(23) Kincheloe, R. D., Schemer, J. H. (to Cook Paint & Varnish Co.), U. S. Patent 3,644,566 (Feb 22, 1972). (24) Ledger, W. A. (to Coates Brothers and Co. Ltd.), U. S. Patent 3,491,065 (Jan 20, 1970). (25) hlarx, M., Hofmann, H. P., Klug, H. (to Badische Anilinund Soda-Fabrik Akt.), U. S. Patent 3,607,834 (Sept 21, 1971). (26) Marx, M., Pohlemann, H., Klug, H. (to Badische Anilinund Soda-Fabrik Akt.), U. S. Patent 3,644,316 (Feb 22, 1972). (27) Miller, L. I., Walus, A. K. (to E. I. du Pont de Nemours & Co.), U. S.Patent 3,585,160 (June 15, 1971). 128) Noll. W.. "Chemistrv and Technoloev of Silicones."
(30) Plueddemann, E. P. '(to 3,453,230 (July 1, 1969). (31) Purcell, R. F. (to Commercial Solvents Corp.), U. S. Patent 3,248,397 (April 26, 1966). (32) Richards, J. C., 111,Adhes. Age, 13,22 (1970).
(33) Saam, J. C. (to Dow Corning Corp.), U. S. Patent 3,655,633 (April 11, 1972). (34) Sample, J. H., Williams, C. H. (to Sherwin-Williams Co.), U. S. Patent 2,890,185 (June 9, 1959). (35) Solomon, D. H., "The Chemistry of Organic Film Formers," Wiley, New York, N. Y., 1967. (36) Speier, J. L., Keil, J. W., Gowdy, W. G. (to Dow Corning Corp.), U. S. Patent 3,440,302 (April 22, 1969). (37) Spotlights, 37, 2 (July 1969). (38) Staudinger, H., Huberle, M. (to Lonza Elektrizitatwerke and Chemische Fabriken A.-G.), German Patent 940,680 (March 22, 1956). (39) The Anchor Chemical Co. Ltd., British Patent 1,197,594 (July 8,1970). (40) Vandenburg J. T., J. Paint Techno/., 44, 77 (1972), and references thereh. RECEIVED for review February 2, 1973 ACCEPTED May 22, 1973
Ammonia Formation in the Nitric Oxide-Methane Reaction James C. Vartuli and Richard D. Gonzalez" Department of Chemistry, University of Rhode Island, Kingston, R. I . 02881
The catalytic reduction of nitric oxide using methane was studied in a static system over a silica-supported platinum catalyst. At 350", ammonia was the primary product while a t 250" nitrogen was predominant. It was found that an increase in nitric oxide concentration increased the rate of nitrogen formation a t 250" while a t 350" the ammonia formation was increased. The complete absence of nitrous oxide from the reaction products tends to exclude HNO as a precursor to ammonia formation. It was concluded that a catalyst dissociative chemisorption of hydrocarbons promotes ammonia formation.
