The Journal of
Physical Chemistry
0 Copyright, 1990, by the American Chemical Society
VOLUME 94, NUMBER 16 AUGUST 9,1990
LETTERS Stoichiometric Catalytic Decomposition of Nitric Oxide over Cu-ZSM-5 Catalysts Yuejin Li and W. Keith Hall* Department of Chemistry, University of Pittsburgh, Pittsburgh, Pennsylvania 15260 (Received: March 15, 1990; In Final Form: May 7 , 1990)
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Copper cation exchanged ZSM-5 zeolites are effective catalysts for the NO decomposition reaction. The present studies is intrinsically stoichiometric, and complete conversion have revealed that the conversion of the reaction NO '/2N2+ '/202 can be obtained over a Cu-ZSM-5 catalyst. Operated in this way the undesirable side reaction NO + ' / 2 0 2 NO2 can be avoided. The decomposition reaction was inhibited, but not poisoned, by 200 Torr of 02.
Introduction
With the linking of the nitric oxides to urban smog formation, acid rain, and, hence, harmful health effects, the removal of NO from the exhaust stream of various combustion sources has become increasingly important. Except at very high temperatures, nitric oxide is thermodynamically unstable with respect to its molecular elements, O2and N2. Therefore, direct catalytic decomposition of N O should be possible and certainly is desirable. For several decades, effective catalysts for NO decomposition have been sought without notable success. Numerous data are available for a variety of metal and metal oxide catalysts. However, until recently none of these have been sufficiently active to have industrial significance. The earlier studies on catalytic decomposition of N O have been reviewed by Hightower and Van Leirsburg.' Other pertinent more recent work is discussed below. The conventional view regarding the low catalytic activity for NO decomposition is that the oxygen atoms produced by N O decomposition remain strongly adsorbed on the catalyst, thus self-poisoning further decomposition. Amirnazmi et aL2 reported fairly low activities for this reaction over Pt/A1203 and CuO at (1) Hightower, J . W.; Van Leirsburg, D. A. In The Catalytic Chemistry of Nitrogen Oxides; Klimisch, R. W., Larson, J. G., Eds.; Plenum: London, 1975; p 63. (2) Amirnazmi, A.; Benson, J . E.; Boudart, M. J . Card. 1973, 30, 55.
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600 OC and pointed out that the product O2had a strong inhibiting effect and, to further complicate matters, reacted with NO producing NO2. Consequently, an effective catalyst for NO decomposition must be active enough to cleave the N-O bond, but the surface-0 bond must be weak enough to allow oxygen to desorb at reasonably low temperatures, Le., below 600 OC. Evidently, materials that are radically different from the conventional catalysts are needed to fulfull this requirement. Iwamoto and co-workers have been active in this area for nearly a They have tested a number of zeolites containing altervalent baseexchange cations, but of these, only the Cu2+forms showed appreciable a ~ t i v i t y .They ~ reported that CuZ+-ZSM-5 (3) Kagawa, S.;Yokoo, S.;Iwamoto, M. J . Chem. Soc., Chem. Commun. 1978, 1058. (4) Iwamoto, M.; Yokoo, S.;Sakai, K.; Kagawa, S.J . Chem. SOC.,Faraday Trans. 1 1981, 77, 1629. (5) Iwamoto, M.; Nakamura, M.; Nagano, H.; Kagawa, S.J. Phys. Chem. 1982, 86, 153.
(6) Iwamoto, M.; Maruyama, K.; Yamazoe, N.; Seiyama, T. J . Phys. Chem. 1977.81, 622; J . Chem. SOC.,Chem. Commun. 1980, 842. (7) Iwamoto, M.; Furukawa, H.; Kagawa, S. In New Deuelopmenfs in Zeolire Science Technology; Murikami, Y., Ed.; Elsevier Publishers: New York, 1986; p 943 ff. (8) Iwamoto, M.; Yahiro, H.; Tanoda, K. To be published in the 1988 edition of Studies in Surface Science and Cafdysis(private communication). (9) Iwamoto, M.; Yahiro, H.; Mine, Y.; Kagawa, S.Chem. Left. 1989, 213.
