CATALYTIC EQUILIBRIUM BETWEEN ACETALDEHYDE AND ALCOHOL* B Y WILDER D. BANCROFT ASD AYERY B. GEORGE
Introduction The catalytic decomposition of ethyl alcohol has been studied by several investigators, and the products obtained have been found to be dependent on the nature of the catalytic agent employed. It was first observed by Berthelot and Jungfleisch' that when ethyl alcohol is passed through a heated taube,it begins to decompose a t about joo°C, the decomposition consisting of two distinct superposed reactions, the dehydration to ethylene and water, and the dehydrogenation to acetaldehyde. In contact with various catalytic agents it is now known that dehydration may take a third course, or rather, stop a t an intermediate stage giving ether. Ipatiev showed that the proportion in which alcohol undergoes these changes is very much modified by the presence of catalysts. With different catalysts, the decomposition products of ethyl alcohol vary by widely different factors, and Sabatier has drawn up a table in which the relative dehydrating and dehydrogenating influences of various oxides are compared. This work was done in the later stages of Sabatier's investigations on the catalysis of organic compounds at solid surfaces. This aspect of the subject was especially studied by Sabatier and Mailhe,* who found that, for example, ethyl alcohol was converted into ethylene and water or, alternatively, into hydrogen and acetaldehyde, by the oxides quoted in Table I. The direction and importance of the activity of the various oxides can be shown clearly by a comparison of the volume and composition of the gas evolved by them, when equal volumes of them are used at 340'-3 jo°C with the same amount of ethyl alcohol. All of the oxides have been prepared below 350°C. From Table I it is clear that thoria, alumina, and the oxides of tungsten and chromium are pre-eminently dehydrating catalysts, whilst a large number of others promote both types of decomposition. Finally, a few oxides, notably those of zinc, tin, cadmium, manganese and magnesium are almost exclusively catalysts of dehydrogenation. It is necessary, especially from a technical standpoint, to consider not only the relative proportions of the two types of reaction, but also to bear in mind the intensity of the change set up by any given oxide. .Is a matter of fact, the dehydrogenating oxides are almost wholly of feeble activity, and consequently have not been utilized in technical practice, in the production of an *This work is done under the programme n o v tieing rarried out at Cornell University and supported in part by a grant from the Heckscher Founda!ion for the Advancement of Research established by August Heckscher at Cornell Gniversity. Berthelot and Jungfleisch: "Trait6 6Mmentaire de Chimie organique," 1. 2 j 6 (1886). Ann. Chim. Phys., (8) 20, 289 (1910).
'
EQUILIBRIUM BETWEEN ACETALDEHYDE AND ALCOHOL
2'95
TABLE I Vol of gas Oxide
in cc./min.
ThOz
31. o
A1203
2 1 .o
1vzos
57 . o
Cr203 Si02 TiOn Be0
4.2
Zr02 u02 MOZOS Fez03 VzOa ZnO SnO (Initial) CdO (Initial) MnO MgO
cu
Composition of
70 C J L IO0
gas 70
Ha
trace
14.0
98.5 98.5 91.o 84.0 63 .o 45 .o 45.0 24.0
5.0
23 . o
77 .o
14.0 6.0
14.0 9.0 5.0
86.0 91. o
45.0
0.0
100.0
0.9
7 .o I .o
I .o
32.0
1.5
1.5
9.0 16.0 37 .o 5 s .o 55.0
76.0
95.0
11.2
0.0
100.0
3.5
0.0
100.0
traces
0.0
100.0
I IO
0.0
100.0
aldehyde or ketone from an alcohol. For dehydrogenation processes the finely divided metals such as nickel, copper, platinum and cobalt have been found to be the most active catalysts, the first two named being used to a considerable extent in technical processes. The action of a metallic catalyst is connected to some extent, although not exclusively, with its capacity for adsorbing hydrogen, so that catalysis by metals can usually be placed in the categories of hydrogenation or dehydrogenation. The metallic oxides display little tendency to act simply as hydrogenating or dehydrogenation catalysts; they are more usually concerned either in the removal of the elements of water from organic compounds or in processes of more or less oxidation. The dehydrogenation of alcohols was first studied by Ipatiev,l but it is to Sabatier and Senderens2 that we owe a complete understanding of the dehydrogenation of alcohols. They showed that when primary saturated alcohols are passed over finely-divided copper, they are regularly decomposed into aldehydes and hydrogen. The action begins a t about 2oo°C, becomes rapid a t z50°C, and is almost the exclusive reaction up to 30oOC. This is a very advantageous method for the preparation of aliphatic aldehydes, particularly for those which, on account of low volatility, are difficult to prepare by oxidation of the alcohols. The transformation can never be complete, even when a long train of copper is used, since the hydrogen which is formed can be added to the aldehyde by copper above 200'C. Hence the reaction is Ber., 34, 3 j j 9 (1901); 35, 1047 (1902). 136, 921 (1903); Ann. Chim. P h p . , ( 6 ) 4, 463 (190j).