V a r i a b l e s affecting the formation of ammonia in the catalytic removal of nitric oxide from automobile exhausts have recently become the subject of intense investigations. In a dual-bed catalytic muffler, nitrogen oxides are reduced over the first bed. Air is then added to the effluents and the unburned hydrocarbons and carbon monoxide are oxidized over the second bed. This approach appears to be chemically sound; however, analytical determinations of nitric oxide concentrations of exhaust effluents show relatively large concentrations of nitric oxide, whereas concentrations of nitric oxide in t h e effluents from the first bed are relatively low. The problem centers around the reoxidation of ammonia in the second stage. Within the usual temperature range of automobile exhausts, ammonia and nitrogen are the main products in the catalytic reduction of nitric oxide with hydrogen, and since the presence of ammonia is undesirable, the study of the variables which affect the formation of ammonia has become a pressing problem. Both hydrogen and unburned hydrocarbons are present in fairly large concentrations in
automobile exhausts so that reduction over the first bed can be carried out with little difficulty. Products formed in the catalytic reduction of nitric oxide with hydrogen have been the subject of many investigations, however relatively few studies have appeared in the literature in which hydrocarbons are used as a reducing agent. It is to this aspect of the problem that we addressed ourselves in this investigation. The variables which affect ammonia formation when hydrogen is used as a reducing agent have been carefully studied (-4yen and Peters, 1962; Jones, et al., 1971; Shelef and Gandhi, 1972). The production of ammonia depends critically on both temperature and nitric oxide concentration. A t temperatures below 200", besides nitrogen, the major nitrogenbearing product is nitrous oxide with only traces of ammonia being formed. The concentration of ammonia increases rapidly with temperature usually going through a maximum between 400 and 500", depending on the catalyst used, due to the decomposition of ammonia t o molecular nitrogen and hydrogen. Ammonia conversion also increases with decreasing Ind. Eng. Chem. Prod. Res. Develop., Vol. 12,
No. 3, 1973
171
Table 1. logarithms of Equilibrium Constants for the NO-CH, Reactions
sot
log K P Temp,
O K
400 500 600 700
Eq 1
Eq 2
149.08 118.73 98.51 84.22
267.24 209.69 171.43 144.82
50
nitric oxide concentration. This is in line with the idea proposed by Shelef and Gandhi (1972) that in order to form nitrogen, two nitrogen-containing species must be adsorbed on adjacent sites, a condition which can only be satisfied a t high nitric oxide concentrations. Kokes (1966) has shown that when an excess of hydrogen is used a t room temperature, ammonia is the main reduction product. I n fact it is the sole nitrogen-bearing product when the ratio of hydrogen to nitric oxide is 100: 1. Ammonia formation in the reduction of nitric oxide by hydrocarbons has not been detected by all investigators possibly due to varying catalytic systems. Xult and -lye11 (1971), using a series of paraffinic hydrocarbons from methane through octane, detected no ammonia formation on a bariumpromoted copper chromite catalyst in the temperature range 225-525'. This is in agreement with Malling (1963) using methane and Sotoodehnia-Korraai and Kobe (1970) using ethylene as reducing agents. Both obtained only nitrogen as the reduction product. On the other hand, Anderson, et al. (1961), aiid Jaros and Krizek (1967) detected ammonia using methane over noble metal supported catalysts. Other minor products that have been reported in past investigations using both methane and hydrogen include hydroxylamine and hydrocyanic acid (Buttemorth and Partington, 1930). Carbon dioxide and water are important products when methane is used as a reducing gas m t h only traces of carbon monoxide being reported. Thermodynamic Considerations
The two equations governing ammonia and nitrogen formation as reported by Anderson, et al. (1961), are
5CH4
+ 8x0 + 2H20 e 5CO? + 8NH3
(2)
Table I ("JANAF Thermochemical Tables," 1965) lists log K , as a function of temperature for both equations. I t is clear from the data that both reactions are thermodynamically favorable. Experimental Section
Apparatus. -5 static catalytic reactor consisting of a U-shaped Pyrex tube was incorporated into the reaction loop of a conventional high-vacuum system capable of Torr. T h e attaining a n ultimate pressure of 1 X reactor was made from 25-mm Pyrex tubing and was a b o u t 10-cm long. T h e side arms consisted of 10-mm Pyrex tubing connected to the vacuum system through two vacuum stopcocks. The catalyst was held in place by two quartz wool "plugs" which had previously been treated in a boiling solution of 9 -11 nitric acid and washed several times in deionized water. The purpose of this treatment was to remove traces of hydrocarbons present in the quartz wool. The entire reactor was placed in a n oven and the temperature 172 Ind.
