摘要:

G. B. KISTIAKOWSKY AND A. GORDON NICKLE 187 THE PYROLYSIS OF DIBENZYL BY C. HORREX* AND S. E. MILES? Received 5th February, 1951 The pyrolysis of dibenzyl has been studied a t low partial pressures and The products have been shown to be in accord using fractional decompositions. with the equation 3PhCH,CH,Ph -+ 2PhCH, + P h C H d H P h + PhH + P h C H d H , and the process has been discussed in terms of a primary dissociation into benzyl radicals with a first order rate constant k(sec.-l) = 1oSs3 exp (- 48.0 kcal.)/RT. If the energy of activation is taken to indicate the dissociation energy of the central carbon-carbon bond it agrees with anticipations based on a resonance energy of 24.5 kcal. for the benzyl radical ; the temperature independent factor is abnormally low. The possible secondary reactions have been outlined.The determination of the bond dissociation energy for the central C-C bond in dibenzyl is a problem of considerable interest in view of the recent work leading to a resonance energy of 24-5 kcal. €or the benzyl radical.1 It is known that this central bond is shorter than the normal carbon-carbon bond by 0.06 and Szwarc, in a recent discussion of these points, used Skinner’s data on the relationship between bond length and bond strength to deduce that the bond is strengthened by 13-7 kcal. * Present address : Chem. Dept., The University, St. Andrews. -f Present address : Thornton Res. Centre, Shell Refining and Marketing Co., Szwarc, Faraday SOC. Discussion, 1947, 2, 39 ; J . Chem, Physics, 1948, Jeffrey, Proc.Rqy. SOC. A , 1947, 188, 222. Skinner, Trans. Faraduy SOC., 1945, 41, 645. Ltd. 16, 138.I 88 PYROLYSIS OF DIBENZYL From the heats of formation of toluene and dibenzyl, together with the bond dissociation energy for C-H in toluene, he deduced that 47-1 kcal. is the dissociation energy of the central C-C bond of dibenzyl. This value is 11 kcal. greater than that deduced by subtracting z x 24-5 kcal. from the normal C-C bond strength of 85 kcal. in ethane, and Szwarc attributed this 11 kcal. to the bond shortening effect, thus obtaining reasonable agreement between the two methods of estimation. Before the publication of his work, an attempt at a direct determin- ation of the bond dissociation energy in dibenzyl had been started. This was based on earlier observation of Horrex and Szwarc on the pyrolysis of benzyl iodide, that the benzyl radical had considerable stability.Other relevant qualitative data have been summarized by Szwarc.1 The pyrolysis of dibenzyl has not been the subject of kinetic study under conditions where it might be Fossible to derive the mechanism of the decomposition. Barbier 4 distilled dibenzyl through a red hot tube and decomposed it into toluene and stilbene, with smaller amounts of phenanthrene. The molecular ratio toluene/stilbene was observed to be approximately 211 and this was confirmed by Graebe.6 The most suitable experimental conditions for the investigation of the primary step in the pyrolysis should be where only small fractional decompositions result (cp. Butler and Polanyi 6 for organic iodides). A conventional flow technique is required for such an investigation and the mechanism has to be derived from the analysis of stable end-products.Since the benzyl radical has a high degree of stability, it was considered worth while to see if the probable primary dissociation of dibenzyl into two benzyl radicals could be observed more directly. For this purpose, very low pressures and moderate temperatures would be favourable, and the use of a quartz fibre manometer in a heated closed system con- taining dibenzyl vapour was therefore considered. Results derived from the application of both techniques are described below. The manometric technique was applied first and showed that secondary reactions were proceeding ; this necessitated a change to the flow technique.Experimental 1. Quartz Fibre Manometer Experiments Haber and Kerschbaum' showed that the rate of damping of the vibra- tions of a quartz fibre in a gas a t a pressure such that the mean free path of the molecules was much greater than the fibre diameter, was given by the equation b/t = u +pM*, where t is the time 01 half damping, a and b are constants characteristic of the fibre and the temperature, M is the molecular weight of the gas and p its pressure. For use as a manometer, direct calibration of t against a McLeod gauge has to be made and a and b determined empirically. It seems that no use has been made of the instrument a t elevated temperatures. was constructed and calibrated for pressures of 0-2400 p air and a t temperatures from room t o 1000' K.The detailed results, which are t o be given and discussed elsewhere,O show that the instrument has distinct disadvantages for our purpose in that a and b vary markedly with temperature, the variation of a being particularly serious in regions which we expected to be of interest to us (Fig. I). At con- stant temperature the instrument obeyed the above equation w7ell up to pres- sures of 50 p of dry air a t which the mean free path is about 10 times the fibre diameter. In using the calibration with air when measuring pressures of other gases, the accommodation coefficient was assumed substantially identical for all particles involved in momentum transfer with the vibrating fibre. A bifilar type of the manometer as described by Coolidge Barbier, Compt.rend., 1873, 1770. Haber and Kerschbaum, 2. Elektrochem., 1914. 20, 296. Coolidge, I Amer. Chew. SOC.. 1923, 45, 1637. Horrex and Miles (to be published). 6 Graebe, Annalen, 167, 161. fi Butler and Polanyi, Trans. Faraday SOC., 1943, 39, 19.C. HORREX AND S. E. MILES 189 The tube containing the quartz fibre manometer projected from its surround- ing furnace and was joined to a sample container through a breakable seal. The samples were held a t constant temperature and the products of pyrolysis could be removed, if required, through a valve. As a check on the technique, vapour pressures of iodine and dibenzyl a t ordinary temperatures were deter- mined and the dissociation of iodine was followed using furnace temperatures from 523 t o 983' K and vapour pressures of 41 to 26 p.The observation of t and use of the relevant a and b values for that temperature permitted PM*, and hence M , to be evaluated. The progressive dissociation of the iodine a t increasing temperatures was clearly shown and the calculated equilibrium constants were in broad agreement with Bodenstein's values, but the heat of dissociation calculated from our data was some 30 yo low. Two sources of error probably contributed to this. The previously mentioned assumption of constant accom- modation coefficients cannot be satisfactory a t high percentages of dissociation ; also the derived value of the equilibrium constant was very sensitive to small ob- servational errors in t . It was concluded from these experiments with iodine that the dissociation process could be shown with sufficient accuracy t o warrant experiments with dibenzyl, but that any heat of dissociation would be subject to considerable uncertainties.Initial experiments with di- benzyl were made by raising the furnace temperature in steps and taking frequent readings at each temperature and using pressures of dibenzyl of about 5 p. These readings showed that up to 460 & I O O C the values of pM*, calculated with the aid of the FIG. I.-Va,riation of a for the quartz-fibre manometer. appropriate u and b values (Fig. I), remained steady, but above this temperature a marked increase in $M* was evident and this was more pronounced the higher the temperature and the longer the time of contact. The increase was irreversible and unaffected by subsequent lowering of the temperature below 460°C.A simple dissoci- ation process should have given a decrease as with iodine and it was concluded that new molecules were being produced by pyrolysis and each was contributing its part to the sum ZflM4. An attempt was made to discover whether the value of pM* was lower than normal immediately after dibenzyl reached the furnace by studying the variation of its value with contact time. The quantity (pM*),/(pM*), was plotted against the time of reaction, but extrapolation was made uncertain by the difficulty of fixing the zero of time when the dibenzyl had developed its full vapour pressure after sudden admission to the furnace. It was clear, however, that any radical dissociation products underwent quite rapid secondary reactions even a t the low pressures employed.Some evidence as t o the nature of the products produced was obtained by noting, in an experiment a t 725" C, that pM* fell from 306 to 26 when liquid air was applied to the sample tube. By measuring the pressure of this non- condensible gas by a McLeod gauge, and correcting for the expansion, the value of p in the manometer system was found to be r 1-7 p and hence M = 4-9. Thus a great part of the gas must have been hydrogen. To obtain further information on the nature and rate of the decomposition a flow technique was adopted. 11. Flow Technique Experiments Procedure.-The apparatus shown diagrammatically in Fig. z was con- structed of Pyrex glass with the exception of the furnace tubes which were madeI 90 PYROLYSIS OF DIBENZYL of silica.In essence, a stream of pure nitrogen picked up tne vapour of di- benzyl, transported it through a furnace, was freed from the products of reaction and unchanged dibenzyl, and returned via a device for measurement of its rate of flow to the circulating pump. Cylinder nitrogen was cooled to the temperature of liquid oxygen to free it from volatile matter and passed very slowly over sodium in two traps heated to 300'C in order to free i t from oxygen. The purified material was stored in sufficient quantity to provide a supply for all the experiments. The purified nitrogen a t about 2 to 3 mm. pressure was circulated by the mercury pump P and freed from mercury vapour by passage through a liquid air trap followed by an arrangement W permitting alternate heating and cooling of the gas in order to remove any slight mercury mist.The dibenzyl could be placed in U-tube 3, or in a metal boat in the preheated section Q when higher partial pressures were required. The partial pressures of dibenzyl employed were I O - ~ to 4 x 10-1 mm. and to prevent condensation in the inlet tubes the latter were arranged to be suitably heated. The reaction vessel (diam. 2-7 cm.) was in a steel tube furnace with tapped windings arranged to provide a length of 32 cm. with a temperature constant to 2 O C and with a sharp fall a t each 1 FIG. 2.-Diagram of apparatus. extremity. The use of a constant voltage source on the furnace windings per- mitted the mean temperature of a run to be maintained constant to f I' C by hand control. The time of contact (0.25 to I sec.) in the furnace and tem- perature were chosen to give decompositions ranging in the main from I to 15 yo.Limitations of the analytical techniques developed prevented smaller decom- positions being employed and runs lasted about Q to 2 hr. The gas issuing from the furnace passed through trap I maintained at a tem- perature of oo C to - 22' C, which was suitable for the removal of the less volatile products from the stream. All the remaining products save hydrogen and methane were removed in trap 2 cooled in liquid air. The hydrogen and methane passed through a furnace containing copper oxide, the temperature of this being adjustable so that the combustion of H, could be achieved whle pyrolysis pro- ceeded. The methane was determined separately by raising the temperature and recirculating later.The products of these combustions were removed in U-tube 4, and the flow rate of the carrier gas deduced from the pressure drop across a suitable capillary system D inserted in the return line to the circulation pump. Constant flow rates were achieved by operating this pump a t a suitable constant temperature by using a Woods metal bath with electrical heating con- trolled by a Sunvic energy regulator. Since the dibenzyl was not separated by taps from the furnace, checks were made on the loss sustained during the preliminary evacuation of the apparatus with the dibenzyl at room temperature and of the loss due to diffusion during a period of 70 min. in presence of 3 mm. N, but without circulation of the latter.The figures were 0-4 yo and 0-16 yo of the amount transferred through the furnace during the experiments. Thus the error introduced a t the beginning and end of an experiment was negligible. Analysis of Products .-For convenience these are divided into three groups : (i) gaseous, (ii) liquid and (iii) solid. (i) THE GASEOUS PRODUCTS include H,, CH, and any C, hydrocarbons. The latter would condense in trap 2 when their partial pressures exceeded about 0.05 mm., which in the known volume of the apparatus corresponds to 10-5C. HORREX AND S. E. MILES mole. By isolating trap 2 a t the end of an experiment, raising the mercury to mark C and replacing the liquid air by a bath a t - 78" C, the rise in pressure in the calibrated volume (after correction for temperature change), permitted an estimation of the condensed C, hydrocarbons.In expt. 39 where 5-55 x I O - ~ mole dibenzyl were passed through the furnace and 15 yo was decomposed to toluene, only 5-2 x I O - ~ mole C , hydrocarbons were found, i.e. less than 0-2 yo decomposition. The amount of products combusted by the copper oxide could be determined by isolating and warming U-tube 4 successively to - 78" C and room tem- perature and noting the pressure developed in the calibrated volume. In almost all cases the hydrogen production in the reaction was small and saturation pressures of water vapour were not reached. A check on the CH, production was made by determinations of the total pressure rise in the circulation system at the end of an experiment using carefully standardized conditions of com- parison with initial conditions.The results by both methods were in fair agreement, showing methane production to be negligible. (ii) THE LIQUID PRODUCTS collected in trap 2 and were transferred to the detachable limb R by vacuum distillation. By raising the mercury to A and placing baths a t a series of temperatures on R the vapour pressure curve could be obtained and served as an indication of substances which might possibly be present. By lowering the mercury to B or C the liquid could, in some cases, be completely vaporized and the pressure exerted in the fixed volume recorded. By recondensing the liquid in R the latter could be weighed and the total amount of liquid products found. The quantity available from a single run was about By consideration of the probable modes of fission of dibenzyl the liquid pro- ducts might contain benzene, toluene, ethyl benzene and styrene.The vapour pressure curve from - 38°C to 25°C lay about midway between the vapour pressure against temperature curves of benzene and tduene, and well above those of ethyl benzene and styrene. A further qualitative indication was that the melting point of the mixture seemed to be over the range - 120' to - 108O C . The melting point of benzene is 5" C, toluene - 102" C, ethyl benzene - 93" C and styrene - 33"C, so that the two pieces of evidence suggested that a con- siderable amount of toluene was present and probably some benzene, but did not enable any decision on the others to be made.By arranging a suitable additional attachment a t R the liquid products were vacuum distilled into previously degassed sulphuric acid, allowed to react a t room temperature for 5 min. and then transferred back to the original con- tainer. By measurement of the pressure of the completely vaporized liquid before and after such treatment i t was found that some I I to 25 % of the total had been absorbed by the acid. Blank tests on benzene, toluene and ethyl- benzene showed less than 2 to 5 "/p loss. The result suggested that styrene might be present and further qualitative indications were obtained from re- fractive index data. The refractive index of the liquid products was in the range 1-5050-1-5136 ; styrene has a value of 1-5379 and the benzene, toluene, ethyl benzene values are rather similar and lower, viz.1-500g-1.4g70 a t 17.2" C. The use of synthetic mixtures suggested that the styrene proportion was definitely 20 yo or more. By using bromine in glacial acetic acid under conditions ( I hr. at room temperature in the dark) such that styrene was estimated with 95 yo accuracy and reaction with the saturated hydrocarbons was negligible, it was found that the styrene present in the liquid products was 21 f 4 yo weightlweight. These analyses covered the whole range of decompositions studied and no significant trend in analysis with the percentage decomposition was evident. Morton and Mahoney's lo technique for fractional distillation of small amounts of liquid was next applied. The fractionating tubes used were 7.5 cm.x 1.5 mm. and packed with powdered glass wool, and the boiling point of the distillate samples was determined by Emich's method. The results, examples of which are given in Fig. 3 and 4, were of the same type whether the pyrolysis temperature had resulted in 15 yo or 0.5 yo decomposition, or if even lower temperatures had been used with a source of benzyl radicals present. For comparison purposes, blank runs on synthetic mixtures containing various proportions of benzene, toluene, styrene and ethylbenzene were carried out. The distillation curves of the pyrolysis products indicated that benzene (b.p. 80-81'), toluene (b.p. 108-IIO"), and styrene (b.p. 144-147' ) might be present, loMorton and Mahoney, Ind. Eng. Chem. (Anal.), 1941, 13, 494. 