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Kinetics of the reaction of hydrogen with thin films of carbon
KINETICS OF THE REACTION OF HYDROGEN WITH THIN FILMS OF CARBON J.-R. CAO and M. H. BACK Chemistry Department, University of Ottawa, Ottawa, Canada KIN 984 (Receioed 28 December 1981) Abstract-The kinetics of the reaction of hydrogen with thin films of carbon has been studied over the temperature range 870-1150K and at pressures of hydrogen from 50-3OOTorr(6.7-40KPa). Thin films of carbon of average thickness about 30 nm were deposited on the surface of a quartz reactor by the pyrolysis of methane at I IO0K and the kinetics were studied in a static system. The products of the reaction were methane, ethane and ethylene. formed in successive hydrogenation steps, which in the low temperature region occurred largely on the surface of the carbon. In this region the activation energy of the rate of formation of methane was 6.5 kcalimole. At temperatures above about 1050K the thermal dissociation of hydrogen provided a source of radicals which caused a rapid increase in the rate of hydrogenation. both heterogeneous and homogeneous, giving an activation energy for the rate of formation of methane of 51 kcal/mole. A self-inhibition was observed, probably caused by a heteroeeneous .nolvmerization reaction leadine. to the formation of higher molecular weight products which . remained adsorbed on the surface.
1. INTRODUCTION
The gasification of carbon by hydrogen is a deceptively simple reaction. In most studies methane is the main, and in some cases the only, product reported but in spite of this simple stoichiometry the details of the mechanism by which the hydrogenation occurs are not well understood. The reaction has received more attention in recent years because of the interest in the production of substitute natural gas. Although for this purpose most of the studies involve the reaction with coal and coal chars [I, 21, or the reaction with carbon catalyzed by metals[3,4], the reaction with pure carbon provides fundamental information necessary to the understanding of the hydrogenation process. Early work has been reviewed by Walker, Rusinko and Austin[S] and more recent work by Lewis[6] and by Krakowski and Olander[7]. Among the first suggestions concerning the mechanism was that of Zielke and Gorin[8] in which successive hydrogenation of neigh-
bouring HLCLH
edge carbon H , -C/H, ‘H
atoms occurred
giving =C-,
and finally methane. This basic
mechanism has been modified but not effectively changed by most subsequent work[1,9-121. One of the difficulties in formulating a mechanism and a rate equation that could have general applicability is the wide variation in the reported kinetic parameters of the rate. There appears to be no general agreement concerning the order of the rate with respect to hydrogen and, probably as a consequence, no agreement on activation energy. This is partly due to the variation in reactivity of the carbon used as the reactant and partly due to the wide range of conditions of temperature, pressure and conversion used in the study of the reaction. It is not
surprising that attempts to derive a rate equation for the reaction have had limited applicability. Recently an extensive analysis of data by Shaw[l3] has shown that the kinetics of the formation and decomposition of methane in the presence of char and hydrogen may be described by a two-step hydrogenation process on the carbon surface leading directly to methane. None of the mechanisms, however, consider the possibility of the formation of products intermediate to the formation of methane, for example, Cz hydrocarbons, and the consequence of their desorption from the surface, although some studies have reported the formation of higher hydrocarbon products[9, 141, particularly at higher temperatures(lS]. Such products would be less stable than methane and would rapidly disappear in secondary reactions, possibly leading to radical reactions in the gas phase. The purpose of the present study was to search for the formation of such products, if possible to distinguish between primary and secondary products, and to consider the related question of the importance of hydrogenation in the gas phase as distinct from the processes occurring on the surface. From the results of the study a mechanism for the hydrogenation has been proposed and the importance of radical reactions in the gas phase has been demonstrated. We have recently developed a technique for the study of reaction of gases with carbon which offers some advantages for kinetic studies [ 161.A thin film of carbon is deposited on the surface of the reaction vessel by the pyrolysis of methane. The concentration of carbon on the surface, in the range lo-‘-lO-~^ gem-‘. may be determined from the optical density of the film using a He-Ne laser as the analyzing beam. The thin film combined with a static reaction system provides carefully controlled conditions of reaction time, temperature, and pressure, and pore-diffusion effects are eliminated. The carbon formed from methane is virtually free of heavy metal impurities. The reaction was studied at pressures 505
J.-R. CAO and M. H. BACK
506
of hydrogen below atmospheric and over the ternperature range 870-l 150K.
a reaction a reproducible surface was obtained, which gave consistent measurements of the rate. 3. RESULTS
2. EXPERIMENTAL
The reaction was studied in an apparatus similar to that described previously[l6]. The reaction vessel was a quartz cylinder, 4.4 cm id. and 11.Ocm long with plane fused windows. The volume was 182cm’ and the surface area 256cm2. The vessel was placed in the center of a Lindberg three-zone tubular furnace. The analyzing beam, from a 2 mW He-Ne laser, passed axially through the vessel and the signal was measured with photodiode and electrometer. At either end of the furnace the beam passed through evacuated cylinders which minimized convection currents and effectively reduced the light from the furnace. The absorbance of the film was calibrated by complete combustion of the carbon to carbon dioxide. The film of carbon was prepared by pyrolyzing methane (- 200Torr; - 27 KPa) at about 1100K. The characteristics of the carbon film prepared in this way have been described in previous studies. Measurements of the total surface area by low-temperature adsorption of nitrogen have shown that the surface area of the film for concentrations up to 6 X 10-6gcm-2 is not significantly different from the surface area of the clean quartz vessel and a mean value of - 30 m* gg’ has been estimated[l7]. The films have been examined with an electron microscope and no details or structure were observed at these low concentrations. Assuming a uniform deposit and a density of carbon of 1.9 g cm-‘, the average thickness of the film used in the present experiments was about 20nm. Measurements of the active surface area of the films with respect to the concentration of strongly-bound complexes of oxygen have shown that such films have great reactivity[l7]. Methane was Research Grade (99.99%) obtained from Matheson of Canada, Whitby, Ont. and used without further purification. Hydrogen was purified by passage through Molecular Seive 5A at liquid nitrogen temperature with a residence time of l-2 sec. Such treatment is usually sufficient to reduce the oxygen impurity to less than I ppm[lS], although the present analysis could not detect less than 6 ppm. The products were analyzed by expansion into a Toepler pump from which duplicate analyses were made by gas chromatography. A Hewlett-Packard Model 5750with FID detector was used to analyze the hydrocarbon products with a 1.2 m Durapak (Phenylisocyanate on Porasil) column maintained at 30°C. A GOW-MAC thermal conductivity detector was used for analysis of CO and CO, with a 4 x 800 mm silica gel column at 25°C. For a series of experiments a carbon film was first produced from pyrolysis of methane at 1100K. Most of the measurements described here were made with carbon films of average concentration 5 X 10mhg cm-‘. Several rate measurements were made with one carbon film since the consumption of carbon in each experiment was very small. When about 5-10% of the carbon had been consumed, the film was removed and a new one prepared. Provided sufficient time was allowed for evacuation after
Products of the reaction
Three products, methane, ethane and ethylene were detected and measured. Yields of the products as a function of time were measured over the temperature range 870-1150K and at pressures of hydrogen from 50-300 Torr (6.7-40 KPr). Yield-time curves for methane, ethane and ethylene are shown in Figs. l-3. Methane was the main product and the C2 hydrocarbons were always minor. Traces of acetylene and propane were observed under some conditions, but no higher hydrocarbons were detected. The yield of ethylene attained a maximum value, which approached 10% of the yield of methane, at an early stage of the reaction and fell rapidly thereafter. At high temperatures the yield of ethane fell rapidly with time and was less than the yield of ethylene, but at low temperatures the yield increased with time in a manner
0
4-
8
12
16
20
24
-I
Time/min
Fig. 1. Yield of methane as a function of time. Pressure of hydrogen177Torr (24 KPa). 0,908 K; 0, 1010K; 0, 1061K; A, 1111K.
