THE PYROLYSIS OF ORGANIC ESTERS

DISSERTATION

Presented in Partial Fulfillment of the Requirements for the Degree Doctor of Philosophy in the Graduate School of The Ohio State University

RICHARD JUI-FU LEE, B.S

y l f

.■) ®rV'«\ a.i @@ :&f a■ '

Adviser * Department of Chemistry -i-

ACKMOWLEDGMEMT

My sincere and profound thanks to Dr. Christo­

pher L. Wilson for the suggestion of the problem

and his inspiring guidance throughout the research.

My gratitude to the B. F. Goodrich Company, whose

Research Fellowship I held for the entire span of

this work. Finally, but not least, my deepest appreciation to Mr. Daniel Loughran for his ever cooperative assistance in making the infra-red analyses. -ii-

TABLE OF CONTENTS

Page

I. INTRODUCTION AND STATEMENT OF PROBLEM 1

II. HISTORICAL INTRODUCTION...... 5

1. Pyrolysis of Organic Compounds...... 5

A. Berthelot's Theory...... 5

B . Haber' s Rule ...... 6

C. Rule of Least Molecular Deformation...... 7

D. Bredt's Rule...... 8

E. Blanc’s Rule...... 8

F. Nef's Theory...... 9

2. Theory of Pyrolysis of Esters...... 10

A. Bilger and Hibbert's Mechanism.... 11

B. Hurd's "Hydrogen Bridge" Cyclic Mechanism...... 13

C. Ritchie's Dual Reactions...... 16

D. Htickel's Concept (Tsugaev Reaction) 19

E. Alexander's Cis-Elimination Studies 20

F. Kinetics on the Pyrolysis of Esters 22

G. Some Contributions from the Kinetic Study of Dehydrochlorination of Chloroparaffins ...... 24

3. The Cyclic Transition State in Other Processes...... 32

A. General Considerations...... 32

B. The Concept of O'Connor and Nace... 34

C. The Pyrolysis of Sulfites...... 37 iii

£a^e

III. Discussion of Earlier Work,...... 42

1. Kinetics as a Method for Mechanism Study.. 42

2. Substitution Effects on the Cyclic Transition Complex...... 48

3. The Excited Triplet State as a Transi­ tion State in a Special Case ...... 63

IV. EXPERIMENTAL...... 88

1. Pyrolysis of Diacetyl Cyanide (D.A.C.).... 88

A. Reagents...... 88

B. Apparatus * 89

C. Decomposition of D.A.C...... 92

D. Isolation and Estimation of Products.. 94

E. Results...... 96

F. By-Products...... 101

G. Preparation and Pyrolysis of Pyruvic Nitrile ...... 102

2. Pyrolysis of Di-esters...... 104

A. Preparation of Materials...... 104

B. Procedure for Pyrolysis ...... 105

C. Collection and Identification of Products ...... 106

3. Pyrolysis of Cellosolve Acetates...... 108

A. Preparation of Materials...... 108

B. Pyrolysis Procedure...... 108

C. Collection and Identification of Products ...... 109

V,. INFRA-RED SPECTROGRAMS...... 110 -iv-

Page

VI. S UMMARY...... 128

VII. BIBLIOGRAPHY...... 129

VIII. AUTOBIOGRAPHY...... 135 INTRODUCTION AND STATEMENT OF PROBLEM

Although much is known concerning the pyrolysis of esters, most of the evidence has not been correlated.

The known facts about thermal decomposition of simple esters were recorded almost as far back as a century ago by Oppenheim and Precht (l). They found that when esters containing beta-hydrogen in the alcohol portion of the molecule were decomposed thermally, acids and olefins were generally the products of pyrolysis. Their work was principally confined to the acetates. Modifications in the acid and alcohol portions of the ester molecule were investigated in succeeding researches such as those of Engler and Low (2),

Peytral (3), Bilger and Hibbert (4 ) and Hurd and

Blunck (5). These workers changed the structure of the ester by varying the number of beta hydrogens available and the nature of the acidsj such as acetates to benzoates, phenyl-acetates, propionates, formates, etc. From these studies they interpreted the behavior of esters on pyrolysis.

In the above cases the temperature entered as a factor to be considered in that it determined the nature of products to some considerable extent. These investigators did not significantly examine the -2- temperature dependency of the primary or succeeding processes of the reaction in the cases studied, having in mind the mechanism. The collection and detection of all products formed were not too explicitly and completely reported since the secondary decompositions accompanying a rise in temperature complicated, in most cases, the results obtained.

Perhaps it was due to this factor that most investigators confined themselves to simple esters, or perhaps it was the simplicity of operations that led the workers to concentrate on the beta-hydrogen effect. Nevertheless the succeeding studies by Ettckel,

Tappe, and Lengutke (6), Stevens and Richmond (7),

Houtman, van Steenis and Heertijis (8), Wibaut and van

Pelt (9), Ritchie, Jones and Burns (10), and many

others, till the very recent studies of Bailey (ll) have been concentrated largely on the selectivity of

the beta-hydrogen elimination phenomenon.

Chronologically there has been no cessation of

effort in the field of ester decomposition. But there were more prominent fads through certain periods of

history than others, and these peaks gave welcomed

respite for the researcher to stop and take inventory.

As with all other well practiced reactions, pyrolysis

of esters enjoyed and is enjoying today a very practi­ cal and industrialized position. The patents of Filachione, Fisher, Ratchford, Rehberg, Fein and

Smith (12) and many others bear out that this unique elimination has been much favored for synthesis.

It is not clear, however, why theoretical con­ siderations of ester pyrolysis have not kept pace with this practical application. Recently there has been some quantitative work, but very little of it concerns carboxylic esters.

Some studies have been made of surface catalysis (8) rate measurements (13), homogeneity of the system and activation energies required (14), however some skepti­ cism remains about applying these data to the decompositi

of esters in general. Hence prediction of products from the molecular structure of the ester or from a knoifledge of the mechanism of decomposition is still whimsical.

We have tried in the present work to broaden

the scope of the now fairly acceptable postulate of

intra-molecular decomposition of esters by testing such

a mechanism on species that do not undergo normal elimination, but decompose rather with some degree of rearrangement. The simple but different products from

these reactions, taken together with what the litera­

ture has revealed concerning normal eliminations, should enlighten us about the transition state. It -4- should prove interesting to observe whether any of the postulated transition states will fulfill the needs of both normal and abnormal eliminations. -5-

HISTORICAL INTRODUCTION

Pyrolysis of Organic Compounds

A . Berthelot*s Theory

If one wishes to decide what was the beginning of

the theory of pyrogenic reactions, Berthelot (15)

may well be considered as the pioneer. His general

theory may be summarized in three parts (l6 )~:

(1 ) In addition to pyrolytic reactions of decom­

position, there are also reactions of synthesis. In

the latter, there is progressive hydrogen elimination,

accompanied by the gradual formation of complex hydro­

carbons which eventually may result in the deposition I of carbon.

(2) The building-up processes and the tearing-

down processes are considered to limit or oppose each

other since the lower ones reform to produce the higher

one. This leads to a complicated equilibrium between

/ an increasing number of hydrocarbons. Berthelot*s

conception of the decomposition of methane is as follows:

2CH4 - ♦ CH2 =CH2 + 2H2

2CH4 - * chhch; + 3H2

ch2=ch2 ♦ GH=CH + H2

HChCH + h2 —* ch2 =ch2 -6-

GH2=CH2 + H2 ------f CH3-CH3

2 CH3-CH3 ______, 2 CH^ + HC=CH + H2

(3 ) These reactions occur whether the hydrocarbon is in contact with hydrogen or with other hydrocarbons.

B . Haber's Rule

Considerable criticism of Berthelot's theory was made by Haber (17). He pointed out that Berthelot's first proposal was an arbitrary interpretation of the observation that in the gasification of hydrocarbons at progressively high temperatures, graphite always appeared; however the carbon was never free from hydro­ gen. Furthermore, am: equilibrium postulated by Berthe­ lot indicated a permanent slate eventually reached because of a permanent set of external conditions. Since the effect of temperature on the equilibrium was obscure, it is hard to accept Berthelot's "methane equilibria".

Furthermore, the irreversibility of the formation of benzene from acetylene makes the explanation awkward.

To supplant Berthelot's ideas Haber gave a general rule of his own. He concluded that the C-C linkage in the aromatic series is more stable than the C-H linkage and the reverse is true in the aliphatic series.

This rule, however, failed to explain many embarrassing data such as the hydrogen production in preparing propylene from propane, the pyrolysis of diphenyl ethane, -7- and the pyrogenic reaction of ethane to yield hydrogen.

Many more cases that the Haber Rule cannot predict make the usefulness of the concept limited.

G. Rule of Least Molecular Deformation.

The trend toward the importance of molecular structure initiated by the previous considerations led to the concept of "Least Molecular Deformation". Al­ though this general rule could have been demonstrated by the earlier works on pyrolysis it was not formulated until the early part of this century. It was Mile.

Peytral (18), whose quantitative studies on the pyrolysis of esters were indeed before their time, who introduced this rule. She observed that the decompo­ sition by heat would follow that reaction which required the least possible deformation of the molecule. Thus in a series of reactions the atomic bonds in the products tend to be as nearly identical as possible with those in the original compound.

CH2-CH2 I I (I) flf HH;HG1 -----* HC=CH; CH3CN; NH3HCI (19) C ' CH3

For example, in the decomposition of 2-methyl-imidaz0- line (I), it is possible to predict the nature of the products. However the inconclusiveness of the quan­ titative aspect leaves much to be desired. - 8-

D. Bredt's Rule (20)

A more limited rule is that of Bredt. He observed the formation of isodehydrocamphoric anhydride instead of dehydrocamphoric anhydride (II),

o. c6 co •

'H (III) H Bi

rf and many other examples of ring closure in ketonic acids (21). He explained these unexpected results by stating that it appeared impossible to realize a re­ action which would lead to a double bond formation at a bridge head position. This rule is still very useful and has theoretical basis.

E . Blanch Rule (22)

Like Bredt, Blanc also observed irregularities in the formation of cyclic ketones from the pyrolysis of acid anhydrides. For example carbon dioxide was not produced when hexahydrohomophthalic anhydride (IV) was heated for 12 hours at 220°C. but only an equilibrium mixture of cis and trans anhydrides was obtained. 9-

(iv) v V v «*' 'TN. „

One of the results of a combination of Bredt1s and Blanc’s

Rules can be well illustrated by the example of thermal

decomposition of barium and calcium salts of tereph-

thalic acid and hexahydrophthalic acid (24). Only in

the second case was there a production of small amount

of ketone, bicyclo-1 ,2 ,2-heptanone.

F. Hef’s Theory

These attempts t o .understand the pyrogenic reactions

were capped by the final proposal of Kef (25). Instead

of an empirical rule Kef proposed a theory of dissocia­

tion where the original compound was always considered

to dissociate into an ephemeral molecule of bivalent

carbon which might then react with itself or other

neighboring substances. At the risk of doing injustice

to Kef’s theory we will, nevertheless, give only the

example of the pyrolysis of acetaldehyde and acetone

to illustrate the concept.

(i) CH3-CHO > (CH2 = ) + (CH20) -- *

(CH2 = ) + 2 (H) + QO

(CH2 = ) + 2(H) » CH4 .

(ii) CH3COCH3 ---). (CH2 = ) + (CH3CHO)

CH^CHO ----* CO + CH. by route (1) 4 -10-

With our actual knowledge of the products of the pri­ mary dissociation of acetone, i.e.

CH3CQCH3 = GH2=C=0 + CH4

Nef's proposal becomes only of historical interest, unless it is believed that

ch2= + CO ch2 =c=o

However, much of.the origin of our theory of free radical mechanism may be traced back to Nef's intriguing

concept of dissociation.

2* Theory of Pyrolysis of Esters

In order to discuss some of the modern ideas on

pyrolysis, the scope of the reaction will now be con­

fined to esters. This will enable'us to relate a more

precise and correlated picture to attempts to predict

products .from structure, energy relationships, and rates

of reactions.

The expansion of Nef's theories came quite a few

years later in the form of free radical concepts (26).

This development considered a primary dissociation to

radicals occurring when compounds were pyrolyzed.

Accompanying these concepts was another for the mechan­

ism of decomposition of esters by heat. This envisaged

an intermediate of an unstable intracyclic nature,

isomeric with the ester, and which was believed to be

the primary product (27). -11-

A . Bilger and HibberVs Mechanism

These two rival theories were explored by Bilger and Hibbert (4 ). Four possibilities of intraoyelic

intermediates were considered, namely, oxonium (V),

chelate (VI), hemiacetal (5, 4 or 3-membered) (VII)

and acetal (6 ,5,4-memhered) rings (VIII) (28). The

case of n-propyl acetate was used to illustrate the

reaction.

cw-co J chi. o/Hv|cH-tH3 0 \ ' II I1

(V) (VI)

o © 4 c.i4 -c-Ha i ( i C.H, til °l

oh oB OH I (Vila) (Vllb) (VIIc)

/O-fcth. ^ C % C.U- I C4f^ ^"CrH I I iCHcvyLttj X 0 -}«•< 0 "" (villa) (V U I b ) (VIIIc)

Of the four forms, only (VIII) was given serious con­

sideration by Bilger and Hibbert but they showed that

even this acetal formation was not allowable as an inter­

mediate (except perhaps unknown (VIIIc)) in the pyrolysis

of esters since cyclic acetals are too stable at temperatures of ester pyrolysis. The difference between the products formed by the decomposition of n-butyl acetate and ethyl n-butyrate also gave conclusive evi­ dence against the formation of cyclic acetals as inter­ mediate. In addition Bilger and Hibbert showed that ether-intermediates were not consistent with their experimental evidence and furthermore that a reverse

Tischtschenko Reaction (29) of ester to aldehyde was not acceptable. They did favor, however, the free radical explanation mostly on the strength of work done with benzyl benzoate. They also pointed out that ester decomposition into acid and olefin through pyrolysis was

"normal”* when hydrogen was present on the beta-carbon atom of the alkoxy-radical.

Hurd and Bennett (30) substantiated Hibbert's studies on the pyrolysis of benzyl benzoate by a study of the thermal decomposition of benzaldehyde. In the pyrolysis of 42.5 g. benzyl benzoate at 340-350°G. they found unchanged ester (21.5 g.), toluene (3 g.), benzaldehyde (2 g.), benzoic acid (2 g.), benzoic an­ hydride (3 g.) and tar (10 g.), with no signs of dibenzyl ether, stilbene, diphenyl or phenanthrene. Although his experiments disagreed with Hibbert's work, Hurd did not choose to explain the mechanism of the reaction by

^"normal" means decomposition into olefin and acid. -13- a free radical scheme (31). Instead he postulated a bimolecular decomposition of the ester to give dibenzyl ether as intermediate. This subsequently decomposed quantitatively giving benzaldehyde and toluene. Thus

Hurd was able to account for the production of benzoic anhydride and the presence of toluene in his experi­ ments.

B. Hurd's "Hydrogen Bridge11 Cyclic Mechanism

Somewhat later Hurd and Blunck (5) came forth with a comprehensive study on the pyrolysis of esters. They

gave quantitative information on seven esters: ethyl

acetate, iso-propyl acetate, tert-butyl acetate, iso­

butyl acetate, phenyl acetate, ethyl a-phenylacetate,

and methyl a-phenylacetate. From the data Hurd and

Blunck concluded that great instability was contributed

in acetates by a tertiary alkyl group, in the alcohol.

Secondary alkyls (iso-propyl acetate) exhibited inter­ mediate stability and primary alkyls showed comparatively

stable characteristics. He found small amounts of

acetaldehyde at 525-550°C. in the pyrolysis of ethyl

acetate and acetone; and propionaldehyde and carbon

monoxide in the decomposition of isopropyl acetate at

430-460°G. Similarly iso-butyraldehyde was found at both 420 and 650°G. when isobutyl acetate was pyrolyzed.

The three aromatic esters required higher temperatures and did not decompose "normally" as in those cases when -14- beta-H was present.. In order to explain these results

Hurd and Blunck attempted to give a single mechanism but realized that it would fail to account for all the facts collected. Hence they assumed three types of processes were involved in the pyrolysis studies. One of these was taken to be a chain mechanism with univalent radi­ cals as intermediates. To tdst this Hurd investigated tert-butyl acetate and found that the predicted products: ketene, toluene, and methane were totally missing. In a second example the predicted formation of formaldehyde, methane, phenyl ketene, toluene, carbon monoxide and acetaldehyde from the pyrolysis of ethyl-phenyl acetate was not substantiated by results. However in the case of ethyl acetate there was some apparent consistency.

