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Studies in the thermal decompositions of isomeric pentanes Chrysochoos, John 1962

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STUDIES IN THE THERMAL DECOMPOSITIONS OF ISOMERIC PENTANES by JOHN CHRYSOCHOOS Dipl. of Chem. (Athens) 1957 A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE. in the Department of Chemistry accept this thesis as conforming to the required standard. THE UNIVERSITY OF BRITISH COLUMBIA April, 1962 In presenting this thesis in partial fulfilment of the requirements for an advanced degree at the University of British Columbia, I agree that the Library shall make i t freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the Head of my Department or by his representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission. Department of The University of British Columbia, Vancouver 8, Canada. Date 7 - 5" - l j 6 2 T L An investigation has been made of the pyr o l y s i s of n-pentane, isopentane and neo-pentane i n a s t a t i c system at temperatures near 500°C. Both uninhibited and i n h i b i t e d reactions were investigated. Of p r i n c i p a l concern has been the effect of v a r i a t i o n of i n i t i a l pressure of hydrocarbon on the decomposition products f o r the uninhibited reaction. The effect of v a r i a t i o n i n surface-to-volume r a t i o on rates and on the d i s t r i b u t i o n of the products has been also a point of concern. Structural effects on rates as w e l l as the v a r i a -t i o n of energy of a c t i v a t i o n and frequency factors with pressure have been considered of importance. Investigations of orders of reaction provide information as f a r as the mechanism i s concerned. As a f i n a l point the most important task f o r the uninhibited reaction was a mechanism explaining the r e s u l t s obtained, and off e r i n g l o g i c a l reasons f o r the s i m i l a r i t i e s and differences between the isomeric pentanes. For the i n h i b i t e d reaction the points of p r i n c i p a l concern considered have been: the effect of n i t r i c oxide on the product d i s t r i b u t i o n ; the effect of packing on both rates and products, the behavior of NO. The investigation has as a main purpose to determine whether the role of n i t r i c oxic5e as an i n h i b i t o r involves homogeneous or heterogeneous reactions. Whether W was consumed or not was also an important question. As a f i n a l point a mechanism i s also II'-proposed f o r i n h i b i t e d reactions which accounts f o r the experi-mental re s u l t s and attempts to give a l o g i c a l explanation of the i n h i b i t i o n phenomenon. IX ACKNOV/LEDGEMBNTS The present i n v e s t i g a t i o n i n the thermal decomposition of isomeric pentanes, has been made as a p a r t i a l f u l f i l m e n t of the requirements for the M. Sc. deg-ree, under the general supervision of Dr. W.A. Bryce, to whom the author i s greatly indebted f o r h i s advice i n both the experimental work and the w r i t i n g of t h i s thesis. The author i s also indebted to the University of B r i t i s h Colum-bia f o r the offered assistantship during the completion of the present study. He i s also greatly indebted to Mrs. Z e l l f o r the valuable f a c i l i t i e s she offered to him f o r the use of infra - r e d instruments i n her laboratory. A p r i l , 1962. I l l " TABLE OF CONTENTS Page PREFACE INTRODUCTION CHAPTER I The Uninhibited decomposition 2 Heterogeneous Reaction 7 The Inhibited Reaction 11 EXPERIMENTAL CHAPTER I I Materials and methods 24 Gas chromatographic equipment 25 Detector 25 Inf r a Red Analysis 25 The pyrolysis apparatus 26 Reaction vessel and furnace 26 Description of a t y p i c a l experiment 28 RESULTS CHAPTER I I I Uninhibited p y r o l y s i s 33 Rates of reaction 33 Rates of reaction for isopentane 36 Vari a t i o n of rate with c o l l i s i o n frequencies 36 The effect of packing for isopentane 40 Rates f o r n-pentane 42 The effect of packing 42 Rates f o r neo-pentane 42 The effect of Packing 42 Order of reaction 46 Order of reaction f o r n-pentane 46 Effect of packing for n-pentane 46 Order of reaction f or isopentane 46 Order of reaction f or neo-pentane 50 Effect of packing 50 Energy of a c t i v a t i o n and frequency factors 50 - n-pentane 52 - isopentane 53 - neopentane 54 An a l y t i c a l r e s u l t s f o r - isopentane 55 - n-pentane 56 - neo-pentane 56 I F Variation of the product with - time 56 - i n i t i a l pressure 58 Variation of product r a t i o s with - time 58 - i n i t i a l pressure 63 Variation of C2H4/C2H6 with pressure 63 The effect of packing 71 Variation of i-C4Hs/C3H6 with pressure 75 The eff e c t of chain branching on rate 77 Inhibited reaction 79 N i t r i c oxided concentration f o r the f u l l y i n h i b i t e d reaction f o r - isopentane 79 - n-pentane 83 - neo-pentane 83 Variation of products with time 84 Variation of the decomposition products with IflrO) as S/V" r a t i o 99 Variati o n of product r a t i o s with P$T0)and S/V 103 - v a r i a t i o n of C2H4/C2H6 103 - v a r i a t i o n of CH4/C2H6 103 - v a r i a t i o n of O2H4+O2H6/C3H6 103 - v a r i a t i o n of CoHA^Hg/C^s 1 1 0 - v a r i a t i o n of CH4/C2H4 with P&TO) 110 - v a r i a t i o n of C 2H6/C 2M0 with BJfO) 110 - v a r i a t i o n of i-C^Ho/C^fig 110 consumption of N i t r i c oxide 110 DISCUSSION CHAPTER IV/ Uninhibited decomposition 118 - isopentane 118 The o v e r a l l mechanism f o r isopentane 126 - n-pentane 128 The o v e r a l l mechanism f o r n-pentane 132 - neo-pentane 133 Inhibited decompositions 138 - isopentane 139 - n-pentane 149 - neo-pentane 151 CHAPTER V Kinet i c laws the thermal decomposition of isopentane 153 Uninhibited reaction 153 Inhibited reaction 162 References 164 Appendix I Calculation of a c t i v a t i o n energies and frequency-factors f o r n-, i s o , and neo-pentane 170 Appendix I I A n a l y t i c a l balance f o r the decomposition products of n-pentane 173 VI Page LIST OP TABLES I Varia t i o n of E^ and A with P^ f o r n-pentane 53 I I V a r i a t i o n of E A and A with P-j_ f o r isopentane ... 53 I I I V a r i a t i o n of E^ with P-j f o r neo-pentane 54 I V A n a l y t i c a l r e s u l t s f o r isopentane 55 VI Variation of the products of decomposition of isopentane with time 58 VII Variation of product r a t i o s with time 62 VIII Variation of C_H./CpH,. with pressure and temperature . 7. 75 IX' Var i a t i o n of C3H6 /I-C4H8, EC2/C3H6, •^02/1-0^8 r a t i o s with pressure f o r n-pentane 77 X Variation of P with time f o r various PJTO f o r neo-pentane 84 XI Consumption of NO 117 XII A n a l y t i c a l r e s u l t s f o r isopentane. Inhibited and uninhibited reaction 139 XIII Variation of C2H4/C2H5N0 with NO 142 XIV" Rate constants f o r various pressures and temperatures f o r isopentane 170 XV" E^ and A f o r isopentane 170 XVI Rate constants for various pressures and temperatures f o r n-pentane 172 XVII E A and A f o r n-pentane 172 XVIII Rate constants for various pressures and temperatures f o r neo-pentane 172 XIX E A f o r neo-pentane 172 XX A n a l y t i c a l r e s u l t s f o r n-pentane 173 VII' LIST OF FIGURES PAGE 1. Variation of energy with the reaction coordinate.. 3 2. Variation of po t e n t i a l energy with the reaction coordinate f or a unimolecular reaction 3 3. Shematic outline of thermal decomposition apparatus 27 4. Furnace heater c i r c u i t 29 5. Chromatogram of the products on HMPA column ...... 30 6. Chromatogram of the products on s i l i c a gel column. 31 7. Variation of pressure with time f o r isomeric pentanes 34 8. Variation of the rate (dp/dt) with time at various values of P-|_ f o r isopentane 35 9. Variation of K o v e r a l l with i n i t . pressure f o r isopentane 37 10. Va r i a t i o n of i n i t i a l rate with pressure f o r i s o ^ pentane 38 11. R-|_ vs Z]y2;i_4 ( c o l l i s i o n numbers) f o r isopentane .. 39 12. k o v e r a l l vs Pp f o r packed and unpacked vessel ... for isopentane 41 13. Variation of R-|_ with i n i t i a l pressure f o r n-pentane 43 14. Effect of packing on i n i t i a l rate f or n-pentane .. 44 15. R-j_ vs Pp f o r neo-pentane 45 16. Order or reaction f o r unpacked vessel f o r n-pentane 47 17. Order of reaction i n packed and unpacked vessels fo r n-pentane 48 18. Order of reaction f o r isopentane 49 19. Order of reaction f o r neo-pentane ; 51 20. Product y i e l d i n mole. $ vs time for n-pentane ... 57 21. Product y i e l d i n mole.$ vs i n i t i a l pressure for n-pentane 59 22a. Variation of decomposition products with pressure fo r iso-pentane 60 22b. Variation of the products with i n i t i a l pressure f o r iso-pentane 61 23. Va r i a t i o n of C2H6/Q2H4 ra'tio with time f o r n-pentane 66 24. Variation of CH4/C2H6 r a t i o with P i at various temperatures f o r isopentane 67 25. CH4/C2H6 vs P i n i t i a l f o r n-pentane 68 26. Variation of C2H4/C2H6 with P-[_ at various temperatures f o r isopentane 69 27. Variation of CpH 4/C 2H 6 r a t i o w i t h P x f o r n- and isopentane 70 28. Effect of packing on the C2H4/C2H6 r a t i o at various pressures f o r n-pentane 72 29. Effect of packing on the C2H4/C2H6 r a t i o at various pressures f o r iso-pentane 73 29b. Influence of packing on the v a r i a t i o n of C2H4/C2H6 with P-j_ f o r n-pentane 74 n i l 30. V a r i a t i o n of i-C.H^C^Hg r a t i o with i n i t i a l pressure fo r v a r i a t i o n values f o r T, f o r isopentane 76 31. Structural effect of isomeric pentanes on the i n i t i a l rate 78 32. Rate ( i n i t i a l ) vs P^o i n unpacked reaction f o r i s o -pentane * 80 33. K o v e r a l l vs PNO at various S/V" r a t i o s f o r isopentane.. 81 34. Effect of PNO on rate ( overall) f o r n-pentane; packed vessel 82 35. Neo-pentane k over vs PNO 85 36. Y i e l d of C2H5NO and C2H6 with time f o r isopentane; f u l l y i n h i b i t e d reaction; S/V 2.5 86 37. C2H4 vs time at various S/v r a t i o s and dif f e r e n t PNO f o r isopentane; f u l l y i n h i b i t e d reaction .... 87 38. i-C.Hg vs time at dif f e r e n t PNO for isopentane 88 39. Product y i e l d vs time i n the f u l l y i n h i b i t e d reaction f o r iso-pentane 89 40. Y i e l d of C2H5 vs time at various S/V r a t i o s and PNO f o r isopentane; f u l l y i n h i b i t e d reaction.... 90 41. Product y i e l d vs PNO f o r n-pentane 92 42. Effect of NO i n the products of C3H6 at various S/V r a t i o s f o r isopentane 93 43. ±-C^EQ VS PNO i n packed and unpacked vessels f o r i s o -pentane 94 44. Product y i e l d vs PNO f o r neo-pentane 95 45. C2 H6 v s PNO at various S/V r a t i o s f o r isopentane 96 46. OEA VS PNO at various S/V, r a t i o s f o r isopentane 97 47. Y i e l d of C2H4 vs PNO at various S/V r a t i o s f o r isopentane 98 48. C2H5N0 vs PNO f o r isopentane 102 49. C2H4/C2H6 vs PNO f o r isopentane 104 50. C2H4/O2H5 vs PNO atrvarious S/V r a t i o s f o r isopentane. 105 51. CoHz/CpHg vs PNO f o r n-pentane 106 52. CH /C H, vs PNO at various S/V r a t i o s f o r isopentane.. 107 53. CHg/CIHg vs PNO f o r n-pentane 108 54. C2H4+ C2H6/C3H6 vs PNO f o r n- and isopentane 109 55. C 2H 4 + C2H6/C4H8 vs PNO f o r n- and isopentane I l l 56. CH4/C2H4 vs PNO f o r n-pentane 112 57. C2H6/C2R5N0 vs PNO at various reaction extent f o r isopentane 113 58. i-C4H8/C5H6 vs PNO at various S/V r a t i o s f o r isopentane 114 59. NO consumption vs PNO for isopentane 116 X PREFACE The present study, i s divided into three main sections: introduction, experimental and discussion. The introduction provides a h i s t o r i c a l survey of the subject. The experimental, which consists of two chapters, I I and I I I , deals with the description of the apparatus and materials and the experimental r e s u l t s . In the t h i r d section, consis-t i n g of two chapters IT and V, mechanisms and k i n e t i c laws are proposed to explain the experimental r e s u l t s . Each chapter i s divided into two main d i v i s i o n s , one dealing with the uninhibited decomposition and the other with the i n -h i b i t e d reaction, and each d i v i s i o n into three more sub-d i v i s i o n s , dealing with isopentane, n-pentane and neo-pentane respectively. A p r i l , 1962. - 1 -CHAPTER I INTRODUCTION The thermal decomposition of hydrocarbons has been subject of investigation f o r several decades. The existence of free r a d i c a l s was proved i n many of these decompositions by the Paneth metallic mirror technique ( 1 , 2 , 3 * 4 ) as well as by mass spectrometric methods ( 5 > 6 , 7 » 8 ) . I t i s also w e l l known that the decomposition of several hydrocarbons i s accelerated by methyl r a d i c a l (9»10) as w e l l as by some traces of impuri-t i e s such as Q>2 e " t c « ( 1 1 » 1 2 , 1 3 ) . Rice was the f i r s t to treat the decomposition of a simple p a r a f f i n , C2H^ as a free r a d i c a l mechanism (14) . The Rice-Hertzfeld mechanisms along with the alternative proposed by llilhle and Therle (15) formed the basis f o r the invest i g a t i o n of thermal decomposition of hydrocarbons. • Most of these decompositions show common charac-t e r i s t i c s . A marked feature i s the chain nature of the mecha-nism as w e l l as the phenomenon of i n h i b i t i o n . The rate of decomposition of almost a l l of them i s reduced by several i n h i b i t o r s of which the most important are n i t r i c oxide, propylene and C^Hs. The i n h i b i t i o n by NO was discovered and • exhaustively investigated by Hinshelwood and h i s co-workers fo r Ctj hydrocarbons. They compared the effect of NO and C-^ Hg and found that the same i n h i b i t i o n i s attained i f the r e l a t i v e concentrations of NO/C^Hg i s 1/12, which means, NO i s 12 times more effective than propylene ( 1 6 ) . Stepukhovitz has compared qua n t i t a t i v e l y the action of isobutylene and - 2 -propylene f o r several hydrocarbons (17,18,19) and found that the same l i m i t i n g rate i s always reached f o r these two i n -h i b i t o r s . I'" The Uninhibited decomposition In the normal decomposition, the main problem a r i s i n g i s the nature of the mechanism, and of course an examination of the chain reaction i s f i r s t of a l l necessary. I t has been known f o r a long time that the majority of gas-phase reactions involve considerable a c t i v a t i o n energy i n the range from 30 - 50 KCal/Mole. On c o r r e l a t i n g the a c t i -vation energy with the energy b a r r i e r Eo and the heat of was found that, . , . _, ^  , _ reaction (20; ,^  often.these quantities were not related. Every molecular decomposition proceeds through a t r a n s i t i o n state (Pig. l ) and f o r endothermic reactions Por free r a d i c a l reaction Eo = A - a|q| where A i s a constant quantity approximately 11 KCal/Mole., s l i g h t l y dependent on the type of the reaction and, a, i s another constant approxi-mately 0.25 (20). Polanyfand Evans (21) found the existence of a formula f o r exothermic reactions &Eo=- a&q. The relationship between these symbols i s given i n (Pig. l ) . Whether the reaction i s a molecular or a free r a d i c a l mecha-nism can now be discussed. Consider the reaction E^a Eo+q, but for exothermic reaction E = Eo The rate i s R^molec." ^ C-^ JP] where k i s a bimolecular rate constant. Hence fcmoiec.= l O 1 0 ^ 5 1 [A] [B] (20) when 1 0 1 0 / i s the preexponential factor f o r the bimolecular reaction. Supposing that the above reaction proceeds through a free r a d i c a l process Rj+B —* M!>R2 ( x) R2*A — > M2+H1 (2) etc. where R-j_ and R 2 are r a d i c a l s produced from A, and M]_, M2 are molecules, we get: fcchain= k i [ R i l t B ] = k 2 [R 21 CAI & lO^e'^™ [Hjjfcl where E]_ i s the a c t i v a t i o n energy f o r the step with the lowest rate. Then by forming the r a t i o -10 -Ei/RT = C % i l e ^ ( i y The concentration of R-j_ i s calculated by the r a t i o of the rate of formation over the rate of disappearance. I f I i s the i n t e n s i t y of l i g h t for a photoinitiated reaction A J L > Rj+ R 2 ^ R 1 + R 1_^ M 2 (3) so I s K 3 [ R f ] 2 or [ R£| = ( I / K ^ ) ^ Substituting i n (l)'we get For molecular reaction E^ i s high around 40 KCal/Mole. For chains reaction E]_ i s small around 5-10 KCal/Mole. so - 5 -X ~ W X ^ 1 ' ^ E ^RT^ . I varies from l O 1 ? ^ 1 0 1 5 and i s usually very much less than unity. Hence U M C l O ^ / l ) * e - 3 5°°%T -4 -6 Values of ^ , under normal conditions vary from 10 ^ 10 f o r hydrocarbons decomposition (20). Thus reaction i n the gas phases proceed much more through a chain mechanism than by the molecular mechanism. In the case of \+ H £ —9- 2 BrH }f = o . l while f o r 2 •Tg + H 2 —>-2TH X =2.10 i . e . the main process i s a molecular process. In a chain type mechanism three stages are d i s t i n c t . I n i t i a t i o n , propagation and termination. In thermal decompo-s i t i o n of hydrocarbons, i n i t i a t i o n can take place through a C-C s p l i t , provided that D(C-C)'is less than D(C-H). This reaction can take place either unimolecularly or bimolecularly. Reactants proceed to products through an activated complex. In A£^B we can follow the scheme (22) A The reaction w i l l take place through two excited states, with E* greater than that required f o r the reaction. The p o t e n t i a l energy of such a system i s given i n (Eig. 2). The frequency factor can be considered as a c r i t e r i o n f o r the evidence of a unimolecular reaction. I t usually has a value -6-13 around 10 J sec. f o r these reactions. This corresponds approximately to the frequency of the bond breaking. In the case of r a d i c a l decomposition i t i s assumed that the pre-exponential factor does not vary much from those of the molecules. For bimolecular reaction the pre-exponential f a c -tor i s close to the number of c o l l i s i o n s per ^ y s e c . and has -10 * , an average value A= 10 cm-ysec. Actually A= P^where P i s the s t e r i c factor, found by several workers to be - 3 - 4 10.r~10 f o r hydrocarbons. By comparing a unimolecular de-composition with a bimolecular one, of the same a c t i v a t i o n energy we get 3 19 n i s the number of molecules per cm = 10 Thus RcflO* andR j-10^ R2« The unimolecular reaction i s much more preferred than the bimolecular process. The u n i -molecular i n i t i a t i o n process w i l l thus be of major importance i n t h i s study. A molecule can s p l i t into r a d i c a l s either by passing through an activated complex unimolecularly, or by c o l l i d i n g with other molecules, bimolecularly. Unimolecular reaction i s favored at lower pressures where energy loss by c o l l i s i o n s i s low. In these cases molecules can use a l l t h e i r k i n e t i c energy and pass the energy b a r r i e r leading to decomposition. Heterogeneous phenomena w i l l be important since the energy needed for a heterogeneous reaction i s lower than that f o r homogeneous. Heterogeneous reactions Heterogeneous i n i t i a t i o n was f i r s t studied by TRIFONOV (23Y f o r the reaction 01 2 + H a Photoflh> J?CLH fo r vessels with different diameter. At low pressures and low diameters (14, 27 mm) he found the r a t i o of the rates R 2 / R l o C d 2 / d l 2 which relates the r a t i o of the rates to the rates of d i f -fusion to the vessel. At higher pressures, dependence of observed rates on diameters was s t i l l , but at higher diemeters no regularity existed. The rates are higher f o r larger surface because less energy i s needed. The heat of reaction i s given A H = D(C-C)-Q where D(G-C) i s the bond di s s o c i a t i o n energy and Q the heat of adsorption. Chaykin (20) using the concept of d i f f e r e n t i a l calorimetry by which one can measure the heat of reaction by a d i f f e r e n t i a l