I
10 t
Time/min Fig. 2. Yield of ethane as a function of time. Pressure of hydrogen 177Torr (24 KPa). A, 908K; 0, 11I I K.
Kinetics of the reaction of hydrogen with thin films of carbon
0”
1
1
10
I
20
1
30
I
I
40
50
Time / min
Fig. 3. Yield of ethylene as a function of time. Pressure of hydrogen 177Torr (24 K_Pa).0,908 K; A, 1061K; 0, II 11K.
507
Because of the rapid disappearance of ethane and ethylene it was not possible to measure an initial rate of formation for these products. For methane an initial rate may be estimated by extrapolation of the rate as a function of time. A log-log plot of these rates as a function of the initial pressure of hydrogen is shown in Fig. 5. The order of the rate is about OS-O.6 over the whole range of temperature. An Arrhenius plot of the initial rate at a pressure of hydrogen of 177Torr (24 KPa) is shown in Fig. 6. The observed activation energy is 51 kcal/mole over the temperature range 1050-l 150 K and decreases to 6.5 kcal/mole between 1010K and 870 K. Insofar as the initial rate of formation of ethylene could be estimated, its variation with temperature showed the same trend as the rate of formation of methane. 4. DISCUSSION
Mechanism of the reaction
The observed change in activation energy indicates a analogous to that of methane and became greater than
the yield of ethylene. The rate of formation of methane decreased slowly with increasing time of reaction and appeared to approach a limiting value. This value was much lower than the equilibrium concentration of methane at these conditions (2 x lo-’ mole/l at 177Torr (24 KPa) and 1111K) so the reverse formation of carbon from methane is not important in this system. The ratio CH,/C,H, as a function of time is shown in Fig. 4 for three temperatures. The increase in this ratio with time of reaction at high temperatures suggests that secondary conversion of ethylene to methane takes place. If methane were formed entirely through such reactions the ratio should approach zero at zero time. The fact that the ratio appears to extrapolate to a finite value indicates that a certain fraction of the methane is formed directly on the surface. Because of the difficulty in making measurements at short times of reaction the value of the extrapolation is probably not very precise. At the low temperature the ratio changed very little with time of reaction showing that secondary reactions of ethylene are not important in this region.
10
I
20
Time/
30
40
50
min
Fig. 4. The ratio CHJC2H4 as a function of time. Pressure of hydrogen 177Torr (24 KPa). 0,908 K; A, 1061K; 0, Ill I K.
Fig. 5. Logarithm of the initial rate of formation of methane against logarithm of the pressure of hydrogen. 0, 908K; 0, 961K; !I, 1061K; A, III1 K.
0’1 0.9
0.95
100
105
110
I 15
lOX”K Fig. 6. Logarithm of the initial rate of formation of methane as a function of IJT.
J.-R. CAOand M. H. BACK
508
Since the work of Zielke and Gorin further knowledge has accumulated on the nature of adsorbed species on the carbon surface [22]. In particular various experiments have shown that atoms and radicals probably have a degree of mobility on the surface[23]. The simplified scheme written above should be modified so that at each stage of hydrogenation the radical may migrate over the lattice and possibly desorb, at least in the high temperature region. For example,
fundamental change in the mechanism of the reaction over the range of temperature studied. Heterogeneous reactions often show a gradually increasing activation energy with increasing temperature and this behaviour has been called the compensation effect [19]. In general terms processes with higher activation energy become increasingly important as the temperature is increased. The effect is more noticeable with surface reactions because these reactions usually involve a complex series of concurrent and consecutive elementary processes whose relative rates continuously change as the temperature is increased. In the present reaction there appear to be two regions of fairly well-defined activation energy-a low-temperature region from 900-1010 K with E = 6.5 kcal/mole and a high temperature region from 1050-1150K where the activation energy is 51 kcal/mole. A similar change in activation energy in this temperature region was reported by Corney and Thomas[20]. Most studies have been made at temperatures above 900 K and activation energies ranging from 30 to 85 kcal/mole have been measured [7]. The large change in activation energy and the apparently constant value in the high-temperature region suggest that the change in mechanism involves a transition from mainly heterogeneous to mainly homogeneous processes. The results will be examined to assess the consistency of this interpretation
C-C,
H ;C=C(
H
+
H
H +
7 ‘:
'C-C'
c'
;c=c-,
-
H ';1 'I'H ,c-c:
or
+
(,\C
H
H2
'r' 7
+
H 8p
---_)
'C
I C
H
*
/ ii
C
2 C
(8)
(CHX),,+ Hz + CH, + H
(9) (10)
(2) H
= c.; Ia
'c
H\;1,H ',':
(U-U,, f (CHJ., -+ (CZHJa + C,H,
(1)
--_)
+
(7)
_H
H2
H
(CI-U, + I-I, + WI,), f I-I
This modification removes the restrictions that the Cz hydrocarbons must be formed by hydrogenation of neighbouring carbon atoms but a similar distribution of products could be obtained by either scheme. According to this picture ethylene, ethane and methane are all primary products whose distribution depends on the competition between abstraction from hydrogen or dimerization of the intermediate radicals. The extent of hydrogenation of the final product also depends on the rate of desorption of the product molecule. In this regard it is significant that acetylene was not detected at these temperatures, but has been reported at higher temperatures. It is possible that acetylene is strongly adsorbed and further hydrogenation competes effectively with desorption, or, in terms of reactions (6)-(lo), the CH radical may have less mobility than the more hydrogenated radicals and the dimerization of CH does not compete with reaction with hydrogen. The over-all activation energy for this series of hydrogenation steps will be the difference between the heat gained following the adsorption processes and the energy required for rearrangement and desorption. While these values are not known, either individually or collectively, a low value, such as observed, is reasonable.
H. 2
(6)
(CHJ, f (CH,). + (CzH,),,+CzHh.
