T H E INTERMEDIATE STAGES O F ALDEHYDE OXIDATION. 11 KINETICSOF
THE
INTERACTION BETWEEN PERACETIC ACIDAND ALDEHYDES
G. D. LUBARSKY
AND
THE
M. J. KAGAN
Received August 19, 198.4
The study of the intermediate stages of aldehyde oxidation was limited chiefly to investigations of the first stage of oxidation, i.e., the formation of peracid from the aldehyde molecule and the oxygen molecule. The kinetics of the second stage of the process, the interaction between the peracid and aldehyde, RCOOOH
+ RCHO + ZRCOOH
remained unstudied. In one of his recent papers Wieland (5) studied the process of interaction between the peracids and aldehyde. He came to the conclusion that in case of a reaction between peracetic acid and aldehyde, the reaction may be carried out with a marked velocity only in the presence of water, when aldehyde hydrate is formed. The latter is subsequently dehydrogenated by the oxygen of the peracetic acid.
0
//
HrCC
\
H
/
OH
+ H20 = HaCC-OH \
0
/OH ; HaCC-0 H \
H
//
+ 0 = HsCC
'H
\
+ HzO
OH
I n the present paper, the kinetics of interaction between peracetic acid and aldehyde is studied. It was found that the formation of two molecules of acid from a molecule of aldehyde and a molecule of peracetic acid involves an intermediate reaction, viz., the formation of an addition product of the reacting molecules having a peroxide character, and that water catalytically accelerates the decomposition of this product. At the same time the kinetics of each intermediate stage was studied. I. EXPERIMENTAL PROCEDURE
Acetaldehyde was obtained by the oxidation of alcohol by air on the surface of a copper catalyst. The pure aldehyde, redistilled several times and boiling a t 20.5"C., was sealed in tared thin-walled glass ampullae. 847
848
G . D. LUBARSKY AND M. J. KAGAN
The ampullae were weighed with the capillary glass tips sealed off; thus the weight of the aldehyde was determined to the fourth decimal. The sealed ampulla was used for the reaction during the same day; in this manner freshly distilled aldehyde was always used. The peracetic acid was obtained by action of concentrated hydrogen peroxide on acetic anhydride (3) in presence of a few drops of sulfuric acid. At the beginning the reaction proceeds very rapidly, therefore it is necessary to cool the acetic anhydride to 0°C.and then gradually to add the hydrogen peroxide drop by drop especially a t the beginning. The hydrogen peroxide (Kahlbaum, 27 per cent HzOz)was first concentrated by distillation. The residue in the distilling flask was generally found t o be 80 per cent hydrogen peroxide, which was used in preparing peracetic acid. Peracetic acid was carefully distilled in vacuum (10-15 mm.) a t 24-25"C., the middle fraction being collected; in this manner a 50 per cent solution of peracetic acid in acetic acid was generally obtained, which contained almost no diacetyl peroxide. When kept in the dark a t room temperature such a solution decomposes very slowly, so that even after several months the concentration changes very little. The glacial acetic acid (f. p., 16.4"C.)used as solvent was obtained by repeated crystallization of a distilled commercial glacial acetic acid. The other solvents-benzene, chloroform, nitrobenzene, toluene-were obtained chemically pure from Kahlbaum. The investigation of the interaction between peracetic acid and aldehyde was carried out in a small closed cylindrical vessel about 120 mm. in height and 20 mm. in diameter. The vessel was placed into a thermostat, the temperature of which was regulated with an accuracy of 0.02"C. The small cross section of the reaction vessel facilitated the maintenance of a uniform temperature of the liquid. The vessel was filled with a definite amount of peracetic acid solution. I n the reaction vessel was placed a sealed ampulla with acetaldehyde, which was broken a t a given moment taken as the beginning of the reaction. At the same time another vessel with a solution of peracid of the same concentration was placed in the thermostat and stability of the acid was periodically tested, The portions from the reaction and the test vessels were pipetted with a micropipet (generally 0.1 cc.) into a previously weighed beaker containing a 10 per cent solution of potassium iodide, and after weighing (which never exceeded three minutes) the iodine evolved was titrated with 0.02 N sodium thiosulfate solution (2). The small additional amount of iodine evolved in standing under action of an admixture of diacetyl peroxide was generally the same for all portions, and was disregarded. When experimenting in neutral solutions (benzene, toluene, etc.) 2 CC. of 1 N sulfuric acid was added to the potassium iodide solution. The
849
INTERMEDIATE STAGES OF ALDEHYDE OXIDATION. I1
concentration of acetaldehyde in the solution generally did not exceed 2-3 per cent, which corresponds to a vapor pressure of 20-60 mm. in the temperature range used (1O-3O0C.). These results remained unchanged if the reaction was carried out in an atmosphere of nitrogen. The investigation was carried out under ordinary conditions in a closed vessel.
