Some Problems in Chemical Kinetics and Reactivity, Volume 1 9781400887712

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SOME PROBLEMS IN CHEMICAL KINETICS AND REACTIVITY

SOME PROBLEMS IN CHEMICAL KINETICS AND REACTIVITY VOLUME I

By N. N. SEMENOV Translated by Michel Boudart

1958 Princeton University Press Princeton7 New Jersey

Copyright © 1958, by Princeton University Press All Rights Reserved L. C. Card 57-10321

Princeton Legacy Library edition 2017 Paperback ISBN: 978-0-691-62642-0 Hardcover ISBN: 978-0-691-62881-3

This translation was prepared under contract with the National Science Foundation. Officers, agents, and employees of the United States Government, acting within the scope of their official capacities, are granted an irrevocable, royalty-free, nonexclusive right and license to reproduce, use and publish or have reproduced, used, and published, in the original or other language, for any governmental purposes, all or any portion of the translation.

Printed in the United States of America

Iv

FOREWORD

In the four decades since World War I a revolution has been accomplished in chemical kinetics, that branch of chemistry in which the element of time, speed, is introduced into the description of the chemical process. Prior to World War I, the dean of experimental kinsticista, Bodenstein, had demonstrated the existence of simple blmolecular processes in the hydrogen-iodine-hydrogen iodide system; with hydrogen and bromine he had provided, with Lind, the classical example of complex bimolecular kinetics; he had formulated the first basic rates of surface reactions, notably in the contact sulphuric acid catalytic reaction at platinum sur­ faces; finally, just prior to World War I, he had demonstrated, with the hydrogen-chlorine reaction, the first long chain reactions initiated by light and by swift particles, the α-particles from Radon and Radium - C. In the early years of the 192 ο's Langnruir formulated the acbepted lcinetles of surface reactions. Polanyl, Christiansen, and Herzfeld ex­ plained the hydrogen-bromine reaction in terms of a sequence of atommolecule reactions. Polanyi subsequently engaged experimentally in de­ termining the speed of such reactions. It emerged that atoms and radicals reacted much more efficiently with molecules than do molecules with each other. These intermediates were shown to play a dominant role in chain processes, soon extended to include oxidation, polymerization, and even thermal explosions and decomposition. Largely from the studies of Hinshelwood and the author of the present volume the concept of branched chains was developed to Interpret oxidation processes leading to explosions. A golden decade began In the 1930's with the use of parahydrogen and deuterium to explore kinetic phenomena and the development of absolute rate theory, principally by Eyring, to explain the inertia of chemical processes in terms of a potential energy barrier separating reactants and products. Atom or radical-molecule reactions revealed significantly low 'activation energies'. The part played by zero point energies in deter­ mining rates was elucidated. Gradually, the dominant role of atom-radical intermediates in a wide variety of chemical processes was revealed. The empirical analyses of Arrhenius and the formulation of rate processes ν

FOREWORD In terms of collisions between 'hard spheres' were modified to yield a more sophisticated treatment in which many of the abnormalities of the classical approach were given an acceptable interpretation in terms of the properties of an intermediate complex of definable energy and entropy. Chemical kinetics was on the highway to a refined theoretical formulation. The present volume and its successor present the fine fruits of that scientific advance. In this production Its distinguished author offers the harvest of fifteen years of meditation, research, and teaching, his personal reactions to the resialts accumulated since first he formulated his news twenty years ago in his earlier volume, 'Chain Reactions'. The books are at once a revealing index of the enormous advances that have been secured and at the same time, as the author insists, a demonstration that 'the fundamental problems have not yet been solved'. Hs pleads eloquently for a co-operative planning of investigation on a comprehensive plans rather than for 'disjointed observations on the course of this or that re­ action'. This will require 'a friendly collaboration between scientists of all countries'. In echoing this plea of Professor Semenov the present writer can testify to the stimulus that has come to him from reading this 'panoramic' survey. Young scientists entering this field of research will doubtless garner a richer reward for their effort. Those who do so from this English text will owe a debt of gratitude to the translator, Professor M. Boudart, for making the text available to χω. Also, to the Princeton University Press, following In the footsteps of the more venerable Oxford University Press, which first published the author's Initial effort in 1935» thanks are due for the publication; nor would this have been possible without the financial support of the National Science Foundation.

Hugh Taylor

PREFACE The first edition of this book was published In 1951+ by the Academy of Sciences. Its history Is the following. The Chemistry Division of the Academy of Sciences decided to hold in 1955 a National Symposlim on Chemical Kinetics and Reactivity. I was asked to present the opening conmunlcatlon- Working on it, I concluded that it would be useful to write and print a small book containing a survey of the literature as well as my personal views on the subjects to be discussed. The book appeared in 195¾ and during the first part of 1955, the Acadany of Sciences also issued a volume containing 69 papers presented at the meeting by various authors from all parts of the Soviet Union. The Symposium took place on June 20-25, 1955 and consisted In a general discussion of my book and of the volume of proceedings. At the Symposium, many Important problems relative to the science of chemical kinetics and reactivity were discussed. Only a small number of copies of the book were printed for the Symposium and soon the book was out of print. The Academy a3ked me to prepare a new edition. The book has been considerably expanded, taking into account new data available in the World literature and at the Insti­ tute of Chemical Physics. Besides, I wrote a new Section IV on 'Branched Chain Reactions with also a chapter on Thermal Explosions. The book contains various thoughts that have occurred to me and calculations that, together with a few collaborators at the Institute of Chemical Physics, I have gathered during the past ten or fifteen years, In research and teaching on Chemical Kinetics. The book has therefore a personal flavor, it does not include all studies on reaction mechanism and has no claim at being a Treatise or a Textbook. It treats mainly radical and radical-chain reactions. More than twenty years have elapsed since my previous book on 'Chain Reactions1. The field of chemical kinetics and chain reactions has grown enonnously and the World literature contains a huge number of papers on this subject. Nevertheless, as It seems to me, the majority of the fundamental problems have not been completely solved as yet. What a contrast with the situation In physics where the mechanism of nuclear reactions, In particular of fission, discovered in 1939, has

PREFACE been clarified In detail. With respect to the mechanism of chemical re­ actions, the situation is not very satisfying, in my opinion because re­ search In all countries concerns itself with disjointed observations on the course of this or that reaction and not with a plan of investigations from all sides. At the present time, such disjointed studies are mostly unsatisfactory and sometimes, maybe, useless to the development of the theory viewed as the elucidation of the complex mechanism of a particular reaction or a solution of the general problems of chemical kinetics and reactivity. Separate investigations without a panoramic view of partial and general problems may be conqpared to determining the coordinates of a few points on a surface the shape of which has to be visualized. Evi­ dently, If the surface is complex, it cannot be done in this fashion with­ out the help of hypotheses having little or no foundation. Only the concerted effort of physical chemists, organlcists and inorganicists, having as a goal a converging attack on the general prob­ lems of mechanism even of the main classes of reaction only and on the problems of reactivity associated with them, may lead to a real advance of chemical kinetics. Here also, friendly collaboration between scientists of all ( countries Is a necessity. I shall be happy, if this book is of some in­ terest to chemists and attracts their attention to the problems to be solved In chemical kinetics and reactivity. grateful acknowledgement to S. S. Polyak, A. B. Nalbandyan, N. S. Enikolopyan, D. G. Knorre, A. E. Shilov and many other collaborators at the Institute of Chemical Physics, Including A. N. Pravednlkov. They have helped In writing the new chapters and the additions to the old chapters, for this second edition.

N. Semenov

TRANSLATOR'S PREFACE Semenov's new book on the chemical kinetics of chain reactions first appeared in 195^· It was written as the prelude to a Symposiiim held in Moscow. Hie book came to my attention in 1955 and X enjoyed it so much that I considered translating it. The translation was ready to go to press in 1957 when some revisions and considerable additions were received from the author. These are incorporated in the present volume. No attempt has been made to edit the original text or to alter its Russian flavor. As the book grew during the past year, it was decided to publish it In two separate volumes. The first volume contains the first two parts of the work. The third and fourth parts on the 'Kinetics of Chain Re­ actions' will appear in Volume II. My thanks are due to Sir Hugh Taylor for encouraging this project and to the National Science Foundation which made it possible.

M. Boudart Associate Professor of Chemical Engineering Princeton, N. J. June, 1958

TABLE OP CONTENTS Page CHAPTER I: REACTIVITY OP MONORADICALS 1. Main Types of Radical Reactions 2. Activation Energy and its Experimental Determination 3- Data on Bond Energies U. Relationship Between Activation Energy and Heat of Reaction 5- Activity of Radicals and Molecules 6. Empirical Formulae to Calculate Bond Energies in Organic Molecules 7- Addition to the Double Bond 8. Decomposition of Radicals and Energy of the it-bond.. 9. Isomerization of Radicals 10. Additional Remarks on Substitution Reactions 11. Role of polar factors in Organic Reactions 12. Role of Polar Factors in Polymerization

1 1 4 10 29 33

57 65 68 7k 75 85

CHAPTER II: COMPETITION BETWEEN MONORADICAL REACTIONS 95 1. Competition Between Different Radical Reactions 95 2. Effect of Temperature and Pressure on the Competition Between Radical Reactions 99 3• Intermediate and Final Products of Chain Reactions.. 108 CHAPTER Ills REACTIONS OF DIRADICALS 1. Transition of Atoms to a Valence-active State 2. Reactivity of 0 2 , S2, Se2 Molecules 3- The Divalent State of Carbon 4. Diradlcals of Complex Structure 5. Production and Reactions of the Dlradicals 6. Diradical Participation in Chain Reactions

1^5 145 ik6 iM 149

159

CHAPTER TV: DISSOCIATION OF MOLECULES AND RECOMBINATION OF RADICALS 169 1. Homogeneous Initiation of Chain Reactions 169 2. Homogeneous Recombination of Radicals 175 3- Dependence of the Overall Kinetics on the Nature of Chain Termination 181 it. Peculiarities of Radical Generation and Recombination in Liquids i8it 5. Inhibition of Chain Reactions by Certain Additives.. 186 xl

COHTEiNTS Page CHAPTER V: INITIATION OP CHAIN REACTIONS BT IONS OP VARIABLE VALENCE 1. Pormatlon of Radicals and Radical-Ions byElectron Transfer 2. Thermodynamic and Kinetic Characters of Production of Free Radicals by Ions of Variable Valence 3· Initiation of Radical Chain Reactions by Ions of Variable Valence CHAPTER VI: WALL INITIATION AND TERMINATION OP CHAIN REACTIONS 1. Wall Destruction and Generation of Radicals 2. The Method of Differential Calorlnietry 3· Pree Valences on Reactor Walls . Conditions for Wall Activity in the Process of Radical Generation 5· Heterogeneous Radical Initiation with Impurities in the Gas Phase 6 . Heterogeneous Initiation Due to Molecular Reactions Producing Free Radicals 7· An Attempt to Apply the New Ideas to Ifeterogeneous Catalysis

Xii

197 197 200 ζ ok

211 211 216 220 225 227 228 230

PART OME RADICAL REACTIONS (Reactions of Chain Propagation and Branching)

CHAPTER I; REACTIVITY" OP MONORADICALS

i. Main Types of Radical Reactions Free radical reactions are of three different types: 1) Substitution reactions:

where R1, Ra, R^ are atoms on free radicals. Let us note that the majority of substitution reactions investigated so far concern cases where R 2 is an atom. 2) Reactions of addition to a multiple bond or to an atom with impaired electrons: (a) or (b) A particular case of type (a) Is polymerization:

2') Decomposition of radicals (the reverse of addition): (a) or (b)

reactions of the type

1

2

I. REACTIVITY OP MONORADICALS 3) Reactions of Isomerlzation of free radicals, for example:

All these elementary reactions (except processes of the types 2b and 2'b) necessitate the rupture of one bond In thereactingparticles and the formation of one new bond. The rate of all these processes is expressed by the formulae: (1)

(monomolecular decomposition, isomerization of radicals) (2) (blmolecular reactions of radicals with molecules in substitution or addition). The magnitude of the rate constant is a numerical expression of molecular reactivity for attack of a given bond by a given radical. The quantity € — the activation energy of the elementary step — gives the minimum energy for reaction required by the reacting system (a molecule or a radical in the case-of a monomolecular decomposition, or both reacting partners taken together In the case of a blmolecular process). Let us represent graphically the state of the reacting system by plotting its potential energy versus the 'reaction path'. Then the energy of activation of the process is determined by the height of the energy barrier which the reacting system must surmount. If, as shown on Figure 1, the reaction is endothermlc, the energy Activated State

Energy

Pinal State Initial State

Reaction Coordinate FIGURE 1

1. TXPES OP REACTIONS

3

of activation.is the sum of a specific activation barrier eQ and the energy difference between Initial and final states, i.e., the absolute value of the heat of reaction q (3)

β - «ο + kl

·

If the process is exothermic, e.g., the reverse of that pictured on Figure 1, the energy of activation is simply the height of the po­ eQ. tential barrier The difference between the activation energies of the backward and forward reactions is equal to the heat of the forward process. I O

_ 1

The pre-exponential factor is of the order of 1 ο sec for the majority of monomolecular reactions. (This is of the order of the vibration frequency of the bond being broken.) It is assumed that the de­ composition of radicals does not differ In this respect from the decompo­ sition of molecules which has been directly Investigated by Polanyl et al t1] for the case of iodides and by Szwarc [2] for a variety of organic compounds. In these cases, a? was found to be of the order of 1O1^ sec"1. Hie pre-exponential factor a2 for bimolecular reactions is fre­ quently close to the number of collisions 2 between radicals and molecules i.e., it is a quantity of the order of 10~10 cm^/seo or 101 ** cm^/sec-mole depending on the choice of concentration units. In fact, however, not every collision between a molecule and a radical leads to reaction even If the energy of the colliding particles suffices to surmount the activation barrier and therefore a2 = fz where f is the so-called steric factor, a quantity usually smaller than unity. Por many reactions, f varies In the range o.i to 1. Lately, a number of data have been obtained, mainly by Steacie and his school [ 3 , 4, 5 ] , which show values of f of the order of 1 o~^ - 1 o-1* In some hydrogen atom abstraction reactions from paraffin molecules. The question of the steric factor for these reactions must not, however, be considered as settled. Thus at the September 195a Discussion of the Faraday Society, contradictory experimental results were presented. Por the reaction H + C2Hg, according to Berlie and Le Roy 16], the activation energy is 6.8 kcal (the abbreviation 'kcal' for 'kcal/mole' will be used from now on) and the steric factor 4.8.1O-^. For the same reaction (D + C2Hg), Darwent and Roberts [7] found ε = 9 kcal and f = 0.6. Subsequently, we shall use these values since we see no theoretical basis for a veiy low value of f in substitution reactions.1 Lately, several papers have appeared where values of the steric factor of the order - CH3 + CO, namely 17 kcal. Conse­ quently eq = 1 kcal, a value considerably smaller than our calculated figure 6q = 7 kcal. Then, on Figure 5, the point of coordinates q = 17 kcal and eQ = 1 kcal would lie substantially below the straight

63

7. DOUBLE BOND ADDITION Thus the radical foimed during decomposition of ditertlarybutylperoxide in the presence of dimethyleneimine follows two reaction paths [120]: 1) d e c o m p o s i t i o n w i t h of acetone and methyl.

formation

2) a metathetic reaction: with formation of tertiary butyl aiconoi. In spite of the relative ease of elimination of an H atom from the NH group, the decomposition reaction predominates. Alcohol formation proceeds at a speed comparable to that of decomposition only at high pressures of the inline. The activation energy for the decomposition of the tertiary butoxy radical into acetone and CH^ was determined by Volman [116] in a study of the polymerization of butadiene Induced by the photolysis of ditertiarybutylperoxide. The value of € was found to be 11.2+2 kcal. This is less than the figure of 1 7 + 3 kcal obtained earlier by Volman [120] in a comparison of the rates ofreactions1) and 2). The high value is not very probable since a dubious assumption was made in Its calculation, namely the equality of the activation energies of the reactions:

Let us therefore take e = 1l kcal for the activation energy of the decomposition of (CHj)jCO. The heat of reaction q has been calculated by Gray [50], who used for this purpose activation energies and heats of formation relative to some organic nitrites and nitrates:

Then the activation barrier

eQ of this reaction is given by:

Since the reverse reaction Is the addition of CH^ to a C - 0 bond:

64

I. REACTIVITY OP M0N0RADICALS

the heat of that reaction is q = 3 kcal and its activation energy is €q = 8 kcal. Por the decomposition of the ethoxy radical CgH^O » CHj + CH2O, the literature [121] gives an activation energy of circa 2 ο kcal. The energy required for abstraction of a CHj radical was calcu­ lated by Gray [50]. It is 13 kcal. Then e0=e-q=20-l3=7 kcal. Therefore, the addition of CH5 to the CO bond: CHj + CHgO »- CH3CHgO has a heat of reaction q = 13 kcal and an activation energy €Q = 7 kcal. If we trust the data given above, it appears that the activation energy for addition of CH5 to a RC = 0 bond is relatively high. This may be due to the small value of the heat of reaction for this type of addition.

kcal mole

FIGURE 5. Relailon between heat of reaction and activation energy for addition reactions This suggests the hypothesis that the ease of addition of atoms and radicals to a multiple bond is due to the relatively small energy re­ quired to open a it-bond. The energy of formation is usually substantially higher and consequently the addition of radicals to multiple bonds is often strongly exothermic. It was shown earlier, in the case of substitution reactions, how the activation energy eQ decreases when the heat of re­ action increases. The hypothesis that a relation of the type eQ = A - aq also holds for addition reactions is not in contradiction with the facts (Figure 5 and Table 2k). Let us note that the values of €Q and q for addition of CHj to a C=O bond are very imprecise and therefore the large values of eQ for small values of q are not adequate. Because of

8. DECOMPOSITION OF RADICALS

65

the lack of adequate and extensive experimental material, it is impossible to determine precisely the values of A and a. It appears, however, that they are close to the corresponding values for substitution reactions. A study of polymerization and copolymerization might reveal a relation between the activity of radicals and molecules. It would seem that the theoretical considerations presented above in connection with substitution reactions are completely applicable to this group of processes. A related theory has been put forward by Kh. S. Bagdasaryan [115]. If the resonance in the free radical is high, i.e., if the free electron is pulled appreciably Inside the particle, the free radical Is inactive and the corresponding monomer molecule-is active. A series of radicals arranged in the order of increasing activity corresponds to a series of molecules in the order of decreasing activity.1 TABLE 2b Heats of Reaction and Activation Energies for Elementary Addition Reactions Reaction

q kcal

e kcal

Reference

47

0

35 38

a.25

1 1-2 1-2

[10] [9] [10]

20

3.4

[122]

11 ?

77

13

7

3

[10]

[121] [120]

8. Decomposition of Radicals and Energy of the it-bond We may now use thermochemical data on the decomposition of alkyl radicals and their halogen derivatives in order to determine Q , the energy of formation of a it-bond between carbon atoms In olefins. Prom the C - H bond energy in ethane (Q = 98 kcal), it is easy to calculate the heat of formation of the radical which is See also Sections 11 and 12.

66

I. REACTIVITY' OF MONORADICALS kcal. Tables give the heat of formation of ethylene aHq

h

and hydrogen atoms AHH. Prom these data, we get the heat of the reaction: We have: kcal. Assuming that the energies required to break a C - H bond in ethane and in the radical are approximately equal, we calculate the value of Q by means of the relation: 98 - Q n « 38 kcal. Thus, kcal. A similar calculation using data on the decomposition of the propyl radical into propylene and an H atom, gives the value Q^ = 54 kcal. The quantity Q n may also be found by means of the reaction , using thermochemlcal data onthe decomposition of propane into methyl and ethyl radicals: q. The heats of formation of fore, the C - C bond energy in propane (tables give q = - 82 kcal [40]).

i

s

:

are known. Therek c a l

We may now calculate the heat ofreactioncorresponding to the decomposition of propyl radicals: The heat of formation of the radical CjH^ is obtained from the process: kcal. = 95 - 24.8 - 51.9 - 18 kcal. Tables [25] give: = 32 kcal. Then ILCUJ.. The heat of the reaction is equal to the difference between the C - C bond energy and the energy As sinning that the C - C bond energy in the propyl radical Is the same as in propane (83 kcal), we get:

We have therefore obtained values for Qn in three independent way3. The figure obtained is practically the same in all three cases and kcal. The radicals are active radicals in which the electrons and, in particular, the valence electron may be considered as localized. It is expected that in such radicals the C - C and C - H bond energies are practically the same as In the corresponding molecules and therefore the value of it is almost the same for all cases considered. The energy of the it-bond may also be calculated in a similar way from thermochemical data on the decomposition of halogen derivatives of alkyl radicals. The literature [123] gives the approximate heats of reaction for:

8. DECOMPOSITION OP RADICALS

67

We have q, = - 26 kcal and q 2 s - 13 koal. We may write on the other h a n d : a n d Let us assume that the free valence on the GHg group has little effect on the C - X bond energy in the CHgX group. In other words, let us assume that the C - X bond energies in these radicals are the same as in the molecules Then since we have, for the molecules, io kcal a n d 6 5 kcal Uo], we find:

The slightly lower value of Q n obtained in this fashion is due to the fact that the halogen atom in the radicals 6h2 - CHgX increases the resonance energy of the 'free' electron above the value it has In the ethyl radical. This leads to a higher value o f i n the radical than In the molecule CgH^Cig. Let us mention the paradoxical result of Steacle and others [12M following a study of hydrocarbon decomposition sensitized by mercury. Steacle found that the rate constant for the decomposition of C,H, into 8 ^ ' CHg and CgH^ was k1 = 10 exp(- 19000/RT). A pre-exponentlal factor of 108 is rather unlikely since for decomposition reactions, it is usually equal to 1013. Note that for the decomposition of n - C^H^ Into CjHg and H, Steacle obtained a normal figure for the pre-exponential factor: k 2 = 10 1 3 exp(- 4o,ooo/RT). A small value for the pre-exponential factor cannot be justified at all theoretically. It is not difficult however, to be convinced of an error In Steacle's result. Indeed, we have concluded that the reaction required an energy equal to Q = 26 kcal. This was based on thermochemlcal data and C - H bond energies in methane and propane. Thus, the activation energy of this reaction kcal. By contrast, Steacle obtained c = 19 kcal which, In last analysis, is Impossible. 1 In the new edition of 'Atomic and Free Radical Reactions' (195^), Steacle [72] made athermochemlcal calculation concerning the decomposition of the radical into and CHj. He found that the endothermiclty of this decomposition is 23 kcal in the case of and 26 kcal in the case of lso and that therefore a value of 1 9 - 2 0 kcal for the activation energy of this process as found by him and others cannot be correct. Steacle and Mandelcorn found a relatively hign vcu.ua — 7 kcal — for the activation energy of the reverse process: addition of the CHj radical to ethylene. This result came from a study of acetone photolysis in the presence of ethylene. Prior to this, all authors had usually accepted an activation energy of 2 - 3 kcal for this addition.

