T H E INTERMEDIATE STAGES OF ALDEHYDE OXIDATION. I
THE CATALYTIC ACTIONOF MANGANESE CATALYST IN THE VARIOUS OF THE PROCESS OF ACETALDEHYDE OXIDATION STAGES M. J. KAGAN AND G. D. LUBARSKY Received August IS, I934
In the autoxidation of acetaldehyde we are concerned with the reaction of the formation of peracetic acid. According to Bodenstein (2), the following chain mechanism takes place 0
0-
CH8CH II -+ C H 3 ~ H
( 1)
I 0
0CHaAH
I
0
/ \ 0 \ /
CH3CH
+
0 2
/ \ 0 = CH&H \ /
(2)
0
0
II + CHICH
0-
I I
= CHICH
0
0
II + CH&OOH
(3)
Molecular oxygen adds to the active form of aldehyde (l),producing the active form of peracetic acid (2), in agreement with the theory of Bach (1) and Engler (4). The latter reacts with a normal molecule of acetaldehyde forming again an active aldehyde molecule and the final product (3). The peracetic acid obtained can in turn oxidize aldehyde further, forming acetic acid. The process of oxidizing acetaldehyde by oxygen directly to acetic acid in the presence of a catalyst is widely used in industry. For the investigation of the mechanism of this process, the study of the kinetics of not only the first stage of the formation of peracetic acid is important, but also the second stage, the interaction of peracetic acid with acetaldehyde. We have made a special study of th- second stage of this process (see the following article), and have found that the interaction between peracid and aldehyde proceeds also in two stages (a) and (b): (a) The addition of a molecule of aldehyde to the peracetic acid and the formation of an inter837
838
M. J. KAGAN AND G. D. LUBARSKY
mediate product of a peroxide character. The activation energy of this bimolecular reaction is 7000 calories, and its rate is higher than that of the second stage. (b) The decomposition of the intermediate product with the formation of two molecules of acetic acid. The activation energy of this monomolecular reaction is 15,000 to 16,000 calories. The oxidation of acetaldehyde by oxygen to acetic acid at both low and room temperatures takes place at a slow rate, and in the absence of catalysts leads to the accumulation of peroxides. Hence it follows that the transformation of the peroxides to acetic acid proceeds more slowly than their formation. Bodenstein determined the activation energy of the first stage of acetaldehyde oxidation-the chain reaction of peracetic acid formation in the gaseous phase at about 10,000 calories. In view of these facts we came to the conclusion that the study of the action of manganese catalyst, which is successfully applied for the oxidation of acetaldehyde to acetic acid, must follow the line of investigating the kinetics of the entire process of oxidation of acetaldehyde to acetic acid in the presence of a catalyst and of the influence of the catalyst on the kinetics of the separate stages of the process. I. THE OXIDATION OF ACETALDEHYDE TO ACETIC ACID USING A MANGANESE
CATALYST.
ELIMINATION OF THE INDUCTION PERIOD O F THE REACTION
The experiments were carried out in the following manner: A known solution of acetaldehyde in acetic acid was sucked into a small vessel, provided with a stopcock and a small blank tube, where a small amount of the catalyst (a manganese salt) was placed. The vessel was previously evacuated and tared, and then connected by a piece of thin rubber tubing with a bomb of oxygen (2 liters), equipped with a manometer. Both the bomb and the reaction vessel were placed in a water thermostat. The vessel was then fastened to a vertical shaking machine (length of stroke, 5 cm.), and shaken in the thermostat at a rate of about 500 strokes per minute. The particles of the catalyst were washed down at the first stroke, and this moment was fixed as the beginning of the experiment. As the reaction proceeded, fresh portions of oxygen were automatically sucked from the bomb into the vessel, owing to the decrease in pressure above the liquid in the vessel. In the experiments which were carried o u t at constant pressure, the readings were taken on a buret with an automatic clamp connected with a levelling vessel (figure 1). At the beginning of the experiment the stopcock in the right bend of the manometer (C) is closed. In proportion to the absorption of oxygen, the pressure of the system falls, making contact (K), and the electromagnet E weakens the clamp. Then a small amount of water runs from the buret (F) into the beaker (D), and the pressure in the system increases to the given level, after which the mercury in the left bend falls and the electromagnet breaks contact.