0'1 to 0'2 g.192 PYROLYSIS O F DIBENZYI, but the possibility of some ethylbenzene (b.p.134-136') being in the last 30 yo or SO of the distillate could not be eliminated. The results of blank experiments using 49.7 % benzene + 29.2 % ethylbenzene + 21.1 % styrene, and 50-5 % ben- zene + 49.5 yo ethylbenzene are superimposed in Fig. 4 and show that the central portion of the I' product data " is given by a genuine fractionation separating - /40 . B.Pt C 'C E - 0 f40 1 B.Pt - "C - /20 I FIG. 4. 0 Run 58 ; - 50'5 "/o C6H6, 49'5 O/o CSHI,. 49'7 "/o c6H6, 29'2 "/o C ~ H I ~ , 21.1 "/o CeH,. - - - - FIGS. 3 and +-Micro fractionation curves for liquid products and synthetic mixtures. toluene a t about 108~-112~ C. This is emphasized in Fig. 3 by the curve given by a blank of composition 20 yo benzene + 59 yo toluene + 21 % styrene, and it also appears from comparison with the latter that the materials in the pyrolysis products which were less volatile than toluene might amount to 25 yo.Due to the small amounts of materials involved (0*020-0.040 g.) and errors caused by fractionation during the determination of the boiling points of the samples,C. HORREX AND S. E. MILES 1 93 there is distinct scatter in the observations but comparisons with blanks sug- gested that the composition of the liquid products of pyrolysis was 20 f 5 Yo benzene, 25 f 5 yo styrene (with ethylbenzene), with residual toluene 55 & 10 yo uy weight. At a late stage in the work a Beckman Model D.U. photoelectric spectro- photometer became available and the u.-v.absorption curves of Fig. 5 for typical tiquid pyrolysis products were obtained. Also included on the graph are the curves given by a blank containing by weight 19 Yo benzene, 59 yo toluene and 21 yo styrene and another for 30 yo benzene, 40 yo toluene and 30 yo ethyl- benzene. The correspondence between the solution containing styrene and the unknowns is close ; due to the much higher extinction coefficient of styrene the characteristics of the first three graphs are essentially those of styrene and the small halt a t 2680 A is the main indication of the presence of other hydro- carbons. It is presumably due to the peak present in that region for both toluene and ethylbenzene. To be assured of the accuracy of these deductions, the calculated optical density of the known solution was compared with the observed a t suitable intervals, using calibration curves foi the pure component obtained on the same instrument.Good agreement resulted ; a t 2680 wheIe the contribution of toluene to the total is proportionally greatest, the result was D (obs.) = 0.14, while D (calc.) = 0-00084 (benzene) f 0.0522 (toluene) + 0.0889 (styrene) = 0.1419. The two unknown solutions show an absorption above 2940 A which is not found in the blank. It is known that the solid products, from which the liquid products had been separated by high-vacuum distillation, contained stilbene which absorbs strongly in this region. The presence of a trace of this would produce an error in evaluating styrene concentrations from the absorption maxima a t 2810 and 2910 A.As styrene does not absorb at 3100 A, a stilbene estimation (c = 20,ooo) giving 0~000108 g./L tor Iun 80, indicated that the 0.0348 g./L liquid products present in the solution contained 0-29 yo stilbene as impurity. At 2910 and 2810 di this would lead, with E (styrene) 446 and 740 and (stilbene) 21,400 and 19,050, to weight % styrene values of 26.6 and 25.3 for the liquid product. Without this correction for stilbene, and assuming the total absorption a t the two peaks was due to styrene, an average value of 34 weight Yo was obtained. Run 81 gave 20 yo with the correction and 25 yo without it. Since the evidence from refractive indices, bromination and u.-v. data favoured a styrene concentration of 20 to 25 weight yo in the liquid products, i t was concluded that the 25 f 5 yo fraction found by micro-fractionation was largely styrene and the ethylbenzene was not appreciable.(iii) THE SOLID PRODUCTS of pyrolysis condensed with a large excess of unchanged dibenzyl as a white solid ; there was no evidence of carbon in this nor in the reaction zone. In view of the earlier work of Graebe, and some evidence concerning the action of methyl radicals on dibenzyl, stilbene was considered as a possible product. Due to its low concentration in the solid fraction from normal runs, crystallization and sublimation brought about little separation. From runs where a high proportion of dibenzyl was decomposed (by means of benzyl radicals produced from benzyl iodide), stilbene was isolated and checked by mixed melting point.Bxomine in glacial acetic acid was used to estimate the unsaturated fraction in the solid products, blank tests having been done on dibenzyl and phen- anthrene under the same conditions. The u.-v. absorption curve for the solid products in ethanol (Fig. 6) combined characteristics of stilbene and dibenzyl. Phenanthrene, which absorbs in the same region as these substances, continues to absorb up to 3750 A and the absence of its characteristic peaks after 3400 A was taken to indicate its absence from the solid. At 3940 the absorption of the ethanol solution of the solid (0.318 g./1.) from Run 80 was ascribed to stilbene (E molar = z3,5co) and this gave 3-5 weight yo stilbene in the solid. Bromination of the same solid gave a value of 3-1 yo while for Run 81 the u.-v.absorption data gave 1-6 yo and bromin- ation 1-5 yo. The analyses based on bromination have been used for the results quoted below as the spectrophotometer was not available sufficiently early in the work. From the composition of the liquid products we derive a molecular ratio of 1-0 f 0*23/2-3 f 0*42/o-g2 f 0.2 for benzene/toluene/styrene. The toluene/ stilbene ratio varied more widely, from 1*51/1 to 311, and we ascribe this to difficulties in determining the small amounts of stilbene by bromination. Since GJ 91 PYROLYSIS OF DIBENZYL the gaseous products were negligible we conclude that the overall decomposj tion can be represented by 3PhCHzCHzPh -+ 2PhCH, + PhH + PhCH=CH, + PhCH=CHPh. In order to justify the assumption that toluene is formed by benzyl radicals attacking dibenzyl, runs were conducted with the flow technique as previously described but using lower temperatures and dibenzyl plus benzyl iodide as the reacting system.From previous work i t is known that even at 450' C the benzyl iodide is 60 yo decomposed in the time which produced about 2 yo de- composition of dibenzyl a t 700° C. Using equimolecular proportions of dibenzyl and benzyl iodide at 701' C (partial pressures 0.045 mm.), the liquid products increased 25-fold and seemed by distillation characteristics to be the same as noted for dibenzyl alone, toluene being a major constituent. The use of a carrier gas containing 53 yo nitric oxide in nitrogen instead of pure nitrogen produced no significant change in the rate of decomposition of dibenzyl a t 752' C and was not further investigated. I I t I I FIG.5.-U.-v. absorption curves for liquid products and synthetic mixtures in ethanol. FIG. 6.-U.-v. absorption curves for solid products in ethanol. Kinetic Data on the Reaction.-A test of the homogeneity of the reaction was made by packing the furnace with silica wool of approximate diameter 0-01 mm. and thereby increasing the surface exposed by 16 f 5 times. The products of the reaction seemed identical with those from the open tube and the liquid portion on micro-fractionation gave a curve indistinguishable from normal runs. The values of a calculated first order constant were about 60 % higher than the value derived without packing (Runs 55 and 77) ; this indicates that with the normal furnace the reaction was substantially homogeneous and rather less than 4 yo of the total rate might have been due to a surface reaction.The partial pressure of dibenzyl was varied by a factor of fifteen (o*ozz to 0.35 mm.) and the time of contact by nearly five times (0.17 to 0.80 sec.). With partial pressures of less than 0.07 mm. of dibenzyl reproducibility decreased as analysis errors became more important due to the small amounts of products available. The data were analyzed on a first order basis assuming that the toluene found was a measure of the initially formed benzyl radicals. The toluene was taken to be 55 yo of the weight of liquid products as a result of the previous work. As can be seen from Table I by comparing runs a t about the same temperature, the first order assumption seemed valid.To illustrate the effect of partial pressure variations, the following pairs of runs can be taken,C. HORREX AND S. E. MILES I95 No. 45, 61, No. 47, 74, No. 76, 68, while for variation of contact times and pressures jointly, No. 59, 58, No. 74, 80, No. 45, 44 are suitable. The data for log k against r/T0 K are plotted in Fig. 7 and the method of least mean squares has been applied to the values for experiments a t partial pressures greater than 0.07 mm. in order to obtain the best straight line. The data at lower pressures clearly are in general agreement with the other data but as indicated above the analytical uncertainties were known to be much greater. The results are repre- sented by the equation, (48.0 & 1.0 kcal.) x 103 2'303 RT log,, k = g 29 f 0'22 - FIG.7.-Variation of velocity constants with temperature. Line refers to high partial pressure data. High partial pressure- Low partial pressure- A Long contact time. v Long contact time. Short contact time. 0 Short contact time. Discussion Mechanism of the Decomposition.-Three possibilities can be con- (4 PhCH,CH,Ph -+ ZPhCH, (b) (4 The activation energy for reaction (a) js 84-4 + 13-7 - 2 x resonance energy of benzyl radical in kcal. as indicated earlier, and with the recent value of 24-5 for the last term, the dissocia.tion would require 49 kcal. For reaction (b) we consider that the activation energy will be rather less than 77-5 kcal. found for the dissociation of a hydrogen atom from toluene. In (b) we are concerned with a secondary hydrogen atom, and the radical produced should have a resonance energy of a t least the same magnitude as the benzyl radical ; moreover it appears that the benzyl radical can abstract hydrogen atoms from dibenzyl.The Ph-C bond in dibenzyl is 0.04 A shorter than the usual C-C distance which implies appreciable sidered for the primary decomposition : -+ PhCH,Ph + H(- Q kcal.) -+ Ph + - CH,CH,Ph.1 96 PYROLYSIS OF DIBENZYL strengthening and, as in toluene, the breaking of this bond should prove more difficult than the loss of a hydrogen atom. The most favourable process is therefore dissociation into two benzyl radicals and we have provided evidence €or believing that this is followed by (d) PhCH,* + PhCH,CH,Ph --t PhCH, + PhCHCH,Ph + (77-Q).The subsequent reactions produced stilbene, styrene and benzene in ap- proximately equimolecular proportions and while insufficient evidence exists to establish a mechanism firmly some features of the problem can be discussed. Stilbene could result from the radical decomposition (e) or disproportionation (f), but the latter offers no route to the production of styrene and benzene while the former does since it can be followed by the further reactions (g) and (h). PhCHCH,Ph -+ PhCH=CHPh + H - (129-Q) kcal. (4 (f) ( g ) H+PhCk,CH,Ph + PhCH,CH,. + PhH + 12 kcal. ( h ) (i) ( j ) 2PhCHCH,Ph 3 PhCH=CHPh+PhCH,CH,Ph+ (2Q- 130) kcal. PhCH,CH, --f PhCH=CH, + H - 37-3 kcal. An alternative sequence of reactions could be PhCHCH,Ph 3 PhCH=CH, + Ph - (128-Q) Ph + PhCH,CH,Ph --f PhH + PhCHCH,Ph + (101.6-Q).Since stilbene does not result from this last alternative, the reactions (2) and ( j ) can only be regarded as proceeding parallel with (e) or (f). The heats of the reactions have been deduced using standard heats of forma- tion for the molecules and derived radical heats of formation.ll The value of Q we estimate to be about 70 kcal. The chain reaction possi- bilities of the schemes ( e ) , (g), (h) and (i), ( j ) would seem to be excluded by the endothermicities involved. The small amount of gaseous products obtained is noteworthy, par- ticularly when compared with the data for toluene. In that case the reactions PhCH, + H -+ PhCH,- + H,, PhCH, + H -+ PhH + CH,- were involved and hydrogen and methane found as major products. Step (g) is analogous to the latter and ( k ) to the former.The absence of appreciable hydrogen formation shows that ( K ) cannot occur to any extent. (I) PhCH,CH,Ph + H + PhCH, + PhCH,. + 30 kcal. would seem to be ruled out by the fact that the proportion of toluene in the products can be accounted for by (d). As it is possible to produce the PhCHCH,Ph radical in other ways a study of its behaviour i s being made. The available evidence would appear to support the view that the initial bond fission produces two benzyl radicals and if we take the observed energy of activation as the dissociation energy of the central bond in dibenzyl we find the value agrees very well with that anticipated on the basis of a resonance energy of 24 kcal. €or the benzyl radical. It must be noted, however, that the temperature independent factor is much lower than 1013, the normal value €or a bond dissociation process.It is not proposed to discuss this feature at present as more sensitive methods of analysis may permit work in the near future with lower percentages of decomposition. But even if the present kinetic analysis was in error the magnitude of the decomposition at particular temperatures remains, and unless an extensive recombination reaction of benzyl radicals is 11 Roberts and Skinner, Trans. Farday Soc., 1949, 45, 339. ( K ) PhCH,CH,Ph + H --t PhCHCH,Ph + H, + (103-Q) kcal. The additional possibility of (I)Run 42 43 55 77 41 72 73 70 45 81 57 80 58 40 39 71 44 75 56 74 59 60 35 37 38 47 48 52 68 61 67 66 63 69 64 76 635 647 693 720 C.HORREX AND S. E. MILES Temp. "C 630 642 644 675 695 700 720 743 750 75 7 793 668 692 712 727 741 753 780 699 700 732 740 747 790 675 699 710 738 751 761 774 668 Packed Furnace :ontact Time (sec.) TABLE I 0.80 0.72 0.70 0.69 0.70 0.65 0.66 0-72 0.64 0.70 0'79 0'74 0.19 0'22 0'22 0'22 0'2 I 0'2 I 0'21 0'20 0'2 I 0'20 0'2 I 0.25 0.17 0.67 0.64 0.66 0.66 0.68 0.69 0.61 0-41 0.46 0.60 0.54 Partial Press. (mm. Hg) 0.27 0.25 0.37 0.3 I 0.30 0.3 I 0.3 I 0.3 I 0.35 0.32 0.28 0.23 0.126 0.074 0.116 0.126 0.127 0.126 0.139 0.039 0.0356 0.0326 0.0267 0-0225 0.023 0.039 0.052 0.039 0.039 0.069 0.038 0.052 0.024 0.38 0.24 0-29 % Decomposition 0.40 0.79 0.54 1-08 1-65 3'35 6-50 7-50 9-15 0.242 0.53 0.89 1-62 2'47 3-98 0.945 0'991 1-55 1.525 1-84 4'37 1'45 1-65 3-26 9'25 5'25 6-00 7-20 0.904 1.045 2-50 3'52 1'1 I 2-1 I 14.8 2'0 I 197 k (sec -1) 0.00464 0'01 I0 0.0078 0.0157 0.0159 0.0256 0.0473 0.105 0'0323 0'112 0'122 0.218 0.0124 0.0241 0.0412 0.0647 0.0925 0.119 0.193 0.045 0.049 0.062 0.072 0.0949 0'212 0-021g 0.0260 0.146 0.0795 O*Ogo9 0.123 0.0507 0'0222 0-0228 0-0422 0.0664 envisaged the primary dissociation velocity constant must be close to our value. If we use the method of Butler and Polanyi and analyze by assuming k = 1013 exp (-E/RT), and use for example our value of K at 750° C, we find E = 65 kcal. Equating this to 98 - 2 x resonance energy of the benzyl radical, we obtain 16.5 kcal. for the benzyl radical resonance energy. As this is about the value expected from theoretical calculation, the experimental value of the temperature independent factor is of considerable importance. The authors wish to thank Dr. A. S. C. Lawrence for valuable dis- cussions and t o acknowledge the award to one of them (S.E.M.) of an Ellison Fellowship during the tenure of which this work was carried out. The University, Shefield. In addition we estimate the heat of formation of the PhCH,CH, radical as - 49.4 kcal. by taking 50 kcal. as the value for the dissociation energy of the carbon-iodine bond in PhCH,CH,I (Butler, Mandel and Polanyi, Trans. Faraduy SOC., 1945, 41, 298) and 95 kcal. for the carbon-hydrogen bond. PhCH,CH,-H.

