The Mechanism of the Acetaldehyde Pyrolysis

The Mechanism of the Acetaldehyde PyrolysisAuthor(s): K. J. Laidler and M. T. H. LiuSource: Proceedings of the Royal Society of London. Series A, Mathematical and PhysicalSciences, Vol. 297, No. 1450 (Mar. 7, 1967), pp. 365-375Published by: The Royal SocietyStable URL: http://www.jstor.org/stable/2415817 .

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The mechanism of the acetaldehyde pyrolysis

BY K. J. LAIDLER* AND M. T. H. Liu

Department of Chemistry, The University of Ottawa, Canada

(Communicated by F. S. Dainton, F.R.S.-Received 29 June 1966)

Previous work on the acetaldehyde pyrolysis is shown to be vitiated by the presence, in the acetaldehyde, of impurities, mainly ethanol and crotonaldehyde. The reaction has been reinvestigated with the use of acetaldehyde, prepared from paraldehyde, which is free from these and other impurities. On the basis of a study of the kinetics of formation of the major products (methane and carbon monoxide) and of a number of minor products (hydrogen, acetone, propionaldehyde, ethane and ethylene) a reaction mechanism is proposed. This includes all of the reactions in the original Rice-Herzfeld scheme, together with a number of other elementary processes, in particular

CH3 + CH3CHO -* CH4 + CH2CHO.

The decomposition of the radical CH2CHO into CH2CO and H provides an additional source of hydrogen, the rate of production of which is therefore not a measure of the rate of the initiation process. Acetone is believed to arise mainly by the reaction

CH3 + CH3CHO -* CH3COCH3 + H,

and only to a negligible extent by the combination of CH3 and CH3CO. The main chain- ending step is concluded to be

CH3 + CH3-* C2H6, with a small contribution from

CH3 + CH2CHO -* CH3CH2CHO.

The work provides further evidence for the falling off, at low pressures, of the second order coefficient for the combination of methyl radicals. Rate constants for various elementary processes are deduced from the rates of formation of the various products, and are shown to be consistent with values obtained directly.

INTRODUCTION

In spite of a considerable amount of work on the thermal decomposition of acetaldehyde, many details of the mechanism have remained obscure and confusing. Some of the results are consistent with the original mechanism of Rice & Herzfeld

(I934): C H3CH0 CH3 +0CH0,

CHO CO+1H,

H + C13H 0 - 1CH3C0+ H2,

CH3+CH3CHO -0 CH4 + CH3C0,

CH300 CH3 + C0,

CH3 +0CH3 C-26

This scheme successfully accounts for the following features of the reaction: (1) The main reaction products are methane and carbon monoxide; these are

produced in the chain-propagating steps. * Now at the Chemical Laboratory, University of Sussex, Falmer, Brighton (1966-67).

[ 365 ]



366 K. J. Laidler and M. T. H. Liu

(2) The order of the reaction is three-halves; this is explained if the order of the termination reaction is one greater than the order of the initiation reaction (e.g. if initiation is first order and termination is second order).

There are, however, a number of experimental results that cannot be reconciled with this mechanism, such as the following:

(1) According to the mechanism the only minor products should be hydrogen and ethane. In addition, however, significant amounts of acetone, propionaldehyde and ethylene are formed (Trenwith I963; Dexter & Trenwith I964). Indeed Dexter & Trenwith (I964) found that the initial rate of acetone production is much greater than that of ethane production, from which they concluded that the methyl radical combination cannot be the main termination step.

(2) Eusuf & Laidler (I964), in agreement with the results of Bril, Goldfinger, Letort, Mattys & Niclause (I950), found that the overall rate of reaction is reduced by the addition of inert gases, and from this they concluded that the methyl radical combination must be in its pressure-dependent region. The recent work of Quinn (I963) and of Lin & Back (I966) has, however, indicated that the methyl radical combination should be in its second order region under most of the conditions used in the acetaldehyde pyrolysis. There is also some confusion on the experimental side about inert gas effects; Niclause (I954) later reported that there was no effect on the overall rate, and Dexter & Trenwith (I964) found an effect on the rates of both the initiation and the termination reactions, but no effect on the initial overall rate.

(3) Trenwith (I963) found that the hydrogen production was second order in acetaldehyde. If the Rice-Herzfeld scheme is accepted this can only mean that initiation is second order in acetaldehyde; this, however, leads to the incorrect overall order if the termination process is taken to be second order.