The interaction of radical (R*) with ethyl acetate could lead to the following species:

(ix) •ch2-co-o-ch2ch3 (x) ch3-co-.o-ch-ch3

(xi) ch3-c'o-o-ch2ch2 .

These on decomposition would be expected to give (IX) formaldehyde, ketene, and methane, (X) acetaldehyde,

carbon monoxide and methane, and (XI) ethylene and acetic acid. These products were indeed found although not in the anticipated amounts. Furthermore the large amount

of acetic acid recovered made an "a priori1” supposition of the equal importance of (IX), (X) and(Xl) unsuitable.

Hurd therefore proposed that the yield of acetic acid -15-

may well be from some other source. Thus as a result

of the low temperature of decomposition of tert-butyl

acetate and the lack of appreciable amounts of other

products caused Hurd and Blunck to propose the "Cyclic

Hydrogen Bridge" mechanism, which they believed opera­

tive in some instances such as this. To amplify this

further the concept of hydrogen bonds (32) developed

to explain the association of water or the acidity of

weak organic acids was utilized by Hurd in the series

of moderately high temperature pyrolyses. He postu­

lated that any ester possessing a beta-hydrogen in the

alkyl group may undergo a chelated type of 6-atom

ring closure by way of a "hydrogen bridge". Following

this, an electron redistribution occurred. Illustrated

as follows:

1- ■ H H N Jv/ \ ■ H : 0: CR2 ;0 CR2 :0 cr2 II I ---- y 11 I ^ I ti R-C CR2 C CR2 C. " \ / / v■ / / o + cr2 0• * itO R

By a combination of this and a chain radical type of

mechanism Hurd was able to reconcile the experimental

findings. In special cases, however, it was necessary

to introduce a third intermediate, the methylene radical, as

with methyl a-phenylacetate which gives benzaldehyde. 7 -16-

C. Ritchie's Dual Reactions

Contemporary with this, studies were made by

P. D. Ritchie (10) and coworkers. Their interest centered in the pyrolysis of acetylated hydroxy esters.

From a generalized point of view their work concerned’ the acid rather than the alkoxy portion of the ester.

Ritchie proposed a mechanism of the two competitive re­ actions as follows:

GH2“CH-C02CnH2n+1 CH^ qAc Ch3 0Ac

j h v y '

CH3G00H H G0-°-CnH2n+l H C02H

He found that route (l) was exclusively followed when n=l and that both (l) and (2 ) were simultaneously

operative in the case of n=2 at 500°C. However with

the change of temperature (l) may be selectively opera­

tive with respect to (2) or vice versa. When n=4 route

(2) became predominant. The products formed were acetaldehyde, butylene and carbon monoxide in large

amounts. The formation of these compounds was explained

by pyrolysis of secondary substances (the alpha-acetoxy-

acids). When phenyl alpha-acetoxypropionate was pyrolyzed

Ritchie found an adherence to route (l) with the exclusion

of (2). He explained this by the known thermal stability

of benzyl radicals. This led us to deduce that Ritchie may also have favored a free radical explanation though -17- he did not choose to express himself thus for the primary- process.

The alpha-acetylated hydroxynitriles were also studied by Ritchie and coworkers. In general they pre­ ferred a dual route of decomposition as follows:

R R i ch2 =c-cn CH3 ® GHq-C + HCN

CH3COOH R ON

At approximately 450°C. route (3) predominated with a little route (4-) • The absence of ketene per se was attributed to the formation of acetic anhydride. In the case of 1-acetoxy-l-cyanocyclohexane (XII)

Ritchie proposed the above reactions and was able to identify cyelohexanone (XIV) by 2,4-dinitrophenylhydra- zine derivative. He found, however, neither ketene nor acetic anhydride as predicted.

Ritchie stated that the course of decomposition of acetylated hydroxy-esters is mainly influenced by:

(a) the relative positions of negative acidic groups and the carbon atom to which the acetoxy group -18- is attached, and

(b) the degree of substitution at this carbon atom.

Ritchie illustrated effect (a) by a series of three esters.

(XV) CH3CH20Ac --> CH2=CH2 + AcOH

(XVI) CH^CH(CN)oA c --- * GH2=CHCN + AcOH

(XVII) CH2 (Cr)CH20Ac --- * CH2=CH(CH) + AcOH

In connection with the parent hydroxy-eompounds Ritchie predicted the ease of deacetylation to be (XVII)> (XVl)>

(XV). His experiments showed that it was indeed so.

In. addition to the results above he showed that (XIX) pyrolyzed easier than (XVIIl).

(XVIII) < T ) - 0 4 t (XIX)

He intimated that the fact the alpha-carbon has changed from secondary to primary may have a subordinate effect.

Some confirmation of Ritchie's work came from the studies of Wibaut (9) and coworlcers. They were interested primarily on the preparation of 2 ,2-dimethylpentenes through the pyrolysis of the corresponding acetates.

Wibaut was able to show the effect of substitution on the alpha-carbon atom on the ease of pyrolysis. Thus he found only one dimethylpentene in the pyrolysis of

2,2-dimethylpentanol-3-acetate. Furthermore, the pyroly­ sis of 3,3-dimethylbutanol-2-acetate and 2,2,4-tri- -19- methylpentanol-3-acetate gave exclusively 3 ,3-dimethyl- butene-1 and 2,4>4-trimethylpentene-2 (33). However,

Cramer and Miller (34) disputed these findings and reported 7% 2 ,3-dime thylpentene and 2 ,4-d'imethylpentene from the pyrolysis of 2 ,2-dimethylpentanol-3-acetate.

The Gramer-Wibaut disagreement (35) on the possible isomerization of the carbon skeleton contributes little to the mechanism of pyrolysis but it did later help in studies on the preferential cis-elimination. This will be considered in the following section.

D. H&ckel's Concept (Tsugaev Keadtion).

The theory of cis-elimination and steric contri­ bution came primarily from an investigation of the Tsugaev

Reaction. Earlier workers such as Hllckel, Tappe, and

Lengutke, Stevens and Richmond, and others favored a cyclic mechanism as follows: © CH' h 2c 0 = c = s \ c=sC X CH. HS H2C: HS - CH. \ CH.

According to Htickel the methylthio-group (CH^S^.) plays the same role as an alcoholate ion does in bi- molecular base-catalyzed elimination (36).

CR2 II + Y“ + XH CRo -20-

The proton splits from the carbon heteroly;fcically whereas the C-0 bond is split simultaneously retaining a pair of electrons on the oxygen. The olefin formed is primarily in a completely polarized state. A slight modification by SteVens was the inclusion of a possible intramolecular hydrogen bonding prior to the proton elimination. He depicted the mechanism as follows:

R - CH - CHR R - CH - CHR R - CH - CHR / I / " + ( " 0 H °, I ' \ / - \ + .— » C = s G — S!H ; R -CH = CHR 1 j. ! S - CH- iS - CHJ .JJ

From this concept Stevens extrapolated to the esters of acetates, oxalates, benzoates and phthalates as they also showed a preference for cis-elimination.

E . Alexander's Cis-Elimination Studies

In addition to the evidence of cis-elimination offered by Stevens (37) and Hllckel (6 ), Alexander and

Mudrack (3S) showed methyl cis-2-phenylcyclohexyl xanthate gave 93 per cent 3-phenylcyclohexene and 7 per cent 1-phenylcyclohexene, whereas methyl trans-2- phenylcyclohexyl xanthate gave 14 per cent 3-phenyl­ cyclohexene and 86 per cent 1-phenylcyclohexene. Com­ parable results were obtained by Marvel and Williams (39) in the pyrolysis of carboxylic esters. They also -21- proposed a cyclic mechanism:

'Stlx > \ h C :•> o H'c -c-'^'Vs -c. ’o ✓ I I I II''' I I + I -a N Q ^ ' P - L ^ '-'ey'®' i

The significance of this mechanism was, as pointed out by Alexander (40), that the departing group is not dis­ placed by the rearward attqck of an unshared electron pair.

Two more examples were given by Alexander (40) to

show conclusively the preferential cis-elimination of

carboxylic acids ffom esters. This evidence stressed

the importance of a cyclic mechanism, although the

claims by Cramer supported by Houtman, Heertijis, and van Steenis (8 ) do not totally bear out an intramolecular mechanism of this nature as the sole channel for the

decomposition of esters. Houtman and coworkers claimed that carbon deposits were necessary if lower-temperature

pyrolysis were to be obtained. They claimed that the

small trace of acid present in es.ters initiated decompo­

sition and that the carbonized material then coated the

surface of pyrolytic columns so as to enhance the ease

of ester cleavage. This surface effect was believed

necessary by Houtman et.al. with esters which possess

$ strong a -C-O-R group containing a/C-0 bond. In contrast to -2 2- earlier investigations they found some isomerization of double bonds. From the pyrolysis of butanol-1- acetate, 1.5 per cent of butene-2 was detected, while in the pyrolysis product of 2-giet£yl butanol-l-acetate some indication of 2-methyl butene-2 was evident. How­ ever structural isomerization did not occur in their experiments. Houtman, van Steenis and Heertijis stated that in the case of a secondary allcyl-group the ease of beta-hydrogen release will determine the ratio of isomers.

For example 2-acetoxybutane yielded .56 per cent butene-1 and 44 per cent butene-2 at 500°C. but only 25 per cent butene-1 at"400°G. The equilibrium quantity will not appreciably change at temperatures above 500°C.; therefore the ratio of rates of beta-hydrogen split will determine the respective amounts of the butenes. They concluded that a primary carbon atom will release a hydrogen atom much easier than a secondary carbon atom and that a tertiary carbon atom will lose a hydrogen atom only with great difficulty, hence the yield of 2-methylbutene-l is lower than 3-methylbutene-1 under the same conditions.

F. Kinetics of the Pyrolysis of Esters

It appears, at this point, more beneficial to con­ sider an aspect that thus far escaped the investigators.

In 1946 Rudy and Fugassi (13) studied the thermal de­ composition of tertiary butyl acetate. Their system of pyrolysis was a static method at reduced pressures. At -23- a range of 243 - 303°C. they found that heterogeneous and homogeneous decompositions occurred. By letting the products of the pyrolysis remain in the vessel for a period of hours the heterogeneous reaction disappeared and the reproducibility of rates became stabilized. The homogeneity of the reaction was tested by a plot of the log k versus l/T. In the stabilized condition the re­ action exhibited first order kinetics and at 303°G. and

132 mm. of mercury the activation energy was 40.5 Kcal.

Indications of an equilibrium were observed, Ke being 5 approximately 10 . This showed that the reaction was essentially complete. Fugassi concluded that a free radical mechanism was not suitable since the simplicity of the products rendered such a postulate difficult to maintain.

On the other hand a cyclic intramolecular mechanism, on the basis of the transition state theory, should re­ sult in an abnormal frequency factor, but this was not in accord with the experiments represented by the equa­ tion:

log k = 13.342 - 40500/RT. (41)

A second study by Fugassi and Warrick (41) was performed on the decomposition of tert-butyl propionate.

The system was again stabilized with the pyrolysis pro­ ducts and the homogeneity of the reaction was confirmed by plots of log k versus l/T to yield stra ight lines. -24-

The kinetcis of the reaction followed. the rate equation:

log k = 12.794 - 39160/RT..

Again, estimates of the equilibrium constanttshowed that the reaction \*as essentially complete (Ke = 3 x 10 ).

The lowering of the frequency factor was attributed to an inductive effect of the propionyl group.

G. Some Contributions from the Kinetic Study of

Dehydrochlorination of Ghloronaraffins

In parallel with the pyrolysis of esters a thorough and fundamental study of the dehydrochlorination of chloroparaffins has been undertaken by D. H. R. Barton and his coworkers (42). In repeating the work of Biltz and Kupper (43) Barton found that he was able to increase the rate of reaction immensely with small traces of oxygen as catalyst. By carefully purifying the commer­ cial reagent (l,2-dichloroethane) he showed that the unintentional reduction in rate was caused by the inhibi­ tion of the oxygen effect due to the presence of small amounts of chlorohydrin. By removing this impurity

Barton was able to show that small amounts of oxygen,

chlorine and other halogens could significantly influence the rate of the decomposition; and that the variation

of surface-area-to-volume ratio could change the rate

of pyrolysis. Barton discovered that irreproducible results were caused by clean glass walls but that by allowing the reaction product to remain in the reactor -2 5- over night the heterogeneity of the system was suppressed.

In the nearly homogeneous system Barton found that all halides observed followed conveniently a first order rate equation:

k = A exp. -E/r T.

From this relationship and the observed velocity of the decomposition Barton was able to calculate the activa­ tion energy of each halide studied. For example:

Ethyl chloride A = 3.80 x 107 E„ = 32 Kcal a 1.1-Dichloroethane 5.89 x 107 33.9

1.2-Diehloroethane 1.59x 106 2711

To gain insight into the mechanism of non­

induced reactions, Barton and Howlett (4-4) studied the

decomposition of 1 ,2-dichloroethane in a nearly homo­

geneous system at a temperature range of 362 to 485°C.

The rate showed a first order reaction, and the products

were pure vinyl chloride and hydrogen chloride with a

trace of acetylene. In addition to the activation energy

obtained (47 Kdal.) and the frequency factor (10^®) Bar­

ton observed that a small amount of propylene greatly

inhibited the rate of decomposition. Thus he showed

that a non-inhibited reaction required an induction period, which, when treated as a rate, gave an activation energy

of 78ll0 Kcal. The fully inhibited reaction on the other hand could be expressed conveniently by the following -26-

equation:

k = 6.3 x 10^ exp. -4-8 Kcal/RT

With the aid of these facts Barton considered four

possible radical chain schemes. The energetically

favored route was selected by Barton and the following

chair mechanism was proposed:

i. CH2C1CH2C1 ---* C2H4+ C12 Ea (Kcal»)

ii. Cl2 + C2H4C12 — * HC1 + 01 + C2H3C12 smatl iii. Cl + C2H^C12 — - HC1 + CgS^Clg 5

iv«. C2H3C12 ----CH2=CHC1 + Cl 21.5

v. Cl + C2H3C12 --- ► HC1 + C2H2C12 3

The energy assignments were calculated from values given

by ICistiakowsky (45) and Steacie (4 6 ) and the application

of the stationary state to the above scheme gave

dp/dt = ( k ^ / k k^)1/2 (CH2C1 CHgCl). Consequently

Ea = 1/2 (B1 + + E5) = 48 Kcal.

The kinetics for 1,1-dichloroethane and ethyl

chloride were also studied by Barton and Howlett (47).

These halides were found to follow first order kinetics

in a nearly homogeneous system and exhibited no inhibi­

tion by propylene at considerably high pressures. Barton

considered these decompositions as truly unimolecular

homogeneous reactions. Consideration of the energy

relationships of the reverse reaction indicated that

the activation energy is approximately 63 Kcal. -27-

The calculation according to the transition state theory showed the following equations to be representative of facts.

^ethyl chloride = 1.6 x 1 0 ^ exp. -.59,500/RT

^1,1-dichloroethane = 1.2 x 10"^ exp. -49,500/RT

The study of the decomposition of tert-butyl chloride by Barton and Onyon (4-8) was probably the most significant and representative of the unimolecular homo­ geneous reactions. Kistiakowsky (49) earlier had established the rate equation to be:

k = io13’9-0,7 exp. -45,000 t 1900/RT

However the work of Barton and Onyon was done with more refinement and as a consequence the frequency factor was established at i o ^ » 4-0.02 the B„ was found a to be -41,400 ± 600 Kcal. Barton tested his system thoroughly in an effort to establish the heterogeneous reaction. He demonstrated that serious contamination of the homogeneous decomposition could be present if the system was not carefully isolated from air and thus may have caused the inaccuracy in Kistiakowslcy * s work.

The heterogeneous reaction was found to be con­ siderably faster. When the reactor was clean the follow­ ing average velocity was found:

k = 1.44 x 1CT6 sec.

And the activation energy was approximately 15,000 * 6000 cal.

In addition to these findings Barton found a relationship -28- of the glass-surface-to-volume ratio with the kinetic order observed.