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TABLE I: Catalyst Analysis catalvst Si/AI CU-ZS M- 5-24-76 24.2 CU-ZSM-5-26-166 26.4 CU-ZSM-5-12-140 12.3
Cu/AI 0.38 0.83 0.70
Letters
Cu content" 0.25 0.49 0.83
I n units of millimoles of Cu per gram of zeolite.
TABLE 11: Catalyst Preparation concn of solution X catalyst IO', M C U-ZS M - 5-2 4-7 6 25 CU-ZSM-5-26-166 50 CU-ZSM-5-1 2-140 50
;;;;;
to oxygen
zeolite/ solution, g/L 81 50 3
exchange time, h 16" 16" 123b
drying temp, OC 80 80 25
2O!fo;
0
was especially active for N O decomposition.' At a contact time of 8 s, conversion levels of >90% were reported, and no deterioration of the effectiveness of the catalyst was observed over a period of 30 h. However, their data also showed that stoichiometric amounts of O2and N 2 were not produced from the N O which was reacted, i.e., that the amount of N 2 was significantly lower than that needed to correspond to the N O consumed, and the O2 deficiency was even greater. It was suggested that NO2 was produced, and this was detected by mass spectrometry but not quantified. It is evident that to understand the system and the chemistry involved a complete mass balance is necessary. Moreover, if NOz is indeed formed, is this process homogeneous, catalytic, or on the GC columns during analysis?2 And, is NO2 formation intimately and inseparably related to the NO decomposition reaction? These questions are addressed herein.
Experimental Section The catalysts used were Cu2+-exchanged ZSM-5. These samples are identified in the following fashion: cation-type of zeolite-Si/Al ratio-Cu exchange level in percentage, e.g., CuZSM-5-26-166. The available information about these catalysts is listed in Table I. The elemental analyses were provided by Air Products & Chemicals Co. (Cu-ZSM-5-24-76 and CuZSM-5-12-140) and by Galbraith Laboratories (Cu-ZSM-526-166). Note that two of the catalysts (Cu-ZSM-5-26-166 and CU-ZSM-5-1 2-1 40) contain more than the stoichiometric amount of Cu2+ required to satisfy the base-exchange capacity of the zeolite. (Stoichiometric exchange occurs when 1 Cu2+replaces 2 Na+ and the molar ratio Cu/AI = 0.5.) The starting materials were the Na+ form of ZSM-5 zeolites made by template free preparation.'0 Dilute aqueous solutions of copper(I1) acetate were used for the ion exchanges. The details concerning the preparation are listed in Table 11. After the ion exchanges the preparations were washed twice, with stirring, in water for 2 h. The standard in situ pretreatment of these catalysts for activity measurement was drying in He at 550 OC for 2 h. The activity measurements were made with a microcatalytic reactor in a steady-state flow mode. The reactor was a Pyrex 4-mm4.d. glass tube with an expanded section (8-13-mm i.d.) as a catalyst bed. The weight of samples varied from a few tenths of a gram to one gram. The contact time was varied between 1 and 10 s to achieve desired conversion levels. The reactor was surrounded by a temperature-controlled furnace. The temperature was monitored by a chromel-alumel thermocouple which was in contact with the catalyst bed. The reactant gas was a mixture of 4% NO in He provided by Air Products & Chemicals Co. Product analysis was obtained using an on-line gas chromatograph (GOW MAC 69-500) with a thermal conductivity detector. The separation column was packed with 5A molecular sieve (80/100 mesh); it was 4 ft long, '/*-in. o.d., and was home-made. The ~~~
O
O
hz-
TimeNoui
"Exchanged once at 25 "C. bExchanged twice (48 h for the first exchange) at 25 OC.