(4) Eusuf & Laidler (I964) found the rate of ethane production, which on the basis of the mechanism must be equal to the rate of the initiation process, to be second order in acetaldehyde and to be increased by the addition of inert gas. Again, this result is inconsistent with the overall order if the termination reaction is second order.

It is evident that several of the reported experimental results are inconsistent with the Rice-Herzfeld mechanism, and that some of these (particularly those relating to inert gas effects) are mutually inconsistent. The present work was undertaken with the object of resolving these difficulties. It was suspected that some of the experimental anomalies, especially those relating to the rates of production of minor products, were due to traces of impurities in the acetaldehyde employed, and great care was therefore taken to test, by vapour phase chromatography, nuclear magnetic resonance spectroscopy and mass spectrometry, for such impurities and to work with samples of acetaldehyde that were substantially purer than had previously been employed. When this was done the kinetic results were somewhat different from those previously obtained. They could be reconciled with a scheme of reactions based on the original Rice-Herzfeld mechanism but including a number of important additional reactions.



The mechanism of the acetaldehyde pyrolysis 367

EXPERIMENTAL

Materials Acetaldehyde obtained from Eastman Organic Chemicals and from the Matheson

Company, and purified by a number of bulb-to-bulb distillations, showed identical kinetic behaviour, even as far as the minor products were concerned. However, the fact that there was a high initial rate of ethylene production, as found by Tren- with (I963), led to the suspicion that ethylene was being formed from an impurity. This was confirmed by pyrolysing partially decomposed acetaldehyde, which produced ethylene at a considerably lower rate. Vapour phase chromatography and nuclear magnetic resonance studies indicated that the main impurities were ethanol and crotonaldehyde. These impurities are not removed from the acetaldehyde by distillation.

It was found that impurities were not present in acetaldehyde prepared from paraldehyde, and this was the material used in the remainder of the investigations. East- man Research Grade paraldehyde was distilled, and the fraction of b.p. 1240C was collected. This was distilled under nitrogen with addition of a small quantity of concentrated sulphuric acid, and acetaldehyde boiling at 21 ?C was collected and purified by trap-to-trap distillation and degassed at - 160 TC to remove traces of oxygen. It was stored at -78 ?C and redistilled each time before use. No trace of any impurity in this material was revealed by vapour phase chromatography, nuclear magnetic resonance spectroscopy or mass spectrometry.

Apparatus

The decomposition was carried out in a conventional static system. The reaction vessel was a quartz sphere 515 ml. in volume, with a S/ V ratio of 0 6 cm-1, enclosed in a furnace consisting of an electrically heated steel cylinder 15 cm thick. The temperature of the vessel was controlled to within 0-2 degC by means of a thermo- electric thermoregulator. The temperature was measured using a chromel-alumel thermocouple.

Procedure

Before each experiment the system was evacuated to a pressure of 10-5mmHg. At the end of each run the products and reactant were fractionated through traps at 130, - 150 and - 210 'C. Acetone was measured by transferring all the conden- sable products to a sampling U tube, followed by gas-chromotographic analysis on a 5 m column of 20 % tetraethylene glycol dimethyl ether on Chromosorb P (hexamethyldisilazane treated) at 30 'C. The area of the acetone peak was measured. The column was calibrated by using a measured amount of acetone mixed with an appropriate quantity of the reactant used in the pyrolyses. The non-condensables, hydrogen, carbon monoxide and methane, were measured in a gas burette and analysed on a 4m 100mesh silica-gel column maintained at 0?C. Analysis on a 1 m silica-gel column at 30 'C showed that the - 210 0C fraction contained ethane, carbon dioxide, ethylene and the residual methane. The volumes of these gases were also measured by gas burette. Helium was used as carrier gas in all cases.




RESULTS AND DISCUSSION

Rates of production of the major product methane and of the minor products hydrogen, acetone, ethane, ethylene, propionaldehyde and carbon dioxide were measured at 523 0C at pressures from 11 to 558 mmHg. Some results are shown in figure 1, and initial rates are given in table 1. These rates are the averages from

06 -

O |3 COCH3

0

04 - ?

0~~~~~~~ 0

0-2 0

CA

CH3CH2CHO 7x

0 200 400 600

time (see)

FiGIJRE 1. Products of pyrolysis at 523 'C; acetaldehyde pressure =117 mmllg.

