DECOMPOSITION OF NITROUS OXIDE ON ZINC OXIDE
1037
Photocatalytic Reactions on Semiconductor Surfaces. I.
Decomposition of
Nitrous Oxide on Zinc Oxide
by Ken-ichi Tanaka1 and George Blyholder* Department of Chemistry, University of Arkansas, Fayetteville, Arkansas 78701
(Receiued June 11, 1970)
Publication costs borne completely by The Journal of Physical Chemistry
The steady-state photocatalytic decomposition of nitrous oxide was studied on zinc oxide. The reaction was At these temperatures, both thermal catalytic and carried out over the temperature range 371 to -431'. photocatalytic decomposition occur. The thermal catalytic reaction obeys a first-order rate equation, Le., r = ICPN,~;however, the photocatalytic reaction obeys a completely different rate equation r = k P N z O / ( l f ~ I P N ~kzPo2'/2). O The kinetic equation for the photocatalytic decomposition of nitrous oxide is reasonably explained by considering the electron concentration at the surface during photocatalysis. Overall reaction rate under illumination is given by the sum of the photo- and thermal reactions. It was concluded that the photocatalytic decomposition occurs by way of NzO- intermediate but that thermal catalytic decomposition probably occurs without an electron transfer process.
+
Introduction The catalytic decomposition of nitrous oxide has been used as a test reaction to investigate the catalytic action of oxide surfaces. Since Dell, Stone, and Tiley2 pointed out a relation between the catalytic activities of oxides and their electronic properties, the importance of an electron transfer process in the catalytic decomposition of nitrous oxide has been empha~ized.~It has been extensively reported that the adsorption of oxygen on semiconductor-type oxides changes the conductivity of the oxides, indicating thereby electron transfer from the oxide to adsorbed oxygen. Since the photoeffect on the conductivity of zinc oxide was shown by Myasnikov and P~hezhetskii,~ the photoeffect on adsorption and/or desorption of oxygen has received much attention. Fujita and Kwan5 and Barry and Stonea found both photoadsorption and photodesorption of oxygen for zinc oxide depended on the immediate pretreatment of the oxide. A comparison of their results indicates that the nature of the photoresponse also depends to a considerable extent on the history of the sample and its long-term pretreatment. Both groups found that the photoeffect decreased with increasing temperature and was not detectable at the temperatures above 2005 or above 30L6 It has been inferred that the oxygen adsorbed on highly evacuated oxides is not necessarily the same as that which is adsorbed as an intermediate during catalytic decomposition of NzO.' Accordingly, it is desirable to better define the behavior of intermediates and the roles of electrons and holes during catalysis at dark and a t illumination. Morrison and Freunds have recently made progress in this direction for steadystate electrode reactions under illumination, but their technique is difficult to apply to other than to liquid
phase reactions. However, the framework they are establishing to consider reactions at semiconductor surfaces is most useful. The aim of this work is to study the kinetics of the steady-state photocatalytic decomposition of nitrous oxide on zinc oxide and to show the roles of electrons and holes in photocatalysis and thermal catalysis. For this purpose, the photocatalytic decomposition was carried out in the same temperature range in which the thermal catalytic reaction was observed.