MOLES PER LITER TIME INTERVAL IN MINUTES
NO, OF PORTION
Peracid
K (mole-1om.asec-1)
I
Aldehyde
I
0.9043 0.8488 0.7426 0.6725 0.6284 0.6097
t = 288°C.
5
I
1.5 30 57 54 31
1
0.4266 0.3711 0.2649 0.1948 0.1507 0.1320
I
1.77 0.236 0.127 0.122 0.115
t = 293°C. 0
1 2 3 4
21 20 20 15
0,1235 0.1071 0.09443 0.08247 0.07462
0.6006 0.5848 0.5722 '0.5612 0.5533
0.192 0,181 0.202 0.200
11. KINETICS OF THE INTERACTION OF PERACTIC ACID WITH ACETALDEHYDE I N NON-AQUEOUS SOLVENTS
A . Reaction in acetic acid solutions The interaction between peracetic acid and acetaldehyde was studied a t a temperature of 10-30°C., the concentration of the solution of peracid being varied from 0.5 to 3.3. peracid The results obtained can be described with a sufficient degree of accuracy by a second-order kinetic equation if we disregard the values obtained by calculation of the first portions taken two to ten minutes after the beginning of the reaction; as a rule these values are very high as compared with the following ones. We shall return to this point below. Some typical experiments in acetic acid solution are given in table 1. In table 2 the constants of the reaction rate in dependence on the temperature are given (the initial constants being disregarded). and the ratio
850
G . D. LUBARSKY AND M. J. KAGAN
The temperature effect is expressed by the Arrhenius equation (figure 1). The activation energy of the process is 15,400 7 300 calories.
B. Reaction in other non-aqueous solutions I n order to determine the specific influence of the solvent on the reaction rate, a series of experiments was carried out using benzene, chloroform, and nitrobenzene solutions. Preliminary test experiments showed that in these
TABLE 2 Kinetics of interaction between acetaldehyde and peracid in a n acetic acid solution t
ER CENT OF PERACII
4c. 10 10 15 15 20 20 25 25 30 30
2.96 3.45 4.91 3.38 3.12 3.38 4.8 4.77 2.5 4.32
ALDEHYDE PERACID
'IME OB EXPERIMENl
K (MEAN)
minutes
1.8 1.51 1.35 2.12 2.2 2.1 1.45 1.61 2.43 1.39
120 105 180 174 200 76 230 300 312 215
0.0754 0.0756 0.123 0.122 0.196 0.195 0.290 0.288 0.475 0.483
2.5
p.io3
10 20 30 40 55 60 IO 80 90 100 110 120 1% MINUTES
140
P
FIG.1 FIG.2 FIG. 1. VALUESOF LN K PLOTTED AGAINST 1/T I, solvent, acetic acid; 11, solvent, nitrobenzene; 111, solvent, chloroform; IV, solvent, benzene; V, solvent, water (92 per cent); VI, solvent, toluene. FIG. 2. THE REACTION VELOCITY CONSTANTS AS A FUNCTION OF TIME I, solvent, benzene at 20°C.; 11, solvent, chloroform a t 30°C.
solvents the peracetic acid is quite stable, and is not sensibly decomposed after standing a considerable'time. Table 3 shows some of the results obtained. (Initial constants are disregarded.) It will be noted from this table that the rate of the reaction in all the solvents is approximately equal, being lower than in the acetic acid solution. Curves obtained by plotting In K against 1/T are straight lines (figure 1). The activation energy calculated from these data is approxi-
851
INTERMEDIATE STAGES OF ALDEHYDE OXIDATION. I1
mately equal for this reaction in all solvents. Table 4 shows constants A and E of the Arrhenius equation, K = ACE'RT,for various solvents. It will be noted that in all non-aqueous solutions the velocity constants calculated from the data of the first portions (one to five minutes after commencement of the reaction) are always considerably higher than the TABLE 3 Kinetics of interaction between acetaldehyde and peracid in various non-aqueous solvents 1 IN
"C.