ISSN:0366-9033    

DOI:10.1039/DF9511000187

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

KINETICS OF DIENE REACTIONS AT HIGH TEMPERATURES BY D. ROWLEY AND H. STEINER Received 5th February, 1951 The rates of the dimerization of butadiene t o form vinylcyclohexene and of the reaction of butadiene with ethylene to form cyclohexene have been measured in the temperature range of 400°-6000C, and a t low converions. The reactions were found to be homogeneous bimolecular associations, leading essentially to the products indicated. Combination of the rate data for the dimerization reaction with data obtained by other workers at lower tem- peratures shows that the activation energy is temperature dependent. This fact is correlated with statistical rate calculations. Such calculations are also carried out for the cyclohexene reaction, the reverse reaction of decomposition of cyclohexene into butadiene and ethylene and the resulting equilibrium. I t is shown that a cyclic tran-ition complex accounts best for all available rate data, provided adjustments are made for the vibrational frequencies assocated mith bonds to be formed or to be broken.It is assumed that these frequencies are lowered because of the extension of these bonds in the transition configurations. The reason for the study of diene reactions at high temperatures was their importance in certain cracking reactions, which, starting from paraffins and naphthenes, lead to aromatic products. Wheeler znd Woods l were the first to point out the correct key reaction for the forma- tion of cyclic bodies under these conditions by demonstrating the reaction of ethylene and butadiene to give cyclohexene.A similar reaction, namely the dimerization of butadiene to give vinylcyclohexene had been investigated previously, and the kinetics studied in a lower temperature range. We have studied both the kinetics and rate of the dimerization reaction oi butadiene, and that of ethylene and butadiene in a temperature range of 400-600° C. These experiments, and the general conclusions derived from these rate measurements, are rcported hcre. The Dimerization of Butadiene. The dimerization of butadiene has been studied by several authors under various conditions. The general conclusion derived from a number of investigations is that i t is a homogeneous bimolecular reaction, not catalyzed by oxygen, and unaffected by anti-oxidants. The primaq' product was identified as vinylcyclohexene.This compound can react with a further molecule of butadiene to give a trimer which was identified as octahydrodiphenyL2 At low temperatures in the liquid phase and under the influence of peroxides or of oxygen, butadiene forms long chain polymers, but this reaction proceeds parallel to and independent of the dimerization and can be suppressed by adding antioxidants. The rate of the dimerization reaction was studied first by Vaughan at temper- atures frcm 330" C to 400' C. It was later reinvestigated by Kistiakowsky and Ransom in the range frcm r70-3g0° C in an extended and very accurate study. , In the liquid state, the rate of the reaction was measured by Lebedew Wheeler and Woods, J . Chem. Soc., 1930, 1819. and by Robey and collaborators.6 'Wheeler, J .Chew.. SOC., 1929, 378. 3Vaughan, J . Amer. Chem. SOL, 1932, 54, 3863. 6 Robey, Wiese and Morrell, Ind. Eng. Chem., 1944, 36, 3. Alder and Richert, Ber., 1938, 71, 373. Kistiakowsky and Ransom, J . Chem. Physics, 1939, 7, 725. Lebedew and Sergienko, Compt. rend., U.S.S.R., 1935, 3, 79. T 98D. ROWLEY AND H. STEINER I99 Finally at high temperatures from 400' C to 600' C, i.e. in the cracking range, the rate of the reaction was measured by Moore, Strigaleva and Shjljaeva.' This investigation, however, can only be classed as semi- quantitative, since both temperature and reaction time were ill defined in the flow system used by these authors. Moreover, their experiments were carried out at large conversions, when extensive secondary reactions takc place, especially at the higher temperatures.In the present investigation we have tried to use well-defined con- ditions of contact time and temperature, and have restricted the reaction to only small conversions t o avoid secondary reactions. Experiment a1 Materials .-BUTAD1ENE.-crude butadiene was prepared according to the method of Kuhoff * by cracking cyclohexene vapour. It was purified further through the cuprous chloride addition compound. The final product contained not more than z yo impurities ; this small contamination being mainly ethylene. NITRoGEN.-Special oxygen-free nitrogen purified over hydroquinone was used. Apparatus.-Because of the short reaction times required for small con- versions at the elevated temperatures of the present experiments we chose a flow system.All experiments were carried out a t approximately atmospheric pressure, the partial pressure of butadiene being varied by dilution with nitrogen. The apparatus consisted of the necessary combination of flowmeters, purifying train, preheater, reactor tube, etc. Traces of oxygen were removed by passing the gases over reduced copper turnings a t 300°C. They were then dried over calcium chloride and Ascarite and led to the reactor system. In order to achieve well-defined conditions of temperature, metal reactor tubes with thick walls were used ; some were fabricated from mild steel, others from phosphor bronze. The results were identical in both types of reactor, neither was the rate affected by an increase in surface.It was possible to reduce the temperature gradient in the reaction vessel to 1.5" a t low gas flows and to 3-2' under the most unfavourable conditions of the highest gas flows. The average temperature was taken as that of the experiment. The average duration of a run was from 60-90 min. During this period the temperature could be kept constant to within IO C. The gases leaving the furnace were cooled to - 80" C to recover the products of the reaction and most of the butadiene. The uncondensed gases passed out of the system through a gas meter. The essential parts of the apparatus could be evacuated by a mercury pump backed by an oil pump to I O - ~ mm. Hg. At the end of a run, the gases condensed a t - 80' C were separated from the dimer by fractionation through a small column fitted to the product trap and provided with a reflux head cooled with solid CO,.In a trial experiment, 0.10 ml. of benzene was separated from 30 ml. of liquid butene with a loss of less than I yo of the benzene and its purity unchanged. The amount of liquid product was found by weighing. In nearly all experiments, the refractive index, the density and the bromine number were determined also, Results Products of Reaction.-The analytical results on the liquid products from the various runs are given in Table I. The last line of Table I gives the literature values for the density and refractive index, and the calculated bromine number for vinylcyclohexene. It is seen generally that the analytical data of the products approximate closely to those of pure vinylcyclohexene.There are indications from Kistiakowsky and Ransom's work that the rate of formation of trimer from butadiene and vinylcyclohexene increases rapidly with temperature. However, from an inspection of bromine numbers and refractive indices, and from a micro-fractionation of the product of the ethylene- butadiene reaction (see Fig. 3), one can conclude that trimer, if formed a t all, must be present in quantities well below 10 yo. The product gases were analyzed in Expt. 12 and IG on a Podbielniak and Bone and Wheeler apparatus. Apart from a small amount of ethylene due to contamination of the feed gas, only Moore, Strigaleva and Shiljaeva, J . Gen. Chem. U.S.S.R., 1938, 5 , 518. * Ruhoff, Orqanic Syntheses, 17, 25. 9 Podbielniak, Ind. Eng. Chem.(And.), 1931, 3, 177 ; 1933, 5, 119.2 00 DIENE REACTIONS z5 butadiene was found. A weight balance for all ingoing and outgoing products tallied very well in both experiments. We conclude from all this evidence that within the range of conversions of the present experiments the product of the dimerization reaction is vinyl- cyclohexene. Thus the reaction even at these elevated temperatures leads to the same products as were obtained previously in lower temperature ranges. Cracking reactions of butadiene are absent. "44 TABLE I.-ANALYSIS OF LIQUID PRODUCTS Expt. No. I7 16* I5 I4 8 I2* I9 I 1 22 ^^ Temp. C 445 492 544 544 588 592 593 600 650 c % Conversion 2.90 6.95 3-80 5'35 2'42 2-90 4'92 4'73 4-03 3-06 0 0.830 0.830 0.830 0.832 0.828 0.829 0.833 0.839 - - 0.8304 1.4640 1.4640 1-4640 1.4640 1.46450 - 1'4650 1-4650 1.4660 1.4626 - Bromine Number k.Brlroo g.) 2 90 290 288 2 90 2 80 28.5 290* 283 288 286 296 * Podbielniak fractionation carried out on gas. Rate Measurements .-The rate constant of the dimerization reaction was calculated according to the following formula : * where a stands for the concentration of butadiene (mole~/cm.~), Aa is the drop in concentration during the time of contact At (sec.). The results of all rate measurements are given in Table I1 : Comparison of Expt. 14 and 15, and of 11 and 12, where the partial pressure of butadiene was varied considerably, show that the reaction is of the second order. Within the range studied, the rate constants are independent of con- version as shown by a comparison of Expt.14 and 15, and 4, 10, 11 and 12. In Expt. 21, a reactor (No. 111) packed with a roll of thin iron sheeting was used, which increased the surface to volume ratio about 10 times over that of reactor 11. Comparison of Expt. 21 with 16, where reactor I1 was used, shows that the packing does not affect the rate. The reaction therefore can be con- sidered to be homogeneous. This is also borne out by the consistency of the results obtained in a number of reactors of varying volume and material (see list of reactors, Table 11). Discussion Within experimental error the rate constants if plotted in the customary way as log k against the reciprocal of the absolute temperature lie on a straight line. This is shown in Fig. I where both the present and Kistiakowsky and Ransom's 4 results are plotted.The upper part of the graph (crosses) representing the present experiments leads t o the following equation : log,, K = 11-14 - - 26800 2'30RT' * (1) 1°Lebedew and Sergienko, Zhur. Obshech. Khim, 1935, 5, 1839. Aschan, * This formula neglects diffusion against the gas stream (see Bodenstein However, a t the very low con- Bere, '9'4, 57, '959. and Wolgast, Z . physik. Cheulz., 1908, 61, 422). versions used, the error introduced by this factor is very small.D. ROWLEY AND H. STEINER 20 I in the lower temperature range Whilst Kistiakowsky and Ransom obtained 23690 2-30 RT' log,, K = 9'95 - ~ TABLE II.-RATE MEASUREMENTS-BUTADIENE DIMERIZATION ~ Expt. No. 24 I7 I6t 21 15 I4 8 9 I 0 I1 I 2 t I9 23 22 Reactor V 11 I1 III* I1 I1 I I I I I I1 IV IV Temp.("C) 418 445 492 494 544 544 592 592 592 593 600 644 650 588 Contact Time (sec.) 9.00 1-96 1-68 2-48 0.98 0.98 0.374 0.276 0.540 0.730 0.633 0.630 0.295 0.387 Initial Partial Pressure Butadiene (-. Hg) yo Conversion 6-17 2-87 7.00 3-80 5.26 2.42 1-28 2-90 4-92 4'73 3-06 3'96 I 1-3 2'12 k x 10-2 (moles-1 cm.8 sec.-1) 4-13 9-05 28-3 31.0 83.0 78.5 3 40 236 20 I 208 232 235 563 520 Reactors I Mild steel 7-46 ~ r n . ~ I1 ), 12-6 ) ) I11 ,) 18.2 ,, -packed tube 10 times surface of reactor 11. IVPhosphor 14-0 ,, VRronze 61.0 ,, * Packed tube. t Podbielniak analysis carried out. It is seen from Fig. I that the experiments of Kistiakowsky and Ransom in the lower part of the temperature range which they investigated x Rowley and Steiner 0 Kistiakowsky and Ransom. FIG.1.-Butadiene dimerization. can well be represented by eqn. (z), b u t deviations are noticeable at higher temperatures. It seems significant, however, that their points lead smoothly into the straight line relation obtained from the present experi- ments. These facts suggest that the activation energy increases over202 DIENE REACTIONS the combined temperature range of the two investigations. Such a tendency was noted already by Kistiakowsky and Ransom,' but partly attributed to intrusion of the trimerization reaction. The present results do not support this last conclusion. In order to establish with greater certainty that the deviations from the Arrhenius equation are outside the experimental error, we calculated the deviations of all experimental points from eqn.(I), i.e., A log K (expt.) - log k (eqn. ( I ) ) . The results are plotted in Fig. 2. It is seen clearly that whilst the present results, if taken by themselves, can be represented by an Arrhenius type equation, if taken together with Kistiakowsky and Ransom's points the deviations are outside the experimental error. This is brought out further by the measurements of Robey, Wiese and Morel1 * at still lower temperatures in the liquid phase, which is included in this graph, though comparison of rate data obtained in the gaseous and liquid phase respectively may not be fully significant. x Rowley and Steiner. 0 Kistiakowsky and Ranscm. m Robey, Weise and Morell. FIG. 2.-Butadiene dimerization, deviations from Arrhenius equation.Calculations applying the transition state theory to the dimerization reaction of butadiene were carried out first by Kistiakowsky and Ransom.4 Their main purpose was to account for the low temperature-independent factor, but at the same time they showed that the transition state theory might account for the deviations from a simple Arrhenius type temper- ature dependence. It is therefore of interest to extend these calculations and apply them to the present results. Kistiakowsky assumed for his calculations the average hydrocarbon vibration frequencies proposed by These authors give their results in yo butadiene dimerized per hour ; from their values the rate constants for dimerization in moles-' ~ m . ~ sec.-l were calculated using the following formula : * Data from Robey.Wiese and Morell, Ind. Eng. Chem., 1944, 3, 36. s 1 54 h i , = - I00 * 3600 - d' where S stands for the percentage dimerized per hour and d for the density of liquid butadiene at the temperature of the experiment. These data were taken from Scott et al., J . Bur. Nat. Res. Stand., 1945, 35, 39.T'K Open Chain O0 22,800 500' 23,770 600°* 24,770 900° 25,900 * Reference temperature. The experj- mental data given for comparison were interpolated for the temperatures stated from our and Kistiakowsky and Ransom's data. It is seen that the rates calculated under the assumption of a cyclic complex are too high by a factor of 4, whereas those based on a straight chain complex are too high by a factor of 10. Both these discrepancies must be considered within the errors of the method because of the un- certainties in the molecular constants.In order t o compare further to what extent the calculations can repre- sent the temperature dependence of the reaction rate we adopted the l1 Pitzer, J. Chem. Physics, 1937, 5, 469, 473. l2 Wassermann, J. Chem. SOC., 1942, 612. l3 See, for instance, Eyring, Glasstone, Laidler, The Theory of Rute ~'rocesses. The calculated rate constants are listed in Table IV. Cyclic Complex Obs. - 2 I ,400 23,070 2 1,500 24s770 24,700 26.500 26,800204 DIENE REACTIONS following method. By adjustment both of the activation energy and the temperature independent factor we made the calculation agree with the rate determined experimentally at 600' K. The adjusted equations are given below, (3) and (4) : hcyclic = 2-23 X I d 3 X T-5'2.fvib exp (- 22380/RT) . (3) (4) hstzafghtchain = 0.85 x IO" x T-s'2. flvib eXp (- Z 4 I O O / R T ) . where f and f' stand for the vibrational contribution to the partition function. We then calculated the deviations of these equations from the standard Arrhenius equation (I) drawn through the points of the present experiments (Fig. I). TABLE IV.