Catalysts for the Equilibrium, EthyleneWater-Ethanol A. J. PAIK, SHERLOCK SWANN,JR., A N D D. B. KEYES University of Illinois, Urbana, Ill.
T
The vapors leaving the catalyst HE problem of the hydraThe aerogels tried gave negative results chamber passed through two traps, tion of ethyiene to ethanol, as catalysts at low temperatures. The one cooled by an ice-salt mixture ahhough not new, is still and the other by solid carbon dibest catalyst studied was 5 per cent silver of scientific and industrial inoxide and acetone. Any condensulfate in sulfuric acid on pumice. The sable material was removed at this terest. At high temperatures point. The noncondensable gases silver sulfate catalyzed the following where suitable catalysts will were metered by a wet test meter, operate, the vapor-phase dehyreaction : K , attached t o the end of the train. dration of ethanol i s quite comThree traps are shown at J ; CQHsHS04 CzH4 HzS04 plete, and the amount of ethanol the first was used while bringing the system to equilibrium before remaining at equilibrium cannot It was inactive for the following reaction : making a run. The following be accurately d e t e r m i n e d . two traps were removable and Previous investigators have CzHsHS04 H2O CzHsOH H2S04 made it possible for the conshown that, although the use densate collected during a run to be weighed. of high pressures along with A constant rate of feeding the alcorhol was maintained by disthe high temperatures increases the concentration of the placing it from a buret, A , with mercury from reservoir D, the ethanol at equilibrium, at the same time side reactions are rate of flow of the mercury being controlled by a capillary, C. promoted to a greater extent. The buret was filled with alcohol from reservoir B before each run. It was therefore proposed to study catalysts for establishThe catalyst chamber was submerged in a large test tube with ing the equilibrium a t low temperatures and atmospheric side-arm overflow L. Glycerol was pumped into the tube from pressure. Later i t became desirable to study the mechanism thermostat G by means of an oil pump, H , driven by motor I. by which the most successful catalyst operated. The glycerol returned to the thermostat through the side arm There are several excellent literature summaries of the work of the tube. The thermostat was electrically heated and was held at the desired temperature by use of a mercury regulator done on the conversion of ethylene to ethanol from both the system. commercial and the theoretical aspects. Neumann (IO) The reason for the use of this type of thermostat was to faciliprobably gave the most comprehensive survey up to 1924. tate better heat transfer and greater ease of temperature control Gallagher ( 2 ) reviewed the work from 1828 to 1934. The than is usually found in the case of chambers heated by an electric furnace. The flow of glycerol could be momentarily liberal footnotes to the second and seventh chapters of the stopped, giving a view of the catalyst to ascertain whether fouling book by Marek and Hahn (8)cover the field thoroughly. occurred. This apparatus operated excellently except for the More recent investigations covering catalyst selection and excessive glycerol fumes at temperatures around 175' C. These free energy determinations are those of Sanders and Dodge were eliminated by placing a hood above the apparatus. The catalyst chamber shown in Figure 2 is an annular space ( l l ) ,Stanley, Youell. and Dymock ( l a ) , Gilliland, Gunness, between two Pyrex glass tubes sealed together a t each end. and Bowles ( S ) , and Bliss and Dodge ( 1 ) . Suitable entrance and exit tubes for the vapors were added. The ethanol vapors were fed into the bottom of the chamber, Apparatus allowing them to come to the temperature of the thermostat before contact with the catalyst. The chamber had a capacity The apparatus used for the catalyst study is shown in of 250 cc. The outer tube was 40 mm. in diameter, and the Figure 1: distance between the inner and outer tubes was about 6 mm. Ethanol was fed at a constant rate to a vaporizer coil in a The annular design was adopted to facilitate better heat transfer heated glycerol bath, E, and then into the catalyst chamber, F . between catalyst and thermostat.
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AN AIRPLANE VIEW O F THE CARBIDE AND CARBON CHEMICALS CORPORATION'S SYNTHETIC ETHYLALCOHOL PLANT UTILIZINGTHE SULFURIC ACIDLIQUIDPHASEPROCESS
INDUSTRIAL AND ENGINEERING CHEMISTRY
174
Experimental Procedure The dehydration was first attempted over the following ten aerogels (6, 6, 7): Alumina (pure), alumina with 29.9 per cent phosphoric acid (H3P04), alumina with 15 per cent boric acid (H3B08) and 15 per cent manganese oxide, alumina with 15 per cent manganese oxide, alumina with 15.1
VOL. 30, NO. 2
next best catalysts, since these metals are in the same periodic group with The results are expressed as the per cent of theoretical Stanley, Youell, and Dymock (12) used a similar silver sulfate-sulfuric acid catalyst in measuring the dehydration
D
FIGURE 1. APPARATUS FOR CATALYST STUDY per cent manganese oxide, alumina with 15.1 per cent boric acid, alumina with 6.9 per cent phosphoric acid, silica (pure), silica with 17.4 per cent phosphoric acid, and especially pure alumina. These attempts were entirely unsuccessful, since no reaction took place a t 175" C. or below. The MnOBO.3HsP04 catalyst used in one run is an attempted duplication of the catalyst of Stanley, Youell, and Dymock (12). This catalyst was also inoperable a t the low temperature of 175" C.