4.1 Heterogeneous reactions In the low temperature region the ratio CH,/C,H, (Fig. 4) was essentially constant with time of reaction indicating that both products are formed directly through intermediates on the surface. The formation of C2 hydrocarbons as intermediates in the formation of methane is a reasonable consequence of the mechanism for hydrogenation proposed by Zielke and Gorin[21]. In this sequence of reactions neighbouring carbon atoms in the graphite structure are hydrogenated until CH7 groups are attained. Further reaction of the CH, group with hydrogen gives methane. If this mechanism is modified to allow the intermediate formation of ethylene and ethane, which have a certain probability for desorption, the present results are qualitatively explained. The mechanism of Zielke and Gorin, with these modifications, may be written as follows:
,
(CH), + Hz + (CH,), t H
t-t,
C2H4
(3)
H& F C
+
---$
2CH4
d
C2H6
(4)
(9
509
Kinetics of the reaction of hydrogen with thin films of carbon Homogeneous reactions 4.2.1 Secondary reactions of ethane. The most notice-
1
able distinguishing feature between the low-temperature and high-temperature regions is the relative importance of the products methane and ethane. At low temperature ethane is an important product but as the temperature is increased it becomes minor and is scarcely measurable at the highest temperature. As ethane is reduced methane becomes predominant. This change may be accounted for by the secondary thermal dissociation of ethane C,H, + 2CHT
I
I
(11)
which is unimportant in the low temperature region but in the high temperature region is sufficiently rapid to prevent accumulation of ethane in the gas phase. The rate constant for this homogeneous dissociation is well established[24]. It was estimated that at a pressure of 177Torr (24KPa) of hydrogen the ratio k,,/k;, was about 0.58, giving a value for k,, at 1111K of 0.107 s ‘. At the maximum yield of ethane under these conditions, 4 x lo-“’ mole/l, the rate of dissociation is: - $ (C2H,,)= k, ,(CZHh)max = 4.3 x IO ” mole/l sec. This rate therefore represents a lower limit to the initial rate of formation of ethane. The measured rate at 2 min reaction time (the shortest reaction time) is less than this, 2.3 x IO-l2 mole/l set, and it is apparent from the shape of the curve that the initial rate will be higher. If methane is formed by hydrogen abstraction from hydrogen by methyl radicals following dissociation of ethane, the rate of formation of methane at the maximum concentration of ethane would be 8.6 x IO-” mole/l sec. The observed rate at this time (3 min) is 9.2 x lo-” mole/l sec. In other words at 1111K the homogeneous thermal dissociation of the product ethane is sufficiently rapid to account for virtually all the methane formed in the reaction. In the low temperature region the rate of dissociation of ethane becomes negligible compared to the observed rate of formation of methane. At 908 K k,, at 177Torr (24 KPa) was estimated as 1.29x 10m5s-’ and the rate of dissociation of ethane at ten minutes,
1!--L_L_
IO
20
30
Time/
min
40
50
Fig. 7. Yield of methane as a function of time. Pressure of hydrogen 177Torr (24 KPa). 0, Pure hydrogen; A. hydrogen t 41 ppm CH.,; 0, A-O. 1.1 x lo-’ mole/l. The results, shown in Fig. 7 show that methane, in an amount comparable to that formed during a reaction, has no effect on the rate of formation of itself or of the other products. We may conclude that homogeneous secondary dissociation of methane is not important under these conditions, in agreement with the measured rate of the thermal dissociation of methane[25]. Further series of experiments were made at 1I1 1 and 908 K using hydrogen containing two different concentrations of -ethylene: 9 and 43 ppm. At each temperature the ethylene was rapidly consumed. At 1I II K (Fig. 8) ethylene was converted almost entirely to
3.0
_-1..
L
I
I
I
- $ (C2Hh)= 3.9 x IO-l5 mole/l sec. The rate of formation of methane at this time is 3.1 x lo-” mole/l set and it is evident that in this region secondary dissociation of ethane is not sufficiently rapid to account for the rate of formation of methane. and 4.2.2 Secondary reactions of methane ethylene. Participation of methane and ethylene in secondary gas-phase reactions was tested by studies of the effect of additions of small amounts of these products to the reactant hydrogen. A series of experiments was made at 1111K with a total pressure of 177Torr (24 KPa) of hydrogen containing 41 ppm CH,, corresponding to a concentration of
Time / min
Fig. 8. Yield of products as a function of time. Pressure of hydrogen 177Torr (24 KPa) with 43 ppm CzH4. 11I I K. 0, CH,; 0, CzH4;A, C2H6;I, X2 (eqn 12).
510
J.-R. CAO and M. H.
BACK
hydrogen. This sum is shown in both figures. This result clearly shows the importance of the secondary hydrogenation of ethylene. To test the relative importance of homogeneous and heterogeneous reactions in the hydrogenation process, further series of experiments were made with hydrogen containing 43 ppm ethylene in the absence of the carbon film. The results are illustrated in Figs. 10 and Il. At 1111K the rate of conversion of ethylene to methane was almost identical to that found in the presence of carbon. At 908 K the disappearance of ethylene was much faster in the absence of carbon. Without carbon, of course, there is no concurrent formation of ethylene and the only reaction occurring is the homogeneous hydrogenation which consumes ethylene. These results provide convincing evidence of the importance of the homogeneous hydrogenation reaction in the high temperature region. A simple mechanism involving initiation by dissociation of hydrogen may be written as follows:
Time / min
Fig. 9. Yield of products as a function of time. Pressure of hydrogen 177Torr (24KPa) with 43ppm C*H,.908K. 0, CH,; 0, C2H4;A, C2H6;I, XC2(eqn 12).
methane; ethane attained only a low concentration. By contrast at 908 K (Fig. 9) ethylene was converted mainly to ethane and the yield of methane was minor. A carbon balance for the reaction is given by the following expression:
xc2 = ; (CH, - CH,.) t (CzHs - GH,.) t (CzH4- CzH4*)
H,*2H
(13)
H t CzH4+ C,H,
(14)
C,H, t Hz -+&He t H
(15)
CzHc+2CH,
(16)
CH,tH,+CH,tH.
(17)
followed by
The rate of disappearance of ethylene is given as
(12) -$
(GH,) = M-MC&).
(18)
where the starred quantity refers to the yield from pure
I
If an equilibrium concentration of hydrogen atoms is achieved -;
(CzH,) = k,dK ;?H2”‘(C,H,).
(19)
The value of k14is obtained from the following[26]t log kld(1 mole-’ s-l) = 10.3- $$+.
IO
20
30
40
Time/min
Fig. 10. Yield of products as a function of time. 1111K. Pressure of hydrogen 177Torr (24KPa) with 43 ppm C*H,. Filled symbols refer to experiments with the carbon fdm. Open symbols refer to experiments without the carbon film.
tThis value refers to the abstraction reaction which may be faster than the addition at these temperatures. The rate measured is simply the disappearance of ethylene.
(20)
At 1111K, K13= 1 x lo-” atm. and with a pressure of hydrogen of 177Torr (24KPa) the equilibrium concentration of hydrogen atoms is 1.67 x lo-“‘mole/l. With 43 ppm ethylene the calculated initial rate of disappearance of ethylene is 1.5 X IO-* mole/l sec. The observed rate, measured over the shortest time interval of 0.5 min is 1.6 x 10d9mole/l sec. The slower rate of reaction probably indicates that the equilibrium concentration of H atoms is not achieved, although the agreement is probably within the uncertainty in the measurement of the initial rate and in the rate constant, k14,at these temperatures. A similar calculation at 908 K shows slightly better agreement. The calculated rate is 3.3 X lO~‘mole/l set, while the observed value is 5.2 x lo-” mole/l sec.
Kinetics of the reaction of hydrogen with thin films of carbon
(II
therefore, that at temperatures above about 1100K thermal dissociation of hydrogen provides an important source of radicals which increases the rate of the surface hydrogenation and the rate of secondary hydrogenation in the gas phase. Initiation by dissociation of hydrogen may be augmented by dissociation of ethane. A radical concentration may be calculated from the following expression
where k, is a radical termination rate constant. At I I I I K and 177 Torr (24 KPa) of hydrogen, assuming a value of lO”‘I/mole sec. for k,, [I?] will equal the equilibrium
Time / min
Fig. I I. Yield of products as a function of time. 908K. Pressure of hydrogen 177Torr (24 KPa) with 43 ppm C2H4.Filled symbols refer to experiments with the carbon film. Open symbols refer to experiments without the carbon film.