TABLE 1. INITIAL RATES OF PROD-UCT FORMATION

(T = 5 23 'C; all rates are in moles ml. - s -1) pressure of

acetyldehyde (mmllg) 1010 vC4 1013 VH 2 1013 V0113000113 1013 V C2116 1013 VC2H15CHO0

11I 1.10 /022 20 2*28 5.6 -0-54 30.5 3.50 0 _ 095 40 6-20 13*6 -1.65 55 10-0 224 64 5 2631 77 17.5 32.5 15*6 3-8 cca. I- 5

117 321 6 61P2 27? 2 5*9 oct. 1 5 150 48.0 82*5 37.5 7.3 oct. 2 195 71T 0 129 57.5 9R5 Oc. 3 292 129.0 230 -157 558 3600 555 - 307

3 to 8 experiments, and were obtained not from initial slopes but by extrapolating rates to zero time. It is to be seen from figure 1 that for the hydrogen production there is no sign of the induction period found by Trenwith (1963). There is also no



The mechanism of the acetaldehyde pyrolysi,s 369

initial rapid formation of ethylene, which was found by Trenwith (i 963) and by us when we used impure acetaldehyde.

The order of methane production was 3 at pressures above 100 mmHg, and some falling off of the rate constant was noted at lower pressures. The -a order rate constant was 0 91 ml. mole' s-', to be compared with 0 99 obtained by Eusuf & Laidler (I964) and 0.59 ml.i mole- s-1 obtained by Letort (I937).

To explain these results we suggest the following mechanism, which is the Rice-Herzfeld scheme to which a number of reactions have been added:

[1] CH3CHO CH3+CHO,

[2] CHO CO + H,

[3] H + CH3CHO H2 + CH3CO,

[4] CH3 + CH3CHO CH4 + CH3CO,

[5] CH3CO CH3 + CO,

[6] CH3 +CH3CHO CH4 + CH2CHO, [7] CH2CHO CH2CO + H,

[8] CH3 + CH3CHO CH3COCH3 + H,

[9] CH3 + CH3 C2H6,

[10] CH3+CH2CHO CH3CH2CHO.

Reaction [1], the initiation reaction, is concluded to be in its first order region. The rate of ethane formation is significantly greater than for propionaldehyde formation, so that the main termination step is [9]; a variety of evidence (Quinn I963;

Lin & Back I966) has shown that this reaction is second order under the usual pyrolysis conditions.

First order initiation and simple fl,8 termination (Goldfinger, Letort & Niclause 1948) lead to 3 order kinetics, in agreement with the results of many workers for this reaction; in the present investigation it was confirmed that the methane and carbon monoxide production is strictly 3 order in acetaldehyde. If the minor termination step [10] is ignored the rate of disappearance of acetaldehyde, approxi- mately equal to the rates of formation of carbon monoxide and methane, is found to be VCu3CHO = Vco = VI = 14(kl/kg)1 [CH3CHO]3. (1)

From the overall rate constant of 0-91 ml.1 mole s-1 it is possible to calculate

kc, by using the value of k4 (see below) and the kg value of Shepp (I956); the result is

ki = 2-5 x 10-7 s-1.

Ethane production

The logarithm of the rate of ethane formation is plotted against the logarithm of the acetaldehyde pressure in figure 2. The order is unity at the higher pressures, but there is a significant falling off from linearity at pressures below 80 mmHlg. If reaction [4] is the only source of methane and reaction [9] the only source of ethane l (

ks = Ve 2H[CI3CHO]. (2) 24 lCI24

24 Vol. 297. ~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~A.




X / 30

bID 0 - _

-1 42

IIII I I I 1 2 3 10 15 20 25

log10 (P mmHg) log , (P mmHg) FIGURE 2 FIGURE 3

FIGURE 2. Plot of log10 VC2R6 against log10 PC113CRO (mmHg); T = 523 ?C. The dotted line corresponds to log10 v1, with k1 equal to 3-0 x 10-7 s-1.

FIGURE 3. Plot of log1o (kQ/k4) against log1o P (mm Hg); T = 523 ?C.

Figure 3 shows a plot of the logarithm of the right hand side of this equation against the logarithm of the acetaldehyde pressure; again the falling off of kg at low pressures is observed. In the ethane pyrolysis at 550 ?C Lin & Back (i 966) have observed a falloff of the rate coefficient for the methyl radical combination below about 200 mmHg pressure, while in the mercury-photosensitized decomposition of dimethyl ether at 200 to 300 ?C Loucks & Laidler (I966) have observed a falloff at 100 mmHg pressure. It appears that for this reaction acetaldehyde is a somewhat more efficient third body than ethane or dimethyl ether.