Reactor Surface/Volume Order of Kinetics

A (clean) 2.72 cm"-*- 0 B (clean) 9.98 cm*"*! o

A(after 14 runs) 1 to 2 B(after 14 runs) 1 predominantly

In the work that followed Barton and Onyon (50) studied 1,1,1-trichloroethane. They eliminated the heterogeneous reaction by stabilizing the system with continuous decomposition of alkyl halides until the rate became reproducible. Under these conditions they found two simultaneous operative reactions occurred. One of these could be suppressed by the addition of propylene and under maximal inhibition the reaction Slosely follows a course, predicted by first order.kineticsj in fact, the type of reaction is characteristic of. unimolecular de­ composition (A = 10^-2 ), Since the overall velocity of the reaction showed also very little deviation from first order rate Barton subtracted the rate constant of the unimolecular reaction from that of the overall reaction and obtained the following data:

-o“&i = K = 1012»53 exp. -47.9/RT (empty reactor)

^o"^i = R 1 = 1012.51 exp. -48 .4/RT (packed reactor)

In view of these facts Barton proposed for the inhibited reaction a free radical chain mechanism. -29-

(1) The Initiation Process kl CI3C-CH3 — =*- CH3CCI2 + Cl k2 — CH3 + CC13 k3 — CH2CCl3 + H

is not energetically favored.

K2 and both may be operative.

(2) The Propagation Process

ci + ch3cci3 > hci + ch2cci3 CHgCCl^ —---* CH2=GC12 + Cl

(3) The T ermination Process

Cl + CH2CC13 -----* products

and also

Cl + CH^CCl^ + M products + M

This was quantitatively shown to be in agree­

ment with experiment.

Further work by Barton, Head and Williams (51) showed the decomposition of 1 ,1 ,1 ,2- and 1 ,1 ,2 ,2-tetra- chloroethanes proceeded by radical chain mechanisms exclusively. These compounds possessed a well-defined induction period which was expressed by Barton as

I = 1 0 ^ exp. 44,000/RT min. and it was attributed by Hoi^lett (52) to a slow building up to the steady state. He further established that the existence of a homogeneous radical chain mechanism in such a system was real. Other data for unimolecular homogeneous decomposition may be summarized as follows: -30-

Compound Temperature A (Sec“^) a n-propyl chloride (51) 693-751°K 1013.45i.35 55^1.2 n-butyl chloride (51) 700-714°K lO1^-*1 57i.4

Isopropyl

1013.4l.2 1 0 • Ul . VJ

Chloride (53) 640-679°K + 1

1 ,2-dichloro propane (53) 689-725°K 10!3.8 54.9

^Activation energies measured in Kcal.

The general mechanism for the decomposition of alkyl halides observed by Barton may now be summarized as follows:

(1) Heterogeneous decomposition on glass surface.

(2) Homogeneous first order unimolecular decompo­ sition.

(3) Homogeneous first order decomposition by radical chains.

The homogene-ous first order unimolecular decomposition was regarded by Barton as a good example of the four center planar transition state mechanism (54). The radical chain mechanism followed four general steps which were outlined by Barton as follows:

(1) Initiating step or steps (kinetically first order) leading to production of Cl atoms.

(2) Cl atoms attacking substrate, abstracting hydro­ s'® a , giving large chlorine-conta ining radical. -31-

(3) DecompoS'ition of the large radical to give olefin or chloro-olefin and- a chlorine atom to carry on chain.

(4) Termination to give non-radical products by bimolecular reaction of Cl radical with large radical containing Cl.

Most frequently the unimolecular homogeneous decompo­ sition is operative but it may be accompanied by a radical chair mechanism. To correlate the influence of the concerted mechanism by radicals Barton formulated the general rule:

A chloro-compound will decompose by a

radical chain mechanism only if neither the

compound itself nor the reaction products

are good inhibitors for the chain.

Furthermore only radicals which can decompose to give chain carrying species (as of above scheme step 3 ) can produce an operative radical chain mechanism. Any re­ sonance stabilized radicals (55) in step (3 ) will tend to terminate the reaction, in favor of the competitive unimolecular homogeneous decomposition. This was tested by Barton in the case of dichloroethyl ether (51).

The cyclic ^transition state of .Barton was extended to xanthate and ester pyrolysis. In a sequence of studies pertinent to the stereochemistry of cis eliminations

Barton and Rosenfelder (56) showed that the original -32- proposal by Hurd, later modified by Hiickel, and by Alex­ ander, was indeed plausible. However, an additional explanation was necessary to account for the selective reactions observed at variance with Bredt's Rule, since in some cases elimination toward the bridgehead of a molecular structure could happen. Barton's reasoning made use of the simple inductive effect. He stated that since the homolytic bond dissociation energies, for

-C-X, R2 =C-X, and R 3sC-X, where X = OH, H, Cl, Br, I, etc., are in decreasing order, the comparable H-X unimolecular elimination in the series -C-X, R2=C-X,

R^-C-X, where X is the same, must also require de­ creasing energies. Thus with the same frequency factor the lowering of the activation energy would necessarily require a selectivity of the bridgehead (hydrogen for elimination. These were exemplified by a sequence of structural assignment studies on the pyrolysis of acetates and benzoates of steroids (56).

3. The Cyclic Transition State in Other Processes

A . General Considerations

More evidence for the cyclic transition state in ester pyrolysis may be cited from other related fields.

In the study of the thermal decomposition'of acetic anhydride (XX), Szwarc and Murowski (57) obtained a rate equation: k = 1 x 1012 exp. -34,500/RT -33-

They tested the possibility of free radical chain mechanism with the aid of toluene and i^ere convinced that the formation of ketene and acetic acid was due to a unimolecular decomposition without radical inter­ vention.

CHg - CH2 =C = 0 (XX) (ch3 co)2o ---- > , ^ 0 _ 0/° >

X ° H 3 CH 3 COOH

The pyrolysis of acetyl bromide (XXI), by the same authors showed that a concerted mechanism was operative even at very high temperatures (600-800°C.). Differenti­ ation between the two possible unimolecular routes was difficult,

k - CH 2 =C=0 + HBr

(XXI) CH3CO Br

lc, CH-Br + CO

but a calculation of the ratio of CO/HBr formation in­ dicated that was preferred.

The rearrangement of allyl vinyl ethers (XXII), by

Schuler and Murphy (58) may also be mentioned. The gaseous pyrolysis of allyl vinyl ether yielded only allyl acetaldehyde. The authors observed a frequency factor of 5 x 1 0 ^ a nd Ea = 30.6 Kcal. with a negative entropy of activation of -5 e.u. The proposed mechanism 34' was as indicated

tflL ~= CH • <- V \ (XXII) GH2 a GH - CH2 - CH2 - GHO

The pyrolysis of alkyl borates by Brandenberg and

Galat (59) constituted another example of the cyclic transition state. However, great significance cannot be attributed to this reaction since the authors were not sure whether isomerization had occurred in their olefins.

B. The Concept of O'Connor and Nace

More significant in this series of comparative studies is the work of O ’Connor and Nace (60). They studied the decomposition of beta-cholestanyl methyl xanthate and cholestenyl methyl xanthate. These reactions were found to possess the expected first order unimole­ cular kinetics, a frequency factor of 10“*^ and an Ea of

33.8 Kcal. and 32.9 Kcal. respectively. O'Connor and

Nace, following the general kinetic interpretation, believed that the negative entropy of activation was due to a highly ordered transition state and hence postulated that the six membered methyl xanthate ring must be resonance stabilized like that of a benzene nucleus with six If electrons at its disposal. The authors were also of the opinion that the pyrolysis mixture at a given stage was composed of a hybrid of this stable -35- tra nsition state and the starting material. Thus in decomposing only a shift of the nuclei was required, followed by the redistribution of the electrons to give the necessary products. They further stated (however without evidence) that the shift of electrons need not

involve a pair but only one by free radical reaction.

In order to further demonstrate the stability of

this six membered transition state O'Connor and Nace (60) made a series of comparative rate studies of carbonates,

acetates, cliloroacetates, benzoates and xanthates. In

all cases they found a first order kinetics was followed 13+i and a frequency factor of 10 ”*L as prescribed by the

transition state theory for unimolecular decompositions.

However the variation of activation energy was interest­

ing. Thus in comparing cholesteryl acetate, ethyl car­

bonate and S-methyl xanthate Nace found a decrease in

Ea values: 44.1* 41.0, and 32.9 Kcal. respectively.

This information served the authors well in explaining

the relative stability of acetates.

Taking the general formula

-t H‘? i' ?11 — ' i — /• . + //I D■ ’I 0/ R

R* = R, Ar, SR, OR, NH0. * Cl

Nace was able to show the influence of inductive effects

on the stability of the transition state. He observed in general (l) the nucleophilic character of the carbonyl oxygen determined to a great extent the stability of the esterj the more nucleophilic these atoms the more un­ stable were the compounds. Hence, the more electroposi­ tive R* the less stable is the ester. Consequently in the series phenyl carbamate, ethyl carbonate, ethyl benzoate, and ethyl acetate stability increased in this order. The reason for the instability of benzoates was expressed as follows:

The stability contributed to the transition state through resonance made the ester less stable. However the lower stability of the chloroacetates was hot ex­ plained. (2 ) The second factor which influences the transition state could only be localized to the xanthate esters. The arrangement of oxygen and sulfur gave maxi­ mum instability; hence the additional driving force by the energy gained from a -0-C=S linkage to -S-C=0 linkage enhanced the ease of decomposition of the xanthate esters. This was postulated by Tarbell and Harnish for other reasons (6l). The influence of the R ‘ group in the xanthates was just the reverse of the carboxylic esters. The electronegativity of the R* here favored instability of the compounds. This apparent anomaly -37- was regarded by Nace as due to the stronger influence of the sulfur-oxygen group rendering the inductive ef­ fect relatively unimportant.

The final statement of O'Connor and Nace was only an empirical observation but it is interesting to note that "an electron deficiency about the central carbon lowers the activation energy for the formation of the cyclic transition state by facilitating partial bond making with respect to bond breaking in the linkage

-0-C=S." Thus the more electronegative R* the greater is the tendency for bond making to proceed rather than bond breaking in the transition state. This in effect is supplying excess energy from the partial C=0 bond formation as a driving force (62).

C. The Pyrolysis of Sulfites

The last comparative study we wish to single out in order to illustrate the concept of a cyclic transi­ tion state is the work on the pyrolysis ofsulfites by

Price and Berti. They showed in a series of pyrolyses of sulfites that the major reaction for decomposition was (l) olefin and alcohol formation, (2 ) ether and sulfur dioxide, and (3 ) rearranged products and sulfur dioxide. In class (l) the pyrolysis was limited to the primary and secondary alkyl sulfites. Of the six sulfites studied by Price and Berti (63) the general equation may be written as follows: -38-

rii_CH2-GH^OSOOR _____+ S02 + ROH +. R" CH=CHR» i R>

The temperature for rapid decomposition was a-methyl— benzyl sulfite; 130°C.; n-amyl sulfite, 200°C. and a-c carbethoxybenzyl sulfite, 240-260°C. The marked stability for a-carbethoxybenzyl sulfite was-thought to involve an ion pair contributing to the transition state as with chlorosulfinates (64). Thus ionization in a sense + R-OSOOR R OSOOR would be favored by the presence of a methyl group as in a-methylbenzyl sulfite and hindered by the carbethoxy group as in a-carbethoxybenzyl sulfite. Further evi­ dence of such contributions of ionic nature was observed in the decomposition of phenyl cyclohexylmethyl sulfites.

In the second and third classes the glycol sulfites were representstive. In addition to the work on 2,3- butylene glycol sulfite (65), Price and Berti (66) studied the pyrolysis of the cyclohexanediol sulfite

(XXIII), (XXIV) and the hydrobenzoin sulfites (XXV),

(XXVI). In the case of the cyclohexanediol sulfites the isomers were represented as follows:

(cis)

(XXIII) H - 3 9 ‘ (trans) cyclopentaldehyde

(XXIV) and in the case of hydrobenzoin sulfite:

(ois) \ _ / *0 £ /tj> 4 4>

\ / ° / \ /C ^ H- H uoSO, . i (XXV) (]>-£ -aUt-cj)

H 4> i ' , /' ri- , (trans) C. - C A r V - - ® - - 1* '/' ' ^ ( b ° 0 H / ,\b s-So/ O30z 1+oSo^ ^ t

(x x v i) oh«l- c »

It is interesting to observe that the presence of BaO causes a rearrangement in n-amyl sulfite. This was re­ presented by Price and Berti as follows:

. BaO (R0)2S0 ----- «RS020R

The mixed alkyl sulfites containing a methyl as one of the alkyl groups were thoroughly investigated by

Brice and Berti (66) as a means of preparation of olefins.

Thus methyl 3-phenylpropyl sulfite, methyl a-methylphen- ethyl sulfite, methyl cis-2-phenylcyclohexyl sulfite, -40-

methyl trans-2-phenylcyclohexyl sulfite were pyrolized.

The scope of this reaction was found to be somewhat

limited to secondary alkyl groups since primary alkyl

methyl sulfite gave a disproportionstion reaction

yielding dimethyl sulfite and symmetrical alkyl sulfitej

consequently the yield of olefins was poor in comparison

with secondary alkyl sulfites. Price and Berti were

encouraged by the analysis of products to propose a free

radical mechanism for this reaction in competition with

the cyclic transition state type of decomposition.

( i.) 0-CH2-CH2-CH2-SO2-.OCH3 --> ^-CHg-CHg-Cttg• + S6 -0CH3

so2+gh3o.

(ii.) 0-CH?-CH„-CH » ----► polymer . ^ disproportionation reaction products

(lii. ) CH30* + EH ^ CH^OH + R*

(iv.) 2 CH^O* --- ^ GH2 O + CH^OH was not identified

The exchange reaction was found to be sensitive to basic

catalysis. The olefin formed showed little selectivity,

as pointed out by the authors, from this conjugative

effect; for example in the case of phenylpropyl methyl

sulfite both propenyl benzene and allyl benzene (60 to

40 ratio) were formed.

The complete disagreement with xanthate pyrolysis

in the case of cis- and trans-phenylcyclohexyl methyl

sulfite showed that the transition states of the two -41- reactions were somewhat dissimilar. In the sulfites the predominant product for both the trans- and cis-isomer was 1-phenylcyclohexene. This and other evidences

caused the authors to predict that some non-classical bridged-ions were present in the process of decompo­

sition of sulfites. One case of agreement with the

Chugaev Reaction was noted, however: the production

of methane in the decomposition of 1-menthyl methyl

sulfite was close to that of xanthate pyrolysis'i. -42-

DISCPSSION OF EARLIER WORK

1. Kinetics as a Method for Mechanism Study

The introduction of kinetics into the field of

pyrolysis has resulted in a better understanding of

the chemical nature of dissociation in organic compounds.

On the whole it has made possible a definite differenti­

ation of systems that earlier workers had not observed.

Pyrolysis today is specified as occurring either in

liquid or vapor phase. Thermal decomposition in a

flask, as in the case of xanthates, and an ester pyroly­

sis in a flow reactor, as in the cases studied here are,

indeed of two entirely different types. The gaseous

decomposition which is generally susceptible to kinetic

studies may yield not at all the same results when per­

formed under pseudo-vapor phase conditions or partially

in the vapor phase. Such confusion must be absent if

clear and definite information is to be obtained. It

is evident from the work of Barton that the vapor phase

pyrolysis can be of two kinds, heterogeneous and homo­

geneous. The system must be confined to homogeneous vapor

phase decomposition or inevitably a complex rate of re­

action will be observed. We shall now summarize in a

general manner the following list of compounds which

are first order gaseous decompositions at 285°C. -43-

Table I.

Some First Order Gaseous Decompositions at 285°C.

*Part (A) A x 1013 Ea si Compound (sec.-1) (Kcal.) (e.u.)

