to o

C..3

Oil-;-!:



UNITED STATES DEPARTMENT OF COMMERCE • Peter G. Peterson, Secretary W OCC - NATIONAL BUREAU OF STANDARDS • Lawrence M. Kushnek, Acting Director (S C(oo

i([0.'3^'^ The Mechanisms of Pyrolysis, Oxidation, and

^ ' ^ Burning of Organic Materials

Based on Invited Papers and Discussion

4th Materials Research Symposium

held at NBS, Gaithersburg, Maryland

October 26-29, 1970

Leo A. Wall, Editor

Institute for Materials Research

National Bureau of Standards

Washington, D.C. 20234

, National Bureau of Standards Special Publication 357

Nat. Bur. Stand. (U.S.), Spec. PuM. 357, 199 pages (June 1972)

CODEN: XNBSAV

Issued June 1972

For sale by the Superintendent of Documents, U.S. Government Prjnting Office.,Washington, D.C. 20402

(Order by SD Catalog No. C 13.10), Price $3.25 (Buckram)

Stock Number 0303-0928
Abstract

A Symposium on The Mechanisms of Pyrolysis, Oxidation, and Burning of Organic Materials was held at the National Bureau of Standards in October 1970. This volume contains the nineteen

papers presented and much of the discussion which followed.

These papers review and discuss the current status of kinetic studies on the reactions of

organic materials in both gas and condensed phases. The topics covered include: pyrolysis of

hydrocarbons, pyrolysis of polymers, oxidation of polymers, oxidation of organic compounds,

burning of organic compounds and burning of polymers. Particular emphasis is placed on the

elucidation of the mechanisms of reaction in terms of free radicals or other transient species and

physical effects.

Key words: Burning; hydrocarbons; organic materials; oxidation; polymers; pyrolysis.

Library of Congress Catalog Card Number: 76-181873

Identification of some commercial materials and equipment has been necessary in this publication. In no case does such identification imply recommendation or endorsement by the National

Bureau of Standards, nor does it imply that the material or equipment is necessarily the best

available for the purpose.

ii
Foreword

The Institute for Materials Research of the National Bureau of Standards has among its responsibilities the investigation of the chemical and physical properties of materials, the development of metrological techniques for measuring such properties and characterizing materials, and

the dissemination of data and information of importance for the advance of industry and commerce.

In many important areas there are phenomena where measurement methods are nonexistent or inadequate, and where the critical properties needed for characterization of the process or material

are not clearly identified. The burning of polymers, which is a complex process involving the pyrolysis and oxidation of polymers and their gaseous decomposition products, had for a time been in

this category, but recent work has shown encouraging signs of definite progress.

We are therefore particularly pleased to have as the topic for our 4th Materials Research

Symposium, "The Mechanisms of Pyrolysis, Oxidation and Burning of Organic Materials." It is also particularly timely because it is apparent that many of our most pressing and difficult national

problems from solid waste disposal to product safety involve or require knowledge of the combustion, burning and deterioration of organic materials. This symposium was organized for the

purpose of intensively reviewing, examining, evaluating and disseminating available knowledge

and data on these complicated physical and chemical processes.

It is hoped that the pubUcation of these proceedings will be of value to specialized investigators in the particular areas discussed, as well as those actively engaged in attacking such current

problems as flammable fabrics, fire safety, air pollution, waste disposal, recycling of organic materials and the design of safe products.

Such problems will require considerable effort in utilizing science for the development, where

possible, of technological solutions for social and national needs.

J. D. HOFFMAN, Director

Institute for Materials Research

National Bureau of Standards

iii
Preface

This symposium, "The Mechanisms of Pyrolysis, Oxidation, and Burning of Organic Materials," germinated from a suggestion made about six years ago by Dr. Alvin S. Gordon of the

Naval Weapons Center, China Lake, California, that a conference on both gas phase and con- densed phase pyrolysis would be very productive. This idea then developed naturally to include

oxidation and burning. Thus the symposium covers six areas of specialization or subdisciplines,

all of which have the common thread of free radical reactions.

Starting with the historical paper by Professor F. O. Rice, this publication reviews and dis- cusses the current status of the fields of gas phase pyrolysis, polymer pyrolysis, polymer oxidation,

gas phase oxidation, gas phase burning, and the burning of polymers, with emphasis on the physical

and chemical mechanisms of the processes.

The arrangement of the subject matter is intended to gradually and systematically introduce

the reader to a very complex and relatively unexplored area of research, the burning of polymers

or organic materials. It is also hoped that the organization of the various topics will stimulate new

work, identify critical problems and lead to fruitful investigations.

The efforts and kind cooperation of the authors, discussants, and all the participants in this

symposium are greatly appreciated.

Leo a. Wall

Editor and Chairman

iv
Contents

Page

Foreword iii

Preface

IV The Genesis of Free Radical Chemistry

Francis O. Rice 1

Pyrolysis of Organic Compounds in the Gas Phase

Alvin S. Gordon and R. H. Knipe 7

Discussion 15

Pyrolysis of Hydrocarbons

Margaret H. Back 17

Discussion 25

Factors Involved in the Degradation of Polymers in Melts

C. R. Patrick 33

Discussion 44

Pyrolysis of Polymers

Leo A. Wall 47

Discussion 60

Patterns and Problems in the Pyrolysis Behaviour of Synthetic Addition Polymers

G. G. Cameron 61

Discussion 70

On Certain Problems Connected with the Inhibited Oxidation Theory

Yu. A. Shlyapnikov 73

Discussion 78

Mechanism of Peroxydation of Polymers

Adolphe Chapiro 79

Discussion 82

Oxidative Degradation of Polymers and Organic Compounds via Unimolecular

Decomposition of Peroxy Radicals

Jean Marchal 85

Discussion 92

Mechanisms of Oxidation of Polyolefins Below 150 °C

Frank R. Mayo 95

Discussion 99

Polymer-Gas Reactions (Air Pollutants: NO2 and SO2) as Function of Pressure,

UV Light, Temperature and Morphology—A Survey

H. H. G. JelUnek 101

Discussion 109

Controlled Gaseous Oxidations of Organic Compounds

C. F. Cullis 111

Discussion 117

Some Current Problems in Oxidation Kinetics

Sidney W. Benson 121

Discussion 129

Kinetics and Mechanisms of Flames and Flame Suppression

R. M. Fristrom 131

Discussion 140

Carbon Formation in Premixed Flames

K. H. Homann 143

Discussion 150

Fluorocarbon Combustion, Fluorine Supported Combustion Kinetics?

Edward A. Fletcher 153

Burning of Polymers

C. P. Fenimore and F. J. Martin 159

Discussion 168

Pyrolysis and Combustion of Cellulosic Materials

Raymond Alger 171

Discussion 178

The Surface Pyrolysis Boundary Condition for the Combustion of Polymers

R. F. McAlevy 111 and W. S. Blazowski 185

Discussion 192
Francis Owen Rice

The career of Professor Francis Owen Rice as a chemist, scientist and educator covers over a half century. He

was born in Liverpool, England on May 20, 1890 and studied at the University of Liverpool where he received the

B.Sc. degree in 1911, the M.Sc. in 1912 and the D.Sc. in 1916. An 1851 Exhibition Fellowship brought him to Princeton

University and the U.S.A. in 1919. After teaching at New York University, 1919-24, and Johns Hopkins University,

1924-38, he became Professor and Head of the Chemistry Department at Catholic University. Upon retiring from

Catholic University in 1959 he served as Professor and Chairman of the Chemistry Department at Georgetown

University till 1962, and then as Principle Research Scientist, Radiation Laboratory, University of Notre Dame,

1962-68. Since 1968 he has been a Fellow-by-Courtesy in the Chemistry Department at the Johns Hopkins University and also actively continues research in his home laboratory.

Professor Rice is best known for his trail-blazing work on transient free radicals and their role in the mechanism

of organic reactions, such as pyrolysis, oxidation and burning. Subsequently he pioneered studies of trapped free radicals and utilized the results of his free radical studies in developing hypotheses for a rational explanation of

extraterrestrial phenomena such as the colors of the planet, Jupiter. At the present time he is interested in the possibility of reacting solids by mechanical methods.

Thus it is particularly fitting and appropriate that Professor Rice presented the introductory paper at the symposium on the Mechanisms of Pyrolysis, Oxidation and Burning of Organic Materials.

vi
NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

The Genesis of Free Radical Chemistry

Francis O. Rice

Department of Chemistry, Johns Hopkins University, Baltimore, Md. 21218

Modern chemistry is a late-comer among the sciences and came into being only at the beginning of the 19th Century when the phlogiston theory was abandoned to be replaced by Dalton's

Atomic Theory. As Dalton's Theory became firmly established there grew with it the belief known

as the doctrine of valency which resulted, by the middle of the 19th Centry, in the denial of the

very existence of free radicals. In the early 20th Century, the work of many scientists—Hinshelwood, Bodenstein, Taylor-

in the fields of photochemistry and kinetics led to the belief that free radicals not only existed but

played an important part in the mechanism of chemical transformations. Paneth's discovery

that free radicals could be detected by their removal of certain metallic mirrors was followed by

Rice's demonstration that organic compounds did decompose thermally into free radicals, and

then, the Rice-Herzfeld theory that showed it was possible to explain quantitatively the mechanism of the thermal decomposition of organic compounds. In The Annual Review of Physical

Chemistry (1971) a chapter on Paraffin Pyrolysis contains the following sentence: "All in this

field are now agreed that the reactions can basically be described in terms of the Rice-Herzfeld

radical chain mechanisms and that, in general, chains are long."

Key words: Alchemy; free radicals; genesis; phlogiston; valency.

Among the sciences chemistry made a late start. Even at the end of the 1700's when its sister

sciences—mathematics, physics and biology—had

attained a modest degree of maturity, chemical

transformations were explained on the basis of the

mythical phlogiston. How could this have occurred?

Except for medicine, chemistry is the most applicable

of all the sciences and from this standpoint it might

be supposed that its development would come about

sooner than that of the other sciences. Doubtless in

prehistoric days many chemical discoveries were made independently in different parts of the world

and kept as closely guarded family secrets. Thus,

the reduction of ores in fires, the formation of glass

by fusion of pot-ashes with sand and the production

of crude soap by hydrolysis of fats by hot aqueous

potash solutions are all examples of discoveries that

may well have been made before the dawn of history.

Although chemical history does not commence

with Greek science, the Greeks really determined the

early history of chemistry and unwittingly set it off

in a wrong direction. We can detect three main lines

of chemical thought among Greek philosophers:

(1) Ignoble and sordid, unworthy of written record,

is every sort of scientific activity directed to use

or profit. The foregoing quotation is a paraphrase

of written thoughts of many Greek philosophers.

(2) To the mind trained in epistemology, dialectics,

and etiology, the experiment is superfluous.

The menial work of the experimentalist should

be left to slaves and the common man.

(3) Properties are not inherent in a substance. By

suitable manipulation properties may be modified so that a substance A may be endowed with

all the properties of a substance B.

It should be realized that not all Greek scientists

subscribed to all, or even to one, of these postulates.

As with everything else, where large numbers are involved there is always a distribution, commonly

known to chemists by the title of Maxwell's Distribution Law. There were notable experimenters

among Greek scientists and glimmerings of Dalton's

Atomic Theory appear in the writings of several

Greek philosophers. However, by the beginning of

the Christian era chemical opinion had crystalized

fairly definitely along the three lines indicated.

Indeed there seemed evidence for all of them,

especially the third. If one takes a piece of ordinary

sulfur, a brittle yellow solid, and melts it, and

continues heating until it almost reaches the boiling

point, a thick, dark-reddish liquid results. When this

liquid is poured into cold water, instead of getting

back the brittle yellow sulfer, one obtains an ambercolored solid with the properties of rubber. Surely it was reasonable to think that one could take a cheap

base metal such as lead and endow it, not only with

the yellow color, but with all the other properties of

gold. The period of alchemistry lasted for centuries,

during which chemists struggled in vain to perform

such feats as to change lead into gold. We know now

that it is possible to change one element into another,

but the technique is far beyond anything available

to the alchemist—and the cost far outweighs the

value of the product.

1
Although alchemical work continued during the

first ten or twelve centuries of the Christian era, iatrochemistry—the study of chemistry in medicine—developed simultaneously. The iatrochemists

hoped to find a magic drug, the so-called philosophers'

stone, that would cure all diseases and would indeed

confer eternal life. Magic prescriptions and potions

abounded, most of them without any merit whatsoever. However, here and there evidence appeared

often in some quite nauseating prescriptions, that

the search was not necessarily in vain. For example,

among the hundreds of recommendations for treating

infections, one was to make a poultice of the green

mold on stale bread and apply it to the wound. This

green mold is rich in penicillin

About the middle of the 1600's there appeared the

strangest theory of all, the Phlogiston theory

designed in the first place to explain the chemistry

of burning. This wrong theory lasted until about

1800, the birth date of modern chemistry.

In order to understand the acceptance of the

experimental work of Lavoisier and the theory of

John Dalton, which at long last (about 1800 A.D.)

set the mainstream of chemistry in the right direction, we have to go back many hundreds of years to

the work of Roger Bacon, a Franciscian friar who

lived in England during the 13th Century and is commonly regarded as the intellectual originator of

the experimental method. However, it was only

many centuries later that the experiment took its rightful place in natural science and in all disputes

became generally accepted as the court of last resort. Even in the 16th Century the attitude of many

scientists may be illustrated by an extract from the

writings of a Frenchman, Jean Rey. He was considering the question, whether matter could be

made to disappear during the course of a chemical

transformation, and wrote "I affirm that the

examination of weights that is made by the balance

differs greatly from that made by reason. The latter

is only employed by the judicious, the former can be

practiced by the veriest clown; the latter is always

exact; the former is seldom without deception."

By an extraordinary piece of sophistry he "proved"

(correctly) that matter does not disappear during

the course of an ordinary chemical reaction.

The combination of the experimental work of

Lavoisier and the theoretical work of Dalton

resulted in a great flowering of the science, particularly of organic chemistry. However, in one vital area chemists were still misled: their work seemed to

disprove the very existence of free radicals. The

doctrine of valency became generally accepted.

Hydrogen was monovalent, oxygen divalent, nitrogen

trivalent, and carbon tetravalent. The existence of

hydrogen peroxide, hydrazine, ethylene and acetylene

could be explained readily by writing respectively

H—O—O—H;

H

H

N—

H

H

H

H

C=C

H

H

H—C = C—H.

The existence of CO and the many metallic salts

obviously of different valencies, was more difficult to

explain, but even here the idea of latent or "sleeping" valencies gave an acceptable explanation.

Even Kekule's formula for the structure of benzene

fitted into the valency theory. If one writes benzene

as (A) with fixed double bonds, there should be two

ortho disubstituted compounds, (B) and (C); however, only one could be found. Kekule must have

had some sort of intuitive understanding of what

we now call resonance when he postulated the migrating double bond to explain the existence of only one ortho disubstituted benzene.

H

H C H

\ / \ /

C C

II I C C

/ \ y \

H C H

H

(A)

H

A

I C

c c

B

C C

/ \ ^

' H C

H

(B)

H

H

H

C

II C

H

C

I c

C B

H

(C)

The situation in chemistry at the end of the 19th

Century may be illustrated by a quotation from the

2
eminent physical chemist Ostwald [l]': "It took a long time before it was finally recognized that the

very nature of the organic radicals is inherently

such as to preclude the possibility of isolating

them." However, some few chemists persisted in

considering the possibility that free radicals did exist

and play a part in chemical reactions. Nef [2^, in a

series of papers published in the 1890's, concluded

that a wide variety of organic transformations could

be explained by postulating the intermediate

formation of a divalent carbon radical. At that time

little was known about relative bond strengths, and

many of his proposed mechanisms are actually

completely impossible.

The situation during the first two decades of the

20th Century may be illustrated by the following

extracts from chemical textbooks published during

that period:^ "The assumption of the existence of

radicals, capable of existing alone and playing a

special role in chemical reactions, has long been

abandoned." "It is evidently quite out of the

question to isolate free alkyl groups." "Alkyl

radicals do not exist independently. They occur

only in combination with other elements or groups."

"Carbon is quadrivalent, and on this foundation the

whole superstructure of organic chemistry rests."

It is ironical that during the first part of the

20th Century physicists, especially spectroscopists,

not knowing at all that free radicals did not exist, were actively investigating them. Their interaction

with electromagnetic fields (wave lengths of absorption or emission, intensities and widths of lines) and

the general theoretical understanding of the potential

energy curve of two particles, made it possible to

obtain with great accuracy for diatomic particles

the equilibrium distance and the energy of dissociation. Unfortunately, at that time there was not the

close liaison between physics and chemistry which

exists today.

Since free radicals are short-lived entities, it is not surprising that chemical evidence of their

existence was first pointed out by kineticists.

Bodenstein in Germany, studying the hydrogenhlorine reaction, finally became convinced that the

lechanism of the reaction could best be described

uy an atomic chain. Hinshelwood, in England,

succeeded in explaining the complicated mechanism

of the hydrogen-oxygen reaction by a chain mechanism involving free radicals.

However, it was not until 1926 that experimental

evidence for the existence of free radicals was first described [3]. Taylor and Jones [4] found that

pure ethylene heated in a quartz flask at 260 °C

remained stable indefinitely. However, when a mixture of ethylene and either mercury or lead

alkyls was heated under the same conditions, a reaction occurred resulting in the formation of oils. Taylor and Jones explained the results by postulating

1 Figures in brackets indicate the literature references at the end of

the paper.

2 Since these quotations are taken from ephemeral literature, references

are not given. However, these or similar statements may be found in every

chemical textbook published during this period.

a primary dissociation of the metal alkyls:

Pb(CH3)4->Pb+4CH3;

followed by reaction of the methyl radicals with

ethylene as follows:

CH3-|-CH2= CH2—>CH3CH2CH2

CH3CH2CH2-|-CH2= CH2

^CH3CH2CH2CH2CH2 etc.

Combination or disproportionation of the radicals

resulted in the formation of oils. Although this work

did not give any conclusive proof of the existence of

free radicals, the explanation of Taylor and Jones

proved to be essentially correct and was a significant

factor in gaining acceptance for the existence and

importance of free radicals in chemistry.

Just a few years later (1929) a paper appeared by

Paneth and Hofeditz [5] showing how the methyl

radical could be isolated and studied. Their method

was to carry tetramethyl lead in a stream of an inert

gas at low pressure, passing the gases through a furnace and at the end of the furnace bringing them

in contact with a cold lead mirror. They thus showed

conclusively that the tetramethyl lead decomposed

into lead and methyl radicals. The latter combined

with the cold lead mirror, reforming tetramethyl

lead. Although the work of Paneth and Hofeditz

seemed quite conclusive, it did not by any means

gain general acceptance immediately. This was

partly because the experiment itself is difficult: the

lead mirror must be quite clean, and the flow of the

gases, owing to the short life of the methyl radical,

must be very rapid. Also, about a year after publication of the original paper, there appeared a paper by

MuUer and Schultze [6] suggesting that in the

Paneth and Hofeditz experiment the removal of the

lead mirror was merely due to the formation of

hydrogen atoms.

Paneth []7] lost no time in replying and it seems

worthwhile to illustrate the climate of the time by

giving two extracts from his paper:

(1) "It will be shown that the experiments of

Messrs. Schultze and MuUer do not constitute

any proof against the existence of free methyl."

(2) Footnote: "It is a fact well known among

logicians that the undependability of such a conclusion is very easily recognized when it leads to an absurdity. As a striking example of

their erroneous logic we may give the following

example: All geese have 2 legs. This girl has

2 legs. Therefore it follows that this girl is a goose."

My own interest in free radicals started two or

three years after I joined the Department of Chemistry at The Johns Hopkins University in 1924.

Hinshelwood [8] had already studied the kinetics of

the thermal decomposition of acetone and found that

the decomposition followed a unimolecular law. He

assumed that the decomposition was represented by

3
the equation

CHaCOCHs^CO+CaHe

One of my students and I (9) soon found that the

decomposition of acetone was really represented by

the equation

CH3COCH3^CH2= CO+CH4

At that time chemists thought that organic transformations occurred through a series of migrations of

atoms or radicals, without the intermediate formation of free radicals. Thus, the thermal decomposition of acetone into ketene and methane would be

represented by the migration of a hydrogen atom

from one methyl group of the acetone molecule to

the other, followed by rupture of the acetone

molecule, thus producing ketene and methane.

This seemed to me to be asking a rather impossible

behavior of a hydrogen atom. In the gaseous state,

at atmospheric pressure or lower, an acetone

molecule spends most of its time out of range of the

force fields of other molecules. It seemed reasonable

to assume that the heat energy flowing into the

molecule would finally result in the rupture of the

weakest bond, with the consequent formation of two

free radicals, which would then fly apart and each

react with the surrounding acetone molecules. It turned out that considerations of this kind, guided

by our knowledge of bond strengths and the principal of least motion,^ led to a chain reaction which

gave the products ketene and methane. Having

established the theory for acetone, it was easy to test

it by comparing the theoretically predictable

products of decomposition of hydrocarbons, other

ketones, ethers, etc., with those actually produced

experimentally. Such comparisons showed that free

radical mechanisms provided a satisfactory ex- planation for the thermal decomposition of aliphatic

organic compounds.

However, when the theory was first proposed we

had no evidence at all of the actual existence of free

radicals, during the thermal decomposition of

organic compounds, and there thus was little hope of

gaining acceptance of our postulated mechanisms.

It was at this very time, when we needed it most,

that the paper by Paneth and Hofeditz [5] appeared.

To be sure, they had only discovered a method of

detection of CH3 radicals by their removal of lead

mirrors, and it remained for us therefore to determine that the thermal decomposition of ordinary

organic compounds, such as hydrocarbons, ketones,

ethers, etc., resulted in the formation of free radicals.

Our experiments were successful, and we found that

all these compounds readily decomposed in this

manner. We thereupon published the results of our work [11].

In the years that followed, the reception accorded our interpretations is best illustrated by such

excerpts as the following:

(1) From a paper by R. E. Burk [12]:

"The thermal decomposition of straight chain

' These matters were clarified in a later paper by Kice and Teller [10].

paraffins is of outstanding interest for at least two

reasons. Firstly, the reactions are basic ones in the

technical "cracking" of petroleum fractions, practiced on a huge scale; secondly, since the C—

C

bonds are probably all alike energetically, straight

chain paraffin molecules are almost ideal ones for

mechanical and statistical calculation of reaction

rates ... A paper on the decomposition of paraffins

has just appeared, in which the first step is assumed

to be free radical formation. The decomposition of

radicals themselves

CH3CH2CH2CH3->2CH3CH2

was thought to play an important part ... It seems

unlikely that free radicals should be formed by

thermal decomposition of hydrocarbons, because

reaction 24 can occur with the evolution of some

70,000 or 100,000 calories, and with no

CH3CH2 +CH3CH2—>C2H6 +CH2^CH2

evident kinetic impediment, before the radicals leave

the sphere of reaction."

(2) From a paper by Kassel [13]:

"The decomposition of paraffin hydrocarbons has

been discussed by F. O. Rice and by Burk; Burk

upholds the view that a paraffin is split out in the

initial act, while Rice assumes free radical formation.

Both attempt to obtain evidence from the composition of the reaction products but without any

very striking successes." "There is a second

explanation for the results of Rice and Evering. It seems possible that the primary process in these

reactions may be of the type

/

CH3CH2CH3—>CH4-|-CH3CH\

This view has the advantage that the methyl group

in splitting off captures a hydrogen atom from the

carbon to which it was itself attached, rather than

from a more distant atom, as is the case when ethylene is the initial product ... It is likely that

CH3CH would behave like a free radical in a Paneth

\

test such as used by Rice and Evering."

(3) Still another extract may be given: Simons

and Dull

"On the Reaction of Gaseous Methyl with Carbon

Tetraiodide" [14].

". . . the identification of these as free radicals

(odd molecules) is still open to question . . . Its reaction with carbon tetra -iodide to form methyl

iodide is an argument against the assumption of

Kassel that the active particles are alkylidene

radicals, for in that case di-iodomethane should

have been formed."

It finally became obvious to organic chemists that

free radicals were important intermediaries in many

organic transformations. There remained, however,

the problem that troubled physical chemists, principally kineticists, that thermal decompositions of

organic compounds followed simple laws, of the first

4
or second order, and this seemed quite impossible on

the basis of the mechanisms I was proposing.

Fortunately for me, at that time a colleague in the

Physics Department, Karl Herzfeld, was also

interested in reaction mechanisms. Together we

made an examination of the free radical theory which

I had proposed, with the object of finding out what

kind of kinetics it would lead to. It turned out that

the problem was a very messy one algebraically, but

our calculations finally led to a simple solution,

depending on the chain ending [15J. It was found

that thermal decompositions such as, for example,

that of acetone, should follow a unimolecular law,

whereas the thermal decomposition of acetaldehyde

should follow a 1.5 order. It was actually found that

published data for the latter reaction fitted better a

1.5 order of decomposition than a second order. We

submitted this paper to the American Chemical

Society for publication and it seems worthwhile to

give some extracts of our correspondence with the

late Professor Arthur B. Lamb, who was then editor.

Comments of referee I (in part)

"The paper is important and is concisely and

rather elegantly expressed; and I have concluded

that I should recommend its publication in full . . . The concept of free radicals in kinetics is a very live

subject now, and in my opinion this paper is a distinct contribution to this field."

Professor Lamb submitted the paper to another

scientist who preferred not to act as a referee and

wrote:

"... I know that we should be criticized for

publishing such a long paper with so many equations

and I believe it would be better to have the judgement of another referee."

The comments of the other referee were as follows

(in part)

"I am very decidedly of the opinion that the

material involved should be published, but I cannot

agree with the judgement of your first referee that

it is concisely expressed . . . ... So far as this particular paper is concerned, I think you would be justified in requiring that it be

shortened by 10 to 12 pages."

Then followed the Editor's decision (in part)

"There is no question as to our desire to publish

this manuscript as a valuable contribution to the

subject of free radicals ... I am glad to accept your

manuscript for publication provided it can be con- densed from its present 27 typewritten pages to not

over 17 typewritten pages."

The Editor's decision posed quite a problem for us, which we solved in the following manner. We went

over the paper carefully, but succeeded in shortening

it only by several pages. As it was necessary to retype

the paper, and as we were unable to shorten it by

the required 10 pages without emasculating it altogether, we resorted to the use of a typewriter

with somewhat smaller type, and used narrower

margins. In this way we met the Editor's requirements.

The publication of the Rice-Herzfeld paper [15]

removed the final objections to the theory and

thereafter the progress of free radical chemistry

proceeded exponentially.

It is extraordinarily difficult to select for mention

any names among those of the many distinguished

scientists who contributed to the advance of free

radical chemistry. However, there are a few lines of

free radical research that call for special mention.

About 1932 I had become interested in the methylene

radical and my former student, Glasebrook, and

I [16J discovered that although it does not combine

with metallic lead it can be studied through its combination with tellurium. We had no notion at the

time of the properties of this extraordinary particle

and it remained for the physical chemist, George

Kistiakowsky [IT], at Harvard and the organic

chemist, Doering [IBJ, then at Yale, to uncover

its properties.

At Rochester W. A. Noyes, Jr. [19] founded a school interested principally in a study of the photochemical production of free radicals. The work of

Porter and Norrish, in England [20] demonstrated

that it was possible through flash photolysis, actually

to obtain the spectra of free radicals.

Lastly the work of Dr. E. W. R. Steacie [21] at

Ottawa should be mentioned. Canada does not have

as many centers for free radical research, nor does it produce a volume of work as great as do some other

countries, but in proportion to its population and

wealth its production far outranks every other

country; this is due in large measure to the foresight

and planning of E. W. R. Steacie.

Finally, since this is essentially an historical paper,

we should ask ourselves "what will future historians

have to say concerning free radical science today?"

What will scientists think, 200 years from now, when

reading the papers presented at this meeting? Are

we going in the right direction? Have we finally

arrived, by and large, at a true and full understanding of free radical reactions? In my estimation,

the answer will be NO.

As an example, for many years I have had a

feeling of uneasiness when reading papers describing

measurements of the speed of chemical reactions, in

which it is mentioned that the surface of the vessel

must be "conditioned" before reproducible results

could be obtained. However, reproducibility can

hardly be regarded as a safe criterion of truth: the

phlogiston theory gave consistently reproducible

results—which were exactly opposite to the truth.

There is a second Rice-Herzfeld paper [22] in

which this matter is discussed. There is the possibility that chains may start from the wall, go out

into the gas phase and terminate on the wall.

Only a few decades ago the difficulties of travel were such as practically to preclude a Symposium

such as this, which provides an opportunity for

scientists from far distant countries to engage in

informal discussion. Out of such deliberations are

likely to emerge many new lines of attack on the

challenging and important problem of controlling

the pyrolysis, oxidation and burning of organic

materials.

5
References

[1] Ostwald, C. W. W., Existence of Free Radicals. Quoted

by W. E. Bachman in "Organic Chemistry." H. Gilman,

p. 582. Wiley, New York, 1938.

[2] Nef, J. U., Annalen 270, 267 (1892); 280, 291 (1894):

287, 265 (1895): 298, 202 (1897).

[3] Taylor, H. S., Proc. Am. Phil. Soc, 65, 90 (1926).

[4] Taylor, H. S. and Jones, J., J. Am. Chem. Soc, 52,

1111 (1930).

[5] Paneth, F. and Hofeditz, F., Ber. 62, 1335 (1929).

[6] MuUer, E. and Schultze, G., Z. physik. Chem. B6, 267

(1929).

[7] Paneth, F., Zeit. physik. Chem., 7, 155 (1930), Translation. Remarks on the work of Messrs. G. Schultze

and E. Muller on lead hydride, by F. Paneth.

[8] Hinshelwood, C. N. and Hutchinson, W. K., Proc. Roy.

Sox., lllA, 245 (1926).

[9] Rice, F. O. and Vollrath, R. E., Proc. Nat. Acad. Sci.,

15, 702 (1929).

[10] Rice, F. O. and Teller, E., J. Chem. Physics, 6, 489

(1938).

[11] Rice, F. O., Johnston, W. R. and Evering, B., J. Am.

Chem. Soc. 54, 3510 (1932).

[12] Burk, R. E., J. Phys. Chem. 35, 2448 (1931).

[13] Kassel, L., Chem. Rev. 10, 16 (1932).

[14] Simons, J. H. and Dull, M. F., J. Am. Chem. Soc. 55,

2696 (1933).

[15] Rice, F. O. and Herzfeld, K., J. Am. Chem. Soc. 56,

284 (1934).

[16] Rice, F. O. and Glasebrook, A., J. Am. Chem. Soc.

(a) 55, 5329 (1933); (b) 56, 2381 (1934).

[17] Kistiakowski, G. B. and Sauer, K., J. Am. Chem. Soc.

78, 5699 (1956).

[18] Doering, W. von E., Buttery, R. G., Laughlin, R. G.

and Chanduri, N., J. Am. Chem. Soc, 78, 3224 (1956).

[19] Noyes, W. A. Jr., Hammond, G. S., and Pitts, J. N.,

Advances in Photochem., Interscience Publishers, 1963,

p. v.

[20] Norrish & Porter, See Annual Rev. of Physical Chem.,

p. 17 (1962), Published by Annual Rev., Inc.

[21] Steacie, E. W. R., Atomic and Free Radical Reactions,

1954, Reinhold Publishing Co., New York.

[22] Rice, F. O. and Herzfeld, K., J. Chem. Phys., 7, 671

(1939).

6
NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Pyrolysis of Organic Compounds in the Gas Phase

Alvin S. Gordon and R. H. Knipe

Research Department, Chemistry Division, Naval Weapons Center, China Lake, Calif. 93555

A critical review of the pyrolysis of organic compounds is made. The review is divided into

hydrocarbons, halides, and oxy compounds. Most hydrocarbon reactions have an important unimolecular component, either the initial step in the case of chain reactions (e.g., saturated hydrocarbon decomposition) or a one step reaction which may involve a transient biradical structure. There are many types of pyrolysis reactions in the latter category.

For halocarbon pyrolysis there are two reaction paths, one involving an unimolecular rupture

of the carbon-halogen bond and the other an unimolecular elimination of the hydrogen halide

molecule via a polarized activated complex. The preferred path is determined by the self-quenching

efficiency of the radical path. If highly efficient, the molecular elimination path will dominate

and vise versa. In some systems, the two paths are competitive.

Oxy hydrocarbons such as acids and esters pyrolyze via the polarized complex molecular

elimination path. Ethers appear to pyrolyze via a free radical chain mechanism while ethylene

oxide is initiated by a unimolecular rearrangement to acetaldehyde with excess energy. The

vibrationally excited molecule can decompose into two radicals and initiate chain decomposition.

Key words: Decomposition; kinetics of decomposition; pyrolysis of hydrocarbons; radical mechanisms; unimolecular decomposition; molecular mechanisms.

1. Introduction

It is fair to inquire how the pyrolysis of organic

compounds is related to flamibility of organic

materials. The relation is that when organic materials

burn in a diffusion flame, the condensed phase is heated by the flame and provides gaseous molecules,

with and without decomposition of the organic

material; these gaseous molecules form the fuel for

the flame. It has been demonstrated that the fuel

molecules inside the flame mantle are extensively

pyrolyzed with only a trivial amount of oxygen

being present and probably acting as an efficient

catalyst [1,

2].i

The gross aspects of the pyrolysis of most organic

compounds were delineated by F. O. Rice in the

early 1930's [3^. For the bulk of such molecules, the

reaction is initiated by the breaking of the weakest

C—C bond to form two radicals. These radicals and

their daughter radicals, abstract an atom from the

parent compound to produce product molecules and

new radicals in a chain mechanism. In the case where large radicals are formed, these will isomerize

to other radicals as well as pyrolyze by unimolecular

reactions to olefins and smaller radicals. The chain

reactions are quenched by radical-radical combination and disproportionation reactions.

In addition to the above chain decomposition

reaction, organic molecules may decompose in a unimolecular reaction to give two molecules without

free radical intervention. Other reaction paths are unimolecular isomerization and bimolecular re-

^ Figures in brackets indicate the literature references at the end of this paper.

actions involving one or two reactants. For these

last processes, the Woodward-Hoffman orbital

correlation rules have been able to provide simple

rationalizations of the products in many cases [4].

Usually, one of the many possible mechanisms

dominates. Bond strengths, molecular structure and

polar effects have been used to explain why a particular mechanism is dominant in a given temperature

range.

For crystalline solids, the pyrolysis reaction rate

may be determined at the interface between reacted

material and the unreacted crystalline material as well as by reactions in the solid volume. For liquids,

because of their comparatively high dielectric constants, the competitive reactions may include

ionized species which are unlikely in gas phase

reactions except at temperatures well above 1000 K.

In condensed media the original two radicals, formed

by decomposing a molecule remain in close proximity

for a period of time because of hindrance to diffusion.

This "cage" effect leads to higher radical-radical

reaction rates than for the same average radical

concentration uniformly distributed.

In this paper we shall restrict our review to

homogeneous gas phase reactions and include

examples of hydrocarbons, halocarbons and some

oxygenated compounds. Although we recognize that

pyrolysis reactions of economic significance such as occur in flames, production of commercially useful

hydrocarbons from petroleum feedstocks, and lampblack manufacture all involve very deep pyrolysis,

we shall focus on the early stages of reaction before

secondary and higher products play any significant

role.

7
2. Unimolecular Reactions

Since unimolecular reactions are so prominent in

the gas phase decomposition of organic molecules,

they have received much theoretical and experimental attention. There is an extensive recent

review of thermal unimolecular reactions of hydrocarbons and we will attempt not to extensively

overlap with this work [5 J. Grossly, a thermal unimolecular reaction is a consequence of an occasional molecule acquiring

enough energy to undergo chemical change as a

result of collision with other molecules in its environment. If the collision rate is sufficiently rapid relative

to the rate of reaction of these energetic molecules,

their steady state concentration will approximate

the limiting Maxwell-Boltzmann distribution with

the result that the rate of reaction at a particular

temperature is a linear function of the concentration

of reactant molecules. At lower collision rates the

concentration of sufficiently energetic molecules falls significantly below the Maxwell-Boltzmann value

with an accompanying "fall-off" in the rate of

reaction. The pressure at which "fall-off" becomes

significant decreases as the complexity of the

molecule increases. As the temperature is increased,

the pressure at which "fall-off" becomes significant

increases. At flame temperatures, "fall-off" will be of

importance for unimolecular reactions involving even moderately complex molecules.

Because of the principle of microscopic reversibility, the pressure independence of the rate of

breaking a bond also applies to the making of a bond

under the same conditions. The experimental

demonstration that the rate of ethane decomposing

to two methyls is first order in ethane concentration,

indicates that under the same conditions two methyls

combine without need for a third body. Because of

the thermodynamic relationship involving the forward and reverse elementary reaction, it is often

more convenient to study the reverse reaction to

generate information about the forward reaction.

Thus, Benson and Bose have studied the bimolecular

addition of HI to olefins in order to gain insight into

the unimolecular decomposition of alkyl iodides [^6^.

There are two general approaches that have been

widely employed to fit experimental data on unimolecular reactions. The statistical approach is exemplified by the RRKM formulation [7]. The

stochastic approach has been diversely applied to

non-thermal initiation of chemical reaction Q83 and

as a tool to investigate the limits of validity of the

statistical approach [9]. The Marcus improvement

on the older RRK approach consists of using the

actual molecular frequencies and anharmonicities

where they are known and replacing the empirical

frequency factor by an Eyring-like transmission

coefficient for the activated molecular states [10].

As a consequence, the RRKM formulation leads to

ostensibly apriori estimates of the high pressure

limit of the unimolecular rate constant which

correspond to the estimates of transition state theory

as well as fitting "fall-off" with decreasing pressure.

There remains an element of adjustability in that

the molecular degrees of freedom are divided into

adiabatic (do not contribute energy to the process)

and active. The "fall-off" behavior is influenced by

the manner in which this division is made. Also, the

frequencies and/or anharmonicities in the activated

molecules are not precisely known and hence may be

subjected to empirical variation to adjust both the

high pressure limit and the "fall-off" behavior. The

"fall-off" behavior can be further adjusted by

variation of the collision efficiency. Recently, it has

been shown that for the case of thermal isomerizations, inclusion of back reaction from excited isomer

configurations can lead to descriptions of the "falloff" behavior with physically more acceptable

collision efficiencies that would otherwise be neces- sary [11].

A major disadvantage of the RRKM approach

over the RRK is the complexity of the computations

which necessitates the use of large scale computers.

The advantage is that it is possible to obtain a description of the "fall-off" in the unimolecular rate

constant which may be valid over a considerable

range of temperature [12]. More sophisticated

treatment of the mechanics of the dissociation process suggests that parametric adjustments used in

fitting the observed kinetics may reveal some qualitative features of the nature of the excited molecular configuration which is undergoing reaction [13,

14]. However, it must be recognized that the principal value of the RRKM approach is in correlating

available data. The wide variability possible in

RRKM estimates of "fall-off" behavior has recently

been illustrated for methyl radical combination by

Casas et al. who suggest that the data for this

reaction should be fit with a somewhat tighter

activated complex than had been previously supposed [15].

T

CH2= CH —CH=CH2

In what follows we will discuss the pyrolysis of

various types of hydrocarbons which follow a unimolecular path. For cyclic saturated hydrocarbons,

the pyrolysis path is overall molecular, although a

transient biradical intermediate may be involved.

The simplest type of reaction which involves electronic rearrangement may be illustrated by the work

of Hauser and Walters on cyclobutene pyrolysis

[16]. A C—C bond breaks and the electrons rear- range to form 1,3 butadiene. The cooperative reac- tion has an ^ = 10^3-6 g-i £ = 32.7 kcal/mole.^

The preexponential factor corresponds to a small

entropy change in going to the activated complex.

2 Throughout this paper experimental rate constants will be denoted by

their preexponential factor A and activation energy E.

8
indicating the complex is cyclic with no additional

degree of rotation over the normal molecule.

Frey and co-workers have pyrolyzed vinyl cyclopropane and alkyl substituted vinyl cyclopropane.^

In all instances there is no evidence of any H migration but an isomerization initiated by opening the

allylic C=C bond; e.g..

R

R

R'^

c:'

R

c—

c

R'C

R

C —C

// I

where R and R' may be an H or an alkyl group. All

these reactions are unimolecular with

kcal/mole, and ^'^lO^* The presence of the

allylic C—C bond lowers the energy of activation over opening a cyclopropane ring. If it is assumed

that all the allylic energy is in the activated complex,

which is quite imlikely, then a value for the allylic resonance of 13-15 kcal/mole may be obtained by

comparison of the energy of activation for opening

the non-allylic and the allylic ring. It would be

perhaps more realistic to say that the value is the

minimum allylic energy. The rotational energy of the

vinyl group is lost when going to its resonance

stabilized activated complex, thus causing a negative

entropy of activation. The activated complex of

cyclopropane has less hindered rotation than that of

vinyl cyclopropane with a resulting higher preexponential factor. Pyrolysis of the isomer of vinyl

cyclopropane, methylene cyclobutane has also been

investigated. The molecule cannot easily rearrange

to cyclopentene and falls apart unimolecularly to

allene and ethylene.

Another class of molecular isomerizations which

are closely related to vinyl cyclopropane isomerization are the sigmatropic rearrangements which

involve the migration of a sigma bond flanked by one

or more double bonds. The Cope rearrangement is an example of such a rearrangement. It may be

illustrated by the reaction

In a cooperative process, Ci and Ce atoms bond

while the C3—C4 bond is broken and the electrons

rearrange. Since some rotational degrees of freedom

are lost in going to the cyclic activated complex, the

entropy of activation is negative [17, 18]. The re- action mechanism is easily followed by using deuterium marked 1,6 groups in the reactant or by using an appropriately alkyl substituted compound. Typically, ^'^10" s~\ and E^36 kcal/mole.

If C—H bonds are in the right configuration to be

abstracted cooperatively with an electronic shift,

then the molecule will rearrange in a unimolecular

reaction analogous to the Cope rearrangement. A

typical example of many reactions that have been

studied is [19 J.

H

./

H

For these reactions typically A'—'10^'^ s~' and E^^SS

kcal/mole indicating, as would be anticipated, that

the formation of the cyclic activated complex reduces

the number of free rotations and results in a negative

entropy of activation. These reactions are similar to

the Cope rearrangement in that both cases involve a

cyclic activated complex in which one bond is broken

and another made with associated electronic re- arrangement. In both of these cases there is little or no AH of reaction; hence the rates of the forward and

backward reaction are the same. Cyclopropane pyrolysis has been extensively

studied. The kinetic data may be interpreted as resulting from the intermediate formation of a biradical by scission of one of the C—C bonds or the

reaction path may be via an expanded ring. We will

examine the biradical mechanism first. Since the opening of the ring is over 60 kcal

endothermic, any subsequent rate processes with a much higher rate constant such as the 1 ,2 H shift to

form propylene occurs very rapidly in the thermal

environment necessary for the formation of the

biradical. A most informative study of cyclopropane

has been reported by Rabinovitch, Schlag, and

Wilberg [20]. It involved the geometrical isomerization of cis 1 ,2-dideutero cyclopropane which exists

in the cis and trans form because H (or D) atoms

cannot rotate through the ring. They studied the

simultaneous geometrical and structural isomerization and found that the overall energy of activation

for the geometric is 1.3 kcal/mole lower than for the

structural while the preexponential factor for the

geometric isomerization is about seven times that for

the structural. Benson has interpreted these data in

terms of the same slow step in both processes which

is the opening of the ring to form a trimethylene

biradical. This allows free rotation of the methylenes

[21].

The 1,2 H shift from trimethylene to form

propylene has a low E^ct (calculated) of about 9.5

kcal/mole. On the other hand, Jackson and McNesby

9
showed that 1,2 H shift in n-propyl radicals did not

complete with C—C bond scission, [22] while studies of a similar reaction of ethyl radicals in our laboratory indicates that the 1 , 2 H atom shift is about 10

percent of the loss of H atom at 500 °C [23]. The

striking difference in the behavior of various 1 ,2 H

shifts may be related to the AH of the reaction.

The difference in the heats of formation of the

biradical and cyclopropane at about 450 °C (a con- venient temperature for cyclopropane pyrolysis

studies) is about 56 kcal/mole. On the other hand

the experimental energy of activation is about 64

kcal/mole, indicating that the trimethylene has an

energy about 8 kcal/mole less than the activated

complex, so that the trimethylene must climb an

8 kcal/mole barrier to reform the ring. Since the

back reaction must be many times faster than the

forward, the two species are likely to be in equilibrium. The various reaction paths are

1 CH2CH2CH2 ^ CH3CH=CH2

From the AS° of the reaction and the experimental

preexponential factor for the forward reaction as well

as the above data, ^2 and ks are readily calculated

[24]. The various self consistent specific rate constants are

A:i=10i«10-8*'""'/4-6^

A;2 = 101310-8200/4-6r,

It should be added that an equally consistent set of

kinetic parameters can be formulated if it is assumed

that the trimethylene biradical plays no role, but

that the activated complex allows free rotation and

is the species which leads to both geometric and

structural isomerization. Rabinovitch et al. refer to

such an activated complex as an expanded ring.

Two papers in the recent literature [25, 26] at- tempt to resolve the two mechanisms, one group by

a study of

A

CD3 / \cH3

CH3 "CD

and the other group by using

In these compounds the change in optical activity is used to gain details of the mechanism. One team

[25] concludes that much of the reaction path is via an expanded ring, while the other workers [26] con- clude that the biradical path is most likely.

It should be noted that all the above types of

unimolecular processes are isomerization reactions,

mainly structural and some geometrical. Many

important unimolecular cyclic hydrocarbon pyrolysis

reactions occur either by a cooperative process to

give final products, or by breaking C—C bonds. This

process may proceed via biradical or expanded ring

cooperative processes which are not readily differentiable as noted for cyclopropane above. Examples

of reactions giving products in one step are illustrated below.

The pyrolysis of cis and trans 1,2 dimethyl

cyclobutane are examples of unimolecular reactions

which result in molecular products [27]. Various

simultaneous reactions may be studied: formation of

ethene and butene-2, formation of 2 moles of propene,

and isomerization of the parent compound. Also,

the butene-2 may be analyzed for cis and trans

components. The kinetic parameters for the various

reactions paths are listed below:

A (s-i) E (kcal/mole)

Cis-^Trans 6AXW 60.1

Transacts 3.7Xl0i* 61.3

Cis^Ethane+Butene-2 S.lXW^ 63.0

rrans^Ethene+ 2.9X10'^ 63.4

Butene-2

Cis^Propene 3.0X101^ 60.4

rrans->Propene 2.8X101^ 61.6

Unlike the cyclopropane (and dimethyl cyclopropane) pyrolysis, the rate of isomerization of the

parent compound is much slower than the rate of

formation of products. The butene-2 product largely

reflects the steric configuration of the 1,2 dimethyl

cyclobutane precursor.

Although Woodward-Hoffman rules are favorable

for the reaction going by a biradical mechanism and

unfavorable for a cooperative cleavage to final

products, the qualitive prediction does not take into

account that the paths may have different endothermicity. Thus, the AH reaction ~40 kcal/mole

for a cooperative ring break to final products is consistent with the observed Ee,ct of over 60 kcal/mole.

This path would be stereospecific. For the biradical

path there would be a competition between the back

reaction to reform the cyclobutane ring and the

breaking of the C—C bond to form two olefins. The

back reaction to reform the ring has a negative

entropy of activation while that to break the C—

C

bond is zero or positive. Additionally, the back

reaction, by analogy with trimethylene going to

cyclopropane, has a 5-7 kcal/mole barrier, while the

breaking of the C—C bond has a barrier very close to

zero. Therefore, the C—C bond break should be

many hundreds of times more rapid than the re- forming of the cyclobutane ring. When the biradical

is in the activated complex configuration to form the

olefins, it may have some olefinic character hindering

the rotation around the C—C bonds which preserves

the original isomer configuration. Thus both the

cooperative expanded ring and the biradical decomposition paths are consistent with the results.

The elimination of molecular hydrogen from a hydrocarbon is a rare process. It has been quite well

10
established in a number of compounds such as cyclopentene, and 1 ,4-cyclohexadiene. For cyclopentene,

pyrolysis of the deuterium compound marked on the

3 position has shown that most of the eUmination is 1,3 rather than 1,2 [28]. The results are in accord

with the predictions of the Woodward-Hoffman

orbital correlation rules [4]. The 1,2 elimination is predicted to involve H atoms which are trans to

each other and thus must react through the ring as well as a 4 membered cyclic activated complex. Such

a reaction is very difficult compared with the 1,3

which involves elimination of cis H's from a six

membered cyclic activated complex.

1 ,4-cyclohexadiene has been shown by a number

of investigators to pyrolyze cleanly to benzene and

molecular H2, while 1 ,3-cyclohexadiene is much

more stable to pyrolysis and seems not to eliminate

molecular H2 [29, 30, 31]. Steric considerations for

both the 1,3 hydrogen elimination from cyclopentene and hydrogen elimination from 1 ,4-cyclohexadiene suggest that the activated complex must involve considerable structural distortion.

The Woodward-Hoffman rules have been used to

rationalize the products from a variety of unimolecular reactions [32]. Orbital correlation has not

been widely considered relative to radical reactions.

Longuet-Higgins and Abrahamson have discussed

the case of the cyclopropyl radical, ion and cation.^^

Since for either the disrotary mode or the contrarotary mode a change in the overall symmetry of the

electronic ground state is involved in the isomerization of cyclopropyl radial to the allyl radical, these

authors suggest that the reaction should have a large

activation energy. McNesby and Gordon have

presented evidence that this isomerization does in

fact proceed at an appreciable rate at temperatures

below 400°C [34]. Relative to the symmetry consideration, H2 elimination from the cyclopentyl

radical is analogous to the disrotary cyclopropyl

isomerization. Gordon has presented evidence for

this reaction at relatively low temperature [35].

Thus, it appears that restrictions due to symmetry

correlations may not be as severe for odd-number

electron systems as for the even-number electron

system.

3. Halo and Oxy, Carbon-Hydrogen

Compounds

The pyrolysis of alkyl halides have been reviewed

recently by MaccoU [36]. The pyrolysis of alkyl

fluorides goes at a high enough temperature so that

the walls of the reaction vessel are strongly attacked,

leading to difficulties of interpretation. Pritchard,

at al. have shown that the reverse reaction forming an excited partially fluorinated ethane molecule via

combination of two CH2F radicals eliminates HF

and gives CHF=CH2 [37]. It should be noted that

this experiment is performed at room temperature

and the C—C formed in making ethane has more

than enough energy to surmount the energy of activation barrier for eliminating HF. The study involves

super excited molecules relative to the thermal pyrolysis of 1 ,2 difluorethane. It is interesting that the

analogous reactions of CH3 radicals has never re- vealed the slightest evidence of H2 elimination from

the excited ethane, even though both thermal processes are quite analogous. It can be concluded that

there is a fundamental difference in the mechanism

for HF and H2 molecular eliminations from a molecule. As will be seen below, HCl, HBr, and HI are molecularly eliminated from the corresponding alkyl

halides as is the organic acid from an ester when

pyrolyzed. Both xanthates and carbonates also

pyrolyze via molecular elimination.

A very striking factor in the pyrolysis of alkyl

halides (methyl halide is a special case) is the large

effect of substitution of a hydrogen on the a carbon

atom (which contains the halogen) by an alkyl

group. There is a much smaller effect of alkyl substitution on the 0 carbon. MaccoU and Thomas first pointed out that the homolytic bond (RX—*^R +X)

strengths of the various C—X bonds did not correlate

with the energies of activation for elimination [38].

On the other hand, there is a strong correlation

with the heterolytic bond dissociation energies

(RX—>R++X~). MaccoU put forth the explanation

that the activated complex was associated with a highly polarized C—X and some small polarization

of the C—H of the 4 atom complex. Benson and

Bose have extended the concept of MaccoU to a complex where the H and X have a one electron

bond with formal changes of +| and — ^ respectively [39]. Calculations based on the model including a polarization term as well as charge attraction

and repulsion to give the minimum energy of the ion

pair have resulted in values in moderately good

agreement with observation. Benson and Haugen

have tried to include molecular Ho elimination in

this model, but the experimental validation is not

very convincing [40]. In addition, there is no evidence that an alkyl group substituted on the a car- bon in hydrocarbons has the dramatic effect on

pyrolysis parameters noted with alkyl halides.

The molecular elimination of HX is complicated

by the possibility of a radical reaction initiated by a C—X scission. In his review, MaccoU quotes Barton

and Oryon, and Howlett that a chain reaction is the preferred path for decomposition of a chloro

compound [36, 41, 42]. Since the molecular elimination has a much lower Eact ('^60 kcal/mole) than

the breaking of a C—C bond ('-^80 kcal/mole), the

radical reaction can only be favored by long chains

resulting from the original break. It will decompose

by the competitive molecular mechanism if the

reactant or the products are strong radical quenching

molecules, or if the preferred abstraction path leads

to a radical intermediate which cannot stabilize

itself by breaking a C—CI bond. As an example,

CH2CICH2CI decomposes via a radical chain because

11
anc

Cl+CHaClCHzCl^HCl+CHClCHaCl

CHC1CH2C1^CHC1CH2+ CI

On the other hand, CH2CICH3 and CHCI2CH3

decompose via molecular mechanisms since the preferred abstraction position on the parent results in

CHCICH3 and CCI2CH3 radicals which cannot easily

decompose to continue the chain. The last mechanism

is substantiated by study of chlorine atom attack on

ethyl chloride, resulting in 1 , 1 dichloroethane and

1,1,1 trichloroethane [43]]. The AJEact for abstraction of H from the a and |8 carbon position thus

strongly favors the a position. To account for the

results at 600 K, we can say that at least 99 percent

of the chlorine atoms must abstract an aH relative

to a /3H from the parent compound. Taking into consideration that there are 3 j8H atoms and 2 aH atoms

in ethylchloride

EalRT

and

£^;3 — £^a>6kcal/mole

For the corresponding bromides, the radical chain

process is favored over the molecular elimination,

indicating that for Br abstraction of H from the a and |3 positions in CH3CH2Br kinetically is more

favorable for the position than for the corresponding chloro compound. When both chlorine and

bromine are attached to carbon atoms, the rate of

H abstraction by Br atoms for an analogous series are in the order CH3C1> CHsBr > CH4. A very

recent paper on the kinetics of the thermal decomposition of 1 bromo, 2 chloro ethane is in line with

the above observations [44]. The main products are HBr and CH2CHCI formed via a radical chain:

CH2ClCH2Br->Br+CH2ClCH2

CH2CICH2ACI+CH2CH2

Br/Cl +CH2ClCH2Br^HBr/HCl+CHClCH2Br

CHClCHsBr^Br+CHsCHCl

The chain is then set up via reactions 3 and 4 and

quenched by various radical-radical reactions.

For iodo compounds, it is likely that both a radical

as well as a molecular mechanism path occur concurrently. Benson and Bose ingeniously studied the

reverse reaction adding HI to an olefin to establish

the unimolecular reaction for alkyliodide decomposition [39]. Analogous studies with olefins and

HBr or HCl were not feasible because in the temperature region where these reactions occur with reasonable rates, the equilibrium is strongly on the side

where the reactants predominate.

The pyrolysis of various organic esters of acetic

and proprionic acid in the gas phase have been

studied. The reactions are homogeneous and imimolecular resulting in an olefin and the acid as products.

0

RCOCH2CH2R'

'"V 0/ iCHR

\' I RC ^£\\2

0

RCOH + R'CH=CH2

This reaction complex is a six-membered ring for

elimination rather than the four-centered complex of

the alkyl halides.

In a study of the pyrolysis of many acetates in the

gas phase, Scheer et al. found that substitution of

alkyl groups in the a position increased the decomposition rate much more than substitution in the /3 position [45]. The effect of substituting both H's in

the a position by alkyl groups increased the rate by

as much as almost 16,000 times the rate of the completely unsubstituted compound. Interestingly, the

substitution of a second alkyl group increased the

rate by as much as 200 times the rate for the mono

alkyl substituted compound. This behavior, some- what analogous to the molecular elimination re- actions of alkyl halides, suggests very stronglv that

the activated complex is quite polar in nature. One

interesting observation is that substitution in the a

or /3 position of an electron withdrawing group (CI)

slows the rate of decomposition. This may be interpreted in terms of a large positive charge on the a

carbon atom which donates the H atom, increasing

the energy of the polar reaction complex.

There is one unanswered question: What is the

effect of the methyl group on the carbonyl carbon?

The question could be answered by studying the

decomposition of a formic acid ester relative to the

same acetic acid ester. Gordon and Norris studied the pyrolysis of

methyl-ethyl carbonate and diethyl carbonate [46].

The reaction is homogeneous and unimolecular and

probably goes through a six-membered activated

complex.

12
ROCOCH2CH3 > ROCOOH + C2H4

ROCOOH then decomposes to CO2 and ROH,

probably on the surface. Although the effect of

substituents on the carbonyl containing carbon

cannot be experimentally tested since HOCOC2H5 is

II O

not stable, this reaction by analogy also proceeds

via a polar complex. Gordon and Norris saw no

evidence that the hydrogen from the methyl group

was transferred in dimethyl carbonate pyrolysis. It

is not surprising because a more strained fivemembered complex would be involved in the re- action and the AH of the reaction would be about

65 kcal higher since no olefin would result.

Similarly the pyrolysis of Xanthates which also

decomposes via a molecular mechanism probably

occurs via a polar activated complex.

Laidler and McKenney have pyrolyzed diethyl

and dimethyl ether [47]. For dimethyl ether, they

found the reaction to be free radical initiated by a break of one of the C—O bonds. Diethyl ether was

reported as being a combination of a free radical

and a molecular elimination reaction.

H

C H OCH CH - 2 5 2 3 C H 0

2 5

CH,

-) C^H OH + C^^H,

2 5 2 4

The pyrolysis of ethylene oxide resembles the initial

step of cyclopropane pyrolysis.

CH CH > CH^CH^ > CH3CHO*

<2/ 2

I ^ ^ 0 0*

The acetaldehyde may have as much as 85 kcal

excess energy over thermal energy because of the

exothermicity of the isomerization of the biradical.

The excess energy coupled with the average thermal

energy is more than enough to break the C—C bond.

As was first noted by Neufeld and Blades and

developed by Benson, the behavior of the pyrolysis of

ethylene oxide may be rationalized as a combination

of free radical reactions induced by the decomposition

of the excited acetaldehyde and the stabilization of

the acetaldehyde by collisions [48, 49]. The two

processes explain the experimental behavior as a function of pressure and of temperature.

The Claisen rearrangement is analogous to the

Cope rearrangement differing in that an oxygen atom

replaces a carbon atom in the skeleton of the molecule. The rearrangement of vinylalkylether is a

typical example:

for which A = W^ '' sec ^ and £' = 30.6 kcal/mole

[50].

4. Pyrolysis of Saturated Linear

Hydrocarbons

Linear saturated hydrocarbons decompose by

chain reactions, initiated by the unimolecular

decomposition of the parent molecule into two

radicals. For a chain reaction, the other elementary

reactions in the mechanism can involve imimolecular,

bimolecular and, for small radical-radical combinations, termolecular reactions.

The pyrolysis of a saturated hydrocarbon may be

illustrated by ethane. Many studies of this reaction

have been conducted with the mechanism generally

agreed to be the following [51]:

C2H6-^2CH3 (1)

CH3+C2H6->CH4+C2H5 (2)

C2H5+M->C2H4+H+M (3)

. H+C2H6-^C2H5 +H2 (4)

The reaction can be termined by

H+H +M-^H2+M (5)

H+C2H5+M->C2H6+M (6a)

H+C2H5->H2 +C2H4 (6b)

C2H5+C2H5^C4Hio (7a)

^C2H6+C2H4 (7b)

13
Of the termination reactions, the bimolecular

combination of two ethyl radicals would have the

largest rate constant since none of the reactions has

any appreciable energy of activation and the other

terminations are termolecular. In addition, the

unimolecular pyrolysis of the ethyl radical has a

fairly high energy of activation and is probably in

its "fall-off" region in the temperature regime where

ethane pyrolyzes, so that reaction (3) is much slower

than reaction (4) . Thus, the ethyl radical has the

largest steady state radical concentration. Methane

has been used to monitor the unimolecular ethane

decomposition since it is completely stable in a temperature region where ethane pyrolyzes very

rapidly. In our laboratory, we have investigated the

effect of inert gases on the pyrolysis of ethane. These

gases cause the overall rate of decomposition of

ethane to increase appreciably; about 500 torr of

Xenon pressure doubles the overall rate of ethane

decomposition [52]. The decomposition of pure

ethane has been shown in many laboratories, including ours, to be about first order relative to ethane

disappearance, and the rate of formation of methane

has been shown to be very closely first order in

ethane concentration. The increase in overall rate

by the addition of inert gases may be understood in

terms of the pressure dependence of ethyl radical

decomposition, as well as any effect of slowing the

diffusion of H atoms to the wall where they may

be quenched.

In any event, the H atom steady state concentration increases because this determines the rate

of H2 production (reaction 4) which is equal to the

total rate of ethane decomposition within about one

percent. However, we have also noted that the rate

of methane increases by about ten percent under the

same experimental condition in the presence of

inert gases. There is no explanation for this in the

mechanism noted above. The most likely reaction

for the extra formation of methane is that between

H and C2H5 radicals which are not stabilized because

the C—C bond is about 85 kcal/mole, while the

C—H bond is 98 kcal/mole, leaving 13 kcal/mole

extra energy in addition to thermal energy which

must be removed by collisions before the molecule can be stabilized. Rabinovitch and Setser have

concluded from a theoretical study that there is a high probability that the excited molecule will split

into two methyl radicales [53 J. The extraneous

source of methane clouds the experimental support

for a preexponential factor of as much as lO^'' for

the unimolecular decomposition of ethane.

5. Bimolecular Mechanisms

For all of the radical chain reactions as discussed

in the preceding section, after the original unimolecular scission of the reactant the resulting radicals (or their breakdown radicals) will abstract an atom from an unreacted molecule in a bimolecular

reaction. This type of reaction is very important in

most pyrolyses of organic molecules. There is an

extensive literature of the kinetics of alkyl radicals,

particularly methyl and perfluoro methyl, abstracting H as well as halo atoms [543-

Diels-Alder reactions have been extensively

studied. They are bimolecular reactions between

two olefin molecules, one of which must be a 1,3

diolefin to conform with the Woodward-Hoffman

orbital correlation rules [4]. The activated complex

is a six-membered ring:

CH

/

R'

CH

R

, ,CH=CH CH^CH

R'"

R

R' ' CH

\

CH

CH

:hr'

:hr' "

The pyrolysis rates of both olefins and acetylenes

have been reported to be bimolecular. These re- actions are discussed in the paper by Dr. M. Back.

6. Practical Pyrolysis Applications

In the work presented above, only the initial

pyrolysis reactions were discussed. In a practical

situation, such as the pyrolysis of hydrocarbon feed

stocks to make more desirable hydrocarbons, or combustion reactions in a diffusion flame, reactions

of the initial products will give rise to other products

which will also participate, etc. All these products

will also interact with the radicals formed in the

original reactions giving rise to a most complicated

situation. With the advent of computers with large

capacities, it is possible to solve such an involved

situation. Computer programs are now becoming

commercially available [55]. The use of such

programs requires that reliable kinetics for all the

elementary steps be established. Such kinetic data is only attainable by careful studies of systems where

secondary reaction complications are eliminated or

controlled.

7. References

[1] Prescott, R. Hudson, R. L. Foner, S. N. and Avery,

W. H. J. Chem. Phys., 22, 145 (1954).

[2] Smith, S. R. and Gordon, A. S., J. Chem. Phys. 22, 1150

(1954).

[3] Rice, F. O., J. Am. Chem. Soc, 53, 1959 (1931); and

55, 3035 (1933).

14
[4] Hoffman, R. and Woodward, R. B., Accounts of Chemical

Research, 1, 17 (1968).

[5] Frey, H. M. and Walsh, R., Chem. Rev., 69, 103 (1969).

(1963).

[6] Benson, S. W. and Rose, A. N., J. Chem. Phys. 39, 3463

(1963).

[7] Marcus, R. A., J. Chem. Phys. 20, 339 (1952); G. M

Wieder and R. A. Marcus, ibid, 37, 1835 (1962); and

R. A. Marcus, ibid, 43, 2658 (1965).

[8] Kohlmaier, G. H. and Rabinovitch, B. S., J. Chem.

Phys., 38, 1692, 1709 (1963).

[9] Valance, W. G. and Schlag, E. W., J. Chem. Phys., 45,

216, 1280 (1966).

[10] The transition from the RRK to RRKM approach has

been discussed in considerable detail by D. L. Bunker,

Theory of Elementary Gas Reaction Rates, Chapter 3, Pergamon Press (1966).

[11] Lin, M. C. and Laidler, K. J., Trans. Far. Soc, 64, 94

(1968).

[12] Plaezek, D. W., Rabinovitch, B. S., Whitten, G. Z., and

Tschuikow-Roux, E., J. Chem. Phys., 43, 4071 (1965).

[13] Bunker, D. L. and Pettergill, O. M., J. Chem. Phys.,

48, 772 (1968).

[14] Marcus, R. A., J. Chem. Phys., 41, 2614, 2624 (1964).

[15] Casas, F., Previtali, C, Grotewold, J., and Lissa, E. A.

J. Chem. Soc. A, 1001 (1970).

[16] Hauser, W. P. and Walters, W. D., J. Phys. Chem., 67,

1328 (1963).

[17] Church, J. P., J. Phys. Chem., 65, 2170 (1961).

[18] Brandaur, R. L., Short, B., and Kellner, S. M. E., J. Phys. Chem., 65, 2269 (1961).

[19] Glass, D. J., Boikees, R. S., and Winston, J., Tetrahedron

Letters, 999 (1966).

[20] Rabinovitch, B. S., Schlag, E. W., and Wilberg, K. B.,

J. Chem. Phys., 28, 504 (1958).

[21] Benson, S. W., J. Chem. Phys., 34, 521 (1961).

[22] Jackson, W. M. and McNesby, J. R., J. Am. Chem.

Soc, 83, 4891 (1961); J. Chem. Phys., 36, 2272 (1962)

and W. M. Jackson, J. R. McNesby and B. deB.

Darwent, ibid, 37, 2256 (1962).

[23] Gordon, A. S., unpublished.

[24] Benson, S. W. and Nangia, P. S., J. Chem. Phys., 38,

18 (1963).

[25] Benson, J. A. and Balquist, J. M., J. Am. Chem. Soc,

90, 7343 (1968).

[26] Carter, W. L. and Bergman, R. G., J. Am. Chem. Soc,

90, 7344 (1968).

[27] Gerberich, H. R. and Walters, W. D., J. Am. Chem. Soc,

83, 3935, 4884 (1961).

[28] Baldwin, J. E., Tetrahedron Letters, 26, 2953 (1966).

[29] Benson, S. W. and Shaw, R., Trans. Far. Soc, 63, 885

(1967).

[30] Ellis, R. N. and Frey, H. M., J. Chem. Soc, A, 553

(1966).

[31] Gordon, A. S. and Norris, W. P., unpublished.

[32] Gill, G. B., Quarterly Rev., 22, 338 (1968).

[33] Longuet-Higgins, H. C. and Abrahamson, E. W., J. Am.

Chem. Soc, 87, 2045 (1965).

[34] McNesby, J., R. and Gordon, A. S., J. Am. Chem. Soc,

79, 825 (1957).

[35] Gordon, A. S., Canad. J. Chem., 43, 570 (1965).

[36] Maccoll, A., Chem. Rev., 69, 33 (1969).

[37] Pritchard, G. O., Venujopalan, M., and Grehen, T. F.,

J. Phys. Chem., 68, 1786 (1964).

[38] Maccoll, A. and Thomas, P. J., Nature, 176, 392 (1955).

[39] Benson, S. W. and Bose, A. N., J. Chem. Phys., 39, 3463

(1963).

[40] Benson, S. W. and Haugen, G., J. Am. Chem. Soc, 87,

4036 (1963).

[41] Barton, D. H. R. and Oryon, F. P., Trans. Far. Soc,

45, 725 (1949).

[42] Hewlett, K. E., Trans. Far. Soc, 48, 25 (1952).

[43] Vaughan, W. E. and Rust, F. F., J. Organ. Chem., 5, 449 (1940).

[44] King, K. D. and Swinbourne, E. J., J. Chem. Soc, B,

687 (1970).

[45] Scheer, J. C, Kooyman, E. C, and Sixma, F. L. J., Rec. Trav. Chim., 82, 1123 (1963).

[46] Gordon, A. S. and Norris, W. P., J. Phys. Chem., 69,

3013 (1965).

[47] Laidler, K. J. and McKenney, D. J., Proc Roy. Soc

(London), A278, 505 (1964); and D. J. McKenney and

K. J. Laidler, Canad. J. Chem., 41, 1984 (1963).

[48] Neufeld, M. L. and Blades, A. T., Canad. J. Chem., 41,

2956 (1963).

[49] Benson, S. W., J. Chem. Phys., 40, 105 (1964).

[50] Shuler, F. W. and Murphy, G. W., J. Am. Chem. Soc,

72, 3155 (1950).

[51] Quinn, C. P., Proc. Roy. Soc (London), A275, 190

(1963).

[52] Gordon, A. S., unpublished.

[53] Rabinovitch, B. S. and Setser, D. W., Adv. Photochem.

3, 1 (1964).

[54] Trotman-Dickenson, A. F., Adv. Free-Radical Chem.

1, 1 (1965).

[55] Bard, Y. and Ladidus, L., Catal. Rev., 2, 67 (1968).

Discussion

R. S. Konar (Catholic University):

Dr. Gordon has raised a serious qoiestion on the

vahdity of the measurement of the initiation rate in

ethane pyrolysis via the initial rate of methane

formation in the iminhibited decomposition (i.e.,

before the radical-olefin reactions set in). According

to his mechanism (p. 1 of preprints) CH3 radicals are not only produced in the reaction (1),

C2H6^2CH3, but also in the atom cracking reaction,

viz:

H+C2H5^C2H6*^2CH3 (8)

C2H6*+]V[-^C2H6+]V[. (8a)

If the reaction (8) is really significant under the

normal experimental conditions (at or above 100

torr C2H6, and at or above 540 °C), then in principle

the measurement of total CH4 at the initial stages

would not be a measure of the initiation reaction

(1). However if the initiation reaction (1) is first order, and all the elementary reactions are homogeneous, then it is not clear how an inert gas can

accelerate the rate of methane formation at a given

pressure of the reactant. On the contrary, it should

decrease the rate of CH4 formation because of the occurrence of the reaction (8a). Dr. Gordon has

found that the inert gas increases the rate of CH4

production by about 10 percent. Does he refer to the

initial rate of CH4 formation? I wonder whether

Dr. Gordon could illustrate this point in greater

detail. Incidentally, it may be relevant to mention

here that the calculations of Setser and Rabinovitz on ethane dissociation and on the recombination of

CH3 radicals (Adv. Photochem. 3, 1, 1964) do not

agree with the experimental data (Lin and Back,

Can. J. Chem. 44, 2360-61, 1966; Van Den Berg,

Callear, and Norstrom, Chem. Phys. Letters 4,

101, 1969).

15
It may be relevant to mention here that it has

been found experimentally that diallyl is a product

in the cadmium photosensitized decomposition of

cyclopropane at 265 °C (Konar and Darwent, to be

published) and it is believed that this diallyl is formed due to the interactions of allyl radicals

originated almost entirely from the isomerization of

cyclopropyl radicals produced in the system. This

confirms the observation of McNesby and Gordon

(page 11) that the isomerization of cyclopropyl radicals to allyl radicals does occur even below 400 °C, al- though this reaction is symmetry forbidden (LonguetHiggins and Abrahamson, J. Am. Chem. Soc. 87,

2045, 1965).

Alvin S. Gordon and R. H. Knipe:

Dr. Konar has raised the question of how a first order initial split of ethane and subsequent homogeneous elementary steps can effect the rate of the

cracking reaction (8). If there is any effect of

pressure he reasons that the cracking reaction should

decrease because of reaction 8(a)

CC2H6*+M-^C2H6+M].

The answer to his question lies in reaction (3)

C2H5+M-^C2H4+H + M.

If this reaction is in its "fall-off" region in the

pressure range of the studies, then increasing (M)

will increase reaction (3) , resulting in a higher steady

state concentration of H atoms. The rate of H2

formation is determined by reaction (4) and reflects

the increased H atom concentration in the presence

of inert gases, since the rate can be doubled with

sufficient inert gas addition. The ethyl radical concentration, according to the mechanism is proportional to (CsHe)"^ independent of (M). Thus,

reaction (8) increases proportional to (H) with

changes in (M) for a given (C2H6). An interesting

question which can be raised is why the rate of

methane formation seems to be first order in ethane

if the cracking reaction contributes? Here the (C2H6)

is proportional to (C2H6)^'^ and the (H) is independent of (C2H6) if no inert gases are present

since the rate of H2 is first order in (C2H6) and is formed via reaction (4) . Thus, there is a component

of CII4 rate which is proportional to (C2H6)^'^. The

fact that it does not show experimentally is because

it is not a high percentage of the total methane rate

and is absorbed in the experimental error. The above analysis is on the basis that reaction

8(a) in Konar's question plays only a minor role. It is difficult to evaluate the significance of reaction

8(a). Since the ethane molecule has about 98

kcal/mol plus its normal thermal energy and re- quires about 85 kcal/mol to break the C—C bond,

it can survive a number of collisions and still have

enough excess energy to break the C—C bond. In

the paper we have referred to calculations by

Rabinovitch and Setzer which predict that the

C2H5* molecule is so short lived that it will not be

quenched by collision. It is true, as pointed out by

Konar, that calculations on the pressure at which

recombination of methyl radicals becomes pressure

dependent are not in accord with experiments, so that calculations of this type are not exact. There

may be some quenching of the excited ethane, but

our experimental results indicate that the cracking

reaction is important in our pressure range.

16
NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Pyrolysis of Hydrocarbons

Margaret H. Back

Chemistry Department, University of Ottawa, Ottawa, Canada

Some current problems arising from studies of the pyrolysis of hydrocarbons are discussed

and two stages of these studies are distinguished. The first stage concerns the description of the

mechanism of the pyrolysis in terms of a series of elementary reactions. The types of reactions

occurring during the pyrolysis of olefins are discussed from this point of view and the effects of

temperature and pressure on these elementary reactions are shown to produce predictable changes

in the over-all pyrolysis. Some problems in the mechanism for the pyrolysis of acetylene are dis- cussed and a mechanism proposed which relates the reactions occurring in both the high and low

temperature regions. The second stage in pyrolysis studies is attained when the mechanism of the

reaction may be satisfactorily described in detail and experiments may then be designed to measure

the rate constants of the elementary processes. Several studies on the paraffins have attained this goal and the results are examined for consistency with related measurements.

Key words: Acetylene; dissociation; elementary reactions; ethane; ethylene; hydrocarbons;

kinetics; mechanism; methane; olefins; paraffins; propylene; pyrolysis; rate constants; unimolecular

theory.

1. Introduction

The mechanisms by which hydrocarbons pyrolyze

may be broadly divided into those which involve one elementary rearrangement process to give one or perhaps two products and those which include a

series of radical reactions and yield several products.

Ring compounds usually fall in the first category

[^l]^ and the large majority of straight and branched

chain hydrocarbons fall in the latter group [2]. This

discussion will be confined to the second class of

mechanisms, those involving radical reactions.

Pyrolysis studies of the latter group may be

considered to involve two stages. The first stage is the elucidation of the mechanism in terms of a series

of elementary processes. Studies on the paraffins

have, during the past several years, succeeded in

completing this stage, with the notable exception of

the simplest member of the series, methane, which

still defies explanation on several aspects. When the

elementary processes are known the system is ready

for the second stage which involves the measure- ment of the rate constants for the elementary

reactions. This stage constitutes the ultimate goal in

pyrolysis studies, as perhaps in any kinetic study,

where the measurement of an elementary rate constant and its variation with temperature and pres- sure may be considered the final achievement.

The discussion will begin with a brief review of

some of the techniques used in pyrolysis studies and

the relative merits of each. The paraffin systems will

be discussed next, as examples of studies in stage

two, and several measurements of elementary rate

^ Figures in brackets indicate the literature references at the end of

this paper.

constants will be described. Finally, the type of

reactions occurring during the pyrolysis of the

simple olefins will be described and compared with

those occurring during the pyrolysis of acetylene.

2. Pyrolysis Techniques

The techniques used for pyrolysis studies are of

three basic types, depending on the time resolution

required for the experiments. The latter in turn is determined by the region of temperature where the

reaction is studied. At temperatures up to about

900 K a static system may be used. The reactant is admitted to a vessel and the reaction allowed to

proceed for a measurable time. For shorter reaction

times, from approximately 1 to 0.01 seconds, a flow

system is used and temperatures commonly range

from about 900-1200 K. The technique of shock tube

pyrolysis generally produces temperatures from

1000-2000 K and detection methods may have time

resolutions as low as 10 jus. Each method has its own limitations. The static

system provides the most controllable conditions of

temperature and reaction time but is most sus- ceptible to surface reactions. The main problem in a

flow system is the definition of the temperature

profile and hence the time of the reaction. There is also the difficulty in obtaining a thorough test of the

order of the reaction, since wide variations in the

concentration of reactant are not usually feasible. An

interesting flow system has been developed in

which the reacting gas is surrounded by a shell of

inert gas, thus isolating the reactant from the walls

of the reaction vessel. Measurements of the rate of

the initiation step in the pyrolysis of neopentane were

17
reported and shown [4] to be in reasonable agreement with previous studies in a static system [5].

Another interesting technique which has proved

useful in measurements of elementary dissociation

reactions and which may provide data on other

fundamental processes is the very low pressure

pyrolysis method developed by Benson and colleagues [6]. This technique uses a flow system with

pressures sufficiently low that molecules are excited

primarily by collisions with the wall. The rate of

flow and hence the steady state concentration of

molecules in the reactor depend on the rate of

effusion from the orifice and can be accurately

calculated. Measurements can be made over a wide temperature range and the method should

provide useful tests for unimolecular theories.

Several measurements of elementary rate constants

have been made, among them the rate constants for

dissociation and combination of allyl radicals [7], a result difficult to obtain by other methods.

Reactions in shock tubes are free of heterogeneous

complications but have encountered difficulties in

defining the temperature exactly and in describing

the mechanism satisfactorily. Nevertheless several

reliable measurements of rate constants for elementary reactions have been reported using the shock

tube technique and among the hydrocarbon studies

may be cited those of Tsang [8]. The special comparative technique used by Tsang allows the ac- curate measurement of relative rates but the absolute

values depend on the accuracy of the measurements

for the standard.

3. Parameters obtained from Pyrolysis

Measurements

The measurement of the rate constant for an

elementary reaction as a function of temperature

gives an Arrhenius plot from which the activation

energy and the frequency factor may be determined.

In order to obtain the enthalpy change of the re- action the activation energy of the reverse reaction

must be known. If the reverse reaction is a radicalradical reaction and its activation energy negligible

then the measured activation energy will be within

RT of the bond dissociation energy. The frequency

factor gives an indication of the entropy change

associated with the formation of the activated

complex and allows some deductions about the

structure of the complex.

The interpretation of the kinetic parameters in

this way depends on the assumption that the re- action occurs homogeneously. For a reaction oc- curring on the surface the parameters would have

quite different meanings. The problem of surface

reactions is a thorny one for most gas kineticists.

The effects of heterogeneous reactions are often quite

impredictable and the various tests and treatments

which have been used are sometimes ambiguous.

The only case for which a meaningful test may be

made is for the measurement of an elementary

reaction. When the rate of an elementary reaction is measured any change in conditions of the reaction

will be immediately observable in the measurement

if such changes affect the rate. If the concentration

of surface sites is increased (as is usually done by

packing the vessel) and the rate is vmaffected then it may be concluded that the particular elementary

reaction does not occur on the surface. Such tests may be quite ambiguous if a series of elementary

reactions is involved, since each reaction may be

affected in a different way at the surface and no

net change may occur [9^.

The comparison of rate constants over a wide

range of temperature can provide useful information

for tests of kinetic theories. The combination of

results from the three techniques gives data spanning

a very large temperature range. Of particular

importance is data on rate constants in the pressure

dependent region. The existing theories, in particular

the RRKM formulation [lOJ, are able to describe

adequately the variation of the first order rate

constant with pressure [11'], but the way in which

the fall-off curves change with temperature provides

a more stringent test for theories of energy transfer.

Comparison of mechanism over a wide range of

temperature can also be useful. Most mechanisms

will be significantly altered by a change in temperature of one or two hundred degrees and if such

changes can be predicted and confirmed experimentally, the mechanism may be well established.

4. Pyrolysis of Paraffins

Studies of the pyrolysis of the simple paraffins,

with the exception of methane, have brought our

understanding of these systems to a reasonably

satisfactory stage. Much of this work has been done

by Purnell, Quinn, and their co-workers at Cambridge and Swansea and by Niclause and colleagues

at Nancy, France. The situation may be illustrated

and summarized by the most thoroughly-studied

paraffin, ethane |]12-16J. The series of reactions

which occur in the initial stage of the pyrolysis of

ethane in the temperature range 800^900 K are

as follows:

C2H6—^2CH3

CH3+C2H6—^CH4+C2H6

C2H5—>C2H4+H

H+CaHe— +C2HB

2C2H5—>C4Hio or C2H4+C2H6

(1)

(2)

(3)

(4)

(5a) or

(5b)

In this type of chain mechanism, initiation occurs by unimolecular dissociation of the reactant, the

rate-controlling propagation step is the dissociation

of a radical, and termination occurs by the bimolecular combination of radicals.

Similar mechanisms have been suggested for the

pyrolysis of propane [11], butane [183, isobutane

[19], neopentane [20, 21] and some Ce to Cg alkanes

1^22]. As the hydrocarbon becomes larger the number

and variety of radicals increases and the variations on the propagation reactions, isomerization for

18
X Quinn A Gordon O Lin a Back • Trenwith

1.15 1.20

I / T x lO'

Figure 1. Arrhenius Plot of ki.

example, can cause considerable complexity in the

mechanism. Without summarizing all the arguments

which led to the acceptance of these mechanisms I shall instead describe the measurements which have

been made of some of the rate constants of these

elementary reactions.

In the pyrolysis of ethane the following relations

may be derived by applying the steady state treat- ment.

CHiRH2 A-3

\l/2 {Rcm.V" (ha)

where R refers to the initial rate of formation of

product. Measurements of the rate of formation of

methane allow the determination of ki and measure- ments of the rates of formation of hydrogen and

butane allow the determination of ks, using the

previously measured value for k&a [23].

5. Measurements of ki

Several measurements of the initial rate of

formation of methane have been made as a function

of temperature and of pressure. Considering the

temperature variation first, figure 1 shows an

Arrhenius plot of measurements by various workers.

The maximum deviation from the line drawn is 40 percent and most of the measurements are within

20 percent. These measurements were made in static

systems and the yield of methane as a function of

time was extrapolated to zero time to obtain the

initial rate. In spite of this close agreement the

Arrhenius parameters obtained from each study

show much wider variations.

Log A E

Kcal/mole Reference

Temperature

range

16.3 86.0 Davis and Williamson 948-1048

17.45 91.7 Quinn 837-881

14.5 81.0 Gordon 748-873

16.00 86.0 Lin and Back 823-893

16.3 88.0 Trenwith 839-873

The rate of dissociation of ethane has also been

studied as a function of pressure [24, 16]. The first order rate constant begins to decrease at pressures

below about 100 Torr at temperatures above about

880 K. The experimental results have been interpreted with the RRKM theory [25] and a reasonable

model for the activated complex has been found to

give the observed variation of A:i with pressure.

This dissociation reaction is important not only

because it is the basic C-C bond but also because

many measurements of the reverse reaction, the

combination of methyl radicals, have been made,

and since the equilibrium constant may be calculated with some precision the two rate constants

may be compared over a wide range of temperature.

The rate of combination of radicals has been measured in the range 300-776 K while the dissociation

of ethane has been measured between 820 and

1050 K. Two aspects of this comparison are important: first the absolute magnitude of the rate constants in the high pressure region, and secondly the

characteristics of the change in the first order rate

constant with pressure.

Measurements of k-\ in the high pressure region

in the temperature range 300-500 K gave a value of

2.2.x 1013 m-'s-i [26, 27], apparently independent of temperature. An average of ki obtained

from the pyrolysis of ethane,

1 / 88000

log Ai= 16.5 -2^3^

together with the equilibrium constant gave a value

for at 500 K of 3.8X10^2 ^^^s m-h-' [24]. The

discrepancy of a factor of 6 is outside the estimated

errors in each of the quantities ki, k-i, AS° or AH".

The four measurements of Ai in the range 820-870 K

were within 50 percent of a mean and if this range is fixed for ki one cannot obtain a frequency factor for

ki high enough to give agreement with k-i without

obtaining an unreasonably high value for Ei and

hence AH. The value of A obtained by Quinn [13]

does agree with the measured value of but the

activation energy of 91.7 kcal/mole implies a heat of

formation of the methyl radical of 35.6 kcal/mole

at 298 K which requires a value for D(CH3-H) of

105.6 kcal/mole. It may be pointed that that any

positive temperature coefficient of would make

its value even greater at the temperatures at which

ki was measured.

A similar discrepancy was observed in the study

of neopentane. The average value for the rate constant for the dissociation of neopentane,

C5H12—>CH3-|-t— C4H9

from studies in the ranges 723-803 K [20] and

1070-1245 K [28] gave a rate constant for the

reverse combination of radicals which was about

10 times less than the measured value for the combination of methyl radicals. The rate constant

measured over the range 713-823 K [21], however,

did not show this discrepancy but gave very good

agreement with measurements of radical combinations.

19
LOG T

Figure 2. Combination of methyl radicals; variation of the medium fall-off pressure with temperature.

[24]; A [161; X [30J; V [31]; + [32]; A [27]; [26]; [33]; O [29]; • [34].

It may be argued that the extrapolation of the

rate constants measured in pyrolysis studies into a region of temperature so far from where they were measured carries with it enough error to account for

the discrepancy observed. This is possible but still implies more error in the four quantities than is currently admitted. Such a wide temperature range

is an exacting test for the validity of rate constant

measurements and may help to reveal sources of

error which might otherwise have remained hidden.

The measurements of the pressure dependence of

the combination of methyl radicals over the temperature range 300-1000 K have been summarized

by Casas, Previtali, Grotewold, and Lissi [293-

Their summary of the variation with temperature

of the pressure at which the first-order rate constant

has fallen to half its high pressure value is reproduced

in figure 2. Grotewold and colleagues have shown

that a transition state in which the free rotation of

the methyl radicals, suggested by Rabinovitch and

Setser [11 J and designated by the dotted line, has

been replaced by low frequency bending, designated

by the solid curve, gives better agreement with their

results at low temperature and also with the measure- ments in the higher temperature range. This model

also reproduces the general shape of the curve of rate

constant as a function of pressure obtained for the

ethane dissociation. The values of pi/2 obtained from

the ethane pyrolysis are, however, still disturbingly

far from the values obtained from the combination

of methyl radicals.

The rate constants for the dissociation of isobutane

and of neopentane have also been measured using

similar techniques, and the results summarized

as follows:

E

log A (s~i) (kcal/mole) Reference

Isobutane 17.8 82.5 19

Neopentane 18.1 85.9 21

16.8 82.0 20

The high frequency factors appear to be characteristic

of these dissociation reactions.

6. Radical dissociations

It had been suggested that the dissociation of the

ethyl radical was pressure dependent in the temperature region of the ethane pyrolysis studies [13].

Measurements of the yields of hydrogen and butane

from the pyrolysis of ethane confirmed this [15] and

the rate constant obtained, extrapolated to the high

pressure limit, agreed with that from the mercury

photosensitized reaction of ethane over a lower

temperature range [35].

The rate constant for the dissociation of the

isopropyl radical into ethylene and a methyl radical

was obtained from the study of the isobutane pyrolysis [36] and compared satisfactorily with the value

measured in a lower temperature region [37].

Dissociation modes requiring higher activation

energies are more easily measured in the higher

temperature region of pyrolysis experiments while

the rate requiring the lower activation energy is usually obtained from photolysis experiments in a lower temperature region.

These rate constants, of which only a few examples

have been given, have been obtained with techniques

and measurements which are not difficult. It is encouraging that simple methods can yield reliable

and useful data.

7. Secondary reactions

The discussion so far has considered only those

reactions occurring in the initial stages of the

reaction, which means those reactions which do not

involve any of the products. In the paraffin systems

the olefin products react rapidly with radicals,

producing secondary products and adding many

new reactions to the mechanism [38]. This is a general consequence of forming products more

reactive than the reactant and has caused considerable trouble in the studies of the paraffins. Most

of the rate constants for the elementary reactions

have been measured in the region where secondary

reactions are not serious. Some attempts have been

made to treat systems where the rates of formation

of the products are changing with time [39], but

because such systems generally involve a large

number of elementary reactions very little quantitative data on rate constants can be obtained. Computer solutions to the equations may be the answer

to this type of problem, although such solutions

depend heavily on accurate data for the rate constants involved.

8. Methane

Since methane is the least reactive of the paraffins

secondary reactions of the products become important at very early stages in the reaction, making

studies of the primary reaction very difficult. The

20
overall course of the reaction may be written as

follows:

2 CH4^C2H6+H2^C2H4+H2^C2H2+H2^C +H2

but the series of elementary reactions by which the

products form has not been established [^40]. No

satisfactory measurement of the initiation step has

been made in the lower temperature range of the

studies in static and flow systems. Indeed, it has not

been possible to identify the initiation step, partly

because the two most likely possibilities,

CH4^CH3+H

CH4^CH2+H2

have very similar endothermicities. Studies in shock

tubes [41, 42] have made the closest approach to

measurements of the initiation step. Either of these

initiation steps ought to show pressure dependence

at fairly high pressures and hence could provide

some interesting data for unimolecular theories. For

identification and measurement of the initiation step

the rate of formation of the primary products must

be measured before secondary reactions become

important. Perhaps the best hope for this is from

the improved sensitivity of techniques for measuring

products which may make it possible to study the

decomposition of methane in the lower temperature

range.

9. Pyrolysis of olefins

The thermal reactions of olefins are more complex

than those of the paraffins but the general features of

the mechanisms now seem moderately well understood. Both decomposition and condensation processes occur, yielding a wide range and variety of

products. Induction periods are usually observed, at

least in the lower temperature region. The kinetics are thus complex and quantiative measurements

of elementary rate constants have not been reported

from studies of the olefins.

Early work [2 J on the pyrolysis of ethylene

established that the rate of disappearance of ethylene

was second order and the activation energy had the

low value of 40 kcal/mole. The major products were propylene, butene-1, butadiene, ethane and

polymer. A variety of minor products were reported

depending on the extent of the reaction. Because an

induction period was observed primary and secondary

products were not always distinguished and the

mechanisms suggested were somewhat speculative.

The relative importance of radical and molecular

processes was also uncertain.

Two recent studies of the thermal reactions of

ethylene in the lower temperature region, 750-900 K,

have provided a general framework for the mechanism of simultaneous polymerization and decomposition, although the results appeared at first to be

conflicting. The first study by Halstead and Quinn

(43), which covered the range 798-923 K and

10-270 torr, reported that all products showed self- acceleration and no changes in total pressure were

observed. They suggested the reaction was a degenerately branched chain reaction with butene-1

as the chain branching agent. The main chain

propagating radicals were ethyl and vinyl and

addition, decomposition and abstraction reactions

of these radicals were able to account quantitatively

for the time dependence of the yield of the major

products.

The second study by Boyd, Wu, and Back [44]

at somewhat higher pressures (150-600 torr) and

lower temperatures (773-873 K) showed that ethane

and a polymer were the only primary products and

all other products were formed with an induction

period. The formation of the polymer was deduced

from the discrepancy between the measured pressure

change and the pressure calculated from the measured

products. The rate of formation of ethane was

approximately second order with an activation

energy of 42 kcal/mole. The mechanism suggested

was a radical chain polymerization initiated by the

bimolecular reaction of ethylene to give a vinyl and

an ethyl radical. Initially the only two products were ethane and polymer.