These calculations and the results of the experiments with added ethylene show that homogeneous hydrogenation of ethylene is rapid under the conditions of the present experiments and that the thermal dissociation of hydrogen provides a sufficient radical concentration for the chain reaction. In the high temperature region therefore, the yields of ethylene and ethane achieved only low steady-state values while being rapidly converted to methane. The major portion of the methane formed, however, is not the result of this homogeneous hydrogenation but rather arises through processes occurring on the carbon surface. This was evident from the extrapolation of the ratio CH&H4 to zero time, where the value was about 10 at all temperatures. The surface processes occurring at I I I I K are nevertheless not the same as those occurring at 908 K. Much more methane is formed at I I I1 K than expected from an extrapolation of the Arrhenius relation in the region 870-1010 K. This rapid increase in rate of the surface reaction can only arise through an increase in initiation and the observed activation energy of 51 kcal/mole for the rate of formation of methane suggests that the thermal dissociation of hydrogen provides this initiation. Hydrogen atoms may react directly with the carbon surface or with radicals or molecules adsorbed on the surface leading to an increased rate of formation of products. For example, methane may be formed by the following series of reactions: U-U., +
H + U-M.,
(21)
(CHZ):,+ H --f(CH,),,
(22)
(CH&, t H --)(CH,);, + CH,.
(23)
The reactions with hydrogen atoms would involve little or no activation energy and the over-all activation energy would therefore be close to one-half the heat of dissociation of hydrogen[27], or 53 kcal/mole. We conclude,
concentration of hydrogen atoms when C2H,>= 2.6 x 10 ‘mole/l. From Fig. 2 it may be seen that this concentration of ethane is not achieved at this temperature and initiation is therefore largely due to dissociation of hydrogen. It may be suggested that the conversion of ethylene to ethane may occur by a molecular addition of hydrogen. CzH, t H, +C,H,.
(24)
The rate constant for this reaction may be estimated from the results of the reaction of 177Torr (24 KPa) of hydrogen with 43 ppm ethylene without the carbon film at 908 K, where ethane is relatively stable. The initial rate of formation of ethane was about 5 x IO I” mole/l sec. giving a value of 1.5I/mole set for kz4 at 908K. Assuming a value of IO’ for AZ4 (a reasonable value for this type of 4-center activated complex[28]), the activation energy is 32 kcal/mole. This value is probably too low for a forbidden molecular addition reaction, and also would imply that the reverse reaction, a molecular dehydrogenation reaction, would have an activation energy of about 67 kcal/mole. This reaction has not been observed in the thermal decomposition of ethane where a radical chain mechanism is firmly established. It may be concluded that the conversion of ethylene to ethane takes place by homogeneous radical reactions. 4.3 Inhibition by products This mechanism provides a general picture of a product distribution involving methane, ethane and ethylene formed in surface reactions followed by further hydrogenation in the gas phase at suffikiently high temperatures. The mechanism does not predict a decrease in the yield of ethylene and ethane with time beyond the maximum value or a decrease in the rate of formation of methane at longer times. In a series of consecutive reactions such a decrease is observed only as the reactant is appreciably consumed; this is not the case in the present studies where conversion of hydrogen was always less than 0.1%. Such an effect may occur if the products inhibit the reaction. The series of experiments with additions of methane and ethylene to the reactant hydrogen showed that neither of these products had an inhibiting effect on the hydrogenation reaction. This is clearly evident in the
J.-R. CAO and M. H. BACK
512
case of ethylene by the sum X2, given by eqn (12), which indicates that the products from pure hydrogen are not reduced when the reactant is hydrogen with added ethylene. It must be concluded that inhibition of the reaction occurs before the products are desorbed from the surface. In the series of reactions represented by (6)-(IO) further addition of radicals or atoms to the adsorbed ethylene may occur to give higher molecular weight products. This may be regarded as a surface polymerization process and could lead to the formation of ring compounds, partially hydrogenated, which might remain strongly adsorbed on the surface. Such compounds could remove active carbon atoms and inhibit the reactions leading to formation of ethylene. Indirect evidence for this type of reaction was obtained. During the course of the experiments it was observed that a reproducible surface was obtained only after sufficient evacuation time after each reaction. In fact the rate measured immediately after an experiment was lower than the original rate, but returned to the original value as the evacuation time was increased. When the material removed by pumping was trapped with liquid nitrogen analysis showed the presence of small amounts of CZ and C, hydrocarbons but no higher molecular weight compounds. These observations suggest that higher molecular weight compounds are formed on the surface during the course of the reaction which desorb slowly at the reaction temperature and after desorption decompose in the furnace to lower molecular weight compounds. It was not possible to obtain a direct analysis of the high molecular weight products or to deduce their composition from the products collected after desorption. The results indicate only that, the surface hydrogenation reactions probably form higher molecular weight products which remain adsorbed. Similar observations of inhibition have been reported where, for example, in a flow system a constant rate of formation of methane was achieved only after a certain duration of the experiment[l2]. The formation of higher molecular weight products has also been reported, although not specifically linked with an inhibition effect [ 141. CONCLUSIONS
The hydrogenation of carbon by hydrogen over the temperature range 870-115OK has been shown to be a
process involving both heterogeneous and homogeneous reactions whose relative importance depends on the temperature. The consecutive steps in the hydrogenation, from ethylene to ethane to methane have been clearly demonstrated. It has also been concluded that an important inhibition process occurs on the carbon surface, probably involving the formation of higher molecular weight products which remain strongly adsorbed. Acknowledgement-This work was supported by the Natural Science and Engineering Research Council of Canada. REFERENCES J. L. Johnson, Adu. Chem. Ser. 131, 145 (1974). f : A. Tomita, 0. P. Mahajan and P. L. Walker, Fuel 56, 137 (1977). 3. J. Weber and M. Bastick, 14th Bjennial Carbon Co& p. 147. Pennsylvania State University (1979). 4. J. G. McCarty and H. Wise, J. Catal. 57,406 (1979). 5. P. L. Walker. F. Rusinko and L. G. Austin. Adu. Catalvsis 11, 138(1959j. 6. J. B. Lewis, Modern Aspects of Graphite Technoiogy (Edited by L. C. F. Blackman), p. 129.Academic Press (1970). 7. R. A. Krakowski and D. R. Olander, Lr.S.At. Energy Comm. UCRL 19149(1970). 8. C. W. Zielke and E. Gorin, Ind. & Engng Chem. 47, 820 (1955). 9. F. Moseley and D. Patterson, J. Inst. Fuel 38, 13 (t%S). 10. J. D. Blackwood, Aust. J. Chem. 15,397 (1962). 11. H. Imai, S. Nomura and Y. Sasaki, Carbon 13, 333 (1975). 12. A. de Koranyi, N. D. Parkyns and S. J. Peacock, 5th ht. Conf. on Carbon and Graohiie. D. 139.London (19781. 13. J. T: Shaw, National Coil Bo&d Gasification hote’No. 29 (1977). 14. P. Breisacher and C. P. Marx, J. Am. Chem. Sot. 85, 3518 (1963).