839
INTERMEDIATE STAGES OF ALDEHYDE OXIDATION. I
In spite of the large number of experiments which were carried out, it was not possible to determine exactly the reaction velocity constants, owing to the turbulent character of the reaction which took place highly exothermically in a short time so that the vessel had no time to cool down to the temperature of the thermostat. This accounts for the qualitative character of the results obtained, which nevertheless are of fundamental interest. The commencement of the decrease of the oxygen pressure always coincided with the appearance of a dark brown coloration of the solution.' In some experiments the coloration did not appear at all, and the reaction did not take place; in others it appeared some time after the beginning OMPRESSED AIR
I
'
0
I
2
FIG.1
3
5 6 MINUTES
4
7
8
3
10
11
FIG.2 FIG. 1. THEAPPARATUS V, to vacuum pump
FIG. 2. THE OXIDATIONOF ACETALDEHYDE IN ACETIC ACID SOLUTION I, catalyst Mn(NO& a t 25"C., induction period 43 minutes; 11,catalyst Mn(NO& a t 25"C., induction period 73 minutes; 111, catalyst (CH&OO)sMn a t 15°C.; IV, catalyst (CHsC00)3Mna t 15OC. in acetic acid and 10 per cent water solution, induction period 62 minutes.
of the shaking (3-75 minutes), and only after its appearance did the reaction and a sharp decrease of pressure in the system take place. Thus the added salts of bivalent manganese (acetate, nitrate2) do not appear to be catalysts. On the contrary the catalyst is formed in the aldehyde solution which is being oxidized, and only after its formation does the high reaction rate take place. The composition of the brown solution acting as catalyst 1 The formation of the brown solution had been observed by Baeyer and Williger when introducing manganese acetate into a solution of perbensoic acid. * In a paper by M. J. Kagan and N. M. Moroaov (6) it was shown that the activity of manganese nitrate as a catalyst of this process is 50 per cent greater than t h a t of manganese acetate.
840
nf.
J. KAGAN AND G. D. LUBARSKY
in this reaction was found to be an acetic acid salt of trivalent manganese. I n the presence of this compound no induction period, which occurred in other experiments, was observed, as is shown by curves in figure 2. The accumulation of this compound evidently takes place during the induction period. It is obvious that anything causing the decomposition of manganic acetate, (CH3C00)3Mn, must tend to lengthen the induction period. Thus, it was shown that the presence of water causes an induction period. Water decomposes maganic acetate, giving a precipitate of manganese hydroperoxide. When about 10 per cent of water was added to the solvent (acetic acid), a prolonged induction period was produced, even in the case when manganic acetate was used as catalyst (see figure 2, curve IV). The oxidation of bivalent manganese to the trivalent form occurs under the action of peracetic acid, which is formed during the oxidation of acetaldehyde by oxygen. An acetic acid solution of manganic acetate may be readily prepared by the action of peracetic acid on manganous acetate, the peracetic acid being decomposed very turbulently. This fact cannot but influence the mechanism of the oxidation process (see below). The preparation of manganic acetate and its isolation in a pure state have been stated in detail by Christensen (3). 11. INFLUENCE OF THE CATALYST ON SEPARATE STAGES OF THE PROCESS
A . T h e injluence of the catalyst o n the rate of the interaction between peracetic acid and acetaldehyde The determinations of the velocity of the summary reaction CH3COOOH + CHsCHO + 2 CH3COOH with manganese catalyst are only of a qualitative character, owing to the poor reproducibility of the numerical values of the velocity constants in different experiments. However a sharp increase of the velocity with the addition of very small amounts of the catalyst has been positively determined (see figure 3). As shown above in the case of interaction between aldehyde and peracetic acid, the formation of the addition product is the faster, while the slower stage, limiting the reaction rates, is the decomposition of the intermediate product into two molecules of acetic acid. By studying the kinetics of the decomposition of this product in the presence of the manganese catalyst, we were able to obtain more reproducible quantitative results (shown in table 1). The experiments were carried out in the following manner: The mixture of peracetic acid and aldehyde dissolved in toluene was left standing for several hours at a temperature of -3O"C., as a result of which an addition product of a peroxide character was formed. For the quantitative determination of this product in the presence of peracetic acid see paper 11. The kinetics of the decomposition
INTERMEDIATE STAGES O F ALDEHYDE OXIDATION. I
84 1
of this product at room temperature without a catalyst was studied. Then, a t a definite moment, a solution of manganic acetate in chloroform
l
.