The high pressure value of k4/k , obtained from (2), is consistent with values of k4 and kg obtained directly. Thus use of the 1c9 value of Shepp (I956) together with our results gives rise to 8-6 x 109 moles ml.-1 s-1 for k4 at 523 ?C, in excellent agreement with the value of 8-7 x 109 calculated at this temperature from the data of Brinton & Volman (I 952). The proposed mechanism is therefore entirely consistent with the observed relative rates of formation of methane and ethane.

The steady state equation for the total radical concentration is

V=V + vvo (3)

whence vi = VC2H6 + vC2H5CHO- (4)

The half-shaded circles in figure 2 denote the values of Vc216 to which have been added the rates of propionaldehyde production. The sum of these rates corresponds to a k1 value of 3 0 x 10- s-1, which agrees very well with the value of 2-5 x 1O-7 S-1

deduced above from the overall rate.

Acetone production Trenwith (I963) found that the rate of acetone production was greater than the

rate of ethane production, and the present work confirms this. Trenwith assumed




that acetone is formed by the reaction

[11] CH3CO + CH3 ->CH3COCH3,

which he therefore concluded to be the main termination step. Eusuf & Laidler (I964) pointed out that the overall kinetic behaviour cannot be explained if this is the main termination step, and suggested that acetone is formed by reaction [8], a reaction which may involve an intermediate of structure

H

H3C-C-O

CH3

There are other objections to reaction [11] as an important termination step. If this reaction occurred, acetyl radicals would dimerize with the formation of biacetyl, and a rough estimate shows that the rate of formation of biacetyl should be comparable to that of acetone. Such amounts of biacetyl would be easily detectable using the techniques of the present investigation, but none was found. It is also possible to estimate the relative concentrations of CH3CO and CH3 in the reaction system. The steady state condition leads to

[CH3CO] _ k4[CH3CHO] (5) [CH3] k5

where k5 is a constant (kcr) at high pressures and becomes proportional to pressure (kO[CH3CHO]) at low pressures. O'Neal & Benson (I962) have obtained values of ko and k?, and the use of these values, together with our value of k4 or those of Kerr & Calvert (I965) and Brinton & Volman (I952), leads to the conclusion that [CH3CO]/[CH3] is 10-2 to 10-3 under the conditions of the present experiments. The methyl-methyl combination will therefore be by far the most important termination process.

This conclusion is supported by a consideration of the rates of formation of methane found in the present experiments. If acetone were formed mainly by reaction [11] it can be shown that

VCl3coH3 kk cll 4cH (6) k3k~ 0114 C21a6~ 6 5 9

The use of O'Neal & Benson's value for k5 and of that of Shepp (I956) for kg, with k11 assumed equal to kg, leads to a rate of acetone production which is two to three orders of magnitude lower than that observed.

The alternative proposal that acetone is formed by reaction [8] seems to be consistent with all of the facts. Thus the value of k8 can be calculated from the data in two completely different ways, with excellent agreement:

(1) From the rates of formation of acetone and ethane, and using kg from Shepp

(I956): VC03COCH3 -8 [CH300I-3](

vC2H6 k9

whence k8 = 7-1 x 106 ml. mole-1 s-1.

24-2




(2) From the rates of formation of acetone and methane, and using our value of k4, VCH3COCII3 _ k8

f-I TT 4k ~~~~~~~~(8)

whence k8 = 70 x 106 ml. mole-1 s-1.

Hydrogen production Trenwith (i 963) reported that hydrogen production showed an induction period,

and that the process was second order in acetaldehyde. In the present study no induction period was found, and the order of hydrogen production was 1P4, as shown by the double logarithmic plot in figure 4. If the simple Rice-Herzfeld mechanism were valid the hydrogen production would be a measure of the rate of the initiation process, and this was Trenwith's conclusion. In terms of the mechanism proposed, however, hydrogen can result from reactions [3], [7] and [8], and the rate of its formation is not simply related to the rate of initiation.

2 -

bH 4 -

+

1 2 3 0 10 20 30

log1o (P mmHg) 104 [CH3CHO]l (molei ml.-i)

FIGURE 4. Plot of log1o vH2 against FIGURE 5. Plot of vH2/[CHI3CHO] against log10 P (mmHg); T = 523 ?C. rCH3CHOI! at 523 OC.