C2H4 +C2H4^C2H3 + C2H5 (1)

C2H5 +C2H4—>C2H6 +C2H3 (2)

-^C4H9 (3)

C4H9+C2H4^C6Hi3 (4)

C„H2„+1 +C2H4-^C2H6 +C„H2„ (5)

C2H3+C2H4^C4Hv (6)

C4H7 +C2H4—^-CeHii (7)

C„H2n_l +C2H4—>C2H5 +CftH2n-2 (8)

2C2H5—>C4Hio

Subsequent decomposition of the unsaturated product formed in (5) and (8) led to the formation of

the other products and the induction period observed

was caused by the accumulation of sufficient unsaturated product in the system. The rate of formation of ethane, for a long chain process, is

The endothermicity of the initiation process is 65

kcal/mole so that the activation energy for formation

of ethane is predicted to be about 40 kcal/mole, in

satisfactory agreement with the experimental result.

The two mechanisms can be seen to be essentially

complementary when the effect of pressure on each is considered [45]. The rate of the bimolecular initiation is proportional to the square of the concentration

of ethylene, and at low pressures the secondary

dissociation of higher molecular weight products

will compete at an early stage. Also at low pressures

the reversal of the addition steps will become increasingly important, the average molecular weight

of the polymer will decrease and the yield of ethane

in the initial stages will decrease. Thus the char21
acteristics of the second study would change over to those of the first study as the pressure was lowered.

The key difference between the two mechanisms is the extent of the polymerization that occurs before

the dissociation of radicals competes with the

addition, or dissociation of products creates a new

source of radicals.

At the high temperature encountered in shock tube

studies ^46] polymerization does not occur, the

vinyl and ethyl radicals rapidly dissociate and the

major products are acetylene, hydrogen and butadiene. At first glance it might appear that a logical

extrapolation of the low temperature radical chain

mechanism could account for these products.

2C2H4—>C2H3 +C2H5

C2H3—>^C2H2 +

C2H6—>C2H4 +

H+C2H4-^H2 +C2H3

2C2H3—>C4H6

The observed activation energy of the rate (50

kcal/mole) is, however, lower than required by this

mechanism (65 kcal/mole). A molecular process

could occur at the high temperature of shock

pyrolysis but would not explain the rapid formation

of HD in experiments with mixtures of C2H4 and

C2D4 (47).

10. Propylene

Polymerization occurs in the pyrolysis of propylene

below about 550 °C but by 600 °C decomposition is the major process. The main products in the latter

region are methane and ethylene and a variety of

minor products of higher molecular weight is also

formed [48]. As in the ethylene system an induction

period was observed, and a radical chain mechanism

appeared very probable. A recent study [49] showed

that in the lower temperature region (740-820 K)

in the initial stages of the reaction the distribution of

products was simplified to methane, ethylene,

propane and a group of Ce compounds, and that

initial rates of formation of these products were

measurable before the onset of secondary reactions.

The initial rates of formation of the Ce products were

second order in propylene with activation energies of

between 45 and 50 kcal/mole. This suggested that

the bimolecular reaction of propylene was the

initiation process.

2C3H6—^C3H5 +C3H7

and addition and abstraction processes of the allyl

and propyl radicals could form the Ce products

observed. It was also shown that the rate of (1) was

comparable to the rate of formation of the major

products and it was concluded that the reaction in

the initial stages was a radical non-chain reaction. In

agreement with this, the activation energy for the

rate of formation of the Ce products was close to the

activation energy for reaction (1) of 52 kcal/mole.

The reason for the difference in chain length from

that in the ethylene system is that the bimolecular

initiation reaction of propylene is much faster than

the corresponding reaction of ethylene, and since the

propagation reactions have similar rates the chain

length should be much shorter in the propylene

system. As in the ethylene system the dissociation

of higher molecular weight products contributed to

secondary initiation and the rate of formation of the

more stable products increased rapidly after the

induction period.

The reactions in olefin systems are dominated at

low temperatures by addition reactions and as the

size of the radical grows isomerization processes may

occur, leading to a wide variety of products. In the

propylene system 1-methylcyclopentene and 3-methylcyclopentene were major components in the Ce

fraction. It was suggested that cyclisation of the

CeHii radical took place forming the methyl cyclopentyl radical, which then lost a hydrogen atom to

form methylcyclopentene.

CH3

\

CH—CH2

\

CH2

/

H2C = CH

CH3

\

CH—CH2

/ \

H2C CH2

. /

CH

In the ethylene reaction benzene was a major

secondary product and in this case the favourable

cyclization process of the Ce radical would be the

formation of the Ce ring rather than the C5 ring.

A question arises concerning the reaction path for

the bimolecular reactions of olefins. The process may

be a simple hydrogen transfer reaction similar to

disproportionation reactions of radicals, or it may

involve the intermediate formation of a cyclobutane

ring.

2C2H4—>C2H3 +C2H6

2 C2H4—>c — C4H8

The rates of these two processes may be calculated

using known thermodynamic and kinetic data [50].

In the case of ethylene the rate of formation of

cyclobutane is about 300 times faster than the rate

of formation of vinyl and ethyl radicals at 800 K

and 600 torr, but in the case of propylene the two

rates are almost identical. Whether cyclobutane

may be an intermediate in the bimolecular reaction

depends both on its rate of formation and on its rate

of decomposition to yield the radicals formed by the

reaction of the olefin. The simple cyclobutanes

decompose almost exclusively by a molecular

mechanism [51] and homogeneous dissociation into

radicals has not been quantitatively measured. An

empirical upper limit to the fraction of the decomposition of cyclobutane which gives radicals below

733 K was given as 0.1 percent. This makes the rate

of formation of radicals from cyclobutane about three

times less than the rate of formation of radicals from

two molecules of ethylene at about 750 K. Given

22
the errors involved in measurement and calculations,

it is just possible that radicals are formed through

cyclobutane at this temperature.

In the propylene system the rates of the two

reactions are about equal and since the formation of

radicals in the decomposition of dimethylcyclobutane

[523 was no more important than in the decomposition of cyclobutane the cyclic intermediate

cannot contribute significantly to the formation of

radicals in this system.

11. Acetylene

The mechanism of the acetylene pyrolysis has

provoked much controversy, partly because of the

difficulty of obtaining reliable data. Modern analysis

has overcome much of this trouble and there is now

general agreement concerning the nature of the

products under various conditions (53). Polymerization is dominant at low temperatures and as the temperature is increased ethylene, ethane,

hydrogen, methane, vinylacetylene, diacetylene,

benzene are among the products found in varying

yields depending on temperature and conversion. In

several aspects the acetylene pyrolysis bears a strong

resemblance to the ethylene pyrolysis. The activation

energy of the reaction is low, between 45 and 50

kcal/mole, and induction periods are observed. Both

the initial rates and the maximum rates after the

induction period are second order in acetylene. Most

of the mechanisms suggest (54) that the initial step

is a bimolecular reaction of acetylene to give an

excited molecule of acetylene in the triplet state or a

triplet or biradical species C4H4* [55]. The rate of

formation of excited molecules in thermal systems,

however, is somewhat uncertain; the frequency factor

may be very low.

The bimolecular reaction of acetylene to give a vinyl and an ethynyl radical has been suggested

[56], but has not been followed up with a complete

mechanism. With this initiation the following

reactions represent an outline of the most probable

types of radical reactions which may occur during

the pyrolysis of acetylene.

High

temperature

Intermediate

temperature

C2H2 +C2H2—>C2H +C2II3

C2H3—>C2H2 +

C2H +C2H2—>C4H3

C4H3^C4H2 +H

H +C2H2—>H2 +C2H

2C2H—>C4H2

C4H3 +C2II2—>C4H4 +C2H

C2H3+C2H2

>C4Hb

C4H5—>C4H4-|-H

C2H3 +C2H2—>C2H4 +C2H

(1)

(2)

(3)

(4)

(5)

(6)

(7)

(8)

(9)

(10)

Low

temperature

2C4H3—^^CgHe

2C2H3—>C4H6

C4H3 + C2H2^C6Hb

C4H6 +C2H2—^CeHy

(11)

(12)

(13)

(14)

A mechanism similar to this in many respects was

recently suggested by CuUis and Read [57]. They

suggested initiation occurred by the following

reaction

2C2H2^C4H3+H (15)

Kinetically these two initiation reactions cannot

be distinguished but reaction (1) is perhaps the

simpler.

The products formed depend on the temperature

region where the reaction occurs. At high temperatures, as found in shock tubes, decomposition will

compete with addition and the C4H3 radical will

disappear to form C4H2+H. The vinyl radical will be

largely dissociated. This region is illustrated by the

work of Gay, Kistiakowsky, Michael, and Niki

[58]. At lower temperatures the C4H3 radical may

undergo abstraction to form C4H4, and reactions of

the vinyl radical will become important. In this

region vinylacetylene, ethylene and butadiene are products. At still lower temperatures further addition

reactions will build up high molecular weight

products.

Taking AHXC2H) as 114 kcal/mole [59],

A/f/CCsHs) as 64 kcal/mole, S^CCzH) as 51. e.u. [60]

and log A:_i(cm^ m~'s~^) = 12.5 [23], the rate constant

for reaction ( 1 ) may be expressed as

logA:i(cm^m ^s ')=15.0 —

70000

2.3i?r

Assuming a long chain reaction the rates of formation

of the products may be expressed as follows.

(A) High temperature region:

Rate C ,H..(|)- (C2H2)2

Taking reasonable values for reaction (3) [61], the

activation energy of the rate is 40 kcal/mole and the

frequency factor is 3 X 10^- cm^ m~^s~'. The activation

energy is in reasonable agreement with the experimental value. The frequency factor is about a factor

of 20 lower than suggested by Gay et al. [58], who

averaged the results of ten different investigations

but still fairly close considering the oversimplification

of the mechanism.

(B) Intermediate temperature region:

Vinyl acetylene may be formed by either of the

sequences (3) followed by (7) or (8) followed by

(9), depending on the relative concentration of the

C2H3 and C2H radicals. In either case the rate- controlling propagation step will be the addition of

the radical to acetylene.

Rate C (C2H2)^

Thus in these two limiting cases represented by the

23
dominant formation of C4H2 or by C4H4 the rate

expressions are the same or very similar.

(C) Low temperature region:

The primary products in this region are somewhat

more uncertain but a general expression similar to

those for the other regions may be written

Rate polymer ~ ( ^ ) ( C2H2 ) ^

where ka is an average rate constant for addition and

kt represents the termination rate constant for the

radicals involved.

Such a mechanism can therefore lead to a rate

expression which will remain essentially the same

over a wide range of temperature. The requirements

are only that the addition of radicals to acetylene is the rate-determining step and that the termination

steps involving the various radicals have similar rate

constants. The mechanism is intended to show, not

that excited molecules of acetylene may not be

formed, but that there is no need to postulate their

formation. First, the bimolecular initiation step can

account for the low activation energy. Secondly, the

rather erratic effects of surfaces observed may be

expected in a long chain reaction where the propagation reaction of radicals may occur following adsorption on surface sites. Thirdly, the importance of

cyclic compounds in the products, particularly

benzene, is fully in accord with the present knowledge

of the rapid cyclisation processes of radicals containing six or more carbon atoms. Earlier work,

having suggested that polymerization occurred by

condensation reactions of acetylene to form linear

molecules or diradicals, was concerned over the

difficulty of forming cyclic compounds from such

species. The formation of products by the radical

mechanism removes this difficulty since the addition

reaction destroys the sp hybridization and the sphybridization will give bent radicals. Fourthly, the

formation of methane and compounds containing an odd number of carbon atoms may be most easily

accounted for by the dissociation of radicals. For

example,

C2H2+C2H3^CH2=CH—CH=CH

CH2=CH—CH=CH^CH3—CH—C=CH

The terminal methyl radical may split off and

form methane by abstraction. Similar isomerizations

may occur with larger radicals and other radicals

containing an odd number of carbon atoms may be

formed. Finally, the radical mechanism accounts

adequately for the products formed when halogens

are added to the reacting acetylene [62 J.

12. The induction period.

As in the reactions of ethylene and propylene the

higher molecular weight products are more unstable

than the reacting olefin and as these accumulate

their dissociation soon contributes appreciably to the

reaction. In the reaction of ethylene the rates of

formation of the main products formed after the

induction period were second order in ethylene and

this was interpreted to show that the main initiation

both in the initial and in the later stages was the

bimolecular reaction of ethylene [44j. The dissociation of the high molecular weight products

contributed new products and new propagation

steps, but not significantly increased initiation. This

is probably the situation in the acetylene pyrolysis,

where the kinetics are second order both in the initial

stages and in the region of maximum rate. In both

regions the main initiation step will be the bimolecular reaction of acetylene to give the vinyl and ethynyl

radicals.

13. Reaction of C2H2 and D2

A related reaction with some puzzling features is the H-D exchange in acetylene. Experiments in

shock tubes with mixtures of C2H2 and C2D2 [58]

and mixtures of C2H2 and D2 [63] had shown that

the isotopic exchange in acetylene was orders of

magnitude faster than the conversion to other

products. The mechanism suggested, in line with

similar studies with other molecules, that exchange

occurred through dissociation of a molecular complex

between C2H2 and D2. An atomic exchange mechanism for the reaction was proposed by Benson and

Haugen [64], but several difficulties with this were

pointed out [65], the main ones being the dependence

of the rate on the acetylene concentration, the low

activation energy and the time required to reach the

steady state radical concentration.

A consideration of the mechanism suggested here

for the reactions occurring during the pyrolysis of

acetylene would lead to the prediction that exchange

of H and D in acetylene would occur very rapidly by

the addition of atoms to acetylene followed by rapid

reversal. Exchange would not require stabilization of

the vinyl radical, which would undoubtedly be a

third order reaction under shock tube conditions. If

initiation occurs by the bimolecular reaction of

acetylene instead of by the similar reaction of

acetylene and Do, an atomic exchange mechanism

can be written which correctly predicts the observed

order of the rate with respect to acetylene and

deuterium and the observed activation energy. At

1500 K the former initiation is about six times faster

than the latter, for equimolar mixtures.

The main difficulty with an atomic mechanism is that the time required for the steady state concentration of radicals to be attained (about 500 /xs) is long compared to the time resolution of the experiments (10 /us) , but no induction period was observed.

This discrepancy is puzzling but must be resolved

before an atomic mechanism could be accepted.

14. References

[1] H. M. Frey. Adv. Phys. Org. Chem. 4, 147 (1966).

[2] E. W. R. Steacie. Atomic and Free Radical Reactions.

Reinhold (1954).

[3] J. E. Taylor, D. A. Hutchings and K. J. Freeh. J. Am.

Chem. Soc. 91, 2215 (1969).

24
[4] F. Baronnet and M. Niclause. Compt. Rend. Acad. Sc. Paris. 269C, 513 (1969).

[5] J. Engel, A. Combe, M. Letort and M. Niclause. Compt.

Rend. Acad. Sc. Paris, 244, 453 (1957).

[6] S. W. Benson and G. N. Spokes. J. Am. Cham. Soc. 89,

2525 (1967).

[7] D. M. Golden, N. A. Gac and S. W. Benson. J. A.m

Chem. Soc. 91, 2136 (1969).

[8] W. Tsang. J. Chem. Phys. 46, 2817 (1967).

[9] B. W. Wojciechowski and K. J. Laidler. Trans. Faraday

Soc. 59, 369 (1963).

[10] G. M. Wieder and R. A. Marcus. J. Chem. Phys. 37,

1835 (1962).

[11] B. S. Rabinovitch and D. W. Setser. Adv. in Photochemistry. 3, 1 (1964).

[12] H. G. Davis and K. D. Williamson. Fifth World Pe- troleum Congress (1959) Section IV, p. 37.

[13] C. P. Quinn. Proc. Roy. Soc. A275, 190 (1963).

[14] A. S. Gordon, C.I.C. Symposium on the Kinetics of

Pyrolytic Reactions. Ottawa (1964).

[15] M. C. Lin and M. H. Back. Can. J. Chem. 44, 505 (1966).

[16] A. B. Trenwith. Trans. Faraday Soc. 62, 1538 (1966).

[17] D. A. Leathard and J. H. Purnell. Proc. Roy. Soc. A305,

517 (1968).

[18] J. H. Purnell and C. P. Quinn. Proc. Roy. Soc. A270,

267 (1962).

[19] R. S. Konar, J. H. Purnell and C. P. Quinn. Trans. Faraday Soc. 64, 1319 (1968).

[20] F. Baronnet, M. Dzierzynski, R. Martin and M. Niclause.

C. R. Acad. Sc. Paris, 267, 937 (1968).

[21] M. P. Halstead, R. S. Konar, D. A. Leathard, R. M.

Marshall and J. H. Purnell. Proc. Roy. Soc. A310, 525

(1969).

[22] F. Done and G. Guiochon. J. Phys. Chem. 73, 2804

(1969).

[23] A. Sheepp and K. O. Kutschke. J. Chem. Phys. 26, 1020

(1957).

[24] M. C. Lin and M. H. Back. Can. J. Chem. 44, 2357

(1966) . [25] M. C. Lin and K. J. Laidler. Trans. Faraday Soc. 64, 79

(1968).

[26] G. B. Kistiakowsky and E. K. Roberts. J. Chem. Phys.

21, 1637 (1953).

[27] R. E. Dodd and E. W. R. Steacie. Proc. Roy. Soc. A223,

283 (1954).

[28] W. Tsang. J. Chem. Phys. 44, 4283 (1966).

[29] F. Casas, C. Previtali, J. Grotewold and E. A. Lissi.

J. Chem. Soc. (A) 1001 (1970).

[30] M. T. Liu and K. J. Laidler. Can. J. Chem. 45, 157

(1967) . [31] K. J. Hole and M. F. R. Mulcahy. J. Phys. Chem. 73,

177 (1969).

[32] L. F. Loucks. Can. J. Chem. 45, 2775 (1967).

[33] S. Toby and B. H. Weiss. J. Phys. Chem. 68, 2492

(1964).

[34] J. Grotewold, E. A. Lissi and M. G. Neumann. J. Chem.

Soc. (A) 375 (1968).

[35] L. F. Loucks and K. J. Laidler. Can. J. Chem. 45, 2795

(1967).

[36] R. S. Konar, R. M. Marshall and J. H. Purnell. Trans. Faraday Soc. 64, 405 (1968).

[37] C. A. Heller and A. S. Gordon. J. Phys. Chem. 62, 709

(1958).

[38] M. C. Lin and M. H. Back. Can. J. Chem. 44, 2369

(1966).

[39] R. M. Marshall. Int. Sym. on Gas Kinetics. Szeged,

Hungary (1969).

[40] H. B. Palmer, J. Lahaye and K. C. Hou. J. Phys. Chem.

72, 348 (1968).

[41] G. B. Skinner and R. A. Ruehrwein. J. Phys. Chem. 63,

1736 (1959).

[42] V. Kevorkian, C. E. Heath and M. Boudart. J. Phys.

Chem. 64, 94 (1960).

[43] M. P. Halstead and C. P. Quinn. Trans. Faraday Soc.

64, 103 (1968).

[44] M. L. Boyd, T.-M. Wu and M. H. Back. Can. J. Chem.

46, 2415 (1968).

[45] M. Simon and M. H. Back. Can. J. Chem. 47, 251 (1969).

[46] G. B. Skinner and E. M. Sokolovski. J. Phys. Chem. 64,

1028 (I960).

[47] I. D. Gay, R. D. Kern, G. B. Kistiakowsky and H. Niki.

J. Chem. Phys. 45, 2371 (1966).

[48] A. S. Kallend, J. H. Purnell and B. C. Shurlock. Proc. Roy. Soc. A300, 120 (1967).

[49] M. Simon and M. H. Back. Can. J. Chem. 48, 317

(1970).

[50] M. H. Back. Int. J. Chem. Kinetics, 2, 409 (1970).

[51] R. W. Carr and W. D. Walters. J. Phys. Chem. 67,

1370 (1963).

[52] H. R. Gerberich and W. D. Walters. J. Am. Chem. Soc.

83, 4884 (1961).

[53] C. F. Cullis and N. H. Frankhn. Proc. Roy. Soc. A280,

139 (1964).

[54] G. J. Minkoff. Can. J. Chem. 36, 131 (1958).

[55] H. B. Palmer and F. L. Dormish. J. Phys. Chem. 68,

1553 (1964).

[56] F. C. Stehling, J. D. Frazee, R. C. Anderson. Eighth

International Symposium on Combustion. Cal. Inst. Tech. (1960).

[57] C. F. Cullis and I. A. Read. Trans. Faraday Soc. 66, 290

(1970).

[58] I. D. Gay, G. B. Kistiakowsky, J. V. Michael and H.

Niki. J. Chem. Phys. 43, 1720 (1965).

[59] A. Williams and D. B. Smith. Chem. Rev. 70, 267 (1970).

[60] estimated as outlined by S. W. Benson, Thermochemical

Kinetics, Wiley (1969).

[61] J. A. G. Dominguez and A. F. Trotman-Dickenson. J. Chem. Soc. 940 (1962).

[62] M. S. B. Munson and R. C. Anderson. J. Phys. Chem.

67, 1582 (1963).

[63] K. Kuratini and S. H. Bauer. J. Am. Chem. Soc. 87,

150 (1965).

[64] S. W. Benson and G. R. Haugen. J. Phys. Chem. 71,

4404 (1967).

[65] S. H. Bauer, P. Jeffers and N. Zevos. J. Phys. Chem.

71, 4412 (1967).

Discussion

C. F. Cullis (The City University, London)

Dr. Back has included in her paper a very helpful

summary of the present situation regarding our understanding of the mechanism of pyrolysis of

acetylene. As she is at pains to point out, however,

there are a number of questions still unanswered and

it is not at present possible, for example, to distinguish between various alternative routes in the

radical chain mechanism. Thus, vinylacetylene,

which appears to be the primary product of the

breakdown of acetylene below ca. 1000 °C, might be

formed either by reactions:

C2H +C2H2—^C4H3

C4H3 +C2H2—>C4H4 +C2H

or by reactions:

C2H3 +C2H2—>C4H5

C4H5—>C4H4+H

It is probably not going to be possible to determine

the extents to which such alternative reaction

25
sequences operate simply by means of more detailed

studies of the associated overall molecular reactions.

Further useful information may, however, be obtained from independent studies of the abstractive

and additive reactions of the radicals concerned.

Thus, for example, Tarr, Strausz, and Gunning have

shown^ that the reaction

C2H +C2H4^C4H4 +H

is faster at room temperature than the alternative

reaction,

C2H +C2H4—>C2H2 +C2H3.

At present therefore, Mr. J. V. Shepherd and Dr.

D. J. Hucknall are attempting, at the City University, to study the kinetics of some of the reactions

of ethynyl, vinyl and butenynyl radicals, with

particular emphasis on their interaction with small

alkyne molecules. It is hoped that such studies will

throw light on the mechanism not only of the pyrolysis of acetylene and its homologues at relatively low

temperatures but also of the formation of carbon at

somewhat higher temperatures.

M. H. Back:

I agree with the remarks by Professor CuUis both

with respect to the mechanism of formation of

vinylacetylene and to the view that measurements

of the rate constants for the elementary reactions

involved would be very useful. Below about 1000 °C

vinylacetylene may indeed arise from reactions of the

C2H or the C2H3 radical. The relative importance of

each sequence would depend on the relative concentrations of CoH and C2H3 radicals as well as on the

rate constants of the reactions involved.

In the mechanism outlined it has been assumed

that vinylacetylene is the main primary product of

the vinyl radicals while diacetylene is the main

primary product of the C2H radicals. This is an oversimplification which may be partly true but cannot

be assessed with confidence imtil the rate constants

for the elementary reactions and heats of formation

of the radicals are known more accurately than

at present.

R. S. Konar (Catholic University):

I would like to make some comments on the

dissociations of ethane and of neopentane.

Ethane Dissociation

The overwhelming experimental evidences [1-5]

now show that the high pressure rate constant

(k^) for the reaction 2CH3^C2H6 in the gas phase is greater than 10^^ exp( — 0//?T) cm^-m~'s~\ and

probably Shepp's number (/c«, = 2.20X10^^ cm^*

m^^s"^) represents the lower limit for k„. Thus in

• A. M. Tarr, O. P. Strausz, and H. E. Gunning. Trans. Faraday Soc. (1966) 62, 1221.

ethane dissociation,

AS298K: Ai/298 i5: (kcal m~i)

ki C2H6 <^ 2CH3

k-i 37.1(±1.0), 88.20(±3.0)

[where the entropy (S) is in cal-deg-^ mole-i.

Standard state ideal gas at 1 atmos; thermochemical

data are taken from Benson [6], and Purnell and

Quinn [7]]. = 1017 08 exp(-87600/i?r) s-i

if = 1013 -30 exp i-O/RT) cm^ • m-i • s'K

By measuring the initial rate of CH4 formation,

Quinn [8] found

A;i = 10i7-«exp(-91700//?r) s-i,

while Lin and Back [9] first found

ki = W'^ exp{-86000/RT) s~\

and later on

ki = W^-'^ expi-mOOO/RT) s-K

(see Lethard and Purnell [10] for a discussion).

It is suggested here that these discrepancies are probably due to the differences in the experimental

approach to the problem.

Quinn measured methane by taking samples

directly from the reaction vessel, and injecting it into

the G.C. (integrated with the reaction system via an

automatic gas sampling valve) . He measured the

initial rate of CH4 formation before the secondary

reactions started in the system (fig. 2 of reference 8).

The Arrhenius parameters as reported by Quinn are somewhat higher than the average values as shown

above, and this may well be partly due to the error

in the absolute calibration of CII4 in the G.C.

columns.

Lin and Back [9] measured CH4 indirectly, i.e. they froze all the products and the unreacted

reactant at the liquid N2 temperature, and then

toeplered the noncondensables. It is found that under

such conditions, all CII4 is not collected in the

toepler pump although in principle it should be

collected by the pump. The situation gets worse if there is hydrocarbon or inert gas matrix which

retains a considerable amount of CII4 by trapping.

H2 and CH4 were then analyzed by combustion of

hydrogen in a copper oxide furnace maintained at

280 °C (for details see page 506 of reference 9).

It may well be that Lin and Back did not collect all the CH4, and as a consequence might have underestimated the initial rate of CH4 formation. This is reflected in the low value of Ai.

The initiation reaction in ethane dissociation can

also be studied by measuring the initial rates of

n-butane and of ethyl formations coming from the

reactions (5a) and (5b) (see page MO). Lin and Back

did find that at high pressures,

J(n-C4Hio)M=A;(C2H6)

26
where /c = ^ exp( -89800/RT) s-\ and n-C4Hio

was measured by L. E. Reid and D. J. LeRoy [11]

have measured (k^/kc) of ethyl radicals, which is given by:

logioCfcdAc) = -0.213±0.014- (906zt30)/(2.3J?T)

Hence £'i = 90.7 kcal/mole. Thus,

A;i = 1016-5 exp(-88000/i?T) s'^ = ^iexp(-90700/i?T) s'K

Hence Ai=W-^'^s~^. It is interesting to note that

O'Neal and Benson [12] estimated

A-i = 1016-75 ( - 89500/iir) s-i,

assuming

fe_i= 1012-9 cm^-m-i-s-i

for which the experimental evidence is very poor.

It is clear that because of analytical errors, one can

hardly expect the expected Arrhenius parameters

from such complex pyrolytic reactions. Thus from

the experimental standpoints, the best Arrhenius

parameters for h are ki = 10'^-TM exp ( - 90700/i?T) s-i.

References

[1] A. Shepp. J. Chem. Phys. 24, 939, 1956.

[2] R. E. March and J. C. Polanyi. Proc. Roy. Soc. London.

A273, 360 (1963).

[3] H. E. Van Den Berg, A. B. Callear and R. J. Norstrom.

Chem. Phys. Letters. 4, 101 (1969).

[4] N. Basco, D. G. L. James and R. D. Stuart. Inter. J. Chem. Kinetics. 2, 215 (1970).

[5] T. C. Clark, T. P. J. Izod, M. A. DiValentin and J. E. Dove. J. Chem. Phys. 53, 2982 (1970).

[6a] S. W. Benson. Thermochemical Kinetics. John Wiley

and Sons, N.Y. 1968.

[6b] H. E. O'Neal and S. W. Benson. Inter. J. Chem. Kinetics.

1, 221 (1969).

[7] J. H. Purnell and C. P. Quinn. J. Chem. Soc. (London).

4049 (1964).

[8] C. P. Quinn. Proc. Roy. Soc. (London). A275, 190

(1963).

[9] M. C. Lin and M. H. Back. Can. J. Chem. 44, 505

(1966).

[10] D. A. Leathard and J. H. Purnell. Ann. Rev. Phys.

Chem. 21, (1970).

[11] L. E. Reid and D. J. LeRoy. Can. J. Chem. 46, 3275

(1968).

[12] H. E. O'Neal and S. W. Benson. Kinetic Data on Gas

Phase Unimolecular Reactions. NSRDS-NBS-21

(1970), pages 383-384.

Neopentane

By measuring the initial rate of C2H6 in neopentane

pyrolysis, Halstead et al. [1] found ki as

fei= 1018-05 exp(-85800//?r) g-i

(CH3)4C (CH3)3C +CH3

k-1 The estimated endothermicity of [1] at 298 K is (80.4±3.0) kcal mole-i, if AH^ of (CH3)3C radical

at 298 K is 6.7±1.0 kcal mole-i, and is (82.5±3.0)

kcal 'mole-i if the AHf of the (CH3)3C radical is (9.1 ±1.0) kcal mole~i as reported by Tsang.^

If we want to correlate the estimated endothermicity

(AHt) with the activation energy of the process

(El), it is necessary to know the value of E-i, since

Some kineticists believe that E-i is zero, some think

E-i is a positive quantity, while others suggest that

E-i is a negative quantity [2]. However all gas

kineticists accept that the experimentally measured

(kd/kc) ratio of alkyl radicals given below is probably

correct.

2C2H5, (CH3+C2H5), (CHj+n-CsHv),

(CH3 +/-C3H7), (CH3+t-C4H9),

0.13, 0.36, 0.058 0.163 0.70

Kc298 K

and that for two ter-C4H9 is 2.30, (Terry and Futrell

[3], Trotman-Dickenson and IMUne [43).

If E-i is zero, then it is very hard to explain all these experimental data. The Ec for the two ethyl

recombination has been measured by Shepp and

Kutsche [5] as 2.0 kcal mole~i, while (kd/kc) by

Reid and LeRoy [6],

\ogio{kd/kc) = -0.213- 906/ (2.3i?T),

i.e., Ed for the two ethyl is about 3.0 kcal/mole, (see

also Ivin and Steacie [7]). Comparing the (kd/kc)

of two ethyl radicals and of (CH3+ter-C4H9) , I am

inclined to think that Ec (i.e. E^i) of (CH3+ter-C4H9)

cannot be zero, and must be greater than the Ec

of two ethyls. Thus E-i may be as high as 3 to 4 kcal

mole~i due to steric barriers [8]. On this basis the

measured JEi = 85.8 kcal/mole is not inconsistent

with the estimated endothermicity of the process

[1], which is certainly not better than ±3.0

kcal/mole.

Tsang [2] has himself questioned his own shock

tube results. He believes that he might have studied

the reaction (1) in the slightly pressure dependent

region, and there may well be a molecular process

such as (CH3) 4C^iS0-C4H8+CH4

under his experimental conditions [2]. In that case, the measurement of total isobutene would not give

ki.Taylor et al. [11] studied neopentane pyrolysis in a flow reactor, and from the product analysis estimated the overall rate of neopentane disappearance.

They found experimentally that at or above 1 percent conversion,

-d(C,U^,)/dt=k(C,Ili2)

where A; = lOi^-^ exp( -80500/i?r) s-i. They then

assumed that k must be ku since the overall order is unity, and the overall activation energy is very close

to the estimated endothermicity of the process (1) at 298 K.

From the analytical data of Taylor et al, it is clear

that they have studied the isobutene inhibited

decomposition of neopentane, and not neopentane

27
itself. The overall order (unity) and overall low

activation energy (80.5 kcal/mole) may well be the

consequence of self-inhibition (see ref. 1).

Taylor et al. did not produce any data below

1 percent conversion and the extrapolation of —log (percent conversion) versus time to zero time

would not give the correct rate constant, and hence

the initial rate, in the inhibited decompositions if the

conversion is high (1 percent or above) . The presence

of olefins masks the normal reactions [1, 14]. It has

been found experimentally that when neopentane is pyrolysed at 500 °C with 1 percent isobutene, CH4

is formed at a constant rate initially and the pressuretime curve when extrapolated goes through the

origin; but it does not give the kinetics of neopentane

pyrolysis. In the absence of olefins (added), CH4 rate

decreases with time [1]. This is also the case in the

flow reactor of Taylor et al. (fig. 4, ref. 11). They

have claimed that they have studied the homogeneous reaction, but did not measure H2, a major

product, nor did show its distribution inside the

reactor. It appears H atoms will reach the reactor

wall under their experimental conditions. It was

found that neopentane diffused to the wall in about

0.3 sees, and the time of residence was up to 0.4 sees. Thus there is considerable doubt about the validity

of the measurement of ki by the method of Taylor

et al., and so their results have to be accepted with

considerable reservation.

References

[1] M. P. Halstead, R. S. Konar, D. A. Leathard, R. M.

Marshall and J. H. Purnell. Proc. Roy. Soc. London.

A310, 525 (1969).

[2] W. Tsang. J. Chem. Phys. 44, 4283, (1966); Inter. J. Chem. Kinetics. 1, 245 (1969).

[3] J. O. Terry and J. H. Futrell. Can. J. Chem. 45, 2327

(1967) ; 46, 664 (1968).

[4] A. F. Trotraan-Dickenson and G. S. Milne. Tables of

Bimoleeular Gas Reactions, NSRDS-NBS-9.

[5] A. Shepp and K. O. Kutschke. J. Chem. Phys. 26, 1020

(1957).

[6] L. E. Reid and D. J. LeRoy. Can. J. Chem. 46, 3275

(1968) .

[7] K. J. Ivin and E. W. R. Steacie. Proc. Roy. Soc. London.

A208, 25 (1951).

[8] R. S. Lonar. Inter. J. Chem. Kinetics. 2, 1970 (in press).

[9] F. Baronnet, M. Dziersynski, R. Martin and M. Niclause.

C. R. Acad. Sci. (Paris). 267, 937 (1968).

[10] D. A. Leathard and J. H. Purnell. Ann. Rev. Phys.

Chem. 21, (1970).

[11] J. E. Taylor, D. A. Hutchings and K. J. Freeh. J.A.C.S.

91, 2215 (1969).

[12] S. W. Benson. The Foundations of Chemical Kinetics.

McGraw Hill, N.Y. page 354 (1960).

[13] H. E. O'Neal and S. W. Benson. Inter. J. Chem. Kinetics.

1, 221 (1969).

[14] R. L. Stock and H. E. Gunning. Can. J. Chem. 38, 2295

(1960).

M. H. Back:

Dr. Konar has provided arguments in support of

the rate constant reported by Quinn for the elementary dissociation of ethane. The frequency factor

measured by Quinn was in good agreement with the

rate constant for the reverse combination of methyl

radicals but the activation energy for the dissociation

which he measured gave a heat of formation of the

methyl radical of 35.6 kcal/mole at 298 K, which in

turn leads to a dissociation energy of methane of

105.6 kcal/mole. This value is probably too high and

it appears not possible to find agreement among the

four quantities ki, k^i, AS and AH for the system

1 C2H6 ^ 2CH3

—1

Regarding the measurements of the rate constant

for the dissociation of neopentane. Dr. Konar's

caution in accepting measurements at conversions

greater than 1 percent may be well-founded. Indeed

very low conversions were needed in the work done in

the static system. The data reported by Taylor

et al. from the experiments in the flow reactor shows,

however, that they were measuring initial rates of

formation of methane, although the accuracy did not

appear very good. Perhaps it is simply a combination

of errors which leads to the low value of 80.5

kcal/mole for the dissociation of neopentane. Their

work was discussed in relation to the previous

measurements on neopentane by Baronnet and

Niclause (ref. 4 of the discussion)

A. S. Gordon (Naval Weapons Center):

Referring to the mechanism of ethene pyrolysis, I would be interested in any experimental evidence

that a vinyl radical (or a homologue) will disproportionate with about zero energy of activation.

Wly interest is aroused because an internal hydrogen

transfer reaction,

D

/ \

CD2 CH2

I il ^CD2=CHCH=CHCH2D

CH CH

has AH= 0, but an Eact^ 35 kcal/mole. This reaction

bears some formal relation to the hydrogen transfer

reaction between two ethylenes. In addition,

endothermic abstraction reactions, e.g. allyl radical

abstracting from a saturated hydrocarbon, have a barrier in addition to the endothermicity of the

reaction. The two types of reactions suggest that the

H transfer reaction between two ethylenes might

have a larger barrier than the endothermicity.

M. H. Back:

There is to my knowledge no direct evidence that

disproportionation between a vinyl and an ethyl

radical occurs without activation energy. This

assumption was made on the basis of what is known

about the rates of radical-radical combination and

disproportionation reactions involving small hydrocarbon radicals. Almost all the evidence leads to the

conclusion that disproportionation reactions between

such radicals require no activation energy. In view

of the considerable exothermicity of these reactions

28
this result is not surprising. The reactions discussed

by Dr. Gordon, on the other hand, are hydrogen

transfer reactions which often, as he points out,

have heats of reaction close to zero and which may

involve a considerable activation energy. Now two

important differences between these types of

reactions could lead to considerable differences in

the activation energies involved. (1) Disproportionation reactions are highly exothermic while

abstraction reactions may be close to thermal

neutral. For example,

2C2H6^C2H4+C2H6

Aif=59 kcal/mole

C2H3 +C2HB—>2 C2H4

AH=65 kcal/mole

CH3+ C4H10—>CH4 +C4H9

AH=5.3 kcal/mole

CD3CH^CHCH=CH2->CD2=CHCH=CHCH2D

AH=0

(2) The dissociation energy of the C-H bond

broken is very different in the two cases. In the

disproportionation reaction the C-H bond dissociation energy is probably reduced to about 40

kcal/mole, while the C-H bond broken in the

abstraction process very likely has close to its normal strength of about 100 kcal/mole.

Dr. Gordon's second point is that an endothermic

reaction involving radicals may require an activation

energy considerably greater than the endothermicity.

This of course depends entirely on the nature of the

reverse reaction and in the case of our proposed

initiation reaction in the ethylene system

2C2H4—>C2H5 +C2H3

we equate the activation energy for this reaction to

the endothermicity by assuming that the disproportionation between a vinyl and an ethyl radical

has the same energy requirements as the disproportionation between two ethyl radicals.

W. Tsang (National Bureau of Standards):

The mechanism, rate constants, and parameters of

the initial step in the gas phase pyrolysis of most

organic molecules can be conveniently and accurately

determined using the comparative rate single pulse

shock tube technique^ This is because the shock

tube has all the characteristics of an ideal pyrolytic

reactor. Wall reactions are eliminated and the short

reaction time (1 us) excludes the possibility of chain

decomposition. Using this technique rate parameters

for the initial step in the decomposition of over 50

hydrocarbons (up to C12) have been determined.

With respect to the alkanes they show striking

regularities. There are no molecular modes of

> W. Tsang, J. Chem. Phya. 43, 352 (1965); 44, 4283 (1966); 46, 2187

(1967). Int. J. Chem. Kinetics 1, 245 (1969).

reaction. All variations in rates are due to activation

energy changes. They are generally in good agreement with presently accepted^ values of bond

energies. The preexponential factor per C-C bond

being ruptured appears to be a constant with the

appropriate number being 2 X 10^^ s~^ Calculated

preexponential factors from measured radical recombination rates in the literature are however a factor

of 10 to 30 higher. This is the same phenomena as observed by Professor Back.

The shock tube technique is applicable to the

study of the pyrolysis of any compound with a vapor

pressure of about 10 ^um or above. For really large

hydrocarbons a heated shock tube is necessary. This

is not a problem. On the other hand from the general

trends already observed, startlingly new behavior

should not be expected.

From the data on ethane as summarized by

Professor Back it appears that for such simple

systems the classical static technique may give

acceptable rate constants, they do not yield dependable rate parameters. Thus any set of rate

parameters derived from a given static experiment

must be regarded with extreme caution.

M. H. Back:

Any one set of rate parameters from a given

experiment, whether by static, flow or shock tube

technique, must always be regarded with caution.

The rate parameters for the dissociation of ethane

which I have discussed were obtained from several

sets of experiments and indeed the parameters

obtained from the static technique were in good

agreement with those obtained by Dr. Tsang from

his experiments in shock tubes. The data from both

techniques show a discrepancy with the radical

recombination rates and it is this discrepancy which

I wished to call attention to in my discussion. The

reasons for it are not yet fully understood.

A. S. Gordon and R. H. Knipe:

We believe that the saturated hydrocarbons

(alkanes) have been shown to decompose exclusively

via a chain mechanism even previous to the work of

Tsang. However, there are other classes of hydrocarbons which do decompose exclusively via a molecular process (or biradical which reacts so fast

that it is indistinguishable) as discussed in our paper.

As for Dr. Tsang's technique: while it is mostly

free of wall, Kistiakowsky and his students have

shown that small quantities of gases are desorbed

from the walls of shock tubes and influence the

reaction. Also, the shaping of the pulse to give a step

rise to a given temperature and a similar decrease at

the end of pulse leaves some error in the results.

The most important criticism of his technique is that

he must rely on toluene (or some similar agent) to

completely quench the free radicals. When acting in

this manner the benzyl radical is assumed not to

2 J. A. Kerr, Chem. Rev. 66, 465 (1966).

29
abstract hydrogen from the parent compound. The

assumption is not correct; it will abstract with an

activation energy greater than the alkyl radical

because of its resonance energy. The higher the

temperature of the reaction the less effective the

technique. Actually, the technique is not restricted

to the shock tube but has been applied in static and

flow systems, where it was pioneered by M. Szwarc.

The single pulse shock tube has the advantage

that it can reach higher temperatures than the

classical techniques. In using the technique at higher

temperatures it must be carefully established

whether the hydrocarbon is in its "fall-off" region.

For the lower temperature region, the classical

techniques of a static system where it can be experimentally established that the reaction is homogeneous, allow the various parameters such as pressure,

inert gases, and temperature to be more conveniently

explored than in a shock tube. As is so often the

case, no one experimental or analytical technique can be used to solve the research problem area.

W. Tsang:

Drs. Gordon and Knipe have made the common

assumption that the only special attributes of shock

tube studies are high temperature and relative

homogeneity of reactions. They have failed to draw

the proper consequences that arise from the short

reaction time {'—^1 ms) and low reactant concentrations (^.01 percent in argon or 3 X 10~® moles/liter

tmder experimental conditions). More than anything

else these isolate the initial processes for study.

Consider, for example, a methyl radical in this

environment; the rate expression for its abstraction

reaction with a typical hydrocarbon (say n -butane)

as recommended by A. F. Trotman-Dickenson and

G. S. Milne [^Table of Biomolecular Gas Reactions,

NSRDS-NBS 9, U.S. Government Printing Office] is

/c(CH3 +riC4Hio^CH4 +C4H9) = lO^ o exp ( - 7,600) /i?ri/mole s. At 1100 K (the usual shock tube temperature)

this is 10^^ 1/mole s. With a hydrocarbon concentration of 3 X 10"'' moles/liter this is equivalent to a methyl radical lifetime (with respect to abstraction)

of 0.1 s. Remembering that the total reaction time is only 0.001 s, this means that methyl radicals are

stable under shock tube conditions. Since there is no

question that the comparable benzyl radical reaction

must be even slower, it is clear that the suggestion

by Drs. Gordon and Knipe is not applicable to

these shock tube studies.

Note that it may be (and has been) used in

explaining the serious errors in the original toluene

carrier work of Szwarc. These are carried out with

much higher concentrations and longer residence

times. More complete discussions of this and the

other minor points may be found in our publications

[_W. Tsang, Internat. J. of Chemical Kinetics 2,

311 (1970) and references therein].

It should be noted that the basic correctness of the

rate constants determined by the single pulse shock

tube studies have now been verified by flow [J. E.

Taylor, D. A. Hutchings, and K. J. French, J. Amer. Chem. Soc. 91, 2215 (1969)]; static [M. P.

Halstead, R. Martin, D. A. Leathard, R. M. Marshall,

and J. H. Purnell, Proc. Roy. Soc. Ser. A. 310, 525

(1969)]; and very low pressure pyrolysis studies

[N. Spokes, personal communication].

Leo A. Wall (National Bureau of Standards):

Drs. Gordon and Back have given very excellent

discussions of thermal decomposition reactions;

however, they confined their remarks to very small

molecules. I would like to turn their attention to

large molecules, for example alkanes with ten or more carbon atoms, and ask them to comment on

the following:

F. O. Rice discussed the thermal decomposition of

large alkane molecules many years ago (The Aliphatic

Free Radicals, Rice and Rice, The Johns Hopkins

Press, 1935) and later the isomerization of large free

radicals (A. Kossiakoff and F. O. Rice, J. Am. Chem.

Soc. 65: 590 (1943)) was invoked to improve the

fit of the original free radical theory with experiment.

A premise of the free radical theory was that large

radicals were unstable and decomposed to small

ones. It has occurred to me that actually large

radicals are quite stable and in a sense inert. In the

polymer field the conformation of molecules is very

important. In molecules containing ten to forty or more carbon atoms, one would suspect that con- formation would play an important role in gas

pyrolysis. In a twenty or so carbon atom molecule an

active radical end could be shielded or in effect

buried and not easily reacted with other molecules,

hence, except for isomerization, decomposition to

smaller species would be the only course they could

follow.

M. H. Back:

The reactivity of radicals in an energetic sense does not appear to change very much after about C5

or Ce, at least for straight chain radicals. Long chains

are remote from the reactive part of the radical and

will have very little eff^ect on the energetics of the

radical reaction. Dr. Wall has suggested that there

may be other important influences besides energy

requirements on the rates of reactions of long chain

polymer radicals, for example, conformational effects.

The reactive end of a radical may become shielded

or buried and bimolecular reactions might be

hindered relative to unimolecular reactions. The

effect would be manifested in the entropy of activation of the transfer process, while the energy of

activation would show little change. The entropy

term in the thermodynamics of the polymerization

process is responsible for the observation of a ceiling

temperature for the reaction. A large entropy term

arising from conformational effects on the reactions

of the polymer radical would thus have an effect on

the ceiling temperature. A comparison of ceiling

temperatures for various types of polymerizations

30
might give some indication of the importance of

conformational effects.

A. S. Gordon and R. H. Knipe:

We agree with Dr. Wall that a large hydrocarbon

radical would be less likely to stabilize itself by H

abstraction from another molecule or radical than a small radical. Since the unimolecular isomerization

reactions are probably the fastest reaction paths in

the gas phase, it is likely that the various isomers of a

large radical would tend to be present in their

equilibrium proportions. In the case of the large

radical many of the isomers would have the free

radical end buried inside the coils of the structure,

so that these would not be able to abstract H from

other molecules. In a practical situation large

radicals could exist in the gas phase only in very low

concentrations except at high temperatures because

of vapor pressure limitations. At high temperatures

the unimolecular cracking reactions would be very

31

NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Factors Involved in the Degradation of Polymers in Melts

C. R. Patrick

The Department of Chemistry, University of Birmingham, Birmingham, England

The full interpretation of degradation of molten polymer would be facilitated by a better

understanding of the elementary chemical processes involved. Suitable analogues of these reactions

may often be studied in detail only in the gas phase. Relationships between the kinetics of ele- mentary reactions in gas and liquid phases are desirable. On the basis of thermodynamic argu- ments it is concluded that Arrhenius parameters for bimolecular reactions in the two phases

should differ, to a small but significant extent, greater than usually accepted, whilst those for unimolecular processes, commonly should not. Diffusion may determine the rates of some re- actions in liquids particularly of reactions between radicals. Evidence on the mobilities of large

and small molecules, models for radicals, in molten polymers, is summarized. Expressions for

diffusion controlled reactions are compared, and applied to reactions in molten polymers. Complicating effects in the size dependence of the rates of diffusion controlled reactions between macroradicals are surveyed. Other complicating effects, notably of geminate reactions between radicals, are discussed. A case is made out for comparing more extensively the kinetics of thermal degradation of gaseous paraffins, and of molten paraffins and polyethylenes.

Key words: Degradation of polymers in melts; diffusion-controlled reactions between macroradicals; diffusivities of molecules and radicals in molten polymers; geminate recombination of

radicals; pyrolysis of hydrocarbons; rates of reaction in liquids and gases; reaction equilibria in liquids and gases.

1. Introduction

This paper examines the bases upon which

kinetic parameters, which for simple substances

may be characterized in the gas phase, may be

related to kinetic parameters for the same, or similar,

processes in the liquid phase. The comparison is made on a thermodynamic basis. It is also necessary

to acknowledge that the rates of some reactions in

liquids, such as those involving free radicals that are of importance in the chain degradation of addition

polymers, are regulated by diffusion, rather than by

chemical factors. Such diffusion-controlled reactions

will be discussed, and outstanding problems will be

indicated.

2. Thermodynamic Considerations of

Reactions in the Gas and Liquid

Phases

We will consider the thermodynamics for the

equilibrium between a set of reactants and products,

and their relation to the thermodynamics for the

same equilibrium in the liquid phase. We assume

that the components form thermodynamically ideal

mixtures in both phases. Equilibrium constants for

reactions in the gas phase are generally formulated in

terms of the partial pressures of the reactants and

products, using reference states of one atmosphere

pressure for the respective substances. Equilibrium

constants in the liquid phase are most appropriately

formulated in terms of the mole fractions of the

components, that is, with reference to the (real or hypothetical) pure liquids (strictly at a defined

pressure, which, for the present purposes, will be

assumed to be a low pressure—say one atmosphere)

The heats of reaction in the gas and liquid phases

(Ai/fi g and AHii,i, respectively) are related by the

heats of vaporisation (AHi,) of the component

liquids in their reference states to the reference

states of the gases (P= 1 atm) . Thus:

products reactants

Similarly the standard entropies of reaction are related by

products reactants

where the entropies of vaporisation, ASo, relate to

changes of state specified for the heats of vaporisation.

The first of these relationships has often been

presented 2, 3, but the second has not been

exploited fully. The detailed development of these

equations cannot be pursued in this paper. At

temperatures below, and remote from, the critical

temperatures of the liquids, the appropriate heats

and entropies of vaporisation are, in principle, easily

determined. In the neighbourhood of the critical

temperature the quantities required are not easily

obtained, and the major contributions to them are due to the non-ideality of the vapors, rather than to

^ Figures in brackets indicate the literature references at the end of this paper.

33
the vaporisation of the liquids. Above the critical

temperatures of the liquids formally identical

relationships still obtain, but the terms for the

vaporisation of the substances need strictly to be

redefined. The solution of a gas in a solvent with

which it has no strong interaction is not accompanied

by a change in heat that is comparable in magnitude

with the heat of vaporisation of the same substance

as a liquid [5j]. There is, however, a significant

change in entropy when a gas is dissolved (at constant pressure) in a liquid, and this, when extrapolated to the limit at which the liquid is composed

of the hypothetic liquified gas, is very similar in

magnitude to that for the vaporisation of a normal

liquid to the vapour at the same pressure and temperature (that is about 22 cal-deg"' mole~^ for

vaporisation to P=l atm at 25°). This result was

embodied in earlier work [5] but was only recently

made explicit [6J. It is easily understood. In the gas

the motions of the molecules are limited only by the

total volume of the system. In the liquid, motions

are confined to a very small volume, commonly

referred to as the free volume. The change in entropy

on this account is similar for all molecules, because,

in simple terms it is only the change in volume

available for the molecular motion that matters.

Some generalisations follow from these relationships. In equilibria that involve no change in

molecularity, it is likely that the summed heats of

vaporisation will equal very nearly those of the

products, so that A/i^R . g — Aifr , i will be small. In

associative equilibria, of the type A-\-B^AB, [7]

the disparity between the sums of heats of vaporisation of reactants and products is likely to be greater, so that, unless some of the species are, at the temperature of interest, real gases, AHr ^ — AHr i will

commonly be significant and negative.

Troutons rule [7, 8] teaches that the entropies of

vaporisation of many substances are approximately

equal at all temperatures, and are about 21 cal deg~^

mole~^ Thus for equilibria involving no change in

molecularity, ASR,g— ASr.i should be small, and in

associative equilibria, as defined above, ASr,^— ASr ,\ should generally be about —21 cal deg~' mole~^

The predictions relating to equilibria involving no

change in molecularity are demonstrated by measure- ments on equilibria between halogens and diatomic

interhalogen compounds [2, 9]. Predictions for

associative equilibria are less well illustrated. The

entropies of dimerisation of carboxylic acids in the

gas phase (— (34 to 36) cal deg~^ mole~i) are greater

than those in solution ( — (10 to 14) cal deg~^ mole"')

[10], and the differences are about —(20-23)

cal deg~' mole"') as expected. The heats of dimerisation in the gas phase (— (14 to 16) kcal mole~') are greater than those in the liquid phases ( — (8 to 10)

kcal mole~') [10]. These results accord with our

simple theory, and do not imply, as has been suggested [lOj that the reactants interact in a special

way with the solvents used. Direct experimental

evidence to test our predictions for associative

equilibria are sparse. With one notable exception, the

dimerisation of nitrogen dioxide [11] the predictions

are, however, supported reasonably well by a body of

evidence, too extensive to review here, but which

deserves further assessment.

3. Application of Thermodynamic

Results to Reaction Kinetics

Reaction rate coefficients are conventionally

defined in terms of concentrations and changes in

concentration. For any chemical process (for

example A-\-B^AB) we may define rate coefficients

kf and kr for the forward and reverse reactions

respectively. These are related to the equilibrium

constant for the reaction, expressed in terms of

concentrations: Kc = kf/kr [12]. We continue to

consider thermodynamically ideal systems. More

generally the relationship between rate coefficients

and equilibrium constant still obtains at equilibrium,

but the rate coefficients, since they have been

defined in terms of concentrations, rather than

activities, may vary with concentration. Considering

further reactions in liquids, in which the thermodynamic equilibrium constant Kx is defined in terms

of mole fractions [12] we have

where An is the difference between the number of

molecules of products and the number of molecules

of reactants in the stoicheiometric equation for the

reaction, and Vs is the volume of liquid that contains one mole of substance. We may write:

Kx= exp[AS«,i].exp[-AH«,i/i{T]

and

K, = Bc exp[-AHfl,ue)/i?T]

To a commonly adequate approximation [13]

AHr,, = AHr,im and fi,= exp[ASi?,i/i?]

The preexponential factors, Af and Ar, for the

forward and reverse rate coefficients, are related by

Bc = A//Ar, and the corresponding energies of

activation are related by AHr,i(c) = (Er—E/).

Taken together with the conclusions of the

previous section these equations offer the possibility

of making predictions about, and testing the consistency of, rate coefficients in liquids, particularly

since it is possible, with reasonable confidence, to

predict the entropies of reactions in the gas phase

tl4].

Standard entropy changes for associative reactions

(A-]-B^AB) in the gas phase lie in the approximate

range — (37 to 48) cal deg~' mole~', so that those in

the liquid phase might be expected to lie in the

range —(15 to 28) cal deg"' mole"'. In a solvent

composed of small molecules, Vs is about 0.1 1 mole"*,

so that Be is about 10"^=*=' 1 mole~'. In a polymeric

matrix, having an average molecular weight of

5X10^ Vs may be 50 1 mole"', so that B^ is about

10 -*±'1 mole-'.

34
4. Comments on Rate Coefficients in

Gas and Liquid Phases

Equilibrium constants for unimolecular isomerisations should be nearly equal in the two phases, so that the rate coefficients must be nearly equal, or change by approximately equal factors in passing

from one phase to the other. Since for associative

reactions the equilibrium constants differ for the two

phases, either or both the rate coefficients involved

must change in passing from one phase to the other.

Arguments based upon the transition state theory of

reaction rates, and a weak argument based upon the

likely behaviour of rate coefficients in unimolecular

isomerisations, suggest that the Arrhenius parameters

for rate coefficients in the gas and liquid phases may

differ significantly for bimolecular processes, but less

so, or only very slightly, for unimolecular reactions

(saving those influenced by diffusion). Other

authors [2, 15, 16] have reached similar conclusions

using detailed treatments of rate coefficients, but

have made approximations which, whilst they might

appear plausible are not strictly thermodynamically

consistent. Kinetic results, which are very sparse,

particularly for bimolecular reactions, do not support

the predictions. Data most frequently cited for

bimolecular reactions in this context QITJ, namely

those for the dimerisation of cyclopentadiene, are not thermodynamically consistent.

For illustration we apply our principles to

the decomposition of t-butyl peroxide. The

rate coefficient for the decomposition in solution is 3XlOi^exp[-35,000/i?r>-i [18]. (In this and

subsequent like expressions R is in cal deg~^ mole~')

The rate coefficient for the association of t-butoxyradicals in solution is 1.8 XlO^ exp[±2,000/i?T]

1 mole's"^ [19], say about 10^=*^^ 1 mole~^s~'. Thus

Be is about 10~^±^ 1 mole"'. The standard entropy

change for reaction in the gas phase is about —40

cal deg~' mole"' [20]. The entropy change in the

liquid, therefore, should be about —20 cal deg~'

mole"'. The discrepancy is disappointing, but it should be remembered that rate coefficients, and

more especially their Arrhenius parameter, are often

subject to greater errors than are admitted. The

preexponential factor for the gas phase decomposition

of i-butyl peroxide is close to 10'^, s~' greater than

that in solution, whilst the energy of activation

(about 38 kcal/mole) is also greater than that

quoted for solution [21]. On the other hand the rate

coefficients for the recombination of the radicals

appear to be very similar in the two phases (ca. 10^

1 mole"' s~') [19, 21]. The example has been chosen

to illustrate difficulties—either in the uncertainties

in the experimental data, or in the evaluation of

the theory.

5. The Effects of Non-Ideality of

Thermodynamic Behaviour

Thermodynamically ideal behaviour is rare in

liquid mixtures. The standard entropy of vaporisation will differ in nonideal solutions, but must

commonly remain of the same order, except where

the non-ideality results from very strong interactions

which may be regarded as chemical in nature. Our

arguments, saving such exceptions, must provide

useful, if sometimes poor, approximations for nonideal solutions. It is interesting to speculate on

whether free radicals will form markedly non-ideal

solutions.

6. The Possible Effects of Diffusion on

the Rates of Thermal Degradation of

Polymers

Diffusion of reactants and products within

polymer melts may effect the rate of degradation in

two principal ways. Firstly, some products of

reaction may diffuse so slowly out of the system that

they interfere with the main course of the reaction.

This was noted in the degradation of polymethylmethacrylate, where the major product, monomer,

remained in the polymer in sufficient concentrations

to cause a back-reaction [22]. The detailed analysis

of such a situation has been given [23]. Secondly the

rates of reaction together of free radicals may be

determined by rates of diffusion as in polymerising

systems [26, 15, 64] and decomposition of polymers

in solution [25].

7. Diffusivities of Molten Polymers

In view of the difficulties involved it is not sur- prising that few measurements have been reported

of diffusivities in polymers at high temperatures

particularly in the range (200-450°) where thermal

degradation is important. It is necessary, therefore,

to estimate values using theoretical or semiempirical

relationships. We shall consider the behaviour of two

extreme classes of molecules, namely of macromolecules, comparable in size with those which

compose the greater part of the mass of polymer, and

small molecules, comparable in size with monomeric

imits, or other small products of decomposition.

The viscosities of polymer melts are more easily

measured than are the diffusivities of the constitutent

molecules. It is convenient, therefore, to use the

relationship between the viscosity, rj, and the self- diffusivity, D, of the system of macromolecules:

[24, 27].

Drj = A{T) = (pNav/36) {RyM)hT

where is the mean square end-to-end distance of

the polymer chain, M is the molecular weight, p is the density, and Nav is Avogadro's number. The

validity of the relationship has been demonstrated,

but chiefly for bulk or swollen polymers at temperatures of up to about 50° above their glass transition

temperatures [28]. The greatest dependence of

A (T) upon T is through the term kT.

The term (R^/M) is not very sensitive to temperature, or to changes in molecular weight. Although

it depends upon the structure of the repeating unit

it is not very markedly so. Experimental values of A

35
are about (4-8) XlO-» cm^ at about 50° [28]. The

equation was derived for, and is best understood if appHed to, systems composed of polymers of uniform

molecular weight. Most real systems contain a wide

range of molecular weight, so that the value of D

to be associated with 77 in the above equation has no

simple significance.

Viscosities of melts have not always been measured

for the molecular weights or temperatures in which

one is interested. In order to estimate the required

properties the following relationships may be used:

logTj = 3.5 logM+constant [30, 29]

logTj = - 8.86 {T-Ts)/{ 101.6+ T- T,)

[24, 27, 31, 32]

where Ts= (Tg+50), and Tg is the glass transition

temperature of the polymer. The second relationship

holds only for molecular weights greater than a certain value, characteristic of the polymer, but

generally about 10^-10^ [33]. For polymers of lower

molecular weight, which do not entangle with one

another to such great effect, i] is approximately

proportional to M [34, 35, 27]. (There is some

evidence that at very high molecular weights ( > 10")

77 varies more rapidly with M than M^-^ [36]. M

may be an average (weight-average) molecular

weight [27, 32].

The temperature-dependence of viscosities of

melts, following the expression quoted, does not

approach the Arrhenius form except at high temperatures [27]. The limiting approximately Arrhenius

form indicated by the expression requires an energy

of activation that is independent of molecular

constitution and is about 4 kcal/mole [32]. This is not true of liquids in general [32]. The equation

may not apply at high temperatures. An alternative

expression has been suggested but not justified for

extrapolation to temperatures far in excess of

(Tg+100) [30].

A significant implication of these equations is that

the temperature-dependences of -q and D are independent of molecular weight, at least to a good

approximation. The constants of proportionality in

these expressions vary considerably from one polymer to another, so that interpolations of data

between different polymers is inadvisable [30].

Amongst polymers whose thermal degradation has

been of interest, polytetrafluoroethylene (PTFE) is notable in possessing a very high melt-viscosity

[37], which is in the region of 10^^ poise, in the

temperature range 330-390 °C. The polymer is generally of high molecular weight ( > 10^) . The

viscous relaxation of PTFE melts at 380°, having M

in the range 10^-10^, shows a dependence on M°-^

rather than M^-^ [38]. If, as is questionable in view

of the latter observations, we may assume Dr] = A{T)

and use a typical value for A ( T) (say 10 X 10~^ cm^)

we estimate that in the PTFE melts described, the

effective self-diffusivities lie between 10~^^ and

10^^^ cm^ s~^ [39]. The apparent energy of activation

for viscous flow in this temperature region is 36

kcal/mole [37], although it has been suggested the

energy of activation varies over the temperature

range [40]. Over the range of temperature of interest

for the thermal degradation of the polymer (380-

450 °C) [23], the diffusivity is likely to vary by a factor of less than 15.

The diffusivity of radioactively labelled polystyrene (M=8X10*) in polystyrene of higher

molecular weight (10^X10") at 237 °C is iXlO-i" cm- s~^ [41] and independent of molecular weight of

the substrate [42]. Whilst these results do not prove

the point, they are consistent with the view that,

provided M is large, diffusivities of macromolecules

are proportional to M~^-^ independently of the

composition of the matrix. M is here the molecular

weight of the diffusing polymer, and not an average

for the matrix. The measurements were made at a temperature of about 140° above the glass transition

temperature of the polymer [43, 44]. The thermal

degradation of the polymer is usually studied in the

temperature range 300-350 °C [45]. In this temperature range the self-diffusivity of the macromolecules would be greater than the value measured

by a factor of between about 6 and 16, assuming that

the viscosity-temperature relationship may be extrapolated so far.

The viscosity of polymethylmethacrylate

(M=6X10*) at 150 °C. (About 45-50° above the

glass transition temperature of the polymer [44]

is 10' poise [35]. The self-diffusivity of the macromolecules is then in the neighbourhood of 10~" cm- s~^ At about 240 °C we calculate that the

diffusivity should be 0.7 X 10~^° cm^ s~\ very similar

to the value obtained for polystyrene. This comparison assumes that the polymethylmethacrylate

sample had a narrow range of molecular weight. The

correction for the differing molecular weights

(80,000/63,000)3 5 is about 2.3. The diffusivities in

the range used for studies of thermal degradation

(250-350°) [46] will be in the range 6XlO-i»

cm^ s~^ to 10~^ cm- s^^

The diffusivities of polymers of lower molecular

weight are as expected, smaller. The diffusivity of

polymethylene in a matrix having molecular weight

5,800 is 1 X 10-' cm^ s"! at 150 °C [47].

8. Diffusivities of Small Molecules in

Molten Polymers

As with macromolecules in bulk polymer, glass

transition temperatures mark a change in the

character of diffusion of small molecules in polymer

[48]. Values of diffusivities should not, therefore,

be extrapolated from measurements made below the

transition temperature. The diffusivities of small

molecules in crystalline polymers appear to be

smaller than those in amorphous polymers, and to

depend on the degree of crystallinity in a manner

that is not well-established [49]. It seems unwise,

therefore, to attempt to extrapolate diffusivities for

melts from measurements made on crystalline

polymers.

The diffusivities of nitrogen and methane in

polyethylene (of unspecified molecular weight) are

36
similar, and lie between 2.5 XlQ-^ to 5.5X10-=

cm^ s"^ in the temperature range 125 to 255 °C.

[50, 51], showing little variation with temperature.

In contrast the diffusivities of methane in polystyrene in the same temperature range lie between

5 X 10~' and 5 X 10~^ cm^ s^^ [50], and vary markedly

with temperature. The diffusivities of n-pentane in

polystyrene (M= 412,000) in the temperature range

140-170 °C. lie between 2.6X10-8 and 1.7X10"'

cm^ s~^ Even over this small range of temperature

the diffusivities cannot be represented by an equation of the Arrhenius form. Similar behaviour has

been reported for the diffusion of other small organic

molecules in polystyrene [52] at these temperatures.

The differences in behaviour shown by the diffusivities of methane in polystyrene and polyethylene

are easily rationalised. The observations for polystyrene were made little above the glass transition

temperature of the polymer (100 °C) [44], whilst

those for polyethylene were made well above the

glass transition teniperature of the polymer ( —21 to

-125 °C) [53, 30J. It would seem plausible, and

has been suggested that [54] an equation of the form

applied to the viscosity of macromolecules in

polymer melts, might account for the temperaturedependence of diffusivities of small organic molecules

in amorphous polymers at higher temperatures. Not

all experimental evidence would support the view

that there is a simple universal relationship involving

Tg for the diffusivities of small molecules in polymers

[55]. For members of an homologous series, D

should vary approximately inversely with M

[32, 56, 57].

Measurements of the diffusivities of small molecules

into polymers are most commonly made at, or a

little above, room temperature, so that extrapolations need to be made over long ranges of temperature in order to obtain estimates in the region

where thermal decomposition is of importance. In

summary we believe that at high temperatures

(>250 °C) diffusivities of molecules of about the

size of monomers have diffusivities of 10"^ to

5 X 10"^ cm^ s"^, and possess small temperature

coefficients (probably of the Arrhenius form, with

energies of activation of about 4 kcal/mole).

9. Effects of Cross-linking and

Swelling of Polymers

Cross-linking of polymers, even to a low degree

will effect markedly their diffusion. It also effects

significantly the diffusivities of small molecules in

them [58]. Conversely the swelling of polymers, or more strictly in the present context, the solution in

them of substances of low molecular weight, greatly

enhances the diffusivities of both small and large

molecules.

10. Rate Coefficients for Diffusioncontrolled Reactions

Theoretical expressions for rate coefficients for

bimolecular diffusion-controlled reactions are based

upon two types of physical model. The first, devised

by Smoluchowski [59] and developed by others,

[60, 66, 67], evaluates the flux of other reactants

towards a chosen reactant molecule. The second,

introduced by Rabinowitch, [61] considers the

imitary diffusive displacements of the reactant

molecules that lead to encounter and reaction.

Rabinowitch assumed a quasi-lattice model for the

liquid, limiting his treatment of systems of small

molecules. Some limitations of this assumption have

been avoided in a later "random walk" treatment

[62]. Models of the second class suffer the disadvantage of requiring the average length of the

unitary diffusive displacement, which is commonly

identified with intermolecular separations of small

molecules in the liquid state [63].

The Smoluchowski equation is:

kA.B{\ mole-1 s-i) =4X10-Wav-D.4,bR

where R (cm) is the collision diameter for the

colliding particles, A'^av is Avogadros number, and

Da,b (cm^ s^^) is the mutual diffusivity of A with

respect to B. Da.b is the sum of the diffusivities of

A and B measured with respect to a common origin.

The Rabinowitch equation may be written:

kA.B{\ mole-1 s-i) =2zy-^DA.BNKy/No P

where z in the coordination number for the lattice

system (12 for hexagonally close-packed systems),

7, (11/6 for hexagonally close-packed systems

[61]), takes account of the fact that not all difusive

jumps lead to the separation of the particles from

the state of encounter, A^o is the number of lattice

sites in one litre, and P is the mean square imitary

displacement of A with respect to B. Multiplying

through by 1, and assuming that 1 P = P = 2i/2( J/^/AT)

[63], we have

&^,b(1 mole-i s-i) =2X2-i/2xlO-3z7-iiVAv 1

If the lattice model is acceptable the approximations

made in deriving this expression are not serious.

The equation for the random walk model is

feA,B(l mole-i s-') = l<d-hjNAyA 1 where A is the collision cross section for the reaction,

and Vj is the mean jumping frequency. Elementary

laws of diffusion require that Vj = 6DA,B/V [65j.

This expression for Ua.b, and that for Vj, involve the

mean relative jump length and mean square relative

jump length, for the diffusion of A relative to B,

respectively. In accordance with a simple model,

appropriate for small molecules, or for a system that can be regarded as composed of large molecules made

up of small segments, we may write A = P/4. Assuming that P= (1)- [63] we have

fex.s(l mole-i s-i) = (6/4)irXlO-3£)A,BA^Avi

Inserting numerical values into these equations we

have (in 1 mole~^ s~' in each case)

Smoluchowski fe^ , s = 7 . 6 X 1 0-^RDa , b = 1.6X10'' I Da.b

37
Rabinowitch A:^ ,b = 5.5 X lO^' IDa.b

Random Walk kA.B = iXW \ Da.b

In order to stress the similarity between the equations we have assumed that R = \, that is the average

distance between the centres of mass of the molecules.

This is an instructive comparison of the numerical

estimates given by the three models, although the

simplifying assumptions that have been made are

specific to small molecules or to particular models to

be discussed later. Even when applied to small

molecules, however, some of the assumptions are arguable in relation to real systems. Particular

difficulties are involved with the choice of 1, which

we have based on a particular lattice model, the

relation of 1 as we have defined it, to /?, and the

relationship of 1 with average diffusive displacements

to be related to the diffusivity and distances of

separation of reactants during encounter. Only crude

estimates of Ua.b may be obtained in a priori calculations. This is stressed by our rounding off

numberical factors in the above equations.

These expressions hold only for extreme diffusion

controlled reactions, and in so far as they are based

upon correct assumptions, give maximum values for

rate coefficients in liquids.

11. Diffusion-controlled Bimolecular

Reactions and their Converse

Consider the sequence of processes: A-^B A. .B

AB, where A. .B represents the pair of reactants

A and B in a state of encounter, or cage, which may

be formed at a rate kD[_A'^\_B'] from A and B, or at a rate kd\_AB~\ from AB. The state of encounter may

be destroyed at a rate k-D\_A .. B~\ to form A and

B, or a rate kc\_A . . B^ to form AB. ko and k^D are determined by diffusive factors, ko is identical with

kA.B in the preceding section. Application of the

steady -state condition for A . . B, we obtain the

rate coefficient for the dissociation of AB to

A^d i33 = k-okd/ ikc+k^D); kass = kDkc/ {kc-i-k-o)

The equilibrium constant for the reaction is given by

Ke=[^J5]/[^][B] = A:assAdiss= {kc/h) {ko/k.D)

Kc is determined by thermodynamic factors only,

and so are the quantities (kc/kd) and {ko/k-o). Rate

coefficients are diffusion-controlled in cases where kc

is comparable with, or greater than fe_x). In these

circumstances both knas and kasa are affected by

diffusion.

The following question arises, having special

significance in macromolecular systems. For diffusioncontrolled reactions both fcdiss and A^ass have a dominant dependence upon the mutual diffusivities

of the reactants. The terms ko and k^o in which

these dependences appear will be greater for a molecule of given size, if it is formed from (or

decomposes into) a large and a small portion, than

if it is from (or decomposes into) two parts of

comparable size. This arises because, although the

form of dependence may be in dispute, diffusivities

of large molecules (and radicals), are inversely

dependent upon some power of the molecular

weight. The suggestion holds when imiform free

diffusion of the particles is allowed. The effect may

be enhanced if the macromolecules and their decomposition products are entangled. Such entanglements,

of course, directly influence the diffusivities of

macromolecules. Provided then that ko and kd are

less sensitive to the sizes of the reactants than are ko and A:_o, kdiaa may be greater for reactions that

produce a large radical and a small radical, than for

those that produce from the same macromolecule,

two radicals of comparable size. Does this account

for the frequent claims that free radical degradations

are initiated by decomposition of end-groups? It appears that random initiation (by breaking of

bonds throughout the lengths of the chains) becomes

more important at higher temperatures [68, 69].

This could be accoimted for on other bases, but is not inconsistent with the hypothesis implied in the

present question. Very small radicals would have

much smaller energies of activation for diffusion

than do larger radicals. Beyond a certain (relatively

small) size, however, all radicals would have similar

energies of activation for diffusion, although their

diffusivities would differ. We will discuss below the

nature of the diffusional processes that being into

contact the reactive portions of macroradicals, but

supposing for the moment that these are determined

dominantly by translational diffusion, it is interesting to recall that (translational) diffusivities of

small molecules in polymeric systems vary less

rapidly with molecular size (as M^^) than do those

for larger radicals (which vary as M~^-^).

It would, of course, be rash to suppose that kc

and kd are completely independent of those factors

that effect ko and kd can be expected to be

determined by the amplitude and frequency of

vibrations of a molecule, or a portion of a molecule,

about its average position. Such vibrations may

affect the intimacy of the interactions between

neighbours, and so affect the rate of reaction between

such neighbours. The frequency of such vibrations

may be determined to some degree by the size of the

molecule, and the character of the matrix in which

it is situated. Such effects must be small compared

with those upon diffusion, but may not be negligible.

12. Estimates of Rate Coefficients for

Diffusion-controlled Reactions in

Polymer Melts

The expressions for diffusion-controlled rate

coefficients may readily be applied to reactions

between small molecules. Even if for such reactions

there is a requirement that the reactants must

assume a certain relative orientation in order that

reaction may proceed, it is likely that the necessary

condition can be fulfilled during the state of en- counter as a result of the rotations which the reactant

molecules are able to undergo. If, however, two

macromolecules come into contact at a point remote

from their reactive centres, it is unlikely that re38
action will take place during the state of encounter.

We must consider, therefore, the relative motions of

the reactive centres, which are small in comparison

with the sizes of the macroradicals. The diffusivities

that we have discussed, and which are most commonly studied, describe the displacements of the

centres of mass of the molecules. The overall displacement of any part of a molecule during any lapse

of time must be the sum of the displacement of the

centre of mass of the molecule, and of the displacements relative to the centre of mass (strictly,

treated in a vectorial sense). In rigid molecules,

displacements of the latter type are due simply to

rotations about the centre of mass. In flexible

molecules, such as we are concerned with, the

effective rotational displacements will commonly be

accompanied by significant changes in molecular

conformation. We are, therefore, confronted with

the problem of the character of the motions of

macromolecules that are flexible. The internal

motions of flexible macromolecules which lead to

changes in conformation, and those which lead to

displacements of the whole molecule, are closely

interrelated. This question arises in the analysis of

the diffusion-controlled reactions of macroradicals in

free radical polymerisations. It has been concluded

[72] that in such systems the assumption that the

mutual diffusivities of the macroradicals' reactive

segments may be identified with the mutual diffusivities of the centres of mass of the macroradicals, fails

to account for the observed behaviour. It is cus- tomary and is generally satisfactory, to regard

macromolecules in dilute solution as (equivalent)

rigid hydrodynamic spheres [71, 70], in order to

account for transport properties, including diffusion.

In the present context we require a more intimate

understanding of the molecular motions, and it is more satisfying to consider the segment-by-segment

displacements [62]. The situation regarding macroradicals in dilute solution in simple liquids remains

uncertain, if not "something of a mystery" [67].

Ironically the situation may not be so complicated

in systems composed entirely, or almost entirely, of

macromolecules. The approved theory for the

diffusion of large molecules in such systems [27] seems to be consistent with a sinuous step-by-step

movement of the molecules, so that it might be

reasonable to use translational diffusivities of the

whole molecule to represent the diffusion of constituent segments. If this assumption is incorrect, the

estimates of rate coefficients made on its basis, will,

insofar as the expressions are otherwise correct, be

too small. It is our opinion that any error will not

be great.

We have recognised that the reactive centres in

macroradicals are small. In estimating diffusioncontrolled rate coefficients these "targets" can be

regarded as having the dimensions of small molecules,

and diameters of about (3-5) XlO"^ cm. In applying

the expressions, and regarding the "targets" as spheres we overlook that the reactive centres are not

accessible from all directions. Thus our choice of e might in any model exaggerate Ua.b- Using estimates

we have given earlier for diffusivities we find that

diffusion-controlled radical-radical rate coefficients

for large radicals in melts of decomposing polystyrene

are about 10^±^ 1 mole~^ s~'. The (effectively aver- aged) experimental values for rate coefficients for

radical-radical reactions in the corresponding polymerising systems (25-90 °C) are greater, being about

10« to lOMmole-is-i [73]. The rate coefficients

for small radicals may be about 10^ to 10^ 1 mole~^ s~\

comparable with those in simple organic liquids.

13. Free Radical Concentration in

Thermally Degrading Polymers

If a substance C decomposes to form radicals Ri,

which are only able to recombine to form C, the

[Ri]c= ([C]/isLc) (following our conventions

[12]). [Ri]<, is an equilibrium concentration. We

assume that such an equilibrium, and the steadystate concentrations of radicals in the system of

interest are attained very rapidly. The utility of this

equation will be enhanced if the thermodynamic

principles presented earlier in this paper could be

consolidated and exploited. In cases of interest the

free radicals first formed react and are transformed

into other species. Consider first a simple chain

reaction, in which a succession of radicals (designated 1, 2, ) are formed. The type of reaction

is illustrated by polymerisation (without transfer)

in which the successively formed radicals are of

increasing size, or by degradation reactions (without

transfer) in which the successive radicals are smaller.

The steady-state concentrations of radicals are given by [74].

i 3 Let the total concentration of radicals be [R]r, and

[R,]=j,[R]r (Eiri=l)'then

[R]r E E fc«,.wy = fc.n[C]

i 3 where A:,„ is the rate coefficient for the dissociation

of C into freely dispersed radicals. We may define an average value [75] such that

2(/c,)av[R]^ = fe,-.[C]

kt,ij and kt.a and ktjj for diffusion controlled re- actions are related [76, 77, 78]:

2A:(,iy= kt,ii~{'kt,jj

at least to an adequate approximation, following

from the dependence of all expressions for fc^.s upon Da.b{ = Da-\-Db) • The rate of reaction is commonly directly proportional to [R]r. Comparing

this quantity with the estimate of [Ri]e we have

[R]r/[Ril = /cMV(fcOav

Thus [R]r will be larger or smaller than [Ri]e

according as to whether kt,ii decreases or increases

with i. In polymerising systems kt,ii falls with the

decreasing mobility accompanying the increasing

size of the radicals [70]. [RJr therefore exceeds

39
[Ri]]e. The converse is true in systems in which

macroradicals are degrading by the loss of monomer

or larger fragments. The effect will probably be

more marked in degrading systems than polymerising systems in dilute solution. Values of kt are proportional to diffusivities. In polymerising

systems these would vary approximately as M~^'-

if translational diffusion of centres of mass

were dominant, but may depend less markedly on M

[72]. In polymer melts diffusivities vary as about

(^fQj. larger macroradicals) . Unless, contrary to our opinion, complications in the motions of the

macroradicals reduce the dominant dependence of

diffusion-controlled rate coefficients upon the translational diffusivities of the reactants, we must anticipate some consequences in degradation kinetics on

this account. Complications must also follow from

changes in D as reaction proceeds.

These generalisations relate to systems in which

there are no transfer reactions. When transfer

reactions intervene it is more difficult to make

generalisations, for it is necessary to know whether

the radicals formed by transfer have greater or lesser

mobilities than the radicals from which they are formed. In degrading polymers transfer reactions

will commonly involve macromolecules, and will

produce macroradicals comparable in size with, or

larger than, those from which they are produced.

This will narrow the range of molecular sizes of the

macroradicals, and so minimise the range of values

of kt,ii which determine {kt)av There could remain,

however, a dependence of {kt)av on the molecular

weights of the macromolecule from which the free

radicals are derived [79].

14. Geminate Recombination

Our treatment of diffusion-controlled reactions has

assumed a random distribution of reactants throughout the volume of the system. Free radicals produced

by thermal or photo-decompositions are inevitably

produced in pairs. This has the effect already referred

to as the converse of diffusion-controlled reaction

[61]. If, however, a free radical diffuses away from

the primary cage in which it was formed, it will for

the earlier part of its "life" lie closer to its geminate

partner than to any other free radical. During this

time there is significant probability that the twin

radicals will encounter and react with one another.

Both effects have been called "geminate" recombination [80]. It must be recognised that geminate

recombination includes not only the reaction of

radicals that have not initiated reaction chains, but

also the reaction, with destruction, of chain carriers

that have been formed by the successful reaction of

free radicals that were created together [81, 82].

Mutual destruction of geminate free radicals that

have propagated reaction has an effect upon the

kinetics of the overall reaction [81, 83]. The reaction

between (nongeminate) free radicals in general is second order with respect to their concentration.

The mutual destruction of a geminate pair of free

radicals is of first order with respect to their concentration [81, 83]. The second-order process (if

exclusive) leads to a dependence of radical concentration upon the square root of the rate of initiation.

Geminate recombination (if exclusive) leads to a

first power dependence of the concentration of

chain-carriers upon rate of initiation [81, 83]. It is likely that the two processes compete [83].

It has been suggested that geminate termination

of chains occurs in the thermal degradation of some

polymers [84, 85]. In particular it has been invoked

to account for an apparent first order termination

step observed when degradation of polymer was

initiated by a portion of polymer containing deliberately introduced weak bonds [84]. It is desirable

that we should examine some characteristics of

geminate reactions.

The probability that a free radical will reencounter

its twin falls rapidly as the two become more widely

separated. Most diffusive jumps lead to increasing

separation of the particles. Imagine one of a pair of

geminate particles. Only displacements which lie in a direction contained within a small solid angle subtended by the twin particle will lead towards possible

reencounter. A crude argument developed on this

basis [82] suggests that only about l/16th of the

radicals that escape the primary cage are involved in

geminate combination. An important assumption

underlying this and other treatments, is that the

medium is uniform. Mathematical difficulties arise

because we are concerned with random displacements

that are not very large compared with the distances

to be traversed whilst the geminate partners are

close together. Of some importance also might be the

effect of the local structure of the medium. Noyes

[80, 66] has considered the problem in some detail,

and recognised these difficulties. He does not estimate

the fraction of radicals that react geminately after

leaving the primary cage, but implies that this is small. There has, with some justification, been a

call for a better theoretical model [87].

Space does not allow an extensive examination of

the subject. We hold the opinion, however, that in

uniform systems geminate recombination has only a small effect on the kinetics of propagation. Those

radicals that recombine geminately will seldom have

propagated the reaction to a great extent. The

opinion is based on the frequent observation that

rates of reaction are predominantly proportional to

the square root of the rate of initiation. Geminate

recombination is predominantly a problem in

geometry, being the consequence of the distribution

of the reactants in the system, and is independent of

diffusivity. It may be influenced by local structure,

and the occlusion of radicals that has been proposed

to account for phenomena in very viscous systems,

(more likely to occur in glassy systems than in

melts) , may result from non-uniformity of structure,

about which little is understood [88].

Further support for our view comes from the

generation of free radicals from mixtures of initiators

[89]. If two initiators, one generating radicals of

type Ri and the other radicals of type R2, then

products of the type R1R2 can only be formed by

40
radicals that have escaped geminate reactions

(either in the primary cage, or secondary reactions).

If it is supposed that the chances of encounter and

reactions between pairs of freely diffusing radicals is statistical, as in the gas phase, and as it is likely to

be in solution is the radicals Ri and R2 are sufficiently

similar to have almost equal diffusivities, then we

would expect that the distribution of products would

be such that, for the products of reaction

[RiR2]/[RiRi]'/TR2R2]"- = 2

Studies in solution have shown that the products do

not conform to this statistical requirement. If

the deviations from statistical requirement are attributed to geminate reactions. In different types

of measurement, Szwarc and his co-workers [_81^

have shown that up to 70 percent of the radicals

formed from azo-ethanes undergo geminate recombinations. These proportions are unusually high,

compared with initiators used in polymerisation

[90, 91]. These measurements include the proportion

of radicals that react in the primary cage, and which,

because they react mutually so quickly, are unable

to propagate the reaction. They do not, therefore,

invalidate the opinion just given.

The results of a careful investigation of the

polymerisation of vinyl acetate [92] show that at

low conversions about 10 percent of propagating

radicals terminate by a first order process, that is probably geminate [83]. At higher conversions

(above about 40 percent) in the same reaction, when

the system is very viscous, a different type of first order destruction of radicals takes place, probably

due to microheterogeneity of the gel [83]. The contribution made to the overall reaction by free

radicals that react geminately clearly depends upon

their reactivity. Radicals may react geminately with

or without having propagated reaction. In degrading

systems the most significant reaction of the radicals

is to eliminate a portion of themselves. This has the

effect of increasing the distance, and number of

molecules, between the geminate partners. The

small products of the reaction may, of course,

diffuse away rapidly. The effect may reduce, but

certainly not enhance, the contribution of geminate

reactions to the overall kinetics.

It is of interest to examine whether the diffusive

displacements undergone by a degrading macroradical compare with the displacements of the

reactive centre resulting from the extrusion of a fragment from the radical. The rate coefficient for

the expulsion of a monomeric (or larger) unit

from a macroradical is unlikely to exceed about

IQi* exp[- 25,000//? T] s-i {R in cal deg-i mole-i)

[96]. At about 400 °C the mean lifetime of the radical

with respect to this process is about 10~^ s. Assuming

a low value of diffusivity, about 10"'" cm^ s~\ the

rms displacement due to diffusion is about 10~^ cm.

This is small compared with the contraction in size

of the radical, and may be interpreted as indicating

that only a fraction of the radicals (or, more strictly,

of their centres of mass) have moved significantly

during the time required for reaction. The figures

chosen for this calculation are extreme. It might be

that conditions can sometimes be realised in which

there is competition between diffusion and decomposition in displacing reactive centres. The situation seems to differ from most polymerisation systems.

15. Loss of Radicals by Evaporation

Studies of the degradation of polymers are fre- quently made using thin films, often of thickness of

10~^ to 10~^ cm. It is proper to ask whether in such

circumstances the more mobile free radicals may

evaporate from the system, so contributing effectively a first order process for the destruction of

radicals. The suggestion has frequently been made.

Recently [86] a detailed mathematical analysis has

been made of the problem. A simple approach is now

made to the assessment of the likely importance of

the process. The radicals will most readily be lost by

this process if they evaporate so efficiently from the

surface that their concentration at the surface is zero. Consider a slab of thickness b, from which

evaporation may take place at one surface {x = b) and

within which the radicals are produced at a uniform

rate / (radicals per unit volume) [93, 94]. The

average steady-state radical concentration is Ib-/3D.

The average lifetime of radicals for this process

(average concentration/rate of production) is b^/ZD. The highest diffusivities likely to be encountered are about 10~^ cm^ s~'. If 6 = 5XlO~^ cm,

the lifetime is about 10~' to 1 s. The lifetime with

respect to homogeneous second order destruction of

radicals is about (/c([R])~^ (Here [R] is the total

concentration of radicals). Values of ki, even for

reactions involving small radicals, are unlikely to

exceed 10^ 1 mole-^ s-^. We take [R] = 10"^ mole l"' [95]. The lifetime for destruction of the small

radicals by this process must be larger than 10~' s. The two processes could be important where degradation produces small radicals.

16. Final Comments

The study of reactions of small molecules, in the

gas phase, provides information about the simple

reaction steps, such as transfer reactions, and

elimination of olefins, that are involved in free radical

degradation of polymers in melts. In few cases, however, do the studies of the thermal decomposition

in the gas phase provide information that may be

compared directly with the kinetics of degradation

of polymers in melts. There is one case in which a

direct comparison might be profitable. This is the

pyrolysis of larger hydrocarbons, which have been

studied for its industrial importance, and which

might usefully be compared with the thermal

decomposition of polymethylenes. There is evidence

that the decomposition of hydrocarbons in the range

Cio to C30 is not a first order process, and the overall

kinetics may be interpreted [97] on the basis of

proportionality to [hydrocarbon]^^. This indicates

the complexity of the reaction, but is an empirical

simplication. Clearly the pyrolysis involves many

41
reaction steps, which include transfer, and a diversity

of termination reactions. This is also true of the

decomposition of polymethylenes. Isomerisations of

free radicals which are important in both decompositions, and which have been discussed recently

[98, 99] may differ, on account of reduced mobility

in the melts.

17. Notes and References

[1] Moelwyn-Hughes, E. A., Kinetics of Reaction in Solution,

Oxford Univ. Press (1947).

[2] Frost, A. A. and Pearson, R. G., Kinetics and Mechanism, Wiley, New York (1961).

[3] These expressions are embodied in a less explicit form,

derived by van't Hoff, Lectures on Theoretical and

Physical Chemistry, Vol. I, p. 221 (1904). (ref. 4)

[4] Benford, G. A. and wasserman. A., J. Chem. Soc, 362

(1939).

[5] Hildebrand, J. H. and Scott, R. L., Regular Solutions,

Prentice Hall, New Jersey, pp. 41-48 (1962).

[6] Linford, R. G. and Hildebrand, J. H., Trans. Farad.

Soc, 66, 577 (1970).

[7] Trouton, F. T., Phil. Mag., (5) 18, 54 (1884).

[8] Hildebrand, J. H., and Scott, R. L., Solubility of Nonelectrolytes, Reinhold, N.Y. (1950).

[9] Blair, G. M. and Yost, D. M., J.A.C.S., 55, 4489 (1933).

[10] Allen, G. and Caldin, E. F., Trans. Farad. Soc, 49, 895

(1953); Quart. Rev. 7, 255 (1953).

[11] Much of the evidence on the association of NO2 in the

liquid phase comes from early work (Cundall, J. F., Trans. Chem. Soc, 59, 1076 (1891); 67, 794 (1895)).

Some assumptions made in this, and some recent work

(Steese, C. M. and Whittaker, A. G., J. Chem. Phys.,

24, 776 (1956); Addison, C. C, and Sheldon, J. C,

J. Chem. Soc, 1937 (1957)), concerned with the ex- tinction coefficients of NO2 in solution, seem open to question (Patrick, C. R. and Smith, P., unpublished

work).

[12] Consider a system comprising re, moles of each of a set of components (i = 1, 2, . . .) having partial molar vol- umes r i. The total volume is given by ^^r=/_,t Vi,

where the summation, here and later, is made over all components. The total number of moles in the system

is TiT =y^. i Tii. We may define an average molar volume

for the system Vg=VT/nT- The mole fractions are defined by Xi=ni/nT. The concentrations may be

written Ci = ni/VT = ni/nTVs=Xi/Vs- Comparison of

the expressions for Kc and Kx give the result quoted

in the text. We have consistently written Kc

{=cab/cacb)='{AB]/[A][B]; Kx=xab/xaxb. ^^e. are aware of the practice in degradation studies of ex- pressing concentrations in alternative forms, e.g. in terms of mole per base mole. We have not considered

it worthwhile to develop our arguments in these al- ternative terms.