-COMPARISON OF CALCULATED AND EXPERIMENTAL RATE CONSTANTS Temp. O K 500 600 700 800 900 Complex Cyclic open chain Cyclic open chain Cyclic open chain Cyclic open chain Cyclic open chain k calc. 1 k expt. (mole+ cm.8 sec.-1) 0.885 4'64 49'0 274 9-11 x 102 5-03 x 103 9-14 x 103 5-45 x 104 5-58 x 104 33-3 x 104 0*446* 22-4* 5-12 x 102 5.63 x 103 4.0 x 104 k adjust 0.337 0.303 22'2 22.4 4.58 x I O ~ 4-82 x 102 5-36 x 103 5-91 x 103 3-96 x IO* 3-27 x 104 * From Kistiakowsy and Ransom's data.A plot of the curve representing these deviations can be found in Fig. z where both the values for the straight chain and the cyclic complex are plotted. Marked deviations from the Arrhenius straight line are notice- able, and we note that both equations follow, though not quantitatively, the trend of the experimental points. The calculations assuming the cyclic complex agree somewhat better than those for the straight chain one, but no matter what complex is assumed, it is clear that the transition state calculations lead to the right type of deviation from the Arrhenius straight line. The Reaction of Ethylene and Butadiene It has been pointed out already that the occurrence of a reaction between ethylene and butadiene was first demonstrated by Wheeler and W0od.l In what follows, we shall describe the determination of the rate of this reaction, which was carried out in a manner very similar to that of the dimerization of butadiene described above. The reverse reaction of dissociation of cyclohexene into ethylene and butadiene has been studied by KU~h1er.l~ Thus the rate of formation and decomposition of cyclohexene, and the resulting equilibrium, which can also be calculated from thermal data, can all be compared.More- over, these reactions are so similar to the dimerization of butadiene that the transition state models used in the previous case can, with slight modifications, be applied here. l4 Kuchler, Nach.Ges. Wiss. Gottingen, 1939, I, 231.D. ROWLEY AND H. STEINER 205 Experimental Ethylene was obtained from the British Oxygen Company as " pure ethylene for medical purposes ". It was subjected to no further purification except to remove traces of oxygen as described before, Butadiene was obtained as de- scribed above. The same apparatus was used for the rate measurements and with a few modifications the method of operation remained unaltered. Method of Analysis.-The liquid product was analyzed in the manner described above, and by determination of the bromine number and of the re- fractive index. In the present experiments, the liquid product was anticipated to be a mixture of cyclohexene and vinylcyclohexene, the latter arising from the simultaneous dimerization of butadiene.The bromine number of this mixture was relied upon to give the exact amount of cyclohexene formed, the values for the two pure compounds being sufficiently different from each other. This was shown to be accurate by fractionating the combined products from several experiments (about I ml.) on a micro-column of the rotating band type. Results Products of Reaction.-A graph reproducing the distillation of the com- bined products from several runs on the micro-column is shown in Fig. 3 where the volume distilled is plotted against the distillation temperature and the Fraction I . Fraction I1 . Cyclohexene) . (Vinylcyclohexene) . FIG. 3.-Fractionation of liquid products. refractive index of the distillate. It is seen clearly that two products only are present corresponding to the two plateaux of the curve, the amount boiling below the first plateau being negligible.The characteristics of the two fractions are given in Table V. TABLE V.-ANALYSIS OF FRACTIONS FROM MICRO-DISTILLATION 84 1-4440 (83) (1'4440) 130 1'4645 (130) (1 -4650)206 DIENE REACTIONS The distillation also confirms the method of analysis based on the deter- mination of the bromine number. The bromine number of the mixture before distillation was found to be 230 g. Br/Ioo g. which, assuming only vinylcyclo- hexene and cyclohexene to be present, corresponds to a cyclohexene content of 66 yo, whereas from the distillation graph we find 66.5 yo. A full analysis of all product gases was carried out for Expt. V and XI. This showed that only ethylene and butadiene were present.A weight balance for all icgoing and outgoing materials tallied well. Rate Measurements.-The rate constant was calculated according to the equation : A (cyclohexene) At = k , (ethylene) (butadiene), where the expressions in brackets stand for the respective concentrations, and At for the contact time. TABLE VI.-RATES OF ETHYLENE-BUTADIENE REACTION Expt. No. XIV IX XI * VIII XI1 VI I1 Etr V* IV XI11 VII teactor I11 I11 I11 I11 I11 I1 IIB I11 I1 I1 I1 I I1 Reactors : Temp. (" C ) 487 507 530 559 549 5 50 553 5 90 593 594 603 62 I 648 Contact Time (sec.) 5.02 3-16 3.16 2'35 2-40 0-97 0.94 1-62 0-72 0.77 0.76 0.834 0.436 Partial Pressure (-. Hg) 3 t h ylene Buta- diene 69 69 69 69 69 69 109 79 109 69 69 69 69 - I 14-0 cm.3 phosphor bronze.[I 12-6 cm.3 mild steel. Total :onversion % 5.68 8-40 10.25 9'5 4'33 7'25 7-10 7'52 I 1-87 8.86 5-60 3-48 8'44 yo Cyclo- hexene in Liquid Product # 73 67 71 71 73 68 66 40 63 68 70 70 70 ka (mole-1 m.3 sec-1) 384 560 938 I565 1340 1230 1590 3120 4050 3470 3740 5620 8200 IIB 13.0 ~ m . ~ mild steel packed (surface of reactor II x 11). I11 61 ~ r n . ~ phosphor bronze. * Podbielnak analysis carried out on product gases. 3 % Vinylcyclohexene = IOO yo - cyclohexene. Diluted with nitrogen joo mm. Hg. The results from all experiments are collected in Table VI. In Expt. VIII, XII, VI and I1 the contact time was varied keeping the temperature as con- stant as possible ; i t is seen that the variations of rate which occur are compara- tively small and are not systematic.In Expt. X, I11 and V, the partial pressure of ethylene was varied by a factor of three with no systematic effect on the mag- nitude of the bimolecular rate constant. To obtain a variation of the partial pressure of butadiene was more difficult, because of the small rates of flow of butadiene which i t would have been necessary to use. In Expt. I11 acd V, the partial pressure of butadiene was varied from 69 to 109 mm. Hg, this repre- sents a variation of 45 yo while the two rate constants differ by 15 yo. W'hilst this is admittedly not as conclusive as would be desirable to show that the rate is proportional to the partial pressure of butadiene, nevertheless in conjunction with the rest of the experimental evidence there is no reason to doubt that the reaction proceeds according to this simple bimolecular mechanism.In Expt. I1 a reaction vessel packed with a roll of mild steel sheet was used to test for surface effects, but as comparison with VIII, XII, VI shows, no such effects are noticeable. The amount of butadiene dimer found in the liquid product corresponded quantitatively to that to be expected from the rate data obtained above.TI. ROWLEY AND 13. STEINER 207 The experimental accivation energy was obtained from a plot of log R against I / T (reproduced in Fig. 4) where the points lie reasonably well along a straight line. The activation energy calculated from this graph is 27,500 cal. and the rate as a function of temperature can be expressed by k, = 3-0 x 101Oexp (- 27,500 T ) moles-' ~ r n .~ sec.-l/ . - (5) Eqn. ( 5 ) shows the present reactio k to be a typical diene association, as demonstrated by the low temperature-independent factor, which is 103-104 times smaller than that of " normal " bimolecular reactions, and also by the low activation energy of 27,500 cal./mole. In both these features, the reaction resembles clo ;ely th2 butsdiene dimerization dealt with above. FIG. +-Temperature dependence of rate of cyclohexene formatioi, Discussion Comparison of Experimental and Calculated Activation Energies.- Evans and Warhurst l6 calculated the activation energy of the reaction of butadiene and ethylene from potential energy diagrams constructed from spectroscopic data, and from an estimation of the resonance energy of the cyclic transition complex.The calculated value for the activation energy of this reaction was 17 kcal./mole. The experimental value for the activation energy found here at 800' K was 27.5 kcal., and if reduced to the absolute zero using standard formulae 13 and the data listed in Table X one obtains E , = 25.1 kcal. This value substantially confirms the calculations of Evans and Warhurst. The Equilibrium C,H4 + C4H, C,H,,.-The monomolecular rate constant of the cyclohexeiie decomposition according to Kuchler l4 can be represented by Combining eqn. (5) and (6), we obtain eqn. (7) for the equilibrium constant of the reaction C2H4 + C4H, 2 C,H,, k d = 9.0 x 1o12exp (- 57,5oo/RT) (sec.-l). . - (6) where k, stands for the rate of the association reaction. also be calculated from thermal and spectroscopic data.The equilibrium between cyclohexene, butadiene and ethylene can To obtain the Evans and Warhurst, Trans. Faraday SOC., 1938, 34, 614.208 DIENE REACTIONS 3 3 41 47 l - -- I heat of reaction, the heats of formation of benzene, l7 ethane Is and butane were combined with the heat of hydrogenation of benzene l9 and cyclo- hexene l9 respectively to cyclohexane, and ethylene to ethane and buta- diene to butane. This allowed us to calculate the heats of formation of 294 163 136 2-56 X I the reactants involved at 355" K.* 290 I45 I45 2-47 x 10-57 I2 Cyclohexene (g) . . -2820 cal./mole Ethylene (g) * 11,614 ,, Butadiene (g) . * 25,490 3 , The heat of reaction at 800° K was calculatea using these heats of formation corrected for the increase in vibrational heat content.This correction was calculated from the following sources : ethylene from Guggenheim, 2 o butadiene from Wassermann, l2 cyclohexene from Table X, using Pitzer's rules.'l In this way one arrives at a value of AH,",, of 37,860 cal./mole, and of AEioo = 36,280 cal./mole. The latter value can be compared directly with the one derived from our and Kiichler's activation energies, AEio0 expt. = 30,000 cal. /mole. The disagreement is cowiderable, but it is unlikely that the thermal value is in error by more than a few hundred calories. The value obtained by us for the activation energy of the association reaction is as expected of the order of that of the butadiene dimerization. This fact suggests that Kuchler's activation energy for the dissociation may be in error.The absolute value of the equilibrium constant was calculated from the calculated heat of reaction and partition functions, using spectro- scopic data. The data for cyclohexene are summarized in Table VII-X, and those for butadiene and ethylene were taken from the sources men- tioned.12, 2 O TABLE VIL-DIVISION OF DEGREES OF FREEDOM Translation . Rotation . Int. rotation Vibration . C y clohexene Cyclic Complex Str. Chain Complex 3 3 3 38 - 47 - TABLE VIII.-MOMENTS OF INERTIA AND SYMMETRY NUMBERS Benzene (g. cm.3 x 1040) Cyclohexene and Cyclic Complex (g. cm.2 x 1040) Str. Chain Complex (g. cm.a x 1040) 629 629 2-20 x I O - ~ ~ 12'2 1 17 Prosser, Gilmont and Rossini, J . Res. Nut. Bur. Stand., 1945, 34, 65. 18 Rossini, J. Res. Nut. Buy. Stand., 1934, 21, 13.l9 Kistiakowsky, Ruhoff, Smith and Vaughan, J . Amer. Chem. Soc., 1936, 146. * The data used were adjusted from 298" K to 355" K by using specific heat data listed by Parks and Huffmann, Free Energy of Organic Substances (The Chemical Catalogue Co., New York, 1932), pp. 68, 93. 2O Guggenheim, Trans. Farnday Soc., 1941, 37, 101.D. ROWLEY AND H. STEINER 209 Rotation around Torsional Vibration Bond around Bond 1-2 5-6 4-5 - 1, 1 2 2-3 I 3 3-4 TABLE IX.-MOMENTS OF INERTIA OF INTERNAL ROTATIONS AND TORSIONAL VIBRATIONS (STR. CHAIN COMPLEX) I (g. cm.2) 3.0 x I O - ~ O 28 x I O - ~ O 50 x I O - ~ O Mode of Vibration I. C-H stretching . 2' ">C bending . H 3. H-4-C H--C---C ">,-C bending . 13 H JC-C 4. C-C stretching . 5. k C stretching . 6. C-C stretching .7. c-c--c c-c-c C-C-C bending . C--c-c 8. C-C-C bending . c-c-c 9. C---C torsion * 10. C - - - C torsiont . Frequency (cm-1) - 3000 I440 950 1000 I 600 I335 320 190 407 190 * Between atoms 5-6 of stra t Between atoms 4-5 of stra No. of Modes in : yclohexene I0 4 16 5 I - 6 - - - Cyclic Complex I 0 4 16 3 I - - 7 - - 3pen Chain Complex I 0 4 I4 2 2 4 - I I h t chain complex. ht chain complex. This calculation leads t o a value of K, = 6-65 x 104 mole-1 ~ m . ~ at SooOC, compared with the experimental value of 5-9 x 105. This cal- culated value was arrived at using Pitzer's frequencies (320 cm.-1) l1 for the bending vibration of cyclohexene ; according to Wassermann much lower bending frequencies occur in cyclic molecules. Using a value of 190 cm.-l for the latter frequencies, and making the necessary adjust- ments, we obtain for K, (low frequencies) = 1-55 x I O ~ .Comparison2 I 0 DIENE REACTIONS with the experimental equilibrium constant shows that the normal fre- quency value is about g times too low, and the low frequency value 3 times too high. These deviations, though large, probably fall within the error introduced by the rather crude molecular constants uscd. Statistical Rate Calculations.-The present reaction of butadiene and ethylene is very similar to the butadiene dimerization, and calculation of its rate according to thz transition state mcthod can be carried out using suitably modified complexes of the cyclic or straight chain free radical type (cf. Fig. 5 ) . Moreover, using the molecular models and par- tition functions for cyclohexeiie derived in the last paragraph, the rate of Free radical complex I 2 3 4 5 6 I 1 I I I 1 CHZ-CH -CH,-CH=-- CET---CH, I Cyclic complex FIG.5.-Configuration of transition state complexes. decomposition of cyclohexene can be calculated and com- pared with Kiichler’s values. The molecular constants used for the two complexes are given in Tables VII-X. They were calculated strictly analogous to the method de- scribed in detail by Wasser- mann for the butadiene dimerization.ll For the cyclic complex, Wassermann’s as- sumption of the bending frequencies (190 cm.-1) was adopted, but calculations were also carried out using Pitzer’s value of 320 cm.-l. The torsional vibrations of the straight chain complex were calculated according t o Ki~tiakowskv.