Application of the steady state treatment leads to the following expression for the rate of hydrogen production:

=H2 kj[CH3CHO] + (k6 + k8) (kcl/kg) [CH3CHIO]& (9)

According to this, a plot of vH1/[CIH3CHO] against [CH3CHlO] should be linear; figure 5 shows that this is the case, and from the intercept on the ordinate it is concluded that k1 at 523?C is 5 0 x 10-7s-1. This is somewhat higher than the value of 2-5 x 10-7 S-1 obtained from the overall rate; it is difficult at the present stage to say which value is the more reliable.

The slope of the plot in figure 5, together with the values of k1 and kg, allows k6 + k8 to be obtained: the value is 8 8 x 106 ml. mole-1 s-1. From the rate of acetone production we deduced that k8 is 7 0 x 106 ml. mole-1 s-1, so that it follows that k6 is




1.8 x 106 ml. mole-' s-1. It is to be noted that k6/k4 is 2 1 x 10-4. This estimate of k6 is considered to be less reliable than the one made in a later section, but the dis- crepancy is not great.

Propionaldehyde production The fact that propionaldehyde is formed as an initial product suggests the pre-

sence of the radical CH2CHO, which is presumably formed by reaction [6]. The participation of this radical was previously suggested by Wall & Moore (I95I) to account for the mixed products formed in the pyrolysis of a mixture of CH3CHO and CD3CD0.

Because propionaldehyde decomposes rapidly at the temperatures of these experiments (Eusuf & Laidler I965) it is difficult to make a reliable estimate of its initial rate of formation; the rate is similar to, but smaller than, the rate of ethane production (cf. table 1). At a pressure of 117 mmHg, the rate of propionaldehyde formation is estimated to be 1-5 x 10-13 mole ml.- s-1, as compared with a rate of 59 x 10-13 mole ml.-1 s-1 for the rate of ethane production.

Besides undergoing reactions [7] and [10], the CH2CHO radical might undergo the isomerization reaction

[12] CH2CHO + CH3CHO 1 CH3CHO + CH3CO.

Since, however, the CH2CHO radical will be considerably stabilized by resonance this reaction will not be very important; if its activation energy is greater than 10kcal/mole it will be much less important than [7] under the conditions of the present work.

Ketene production Ketene decomposes rapidly under the conditions of the present experimelnts,

so that it was not possible to make a direct measurement of its rate of formation. This rate can, however, be estimated on the basis of a material-balance equation

V12 = VCH3COCH3 + VC2HG + VC2H,CHO + VCH2CO, (10)

which is readily derived from the steady state equations. At an initial acetaldehyde pressure of 117 mmHg the rate of ketene formation is found to be, from the rates in table 1, VCH2CO= 266 x 10-13 mole m1-1 s-1.

The steady state equation for the radical CH2CHO is

V6 V7+V10 = VCH,CO+VC2H5CHO, (11)

so that, at 117 mmHg, v 28.1 x 10-13mole ml.-1 s-1.

The ratio of the rates of reactions [6] and [4] is therefore

v6_28*1 xl 101 _ ________ = 8&6 x 10 w'

V4 32*6 x 10-10

This is the ratio of the rate constants and it follows that k6 = 7-5 x 106. If the fre- quency factors of reactions [4] and [6] are the same the activation energy of [6] is 11 2keal/mole higher than that of [4], or is 18-7kcal/mole. This difference in




activation energies is of interest in view of the fact that D(H-CH2CHO) is probably about 102 keal/mole, which is 15 kcal/mole greater than D(CH3C0-H) (Benson I965).

The relative concentrations of [CH2CHO] and [CH3] can be estimated from the relations VC2H5CHO k1o[CH31] [CH2CH0] (12)

and VC2H6

= kg[CH3]2. (13)

It follows that [0H20H0] k9VCHCH(

[CH3CHO k10VC:2H6 (14)

If k9 and k1o are equal [CH2CH0] = 15 x 1013

[0113] -5*9 x 10 0-025. (15)

It was noted above that [CH3C01/[CH31 is 10-2 to 10-3. Although formed much less rapidly CH2CHO is present in much higher concentrations than CH3C0, because of its lower reactivity.

Other minor products The earlier work (Trenwith I963) appeared to indicate that ethylene was a

primary product of the reaction. The present results show a much lower rate of ethylene production with no detectable induction period. This result is consistent with the conclusion that the ketene formed in reaction [7] breaks down very rapidly at these temperatures (Guenther & Walters I959).