* Compt. rend.,
2196
WILDER D. BANCROFT AND AVERY B. GEORGE
limited, but the conditions are favorable to the decomposition because the operation is carried on in the presence of a small concentration of hydrogen. CzHsOH
a CHaCHO + Hz
The equilibrium is favorable to hydrogenation a t a low temperature, dehydrogenation becoming more pronounced as the temperature is raised. An increase of pressure, as would be expected from Le Chatelier's rule, is favorable for the hydrogenation process. By operating under reduced pressure, there is the double advantage of a more readily volatilization of the alcohols and a diminishing of the reverse action of hydrogenation, so consequently increasing the practical yield. At higher temperatures the aldehydes begin to decompose into carbon monoxide and a hydrocarbon. In the case of ethyl alcohol a t 42ooC, 16% of the acetaldehyde is destroyed, and the gas collected contains carbon monoxide and methane as well as hydrogen. CHICHO +CO
+ CH,
Reduced nickel acts similarly, but more energetically and at a lower temperature than does copper. Decomposition takes place from 15o0C up, while the reaction is rapid above 2 3 0 T . At 180' almost a third of the aldehyde formed is decomposed, and a t 3 3 O o C its destruction is complete. Platinum sponge acts on alcohols as does nickel, but its action does not begin till above z50°C, a t which point the acetaldehyde formed is mostly decomposed into carbon monoxide and methane. The reverse reaction, the hydrogenation of acetaldehyde, takes place readilyat 140'C with active nickel, while at 18o"the formation of methane and carbon monoxide begins. Copper on account of its weak activity, relative to that of nickel, requires a higher temperature before it exerts any action, in which case dehydrogenation of the alcohol becomes the predominating change, a t the temperature necessary for hydrogenation. For example, alcohols in contact with copper are not attacked below 230'C, a t which temperature, under ordinary pressures of hydrogen, dehydrogenation takes place. Finely divided platinum is unsuitable for the regular transformation of aldehydes into the alcohols by hydrogenation, since a t the temperatures necessary for action, which are above zoo°C, the metal acts powerfully to break up the aldehyde molecule into carbon monoxide and hydrocarbon. There has been much work done on the behavior of ethyl alcohol a t a copper surface, while with nickel there has been very little done since the mechanism was first put forward by Sabatier. Palmer' and his associates have made a rather extensive study of the activity of a copper catalyst. Thinking that a pure metal was necessary for high activity, he prepared pure metallic copper by the electrolysis of a copper salt, but found this to be catalytically inactive. However copper obtained by reduction of the oxide is active in the dehydrogenation of alcohol, and he considers that the activity Proc. Roy. 107A72.55
270
Soc., 98A, 13 (1920);99A,412 (1921);101A, 175 (1921);106A,250 (1924); (1925).
EQUILIBRIUM BETWEEN ACETALDEHYDE AND ALCOHOL
2197
is due to a kind of copper produced from cuprous compounds. There would, of course, only be “cupric” copper in the deposit on an electrode and hence, according to Palmer, this would be inactive as a catalyst. At temperatures below 3 0 0 T Palmer‘ found that pulverulent copper decomposes ethyl alcohol into acetaldehyde and hydrogen, with no formation of secondary products; but if a mixture of acetaldehyde and hydrogen is passed over the copper catalyst between 2 j o o and 3oo0C, much of the aldehyde is decomposed. The author accounts for this apparent discrepancy by assuming that alcohol is adsorbed selectively by copper from a mixture of alcohol and aldehyde vapors, so that the surface of the copper during dehydrogenation is covered by a layer of alcohol molecules, these latter preventing the adsorption and consequent destruction of the aldehyde. Palmer’ believes that the dehydrogenation of alcohol involves three stages. (I) Adsorption of alcohol molecules over the surface of the catalyst. Activation of certain alcohol molecules by adsorption of energy. (2) (3) Evaporation of acetaldehyde and hydrogen away from the adsorption surface into the alcohol stream flowing past the catalyst. Bancroft2has suggested that such cases might be called “protective poisoning.” The velocity of decomposition of alcohol at the surface of a solid catalyst has been shown by =deal* to be very much higher than the reverse reaction of hydrogenation. He states that “although the dissociation of the alcohol to aldehyde and hydrogen takes place relatively rapidly a t the commencement, yet the reverse reaction proceeds but slowly, and a long period of contact is necessary even to approximate to equilibrium conditions. According to the general concept of catalytic mechanism, an alcohol molecule striking the surface of the catalyst may either undergo decomposition or it may evaporate from the surface unchanged. For the combination of aldehyde and hydrogen, a molecule of each species has to strike adjacent molecules of the catalytic surface, a much less frequent phenomenon.” Rideal makes a calculation to show that, unless the number of alcohol molecules striking the surface and evaporating again unchanged is extraordinarily large, the reverse reaction will proceed much more slowly. Since partial poisoning may be secured, it is evidently possible that use may be made of it to achieve certain reactions whilst obviating others. A very obvious application is in the case where a catalytic process proceeds too far with an active catalyst. By minimizing the activity of such a catalyst or poisoning the more highly reactive patches with the requisite poison, the reaction may be stopped a t the desired stage. An example of this has been found by Armstrong and Hilditch? who have shown that in the catalytic dehydrogenation of ethyl alcohol in the presence of copper, the presence of water in the alcohol improves the yield of acetaldehyde relative to that of hydrogen. Proc. Roy. SOC.,98A, 1 3 (1920). J. Ind. Eng. Chem., 14, 545 (1922). Proc. Roy. SOC., 99A, 153 (1921). Proc. Roy. Soc., 97A, 259 (1920).