~~
Ibj
~
(IO) Sh~ralkar,V P , Clearfield, A. Zeolife 1989, 9, 363 ( 1 1 ) Harrison, B : Wyatt, M ; Gough, K G In Catalysis: Royal Society of Chemistry London, 1982, Voi 5 . p 127
Figure 1 . Conversions of NO decomposition as functions of time over Cu-ZSM-5-26-166 at 773 K and a contact time of 2 s. The lines taken from ref 7 were obtained on Cu-ZSM-5-25-54 at the same temperature and a contact time of 8 s.
J
50 -
40
-
-. 5 30 =
.-
=9 v)
2
20-
+
I
mtrogen
7 nitric oxide
0 0
10
20
30
40
50
60
Time/" Figure 2. Mass spectroscopy signals of the product stream of NO decomposition reaction over Cu-ZSM-5-26-166 at 773 K and a contact time of 1 s. The background signals of O2and N2 were 4 and 16 (arbitrary units), respectively. Shown are the uncorrected data. The instrument used was not calibrated for sensitivity, and so the data are presented as arbitrary units (au).
column temperature was 25 O C . The flow rate of the carrier gas was 30 cm3/min. The data were integrated and plotted with a Hewlett-Packard intergrater (3390A). A cold trap thermostated in dry ice and acetone slurry was placed between the sampling valve and the GC column to remove any NO2 formed. It was possible to bypass the trap to connect a mass spectrometer (QUAD 250A residual gas analyzer). However, quantitative measurements of NO2 were made volumetrically by transferring the NO2 from the trap into a stainless steel vacuum system. The pressure was read from an electric Baratron gauge with a best accuracy of 0.001 Torr. (Because NOz reacts with Hg, a conventional BET system could not be used for this purpose.)
Results and Discussion The activities for NO decomposition over Cu-ZSM-5-26-166 catalyst a t 773 K are presented in Figure 1 as functions of time where they are compared with similar results of Iwamoto et al.' The excellent agreement shown confirms their earlier observations and establishes that a common problem is being investigated. The conversions of NO (the percentages of NO which disappeared) were over 90%, and the conversions to N2and O2(as percentages of NO in the feed converted to N2 or 02)were about 60% and 20%, respectively. The activity increased slowly with reaction time. After 24 h of continuous operation, the catalyst was actually more active than initially. In a similar experiment, the product stream was analyzed by a mass spectrometer at the beginning of the reaction and, occa-
Letters sionally, later. As shown in Figure 2, 02,N,, NO, NO2, and N 2 0 were detected. NO2 formation, which was negligible initially, increased with time and reached a steady state after about 30 min. On the other hand, N 2 0 , initially formed in substantial amounts, drastically decreased with time and was not detectable in the steady state. Conceivably, N O was decomposed most completely at the beginning of the reaction, and this was evidenced by the largest amount of N 2 formed. However, the O2formed from the decomposition in the early stages reacted with the catalyst; hence, more O2was released from the catalyst as it came into the steady state. In a separate experiment, when the reaction reached a steady state, a complete mass balance was achieved by combination of chromatographic and volumetric techniques. NO, 02,and N, were analyzed by the gas chromatograph; NO2 was trapped upstream and measured volumetrically, providing a mass balance of 99 f 5% on N2 and 98 f 5% on 02.Thus, it was confirmed that the gaps among the conversion of NO, conversions to N2 and to O2shown in Figure I , are bridged by NO2 formation. Two of the possible ways in which NO2 could be formed were in the catalytic reaction itself or on the chromatographic column during analysis.2 The latter was not important in the present case as evidenced by the closure of the mass balance mentioned above. Another was homogeneously