THE ADSORPTION OF HYDROGEN, ETHYLENE, ACETYLENE AKD E T H A S E BY STAKYOUS OXIDE J. N. PEARCE AXD SYLVIA M. GOERGEX
The phenomenon of heterogeneous catalysis has been the subject of numerous investigations. Through the agency of these researches several facts have been established regarding the nature of contact catalysis. The more important of these may be briefly mentioned. ( I ) h solid cannot catalyze a reaction unless one or more of the reactants are adsorbed on its surface.' ( 2 ) While the activity of the catalyst and its power to adsorb reacting vapors on its surface do appear to go hand in hand, there is no quantitative relationship between the two phenomena" (3) Substances which are strongly adsorbed by a catalyst act as "poisons," in that they prevent the reactants reaching the surface of the catalyst. These poisons include foreign substances, -impurities mixed with the reactants. The poison may be one of the products of the reaction; indeed, either of the reactants may act as a poison, if it is so strongly and selectively adsorbed as to prevent contact of the second reactant with the catalyst surface. (4) The activity of the catalyst is markedly influenced by the nature of the heat treatment to which it has been subjected. ( 5 ) The maximum catalytic activity does not necessarily occur a t the temperature of maximum adsorption; it may only begin where the adsorption is barely measurable. The adsorption capacity may be considereds as a n index of the temperature at which the reaction can be induced. In other words, a lower temperature will initiate a reaction when adsorption is strong. An extended investigation of the activity of oxide catalysts a t temperatures betwen 300' and 400' by Sabatier and Mailhe4 has shown that a majority of the oxides promote two reactions simultaneously, one the process of dehydration, the other, the process of dehydrogenation. The analyses of the gaseous products obtained when the vapor of ethyl alcohol was passed over various oxides at 340' to 350' showed that with certain oxides ethylene only is produced. I n general, however, the effluent vapor is a mixture of ethylene and hydrogen in varying proportions, depending on the catalysts used. Thoria behaves almost exclusively as a dehydrating agent under the given conditions; alumina and tungstic oxide are only slightly less active dehydrating catalysts. The oxides of manganese, cadmium, and tin were found to be dehydrogenating catalysts exclusively; their dehydrating power is entirely surpressed. Between the two extremes the dehydrating and deBancroft: J. Phys. Chem., 21, 573, 644,734 (1917). Taylor: Third Report of Committee on Contact Catalysis, J. Phys. Chem., 28, 898 (1924). Pease: J. Am. Chem. Soc., 45,2296 (1923). Sabatier and Mailhe: Ann. Chim. Phys., (a), 20, 341 (1910). l
J
I424
J. N. PEARCE AND SYLVIA hl. G O E R G E S
hydrogenating activity of the oxide catalysts vary widely. I n general, the greater the dehydrating activity of the oxide, the less is its dehydrogenating activity toward alcohol. Benton' has studied the adsorption of a number of gases by several oxides. I n these papers he distinguishes between primary and secondary adsorption, as proposed by Langmuir. According to him, secondary adsorption is exhibited by inert adsorbents, such as charcoal, mica and silica gel, where the tendency of the different liquids to be adsorbed is in the same order as their boiling points, or as their freezing points. Secondary adsorption decreases continuously with rise in temperature, and it does not reach saturation value until a high pressure is attained. Primary adsorption is selective and is due to primary valences. From certain assumptions based on his own data, Benton has devised a method for calculating the relative magnitudes of primary and secondary adsorption. He finds that neither the secondary adsorption nor the total adsorption by oxide catalysts bear any relation to their catalytic activities. The primary adsorption of carbon monoxide is, however, in the same order as the catalytic activity of the oxides. Bischoff and ildkins2 have measured the adsorption of hydrogen, ethylene, and ethane by titanium oxide a t zoo and 97'. While the titanium oxides from different sources differed greatly in absolute and relative activities, they showed no similar differences in the adsorption of the gases which were the products of the reactions catalyzed. For a given temperature the volume