thermocouple both i n the center of the vessel and on the w a l l , proved that generation and termination of the chain take place on the walls i n the thermal reaction. H 2 4.Cl2 —* 2 HCL The hypothesis of heterogeneous - homogeneous c a t a l y s i s f i r s t proposed by P0LYAK0V (24), found application to the i n v e s t i -gation of oxidation reactions. The c a t a l y s i s depends on both the nature of the w a l l and the nature of the compound. -8-Actually the p r o b a b i l i t y f o r the capture of a r a d i c a l or atom by a surface s i t e was found by several workers (25-26, to be - 5 27,28)Abetween 10 and 1. In the case of w a l l reactions SEMENOV (20) d i s -cusses the p o s s i b i l i t i e s of bond termination between the free r a d i c a l or atom and the unbalanced charge of the surface. He correlates the magnitude of heterogeneous reaction with the required energy f o r electron transfer i n the surface with formation of a surface with unbalanced charges.If u i s the p o t e n t i a l energy of such a surface then homogeneous i n i t i a -t i o n i s favored i f u i s more than D(C-C) and heterogeneous i s favored i f u i s less than D(C-C). Heterogeneous i n i t i a -t i o n , even i f not marked^ at fecest appears to promote the i n i t i a t i o n action of impurities. POLTORAK and VOEVODSKII (29) proved that acceleration i n the cracking of propane at low pressures by oxygen strongly depended on the pretreatment of the vessel. Many workers have t r i e d to show the existence or not of w a l l effects i n hydrocarbon decomposition. Many re s u l t s are contradictory. In most of them the same basic error has been made. They used two d i f f e r e n t vessels, of the same size one packed and the other unpacked. But i t i s not the which size-e-f- i s the important factor, but the number of active centers. I t i s possible that t h i s difference some times i s so important that, packing may merely equalize the number of centers. LAIDLER and WOJCIECHOWSKII (30,31) t r i e d to show surface effect i n ethane decompositions with negative r e s u l t s . They found rates to be a l i t t l e decreased by an increase i n S/V r a t i o s . PURNELL AND QUINN (32) found no effect f o r differe n t w a l l coatings f o r ti-butane. ' In cnntrast with the above authors WALL AND MOORE (33) i n the mixed py r o l y s i s of C2R5 and > following the re s u l t s by mass spectrograph*^ analysis, have found evidence of c a t a l y s i s by the surface f o r packed vessel. This work along with the studies of the Russian investigators mentioned above^ are i n contrast with the previously mentioned. To determine wether or not there i s w a l l effect ;was one of the main tasks of t h i s present work. • . -I n i t i a t i o n takir$|place on the wall} i s expected to increase rates more at low pressures than at high pressures. But termination on the walls i s expected to favor more high pressures than low, because for the high pressure, the chain length i s larger than at low pressures. As f a r as the propa-gation i s concerned i t can take place by three main mechanisms Substitution reaction % R^ + R2R3—?-R^R^+ R3 (a) decomposition of r a d i c a l s S R A^Ag—> R+A^= A2 (b) and isomerization reaction ; CH3CHCH3—* CH3CHJPH2 (.c ) Reaction (b) and (c) are unimolecular and (a) bimolecular. For reaction (a) the a c t i v a t i o n energy can be given as R-L+ R 2R 3—> R-^+Rj E A = AH f(R!R 2) +AH }(R 3) - AHj(Ri) - A H ^ R ^ ) Enthalpies of fomation as well as heat capacities are we l l -10-known f o r molecules but not f o r r a d i c a l s . TocciHy w e use the values of Mf f o r radicals, found for lower temperatures, assuming they remain the same, or that, we can calculate them, by using the l i t e r a t u r e and Voevodskii's empirical formula. (33). 3>(C-H) p r i m # = D(0-H) a e c4. B = D(C+-H) t e r t >* 2 B This method would be followed i n connection with thermochemical data and d i s s o c i a t i o n energies found experimentally to calcu-l a t e a c t i v a t i o n energies f o r i n d i v i d u a l steps i n the proposed mechanism. For reaction (b) more than one -mo^ k of decomposition can be accepted. The choice between them i s a function of t h e i r a c t i v a t i o n energies,fteguencyfactors, etc. Their r e l a t i v e s p r o b a b i l i t i e s commonly are based on product analysis. Methyl and ethyl r a d i c a l are l i k e l y to be present i n the reaction vessel i n s i g n i f i c a n t amounts, .since the energy of a c t i v a t i o n i n these decompositions i s small. Isomerizations as i n reaction (c) are very important, as f a r as the most favorable angle of attack i s concerned. Rough quantun mechani-c a l calculations (20) lead to the q u a l i t a t i v e r e s u l t that the attack on a C-bond i s more favorable when the approaching r a d i c a l and the bond attacked, are collitofemT. These c a l c u l a -tions indicate that twice as much energy i s required f o r perpendicular attack as f o r a l i n e a r attack. In the case of a TT-bondja d i r e c t i o n perpendicular to the bond axis i s the most favorable. In isomerization reactions, which diff e r ' from C-H or from C-C bond s h i f t s and forms a pair :" ".--11-with the free valence electron, forming a new TT - "bond. At the same time the 6"-bond i s ruptured and the molecule assumes the most favorable configuration f o r the attack. The problem of isomerization of free r a d i c a l s , as common property, has been discussed (34,35). Isomerization processes w i l l be considered as important i n the following study. The i n h i b i t e d reaction As we have seen before, the decomposition of hydrocarbons follow a free r a d i c a l mechanism. The compo-s i t i o n of the products i s a function of the chain length which i s given by the r e l a t i v e importance of the propagation against the termination steps. For t h i s reason the product d i s t r i -bution should be a function of pressure since the chain length depends on pressure also. The recombination of r a d i c a l s can cause the in t e r r u p t i o n of the chain. The longer the chain the higher the rate because propagation reactions are very f a s t , having low a c t i v a t i o n energies, while the i n i t i a t i o n i s rather slow. The in t e r r u p t i o n of a chain can be done not only by r a d i c a l recombinations '.but also by other ways. The capture of the propagating r a d i c a l by an i n h i b i t o r i n e v i t a b l y stops the chain. As i n h i b i t o r s we use either molecules which are ra d i c a l s i n themselves such as NO with an unpaired electron, or compounds such as o l e f i n s . The i n h i b i t i o n by o l e f i n s i s ;• often termed"" - s e l f - i n h i b i t i o n ) since these compounds appear i n the products, and t h e i r effect i s obvious, from the shape of either pressure-time curves or the behavior of products with time. -12-N i t r i c oxide i n s l i g h t amounts suppresses markedly the rate of reaction, and a f t e r that a l i m i t i n g rate i s a t t a i -ned. The nature of the residual reaction was a point of con-troversy, f o r many years, and even now i t cannot be said that i s completely known. One point that i s established i s that f o r each hydrocarbon there exists under c e r t a i n conditions, a d e f i n i t e l i m i t i n g i n h i b i t e d rate of decomposition, depending only on the addition of s u f f i c i e n t amount of i n h i b i t o r , but independent of the nature of the l a t t e r . A thorough i n v e s t i -gation of the i n h i b i t i o n f o r a l l hydrocarbons under the same experimental conditions could give information on whether fir i n h i b i t i o n i s related to the nature of the hydrocarbon vto the conditions of the experiment. Hinshelwood and h i s co-worker's supposed that chain reaction disappears completely and that residual reaction i s simply a molecular reaction(36). Studying the i n h i b i t i o n of bttfcttrit by NO they proposed the following scheme. C4HJD-KJ_5> 2 R (a) RJ-C^HXO -^ 2-> R +.products (b) 2 R x (c) R * NO s (d) 5 l l * R + NO (e) I f (d) and (c) come into equilibrium at once then they w i l l not affect the stationary concentration of R and ho i n h i b i t o r y effect w i l l be observed. This has been also assumed by ECHOLS and PEASE (37) f o r the l a t e r stages of the reaction, when i n h i b i t i o n i s no longer e f f e c t i v e . Why t h i s equilibrium should occur i n the l a t e r stages and not i n the beginning i s -13-open to question. But Hinshelwood adds another step for the beginning of the reaction to explain the i n h i b i t i o n . He as-sumes that S i s destroyed i r r e v e r s i b l y and the rate d i f f e r s only by a constant by t h i s addition from that obtained i f reaction (d) i s also i r r e v e r s i b l e . I f equilibrium i s attained slowly there i s some explanation of the i n h i b i t i o n on the beginning'but s t i l l the problem i s not clear. Oti*r. hypothesis that at maximum i n h i b i t i o n the complex S, regen-erates r a d i c a l s as fast as i t destroys them, i s considered i n v a l i d , since addition of more NO should have some effect on the amount S and of course on i n h i b i t i o n , but t h i s i s not found to be the case. The solution of the above set of equation i s given by ECHOLS and PEASE (37) as follows: d[R/ d t = k-L [C 4 H i q j - k 3 t R l - K 4 [ R ] 6 K ) > k 5 £ 0 = 0 ( l ) IRJ= k l C c4 H10l+ V S ] ^ 3 ^ 4 [NOl (2) d C 5 J / d t = k4(R]tNQ)- k ^ S l (3) but 5 i s not stationary so sohin^(j>) and using (2) xve get ^ J / d t . k 4 ^ C 0 4 H 1 0 > ^ A ? + k m j _ ^ and substituting (4) back into (2) we get [R]. k i p 4 H 1 0 y k 3 . k 4 f N 0 J + R l k 4 C°4H10l P 0 ^ ( k , + k 4 BfO]) ( ( 1 . e " k 3 k 5 t A ^ 4 [ H O ] ) ( 5 j Prom (5) at t « o DO. ^ [ o + H ^ ^ ( 6 ) 14-Thus the concentration of R i s dependent upon the concentr-t i o n of NO. At larger values of t increased NO which means that the more NO we add the more abrupt i s the reduction i n R. No l i m i t i n g rate can be con-sidered i n that case and even i f we suppose that the t o t a l concentration of R i s suppressed by considerable amount the mechanism does not explain why more r a d i c a l s are not generated. Also, assumption of a long reaction time as an explanation of the l i m i t i n g reaction rate i s uncertain. On the basis of these ideas Hinshelwood introduced the concept of the apparent chain length. This chain length was assumed to be given by the r a t i o -The formula i s based on the hypothesis that a l l chains are suppressed by NO and the residual reaction i s a molecular one. However chain lengths were found by other methods to be much longer than those predicted by the above formula. For i n h i b i t e d reaction chain length was also found to be a function of pressure (38). By cai^'t^the rate of uninhibited reaction To> the rate of the f u l l y i n h i b i t e d too and t the rate f o r any NO concentration the r a t i o Immediately two questions a r i s e . From £^g»i-(6) the concentration of R i s always decreasing by rate i n the absence of i n h i b i t o r Tate i n .the pres.erfc.e of', maximum amount of i n h i b i t o r -2- t -15-the l i m i t i n g concentration i n HO as does also.the simple pl o t of f ^ i n i t . against P$T0). Hinshelwood and h i s school, at-tach much importance i n the v a r i a t i o n of t h i s r a t i o , with NO by connecting the v a r i a t i o n of chain length with NO concentration, but l i t t l e information comes from these relationships so f a r as the r e a l nature of i n h i b i t i o n i s concerned. They found (39) that where y i s a constant diff e r e n t f o r different hydro-carbons, Hinshelwood^" ntrying to explain the fact that the products appear' to be i d e n t i c a l i n the normal and f u l l y i n h i b i t e d decomposition suggestedthat two separate free r a d i c a l mechanisms might occur simultaneously, one suppressed by NO and the other remaining (40) This hypothesis i s a r b i t r a r y . The property of chains i s to be suppressed and there i s no obvious reason as to why one chain reaction should be suppressed and the other not. The second hypothesis which comes from the Rice-Herzfeld mechanism has been used by several workers Ctoui4 has bcommon points with the previous one.tfs a. general mechanism, the following i s proposed: % — > R i + R 2 (1) R l + M 1 - ^ B 1 E + E 2 ( 2) R 2 —5 R i * M 2 (3) R 1 +R 2 — ( 4 ) R.j+ NO — ^ Products (5) R 2* NO — ^ Products (6) -16-2 % ->M4 " ( 7 ) 2 R 2 ym5 (8) M 1 ?products (9) ! To solve such a system we need some s i m p l i f i c a t i o n s . Two p o s s i b i l i t i e s e x i s t . Chains can be broken either by R-j_ or by R2. I f they are broken by R 2 M-L —^ R-j+Rg (1) R + M]_—* R1H*R2 (2) R2—*-R1+M2 (3) R 1 +,R 2 9 R 1R 2 (4) R 2 4. NO y Products (6) Mi 9- Products (9) This case was examined i n d e t a i l by Hobbs (41) with the 1 following r e s u l t s , assuming the steady state: D ^ = *LP * L > K 3 P 2 ] - ^ M f t t f - k 4 ^ l ] fe3 = 0 d M = k l p y * k 2 [%][%] - k 3 [ R 2 ] - k4pt£J|R3- k 6[R2l[N0j d-t Where [ R j = k 2k 6[N0]+ ( k 2 k 6 [ N O j 2 + 8k2k 4(k!k6 NO + 2k 1k 5jj i 4 k 2 k 4 Reaction ( l ) i s considered n e g l i g i b l e i n comparison with reaction (2) rate and by using the value f o r the concentrate of R]_ from the previous formula we get f o r the rate - 1 7 -rate = k^M-J * k 2 ^ ^ ^ k g prcTp ^  8 k 2 k 4 k^pTO/* k 2 k 6 HTO] j where, 0 = k ^ > ^ [ M ^ / ^ ( l 6 k ^ k ^ f and = kg [Mj Hence ^ - z ^ / ^ . . ^ = (a £N0] * B[N0j+<l) 2- a/NOj where a, and b are the combinations of various values of k. The second case, which suggests that chains are broken by has been investigated by HOBBS AND HINSHELWOOD (42) and l e d to a r e l a t i o n analagous to the above In these two cases stationary state conditions were assumed and t h i s i s open to question. The concentration NO must remain the same, i . e . no consumption, and the nature of the r a d i c a l involved requires consumption of NO. TOMSON and NEISSNEZ (43) photographed the absorption spectrum of mixtures of NO with organic components during decomposition. Other workers (44,45) proved the existence of CH^ NO and CH2!s NOH, i n products of decomposition i n the presence of NO. The formation of HNO i s considered probable (46) a :%nat NO reacts with some ra d i c a l s must be considered as cer t a i n . The most c r i t i c a l point a r i s i n g from the above suggestionsis the following. Is the residual reaction a mole-the cular one or not? F i r s t of a l l products are considered to -18-to be i d e n t i c a l f o r both i n h i b i t e d and non-inhibited reaction. (12) Poltorak found i d e n t i t y i n the products of cracking of C^Hg, but from h i s table i t i s obvious that some v a r i a t i o n exists i n h i s r e s u l t s although he presents r e s u l t s f o r only the uninhibited and f u l l y i n h i b i t e d reaction. Hinshelwood ascribes t h i s supposed 'identity to be a coincidence. In contrast with Hinshelwood,"Rice and Rolley (47) as we l l as G-olanskii (48)-use the hypothesis that molecules capable of ending chains are at the same time able to regenerate them. The explanation makes molecular reaction improbable but i t i s not e n t i r e l y s a t i s f a c t o r y by i t s e l f . Whf^fHe same i n h i b i t e d rate exist f o r d i f f e r e n t i n h i b i t o r s ? A direct attempt to check the idea of the molecular reaction was made by Wall and Moore (43)• They decomposed thermally a mixture of C2H5 and C2D6 i n both the absence of NO and i n the presence of 2.5$ NO. They found i d e n t i c a l products and among them H>>, D2 and HL\ Por the uninhibited reaction t h i s i s l o g i c a l but f o r the i n h i b i t e d reaction i f a molecular one takes place HD would not be ex-pected. Unfortunately, they did not extend the inv e s t i g a t i o n over a broad range of NO pressures, and Hinshelwood questioned the r e s u l t , since i n s u f f i c i e n t P$F0) was used (50). But accor-ding to the work of Hinshelwood himself i t i s obvious that maximum i n h i b i t i o n i s attained very quickly and f o r a small amount of NO, so even i f NO was not s u f f i c i e n t i n WALL and MOORE'S experiment the products ought to be dif f e r e n t and especially the amount of HD highly reduced f o r the i n h i b i t e d reaction. Poltorak and Voevodskii t r i e d to c l a r i f y t h i s -19-work (29). They studied p y r o l y s i s of C2Hg i n the presence of D2. The exchange of D2 could give information about the chain mechanism. A lack of exchange could validate the molecular reaction. Exchange of D was found i n both uninhibited and i n h i b i t e d reaction,,Rice and Varnerin (51) completed the previous r e s u l t s by decomposing CjpDg i n the presence of CH4. The CH3D/CB4 r a t i o was found to be i d e n t i c a l in*both i n h i b i t e d and uninhibited reaction. Prom theal3'°"information the problem of molecular reaction seems to be excluded. The problem of the identity of the products remains. Recent efforts to explain i n h i b i t i o n , and l i m i t i n g rates has - been made by Voevodskii (52) and / Voevodskii and Poltorak (29). They t r i e d to correlate heter^eneous phenomena with i n h i b i t i o n . They postulate the existence of i r r e v e r s i b l e decomposition processes on the w a l l . They showed that i n i t i a l rate i n the presence of small amounts of 0 2, strongly depends on the state of the surface. They treated the vessel with HP to increase the effect of the surface, and f i n a l l y decomposed C^Hg i n such a vessel i n the pressence of small amounts of 0 2. Por low pressures they found acceleration. S i m i l a r l y they found reduced effect of NO i n vessels with active centers. By covering the surface with several oxides, they found s i m i l a r r e s u l t s . The t r e a t -ment of the surface may not have been the most e f f e c t i v e . The active surface might have been increased but might have been decreased as w e l l . And no information about the S/V 20 r a t i o i s given. Voevodskii questioning the coincidence that a molecular reaction and a chain^couicPgive the same products introduced the suggestion that decomposition might be i n i t i a t e d and terminated on the walls. This leads to a reversal of ideas on i n h i b i t i o n , since the l i m i t i n g pftfo) might be related with the vessel and not to the hydrocarbons. The l i m i t i n g pressure of P^0)has been always given as percent of the hydrocarbon. At the same time Voevodskii accepts the existence of two heterogeneous chain i n i t i a t i o n s , one reversible and one i r r e v e r s i b l e but t h i s looks a r b i t r a r y and without much v a l i d i t y . A conclusion of Voevodskii's i s that no effect i s expected at higher pressure, which does not agree with the concept of heterogeneous termination. Recently some new material from a t h e o r e t i c a l point of view have been presented by l a i d l e r and Wochice-chowskii (53). This material .is related to the i n h i b i t e d reaction of C2Hg, a hydrocarbon exhaustively investigated, but i t i s rather hard to apply - • to more complicated systems. A free r a d i c a l mechanism i s proposed involving the p a r t i c i p a t i o n of i n h i b i t o r molecules. The i n h i b i t o r can react i n some H - abstraction reactions such as C 2H 6 + MO — ^ C2Hjj-+..HN0 Chain termination may also involve C 2H 5 + HNO — » C 2H 6 4- W These ideas applied to C2Hg thermal decomposition give the mechanism;-21 C2H6+ NO k J — ^ C2H5+ HNO C 2H 5 JE2^ C2H4+ H H + C 2H6 k3-» C2H5+ H 2 H + NO HNO C2H5+HNO C2H6+NO The rate i s expressed as: rate = (k nk 0k-zk „ / , i 1 2 3 - 4 / k - l k 4 ) ^ [ c 2 H 6 3 This mechanism leads to f i r s t order k i n e t i c s as found experimentally (54»55). The a c t i v a t i o n energy of d i s s o c i a t i o n of HNO i s 44 KCal/Mole. which leads to an ac t i v a t i o n energy f o r reaction .