.
M 20
.
M
.
.
.
.
.
.
40 50 60 70 'b0 90
.
MO
.
.
.
110 120 13
10 20
33 10 50 60 70 80 90 100 110 120 MINUTES
IIINUTES
FIG.3 FIG. 4 FIQ.3. THECATALYSIS OF INTERACTION BETWEEN PERACETIC ACID AND ACETALDEHYDE Points A correspond to the moments of pouring of the catalyst, 6 X g. FIG. 4. CATALYTIC DECOMPOSITION O F THE ADDITIONPRODUCT (PERACETIC ACID AND BENZALDEHYDE) I N TOLUENE SOLUTION I, experiment No. 119 at 10°C.; 11, experiment No, 120 a t 20°C.; 111, experiment No. 121 a t 30°C. (high concentration). TABLE 1 Decomposition of the intermediate product (peroxide) obtained at -30°C. i n a toluene solution t =
Time from the
+ 20°C.
II
t =
+ 10°C.
Time from the
'
minutes
(a) Without catalyst
0 13 73 123
0.4015 0.3791 0.2895 0.2334
7.37 X 7.49 X 10-6 7.18 X 10-5
(b) With catalyst (manganic acetate = 0.0008 g.)
0 4 8 17
0.2005 0.1504 0.1169 0.06466
1.2 X 10-3 1.05 X 10-3 1.1 X 10-3
(a) Without catalyst 0 21 51 86 116
0,2840 0.2711 0.2597 0,2457 0.2351
2.644 X 10-6 2.380 X 10-6 2.641 X 10-5 2.436 x 10-6
(b) With catalyst (manganic acetate = 0.0009 g.) 0 4 10 22
0.2130 0.2004 0.1830 0.1548
2.545 X 10-4 2.825 X 10-4 2.325 X lo-'
was introduced. The reaction rate was estimated by the disappearance of the peroxide, the content of the latter being determined iodometrically.
842
M. J. KAGAN AND G . D. LUBARSKY
These data show a considerable increase in the reaction rate of this stage under the influence of the catalyst, and we are able to state quite definitely the catalytic effect of manganic acetate on the rate of this very slow reaction. Thus we have reason to suppose that the part played by the catalyst consists in accelerating the rate of decomposition of the intermediate peroxide product with the resulting speeding up of the whole oxidation process. On the other hand, a still greater catalytic effect is observed during the decomposition of the peracid, which points to the possibility of another mechanism of the catalysis.