2198
WILDER D. BASCROFT AND AVERY B. GEORGE
Thus at 3oo0C, whereas anhydrous alcohol only gave an aldehyde-hydrogen ratio of 67 percent, alcohol containing eight percent water gave a 95 percent ratio. With the anhydrous alcohol, there is a marked increase in the yield of gaseous products, so it is evident that secondary decomposition of aldehyde is much more pronounced than with alcohol containing water. Evidently the water acts as a beneficial poison in that it poisons the catalyst for such secondary decompositions. Armstrong and Hilditch showed, by passing aldehyde vapor together with water over a copper catalyst, that the water had a considerable protective influence on the aldehyde. In the hydrogenation whilst water protects the aldehyde from decomposition] the effect of even a small proportion of water is to retard hydrogenation considerably] in fact to render it almost negligible. The authors state in regard to this that, “the protection afforded by water in the dehydrogenation process may be ascribed to the influence it exercises on the removal of the molecules of aldehyde from the sphere of action. The reason for no influence on the reverse process may be that the conditions cannot well he made the same in the two cases; in other words the behavior of a surface at which the aldehyde is being produced towards hydrogen and water may well be different from that of a surface saturated with water towards a mixture of aldehyde and hydrogen.” During the years 19 I 7 - I 9 I 9 this principle of protective poisoning was used for the production of ethylene from alcohol in presence of kaolin. Instead of a low temperature process, the reaction was conducted a t red heat and the products were protected from decomposition by using alcohol-water mixtures, even though the process was dehydration. Hoover and Rideall find that the presence of water affects the ratio in which the competing dehydration and dehydrogenation reactions occur at the surface of thoria. Armstrong and Hilditch obtained some data on the dehydrogenation of a 92 percent ethyl alcohol-water solution over nickel a t 25ooC, the aldehydehydrogen ratio being only 35.7 percent, and the evolved gas contained 60 percent hydrogen] 20 percent carbon monoxide] and 15-17 percent methane. Evidently, in this case, the 8 percent of water in the alcohol used was not sufficient to protect the aldehyde produced. Russell and Marschner* studied the effect of water on the decomposition of ethyl alcohol a t a nickel surface. They worked at 2oo0C, and found that over a wide range of concentrations the effect of water was, (a) to increase the amount of alcohol undergoing reaction, and (b) to decrezse the percentage of the aldehyde decomposed. It was the object in this work to determine the point of equilibrium between ethyl alcohol and acetaldehyde] employing finely-divided nickel as the catalyst. The effect of a platinized asbestos catalyst on this reaction was also studied under the same conditions. Then there was a general study of the reaction] bringing out certain points. The temperature employed was 140’145°c1 in which range nickel does not cause any appreciable decomposition of the acetaldehyde. 2
J. Am Chem. Soc., 49, 114 (1927). J. Phys. Chem., 34, 2554 (1930).
EQCILIBRIUY BETWEEN ACETALDEHYDE AND ALCOHOL
2199
Experimental Procedure dfnterials. Electrolytic hydrogen was used in this work, and it was further purified by passage through a tube containing platinized asbestos at a dull red heat, then over solid potassium hydroxide. Finally, to dry the hydrogen it was passed through a tube containing soda-lime and then phosphorus pentoxide. Anhydrous ethyl alcohol was used, which gave no test for water by the anhydrous copper sulphate test. The acetaldehyde employed was obtained by distilling paraldehyde after having added a few drops of concentrated sulphuric acid. The acid depolymerizes the paraldehyde to acetaldehyde, which is distilled over and condensed. The alcohol.water and the acetaldehyde-water mixtures used were prepared by diluting these liquids with distilled water. The finely-divided nickel and also the platinized asbest,os catalysts employed in this work, were very kindly furnished by Dr. Stine of E. I. du Pont de Kemours &? Company, to whom our thanks are given. Apparaizts and Procedure. The apparatus consisted of a purifying system for the hydrogen, with a flow meter interposed in the path of the hydrogen, then a flask for vaporizing the liquid, this being connected to a reaction tube heated by an electric furnace, with a system for condensing the product a t the other end. A thin layer of the catalyst was spread along the reaction tube for a length of 39 cm. Before making a run the furnace was heated to 300°C with a slow stream of hydrogen flowing through the apparatus, and this condition was maintained for thirty minutes. Then the temperature of the tube was lowered to I ~ o O - I ~ ~ O Cat, which temperature all of the runs mere made. In the case of the hydrogenation, a measured amount of the acetaldehyde was put into the flask through a funnel at the top. At the end of a run the amount of liquid left was measured, then the amount of acetaldehyde actually used could be determined. The flask containing the acetaldehyde was immersed in an ice bath, a t which temperature the vapor pressure of the aldehyde was 331 mm. The flow of hydrogen was adjusted to the desired rate, and with the delivery tube extended nearly to the bottom of the flask the run was started. The hydrogen admixed with the vapor of acetaldehyde was passed through the catalyst tube, the product being condensed and collected on emergence from the heated tube. The experiments were carried out for a period of two and one half hours, and the rate of hydrogen w1.s varied over a fairly wide range. After the vapors passing through were condensed, the liquid was collected in a small flask which mas surrounded by an ice bath. The products obtained consisting of, a mixture of alcohol and acetaldehyde were analyzed, after the weight of the product had been determined. Some work was done on devising a method for analyzing the product by specific gravity and index of refraction determinations. However the relations obtained were only qualitative in most cases, so these methods were not used. An attempt to separate the acetaldehyde-alcohol mixtures by fractional distillation did not yield quantitative results. So the acetaldehyde