of gas adsorbed was greatest for ethylene and least for hydrogen. The adsorption of hydrogen and ethylene by zinc oxide and ferric oxide has been studied by Lazier and Adkim3 Their results indicate that, while there is a qualitative agreement, there is apparently no quantitative relationship between total adsorption and catalytic activity. Ridea14 has measured the adsorption of the same two gases by thoria. Taylor and Kistiakowsky5 find that a t oo and 100' both zinc oxide and ZnO-Cr201, per unit weight, have greater adsorption capacities for hydrogen and carbon monoxide than d o most metal catalysts. The adsorption of both gases is pronounced at very low pressures and rapidly reaches saturation capacity independent of further increase in pressure. Thus far, investigations have been limited to oxides which are either wholly dehydrating catalysts, or to those which show both dehydrating and dehydrogenating activity. It seemed advisable, therefore, to continue these studies with a n oxide which, according to Sabatier and hlailhe, is a dehydrogenating catalyst only. The present investigation involves the study of the adsorption of hydrogen, acetylene, ethylene and ethane by stannous oxide a t oo, 78.5' and 100'. 'Benton: J. .Im. Chem. Soc., 45, 887, 900 (1923). * Bischoff and Adkins: J. Am. Chem. Soc., 47, 807 (19zj). 3Lazier and Adkins: J. Phys. Chem., 30, 3 j 3 (1926). Rideal: J. Am. Chem. Soc., 49, 116 (1927). 5Taylor and Iiistiakowsky: J. Am. Chem. Soc., 49, 2486 (192jl.
ADSORPTION BY STANNOUS OXIDE
I425
I n a later work, Sabatier’ states that tin oxide acts catalytically upon ethly alcohol above 300’ as a dehydrating agent. As we might expect a t this temperature, the oxide was found to be reduced to the metallic state. The metal itself exhibits marked catalytic activity above its melting point,2 but owing to the gradual growth of the large drops of the liquid a t the expense of the smaller ones, the active surface of the catalyst is slowly reduced. Brown and Henke3 have also found that tin is an excellent reduction catalyst for nitro-aromatic compounds. Materials and Apparatus The tin oxide was prepared by a method similar to that used by Bury and Partington‘ from an “analyzed” sample of SnC12.zH20. The label attached guaranteed it to give water-clear solutions and to contain only the following negligible impurities: “Cu, nil; Fe, 0.003; SOS, 0.0001;Koa, nil.” It was therefore used without further purification. 4 j grams of the S n C L 2H20 were first dissolved in zoo cc. of distilled water, clarified by the addition of a minimum amount of constant boiling hydrochloric acid, and then heated in contact with metallic tin. This solution was then slowly added, with rapid stirring, to a solution containing an amount of ?;a2COa just sufficient to neutralize the free acid and to precipitate the tin as Sn(OH)*. The precipitate was repeatedly washed by decantation until it no longer setled on standing. It was the filtered on a Biichner funnel, washed, dried a t room temperature, and finally stored in a vacuum desiccator over solid KOH. Bury and Partington report that stannous hydroxide reacts with air only slightly, if at all. They state that their dried product corresponds to the formula : 3 SnO.zH20. The technique employed in this work differs from that of Bury and Partington in that they reversed the procedure and added the Ka2C03solution to that of the tin chloride. When we followed their method we found that the reaction product was frequently contaminated by traces of SnC12. Further, the product exhibits a considerable tendency to turn gray, and even black during the process of decantation, especially if the walls of the vessel are scratched. Using our procedure, the product is always white. When heated in air the oxide readily takes fire and becomes incandescent during the oxidation. The dried stannous oxide thus prepared was ground to pass through an loo-mesh sieve, and then heated in an atmosphere of C 0 2 at the temperature of boiling aniline until the evolution of moisture ceased. The final product was yellow. The hydrogen used was prepared by electrolyzing a concentrated solution of KOH between nickel electrodes. The gas was then passed through an Sabatier: “Catalyse”, IOO (1927). Translation by Dr. B. Finkelstein. and Taylor: “Catalysis in Theory and Practice”, 214 (1919). a Brown and Henke: J. Phys. Chem., 27, 739 (1923). Bury and Partington: J. Chem. Soc., 121, 1998 (1922).