at the reaction temperature (typically, 773 K). Finally, NO is known to react with 0, in the homogeneous gas phase at low temperatures (e.g., room temperature in the exit lines from the reactor). The high-temperature gas-phase reaction should be sensitive to reactor design, viz., the volume of the hot postcatalytic zone. To test this, a microreactor was designed such that reactant could be made to flow either in the regular or in the reverse direction, but with the volume of the postcatalytic zone 60 times smaller (a capillary tube) in the regular flow direction than that in the reverse flow direction. Thus, it would be expected that if any gas-phase reaction occurred in the postcatalytic zone, the O2 formation would be lower and the NO consumption would be higher when the reverse flow was employed than when the regular flow direction was employed. However, the results obtained from both flow directions were the same within an uncertainty of 1%. This demonstrated that a high-temperature gas-phase reaction was not responsible for NO2 formation. A low-temperature gas-phase NO oxidation reaction was tested in blank experiments by flowing a mixture of NO (4% N O in He) and 0, (4% 0, in He) through the system in the absence of catalyst. An empty glass U-tube at room temperature replaced the catalytic reactor. Substantial amounts of NO2 were formed and collected downstream in a cold trap. This was evidenced visually by the appearance of a very dense brown color in the glass lines leading to the trap and in the trap itself when it was warmed up. Quantitative analysis of the contents of the trap allowed us to conclude that this low-temperature process was responsible for the distortion of the mass balance from the simple stoichiometry, 2 N 0 N, + 02,as depicted in Figure 1 and, more important, that the conversions of NO to N2 measure the extent of the decomposition reaction, whereas the disappearance of N O measures the sum of this reaction and the further reaction with the 0, produced. Thus, the former may be correlated to the 0 2 / N 2 ratios detected by the gas chromatograph. A series of conversion levels for N O decomposition over CuZSM-5-26-166 and Cu-ZSM-5-12-140 catalysts were obtained at a variety of reaction temperatures and reactant flow rates. These results are shown in Figure 3. The 0 2 / N 2 ratio increased monotonically with the conversion of N O to N2 on a common line independent of the external variables. Moreover, at a conversion level of 100% the 0 2 / N 2 ratio became unity (stoichiometric). Below 40% of conversion, O2was not detected. At higher conversions, the amounts of N O left in the tail gas were small, and the 02 consumptions to form NO2 were therefore low. This clearly demonstrates that the reaction 2 N 0 + O2 2 N 0 2 takes place in the reactor lines in the homogeneous gas phase at ordinary temperatures and that the 0 2 / N 2 ratio is limited only by the availability of NO in the tail gas. The fact that the relationship shown in Figure 3 is independent of the reaction temperature and
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The Journal of Physical Chemistry, Vol. 94, No. 16, 1990 6147 0
1.01 0 .O
$ .G
0.5-
0.3
BO
0.2
6
0.1
3
0
0.4-
8
-
T=773K A
0
..
T=823K T=748K T=723K
0 0
T=698K T=773K, (-12-140)
Conversion of NO to Nitrogen (%)
Figure 3. Oxygen/nitrogen ratio detected by gas chromatograph versus the conversion of nitric oxide to nitrogen. Except the point at the full conversion level (on CU-ZSM-5-1 2-140), all other points were obtained on CU-ZSM-5-26-166.
TABLE 111: Equilibrium Constant, K,, versus Temperature for the Reaction 2N0 + O2 == 2N02"
temp, K 298 400 500 600 a
KP 2.43 X 1.80 x 1.69 x 1.57 x
10l2 107
104 102
temp, K 700 773 800 900
KP 5.51 8.20 4.46 6.29
X X X X
IOo
IO-' IO-' IO-*
Taken from ref 1 1.