Propylene oxide 12 58.0 6 Ethyl peroxide 51 31.5 8 n-Propyl peroxide 230 36.5 11 Diacetyl 870 63.2 14 Methyl nitrite 1.8 36.4 2 Ethyl nitrite 1.4 37.7 6 n-Propyl nitrite 28 37.7 7 Isopropyl nitrite 13 37.0 6 Azomethane 3500 52.5 17 Isopropyl iodide 1.6 42.9 1 t-Butyl bromide 2 40.5 2 Silicon tetramethyl 160 78.0 11 Trioxymethylene 1500 47.4 15 Nitrogen tetroxide 80 13.9 9 Triehioromethyl chloroformate 1.4 41.5 1 Methyl azide 300 43.5 12 Ethyl azide 20 39.7 7 Azopropane 5.7 40.9 4 Nitromethane 4.1 50.6 2 Trimethyl acetic acid 48 65.5 8 Dimethyl ethyl acetic acid 3.3 60.0 3 Paraldehyde 130 44.2 10 Parabutyraldehyde 24 42.0 7 Dicyelopentadiene 1.0 33.7 0 Isopropyl bromide 4.0 47.8 3 Benzyl bromide 1.0 50.5 1 Ethyl chloride 16 59.5 6 Ethyl nitrate 630 39.9 14

^Frorn Kinetics and Mechanism" by Frost and Pearson, Chapter 6 , p. 104-105. John Wiley & Sons, 1953. -44“

Table I, cont.

*Part (B) 10 E _± A x 10 a S + Compound (sec. ) (Kcal.) (e.u.

Ethyl chlorocarbonate 0.092 29.1 -18 Methylidene diacetate 0.17 33.0 -17 Methylidene dibutyrate 0.17 33.0 -17 Ethylidene dibutyrate 1.8 33.0 -12 Ethylidene diacetate 2 32.9 -12 Iieptylidene diacetate 3 33.0 -11 Trichloro ethylidene diacetate 1.3 33.0 -13 Furfurylidene diacetate 13 33.0 - 8 Glyoxal tetraacetate 180 39.2 “ 5 Methyl iodide 390 43.0 - 1 Vinyl allyl ether 50 - 30.6 - 5 Acetic anhydridd 100 34.5 - 4 Allyl bromide 210 45.5 - 3 1-Butene 500 63.0 - 1 Ethylidene dichloride 120 49.5 “ 4 t-Butyl chloride 2 50 41.4 - 2

All of these carefully selected compounds involve an unimolecular decomposition or a rearrangement as the probable first step. However to determine this from kinetics alone sometimes becomes difficult. This is particularly true in our special case of the pyrolysis

of esters. It has been quite clearly established in the work of Barton how a radical chain mechanism can be

differentiated from a homogeneous unimolecular decompo­

sition through an intromolecular rearrangement. We have

# See footnote page 43. -45- also shown other cases of such postulates as that of the acetic anhydride pyrolysis and allyl vinyl ether decomposition. However the prediction of such mechan­ isms from/structure and without experimentation requires more study.

It is generally understood through the application of the quantum mechanical theory to the derivation of rate equations (67), that the frequency factor of a unimolecular decomposition is usually in the neighbor— 13 n hood of 10 sec , if the entropy of activation is zero. Therefore, a cyclic intermediate from a non- cyclic reactant would require a decrease in the entropy of activation; hence a negative AS will result. This consequently, from general experience with dimeriza- tion reactions, would reduce the frequency factor, lowering 1 0+1 it markedly from the 10 value. However some ambiguities concerning this deduction have risen. As it was pointed out by Benson (68) in order to have a reaction appear unimolecular it is necessary that k2 (A^lc^ / ■ k2 (where k9 refers to the reaction A* + A v A + A and

k^ refers to the reaction A*, — > B + C ) but under ordinary conditions k2 (A) is not greater than

10^° sec."'*-. The anomaly is explained by the theory of

Rice, Ramsperger and Kassel, and the deviation from the quantum mechanical value has been discussed by Trotman- -4-6-

Diclcinson (69). This author stated that the low frequency t fao'tor may be looked upon as comparable to the cis- and trans-isomerization. Such a first order reaction has a

low value for its frequency factor due to a I d w trans­ mission coefficient. The assumed mechanism involved a

transition to an excited triplet state of the molecule.

Such a transition as observed by Magee, Shand, and Eyring

(70) will have a low probability due to the change in

electronmultiplicity. Since the triplet state depends

little on rotation the loss of rotational entropy would

conform to the observed negative values in the compounds

cited by Frost and Pearson (71), and in Table I. Fur­

ther theoretical discussion here is not of great value

and we may conclude that the deviation of the frequency 13 factor from a value 10 will indicate one of two paths

of decomposition.

(1) A large amount of energy is distributed over a

number of degrees of freedom, and if not deactivated,

would result in the fission of one or more bonds. This

process will show high activation energy and high fre­

quency factor. Because of deactivation by collision

this process will become more important at low pressures.

(2) A small amount of energy is placed more precisely

at given bonds which are broken and result in rapid re­

action. This process would predict low activation energy

and low frequency factor since the activating collision -47- must be accurately oriented and only one or a few degrees of freedom are involved. Because of the rapidity of decomposition of the activated species collisional de­ activation is not important band the process can operate at all pressures.

It is our present belief that kinetics can serve as a tool in the differentiation of reaction mechanism.

It is also our hope that the general kinetic result will serve as a signpost in suggesting the route of ester de­ composition. It has been pointed out that the general trend of variation of the frequency factor in respect with the entropy of activation can aid one's prediction of the primary decomposition step as being a radical chain or an unimolecular concerted mechanism. Because of the general rule proposed hy Barton on the feasibility of a radical chain mechanism under given conditions* and because of the soundness of Brown’s concept on the stabil­ ity or reactivity of a given radical in regard to chain propagation in a special system {55), we feel that the radical chain mechanism can be somewhat de-emphasized in our discussion. It is our observation and that of many others that the usual route of ester decomposition is an unimolecular one. We shall now consider this transition state in detail.

*See page 31. -48-

2. Substitution Effects on the Cyclic Transition

Complex

The cyclic transition state as originated by Hurd

has suffered many variations but its general usefulness

has always been accepted in explaining pyrolysis of

esters. One of the principal features in the postulate

of a cyclic transition state is the possibility of ex-,

plaining the high selectivity in the elimination. The

lack of any substantial amounts of olefinic isomers in ■

the pyrolysis products has always encouraged the theorists

concerning the soundness of the theory. The most encour­

aging evidence was observed by O ’Connor and Nace recently

on the effects of substitution on a cyclic mechanism.

Nace was able to show that the central carbon atom in

an ester molecule has great influence on the stability

of the molecule toward heat.

0 II A - 0 - Q - R» Central carbon atom

Thus, in the above molecule the inductive effect of E*

on the central carbon atom would affect the ease of •i passing to the transition state. On the other hand, in

establishing the selectivity of certain eliminations to

occur towards the bridgehead carbon atom, Barton showed

that substitution effects on the alkoxy-side of the

molecule might also contribute to the transition state. For the elimination of HX (where X is halogen) he pre­ dicted a decrease in the energy of activation in the series

-C-H ----* =C-H * =C-H with X on the adjacent carbon atom. Comparing with the ester-transition state, Barton's four-centered planar configuration (XXVII), for an unimolecular reaction is essentially the same as that (XXVIII), in the proposed ester decomposition. t (XXVII) -0 - C- (XXVIII) -cx 0 I II Hl~' i jX i 1

Therefore, one may expect the ease of homolytic dissoci­ ation to be in the sequence of tertiary, secondary, primary for HOAc cleavage. It must be remarked, however, in applying this to alicyclic compounds, the effect is minor. Furthermore, such a postulate directly contra­ dicts the statement of Houtman, Heertijis and van Steenis who said that a beta-tertiary hydrogen atom in the allcoxy portion of an ester was very difficult to remove, and the ease of a proton release from a primary carbon atom was comparatively greater than that from a secondary.

Despite the controversy, the final structure of the transition state, (St), below shows that both starred carbons have influence in the formation of the transi­ tion state following the reasoning of Swain. It now -50- - seems obvious that one should question whether the third carbon atom in the ring can also affect the results of elimination. This carbon which we shall now refer to as the ’’a1' carbon atom may be substituted with various groups: R- (alkyl), AcO-, -COOK, R-CH20-, CH =,

CH2=CH-, CH2=CH-CH2-, (cyclic-CH2 )x-CH=, -CN, -CH2-CN, or any two of these combined, etc. Let us now consider

r." u S * a group,' the aliphatic alkyls (R-).

V L jl-K

Table II. Pyrolytic Decompo sition of Certa in a--Substituted Ace ta tes of ' A* Compounds Substi tution Substitution Refs. Temp. Decomp. (Group) on C** on a C R R ,t X x»

XXIX H H Me H 4,5 430- 37.7(d) 460- 55.0(d) 470 97.4 XXX H H Et H 5,8 460- 80. 0- 500 90.2 XXX a Me H Et H 9 430 71.0 XXXI H H t-bu H 33,35 400 8 .0 (d) XXXI a •Me H t-bu H 34 349 11.5 401 7.7 XXXIb Me Me t-bu H 9 460 75.0 XXXII H H Me Me 8,5, 360 2.7(d) 13 350 78 XXXIIIa Me H Me Me 72 155* 96 XXXIIIb Me Me Isopr. H 9 450 88 XXXIV Me Me H H 4,5 500 77,8 420 13.6 650 98.1

(d) = true measure of decomposition * This is not a true vapor phase pyrolysis. - 51-

In the series substituted by an aliphatic alkyl on the a-C atom it is hard to determine how much influ­ ence is exerted by the or by the a-C atom to the transition state. The obvious reason lies in the fact a that a rotation about the -C-0- bond would interchange the position of other £-C atoms with the atom. Thus any contrast that one may present such as groups XXIX to XXXIV, XXXa, XXXIa, to XXXIIIa, or XXXIb to XXXIIIb, cannot signify unambiguous selectivity due to one special effect. In order to confirm or deny either the proposal by Heertijis, van Steenis and Houtman or the reasoning of Barton in explaining the selectivity of olefin forma­ tion, one must obtain some correlation of the structure of the ester with thermal stability. Unfortunately the above ten compounds show only a qualitative trend. Since the data by the various authors do not allow a critical comparison we, therefore, must await other information.

One significant inference can be drawn, i.e. regardless of the nature of the C-H bond in the alkoxy-group it will undoubtedly affect the formation of the transition state.

Therefore the individual effects of both the C# and the G*w must be favorable in order to predict a low energy of activation.,

In the series of experimonts by Alexander and Mud- rack (4-0 ) the high selectivity of elimination can mainly -52- be attributed to steric factors. The transition state in these cyclic compounds can have one or two atoms but the configuration resulting by substitution on the p-C atom makes any study of inductive effects contri­ buting to the transition state impossible, since a con­ sideration of polar and equatorial bonds make the selectivity confused..

(73)

.0. - C H.

trans (XXXV) trans (XXXVI)

In comparing the above compounds, (XXXV) and (XXXVI), one can record that the trans-isomer of (XXXV) yielded

86.5$ 1-phenylcyclohexene at 575° to 600°C. in a total yield of 68.0$ olefin, while the trans-isomer of (XXXVI) gave only 25^5$ 1-methylcyclohexene at 490° to 500°C. in a total yield of 76.5$ olefin. It seems therefore the "H" atom opposite the phenyl group is more reactive than the T'H" atom opposite the CH^-group. This result may be explained by a resonance contribution. (XXXV)a (XXXVI)a favoring (St) formation unfavorable (St) formation

In the case of pyrolysis of trans-2-methyl-l-tetralol acetate at 550°C. (40) only 2-methyl-3,4-dihydronaph- thalene was obtained. This is obvious since only one choice was left for the cyclic transition state. The behavior of the cis-isomer confirms this as it does also in the case of cis-2-»phenylcyclohexyl acetate (92,8%

3-phenylcyclohexene). These considerations do not solve the controversy. More experimentation must be performed before a definite opinion can be stated.

The influence of the ether-group on the (3-C atom is also obscure, but some isolated cases are of interest.

(XXXVII) dihydrofuran

(XXXV:

(XXXIX)

Thus in the pyrolysis of 3-acetoxytetrahydrofuran (XXXVII)

(74) Olsen reported 19.1% furan and 68,8% dihydrofuran. -54-

The isomeric nature of the dihydrofuran was not speci­ fied. The pyrolysis of (XXXVIII) at 500° to 600°C. on quartz gave an<^ °7Ii10 was postulated by the author that the oxygen containing compound was a mixture of (XXXIX) and (XL), but no identification was made and (XLI) was not suggested. One may suspect that the tetrahydrofuran ester (XXXVII) would not yield

2 ,3-dihydrofuran under the conditions employed (83$ yield at 420° to 430°C.) (76). Therefore the product must have been the 2,5-isomer., This would lead one to think that in the case of (XXXVIII) the product would be entirely (XXXIX), But a conformation analysis shows that (XXXIX) should not be at all favored (page 53).

Hence (XL) and some (XLI) should be the products.

Isomer (XL) is more reasonable for the explanation of the C^H^0, secondary product, found.

Possible Decomposition Mechanism

for Cellosolve Esters

R - 0 - - C.H*. -OA-s_

cri-jC^O — p t CO -55-

g - o - <^z - C'Ui 'OAa-

L*i

HaC — d P p: $ S K / V

I

/? 0 -CO-C H3 -t C-^3

luheiJ £ = &■/•

The pyrolysis of cellosolve acetates and propionates in the present work showed ethyl cellosolve acetate decomposed only to an extent of 17,5% at 450°C. How­ ever at 500°C. about 50% of the ester had decomposed.

The products from 80 grams of ester were ethyl vinyl ether, acetaldehyde, acetic acid', and seventeen liters of gases. The. principal components of the gases were

CO, OOg, OH^, and ethylene. The origin of these was

Checked by the pyrolysis of ethyl vinyl ether. Twenty grams of the carefully dried ether yielded approximately ten grams of acetaldehyde and five and a half liters of gas which consisted of CO, CH^, and ethylene only. The carbon dioxide must have, therefore, been contributed -56-

principally by the decomposition of acetic acid.

Methyl cellosolve acetate, 56 grams, decomposed

at 500°C. to yield 37% decomposition products which

consisted of acetaldehyde, methyl vinyl ether, methyl

formate (•), and acetic acid and seven and one-fourth

liters of gases principally CO, CO2 , CH^, and ethylene.

The pyrolysis of ethyl cellosolve propionate (90 grams)

. , o gave almost 65% decomposition at 500 C. The products

were also identified as ethyl vinyl ether, propionic

acid, acealdehyde, CO, COg, methane and ethylene.

The data suggest that the process of decomposi­

tion for the cellosolve esters proceeds essentially

through a cyclic transition state. The decomposition

in comparison with ethyl acetate and n-propyl acetate

is summarized in Table III.

Table III

Decomposition of Cellosolve Esters

*Ethyl acetate 525-550°C.. 63.7-90.1% Decomposition w*n-Propyl acetate 470°C. 49.5% Decomposition Ethyl cellosolve propionate 500°C. 65.0% Ethyl cellosolve acetate 500°C. 50.0% Methyl cellosolve acetate 500°C. 37.0% Ethyl cellosolve acetate 450°C. 17.5%

R*f- * (5)j (4). -57-

The higher percentage of decomposition for propionates is in agreement■with Fugassi and Harrick's finding of the influence of the E group on the central carbon atom in the acid portion of the ester. Similarly, we ob­ served the changing of a methoxy group for ethoxy in­ creased the rate of decomposition. However one must necessarily point out the presence of the oxygen atom has introducted a certain stability in the ester and this may be explained by a bond order change in the 2 linkage, the carbon atom changing to a sp - 3 bonding from a sp hybrization. This would increase the bond angle up to 120° if the transition is complete.

However a small distortion of the angle would cause sufficient resistance to the formation of the transi- 4 tion state and hence lessen the decomposition of the ester. Other support for this idea will come later as we consider various groups in the substitution of the a-C atom in the alkoxy-portion of esters.

single sp^ hybrid atomic orbital.

(b) Three sp^ hybrid x orbitals. -.58-

(a) Gross section of a single sp3 hybrid atomic orbital.

(b) Perspective of four sp3 hybrid orbitals.

The substitution of a -COOR group on the a-C atom has been studied by others extensively though not for mechanistic reasons. The 2-acetoxypropionates of some thirty different alcohols have been investigated. In most cases the elimination of acetic acid vas found to be the principal reaction. The stability of these com­ pounds with respect to variation of R- in the general formula below may be observed in Table IV.