* Rideal
1426
J. S . PEARCE AND SYLVIA M. GOERGEN
alkaline solution of pyrogallol, concentrated H2SOa, and finally through a tube containing P205. The commercial ethylene showed by repeated analysis a 99.7 percent absorption in fuming H2S04. It was passed over fused CaClz and then through a trap immersed in a freezing mixture of solid COz and ether. The ethane was made by the Grignard method.' It was further purified by passing it successively through alcoholic hgKOs and concentrated H2S04. A
F
B
C
D
I
E
FIG.I
Acetylene prepared from calcium carbide and water was first allowed ----_.&l.- - l . . L l - : --..I :&:---A 4L C..-&l.--..-: C - A by passing it successively through a 40 percent solution of KOH, a I N solution of CuSO,, concentrated H2S01,and over Pz05. The nitrogen was made from equivalent quantities of NHaCI and NaN03. It was dried by passing through concentrated sulphuric acid and phosphorus pentoxide. &-..A
Grignard and Tissier: Compte8 rend., 132, 835 (1901).
--
ADSORPTIOS BY STANNOUS OXIDE
I427
The helium used in determining the dead space was purified by passing it through a liquid air trap containing an activated charcoal which had been evacuated previously a t red heat. The apparatus, Fig. I , is very similar to that used by Pease.l The bulb A containing the tin oxide is connected with the apparatus by capillary tubing. The manometer F, provided with a meter stick, is used to measure the pressure in the bulb A. The accurately calibrated burette B is surrounded by a water jacket containing a completely immersed standard thermometer. To it is attached a small manometer which permits accurate adjustment of the pressure within the burette. The large tube C is a gas reservoir for storing pure gases or their mixtures. The U-tube D serves as a valve for eliminating the pumps, and E is the hIcLeod gauge. IIercury is used as the displacing liquid throughout the apparatus. Experimental Procedure The entire apparatus is evacuated by means of a mercury vapor condensation pump in series with a “Hyvac” oil pump until the McLeod gauge indicates a pressure of less than 0.001 mm. The evacuation of the bulb is made difficult because of the tendency of the oxide to “puff” over. To avoid this completely the oxide bulb is first cautiously opened to the previously evacuated manometer F. By allowing the mercury to drop slowly in F the gases in A are gradually removed. During this evacuation the bulb -1is surrounded by the vapor of boiling aniline. When the pressure of the gases has been reduced to a few mms. the bulb is then opened directly to the pumps and is evacuated with the remainder of the apparatus. When the pressure as indicated by the IIcLeod gauge is sufficiently low and constant the well a t D is raised and the stopcock is closed. The pure gas to be studied is then introduced into B and carefully adjusted to atmospheric pressure. I n the meantime, the bulb A has been immersed in a Dewar flask containing ice and water, or in the vapor of a boiling liquid, depending on the temperature desired. W-hen thermal equilibrium has been attained a small amount of gas is admitted from B and its equilibrium pressure is determined. This procedure is repeated until the pressure within the bulb is approximately that of the atmosphere. From to 9 additions of gas are made a t each temperature. All gas volumes are corrected to oo and 760 mm. The “dead space” was determined in the usual way by means of helium. TO this end we carefully determined and plotted the pressure-volume isotherms of helium and the several gases a t ooJ 78.5’ and 100’. These graphs were readable to f 0 . 0 1 cc. The values used in making our calculations were read from these P-V isotherms a t even pressures. If we assume that the helium is not adsorbed by the oxide, then the difference between the number of CC. of helium and that of any other gas necessary to produce the same pressure a t the given temperature gives directly the volume of the gas adsorbed. Since we did not use exactly I O g. of oxide, these volumes were recalculated to give the volumes adsorbed by I O g. a t even pressures. The data thus accumulated are given in Tables I-V. I
Pease: J Am Chem. SOC 45,
1195 (1923).