15. S. J. Steck, G. A. Pressley, S. S. Lin and F. F. Stafford, J. Chem. Phys. 50(S), 31% (1969). 16. C. J. Chen and M. H. Back, Carbon 17,495 (1979). 17. S. B. Tong, P. Pareja and M. H. Back, Carbon, in press. 18. C. R. Baker and R. S. Paul, Chem+Engng Prog. 59,61 (1963). 19. A. K. Galwev. Adu. Cat. 26.247 iI977). 20. N. S. Corney and R. B. Thomas, FERN-CUR-2502(1958). 21. C. W. Zielke and E. Gorin, Ind. Engng Chem. 47, 820 (1955). 22. H. Marsh, Chem. Sot. Spec. Pub. No. 32, p. 133. London (1978). 23. H. Marsh and A. D. Foord, Carbon 11,421 (1973). 24. A. B. Trenwith, I. Chem. Sot. Fur Trans. 175,614 (1979). 25. C-J. Chen, M. H. Back and R. A. Back, Can. J. Chem. 53,358O (1975). 26. V. V. Voevodsky and V. N. Kondratiev, Proa. Reaction Kinetics 1,41 (1961). 27. S. W. Benson. The Foundai~oas of ~hemicul Kinetics. ’ McGraw-Hill, Toronto (1960). 28. S. W. Benson, Thermochemica/ Kinetics. Wiley, New York (1968).
1. INTRODUCTION
The gasification of carbon by hydrogen is a deceptively simple reaction. In most studies methane is the main, and in some cases the only, product reported but in spite of this simple stoichiometry the details of the mechanism by which the hydrogenation occurs are not well understood. The reaction has received more attention in recent years because of the interest in the production of substitute natural gas. Although for this purpose most of the studies involve the reaction with coal and coal chars [I, 21, or the reaction with carbon catalyzed by metals[3,4], the reaction with pure carbon provides fundamental information necessary to the understanding of the hydrogenation process. Early work has been reviewed by Walker, Rusinko and Austin[S] and more recent work by Lewis[6] and by Krakowski and Olander[7]. Among the first suggestions concerning the mechanism was that of Zielke and Gorin[8] in which successive hydrogenation of neigh-
bouring HLCLH
edge carbon H , -C/H, ‘H
atoms occurred
giving =C-,
and finally methane. This basic
mechanism has been modified but not effectively changed by most subsequent work[1,9-121. One of the difficulties in formulating a mechanism and a rate equation that could have general applicability is the wide variation in the reported kinetic parameters of the rate. There appears to be no general agreement concerning the order of the rate with respect to hydrogen and, probably as a consequence, no agreement on activation energy. This is partly due to the variation in reactivity of the carbon used as the reactant and partly due to the wide range of conditions of temperature, pressure and conversion used in the study of the reaction. It is not
surprising that attempts to derive a rate equation for the reaction have had limited applicability. Recently an extensive analysis of data by Shaw[l3] has shown that the kinetics of the formation and decomposition of methane in the presence of char and hydrogen may be described by a two-step hydrogenation process on the carbon surface leading directly to methane. None of the mechanisms, however, consider the possibility of the formation of products intermediate to the formation of methane, for example, Cz hydrocarbons, and the consequence of their desorption from the surface, although some studies have reported the formation of higher hydrocarbon products[9, 141, particularly at higher temperatures(lS]. Such products would be less stable than methane and would rapidly disappear in secondary reactions, possibly leading to radical reactions in the gas phase. The purpose of the present study was to search for the formation of such products, if possible to distinguish between primary and secondary products, and to consider the related question of the importance of hydrogenation in the gas phase as distinct from the processes occurring on the surface. From the results of the study a mechanism for the hydrogenation has been proposed and the importance of radical reactions in the gas phase has been demonstrated. We have recently developed a technique for the study of reaction of gases with carbon which offers some advantages for kinetic studies [ 161.A thin film of carbon is deposited on the surface of the reaction vessel by the pyrolysis of methane. The concentration of carbon on the surface, in the range lo-‘-lO-~^ gem-‘. may be determined from the optical density of the film using a He-Ne laser as the analyzing beam. The thin film combined with a static reaction system provides carefully controlled conditions of reaction time, temperature, and pressure, and pore-diffusion effects are eliminated. The carbon formed from methane is virtually free of heavy metal impurities. The reaction was studied at pressures 505
J.-R. CAO and M. H. BACK
506
of hydrogen below atmospheric and over the ternperature range 870-l 150K.
a reaction a reproducible surface was obtained, which gave consistent measurements of the rate. 3. RESULTS
2. EXPERIMENTAL
The reaction was studied in an apparatus similar to that described previously[l6]. The reaction vessel was a quartz cylinder, 4.4 cm id. and 11.Ocm long with plane fused windows. The volume was 182cm’ and the surface area 256cm2. The vessel was placed in the center of a Lindberg three-zone tubular furnace. The analyzing beam, from a 2 mW He-Ne laser, passed axially through the vessel and the signal was measured with photodiode and electrometer. At either end of the furnace the beam passed through evacuated cylinders which minimized convection currents and effectively reduced the light from the furnace. The absorbance of the film was calibrated by complete combustion of the carbon to carbon dioxide. The film of carbon was prepared by pyrolyzing methane (- 200Torr; - 27 KPa) at about 1100K. The characteristics of the carbon film prepared in this way have been described in previous studies. Measurements of the total surface area by low-temperature adsorption of nitrogen have shown that the surface area of the film for concentrations up to 6 X 10-6gcm-2 is not significantly different from the surface area of the clean quartz vessel and a mean value of - 30 m* gg’ has been estimated[l7]. The films have been examined with an electron microscope and no details or structure were observed at these low concentrations. Assuming a uniform deposit and a density of carbon of 1.9 g cm-‘, the average thickness of the film used in the present experiments was about 20nm. Measurements of the active surface area of the films with respect to the concentration of strongly-bound complexes of oxygen have shown that such films have great reactivity[l7]. Methane was Research Grade (99.99%) obtained from Matheson of Canada, Whitby, Ont. and used without further purification. Hydrogen was purified by passage through Molecular Seive 5A at liquid nitrogen temperature with a residence time of l-2 sec. Such treatment is usually sufficient to reduce the oxygen impurity to less than I ppm[lS], although the present analysis could not detect less than 6 ppm. The products were analyzed by expansion into a Toepler pump from which duplicate analyses were made by gas chromatography. A Hewlett-Packard Model 5750with FID detector was used to analyze the hydrocarbon products with a 1.2 m Durapak (Phenylisocyanate on Porasil) column maintained at 30°C. A GOW-MAC thermal conductivity detector was used for analysis of CO and CO, with a 4 x 800 mm silica gel column at 25°C. For a series of experiments a carbon film was first produced from pyrolysis of methane at 1100K. Most of the measurements described here were made with carbon films of average concentration 5 X 10mhg cm-‘. Several rate measurements were made with one carbon film since the consumption of carbon in each experiment was very small. When about 5-10% of the carbon had been consumed, the film was removed and a new one prepared. Provided sufficient time was allowed for evacuation after
Products of the reaction
Three products, methane, ethane and ethylene were detected and measured. Yields of the products as a function of time were measured over the temperature range 870-1150K and at pressures of hydrogen from 50-300 Torr (6.7-40 KPr). Yield-time curves for methane, ethane and ethylene are shown in Figs. l-3. Methane was the main product and the C2 hydrocarbons were always minor. Traces of acetylene and propane were observed under some conditions, but no higher hydrocarbons were detected. The yield of ethylene attained a maximum value, which approached 10% of the yield of methane, at an early stage of the reaction and fell rapidly thereafter. At high temperatures the yield of ethane fell rapidly with time and was less than the yield of ethylene, but at low temperatures the yield increased with time in a manner
0
4-
8
12
16
20
24
-I
Time/min
Fig. 1. Yield of methane as a function of time. Pressure of hydrogen177Torr (24 KPa). 0,908 K; 0, 1010K; 0, 1061K; A, 1111K.