(l) of 56 KCal/Mole. E = i ( E-,_+E2+E3+E_4 -E_^-E 4j = = i(56+40+10+44-0-o)= 75 value which i s near the experimental value of 74 kCal/Mole. Eor i n h i b i t e d decomposition the i n i t i a l rate has been shown to have an order of 3/2 and t h i s i s explained by the addition to the above set of reactions of the reaction H+HNO k 6 > H2+ NO 22 JL 3/2 Hence: rate = k^ (^k^/k^kg) 2 [CgH^ An i n i t i a t i o n l i k e the above can be considered as reasonable f o r systems l i k e C2H5 or CH^ with high a c t i v a t i o n energy, but l i t t l e help seems to come from that point f o r more complex systems. The appearance of c a t a l y s i s by NO, i n high NO concentrations, i n the decomposition of CH^COCH^, provides some support f o r reaction ( l ) . But i n the case of CH^COCH^ the H-atoms are highly activated and abstraction reaction requires lower energy. Resume. The following r e s u l t s , involving controversies , are provided from a l l previous investigations: 1. The nature of the i n h i b i t i o n , heterogeneous or homogeneous. 2. The effect of the P^ q on the products. 3. The effect of S/V r a t i o on the products and rates and 4. The mechanism of both i n h i b i t e d and uninhibited reaction. An invest i g a t i o n of these problems was-- the main objective of the present study. The extensive use of a n a l y t i c a l r e -sult s should provide r e l i a b l e data. The use of the same vessel f o r both packed and unpacked • s.ystems: - would 23 eliminate possible effects of surfaces between vessels of the same S/V' r a t i o . The in v e s t i g a t i o n was extended to a l l the isomeric pentanes, to determine the possible effect of the structure of the hydrocarbons on the decom-po s i t i o n mechanism. -24-CHAPTER I I EXPERIMENTAL Materials and Methods 1) Hydrocarbons The hydrocarbons used were l i q u i d s (n-pentane, and i s o -pentane) and gases (neopentane and a l l the C^ -C^ , hydrocarbons which were used to cali b r a t e gas chromatography and infra-red spectra). A l l these hydrocarbons • i ' . from Research grades of P h i l l i p s Petroleum Company, B a r t l e s v i l l e , Oklahoma, U.S.A. and t h e i r p u r i t y as supplied varied from 99*5 - 9 9 . 9 $ . For further p u r i f i c a t i o n successive d i s t i l l a t i o n s were used, from trap-to-trap, using two temperatures,- 196° C, l i g . Ng, so that traces of v o l a t i l e compounds l i k e H etc. were removed, and - 78° C, ( s o l i d C0 2 and methanol, so that less v o l a t i l e compounds were kept in s o l i d condition). The hydrocarbons were f i n a l l y stored i n globes on the vacuum system. 2) N i t r i c oxide N i t r i c oxide was supplied i n a cylinder by the Matheson Company, East Rutherford, N.Y., i n gas phase. By successive the trap-to-trap d i s t i l l a t i o n A n i t r i c oxi<fe was p u r i f i e d from more v o l a t i l e impurities by condensing i t at - 1 9 6 0 C and from N0 2 which was kept i n s o l i d state at - 7 8 0 C. Special care was paid to i t s p u r i f i c a t i o n , because impurities l i k e N0 2, affe c t markedly the since i t decomposition of hydrocarbons^generates traces of 0 2 at high pressures. 3) Mercury dimethyl H g(CH 3) £ A standard sample from the American Petroleum I n s t i t u t e . -25-Pittsburgh, Penn., was used. I t was not p u r i f i e d further. Gas Chromatographic Equipment Columns Two d i f f e r e n t columns were used i n the analysis of a l l products. The f i r s t was a s i l i c a gel 60/80 mesh column, 2 meters long, used for the analysis of c 1 - c 2 hydrocarbons and inorganic gases. The other column used f o r c 2 " - c 6 compounds was a 30$ HMPA (Hexo-methyl phosphoramide) on 60/80 mesh columpack, obtained from Fisher S c i e n t i f i c Company. The container was copper tubing 3 meters long. Detector For the gas chromatography a thermal conductivity detector (2) was used, designed by Ryce et a l . The c e l l current was always kept at 90 mA. C e l l temperature and recorder s e n s i t i v i t y were constant for the whole inve s t i g a t i o n , at 44° C and 1 mv respectively. The s e n s i t i v i t y corresponding to 1 mv was the maximum possible. of The s e n s i t i v i t y was checked day by day, using a sample a i r . By keeping the c e l l temperature, dolumn temperature and c e l l current s t r i c t l y constant, the s e n s i t i v i t y was kept constant . , within the l i m i t s of experimental error. This was important f o r the r e p r o d u c i b i l i t y of the a n a l y t i c a l r e s u l t s . The temperature of the columns were kept constant, 0° C and 3^° C for HMPA and s i l i c a gel respectively. Infra-red analysis For the i d e n t i f i c a t i o n of the products, infra-red spectra were taken for a l l products, on a Perkin Elmer Model 21 instrument. The spectra of a l l compounds were taken i n the region of 650-^000 cm"1. - 2 6 -The gaseous sample was kept i n a glass c e l l with NaCl windows polished time by time to remove any deposited compounds. The c e l l was evacuated to about 10"^ mm and i t was possible to maintain the vacuum for 5-6 hours. Pyeene cement was used to seal the NaCl windows into the c e l l . By comparing the infra-red spectra of the products with the spectra of i n d i v i d u a l compounds and espe c i a l l y for absorption bands i n the finger p r i n t region, 1300-650 cm"1, the products were i d e n t i f i e d . The P y r o l y s i s Apparatus  The vacuum system The apparatus used was a conventional s t a t i c vacuum system shown schematically i n F i g . 3 . The hydrocarbons being studied were stored i n globes, G, and could be admitted to the quartz reaction vessel, V, through a stopcock, S-^ . The pressure i n the reaction vessel could be read on a monometer, M. Samples for gas chromat-ographic analysis were withdrawn into a sample pipette, P, through stopcock S2« A Toepler pump, B, was used to transfer a l l hydrocarbons into the s t i r r i n g v e s s f l , K, driven by a magnetic s t i r r i n g motor, for premixing where necessary. The globes were f i l l e d by freezing the hydrocarbons from cylinder into a trap T^, from which a i r was removed. Trap-to-trap d i s t i l l a t i o n from T-^  to T g could be used for further p u r i f i c a t i o n where necessary. Liquid compounds were transferred to the globes, by freezing into a cold-finger, I . Reaction vessel and furnace The reaction vessel consisted of a c y l i n d r i c a l quartz vessel with a volume around 300 cm^. I t was placed i n a brass cylinder wound with heating wire and covered with cement. Several Inches of D I F F U S I O N P U M P C O L D T R A P T O T O E P L E R B T O T O E P L E R B rIG. 3 SCHEMATIC OUTLINE OF THERMAL DECOMPOSITION APPARATUS-Front Sit -28-vermiculite provided thermal i n s u l a t i o n . The furnace was modified so that, the reaction vessel was removable, by removal of i t s upper section. A quartz c a p i l l a r y tube connected the reaction vessel with the vacuum system. A variable voltage was supplied to the heating element of the furnace by means of a con-t r o l l e d power supply, shown i n F i g . Two chromel-alumel thermo-couples were placed In d r i l l - h o l e s one i n the cover of the furnace at h a l f the length and on the l i d . A Honeywell Electronic M i l l i v o l t c o n t r o l l e r was used to control the temperature within one degree at 500° C. The reaction products were withdrawn a f t e r pyrolysis and were analysed on the HMPA and s i l i c a gel columns and by I.E. for t h e i r i d e n t i f i c a t i o n s . In F i g . 5 and 6 t y p i c a l gas chromatograms are shown. The spectra of i n d i v i d u a l hydrocarbons, as used for the i d e n t i f i c a t i o n of the products are given i n Appendix IV. Description of t y p i c a l experiment After pumping the reaction vessel, mixing system and pre-expansion volume to a "black vacuum", the substance being used to be . pyrolysed was admitted to the mixing system. The pressure was measured before the entrance of the compound into the mixing system, so constant pressure could be used. In the case of i n -h i b i t e d decomposition, the preceding procedure, was performed f i r s t for the i n h i b i t o r which was used i n r e l a t i v e l y small amounts. The apparatus was pumped down again, and the hydrocarbon was admitted to the vacuum system. By the same procedure, the hydrocarbon, was pushed into the mixing system, where an homogeneous mixture was formed. The p a r t i a l pressures of each species has been pre-H O N E Y W E L L MILLIVOLT C O N T R O L L E R n i i i Lx_ s FIG. 4 FURNACE F U R N A C E H E A T E R R H E O S T A T T H E R M O -C O U P L E S S E L E C T O R S W I T C H M E A S U R I N G P O T E N T I O -M E T E R S HEATER CIRCUIT L / F i g . 5 : Chromatogram with HMPA 30#6Q/8(fcolumn of product mixture for inhibited decomposition of isopentane. V5) -32 -measured, before i t entered the mixing system. The pressure of was hydrocarbon was kept constant while the i n h i b i t o r A v a r i e d . After the evacuation of the system, the mixture was admitted into the reaction vessel. A manometer d i r e c t l y connected with the reaction vessel, was used to measure the i n i t i a l pressure, the f i n a l pressure, and the v a r i a t i o n of pressure with the time. After the desired extent of decomposition, the products were expanded from the reaction vessel to an evacuated f l a s k which was connected through a two-way trap with both the vacuum system and the reaction vessel. The sample i n the evacuated f l a s k was used for the analysis. The gas chromatography apparatus was evacuated to a less degree than that of the reaction vessel. By using a Toepler i n the case of small pressures of samples the whole sample could be used f o r analysis. S i l i c a column, used for analysis of c i " G 2 c o m P o u n d s a n d a f t e r that i t was heated, to 150° so compounds l i k e hydrocarbons, and NO , which were kept i n the column, were removed, and the column was keeping i t s e f f i c i e n c y . Helium was used as c a r r i e r gas, and the temperature of the c e l l and column was always above the max-imum possible room temperature, to avoid any ef f e c t . The helium was p u r i f i e d by a system of two charcoal traps, and two traps of CaClg. A reducing valve could regulate the available pressure. The flow rate was kept constant for each column, so that e f f e c t s due to di f f e r e n t absorption were avoided. - 3 3 -CHAPTEE I I I RESULTS I) Uninhibited pyrolyses The pyrolyses of the isomeric compounds isopentane, n-pentane and neo-pentane were done at temperatures near 500° C. The following investigations were made for the uninhibited or normal decomposition. 1) The effect of v a r i a t i o n i n i n i t i a l pressure of hydrocarbon on the rates and on product., d i s t r i b u t i o n for constant extent of reaction. i i ) The effect of v a r i a t i o n i n extent of reaction on product d i s t r i b u t i o n for constant i n i t i a l pressure of hydrocarbon. i i i ) The effect of v a r i a t i o n i n surface to volume r a t i o i n the vessel on rate of reaction and on product d i s t r i b u t i o n at various temperatures. A) Rates of Reaction The pressure-time curves ( F i g . 7) do not have a sigmoid shape but show that reaction rates f a l l a fter a certai n reaction time. This can be attributed to s e l f - i n h i b i t i o n by the products. Py r o l y s i s at d i f f e r e n t pressures showed the same feature but i n varying degrees. Isopentane, neopentane, and n-pentane have curves of s i m i l a r shape. A plot for isopentane of dp/dt for both o v e r a l l rateft(over) ? and i n i t i a l rate, (R3),against time f o r various pressures (Fig. 8) shows a sharp reduction with time for the higher pressures, but the reduction i s very s l i g h t f or the lower pressures. Rates were calculated from pressure measurements and not from a n a l y t i c a l r e s u l t s . I t has been established by other workers (Hinshelwood etc.) 2 4 6 8 1Q 12 14 16 time min P i g . 7: Variation of pressure with time for isomeric pentanes. that pressure measurements give an accurate measure of reaction rates i n p a r a f f i n p y r o l y s i s . This method was chosen because of i t s convenience and because i t gives a more uniform rate curve. Both o v e r a l l and i n i t i a l rates were used i n c a l c u l a t i n g rate constants, energy of a c t i v a t i o n and frequency factors. The var-i a t i o n of rate with pressure was measured for pressures ranging from 10 to 200 mm at several temperatures. A> 1 ; Isopentane In the case of isopentane the plot of o v e r a l l rate against i n i t i a l pressure, was l i n e a r f or pressures ranging from 100 to 40 mm ( F i g . 9). Some curvature was observed at higher pressures. At lower pressures the re l a t i o n s h i p Rate ( o v e r a l l ) s a P i i s obeyed, while for high pressures the r e l a t i o n s h i p becomes Rate ( o v e r a l l )saP1+-bE|. Pyr o l y s i s at four different temperatures from 502 to 530° gave curves of s i m i l a r shape. Extrapolation of the l i n e a r portion shows that the curves diverge from l i n e a r i t y very rapidly as (57) pressure increases. Partington et a l a t t r i b u t e t h i s increase to a simple k i n e t i c f a c t o r , indicated by the second expression for Rate(overall) given above. This re l a t i o n s h i p indicates that second order mechanisms are more important at higher pressures. The plots of R^ against P are approximately l i n e a r ( F i g . 1 0 ) . Variation of rate with c o l l i s i o n frequencies A c o r r e l a t i o n of reaction rates with c o l l i s i o n frequencies at each pressure used has been attempted i n F i g . 11. R. i s 1.5 h P (isopentane ) F i g . 9: Variation of R 0 v e r a l l w i t h I n i t i a l pressure for isopentane. -40-p l o t t e d against Z^/Z^ mm, 14- mm "being the lowest pressure at which decomposition was done. The curve shows an i r r e g u l a r portion. The r a t i o of c o l l i s i o n number Z./Z was found to be given 1 In-d i r e c t l y by the pressures r a t i o Pf/Pf^ c '-v • ':• '• T n e rate at lower pressures appears to be enhanced r e l a t i v e to z J / Z I / J . » Then an almost horizontal region, i s followed by a regular increase of rate at higher pressures. The lower portion indicates that at low pressures the rates predicted by c o l l i s i o n a l processes are much lower than the experimental rates. Heterogeneous i n i t i a t i o n may be responsible ,for the enhanced rate, since the required energy of a c t i v a t i o n i s less for the heterogeneous than for the homogeneous reaction. At high pressures both heterogeneous and homogeneous processes w i l l occur. At low pressures, i n e l a s t i c c o l l i s i o n s take place on the w a l l s , leading to decomposition of molecule, while the c o l l i s i o n which takes place i n the gas phase, are largely e l a s t i c at low pressures. The available energy i n these c o l l i s i o n s i s transferred more to the t r a n s l a t i o n a l mode than to the v i b r a t i o n a l mode, so that the energy for decomposition of molecule i s i n s u f f i c i e n t . The effect of packing To evaluate the effect on rate of packing the vessel was packed with quartz rods, so that the surface to volume r a t i o of 1.2 for the empty vessel became 2.5. The temperature used, was low, 4-90° C, so that, the regular reaction would be slow and any increase due to the packing would be marked. The ov e r a l l rate was increased for both low and high pressures but increase was greater for high pressures. ( P i g . 12). The increase i n rate can be explained only on the assumption of heterogeneous i n i t i a t i o n and termination. -42-For higher pressures longer chains exist due to higher concen-t r a t i o n s , and consequently termination i s suppressed, but not for lower pressures. However heterogeneous i n i t i a t i o n a f f e c t s reaction at both low and high pressures. The net re s u l t i s that packing promotes decomposition at higher pressures to a greater extent than at lower pressures. A- 2 : n-pentane A plot of i n i t i a l rate against I n i t i a l pressure i s given -In Fig . 13. The v a r i a t i o n of i n i t i a l rates with pressure i s similar to that for isopentane. Effect of packing By Increasing the packing by a factor of 3 , with pyrex tubes, no regular effect was observed ( F i g . 14). Lack of s u f f i c i e n t data Is an obvious disadvantage i n t h i s connection. The temperature used was rather high, 520° C, i t might be that i n such a temp-erature w a l l effects are less important. I t i s also possible that the pyrex packing provided a much lower surface concentration of active centres than that provided by the quartz packing used i n the isopentane experiments. A 3 : neopentane A plot of I n i t i a l rates against pressures for neopentane Is given i n F i g . 15. I t i s very s i m i l a r to the corresponding figures for n-pentane and isopentane. Effect of packing By Increasing the packing by a factor of 3 , with pyrex tubes, no regular effect was observed. Some i r r e g u l a r i t i e s appeared i n higher pressures where packing appeared to slow down the rate. P(neo-pentane) F i g . 15: R< VS P. f o r neo-pentane - 4 6 -In the most cases--the} A P - t i m e v a r i a t i o n was the same f o r both packed and unpacked vessel. Lack of s u f f i c i e n t data, i s the main point i n t h i s section too. B ; Order of reaction Some i r r e g u l a r i t i e s were found i n the reaction orders which were more marked as one goes from n-pentane to neo-pentane. These i r r e g u l a r i t i e s seem to be a function of the branching i n the hydrocarbon molecule. Orders were obtained by p l o t t i n g log (dP/dt) against log (P^) that i s by using the common rate equation dP/dt-z KP21 Evaluations of log dP/dt for a l l compounds are given i n appendix I I I . B 1 ; n-pentane The order of reaction calculated for three d i f f e r e n t temper-atures was found to be unity ( P i g . 1 6 ) . Thus the decomposition of n-pentane seems to follow a f i r s t order mechanism for the pressure range 10 to 150 mm. Effect of packing By increasing the surface to volume r a t i o by a factor of approximately 3 , with pyrex tubes, the order for the packed reaction vessel was increased to 3/2 (Pig. 17) . These r e s u l t s were obtained by using only one series with a packed vessel and thus cannot be considered as completely r e l i a b l e . B 2 : Isopentane Some i r r e g u l a r i t i e s i n order were observed for isopentane (Pig. 18). For three different temperatures a value of nr» 1 . 1 2 was found, which agrees with the concept of f i r s t order decomposition but :the-^ points are rather scattered. The curve seems rather to consist of two parts, one for low pressures with n- 1 and another for higher pressures with n > l ( ^ 3 / 2 ) . - 5 0 -B 3 ; Neo-pentane In the t h i r d and more branched isomeric pentane the above i r r e g u l a r i t y i s more marked. ( F i g . 1 9 ) . The o v e r a l l order i s approximately unity but a d i v i s i o n at low pressures into n r l and at higher pressures into n - 3/2 i s apparent. Two d i f f e r e n t temp-eratures were used. Some second order effects might be caused by c o l l l s i o n a l factors c h a r a c t e r i s t i c of higher pressures, (see chapt.IV). Effect of packing One series only of runs was used as was i n the case of n-pentane. There was no effect of packing on the o v e r a l l reaction order. The order was again unity but the curve has the same shape as f o r the unpacked vessel, with a s l i g h t l y higher slope ( F i g . 