxygen species produced on the surface of the catalysts, decomposition of ozone (5000 ppm) was carried out in the presence of CO (10000 ppm), and the result is shown in Table 111. Although excess CO (58.9% X 10000 ppm) over the amount of ozone decomposed (100% X 5000 ppm) was oxidized on AgzO, it was because 6.1% of CO was oxidized in the absence of ozone. Therefore, an approximately stoichiometric amount of CO corresponding to decomposed ozone was oxidized. On the other hand, decomposed ozone only partly contributed to the oxidation of CO (NiO, Co304)or was completely inactive (Mn203,Fe203). Ozone itself was inactive for CO oxidation as is shown by the reaction on quartz sand. Therefore, the nature of the oxygen species formed on Ag,O and that of the oxygen species on other transition-metal oxides were different, which made AgzO an interesting object of investigation. Decomposition of Ozone on AgZ and NiZ. Ag20, NiO, and FezOs had high activity in decomposing ozone. However, as the surface area of Ag,O was the smallest among the three, its inherent activity seemed to be the highest. Therefore, Ag was exchanged on Na-Y zeolite (AgZ) in order to increase the degree of dispersion, and its catalytic action was investigated. Ni was also exchanged on Na-Y zeolite (NiZ), and its activity was compared with that of AgZ. Table IV shows the result of ozone decomposition at 0 "C.The activity of AgZ increased during the reaction, and the X-ray diffraction analysis indicated that
220 Ind. Eng. Chem. Res., Vol. 30, No. 1, 1991
44 ,4
'
\
*
(2,
3
0
A
o
Temp ('C)
Figure 1. TPD of gaseous compounds from Ag20 "02or C02a t the exit of the T P D reactor; (-) O2 desorbed from Ag20 after introduction of 5000 ppm of O3in O2 (150mL/min) a t room temperature for 60 min; (- -) O2 and (- - -) COP from Ag20 treated as above, followed by a contact with 10% CO in N2 for 10 min a t 50 O C .
-
peaks due to Ago became appreciable in addition to those for Ag20 after prolonged contact with O3 Ag changed gradually to more oxidized states. Although the activity of NiZ also increased during the reaction, it was always lower than that of AgZ. As the conversion of ozone was almost independent of the particle size of the catalysts at the same reaction time, it was assumed that the reactions were not affected by any mass-transfer limitation. Therefore, Ag seemed to have the highest activity among the catalysts investigated in this work. Effect of Water on the Activity of Ag. Exhaust gases may contain water vapor as well as residual ozone when ozone is applied under humid conditions such as that at water-treatment facilities. As water may be detrimental for the catalytic decomposition of ozone, the effect of water on the activity of the Ag catalyst (Ag20mixed with quartz sand) was investigated. The reaction condition was the same as that described in Table I except for the reaction temperature of 35 "C. A 0.65% water content had no effect on the activity of the Ag for more than 5 h; the conversion of ozone was 100%. A 4.2% water content (a saturated concentration at 30 "C)also had no influence. The reactor was kept at 0 "C for 10 min under 4.2% water vapor to allow liquid water to condense on the surface of the catalyst, and then, the temperature of the reactor was again increased to 35 "C. By this procedure, the conversion of ozone dropped to 57.6%. After the reactor was heated at 150 "C for 20 min to eliminate the surface water, the conversion of ozone recovered to 100% at 35 "C. These results shows that water is not so detrimental unless a large amount of water covers the surface of the Ag and prevents ozone from approaching it. Naturally the decomposition of ozone on any catalyst will be suppressed if the surface of the catalyst is covered with liquid water. Therefore, the decomposition of ozone on the Ag catalyst will proceed smoothly by employing a condition where liquid water does not condense on its surface. Consideration on the Surface Oxygen Species on Ag. The TPD analysis of Ag,O (without support) after contact with ozone was carried out, and desorbed gaseous species were monitored with a mass spectrometer (Figure 1). Ag20 without Ozone treatment did not show any sharp desorption peak below 300 "C. An introduction of ozone (500 ppm, 150 mL/min) for 60 min resulted in the appearance of a sharp O2desorption peak at 140 "C. As the amount of this oxygen was calculated to be about 1% of the total lattice oxygen of Ag20, it was attributed to the adsorbed oxygen on the surface of Ag20. When Ag20 treated with ozone was allowed to contact with 10% CO
c
532
6
528
Binding energy (eV)
Figure 2. ESCA spectra of (a) Ag20 and (b) Ag20 after introduction of 5000 ppm of O3in O2(150 mL/min) a t room temperature for 5 h.