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Catalyst Inipregnated on 250 Cc. of Pumice
2 g. AgzS04 10 cc. H&04 (1
+ 1.4 g. AgZSOr + 0.1 g. HaBOa
96% CpHsOH per Hr.
Dehydrstion
cc.
Cc.
%
296 312
24.4 26.6
81.1 86.9
280 252 188 209 268 283 235 271
iii
li:7 14.6 24.0 28.6 26.1 25.3 24.6 28.0 27.4
0.0 35.1 76.7 68.9 51.5 57.3 73.4 77.6 64.3 74.2
3iO
23:1
84.9
~
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C2HsOH e C~H-H~HSO~ HzO C2HsHS04 ;t CzH4 H2S04
OF CATALYSTS (178" C.) TABLEI. COMPARISON
CaH4a per Cc. of 95% CaHsOH
equilibrium a t temperatures as low as 145' C. No account was given of its performance. Bliss and Dodge (1) listed the same catalyst in a group whose activity they studied, but in their discussion of results this catalyst was not mentioned. At this point in the investigation it was decided to study the mechanism of ethanol dehydration over the silver sulfatesulfuric acid catalyst in the hope of establishing a background for the future selection of catalysts. The accepted mechanism for the formation of ethylene from ethanol and sulfuric acid is as follows:
0.0
At standard conditions. I
Another group of catalysts was prepared by impregnating granular pumice with sulfuric acid containing aluminum sulfate, boric acid, manganese carbonate, platinum chloride, gold chloride, silver sulfate, and various combinations of them. The results obtained are shown in Table I. Silver sulfate with sulfuric acid was the best catalyst tried, and attempts to promote it were unsuccessful. The second best catalyst was a mixture of 18.3 per cent by weight of aluminum sulfate with sulfuric acid. It is interesting to note that the gold and lithium salts with sulfuric acid are the
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(1) (2)
Lommel and Engelhardt (8) and Tropsch and Mattox (13) found that the reverse of reaction 2 was greatly aided by the presence of silver sulfate. The problem remaining was to ascertain the effect of silver sulfate on reaction 1 from both directions and on the decomposition of ethylsulfuric acid according to reaction 2. Ethylsulfuric acid was made by bubbling ethylene (obtained from the U. S. Industrial Chemical Company) through a sintered glass diaphragm into a 70-mm. test tube containing c. P. 95 per cent sulfuric acid heated to approximately 50" C. in a n oil bath. The ethylene was purified by passing i t through a train of three sulfuric acid scrubbers and a trap cooled by solid carbon dioxide and acetone before absorbing it in the sulfuric acid. The concentration of ethylsulfuric acid is expressed in per cent by weight based on the increase in weight of the sulfuric acid due to absorbed ethylene. The effect of silver sulfate on the decomposition of ethylsulfuric acid was studied by heating two equal portions to 140" C. and collecting the ethylene formed. Silver sulfate was added to one of the portions. I n one of the tests fuming sulfuric acid was added to the ester a t room temperature in a n amount sufficient to react with all the water brought in with the 95 per cent sulfuric acid used in producing the ester. This was done to make certain that hydrolysis of: ethylsulfuric acid to ethanol with a subsequent dehydration of the latter compound was not occurring. Excessive carbonization
INDUSTRIAL AND ENGINEERING CHEMISTRY
FEBRUARY, 1938
175
shown in Table 111. The silver sulfate had no effect on this reaction, as the results with and without silver sulfate are identical within experimental limits. The reverse reaction was studied by the use of two series of sealed tubes placed in a thermostat. Each tube contained the same amount of ethylsulfuric acid to which was added water in the molecular ratio of ten to one. One series conTABLE11. EFFECTOF SILVERSULFATEON THE REACTION tained silver sulfate. All the tubes were placed in the therCzHsHS04 CzH4 mostat a t the same time, and corresponding tubes of each (Temperature, 140" C.; time of runs, 3 hours) series were removed together at definite intervals. After CZH4 a t cooling in an ice bath, aliquot portions of the contents of each Standard Decompn. of Conditions CzHoHS04 Flask Content tube were titrated (2-cc. samples) ; the reaction progress was cc. % again indicated by the free acid. The reaction was run a t 100 cc. (147 9.) CzHaHSOa, 69%: 99 CC. 60" C. so that its progress would be measurable over a period (194g ) 157 fuming HzSOa 512 2.9 100 cc. (i47 g.? CzHpHS04, 69%; 77 cc. of several hours. (151 g ) 157 fuming HzSOr. 10 g. AgzSOh 980 5.5 100 cc. (i48g.7 CzHaHSOa, 70:7%; 10 cc. The results are shown in Table IV and indicate that the 4710 25.3 (109.1 Hz0 silver sulfate had no effect on the reaction. The free acid in 100 CC. (1489.) CzHsHSOa 70.7%; 10 CC. (10g.) HzO; 10 g. A g z h 7941 42.8 the previous runs is expressed as cubic centimeters of alkali needed for titration, since it was not convenient to determine the density of each sample. Forward reaction 1 was studied by making two identical mixtures of equivalent amounts of ethanol and sulfuric acid. T o one of these mixtures silver sulfate was added. Samples TABLEIV. PROQRESB OF REACTION C2HsHS04 H20= C2HsOH HzS04AT 60' C. WITH AND WITHOUT SILVER SULFATE
prevented the direct addition of ethylene to 100 per cent sulfuric acid a t 50' C. The results are shown in Table 11; the presence of silver sulfate almost doubles the amount of ethylene formed over a given period.