1 9 ) . C : Energy of a c t i v a t i o n and frequency factors Energies of a c t i v a t i o n and frequency factors were calculated from the Arrhenius equation. The formula E r 2.303 (T,T ?/ ) log (K9/K-i ) K cal/mol 1 * T 2 - * i was used. S i m i l a r l y A was calculated from the expression K=Ae _ E A/ST. Details of these calculations are given i n appendix I I . Energies of a c t i v a t i o n as wel l as frequency factors showed a v a r i a t i o n with pressure reaching a l i m i t i n g lower value at high pressures. The E A f a l l s as the pressure increases. Such a v a r i a t i o n with pressure was found by other worker ^ 8 ) $Qr 2 and 3-methyl pentanes. One common point i n a l l isomeric pentanes i s that the frequency factors f or the lower pressure range agrees with the expected value for unimolecular reaction, A - 1 0 1 3 S e c _ 1 . Neo-pentane shows i r r e g u l a r i t i e s as far as t h i s value i s concerned. In the above case \ Log (P neo-pentane) F i g . 19: Order of r e a c t i o n f o r neo-pentane. we expect that, unimolecular decomposition should be predominant fo r low pressures and should also p a r t i c i p a t e at higher pressures. Provided that c o l l i s i o n s are rare at lower pressures, molecules r e t a i n t h e i r c o l l i s i o n a l energy, so they can e a s i l y reach the higher values required for a unimolecular reaction. Wall effects i n such cases are very important because the required energy i s reduced by. an amount app.'ro.x imat e l y of that:; .equal to the absorption energy. I f D i s the d i s s o c i a t i o n (C-C) energy, and Q i s the adsorption heat, which i s a function of the nature of the surface then the required energy "for.. -;. decomposition i s E (heterog) D-Q,. Several workers have used these concepts of unimolecular and bimolecular decomposition i n both phases, c o r r e l a t i n g the v a r i a t i o n of with the pressure, and t r y i n g to f i n d support, from t h i s v a r i a t i o n of E^. Cundall and Palmer (59) studying the isomerization of cis-butene-2 found analogous v a r i a t i o n of E A. At low pressures a complete unimolecular reaction appears to take place. For higher pressures they found side reactions and wa l l e f f e c t s . Rabinovitch and M i t c h e l l found the same va r i a t i o n for the above isomer-i z a t i o n . C 1 :. n-pentane Overall rates were used to calculate rate constants by drawing tangents to the Rate (overall) VS TP^ curves at two temper-atures only. A greater v a r i a t i o n i n temperature was not used because exact values f o r a c t i v a t i o n energies were not sought (given by other workers). The v a r i a t i o n with the pressure was the major in t e r e s t . The r e s u l t s are tabulated i n Table I . -53-Table I Variation of and A with P 1 for n-pentane I n i t i a l a c t i v a t i o n energy frequency pressure mm K cal/mole factor $ec"J 10 66.5 3.08 x 1 0 1 5 20 66.5 3.08 x 1 0 1 3 30 66.5 3.08 x 1 0 1 3 50 66.5 I . 6 7 x 1 0 1 2 100 56.5 1.1 x 1 0 1 1 120 55.5 8 .6 x 1 0 1 0 Average values are not to be used because they are considered as meaningless. The E^ determined by Hinshelwood et a l ^ 8 ) r a n g e a from 65.8 K c a l / at 50 mm to 63 .7 K °al at 100 mm. Prey (61) and Hepp v 'give the following expression for the rate constant. K-2 . 5 x 1 0 1 3 e - 6 l - 2 0 ° / a T sec-1 C !L : Isopentane In the case of isopentane three d i f f e r e n t temperatures were used. A s i m i l a r v a r i a t i o n of E A and A with i n i t i a l pressure was found as for n-pentane. The r e s u l t s are given i n Table I I . Table I I Variation of E A and A with P i for isopentane I n i t i a l a c t i v a t i o n energy frequency -pressure mm K cal/mole factor sec 10 64 . 1 3 . 3 x l O ^ 20 64 . 1 3 .3 x 1 0 1 3 30 64 . 1 3 . 3 x 1 0 1 3 40 64 . 1 3 .3 x l O 3 ^ 50 64 .5 3 .3 x 1 0 1 3 ....cont. -54-(Table I I cont.) I n i t i a l a c t i v a t i o n energy frequency _ L pressure mm K cal/mole factor Sec 70 59.1 1-9 x 1 0 i 2 100 55.6 2.23 x 1 0 1 1 150 56.8 3.82 x 1 0 1 1 200 56 .3 3.17 x 1 0 1 1 The EA'S determined by Hlnshelwood et a l (62) varied from 67.8 K c a l at 8 .45 mm to 70 (average) K c a l at 50-400 mm. The rate expression given by Prey and Hepp (^1) for isopentane Is K . 7 . 9 x 1012 e-58600/BT sec - J-. C 3 : neo-pentane Neo-pentane was i n s u f f i c i e n t l y studied. Energies of a c t i v -ation are not very r e l i a b l e . Values of A show very low values. For low pressure, for Instance A ^ I O 1 1 and Is s t i l l lower f or higher pressure. The lack of s u f f i c i e n t data unable us to make any comment for t h i s p e c u l i a r i t y . In any case the accuracy of the values of A, i s a function of the accuracy of the values E^. A v a r i a t i o n of E A with the pressure i s again profound. Results are given i n Table I I I . Table I I I Variation of E^ with the P^ for neo-pentane I n i t i a l energy of a c t i v a t i o n pressure mm K cal/mole 10 57 20 57 30 V 57 40 57 . 50 57 70 49 100 47 . c, 55 D. A n a l y t i c a l r e s u l t s The products of decomposition were analysed by gas chromatography, by using a s i l i c a gel 60/80 column f o r C'-^ -Cg, hydrocarbons, N i t r i c oxide, etc. and HMPA column f o r hydrocarbons. I d e n t i f i c a t i o n was completed where necessary by infra-red spectroscopic analysis. Dl: Isopentane The main products i n isopentane decomposition are CH^, i-C^Hg and C^Hg. C2H5 and C2H4 are formed i n lower amounts. C^Hg was detected only i n traces and i t was measur-able only f o r high pressures and high temperatures. By increasing both the pressure and the temperature some traces of cis-butene - 2 were found by I.R. analysis. The trans isomer was not detected. The high temperature of the reaction provides a highly energetic medium which favors the formation of the cis-isomer. Hydrogen was detected i n trace amounts only and at very high temperatures and pressures. The I.R. Spectra confirmed the i d e n t i f i c a t i o n of CH^ C2H4> C3H6> i-C^Hg, and cis-butene-2 but they didn't give any unique ch a r a c t e r i s t i c peak f o r C2Hg and C5H3. A complete mass balance was obtained as i s shown by the a n a l y t i c a l r e s u l t s given i n Table I T TABLE IV An a l y t i c a l r e s u l t s (mole.$) for the pyrolysis of isopentane P i =50 mm.T=510uC. Pi-73 mm.T=510°C. P i =100 mm.T=510uC, Products P/P=0.415 P/P=0.425 P/P=0.41 1-C5H12 47.1 48.5 42.2 C3H6 18.1 18.5 23.1 i-C/Hs 10.3 11.8 13.1 CHj 19.1 15.3 12.8 C 2 H 4 C 2 H 6 33.1 3.30 4 . 4 2.2 2 .2 3-9 CJHg traces traces traces 02: Total 9 9 . 9 ^ 99 .6$ 99.5$ The main products again appear to be CH^, Co H^> l~ C 4 Hg» C 2 H 6 a n d G2H2j, w i t h small amounts of Hg. Some traces of cis-butene - 2 were found for higher pressures and temperatures by only I.E. spectro-scopy. C 2 H 6 a n d C 2 H 4 a r e i n c r e a s e d i n comparison with the corres-ponding amount i n isopentane. The main difference with isopentane i s the production of 1-butene and not iso-butene. Mass balance for the analysis of the products i s given i n appendix H. D 3 : Neo-pentane Two main products appeared i n neo-pentane p y r o l y s i s : CH^ and i-C^Hg; C^H^ i n very small amounts. G / j , H l 0 w a s n o t checked, because neo-pentane had nearly the same elutio n time, so i f there was any butane i t was masked by the neo-pentane peak. Hydrogen was not analysed. Some very s l i g h t traces of compounds appeared some-times. The c h a r a c t e r i s t i c predominance of CH^ and i-C^Hg over the other products suggests a different decomposition mechanism from those of the other two pentanes. E : The v a r i a t i o n of products 1) Variation of the products with time The v a r i a t i o n of the products with time f o r a giv#n was measured at i n t e r v a l s for a given time period up to 1 hour. Results were obtained i n d e t a i l for n-pentane only;these result s are presented i n Pig. 20 for n-pentane. -58-The y i e l d s of CH^ and follow Anearly l i n e a r r e l a t i o n with time, while those f o r C^i and l-C^H g have more curvature. The production of C H increases markedly with time to a nearly 3 6 l i m i t i n g value. Corresponding r e s u l t s for Isopentane are given i n tables VI. Table VI Variation mole3 % of the products of decomposition of isopentane with time time min C 2 H 6 C 3 H 6 i - C 4 H 8 ; 26 60 120 7-5 13.3 1.39 2.98 1.04 2 .46 11 .4 2 0 . 4 25 7.6 13 .4 15.6 2) The v a r i a t i o n of Individual products with P^ The complete v a r i a t i o n of products was investigated with pressure i n both isopentane and n-pentane. In n-pentane C^H^ increases s t e a d i l y , 1-C^Hg i s nearly constant, while CHj^, CgH^ and C show increase i n the beginning and then reach a l i m i t i n g 2 o value P i g . 21 . For the isopentane series the siuation i s d i f f e r e n t . C^H^ and i-C^Hg show a s l i g h t but constant increase. C 2 H 6 increases steadily and C 2H^ f a l l s i n the beginning and then reaches a l i m i t i n g value. CH^ increases but not so rapidly as c 2 H i j , 22.o:'yb 3) Variation of product r a t i o with time The r a t i o s : C ^ / C ^ , I C ^ C ^ , iCg/i-C^Hg, C^/i-C^Hg, CH../C H and CH,,/C0R\ were obtained as a function of time for 4' 2 4 4' 2 6 d i f f e r e n t experimental conditions. The re s u l t s are given i n table VIII' for diff e r e n t i n i t i a l pressures of n-pentane at 5 0 9 0 C. T?520° C t-. 2 h i o h / / 40 80 120 160 200 P (isopentane) (mm) F i g . 22a": Variation of the decomposition products with i n i t i a l °' pressure f o r isopentane. -62 -Table VI IS Va r i a t i o n of product r a t i o s with time Pn-pentane = 76 * 1mm T - 509° C time n T T in T T min C 2*V C 2 H 6 IC 2/G 3H 6 t C 2 / i - C / f H 8 C V ° 2 H 4 C V ° 2 H 6 5 0.41 1.41 5 3.55 1.93? 0.79 13 0.68 1.1 - - 0.5 0.34 17 0.59 0.8 4.5 5.3 0.72 0.42 27 0.49 0.875 - - 0.765 0.37 36 0.66 0.870 3.46 4 O.8.74 0.58 44 0.685 1.04 - - 0.99 0.666 55 0.78 0.932 4.25 4.6 0.85 0.573 Pn-pentane - 100 i 1mm T= 509° C 4 1.02 3.3 3.21 - -9 0.66 0.97 3.1 3.21 1.02 0.67 15 0.91 1.1 3.4 3.08 0;34? 0.318 25 1.03 3.2 3.2 - -32 0.67 - - - 0.815 0.542 38 0.73 0.9 - 4.08 O.65 0.46 49 I . 0 7 3.64 3.76 1.19 0.592 60 0.85 1.1 0.94 0.795 Pn-pentane =57-: 1mm T - 509°C 14 0.61 1.66 - - O.98 0.6 19 0.71 1.21 6.08 5 0.72 0.515 25 0.76 1.13 4.25 3.78 0.744 0.556 31 0.71 1.1 - - 0.83 0.582 45 0.75 1.23 4.37 4.30 0.9 O.67 55 1.07 4.65 4.34 - -65 0.89 1.13 5.25 4.62 0.815 0.711? -63-From the above table we see that the r a t i o C^H^/C^Hg increases regularly with time. This regular v a r i a t i o n i s more c l e a r l y shown i n Fig . 23. The conclusion i s that the more extensive the reaction the greater the production of C2R4 r e l a t i v e to C^Hg and CH4. The r a t i o s C^/C^Hg and C2/l-C4Hg seem to be approxi-mately constant. This means that the mechanism leading to and C'4 products on one hand and to O2 products on the other, i s not affected by the extent of the reaction. Since, as w i l l be shown, the l a t t e r two r a t i o s are unaffected by v a r i a t i o n of pressure, i t means that the mechanism leading to the formation of these products vary with temperature only. 4) The v a r i a t i o n of product r a t i o s with P i I) CH/j/C2Hg r a t i o : The v a r i a t i o n of CH^/C^Hg with Pj_ i s s i m i l a r i n both isopentane and n-pentane decomposition but i t i s much more marked i n isopentane where i n the pressure range 30-200 mm the r a t i o f a l l s by a factor of nearly 2. (Fig. 24.) In n-pentane for the same range of pressure the f a l l i s approximately 1.5. (Fig. 25.) This difference i s explained by the fact that i n n-pentane 02*% appears as a main product. So an increase of pressure has less influence on i t than i n isopentane (see chap-te r IV.) Neo-pentane did not give any information about t h i s r a t i o since C2Hg could not be detected i n the products of the py r o l y s i s . I I ) The v a r i a t i o n of CpH^/OpHg with pressure The v a r i a t i o n of t h i s r a t i o with pressure f o r both cases was s i m i l a r to that above Fig . 26 and Fig . 27. The v a r i a t i o n i s again more marked f o r isopentane changing by a factor greater than 2 CM o '<2)P n-pent = 75+ 1 O P n-pent=99 - 1 T=509° C |(mm) (mm) Fig„ 23: V a r i a t i o n of ^ H g / C ^ r a t i o w i t h time f o r n-pentane 0 . 6 NO 35 CM O •40.4 35 O 0 0 . ? i 1 5 1 9 ° c 509° C i <y> ^0 80 120 1 6 o P(n-pentane) (mm) F i g . 2 5 : CPi^/C 2Hg VS P. i n U . f o r n-pent. 502° c I 1 I I 1 » 1+0 80 120 160 200 P(isopentane ) P i g . 26: V a r i a t i o n of CpH^/C wit h F 1 a t v a r i o u s temperatures f o r isoDentane. -71-while for n-pentane the r a t i o f a l l s by a l i t t l e less than a factor of 2. The v a r i a t i o n was si m i l a r at each of several temp-eratures. In the case of n-pentane only two dif f e r e n t temper-atures were used, since 4 d i f f e r e n t isothermals i n the case of isopentane removed any doubt about the v a l i d i t y of t h i s change i neopentane i s not included, due to the lack of a n a l y t i c a l r e s u l t s . The effect of packing The v a r i a t i o n of CgH^/CgHg with pressure was compared for packed and unpacked vessel, i n both n-pentane and Isopentane. For n-pentane the unpacked vessel was packed with pyres tubes to change the s/v r a t i o by a factor of nearly 3 , and i n the case of isopentane, the vessel was packed with quartz rods to change the s/v r a t i o by a factor of approximately 2 . 5 . In both cases no marked effect was i n the r a t i o s found. Points f a l l near a common curve which could be considered as v a l i d f o r both packed and un-packed vessel F i g . 28 and 29. At f i r s t sight i t seems that either packing does not affect C^H^ and C 2 H 6 o r ^ a f f e c t s both of them to the same extent. The problem w i l l be examined i n d e t a i l i n the in h i b i t e d reaction. The v a r i a t i o n of C2 H2j/ G2 H 6 w l t h temperature for constant pressure and with pressure f o r constant temperature i s given i n the following table 1.3 T^519° C AE= 20% 3 T= 490° C &Pr20>* a: CM unpacked 50 • 100 1^ 0 P(isopentane)( ram) F i r . 29: E f f e c t of packing on CgH^/CgH, r a t i o at v a r i o u s pressures f o r Isopentane. 1.3 T = 529° C &P=20 % 1.1 o . 9 0 . 7 h ° acked 0 . 5 U0 F i g . 2 9 a : I n f l u e n c e of packing on f o r n-pentane. 80 120 160 P(n-pentane)(mm) he v a r i a t i o n of C~H. /C H r with P ^ 4' 2 6 i Table VTII Variation of C^/C^Hg with pressure and temperature Temp ° C pressure mm 490° •502 509 519 529 30 - 1.55 1.5 2.31 40 - 1.47 1.8 1.88 2 .06 50 2 ? 1.63 1.28 1.76 1.53 70 1.2 1.42 1.26 1 1.14 100 1.32 1.28 1.64 1 .09 -150 0.69 1.05 1 0.89 -200 - 0 .92 1 .32 0.82 1.16 For 30 and 40 mm pressure there Is a regular Increase of C^ H^ /G with temperatures, while f or other pressures the values remain around a constant value, d i f f e r e n t for every pressure. Higher pressures seem to have more effect than higher temperatures, so no regular v a r i a t i o n i s observed with temperature for higher pressures. I l l ) The v a r i a t i o n of l-C;,,Hg/C^ H^  r a t i o with pressure This r a t i o was available only for isopentane. No marked effect of pressure on i t i s observed and the curves show a very s l i g h t increase F i g . 3 0 . C^Hg and i-C^Hg seem to be produced mainly by the primary process which i s unaffected by pressure i n contrast with CgHg and C^H^ which might be produced In secondary processes more than i n primary. In the case of n-pentane, the va r i a t i o n of the r a t i o s C H /1-C.Hp, I C 0 / C and T_C~/1-C,,H0 3 8 *r  0 ^ 3 6 c o with pressure indicates that C^H^ and 1-C^Hg as well as C 2 com-pounds come from processes which are unaffected by pressure. 50 l^n 150 P(isopent)( ram) Pig. 30: Variation of i-Cj^/C-H, r a t i o with i n i t i a l pressure for various values of 1T fo? isonentan^ Pressure lor - 7 7 -The values of these r a t i o s at various pressures appear i n the following table. Table IX Variation of C H^/l-C^Hg, ^ C 2 / C 3 H 6 3 2 1 ( 1 :£C2^1"CI^B r ^ t l o s w i t h pressure f o r n-pentane. Pn-pentane mm C 3H 6 /1-C 4H 8 I C 2/C 3H 6 22 - - -35 2.98 - -46 3.33 1.05 3.5 70 - - -98 3.9 1.38 ? 5.4 150 3 .4 1.13 3.86 203 3.9 1.2 4 . 7 F: Effect of chain branching on rate The effect of the structure i n the decomposition of hydro-carbons i s attempted i n P i g . 31 . The i n i t i a l rates of the three isomeric compounds obtained under the same experimental conditions, are plotted VS Pj_.< • The more the branching, the lower the rate. This i s i n agreement with r e l a t i v e data given by various i n v e s t i g -ators (H 8> 119, 120) for the' oxidation of par a f f i n s . This prob-. view lem i s approached i n chapter 4, from a more th e o r e t i c a l point. The difference i n effect can be attributed p a r t i a l l y to a s t a t i s -t i c a l base. Normal pentane with 6 primary and 6 secondary hydro-gens i s more vulnerable to the abstraction of H by a r a d i c a l than isopentane with 1 t e r t i e r y , 2 secondary and 9 |>fM&a\ry, and much more than neo-pentane with 12 primary hydrogens. The size of the molecule i s one more important factor, as f a r as c o l l i s i o n s are 50 4P/ri-pent ane iso-pentane neo-pentane P^L (;mm) 100 150 F i g . 31 : Structural effect of isomeric pentanes on the i n i t i a l concerned. The importance of molecular size and the c o l l i s i o n a l area i s explained i n chapter IV II) Inhibited reaction The thermal decomposition of a l l isomeric pentanes, i s in h i b i t e d by NO, yet differences appear from member to member. I n h i b i t i o n i n n-pentane decomposition has been investigated suff-i c i e n t l y by various workers i n the past. In normal "inhibition the rate f a l l s r apidly for a small amount of NO and f i n a l l y , approaches a l i m i t i n g value. By adding more and more n i t r i c oxide no more i n h i b i t i o n i s observed. A s l i g h t acceleration occurs at higher n i t r i c oxide pressures. The necessary amount of NO to reach the l i m i t i n g rate, has been found by other investigators to be about 10$ of the hydrocarbon. A : N i t r i c oxide concentration for the f u l l y i n h i b i t e d reaction 1 : Isopentane The effect of NO was examined at dif f e r e n t pressures of the hydrocarbon . to determine whether the l i m i t i n g P N 0 i s related to the hydrocarbon pressure. In Pig. 3 2 , the i n i t i a l rate i s plotted against PJJQ. A common l i m i t at P N 0 10 mm appears at a l l three pressures, or at least i t i s apparent that the l i m i t i n g pressure of NO i s independent of the p a r t i a l pressure of hydrocarbon. To determine whether or not the l i m i t i n g P N Q i s a function of the surface-to-volume r a t i o for the vessel, the l a t t e r was repacked twice so that the S/V r a t i o was increased from 1.2 to 2.5 and then 5 . 0 . By using constant amounts of hydrocarbon (101*lmm) the change of the rate with P N 0 was examined. A s h i f t of the l i m i t to higher values appeared with increasing packing Pig. 33 . For the three 16 14 Tr519° C 100 (miiy 0 - 80 (mm) ^ <D 50 (mm J 1 0 8 12 16 20 24 P(NO) (ram) 28 F i g . 32: Rate ( i n i t i a l ) VS p(NO )in unpacked r e a c t i o n f o r l^noootane. I 1 I I I ! I I I 0 4 8 lm 12 Ira 16 ?0 24 P(NO)(nra) l m F3g. 3 3 : ftoverall VS P(NO)at v a r i o u s S/V r a t i o s f o r isopentane. -82-Fi g . 