[13] The properly defined free energy change for the reaction

is (using simplified notation) AGj = — jRT In Kj, and

the proper enthalpy and entropy changes are Af/i = RT\d In KJdT) and ASx = d[RT In K^/dT. Hthe heat

of reaction is determined from the equilibrium constant

expressed in terms of concentrations we have A/fc =

R'h{d In Kc/dT)=AHr+ocAnRT\ where

a=VrKdVs/dT)

is the coefficient of expansion of the solution. The

corresponding entropy change is

ASc = ASx+AnR In V, +AnaRT.

The terms containing a are often small, and may be

neglected in comparison with experimental error. The

equation for AHc was given by Guggenheim, E. A., Trans. Farad. Soc, 33, 607 (1937).

[14] Benson, S. W., Thermochemical Kinetics, Wiley, New

York, (1968).

[15] Benson, S. W. The Foundations of Chemical Kinetics,

McGraw Hill, New York (1960).

[16] Martin, H., Angew. Chem., (Int. Edn.), 5, 78 (1966).

[17] Khambata, B. S., and Wassermann, A., J. Chem. Soc,

371 (1939); Benford, G. A.; Kaufman, H., Khambata,

B. S., and Wassermann, A., J. Chem. Soc, 381 (1939);

Kaufman, H. and Wassermann, A., J. Chem. Soc, 870

(1939).

[18] Offenback, J. A., and Tobolsky, A. V., J.A.C.S., 79, 278

(1957).

[19] Carlsson, J., Howard, J. A., and Ingold, K. U., J.A.C.S.

88, 4726 (1966).

[20] Batt, L. and Benson, S. W., J. Chem. Phys., 36, 895

(1962) see also ref. (14)

[21] Pritchard, G. O., Pritchard, H. O., and TrotmanDickenson, A. F., J. Chem. Soc, 1425 (1954); Blake,

A. R., and Kutsche, K. O., Cna. J. Chem., 37, 1462

(1959); Long, J., Skirrow G., Trans. Farad. Soc, 58,

1403 (1962).

[22] Cowley, P. R. E. J., and Melville, H. W., Proc. Roy.

Soc, A210, 461 (1951); A211, 320 (1952).

[23] Florin, R. E., Parker, M. S., and Wall, L. A., J.Res. NBS, 70A, 115 (1966).

[24] Bueche, F., J. Chem. Phys., 20, 1959 (1952).

[25] Jellinek, H. H. G. and Luh, M. D., European Polymer

J., Supplement, 149 (1969).

[26] Benson, S. W., and North, A. M., J.A.C.S., 81, 1339

(1959).

[27] Bueche, F., "Physical Properties of Polymers", Interscience, New York, (1962).

[28] Bueche, F., Cashin, W. M., and Debye, P., J. Chem.

Phys., 20, 1956 (1952).

[29] Fox, T. G., and Loshaek, S., J. Appl. Phys., 26, 1080

(1955).

[30] Tobolsky, A. V., "Structure and Properties of Polymers",

Wiley, New York, (1960).

[31] Williams, M. L., Landel, R. F., and Ferry, J. D., J.A.C.S.

77, 3701 (1955).

[32] Ferry, J. £)., "Viscoelastic Properties of Polymers",

Wiley, New York (1961).

[33] The inclusion of Tg in this expression does not invalidate

its applications to polymers that highly crystalline at low temperatures. We are interested in melts.

[34] Fox, T. G., and Flory, P. J., J. Polymer Sci., 14, 314

(1954).

[35] Bueche, J. Appl. Phys., 26, 738 (1955); 24, 423 (1953).

[36] Tobolsky, A. V., Aktomis, J. J., and Akavali, G., J. Chem. Phys., 42, 223 (1965).

[37] Nishioka, A., and Watanabe, M., J. Polymer Sci., 24,

298 (1957).

[38] Tobolsky, A. V., Katz, D., and Eisenberg, A., J. Appl.

Polymer Sci., 7, 469 (1963).

[39] A(T) includes the term (R^/M). It is often asserted that

PTFE molecules exist in a more extended form than

do other macromolecules. Assuming that the expression

D-n=A{T) holds for this polymer, which we have

questioned, A{T) might have a larger value than for other polymers. Estimates of D, might on this account,

be small, but by less than an order of magnitude.

[40] Case, L. C, J. Appl. Polymer Sci., 3, 254 (1960).

[41] Bueche, F., J. Chem. Phys., 48, 1410 (1968).

[42] The labelled material generally constituted less than

10 percent of the whole in this work.

[43] The values quoted for Tg are generally those for polymers

of high molecular weight. Since there is a dependence

of Tg upon molecular weight [30, 32], the values may

be too low for polymers of low molecular weight (less than 10'^).

[44] Wood, L. A., J. Polymer Sci., 28, 319 (1958).

[45] Jellinek, H. H. G., J. Polymer Sci., 4, 13 (1949); Madorsky, S. L., J. Polymer Sci., 9, 133 (1952).

[46] Madorsky, S. L., J. Polymer Sci., 11, 491 (1953).

[47] McCall, D. W., and Douglass, D. C, J. Chem. Phys.,

31, 860 (1959). These results for cases of M= 5800

and 4100 imply that aM'^ as required but that DaM'"-^

which is not in accord with Dr) = A(T).

42
[48] Meares, P., J.A.C.S., 76, 3415 (1954); Trans. Farad.

Soc, 53, 101 (1957).

[49] For a brief survey see Meares, P., Polymers, Structure and Bulk Properties, van Nostrand, London, (1965).

[50] Lundberg, J. L., Wilk, M. B., and Huyett, M. J., J. Polymer Sci., 57, 275 (1962).

[51] Lundberg, J. L., Wilk, M. B., and Huyett, M. J., J. Appl. Phys., 31, 1131 (1960); Newitt, D. M., and

Weale, K. E., J. Chem. Soc, 1541 (1948).

[52] Zhurko, S. N., and Ryskin, G. Ya., J. Tech. Phys.,

USSR., 24, 797 (1954).

[53] Dannis, M. L., J. Applied Polymer Sci., 1, 121 (1951).

[54] Fujita, H., Kishimoto, A., and Matsumoto, K., Trans. Farad. Soc, 56, 424 (1960).

[55] Alexopoulis, J. B., Barrie, J. A., Tye, J. C. and Fredrickson, M. Polymer, 9, 56 (1968).

[56] Auerback, L, Gehman, S. ID., Miller, W. R., and Kuryla,

W. C, J. Polymer Sci., 28, 129 (1958).

[57] It is to be expected from theoretical analysis of the problem [27] that until the diffusing molecules become so

large as to entangle with molecules composing the

matrix, D will vary inversely as M. The upper limit

for this behaviour is M= 10^-10^

[58] Hayes, M. J. and Park, G., Trans. Farad. Soc, 51, 1134

(1955).

[59] Smoluchowski, M. von, Z. Phys. Chem., 92, 129 (1919).

[60] Noyes, R. M., Progress in Reaction Kinetics, Vol. I., ed. G. Porter, Pergamon, London (1961); also for example Yguerabide, J., Dillon, M. A., J. Chem. Phys.,

40, 3040 (1964)).

[61] Rabinowitch, E., Trans. Farad. Soc, 33, 1225 (1937).

[62] Allen, P. E. M. and Patrick, C. R., Macromel. Chem.,

47, 154 (1961).

[63] It is arguable that the mean unitary diffusive displacements are smaller than the dimensions even of small

molecules. (Watts, H., Alder, B. J., and Hildebrand,

J. H., J. Chem. Phys., 23, 659 (1955)). There is little doubt that the assumption that all unitary displacements are equal is an over -simplification. The simplest

plausible assumption regarding intermolecular dis- tances is to suppose that the molecules are spherical,

and are hexagonally close-packed. A cube of side 1/2

contains 4 molecules, so that V/N = H/2^'''. For a simple liquid the molecular concentration is about 10

mofe 1~\ so that 1 is about 5X 10~* cm. Diffusive jumps

could plausibly be smaller than this. Clearly e varies

with temperature. The assumptions that V = (l)^ etc., are very crude, for they apply only to strict lattice models. Jump lengths give most difficulties in random

walk models.

[64] Allen, P. E. M., and Patrick, C. R., Makromol. Chem.,

47, 154 (1961).

[65] The expression x'^=2 Dt is the most familiar form for the mean square displacement of a particle. It relates to displacements in one dimension (A. Einstein, Z. fiir Electrochem., 14, 235 (1908). The expression r'^ = 6 Dt

relates to displacements in three dimensions (A. Einstein, Ann. Phys., 17, 549 (1905).

[66] Noyes, R. M., Progress in Chemical Kinetics, Ed. G.

Porter, Pergamon, London (1961), p. 129.

[67] North, A. M., The Collision Theory of Chemical Reactions in Liquids, Methuen, London, (1964).

[68] Bagby, G., Lehrle, R. S., and Robb, J. C, Polymer, 10,

683 (1969); Makromol. Chem., 119, 122 (1968).

[69] Wall, L. A., Florin, R. E., Flynn, J. H., and Mclntyre,

D., paper seen in manuscript,

[70] Tanford, C, Physical Chemistry of Macromolecules,

Wiley, New York, (1961).

[71] Flory, P. J., Principles of Polymer Chemistry, Cornell

Univ. Press, Ithaca, N.Y. (1953).

[72] Benson, S. W. and North, A. M., J.A.C.S., 84, 955

(1962). This work assumes as one of its models, that

fci,A, B, =constant(RA-"^ — nB-"^) where and ub are measures of the sizes of the respective macroradicals,

consistent with observations that approximately D =

constant re"''^, relating to translational diffusion [64].

[73] see, for example Bamford, C. H., Barb, W. G., Jenkins, A. D., Onyon, P. F., The Kinetics Vinyl Polymerisation

by Radical Mechanisms, Butterworths, London (1958).

[74] Consistent with this notation Kc = 2kt,ii/ki„; we follow

the convention +d[Ri]/dt = 2kt,ii[Ri]^ in the reaction

Ri+Ri.

[75] Consistent with our assumption, write kt,ii=gj; then

(fc«)av=LE(g. -&)rfJ;=2X)^ giyiJy=2X; giji

' i i i i

If gi+i>gi, all i, then {kt)av>kt,ii; and conversely.

[76] Poison, A., J. Phys. Chem., 54, 649 (1950); Schick, A. F., and Singer, J., J. Phys. Chem., 54, 1028 (1950).

[77] Allen, P. E. M. and Patrick, C. R., Makromol. Chem.,

48, 89 (1961).

[78] An alternative relationship often assumed is of the form

kij = {kiikjiy'^4>ij. This is inappropriate for diffusion- controlled reactions [64]. [79] It has recently been suggested (Ito, K., J. Polymer Sci.,

lA 7, 827 (1969)) that the effect of changing radical

size upon diffusivities, and in turn upon diffusion- controlled rate coefficients (specifically in polymerisation systems), can be taken into account by the

solution of the (Smoluchowski) equation for diffusioncontrolled reaction using a time-dependent expression

for the diffusivities. Plausible as this may appear it is not so correct as that described in ref . [72] in which the

distribution of macroradical sizes is set down and used

to evaluate the required average, (fc()av. [80] Noyes, R. M., J.A.C.S., 77, 2042 (1955).

[81] Allen, P. E. M. and Patrick, C. R., Nature, 191, 1194

(1961).

[82] Allen, P. E. M., and Patrick, C. R., European Polymer

J., 1, 247 (1965).

[83] Allen, P. E. M. and Patrick, C. R., Trans. Farad. Soc,

59, 1819 (1963).

[84] Richards, D. H., and Salter, D. A., Polymer, 8, 139

(1967).

[85] Penski, E. L. and Goldfarb, J., J. Polymer Letters, 2, 55

(1964).

[86] Florin, R. E. and Wall, L. A., paper communicated privately in manuscript.

[87] Dobis, O., Pearson, J. M. and Szwarc, M., J.A.C.S., 90,

278 (1968); Pearson, J. M. and Szwarc, M., J.A.C.S.,

90, 283 (1968).

[88] Burnett, G. M. and Loan, L. D., CoU. Czek. Chem.

Comm., 22, 113 (1957).

[89] Bowles, Rm., Majer, J. R., and Robb, J. C, Nature,

187, 314 (1960).

[90] Bevington, J. C, Radical Polymerisation, Academic Press London (1961).

[91] Bevington, J. C, and Stamper, D. J., Trans. Farad. Soc,

66, 688 (1970).

[92] Bengough, L. W., Trans. Farad. Soc, 58, 716 (1962).

[93] The equation is Dd'^[R]/dx''-= ^I, which needs to be

solved for the boundary condition

J{b) = +D{d[Rydx)i=+Ib.

Total loss of radicals by evaporation must equal rate of

generation. The solution is [/i] = /(6^+a;'^)/2£).

[94] It is tempting to ask whether in such films the surface on which the film is deposited may influence the

reaction.

[95] This is a guess, in our experience [jR] seldom is much

greater than 10"* mole 1~^

[96] Preexponential factors for most free radical decomposition (of the type we are considering lie in the range

10'^ to 10^^s""S and are often closer to the lower Umit.

They are usually associated with larger energies of

activation (closer to 30 kcal mole"'). For recent values

see Kerr, J. A., Ann. Rep. C.S., 65, 189, 64, 73.

[97] Woinsky, S. G., Ind. Eng. Chem. Proc. Des. Developments, 7, 529 (1968).

[98] Kysel, O., Romanov, A., and Durd'ovic, V., European

Polymer J., Supplement (Bratislava Conference), 323

(1969).

[99] For example. Done, F., and Guiogion, J. Chim. Phys. 65,

395 (1968).

43
Discussion

R. E. Florin (National Bureau of Standards)

It would appear that many discussions of diffusion

controlled reactions of polymers focus on the motion

of chain ends. Segments in the middle of a polymer

chain ought to have lower diffusion-controlled

reaction rates than those at chain ends. Computer

simulations of polymer segment diffusion [P. H.

Verdier, J. C.hem. Phys. 52, 5512 (1970); J. Computational Phys. 4, 204 (1969)] could in time lead

to some estimate of the difference. I would be

interested in Dr. Patrick's view of this question.

Dr. C. R. Patrick, University of Birmingham,

England (communicated).

In my paper I have discussed briefly the diffusive

motions of macromolecules. Whilst admitting the

possibility that the overall diffusive displacements of

the segment of a macromolecule in a given time may

be greater than those of the centre of mass of the

whole molecule, I ventured the opinion that the

diffusion of a segment of a macromolecule may be

described by the diffusivity of the centre of mass of

the molecule without too great and error. This

practice was followed in subsequent calculations.

Some further comments on this question may be

appropriate.

Macromolecules or macroradicals of interest are

flexible and able to take up a large number of

alternative configurations. In detailed treatments of

dynamic properties of macromolecules, particularly

those relating specifically to diffusion and changes in

configuration, a macromolecule is regarded as composed of a set of segments, linked together, but each

of which may move independently of other segments.

Some models regard the molecule as composed of

segments, all equal in size, arranged so that the

centre of mass of each segment is located on a lattice

point. Diffusion occurs by the jumping of a segment

to an adjacent site on the lattice. It may be shown

[1] that such a model requires that the diffusivity

of the centre of mass of the macromolecule, Dc.o.m.,

is related to that of any segment, ^segment, by

^^c.o.m. -t^segment ^q 1

where is the number of segments in the macromolecule. The detailed proof of this equation, as usually given, follows from the expression of the

diffusivity ascribed to each segment in terms of the

friction coefficient involved [1, 2]. Whilst this is clearly correct, the equation may be derived more

directly. Consider a macromolecule composed of

segments, each of which has a diffusivity Dsegment.

It may readily be shown that any displacement, 8, undergone by a segment, must cause a displacement

i5/N) of the centre of mass of the molecule [2].

Assuming that all unitary length, we may write

jDsegment = vd'^/ 6

where v is the jumping frequency, so that for the

centre of mass

Do.o.n.=pN{5/Ny/6

from which eq 1 follows immediately.

It is normally assumed that the segments are equally free to move, and are of equal size. These

assumptions may not always be realistic. It is possible, for example, that the motions of some

segments may be impeded by other portions of the

same molecule. If this possibility is recognised,

eq 1 may be generalised to give

-^ZJc.o.m. (-Dsegment) av where the latter term is an average of all segmental

diffusivities. In such circumstances it would be

difficult to know the diffusivity of a particular

segment which contains a centre of chemical re- activity.

The definition of a segment as a portion of a macromolecule that is able to move independently, is a convenient mathematical device. A segment

cannot readily be identified with a specific chemical

unit within the molecule. It cannot, for example, be

identified with a monomer imit in a polyvinyl

compound. The property of independent movement

ascribed to segment imposes a severe stricture upon

the motions of segments within the polymer chaia.

Such motions are limited to changes of conformation

of a segment without moving the positions of its ends,

and to rotational motions of a segment about the

line joining its ends. It is conceivable that a particular atom in a macromolecule may move (in a diffusional manner) with different combinations of other

atoms in successive jumps or sequences of jumps.

The definition of a segment, when generalised, has a

statistical character, similar to that of the subimit

of a macromolecule considered in the enumerations

of the statistics of chain configurations. It would be

unwise, however, generally to identify a segment in

the present context with the unit involved in chain

configurations. Evidence for the sizes of segments,

and the relation of N to molecular weight, must be

sought elsewhere, and may, perhaps be found in the

analysis of other dynamic properties of polymers.

Equation 1 has the satifying property that it allows for the mobility of segments within a very

large molecule, such as may occur in a cross-linked

system. The movement of the centre of mass of

such a molecule may be barely perceptible although

displacements of its constituent parts may be

significant (as may be seen by setting A'^ equal to

infinity). Clearly the setting of TV equal to unity is equivalent to regarding the molecule as rigid.

Alternatively, as an excellent approximation we may

write Do o m. = i)segment for the special case, referred

to in the paper, in which the diffusion of a macromolecule proceeds by segment-by-segment displacements, in which the end segment of the molecule

may move forward and drag all other segments

behind it, so that each moves to the space previously

occupied by its nearest neighbour. Since the end of

44
the molecule first to move in a displacement may

move into any of a (possibly small) number of

positions, each displacement must lead to a new

conformation of the molecule, so that over the

course of time the molecule may take up all the

accessible conformations.

The consequences of eq 1 in radical-radical

reaction kinetics need to be considered (for the

possible cases that N^l). Assessment is needed of

the error involved in making the assumption,

previously made, that -Dc.o.m. = ^^segment. It is needed

both for reactions between macroradicals in melts

and in solution in simple liquids. The problems in the

two cases may differ. For example the diffusivities of

macromolecules in simple liquids vary approximately

as (molecular weight) (see references in paper).

Since (imless A'^=l) we would expect N to be

proportional to molecular weight, it follows that

£^segment should approximately be proportional to

(molecular weight)''', which appears absurd, and

poses a problem for closer examination. The difficulty

is related to the fact that eq 1 is essentially related

to the so-called "free-draining" model for macromolecules in solution p], whilst the simple transport

properties of macromolecules in solution are better

accounted for in terms of "equivalent hydronamic

spheres" |^4, 5^. The situation is further complicated

by the finding that viscoelastic properties of polymers

in dilute solution are satisfactorily accounted for in

terms of the free-draining model [6].

References

[1] Verdier, P. H., J. Chem. Phys., 52, 5512, (1970); 45,

2118 (1966).

[2] Bueche, F., Physical Properties of Polymers, Interscience,

New York, 1962.

[3] Debye, P., J. Chem. Phys., 14, 636 (1946).

[4] Kirkwood, J. G., and Riseman, J., J. Chem. Phys., 16,

565 (1948).

[5] Flory, P. J., Principles of Polymer Chemistry, Cornell

University Press, 1953.

[6] Ferry, J. D., Viscoelastic Properties of Polymers, Wiley,

London, 1961.

45

NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Pyrolysis of Polymers

Leo A. Wall

Institute for Materials Research, National Bureau of Standards, Washington, D.C. 20234

The nature of their pyrolysis products, monomer, oligomers and carhonaceous residues enables one to group polymers into three classes, those that decompose by net main chain scission; by stripping of the main chain, for example the thermal dehydrochlorination of polyvinyl chloride;

and by crosslinking of the main chain followed by some production of volatiles. Highly unsaturated

or aromatic chains tend to follow the latter course. At the present time a theoretical framework exists which permits, provided adequate experiments are performed, the elucidation of the decompositions of the first type. For most of the well- known polymers in this class, this framework of knowledge gives very acceptable mechanistic

explanations or interpretations of the decomposition process based on observations of rate of

weight loss, of molecular weight changes and composition of volatile products.

Knowledge of the decomposition of the second class of polymers is at an intermediate state development. However, an important practical research objective would be the acquisition of

methods for converting the mechanisms of decomposition of class one substances to that for

class two type.

For the third class of materials, knowledge of their pyrolysis mechanisms is nonexistent.

This is largely due to the fact that methods for quantitatively following soUd-phase processes of

decomposition are relatively difficult and unsatisfactory.

Key words: Chain scission; crosslinking; polymers; pyrolysis; stripping.

Under vacuum or moving inert atmospheres

polymers decompose to produce volatile products

and char to various degrees [1, 2, 3j.^ Polymers can be divided into three groups according to the

characteristics of their decomposition process (see

table 1).

1. Saturated Polymers

The saturated vinyl type polymer decomposes

either by some form of chain scission or by the loss of

some substituent group through a chain stripping

reaction. The type of reaction is determined by which

bonds in the molecule are most easily thermolyzed.

If the bond to a substituent is easily ruptured, we

get a dehydrohalogenation type reaction leaving an

unsaturated and possibly aromatic carbon system

which then crosslinks to form

-C

X

-cYor

X X

-c —cX Y

eventually a carbonaceous char. On the other hand,

if a carbon-carbon bond ruptures, we get sometimes

monomer, but more often a variety of products,

including dimer, trimer, etc. Table 2 lists monomer

yields for some linear polymers and heats of poly-

* Figures in brackets indicate the literature references at the end of this paper.

merization. It is clear that the more substituted

structures give higher monomer yields. They also

have low heats of polymerization, degrade under

high energy radiation, and are the least thermally

stable. This generalization is true only if we restrict

ourselves to hydrocarbon polymers. Polytetrafluoroethylene yields, imder high vacuum conditions,

100 percent monomer and has a heat of polymerization of 40 kcal. The precise mechanism of

decomposition for substantially pure polymer would,

of course, depend on temperature among other

things.

It would also depend on the pyrolytic procedure

and the geometrical arrangement of sample and

furnace. In order to emphasize the effect of molecular

structure, our studies are usually carried out by

following the rates of volatilization or vaporization

under high vacuum conditions by recording the

weight of sample lost (gravimetry)

Vaporization phenomena can be divided into three

categories depending on the particular step in the

process that is rate controlling. The three categories

are: (1) rate proportional to surface area, (2) rate

proportional to surface area and inversely proportional to thickness, and (3) rate proportional to

weight of sample. The first category is simple

molecular vaporization in which the rate determining

step is the escape of well-defined molecules from

their liquid or solid phase into the gaseous phase.

At a given temperature the rate of volatilization

depends on the fraction of molecules in the surface

layer that have an energy equal to or greater than

47
Table 1. Types of thermal decomposition

Type Products Examples

Chain scission <99% volatile monomer

Species <2000

M„

-CH2CH2-

alkenes

-CH2CH-

>CF2CF2

->alkanes +

H

^CH2 = C, Di,

Stripping Volatiles and

char

trimer, etc.

-CH2CHC1~-^HC1 +

aromatics +char

-'CH2CF2~-^HF+

aromatics +char

Crosslinking Little volatiles,

much char

the minimum kinetic energy needed for escape from

the surface. In the case of water, or pentane, the

kinetic energy must be greater than the energy

resulting from the Van der Waals type of attractive

forces. In the case of a metal the process is the same

except the attractive forces between the metallic

atoms are larger than the usual Van der Waals

forces. One may also consider diamond volatilizing

as C atoms, except that diamond converts to graphite

between 500 to 1000 °C and does not volatilize to

give 100 percent pure C atoms but probably gives

various amounts of C2, C3 species, depending on

conditions.

These systems differ basically in the amount of

energy involved in the evaporation process. In each

the overall process is:

^liquid or solid ' ^gasFew, if any, studies of polymer degradation are concerned with establishing whether the process

under study is obscured by molecular vaporization

or by diffusion processes. The molecular vaporization

of large organic molecules has been examined by us in

order to determine the upper molecular weight

limit for molecular vaporization without decomposition for various organic structures. We believe

information in this area is critical for the full elucidation of polymer pyrolysis.

2. Molecular Vaporization

The molecular vaporization of three groups of

compoimds have recently been investigated. They

(1) various organic plasticizers [4], (2) normal

(3) various triglycerides and

are alkanes [5], and

sebacate esters [6].

According to kinetic theory the rate of condensation, dn/dt, of a vapor in moles per sec is given

by [7]

dn

dt

ocpA

(27rM/?T)i/2 (1)

where a is the accommodation coefficient, assumed

here to be one; p the pressure, dynes per cm^, of the

substance in the gas phase; A the surface area, cm^; M the molecular weight; R the gas constant in

ergs per degree Kelvin; and T the absolute temperature. Since, when a vapor and liquid phase are

in equilibrium, the rate of vaporization equals the

rate of condensation, we can write the same equation

for the rate of vaporization except that now p is the

equilibrium vapor pressure at the temperature T.

The rate of vaporization we are concerned with is that measured under vacuum conditions and is sometimes referred to as the maximum rate of

vaporization. Since the vapor pressure, in atmospheres, is the equilibrium constant for the process,

we may write

(2)

For one mole of vapor, PAV=RT, and hence the

rate of volatilization, can be written as dw

dt

16.25 jIqASvIRq-AEvIRT^

(3)

where dw/dt is the rate of vaporization in grams per

second assuming that a, the accomodation coefficient,

is one. From the preceding relationship, the Arrhenius

slope is found to be

d \n{dn/dt)

d{l/T)

T

2 R

(4)

From our rate measurements then we can obtain,

according to kinetic theory, the internal energy,

A£^„, for vaporization and also the enthalpy Ai/„. If

the apparent Arrhenius activation energy is designated as S„, then

and

smce

A£„ = 8„ +

RT

3RT

-Rd \n (dn/dt)

dil/T) (5)

Using these relations and our observed apparent

activation energies, values were calculated for the

48
enthalpies of vaporization AH^, at the mean temperature of our experiments. For n-pristane, ra-tetra- cosane, n-hexatricontane and n-tetranonacontane

the values were 18, 23, 29, and 47 kcal/mol, respectively, and did not appear to be in line with literature

values at first sight.

In order to evaluate our experimental data and

hence our technique, we reviewed and recalculated

values of Aif„ from data [8-17] in the published

literature. The most extensive source of original

data was the compilation made by Stull [8]. Direct

calorimetric or Knudsen technique measurements

were available for only a limited number of alkanes,

Ct to Ci8, and only at a single temperature. The bulk

of the AHv values obtained were from vapor pressure versus temperature data either in its original form or

in the form of the Antoine equation. Examples of the

latter are those found in reference [9].

Greatest weight was given to AHv values determined from original vapor pressure data through the

slopes of logp„ versus l/T plots at selected temperature, viz., 0, 50, 100, 150, 200, and 250 °C using

the Clapeyron equation corrected for the compressibility factor AZ. The relationship is

AH. = - 2.303/?AZ . (6) al/i

Estimates of AZ as a function of l/T are tabulated

[18] in the literature. This tabulation is apparently

based on the relationship

1^ =1.0-0.05P(atm) (7)

at 25 °C. As no values of p versus T data above

1 atm pressure were used in deriving AHv values,

AZ was always >0.95. The error associated with

this method should be no greater than 1.5 percent

at 1 atm.

The calculated AH„ values were first plotted as a

linear fimction of the number of carbon atoms, n, in the alkanes. Over a range of species with 4 to 24

carbon atoms the plots for several different temperatures did not yield straight lines. The data

appear to be best fitted by straight lines for every

eight or so consecutive species as shown. The breaks

in linearity seem to occur at about every 7-8 carbon

atoms. This suggested a coiling or folding of the

vaporizing molecules as was predicted by Langmuir

[19]. Also, Muggins [20] in 1939 predicted that the

heats of vaporization of linear alkanes should vary

as the 2/3 power of the number of carbon atoms.

Therefore, we plotted the data as a function of n^'^ and obtained excellent straight lines. It was evident

that at a given temperature the variation in heats of

vaporization for the linear alkanes is very closely

approximated by a linear function of the 2/3 power

of the number of carbon atoms. The straight lines

obtained are given by the relationship

Ai/„ = 3.21n2/3-0.0193r+2.92, kcal/mol, (8)

where T is degrees Kelvin.

Table 2. Monomer yield from the decomposition of polymers

Polymer Structure Monomer

yield

Wt. %

Heat of

polymerization

Kcal

CH3

a methyl styrene

1 —CH2—

C

U

D

99 9

a deuterostyrene

1 —CH2—C—

1 U

H

70 17

styrene

1 —CH2—C—

1 U

CH3

42 17

isobutylene

1 —CH2—c—

1 CH3

25 15

propylene

1 —CH2—C—

1 H

2 17

ethylene —CH2CH2

0.1 22

The n^'^ function is highly acceptable on the

physical basis that a linear molecule will find it energetically easier to vaporize when coiled into a sphere than as a long extended molecule. The

spherical model entails the rupturing of a minimum

amount of Van der Waals bonding during the

vaporization process. According to this model and

the above equation, the increment in AH„ per

carbon atom or methylene group decreases with the

size of the molecule. By differentiating eq (8) with

respect to n, it follows that the increment in AH.

per carbon atom is found.

A(Ai/„) =2.Un-'i\

Thus, the increment at 10^ carbon atoms is 214

calories per mole.

In table 3 we summarize our experimental data

and show the numerical agreement between the

experimental activation energies and values calculated from the equation

8„ = 3.21fi2/3_o.0223T+2.92, kcal/mol. (9)

Both the experimental and calculated values are

49
Figure 1. Rate of vaporization of tetranonacontane (C94Hj9o) and Arrhenius plot of data, activation energy 46 kcal/mol (quartz

bucket 10X4 mm i.d., 10 mg sample) [5].

compared at the mean experimental temperature,

T; see last column. The agreement is very good

particularly for the three smaller compounds and

supports our premise that correct values of the

energies of vaporization can be obtained quickly

with the presently available equipment. The low

result for the largest species, the C94, is still quite

good considering the degree of extrapolation involved. At such a high n value, one suspects that

eq (8) will give high values since one anticipates

spheres of greater density, the greater the value

of n. Experimental rates of vaporization at the

temperature given in the last column are also listed.

The rate of decomposition pij of polyethylene is usually measured in the region of 400 °C and is proportional to the weight of polymer, not to surface

area. The products of thermal volatilization of

linear polyethylene assuming a random mechanism

indicate [41] that 72 carbon atom species are about

the minimum size that decompose before evaporating.

From eq (8) it follows that the internal energy

/^Ev for vaporization at 400 °C is for n = 72, 43.5 kcal;

for 71 = 129, 70 kcal; and for fi = 154, 80 kcal. The

value 43.5 kcal is closer to the activation energy

values of the possible free radical decomposition

reactions constituting part of the overall thermal

decomposition mechanism. The actual overall ac- tivation energy for thermal decomposition is about

70 kcal and the usual carbon-carbon bond dissociation energy is 80 kcal. These results show that

molecular vaporization of linear alkanes comprised

of 100 or so carbon atoms can occur without decomposition. For several experiments with the C94

compound a gravimetric apparatus with an attached

mass spectrometer, capable of monitoring masses from 12 to 120, was u ed and no hydrocarbons from

the C94 sample were observed. This supports, as does

the observed kinetic behavior, the conclusion that

the C94 volatilized molecularly without appreciable

thermal decomposition. In figure 1 we show both the

rate of vaporization for tetranonacontane (C94H190) as a fimction of fraction vaporized, the horizontal

line typical of molecular vaporization, and the

Arrhenius plot of the data, the open circles.

Up to molecular weights of '^900 the energies of

vaporization of both the branched triglycerides and

the linear sebacate esters follow rather closely the curve for the linear alkanes, i.e., eq (8) [6]. This

again is presujjably a result of the coiling of the

molecules, both branched and linear, into spheres of

similar size during the vaporization process.

3. Diffusion Controlled Vaporization

Vaporization may in some cases be diffusion controlled, the second category. In the extreme limit, if every vaporizable molecule leaves the exposed sur- face immediately, then the rate at which they

arrived at the surface determines the rate of vaporization, for example, a volatile organic liquid in a

plastic vessel. The rate of vaporization would, in the

simplest case, be proportional to the outer surface area and inversely proportional to the wall thickness

of the vessel. This kind of process is rare in studies

of polymer pyrolysis, where sample size and thickness

are kept small in order to eliminate the effect. It can produce delayed effects during the start up [22]

or stoppage of pyrolytic processes. For example,

monomer [22] or other products must build up to a steady state concentration before the steady rate

of volatilization is achieved. Here we have essentially

digressed to the questions of ordinary diffusion of

50
small molecules through polymer. Later on we shall

mention several ways diffusion can alter the competitive situation between the various elementary

reactions in the degradation mechanisms.

4. Degradative Vaporization

The third category, in which we have been most

interested in the past, applies in general to practically

all high polymers. Because polymeric solids or

liquids are, in the main, long linear molecules, the

vaporization phenomenon is basically different from

say graphite or water. Polymer molecules are, in

general, too large for vaporization to occur without

decomposition of the molecules. Thus, high polymers

degrade to relatively low molecular species in order to

volatilize. Chemical bonds must break, which they

do by a variety of mechanisms. Regardless of the

mechanism, however, the rate of volatilization is determined basically by the rate of bond rupture.

Since the number of bonds is proportional to the

weight of material, rates of volatilization for polymers

are, ideally, proportional to the weight of polymer,

not to the surface area or thickness. The rate

determining step in the volatilization process for

high polymers having no low molecular weight

impurity is the degradation process. This is not to

say that, depending on various factors and conditions, diffusion phenomena may not play some role

in retarding the overall evaporation or altering the

mechanism and products of polymer degradation.

Schematically, then:

polymer degradation^ Small

molecules (cond. phase) diffusion

gjjjall molecules (at

surface) evaporation

gmaH molecules (gas)

In a particular system, diffusion or evaporation can be rate controlling. With high polymers, degradation is usually rate controlling but this should be

confirmed experimentally.

The critical size, L — 1, above which degradation

should occur, has been roughly estimated in the past

by equating the heat of vaporization to the bond

strength, 9.92 + 1.18(L-l-8) =80, where 80 is the

usual carbon-carbon single bond strength in kilocalories, 9.92 is the heat of vaporization of octane,

and 1.18 is the increment above ra-octane in the heat

of vaporization per methylene unit. The value thus

obtained for L — 1 is 67 which is close to the value of

72 deduced from experimental data on the thermal

volatilization of polyethylene. Recent work has

shown that the heats of vaporization of the n-alkanes

are not linearly proportional to the size of the alkanes.

Based on the new studies the heat of vaporization of

72 carbon atom alkanes is approximately 44 kcal/mol.

Experimentally, we define the conversion, i.e., the fractional weight lost, in terms of the weight,

W, of the polymer

cw=i-^. (10)

In theoretical treatments for the volatilization of

monodispersed molecules, each containing A'^ units,

we write

C{t) = lE iQ^{t)

NQMiO)

' (11)

It is readily seen from the term with the summation

sign that L is defined as the smallest species which

must decompose in order to evaporate. For the same

system the number average degree of polymerization

is given by

Pn{t)

E iQ^it)

EQiit)

L

(12)

Rates of decomposition of the polymers would, in

general, be measured at temperatures where the

rate of weight lost is approximately one percent per

minute. In the vicinity of these rates the mechanisms

of decomposition involve a competition between the

monomolecular dissociation of intermediate radicals

to monomer and smaller radicals and the bimolecular

transfer reaction illustrated below [23, 24]. Since in

what follows only cases with one type of initiation

and one type of bimolecular termination are treated,

the rate constants for the particular type of reaction

are not given unique symbols.

Depolymerization Mechanisms

Initiation

End

Random

Propagation

Qi^R._;-R;

R.^Ri_i-M

2fei L.Q.-

Intermolecular R.+Qj—»Q.+Qj-*+R* v>ER»EyQj

Transfer

Termination

Disproportion - ation

Ri+Ri^Q.+Q, ^ [E Rd^

Combination Ri +R;—^Qi+j

^ [E R.-]^

Monomolecular Ri-^Q.-,+R,T /C4,E R.

Caged Combination

Ri+R;—^Qi+j kicE Ri

Caged Disproportionation

R<+R;^Qi+Q, fc4dE R»

Polymer molecules and radicals are, respectively,

designated as Qi, Qy, R,, Ry, etc., where the subscripts indicate the number of monomer units in the

molecule. Monomer is given the symbol, M.

No closed form relation can be derived for the

entire set of equations shown above. To evaluate

experiments, one must use a computer to numerically

evaluate specified cases [25, 26]. However, many

51
Table 3. Activation energies for the molecular vaporization of

n-alkanes [5].

A 11 Alkane n M

Rate,

/ 9 g/cm' s X 10^

8j), kcal/mol

Lxpt. Eq (9)

I L

10 268 6 0 18 18.9

n-tetracosane 24 338 6.8 21 21^9 75

7i-hexatricontane 36 506 7.5 29 28.6 145

n-tetra- 94 1318 12.0 46 55.6 345

nonacontane

polymer decomposition mechanisms can be evaluated

by comparison with theoretical relations derived

from simplified mechanisms which are special cases of the more complete system of reactions. At least

25 useful and easily deduced special cases can be

treated.

5. Decomposition of Polystyrene

In what follows, we will present only two cases; one which is applied to polystyrene data follows

immediately. Later on, a second case is presented

and applied to polytetrafluoroethylene data.

Assume new polymer molecules are only produced

by transfer.

dQ ^

hR(t)P4t)Q{t)

dt V{t)

where Q represents all polymer molecules, P„ the

number average degree of polymerization, and V

the volume of polymer. All these quantities depend

on time. It can be shown that

(13)

at y{t)m m

(14)

Here, p is the density and m the molecular weight of

monomer. The rate of conversion to volatiles,

assumed here to be only monomer, is given by

dC m t?M w , „ , — , k2R{t).

dt Wo dt wq

Combining eqs (14) and (15) gives

dQ _ kzpwg dC kzni

which integrates to

2

'

Since Q (t) w(t) /mPnit) , eq (5) becomes

1-C 1 ksp

Pnit) P„(0) hm C,

or

Pnit) 1-C

Pn{0)hpC

(15)

(16)

(17)

(18)

(19)

1+

kzTTl

Assuming a steady state and equating end initiation

with termination by disproportionation,

kipV(t) _ kiR^ty

rnPnit) V(t)

we find

R(0 =

hp 1/2

kimPnit)

Vit),

(20)

(21)

and hence

dC , d^=^-'

(kim

kiP

1 +

Pn{0)k,pC

k2m

1/2

(22)

In figure 2 we see how eq (19) fits data available for

polystyrene decomposition. Equation (22) also

reproduces much of the rate of volatilization behavior of polystyrene. See figure 3 where we show

the variation in the relative rate of volatilization

{dC/dt)

I jn ^ lunction of conversion, dashed {dLj/dt) max

lines. The full lines are based on an equation, in

which the square root is absent, devised by assuming

termination to be monomolecular. Caged recombination would be one type of termination which is

Figure 2. Pyrolysis of polystyrene at 302 °C, relative number

average degree of polymerization as a function of conversion.

Full lines theory eq (19) numbers on curves equal values of factor

(

f*n( ) ^3P \ equation. The initial number average molecular

weights, Afn(O), are: A 5,300 37.000 I 130,000 O 170,100 1 2 X 10»

52
kinetically monomolecular and may be involved in

the degradation of polystyrene.

Study of the initial rates in the 0-0.03 conversion

range of monodispersed or polydispersed volatilization reveals no correlation with the reciprocal of the

molecular weight, i.e., number of ends and hence

renders suspect that the mechanisms of initiation is at the molecular ends. The initial rates appear to be

essentially independent of molecular weight.

The molecular weight distribution of degrading

polystyrenes, both mono- and polydispersed, has

been examined carefully [27] and found to approach

the weight-to-number average DP ratio of 1.5

regardless of whether the initial ratio is one (fig. 4) or larger than two. This distribution is produced by

pure coupling mechanisms. Statistically coupling of

species in the "most probable" distribution produces narrower distributions [28, 29]. It would seem

obvious that any random splitting into fragments

followed by a statistical recoupling of the said

fragments would approach a 1.5 weight-to-number

average DP ratio, whatever the initial molecular

weight distribution of the system. The mechanism of

polystyrene decomposition continues to have many

puzzling aspects. The mechanisms of initiation and

termination are in doubt. The presence of a very

short kinetic chain length and intermolecular

transfer, however, is unquestionable.

6. Decomposition of

Polytetrafluoroethylene

This polymer decomposes with a very short

kinetic chain length, random initiation and termination by combination with perhaps a small amount of

disproportionation occurring [30, 31]. In the treatFiGURE 3. Pyrolysis of polystyrene, rates of volatizization as a

function of conversion.

Dashed lines theory himolecular termination based on eq (22); full lines theory for monomolecular termination. Experimental data and theory all normalized so that (<iC/dt)max = 1.0. Experimental points:

A

o

M„(0)

5,300 170,100

170,100

T°C

343

331 310

Figure 4. Variation of the ratio of the weight to number average

degrees of polymerization of polystyrene during thermalde gra- dation at 300 °C as a function of conversion.

O NBS 706 P„(0) = 1.31 X 10^ NBS 705 P„(0) = 1.64 X lO' A Fn(0) = 3.58 X 102

ment that follows, /c4 is the rate constant for total

mutual termination, i.e., ki = kic-\-k4d, where kic and

kid are the rate constants for termination by combination and disproportionation, respectively. The

steady state is then obtained by equating the rate of

initiation to that for termination.

hpV{t)

m V{t)

\kimj Vit).

(23)

(24)

Here, as before, R(t) is the number of moles of

radicals in a sample of weight, pV{t). It is therefore

not a concentration and the volume V{t) comes into

the second order rate expressions. The rate of conversion to volatile monomer is therefore given by

dt

dC

dt

mdM

Wodt = -k2R{t),

Wo

kim

kip

1/2 (1-C).

(25)

(26)

In figure 5 we show some characteristic rate of

volatilization curves for polytetrafluoroethylene

which have the properties predicted by eq [26].

We now neglect the molecular vaporization of

small polymer molecules. Then polymer molecules,

Q, increase only by the disproportionation process.

53
dc.

dt

^ 500°C

-on..

0

c

1.0

Figure 5. Rate of depolvmerization of polytetrafluoroethyleneas a function of conversion.

At 520°C the points are from three triplicate runs. At 500°C the points

indicate results for samples with different prior dosages of gamma

radiation.

Thus,

dt V{t)

Combining eqs [25] and [27],

(27)

1 1/2 dC ~ mki V{t) ki m k2\ m )

'

The Zip length, Z =

2^ (-^ I . Integ

eq (28) gives

ration of

(29)

which becomes

m;(0)(1-C) Wo

mPnit) mP„(0) (

k4d\ Wo

ki J m 2Z

or 1-C ]_ ^ /^\

Pnit) Pn{0) \kJ2Z

vaUd for Z«P„

(30)

(31)

0.2 0.4 0.6 0.8

C

Figure 6. Variation of relative degree of polymerization as a

function of conversion.

Points experiment at 480°C. Line theoretical from eq (32).

where P„ is the number average degree of polymerization.

Equation (31) may be rearranged to

Pn{t) _ (1-C)

Pn{0)

(32)

In figure 6 we show theoretical lines based on eq

(32). Since we know the zip length [30], Z=13 at

480 °C, and the degree of polymerization initially

and for degraded polymer [22] we find a value of

^^0.02 for kid/ ikid^k4c) . The experimental observation can easily be due to organic impurities, thus, it

is likely that termination in pure polytetrafluoroethylene would be entirely by combination.

7. Diffusion Effects Enhancing Rates of

Initiation or Termination

During studies of the radiolytic depolymerization

of polytetrafluoroethylene [30], it was observed that

rates were enhanced as the samples were made

thinner. The enhancement was also greater with

greater dose-rates as shown in figure 7.

IDO

0.0

LOGio (thickness, ;im)

Figure 7. Gamma ray induced depolymerization of polytetrafluoroethylene at °C.

Full lines observed rates of monomer production as a function of sample

thickness. Numbers to right of curves are dose rates, Mrad/hr. Dashed

line is theoretical for back reaction of monomer. [30].

In large samples it is obvious that sample dimensions can influence the decomposition of polymers by

way of retarded monomer escape, or of heat transfer

restrictions. Some possible causes of sample size or

surface influence have been put forward [32, 33, 34]

in the past. The radiolytic depolymerization result is evidence of another mechanism for dimensional

54
Figure 8. Schematic drawing of the variation in the concentration of large, R, and small, r, radicals in an infinitely long

slab of polymer having thickness I.

For dashed curves, I is smaller than for full curves.

effects on rates which is operative at considerably

smaller sample size, and which should be of general

importance. The basic feature is that in cases where

both large and small radicals or other small species

participate in the reaction, the latter can escape at

the surface, thus modifying the steady -state concentrations of all radicals in the shallower layers, and

in particular that of the decomposing polymer

radicals. In earlier studies of the trapping of small

radicals at 4 K, it was observed that the radiolysis

of solid methane produced methyl radicals and

hydrogen atoms of different concentrations at

long times [35].

A theoretical treatment [35] of this situation

required consideration of the competitive diffusion of

the two species. The specific process treated [35] is that of the radiation-induced decomposition of

polytetrafluoroethylene at 400 to 450 °C.

It has been shown that the back reaction of

monomer is inadequate to explain the observed

trend of decomposition rate as a function of sample

thickness [30]. Slow diffusion of the monomer

through the sample has been considered over twenty

years ago [29]. However, the variation of rate as a function of thickness was not calculated. The effect

of introducing the back reaction,

k-2 R,+M >R,+i,

into the usually depolymerization reaction as shown

by the dashed curve in figure 7. It should also be

pointed out that there is no effect in the pure thermal

decomposition

In figure 8 we schematically show the effect of a mechanism involving large polymer radicals, R, and

small, r, on the concentration of R in a slab of

infinite length. It is very clear that [R] and hence the

rate of depolymerization increases the thinner

the slab.

A survey of relative diffusion rates of polymeric

molecules and small molecules, approximately 1 to

10X10~^ and IXIO^^ cm^ s~', respectively, suggests

that it is permissible to ignore the migration of

polymeric radicals. Relative termination rates are involved in the theory as well. Many of the known

termination rate constants for small plus small and

large plus large radicals are in the range 1 to 6 X 10^

and 0.2 to 7 X 10' 1 mol~' s~\ respectively. It must be

cautioned that all these values are for dilute solution near room temperature, while the medium under

discussion here is molten polymer at several himdred

degrees Celsius. Wider ranges of variation are suggested by values of diffusivity of small molecules

in elastomers, of the order of 10~^ and 10~' cm^ s~\

and by one termination constant 4.0 X 10* 1 mol"~^ s~' determined in a decomposing polymer melt. As a hypothesis concerning large plus small termination

rates we rely on a formulation given originally in

Smoluchowski's diffusion theory for spherical molecules, but often accepted in more complicated cases

as well:

ki2 =

47r/Vo

1000

(Di+D2)ri2. (33)

Here ki2 represents the rate constant for the reaction

of small with large molecules, Nq is the Avogadro

number, Di and D2 are diffusivities, and is the

encounter distance in cm. Units of ku are 1 mol~^ s~^

It is likely that for actual small molecules relative

to large polymeric molecules, the encounter distances

are essentially identical, and diffusivities are greatly

different. Thus, setting ri2 = rii and £)2<$CjDi, it follows that:

ku^^kn, (34)

where ku is the rate constant for reaction of two

small molecules.

The above amounts to the estimate that large

plus small termination rate constants are of the same

order of magnitude as small plus small, and roughly

about half as great. Polymeric radicals thus terminate

with each other slowly, while small radicals terminate

with each other or with large radicals much more

rapidly. In layers near the surface, small radicals are depleted and only the slow termination of large

plus large radicals remains. There is thus a high

surface concentration of large radicals and a con- sequent high surface rate of decomposition. These

considerations are embodied in the scheme below,

where only random initiations are considered.

Interesting special cases containing thermolytic end

initiation can be developed, and at least one such

will be mentioned later in this article.

Reaction Rate

P-^R

P-*S

Ry+2—^R3 +monomer

R+R-^P

R+S->R-S

S +Byproducts

Vi, or kipm-\-^iIp/m

V2, or 4>2/p/

fcsR

k^W

fcsRS

Six reactions appear in the scheme, but the

propagation, A:2R, is involved only perfunctorily.

55
0.0

1.0

L ,M

- 100,0.5

1000,1.0

0.0 1.0 log

10 ^0

2.0 3.0

Figure 9. Theoretical rates of monomer production as a function

of sample thickness.

Pairs of numbers at right are values for the parameters L = fc5(fe4fe6)0*/2. M = vi/ui. The thickness parameter is the product of the actual

thickness and the constant factor, J -jj- \ {vik^yi^, where D is the

diffusion constant and v\ the rate of production of large radicals. Horizontal dashed lines give limiting rate value, y, at infinite thickness [31].

P, R, and S refer to polymer segments, polymer

radicals and small radicals, respectively. The

symbols Vi and V2 are radical generation rates; the

alternative expressions given involve a random

thermal initiation rate constant fci, density p, segment molecular weight m, radiation dose rate /,

and radiation efficiency factors $1 and <i>2. Because

the small radical diffusivity D, cm^ s~\ will be used,

concentration and rate-constant units are in moles,

cm^ and s. We consider the time-independent steady state in a strip sample of thickness X cm with positionvariable X cm below the free surface and an impermeable bottom. Evaporation of small radicals S

keeps a surface concentration, S = 0, at x = 0. The

steady state with diffusion is then

= 0 = j;2-/c6S2-/£5RS+D

dt

at

XX'- (35)

(36)

The solutions of the above equations are given

elsewhere pi]. Figure 9 gives the same representative results graphically. Comparison of these with

the experimental curves in figure 7 shows that the

theory fits the main characteristics of the experimental observation.

It is at first surprising that the mechanism has

a transition from thin-film to thick-film behavior at

thicknesses much less than those involved in retarded

monomer escape. The latter experimentally appears

to cause no change of regime until thicknesses of

tenths of a millimeter, under roughly comparable

conditions. Superficially, this difference in location

of the transition region can be rationalized by noting

that although the diffusivities, Z), may be similar for

monomer and small radicals, S, e.g., 10"'' cm^ s~\

the pertinent concentrations are very much smaller

for the radicals. This suggests, by the inversion of

Pick's first law, {dC/dx) = {l/D) (flux rate), that

at similar flux rates the radicals can go from near zero

to near their maximum permissible concentration over a shorter interval. The full truth is not so simple, and would require comparison of generation

and destruction rates as well. Although the treatment

[31] was developed in order to elucidate 7 -ray

induced depolymerizations, it would be expected to

apply to thermal depolymerizations initiated at the

ends of polymer molecules depending on the size and

diffusivity of the small radical produced.

Polystyrene [36] is thought to initiate with

production of small radicals; however, it also has a very short kinetic chain length [37] and hence the

effect would probably be very difficult to detect.

Essentially, the same mechanism can conceivably

lead to slower rates with thinner films. Consider the

following types of pseudo monomolecular termination processes:

CH3

-^CH2-C - X

H H

-C-CH2-C-CH2-CH

XXX

CH2

II CH2-C+H

X

H

C-CH2C = CH2+CH2 XX X

which would operate most efficiently at the surface of

decomposing polymer melt or film. Thus, diffusion

situations between large and small radicals or molecular species are likely to produce larger or smaller rates in thermolysis or radiolysis of small

thin samples of polymeric materials.

However, when the modes of initiation and

termination just discussed exist together and are equally affected by diffusion, the net result of

changes in thickness is merely a change of kinetic

chain length with no change in rate. To illustrate,

let the steps be designated as follows:

Initiation

Q—Ur+S T Propagation

R;+l—UR;+M T Termination

Rwhere Q, M, R, and S denote concentrations of

polymer molecules, monomer, large radicals, and

small radicals, and let the net effect of reduction in

thickness be to double ki and k^. Then the rate,

dlA/dt, is unchanged:

/C2/C1Q

smce,

ki

56
while the kinetic chain length, z, is reduced to one half:

smce,

A;4R Some of these effects of thickness may have been

encountered in other polymers. Films of polymethylmethacrylate have been reported to give

higher rates at such small thicknesses, 40 nm,

that the mechanism treated here could be expected

there as well. Their results in view of the high temperatures, short times, and procedures used, suggest

in our opinion many other possible explanations.

Qualitatively, the interaction of large and small

radicals at various depths could be important in

still other reactions. The cosmic dust particles discussed by Donn p9[], containing hydrocarbon and

exposed to ultraviolet, are small enough, 1 to 10 nm

thickness, so that efficient escape of H atoms would

occur from a large portion or all of a particle. The

radiation crosslinking of crystalline polyethylene is known to depend on thickness []40J, as follows

readily from our analysis if unescaped H atoms

scavenge polymer radicals which would otherwise

combine as crosslinks. The mechanistic consequences

of diffusion and thickness would also be of obvious

importance in oxidative degradation.

8. Diffusion and the Cage Effect

In a relatively few cases only, have values for the

activation energy for initiation and other elementary

steps been deduced in polymer depolymerizations.

The rates of elementary processes in polytetrafluoroethylene decomposition [30] are:

Initiation fci = 2.1 X 1028e-"«''5o»/«^ s"'

Propagation A:2 = 7 X lO^e-^^ ^oo/flr g-i

Termination ^4 = 3 X W^e-^^fi^^it^T \ ^-i g-i.

In the depolymerization of polymethylmethacrylate

[44] it appears that

A;4 = 3Xl0i%-20'°'"'/«^.

For poly-a-methylstyrene [43],

/ci = 4XlOi8e-'55.oo''/«?'.

In table 4 we present the ^4 -factors and activation

energies for the overall rate of decomposition

(volatilization) of several polymers. Very high

y4 -factors are evident for these systems. In two cases.

it is indicated that the ^-factor for the initiation

process is very high. It is likely that in the other

cases the initiation step has high ^4 -values and hence

-values for the net processes are high.

Since the early work of Cowley and Melville [44],

it has been known that the mutual termination of

radicals in pyrolyzing melts is diffusion controlled

and the rates of termination have experimentally

large (>;20) activation energies. In the case of

polytetrafluoroethylene the activation energy and

.4 -value for initiation also demonstrate that the

effect of diffusion (cage effect) must be considered

when discussing organic mechanisms in hot melts. It has been pointed out [45] that in the melt the

dissociation energy of a bond in the interior of a

large polymer molecule would be the difference of

the activation energies for initiation and termination:

D(R — R)m = £'i — £4. In the gas phase £4 = 0, hence,

D(R — R)g = £'i(gas). For polytetrafluoroethylene,

then the activation energy to produce active radicals

is D(R — R)m+£'4 = 118.6 using the data above. The

concomitant high yl-values indicate a highly disordered transition state which must encompass

'^5-10 CF2 units. Numerically, it requires ASJ

equivalent to the vaporization of >10 CF2 units.

These results and conclusions raise certain questions concerning the relationship between the dissociation energy in the melt, D(R — R)m and that in

the gas phase Z)(R — R)g. Of course, since these

large molecules cannot really vaporize, the latter

quantity is for an idealized process. Considering

the cycle:

D(R-R)

-> 2R:

AH^(R-R)

R-R„

d(r-r;

s

2AH^(R)

m

m * 2RIf, as might be reasonable, we assume: D(R — R)ni =

I>(R-R)g, then, AH„(R-R) =2A//„(R) . This

would be a very useful situation but we reject it because at least in the case of the ri-alkanes it is not

compatible with known data, see eq (8). From the

above cycle we can derive the relation for dissociation of the central interior bond in a large

rt-alkane molecule.

Table 4. Preexponential factors and activation energies for the thermal decomposition of polymers [45]

Polymer ^, s ' kcal/m

a-methyl styrene 10'^ 65

tetrafluoroethylene 10^' 81

a/3/3-trifluoro8tyrene 10'' 64

ethylene (linear) IQi* 72

methyl methacrylate lO'-lO'^ 30-52

styrene 10'^ 55

D(R-R)g

= D(R-R)^+2AH„(R-)-Ai/,(R-R).

From eq (8) it follows: D(R-R)g = D(R-R)„,+

0.26(3.21^2/3) -0.01937+2.92. Here, n is the number

of carbon atoms in R — R, the original molecule. For

a large linear n-alkane of 1000 carbon atoms we find

that at 27 °C, Z)(R-R)g = D(R-R)„,+80.59 kcal.

One concludes from the above result that the

"true or ideal" dissociation energy for large molecules

57
Table 5. Zip lengths for polymer depolymerization at

1 percent/min rates Table 7. Comparison of decomposition and ceiling temperature

Folyraer Zip lengths

methyl methacrylate ~103

a-metbyl styrene ^103

styrene ~5

tetrafluoroethylene ~3

should come from measurements in the melt phase or more precisely in a theta condition []46j. This

follows since it seems clear that Z)(R — R)m is the

energy characteristic of the chemical bond. One

suspects that there may be an appreciable activation

energy for recombination of large molecules in the

gas phase, since in the formation of the big molecules

a certain amount of reorganization would occur.

9. Thermodynamics of Polymer

Decomposition

The decomposition of organic polymers is, in

general, more a function of the dissociation rates of

bonds in the molecular structure, rather than the

thermodynamic properties of the material, since

photo and gamma rays will often initiate rapid

decomposition of some polymers at certain temperatures. Such effects will occur more strongly the

longer the zip length of the polymer decomposition

mechanism.

Zip length =

Depropagation to Monomer

Termination +Transfer

Table 6. Thermodynamic quantities for decomposition of

polyethylene and polytetrafluoroethylene.

AH AS AF

kcal/mole kcal/mole kcal/mole

* With the exception of graphite all the substances are taken to be in the "ideal" gas state at 298.2 K. For the high polymers this "ideal" gas state is evidently for the polymer in its most extended conformation [5].

Folymer

X c Li d 11 uui my icuc OoU

j.nn

J- LwLiy Italic

^"t kj i« Vrf*Tl y L \^n^A zou Methyl methacrylate 330 220

IVXC tUctCK Y lull ILL lie 1 77 III Isobutylene 340 50

a-methyl styrene 290 7 Formaldehyde 120 35

~CH2CH2~ACH2CH2 22.35 34.07 12.19

b—^l/SCeHe —H2 16.46 34.28 6.23

A2C(Graphite)

-2H2

9.85 46.76 +4.10

d—>re-Alkene8 [(1/18)CH3

(CH2)33CH =

CH2]

1.09 2.03 0,49 dc

dt

e—>C (Graphite)

-CH4

-8.04 27.49 +16.23

~CF2CF2—^CF2CF2 46 45 33

-^2C(Graphite)

-2F2

194 71 172

h—>C(Graphite)

-CF4

-27 35 +38

That is, the average number of monomer units

produced by a decomposing radical before the

radical is deactivated either by termination or

transfer. See table 5 where some estimated zip

lengths are listed.

Consideration of table 6, which gives values of the

various thermodynamic fimctions for the decomposition of polyethylene and polytetrafluoroethylene

by several different processes into quite different

products, illustrates how metastable organic polymers

are. It is interesting to note that certain of the

conceivable mechanisms are very exothermic, but

that the polymers normally decompose by reactions

which are not thermodynamically favored at ambient

temperatures. It would seem that the possibilities

for the incorporation of catalyst for controlling the

decomposition of polymers to give non-combustible

volatiles could be extensive.

Polymers are then in general metastable. We see

this fact also in table 7. Most polymers listed show

greater stability than one would expect from their

ceiling temperature Tc. This is the temperature

where the rate of propagation equals the rate of

depropagation or the equilibrium constant for the

reaction, Rj +M<:^Ri+i, equals one. Thermodynami3.0

2 0

-

Figure 10. Pyrolysis of polycarbonate, rates of volatilization as a function of conversion, activation energy based on maximum rates, 53 kcal/mole.

58
cally then, polymers can decompose above Tc,

which can be estimated from the following relation:

where the quantities AHp and ASp refer to the

standard enthalpy and entropy of polymerization.

10. Aromatic Polymers

Aromatic polymers, such as polycarbonate, polymellitimide and polydimethylphenyleneoxide, tend

to decompose in varying degrees to a char, i.e., a carbonaceous residue. Most of these polymers are

insoluble and hence study of the decomposition

processes, particularly those producing crosslinks or

scission is extremely difficult if not impossible. One

polymer that is soluble is the polycarbonate which

dissolves in dioxane.

On pyrolysis, polycarbonate volatilizes

rapidly at first but quickly drops to a very slow rate

in the vicinity of 0.50 conversion (see fig. 10). At

this stage ~0.50 conversion, the residue is a car- bonaceous char.

At lower temperatures —^300°, it is found that the

percent gel in the polymer increases with conversion

to values of ~90 percent at —-10-20 percent conversion. The swelling ratio of the gel decreases with

the conversion, but less strongly the higher the

temperature. These observations indicate a competition between net scission and crosslinking

processes as a fimction of conversion. The relative

ratio of the two depends on temperature, the thick- ness of the sample and also on the amoimt of oxygen

present. In figure 11 we show the effect of thickness

and air. The thicker the polymer the slower the rate

at the longer times. These curves also suggest that

the thicker samples as might be expected favor

crosslinking. In figure 12 we compare the rates of

volatilization at 410 °C of various samples preheated

0 6

C

0.3

/

,' 100 mil,/'

1 air /

7 -r^—

/ VQC

30 mil,^-^^

^^-^00 mil, vac

iZmil, ' /

; air ' /

1 1 / /

\ ' / / ' / /

I ' / / /

1 1/ / /

1 m / \

' / 1/'

/ r

1

' / / ' '/ // ' '/// '

r // ' /// ' // /

' V I / / ' At' //'I

// 'J 1 // 1 '

'/f /

1 30 60 120

t, min

Figure 11. Pyrolysis of polycarbonate, effect of thickness on

conversion as a function of time.

2.0

0 .5 C Figure 12. Pyrolysis of polycarbonate at 410 °C.

Effect of preheating at 315 °C on rates as a function of conversion. Numbers in curves give percent weight loss during the preheat

period.

at 315 °C to the percent weight losses indicated by

the numbers on the curve. The conclusion is that the

greater the crosslinking the less the rate of volatilization. Thus, the aromatic chain polymers, in that they

tend to crosslink and eventually char, are mechanistically different from the saturated line polymers

where scission processes predominate. As yet the

understanding of the degradative process in aromatic

polymers is so limited that theoretical developments

are not feasible.

11. References

[1] L. A. Wall, Pyrolysis, Chap. V, Analytical Chemistry of

Polymers, G. M. Kline, ed. (Interscience, New York,

1962).

[2] L. A. Wall and J. H. Flynn, Rubber Chemistry and

Technology 35, 1157 (1962).

[3] M. Tryon and L. A. Wall, Oxidation of Polymers, Chap.

19, Vol. 2, Autoxidation and Antioxidants, W. O.

Lundberg, ed. (Interscience, New York, 1962).

[4] L. A. Wall, J. H. Flynn, and S. Straus, Polymer Eng.

and Sci. 10, 19 (1970).

[5] L. A. Wall, J. H. Flynn, and S. Straus, J. Phys. Chem.

74, 3237 (1970).

[6] L. A. Wall, G. Mattamal, and M. H. Aldridge, unpublished.

[7] I. Langmuir, Phys. 8, 149 (1916).

[8] D. R. Stull, Ind. Eng. Chem. 39, 517 (1947).

[9] F. D. Rossini, et al., API Res. Proj. 41 (1952).

[10] W. M. Mazee, Rec. Trav. Chim. 67, 197 (1949).

[11] E. Morawetz and S. Sunner, Acta. Chemica Scand. 17- 473 (1963).

[12] C. B. WilUngham, W. J. Taylor, J. M. Piznocco, and

F. D. Rossini, J. Res. Nat. Bur. Stand. 35, 219 (1945).

[13] R. W. Sbiessler and F. C. Whitmore, Ind. Eng. Chem.

47, 1660 (1955).

[14] A. F. Forziatti, W. R. Norris, and F. D. Rossini, J. Res. Nat. Bur. Stand. 43, 555 (1949).

[15] H. T. Coach, W. Kozicki, and B. H. Sage, J. Chem.

Eng. Data 8, 347 (1963).

59
[16] R. S. Bradley and A. D. Shellard. Proc. Roy. Soc. (London) A198. 239 (1949).

[17] H. S. Myers and M. R. Fenski, Ind. Eng. Chem. 47, 1652

(1955).

[18] S. H. Fishtine, Ind. Eng. Chem. 55, 51 (1963).

[19] I. Langmuir, Third Colloid Symposium Monograph, pp

53-54, The Chemical Catalog Co., Inc. New York,

1925.

[20] M. L. Huggins, J. Phvs. Chem. 43, 1083 (1939).

[21] L. A. Wall, S. L. Madorskv, D. W. BroM-n. S. Straus. and R. Simha, J. Am. Chem. Soc. 76, 3430 (1954).

[22] T. C. Siegle, L. T. Muus, Tung-Po Lin, and H. A.

Larsen. J. Polymer Sci. A2, 391 (1964).

[23] R. Simha, L. A. Wall, and P. J. Blatz. J. Potymer Sci.

5, 615 (1950).

[24] R. Simha and L. A. Wall, J. Phys. Chem. 56, 707 (1952).

[25] R. Simha. L. A. Wall, and J. Bram, J. Chem. Phvs. 29,

894 (1958).

[26] L. A. Wall, S. Straus, J. H. Flynn, D. Mclntvre, and

R. Simha, J. Phys. Chem. 70. 56 (1966).

[27] L. A. Wall. S. Straus, and L. J. Fetters. Polymer Preprints 10, 1472 (1969).

[28] P. J. Flory, Principles of Polymer Chemistry, Cornell

Univ. Press, p. 333, 1953.

[29] G. M. Burnett, Mechanisms of Polymer Reactions, Interscience, pp. 76-105, 1954.

[30] R. E. Florin. M. S. Parker, and L. A. Wall. J. Res. Nat. Bur. Stand. 70A, 115 (1966).

[31] R. E. Florin and L. A. Wall. Macromolecules 3, 560

(1970).

W. Tsang (National Bureau of Standards)

In the decomposition of medium sized (C5 or above) olefinic and acetylenic compounds in a single

pulse shock tube at about 1100 K, concurrent with

the bond brealiing processes, there appears to be a molecular process of the type observed by Bailey

[W. J. Bailey, SPE Trans'. 5, 1 (1965)]. For the

acetylenic compound 4-methylpentyne-l the rate

expression for the molecular process is [W. Tsang,

Int. J. Chem. Kinetics 2, 23 (1970)]

k ( 4-methylpentyne-l^allene propylene

)

= exp ( - 58,500/i?r) s'K

Furthermore, it appears that the molecular reaction

for olefin decomposition follows a similar expression.

On the other hand, for the cleavage of a normal

alkane, the shock tube data suggest a rate expression

of the form A-[n-alkane-^ (Radical) 1+ (Radical) 2]^^

101^-3 exp (-82,000 i?r) s-i. If such data are applicable to the pyrolysis of polyethylene or any other

polyolefin at 100 K, it would appear that unless

unsaturated impurities are less than 1 part in 100,000,

[32] L. A. Wall, L. J. Fetters, and S. Straus, Polymer Letters

5, 721 (1967).

[33] H. H. G. Jellinek and H. Kachi, Polymer Eng. Sci. 5,

1 (1965).

[34] S. Straus and L. A. Wall. J. Res. Nat. Bur. Stand. 65A,

223 (1961).

[35] D. Brown. R. E. Florin, and L. A. Wall, J. Phys.

Chem. 66, 2602 (1962).

[36] G. G. Cameron. Makromol. Chem. 100, 260 (1967).

[37] L. A. Wall, S. Straus. J. H. Flynn, and D. Mclntyre,

J. Phys. Chem. 70, 53 (1966).

[38] A. Barlow. R. S. Lehrle. and J. C. Robb, Makromol.

Chem. 54, 230 (1962).

[39] B. Donn, "Formation and Trapping of Free Radicals", A. M. Bass and H. P. Broida, eds.. Academic Press,

N.Y., 1960, p. 347.

[40] R. Salovey and M. Y. Helbnan, ACS Polymer Preprints

9, 1124 (Sept. 1968).

[41] R. Simha and L. A. Wall, J. Polymer Sci. 6, 39 (1951).

[42] L. A. Wall, High Temperature Resistance and Thermal

Degradation of Polymers, S.C.I. Monograph No. 13,

p. 146, Society of Chemical Industry-. London. 1961.

[43] D. W. Brown and L. A. Wall. J. Ph'ys. Chem. 62. 848

(1958).

[44] P. R. E. J. Cowley and H. W. Melville, Proc. Roy. Soc. A210, 461 (1951); A211, 320 (1952).

[45] L. A. Wall, S.P.E.J. 16, Nos. 8 and 9 (1960).

[46] P. J. Flory. "Principles of Polymer Chemistry", Cornell

Univ. Press, Ithaca, N.Y., 1953).

[47] L. A. Wall and S. Straus, Unpublished.

the initial degradation process must also involve a molecular reaction.

LeoA. WaU:

Reactions of unsaturated substances of the tvpe

that Dr. Tsang calls to our attention may be occurring in the pyrolysis of polyethylene: however,

experimental evidence indicates they are of negligible

importance. During the thermal degradation of

polvethvlene vinyl end groups are produced (Oakes

and Richards, J. 'Chem. Soc. 2929 (1949)) by a ran- dom cleavage reaction which has the rate constant

X0i7.6g--2,ooo/5r^ (S. Straus and L. A. Wall. J. Res.

NBS 65A, 221 (1961) deduced from the maximum

rate that occurs in the vicinity of 26 percent weight

loss. From the thermal stress relaxation of crosslinked

polyethylene in the vicinity of < 1 percent weight

loss, a similar value, IQ-^e^'^'"TM^*^, for the cleavage

process is deduced.

One of the still to be determined questions about

the mechanism of the thermal decoinposition reaction

of polyethylene between 300—450 °C is whether this

cleavage process is or is not a chain reaction. If it is, then it would appear that polyethylene decomposition is initiated by the dissociation of carboncarbon bonds.

60
NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrglysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Patterns and Problems in the Pyrolysis Behaviour of Synthetic

Addition Polymers

G. G. Cameron

University of Aberdeen, Aberdeen, Scotland

The types of pyrolytic decomposition reactions undergone by synthetic addition polymers

are outlined. Particular attention is given to two classes of polymer—those which undergo unzipping to monomer and those which fragment randomly along the backbone to yield low polymer.

Free radical transfer reactions, usually involving tertiary hydrogen atoms, can play an important

part in these degradation reactions. The thermal stabiUty of addition polymers is not always

predictable from known reactions of small molecules. This is often due to the presence of labile irregularities which are chemically incorporated in the macroraolecules and which can initiate thermal decomposition. Also, since polymers are susceptible to chain reactions and neighbouring

group effects, these abnormalities can affect the stability of a large proportion of the sample.

The effects on pyrolysis of abnormalities such as unsaturation, chain branches, head-to-head links and oxidation products are discussed.

Key words: Free radicals; polymer abnormalities; polymer degradation; polymer pyrolysis; transfer reactions.

1. Patterns of Decomposition

The vacuum pyrolysis of addition polymers has

been the subject of many detailed studies during the

past twenty years. Consequently it is now possible

to make certain generalizations regarding the link

between structure and decomposition mechanism.

Most of the polymers considered here can be represented by the generalized structure

r x"i

I -v^-CH^-C---v^

which shows the repeat unit of the initial monomer

X

Y

The pendant groups, X and Y, may be hydrogen or some other substituent such as a halogen, nitrile

group, etc. Three distinct types of degradation

reaction can occur with such polymers: (a) depolymerization to yield monomer; (b) random breaking

of the polymer backbone to yield small polymer

fragments; (c) reactions involving substituents

X and Y.

Polymers in categories (a) and (b) can usually be

converted completely to volatile materials, while

those in (c) generally form a char or residue. While

some polymers have degradation patterns which fit almost exclusively into one of these categories many

show intermediate behaviour. For example,

poly (methyl methacrylate) gives —^98 percent

monomer on pyrolysis and therefore falls into

category (a), but polystyrene gives a mixture of

monomer (^^45 percent by wt.) and low polymer

and therefore shows characteristics of both (a) and

(b). Poly (methyl acrylate) yields low polymer

fragments as well as CO2 and methanol so that it shows characteristics of (b) and (c). A few polymers

also crosslink during pyrolysis. Table 1 shows a list of some addition polymers in categories (a) and (b)

\vdth the monomer yield on pyrolysis.

This table highlights behavioural patterns. When

neither X nor Y are hydrogen atoms, yields of

monomer are very high. When X is replaced by a hydrogen atom there is a dramatic fall in monomer

yield (compare poly (methyl methacrylate) with

poly (methyl acrylate) and polymethacrylonitrile

with polyacrylonitrUe) . Similarly among the polystyrene-types, replacement of the a-hydrogen atom, even by deuterium, increases the monomer yield.

Clearly the a-hydrogen atoms play a key role in the

decomposition pattern. This fact can be rationalized

to a large extent if we assume that decomposition

involves a free radical chain reaction. This assumption has been shown to be vaUd in a number of

systems and it is reasonable to extend it to others

showing similar behaviour.

61
Table 1. Yields of monomer on pyrolysis of some addition polymers and transfer constants of polymer radicals to toluene at 60°C

X

n 1 Polymer X and Y in —CH2—C

1 Y

Monomer

yield

Wt. %

to toluene

V_-JJ.3 ^yo i .ZD X iO ^ XPolUiyT7 f^^IIICLU-Yl TTi n 1 aaLLiJLJcL iTr^n a 1"f UC i^

JT Ol y iiic Liidi-r y luxii lljjc CN on Poly (methyl acrylate) H CbaCHs trace 0.56

Pnlvji pi*TrlnTTit"nlf H CN n 7Q

Polyethylene H H 0

Polypropylene H CH3 2

Polystyrene H C6H5 45 2.11X10-3

Poly(a-methyl styrene) CHs CeHs 95

Poly (a -deuterostyrene) D CeHj 70

Poly(;3-deuterostyrene) H Cells 42

The a-hydrogen atoms can take part in transfer reactions of the type:

H Hi

I

I

H

I CH2-C-CH2-C- -CHg-C- -CHg-C • -» -v-CHg-C-CHp-C- -CHg

n

H

I •CI XL

I X

i-i CHg-C

X-J

H- -il

I XCHg—CHg

X

(1)

n

which shift the radical centre from a chain end

(the propagating radical) to an intrachain position.

Scission at the new radical centre leads to fragmentation of the original polymer chain. Such

transfer reactions may be intramolecular (Xo a hydrogen atom on the same chain ) or mfermolecular

(to a hydrogen atom on another chain ) . When

a-hydrogen atoms are replaced by methyl or other

groups transfer reactions are less favoured and an

unzipping or reverse polymerization process takes

over, leading to monomer as a major degradation

product. This explains the marked difference in

behaviour of poly (methyl methacrylate) and

poly^methyl acrylate j for example.

Ready availability of a-hydrogen atoms is not,

however, the only criterion in determining the course

of decomposition, otherwise we should expect

polystyrene to produce little, if any, monomer.

Another factor which must be taken into account is the reactivity of the propagating radical as a transferring species. On this basis the relative transfer

constants for a series of polymer radicals measured

imder identical conditions should follow a pattern

similar to that observed for the yield of monomer on

thermal degradation. Such a set of transfer constants

of several polymer radicals towards toluene [1]^ at

^ Figures in brackets indicate tlie literature references at the end of this paper.

60° is also shown in table 1. Correlation is fairly good.

Polymer radicals with low values of i^trans give high

yields of monomer and vice versa. The polystyryl

radical does not transfer very frequently, despite the

abundance of a-hydrogen atoms, because it is relatively imreactive. This is presumably a consequence

of the stabilizing influence of the benzene ring. These

generalizations, though rather oversimplified, provide a convenient framework for considering the

overall pyrolysis behaviour of polymers in categories

(a) and (b).

It should be made clear that while free radical

intermediates are involved in most unzipping and

random fragmentation reactions it is by no means

certain that this generalization can be applied to

all thermal degradations, particularly those involving

side chain decomposition.

2. Factors Affecting Thermal Stability

of Polymers

The thermal decomposition reactions of polymers

are not always predictable from known reactions of

small molecules. This is often due to the macromolecular nature of the reacting species in which

the decomposition of one unit in the chain can

activate the decomposition of an adjacent unit. For

example, in the deacetylation of poly (vinyl acetate j

62
once the first molecule of acetic acid has been

eliminated the double bond formed in the backbone

weakens the next ester group and facilitates further

reaction resulting in long sequences of conjugated

double bonds [2]].

The unzipping process, which produces monomer,

may occur more rapidly and at much lower temperatures than a model reaction. Thus poly (ethyl

methacrylate) liberates monomer in high yield at

about 250 °C while primary ethyl esters are stable

at this temperature and eventually decompose into

ethylene and carboxylic acid at about 450 °C.

In the case of the polyolefins thermal decomposition occurs as in model short-chain paraffins.

Polyethylene yields shorter chain monoolefins.

However, significant decomposition occurs some

100 °C below that observed in the model. A similar

effect is observed with poly (vinyl chloride) which

decomposes by loss of HCI leaving an olefinic

residue. This reaction is in accord with predictions

from simple chloroparaffins but again occurs at a much lower temperature than predicted by the

model. Effects of this type are by no means rare and

can often be attributed to the presence of labile

irregularities in the polymer chains. These abnormal

structures can be formed during or after polymerization and are partly responsible for the varying

thermal stabilities of different samples of the same

polymer.

From a technical point of view it is obviously

desirable to improve, as far as possible, the thermal

stability of polymeric materials. It is therefore

pertinent to consider how thermolabile irregularities are incorporated into the polymer structure and how

they can affect degradation reactions. It is important

at this point to distinguish clearly between irregular

structures chemically bonded to the polymer chains

and adventitious contaminants physically mixed

with the polymer. In this article we are concerned

only with the former. However, there are many

examples in the literature of the destabilizing effects

of impurities, particularly residues of polymerization

initiators. The doubtful purity of many polymer

samples studied in the past is probably responsible

for much of the conflicting data which has appeared

in the literature.

Some examples of chemically bonded abnormalities, incorporated both during polymerization and

as a consequence of post-polymerization treatment,

and their behaviour on pyrolysis are discussed in

more detail below. The list is illustrative rather than

exhaustive and raises a number of contentious

points.

2.1. Unsaturation

One of the earliest studies on polymer degradation

concerned poly (methyl methacrylate). This work

established that there are at least two initiation

reactions which can lead to unzipping of this polymer

to monomer [3]. In samples prepared by free radical

routes it was shown that approximately half of the

polymer chains could be decomposed at temperatures

about 220°C while the remainder required to be

heated to 270 °C before significant decomposition

occurred. During free radical polymerization of

methyl methacrylate termination is mainly by

disproportionation

CH^

-V-CH2-C-CH2—C'

CH^

C-CH.

COoCH^

-^-CHo-C-CH

CO^CH^

COgCHj

CH,

CH^

-C-CH.

I CO2CHJ COgCHj

CH

= C

3 1 3

(A]

COgCH^

+ H-C-CHo—C-CHo-^

CO^CH^ ^3 COoCH^ ^3

(2)

leading to about 50 percent of polymer chains with

unsaturated end groups, the remainder being wholly

saturated. It was proposed that the unsaturated

groups constituted weak spots in the chains and were

responsible for the low temperature reaction by

facilitating scission at one of the bonds indicated to

form an allylic radical. The kinetics of the low

temperature reaction were found to be consistent

with a chain-end initiated reaction.

More recently it has been shown that the unsaturated end formed by disproportionation has the

structure

CH^

I ,3

•-^-CH^-Ci-CH^—

II 2

(B)

co^cn^ cOgCH^

which leads directly to the depropagating radical

CH^

I 3

CHo-C • ^ I CO„CH,

3 on decomposition. This destabilizing abnormality

is a consequence of the polymerization mechanism

and is difficult to eliminate in free radical polymerizations. Poly (methyl methacrylate) polymerized

anionically contains no unsaturation and is noticeably more thermostable than its free radical counterpart undergoing depolymerization after high temperature initiation.

The high temperature reaction appears to occur by

random homolytic scission.

9'b

CH^

••CHg-c-CH^c-cn^-

COpCH^ COgCH^

-^-CHg-C-CH^

COgCH^

CH,

• C-CHo

I ^

(3)

It is normally assumed that both radicals are capable

of imzipping to monomer and certainly the available

evidence suggests that if the initial polymer chains

are short enough they are completely converted to

63
monomer once initiated. The two radicals produced

in this manner must have very different reactivities

and in view of the earher comments it is possible

that the radical (C) vmdergoes some rapid reaction

Photolysis of poly (methyl

produces random homolytic

examination of the degraded

glasses shows no evidence of radical (C) [4]. The

spectrum of the propagating radical can be identified

along with a weaker four-line spectrum which has

been attributed to an allylic radical formed from

(C) as follows:

besides unzippmg.

methacrvlate) also

scission, but E.S.R

I 3 = + HCO2CHJ

(C) (4)

It is possible that during thermal degradation

radical (C) reacts in a similar fashion, e.g..

CH,

^2 ""5

(5)

The structure (D) would constitute a thermolabile

group which could lead to an element of chain-end

initiation superimposed on the random kinetics. A

similar conclusion has been reached on the basis that

radical (C) would transfer rather than unzip Q5j.

When the polymer chains are long cf. the zip

length a bimolecidar termination step occurs. If

this step is by disproportionation, as in polymerization, imsaturated end groups of type (B) are formed.

The partially degraded polymer is therefore less

stable than the original and the kinetics of degradation are complicated by an element of chain-end

initiation superimposed on the original random

process. A reexamination of some earlier rate data

for poly (methyl methacrylate) bears out this comment and shows that the kinetics and mechanism of

the randomly initiated depolymerization are more

complicated than is commonly supposed \_()~\. In the

structurally similar polymer, poly (methyl atropate),

random initiation of unzipping occurs in the early

stages of pyrolysis but beyond about 45 percent

conversion chain-end initiation predominates

The accmnulation of labile chain ends as discussed

above may accoimt for this.

The above discussion shows how a thermolabile

unsaturated end group may be formed in a previously saturated macromolecule as a consequence of

thermal degradation. Similar structures could be

formed by any process which leads to chain scission,

e.g., mechanical degradation or exposure to radiation.

The latter processes may be carried out below the

ceiHng temperature of the polymer without evolving

monomer and provide examples of abnormalities

introduced by post-polymerization treatment.

Unsaturated chain ends are likely to be a source of

weakness in any polymer system but do not necessarily initiate depolymerization via a radical chain

mechanism. In the case of polyisobutylene random

degradation occurs and appreciable quantities of

monomer appear in the volatile products. Some of

this monomer is almost certainly produced via free

radical intermediates, but it has been shown by

deuterimn labelling studies that imsaturated end

groups (formed in proton expulsion termination of

polymerization) can evolve monomer via a cychc

mechanism |^8 j.

^Er.

^E

CH^

II ^ 1 5 II ^

C-CHg- • C—CHj — -^Hg-C-CHg-ip +

CH^ CH^ CH^ CH, CH^

(6)

CH^

Any unsaturation produced by a scission reaction

would behave in a similar manner. The effects of

unsaturation on the decomposition of other polymeric hydrocarbons has been studied by Bailey [8J.

Unsaturation within the maia chain can also occur in certain polymers. McrSeill and coworkers

found that polvstyrene prepared by a free radical

route contains about 1.5 double bonds per 1000

monomer units No terminal unsaturation was

present when benzoyl peroxide was used as initiator.

This is consistent with the fact that, unlike the case of methyl methacrylate, termination ia styrene

polymerization is predominantly by combination.

H H H H

Ph Ph Ph Ph

(The effect of the resulting head-to-head structure is McXeill has suggested [9] that it could arise from

considered later). The source of main chain un- transfer to polymer:

saturation in free radical polystyrene is obscure but

64
H

I H

2 I Ph

I Ph

'i-^CH^-C-CH^-C-CHg"^ + R

The radical (E) could reinitiate polymerization

giving rise to a branch, or the adjacent methylene

group could eliminate a hydrogen atom to form a double bond as in (F). The bonds indicated in

structure (F) would be weakened compared with

normal backbone bonds. If this mechanism of double

bond formation is correct, then unsaturation in the

backbone should increase with conversion during

polymerization. Consequently polystyrene samples

prepared to high conversion should be less thermally

stable than low conversion polymers. This has been

shown to be the case in pyrolysis studies of polystyrenes prepared under different conditions [^lOj.

In this work the degradation temperature was kept

below the ceiling temperature so that bond scission

could be studied without the complication of significant formation of volatiles. Figure 1 shows plots of

the fraction of bonds broken (a) as a function of

time of heating at 287° for two samples of polypi-o

4 0 6 0

Time, hr 00

Figure 1. Degree of degradation, a, versus time of heating at 287°. A = TPS/5 (12.5 percent conv.) O = TPS/6 (60 percent conv.) H

\

•CH.

H

I RH + 'T-CH2-C-CH2-C-CK:2-

Ph Ph

(E)

CH^

C

I Ph Ph

(8)

+ Hstyrene prepared imder identical conditions except

that the conversions of monomer to polymer for

TPS 5 and TPS 6 were 12.5 and 60 percent respectively. The high conversion polymer, TPS/6, is markedly more thermolabile. This evidence is not

unequivocal proof that TPS 6 contains more imsaturation than TPS 5. An increase in concentration

of any abnormalitv could bring about the same

effect. However, branching is the only other abnormality which would be expected to be conversion

dependent but from a knowledge of the likely branch

density the authors concluded that increased unsaturation was a more acceptable explanation. The

activation energy for random bond scission in free

radical polystyrene lies in the range 49-54 kcal

mole^^ which is some 15-20 kcal lower than the

estimated bond dissociation energy of normal backbone bonds. Unsaturation, as in structure (F), would

lower the bond energy of the j3 C—C bonds by

about 17 kcal. Thus, a contribution to degradation

by decomposition at double bonds is consistent

with the measured activation energy. Polystyrenes

prepared by anionic routes are more thermostable

than free radical polymers and have activation

energies for random scission in the region of 70 kcal

[11]. Propagation in anionic polymerization seems

to be comparatively free from abnormal steps and

termination, normally by proton addition, yields a more stable end structure. In the case of cationic

polymerization of styrene proton expulsion leads

to unsaturated chain ends but no main chain unsaturation [12]. The effects of these unsaturated

ends has not been investigated in detail but would be

expected to destabilize the polymer.

2.2. Branching

Chain branches on polymers normally arise by

transfer to polymer during free radical polymerization. Transfer may be intermolecular yielding long

branches. Short-chain branches may be formed by

intramolecular transfer, a process which is apparently

relatively frequent during the high pressure polymerization of ethylene.

CE

"CH

-CH.

"CH

CHg—CHg

(9)

65
High pressure polyethylene is characterized by the

frequency of these short branches. The precise

effect of branches on the thermal stability of high

pressure polyethylene is not absolutely clear since

other abnormalities, mainly oxygenated groups, are generally present also. It has been suggested that

branching introduces more intramolecular transfer

at the expense of intermolecular transfer due to the

more compact form of highly branched molecules,

particularly when the branches are long [13, 14].

Branched structures also occur during free radical

polymerization of poly (methyl acrylate ) , probably by

transfer to the relatively labile a-hydrogen atoms.

The effect of branches on the pyrolysis behaviour of

this polymer is rather different from that operating

in polyethylene. Many features of the degradation

of poly (methyl acrylate j are unusual; among these

is the formation of appreciable quantities of insoluble gel in the polymeric residue [15, 16]. Gel

formation is a consequence of some crosslinking

reaction ^hich may be a combination of radicals.

COoCH^ COoCH^

(G)

+

'CH2-C-CH2—c-^

COpGHj CCpCH-j

CO„CH^ COoCH^

I 2 3 I 2 3

C Ho-C-GHo—

H

H

I (10)

G02CH^ C02CH^

The appearance of large amounts of low polymer in

the volatiles and the rapid drop in molecular weight

of the soluble portion of the polymeric residue

indicate the frequent occurrence of transfer reactions

and the existence of radicals of type (G) is fairly

certain. The crosslinking reaction competes with

chain scission and, if it occurs frequently enough,

can lead to the formation of an infinite network.

Examination of a range of poly (methyl acrylate)

samples indicated that gel formation is dependent on

two main factors in the preparative history of the

sample—the temperature of polymerization and the

degree of conversion of monomer to polymer. The

effect of polymerization temperature on gel formation

is showTi in figure 2 in which the most notable

feature is the minimum at ~50 °C. At 238 °C

virtually no volatUes are evolved. Measurements of

the second virial coefficients of poly (methyl acrylate)

samples prepared at different temperatures ffig. 3)

show that the degree of branching also goes through a minimum in the region of 50 °C. The reason for this

behaviour is obscure at present. Although corre- spondence between figures 2 and 3 is not exact the

appearance of minima in each suggests that highly

branched samples of this polymer have the strongest

tendency to form gel during pyrolysis. As the temperature of poh"merization is lowered the polymer

has a higher molecular weight which would tend to

push up the gel-forming tendency of the sample

prepared at 40 °C. The correlation between gel

formation and branch density is reinforced by the

observation that the higher the conversion on

polymerization, the higher the gel content on

degradation. Branching in this polymer is an imavoidable consequence of free radical polymerization.

It would be interesting to confirm the above view by

degrading a truly linear sample of poly (methyl

acrylate), possibly prepared by anionic poljTnerization.

In general the effects of chain branches on thermal

stability are not as clearcut as those of unsaturation.

There is some doubt, for example, as to the possible

destabilizing role of branch points on P^ C dehydrochlorination [17].

50 60 70

Temperature of Polymerisation CO

Figure 2. Percent insoluble gel after 1 hr at 238° as a function

of polymerization temperature.

66
o

»

Eo

in 3

O

Q.0 2

V)«

uo

h. 1 1

40 50 60 70

Temperature of Polymerisation CO

Figure 3. Reciprocal osmotic slopes of poly{methyl acrylate)

samples as a function of polymerization temperature.

2.3. Head-to-head Links

It is generally accepted that addition polymerization occurs predominantly in the head-to-tail

fashion.

X

1 '-<;Ho-c--cHo-c- + CH, 2 I Y

X

I c

I Y

X

-H-CH„-CY.

X

I CHo-C-

^ I

(11)

This mode of addition is strongly favoured on steric

and energetic grounds. Nevertheless the possibility

of an occasional head-to-head addition can never be

completely discounted. The resulting structure

X X

I I —CHg-C-C-CHg"

Y Y

would be expected to be more thermolabile than

normal backbone bonds and could provide a weak

point in the chain.

The role of such linkages in the pyrolytic decomposition of polystyrene has received considerable

attention against the background of a lively controversy on the question of "weak links" in this

polymer [18]. Since termination of the free radical

polymerization of styrene is by combination (see

eq 7) each chain in the sample will contain at least one such abnormality [193. (The additional concentration of such groups due to head-to-head addition

during propagation is likely to be very small.) The

proportion of head-to-head units can be substantially

reduced if polymerization is conducted in the

presence of an active transfer agent or if the polymerization temperature is increased to about 100 °C

when disproportionation occurs. Despite their probable low concentration head-to-head units un- doubtedly weaken the polymer chain. This has been

proved by two independent studies. In the first of

these, polystyrenes containing controlled amounts

of head-to-head structures were prepared by coupling

"living" polystyrene anions with iodine [20]].

H H

I l_ +

Ph Ph

H

I

Ph

H

I -CI

H

I -CI

H

CH,

Ph Ph

2-f

Ph

--OH.