PERCENTAQE OF PERACID
ALDEHYDE PERACID
K (mole-! sec.-I cm.3)
2.59 2.48 2.5
0.0220 0.0349 0.0610
In benzene 10 15 20
2.67 2.88 2.71 In chloroform
20 30
0.0993 0.250 In nitrobenzene
20 30
0.1054 0.266 In toluene ~
20
I
2.58
I
3.1
I
0.0974
TABLE 4 T h e values of the constants of the Arrhenius equation jor the reaction in various nonaqueous solvents SOLVENT
Acetic acid ........................................ Benzene. ......................................... Chloroform. ...................................... Nitrobenzene ..................................... Wafer .............................................
I
IE
7 X 10'0 8 . 5 X 1010 6 . 6 X 1010 1 . 5 2 X 1011 4.56 X 1012
INCALORIESPER GIRAM-MOLE
15,500 16,260 16,200 16,280 16,230
following normal constants. When plotting the values of the velocity constants calculated according to the second-order equation, curves are obtained with a pronounced sharp inflection, after which the process goes on uniformly corresponding to the equation. This phenomenon suggested that the sharp inflection of the reaction
*
852
G. D. LUBARSKY AND M. 3. KAGAN
velocity curve may be attributed to the presence of an intermediate stage, the first stage being carried out a t a high rate but within the limits of the equilibrium, which is gradually shifted as the second slower stage of this reaction. This may be expressed schematically as follows:
-
AB
A+B (rapid stage)
2C (slow stage)
On this ground we might infer that if the temperature of the reaction is considerably lowered, the second stage will be slowed down in such a degree as to be almost imperceptible, while the rate of the first stage of the reaction TABLE 5 Kinetics of interaction between aldehyde and peracid i n a toluene solution 1 IN
"c.
PERCHlNTAQE OF PERACID
-41
I -
3.14 3.95 3.44
-51 -61
25
\, \
I
1.5
\
,.'\~,
0.5 ,
,
,
10 20 30 40 Y) 60 '70
,
,
,
,
,
, I
BO 90 100 110 I20 130 140
MINUTES
4,. .lo'
0.132 0.066 0.031
i
2.0
1.0
1.O 0.5
CM.8 SEC.-I) (MEAN)
(MOLE-1
3.0 2.79 3.16
25 1.5
K
ALDEHYDE PERACID
' ',
IO 20 M
,
.
90 50 60 70 Bo
I
I
90 100 XU 120 I30 140
MINUTES
FIG.3 FIQ.4 PIG. 3. INTERACTION BETWEEN PERACID AND ALDEHYDE IN TOLUENE FIG.4. KINETICS OF INTERACTION BETWEEN VALERALDEHYDE AND PERACETIC ACID
will remain slow but quite perceptible. The reaction being obviously exothermic, the equilibrium concentration of the intermediate product will increase with cooling. Experiments absolutely confirmed this conclusion.
C . Kinetics of the reaction at low temperatures The series of experiments at low temperatures was carried out in a toluene solution, the reaction vessel being immersed in a mixture of acetone and solid carbon dioxide in a Dewar flask. The temperature was held a t f0.5'C. by means of the periodical addition of small bits of solid carbon dioxide. The test portions were taken out with a micropipet and
INTERMEDIATE STAGES O F ALDEHYDE OXIDATION. I1
853
added to a solution of potassium iodide acidified with 1 N sulfuric acid. Table 5 characterizes the most typical experiments. Table 5 shows that the reaction rate at -41°C. is higher than the normal rate (corresponding to an established constant) at +20°C., i.e., a t a low temperature we have quite another process. Curves on figure 3 give a graphical illustration of these reactions. Plotting values of In K against 1 / T (figure 1) we obtain a straight line. The low activation energy, 7000 calories (notwithstanding the abnormally small value of the coefficient A of the Arrhenius equation, 5.2 X lo6),indicates the facility with which the process takes place. This is also shown by the high values of the velocity constants : the calculation of the velocity constants at 20°C. using the values at these low temperatures gives the value 3.1. Thus the presence of two stages in the interaction between peracid and aldehyde cannot be doubted.