~ The rate of the dissociation reaction of cyclohexene was calculated for the same alter- natives of straight chain and cyclic complex, the latter with high and low bending frequencies respectively.In Table XI are listed the rate constants calculated under the three alternative assumptions. Inspection of Table XI reveals clearly that the rates based on the assumption of a straight chain complex, both for the association and dissociation reaction, are too high by factors ranging from 102 t o 104. Since there is no doubt that such discrepancies exceed even the consider- able error of the present calculations, these results show that in the present reaction the straight chain complex cannot be the correct one. A model stiffer than the present one would agree better with experi- ment, but it is difficult to see how this could be visualized without intro- ducing quite arbitrary assumptions.On the other hand the rate constants assuming a cyclic model all agree within a factor of m with the experi- mental values (see Table XI). In the present case it is even unnecessary to use the low bending frequencies proposed by Wassermann, since quite good agreement can be obtained with Pitzer’s frequencies, as was also the case for the equilibrium between cyclohexene, ethylene and butadiene discussed previously. On the other hand it will be recalled that in the butadiene dimerization the cyclic complex led t o satisfactory values for the rate constant only if the low bending frequencies were used. One can hardly assume that these bending frequencies occur in the butadiene dimer complex but not in the very similar cyclohexene complex and in the fully formed cyclo- hexene molecule, and one is led therefore to the following alternatives.Either the cyclic complex with low bending frequencies is the correct transition state in both reactions, and such low frequencies occur also in the fully formed cyclohexene molecule, or a cyclic complex with normal bending frequencies is operative in the cyclohexene formation, and aD. ROWLEY AND IH. STEINER 21 I straight chain one in the butadiene dimerization. The second alternative implies a considerable change in mechanism in going from the one to the other reaction which, while it cannot be excluded, seems unlikely. If the first alternative is preferred, it is necessary to examine the evidence for low bending frequencies in cyclic structures.The Raman spectrum of cyclohexene has been examined by a number of authors,21 amongst tht m Weiler, who observed lines corresponding to frequencies of 176 and 273 cm.-'. These are frequencies of the magnitude required by Wassermann. However, the intensity of the lines is given as very weak and by analogy to similar lines found in the benzene spectrum i t is most probable that they have to be assigned to combination frequencizs and not to fundamentals. The line of lowest frequency having an appreciable intensity is that at 396 cm.-l, which corresponds more nearly with Pjtzer's estimation of 320 crn.-l for C-C-C bending vibrations. The evidence therefore from spectroscopic data is doubtful ; one cannot say that it clearly supports the assumption of low bending frequencies.TABLE XI.-CALCULATED RATE CONSTANTS AT 800' K Association reaction. k (mole-l ~ r n . ~ sec.-l) Cyclic Complex I Str. Chain Complex Expt. 1-45 x 10' (v = 190 crn.-l) (v = 320 cm.-l) 5-8 x 102 15'3 8.7 x I 0 2 Dissociation reaction. k (sec.-l) 1-33 X I O a 1-66 x I O - ~ (U = 190 cm.-l) 2-25 x I O - ~ (U = 320 cm.-l) 1-30 x I O - ~ Apart from spectroscopic data, an accurate determination of the entropy of cyclohexene is available. 4 2 From this, using standard formulae and the molecular data of Tables VII-X, the vibrational entropy of cyclo- hexene was calculated. This " calorimetric " value is compared in Table XI1 with the calculated value using (a) Pitzer's frequencies for all vibrations, and (b) Wassermann's frequencies for the bending vibrations of C-C-C and C-C=C bonds (190 cm.-l) and Pitzer's values for all other vibrations.It is obvious that Pitzer's values agree much bettef than Wassermann's value with the experimental thermal determination. Since the evidence from the thermal determination of the vibrational entropy of cyclohexene strongly supports Pitzer's frequencies, it does not appear possible to account for the reaction rates by the adoption of a uniformly low frequency of the order of 190 cm.-l for aZZ bending fre- quencies in the transition complexes. However, it seems to us that the following compromise solution not only has the merit of agreeing well with all experimental data, as will be seen below, but seems in itself logically more consistent than the previous assumption of uniformly low bending frequencies.We propose to retain Pitzer's frequencies for all C-C-C and C-C=C bending vibrations of the stable molecules cyclo- hexene and vinylcyclohexene respectively, and equally for all bending modes of the reaction complexes, which are not directly involved in the formation or in the breaking of new bonds. On the other hand, for all 21 For references see Hibben, The Raman Effect and its Application (Reinhold, New York, 1g3g), p. 218. 22 Parks and Huffman, J. Anrev. Chem. Sot., 1930, 52, 4381.212 DIENE REACTIONS bending frequencies involving bonds to be formed or to be broken we shall assume Wassermann’s value of xgo cm.-l. These assumptions seem to us justified because in the transition complex the atoms linked by bonds to be formed or to be broken are displaced from their equilibrium positions.Consequently the binding forces must be weakened considerably, which should lead to a reduction of the frequencies of all vibrational modes affected. TABLE XII.-COMPARISON OF VIBRATIONAL ENTROPY OF CYCLOHEXENE (Units : cal. deg. -1 mole -l) Calculated Value I Calorimetric Value Pitzer’s Frequencies I ~~~~~~ I I 10’2 1 9’7 1 15‘4 In Table XI11 we indicate the splitting of the vibrational bending frequencies in the cyclic complexes into those which are stable and thus will be assumed to have Pitzer’s frequency of 320 cm.-l, and those which are associated with new bonds to be formed or to be broken, and therefore will be assumed to have a reduced frequency of 190 cm.-l.TABLE XIII.-BENDING FREQUENCIES IN MODIFIED CYCLIC COMPLEXES Butadiene I I Cyclohexene Frequency (cm.-l) Number of bending modes 4 2 5 4 Assuming a cyclic complex, and basing it on this new splitting of bending frequencies, the rates of the association and dissociation reaction of cyclohexene and the rate of dimerization of butadiene were recalculated. The results are tabulated in Table XIV, where they are compared with the experimental values both for the cyclohexene and butadiene dimer- ization reaction already given in earlier tables. The results given in Table XIV show that the rate constants calculated under these last assumptions, with the exception of the dissociation re- action agree within a factor of 10 with the experimental values.The rate constant of the dissociation reaction is 60 times greater than the experimental value. However, i t may be significant that very good agreement is obtained when Kiichler’s value for the activation energy (57,500 cal.) which, as mentioned previously, may be in error, is adjusted to agree with the calculated heat of reaction and the activation energy of the association reaction (adjusted activation energy 63,800 cal.). We note too that under the present assumptions, i.e., bending fre- quencies of 320 cm.-l for all vibrations of “ stable ” bonds, these last frequencies have to be used for the calculation of the vibrational par- tition function of the stable cyclohexene molecule. We recall that the equilibrium constant of the cyclohexene equilibrium calculated in this way also agreed within a factor of 9 with the experimental value. We can conclude, therefore, that these last assumptions regarding the fre- quencies of the transition complex, namely “ normal ” frequencies for all vibrations not affected by the reaction and “ low ” frequencies for all vibrations loosened due to the reaction, leads to quite reasonable agreement with the available data.D. ROWLEY AND H. STEINER 213 The present results confirm that the assumption of a similarity of transition state and final reaction product in diene association reaction is in the main correct. However, because of its greater looseness, the transition complex does not correspond in all its characteristics to the final rnolccule. Significant deviations occur, of which the lowering of the frequencies of all bending vibrations associated with bonds to be formed or to be broken seem to be the most important. TABLE XIV.-CALCULATED RATE CONSTANTS ASSUMING MODIFIED CYCLIC COMPLEXES Association reaction k, calc. (mole-1 ~ r n . ~ sec.-l) k , expt. (mole -1 ~ r n . ~ sec. -l) 1-66 x 103 Dissociation reaction I 8-70 x I O ~ kd calc. (sec. -l) 1-66 x I O - ~ I 1-02 x 10'1 k, expt. (sec. -l) But adiene di merization kCF& (mole -1 ~ m . ~ sec. -l) kexpt (mole -l ~ m . ~ sec. -l) 104.0 One of us (D. R.) wishes to thank the Central Research Committee of the University of London for a grant in support of the work recorded in this paper. We are also very grateful to Prof. M. G. Evans, F.R.S., and to Dr. A. Wassermann for several helpful discussions. Petrocarbon Limited, Twining Road, Traffoord Park, Manchester, I 7. 22-4

ISSN:0366-9033    

DOI:10.1039/DF9511000198

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

D. ROWLEY AND H. STEINER 213 11. HYDROCARBON REACTIONS GENERAL DISCUSSION A. THERMAL REACTIONS Dr. P. Torkington (Brit. Rayon Res. Assoc.) (communicated) Map I point out rather an interesting correlation between results quoted in papers by Sir Cyril Hinshelwood and Prof. Glockler ? The latter shows that the C-H and C-C bond energies in n-paraffin chains increase in passing from the centre towards the ends ; the former states that in the thermal chain fission of n-paraffins the probability of C-C fission (and GENERAL DISCUSSION * 214 hence also of C-H fission accompanying H-transfer) also increases to- wards the ends. I suggested earlier in this Discussion that Prof. Glockler’s results might be interpreted as evidence for the existence of molecular orbitals (overall but not necessarily delocalized in the Sense of implying electron mobility) favouring the chain ends.The correlation here implies that the presumably higher terminal electron density is associated with greater activity i.e. electron mobility an interesting negative analogy with true bimolecular reactions. Regarding the concentration of energy via vibrational coupling I should like to draw attention to another previous remark of mine following Prof. Ubbelohde’s paper. The envisaged concentration of energy could surely never occur with “ perfect ” coupling ; it seems that if the occur- rence of energy concentration by vibrational coupling is accepted as proved then so also must the occurrence of two or more collisions pro- viding the energy sources.Transfer from two or more such sources by coupling might well lead to a concentration of energy favourable to re- action by’ normal interference. It seems quite probable that a long-chain molecule might be involved in several simultaneous collisions but against fundamental principles for a concentration of energy to arise from a single collision. Finally is it possible for a little more non-mathematical commonsense to make a little more headway against a little more formalism ? How does symmetry lead to electro delocalization 9 Why because one can write functions 4 = ,Zai$i because of symmetry does an electronic system become mobile ? Perhaps too much n-bond theory has preceded a study’ of normal molecules. In an unsymmetrically-substituted benzene de- rivative there is no symmetry but still the same delocalization as in the parent compound.The point is obvious and not worth developing. at any link. Dr. G. Gee (Welwyn Garden City) said I should like to outline a possible explanation of the observed thermal decomposition products of paraffins. Considering first radical mechanisms to a first approxima- tion the thermal dissociation of a long paraffin chain will occur randomly CH,(CHz)nCH + CH3(CHa)z + CH,(CHa)L - - ( 1 ) The resulting radicals will then tend to degrade by the facile process of losing a series of ethylene molecules from the end CH3(CH,),* + CH3(CH2)g2 + CH*CH2 and this process will continue until a methyl or ethyl radical remains. Thus primary dissociation of long chains will result mainly in the produc- tion of ethylene and ethyl and methyl radicals and the ratio of these primary products will not be markedly temperature dependent.+ CH,(CH,),CH=CH,+CH,(CH,),-, (2) ( 3 ) The most probable secondary reaction of a methyl or ethyl radical appears to be the dehydrogenation of a paraffin molecule a process which again will occur nearly randomly along the chain (3) and will be followed by dissociation of the resulting radical (4) CH3* + CH,(CH,)nCH + CH + CH,(CHa),*CH(CHa),-,,CH CH3(CH2).*CH(CH2),-,,CH3 (4) The situation now becomes more complex in that the olefin produced will be more reactive toward further methyl and ethyl radicals than is the original paraffin. Without seeking to follow o u t these processes in detail i t can be seen that a further series of degradative processes will follow from the secondary radicals in which the principal final products will again be ethylene methane and ethane.Thus i t is easy to see why these products predominate and also why their relative yields are largely independent of conditions. It also follows from these considerations that in presence of nitric oxide the main hydrocarbon product to be expected from radical dissoci- ation is ethylene since nitric oxide would compete more successfully with GENERAL DISCUSSION 215 (3) than with ( 2 ) . The fact that methane and ethane are still produced in presence of nitric oxide therefore argues in favour of Stubbs and Hinshelwood’s contention that a molecular reaction occurs.The mechanism of this reaction does not appear very clear but it may be plausibly sug- gested that ethylene could be eliminated from the chain at any point through the formation of a cyclobutane type of activated complex CH,R CH,-R -+ I I + CHeCH,+R-R CH,-.R CH,-R . - ( 5 ) where R R are alkyl radicals or hydrogen. It is not a t first sight clear why this reaction should lead to methane and ethane rather than to larger paraffins but it appears likely that increase in the size of the radicals R and R would lead to steric difficulties. If this is indeed the case each successive step in the degradation of a long chain (by removal of ethylene) will lead to a more readily dissociable paraffin and-as in the radical decomposition-the principal final products will be ethylene methane and ethane.Dr. F. J. Stubbs and Prof. C. N. Hinshelwood (Oxford) (co.mununi- cated) In connection with Dr. Gee’s interesting suggestions we may state for comparison the views which we have formed about the reason for the similarity of products from a chain reaction and a molecular process. For the typical case of n-hexane the molecular process would be CH,CH,CH,CH,CH,CH -+ CH + CH CHCH,CH,CH CH,CH,CH,CH,CH,CH -z C,H + CH CHCH,CH, where El and El’ are the activation energies of the two reactions. The various steps in the chain reaction would be Initia- (1) (1’) in the molecule CH,CH,CH,CH,CH,CH,. El El’ CH,CH,CH,CH,CH,CH 3 X + Y tion { X+CH,CH,CH,CH,CH,CH + XH+’CH,CH,CH,CH,CH,CH DCH,CH2CH2CH,CH,CH -+ CH;+CH CHCH,CH,CH *CH,CH,CH,CH,CH,CH -+ C,Hi+CH CHCH,CH (2) CH;+CH,CH,CH,CH,CH,CH +- CH4+’CH,CH2CH,CH,CH,CH (2’) C,H; +CH,CH,CH,CH,CH,CH -f C,H,+’CH,CH,CH,CH,CH,CH where e, e, etc.are the respective activation energies. With ling chains the ielative proportions of methane and ethane formed will depend only on the repeating steps ( I ) and (1’) in the above scheme and their sequels ( 2 ) and (2’). The nature of the initiating re- action is unimportant for this purpose. For the products from the molec- ular process and chain reaction to be in the same proportions all that is necessary is that (&-El’) should equal (el-el’). This seems quite likely since the alternative steps in the molecular process are closely parallel with the steps (I) and (1’) respectively the two positions of rupture being in the same relative position in the radical CH,CH,CH,CH,CH,CH; as Dr.M. Szwarc (Manchester l7niversit-y) said Stubbs and Hinshelwood proposed in their paper an interesting mechanism for a unimolecular decomposition. I would like to examine some of the conclusions which it seems to me follow from their assumptions. Let us consider a system composed of two components viz. an inert gas X and a reacting gas M which may decompose unimolecularly. Furthermore let u s choose the experimental conditions in such a way that the energy supplied to molecules M arises from collisions between M and X. According to Hinshelwood and Stubhs we can observe two kinds of collisions. (if One type of collisions between M and X lead to the formation of energized molecules M* i.e.molecular species which contain an amount of energy at least sufficient for their decomposition GENERAL DISCUSSION 2 16 but distributed unfavourably amongst various internal degrees of freedom. This process we represent by the equation X + M + M* + X (collision of the 1st type). The decomposition of energized molecules requires a reshuffling of their energy which has to be accumulated eventually in the reacting centre i.e. M* +- MS + Products MS denoting an " activated " molecule i.e. a species which contains at least the necessary amount of energy localized in the reacting centre. Although collisions of the first type are frequent the probability for the energy redistribution is low and consequently the reaction caused by these collisions attains a first-order character at comparatively low pressure of the gas X.Plotting the rate of this decomposition as a function of partial pressure of gas X we obtain the curve I shown in Fig. I . (ii) Collisions of the second type lead directly to the formation of I ' activated " species MS X + M -t MS + X (collisions of the 2nd type). Since the chance of the occurrence of such a collision is very low and the probability for the subsequent decomposition of Mt is very high the reaction caused by collisions of the second type retains the second- order character up to comparatively high pressures of the gas X. Plotting the rate of the latter reactions as a function of pressure of the gas X we obtain the curve 2 shown in Fig.I. FIG. I. Fig. I illustrates the features of the decomposition claimed by Hinshelwood and Stubbs. There is a comparatively low pressure region (C-D) in which the first type of collision predominates and in which the reaction obeys first-order kinetics. There is a comparatively high pressure region (E-F) in which collisions of the second type are predominant and conseqfiently in this region the reaction obeys second- order kinetics. However in addition to these two regions there is a very low pressue region (A-B) in which both types of collision call for a decomposition of the second order and finally there is a very high pressure region (G-H) in which both types of collision demand first-order kinetics of decomposition.If we plot therefore the order of reaction as a function of the partial pressure of gas X we obtain a curve represented in Fig. 2. I would like to ask Prof. Hinshelwood if he agrees that this behaviour of decomposition is a consequence of his assumption. There is however another point which I consider to be important. If the reaction takes place in a very high pressure range say G-H then the first type of collision must lead to an equilibrium between the energized molecules M* and the ordinary molecules M i.e. X + M + M* + X (collisions of the 1st type). GENERAL DISCUSSION 217 I n the same pressure region collisions of the second type must lead to an equilibrium between the " activated " molecules MS and the ordinary molecules M i.e.X + M $ MS + X (collisions of the 2nd type). One has to realize that any type of " activated " molecules MZ can be produced either by collision of the 1st type via M* or directly by collision of the 2nd type. If the latter modes of activation leads to the state of equilibrium between ordinary molecules M and activated species MZ then i t would be expected that the activation through collisions of the first type should also lead to an equilibrium between these two species (this follows from the fact that both types of collisions occur in the same gas phase). Consequently the following scheme should be valid for a reaction resulting from collisions of the first type M + X + M* + X MZ -f Products Px I Jr and the equilibrium concentration of MZ should be the same as in the case of collisions of the 2nd type.That however would contradict the as- sumptions on which Fig. z was obtained. FIG. 2. One can argue that the collisions of the second type do not lead to an equilibrium concentration of Mi since their rate of decomposition is too high being of the order of 1013 (MS) sec.-l. This seems to be quite correct. Indeed collisions of the second type can never be of any significance and if they have to maintain a first-order character of decomposition a pressure of the order of 105 atm. would be required. It seems that the best method of approaching the problem is t o sum up aZZ the contributions arising from all possible collisions (appropriately weighed). The reaction would be of the second order a t low pressures and would attain the first-order character at sufficiently high pressures In the latter case the activation energy of decomposition would be given by the activation energy of the process M +M$.Dr. F. J. Stubbs and Prof. C. N. Hinshelwood (Oxford) (communic- ated) Given the hypothesis that the nitric-oxide inhibited reaction is a definite molecular reaction (which as we have explained there is a good prima facie case for accepting) the changes of order referred to by Dr. Szwarc are the basic experimental facts upon which our views about the kinetically composite nature of the reaction are founded. The tentative interpretation which we have advanced is not in any way built upon a preconceived theory and may be modified in the light of further experi- ments now in progress.In the meantime we would prefer not to regard the facts about the order as deductions from a theory as Dr. Szwarc 218 GENERAL DISCUSSION implies but as observations calling for explanation. What we have suggested is that the normal theory requires some degree of amplification. Prof. S. W. Benson (California) (conznzunicated) The data presented by Ingold Stubbs and Hinshelwood and quoted here seem to indicate that even in the " fully inhibited " pyrolysis there is a surface effect which increases with decreasing temperature and is virtually independent of pressure. It seems to me that until this effect is further elucidated i t is difficult to accept the discussion and rate constants quoted as anything but tentative.There still remains the possibility that there is a heterogeneous reaction initiated on the walls which starts chains which also end at the wall. Rice and Herzfeld 1 have recently shown that under such conditions i t is possible that depending on the mechanism the rate may show any de- pendence on the surface volume ratio between complete independence and direct proportionality. Even in a direct decomposition to molecular products i t is conceivable that heterogeneous decomposition competes with homogeneous decom- position. In such case the activation energy of the two would be ex- pected to be different and since the relative amount of each would be expected to show a monotonic if not simple pressure dependence there would be observed an effect of pressure on experimental activation energy.For n-pentane i t is difficult to believe that when the experimental activation energy is found to vary between 80 and g o kcal./mole in the pressure range 50-150 mm. Hg (inhibited reaction) t h a t the free radical mechanism of splitting of the C-C bond which has an activation energy of 80-85 kcal. is not proceeding at a competitive rate. Finally it is difficult to reconcile the observed decrease of activation energy with increasing pressure with any reasonable behaviour of mole- cules. In the Rice-Ramsperger-Kassel theories of the unimolecular decomposition the activation energy is supposed to decrease at the lower pressures where the higher energy states can no longer make their com- plete contribution to the overall rate.The explanation offered by the authors to the effect that there may be a single mode which requires a very long time to reach activation even when the energy is present in other modes seems completely ad hoc and unsupported by any other experimental or theoretical evidence. It also leans upon an interpreta- tion of a molecular collision that seems in immediate conflict with the law of microscopic reversibility. That is there seems to be an implicit assump- tion that an active molecule can only lose energy on collision and never gain it. If there is a mode which is only slowly activated then the chances are that i t is just as slowly deactivated and collisions do not necessarily play any decisive role in the time of decomposition.The example of ethane where the kinetics seem to go through a second to first order transition while the activation energy changes but little (or possibly is higher a t the lower pressures) seems a striking example of the above-mentioned contradiction. I doubt very much that these pyrolyses will be very tractable theoreticallv until the nature of the surface efiects and the inhibition axe better understood. Dr. F . J. Stubbs and Prof. C. N. Hinshelwood (0,vford) (conz~lzunic- nted) Chain reactions which are initiated by a decomposition are not at all likely to start and stop at a wall. If the molecule dissociates on the wall radicals are likely t o remain adsorbed. In any case chain reactions staxting at a wall should be subject to the usual kinds of inhibition just like any other chain reactions.This assumption of Prof. Benson seems to us to be improbable. We are unable to understand his fourth paragraph. In pentane we have always believed the chain reaction normally to proceed a t a " competitive rate " but we state a prima facie case for concluding that 1 Symposium on Abnormalities in Reaction Kinetics Amer. Chem. SOC. Minn. hfinn. Spring 1950. GENERAL DISCUSSION I 80 CH3-CH3 1 219 this competing reaction has in the conditions of the experiments been suppressed. In his penultimate paragraph we do not recognize our views very clearly and as to the last paragraph we can assure him that our approach continres to be essentially experimental and that all the details of these admittedly rather puzzling phenomena are under examination still and from several new points of view.Dr. L. Bateman (Welwyn Garden City) (communicated) The form and relative position of the paraffin and olefin decomposition rate curves in Fig. I of Stubbs and Hinshelwood’s paper are just as would have been predicted on the basis of A H changes on breaking the most easily broken bonds so that some reference to this mode of primary bond scission is perhaps of interest. The AH changes (in round figures) for the C,-C compounds are summarized in the following Table. Ethylene should be I00 80 < 60 much more stable than ethane propane and propylene should be of similar stability and the higher members of the two series should exhibit a difference reflecting the partial utilization of allylic resonance energy in the bond-breaking process.These features are evident experimentally but ethane and I-butene appear to be relatively more stable than might be expected and may serve to illustrate how secondary processes affect the overall decomposition. The common factor is the primary formation of methyl radicals and not higher alkyl radicals capable of ready decom- position into smaller fragments. Hence in ethane recombination will be specially favoured since the only alternative reaction is the rather difficult hydrogen exchange In Me I-butene + CH the CH corresponding . CH,Me -f MeH reaction + CH to the CH latter .CH . Me ( E - 5 kcal.) is much more facile but the resulting substituted allyl radical like that initially produced will be far less active than an alkyl radical in inducing further decomposition. An important aspect of these considerations is that all A1-olefins are required to yield simple allyl radicals as primary fission products. The reactivity of these radicals under the decomposition conditions their ultimate products and possibly the estimation of their extent of forma- tion might well be stcdied experimentally by carrying out co-decomposi- tions with I 5-hesadiene as a foreign source of allyl radicals. Informa- tion on the third point would establish the contribution of allylic bond primary scission to the evidently composite overall decompositi m.Dr. R. G. Partington (Oxford) (communicated) In their paper Stubbs and Hinshelwood state that Steacie and Folkins suppnrt the view that the reaction taking place In the presence of sufficient nitric oxide to reduce the rate to a minimum (the “ residual reaction ”) is “ the primary process of what in the absence of inhibitors would be a chain reaction ”. This is I think a mistaken interpretation of Steacie and Folkins’ conclusions I 1 slightly ( 8 0 CH . CH . CH,-H CH3. CH,-CH3 CH CH-H 1 due to second- CH CH . CH,-CH3 60 CH3. CH,-CH,. CH3 CH . CH,-CH . CH . CH J order effects CH CH . CH,-CH . CH3 ’lightly Me + C,H + MeH + C,H (E N 10-15 kcal.). Steacie and Folkins Can. J . Res. B 1940 18 I. CH,-CH,-CH,-CH,-CH3 CH3+N0 = CH3N0 CH3-CH2-CH,-CH,+ 2 20 which read '' the addition of nitric oxide then merely diminishes the chain length without completely suppressing the chains ".Had the suggestion been made that all the initidly formed radicals reacted with nitric oxide one would have to consider the possibility of the formation of such a molecule as (I) in the following scheme e.g. (1) which might well participate in further reaction (starting by its decom- position) in a similar way to the large radical itself. A further paper by Steacie and Folkins 8 is of interest in this connection. It is concerned with the decomposition of n-butane initiated by free C A FIG. 3 . (19391 P. 700. There are many possible sources and convenient methods of producing the appropriate radicals either thermally or photochemically and in connection with the possibility of making such a study some experiments of mine on the chemical analysis of small concentrations of nitric oxide in gas samples of the size usually handled in work of this character might be of value.A gas sample (say 50ml. of gas at N.T.P.) containing say Q to I yo of nitric oxide is enclosed in the holder A (Fig. 3 ) . This is then con- nected to a flask B containing air and small amount of an aqueous solution of sodium or potassium hydroxide free of nitrite. On opening the wide- bore stop-cock C and gently shaking the reactions take place. Reaction ( 3 ) is slow and the conversion to nitrite is almost complete. The solution of nitrite can then be analyzed with great accuracy by the standard colorimetric method^.^ The conversion factor NO -f nitrite is reproducible for a given nitric oxide concentration and its value could be determined over the desired range of concentrations by making up mixtures.Can. J . Res. B 1939 17 gg. 4 Trotman-Dickenson and Steacie J . Amer. Ckem. SOC. 1950 72 2310 ; J . Chem. Physics 1950 18 1097 and following papers. 