T h e decomposition of peracetic acid under the action of manganous acetate The study of the decomposition of peracetic acid in the presence of a manganese salt is of some interest in explaining the partial formation of carbon dioxide in the industrial oxidation of acetaldehyde. Moreover, as was pointed out, it may considerably influence our conception of the TABLE 2 Decomposition of peracetic acid GAS CONTENT
NO. OF EXPERIMENT
coz
co
per cenl
per cent
87.4 88.9 88.3 93.0
11.8 11.1 11.1 7.0
mechanism of the oxidizing action of peracetic acid. The decomposition was carried out at room temperature in acetic acid and water solutions in a small vessel of about 10-cc. capacity with a stopcock, connected by a piece of rubber tubing with a Hempel buret. A definite amount of the peracid solution was poured into the vessel, while manganese acetate particles were placed in the closed side-arm tube. On shaking the vessel the particles of catalyst dropped into the solution, whereupon the evolution of gas began and the liquid became strongly heated. The gas was collected in the Hempel buret and afterwards analyzed for carbon dioxide, oxygen, and carbon monoxide. The results of the peracid decomposition in an acetic acid solution are very interesting. The gas evolved during the decomposition contained no oxygen, being composed chiefly of carbon dioxide and carbon monoxide, with a small percentage of hydrocarbons. The data given in table 2 are characteristic of such a process of decomposition. Immediately after coming into contact with manganous acetate, the solution acquired a dark brown color,
INTERMEDIATE STAGES O F ALDEHYDE OXIDATION. I
843
due to the formation of manganic acetate. On diluting this solution with water, a brown-colored residue of manganese hydroperoxide is instantly formed A quite different process of peracetic acid decomposition was observed in a water solution; the process took place smoothly, slowly, and without any perceptible heating. When small quantities of manganous acetate were added, no manganese dioxide residue was formed, but the solution acquired the characteristic permanganate color. In case of an excess of manganous acetate the solution acquired an orange-red color, and after standing several hours or days (according to the quantity of manganous acetate) a precipitate formed. It must be noted that the decomposition of the peracid in the presence of potassium permanganate solution also proceeds very slowly. The titration of the peracid solution decomposed in the presence of potassium permanganate was as follows: 0.2 CG. of peracid solution = 4.5 cc. of thiosulfate solution (0.02 N ) ; on the following day = 3.1 cc.; on the third day = 2.6 cc. The gas evolved during the reaction
.
TABLE 3 Decomposition of peracetic acid QAS CONTENT
c02
co
0 2
per cent
per cent
per cent
27.1 30.8 25.1
15.6 10.7 11.3
57.2 58.4 63.5
under these conditions was composed chiefly of oxygen, i.e., it had quite a different composition from the gas evolved during the decomposition in the acetic acid solution (see table 3). Therefore, we may conclude that oxygen forms the initial decomposition product of the peracid decomposition, which during its slow evolution is able to combine into molecules of oxygen, whereas during a turbulent evolution it oxidizes the molecules of organic compounds to carbon dioxide and carbon monoxide :
o+o=oz
:H3COOOH + CHaCOOH
7
+0
I
CH3COOH
H2O
+
1
+ 3 0 = C02 + HzO + HCOOH
These data are in agreement with Hatcher and Tool (5), who observed that peracetic acid is decomposed when heated, forming carbon dioxide, formic acid, and glycolic acid.
CO
844
M. J. KAGAN AND G . D. LUBARSKT
The determination of the kinetics of the decomposition of peracetic acid in an aqueous solution in the presence of manganous acetate showed that this process proceeds as a reaction of the first order and a t a very slow rate, the decomposition beginning a t the moment of the appearance of a pink hue of the solution (apparently due to the formation of permanganic acid). Comparison of the kinetic data obtained may be seen from table 4. The activation energy of this process is expressed by a value approximating 12 kg-cal. TABLE 4 Decomposition of peracid i n aqueous solution i n presence of manganous acetate (0.0048 8 . ) Experiment 88
0 19 85 98 138
0.3668 0.3241 0.2273 0.2116 0.1673
10 9 8.96 9.17 9.8
t = f20"C. 0 35 77 113
0.1217 0.1177 0,1111 0.1058
1.61 2.28 2 26
Observations analogous to ours were made by Milas (7) on investigating the decomposition of furo-peracid:
Solid furo-peracid was decomposed in a vacuum with the formation of a gas consisting only of carbon dioxide. In the residue was furo-acid and tarry polymer. In a weak solution of chloroform a monomolecular decomposition took place with the formation of oxygen and furo-acid. The energy of activation = 15,800 calories. I n comparing the above experimental data in the first as well as in the second version, the great influence exerted by the catalyst is observed. In the first version, in the absence of the catalyst, we observe a high rate
INTERMEDIATE STAGES OF ALDEHYDE OXIDATION. I
845
for the reaction of the addition of aldehyde to peracid, and a low rate of the decomposition of the addition product into 2 molecules of acetic acid. With manganese catalyst this slow process is accelerated about twenty times. It seems logical to infer that the action of the catalyst consists in the acceleration of the slowest stage, and that the scheme of the oxidation process remains the same whether it takes place in the presence or the absence of the catalyst.
