mate agreement with the results obtained by Bond et al. (4) in a similar apparatus. However, the -325-mesh coal was still not completely devolatilized, although it is estimated that coal temperatures of about 2900' to 6300" C. would have been reached if thermal equilibrium had been attained. Total yields and carbon and ash recoveries were low in experiments 40 and 41 (Table 11), possibly because of losses of extremely fine solid residue that filtered through the recovery system. In all three experiments, the recoveries of hydrogen and oxygen were high, and in experiments 38 and 40, the yields of carbon monoxide obtained would not be possible even if all the oxygen in the coal had been converted to carbon monoxide. I t is likely that some cooling water leaked into the plasma. The reaction of this water with coal or coal decomposition products would account for the high hydrogen recoveries and the high yields of carbon monoxide. The predominance of acetylene in the gaseous products shown in Tables I and I1 and the absence of liquid products indicate that, although the coal was heated only to temperatures high enough for partial devolatilization to occur, the volatiles released attained much higher temperatures.
obtained in yields as high as 1 5 iceight Yc of maf coal. HOWever, the coal was not heated to temperatures high enough for complete devolatilization. To obtain higher coal temperature and yields of gaseous products, future ivork by the bureau will be concerned Lcith changes in plasma generator design and;'or operating techniques that may result in more efficient utilization of the heat available in plasmas. Higher coal temperatures should be attainable by feeding smaller coal particles and increasing the residence time of the coal in the plasmas.
literature Cited
(1) hmmann, P. R., et al.. Chem. E n g . Progr. 60, 52-7 (1964). (2) Baddour; R. F., Blanchet, J. L., ISD. E s c . CHEM.PROCESS
DESIGK DEVELOP. 3,258-66 (1964). (3) Baddour, R. F.. Iwasyk, J. M., Ibid., 1, 169-76 (1962). (4) Bond, R. L . >et al.. .Vature 200, 1313-14 (Dec. 28, 1964). (5) Freeman. M. P., Skrivan, 3. F., Hydrocarbon Process. Petrol. Rejner 41, 124-8 (.August 1962). (6) Leutner, H. LV., Stokes, C. S.: Znd. Eng. Chem. 53, 341-2 (1961).
RECEIVED for review January 21. 1965 ACCEPTEDAugust 13, 1965
Conclusions
Acetylene was the principal hydrocarbon gas produced when hvab coal was injected into argon plasmas and was
Division of Fuel Chemistry, 148th Meeting, ACS, Chicago. Ill., September 1964.
CATALYTIC OXIDATION OF METHANE R E l J I M E Z A K I A N D C H A R L E S C . W A T S O N Department of Chemical Engineering, Yniuersity of Wisconsin, Madison, Wis.
The catalytic oxidation of methane was investigated b y experiments conducted with a flow-type integral reactor. The reaction temperature ranged from 320" to 380" C. a t a pressure of 1 atm. The catalyst employed was palladium on alumina. The experimental results indicated that the reaction rate could be controlled b y a surface reaction in which the gaseous methane reacts with the adsorbed oxygen, producing adsorbed carbon dioxide and adsorbed water.