68

I. REACTIVITY· OP MONORADICALS

The error probably lies in the fact that Steacle measured the rate of the decomposition of Into CgHll and CH^ by means of the ethylene yield. However, photosensitized cracking is a complex process and Steacie was able to relate the rate of ethylene formation to the rate of the ele­ mentary decomposition step, only by making definite and apparently in­ correct assumptions. It is clear that the factor of 1o® does not corre­ spond to the elementary step but to an overall complex process. Note that the rate constant of Steacie for the decomposition C3Hg + H is not open to this criticism. This reaction is endothermlc by 38 kcal and the activation energy (~ kcal) is quite reason­ able, giving ε0 ~ 2 kcal, a value In agreement with that found for the addition of an H atom to propylene. It appears that for all decompositions of alkyl radicals, the pre-exponential factor is approximately 101 Jλ and the activation barrier eQ is a small quantity. 9.

Isomerlzation of Radicals

Let us now turn our attention to the third class of radical re­ actions — lsomerization. Before we do this however, let us consider the question of the most favorable angle of attack of a given bond in a mole­ cule by a radical. Very approximate quantum-mechanical calculations lead to the qualitative result that the attack of a σ-bond is most favorable when the radical approaches the bond under attack In such a manner that the atom bearing the free valence and the attacked bond are colinear. If, for Instance, the atom A attacks a diatomic molecule BC, the most favor­ able configuration for attack is that whem all three atoms are on a straight line. It is difficult to say how much the activation energy e is raised for a perpendicular attack. Rough quantum-mechanical calculations (Appendix II) indicate a very high value — twice the activation energy for linear attack. Unfortunately this question has not yet been solved ex­ perimentally. In the case of a π-bond, when the electron cloud is dis­ posed perpendicularly to the bond axis, the most favorable attack appears to be In a direction perpendicular to the bond axis. For substitution and addition reactions, the direction of attack does not matter much since it is evident that reaction occurs following collisions corresponding to the most favorable angle of attack. This only leads to the appearance of a certain steric factor expressing the fact that not all collisions are uniformly effective. In the case of free radical lsomerization however, the direction of attack is quite important.

69

9. ISOMERIZATION OF RADICALS

We call lsomerizatlon the intramolecular reactions of a free radical where a free valence 'attacks' any bond of the radical itself. Por instance: H I . C - C I I H H

H I C

H

H

H - C- L C - H I I H H

During such a reaction, an atom (In this case an H atom) mi­ grates and the free valence changes place (displacement of a reaction center). These reactions therefore differ from decomposition reactions treated above (type 2'), for instance: a)

H 3 C -CH-CH 3

H + H 2 C = CH - CH 3

H 2 C-CH 2 - CH 2

CH 2 = CH 2 + CH 3

or b)

In such decompositions, under the action of a free valence electron, one of the electrons of a σ-bond, either C-H (a) or C-C (b) shifts and forms a pair with the free valence electron, making a new «-bond. At the same time the previously existing σ-bond is ruptured. The electron shift is not accompanied by any displacement of an atom from one place to an­ other. The rupture of the σ-bond leads, however, to decomposition of the radical. In lsomerizatlon, the electron of a σ-bond also shifts under the in f l u e n c e o f a f r e e v a l e n c e e l e c t r o n . T h e σ - b o n d i n q u e s t i o n i s a C - H bond in the example given above. In this case, however, the electron shift is accompanied by an atomic displacement as a reSiiLt of which a σ-bond breaks and another one is formed, just as in a substitution reaction. The difference between substitution and lsomerizatlon is that in the first case the radical abstracts an atom from another molecule while in the second place it abstracts the atom from itself. Kinetically, there are two differences between lsomerizatlon and substitution: 1 ) lsomerizatlon is monomolecular while substitution is bimolecular. The rate of lsomerizatlon is given by:

Wi3i= IO13Cexp (- E1ZRTJJ(R)

70

I. REACTIVITY OP MONORADICALS

The rate of substitution is given by:

In these expressions (R) is the radical concentration, (M) the concentration of molecules. At atmospheric pressure, . . Then If e1 and e2 have the same value, isomerization takes place i CT times faster than substitution as was already indicated. 2) Since isomerization requires a bond 1 attack' at an angle different from zero, it is very likely that - c2 = Ae > o. At high temperatures, when the exponent (Ae/RT) is small, isomerization will be more favorable than substitution because of the favorable ratio of preexponential factors. At low temperatures, the reverse is true because of the difference in activation energies As. The question of the existence of isomerization of free radicals and of isomerization being a general property of the latter has been raised recently by the studies of Soviet scientists [1261.1 The hypothesis of Isomerization was forced upon us by some of our observations in hydrocarbon oxidation. The kinetics of formation of various products (alcohols, aldehydes, unsaturates, etc.) a3 well as the complex composition of the products could not be explained without the hypothesis of isomerization. Let \xs illustrate this by means of a very simple example. A. B. Nalbandyan sensitized photo-oxidation tem. They have shown that rates, the unique reaction

and N. V. Fok [128-130] have studied the mercury of methane, ethane and propane in a flow sysat low temperatures and sufficiently high flow product is a peroxide formed by a chain reaction:

or generally

At higher temperatures however, above 100"C, aldehydes are formed besides the peroxides. All kinetic data suggest that aldehydes are not the result of a secondary reaction but, under these conditions, are obtained In 1936, Glazebrook and Pearson [1271 studied the photolysis of di-isopropyl ketone. The radicals were removed by metallic mercury. After treatment with Hgl2 n-propylmercuriciodide was unexpectedly obtained. The authors think that the iso - C^H^ radicals formed by the decomposition Itself are isomerized by migration of an atom. Moreover, isomerization, they believe, takes place when the radical reacts with mercury since in the absence of the latter., no isomerization is observed.

9.

ISOMERIZATION OP RADICALS

71

simultaneously with the peroxides. At still higher temperatures, practically all the products consist of aldehyde. Obviously, these facts can be explained if the radical R02, at high temperatures, decomposes into aldehyde and OH before it succeeds in entering into a substitution reaction to form peroxide:

How can such a decomposition take place? merization:

The only explanation is iso-

followed by decomposition into CHgO and OH. The OH radical reacts with CH^ regenerating CH^ with formation of water (substitution reaction). Because of the higher activation energy for isomerization a3 compared to that for substitution, aldehydes are only formed as primary products at higher temperatures. Many reactions, written without reflection by different authors in connection with a reaction scheme, are, in fact, impossible without an isomerization step. Thus, for instance, the chain decomposition of ethylene oxide is propagated by the radical C£HJ0 [131]To explain the composition of reaction products, it is necessary to postulate the decomposition of C 2 H 3 0 into CHj and CO. This last reaction can hardly be visualized without a radical isomerization:

The reaction is the main step in the scheme proposed by Lossing arid Ingold [131] to aqcount for the chain pyrolysls of ethylene oxide. The reaction was investigated in a stream of helium containing very small quantities of ethylene oxide. The analysis was mass spectrometric. The conditions were: 8oo -. 1ooo°C, p H e = 15 mm mm Hg, contact time about sec. In the reaction products, besides the stable substances CH^, CO, and small quantities of C2Hg and CHgCO, there was a large amount of radicals. No methylene radicals could be observed. The quantity of radicals and methane was close to the amount of CO.

72

I. REACTIVITY OP MONORADICALS

The carbon balance .accounted for oxygen balance reached 100$.

92

to

99$ of the carbon, the

That Isomerization of alkyl radicals proceeds even at room temperature follows also, apparently, from the work of V. V. Voevodskii and R. E. Mardaleishvili [132] who measured the rate of deuterium exchange of the radicals CHgCH(CHj)a, CHgCHgCHj and oyclo-CgH,,• The reaction was conducted in a mixture of H + D g + olefin. They found that the recombination products of the radicals formed were heavily deuterated (up to 50$). This lsotopic composition cannot be explained by a sequential mechanism of the type:

Since it was shown that saturation of the molecule with deuterium D 2 is practically absent under these conditions [133I. Therefore to explain the appearance of heavily deuterated products, it is necessary to assume that the radical C Hj , once formed, is able to exchange all its H atoms for D in a succession of steps, being transformed into a saturated molecule only after exchange has taken place.1 It is quite natural that a free valence in alkyl radicals may facilitate exchange only in its immediate sphere of action. The fact that the cyclohexyl and (CH^CHCHg radicals can exchange heavily, may be used as a proof that the free valence in these radicals can move from one place to another In these reactions. In contrast with the results of V. V. Voevodskii et al., Steacie and co-workers [134] found that CH? and C2H5 radicals in the presence of molecular deuterium do not exchange heavily but that only CH D and appear or products of subsequent substitutions of these products. Tne radicals were produced by photolysis of the corresponding ketones, azo compounds and dialkylmercuiy. N. V. Fok and A. Chernycheva have shown that monodeuterated methane is obtained when CH? radicals are produced photochemically not only at high temperatures (100 - ivOO'C) and pressures (~100mmHg), the conditions 01 bteacie s work but also at room temperature and at pressures around 0.5 mm Hg., i.e., under conditions similar to those prevailing in a discharge tube. The discrepancy between the results of V. V. Voevodskii and those of Steacie is still unexplained and must be related to some peculiarities in the experimental conditions of V. V. Voevodskii et al. At any rate, the experiments show that under certain conditions, deuterium exchange of radicals takes place, even at a very fast rate. Although the absence of heavy deuteration cannot be used as a proof against the existence of radicals, the appearance of polydeuteration products in quantities not explainable by stepwise substitution, may be used as a criterion for the appearance of free valences In a given system.

9-

ISOMERIZATION OP RADICALS

73

The same authors have shown that the n-propyl radical CHgCHgCH^ exchanges more than two H atoms. This supports the possibility of an H atom transfer between the end C atoms: The fact that the ethyl radical could exchange only two H atoms shows that in that case an H atom transfer between adjacent carbon atoms is unlikely. In the n-propyl radical C^H^ as well as in C^H^, these transfers take place between two C atoms separated by another carbon and extensive deuteration is possible. Let las recall that C atoms in propane are arranged at an angle:

As a result, in the n-propyl radical, the H atom most susceptible of being attacked Is situated on an end carbon rather than on an adjacent carbon:

The 3ame is true of the isobutyl radical where the CH^ groups are more favorably attacked than the CH group. From purely steric arguments, transfer to the neighboring Pposition must take place more readily. Thus, for instance, when isomerization in p-position is hard, heavy exchange fails to take place. Thus, in the system H + (CH? )2C = C(CHJ)2 + D 2 , no heavy exchange is observed since the radical formed by H atom addition:

is such that isomerization in P-posltion cannot take place since its endothermicity is high (~ 10 - 12 kcal)• T.his displacement of free valences in alkyl radicals is completely equivalent to the process of Isomerization of oxygenated radicals, that we discussed earlier.

I. REACTIVITY OP MONORADICALS 10. Additional Remarks on Substitution Reactions To complete our survey of free radical reactions, we must add a few words on substitution reactions. So far we have considered only a type of substitution where the free radical saturates Its free valence by abstracting an atom from a given molecule. Nothing was said about abstraction of a more or less complex radical. Consider now, these reactions. For Instance:

The simplest way of visualizing these reactions is to picture the radical as attacking a a-type C - C bond from a perpendicular direction. But this scheme Is not very probable since, as we have seen, this type of attack is difficult. An end attack (reaction of type 1)):

is also made difficult by the repulsion of the H atoms of the methyl group which ought to suffer a Walden inversion if the reaction is to succeed. Such a reaction, as shown by experiments, has a rather high activation energy equal to 14 to 16 kcal. Thus, Polanyi [135] has studied the reaction of iodine with lsobutyl iodide. This involves a Walden inversion and the activation energy was found to be equal to 14 kcal:

This reaction path is also that of the ionic exchange of radioactive and non-radioactive bromine atoms with bromine substituted methane molecules. The activation energy of these reactions ranges between 16 and 26 kcal

[136]. Therefore, this type of substitution is not expected very

11.

75

POLAR FACTORS

frequently and then only at high temperatures. There are reasons to believe that reactions of type 3) and 4) involving a molecule with a double bond, proceed in two steps. The first step is the addition of the radical to the double bond of the molecule while the radical thus formed decomposes in a second step by rupture of a C-C bond. This scheme is able to represent the addition of an H atom to cyanogen: H + N

S

C - C = N

H «- N = C - C S N

H •N = C + C N .

Reactions 3 ) and 4) proceed in a similar way. Thus we may write for re­ action 3 ): CH3 CH5 + CH3 -C-C- CH, 3

ρ

0

ρ

0

3

- CHs

CH3

-A -! CI -I C H , 0

0

» CHa

A + CI - CH,

^ll

0

.

1

0

The activation energy of this reaction was determined by Blacet and Bell in a study of photolysis of diacetyl. It is equal to 5-6 kcal [137]· This interpretation of the mechanism of such reactions was first proposed by Daxvent [138] at the 1952 Discussion of the Faraday Society. 11·

Role of polar factors in Organic Reactions

In organic chemistry, there is a widely used concept according to which selectivity and rates of chemical reactions are determined by the magnitude and distribution of electronic density in molecules. A reaction takes place most readily at a site of maximum (for electrophilic reactions) or minimum (for nucleophilic reactions) electronic density. The reaction rate is fastest when the most positive and negative parts of two molecules are brought together. This essentially purely electrostatic treatment of reactivity has no theoretical basis since there are no binding reasons to believe that reaction rate is entirely determined by the magnitude of electronic densities of the reactant molecules. An experimental verification of this theory is still quite difficult because the correct distribution of charge density in most molecules is not known. Frequently electronic distribu­ tions are imagined in a rather arbitrary way. This often leads to un­ supported conclusions concerning the course of a given reaction.

76

I. REACTIVITy OF MONORADICALS

The concepts of electron theory In organic chemistry were de­ veloped from empirical reactivity rules applicable to heterolytic re­ actions: addition, elimination, substitution, etc. and especially substi­ tution in aromatic compounds. It must be emphasized that the heats of the elementary 3teps of these reactions depend on the structure of reactant and product molecules and in particular on distributions of electron densities. Therefore, the introduction in a given molecule of certain polar groups must naturally affect heats of reaction. It may be that the observed rules of reactivity may be explained by changes in activation energies resulting from changes in heats of re­ action, as was shown to be the case with radical reactions. Unfortunately, in most cases, the heats of reaction for the elementaiy steps Involving molecules and ions, are unknown. As a result, only qualitative interpretations of the experimentally observed rules can be presented. Consider, for instance, the elementary step consisting of the addition of an ion X+ (e.g., NOg+) to an aromatic molecule1 with one substltuent Y:

X If Y is an electronegative substltuent (e.g., Ci), when the C-X bond is formed, the transfer of electrons is more difficult than for the case of unsubstituted benzene. Indeed, whatever may be the nature of that bond (a σ-bond or a π-complex), it will be weaker and the energy re­ leased will be smaller as the ring electrons are more strongly attracted to the electron cloud of the electronegative substltuent. If also, as in the case of radical reactions, the activation energy increases as the heat of reaction decreases, the presence of such a polar substltuent must de­ crease the reaction rate, In agreement with observation. This treatment Is, to a certain extent, analogous to our treat­ ment of radical reactions. For instance, the unreactivlty of the benzyl radical as compared to the reactivity of the methyl radical, was explained by saying that the electron of the benzyl radical, during a reaction of T As shown by Melander [139J by means of isotoplc effects, the addition step is apparently rate determining in electrophilic substitutions, at any rate for nitration.

11. POLAR FACTORS

77

the type CgH^CHg + RH CgH^CH^ + R must be extracted from the phenyl ring In order to form the new bond. The ring plays the role of an 'electronegative' substltuent. As a result, the energy released upon bond formation is smaller and the activation energy higher. In this fashion, for both heterolytlc and radical reactions, a smaller reaction rate is explained by the endothennlc process of electron transfer required to form a new bond. This energy consideration, based on the rule of Polanyi, applied to this special case of the addition of an ion X + , coincides quantitatively with the concepts of the electron theory according to which the decrease in reaction rate is due to a diminution of electron density. However, the reasons advocated in both cases are essentially different. These arguments In favor of the role of heats of reaction in the elementary steps of heterolytlc processes, find additional support, as it seems to us, In the rule of Hammett [i4o, iM]. This rule, verified for a large number of heterolytlc reactions involving the side chain of a benzene ring substituted in ortho and para positions, is expressed by the relation:

where k^ and k Q are rate constants for the substituted and unsubstituted compound, a is a constant characterizing the substltuent only and p is a constant for a given series of reactions. If the entropy factors axe the same, this relation means that the changes in activation energies of various reactions involving the side-chain of a benzene ring are proportional to each other when given polar substituents are introduced in the ring. Indeed, since in a given series of reaction, p does not change, and the pre-exponential factors are assumed to stay the same, one has:

Since

we get:

I. REACTIVITY OF MONORADICALS

On the other hand, It may be expected that heats of reaction for the elementary step will also change in a proportional manner upon intro­ duction of polar substltuents. Indeed, elementary steps Involve breaking or making of a new bond in the side chain of the aromatic compound. For instance, the slow step of the base hydrolysis of complex esters appears to be the addition of 0H~ to the carbonyl group of the side chain: OH

The simplest example consists in the dissociation of substituted benzoic acld3: COOH

In such cases, the energy involved is determined by the type of bond and the effect on this bond of the various groups present in the molecule. Since the reacting substances considered differ only by the nature of a substituent, the difference in bond energies will be determined only by the effect of the substituents. The magnitude of this effect will be de­ termined, among other things, by the length of the side chain. The effect of the ring on a given bond decreases as this bond is further removed from the ring. In first approximation, the effect may be expressed by the product 7X where X is determined by the substituent itself and y by the position and nature of the bond in the side chain. If is the bond energy In the unsubatituted compound, the bond energy becomes -7 for a given substituent, - 7Z*2 fol> anOtller one· The difference is 7,(X, -X2). For another type of reaction, we have similarly - 72X, and - 72X£ and a difference 72(X, -X2). In this fashion the differ­ ence between heats of reaction is determined by the difference between bond energies: they are proportional and the coefficient of proportionality is This relation finds strong support In the fact that the rule of Hwnmett is also applicable to changes in equilibrium constants A similar relation obtains when logaritlsns of pre-exponential factors change lineairly with activation energies, a frequently observed state of affairs, especially in the case of liquid phase reactions.

11. PQLflH FACTORS •

79

Ki

l0g ^

where and K q are the dissociation constants of substituted and unsubstituted benzoic acids. Thus the changes in heats of reaction are pro­ portional to each other. Consequently, Introduction of polar substituents is accompanied, according to experimental observations, by proportional changes in activa­ tion energies and, apparently, by proportional changes in heats of re­ action of the elementary steps. This means that these changes are also proportional to each other. Thus ΔΕ^ = - PAq^ and = A - Pqi1 and heterolytic reactions are characterized by a relation similar to that of Polanyi for radical reactions. In this fashion, polarity affects reaction rates by its effect on heats of reaction. Qualitatively, this effect shows up quite often. It may be recalled that the polarity of a group is frequently assessed by its effect on acidity or basicity, i.e., from a purely energetic view point. However, it is known that quantitative relationships cannot always be found. The rule of Hammett breaks down already in the case of orthoderivatives. This may be due to the role of steric factors so that entropy terms are not constant any longer, following the introduction of various polar substituents. A striking illustration Is given by the ortho sub­ stituted ben2ene derivatives in which the substltuent is in the immediate vicinity of the locus of reaction [140], The rule of Hammett is not obey­ ed. Sterlc effects affect not only the entropy but also the energy of activation. Sometimes it is possible to use this effect to separate steric from polar factors. This has been carried out for certain reactions, for Instance, by Taft [142]. It turns out that changes in rate constants that are due solely to the polarity of the substltuent, obey Hammett1s rule. This means that the linear relation between actxvation energies and heats of reaction is apparently of wide applicability in chemistry. In radical reactions, the effect of polarity is, in many cases, quite different from the corresponding effect in heterolytic reactions. For Instance, the rules found in the latter cases are not obeyed by re­ actions between free radicals and aromatic compounds. In homolytic re­ actions, substitution has about the same effect in all positions of the benzene ring. The absence of a clear orienting effect of substituents is often quoted as a proof of the radical nature of the reaction. Thus, it has been shown [143] that hoaolytic substitution of H atoms by phenyl radicals in aromatic compounds (fIuor-, chlor- brom- and chlorobenzene) 1 The minus sign represents the fact that, while rate constants and equi­ librium constants change in the same direction, activation energies and heats of reaction change in the opposite direction.