Eng. Chem. Prod. Res. Develop., Vol. 12,
No. 3, 1973
2
4
TIME
6
8
- HOURS
1 0 1 2
Figure 1. Ratio of NH3 and Nf to concentration of function of time at 250"
NO
as a
was nionitored by a thermocouple inserted in a thermocouple well made in the catalytic bed. The temperature could be cont,rolledto within 5" over the entire catalytic bed. Materials. The catalyst was prepared b y impregnating Cab-o-Sil, Grade >I-8, obtained from the Cabot Corp., Boston, Xass., with chloroplatinic acid obtained from t h e Engelhard Industries, Semark, N . J. Spark spectra analysis of Cab-o-Sil, a high-grade silica, indicated a total metallic impurity content of less t h a n 100 ppm. The chloroplatinic acid solution was prepared by dissolving the appropriate amount of chloroplatinic acid in deionized iyater. The catalyst used in this st'udy had a nominal 2 7 , platinum content. The mixture was dried in an oven a t 110" and stirred regularly to retain uniformity. The dried mixture was then ground and screened to 38-60 mesh and a 2-g aliquot was sealed into the reactor. The metal surface area was determined by H2-Dz exchaiige.oThe particle size was determined to be between 10 and 30 A. Ultrapure methane (99.977,) and technical grade nit'ric oxide (98.5)) were obtained from the Xatheson Co. The methane \\-as passed through a Dry Ice trap prior to storage in the vacuum system. The nitric oxide mas purified using the freeze-thaw technique. It was found that a t least five thaws were necessary to remove the nitrogen. Prior to storage, the nitric oxide was subjected to a vacuum distillation treatment in which only the middle cut was retained. Carbon dioxide and ammonia used in the analytical calibrat'ions were obtained locally aiid were purified by standard vacuum diatillation techniques. Regular tank hydrogen was used in the reduction of the catalyst. It was purified by passing it sequentially through a deoxo unit, t'o remove traces of oxygen, and then a molecular sieve and a liquid nitrogen trap to remove water. A11 gases were maintained a t better than 99% purity by periodically checking them with a mass spectrometer. Techniques and Analytical Methods. The catalyst was briefly outgassed a t room temperature and then the temperature was gradually increased to 300" in flowing hydrogen (150 cm3/min). Reduct'ion time a t 300" was for 4 hr. The hydrogen was then evacuated and the catalyst gradually heated t o 400" and outgassed for 4 hr a t that' temperature. T h e catalyst was then cooled to t h e temperature of the run. Separate predetermined amounts of nitric oxide and methane were introduced into the reaction cell a t the start of the reaction. I n an attempt to understand the importance of Fvater in the production of ammonia, runs were made with and without a Dry Ice trap in the reaction loop.
I
70t
I
1
* N2/NO 3500c
70 REACTION TIME
' NH3/NO
t.
6o 5ot.
< 40 O
4
I
* N?/NO
12 HOURS
140
COZ/NO
..
L
30.
ae
20..
lot.
I.