2200
WILDER D. BANCROFT AND AVERY B . GEORGE
was determined in an aliquot portion of the sample by the method of Ripper.' On later comparison it was found that the results obtained from one set of specific gravity data agreed very well with the amount of acetaldehyde as found in the product by using the procedure of Ripper. So a method of analysis along this line could no doubt be worked out for this case. The method of analysis actually used consisted essentially in adding a measured excess of sodium bisulphite solution to the sample, and after a t least fifteen minutes the amount of excess is determined with iodine solution. This solution was standardized against sodium thiosulphate solution. Under the conditions of these experiments, the alcohol present in the product was obtained by difference. I n the case of the dehydrogenation of the ethyl alcohol, practically the same procedure was followed as given above. Hydrogen was passed through the apparatus until the measured quantity of alcohol was put into the flask, then the flow of hydrogen was shut off. The flask containing the alcohol was immersed in an oil bath, and the rate at which the alcohol vapors were allowed to pass over the catalyst, was controlled by varying the temperature of the oil bath. The product was collected as before in a small flask surrounded by ice, and then a delivery tube led from this into a solution of sodium bisulphite, which served to absorb any acetaldehyde that may have passed over in the vapor. An aliquot portion of the product was analyzed for acetaldehyde as before, and the necessary data obtained.
Hydrogenation of Acetaldehyde A series of runs were made on the hydrogenation of acetaldehyde employing finely-divided nickel as the catalyst, and the rate of hydrogen flow was varied between 56 cc./min. to 148 cc./min. The principle involved was to plot the percentage yield of alcohol obtained against the rate of hydrogen flow, and on extrapolating the curve, the percentage of alcohol a t zero rate or the equilibrium value could be determined. A few runs were made using the same catalyst each time, the assumption being that the activity of the nickel catalyst was the same in each case. The procedure followed was to heat the catalyst to 300°, and to maintain this temperatu.re for thirty minutes, then to lower the temperature to 14oO-145~C before starting a run, the catalyst also being cooled in an atmosphere of hydrogen after the completion of a run. The results obtained were not consistent, and they seemed to indicate that the activity of the catalyst must have undergone some change. In order to get away from this difficulty, a series of runs were made using a fresh portion of catalyst each time, in which case the various runs were comparable. The results obtained are given in the following table. Rate of Hydrogen cc./rnin.
yoAlcohol
56
94.0 91.2 90.0
74 95 148 I
Monatsheft, 21, 1079 (1900).
85.9
EQUILIBRIUM BETWEEN ACETALDEHYDE AND ALCOHOL
2201
These data were plotted and a straight line relation obtained. The mechanism of this decrease in activity of the catalyst, and to what extent the poisoning took place, was an interesting point for consideration. From some of the data obtained, the indications were that there was a rapid initial poisoning action, after which the activity of the nickel surface either remained constant or changed but slightly. This was further realized experimentally by a series of runs using the same catalyst for each run, the only treatment being the heating and cooling of the nickel surface in an atmosphere of hydrogen, before and after each experiment. The following set of data was obtained. Rate of Hydrogen cc./min.
f Alcohol
56
88.8 88.3
95
87.3 8j.6
I48
79.8
79.3
These runs were made employing a nickel catalyst which had been previously used, so it had lost its initial activity. Since the reaction takes place in or at the surface, it follows that any substance, which cuts down the rate at which the reacting substances reach the catalytic surface' or which prevents them from reaching it, will decrease the reaction velocity and may destroy the catalytic action entirely. I n this case it seems that the nickel surface possesses spheres of different activity. The more active points would be those showing the greatest adsorption and activation of the reactants. During its initial run certain points of the nickel surface must become altered and thus poisoned to any further reaction a t these points. On preparing the catalyst for another run by the treatment outlined above, certain spheres of the surface, which are no doubt the most active portions, remain unaltered during this procedure, thus giving the c a t a l p t a lower activity for this experiment. Thus after the initial run the reaction must take place a t the unaltered points of the nickel surface, these points having a lower adsorption so the activity of this surface would be less than that of the original surface. So from this it must mean that the activity is gradually decreased during the initial run with the nickel catalyst. This point has been realized experimentally, for by taking samples a t intervals during an initial run, it was found that the amount of alcohol formed decreased with increase of time. I n the second case, after this initial poisoning had taken place, the amount of alcohol formed remained practically the same throughout a run. Considerable work has been done by Palmer and Constable2 on the activity of a copper catalyst, using this reaction for study. The work dealt primarily with the activity of the catalysts with temperature. 'Taylor: Trans. Am. Electrochem. Soc., 36, 149 (1919). *Proc. Roy. Soc., PPA, 412 (1921); 101A, 1 7 5 (1922).