34: Effect of P(NO$on rate(overall) for n-pentane;packed vessel. d i f f e r e n t S/V r a t i o s the l i m i t was approximately 10, 13, and 15 mm respectively. These numbers are not absolute, but what i s obvious i s that the following inequality holds. Lim. PN0(S/V=1.2) <lim.PN0(S/Vr2.5Klim.P 0(S/V=5.0) This the indicates c l e a r l y the relationship of i n h i b i t i o n process to the surface-area of the reaction vessel and shows the importance of heterogeneity i n the o v e r a l l i n h i b i t i o n . I I 1 N-pentane In normal pentane no comparison l i k e the above was made, because the l i m i t i n g P N Q for decomposition was considered as w e l l known from previous work, to be 10 mm. The vessel was packed u n t i l the S/V r a t i o became approximately 3. The change of the o v e r a l l rate against P N 0 i s given i n F i g . 34 which indicates, but not p o s i t i v e l y , that lim.P N 0(unpacked vessel) ^ lim.P N 0(packed). The fact that the vessel was packed with pyrex tubes and not quartz rods, might be an explanation of the reduced surface effect as compared with Isopentane. The surface a c t i v i t y of these two packings may^not be the same. I l l ) Neo-pentane Neo-pentane show differences from the other two members of the series. N i t r i c oxide has no s i g n i f i c a n t effect on the i n i t i a l rate of decomposition of neo-pentane. There may be i n fact an accelerating, rather than i n h i b i t i o n . The experiments done were i n s u f f i c i e n t to estimate t h i s effect with ce r t a i n t y . These r e s u l t s are given i n Table XV.. -84-Table XI Vari a t i o n of Ap with time for various P N Q for neo-pentane Pfreo-pent) ~ 76 2 1 mm T^540° c & P mm time min PN0° mm mm EN0 7 mm P N 0 10 mm mm PN0 19 mm 0 0 0 0 0 0 0 1 1.5 2 1 2 1.5 1 2 2.5 3 1.5 3 2.5 2 3 3.5 3 .5 2 3.5 3 3 4 4 3 4 4 3.5 5 4.5 5 3.5 5 4.5 4 10 7.5 7.5 5.5 7 6 11 7 I f decomposition i s extended to some 20$ of the i n i t i a l pressure, the o v e r a l l rate shows a d i s t i n c t I n h i b i t i o n . For two different pressures of neo-pentane, 75 and 100 mm the l i m i t i n g pressure of at NO appears to be fixed A10 mm. By packing the vessel with pyrex tubes so that S/V rate i s changed by a factor of nearly 3 , the l i m i t was shifted to approximately 12 mm, an effect s i m i l a r to that found f o r isopentane P i g . 35 . B) Variation of products with time The v a r i a t i o n of product d i s t r i b u t i o n with the time was investigated for the f u l l y i n h i b i t e d reaction for isopentane Pig. 36-41. The percentage of each product i n the reaction mixture varied l i n e a r l y with time, i n contrast with the uninhib-i t e d reaction. Assuming that a f u l l y i n h i b i t e d decomposition occured i n the l a t t e r case, t h i s difference can be explained, by s/v 3.0 O S/V 1.2 (3 S/V 1.2 (D Tc540° C packed 100 (mm) unpacked 100 (mm) J—unpacked 77 (ram) X 8 lm 12 16 PC NO) (mm) 20 24 Ncopentane : RQVeraU v $ ' P ( N 0 ) 1 I I I ! L 24.0 70 100 1™ 160 190 220 time (rain) i g . 36: Y i f l J , o f £2#H5NO and C 2 H £ w i t h time f o r isopentanes fully.- i n h i b i t e d ff /Tf z. 37:C 2H 4 VS time at various S/V ra t i o s and different p(No)for isopentane. F u l l y inhibited reaction. F i g . 38: i-C^Hg VS time at different pfNOJfor Isopentane. 37 28 24 v. 20 0) r-l O e t* 16 CO o 'a 2 12 P NO" 2.5(mm) -T c 5 1 9 °C P- i s o p e n t ^ 1012 1 (ram) o ---^ ^ ^ ^ r - ^ T ^ 11 i i • 0 30 60 90 1 2 0 time (min) 150 1.80 P i g . 39: Product y i e l d VS time i n the f u l l y i n h i b i t e d isopentane. r e a c t i o n f o r assuming that, i n the i n h i b i t e d reaction the effects of heter-ogeity were completely counterbalanced by the NO so the reaction followed a homogeneous mechanism, while i n t h e ^ n h i b i t e d reaction heterogeneous effects participated with the homogeneous reaction. By using high pressures of i n h i b i t o r , a nitroso compound was detected, and i t s i d e n t i f i c a t i o n has been attempted by i t s I.E. spectrum. I t gave an absorption peak at 695-700 cm"1. Liter a t u r e data on I.E. spectra * gave an absorption peak at 690-695 for a l l nitroso compounds. A second l i n e of evidence was obtained from the fact t h a t ; the unknown substance was eluted the from s i l i c a gel column i n which a l l hydrocarbons but C^-C2 and th e i r derivatives are retained. The substance appears to be either CELNO or C^HJTO. CH NO was excluded by pyrolysing J c 5 3 Hg(CH ) In the presence of NO. Analysis of the products on a o 3 2 s i l i c a gel column gave no peak at an el u t i o n time corresponding to that of the unknown compound. I t i s assumed that t h i s substance i s CgH^NO. C2H^N0 i s much more stable thermally than CH^ NO and i t s presence i s more reasonable at the temperature used i n the present study. In F i g . 36 the v a r i a t i o n of t h i s assumed CgH^NO with time i s compared with that of C2H^. At the beginning of the reaction C2HjjN0 i s favoured and C2H^ appears i n small amounts. As the extent of reaction increases, C2H^N0 increases very slowly, while C2H^ increases very r a p i d l y . The i n s t a b i l i t y of CgH^NO does not allow i t to Increase markedly with time. The actual increase of C2H^N0 i s by a factor of 1.5 over the time of reaction while C2H^ for the same period i s increased by a factor at least 9» C3 H6 ® C2H4 C2 H6 T = 529° C 1-n-pentane = 130+ 2mm| P = 20/. 28 8 a c •H o T-519 0 0 P isopent*1011 1 (mm) &P r 20# Q) S/V* 1.2 ± 12 16 p(N0) (mm) 20 24 28 32 F i g . 4 2 : e f f e c t ; o f NO i n the p r o d u c t i o n of C^H^ at various S/V r a t i o s f o r isopentane. • T r 519° c 1 8 10 12 p (NO)(mm) 14 16 k • 6 u -Pig. 43: I-C^H8 VS PfNO)in paoked and oked vessel for isopentane. I t \ I I I _ l L_ 0 4 8 12 16 20 24 28 P(NO;(ram) Fi g . 45; C2H£ VS P(NOjat various S/V ra t i o s for isopentane. 1 -99-In P i g . 37 the v a r i a t i o n of CgH^ with time at various S/V r a t i o s and P N Q i s shown. At constant reaction time the effect of NO and surface i s obvious. The curve corresponding to higher packing i s not s t r i c t l y a straight l i n e because the P N 0 used (15 mm) i s exactly the l i m i t i n g pressure for t h i s S/V r a t i o . In F i g . 38 and 39 the variations with time of i-C^Hg, C^H^ and CH^ are presented. At constant hydrocarbon concentration and constant P K 0 and S/V r a t i o , f or a l l products except C2H£(Fig. 40) the equation holds d(products) „K (constant) dt I n r e a l i t y P N o does not remain constant, as w i l l be shown l a t e r on. There appears to be a close r e l a t i o n s h i p between the production of C2H£ and the consumption of NO. C 2 H 6 a n < i C 2 H 5 W 0 a P P e a r to be formed i n competing reactions. C 2 H 6 i s f a v o u r e d a s t n e reaction i s extended. C) Variation of the decomposition products with NO and surface- to-volume r a t i o Previous workers postulated that the product dis t r i b u t i o n s with the i n h i b i t e d and uninhibited decomposition were i d e n t i c a l . A caref u l analysis i n the present work,by gas chromatography of more than 1000 samples obtained from pyrolysis under d i f f e r e n t conditions, from p a r t i a l l y or f u l l y i n h i b i t e d decompositions, proves that t h i s i s not the case. The concentrations of most products were found to be dependent upon the NO pressure and reached a l i m i t i n g value when the NO pressure exceeded approximately 15 mm for a l l isomers (F i g . 4 1 - 4 7 ) . The products are i n general the same for both the i n h i b i t e d and uninhibited decomposition, but with some differences. The y i e l d of saturated hydrocarbon i s reduced with the increasing of P N 0- Some differences are encountered for the three pentanes -100-as w e l l as some s i m i l a r i t i e s . C 3 H g i s decreased with P N Q i n the case of the case of n-pentane while i t i s independent of P N 0 in^iso-pentane Pig. 41 and 4 2 . C^Hg i s decreased i n both n-pentane and i s o -pentane but increases rap i d l y with P'NQ for neopentane (Pig. 4 1 , 4 3 , 4 4 ) . A c a t a l y t i c effect of NO on the production of i-C^Hg i s observed i n neo-pentane. A common feature for a l l isomeric pentane i s that C2Hg decreases with increasing P NQ a n ( 3 : that and CH^ increase proportionally with P N Q (Pig. 4 3 , 4 5 , 46 and 4 7 ) . The y i e l d of C„H, from isopentane (Pig. 42) appears j 6 independent of both P N 0 and S/V r a t i o while i-C^Hg appears indep-endent of S/V r a t i o but s l i g h t l y dependent on P J J O Both NO and packing seem to have no effect on the main mode of decomposition. In ( F i g . 45) the dependence of CgHg on both P N 0 and S/V r a t i o i s presented. C2Hg decreases slowly with P N 0 reaching a l i m i t i n g value, a f t e r the P N 0 passes the value for f u l l i n h i b i t i o n . The var i a t i o n i s clear for both packed and unpacked vessel. Packing seems to favour C 2 H 6 highly. An Increase of S/V r a t i o by a factor of approximately 2 , nearly doubles the y i e l d of C2Hg. The addition of NO was extended to some 25 mm, nearly 2 . 5 times the l i m i t i n g value to see whether a l i m i t i n g value of C2Hg i s reached. The y i e l d of C increases with P NQ to a l i m i t i n g value at high of P N 0 (Pig* ^7) An increase by a factor / tnearly 1.5 i s observed for both packed and unpacked vessel. Higher S/V r a t i o s favour the y i e l d of CgH^ more than C2H^. A change i n S/V r a t i o of 5 causes a change i n C2H^ of a factor°A"approximately 2 , h a l f of the corresponding effect for C"2Hg. The var i a t i o n of CH/j, i s shown, i n (Pi g . 46) with both P N 0 and S/V r a t i o City,, l i n e CgH^, increases -101-with P N O , for both packed and unpacked vessels reaching also a l i m i t i n g value for higher P N Q. The effect on CH^ i s less marked than for CJi^. H igher S/V r a t i o s favour the y i e l d of CHjj, but not as much as i n the case of CgH^ and CgH^. As a f i n a l which ,^ , point comes the- increasing surface area, favours the product formation i n the following sequence C2 H6> C2HZj.) C H2f •••• C 3 H 6 ' i" C4 H8 no effect Increasing the P N 0 favours the reaction products i n the sequence C2 H4> G H 4 > i " c 4 H 8 > c 3 H 6 C2H£ formation i s reduced by NO. In Pig. 44 the v a r i a t i o n of i-C^Hg and CH. vjith P J J 0 Is shown for neo-pentane. Both products with increasing PNO appear to increase^but the increase i s more marked i n the case of i-C^Hg. I t increases by a factor at least 2, while CH^ changes very s l i g h t l y . A peculiar point i s that i n spite of the fact that PJJO passes i t s l i m i t i n g value, for f u l l i n h i b i t i o n , l-C^Hg continue to increase regularly. Thus 1-C^Hg must ':.: r e s u l t , not froma free r a d i c a l reaction, which i s subject to i n h i b i t i o n , but f^emanother process such as molecular s p l i t t i n g unaffected by NO. In ( F i g . 48) the va r i a t i o n of CgH^ NO with NO i s shown for i s o -pentane. An attempt was made to increase the y i e l d of C2H^N0 by greatly increasing the NO pressure up to 75 mm. C2H^N0 i s favoured by increased PJJQ but not proportionately. This f a i l u r e to promote C2H^N0 formation s i g n i f i c a n t l y can be attributed to the i n s t a b i l i t y of C2H^N0 at the temperatures of 500° C and higher. -103-D) Variation of product r a t i o s with Pwn and S/V I) Variation of GpH^/G^H^ with 2 m and S/V r a t i o The change of the CgH^/CgHg r a t i o was considered as the most important r a t i o since i t indicated d i r e c t l y a di f f e r e n t behav-iour of both saturated and unsaturated products, with PJJO* This v a r i a t i o n i s s i m i l a r for both n-pentane and isopentane (Fig. 4-9, 5 0 , 5 1 ) . In both cases the r a t i o changes by a factor of nearly 2 and reaches a l i m i t i n g value when P NQ exceeds i t s l i m i t i n g pressure. The v a r i a t i o n of C2H/j,/C2Hg with both P N 0 a n d S/V r a t i o i s very important because i t gives an i n d i c a t i o n i n the in h i b i t e d reaction of the importance of heterogeneous processes. The r a t i o shows a v a r i a t i o n with P N 0 s i m i l a r i n kind, but d i f f e r i n g i n magnitude for three diff e r e n t values of S/V ( F i g . 5 0 ) . For con-stant PJJQ the C^H^/CgHg r a t i o i s r e c i p r o c a l l y proportional to S/V r a t i o . This behaviour indicates that increased surface favours C2Hg to a greater extent than CgH^. II ) Variation of CH,,/CoH6 In both n-pentane and isopentane CH^/C2Hg i s s i g n i f i c a n t l y increased by increasing PJJQ ( F i g . 52 and 5 3 ) . In the case of isopentane the var i a t i o n of CH^/CgH^ r a t i o i s reduced as the S/V r a t i o increases, at constant PJJO? - ThiS' i s another i n d i c a t i o n that increased surface, for constant P N 0 , favours highly the formation of CgHg. I I I ) Variation of CQH^+COH^/C^H^ wlth PN0 Variation i n PHO affects only s l i g h t l y the r a t i o C^^+CgHg/C^Hg (Fi g . 54) for n-pentane. The r a t i o increases s l i g h l y with P N Q. In contrast to/, iso-pentane the r a t i o i s constant for isopentane. \ 1 I 1 I I L I I 1 0 ^ 8 12' 16 20 24 26 32 PfNOj(ram) Fi g . 5 0 : C2H./ C-H, VS PfNOjat various S/V r a t i o s for isopentane.^ F i g . 51: C 2H 4/C 2H 6 VS p f N 0 ) for n-pentane. — © ^ s/v ~ l . g — •5^0 0 0 ^ ^ ~ s / v ~ 5 . o L 6/ P - lOltl(nm). , i-isopent - / %y &P^20$ i 1 1 1 1 1 1 1 8 12 16 20 P(NO) (ram) 24 28 32 F i g . 52: CH^/C2H6 VS p(No)at various S/V r a t i o s for isopentane. ) p(NO)(mm) F i g . 5^'^2nl^C2u6^CZE6 V S P(NO) f or«yj-pentane and isopentano. -110-IV) Variation of C J j i C o H V C H, with P H Q The v a r i a t i o n of CgH^-f CgHg/C^Hg changes slowly with P N Q for both isopentane and n-pentane ( F i g . £5). C^Hg i s of course d i f f e r e n t i n these two cases, l-C^Hg i n n-pentane and i-C^Hg i n isopentane. This difference i s unimportant because both products are a consequence of a CH^-C^H^ primary decomposition. V) Variation of Clfy/CgHj, with P The r a t i o was examined only f or$3:-pentane series. I t increased very s l i g h t l y with P N Q ( F i g . 56) i n contrast with the rapid change of both CgH^/CgH^ and CH./CgH,. I t appears that NO favours and CpHg both CH^so that the ratios remain nearly constant. VI) Variation of C^Hg/C^JTO with NO This r a t i o was examined for Puo^^NO ( l i m i t i n S ) because for lower PJJQ CgH^NO i s not detected. By increasing the P^Q "to high values formation of C HJO i s accelerated but steadily reaches a l i m i t i n g value (Fig. 5 7 ) . I t i s rather improbable that more NO does not give r i s e to the formation of more CgH^NO, but the concentration of CgH^ NO appears to reach a steady state. For higher extent of decomposition (&P 3 0 $ ) , the decrease i n the r a t i o i . e . the formation of CgH^NO, i s more pronounced. This. nitroso-compound was detected with n-pentane. VII) Variation of l-C^Hg/C^ with P M 0 and S/V r a t i o This r a t i o , which was examined only with isopentane, decreases were s i g n i f i c a n t l y with P^QJ "but n o regular variations^observed with S A r a t i o ( F i g . 5 8 ) . E) Consumption of NO In previous studies, NO, was either reported to be s l i g h t l y - i ' i 2 -o ho CM O o CM B E CM +« O r- l <D CD 4-5 C © c t-i A H CM CM 00 ho a a o CM 0) c CO +J c ft I R O •"•too CO > VO CM CN 00 CM CM CM r- l CO o 60 •H T =550° C p i - l s o p =100±l(rara) 1.6 — 1.2 • o K-X CM 0.8 — ^ 0 0 \ ^ \ \D EC CM O < 0.4 ^ \ 30$ tp 15% A P 0 1 1 1 l i l t 20> 30 40 50 60 70 80 90 P NO more than P(NQ)(limiting) (mm) Pig. 5 7 : C X / C H NO VS p(N0)at various reaction extent for isopentane. 8 10 FCNCX mm) 12 P i g . 5 8 : /C~H, VS p(NO)at various S/V r a t i o s f o r isopentane. -115-consumed, or not consumed at a l l . The chemical analysis used for NO were l i a b l e to considerable error, since NO i s very unstable i n the presence of 0 2. In the present study, .a very c a r e f u l estimation of NO was attempted for a l l p y r o l y s i s , by gas chromatography, using a s i l i c a l gel column. The fact that NO was consumed has been established f o r high concentrations of NO ( F i g . 5 9 ) . The a n a l y t i c a l r e s u l t s for NO pressures below 8 mm are not r e l i a b l e . 20-40$ of the n i t r i c oxide used, was found to be consumed i n many cases. Some of t h i s consumption may be due to traces of 0 2. However, the resul t s c l e a r l y show that the comsumption of NO increases with increasing P N Q 3 and decreases with increasing S/V r a t i o s . For small concentrations of NO, the consumption would thus be low. Above the l i m i t i n g NO pressure the consumption Increases ra p i d l y . For higher S/V rat lo$consump-t i o n i s lower. Thus, NO consumption, i s not a simple l i n e a r function of i n i t i a l PJJO« P O R pN0^ pN0 ( l i f t i n g ) , NO seems to be less active while f o r PJJQ^NO ( l i m i t i n g ) i t i s consumed rap i d l y . The f a c t that increasing . ' the surface, reduces the consumption, means that NO i s retained on the surface and consequently prot-ected from being consumed. Afte r the complete covering of the active centers of the surface, NO establishes a dynamic equilibrium with molecules i n the gas phase, being subject to consumption, > . . much more than those held on the surface. In a number of experiments, the extent of decomposition was increased from 15 to 30$. A s l i g h t increase i n NO consumption occurred. The consumption of NO for various i n i t i a l pressures of NO, and f o r a 6P - 2 0 $ i s given i n table XI. -116-pawnsuoo on mai -117-Table XI Consumption of N0$ of the P N 0 &P* 20$ isopentane n-pentane P N0 mm consumption %' P N 0 mm consumption % 9.5 29 4.7 36 1.4 27 5.1 32.8 5.1 5 K ? ) 5 34.3 10 16.6 7.45 34.8 12 43 8.1 28.9 8.5 37 3.72 30 7.43 40 -118-CHAPTER IV DISCUSSION A. Uninhibited Reaction A great deal of work has been done i n the past on the pyr o l y s i s of the pentanes and especially of n-pentane. In spite of t h i s , r e l a t i v e l y l i t t l e i s known about the mecha-nisms of these reactions. The lack of s u f f i c i e n t experimental data i s a common feature of the vu'rveti'cs of decomposition of higher hydrocarbons. Investigations of the decomposition products of n-pentane and iso-pentane were done by Prey and Hepp (61)"' i n 1933. They proved the following rate expressions f o r the two hydrocarbons. n-pentane: k = 2.5xl0 1 5e- 6 1-200/RT iso-pentane: k = 7 . 