Experimental Section As shown in Figure 1, the apparatus used in this work is a closed circulating system with a total volume of 141 ml (VI = 112 ml and V2 = 29 ml). The light source, a 500-W super pressure mercury lamp (OSRAM HBO-500), and a flowing water filter (55 mm length with 3 mm thick Pyrex windows) were set as shown in Figure 1. The reactor was constructed from an optically flat Pyrex glass window (63-mm diameter and (1) On leave from Tokyo Institute of Technology. (2) R. M . Dell, F. S. Stone, and P. F. Tiley, Trans. Faraday SOC.,49,
201 (1953). (3) E. R. 8. Winter, Discuss. Faraday Soc., 28, 183 (1959); J . Catal., 15, 144 (1969); K.Hauffe, R. Clang, and H. J. Engell, Z. Phys. Chem. (Leipzig), 201, 223 (1952); G. M. Schwab and J . Block, Z . Phys. Chem. (Frankfurt a m M a i n ) , 1, 42 (1954); H.J. Engell and K. Hauffe, 2. Elektrochem., 57, 776 (1953). (4) I. A . Myasnikov and 8. Ya. Pshezhetskii, Dokl. Akad. Nauk S S S R , 99, 125 (1954). (5) Y. Fujita and T. Kwan, Bull. Chem. SOC.Jap., 31, 379 (1958). (6) T. I. Barry and F. S . Stone, Proc. Roy. SOC., Ser. A , 255, 124 (1960). (7) K. Tanaka and A . Ozaki, J . Catal., 8, 307 (1967); K. Tanaka and A. Oraki, Bull. Chem. SOC.Jap., 40, 420 (1967). (8) 8. R. Morrison and T. Freund, J . Chem. Phys., 47, 1543 (1967); T. Freund and 8. R . Morrison, Surface Sci., 9, 119 (1968); 8. R. Morrison, ibid., 10, 459 (1968); T . Freund and S . R . Morrison, Electrochim. Acta, 13, 1343 (1968); S. R. Morrison, Surface Sci., 15, 363 (1969). The JOUTnal of Physical Chemistry, Vol. 76,No. 8,1971
KEN-ICHITANAKA AND GEORGEBLYHOLDER
1038
V
& _ I
63 m m
time hicn.)
(A1 Figure 1. (A) Schematic drawing of the apparatus: C, flowing water filter; M, front surface mirror; F, furnace; R, reactor; G, to gas chromatograph; P, circulation pump; V, vacuum line; S, sample supply. (B) Details of reactor.
3-mm thickness). It was mounted in a furnace with a Pyrex window at the bottom. Illuminating light was applied to the reactor by reflection a t a front surface mirror as shown in Figure 1. A blank test of the reaction under illumination without catalyst was carried out to be sure that no detectable reaction occurred at 390" either in the gas phase or on the cell walls. The oxide used in the present experiment was a powdered sample of Kadox-25 obtained from New Jersey Zinc Co. A 2.38-g sample of zinc oxide was placed in the reaction cell and was activated by more than five treatments of overnight oxidation and several hours' evacuation at about 410" to remove carbonate impurities. The oxide was pretreated by contacting with nitrous oxide at the reaction conditions (in photocatalytic experiments the pretreatment was carried out under illumination) for more than 30 min, followed by about a 10-min evacuation. Such a pretreatment was carried out before each experimental run and gave a stable and reproducible surface in the present experiments on zinc oxide for both dark and photoreactions. Similar pretreatment schemes have been used on other oxides'tg to obtain reproducible surfaces for nitrous oxide decomposition. Before each experimental run the zinc oxide powder was checked for uniform spreading over the reactor window. Nitrous oxide from a commercial cylinder was purified by freezing at liquid air temperature and sublimating a t Dry Ice-methanol temperature. Oxygen from a commercial cylinder was purified by passage through a liquid air trap. The effect of oxygen pressure on the reaction rate was studied by mixing a certain amount of oxygen with the nitrous oxide flowing The Journal o j Physical Chemistry, Vol. 76, No. 8,1071
Figure 2. Time course of nitrogen evolution in the thermal catalytic reaction. The temperatures and the initial pressures of nitrous oxide are (1)371°, 45.7 mm; (2) 405", 46.9 mm; and (3) 426O, 45.7 mm.
through the reactor. The product mixtures, oxygen and nitrogen, were analyzed in an on-line gas chromatograph using a molecular sieve 5A column. The reaction rates were obtained from plots of the nitrogen evolution time course determined by gas chromatographic analysis.