H

I -C-CH,

I Ph + H

H

I o -CI + Nal

Ph

H

I -CH„-C-CHo-C Na

2 I 2 I

(12)

Ph Ph

+ Nal

67
The resulting polymers were found to be more

thermolabile than control samples also polymerized

anionically but "killed" with methanol. In the

second study [10] head-to-head structures were

introduced by incorporating small amounts of

stilbene during free radical polymerizations. A

copolymerized stilbene unit produces the structure

H H H H III I

Ph Ph Ph Ph

which, if anything, should be slightly weaker than a head-to-head unit. Rates of bond scission at subvolatilization temperatures for a series of such

copolymers with different stilbene contents were compared with a control homopolystyrene. The

results are shown in figure 4. At 280 °C the control

homopolymer does not undergo appreciable chain

scission. The copolymers, however, decompose

rapidly at this temperature and the rate of bond

scission increases with stilbene content.

An interesting preliminary study has been made of

the comparative thermal stabilities of head-to-head

and conventional poly (vinyl chloride) [21]. The

former was prepared by chlorination of 1,4-transpolybutadiene and therefore all the vinyl chloride

units in the chain were linked head-to-head. Using

4 0 &0

Time.hr

100

Figure 4. Degree of degradation, a, versus time of heating for

styrene homopolymer {287°) and styrene-stilbene copolymers

(280°).

9 Styrene homopolymer.

G 157, © 116, ® 72, Styrene-stilbene molar ratios in copolymers.

thermogravimetric analysis it was shown that the

head-to-head polymer began to lose HCl at an

appreciably lower temperature than the conventional one, although at a lower rate. It follows that

any head-to-head units in normal PVC could act as

initiation sites for dehydrochlorination. By way of

contrast, head-to-head poly (vinylidene chloride)

(prepared by chlorination of polydichlorobutadiene

[21 Jis more stable than its head-to-tail counterpart.

It was suggested that this improved stability in the

head-to-head polymer was due to the smaller number

of hydrogen atoms adjacent to each chlorine atom

and to the freedom from stripping of HCl from

sequences of monomer units.

2.4. Oxidation Products

The oxidative degradation of polymers is treated

in detail in other contributions to this symposium.

In this section a few examples will be discussed to

illustrate the effects of oxygen-containing groups in

subsequent thermal degradation. Many polymers are subject to slow oxidative degradation at ambient

temperatures, particularly if exposed to ultraviolet

radiation.

It is widely accepted that hydroperoxides play a

vital role in the early stages of oxidation of saturated

materials. Numerous secondary products accumulate

with the decomposition of hydroperoxide. The

hydroperoxide and the secondary products may be

regarded as abnormalities in the polymer chains and

may initiate thermal decomposition in a number of

ways. Homolytic decomposition of hydroperoxide

gives rise to hydroxy! and alkoxyl radicals which can

cause chain scission reactions at abnormally low

temperatures, or accelerate decomposition at higher

temperatures. In the case of polyethylene, it has

been suggested that the combination of oxygenated

structures and chain branches is largely responsible

for its relative instability [22]. These oxygenated

abnormalities may be introduced by incorporation

of impurities during preparation or by subsequent

oxidation. The degradation of polypropylene is also

influenced by oxidation products. Some investigators

have suggested that these include hydroperoxides

[23] presumably formed by oxidation at the tertiary

hydrogen atoms. Indeed, tertiary hydrogen atoms

are among the more sensitive sites in polymers for

oxidation. It is therefore not surprising that

poly (methyl acrylate) is subject to oxidative ageing

during storage and is considerably more thermally

unstable as a result [24].

The effects of oxygenated structures on the thermal

stability of polymers has not always been appreciated. According to Grassie [25] much of the

published work on PVC degradation is complicated

by the existence of oxidation products in the polymer

samples.

Oxygen may be incorporated directly into a few

polymers during free radical polymerization. For

example, in polystyrene it is believed to form a peroxide link [26, 27].

68
H

I 2 I Ph

H

I -C-0-0 •

I Ph

H

I

+ CH.

H

H

I P

I Ph

CH,

—CH^-C-O-O-CHo-C "2 2

Ph Ph

(13)

The resulting structure must be highly thermolabile

and can probably form in low concentration if air is not rigorously excluded from the polymerization

vessel. In figures 1 and 4 it will be noticed that the

plots of a (fraction of bonds broken) versus time

show a well defined ordinate intercept at zero time.

This behaviour has been interpreted as resulting

from a few highly reactive weak points in the polymer

backbone and the possibility that these are peroxide

units has been proposed [10]. This suggestion

remains unproved at present but it is noteworthy

that analogous data from anionically synthesized

polystyrene do not show this initial rapid decomposition [11].

Because of the stringent purity conditions re- quired and the nature of the propagating species it

is improbable that oxygen would be incorporated

in a similar fashion during anionic polymerization.

Weir and Lawrence have attributed the increased

photolytic instability of free radical polystyrenes to

the same source [28].

It should be made clear that the initial precipitous

drop in molecular weight of polystyrene during

pyrolysis at temperatures high enough to liberate

volatile material is not due entirely to the existence

of weak links of any structure. This behaviour is quite characteristic of any random fragmentation

process and also occurs in anionically synthesized

polystyrene. The effect of any weak bonds is to

accelerate this decomposition or to induce it at lower

temperatures.

2.5. Other Abnormalities

Almost all of the abnormalities discussed so far

reduce the pyrolytic stability of polymers. This,

however, is not an invariant rule and there are one

or two striking cases of abnormal structures which

actually stabilize the polymer chains. An example

of this concerns polymethacrylonitrile. Samples of

this polymer prepared with very pure monomer can

evolve nearly quantitative yields of monomer if heated rapidly to about 300 °C. This polymer is, however, sensitive to coloration reactions [29] which

occur fairly rapidly around 200 °C and even lower

when traces of methacrylic acid are present as comonomer. The coloration reaction is due to the

formation of conjugated carbon-nitrogen double

bonds.

C

I 2\ I ' C "

I c .

CH,2\

N

CH,

I -

C "

I C

CH.

The cyclic or ladder structure produced blocks the

unzipping reaction which is the source of monomer.

Similarly, the formation of anhydride structures in

higher polymethacrylates (as a consequence of

olefin elimination) blocks the unzipping process

which produces monomer [30].

3. Conclusions

The examples discussed above illustrate some of

the effects of abnormalities on the degradation

mechanism. In many cases, though not all, the effect

is on the initiation step of the degradation but

because of the susceptibility of polymers to radical

chain reactions (e.g. unzipping of poly (methyl

methacrylate) ) or to neighbouring group effects

(e.g. elimination of acetic acid from poly (vinyl

acetate)) the abnormality can effectively lower the

thermal stability of a large proportion of the sample.

It is also noticeable that much of the evidence for

the existence of labile abnormalities is indirect. This

is particularly true of polystyrene. The anionically

synthesized polymer is considerably more stable than

free radical material and it is a reasonable conclusion

that thermolabile structures exist in the latter, but

it has not been possible to follow the concentration

of these with time of pyrolysis. Indeed, there has

been no direct identification of thermolabile irregularities in this polymer. This point underlines one of

the practical difficulties in pyrolysis studies of

polymers, namely, the sensitivity of the polymer to

concentrations of abnormal structures which are often too low to be detected by conventional means.

Nevertheless, for meaningful kinetic and mechanistic

studies the starting material must be as fully characterized as possible, and it is the responsibility of

the investigator to consider irregularities in this

context.

4. References

[1] Bamford, C. H., Jenkins, A. D., and Johnstone, R., Trans. Faraday Soc. 55, 418 (1959).

[2] Grassie, N., Trans. Faraday Soc. 48, 379 (1952).

[3] Grassie, N., and Melville, H. W., Proe. Roy. Soc. (London) A199, 1, 14, 24, 39 (1949).

[4] Bullock, A. T., and Sutcliffe, L. H., Trans. Faraday Soc.

60, 625 (1964).

[5] MacCallum, J. R., Makromolek. Chem. 99, 282 (1966).

[6] Cameron, G. G., and Kerr, G. P., Makromolek. Chem.

115, 268 (1968).

[7] Cameron, G. G., and Kerr, G. P., J. Polym. Sci. A-1 7, 3067 (1969).

[8] Bailey, W. J., SPE Trans. 5, 1 (1965).

[9] McNeill, I. C, and Makhdumi, T. M., European Polymer

J. 3, 637 (1967).

[10] Cameron, G. G., and Kerr, G. P., European Polymer J.

6, 423 (1970).

[11] Cameron, G. G., and Kerr, G. P., European Polymer J.

4, 709 (1968).

69
[12] McNeill, I C, and Haider, S. I., European Polymer J. 3, 551 (1967).

[13] Wall, L. A., Thermal Degradation of Polymers, See. of

Chem. Ind. Monograph 13, London, 1961.

[14] Wall, L. A., and Straus, S., J. Polym. Sci. 44, 313 (1960).

[15] Cameron, G. G., and Kane, D. R., J. Polym. Sci. B2,

693 (1964).

[16] Cameron, G. G., Davie, F., and Kane, D. R., Makromolek. Chem. 135, 137 (1970).

[17] Grassie, N., in Chemical Reactions of Polymers, Ed.

E. M. Fettes, pp 637-638 (Interscience, N.Y., 1964).

[18] Cameron, G. G., and MacCallum, J. R., J. Macromol.

Sci. CI, 327 (1967).

[19] Bevington, J. C, Melville, H. W., and Taylor, R. P.,

J. Polym. Sci. 12, 449 (1954); ibid, 14, 463 (1954).

[20] Richards, D. H., and Salter, D. A., Polymer 8, 139

(1967).

[21] Amagi. Y., and Murayama, N., J. Polym. Sci. B4, 115

(1966).

[22] Oakes, W. G., and Richards, R. B., J. Chem. Soc. 2929

(1949).

[23] Van Schooten, J. W., and Wijga, P. W. O., Ref. 13, p. 432.

[24] Cameron, G. G., and Kane, D. R., unpublished work.

[25] Grassie, N., ref. 17, p. 634.

[26] Bovey, F. A., and Kolthofif, I. M., J. Amer. Chem. Soc.

69, 2143 (1947).

[27] Mayo, F. R., and Miller, A. A., J. Amer. Chem. Soc. 78,

1017 (1956).

[28] Weir, N. A., and Lawrence, J. B., Proc. I.U.P.A.C., on

Macromolec. Chem., Budapest, 1969, Vol. V, p. 323.

[29] Grassie, N., and McNeill, I. C, J. Polym. Sci. 27, 207

(1958).

[30] Grassie, N., and MacCallum, J. R., J. Polym. Sci. A2,

983 (1964).

Discussion

JVI. J. Brock (The Firestone Tire and Rubber

Company)

I was greatly interested in the comments of

Gordon Cameron with respect to thermolabile

irregularities which can be incorporated into the

polymer structure during polymerization and how

they can affect degradation reactions. We have

recently observed that the microstructure of polyisoprene polymers has a great affect upon some of the

products obtained during pyrolysis.

The polyisoprene system is convenient for study

since a wide variety of microstructures can be

prepared as shown below by varying the catalyst

system.

Percent Percent Percent Percent

cis 1,4 trans 1,4 1,2 3,4

units units units units

Natural Rubber (cis PI) 97.6 0 0 2.4

trans PI 0 98.0 0 2.0

Ziegler PI 96.8 0 0 3.2

Lithium PI 93.8 0 0 6.2

Alfin PI 10.9 59.2 5.1 24.9

Emulsion PI 24.4 62.0 6.1 7.5

Sodium PI 0.0 44.0 6.0 50.0

Li, PI, ether system 11.7 25.0 5.2 58.0

Everyone is familiar with Dr. Wall's work on the

pyrolysis of natural rubber, cis polyisoprene, to

produce the monomer, isoprene, by an unzipping

mechanism. If the unzipping reaction involves two

isoprene imits at a time, then dimers could be

produced also during thermal degradation. There are four possible alkyl substituted cyclohexene dimers

of isoprene.

All of these dimers are foimd in varying degrees in

the pyrolysis products of polyisoprenes having

different microstructures. Examination of the data

obtained indicates that cis and trans 1,4 polymers

produce mostly dimer III and some dimer II. Polymers containing 1,2 and 3,4 vinyl structures

produce increasing amounts of dimer II depending

upon the vinyl content. Extremely high vinyl

polymers produce mostly dimers III and IV. Dimer I was produced in significant quantities only from

Alfin, emulsion, and sodium polyisoprenes. Assuming

a two isoprene unit "backbiting" mechanism, dimer

I can only be produced from head-head structures

in polyisoprene.

1,4-Dlmethyl4-vlnylcyclohexene-1

III

1,5-Dlra3thyl5-vinylcyclohexene-1

Dlprene

70
Microozonolysis data have been obtained on these

samples and only Alfin, emulsion and sodium polyisoprenes give substantial quantities of 2,5 hexanedione to indicate the presence of head-head

isoprene units in these polyisoprene polymers.

These results are further evidence that pyrolysis

products and probably pyrolysis kinetics are affected

by irregularities incorporated into the polymer

structure during polymerization. I would like to reemphasize the statements of Gordon Cameron.

Complete characterization of polymer composition

and structure is extremely important where kinetic

and mechanistic studies are being contemplated.

G. G. Cameron:

Dr. Brock's contribution is a very elegant illustration of how the fine structure of a polymer can affect

the products of degradation. It also underlines the

importance of considering the preparative history of

each sample before making any generalizations.

S. E. Wentworth (Army Materials and Mechanics

Research Center)

Dr. Cameron, during your talk you mentioned the

importance of oxidation in the thermal degradation

of polymers. You further mentioned that tertiary

protons on the polymer backbone were most sus- ceptible to attack by oxygen. I am aware of at least one study on polyacrylonitrile where it is stated that

secondary protons are attacked in preference to

tertiary. Would you comment on this?

The study cited in the question is by J. Brandrup

and L. H. Peebles, Jr., Macromolecules 1, 64 (1968).

G. G. Cameron:

In hydrocarbon polymers tertiary protons are most susceptible to attack by oxygen. When strongly

electron withdrawing groups are attached to the

tertiary carbon atom, however, oxygen attack on the

proton is diminished. Under these conditions

oxidative attack may occur at secondary protons.

This point is discussed in detail by Brandrup and

Peebles [Macromolecules 1, 66-69 (1968)].

J. H. Magill (University of Pittsburgh)

In earlier papers of the sessions, notably those of

L. Wall and G. G. Cameron, much attention was

given to question of polymer molecular weight but

it was Dr. Brock of Firestone (a speaker from the

floor) who pointed out that the polymer microstructure was very important. He illustrated his

remarks with polyisoprenes. In this connection, I wish to inquire if either L. Wall or G. G. Cameron

have compared the relative decomposition rates of

isotactic with atactic well-characterized polymer

systems (e.g. polystyrenes). It seems at first sight

that the kinetics would be affected by the polymer

tacticity.

G. G. Cameron:

I am not aware that "much attention was given to

polymer molecular weight" in my paper. I have not

carried out studies of the effects of tacticity on polymer degradation but from the limited data in

the literature it seems that changes in tacticity do

not have a significant effect on the mechanism of

decomposition of polymers in categories (a) and (b).

Some differences in energies of activation or kinetic

coefficients are to be expected, however. For example

it has been shown that for poly (methyl methacrylate)

the kinetic chain lengths of the imzipping reaction

are greater in the atactic and syndiotactic polymers

than in the isotactic polymer. (Jellinek in The

Stereochemistry of Macromolecules, Vol. 3, Ed.

A. D. Kettley, Dekker, N.Y. 1968, pp. 378-385).

In studying effects of tacticity on the detailed

kinetics of degradation care must be taken to ensure

that the effects of tacticity are not obscured by the

presence or absence of abnormal structures in the

polymer samples prepared by different synthetic

methods.

L. A. Wall (National Bureau of Standards)

The tacticity of a polymer chain would be expected

to have only small effects on the kinetics of thermal

decomposition, assuming that one would be comparing the decompositions of both the atactic and

tactic polymers in the melt. The differences in the

energies of formation between the various tactic

structures are only a few hundred small calories.

Thus, large kinetic effects, for instance on the

thermal stability, are not anticipated. A good discussion of the current literature on this point can be

found in the article mentioned above: Degradation

of Stereoregular Polymers, by H. H. G. Jellinek

(Chapter 10, The Stereochemistry of Macromolecules, Vol. 3, editor Ketley, publisher Dekker).

F. H. Winslow (Bell Telephone Laboratories):

In the first section of Dr. Cameron's paper, the

relatively high monomer yield from polystyrene

pyrolysis was attributed to the low transfer rate of

the polystyryl radical. It was inferred that a-hydrogen atoms were readily available for reaction.

Most likely the phenyl ring does have a stabilizing

effect on the propagating radical but the enormous

stabilizing effect it has on the a-hydrogens may be

much more important as your work has so clearly

shown. For example, cumene and the isopropyl

groups in poly (p-isopropylstyrene) oxidize readily

in the dark at 80 °C. But the a-hydrogens in polystyrene are remarkably resistant to oxidation at 140 °C. If the a-hydrogens on the polymer chain were typical allylic hydrogens readily accessible to

molecular oxygen or other oxidation intermediates,

the oxidation rate of the polymer would be similar to

that of cumene or at least to that of poly(p-isopropylstyrene)

71
G. G. Cameron:

I agree with Dr. Winslow that steric effects must

also play some part in the degradation mechanism.

This is an important point and I am glad that Dr.

Winslow has raised it. This does not detract from

my own observation, however, that the reactivity of

the abstracting radical has to be considered along

with other factors.

72
NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

On Certain Problems Connected With the Inhibited Oxidation Theory

Yu. A. Shlyapnikov

Institute of Chemical Physics, Academy of Sciences of the USSR, Moscow

The oxidative reactions of organic compounds and polymers are discussed in terms of free radical mechanisms, with particular emphasis on initiation, inhibition, and stabilization phe- nomena.

Key words: Antioxidants; inhibition of oxidation; initiation of oxidation; oxidation mechanism;

oxidation theory.

The main reason for aging of organic compounds is their interaction with air oxygen. The oxidation of

hydrocarbons is that investigated most extensively.

It occurs also with compounds containing molecular

hydrocarbon fragments. The mechanism of hydrocarbon oxidation had been essentially elucidated by

1946. It was found to involve two alternating steps

of chain propagation [1,

R-+02-^R02 (1)

R02+RH^R00H+R (2)

Here RH is the hydrocarbon and R the relevant

free radical.

Nevertheless the mechanism of hydrocarbon

oxidation cannot be considered as completely

established.

The sources of free radicals initiating oxidation

usually represent certain additives, various types of

radiation including light, and mechanical degradation. In the case of "noninhibited" oxidation the

main sources of free radicals are suggested to be the

reactions involving various impurities. The author

has no knowledge of papers on measurements of the

rates of direct interaction between hydrocarbons and

molecular oxygen. Such interaction is considered to

be the initiating step:

RH+02-*R +HO2 (3)

When hydroperoxide is formed in the course of

oxidation, its decomposition becomes the main

source of free radicals. Without discussing yet the

mechanism of this reaction in detail, let us write it in the form

(RH...)

ROOH >5R +inactive products (4)

Chain termination may occur by recombination of

two radicals, for example

* Figures in brackets indicate the literature references at the end of this paper.

R +R ->R2 (hydrocarbon) (5)

RO2+RO2—^inactive products

Recombiaation reactions yielding peroxide, such as R02+R->R00R (6) or 2R02->ROOR+02 (6a)

terminate the chain only when further peroxide

decomposition yields on an average less than two

free radicals per peroxide group.

The hydroperoxide group decomposition shoidd

be discussed in detail.

Decomposition of a hydroperoxide group to

two radicals

ROOH^RO + OH (7)

requires a high activation energy. In the solid or

liquid phase this reaction will be negligible because

of the cage effect that considerably decreases the

rate of reaction (7). Another reason is the competition of more fast reactions, such as [3, 4]

RO +R +H2O

/

ROOH+RH (8) \

ROH+R +0H

Intracage interaction of radicals occurring by

reaction (8) would not affect the reaction rate, but

would only decrease the average number of radicals

emerging from the cage.

When foreign molecules, such as O2, are present in

the cage, the probability of emergence of a free

radical becomes higher, as recombination of the

resulting radical RO2 with RO or OH, as well as that

of interaction between RO2 and the surroimding RH,

would differ from the rates of similar R reactions.

Let us denote by X the cage containing the radical

pair. Then the scheme of chain branching will be

73
R

ROOH+RH

where po and pi are the probabiHties of the emergence

of one radical out of a cage containing no (po) and

containing O2.

In accordance with (9)

fc8[R00H][RH] k',+k','[_02'] ^ ^

or for low concentrations of O2, for which k'^^k'g' [O2]

5=/>o+/>i^ CO.,] = 5o+5i[02] (10a)

"9

The relation (10a) was confirmed experimentally

when studying the degradation of hydroperoxide

groups of polypropylene in O2. Use was made of

radicals of a definite type Ry, such as CHj or microradicals with end free valence, capable of depolymerization by scheme —CH2—CH ( CH3) CH2-»CH2 +CH3—CH=CH2

(11)

Both radicals can also react with the surrounding

species RH (with a constant kij) and with 02(^2;).

The resulting Ry concentration would be

5A[RH][R00H]

L^^J- /ci,[RH] +fe2.[02] ^'^^

and at 5y = const would tend to zero with increasing

[O2] pressure. Since the concentration of Ry radicals

tends to a certain limit, dj depends on [O2] ac- cording to (10a). The methane and propylene yields

obtained by first-order Ry conversions tend to

limits differing from zero with increasing [O2]

(fig. 1, curves 1 and 2).

0 4. 0

0 200 m eoo p ,nmi Hg

Figure 1. Yields of methane (a) and propylene

(2) in polypropylene hydroperoxide decomposition as a function 0} oxygen pressure.

100 °C, time of decomposition 60 min, [ROOH] =

3.6.100* mol/g.

Until recently the dependence of degenerate chain

branching probability 8 on O2 pressure was usually

neglected.

Most investigations on the mechanism of hydrocarbon oxidation were carried out at temperatures

up to 150°. One of our papers deals with oxidation of

linear polyethylene at temperatures higher than

200°. At high temperatures (200-400 °C) the

kinetics of oxidation becomes different. Within this

temperature range the oxidation rate dependence on

oxygen pressure becomes first order in (fig. 2).

The effective activation energy decreases from

20-30 to 1—2 kcal/mol and a decrease in oxidation

rate with increasing temperature is observed in

some cases (fig. 3). Obviously, the value 5 increasing

with temperature mainly due to the rapid increase in

do reaches its stoichiometric limit and depends no

more on oxygen pressure.

Figure 2. Curves for oxygen absorption in polyethylene oxidation at 360 °C plotted as log po/p versus time.

Oxygen pressure 20 mm (1); 40 mm (2); 60 mm (3); 80 mm (4) and 100 mm (5).

-L 1 1 1 100 250 iOO ISO iOO TC

Figure 3. Initial rate of polyethylene oxidation

as a function of temperature.

Oxygen pressure 20 ram (1); 40 mm (2); 60 mm (3); 80 ram (4) and 100 mm (5).

At the same time RO2 decomposition at a rate

increasing with temperature (Eeff 20 kcal/mol)

becomes the main route of RO2 conversions. The

decomposition of ROj in polyethylene yields, in

particular, OH radicals and epoxy-groups that can be detected by ir-spectroscopy

74
00*

ICH2-CH-CH2-

•M3

^3

CiLjC OGH2-"

/°\ . ~CH2-CH - CH~ + OH

0

II —CHg-C-CH^^ + OH

(13)

The main route of chain termination is now the

recombination of most mobile radicals OH with

radicals R of various types:

R +OH-^ROH (14)

RH+OH-^R+HaO (15)

Taking into account reaction (15) and assuming that

A:i4[R J<<CA:i5[RH], the rate of high -temperature

oxidation will be

r=A;i[R-][02] = (16)

New facts were observed in investigating the

oxidation of a solid crystalline hydrocarbon polymer

(polypropylene) . Its total crystallinity depends on

its chain regularity and does not depend on the

technique of sample preparation. Nevertheless, the

induction period of oxidation for samples prepared

by rapid cooling of the melt and consisting of small

spherules is far shorter than that for samples consisting of big spherules and prepared by slow cooling.

In both cases oxidation occurs in amorphous

regions consisting mainly of interspherulic substance

for rapidly cooled samples, and of itraspherulic,

interfebrilic substance for slowly cooled samples.

One of the probable explanations would be the

difference in thickness of layers of amorphous

regions. According to Semenov, the rate of chain

termination on the reactor wall [5^ is

JFg = ad-' (17)

Consequently, the length of unbranched oxidation

chains will be

,.= 1+ ^^^^ =l+fe2[RH]<i2a-i (18)

and the factor of autoacceleration (coefficient in

equation W= Wo expfr) is

'p8-

= A:8[RH] {l±k£RE-\d'a-') (po/r^+PiMO^])

k',+k','\:02']

(19)

where 6 is the average lifetime of a chain branching

product, i.e., l/ZcgCRH].

The induction period r is determined from

JF= Wo exp(pr = const, i.e., ^7 = const. Hence

const const

T =

X

ki +k','[_0.2'] mpok£RR:\d'-a-'+po-i)

+&^' (pifc2[RH]d2a-i +pi - 1 ) [O2]

(20)

When the oxygen pressure is high enough, expression

(19) can be written as

[O2] p

(21)

where p is the oxygen pressure.

It will be seen from figure 4 that the function of

type (20) is actually fulfilled, the coefficients A' and

B depending on conditions of sample preparation.

At high degrees of oxidation recombination of free

radicals becomes the main route of chain termination

and the rates of oxygen consumption appear to be

almost identical for different samples.

Inhibitor [IH] added to the oxidizing substance

terminates the kinetic chains, mainly by reaction

with peroxide radicals RO2, e.g., [6]

R02+IH^R00H+I (22)

r.

nun

Figure 4. Induction period for polypropylene oxi- dation as a function of the oxygen pressure

reciprocal (130°).

1—slowly cooled samples; 2—rapidly cooled samples.

75
The inactive radicals I yielded by this reaction

mostly disappear and do not initiate new chains. But

the actual inhibitor function is not confined to chain

termination. Phenols and amines used as inhibitors are more readily oxidized than the main substance,

and initiate oxidation of the latter [[7]. This process

reduces the inhibitor efficiency. But to our opinion

for more important is the participation of the

inhibitor in degenerate chain branching. It was

stated above that most radicals formed from the

hydroperoxide group recombine in the cage. For

instance, the yield of light-weight methyl radicals

formed by decomposition of solid polypropylene

hydroperoxide at 100° is 5 = 7X10~^, and its value in

the melt must be much lower. A particle penetrating

into the cage and destroying one of the free radicals

in a pair will prevent destruction of the second

radical and the emergence of free radicals out of a cage containing such particles will approach unity.

Common inhibitors containing one active group,

though capable of destroying consecutively both

radicals of the pair, will react with the first and

second radicals at different rates. That must also

enhance the radical emergence. It may be considered

approximately that the cage volume is (10 A) ^ =

10"-'' 1. When the inhibitor increases the probability

of radical emergence to 0.5, only 0.2 percent of the

inhibitor content in cages is sufficient for increasing

d from 1 X 10~^ to 2 X 10~^. The inhibitor concentration will then be 102^X2 X 10-^ = 2 X lO^i molecules/litre =3.3 X10~^ mol/litre. The common concentrations of inhibitors are (5-30) XlO^^ mol/litre.

That is sufficient for a great change of 5. The inhibitor effect on the probability of degenerate

chain branching was proved directly by Pudov's

investigations [8] on polypropylene hydroperoxide

decomposition in the presence of ionol.

Taking the above into account, the scheme of

oxidation in the presence of an inhibitor [IH]

will be:

RH+Os^R +HO2

ROOH +RHIH+02^I

R +02->R02

R02+RH->R00H+R

R02+IH^R00H+I

(IHi02)

(RH)

HO2 >/3[R-]

(0)

(23)

(1)

(2)

(22)

-*5R -|- inactive products (24)

When a hydroperoxide reducing agent [Y], for

example dialkyl sulfide [9], is added to the system

along with the inhibitor [IH], another step must be

added to the scheme

given with sufficient accuracy by an expression

similar to (10a)

5 = 5o+5,[IH] (26)

Allowance for the dependence of 5 on [IH]

permits explaining virtually all phenomena observed

in studying inhibited oxidation of hydrocarbons and

polyolefins. Among them of greatest interest are the

critical concentrations. The set of balance equations

for R , RO2 and hydroperoxide will be:

=afeo[RH][02] +/3fc23[IH][02]

+A-i[R-][02]+fc2[R02][RH]+5/c24CROOH][RH]

(27)

=/ci[R-][02]-A-2[R02][RH]

-fe22[R02][IH] (28)

rf [ROOH]

dt

^^2[R02][RH]+A^22[R02][IH]

-A-24[ROOH][IH]-A-25[ROOH][R25] (29)

Using the Bodenstein-Semenov method [10], i.e. assuming that d[B. ~\/dt = 0: t/[RO = 0-, but

d[JX00n]/dt9^Q, we can find the solution:

[R00H] = .4e^«+fi (30)

where, assuming []R2S ] = 0

fc24[RH] , . rTTT-|2

<p = <Po=

^.^^j-jjj-j

[0 Mii.i-ti

+ (6,A2[RH]-A-22 +U22)[IH]+U2[RH]} (31)

and when [R2S Jj^O

+ (5.fc2[RH]-fe22 +U22) [IH] +U2[RH]1

-A-25[R2S] = ^0-A-25[R2S] (32)

The kinetics of oxidation depends on the sign of f

When <(5>0, the hydroperoxide concentration will

rise autocatalytically. When ^<0, the hydroperoxide concentration will be an unambiguous

function of the rate constants and concentrations of

RH, O2 and IH.

Solution of equation v5([IH]) =0 gives the critical

values of inhibitor concentration

ROOH+y^inactive products (25) [IH]cr =

This scheme does not take into account the ter- mination by recombination, as its rate is negligibly

low in the presence of an inhibitor.

The dependence of degenerate chain branching

probability on the inhibitor concentration may be

(l-5Q)fc22-5^fc2[RH]

2^^022

\/5?A-i[RH J- 2 ( 1 - 5o) 5 ,-A:2A22[RH] + ( 1 - hYk^

2ik22

(33)

76
When 5,A-2[RH]-(1- V5o)'fc2i>0, equation

(p([IHJ) =0 has no real solution. In this case (p>0,

independently of the inhibitor concentration, i.e. the process is not stationary at any inhibitor concentration. If 5,7c2[RH]- (1- v^)2A-2i<0, the secondorder equation 95([IH]|) =0 (see 31) can have two

real solutions namely [IH] and [IHJa? which we call

lower and upper critical concentrations. When

[IH] < CIH]i, <p>0 and, in accordance with (30) , the

oxidation process is autocatalytic. If [IH3 = CIH3i,

becomes zero and remains negative up to

[IHJ = [IHJ2. Within this range oxidation is sta- tionary. When [IH]]> [IH]]2, oxidation again becomes autoaccelerated and its mean rate simply

increases. The strong bends of curves for the induction period as a function of the inhibitor concentration correspond to critical concentrations []IHJ

and [IHo] (fig. 5).

In the case of weak antioxidants, i.e. <p[IH]]>0 at any [IH], oxidation can be

stationary by addition of a substance R2S reducing

hydroperoxide. At a sufficient concentration of R2S

equation ip = v?o — ^24CR2SJ = 0 will also have two

positive solutions. Such a conversion of a weak

antioxidant into a strong one by action of substances

reducing hydroperoxides is well known as synergism.

The low critical concentration [IHJi of the weak

antioxidant in the presence of didecyl sulfide is illustrated in figure 6. The effective antioxidant

hinders oxidation converting the autocatalytic

when

made

20 Mftr', moe/ig

Figure 5. Induction period for polypropylene oxidation as a function of the 2 , 2' -niethylene-bis{4-chloro-6-teri-butylphenol)

concentration.

180° (1); ]90° (2) and 200°C (3); PO2 = 300 mm.

Figure 6. Induction period for oxidation of pure polypropylene

(1) and of propylene containing didecyl sulfide, 0.08 mol/kg,

as a function of 2,4,6-triterbutylphenol {200°C).

Figure 7. Distribution of dimensionless concentrations of inhibitor I and active center N {dotted line) in various time intervals.

t = 0 (0); 1 (1); 5 (2); 12.5 (3); and 17.5 conditional units

(4). The uninhibited part size exceeds the critical value.

process into a stationary one. At the same time the

common antioxidants, usually phenols and amines,

that become readily oxidized, initiate oxidation.

However, though the stationary reaction rate is increased by initiation, it remains many times lower

than that of uninhibited oxidation. Thus, the main

function of the antioxidant is the conversion of an

unstationary process into a stationary one, and the

theory of the antioxidant action is essentially that of

critical phenomena in inhibited oxidation. In some

cases the antioxidant efficiency can be predicted

from its molecular structure.

The autocatalytic increase in the reaction rate can

not continue infinitely. Oxidation reaches its limit

due to the rapid increase in the radical recombination

rate and becomes stationary. But the rate of such a stationary reaction is by many orders of magnitude

higher than those of both unstationary and stationary

inhibited reactions.

Some investigations have been carried out recently

on the oxidation of unevenly inhibited samples. It was found |]11] that if one part of the sample is deprived of the inhibitor, a fast reaction initiated in

this part of the sample can propagate over the whole

sample (fig. 7). The minimal size required for the

uninhibited domain to become the center of fast

oxidation depends on the inhibitor concentration in

the remaining part of the sample. For this reason the curves for the induction period as a function of the

average inhibitor concentration are markedly different for uniformly inhibited propylene and for a polymer consisting of spheric particles (0.1 mm)

containing the inhibitor on their surface only (fig. 8)

too

Figure 8. Induction period for polypropylene oxidation as a function of (CH3)2Si(OC6H4HNCoH5)2 concentration, for

uniform {l) and nonuniform mixing (2) {200°C).

11
The questions discussed above certainly do not

cover all theoretical problems connected with

practical stabilization of organic compounds.

References

[1] A. R. Ubbelohde. Proc. Roy. Soc. A152, 345 (1935).

[2] J. L. Bolland, G. Gee, Trans. Faraday Soc. 42, 236, 244

(1946).

[3] D. G. Knorre, Z. K. Maizus, N. M. Emanuel. Dokl.

Akad. Nauk SSSR 123, 123 (1958).

[4] V. S. Pudov, M. B. Neiman. Neftekhimiya 2, 918 (1962).

[5] N. N. Semenov. Tsepnye Reaktsii, Leningrad, ONTI,

(1934).

[6] J. L. Bolland, P. Ten Have. Trans. Faraday Soc. 43, 201

(1947); Disc. Faraday Soc. No. 2, 252 (1947).

[7] H. Winn, J. R. Shelton. Ind. Eng. Chem. 40, 2081 (1948).

[8] V. S. Pudov, L. N. Tatarenko, Vysokomol. Soedineniya

All, 2773 (1969).

^

[9] G. H. Denison. Ind. Eng. Chem. 36, 477 (1944); G. H.

Deni|on, P. C. Gondii. Ind. Eng. Chem. 37, 1102

[10] N. N. Semenov, Zhur. Fiz. Khim. 17, 187 (1943).

[11] L. A. Lovachev, Z. I. Kaganova, Ya. A. Shlyapnikov.

Dokl. Akad. Nauk SSSR, 183, 379 (1968).

Discussion

F. R. IMayo (Stanford Research Institute)

While Dr. Shlyapnikov's manuscript contains

some new and interesting ideas and data, some of the

reactions and equations proposed deserve some comment. Reactions 8 can hardly occur in one step. I know of no adequate precedent for rearrangements of

peroxy radii als by reactions 13' and 13". The molecular products indicated probably arise via alkoxy

radicals and hydroperoxide molecules, respectively.

While 13" has adequate precedent, cyclic ethers with

larger rings (preferably tetrahydrofuran derivatives)

are formed much more easily from n-alkanes.

Reaction 14 violates the general principle that

termination involves mostly the least reactive and

most plentiful radicals; 1 would expect HO radicals

to disappear by 15 before they could find an alkyl

radical for 14.

The set of equations on page 72 neglects at least

two reactions which are probably important, ter- mination by RO2 +I and reinitiation by I +RH.

Without allowance for these, the subsequent rate

expressions cannot be exact although the conclusions

may be qualitatively useful.

R. E. Florin (National Bureau of Standards)

In relating induction period to geometry of

amorphous regions, the Semenov equation (17)

leads to the right answer but its use can be questioned on physical grounds. Wall terminations are important

in gas reactions because of the need to remove excess energy. In condensed systems a preference for wall

termination is harder to understand. IVIight not the

longer induction periods in slowly cooled polymer be

associated instead with a more tortuous diffusion

path for O2 and a consequent local deficiency in the

^sTOs] contribution?

Also, the lines 1 and 2 in figure 4 seem to have a common intercept, implying equal values of B in eq

(21). In contrast, the expansion of eq (20) in powers

of

-| calls for a term B varying with the size

parameter d^. In fluorination, it is easy to observe a process

analogous to direct initiation, RH+O2—>R +IIO2,

reaction (3) in Shlyapnikov's text. If dilute F2 is

allowed to flow over a finely divided polymer in an ESR cavity, free radical spectra appear rapidly.

With polystyrene fluff and 2 percent F2 in He intense

spectra appear in a single scan time (2 min) but are

largely due to peroxy radicals RO2 with only weak

shoulders from hydrocarbon radicals. Polyethylene,

2X10"^^ m film, with 5 percent F2, requires 10 min

to develop full intensity, but its hydrocarbon radical

spectrum is better separated and more intense.

Polymer radical ESR spectra have been produced

previously by the action of H atoms on a fluff of

polystyrene [1], H +RH-^R +H2. With fluorine,

the molecular form is sufficiently reactive, F2+RH-^

R+F+HF. The energy liberated presumably

breaks up pair correlation and delays recombination.

IMolecular fluorine has been used to initiate polymerization [2], and a dissociative addition has been

observed already by ESR [3 J, F2+CH2=CH—CH=CH2

^FCH2CHCH=CH2+F .

In the latter case, only peroxy radicals and no

primary hydrocarbon radicals were observed. The

peroxy radicals observed in these experiments arise

from reaction of primary radicals with oxygen,

usually present in commercial fluorine to the extent

of 0.2 to 2 percent, R +02->R02.

Abstractions by O2, reaction (3) above, must have

higher activation energy than those with F2, but

might be observable in more rigid polymers at higher

temperatures. There is an outside possibility that

many-lined radical spectra observed on heating

polynuclear aromatic hydrocarbons \_^~\ could have

been due to reaction (3) rather than to simple

pyrolysis.

References

[1] R. Ingalls and L. A. Wall, J. Chem. Phys. 35, 370 (1961);

ibid 41, 1120 (1964).

[2] R. E. Florin, L. A. Wall, D. W. Brown, L. A. Hymo, and

J. D. Michaelsen, J. Res. Nat. Bur. Stand.' 53, 121

(1954).

[3] G. A. Kapralova and A. E. Shilov, Kinetika i kataliz 2,

362 (1961).

[4] W. F. Forbes and J. C. Robinson, Nature 214, 80 (1967).

78
NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Mechanism of Peroxydation of Polymers

Adolphe Chapiro

Laboratoire de Chimie des Radiations du C.N.R.S., 92-Bellevue, France

The irradiation of polymers in the presence of air leads to polymeric hydroperoxides POOH and

diperoxides POOP. When such irradiated polymers are used to initiate the polymerization of a monomer, hydroperoxides generate equal amounts of graft copolymer and homopolymer, whereas

diperoxides only generate graft copolymer. This technique was used to study the peroxidation of

various polymers under different irradiation conditions. It was found that polypropylene irradiated

at temperatures ranging from 10 to 60 °C undergoes hydroperoxidation by a chain reaction.

At 0 °C and below (in the vitreous polymer) polymeric radicals are formed but remain trapped.

Poly (vinyl chloride) also leads to hydroperoxides when irradiated at 20 to 60 °C but no chain

process develops below the glass transition temperature. Irradiation of the plasticized polymer

leads to a hydroperoxidation by a short chain mechanism. It is concluded that the propagation

of the hydroperoxidation chain, involving hydrogen abstraction by PO2 radicals, only occurs above 7'„ when the mobility of polymeric segments is sufficient to ensure the interaction of PO2

with neighboring molecules.

Polyethylene irradiated at room temperature chiefly generates diperoxides POOP. The chain

hydroperoxidation only sets in above 35 °C. Below this temperature, recombination of PO2 radicals occurs faster than hydrogen abstraction. The behaviour 01 polymers with respect to peroxidation

is discussed and compared with that of low molecular weight model compounds.

Key words: Hydroperoxides versus diperoxides; influence of polymer phase on peroxidation;

influence of temperature on peroxidation; peroxidation of polyethylene; peroxidation of polymers;

peroxidation of polypropylene; peroxidation of poly (vinyl chloride).

1. Introduction

When subjected to high energy radiation in the

presence of oxygen most polymers undergo a peroxidation process. This reaction follows the

classical scheme which can be written as follows:

Initiation:

PH/vM-P +H (1)

Propagation:

P+02-^P02 (2)

PO^+PH^POzH+P (3)

Termination:

P02 +P-^P02P (4)

P02+P02^P02P+02 (4a)

According to this scheme, polymeric hydroperoxides POOH arise as a result of a chain reaction

(with propagation steps 2 and 3) while diperoxides

POOP (or peroxidic crosslinks) are formed in the

combination of two radicals. When such polymeric

peroxides are used to initiate a graft-copolymerization

reaction, diperoxides generate only graft-copolymer,

while hydroperoxides lead to equivalent amounts of

graft-copolymer and homopolymer as shown below.

+nB

P00P^2P0 > 2F0Bn (5)

graft-copolymer

POOH^PO+HO-^ FOBn (6)

graft-copolymer

+

ROBn

homopolymer

It thus appears that a simple determination of the

relative amounts of homopolymer versus graftcopolymer formed in the reaction provides information on the peroxidation process.

2. Experimental

Films of polymer 0.1 to 0.3 mm thick were subjected to different doses of gamma-rays in the

presence of air at different temperatures. These

films were thereafter sealed under vacuum in glass

tubes containing a large excess of acrylonitrile. The

tubes were then heated in a thermostat at temperatures usually ranging from 80 to 130°. The weight

increase of the films as well as the amount of polyacrylonitrile homopolymer formed were determined

as a function of time. Experiments were conducted

with polyethylene [1, 2, 3],^ polypropylene [2, 4]

and poly (vinyl chloride) [5, 6]. Further experimental details are given in the original papers.

^ Figures in brackets indicate the literature references at the end of this paper.

79
P-R

par heure

Dose ( Mrad ) Figure 1. Rate of grafting of acrylonitrile on preirradiated

polyivinyl chloride) films as a function of dose of gamma-rays

[5, 6].

curve 1: 20 °C; dose-rate 76 rads/min. curve 2: 40 °C; dose-rate 73 rads/min.

curve 3; 60 °C; dose-rate 74 rads/min. curve 4: 40 °C; dose-rate 7 rads/min.

3. Results and Discussion

Early experiments [1, 2] have shown that ir- radiation of polyethylene at 20 °C almost exclusively

leads to diperoxides POOP whereas polypropylene

under similar conditions generates hydroperoxides.

This result was accounted for by assuming that step

3 does not occur in polyethylene at 20° but could

take place in polypropylene on behalf of the labile

tertiary hydrogens present in this polymer. These

data conflicted, however, with the results of Bakh

et al [7]] who found that neither ra-heptane nor

isooctane (which contains tertiary hydrogens)

underwent a chain hydroperoxidation when irradiated

at room temperature. A chain reaction only developed

above 40 °C.

Further work conducted with poly (vinyl chloride)

showed that this polymer generates hydroFlGURE 2. Conversion curves of the graftcopolymerization of

of acrylonitrile at 80°C onto polypropylene films preirradiated

in air at different temperatures {as indicated On the curves).

Dose of preirradiation: 0.3 Megarads [4J.

peroxides under irradiation. However, the peroxide

yield was found to be independent both of irradiation

temperature and dose rate, thus precluding a chain

reaction [6] (fig. 1).

A systematic study was conducted with polypropylene. Films of this polymer were irradiated at

temperatures ranging from —20 to +60 °C and

acrylonitrile was thereafter grafted at 80° [4]. The

conversion curves obtained after a dose of 3 X 10^

rads are shown in figure 2. An Arrhenius plot of the

initial rates is presented in figure 3. It can be seen that for irradiations conducted

between 20 and 60 °C the yield of peroxides (rate of

grafting) steadily increases as required by the

conventional chain peroxidation (steps 1 to 4).

Below 0°, the yields are temperature independent

and the shape of the conversion curves (fig. 2)

denotes a change in mechanism, suggesting that in

polymer films irradiated in this temperature range

the grafting is initiated by trapped radicals. The

P P \ -^-ijxlOO par heure

3,8 A,2

1/Tx10'

Figure 3. Arrhenius plot of the initial rates of grafting of

acrylonitrile onto polypropylene derived from the conversion

curves shown in figure 2 [4].

sudden change in mechanism appears more dramatically on figure 3. It is of interest to note that the

glass transition temperature (Tg) of polypropylene

is reported to be near 0°. The above results thus show

that polypropylene irradiated above it Tg suffers a chain peroxidation the chain length of which steadily

increases with temperature. Below Tg the polymeric

radicals P and PO2 cannot react for lack of mobility

and they remain trapped in the polymer matrix.

These results were confirmed by further experiments with poly (vinyl chloride) (PVC). A hydroperoxidation chain process does not develop in this

polymer at room temperature since its Tg is 70 °C.

The hydroperoxides POOH formed in this polymer

presumably arise as a result of the recombination of

trapped PO2 radicals with small, mobile radicals

such as H - or HO2:

POi+HPOi+HOi-

>P02H

>P02H +02.

(7)

(8)

Experiments conducted with poly (vinyl chloride)

films which had previously been "plasticized" with

ca. 15 percent of chloroform showed that the yield of

peroxides significantly increases. Figure 4 shows the

80
grafting rates as a function of dose of gamma-rays

for the plasticized and unplasticized polymers. If one assumes that the rate of grafting is first order

with respect to peroxide concentration (termination

by radical burying) the data shown in figure 4

suggests a chain length of ca. 10 for the hydroperoxidation in the plasticized polymer.

In spite of the fact that polyethylene has a very

low lying Tg( — 40 °C) irradiation of this polymer at

20 °C chiefly leads to diperoxides, presumably via

reaction step 4. By raising the irradiation temperature it was found Q3j that a chain hydroperoxidation

sets in above 35 °C. An Arrhenius plot of the data is shown in figure 5. At 70 °C the yield of peroxides

exhibits a maximum. It is most probable that above

this temperature a significant decomposition of the

peroxides takes place during irradiation.

A ^_ ^ vlOO par heure

p 100|_ ° 70°C

i

50

^0

30

20

30 32 3^

X 10 =

Figure 5. Arrhenius plot of the rate of grafting of acrylonitrile onto polyethylene films preirradiated to a dose of 0.3 megarads

at different temperatures [3].

P-Po

par heure

3 A Dose ( M ra d )

Figure 4. Influence of dose on the rate of grafting acrylonitrile onto poly{vinyl chloride).

curve 1: unplasticized PVC;

curve 2: PVC plasticized with ca. 15 percent chloroform [6].

It should be pointed out that in this polymer the

change in mechanism at ca. 35° only results from the

competition involving reactions 3 and 4. Reaction 3

requires a higher activation energy and is thus

favored at elevated temperatures. This competition

is further affected by the concentration of PO2

radicals since these reactions are respectively of first and second order with respect to (PO2). It thus

follows that the shift from the formation of peroxides

by a single step reaction such as 4 to a chain hydroperoxidation (i.e., the break in the Arrhenius plot of

figure 5) should depend on the radiation dose-rate.

Such an effect was indeed observed by Saraeva [8^

who found that the onset of the chain hydroperoxidation of heptane and other low molecular weight

compounds was shifted to higher temperatures the

higher the dose-rate. An Arrhenius plot of these data

is schematically shown in figure 6.

It should be further noted that reaction step 4

which involves the interaction of two polymeric

radicals PO2 is necessarily hindered by the low

mobility of such radicals in the polymeric matrix.

The corresponding reaction

ROa+ROi^ROzR+R (9)

is much faster in a low molecular weight Uquid than

in a polymer and this accounts for the observation

that polypropylene undergoes a chain hydroperoxidation under experimental conditions (temperature, dose-rate) under which isooctane does not.

This emphasizes the fact that low molecular liquids

cannot be taken as reliable model compounds for

studying peroxidation of polymers at low temperatures.

4. Conclusions

The data presented in this work all pertain to

radiation initiated peroxidation of polymers. However, since radiation peroxidation can be entirely

accounted for by conventional free radical processes

Log G ( ROOH

)

1, > I2 > I3

1, I2 I3

I/T1

1 I/T3 1 /T

1/Tj

Figure 6. Schematic plot of the radiation chemical yields of

peroxides in low molecular weight hydrocarbons, based on

data of Saraeva [8].

81
the conclusions derived from these results are more

general in character and applv to anv mode of

initiation. It thus appears that when free radicab are generated in a polymer at temperatures ahove its glass transition point, a chain hydroperoxidation

reaction may arise under appiropriate conditions.

The onset of the chain process is governed by a competition for perosidic radicals which may either

abstract hydrogen atoms from the pohTner or combine. As in all chain reactions, higher temperatnres and lower steady state concentrations of

radicals favor the propagation step. Furthermore,

since the combination of radicals occurs much faster

in a mobile liquid than in a viscous pohTner medium,

the propagation step i^ favored in polymers and,

hence, a chain hydroperoxidation process occurs more readily in a pohTner than in a low molecular

weight model compound treated under identical

conditions.

Below the glass transition point of the polymer,

most poKTneric radicals remain trapped and thus

accumulate. Peroxides mav then still arise as a result

of the combination of such trapped radicals with

small radicals which are still mobile in the medium

such as for instance H- and HO;. This leads to the

formation of hydroperoxides bv a one step reaction

which 15 independent of temperature and of the rate

of free radical generation.

5. References

[11 A. Chapixo. J. Polvm. 5ci.. 29. 321 '1938 : 34. 439 ^959).

\1\ A. Chapiro, J. PolVm. Sci.. 4S, 109 '1960 . [3^ A. Chapiro. A. \f. Jendrychowska-Bonamour and J. P. Leca, I.L.P.A.C. Conference on Chemical Traosformations of Polvmers. Bratislava. June 1971. Preprints,

Yol. n, p. 46:

[4] A. Chapiro and T. Czvikovizki. J. Chim. Phv§.. 66. 1-502 '1969.

A. Chapiro and Z. Mankowski. Europ. Polvm. J.. 2. 163

'6 M, Poex and A. M. Jendrychowski-Bonamour. ILTPAC

Sympos. Alacromol. Chem. Budapest 1969; A. Chapiro,

AI. roes and A. M. Jendrvchowska-Bonamour. Europ.

Polvm. J. 7. 1241 a971 .'

[7] K. A.'Bakh, Radiation E.e5earch SuppL, 1. 190 '1959j.

[8] V. V. Saraeva, Linetika i KataLLz, 6, 537 :i965,.

Discussion

J. H. ]\IaglU Lniversity of Pittsburgh :

It is not clear to me whv low molecular weight

compounds cannot be used as model compounds to study reactions in polymers provided that experiments were conducted under isoviscous conditions.

Compare.- for example, the viscosity data on polymers v\ith small molecules. If the Stokes -Einstein

relationship is applicable- to the molten polymeric

state C. Pv, Patrick. ""Factors involved in the

degradation of polymers in melts"' . then the

mobditv of radicals would be comparable if viscosity

or diffusivitv of polvmer segments or molecules of

the model compound ^v'ere the major rate determining factors.

A. Ckapiro;

The peroxidation of a polymer in bulk is governed

bv two reaction steps: 1 propagation of the

peroxidation chain in which a PO; radical abstracts an atom from a neighboring P\-II groups; (2)

termination of the chain by interaction of trwo polymeric radicals PO,—POj or POf—P-. P^eaction

(1) involves the interaction of a pohTneric radical

w-ith a segment of a neighbor molecule. Such a process could perhaps be studied with low molecular

weight model compounds since only a small displacement of poK"mer segments is required for the reaction. However, reaction (2) which requires the interaction of two free radical sites, attached to polymer

molecules, w-ith verv low or even neghgible mobihty

in the reaction medium and which must diffuse over

- See firare 3 in G. C. Bern- and T. G. Fox, AdTances in Polymer Science,

5, 261 (1962).

- There is no experimental evidence whicii suggests that it is widely

applicaile to hifflilj" anisotropic molecales oTer wide intervals of temperature and fluidity.

fairly long distances to meet each other, cannot

be compared with the reactivity of model compounds in a medium of same viscosity siace the

diffusibility of the smaU molecule is always much

larger than that of the poh-meric species. This is actuaUy the reason, w-hy under identical reaction

conditions and this includes identical viscosity a polymer will suffer more severe peroxidation than a low molecular weight model compound: the lifetime

of the free radicals and hence the chain length being

much longer in the former svstem.

C. R. Patrick ITniversity of Birmingham. England

It has been suggested by Dr. Chapiro that in some

degrading svstems in melts the reactive free radical

centres might be displaced more extensively as a consequence of chemical reactions than by physical

diffusion. It has been acknowledged in my paper

that under extreme conditions of low molecular

mobihty such displacements by reaction may

proceed at a rate comparable w-ith that of chsplacement by physical diffusion. Such conditions may

readily prevail in reactions in rigid, partly crystalline or glassy polvmeric matrices, or. indeed, in rubbery

matrices at low temperatures. It remains my opinion

that whilst such conditions may occur in polymeric

melts, and that the possibhity of their existence

should not be overlooked, they will onlv infrequently

justify the neglect of physical diffusion.

Anv chemical reaction undergone by a free radical,

be it a transfer reaction, of the type

R -R"X—R -R"

or a degradation step, in which the radical decomposes by the elimination of an olefin, is accompanied

82
by a displacement of the reactive free radical centre.

The length of the displacement in the first type of

reaction must be greater than length of a chemical

bond, but not by a large amount. A reasonable

estimate would seem to be about (3 to 5) XlO~^ cm,

but most probably would lie close to the lower limit.

In the second type of reaction the main chain of the

free radical produced will be shorter than its progenitor by two carbon atoms. The displacement of the

reactive centre in the second case is likely to be a

little larger than in the first, and is probably about

5 X 10~^ cm, but is unlikely to exceed about 8 X 10~* cm.These displacements will take place on the average

in an interval of time tcbem, given by fchem= (^'/[M])"^

for transfer reactions, and by {tchem) av = kj^ for

elimination reactions, where kf is the rate coefficient

for the transfer reactions, [M] is the molar concentration of centres at which the transfer reactions

may take place, and kd is the rate coefficient for the

elimination of the olefin from the radical. Formally

account should be taken of the possibility that the

radical may undergo different reactions of either

type. In most degrading systems of interest the rate

of formation of olefin is likely to exceed that of the

transfer reactions, although it is evident from the

nature of the products formed in the thermal

degradation of some polymers that the transfer

reactions may proceed at rates that are significant

compared with the elimination reactions.

As stated in the paper, it is unlikely that rate

coefficients for elimination of olefins from radicals

will exceed 10^* exp[-25000/RT] s-i (R = 1.9873

cal deg~^ mole~^) . Thus even at 450 °C, kj is unlikely

to exceed 3X10^ s~^. It would seem that in many

polymeric melts in which degradation takes place,

the diffusivities, even of the larger radicals, would

be greater than above 10~^ cm^ s~^ In a time fchem =

^•7^ = 3X10~' s., the root mean square diffusive

displacement ( = [6Dfchem]''") is about 5X10~* cm.

This is to be compared with estimate of about

5XlO~* cm. for the displacement of the reactive

centre resulting from the chemical reaction. At lower

temperature for example 300°, k^ is unlikely to

exceed 5X10* s~'. In the case of polystyrene at this

temperature, D is about 10"'" cm^ for molecular

weights of interest. The r.m.s. displacement resulting

from translational diffusion during the average time

for one chemical act of eliminating an olefin is therefore IIXIO"^ cm., and larger than the displacement

of the radical centre as a result of the chemical

process.

We conclude that in general displacements due to diffusion occur at a rate that is at least comparable

with the rate of displacement by chemical processes,

and diffusion will normally have an important, if not

dominant role, in determining the rate of encounter,

and therefore of reaction, between macroradicals in

polymer melts. It should be noted that this view is reinforced by the fact that in our calculations we

have considered only the least favourable class of

radical-radical reactions, namely by the encounter

of two large radicals, rather than by, for example a

large radical and a small radical, and that we have

used as estimated for the diffusivities values that

relate to the diffusion of centres of mass, which may

be smaller than those relating to the diffusion of

segments.

83

NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Oxidative Degradation of Polymers and Organic Compounds Via

Unimolecular Decomposition of Peroxy Radicals

Jean Marchal

Centre de Recherches sur les Macromolecules—C.N.R.S., 6 rue Boussingault, 67-Strasbourg,

France

Autoxidation of various polymers and organic compounds is shown to include, already at room temperature, chain scissions involving unimolecular decomposition of secondary and tertiary

peroxy radicals. Rate and product data are consistent with a scheme including a six-membered

ring transition state leading to three fragments, one of them being an alkoxy radical. This oxidative

degradation scheme at room temperature does not correspond to a known mechanism.

Key words: 1,4-dioxane; kinetics; mechanism; organic compound; oxidative degradation;

peroxy radical isomerization; poly (ethylene oxide); polymer; polypropylene; poly (propylene

oxide); polystyrene; polytetrahydrofuran ; room temperature.

1. Introduction

Experiments in the very early stages of autoxidation of various polymers and organic compounds

were done at room temperature by using 7-irradiation for initiation. Whether the radiation -induced

reactions take place in dilute solution or on the pure

compound, the results show that chain scissions occur without any induction period. This oxidative

degradation is described by means of a unimolecular

decomposition process of secondary and tertiary

peroxy radicals leading to three fragments, schematically written:

A - B

/ \

C C -

\ 0-0*

A = B

- c c

W / 0 "0

The alkoxy radicals thus formed are responsible for

propagation and termination of chain reactions of

autoxidegradation. Intramolecular propagation leads

to small fragments. A=B can be an ethylenic compound, in which case it works as an efficient radical

scavenger.

2. Radiation-induced Autoxidation by

7-Irradiation of Dilute Solutions

To study the reactions involved in the autoxidation

of organic compounds, it is very important to ensure

that the process used for initiation does not also

induce decomposition of the reaction products.

Many results obtained by using conventional

initiation processes such as thermal, uv and redox

free radical generators are questionable in this

respect.

This difficulty is overcome when 7-rays are used

to produce active species for the initiation of the

autoxidation. Since the amount of energy which is dissipated by collisions leading to ionization by the

Compton effect is proportional to the electron

density of the irradiated medium, it is clear that the

relative amount of this energy absorbed by each

constituent is proportional to its electron density

and weight fraction. The result is that a relatively

large accumulation of the products of the autoxidation is possible before the limiting effect of their

decomposition by 7-irradiation becomes appreciable.

In fact, this direct effect is still negligible when

secondary reactions of the products of the autoxidation with free radical initiators become noticeable.

For the same reason, it is possible to irradiate a mixture in which the dilution of an organic solute is such that the yield of free radicals resulting from its radiolysis by the direct effect of 7-irradiation can be

neglected compared to that resulting from the

indirect effect due to the H atom abstraction from

the solute by active species produced by the solvent

radiolysis. As a very important consequence, the

solute does not undergo any chain scission which

would be due to the direct effect of 7-irradiation.

A good illustration of the use of those properties

in oxidation reactions is given by the hundreds of

investigations devoted to the quantitative determination of the yield of free radicals and "molecular"

products which is basic to the understanding of the

water radiolysis. To determine the radiolytic yield of

reducing radicals, one way, reviewed by Baxendale

[Ij,^ is to use oxygen as a radical scavenger and to

determine the radiation yield G(H202) of hydrogen

* Figures in brackets indicate the literature references at the end of this paper.

85
peroxide. Another way, reviewed by Bielski and

Allen [2], is to use the radical scavenger made of

oxygen +organic additive such as ethanol or ethylene

glycol (Ahmad et al. [3]) and to determine the

total peroxide yield.

Added to the possibility of working in ordinary

vessels, at any temperature and with very well

determined absorbed dose rates, those properties

make the radiation-induced initiation by 7-irradiation a very versatile and powerful technique to

study autoxidation reactions, mainly in dilute

solutions.

3. Oxidative Degradation

Measurements in

Dilute Solution

The experiments reviewed here are done in our

laboratory with polymers in dilute solution by using

radiocobalt 7-rays to induce oxidative processes at 25 °C. The oxidative degradation occuring during

autoxidation has been determined by viscometry

and gas phase chromatography.

Viscometric properties of polymer coils are not

modified by hydroperoxide formation on the macromolecules since it does not change their overall

dimensions. On the other hand, they are very

sensitive to random chain scissions which are produced by the oxidative degradation. The quantitative determination of the yield of random chain

scissions is possible from the experimental value of

the limiting viscosity number [r;], when the initial

molecular weight distribution of the polymer sample

and the Mark-Houwink law [_7]'] = kMa in the same

solvent are known. We used the Saito [4]-Kotliar

[_5^ method. Irradiation cells were glass tubes topped

by a capillary viscometer (fig. 1).

Viscometric properties of polymer coils are not

significantly modified by fragmentation when the

oxidative degradation concerns the end groups of

macromolecules or by minute fragmentation when

intramolecular propagation follows random chain

scissions. Gas phase chromatography has been used

to identify and to determine the yields of the low

molecular weight products of those scissions.

The yield of hydroperoxides was determined in

irradiated aqueous solution by using titanic ions in

sulfuric acid medium as described by Pobiner [6].

4. Autoxidative Degradation of

Poly (ethylene oxide

)

Experiments are carried out on a standard polymer

sample Hoechst "Polyaethylenglykol 15,000"

Mu, = 15,000=t500: light scattering measurement;

M„ = 12,500 ±500: dosage of alcohol end groups.

Autoxidation of the polymer is radiation -induced in

aqueous or chloroform solutions. Absorbed doses are determined by taking the standard G(Fe^+) =

15.5/100 eV for the Fricke dosimeter in 0,8N H2SO4.

Good oxygenation conditions of the irradiated

solutions exist when the rate of the autoxidation is

Figure 1. Irradiation cell (viscometer).

not dependent on the liquid level in the cells, the

oxygen pressure and absorbed dose rate being

constant. They were determined in experimental

tubes as follows, for the ratio free surface/liquid

volume >1 cm~':

I<400 rads*min~^ for concentration l<c<4

g-100 cc~^ in chloroform solutions saturated with

oxygen in open air,

I<50 rads-min"^ for concentration 2<c<6 g-lOO

cc^^ in aqueous solution saturated with oxygen

under P(02) = 2 kg-cm"^.

4.1. Random Chain Scissions

Kinetic results obtained in chloroform solution by

Gugumus [7J and in aqueous solution versus pH by

Crouzet [8'] have been reviewed in a first paper [_9j.

Viscometric measurements show that chain scissions are formed without any induction period. By

using the Saito-Kotliard method, it has been possible

to demonstrate that the rate of formation of chain

scissions is not dependent on the initial M„ value of

the sample. This is the proof that observed chain

scissions are randomly distributed along polymer

chains. Knowing the initial molecular weight distribution function of the samples, which is very

close to the Schulz-Zimm one, the average number x

of random chain scissions per macromolecule is deduced from the value of the ratio [jjJ/C'jDo of the

limiting viscosity number after and before irradiation.

Kinetic results obtained in each solvent show that

the average number x of random chain scissions per

86
Figure 2. Random chain scissions in aqueous

solution.

c = 2g.l00cc~', 1 = 36 rads.mn"', natural pH~6.

I rads/mm

10 1—

. I I I '

I i l 1—>—.-I » 2 5 10 20 50

Figure 3. Experiments in water.

macromolecule is proportional to the absorbed dose

in the early stages of radiation-induced autoxidation

at 25 °C. An example is given in figure 2. One can see on this figure that the effects of secondary reactions

responsible for the increase of the rate of formation

of random chain scissions begin to be noticeable

when the effects of the primary random chain

scission process are already relatively important.

One can also see from figures 3 and 4 that the initial

rate of formation of random chain scissions per

initial macromolecule obeys a law Vx = kl^>^-^ in each

solvent, € decreasing as the polymer concentration

increases.

The initial rate of formation of random chain

scissions per unit volume of solution Fscissions is given

by the relation:

V' scissions =r..c.^ r X ^

-jyj

where is Avogadro's number.

4.2. Related Chemical Groups

The related chemical groups are formate ester or primary alcohol at the end of the fragments and

I . , , I , , , ,1 I. . 10 50 100 200 400

Figure 4. Experiments in chloroform.

formaldehyde. In dry chloroform solution, Gondet

et al. [lOJ established that the very fast side reaction

of formaldehyde with alcohol end groups of initial

macromolecules explains the accumulation of hemiformal groups identified by Gugumus without any

measurable induction period. In aqueous solutions,

formate esters tend to be hydrolysed.

4.3. Minute Fragmentation

In both solvents, the initial rates of formation of

those chemical groups including formaldehyde are higher than the initial rate of random chain scissions

determined by viscometry. This result has been

explained by Decker [11] who, using gas phase

chromatography, identified 9 low molecular weight

compounds which are also produced without any

measurable induction period. These are small

fragments produced by intramolecular propagation

and termination of the autoxidative degradation of

the polymer. They are: formic acid, ethylene glycol,

its formate esters and aldehyde, diethylene glycol,

its formate esters and aldehyde.

Since gas phase chromatography is used to analyse

the oxidation products, it has been necessary to

verify that none of these products are formed by

thermal decomposition of peroxides in the apparatus.

This has been done by two ways. The first is that the

increase of the temperature of the injection chamber

from 120 °C to 250 °C is followed neither by a change of the yields of the detected products or by

the detection of new compounds. The second is that

during freeze-drying of the polymer from the

irradiated aqueous solutions, the low molecular

weight products are removed and recovered with

the water. Furthermore no signal appears on the

chromatogram of solutions made with the dried

oxidized polymer in fresh water.

These results, their kinetic treatments and

mechanisms have been recently detailed [11].

4.4. Unimolecular Decomposition of

a-Alkoxyalkylperoxy Radicals

It has been pointed out in the first part of this

paper that the homolysis of peroxides by direct

87
effect of 7-irradiation must be negligible in the early

stage of the radiation-induced autoxidation because

their concentration is too low.

Furthermore, the increase in the yield of chain

scissions, which is due to the post-effect initiated by

the thermolysis of peroxides at 25 °C, remains very

low when, after irradiation, the solutions are kept at

this temperature in the dark. This result confirms

that the spontaneous peroxide decomposition is not

able to explain the initial rate of formation of chain

scissions during the autoxidation of the polymer

at 25 °C.

Our rate and product data on several polymers are consistent with a imimolecular decomposition of

alkylperoxy radicals PO2 such as:

/

.0 - CH

CH^ - 0

\

— CH2^

0-0'

= 0 ~ 0 - CH

0 '0

CH^ - CH^-"

(PO2) (P0-)

According to this scheme which, admittedly, does

not correspond to a known mechanism, the primary

alkoxy radicals PiO would be responsible for

propagation of the autoxidative degradation of

the polymer.

Consequently, bimolecular reactions of two PO2

radicals would be less probable. Therefore we have

constructed a scheme for the autoxidegradation of

the polymer at 25 °C where termination reactions

are assumed to occur by bimolecular reactions of

PiO radicals.

In aqueous solutions, PO2 and PiO- radicals being

unable to react with the solvent, the scheme of the

autoxidative degradation of the polymer is simpler.

Initiation*

HjO^ OH- Fi = K(GoH-I)

Transfer and Propagation

OH +PH^H20+P- (fc2) P-+02-^P02 (fe) POj +PH^POjH+P (&4) P02^PiO +P2-O-CHO+HCHO

(unimolecular decomposition)

PiO -PH^PiOH-P (fee)

Termination**

2PiO-^PiOOPi (negligible) (&7) —>Pi =0+PiOH (disproportionation) (fc8)

' Crouzet [8] found that initiation by O2 radicals is not possible because

they do not react with the polymer. This was already found by Bakh [14] for ethers.

** According to our scheme, PiO' are primary alkoxy radicals: their recombination (reaction 7) must be negligible.

4.5. Kinetic Equation for Random Chain

Scissions

From the above theoretical scheme of free-radical

chain reactions where intramolecular propagation

reactions are disregarded, the rate of the formation

of random chain scissions is:

V' scissions

= aKI+i3(KI)i/2

al corresponds to termination reactions, the rate of

which is equal to the rate of initiation under steady

state conditions. /3(KI)^'^ is the rate of propagation

by PiO- radicals.

(a) ot Values

In aqueous solution as well as in chloroform

solution, this kinetic equation for random chain

scissions is in good agreement with experimental

results which lead to a values independent of the

polymer concentration c:

H2O: a = 2.8d=0.2 2<c<6g-100cc-i

Crouzet [8]

CHCI3: a = 6.2±l l<c<4g-100cc-i

Gugumus [7]

Within experimental errors, a values are close to

GoH. of water radiolysis and to Gr. for chloroform

radiolysis.

H2O: GoH- = 2.74±0.08

Bielsky and Allen [2]

CHCI3: Gr=7±2

see [15] reviewed by Gugumus [7]

These results are consistent with kT^h^, meaning

that the termination of autoxidegradation of the

polymer is done by disproportionation of the primary

alkoxy radicals PjO . As was already mentioned

Decker [11] found that low molecular weight PJOradicals are formed by intramolecular propagation of

the autoxidegradation of the polymer. These alkoxy

radicals P^O are: HO-CH2-O , HO-CHo-CHs-Oand HO-CH2-CH0-O-CH2-CH2-O . From the study

of the autoxidegradation of the first homologous

88
polymers of the series, di- and tri-ethylene oxide

a,w-glycols. Decker [11] found that the smaller

the molecular size of these alkoxy radicals, the higher

is their efficiency for termination reactions. This

may be understood in terms of mobility of active

species. The efficiency of HO-CH2-O in disproportionation reaction leading to formic acid is specially high. It can be explained by the mesomeric

effect:

HO—CH2—0-^0=CH—OH

H

These results also support the assumption that

termination reaction without chain scission which

may involve 2PO2 radicals or PO2 +HO2 (or RO2 in

the case of chloroform solutions) are negligible. If

this was not so, we would have found a<GoH.

or a < Gr.

(b) 13 Values

In both solvents, experimental 13 values for random

chain scissions vary linearly with the polymer

concentration c. According to the kinetic equation

for Fscissions, this result indicates that all the radicals

capable of initiating the autoxidation of the polymer

are trapped by the solute in the range of concentrations used. This result also demonstrates the

existence of chain reaction of oxidative degradation.

4.6. Autoxidegradation as Main Chain

Reactions at 25 °C

To conclude, the main feature of this work is that

the termination of the free-radical chain reactions of

the autoxidation of poly (ethylene oxide) induced at

25 °C seems to be performed by alkoxy radicals

produced by the unimolecular decomposition of

a-alkoxyalkylperoxy radicals PO2 and not by

bimolecular reactions of these radicals because those

reactions appear to be less probable; this is contrary

to usual concepts for ether autoxidation.

According to this conclusion, the main chain

reactions are those of "autoxidegradation", hydroperoxidation acting only as side chain reactions

regenerating P macroradicals (retardation re- actions) .

PH linitiationf •

P- + POOH

hydroperoxidation

oxidative deqradation

P,0* (chain

scission'

|propagation|

Iterminationl

)> cm-'

1600 1700

1600 1700 1600

Figure 5. Irradiated solutions of poly{ethylene oxide) in chloro- form, 50 and 96 hours, I = 125 rads.mn^^.

We are pursuing this work in order to confirm and

to describe the unimolecular decomposition of PO2

radicals which has been postulated here.

Recently Decker [12] got the proof that the

oxygen atom in formaldehyde is coming from the

polymer chain and not from oxygen gas. This result

which demonstrates that formaldehyde is produced

by fragmentation of the polymer chains has been

obtained by using almost pure ^^02 gas instead of

'^02 to perform the radiation -induced autoxidation

at 25 °C. Experiments were done in dry chloroform

solutions where isotopic exchange reactions are

negligible. The results were followed by ir spectroscopy on the irradiated solutions. The upper

part of figure 5 corresponds to experiments in the

presence of ^^02. It shows that two absorption bands

are present in the region of vc=o groups. j'~1755 cm~^

has been attributed [7] to ester groups formed in the

polymer chain —O—C— and 1/^1725 cm~^ to the

O

remaining free formaldehyde and to formate esters. If all carbonyl groups formed during autoxidation in

the presence of ^^02 have been labelled with ^^O, the

result would have been a shift of — 30 cm~^ of the

89
entire ir spectrum without any noticeable change of

its shape. The lower part of figure 5 which corre- sponds to the same experiments in the presence of

'^©2 shows that this is not the case, indicating clearly

that one kind of carbonyl groups is not labelled with

^*0. The disappearance of the previously v^llSS

cm^' band is explained by the shift of -30 cm

of vc^o of esters in the chain. The appearance of a band i':::^1695 cm~^ is attributed to the shift of vc=o

of formate esters. Thus f'~1725 cm~^ corresponds to the absorption of labelled esters in the chains and to

nonlabelled formaldehyde. This has been confirmed

by the fact that the addition of a weak amount of

water to irradiated solutions in the presence of

'^02 or ^^02 was enough to induce the decrease of the

absorption j'c=o— 1725 cm~^ because the hydration

reaction of formaldehyde is very fast without

noticeable hydrolysis of formate esters. Since formaldehyde is produced by fragmentation

of polymer chains, it might also be possible to

interpret its formation as the result of the decomposition of a-alkoxy-alkylalkoxy PO radicals:

-O-CH-CH„-0-CH, - 0*

•CH. -0-CH + •CH„-0-CH„-CH

II ^ A z 0

? 25°C

CH^=0 + •CH2-CH2'^

To be kinetically valid, this hypothesis requires

that the PO radicals be produced without any

induction period. Considering the conventional

concepts of autoxidation, this can only be achieved

by the bimolecular reaction that we have so far

neglected:

2P02^2PO+02

This hypothesis also requires that these POradicals disappear only by unimolecular decomposition and that the ^CH2-CH2-0-CH2 radicals

thus formed be in turn too unstable, at 25 °C in

solution satured with oxygen, to lead to any reaction

other than their decomposition producing formaldehyde.

This research is actually undertaken with polymer

samples containing hydroperoxide groups. They are prepared by recovering from aqueous irradiated

solutions by freeze-driing in order to eliminate all low molecular weight products. Then, hydroperoxide

groups are decomposed, in fresh dilute solutions, by

different ways: thermolysis, uv radiation and redox

systems.

Another way to get information is to use ESR

spectroscopy to observe free radicals produced in a liquid flow. We are investigating this technique in

order to obtain direct informations concerning

kinetics and lifetimes of PO2 and alkoxy radicals.

The determination of the orders of the disappearance

rates of these species would provide conclusive

evidence.

Finally, by using the same radiation-induced

technique, we are extending this work, by searching

for other compounds undergoing oxidative degradation, at 25 °C, which can be interpreted as we did

for poly (ethylene oxide). Until now, we found that

this is the case for o-dioxane, poly (propylene oxide),

polytetrahydrofuran, polypropylene and polystyrene.

The experiments mentioned below were done in

chloroform solution and also in the pure state for

p-dioxane.

5. p-Dioxane

For this cyclic dimer of ethylene oxide, autoxidegradation leads to formaldehyde and ethylene glycol

monoformate ester with similar yields [16J. By

using ^^02, Decker [13] reached the same conclusion

concerning the formation of formaldehyde as for

linear poly (ethylene oxide). One can see in figure 6

that autoxidation of p-dioxane in the presence of

^^02 leads to only one vc=o absorption band ~1725

cm~^ and that this absorption band is split into two

bands ~1725 cm~^ and ~1695 cm~^ when autoxidation has been induced in the presence of ^^02. It was

thus easier to interpret this result: the remaining

absorption band ~1725 cm~^ corresponds to unlabelled formaldehyde and the other absorption to

the ^^O labelled carbonyl group of formate esters.

>) cm-^

% 1800 1700 160O

1800 1700 160O

Figure 6. Irradiated p-dioxane, 87 and 159

hours, I = 125 rads-mn"^.

90
Furthermore, this result is qualitatively the same

whether the radiation -induced experiments are done

in dilute chloroform solution or with the pure

compound. This is the proof that the observed

degradation is not due to the direct effect of 7 rays on peroxides, which confirms the general statement

given in the general introduction.

6. Poly (propylene oxide

)

Radiation-induced autoxidation in dilute chloroform solution has been studied on a regular head to

tail but atactic polymer sample (Union Carbide

CH^ - 0 CH^

/ ^ \ / ~0 - C CH - CH^^ -

y \ ^ CH^ 0-0*

1 CH -

/

^0 - CH

\

0 -

poly (propylene oxide)a,w-glycol NIAX 4025, M„ =

3,200).

Gondet \Yl~\ found that all conclusions concerning

poly (ethylene oxide) are valid for this homologous

polymer. But the number of low molecular weight

compounds produced by intramolecular propagation

of the autoxidative degradation is higher because

monomeric units contain an asymetric carbon atom.

The postulated intramolecular decomposition of

PO2 radicals is able to explain the yields of formaldehyde and acetate ester groups and the yields of

acetaldehyde and formate ester groups, depending

whether the PO2 are secondary or tertiary peroxy

radicals:

CH„=0 CH^

z /J

'-^ 0 - C CH -CH^

CH^ 0 -0

CH^ -

CH,

1CHCH_

I ^ CH=0

0 - CH

CH,

CH^ - CH

0*

The yield of formaldehyde is higher than the

yield of acetaldehyde as expected from usual properties of H atom abstraction by free radicals.

7. Polytetrahydrofuran

A polymer sample, Mu,~30,000, has been prepared

by SbCU cationic-induced polymerization in dry

diethylether at low temperature. Thomassin [18j

found that ethylene is formed without any induction

period during the radiation-induced autoxidation of

the polymer in chloroform solutions at 25 °C. This

result may be described as follows:

CH~ — CH„

0 - CH

\

CH^ - 0-^

0-0*

CH2 ~ CH2

--^0 - CH CH^ - 0-

0 •0

In this case, it is interesting to point out that the

low molecular weight product of the unimolecular

decomposition of PO2 is an ethylenic compound

which can work as an efficient free radical scavenger.

8. Polypropylene

Thomassin [18] also found that propylene is formed during the radiation-induced autoxidation of an atactic polymer-sample in chloroform solution at 25 °C. The same oxidative degradation process has

been proposed to explain this result:

CH,

CH,, - CH CH,

CH - C

y \

CH3 0

CH^ - CH

0*

CH,

CH, CH CH,

CH„ - C

^ / W

CH3 0

CH„- CH^

•0

91
9. Polystyrene

Like the above listed polymers, polystyrene

undergoes oxidative degradation during radiationinduced oxidation in dilute chloroform solution at 25 °C. Samples used have been obtained by anionic

CH^ - CH 0

'CH„ - C CH^ - CH0 0"

polymerization. Abadie [19] determined the initial

rate of random chain scissions by using the viscometric method of Kotliar-Saito and found that

styrene is produced with the same initial rate. Those

results support the assumed process:

0

CH^ = CH

CH„ - C

I CH^ - CH^

(6 0

/

•0

Styrene works as a very efficient radical scavenger.

The main differences compared with the behavior

of poly (ethylene oxide) are:

• The oxidative degradation is not a chain reaction:

J^scissions is proportional to the absorbed dose and

does not depend on the polymer concentration

3<c<8g-100cc-i.

• Consequently there are no small fragments which

would be due to intramolecular propagation.

Detailed results will soon be published [20].

10. References

[1] J. H. Baxendale, Radiation Res. Suppl., 4, 114 (1964).

[2] B. H. J. Bielski and A. O. Allen, Int. J. Radiation Phys.

Chem., 1, 153 (1969).

[3] M. Ahmad, M. H. Awan and D. Mohammad, J. Chem.

Soc, B, 9, 945 (1968).

[4] O. Saito, J. Phys. Soc. Japan, 13, 198 (1958).

[5] A. N. KotUar and A. D. Anderson, J. Polym. Sci., 45,

541 (1960).

A. N. KotHar, J. Polym. Sci., 61, suppl 25 (1962).

[6] H. Pobiner, Analytical Chem., 33, 1423 (1961).

[7] F. Guguraus, D.Sc, Strasbourg, Oct. 1965.

F. Gugumus and J. Marchal, J. Polym. Sci., Part C, 16,

3963 (1968).

[8] C. Crouzet, D.Sc, Strasbourg, Dec. 1966.

C. Crouzet and J. Marchal, Kinetics and Mechanism of

Polyreactions, preprints 5, 259 (1969), Budapest.

[9] J. Marchal, Radiation Research 1966 (Ed. by Silini,

North Holland, Amsterdam 1967), p. 327.

[10] J-C Gondet, C. Crouzet and J. Marchal, Kinetics and

Mechanism of Polyreactions, preprints 5, 249 (1969),

Budapest.

[11] C. Decker, D.Sc, Strasbourg, June 1967.

C. Decker and J. Marchal, Kinetics and Mechanism of

Polyreactions, preprints 5, 231 (1969), Budapest.

[12] C. Decker and J. Marchal, C.R.Ac.Sci. Paris, serie C,

270, 990 (1970).

[13] C. Decker and J. Marchal, C.R.Ac.Sci. Paris, serie C,

270, 1102 (1970).

[14] N. A. Bakh, Radiation Res. Suppl., 1, 190 (1959).

[15] W. H. Seitzer and A. V. Tobolsky, J. Am. Chem. Soc,

77, 2687 (1955). E. A. Cherniak, E. Collinson, F. S. Dainton, G. M.

Meaburn and C. D. Walker, Proc. 2nd Int. Conf.

Peaceful Uses of Atomic Energy, 29, 121 (1958),

Geneva.

R. J. Hanrahan, quoted by F. S. Dainton, id. p. 127. Reviewed by F. Gugumus and J. Marchal (7). [16] M. Galin-Vacherot, Bull. Soc. Chim. France, 765 (1969).

[17] J-C. Gondet, D. Ing., Strasbourg, June 1969. J-C. Gondet and J. Marchal, Kinetics and Mechanism of

Polyreactions, preprints 5, 255 (1969), Budapest.

[18] C. Thomassin and J. Marchal, Kinetics and Mechanism

of Polyreactions, preprints 5, 245 (1969), Budapest.

[19] M. Abadie and J. Marchal, C.R.Ac.Sci. Paris, serie C,

266, 520 (1968).

M. Abadie and J. Marchal, Kinetics and Mechanism of

Polyreactions, preprints 5, 239 (1969), Budapest.

[20] M. Abadie and J. Marchal, Makromolekulare Chemie,

141, 299 (1971).

Discussion

C. F. CuUis (The City University, London)

One of Dr. Wall's objects in arranging the present

Symposium was, I think, to seek "bridges" between

those working on systems in which only simple welldefined chemical entities are involved and those

studying the reactions of more complex species such

as polymers. In this connection. Dr. IVfarchal's work

underlining the importance of the imimolecular

decomposition and rearrangement of peroxy radicals

during the oxidation of polymers at room temperature finds some parallel in studies of the gaseous

oxidation of simple hydrocarbons and related compoimds at considerably higher temperatures. Thus,

it has been possible to acquire in recent years

[A. Fish, Quart. Rev. Chem. Soc. 18, 243 (1964)]

a greatly increased tmderstanding of the modes of

gaseous decomposition of alkyl peroxy radicals as well as a realisation of the importance of rearrangement reactions some of which are at least formally

analogous to those suggested by Dr. IVIarchal.

However, it is fairly generally agreed that, during

the combustion of hydrocarbons, unimolecular

reactions of alkylperoxy radicals become important

relative to bimolecular reactions only at rather high

temperatures. Although, from other than energetic

points of view, the conditions existing in the con- densed phase will be more conducive to the occur92
rence of the breakdown and rearrangement of

peroxy radicals, it is perhaps a httle surprising that

the favourable entropy factors should be able to

outweigh the unfavourable energy factors and thus

enable unimolecular decomposition of such radicals

to take place quite extensively at room temperature.

J. Marchal:

The comment of Professor CuUis opens the discussion by raising the question that I was anticipating.

First, may I outline some points concerning the

aim of my paper to this symposium.

-The object of this paper has been very precisely

to report experimental results, for polymers and low

molecular weight compounds, which are inconsistent

with the usual assumption that unimolecular decomposition reactions of peroxy radicals are energetically

difficult and are thought not to occur at room

temperature.

-The fundamental question of the proposal of a mechanism which would be able to justify such a scheme at room temperature was not treated in this

paper because the necessary theoretical information

on peroxy radical chemistry has not yet been

developed. It was thought that reference only to

results obtained for hydrocarbon oxidation at high

temperature would not be helpful.

I should add that we have been working on this

basic question from a theoretical point of view. I am

pleased to have the opportunity to reply to Professor

CuUis before the printing of the proceedings of this

symposium by referring to the beginning of the

publication of a series of papers which propose a way

to reach an answer. In the first paper [1] of this series, a three step

mechanism is proposed to interpret the unimolecular

decomposition scheme of peroxy radicals at room

temperature.

-The first step is a peroxy radical isomerization

reaction leading to the introduction of a peroxide

linkage in the molecular backbone. This reaction is based upon charge transfer interaction:

A—

/ \

C I C

^0—0*

A—B*

/cue

—The second step is the radical-induced decomposition of the peroxidic bond:

^A—B*

•C II C

A—B' /

C III

'0

A—B" A=B

-The third step is the homolytic decomposition

of AC bond of the diradical III:

This three step mechanism is detailed in a communication [2] at the lUPAC conference on Chemical

Transformations of Polymers in Bratislava (to be

published)

References

[1] J. Marchal and L. Valko, C.R.Ac.Sci. Paris, 272 Serie C,

2042 (1971).

[2] L. Valko, V. Laurinc and J. Marchal, Chemical Transformations of Polymers, Bratislava, June 1971, Preprints, III, No. 64.

R. E, Florin (National Bureau of Standards)

In Dr. Marchal's paper, the number of identified

products corresponding to the mechanism is impressive. Regarding the program for observing radical

intermediates, the ESR spectrum of radical P has

been well characterized in flow mixing experiments on TiCls, H2O2 and polyoxyethylene in aqueous

solution [1]. Another radical spectrum of more

ambiguous origin is observed as well, perhaps

^OCHsCHOH or ['-0CH-CH2]+ having g =

2.0024, doublet a = 18.7 Oe, triplet a = 22.2 Oe. The

peroxy radical PO2 spectra are well known, with high

g and no hyperfine structure. The alkoxy radical g

value and hyperfine splitting will be hard to predict.

Although OH- [2], CHaCHsO- [3], and CeHsO' [4]

ESR spectra are reported, the former two are broadlined, anisotropic polycrystalline or glassy spectra

with little detail and doubtful identification, while

the latter, though exact, may be irrelevant because

of extensive electron delocalization.

References

[1] R. E. Florin, F. SiciUo and L. A. Wall, J. Res. NBS 72A,

49 (1968).

[2] S. Siegel, J. M. Flournoy, and L. A. Baum, J. Chem.

Phys. 34, 1782 (1961).

[3] P. J. SulUvan and W. S. Koski, J. Am. Chem. Soc. 86,

159 (1964).

[4] W. T. Dixon and R. O. C. Norman, J. Chem. Soc. 1964,

4857.

J. Marchal:

Dr. Florin refered to ESR measurements of free

radicals produced on poly(ethylene oxide) in flow

mixing experiments. May I give our results on the

same subject

The second radical spectrum described by Dr.

Florin is observed in acidic but not in neutral media

(the reagent being neutralized by NH4OH and the

Ti*+ ions being complexed with EDTA). The spectrum of the radical P" produced by random H atom

abstraction along the polymer chain ^0-CHCH2-O'—' is then the only one which is observed.

93
The origin of the second radical spectrum observed in acidic medium is neither of the two species

mentioned by Dr. Florin, but is '^0-CH0H-CH2.

The formation of this radical, in acid catalysed

reactions, can be rationalized on the basis of the decomposition of radical P' and the isomerization of

the radical '^0-CH-CH20H thus formed. It is proposed that the isomerization reaction proceeds

in the manner described for corresponding radicals

of vinyl ethers [2]:

--CHa-CHs-O -CH2-CH-0~

H3O+

>-'CH2-CH2-OH+HO-CH,-CH-0'-

H3O+

^OC-H-CH2-OH^=:i~0-CHOH-CH2

References

[1] M. A. Herrmann, C. Crouzet, and J. Marchal, to be published.

[2] D. J. Edge, B. C. GUbert, R. O. C. Norman, and P. R. West, J. Chem. Soc, B, 189 (1971).

94
NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972)

Mechanisms of Oxidation of Polyolefins below 150 °C

Frank R. Mayo

Stanford Research Institute, Menlo Park, Calif. 94025

This paper applies what we have learned ahout oxidations of simple aUphatic hydrocarbons

below 150 °C to oxidations of bulk polyolefins. The principal difference between these groups is due to the production of initiating radicals in pairs and to the greatly restricted mobilities of all radicals. Data are given for the efficiency of initiation by di-f-butylperoxy oxalate of the oxidation

of atactic polypropylene at 45 °C and for the effect of peroxide concentration on efficiency and

kinetic chain length. The following other topics are discussed: chain propagation and geminate

and statistical termination by alkylperoxy radicals, nonterminating reactions of two tertiary

alkoxy radicals, formation and reactions of alkoxy radicals, sensitivity of oxidation products to

further oxidation, intramolecular propagation of alkylperoxy radicals, and the question of the

formation of carbonyl groups in oxidations of polyolefins.

Key words: Alkoxy radicals; carbonyl compounds; chain initiation; chain propagation; chain

termination; kinetics; oxidation; peroxy radicals; polyolefins; polypropylene.

1. Introduction

Whether we wish to prevent oxidation of polymers,

to accelerate their destruction, or to control the

introduction of functional groups by oxidation,

better understanding of the mechanisms seems

essential for efficient progress. From our experience

with liquid-phase oxidations of simple hydrocarbons,

the competing reactions that determine the rates and

products of oxidation of polyolefins below 150 °C

have been deduced. The difficulties in measuring

these competitions, and the failure of nearly all published work to meet the conditions necessary to

obtain basic information are discussed. These discussions assume that polyolefin radicals are almost

always generated close together in pairs, that their

mobility is so restricted in bulk polymers that

efficiencies of initiation may approach zero, and that

for every pair of effective initiating radicals, there

must be one, and only one, pair of terminating

groups. While light and ionizing radiation may

cleave chains directly, in oxidations of saturated

aliphatic hydrocarbons the diiference between oxidations in the light and in the dark is assumed to

depend on the ways in which pairs of radicals are generated, not on what happens to them subsequently.

2. Efficiency of Production of Polymer

Radicals

We have used di-f-butylperoxy oxalate (DBPO)

to initiate oxidations of atactic polypropylene [1^

at 45°. This initiator has a convenient half life and

• Figures in brackets indicate the literature references at the end of this paper.

after several half lives the t-butoxy radicals can be

recovered quantitatively and the number of chains

initiated determined. In one experiment with 0.615

mole DBPO/kg atactic polypropylene at 45°, 73.5

percent of the total potential i-butoxy radicals were

recovered as di-t-butyl peroxide, 14.8 percent as t-butanol, and 11.7 percent as i-butoxy groups bound

to polymer (as t-butyl chloride after cleavage of the

ether with anhydrous hydrogen chloride). The

di-i-butyl peroxide measures the f-butoxy radicals

which were wasted in a cage combination reaction.

The 14.8 percent of f-butanol measures the total

hydrogen abstracted from the polymer (polymer

radicals produced) by t-butoxy radicals. The 11.7

percent ether measures the t-butoxy groups wasted

by near-cage reactions of i-butoxy and polymer

radicals. Thus only the f-butanol in excess of the

ether, here 3.1 percent, corresponds to polymer

radicals available for starting oxidation chains and

to the efficiency of initiation. Here, 73.5 percent of

the potential t-butoxy groups were lost in cage

combination and 88 percent of the remainder were wasted in near-cage ether formation.