WILDER D. BANCROFT AND AVERY B. GEORGE
2202
Of course in our case it may be possible to restore the nickel catalyst, t o its original activity, but by the treatment given in this work, a certain portion of the surface was not reactivated. A series of runs mas next made using as the catalyst some of the nickel which had become poisoned, in other words whose initial activity had been lost. The following set of results was obtained. Rate of Hydrogen cc.:'min.
(5:
Alcohol
88
56
j
86. j
74 95 148
e6.j i9 5
1
PO
1
1
I
I
40 60 80 /OO /ZO nRM O f H Y D R O G E N C C / M / N .
140
FIG.I Hydrogenation of Acetaldehyde with Sickel
The data obtained by using some fresh catalyst each time were plotted, the percentage of alcohol against the rate of hydrogen flow, and a straight line relation was obtained. On extrapolating t,he line to zero rate, a value of 96-9 i?: alcohol was obtained which will correspond to the equilibrium value as reached from one end of the reaction. The data given above which were obtained by using some of the poisoned catalyst, were p1ott)edas in the other case, and the straight line resulting, on extrapolating to zero rate, intersected a t practically the same point as did the line obtained from using fresh catalyst each time. This means that the same equilibrium value will be obtained in each case, but that it would be reached more slowly using a poisoned catalyst, than when a fresh catalyst w'as employed. These results are what would be expected for, in t'he case of the poisoned catalyst, the reaction takes place on areas which have a lower adsorptive power than those of the fresh catalyst, and consequently the reaction velocity is less in the case of the poisoned catalyst. Eflect of V-ater. The hydrogenation of a water solution of acetaldehyde was carried out and some interesting results obtained. A 107~ and a zjyG solution of water in acetaldehyde were made up, and these two solutions were hydrogenated, the same teniperature being used as in the above cases, and the rate of hydrogen passing through being 9 j cc./min. The yields of alcohol obtained were approximately 94% and 96Yc respectfully, which shows that the presence of water increased the amount of alcohol formed, for by making a run under similar conditions using anhydrous acetaldehyde a 9oyGyield of alcohol was obtained. The presence of r a t e r was also found to increase the yield of alcohol when using a poisoned catalyst. So the water tends to re-
EQUILIBRIUhl BETWEEN ACETALDEHYDE A S D ALCOHOL
2203
activate the nickel surface, and also the activity of the catalyst does not diminish as rapidly when water is present. It may be that the action of the water is a specific action, and thus involving its preferential adsorption, on the nickel surface, to that of poisons thereby allowing more acetaldehyde to react. According to this then, the reactivation by water would mean that the water displaces the adsorbed poisons on certain points. Dehydrogenation of Ethyl Alcohol The reaction was now approached from the other side, that is, by dehydrogenating the ethyl alcohol. Some runs were made on this reaction by varying the rate of alcohol passage over the nickel catalyst from 7 cc.,:hr. to 26 cc./hr., over which range the yield of acetaldehyde obtained remained practically constant. This is in agreement with what Armstrong' found, for he stat,es that the reaction velocity of the dehydrogenation of ethyl alcohol was found to be independent of the rate of flow of alcohol vapor over the catalyst between j cc. and 3 j cc. of liquid vaporized per hour. In view of this fact all future runs were made a t the higher rate, that is 26 cc./hr. due to the experimental advantages involved. One of the important results of the work of Palmer2and his associates was that, on a given copper catalyst, the rates of dehydrogenation of the primary alcollols, ethyl, propyl, butyl, isobutyl and isoamyl, are all equal within the limits of experimental error, and the temperature coefficient is the same for all. ConstableJ3in the last of the series of these researches, showed that the reaction velocity at 2 jo"C with ethyl and butyl alcohols was independent of the pressure in the range 10-140cm. of mercury. It is evident therefore that, in this pressure range, the surface is practically covered with alcohol molecules and the mean life of the molecule in the activated unimolecular layer changes only slowly with the pressure over the range investigated. Some runs were made on the dehydrogenation of alcohol using a fresh portion of catalyst each time, the temperature being the same as in the hydrogenation of the acetaldehyde, that is, r40°-~45'C. The following table gives the results obtained. Rate of Alcohol cc./hr. % Acetaldehyde 26
2.96
26
3.08
Average-3.07~ Acetaldehyde On making some runs employing a nickel catalyst which had been previously used, a little lower yield of acetaldehyde was obtained than is given in the above table, the average value being about 2 . 6 8 Y c acetaldehyde. So in the dehydrogenation of ethyl alcohol there is also a decrease in the activity of the nickel catalyst, the mechanism of which is probably the same as in the hydrogenation of the acetaldehyde. The heat treatment of the nickel sur-
' Proc. Roy. SOC.,97A,2j9 (1920). * Proc. Roy.
SOC., 107A,255 (192j). Proc. Roy. Soc., 107A,279 (1925).