80

I. REACTIVITY OP MONORADICALS

leads to a mixture of isomeric diphenyls with a predominance (~ 60%) of orthoisomer. Table 25 shows the quantities of ortho, meta and para Isomers formed as a result of the phenylation of various compounds. TABLE 25 Isomer Distribution in the Products of Phenylation (in 56) Compound

Isomer content, in % ortho 5^ 63 b9 57

meta 31 23 33 10

para 15 I1* 18 33

The isomeric distribution of Table 25 is quite different from that obtained in heterolytic processes (e.g., nitration, sulfonation, alkylation). Indeed, nitration of nitrobenzene gives 93$ of the metaisomer while phenylation of the same compound yields less metasubstltuted products than ortho and para isomers. Nitration of chlorobenzene gives a mixture of ortho and para isomers with predominance of the latter, while phenylation of the same compound gives more meta- than para- isomer. Sulfonation and alkylation also yield mainly paraisomers with negligible quantities of the metasubstltuted product. Lately, Hey and others [144] have determined the Isomer distribution following phenylation of various alkylbenzenes. As shorn in Table 26, the distribution is quite different from that obtained in the corresponding heterolytic process — nitration. TABLE 26 Isomer Distribution in Nitration and Phenylation C6H5CH3 nitration ortho meta para phenylation ortho meta para

C6H5C2H5

1soC3H7C6H5

ter .C^CgH,

ko

55 0 ^5

30 7-7 62.3

11 .8 8.7 79-5

66.5 19.2 H.3

53 28 19

31 k2

2k k9

27

27

57 3

Such a difference between heterolytic and radical processes is rather natural. In heterolytic substitution, the hetero- polar substltuent

11. POLAE FACTORS

81

may easily release or attract electrons, thus acquiring a charge. As a result, the heat of reaction may become larger, e.g., In the case of an electropositive substltuent NHg since the Ionization potential of the heteropolar atom Is substantially smaller than that of carbons. Conse­ quently, we have a process of the type:

Let us note here that such reasoning affords a qualitative explanation of some rules of orientation for the benzene ring. Indeed, substituents like NH2, OH etc. are ortho and para directing In electrophilic substitutions since a new structure of the type written above with the charge on the heteroatom can be written only if the new substltuent enters in one of these positions. If it entered in a meta position, this structure would be impossible. In a variety of radical reactions, the ,addition of radical R gives a new free radical:

6· — ό /\

R

H

If the electron is transferred to the substltuent formed may only be:

A, the type of radical

A-

The energy released In this electron transfer is either zero or at any rate much smaller than in the case of heterolytic substitution. Consequently, different substituents exert only a weak Influence on radical reactions. This conclusion is supported by the data of Szwarc [71», 1^5] who measured the C-Br bond energy in a series of substituted bromobenzene and benzylbromlde molecules, and by the work of Prltchard [119] who determined the C-C bond energy in compounds where a hydrogen atom attached to a carbon atom has been replaced by fluorine. Tables 8, 27 and 28 show that

82

I. REACTIVITY OP MONORADICALS

various polar substituents have little effect on C - Br bond energies. TABLE 27 Difference in C - Br Bond Energies in Some Substituted Benzylbromldes [145] Substltuent

Q kcal

Substltuent

Q kcal

ortho-chlor meta-chlor para-chlor meta-brom para-brom ortho-methyl

0.9 0.1

meta-methyl para-methyl meta-nitro para-nitro meta-oyano para-cyano

0.0 1.4

0.4 0.3 0.3 2.0

2.1

1.1

1.4 0.7

TABLE 28 C - C Bond Energies [119] Compound

Q kcal

In spite of what was just said, changes in activation energy may not always be due to changes in heats of reaction. It is quite likely that the distribution of electron density also affects the quantity A in the formula E = A - aq. Although no clear evidence on this point is yet available, a number of facts may be cited in support of this idea. First of all, there is no doubt that in a number of cases, the presence of a charge markedly decreases the activation energy. Thus, positively charged ions of relatively simple structure react, in vacuo, with molecules, there being no noticeable activation energy even when the heats of reaction are close to zero. According to Polanyi's rule, this would mean that A = 0 in these cases. For example, [44, 45, 146, 147]

Further, a large lowering of the activation energy is observed in reactions with halogen atoms. Here also, the quantity A is considerably reduced. Thus, atomic chlorine reacts with methane with an activation

11 .

PCEAR FACTORS

83

energy of 3·9 kcal while the CH3 and CF3 radicals react with methane with an activation energy around 10-12 kcal [148]. Moreover, in the case of CF^, the value of eQ is about 2 kcal lower than in the case of CH3. The heats of reaction, in all these cases, are practically the same· In contrast with some data of Szwarc,1 the studies of Razuvaev et al. [1U-9] have shown that polar substituents in acyl peroxides increase the stability of the radicals formed by their decomposition. For instance, the nitrobenzoyl radical formed by decomposition of para and metanitrobenzoyl peroxide is more stable with respect to elimination of CO2 than the unsubstltuted radical. The chloracetyl radical is also more stable than the unsubstltuted acetyl radical. There is also a polar effect in certain other radical processes in solution e.g., radical initiated chlorination and bromination of meta and para substituted toluene molecules and chlorination of substituted isobutane. It was found that polar substituents affect reaction rates in a definite way (positive substituents Increase the rate, negative ones de­ crease it). The ratio of rates of substituted and unsubstltuted toluenes obey Hammett's rule. For bromination ρ = 1.05, for chlorination ρ = 1.55. With the usual mechanism of bromination: k. 1 ) Br + C6H5CH3 —V CgH5CH2 + HBr k2 2) C6H5CH2 + Br2 —U- C6H5CH2Br + Br and if chains are not too short, the ratio of bromination rate constants in a mixture of substituted and unsubstltuted toluenes must be equal to k|/k1 where k| and k1 are the rate constants of step 1 ) for substituted and unsubstltuted toluenes respectively. On the other hand, the C - H bond energies which must determine the differences in heats of reaction 1) are practically independent of the introduction of polar substituents, as mentioned above. Consequently, we deal here with polar effects that affect only reaction rates, i.e., activation energies and not heats of reaction. Let us remark also that equilibrium constants of reaction 1) should not obey Hammett1S rule a3 is apparently the case for heterolytlc reactions. The only way to explain this contradiction is to assume that, in spite of re­ sults obtained by Szwarc at high temperatures, bond energies may change to the extent of a few kcal because of substituents. This would change re­ action rates of liquid phase reactions substantially (between - 60· and 100°). ^ The data of Szwarc were obtained at high temperatures where a change of 1 - 2 kcal in E is within experimental error. On the other hand, at low temperatures, such a difference plays a decisive role. The contradiction mentioned here and also later must be due to this circumstance.

84

I. REACTIVITY OP MONORADICALS

Data In support of this assumption have appeared recently. Szwarc and Wilmarth [151] have studied the decomposition of tetrazanes into free radicals. These molecules contained polar substituents. The decomposition was carried out In acetone between - 6 0 and - 10°.

It was found that the changes in equilibrium constants brought about by polar substituents obeyed Hammett's rule:

The heats of reaction that are here equal to N - N bond energies, are related linearly to a: (1 ) Since or varies from - 0.7 to + 0.9 in this work, this means that the bond energy N - N depends very strongly on polar substituents. In the same study, rates and activation energies of the dissociation of tetrazane were also measured. Using the values of Q obtained previously, the authors calculated the rates and activation energies of the reverse reaction, the dimerizatlon of radicals. They found that rate constants for this reaction also obey Hammett's rule

and that activation energies are proportional to a: (2) As we have seen before, this means that Polanyi's rule Is also obeyed. Indeed, from (1) and (2), it follows that:

85

12. POLARITY IN POLYMERIZATION E = 12.6 -

0.62q

In this fashion, the study just discussed clearly confirms the relation between the rules of Hammett and Polanyi. Moreover, while we previously used indirect arguments to establish this connection, in this case where the reaction consists in the rupture of a given bond, the connection follows directly from the experiment together with a few theoretical considerations . 12. Role of Polar Factors in Polymerization1 In the course of copolymerization of two monomers, four chain propagation steps are possible:

where A and B are monomers, ~ A and ~ B polymeric free radicals with end groups A and B respectively. Information on the composition of the copolymers thus obtained, permits us to calculate the constants of copolymerization :

1 The role of polar factors in these reactions was pointed out to us by A. N. Pravednikov.

p

Experimental determination of r1 and rg solving graphically [152] the equation

is usually accomplished by

obtained with the assumption that copolymerization can be treated as a steady state process with not too short chains. At small cbnversions,

is determined and equated to

If two runs are made with different ratios of equations with two unknowns r1 and r g .

[A] to

[B], we have two

I. REACTIVTTy OP M0N0RADICALS

86

These constants show how many times faster a monomer adds to a radical of its own kind than to a radical of the other kind. It was shown earlier in the case of radical reactions how activa­ tion energies and heats of reaction change in opposite direction and how rate constants change in the same direction as heats of reaction if steric factors stay the same. The difference between the heats of reaction of (1 ) and ( 2 ) , as well as (4) and (3) Is equal to the difference between the energies re­ quired to convert the monomers A and B into radical end groups A and B. If k^g > kj^, then kfiR > k^. Consequently, T1 will be small­ er than unity and r2 larger than unity (or the reverse If kj^g < kj^). Consider the copolymerization of two monomers A and B. The more reactive monomer is B. Then the radical B obtained from this monomer has little activity. Separately, monomers A and B polymerize relatively rapidly. Dspending on the difference in'reactivity of these monomers, there will exist various types of copolymerization. 1) If monomer B is considerably more reactive than A, 811(1 kfiB ^ kBA' 1·Θ·> is much smaller than unity, r2 is larger than unity. Addition to A of a very small quantity of B changes profoundly the rate of polymerization. Indeed, the A radical ob­ tained from monomer A is quite active and reacts rapidly with monomer B. But the radical B formed in this fashion Is relatively Inactive so that it cannot react with the relatively inactive A molecule. It will not react rapidly either with monomer B since the latter is present in very small concentration. Therefore polymerization practically stops. Higher concentrations of B will increase the rate of polymerization which will approach that of pure monomer B. Then a polymer is obtained that consists exclusively of monomer B. Monomer A does not enter appreciably into the reaction so that no copolymer is formed in this case, which is exemplified by the system vinylacetate (A) - styrene (B). The latter is considerably more active than vinylacetate. It is found experimentally [153] that T1 = 0.02 and T2 = 55· Figure 6 shows the dependence of the rate of simultaneous polymerization of vinylacetate and styrene on the composition of the mixture. As can be seen, the rate of polymerization passes through a sharp minimum at extremely small concentrations of styrene. kAA ^

kAB

2 ) It is also possible that the reactivities of both monomers are close to each other. Then T1 will be slightly smaller than unity and rg slightly larger than unity. In this case, addition of monomer B to A decreases only slightly the polymerization rate. If the concentration of B is raised, the rate Increases. Copolymer Is now formed, as in the cases of simultaneous polymerization of styrene and butadiene [15¾-]

12. POLARITY IN POLYMERIZATION 12

FIGURE 6. Rate of co-polymerization of vinylacetate (A) with styrene (B) versus composition of monomer mixture

/hour

FIGURE 7- Rate of polymerization of vinyl chloride (A) and vinylidene chloride versus composition of monomer mixture

I. REACTIVITY OF MONORADICALS

88

(r, = 0 . 7 8 and [155] (r, = 0.3 [153] (r1 = 0.2 7-

r2 = 1-39), methylmethacrylate and methyl-a-chloracrylate and rg = 1.2), vinyl chloride and vinylidene chloride and r2 = ^.5)· This last example is illustrated on Figure

For a large variety of compounds investigated, copolymerization proceeds following either type (1) or type (2), i.e., in accordance with the rule of radical activity. However, there exists a remarkably large number of cases where copolymers with alternating monomers tend to be formed. This should not happen on the basis of the simple considerations ju3t discussed. In these cases and are abnormally increased and both r^ and r2 are smaller than unity. This contradicts our ideas concerning activity of radicals since monomers A and B will react preferentially with the 'foreign' radical, irrespective of its nature. These cases are generally observed when molecules interact with polar groups exerting an opposite action. Thus [153] when vinylcyanide (A), a compound with an electron accepting group, is polymerized with styrene (B), it is found experi­ mentally that T1 = 0.03 and r2 = 0.36. In this case (see Figure 8) the rate of simultaneous polymerization goes through a maximum and is higher than that of the monomers taken separately. This phenomenon is particular­ ly striking when methylmethacrylate and styrene, containing various polar substltuents In the ring, are polymerized together [156]· The quantity Tj increases regularly as the transition is made from the most electropositive substltuents (N(CH3)2) to the most electronegative one (CN), as shown in Table 2 9 . TABLE

29

'

Simultaneous Polymerization of MethylmethaciTlate and Substituted Styrenes at 6 0 ° Substituted Styrene p-dimethylaminostyrene p-methoxystyrene p-methylstyrene styrene p-chlorostyrene ρ-bromostyrene p-cyanostyrene

r

2

r

1

0.205

0.1 1

0. 2 9

0.32

0 Λ6 0.415 0.395

0.52

1.10

0.22

1 .41

0.89

This type of copolymerization may be explained In two ways: either the activation energy has a value other than the one it should have If the heats of reaction determined the rates, or the presence of polar groups Increase abnormally the heats of reaction when alternation takes place.

12. POLARITY IN POLYMERIZATION

89

1,0

3,0 2,0

A

FIGURE 8. Rate of polymerization of vinylcyanide (A) and styrene versus composition of monomer mixture

Considering the various factors affecting the heat of reaction, our starting point was the possibility that the free electron interacts only with a molecule of a given monomer but we did not consider its inter­ action with the neighboring monomeric groups of the polymer molecules. This interaction may occur and even be significant when monomers alternate while it is absent during polymerization of the pure monomer. As a result the heat of the reactions B + A and Λ + B may be enhanced. These views must, of course, be confirmed experimentally. To do this it is necessary to find some way to determine directly the heat of the reactions A + B and B + A. Some indication on this point may be provided by a determination of the heat of simultaneous polymerization. If, for instance, the heat of simultaneous polymerization turned out to be smaller than the heats of polymerization of each monomer taken separately, one might argue in favor of a purely polar effect — that of the polar substituent on the activation energy alone. The opposite result would tend to favor the other viewpoint — an.abnormal increase in the heats of reaction. REFERENCES [11 E. T. Butler and M. Polanyi, Trans. Far. Soc., 3£, 19, (19U3). [a] M. Ladacki and M. Szwarc, J. Chem. Phys., 20, 1814, (1952); M. Szwarc and D. Williams, J. Chem. Pnys.,~"?o, 1171, (1952).

90

I. REACTIVITY OP MONORADICALS

[31 A. P. Trotman-Dickenson, J- R« Blrchard and E. W. R. Steacle, J. Chem. Phys., 19, 162, (1951). [4] A. P. Trotman-Dickenson and E. W. R. Steacle, J. Chem. Phys., 19, 169, (1951 ). [5] W. R. Trost and E. W. R. Steacle, J. Chem. Phys., 16, 361, (1948)} R. E. Rebbert and E. W. R. Steacle, J. Chem. Phys., 2J_, 1723 (1953). [6] M. R. Berlie and D. J. LeRoy, Disc. Par. Soc., _1_4, 50, (1953). t7l B. de B. Darwent and R. Roberts, Disc. Par. Soc., Jjt, 55, (1953). [8] A. D. Stepukhovich, Uspekhl Khlm., 25, 263, (1956). [9] H. W. Melville and J. C. Robb, Proc. Roy. Soc., (London) A 202, 181,

(1950).

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[70]

(195*).

~

J. L. Franklin and J. E. Lumpkin, J. Amer. Chem. Soc., 74, 1023, (1952). ~~

9a

I. REACTIVITY OP MONORADICALS

[71] E. W. R. Steacie, Atomic and Free Radical Reactions, Reinhold Publishing Corporation, New York, 195*. [72] J. B. Farmer and F. P. Losslng, Can. J. Chem., 33., 861, (1956). t73] M. Szwarc and D. Williams, Proc. Roy. Soc., (London), A 219, 353, (1953). [7*1 A. A. Zilberman-Granovskaya and E. A. Shugam, Zhurn. Flz. Khim., _1_4, 1004,

(19*0).

[75] K. Hartley, H. 0. Pritchard and H. A. Skinner, Trans. Far. Soc., itl, 25*, (1951 )• [76] T. L. Cottrell, The Strengths of Chemical Bonds, London, 1954, (Russian translation, Moscow, 1956 ). [77] G. Herzberg, Spectra and Structure of Diatomic Molecules, (Russian translation, Moscow, 1949). [78] A. G. Gaydon, Dissociation Energies and Spectra of Diatomic Molecules, (Russian translation, Moscow, 1949). [79] L. I. Avramenko and V. N. Kondratlev, Acta Physicochimica, URSS, I, 567, (1928). [80] P. Brix and G. Herzberg, j. Chem. Phys., 21_, 2240, (1953). [81] R. Edse, Third Symposium on Combustion, Baltimore, 6 1 1 , (1949). [82] J. P. Toennis and E. F. Greene, J. Chem. Phys., £6, 655, (1957). [83] H. Wise, Phys. Rev., 92, 532, (1953). [84] G. S. Pierre and J. Chipman, J. A. C. S., 76, 4787, (195*)[85] A. G. Gaydon, Dissociation Energies and Spectra of Diatomic Molecules, London, 1953 [86] S. N. Foner and R. L. Hudson, J. Chem. Phys., 23, 1364, (1955)• [87] A. P. Altshuler, J. Chem. Phys., 22, 19*7, (195*)[88] E. D. Coon, Proc. N. Dakota Acad. Sci., 7, 46, (1953). [89] K. Wieland, Helv. Chim. Acta, 24, 1285, (1941). [90] K. K. Kelley, Bull. U. S. Bur. Min., No. 383, (1935). [91] J- G. Winans and M. P. Hertz, Z. Phys., 135, *06, (1953). [92] B. G. Gowenlock, J. C. Polanyi and E. Warhurst, Proc. Roy. Soc., (London), A 219, 270, (1953 ). [933 F- P. Losslng, K. U. Ingold and I. H. S. Henderson, J. Chem. Phys., 22' 1*89> (195*). [9*] C. 0. Pritchard, H. 0. Pritchard, H. I. Shlff and A. P. TrotmanDlckenson, Chem. and Ind., 896, (1955). [95] B. S. Rabinowitch and J. P. Reed, J. Chem. Phys., 2£, 2092, (1954). [96] F. W. Kirkbride and F. G. Davidson, Nature, (London), 174, 79, (1954). [97] V. H. Dibeler, R. M. Reese and F. L. Mohler, J. Chem. Phys., 20, 751, (1952). ~ [98] V. H. Dibeler, R. M. Reese and F. L. Mohler, Jour. Res. NBS, 57, 113, (1956). ~ [99] M. G. Evans and M. Polanyi, Trans. Par. Soc., 3*, 11, (1938). [100] Kh. S. Bagdasaryan, Zhurn. Piz. Khifli., £3, 1375, (19*9). [101] N. N. Tikhomirova and V. V. Voevodskii, Dokl. Akad. Nauk U.S.S.R. 79, 993, (1951); (English translation: Canada, Nat. Res. Council, TT- 260, (1951 )• [102] G. A. Razuvaev and Yu. A. Oldekop, Zhurn. Obshch. Khim., 19, 736, (1949 )j (English translation: Canada Nat. Res. Council, TT^132, (1950)j G. A. Razuvaev and Yu. A. Oldekop, Zhurn. Obschch. Khim., 19, 1483, (1949).

12. POLARITY IN POLYMERIZATION [103]

93

G. A. Razuvaev and Yu. A. Oldekop, Zhurn. 0b3hch. Khlm., 20, 181,

(1950).

[104] G. A. Razuvaev, Yu. A. Oldekop, and N. S. Vyazankln, Zhurn. Obshch. Khlm., 2J_, 1283, (1951 )• [105] S. N. Poner and R. L. Hudson, J. Chem. Phys., 2l_, 1608, (1953). [106] Linrett, Poe, Trans. Par. Soc., 47, 1033, (1951 ) [107] C. P. Cullis, C. N. Hlnshelwood and M. F. R. Mulcahy, Proc. Roy. Soc., (London), A 196, 160,

[108] [109] [110] [111] [112]

V. J. J. H.

(1949).

V. Voevodskli, Dokl. Akad. Nauk U.S.S.R., 79, 455, (1951). L. Franklin, J. Chem. Phys., £1_, 2029, (1953 ). L. Franklin, Ind. Eng. Chem., h±, 1070, (1949). Schmltz, H. J. Schumacher and A. Jager, Zelt. Phys. Chem.,

B 51, 2 8 1 ,

(1942).