The efficiency of initiation increases with increasing

temperature or with addition of small amounts of

solvents, both of which facilitate separation of caged

radicals. Knowledge of the efficiency of initiation is necessary for establishing kinetic chain lengths and

the proportions of oxidation products corresponding

to chain termination.

3. Chain Propagation and Termination

Reactions of Alkylperoxy Radicals

In the presence of sufficient oxygen, all the

substrate is brought into reaction by alkylperoxy or

95
alkoxy radicals. Propagation reactions of the

resulting alkylperoxy radicals in oxidations of alkanes

are relatively difficult. The activation energy for

propagation of f-alkylperoxy radicals with tertiary

aliphatic hydrogen atoms is about 16 kcal/mole

[2, 3], with secondary hydrogen atoms, about 17

kcal/mole. Reactions of peroxy radicals with

primary hydrogen atoms (methyl groups) are

negligible below 150°. The only termination known

for two tertiary peroxy radicals is 2R02^R202+02

for which the activation energy is about 9 kcal/mole

[3]. Competing with this termination is the nonterminating reaction

2R02^2RO +O2

with the same or a slightly higher activation energy

[4]. In the absence of peroxide decomposition by

heat, light, or transition metals, this latter reaction

is the principal source of alkoxy radicals, which are considerably more reactive than alkylperoxy radicals

and can attack methyl groups. With f-butylperoxy

radicals in liquid isobutane at 80°, about 10 nonterminating interactions occur for each terminating

interaction. The fraction of nonterminating interactions presumably decreases, as for alkoxy radicals, as the viscosity of the medium increases.

Terminations that involve one or more primary or secondary peroxy radicals are much faster than those

for two tertiary peroxy radicals,'' have an activation

energy close to that for diffusion, show a deuterium

isotope effect for a-substitution [5, 6], and give

mostly an alcohol and an aldehyde or ketone as termination products. Hiatt [4] showed that interactions of such radicals are almost always ter- minating in liquid systems below 100° and hence

give few, if any, alkoxy radicals.

4. Kinetics of Oxidation

The oxygen consumed in initiation and termination alone in the oxidation of polypropylene (PH)

is directly proportional to the number of initiating

radicals:

2i-BuO +2PH^2t-BuOH+2P

2P+202->2P02

2P02—>termination products +O2

2t-BuO +2PH+O2

-^2t-BuOH+termination products

We are now concerned with the additional oxygen

absorption, all associated with chain propagation.

When the two radicals that terminate are descendants of the same initiating pair (geminate radicals),

then this oxygen absorption is first order in rate of

initiation and radical pairs. When the two radicals

that terminate are descendants of different initiating

pairs (statistical combination), then this oxygen

absorption is one-half order in rate of initiation.

The distinction between geminate and statistical

combination is important. The former is probably

associated with short kinetic chains that may be

difficult or impossible to inhibit. The latter is probably associated with much longer kinetic chains

that may be easy to inhibit. Besides indicating the

direction of further research on inhibitors, the distinction is important for correlating initiation with

oxygen absorption and for devising and interpreting

accelerated aging tests. Table 1 summarizes some recent experiments

which illustrate our difficulties in determining the

kinetics of oxidation. The major difficulty arises

from an increase in efficiency of initiation with the

concentration of DBPO, probably due to the

plasticizing effect of DBPO and its decomposition

products. Some interesting but useless coincidences

result: The total oxygen absorption is closely proportional to the concentration of initiator but to the

one-half power of the number of initiating radicals.

However, the oxygen absorption in chain propagation

alone (last line) is not between one-half and first order as expected, but is about 0.3 order. We con- clude that the increasing concentration of DBPO

which results in increasing efficiency of initiation

also favors termination of oxidation chains more than

it favors propagation of chains. This conclusion

suggests that oxygen absorption in propagation

should be measured at various concentrations of

DBPO in mixtures adjusted with solvent so that

the efficiency of initiation is constant.

5. Reactions of Alkoxy Radicals

Alkoxy radicals are important in polymer oxidations because (except for radiation) they provide the

only free radical route to chain cleavage below

150°. At higher temperatures, or in the presence of

light or transition metals, alkoxy radicals may arise

from decompositions of peroxides, but in the dark

below 100° they are formed mostly from interactions

of two tertiary peroxy radicals. Thus, chain cleavage

in the dark at moderate temperatures depends on the

proportions of secondary and tertiary peroxy

radicals that are present, the competition between

terminating and nonterminating reactions of two

tertiary peroxy radicals, and the competition between propagation and cleavage of the resulting

alkoxy radicals. Whereas some of these ratios are known for simple alkanes, I know of no published

values for polymers.

Table 1. Oxidations of polypropylene at 45°

(Concentrations in moles/kg)

[DBPO]o 0 .317 0 ,671 1 .269

Efficiency (e), % 6 .1 10 ,8 21 .2 AP =2e[DBPO]o » 0 ,0387 0 .145 0 .538

0 .153 0 ,296 0 .563 A02/[DBPO]o 0 .48 0 .44 0 .44 A02/(AP-)i'2 0,,778 0 .777 0 .768

(A02-c[DBPO]o)/(AP-)i'' 0,,68 0,.59 0 .40

a Total P- radicale generated,

b Total O2 absorbed.

96
Some results on the oxidation of cumene are

relevant. Most of the chain termination in the

oxidation of cumene involves methylperoxy radicals,

present in very low concentrations compared with

cumylperoxy radicals, but much more reactive in

chain termination. The methylperoxy radicals come

from cleavage of cumyloxy radicals, formed from

nonterminating interactions of cumylperoxy radicals.

Participation of methylperoxy radicals was established by Traylor and Russell [7]; they found that

addition of i-butyl hydroperoxide accelerated the

ABN-initiated oxidation of cumene by converting

methylperoxy and cumyloxy radicals to f-butylperoxy radicals. At low temperatures (which discourage cleavage of cumyloxy radicals) and low

rates of oxidation (where radical interactions are

less probable), Hendry showed that methyl ter- mination is largely eliminated [8J.

In oxidations of polypropylene and poly ( 1 -butene )

,

alkylperoxy radicals presumably attack mostly the

tertiary hydrogen atoms and the resulting tertiary

peroxy radicals terminate relatively slowly. However,

any accompanying nonterminating interactions produce alkoxy radicals, and cleavage of the latter

radicals then can produce primary alkyl and alkylperoxy radicals. The latter accelerate chain ter- mination, slow the oxidation, and reduce the

nonterminating interactions that produce cleavage

and more primary radicals.

6. Implications of Other Oxidations

Among alkanes, oxidation of isobutane [3] is the

best understood. The unique characteristics of this

oxidation illustrate the serious difficulties encountered in studying oxidations of all other alkanes and

polyolefins. The principal products from isobutane

below 200°, depending on conditions, are i-butyl

hydroperoxide, f-butanol, di-t-butyl peroxide, ace- tone, and oxidation products of methyl radicals.

Except for products from methyl radicals, all of

these products are surrounded by unreactive methyl

groups and are reasonably resistant to further

oxidation. Thus, oxidation of isobutane can be carried

to 5 to 10 percent conversion without much interference from secondary oxidations. Whereas the

methyl radicals from cleavage cause some complications, f-butoxy radicals are stable enough that

effects of cleavage and the resulting methyl radicals are not serious below 50°.

However, i-butoxy radicals are unique in that they

contain only methyl groups, which are difficult to

cleave. Hence, all other alkoxy radicals cleave more

easily than t-butoxy [9, 10], and complications from

the (usually) primary cleaved alkyl radicals greatly

affect the rates and products of reaction. (Phenyl

radicals are less easily lost than methyl radicals;

hence the oxidation of cumene is also relatively

simple.

)

Whenever primary or secondary alcohols or aldehy des or ketones are primary oxidation products,

their further oxidation presents complications.

From labeling experiments in the oxidation of

cyclohexane at 150°, the relative reactivities of

cyclohexane, cyclohexanol, and cyclohexanone are found to be in the ratio 1:22:29 on a molecular

basis, 1:264:87 on a per-active-hydrogen basis

[11]. (The per-hydrogen results are probably a

little high because they neglect the reactivity of all other hydrogen atoms.) These differences should be

larger at lower temperatures. Thus, secondary

reactions of primary products may complicate

analyses as well as their interpretation, even at conversions as low at 0.1 percent. In polymers where

oxidation depends on autocatalysis or photolysis of

previous products, secondary reactions may be

further exaggerated because of high local concentrations of oxidation products that do not diffuse.

Aldehydes, formed by oxidation of primary

radicals from cleavage of alkoxy radicals, can be

particularly troublesome. Benzaldehyde is about 100

times as reactive as toluene [12] toward i-butylperoxy radicals at 30° and probably still more

reactive compared to isobutane. Moreover, the

benzoylperoxy radical appears to be 100 to 1000

times as reactive as alkylperoxy radicals in chain

propagation [13] and presumably less selective.

The aliphatic aldehydes are comparable [13].

Further, aldehydes in general seem able to initiate

oxidation chains.

As a result of all these secondary reactions, the

oxidation of any pure alkane or saturated polyolefin soon begins to involve the relatively very reactive

primary and secondary products, hydroperoxides,

alcohols, aldehydes, and ketones, and distinctions

due to differences in starting hydrocarbon will tend

to disappear earlier at lower temperatures. As a warning to chemists who are concerned with meas- urements of rate constants in oxidations of hydrocarbons, I propose the approximation: at about

1 percent conversion in the oxidation of any saturated aliphatic hydrocarbon at any chosen temperature, the apparent rate constants for chain

propagation and termination in the oxidizing mixture

will be largely independent of the starting hydrocarbon. For the reasons given above, this principle

fails for the oxidation of isobutane. It does not

apply to alkenes and alkylbenzenes, which are reactive enough to give long kinetic chains and to

protect their primary oxidation products.

7. Intramolecular Chain Propogations

The most quoted evidence for intramolecular

chain propagation in oxidation is that of Rust [14].

Oxidation of 2 ,4-dimethylpentane with 35 psi of

oxygen at 115 to 120° gave an 89 percent yield of

hydroperoxide with an estimated 90 percent of the

total oxidation product containing oxygen on both

the 2- and 4-carbon atoms. 2 ,3-Dimethylpentane

and 2 ,4-dimethylhexane gave less but still substantial disubstitution.

Chien, Vandenberg, and Jabloner [15] obtained an oxidized isotactic polypropylene from an ABNinitiated oxidation in chlorobenzene solution at

100°. They determined the hydroperoxide content.

97
and by infrared and other studies concluded that it contained no isolated, but only sequences of, hydroperoxide groups and that there was a strong tendency

toward intramolecular propagation. Although they

may eventually prove to be partly correct, their conclusions are not justified. Their kinetic chain lengths

suggest that their oxidation product may have con- tained as many terminating as hydroperoxide

groups, but they give only hydroperoxide, no total

oxygen, determinations. A simple calculation shows

that if each polymer peroxy radical has an equal

chance of propagating intramolecularly or of ter- minating, most of the hydroperoxide groups would

still be in sequences: one quarter would be isolated; one quarter would be in pairs, and one half would

be in sequences of three or more. I see no plausible

mechanism that would account for the reported

absence of isolated hydroperoxide groups.

Dulog and Kern have studied the oxidations of

isotactic and amorphous polypropylenes [16] and

poly (1-butene) s []17] in solution. Although differ- ences were found in the rates of autocatalytic

oxidations of the polypropylenes at 120 to 140° and

attributed to a greater tendency toward intramolecular propagation in the isotactic polymer, in

ABN-initiated oxidations at 70° no significant

differences were found. At 70°, the indicated kinetic

chain lengths were 1 to 3. Yields of hydroperoxide were about 30 percent on the oxygen absorbed.

Apparently no one has so far obtained kinetic

chains in oxidations of polyolefins in solution that were long enough to justify good yields of hydroperoxides in sequences.

8. Carbonyl Groups in Oxidations of

Polyolefins

Some workers have reported the presence of

carbonyl groups in oxidation products of polypropylene in the dark and some have reported the

absence of such groups. I am seeking further ex- amples. Dulog and Kern Q16]] report that carbonyl

groups were absent in the products of their autocatalytic oxidations at 120 to 140°. Chien et al. [15]

reported a similar result in their ABN-initiated

oxidations at 100°. However, Stivala and Reich [18]

have used carbonyl groups as the principal measure

of rate of oxidation of isotactic polypropylene films

at 120 to 150°. Further, Adams [19] has just re- ported finding several kinds of carbonyl groups in

the thermal oxidation of polypropylene, although

more hydroperoxide and less carbonyl were found in

oxidations after 380 hours at 63° than after 2 hours

at 138°. Our own work at 45° provides ample infrared evidence of carbonyl groups. Since these

experiments lasted only 6—7 hours in the dark, the

carbonyl groups can hardly come from the decomposition of hydroperoxides but must represent

cleavage and termination products.

I have no satisfactory explanation for these

discrepancies. The few cited results suggest that

carbonyl groups are more likely to be fovmd in

oxidations of bulk than of solution polypropylene.

I also point out that alcohol and hydroperoxide

groups, particularly when they are in 1 ,3-positions,

may obscure carbonyl absorption by hydrogen

bonding or hemiketal formation.

9. Conclusions

Essentially all of the published work on the

oxidation of polyolefins fails to provide the needed

kinds of fundamental data for one or more of the

following reasons. (1) It was carried out in solution

and hence neglects the important cage and viscosity

effects in bulk polymers. (2) The rate of initiation

was unknown. (3) The conversion was too high and

secondary reactions interfered. (4) It was carried

out at temperatures too high to be relevant to

practical problems. (5) It reports activation energies

that are inconsistent with simple primary reactions,

meaning that the authors are dealing with an unrecognized change in mechanism. (6) The rate of

oxidation was high enough or the sample was thick

enough that effective oxygen concentrations in the

polymer were diffusion-dependent and imknown.

(7) To the extent that the oxygen supply varies in

different regions of the polymer, so may the rates,

products, and mechanisms of oxidation vary.

The experimental work described here and in

progress was carried out by Dr. Etsuo Niki and

supported by the Army Research Office (Durham),

Amoco Chemicals Corporation, Esso Research and

Engineering Co., Polymer, Solvay & Cie., and

Unilever Limited.

10. References

[1] F. R. Mayo and K. C. Irwin, Polym. Eng. Sci., 9, 282

(1969) .

[2] J. H. Knox, Advan. Chem. Ser., 76, 1 (1968).

[3] D. L. Allara, T. Mill, D. G. Hendry, and F. R. Mayo,

ibid., 40,

[4] R. Hiatt, T. Mill, and F. R. Mayo, J. Org. Chem., 33,

1416 (1968).

[5] K. U. Ingold, Accounts of Chemical Research, 2, 1 (1969).

[6] G. A. Russell, J. Am. Chem. Soc, 79, 3871 (1957).

[7] T. G. Traylor and C. A. Russell, ibid., 87, 3698 (1965).

[8] D. G. Hendry, ibid., 89, 3433 (1967).

[9] C. Walling and A. Padwa, ibid., 85, 1593 (1963).

[10] J. D. Bacha and J. K. Kochi, J. Org. Chem., 30, 3272

(1965).

[11] I. V. Berezin, E. T. Denisov, and N. M. Emanuel, "The

Oxidation of Cyclohexane," English translation by

Pergamon Press, New York, 1966.

[12] J. A. Howard and S. Korcek, Can. J. Chem., 48, 2165

(1970) . [13] G. E. Zaikov, J. A. Howard, and K. U. Ingold, ibid., 47,

3017 (1969).

[14] F. F. Rust, J. Am. Chem. Soc, 79, 4000 (1957).

[15] J. C. W. Chien, E. J. Vandenberg, and H. Jabloner,

J. Polym. Sci., Al, 6, 381 (1968).

[16] L. Dulog, E. Radlman, and W. Kern, Makromol. Chem.,

60, 1 (1963).

[17] L. Dulog, J. K. Weise, and W. Kern, ibid., 118, 66 (1968).

[18] S. S. Stivala and L. Reich, Polym. Eng. Sci., 5, 1 (1965).

[19] J. H. Adams, J. Polymer Sci., A-1, 5, 1077 (1970).

98
Discussion

F. H. Winslow (Bell Telephone Laboratories)

In section 1 of his paper. Mayo indicated that

both thermal and photooxidations of saturated

aliphatic hydrocarbons proceeded primarily by

radical mechanisms. But, there is abundant evidence

that chain cleavage by a Norrish II (non radical)

process is a major mechanism in the oxidative photodegradation of polyethylene. In addition, the hydroperoxide intermediates, so important in thermal oxidations of hydrocarbons, are absent or at least their

presence has not been demonstrated in photo-oxidations. In fact, our studies have shown that hydroperoxides, formed by autoxidation of polyethylene,

decompose so fast on exposure to 3000 A radiation

that significant concentrations cannot develop during photooxidation. Some Norrish I or other radical

reactions undoubtedly occur in photo-oxidations at moderate temperatures but radical-chain reactions

are absent or at least insignificant since: (1) photooxidations of hydrocarbons at room temperature or below are essentially non-autocatalytic and exhibit no induction period. (2) Chain terminating antioxidants such as hindered phenols have no significant effect on the degradation rate.

D. W. Brown (National Bureau of Standards)

Polymers of 3 ,3 ,3-trifluoropropene; 2,3,3,3-

tetrafluoropropene; and 3 ,3 ,4 ,4 ,5 ,5 ,5-heptafluoropentene-1 were aged at 225 °C in air and in

preevacuated tubes. At various times samples were

dissolved in hexafluorobenzene and the intrinsic

viscosities were determined at 29.8 °C. In 87 hours none of the samples heated in preevacuated tubes

changed in intrinsic viscosity. With the air-aged

specimens, the intrinsic viscosity of poly (3,3,3-

trifluoropropene) decreased rapidly whereas values

for the other two polymers showed little change.

Presumably, the relatively great resistance of poly-

(2 ,3 ,3 ,3-tetrafluoropropene) to oxidation is due to

its lack of a tertiary hydrogen atom. Our speculation

as to the cause of the resistance to oxidation of

poly(3 ,3 ,4 ,4 , 5 , 5 , 5-heptafluoropentene-l) is con- nected with its solubility. Although this polymer has

a hydrocarbon backbone, it is soluble in perfluorohexane, whereas neither of the other polymers is so

soluble. In acetone, the polypropenes are soluble and

the polypentene insoluble. Possibly, the long fluorocarbon side chain associates with the hydrocarbon

backbone and confers fluorocarbon-like behavior, in

oxygen as well as in solution.

99

NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Polymer-Gas Reactions (Air Pollutants: NO2 and SO2) as Function of

Pressure, UV Light, Temperature, and Morphology—A Survey

H. H. G. Jellinek

Clarkson College of Technology, Potsdam, N.Y. 13676

Reactions of various polymers, such as polystyrene and its stereo-specific isomers, butylrubber, nylon, etc., with nitrogen dioxide and sulfur-dioxide were studied by the author and

co-workers in recent years. Mechanisms of these reactions in presence of these gases, near ultraviolet light and oxygen were elucidated. Diffusion phenomena have also been investigated and

their importance has been pointed out. More recently, work has been initiated on the influence of

polymer morphology on degradation of polymers in presence of these gases, near uv radiation

and oxygen. Unexpected effects have been observed during chain scission near room temperature.

Thus, for instance, isotactic polystyrene of various crystallinities, as far as extent and type are concerned, show marked differences in their degradation characteristics. Thus, for instance,

crystalline polymers show faster degradation than amorphous ones, which seems to be contrary

to expectations. However, this phenomenon can be explained in quite a consistent manner. The

importance of all these reactions in connection with air pollution is briefly discussed.

Key words: Air pollutants; butylrubber, degradation, effect of NO2; degradation, effect of morphology; effect on polymers of NO2, SO2, air and U.V. light; nylon, degradation, effect of NO2;

polymer-gas reactions; polystyrene, atactic, degradation, effect of NO2; polystyrene, diffusion of

NO2 and degradation; polystyrene, isotactic, degradation, effect of SO2.

1. Introduction

We have studied reactions between polymer films

and gases such as NO2, SO2, O2 (air), irradiated with

UV light, in recent years. The kinetics and mechanisms were elucidated. We have also started investigations on the effect of polymer morphology

(extent of crystallinity) on the stability of various

polymers under such "environmental" conditions.

Quite unexpected results were obtained as function

of morphology, contrary to expectations, which,

however, after some reflection, appear to be quite

reasonable and can be at present, qualitatively

explained.

The films which were used for these investigations were usually quite thin (ca. 20 iim to 25 iim), this

eliminates diffusion phenomena. However, for thicker

films, diffusion has to be considered and we have

done some research in this direction. The significance

of all this work is not only of fundamental but is also

of practical value with respect to air pollution.

2. Some Reactions of NO2 and SO2

It will be opportune to discuss briefly some re- actions of these gases with O2 and small hydrocarbons

in presence of ultraviolet light; these processes are relevant to the reactions with polymers which will

be discussed in this paper.

NO2 is an odd electron compound, hence very

reactive—a strong oxidizing agent. It reacts in

presence and absence of oxygen (in the dark) with

imsaturated hydrocarbons such as cyclohexene.

1-3 butadiene and 1-hexene. It is suspected that

amongst the reaction products, there are small

amounts of powerful (effective in parts per billion)

eye irritants of the type

R R

\ /

C=C

/ \

NO2 R

In presence of uv light, NO2 reacts very strongly

with hydrocarbons forming aldehydes, ketones and

epoxides [1].^ It is suspected that ozone is formed

during some of these reactions,

N02+/jr^N0+0

0+N0+M^N02+M

0+N02->NO+02

0+N02+M^N03+M

0+02+M^03+M

NO+N03->2N02

NO+03-^N02+02

If this is the case, much of the deterioration of

elastomers (e.g. rubber) will be due to ozone. Thus

olefins (e.g. CH3CH=CHCH3) can react with

^ Figures in brackets indicate the literature references at the end of this paper.

101
ozone as follows,

H 6

CHsCH^CHCHs+Os^CHsC—CHCH3

00-

H

-^CHsCOO+CHsCHO

The reaction of SO2 with saturated and unsaturated hydrocarbons in presence of uv light leads

to sulfinic acid; Dainton and Ivin [2 3 suggested the

following reaction sequence,

SO,*-^SOo

SO?- -SOa^SOo- -SOo

S02*+RH^S02+RH

S02*+RH-^RS02H

Also here ozone may be one of the reaction

products,

SO2 +202—^803+03

Thus NO2 and SO2 are quite aggressive under

various conditions; it is not surprising that polymers

are aifected by these gases.

3. Reaction of NO2 and SO2 with

Polymers [3] A survey was first made of the effect of SO2 and

NO2, respectively, in absence or presence of O2 and

near uv radiation on various polymer films. The high

vacuum apparatus will not be described here. Two

classes of polymers stand out quite distinctly with

respect to main chain scission due to these gases:

(1) saturated polymers and (2) unsaturated

polymers (elastomers). The latter are very much

more reactive, than the former, as is to be expected.

The saturated polymers show a variety of different

reactivities according to their chemical structure.

Polycarbonate, for instance, is fairly stable, polyesters undergo chain scission and cross-linking. All

reactions were carried out at about 30 °C.

Some reactions were studied in detail and a short

survey will be given here, indicating their most

prominent features.

4. Reaction of NO2 with Atactic

Polystyrene Films [4]. High vacuum apparatuses were used in this and

subsequent investigations in conjunction with

medium pressure uv light sources; all light below

about 2800 to 3000 A was cut off by pyrex glass.

The reaction vessels, arranged in a circle around this

light source, were thermostated. Devices were

incorporated in the reaction vessels so that films on

glass -slides could be exactly located with respect to the light source. At the end of each film exposure,

the film was evacuated to prevent any continuation

of the reaction. Intrinsic viscosities were measured

at definite time intervals and eventually converted

to number average chain lengths.

It is of importance to note that there exists an

equilibrium between the monomer (NO2) and its dimer (N2O4). Below about 1 cm Hg of NO2 pressure,

this equilibrium can be neglected as it lies for low

pressures practically completely on the side of the

monomer, which is the reactive species,

N204?^2N02

The partial pressure of NO2, Pno2'> is given by,

Pno, = [ ( A'2p+4XpPtotai) -K^'] ( 1

)

Here, Kp, is the equilibrium constant and Ptotai =

Pno2+Pn204 the total gas pressure. Films prepared

by casting from solution [4] were exposed over a range of temperatures from 25 to 55 °C and pressures

from 11 to 60 cm (total) . The results are expressed in

degrees of degradation a = — -f~ versus ex- DP„, DP,n,0 posure time, t. A typical plot for 30 cm NO2 as a

function of temperature (dark reaction, no air) is shown in figure 1. DP„ t and i)P„,o are the number

average chain lengths at time t and t = 0, respectively. The Arrhenius energy of activation for this

reaction amounts to ca. 12.2 kcal/m and the preexponential factor is about 2 X lOHo 1 X 10^ hr~^.

The films were not treated especially during their

formation as far as their morphology is concerned

and they are most likely amorphous. At this stage

of the work, morphology was not investigated.

Quite an interesting and noteworthy phenomenon

was observed during the initial stages of the degradation reaction. Intrinsic viscosity measurements

require straight lines for their plots of Tjsp/c versus c. However, after relatively short exposure times of

the films, straight lines could not any more be

~T r

lO 26 30 40 50 60 70 80

TIME (HRS)

Figure 1. Degradation of atactic polystyrene as function of temperature {30 cm Hg of NO2).

102
obtained in benzene solutions. Curves were obtained on dilution, which suggested that disaggregation of

polymer molecules took place. Apparently, NO2 and

peroxide side groups are formed which make the

chains polar, causing them to aggregate in benzene

solution. This effect disappeared on further exposure.

Straight lines, however, could be obtained for these

plots in dioxane solution at all stages of the degradation process.

Degradation data can be satisfactorily accounted

for by a mechanism, which can be considered to be

the analog of Holland's [5] oxidation mechanism,

RH+N02-^R+HN02

Fast R-+N02-

RNO2-

RNO2-

(N02+N204)+CageChain Scission:

Cage+(N02+N204)-

•RNO2+RH-

•R,N02+N02-

•Ry+'R:C?

-^RN02

>Inert Products

>(-R.N02Ri) Cage

»RN02+(N02+N204)

>-R,+-R.N02+(N02+N204)

>HRN02+R-

>02NR,N02

iCage^P,+P.

RH stands for polymer, R- etc. for polymer radicals

and P for 'dead' polymer.

Evaluation of this kinetic scheme leads to.

a = KmPtot&iP:NO2 (2)

The final result actually contains some refinements, which will not be discussed here. A few

words may be added about "cages". It is generally

agreed now that chain scission and also other steps

in polymerization and degradation reactions, proceed

via solvent or medium cages. The problem in the

case of polymerization was treated by Benson and

North [6] and North [73- A cage is considered

which contains potentially two polymer radicals.

The escape of these radicals from the cage is a

diffusion process (we discuss here the degradation

case) and can be expressed by considering its rate

constant, fes.

(3)

Rab is the maximum distance of the centers of gravity

of the two imprisoned radicals in random coil

configuration, at which reversal of the process to the

original polymer can still occur. Da and Db are the

respective translational diffusion coefficients of the

radicals in the surrounding medium. On the average,

two radicals of equal length are produced, hence,

eq. (3) simplifies to,

fe3 = 47ri?AB(2D) (4)

If the Stokes-Einstein relation holds, D is given by,

D = kT/TrB'7,R (5)

k is Boltzmann's constant, B' a constant equal to six

for large particles, ri is the medium viscosity and jR

the radius of one of the radical coils. The latter

is given by,

R% = R\+R% = 2m (6)

Introducing equations (5) and (6) into the original

eq. (4) yields.

^3 =

8\/2kT

B'v

(7)

Equation (7) is well obeyed in the case of poly-amethylstyrene in solutions of different viscosities

The diffusion constant D can be obtained from

eq (5) (random coils). The root mean square radius

of a polymer radical in a cage in the configuration

of a random coil is given by.

R={lDPn.t.cy'H/{B') 111 (8)

DPn,t,c is the critical chain length where the average

kinetic chain length e has just become equal to that

of the polymer chains, I is the length of a structural

unit of the chain, {B' = 6). Calculated and experimental values for poly-a-methylstyrene in solution are quite reasonable [8]].

In summary, the whole degradation process in case of polystyrene consists of a random chain

scission process and incorporation of NO2 as side

groups. IR spectra for polyethylene show this

incorporation quite clearly [9]. The chain scission

is controlled by diffusion out of cages.

3.0

30 40 50

t (hrs.)

6 0 7 0

Figure 2. Degree of degradation a of butyl rubber as a function

of time for various NO2 pressures in absence of air at 35 °C:

(1) 1 mm Hg, (2) 0.5 mm Hg, (3) 0.35 mm Hg, (4) 0.20 mm Hg. The

lines are calculated, the points are experimental data.

103
5. Butylrubber and NO2 [10,11]

A typical elastomer (Butylrubber) was also

investigated. This polymer had an isoprene content

of 1.75 percent by weight. Polymers having larger

amounts of isoprene, become too sensitive for proper

handling.

Here again, cast films (20-25 fira thick) were

exposed under various conditions. The NO2 pressures were quite small 0.01 to 1 mm Hg and the temperature range was 25 to 65 °C. These small N02-pressures were obtained in the apparatus by keeping

liquid NO2 at certain temperatures for which its vapor pressures were of the desired magnitude.

Experiments were performed in absence or presence,

respectively, of 1 atm of air. Some typical degradation curves are shown in

figure 2 at various NOo pressures (no air).

The rates of chain scission in presence of air are about five times larger than those in its absence

(see fig. 3). There are a number of weak structures

in butylrubber, more susceptible to chain scission

than the normal ones. The chain scission reaction

with NO2 alone can be pictured as follows:

—C=C—hNOo?

CH3 H

k2

(-

NO2

-C—C—

)

Cage

CH3 H

+NO2U3

Chain Scission +NO2

The double bond is directly attacked by NO2

whereas in the case of Oo the a hydrogen atom reacts

first and the unsaturation remains initially intact.

Evaluation leads to.

a = (9)

Kexp is a function of the NO2 concentration.

If O2 is also present, three reactions take place

with respect to chain scission: (1) Chain scission

due to NO2 alone, (2) Chain scission due to O2 alone,

(3) Synergistic reaction due to NO2 and O2.

6. Mechanism

(1) Reaction with NO2 only:

RH+N02-^Cagei

Cageik2

Cagei+N02 >Chain Scission +NO2

(2) Reaction with O2 only:

RH+O2 >R +HO2

ki, R+O2 >R0o

•RO2+RH >ROOH+R

I 0 20 3 0 Hours

4 0 5 o 6 0

Figure 3. Degree of degradation of butyl rubber in presence of

1 atm of air at 35 °C as function of NO2 pressure and time.

(1) 1 mm Hg, (2) 0.5 ram Hg, (3) 0.02 mm Hg, (4) 0.01 mm Hg.

ki ROOH »Inert Products

ks ROOH >Cage2

ks Cagco >ROOH

km

Cage2+02 >Chain Scission +O2

(3) Synergistic Reaction of NO2 and O2:

ROOH +NO2—^N02- ROOH

NOa-ROOH-^Cages

kis Cages

kii ^N02-R00H

Cages >Chain Scission

il6 R+R- >Cage4

Cage4+02 >R- +R-+02

kn

Cage4 >R — R

If such a mechanism is evaluated, it leads to an

104
0 10 20 30 40 50 60

t (hrs)

Figure 4. a versus tfor 1 mm Hg of NO2 at 35 "C with and tvithout near uv irradiation {no O2), (i) without, {2} with irradiation.

equation as follows:

a = Kexpt+ w (1 - e-^y^ (10)

This equation represents the experimental data well.

Studies were extended to include the effect of near

u.v. light QllJ. Figure 4 shows a comparison of a

versus t in presence and absence of near uv radiation

(1 mm Hg of NO2, 35 °C, no O2). Actually NO2

retards chain scission when near uv light is present.

Figure 5 gives data for degradation in presence of uv radiation, 1 atm of air and various NO2 pressures

at 35 °C. One curve in the absence of radiation is also shown.

As pointed out above, NO2 reacts directly with a NO2

double bond forming a radical -^C—C--^ which is

CH3 H

located in a cage. The actual mechanism suggested

is as follows.

Reaction with double bond:

negligible on exposure to near uv radiation.

RR+hp >R+HR +NO2 >RN02 Inert Product

Reaction with

double bond: RH+N02<i±(RHN02) Cage

k2

ks (RN02)+/m' >Inert Product

k 4 (RNO2) >Chain Scission

Evaluation gives,

«1=_ fciA:4[n^]o[N02]f^ = ^ ^, P A.exp,lf - •^.exp,lLJ^02J^

L"Jo

(11)

and [n]o are the initial isoprene and total

monomer concentration, respectively. Aexp.i is a function of NO2 pressure. In absence of light the

corresponding relation is a = Kexp.it.

Photo-oxidation (no NO2) leads to,

Ln'^okihh[P2y _ J. ,2- IT' rn 1/2 (io\ <X2— r n — -'^-exp,2r — Aexp,2L^2jt {LA)

The tentative mechanism is as follows:

RH+fej/^R+HR+02-^R02

•R02+RH^R00H+RROOH^iCage4

ks

ke Cage4 +02 >Chain Scission +O2

Addition of NO2 gives 0:3 = Kexp.s^^; Kexp.s here is a function of O2 and NO2 pressure, respectively. If the

NO2 pressure is increased, the mechanism changes.

The experimental curves change from convex to concave towards the time axis. An additional step seems to become significant, where NO2 reacts

directly with the double bond, which was the case when this gas is present alone: R +NO2—>RN02

(Inert). The degree of degradation in this case

is given by,

« = Kexp ,3f' +Ki[N02] ( 1 -e-*3« ) ( 13

)

For sufficiently high NO2 pressures a reduces to the

second term and for sufficiently small NO2 pressures

only the first term has to be considered.

7. Diffusion of NO2 in Atactic

Polystyrene Films [12]

A detailed study was carried out on diffusion of

NO2 through several polystyrene films. Diffusion

105
X 0 10 20 30 40 50 60

t (hrs)

Figure 5. a versus t for 1 atm air at 35 °C near uv irradiation {except curve 1) and

various NO2 pressures in mm Hg:

(1) dark reaction, 1; (2) 1; (3) 0.5; (4) 0.2.

coefficients were determined by two different

methods, whose data agree quite satisfactorily:

(1) Penetration and (2) Sorption-Desorption experiments. A special device was constructed for the first method which is illustrated in figure 6. The polymer

fdms were arranged in the form of a tightly packed

deck of cards. After diffusion time, each fdm was

analyzed and its a value was determined. (This is the first time that a was taken as a indicator of

diffusion). The plot of a versus distance from the

exposed film surface was used for evaluation of the

diffusion coefficients. Sorption-Desorption curves

are shown in figure 7. The penetration experiments were evaluated according to

Cno....„ "^'"*^2(Z),f)^/^ ^^^^

Cnoz.i and Cno2.x=o are the NOo concentrations at distances x and .i: = 0 from the exposed polymer

surface, Dp is the diffusion coefficient and t the

diffusion time. It was found experimentally that the

degradation and incorporation of NO2 can be

neglected.

The sorption-desorption experiments were evaluated by a relation as follows,

8 " 1 w,/w =1— — y —=

Here Wt is the weight sorbed or desorbed, respectively, at time t and that at saturation. Ds and

Dd are the respective diffusion coefficients for

sorption or desorption, respectively, (usually the

average of the two coefficients is taken). The

diffusion coefficients are of the order of 10~^ cmYs.

The Arrhenius energy of activation for all diffusion

coefficients (Penetration and Sorption-desorption)

are: Ptotai>10 cm: (8.6±1.8) kcal/m, and for

Ptotai = 2 cm: (9.4±5.3) kcal/m. (The =t values

are standard deviations.)

8. Isotactic Polystyrene and Sulfur

Dioxide [13,14]

An investigation of the effect of sulfur dioxide on

isotactic polystyrene was carried out. This reaction

8 8mm

Figure 6. Diffusion cell. B metal bellows K screw for tightening bellows

J special connection to high vacuum at far side of films F set of films (ca. 40 jum each) O Viton-O-ring

T Teflon joint P Teflon covered back plate S screws

106
3. O

Figure 7. Sorption-desorption curves of {NO2 + NiOi) for

polystyrene at 45 °C.

Sorption: (1) 60 cm, 69.1 fi; (2) 30 cm, 78.0 m; (3) 15 cm, 78.0 m- Desorption: (4) as (1); (5) aa (2); (6) as (3).

was studied in the presence of constant light intensity,

oxygen and SO2 pressures as function of temperature

(30 °C to 57 °C). Small pressures (e.g. 0.85 mm Hg)

of SO2 have quite an appreciable effect on the

polymer, leading to chain scission in presence of

150 mm Hgof O2 and near uv light. (SO2 alone at thisn

pressure has no effect). Films were cast from

chloroform solutions (ca. 25 ;um thick). Random

chain scission takes place, which can be expressed by

a- DPn,t DPn,

(16)

The Arrhenius energy of activation amounts to

4.5 kcal/m, the preexponential factor is very small

(1.02X10~- hr"'). The light intensity exponent was

found to be 1.2. A mechanism which accounts satisfactorily for the

experimental results was suggested:

hp

hv R00H+S02<^Cagei

CagCi >Chain Scission

R00H?±Cage2

Cage2 >Chain Scission

kg

Jiio, Jin

R+SO2-

•RSO2 +RHRSO2Hki3 ku

ki5 The degree of degradation a for this reaction scheme

becomes

->-RS02

^•R+RSOsH

-^Inert Products

ih +klSO^']) (fc3C02] +/Cl3[S02])

;i7)

Figure 8 gives some rate constants for isotactic

polystyrene as function of SO2 pressure (15 cm Hg

of02) [14].

9. The Effect of Morphology on

Degradation in Presence of SO2

and NO2, respectively [15,16]

Still more recently, work has been initiated on the

effect of morphology (the percentage of crystallinity

was changed) on the stability of various polymers

such as isotactic polystyrene [15] and Nylon [16].

The amount and type of crystallinity depends on

the mode of preparation of a particular polymer.

Melting a polymer and slow cooling usually produce

crystallinity. Also slow precipitation from hot

solution yields crystalline polymers. The crystalline

morphology is usually present in a variety of stages,

sizes of crystallites and crystalline size distribution:

needles, single crystals, spherulites, fibrils and so on. Orientation of polymer films and fibers also leads to

crystallinity. The arrangement of individual chains

in the crystallites is of great importance. Chain

folding was found in this work to be of great signifiRH+02- ^R -t-H02 h

S02-

hp -^S02* K

^exp

10'. hr'

S02*- >S02 k^

2.5; RH+S02*-—>R + HS02 h

R+02-—>R02 h

2.0-

•R02+RH —>R00H+R- ki

1.5: ROOM ——^Inert Products

ROOH+SO2 >Inert Products

(e.g. ROSO2H)

h

1.0 f

0.5

6 p^^.mmHg

Figure 8. Experimental rate constants for isotactic polystyrene

as function of sulfur dioxide pressure (15 cm Hgof O2).

107
cance for chain scission. The main chain links in the

folds due to their additional energy content are more

susceptible to chain scission than the links in the

straight chain portions. The type of crystallites and

their structure are best illustrated by some characteristic drawings (figs. 9, 10).

The polymer films prepared in this work, were

characterized under a polarizing microscope and the

approximate percentage crystallinity was obtained

from density measurements. Isotactic polystyrene

in practically amorphous form and films of various

degrees of crystallinity (fibrils) were degraded in

presence of 150 mm Hg of O2, constant light intensity

(\>2800 A) and various SO2 pressures over a range

of temperatures. Random degradation took place.

Degradation was slowest for the amorphous sample

(0.93 Xl0~^ s~^), the crystalline films degraded at about twice this rate for 1.0 mm Hg of SO2. At higher

SO2 pressures the discrepancy increased: 3:1 for

amorphous : crystalline. The energy of activation for

amorphous films is practically zero whereas for

crystalline films (ca. 5 percent w/w), it amounts to

4.5 kcal/m. The increase in the experimental rate

constants with w/w percentage crystallinity is shown

in Fig. 11. It follows approximately a straight line. It was ascertained that the increase in these rate

constants is not simply a concentration effect, by

assuming that SOo reacts only with the amorphous

polymer portion. Apparently the difference in

stability is due to the many folds of the chains in the

A, FOLDED CHAIN THEORY

o 00 Q on

<<

11

CJ

Ti n

n n n

(J O U O

r>r? n

o

o

o

o

B. FRINGE- MICELLE THEORY

Figure 9. Morphology of crystalline polymers.

A. Schematic diagram of polymer crystallite made up of

folded chains. B. Crystallites and amorphous regions ac- cording to the fringe-micelle theory.

Figure 10. Crystallization from solution according to the pic- the picture of Schramm, Trillat and Sella, and Dolmetsch

(a) Coil model, (b) After coiling, to an extent depending on chain stiff- ness, the molecule assumes a zig-zag form, (c) Several folded molecules

aggregate end-to-end and form an elementary fibril.

crystalline polymer which have additional energy

content because of strain and are better able to

escape from cages. The susceptibility of main chain

links in folds to degradation was also noticed in the

case of reaction of nitric acid with polyethylene

[17]]. The reaction of O2 is diffusion controlled under

the experimental conditions; this was found in

previous work on oxidative degradation of isotactic

polystyrene [18]. However, SO2 is quite soluble in

isotactic polystyrene and diffusion does not play a

role for this gas. Similar experiments were carried

out in the presence of NO2 This gas has an

appreciably different effect; while in the case of SO2

the rate constants increase approximately linearly

with weight percentage of crystallinity, in the case of

NO2, these constants only increase initially until a

relatively small percentage of crystallinity (8 percent

w/w) is reached, but then remain constant with

further increase in crystallinity. The explanation for

this behavior is only tentative. Here again the main

chain links in folds may be assumed to degrade

faster than those in the straight line parts. It was

pointed out that oxidative degradation is diffusion

controlled; the diffusion coefficient for O2 into polystyrene is actually larger than that for NO2. Hence

the reaction with NO2 must also be diffusion controlled. The diffusion rates will decrease with

increasing crystallinity, at the same time the rate of

reaction will increase with crystallinity. There must

come a point where the reaction becomes completely

diffusion controlled and the rate constants from this

stage on remain constant.

Some studies are now being carried out on Nylon

[16] (amorphous and samples containing sperulites)

The crystallinity was again characterized with the

polarizing microscope and by density measurements.

Exposure to NO2 alone (0.5 mm Hg) leads to chain

scission, which ceases at a definite time. The crystalline material shows about twice the rate of that of the

amorphous polymer. The Arrhenius parameters for

either sample are practically identical. Oxygen and uv light accelerate the reaction. The rate in this latter

108
0 10 20 30 40

% */w CRYSTALLINITY

Figure 11. Experimental rate constants as function of crystalUnity for isotactic polystyrene as function of percentage of

crystallinity.

Dotted line: rate constants as function of SO2 concentration in amor- phous parts of films only, zero concentration in crystalline parts.

case is also faster for the crystalline material. These

studies have not yet been concluded; however, it seems to be the case that a concentration effect is operative. The number of susceptible folds and

hence the number of "weak" links is larger for the

crystalline samples than for the practically amorphous ones.

10. Summary

Atactic, isotactic polystyrene and butyl rubber

have been degraded by ]N02 gas alone near room

temperature. In presence of oxygen, the chain

scission rate is enhanced. In some instances NO2 gas can actually function as retarder if near uv light is present. In contrast to oxygen, NO2 attacks double

bonds in elastomers directly. Usually, diffusion does

not play a role with thin films of the order of 20 jim.

Isotactic polystyrene seems to be an exception.

SO2 alone is not reactive (0.85 mm Hg). Diffusion

was studied for thick films; the diffusion coefficient

from penetration and sorption-desorption experiments, taking the degree of degradation as indicator

for the progress of diffusion, agree well with each

other. The order of magnitude of the diffusion

coefficients is similar to those usually obtained for

gas diffusing through polymers. Polymer morphology has appreciable influence on the stability of

polymers with respect to NO2 and SO2 in conjunction

with oxygen and near ultraviolet light. Contrary to

expectations, the rate of degradation in crystalline

isotactic polystyrene (SO2 and NO2, respectively)

and Nylon (NO2 only) are about twice those corre- sponding to amorphous polymers. Indications are that main chain links in chain folds of the polymer

degrade faster than links in the straight chain

portions, because of excess energy in the folds.

This work was supported by a Grant from PHS,

Division of Air Pollution, and another from NASA.

11. References

[1] R. J. Cvetanovic, Journ. Air Poll. Contr. Assoc., 14,

No. 6, 208, (1964); see also P. A. Leighton, Photochemistry of Air Pollution, Academic Press, Inc., New York,

1961 and Chemical Reactions in the Lower and Upper

Atmosphere, Proc. Internatl. Symp. San Francisco,

April 18-20, 1961, Interscience Publ., New York, 1961.

[2] Dainton and Ivin, Trans. Far. Soc, 46, 473, 382, (1950).

[3] H. H. G. Jellinek, F. Flajsman and F. J. Kryman, Journ. Appl. Polym. Sci., 13, 107, (1969); H. H. G. Jellinek

and F. J. Kryman, ibid., 13, 2504, (1969); H. H. G.

Jellinek, Journ. Air Poll. Contr. Assoc., 20, No. 20, 612,

(1970).

[4] H. H. G. Jellinek and F. Flajsman, Journ. Polym. Sci.,

A-l, 7, 1153, (1969); see also H. H. G. Jellinek and

Y. Toyoshima, ibid., A-l, 5, 3212, (1967).

[5] J. L. BoUand, Quart. Rev., 3, 1, (1949).

[6] S. W. Benson and A. M. North, J. Am. Chem. Soc, 84,

935, (1962).

[7] A. M. North, The Collision Theory of Chemical Reactions

in Liquids, Methuen, London, (1964); also A. M. North,

The Kinetics of Free Radical Polymerization, Pergamon

Press, New York, 1966; A. M. North, Quart. Rev., 20,

421, (1966); A. M. North, Diffusion Control of Homofeneous Reactions, Vol. 2, p. 95 Progress in Polymer

cience, ed. J. Jenkins, Pergamon Press, New York,

1969.

[8] H. H. G. Jellinek and M. D. Luh, Europ. Polym. Journ.,

Supplement, 149, (1969).

[9] T. Ogishara, S. Tsuchiya and K. Juratani, Bull. Chem.

Soc. Jap., 38, 978, (1968).

[10] H. H. G. Jellinek and F. Flajsman, Journ. Polym. Sci.,

A-l, 8, 711, (1970).

[11] H. H. G. Jellinek and P. Hrdlovic, Journ. Polym. Sci., in press, 1970.

[12] H. H. G. Jellinek and S. Igarashi, Journ. Phys. Chem.,

74, No. 7, 1409, (1970).

[13] H. H. G. Jellinek and J. F. Kryman, Symp. Anaheim,

October, 1969, Proc. in Plenum Press, 1970, p. 85; see

also J. F. Kryman, M.S. Thesis, Clarkson, 1969.

[14] H. H. G. Jellinek and J. Pavlinec, ibid., 1970, p. 91.

[15] P. Hrdlovic, J. Pavlinec and H. H. G. Jellinek, Journ. Polym. Sci., in press, 1971.

[16] H. H. G. Jellinek and A. Chaudhuri, in preparation, 1971.

[17] I. M. Ward and T. Williams, Journ. Polym. Sci., A-2, 7, 1585, (1969), also K. H. Illers, Makrom. Chem., 118,

88, (1968).

[18] H. H. G. Jellinek and S. N. Lipovac, Macromolecules,

3, 231, 237 (1970).

Discussion

Frank R. IMayo (Stanford Research Institute)

While this manuscript contains some interesting

new data, some of the mechanisms presented in

explanation cannot be taken very seriously. For

example, reaction 16 on page 96^ says that two caged

^ Page numbers here refer to those in preprints.

polymer radicals can collide with oxygen in a polymer matrix without reacting with either oxygen or each other. This reaction cannot be mechanistically

significant. I know of no precedent for reaction 4,

page 97; the usual cleavage occurs at the 0-0 bond.

Reaction 9 in the same set seems most improbable;

a hydrogen atom should readily abstract a hydrogen

atom from the polymer before it could find a peroxy

109
radical. The set of reactions on page 99 neglects the

O

reaction of R—S radicals with oxygen and ignition

O

by RO2H +SO2.

H. H. G. Jellinek:

In answer to Dr. Mayo's remarks I should like to

make the following observations. He maintains that

"some of the mechanism cannot be taken very

seriously." In answer to this I wish to say that in

general all reaction mechanisms contain considerable

elements of doubts; for example Bodenstein's

mechanism for hydrogen iodide formation held up

very well for about sixty years. As far as the specific

comments are concerned my answers are as follows:

(I) reaction 16: this reaction does not seem to be

unreasonable to me as some additional energy is needed and collision with the reactive sites of the

polymer radicals will be very much more unlikely

than collision with the cage system as such. I agree

that reaction 4 is not correct. I have omitted reactions (4), (5) and (9) in that mechanism. From a kinetic point of view they are superfluous anyway.

Reaction (9) was already eliminated when finalizing

the paper on this topic (ref. 11). Reaction (2) actually

implies that S0| is the reacting species. I have

formulated this more specifically now. The reaction

R-SO2 +O2 does not enter into calculations which

have any connection with the chain scission reaction

and does not need to be considered.

W. Lincoln Hawkins (Bell Laboratories)

The reactions of isotactic polystyrene with SO2 in

the presence of oxygen may be important in the

oxidation of polymers inhibited with various sulfur

compounds. It has been shown that aromatic disulfides inhibit the oxidation of polyolefins, presumably

by catalyzing the decomposition of hydroperoxides

into inert products. In the course of the reaction of

naphthyl disulfide with organic peroxides, it has been

determined that traces of SO2 are formed. The

catalytic effect of SO2 in peroxide decomposition has

been demonstrated in these systems, supporting the

reaction

ROOH^inert products

shown in the kinetic scheme for the oxidation of

isotactic polystyrene.

110
NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Controlled Gaseous Oxidations of Organic Compounds

C. F. CuUis

The City University, London, England

The homogeneous gaseous oxidation of organic compounds generally leads to a wide variety

of products and this may be ascribed not only to the occurrence of a number of consecutive and

concurrent reactions but also to the fact that the various constituent elementary steps tend to interact with one another in complex ways. Several methods are outlined for minimising the extent

of unwanted consecutive reactions but it appears to be much more difficult to restrict undesirable concurrent reactions. If the mutual influences of the various part -reactions are to be understood

and controlled, it is important to elucidate the nature and role of the intermediate molecular

products. Although here carefully planned experiments involving the artificial addition of reactive

intermediates can yield some information, a much fuller understanding of the quantitative importance of the different competing reaction paths can be obtained from experiments with iso- topically -labelled intermediates. Several examples are given of the use of isotopic tracer techniques

in the unravelling of the mechanism of complex gaseous oxidation reactions.

Key words: n-butane; 1-butene; 2-butene; concurrent reactions; consecutive reactions; controlled

oxidation; conventional additive experiments; cool-flames; re-heptane; intermediate products;

isobutane; isotopic tracer techniques; isotopically -labelled additives; macroscopic stages; mutual

influences; O-heterocycles; radio gas chromatographic analysis; selective inhibition; selectivity;

self-inhibition; thermal control.

1. Introduction

The oxidation of organic compounds by oxygen is one of the most widespread of all chemical processes

used by man. Thus the combustion of hydrocarbons

in air is still among the most important sources of

heat or mechanical energy, but if this energy is to be

used to the best advantage, the chemical reactions

taking place must be suitably controlled. One of the

crudest forms of control consists in more or less

complete inhibition as exemplified by the use of

"anti-knock" additives to eliminate "knock" in

internal combustion engines. Another example of

similar rigid control is to be foimd in the ability of

suitable antioxidants to prevent or severely minimise

the oxidation by air of solid or liquid organic compounds at or near ambient temperatures.

Nowadays the various fractions of petroleum,

which consist almost exclusively of hydrocarbons,

are being used to a rapidly increasing extent as starting materials for the production of other

chemicals, and controlled oxidation by air or oxygen

is one of the most important reactions to which such

compounds are subjected. In this case, what is wanted is no longer to inhibit reaction completely

but rather to exercise a more subtle and refined form

of control, so that oxidation is encouraged to proceed

selectively along a particular desired route giving

high yields of specific intermediate products.

Although such controlled oxidation of organic compounds can be effected either in the gas phase or in

the liquid phase and may be assisted by both

homogeneous and heterogeneous catalysts, this

paper is concerned primarily with possible ways of

'directing', as it were, the purely gaseous processes

taking place at moderately elevated temperatures.

2. The Principal Complicating Factors

The homogeneous gaseous oxidation of organic

compounds provides an example, par excellence, of a system involving a large number of consecutive and

concurrent stages. Now, even in the case of a very

simple chemical system consisting of, say, two

consecutive first-order reactions:

ki B-*C

the kinetics are relatively complicated. Harcourt and

Esson [[1],^ over a hundred years ago, first integrated

the differential equations concerned and showed that

the concentration of the final product, C, is given as a fimction of time, t, by the expression

C = Ao 1 +

ki —h

k2t. kic -kit^

Various possible combinations of consecutive first- order and second order reactions were later considered in detail by Chien pj, who carried out the

difficult integrations involved. In general, however,

for reactions more complex than these "ideal" cases, the relevant equations can be handled mathematically only by making certain (often rather

' Figures in brackets indicate the literature references at the end of this paper.

Ill
drastic) simplifying assumptions. The complexity in

the case of real systems involving the oxidation of even the simplest organic compounds is of course enormously greater; for not only are the various

part-reactions themselves complex free radical chain

processes but many of the intermediate products

formed exert specific effects (the nature and magnitude of which are highly dependent on the prevailing experimental conditions) on the reactions of

the other species present. Thus, in order to isolate

high yields of a specific intermediate product, such

as C, in a system of the type:

^ D E

F >- G

it is clearly necessary to control not only unwanted

consecutive and concurrent steps but also the

mutual interactions of the various reactions involved.

3. Control of Consecutive Reactions

Perhaps the most obvious way of eliminating or at

least minimising the further reaction of desired

intermediate products is to carry out oxidation at

very low conversions. Thus, for example, Knox [3]

has shown that, although the oxidation of propane

leads eventually to a complex mixture of products,

nevertheless if measurements are confined to the

first 1-2 percent of reaction, propylene is effectively

the sole organic product formed.

In other systems, including some where complications might at first sight be expected to be rather

severe. Nature comes to one's aid in the sense that

the very presence of certain intermediate or final

products simplifies the overall behaviour. Thus in

the system:

k

^

B >

A may inhibit the conversion of B to C. One example

is provided by the oxidation of primary aliphatic

alcohols [4]. Here the initial product formed is the

corresponding aldehyde, which would be expected,

at the relatively high temperatures concerned, to

undergo rapid further oxidation and in so doing

perhaps act as a degenerate-branching agent leading

to other modes of attack on the original alcohol.

However, in practice, the alcohol present inhibits the

further oxidation of the aldehyde [5], so that, at

least in the earlier stages of reaction, the conversion:

RCH2OH +02^RCH0+H2O2

takes place more or less quantitatively.

Several examples are known too of the opposite

type of behaviour in which the overlapping of

consecutive reactions is limited by the inhibition by

the products of initial attack of the starting compound. In other words, the oxidation of the original

organic compoimd becomes self-inhibited, so that

reaction ceases while considerable amounts of both

initial reactants are still present. Self-inhibition is quite frequently observed, for instance, during the

oxidation of alkenes [6J, amines [7 J, and organic

sulphur compounds |]8, 9]. Another example is provided by cool-flames which are, by their very

nature, clearly self-inhibiting. Here the "product"

responsible for the inhibition has not been imambiguously identified but there is reason to believe

that this may be largely the heat of reaction [10]

rather than a particular chemical species.

In other cases, it may be possible to reduce the

extent of unwanted consecutive reactions by the use

of additives which cause selective inhibition of certain

stages. Very little work has been done in this field but

perhaps one example may show that the subject is at

least worthy of further attention. Thus tertiary

amines are powerful inhibitors of the oxidation of

aldehydes [11] but have little or no effect on the

oxidation of ethers under similar conditions [12].

Now clearly, if an additive can be found which

inhibits the further oxidation of an intermediate

product without affecting its rate of formation, the

isolable yields of this product will be increased.

Other additives, although not necessarily inhibiting so completely certain stages of complex

reactions nevertheless minimise the effects of

consecutive reactions by causing oxidation to take

place in well-defined macroscopic stages. A good

example is to be found in the detailed studies by

Emanuel and his co-workers of the catalysis of the

gaseous oxidation of hydrocarbons by hydrogen

bromide. In particular, investigations were made of

the reaction of equimolar mixtures of propane and

oxygen containing quite large amounts (3-20

percent) of hydrogen bromide at ca. 200 °C [13].

Under these conditions, the main product is acetone,

the production of which obeys a quasi-unimolecular

law, viz:

where At and A^ represent the concentrations of

acetone after time, t, and at the end of the reaction

respectively. This law holds with different mixtures

over a wide range of conditions, although A„ is dependent on mixture composition. In order to

explain the above relationship, it was suggested that

there was rapid formation of an intermediate

product, C, which was a powerful catalyst for the

production of acetone but itself decomposed slowly.

If the decomposition of the catalyst took place

according to a first-order law and each decomposition

initiated a non-branching chain leading to the

production of acetone, then the observed kinetics

could be accounted for. It was indeed later shown experimentally that

acetone (the product) does inhibit the formation of

the catalyst which is responsible for its production.

Thus, if acetone is added in advance in sufficient

concentration, there is indeed complete inhibition;

but if this compound is introduced after the reaction

112
has started, it has no effect. It is therefore clear that

acetone inhibits the formation of the catalyst but

does not affect its decomposition. Artificial addition

of acetone at various stages of the initiating re- action—along with determinations of the final

yields of acetone, being a measure of the amount

of catalyst formed—makes it possible for the kinetics

of the initiating reaction to be determined on their own.The oxidation of propane, catalysed by hydrogen

bromide, is therefore a reaction which can occur in

well-defined stages, and this minimises the mutual

effects which consecutive reactions normally exert on

one another. The occurrence of more or less discreet

stages of reaction is indeed confirmed by measure- ments of the accompanying temperature rise which

shows two maxima, well separated in time, implying

two different modes of heat production.

In the same way, the HBr-catalysed oxidation of

isobutane also proceeds in distinct macroscopic

stages [14]. Thus at 170 °C, the first step is the

conversion of the hydrocarbon to tert -butyl-hydroperoxide:

(CH3) 3CH +02-^ ( CH3) 3COOH

and the second step is characterised by the complete

cessation of oxygen consumption and the oxidation

of further isobutane solely by the hydroperoxide:

(CH3)3CH+(CH3)3COOH->2(CH3)3COH.

Again due to the (little understood) catalytic

influence which the primary products evidently

exert on the reactions in question, oxidation takes

place in distinct stages rather than by a series of

more closely interdependent steps.

4. Control of Concurrent Reactions

The effective control of unwanted concurrent

reactions is in general more difficult than the corre- sponding problem with consecutive reactions. In

order to discuss this question more fully, it is neces- sary to consider briefly the basic mechanism of the

gaseous oxidation of organic compounds. Although

it is difficult to make broad generalisations, the

initial attack of the organic compound at relatively

low temperatures (<400 °C) appears invariably to

involve the abstraction of hydrogen to yield a free

radical:

X+RH-^R+XH.

Now this step, although at first sight relatively

simple, may in fact be quite complex, since the

various hydrogen atoms in the organic compound

will not usually all be equivalent so that a number of

different radicals, R, may be formed concurrently.

Thus, for example, in the gaseous oxidation of

alkanes at 300-400 °C, the relative rates of attack at

primary, secondary, and tertiary centres [15, 16] are, very roughly, 1:3:15. Since, however, these figures

refer to relative rates per hydrogen atom, the

differences in vulnerability of different parts of

organic molecules may be much reduced, so that

initial attack is usually a very random process even

at the lowest convenient reaction temperatures.

One possible way of increasing the selectivity of

attack is to arrange that the species, X, which

normally reacts with a starting compound, is replaced

by some other species which attacks it with greater

specificity at the most vulnerable positions. Direct

studies of the reactions of atoms and free radicals

with hydrocarbons have shown that the selectivity

varies considerably with the nature of the attacking

species and that bromine atoms are exceptionally

selective in their attack, the relative rates of hydrogen

abstraction from primary, secondary, and tertiary

centres [17, 18] at ca. 100 °C being 1:250:6300.

Since under normal conditions of gaseous oxidation,

the hydrocarbon is probably attacked mainly by

relatively unselective species such as hydroxyl

radicals, the selectivity would be greatly increased

if these were replaced by bromine atoms. This could

perhaps be done by using hydrogen bromide as a

catalyst when as a result of the reactions:

0H+HBr^H20+Br

Br+RH->R+HBr

effectively only one radical, R, should be formed

from an organic compound containing one preferred

point of attack. Studies of the HBr-catalysed

oxidation of hydrocarbons indeed tend to show rather

high selectivity of product formation [19-21].

Even if the initial attack can be made to take

place selectively, however, the free radicals so produced may react further in a variety of ways. On

the whole it seems likely that, except at rather high

temperatures, the predominant fate of R will be to

react with oxygen. In this case, the first step must be

the formation of some loose complex R . . . . O2.

This may then decompose more or less immediately

to regenerate the original reactants or to give, when

RH is an alkane, an alkene and an HO2 radical.

Alternatively, if the energy can be rapidly distributed between the various bonds, the complex

may stabilise itself to form an RO2 or alkylperoxy

radical.

Alkene +HO2

/

R+Os^R O2

\

R02

The balance between the abstractive and additive

reactions will presumably not be affected by such

factors as mixture composition but some control

might be exerted for example by adjustment of

temperature, since the activation energy is at least

10 kcal mole~^ (42 kjoule mole~^) higher for the

abstractive reaction, the use of relatively low

temperatures thus favouring the additive reaction.

Even if the balance between the two concurrent

reactions of the radical, R, with oxygen can be

controlled, there is nevertheless a variety of possible

concurrent fates which the products of these and

subsequent reactions can imdergo. Thus there are, to say the least, formidable difficulties in controlling

113
the sequence of concurrent reactions so that oxidation

follows predominantly one particular route and

leads to the formation of only a few specific products.

5. Control of Mutual Influences of

Species Present

A number of ways have been discussed in which

it may be possible to prevent or limit the occurrence

of unwanted consecutive and concurrent stages

during the gaseous oxidation of organic compounds

and hence to exert some control on the course of the

reaction as a whole. However, it is only in a very few

cases that oxidation can be made to proceed solely

along particular desired routes and in the majority

of real systems it is necessary to investigate also the

mutual influences which the various species present

exert on the reactions of one another. Now, if a study is being made of the gaseous oxidation of a compound, A, this will in general lead to the formation of a number of products Bi, Bi, B3 . . . . , some

of which may be stable final products and others

intermediate products. Certain of these intermediates may be oxidised further quite independently

but some may react in such a way as to affect quite

profoundly other reactions in the overall scheme. It

is important, wherever possible, to discover the

precise nature and role of such intermediate products,

although in practice the required information is not

always easy to obtain. The mere identification of a particular compound among the products of oxidation of A gives little clue as to the role of this compound in the reaction.

Furthermore, the failure to detect a given compound does not prove unequivocally that it is not

formed or that it does not play an important part in

the overall reaction. Nevertheless, there are methods

for investigating the nature and role of intermediate

molecular products and the information gained

therefrom may be of considerable help in the unravelling of the imderlying mechanism.

6. Methods of Studying Nature and

Role of Intermediates

6.1. Use of Conventional Additive Experiments

One way in which it may be possible to draw useful

conclusions regarding the intermediate reaction

products is by means of experiments in which small

amounts of possible intermediates are added to the

system and their effects measured. However, experiments of this kind, unless carefully planned, are often not very informative and deductions made

from them are by no means always justified. In

general, significant information can be obtained

only if attempts are made to seek quantitative

correlations between the effect of compounds when

artifically introduced into a system and the amounts

of these same compounds normally formed in the

system.

An example of a well-designed conventional

additive experiment of this sort is provided by the

work of Burgess and Laughlin [22^, who studied the

role of heptyl hydroperoxides during the cool-flame

combustion of n-heptane and showed that the concentration of these compounds increased steadily

prior to the passage of the cool flame but fell abruptly

as the flame traversed the mixture. On the other

hand, when small quantities of 2 -heptyl hydroperoxide were added to the system, this compound

reduced the induction period preceding the cool

flame but did not affect the accompanying pressure

change unless more than a certain amount was

added. However the significant observation was

made that the minimum concentration of peroxide

required to make this occur and that normally

present at the instant of passage of the cool flame were almost exactly equal; this strongly suggests

that one or more heptyl hydroperoxides are the

degenerate branching agents during the oxidation of

n-heptane and that a critical concentration of these

is needed to enable a cool flame to pass.

6.2. Use of Isotopically-labelled Additives

Perhaps even more significant information re- garding the role of intermediates can be obtained by

the use of isotopic tracer techniques and probably one of the most important and interesting cases is that where the labelled compound is an intermediate

product of combustion of the fuel. Now suppose

that in the system:

Organic compound (x) Products (Z)

Other

intermediates

114
a small concentration of labelled Y is added. Measurements can then be made of the variation with

time of the specific activities of Y and Z and, by

quite a simple mathematical analysis, it can be

shown to what extent the conversion of X to Z takes

place via Y. This technique was first applied by

Neiman and his co-workers and, as an illustration,

mention may be made of their studies of the oxidation of methane [23], where the results show that:

100 percent of CO is formed via HCHO

25 percent of CO2 is formed via CO.

Similarly, when propylene reacts with oxygen

[24],

30 percent of CO is formed via CHO group

of CH3CHO

50 percent of CO2 is formed via CHO group

of CH3CHO

75 percent of CH3OH is formed via CH3 group

of CH3CHO

<5 percent of HCHO is formed via CH3OH.

Many of these findings are surprising and indeed

show the necessity for drastic revision of previously

proposed (and often widely accepted) reaction

mechanisms.

Another example of the use of isotopically labelled

additives in elucidating the role of intermediates in

oxidation processes is concerned with the old and

rather celebrated controversy regarding the extent

to which the combustion of hydrocarbons (and

indeed other organic compounds) proceeds via the

formation of organic peroxides. One experimental

finding which has tended to support the protagonists

of theories which ascribe to alkyl hydroperoxides,

ROOH, an important role as intermediate products

during oxidation of the parent hydrocarbon, RH, is the similarity of the products of hydroperoxide

decomposition to those of hydrocarbon oxidation.

However, in order to establish that peroxides are true intermediates, it is necessary to show that the

mechanisms of formation of given products are the

same in the two reactions. Some studies were therefore made of the oxidation of isobutane and then,

under as nearly as possible the same conditions, of

the thermal decomposition of tert-butyl hydroperoxide. One product, albeit a minor one, which is formed in both reactions is methyl ethyl ketone.

There are two likely routes for the production of

this compound from a branched C4 skeleton. It might

be formed either by further reactions of acetone

(which is an important product from both isobutane

and tertbutyl hydroperoxide):

-H +CH8

CH3COCH3 ^CHsCOCHz ^CHsCOCHzCHs

or directly from the C4 starting compoimd by some

process involving migration of a methyl group, e.g..

(CH-j)2CHCH-j —2^ (CHj)2CCHj

^

(CHj)2CCH2

0-0

1

CH3COCH2CH3 ^^^g'^tlon (CH3)2CCH2.^^ (CH3)2CCH2-

0-' 0-OH

Now, if the oxidation of isobutane is carried out

in the presence of ^^C-labelled acetone [25], the

methyl ethyl ketone formed is almost completely

inactive, i.e., it is nearly all derived from isobutane.

If, on the other hand, tert-butyl hydroperoxide is

allowed to decompose in the presence of labelled

acetone [26], the greater part of the methyl ethyl

ketone is active, i.e. it is formed via acetone. The

contrast between the two mechanisms of formation

of methyl ethyl ketone:

(CH ) CH

(CHj)jC'

(CH,)^COO- (CH^)^COOH

lntramolr>\^ acetone

rearrangement

^

addn.

CH3COCH2CH-J

115
strongly suggests that the oxidation of isobutane

does not, in the main, proceed via the intermediate

formation of tert-butyl hydroperoxide.

Somewhat more recently, isotopically-labelled

additives have been used to throw light on another

point of fairly fundamental importance regarding

the role of intermediates in hydrocarbon oxidation.

For a time it was fairly generally agreed that, whereas

at low temperatures alkyl radicals react with oxygen

to give alkylperoxy radicals:

R+02^R02,

at higher temperatures this reaction leads directly to

the conjugate alkene and an HO2 radical:

R+02-^alkene +H02.

It has lately been proposed, however, that during

the oxidation of most alkanes, even at temperatures

as low as 300 °C, the second reaction represents the

major fate of alkyl radicals and that many, if not

most, of the oxygenated products arise by the

further oxidation of the intermediately-formed

alkenes p7j. Experiments involving the addition of

small quantities of '''C-labelled 1-butene and 2-butene

during the combustion of n-butane [28] make it possible to assess quantitatively the role of these

compounds as intermediates. By radio gas chromatographic analysis of the products formed, i.e. by

determination of both the amounts of the various

products and the activity associated with them, it is possible to discover the extent to which the oxidation

of n-butane proceeds via 1-butene and 2-butene and

also which and how much of the various oxygenated

products arise from further reactions of the two

C4 alkenes.

The general conclusion is that, during the oxidation

of n-butane at about 320 °C, the butenes play an important but not overwhelming role. The results

show, for example, that at a point when the reaction

rate is beginning to decrease, at least 35 percent of

the n-butane has been oxidised to butenes and about

60 percent of the butenes has reacted further,

2-butene being in general a more important intermediate than 1-butene. However, that many

products are being formed by a mechanism not

involving butenes is shown by radio gas chromatographic analysis of the 0-heterocyclic products

formed. The C4 epoxides could plausibly be formed

by further oxidation of the alkenes (as in epoxidation

by other reagents such as peroxyacids) but it is more

difficult to see how larger ring compounds could be

formed in this way and it seems more likely that

these other 0-heterocycles are produced by some

process involving intramolecular rearrangement of

the alkylperoxy radicals. The results with isotopically-labelled additives confirm these expectations.

Thus, during the oxidation of n-butane in the

presence of labelled 1-butene, the corresponding

epoxide is radioactive, its activity indicating that

between one third and one half of it arises by further

oxidation of 1-butene. On the other hand, the larger

ring compounds are completely inactive, showing

that 1-butylperoxy radicals can be formed only by

direct addition of oxygen to 1 -butyl radicals and

are not produced at all by the 1-butene route:

CH-jCH2CH2CH2'

+0^

CHgCHgCHgCHgOOH

-OH

CH-jCH2CH2CH200'

CH-JCHCH2CH2OOH

-OH

C C H2 C H — C H2

+H0^

CHJCH2CHCH2OOH

-OH

GH2CH2CH2CH2 CHjCHCH2 CH2 CHJCH2CHCH2

The use as additives of isotopically-labelled

intermediates has also other advantages. For

example, experiments of this kind can yield values

for the separate rates of formation and destruction

of intermediates during complex oxidation processes

[29, 30]. Conventional analytical data of course provide information only about the net rates of

accumulation of such compounds.

7. Other Forms of Control

Apart from chemical control, other types of control are clearly needed if gaseous oxidation reactions

are to be successfully 'tamed' to give high yields of

desired intermediate products. Thus the interaction

of organic compounds with oxygen is nearly always an exothermic process, so that the efficient removal

of the heat of reaction may sometimes be an im116
portant factor governing the amounts of intermediates which can be isolated. In some cases, a large temperature rise may be prevented and the

yields of certain partial products increased if a

'rain' of supposedly inert solid particles is allowed to

fall through the system pi3. However, solid surfaces

constitute perhaps the most enigmatic of the many

factors which govern the course of gaseous oxidation

reactions and there seems little doubt that in most

systems the surface of the containing vessel plays an important part, perhaps causing both initiation and

termination of chains. Indeed, in certain circumstances, many of the characteristic reaction products

formed are believed to result from heterogeneous

processes [32 J, so that the above method of exercising thermal control may well have some influence

also on the imderlying chemical mechanism.

8. Conclusions

Whatever the nature of the reactants and of the

desired products and whatever the type of system in

which the required conversion is to be carried out,

effective control cannot in general be achieved unless

the reaction mechanism is fully understood. There

remains therefore a pressing need for further elucidation of the processes involved during the gaseous

oxidation of many organic compounds. Without a knowledge of the identity of the intermediate

molecular and free radical species taking part, of the

kinetic parameters of their reactions, and of the

extents to which various possible consecutive and

concurrent stages operate and influence one another,

it is difficult to devise means of encouraging certain

stages and suppressing others. It is only in this way

that further useful developments may be expected

which will make it possible to obtain high yields of

specific intermediates instead of the complexity of

products which normally results from the uncontrolled gaseous oxidation of organic compounds.

9. References

[1] Harcourt, A. V., and Esson, W., Phil. Trans. Roy. Soc. 156, 193 (1966).

[2] Chien, J. Y., J. Am. Chem. Soc. 70, 2256 (1948).

[3] Knox, J. H., Trans. Faraday Soc. 56, 1225 (1960).

[4] Cullis, C. F., Newitt,E. J., et al., Proc. Roy. Soc. A.237,

530 (1956); A.242, 516 (1957); A.257, 402 (1960);

A.264, 392 (1961).

[5] Farmer, J. B., and Macdowell, C. A., Trans. Faraday Soc. 48, 624 (1952).

[6] Cullis, C. F., Fish, A., and Turner, D. W., Proc. Roy.

Soc. A.262, 318 (1961).

[7] Cullis, C. F., and Waddington, D. J., Proc. Roy. Soc. A.244, 110 (1958); A.246, 91; A.248, 136.

[8] Cullis, C. F., and Roselaar, L. C, Trans. Faraday Soc.

55, 272 (1959).

[9] Harkness, A. C, and Murray, F. E., Atmospheric Environment 10, 245 (1966).

[10] Gray, B. F., and Yang, C. F., Trans. Faraday Soc. 65,

1603, 1614, 2133 (1969).

[11] Cullis, C. F., and Waddington, D. J., Trans. Faraday Soc.

53, 1317 (1957).

[12] Waddington, D. J., Seventh International Symposium

on Combustion, p. 165 (Butterworth's, London, 1959).

[13] Maizus, Z. K., Emanuel, N. M., et al., Dokl. Akad. Nauk.

S.S.S.R., 83, 717 (1952); 87, 241, 437, 801 (1952); 89,

1049 (1953).

[14] Babaeva, A. A., Maizus, Z. K., and Emanuel, N. M.,

Izvest. Akad. Nauk. S.S.S.R. Otdel. Khim. Nauk. 1913

(1959); Dokl. Akad. Nauk. Azerbaidzhan 15, 1009

(1959).

[15] Cullis, C. F., Hardy, F. R. F., and Turner, D. W., Proc. Roy. Soc. A.251, 265 (1959).

[16] Falconer, W. E., Knox, J. H., and Trotman -Dickenson, A. F., J. Chem. Soc. 4285 (1961).

[17] Anson, P. C, Fredericks, P. S., and Tedder, J. M., J. Chem. Soc. 918 (1959).

[18] Fettis, G. C, Knox, J. H., and Trotman-Dickenson,

A. F., J. Chem. Soc. 4177 (1960).

[19] Rust, F. F., Vaughan, W. E., et al., Industr. Engng. Chem.

41, 2595, 2597, 2604, 2609 (1949).

[20] Allen, E. R., and Tipper, C. F. H., Proc. Roy. Soc. A.258,

251 (1960).

[21] Bastow, A. W., and Cullis, C. F., unpublished work.

[22] Burgess, A. R., and Laughlin, R. G. W., Chem. Comm.

769 (1967).

[23] Neiman, M. B., Nalbandyan, A. B., Moshkina, R. I., et

al., Izvest. Akad. Nauk. S.S.S.R. Otdel. Khim. Nauk.

789 (1955); 821 (1957).

[24] Neiman, M. B., Eframov, V. Ya., and Serdyuk, N. K.,

Izvest. Akad. Nauk. S.S.S.R. Otdel. Khim. Nauk. 397

(1956); Kinetika i Kataliz. 1, 345 (1960).

[25] Cullis, C. F., Fish, A., and Trimm, D. L., Proc. Roy. Soc. A.276, 427 (1963).

[26] Cullis, C. F., Garcia-Dominguez, J. A., Kiraly, D., and

Trimm, D. L., Proc. Roy. Soc. A.291, 235 (1966).

[27] Knox, J. H., Combustion and Flame 9, 297 (1965).

[28] Berry, T., Cullis, C. F., and Trimm, D. L., Proc. Roy.

Soc. A.316, 377 (1970).

[29] Neiman, M. B., and Feklisov, G. I., Zhur. Fiz. Khim.

28, 1235, 1737 (1954); 30, 1126 (1956).

[30] Horscroft, R. C, J. Chem. Soc. 5247 (1963).

[31] Jones, J. H., Fenske, M. R.,et al., Industr. Engng. Chem.

51, 262 (1959); J. Chem. Eng. Data 6, 623 (1961); Industr. Engng. Chem. Proc. Des. Develop. 8, 17, 196

(1969).

[32] Hay, J., Knox, J. H., and Turner, J. M. C, Tenth Inter- national Symposium on Combustion, p. 331 (The Combustion Institute, 1965).

Discussion

Alvin S. Gordon (Naval Weapons Center, China

Lake, Calif.)

The use of isotopically labelled compounds to

detect reaction mechanisms has been exploited very

successfully by Dr. Cullis. One mechanism he has

put forth states that methyl groups migrate during

the course of the homogeneous gas phase reaction.

Although some of the evidence presented in defense

of the mechanism is quite persuasive, I find it very

difficult to conceive of methyl (or alkyl) radicals

migrating without the intermediary of free radicals.

This is because hydrogen atoms form a very protective shield around a C-C bond, preventing a free

radical from being able to react with the orbitals of

the C-C bond. As an example, even H atoms cannot

attack the C-C bond of a saturated hydrocarbon to

form a smaller hydrocarbon and a free radical, even though the process is quite exothermic.

117
C. F. Cullis:

I agree with Dr. Gordon that care must be

exercised when drawing detailed mechanistic deductions from the results of experiments involving

isotopically-labelled compounds. However, in some

studies of the gaseous oxidation of relatively large

hydrocarbon molecules, the only plausible conclusion

appears to be that alkyl groups do migrate, even though this may involve the intermediate formation

of free radicals. Perhaps I might mention two

examples taken from our earlier work.

The first concerns the slow combustion of

specifically -labelled 2-methylpentanes (Cullis,

Hardy and Turner, Proc. Roy. Soc. A.251: 265

(1959)). When the 4- and 5-positions are labelled with

carbon-14, the labelled atom does not appear in the

resulting methyl ethyl ketone whereas, when the 2-

and 3 -positions are labelled, the specific activities of

this product are the same as those of the original

labelled hydrocarbons. Both the 2- and 3 -carbon

atoms (but not the 4- and 5-carbon atoms) are therefore always involved in the formation of methyl

ethyl ketone and, since degradation experiments

show conclusively that the carbon atom in the

2 -position forms the carbonyl group of the ketone,

the 3 -carbon atom must form the adjacent methylene

group:

r

CH3

CH3CHCH2CH2CH3-

1 2 3 4 5

>CH3COCH2CH3 12 3 4

It seems highly unlikely that both of the two attached

methyl groups should become separated from the

tertiary carbon atom, so that the 1 -carbon atom

becomes the a-methyl group in the ketone. There is no direct evidence from the labelling experiments

that the 1 '-carbon atom becomes the /3-methyl group

in methyl ethyl ketone but the fact remains that the

I'-carbon atom is split off and somehow or other a methyl group becomes attached to the methylene

group which was originally in the 3 -position.

A second example is provided by the formation of

fert-butyl alcohol during the slow combustion of

2-methyl-but-2-ene (Cullis, Fish and Turner, Proc.

Roy. Soc. A.267, 433 (1962)). When the 2-position

is labelled with carbon-14, this appears as the

a-carbon atom in the alcohol whereas, when the

4-position is labelled, the methyl groups in the

resulting tert-hutjl alcohol have some activity. The

second finding makes it difficult to avoid the conclusion that the 4-carbon atom migrates so that it becomes directly attached to the 2-position:

GH„ C.OH

CH3

although the precise mechanism of this migration

remains a matter for conjecture.

More recently evidence has been obtained that

ethyl groups can migrate during the gaseous oxidation of 3-ethylpentane (Barat, Cullis and Pollard,

Thirteenth Symposium on Combustion, p. 179.

Pittsburgh, The Combustion Institute, 1971) but

this work did not involve the use of isotopic tracer

techniques.

A. S. Gordon (Naval Weapons Center, China Lake,

Calif.):

As a speculation of an alternative mechanism to

that proposed by Cullis, Hardy, and Turner (Proc.

Roy. Soc, (1959)), for the oxidation of 2-methyl

protane to acetone, I would like Dr. Cullis' comments on

CH- ^

c=c

CH,

H

0

:C-C-C-C

2

• ^ ^^-C-C-C + HOo

0

I 0

:C-C-C-C

intraradical

isomerization

H

I 0

c J • p^-c-c-c

H

I 0

6

-> + CHoCH=CH2

OH + Acetone

118
The mechanism postulates a six-membered ring for

internal abstraction, which should be fairly strain

free. All the energetics would be favorable. By

\ . analogy, starting with the C-C-C-C radical, one

would form ethylene and isobutyraldehyde, an isomer

of methyl ethyl ketone. Would the analyses used by

Cullis et al. be able to discriminate methyl ethyl

ketone and isobutyraldehyde?

C. F. Cullis:

The components of the reaction mixture were

identified by carrying out gas-chromatographic

analysis not only on different columns and at

different column temperatures but also after artificial

addition of each of the components believed to be

present. In addition the combined products of

several identical runs were passed through a large

preparative chromatographic column; the effluent

fractions were bubbled through Brady's solution and

the presence of individual carbonyl compounds was

confirmed from the melting points of the precipitated

2 :4-dinitrophenylhydrazones.

However, although there is no doubt whatever

about the presence of methyl ethyl ketone in the

products, one would not want to be too dogmatic

about the absence of isobutyraldehyde. At the time

when the experiments referred to were carried out,

gas-chromatographic columns were not nearly as well developed as they are now, when it is a very

easy matter to distinguish between two compounds

such as methyl ethyl ketone and isobutyraldehyde.

Questions from Audience:

Professor Cullis, please explain what you mean on

the last page of your paper by using a "rain" of

solids to control the temperature in an oxidation

reaction.

C. F. Cullis:

I understand that Professor Jones is in the

audience and if so perhaps he would like to give the

explanation. Professors Jones and Fenske at the

Pennsylvania State University first reported on

this technique.

J. H. Jones (Pennsylvania State University):

Thank you, Dr. Cullis, for allowing me to have this

pleasant task. The reaction vessel we use at Penn

State consists simply of a 26-inch length of steel

pipe 1| inches in diameter supported in a vertical

position. Vapors of hydrocarbon or of any organic

compound are fed into the pipe bottom and pass

upwardly where they meet oxygen which is added at

as many as five separate points along the length of

the reaction vessel. The oxygen added at each inlet

is completely consumed before the next inlet is reached. If no means were available to remove the

reaction heat, the mixture of hydrocarbon and

oxygen would catch on fire. Heat is removed,

however, and good temperature control is achieved

by allowing fine (300 to 800 fj.m in size) inert solids

to fall through the reaction zone. The amount of

solids required to remove the heat is small and

occupy less than 0.5 percent of the total reactor

volume. Thus, the oxidation is permitted to occur in

what is essentially an open tube. In fact, for the case of C5 to C16 hydrocarbons the results are identical to

those obtained when using an open tube and 90

percent of an inert diluent. (The large excess of

inert diluent is required in an open tube reactor in

order to maintain the reaction temperature within

reasonable control.)

As noted above, the amount of solids needed for

good temperature control is small and has no

noticeable effect on the oxidation products. However,

if the amount of solids introduced is appreciably

increased, then oxygenated compounds including

undesirable carbon oxides are favored and finally,

when an even larger excess of solids is used, the

oxidation reaction can be completely inhibited.

In an early version of our reactor fitted with a window for observation, it was noted that five

separate cool flames (one above each oxygen inlet) were present, each of which was at the uniform

reactor temperature. Thus, such a reactor is useful in

preventing the temperature from rising into the

higher range where a negative temperature effect

is observed.

The usefulness of the "rain" of solids for temperature control also works for chlorination re- actions; thus, methane can be converted directly to

carbon tetrachloride in one pass through the reactor.

If solids are not employed, then such a mixture

catches fire.

119

NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Some Current Problems in Oxidation Kinetics'

Sidney W. Benson, Chairman

Department of Thermochemistry and Chemical Kinetics, Stanford Research Institute,

Menlo Park, Cahf. 94025

Experimental data on low temperature (<150°C) and high temperature (>250 °C) oxidations are examined from the point of view of reported quantitative inconsistencies. Activation

energies for tBuOo metathesis reactions with alkanes appear to be 7 kcal/mole higher than for comparable reactions of HO2. Related isomerization reactions are examined in the light of these

differences without reaching any simple conclusions. The Russell mechanism for a 6-membered,

cyclic, transition state for termination of primary and secondary alkyl peroxy radicals is shown

to be either inconsistent with thermochemical data, or else unique to solution reactions. Addition

reactions of O3 with olefins and acetylenes are shown thermochemically to have the possibility of

following concerted and biradical pathways, respectively.

Recent data showing strong inhibition by PbO-coated surfaces of both oxidation and pyrolysis

of i—C4H10 are examined in terms of mechanism.

Key words: Acetylene reaction with O3; activation energy for oxidation; O3 addition to unsaturates; chain lengths in oxidation; elementary steps in oxidation; epoxidation; gas-phase oxidation;

kinetics; liquid-phase oxidation; mechanism; oxetane formation in oxidation; oxidation; peroxy

radicals; radical isomerization; surface effects in oxidation; termination rates; transition state for termination.

1. Introduction

In 1967 an International Symposium on Combustion, held in San Francisco |]1]' seemed to offer

for the first time a consistent qualitative chemical

scheme tying together high and low temperature

oxidation of organic molecules, and also cool flame

phenomena. At that time, however, some quantitative objections were raised to the simple scheme. It

is my purpose in the present manuscript to review

these objections, and also the relevant kinetic data,

which has been assembled since then, and to assess the current situation on oxidation mechanisms.

To facilitate the discussion, let us summarize

briefly the simple aspects of the oxidation mechanism. At temperatures below 150 °C, the oxidation

of most organic molecules will not proceed measurably in the absence of either catalysts or a radical

initiation source. With such initiation, the major,

and in some cases quantitative, product of the

reaction is the hydroperoxide. The reaction can be

represented by the exothermic equation

RH+Oz^ROsH H 20 kcal ( 1

)

with a propagation mechanism:

• This work has been supported in part by Grant No. AP 00353-06 from

the Public Health Service, Division of Air Pollution, and in part by Contract NAS7-472 with the National Aeronautics and Space Administration.

^ Figures in brackets indicate the literature references at the end of this paper.

R +02;=iR02

—1

RO2 +RH (or SH)iR02H+R (or S ) (2)

Reaction 1 in this sequence can be looked upon as a

radical recombination with no activation energy.

Below 200 °C, for most radicals, R , it is essentially

irreversible [2, 3]. Its absolute rate has been measured in the gas phase for CH3 and C2H5 radicals and

it occurs at about 1 in 100 collisions. In solution, it

is a diffusion-controlled rate constant with k (solution) ~10^ 1/mole-s.

Between 200 and 300 °C, hydrocarbon oxidations,

in the absence of initiators, show pecuHar induction

periods, surface sensitivity, and auto-catalysis.

These phenomena, while occupying much attention,

have tended to obscure the quantitative study of

oxidation in this temperature range. They have

demonstrated that in the course of oxidation,

products are accumulated which accelerate the rate

of oxidation. It is most likely that the essential

auto-catalytic product is RO2H. The half-life for its homolytic fission, ROoH^RO +H0 is about 500

seconds at 250 °C and about 2 seconds at 330 °C.

RO2H provides the "degenerate, secondary, chainbranching agent" postulated by Semenov, to

account for the "cool flame" phenomena observed

between 250 and 320 °C. It is by now well documented that these are real flames and by no means

cool. Temperature rises of between 100 and 200 °C

have been observed during cool flames. The con121
tribution of hydroperoxides to branching is obscured

by the added contribution of their breakdown

products, aldehydes. These latter can increase

oxidation rates by providing much faster H-atom

donors than alkanes in the slow propagation step

2. They can also contribute to faster initiation, but

this must be a heterogeneous process below 300 °C

and negligibly slow as a homogeneous reaction

compared to RO2H fission. It should not be over- looked that the end point of aldehyde supported

propagation (step 2) is a hydroperoxide! The

relevant mechanism can be illustrated in the case

of C2H5OOH

C2H5OOHAC2H6O + OH

C2H5O +02^CH3CHO+H02

C2H50^CH20+ CH3-^->CH300H (3)

What has been astonishing about the cool flame

regime and the related phenomena of multiple-stage

ingition (i.e., real flame) which may accompany cool

flames is the fact that cool flames are self-quenching.

The equivalent phenomenon observed in flow systems

is the negative temperature coeflicient for hydrocarbon oxidation. Above 320 °C oxidation rates in

hydrocarbon-oxygen mixtures decrease with rising

temperature. The decrease in rate between 300 and

400 °C can be as much as a factor of 20. Very similar

mechanisms for this negative temperature coefficient

have been proposed independently by Knox

and the author []3] in terms of the increasing reversibility of reaction 1 (eq. 2) above 250 °C and a competing reaction leading to more stable products:

R+Ozi^olefin +HOj

H02+RHi±H202+R (4)

In contrast to the unstable RO2H, the product of

the high temperature reaction, HO2H, is stable up

to 450 °C and doesn't serve as a new radical initiation source much below 450 °C.

Although the negative temperature coefficient of

oxidation reaction has been attributed to the

decreasing production of RO2H with increased

temperatures |^4], it is also true that the concomitant

turnover in the nature of the propagating radical

from RO2 to HO2 leads to an enormous increase in

termination rate. Where RO2 termination rate

constants are supposed to be in the range of 10^ to

10® 1/mole-sec in gas phase reactions (2tBu02 can't

terminate at all), termination of IIO2 with either

RO2 or another HO2 by simple disproportionation is expected to have kt^lO^ 1/mole-sec.

One final point should be commented upon, namely

the heterogeneity or sensitivity of the oxidation

rates to surfaces. I believe that primary, spontaneous initiation is a heterogeneous process in

almost all systems. On the other hand, above 250 °C

most radical reactions, such as fission of RO2—>R +O2

epoxidation of hydroperoxy alkyl radicals and some

metathesis reactions of HO2 or RO2 (with Ea<l6

kcal) , can be shown to be too fast for diffusion to the

surface to be competitive above a total pressure of

100 torr.

2. The Activation Energy for Chain

Propogation

There are no direct measurements of the absolute

rate constants or Arrhenius parameters for the

metathesis reaction of HO2 of RO2 radicals with

alkanes. Two types of indirect study have yielded

approximate values for these important reactions.

One is the low temperature (50 to 150 °C) oxidation

of hydrocarbons, while the other is the high temperature (440 °C) co-oxidation of CH2O with

alkanes. The latter yields values for rate of metathesis reactions of HO2 with alkanes, rather than for

RO2, but the two radicals are expected to show

similar behavior. Unfortunately, the results show

considerable disagreement.