-H
50G
Figure 3. ESR spectra of (a) AgZ contacted with 5 Torr of O3a t room temperature [g,, = 2.040, g:,y = 2.009, and g,, = 2.0031 and (b) AgZ treated as in (a), followed by a contact with 200 Torr of CO a t 50 O C . Both spectra were observed a t -198 "C.
in nitrogen for 10 min at 50 "C,about 40% of the surface oxygen (corresponding to the desorption peak at 140 "C) decreased and a new C02desorption peak appeared at 110 "C (only the peaks below 200 "C are shown in the figure). This indicated that the surface oxygen accumulated during the decomposition of ozone was reactive and oxidized CO. The Ag20 (without support) showed an O(1s) peak at 529.9 and 531.2 eV as observed by an ESCA spectrophotometer (Figure 2). When it was treated with 5000 ppm of ozone (150 mL/min) for 5 h, the peak at 529.9 eV became very small and a large peak was observed at 531.6 eV. In addition, a small peak was observed at 533.8 eV. The peak at 529.9 eV was assumed to be due to the lattice oxygen, 02-(Schon, 1973). The decrease in the intensity of the 529.9-eV peak, the increase in the intensity of the 531.6-eV peak, and, in addition, the appearance of a new peak at higher binding energy (533.8 eV) indicated that oxygen species less negatively charged than 02became abundant. When AgZ was evacuated at 150 "C for 30 min followed by an introduction of 5 Torr of ozone at room temperature, a broad asymmetric signal was observed by an ESR spectrometer (Figure 3). The same procedure using oxygen did not present this signal. This paramagnetic species was stable even on evacuation of the sample at 150 "C. However, an introduction of 200 Torr of CO at 50 "C resulted in the disappearance of this signal except for a small signal which was always observed for Ag20 and was assumed to result from lattice defect of silver oxides. Superoxide ion, Of, is observed on the surfaces of various inorganic materials including silver oxides (Che and Tench,
Ind. Eng. Chem. Res., Vol. 30, No. 1, 1991 221 employing Ag catalyst. If ozone is used for destruction of gaseous pollutants such as for deodorization, use of Ag catalyst may be quite effective because of the production of highly reactive oxidizing species from ozone. Acknowledgment
-900
-50
0
50
Temp (TI
Figure 4. Decomposition of O3and oxidation of CO on AgZ: AgZ = 1 mL, [O,] = 5000 ppm, [CO] = 10000 ppm, [02] = 89.5%, [N2] = 9%, and SV = 9000 h-'; O[CO oxidized]/[O, decomposed] X 100 (mol %).
1983a). It is stable at relatively high temperature (160 "C) but is reactive toward CO (Clarkson and Cirillo, 1974). The ESR signal shown in Figure 3 is considered to be due to superoxide ion, judging from its g factors: g,, = 2.040, gyy= 2.009, and g,, = 2.003 (Che and Tench, 1983a). The formation of negatively charged oxygen species (result of conductance change) whose valence states are more pos(result of ESCA analysis) also supitive than that of 02ports this deduction. Figure 4 shows the temperature dependence of the decomposition of ozone and oxidation of CO on AgZ. More than 90% of ozone was decomposed at as low as -78 "C, and a nearly stoichiometric amount of CO was oxidized to COPin the whole temperature range examined. Although a more than stoichiometric amount of CO over decomposed ozone was oxidized below -50 "C, the reason was not clarified. Superoxide ion stabilized on metal or metal oxide surface shows only moderate reactivity toward CO even above room temperature (Clarkson and Cirillo, 1974). Therefore, the oxygen species that oxidized CO at low temperatures must be different from the superoxide ion stabilized on the catalyst surface. Since the decomposition of O3 into the stabilized 0; on the catalyst surface does not seem to proceed through a simple one-step reaction, it is reasonable to presume that oxygen species e.g., oxygen atom (01,ozonide ion (03-), other than 02-, and oxygen ion (0-1,are also formed as intermediates. Especially, 0- is known to have the highest activity among various oxygen species (Che and Tench, 1983b). Therefore, 0- seemed to be the main species responsible for the lowtemperature oxidation of CO. We could not directly confirm the presence of this species by ESR spectroscopy, however; some kinds of 0- are undetectable by an ESR technique (Aika et al., 1984; Che and Tench, 1982). In conclusion, Ag was the most effective catalyst in decomposing ozone. Some transition-metal oxides are deactivated during the reaction due to accumulation of oxygen species that cover their active site (Calderbank and Lewis, 1976). By contrast, the reactivity of Ag increased as the reaction proceeded. The oxygen species produced on Ag are remarkably active in the oxidative of CO compared with those on other transition-metal oxides, and these reactive oxygen species may have reacted also with ozone and accelerated its decomposition. The high reactivity of the oxygen species on Ag is probably due to the especially weak bond between oxygen and silver compared with those between oxygen and other metals (Imamura et al., 1988). This may bring about another advantage of
We thank Dr. S. Ishida and K. Utani of the Kyoto Institute of Technology for their kind cooperation in carrying out this work. Registry No. Ag20, 20667-12-3; NiO, 1313-99-1; Fe203, 1309-37-1; Co304,1308-06-1; CeOz,1306-38-3;Mn203, 1317-34-6; CuO, 1317-38-0; Pbz03, 1314-27-8; Bi203, 1304-76-3; SnOz, 18282-10-5; Moo3, 1313-27-5; V205, 1314-62-1; Si02,7631-86-9; CO, 630-08-0; 02,7782-44-7; ozone, 10028-15-6; water, 7732-18-5.