9 x l 0 1 2 e ~ 5 8 , 6 0 0 / R ! r ' These values should be i n general agreement with the present r e s u l t s i f one refers to an average energy of ac-t i v a t i o n . Whether or not one can use average values f o r E^ f o r the whole pressure range, where diffe r e n t modes of reac-t i o n might take place, i s open to question. The decomposition proceeds through a free r a d i c a l mechanism, as i t was proved by Rice et a l (64-65). The mechanisms of the uninhibited decomposition of the three isomeric pentanes w i l l be discussed. I. Isopentane The decomposition of isopentane involves a free - J i g -r a d i c a l processes has been shown by the action of i n h i b i t o r s . A s i g n i f i c a n t point, open to question i s whether homogeneous i n i t i a t i o n of the r a d i c a l reaction, or heterogeneous i n i t i a -t i o n i s predominant. At low pressures, as we have shown before, molecules can accumulate energy without c o l l i s i o n a l deactiva-tions, and hence can pass the energy b a r r i e r f o r unimolecular reaction. I f the reaction i s i n i t i a t e d on the walls, the required energy i s diminished. The surface possesses active centers, formed during the production of the vessel. The number of centers vary with the method of production of the vessel, but they may be s t i l l formed during the reaction at high temperature. According to PARAVANO (66), physical and chemical changes take place, i n a number of semi-conductors and c e r t a i n oxides at higher temperatures. An equilibrium i s attained f o r dislocations or migrations of elements from the i n t e r i o r to the surface, leading to a heterogeneous surface. On such a surface with active centers, molecules with an excess of free vafiencicjj unsaturated hydrocarbons and free r a d i c a l s , can be held, where they are subject to c o l l i s i o n s with other molecules. These molecules are not strongly attached but they a t t a i n a dynamic equilibrium f o l -lowing the laws of adsorption. On such a surface a molecule of hydrocarbon can dissociate, thus causing an i n i t i a t i o n reaction. At low pressures, where homogeneous i n i t i a t i o n through binary c o l l i s i o n s i s n e g l i g i b l e , heterogeneous i n i t i a -t i o n must be considered as the predominant process, except f o r higher temperatures, where the available energy i s high enough to enable unimoleculomr decomposition to occur. Por ethane the difference between B(C-C) and D(C-H) i s nearly -120-20 KCal/Mole ( 2°). I t i s obvious that f o r higher hydrocarbons some v a r i a t i o n i n t h i s difference occurs, but the greater part of difference remains. Prom Semenov's data (20) we get the following values: D(*t-05Hy-CH5) = 79 KCal/Mole DCi-C^-GH^) - 74 KCal/Mole (?) The production of a r a d i c a l , from a CH3 group joined to a secondary carbon atom, w i l l thus be favored as an i n i t i a t i o n reaction: CH3-CH-CH2-CH3 k°m°£-> CH,+ CH*CHCH?CH,(Ia) CH3 or heterog. •> J •> A second decomposition mode involves the loss of the methyl group with the formation of the isobutyl r a d i c a l . CH3-CH-CH2-CH3 > CH3-CH-CH2°+;CH3 (lb) CH3 CH3 The difference i n f o r these two processes i s about 4 KCal/Mole and thus (la) w i l l be approximately 12 times f a s t e r than reaction ( l b ) . A t h i r d p o s s i b i l i t y would involve decomposition into ethyl and isopropyl r a d i c a l s j t h e endothermicity of t h i s reaction can be calculated as follows: &Hj-(0H 3) = 31.5 KCal/Mole while &Hf(C'2H"5) = 26 KCal/Mole. So examirringthe following process iso-CgB^"' V C2H'5+ CH3CHCH3 (Ic)' and comparing i t with (la) we get the following. D.(CH3-9H" CH3) = AHf (CH3)S+^HiCH3CHCH2CH3)'-CH2-CH3 -AHf(i-C 5Hi 2). and D((CR3)2CPr— C2H^) = &Hf(O2H-5-)' +6Hf(CH3CHCH3) -- bWf ( i - C 5 H i 2 ) : -121-So D(i-C4Hc^- CH5): - D((OH'3)2C3H-- C^Y. = : A%(QH 5) - AHj(C 2H 5) .+:6H1F (OH3OHOH2OH3)' -AHfcCCHjCHCHsy . 31.5 - 26 +.8 - 12.2 Of If. KCal/Mole So D(i-C 4H 9-CH3) - D((CH3) 2CH-C 2H 5) ;= I KCal/Mole Average of AH^for t-butyl and n-butyl r a d i c a l s 4.5 and 12 KCal/Mole respectively were taken, f®r ftHf (CH^CHCH^CH^)^ following the empirical formula of Vbevodskii, f o r the c a l -c u l ation of bond di s s o c i a t i o n energies. t>(C.-H) (prim.) = D(C-H')(second. )•+, B'- D(C-H) (tertiary)-nB So n(0-H)-Beoond.» D ( c - H ' ) primary E ( C - H ) t e r t ./2 The above ca l c u l a t i o n indicates that reaction (Ic) can make a s i g n i f i c a n t contribution to the i n i t i a t i o n process. Error i n c a l c u l a t i o n i s due to approximation ej &Ef ((isobutyl). I n i t i a t i o n can be attributed to these three decompositions. Of course the contribution of these reactions to the o v e r a l l process would be ne g l i g i b l e i f the chains are long-, but w i l l be very important i f the chains are short. The d i s t r i b u t i o n of the products, i n the present study shows that reactions (la) and (lb) make a greater contribution than reaction ( i c y . I f these three reaction mechanisms have nearly the same energy of a c t i v a t i o n , then reaction (IC)' must have a lower s t e r i c f a c t o r . The r a t i o of amounts of OH" and CLHV 4 do which must be at least proportional to the concentrations of CE3, and C/jHtj respectively, varies markedly with pressure. The r a t i o CH^/ C2Hg i s nearly 9 f o r 50(mm)initial pressure, but f a l l s to 7 f o r 70(mm)and to 4 f o r 100(mm)initial pressure. -122-These differences do not require that the s t e r i c factors vary with pressures but merely that C2H5 i s converted more to C2H6;. at higher pressures than at lower pressures. The difference of 9-10 KCal/Mole f o r D( C - & ) p r i n u and DCC-H1), j w i H favor the abstraction of hydrogen by CH^and C^H^ r a d i c a l s from the t e r t i a r y carbon. The fact that the increase i n CgHgis accompanied by decrease of CH^indicates the competition between CH3 and C2H5 r a d i c a l s f o r the propo-gation of the reaction. The propagation w i l l take place as follows: CH5-CH-CH2CH3+CH5 » CH4+ CH^-CHgCH^ (S)) C H3' CH,-CH^CHi-CH 2 CH 3 + C2H5- > C2Hg •< CH^C-CH^CH^. (3)) CH 5 Cff 5 5 The a c t i v a t i o n energies f o r the reaction of OKy r a d i c a l with paraffins i s approximately (£7) 6 - 8 KCal/Mole. Trotman-Dickenson and Steacie calculated that 1 the abstraction of a secondary hydrogen by CH3 requires 2 KCal/Mole. les s than a primary hydrogen. Extending t h i s c a l c u l a t i o n to t e r t i a r y hydrogen, and provided th&t f o r the reaction CH5, + <f£-C5I%2 > 0H 4 +tl-C 5H i : L (68) E A = 8.1 KCal/Mole. and the s t e r i c factor Ps 3 x lO""4" we obtain f o r reaction (2) that E A w i l l be d i f f e r e n t , approximately b> 5 KCal/Mole. with the same s t e r i c f actor P = 3'x 10~4 . For reaction (3) i f D(CH5-H) = 101 KCal/Mole. and D(C2H5-H)= 98 KCal/Mole., t h i s difference of 3 KCal/Mole. would be encountered i n reactions l i k e -123-RH 4- C2H5 * C2H6+ R. BET + CH3 » CH 4 * R. So f o r reaction (3) i s expected to be around 8-9 KCal/Mole., but the s t e r i c factor i s unknown. So from the point of view of a c t i v a t i o n energies the following r e l a t i v e reaction con-stants f o r reaction (2) and (3) should be found k 2 / k r e 3.000/RT'D e 3.000/2 x 800 ^  e2 Thus k 2e7.4 k^; f o r T = 529° C unless the s t e r i c f a c t o r * are greatly di f f e r e n t and counterbalance the difference i n ac t i v a t i o n energy. Reaction (2)' i s expected to proceed at least 7 times f a s t e r than reaction (3). The free r a d i c a l produced CH^S-CH^CH^ i s very unstable and CH3 can decompose by two dif f e r e n t mechanisms CH~—C—0HoCHv 3 , 2 3 f 3 . yrz » CH5+CH2=C-CH5 (4a) CH3 " ^ * C 2 H 5 + C3 H6 The r e l a t i v e p r o b a b i l i t i e s of (4a) and (4b) were found to be 1/4 (69)'but the present r e s u l t s indicate that the value i s closer to 1/2. The r a t i o i-C^R^^Hg seems to be l a r g e l y unaffected by pressure. Higher temperature appears to favor reaction (4b). The r a t i o OH4/C2H5 on the other hand indicates the competitive reaction with isopentane of CH3 and C2Hi>-r a d i c a l s . -184-The r a d i c a l s which, were formed i n the i n i t i a t i o n process, also take part i n the propagation. CH3-CH2-CH-CH3 $rCH5CH .. 0H"2+ CH3 ( 5 ) CH3-CH-CH2 » CH5CH = CH2+CH3 (6T CH3 CH3-CHCH3 » O2H4.4.CH3 (7) Isomerization can change i s o b u t y l into n-butyl. This r a d i c a l decomposes as follows < * - C 4 H 9 — ^ "* CjR6+ OH3 E A = 23 KCal/Mole. (5a) -» OgH^+CgHj E A = 23 KCal/Mole. (5b) For the reaction ( 7 ) 'the a c t i v a t i o n energy i s 26^4 KCal/Mole. iiit.+Ke p r o d u c t s d s w e l l a s t * » « f v « L < l o » r > i r \ a » u : e The absence of any l-C^SQ^of" C^Hg over i-C 4Hg gives i n d i r e c t support f o r the reactions (5)'and (6). The termination process takes place i n both homogeneous and heterogeneous phases. The following might be included: CH3+CH3 > C 2H 6 (8) CHyC 2H 5 » CH 4+C 2H 4 (9) CH3+C2H5 > C 3H 8 (10) y > C:2H4+C2H6 (11a) 2C'2H5 < \ >C 4H 1 0 (Hb) C 2 H 5 — » C 2H 4 «• H* (12) 2H >H 2 (13) Since CH^ can react very e a s i l y with other molecules i t s concentration on the system w i l l always be very low and hence, i t i s rather improbable to have reaction of CH^ with CH^ i n homogeneous phase. Reaction (8) w i l l probably take place more e a s i l y on the w a l l , where adsorbed CH^ can c o l l i d e with GH^ s t r i k i n g the w a l l . The recombination has zero energy of a c t i v a t i o n . Reaction (10) seems to be much more favored than (9) since i t s energy of a c t i v a t i o n would be zero, but / Vit w i l l the take place in,|homogeneous phase, and f o r t h i s reason, i t i s rather rare due to the small number of c o l l i s i o n s . In con-t r a s t reaction (9).for which some E^ i s required, probably takes place on the w a l l . The photolysis of mercury d i e t h y l (72) gives E A ( l l a ) = E A ( l l b ) = 0.8- 0.2 KCal/mole. A d i f f e r e n t i n v e s t i g a t i o n (73) reports that E A ( l l a ) - E A ( l l b ) = 4.8 KCal/mole. and that the v a r i a t i o n of k l l a / j Q l b with temperature i s as follows T °C 75 100 150 200 k l l a / k l l b 0.36 0.38 0.40 0.46 Extrapolating these values to 500° we expect that reaction ( l l a ) would be half as fast as reaction ( l i b ) . However, re s u l t s of the present study show that C^H]_Q i s not formed / x place and thus reaction ( l i b ) does not appear to take/^ i n the temperature Auled. Propane i s formed i n small amounts and hence the p r o b a b i l i t y of c o l l i s i o n between CH3 and C2H5 i s higher than between C2Hjj and C2H5. Por reaction (12) a high a c t i v a t i o n energy i s required. -126-By using mercury-photosensitized decomposition of ethane at high temperatures (74) the energy of a c t i v a t i o n was found to he 39.5 KCal/Mole. Hence reaction (12) takes place on the w a l l . The increase of C2Hg over C2H^ f o r higher pressures a is Aconsequence of competition between reaction (11a) and (9) on the one hand, and (3).on the other. The production of C^Hg cannot be explained only by reaction (10) so that p a r t i c i p a t i o n of CH3CHCH3 i n another reaction besides (7)) i s expected at high pressures. i - C 5 H 1 2 + CH3CHCH3 » CJKQ * i - C 5 H n (7a) F i n a l l y the production of cis-butene-2 can be explained as follows at high pressures. OH'jCHCHgCHj+'M >-CH3CH-CHCH3+ H"+ M (5c) The simultaneous appearance of cis-butene-2 and H 2 at high temperature and pressure can be explained by reaction (5c) or by an alternative CH3CHCH2CH3 * H c i s _ CH3CH =CHCH34.H2 OVERALL MECHANISM The o v e r a l l mechanism can be summarized as follows: CH3CH"CH2CH5 > CH5+CH3CHCH2CH3 EA= -D(C-C) s e c > ( l a ) CH3 CH3-CH CH2CH3 » CHjCH-OH^CH^s E ( C - C ) p r i m . (lb) CH3 CH3 CR^-CH-CHg-CHj -—» CH3CHCH3+CH3 E A =D(C-C) g+ cl ( i c ) CH3 -127-CH5-CH-CH2-CH3+ OH3 * CH4*-CH5C-CH2CH3 EA= 5 KCal/Mole. (f2) CH 5 CH3 CH3CH-CH2-CH3 + C 2H 5 * CH3C-CH2CH3 * C2H6 E A = 8-9 KCal/Mole. (3) CH3 OH3 CH, / > CH,+CH2=ia.CH3 (4a) CH,_C-CH9CH, ( ^ I ^ ^ N -> C 2H 5 + C 5H 6 (4b) CH3 CH3-6: -CH2CH3 + . M. » CH3CH2CH2CH2CP^M CH3 *L- G,H, y 5> C5H6+ CH3 E A = 23 KCal/Mole. (5a) 4 9 ^ > C2H4+ C 2 % E A = 23 KCal/Mole. (5b) CH5-C-CH2 * CH3+ CH^CHiCHg (6) CH3 CH3CH-CH3 > C 2H 4 + CH3 E A = 26* 4 KCal/mole. (7a) CH3-CH-CH3 + i - C 5 H 1 2 » C 3 H 8 * i-°5 Hll ^ 7 b) CHj+CH^ H a l ^ C' 2H 6 E A = 0 (8) CH34.C2H5 i m l i s c 2H 4*C 2H6 E A ^ O (9) CH 5*C 2H 5 Q^THQg> C3H8 E A = 0 (10) jPaUs> C2H54.C2H/ E A = 4.8 KCal/Mole. ( l l a ) aC 2H 5, < W EA= 0 ( l i b ) walls C2H5 -> G2H4 + H E A = 39.5 KCal/Mole. (12) -128-CHjCHCHgCHj* H Wftllffc CIS- CH3CH=CIiCH3+H2 E A P o s i t i v e (13a) H~* H T w a l l H 2 (13b) I I W - Pentane N - Pentane resembles isopentane i n many respects as f a r as the mechanism i s concerned. The differences i n products are small. As has been shown previously the decomposition proceeds through a free r a d i c a l mechanism. The molecules of <ft - pentane show: preference f o r breaking at the 1 - 2 p o s i t i o n than at the 2 - 3 p o s i t i o n . Hinshelwood and Stubbs (36) had c a l -culated the r e l a t i v e p r o b a b i l i t i e s to be I/Q.78* They suggest that - pentane i s les s stable than the more branched isomers, because of an inductive e f f e c t . They claim that the C-C bond i n a branched isomer i s strengthened because of the weakness of the adjacent C-H bond. Thus normal pentane, with four secondary hydrogens adjacent to the 2 - 3 p o s i t i o n w i l l than ^ . . ^ tend to decompose r e l a t i v e l y less Athe 1 - 2 pos i t i o n , with two secondary hydrogens adjacent. Hence, decomposition w i l l s tart as follows. CH3CH2CH2CH2CR * CH3+CH2CH2CH2CR (1) (2a) CH3CH2CH2CH2 \ •> C2H'3+C2H4 (2b) Reactions (2a) and (2b) have the same a c t i v a t i o n energy E A = 23 KCal/Mole. (75)' while the reaction —429-CH5CH2CH2CH2 > CH^CH^CHsCH^ + H" (3) with E A« 40 KCal/Mole. (75) i s of minor importance. At 500°C. the medium i s highly energetic so that reaction ( 3 ) c:an p a r t i c i p a t e . Actually Prey and Hepp (76)' found that the H - C^H^ r a d i c a l decomposes three quarters "by mode ( '2b). The r e l a t i v e l y small amounts of $ - C^Hg i n the present study, indicate that equation ( 3 ) i s s t i l l important. The decomposi-t i o n i s propagated as follows: CH3. +« CH5CH2CH2CH2CH3: »• C H 4 * CHJCHCH2CH2OH3 (4a) CH 3. +• CH3CH2CH2CH2CH3 > OH^*. CH3CH2OH2C3H2CH2; (4b) Equation (4a) participates more than (4b) because of the s t a b i l i t y of the primary hydrogens, pressure At higher C2H5; would compete with CHj; i n the propagation. C H 3 C H 2 C H 2 C H 2 C H 3 +, C2Kj -> C2Hg- *, CH3CHCH2CH2CH3 (5a) CH3CH2CH2CH2CH3 +1C2H5 » C2He + , CH3CH2CH2CH2CH2' (5b) The v a r i a t i o n of CH^/C^&g with pressure i s a j u s t i f i c a t i o n of the above equations. The f a c t , no 4SL-C'CJHJQ was detected excludes the reaction 2CH3CH2CH2CH2CH2 * 4 i - C5H12 * 4 l - C 5 H i 0 f o r the experimental conditions of the present study, perhaps because n - pentyl r a d i c a l s are very unstable and are u n l i k e l y to react with themselves v/ithout decomposition to smaller fragments. The r e l a t i v e l y high amounts of C 2 H4 +>• C^ H'g-t- C3H6 that were found i n comparison with -l-C^Rg+i CH4- shows a pre-ference f o r the decomposition of *q, - C^H^ into C2H5. + C3H6 - 1 3 0 -This point i s i n agreement with a n a l y t i c a l r e s u l t s of other previous workers (77). The reaction thus could be propagated as follows: — » C 2H 5. t- C5H6 (6a) CH3CH2CH2CH2CH2. x » CH 3 * S-C4H8 (6b) For the isopentyl r a d i c a l three p o s s i b i l i t i e s e x i s t : 2CH5CHCH2CH2CH3 » CR^CHsCHCH^ CH3+. -ft-C 5 H 1 2 (7a) TL-Q'5K±2 * CHjCBCHgGHgCHj »-ft-C^H^ + CE^CH^GR^CH^CH^ (7b) 2CH3CHCH2CH3 • C'10H22 (7c). No i n d i c a t i o n of reaction (7a) was found. Reaction (7c) was not checked, since i f some Cio^-22 w a s f o r m e < i would probably remain i n the chromatography column. Reaction (7b) seems to be the more reasonable. In addition i t i s i n agree-ment with isomerization theory (78)' which suggests that i s o - r a d i c a l s can isomerize to a normal r a d i c a l by c o l l i s i o n s with other molecules. N-pentyl radicals can also react with methyl radicals leading to the termination of chains. CH 5 4.«*K-C5Hii > CH4 +<.I-0,5H]Lo This reaction can be detected at temperatures below 110° C. (79). At temperatures above 500° C. i t must be highly im-probable due to the high i n s t a b i l i t y of Tt-CcjHn r a d i c a l . Methyl r a d i c a l s of course react with the products of primary decomposition such as I-butene (80) but t h i s i n v e s t i g a t i o n has been done larg e l y at temperatures below 220° C. -131-Termination can take place as follows: wall, C 2H 5. +« CH 5 > CH 4 + 0 2H 6 (8a) ho^ot.°3H8 (8bX Reaction (8a) seems to predominate since i n s i g n i f i c a n t amounts of O3H3 were detected. Another termination step would he toaogs, a H . ( 9 a) 2 C'pHc- ' 4 ^ 0 X * C ?H^. G'oHg (9b) w a l l 2 4 ^ Reaction (9a) has a zero a c t i v a t i o n energy, but i s ruled out by the fact that no C^H^Q was found. Reaction (9b) has an ac t i v a t i o n energy approximately 5 KCal/Mole. The reaction C 2H 5 > C 2H 4 + Hr (10) although i t has high a c t i v a t i o n energy, approximately 40 KCal/Mole., cannot be excluded, since i t provides a source of H atoms f o r the production of Eg. The reaction should take place more read i l y under heterogeneous conditions where lower energy i s required. The low concentration of C2H5. makes reaction (9b). l e s s probable. Hence decomposition of C2H|by reaction (10) w i l l be favored. I f t h i s reaction takes place on the walls i t needs less energy. Methyl can disappear by recombination on the w a l l CH3+, CH3 ^ i i > C 2F 6 (11) In spite of the fa c t that (11) has zero a c t i v a t i o n energy, i t probably takes place on the wall due to the low concen-t r a t i o n of CH^ and the low frequency of c o l l i s i o n s . Reactions (9b), (10) and (11) w i l l probably take place almost exclusively on the walls, because i n the gas phase, they w i l l be displaced -132-by (4a), (4b), (5a) and (5b). F i n a l l y another termination takes place i n homogeneous phase. CH 4 * . l - C 4H 8 (12) C5H12 (13) The r e l a t i v e rate constants f o r these reactions are given as CH3* <tt-C4H6 ( * \ > k 1 2 / k 1 5 = 0.15 (81) These reactions would be of minor importance at higher temperatures but they cannot be excluded. The products do not follow f i r s t order k i n e t i c s with time, which suggests some complexity i n t h e i r mechanism of formation. The main mode i s rather (6a) and (6b), especi-a l l y i f the chains are long. The competition i n the produc-t i o n of CH 4 and C^Hg and the v a r i a t i o n of CH^/C^Hg i s explained by reactions (4a) and (4b) on the one hand and (5a*) and (5b) on the other. The v a r i a t i o n of C^Hg with pressure can be explained by the reaction M4.CH3CH2CH2CH2. ^ M 4. CH34.CH2rsCH.CH3 (3a) 1 which w i l l take place i n homogeneous phase, and of course i s favored by increased pressure. OVERALL MECHANISM The t o t a l mechanism can be summarized as follows: CH3CH2CH2CH2CH3 > CH34-CH3CH2CH2CH2. E A = D(C-C) ( l ) . »CH*+ CH3CH =CH 2 EA-s23 KCal/Mole. (2a) CH3CH2CH2CH2. ( \ ^ C 2H 5 4. C 2H 4 E A^23 KCal/Mole. (2b) CH3CH2CH2CH24 HS3t^ CH3CH2CHSCH24. H E A = 40 KCal/Mole. (3) M 4. CH3CH2CH2CH2. ^°m°#- M4.CH3. 4. CH3CH=CH2 (3a) -133-A > CHA4.CH3CH2CH2CH2CH2. (4a) CH^CHaCHoCHoCHa+CHi. < 4 > ^ CH4+CH5GHCH2CH2CH5- (4b) , > 0pH64.CH3CH2CH2CH2CH2' (5a) CH,CHpCH2CHpCH-z4.C2H5 ( i > C2H64.CH3CH-CH-2CH2GH5 (5b) 2 CH3CH2CH2CH2-—» *i-C 5Hi2+%-C5Hi0 (?) . *C 2H 5. + C 3H 6 (6a) CH-zGHpCHpCHpCHp — — v ^ ^ *,CH3+ S-C 4H 8 (6b) 2 CH3CH CH2CH 2CH3 CH30H = CHCH3+-«^C 5Hi 2 (7a) M4.CH3CHCH 2CH 2CH 3 * M+CH3CH2CH2CH2CH2• (7b) 2CH^CH CH^OE^CH^ ~> ^10^22 ) ™2i4 s CH44.C2H4 EA>0 KCal/Mole. (8a) C2H54.CH3-2 C2H'5-hS=5|. °3H8 EA= 0 ^ C 4H 1 0 E A = 0 ( 9 a ) — C 9 H A J + C 2H6EA= 4£ KCal/Mole. (9b) walls ^ 4 C 2H^ w a l^ s> C2H4,4.H- E A = 40 KCal/Mole. (10) CH3 4. CH3 C 2H 6 E A = 0 (11) j * CH 4 + 4-C 4H 8 E> 0 (12) OH, + ta-C4Hb ( * N * <ft-C5Hi2 E=0 (13) I I I NEO-PENTANE The decomposition of neo-pentane was investigated i n s u f f i c i e n t l y . I t s energy of a c t i v a t i o n 53.5 KCal/Mole. -134-whieh was found only from two temperatures, appears to be rather low. I f the more branched isomer has higher thermal s t a b i l i t y and , the order of a c t i v a t i o n energies i s reversed, that i s E A(n-pentane)y E A ( i s o p . ) ^ E^(neo-pentane) the s t a b i l i t y order must be governed by the s t e r i c f a c t o r . The following order must be v a l i d to explain the r e l a t i v e s t a b i l i t y i n spite of the reversal of a c t i v a t i o n energies. ^ neo-pent. ^  ^ isopent .