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Cc. of 1.015 N NEOH Without With 5% AgPSOa AgzSOa 15.35 15.35 15.70 15.70 16.40 16.25 16.60 16.60
Time Interval, Hr. 0.5 1.0 1.6 2.0
,
Cc. of 1.015 N NEOH Time Without With 5% Interval, AgzS04 AgBO4 Hr. 17.05 17.00 2.5 17.40 17.20 3.0 18.05 18.10 5.0 18.65 18.60 7.0
The results can be summarized by the statement that silver sulfate has no effect on the rate of attainment of equilibrium in the system CzHbOH HzS04 e CzHsHS04 HzO, but it greatly hastens the rate in the system CzHsHS04 CzH4 HzS04. This shows conclusively, then, that the silver sulfate-sulfuric acid catalyst does not promote the direct dehydration of ethanol. Instead, the ethanol forms the ester with the sulfuric acid, and the silver sulfate causes the rapid decomposition of the ester into ethylene. Ipatieff (4)showed that ethylene forms the monoethyl ester with phosphoric acid. It can, therefore, be safely deduced that the successful phosphoric acid catalysts-for example, the Stanley catalyst-behave in a similar manner through the formation and decomposition of the ester. The results of this investigation show that catalysts capable of sorbing ethylene should be sought for the hydration of ethylene to ethanol .
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Acknowledgment Several of the aerogels for this investigation were supplied by K . K . Kearby.
FIQURE 2. CATALYBT CHAMBER
Literature Cited (2 cc.) were taken from each flask a t definite intervals and titrated. The amount of free acid was a n index of the progress of the reaction. The reaction was run a t tap water temperature to decrease the reaction velocity. Results are
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TABLE 111. PROGRESS OF REACTION CtH60H HzS04 = C2H6HS04 H 2 0AT 15.9" C . WITH AND WITHOUT SILVER SULFATE
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Cc. of 1.015 N NaOH Without With 5% AgaSOa AgaS04 31.50 31.50 29.40 29.00 27.90 28.20 27.20 27.05 26.50 26.50
Time Interval, Min. 0 20 40 60 80
Cc. of 1.015 N NaOH Without With 5% AgrSO4 AgaSOn 25.90 25.90 25.70 25.75 25.20 25.30 25.10 25.20
Time Interval, Min. 100 120 260 300
(1) Bliss and Dodge, IND. EKG.CHEX.,29, 19 (1937). (2) Gallagher, Ph.D. thesis, Univ. Ill., 1934.
(3) Gilliland, Gunness, and Bowles, IND.ENG. CHEM.,28, 370 (1936). (4) Ipatieff, Trans. Electrochem. Soc., 71, 333 (1937). (5) Kearby, K.K., Ph.D. thesis, Univ. Ill., 1937. (6) Kistler, S. S.,S. Phvs. Chem.,36, 32 (1932). ESG. CHEX.,26, 388 (1934). (7) Kistler, Swann, and Appel, IND. (8) Lommel and Engelhardt, Ber., 57, 848 (1924). (9) Marek and Hahn, "Catalytic Oxidation of Organic Compounds in the Vapor Phase," A. C. S. Monograph 61, New York, Chemical Catalog Co., 1932. (10) Neumann, Gas- u . Wasserfach, 67, 1, 14,53 (1924). ENG.CHBM.,26, 208 (1934). (11) Sanders and Dodge, IND. (12) Stanley, Youell, and Dymock, J . SOC. Chem. Znd., 53, 205T (1934). (13) Tropsch and Mattox, IND.ENG.CHEM.,Anal. Ed., 6,404 (1934). REOEIYED
July 19, 1937.