D. The mechanism of the reaction and the method for the quantitative determination of the intermediate stage The intermediate product which is here formed might be thought of as a direct addition product of a molecule of peracid to a molecule of aldehyde according to the scheme: 0 HrC-C
// \
---to
\\ C-CH3 /
+
0
Hac-C
0-OH H
// \
OH
I
O-O-C-CH3
(a>
H forming a new peroxide compound, hydroxyethylacetyl peroxide. The peroxide intermediate product shown in scheme a slowly decomposes, forming 2 molecules of acetic acid
HsC-C
// \
0 OH
I
+
2CH3COOH
0-0-C-CH,
H In the course of the process b, the equilibrium in a is shifted to the right, and thus the reaction rate in a is determined by the reaction rate of the second stage (b). The possibility of the formation of such a peroxide intermediate compound by the interaction of peracid with aldehyde was pointed out by Wieland ( 5 ) . He was, however, not able to demonstrate by direct experi-
854
G. D. LUBARSKY AND M. J. KAGAN
ments the formation of the addition product of peracetic acid with acetaldehyde, and traced the disappearance of the peracetic acid only in the aqueous solution of aldehyde. Hence Wieland inferred that the oxidation of aldehyde by the peracid is due to dehydrogenation of aldehyde hydrate. Wieland and B. Wingler (6) obtained a compound like (a) by the action of hydrogen peroxide on acetaldehyde. Rieche (4) describes such a compound obtained by the action of ethyl-hydroperoxide on aldehyde:
0
HsC-C
//
/
OH
+ H202 + H3C-C-O.OH
\
H
OH
/ H3C-C-0 OH \ H
0
+
\\ C-CH3 /
H
-+
/
OH
H&-C-O-O-k-CH3
\
H
OH
I
H
We found the means of using the iodine evolved in the neutral potassium iodide solution (aqueous or methanol) for the determination of the whole amount of active oxygen contained in the peracetic acid and in the hydroxyethylacetyl peroxide. On the other hand, free peracetic acid may be determined in the acid solution of potassium iodide; under the action of the acid, the intermediate peroxide product is readily decomposed and the iodine evolved corresponds t o the free peracetic acid only. Thus if the reaction is carried out a t low temperatures and portions are introduced into a neutral potassium iodide solution, we are able to state that the amount of active oxygen remains practically unchanged throughout the experiment, so that it may seem that no process takes place a t all. But when the same portions are brought into an acid solution of potassium iodide a gradual decrease of the amount of iodine evolved may be observed, which corresponds to the expenditure of peracetic acid for the formation of the intermediate peroxide product according to equation a. The decomposition of the intermediate product according to equation b must be a monomolecular reaction, and this type of reaction is actually observed. The data in table 6 may serve as an illustration for the above. The experiments as well as other similar ones enable us to calculate the activation energy for the decomposition of the intermediate product. Thus, based on data of this experiment we find the latter to have a value of 18,800 calories, and from the data of another experiment 16,100. These values closely approach the values of the activation energy of the summary process of the reaction
CH3COOOH
+ CHSCHO + 2CH3COOH
INTERMEDIATE STAGES OF ALDEHYDE OXIDATION. I1
855
at 10-30°C. shown in table 4 and amounting on the average to 16,000 calories. Quite analogous data were also obtained in a chloroform solution. We were able also t o demonstrate analytically the formation of the intermediate peroxide compound by preparing a highly concentrated solution of this compound. Owing to the interaction between acetaldehyde and peracetic acid in a toluene solution at -30°C., and the subsequent cooling of the solution down t o -40°C.,white crystals precipitated from the soluTABLE 6 Formation of the intermediate product t = -40°C. PERACETIC ACID I N MOLEB TIME FROM T E E BEQINNINQ I N MINUTES
Acid KI solution
0 10 20 35
'IME
THE MINUTES
0.001946 0.001305 0.000903 0.000601
IN
I
I
Neutral KI solution
0.001941 0.001937 0.001957 0.001907
INTERMEDIATE PRODVCT I N MOLES
0 20 46 81
0.001881 0.001821 0.001748 0.001653
2.70 2.62 2.66
0 30 82 107
0.001204 0.001034 0.000795 0.000697
8.43 8.42 8.38
tion which could easily be separated from the mother liquor. These crystals readily melt at - 20°C., forming a transparent liquid with a peroxide concentration 4.5 times higher than that of the solution decanted from the crystals. Thus an addition product of acetaldehyde and peracid of a peroxide character actually seems to exist. I n this manner the proposed scheme of the mechanism of the process is fully supported kinetically and analytically. The proposed method enables us to study each stage of this process separately. THE JOURNAL OF PHYSICAL CHI~UIBTRY, VOL.