I
10 t
Time/min Fig. 2. Yield of ethane as a function of time. Pressure of hydrogen 177Torr (24 KPa). A, 908K; 0, 11I I K.
Kinetics of the reaction of hydrogen with thin films of carbon
0”
1
1
10
I
20
1
30
I
I
40
50
Time / min
Fig. 3. Yield of ethylene as a function of time. Pressure of hydrogen 177Torr (24 K_Pa).0,908 K; A, 1061K; 0, II 11K.
507
Because of the rapid disappearance of ethane and ethylene it was not possible to measure an initial rate of formation for these products. For methane an initial rate may be estimated by extrapolation of the rate as a function of time. A log-log plot of these rates as a function of the initial pressure of hydrogen is shown in Fig. 5. The order of the rate is about OS-O.6 over the whole range of temperature. An Arrhenius plot of the initial rate at a pressure of hydrogen of 177Torr (24 KPa) is shown in Fig. 6. The observed activation energy is 51 kcal/mole over the temperature range 1050-l 150 K and decreases to 6.5 kcal/mole between 1010K and 870 K. Insofar as the initial rate of formation of ethylene could be estimated, its variation with temperature showed the same trend as the rate of formation of methane. 4. DISCUSSION
Mechanism of the reaction
The observed change in activation energy indicates a analogous to that of methane and became greater than
the yield of ethylene. The rate of formation of methane decreased slowly with increasing time of reaction and appeared to approach a limiting value. This value was much lower than the equilibrium concentration of methane at these conditions (2 x lo-’ mole/l at 177Torr (24 KPa) and 1111K) so the reverse formation of carbon from methane is not important in this system. The ratio CH,/C,H, as a function of time is shown in Fig. 4 for three temperatures. The increase in this ratio with time of reaction at high temperatures suggests that secondary conversion of ethylene to methane takes place. If methane were formed entirely through such reactions the ratio should approach zero at zero time. The fact that the ratio appears to extrapolate to a finite value indicates that a certain fraction of the methane is formed directly on the surface. Because of the difficulty in making measurements at short times of reaction the value of the extrapolation is probably not very precise. At the low temperature the ratio changed very little with time of reaction showing that secondary reactions of ethylene are not important in this region.
10
I
20
Time/
30
40
50
min
Fig. 4. The ratio CHJC2H4 as a function of time. Pressure of hydrogen 177Torr (24 KPa). 0,908 K; A, 1061K; 0, Ill I K.
Fig. 5. Logarithm of the initial rate of formation of methane against logarithm of the pressure of hydrogen. 0, 908K; 0, 961K; !I, 1061K; A, III1 K.
0’1 0.9
0.95
100
105
110
I 15
lOX”K Fig. 6. Logarithm of the initial rate of formation of methane as a function of IJT.
J.-R. CAOand M. H. BACK
508
Since the work of Zielke and Gorin further knowledge has accumulated on the nature of adsorbed species on the carbon surface [22]. In particular various experiments have shown that atoms and radicals probably have a degree of mobility on the surface[23]. The simplified scheme written above should be modified so that at each stage of hydrogenation the radical may migrate over the lattice and possibly desorb, at least in the high temperature region. For example,
fundamental change in the mechanism of the reaction over the range of temperature studied. Heterogeneous reactions often show a gradually increasing activation energy with increasing temperature and this behaviour has been called the compensation effect [19]. In general terms processes with higher activation energy become increasingly important as the temperature is increased. The effect is more noticeable with surface reactions because these reactions usually involve a complex series of concurrent and consecutive elementary processes whose relative rates continuously change as the temperature is increased. In the present reaction there appear to be two regions of fairly well-defined activation energy-a low-temperature region from 900-1010 K with E = 6.5 kcal/mole and a high temperature region from 1050-1150K where the activation energy is 51 kcal/mole. A similar change in activation energy in this temperature region was reported by Corney and Thomas[20]. Most studies have been made at temperatures above 900 K and activation energies ranging from 30 to 85 kcal/mole have been measured [7]. The large change in activation energy and the apparently constant value in the high-temperature region suggest that the change in mechanism involves a transition from mainly heterogeneous to mainly homogeneous processes. The results will be examined to assess the consistency of this interpretation
C-C,
H ;C=C(
H
+
H
H +
7 ‘:
'C-C'
c'
;c=c-,
-
H ';1 'I'H ,c-c:
or
+
(,\C
H
H2
'r' 7
+
H 8p
---_)
'C
I C
H
*
/ ii
C
2 C
(8)
(CHX),,+ Hz + CH, + H
(9) (10)
(2) H
= c.; Ia
'c
H\;1,H ',':
(U-U,, f (CHJ., -+ (CZHJa + C,H,
(1)
--_)
+
(7)
_H
H2
H
(CI-U, + I-I, + WI,), f I-I
This modification removes the restrictions that the Cz hydrocarbons must be formed by hydrogenation of neighbouring carbon atoms but a similar distribution of products could be obtained by either scheme. According to this picture ethylene, ethane and methane are all primary products whose distribution depends on the competition between abstraction from hydrogen or dimerization of the intermediate radicals. The extent of hydrogenation of the final product also depends on the rate of desorption of the product molecule. In this regard it is significant that acetylene was not detected at these temperatures, but has been reported at higher temperatures. It is possible that acetylene is strongly adsorbed and further hydrogenation competes effectively with desorption, or, in terms of reactions (6)-(lo), the CH radical may have less mobility than the more hydrogenated radicals and the dimerization of CH does not compete with reaction with hydrogen. The over-all activation energy for this series of hydrogenation steps will be the difference between the heat gained following the adsorption processes and the energy required for rearrangement and desorption. While these values are not known, either individually or collectively, a low value, such as observed, is reasonable.
H. 2
(6)
(CHJ, f (CH,). + (CzH,),,+CzHh.