1428
J. S . P E A R C E AXD SYLVIA M. GOERGES
TABLE I Volumes of Various Gases adsorbed by I O Grams of Stannous Oxide a t (Sample I). P. Hz CzH4 c2h5 cm.
cc.
5 IO
0.06 0.06
20
0.0;
CC.
CC.
0.29 0.59 1.16 1.73
0.70 1,38 2.51 3.50
4.50 5.37 6.21 7.04
30
0. I2
40
0.
I8
2.27
50 60 70
0.20
2.73
0.22 0.25
13
0.26
3.19 3,64 3.84
--
TABLE I1 Volumes of Hydrogen and Ethylene adsorbed by (Sample I).
IO
....
Grams of Oxide a t
P.
H2
cm.
cc.
cc.
5
0.05
IO
100'.
C2H4
0. I1
0.18 0.34
20
0.16
0.74
30
0.20
I.IO
40
0.25
1 , j0
50 60
0.26 0.25
io
0.25
75
0 . 2 5
1.91 2.29 2.67 2.82
T4BLE
111
Volumes of Various Gases adsorbed by I O Grams of Stannous Oxide at (Sample 11). P. cm.
Hz
C2Hz
CC.
cc.
0.86 1.49
20
0.03 0.06 0.09
30
0 .I O
40
0.13 0 . I8
3 ' 93 4.99 6.02 7.00 7.84 8,27
5 IO
50 60 70
0.21
75
0.2;
0.24
2.75
0'.
CZH4
0'.
C~HB
CC.
cc.
0.71 1.31 2.46 3.47
0.86 1.61
4.42 5.27
6.14 6.87 7.20
3.02
4.23 5.33 6.45 7.48 8.53
....
ADSORPTION BY STANNOUS OXIDE
I429
TABLEIV Volumes of Various Gases adsorbed by I O Grams of Stannous Oxide a t 78.5'. (Sample 11). P. Hz CzH2 CzH, CzHe cm.
cc.
cc.
cc.
5
0.20
0.14
IO
0.06 0.09
cc.
0.42
0.29
0.41
20
0.18
0.73
0.52
0.70
30
0.26
I .'08
0.74
0.96
40
0.27
1.41
1.21
50
0.29
60
0.28
1 ' 74 2.05
0.97 1.18 I,39
io
0.26
2.38
1.69 1.96
75
0.27
2.54
0.21
1.47
1.63 1.76
2.13
TABLE V Yolumes of \-arious Gases adsorbed by I O Grams of Stannous Oxide a t (Sample 11). P.
Hz
CIH,
em.
cc.
cc.
C2H6 cc.
5
0.04
0.10
0 .xo
IO
0.07
0.20
0.19
20
0.11
30 40 50
0.17 0.20
60 70 75
0.27 0.26
0.38 0.57 0.69 0.84 0.96 I .06
0.28
1.12
0.22
100'.
0.36 0.54 0.70
0.90 I .07 I.2j
I .36
Duplicate series of adsorption measurements were made for each gas a t a given temperature, and the P-V isotherms were carefully plotted. In general, the points obtained for the two series were found to lie almost exactly on the same smooth curve. When the two curves did not coincide, the adsorption measurements were repeated. On the whole, the adsorption data for the different gases are easily reproducible at these temperatures. The relative adsorption magnitudes are most clearly indicated by Figs. z and 3 , representing adsorption on samples I and 11, respectively. Because of the fact that the adsorption of hydrogen by both samples is so small and so nearly identical a t all temperatures, only one curve for hydrogen has been plotted. The adsorption of the other three gases is large, and especially so at 0'. TVith rise in temperature the magnitude of the adsorption decrease. rapidly. The effect of rise in temperature upon the volunie of the hydrocarbons adsorbed at any pressure is greatest for ethane and least for ethylene. At the higher temperatures the isothernu f!,ltten and beconie almost rectilinear
I430
J. N. PEARCE AND SYLVIA M. GOERGEN
a t the higher pressures. At oo ethane is the most highly adsorbed. With this exception the magnitude of the adsorption decreases a t all temperatures in the order: C*Hz, CzHs,C2H4,Hz. S o isotherm could be obtained for acetylene at rooo. When acetylene was admitted to the catalyst bulb it showed no tendency to attain an equilibrium pressure, even over a period of six days. From the time that the gas was admitted the pressure began to drop slowly and a t a continuously decreasing rate. The acetylene must, therefore, either decompose into carbon and hydrogen, and the hydrogen then react with the oxide, or it must undergo polymerization. The adsorption isotherm for hydrogen at 100' mas then
P
FIG.2 Adsorption Isotherms at oo and
100'.