One of the most careful quantitative studies of

low temperature hydrocarbon oxidation has been

the work of AUara et al. [5] on both the liquid and

gas phase oxidation of i-butane. They used ABN or DTBP as initiators and obtained reasonably good

mass balances on Oo consumption (90 to 97 percent)

p Based on a single propagation step, t-Bu02+i-BuH-^

f-Bu02+i-Bu and a single termination, 2f-Bu02-^

(f-BuO)2+02, they estimated from the temperature

coefficient of the rate, that Ep — ^Et = 11.3±l

kcal/mole. Et had been reported earlier in the range

7.6 to 10 kcal/mole, so that the combination gives

£^p=16±2 kcal/mole, an unexpectedly high value

for a nearly thermoneutral reaction ( Ai/p is ~2 kcal)

This value is consistent with fc(~10^ l/mole-sec,

which has been reported for sector methods in

CeHe solution [6]. On the other hand, if A;( = 10^

1/mole-sec due to contributions from other radicals,

then £Jp=12±l kcal/mole. If we assign ^p=10^^

I/mole-sec and use the reported values at 50 °C in

the pure liquid of kp/k]''^ and k\'-, then we find

£^^=14.0 kcal/mole. Because of difficulties with the

concentration dependence varying from (RH)^"

to (RH)^^ with decreasing (RH), it is not certain

how much confidence to place in these values.

In a completely different type of study, Baldwin

et al. [7j have studied the high temperature (400

to 550 °C) oxidation of aldehydes in aged, boric

acid-coated pyrex vessels. For CH2O, the mechanism

is:

CHsO +Oa^CHO+HOa

CHO+02^CO+H02

CHO +02iHCO(02) (Keq~10 atm-i)

HO2 +CH20^H202 +HCO

122
2H02^H202+02

H202+M^26h+M

HO +CH20^H0H+HC0

HCO(Oi)+CH20^HC03H+HCO (5)

Similar studies lead to values for fci^lO^~^='^

1/mole-sec for the more complex C2H5CHO oxidation

(0 = 2.303 RT in kcal/mole). The uncertainty in

Arrhenius parameters was reported as 45 ±5

kcal/mole where the AH for the reaction is 40±2

kcal/mole. For kp{EtC}iO) /kV^ they report a value

of 43 (1/mole-sec) 440 °C and fcp(EtCHO)~

109-10/9 l/mole-sec. This leads to A-( = lO^" 1/mole-sec.

Use of the more widely quoted value of kt = 10^ \ the

value kp/kY'^ = 43 and an assigned .4 -factor of

108 5 1/mole-sec yields ^^(EtCHO+HOa) = 7.5

kcal/mole. This latter reaction is exothermic by

3 kcal.

When they introduce small amounts of alkanes

into the CH20 +air mixture, they can follow the

disappearance of alkane relative to CH2O, presumably by reactions with the HO2 radical. This

leads to metathesis rates of HO2 with C2H6, CsHg and

i-C4Hio relative to CH2O at 440 °C of 0.028, 0.088,

0.155. This is an astonishingly high selectivity and

rules out the possibility that the propagating

radical for these species is an active (hence less selective) radical, such as OH or O.

If we assume that the absolute rate of HO2 +CH2O

is twice that of HO2 + C2H5CHO, or 10** ^ 1/mole-sec,

then at 440 °C the absolute rate constants for

H02+the above alkanes are, respectively, 10''%

10^^, and 10^ * 1/mole-sec. If we estimate from these

rates the primary, secondary, and tertiary H-atom

abstraction contributions by additivity, and assign

reasonable ^-factors, we can estimate the following

rate parameters for HO2 abstraction:

(primary H only)

(secondary H only)

(tertiary H only)

kpiCMs)=lO'-°-'-"'

It is hard to estimate the uncertainty in the above

activation energies, but they are probably not in excess of +3 to —1.5 kcal/mole.

Ep for the metathesis of HO2 with r-C4Hio from

these gas-phase high temperature studies is thus

7.5 kcal/mole with a possible upper limit of 10.5

kcal/mole. This is in contrast with the solution

values for f-Bu02+i-C4Hio at 50 °C for which Ep is estimated to lie in the range 12 to 16 kcal/mole.

We can only conclude that either additional steric

strain in the transition states, or solvation effects, or experimental errors, or errors in interpretation of the

mechanisms are the source of these differences. At

the moment we are inclined to favor somewhat

increased gas-phase parameters and assume that

adventitious terminations are complicating the solution results. A value of £'p(H02 +i-C4Hio) = 10

kcal/mole might not be an unreasonable compromise

of both with a slightly higher value by 1 or 2 kcal for

t-Bu02 +i-C4Hio.

3. Chain Lengths in the High

Temperature Oxidation

In examining the proposed high temperature

mechanism propagated by HO2, Knox [2] noted that

the observed, maximum rates of reaction put some

severe requirements on the propagation rate constant. Thus, Rp, the rate of reaction if HO2 is the

propagating radical, is given by:

Rp = kp{Rli){nOi)

subject to the termination step:

Rt = 2kt{llOiV

The kinetic chain length, X, is given by Rp/Rt, or: Rp kp (RH)

2k,

k%\ RH)

(HO2) 2/c,[/?p/(RH)]

At 600 °K, typical values of i?p/ (RH) at the maximum in the rate curve are about 10~^ s~^ If in the

case of CoHe we set A;p= lO*-^"^'* 1/mole-s, and

A,= 109 2 1/mole-s, then (H02)max = 10-Vfcp =

Xn,ax = fcMRH)Xl0V2A-(

If we use A-( = 10^^ 1/mole-s as suggested by

Baldwin, et al. [7], we find (at (RH) ~100 torr)

that in order for X>1, jEp<ll kcal/mole. This

suggests that in the C2H6 oxidation, either the rates

suggested by Baldwin's work have too high Arrhenius

parameters, or X<1, or HO2 is not the main propagating radical. For CsHg or i-C4Hio, there is no

problem. The rate constants for HO2 metathesis

listed are in agreement with the assumption that

X>1.

A reasonable solution to this problem is that in the

case of CoHe oxidation, the HO2 radical makes only a minor contribution to the overall rate (e.g., —^20 percent) and that a less selective radical, such as OH,

does indeed do the main propagation. We could still,

under these circumstances, observe H2O2 as a major

product if nearly every C2H5 radical produced

reacted with O2 to form HO2 +C2H4 and the HOj

produced H2O2 either by disproportionation with

itself or by abstraction from CoHe or one of the more

reactive aldehyde side products. These latter can be

produced by reactions such as:

C2Hb-)-02?=^C2H602

2C2H502^2C2H50 +O2

C2H5O +02-^CH3CHO +HO2

M+C2H50^ CH3+CH2O+M

CH3CHO+H02-^CH3CO +H2O2

CH3C0+M^ CH3+CO+M (fast)

CH3+O2 +M;:±CH302 +

123
CH3O2 +CU.Oi^CRzO +C2H5O +O2

CH3O +02-^CH20+H02

ROj +CO-^RO +CO2 (6)

The overall stoichiometry in such a scheme is such

that without considering propagation:

C2H6+t02-^2CO+3H202

The auto-catalytic nature of the reaction could

then be accounted for in terms of secondary initiation

by R02H(R = Me or Et) produced. If RO2H decomposed into RO + OH, then each OH could account

for 3H2O2 production via complete consumption of

secondary products. It does not appear possible to reconcile the very high activation energies for

Ep{JiOi) with long chain lengths in the alkane

oxidation. However, in all fairness, it must be

pointed out that X has never been measured in the

alkane oxidations!

Knox has commented [2, 14], on the fact that in

his studies there was no effect of alkane pressure on

the alkene yield, such as would be required by the

simple mechanisms, (2) and (4) [3]. However,

Knox has covered a relatively small range in conditions where such effects are easily obscured. In their

review article on oxidation. Medley and Cooley [9]

have described very dramatic studies in which at

600 K the yield of C2H4 from C2H6 went from over

80 percent to zero on raising the pressure from 7 to

140 atm. Even at 600 °C (~900 K) and 140 atm,

the C2H4 was only 46 percent. In similar fashion, at 265 °C and 150 atm pressure the yield of i-C4H8 from

i-C4Hio goes to zero and at 280 °C and 20 atm pressure, no C3H6 is detectable from C3H8+O2.

4. Isomerization of ROj Radicals

Some additional information on the activation

energy from RO2 abstraction is provided by isomerization reactions. Baldwin et al. have made direct

measurements on the rates of reaction of alkanes

introduced in small amounts (<0.1 mole percent)

into slowly reacting 2H2+O2 mixtures at 480 °C.

They follow RH disappearance and product appear- ance directly by GLC.

With neopentane, they find CH4, 3,3-dimethyl

oxetane, and i-butene, as major initial products.

Acetone, C2H4, and C3H6 are all secondary products!

There appears to be only one reasonable mechanism

to account for this behavior:

(CH3)3C(CH3) + X (CH3)3C-CH2 + XH

(X = 0,H,OH)

(CH3)3C-CH2

+0

CH

(CH3)2C

^ CH2-

2

2

.0 + OH

i-butene + CH,

(CH3)3C-CH200'

.CH.

(CH3)2C

(7)

^CHaOOH

From the initial measured rates of appearance of

oxetane and olefin, respectively, the author estimates

that feoxet.Aoief.-lO- ^ 1/mole (440 °C). But fcoief. has been measured [10] at these conditions and is

]^03 3± 3 g-i xhe half-life for oxetane under these

conditions is about 70 seconds [11], which leads us

to correct their ratio by a factor of 2 for lost oxetane

to 10-^ 1/mole. This then gives fcoxet.'^lO'^^^^^ 1/mole-s.

Now, the steady state treatment of the mechanism

gives

_

K^Kzkj

There are three limiting cases for /coxet.- If k^<k^3

and k3<k-2, then step 4 is rate limiting and

feoxet. = KiKski

If /C4>A_3 but /c3<A:-2, then step 3 is rate limiting and

fcoxet =^2^3. Finally, if kA>k-z and k3>k-2, step 2 is

rate limiting and froxet. = fc2. We can estimate by

analogy with Me and Et that fea^lO^ 1/mole-s, so that this last possibility is about 10=* too fast. From

thermochemical data [12] we can estimate K^'^W

1/mole at 750 K, so that in order for h to be rate

limiting, its rate would have to be 10^ ^ s-^ A3 can be estimated at lO^'^^ s-\ which would require

£3 = 28 kcal/mole, an impossibly high value from

any standpoint [13]. K3 can be similarly estimated

at 10-1 *^ at 750 K, so that for the remaining case we

conclude that ^4=10^^ s-^. Since A^^W'-^ s-\ this

leads to £4^21 kcal/mole with an uncertainty of

about ±3 kcal/mole.

124
The comparable 3-membered ring epoxidation of

the radical obtained from dtBP has an activation

energy estimated at less than 16 kcal/mole, but its reaction is more exothermic (17 kcal versus 12 kcal).

One further piece of evidence on Ep comes from

the relative isomerization rates of branched and

unbranched peroxy radicals. Rust observed [14]

that in the oxidation of 2 ,4-dimethylpentane in the

neat liquid at about 150 °C, a 90 percent yield of

2 ,4-dihydroperoxide was formed. The ratio of

mono- to diperoxide was about 1 : 10 and this has

been since confirmed by independent measurements

[15].

In direct contrast to this, the oxidation of n -octane

is foimd to proceed with less than 10 percent of

internal abstraction [15]. Since all peroxy radicals

of n-octane can always find a 7-H atom to abstract,

this suggests a factor of about 100, favoring internal

abstraction in the case of the branched hydrocarbon

over the unbranched. A factor of close to 10 can be

accounted for by a higher ^-factor for external

abstraction for long chain unbranched alkanes

compared to 2 ,4-dimethyl pentane. This arises from

the statistical factor that it has only 2 abstractable H

atoms per mole while n-octane has 14. But no set of

activation energies, Ep, for H-abstraction by RO2,

high or low, can reconcile these data, and it must be

assumed that at 120 to 150 °C, RO2 abstraction

from a secondary C-H position is sufficiently slow

that other H-donors provide the propagation path

and so inhibit internal abstraction.

A last point to consider in radical isomerization

has to do with the reaction 02 +R^H02+olefin.

Knox [16] has suggested that this comes from a competition of isomerization and dissociation. For

the CH3CH2O2 radical, for example, he has suggested

C2Hb +02<^CH3CH200

CH3CH20o4^CH2CH200H-^C2H4+H02 (8)

Now Baldwin et al. [7] have reported for reaction

of C2H5, n-CsH?, and t-butyl radicals with O2, almost

quantitative reaction to give olefin +HO2 with rate

constants of 8.2X10' 1/mole-s for C2H5 and 5.8X10'

for n-propyl at 440 °C. If we assume these come from

a direct metathesis reaction with an ,4 -factor of about

10^'' 1/mole-s, this yields the quite reasonable

activation energy of about 4 kcal/mole, in agreement

with other observations at lower temperatures [17].

If, however, we interpret the rate in terms of the

above isomerization and use 16 kcal/mole for the

RO2 primary H atom abstraction and a strain

energy of 7 kcal in the 5-membered ring, we find that

step 2 is rate determining with a rate constant,

^olefin, equal to Kih. Since A'l = 10"^ o+^'/f at 600 K

and with ko^W^-^-im g-i this gives A:oiefin = 10'-3+*/''. This becomes 10« « at 600 K and 10^ ^ i/mole-s at

713 K, in disagreement with the observed data. To

make the rates agree would require E2 = 25 kcal/mole

and make O2 addition to CH2CH2OOH faster than

step 2. The value £^2 = 33 would result in a rate

10^-fold too small! Another objection to the mechanism, which has already been noted, is that the

predominant mode of decomposition for (3-hydroperoxy alkyl radicals is via the epoxide, not the

alkene. More details of this mechanism have appeared earlier [4].

5. Termination Reactions of Peroxy

Radicals

Perhaps one of the largest discrepancies between

gas-phase and liquid-phase oxidation studies comes

from the termination reactions of peroxy radicals.

For t-Bu02 radicals, there is no termination in the

gas phase. Termination in solution occurs via the

trapping of a caged pair of f-BuO radicals.

2t-BuO,' ^ (t-BuOOOOt-Bu)

-1

(t-BuOOOOt-Bu)

[rO' + RO,']

cage

(9)

[ 2R0' + O9]

cage

2R0' + O2

R2O2 + O2

125
Recent ESR measurements by Ingold et al. [18]

at low temperatures have given equilibrium constants

and overall rate constants which are in good agreement with the thermochemical estimates [12 J for

steps 1, 2', and 4. Their measurements suggest that

the dissociation energy of the RO2-O2R bond is about 8 kcal/mole, rather than the value 6, estimated

from bond additivity rules [12].

This is an analogue of the 1,5-H-shift reactions

which have been much studied recently. The

termination rate constant for such a mechanism

would be given by:

From thermochemical data [11] the AS" for step a is —37.4 gibbs/mole, while the minimum entropy

loss in step b accompanying the loss of 4 hindered

rotors is 14.4 gibbs/mole and more likely 16.0

gibbs/mole [21, 22]. If step b is rate determining,

this gives a maximum ^4 -factor for kt of ^( = 10*

1/mole-s at 300 K and an expected value oi At = 10^ ^

1/mole-s.

If we accept the data on Ka which give AHa'^S

kcal/mole, then the absolute values of the reported

termination rate constants could only be reconciled

with the thermochemical estimates, if £fc < 1 kcal/mole. However, under these circumstances

k^a<kb at —40 °C and step a, not step b, becomes

rate limiting. This would also require a negative

activation energy for recombination, which has not

been observed.

A value of Eb^\ kcal/mole would be extraordinarily low for a reaction which involves a 1,5-

H-atom shift. On the other hand, the overall step is 58 kcal/mole exothermic even with (^A)02 as product, and we may have no adequate analogue for

such an energetic reaction. Perhaps the main factor

If, however, we turn to primary or secondary

peroxides, the situation is not so good. They appear

to have termination constants [19] appreciably

higher than those for tertiary peroxy radicals. Thus,

secondary radicals have /cf'^lO^ to 10^ 1/mole-sec,

while primary radicals have kt from 10' to 10*-^ 1/mole-sec. It has been assumed by many authors

that this arises from a cyclic complex transition

state originally proposed by Russell [20].

(10)

mitigating against such a reaction is that it does not

appear to occur in the gas phase. It is not out of the

question that ionic, or at least very polar, reaction

paths specific to solutions are involved in such steps,

in which case the present type of analysis is irrelevant.

Some indication in this direction is provided by the

observation [23] that the pyrolysis of (sec-BuO)2

yields about 30 percent H2 in solution-phase pyrolysis. No such pathway or product has ever been

observed in a gas-phase pyrolysis of a secondary

peroxide, and it would seem to involve a solutionfavored 6-centered transition state. Once again, the

overall reaction is very exothermic ('^60 kcal/mole)

H2 evolution had been reported [24] in the

decomposition of a-OH peroxides in solution, but

these seem more likely to follow ionic pathways than

simple peroxides or polyoxides.

6. Ozone Reactions

Although O atom and O3 reactions are not of

importance in oxidation reactions below 500 °C, they

are of interest above these temperatures (e.g., in

flames) and of special interest in the closely allied

photochemical oxidations which occur in the upper

atmosphere and in smog.

Recent studies [25] of the absolute rates of the

gas-phase reactions of O3 with C2H4 and with C2H2

have yielded the somewhat surprising results that the

latter is appreciably less reactive than the former.

The Arrhenius parameters for the bimolecular rates

are: A:(C2H4+03) =10^-=*-4''"' 1/mole-s

^(CoHa+Os) =108-s-io-8/'' 1/mole-s (11)

These parameters for the C2H4 reaction are quite

consistent with a concerted addition to form,

exothermically, a 1 ,2 ,3-trioxocyclopentane which

can then easily rearrange to form the 1 ,2 ,4-trioxide or ozonide. This has been discussed in some detail

a b

2R2CHO2 * t(R2.CH)2 0 ^ t

0* ^0

R2C

CHR,

•H •

i

RgC^O + O2 + RgCHOH

126
[26] and it has been shown that a biradical intermediate CH2-CH2OOO would require an activation

energy in excess of 17.0 kcal/mole.

If the concerted path requires formation of a semi-ion pair transition state, then it is not easily

understandable why acetylene doesn't react with

nearly the same parameters as C2H4, since it appears

to be reactive in concerted I2 or HCl addition

reactions [27]. Assuming that this is indeed the

case, however, the reported parameters can then

only be compatible with a biradical intermediate.

For the reaction

C2H2 + O3 Hc=-c

"0—0

H

We can estimate from thermochemical data [21]

that AH°<U.l kcal/mole and AS2=-23.7

gibbs/mole. The reported parameters (11) are consistent with these values, if we assume zero or small activation energy for the reverse reaction

(—a). In actual fact, the biradical may be partly

stabilized by valence bond structures, such as:

so that we have underestimated its AHf and step

( — a) may have some activation energy. If it does,

the next step, which is 1 ,2 ,3-trioxocyclopentene

formation, may be in competition with it. The final

products appear to be compatible with the trioxide

formation for C2H2, MeC2H, and MeCsMe. A

generalized mechanism is:

C2R2 + ^3

0:

R

r

0-

0.

^0.

'C=C*

' \R

—

R

(12)

stable

products

7. Effects of Surfaces on Oxidation and

Pyrolysis

A recent work of interest in connection with the

work on isomerization of peroxy radicals is that by

Niclause et al. [28] on the effect of traces (0.5 to

5.0 percent) of O2 on the oxidation of i-butane at

490 °C. These authors find that in the usual reaction

systems of quartz or pyrex there is an initial fast

reaction during which O2 is consumed and after

which the rate becomes the normal pyrolysis rate. With increasing surface, this acceleration decreases

and it also decreases on coating the surface with

PbO. With a very large surface/volume ratio and

127
PbO coating, there is attained a nearly zero O2

consumption and vanishingly small pyrolysis rate!

If we assume that the Oo-free pyrolysis is largely

homogeneous, then it appears that 1.0 percent of O2

plus enough PbO surface can destroy essentially all the hydrocarbon chains. At the surface /volume ratio

required to do this, it can be shown that CH3 radicals

will collide at the walls more readily than they will

abstract H from the i-C4Hio. H atoms, on the

contrary, still abstract more readily than they

collide with the walls. However, those radicals which

can give rise to H atoms, t-butyl, will themselves

strike the walls more readily than the competing

reaction to form i-C4H8+H. Also, under the condia tions employed the equilibrium, R +Oo5=±R02, is far

to the left (K„^10-2 torr'M so that (R )»(R02).

The only exception to this is H+02<=^H02 for which

Jteq^lO^ torr"'^ However, the formation of HO2

from H is very slow relative both to the abstraction

H+RH—>H2 +R and to collision with the walls.

fBu +O2—*^H02 +i-C4H8 is slower than wall collisions.

We see that the striking inhibition by 02 +PbO

surface of both the pyrolysis and the high temperature oxidation is consistent with a very efficient

transformation of either i-butyl or CH3 radicals into

nonradical products. Since this occurs in the presence

of about 0.5 torr O2, but not in its absence, we must

conclude that there is, in part at least, a surface

reaction forming a surface species that is both

stable at 700 K and capable of oxidizing alkyl

radicals to molecular products. This is not Pb02(s)

which is unstable relative to PbO under these

conditions, but may be the very stable Pb304. The

authors note, however, that KCl covered surfaces

have a similar behavior and conclude that it is the

destruction of HO2 or OH at the walls which leads to

the inhibition. HO2 is most probable.

Another finding of these authors is that under

these extreme inhibiting conditions, the normal

pyrolysis stoichiometry which produces about 71

percent i-C4H8+H2 and 29 percent CH4+C3H6,

changes (at 1 percent O2) to produce slightly more

of the latter (^-^40 percent) while the H2 is largely

(90 percent) converted to H2O and H2O2. The

authors have attributed this to a more selective

oxidation of H atoms or f-butyl radicals than CH3

by O2. The reaction fBu +02^i-C4H8+H02 either on the surface or in the gas phase could account for

the H2 oxidation. However, if we assume that the

observed products still arise from gas-phase radical

reactions, we must further account for a further

rapid destruction of f-butyl radicals at the surface,

since both H atoms and CH3 radicals will still form

about 50 percent each of i-butyl radicals and i-butyl

radicals in abstractions from i-C4Hio. i-butyl radicals, once formed, will split into CH3 +C3H6 more rapidly

than they will strike the surfaces.

The observation on the small shift in stoichiometry can be understood if some nonselective radical, such

as OH, is in part responsible for the observed

products and/or, in addition, i-C4H8 is more easily

destroyed than i-C4Hio in the system. The latter is to be expected since i-C4H8 has six allylic H atoms

which can be more readily abstracted by "sluggish"

radicals, such as RO2. Once methyl-allylic radicals are formed, they may be expected to imdergo

oxidation by the following sequence:

CH. CH.

C—CH2 + O2 ^ /

CH,

.CH2O2 CHgOOH

+ 0,

CHg.,

X,

CHO

:h20oh

+ OH

HOOH CH. CH„OOH / 2

CH2OOH CHaOO*

CHO

CH2 = C + OH

Y

CH2 = C

CH2O*

^CHO

+ H

^CHO

128
Despite the fact that the equiUbria in steps 1 and

3 are very unfavorable for RO2 formation (A^i'-^IO"^ torr~' at 750 K), the competing isomerization steps

(2 and 4, respectively) are as rapid as their inverse

fission steps, so that one may expect further oxidation of the methyl-allyl radicals.

This sequence is peculiar to branched allyl radicals

in contrast to unbranched, since only the former can undergo rapid internal isomerization via a 6-membered ring. CH3-CH = CH-CH2OO radicals, for

example, can only isomerize via a much less favorable

7-membered ring. The above sequence also produces

the very reactive, nonselective, OH radicals which

will react with parent i-C4Hio more rapidly than

they will diffuse to the walls and will also produce

more i-butyl than ^-butyl (3:2) radicals [293-

8. References

[1] The proceedings are summarized in 3 volumes. Advances

in Chemistry Series, 75, 76, 77, Amer. Chem. See,

Washington, D.C. (1968).

[2] J. H. Knox, Combus. Flame. 9, 297 (1965).

[3] S. W. Benson, J. Amer. Chem. Soc. 87, 972 (1965).

[4] S. W. Benson, Adv. in Chem. 76, 143 (1968).

[5] D. L. Allara, T. Mill, D. G. Hendry, and F. R. Mayo,

Adv. in Chem. 76, 40 (1968).

[6] J. R. Thomas and K. U. Ingold, Adv. in Chem., 75, 258

(1968).

[7] R. R. Baldwin, D. H. Langford, M. J. Matchan, R. W.

Walker, and D. A. Yorke, Xlllth Symp. on Comb.,

Salt Lake City, Utah (1970).

[8] R. R. Baldwin, C. J. Everett, D. E. Hopkins, and R. W.

Walker, Adv. in Chem., 76, 124 (1968).

[9] H. D. Medley and J. D. Cooley, Adv. in Petr. Chem. Refining, 3, 309 (1960).

[10] S. W. Benson and K. H. Anderson, J. Chem. Phys., 40,

3747 (1964).

[11] G. F. Choe and W. D. Walters, J. Phys. Chem., 71, 2326

(1967).

[12] S. W. Benson and R. Shaw, Organic Peroxides, Chap. 11,

edited by D. Swern, Wiley-Interscience, New York

(1970).

[13] J. H. Knox (Xlllth Int'l Symp. on Comb., Salt Lake

City, 1970) has proposed that such isomerization reac- tions have activation energies in excess of 33 kcal/mole.

Such a value would yield feoxetane = 10* ' which would be

100-fold too slow!

[14] F. F. Rust, J. Amer. Chem. Soc, 79, 4000 (1957).

[15] D. Hendry, T. Mill, and F. R. Mayo (private communication).

[16] J. H. Knox, Adv. in Chem., 76, 1 (1968).

[17] D. P. Dingledy and J. G. Calvert, J. Amer. Chem. Soc,

85, 856 (1963).

[18] K. Adamic, J. A. Howard, and K. U. Ingold, Chem.

Comm., 585 (1969); see also J. E. Bennett, D. M. Braun,

and B. Mile, Ibid, 504 (1969).

[19] K. U. Ingold, Accounts Chem. Res., 2, 1 (1969).

[20] G. A. Russell, J. Amer. Chem. Soc, 89, 4799 (1967).

[21] S. W. Benson, Thermochemical Kinetics, John Wiley and

Sons, Inc., pp. 52-54 (1968).

[22] H. E. O'Neal and S. W. Benson, J. Chem. and Eng. Data,

15, 266 (1970).

[23] R. Hiatt and S. Szilagyi, Can. J. Chem., 48, 615 (1970).

[24] L. J. Durham, C. F. Wurster, Jr., and H. S. Mosher,

J. Amer. Chem. Soc, 80, 327, 332 (1958).

[25] W. B. deMore, Int. J. Chem. Kin., 1, 209 (1969).

[26] Ref. 21, p. 173.

[27] G. R. Haugen and S. W. Benson, J. Phys. Chem., 70, 3751

(1966).

[28] M. Niclause, R. Martin, A. Combes, J. Fusy, and M.

Dzierzynski, Ind. Chim. Beige, 32 (special number.

Vol. II), 674 (1967); M. Niclause, R. Martin, A. Combes, M. Dzierzynski, Can. J. Chem., 43, 1120 (1965).

See, also, J. Fusy, R. Martin, M. Dzierzynski, and

M. Niclause, Bull. Soc. Chim. Fr., 3783 (1966); and

J. Fusy, G. Sacchi, R. Martin, A. Combes, and M.

Niclause, Compt. Rend., 261, 2223 (1965) for surface

effects in the absence of O2.

[29] N. R. Greiner, J. Chem. Phys., 53, 1070 (1970).

Discussion

R. S. Konar (Catholic University)

Question 1:

Dr. Benson has estimated the energetics of the

H-abstraction reactions by HO2 radicals from

paraffins containing primary, secondary and tertiary

H atoms, (page 120 of preprints) , and it follows that

kp/kt = 10 exp ( - 5200/i?r)

and

kp/ks = 10° -'^ exp(-2900/i?T)

where fep, ks and kt are the rate constants for the

primary, secondary and tertiary H-atom abstraction

reactions by HO2 radical respectively.

In isobutane pyrolysis, (10 to 150 torr, 497 to

582 °C, R. S. Konar Ph.D. Thesis, Cambridge

University, U.K., 1966), it has been found experimentally that

CH3(or, H) +iso-C4HioicH4(or, H2) +rer-C4H9

k p ^CH4(or, H2) +iso-C4H9

(kp/kt) = 101 .o8±o.3 exp[ ( - 5500± 1000) /i?T]

While from propane pyrolysis, (Leathard, Ph.D.

Thesis, Cambridge University, U.K. 1967; Leathard

and Purnell, Proc. Roy. Soc. London. A 305, 517,

1968) , it has been foimd that

kj,/h = W-'' exp(-2700//?T)

for CH3 or H atoms. It is evident that there is a good

agreement between our experimental results and

those of Dr. Benson's estimates. I wonder whether

the ratio (kp/ks) or (kp/kt) could well be independent

of the nature of the atom or radicals involved in the

H-abstraction reactions in thermal systems.

Question 2

In the decomposition of 'Chemically activated

isobutane' (Simons et al., J. Chem. Phys. 52, 3911,

1970), it is believed that the reactions (a) to (d)

will occur,

(a) CH2N2+/^!'^lCH2+N2, (X = 4358, or 3660 A).

(b) 1CH2+C3H8—>iso-C4H*o (vibrationally excited,

120 to 125 kcal/

mole)

129
(c) iso-C4Hfo—^CHs+Iso-CsHt

(d) iso-C4Hfo+M^iso.C4Hio+M

but not the reactions, (e) and (f),

(e) iso-C4Hfo-^ter-C4H9 +H

(f) iso-C4Hfo+02^product8.

There was incomplete product analysis, and the

authors simplified the situation by pumping out the

noncondensables at the liquid nitrogen temperature.

If (d) occurs, why not (f)? I wonder whether Dr.

Benson would make some comments on this oxidation process (f ) at 298 K.

S. W. Benson:

1. I think the striking parallel between relative

rates of attack on primary, secondary, and tertiary

H-atoms by CH3 and HO2 is just a fortuitous circumstance. Very selective radicals like I atoms and

Br atoms will show about the same behavior while

very unselective species like F, CI, OH, and NH2 will

probably show little or no diiference in activation

energy and only a statistical difference in H-factors

of 3:2:1.

2. One can estimate using RRK or RRKM theory,

the relative rates of disappearance of a chemically

activated species via each of several competing

routes. For two competing paths 1 and 2 with

Arrhenius parameters Aie"^^'^'^ and A^e"^'^"^'^, this

ratio is:

Rate (path 1) ^ Ai{E-Ei)'-^

Rate (path 2) ~ A^iE-E^y-^

where s is the effective number of internal oscillators

in the excited species.

For the paths c and e suggested by Dr. Konar,

Ec^SO kcal/mole and A,^W^-^ s-^ while Ee^9l

kcal/mole and ^^'^lO^*-^ s~^ so that path e will never be seen. The competing path F won't be seen for other

reasons, namely that its ^-factor is about 10* '' 1/mole-s so that even if it has no activation energy.

it is unlikely to compete with the much faster bimolecular step d which has a rate constant of about

10" l/mole-8.

What does surprise me is that the authors do not

report the more favored path:

g r-C4Hio—>i-C4H8+H2

which has £Jj~75 kcal/mole and ^j~10^^-* s"^

M. H. Back (University of Ottawa)

I would like to ask Professor Benson whether the

large difference in activation energy which he points

out between the hydrogen abstraction reactions of

the i-butylperoxy radical and the HO2 radical is also

observed in the termination reactions of the two

radicals. He mentioned an activation energy of about

8 to 10 kcal/mole had been measured for the ter- mination of t-butylperoxy radicals. If the HO2

radicals terminate with no activation energy then

the large difference would appear to persist in the

termination reactions. Such a difference would then

seem to have some basis in the structural or other

differences between the two radicals. On the other

hand if the activation energy for termination is the

same for both radicals then the observed differences

in the abstraction reactions can probably be attributed to experimental error or other artefact.

S. W. Benson:

The disproportionation of HO2 radicals has never been measured directly and its rate constant has been

estimated from very complex overall kinetics in the

H2O2+H2+O2 system. It appears to be similar to

the disproportionation of C2H5 radicals but about

3-fold slower. Like alkyl radical disproportionation,

it does not appear to have an activation energy.

tBu02 radicals cannot terminate in the gas phase.

In solution they form 2tBuO +O2 in a solvent cage,

and the small amount of observed termination is due

to this latter mode (eq 9). No such cage recombination can occur in the gas phase.

130
NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Kinetics and Mechanisms of Flames and Flame Suppression*

R. M. Fristrom

Applied Physics Laboratory, The Johns Hopkins University, Silver Spring, Md. 20910

A survey is given of the processes important in gase phase combustion with emphasis on

chemical kinetics. Gas phase combustion processes are important to the burning of polymers because the major

exothermic reactions occur there. The combustion of many plastics can be roughly divided into a

pyrolysis stage occurring at or close to the surface producing combustible gases and a gas phase

oxidation stage in the surrounding atmosphere where the evolved gases are burned. Since the

pyrolysis is usually endothermic, it is driven by the heat transferred from the gas phase flame by

conduction and radiation.

Reaction schemes are outlined for some typical simple gas phase flames and these are related

to the complex chemistry of the polymer combustion. A discussion is given of possible mechanisms

for gas phase flame inhibition. Sources of information on chemical kinetics are listed and a bibli- ography is given of compilations of gas phase chemical kinetic rate parameters.

Key words: Chemical kinetics; combustion, gas phase; flame kinetics; flame mechanisms; flame

suppression; inhibition; suppression.

1. Introduction

This survey presents an outline of combustion

processes with an emphasis on the kinetics of gas

phase flames and flame inhibition. Models will be

outlined and sources of transport thermodynamic

and kinetic information given. Gas phase reactions

are important to the oxidation of organic solids

because the major exothermic steps appear to be

located there. Simplistically, one may look upon a burning solid (liquid) as a fuel source for a gas phase

flame. The flame in turn feeds back thermal energy

and reactive species to induce gasification. Gas

phase combustion and condensed phase degradation

are thus strongly coupled. The elementary processes

of gas phase combustion are well understood and

much quantitative information is available. Flame

mechanisms have been intensively studied and some

of the simpler flames are quantitatively understood

[1, 2].' The surface phenomena in the oxidation of

organic materials are complex and less well understood.

2. Combustion Processes and Models

Fire is the common term for combustion, especially

when it is uncontrolled. Combustion is an exothermic

reaction, usually an oxidation, which has the ability

to propagate through space. Reaction requires that

the fuel and oxidizer be mixed on the molecular level

and that a source of ignition be present. Fires may be

* This work was supported by the National Science Foundation under

Grant GI-12.

* Figures in brackets indicate the literature references at the end of this paper.

classified according to: (1) the initial phases of the

fuel and oxidizer; (2) the character of the initial

distribution of fuel and oxidizer; and (3) the dominant processes by which mixing occurs. The propagation of such systems may be limited by reaction

rates, by mixing rates, or by the external rate of

introduction of fuel or oxidizer (fig. 1). Detonations

may be considered a special case of fire in which the

reaction processes are so rapid that the propagation

rate is supersonic relative to the cold reactants.

A number of strongly coupled chemical and

physical processes are involved. They fall into three

categories: (1) heat transfer, (2) overall mass

transfer, and (3) species transfer. Heat can be

transferred by convection, thermal conduction, and

radiation. Mass can be transferred by convection

and molecular diffusion. Species transfer, or exchange,

as distinguished from overall mass transport, can be

accomplished by turbulent transport, molecular

diffusion, and chemical reaction. Chemical reaction

furnishes the energy source in fires and is, therefore,

the ultimate driving force.

In the gas phase fires are referred to as flames.

Flames are an important component of most fire systems because the exothermic stages of reaction

are usually located in the gas phase. This seems

intuitively reasonable because mixing is most easily

accomplished in the gas phase. Therefore, unless a system is deliberately premixed, reaction will con- centrate in the gas phase or at a boundary between

two phases. A contributing factor is the slowness of

diffusion as compared with thermal conduction in

condensed phases. This sharply reduces the transport

of active species. Thermal flames are possible in such

systems, but intimate mixing is required. As a result,

131
INITIAL MIXING

INITIAL PHASE PREMIXED NON-PREMIXED

GAS-GAS LABORATORY

BUNSEN FLAME

GAS FURNACE

GAS-LIQUID MIST EXPLQSIONSIal OIL BURNER

GAS-SOLID DUST EXPLQSIONSlal COAL FURNACE

LIQUID-LIQUID METHL NITRATE'bl DECOMPOSITION FLAME

LIQUID O2 KEROSENE

ROCKET MOTOR

LIQUID-SOLID FUEL OIL-NH4NO3 EXPLOSIVE

H2O2 -POLYETHYLENE HYBRED ROCKET MOTOR

SOLID-SOLID TNTlb) GUN POWDER

la) COLLOIDAL DISPERSIONS

(b) THESE ARE SPECIAL CASES IN WHICH FUEL AND OXIDIZER ARE

GROUPS ON THE SAME MOLECULES. MANY SUCH MATERIALS CAN BE DETONEATED.

Figure 1. Examples of types of fires.

most condensed phase flames are of the decomposition type where the oxidizer and fuel occupy the

same molecule. They produce detonations when

the pressure drop from rapid volatilization of a solid

(liquid) produces supersonic flow. Most species will

be gaseous at the temperatures involved.

The combustion of liquids and solids is often

limited by the gasification mechanism. This may be:

vaporization (sublimation), or pyrolysis. Pyrolysis can result from pure thermal decomposition or be

induced by chemically active species.

The most common fuels are compounds of

hydrogen-carbon and the most common oxidizer is atmospheric oxygen. Other combinations, however,

are combustible [3, p. 29J.

2.1. Flames

Flames are gas phase chemical reactions with the

ability to propagate through space. This results

from a coupling between chemical reaction and

molecular transport. To produce gradients sufficient

for coupling, a reaction must be rapid compared

with transport processes. This is usually rapid by

laboratory standards. As a result flames are localized

in space. The reaction region is called flame front,

reaction wave, or reaction zone. There are many types of flames (fig. 2). Most,

though not all, flames occur in the gas phase. They

can be classified according to the character of the

mixing process which must take place before re- action. Thus one has: decomposition flames where

fuel occupy the same molecule;

where the fuel and oxidizer are molecular level; diffusion flames

oxidizer mix by diffusion; and

where they mix through eddy

final mixing must be on the

the oxidizer and

premixed flames

premixed on the

where fuel and

turbulent flames

difl^usivity. Since

molecular level, the reaction zones of these flames of

differing appearance involve the same basic physical

and chemical processes. They differ only in the mode

of introduction and removal of species (fig. 3).

Flames can be quantitatively described by a set of

equations [_3, (p. 74), 4], combining the mechanics

of the fluid flow with the constraints of conservation

of energy, conservation of atomic species, and

conservation of momentum, and the differential

equations describing: molecular diffusion, thermal

conduction, thermal diffusion, and chemical re- action. This set of simultaneous differential equations

form a boundary problem [4] in which the rate of

propagation (commonly called burning velocity)

plays a role of eigenvalue. In the general case the

eigenvalue is a tensor which is both space dependent

(diffusion flames), and/or time dependent (turbulent flames) . This is very difficult to visualize.

2.2. Premixed Flames

The simplest form of the flame equations is the

one-dimensional, time-independent case [3, 4]. This

quantitatively describes most premixed laboratory

flames [3]. The eigenvalue can be identified with the

propagation rate (commonly called burning velocity). These flames conform to the geometry imposed

by the flow system balancing the burning velocity

with the local component of inlet gas flow (fig. 4)

2.3. Flame Reaction Zone Structure

All of the processes important to the .general case

occur, with the exception of the time dependence,

which can become important in ignition, extinction,

and in turbulent flames.

The chemistry of flames is complex, although the

reactions are principally those of a limited number of

radical species (H, O, OH, and CH3 in the carbon

hydrogen-oxygen system).

Several simplifying approximations appear applicable to many flame reaction zones. (1) There is a dominant pair of reaction paths (one for the oxidizer,

DIFFUSION FLAMES

TURBULENT DIFFUSION

FLAMES

_j MULTIPLE L lOIFFUSION FLAMESI

REMIXED LAMINAR FLAME

CELLULAR FLAME

PREMIXED LAMINAR

FLAMES

mTURBULENT FLAME JET ROCKET MOTOR

t—"

1 TURBULENCE PERTURBED FLAME

\^0

TURBULENCE LEVEI LOW HIGH

RAM JET

PREMIXED TURBULENT

FLAMES

LAMINAR FLOW TURBULENT FLOW

FiGUKE 2. Examples offlame types.

132
Figure 3A. The structure of simpleflames—Premixed laminar flame.

Figure 3B. The structure of simple flames—Laminar diffusion flame.

133
too 300 500 700 TIME (n ssc)

(a) TURBULENT FLAME BALLS 1 ATM STOICHIOMETRIC CjHg-AIR

1000

REACTANTS

Vg INSTANTANEOUS PROPAGATION VECTOR

REACTION ZONE

(LUMINOUS)

PRODUCTS

(b) INSTANTANEOUS DIAGRAM OF REACTION REGION

, 2000 TEMPERATURE (°K) 1500 1000 500

10"

10"

u

< 1

1 1

H20

C0^\^ X.

" LUMINOUS^'^

REGION

1 1 1 / 1 1 1 1

3 2 1 0-1-2-3-4 -5 DISTANCE, Z (cm x 10^)

(c) MICROSTRUCTURE (schematic)

Figure 3C. The structure of simpleflames—Turbulent flame (premixed).

one for the fuel). (2) In the sequential reaction path

one may find several spatially separated stages

coupled principally by convection. (3) In each stage

reaction will be concentrated at the hot boundary.

(4) Propagation rate is controlled by the slowest

step in the most rapid path. Details of the arguments

leading to these conclusions may be found in [5].

3. The Thermodynamics of

Combustion

Thermodynamics deals with the equilibrium

properties. Flames are approximately adiabatic

overall, and precise thermodynamic data are available which allow reliable calculations to be made of

the temperatures and compositions of most common

flame systems. The calculations are straightforward,

but chemical complexity and the presence of radicals

in the high temperature gases preclude writing

stoichiometric relations without a prior knowledge of

the flame temperature. The calculations must be

made iteratively.

Two equivalent formulations have been developed

for calculations: the equilibrium constant method,

and the free energy minimization method. The

former is found in many texts [6J. The latter is a recent development. Machine programs are available

for both methods [1, 8]. The most reliable source of

data for such calculations is the JANAF Tables [9].

Machines make these tedious calculations economically in a few seconds.

The insight provided by thermodynamic calculations is useful for predicting combustion gas

properties, evaluating fire hazards [^10 J, and in

simple flame theories [5]]. Nonequilibrium systems

often can be treated by truncation to calculate

partial or frozen equilibria [2, p. 34j, in which only

certain reactions are allowed to go to completion.

Such calculations are useful in estimating the

properties of systems, such as turbulent flames,

rocket nozzles, and flame reactions.

4. Transport Properties

An understanding of fire systems requires a quantitative knowledge of transport processes.

These are: thermal conduction, the transport of heat

by temperature gradients; and diffusion, the transport of matter by temperature gradients and the

second order processes, thermal diffusion, and

diffusion thermal effects.

Viscosity is unimportant in flames (but important

in detonations); thermal diffusion is important for

very light species (H, H2 and He), and the diffusion

thermal effect is negligible. A discussion of these

processes and compilation of data for use in gas

phase flame calculations can be found in [2 and 11].

In condensed phases diffusion is negligible on the

time scale involved in fires and by comparison with

thermal conduction. Collections of evaluated transport information for condensed phases can be

obtained from [12].

134
5. Combustion Kinetics perature dependence. They are often fitted to a power law:

The energy source of combustion systems is chemical reaction. The time history and characteristics of flames and fires are controlled by the

kinetics of the reactions. This branch of chemistry

involves the study of the rates and detailed paths by

which chemical reactions occur. The discussion can be conveniently divided into two parts: the mechanisms or qualitative description, and the elementary

rate constants which provide the quantitative data.

Mechanisms can be given on three levels: the overall

stoichiometric description, often called global re- action; the elementary reaction scheme; and the

chemical dynamic level which describes the detailed

internal processes and energy levels involved in

reaction. At present the first two levels are the most

useful for combustion systems. The goal of any

mechanism is to provide a quantitative description

of a reacting system. The difference between these

three levels can be appreciated by considering the

hydrogen-oxygen flame reaction.

The "global" or stoichiometric description indicates that two molecules of hydrogen and one

molecule of oxygen are converted into two molecules

of water, and the rate is a function of the concentrations and temperature.

i?m+o.=/(H2, O2, T) e.g., -H?Otr<' (1)

This is an empirical result and should not be extrapolated beyond its region of validation.

The elementary reaction formulation uses a detailed mechanism of elementary steps. For

example, in combustion systems molecular hydrogen

does not react directly with molecular oxygen. For

the hydrogen-oxygen system the accepted mechanism is a set of six forward and reverse reactions

(table 2). This requires six reaction rate parameters

and six equilibrium constants for the reverse re- actions.

Rm+o.=f(h ...h;K,...K,;N^... iV„; T)

=f{Ri . . . (2)

where

R = Rate of Reaction

k = Rate Constant

K= Equilibrium Constant

N= Concentration

A; = Indices

The virtue of this formulation is that the functional forms are known (see below) Since this

^ Bimolecular chemical kinetic rate constants are normally

given in the Arrhenius form:

fc.7 = Ai,Tlfexp (-Eij/*RT)

Because of the exponential relationship dependence of reaction rate on temperature, most experimental data is not of sufficient precision to allow the reliable determination of b and

it is often set arbitrarily at zero. Termolecular reactions show a much lower negative temThese rates show a cubic dependence on pressure. In these

equations: A = Frequency Factor; E = Activation Energy;

k = Rate Constant; m = Temperature Exponent; *R =

Molar Gas Constant; T = Temperature.

description is not empirical, it can safely be extrapolated outside the temperature or time area of study.

The difficulty with the method is that many more

variables (concentrations) are involved, which often

are implicit functions not only of other species, but

also time and temperature. In many cases parameters

for the steps may not be available. The complexity

of such a description usually exceeds the empirical

global formulation. In favorable cases it is possible

to reduce this complexity [5], but this is an art. Often the description is amenable only to machine

processing. This should not always be a deterrent

since the complexity is simply one of numbers of

items, and the capacity of modern computers has

been doubling every eighteen months for the last

two decades and shows every prospect of continuing

this trend [13]. Therefore, the elementary reaction

attack on flame problems should ultimately prove

the most fruitful approach.

The chemical dynamic level involves the specification of the reaction rate or cross-sections for the

individual states of the molecules involved. This

detail is available for very few systems. An example

is the reaction of hydrogen atoms with molecular

bromide [M], where the distribution between the

various excited states is known.

fc[H+Br2] = fci[H+Br2^HBr,<4+Br*(2Pi/2)]

+A-2[H+Br2^HBr.<6+Br*(2P3/2)]+ ... (3)

This detail is available for so few systems that at

present its utility is severely limited. As such

studies become more common they may provide

considerable information on nonequilibrium radiation. The description is even more complex than the

elementary reaction schemes since each individual

state of a molecule (vibrational, rotational, and

electronic) is considered a separate species. It seems

probable that by the time sufficient data are available

for use in practical systems, the data handling

capabilities will have expanded sufficiently to

accommodate the information. The practical utilization of such detail appears to be a number of decades

away.

There are three aspects of flames: ignition,

extinction, and propagation. These areas share a common chemistry since the set of elementary

reactions required for a complete description depends only on the species involved. In practice,

when one wishes to select a quantitative mechanism

with a minimal number of reactions, the reaction

conditions differ sufficiently so that it is convenient

to discuss their mechanisms separately. The complexity and availability of information varies from

area to area so the coverage will be rather unequal.

The elementary rate data are applicable to all of the

135
areas so they will be discussed separately from the

mechanisms.

Flames can be ignited by thermal means and by

active species transfer. In thermal ignition a critical

volume of gas is raised to the point of thermal

explosion ^15] where heat is generated more rapidly

than it can be dissipated. This may be done by

external heating by a hot surface or by internally

generated heat. The latter process is known as spontaneous ignition. Since the system passes

continuously from cold gas to flame temperature

with many intermediates, the chemistry is complex

and time dependent. The hydrogen-oxygen thermal

ignition is moderately well understood [15], but

other systems are more complex and less well understood. In the case of carbon-hydrogen-oxygen

system oxygenated intermediates and peroxy radicals

play important roles. The area of hydrocarbon

oxidation has been surveyed [16^. Ignition by

transfer of active species occurs in spark ignition or

direct contact with a pilot flame. In both cases high

temperatures and radical concentrations are involved. This chemistry is similar to that of propagating flames discussed below. Low temperature

ignition by radicals is probably possible, but is a less

common occurrence. Propagating flames involve principally the high

temperature chemistry of a restricted number of

radical species. In the carbon -hydrogen -oxygen

system these are H, O, OH, and CH3. Flames involving the chemistries of compounds containing

nitrogen, sulfur, boron, etc., have received less

study.

The chemistry of flame extinction involves

reactions which reduce the concentrations of active

radical species. Active radical concentration can be

reduced by three body recombination, by surface

recombination, or by substitution of an unreactive

for an active radical species. This is usually accomplished by the addition of particles (for surface

recombination) or by the addition of foreign species

which promote recombination or substitution. The

particles are generally alkali metal salts which may

promote heterogeneous reactions. Foreign species are commonly halogen or halogen containing compounds so that hydrogen-halogen chemistry is important. This area is not completely understood

although a number of important reactions have been

identified [17].

5.1. Reaction Mechanisms

The best understood flame chemistry is that of

steady state premixed flames of the carbon-hydrogenoxygen system. This involves the high temperature

reaction of H, O, OH, and CH3. The system is most

easily visualized as a hierarchy of reactions. Since

oxygen is a divalent radical, one must consider the

chemistry of both oxygen atoms and hydroxyl

radicals. The high temperatures in flames appear to

make ozone, hydrogen peroxide and HO2 radicals

quantitatively unimportant. In addition, hydrogen

atoms and methyl radicals must also be considered.

Present evidence indicates that hydrocarbon and

peroxy radicals, with the exception of CH3, are thermally unstable under flame conditions. Hydrocarbon radicals can be important in low temperature

oxidations and ignitions; and HO2 radical is important in ignition and perhaps in low temperature

hydrogen flames.

5.2. The Ozone Decomposition Flame

This is the simplest flame which can be realized in

the laboratory. Three species are involved and the

reactions can be written by inspection (table I).

Because of its simplicity it is the favorite of the

theoreticians. There is reasonable agreement between measured burning velocities and computations

based on elementary rates.

Table I. Ozoneflame

[See Table VIII, #3, 15, 17, 20, for sources of kinetic

constants]

O3 + M* O + O2 + M

0 + 03^ 202

O + O + M O2 + M*

5.3. The Hydrogen Oxygen Flame

Eight reactions are required to describe this

intensively studied flame. They can be written by

inspection and most of them have received some

study. The mechanism assumes that H2O2, O3 and

HO2 are quantitatively absent. This would be

expected from their low thermal stability. Detailed

flame studies have confirmed this assumption except

in the case of certain low temperature, high pressure

flames where HO2 reactions become important

[18] (table II).

Table II. Hydrogen oxygen flame

[See Table VIII, #3, 17, for sources of kinetic constants]

H + O2 OH + O

H + H2O f± H2 + OH

O + H2 ?^ OH + H H+O+M^OH+M*

H + OH + M <^ H2O + M*

O + H2O ?± 2 OH H+H+M^=iH2 + M* 0+0+M^02+M*

5.4. The Carbon Monoxide Flame

A key peculiarity of this flame system is its negligible propagation rate in the absence of traces

of a hydrogen containing compound. This catalysis

is independent of the molecular form, water being as effective as hydrogen in trace amounts [2]. This

implies that reaction of CO with O is slow under

flame conditions and the principle attack of CO must

be by hydroxyl radicals. Thus a scheme for the carbon

monoxide flame should include the reactions of the

136
Table III. Carbon monoxideflame

[See Table VIII, #2, 3, 17, for sources of kinetic constants]

CO + OH <^ CO2 + H

hydrogen oxygen system and an additional reaction

for the attack of CO (table III).

5.5. Hydrocarbon Flames

The chemistry of hydrocarbon flames is complex,

but the general outline seems clear. Complications

include the differing chemistries of fuel lean and fuel

rich systems; the effect of changing chain length in an homologous species; and the effect of unsaturation. Saturated hydrocarbons show substantially

similar burning velocity curves if compositions are expressed in terms of the stoichiometry. As might be

expected agreement is poor for methane and highly

branched compounds which have rather different

C-H ratios. Unsaturated compounds burn more

rapidly than their saturated analogs, but the effect

is less pronounced in the higher weight molecules.

In fuel lean systems one finds as intermediates only

hydrocarbons which can be obtained from the parent

fuel molecule by degradation. Fuel rich flames by

contrast show intermediates higher in molecular

weight than the parent fuel. This is, no doubt,

connected with reactions of methyl radical which is stable in such environments.

A correlation between hydrogen atom concentration and burning velocity has been discussed at

length in the literature [19J. It appears to apply to

hydrogen, carbon monoxide and saturated hydrocarbon flames and less well to olefin and acetylenic

flames [17]. An interpretation of these facts which

appears reasonable is that the reaction of oxygen,

which is common to all of these systems, is principally with H atoms [2, 20J. This is the principle

branching step which furnishes the radicals which are necessary for flame propagation, and it is the

limiting slow step which controls the propagation

rate [2, 5].

In fuel lean hydrocarbon flames the initial attack

of the molecule appears to be by hydroxyl radical

forming water and the corresponding radical.

Hydrocarbon radicals higher tlian methyl are thermally unstable under flame conditions and

appear to undergo unimolecular decomposition

rather than radical attack. Two modes of decomposition occur: a fission into a radical and an olefin

and splitting to form methyl radical and an olefin.

This breakdown continues so that final oxidation is principally of ethylene and methyl radicals.

Fuel Lean Hydrocarbon Attack [See table VIII,

#22 for sources of kinetic constants]

C„H2,,+2+OH—>H20+ C„H2„+i

CnH^n+l—>'CH3 +C„_iH2n_2

"C„H2n+l—>"CH2m+l+Cn—mH, in —m)

In fuel rich hydrocarbons hydrogen atoms are the

dominant radical and attack the hydrocarbon. The

formation of hydrocarbons higher in molecular

weight than the fuel are observed in such flames.

This is attributed to the low concentrations of

oxidizing species and the concentration of methyl

radicals can be quite high and additional reactions can take place.

Fuel Rich Hydrocarbon Attack [See table VIII,

#22 for sources of kinetic constants]

C„H2n+2 +H—>H2+ •C„H2„+l

Olefins undergo the same radical attack sustained

by saturated hydrocarbons, but in addition they can be attacked at the double bond by oxygen atoms. In

ethylene flames this appears to result in the formation

of formaldehyde which in turn is rapidly attacked by

hydroxyl radicals or oxygen atoms with the formation of carbon monoxide.

Acetylenic compounds also can be attacked by

oxygen atoms at the site of the triple bonds; again in

acetylene flames some formaldehyde is formed.

The mechanism of these reactions is not presently

clear.

Unsaturated Hydrocarbons [See table VIII, #23

for sources of kinetic constants]

H, 0H+C„H2«—>H2, H2O C„H2„_i

O +C„H2,i—>CH20 +C„H2„_2

H, OH+C„H2„_2^H2, H2O+ C„H2„_3

5.6. The Inhibition of Flames by Halogen

Compounds

Flames and fires can be inhibited or extinguished

by the introduction of halogen containing compounds. The effect of these species is so pronounced

that the action is assumed to be chemical rather than

physical [17]. It appears that the principal action is the modification of radical concentrations and/or

distribution in the reaction zone. This interferes with

the propagating and branching steps. Halogen compounds appear to act by substituting a radical of

low reactivity for the propagating radicals.

Typical inhibitors are the halogens, or halogen

substituted hydrogen or carbon compounds. Corrosion and toxicity considerations limit practical

extinguishments to halogenated hydrocarbons (e.g.,

CFsBr). The chemistry of halogenated hydrocarbons

includes that of the halogens and halogen acids.

The effectiveness of these inhibitors appears to be

in the order of decreasing reactivity of the atom

species, F<Cl<Br<I, and is proportional to the

number of halogen atoms in the inhibitor molecule.

Fluorocarbon radicals also appear to function as unreactive radical exchange species.

All of the halogens, with the exception of iodine, can sustain hydrogen flames. The reactions for these

systems are similar and can be written by inspection

(table IV).

Halogen molecules and the halogen acids (excepting fluorine) react rapidly with both O atoms and

OH radicals with the formation of the corresponding

137
Figure 4. Burning velocity.

The rate at which a plane flame front advances into the unburned gas is called burning velocity. It depends principally on initial composition,

temperature and pressure. At any point along a flame front the propagation velocity balances the cold gas approach velocity. For conical bunsen

flames there is a simple relation between cone angle, gas velocity and

burning velocity. (See above).

Table IV. Hydrogen -halogen flame reactions

[See table VIII, # 18, 22, 23, for sources of kinetic constants]

H + X2 ?^ HX + X

X + H2 HX + X

X+X+M<=±X2 + M* H+H+MfiHj + M*

H+H+M?=±H2 + M* H+X+M<=iHX+M*

halogen atoms. In the case of Br and I the atoms are unreactive with respect to hydrogen, molecular

oxygen and the common hydrocarbon fuels. The

result is the substitution of a stable unreactive

radical. CI atom reactivity is sufficient so that

inhibition is less pronounced and in some cases a promotion of reaction can occur (table V).

Table V. Reactions of halogen species with oxygen atoms and

hydroxyl radicals

[See Table VIII, # 18, 23, for sources of kinetic constants]

0 + X2 OX + X*

OH + X2 HOX + X (OX + HX)t

0 + HX OH + X (OX + H)t

OH + HX H2O + X

t Thermodynamically unfavorable.

Halogenated hydrocarbons are rapidly attacked

under flame conditions to form halogen acids and

atoms. This allows secondary inhibition reactions

described above to take place (table VI).

Table VI. Reactions of halogenated hydrocarbons with flame

Radicals

[See Table VIII, #18, 23, for sources of kinetic constants]

R- X + H ^ HX + RR- X + O ^ OX + RTable VII. Groups and programs concerned with the gathering

and evaluation of chemical kinetic data

(1) CODATA (Committee on Data for Science and Technology). Dr. C. Schafer, Westend Strasse 19, Frankfurt/

Main, Germany. This organization was created by the

International Council of Scientific Unions. A Task Group

on Kinetics was charged with finding out what is being

compiled. The group is headed by Dr. S. W. Benson,

Stanford Research Institute, Stanford, Cal.

(2) (U.K.) Office of Scientific and Technical Information.

(3) (USSR) Commission on Compilation of Rate Coefficients. Organized February 1970 by Prof. V. N. Kondratiev to

collect rate coefficients. A relevant bulletin will be published annually. Eleven Working Groups have been established: 1) Exchange, combination, and disintegration

reactions of simple molecules (Chairman: A. I. Porojkova); 2) Monomolecular reactions (V. I. Vedeneev); 3) Reactions of radioactively contaminated particles (G. V. Karacbevtsev); 4) Hemolytic liquid-phase

reactions (E. T. Denisov); 5) Heterolytic reactions

(M. I. Vinnik); 6) Topochemical reactions (O. V. Krylov); 7) Reactions of solid inorganic compounds; 8) Reactions of solid organic compounds (Ya. S. Lebedev);

9) Polymerization reactions (A. A. Berlin); 10) Energy

exchange processes (A. M. Tchaijkin); 11) Theoretical

Group (E. E. Nikitin).

(4) National Acaderrxy of Sciences (U.S.) Committee on Data for Science and Technology, F. D. Rossini, Chairman; Ad Hoc Committee on Chemical Kinetics, S. W.

Benson.

(5) JANAF Committee on Thermodynamics and Chemical Kinetics (formerly ICRPG Thermochemistry

Working Group), J. Massey, Chairman.

(6) Landolt-Bornstein Tabellen, Technische Hochschule,

Institut fur Technische Physik, 61 Darmstat, Germany.

Dr. K.-H. Hellwege, General Editor.

(7) Tables de Constantes Selectionnees, 250 rue St. Jacques, Paris F^, France, Dr. P. Khodadd, Director.

(8) Atomic Collision Information Center, Joint Insti- tute for Laboratory Astrophysics, University of Colorado,

Boulder, Colorado, U.S.A. Dr. L. J. Kieffer, Director.

Collects, evaluates and compiles data on collisions of

electrons, and photons, with ions, atoms and molecules.

The center is part of the U.S. National Standard Refer- ence Data Systems. Ion-Molecule reaction rate data are

included.

(9) Atomic and Molecular Processes Information

Center, Oak Ridge National Laboratory, Oak Ridge,

Tennessee. Dr. C. F. Barnett, Director. Directory of

Federally Supported Information Analysis Centers

(Committee on Scientific and Technical Information,

Washington, D.C., COSATI-70-1, PB 189300 (1970),

No. 7.

(10) Chemical Kinetics Information Center, National

Bureau of Standards, Washington, D. C, U.S.A. Dr.

David Garvin, Director. Collects and indexes published

reactions in the gaseous, liquid and solid phases, photochemistry and inelastic scattering. It is part of the U.S.

National Standard Reference Data System. Prepares

bibliographies for authors of reviews in the NSRDS

critical review series and for the scientific public, both on

request and on its own initiative. Evaluations of rate data and tables of data occasionally issued.

(11) High Temperature Reaction Rate Data Centre,

Department of Physical Chemistry, Leeds University,

Leeds, England, Dr. D. L. Baulch, Director. Evaluates

rate data for gas phase reactions that are of importance

at high temperature. Publishes results as "High Temperature Reaction Rate Data Reports". Nos. 1-5 issued as of 1970.

138
Table VIII. Sources of gas phase chemical kinetic data

1. Anon. "Bibliography of Atomic and Molecular Processes,"

issued by the Atomic and Molecular Processes Information Center, Oak Ridge National Laboratory, Oak

Ridge, Tennessee. Bibliographies listing papers and re- ports during stated time periods, issued semiannually

starting with the period July-December 1965. Available

from the Clearinghouse for Scientific and Technical

Information, Springfield, Va., USA. Hard copy $3.00,

microfiche $0.65.

2. Bahn, G. S. Pyrodynamics—(Chemical Kinetics Section).

4 371-81 (1967) Reactions in the F-Cl-(H-O-N) System;

5 49-74 (1967) Detailed Compilation of Reaction Rate

Information for the H-O-N System; 5 221-72 (1967)

Reactions Involving O and H (271-65); Reactions In- volving N and H (265-72); 6 Reactions in the C-O-H

System; Reactions for the Conversion of CO and at CO

and CO2. Western States Section of the Combustion

Institute Paper 67-26, "Reactions of the B-O System for Dry Air," (1967). American Institute of Aeronautics

and Astronautics Paper 67-182, "Reactions in the Al-HCl-F and B-H-Cl-F Systems," (1968).

3. Baulch, D. L., Drysdale, D. D., Lloyd, A. C, No. 1 (May

1968), High Temperature Reaction Rate Data Series

(issued by the High Temperature Reaction Rate Data

Centre, Leeds University, Leeds, England). C-H-O-N

reactions.

4. Barnet, C. F., Ray, J. A., Thompson, J. C, Atomic and

Molecular Collision Cross-Sections of Interest in Controlled Thermonuclear Research, ORNL-3113, Revised, Oak

Ridge National Laboratory, Oak Ridge, Tenn., August

1964.

5. Benson, S. W., O'Neal, H. E., "Kinetic Data on Gas Phase

Unimolecular Reactions," NSRDS-NBS-21, xvi, 626

pp, (1970).

6. Bonner, B. H., Reaction Rate Data for Some Nitrogen/Oxygen Species, Rocket Propulsion Establishment Tech.

Memo 425, Westcott, England, December 1966.

7. Bortner, M. H., Chemical Kinetics in a Re-Entry Flow

Field, Report No. Space R63SD63, General Electric

Company, Missile and Space Division, Sciences Laboratory, King-of-Prussia, Pennsylvania, August 1963.

8. Brocklenhurst, B., Jennings, K. R., "Reactions of Nitrogen Atoms in the Gas Phase," Porter, G. (Ed.), Progress

in Reaction Kinetics, Vol. 4, 1967, Pergamon Press, Pg.

3.

9. Camicom, Mina L., Reaction Rates for High -Temperature

Air, SC-R-66-885, Aero-Thermodynamics, Monograph,

Sandia Corporation, May 1966, 58 pp, 13.00 from Clearinghouse, National Bureau of Standards, Springfield,

Va.

10. Cherry, S. S., Gold, P. I., and Van Nice, L. J., "Phase I Final Report Screening of Reaction Rates," TRW Systems, Redondo Beach, California, Report No.

0883206001-TOOOO (1967).

11. Fenimore, C. P., Chemistry in Premixed Flames, Pergamon

Press, N. Y., (1964).

12. Fristrom, R. M., Westenberg, A. A., Flame Structure,

McGraw-Hill Inc., N. Y., (1965).

5.7. Fluorocarbon Radicals

Mixed fluorocarbon -halogen molecules (F2 and

HF do not inhibit flames) are better inhibitors than

would be predicted on the basis of the halogen atom

concentration. This is probably associated with the

stability of the fluorocarbon radicals under flame

conditions. Some plausible reactions can be written,

but little reliable kinetic information is available.

The principle reactions should be with hydrogen

atoms. Reactions with oxygen atoms are thermodynamically favorable, but sterically difficult.

CF2 has been observed to react with molecular

13. Grewer, T., Kuchler, L., Zeininger, H., (Editors), "Reaktions-geschwindigkeiten," Landolt-Bornstein Tabellen,

Zahlenwerte und Funktionen aus Physik-ChemieAstronomie-Geophysik und Technik, 6th Ed. II Band.

Eigenschaften der Materia in Ihren Aggregatzustanden,

5, Teil, Bandteil b. Transportphanomene-Kinetik

Homogen Gasgleichgewichte," (Springer-Verlag, Berlin), (1968). Tabulation of rate constants, mainly for atom-molecule reactions.

14. Hochstim, A. R., Ed., Berman, M., Gilardi, R. D., Goel,

N. S., Povard, V., Riegler, G. R. Bibliography of Chemical Kinetics and Collision Processes, Plenum Pub.

Corp., New York.

15. Johnston, H. S., "Gas Phase Reaction Kinetics of Neutral

Oxygen Species," NSRD-NBS-20, vi, 49 pp, (1968).

16. Jost, W., Ed., Low Temperature Oxidation, Gordon and

Breach, New York, (1965).

17. KaufTman, F., "Oxygen Atom Chemistry," Progress in Reaction Kinetics, Vol. 1, Porter, G. (Ed.), (1961),

Pergamon Press, New York.

18. Kondratiev, V. N., "Konstanty Skorosti Gasofaznykh

Reaktsij Spravochnik," (Handbook of Kinetic Constants of Gaseous Reactions). Izdatel'stvo "Nauka,"

Moskva, 1970). Extensive summary of gas phase

kinetics.

19. Minkoff, G. O., Tipper, C. F. A., The Chemistry of Combustion Reactions, Butterworths, Washington, D. C,

(1962).

20. Schofield, K., "An Evaluation of Kinetic Rate Data for Reactions of Neutrals of Atmospheric Interest," Planet.

Space Science 15, 643, (1967).

21. Tedder, J. M., Walton, J. C, "Reactions of Halogenomethyl Radicals," Porter, G. (Ed.), Progress in Reaction

Kinetics, Vol. 4, 1967, Pergamon Press, p. 37.

22. Thrush, B., "Reactions of Hydrogen Atoms in the Gas

Phase," Porter, G. (Ed.), Progress in Reaction Kinetics,

Vol. 3, 1965, Pergamon Press, p. 63. 23. Trotman -Dickenson, A. F., and Milne, G. S., "Tables of

Bimolecular Gas Phase Reactions," NSRDS-NBS-9,

vi, 129 pp. (1967). (Revised 1970).

24. Tunder, R., Mayer, S., Cook, E., and Schieler, L., Compilation of Reaction Rate Data for Nonequilibrium Perfor- mance and Re-Entry Calculation Programs, Aerospace

Corporation, Los Angeles, Cal., Jan. 1967, 71 pp. Air Force Aerospace SSD TR67-45, TRIOOI, (9210-02)-l.

Calculations based on BEBO Method.

25. Westley, F., "A bibliography of Kinetic Data on Gas

Phase Reactions of Nitrogen, Oxygen and the Nitrogen

Oxides." NBS Report 10241 (May 1970).

26. Wilson, W. E., Jr., Kinetic Tables for the Reactions N +

O2 - NO + O, CH4 + OH - CH3 + OH2 and CO +

OH - CO2 + H. ICRPG Thermochemistry Working

Group, Fifth Meeting Bulletin, Vol. 1, March 15-18,

1967, CPIA Supplication # 146, Vol. 1, p. 147. (Revised

for NSRDS 1970).

27. Wray, K. L., "Chemical Kinetics of High Temperature

Air," Hypersonic Flow Research 7, 191 (1962).

oxygen in a shock tube at 2000 K pi], but the

elementary steps are not clear.

Thermodynamically Favorable Reactions of CF3 and CF2

Radicals in Flames

CF3 + X + H ^ HX + CFs

CF3 + H ^ CF2 + HF

CF2 + H ^ CF + H

CF2 + O CO + 2F

CF2 + O2 CO + F2 + O

139
5.8. Sources of Quantitative Chemical Kinetic

Data

In recent years a number of groups and individuals

in this country and abroad have undertaken the

compilation and evaluation of kinetic data (table

VII) . The results have been a substantial number of

bibliographic and critical evaluations (table VIII).

Two recent extensive compilations have been made

by Trottman-Dickenson and Milne (Table VIII,

#23), and Kondratiev (table VIII, #18). These

volumes present the gas phase chemical kinetic

literature in a compact form.

Extensive critical evaluations have been made by

Schofield, the High Temperature Kinetics Group in

Great Britain, and the National Standard Reference

Group in the U. S. (table VIII, #3, 5, 15, 20). These

publications offer critically evaluated kinetic data on

many flame reactions. References to sources of

information in this table have been given following

the various reactions discussed throughout this

paper.

5.9. Summary

The chemistry of flames is a complex subject, but

progress has been made in the understanding of this

phenomena in terms of elementary processes. The

quantitative utilization of this understanding will

require the accumulation and evaluation of much

more information than exists at present, but a substantial start has been made.

The author would like to thank Dr. David Garvin,

Director of the Chemical Kinetics Program of the

National Standards Reference Data Program of the

National Bureau of Standards for bringing the sources of chemical kinetic information to his attention, and

providing a bibliography of recent reaction studies.

He would also like to thank Dr. John Herron of the

National Bureau of Standards for providing information on O atom reactions. The assistance of

Mr. Mike Robison and Mrs. Nell Blake is gratefully

acknowledged. Thanks are also due to the very

capable APL Art Group, under the direction of Mr.

Emanuel Sultana, for the preparation of the figures.

6. References

[1] C. p. Fenimore, Chemistry in Premixed Flames (Pergamon Press, Maemillan, N.Y., 1964).

[2] R. M. Fristrom, The Chemistry of Flames, Sec. 2 of

Dr. R. S. Konar (Catholic University)

We face considerable difficulties in detecting and

estimating HCHO quantitatively when it is found to

be a product in the gas phase reactions. I would like

to know in detail how Dr. Fristrom detects and

estimates quantitatively formaldehyde in flames.

Survey of Progress in Chemistry, (Ed. A. Scott), pp. 55-85 (Academic Press, N.Y., 1966).

[3] R. M. Fristrom and A. A. Westenberg, Flame Structure (McGraw Hill Book Co., N.Y., 1965).

[4] J. O. Hirschfelder, C. F. Curtiss, and R. B. Bird, Molecular Theory of Gases and Liquids (John Wiley & Sons,

Inc., N.Y;, pp. 756ff, 1965).

[5] R. M. Fristrom, Simplified Chemical Models for Flames

and Fires, in preparation for submission to Comb.and

Flame.

[6] J. G. Kirkwood, and I. Oppenheim, Chemical Thermodynamics (McGraw Hill Book Co., N.Y., 1961).

[7] S. Gordon, Complex Chemical Equilibrium Calculations,

(see bibliography). Sec. 1 of Kinetics and Thermodynamics in High Temperature Gases, NASA SP-239, for

sale by the National Technical Information Service,

Springfield, Va., 22151 ($3.00).

[8] H. Browne and M. Williams, Theoretical Computation of

Equilibrium Compositions, Thermodynamic Properties

and Performance of Propellant Systems, Nav. Weps

Report 7034, U. S. Naval Test Station, China Lake,

Cal.

[9] D. Stull, et al., JANAF Thermochemical Tables, Dow

Chemical Co., Midland, Mich., Available through National Technical Information Service, Springfield, Va.,

22151, in loose leaf form and in tape form for machine

use.

[10] D. Stull, The Use of Thermodynamics for the Prediction

of Fire Hazard Class, Presented at the American Chemical Soc. Meeting, 1970, Reviewed in Fire Research Abstracts and Reviews, 12, No. 2, p. 150, (1970).

[11] R. M. Fristrom and A. A. Westenberg, Molecular Transport Properties for Flame Studies, Fire Research Abstracts and Reviews, 8, 155-83, (1966).

[12] Y. S. Touloukian, C. Y. Ho, P. E. Liley, R. W. Powell

and S. C. Saxena, Thermophysical Properties of Materials, Vols. 1-3 (Thermal Conductivity), Plenum

Data Corp., 227 W. 17th St., New York, N.Y., 10011,

(1970).

[13] R. M. Fristrom, Chemical Kinetic Information Programs:

Status in 1968, CPIA Publication No. 173, Aug. 1968, pp. 1-16, Chemical Propellant Information Agency,

Applied Physics Laboratory, The Johns Hopkins Univ.,

8621 Georgia Ave., Silver Spring, Md., 20910.

[14] J. Airey, P. Pacey and J. Polanyi, Infrared Chemiluminescence in the Systems H + Br2 and H + HBr, 12th

Symposium (Int.) on Combustion, p. 85, Combustion

Institute, Pittsburgh, Pa., 15219, (1967).

[15] R. Strehlow, Fundamentals of Combustion, p. 168, Inter- national Textbook Co., Scranton, Pa., (1968).

[16] G. J. Minkoff and C. F. H. Tipper, The Chemistry of

Combustion Reactions (Butterworths, London and

Washington, D. C, 1962).

[17] W. E. Wilson, J. T. O'Donovan, and R. M. Fristrom,

Flame Inhibition by Halogen Compounds, 12th Symposium (Int.) on Combustion. The Combustion Insti- ture, Pittsburgh, Pa., 929, (1969).

[18] G. Dixon-Lewis and A. Williams, The Role of HO2 in Ha-Oj Flames, Nature 196, 1309, (1962).

[19] A. Gaydon and H. Wolfhard, Flames (Chapman & Hall

Ltd., London, 1953).

[20] E. T. McHale, Survey of Vapor Phase Chemical Agents

for Combustion Suppression, FRAR, 11, 90, (1969).

[21] A. Modica and J. LaGraff, Decomposition and Oxidation

of C2F4 in Shock Waves, J. Chem. Phys. 43, 3383,

(1965).

I wonder whether Dr. Fristrom could give us any

idea about the magnitude of the rate constant of the

reaction

H+CH3+M-^CH4+M

where M may be N2 or any other inert gas.

140
R. M. Fristrom:

The analysis of formaldehyde offers difficulties

because of its tendency to polymerize to paraformaldehyde, a six-membered ring structure. This

makes it difficult to prepare standard samples for

instrument calibration. Quantitative sampling re- quires conditions such that condensation is unimportant. Where these conditions are met, mass

spectrometry and presumably other standard analytical techniques can be used.

In the case of our own flame studies, we found that

formaldehyde could be reproduceably sampled

using microprobe techniques with a continuous flow

inlet into the mass spectrometer. Condensation,

adsorption and reaction were negligible. The transfer

lines were teflon coated, held at room temperature

and at a pressure of the order of 10~^ torr. With

batch sampling in glass bottles at about 1 torr

condensation occurred, but sufficient formaldehyde

remained after eight hours to allow identification by

microwave spectroscopy. To calibrate the mass

spectrometer we thermally depolymerized paraformaldehyde in a stream of argon, measured the

total pressure and deduced the partial pressure of

formaldehyde by difference, using the known

sensitivity of the instrument to argon.

I am not aware of any direct studies of the three

body recombination of methyl radical with hydrogen

atoms. The reaction is probably hard to isolate from

competing reactions such as H +CH3-^CH2+H2,

followed by CH2 +CH3-^C2H6, etc. In a fuel rich

(reducing) system stable products such as ethylene

and ethane appear.

W. Tsang (National Bureau of Standards)

Compilations of high temperature rate data show

in general large scatter and uncertainties. The fact

that flame properties may be predicted on the basis

of such information suggests that they are in many

cases insensitive to the rate processes. Conversely,

this will raise questions on the accuracy of kinetic

information as extracted from flame properties.

R. M. Fristrom:

This is a question of which reactions are under

discussion and which reactions are relevant. Recent

critical compilations such as those of the High

Temperature Reaction Kinetic program at the

University of Leeds, England, and others mentioned

in Table VII list a number of reactions where agreement among recent reliable experimental studies lies within a factor of two. On the other hand many

or most reactions are still in doubt by several orders

of magnitude. The situation is improving. If one is interested in an overall flame property such as burning velocity or a single parameter such as engine

thrust, studies have shown that these processes are controlled by a few (<10) of the hundreds of

reactions which can be written. If the critical re- actions are known with precision (and these are often

well studied for just this reason), then it makes little difference whether the parallel reactions are in

doubt. If one is interested in trace species or detailed

processes, however, the situation is very clouded.

The reverse side of this coin is the use of flame

studies to deduce kinetic parameters.

Here, if one can isolate a flame process which

depends principally on a single reaction, then flames

offer a good media for studying high temperature

reactions (ref. 2). Otherwise, the situation is hopeless. It should be pointed out that this is generally

true of kinetic systems.

141

NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pykolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Carbon Formation in Premixed Flames

K. H. Homann

Institut fiir Physikalische Chemie der Universitat Gottingen, Gottingen, West Germany

In this paper a survey is given on the processes leading to the formation of solid soot particles

in premixed flames. Like most combustion products carbon formation is not exclusively determined

by flame chemistry but is also dependent on transport processes. A local increase in the overall

C/0 ratio in the burning zone which is connected to the back diffusion of hydrogen and hydrogen

atoms can reinforce carbon formation. This gives rise to the different types of carbon appearance in burner flames. It is shown that the C/0 ratio at the site of carbon formation can be very different

from that of the initial mixture.

For the formation of soot particles the presence of both highly unsaturated hydrocarbons

and hydrocarbon radicals is necessary. The process of particle growth which cannot be described

by equilibrium condensation theory shows two distinct phases: (1) A first phase of rapid growth

which is due to the addition of higher unsaturated hydrocarbons to a particle surface with radical

character. During this phase the number density of particles decreases while the individual particles retain their almost spherical shape. This is followed by (2) a phase of slow growth governed

by the heterogeneous decomposition of a variety of unsaturated hydrocarbons on the particle

surface which has lost most of its radical-like properties.

Key words: Acetylene-oxygen flame; benzene flames; bunsen flames; carbon limit; concentration

profiles; cyanogen-oxygen flame; flame front curvature; hydrocarbon radicals; hydrogen diffusion; carbon formation; oxidation reactions; particle size; transport processes.

1. Introduction

Particles of solid carbon, or rather soot particles, are among those products of combustion processes

which have been studied most intensely and under

many different aspects. A large number of facts

about carbon formation and soot properties has been

obtained since M. Faraday's famous lecture, "The

Chemical History of a Candle" [1].' In this lecture

he pointed out that carbon particles are set free by

the decomposition of the vapor of the wax, that they

rise up heated and glowing and are totally burned

when they come into contact with the surrounding

air. These three phases: formation, radiation, and

combustion of carbon particles are still subjects of

intense studies. Interest in them has risen to the

same extent as the industrial and ecological importance of carbon formation has grown in the last few

decades. The greatest efforts have focussed on the

control of carbon formation and combustion, ranging

from zero yield to the largest possible production of

soot imder given conditions. These studies are complicated by the fact that carbon formation is very

sensitive to many kinds of inhomogeneities that can

be present in a combustion system, such as those of

composition, temperature, flow conditions, walls etc. In the following we wish to restrict ourselves to the

formation of carbon in premixed laminar flames of

various fuels, mostly hydrocarbons.

^ Figures in brackets indicate the literature references at the end of this paper.

2. Appearance of Soot-Forming Burner

Flames

The sensitivity of carbon formation to inhomogeneities in the system gives rise to different types of

sooting flames that can be observed on Bunsen or nozzle burners: From the classification of these

flames as given by Behrens [2J (1951) and by Street

and Thomas [3] (1955) two groups shall be mentioned :

(1) The acetylene type flame where carbon appears

equally distributed above the whole flame cone. (See fig. 1)

(2) Other types of flames, where carbon formation is seen only at the tip of the flame cone, either

close to the flame front (e.g. benzene) or a short

distance above (e.g. cyclohexane) . (See figs.

2a b.)

In most hydrocarbon-air or hydrocarbon-oxygen

flames, except for very light fuels such as acetylene

and ethylene, carbon appears under normal conditions in the way mentioned under (2)

The limits of carbon formation, i.e., the ratios of

carbon to oxygen in the unburnt gas at which soot

formation is just observed, are also based to a large

extent on flames of type (2). These limits were measured by Street and Thomas [3].

However, the limits and the amount of carbon

formation for a certain fuel-oxidizer pair are determined not only by the mixture ratio. Also the flow

conditions at a given temperature and pressure can

143
Figure 1. Acetylene type.

markedly influence the carbon formation. Using a nozzle burner it can be demonstrated that a small

variation in flow which results in a change of the

shape of the flame can reinforce or weaken carbon

formation. A normally cone shaped benzene-air

flame forming carbon at its tip can be made to switch

to a nonluminous flame forming a button-shaped

flame front without alteration of the mixture ratio or the through-put of gas [4^. Compared to the cone shaped flame this flame lacks a strong curvature of

its combustion zone since it does not have a sharp

tip.

Other examples of reinforced carbon formation

above strongly convex portions of the flame front

(concave relative to direction of gas flow) are the

appearance of carbon streaks above the ridges of

disturbed flames of the polyhedral type and above

the many small tips of cellular flames [^5].

The influence of flame front curvature on carbon

formation is not restricted to stationary flames.

Gaydon describes a Bunsen flame burning without

carbon formation but near the limit thereof. If one now varies the flow velocity periodically then the

carbon streaks are emitted from the tip in phase with

the variation of the flow which also results in a variation of the flame curvature at the tip []6]. In

turbulent flames there are turbulence balls with

rapidly varying curvatures of the local flame front

which also will influence carbon formation in the

same sense.

3. Influence of Transport Processes

What is the reason for such localized appearance of

carbon formation in many flames? The examples

0,350

30 1,0 [mm]

Height above the burner

Figure 3. C/0 ratio for an n-heptane-airflame burning at the limit of carbon formation as function of

the height above the burner, measured along the axis

of the flame cone.

given above seem to indicate that the same reasons

that give rise to polyhedral flames are also related to

the locahzation of carbon formation. This would

mean that the mixtures entering the various parts of

tlie flame front are not the same and are different

from the gas mixture fed to the burner. Accordingly

the burning conditions are not the same.

It has been shown by Jost and co-workers that the

C/0 ratios at the tip of a Bunsen cone and above the

ridges of polyhedral flames are larger than those

above other parts of the flame [7]. More recently,

measurements along these lines have been carried

out on an n -heptane-air-flame burning just slightly

below the limit of carbon formation [8]. It was a polyhedral flame of 27 mm height having 9 ridges.

Figure 3 shows the variation of the C/0 and C/H

ratios along the axis of the flame. There is a steady

increase of the C/0 ratio towards the tip of the

flarne, becoming rather pronounced near the re- action zone. This is accompanied by a decrease in

the C/H ratio. Behind the flame front these ratios

tend towards their original value. In still richer

mixtures this increase in C/0 ratio develops so far

1 1 :

Flame '

1 1

H^O

j

CO,

P^-*-=5^o. o, 0

-

f

1 1

Figure 2. a. Benzene type; b. Cyclohexane type.

10 20 30 1,0 cmm]

Hohe uber Brenner

Figure 4. Concentration profiles along the axis of the flame cone of the flame from figure 3.

144
that the limit of flamability is surpassed and the

flame front breaks up at the tip.

A more detailed picture is given by the concentration profdes in figure 4. It is seen that the decrease

in the C/H ratio is mainly due to the comparatively

large amount of hydrogen diffusing back into the

rmburnt gas. A similar situation is found when

measuring concentration profiles through ridges of

the flame front. From concentration measurements

on flat hydrocarbon-oxygen flames it was known that

hydrogen diffuses farthest back into the unburnt gas

[9, 10]. However, this concentration of hydrogen

cannot be explained simply by diffusion back from

the tip of the flame.

Because of the complexity of the interdependence

of transport processes and chemical reaction in a Bunsen flame only a qualitative and necessarily

incomplete picture can be given here of the phenomena that cause these variations in mixture ratios, see also [11]. Let us first consider the situation in an

ideally cone shaped flame front, a cross section of a part of which is shown in figure 5. The straight lines

burning zone stream tubes

Figure 5. Diffusion of hydrogen in the burning zone of an

ideally cone-shaped hydrocarbon -air flame.

indicating an extended flame front shall represent

the lines of constant hydrogen concentration, a

profile of which is represented schematically at the

upper ends of the lines. The stream tubes of the gas

stream go through the burning zone as indicated.

Hydrogen diffuses back against the gas stream along

its concentration gradient. Thus, hydrogen diffuses

not only out of one volume element—arbitrarily

defined by two lines of constant H2 concentration and

two stream tube boundaries—into the element

directly upstream from it, but also into its neighboring volume element at the side towards the tip of

the flame. But through this there cannot be an

appreciable enrichment of H2 towards the tip, since

the same amount of hydrogen is diffusing into the

volume element from its other neighbor at the side

towards the base of flame.

However, this situation cannot be stationary if a

slight increase in hydrogen concentration in one volume element decreases the local burning velocity.

This is just what happens in mixtures of hydrocarbons and air near the limit of carbon formation.

As a consequence the flame front curves. If the lines

of constant hydrogen concentration would curve in

the same way there would be a net flow of hydrogen

towards the tip which again would cause a further

curvature of the flame front leading to an unstable

situation. Since there is a stable flame, the lines of

constant H2 concentration will assume the form

indicated by the dotted lines of figure 6. This again

T- profile ^Hj- profile

burning zone stream tubes

Figure 6. Diffusion of hydrogen in a curved flame front of a

hydrocarbon-air flame.

balances the transport of hydrogen in and out of the

imaginary volume elements, but causes an increasing

H2 concentration towards the tip when measured

along the axis of the flame.

Since the pressure in the flame does not change,

this diffusion of hydrogen towards the tip and the

axis must be balanced by a diffusion of unburnt gas

in the opposite direction. If the diffusion coefficients

of fuel and oxygen in the mLxture are sufficiently

different as in the case of n -heptane-oxygen, oxygen

diffuses faster towards the flame zone and causes an

increase in the C/O-ratio measured along the flame

axis. Since this increase is strongest towards the tip

of the flame it is here where the carbon formation

sets in. That the back diffusion of hydrogen is related to the appearance of carbon can be shown

in a simple experiment.

It is possible to eliminate that part of the hydrogen

diffusion which is caused by molecular hydrogen by

adding to the unburnt mixture a certain amount of

hydrogen so that the mole fraction of it in the

unburnt gas becomes nearly equal to that in the

burnt gas. The appearance of such a flame when

sooting is very similar to that of an acetylene type

flame: carbon luminosity appears above the whole

flame cone; furthermore, the slightly bent flame front

straightens out and resembles more that of an ideal cone. It is to be noted, however, that the carbon

luminosity is still strongest at the tip [4]. These few

examples show that transport processes, especially

the back diffusion of hydrogen from parts of the

burning zone, have a strong influence on the carbon

formation. The appearance of carbon at regions of

high curvature of the flame front, however, cannot

be attributed to a single effect. For a more complete

145
picture the diffusion of hydrogen atoms, and other

light radicals, must be taken into account. Furthermore, the steady rise of temperature may cause

pyrolysis ahead of the visible burning zone. The

laeating of the unburnt gas would be particularly

effective within the cavity of a convex portion of

the flame front.

Thus, in most cases carbon formation is not only

controlled by the flame chemistry but also by the

burning conditions. At the local site of carbon

formation the mixture in terms of the C/0 and C/H

ratios may be very much different from that in the

fresh gas. It shows that one should be careful in using

limiting C/0 ratios measured in the unburnt gas for a discussion of the chemistry or even mechanistic

aspects of carbon formation. This does, however, not

diminish the great practical value of these measured

figures.

4. Theoretical and Practical Carbon

Limits

The critical C/0 ratio in the unburnt gas at which

carbon appears in practical flames ranges from 0.4

for heavy fuels to about 0.9 for light fuels such as acetylene, as measured by Street and Thomas [3]

and others. For solid carbon in thermodynamic

equilibrium with the flame gases, this ratio should be

fairly independent of the fuel and should have a value

very close to unity, changing somewhat with the

temperature, e.g. [12]. The measurements of the

variation of the C/0 ratio of the unburnt gas along

the axis of a sooting Bunsen flame shows that this

ratio could increase quite steeply. This raises the

question whether it might reach the theoretical value

of about unity in the actual combustion zone at the

tip where the carbon is formed. Measurements on a flame burning at 1 atm could not be carried out

directly in the flame front because of interference of

the probe with the narrow combustion zone. However, concentration measurements in the combustion

zone of flat flames at lower pressure showed that the

ratio does not reach its theoretical value at the point

where sooting begins [10, 13]. The practical carbon

limits in flat flames are on the average only 10-13

percent higher than those in Bunsen flames [14]. If

the fuel-oxidizer mixture is burned in a well-stirred

reactor, the carbon limits are 10-20 percent higher

than in conical flames [15].

These limits of carbon formation and more detailed

measurements of the composition of sooting hydrocarbon flames show that the process of carbon

formation takes place under non-equilibrium conditions. It is, therefore, intrinsically different from the

condensation of a normal vapor where the liquid

droplets represent an equilibrium state [16]. There

is no critical size of carbon particles above which the

rate of growth (as in condensation) is greater than

the rate of vaporization and which could be corre- lated with a certain degree of supersaturation of

carbon vapor. The only flame known in which carbon

formation seems to take place almost under equilibrium conditions is the very hot cyanogen-oxygen

flame [17].

5. Chemistry of Carbon Formation

Since carbon formation in hydrocarbon flames is recognized to be governed by chemical kinetics

rather than through thermodynamic equilibrium one

FIG 7q

FLOW OF BURNT GAS

20 30 iO 50 [mm] HEIGHT ABOVE THE BURNER

CONCENTRATION

/ \

FIG 7c

/ \

/ \^ solid Carbon

H2

FIG 7d

50 [mm] HEIGHT ABOVE THE BURNER

d [A]

100 150 200 250

Figure 7. Particle history in aflat CiHi-Oiflame burning at 20

mm Hg^ C2H2/O2 approx, 1 ,4.

(a,b) Scheme of the flame and schematic representation of the concentration profiles of main components. The mole fractions at a distance of 40 mm above the burner are approximately: CO: 0.56, H2O: 0.05, C2H2:

0.12. Curves are not to scale among each other.

(c) Schematic representation of various products. The mole fractions at 40 mm above the burner are approximately: C4H2: 0.016, C14H8: 6* 10"®;

solid carbon: 0.25 percent by weight. Precursors of carbon particles p = 10~^ at 14 mm; ••••• represents the change of number density of carbon particles with diameters around 100 A; represents the change in number density of carbon particles with 200 A diameter.

(d) Distribution functions: The number of particles/cm^ in a certain diameter interval plotted against the diameter for various heights above

the burner. The height is given as parameter. The function at 14 mm

represents the precursors p. The number of acetylene molecules/cm^

which is transferred into carbon is ca, 4' 10^** at burned gas density.