Results I. Thermal Catalytic Decomposition of Nitrous Oxide. The thermal catalytic decomposition of nitrous oxide was studied over the temperature range 371426". Figure 2 shows typical time courses of nitrogen evolution obtained at 371,405, and 426". The oxygen evolved is not shown in this figure; however, it is stoichiometric within experimental error. Similar time courses of the reaction were obtained with various initial pressures of nitrous oxide. Figure 3A shows the relation between the initial pressures and the reaction rates obtained from the slopes of the time courses at 424". The effect of oxygen on the reaction rate was studied at the initial oxygen pressures of 0, 1.83, and 2.93 Torr with an initial pressure of 6.3 Torr of nitrous oxide at 424". The result shown in Figure 3B indicates that the addition of oxygen does not change the reaction rate. These results indicate that the kinetics of the thermal catalytic decomposition of nitrous oxide on zinc oxide are first order in nitrous oxide pressure and zero order over the range in oxygen pressure, that is, r = ICPN,~, studied. The logarithms of the rate constants a t various temperatures are plotted against the reciprocal of the temperatures in Figure 4. The apparent activation energy of the thermal catalytic decomposition of nitrous oxide is calculated as 35 kcal/mol. 2. Photocatalytic Decomposition of Nitrous Oxide. (9) L. Rheaume and G. Parravano,
J. Phys. Chem., 63, 264 (1968).
1039
DECOMPOSITION OF NITROUS OXIDEON ZINC OXIDE
45.0* 0
I
[
2,01 1 .o
1 .o
0.5
N
0
0
x
5
10
15
20
E
25 time (min.)
Figure 5. Time course of the photocatalytic reaction a t 400'. The initial pressures of nitrous oxide are: (1) 19.0 mm; (2) 46.8 mm; (3) 107.8 mm; and (4) 158.6 mm. Solid lines indicate nitrogen pressures and dotted lines indicate oxygen pressures. PNpn (cmHg)
Figure 3. (A) Nitrous oxide pressure dependence of thermal catalytic reaction rate a t 424'. Solid circle shows the value with oxygen added. (B) Effect of oxygen pressure on thermal catalytic reaction rate. PN*O= 6.3 cm and 424'.
r
10.0
Y v
1.4
1.6
1.5 + X
5.0
10.0
15.0
10)
Figure 4. Arrhenius plot of rate constants: reaction; D, thermal catalytic reaction.
L,photocatalytic
The photocatalytic decomposition of nitrous oxide was carried out over the temperature range of 371-431". Figure 5 shows typical results of the photocatalytic decomposition of nitrous oxide at 400" with various initial pressures of nitrous oxide. It is characteristic of the photocatalytic reaction that the slope of the time course gradually decreases as the reaction proceeds. Figure 6 shows a typical pressure dependence of the reaction rate obtained from the initial slope of the time course curve under illumination a t 400 and 418". The
Figure 6. Nitrous oxide pressure dependence of the reaction rate under illumination: solid line, 418; dotted line, 400'. Straight lines show the expected thermal catalytic reaction rates.
straight lines in the figure show the expected thermal catalytic reaction rates at the given temperatures. It is found that the pressure dependence of the reaction rate is markedly different from that of the thermal catalytic reaction. This figure shows that the reaction rate under illumination does not increase in direct proportion to the nitrous oxide pressure but that the slope decreases with increasing nitrous oxide pressure, and The Journal of Physical Chemistry, Vol. Y6, No. 8,1971
KEN-ICHITANAKA AND GEORGE BLYHOLDER
1040
...
N 0
---------------- - I
0 0
1
2
3
4
5
I 6
PC2 (cmHg1 G y
0
Figure 7. Effect of oxygen pressure on the reaction rate under illumination. PN*O= 4.6 cm: (1) 431"; (2) 399". Straight lines show the expected thermal catalytic rate.
that at higher pressures the slope appears to approach that of the thermal catalytic reaction. This behavior of the reaction under illumination was also found at 371,378,387, 391,400, and 418". The effect of the added oxygen on the reaction rate under illumination was studied at 371, 399, and 431". While the oxygen pressure has no effect on the thermal catalytic decomposition of nitrous oxide, it was found that the reaction under illumination is markedly retarded by oxygen. Figure 7 shows the results of the oxygen effect on the reaction rate under illumination at 399 and 431". The result at 371" is similar. The straight lines in the figure show the expected thermal catalytic reaction rates upon which oxygen has no effect. This figure shows that the reaction rate under illumination is markedly reduced in the presence of a small amount of oxygen. The results shown in Figures 6 and 7 reveal that the reaction under illumination consists of a photocatalytic reaction plus a thermal catalytic reaction. The photocatalytic reaction may be obtained by subtracting the thermal catalytic reaction rate from the overall reaction rate obtained under illumination. Figure 8 shows this photocatalytic reaction rate at various temperatures as a function of nitrous oxide pressure. It is seen from this figure that the photocatalytic reaction approaches zero-order kinetics at higher nitrous oxide pressures. By the same procedure, the oxygen effect on the photocatalytic reaction rate may be obtained for the reactions at 371,399, and431".