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Soc., 48, 220,

(1952).

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17 I. REACTIVITY OP MONORADICALS [135] M. B. Neiman, D. A. Kuznetsov and Yu. M. Shapovalov, Dokl. Akad. Nauk U.S.S.R., 92, 611, (1953). [136] P. E. Blacet and W. E. Bell, Disc. Par. Soc., 70, (1953). [137] B. de B. Darwent, Disc. Par. Soc., J_4, 129, (1953 ). [138] L. Melander, Nature, 163, 599, (1949)[139] L. P. Hammett, Physical Organic Chemistry, Intern. Chem. Series, McGraw Hill Book Co., New York, London, 1940. [140] H. H. Jaffe, Chem. Rev., 53, 191, (1953). [141] R. W. Taft, J. Am. Chem. Soc., 74, 2729, 3121, (1952). f142] D. H. Hey and G- H. Williams, Disc. Par. Soc., No. 14, 216, (1953). [143] D. H. Hey, B. W. Pengilly and G. H. Williams, Journ. Chem. Soc. 1463, (1956). [144] M. Szwarc, C. H. Leigh and A. H. Sehon, J. Chem. Phys., 19, 657, (1951)~ [145] V. L. Talrose, Candidate's Dissertation, M., 1952. [146] D. P. Stevenson and D. 0. Shissler, J. Chem. Phys., 23, 1353, (1955); D. 0. Shissler and D. P. Stevenson, J. Chem. Phys., IT, 926, (1956); S. H. Field, J. L. Franklin and S. W. Lamp, Journ. Am. Chem. Soc., 18, 5697, (1956). [147] G. 0. Pritchard, H. 0. Pritchard, H. I. Schiff and A. F. TrotmanDickenson, Trans. Par. Soc., 52, 849, (1956). [148] G. A. Razuvaev, Yu. A." Oldekop and V. N. Latyaeva, Zhur. Obshch. Khim., 26, 1110, (1956); G. A. Razuvaev and Yu. A. Oldekop, Zhur. Obshch."Khim., 27, 196, (1957). [149] E. C. Kooyman, R. van Helden and A. F. Bickel, Koninkl. Nederl. Akad. Wetensch., B-56, 75, (1953); R. van Helden and E. C. Kooyman, Rec. Trav. Chim. Pays-Bas, 73, 269, (1954). [150] W. K. Wilmarth and N. Schwartz, J. Am. Chem. Soc., 77, 4543, 4551, (1955). ~ [151] A. D. Abkin, Dissertation, M. 1951; T. Alfrey, J. Bohrer and GMark, Copolymerization, Moscow, 1953. [152] A. D. Abkin, Sbornik, "Voprosy khlmicheskol kinetiki, katalyza 1 reaktzionnoi sposobnosti, Akad. Nauk U.S.S.R., M. 1955, p. 338. [153] L. B. Sokolov and A. D. Abkin, Zhur. Fiz. Khim., in press. [154] T. Alfrey, J. Bohrer, H. Haas and C. J. Lewis, Journ. Polym. Scl., 5, 719, (1950). [155] C. Walling, E. Briggs, K. Wolf3tirn and F. Mayo, J. Am. Chem. Soc., TO, 1573, (1948).

CHAPTER II.

COMPETITION BETWEEN MONORADICAL REACTIONS

ι.

Competition Between Different Radical Reactions

Any analysis of kinetic data leads at once to the question con­ cerning the nature of the reaction products, i.e., the direction of re­ action. In radical reactions, the appearance of certain products is de­ termined first of all by the type of bond with which a given radical can react most easily.

For instance, the studies of Parmer [i] have shown

that oxidation of unsaturated compounds by oxygen leads to peroxides in which the

OOH

group is in α-position to the double bond.

Chain oxida­

tions of hydrocarbons at low temperatures follow the scheme: 1 ) R + O2 2) RO2 + RH

• ROO >- ROOH + R

.

The peroxide radical RO2 naturally abstracts the H atom which is most weakly bound to carbon. It was shown above that the bond dissociation energy of an CH^

H

atom in α-position is quite low (to abstract

group in propylene requires

102 - ι(A kcal

for the

CH2

77 kcal

versus

group in propylene).

95 kcal

H

from the

In propane and

Therefore, the orienta­

tion of the reaction just mentioned receives a natural explanation (if the relation between c and q Is taken into account). Aldehyde oxidation always gives peroxyaclds of the type CH, - C - OOH

and not of the type HOO - CH0 - C ^

nH

Here we also have to deal with a chain reaction following the same pattern as above and the RO2 radical will abstract an H atom from the aldehyde group and not from the

CH^

group.

Indeed, as was seen earlier, the

96

II. COMPETITION BETWEEN MONORADICAL REACTIONS

C - H bond energy In the aldehyde group (80 - 85 kcal) la smaller than in a CHj group (90 - 95 kcal). For the same reason, low temperature oxidation of paraffins usually gives peroxides of lso-structure since less energy (about 4 or 5 kcal) is required to abstract an H atom from a CHg group than from a CH^ group. As an example, let us discuss the work of K. I. Ivanov [2-4], who studied hydrocarbon photooxidation in the liquid phase. This investigation shows that oxygen attacks the tertiary C - H bonds of hydrocarbons containing a tertiaiy carbon atom. Thus, oxidation of sec-butylbenzene [3] gives a hydroperoxide as follows:

Oxidation of 2,7-dimethyloctane gives 2,7-dlmethyloctylhydroperoxide [U];

In the absence of a tertiary group, attack takes place on a secondary C - H bond, in a CHg group in a position to the benzene ring. Thus, oxidation of n-butylbenzene proceeds as follows:

It is well known that halogenated paraffins (with bromine and chlorine) decompose to give the corresponding acid and an olefin: [5]. Iodine derivatives might similarly decompose Into an olefin and hydrogen iodide. This, however, does not happen frequently. In the vast majority of cases, iodides decompose into Iodine, an olefin and a paraffin, the overall reaction being:

1. COMPETITIVE RADICAL REACTIONS

97

This process competes with that corresponding to the overall reaction: C3H7I HI + C^Hg, because of the competition between the steps: I + C3H7I I + C3H7I

HI + C3H6I - Ci1 , » I2 + C3H7 - qj

The dissociation energy of HI is only 7 0 kcal. Since we have 9 5 kcal for the C-H bond energy in propane, we make only a small error in assuming that the C-H bond energy in n-propyllodide is 90 kcal. Then the reaction I + C3H7I >- HI + C3HgI requires q, = 9 0 - 7 0 = 2 0 kcal. The competing step I + C3H7I —»- I2 + C3H7 requires only q' = 50 - 35.5 = 14.5 kcal. Indeed, the dissociation energy of I2 is 35.5 kcal and the C-I bond energy In propyl iodide is 50 kcal. We see that the second step in the decomposition of n-propyliodlde is more favorable than the first, since it requires a smaller amount of energy (5-5 kcal). It will also take place faster and this explains why iodides generally decompose with elimination of iodine and not of hydrogen iodide. In the case of bromides and chlorides, a similar comparison of the two possible reaction paths leads to the opposite conclusion. Indeed, the reaction Br + C3H7Br •- HBr + C3HgBr requires only about 5 kcal since the C-H bond energy is of the order of 9 0 kcal while the dissociation energy of HBr is equal to 85 kcal. On the other hand, the reaction Br + C3H7Br—*• C3H7 + Brg requires 17-20 kcal, the difference be­ tween the C-Br bond energy (62 - 65 kcal) and the absolute value of the dissociation energy of Br2 (45 kcal). Consequently, In the case of the bromides — and also the chlorides — the decomposition with elimina­ tion of HBr (or HCi) is many times faster than the decomposition with elimination of Br2 (or Ci2). Dlchloro-derivatives of hydrocarbons usually decompose into an olefin and HCi, according to the overall process: CH2Ci - CH2Ci

• HCi + CHCi = CH2

,

whereas the corresponding iodine derivatives yield an olefin and iodine: CH2I - CH2I

I2 + CH2 = CH2

The competition between both paths is due here to the competition between two elementary steps:

98

II. COMPETITION BETWEEN MONORADICAL REACTIONS 1 ) X + CH2X - CH2X

- X2 + δH2 - CH2X ,

2) X + CH2X - CH2X

» HX + 0ΗΧ - CH2X

With the chlorine derivatives, reaction l) is endothermic by about 23 kcal. Indeed, in this reaction, Ci2 is formed with evolution of 57 kcal and a C-Ci bond is broken, the energy of which is probably close to that of C-C^ in CHjCi (80 kcal). Thus q1 = 57 - 80 = - 23. In reaction 2 ) , HCi is formed with evolution of 102 kcal and a C - H bond is broken, the energy of which appears to be somewhat smaller than that of C-H in ethane, because of the effect of chlorine. V. V. Voevodskii has calculated for the halogen derivatives of methane (Chapter I, Section 3, Table 8), that introduction of a Ci atom in methane lowers the C-H bond energy from 101 down to 97Λ kcal. Let us assume here that the C - H bond energy is 3 to 4 kcal less than that of C - H in ethane, namely, 9k kcal. Then the exothermicity of reaction 2) is 102-91 = 11 kcal. In this fashion, the decomposition will yield only HCi. In the case of di-iododerivatives (e.g., ICH2CH2I) reaction 1 ) is endothermic (14.5 kcal) to a smaller extent than reaction 2) (endothermlc i t y : 2 0 k c a l ) . T h e r e f o r e , r e a c t i o n 1) p r o c e e d s f a s t e r t h a n r e a c t i o n 2 ) and the compounds give mainly iodine upon decomposition. Photolysis, of dlphenylmercuiy, as shown by G. A. Razuvaev and Yu. A. Ol'dekop [6], yields phenyl radicals in chloroform solution that may react in two possible ways: they abstract from the solvent either an H atom to give benzene or a Ci atom to give chlorobenzene. Reaction products show only benzene. This appears to be due to the larger evolution of heat obtained by forming benzene than by forming chlorobenzene. An approximate calculation shows a difference of about 3 to 5 kcal. At the low tempera­ tures at which these reactions are carried out, this difference means that the abstraction of an H atom will predominate. Photolysis of diphenylmercury also yields the CgH5Hg radical which, by contrast, abstracts a Ci atom from the solvent to give CgH^HgGi, a compound considerably more stable than the hypothetical CgH^HgH which would be formed by H atom abstraction. G. A. Razuvaev and G. G. Petukhov have shown that CgH5HgCi is not affected by illumination, even of long duration [71. To sum up, there are good reasons to believe that the direction of a process will be mainly determined by the heat effect of radical re­ actions (i.e., by the bond energies). This is in agreement with the re­ lation discussed earlier between heat of reaction and activation energy. In this connection, it is interesting to consider reactions of addition of compounds of the type R1H where R1 is Br or SH

2. TEMPERATURE AND PRESSURE EFFECTS

99

(hydrogen halldes, thiols) to olefins. Addition promoted by light or peroxides is opposed to the rule of Markovnlkov. The radical R1 adds to carbon atoms vith the largest number of H atoms. These reactions (due to light or peroxides) have a chain mechanism: 1 ) CH3 - CH = CH2 + Br a) C H 3 - C H - C H 2 B r + H B r

• CH3 - 0H - CH2Br + 13 kcal , «- C H 3 C H 2 C H 2 B r + f i r + 5 k c a l .

Addition of a Br atom to the CH2 group of the double bond is more exothermic by a few kcal than addition to the CH group. There are no direct data on bond energies but it is known that H addition to propy­ lene is more exothermic by 5 to 6 kcal If the product is η - C^H^ than if it Is iso - C3H17. It may be assumed that a similar difference holds for the double bond addition of bromine, accounting for the viola­ tion of Markovnikov's rule. Vlhen the reaction is carried out in solution, in the dark and in the absence of peroxides, the reaction path appears to be purely ionic and the propylene molecule is. attacked by an H+ ion with formation of a carbonium ion. The latter subsequently recombines with a Br" ion to give the normal addition product following the rule of Markovnlkov. In this case the H+ ion goes to the CH2 group to form a normal C-H bond with charge transfer to the neighboring carbon atom. This explanation evidently has a hypothetical character but seems to be a likely one. At higher temperatures, a mixture of products is always obtained corresponding to addition following and against Markovnikov's rule. 2.

Effect of Temperature and Pressure on the Competition Between Radical Reactions

The rules established JLn organic chemistry concern, for the most part, reactions taking place at low temperatures, usually not in excess of IOO0C. At these temperatures, a difference of 2 or 3 kcal In activation energies determines almost completely the direction of a chemical change. Indeed, the ratio of reaction rates In this case is W

rp- = exp(- 2000/RT) 2

provided that sterlc factors are equal. At 373°K this ratio is equal to 7.10-2 i.e., 9356 of the products will be obtained by one reaction path and only 7% by the other. If e, - ε2 = 3 kcal, the reaction will

100

II. COMPETITION BETWEEN MONORADICAL REACTIONS

proceed practically along a single path. At higher temperatures, a difference of '2 or 3 kcal In activation energies ceases to play a decisive role. At 700°K, the ratio of rate constants of two reactions with a difference of 2 kcal In activation energies, will be

i.e., 25jf of the reaction will follow one of the two paths. Let us consider the effect of temperature on the orientation of chemical reactions. There are many cases where two different bimolecular processes can take place in a system, one with a low activation energy and a low steric factor, the other with a high activation energy and a large steric factor. At low temperatures, the first process will predominate; at higher temperatures both reactions will take place concurrently; at still higher temperatures the second reaction will take over entirely. This situation is exemplified by the chlorination of unsaturated compounds. Rust and Vaughan [8] have shown that In the case of chlorination of olefins, ethylene in particular, addition reactions giving dlchloroethane begin to proceed at an appreciable rate around 235°C. In the range between 250 and 350°C, the reaction starts shifting: addition decreases in importance while substitution, giving vlnylchloride, becomes more and more significant. About 400oC, substitution predominates almost exclusively. Chlorination of olefins proceeds by a chain mechanism as follows: I. Addition

II. Substitution

Reaction 1) is exothermic to the extent of 26 kcal and Its activation barrier is about 1 or 2 kcal. As in all addition reactions, the steric factor of this process is 3mall, apparently 1 - 1 o-lf. The competing step is la), an approximately thermoneutral reaction with an appreciable activation energy. There are no direct data concerning the activation energy of this reaction but the activation energy for chlorination

101

2. TEMPERATURE AMD PRESSURE EFFECTS

of methane: is equal to 6.2 kcal [9]. For ethylene, a slightly higher value can be expected. The steric factor of reaction la) may be taken as 0.1-1. Therefore, at high temperatures, the substitution reaction becomes predominant. Competition between reactions 1 ) and la) naturally affects the ratio of chlorination products: dichloroethane and vinyl chloride. Thus at 3 08°C, addition products amount to about 65$ and there is about 20$ of vinyl chloride. At 3*3°C one already obtains ~ 5356 of vinyl chloride and only ~ 2 0$ of dichloroethane. Often there is competition between a bimolecular and a monomolecular step. Thus for Instance, in hydrocarbon oxidation, aldehydes are formed besides peroxides. A. B. Nalbandyan and N. V. Fok [10-12] have studied the mercury photosensitized oxidation of propane in a flow system In the temperature range 25 - 300°C. Mercury atoms excited by a quantum of light collide with a propane molecule and abstract an H atom to form normal and isopropyl radicals: CHj - CHg - 6h2 and CH3 - 6h - CH3. These radicals unite with oxygen to give peroxide radicals which react with the hydrocarbon, forming a peroxide and regenerating the C^H^ radical (chain):

etc.

With the isopropyl radical, we have the following chain:

Normal- and iso-propylhydroperoxides are formed in this manner. At room temperature, according to the data of A. B. Nalbandyan and N. V. Fok [11], the sole product of propane oxidation is propylhydroperoxide. These authors did not observe any aldehyde or other products. We know that abstraction of an H atom from the CHg group of propane requires 5 kcal less than abstraction from a CH^ group. Therefore, in reactions 2) and 21 ) radicals of iso structure must be obtained predominantly. Reactions 2) and 2 1 ) are exothermic and the difference

II.

102

COMPETITION BETWEEN MONORADICAL REACTIONS

between their heats of reaction ( 5 kcal) gives the difference between their activation energies: 0.25.5 = 1 -25 kcal. As a result, at room temperature, the quantity of isopropylhydroperoxide ought to exceed that of normal propyl peroxide by a factor of exp(1250/R.300) = 8. A. B. Nalbandyan and Ν. V. Pok [11] carried out the reaction at room temperature and small contact times (up to 1 ο sec) and within the limits of analytical precision, were unable to find anything but Isopropylhydroperoxide. At higher temperatures, up to IOOeC, aldehydes begin to appear as primary products. At still higher temperatures, above 3oo°C, peroxides cease to be observed and the primary products are aldehydes and their oxidation products, CO and CO2- The experiments of A. B. Nalbandyan and Ν. V. Pok [12] and of A. I. Porolkova (Inst. Chem. Phys.) who studied the oxi­ dation of propane by oxygen In the presence of ammonia and under illumina­ tion of a mercury quartz lamp, show that aldehydes are not formed by decomposition of hydroperoxides. These authors obtained kinetic curves describing the accumula­ tion of peroxides and aldehydes during the course of the reaction. Figure 28 shows the data of Poroikova at 120 and 220°. The effect of contact time on the yields of peroxide and aldehyde shows that they are formed along parallel paths. Indeed, If aldehydes were formed from peroxides, the rate of accumulation of aldehydes should be zero at the beginning of reaction and the accumulation curve would be S-shaped. Figure 9 shews that this Is not the case: the kinetic curves are characteristic of parallel processes. Undoubtedly, this is also true in thermal oxidation. When oxidation can proceed at low temperatures, peroxide is formed in

/ / X

PIGORE 9· Isopropylhydroperoxide and acetaldehyde yields versus contact time of propane oxidation in the presence of ammonia Ptotal

=

100

™

«8» C3hS : °2

:

Nh3

7: 1 : 2 χ

peroxide

.

aldehyde

curve e

1

and

at

2200

3

and 4 at

170°

2

2*/

S

I ι

ο a ο Ό Φ 03 «β

1

// U

0,2

/A

•

3/

0,1

10

20

30

Contact time, sec*

2. relatively

103

TEMPERATURE AMD PRESSURE EFFECTS

large quantities and many data indicate that aldehydes and

ketones are obtained from peroxides.

It is then also certain that de­

generate branching is associated with the decomposition of peroxides. The situation is different when hydrocarbon oxidation takes place, as usual, around

300 - 400°.

Then, as is well known, only very

small quantities of organic peroxides are formed.

In the experiments of

Stern (14, 15] who studied the oxidation of propane and propylene, it was s h o w n t h a t p e r o x i d e s a r e f o r m e d i n n e g l i g i b l e q u a n t i t i e s ( 0 . 3 - ο the original mixture) and moreover, consists of hydrogen peroxide.

70-80#

o

f

of the peroxidic product

Norrish [16] investigated the oxidation

of propane and concluded that alkylhydroperoxides are not formed at all in the reaction zone and that the peroxide detected by him was hydrogen peroxide. There are, then, good reasons to believe that the aldehydes formed do not come from peroxides.

Moreover, it was shown that the branch­

ing species in the oxidation of propane and propylene, was an aldehyde and not of

a peroxide. To prove this point, Shtern [15, 17] interrupted reaction a reacting mixture and kept the latter in a vessel in contact with

mercury at room temperature, during various periods of time (50 to

68 hours).

ture and pressure conditions of the interrupted process, reaction after an identical, small period of induction. (especially

seconds

In all cases, when the mixture was brought back to tempera­ started

It was shown that peroxides

H2O2) kept a long time in contact with mercury were destroy­

ed and that the mixture did not contain any after that treatment.

There­

fore the branching species could only be an aldehyde and additional experi­ ments showed that it was acetaldehyde in the case of the hydrocarbons studied.

More recently,

similar experiments have been conducted by Ridge

and co-workers [18] who concluded also that, in the oxidation of propy­ lene and isobutane, the peroxides are not responsible

for branching.

It may therefore be stated that peroxides formed, if at all, in small quantities in the gas phase oxidation of hydrocarbons, do not affect the kinetics of the reaction.

Thus, one may think that aldehydes

and peroxides are formed in parallel and independently.