2204
WILDER D. BANCROFT AND AVERY B. GEORGE
face before and after a run was the same in the hydrogenation and the dehydrogenation processes. E$eet of Water. The effect of the presence of water on this reaction was determined by dehydrogenating a solution containing 10% water and 9 0 7 ~ ethyl alcohol, the following results being obtained. Rate of Solution Passed cc./hr.
26
R Acetaldehyde 2.97
26
3.11
Average-3 .o4Oc Acetaldehyde
So from these data it can be seen that the equilibrium in this case is not displaced by the presence of ten percent water. Armstrong’ has shown that water has a considerable protective influence on the acetaldehyde, but under the conditions of these experiments the amount of acetaldehyde that is decomposed would be very small, thus the presence of a small amount of water should have but little effect in protecting the acetaldehyde against decomposition. However the presence of water does tend to reactivate the nickel surface, as can be shown in the following manner. A nickel catalyst which had been used once, thus having lost its initial activity, was employed in the dehydrogenation of a 10% water and 90% alcohol solution. The results instead of showing about a 2.68% yield of acetaldehyde as obtained in the case of a poisoned catalyst, the percentage yield of acetaldehyde approached very closely to the value obtained when using a fresh portion of catalyst. So under these conditions the presence of a small amount of water will bring an apparently poisoned nickel catalyst approximately back to its initial activity. This was also stated by Russell and Marschner2in their work on the effect of water on the decomposition of ethyl alcohol in the presence of a nickel catalyst. They state that, “increasing the water concentration causes a considerable reactivation of the catalyst, but not as complete as with the hydrogen treatment a t 350’ C.” In this case the poisoning seems to be most complete with the anhydrous alcohol, so in the presence of water more of the active areas may be restored to their original activity due to the preferential adsorption of the water. The behavior of ethyl alcohol at a nickel surface is similar to its behavior with other catalysts. Constable3 has found that the activity of a copper catalyst decays with time when exposed to the vapors of pure ethyl alcohol a t temperatures above 28oOC. The curves given in the work of Russell and Marschner by plotting the yield of aldehyde against time are entirely similar to those obtained by Constable in the temperature range of 3o0°-3300C. The same initial rapid decrease in aldehyde yield, followed by a much slower rate of decrease, was found. Palmer‘ also observed that the catalytic activity of a supported copper catalyst decreased with use although everything was pure, ‘Proc. Roy. SOC., 97A, 2j9 (1920). J. Phys. Chem., 34, 2554 (1930). J. Chem. SOC.,1927 11, 2995. Proc. Roy. Soc., PSA, 13 (1920).
EQUILIBRIUM BETWEEN ACETALDEHYDE AND ALCOHOL
2205
however, after oxidation and subsequent reduction the initial activity again returned for about an hour of use. The curves obtained by Adkins and Millington’ show a similar behavior for ethyl alcohol with oxide catalysts. They noted that “reheating a catalyst for thirty minutes in dry air after it has been in use for an hour or two restored the percentage of ethylene to approximately its former level.” In view of the fact that, in the dehydrogenation of ethyl alcohol, the reaction velocity is independent of the rate of flow of alcohol vapor, over a fairly large range, which is not true in the hydrogenation of acetaldehyde, the dehydrogenation reaction must proceed a t a faster rate than does the hydrogenation of acetaldehyde. So in this case the equilibrium value is obtained at a much faster rate in the dehydrogenation process. The 3 percent acetaldehyde obtained by approaching the reaction from this end, corresponds very well to the equilibrium value obtained by extrapolating the hydrogenation curve, which gave 96-97 per cent ethyl alcohol. A high temperature favors the dehydrogenation reaction, so a t the relatively low temperature employed in this work 140°-1450,the equilibrium would be expected to be far on the ethyl alcohol side. The Eflect of Platinized Asbestos as Catalyst. The hydrogenation of acetaldehyde employing platinized asbestos as the catalyst, was tried under the same conditions as the above work was carried out, the temperature being 14oO-145~C. The rate of hydrogen flow was varied over a fairly wide range, but no ethyl alcohol was obtained in any case. It was thought that perhaps water would have some effect on this reaction; but, on attempting the hydrogenation of a water solution of acetaldehyde, no alcohol was formed. So under these conditions the presence of water has no effect on the hydrogenation process. However, with a platinized asbestos catalyst a t much higher temperatures, when there is a partial hydrogenation of the acetaldehyde, it might be that the presence of water would exert some appreciable effect. As no hydrogenation took place with platinum under the above conditions, it was desirable to run an experiment in order to determine whether hydrogen was taken up by an acetaldehyde solution in the presence of platinum. This was carried out in the following manner. A 50 percent solution of acetaldehyde and water was made up, and about 250 cc. of this solution put into a j o o cc. round-bottom flask. Then after platinizing a platinum cylindrical electrode, this was put into the flask so as to have about two-thirds of it immersed in the solution. The air was displaced from the flask by passing in hydrogen, then a mercury manometer was attached to the flask, into which hydrogen was now forced, and the pressure read on the manometer. The apparatus mas allowed to stand, and if any hydrogen was taken up, it could be recorded by the drop in pressure as shown by the manometer. The results were, that hydrogen was apparently taken up, as shown by the drop in pressure of the manometer, but very slowly, the pressure drop being about one cm. of mercury 1