ZOO
TIME-HOURS
250
35C
300
TEMPERATURE OC
Figure 2. Ratio of NH3 and Nz to concentration of function of time a t 350"
NO
as a
It was felt that if water was a reactant in the formation of ammonia, the concentrations of the products \vould be different if the Dry Ice trap was present. The reaction-h were stopped by placing a liquid nitrogen trap in the reaction loop to remove condensable gases. Soncondensable gases were transferred to a measured volume using a Toepler pump. The reactor was then sealed off and the remaining condensable gases were pumped into the measured volume. -411 aliquot sample was taken and analyzed mass spectrometrically on a CEC (Ilu Pont) Model 21-101 mass spectrometer equipped with ail electron mult,iplier. Calibration curves relating the ratios of mass spectral lines of the five major compoiients to their absolute values were obtained and the raw data were programmed and processed in a digital computer. Positive product identification was obtained using a Perkin-Elmer hlodel 521 Grating infrared spectrophotometer. Ainalytical data for a typical 12-hr run are listed in the I p p e n d i s . Results
Effect of Time. A t constant initial coiicentrat'ions of nit'ric oside and methane, several runs lvere made t o investigate the effect of per cent conversion on t h e distribution of products. T h e total initial pressure in t h e reactor was approximately 120 Torr. Initial partial pressures of nitric oxide and methane n-ere approximately 53 a n d 6 i Torr, respectmirely.T h e result's are shown i n Figures 1 a n d 2. Conversion of nitric oxide, based on initial concentrations] increased from 607, for a 2-hr run to over 807, for the 12-hr run. Of significance here is that the concentration of ammonia does not go through a maximum and decrease with time. This implies that' the ammonia does not decompose or react further witli nitric oside. This is rather surprising in view of the results of Otto. et ai.(1970, l g i l ) , who have shown that ammonia reacts readily with nitric oxide over a supported platinum catalyst in the 200-250" temperature range. TTe' confirmed these results on a blank run in which nitric oside and ammonia were allowed to react over the catalyst for a period of 4 hr. .Zn analysis of the products showed large quantities of nitrogen and nitrous oxide in agreement with the results of Otto, el al. (1970, 1 9 i l ) . Sitrous oside was never observed as a product in the reduction with methane; therefore] we must' surmise that under the conditions of our esperiment, ammonia must be excluded from the surface either by methane or by nitric oxide. The concentration of a:mmonia always increased with time. This would preclude an:; secondary reactions involving ammonia. Effect of Temperature. T h e reaction was st,udied over a 200-350" temperature range a t a constant initial concen-
Figure 3.Ratio of products to initial concentration of NO as a function of temperature
70.REACTION TIME
*
NdNO
'
C02/NO NH3/NO
12 HOURS
60.
200
250 300 TEMPERATURE ' C
35C
Figure 4. Ratio of products to initial concentration of a function of temperature; Dry Ice trap included
Table
400 500
as
II. logarithms of Equilibrium Constants for CO
Temp,
NO
O K
+ Hz e Hz +
log K,
3.19 2.13
Temp,
COz O K
600 '700
log K P
1.46 1.12
'
t'ration ratio of nitric oside to methane of 4 : 5 . Reaction times for all runs were 12 hr. The results are shown i n Figure 3. I t is apparent' t h a t high temperatures favor t h e formation of ammonia a t the expense of a decrease in the amount of nitrogen produced. Figure 4 sho\\-s an ident,ical set of runs with a Dry Ice trap placed in the reaction loop to remove water from the reaction. The effect is to enhance nitrogen formation, especially a t the higher temperatures, a t the expense of ammonia formation. Carbon dioside formation was also depressed in the presence of the Dry Ice trap. Lower methane conyersions, based on initial methane concentrations, coincided with this observation. A glance a t Table I1 sho~vvs that a t the temperatures of this investigation, the equilibrium for the water gas shift reaction is thermodynamically favorable for the production of additional hydrogen. I n fact Klimisch and Barnes (1972) have shown that on a copperpromoted chromia-alumina catalyst a t temperatures over 200' the water gas shift reaction provides a n additional source of hydrogen and enhances rates of ammonia formation. Our reInd. Eng. Chem. Prod. Res. Develop., Vol. 1 2 , No. 3, 1973
173
"/": 70' 260 * C
* N ~ N O
' NHI/NO
30
60.
100 150 NO C O N C E N T R A T I O N
.-.-'
I
I .