Discussion The most remarkable result in this work is that the thermal catalytic and photocatalytic reactions obey completely different kinetics, and these two reactions appear to be simultaneously occurring on the zinc oxide under illumination. The zinc oxide used in the present experiment is well oxidized in the reaction cell and is sufficiently acclimatized to the reaction condition by the pretreatment The Journal of Physical Chemistry, Vol. 76, N o . 8,1971
I
5.0
10.0
15.0
Figure 8. Pressure dependence of the photocatalytic reaction rate a t various temperatures: (1)418"; (2) 400"; (3) 391"; (4)387"; (5) 378"; (6) 371".
scheme to give reproducible results.
Cunningham,
et al.,1° recently reported that zinc oxide, evacuated at
400" for 16 hr, decomposed nitrous oxide very rapidly a t room temperature, but in this case 95% or more of the product was nitrogen and the reaction did not proceed catalytically. They proposed that the very low concentration of holes in zinc oxide greatly reduced the rate of neutralization of product oxygen ions at room temperature so that the desorption of this oxygen would be the slow step in the catalytic decomposition of nitrous oxide at higher temperatures. It is, however, highly doubtful that the oxygen strongly held on the evacuated zinc oxide is the same type of oxygen as that adsorbed as an intermediate in the catalytic reaction at higher temperatures because the chemisorption of oxygen on highly evacuated zinc oxide is particularly strong and the form of the adsorbed oxygen is expected to change as the temperature and coverage changes. The heterogeneous character of the surface is readily shown by a comparison of the work done by various investigators. For example, the isotopic equilibration between IaOzand IeOzoccurs rapidly on zinc oxide even a t low temperatures.6Sl1 This fact indicates that dissociative adsorption and desorption of oxygen occurs rapidly at temperatures at which strongly held oxygen such as produced by the decomposition of NzO on a highly evacuated surface at room temperature does not desorb. Furthermore, the rate of adsorption and deTorr obey sorption of oxygen at pressures around the Elovich equationj6 but the rate of evolution of oxygen during catalysis increases with nitrous oxide pressure according to first-order kinetics (Figure 3). The added oxygen has no effect on the reaction rate and the time courses of the thermal catalytic reaction do (10) J. Cunningham, J. J. Kelly, and A. L. Penny, J. Phys. Chern., 74, 1992 (1970). (11) E.R. S. Winter, J. Chern. Soc., 1522 (1954).
DECOMPOSITION OF NITROUSOXIDEON ZINCOXIDE
1041 into the conduction band are repelled from the surface by the potential barrier (a)a t the surface. Thus, the number of electrons per unit volume at the surface, e,, is given by
where eT is the total number of conduction band electrons in the solid particle of volume V , p is the electronic charge, and a is the potential barrier at the surface. Since most of the holes produced by illumination are attracted to the surface by the surface potential,8 the number of holes per unit volume at the surface, h,, is given by
hT
-
h,
Vh
Figure 9. Model for setting up the rate equations: D, dark state; L, illuminated state.