Sample I.
carefully determined. As may be seen in Fig. 3 , there is no distinct evidence of any reaction between the oxide and hydrogen a t this temperature. This, then, eliminates the possibility of the formation of water vapor and the subsequent decrease in pressure due to its adsorption or condensation. If free hydrogen were formed by the deconiposition of the acetylene and it remained as a gas, the pressure in the oxide bulb should increase with time, since it is less adsorbed than the acetylene from which it would be formed. If, again, the hydrogen which might thus be formed were to reduce other acetylene molecules to ethylene, or possibly to ethane, the pressure should decrease as the reduction proceeds. In either case amorphous carbon should be formed and its presence would be indicated by a darkening of the oxide surface. SO
ADSORPTION BY STASNOUS OXIDE
1431
such discoloration was observed. We must conclude, therefore, that acetylene is polymerized slowly by stannous oxide a t 100'. I n a study of the effect of a large number of catalysts upon the reactions of ethylene, Walker' has found that acetylene in glass is stable toward heat a t temperatures up to 450'. Above this temperature acetylene is polymerized to a brown fluorescent liquid; it decomposes also to a slight extent into carbon, hydrogen and methane and some ethylene is formed by hydrogenation.
P FIQ.3 Adsorption Isotherms at oo, 78. j" and
100'.
SampIe 11.
Likewise, under similar conditions, he found that a temperature of 6 j o o is necessary to polymerize and decompose ethylene. For reasons to be mentioned later two samples of the stannous oxide were used. With sample I the isotherms were determined for C2H4, Hz and C2H6 a t oo,and for H Pand CPH,a t 100'. With either sample it was found possible to use the oxide first with one gas, then with a second, and finally back to the first without showing any impairment in adsorptive properties, so long as the oxide was not reduced. Both samples were prepared up to the Sn(OH)2 \Talker: J. Phys. Chern., 31, 961 (1927).
I432
J. S . P E A R C E AND SYLVIA & GOERGEN I.
stage a t the same time. Sample I was taken immediately after formation and subjected to heating and evacuation. Sample I1 was allowed to stand in a vacuum desiccator over solid potassium hydroxide for about six months, during which it acquired a light yellow surface film. It was subjected to the same heat treatment as sample I before using. Both samples became yellow throughout in the final heat treatment. While the order of the adsorption isotherms for the various gases is the same for both samples, the actual values for any particular pressure differ markedly. The adsorption by sample I is less than that by sample I1 a t 0 ' ; exactly the reverse is true a t IOO', the adsorption by I1 being less than that by I. I t is evident, therefore, that the rate of dehydration of the hydroxide does exert a marked influence on the surface properties of the oxide. The possibility that stannous oxide might catalyze the hydrogenation of et'hylene led us to attempt to determine the velocity of the hydrogenation process. h mixture containing equal volumes of hydrogen and ethylene was introduced into the adsorption bulb a t oo and the pressure readings were taken a t frequent intervals. The very slight decrease in pressure within the bulb over a 2 4 hour interval was so small that the reaction, if any, was inappreciable. A second experiment a t ob verified this conclusion. Duplicate experiments with similar j o percent volume mixtures were repeated a t 100'. Here also no measurable reaction was detected. Because of the relatively great difference in the adsorption of hydrogen and ethylene, it seemed advisable to increase the proportion of hydrogen in the reaction mixture. Hence, a mixture cont'aining i o . j percent by volume of hydrogen was made up in the gas burette and accurately checked by analysis. When this mixture was introduced into the catalyst bulb a t o o , the final equilibrium pressure after thirty-two hours was only 0.8 cm. lower than it would have