This implies

that two reaction

The first path,

paths are open to peroxidic radicals.

discussed above, gives peroxide by reaction 2'): 2·)

C3H7OS + C3H8

• C3H7OOH + C3H7

and the second path consists in the decomposition of the peroxide radical: 3 1 ) CH3CHO + CH3O 3)

C3H7OO3") HCHO + C2H5O

21 II. COMPETITION BETWEEN MONORADICAL REACTIONS This accounts for the primary formation of aldehydes. Therefore, we have to deal with a competition between a bimolecular step 2) and a monomolecular step 3)• Reaction 3) always necessitates the isomerization of the peroxide radical. For example:

In this case, the free valence of oxygen attacks the C - C linkage. The possibility of isomerizatlons of this type was first put forward by V. Ya. Shtern and his collaborators [13, 14]. The isomeric radical

breaks down with a small activation energy into CH^CHO and CH^O (acetaldehyde and the radical of methylalcohol). The limiting step in the monomolecular decomposition of the peroxide radical is the isomerization which, as we have seen, requires an appreciable activation energy e^• The bimolecular step 2) is exothermic since the energy of the 0 - H bond formed Is always larger than that of the C - H bond broken. The activation energy of this step is apparently small and as is usual for reactions between radicals and molecules, is equal to 5 - 1 0 kcal. Reaction 2) will proceed at a rate w g = f • lo"10exp(- e2/RT)(RH)(N) where (RH) is the number of hydrocarbon molecules per cc and (N) is the number of R02 radicals per cc. We also have f ~ 0.1 and €g = 7 - 10 kcal. The rate of the monomolecular isomerization is

where The ratio of rates determines the relative yields of aldehydes and peroxides:

S. TEMPERATURE AMD PRESSURE EFFECTS

105

The difference Ae Is positive since the activation energy for Isomerlzatlon is usiially larger than that for substitution (Chapter I, Section 8). Equation (1) shows that the aldehyde yield Increases with tempera­ ture and with decreasing pressure. Experimental data [12] indicate that at room temperature only peroxide is formed while at 300°C aldehydes are preferentially obtained. Assume that £e = 1 2 . 5 kcal. This appears to be a good guess. Then at room temperature T = 3oo°K with (RH) = 3 .101 ft : aldehyde yield peroxide yield

s

3.,05 exp(-

At T = 2OO0C = 4730K with (RH) ~ aldehyde yield _ peroxide yield

1 2 5 0 0 / 6 0 0 ) = 3 .10-1*

2 .1018:

5-105 exp(_ 1 2 5 0 0 / 9 ^ 6 )

=» 0 . 8

.

Therefore, if we assume Ae = 1 2 . 5 kcal, we predict, in agreement with experimental data, that at room temperature there must be 99-97$ peroxides and 0.03$ aldehydes while at 200°C the amounts of aldehydes and per­ oxides are about equal. We then get: = €g + Δε ~ 20 kcal. Similar re­ sults have been obtained by G. B. Sergeev and V. Ya. Shtern [19] and R. A. Kalinenko and V. V. Voevodskii [20]. These authors have photooxidized pro­ pane in the presence of bromine. At low temperatures, the reaction product consisted exclusively of isopropylhydroperoxlde. Another example of competition between monomolecular and bimolecular steps with formation of different products, Is offered by the thermal reactions of propylene. As is well known, propylene does not polymerize at low tempera­ tures. This is due to the fact that the allyl radical formed from propy­ lene has very little activity as discussed earlier and is not able to add to a propylene molecule at ordinary temperatures. At higher temperatures (300 - Itoo0C) propylene starts to polymerize and at still higher tempera­ tures (700 - 800°C) cracking becomes important. The Investigations of M. S. Nemtsov et al. [21] have shown that propylene polymerization takes place in the range 330 - 43 o°C at high pressures (about 90 atm) while Pease [22] has observed the same reaction in the range 450 - 6oo°C at lower pressures (1 atm). Propylene polymerization probably proceeds via a chain mechanism as follows:

106

II. COMPETITION BETWEEN MONORADICAL REACTIONS

chain chain transfer termination

As In all polymerization processes, the propagation step has a small activation energy, about 5 - 1 0 kcal and also a very low steric factor f = 1 0 _ l t . The polymer chain is terminated by a chain transfer process of type 3). Chain length is determined by the competition between the chain propagation step 1) and the termination step. The experimental data of M. S. Nemtsov indicate a polymer chain length of about 4 to 5 segments. Szwarc [23], on the other hand, has proposed the following scheme for propylene cracking:

Another reaction path is open to the H atoms:

Szwarc has shown that, in accordance with this scheme, methane and ethylene are formed In about equal amounts while the quantity of allene is about equal to that of Hg. Since the quantity of ethylene slightly exceeds that of allene, the process preferentially goes via 2b). We see that the allyl radical appears in both polymerization and cracking. In polymerization, this radical either adds to a propylene molecule, giving CgH11 [reaction 1)] or abstracts an H atom from C^Hg to regenerate C^Hg and C^H^ [reaction 3)]. In cracking, the allyl radical decomposes into allene and an H atom. Competition between these two modes of reaction of the allyl radical will determine product distribution in cracking and polymerization at various temperatures. As was shown earlier, the ration of rates of bimolecular polymerization and monomolecular decomposition is equal to

where

e, is the activation energy for addition of the allyl radical to a

2. molecule

RH

TEMPERATURE AMD PRESSURE EFFECTS

(assume

e1 «• 10 kcal) and

the energy required to remove an proposes we put

e„ = 66 kcal

f = 10 ρ = 6

[23] at

and to

H

e2

107

Is practically equal to

atom from the allyl radical.

Szwarc

following a kinetic study of the decomposition. (RH) = 10

8

1 rJ

If

(Szwarc1s experiments were conducted

mm Hg), we obtain equal rates at a temperature of

900°C.

In fact, decomposition of propylene mainly takes place already at

8000C.

The discrepancy is due to the use of incorrect values for

e

and

f. Pressure changes may also affect competition between reactions. Thus, for Instance, as shown by V. Ya. Shtern, in the case of propane oxi­ dation, elevation of pressure increases the rate of the bimolecular forma­ tion of peroxide: RO2 + RH • ROOH + R. Changing the pressure of a CJH8 + O2 twofold.

mixture from

300

to

700

mm Hg

increased the peroxide yield

Consider still another example — the oxidation of hydrogen.

The

effect of pressure on the oxidation of hydrogen is due to the competition between two steps: 1 ) H + O2 ,

OH + 0 - 15 kcal

2 ) H + O2 + M

HO2 + M + ~ ^7 kcal .

Reaction 1) has a rather large activation energy equal to

18 kcal.

According to recent data, E1 = 15.1 - 15.4 kcal (see Chapter VIII). Therefore, at low temperatures only reaction 2) can proceed since it has no activation energy.

Following reaction 2) is the process HO2 + H2

giving hydrogen peroxide.

• H2O2 + H

As temperature is raised, the rate of reaction 1)

increases and at some temperature the sole reaction product will be (produced by the step:

OH + H2

>- H2O + H).

HgO

If, at that temperature,

the pressure is now raised, reaction 2) soon dominates again and hydrogen peroxide is once more a reaction product. The measured small activation barrier

€Q «15.4-15 =0.4

raction 1) is quite easy to understand theoretically.

of

This is due to the

fact that reaction 2) takes place practically on every triple collision without activation energy.

Since the third body

M

is required only to

remove the excess energy, then the short-lived combination of H and O2 should also proceed without any appreciable activation energy. The complex formed, if its energy is sufficient (i.e., more than 15 kcal), will de­ compose without activation barrier into either H and O2 or into OH + 0.

108

II.

COMPETITION BETWEEN MONORADICAL REACTIONS 3· Intermediate and Final PiOducts of Chain Reactions

The existence of competing reactions leads to a situation where, In a given range of temperatures and pressures, a series of radical re­ actions yield a variety of products. This multiplicity of products, for instance In oxidation, polymerization, and cracking complicates the task of unraveling the reaction mechanism. In such cases, a powerful method consists in determining quantitatively all the stable products during the course of the process. A number of important facts are established in this way and they contribute to the elucidation of the mechanism. V. Ya. Shtern et al. [13, 14-, 2 k , 2 5 ] have carried out a very detailed and precise study of the oxidation of propane and propylene. They have investigated the oxidation of mixtures of propane and oxygen [13, 14, 2k] (CjHg + O2 and iC^Hg + O2) between 200 and k6$°C in a pressure range from 280 to too mm Hg. The reaction products were aldehydes (HCH0 and CH3CHO), peroxides, methyl alcohol, ethyl alcohol, acids, propylene, ethylene, methane, hydrogen, CO, CO2 and H2O. A balance was made for the stable reaction products. The experimental data of V. Ya. Shtern et al. [13, I1»·] concerning the kinetics of formation of products during the oxidation of equimolar propane-oxygen mixtures at ρ = 282 mm Hg and T = 350°C, are shown in Table 30. Thermal oxidation of hydrocarbons, just like the photochemical process, occurs via free radicals. To observe directly these radicals during a slow oxidation is very difficult. However, all known facts leave no doubt as to the existence of radicals. Indeed, artificial introduction of free radicals in a system susceptible of oxidation (by means of light, peroxides, hexaphenylethane, ions of variable valence) brings about the oxi­ dation of hydrocarbons, aldehydes and other organic materials. In many oxi­ dation processes (in particular with hydrocarbons), one observes a very marked inhibition of the reaction by trace quantities of specific additives. The chain propagating radicals in propane oxidation, are, as was seen earlier, the normal- and lsopropyl radicals: CH^ - CH2 - 0H2 and CH^ -CH- CH^· The results of propane photo-oxidation discussed earlier, show that besides these two radicals, normal and isoperoxidlc radicals also propagate the chain as a result of the addition of oxygen to the normal and lsopropyl radicals: CH^ - CHp - CHp I - 0 - 0

and

CH, -CH- CH, 3 3 I - 0 - 0

3-

INTERMEDIATE AND PINAL PRODUCTS

109

TABLE 301 Composition of Reaction Mixtures During Oxidation of Propane by Oxygen [14] t sec after Introduction of the Mixture Ap in the Vessel mmHg 57 71.5

Concentration of Components mm Hg Peroxides Acids HCHO

CH,CH0 CH,0H C,Hfi C„H. 3 J J ° *

3

0.24

0.5

1.35

0.5

1.8

6.25

2.8

10

0.75

0.7

4

1.34

2.3

7.6

3.3

76.5

17

1.1

0-9

7.65

2.7

5-6

8

3.5

79.2

22

1.4

1.0

8.65

3.8

6.5

11

4.5

84

30 4o

1.8

1.2 1.35

11.5

4.5 4.5

8.5

1.6

15.7

13.7 15-7

5-7 7

60

0.2

1.8

11.5

4.5

25

19

8

92.5 150

11.5

TABLE 30 (cont. ) t, sec after Introduction of the Mixture Ap in the Vessel mm Hg 57 71.5

3 10

76.5

17

79.2

22

84

30

92.5 150

Concentration of Components mm Hg CH^ Hg

CO

C02

0.5

1.5

0

3.4

0.4

128

137.3

0.7

4

2

7.6

1

121

128

12

0.9

5

3.8

114

115

20

6

13

4.5

109

110

24.4

1.2

9.2

4 5 7 9

11

1.0

20

5-7 7-7 13

Peroxides Acids 0.24 O.75 1.1 1.4 1.8

40

1.6

1.35

11-8

60

0.2

1.8

12

11

35 63

C3H8

99

02

HgO 4.5

95-7 34.6

87

7 3 - 4 54.2

68

30

93

We have seen (page 105) that above 300°C, the peroxide radicals decompose practically completely before they can form peroxides by abstraction of an H atom from propane. In fact, Table 22 shows that at 38o°C only about ijt peroxide is formed in the early stages of the process while at the end only 0.1$ remains. V. Ya. Shtern et al. [13, 14] have also shown that these peroxides consist of 70 to 8056 hydrogen peroxide and only 20 to 30$ organic peroxides. Furthermore, the possible decomposition products of peroxides were not found (acetone in the case of isopropylhydroperoxide and propionaldehyde in the case of n-propylhydroperoxide). 1 The quantity of water formed was calculated from carbon, hydrogen and oxygen balances.

110

II.

COMPETITION BETWEEN MONORADICAL REACTIONS

Now let us present the theory of hydrocarbon oxidation proposed by V. Ya. Shtern et al. [13, 1M· It Is based on their data for propane and propylene. We will Introduce some corrections to this theory. Consider the possible reactions which may propagate the chain. The peroxide radical of isopropyl may decompose in only two ways: 1 ) The oxygen atom bearing the free valence attacks a C-C bond in the radical itself to give, by isotnerization, the radical CHj -CH-O- OCHj (see page 68). The latter obviously decomposes rapid­ ly into acetaldehyde CH3Cj \H and the radical of methanol H3C - 0. The first step - radical isotneriza­ tion — is almost thermoneutral since a C-O bond is made at the expense of a C-C bond. The second step - decomposition of the isomerized radical - is exothermic since a C=O double bond l~ 75 kcal) is formed at the expense of* an 0-0 bond Cto to 50 kcal). The activation energy of this second step is probably relatively small. Therefore, the overall rate will be determihed by that of the first step. Thi3 step in­ volves a considerable activation energy normally associated with isomerization and we have determined above a possible value for this activation energy: ε ~ 20 kcal. The radical of methanol formed as a result of the decomposition of the isoperoxidic radical is not likely to isomerize since the geometric arrangement of H and 0 atoms in the particle H I H - C - 0*

is quite unfavorable to isomerization. The decomposition of the radical CH3O* is also unlikely since such a process is strongly endothermlc. [A C-H bond Is broken (95 - 110 kcal) and a C=O double bond is formed (75 kcal).]1 The experiments of Rust, Seabold and Vaughan [28] who studied the decomposition of ditertlarybutylperoxide, show that, at 195°C, methyl alcohol accounts for 99*8# of the products originating in the methoxy radical. This leaves only 0.256 of other products due to its decomposition. Therefore, the CH3O" radical will react almost exclusively with RH, abstracting an H atom, forming methanol and regenerating the chain Using Steacie1s values for the decomposition of nitrites, Gray [27] cal­ culated the heat of formation of the radical CH,0 and then the endoi thermicity of the reaction: CH3O

» CHgO + H - 25 kcal.

3. INTERMEDIATE AND PINAL PRODUCTS radical n- or lso-propyl: CH3O' + C3H8

CH3OH + δ3Ηγ

This process is exothermic since a C-H bond of propane is broken ( 9 5 kcal) with simultaneous formation of an O-H bond (11ο kcal). In this fashion, this first reaction path of the isoperoxidic radical gives equal amounts of acetaldehyde and methyl alcohol, with regeneration of the start­ ing radical C3H7. 2 ) The oxygen atom bearing the free valence may attack a C-H bond of the CH3 group of the radical Itself. As a result, the following exothermic isomerization takes place:

CH3-CH-CH3 -0-0

C H 2 - C H - CH3 HO-O

The new radical breaks down Into propylene and the HOg radical. This necessitates the dissociation energy of the C-O bond while the double bond energy of C = C (~ 57 kcal) is released. Thus, the process should be endothermlc by about 2 0 kcal. It must be kept In mind, however, that the resonance of the 'free1 electron in the radical C H 2 - C H - CH3 HO- A is considerably less than that of the free electron In the HO2 radical. It is known that the latter is an Inactive species. Its activity might be 15 to 20 kcal lower than that of C3HgOOH. Consequently, the de­ composition step: C H 2 - C H - CH3 ι

HO-O

CH2 = CH - CH3 + "

HOn

is, in fact, thermoneutral or only slightly endothermlc. Therefore, the rate determining step is again the isomerization: C3H7OO C3HgOOH, which, in spite of its exothermlclty, involves a sizeable activation energy, as all isomerizations do. It would seem that this second path for iso­ merization, which gives propylene as a reaction product, requires a larger activation energy than the first path which produces acetaldehyde and methanol. Indeed, Table 30 shows less propylene than methanol in the re­ action products at 35°°C. The experimental data of V. Ya-. Shtern and V. L. Antonovskii [21*] also indicate that the ratio of products of cracking and oxidation increases with temperature.

112

II. COMPETITION BETWEEN MONORADICAL REACTIONS

The radical H02, once formed, may react in two ways: This first possibility yields and regenerates the C^Hf radical. Hydrogen peroxide, at the temperatures considered here, may decompose [29], oxidizing aldehydes at the same time. However, as a result of the endothermicity of the step: , H02, most of the time, will Interact directly with aldehydes, oxidizing them to CO or C02 and H 2 0 with elimination of molecular oxygen and regeneration of the chain radical To sum up, oxidation of the original isopropyl radical leads to the appearance of ; and oxidation products of aldehydes: CO or C02 a n d T h e r a d i c a l is regenerated and may start a new chain. Consider now the fate of the normal propyl radical CHg CHg CH^. Like the Isopropyl radical, it adds oxygen and forms the normal peroxidic radical:

The latter may also isomerize in two ways, resulting in different products: 1 ) Isomerization involving rupture of a ating a C - 0 bond:

C - C bond and cre-

The new radical breaks down to formaldehyde and the ethoxy radical:

The latter reacts with RH to form ethyl alcohol. A. P. Revzin and V. Ya. Shtern [30] have indeed indentified ethyl alcohol among the oxidation products of propane but only at lower temperatures (285°C). At 350°C, no ethyl alcohol is observed. The point is that the ethoxy radical may suffer a competing reaction, its decomposition into CHgO and a methyl radical:

3.

INTERMEDIATE A M D PINAL PRODUCTS

This reaction is tnonomolecular,

1

113

and m a y therefore proceed faster than the

bimolecular step leading to ethyl a l c o h o l . The methyl radical, once formed, m a y react with R H to give CH^ + R , or with oxygen to give the peroxide radical CH^OO w h i c h , after isomerization decomposes into CHgO and OH (see page 68). The OH radical abstracts a n H atom from p r o p a n e . W a t e r is formed and the chain radical C^H^, regenerated. 2)

The second isomerization path f o r the normal peroxide radical

is: C H 3 - CHg - C H 2

- C H 2 - C H 2 - CHg

- 0 - 0

HO - 0

This radical cannot give propylene b y direct elimination of H 0 2 decompose only into ethylene, formaldehyde and the O H radical: CH2 - C ^

- CHg

» C H 2 = CHg + CHgO + O H

but may

.

HO - 0 p The overall process is e x o t h e r m i c . According to Gray's calculations [27] the decomposition of CgH^O into CHgO and CH^, is endothermic (13 k c a l ) and its activation energy is ~ 20 k c a l . Since e = 7 kcal and f = 0.1 f o r the formation of alcohol C 2 H 5 0 + C3Hg C 2 H 5 O H + CjH^, at 285°C the ratio of rates of m o n o molecular decomposition of CgH^.0 and abstraction is equal to: 1 0 1 3 exp(- 2 0 , 0 0 0 / 1 1 1 6 ) • ( C J 5 O ) •TTC : : • = r = 3.6 1 0 ~ l u • 0 . 1 • exp(- 7 0 0 0 / 1 1 1 6 ) • ( C 2 H 5 0 ) ( M ) 1 ft with (M) = 2.7 • 10 m o l e c u l e / c c . Thus ethyl alcohol is also f o r m e d . But at 350°C, 7 = 1 2 . 3 and the decomposition of CgH.O overshadows the substitution reaction. ' If the steric f a c t o r is assumed to be 1 ( a s currently accepted b y Steacie [26] ), it is easy to show that u t 2 8 5 ° , 7 should be equal to 360 i.e., alcohol should not be d e t e c t e d , in contradiction w i t h experim e n t a l observations. 2 The scheme b y w h i c h propylene and ethylene are formed as a result of the decomposition of peroxy radicals (reaction b)): CH. t3 - C H CHj - C H - CHj 0 0-

CHj

>- CH^CHO + C ^ O -

0 - 0 \

CHj - C H - CHg 00H

CHj - C H = CHg + H 0 g

114

II.

COMPETITION B E T W E E N M O N O R A D I C A L R E A C T I O N S

To sum u p , the reactions of the n - p r o p y l radical g i v e f o r m a l d e h y d e , ethylene and w a t e r . or

C02

F o r m a l d e h y d e c a n b e oxidized f u r t h e r into

CO

and w a t e r . The reaction

m e c h a n i s m f o r p r o p a n e o x i d a t i o n is represented o n

the diagram b e l o w . Step b y step a p p l i c a t i o n of t h e theory succeeds i n p r e d i c t i n g t h e f o r m a t i o n from r a d i c a l s of the p r o d u c t s a c t u a l l y f o u n d and identified b y V . Y a . S h t e r n , at t h e e x c l u s i o n of a n y o t h e r p r o d u c t . W e also p r o v i d e h e r e a n a t u r a l e x p l a n a t i o n of t h e easy a p p e a r a n c e d u r i n g o x i d a t i o n of c r a c k i n g

2

(continued)

r e a d i l y explains w h y cracking i n the p r e s e n c e of o x y g e n (and b e s i d e s t h e o x i d a t i o n r e a c t i o n i t s e l f ) p r o c e e d s c o n s i d e r a b l y m o r e e a s i l y (at "lower t e m p e r a t u r e s ) t h a n w i t h p u r e h y d r o c a r b o n s . H o w e v e r , as shown b y S h t e r n [17J, this 3imple scheme offers some d i f f i c u l t y b e c a u s e , i n the c a s e of p r o p a n e , the r a t i o of t h e c r a c k i n g p r o d u c t s to the o x i d a t i o n p r o d u c t s increases too r a p i d l y w i t h t e m p e r a t u r e . This i n d i c a t e s that the a c t i v a t i o n e n e r g y is 13 k c a l h i g h e r f o r r e a c t i o n a ) t h a n f o r r e a c t i o n b ) . T h e n to e x p l a i n the o b s e r v e d r a t i o of c r a c k i n g p r o d u c t s ( .3 to 2.k) to oxid a t i o n p r o d u c t s , it is n e c e s s a r y to a s s u m e that t h e p r e - e x p o n e n t l a l factors of the two reactions involving t h e same r a d i c a l d i f f e r b y a f a c t o r of

10

5

6

- 10 .

This is d i f f i c u l t to e x p l a i n .

3.