not bend (Figure 2). These facts imply that the oxygen adsorbed on the surface as an intermediate of the thermal catalytic decomposition of NZO is highly labile and weakly held oxygen like an intermediate of the equilibration reaction. Accordingly, the low catalytic activity at lower temperatures may not be due to a low concentration of holes, i e . , slow oxygen desorption, but rather to a slow decomposition of adsorbed N2O. The detail reaction mechanism will be discussed later. While the electronic structure of the actual surface is presumably complicated by nonuniformities, we propose to describe that part of the surface on which the reaction proceeds catalytically in terms of the band model shown in Figure 9. Electrons in the conduction band are probably very few at our dark experimental conditions for the following reasons. (1) Most of the donor defects have been destroyed by thermal oxidation. l 2 (2) The majority of the electrons in the conduction band are trapped by adsorbed oxygen whose ions are presumably different from an intermediate of the decomposition reaction as discussed above. When the oxide is exposed to light of energy corresponding to the energy gap between the conduction band and the valence band, i e . , 385 mp or shorter wavelengths, electrons in the valence band are excited into the conduction band as shown by process 1 in Figure 9. Electrons excited
where h~ is the total number of holes and Vh is that volume next to the surface in which the holes are distributed. When the only surface acceptor is nitrous oxide, which is known to be an efficient electron scavenger, electron transfer between zinc oxide and NzO and/or NzO- will be established. Assuming that direct electron transfer from the valence band to nitrous oxide molecules rarely occurs a t the experimental conditions, the electron transfer processes established on the surface are described as follows (these processes correspond to 2 and 3 in Figure 9). The electron transfer rate of each process is N2O(g)
+ e,
NzO-(a)
k2
NWa)
+ h, -$ NzO(g)
expressed as V2
= lCzesPNzo
- kz'(NzO-),
75 = ksh,(NzO-),
(3)
(4)
If NzO- is the only charged molecule being considered in a total charge balance, then the following charge balance is required h~ =
eT
+ (Nz0-1~
(5)
where (NtO-)T designates the total number of NzOions on the surface. Combining eq 1, 2, and 5 with the notation that the number of NzO- ions per unit surface area, (NzO-),, is equal to (N~O-)T/S, where S is the surface area, gives the value of e, as
At a steady state under illumination, the rate of elec(12) K. M . Sancier, J. Catal., 5, 314 (19613).
The Journal of Physical Chemistry, Vol. 76,No. 8,1071
KEN-ICHITANAKA AND GEORGE BLYHOLDER
1042 tron capture equals the rate of hole capture; Le., =
Vz
V3.
Accordingly, at the initial stage of the reaction, that is, the only surface acceptor is nitrous oxide, the following relation is established
k 2 P ~ , 0 es kz'(NzO-), = kah,(N20-),
(7)
Substituting e, from eq 6 into eq 7 and solving for (NzO-), gives
(NzO-),
=
(kz'
+
kzvhh~p~zo k3h,)Veqa/kT ICZSPN~O (8)
+
Assuming that the slow step of the photocatalytic reaction is the decomposition of nitrous oxide ion, the initial rate of the reaction, proceeding on the surface on which nitrous oxide is the only molecule capturing electron, is described by
Figure 10. Plot of I/rate us. 1 / P ~ , oa t various temperatures: (1) 371"; (2) 378"; (3) 387'; (4)391'; (5) 400'; (6) 418".
This equation implies that the reciprocal of the reaction rate vs. l / P ~ , owill be a straight line if the variation of a,which in general is a function of P N ~ ois, negligible. Figure 10 shows the expected linear relation a t various temperatures. It may be noted that the term exp. (qa/kT) plays a role in the kinetics in the low nitrous oxide pressure range. When two kinds of surface acceptors are on the surface, oxygen and nitrous oxide, we derive similar equations for the electron transfer processes corresponding to (4) and (5) shown in Figure 9 as follows
V4 = 1c4POz1/Zes - ]C~'(O-), Vs ksha(O-)a
(10) (11)
I n this equation, 0- is assumed to be an important adsorbed state of oxygen during the photocatalytic reaction a t around 400". + + At a steady state under illumination, V 2 = Va, V4 = VS,and hT = eT ( N 2 0 - ) ~ O-T. Then, the electron concentration a t the surface e, is given by
+
+
(12) and the concentration of NzO- is given by
Accordingly, the reaction rate, r , is given by
kkzh,Vh
-
r = k(N2O-)s (14) The Journal of Physical Chemistry, Vol. 76,No. 8,1971
Figure 11. Plot of l/rateus. Z/g2 a t various temperatures; PN*O= 4.6 cm; (1) 371"; (2) 399"; (3) 431'.