R-OOCH-CH3 1 (general formula for acelated lactates) OAc -59-

Table IV

Decomposition of Lactates

R- Temp.°C. % % yield Decomn. Acrylate Ref

Phenyl 550 100 66 12b Methoxy-Et. 496-503 46-90 21.9-23.3 12b Ethoxy-Et. 490-500 42-87 16.8-18.0 12b Butoxy-Et. 490 52.3 17.6 12b tetrahydro- furfuryl 495 55.7 21.6 12d Ethyl 500 43.0 19.0 12d n-Propyl 500 27.0 25.0 12d Isopropyl 500 99.0 0.30 ‘12 d Isobutyl 500 31.0 40.0 12d n-Butyl 500 61.0 27.0 12d 2-Et~butyl 500 57.0 31.0 12d 2-Et-hexyl 500 38.0 20.0 12d Cyclohexyl 500 96.0 0 l2d p-Chloroethyl 500 46.0 51.0 12d £-Chloroallyl 52 5 68. 0 29.0 12e Crotyl 500 85.0 4.0 12e Methallyl 525 87.0 ' 6.0 12e p-Chlorophenyl 522 55.0 77.0 12 j M-Tolyl 550 92.0 59.0 12 j o-allyl- phenyl 499 33.0 76.0 12 j 555 100.0 52.0 12 3 p-t-Butyl- phenyl 549 85.0 80.0 j23 p-t-Amyl- phenyl 584 4 6 . 0 31.0 12 3 p-Cyclohexyl 550 90.0 71.0 i2j Allyl 546 71.0 43.0 12a Methallyl 545 75.0 41.0 12a Methyl 478 — .i 63.4 12a 475-485 80.0 88. 0 12a O-Tolyl 547 100.0 67.0 12h Methyla ctate 500 47.0 14.0 12h Methylglycolate 500 "48.0 54.0 12h Ethyl 450 20.0 76.0 10 500 69.0 10 Butyl 560-580 89.0 18.0 10 Benzyl 530 82.0 75.0 10 -60r-

The isobutyrates may also be tabulated according to the general formula below.

OR, i 3 . CH. -C'-C-O-R (General formula) 3 H

Table V.

Decomposition of Isobutyrates

R_ Temp. °C % % Yield Meth- Decomp. acrylate Refi

Methyl 4-80-500 95.0 92.0 10 Methoxy-Et. 450 100.0 76.0 10 Methylglyco- late 501 100.0 81.0 12h Methy1- lactate 500 100.0 56.0 12h Ethyl- lactate 502 100.0 26.0 12h Allyl-lactate 446 79.0 82.0 12i Methallyl- lactate 450 63.0 85.0 12 i Allylglyco- late 450 69.0 78.0 12i Phenyl 450 80.0 90.0 12 j Ethyl 475 78.7 71.0 12 d Methallyl 500 --- 73.0 12a

The lesser stability of the isobutyrates compared with the propionates shows that an additional methyl group does statistically aid the formation of transition states as was observed in the case of isopropyl acetate compared with ethyl acetate. The influence of the

-COOR-group on the transition state is not clear.

However in ethyl a-acetoxy-propionate (XLIl) -61-

(XLII)

ring "A" in this case is more stable than ring "Bu, since at 500°C. "A" decomposes to 43$ and "B" 56$.

Comparing (XLIl) with ethyl acetate which gave 63.7$ decomposition at 525°C. it is hard to deduce whether ring "A" has exerted influence on C* atom or that ring

"B" has exerted its influence on the a-C atom of ring

»A»

In the case of (XLIII)

(XLIII)

ring UAU at 503°C. was decomposed to an extent'of 90$ whereas ring "B" was decomposed to an extent of 74$; and the ethoxy-compound showed ring "A" gave S7$ de­

composition at 500°C. and ring "B" yielded 79$ decompo­

sition. These facts are in good agreement with our

cellosolve acetate data presented before, that the thermal

stability of. an ester is increased by the ^-substitution

of an alkoxy group on the alcohol portion. Let us now

re-examine the alkyl substitution effect by this type

of ring analysis. -62-

Table VI.

Decomposition of Esters

R(alkyl) on Ring "A" Ring "B" Temp. Ring "B" {% Decomp.) (% Decomp,) °C. Refs.

Ethyl 43 56 500 12d n-Propyl 27 7.4 500 it Isopropyl 99 99.7 500 ti Isobutyl 31 0 500 11 n-Butyl 61 55.8 500 11 2-Ethyl-butyl 57 45.7 500 11 2-Ethyl-hexyl 38 47.3 500 11 {3-Chloroethyl 46 0 500 11 Cyclohexyl 96 100 500 11 Crotyl 85 95 500 12e

Table VI shows that, in general, any substitution on the C*# atom will stabilize the ester against heat, and any substitution of the a-C will result in a lowering of the activation energy for the transition state pro­ vided the substituent is allcyl or -COOR. The effect of other substituents cannot be inferred until more data are available. One further point must be made here.

The decomposition of ring nBH is .enhanced by the acrylic acid resonance, therefore, one should expect this decomposition favored. Since the variation of R (alkyl) does change the stability, the true alkyl effect on the

C**# must be even stronger than observed by the double ring analysis. -63-

3. The Excited Triplet State as a Transition State

in a Special Case.

Having qualitatively differentiated the effects of substitution on the a-C and C** atoms we may now observe the influence on the transition state by a CH2 = or CH2 =CH_ group on the a-C atom.

Table VII

Decomposition of Some Unsaturated Esters

Compounds Temp. $ Yield of °C. Decomp. Dilcetones Product Refs,

Enol benzoate of 500 8 88 Caproylacto- 77 methyl,n-amyl- phenone 10$ lcetone n-butyl benzoyl- acetone 60$

2-acetoxy 500 ? 85 acetylace- 77 propene-1 tone 78

Enol acetate of 500 71 64 propionylace- 77 methyl ethyl tone 65$; ketone metl^a. acety- lacetone 35$

Enol acetate of 500 38 33 isovalerylace- 77 isobutyl methyl tone 15%', iso- ketone propylacety- lacetone 25$ a-acetoxystyrene 500 77 60.8 benzoylacetone 77

Enol benzoate 500 60 45 benzoylacetone 77 of acetone

Ehol acetate of 500 61 — benzoylacetone 77 a cetophenone

4-acetoxy- 450 66* dienes 9 hexene-2

2-isopropyl, 500 — 23* dienes 79 3-acetoxy- butene-1 -64-t

Table VII, cont

Compounds Temp. % Yield of °C. Decomp. Diketones Product Refs

5-acetoxy-penten-l 560 100 60* divinyl methane BO

2-propenyl- - 500 61.6 5-ethyl,2,4- 77 ethyl hexanoate nonandione

#No diketone

These twelve compounds in Table VII have been care­ fully selected from a great number of unsaturated esters.

The principal idea is to select compounds in which the

double bond is close enough to the a-C atom so that its

effect.may be observed in the decomposition. In general,

a RCH= substitution on the a-C atom causes a change of

reaction mechanism. This is explicable on the basis that

the cyclic transition state is not then permissible.

The a-C atom in all these RCH= cases is definitely not

tetrahedral but rather sp*, Since the a-C is the pivot

point in the cyclic transition state an enlargement of,

the angle would mean a wide separation of the H atom from

the C=0 group. Thus the transition cannot be allowed

through a cyclic configuration. On the other hand the

activation of such a molecule followed by the formation

of products may require even less energy than would be

required with a ring. It may be suggested that the acti­

vation of a double or triple bond may be the transition -65- state in these cases. Such a transition complex may wel-^ be of the type similar to a cis- and trans-isomerization reaction. The triplet state employed in the isomeriza­ tion was indicated earlier in the discussion of kinetics but it must be stipulated that this transition state in­ volving an excitation of an unsaturated bond may be some­ what more than a mere unpairing of electrons or "a change of electron multiplicity”. In the work of R. D. Brown (81) he stated that cis-trans-isomerizations possessed a transition state which in effect caused a localization of electrons. Because of this localizing effect the planar model of the molecule now transforms into a folded model with the two planes at 90° angles. This was explained by k, D. Walsh (82) as the promotion of an electron from the bonding orbital to an anti-bonding orbital. In this process the rigidity of the double bond is lost and a most stable position is assumed by the two carbon atoms o at an angle of 90 to each other (83). From a practical point of view this means that a diradical (in the extreme sense of the excited triplet state) has been formed which in its very short life span undergoes a reaction and gives products principally of the rearrangement t?/pe. We may postulate a tra nsition state in the case of 2-acetoxy- propylene as follows: -66-

/

This essentially is comparable to the four center planar configuration of Barton’s alkyl halide pyrolysis, how­ ever one would expect the activation energy to be even lower since only the movement of one electron is required to enter the transition state complex after the initial unpairing. The energetically favored 0=0 formation offers additional driving force also. The mixture of products in the cases observed in the unsaturated ester decompositions may also be explained by this transition state complex. , C.U 3 Cc>-C-vV- -C.OCH, p.

V UK.-e II

Resonance Contributing Forms of the Transition State.Complex

The application of these ideas will be made to the data

obtained on cyano-esters studied by this author and also -67- by others. One interesting coincidence may be stressed.

O'Connor and Nace's speculation of a cyclic free radical mechanism is indeed very pertinent now.

In discussing the cyano-group effect we -intend to

analyze in detail the results of the decomposition of

diacetyl cyanide (XLIV). The stoichiometry for the normal mode of decomposition of this ester follows the reaction:

(1) GHoCO.0*C - CH3 ----> CHoCO*OH + CH0 = G 3 nCIM 3 . 2 "CN (XLIV) (VCW)

However it was noticed that small amounts of pyruvic

nitrile, diacetyl, CO, and acetonitrile were present.

This led us to believe that perhaps route (2) is also

operative. .ON vo (2) CHoCO* 0* G - GHoJHo ---- 2 CHoCOCNCHoCOCN -----“ > CHoCN J kt ' J +oo ‘CBT

The presence of HCN has been shown to be primarily due

to the lack of purity of the diactyl cyanide. However,

even pure starting materials give small amounts of IICH

under high temperature pyrolysis. This may be due to a

free radical mode of decomposition of diacetyl cyanide.

The extent of this reaction can be controlled as the data

present in graphs I and II. GrtPts 30 3)0 C«> 0 1 4 A. Rctoi f f t if o t c R . .C .A D OH .O A 9 4 % D - 48- « 9 IT l C * l t 310 G** ph x 450 (.lTtK.3 /«* Hoo*. to so 0-4 CO. &.A.C. to M - A c OH VCM -70-

In an attempt to show that CO and CH^CN were pro­ duced from the pyrolysis of pyruvic nitrile we, therefore, synthesized this material (8 4 ). The pyrolysis of pure o CH^COCH at 400 0. gave almost pure CH^CN and a gas con­ taining only CO. Very small amounts of diacetyl were also obtained (about 1%). These results are very similar to those of Szwarc and Murowski on the pyrolysis of CH^COBr.

Our efforts in establishing the presence of cyanogen were fruitless but Szwarc similarly did not identify any free bromine molecules.

The temperature effect in this ester was studied with respect to the per cent of decomposition. As we would predict according to the cyclic transition state mechanism the resistance to heat must be very much lowered by the disubstitution on the a-C atom of the ring of the transi­ tion state. This was shown to be true, as at 370°C. there was 14.8$ decomposition; at 390°C, 70.0%; at 415°C., 85%; and at 450°C., 100%. With the increase in temperature there was an initial increase in olefin formation, but a slight decrease was found between 390°C. and 500°C. in the production of olefins.

The effect of contact time was complicated by the addition of inert gas. In general an increase of inert gas flow increased the rate of decomposition. An examina­ tion of the physical parameters showed that the expected effects due to the increase of carrier gas flow are as -71- follows:

(a) The pressure increase in the system due to the increase of nitrogen flow rate ranged from 5 to 10 mm. of mercury. This is approximately a 1-5% change in the overall pressure. The temperature decrease expected was not observed (p../oo). This may be interpreted as due to efficient instrumental compensation. The decrease of ester.concentration in the system due to the increase of nitrogen pressure, in the extreme case, introduced a factor not larger than ten. The increase of olefin production is considerably short of the observed increase in the degree of decomposition with shorter contact time. (see

Table VIII)

(b) Greater detail of the control of experimental conditions can be found in the experimental section.

However, it may well be emphasized here, the temperature and pressure changes have been conditioned to give the minimum influence to the rate of pyrolysis at a chosen contact time.

(5) The exclusion of the above mentioned variables in the consideration of increased rate of decomposition at shorter contact time suggests a quick efficient heat equilibration when inert gas flow is enhanced.

The following general equation for the diacetyl cyanide (D.A.C.) pyrolysis may be anticipated from the above data and the considerations of ester pyrolysis -72- already made. A detailed discussion x^ill follow 011 the product formation and the mode of decomposition of D.A.C. under the influence of various inert gases and different contact times.

+ CfJ tu tkc = I Cll3 |c. t £> t; ■MG | II * ■+.. 1 0 eoJ (y l w ) ctt3 Acoif

t k

Excited triplet state

From the preceding equation the rate of D.A.C'. decompo­ sition to the right is taken as represented by but the reaction to the left is made more complex by'the collisional deactivation k2 rate. We shall at present consider lc^ as a net rate which, depending on the rela­ tive rates of k-^ to lc^, can be estimated either as k-^ or k 'iKe" At a given temperature and flow rate, lcT .and lc^ will have a ratio which will depend on the specific rates of both of the competing reactions. When a faster flow of inert gas is introduced the rati© of must necessarily change. Let us now study Table VIII. -73-

Table VIII

Pyrolysis of Diacetyl Cyanide at 46ol5°C.

and at Constant Rate of Ester Addition.

Exp. NLRa te % % Conv. to . % Conv. to No. (l./hr.) Decomp. Olefin (kip) Monomer (k-^) lcl/k$ 17 17 '53.33 54.66 24.08 0.440 20 27 60.00 55.98 24.41 0.436 19 30 56.66 54.40 25.84 0.475 21 43.5 66. 66 67.28 21.81 0.324 22 60 91.66 76.44 30.60 0.400 24 - 47(0H) 58.33 62.49 21.88 0.350 25 22.4(Cl/) 53.33 68.33 16.89 0.247 26 50(C02) . 61.66 72.72 16.81 0.231

It can be noted that an increase in the nitrogen input enhances both and k^. However, the increase in the krj> is much greater and the k^/lc^ has a negative slope. This phenomenon is just the opposite of the temperature- effect, in which case d(k]_/k^)/dT is positive.

The effect of the inert gas varies with its nature.

Thus at 50 l./hr. for nitrogen the k^/k^ value is = 0.425 while for carbon dioxide it is = 0 .231; at 22,4 l/hr., k1/kT for nitrogen = O.46 , for chlor.obenzene = 0.247; at

47 l./hr., k^/k^, for nitrogen = 0.43 and for benzene = 0.35.

These data show clearly that the structure of the inert gas is very important and must, therefore, be connected with activation. The main effect of these collisions is to be found in the variation of kT . Thus we can postulate that the k^ rate is dependent on the speed and ease of transfer of energy to the ester molecule so that redis­ tribution can allow the transition complex to form. Its -74- passage over the•potential energy barrier is probably not the rate controlling step. The influence of the nature of the colliding species on the transfer of energy (85) is dependent also on the nature of the ester molecule.

This is expressed by a transmission coefficient. When this constant is nearly unity the transfer of energy from collision is virtually 100 per cent effective. Hence under these conditions a "favorableT, colliding molecule can transmit all its energy to the receiving molecule which on redistribution can enter into the transition state by a single collision. Such may happen in the reaction following the ky route in D.A.C. pyrolysis.

On the other hand the'transfer of energy may be hindered greatly if the activation depended on a localized collision to promote an excitation of one or two bonds.

Because of this geometric consideration the transmission coefficient is considered low and thus the specificity of the type of energy enrichment will not be found in these reactions. The controlling factor here for an increase of population in the excited state can only be a statis­ tical contribution. Thus in the substitution of nitrogen for ester molecules the effect on k^ is only an increase in the frequency of collisions due to the larger trans­ lation energy possessed by the Ng molecules as compared to the ester molecules at a given temperature. This idea well explains the finding of high decomposition .data at high -75- temperatures and the small decrease of k^ rate at the higher temperatures of pyrolysis. The equilibrium postulated (page 72) in the transition to the excited triplet state is in good accord with this explanation of the temperature effect.

The increase in flow rate must also have its limi­ tations in the enhancement of certain rates. In order to show the limitation and scope of this effect we per­

formed the following test. By fixing the volume concen­ tration of ester at a constant value (2.23 x 10“^ g/cc) in each experiment the ki/kj ratio was determined at different ratios of flow. These values are given in

Table IX.