been, assuming that no reaction had taken place. Another portion of the same mixture was then introduced into the previously evacuated bulb a t 100' and the change in pressure was observed over a period of twelve days. During this period the total decrease in pressure amounted to 11 cm.; even then there was no indication of the attainment of a constant equilibrium pressure. The j r e e gas was pumped off and its ethylene content was determined in the usual way by absorption in fuming sulphuric acid. Analysis showed 23.9 percent C,H,, as compared with 29.5 percent in theoriginalmixture. To be sure, the change in the percentage of ethylene was not entirely due to chemical reaction, but was due in part to a small amount of adsorbed ethylene which was not removed by complete evacuation. The fraction of ethylene remaining adsorbed under these conditions a t would be too small to account for the total decrease in the proportion of ethylene. This, and t'he continued drop in pressure within the bulb leaves little doubt that a very slow reaction mas taking place. Knowing that a rise in temperature should increase the reaction velocity, we submitted another sample of the (70. j-29.5) mixture to a temperature of 183'. Again t,he pressure continued to drop, but this time more rapidly. After contact with the oxide for I j hours the mixture was pumped IOOO
ADSORPTION BY STAKXOUS OXIDE
I433
out and analyzed. This time the analysis showed 4 1 . 4 percent of CzH4,-an increase in the proportion of ethylene. This behavior indicated that the hydrogen must be reacting with the stannous oxide. To verify this supposition hydrogen alone was admitted to the bulb a t oo and the equilibrium pressure was read. The temperature was then raised to 100' and again a definite equilibrium pressure was obtained. Finally, the bulb was heated to 183'. At this temperature no equilibrium was attained: the pressure within the bulb continued to drop over a period of 1 2 hours. The bulb was then cooled to o ' , and the pressure was found to be 37.85 cm., whereas the initial pressure a t oo was 47.26 cni. This decrease in pressure indicates conclusively that stannous oxide is reduced by hydrogen a t temperatures as low as 183'. Since the composition of the oxide was unquestionably changed by the reduction, the material was rejected and the adsorption experiments were continued on sample 11. This accounts for the apparent incompleteness of the adsorption data for sample I. The surprisingly low reduction temperature of stannous oxide brought the study of the hydrogenation velocity t o a sudden end. The experimental results do show, however, that stannous oxide does slowly catalyze the reaction at IOO', when the mixture of hydrogen and ethylenc is approximately 3 to I . A j o percent mixture gives no detectable reaction velocity. This is not surprising when we consider the relatively high adsorption of the ethylene a t 100'. As the mixture becomes richer in hydrogen the possibility of its being adsorbed even in the presence of the more readily adsorbed ethylene is increased. From the results obtained it would also appear that the simultaneous adsorption of both hydrogen and ethylene is necessary before hydrogenation can take place. However, the ethylene may be adsorbed so highly that it will inhibit the adsorption of the hydrogen to an appreciable extent.
Summary The adsorption isotherms of Hz, C2H2,CzH4and C2H4upon stannous oxide have been determined at oo, 78.5' and 100'. The adsorption of,hydrogen is very small a t all temperatures. The adsorption of the three hydrocarbons is relatively large and decreases rapidly with rise in temperature. Hydrogen reduces stannous oxide a t temperatures as low as 183'. The acetylene is polymerized slowly in contact with the oxide at 100'. Hydrogenation can take place at 100' only when the proportion of the hydrogen in the mixture is relatively very large. This would indicate that the adsorption of hydrogen is necessary before the reaction can take place. Physical Chemistry Laboratory, State Cnic,ersity of I o w a . M a y 5 , 1988.