The rate of carbon dioxide production is initially low or zero, but increases and finally decreases. The results are consistent with the hypothesis that the carbon dioxide is formed from ketene, which according to Guenther & Walters (I959) rapidly decomposes into two sets of products:

2CH2C00 0C3H4 + C02,

2CH2C0 = C2H4 + 2C0.

The rate of the first process is, at this temperature, about 5 times that of the second. No other process in the present system seems likely to produce carbon dioxide.

CONCLIUSIONS

The scheme of reactions proposed for the acetaldehyde pyrolysis appears to give a satisfactory explanation of the results, in that the rate constants deduced on the basis of the rates of formation of products are self-consistent and in good agreement with directly observed values when these are available. Table 2 summarizes the rate constants determined from the present and from other investigations.




TABLE 2. SUMMARY OF RATE CONSTANTS FOR ELEMENTARY REACTIONS,

OBTAINED IN THIS AND OTHER WORK

k (523 ?C) (s-1 or ml. mole-' s-1)

present other reaction work values reference

[1] CH3CHO -- CH3 +CHO 2-5 x 10-7 1P5 x 10-7 Benson (I960) [4] CH3 + CH3CHO -+ CH4 + CH3CO 8 6 x 109 8 7 x 109 Brinton & Volman (1952)

4*4 x 109 Kerr & Calvert (I965)

[5] CH3CO CH3 +CO 1- 6 x 106 (ko) O'Neal & Benson (i 962) CH3CO-+CH3+CO - ~~~~~1 6x IO" (k0)J

[6] CH3 + CH3CHO CH4 + CH2CHO 7*5 x 106 [8] CH3 + CH3CHO CH3COCH3 + H 7*0 x 106 -

[9] CH3 + CH3 -- C2H6 2-2 x 1013 Shepp (1956)

The authors are indebted to the National Research Council of Canada and the Petroleum Research Fund (administered by the American Chemical Society) for grants in support of this work. They thank Dr Margaret H. Back for valuable discussions and assistance, Dr R. R. Fraser for nuclear magnetic resonance analyses and Dr J. L. Holmes for mass spectrometric analyses. Thanks are also due to Dr A. B. Trenwith for valuable comments.

REFERENCES

Benson, S. W. I960 The foundations of chemical kinetics, p. 379. New York: McGraw-Hill. Benson, S.W. I965 J. Chem. Ed. 42, 502. Bril, K., Goldfinger, P., Letort, M., Mattys, H. & Niclause, M. 1950 Bull. Soc. chim. Belg. 59,

263. Brinton, R. K. & Volman, D. H. I952 J. Chem. Phys. 20, 1053. Dexter, R. W. & Trenwith, A. B. I964 J. Chem. Soc., p. 5459. Eusuf, M. & Laidler, K. J. I964 Can. J. Chem. 42, 1851. Eusuf, M. & Laidler, K. J. I965 Can. J. Chem. 43, 268. Goldfinger, P., Letort, M. & Niclause, M. 1948 Volume Commemoratif Victor Henri: Con-

tribution a 1'etude de la structure moleculaire, p. 283. Liege: Desoer. Guenther, W. B. & Walters, W. D. 1959 J. Am. Chem. Soc. 81, 1310. Kerr, J. A. & Calvert, J. G. I965 J. Phys. Chem. 69, 1022. Letort, M. 1937 J. Chim. Phys. 34, 265, 355, 428. Lin, M. C. & Back, M. H. I966 Can. J. Chem. 44, 505, 2357. Loucks, L. F. & Laidler, K. J. I966 Canad. J. Chem. (To be published). Niclause, M. 1954 Rev. Inst. Fr. Petrole 9, 327, 419. O'Neal, E. & Benson, S.W. I962 J. Chem. Phys. 36, 2196. Quinn, C. P. I963 Proc. Roy. Soc. A 275, 190. Rice, F. 0. & Herzfeld, K. F. 1934 J. Am. Chem. Soc. 56, 284. Shepp, A. 1956 J. Chem. Phys. 24, 939. Trenwith, A. B. I963 J. Chem. Soc., p. 4426. Wall, L. A. & Moore, W. J. I951 J. Phys. Colloid Chem. 55, 965.



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The Mechanism of the Acetaldehyde Pyrolysis

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