30,
NO.
6
856
G. D. LUBARSKY AND M. J. KAOAN
11. KINETICS OF INTERACTION BETWEEN PERACETIC ACID AND OTHER ALDEHYDES
It was interesting to determine whether the proposed scheme holds for reactions with other aldehydes or whether the behavior of the acetaldehyde is specific. For this purpose a series of preliminary experiments was started with isovaleraldehyde and peracetic acid. The kinetics of this reaction is quite analogous to that of the acetaldehyde reaction. The results obtained show the same characteristic features as the reactions with acetaldehyde, a sharp inflection of the reaction velocity curve near the beginning indicating the existence of an intermediate stage going on with a considerably high rate even a t -30°C. Thus we have reason t o infer that the proposed mechanism of interaction between peracetic acid and acetaldehyde is characteristic for all aliphatic aldehydes. The effect of water on the reaction rate between peracetic acid and acetaldehyde was also studied. We have found that the accelerating effect of water is connected with the acceleration of the second stage of this process, i.e., the decomposition of the intermediate peroxide. These experiments confirmed the proposed mechanism. Details of this work will be given in another paper. 111. SUMMARY
1. The kinetics of interaction between peracetic acid and acetaldehyde were studied in acetic acid, benzene, nitrobenzene, chloroform and toluene solution a t temperatures varying between 10-30°C. The activation energy varied between 15,400 and 16,300 calories in various solvents. 2. It was found that the reaction follows the second-order equation, but in the beginning there is a short period of high reaction velocity, due to the intermediate stage. 3. A mechanism of the process is proposed according to which a peroxide compound (hydroxyethylacetyl peroxide) is formed as intermediate product. 4. A method for the quantitative determination of the intermediate peroxide in presence of peracetic acid is given. 5 . Experiments were carried out with the purpose of obtaining (chiefly) the intermediate peroxide as a product of the reaction a t temperatures between '-40" and -60°C. The activation energy of this reaction amounts to 7000 calories. The coefficient A of the Arrhenius equation K = ALE'RThas the abnormally 1ow.value of 5.2 X 106. 6, A series of experiments was carried out in order t o determine the kinetics of the monomolecular decomposition of the intermediate product. 7. The mechanism proposed was extended to other aliphatic aldehydes
INTERMEDIATE STAGES OF ALDEHYDE OXIDATION. I1
857
by testing its applicability to the interaction of peracetic acid with isovaleraldehyde. REFERENCES (1) ARBUSOV: J. prakt. Chem. 131, 357 (1932). (2) CLOVERAND HOUGHTON: Am. Chem. J. 32, 43 (1904). (3) D’ANs AND FREY:Ber. 46, 1845 (1912); 2.anorg. Chem. 84, 145 (1914). (4) RIECHE,A: Ber. 64, 2328 (1931). Ann. 496, 284 (1932). (5) WIELAND: (6) WIELANDAND WINQLER: Ann. 431, 312 (1923).