Literature Cited Aika, K.; Tajima, M.; Isobe, M.; Onishi, T. Surface Reaction of 0Ions with C2Hs and the Relation to a Catalytic Reaction between C2Hs and N20 over Co-MgO. Proceedings of the 8th International Congress on Catalysis, Berlin, III; Verlag Chemie: Weinheim, 1984; p 335. Calderbank, P. H.; Lewis, J. M. 0. Ozone-Decomposition Catalysis. Chem. Eng. Sci. 1976, 31, 1216. Che, M.; Tench, A. J. Characterization and Reactivity of Mononuclear Oxygen Species on Oxide Surfaces. In Advances in Catalysis; Eley, D. D., Pines, H., Weisz, P. B., Eds.; Academic Press: New York, 1982; Vol. 31, Chapter 2, pp 77-78. Che, M.; Tench, A. J. Characterization and Reactivity of Molecular Oxygen Species on Oxide Surfaces. In Advances in Catalysis; Eley, D. D., Pines, H., Weisz, P. B., Eds.; Acadamic Press: New York, 1983a; Vol. 32, Chapter 1, pp 29-29. Che, M.; Tench, A. J. Characterization and Reactivity of Molecular Oxygen Species on Oxide Surfaces. In Advances in Catalysis; Eley, D. D., Pines, H., Weisz, P. B., Eds.; Academic Press: New York, 1983b; Vol. 32, Chapter 1, pp 116-118. Clarkson, R. B.; Cirillo, A. C., Jr. The Formation and Reactivity of Oxygen as 02on a Supported Silver Surface. J . Catal. 1974,33, 392-401. Emel'yanova, G. I.; Atyaksheva, L. F. The Interaction of Activated Charcoal with Ozone at Elevated Temperatures. Russian J. Phys. Chem. 1979,53, 1588-1590. Gould, J. P.; Weber, W. J., Jr. Oxidation of Phenols by Ozone. J . Water Pollut. Control Fed. 1976,48,47-60. Hattori, K. Structure and Catalytic Action of Metal Oxides. In Shokubai Kogaku Koza; Shiba, T., Ed.; Chijin Shokan: Tokyo, 1966; Vol. 2, Chapter 2, pp 140-142. Imamura, S.; Sawada, H.; Uemura, K.; Ishida, S. Oxidation of Carbon Monoxide Catalyzed by Manganese-Silver Composite Oxides. J . Catal. 1988, 109, 198-205. Koga, 0.;Hori, Y.; Suzuki, S. Ozone Sensor with the Use of Thin Films of Transition Metal Oxides. Nippon Kagaku Kaishi 1987, 147-151. Rovero, G.; Piccinini, N.; Grace, J. R.; Epstein, N.; Brereton, C. M. H. Gas Phase Solid-Catalyzed Chemical Reaction in Spouted Beds. Chem. Eng. Sci. 1983,38, 557-566. S c h h , G. ESCA Studies of Ag, Ag20 and Ago. Acta Chem. Scand. 1973,27, 2623-2633. Stevens, R. D. S.; Brown, P. M. Method for Removing Malodorous or Toxic Gases from an Air Stream. Eur. Pat. Appl. E P 261987, 1988; Chem. Abstr. 1988, 108, 226230k. Suzuki, S.; Hori, Y.; Koga, 0. Decomposition of Ozone on Natural Sand. Bull. Chem. SOC. Jpn. 1979,52, 3103-3104. Teramoto, M.; Imamura, S.; Yatagai, N.,; Nishikawa, Y.; Teranishi, H. Kinetics of the Self-Decomposition of Ozone and the Ozonation of Cyanide Ion and Dyes in Aqueous Solutions. J . Chem. Eng. Jpn. 1981, 14, 383-388. Tsuruta, S. Apparatus for Sterilization with Ozone and Ultraviolet-Ray Radiation. Jpn. Kokai Tokkyo Koho JP 62227358,1988; Chem. Abstr. 1988, 108, 7988k. Received for review December 21, 1989 Accepted August 4, 1990