catalyst further indicates that the NO2 formation is not catalytic. The data of Figure 3 are sufficient to show that the catalytic decomposition of NO to 1/2N2+ 1/202 is intrinsically stoichiometric. The complete conversion of N O to N, and O2 was obtained on Cu-ZSM-5-12-140 at 500 OC using a 0.75-g of sample and a flow rate of 11 cm3/min of 4% N O in He. This result was reproduced several times under identical conditions. It is possible, however, that a complete conversion of N O to N 2 and 0, can be achieved over some other catalysts, e.g., Cu-ZSM-5-26-166, if the contact times are sufficiently long. The reaction 2 N 0 + 0, 2N02 is thermodynamically favored by low temperatures. The equilibrium constants, Kp, are listed for various temperatures in Table 111. Interestingly, when the N O decomposition reaction is carried out above 773 K, NO2 formation becomes limited by thermodynamics, and if the reaction is complete, NO2 formation can be completely avoided. Otherwise, the secondary reaction in the cold part of the system is kinetically controlled. The rate equation of this reaction is generally thought to be a third-order process, d[N02]/dt = k[NOI2[O2],and its rate constnat has been reported by several authors.12J3 A typical value, calculated for ppm concentrations of N O by Morrison et al.,12is 1.3 X lo4 (L2/(mo12s) at 26.5 O C . Thus, it is probable that the NO2 formation rate will remain low if the concentration of NO is small, e.g., ppm levels. Of course, the walls of the reactor system may act as the third body. In a practical application, NO2 formation can be minimized by a thoughtful design of the catalytic converter (reactor) such that any mixture containing unconverted N O is quickly released to the atmosphere. Of course, the best way to minimize NO, formation in this process is to maximize N O conversion. The effects of addition of O2and H 2 0 vapor to the feed for the decomposition reaction were tested. Addition of from 1% to 20% of O2to the feed led to mild inhibition of the decomposition (12) Morrison, M. E.;Rinker, R. G.;Corcoran, W . H. Ind. Eng. Chem. Fundam. 1986, 5 , 175. ( 1 3 ) Ashmore, P. G.;Murnett, M. G.; Tyler, B. J. Trans. Faraday SOC. 1963, 58, 685.
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J . Phys. Chem. 1990, 94, 6148-6150
way to compensate for the decrease of the conversion due to O2 and H 2 0 might be to raise the reaction temperature.
to N2. On our most active catalyst, Cu-ZSM-5-12-140, a conversion of 97% was obtained at 500 OC without adding O2to the feed. On the same catalyst and at the same reaction temperature a conversion of 80% was obtained after adding 10%O2into the feed for 1 h. However, the conversion was restored to the original level, 97%, in 30 min after elimination of O2 in the feed. Adding 2% H 2 0 vapor reduced the conversion from 97% to 72%, but the recovery of the conversion took longer than with 02.One possible
Acknowledgment. Thanks are due to the Air Products and Chemicals Co. for support of this research and in particular to Dr. John Armor for his continual interest and encouragement. We thank T. R. Gaffney and C . Coe for supplying the catalysts studied.
Non-RRKM Kinetics in Gas-Phase SN2 Nucleophilic Substitution Scott R. Vande Linde and William L. Hase* Department of Chemistry, Wayne State University, Detroit, Michigan 48202 (Received: March 20, 1990; In Final Form: June 8, 1990)
Classical trajectories are used to simulate the intramolecular and unimolecular dynamics of CI--CH3CI complexes formed by collisions between CI- and CH3CI. Unimolecular rate constants for decomposition of the complexes are decidedly non-RRKM and are consistent with a model in which only three or fewer modes in the complex are active in intramolecular vibrational energy redistribution.