6 See for example Vogel A Text Book of Quantitative Inorganic Analysis (3) (4 2NO + 0 = 2N0 . (NO + NO,) + 2NaOH = 2NaN0 + H,O . - (4) GENERAL DISCUSSION = CH,+CH,-CH2-CH2-CH3 NO = CHs-CH~-CH~-CH2-NO radicals from ethylene oxide and the authors conclude that their results " cast some suspicion upon the idea that maximum inhibition by nitric oxide in all cases corresponds to complete suppression of chains *'.The whole problem of the action of nitric oxide as ail inhibitor (and the foregoing remarks were largely included to justify the opinion that the problem still exists) seems to depend on the relative rates of the two reactions viz. reaction of radicals with the hydrocarbons e.g. * R + RH = R,H + R . and reaction of radicals with the inhibitor e.g. * (1) R + NO = R,NO. . Reactions of the type (I) have recently been studied and similar specific attention to reactions of the type ( z ) would undoubtedly be of the greatest value in helping to solve the above problem and might be less subject to ambiguity than experiments on a system in which reactions of this type are among the very many taking place.* 22 I GENERAL DISCUSSION . - (20) * CH + (CH,),CO -+ CH + residue. Since the agreement of values of El based on the two separate assumptions was good one conclusion was that the two sets of results were consistent. The more recent value of Dodd (E20 = 10.7 f 0.5 kcal.) is simply not reconcilable with the older value (8.6 kcal.) of Grahame and Rollefson. It is consistent with the results of Trotman-Dickenson and Steacie only if it is assumed that our measurements of E differences are inaccurate to about 2 kcal. We are cognizant of the difficulties of this work. Never- theless the new evidence presented by us does support a high value of El and Dodd’s results emphasize that fact. Dr. A. F. Trotman-Dickenson (Manchester) said I am not clear why Anderson Davison and Burton have disregarded the recent work of Dodd in selecting 8.6 kcal.as the best available value for the activation energy of the reaction CH + CH,CHO -f CH + CH,CO. Dodd has shown that the results of all previous workers (including those of Rollefson and Grahame) may be interpeted as giving a value for the activation energy of 10.7 f 0.5 kcal. This value makes the agree- ment between the acetaldehyde and acetone results much worse (15.8 and 13.2 kcal. respectively). Dr. Steacie and I recently reviewed the data on the reaction CH + H +CH4 + H El = 8.8 kcal. steric factor = 4 f 2 x I O - ~ (oR2 = 2-8 A ucas = 3.5 A) and concluded that the best work (three separate determinations) gave and K / K t at 182O C = 50 x 10-1 molecules-* cm.812 sec.-4.The results of Anderson Davison and Burton give K,/kt2 at 182O C = I x 10-1 molecules-* ~ m . ~ / e set.-* if the same collision diameters were used as above. No satisfactory reasons for these serious discrepancies have been given. Prof. M. Burton (Notre Dame) (communicated) Mr. Trotman- Dickenson mLst not have realized that the full paper by Dodd did not reach our laboratory until after the paper by Anderson Davison and myself was submitted for publication. Otherwise we would certainly have commented on it. Our actual work was concerred with establishment of activation energy differences. Calculation of E for - (1) CH,+H + CH4+H . involved also assumptions of reliability of E values of Grahame and Rollefson for CH + CH,CHO -+ CH + CH,CO and of Trotman-Dickenson and Steacie for CH + H = CH + H .(18) Regarding the relative merits of the results of Dodd and of Grahame and Rollefson we have no firm opinion. However we may note that E, >El* is a new idea. The reverse order of activation energies ac- cepted prior to the work of Dodd is consistent with the usual notion that since in reaction (20) a formyl H is involved while in reaction (18) it is a methyl H reaction (20) probably has a lower E,. Mr. Trotman-Dickenson’s remark concerning discrepancies between values of kl/kl24 from old experimental data and from calculation involving our El and s emphasizes the dangers inherent in acceptance of old cal- culations based on experiments which may contain innate error e.g.the general difficulty of interpretation of experiments of involved mechanism. Prof. W. A. Noyes Jr. (Rochester) said It is interesting to note in the paper by Prof. Burton and co-workers that for the reaction (1) Dodd J . Chem. Physics 1950 18 234 ; Trans. Faraday SOC. 1951 47 56. 7 Trotman-Dickenson and Steacie J . Physic. Chem. (in press). GENERAL DISCUSSION * 222 the activation energy is 13-14 kcal. and the steric factor about I O - ~ . These results are based on studies in which reaction (I) could compete with reactions of the type CH + RH = CH + R. . The activation energies for (2) depend on R but generally fall in the region of 8-10 kcal. (e.g. if R is CH,COCH,- at about 9-7 kcal.).Thus if steric factors for (I) and (2) were identical (2) would proceed about 102 times as f a s t as (I) at equal pressures of H and of RH at temperatures of 200-300° C. A ratio of rates as high as 102 would invalidate most conclusions due to attendant experimental difficulties. The very fact of successful com- petition under experimental conditions used by most authors indicates quite strongly that steric factors for (2) are usually 10-1000 times lower than for (I). (3) (2) This fact and other related facts raises questions about “ hot ” radical conclusions. If a radical separated from a parent molecule by absorption of energy much greater than that necessary to break the bond retains energy solely as kinetic energy i t could hardly undergo 103-105 collisions without being reduced essentially to thermal equilibrium with its sur- roundings.One must conclude either that hot radical effects would be unobservable or that the so-called steric factor is a very pronounced function of kinetic energy. Theory concerning this matter is very vague and qualitative although one is led to predict an increase in steric factor with increase in kinetic energy. An examination of published data on hot radical effects for reactions of the type of (2) reveals that positive conclusions are not warranted. Uncertainties exist either due to scatter in the data themselves or to effects other than hot radicals which afford equally satisfactory explanations. This is true particularly in those cases such as CH,COCH3 and Hg(CH,) where the full nature and yield of the primary photochemical process have not been elucidated as a function of temperature and other variables.A word should be said also about the reaction CH + CH = CpH6. . All workers seem agreed that activation energies for radical combination reactions of the type of (3) are very low. Less agreement is found con- cerning the steric factor but the safest conclusion seems to be that i t is high and probably near unity. The necessity for a third body has not been clearly shown but an examination of the data a t sufficiently low pressures will indicate a trend which might show a third-body effect More data on this point will be found in a forthcoming article by Dr. A. J. C. Nicholson as well as in a recent article by Gomer and myself.* A third body may be introduced in several ways but i t is not satisfactory to introduce a mere triple collision.One way would be (4) 1 4 * CH + X c Z ks CH,X CH,X + CH = C,H6 + X . in which case the rate of ethane formation would be ( 5 ) Under conditions of high radical concentrations (high intensities) eqn. (6) reduces to an equation which is satisfactory for part of the ethane formation during photolysis of mercury dimethyl but which is not satisfactory under con- ditions so far studied in acetone. The trend in acetone at low pressures may indicate the necessity for a three-body collision. J . Amer. Chem. SOL 1950 71 101. GENERAL DISCUSSION * (8) 223 The most satisfactory way to obtain the right rate equation for the majority of experimental conditions is that used by Gomer and others viz.the reversible formation of an intermediate complex from two methyls the intermediate complex being stabilized to ethane by collisions. At sufficiently high pressures the rate-controlling step is the formation of the intermediate complex a step whose rate depends only on (CH3)2. One comes to the conclusion therefore that the nature of the third body may determine the way in which i t acts in agreement with the classical experiments of Rabinowitch and Wood on the recombination of bromine atoms. Under most conditions the rate of ethane formation is given by RQG = k(CH,)' in which h may be a complex constant which indicates nevertheless the rate of (3) to be very high.Dr. P. Torkington (Brit. Rayon Res. Assoc.) (corutmuvzicated) Is the steric factor temperature-dependent ? If so its variation might possibly be related to the proportion of molecules (in this case radicals) undergoing in the reaction CH3- + H, it might be thought that approach along the a mode of vibration more favourable to reaction than other modes. Thus line C . . . H-H coinciding with opening of the CH " umbrella " at the correct stage (the hydrogen molecule lying on the opposite side to the hydrogen atoms of the methyl group) would have a high probability of successful reaction. I n this case then the favourable vibration is the symmetrical breathing frequency of the methyl group. The argument is not essentially altered if the group is planar as has been fairly recently suggested though the favouring would not be so noticeable.Possibly the principle could be applied generally. As regards feasibility with a frequency of 1500 cm.-l and amplitude of the order 10-g cm. the hydrogen atoms in the breathing mode of a methyl group have a mean velocity of the order 105 cm./sec. ; the root-mean-square velocity of hydrogen molecules at 2 5 O C is about 2 x 105 cm./sec. The example shows that vibrations might co-operate. Dr. George Porter (Cambridge) (communicated) The high value for the bimolecular rate constant of methyl radical combination at room temperature of about 5 x 1013 cm.a mole-I sec.-l quoted by Sztvarc and others in the discussion now seems well established having been obtained by several workers using the sector method There appears a discrepancy however with the results of mirror experiments which cannot be dis- regarded on the grounds that radical removal under these conditions is known to be a wall reaction.Whatever the mechanism by which the radicals disappear it is the rate-determining one and thus sets an upper limit for the homogeneous bimolecular rate constant. If the concentration of radicals is low as in the photochemical experi- ments of Norrish and Porter O the rate constant is found to be high and no discrepancy appears but as the concentration is increased so is Kmax in- creased for the radical lifetime does not decrease proportionately. Thus Forsyth's values 10 give kmax = 6 x 1012 cm.3 g .mole-1 sec.-l and although other workers have not recorded concentration measurements specifically i t seems certain from a consideration of mirror removal times that even higher concentrations were involved in the experiments of Rice Johnston and Evering l1 and of Paneth Hofeditz and Wunsch l2 who also increased the life-time to 0.1 sec. by decreasing the rate of the wall reaction. As the rate-determining reaction is heterogeneous i t appears that the velocity of the bimolecular gas phase reaction determined in this way is consider- ably less than that corresponding to unit collision efficiency. Norrish and Porter Faraday SOC. Discussions 1947 2 97. Forsyth Trans. Faruday Soc. 1941 37 312. l1 Rice Johnston and Evering J. Amer. Chem. Soc.1932 54 3529. la Paneth Hofeditz and Wunsch J. Chem. Soc. 1935 372. GENERAL DISCUSSION The value of kmax is obtained from t* and the concentration the latter being determined from the weight of mirror removed or metal alkyl formed in unit time which can be measured to within 10 yo without great difficulty. The half-life is also a reasonably accurate determination and the well-known difficulties of the method such as mirror poisoning and irreproducible wall conditions can produce only a scatter and not a general trend of this kind. It appears therefore that this discrepancy must be considered seriously unless a reason can be given for doubting the basic assumptions of the mechanism of mirror removal. The bimolecular rate constant must eventually decrease at low pressures when the collision life becomes greater than the lifetime of the collision complex.A possible explanation therefore lies in the lower pressure of the mirror experiments which is about I mm. whereas the pressure in the sector experiments was usually much higher and always above 5 mm. Dr. E. J. Bowen (Oxford) said The importance of the amount of delocalization energy in the products of a dissociation affecting the ease with which it occurs is also shown by work carried out by Miss Rohatgi on the photochemical reaction of anthracene with liquid chlorinated hydrocarbons. In light absorbed by anthracene the reaction is AX + RC1 -+ ARC1 (derivative of dihydroanthracene). . 0.018 The following quantum efficiencies were found Carbon tetrachloride CC Pentachlorethane CHCl,CCl .Tetrachlorethylene CCl,=CCl Tetrachlorethane CHCl ,CHCl . Trichlorethylene CHC1-CC1 a . Ethylene dichloride CHCl=CHCl . Chloroform CHCl . 0'4 . 0.27 . 0.07 . 0.008 . 0.0076 . 0.0032 High quantum efficiencies appear only where the CCl radical might be involved. Prof. S. W. Benson (California) said It seems rather strange that the Br atoms liberated in the initial step do not undergo addition and hydrogen abstraction reactions with the cyclohexene. I wonder if Dr. Robb has any information on this point. Dr. E. C. Kooyman (Amsterdam) (communicated) It seems likely that the main differences between the mechanisms of the addition of bromotrichloromethane and that of carbon tetrachloride to olefins arise from the far greater reactivity of the bromine atom in CBrCl as compared with that of the chlorine atoms in CCl,.Kharasch and Friedlander la found bromotrichloromethane to react rapidly with styrene at 50°C under the influence of ultra-violet light. However little reaction occurred at 20' C ; at this temperature styrene was found to inhibit the addition of CBrCl to other olefins. These facts were interpreted by Kharasch and Friedlander as resulting from the low reactivity of the benzyl type radical formed by preferential addition of CCl to the styrene double bond -CCl + CH,=CH-Ph -+ CC1 ,CH,CH-Ph. These data suggest the abstraction of a bromine atom from CBrCl to be a slow reaction at low temperatures when the attacking radical is re1 atively stable.In detailed analyses of the reaction between cyclohexene CCl and benzoyl peroxide at 78" C Kooyman and Farenhorst l4 found 60-70 ?& of 3 3-dicyclohexenyl on the basis of peroxide decomposed. In our in- vestigation on a-methylenic reactivity l5 the " retardation constant '' for cyclohexene was found to be 11 x I O - ~ ; in view of the value 0 - 2 reported l 3 Kharasch and Friedlander J . Org. Chem. 1949 14 239. l4 Kooyman and Farenhorst Rec. trav. clzim. 1951 (in press). 