J. Am. Chem. SOC., 51, 2455 (1929).
2206
WILDER D. BANCROFT AND AVERY B. GEORGE
in three to four hours. However, this did not tell exactly how the hydrogen was taken up, so another experiment was made in the same way as the above one except that no platinum electrode was used. The object of this experiment was to show whether hydrogen was absorbed by the acetaldehyde-water solution. It was found that hydrogen was taken up, but more slowly than in the first case, the pressure drop being about one cm. of mercury in seven hours. It is also known that platinized platinum adsorbs hydrogen, but even though hydrogen was taken up in these two cases, the process was slow. So apparently the hydrogen did not reduce any of the acetaldehyde to ethyl alcohol, for no test could be obtained for alcohol from the solution, after being in an atmosphere of hydrogen for a long time. The reduction of acetaldehyde electrolytically was next taken up. Formaldehyde and acetaldehyde, the simplest of the aliphatic aldehydes, under certain circumstances are almost quantitatively reduced in alkaline solution a t a copper or silver cathode. On the other hand in an acid solution they are not so easily reduced. Platinum and silver cathodes are without action, but copper or mercury cathodes lead to the formation of small quantities of methane when formaldehyde is the starting material, or to ethane when the acetaldehyde is used. This formation of a hydrocarbon is enhanced by the use of a cadmium cathode, and so pure propane may be prepared in good quantities from the propionaldehyde, but not in such large quantities as are obtained by the reduction of acetone.' In the reduction of aliphatic aldehydes, mercury and lead electrodes as a rule give excellent results, but the yields are always lower than those obtained with a cadmium electrode. In the use of lead and mercury cathodes with aldehydes no metal alkyl compounds are formed, as in the reduction of acetone at these electrodes.? The industrial production of acetaldehyde and paraldehyde from acetylene has led to the development of several interesting methods for their reduction to a variety of products, that are of immense value, especially alcohol and ethyl acetate. These methods are covered by patents, and the electrolyte is 5-10 percent sulphuric acid, sodium sulphate, or orthophosphoric acid. The presence of a mercury salt, which acts as a catalyst, in the solution permits the passage of the acetylene directly into the electrolyte without being converted into acetaldehyde by a previous and separate p r o ~ e s s . ~The formation of alcohol requires the use of a diaphragm to prevent anodic oxidation, but ethyl acetate is so stable that no diaphragm is needed when it is the end product. The cathode material should be of lead or mercury. The temperature, the acidity, the concentration of the unreduced aldehyde, the current density and the duration of the reaction must be held as low as possible, as otherwise there will be too large a formation of undesired by-products. In this work an attempt was made to reduce acetaldehyde electrolytically a t a platinized-platinum cathode. Two experiments were made, one using a 50 percent and the other a I O percent solution of acetaldehyde in water, with F. Muller: Dissertation, Dresden (1921).
* Brockman: "Electro-Organic
Chemistry," page 28 j . British Patents, 140115(1918); 140527 (1919);Pascal: Swiss Patent, 88188 (1921).
EQUILIBRIUM BETWEEN ACETALDEHYDE AND ALCOHOL
2207
a little I O percent sulphuric acid as the electrolyte. The acetaldehyde solution was put into a porous cup, with a cylindrical platinized-platinum electrode immersed in the liquid. A I O percent sulphuric acid solution was used as the anode liquid, with a platinum flag as electrode. On passing a current through the solution, gas was evolved immediately a t the two electrodes, this occurring in both concentrations employed. Of course this means that if there is any reduction it must be very inefficient, but these runs were allowed to go for a sufficient length of time to get some reduction, if such a process took place. A current density of about 0.01amperes/sq. cm. of cathode surface was used. The cathode liquid was tested for ethyl alcohol by using the Schotten-Baumann reaction, the odor of ethyl benzoate being detected in the presence of ethyl alcohol. However, no positive test could be obtained, so it seems that the acetadehyde is not reduced electrolytically, to ethyl alcohol, with a platinized-platinum cathode. The dehydrogenation of ethyl alcohol over a platinized-asbestos catalyst a t 140'-145' C was carried out. However, the alcohol passed over the catalyst apparently unchanged, as no acetaldehyde could be detected in the resulting product. So a platinized asbestos catalyst will cause neither the hydrogenation of acetaldehyde, nor the dehydrogenation of ethyl alcohol a t the temperature employed in this investigation. T h e Effect of adding Oxygen in the Dehydrogenation of Ethyl Alcohol using a Nickel Catalyst. The dehydrogenation of ethyl alcohol takes place according to the following reaction.
+
C2HSOH CHsCHO Hz It was thought that by adding oxygen to the reaction, the oxygen and the hydrogen would unite to form water, so by removing one of the products of the reaction it would be displaced towards the right with the formation of more acetaldehyde than the equilibrium value. A few experiments were carried out, in the same way and under the same conditions as the dehydrogenation runs above, except that a slow rate of oxygen was passed into the reaction tube during the runs. The following data shows how the yield of acetaldehyde obtained varied with the rate of oxygen passage, all other conditions remaining the same. Rate of Oxygen
Acetaldehyde
8 cc./min.