A
+ o2= P; A + P = AP; AP
= 253
However, according to the above kinetics data, the decomposition of peracetic acid, using a manganese catalyst, proceeds with evolution of oxygen; only'in the absence of aldehyde, which is the acceptor of oxygen, does the combustion of acetic acid take place. The decomposition of peracetic acid dissolved in glacial acetic acid occurs at an exceptionally high rate, and when aldehyde is present it may be oxidized very rapidly to acetic acid. It follows that the catalytic process may also proceed according to another scheme viz. : A
+
0 2
=
P; P = S + 0 ;A
+0 =S
We showed above the scheme of Bodenstein for the chain mechanism of aldehyde autoxidation. The property of peracetic acid of decomposition under the action of manganese catalyst cannot but influence the chain process. At the same time the generation of the oxygen atoms may lead to a mechanism of a type differing from that proposed by Bodenstein. The elimination of the induction period in presence of manganic acetate catalyst is of great interest. The induction period consists in the slowing down of the oxygen absorption, Le., the delaying of the peracetic acid formation (which oxidizes Mn++ to 'Mn+++). By introducing manganic acetate, the conditions are created for the reaction of the further transformation (decomposition) of peracetic acid, at the same time eliminating the induction period. It is hoped to resume this problem later in case of other oxidation reactions. SUMMARY
1. The oxidation of acetaldehyde by oxygen in an acetic acid solution in the presence of manganese salts proceeds vigorously after the formation of a dark brown catalytic product. 2. In this process the catalytic product is trivalent manganese acetate obtained by action of peracetic acid on &In++. 3. One of the causes of the induction period of the reaction is the presence of water which decomposes the catalyst manganic acetate, forming 3
A = aldehyde; P
=
peracetic acid; S = acetic acid.
846
M. J. KAGAN AND G . D. LUBARSKY
manganese hydroperoxide. In the absence of water the induction period is eliminated by the addition of a trivalent manganese salt as catalyst. 4. The rdles of the manganese catalyst may be two: on one hand it hastens the decomposition of the intermediate peroxide (product of addition of peracetic acid to aldehyde) leading to the formation of acetic acid; on the other hand it decomposes peracetic acid with the separation of active oxygen, which directly oxidizes aldehyde. 5. The formation of carbon dioxide during the oxidation of acetaldehyde by oxygen is due to the partial decomposition of peracetic acid under action of trivalent manganese. The gaseous products of peracetic acid decomposition consist chiefly of carbon dioxide ( m 90 per cent). The decomposition in the acetic acid solution takes place very turbulently and exothermically. 6. The decomposition of the peracid in the aqueous solution after the addition of manganese acetate takes place at a very slow rate. The gaseous decomposition products are chiefly oxygen (- 60 per cent). The decomposition process is a first-order reaction and requires an activation energy approximating 12,000 calories. 7. Owing to the catalytic decomposition of peracid in the presence of aldehyde in the acetic acid solution no combustion product except acetic acid is formed, i.e., the oxygen atom oxidizes the aldehyde molecule. There seem to be good reasons for believing that such a scheme of the catalytic oxidation of aldehyde is very probable. REFERENCES (1) BACH:Compt. rend. 124, 951 (1897). (2) BODENSTEIN: Ber. Berl. Acad. 111,.71 (1931). (3) CHRISTENSEN: Z. anorg. Chem. 27, 321 (1901). (4) ENOLER: Ber. 30, 1669 (1897); Z . Elektrochem. 18, 948 (1912). (5) HATCHER AND TOOL:Trans. Roy. SOC. Canada [3] 20, 415 (1928); Canadian J. Research 7, 149 (1932). (6) KAOANAND MOROZOV: Zhur. PrikladnoI Khim. 6, 400 (1932). (7) MILAS:J. Am, Chem. SOC. 66, 1221 (1934).