Table IX

Inert Gas Effect on the Pvrolvsis

of Diacetvl Cyanide

Exp. Temp. Benzene Ester % fa No. °C. Rate Rate Consump. Conv. k^ k-^/krp cc/min g/m (kj)

33 45515 560 0.75 93.33 72.4 25.75 .3556

34 45515 1866 2.5 68.33 61.51 26.30 .4275

35 45515 320 0.43 98.33 79.83 19.83 .2484

It is apparent that the k^/k^ ratio increase with the increased rate of input but is approaching a maximum value. Since the decrease of the consumption is per­ ceptibly faster than the increase of k^/bp, its maximum -76- value can at best be betx^een 0.5 to 0.6 before the per cent of decomposition becomes immeasurable. This is consistent with our general equation for the decomposi­ tion of D.A.C. (p. 72) Since the increase of the ratio of lc^/kj is explicable by either a decrease of k^ or an increase of k^ the limitation may be more reasonably ex-* pected from a change of lc-j_ to 1c* 2.Ke • Therefore at the limit the only variable of the ratio will be k^ which was shown previously to be dependent primarily on the type of colliding species. Since the concentration of ester in the above Table IX is fixed, a definite limit must exist in the k^/ky variations.

•Note: kj and km denote not only the reaction rate constant but also will be employed as a term to differenti­ ate the types of reaction in the following section.

The established relationship of k^ and krp can now be used for other reactions of this type. In separating k^ from k^ we emphasized each specific reaction. In reality the magnitude of the ratio of k]_ to krj, is more connected with physical conditions. Just as input of an inert gas can enhance k^ and kj or selectively more of the one than the other, in other systems k-^ and k^ may go in opposite directions and ultimately favor one reaction with the complete suppression of the other. To illustrate this we may observe a series of compounds whose thermal decompositions are channeled into two paths k^ and k^. -77-

Table X gives the values of the ratio k^/krp.

Table X

Decomposition of Cvano-■E sters

Compounds

1. a-Cyanoethyl 540-550 95 (ca.) 2.0 36+3 0.02 10,12a acetate

2. 1-acetoxy-l- 340-350 80 (oa.) 2.2 11 .025 10 cyanocyclo- 440-450 81 . 8.7 46 .407 hexene 540-550 80 20 100 .250

3. 2-cyano-iso- 425-435 96 3.5 87 10 propylacetate 550-560 17 100

4* p-cyanoethyl acetate 540-560 63 33 88 10

5. 4-isopropyl- 575-600 74 26 100 .351 86 1-cyanocy- 500 60 39.6 69 .660 clohexylace- 565-575 62 38 100 .611 tate

6. a-Cyanocrotyl 400-600 — l-Cyano-1,3-butadiene 87 acetate

7.7a-cyano, a-vinyl- ethylacetate 475 27.6 52 80 (ca .) 12a,88

8. a-Cyano, a-vinyl- ethylbenzoate 550 10 8 90 88

In the cases (2) and (5) of Table X the steric effect

is operative to some extent . The ki/kT ratio increases

with temperature showing the same effect as the diacetyl

cyanide (D.A.C.) case. The yield of cy'clohexanone was

not reported. However this product cannot be predominant

in the k-p route of decomposition since the steric factor

is against forming an exo-unsaturation. This may explain -78- the HCN formation. In the cases (l), (3) and (4 ) the analogy with diacetyl cyanide is closer. In compound (l) the simple a-CH-group has little effect on the a-C atom since there is little resonance stablization; there is even less in the case of D.A.C. Therefore k^/k^ should he much smaller than one, as is the case.

/ N = not V allowable. Oftc

Table X also shows that reaction k^ is, on the whole, rare and the rate of decomposition in general is very similar to that of ethyl acetate. In compound (4) the dis­ tortion of the p-carbon atom electron hybridization is not as effective as that of the a-C would be. The impeding effect for the formation of the cyclic transition complex, therefore, comes only from the lack of a p-hydrogen. Hence k^/kj should be less than one. However the results given are not very suitable for our purpose since a composite figure from several experiments was quoted by the author.

In (3), the rearrangement reaction was found to be present to some extent at high temperatures. This may be explained, comparing with (l), by the two methyl groups with the cyano- group sterically affecting the acetate group. But the relatively low temperqture experiments showed small amounts of rearrangement products as compared to diacetyl cyanide. -79-

In the cases (6), (7) and (8) the lack of p-H atoms and the significant modification of the a-carbon electron configuration make the normal reaction, krj., most impro­ bable, hence k-^/krp 1. Even with the benzoate, enhance­ ment of normal elimination, k^ still dominates product formation. The mechanism for (6) may be represented as follows:

' " r ' / Ac Ac C.U'-C.H - and for (7)

u oft‘

Both systems are hard to recognize with respect to normal eliminations and rearrangements. Exact analysis of products in these cases is of the utmost importance since the forma­ tion of an olefin is the result of both k^ and ky reactions.

We have now discussed the extremes of the lc^ and k^ reactions i.e. those compounds which decompose by purely normal eliminations and those by pure rearrangement. We have also observed examples x^hich simultaneously decompose by both k^ and k,p reactions. There is still one more type.

This is that resulting in rapid consecutive reaction, i.e., kj followed by k^ or vice versa. The compounds that best examplify this kind of pyrolysis are di-esters. Let us observe the esters in thermal decomposition given in Table XI. -80-

Table XI

Decomposition of Some Di-Eaters

Temp. % De- Compound Ref. ■ °0. comp. Products

1. Ethylidene 89 285 100 Aldehyde Diacetate Anhydride 2a. Trimethylene 90 565 ? Allyl acetate (50%) 1,3-diacetate (based on total decomp.) 2b. 1,3-diacetoxy 91 575 9 1,4-pentadiene(7,6%) pentane 600 Cis-trans piper- lene (7456) 3a. 1,4-diacetoxy 92 600 ? Total olefin (86%) hexane 1.3-diene (31%) 1.4-diene (28$S) 2.4-diene (24%)

3b. l,4£diacetoxy- - 93 595- 77.3 1,3-diene (54%) 2-isopropyl 605 butane

3c. 1,4-diacetoxy- 94 600 ? 1,3-diene (18%) 2-methyl-3-ethyl butane

4a. 1,5-diacetoxy 95 575i5 ? 1,4-diene (96%) pentane 96 97

4b. 1,5-diacetoxy 92 560 ? Total olefin (92%) hexane 1,4-diene (65%) 1.3-diene (12%) 2.4-diene (12%) 1,3-pentadiene(4%)

5a. 2 ,3-diacetoxy- 88 595 100 Butadiene (85.4%) butane E thy lme thy lice tone (11.2%) 2-ethoxy butene-2 Methyl acetyl acetone

5 U 1,2-diacetoxy- 95 .590+5 9 Diene (65%) pentane Aldehyde present

5c. 1,2-diacetoxy- 90 575 ? . Allyl acetate (16.1%) propane Aldehyde present

6, 7 and 8 are recorded in Table XIV -81-

In the case of etliylidene diacetate (l) no analysis of products was given, although kinetic measurements are available and data were offered in the cases of ethyli- dene dipropionate and butylidene diacetate. There can be no assurance that the increase of pressure (the authors followed the reaction by pressure measurements) is an affirmation of aldehyde and anhydride formation. An olefin and acid formation will also give similar pressure increases. Cases (2a) and (2b) are somewhat useful but the lack of the exact data on amount of products makes analysis awkward. The yield of allyl acetate gives good indication of a normal elimination, however, the scarcity of detailed data renders it impossible to identify con­ secutive reactions. One would expect some production of methyl ethyl ketone, but this was not reported. Case

(2b) is promising. The large production of cis-trans piperylene shows clearly that both 2-acetoxypentene-3 and 2-acetoxy-pentene-1 could have been formed as an initial step. These two isomeric pentene acetates will favor decomposition into conjugated diolefins. Hence the small yield of 1,4-pentadiene produced.

As the acetate groups are further separated from each other by methylene groups the effect of a double bond, in the intermediate ester in the pyrolysis, no longer will contribute to the k^_ route, hence, a consecutive reaction if present, will be difficult to follow both kinetically I

Page(s) 2 *L> missing in number only UNIVERSITY MICROFILMS -33- and chemically. This is amply shown in the cases (3a),

(3b), (3c) and(4a) and (4b). No discussion can be made about these compounds, since the yields of olefin are uncertain and the per cent of decomposition were not re­ ported. Ho kinetic data were offered in these cases nor was a careful examination of products given. However,

some qualitative information can be obtained from compounds

(5a), (5b) and (5c). Compound (5a) at almost total de­

composition gave 85.4$ butadiene and 11.2% ethyl methyl ketone. Since this ketone is most probably formed by the following scheme i

C H , - C I V - < U - C. H j \iT ^ cH,-CH-C4|-=t4l, J f j . /c U - c H - ) i * — — > 3 . ^ V- 1 'z

y £ . r \ c w

Act.0 + CM 3 CII--CH.-CH3 \ c-c.H- 3 (*U*J 111) ofk. AcO 1 \.U) (Uni)

lV ' . &

U ) 0 0 -cH, c#,co-c«,. 0-“) J V " 0

one must necessarily conclude that k^ yields better than

85.4$ of a-vinyl ethyl acetate. The conversion of (XLV)

to (XLVII) is very much favored as can be seen from the

previous consideration of the decomposition of a-cyano-

crotyl acetate. Thus the rate ^k^ cannot be unduly

complexed by any other competitive reactions. In the con­

sideration of k^ the production of ethyl methyl ketone may well be taken as the rate. The lack of anhydride is -84- understandable, since the temperature used is high enough for its complete decomposition. The other competitive r reaction by kj nrnst be very small, since an oxygen .... ' - deactivating effect is present and the hydrogen atom in­ volved is tertiary. For a qualitative estimate let us

assume k^ to be 86.3$, based on the possibility that

2 fop is about 99$ efficient. Therefore k|, by difference must be 2.5$. Since 2-acetoxy butene-2 (LI) was detected,

the isomeric diketone formation must favor k^ as suggested

earlier by the discussion on (H C=) or (H C=CH~) sub- 2 2 stitution on the a-C atom. Hence the yield of (LIV) is

probably 2$, and the remainder 0.5$ is (XII). The im-

portance of the fop rate need not be over emphasized.

The decomposition of (5b) and (5c) showed the presence

of aldehydes. Since the analysis of products is not very

good we can only note the yield of allyl acetate is very

low and anticipate the consecutive reaction to be very

efficient and its rate very rapid.

In the last three cases to be discussed we are able

to explain with more certainty. It becomes apparent that

the adjacent diacetates will proceed by both Ic^ and lc^

routes with comparable ease. However, when the acid por­

tion of the molecule is modified to enhance the normal

elimination, the lc^ reaction will decrease accordingly a'S

a calculation of k-p/foj will show: (k-p/kp) of diacetate =

0.550, (kq/fop) of dipropionate = 0.410. Let us now -85- examine more quantitatively the pyrolysis of glycol di­ acetate. The comparative stability toward heat may be understood by the deactivating effect of the neighboring

oxygen atom. Because of this, the decomposition of 57.54$ o at 500 C. has shown some enhancement of the k-^ reaction in preference to normal elimination. The yield of 64.48$

of vinyl acetate is maximum if we discount any consecu­ tive decomposition present. However, the existence of acetylene in the gaseous fraction shows this is probably low. On the other hand the production of 35.47$ acetr? aldehyde also is not maximum. Methane and carbon monoxide

are present. Acetic acid was shown to be formed to the

extent of 78.7$. This together with the amount of carbon

dioxide in the gas fraction gives 15.3 g. of acetic acid.

Since the acetic anhydride yield is only 13.6$ about 3.8 g.

of the acetic acid must have- been contributed by the following reaction.

0° GH — G 3 ^ CH3COOH + ch2=c=o CH- - (/ 0 Thus 11.5 g. of acetic acid have been produced by the normal elimination. This would correspond to a 66.54$ yield. The difference between the observed yield of vinyl acetate and this calculated yield is 2$. This can be at­ tributed to the consecutive reaction according to the following: - 86- kf ki (-CH2-)2 (-QAo)2 -----* CH2 =CH-0A.c #HC3 CH +A c OH

#The pyrolysis of vinyl acetate gives essentially, acetylene and acetic acid. (C. L. V/ilson unpub­ lished data.)

This calculation is confirmed by the presence of propy­ lene and ethylene from the decomposition of ketene. The

carbon monoxide contribution from ketene is equivalent

to 1.4 liters, hence 2,3 liters of gases must consist

of unsaturated hydrocarbons, methane and carbon monoxide.

Gas analysis (from 73 g. diacetate) indicated an amount

of^0.3 liters of unsaturated hydrocarbon gas, hence 2

liters remain which can be attributed to a consecutive

reaction from the competitive k-^ route. If acetaldehyde

is taken to give equal amounts of methane and carbon

monoxide, then the secondary decomposition of 'acetalde­

hyde is equivalent to 0.2 g. which is about 4.7$ of the

aldehyde produced by k-^. We may now summarize the overall

reaction of the decomposition of glycol diacetate by the

following diagram. A^Cj- c.^-c.i k -OA-c.

C.UX= C U- o /W.

^‘Calculate values are based on the % of decomposition and $ AcOH, vinylacetate, CO, CH^ and CH2=CH2 found. -87-

The glycol dipropionate is essentially different in its mode of decomposition. The enhanced reaction not only causes more decomposition (68.60$), hut also suppresses the lc^ reaction. The severe shortage of propionic anhydride may be due to three factors, (l) the k^ route being suppressed, (2) the comparative ease of decomposition of propionic anhydride, (3) a steric factor preventing the formation of anhydrides. The decrease in the ratio of k^/kg? was not explained by an increase in the production of'vinyl propionate. The consecutive reaction k£, must therefore, be more efficient with the propionates than that with the acetates. A calculation of the propionic acid formed shows there is 2.25 times as much acid as vinyl gster formed. Since the aldehyde yield is extremely low and the amount of anhydride is barely detectable, the calculated 28.15$ k-^ is much too high. The significance of knowing the k,j, rate is evi­ dent. Hence this illustrates well the presence of con­ secutive reactions. An approximation of k£ is of the order 80$.

The information on the pyrolysis of glycol diformates is not suitable for discussion. The formates are so very different in behavior, they are to be regarded as ab­ normal from the ester. For a survey on the modes of de­ composition of formates one can find a good presentation in Hurd’s classic work, wThe Ryrolysis of Organic

Compounds'1, ACS Monograph. -88-

EXPERIMENTAL

1. Pyrolysis of 1,1 Dicyano Ethyl Acetate (Diacetyl

Cyanide).

A . Reagents:

Diacetylcyanide (D.A.C.) obtained from the B. F.

Goodrich Research Department, was purified by fractional crystallization: In a 500 ml. beaker the crude diacetyl­ cyanide (D.A.C.) was melted in an oil bath thermally con­ trolled by a resistance heater and regulated at 60-65°C. by the use of a variac. After the D.A.C. was completely molten the beaker was covered with an evaporating dish and elevated so that an inch of the liquid level was above the oil level in the bath. After 12 hours a big crystalline cylinder was formed. This was lifted out of the beaker by gently warming the sides of the beaker with a flame. The purified D.A.C. from several operations was combined and recrystallized in a similar manner. After three recrystal­ lizations the product was obtained m.p.; 70.05°C. This is comparable to the very pure sample of D.A.C. obtained from the Goodrich Research Department} m.p. 70-71°C. This method was compared with sublimation of D.A.C. which yields very pure material in small crystals, m.p. 70.5°C. Infra­ red spectra of the products were taken.

U.S.P. carbon tetrachloride was used for the separa­ tion of pyrolysis products.

One molar silver nitrate solution was used for -89- hydrogen cyanide gas absorption, and standard base for acetic acid titrations (phenol red as indicator).

Other solvents: benzene, chlorobenzene, acetic acid, acetic anhydride, acetonitrile, diacetyl and acetyl bromide were all EE white label reagents.

Fuming sulfuric acid and 30$ potassium hydroxide were used for gas analysis of unsaturated components and carbon dioxide respectively.

B. Apparatus:

1. Furnace and Related Equipment

(a) Dropping Funnel. (Figure l) A constant head dropping funnel was used to pass the molten material into the pyrolysis tube at a uniform rate. The rate controlling device which consisted of a ground plunger with a wire insert was embodied in a circulatory heating jacket so that the amount of D.A.C. remains at a eorstant tempera­

ture. Fig. 1 gives a detailed description of the

apparatus.