c
the total catalytic oxidation of light hydrocarbons has become increasingly important with respect to the problem of combustion removal of hydrocarbons in process tail gases such as the exhausts of internal combustion engines. The hydrocarbons from the exhausts have been said to contribute to the generation of smog. Combustion elimination of hydrocarbons is, therefore, one of the challenging problems to be solved for the prevention of smog formation. While recent development of automotive afterburner devices may be a sufficient remedy a t least to defer this problem, industrial exhaust gases containing unburned hydrocarbons and other organic contaminants will continue as a field for application of catalytic combustion processes. Since methane is a representative compound in the light hydrocarbon family and is the most difficult member of this family to oxidize, it was selected for this investigation. Many investigators ( 7 , 7-70) have reported the experimental results of the catalytic oxidation of methane. Their results are, however, useful only for the selection of a suitable oxidation catalyst. For use in equipment design and development studies, the results of kinetics research ought to be expressed in terms of practical models, as realistic as possible, to facilitate extrapURRENTLY
62
I & E C PROCESS D E S I G N A N D D E V E L O P M E N T
olation. Especially useful relations were developed by Langmuir and Hinshelwood, and are commonly called LangmuirHinshelwood models. These models, or reaction rate equations, result from a simple kinetic treatment in which a t least one reactant species must react chemically with one or more uniform sites upon the surface of a solid catalyst. Techniques of applying these concepts to the analysis of rate data of industrial importance were developed by Hougen and LYatson ( 4 ) . and the resulting rate expressions are sometimes also referred to as Hougen-Watson models. In the present work the kinetics of the total oxidation of methane was investigated by consideration of various possible Langmuir-Hinshelwood type models and a satisfactory rate equation was developed. Equipment
The experimental equipment is described in detail by Mezaki ( 5 ) . Figure 1 shows a schematic flow diagram of the entire equipment. Air, nitrogen, carbon dioxide, and methane were measured through previously calibrated flowmeters into the reactor
system. All gas flows were controlled by pressure-regulating valves and needle valves. Commercially available cylinder gases and compressed air from the laboratory lines were used. Water vapor was introduced into the air-nitrogen mixture, when desired, by passing the gas mixture through a humidification column. Water vapor fraction was regulated by changing the water temperature. The line between the humidification column and the reactor was heated so that condensation of water vapor was avoided. T h e humidification column, 15-cm. i.d. glass tube, 5 feet long, contained a 10-inch layer of water, heated by a n immersion heater. Wet- and dry-bulb thermometers were mounted a t the outlet of the column. The reactor was a IO-mm. i.d. quartz tube, 30 inches long, which consisted essentially of two sections: preheater and catalyst zone. T h e preheater section was a 25-inch empty tube mountea in an electric furnace. I n the catalyst zone the catalyst was supported on a '/(-inch layer of quartz chips. This zone was also surrounded by an auxiliary furnace. To provide precise temperature control of the reactor, compressed air was passed through an air jacket made of a concentric 1inch 0.d. quartz tube. Two thermocouple wells \cere mounted a t the top and bottom of the catalyst layer. The capacities of the main and auxiliary furnaces were 1000 and 750 watts, respectively. Prior to analysis of the reactor exit gas, water vapor and carbon dioxide \cere removed. Drierite (anhydrous C a S 0 4 ) was used to remove water. Carbon dioxide was adsorbed by 2x7 KOH solution and by Ascarite (sodium hydride). Gases
Matheson C.P. grade methane was used, stated by the supplier to contain 99.05% methane, 0.12y0 ethane, 0.2070 carbon dioxide, 0.60% nitrogen, 0.03% propane, and 50 p.p.m. oxygen. Nitrogen was supplied by the Linde Nitrogen Co. This high purity nitrogen contained a trace amount of argon and 5 to 9 grains of water per 1000 cu. feet a t STP. Carbon dioxide was purchased from the Midwest Carbonic Co. The purity was 99.97,. Catalyst
The 0.57, palladium catalyst was supplied by Engelhard Industries, Inc. Its physical properties were : bulk density, 56 pounds per cu. foot; surface area, 120 sq. meters per gram; x-ray pattern, a gamma-type alumina; physical form, X ' / g inch extruded cylinders. The '/'g-inch particles \cere ground and screened by a U. S. sieve series set. The fraction between 12- and 14-mesh was employed for this investigation. In the catalytic oxidation of methane, it was found that the activity of palladium dropped during the operation. T h e decline of activity was suspected to be caused mainly by the adsorption of water which was contained in the feed gas stream and was produced by the oxidation reaction. The catalyst lost considerable activity during the 0- to 20-hour service periods in the presence of water vapor. After these periods the activity reached a fairly constant value. All experimental data were taken in the constant region of activity.
HEUUH
CYUNER
Figure 1 .
The desired amount of catalyst was charged into the reactor, which was then mounted in the two electric furnaces. The main furnace was turned on and the temperature adjusted with the Variac. During the heating period, cooling air was passed through the cooling jacket of the reactor until the reactor system reached the operating temperature. T h e heater in the humidification column was turned on if water was to be introduced in the reactant gas. T h e pipeline heater was also turned on. During the heating period, the gas chromatograph was calibrated. After the reactor system had reached a n operating temperature, the desired flow rate of air was passed through the reactor and the cooling air was turned off. After the temperature of the catalyst bed reached a constant level, methane was added to the reaction system. Nitrogen and carbon dioxide were added to the system, if required.