There is considerable interest in identifying for what molecules and under what experimental conditions is the RRKM theory inadequate for describing the rates of unimolecular reactions.' Extensive exmrimental and theoretical studies have shown that RRKM thedry correctly describes a wide array of unimolecular reactions2 This is because the basic assumptions of the theory may be valid and/or the collision and energy averaging of many experiments makes the analysis of experimental results insensitive to the theory's assumption^.^*^ Though RRKM theory has been highly successful, there are an important number of situations where it does not agree with experiment. For the most part RRKM theory adequately describes the pressure and temperature dependence of thermal unimolecular reactiom2 However, even for relatively complex molecules, it has been found that the pressure and temperature dependence of isomerization rate constants are not in accord with RRKM theory if the isomerization barrier is sufficiently small.s Another case involves van der Waals molecules, for which usually neither the variation in the unimolecular rate constant with energy nor the magnitude of the rate constant agrees with RRKM theory.6-8 Finally, molecules which exhibit RRKM behavior when studied in energy- and collision-averaged macroscopic settings may have state-specific non-RRKM rate constants when studied at the microscopic leve1.+l2 ( I ) H a s , W. L. In Potential Energy Surfaces and Dynamics Calculations; Truhlar, D. G., Ed.; Plenum: New York, 1981; p 1. (2) Oref, I.; Rabinovitch, B. S.Ace. Chem. Res. 1979, 12, 166. (3) Marcus, R. A.; Hase, W. L.; Swamy, K. N. J. Phys. Chem. 1984,88, 6717. (4) Lu, D.-h.; Hase, W. L. J . Phys. Chem. 1989, 93, 1681. Lu, D.-h.; Hase, W. L. J. Chem. Phys. 1989, 90, 1557. ( 5 ) Borchardt, D. B.; Bauer, S. H. J. Chem. Phys. 1986.85, 4980. (6) Johnson, K. E.; Wharton, L.; Levy, D. H. J . Chem. Phys. 1978, 69, 7719 -. ... (7) Ewing, G. E. In Potential Energy Surfaces and Dynamics Calculations; Truhlar, D. G . , Ed.; Plenum: New York, 1981; p 75. ( 8 ) Huang, Z . S.; Miller, R. E. J. Chem. Phys. 1987,86, 6059. (9) Polik, W. F.; Moore, C. B.; Miller, W. H. J. Chem. Phys. 1988, 89, 3584.
(IO) Zimmerman, Th.; Koppel, H.; Cederbaum, L. S. J . Chem. Phys. 1989, 91, 3934.
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In recent researchI3+l4we have used classical trajectories to study the dynamics of the SN2 nucleophilic substitution reaction
c1,- + CH3Clb
-+
CI,CH3
+ Clb-
(1)
on a multidimensional analytic potential energy function,15derived in part from ab initio calculations. Selectively exciting the c-cb stretch normal mode of CH3Clbwas foundI3 to open up a direct mechanism for nucleophilic substitution, in which the reactive system is not temporarily trapped in either the minima for the CIa---CH3Clbor ClaCH3.-CIb- ion-dipole complex. The cross section for this direct substitution is several orders of magnitude larger than that for indirect substitution, which occurs without c-clb stretch excitation. For this latter indirect substitution, the reactive system becomes trapped in the ion-dipole minima.I4 The results reported here show that the lifetimes of the ion-dipole complexes do not agree with RRKM theory. The standard quasiclassical trajectory method'618 was used to simulate collisions between CI; and CH3Clb as a function of relative translational energy E,, and CH3Clbvibrational and rotational energy. The general chemical dynamics computer program VENUS^^ was used for the calculations. The time the reactive system remains trapped in the ion-dipole minima is taken as the difference between the time of the first C1; CH3Clb inner turning point and the last inner turning point for the dissociating
+
( I I ) Hase, W. L.; Cho, S.-W.; Lu, D.-h.; Swamy, K.N. Chem. Phys. 1989, 139, 1. (12) Polik, W. F.; Guyer, D. R.; Miller, W. H.; Moore, C. B. J . Chem. Phys. 1990, 92, 3471. (13) Vinde Linde, S.R.; Hase, W. L. J. Am. Chem. Soc. 1989,111,2349. (14) Vande Linde, S. R.; Hase, W. L. J . Chem. Phys., submitted for
publication. (15) Vande Linde, S. R.; Hase, W. L. J . Phys. Chem. 1990, 94, 2778. ( I 6) Porter, R. N.; Raff, L. M. In Dynamics of Molecular Collisions, Part B Miller, W. H., Ed.; Plenum: New York, 1976; p I . (17) Truhlar, D. G.; Muckerman, J. T. In Atom-Molecule Collision Theory; Bernstein, R. B., Ed.; Plenum: New York, 1979. (18) Hase, W. L.: Ludlow, D. M.; Wolf, R. J.; Schlick, T. J. Phys. Chem. 1981, 85, 958. (19) Hase. W. L.: Duchovic. R. J.; Lu. D.-h.; Swamv, K. N.; Vande Linde,
S. R.; Wolf, R. J. Quantum Chemistry Program Exchange, to be submitted for publication.
0 1990 American Chemical Societv