The ground surface served as a stopcock. The molten

D.A.C. may be ;shut.off whenever desired by plunging down

the center glass shaft. The stainless steel wire regu­

lated the rate of addition of D.A.C. by the length of the protruding section through the capillary, since the wire is tapered.

For the introduction of liquids a plain constant -qc.-

',T*i‘u/tst & T n t mt/re

Co*tiT*-*T bAoP/'t*& fZuUHGL.

1« CWxJiWf

ffor iuUf-

7w frciriwt. -91- head dropping funnel was used.* This funnel is essen­ tially the same except a stopcock was used instead of the glass plunger to control the flow of fluids. The rate is again regulated by a wire insert.

• (b) Gas Flow. The rate at which nitrogen carrier gas was admitted to the pyrolysis tube was measured with' a U-tube flow meter, filled with dibutyl phthalate and

calibrated vrith a wet test meter. The exit gas was also measured by a similar meter at the end of a train of

cold traps.

When solvents were used, the flow of inert gas was

calculated by the corresponding volume of the rate of

solvents dropped into the furnace.

(c) Furnace:* A specially designed heating device was used. Ho detailed description will be given here.

The temperature was regulated by a Minneapolis-Honeywell

Regulator Company Pyrometer.

(d) Pvrolvsis Tube. The pyrolysis tube consisted I of a pyrex cylinder, 136 cm. long and 4.5 cm. in diameter.

It contained a heated zone 90 cm. long, which was packed with glass chips and which could be maintained within

15° of the desired temperature. The free space in this

zone was 660 cc. A small thermocouple tube extended

through the center of the pyrolysis tube packing.

*Luke: Ph.D. Dissertation, N.D. University, Notre Dame, Indiana. . • -92-

(e) Receiver and Traps. The receiver was a 500 ml. ground joint, round bottom flask with a side arm. The

issuing gas from the furnace was baffled by an inner tube

in the flask for efficient condensation. The receiver

was packed in a dry ice and acetone mixture. A dry ice

trap (i) was placed in sequence followed by a liquid

air trap (II). The effluent from these traps was

measured by a U-tube manometer as described in part (b).

(f) Gas Storage. The condensed gases in the liquid

air trap were stored over a sodium chloride solution,

containing 1% hydrochloric acid, in a fifty liter bottle.

The pressure of this storage was regulated by an open

mercury manometer and an overflow. Aliquot of 200 ccs.

were removed from this tank by displacement of brine at

room temperature and pressure and finally transferred to

the gas analysis apparatus.

C. Decomposition of D.A.C.

With a stream of nitrogen passing through the pyroly­

sis tube the system was permitted to attain the desired

temperature overnight. The temperature readings were taken

at six positions along the tube at the start and end of

each run. No adjustments were made during the experi­

ment. The traps were then placed in position and swept

out with nitrogen.

A weighed amount of D.A.C. was melted in a beaker on

a steam bath, and introduced into the dropping funnel. -93-

The jacket of the funnel was kept at boiling water tem­ perature. The addition of molten D.A.C. was timed with a stop watch and the rate adjusted by the wire in the funnel. The total time for the addition was also recorded.

After the addition of D.A.C. was completed, the sys­ tem was disrupted. The liquid air trap was replaced by a silver nitrate trap. The flow of nitrogen was con­ tinued for one-half hour. Following this period the dry ice trap and receiver were warmed to room temperature and flushed out with nitrogen for the complete precipita­ tion of hydrogen cyanide.

The liquid air trap was connected to the gas storage and expanded by warming slowly. The volume of brine displaced was calibrated to read in liters at one atmos­ phere.

Before another run was made the furnace was cleaned by heating to 530°C. for no less than,30 hours with a

slow stream of air passing through the pyrolysis tube.

For the experiments which contain solvents the pyroly­

sis column was first heated and flushed with air and followed by 30 hours of flushing with nitrogen at a temperature at which the experiment was to be performed.

Prior to the addition of the D.A.C. solution a few cubic

centimeters of pure solvent were dropped through the tube.

At the end of each run nitrogen flushings were employed to remove all the products. -94-

D . Isolation and .Estimation, .of Products

Procedure I..

The products in the receiver were first rinsed three

times with cold distilled water. This was added to the

rinse of the traps and made up to 500 cc. in a volumetric

flask for analysis. Titration of a 25 cc. aliquot with

sodium hydroxide and phenol red indicator was assumed

to measure acetic acid.

The residues of the receiver and traps were boiled

with carbon tetrachloride. Three extractions were made

and each operation was filtered hot to separate the poly­

mer of vinylidene cyanide (VCW). The polymer on drying

was weighed as (VCH) m.p. 250°C. (It decomposed to give

a lachrymatory gas.)

The carbon tetrachloride solution was evaporated to

dryness and D.A.C. was removed. After recrystallizations

from CCl^ the m.p. of 70°C. was obtained.

The silver nitrate trap was carefully rinsed with

very dilute nitric acid solution and the precipitate

filtered and dried in an 80° oven for four hours. The

weight of the Ag CN was used to calculate hydrogen cyanide.

Gas analysis was performed as follows: A 200 ml,

gas sample from the gas storage was taken into a sampler

and this was connected to a Fisher Precision Gas Analyzer.

One hundred milliliters of gas was withdrawn into the

gas buret from the sampler and analyzed. Absorbents used -95-

were 30$ aqueous potassium hydroxide for carbon dioxide,

fuming sulphuric acid for unsaturated hydrocarbons, a

copper oxide tube heated to 300° for CO and a Burrell.

Perma-Therm Catalytic Heater for the combustion of hydro­

carbons .

Procedure II.

The gases condensed in the liquid air trap were ex­

panded into the storage while HCH was collected in the AgNO

scrubber. The products in the receiving flask were care­

fully warmed to room temperature and the gases evolved

were combined vrith those in the gas storage tank. The

liquid was then distilled under reduced pressure (l-5 mm.)

until all of the solvents and volatile products were

collected in a dry ice cooled receiver. This effectively

separated any residual D.A.C. since the tarry residue

washed with CCl^ gave pure D.A.C.

The distillate was rectified through a fractionating

column (about 30 plates). The initial fraction boiling

' between 68 to 85°C., depending on the solvent, contained

diacetyl and acetyl cyanide. ' This portion of the distil­

late was quickly made into 2,4-dinitrophenyl-hydrazone

derivatives according to the method in "Identification of

Organic Compounds" by Shriner and Fuson, John I/iley and

Sons, 194-0. The osazone derivative can be easily separa­

ted from the liydrazone derivative since the former is

only soluble in nitrobenzene while the latter is easily -96- ' dissolved in 95% alcohol. The derivatives were further purified by recrystallization with the above mentioned solvents. The second fraction, b.p. 85-120°, usually contained some solvent along with the monomer. It was weighed and diluted according to the procedure of the

B. F. Goodrich Research Analysis of VNC.# The spectro- graphic data were calculated to give the yields of the adduct of VCN and anthracene from optical densities of known solutions. The residue from the fractionation contained usually a small amount of VCN polymer and some tar. These were extracted again with CCl^. The identi­ fication of the derivatives of these by-products will be discussed under the By-Products section.

The estimation of acetic acid by this method is

only semi-quantitative. Since the monomer is sensitive

to water no total titration of the acidity of the pro­ ducts can be made. Furthermore the acetyl cyanide de­ composes readily with water which would lead to erroneously high yield of HCN. Attempts to take aliquots for water extraction were found useless.

E . Results

1. Variation of Temperature on the Formation of Products.

The starting compound pyrolyzed in these experi­ ments was the purified D.A.C. from Goodrich. (See "A" -

Reagents). Each experiment was performed with 30 grams

Private Communication -97- of D.A.C. with the flow of carrier gas, nitrogen, regulated at 10 to 12 liters per hour. The rate of addition of D.A.C. was 3 drops per 10 seconds. (See

Table XII).

2. Effect of the Rate of ITitrogen Flow on the

Decomposition of D.A.C. at A00°C.

This series of experiments was performed at a con­ stant temperature (400 1 5°C.). The rate of addition of

D.A.C, was held at 10 drops per 10 seconds. The issuing gases were passed through a series of traps including a liquid air trap in order to obtain the condensible gases for analyses. The amount of D.A.C. pyrolyaed per experi­ ment was 30 grams. (See Table XIII and Xllla)

3. Test of Possible Reversal of Fyrolvsis

The data in Tables XIII and Xllla might be explained if the main pyrolysis reaction was in effect an equili­ brium. Attempts were made therefore to add acetic acid to VCN.

Experiment No. 27: - To a constant flow of Ng (10 l/hr.) • acetic acid was introduced at a rate of one drop per second.

After 5 minutes VCN was injected by means of a syringe at the top of the reaction tube (which was kept at 400°C.).

The rate was controlled manually to last 30 minutes.

Experiment No. 28: - Similar to Exp. 27 except rate of

VCN was doubled. Table XII

Products from Pyrolysis of D.A.C.on Glass at Constant N2 Flow (10-12 liter per hour) and Rate of Addition of D.A.C. (18 to 20 drops per minute) (See Graph I)

Exp. Temp. D.A.C. V(CN) AcOH HCN C02 % Conv. on Wo.. ±5°C. (rec^d) gms .. gms gms gms Consump.n * ester De­ kl gms comp.

12 370 23.0 0.5 0.6 0.14 *20.7 14.0 17.8 17.8/14=1.27 9 390 9.0 3.5 4.1 0.84 3.3 70.0 28.0 50.4 50.4/28=1.80 10 415 4.5 3.2 6.00 1.4 — 85.0 21.1 67.1 67.1/211=3.18 8 450 none 3.0 10,00 2.0 0.7 100.0 16.8 83.2 83.2/16.8=4.9 5 11 500 none 3.2 7.80 0.56 100.0 17.9 82.1 82.1/17.9=4.59

* Only 27 grams of D.A.C. pyrolyzed. **k^=$ consumption - ($ Conversion times % consumption) Table XIII

Products from Pyrolysis of D.A.C. on Glass at Constant Temperature 400*5°C. Rate of Addition of D.A.C. (60 drops per minute.) was virtually the same for all runs.

Exp. N? Rate VCN D. A. C. AcOH HCN co2 CO Nnsat. Methane No. l./hr. gms - gms . gms gms gms gms Comp, as gms. H2C-CH2 gms 17 17 5.2 14.0 3.6 0.75 1.95 0.009 0.008 0.51 20 27 6.0 12.0 3.65 0.14 — - — — 19 30 5.5 13.0 4.6 0.14 0.88 none none 0.29 18 37 3.0 13.0 3.25 0.10 1.82 0.7 1.12 21 43.5 8.0 10.0 3.4 8 0.02 0.36 0.000 0.000 0.30 22 60 12.5 2.5 3.00 0.01 0.13 0.000 none 1.31

^Contaminated with CO, value not taken

Table Xllla

Exp. Inert VCN D.A.C*. No. gas Rate gms gms l./hr.

33 *33.5° 12.7 2.0 34 *112 c 7.5 9.5 35 *19.2C 14.0 0.5

24 ' *47 6.5 12.5 25 #22.4 6.5 14.0 26 **50 8.0 11.5

* Benzene as carrier gas$ c=ester concentration at 2.23x10“^ g/cc/eec.; #=Chlorobenzene as carrier gasj ** CO2 as carrier gas. -100-

Experiment No. 29: - Same as in Exp. 28 except no N2

carrier gas.

Experiment No.. 3 0 : - 30 gms, of D.A.C. in 80 cc. of

chlorobenzene were passed through the furnace tube at

400°C. at flow rate of 50 l/hr. The products were

quickly distilled under reduced pressure. The residue

contained 12.5 gms. of D.A.C. The distillate containing

VCN (monomer) were added to a mixture of acetic acid and

V'CN monomer calculated to give 12.5 gms. of D.A.'G. if

the reaction D.A.C, ----■> AcOH + VCN is the only process.

This mixture was passed through the furnace tube at 4-00°

at a rate of 50 l/hr.

None of the 4 experiments yielded any D.A.C. The

apparent difficulty in the Exps. 27, 28 and 29 was the

polymerization of VCN on contact with the hot glass above

the 400°C;. zone. In experiment 30 dilution by chloro­

benzene may make reversibility difficult to recognize.

3a. Effect of Flow Kate on Temperature

The furnace tube may be regarded as composed of 3

zones} top, middle and bottom. Each of these zones was

regulated to 400°C. The sudden introduction of cold gas

gave the following results after one minute.

Exp. 31 Co2 gas 40 l/hr. Top - 400° Middle - 3.97° Bottom - 400°

60 l/hr. Top - 395° Middle - 401° Bottom - 400° -101-

Exp. 32 Benzene (D.A.C.) 80 l/hr. T o d - 4-00° . . Middle - 4.05° Bottom - 4.05°

E. By-Products

1. Diacetyl:

The separation of this compound was mentioned under

the section (D) on Estimation and Separation of Products.

In order to obtain pure d i a c e t y l the benzene solution of

diacetyl, pyruvic nitrile and acetic acid was frozen.

On warming a green yellowish liquid separated from the

solid benzene and acid mixture. This liquid was decanted

and distilled to separate acetyl cyanide. The diacetyl o had b.p. 93-95 C. The osazone was made and infra-red

spectra were taken. The mixed melting point of this deri­

vative with that of an authentic sample of the osazone

of diacetyl varied slightly (m.p. 314.5°C.). The spectra

of the authentic sample, that of the product from pyrolysis

and that which is reported in Ind. Eng. Chera. (Analyti­

cal Edition)(Sept. 1953) all confirm the identity of the

compound. The chemical analysis of C, H and II gave:

c 16h 14°8w8 Calcd.: C, 43.05; H, 3.14; H, 24.10;

0, 28.71. Found: C, 43.05; H, 3.34, 3.62; N, 2 4 .4.8 ,

24.72. (Galbraith Microanalytical Laboratories).

A semiquantitative estimate of the amount of diacetyl

in a 30 gram pyrolysis indicated about 6% by weight.

2. Pvruvic Nitrile:

This component could not be separated easily since it decomposes on contact with moisture and polymerizes on the surface of etched glass to give D.A.C. even in the cold. Even pure acetyl cyanide yields some D.A.C. on standing in a flask in the cold with total absence of light. The only method of separation employed was mentioned under Section D. The 2,4--<3initrophenylhydra- zone derivative of acetyl cyanide is a bright yellow crystalline material. The melting point of the authentic sample of acetylcyanide hydrazone made from reacting acetyl bromide and cuprous cyanide in the presence of acetic anhydride differed by a degree with that found in the pyrolysis product. The mixed melting point gave o 186 which is slightly higher than expected. However, the infra-red spectra of the authentic hydrazone and that of the pyrolysis product gave identical absorption peaks. Chemical analysis for cyanides as silver cyanide gave confirmatory results. Approximately 0.5 g. of acetyl

cyanide was produced calculated from the yield of AgGN.

G.. Preparation and Pyrolysis of Pyruvic Hitrile

1. Synthesis of Pyruvic Hitrile:

To a 1:1 mixture of acetic anhydride and glacial acetic acid was added slowly an equimolecular amount of phosphorous tribromide. When the addition was complete the solution was vigorously stirred and heated until the reflux temperature reached that of acetyl bromide. Re- -103- flux was continued for two hours $ then the solution was allowed to cool. Two phases separate from this solution and the acetyl bromide— acetic acid solution was decanted.

A quick rectification through a column of 50 plates gave

pure acetyl bromide (b.p. 76°C. - 757 mm.).

The above prepared bromide was slowly added to a

boiling mixture of acetic anhydride and cuprous cyanide.

The reflux temperature after long stirring increased to

about 90°C. An additional 30 minutes were given to the

reaction and the products fractionally distilled. Pure

acetyl cyanide boiled at 93°. Some bromocyanogen was

formed.

2. Pyrolysis of Pyruvic Nitrile t

Thirty-four grams of pure acetyl cyanide was passed

through the pyrolysis tube at 4.00°C. and a rate of 1 drop

per second. A semiquantitative estimate of the products

gave about 10 grams of CO and 27 grams of liquid. The

liquid fractional distilled gave acetonitrile and a resi­

due of tarry liquid. The acetonitrile fraction amounted

to about 20 grams. Its identity was confirmed with infra­

red by comparison with an authentic sample. The tarry

residue was reacted with 2,4-dinitrophenylhydrazine.