200
50
,UM/L
Figure 5. Ratio of NH3 and Nz to concentration of N O as a function of N O concentration at 250"
50 CH4
100 I50 CONCENTRATION
20C
/IWL
Figure 7. Ratio of NH3 and Nz to concentration of N O as a function of CH4 concentration at 250' ~
~~
* 6
Nzh0 NH3/NO
6ot
* N ~ N O "$/NO
101
50 NO
100 150 CONCENTRATION
200 /M/L
CH4 CONCENTRATION
pM/L
Figure 6. Ratio of NH3 and Nz to concentration of N O as a function of NO concentration at 350"
Figure 8. Ratio of NH3 and Nz to concentration of N O as a function of CH4concentration at 350'
sults using the Dry Ice trap also suggest that the water gas shift reaction may be a n important source of hydrogen for ammonia formation. In any case it is clear that water is a cocatalyst in the formation of ammonia. Effect of Reactant Concentrations. Both methane and nitric oxide concentrations were varied in a n a t t e m p t t o study concentration effects on the rate of ammonia and nitrogen formation. Two temperatures (250 and 350') were selected for this study. Four-hour reaction times were used in all runs. T h e effect of increasing the nitric oxide concentration is shown in Figures 5 and 6. A t 250' there is a noticeable increase i n t h e rate of nitrogen formation while a t 350' t h e nitrogen formation is almost independent of nitric oxide Concentration. T h e rate of ammonia formation decreases with increasing nitric oxide concentration at 250', this trend being reversed a t 350'. Similar data showing the effect of increasing methane concentration are shown in Figures 7 and 8. A t 250' both nitrogen and ammonia formations increase with increasing methane concentration while a t 350°, nitrogen formation goes through a maximum before decreasing a t higher methane concentrations.
bonded together), two molecules of nitric oxide must be adsorbed on adjacent surface sites. If this is true, the rate of nitrogen formation should increase with increased nitric oxide pressure provided of course that reaction occurs between two adjacent chemisorbed nitric oxide molecules. Our results clearly bear this out with the exception of the hightemperature run in which nitrogen formation goes through a maximum and then decreases. We feel this is in part due to a decrease in nitric oxide adsorption a t the higher temperature in addition to a n increase in the dissociative adsorption of methane at 350'. Hydrogen atoms compete with nitric oxide molecules for surface sites under these conditions with a resulting decrease in the rate of formation of unfixed nitrogen. One would then expect a n increase in the rate of ammonia formation with increasing methane concentration since presumably some kind of half-hydrogenated intermediate is necessary for ammonia production. This is clearly shown in Figures 7 and 8. The rate of ammonia formation changes rather slowly with increasing methane pressure a t 250" but shows a sharp increase a t 350'. This is due to the relatively small amount of dissociative methane adsorption a t 250' compared to extensive surface dissociation of methane a t 350". An increase in the rate of ammonia formation with temperature provides additional evidence that dissociative methane chemisorption is necessary for ammonia formation. Kokes (1966) has shown that even a t room temperature a high hydrogen to nitric oxide ratio results in extensive ammonia formation.
Discussion
Our results are in agreement with several features of the nitric oxide-hydrogen reaction scheme. Shelef and Gandhi (1972) have suggested that in order to obtain unfixed nitrogen (defined here as nitrogen species with two nitrogen atoms 174 Ind. Eng. Chem. Prod. Res. Develop., Vol. 12, No. 3, 1973
The time-dependent runs again bear out the importance of the methane surface coverage. Ammonia was found not to be a reducing agent for nitric oxide in the presence of methane while in the absence of methane i t can reduce nitric oxide effectively. We feel t h a t in the methane reduction, ammonia cannot compete effectively for suface sites. Once i t i t is formed, i t is displaced from the surface inhibiting any reaction with nitric oxide. T h e complete absence of nitrous oxide in the reduction products was rather surprising. T h a t nitrous oxide was perhaps a very reactive intermediate in the formation of nitrogen or ammonia was ruled out as runs in which nitrous oxide was treated with methane and hydrogen under identical conditions yielded no ammonia. This not only rules out the possibility of nitrous oxide being a n intermediate but also rejects the idea of a common intermediate for nitrous oxide and ammonia formation. The formation of H N O as a precursor to nitrous oxide formation in the low-temperature reduction of nitric oxide with hydrogen has been suggested b y Kokes (1966) and Gonzalez and hudette (1970). Kokes (1966) has also suggested that the reaction. HKO
5 NHs + HzO
might occur t o account for the ammonia production. The complete absence of nitrous oxide from the reduction products should tend to exclude €IN0 as a n intermediate in the hightemperature reduction of nitric oxide with methane, although trace amounts of N 2 0 are formed when hydrogen is used as the reducing gas. The identity of reaction intermediates is presently being investigated in this laboratory. R e realize t h a t extrapolation of results obtained on a static system with much greater gas concentration to the dynamic transient flow system in the automobile is dangerous and must be made with care. The choice of a static reactor was made as we wanted to draw a direct comparison between the results of nitric oxide reduction with methane and those obtained by Shelef and Gandhi (1972) using hydrogen. We also feel that the identity of reaction intermediates could best be studied spectroscopically under static conditions and could be compared more productively to the product distribution in the catalytic reduction under similar conditions. I n conclusion, i t is evident that if ammonia formation is to be minimized in the design of a n effective catalytic converter, a catalyst which minimizes the dissociative chemisorption of hydrocarbons would be a good choice due to the relatively large amounts of unburned hydrocarbons in the automobile emissions.