146
must ask: What happens to hydrocarbon molecules

in flames until their carbon atoms appear in solid

particles. Some insight into the various steps of that

process has been obtained from investigations of flat hydrocarbon flames burning at reduced pressure.

The use of reduced pressure was necessary to obtain

sufficiently extended reaction zones [13, 18j.

In order to follow the formation of solid particles

the following questions should be considered:

(1) What are the main products formed in the

oxidation zone of the flame?

(2) Which intermediates are formed during the

combustion of the fuel?

(3) Which higher carbon containing species are formed in the reaction zone and what are to

be considered the precursors of carbon

particles?

(4) How do the carbon particles grow?

(5) How does the total amount of carbon increase

in the particles and what are their properties

in the various stages of their growth?

As an example, some experimental answers to these

questions obtained from acetylene -oxygen flames

burning at pressures around 20 mm Hg will be

discussed (see fig. 7).

The main reaction products are CO, H2O, CO2, H2

and C2H2 (fig. 7b). Another group of products, the

concentrations of which show maxima at the end of

the oxidation zone, are polyacetylenes (C2nH2). In

the burnt gas their concentrations remain practically

constant at this pressure. Together with C2H2 and H2

they form a kind of equilibrium. A very similar

situation arises in flames of other hydrocarbons. In

figure 7c concentration profiles of polyacetylenes are represented by that of C4H2, their maximum concentration decrease rapidly with increasing numbers of

carbon atoms in the chain. In the oxidation zone which ends with the consumption of oxygen a variety

of other hydrocarbons are found, such as propylene,

methylacetylene, dimethylacetylene, vinylacetylene,

benzene, methane and others. They all reach a concentration maximum inside the oxidation zone and

have practically disappeared when oxygen is con- sumed.

In addition to the stable components mentioned

several radicals such as C2H3, C2H, C2, CH, H, O,

OH etc. have been analyzed in the reaction zone of an

acetylene-oxygen flame [10, 13]. The large number

of various positive and negative hydrocarbon ions

that can appear in hydrocarbon flames may be seen from measurements of Sugden and co-workers [19].

Another characteristic group of hydrocarbons

formed in that flame is the polycyclic aromatic

hydrocarbons such as naphtalene, phenanthrene,

pyrene, 3,4 benz-pyrene, coronene and related

species. The concentration profile of one of them

(ChHs) is shown in figure 7c. These species start to

appear at the end of the oxidation zone and keep

steadily increasing in concentration in the burnt gas, even in that region where carbon formation has

become immeasurably slow under low pressure

conditions.

Still another group of hydrocarbons was detected

in the gas phase (P in figure 7c) . They have molecular

weights between about 150 and 550 (550 was the

limit of the mass range of the analyzing mass

spectrometer). In contrast to the aromatics mentioned above their concentration shows a maximum

between 12 and 14 mm above the burner and they

disappear again within the zone where the carbon

particles grow [20]. The number of carbon atoms in

them ranges from about 15 to above 40. Their

hydrogen content is larger than that for condensed

aromatic species of the same number of carbon

atoms. Although they have not yet been identified

individually, many of them seem to contain side

chains. With increasing height in the flame the

relative concentration maximum within this group is shifted to higher molecular masses, while their

absolute concentration decreases steadily. They are considered to be important intermediates for the

formation of solid particles.

The individual concentration of the P type species

while further growing become too low for a direct

analysis until the first small carbon particles, 20 to

30 A in diameter, can be perceived by aid of electron

microscopy a few mm further in the burnt gas.

Sampling of carbon particles from various heights in

the flame gave the following results:

(1) The size of the particles increases continuously up to about 50 mm above the

burner.

(2) The number of particles decreases in that

range.

(3) The total amount of solid carbon increases

during this process.

(4) The agglomeration of individual particles into

the well-known grape or bead-like clusters occurs in a later stage.

At greater heights in the flame no marked change

in particle size and total amount of carbon was

observed. This means that a continued formation of

carbon in the burnt gas (which would have to be

measured in flames burning at higher pressure) must

be much slower than the process described. The distribution functions of particle size at different

heights in the flame show that the number of

particles with diameter less than 100 A has become

very small even at 20 mra above the burner. This

means that the rate at which new small particles are formed at this height of the flame is negligibly small.

From figure 7d one can see that the intermediate

hydrocarbons (P) which are considered here as "carbon particles" of 8 to 10 A diameter fit quite

well into the scheme.

It is not only the size of the particles that changes

during this phase of rapid growth but also their

chemical and physical properties. Their H/C ratio,

for example, steadily decreases towards greater

heights above the burner. E.S.R. signals obtained

from carbon particles sampled at various heights

show a continuous and rather strong decrease from

"young" to "old" carbon particles. When rapidly

cooling a flame gas sample containing carbon

147
particles a large numbers of higher hydrocarbons

will condense on the particles. They can be evaporated again when the carbon is heated under vacuum.

In agreement with the direct gas phase analysis there

are mainly aromatics with the carbon particles from

greater heights in the flame while hydrocarbons of

type P are collected with the "young" particles

[21].

6. Discussion of the Process of Carbon

Formation from Aliphatic Fuels

What can be said about the essential steps of the

process of carbon formation on the basis of these

experimental results? The initiation of carbon

formation in premixed flames under the conditions

discussed is definitely related to the chemical re- actions in the burning zone where the oxidation

process gives rise to various types of radicals in

rather high concentrations. This is in contrast to the

mainly pyrolytic decomposition of the fuel and

subsequent particle formation in diffusion flames.

Both the production of hydrocarbon radicals and

of highly unsaturated hydrocarbon molecules is essential for the initiating steps of very complicated

sequences of reactions which finally lead to the

formation of solid particles.

The unsaturated hydrocarbons which are formed

as intermediates or final products in the combustion

of different lower aliphatic fuels are essentially the same. They differ only in concentration depending

mainly on the nature of the fuel molecules. Very

often their relative concentrations are almost

identical if the flame temperatures do not differ

too much.

The main rmsaturated species formed in the

oxidation zone are acetylene, polyacetylenes, methylacetylene, ethylene, propene, vinylacetylene, and

butadiene. Of these, however, mainly acetylene and

polyacetylenes are still present at the end of the

oxidation zone where carbon formation begins. The

other hydrocarbons disappear again at an earlier

stage in the oxidation zone. In view of this situation

we shall limit our discussion of the first steps leading

to carbon formation to those in which acetylene and

polyacetylenes are involved.

That the presence of radicals which are generated

by the oxidation reaction is necessary for a rapid

formation of carbon can be clearly seen from the

situation in the burnt gas. Here, acetylene, polyacetylenes and polycyclic aromatics are still present

in appreciable concentrations but the formation of

new particles or the increase of soot has apparently

become very slow. Thus, carbon formation is only

rapid in the vicinity of the radical-rich oxidation zone but it slows down with decreasing concentration

of small hydrocarbon radicals.

In premixed flames hydrocarbon radicals are mainly formed by reactions of hydrocarbons with

active species such as O, OH, and hydrogen atoms.

In connection with carbon formation we will only

consider those with acetylene and polyacetylenes for

the same reasons as above.

The study of the reaction of 0-atoms with acetylene over a temperature range of a few hvmdred degrees

shows that the main reaction proceeds via [22]

0+C2H2^CH2+CO

{k ^ 10i3e-28oo/flr ci„3 mole-i sec-i)

H atoms and OH radicals react with acetylene to

give

H +CsHs^CaHa [23]

and

OH+CaHa-^CO +CHs, respectively [24].

These reactions give rise to the radicals CH2, C2H3

and CH3. Furthermore, C2H radicals, the source of

which is not yet certain, have been observed in

acetylene flames.

Primary products of the corresponding reactions

of diacetylene are not known. It is not unreasonable,

however, to assume that CO is also one of the primary products which makes probable the formation

of C3H2 and C3H3 for the reaction with O and OH,

respectively. Reaction with H-atoms will certainly

involve the formation of C4H3. The fate of these

radicals will be oxidation in flames rich enough in

oxygen to prevent carbon formation. In sufficiently

fuel rich flames the concentration of acetylene and

polyacetylenes are high enough so that addition

reactions of these radicals to the acetylenes are

faster than the oxidation reactions.

In an excess of acetylene a consecutive reaction

of the CH2 radical gives

CH2" +C2H2—^CsH^'

which at lower temperatures will stabilize to

methylacetylene. In analogy to this reaction,

CH2 -}-C4H2^C5H4' is probable. In a similar way

radicals such as CH3, C2H3, C2H can add to acetylenic hydrocarbons and produce higher hydrocarbon

radicals. As pointed out by Porter [25] the presence

of double and triple bonds is very important because

it enables these particles to maintain their radical

character during further addition of small highly

unsaturated hydrocarbons. The radical character is to a large extent still present in "young" carbon

particles.

The addition of small radicals to acetylene and to

terminal carbon atoms of polyacetylenes causes the

formation of diacetylene and higher polyacetylenes,

respectively. Measurements of polyacetylene concentrations in flames under different conditions suggest

that the single addition reactions are fairly well

equilibrated. This leads to a quasi -equilibrium

between acetylene, the polyacetylenes, and hydrogen

as long as radicals are present.

Addition of a radical to a nonterminal carbon atom

of a polyacetylene, however, does not easily lead to

a stable branched molecule so that by further

addition of acetylene and polyacetylenes a highly

branched radical can be formed. For this mechanism,

growing biradicals such as CH2 or C3H2 are much

more effective than normal radicals. The larger those

148
H/C - atomic ratio of carbon

C2H2/O2 = 1.7;

20 Torn

10 20 30 AO mm

Height above the burner

Figure 8. H/C-ratio in the carbon particles of a C2//2-O2

flame with CiHi/Oi = 1.7 as a function of the distance

above the burner.

(p = 20 mm Hg, vo = 58.5 cm/e)

radicals grow the more easily can they make internal

additions i.e. ring closures while the whole particle

retains its radical character. Through recombination

of two such radicals or through "internal recombination" (ring closure of a branched biradical with loss

of the radical character) hydrocarbons of the type

P and also the much more stable polynuclear aro- matic species can be formed. From the concentration

behavior of the aromatics in the flame and also from

their known comparatively high thermal stability it must be assumed that they are byproducts in the

process of carbon formation rather than important

intermediates [26, 21~\.

These large branched and ring-containing hydrocarbon radicals add predominantly acetylene and

diacetylene first without losing hydrogen. This is indicated by the atomic H/C-ratio in the first carbon

particles (fig. 8). During the formation of carbon

both the large hydrocarbon molecules and the small

soot particles between which there is a continuous

transition decrease in number concentration while

the total amount of soot increases. Apparently two

things happen: (1) small carbon particles (and also

large hydrocarbon radicals) which come together

will coalesce; (2) these particles continue to add

hydrocarbons from the gas phase. In a later phase of

growth the carbon particles—by a kind of internal

tempering process—give off hydrogen and loose

much of their radical character. This prevents a further rapid growth just by addition of hydrocarbon molecules. As a consequence the increase in

the amount of carbon slows down and continues

only by heterogeneous decomposition of hydrocarbons on the surface of the carbon particles. This,

however, is a slow process as compared to the one

described.

7. Carbon Formation from Aromatic

Fuels

In flames of benzene and related aromatic fuels the

situation is different in so far as even at the beginning

of the process rather large hydrocarbon molecules

are present as "building bricks" for carbon particles.

This increases the rate of forming higher hydrocarbons relative to the rate of oxidation. It causes the zone where the small soot particles begin to grow

to overlap more with the oxidation zone of the flame.

Thus, in addition to the presence of large hydrocarbon molecules from the beginning, carbon

formation is reinforced here because it starts even

in a flame zone where the concentration of small

radicals is relatively higher than in the zones where

particle formation starts in flames of lower aliphatic

fuels [26].

In zones of rapid particle growth in a benzene

flame similar polycyclic hydrocarbons of type P, as

in acetylene flames, are found. Again condensed

aromatics that are formed behave like comparatively inert byproducts in the process of carbon

formation.

Although the concentration of higher hydrocarbons and hydrocarbon radicals may be quantitatively different in the carbon forming zones of flames

of different fuels, qualitatively they are similar, even

for fuels such as alcohols.

The examples discussed in this paper naturally cover only a part of the problem of the chemistry of

carbon formation. Even for the phase of formation

of carbon there are many aspects that could not be

discussed here, such as carbon formation in diffusion

flames, the influence of additives and of electrical

fields on the carbon content of flames, to mention

only a few. Many details of carbon formation not

mentioned here are contained in recent review

articles by Gaydon and Wolfhard [6], Palmer and

Cullis [28], Thomas [29], and Homann [4, 21].

I wish to express my sincere gratitude to my

academic teachers Professors W. Jost and H. Gg.

Wagner.

8. References

[1] Faraday, M.: The Chemical History of a Candle, p. 36,

(Hutchinson & Co., London:, 1907).

[2] Behrens, H.: Z. physik. Chem. 196, 78 (1950); ibid. 199,

1 (1952).

[3] Street, J. C. & Thomas, A.: Fuel, Lond. 34, 4 (1955).

[4] Homann, K. H.: Angewandte Chemie Int. Ed., 7, 414

(1968).

[5J Flossdorf, J. & Wagner H. Gg.: Z. physik. Chem. N.F.

54, 113 (1967).

[6] Gaydon, A. G. & Wolfhard, H. G.: Flames, their structure, radiation and temperature, (2nd ed.) p. 171,

(London: Chapman & Hall 1960).

[7] Jost, W., Krug, J. & Sieg, L.: Fourth Symposium (Int.) on Combustion, p. 535, Baltimore: Williams & Wilkins

1953.

Jost, W. & Krischer, B.: Z. physik. Chem. N.F. 5, 398

(1955).

[8] Flossdorf, J.: Dissertation, University of Gottingen, 1968.

[9] Fristrom, R. M. & Westenberg, A. A.: Flame structure,

(New York: McGraw-Hill 1965).

[10] Homann, K. H., Mochizuki, M. & Wagner H.Gg.: Z. physik. Chem. N.F. 37, 299 (1963).

[11] Markstein, G. H.: Nonsteady flame propagation, Oxford

Pergamon 1964, Section D.

[12] Gay, N. R., Agnew, J. T., Witzell, O. W. & KarabeU,

C. E. Combust. Flame 5, 257 (1961).

149
[13] Homann, K. H. & Wagner, H.Gg.: Ber. Bunsengesellschaft Phys. Chem. 69, 20 (1965).

[14] Ryan, P.: Diplomarbeit, University of Gottingen 1968.

[15] Wright, F. J.: Twelfth Symposium (Int.) on Combustion,

p. 867, Pittsburgh: The Combustion Institute, 1969.

[16] Becker, R. & D6ring, W.: Annln. Phys. 24, 719 (1935).

[17] Tiscber, R. L. & Scheller, K.: Tech. Rep. WSCI-68-4,

Aerospace Res. Lab. Wright-Patterson AFB, Ohio 1968.

[18] Bonne, U., Homann, K. H. & Wagner, H.Gg.: Tenth

Symposium (Int.) on Combustion, p. 503, Pittsburgh:

The Combustion Institute 1965.

[19] Sugden, T. M.: Tenth Symposium (Int.) of Combustion,

p. 539, Pittsburgh: The Combustion Institute 1965.

[20] Homann, K. H. & Wagner, H.Gg.: Eleventh Symposium

(Int.) on Combustion, p. 371, Pittsburgh: The Combustion Institute 1967.

[21] Homann, K. H.: Combust. Flame 11, 265 (1967).

[22] Hoyermann, K. H., Wagner, H.Gg. & Wolfrum, J.: Z. physik. Chem. N.F. 55, 72 (1967).

[23] Hoyermann, K., Wagner, H.Gg. & Wolfrum, J.: Ber. Bunsengesellschaft Phys. Chem. 72, 1004 (1968).'

[24] Gehring, M., Hoyermann, K., Wagner, H.Gg. & Wolfrum,

J. Z. Naturforsch. 25a, 675 (1970).

[25] Porter, G.: Fourth Symposium (Int.) on Combustion,

p. 248, Bahimore: Williams & Wilkins 1953.

[26] Homann, K. H. & Wagner, H.Gg: Proc. Roy. Soc. A307,

141 (1968).

[27] Long, R. & Tompkins, E. E.: Nature, London 213, 1011

(1967).

[28] Palmer, H. B. & Cullis, C. F.: in "Chemistry and Physics

of Carbon" (ed. Walker P.L.) 1, 265 New York: Marcel

Dekker 1965.

[29] Thomas, A.: Combust. Flame 6, 46 (1962).

Discussion

Henry A. Hill (Riverside Research Laboratory,

Inc.):

An initial comment related the finding of the cyclic

carcinogens (benzpyrenes) in vapors formed by fats

dropping onto hot wood or coal carbons in the

charcoal broiling of foods.

The second comment asked the effect of a solid

phase of red hot carbon in the burning zone on

carbon formation and the backward diffusion of

hydrogen. From this the next question tied together

Dr. Wall's question about the presence of a 1600

molecular weight fragment of polymer in vapor

phase in flame to the presence of added mechanically

propelled (not vaporized) carbon particles into the

flame as nuclei for growth or as fuel for combustion.

What happens to added carbon particles in the zone

of hydrogen diffusion?

K. H. Homann:

As far as I know there is no direct experimental

information on the influence of artificially added

carbon particles on the formation of new particles.

According to the experiments on burning carbon

particles one may guess that most of the carbon

particles will be oxidized in the burnt gas of the flame

if its mixture ratio is below the limit of carbon

formation. If the flame forms soot by itself injected

carbon particles will not influence the generation of

new small particles but will grow by heterogeneous

decomposition of hydrocarbons on their surfaces.

C. F. Cullis (The City University, London)

The kinetic parameters of the reaction:

C2H+C2H2-^C4H2 +H (1)

which has been shown to be fast [1] are at present

being measured at the City University, where, by

the use of acetylene-d2, it is possible to distinguish

reaction (1) from the hydrogen-transfer reaction:

C2H +C2H2^C2H2 +C2H. (2)

The widespread occurrence of reaction (1) during

the formation of carbon from acetylene also receives

indirect support, however, from evidence that

diacetylene (a product of reaction (1)) is an important intermediate in the carbon -forming process.

Thus, although vinylacetylene is under a wide range

of conditions the principal primary product of the

breakdown of acetylene [2-6j, kinetic and analytical

studies of the pyrolysis of the former compound [7],

together with isotopic tracer investigations [8],

suggest that the dimer is not an important intermediate in the formation of carbon from acetylene.

On the other hand, diacetylene, which produces

carbon much faster than does acetylene imder

comparable conditions [9-10] almost certainly

plays a much more prominent role.

References

[1] Lange, J., Diplomarbeit, Univ. of Gottingen, 1969.

[2] Skinner, G. B., and Sokoloski, E. M., J. Phys. Chem.

64, 1952 (1960).

[3] Bradley, J. N., and Kistiakowsky, G. B., J. Chem. Phys.

35, 264 (1961).

[4] Hou, K. C, and Anderson, R. C, J. Phys. Chem. 67, 1579

(1963).

[5] Palmer, H. B., and Dormish, F. L., J. Phys. Chem. 68,

1553 (1964).

[6] Cullis, C. F., and Franklin, N. H., Proc. Roy. Soc. (London A. 280, 139 (1964).

[7] Cullis, C. F., and Read, I. A., Trans. Faraday Soc. 66,

920 (1970).

[8] Cullis, C. F., Read, I. A., and Trimm, D. L., Eleventh

International Symposium on Combustion, p. 391 (The

Combustion Institute, 1967).

[9] Kinney, C. R., and Slysh, R. S., Proc. Fourth Conf.

Carbon, p. 301 (1960).

[10] Slysh, R. S., and Kinney, C. R., J. Phys. Chem. 65, 1044

(1961).

K. H. Homann:

Preliminary measurements on the reaction

C2H +C2H2-^C4H2 +H (1)

at a temperature of about 500 K have rendered a fci~2'10'^ cmVmol s [1]. This figure can be compared with a rough estimate from measurements of

the C4H2 profile in an C2H2-O2 flame and the C2H

concentration as estimated in such a flame [2], If

it is assumed that at the maximum of the C4H2

I
profile the rate of formation of C4H2 by reaction ( 1

)

equals the rate of consumption through O atoms and

OH radicals a ki of (3 to 4)-10'2 cmVmol s at a temperature of about 1900 K is obtained. This

estimate also implies the assumption that according

to measurements of Niki [3] C4H2 reacts 6 to 8 times

faster with O atoms (and also with OH?) than does

C2H2. Although this is a very rough estimate it shows at least agreement within much less than an

order of magnitude. There is evidence, however, that

also the reaction C2H3 +C2H2 can contribute to the

formation of C4H2 in flames.

A still unsolved problem is the source for C2H in

C2H2-O2 flames. All experiments on reactions with

acetylene and active particles seem to indicate that

the source must be a reaction with an appreciable

activation energy that can only be observed at

higher temperatures. At lower temperatures other

reactions not leading to C4H2 are much faster.

References

[1] K. H. Homann, W. Lange, and H.Gg. Wagner, Ber. Bunsengesellschaft Physik. Chem. 75, 121 (1971).

[2] K. H. Homann and H.Gg. Wagner, Ber. Bunsengesellschaft Physik. Chem. 69, 20 (1965).

[3] H. Niki, J. Chem. Physics 47, 3102 (1967).

H. B. Palmer (Pennsylvania State University)

The observation that flame inhibitors frequently

enhance carbon formation can be qualitatively

understood by noting that in any hydrocarbon flame

there will be competition between oxidation reactions

and "cracking"-type reactions. The latter are the

steps that may ultimately produce carbon. Anything

that slows the oxidation is apt to enhance the relative

significance of cracking. Work by Roger Millikan

several years ago showed that the critical question in

carbon formation probably is the competition

between oxidation of some intermediate by OH and

the reaction of that intermediate with other species

that can lead to carbon formation. But in flame

propagation, we know that OH also plays an important part in carrying chains, and inhibitors may

function by removing OH. Thus, in general, inhibition can be expected to increase the carbon-forming

tendency of a flame, because of the significance of

OH both in flame propagation and in the prevention

of sooting.

K. H. Homann:

The point of view that carbon formation or no

carbon formation in premixed flames is the result of a competition between oxidation reactions in which

OH and O atoms play a role and "cracking"-type

reactions is in principal correct—although these

"cracking"-type reactions can be quite different from

those known to occur in cracking or in pyrolysis.

Considering the role of inhibitors in influencing the

burning velocity by removing OH radicals or other

active species and increasing the carbon forming

tendency one has to differentiate. It depends very

much on the conditions of burning, on the nature of

the fuel, on the temperature where carbon formation

begins and on other parameters whether the inhibitory effect is paralleled by an increased carbon

formation. In oxyacetylene flames, for example, a

relatively large concentration of OH in the oxidation zone before the temperature reaches its maximum

leads to formation of CO2 and H2O in excess of those

concentrations which correspond to the equilibrium

at the final flame temperature. Thus, too much

oxygen is bound in products other than CO and is lost for the prevention of carbon formation. An

inhibitor which serves for suppressing the OH

concentration in lower temperature zones of the

flame and forces the combustion to take place at a higher temperature on the average will prevent this.

In this way carbon formation will be reduced and

not enhanced.

Donald Aldrich (Massachusetts Institute of Technology) : You implied that carbon formation in premixed

flames follows a different mechanism from that in

difi^usion flames. This, of course, is imderstandable

if a different fuel is used in the two types of flames,

but it is not clear if the same fuel is used in both

cases. There are obvious physical differences, but

what is the difference in the chemical mechanism of

carbon formation in premixed and turbulent flames

if one considers this formation to result from the

same chemical species in both cases?

K. H. Homann:

The main difference in the "mechanisms" of

carbon formation in diffusion flames and premixed

flames lies in the early stages of carbon formation.

In premixed flames two stages of carbon formation can be distinguished. The first and faster one is strongly coupled to the oxidation reactions and the

small hydrocarbon radicals formed herein. The

second consists mainly in heterogeneous decomposition of hydrocarbons on the surface of soot

particles. This second process can in principle also occur in diffusion flames. If carbon is formed in the

absence of oxygen which can be the case in diffusion

flames the precursors of carbon particles must be

products from pyrolytic reactions of the fuel.

Pyrolytic reactions, however, are quite different

from those reactions in the oxidation zone of premixed flames leading to higher unsaturated hydrocarbons. A phase of rapid particle formation comparable to that directly behind the oxidation zone of

premixed flames which is due to a large concentration

of highly unsaturated hydrocarbons and a relative

large concentration of hydrocarbon radicals has not

been found with carbon formation in the absence

of oxygen.

In turbulent flames, of course, both "mechanisms"

are operative depending on the degree of premixing

and on the condition of combustion.

151

NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pykolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Fluorocarbon Combustion, Fluorine Supported Combustion Kinetics?

Edward A. Fletcher

Department of Mechanical Engineering, University of Minnesota, Minneapolis, Minn. 55455

Some interesting characteristics of the combustion behavior of freons, and fluorocarbons are presented. The flame speeds and spontaneous ignition hmits of chlorine burning in fluorine are

easily described by simple models using well characterized kinetics. The flame speeds of CF3CI-F2

and CFsBr-Fj mixtures can be discussed in terms of simple thermal theories. The striking differ- ences in the combustion behavior of homologous fluorocarbons are shown to depend on big differ- ences in the rates at which the fuel molecules are initially attacked by the competitive combustion

reactions, but the kinetics are not yet understood.

Key words: Chlorine; combustion; explosion limits; flame speed; fluorine; fluorocarbons; freons; ignition limits; kinetics.

1. Introduction

^'^hen Dr. Wall invited me to talk about combustion kinetics at this meeting, I was honored and

delighted to accept. I had planned to talk about

combustion phenomena using the usual textbook

approaches to the kinetics, i.e., global kinetics, and

to illustrate some of the points from our own inprogress work with fluorine supported and fluorocarbon combustion studies. When I finally got the

printed program and book of abstracts I realized to

my dismay that Sidney Benson and Bob Fristrom

would probably scoop everything that I had planned

to say about kinetics in a general way. Besides, I felt a little like the victim of the Johnstown flood who

agreed to give, in heaven, a seminar on his experiences, only to discover that Noah was to be in the

audience. Nevertheless, despite the hazards, my

subject still will be as announced in the program.

Flames have been called the poorest of all possible

systems in which one can do kinetic studies^ [^Ij.

Whether that statement is true or not depends on what one means by kinetic studies. Although the

rates and mechanisms of chemical reactions are

i extremely important in determining how pyrolizing or burning systems behave, kineticists and combustion men are rarely talking about the same thing

when they use the same words. Sometimes even a given individual uses the same words to mean

different things, depending on whether he's talking

as a kineticist or a combustion man [2, 3]. On rare

^ Figures in brackets indicate the literature references at the end of this paper.

occasions, realistic kinetics has been applied to

compute flame structures and speeds fS, 6] and

ignition limits [7j, but the techniques one has to

resort to are not very useful, whereas the most

useful models use techniques that are repugnant to

the kineticist. This uncomfortable state of affairs

exists because many models that are, at best, heuristic work quite well, and in science and technology

we frequently pragmatise: If it works it must be

right.

Thus, any number of flame speed "theories" that,

in their fashion, work, have been proposed [8, 9, lOj.

Most of them include an energy equation which

contains a "global" reaction rate term. The rate is given by a rate constant which varies exponentially

with the temperature; the variation depends on an

"activation energy". Having written the equation,

the theoretician devotes his effort to solving it, along

with momentum and continuity equations, revealing

much ingenuity and sometimes brilliance, having

forgotten, in the meantime, that the exponential

term that is causing his difficulties is not real anyway.

To a kineticist, an activation energy has, he hopes, a

real, physical significance related to the mechanism

of a specific elementary reaction. In combustion

processes, it's just a device that gives you a term

that looks like an elementary reaction rate and can sometimes fool you into believing that you have

gotten insights when, in fact, you haven't. The

combustion man hopes it will work, i.e., have

applications; it frequently does, because the "rate

constants" of complicated reactions can often be

represented as products of the rate constants of

153
elementary reactions.^ The fact is interpreted to

mean that the kinetics and mechanism of the combustion process are understood, at least "in principle."

This kind of treatment of matters dear to the

hearts of kineticists has led Fristrom and Westenberg

to apply the adjective "execrable" to the "global"

kinetics so dear to the hearts of combustion and

pyrolysis men. But the latter aren't the only ones who are guilty of this chicanery. If this be sin, we

can derive consolation from the knowledge that we

sinners are probably in the majority.

On the other hand, one might, with some justification argue that the formalism of chemical kinetics is unrealistic and that even the simplest of the chemical

reactions in which there is intrinsic interest are too

complicated for the formalism of classical kinetics.

The kinetics related sciences make use of many

kinds of ideas applied to solve a vast conglomeration

of problems ranging from intra-atomic energy

transitions to reaction schemes of literally global

magnitude. What is basic depends on where one

chooses to enter the conglomeration. We may be

trying to push a camel through a needle's eye if we

insist that we describe the processes that take place

in flames and in pyrolizing solids using the same

concepts that can be applied to the homogeneous,

gradient free systems in which kinetic parameters are usually defined. Ultimately, there should be some

kind of convergence—a sort of correspondence

principle—but I think that combustion and pyrolysis

are often the losers if we insist on a blind allegiance

to the formalism of kinetics as a necessary prerequisite to a claim that we understand combustion.

If the kineticist gets too hard to get along with, we

can always point out that his assumptions about the

state of his molecules look worse and worse as we

take a closer and closer look at them. An ox has

fleas to plague him, and the flea has other fleas to

plague him ad infinitum.

^ For example: (B. E. Dahnecke, S. J. Wiersma, and E. A.

Fletcher, unpublished work). The rate of formation of GIF

from CI2 + F2 is well represented by the equation:

d[ClF] 2kA[F2]W

dt kelCh] kc[ClF]\

[Fd ^ [F2] /

where

kA = (.hkikz/kiY'^

kB = kike/kiJci

kc = h/kz

The numbered fc's are specific rate constants for the ele- mentary reactions that go to make up the mechanism of the

overall reaction. "Activation energies" associated with the

lettered fc's can be computed from the activation energies of

the numbered fc's, but the numbered fc's are the only ones whose activation energies have mechanistic significance. An

"activation energy" of the overall reaction can be formally

computed, but its magnitude depends on the initial mixture

composition and the extent of the reaction. Nevertheless, it may be a very useful thing to know, if one wants to know

something about the behavior of chlorine-fluorine mixtures

spontaneous ignition limits, for example.

We adopt, then the point of view that the study of

the kinetics of combustion processes is a useful,

worthy discipline in its own right. It is best studied

in combustion systems. The kinetics of other kinds

of complex processes are best studied in the systems

to which these processes apply. Whenever we can, we

try to relate the kinetics to the formalism of classical

chemical kinetics. In complicated systems, it is questionable whether information about the rates

and mechanisms of the so-called elementary reactions

is more important or useful than other kinds of

information in helping us understand many kinds of

phenomena such as ignition, flame propagation, or quenching, although as was pointed out before, such

information has been used in a few cases to rationalize

experimental observations.

The formal definitions of transport coefficients

and rate constants don't really apply to the systems

we deal with. How can you talk about the variation

of thermal conductivity or a rate constant—in a reacting system with concentration gradients in it,

when the concept of temperature, especially as it applies to conductivity or reaction kinetics is not

clearly defined for such a system—without compromising your principles? You can't. All sorts of

dodges, many of them useful, have been employed

to permit us to retain our formalisms, but they don't

change the fact that many of the formalisms we are used to using become intractible—indeed, are not

realistic—when we try to apply them to real systems.

Keeping that sort of thing in mind, my students

and I set out several years ago to study some combustion phenomena in what we hoped would be

simple systems. Fluorine (occasionally CIF3) was

our oxidant. The studies are still in progress and it's sometimes foolish to write about things while you're

still thinking about them. We know we've sinned.

Please forgive us.

2. Combustion Phenomena Studies

The studies fall conveniently into three categories:

(1) Fluorocarbon combustion studies in which

variations of the molecular structure of the fuel are found to be very important in determining flame

speeds and flammability and detonation limits, and

in which the kinetics seems to depend on the ease of an initial attack on the fuel molecule. Larry Ambs,

Jerry Hinderman and David Kittelson did most of

the work here.

(2) Freon combustion studies, in which it appears

that:

(a) Flame speeds of CF3CI-F2 mixtures seem to be

much more sensitive to the preexponential terms

than they are to the exponential term which is the

darling of the thermal theories.

(b) Flame speeds and quenching behavior of

BrCFs-Fo mixtures are much more dependent on

transport phenomena than they are on chemical

kinetics. Donald Parks did this work.

(3) Chlorine combustion studies in which it appears

that even when such phenomena as spontaneous

154
ignition limits and flame speeds can be well correlated

with the formalisms of chemical kinetics the preexponential terms and activation energies are not

necessarily the same, nor do they necessarily have

the same significance. Larry Ambs, Barton Dahneke

and Steve Wiersma contributed to this part.

2.1. Fluorocarbon Combustion Studies

Fluorocarbons burn nicely in chlorine trifluoride

l^llj or fluorine [12, 13, 15] and some will burn

nicely in oxygen as well [13, 14]. The oxygen flames

are probably more complicated than those supported

by fluorine but even these are much simpler than

hydrocarbon flames [16]. When hydrocarbons burn

in air, their combustion behavior, flame speeds,

ignition limits and quenching distances, for example,

are all pretty much the same [17] suggesting that

the nature of the individual fuel is of little importance. Differences which might be caused by variations in the kinetics are washed out by other factors.

Even finely divided non-volatile liquid hydrocarbon

aerosols have about the same flame speeds as those

of the gaseous hydrocarbons. There are so many

reactions possible and going on at once that the

important kinetics in hydrocarbon oxidation re- actions apparently takes place in that part of the

flame that consists of the partially oxidized fragments of the fuel molecules in a melange that is pretty much the same regardless of where the components came from originally. In contrast, flame

speeds, ignition limits and energies, and quenching

distances are quite different in different fluorocarbons [11].

Maximum flame speeds in hydrocarbon-air mixtures occur in compositions near those which give

maximum flame temperatures [17]. In fluorocarbonoxygen [14, 19] and freon-fluorine flames [20],

maximum flame speeds seem to be favored by

equimolar concentrations of reactants even though

this composition may be quite far from that which

gives the maximum flame temperature. This observation suggests again that there is something quite

different in the way kinetics enters into fluorocarbon

combustion in comparison with hydrocarbon combustion.

Perfluoroethane-ClFs can be ignited only with

some difficulty and burns only between relatively narrow composition limits. The flame speed is low

of the order of 10 cm/s. Ferfluoropropane-ClFa, burns

over a range of compositions comparable to that of

hydrocarbons and at hydrocarbon-air flame speeds.

Perfluorocyclobutane-ClFs and cyclo-C4F8-F2 mixtures are very easy to ignite over a wide range of

compositions. They detonate over practically the

entire combustible range. The problem is frequently

to keep them from igniting prematurely. The flame

temperatures of the three fuels are in the same order

as their reactivities so one is tempted to attribute

their combustion behavior only to the effect of flame

temperature on the "global" rate. Many hydrocarbon-based fuels behave as if the details of the

reaction mechanisms are unimportant—the flame

speeds depends only on the temperatures [21] and

global kinetics works well.

Nevertheless, Jerry Hinderman sought to get

some idea of the relative "global" reaction rates of

these fuels burning in fluorine. His technique was to

burn equimolar mixtures of the fluorocarbons with

varying amounts of fluorine present in amounts

insufficient to burn the fuels to completion and to examine the combustion products gas chromatographically. The results were surprising. We weren't

able to deduce relative rates. The differences were

so great that one of a pair of fuels was usually

completely used up before the other had started

reacting. An example of the kind of results that we

observed is shown in figure 1.

c

(D>

D" U 6

0)a

u

xsO 4

1_ CL

too2 2

\\

1 v

\ V

\ V

\ X A

^.THEORETICAL

EQUILIBRIUM CF4

[\:rEMPERATURE

//

//

' /

L-s THEOR^^

'V 1 EQ. CARBON

3S00

3^00

3200

3000°

1_D

2600 a

1_ 2100 ^

I 2 3 4 5 6 7 8 9 10 II 12 13 Equivalence Ratio

Figure 1. Product-composition from a mixture containing

equimolar amounts oj dFe, CiF^ and cycle CiFs

Figure 1 presents the product composition which

resulted from the explosion of equimolar mixtures of

C2F6, CsFg and cyclo-C4F8 with various amounts of

fluorine. The ordinate shows the number of moles of

each of the products which was produced from an amount of mixture which contained one mole of

each of the reactants, i.e., 9 gram-atoms of carbon.

The amount of fluorine is shown on the abscissa,

which gives the equivalence ratio. For example, an

equivalence ratio of 2 means that there was half as much fluorine present as would have been necessary

to burn the fuels completely to CF4 in an isothermal

reaction. (In the adiabatic reaction, dissociation

products are, of course, present.) The figure also

shows the calculated adiabatic constant volume

equilibrium flame temperature's variation with

composition and the low temperature equilibrium

composition of the products. The results are interesting for several reasons.

(1) They clearly illustrate that the differences in

the combustion behavior are associated with vast

155
differences in the rates at which there is an initial

disruption of the molecule under the conditions of

flame propagation. As the fluorine concentration

(and flame temperature) is reduced, C-jF^ stops

reacting first although a 2200 K flame which con- sumes all of the CsFg passes through it. Then CsFg

stops reacting when the flame temperature is much

lower (because adiabatic equilibrium conditions are not achieved) although almost all of the C4F8 is consumed, and finally, unreacted C4F8 begins to appear among the products when the temperature is quite low. Clearly, because of the limited number of

things a fluorocarbon molecule can do in reacting

with fluorine the course of the action depends on

first breaking a carbon-carbon bond. Once that has

happened, what follows happens quickly.

(2) The results also suggest that at lower temperatures (and perhaps early in the reaction) the

principle reaction is C—C scission followed by

recombination of the scission fragments and/or

reaction with available fluorine. There seems to be

very little evidence for breaking of C-F bonds which

would result in disproportunation and formation of

CF4 and carbon. Thus, CsFg decomposes to yield,

ultimately one mole of CoFe and cyclo-C4F8 yields

another. At the lower temperatures, C4F8 yields

primarily the scission products, CsFe and C2F4, just

before, with further reduction in fluorine content the

mixtures stop burning or support only very weak

flames. At an intermediate composition CsFe and

C2r4 are produced in just the right proportions to

suggest that the cyclo C4F8 ring breaks to produce

CsFe and a CH2 radical.

(3) With the higher fluorine concentrations (and

temperatures) the C-F scission reactions which lead

to disproportionation and chemical equilibration

become important as is evidenced by the formation

almost exclusively of CF4 and carbon.

Fm not sure what a global reaction rate means

with this particular family of compounds. The

combustion behavior clearly must depend primarily on the initial pyrolysis of the fuel molecule. There is no way to relate a "global" reaction rate or an

instantaneous reaction rate to temperature because

the concept of an "extent to which the reaction has

gone to completion" is meaningless insofar as an

application to the kinetics of the reaction is concerned, and the rates of the reactions which follow

the initial C-C bond rupture are probably fast

enough that they are relatively unimportant in

determining combustion behavior. When the oxidant

is oxygen, however, the other reactions probably

become important as is suggested by the fact that

the maximum fundamental flame speed in cyclo

C4F8-O2 mixtures occurs not where the flame temperature is highest, but where the equilibrium F

concentration is a maximum [143-

2.2. Freon Combustion

The fluorine supported combustion of CF3CI

demonstrates how differently the kinetics and

transport coefficients interact in deflagration through

hydrocarbon-air and through C] 3CI-F2 mixtures;

the fluorine supported combustion of argon and

hehum diluted CFaBr might be cited to "prove"

that these propagation processes' behavior is almost

completely dominated by the transport coefficients.

Some of the characteristics of CF3CI-F2 flames are shown in figure 2.

% CF,Cl by volume

Figure 2. Burning velocities, equilibrium temperatures, and

product compositions of CF3CI-F1 flames.

The flame temperature maximum occurs near the

rich limit, and the maximum flame speed at a composition corresponding to equimolar amounts of

the reactants. Evidently the augmentation of the

reaction rate constants associated with the higher

flame temperature is not great enough to compensate

for the adverse effect on the stoichiometry in the

colder part of the flame or possibly the higher concentration of the low-molecular-weight F in the

leaner mixtures. In hydrocarbon-air, highest flame

speeds are associated with highest temperatures.

If one were to examine the results for the very

dilute flames of CF3Br-F2 mixtures separately, one might be tempted to say that global kinetics works

nicely in describing the effect of the diluent on the

behavior of CF3Br-F2-Ar and CFaBr-Fs-He mixtures. But the variation we have chosen to describe

is completely dominated by transport properties, as

is evident from an examination of all the data. A

knowledge of the kinetics is unimportant for an156
60

50

1^0

0)E 30

^TM20

a;E

ra ° 10

rj Ll oL 60,

50

XI '^1 30 o o

q: ^ 20

si 8 10

0

Unstable

propagations

CF,

Br

Br,

_ BrF

1

1 50 60 70 % Inert gas dilution

80

Figure 3. Burning velocities, equilibrium temperatures and noninert product compositions oj 65 per cent CF^Br-Fi flames

diluted with argon {lower curve, circles) and helium {upper

curve, x's).

swering only the question we've asked. Figure 3

shows the behavior and properties of such flames.

Mixtures which differ from each other only in the

nature of the monatomic-inert diluent have the same

flame temperatures, product composition, and, very

likely, "reaction-rates". Yet, as can be seen in figure

3, the fundamental flame speeds are quite different.

Dilution by the low-molecular-weight helium at first increases the flame speed because it increases the

rate of heat transfer from the products to reactants.

Further dilution decreases flame speed because it reduces both temperature and reactant concentrations. Argon-diluted mixtures have much lower

flame speeds. Deflagration theories, such as the Zeldovitch-Frank-Kemenetsky-Semenov theory,

which use average (global) reaction rates predict

that, if the global rates are the same, flame speeds

should be proportional to thermal conductivities,

and if one compares the very dilute argon and helium

mixtures described here, the "confirmation" of the

theory is quite striking. One might even be tempted

to use such a flame for measuring the rate of the

reaction.

2.3. Chlorine-Fluorine Studies

The simplest fluorine supported combustion

system we have worked with is the CI2-F2 system.

We have studied low temperature kinetics, spontaneous ignition limits, flame structure and some

fundamental flame speeds. The kinetics of the

reaction is well described by seven elementary

reactions:

I.

II.

III.

IV.

V.

VI.

VII.

F2+M^2F+M

F+C12^C1F+C1

kZ C1 +F2->C1F + F

Cl+F+M^CIF+M

F+F+M-iF2+M

Cl+Cl+Mici2+M

C1+C1F^C12+F

AH = 38 kcal

-3

-23

-61

-38

-58

+3

The rate of either the isothermal reaction or that

used in the interpretation of the spontaneous

explosion limits is well represented by

^[CIF]

dt

(

M-6[Cl2] fcv[ClF]Y/^

In applying the rate expression to the explosion

limits, one makes use of the fact that the system is committed or not committed to an explosion before

the reaction has gone to the extent of 1 percent. The

third term in the denomenator is therefore un- important and if one has started with a stoichiometric mixture, for example, the "rate constant" of

the reaction is

and the activation -energy which can be computed

from tabulated [22] values of the activation energies or reasonable estimates of kl-k^ turns out to be about

one-half of the activation energy of reaction I, which

is in good agreement with the value deduced from

the application of Semenov thermal explosion

theory to the experimental observations, 18.65 kcal.

The magnitude of the activation energy thus has an

appealing mechanistic significance—it is a nice

activation energy. When we consider the isothermal

reaction and wish to describe the kinetic behavior

over a wide range of compositions, the reaction stops

being nice. The "activation-energy" may be said to

depend on the extent of reaction as well as on the

initial composition of the mixture if we try to find a simple formula that will give us a temperature

dependence of the reaction rate. If we assume that

157
the third term in the denomenator is large, we get

the rate expression

it. We can probably find a set of coordinates some- where that makes it basic if we look for it.

dt [GIF]

which correlates the experimental observations with

a correlation-coefficient of about 0.99 for the appropriate values of k'. But if we make a conventional

In & versus 1/T plot to determine the activation

energy of this reaction, it is about 25 kcal. The

difference here results not from a change in mechanism but because in one case our "global kinetics"

describes what is important in the process we are dealing with, and, in the other, what is important just

happens to be well represented by a formula that has

the same form as a simple kinetic equation. It's ironic here that the more appealing result comes from

an abominable combustion process and the less

appealing result comes from the clean, isothermal

experiment.

It's customary in presenting a paper like this to

draw conclusions. I'm reluctant to do it because the

study of the kinetics of combustion-processes is an

extremely lively one now, and you imply that an end has been reached when you draw a conclusion.

In kinetics, as in combustion, one frequently gets

the idea that we attack fundamental problems simply

because they are there. We usually rationalize our

efforts by saying that we can describe anything when

we really imderstand the fundamentals. I'd say the

state of the art of combustion is not that well

advanced. I think its most useful to make use of

whatever information you can get. Most frequently

the kind of information you will need will depend on

where you want to go from where you are. If you can

describe your system well enough so that you can use the "exact" physical quantities that are defined in

terms of the formalisms we like to use, that's fine.

If you can't, but you have a good formula that

gives useful answers using execreble kinetics—so be

[1

[2

[3

[4

[5

[6:

[7;

[8

[9

[10

[11

[12

[13

[14

[15

[16;

[17

[18

[19

[2o;

[21

[22

3. References

Fenimore, C. P., Chemistry in Premixed Flames, 1 (The

Macmillan Co., New York, 1964).

Lewis, B. and von Elbe, G., Combustion, Flames, and

Explosions of Gases (Academic Press, New York, 1951).

Kondrat'ev, V. N., Chemical Kinetics of Gas Reactions

(Addison-Wesley, 1964).

See for example, Fristrom, R. M. and Westenberg, A. A.,

Flame Structure (McGraw-Hill, 1965).

Gerstein, M., McDonald, G. E., and Schalla, R. L., Fourth Symposium (International) on Combustion,

Wilhams and Wilkins, Baltimore, p. 375, 1953.

Hirschfelder, J., Curtiss, C. F., et al.. Third Symposium

(International) on Combusion, Williams and Wilkens,

Baltimore, p. 121 et. seq., 1949.

Gray, B. F., and Yang, C. H., J. Phys. Chem. 69, 2747,

(1965).

Evans, M. W., Chem. Rev. 51, 363, (1952).

Spalding, D. B., Combustion and Flame 1, 287, 296,

(1957).

Lovachev, L. A., and Kaganeva, Z. I., Dokl. Akad. Nauk.

U.S.S.R. 188, 1087, (1969).

Fletcher, E. A. and Ambs, L. L., Combustion and Flame,

8, 275, (1964).

Fletcher, E. A. and Kittelson, D. B., ibid, 12, 119, (1968).

Fletcher, E. A. and Kittelson, D. B., ibid., 12, 164, (1968).

Fuller, L. E. and Fletcher, E. A., ibid, 13, 434, (1969).

Fletcher, E. A. and Hinderman, J., unpublished work in preparation.

Bauer, S. H., Private communication based on shocktube

studies of pyrolysis reaction.

Barnett, H. C. and Hibbard, R. R., et al.. Report 1300 of

the National Advisory Committee for Aeronautics,

Basic Considerations in the Combustion of Hydrocarbon Fuels with Air, p. 257, 1957. Burgoyne, J. H., Selected Combustion Problems, p. 392

(Butterworths, London, 1954).

Matula, R. A. and Agnew, J. T., Comb, and Flame, 13, 1, 101, (1969).

Parks, D. J. and Fletcher, E. A., Combustion and Flame,

13, 487, (1969).

Bartholome, E. and Hermann, C, Z., Elektrochem 54,

^ 165, (1950).

Kondrat'ev, V. N., Rate Constants of Gas Phase Reactions, Academy of Sciences of the Soviet Union, Moscow, 1970.

158
NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Burning of Polymers

C. P. Fenimore and F. J. Martin

General Electric Corporate Research and Development, Schenectady, New York

After a rod of solid polymer mounted vertically in a slowly rising stream of oxidant/inert gas mixture is ignited at the top end, the oxidant content of the mixture may be varied until that

composition is found which will just support burning down the whole length of the sample. For

many polymers this limiting volume-fraction of oxidant can be determined with a precision of

about one percent and it serves as a convenient measure of flammability of the material.

Experiments show that the simplest model for such burning is one in which the oxidant does

not react directly with the polymer, but is consumed in a nearby flame zone which heats and

pyrolyzes the polymer to yield fuel gases. This model appears to be valid for some systems, but

not for others. Experiments may be carried out in which the compositions of the sample and the oxidizing

atmosphere are varied in a systematic way. The measured flammabilities, together with a simple

model of the burning process permit inferences about the manner in which composition variations

affect the burning. The mode of inhibition by flame retardant additives may be inferred in this way.

For example an inhibitor which interferes directly with the gasification of the condensed

phase should work independently of the flame reactions for the simple two-stage candle-like burning; it should still inhibit when the polymer is burned in N2O rather than O2. Such has been observed when chlorine has been substituted into polyethylene.

By contrast an inhibitor which poisons the flame reactions may be specific for reactions with

oxygen and not work with another oxidant such as N2O. Such is the case for Sb203 added to

slightly chlorinated polyethylene.

Effects of different inert gases, of initial temperature and of pressure have also been observed.

Key words: Burning of polymers; flame retardants; flammability.

1. Introduction

The burning of organic solids in air is a complex

phenomenon involving high -temperature gas phase

reactions, convection and conduction of heat,

diffusion of gaseous species, and chemical reactions

in the condensed phase at and in the vicinity of the

burning surface. In general, burning is not a steadystate phenomenon. Some or all of the component

processes will depend strongly upon sample geometry,

size and dimensions of surrounding enclosure (if

any ) , physical and chemical properties of the sample,

nature of ignition and elapsed time from ignition.

From the technological point of view, it is important to be able to determine and to predict the

effects of these various parameters on the burning

process—and particularly on the rate of burning and

the nature of the products.

It is the purpose of this paper to discuss the

information that can be derived from simple experiments with candle-like flames together with a very

simple model for such flames.

2. Model for Candle-like Flames

If one holds a rod of high-molecular-weight

polyethylene (PE) or cast polymethylmethacrylate

(PMMA) vertically in a suitable gaseous oxidizing

atmosphere and ignites the top end, it is found that

the rod burns much like a candle [1, 2J.^ Such a flame for polyethylene burning in oxygen-deficient

air is shown in figure 1. A clear molten dome at the

top of the rod is surrounded by a thin blue flame with

a yellow cap arising from thermal emission from soot

particles. A few drops of molten polymer have run down the side of the stick and frozen.

Burning experiments of this kind are conveniently

carried out in a vertical Pyrex chimney about 7.5

cm I.D. by 45 cm high, although the dimensions are not critical. The specimen is mounted vertically on

the chimney axis roughly midway up. The oxidizing

atmosphere is prepared outside the chimney by

metering and mLxing the components in a flowing

system; it then enters the chimney from the bottom

through a bed of glass beads which serves to smooth

the flow. The upward gas flow is kept steady in the

range 3-12 cm/s, which serves to sweep smoke and

products out of the chimney. Test equipment of this sort is now commercially available.

The size of the flame varies with the oxygen

content of the air and becomes very large in pure

oxygen. As the oxygen content of the air is decreased,

the flame size decreases until a concentration is

* Figures in brackets indicate the literature references at the end of this paper.

159
Table 1. Selected oxygen indices

Figure 1. Polyethylene burning close to extinction

limit at 1 atm.

Oxygen index = 0.175. Upward gas velocity 8 em/sec.

Stick about 8 mm wide.

GLASS COLUMN (MINIMUM DIMENSION i.5CM H X 7.5CM 10)

15

1. BURNING SPECIMAN

2. CLAMP WITH ROD SUPPORT

3. IGNITER

4. WIRE SCREEN

5. RING STAND

6. GLASS BEADS IN A BED

7. BRASS BASE

8. TEE

9. CUT-OFF VALVE

ORIFICE IN HOLDER

PRESSURE GUAGE

PRECISION PRESSURE REGULATOR

FILTER NEEDLE VALVE ROTAMETER

ALTERNATE FLOW

MEASURING DEVICE

FiGXJKE 2. Typical apparatus for oxygen index determinations.

Hydrogen (4) 0.054

Ethylene (4) 0.105

ra-Paraffins (C3 to Cio) (4) 0.127 to 0.135

Benzene (4) 0.133

Polyoxymethylene 0.157

Delrin or Celcon

Kitchen Candle 0.16

Polymethylmethacrylate 0 . 1 74

Rohm and Haas Plexiglas

Polypropylene 0 . 1 74

Hercules Profax 6505

Asbestos-filled "slow-burining" 0.205

Polypropylene

Union Carbide JMDC-4400

"Self-extinguishing" Polypropylene 0.282

Union Carbide JMDA-9490

Polyethylene

Allied Chem 1220 or Phillips Marlex 5002 0.174

Marlex filled with 50 wt % AI2O3 0 . 196

Marlex filled with 60 wt % Al.Os-SHaO 0 . 302

Polystyrene 0.181

Epoxy* unfilled 0.198

Filled with 50% AI2O3 0.250

Norton 38900

Filled with 60% AI2O3 -SHsO 0 .408

Alcoa C-333

Polyvinyl Alcohol 0 . 225

DuPont Elvanol 70-05

Polyvinyl Fluoride 0.226

DuPont Tedlar

Chlorinated Polyether 0.232

Hercules Penton [—CH2C(CH2C1)2CH20—

]

Polycarbonate 0.26-0.28

General Electric various clear Lexans

Polyphenylene Oxide 0 . 28-0 . 29

General Electric [—C6H2(CH3)20—

]

Silicone Rubber 0.30

SiUca-fiUed General Electric SE9029

Polyvinylidene Fluoride 0.437

Kynar

Polyvinyl Chloride 0.47

Geon-101

Polyvinylidene Chloride 0.60

Dow, Saran

Polytetrafluoroethylene 0 . 95

DuPont, Teflon

Loosely Woven Fabrics

Cotton 0.185 ± .005

Wool 0.238

Synthetic 60/40 = vinyl chloride/ 0.298

acrylonitrile

Dynel

* Epoxy was 100 parts Epon-826, 75 parts hexahydrophthalic anhydride, 1 part DMP30, cured 16 hours at 400 K

160
reached such that the flame after ignition will no

longer propagate down the whole length of the stick.

Experimentally one can determine a limiting oxygen

concentration such that a small flame propagates

down the whole length of the stick in a nearly

steady -state manner; any reduction in oxygen would

result in extinction. This limiting concentration

expressed as the volume fraction of oxygen in the

atmosphere is called the oxygen index of the material

burned. It has proven to be a convenient rating of

polymer flammability, and a standard ASTM test procedure D2863-70 has been set up for determining

it p]. Some oxygen index values are given in table 1. An apparatus sketch taken from the ASTM test instructions is shown in figure 2. The lower the index, the more flammable the

material. There are some experimental limitations

which will be discussed later, but the index is often

reproducible to about one percent, and for many

(but not all) materials largely independent of

specimen dimensions and upward gas velocity over

considerable ranges of these variables. (The same

index was previously defined and used by Wolfhard

and coworkers [4, 5 J as a measure of flammability of

volatile fuels and gases burning as diffusion flames.)

The flame for a rod of cast polymethylmethacrylate

(PMMA) at the lower oxygen limit looks much like

that for PE except that there is less melting and

dripping and more bubbling at the burning surface.

The flame for PMMA propagates down the rod with

very little change in appearance or velocity with

time; and hence closely approaches a true steady

state. When a rod of cured epoxy resin heavily filled with hydrated alumina is burned at the lower oxygen

limit, a zone of glowing alumina works its way

steadily down through the specimen with little visible flame; this perhaps represents an example of a surface combustion process [6].

Such observations suggest a simple, steady-state

model for candle-like flames burning close to the

extinction limit. The gaseous oxidant is supposed

not to react directly with the polymer but is con- sumed in a nearby combustion region which heats

and pyrolyzes the polymer, and thus generates the

fuel for the combustion region. The combustion

region may be a diffusion flame surrounding the

polymer as discussed above for PMMA; or it might

be a region of surface combustion, as with a polymer

filled with a nonburning inorganic material.

Gas analyses and temperature measurements

through the flame above rods of PMMA, and POM

(burning at reduced pressure to spread out the flame)

support the model [7J. The materials are known to

pyrolyze largely to the monomer. Some of the data

are summarized in table 2; experimentally determined temperatures and gas concentrations above

burning PMMA are shown in figure 3. Within the

approximations involved, the estimated heat transfer

by conduction and radiation from the flame to the

polymer surface is sufficient to warm the surface of

the polymer to the observed temperature and to

depolymerize the material to form fuel gases at the

observed rate; in this way steady-state candle-like

burning is sustained. In the case of PE the primary

pyrolysis products are unknown so that it cannot be

certain that the model applies, but it seems likely.

On the other hand, as seen in table 2, when warm

PTFE rods are burned in concentrated 02/Ar

mixtures at 1 atm, the heat conducted and radiated

Table 2. Heat fllow from flame to the condensed phase for polymers burning in Ar-02 atmospheres*

Polymer (cross section)

of rod)

Methylmethacrylate

0.36 cm2

Formaldehyde

0.36 cm2

Tetrafluoroethylene

11.3 cm2

4.9 cm"

2.8 cm2

1.2 cm2

Ethylene

0.36 cm2

Surface

temperature

K

850

750

920

920

920

900

Fraction of Car- bon as Monomer, 0.1 cm

above tip

70

33

21

13

50

30

Burning Rate

cm/sec X 10'

2.6

2.3

5

7

7

8

1.8

Heat Conducted to Polymer

kcal per monomer unit

Obs.

30

15

13

20

Calc*

38

17

50

50

50

50

Observed Heat

Conducted as Fraction of

Total Heat

of Combustion

0.05

0.14

0.12

0.07

* Burned at 0.07 atmos P in 20-22 percent O2 except for polytetrafluoroethylene which was burned at 1 atmos in 88-100 percent

O2 after being preheated to 400 K.

** Heat of gasification to monomer (21.4 kcal for methylmethacrylate (12), 13 for formaldehyde (13), 41.1 for tetrafluoroethyl- ene (14) plus estimated heat to warm to surface T gives this column.

161
from the flame back to the condensed phase falls far

short of that required to further warm and depolymerize the polymer at the rate necessary to sustain

the observed steady-state burning. In this case heterogeneous reactions with species in the gas

phase must contribute to the polymer gasification.

The gaseous species may be O2 or O or more likely in

some cases it is flame-generated F atoms which are calculated to be a major flame species at equilibrium. Thus, the simple candle-like burning model in

which only heat is transferred by conduction and

radiation from the flame to the burning surface is not

valid for PTFE [8].

3. Effects of Different Inert Gases in the

Atmosphere

If one burns in a candle-like manner similar

samples of a given polymer in atmospheres that differ

in the inert gas present, it is found that the mole

ratios oxygen/inert gas at the lower oxygen limit

vary linearly with the heat capacities of the inert

gases when the latter are Ar, N2 or CO2 [9]. Data are

_, I800°K

BOUNDARY OF DIFFUSION FLAME I 05 cm ABOVE TOP OF BURNING POLYMER

C2 HYDROCARBONS PLUS 0.2

1

Ar/2

H2 PLUS 0.2

1300

800

1 2 3 cm DISTANCE ABOVE TIP OF BURNI NG POLYMER

Figure 3. Gas composition and temperature along the central axis above a polymethyl methacrylate rod, 0.36 cm^ cross

section, burning at 5 cm of mercury pressure in argon- oxygen, 20 percent O2.

Distances are measured from the burning tip, and the dashed Hne at

1.05 cm shows the position of the visible diffusion flame. The Hmiting

temperature gradient at the surface is 2500 K cm"'. For clarity, H2

and C2 hydrocarbons are displaced 0.2 upwards.

//o,l545C\

\a,l470C/

^ 1240 C

C0 = C02

I , polyoxymethylene, AHox = l03 kcal/mol O2 a , polymethylmethocrylote , AHox = 90.3 o , polyethylene, AHox =89.4

0 .chlorinated polyethylene, C2H3boC£o20'^Ho)(=89.I

, 96wt.%C2H38oCfo.20 + ^«' % Sb2 03, AHox =89.1

He, indicates Helium used for this point J I I 0 - 2 _ 4 6 Cp/R of inert qos

Figure 4. Effect of heat capacity of inert gas on (oxygen /inert

gas) ratio. AHo\ is estimated heat of combustion per mole of O2 for reactions written

to produce equimolar quantities of CO and CO2. Calculated flame temperature to right of curves.

shown in figure 4. It is noteworthy that curves have

an intercept on the Uok/n , axis.

We propose the following model to explain these

data:

a. For stability the flame must attain a minimum

temperature which is independent of the inert

gas in the atmosphere.

b. The stoichiometry of the overall flame reactions

(and hence the heat release per mole of oxygen

consumed) is independent of the inert gas.

c. The flame is nonadiabatic. The mean rate of

heat loss per mole of mixed hot products and

inert gas is independent of the inert gas. By

heat loss we mean heat transferred outside the

flame by radiation or conduction to the at- mosphere. It does not contribute to heating

the preflame gases.

For this model a simple heat balance may be written

as follows:

[nprCp,pr-\-niCp,i']{T'2— Ti) =noxAfl'ox— (ipr+ni)^

(1)

where:

ripr = flux of products from flame zone, moles/s. n, =flux of inert gas through flame zone,

moles/s.

162
Tiox =flux of oxygen into flame zone, moles/sec.

Cp = heat capacity per mole.

T2 = flame temperature.

Ti = initial temperature of sample and at- mosphere.

A//ox = enthalpy decrease for reaction of one mole of O2 at Ti.

Q =mean heat loss per mole from flame zone

gases. It is taken to be constant for flames

of a given material and burning at the

same initial conditions.

Equation (1) may be rearranged to give:

"ox _ 2 T2—T1

rii ~ B

(2)

where B = AHo

iipr

"ox

l-Q+Cp.pr{Ti-T^)-] and

represents the net heat available to warm the inert

gas per mole of O2 consumed. (Here heat loss per

mole from product gases alone is taken to be equal

to that from the whole hot mixture; it is an approximation.) The quantity Q/B represents the fraction

of heat available to the inert gas which is lost from

the flame.

To the three points determined for Ar, N2 and

CO2 one may fit a straight line; this in turn gives

values for Q/B and (T2— Ti)/5. For an assumed

value of the heat release, AHok, one can then solve

for Q and for T2—T1. The temperature values in

figure 4 were obtained in this way. Clearly if the flame were adiabatic, Q would be zero and a curve of

nox/ni versus Cp,, would pass through the origin.

This is not observed.

The derived temperatures and heat losses appear

reasonable and lend support to the assumptions

made in the model. Estimates of the fraction of the

heat of combustion lost run generally between 10

and 40 percent.

Thermal conductivity values for Ar, N2, and COo

(which are generally the major components of the

product mixture) vary by a factor of two or less. On

the other hand, the thermal conductivity of He is higher by a factor of eight or so. The nox/ni values

measured in helium/oxygen for PE, PMMA and

chlorinated PE are 1.50 times those measured in

Ar/02 even though He and Ar have identical molar

heat capacities. For chlorinated PE plus 4 wt percent

Sb203 the corresponding ratio is 1.86. It seems

reasonable to attribute the higher values in helium

to greater heat losses from the flame. The higher

sensitivity of chlorinated PE plus SboOs to He may

be attributed to slower gas phase reactions for this

material and hence relatively greater heat losses.

Other work supports the idea that Sb203 inhibits by

slowing down gas phase reactions.

Finally, from a practical point of view, we mayrate the various inert gases for effectiveness in

quenching flames in the order of ascending nox/n.

values; the higher Uox/rii, the more effective the gas

as a flame quencher. On this basis, CO2 > He > N2 > Ar

for the flames studied. This is the same order of

POLYTETRAFLUOROETHYLENE RODS

OF VARIOUS CROSS-SECTIONAL AREAS

-0.4

-08

POLYCARBONATE SHEET

OF VARIOUS THICKNESSES

POLYMETHYL METHACRYLATE SHEET

OF VARIOUS THICKNESSES

-1.6 -0.8 0 LOG THICKNESS OF SHEET OR

LOG CROSS -SECTIONAL AREA OF RODS

0.8

Figure 5. Effect of size of burning specimen.

Polycarbonate data from Ref. 15.

effectiveness as that reported for increasing the

quenching distance in helium-oxygen-inert gas

mixtures [^10].

4. Effect of Some Variables on the

Oxygen Index

The indices of polymers are nearly independent of

the oxidant gas velocity in the range 3 to 12 cm/s

possibly because flame-generated convection currents

in the neighborhood of the test specimen always

exceed these values anyway. However, there is a small dependence on gas velocity and for comparison

purposes a gas velocity of 4 cm/s has been set for the

ASTM-D-2863-70 test procedure.

For PE, PMMA, polystyrene (PS) and possibly

for most other unfilled thermoplastics, the indices are independent of sample thickness in the range 0.2

to 1.0 cm and decrease only slightly for smaller

thicknesses. However, for polycarbonate (PC) sheet

less than 0.2 cm thick, the index falls rapidly with

decreasing thickness; it varies as thickness to about

the 0.13 power. Data are shown in figure 5. Polytetrafluoroethylene (PTFE) rods show a similar

dependence of index on cross-section area even for

large specimens [8]. It is possible that the greater

dependence of index on sample size is associated

with a mode of burning involving heterogeneous

reactions in the flame zone. Despite the size effect, reproducible values of the

index can be obtained for most thin films and fabrics

163
1.0

0,8

X

LUQ

>-X

o

0.6

0.4

0.2

POLYTETRAFLUOROETHYLENE

POLYVINYL CHLORIDE

60%

AlgOj'BHgO

FILLER

EPOXY

POLYETHYLENE OR

POLYMETHYL METHACRYLATE

200 400 600

INITIAL TEMPERATURE - °K

FiGXJKE 6. Effect of initial temperature.

which do not melt too badly ahead of the flame [11].

This is conveniently done on a strip of material about

15 by 5 cm (dimensions not critical) clamped along

the vertical long edges in a light metal U-frame. The

sample is ignited uniformly along the top open end of

the U; and the index is determined in the same

manner as for solid rods. The metal frame serves to

support the sample and to suppress rapid burning

down along the otherwise free edges. Some results are included in table 1. By jacketing the apparatus and precooling or preheating the gas and polymer sample, the effect of

initial temperature on the index can be studied.

Data are shown in figure 6. By modifying the apparatus to permit burning

under a controlled pressure, the dependence of index on pressure can be determined. In figure 7, the index

for polytetrafluoroethylene is seen to vary nearly

inversely as the | power of the pressure. For POM,

PE, and PMMA the dependence is much smaller.

One might wish to evaluate an index for upward

burning of a sample after ignition at the bottom end.

Such an arrangement would maximize convective

heat transfer to the unburned material, in contrast

to the usual downward burning where convective

heat transfer is at a minimum. However, only a few

materials will withstand the drastic preheating

associated with upward burning long enough to

permit meaningful measurements; more often the

samples drip and sag too badly. For cast PMMA

which can be burned in both directions the usual

index for downward burning is 0.174=b.001, but for

upward burning the index is only 0.151zt.004.

5. Some Comments on Flame

Appearance

Although there are certain similarities in appear- ance of candle-like flames for many different

materials, there is also a wide variation in the

detailed characteristics. Some samples are topped

by a quiescent dome of clear liquid showing no

formation of gaseous pyrolysis products below the

surface; others show extensive bubbling in the

interior. Some samples burn without char or soot

formation; some generate much char and soot—and

the appearance of the chars differs. Some chars are lacy and stand upright in the flame; others are massive (polycarbonates) and enclose considerable

partially reacted polymer. The latter are prone to

topple and fall down along the imburned sample as it softens below the char.

Some flames burn quietly; others very erratically as bubbles of fuel gases burst out through the

burning surface in an unpredictable manner. Some polymers melt and drip badly. The melt may

rim down and refreeze on the sides of the imburned

sample; in this way heat is removed from the flame zone and results in an oxygen index higher than that'

for a chemically similar but non-melting sample. A

high molecular weight polypropylene (Hercules

Profax 6505) burns with an oxygen index of 0.174;

a low molecular weight polypropylene (Epolene

N15) has an index of ^0.23. The latter burns with a

light blue flame which rides down the specimen on a

Polytetrafluoroethylene

o

E

-0.4

-0.8 - o

o

o

-1.2

0.63 cm rod

of 400 »K

1.27 cm rod

at 298 "K

Polymethylmethacrylate

0.63 cm rod

at 298 "K

-0.4 0 0.4

log (P/Pq), Pq '-

I atmos.

Figure 7. Effect of pressure.

164
wave of molten polymer leaving the center standing

unburned. Its high index reflects the extra heat loss

due to easy melting. When the same material is made into a candle with a ceramic wick and burned

off the wick, the index is only 0.177=b.005. Some

materials melt and run so badly that it is not experimentally possible to measure an index; polyethylene

film is an example. Fortunately there are not too

many of these.

As the oxygen content of the atmosphere is reduced gradually below the critical level, the flames

of many polymers which burn without solid residues

(such as polyethylene, polymethylmethacrylate,

polyoxymethylene and polytetrafluoroethylene)

slowly shrink and the leading edge recedes until just

before extinction only a small blue cap exists above

the sample. For other polymers, particularly some

containing antimony oxide (Sb^Os) plus halogen,

only large flames are stable. As the oxygen is reduced,

a violent flicker appears. It looks as though the

flame is suddenly blown back from the leading edge,

then jumps forward again momentarily but cannot

stabilize itself and the phenomenon repeats rapidly

through many cycles until finally the flame although

still large is blown ofi^ completely. The details of the

phenomenon are not understood.

When one considers the great variations in flame

behavior, it is rather amazing and gratifying that

reproducible indices can be determined for so many

different materials.

6. Modes of Inhibiting Polymer

Flammability

Having reviewed the characteristics of candle-like

flames, we will now discuss their use as a tool to

study the burning process. In particular we will be

interested in ways to make polymers more resistant

to burning in air. In the simple candle-like model, heat flux from the

flame to the condensed phase is sufficient to gasify

the polymer at a rate adequate to maintain the

flame. If the polymer is modified to require a greater

heat flux in order to provide the same flow and

quality of fuel gases, a higher flame temperature is required; it is achieved by reduction of the quantity

of inert gas in the atmosphere and hence in an

increase in the oxygen index.

Likewise if the polymer is modified to yield species

in the fuel gases which retard or poison the flame

reactions, a higher flame temperature is required to

maintain the flame in the face of unavoidable heat

losses; in this case too, the higher flame temperature

would be achieved by reducing the N2/O2 ratio in

the atmosphere and thereby increasing the oxygen

index.

These two modes of inhibition should be distinguishable. An inhibitor that reacts with the

condensed phase to alter the pyrolysis reactions

would not be expected to be equally effective in all polymers, since they differ widely in their chemistry;

but, in those polymers where it is effective, it should

work independently of the nature of the flame

reactions. It should inhibit when the polymer is burned in another oxidant than oxygen, for example,

N2O. This is so because only heat is required from

the flame. It will be shown below that phosphorus

compounds in epoxy inhibit in this way.

On the other hand, the chemistry in the flame zones of different polymers burning in nitrogen/

oxygen mixtures is often quite similar; and an

inhibitor which poisons the flame reactions would be

expected to work about equally well in polymers of

different types, although there appear to be some

exceptions. As discussed below, bromine or bromine

compounds inhibit in a wide variety of fuels and,

therefore, we believe they poison the flame reactions.

However, the flame chemistry can be readily

altered by using another gaseous oxidant such as N2O. Then the nitrous oxide index is defined to be

the volume fraction of N2O in the atmosphere

required to sustain steady downward flame propagation under the same conditions specified for the

oxygen index. It is a measure of flammability in

N2O/N2 mixtures. Inhibition in O2/N2, but not in

N2O/N2 flames as measured by relative changes in

the oxygen and nitrous oxide indices would indicate

poisoning of the flame reactions.

We will summarize briefly the conclusions reached

about the mode of inhibition by several classes of

inhibitors.

7. Phosphorous Compounds

Phosphorus compounds inhibit weakly in polyethylene (see figure 8 and references [1, 7]); but

triphenyl phosphine inhibits strongly in epoxy as

0.4 1-

SbaOj

O TRIPHENYL ANTIMONY ^ TRICRESYL PHOSPHATE

NO HALOGEN

0.3 o

o

0.005 0.01 ATOMS OF Sb OR P PER C ATOM.

Figure 8. Phosphorus or antimony compounds added

to polyethylene (Marlex 5002) without halogen.

165
°2/(N2 + 02)

POLYETHYLENE BURNT IN ATMOSPHERES

CONTAINING HCl OR Clj

0.5

Figure 9. Inhibition by chlorine in chlorinated polyethylene.

shown in figure 10. Furthermore triphenyl phosphine

in epoxy is effective for burning both in oxygen and

in nitrous oxide; therefore we conclude that it inhibits primarily by affecting the pyrolysis reactions

in the condensed phase.

8. Chlorine Compounds

Oxygen and nitrous oxide indices for chlorinated

polyethylenes are shown in figure 9. Inhibition is effective in both oxidants and therefore probably

does not involve poisoning of the flame reactions.

This is consistent with the observation that addition

of large amounts of CI2 or HCl to N2/O2 atmospheres

only slightly increase the oxygen index for polyethylene as shown by the bottom curve in figure 9. Gas samples collected just above the burning

surface of chlorinated polyethylene contain no

molecular oxygen, no chlorinated hydrocarbon, but

more benzene and acetylene the greater the degree of

chlorination of the polyethylene. (In all cases about

70 percent of the C atoms formed CO and CO2.)

The chlorinated polyether, penton, and chlorinated

polyethylene having the same Cl/C ratio of 0.4 have

quite different oxygen indices—namely 0.23 and

0.39 respectively. In contrast to chlorinated polyethylene, Penton yielded some chlorinated hydrocarbons in the products.

Thus, all the evidence indicates that CI compounds

inhibit in the condensed phase by affecting the

pyrolysis reactions.

Flames of chlorine-containing samples are sooty

and form much char, which may tend to insulate

the condensed phase from the flame. Whether or not

this effect is important is not known.

9. Bromine Compounds

For n-hexane, polyethylene and for epoxy resin the

addition of bromine or bromine compounds to a like

Br/C ratio results in each case in about the same

fractional increase in the oxygen index. This similar

effect for quite different fuels indicates that bromine

acts primarily as a flame poison.

10. Antimony Compounds

The oxygen index of partly chlorinated polyethylene or of polyethylene containing bromine

compounds is markedly increased by the addition of

antimony in the form of Sb203 or as triphenyl stibine;

but no increase is observed in the nitrous oxide index.

Analysis of the residues from partly burned

samples show that antimony is only effective when

it vaporizes into the flame, and then only when the

oxidant is oxygen. Thus, we conclude that antimony

compounds yield flame poisons [73.

0 0.01 ATOMS OF Sb OR P PER C ATOM

Figure. 10. Triphenyl phosphorus or triphenyl antimony added

to epoxy containing no halogen.

Epoxy Bame as in table 1.

166
It should be noted that phosphorus kills the

antimony-halogen synergism in polyethylene by

preventing the vaporization of antimony.

Antimony may be added in the form of the

triphenyl compound in the absence of halogen. When

added in this manner to polyethylene (no halogen)

the effect is small as shown in figure 8; the maximum

increase in the oxygen index is 15 percent. However,

when added to halogen-free epoxy, a large inhibition

is observed for burning in oxygen/nitrogen at- mospheres, but much less in nitrous oxide/nitrogen.

The maximum increase in oxygen index is 66 percent,

but only 13 percent for the nitrous oxide index.

Again the evidence indicates that the antimony

mainly poisons the flame reactions. Here is a case where a flame poison is not equally effective in

different polymers, i.e., polyethylene versus epoxy.

The implication is that triphenyl stibine vaporizes

from the burning epoxy without the aid of halogen,

but not from burning polyethylene.

11. Inhibition by Fillers

Inorganic fillers accumulate at the burning surface

and appear to smother the flame by interfering with

the heat flux to the polymer. At high enough oxygen

levels, the cap of accumulated residue will split open

or be tipped off to one side, and the flame will flare

up until depressed again by the formation of another

0.8

0.6

0.4

/

I

/

/ FILLER = Al203 '3H20 / ALCOA C -333

o /

°2/(N2+ Oj)

r

/AUOj'SHpO

AUO, , NORTON 38900

1^ a

0.2,

AI2O3 , LINOE A

30 60 WT PERCENT FILLER

Figure 11. Inhibition by fillers in epoxy resin.

residue cap. Despite the cyclical character of the

burning, it is often possible to determine a reproducible oxygen index; it is usually but not always

higher than for the unfilled polymer.

With larger proportions of filler a standing ash

may be formed which roughly maintains the shape

of the original sample; steady burning is generally

observed in these situations; and heterogeneous

reactions may be important or even predominant.

In figure 11 it is seen that inert fillers of alumina

(or quartz) at high concentrations in epoxy may raise

the oxygen index by as much as 40 percent; but the

effect levels off and does not continue as more filler

is added. On the other hand, at concentrations

greater than 40 weight percent, hydrated alumina

continues to show a steadily increasing oxygen index;

at 60 weight percent Al203'3H20, the index is twice

that of the unfilled epoxy. The nitrous oxide index

behaves similarly [6J. The effect presumably arises

from heat absorbed in the partial dehydration of

Al203'3H20. Since the effect is observed in both

oxygen and nitrous oxide atmospheres, it probably

does not involve gas phase reactions with the oxidant.

Very small amounts of soot and char are formed.

12. Concluding Comments

If the above discussion is correct, we conclude

that simple candle-burning experiments in controlled

oxidizing atmospheres can often establish the mode

of inhibition and the optimum quantity of chemical

additives for burning under these conditions.

Appropriate analysis can sometimes be made to

check these conclusions.

Of course, it would be nice to know more. However, this will involve a greatly expanded knowledge

of the details of pyrolysis reactions at and below the

burning surface and knowledge of the complex flame

reactions. These are difficult but challenging areas. In another direction we need to understand better

the relationship between the indices observed for

candle-like burning and other measures of flammability. Hopefully such knowledge will lead

eventually to improved design criteria to control

flammability in complex structures.

13. References

[1] C. p. Fenimore and F. J. Martin, Combustion and Flame,

10, 135 (1966).

[2] C. P. Fenimore and F. J. Martin, Modem Plastics

(Nov. 1966).

[3] J. L. Isaacs, General Electric Co. Report TIS 69-MAL-13

(Aug. 20, 1969).

[4] R. F. Simmons and H. G. Wolfhard, Combustion and

Flame, 1, 155 (1957).

[5] A. R. Hall, J. C. McCoubrey, and H. G. Wolfhard,

Combustion and Flame, 1, 53 (1957).

[6] F. J. Martin and K. R. Price, J. Appl. Poly. Sci., 12, 143

(1968).

[7] C. P. Fenimore and G. W. Jones, Combustion and Flame,

10, 295 (1966).

[8] C. P. Fenimore and G. W. Jones, J. Appl. Poly. Sci., 13,

285 (1969).

[9] F. J. Martin, Combustion and Flame, 12, 125 (1968).

167
[10] C. E. Mellish and J. W. Linnett, Fourth Symposium

(International) on Combustion, p. 413. Wilkins &

Wilkins: Baltimore (1953).

[11] F. J. Martin, General Electric Co. Report 68-C-422

(Nov. 1968).

[12] K. J. Ivin, Trans. Faraday Soc, 51, 1273 (1955).

[13] J. F. Walker, Formaldehyde, Reinhold, N.Y. (1951).

[14] J. C. Siegle, L. T. Muus, Tung-Po Lin, and H. A. Larsen,

J. Pol. Sci., A2, 391 (1964).

[15] K. B. Goldblum, SPE Journal, 25, 50 (1969).

Discussion

Raymond R. (Hooker Research Center)

The conclusion in Dr. Fenimore and Dr. IVIartin's

paper that fire retardant inhibition in the condensed

phase should be independent of the oxidizing medium

implies an assumption that any interactions between

the hot molten surface of the polymer and the

oxidizing gas are unimportant or unchanged by a change in oxidant. Can you comment upon the

validity of this assumption?

Is the temperature profile of the N2O-N2 flame the

same as that observed with the O2-N2 flame?

What are the carbon dioxide and carbon monoxide

concentrations on the surface of burning polyethylene? Do they change in the presence of fire retardant additives?

F. J. IMartin:

1. The simple candle-like model postulates no

direct reaction between the polymer surface and the

oxidizing gas; reaction at the surface occurs as a

result of heating in a local atmosphere of pyrolysis

products and combustion products. Furthermore,

it is assumed that reactions between the combustion

products and the condensed phase do not significantly influence the pyrolysis reactions at the

surface. The question is, "Is this a valid model for

actual situations?"

For many polymers burning in oxygen or nitrous

oxide mixed with an inert gas, we believe the simple

model is valid. Analyses of gases above polymethylmethacrylate, polyethylene, and polyoxymethylene

burning at a reduced pressure (6 cm of Hg) in

argon/oxygen, 80/20 gave compositions expected

for pyrolysis reactions in an inert atmosphere.

Furthermore, the heat flux into the polymer,

estimated from the measured temperature gradient

above the regressing surface, was that required to

warm and gasify the polymer without the necessity

of surface reactions with any product species (Ref. 7)

Presumably the validity in these cases arises from

the similar inert products (primarily H2O, CO2, and

CO) found above the pyrolyzing surface. If one were

to choose an oxidant yielding widely different

combustion products (such as F2 producing HF,

CF4, etc.), the inertness of the products with respect

to the polymer-surface pyrolysis reactions would

have to be established. Even if one were to add

increasing amounts of inhibitor to a polymer, like

samples of which were then burned in two different

oxidants chosen such that for one oxidant there was

no surface-product interaction and for the second

there was reaction, comparable inhibition in both

oxidants as measured by changes in the suitable

oxidant index would be evidence for probable action

in the condensed phase either by a strictly thermal

pyrolysis or surface interaction with products; it would not be possible to distinguish these two possibilities by this test alone. (Of course we have not

here ruled out the possibility, although it seems

remote, that the added inhibitor might be equally

effective in both oxidants as a poison for the gasphase flame reactions.) We have not tried this

experiment.

2. Although temperature profiles above polymers

burning in N2O-N2 mixtures have not been measured,

we anticipate that they would be similar to those for

burning in O2 mixtures.

3. The fractions of carbon measured as CO2 and

CO in gas samples collected 0.1 cm above chlorinated

polyethylenes burning in Ar/02 mixtures at 15 cm of

Hg are as follows (Ref. 7)

Fraction of C as Polymer

Composition CO2 CO

Polyethylene, C2H4 0.38 0.30

Chlorinated to C2H3.SCI0.2 .42 .27

Chlorinated to C2H3.4CI0.6 .50 .20

With chlorine incorporated into the polymer

chain, it is seen that the fraction of carbon appearing

as oxides remains nearly constant at around 0.69

at least until the Cl/C atom ratio is increased to 0.3;

with increasing CI, the fraction appearing as CO2

increases at the expense of the CO. At the same

time, the fractions appearing as C2H2 and benzene

increase.

We do not have such data on other inhibitors in

polyethylene.

Robert F. IVIcAlevy, III (Stevens Institute of

Technology)

Although the "oxygen index" represents a con- venient means of ordering some aspects of polymer

combustion it does not appear to be sufficiently

fundamental to relate it to the actual mechanisms of

ignition, flame spreading and burning of polymers.

These will have to be studied individually, and in

depth, before a rational "inflammability index" can be evolved. The commentor's laboratory is engaged

in these fundamental studies and copies of reports

dealing with this work are available.

F. J. IMartin:

Certainly much work needs to be done to elucidate

the complex processes of ignition and burning. Only

168
time will tell whether or not the oxygen index will

prove to be a useful and meaningful tool in this

difficult field.

William C. Kuryla (Union Carbide Corporation):

Dr. Martin in his excellent paper on the burning of

polymers, particularly as measured by the oxygen

index apparatus, noted on page 168 of the preprint

that: "phosphorus kills the antimony-halogen

synergism in polyethylene by preventing the vaporization of antimony". Has this been noted for any

other class of polymers, and if so to what degree?

I believe this is a very important observation in that one can easily envision a formulator thinking to

himself: "I'll make this article really flame retardant

by not only putting in the usual antimony/halogen

components, but I'll throw in some phosphorus

flame retardant as well!"

F. J. Martin:

We have not studied other polymers to see whether

or not phosphorus kills the antimony-halogen

synergism in them as it does in polyethylene. It would not be surprising to find similar behavior in

other materials. This observation points up the

desirability of learning more about the chemistry

of inhibitor/polymer systems.

Bernard Miller (Textile Research Institute)

The authors mention the desirability of obtaining

oxygen index values for the more vigorous upward

burning of a sample, which should accentuate the

effect of convective heat transfer ahead of the flame

front. Our experience agrees with their statement

that this kind of experiment generally cannot be

carried out with a conventional oxygen index test arrangement. However, we have made what we

believe to be related measurements by arranging for

the gas mixtures to flow in the reverse direction

(down the test chamber) while still burning the test sample from the top. Thus, all products of combustion should be moved ahead of the flame. The

resultant data, for textile fabrics, have proven to be

quite interesting. Instead of the expected straightforward decrease in O.I. as described by the authors

for cast PMMA, we have observed a variety of O.I.

changes that are dependent on the gas flow rate. Using about the recommended flow rate (actually

4.8 cm/sec) we have found that downward gas flow

often gave a higher O.I. than upward flow, viz:

Material Upward flow Downward flow

O.I. O.I.

Cotton

Acrylic

Wool

Cotton +3.6%

(NH4)2HP04

Polyester + 5.4% tris0.180

.182

.252

.228

.238

0.186

.188

.259

.244

.265

In contrast, nylon and polyester have the same

index for gas flow in either direction, but in no case have we found a decrease in O.I. with a downward

gas flow of 4.8 cm/s. As stated by the authors, the normal O.I. value is independent of flow rate over the range 3-12 cm/sec,

but this is not the case for the reverse flow. At twice

the previously mentioned rate, all except two of the

above materials recovered their conventional O.I.

values. Wool dropped down to 0.246, while the

treated cotton (-|-DAP) was still high at 0.236.

So that only in one instance (wool at 9.6 cm/sec

downward gas flow) have we observed a decrease in

O.I. that could possibly be attributed to increased

convective heat transfer. Instead it seems likely that

the burning of most of these materials is inhibited by

the generation of non-combustible gaseous byproducts which are not removed by the lower

velocity downward gas stream as they are at higher

flow rates or when the gas is moving in opposition

to the flame.

The question does arise as to how we can reconcile

the above observations with the indisputable fact

that materials do burn faster upward than downward. If we remind ourselves that in a reverse flow

experiment, the flame itself is still forming above the

burning line, we must see that this is not the same

situation as in upward burning where the flaming

region actually impinges on the unburned portion.

The reverse flow experiments do indicate that convective heat transfer itself may not have a significant

influence on burning. This would be consistent with

other experimental results we have obtained showing

that the flame propagation rate for such materials

does not change as the flame movement direction

varies from horizontal to the candle-like, downward

burning mode. In other words, once the flame itself

cannot actually come in contact with the material

to be burned, heat transfer from it is not consequential.

The reverse flow experiments also are useful in

revealing the relative flame quenching capabilities

of any non-combustible byproducts. Of the examples

given, the most extreme case is the treated cotton,

which has a higher O.I. even at the increased downward flow rates. Nylon and polyester, on the other

hand, are not affected by gas flow changes.

F. J. Martin

The downward airflow over a specimen burning at the top end can result in complex gas flows in the

flame region where buoyancy becomes important.

We have not tried to analyze such flows; rather we

have sought to establish a reproducible, easilyobtained flow pattern and defined the oxygen index

to apply to it. One could define some other set of experimental

conditions as proposed by Miller, but there appears

to be little advantage.

Furthermore, we have sought to interpret our

indices in terms of a very simple model for the

burning process. Such a model cannot explain

169
details of the complex flame stabilization process, but

can be used to distinguish some interesting differ- ences in burning characteristics of different materials. While Mr. Miller's explanation of flame

quenching by "noncombustible byproducts" may

seem quite plausible, a much more detailed analysis

of the flame would be required to justify it.

C. F. CuUis (The City University, London):

The "oxygen index" method for the measurement

of polymer flammability, developed by Dr. Fenimore

and Dr. Martin, is proving to be of considerable

value. Apart from the determination of the readiness

with which flame is sustained at the surface of a burning polymer, however, it is important also to be

able to measure the ease of ignition of polymers.

Mr. A. W. Benbow has been attempting, at the City

University, to determine, by means of an electric arc method, the energy required for ignition of various

polymers. Due to the large excess voltage needed to

produce an arc, this method is not very satisfactory

and experiments are now therefore being carried out

which involve measuring the time taken to ignite

polymers by means of a hydrogen flame source placed at a given distance above the surface of

the polymer.

One of the associated advantages of the "oxygen

index" method is that it makes it possible to distinguish between some of the several mechanisms

whereby additives confer flame retardance on

polymers. Thus, certain additives decrease the rate of

evolution of flammable degradation products, while

others inhibit the subsequent combustion of these

products. Dr. Fenimore and Dr. Martin have shown

that these two broad modes of action can be distinguished by determining whether the inhibiting

effect is dependent on polymer structure on the one hand and on the nature of the gaseous oxidant on the

other. It is necessary to recognise, however, that

flame extinction may be brought about in several

ways. Thus the additive may cause the release into

the gas phase of a compound which reacts with

highly active free radicals, replacing them by less

reactive ones [1]. In some systems finely divided

solid particles may be given off which catalyse

chain-termination processes [2]. In other cases, however, the additive may simply yield a heavy inert

gas which "smothers" the flame by reducing the

supply of gaseous oxidant; this seems to be the

principal mode of action of certain halophosphorus

esters [3]. It seems important to be able to devise

experimental criteria by which these different modes

of action of flame-inhibiting additives may be

distinguished.

There are, of course, as Dr. Fenimore and Dr.

Martin point out, several other ways in which compounds can exert a flame-retardant action. Thus

certain additives (inert fillers) simply act as a heat

sink, while others facilitate melting of the polymer

[4] or cause the formation of a protective layer of

solid which tends to prevent the oxidant from

reaching the polymer surface.

References

[1] Volans, P., Trans. J. Plastics Inst., 1967. (Suppl. Conf.

Flame Resistance in Polymers, 1966).

[2] Thiery, P., Plastiques modernes et elastomeres 20, 143

(1968).

[3] Benbow, A. W., and Cullis, C. F., Conference on Flammability of Polymers, Prague, 1970.

[4] Eichorn, J., J. Appl. Polymer Sci. 8, 2497 (1964).

170
NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

Pyrolysis and Combustion of Cellulosic Materials

Raymond Alger

Naval Ordnance Laboratory, Silver Spring, Md. 20910

Despite centuries of experience with burning cellulosic materials, men still do not agree on

the interpretation of their observations. These divergencies of opinion generally involve mechanisms at the microscopic level where areas of ignorance still abound. This review deals with some

of these opinions and the supporting experimental evidence pertinent to three chronologically

related areas of fire research; namely, pyrolysis, ignition, and fire spread. Some of the ignition

concepts and the models for fire spread are not limited to cellulose but should apply to a variety

of solids.

Key words: Burning of polymers; cellulose; combustion of polymers; decomposition; fire spread;

ignition.

1. Introduction

In the biblical account of a cellulose ignition

experiment conducted by Elijah and the prophets

of Baal, the outcome was unique in the annals of

fire research.^ First the results conclusively established the source of ignition beyond all doubt and

second, the dissenting voices were silenced forever.

In contrast, present day fire research generally

approaches the resolution of divergent viewpoints

and theories through a slow amassing of evidence,

not by design but for lack of a quick definitive

experiment. Usually, the complexity of the problem

permits the results to be rationalized within the

bounds of several competing fire models and theories.

This paper is concerned with three such problem

areas in cellulose fires. The order of presentation

follows the chronological development of a fire commencing with pyrolysis and progressing through

ignition and fire spread. The scope is limited to small

areas where recent results are stimulating new concepts or are providing support for one of the competing models. Our interest is to stimulate and

encourage further work aimed at answering a few

questions rather than to attempt a general review of

cellulose pyrolysis and combustion.