INTERMEDIATE A M D PINAL PRODUCTS

115

products (ethylene, propylene). Let us note that the direct decomposition of C j H ^ into propylene and ethylene at 350°C Is unlikely since the corresponding processes are too endothermic (C^H^ C^Hg + H - 38 k c a l f and C j H ^ «- CgH^ + CHj - 2 5 k c a l ) . The appearance of C O , C 0 2 and HgO Is explained naturally b y the oxidation of aldehydes. Experimental data confirm this viewpoint. Table 30 shows that aldehydes accumulate only to a certain extent and thereafter, their amount does not change to any great extent during the course of the p r o c e s s . The quantities of C O , C 0 2 and H g O , however, steadily increase till the end of reaction. The relative amounts of products formed from the two C^H^ radicals (normal and l s o ) are close to each other, as might b e expected, since their energies of formation are almost the same (a difference of 5 k c a l ) . The corresponding difference in activation energies is equal to 1 . 2 5 k c a l . This figure means that the rate of formation of iso-CjH^ is exp(i250/RT) times greater than that of n - C ^ H ^ . At 350°C, this factor is equal to 2 . 7 . This result m a y b e checked experimentally b y the propylene to ethylene ratio. The data of Table 3 0 give f o r this ratio the value 2.2 to 2 . 3 . If, however, it Is realized that:

and also that one mole of C 3 & ^ gives three moles of CHgO, it follows that the quantities of normal- and iso-propyl radicals formed are almost identical. This is in contrast with the theoretical ratio of 2.7 calculated above, but the discrepancy is due to the omission of the steric factors which are different for the formation of both radicals. In the case of n - CjH^, the probability of abstraction of a n H atom from propane b y a free radical is larger (because there are two g r o u p s ) than in the case of the formation of iso There, a n H atom must be abstracted from the single methylenlc group of p r o p a n e . Steacie [31], while studying the reactions of the CH^ radical with various molecules, found, under similar conditions, a ratio of steric factors equal to t w o . Then, the theoretical ratio for the formation of iso- and n - C^H^ becomes ( 2 . 7 / 2 ) = 1.35. This is close to the ratio found b y comparing the concentrations of C H 2 0 and C H ^ O H . It is possible The mechanism shows that the amount of CHjCHO formed primarily (i.e., without further oxidation) is equal to the amount of m e t h a n o l . As to the amount of formaldehyde formed primarily, it has b e e n calculated b y V . Y a . Shtern [13, H J .

116

II.

COMPETITION BETWEEN M O N O R A D I C A L R E A C T I O N S (2.2 to 2 . 3 )

that the slightly h i g h e r ratio centrations of o

f

i

s

and

found b y comparing the con-

is d u e to the fact that a certain q u a n t i t y

formed from

.

A calculation of V . Y a . S h t e r n

s h o w s , h o w e v e r , that the b u l k of p r o p y l e n e comes from

iso -

The rapid a c c u m u l a t i o n of cracking products observed [14, 24] at h i g h e r temperatures u p to

420°C

n -

around

In the summarizing scheme a b o v e , these p o s s i b i l i t i e s are

4oo°C.

and

m a y b e explained b y the d i r e c t d e -

composition of

represented b y dotted a r r o w s /

iso -

into p r o p y l e n e and e t h y l e n e

This is confirmed b y w o r k of Steacie et a l .

[ 3 2 ] w h o have s h o w n , i n a study of the m e r c u r y p h o t o s e n s i t i z e d of p r o p a n e , that at

4oo°C,

decomposition

appreciable quantities of m e t h a n e a r e already

formed. The m e c h a n i s m p r o p o s e d above m a y b e checked b y comparing t h e amount of w a t e r calculated I n two ways: as formed b y the various p o s t u l a t e d reactions a n d as obtained from the C , 0 a n d H b a l a n c e s . I n this calcul a t i o n , it is n e c e s s a r y to assume that aldehydes n o t only oxidize b u t a l s o p a r t i a l l y d e c o m p o s e as s h o w n b y t h e a p p e a r a n c e of Hg i n the r e a c t i o n p r o d u c t s Table 31 contains the results of this c a l c u l a t i o n f o r the course of propane o x i d a t i o n at C

H

3 8

:

°2

=

1

:

282 m m H g

and

350°C.

T h e Initial m i x t u r e w a s :

''

TABLE 31 M e a s u r e d and Calculated Concentrations of W a t e r f o r O x i d a t i o n of P r o p a n e w i t h Oxygen C o n c e n t r a t i o n of Ap mm Hg

t sec

Prom Balance

HgO, mm Hg

Prom Mechanism

10

71.5

12

17.2

17

76.5

20

23.4

22

79.2

24.4

29.4

30

84

34.6

39.4

40

92.5

60

150

54.2

57.2

93

97

As c a n b e s e e n , the agreement is s a t i s f a c t o r y . The elucidation of o x i d a t i o n m e c h a n i s m s n e c e s s i t a t e s a k n o w l e d g e of

w,

and

w2,

the rates of f o r m a t i o n and d i s a p p e a r a n c e of the m a i n

intermediate p r o d u c t s .

E x p e r i m e n t u s u a l l y gives the d i f f e r e n c e

3. INTERMEDIATE AND PINAL PRODUCTS

117

i.e., the rate of accumulation of intermediates. But it is of exceptional interest to know W1, the rate of formation of say, aldehydes, as If they did not react any further. In their study of the accumulation of the main products of propy­ lene oxidation, Shtern and Polyak [15] were able to establish a balance be­ tween the consumption of fresh reactants and the accumulation of Inter­ mediate and end products. This then permitted the authors to propose a radical chain mechanism explaining product distribution both qualitatively and quantitatively. Prom this mechanism, the quantities of aldehydes formed could be calculated (assuming that they do not react any further. Analytical data give the amount of aldehyde accumulated during reaction. The difference between the two figures gives the quantity of aldehyde re­ acted to form CO, CO2 and H2O. The mechanism of propylene oxidation was then verified by com­ paring calculated quantities of CO, CO2 and HgO to the analytical values. The two figures were found to be almost the same. Until recently, however, no experimental method was available to determine the rates of formation and disappearance of intermediate products. Only recently, at the Institute of Chemical Physics, has a kinetic method been developed by M. B. Neiman: it uses tracers to study the mechanisms of complex reactions and provides an experimental determination of the rate W1 of formation and of the rate W2 of disappearance of intermediate products. If we call (B) the concentration of Intermediate at any time, we have: «>

iSi-",

In order to determine W1 and W2 separately, one more equation is need­ ed and it is provided by the kinetic method. If a certain small quantity of tagged product B* is Introduced in the system, the relative activity P equal to (B*)/(B) will change with time: dp . d(B*)/(B) _ 1 . d(B*) at cT TET

P . d(B) TF7 "3ϊ"

Since, in the present case, the product B* is not formed but only con* sumed, W1 = 0 and W2 = Pwg. Then:

at=

pw i 757

118

II.

COMPETITION HEiTWEEN MONORADICAL REACTIONS

From this equation, it is possible now to determine the rate of formation of the intermediate product: w, = - (B)

.

The net quantity (B1) of intermediate product formed at time equal to t

(4)

(B') =

J ο

t

is

o (B) d log β .

P

W1 dt

=J β

By means of the kinetic method, this quantity can be calcinated directly from experimental data. M. B. Neiman et al. [34] have used this kinetic method to study the mechanism of oxidation of propylene at 315° (C^Hg + O2, Ptota^ = 243 mm Hg)· To the reacting mixture was added ~ iji (2 - 3 mm Hg) of tagged acetaldehyde. The specific activity β and the concentration of acetaldehyde (B) were measured as a function of time and the authors calculated by means of equation (4) the quantity (Bt) of acetaldehyde formed during the reaction. The difference between this value and the amount (B) of acetaldehyde accumulated during the process, gives the quantity that has disappeared to form end products. Figure 1 ο shows curve (1 ) calculated by means of equation (4) and curve (2) drawn through the experimental values on the accumulation of CH^CHO. Figure 10 also shows that curve (1) is in good agreement with curve (3) calculated by V. Ya. Shtern and S. S. Polysik on the basis of their radical chain mechanism of propylene oxidation. Besides the rates of formation and disappearance of intermediate products, the elucidation of reaction mechanisms also requires a knowledge of the order in which the various products appear. This problem can also be solved by the kinetic method of M. B. Neiman that determines the rate of formation of any product B from its various precursors. It is then necessary to tag one of the precursors, say A and to follow with time the change in concentration of B and of the specific activity α and β of the products A and B. Let us call W1 the rate of formation of that from other precursors. Then: ^= v,

+

w„

.

The rate of increase of the specific activity is:

B from

A

and

wg

3.

INTERMEDIATE A M D P I N A L P R O D U C T S

p mm Hg

F I G U R E 10. P r o p y l e n e Oxidation: f o r m a t i o n (curve 1 ) a n d a c c u m u l a t i o n (curve 2 ) of a c e t a l d e h y d e . C u r v e 3 i3 calculated b y S h t e r n and P o l y a k (15)

F I G U R E 11.

S p e c i f i c A c t i v i t y of f o r m a l d e h y d e (1) a n d c a r b o n m o n o x i d e (2) d u r i n g m e t h a n e o x i d a t i o n

119

6

II.

COMPETITION B E T W E E N M O N O R A D I C A L R E A C T I O N S

F I G U R E 12. Specific a c t i v i t y of carbon m o n o x i d e (1 ) and c a r b o n dioxide (2) d u r i n g b u t a n e o x i d a t i o n

3· INTERMEDIATE AND FINAL PRODUCTS

121

If, at the start of reaction, a small quantity of untagged product B is added at t = ο, ρ = o. In this case, β first increases to a maximum value and then decreases since a decreases. The maximum value of β is determined by the condition άβ = 0

i.e.

Γ - β w—-—-J 1 + W2~l = 0 j^a

Hence:

Also, if B is formed only from A, W2 = 0. Then ο/β = 1 and a = β. The curves of activity must then cross each other. But if W2 / 0, they do not cross and from the value of α/β at the maximum, the rates W1 and W2 can be obtained. This method was used by M. B. Neiman, A. B- Nalbandyan et al. [35] to decide whether, in the course of methane oxidation, carbon monoxide is formed only from HCHO or also from some other substances. In order to prove this point, to a mixture of 33# CHk, 6656 air and 0.156 NO, they 1L added 0.0756 tagged C H2O and 0.556 CO. The reaction was carried out at 670°C. During the course of the process the specific activities α and β of the products HCHO and CO were measured. At the point of maximum value of β (see Figure 11 ), a = β. This may be the case if W2 = 0. Therefore, carbon monoxide is produced only from HCHO by the overall scheme CHlf • CHCHO » CO. The kinetic method was also used by M. B. Nelman and A. 7. Lukovnikov to prove rigorously that, in low temperature oxidation of hydro­ carbons, CO2 is formed mainly not from CO but from other compounds, especially, it seems, by decomposition of radicals. These authors oxidized butane in the presence of small quantities of C 1 40 and CO2 and measured the activities a and β of these products as a function of time (see Figure 12). As can be seen Pmax « a. Using equation (6), they concluded that about k0. ~ [43] I. V. Berezin, E. T. Denlsov and N. M. Bnanuel1, Symposium or, "Chemical kinetics, catalysis and reactivity", p. 273, Academy of Sciences of U.S.S.R., Moscow, 1955. [44] British Patent 633354, 12/12, 1949. [45] A. Parkas and E. Passaglia, J. Amer. Chem. Soc., 72, 3333, (1950). [46] I. V. Berezin, B. 1G. Dzantiev, N. F. Kazanskaya, L. I. Sinochkina and N. M. Bnanuel , Zhurn. Piz. Khim., No. 1957[47] N. Brown, M. J. Hartig, M. J. Roedel, A. W. Anderson and C. E. Schweitzer, J. Amer. Chem. Soc., _77, 1756, (1955). [48] W. Pritzkow and K. A. Muller, Berichte 2321, (1956). [49] E. J. Gasson, E. G. E. Hawkins, A. P. Mlllidge and D. C. Qjjlnn, J. Chem. Soc., 2798, (1950). [50] B. I. Makolets, Thesis for Diploma, Moscow State University, 1953. [51] P. 0. Rice and K. K. Rice, The aliphatic free radicals, Baltimore, 1935» (Russian translation, Moscow, 1937). [52] A. V. Prost, Uspekhi Khim., 8, 956, (1939); A. I. Dintses, Uspekhi Khim., 3, 936, (1934). [53] V. V. Voevodskii, Symposium on "Chemical kinetics, catalysis and reactivity", Moscow, Academy of Sciences, 1955. [54] L. Shmerllng, Symposium on "Catalysis in organic chemistry", p. 123, Moscow, 1953. [55] Yu. A. Arbuzov, Scientific reports of Moscow State University Organic Chemistry Section, 89, 7, (1945). [56] N. N. Semenov, Chain reactions (In Russian), Leningrad, 1934. Translated and revised as "Chemical kinetics and chain reactions", Oxford, 1935. [57] V. N. Ipatieff, Berichte, 3j>, 1047, (1902). [58] C. N. Hinshelwood and W. K. Hutchinson, Proc. Roy. Soc., (London), A 111, 380, (1926). [59] P. D. Zemany and M. Burton, J. Phys. Coll. Chem., 942, (I95I). [60] L. A. Wall and W. J. Moore, J. Phys. Coll. Chem., 965, (1951). [61] H. W. Melville, J. C. Robb and R. C. Trutton, Disc. Par. Soc., 14 150, (1953). —

1W

II.

COMPETITION BETWEEN MONORADICAL REACTIONS

[62]

M. S. Kharasch, E. V. Jansen and W. H. Urry, J. Amer. Chem. Soc.,

[63]

R. M. Joyce, W. E. Hanford and J. Harmon, J. Amer. Chetn. Soc., 70,

6g, 1100,

(19*7).

2529, (19*8).

~

[6*] E. C. Kooyman and E. Farehhorst, Rec. Trav. Chlm. des Pays-Bas, jo, 267, -(1951 )• [65] Kh. S. Bagdasaryan and R. I. Mllyutlnskaya, Zhurn. Flz. Khlm., 27, *20, (1953). ~~ [66] W. A. Waters, Trans. Par. Soc., 37, 770, (19*1)} M. T. Jaqulss and M. Szvarc, Disc. Par. Soc., No. ]T, 2k6, (1953). [67] R. I. Mllyutlnskaya and Kh. S. Bagdasaryan, Zhur. Piz. Khim., In print. [68] R. I. Mllyutlnskaya and Kh. S. Bagdasaryan, Zhur. Flz. Khlm., In print.

CHAPTER III: REACTIONS OP DXEiADICALS

1.

Transition of Atoms to a Valence-active State

Typical divalent atoms (0, S, Mg, Ca, Cd) have, in the free state, an even number of electrons. The ground state of 0, S, Se is a triplet (3p) i.e., these atoms have two unpaired electrons. They are very active chemically and react with various valence-saturated compounds as readily as typical univalent atoms. The activation energy e of the reaction 0 + H2 - OH + H is equal to about 6 kcal, while ε = 7 kcal for the reaction H* + H2 —• H*H + H. Consequently, divalent atoms may be pictured as diradicals

ι ι i

— O f S —, S e —

.

An energy AR must be spent to promote atoms 0, S, Se Into a 'valenceactive' state. We use this term to designate the state of the atomic electron cloud that would hold if its configuration in a molecule stayed the same after all bonds have been broken. In reality, such a state may not exist because it is not stable and passes over at once to the normal state characteristic of the isolated atom. It can be Imagined that bond formation takes place in two steps: first the atom is promoted to the valence-active state by means of the ex­ penditure of some energy and then the valences prepared In this way are saturated with evolution of energy. Such a separation of the energy of chemical bonds into two terms is useful. For the 0 atom, ΔΚ - 1¾- kcal. This is shown by the fact that union of an 0 atom with an H atom to form OH releases 103 kcal while addition of a second H atom to form water releases 117 kcal. Naturally, both H atoms in H2O are bound in the same way. Therefore, the difference 117 - 103 = 14 kcal corresponds to the promotion of the 0 atom to its valence active state. In reality, the energy difference between the first and second O-H bonds is still larger since the formation of the second bond necessitates the expenditure of some energy to deform the valence angle between two O-H linkages from its normal value of 90° to IOif090', because of repulsion between 1Ό

III. REACTIONS OP DIRADICALS

146 the H atoms.

It would seem natural to assume that the large chemical activity of oxygen, sulfur and selenium atoms is due to the presence of two unpair­ ed electrons and that the latter completely define diradical behavior. But this is not so. Divalent metal atoms of the second group have a 1S singlet ground state, i.e., their valence electrons are paired. For this reason, the quantity AR is larger here than for atoms of the sixth group. Thus, for instance, addition of a first Ci atom to an Hg atom releases 23 kcal, as compared to 80 kcal for the second chlorine. The promotion energy to the valence-active state is then about 55 kcal. The successive addition of two chlorine atoms to Cd and Zn corresponds, following V. N.Kondrat'ev [1], to energy releases of 51 and 104 kcal, and 50.5 and 105 kcal respectively. The value of AR is now about 50 kcal. Never­ theless, all these atoms exhibit true diradical behavior since the activa­ tion barriers for reactions of these atoms with saturated molecules are as low as in the case of typical single and di-radicals, which have unpaired electrons. Por instance, when zinc vapor reacts with Ci2 : Zn + Cig ZnCi + Ci, the activation energy is 8 kcal. With cadmium: Cd + Cig «CdCi + Ci, the activation energy is 1 2 . 5 kcal. The heat of reaction is 6 kcal, the difference between the bond energies of Cd-Ci (51 kcal) and Ci -Ci (57 kcal). The activation energy € of an endothermlc process is q + eQ where q is the endothermlcity and eQ the activa­ tion barrier. The value of the latter in this case is thus very small — only 6.5 kcal. The quantity AR, therefore, only decreases the heat of reaction and in this manner decreases the activity of atoms of the second group. But the activation barrier itself has a small value whether the valence electrons are paired or not. Therefore, from a chemical viewpoint, second group atoms behave like diradlcals, but d!radicals of small activity (similar to the elngle radicals HO2, CH2 = CH - (1¾

etc.

).

The carbon atom in its ground triplet state is divalent; the energy required to bring it to its quadrivalent state is, however, very large, as shown by calculations. 2.

Reactivity of

O2, S2, Se2 Molecules

Consider now the molecules O2, S2, Se2- Their ground state is a triplet i.e., they have two unpaired electrons, they are para­ magnetic and it would seem that they might exhibit diradical behavior: -0-0-, -S-S -, -Se-Se-. But the O2 molecule, for instance,

2. REACTIVITy OP O2, S2, Se2 doe3 not have dlradical properties, neither from a kinetic nor from a thermochemical standpoint. The energy of the bond formed by the two impair­ ed electrons in the O2 molecule is not smaller but larger than that form­ ed by two electrons with antiparallel spins. This can be appreciated from the ordinary 0-0 bond energy in peroxides (H2O2) which is in the vicinity of 50 kcal, whereas the dissociation energy of 0=0 amounts to 118 kcal. Chemically a molecule of oxygen is very inert. The apparent ease with which O2 enters into some reactions is not explained by the activity of the molecule itself but by the properties of the 0 atom and • the R - 0 - 0 - radical that favor the propagation of chain reactions. This is now quite clear as a result of studies made at the Institute of Chemical Physics. When conditions are not favorable to the propagation of a chain reaction, the O2 molecule is Indeed extremely inert. It may re­ main in contact during several days with vapors of phosphorus, phosphine, sllane, without any trace of reaction, while, If conditions are favorable to the propagation of a chain process, reaction of O2 with these sub­ stances is violent even at room temperature. Oxygen reacts homogeneously with hydrogen only by a chain mech­ anism. There is practically no direct reaction between H2 and O2· The activation energy of the direct reaction is at least 50 kcal, a typical figure for reactions between saturated molecules and an order of magnitude larger than for reactions between radicals and saturated molecules. The chemical, inertness of the oxygen molecule is also illustrated by the absence of dlmerization in oxygen. (Some association may occur in the liquid at low temperatures.) Sulfur and selenium molecules resemble diradicals more, but they are not particularly active. The energy of a single S-S bond is approximately 52 kcal, that of the S=S bonds, 83 kcal [2]. This is evidence that the second bond formed by electrons with parallel spins is appreciably weaker here. Sulfur, as is well known, can polymerize. In sulfur vapor, Sg and Sg molecules are formed, apparently with a cyclic structure. The S atoms in such complexes are held by ordinary bonds. Finally, sulfur may form long polymer chains (polysulfides). There are no direct data concerning the activation energy for reaction of S2 with other molecules. There is some indication however that such reactions pro­ ceed readily when they are exothermic. The P2 molecule has a singlet 1S ground state and a large dissociation energy — about 116 kcal. In spite of this it exhibits dl­ radical behavior just like sulfur. Thus, P2 molecules form Plf dimers (a tetrahedral structure with ordinary bonds) with evolution of 30 kcal. The P„ molecule behaves like an active molecule and there are more C. reasons to regard it as a dlradical than O2, in spite of the fact that

148

III. REACTIONS OP DIRADICALS

the electrons of P2 are paired and those of

O2 are not.