^^n-pent. I t i s well known that frequency factors are d i r e c t l y related to the entropy of a c t i v a t i o n P = e*S~* /R while the c o l l i s i o n number, ^ , i s a function of pressure and temperature. The entropy f o r the l o c a l i z e d (constant volume); gas, can be given by S = R $ lnQ+' T dlnQ \ dT J I f we consider that, molecules of saturated hydrocarbons behave l i k e i d e a l gases with n e g l i g i b l e i n t e r a c t i o n energy, the t o t a l p a r t i t i o n function should be the product of the p a r t i t i o n functions f o r t r a n s l a t i o n , r o t a t i o n etc. so Q tot.= <1 trans.TT * v i b r . T l * r o t . q e l . * mucl. Experimental data show small difference i n r e a c t i v i t y between isopentane and n-pentane, but much greater difference between -135-them and neo-pentane. This point i s i n agreement with experi-mental energies of ac t i v a t i o n . One important point i n the treatment of p a r a f f i n molecules i s that they appear more c o i l e d than straight. Raman spectret of l i q u i d C4-C7 hydro-carbons show more l i n e s than these corresponding to straight configuration (82,83,84). In the gas phase l e s s c o i l i n g i s expected than i n l i q u i d , but c o i l i n g s t i l l e x i s t s . C o i l i n g has an important influence on the c o l l i s i o n cross section, and thus on the r e a c t i v i t y of the molecule. Coiled molecules interact l e s s than l i n e a r molecules to form free r a d i c a l s , and these s t e r i c e f f e cts favor attack at the end of the molecule rather than i n other parts, i n spite of the fact that t e r t i a r y or secon-dary C-PT are more vulnerable (85). McCOUEREY et a l (86) gave values f o r the c o l l i s i o n a l area according to Sutherland's The equation, small difference : . • between n - and isopentane °o 19.6- and 18.9 respectively, • 1 gives an account f o r the s i m i l a r r e a c t i v i t y . E f f e c t i v e c o l l i s i o n diameters given by DARWE1T (87) show r e l a t i v e r e s u l t f o r n - and isopentane but lower f o r neo-pentane. These values are 2.94, 3.52 and 1 1.22 x 10 cm r e s p e c t i v e l y ^ c o l l i s i o n s are much more favored i n the case of n- and iso-pentane than i n neo-pentane. A t h i r d point which relates to the c o i l i n g of molecules i s the r e s t r i c t i o n i n energy transfer i n r o t a t i o n a l and v i b r a t i o n a l modes. These r e s t r i c t i o n s can be examined from thermal con-d u c t i v i t y data (88). Defining the gas> p a r t i t i o n function, f o r no transfer r e s t r i c t i o n s we have the following expressions -136-f o r p a r t i t i o n function: q = ^ t r a n s l . * ' * r o t . + ' <I V i b r . According to the Chapman-Enskog 4 transl.= 2»5^ptransl. and according to the Eucuen's hypothesis * r o t . * ^ 0 r o t . ( 8 9 ) * v i b . " ^ 0 v i l a. w h e r e ^ i s the v i s c o s i t y and C i s the heat capacity. The observed data l i e between two extremes q (no - r e s t r i c t i o n ) ' = /0 (2.5V>C * C . *'0 \ * I „,tran. r o t . ^ vib.) q (vibration r e s t r i c t i o n ) (2.5C t r a n s l #-«- C rot.) For the lower p a r a f f i n s , there i s no evidence f o r r e s t r i c t i o n s but f o r higher, there appear to be a r e s t r i c t i o n , connected with the c o i l i n g of the molecules. So the p a r t i t i o n function due to v i b r a t i o n a l transference appears to be reduced f o r more branched hydrocarbons and s p e c i f i c a l l y f o r neo-pentane. Consequently the t o t a l p a r t i t i o n function, and entropy i n r e l a t i o n with i t , w i l l follow the same rel a t i o n s h i p . Since the P factor i s an exponential function of S, a small change i n S, r e s u l t s i n a large change i n P. Thus P (neo-pentane) « P n _ c 5H 1 2i-C 5R- 1 2 Thus the reduced r e a c t i v i t y can be accounted f o r . The main products of neo-pentane decomposition are i-C 4Rg and CH^. Por t h i s reason i t s decomposition seems to be rather d i f f e r e n t from that of the other isomers. The lack of i n h i b i t i o n at least f o r the f i r s t stages -137-indicates the possible occurrence of two mechanisms, process, a molecular rearrangement The main CH* i 5 c a 3 , CH,-C-CH, ? CHA + CHr>» C' which dominates I * ^CH3-CH3 5 at le a s t i n the i n i t i a l reaction, and a secondary reaction, a free r a d i c a l process, indicated by the s l i g h t i n h i b i t i o n i n o v e r a l l rates. Lack of data do not allow confirmation of these mechanisms. A free r a d i c a l mechanism such as CH3 CH3 OH3-C-CH3 — ^ CHjJr* + CH3- (1)' CH3 CH3 c m CHI, 1 t D CH3-C-CH3 +. C H 3 —^ CH4+ CH2-C-CH3 (2) CH3 CH3 CH3 I CH3 •CHp-C-CH, ^CH3.+ CH2= C / (3) I ^ X C i r , CH3 0 does not explain the traces of C3H5 detected. N i t r i c oxide could react very e a s i l y with the stable r a d i c a l CH3—C—CH3 leading to the formation of (CH^^C'-NO, but t h i s compound was not detected, although i t should be reasonably stable. C3H6 may come from the pyro l y s i s of CH2=C^CH'3 -138-B: INHIBITED DECOMPOSITION N i t r i c oxide was used as the i n h i b i t o r i n a l l cases. I t was preferred to C^Hg f o r two reasons. F i r s t , because i t i s a much more potent i n h i b i t o r and secondly because, C^Hg i s one of the decomposition products, so i t s consumption during reaction could not be checked. Previous workers found i d e n t i c a l products f o r the normal and f u l l y i n h i b i t e d decom-po s i t i o n . In the present study, thajr has been shown to be not the case. The l i m i t i n g pressure of NO.appears to be a function of the reaction vessel rather than of the hydrocarbon concen-t r a t i o n . This fact can be related to the heterogeneous pro-cesses rather than to homogeneous processes. The l i m i t i n g rate appears to be a function of hydrocarbon pressure f o r isopentane ( F i g . 32.) The following v a r i a t i o n of the r a t i o : rate of u n i n h i b i t e d ^ / , u / r a t e of f u l l y i n h i b i t . with pressure was found p isop. 50 mm 80 mm 100 mm ^ 0 / ^ 5.3 4.8 3 This v a r i a t i o n i s contrary to Hinshelwood's ideas, that a l l chains are suppressed by NO. The suppression i s higher f o r lower pressures. I f NO i s related to the surface of the reaction vessel and consequently to heterogeneous processes, then only the heterogeneous ones are to be stopped by NO. For lower pressures therefore, heterogeneous i n i t i a t i o n i s the predominant i n i t i a t i o n reaction, so the rate f a l l s very -139-sharply. For higher pressures, homogeneous i n i t i a t i o n com-petes successfully with heterogeneous i n i t i a t i o n and the reduction of r a t i o i s not so sharp. I Isopentane The required amount of NO, to a t t a i n f u l l i n h i b i -t i o n i s 10 mm f o r the entire pressure range of . 50 - 100 mm. Analysis of the products shows the difference between the normal and f u l l y i n h i b i t e d decompositions. This i s i l l u s t r a t e d i n results i n Table XII. TABLE XII ANALYTICAL RESULTS FOR ISOPENTANE p i s o p = 100(mm) Pft 0 = 13 (mm) T = 569^ AP/ p = 0.2 Products Uninhibited Mble.#'' f u l l y i n h i b i t e d : mole. $ NO estimated NO excluded i - C 5 H 1 2 72 60 64.2 C,3H6 10 10.65 11.25 i-C 4 H 8 6.85 6.1 6.53 CH4 6.95 12.15 13.0 C2H4- 2.16 2.50 2.68 C 2H 6 1.74 0.95 1.02 NO 0 6.50 TOTAL 99.70 98.70 98.68 From these results i t i s obvious that products such as CH4, -140-C2H4, C^Hg are sharply influenced by the NO concentration. The t o t a l i n the case of f u l l y i n h i b i t e d decomposition i s 98.70 which means that either some NO i s l o s t i n the form of N0 2 which i s kept i n the column or i n the form of n i t r o s o - compounds, other than C2H5NO, which were not detected. The energy of a c t i v a t i o n was calculated f o r the f u l l y i n h i b i t e d reaction and was found to be approximately 5 ^KCal/Mole. with an A factor of 1.4xl5" 1 2 Seel" The estimated values for the rate constants were: k519* = T.65xl02mm min 1  k502° = 7.4xl03mm mlrfi Those values are more or l e s s i n agreement with the corres-ponding ones found f o r higher pressures i n uninhibited reac-t i o n ; This seems to be reasonable from the point that i n the f u l l y i n h i b i t e d decomposition a l l heterogeneous processes are stopped and only homogeneous decomposition takes place. This idea i s supported by the v a r i a t i o n of products with time. Prom those curves, presented i n Chapter I I I i t i s concluded that for the l a t e r stages of the reaction d( C2 H4]/dt = k l a ( C H 4 ) / d t = *2. d(C 5H 6)/ d t B ! k 3 and d(i-C4H8)/at = k^ where K]_, k2> k 3, and k 4 are c * t J - r ^constant i n i t i a l pressure of hydrocarbon. Ethane shows a higher order dependence f o r reasons discussed i n Chapter I I I . When NO pressures exceeds the l i m i t i n g point, small amounts of a nitroso-compound were detected, which was believed to be C2H5NO and whose concentration increased with increasing NO. -141-The reaction of NO with a l k y l r a d i c a l s has been discussed i n the l i t e r a t u r e f o r lower temperatures (90,91»92,121). Nitroso-compounds, RNO, were eliminated from the pyrolyses of Hg(R) 2 with NO. Their s t a b i l i t y around 350°C. varies as follows: CH5N0 < C 2H 5N0 <_ isopropyl-NO < t e r t . butyl-NO The temperature used i n the present work i s very unfavorable f o r CH^ NO which does not appear at a l l }while C2H5NO i s not completely decomposed. The fact that f o r ^ NOa 15 mm, there was no detectable C2H5NO although t h i s amount i s much more than the l i m i t i n g value i n PN0, shows that i t s decomposition i s counterbalanced by i t s formation,which i f favored by an increase i n ^ NO up to 75 mm. The competition between NO and other molecules f o r C2H5 ra d i c a l s i s indicated by the v a r i a t i o n of the r a t i o C2H5/C2H5NO with ^NO^as we l l as., by the v a r i a t i o n of (^2^/02'^^' T n e PN0 used up,was higher than the l i m i t i n g value. The v a r i a t i o n of the l a t e r r a t i o i s not regular. Some values with '^ITO are given i n Table X I I I . -142-TABLE X I I I VARIATION OP O2H4/C2H5-NO WITH P^O) C!2H4/C2H5N0 *P/p = 0.15 * % - 0.2. * P / P = 0.3 30 1.37 3.38 40 1.96 1.67 2.9 50 3.9 1.95 2.14 75 1.2 ? 4.1 2.4 Ih chapter I the more important ideas about i n h i b i -t i o n were introduced. These involve mainly i n h i b i t i o n as a homogeneous phenomenon. The re s u l t s of the present study are by accounted f o r v d i f f e r e n t points of views of both heterogeneous and homogeneous reaction. Thus,we are led to a twofold effect of NO. A heterogeneous effect at low concentrations of i n -h i b i t o r with complete coverage of the active centers, i n agreement with Voevodskii's hypothesis. Secondly, at higher concentrations of the i n h i b i t o r both heterogeneous and homo-geneous effects operate. This l a t e r hypothesis i s supported by the detection of a nitroso-compound and by the consumption of NO. The adsorption of NO by the walls suppresses not only heterogeneous i n i t i a t i o n but also heterogeneous terminations. This l a t e r postulate i s supported by the observation that surface effects occur even at high pressures of isopentane. This i s a contention i n disagreement with Voevodskii's ideas. He predicts an influence of surface only f o r low pressures. -143-An effect of different S/V r a t i o i n both low and high pressures was found. That NO i s not f i x e d on the w a l l butjfl? attains a dynamic equilibrium i s shown from the fact that NO i s con-sumed s l i g h t l y even f o r p a r t i a l l y i n h i b i t e d decompositions. Thus,the d i s t i n c t i o n between the two roles of NO i s not sharp. At low pressure heterogeneous effects predominate and provided that heterogeneous i n i t i a t i o n i s the main reaction f o r low pressures,the effect i s profound. This has been shown i n experiments with f u l l i n h i b i t i o n f o r the whole pressure range. For pressures lower than 20(mmjdecomposition was hard to detect. At those pressures only heterogeneous processes appear to par t i c i p a t e and since they are suppressed decompo-s i t i o n does not occur. N i t r i c oxide, i t s e l f a r a d i c a l , i s chemisorbed on quartz. C o l l i s i o n between NO and a l k y l r a d i c a l s especially CH3 and C2H5. are more probable on the w a l l than i n the gas phase. A r e l a t i v e l y stronger N-C bond^causes the rupture of the NO-quartz pseudo-bond and leads to the production of C2H5NO and CH3NO. Nitroso-methane i s very unstable and i t s decomposition i s accelerated by NO ( 9 3 ) . C o l l i s i o n between C2H5, CH3, C2H5NO and CH3NO on the w a l l and i n the gas phase can take place leading to the production of CH4, C2H4 and C2 Hg CH3NO * CH3 > C2Hg4.N0 (a) CH3NO + C 2H 5 9* C2H44.CH4+NO (b) C2H5NO + CH3 ^ C2H4-CH4-NO (c) C2H5N0 + C 2H 5 ^ C2H4+CH4+NO (d) -144-Ih the above set of reactions (a) and (b) are predominant because of the d i f f e r e n t s t a b i l i t y of CH3NO and C2H5NO. The effect of NO remains l a r g e l y as a function of the surface as long as we are below the l i m i t i n g pressure of NO, but f o r higher amounts>another reaction takes place, the homogeneous reaction between the hydrocarbon and NO. In order to explain the fact that the rate i s no longer suppressed, but a c t u a l l y there i s some s l i g h t acceleration, NO i s con-sidered to p a r t i c i p a t e i n both homogeneous i n i t i a t i o n and termination of chains as has been proposed by a number of previous workers. As f a r as,the f i r s t point i s concerned, the fact that the l i m i t i n g amount of NO, was independent of the p a r t i a l pressures of hydrocarbon but i s a function of the surface area provides enough support. By increasing the surface the l i m i t i n g point was s h i f t e d to higher NO values. A l i n e a r relationship between t h i s s h i f t and the S/V r a t i o was not expected, since the important factor i s not the whole surface but the number of active centers. The effect of surface has been examined by other workers also f o r various hydrocarbons with both negative and p o s i t i v e r e s u l t s . PEAP.DS et a l (94) reported some effect of surface which stopped at the early stages. However, they used pressures i n the range 50 mm and up i n which homogeneous reactions predominates. POLTORAK and VOEVODSKII (95) exa-mining the decomposition of propane found the i n h i b i t e d rate approximately of second order f o r an untreated quartz vessel. Treatment of the vessel with HP increased the rates P 1 4 5 -i n the pressure range 30 - 60(mmJ. PURNELL and QUINN (32) using diff e r e n t surfaces i n the pyro l y s i s of n-butane re-port no change i n product with the surface. In the present study v a r i a t i o n i n the product d i s t r i b u t i o n with the surface area was found, p a r t i c u l a r l y i n the i n h i b i t e d reaction. The fact that there i s no r e a l effect of surface on C3H5 or i-C'^Hg; suggests that the main mode of reaction i s not affected by heterogeneous processes. Since heterogeneous i n i t i a t i o n i s completely suppressed, homogeneous i n i t i a t i o n appears to be independent of surface processes. The decom-posi t i o n can thus start as follows CH3 CH^CHoCH< — * CH,- + CH*CHoCHCR* ( l ) ° CH3 0 0 ° CH3- + CH3CH2CHCH3 » CH4+ CH3CH2C-CH3 (2) CH3 CH3 OH3 y > CH34.CH3C'=CH2 (3) CH. OH: 6 (. "n3 - » C 2H 5 4.: CH3CH=CH2 (4) \ ^ 3 CH* V CH3 CHT5 CH3CH2-CH< + C 2H 5—* C2H6+ CH 5CH 2C< (5) CH3 CH3 Extending the isomerization hypothesis proposed for isopropyl r a d i c a l s to isobutyl r a d i c a l s we get: CH3CH2CH-CH3 4- i - C 5 H 1 2 — i - C 5 H 1 2 + CH3CH2CH2CH20 (6) CH3CH2CH2CH£ > CH 5+, C 3H 6 (7) The mechanism i s now i d e n t i c a l with that f o r the uninhibited reaction and no effect of NO i s involved]for higher pressures -146-of NO, jjJome homogeneous i n i t i a t i o n takes place through c o l l i s i o n s as: i-C 5H 1 2+ NO 1—» i - C j H i ! + HNO (la)' CH3+ i - C 4 H 8 (3) i - C 5 H n r \_> C2H54. C 3H 6 (4) An increase i n rate of decomposition i s thus expected f o r higher P(N0). This increase should p a r t i a l l y be counterbalanced by homogeneous termination. The formation of C 3H5 i s unaffected by both NO and S/V r a t i o , which means that reaction (4) i s unaffected by both PN0 and S/V r a t i o . Accordingly reaction (3) must be s i m i l a r l y independent)because the p r o b a b i l i t y i n the decompo-s i t i o n of a r a d i c a l remains unaltered. Actually the concentra-t i o n of i-C^Hg decreases slowly with NO. Thus there must exist another source of i-C'^Hg ^ heterogeneous i n nature } and of minor importance. Ethylene and methane increase with ^NO,while CjpHg decreases. A l l of them are favored by increased surface. This fact requires that termination i s more affected by the surface than propagation. The CH3 and C2H5 r a d i c a l s can enter both homogeneous and heterogeneous processes. C2H5+GH3 —^CH 4+ C 2H 4 E A> 0 (8)' This reaction must predominate i n the homogeneous reaction because of the higher p r o b a b i l i t y of C2HCJ reacting with CH'3 than of CH3 reacting with CH 3 or of C2H5 reacting with O^S.