4.1 Heterogeneous reactions In the low temperature region the ratio CH,/C,H, (Fig. 4) was essentially constant with time of reaction indicating that both products are formed directly through intermediates on the surface. The formation of C2 hydrocarbons as intermediates in the formation of methane is a reasonable consequence of the mechanism for hydrogenation proposed by Zielke and Gorin[21]. In this sequence of reactions neighbouring carbon atoms in the graphite structure are hydrogenated until CH7 groups are attained. Further reaction of the CH, group with hydrogen gives methane. If this mechanism is modified to allow the intermediate formation of ethylene and ethane, which have a certain probability for desorption, the present results are qualitatively explained. The mechanism of Zielke and Gorin, with these modifications, may be written as follows:
,
(CH), + Hz + (CH,), t H
t-t,
C2H4
(3)
H& F C
+
---$
2CH4
d
C2H6
(4)
(9
509
Kinetics of the reaction of hydrogen with thin films of carbon Homogeneous reactions 4.2.1 Secondary reactions of ethane. The most notice-
1
able distinguishing feature between the low-temperature and high-temperature regions is the relative importance of the products methane and ethane. At low temperature ethane is an important product but as the temperature is increased it becomes minor and is scarcely measurable at the highest temperature. As ethane is reduced methane becomes predominant. This change may be accounted for by the secondary thermal dissociation of ethane C,H, + 2CHT
I
I
(11)
which is unimportant in the low temperature region but in the high temperature region is sufficiently rapid to prevent accumulation of ethane in the gas phase. The rate constant for this homogeneous dissociation is well established[24]. It was estimated that at a pressure of 177Torr (24KPa) of hydrogen the ratio k,,/k;, was about 0.58, giving a value for k,, at 1111K of 0.107 s ‘. At the maximum yield of ethane under these conditions, 4 x lo-“’ mole/l, the rate of dissociation is: - $ (C2H,,)= k, ,(CZHh)max = 4.3 x IO ” mole/l sec. This rate therefore represents a lower limit to the initial rate of formation of ethane. The measured rate at 2 min reaction time (the shortest reaction time) is less than this, 2.3 x IO-l2 mole/l set, and it is apparent from the shape of the curve that the initial rate will be higher. If methane is formed by hydrogen abstraction from hydrogen by methyl radicals following dissociation of ethane, the rate of formation of methane at the maximum concentration of ethane would be 8.6 x IO-” mole/l sec. The observed rate at this time (3 min) is 9.2 x lo-” mole/l sec. In other words at 1111K the homogeneous thermal dissociation of the product ethane is sufficiently rapid to account for virtually all the methane formed in the reaction. In the low temperature region the rate of dissociation of ethane becomes negligible compared to the observed rate of formation of methane. At 908 K k,, at 177Torr (24 KPa) was estimated as 1.29x 10m5s-’ and the rate of dissociation of ethane at ten minutes,
1!--L_L_
IO
20
30
Time/
min
40
50
Fig. 7. Yield of methane as a function of time. Pressure of hydrogen 177Torr (24 KPa). 0, Pure hydrogen; A. hydrogen t 41 ppm CH.,; 0, A-O. 1.1 x lo-’ mole/l. The results, shown in Fig. 7 show that methane, in an amount comparable to that formed during a reaction, has no effect on the rate of formation of itself or of the other products. We may conclude that homogeneous secondary dissociation of methane is not important under these conditions, in agreement with the measured rate of the thermal dissociation of methane[25]. Further series of experiments were made at 1I1 1 and 908 K using hydrogen containing two different concentrations of -ethylene: 9 and 43 ppm. At each temperature the ethylene was rapidly consumed. At 1I II K (Fig. 8) ethylene was converted almost entirely to
3.0
_-1..
L
I
I
I
- $ (C2Hh)= 3.9 x IO-l5 mole/l sec. The rate of formation of methane at this time is 3.1 x lo-” mole/l set and it is evident that in this region secondary dissociation of ethane is not sufficiently rapid to account for the rate of formation of methane. and 4.2.2 Secondary reactions of methane ethylene. Participation of methane and ethylene in secondary gas-phase reactions was tested by studies of the effect of additions of small amounts of these products to the reactant hydrogen. A series of experiments was made at 1111K with a total pressure of 177Torr (24 KPa) of hydrogen containing 41 ppm CH,, corresponding to a concentration of
Time / min
Fig. 8. Yield of products as a function of time. Pressure of hydrogen 177Torr (24 KPa) with 43 ppm CzH4. 11I I K. 0, CH,; 0, CzH4;A, C2H6;I, X2 (eqn 12).
510
J.-R. CAO and M. H.
BACK
hydrogen. This sum is shown in both figures. This result clearly shows the importance of the secondary hydrogenation of ethylene. To test the relative importance of homogeneous and heterogeneous reactions in the hydrogenation process, further series of experiments were made with hydrogen containing 43 ppm ethylene in the absence of the carbon film. The results are illustrated in Figs. 10 and Il. At 1111K the rate of conversion of ethylene to methane was almost identical to that found in the presence of carbon. At 908 K the disappearance of ethylene was much faster in the absence of carbon. Without carbon, of course, there is no concurrent formation of ethylene and the only reaction occurring is the homogeneous hydrogenation which consumes ethylene. These results provide convincing evidence of the importance of the homogeneous hydrogenation reaction in the high temperature region. A simple mechanism involving initiation by dissociation of hydrogen may be written as follows:
Time / min
Fig. 9. Yield of products as a function of time. Pressure of hydrogen 177Torr (24KPa) with 43ppm C*H,.908K. 0, CH,; 0, C2H4;A, C2H6;I, XC2(eqn 12).
methane; ethane attained only a low concentration. By contrast at 908 K (Fig. 9) ethylene was converted mainly to ethane and the yield of methane was minor. A carbon balance for the reaction is given by the following expression:
xc2 = ; (CH, - CH,.) t (CzHs - GH,.) t (CzH4- CzH4*)
H,*2H
(13)
H t CzH4+ C,H,
(14)
C,H, t Hz -+&He t H
(15)
CzHc+2CH,
(16)
CH,tH,+CH,tH.
(17)
followed by
The rate of disappearance of ethylene is given as
(12) -$
(GH,) = M-MC&).
(18)
where the starred quantity refers to the yield from pure
I
If an equilibrium concentration of hydrogen atoms is achieved -;
(CzH,) = k,dK ;?H2”‘(C,H,).
(19)
The value of k14is obtained from the following[26]t log kld(1 mole-’ s-l) = 10.3- $$+.
IO
20
30
40
Time/min
Fig. 10. Yield of products as a function of time. 1111K. Pressure of hydrogen 177Torr (24KPa) with 43 ppm C*H,. Filled symbols refer to experiments with the carbon fdm. Open symbols refer to experiments without the carbon film.
tThis value refers to the abstraction reaction which may be faster than the addition at these temperatures. The rate measured is simply the disappearance of ethylene.
(20)
At 1111K, K13= 1 x lo-” atm. and with a pressure of hydrogen of 177Torr (24KPa) the equilibrium concentration of hydrogen atoms is 1.67 x lo-“‘mole/l. With 43 ppm ethylene the calculated initial rate of disappearance of ethylene is 1.5 X IO-* mole/l sec. The observed rate, measured over the shortest time interval of 0.5 min is 1.6 x 10d9mole/l sec. The slower rate of reaction probably indicates that the equilibrium concentration of H atoms is not achieved, although the agreement is probably within the uncertainty in the measurement of the initial rate and in the rate constant, k14,at these temperatures. A similar calculation at 908 K shows slightly better agreement. The calculated rate is 3.3 X lO~‘mole/l set, while the observed value is 5.2 x lo-” mole/l sec.
Kinetics of the reaction of hydrogen with thin films of carbon
(II
therefore, that at temperatures above about 1100K thermal dissociation of hydrogen provides an important source of radicals which increases the rate of the surface hydrogenation and the rate of secondary hydrogenation in the gas phase. Initiation by dissociation of hydrogen may be augmented by dissociation of ethane. A radical concentration may be calculated from the following expression
where k, is a radical termination rate constant. At I I I I K and 177 Torr (24 KPa) of hydrogen, assuming a value of lO”‘I/mole sec. for k,, [I?] will equal the equilibrium
Time / min
Fig. I I. Yield of products as a function of time. 908K. Pressure of hydrogen 177Torr (24 KPa) with 43 ppm C2H4.Filled symbols refer to experiments with the carbon film. Open symbols refer to experiments without the carbon film.