To S-UIii up, with molecules, also, the physical concept of a diradlcal (triplet state, paramagnetism) and the chemical concept (absence of activation harriers, tendency to dimerize, weakness of the second bond or its non-existence) do not coincide in a number of cases. 3· The Divalent State of Carbon As was already pointed out, the carbon atom is a diradical. After saturation of two valences of the C atom by hydrogen atoms, the CH2 particle is formed which, chemically speaking, is also a diradical. The chemical properties of CH2 will be described in detail in a subsequent section. It appears that after saturation of the two valences of the CEirbon atom, two new ones which did not exist before malce their appearance. As a result the CH2 particle is a typical diradical. Among the compounds of divalent carbon with various atoms, only CO does not exhibit the chemical properties of a diradical. The reason for this is not entirely clear but the strengthening of the bond in the CO molecule (256 kcal) as compared to the bond strength of the carbonyl group (150 - 160 kcal) in a variety of organic compounds, shows that the 2 s-electrons of carbon take part in bond formation in the CO molecule.1 In this fashion, the CO molecules represent an exception, even more marked than NO among the single radicals. To transform a CO molecule into a typical diradical, i.e., a valence-active state, requires a large expenditure of energy, 60 to 80 kcal. This is clear from the fact that addition of a chlorine atom to CO to form COCi releases only a few kcal, whereas addition of a second Ci atom to make phosgene releases about 80 kcal. Since the energy of both C-Ci bonds in phosgene Is obviously the same, this shows that the difference between the heats of re­ action for successive addition of two chlorine atoms to CO, represents the energy required to promote CO to a valence-active state. The CS particle which, according to the data of V.'N. Kondrat'ev [4], has a life-time of about 10 sec at room temperature, differs marked­ ly from CO. In particular, it much resembles a diradical as illustrated by its tendency to polymerize. Below 100°C, CS does not react appreciably with O2 but this reaction is observed above ioo°C. It appears that a CS particle is a relatively Inactive radical occupying a position inter­ mediate between those of CH2 and CO. We know little about other compounds of divalent carbon but it is possible to suppose that the majority behaves like diradlcals. This is confirmed by a calculation performed in

1953

[3].

4. COMPLEX DIRADICALS The examples just quoted make it clear that the state of the second electron pair of carbon depends strongly on the nature of the atoms bound to the first two valences. After combination of the two diradicals -S- and - 0 - to form the SO particle, the second electron pair of sulfur, previously bound within the atom, becomes accessible. Indeed, SO exhibits the typical chemical properties of a dlradical and its ground electronic state is a triplet. As shown by Ν· M. Emanuel' [5] SO dimerizes into SgO2 at low temperatures. Also, SO reacts easily with various molecules. With O2 for instance, it gives SO2 and an 0 atom. The phenomenon of liberation of another electron pair after saturation of the first two valences, is sometimes observed among the ele­ ments of the third, fourth and fifth groups of the periodic table. A typical, chemically active, dlradical is then formed. The phenomenon is rarely observed however. In the vast majority of cases, addition of atoms or radicals yields valence-saturated molecules that react with the char­ acteristically high activation energies of such species. 4. Plradlcal3 of Complex Structure Diradicals may be obtained from polyatomic molecules, e.g., hydrocarbons, by abstraction of two H atoms. In practice, this is a very difficult way to form a dlradical. Sometimes, however, diradicals are formed during the course of a reaction, for Instance polymerization of pure olefins and diene addition. In a few Instances it is possible to synthetize complex molecules of special structure, which are diradicals in their ground state. For ex­ ample, 2, 6, 2', 61-tetrachlordiphenyldimethane Ci

Ct

-O-O-

6HS

Ci

may not have a plane configuration because of the chlorine atoms on the benzene rings. This circumstance prevents the formation of a qulnonold structure and the molecule possesses two free valences. Solutions of this compound are strongly colored and react extraordinarily easily with air. In solution, this compound is paramagnetic and its paramagnetism Increases with temperature although the opposite should happen following Curie's law. Paramagnetic susceptibility increases with dilution. These facts Indicate an equilibrium in solution between paramagnetic monomers and diamagnetic dimers or polymers. Therefore, this compound has all the chemical and physical attributes of a dlradical.

150

III. REACTIONS OP DXRADICAL3

Synthesis of similar compounds, In the absence of oxygen and In the dark, Is quite complex and tricky; for this reason relatively few com­ pounds with clear dlradlcal behavior are described in the literature, and we will limit ourselves to the illustration just described. Much more often, one has to deal with molecules with a singlet ground state. However, thermal activation may promote them to an excited dlradlcal state. Obviously, the concentration of dlradlcals obtained In this way will depend essentially on the magnitude of the energy gap be­ tween singlet ground state and triplet excited state. When the energy difference is small, the possibility of participation of dlradlcals in chemical changes must be considered. For instance, in the case of the Chlchibabln hydrocarbon:

(Wa0 =- fWa O - O ... Molecules of p-qulnodimethane

H2c=- Ci2 + M. The equilibrium concentration is determined by the equilibrium constant γ

(Ci)2 (Ci2)

- ———

·

As we have just seen, chlorine atoms are captured by the wall at a certain rate. According to the principle of detailed balance, each act by which a chlorine atom disappears at the wall must be compensated by the reverse process: formation of a Ci atom. Therefore, processes contribute to the establishment of equilibrium In the volume: Ci2 + M 1" Ci + Ci + M Ik0(Ci2) = k,(Ci )2]

1. and at the wall.

RADICAL WALL REACTIONS

213

The following steps take place at the wall:

1) CJ + wall »- [Ci] 2) Ci + [CJ ] —- Ci2 the capture of a the walls:

Ci

atom and Its reverse, the formation of free atoms on

3) Ci2 [Ci] + Ci it) [Ci] —.- Ci . Thermodynamics tells us that the equilibrium concentration of Ci atoms established by volume and surface processes of dissociation and form­ ation of Ci2 molecules and recombination of Ci atoms, corresponds to a fixed value of the degree of dissociation i.e., the presence of the wall cannot shift the equilibrium between Ci£ and Ci (principle of detailed balance). If a concentration of Ci atoms exceeding the equilibrium value is artificially created (for Instance, by illumination), the rate of atom production at the wall stays the same but the rates of heterogeneous ter­ mination 1 )and 2 ) which are proportional to the volume concentration of Ci, increase so that the excess of Ci atoms above the equilibrium amount decreases. On the contrary, let us add to the gas phase a substance of the type of NCi3 that quickly captures Ci atoms and thereby lowers their concentration below the equilibrium value. Then, since the rates of processes 1 )and 2) decrease, the relative importance of the role of the wall in the generation of chlorine atoms increases, so that the volume con­ centration of Ci atoms is now higher than it would be in the absence of the wall. Consequently, in both cases, the wall tends to re-establish the equilibrium atomic concentration in the volume. It must be noted that the rates of the surface processes destroy­ ing or producing Ci atoms are larger than in the volume. In the latter case, two Ci atoms are produced simultaneously and the energy required is that of dissociation, namely 57 kcal. In the former case, however, at the surface, there occur two events: first a Ci atom is produced by process 3 ) and the heat required is 57 - Q where Q is the heat of ad­ sorption; then the second Ci atom evaporates and the heat (¾ required is less than 57 kcal. The occurrence of two steps, each requiring less energy than the direct process of dissociation, considerably Increases the rate of the overall process taking place at the surface. The wall does not shift the thermodynamic equilibrium Ci2 , ' Ci + Ci but accelerates the approach to equilibrium. It is easily calculated that at 6oo°K, it takes 1.5 mln to reach 50# of the equilibrium con­ centration of Ci atoms by homogeneous processes. During the H2 + Ci2 reaction, the concentration of Ci atoms at equilibrium Is the 3ame as in

214

VI. WAH INITIATION AND TERMINATION

pure Ci2, without H2- The rate of formation of HCi Is w = It(Ci)(H2). In fact, experimental data for the H2 + Ci2 reaction show that the equi­ librium Ci2 , ' 2 Ci and therefore the steady-state rate of formation of HCi is reached in less than 1 sec. This can be explained only by the effect of the walls. Prom the well established fact of wall capture of Ci atoms, we conclude that the reverse process, wall production of Ci atoms also takes place. As we have seen, when the equilibrium Ci2 - ' 2 Ci is reached, the presence of the wall does not change the concentration of atomic chlorine and therefore does not change the steady-state rate of H2 + Ci2 in pure gases. If the mixture of H2 and Ci2 contains impurities that start or terminate chains easily, the walls will affect the steady-state reaction rate by diminishing the effect of Initiators and inhibitors be­ cause they tend to re-establish the homogeneous equilibrium concentration of chlorine atoms. A direct proof of the wall generation of Ci atoms was given by A. M. Markevich [4]. He measured the temperature distribution in a vessel where the H2 + Ci2 reaction took place (T = 260 - 36o°C). In pure gases, the temperature distribution corresponded to that calculated from a uniform evolution of heat throughout the reactor with heat conduc­ tion of the heat through the walls. After addition to the H2 + Ci2 mix­ ture of an Inhibitor (oxygen), the heat evolution took place entirely in a narrow zone near the reactor walls. The width of this zone decreased as more oxygen was added to the mixture. This phenomenon is due to the fact that the elementary steps that propagate the chain were confined to the vicinity of the wall because Ci atoms were for the most part generated at the walls. The chains diffusing into the interior were terminated close to the walls by oxygen and consequently the reaction was unable to sustain itself in the central part of the reactor. In 1 9 5 4 - 1 9 5 5 , in the kinetics Laboratory of Moscow State Uni­ versity, Chalkin [51 obtained new results by means of the method of differ­ ential calorimetry. They confirm the conclusion relative to the surface generation and termination of chains In the thermal H2 + Ci2 reaction. In this case, changing the surface to volume ratio should not change the reaction rate. In fact, Chaikln showed that an eightfold Increase of the svirface to volume ration (S/V) changes the rate only by 1 o- VCi + Ci

(5)

atom is thereby released to the gas phase.

Aa for most

free radical reactions, the activation barrier of such a process is likely to be small.

This is the reason for the easy generation of free radicals

by a surface. The mechanism of reaction (5) that we have just represented in symbolic form may be pictured in the following way: approaches a surface radical-ion Ci~Zn++0

with expulsion of a

·Ζη+0 Ci

atom Into the volume.

of reaction (5) is the difference between the energy Ci~Zn++ I

q

=

and the energy

Qyci ~ ^Ci-Ci"

Valencia

V

Ocg.Qj y

We raa

A chlorine molecule

and forms a surface compound Qvci

of'dissociation of a TTi2

The heat

qt

of the bond molecule:

y that during this process, a surface free

sa

disappears while a free valence appears in the volume with

atomic chlorine. The surface compound CiZn++0

3.

223

F R E E V A L E N C E S OH REACTOR W A L L S

o r w i t h a homopolar bond

is represented b y the symbol V C i . The unexcited elements of the lattice are designated b y V 2 . Then the process of excitation of the crystal corresponding to the formation of two free valences, can be w r i t t e n as:

W h e n a gaseous C< atom hits a surface site where there is n o free valence, the following reaction m a y take place (representing chain termination o n the wall):

(6) Thereby, a chemlsorbed valence 2n

++

0

Ci

atom is formed together w i t h a surface f r e e

To fix the ideas, w e m a y imagine that a Ci surface, attracts one of the electrons of 0

a t o m , approaching a . The complex

or

is f o r m e d , depending Ci — heteropolar o r equal to quired to e x c i t e t h e

o n the nature of the strongest bond between 2n and h o m o p o l a r . The energy released i n process (6) is • U w h e r e U , already defined, is the energy recrystal.

A t thermal equilibrium processes (5) and (6) also go i n reverse: and Finally, we must also consider: (recombination of atomic chlorine w i t h a free v a l e n c e ) and (dissociation of V C i into gaseous atomic chlorine and a surface free v a l e n c e . A t equilibrium, w e therefore have four direct and four Inverse processes:

^ The process m a y w e l l take place i n two stages: (a) the gaseous C< atom h i t t i n g the surface is held tfcere b y relatively w e a k forces due f o r instance (following F . F . V o l ' k e n s h t e i n ) to a one-electron b o n d w i t h evolution of a few kcal; (b) this w e a k l y bound Ci atom m a y either evaporate a g a i n o r surmount a n activation barrier and form a surface compound V C i w h i c h Is n o w strongly chemlsorbed.

22k

VI.

W A L L INITIATION A N D TERMINATION

where

where

It is now easy to find the equilibrium ratio of the gaseous concentration (Ci) and the surface concentrations [V], [VCi], [ V 2 ] assuming 2 of course that 1 c m of surface consists entirely of either V or VCi or V 2 i.e.,. the sum of the concentrations of these species is constant. The surface concentration [VCi] is determined b y :

Here the constant

G

obviously depends on the surface to volume ratio.

W e see that the c o n c e n t r a t i o n o f chemisorbed Ci atoms w i l l increase w i t h the strength of the b o n d . If this bond is very strong so that is considerably larger than U and Q^jq^* the concentration of VCi will be so large that the surface becomes entirely covered with a chemisorbed layer. The process ^VCi ~ ^CiCi then exothermic and proceeds quite easily. However, since tEe w a l l is rapidly covered with V C i , the steady-state concentration of free valences V w i l l be quite small, and generation of free atomic chlorine becomes impossible. If, o n the contrary, is considerably smalle r than the surface concentration of VCi will be small. The w a l l remains clean b u t the process is now very endothennic and the 3teady-state rate of generation of chlorine atoms will again b e small. Therefore, f o r a given value of U , smaller than Q Q ^ C i ' the rate of generation w i l l b e maximum for some value of Qyc^, smaller, b u t not too much smaller, than Consequently, not every w a l l is capable of generating free radicals. In reality, however, because of the heterogeneity of all surfaces, different sizes have different values of Qajqi and the activity of a surface in the process of radical generation w i l l be determined, for a given value of U b y the relative number of 3ites w i t h optimum value of 1

A l l the preceding reasoning has ignored the role of entropy i n surface processes. Thus, w h e n gaseous chlorine is captured b y the w a l l , the entropy of the system decreases. This in turn decreases the probability of the chemisorbed state and consequently the equilibrium concentration of V C i .

225

W A L L ACTIVITY IN RADICAL GENERATION

Conditions f o r W a l l Activity In the Process of Radical Generation F o r an intensive surface generation of free atoms, the most essential requirement is that the w a l l be characterized b y a relatively small value of U . The condition f o r a significant wall generation of atomic chlorine is g i v e n b y the inequality If, on the contrary, U > the homogeneous d i s s o c i a t i o n w i l l be more favorable. ++

F o r crystals of the type of Z n 0 , a typical semi-conductor, the quantity 0 is rather small. It amounts to 1 - 2 e . v . , i.e., 2 0 - 5 0 k c a l . To be sure, these values are true for the b u l k of the crystal. A t the surface they m a y be even smaller. According to semi-conductor 2 theory, the concentration of free radicals, i.e., their number p e r c m of surface, is equal to [V] - G exp(- U / 2 R T ) where 1

2

Since there are about io -* particles per c m of surface, the 2 fraction of surface occupied b y free radicals is a = 10~ exp(- U / 2 R T ) . If the endothermicity of the process is 10 k c a l , if w e assume that its e Q value is - 0 and if w e remember that the number 2 of C i 2 molecules colliding p e r sec with i c m of surface, is at atmospheric pressure w e obtain the number of Ci 2 atoms formed per 1 c m of surface per second:

atoms p e r second per The data of A . M . Markevich [4] give the value for glass w a l l s . The rate of homogeneous generation

is n o m o r e than atoms per cc per second at 6oo°K. Note that usually the surface to volume ratio is ~ i (in a cylindrical tube w i t h d = 3 cm, W i t h the figures w e have chosen, the rate of

226

VI. WALL INITIATION AND TERMINATION

generation of atomic chlorine is 3.10k times larger on the surface than in the volume. If we select U = 1.5 e.v., the rate of wall generation is still 300 times larger than in the volume, although it is 100 times smaller than in the former case (U = 1 e.v.). If the process Ci2 + V VCi + Ci is thermoneutral and its activation energy is 3 kcal, even with U •» 1 . 3 e.v. the ratio of sur­ face to volume generation is equal to 105. With ε = 3 kcal, the rates of surface and volume generation will he of the same order of magnitude for U as high as 2 e.v. This is natural since then there is a small difference between the energy of dissociation of chlorine (57 kcal) and the energy required to form surface free valences (k6 kcal). One must not conclude, however, that a material with a very large value of U is not capable of generating Ci atoms. In fact, real crystals never have a perfect structure. Real crystals always have some sort of defects, for instance, vacancies, i.e., lattice positions not occupied by the ion e.g., 0 that should be there. Then, within the radius of action of the defect, the excitation energy U may be much lower than the value of U for the ideal crystal. This means that each such defect constitutes a free valence. In our simplified picture, the absence of an 0 atom makes the neighboring ions singly charged. Evidently these defects may move about the crystal and therefore the free valence may find itself at the surface. Such a surface will generate Ci atoms. If the geometry permits it, it is possible that during the course of the reaction, va­ cancies may be filled by strongly bound chlorine atoms. The latter, how­ ever, because of vacancies in the bulk of the crystal or because of the porosity of the wall, will migrate inside the crystal and liberate the sur­ face until the bulk of the crystal is saturated with chlorine atoms. Such a 'defect' wall may remain able to generate chlorine atoms during a long period at low temperatures, until finally it becomes completely poisoned. We have looked at the problem, using a chlorination reaction as an example. Naturally, a wall may jyst as successfully generate other free radicals e.g., H atoms from alcohols or hydrocarbons. At the same time, several chain reactions are made possible. Depending on the energies of chemlsorptlon and bond dissociation, the volume generation of radicals may be more or less important. Let us note that with hydrocarbons, an H atom will be attracted by the field of O and not of Zn++; a hydroxy1 group is formed: Ζη++0·" + RH

«- Zn++0" - H + R

or Zn++0 H+ + R .

5. WALL INITIATION WITH GAS IMPURITIES

227

At the same time, a hydrocarbon radical Is sent back to the gas phase. Be­ sides zinc oxide, any other dielectric or semi-conductor may be used as an active surface. To sum up, any active surface In contact with any molecule may act as a source of free radicals. By active, we mean a surface of which the crystals have a value of U inferior to the energy of dissociation of the molecule Into free radicals, or more generally, a surface with any (maybe large) value of U but a sufficiently large number of defects. 5·

Heterogeneous Radical Initiation with Impurities in the Gas Phase

Even if a wall exerts no direct action on a reacting molecule, it may considerably improve the Initiating action of an impurity. Por instance, hydroperoxides the 0-0 bond of which has an energy of ^5-50 kcal, do not dissociate sufficiently readily into free radicals homogeneously. They may ultimately disappear without forming free radicals. With a wall, owing to the process V + ROOH «- VOOH + R, the same peroxides may gen­ erate free radicals easily. The accelerating effect of small quantities of oxygen on hydro­ carbon reactions, may perhaps be explained similarly. One may assume that an O2 molecule reacts at the surface as follows: ν + O2 VOO - . Then a strong peroxide radical VOO - would be formed. The latter, in the presence of hydrocarbons, would react with a reactant molecule, say RH, giving a surface peroxide VOOH and a R radical. In the literature, one also finds reference to the accelerating effect of oxygen on the thermal decomposition of several organic compounds. Usually, the addition of small quantities of oxygen Increases the rate of decomposition, but only within certain limits. With large amounts of Og, the rate ceases to depend on the amount added. This effect of small quantities of O2 may be explained by talcing into consideration the role of O2 In heterogeneous chain Initia­ tion. Host Important in this respect is the observation that pretreatment of the reactor with O2 followed by evacuation, markedly accelerates the decomposition of tertiary butyl chloride, which Itself is not accelerated by Oxygen addition. Also more recently, v. A. PoltoraJc [23] has shown that the acceleration by oxygen of propane cracking, strongly depends on the state and treatment of the reactor walls. Thus, In a vessel treated with hydrogen fluoride, the initial rate in propane containing 296 oxygen, is four times larger than In a vessel with a surface covered by a film of hydrocarbon material.