This equation is just like eq 8 with the oxygen term added. This equation predicts that at a constant value of PN%O the reciprocal of the reaction rate will change linearly with the square root of oxygen pressure. Figure 11 shows the results obtained a t various temperatures a t a constant nitrous oxide pressure, P N ~=O4.6 cm, are indeed in accord with this expectation. Thus both the effect of added oxygen and the bending of the time course of the photocatalytic reaction (Figure 5 ) are reasonably explained by the competitive electron scavenge effectof oxygen. I n deriving expressions to fit the somewhat complex kinetics observed, it was convenient to assume that the decomposition of N20- was the rate-controlling step. Further support for this assumption is found in evidence which suggests that desorption of oxygen should not be the rate-controlling step. Several workers have dem-
DECOMPOSITION OF NITROUS OXIDEON ZINC OXIDE onstrated that the photoeffect on the rate of oxygen desorption from ZnO decreased with increasing temperature and was not detectable above 2OOoK or above ~ O E I " . ~However, we find a large photoeffect on the decomposition of NzO and hence conclude oxygen desorption is not rate determining. The results obtained in this paper imply that the thermal catalytic and photocatalytic reaction occur simultaneously over the illuminated zinc oxide. If we consider the difficulty of homogeneous illumination of powder samples, we cannot say for certain whether the thermal catalytic reaction occurs on an illuminated part of the surface or not. Since Dell, et aL12pointed out the importance of electronic properties of oxides in the catalytic activity of nitrous oxide decomposition, the reaction mechanism of the thermal reaction has been explained by an electron transfer mechanism. However, the completely different kinetics of the thermal catalytic reaction and the photocatalytic reaction strongly suggest different mechanisms in these reactions. Just being a photochemical reaction and having an effect by the electron scavenger oxygen fairly definitely leads to an electron transfer mechanism for the photodecomposition. However, it seems clear that the thermal catalytic reaction does not pass through an intermediate such as NzO- which is common to the photocatalytic process. Accordingly, it is suggested that the thermal reaction proceeds via an atomic and molecular mechanism as
-
NzO(8)
NzO(g)
n'ZO(S)
Nz(g)
Ob)
+
O(S)
(slow)
L- 1/202(g)
If the desorption of oxygen were rate controlling, the active surface sites would be covered with oxygen and
1043 the reaction would be less than first order in NzO pressure. Since the reaction is observed to be first order in N20 pressure, the desorption step is not rate controlling. The activation energy of the photocatalytic decomposition of nitrous oxide, which was obtained from the temperature variation of kh,Vh/S, was 33 kcal/mol at temperatures higher than 390" and 13 kcal/mol a t temperatures lower than that. Winter'l found in his isotopic exchange reaction that the exchange between adsorbed oxygen and lattice oxide ions is faster than the surface migration of adsorbed oxygen at temperatures higher than 415". It seems reasonable that these two critical temperatures are related to each other. Below 400" it seems likely that the nature of the surface remains unchanged as the temperature varies so the activation energy of 13 kcal/mol reflects the temperature coefficient of k in the term kh,Vh/S. Above 400" Winter's work indicates the lattice becomes mobile so that the active sites a t the surface can change with temperature. This will result in the term h,Vh/S exhibiting a temperature coefficient above 400" as observed. It may be noted, however, that the basic mechanism, as reflected in the dependence of the reaction on NzOand O2pressures, is not changed. I n our interpretation the activation energy of 35 kcal/mol for the dark reaction represents the barrier to dissociation of a neutral N2O molecule minus the heat of adsorption, while the 13 lmd/mol activation energy for the photocatalytic reaction represents the barrier to dissociation of NzO-. The difference in activation energies is then the result of putting an electron into an antibonding molecular orbital of N20.
Acknowledgment. This investigation was supported in part by PHS Research Grant No. 00818-01 from the National Air Pollution Control Administration.
The Journal of Physicat Chemistry, Vol. 76, No. 8, 1071