Both diacetyl osazones and pyruvyl nitrile phenylhydra-

zone were found. Due to the small amount of the diacetyl

derivative it was estimated that the original compound

could not have been more than 1% of the parent compound. Infra-red analysis also showed that the pyrolysis gas was almost pure CO.

3. Attempts to Detect Cyanogen:

Since this gas is readily polymerized a complete vacuum technique was used to keep the cyanogen gas as dilute as possible. For an authentic sample, silver cyanide was prepared and heated in an enclosed system.

The gases evolved were collected in liquid air and ex­ panded at dry ice temperature since the vapor pressure of cyanogen is about 14 mm., at -80°. This gas was stored in a flask and then transferred into the gas analyzer cell of a Parkin-Elmer Infra-red Machine.

Several primary peaks were detected which agreed with those found in Hertzberg "Infra-red and Raman Spectra"

(194^)) but the pyrolysis product gases using the same technique failed to indicate any cyanogen.

2* Pyrolysis of Di-esters

A . Preparation of Materials

Glycol diacetate, glycol dipropionate, and glycol diformate were obtained from Eastman Kodak Company.

These esters were fractionally distilled on a column on approximately thirty theoretical plates. Only the middle third of the distillate was collected and used for pyroly­ sis. The boiling points respectively for the diacetates, dipropionates, and diformates of glycol at 752 mm. of -105- pressure were 190 1 1°, 211 i .5°, and 174 i .5°G.

Those melting points given in Beilstein are: diacetate

190.5°C. (Beil. II- 142), diformate 174°C. (Beil. II- 23), dipropionate 211 - 2°C. (Beil. II- 242). Infra-red spectra of these purified esters are given in section VIII.

In addition to these esters infra-red spectra of acetic acid, acetic anhydride, acetaldehyde 2,4-dinitro- phenylhydrazone derivatives (Form I and II) were taken for comparison in anticipation of products formed. A gas spectrum of acetylene was recorded. All of these and others will be found also in Section V.

B . Procedure for Pvrolvsis

The apparatus used was essentially the same as in the case of diacetylcyanide (Part I). Since liquids were involved the constant head dropping funnel as described by Luke was used. The liquid air trap was replaced with another dry-ice-acetone trap and the exhaust gases were led directly into the gas storage tank. The rate of pyrolysis was controlled by the input of ester and the flow rate of the displaced brine solutions The pressure of the system was kept virtually one atmosphere through­ out the experiment. The total collection of gaseous products was insured from possible nitrogen contamina­ tion due to flushing by discarding the final few hundred milliliters of gases from the pyrolysis. -106-

C. Collection and Identification of Products

The products were classified into gaseous non- conderisable' (with dry ice and acetone misture), condens- ables, and liquids. The non-condensable gaseous substances were totally collected over brine in the storage tank.

No separations were made on this fraction but a total infra-red analysis on an aliquot was performed. In some instances gas estimation by standard gas analysis was made. The condensable gases were trapped in acidified

2,4-~dinitrophenylhydrazine alcoholic solutions. The derivatives formed were isolated and purified. The liquid fraction was generally fractionated by distilla­ tion on a ten-plate column. The natural temperature dif­ ferences of components were used for the fractionation.

In some instances vacuum distillation was applied and washings with sodium carbonate saturated aqueous solutions were applied to the liquids before ether extractions and drying with annydrous sodium sulfate. Table XIV

Products from Pyrolysis of di-esters on Glass at Constant Temperature 500*5°C.

Ester T e m p .. Gms. Rate Non-cond, Conden, Liquids °C. used g./sec gases Gases I Diacetate 500 73 *059 (ca) 4.5 liters c h 3cho Vinyl acetate of Glycol 15% C02 (9g); Ester (41g) 5% unsat. undecomp. CO, c h 4 , ch2=c h 2 -107 500 73 018 (ca ) 4.3 liters CH^CHO Vinyl acetate acetylene ; (10.6g); acetic propylene; (2.5 g j acid (I3.6g)j , ethylene; CO; acetic anhyd.(4.0g) C02»CH4 undec. ester (31g)

Dipropionate 500 86 0.024(ca) 4.5 litSrs CH CHO Vinyl propionate of Glycol acetylene; small, (I5g)j propionic ethylene; CO; amount acid . (25g)s undecomp, C02 ;CH4 ester (27g)

Diformate 500 63 0.013(ca) 4 liters Vinyl Glycol (24g) of Glycol ethylene (?) formate some formic CO; C02 ; H2 (1.5) acid

All liquids are identified by infra-red analysis. The gases are analyzed by infra-red and standard gas analysis. The aldehydes are identified by 2,4-dinitro- phenylhydrazine derivative and infra-red analysis. Mixed melting points with known samples were also taken for the dinitrophenylhydrazones. -108-

3. Pyrolysis of Cellosolve Acetates

A . Preparation, of' Materials:

Crude methyl cellosolve and ethyl cellosolve were esterified with acetic acid and a trace of sulfuric acid at reflux temperatures. The reaction was allowed to proceed oVer a period of 20 to 30 hours. The product was fractionally distilled over a hundred-plate column and at a 10 to 1 reflux ratio. Only the middle third was collected for use. These .samples when compared with the purchased product (EK white label) showed a slightly higher boiling point, with methyl cellosolve o boiling at 136-138 C. and ethyl cellosolve acetate boiling at 15A“156°C. The literature values respectively are

13S°C.(Beil II- 24.1 (66) and 156.3°C. (Beil. II- 141) .

The infra-red spectrograms showed no differences in the commercial products and our own. Both the ethyl cello­ solve propionate and methyl cellosolve propionate were made and purified essentially in the manner described above.

B. Procedure for Pyrolysis:

The method used is essentially the same as in the di-ester pyrolysis, except the flow rate in some cases was changed. In general the 5£0°C. temperature was maintained. -109-

C. Collection and Identification of Products

The same method of infra-red analysis for gaseous and liquid fractions was employed as in the di-ester case. The accuracy of these analyses is within *5%.

TABLE XV

Decomposition of Cellosolve Acetates

Compound Temp. % Gaseous Liquid °C. Decomp. Products Products

Me Cell. 500° 37. 7.2.5 liters Me vinyl ether Acetate gases: COjCOg methyl acetate (56g/30min.) CEgCH-g acetaldehyde

Et Cell, 450° 17.5 4.1 liters Et vinyl ether acetate CO, CH^, Ethyl acetate (80g/42min.) Ethylene acetic acid

(80g/46min.) 500° 50.0 17 liters same same prod

Et Cell. 500° 65.0 18 liters Pr vinyl ether Propionate CO, methane propionic acid (90g/60min.) . . propylene Pr acetate

Ethyl vinyl 400° 100.0 5.5 liters 10 g. acetald. ether ethylene (20g/10min.) methane, CO INFRA-RED SPECTROGRA MS PERCENT TRANSMITTANCE 40 3 002500 3000 5 4 WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE AE EGH N IRN WV LNT I MICRONS IN LENGTH WAVE MICRONS IN LENGTH WAVE 2000 6 50 40 1300 1400 1500 7 10 8 1200 9 1100 1000 II2 900 WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE 12 800 E H E U H T S 31 15 14 13 700 625 16 100

IITTANCE WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 5000 4000 3000 2500 2000 1500 1400 1300 1200 1100 1000 900 800 625 Y&tf .LwtU> I I 100

p i r o u s i s PRODUCT

ETHILEHE DIFORMATE

FRACTION I 20

WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS

WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 5000 4000 3000 2500 2000 1500 1400 1300 1200 1100 1000 900 800 700 625 100

PIROLTSIS PRODUCT 5 40 of 4 0 1 STHTiare DIFORMATE

FRACTION II

WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS PERCENT TRANSMITTANCE -II3- PERCENT TRANSMITTANCE 100 00 4000 5000 00 4000 5000 2 3 002500 3000 0010 10 13002500 1400 1500 3000 4 WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE AE EGH N MICRONS IN LENGTH WAVE AE EGH N MICRONS IN LENGTH WAVE 2000 2000 6 5 50 40 1300 1400 1500 7 8 1200 1200 9 1100 1100 1000 1000 10 900 900 WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE AE EGH N MICRONS IN LENGTH WAVE AE EGH N MICRONS IN LENGTH WAVE 2 13 12 800 800 TOTI PRODUCT PTROLTSIS O F ETHTL3IE F O FRACTION I FRACTION DIACETATE DIACETATE

14 700 700 516 15 625 625 100 100 *5 PERCENT TRANSMITTANCE -Ilf- PERCENT TRANSMITTANCE 100 20 40 60 00. 00 00 50 2000 2500 3000 4000 . 5000 00 00 00 50 00 50 40 30 20 10 10 90 0 700 800 900 1000 11.00 1200 1300 1400 1500 2000 2500 3000 4000 5000 'I I I II I I I,, I I I. II II I I II I I I I I I I II I I - I I , 1 , 1 . I . I , I ., I , I . I . II . I . I . II . I ,I I . I , I , I 6 7 6 5 4 AE UBR I C ' WV NMES N CM-' IN NUMBERS WAVE - -' CM IN NUMBERS WAVE WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE AE EGH N MICRONS IN LENGTH WAVE AE BGH N MICRONS IN LB4GTH WAVE 6 7 6 5 i m 50 40 30 20 O 10 90 800 900 1000 MOO 1200 1300 1400 1500 WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE AE EGH N MICRONS IN LENGTH WAVE AE EGH N MICRONS IN LENGTH WAVE 2 13 12 2 3 14 13 12 l i M IOII rtuJDUCT PISOLISIS PRODUCT S I S T L O m OF ETHTLBIB OF FBTHIUME M U I H T B OF FRACTION IV FRACTION DIACETATE DIACETATE FRACTION V FRACTION

625 625 PERCENT TRANSMITTANCE ~ / / 5 - PERCENT TRANSMITTANCE 1001— /. 1 tool 40 00 4000 5000 00 4000 5000 2 )

1

1 3 0020 50 40 13,00 1400 1500 2500 3000 2500 4 WAVE NUMBERS IN CM -' -' CM IN NUMBERS WAVE WAVE NUMBERS IN CM -' -' CM IN NUMBERS WAVE AE EGH N MICRONS IN LENGTH WAVE AE EGH N IRN WV LNT I MICRONS IN LENGTH WAVE MICRONS IN LENGTH WAVE 2000 2000 5 6 50 40 13003000 1400 1500 13 7 8 1200 1200 1 01000 1100 9 1100 1000 900 WAVE NUMBERS IN CM -' -' CM IN NUMBERS WAVE WAVE NUMBERS IN CM-* CM-* IN NUMBERS WAVE AE EGH N MICRONS IN LENGTH WAVE 1210 800 PRODUCT FIROIISIS 800 PIROLYS PRODUCT 13 PIROLYS O F ETHYLENE F O O F ETHYLENE F O DIACETATE FRACTION m FRACTION DIACETATE

-V=i

14 700 700 516 15 625 100 c f 0 4 60 100 WAVE NUMBERS IN CM ' WAVE NUMBERS IN CM-' 5000 4000 3000 2500 2000 1500 1400 1300 1200 1100 1000 900 800 700 625 100 100

80

= 60 60 =

53“

2,4 DHHTROPHEHIL KTDRAZOHE OF ACETAIDEHIDE

WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS

WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 5000 4000 3000 25002000 1500 1400 1300 1200 1100 1000 900 800 700 625 100

- -f- = 60 60 =

UUfri'HUrMENXI. H2DRAZZH5 DERIVATIVE OF FIR0LI5I5 PRODUCT OF STHYLEHE DIACETATE

2 3 4 5 6 7 8 910 12 13 14 15 16 WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS PERCENT TRANSMITTANCE PERCENT TRANSMITTANCE

N>

u»

I m vi

a o*

CO

o

h)

S 3

U1

o

PERCENT TRANSMITTANCE PERCENT TRANSMITTANCE WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 50 4000 3000 2500 2000 1500 1400 13001200 1100 1000 900 800 700 100

80

50

40 DIPHOPIOHATE

20

0 3 4 5 6 7 8 9 10 II 13 14 1512 WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS ) I ' I WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 50 4000 3000 2500 2000 1500 1400 1300 1200 1100 1000 900 800 700 100

80

60

40 METHYL CELLOSOLVE ACETATE

20

0 3 4 56 7 8 9 10 II 12 13 14 15 I t WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS PERCENT TRANSMITTANCE - / / * ? - PERCENT TRANSMITTANCE 00 00625 4000 5000 00 0010 10 1300 1400 1500 4000 5000 3000 002500 3000 2500 WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE AE EGH N MICRONS IN LENGTH WAVE AE EGH N MICRONS IN LENGTH WAVE 2000 2000 50 40 1300 1400 1500 1200 1200 1100 1100 1000 1000 900 900 WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE WAVE NUMBERS IN CM-' CM-' IN NUMBERS WAVE AE EGH N MICRONS IN LENGTH WAVE AE EGH N MICRONS IN LENGTH WAVE r AEAE . ACETATE . FPSROHHIS PRODUCT S I H H O R S P OF DERIVATIVE E M Z A R D 2 H iiiun

700 700 625 100 100 WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 5000 4000 3000 2500 20C0 1500 1400 1300 1200 1100 1000 900 800 625 100 ICO

ETKXL CELLOSOLVE ACETATE

WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS

WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 5000 400030002500 2000 1500 1400 1300 1200 1100 1000 900 800 700 625 100

PIH0LTSIS PRODUCT IE40 0? STHTL CELLOSOLVE ACETATE

FRACTION I

2 3 4 5 6 78 9 12 13 14 15 16 WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS WAVE NUMBERS IN CM*' WAVE NUMBERS IN CM*' 5000 4000 3000 2500 2000 1500 1400 1300 12001100 1000 900 800 625 100 100

4 0 £ ui

r a n g ETHTL ETHER

12 13 WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS

WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 5000 4000 3000 25002000 —_ 1500 1400 1300 1200 1100 1000 900800 700 625 100 100

c.it umxinurnuijj* HTDRAZZHE DERIVATIVE OF FIR0U3IS PRODOCT (COffllEKSABIE CASES) OF ETHTL CELIOSOLVE ACETATE

2 3 4 5 6 8 12 13 147 15 16 WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 5000 40003000 25002000 1500 1400 13001200 1100 1000 900800 700 625 100 100

PIROUSIS PRODUCT O F STKXL CELL0S0LVE ACETATE

FRACTION II

WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS

WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 5000 4000 3000 2500 2000 1500 1400 1300 1200 1100 900 800 625 100 100

PIROUSIS PRODUCT OF ETHTL 40 L? CELL0S0L9E ACETATE

FRACTION H I

WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 5000 4000 3000 2500 2000 1500 1400 13001200 MOO 1000 900 800 700 625 100 100

2,4 DUIITROPHEHYL 40 40 t HTDiUZniE DERIVATIVE

WAVE LENGTH IN MICRONS WAVE LENGTH IN MICRONS

WAVE NUMBERS IN CM-' WAVE NUMBERS IN CM-' 3000 2500 2000 1500 1400 1300 1200 1100 1000 900 800 700 625 100 100

!= 40 2 ,4 DIHITROPHam HTDRAZIHE DERIVATIVE OF PYRQLT5IS PRODUCT OF

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The historical development of the under standing of the

mechanisms involved in ester pyrolysis and pyrolyses in

general has been survey and critically discussed. In

conclusion three approaches to the understanding of ester

decomposition have been foundj namely rate determination,

structure analysis and product identification. From a

combination of these three techniques predictions can be

made concerning the thermal stability, mode of decomposition,

and decomposition products of a given ester.

The esters, diacetyl cyanide, glycol diacetate, glycol diformate, glycol dipropionate, methyl cellosolve acetate,

ethyl cellosolve acetate, and ethyl cellosolve propionate

•were pyrolytically studied. The data obtained served to

confirm and complete the understanding of the many phases

of ester pyrolysis. The pyrolysis of acetyl cyanide,

ethyl vinyl ether, vinyl acetate and other by-products were

also investigated. -129-

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AUTOBIOGRAPHY

I, Richard Jui-Fu Leej was born in 1919 of

Chinese ancestry. In 193S I came to the United

States as a student. In the ensuing four years I attended Loras College, Dubuque, Iowa, and Loyola

University of Chicago from which I received the degree Bachelor of Science in 1942. After years of industrial employment and several years of military service in the United States Army Air

Force, I entered the graduate school of Notre

Dame University in 1950. In 1952 I was trans­ ferred to the graduate school of The Ohio State

University.-