228

VI. WAT,Τ, INITIATION AND TERMINATION

The effect of Inhibitors in various hydrocarbon reactions (e.g., NO for cracking) is also characterized by saturation. This fact strongly militates in favor of the heterogeneous character of the action of these additives. The usual explanation of saturation rests on the assumption of two simultaneous reaction mechanisms — chain and molecular. In our opinion, this is not very likely. The action of NO apparently resides In the fact that NO, be­ cause of its radical behavior, exterminates all the active surface sites, thereby preventing heterogeneous chain initiation. In the presence of O2 however, addition of NO catalyzes the reaction. As shown by V. A. Poltorsik [23], this appears to be due to the interaction between NO and oxygen, that facilitates heterogeneous generation of radicals. If the walls are treated with MoOy a substance that gives off oxygen easily, the inhibiting action of NO disappears. A number of cases of 'after-effects' apparently belongs to the same class of phenomena due to heterogeneous initiation in reacting sys­ tems after they have been submitted to an intensive period of chain initia­ tion. This happens for Instance when propane is oxidized at low tempera­ tures in the presence of bromine. R. A· Kalinenko [sb] has found that after cuttlng-off the illumination used for initiation, the reaction continues to proceed during tens of minutes at a rate comparable to that of the photo­ chemical process, although under these conditions, and in the dark, the reaction usually does not occur. Direct experiments have shown that this effect is due to a treatment of the surface by bromine atoms. When the reaction is studied in a flow system, the wall maintains its activity during 2-3 hours. Apparently, the catalytic effect of molecular chlorine on propane oxidation in a flow system is also related to wall behavior. Thus, N- M. Iinanuel1 [25] has found that after interrupting the supply of chlorine, the reaction continues to proceed during a long period at a high rate. These phenomena may be explained by assuming that halogen atoms are able to pene­ trate rather deeply into the wall. As a result, the rate of chain Initia­ tion on the walls changes. Indeed, surface properties also change, in par­ ticular the position of energy levels that determines the ease with which free valences are formed at the surface. This might explain the Increased rate of reaction itself. 6. Heterogeneous Initiation Due to Molecular Reactions Producing Free Radicals In what precedes, when we considered the effect of the wall on pure hydrogen and chlorine, we concluded that, in the case of a pure

229

6. HETEROGENEOUS INITIATION

H2 + Cig mixture, the wall does not perturb the Cig , ' 2Ci equilibrium and therefore does not change the reaction rate which is determined by the homogeneous concentration of chlorine atoms. This is true when, among the processes leading to radical formation both homogeneously and heterogeneously, there are no Irreversible processes related to the evolution and utilization of the heat of reaction and through which final reaction prod­ ucts are obtained. If chemisorption happens to be irreversible, it may produce a homogeneous concentration of atomic chlorine above its equi­ librium value. But if this process is irreversible, any effect of the wall stops after its entire surface becomes covered with a layer of chemlsorbed Ci atoms. We have seen already that the wall accelerates the reverse process if there are homogeneous Impurities (e.g., oxygen) that remove Ci atoms irreversibly. But even in a pure H2 + Ci2 mixture, besides the reversible chain termination step Cf + Ci >- Cig, one must take into account the irreversible process: H + Ci HCi. Because of the small concentration of atomic hydrogen established during the reaction, this irreversible process has a small rate. Even so, this process lowers the homogeneous concentra­ tion of atomic chlorine below its equilibrium value and therefore the wall exerts some accelerating effect even In the absence of impurities in the H2 + Ci2 mixture. Nevertheless, our discussion of the Interaction between Cig and the wall in the presence of H2 molecules, may still be Incomplete. It has been pointed out previously that there exists a possibility for the process: H2 + Ci2-—»- HC£ + H + C£ - 58 kcal. However, because of its large endothermlcity, such a process cannot play any essential role homo­ geneously. But on the wall, e.g., Zn++0 , the same process may require a smaller expenditure of energy because of the evolution of the heat of adsorption of an H atom: Zn++0

+ H2 + Ci2

.- -Zn+O" - H + HCi + Ci - U+ Q0^ - 58

.

A Ci atom leaves the surface and a surface radical ·Ζη+ is formed. The O-H bond of chemlsorbed H is naturally weaker than the O-H chemical bond in alcohols for instance (100 - 110 kcal) but the value of Qq-h may still be appreciable, maybe 60 kcal. If U = 40 kcal, Q = Qq_jj - U = 20 kcal and thus the formation of -Zn+ and Ci requires 58 - 20 = 38 kcal. This can take place at an appreciable rate. This rate evidently will be smaller than the rate of Ci2 + V »- CiV + Ci. The essential advantage of the former process, however, is that it is ir­ reversible since the H2 + Cig mixture reacts irreversibly: Hg + Cig >- 2HCi. By contrast, the process Cig + V is reversible.

230

VI. WALL INITIATION AND TERMINATION

This lrrevereiblllty has two consequences. First, the homogeneous concen­ tration of atomic chlorine will be higher than that determined by the Gl2 ι ' 2Cι equilibrium. Secondly, the surface concentration of free valences V will be larger than that corresponding to the equilibrium ν - Τ — gyWe are thus led to the hypothesis that many reactions may be affected by non-equilibrium processes of free radical generation on the walls as a result of reaction between reactant molecules or Intermediate products (see also h]). 7. An Attempt to Apply the New Ideas to Heterogeneous Catalysis As we have seen earlier, free radicals that initiate radical re­ actions in solutions may be generated by ions of variable valence. We may presume that surfaces of metals and semi-conductors may be the seat of similar phenomena of initiation of radical reactions since free valences are generated easily on such surfaces. We have considered Initiations of volume chains at the surface of a semi-conductor by means of the reaction: A2 + v —— AV + A

(1 )

According to earlier considerations, a molecule Ag strikes a site of the crystal surface containing a free valence V. As a result a free radical A is formed. We have seen that the number of free valences per cm2 is equal to η « C exp(- U/2RT) where C » 1O1^. At T = 6oo°, with TJ = 2 3 kcal we obtain η «• 10^. Then the fraction of the surface occupied with free valences is (10^/10'^) = io~®. Thus, the svirface concentration of free valences V on dielectric and semiconducting materials is relatively small. It must be noted that the atoms A which are formed need not leave immediately for the volume. Atom A may be bound to the surface by relatively weak forces of the type, for Instance, of a one-electron bond (as shown by F. P. Vol'kenshtein [26]. If the energy of this bond is small, substantially smaller than 10 kcal, the surface atom (A) will soon evaporate into the volume (A) » A - q1. However, while atom A is bound to the surface, it may enter into the reaction: (A) + AV

- A2 + V

(2)

7.

NOTES ON HETEROGENEOUS CATALYSIS

231

or (3) In this fashion, not every atom formed In reaction (1) w i l l reach the v o l u m e . A t equilibrium, all these processes w i l l be compensated b y the corresponding reverse processes and will n o t affect the quantity of atoms i n the v o l u m e . This situation radically changes if another gas B 2 is present i n the system and a n irreversible reaction between Ag and B 2 can take place:

In this case, atoms (B) m a y react w i t h

(A), AV:

before evaporating, m a y react w i t h

BV

o r atoms

These reactions are irreversible under the conditions of the experiment since they lead to f i n a l p r o d u c t s . These processes decrease the concentration of (A) and (B) and therefore also the generation of A and B i n the volume b u t a heterogeneous reaction takes place on the crystal surface. If 'physically' adsorbed gases ( A 2 ) and o n the surface, ordinary chains are also possible:

(B2)

are also present

1

and so o n u n t i l reactions 3 ) o r 3 ) take p l a c e , terminating the c h a i n . Moreover, the system of reactions also represents a chain reaction i n w h i c h the valences of the surface take part as active centers of the c h a i n . Termination of this chain occurs through recombination of the free valences, I.e., as a result of these processes:

The interdependence of these two types of chains m u s t b e n o t e d . Termination of the ordinary chain [reactions 3) and 3')] is at the same time the initiating step f o r the other chain since i n these reactions a free valence V is f o r m e d . If bonds

and

are very strong, reactions 3 ) and

232

VI. WALL INITIATION AND TERMINATION

3 1 ) a r e m o r e rapid t h a n reactions 4 ) a n d 4 ' ) . T h e n t h e process d o e s n o t lead to final product AB but to chemlsorbed species AV and BV. This will continue until practically the entire surface is covered with chemisorbed atoms. Then the fraction of free surface becomes so small that reactions 1) and 11) proceed very slowly. Then the rate of the catalytic process tends to zero and the process stops when the surface is completely covered with chemlsorbed gases. If, on the other hand, bonds A-V and B-V are weak, the heats q1 and qj of reactions 1 ) and 1 1 ) axe small since Q1 = Qav - Qaa and q] = Qgy - Qrr where Qav- and Qbv- are the bond energies of —A - V and B-V. "Values of q1 ancT qj are-then negative since Qaa is relatively l a r g e . I t f o l l o w s t h a t t h e energy o f activation o f 1 ) a n d 1 1 ) may be quite large. This also leads to very small catalytic rates. In this fashion, to each pair of values of Qaa and Qrr, there must correspond on the surface certain values of Qav and Qbv which must not be too small or too large. Only in this case will the surface have optimum catalytic properties. Small specific Impurities profoundly modify the electrical con­ ductivity of semi-conductors. This is due to the fact that hetero-atoms are either traps for, or donors of, free electrons and therefore Increase the concentration of electrons or positive holes (i.e., the number of free valences). Thereby the value of the energy U is also lowered. Other lattice defects such as the absence of an atom at a lattice position act in the same direction. It may be that this will also explain the action of promoters and also the different activity of catalysts prepared In differ­ ent ways. Let us remark that according to these considerations, the number of free valences on the surface is determined by the entire catalyst volume since a free valence may move from one place to another. The viewpoint just presented on radical catalytic reactions has recently been considered by P. P. Vol'kenshtein [27] who first put forward a number of propositions in this field, and subsequently by N. N. Semenov and V. V. Voevodskil [28]. More recently, S. Z. Roginskii has shown that a series of phenomena in catalysis could not be explained easily without assuming radicals and surface chains in mechanisms of catalytic reactions. The same viewpoint is also applicable to metallic catalysts. In metals, electrons are unlocalized. None belongs to any par­ ticular metal atom. Energetically, electrons are paired on very closely spaced levels. This system of occupied levels is Immediately followed by a system of vacant levels. Thus, as a result of thermal motion, Indi­ vidual electrons may be easily excited and be transferred to vacant levels.

7. NOTES ON HETEROGENEOUS CATALYSIS

233

This explains the large electrical conductivity of metals. Free unpaired electrons on these higher energy levels give to a metal properties of a polyvalent free radical. This accounts for the possibility of relatively easy reactions between a metal surface and ad­ sorbed molecules, similar to the reaction between free radicals and mole­ cules: Na + Ci2 NaCi + Ci. In metals, free electrons strongly interact with each other and with metal ions. This interaction partially determines the high values of sublimation energies and electronic work functions of metals. Consequent­ ly, the surface of a metal is also a relatively inactive radical. This means that the energy required to promote a metallic electron into a 'valence active state1 is quite large. Indeed, this is a state in which the electron becomes localized around a given atom. This process is re­ quired to suppress the 'resonance' of the electron with all the other metallic electrons. In this sense, a metal surface is a radical of small activity, like the allyl radical or NO for which the interaction of the •free' valence electron with the nuclei and with the other electrons (i.e., resonance) is large. Whereas the surface concentration of radicals is quite small for semi-conductors, a metallic surface is practically entirely covered with 'free valences.' This circumstance strongly increases the rate of process i): 1 )

A2 + V

~ AV + A .

A second difference with semi-conductors is that reactions of the type: 3) A H - V 2

AV + V

do not proceed in a metal since essentially all metallic atoms are free radicals. Therefore the concept of V2 is not applicable here or it may apply to these parts of the surface where there are no free valences. More­ over, the process A + V AV 2") which is very infrequent for a semi­ conductor because of the small concentration of V, here becomes extreme­ ly important. Like all exothermic recombination steps between radicals, it has no activation energy, or, if there is one, it is very small. Therefore, radical A formed by reaction 1) is almost immediately captured by the metal surface. This leads to chemisorption of molecule Ag taking place in two almost simultaneous steps: 1 ) A2 + V 2") A + V

AV + A » AV .

The net reaction is: A2 + 2V

2AV.

If there are two gases A2 and B2 in the system, both will be adsorbed by this mechanism and they will not react together since processes

VI. WALL INITIATION AND TEEiMINATION A + BV ·- AB + V and B + AV » AB + V leading to reaction, cannot take place. However, as the surface becomes progressively covered with compounds AV and BV, the atom A being formed will be more and more often surround­ ed by neighboring BV compounds. Then before reaction 2") can take place, atom A must diffuse to the nearest vacant site. This requires a certain time τ increasing with the fraction of surface covered with chemisorbed atoms AV and BV. After this time τ, either reaction 4) or 41) may take place. Therefore if surface coverage is sufficiently high, a catalytic process becomes possible on the metal surface. As in the case of semi-conductors, too large bond energies AV and BV will be harmful to the reaction since AV and BV will then cover practically the entire surface, leaving no vacant site for reaction 1) to proceed: A2 + V AV + A. On the other hand, too weak AV bonds will also be detrimental to the catalytic reaction since the heat of reaction 1) becomes small. (Since Qav is small, q1 becomes negative). Then the activation energy of process 1) becomes large. Consequently, on metals as on semi-conductors, reaction will be most favored by some intermediate val­ ues for the energy of the bonds AV. 'Physically1 adsorbed gases, as was already shown for semi-conductors, may lead to chain propagation: (A) + (B2) •AB + (B), (B) + (A2) • AB + (A) and so on. Lattice defects (points, vacancies, foreign atoms) will also effect reaction rate since they will change the energy L of electron derealiza­ tion and may therefore change all the heat quantities, in particular the energies of bonds AV and BV. This situation leads to surface hetero­ geneity both for heats of chemisorption and activation energies. Reaction will take place preferentially on surface patches where Qav and Qbv — — have the most favorable value for the reaction rate. Finally, all these factors must be considered as tentative, but we think they are a good starting point from which a theory can be develop­ ed to explain certain types of catalytic processes. In particular, our considerations deal with elementary steps and not with the overall reaction and are therefore unable to cope with such phenomena as, for instance, catalyst modification. We have studied above the mechanism of a metathetic catalytic reaction of the type A2 + B2 •2AB. In practice, addition reactions are quite interesting. The simplest illustration is ethylene hydrogenation. The radical mechanism of addition reactions is basically the same as that of abstraction reactions. But there are some peculiarities. In abstraction reactions, a molecule A2 reacting with a free valence V gave AV (a chemisorbed atom, chemically inert) and another atom, weakly bound to the surface with the properties of a free radical. A molecule with a double bond upon reaction with a surface free valence, will become

7.

235

NOTES ON HETEROGENEOUS CATALYSIS

a free radical one end of w h i c h Is strongly bound to the surface: This chemisorbed radical m a y react following several p a t n s . The first path is the reaction of the radical w i t h This leads to a tightly chemisorbed ethylene m o l e c u l e . molecule

The second p a t h is the reaction of the chemisorbed radical w i t h a H g , coming from the gas phase or f r o m a physically adsorbed layer:

This results i n the formation of a n adsorbed radical A V f o r abstraction reactions.

VCHgCHj

analogous to

The atom H m a y immediately attack the C — V bond of the radical, leading to the f i n a l reaction product ethane and liberation of a f r e e surface valence:

O r the atom

H

m a y react w i t h the surface:

This m e a n s regeneration of a f r e e valence and formation of a chemisorbed atom H . This last process w i l l progressively lead to a surface completely covered w i t h chemisorbed species: H atoms, ethyl radicals and ethylene m o l e c u l e s . H o w e v e r , as surface coverage Increases, chemlsorptlon of H atoms w i l l become less frequent and other processes w i l l take over:

and

I n addition at large values of coverage, the primary radicals w i l l start to react w i t h chemisorbed H atoms:

VCHgCHg -

A l l these processes combined w i l l lead to a quasi-steady coverage of the surface w i t h chemisorbed species and the hydrogenation of ethylene w i l l proceed at a certain r a t e .

236

VI. WALL INITIATION AND TERMINATION

The radical VCH2 - CH2 - may react tuider certain conditions with another ethylene molecule to fonn a dimer or a trlmer. Such processes usually accompany the hydrogenation reaction and poison the catalyst. With more complex hydrocarbons, reaction of the primary radical VCH2 - CH2 is likely and, depending on the catalyst, isomerization, polymerization and cycllzation of hydrocarbons will take place. To illustrate our viewpoint, let us sketch a hypothetical reaction scheme for hydrocarbon synthesis from CO and H2Such complex catalytic syntheses forming large complex molecules from simple ones, may not be explained by the action of adsorption forces on reactants. Organic chemists have long since recognized that such processes may take place only via free radicals. Thus, N· D· Zelinskil and N. I. Shuikin [29] have come to the conclusion that formation of a diradical CH2 is necessary to explain a number of catalytic and organic reactions. Ya. T. Eldus and N. D. Zellnskii [30] have used this concept to explain hydrocarbon synthesis. N. N. Semenov and V. V. Voevodskii have shown that the reaction scheme postulated by N. D. Zellnskii and Ya. T. Eidus may be obtained without assuming large concentrations of CH2 radicals if monoradical reactions are considered together with migration of the free valence in the radical. According to N. N. Semenov and V. V. Voevodskli, during formation from simple reactant gases of molecules containing a linear car­ bon chain, only one atom at the end of the chain must be attached to the surface while new links are added continuously at the other end of the chain. Let us consider a possible scheme by which complex products may be formed from carbon monoxide and hydrogen: 1 ) V + CO

-_ VC ! = O

2) VC = O + H0

•. VC = O + H . H

Atom

H formed in reaction 2) either reacts immediately with VC = O "I H

to form an alcohol radical: 3 ) VC = O + H H H

- VC - OH "I H

or reacts with the surface to give a chemisorbed face valence: 3 1 ) H + V2

HV + V

.

H atom and a free sur­

7.

NOTES ON HETEROGENEOUS CATALYSIS

237

The alcohol radical

obtained I n reaction 3 ) reacts w i t h hydrogen:

Atom

H

either attacks the

C - 0

linkage to form w a t e r and radical

or attacks the V - C linkage to give one of the final products — alcohol (frequently observed during these syntheses):

methyl

The radical

reacts in turn w i t h

CO:

By alternation of processes 2 ) , 3), 4), 5), 6), the carbon chain w i l l grow w i t h formation of hydrocarbon radicals V ( C H g ) n - Radicals of this type react w i t h Hg to give AtomH reacting w i t h the bond gives a saturated hydrocarbon and regenerates 1 a free valence at the surface V . Reactions of type 5 ) w i l l produce higher alcohols.

238

VI.

W A L L INITIATION A N D TERMINATION 1

Participation of H atoms i n reaction 3 ) w i t h formation of chemlsorbed species H V leads to progressive covering of the surface w i t h H V complexes and chemlsorbed alcohol radicals (of the type

1

H o w e v e r , w h e n surface coverage becomes large e n o u g h , reaction 3 ) becomes slowed d o w n and H atoms w i l l m a i n l y enter Into a l l the other reactions of the scheme (3), 5 ) and 5 ' ) ) . M o r e o v e r , w h e n the concentration of H V is l a r g e , radicals formed i n reactions 3 ) , 5 ) and 6 ) w i l l react n o t only w i t h Hg b u t a l s o w i t h H V w i t h addition of a h y d r o g e n a t o m and liberat i o n of a valence V . This brings reaction rate to some steady l e v e l . REFERENCES [1] [2]

V . V . V o e v o d s k l l , D o k l . A k a d . N a u k , U . S . S . R . , 90, 8 1 5 , (1953). Y u . B . K h a r i t o n and Z . V a l t a , Z e i t s . f . P h y s . , 39, 5^7, (1926); N . N . Semenov, Zhurn: R u s s k . P i z . - K h i m . O b s h c h . chasfc. P i z . , 58, 329, (1926)} N . N . S e m e n o v . 'Chain R e a c t i o n s ' , L e n i n g r a d , T 9 3 4 . Translated and revised as Chemical kinetics and c h a i n reactions", Oxford, 1935.

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[5]

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A . E . B i r o n and A . B . N a l b a n d y a n , A c t a Physicochim., U . S . S . R . , 6, 4 3 , (1937). L . I . A v r a m e n k o , Z h u r n . P i z . K h i m . , £ 3 , 7 9 0 , (1949); (Canada: N a t . R e s . Council, T T - 1 2 3 , 1950). Z . K- M a l z u s , A . M . Markevitch and N . M . E m a n u e l ' , D o k l . A k a d . N a u k , U.S.S.R., 8 9 , 1049, (1953)V . I . Urizko and M . V . Polyakov, D o k l . A k a d . N a u k , S.S.S.R., 95, 1239, (1954). A . M . M a r k e v i t c h , Z h u r n . P i z . K h i m . , _3£, 735, ( 1 9 5 6 ) . p . p . Volkenshtein, Electrical conductivity of semiconductors, (in Russian), G o s t a k h i z d a t , 1947.

7.

NOTES ON HETEROGENEOUS CATALYSIS

239

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F- P- Volkenshtein, Z h u r n . P l z . K h l m . , 2 7 , 159, (1953)• p . p . Volkenshtein, Z h u r n . P l z . K h l m . , 27, 167, (1953)' p . P . Volkenshtein, Z h u r n . P l z . K h l m . , 26, 1462, (1952). L . I . Avramenko, M . I . G e r b e r , M . V . N e l m a n and V . A . Shushunov, Z h u r n . P l z . K h l m . , 20, 1347, (1946).

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V . A . Poltorak and V . V . V o e v o d s k l l , D o k l . A k a d . N a u k , U.S.S.R., 91, 589, (1953). R . A . K a l i n e n k o , Thesis f o r Diploma, Moscow State University, 1952. 1 K . E . Kruglyakova and N . M ,.8 Q n a n u e l , Izvest. A k a d . N a u k , U . S . S . R . , Otdel. Khlm. Nauk, N o . » (1957). P . P . Volkenshtein, Z h u r n . P l z . K h l m . , 21, 1317, (1947)} 2 6 , 1462, (1952)~ ~~

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P . P . Volkenshtein, Z h u r n . P i z . K h l m . , 2 7 , 159, 167, (1953); U s p e k h l P l z . N a u k , j>0, 2 5 3 , (1953). V . V . Voevodskll, P . P . Volkenshtein and N . N . Semenov, Sbornik "Voprosy khlmlcheskol kinetiki, catalysa 1 reaktzlon n o l sposobnosly". A k a d . N a u k , U.S.S.R., p . 4 2 3 , (1955)• N . D . Zelinskii and N . I . Shuikin, D o k l . A k a d . N a u k , U.S.S.R., 3, 2 5 5 , (1934). N . D . Zellnskll and Y a . T . E i d u s , Izvest. A k a d . N a u k , U.S.S.R., O t d e l . K h l m . N a u k , 289,(1940); 190, (1942).