4 . IO
cc./min. 5 5 cc./min. 2 2 cc./min.
5.30 3.80
22
5.35
The equilibrium value as found above was 3 percent acetaldehyde, so while the amount of acetaldehyde formed was increased by the addition of oxygen, the extent of this increase was not particularly alarming. The first two rates showed an increase in the yield of acetaldehyde, but with a very fast rate, 5 5 cc.,/min., the yield of aldehyde dropped off. This may be due to two causes, first the oxygen may carry the alcohol over the catalyst at such a rapid
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WILDER D. BANCROFT AND AVERY B . GEORGE
rate as to prevent contact with the catalyst for reaction, and secondly this rate of oxygen is much in excess of the amount of hydrogen given off from the alcohol, so the catalyst is probably acting as a nickel oxide surface rather than that of metallic nickel. This experiment was also made in order to determine whether alcohol was oxidized to acetaldehyde over the catalyst in the presence of an excess of oxygen, but as the yield of acetaldehyde decreased this was not the case. A second run was made at the slower rate in order to determine whether or not check runs could be obtained, and the data show that the run was substantiated. So in this work, it has been shown that a slow rate of oxygen passing into the reaction tube over the catalyst will increase the practical yield of acetaldehyde, but this increase is not very great. T h e E$ect of Excess Hydrogen tn the Dehydrogenatzon of Ethyl Alcohol using a A'zckel Catalys,. The dehydrogenation of ethyl alcohol was carried out a t 180°C,a t which temperature a certain amount of the acetaldehyde formed was decompospd into carbon monoxide and methane. Tk.3 liquid product was collected, as was also the evolved gas, the latter being collected in a large bottle acting as a gas holder. After making an analysis of the products it was found that a little less than one third of the acetaldehyde was decomposed into gaseous products. This is in agreement with the results of Sabatier,' for a t 178°C he obtained some acetaldehyde, but about a third was decomposed into carbon monoxide and methane. Now another run was made, and this time a slow stream of hydrogen was passed over the catalyst, in order to find out whether or not the decomposition of the acetaldehyde was diminished by the presence of the excess of hydrogen. However, the results show that the decomposition of the acetaldehyde is practically the same in each case, so the hydrogen apparently has no protective influence in this case. The percentage yield of acetaldehyde was less in the latter case which would be expected, as an excess of hydrogen would tend to suppress the dehydrogenation reaction, and favor the reverse process. The actual amount of decomposition gases is less in the latter case, but the percentage decomposition is about the same, for there is less acetaldehyde formed in the presence of an excess of hydrogen. summary
I. The equilibrium value of the ethyl alcohol-acetaldehyde reaction was found to correspond to 96-9770 ethyl alcohol, when acetaldehyde was hydrogenated a t different rates, this value being obtained on extrapolating the percent ethyl alcohol-rate of hydrogen flow curve to zero rate of flow. On approaching the reaction from the other side, that is the dehydrogenation of the ethyl alcohol, an equilibrium value of 3 percent acetaldehyde was obtained, which checked very well with the value from approaching it by the hydrogenation of acetaldehyde. The nickel catalyst employed showed a rather rapid initial poisoning 2. action, after which the activity of the nickel surface remained practically Compt. rend., 136, 738 (1903).
EQUILIBRIUM BETWEEN ACETALDEHYDE AXD ALCOHOL
2209
constant. So it seems that the nickel surface must possess spheres of different activity, the more active points would be those showing the greatest adsorption and activation of the reactants. The more active points would be those most easily poisoned, so after the initial run the reaction must take place a t the unaltered points of the nickel surface, these points having a lower adsorptive power, so the activity of this surface would be less than that of the original surface. 3 . On extrapolating the plotted data obtained in the hydrogenation of acetaldehyde, using some catalyst which had lost its initial activity, the line intersected the zero rate line a t practically the same equilibrium value as when using fresh catalyst. This means that the same equilibrium value will be obtained in each case, but that it would be reached more slowly when using a poisoned catalyst. 4. The presence of water has an activating effect on the nickel surface, from whichever side the reaction is approached. 1, the hydrogenation of acetaldehyde the presence of water increased the amount of alcohol formed. It may be that the action of the water is a specific one, and thus involving its preferential adsorption, on the nickel surface, to that of poisons thereby allowing more acetaldehyde to react. In the case of the dehydrogenation, while the presence of I O percent water did not appreciably displace the equilibrium, it did have an activating effect on a used nickel catalyst. 5 . A platinized asbestos catalyst will cause neither the hydrogenation of acetaldehyde nor the dehydrogenation of ethyl alcohol a t the temperature employed. It has also been shown that acetaldehyde is not reduced to ethyl alcohol electrolytically with a platinized-platinum cathode. 6. In the dehydrogenation of ethyl alcohol, the yield of acetaldehyde can be increased by passing a slow rate of oxygen into the reaction tube over the catalyst, but the amount of this increase is not very great. 7. Passing an excess of hydrogen over the nickel catalyst in the dehydrogenation of ethyl alcohol, does not appreciably affect the percentage of acetaldehyde that is decomposed into the gases carbon monoxide and methane, Cornell Unzuerszty