Flow diagram
T h e mixture of air and nitrogen was passed through the humidification column. Maximum temperature deviation through the catalytic bed was +4.5' C. The entire system was allowed to operate for approximately 1 hour. When steady state was reached, the temperature, pressure. f l o ~rates of the gases, and all other necessary data were recorded, and a sample of the reaction product was taken. Analyses. .4 Perkin-Elmer Model 154-C L'apor Fractometer was employed for the analyses of both reactant and product gases. The specifications of the instrument and its operating conditions were: sample volume, 5 cc., carrier gas. helium; type of column, J-column; column length. 2 meters; column pressure, 6 p.s.i.; column temperature, 40' C . ; recorder. Wheelco 0- to 5-mv. Results
The experimental data are summarized in Table I.
(Tables
I and 111 are available from the American Documentation Institute.) All gas compositions were expressed in volume fractions and conversions were calculated on the basis of methane converted. The data were correlated by using Langmuir-Hinshelwood type rate equations. I n an initial survey, 84 such models were linearized and fitted to the experimental rate data by linear least squares. For all models in which adsorption of reactants or desorption of products was the rate-controlling step, this fitting resulted in two or more negative constants. Estimated confidence intervals for each rate constant did not include zero a t 95% level. Therefore, we concluded that the estimated negative constants were significantly different from zero. This led to the rejection of these models. Assuming that the surface reaction is the controlling step, one can postulate various models and derive
Table II. Procedure
il
CHROMATOGRAPH
SOLUTION
Matrix of Surface Reaction Models"
CH,
0 2
1
0 0
GO1 H20 Site 1 0 1 2 1 0 1 1 3 0 1 1 0 1 4 1 1 1 1 2 5 0 1 0 2 1 0 6 1 1 2 0 7 1 1 1 0 2 8 1 1 0 1 2 9 1 1 0 0 2 10 1 0 1 1 2 11 0 1 1 1 2 12 1 0 0 0 1 0 13 1 0 0 1 a 0 and 1. Gaseous and adsorbed state, respectively. Sites 7 and 2. Single and dual site, respectively.
Models 1
VOL. 5
NO. 1
JANUARY 1 9 6 6
63
OBSERVED
Figure 2.
CONVERSION ( %)
Observed and calculated values of conversions
the rate equations. The matrix in Table I1 specifies the assumptions made in developing these 13 models. The rate expressions themselves are shown in Table 111. The thermodynamic equilibrium constant, K , is extremely large a t the temperature investigated (approximately 1070 a t 600' K.). The rate equations, therefore, can be simplified and readily transformed into linear forms. The hypothetical rate equation of model 11 for example, gives
rate equations were easily obtained by eliminating the nitrogen terms in the previously described rate equations. Linear estimation gave three good-fit models. The residual sums of squares for these three models were significantly smaller than for the other models; in fact, they were in the range 0.01 to 0.10 times the sums of squares of the latter. Their parameters are listed in Table IV. The values thus obtained provide usable initial guesses for the rate constants in a nonlinear least
Introducing the following stoichiometry of methane oxidation reaction, aCEl = a(1 - x)n ao, = ( b - 2 ax)?r ace, = (6 ax)r a H I O = (d 2 ax)n
squares calculation. The elimination of the need for differential reaction rates would be desirable for the most precise evaluation of the kinetic parameters. This can be accomplished by the analytical integration of Equation 1, the resulting expression being shown as Equation 5 in Table V (on file with the American Documentation Institute). S o w , a nonlinear least squares calculation (2, 3 ) was applied for Equation 5, evaluating the rate parameters. The calculated results for model 11 are shown in Table VI. T h e same procedure was used also for models 5 and 6 . For model 5 the ratio of mean square of the lack of fit to the estimated mean square error is 7.5. The upper 5YGpoint for the F-distribution with 4 and 17 degrees of freedom is 5.8. This indicated that model 5 was not adequate. For the same reason model 6 was discarded. The ratio of model 11 was 1.0. which showed that the model is adequate. since F4,1s(0.95) = 5.8. Figure 2 represents the comparison of the observed and calculated values of conversions. Figure 3 shows the conversion-space time correlation of selected data. A rather large value was obtained for the adsorption equilibrium constant of oxygen. The adsorption constant of the oxygen-palladium system was determined by Parravano and Bartner (6) by equilibrium adsorption methods. Their experimental results have indicated that the equilibrium constant was extremely large (1.3 X a t 523' K.). The surface properties of palladium used in Parravano's investigation were not identical with those of the authors The adsorption equilibrium constants in the postulated rate expressions refer to the adsorption sites effective in catalysis only, which
~
+
+
axz = en
we derived the linearized form of the rate equation.