Appendix
Analytical data for a typical 12-hr run included the following raw data: run number, 28; temperature of run, 350'; total number of micromoles in product, 490.65. The composition of reactants and products (in micromoles) is shown below.
Nitric oxide Methane Nitrogen Ammonia Carbon dioxide
Reactants
Products
286.31 359.34
4.35 170.81
48.34 166.42 100.74
Material Balance. T h e nitrogen material balance agreed t o within 8%. When the catalyst was heated to 400' for 4 hr, 60 of t h e missing 89 pmol of carbon products and 10 of t h e missing 23 pmol of nitrogen products were recovered. Only SZ,NH8, Con, and CH4 were present in this recovered fraction. An appropriate correction was made to take this aliquot into account. With this correction the material balance was within for nitrogen and 8% for carbon. T h e accuracy of the mass spectral measurements was 5%. Literature Cited
Anderson, H. C., Green, W. J., Steele, D. R., Ind. Eng. Chem., 53, 199 (1961). Ault, J. W., Ayen, R. J., A I C h E J., 17, 265 (1971). Ayen, R. J., Peters, M.S.,Ind. Eng. Chem., Process Des. Develop., 1. 204 (1962). Buiterworth, A. J., Partington, J. R., Trans. Faraday SOC.,26, 144 (1930). Gonzalez, R. D., Audette, D. E., J . Catal., 17, 116 (1970). "JANAF Thermochemical Tables," D. Stull, Ed., Dow Chemical Co., Midland, Mich., 1965. Jaros, S., Krizek, J.; Chem.'Prum., 17, 581 (1967). Jones. S. H.. Kummer. J. T.. Otto. K.. Shelef., M.., Weaver. E. E.. Environ. Sci. Technol.. 5.'790 (1971). Klimish, R. L., Barnes, G. 'J., ibid., 6,'543 (1972). Kokes, R. J., J . Phys. Chem., 70, 296 (1966). Malling, G. F., M.S.Thesis, University of Illinois, Urbana, Ill., 1462
O G i r K . , Shelef, M., Kummer, J. T., J . Phys. Chem., 74, 2690
(1970). Otto, K., Shelef, M., Kummer, J. T., ibid., 75, 875 (1971). Shelef, M., Gandhi, H. S., Ind. Eng. Chem., Prod. Res. Develop., 11, 2 (1972). Sotoodehnia-Korrani, A,, Nobe, K., Ind. Eng. Chem., Process Des. Develop., 9, 455 (1970).
RECEIVED for review September 11, 1972 ACCEPTED May 22, 1973 J. C. V. is grateful for the award of an NDEA Graduate Fellowship.
Ind. Eng. Chem. Prod. Res. Develop., Vol. 12, No. 3, 1973
175