These calculations and the results of the experiments with added ethylene show that homogeneous hydrogenation of ethylene is rapid under the conditions of the present experiments and that the thermal dissociation of hydrogen provides a sufficient radical concentration for the chain reaction. In the high temperature region therefore, the yields of ethylene and ethane achieved only low steady-state values while being rapidly converted to methane. The major portion of the methane formed, however, is not the result of this homogeneous hydrogenation but rather arises through processes occurring on the carbon surface. This was evident from the extrapolation of the ratio CH&H4 to zero time, where the value was about 10 at all temperatures. The surface processes occurring at I I I I K are nevertheless not the same as those occurring at 908 K. Much more methane is formed at I I I1 K than expected from an extrapolation of the Arrhenius relation in the region 870-1010 K. This rapid increase in rate of the surface reaction can only arise through an increase in initiation and the observed activation energy of 51 kcal/mole for the rate of formation of methane suggests that the thermal dissociation of hydrogen provides this initiation. Hydrogen atoms may react directly with the carbon surface or with radicals or molecules adsorbed on the surface leading to an increased rate of formation of products. For example, methane may be formed by the following series of reactions: U-U., +
H + U-M.,
(21)
(CHZ):,+ H --f(CH,),,
(22)
(CH&, t H --)(CH,);, + CH,.
(23)
The reactions with hydrogen atoms would involve little or no activation energy and the over-all activation energy would therefore be close to one-half the heat of dissociation of hydrogen[27], or 53 kcal/mole. We conclude,
concentration of hydrogen atoms when C2H,>= 2.6 x 10 ‘mole/l. From Fig. 2 it may be seen that this concentration of ethane is not achieved at this temperature and initiation is therefore largely due to dissociation of hydrogen. It may be suggested that the conversion of ethylene to ethane may occur by a molecular addition of hydrogen. CzH, t H, +C,H,.
(24)
The rate constant for this reaction may be estimated from the results of the reaction of 177Torr (24 KPa) of hydrogen with 43 ppm ethylene without the carbon film at 908 K, where ethane is relatively stable. The initial rate of formation of ethane was about 5 x IO I” mole/l sec. giving a value of 1.5I/mole set for kz4 at 908K. Assuming a value of IO’ for AZ4 (a reasonable value for this type of 4-center activated complex[28]), the activation energy is 32 kcal/mole. This value is probably too low for a forbidden molecular addition reaction, and also would imply that the reverse reaction, a molecular dehydrogenation reaction, would have an activation energy of about 67 kcal/mole. This reaction has not been observed in the thermal decomposition of ethane where a radical chain mechanism is firmly established. It may be concluded that the conversion of ethylene to ethane takes place by homogeneous radical reactions. 4.3 Inhibition by products This mechanism provides a general picture of a product distribution involving methane, ethane and ethylene formed in surface reactions followed by further hydrogenation in the gas phase at suffikiently high temperatures. The mechanism does not predict a decrease in the yield of ethylene and ethane with time beyond the maximum value or a decrease in the rate of formation of methane at longer times. In a series of consecutive reactions such a decrease is observed only as the reactant is appreciably consumed; this is not the case in the present studies where conversion of hydrogen was always less than 0.1%. Such an effect may occur if the products inhibit the reaction. The series of experiments with additions of methane and ethylene to the reactant hydrogen showed that neither of these products had an inhibiting effect on the hydrogenation reaction. This is clearly evident in the
J.-R. CAO and M. H. BACK
512
case of ethylene by the sum X2, given by eqn (12), which indicates that the products from pure hydrogen are not reduced when the reactant is hydrogen with added ethylene. It must be concluded that inhibition of the reaction occurs before the products are desorbed from the surface. In the series of reactions represented by (6)-(IO) further addition of radicals or atoms to the adsorbed ethylene may occur to give higher molecular weight products. This may be regarded as a surface polymerization process and could lead to the formation of ring compounds, partially hydrogenated, which might remain strongly adsorbed on the surface. Such compounds could remove active carbon atoms and inhibit the reactions leading to formation of ethylene. Indirect evidence for this type of reaction was obtained. During the course of the experiments it was observed that a reproducible surface was obtained only after sufficient evacuation time after each reaction. In fact the rate measured immediately after an experiment was lower than the original rate, but returned to the original value as the evacuation time was increased. When the material removed by pumping was trapped with liquid nitrogen analysis showed the presence of small amounts of CZ and C, hydrocarbons but no higher molecular weight compounds. These observations suggest that higher molecular weight compounds are formed on the surface during the course of the reaction which desorb slowly at the reaction temperature and after desorption decompose in the furnace to lower molecular weight compounds. It was not possible to obtain a direct analysis of the high molecular weight products or to deduce their composition from the products collected after desorption. The results indicate only that, the surface hydrogenation reactions probably form higher molecular weight products which remain adsorbed. Similar observations of inhibition have been reported where, for example, in a flow system a constant rate of formation of methane was achieved only after a certain duration of the experiment[l2]. The formation of higher molecular weight products has also been reported, although not specifically linked with an inhibition effect [ 141. CONCLUSIONS
The hydrogenation of carbon by hydrogen over the temperature range 870-115OK has been shown to be a
process involving both heterogeneous and homogeneous reactions whose relative importance depends on the temperature. The consecutive steps in the hydrogenation, from ethylene to ethane to methane have been clearly demonstrated. It has also been concluded that an important inhibition process occurs on the carbon surface, probably involving the formation of higher molecular weight products which remain strongly adsorbed. Acknowledgement-This work was supported by the Natural Science and Engineering Research Council of Canada. REFERENCES J. L. Johnson, Adu. Chem. Ser. 131, 145 (1974). f : A. Tomita, 0. P. Mahajan and P. L. Walker, Fuel 56, 137 (1977). 3. J. Weber and M. Bastick, 14th Bjennial Carbon Co& p. 147. Pennsylvania State University (1979). 4. J. G. McCarty and H. Wise, J. Catal. 57,406 (1979). 5. P. L. Walker. F. Rusinko and L. G. Austin. Adu. Catalvsis 11, 138(1959j. 6. J. B. Lewis, Modern Aspects of Graphite Technoiogy (Edited by L. C. F. Blackman), p. 129.Academic Press (1970). 7. R. A. Krakowski and D. R. Olander, Lr.S.At. Energy Comm. UCRL 19149(1970). 8. C. W. Zielke and E. Gorin, Ind. & Engng Chem. 47, 820 (1955). 9. F. Moseley and D. Patterson, J. Inst. Fuel 38, 13 (t%S). 10. J. D. Blackwood, Aust. J. Chem. 15,397 (1962). 11. H. Imai, S. Nomura and Y. Sasaki, Carbon 13, 333 (1975). 12. A. de Koranyi, N. D. Parkyns and S. J. Peacock, 5th ht. Conf. on Carbon and Graohiie. D. 139.London (19781. 13. J. T: Shaw, National Coil Bo&d Gasification hote’No. 29 (1977). 14. P. Breisacher and C. P. Marx, J. Am. Chem. Sot. 85, 3518 (1963).
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