(G
+
axh
+
KH20
(d
3
dLS(S
-
+ 2 ax)n
t
1)ksd'oZ KN,
3
ea
(2)
d L S ( S - l)k8RKO2 A linear least squares technique was used to estimate all parameters in the rate equations. The 24 data points a t 320' C. were used for the discrimination of reaction models. In this regression all adsorption equilibrium constants of nitrogen gas were found to be very close to zero. Therefore, it was assumed that nitrogen would not be adsorbed by the catalyst in the range of the experimental conditions. This assumption generated another 13 hypothetical reaction models. Their 64
IbEC PROCESS DESIGN A N D DEVELOPMENT
Table IV. IVO.
G o d - F i t Models and Their Rate Parameters at Experimental Temperature 320’ C.
and adsorbed oxygen. Integral kinetic equations were employcd for correlating experimental data. Nomenclature
of
Modi1
Rate Paramtiers
(LS)ksR koz KH,O
5
= = =
a
2,000 2,150 13.8
%H4
ace,
Kc6, KH,O
= =
54 73
d e
= =
F
=
kSR
=
K Table VI.
320 350 380 1
Calculated Values of Rate Parametersa
79,700 120,000 151,000
Kco, 54 2 50 0 36 3
45,300 30,000 23,200
=
KCH4
KH~O
86.7 60 4 46 5
Kc,,
=
KB20 =
pp
= -
r R
= =
T
=
M’
= =
KN,
Numerzral I d u e s rounded off to threr szgnzjrrrrl: figures
s are genera!ly fewer than those effective for total adsorpticin, as measured in adsorption equilibrium experiments. Thus, the valuer we obtained for Ko2 are not inconsistent with P a m vano’s results. Finally the Arrhenius correlations were oLtained for the four reaction paramrters. The equations are as follows : 36.35
l n o = __ R
R?‘
6.71 In KO, = -
R
In KBzO= -
- 8228
+ 8631 RT
4.66
R
--
(7)
8006
‘-E
No simple method is available to prove that this reaction model is incorrect; the reaction model postulated is not totally unrealistic, and the rate equation derived on the basis of the model fits experimental data reasonably well. These facts lead to thc following conclusions. Conclusions
Data for catalytic total oxidation of methane over a palladium-alumina catalyst were well correlated by a LangmuirHinshelwood type rate model, in which the rate-controlling step could be the surface reaction between gaseous methane
partial pressure of caibon dioxide water vapor oxygen nitrogen oxygen in feed carbon dioxide in feed volume fractic,it of water vapor in feed volume fractiu I of nitrogen in feed mass flow rate of feed reaction rate constant thermodynai,iic equilibrium constant adaurption t qiiilibrium constant of methane on catalytical!\ active sites adsul p i o n equilibrium constant of carbon dioxide adsLrption equilibrium constant of N a t a vapor adsorption equilibrium constant of nitrogen adsorption equilibrium cciistant of oxygen total concentration of active sites reaction rate gas cullJtant numbrr of equidistant aciive centers adjacent to each otter . absolu-r temperature weigh of catalyst conversion of methane total pressure of reaction systcrn (Ls)(s- 1) ksn
c
b
11
=
= partial pres>ure of = partial pre1,ut-e of = partial pre>,ure of = volume frac Lion uf = volume frac 1 0 1 1 of
aH,O a,, aN2
6
= volume fraction of methane in feed = partial pressure of methane
=
x 7
o
=
r
=
=
Acknowledgment
’The authors express their jratitude io 0. A. Hougen for suggesting this project. They are grateful to the Wisconsin Alumni Research Foundation and the National Science Foundation for financial support (NSF-GP 2755). Literature Cited (1) Andersoil, K. B., Stein. k.C., Feenan, J. J., Hofer, L. J . E., Ind. Enp. C ; e m . 53, 809 (1961).
(2) Booth, G. W., Peterson T. I., L. B. M. Share Program, Pa., “Non-Linear Estimation,” 687 WL NL1 (1958). (3) Bus, G. E. P., Ann. 1%‘. Y.Acad. Sci. 86,792 (1960). (4) Hoilgen, 0.A., \\’atson, K. M., “Chemical Process Principles,” Part 111, Wiley, New York, 1950. (5) Mezaki, I