^. In the same time C2H5 and CH3 c o l l i d e with NO on the walls -147-giving r i s e to the formation of C2H5NO and CH3NO. CH3 +N0 * CH3NO (9) C2H5+NO >C2H5NO (10) CH3NO + OH3 > C2H6*N0 (11) CH3NO + C 2H 5 CH4+ C2H4-«. NO (11a) Reaction ( l l a ) provides an explanation f o r the increasing of both CH4 and C2H^ as NO increases. At the same time C2H^ i s retained by NO, and reaction (5) i s suppressed, leading to decrease of 02Hg with increasing NO. Nitrosoethane p a r t i a l l y reacts with both CH3 and C2Hij r a d i c a l s . C2H5N0 + CH3 > C2H4 + OH4+ NO (12) C2H"5NO + C 2H 5  -> C2H4.4. C 2H 6 +N0 (13) -» C 4H 1 0 + NO (14) Thus NO competes with i - C ^ H ^ i n reacting with C2H'5 causing a reduction i n C2Hg and consequently an increase i n CH4 and C 2H 4. Increasing the surface means that more molecules of adsorbed NO are available, so higher amounts of C 2H 4 and CH4 are produced. But reaction ( l l ) i s not s u f f i c i e n t to predict the increase of C2H6 that i s observed f o r the i n -creased surface unless i t i s assumed that NO also reacts with C2HEJ homogeneously as follows: C 2H 5 * NO ^C 2H 5N0 (15) -148-Thus the increased surface reduces the amount of NO available f o r the homogeneous reaction (15). The formation of C2H5NO whose s t a b i l i t y i s enough to allow i t to remain as a f i n a l product f o r higher NO concentration^ shows a s l i g h t increase with NO. I t s formation competes with reaction (12), (13), and (14). I t i s not so stable as (CH5)2CHNO or (CH^CNO and at high temperatures i t s concentration appears constant with the time. N i t r i c oxide i s consumed i n the reaction, but the amount i s rather high i n comparison with the concentration of CgH^NO. This raises the question: Is NO consumed only i n that way? Actually no i n d i c a t i o n of other compounds was found i n the infra-red,but small amounts of nitroso compound, undetec-ted at the s e n s i t i v i t y used, could be responsible f o r the NO consumption. The relationship between C2H6 and C2H5NO gives an i n d i c a t i o n that NO i s consumed i n competition with the formation of C2Hg, which means a reaction between OH3 and C2H5. This i s not proof however that t h i s i s the only way. A l l products, but O2HJ5 increase l i n e a r l y with timej t h i s means, t h e i r formation follows a f i r s t order reaction f o r f u l l y i n h i b i t e d decompositions. Ethane competes with CH3NO and C2HcN0 formation. In the f i r s t stages i t appears rapidly while CH3NO decomposes completely and C2H5NO p a r t i a l l y . Abstraction of hydrogen by NO has been proposed i n reaction ( l a ) . This i s based on the observation that NO, cata-lyzes acetone decomposition at higher NO concentration. The -149-mechanism suggested i s (53) CH3COCH3+NO > CH2COCH3 + NOH The H - atom adjacent to the B - carbonyl are activated. HNO i s believed to enter the equilibrium HNO 1 * H" + NO or transfer H to another r a d i c a l . This decomposition reaction can be considered of minor importance, i f i t w i l l take place at a l l , i n the case of isopentane, since H 2 was detected at most i n traces. CH5CH -CH2CH3 * NO < Z ± CH3C -CH2CH3 4. HNO (la) CH3 CH3 HNO <—> NOo-H (16) CH3+NOH 9-T!H4+N0 (17) Another cause of the decrease in? CgHg comes from the reaction C2H6+NO — ^ C2H5+HNO E A = 56 KCal/Moie^ (18) The l a s t reaction can take place more e a s i l y on the w a l l , where the required energy w i l l be le s s than 56 KCal/Mole. I I n-pentane r In spite of the fact that i n h i b i t i o n i n n-pentane decomposition has been exhaustively examined, r e l a t i v e l y l i t t l e i s known so f a r , as the mechanism i s concerned. The i n h i b i t i o n i n n-pentane decomposition by NO was discovered by Hinshelwood and many investigations have been done by Hinshelwood and his colloborators (96,97,98,55,99,100). was Unfortunately most of the work/based on pressure change -150-r e s u l t s ; i . e . f o r estimation of rates etc. The k i n e t i c s of the reaction and certain a n a l y t i c a l r e s u l t s have been reported by Partington et a l but no v a r i a t i o n of the products with PJJO has been examined. The idea that the products were i d e n t i c a l f o r both i n h i b i t e d and uninhibited decomposition was accepted i n a l l these works. The apparent chain length given by the r a t i o of the uninhibited to the i n h i b i t e d reaction rate, was considered to be very important i n accordance with the idea that the i n h i b i t e d reaction was not a free r a d i c a l process but a molecular rearrangement. This contention i s now be-lie v e d to be i n v a l i d (Chapter I ) . Prom the reported values of apparent chain length (71) f o r isopentane and neo-pentane, i t i s quite obvious that i n h i b i t i o n f a l l s i n the order n»- pentane > iso-pentane > neo-pentane which conform^to the order of t h e i r r e a c t i v i t y . The apparent chain lengths decrease f o r increasing pressure f o r neo-pen-tane. This l a s t observation i s enough to cause doubt about the idea of complete suppression of chains. In the present study of n-pentane, no n i t r o s o -compound was detected, yet NO was found to be consumed. The l a s t point -raises a question about the fate of NO. The effect i n many cases i s the same as i n the case of iso-pen-tane but not so marked, n-pentane i t s e l f i s much more re-active than isopentane and the products which vary highly with NO concentration come from primary processes also. Their changes are of lesser importance. N i t r i c oxide also appears to be involved i n -151-heterogeneous processes, "but no evidence was found f o r homogeneous reaction with NO. The r e l a t i v e l y high y i e l d of nitroso-compound by reaction with these r a d i c a l s . P i n a l l y the main features of the isopentane reactions would take place f o r n-pentane too. I l l Neo-pentane cause a s l i g h t acceleration-of rate i n the f i r s t stages of the reaction, but the o v e r a l l rate i s s l i g h t l y reduced by NO. Previous workers have calculated an apparent chain length nearly unity, which means that the decomposition i s l i g h t l y suppressed by NO. This phenomenon does not conform to a free r a d i c a l process, so r a d i c a l reaction must pa r t i c i p a t e to a very s l i g h t extent. The marked influence of NO on i-C^Hs, on the other, suggests that CH^, i s the main product with i-C^Hs, of a molecular rearrangement, while i-C4H8 appears also to be main product of a free r a d i c a l mechanism. The s l i g h t effect of NO can be explained by an i n i t i a t i o n caused by i t . The a c t i v a t i o n energy for such a process appear to be around 50 KCal/Mole. This i s followed by: C 2 H 5 and CH-2; r a d i c a l s , may cause the decomposition of any The effect of NO on neo-pentane, i f any, i s to CH3 CH^ -C-CH^ -J- NO CH3 CH3-C -CH2* +HN0 CH3 -152-CH3 OH^-C-+CHP * OH, •«• CHp- C< C H 3 0 » ^ CH^ CH 5 3 9H3 <{H3 CH3+ CH3- <j - CH3 ^ CH 4 + CH3 _ G ^  CH 2 CH 3 CH"3 Ih such a mechanism an influence of W on i-C^Hg and CH 4 i s expected but i t does not explain the production of i-C 4Hio found by other workers (94). The present inve s t i g a t i o n of neo-pentane suffers from a lack of experimental r e s u l t s and no r e a l attempt to explain the mechanism can by made. -153-CHAPTER T KINETIC LAWS IN THE THERMAL DECOMPOSITION OP ISOPENTANE I' Uninhibited reaction The mechanism f o r isopentane decomposition has been proposed i n chapter IV. I t s complexity makes quanti-t a t i v e treatment d i f f i c u l t . In order to f i n d k i n e t i c laws connecting the products with concentration of reactants etc. c e r t a i n assumptions and facts are used. These are summarized as follows: Equations (6) (7) and (8) i n chapter I V are consi-dered n e g l i g i b l e f o r CH5 production and consumption i n comparison with (la) (lb) (2) (4a) (9) and (10). Provided that the chain length i s long enoughjthe main production of CH3 comes from (4a) (la) and (lb) and the main consumption from (2) (9) and (10). Bimolecular reactions l i k e (8) con-tr i b u t e s l i g h t l y . The same considerations hold f o r the C2H5radical', reactions ( l c ) , (3), (4b), (9) and (10) i n which i t i s involved, are considered as the more important. For CH3CHCH3 reaction ( l c ) and (7a) are important provided that the pressure i s high. For CH3-C-CM2-CH3, reactions (2), CH3 (3)> (4a)and (4b) are the only important reactions. Isobutene i s considered to be produced only by (4a) while C^Hg i s formed by (4b)and to a small extent by (6) f o r low pressures. The r a t i o k2/v„ i s found to be approximately 7.4 (chapter IV) and the r a t i o CH2- = 0 / C H 3 C H 3 /C 3H 6 i s shown to be 1/2. Also klla/j^Qb i s a P P r o x i m a " f c e l y 0.5. The r a t i o C2H4/C2H.6 attains a nearly constant value approxi-mately 0.85 at high pressures. The r a t i o CH4/C2H6 a t t a i n s a nearly constant value of approximately 3 . Prom these two ratios the r a t i o of CH4/C2H.4has an approximate value at high pressures of 3 . 2 . Other r e s i i l t s or assumption are employed as follows. -»C4Hilo ( 1 " ) ! 2 H4+C 2H 6 (11a) > C2 A+( f o r the reaction of C2H5. Recent work (101) from photolysis data gave k l l a = 4.2 x 10" 1 secT 1 . The value % l a / k i r b appears to take values between 0.12 to 0.40 (101,102,103). Under these conditions a value of approximately 0.20 i s assumed to be probable. The homogeneous reaction CH3+CH3 — > C 2 H 6 (8) c e r t a i n l y occurs but i t s s t e r i c factor i s low —6 P = 8.10 (104) so i t s contribution to the disap-pearance of GH3 i s neglected. The l a s t assumption i s concerned with CH3CHCH2 r a d i c a l . Both i-C^H^ and n-C3H>7 have been e a s i l y detected at room temperatures (105,106,107,108,109) but the ef f o r t to detect them at higher T, f a i l e d absolutely (107,110,111,112, 113). Thus i t i s assumed that n- and iso C3H7 decompose very e a s i l y at high temperatures so the' contribution of reaction (7b) i s considered to be unimportant. For the decomposition 1 -155-n- and i - C ^ > C2H4+CH3 • (7a) -9 the A factor has been given as 3 x 10 but t h i s seems rather low. Under these assumptions a stationary state c a l -c ulation f o r CH3, C3H7 and C^H^Q can be made. ^ T ^ = k l a [ l - C j j H ^ * , k l b [ i - C 5 H l 2 ] - K 2 [ i - C 5 H i j H g . + k 4 a [ i C 5 H ] | - kgUCH^Jc^Hg- k u l H ^ H ^ = 0 ( l ) d[C 2H 5l ~ d F - = klc[i-°5H12]- ^ - C 5 H 1 2 ] [C2H5-] t, k 4 b [ i - C 5 H 1 1 ] - kg [CH3][C4H5] - k 1 0[CH 3][q iH^| = 0 (2) d[C 3H 7 3 dt - k7a[ C3 H7]=° ^ dfi-CcHi-n _ "~dT~"= k 2 [ i - C 5 H 1 2 ] p H 3 l - k 3 & - c 5 H l i L \ H i -From (3) [ C 3 H 7 ] = k l c C i - C 5 H 1 2 l _ k l c k7bC L-°5 Hll f K7a" k 7 b + k 7 A / £ - C 5 H 1 2 ] For high-pressure L C3H7"] reaches a- l i m i t ^ k l c (5) k7b By substracting ( l ) and (2) we get (7) -156-h S 1 - 0 ^ ! ^ *Lb&-C5 H12j- k 2[i-C5H 1 2][CH3] + ^ 1 - 0 ^ -- ^ c [ i - 0 5 ^ i * ^C i-°5 HlIp2H3- ^ b P . - O s H i l j - 0 (6) Prom (4) we get [ i - 0 5 H l % k^HJ) • k 3[G 2H3t[i-C5H 1 2] \ k4a"- k4b * ' By substituting (7) into (6) and we get ( k l a + k l b H3] * ( k 4 < r k 4 b ^ k 2[CH 3> k ^ f a H g l k 4 c t - k 4 b k i c * ^ [°2 I i^r D--°5H12| = 0 or k l a + k l b - k 2 [ c H 5 ] + ( ! ± * k 2[CH 3l4. ^0^5% k4b k l c " k 3 K H 5 ] = 0 ( 8) Provided that CH 2 = C". ^  * and C'3Hg are produced almost CH3 exclusively by C4a>)/4b) we have set d 1> 2 = 0<gHp — T T - 3 " " ^ ^ 1 2 7 d [C 3H 6] k,^ dt 4 b Thus: i C4H8 _ k ^ _ O3H6 ~ £ — ~ 0.5 ( i n accordance with the observed product d i s t r i b u t i o n . No effect of pressure also i s predicted, i n accordance with the a n a l y t i c a l r e s u l t s ) -157-Thus equation (8) becomes 1^+ k l b- k2|CH£J-»- (0 . 5 - l ) k 2[CHg + lqfeHJJD - k l c-k 3[C 1Hgj= 0 0.5+1 or rearranging the above formula 1 .5 k l a+ 1 .5 k l b - 1 . 5 k 2 j ^ J - O.SKgjCH^- 0 . 5 k 3 j C 2 H j 1 . 5 k l c - 1 . 5 k 3[C 2H 5"j r 0 F i n a l l y -2k 2fCH£j- 2 k 5 { c 2 H ^ - 1 . 5 k l c - l . 5 k l b + 1 . 5 k l a and k 2 JCHJ -k^2H^| = 0.75 ( k- L a+k l b-k 1^ but [GHJ|-(&2H53- 0 . 7 5 ./ k^kifr-kio \ k 3 V k 3 but ko k f * 7 ' 4 7.4pH;i-[C^Hg = 0 . 5 7 ( k l a ^ k l b - k l c N k"7 / and [C 2H^| = 7.4[CH3-|- 0 . 5 7 / k l a + k l b - k l c \ k3 / (9) The main source of production of C2R> and C 2H 4 are equations ( 2 ) and (5) especially at high pressures where the long chain length makes ( 2 ) and (3) overwhelmingly predominant over ( 8 ) and (9) and ( l b ) . Thus we can write 4 l S « f i " 1 - ^1^5*12] P ^ l ( 1 0> ^ [ l - C ^ K ) (ID and af t e r integrating and using the l i m i t i n f o r high C 2H 6 -158-pressures we get 0H 4 k2[CH3"] ^ 2.7 (12) C 2H 6" k 3 CC2H5-] " So C C H 3 l ^ r 2.7 ^ c ^ - j ( 1 3 0 k 2 and by substituting (13) into (9) we get f°2H5j= 7.4x2.7 k ^ p 2 H g - 0.57 ( k l a ^ b j ^ l c ^ or [C2H53 - 0 - 5 7 k 2 ( k l a . k l b - k l c ) 20k3-k 3 k 2 (13) accordingly [CH 3-j . 2.7x0.57 j k l a . k l b - k l c ) ( 1 4 ) 20k 3-k 2 The v a r i a t i o n of CH4 and C2H6 with the pressure. I t i s shown that they follow approximately f i r s t order k i n e t i c s with respect to the pressure. Of course i n r e a l i t y the mecha-nism i s f i r s t order only f o r high pressures v/here termination reactions are r e l a t i v e l y reduced. Ethylene comes from (5b), (7a), (9) and ( l l a ) and thus d g 2 H 4 / d t = k 5 [ n - C 4 H ^ + k 7 afOH 3CHCH| + k g f C H j g H g * klla£2H5l2 Isopropyl r a d i c a l i s constant for high pressures but s l i g h t l y dependent on pressure at low pressures, n-butyl -159-r a d i c a l depends on i-C^H^ but at high pressures the c o n t r i -bution of reaction (5) seems to be rather small, because f o r long chains the abstraction of H atom from isopentane would be favored. Under these circumstances dJD-pH^ ydt seems to be approximately independent of isopentane concentration f o r high pressures and s l i g h t l y dependent on i t f o r low pressures. This behavior approximately agrees with experimental r e s u l t s . As f a r as the t o t a l rate i s concerned we can write d [ i - C 5 H 1 2 ^ d t = ( k l a + k l b + k l c ; fi-C 5H 1 2-j + i- k 2[i-C 5H 1 P7jpH3) * k 5 §--C5H12] [C2H5J i-* k 7 b ^ - S V ) t°3H7l ( 1 5 ) Substituting (13), (14) and (5)into (15) we get d k " d t H l ^ = ( K l a * k l b * k i c ) [ i - C 5 H i ^ 1-W*?fa^\r*]r) [ i - C 5 H 1 2 J j + 0 - 5 7 ( 3 ^ ^ ^ 0 ) + k / 20 % - k 2 7b i - c 5 H 1 2 y Li—C5H12I = k i Q * k l b J 20 k 3-k 2 ( 2 0 k 5 - k 2+l» 5 3 k 2 + 0 . 5 7 k 2 } [ i - ^ 5 H 1 2 ^ • k l c f 2 0 k 3 - k 2 - 1 . 5 3 k 2 - 0 . 5 7 \ f V i o T 3 - k 7 " ~ J L c 5 H i 2 j * + ' k 7 b ( k 7 b * k 7 a ^ 5 l ^ ^ - 0 5 H l ^ = -160-:/TKia + klb \f \-[ F o i ^ T A 2 0 k3 - 1 ' 1 ^ j L i - ^ ] * fcic^OKg-g.lKg V n 7 \ 20 k 3-k 2 A 5 1 3 [ i - C 5 H l 2 > Wb+ k 7 a / [ i . C 5 H 1 2 - j / . k l c ( 2 0 k 3 - 3.1*2/1 ^ k l a T k l b 2OK3 - K 2 K7b / L i ^ \ [ i - 0 5 H 1 2 ] k7b* k7a/^.c 5H 1 2- ] k^b i s n e g l i g i b l e while reaction (7a) i s the only route f o r C^Hy reaction, and thus the l a s t |$rm vanishes, and f i r s t ordered k i n e t i c s i s established. For higher pressures the l a s t term makes some contribution, reaction (7b) i s no longer n e g l i -g ible and a s l i g h t l y higher order r e s u l t s . Reactions ( l a ) , (lb) and ( l c ) do not differ^more than 2<-3 KCal/Mole. i n a c t i v a t i o n energy. Reaction (3) i s much more slower than reaction (2) especially at low pressures, so the ov e r a l l rate constant can take the form k over K ' ( k l a / 2 0 k 5 _ k 2 y f i n c e k l a ' klb> k l c a r e approximately the same, f o r the high temperature of 500°C. where = 3 ( 2 0 k 3 - r 2 k 2 J . and provided that reaction (2) and (3) d i f f e r by only 2<~ 3 K Cal/Mole. i t i s assumed that 2OK3 - k2C^ 2 0 k 5 - 2 k 2 so k over - k l a / 3 C r k l a > 161 Thus K corresponds to the slower i n i a t i o n process and over " consequently EA^>ver)^EA ^ l a^ but E A (la) i s the D(C-C). TO calculate the bond energy the following data are used. CH5-CH-CH2CH5 — * CH3CH CH2CH3+CH3 ^ 3 (71) D(C-C')= A^(CH 30HCH 2CH 3) + (CH3 )-AH^ ( i - C 5 H i 2 ) flH^ (CH3) = 31.5 KCal/Mole. ^ Hj(isopentane) = -36.92 KCal/Mole. For & H^(CH3CHCH2CH2); Yoevodskii" empirical rate A t e r t . = "^sec.-^ Aprim." 2 B i s u s e d So Asec.= Aprim.* A t e r t . f o r the D(C-H) values. _ Using the value D(t-C 4H g H) = 89.6 KCal/Moles and D(n-C 4H g H) = 94 KCal/Mole. Thus we f i n d D(sec. C 4H q H) -91.8 KCal/Mole. and using t h i s value i n the reaction CH 3CH 2CH 2CH 3—» CH3CHCH2CH3+ H" we get D(C-H) - (CHjCHCHgCH^) - t^H^df) -^(CHjCR^CR^CI^ So &H^(CH3CHCH2CH3^ = D(C-H) - A^H^CH) * ^ (CH5CH2CH2CH3) 162 (116) (117) = 91.8 -52 - 30.15 = 9.60 Substituting t h i s value into the previous formula f o r D(O-C) we get D(C-O) = 9.6 + 31.5 + 28.70 = 69.8 "KCal/Mole. E A(exp) = 64.1 KCal/Mole. which appears i n reasonable agree-ment with E A (calc) value, considering the inevitable error introduced by using thermochemical data. The value of 69.8 KCal/Mole. which calculated f o r homogeneous processes, must be lower at lower pressures due to heterogeneity^ so the c a l -culated value approaches s t i l l more the experimental. I I Inhibited decomposition At low pressures the main effect of NO i s the suppression of heterogeneous i n i t i a t i o n so k o v e r a i i would correspond to k i n i t i a l (komog) The a c t i v a t i o n energy i s expected to be a l i t t l e higher because surface i s deactivated by NO. At higher pressures of NO, n i t r i c oxide s t a r t s to par-t i c i p a t e i n certain reactions as suggested i n Chapter I I I . The rate must remain unchanged because NO does not in t e r f e r e i n the primary processes which are rate determining. For very high amounts of NO some homogeneous i n i t i a t i o n by NO w i l l take place, so s l i g h t acceleration i s expected, as indicated i n reaction (la) Chapt. IV). 163 The rate can be expressed by - d ( i - C 5 H 1 2 ) / . = (kia+kib+lcic) , / a T 2 0 k 3 - k ~ { 2 0 k 3 + l . l k 2 ) [ i - C 5 H 1 2 ] * k 7 b + ^ a r i - o 3 H 1 2 -but k-^ a i s very small. Isobutene and C3Hg appear unaffected by NO, because reactions (4a) and (4b) are not affected by i t . The rate of production of C2B4 i s given by d f c J 2 I f/dt = k 5[n-C 4H g3 4. k 7 a^H 5CHCH 3'] + kg [CH 3][C 2H 5] + + K l l a ^ C 2 H 5 ] 2 - k 6 ( C H 3 N 0 3[C 2H 5*] + K c[C 2H 5NO] [CHj+ + k d [ C 2 H 5 N 0 ] t 0 2 H 5 ] Normal butyl and CH3CHCH3 are not affected by NO concentration but the concentrations of CH3, and C2R"3 are constant only f o r constant NO, because they react with i t . Thus the rate appears to depend on a NO concentration to a power greater than unity. Methane has a si m i l a r dependence on NO concentration (Pig. 4 7 ) . For C2Hg r e9- c"ti° n ( a) and (b) predict an increase but remembering that the main source of C2H6 i s reaction (3), reactions (a) and (b) simply compete with (3). At low NO concentration ethyl r a d i c a l disappears almost completely through reaction (3) but at higher NO concentration the 164 i r r e v e r s i b l e process C2H5+ NO —> C2H5NO also uses up O2H3 rapidly. Thus the y i e l d of &2^6 ^s reduced by increasing P^o* 165 REFERENCES 1. Rice and Johnson, J.A.C.S. 56-214-19 (1939). 2. Rice, O.F. and Doley, M.O., J.A.C.S. 55-4245 (1933). 3. Rice, O.F. and Johnston, W.R., J.A.C.S. 56-214 (1934). 4. Rice, O.F. and Johnston, W.R., J.A.C.S. 54-3529 (1932). 5. Ingold and Lossing, J", of Can. Chem. 31-30-41 (1953). 6. Eltenton, G.C., J. Chem. Phys. 10-403 (1942). 7. Elenton, G.C., J. Chem. Phys. 15-455 (1947). 8. Elenton, G.C., J. Phys. Chem. 52-463 (1948). 9 . Kebarle, P. and Bryce, W.A., Can. J. of Chem. 35-576 (1957). 10. Prey, P.E., Ind. Eng. Chem. 26-198 (1934). 11. Maiaus, 3. 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Rev. 47-75 (1950). 117. American Petroleum I n s t i t u t e Research Project. 118. C.N. Hinshelwood, Disc. Parad Soc. 2-111 (1947). 119. C u l l i s , C P. and C.N. Hinshelwood, Dis.Par. Soc. 2-117-(1947). 120. Muleany, M.P.R. Disc. Par. Soc. 2-128 (1947). 121. W.A. Bryce, and K.U. Ingold, J. Chem. Phys. 23-10-1968 (1955). 171 APPENDIX I CALCULATION OP ACTIVATION ENERGIES AND FREQUENCY FACTORS 1. Isopentane TABLE XIV RATE CONSTANTS FOR VARIOUS PRESSURES AND TEMPERATURES ( XIO 5) mm-1 Pressure (mm) 775°k 783°k 792°k 0-5. 2.0 3.2 4.9 70 2.5 3.7 5.4 100 2.7 4.6 6.4 150 3.6 5.4 7.5 200 4.2 7.0 9.2 The E_£_ were calculated from Arrhenius equation E A = E ( T l T 2 / T „ . T l ) 2 . ^ g ( * 2 / k l ) Substituting the values reported above we get table XV. TABLE XV" Energies of a c t i v a t i o n with pressure Pressure(mm) K C a l / M o l e . E2 K C a l/Mole. E3 K°al/Mole.E K C a^Mdle 0-50 63.8 61.4 67.3 64.1±2.7 70 59.2 62.3 55.7 59.1_b.4 100 62.8 56.5 46.5 55.6 analogous 172 Frequency factors substituting E^ values i n one of the equations of the system log k_ = log. A - E log e Kxl log k? - log A - E log e we get RT 2 0-50 mm 70 mm 100 mm. log A = log k-j_ + E_ {p^ e so f o r -3 3 log A = log 2x10 +.6.4.1x10-^ 0.434 - -2.7+18 1.98 x 775 • 15.3 so 15 i 13 A = 1.98x10 min" = 3.3x10 sec 14 -1 12 log A = 14.075 so A - 1.19 x 10 min = 1.98x10 se log A = 13.126 so 13 -1 11 _ i A - 1.34 x 10 mm = 2.23 x 10 sec 150 mm log A - 13.36 or A - 2.29 x 10 1 3min" 1- 3.8 x 1 0 1 1 s e c 1 200 mm A = 2.9 x 10 5 min - 4.83 x 1 0 1 1 s e c 1 I I n-pentane TABLE XVI Rate constants f o r various pressures and temperatures ( xlO 2) m i n - 1 Pressure(mm) 795°k 802°K 10 - 30 1.75 2.9 50 1.85 2.95 100 2.7 4.2 120 3.9 5.9 173 TABLE XVII Energy of ac t i v a t i o n with pressure pressure (mm) E A KCal/Mole, 10 - 30 66.5 50 61.5 100 56.5 120 55.5 Frequency factor are treated s i m i l a r l y as above. I l l Neo-Pentane TABLE XVIII Rate constants f o r various pressures and temperatures (xlO^) min- 1 pressure (mm) :'T9J5?k_; .. 815°k 10-50 7.0 17 70 9.0 20.0 100 13.7 30.0 TABLE XIX; Activation energies with pressure pressure (mm) E^ kCal/Mole. 10 - 50 57 70 49 100 47.5 174 APPENDIX I I An a l y t i c a l r e s u l t s f o r n-pentane TABLE XX An a l y t i c a l r e s u l t s f o r n-pentane Mole. $» A P / P = 0.4 T = 519°C P n_ pent.= 7 5 Products Mole, jo 38 n-C 5Hi2 11 C2 H4 C2H'6 10 9.4 I-C4H8 3.6 TOTAL 97.0 

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