In order to clarify the points of contention and

areas of ignorance, it is profitable to outline some of

the physical and chemical properties of the fuel and

to describe the pertinent characteristics of the fire. The first two sections contain this background

information, while the last three parts deal with the

new results in support of the various theories.

2. Chemical and Physical Properties of

the Fuel

The "ideal" cellulose is composed of anhydroglucose or beta-glucopyranose structural units

1 1 KingB, 18.

(CeHioOs) that join together through the 1,4

glucosidic linkages as shown in figure 1 to form a

linear molecule. Any number of units odd or even and ranging up to 5,000 or more can unite to form

the cellulose chain. The number or degree of polymerization (D.P.) provides a basis for classifying

cellulose. Alpha-cellulose consists of the longest

chains, i.e., 96 percent of a good grade of alphacellulose will have a D.P. in excess of 1,000 based on viscosity measurements (Ott and Spurlin-1954)

The boundary between alpha- and beta-cellulose

corresponds to a D.P. of about 200 and gammacellulose has the shortest molecules, i.e., D.P. up to

about 10. Although these categories were initially

based on solubilities in a caustic soda solution^ the

boundaries indicate changes in other properties as

well. For example, mechanical strength is associated

with long chains and this strength is lost when the

D.P. is as low as found in the beta- and gammacellulose.

The starting material in the experiments described

here is alpha-cellulose. Since some of the molecular

bonds will be broken during pyrolysis, the weakest

links are important in determining the course of

events. The C-0 bonds are weaker than the C-C

bonds, therefore the principal candidates for scission are the A, B, C, and D linkages in figure 1,

(Shafizadeh-1969) (Madorsky-1964) . Rupture of

the bonds in the ring structure (A, B) leads to

decomposition of the molecule to yield water,

carbon dioxide, carbon monoxide, and char, whereas

cleavage of the C or D bonds leads to depolymerization and the production of levoglucosan.

In the next stage of aggregation, cellulose molecules

group themselves together to form fibers whose

detailed structure remains a point of debate. X-ray

diffraction patterns of these fibers indicate ordered

2 Beta- and gamma-cellulose are soluble in 17.5 percent NaOH solution but alpha-cellulose is not.

171
CHaOH

NONREDUCING

END GROUP

REDUCING END GROUP

Figure 1. Cellolose chain structure and the anhydroglucose unit (Ott and Spurlin—1954)

regions typical of a crystalline structure along with

disordered regions which are considered to be

amorphous. Presumably the forces that hold the

molecules together, i.e., cross linking hydrogen bonds

and van der Walls attraction are most effective in

the crystalline regions. Several types of evidence

have been used to indicate the relative magnitude of

the crystalline and the amorphous regions, for

example:

1. Chemical reactions involving a selective reagent

that cannot penetrate the crystal regions of

the fiber.

2. Kinetic methods where the observed parameter

exhibits an initial rapid rate followed by a slower final rate characteristics of the amorphous

and crystalline regions respectively.

3. Physical methods such as the relative intensities

of the amorphous and crystalline portions of the

X-ray diffraction pattern (Ott and Spurlin—

1954) , i.e., the crystallinity index CrI defined as

1002 -lam

1002 ^^^^

where 1002 equals the maximum intensity of

the lattice defraction in arbitrary units and

lam is the intensity of the amorphous background defraction in the same units (Parker

and Lipska-1969).

Both methods 2 and 3 play an important role in

the supporting evidence for the pyrolysis models to

be discussed later. While the absolute degree of

crystallinity values obtained in the various techniques cannot be readily compared, the relative

values are in general agreement and provide a useful

means of checking the microscopic behavior of the

fibers. Typical values of the percent crystallinity

ranged from 60 to 90 depending on the source of the

cellulose and the measuring technique (Ott and

Spurlin-1954)

Crystal sizes and shapes have been studied with

the electron microscope and by viscosity measure- ments of dissolved samples. Typical lengths range

from 100 to 300 D.P. depending on the source of

cellulose and the treatment. If the length of an

anhydroglucose unit is 5.15 angstroms, these

crystals would range from about 500 to 1500 angstroms in length and perhaps 40 to 100 angstroms in

width (Ott and Spurlin-1954). Since most of the

individual alpha-cellulose molecules are at least

5,000 angstroms in length, i.e., D.P. of over 1,000,

they can extend through a number of amorphous and

crystalline regions. Numerous authors, for example,

(Ott and Spurlin-1954, Rebenfeld-1965, Warwicker

et al., Manley-1965, and Parker and Lipska-1969)

have discussed the various models for describing the

crystalline-amorphous structure in cellulose fibers,

e.g., micellar theory, the fringed micellar structure,

the fringed fibrillar structure, and the folded chain.

Fortunately, variations in these models are not

crucial to the present state of the pyrolysis models.

All of the fiber structures permit bond rupture and

molecular rearrangement to occur more readily in

the amorphous regions.

3. Stages in Fire Development

3.1. Pyrolysis

When cellulose is pyrolyzed, the principal points

of interest are:

1. The identity of the pyrolysis products

2. Their rate of production

3. The dependence of species and yield on

heating rate and temperature, and

4. The controlling reactions.

Considerable information and agreement exist

with regard to the first three points, but diverging

concepts about the mechanisms are one of the

principal concerns of this paper. During pyrolysis,

cellulose undergoes some very obvious and well

known physical changes, e.g., the color changes from

white to brown then black, the size and weight are reduced, and the flexibility and mechanical strength

are lost. The time and temperature dependence of

the weight loss during pyrolysis has been studied

extensively and the results constitute one of the

principal bits of information regarding the mechanism. Figure 2 shows typical isothermal weight loss curves for two temperatures that bracket the region

of immediate concern. Conjecture about the order of

the reaction during pyrolysis, i.e., first order, zero

order, etc., is based largely on the shape of such

curves and the associated measurements of the

remaining glucosan. We will return to this point.

If pyrolysis is continued at the temperatures of

figure 2, the weight loss eventually reaches about

80 percent and the remaining cellulose has been

converted to a char. Very prolonged heating or exposure to higher temperatures, e.g., 600 °C will

reduce the char fraction to about 9 percent of the

172
Crystotlinity at aSS'C.

.„„ _ (Porker a I.(p«»o, 1969)

O 90 —

-J \ Cryatolllnlly at Z60^C. t .„ _ -0<^

/ (Wainilem a BraiiSa, 1970)

.0001 .001 .01 .1 1.0 to too TIME (Mrs)

Figure 2. Degree of crystallinity and weight loss as a function

of pyrolysis time for several different temperatures.

original cellulose weight. If the crystallinity is measured throughout the course of pyrolysis, the

index will follow the second set of curves in figure 2. The index may be increased slightly during the early

stages, but for the most part there is very little change in crystallinity until the weight loss reaches

about 50 percent. When the crystallinity shows an

increase, it is because lam for the amorphous region

has been reduced and not because of an increase in

the reflections from the crystalline portions of the

fiber.

The chemistry of pyrolysis is sufficiently complex

that the available information consists primarily

of a knowledge about the starting material, the stable

end products, and how some of these products depend

on the thermal treatment. For convenience, the final

products have been divided into three categories:

1. Volatile products of molecular weight below

about 105. These products escape readily during

pyrolysis and at last 59 different compounds have

been observed, of which 37 have been identified.

Typical examples are carbon monoxide, carbon

dioxide, water, acetol, furfural, and a variety of

unsaturated aldehydes and ketons.

2. Tar constituents of higher molecular weight.

Levoglucosan is one of the principal components, and

3. Chars.

The intermediate char corresponding to the 80

percent weight loss has an approximate formula

C3H4O2 which is only slightly carbon enriched over

the original material. Further pyrolysis to the final

9 percent char involves the elimination of water to

give a composition of about 90 percent carbon,

2.7 percent oxygen, and 5.4 percent hydrogen, i.e., a 48 percent increase in the carbon fraction (Parker

I and Lipska-1969) . In a further effort at classification

the degradation reactions responsible for product

categories 1 and 3 have been termed dehydration or char forming reactions, while the tars are generated

in depolymerization processes. Obviously considerable bond rupture and rearrangement are required to

achieve the wide variety of products observed. The

distribution of products depends both on the temperature and the temperal position in the course of

pyrolysis. At the lower temperatures, dehydration

reactions predominate (Martin-1965) while depolymerization with the production of levoglucosan

becomes the dominant mode at higher temperatures,

e.g., at 360 °C. The ratio of tar/char formation has

reached 60/40. Between 275 and 400 °C the pyrolysis

products formed appear to be independent of

temperature. Under isothermal conditions, all the

products are emitted throughout the course of

pyrolysis, although not necessarily at a constant

rate or ratio (Lipska and Wodley-1969)

Realizing that a reasonably complete understanding of the many reactions and interactions

involved in cellulose pyrolysis appears to be a long

way off, our principal question is whether a simplified

model can account for the general features of

pyrolysis, particularly the relative production of

volatiles, tars, and chars. Aside from the fundamental interest, the factors controlling the reaction

path are of considerable practical interest to the

fire community, because most fire retardants for

cellulose enhance char production. If we are to

intervene either to enhance or suppress combustion,

control of the pyrolysis path appears essential.

3.2. Ignition

Ignition in cellulose is similar to other ignitions in

the sense that combustion occurs when a suitable

combination of fuel, oxygen, and heat are available.

Since the fuel concentration is not normally known

for pyrolyzing cellulose, ignition conditions have

generally been measured in terms of

1. The critical irradiance, i.e., the threshold value

of heat input for ignition.

2. The time to ignition for a constant level of

heat input, and

3. The ignition temperature which is really the

temperature of the surface at ignition.

Both the fuel concentration and the temperature

depend on the heating rate, consequently, ignition can occur for various combination of heating rate,

temperature, and time. Martin (1965) classified

cellulosic ignitions, produced by radiation, into three

categories on the basis of the heating rate.

1. Low heating rates where the sample approaches

a uniform temperature and the heat losses depend on

the local air motion and turbulence,

2. Diffusion controlled ignition which corresponds

to intermediate heating rates of 5 to 20 cal cm~^ s~^

The temperature profile in the cellulose is determined

by the diffusion of heat into the solid, and

3. High heat inputs where the ignition behavior is controlled by the ablation of the irradiated surface.

Furthermore the type of ignition, i.e., transient

flaming, sustained flaming, or glowing depends on the

fuel supply and therefore the heating rates. The results discussed here will be limited to the

diffusion controlled case where the production of

173
combustible products is adequate to insure flaming

ignition. Two observations (Butler et al.-1956 and

Martin [8]-1965) have contributed substantially

to our ability to rationalize and predict ignition

behavior from the knowledge of the heating rate and

the thermal and optical properties of the cellulose

sample. First, the time to ignition ti was found to be

inversely proportional to the square of the heating

rate, i.e., ti <^ where H is the irradiance. Second,

a solution to the heat conduction problem for radiant

energy entering a semi-infmite slab gave the same

inverse relationship provided ignition occurred at a constant surface temperature (T), i.e.,

tioz {AT/Hyirkpc.

Where Ape are the thermal conductivity, density,

and specific heat respectively for the slab. Subsequent

measurements over the intermediate heating range

from 4.4 to 11.6 cal cm~- s"^ gave essentially a constant surface temperature of 600 to 650 °C at the

time of ignition, (Martin-1965) and (Alvares-1967)

With this correlation and despite the lack of a

specific ignition mechanism, considerable success has

been achieved in predicting ignition conditions for p variety of cellulosic materials.

Efforts to determine the ignition mechanism have

concentrated on the pyrolysis products in a search

for a species or production rate that could be

correlated with the constant ignition temperature.

Martin (1965) concluded that a threshold in either

the rate or the total yield of the combined combustible products was untenable because at the

ignition time both the rate and the amount varied

depending on the irradiance, i.e., the rate was

proportional to H and the amount to 1/H. A possibility still under consideration involves an active

species that becomes available at the ignition

temperature. This concept will be considered subsequently in connection with the recent ignition

data. Meanwhile, the principal questions remain:

What is the identity of the products responsible for

ignition? And what is the controlling reaction?

3.3. Flame Spread

Each successive stage in fire development becomes

more complicated to analyze because all of the

proceeding stages are involved. In a 1968 review

Friedman [12] noted that "Combustion science is not able to specify mathematically the laws governing

this process (flame spread) and indeed there is uncertainty as to which variables exert the dominant

affects." In addition to a continuing process of

pyrolysis and ignition, flame spread also involves

heat transfer and aerodynamics. At this point the

geometry, size, and orientation of fuel becomes an

important factor contributing to the pattern and

rate of flame spread. Two examples that illustrate

this point are (1) linear flame spread on a card, and

(2) exponential fire growth in a box. If a vertical

3X5 card is ignited at the top, flames will spread

downward at a uniform velocity of about 3| in. per

minute (Parker [13]-1969). Other samples ranging

from sheets of cotton upholstery to wooden barns

300 ft long have also exhibited constant flame spread

rates. All of these cases are geometrically similar,

i.e., the fuel is long compared to the cross sectional

dimensions. Consequently, only the fuel close to the

flame receives heat from the combustion zone. Because of the light fuel loading, the fire quickly

consumes the fuel and reaches a fixed size that

moves as a narrow band of flames from one end of

the fuel element to the other. Most of the combustion

heat escapes without contributing to the pyrolysis

of unburned fuel.

In contrast, the fire inside a box warms all the

inside surfaces. Each segment of surface sees the

entire burning volume and the fire spreads in

proportion to the burning volume. Again, this

behavior has been observed both in small volumes

and in full sized houses. Obviously, these characteristics are not limited to cellulose fires. In fact,

the linear flame spread behavior has been observed

in liquid fires of the proper geometry. However, the

details of heat transfer and volatile fuel production

are a characteristic unique to each material. For

example, the pyrolysis products can influence the

heat transfer to unburned fuel by irradiation. In

the card experiment, cellulose burns with an almost

completely nonluminous flame so that little heat is radiated. The principal objective in flame spread

studies has been to simplify the problem by eliminating as many parameters as possible through the

choice of sample geometry and orientation so that a

solution can be obtained for the heat transfer and

aerodynamic affects.

4. Recent Developments in Proposed

Mechanisms for the Thermal

Decomposition of Cellulose

In conjunction with research on fire retardants,

several models have been developed to explain the

available cellulose pyrolysis data. These are simple

models that seek to determine the principal events

that control the production of combustible volatiles

and chars; however, no attempt is made to treat the

detailed chemical reactions. First we will outline the

models and the supporting experimental evidence

then examine the points of similarity and divergence.

The first model (Parker and Lipska-1969) is for the

temperature range between 275 and 360 °C. This

limitation is imposed because below 275 °C there

appears to be a significant change in the predominant

decomposition process, i.e., the rates of dehydration

and cross linking becomes competitive with depolymerization. At the higher temperatures, the

proposed mechanism is as follows:

1. When pyrolysis commences, thermal vibrations cause the long alpha-cellulose molecules to

break at weak links, thereby forming short

molecules about 300 to 500 D.P. long. Presumably, the scissions occur primarily in the

174
r

o

k.

[J'

(400

1200

1000

800

600 V

Z75'C (Hatpern a Patai, (969J

SIS'C. (L/psko, 1970)

I < TIME (Hrs.l

Figure 3. Degree of polymerization as a function of pyrolysis

time.

amorphous regions of the fiber; consequently,

the short molecules are mostly in ordered

crystalline segments. Evidence of this rapid

reduction in size comes from measurements of

the degree of polymerization as a function of

pyrolysis time. Figure 3 shows typical degree of

polymerization characteristics for two pyrolysis

temperatures, one above and one at the 275 °C

boundary.

Next, the short molecules begin to unzip and

produce severed monomer units. All of the

molecules were assumed to depolymerize simultaneously by starting at one of the ruptured

ends and unzipping progressively through the

molecule. This concept of simultaneous unzipping is based on two bits of evidence. First at

288 °C the remaining glucosan units in figure 4

decrease linearly with time after the initial

rapid adjustment. Second, the weight loss curve exhibits a linear time dependence up to

at least 50 percent of the weight loss, i.e., the

reaction rate is of zero order. In figure 4 there

is an initial 5 to 10 percent decrease in weight

which occurs too rapidly to plot on the time

scale of this figure. Presumably, the rapid

unzipping of molecule ends in the amorphous

region is responsible for this sudden change.

The crystalline regions are assumed to be more

stable and therefore responsible for the subsequent weight loss exhibited in the linear part

of the curve. The degradation products formed in this unzipping process can go either into char forming

reactions where dehydration liberates CO,

CO2, H2O, etc., and leaves a residual char or they can undergo a rearrangement to form

levoglucosan. While the reaction details are not

essential to the development of the model, the

yields of levoglucosan and char, i.e., about

60 percent and 40 percent respectively do

provide evidence for the two competing

processes.

4. When the cellulose is heated with fire retardants

two effects are observed. The rate of degradation is enhanced and the yield of char is increased. In the model, the retardant acts in two

ways. First, it increases the number of short

molecules, thereby increasing the overall rate

of unzipping, and, second, it interacts with

levoglucosan to generate a furanose derivative

which forms char in preference to combustible

volatiles. Evidence for these conjectures hinges on measurements of levoglucosan and furanose

derivative yields as functions of the fire retardant treatments.

The second model (Weinstein and Broido-1970)

is based on isothermal data to 260 °C.

1. Initially the behavior is similar to the first model except that the rapid phase of decomposition becomes constant at a D.P. of about

160 as shown in figure 3. During this constant

period the crystallinity index in figure 2 also

remains nearly constant.

2. This model is based on competing dehydration

and depolymerization reactions, where the

results depend on the heating rate and the

sample purity. Dehydration is favored by cross

links between neighboring units in the crystalline regions. Consequently, the amount of char

produced depends on the original amount of

order and disorder in the sample. Depolymerization is favored by high temperatures and

once started, a molecule is assumed to imzip

completely in a time short compared to the

pyrolysis period. Presumably the number of

molecules unzipping at any one time is proportional to the number present and the reaction

is assumed to be first order. As the pyrolysis

continues the number of cellulose molecules

decreases, but the D.P. and crystallinity

remain unchanged in agreement with evidence

in figures 2 and 3, for the low temperature

samples. Other evidence comes from weight

100

90

80

70

\-^ 60

UjO 50

CtLu 40

CL 30

\ \^ ^^Residual weight of sample

\\\ \^ ^Remaining glucosan units \

Figure 4. Pyrolysis of alpha-cellulose in nitrogen at 288 °C

{Parker and Lipska—1969).

175
loss curves where the rates indicate a first order

reaction.

3. Fire retardants are assumed to catalyze

reactions in the initial stage of bond rupture

where they can affect the active ends of

molecules to enhance cross linking and thus the

char forming reactions.

A slight variation on this model is found in

(Halpern and Patai-1969), where the general

assumptions are the same regarding the rate of

reaction; however, the initial bond ruptures give

some crystalline segments about 160 D.P. long and

some shorter amorphous segments. Dehydration is assumed to predominate in the amorphous segments

leading to char formation, while the crystalline

segments favor depolymerization.

A comparison of these models reveals the following

points of divergence.

1. The differences commence with the initial

rapid segmentation of the fibers. In the first model, the molecules continue to decrease in

length, while in the second case, the D.P.

remains fixed throughout much of the reaction.

2. The pyrolysis rates lead to different orders for

the main portion of the unzipping reaction,

i.e., zero order and first order respectively.

3. While unzipping becomes a dominant mechanism at high temperatures, different modes are invoked in the two cases, namely, slow unzipping of all molecules simultaneously versus rapid and complete unzipping of a few reacting

molecules.

4. The choice between char and tar forming

reactions enters at different stages of degradation and therefore is controlled by different

parameters.

5. Finally different modes of intervention by the

fire retardant materials are required to meet the

difference in the char and tar forming reactions.

In a way, such conflicts are to be expected because

each model is trying to explain a different set of

experimental data. Furthermore, the experimental

conditions, i.e., sample material, temperature, and

measuring procedures were sufficiently varied to

prevent reliable extrapolations from one to the

other. The temperature differences and heating

techniques are probably the most influential parameters. For example the temperature effects are very

noticeable in the weight loss and glucose loss data.

The Parker and Lipska model contains the assumption that these loss curves exhibit three regions of

behavior as shown in figure 4, i.e., the initial rapid

loss, the linear region corresponding to the simultaneous unzipping of all the molecules, and the final

slow approach to the intermediate char. At high

temperatures the linear region includes the majority

of the weight loss, but at lower temperatures it is not as pronounced. Since the whole curve does not

follow zero order kinetics, some attention has been

given to the relative degree of fit to the data that can be obtained for the zero order and a first order

curve. In some preliminary efforts to fit the data for

temperatures above 275 °F, a computer was programmed to perform two tests, (a) to find the

pyrolysis time interval over which zero order and

first order curves would fit the data within the same

allowed deviation, and (b) to find an exponent for

the curve giving the best fit. In the first case, the

zero order gave the longest time interval and in the

second the exponent was a fraction much closer to zero than unity.

Obviously a first step in resolving the conflict over

pyrolysis models involves obtaining compatible data

so that all theories are based on the same experimental observations. If the data does not converge

when common materials and experimental parameters are employed, it may be necessary to return to

Mount Carmel for another decision.

5. Some Recent Ignition Results

For intermediate heating rates the constant

ignition temperature gave rise to speculation

favoring a threshold for a reactive pyrolysis species as the controlling factor in cellulose ignition. In a series

of experiments, Alvares (1967) examined this

possibility along with the classical piloted, thermal,

and free radical self-ignitions. The fuel and external

heating rate were held constant so that the ignition

time could be measured as a function of the en- vironment. Both the composition of the atmosphere

and the pressure were varied. Presumably, the

variations in oxygen concentration at constant

pressure did not alfect the pyrolysis process prior to

ignition. Consequently, the dependence of ignition

temperature on the oxygen concentration shown in

figure 5 was interpreted as evidence against the

reactive product threshold theory. While the results were compatible with the thermal self-ignition model,

the qualitative nature of this study did not rigorously

exclude the piloted or free radical chain branching

mechanisms.

In a recent re-examination of this data, the results were compared to values computed for a simplified

and idealized thermal auto -ignition (Alvares and

1100

<3I000

o

I

I

!

Uj

900

800

700

6O0

500

\ ^^Anolytical solution (VARYING Ot CONCENTRATION

1^ » -^Experimental points yAT I ATMOSPHERE

nalytical solution (VARYING O2 CONCENTRATION

\), y,Experimental points\BY VARYING TOTAL PRESSURE

too 200 300 400 500 600 700

OXYGEN PARTIAL PRESSURE, mm. Hg.

Figure 5. Surface temperature at time of ignition as a function of oxygen partial pressure.

176
Martin-1970) . Ignition times were computed as functions of pressure, oxygen concentration, and

diluent gas composition. In this model, ignition occurs when the heat release rate just exceeds the

rate of heat loss. Simplified expressions were developed for the pertinent parameters, e.g., heat loss,

reaction rate, and surface temperature. A digital

computer provided calculations of ignition times

versus total pressure and oxygen partial pressure.

Typical results in the form of a calculated curve for

ignition temperature versus oxygen pressure are shown in figure 5 along with experimental data

points. At the extremes of high and low oxygen

concentration, some additional factors become

important that were neglected in the simplified

model. The high temperatures associated with low

oxygen concentrations are not observed because the

surface ablates away before such temperatures can be reached. Also the model makes no correction for

radiative heat losses at high temperatures. For

intermediate oxygen concentrations the calculated

values are quite compatible with the experimental

points. At high oxygen concentration, the model

predicts ignition temperatures below the observed

values. Presumably, the minimum ignition temperature corresponds to the onset of sufficient pyrolysis

to form a combustible mixture. Fuel production is assumed to be adequate in the simplified model;

however, a logical extension would include a term

for the kinetics of fuel production. In view of the

favorable results obtained with this model, a more

sophisticated description of the thermal self-ignition

process would appear to be in order.

6. A Simplified Flame Spread Model

The experiments to be described here are submitted as an example of a flame spread geometry

that appears to have sufficient simplicity to make it worth a more rigorous treatment. The desired end

product should be capable of predicting flame spread

velocity from the thermal, chemical, and physical

properties of the fuel and atmosphere without

resorting to empirical parameters. Downward flame

spread along a long thin sheet or card of cellulose fits the criterion for a simple geometry that minimizes

the number of parameters entering the balance of

energy and fuel transport. The flame spread mechanism proposed by Parker (1969) contains two very important observations. First the pyrolysis gases are released behind the flame front. This is in contrast

to the belief that in cellulosic materials the pyrolysis

gases are distilled from the surface ahead of the flame

front. Second, the heat transfer to the surface is principally by gas phase heat conduction.

These mechanisms are the result of the geometry

at the leading edge of the flame. The flame stands

away from the surface and extends downward beyond

the visibly darkened zone where the pyrolysis

products are being generated. The high temperatures

in the flame and the close dimensions favor heat

conduction in preference to convection and the lack

of a luminous flame minimizes radiation as a source

of heat transfer to the unpyrolized material under the

flame overhang. In Parker's paper the flame velocity

was calculated from physical and thermal characteristics of cellulose, the pyrolysis rate as a function of temperature, and measured values of the

overhang and standoff distances. Once the parameters controlling the flame overhang and standoff

distance are understood in terms of the fuel and

environment these dimensional measurements should no longer be required.

This brief examination of some problems and

progress in pyrolysis, ignition, and flame spread

shows that there is promise of significant future developments as well as room for another Elijah with

his definitive experiments.

The author is deeply indebted to collegues Alvares,

Butler, Lipska, and Martin at Stanford Research

Institute and Parker now at NBS who provided the

information and inspiration for this paper. The

differing points of view from the other references are

also appreciated as an essential ingredient to the

theme.

7. References

[1] Alvares, N. J., Oxygen Concentration and Pressure Effects on the Ignition of Cellulosic Fuels by Thermal

Radiation, TR-67-25 U.S. Naval Radiological Defense

Lahoratory, San Francisco, California, 1967.

[2] Alvares, N. J. and Martin, S. B., Mechanisms of Ignition

of Thermally Radiated Cellulose, Combustion Institute

Meeting at Salt Lake City, 1970.

[3] Butler, C. P., Martin, S. B., and Lai, W., Thermal

Radiation Damage to Cellulosic Materials, Part II, Ignition of Alpha Cellulose by Square Wave Exposure,

TR-135, U.S. Naval Radiological Defense Laboratory,

San Francisco, California, Nov. 1956.

[4] Friedman, R., Fire Research Abstracts and Reviews, 10,

1 (1968).

[5] Halpern, Y. and Patai, S.. Israel J. of Chem., 7, 673 (1969).

[6] Lipska, A. E., The Effects of Flame Retardants on Thermal Degradation of Alpha Cellulose in Nitrogen, Stanford Research Institute, Menlo Park, California, to be

published, 1970.

[7] Lipska, A. E. and Wodley, F. A., J. Applied Polymer Sci., 13, 851 (1969).

[8] Madorsky, S. L., Thermal Degradation of Organic Polymers, Polymer Reviews, Interscience Publishers, Editors H. F. Mark and E. H. Immergut 7:238, 1964.

[9] Manley, R. S. J., Nature 204, 1155, (1964).

[10] Martin, S., 10th Symposium (International) on Combustion, p. 877, 1965.

[11] Ott, E. and Spurlin, H. M., Editors, Cellulose and Cellulose Derivatives (Interscience Publishers, New York,

New York, 1954).

[12] Parker, W. J., Flame Spread Model for Cellulosic Materials, Central States Section of the Combustion Insti- ture. University of Minnesota, March 1969.

[13] Parker, W. J. and Lipska, A. E., A Proposed Model for the Decomposition of Cellulose and the Effect of Flame

Retardants, TR-69, U. S. Naval Radiological Defense

Laboratory, San Francisco, California, 1969.

[]4] Rebenfeld, L., Morphological Foundations of Fibers

Properties, Structure and Property of Polymers, Interscience Publishers, Edited by A. V. Tobolsky, 1965.

177
[15] Shafizadeh, F., Pyrolysis and Combustion of Cellulosic

Materials, Recent Advances in Carbohydrate Chemistry 23, 419 (1969).

[16] Warwicker, J. O., Jeffries, R., Colbran, R. H., and Robinson, R. B., A Review of the Literature on the Effects

of Caustic Soda and Other Swelling Agents on the Fine

Structure of Cotton, Shirley Institute Pamphlet No.

93, p. 38, Shirley Institute.

[17] Weinstein, M., and Broido, A., Combustion Sciences and

Technology 1, 287 (1970).

Discussion

J. H. IMagill (University of Pittsburgh)

The role of basic and acidic flame retardants leads

me to inquire about the efficacy of (1) a chemically

bound inhibitor (or retardant) versus (2) the same

chemical element or group present as an integral

part of the polymer.

IVTention was also made (in your paper) about the

retardant effect produced through radiation cross- linking cellulose, but nowhere (it seems) has much

attention been given to the crystallinity of many of

the polymers tested in the Symposium. One would

hardly expect a low and high crystallinity polymer

to behave similarly if there was an appreciable

difference in the enthalpy of fusion. I wish to inquire

of Drs. Alger and IVIcAlevy III, if the crystalline

state of the polymer is important here. Furthermore,

based on the models of Jellinek, figure 10 (p. 106 of

preprints), the susceptibility of the structure to

attack may be explained on the accessibility of

oxygen to the polymer. The nature of the morphology

or polymer architecture can be modified experimentally to test if this model is the most relevant.

R. Alger:

In answer to Dr. IMagill's questions, we have no

data on the retardant efficiency of an inhibitor

present on cellulose compared to the same chemical

element incorporated in the molecule. Lipska has

measured the effect of two halogens on the production of char during the pyrolysis of cellulose. Comparable weights of potassium bromide added to

cellulose and iodine substituted in the cellulose

molecule produce comparable increases in the yield

of char. The second question was concerned with the

importance of the crystalline state of the polymer.

According to the pyrolysis models for cellulose, the

initial bond rupture occurs principally in the

amorphous regions. Consequently the degree of

crystallinity would influence the initial rapid drop in

weight loss, i.e., this initial weight loss should be less

for samples of high crystallinity.

D. C. Aldrich (IVIassachusetts Institute of Technology) : You indicated that the pyrolysis mechanisms for

cellulose proposed by various investigators are inconsistent because they are based upon data

obtained with different sample material, temperature, measuring procedures and techniques. Would

you not also include the differences in sample

geometry and cracking behavior of the char structure

since these effect secondary pyrolysis and heat con- vection by the escaping volatiles? You suggested

that these discrepancies could be eliminated by

choosing a common experimental technique. What

technique do you suggest? If an isothermal technique were so chosen, would it be possible to apply

the data obtained in this manner to large scale

applications?

R. Alger:

In answer to the questions raised by Donald

Aldrich, (1) sample geometry, particularly the

sample thickness, is an important parameter tn attempts to obtain experimental reproducibility,

(2) the feasibility of using isothermal pyrolysis data

for large scale applications will depend on the

particular applications; e.g., the isothermal data can

be used in applications involving fire retardant

treatments and for the identification of pyrolysis

products; however, quantitative yields of products

will depend on the particular time temperature

schedule.

R. F, IMeAlevy, III (Stevens Institute of Technology) :

The author has ignored much of the standard

combustion literature in his survey. Some of the

problems that are cited as still outstanding in the

case of cellulose have, in fact, been solved for any

fuel bed material.

IVIuch of the work was done under the aegis of

aerospace research funds, so the examples used to

illustrate solutions, and titles of papers, deal with

solid propellants.

The mechanism of solid fuel ignition in an oxidant

containing environment (air), was understood 10

years ago. (IMcAlevy, R.F., III, Cowan, P.L., and

Summerfield, M.., The ]VIechanism of Ignition of

Composite Solid Propellants by Hot Gases. Progress

in Astronautics and Rocketry: Solid Propellant

Rocket Research, Vol. 1, p. 652, Academic Press,

1960.) An excellent review of solid fuel ignition

theory was presented a few years ago, "Theory of

ignition of solid propellants," Price, E. W., Bradley,

H. H., Jr., Dehority, G. L. and Ibiricu, M. M.,

AIAA J. 4, 1154-1181 (1966).

The mechanism of flame spreading over a solid fuel

surface has recently been revealed. Flame Spread

Over Fuel Beds: Solid Phase Energy Considerations,

Lastrina, F., ]VIagee, R. S., IMcAlevy, R. F., Ill,

presented at the Thirteenth Symposium (international) on Combustion, Univ. of Utah, Salt

Lake City, Utah, Aug. 24-28, 1970; to be published

in 1971.

178
R. Alger:

Professor McAlevy's point regarding the extensive

ignition literature is well made; however, I cannot

share his conclusion that the ignition problem is solved for cellulose, at least in sufficient detail to

meet current requirements. Much of the rocket

ignition work dealt with the problem of comparing

gas-phase models to the condensed-phase ignition

theory that was prevalent at that time. In the case of

cellulose, the competing processes all involve gasphase reactions. The question is, which species

starts the reaction, i.e., radicals, one of the 59

stable pyrolysis products, or catalyst. When the

Semenov thermal autoignition theory is applied to the problem, the boundary conditions imposed by

the experiment and the criteria for ignition have a strong impact on the solution and the degree of detail

that can be accomodated in the expression for the

conservation of material and energy. The work of

Martin and Alvares was included here because the

boundary conditions include cellulose and a fairly

common procedure for applying external heat. In

other respects, they follow their predecessors by

resorting to simplifying assumptions that are quite

general. Consequently, the number of fitting parameters introduced to compensate for the dearth of

known physical and kinetic constants detracts from

the rigor of the results.

An example of such generalization is the criteria

for ignition. Martin and Alvares assume ignition occurs when the energy released in the elemental

volume exceeds the energy lost. McAlevy, et al., assume ignition occurs when a combustible mixture

of pyrolysis products and oxidant develops.

Obviously, both of these general conditions must be

met, but the details describing how are swallowed up

in the fitting parameters. Additional work will be

required before detailed boundary conditions such

as the production of the various specific pyrolysis

products can be incorporated in the autoignition

model.

With respect to the mechanism of flame spread over a solid surface, Parker, 1969, deRis 1969, and

Lasterina, et al., 1970, have converged upon similar

expressions for the spread velocity, at least for a thin

fuel element. All the models use gas-phase heat

conduction as the principal heat transfer mechanism;

however, there are some basic differences in the

physical models employed. Parker begins with a detailed study of the flame front and concludes that

the flame standoff distance is intimately related to

the heat transfer and that pyrolysis occurs behind

the flame front following a preheat zone that extends

up to the flame front. Although the deRis model

extends the flame to the surface, the pyrolysis zone

is behind the flame front in agreement with Parker's

findings.

R. Friedman (Factory Mutual Research Corporation): (communication)

In regard to the theoretical status of the flame

spread process, it seems valuable to call attention to

the theories of de Ris [1] for thermally thick and

thermally thin materials, which have been remarkably successful without recourse to adjustable

parameters. These theories give absolute prediction

of spread rate, including dependence on oxygen

partial pressure, total pressure, and sample thickness,

without requiring knowledge of the chemical rates. The gas phase reactions are assumed to be infinitely

fast, once fuel and oxygen have diffusively mixed,

and they occur all the way to the surface (i.e., no

flame stand-off distance is assumed). Forward heat

transfer from the flame through the gas phase to the

solid is crucially important for both thin and thick

fuels. It was possible to solve the associated mathematical models, provided one assumes constant

density and transport properties and a uniform

velocity profile.

The thin -fuel solution for the spread velocity, V, is

/Tflame~ T'vap

c T± va

where X is gas thermal conductivity, r is fuel halfthickness, and subscript w stands for the solid-phase.

Correspondingly, for a thermally thick fuel.

pCpK jTflame~Tvap

where Va^l\g{Tn^^.-T^) /pC^T^J'^ is the effective gravitationally -induced gas-phase velocity evaluated one characteristic gas-phase length 2X/ (pCpVa)

above the fuel bed. X„j, is the fuel bed thermal conductivity normal to the surface.

The thin-fuel formula predicts a spread rate

independent of both the opposing gas velocity, Va

and pressure; however the spread rate does increase

with the adiabatic stoichiometric flame temperature

Tfiame- The thick fuel bed result is quite different.

In this case the spread velocity is proportional to

Va, SO that in a quiescient environment with gravity,

the spread rate is proportional to the two-thirds

power of pressure.

The data taken by Royal and reported by Lastrina

et al. [2 J are correlated in the accompanying figure.

The coordinates are chosen so that both the thin and

thick fuel data should correlate if the theories are

valid. The ordinate is proportional to the spread

velocity while the abscissa is proportional to pressure.

The specific choices of property values do not

influence the shape of these curves; they influence

only the relative positions.

One sees that the correlation is very good for

oxygen mole fractions X02 greater than 21 percent.

The data for 21 percent O2 have the correct shape,

but are shifted downward and to the right.

These data lend powerful support to the two

theories. This agreement suggests that gas-phase

reaction kinetics is unimportant in the range of

experimental conditions. One does anticipate, however, that as the pressure or oxygen concentration

decreases, an extinction point will be reached, be179
T

3/2 r

_Tsi

3/2 (Tflame'Tpo)

/3Cp Too

./a

cause of reaction kinetics. This extinction point

should be sensitive to small concentrations of

gas-phase inhibitors.

References

[1] de Ris, J. N,, Spread of a Laminar Diffusion Flame, Twelfth

Symposium (International) on Combustion, p. 241, The

Combustion Institute, 1969.

[2] Lastrina, F. A., Magee, R. S., and McAlevy, R. F., Flame

Spread Over Fuel Beds: Solid Phase Energy Considerations, Thirteenth Symposium (International) on Combustion, 1970. To be published by the Combustion Insti- tute, Pittsburgh, Pa.

Robert H. Barker (Clemson University) (communication) : Mr. Alger has correctly pointed out the importance

of the factors controlling cellulose pyrolysis pathways

and their relevancy to the problem of fire retardance.

Unfortunately the models which he has presented to

describe the pyrolysis mechanism are probably not

the best description of these factors. Although there

are several points where these models do not seem to

provide an adequate explanation for experimental

observations, the most desturbing to me is the

mechanism proposed for levoglucosan formation.

This is a particularly important facet of any cellulose

pyrolysis model because of its direct bearing on the

nature and rate of the chain unzipping reaction and

consequently on the rate of formation on the

flammable volatiles which serve as fuel for flaming.

It is probably at this point that many flame re- tardants exert their influence.

Mr. Alger expressly states in the first model and

implies in the second that the product of the cellulose

unzipping is either glucose or some simple glucosan;

these monomers then undergo conversion to levoglucosan in some subsequent step. Such a proposal

has ample precedent in the older literature but is probably erroneous on several accounts: it does not

account for the large variation in products observed

when cellulose is pyrolyzed at different temperatures; there is no obvious reason why simple cleavages

of glycoside bonds should occur in a systematic

sequence as required for unzipping; the conversion

of glucose and simple glucosans to levoglucosans is usually a very low yield process whereas high yields are often obtained from cellulose; and, finally, the

intermediacy of glucose considerably complicates

the mechanisms which must be proposed to ac- curately describe the action of most flame retardants.

Of course, it is generally recognized that mechanistic

models are only useful approximations rather than

accurate descriptions of real chemical processes, but

in this case there is available a great deal of information which would seem to lead to a more

useful model.

Numerous proposals have been made in the

literature. As early as 1921, Irvine and Oldham [I]

proposed that the first step in pyrolysis was hydrolysis to glucose followed by dehydration of the

glucose to form levoglucosan. Essentially this same

mechanism was proposed again by Heuser |]2] in

1944. Similarly, Tamaru [3], in 1951, proposed that

the mechanism of cellulose pyrolysis involved

depolymerization of the macromolecules to some low

molecular weight species. He contended that some of

this species underwent dehydration and polymerization with continued and increased heating to

yield a high polymeric material, while the remainder

of the species underwent further decomposition.

The polymeric material was believed to be degassed

by heat with resulting aromatization and evolution

of carbon monoxide, carbon dioxide, methane, and

hydrogen.

Esteve and coworkers [4] believed that depolymerization of cellulose to levoglucosan through

scission of 1 ,4-glycosidic linkages, followed by

cyclization to the glucosan occurred during pyrolytic

degradations. It was thought that part of the levo180
glucosan polymerized and aromatized to form chars

while the remaining levoglucosan underwent destructive distillation to form flammable tars, volatile

gases, and water. Ivanov et al. [5], in 1956, independently proposed a similar mechanism by

which cellulose pyrolytically degraded in vacuo

through splitting of the 1 ,4-glycosidic linkages with

concurrent formation of levoglucosan.

Martin [6] and Day [7, 8] suggested that

pyrolytic degradation of cellulose might involve a

free radical process; however, neither investigator

described his suggested mechanism in detail. Furthermore, they were unable to find evidence to support

their postulations although Arthur [9] has recently

reported ESR spectral iiata which may be interpreted as evidence of homolytic degradation. Esteve

[4] calculated the energy of activation for the

pyrolysis of cellulose and found it to be substantially

lower than that required for a free radical process.

Madorsky and coworkers [10], in 1956, proposed

a mechanism based upon random dehydration along

the cellulose chain occurring simultaneously with

scission to produce levoglucosan, water, carbon

monoxide, and carbon dioxide. They based this

theory on the fact that C-0 bonds are thermally

less stable than C-C bonds. In 1958, Madorsky

refined the theory to involve thermal scission of the

C-0 linkages between rings, followed by transfer

of a hydrogen from the primary alcohol to the oxygen

of the broken glycosidic linkage and formation of an oxygen linkage between C-1 and C-6. He deduced

that, when a similar scission took place at a point one

structural unit removed from that of the original

scission, levoglucosan would be formed. Scissions

at points more than one unit removed from the

original should result in nonvolatile multiple unit

chain fragments [11].

Schwenker [12, 13] reasoned that if the pyrolysis

of cellulose involves direct depolymerization to

levoglucosan, then modifying cellulose at the C-6

positions should minimize levoglucosan formation

and decrease flammability. In support of his proposal,

he oxidized approximately half of the primary

hydroxyls of a cellulose sample and found that when

pyrolyzed it yielded only half as much levoglucosan

as did the untreated sample. The sample also

exhibited reduced flammability.

In 1957, Golova et al. [14] investigated the

effects of the packing density of cellulose chains on

the pyrolytic decomposition of cellulose in a vacuum.

They found untreated cellulose to give approximately 60 percent levoglucosan, while mercerized

cellulose yielded only 38 percent, hydrocellulose

from cuprammonium solution about 15 percent,

cellulose hydrate from unoriented viscose staple

about 4 percent and cellulose hydrate from oriented

viscose rayon about 5 percent. With this information

they concluded that a change in the density of the

cellulose chains caused a change in the courses of the

chemical processes occurring during pyrolytic decomposition. Decreased packing density was believed

to cause a decrease in levoglucosan formation, while

increased packing density hindered the side processes

of hydrolysis and dehydration, thus increasing

levoglucosan formation.

Golova also investigated the theories which

considered levoglucosan to be formed through an

intermediate stage in which cellulose was hydrolyzed

to glucose and then dehydrated. Golova found that

pyrolyzed cellulose gave a levoglucosan yield of

60 percent under vacuum whereas a-6-glucose gave

only 5-6 percent. Decreasing the degree of polymerization of cellulose to 150 did not significantly

affect the yield of levoglucosan, yet cellobiose was

found to yield only 3-4 percent levoglucosan. From

these data, Golova et al. [15] concluded that the

glucose intermediate theory was incorrect. They

further proposed that in regions of high packing

density, levoglucosan formation is favored and side

reactions are secondary. In regions of different

physical structure, dehydration and hydrolysis are the primary reactions with the production of

levoghicosan less favored. Based on this information,

Golova and coworkers proposed a theory, similar to

that later proposed by Madorsky [10], in which

levoglucosan was thought to be formed through

scission of the 1 ,4-glycosidic linkages followed by

hydrogen abstraction and cyclization to levoglucosan.

Further, they concluded that there is some limiting

D.P. below which levoglucosan formation is not

favored.

In an effort to gain further support of their

cellulose pyrolysis theory, Golova et al. [16] degraded cotton cellulose through mild hydrolysis to

obtain samples which represented a range of D.P.

They thermally degraded the samples at constant

temperature for successively increasing time increments and found that all of the samples rapidly

degraded to a limiting D.P. of about 200 which

remained unchanged during further decomposition.

They then concluded that the mechanism of pyrolysis

involved initial breakdown to a D.P. of 200, followed

by complete depolymerization of the chains yielding

levoglucosan. Golova [17] later showed that levoglucosan production was low during the initial

stages of pyrolysis and that a large amount of

volatile products was produced. In the later stages

of pyrolysis the yield of levoglucosan increased and

remained constant as degradation continued.

In reporting their work, Golova and coworkers

did not speculate on the reason for the apparent

stability of the cellulose molecule having a 200 D.P., nor did their results establish whether this number

represents an average over a broad spectrum or whether a 200 D.P. was actually a critical size for the

cellulose chains. They did surmise the latter. From

their observations, it might be further hypothesized

that the 200 D.P. represents the average length of

the crystalline regions, where chain packing density

would be greatest. Thus, one might expect the

initial breakdown to be that of the amorphous areas where side reactions would be favored. Breakdown of

the amorphous areas could be followed by complete

depolymerization of the chains within the crystalline

areas, where levoglucosan formation would be

favored. This would be consistent with the recent

181
results of Parker and Lipska [18] and Weinstein

and Broido [19] which were reported by Mr. Alger.

Honeyman [20] investigated cellulose pyrolysis to

gain information applicable to the improvement of

flame retardants for cellulose. He pointed out that,

in order to form levoglucosan from the anhydroglucose units of the cellulose chain, an interconversion from the all-equatorial CI to the higher

energy, all-axial IC conformation must occur (see

Scheme I). If this process were an integral part of

the depolymerization process it would account both

for the difference in products observed at various

temperatures and the sequential nature of the chain

imzipping reaction. Only at higher temperatures

would there be sufficient thermal energy for a

significant number of the cellulose anhydroglucose

units to exist in the strained IC conformation; thus

only at higher temperatures would there be a significant amount of unzipping and levoglucosan formation relative to dehydration, cross-linking, and

aromatization. But in those cases where sufficient

thermal energy was available for the CI to IC interconversion to occur at the reducing end of the

cellulose chain, the glysosidic bond joining the

terminal imit to the one adjacent to it would be

axially disposed relative to the terminal ring and thus

subjected to considerable strain. This would weaken

the bond and facilitate the cleavage of the terminal

unit from the chain. At the same time, the adjacent

unit would be converted into a 1 ,6-anhydro structure

with a IC conformation. Thus the cleavage of each

unit from the reducing end of the chain would

labilize the next unit for cleavage and produce the

observed unzipping reaction.

Mack and Donaldson [21] have recently proposed

a mechanism similar to this to explain the reduction

in flammability which they and others have observed

from samples of cotton cellulose treated with

inorganic bases. They proposed a base-catalyzed

degradation below the minimum temperature neces- sary for the CI to IC interconversion. This catalyzed

reaction should be in a direction other than that

which would produce levoglucosan. We have also

proposed a similar explanation for the flame retardant action of certain phosphates in which phosphorylation of the cellulose occurs at C-6 to block

the unzipping reaction and inhibit levoglucosan

formation. Simultaneously the degradation reactions

are assisted so that they occur more rapidly at lower

temperatures and thus produce predominently

nonflammable gases and char [22]. We have found

the efficiency of various classes of flame retardants to

correlate with their ability to phosphorylate, crosslink, and catalytically dehydrate cellulose as shown

by thermal analysis (23, 24), thermochemical

studies [25] and investigations of model substrate

reactions [26].

The majority of the investigations of cellulose

pyrolysis have been carried out near or below the

cellulose combustion temperature but they are probably valid nevertheless. Holmes and Shaw [27]

investigated the pyrolysis of different cottons in

vacuum and dry air mainly at 418 °C while BerkowitzMattuck and Noguchi [28] investigated the effect

of one-second exposure of cellulose to high radiant

flux levels of 5-25 cal/cm- s. From both investigations

it was concluded that the same mechanism observed

in lower-temperature pyrolyses operated in these

investigations. Reductions in tar formation with

increased char yields were reported to be the resultant

effects of flame retardants in these investigations.

This was also the case in lower-temperature degradations.

The mechanism of cellulose pyrolysis after the

formation of levoglucosan still remains vague.

Berkowitz-Mattuck and Noguchi [28] proposed that

fragmentation of levoglucosan through a free

radical mechanism accounted for the observed gases,

aldehydes, ketones, and unsaturated products. This

mechanism although logical, has not been substantiated.

Tang and Neill [29] used differential thermal

analysis (DTA) and thermogravimetric analysis

(TGA) in their studies concerning the effect of

flame retardants on the pyrolysis and combustion of

a-cellulose. Their TGA studies showed that flame

retardants lowered the threshold temperature for

pyrolysis and raised the char yield of the products.

DTA studies in oxygen showed that although

flaming was reduced, flame retardants often extended

glowing. They explained that flame retardants cause the heat of combustion to be liberated over a wider

temperature range and at a lower maximum intensity. Results of depolymerization and dehydration

of cellulose and of polymerization and aromatization

of pyrolysis products to a graphite-like char gave a reduced heat of pyrolysis.

In some of their most recent work. Tang and i coworkers [30] postulated four major types ofl
reactions occurring in the pyrolysis from observations related to their DTA studies on wood cellulose.

At 220 °C, cellulose undergoes a decomposition

involving loss of water to yield a partially dehydrated

species. Mass spectrometric thermal analysis con- firmed loss of water. Depolymerization of the

dehydrated species and volatilization of the resulting

products begins at about 280 °C. In this process,

previously untreated cellulose breaks into smaller

fragments which are volatilized via an endothermic

process competitive with the dehydration. The

remaining reactions involve decomposition of the

dehydrated cellulose via exothermic reactions into a number of gaseous products and residual char, and

condensation of simpler substances from the decomposition products. Effective flame retardants are reported to minimize the depolymerization producing

flammable tars (levoglucosan) and stimulate dehydration producing water and char.

References

[1] Irvine, J. C. and J. W. H. Oldham., J. Chem. Soc, Trans.

119, 1744, (1921).

[2] Heuser, E., The Chemistry of Cellulose (John Wiley and

Sons, Inc., New York, 1944).

[3] Tamaru, K., Bull, of the Chem. Soc, Japan 24, 164-8,

(1951).

[4] Esteve, R. M., et al.. Progress Reports Nos. 12 and 13,

QM Contract No. DA 44-109-QM-1414, University of

Rhode Island, Kingston, R. I.

[5] Ivanov, et al. Izvest. Akad. Nauk. S.S.S.R., Otdel. Khim.

No. 10, 1266. (1956).

[6] Martin, S. Research and Development Technical Report

USNRDL-TR-102. (1956).

[7] Day, A., et al. Progress Report No. 1, QM Contract No.

DA 19-129-QM-211, University of Penn., Philadelphia, Pa. (1954).

[8] Day, A., et al. Final Report, QM Contract No. DA-19-

129-QM-211, Univ. of Penn., Philadelphia, Pa.

[9] Arthur, J. C. and Hinojosa, O., Text. Res. J., 36, 385

(1966).

[10] Madorsky, S. L., et al. J. Res. Nat. Bur. Std. 56 (6), 343-

54. (1956).

[11] Madorsky, S. L., et al. J. Res. Nat. Bur. Std. 60, 343-9

(1958).

[12] Schwenker, R. F., Jr., and E. Pacsu. Chem. Eng. Data

Ser. 2, 83-8. (1963).

[13] Schwenker, R. F., and L. R. Beck, Jr. J. Poly. Sci., Pt. C,

2, 331-40 (1960).

[14] Golova, O. P. et al. Doklady Akad. Nauk. S.S.S.R. 112,

430-2. (1957).

[15] Golova, O., et al. Doklady Akad. Nauk. S.S.S.R. 115,

1122-5. (1957).

[16] Golova, 6., et al. Doklady Akad. Nauk. S.S.S.R. 116,

419-21. (1957).

[17] Golova, O., et al. Vysokomole-Kulyarni Soedineniya 1, 1295-1304. (1959).

18. Parker, W. J. and Lipska, A. E., TR-69, U. S. Naval

Radiological Defence Laboratory, San Francisco, CaUf.,

1969.

[19] Weinstein, M. and Broido, A., Combustion Sci. and

Tech., 1, 287 (1970).

[20] Honeyman, J. Final Progress Report, Contract No.

UR-E29-(20)-9, Shirley Institute, Didsbury, Manchester, England, April 1964, p. 103-4.

[21] Mack, C. H., and D. J. Donaldson. Textile Res. J. 37

(12), 1063-71. (1967).

[22] Hendrix, J. E., PhD dissertation, Clemson University,

1970.

[23] Hendrix, J. E., et al., J. Appl. Poly. Sci., 14, 1701 (1970).

[24] Wu, H. S. and Barker, R. H., Abstracts of Paper pre- sented SE-SW Joint Regional Meeting, Am. Chem.

Soc, New Orleans, La., 1970; Wu, H. S., M.S. Thesis,

Clemson University, 1970.

[25] Yeh, K. and Barker, R. H., Abstracts of Papers presented

160th National Meeting, Am. Chem. Soc, Chicago, 111., 1970.

[26] Hendrix, J. E. and Barker, R. H., unpublished results.

[27] Holmes, F. H., and C. G. H. Shaw. J. Appl. Chem.

{London) 11, 210-6. (1961).

[28] Berkowitz-Mattuck, J. B., and T. Noguchi. J. Appl.

Polymer Sci. 7, 709-25. (1963).

[29] Tang, W. K., and W. K. Neill. J. Polymer Sci., Pt. C (6), 65-78. (1964).

[30] Tang, W. K., et al. U. S. Forest Service Research Paper

FPL 82, Forest Products Laboratory, Madison, Wisconsin, January 1968.

183

NATIONAL BUREAU OF STANDARDS SPECIAL PUBLICATION 357, The Mechanisms of

Pyrolysis, Oxidation, and Burning of Organic Materials, Proceedings of the 4th

Materials Research Symposium, Oct. 26-29, 1970, Gaithersburg, Md. (Issued June 1972).

The Surface Pyrolysis Boundary Condition for the Combustion

of Polymers*

R. F. McAlevy III** and W. S. Blazowski***

Combustion Laboratory, Stevens Institute of Technology, Hoboken, New Jersey 07030

The pivotal boundary conditions for polymer combustion involve the surface pyrolysisvaporization characteristics. Generally, conditions at the burning surface (e.g., 500 °C or so) are

so different from those encountered in classical bulk pyrolysis experiments that these data cannot

be extrapolated reliably—data must be obtained at the conditions of interest. A critical review

of experiments designed to produce such data, and the possible interpretations of these data are presented.

The authors have employed a combustion self-heating technique. In the experiment a planar

diffusion flame is established between an axi -symmetric, "stagnation -point" flow of oxidant and the

pyrolysis-product vapors emerging from the polymer surface lying below. The surface temperature,

regression rate, and oxidant mass flow rate (the principal control variable) are measured. Recent

results of this experiment are presented and discussed.

Key words: Combustion; flammability; kinetics; polymer burning; pyrolysis; rapid heating;

surface pyrolysis; surface temperature measurement; thermal degradation.

1. Background and Introduction

The combustion behavior of polymeric materials

that are incorporated into solid rocket propellants

has been actively researched for more than a decade.

It was assumed that better control of the rocket

would be possible once this information was in hand.

More recently, the combustion behavior of polymeric materials that are incorporated into fabrics,

human habitats and vehicles has become a focus of

attention. It is assumed that control of fire hazards

would be possible once this information is obtained.

This paper reviews previous studies and describes

a current study of the pivotal boundary condition

common to all polymer combustion situations—the

characteristics of surface decomposition-vaporization

in response to intense surface heating.

During combustion of a polymeric material,

decomposition -product vapors emerge from the

surface and are burned with the oxidant component

of the environment in a nearby flame. To produce

ignition, an external energy source is required to

initiate vaporization and the flame itself. Once

ignition occurs at one site, the spreading of the flame

to engulf adjacent surface locations is conveniently

viewed as a continuous ignition process Ql].^ To

become self-sustaining, some fraction of the chemical

energy carried to the flame by emerging vapors must

* This work is sponsored by the OflBce of Naval Research contract number NONR 263-48.

** Director Epoxy same as in table 1. *** Previously, Ph.D. candidate sponsored by an NDEA fellowship.

Presently, Air Force Aero Propulsion Lab./SFF, Wright-Patterson AFB,

Ohio.

1 Figures in brackets indicate the literature references at the end of this paper.

be fed back to the surface in the form of thermal

energy to support further vaporization.

The combustion theorist formulates the detailed

mass and energy balance appropriate for the situation of interest, and then attempts to solve the

equations to predict ignition delay, flame spreading

rate, steady-state deflagration rate (surface linear

regression rate), etc. But in all cases, the pivotally

important boundary condition is the surface decomposition-vaporization characteristics.

Conditions at the burning surface are outside the

reach of equipment normally found in polymer

research laboratories. Surface heating rates approaching a hundred cal/cm^/s produce surface

temperatures of several hundred °C, subsurface

temperature gradients approaching a thousand

°C/mm, surface regression rates approaching a mm/s, and result in the subsurface material being

heated to surface conditions at several hundred

°C/s.

With other purposes in mind, for example negligible

temperature gradients in test specimens, the typical

polymer research laboratory employs "bulk heating"

techniques—thermogravimetric analysis (TGA),

differential thermal analysis (DTA) and differential

scanning calorimetry (DSC) —which are designed to

produce exposures far removed from typical combustion conditions. For example, extreme operations

of such equipment results in a rate of temperature

increase of only a few °C/s. And even though

temperatures of a few hundred °C can eventually be

reached, the lifetimes of the intermediary decomposition products are so much longer, vis-a-vis the

combustion situation, that it is unrealistic to expect

the same chemical path is followed in both cases.

185
For example, the temperatures at which important

events occur, including melting, and charring, are generally not the same in both cases. And since the

high surface viscous shear stresses of combustion

situations cause "blow away" of the ubiquitous

liquid-like layer at the surface of burning polymers,

information concerning melting is of special interest.

Thus the combustion research community largely

ignores the results of bulk heating decomposition

experiments. Instead they attempt to determine the

required information by means of experiments that

produce conditions at the polymer surface which

approximate combustion conditions. At the high

surface heating rates employed in these experiments

decomposition is restricted to the surface layer of

polymer, with attendant steep temperature gradients.

Frequently this is referred to as "surface pyrolysis".

As the heated surface steadily regresses as a result of

the exposure to heat, the term "linear pyrolysis" is employed interchangeably in describing the phenomenon.

Typical data are measured surface regression rate

(r), and surface temperature (Ts). Due to the steep

subsurface temperature gradients, and other complications, it has been found that unambiguous

measurement of Ts is extremely difficult. Arrhenius

plots of such data, that is, ln(r) versus 1/Ts are frequently presented, and the slope interpreted in

terms of a "rate controlling" surface process, or "surface kinetics". Other models have been proposed

that assume subsurface decomposition processes to

be rate controlling. But the validity of all are open to

serious question. Occasionally, correlations of r with

surface heat flux are attempted to assess the energy

required for vaporization—"surface energetics".

Very rarely is the chemical composition of the

emerging vapors measured.

This paper is intended to be a review of the

experimental methods used to obtain polymer

surface pyrolysis data, and to describe the current

state of understanding of the process as well as a new

approach to its study. The authors caution the

reader that almost all existing data are of a quasisteady nature, so that blind application to essentially

unsteady combustion situations, such as ignition and

flame spreading, could lead to erroneous prediction.

2. Review of Experimental Approaches

for Investigation of Linear Pyrolysis

In this section some of the various techniques

employed to obtain linear pyrolysis data will be

discussed. These will be divided into categories

determined by the principal mode of surface heating

employed—conduction, radiation or convection, as well as controlled self-heating during combustion.

2.1. Conduction

The first attempts to obtain linear pyrolysis data were made by Chaiken [2] and others over 10

years ago. They attempted to obtain surface kinetics

data by means of a technique that involved forcibly

pressing a polymeric strand against a flat, heated

plate, termed the "hot-plate technique". Due to the

intense surface heating rates so generated, linear

pyrolysis rates approaching the range encountered

in combustion situations were produced. A thermocouple placed in the plate was intended to measure

the specimen "surface temperature".

This technique has several shortcomings: (1) The

actual strand surface temperature is different from

the measured plate temperature due to the presence

of a thin vapor layer flowing between the hot plate

and the vaporizing surface [^4 J; (2) if the surface

decomposition process is influenced by the partial

pressure of decomposition product at the surface,

the presence of the solid hot-plate might produce

conditions at the surface sufficiently different from

those encountered during combustion to render the

information useless as a realistic boundary condition

for the combustion process [5]]"; and (3) the vigorously emerging vapor flowing across the specimen

surface might produce "erosion" of the specimen [_6'], which would be especially serious if the material

liquifies as it approaches the surface [1^.

In an attempt to eliminate some of these shortcomings, a porous hot-plate was substituted for the

solid hot-plate [8], and the device was employed to

obtain data with inorganic materials. A one-dimensional gas phase theory has been developed to

describe this experiment [9^- Empirically, large

data scatter resulted even when subliming materials were investigated. Although, to the authors' knowledge, this technique has never been used to investigate polymeric materials, the high probability

of the surface liquid layer flowing into the porousplate prior to vaporization renders this technique

quite unattractive for these materials.

In short, the conductive heating approach appears

to have little future.

2.2. Radiation

Shannon [10, llj and Lincoln [12] have exposed

polymer specimens to radiant heating. A radiation

furnace provided quasi-steady conditions while a "flash-tube" device was used to produce essentially

unsteady exposures (e.g., exposure times of approximately 10 ms)

The basic drawbacks of this approach are that:

(1) the radiation that passes through the surface

and is absorbed in depth results in a subsurface

temperature distribution much different than that

experienced in a burning polymer, where internal

radiation is generally negligible; (2) the fraction of

the total incident radiation that passes into the

surface is generally imknown since polymer surface

spectral properties (i.e., reflectivity), are generally

unknown at the conditions of interest; and (3)

vapors emerging from the surface absorb an unknown

fraction of the total radiation that would otherwise

strike the surface.

The most encouraging aspect of the radiation

' This eCf^ect is discussed in section 5.

186
heating method is the potential it holds for obtaining

transient linear pyrolysis characteristics, but the

drawbacks mentioned weigh heavily against it.

2.3. Convection

Thermoplastic materials have been subjected to

intense convective heating by immersion in the

exhaust of a small laboratory rocket motor [13].

In some experiments, the influence of environmental

chemical composition was investigated by passing

candidate gases up through specially-prepared,

porous specimens, which then emerge at and bathe

the heated surface.

As surface regression rate increases with increasing

enthalpy of the impinging jet, substitution of

electrical heating for the chemical conversion in the

rocket motor—that is, the use of arc-jets [14, 15]

produces polymer regression rates approaching

several mm/s [16]; approximately ten times greater

than the typical maxima with the chemical rocket

apparatus.

Convective heating appears to be superior to

conductive and radiative heating techniques due to

the fact that the experimental condition at the

specimen surface is relatively well-defined, and there

is ready optical access to it so that melting and

charring can be observed.

The major drawback seems to be that the high

velocity jet impacting and running over the heated

surface can "blow away" liquid from the polymer

surface melt layer. This additional mechanism for

surface mass removal makes difficult determination

of the actual surface decomposition-vaporization

rate, and spuriously high values of this quantity are sometimes reported.

2.4. Self-heating During Combustion

A particularly useful form of surface heating

utilizes a situation whereby the controlled combustion of pyrolyzed vapors above the surface

provides heat feedback to drive the decomposition

process. It possesses many of the advantages of the

convective technique while oftentimes having a low

velocity of external flow, and hence reduced influence on the melt layer.

Although mainly concerned with modeling a particular practical combustion situation, many

studies of hybrid rocket motor characteristics

[17, 18, 19] can be considered to be in this category

of "self-heating" pyrolysis experiments. It is possible

that parameters relevant to pyrolysis energetics

could be obtained in this way. However, in most

cases the flow in the combustion region was turbulent

and could not be characterized in a fundamental

manner. And thus surface heat flux predictions have

an unacceptable element of uncertainty, which in

turn is reflected in values of parameters extracted.

Other investigators [7, 20, 21, 22] have employed

self-heating techniques involving better defined fluid

mechanics. Simplistic attempts to calculate heat flux

to the surface have been made [21, 22], but they

have not been sufficient for accurate determination of

surface energetics parameters.

Since surface heat flux depends on combustion of

the pyrolysis product, self-heating techniques usually

do not have as large a range of data as other techniques. The range of data in a self-heating method

depends on the flammability limits of the particular

system studied, while other techniques have limits

depending on the ultimate output of the heating

apparatus employed.

Finally, a significant drawback is that the surface

pyrolysis characteristics of all polymeric materials

cannot be investigated with these techniques. Some,

such as Teflon, cannot be ignited and burned under

ordinary conditions. Others which do ignite but char

as they burn, such as acetate, make measurement of

surface quantities (temperature, and even regression

rate) very difficult.

3. Technique for Measurement of

Important Quantities

The three principle quantities to be measured in

surface pyrolysis experiments are the surface

regression rate, surface temperature, and the heat

flux to the surface. The regression rate measurement

is readily accomplished by measuring the length of

sample pyrolyzed in a given amount of time. Assuming the loss due to blow-away can be taken into

account, it is a relatively simple quantity to determine. Surface temperature and heat flux determination pose much more difficult problems for the

investigator. Techniques used for these measure- ments will be discussed in the following two subsections.

3.1. Surface Temperature Measurement

One common method of determining surface

temperature is by means of an embedded thermocouple that is overrun by the regression surface.

Errors are introduced by disturbances due to

introduction of the thermocouple in the specimen—it

is usually cemented together to hold the thermocouple in place. Thus the bead detects events in the

cemented, rather than original, specimen. Errors are introduced due to the presence of the thermocouple

in the subsurface temperature field—the different

values of thermal conductivity of the metal thermocouple and specimen produce distortions of the heat

flow and local temperature distribution patterns. The

magnitudes of these errors are almost impossible

to predict.

The embedded thermocouple technique has two

other drawbacks: (1) steep temperature gradients

below the surface may cause temperature differences

of, typically, 100 °C across the thermocouple bead;

and (2) the time-response of the thermocouple

might introduce unacceptable errors in the rapid

linear pyrolysis situations. The effect of each of these

problems is minimized as the thermocouple size is decreased [23]. Thermocouples with bead diameters

as small as 15 fjim have been used [24, 25].

187
Although many investigators [_23, 26, 27] have

attempted to provide corrective schemes, the

combined effects of all the errors (those due to

moimting being perhaps the principal contributors)

makes the embedded thermocouple technique unattractive.

Some data has been obtained by optically measuring the surface temperature [7, 13, 20, 28J. In

these experiments infrared radiation emitted by the

burning polymer in a specific wavelength band is recorded and compared to the radiation emitted by

a blackbody source at known temperature. If the

spectral properties of importance (surface reflectivity and material absorptance) are known, the

measurement may be used to calculate surface

temperature. These properties are not known at the

temperatures of interest, so values obtained at lower

temperatures must be used in reducing data.

This technique, however, can only be applied in

the convective or self heating experiments. In

conductive experiments the surface cannot be

observed. In radiative heating experiments the

possibility that some reflected radiation may be

detected makes this technique generally unsuitable.

A possible drawback with this technique is that

the hot gases between the sensor and the polymer

surface may contribute to the radiation level

detected. This is especially serious in the convective

heating case employing rocket motor exhaust gases

containing H2O, COo, and CO, all of which can be

strong emitters in the infrared wavelength regions of

interest. With the self-heating techniques, combustion and formation of hot products occurs in a much thinner layer near the surface and its contribution can generally be neglected.

Simon [16J employed a novel radiative film

technique for measurement of surface temperature.

It involved use of a very thin film embedded in the

polymer and oriented in the same plane as the

regressing surface. Thermal radiation emitted by the

film in the direction away from the pyrolyzing sur- face is measured and compared to a calibration,

previously obtained at lower temperature, for the

determination of surface temperature. The thin film

(fractions of a micron in thickness) is silvered on the

side facing the polymer surface and blackened on the

side facing the sensor. Thus, it is not heated significantly by the radiation that impinges on it from the

pyrolyzing surface, and it emits the maximum

amount of radiation in the direction of the sensor. Although, because of the film's dimensions, problems

of response, temperature difference, and lead loss are minimized with this technique, the complications

due to the mounting technique are still encountered.

3.2. Heat Flux Determination

Surface heat flux is the other important quantity

to be measured. Calorimetric techniques have been

employed for this purpose with both the convective

[13 J and radiative [10] heating methods.

The amount of convective surface heating will be

appreciably greater with the calorimeter in place

because there is no "blowing" of pyrolyzed vapor

PRESSURE REGUUTOR

PRESSURE GAGE

' CRITICAL ORIFICE

/

OXIDAHT

SUPPLY

OXIDAHT

STREA'-t DIFFUSION FWy£

PYROLYSIS VAPORS

COOLING JACKET

FLAT PUTE

FOLYi'CR SM-tPLE

ADVATKEIIEIIT MECHANIS!1

Figure 1: Schematic representation of experimental apparatus,

from the surface (a transpiration cooling effect) in

this case. Various means of correcting the calorimetric data to compensate for this effect can be used,

although they may require some questionable as- sumptions concerning composition of decomposition

products, etc. With radiation heating one must be concerned

w ith the different reflectances of the calorimeter and

polymer surfaces, as well as the absorbed radiation

when pyrolyzed polymeric gases appear above the

surface.

Measurements of heat flux in the conductive and

self-heating experiments at present appear to be

insuperably difficult to obtain. However, it is possible to model analytically the self-heating

combustion situation to calculate the surface heat

flux. This requires that the combustion region be

well-defined and easily characterized fluid dynamically.

4. Surface Pyrolysis by Means of the

Diffusion-Flame Technique

A self-heating experiment used by the authors to

study polymer linear pyrolysis will now be discussed.

In particular, a diffusion -flame technique is employed.

In the experiment, a diffusion-flame is established

between an axisymmetric flow of oxidant and the

decomposition vapors emerging from the polymer

surface lying below". Figure 1 is a schematic representation of the apparatus. An automatic advancement mechanism passes f in diam. rods of the

polymer specimen through a tight-fitting hole in a

flat plate. The pyrolyzing surface is kept even with

the flat plate which is a fixed distance (If in) from

the oxidant nozzle (f in diam.). A cooling jacket on

the underside of the flat plate arrangement insures

that no significant preheating of the sample will occur. The oxidant mass flow rate m is regulated by

controlling the pressure upstream of a choked

orifice. The orifice had been calibrated by a standard

flow meter. In most cases oxygen is used as the

oxidant. In the remainder, oxygen is diluted with

nitrogen, argon, or helium.

Temperature measurements are accomplished

with the optical technique. The monitoring wave188
580

,30

i.20

;,io

.06

SURFACF. PRIGHTOESS TFi'iPFRATURF. Cc)

SGO 5iiG 520 5U0

1 1

i [—

km

° m O2 KET

525; O2 NET + 1(0% 0 NET I ^2 -

X

tux

J I I I I I J L_J L_

1,20 1,25

RECIPROCAL SURFACE BRIGHTNESS TEf'IPERATURE Qy^K^So^

Figure 2: Regression rate versus reciprocal surface bright- ness temperature for Delrin.

length band is chosen to be one in which the absorption coefficient is high (typically 2,000 cm~'). If

this is so, and if the surface reflectivity is very low,

the assumption of unity emissivity (i.e. calculating

"brightness temperatures") may be made without

causing significant error. Since accurate determination of these radiative properties at elevated

temperatures (approaching those of the pyrolysis

situation itself) constitutes a formidable problem,

this laboratory's data has thus far been reduced to

brightness temperatures only. A comparison between

data taken at two separate wavelengths serves to

check the brightness temperature values.

The previously mentioned problems that occur in

self-heating experiments with materials that do not

burn easily and with charring materials, certainly

apply to the diffusion -flame technique. Although the

range of data is not as extensive as in other techniques, regression rate did vary by a factor of 5 in

most cases, (fig. 2).

The data may be presented in two plots. Kinetic

data are typically shown on an Arrhenius plot as

illustrated in figure 2 for Delrin (polyoxmethylene)

The important parameter obtained with such a plot

OXIDAItr FLOW RATE (gm'sEc) Figure 3: Regression rate versus oxidantflow rate for

PMM with 100 percent oxygen.

,50

,30

,20

,10

.06 A 1,0

nviDANT FLO'' RATE (niVSEC)

,1 ,2

Figure 4: Log-log plot of data from figure 3

is the apparent activation energy for a surface

process.^ Delrin, for example, exhibits a value of

approximately 23 kcal/mole. Note that the changing

oxidant composition does not seem to affect the

Arrhenius correlation of these data.

Empirical values concerned more with the energetics of surface decomposition are plotted as regression rate versus oxidant mass flow rate. Such

information is shown for polymethylmethacrylate

(PMM) in figure 3. A power-law correlation (i.e.

expressing the data in the form r = am^) for such

data may be performed on a plot of log regression

rate versus log oxidant flow rate, as shown in figure

4. Results of such a correlation for some polymers

are shown in table 1.

Table 1. Empirical power law (r = arhP) parameters

Material o!(nim/s) /3

Plexiglas 0.23 0.51

Plexiglas 55 .22 .53 Nylon 6-6 .24 .46

ICRPG binder (polyurethane) .28 .44

Delrin .22 .48

Polystyrene .17 .53 Polypropylene .23 .67

A theoretical analysis of the experimental situation

is required for the extraction of energetics parameters from measured values of a. Recently an

analysis has been produced [29] that goes beyond

the unrealistically simple efforts of the past [21, 22].

The combustion situation is modeled by an axisymmetric boundary layer flow where combustion is assumed to occur in an infinitesimally thin "diffusionflame". The analysis requires specification of oxidant

composition, surface heat of gasification, heat release

per unit mass of fuel reaching the flame, and the

stoichiometric ratio for the appropriate combustion

products. The last two quantities describe the gross

details of the chemical reaction at the flame.

Besides providing a basis for the interpretation

of a values for the evaluation of energetic param3 The bases for rationalizing pyrolysis data will be discussed in the

next section.

189
eters (i.e. surface heat of gasification), the analysis

predicts values for surface concentration of decomposition vapor. This concentration may be varied

by changing the oxidant composition, as was done in

generating the Delrin data for figure 2. In this way

the influence of the decomposition product partial

pressure at the surface on the decomposition process can be checked. In the Delrin example, no significant

effect is observed.

The analysis predicts that the regression rate

should vary as the square root of oxygen impingement rate (i.e. (3 = 0.5). Comparison with table 1 shows that all experimental values of this quantity

are reasonably close to the predicted value except

that for polypropylene. Significant deviation from

j8 = 0.5 may indicate a change in the energetic re- quirements of the degradation-vaporization process,

and hence, in the process itself, according to the

theory.

A future publication [30] will contain the kinetic

and energetic polymer data obtained with this

experiment. Further description of the experiment

and technique, including enlightening photographs

of the pyrolyzing surfaces during combustion also

will be contained therein.

5. Theoretical Foundations for

Data Interpretation

This section will concern different ways in which

linear pyrolysis data can be interpreted, based on

different assumed models of the actual mechanism.

Many investigators [31-34] have attempted to

rationalize linear pyrolysis data on the assumption

that the thermal degradation which occurs in a subsurface reaction zone is the rate-limiting step.

Shteinberg and Sokolova [31] point out that this

situation is very similar to the classical case of

gaseous reaction front propagation. So that for a "zeroth" order reaction, the apparent activation

energy of the process will be approximately half the

isothermal bulk activation energy value. And these

authors offer this as a first approximation of the

polymer linear pyrolysis problem.

Rabinovitch [32] used an assumed subsurface

temperature profile coupled with the assumption of a

first order degradation process to derive an integral

expression which correlates the surface temperature,

surface regression rate and the first order process

activation energy and frequency term. The assumed

temperature profile which is for the case of constant

surface regression rate (where the pyrolyzing

polymer is considered to be a semi-infinite, inert

solid with constant properties) follows the relationship [35]

T{x)~To = exi (1)

Where:

X is the distance beneath the surface.

T(x) is the temperature at position x within the

solid phase.

To is the temperature of the solid far from

the surface.

Ts is the surface temperature.

a is the thermal diffusivity of the solid,

assumed constant,

r is the surface linear regression rate.

Since the subsurface decomposition reaction will

involve a heat of reaction, eq (1) is not an accurate

representation of the temperature profile. An

improvement on the Rabinovitch model, then,

would include a heat of reaction term in the analysis.

This ameliorative step, however, requires a much

more complicated analysis.

Lengelle [36] accounted for subsurface reactions

by employing a matched asymptotic expansion

scheme. The outer expansion (for the region away

from the surface) takes into account only the

conduction heat transfer, while with the inner

expansion (the region close to the surface), the heat

of reaction term is taken into account also. A more

accurate mass conservation expression was also used

in this model.

Another common method of explaining pyrolysis

data involves the assumption of an unopposed

surface reaction [2, 13]. In this case correlation of

data is accomplished by employing the following

relationship

(2)

where Ai is a constant, R is the gas constant, and

Ea is the apparent activation energy of the surface

decomposition process. This equation will give nearly

a straight line in a log r versus 1/Ts plot (the linear

Ts dependence is weak and often neglected) . If data

plotted on these same coordinates results in a straight line, empirical values of Ai and Ea may

be found.

As Barrere and Williams [5] point out, however,

it is also possible that equilibrium conditions for the

decomposition process apply at the surface. If this

boundary condition is to be used, the following

relationship must be satisfied.

n (c^--

)

i=l (3)

where fi'/ is the stoichiometric coefficient for decomposition vapor species i, C, the surface concentration of species i, and Ks the equilibrium constant

for concentrations.

Naturally the surface process can also be intermediate between the two extremes of unopposed

and equilibrium surface gasification. If this is the

case, the following relationship governs the decomposition rate

r = A2Ts exp

PX

Pi

(4)

where A2 is a constant, P is the total pressure. Pi the

equilibrium vapor pressure corresponding to Ts' and X the mole fraction of decomposition vapor at the surface.

Equation (4) illustrates a possible hazard of

190
directly applying pyrolysis data to combustion

situations. In many cases (such as the hot plate and

arc jet) total pressure or mole fraction of the decomposition vapor at the surface may not be controllable or known.

6. Comparison Between Data and

Theoretical Foundations

The various techniques employed have generally

yielded conflicting data. Lack of a standard test material makes difficult establishment of their

relative validity—each experimenter used a different

material. For reasons previously discussed, the

authors believe that their diffusion-flame technique

will yield data that is well characterized and contains

less error than that of previous data. At this writing

insufficient data has been produced to allow an unambiguous evaluation of the technique.

Without reliable data, evaluation of the various

theoretical approaches is a futile undertaking. Once

sufficient data is available, however, the investigation of model validity will still be complex, for there

is no reason to expect that one model will be suitable

for all polymers. Much more work is required in this area.

7. Summary

After reviewing the known linear pyrolysis

techniques the authors have decided in favor of the

diffusion flame technique. Perhaps further work

with this technique will yield information useful in

the explaining of discrepencies which now exist

among the data generated by the many types of

surface pyrolysis experiments employed in the past.

A number of theoretical approaches to data

interpretation have been presented. Since past data

have been in serious disagreement, they cannot

provide a basis by which the validity of these models

may be evaluated.

Perhaps one of the causes of discrepencies in the

data is that different experimenters have employed

different materials in their investigations, and it appears no common material was ever tested. For

example, polymethylmethacrylate data are reported

by a number of experimenters [_2, 3, 7, 10, 16J but

some prepared it in their own laboratories while

others obtained it from different commercial sources. Thus, there was no control on curing, etc., and fine

comparison and interpretation of the resulting data

is unwarranted at this time.

The establishment of a standard testing material can be invaluable in helping experimenters calibrate

their equipment, etc. This is a necessary action

which must be taken if the various techniques, and

thus the quality of the data produced, are ever to be

meaningfully evaluated. The authors suggest that

the National Bureau of Standards take a leadership

position in this endeavor.

Once polymer surface pyrolysis is understood it will be possible to design polymers for many specific

combustion related purposes. For example, material

designed to be fire resistant may make use of the

release of bromine, or some other known fire retardant [37^, during the decomposition process. So a rational control of polymer pyrolysis properties can

lead to many improvements in materials for use in

solid rocket propellants, fabrics, building materials,

and many other applications where their combustion

behavior is of interest.

8. References

[1] R. F. McAlevy III, and R. S. Magee, The Mecfianism of

Flame Spreading Over the Surface of Igniting Condensed-Phase Materials, Twelfth Symposium (International) on Combustion, Academic Press, New York,

1969, p. 215.

[2] R. F. Chaiken, et al.. Kinetics of the Surface Degradation

of Polymethylmethacrylate, J. Chem. Phys. 32, 1960, pp. 141-146.

[3] W\ Anderson, et al., A Model Describing Combustion of

SoUd Composite Propellants Containing Ammonium

Nitrate, Combustion and Flame 3, 1959, p. 301.

[4] R. H. Cantrell, Gas Film Effects in the Linear Pyrolysis

of Solids, AIAA Journal 7, No. 1, 1963, p. 154.

[5] M. Barrere and F. A. Williams, Analytical and Experimental Studies of the Steady State Combustion Mechanism of Solid Propellants, in Advances in Tactical

Rocket Propulsion, AGARD Conference proceedings

No. 1, 1966, Technivision Services, Maidenhead

England, pp. 49-137.

[6] R. F. Chaiken, F. J. Cheselske, Surface Rate Processes and Sensitivity of High Explosives, J. Chem. Phys. 43,

No. 9, Nov. 1965, pp. 3228-3236.

[7] R. F. McAlevy III, S. Y. Lee, and W. H. Smith, Linear

Pyrolysis of Polymethylmethacrylate During Combustion, AIAA Journal 6, No. 6, 1968, p. 1137.

[8] R. L. Coates, Linear Pyrolysis Rate Measurements of

Propellant Constituents, AIAA Journal 3, No. 7, 1965,

p. 1257.

[9] W. Nachhar and F. A. Williams, On the Analysis of Linear Pyrolysis Experiments, Ninth Symposium (International) on Combustion, Academic Press, New York,

1963, p. 345.

[10] L. J. Shannon, and J. E. Erickson, Thermal Decomposition of Composite Solid Propellant Binders, 6th ICRPG

Conference Expanded Abstracts and Slides, Vol. 1, 1969, p. 519.

[11] L. J. Shannon, Composite Solid Propellant Ignition

Mechanisms, Final Report, AFOSR-69-1250TR, April

1969.

[12] K. A. Lincoln, Flash Pyrolysis of Solid-Fuel Materials by

Thermal Radiation, WSS5C1 Paper 64-6 presented at the Western States Combustion Institute Meeting,

Stanford University, April 1964.

[13] R. F. McAlevy III, and J. G. Hansel, The Linear Pyrolysis

of Thermoplastics in Chemically Reactive Environments, AIAA Journal 3, No. 2, 1965, p. 244.

[14] H. S. Schwartz, Laboratory Techniques for Studying

Thermally Ablative Plastics, Tech. Doc. Rep. No. ASD

TR 61-517, Dir. of Materials and Processes, WrightPatterson A.F.B., Ohio, January 1962.

[15] R. A. Greenburg and N. H. Kemp, Structure of the

Laminar Ablating Air-Teflon Boundary Layer, pre- sented at AIAA 7th Aerospace Sciences Meeting, Jan. 20-22, 1969, AIAA Paper No. 69-99.

[16] D. R. Simon, Rapid Linear Pyrolysis of Composite Solid

Propellant Ingredients, NASA CR-66856, Final Report,

1970.

[17] T. J. Houser, and M. V. Peck, Research in Hybrid Combustion, in Heterogeneous Combustion, Vol. 15 of Progress in Astronautics and Aeronautics Series, Academic

Press, New York, 1964, p. 559.

[18] G. A. Marxman, C. E. Wooldridge, and R. J. Muzzy,

Fundamentals of Hybrid Boundary-Layer Combustion, in Heterogeneous Combustion, Vol. l5 of Progress

in Astronautics and Aeronautics Series, Academic Press,

New York, 1964, p. 485.

191
[19] G. A. Marxman and C. E. Wooldridge, Research on the

Combustion Mechanism of Hybrid Rockets, in Advances in Tactical Rocket Propulsion, AGARD conference proceedings no. 1, 1966, Technivision Services,

Maidenhead, England, pp. 421-477.

[20] J. Fowling and W. A. Smith, Measurement of the Burning

Surface Temperatures of Propellant Compositions by

Infrared Emission, Combustion and Flame 6, 1962,

p. 173.

[21] R. J. Zabelka, and D. F. Brink, Studies in Hybrid Combustion, Final Report, Contract N123 (60530) 29155A,

Rept JPC-362; F 64-1, Purdue University, 1964.

[22] R. F. McAlevy III, R. A. Battista, and J. V. Havel,

Experimental Studies Bearing on Mechanistic Modeling of Composite Solid Propellant Combustion, Paper

#69-16 presented at 1969 Western States Comb. Inst. Spring Meeting, U.S. Naval Weapons Center, China

Lake, California, 28-39 April.

[23] W. D. Brewer, Effect of Thermocouple Wire Size and Configuration on Internal Temperature Measurements in a Charring Ablator, NASA TN D-3812, 1967.

[24] W. H. Smith, The Linear Pyrolysis of Polymethylmethacrylate, ME Thesis, Department of Mechanical

Engineering, Stevens Institute of Technology, Hoboken, New Jersey, 07030, 1966.

[25] A. J. Sabadell, J. Wenograd, and M. Summerfield, Measurement of Temperature Profiles Through Solid Propellant Flames Using Fine Thermocouples, AIAA Journal

3, No. 9, 1965, p. 1580.

[26] N. P. Sub, C. L. Tsai, Thermocouple Response Characteristics in Deflagrating Low Conductivity Materials,

ASME Paper No. 70-HT-7, 1970.

[27] N. R. Johnson, A. S. Weinstein, and F. Osterle, The Influence of Gradient Temperature Fields on Thermocouple Measurements, ASME paper 57-HT-18, 1957.

[28] P. L. Hanst, and A. C. Walker Jr., The Infrared Emission

Spectra of Plastics Ablating in a Low Enthalpy Air

Stream; Measurement of Surface Temperatures and

Temperature Profiles Behind the Surfaces, Conference on Behavior of Plastics in Advanced Flight Vehicle Environments, WADC Technical Report 60-101, Sept.

1960.

[29] W. Blazowski and R. F. McAlevy III, Investigation of

the DilTusion Flame Between a Burning Polymer and

an Impinging Flow of Gaseous Oxidant, presented at the 1970 Eastern Section Fall Technical Meeting of the

Combustion Institute, 5-6 November, Georgia Institute

of Technology. Also, Blazowski, W. S., Cole, R. B., and McAlevy, R. F. Ill, Effective Heats of Gasification of Polymers Under Combustion Conditions, AIAA

Paper No. 72-34, Presented at the AIAA 10th Aerospace Sciences Meeting, San Diego, Calif. Jan. 17-19,

1972.

[30] R. F. McAlevy III and W. S. Blazowski, An Experimental

Technique for Determining the Fundamental Combustion Characteristics ol Polymers, to be presented at 161st National A.C.S. Meeting in Los Angeles, March

28-April 2, 1971.

[31] A. S. Shteinberg and Sokolova, Linear Pyrolysis of Condensed Substances, Doklady Akademii Nauk SSSR

158, No. 2, 1964, pp. 448-451.

[32] B. Rabinovitch, Regression Rates and the Kinetics of

Polymer Degradation, Tenth Symposium (International) on Combustion, Academic Press, New York,

1965, p. 1395.

[33] T. J. Houser, Kinetics of Polymer Pyrolysis from Surface

Regression Rates, J. Chem. Physics 45, No. 3, 1966.

p. 1031.

[34] N. H. Kemp, Surface Regression Rate of an Ablating

Polymer, AIAA Journal, 6, No. 9, 1968, p. 1790.

[35] D. Rosenthal, The Theory of Moving Sources of Heat and

its Application to Metal Treatments, Trans. Am. Soc,

Mech. Engs. 68, 849, 1946.

[36] G. Lengelle, Thermal Degradation Kinetics and Surface

Pyrolysis of Vinyl Polymers, AIAA Journal 8, No. 10,

November 1970, pp. 1989-1996.

[37] P. Baham, Diagnostic Tests of Some Models of Flame

Spreading Over Condensed Phase Materials, M.E.

Thesis, Stevens Institute of Technology, Hoboken,

New Jersey, 1969.

Discussion

W, C. Kuryla (Union Carbide Corporation):

I wonder if Dr. IVIcAlevy or IMr. Alger could

explain why in certain cases a two second ignition

time in the AATCC standard vertical flame test yields a completely burned fabric while the same

fabric after exposed to a 12 second ignition time

burns only half the way before it extinguishes?

What are the important physical as well as chemical

factors which control this particular situation?

Raymond Alger

While I cannot explain Dr. Kuryla's observation

regarding the effects of ignitor time on the AATCC

tests, I am interested in what attempts have been

made to resolve the question. Two effects of the 10

extra seconds of heating would appear to be a change

in the sample environment, particularly the at- mosphere in the box and a change in the fuel. The

atmosphere could be altered both by the combustion

products from the torch and the increased sample

combustion and pyrolysis products. While fuel

modification is not a very plausible reason for the

observed effect, the longer heating time will generate

more pyrolysis products which can escape and leave

the fuel changed. If the atmospheric environment is responsible, the effect should vanish when the box

is removed. On the other hand, alteration of the fuel

should be independent of the box.

192
FORM NBS-114A (1-71)

U.S. DEPT. OF COMM.

BIBLIOGRAPHIC DATA

SHEET

1. PUBLICATION OR REPORT NO.

NBS SP-357

2. Gov't Accession

No.

3. Recipient's Accession No.

4. TITLE AND SUBTITLE

The Mechanisnas or ryroiysis, uxiaacion, ana Durning ui ut

ganic Materials Proceedings of the 4th Materials

Research Symposium

5. Publication Date June 1972

6. Perfonning Organization Code

7. AUTHOR(S)

Leo A . Wall , Editor

8. Performing Organization

9. PERFORMING ORGANIZATION NAME AND ADDRESS

NATIONAL BUREAU OF STANDARDS

DEPARTMENT OF COMMERCE

WASHINGTON, D.C. 20234

10. Project/Task/ Work Unit No.

1 1. Contract/Grant No.

12* Sp on s or in g Org sn i z ion Nsmc ^nd Acid rc s s

Same as No . 9

13. Type of Report & Period

Covered

Final

14. Sponsoring Agency Code

15. SUPPLEMENTARY NOTES

16. ABSTRACT (A 200-word or less factual summary of most significant information. If document includes a significant

bibliography or literature survey, mention it here.)

A symposium .on The Mechanisms of Pyrolysis, Oxidation, and

Burning of Organic Materials was held at the National Bureau of

Standards in October 1970. This volume contains the nineteen

papers presented and much of the discussion which followed.

These papers review and discuss the current status of kinetic

studies on the reactions of organic materials in both gas and

condensed phases. The topics covered include: pyrolysis of

hydrocarbons, pyrolysis of polymers, oxidation of polymers,

oxidation of organic compounds, burning of organic compounds and

burning of polymers. Particular emphasis is placed on the elucidation

of the mechanisms of reaction in terms of free radicals or other

transient species and physical effects.

17. KEY WORDS (Alphabetical order, separated by semicolons)

Burning; hydrocarbons; organic materials; oxidation; polymers; pyrolysis,

18. AVAILABILITY STATEMENT

fXl UNLIMITED.

I I FOR OFFICIAL DISTRIBUTION. DO NOT RELEASE

TO NTIS.

19. SECURITY CLASS

(THIS REPORT)

UNCL ASSIFIED

20. SECURITY CLASS

(THIS PAGE)

UNCLASSIFIED

21. NO. OF PAGES

199

22. Price

$ 3. 25

USCOMM-DC 66244.P71

ft U.S. GOVERNMENT PRINTING OFFICE : 1972 O - 447-708