15 Kooyman this Discussion. GENERAL DISCUSSION 225 by Kharas3i and Friedlander and by Kharasch and Sage l6 for the ratio of the addition rates of CC1 radicals to the double bond in cyclohexene and n-octene respectively this suggests the abstraction of a-hydrogen atoms in cyclohexene by trichloromethyl radicals to proceed at a rate which is not much smaller than the rate of addition.Finally attention is called to the experiments of Kharasch and Fried- lander with respect to the overall kinetics of the photochemical addition of bromotrichloromethane to various olefins including cyclohexene at 10-50OC. The authors found the reaction to be zero order with respect to olefin. Olefin determinations were made by direct titration with bromide-bromate solution rather than by dilatometry. On the bases of these arguments it is suggested that the rate constants reported by Melville and his co-workers may contain contributions from cyclohexenyl radicals rather than to apply to the trichloromethyl radicals only.Thus termination might consist both of dimerization of trichloro- methyl radicals and of dimerization of cyclohexenyl radicals. The in- fluence of the latter will be of course more pronounced a t lower halide/ olefin ratios. I f 2-5 x IO* 16 Kharasch and Sage J . Org. Chern. 1949 14 537. Prof. H. W. Melville Dr. J. C. Robb and Mr. Tutton (Birminghum) (communicated) The kinetic evidence presented in our paper entirely pre- cludes atatck by trichloro-methyl radicals on the a-methylenic hydrogen atoms to form chloroform and a stable cyclohexenyl radical as suggested by Dr. Kooyman since under all our conditions the rate of reaction is very accurately proportional to (rate of initiation) 4 and also to the concentration of olefin. If such a reaction did occur exclusively as a termination CC1 + CgHlo +CHCl + CBHg .. . . k' the rate of reaction would be given by which is of course not the case. If on the other hand both processes i.e. reactions given by k' and R were operative then the intensity ex- ponent would lie between 0.5 and I. The fact that i t is so accurately 0.5 supports the original kinetic scheme suggested in our paper. Since the discussion an experiment has been done in which cyclo- hexene in presence of benzoyl peroxide as catalyst with excess of carbon tetrachloride was placed in a sealed dilatometer tube. By irradiating with light of suitable wavelength the photo-catalyzed reaction was studied and in agreement with the work reported by Dr. Kooyman this was shown to be dependent directly on the light intensity thus establishing that under his conditions termination is indeed first order with respect to radical concentration.The rate of reaction under these conditions is very much less than when CC1,Br is used. This then raises the point of the reason for the different mechanism under the two different experimental conditions but further experiments are necessary in order to establish the nature of the differences in the mechanism involved and more precise information is required regarding the velocity constants for all possible steps in the reaction. 2k6 Prof. H. W. Melville Dr. J. C. Robb and Mr. Tutton (Birmingham) said Since the paper on the reaction of trichloromethyl radicals with cyclohexene was submitted further values have been obtained for the proposed kinetic steis.These are listed below. 1.3 x 1 0 2 1. mole-1 sec.-l J I R (30" C) k (40" C) 1-5 x I O ~ k (5oOC) 1-90 x I O ~ Eovers1l = E - &Es = 4-5 kcal./mole. 2.0 x 10' 1. mole-1 sec.-l. 2kS H GENERAL DISCUSSION 226 This gives a value of 4 = 5.5 where 2k6 = +V2k4 . 2k5. This low value is in accord with those recorded elsewhere in the dis- cussion by Bateman Gee Morris and Watson. Further attempts to obtain a value for the energy of activation of the termination step by means of experiments conducted using the technique of intermittent illumination at various temperatures has resulted in failure to detect any activation energy for this step. It should be pointed out however that this technique is not sufficiently sensitive or accurate to measure small activation energies less than 3 or 4 kcal.over the rather small temperature range normally available in experiments of this kind. Dr. M. Magat (Paris) said I would like to mention some additional evidence concerning two points raised by Dr. Kooyman in his paper. First concerning so-called " stable " i.e. non-dimerizable radicals some caution is advisable. Dr. Chapiro in our laboratory has for instance observed that the reputedly stable radical I I-diphenyl-2-picrylhydrazyl reacts with double bonds of polymerizable vinyl compounds forming molecules of a molecular weight higher than that of the radical dimer. We are now investigating the kinetics of this process and the nature of the molecules formed.Secondly concerning the reaction of radicals with poly cyclic com- pounds we have investigated the effect of addition of cancerogenic hydro- carbons and their homologues on the rate of thermal polymerization of styrene at 37" C. If the usually present traces of peroxides are destroyed by a preheating under vacuum at this temperature for 2-3 weeks all these cancerogenic compounds slow down the polymerization. It can be shown by persistent fluorescence of the polymer that these compounds do enter the chain. It is remarkable that the absorption and fluorescent spectra are not significantly shifted. The slowing-down efficiency decreases in the order 3 4-benzpyrene eo-methylcholantrene phenantrene a-methylanthracene I 2 5 6-di- benzanthracene.Chrysene and pyrene have no effect at a1l.l' Dr. K. S . Pitzer (Washington D.C.) said Prof. Kistiakowsky has reported in his paper some exceptionally fine experimental work and the minor point of interpretation which I am about to raise does not in any way detract from the principal results of the work. In the text just after Table IV i t is stated that the presently unknown corrections for anharmonicity in the vibrations would raise the calculated heat capacity of ethane more than that of ethylene. I would urge caution in accepting this estimate as certain although i t may well be correct. While i t is true that ethane has more vibrations in a given frequency range than ethylene i t could be that the effect of anharmonicity in the torsional vibrations overshadows all others at moderate temperatures.The present treatment as a restricted internal rotation accounts for the anharmonicity in the torsional motion in ethane while the more highly restricted torsion of ethylene is presently treated as a harmonic oscillator. Thus i t seems possible that the anharmonicity correction for the torsional motion in ethylene might be large enough to make the total for ethylene exceed that for ethane. The same statements can also be applied to propane and propylene. Prof. E. A. Guggenheim (Reading) (comr~zunicated) It j s clear that the measurements of Kistiakowsky and Nickle on the ethane-ethylene equilibrium are appreciably more accurate and more reliable than any previous measurements.It is particularly satisfactory to notice how well these results agree with the most recent theoretical values. It is perhaps not entirely without interest to compare the experimental values with theoretical values calculated before the experiments. The following l7 Bodme and Magat Comfit. rend. 1951,232 1657. GENERAL DISCUSSION 227 Table gives a comparison between experimental measurements made in 1942 and here reported and the values calculated according to a formula la published in 1941 based on Kistiakowskfs own experimental value for the heat of hydrogenation a t 82’ C. 450’ c . EQUILIBRIUM VALUES OF @CIQB[2/p,-,Hb IN ATM. Measured Temperature 1942 . . 5-13 f 0.13 x 10-4 380’ C 4-04 f 0.17 x 10-6 4’3 x 10-6 Calculated 1941 .. 5-2 x 10-4 Dr. L. Bateman (Welwyn Garden City) (commu.nicated) Horrex and Miles’s reference to the bond shortening in dibenzyl requires revision. Cruickshankl9 has re-analyzed Jeffrey’s data and finds that they are actually indicative of only half the contraction quoted. Moreover recent isomerization equilibrium measurements by Dr. J. I. Cunneen and myself 20 fail completely to reveal any chemical effect (AG differences < 0.1 kcal.) which could be attributed to bond shortening in I 5-dienes. I suggest therefore that the agreement claimed between the experi- mental and ‘ I calculated ” bond dissociation energies needs critical reconsideration. Three uncertainties are apparent ((i) the correctness of the observed activation energy absolutely and as regards its identity with the energy rather than say CH .CH,-CH . CH as the “ non-resonance ” refer- of primary bond scission ; (ii) the appropriateness of using CH,-CH ence compound for P h . CH,-CH,. P h ; (iii) the magnitude of the resonance energy of the benzyl radical. Horrex and Miles present con- vincing evidence for the approximate validity of their conclusions con- cerning (i) and this receives support from Bolland and Orr’s 21 investigation into the analogous breakdown of aliphatic I 5-dienes. I question however whether f I kcal. is a fair estimate of the overall uncertainv in the activation energy in view of the complexity of the decomposition process the analytical intricacies and the spread of the points in Fig. 7.Any error associated with (ii) is undoubtedly small but will be such as to tend to reduce the “ standard ” CC-bond energy (by about 1-2 kcal.). The third factor is thus left as the main source of uncertainty. Szwarc has derived the value of 24.5 kcal. by subtracting the activation energy for the pyrolytic decomposition of toluene (this being identified with the energy required to break an acyclic CH-bond) from the CH-bond dissoci- ation energy in methane. In fact there is sufficient uncertainty in this derivation for the formerly accepted value of 19 kcal. still to be tenable. Szw-arc reports the activation energy as 77-5 f 1.3 kcal. but statistical regression analysis of all his tabulated data (for uniform surface condi- tions) leads to the result E = 80.0 f 4-2 kcal.(95 yo limit). Further apart from the small uncertainty in D, in methane (101 f I kcal.) i t would seem more legitimate to compare with an alkane €We where D, in the Me group is undoubtedly lower,2s rather than with the sym- metrical methane molecule. dissociation energy of dibenzyl simply as (84 - 2 x 19) = 46 kca1.-in For the present therefore i t would seem preferable to calculate the satisfactory agreement with the thermochemical and kinetic data- rather than as (85 + 11 - z x 24.5) = 47 kcal. in which a large other- wise unrecognized bond energy term has to be invoked in order to com- pensate for a higher resonance energy which is not definitely specified experimentally. To be published shortly. l9 Cruickshank Acta Cvyst.1949 2 65. l 8 Trans. Favaday Soc. 1940 37 272 a1 Bolland and On I.R.I. Tram. 1943 21 133. 22 Szwarc J . Chem. Physics 1948 16 138. Stevenson J . Chem. Physics 1942 10 291 ; Anderson and Van Artsdalen J Chem. Physics 1944 IS 479. GENERAL DISCUSSION Dr. M. Szwarc (Munchester) (communicated) Dr Bateman's remarks enables me to clear up some details of calculation of activation energy in the pyrolysis of toluene.24 The technique elaborated in this investiga- 0.01 yo. This seems to be the lower limit permissible in this technique tion makes i t possible to measure the rate of pyrolysis down to about and indeed runs 89 88 87 and 94 (crosses in Fig. 4) show that losses of products start to be noticeable when the limiting case of 0.01 yo is reached. Therefore these four results were ignored in calculating activation energy and were omitted in the graph which was presented in the paper but they were included in the Table giving all the results.The value of 80 kcal./mole obtained by Dr. Bateman results from including these four runs; their omission leads to an activation energy of 77.5 kcal./mole as quoted in the original paper. I take this oppor- tunity to include in Fig. 4 the results of pyrolysis of toluene obtained in 1948 by Dr. J. S . Roberts and in 1949 by Mr. J. Murawski. Pyrolysis of toluene. 0 szxTm2 (1947) A Roberts (1948) E = 77'5 kcal./mole. x runs No. 89 88 87 94. Murawski (1949) 228 C-C FIG. 4. In his contribution Dr. Bateman deals with the possible connection existing between the length and the dissociation energy of the central bond in the molecule of dibenzyl.This problem has been discussed previously,26 and here I would like to clarify further certain points which need additional emphasis. (i) The dissociation energy of a bond linking two atoms depends not only on the nature of the two atoms but i t is also greatly influenced by the molecular environment of the bond in question (e.g. Table I in ref. (2)). (ii) The factors which influence the magnitudes of bond dissociation energies can be divided into two groups those connected with the struc- ture of the undissociated molecule and those due to the nature of radicals produced on dissociation. On the whole i t is not possible to ascertain which factor and to what extent is responsible for the observed changes in bond dissociation energies (see however ref.(3) in which an exceptional case is discussed). (iii) Further treatment of the problem of variations in bond dissoci- ation energies requires therefore the introduction of certain simplifying t4Szwarc J. Chem. Physics 1948 16 128. 25 Szwarc Faraday SOC. Discussions 1947 2 39. GENERAL DISCUSSION 229 assumptions. It has been assumed tentatively 26 that the variations in the C-H bond dissociation energies in molecules of the type R . H are due entirely to the factors arising from the nature of the radical R (i.e. i t has been assumed that these variations reflect the changes in the stabilities of various radicals R). Taking the value of D(CH,-H) as the point of reference it is possible zo to build up a system of experimental resonance energies " for various radicals R defined as D(CH,-H) -D(R-H).(iv) It follows from the examination of heats of formation of the relevant compounds that variations in other bond dissociation energies cannot be accounted for by the experimental resonance energies only. For example D(CH,-CH,) - D(R-R,) -+ [D(CH,-H) - D(R-H)] + [D(CH,-H) - D(R,-H)]. I would like to emphasize that this inequality arises from thermochemical data only i.e. the value of [D(CH,-H) - D(R-H)] + [D(CH,-H) - D(R,-H)] - [D(CH,-CH,) - D(R-R,)] is independent of the values of dissociation energies used in this expression. (v) I t has been assumed that the above-mentioned value the " strengthening effect ",26 is related to the length of the relevant R-R bond.For example its value for dibenzyl has been calculated as about 11 kcal./mole and i t has been suggested 2 5 ~ 26 that this value is related to the bond shortening in the molecule of dibenzyl. shortening of the central C-C bond in the molecule of dibenzyl. Since as has been said above the value of II kcal./mole is independent of the value of the C-H bond dissociation energy in toluene the uncertainty (iij) quoted in the communication by Dr. Bateman is irrelevant €or the problem of the C-C (vi) The relation between the C-H bond dissociation energy in toluene and the C-C bond dissociation energy in dibenzyl is given by the heats of formation of toluene dibenzyl and the H atom i.e.ZD(CeH5. CH2-H) - D(C6H5. CH,-CH,. C6H5) = ZAH,(H) -ZAHf(C6H5 . CH,) +AHI(C6HQ . CH . CH . C6H5) = 106 kcal. /mole. Hence the value of D(C6H5. CH,-H) = 77-5 kcal./mole,28 requires D(C6H5. CH,-CH . C6H,) to be 49 kcal./mole while the value of the latter C-C bond dissociation energy is experimentally estimated by Horrex and Miles 29 as 48 kcal./mole. If D(C,H5 . CH,-H) = 80 kcal./mole as Dr. Bateman suggests then D(C6H5 CH,-CH . C6H5) must be 54 kcal./mole i.e. the values D(C,H . CH,-H) = 80 kcal./mole and D(C6H5 . CH,-CH . C6H5) = 46 kcal./mole suggested by Dr. Bateman are incompatible. Dr. B. G . Gowenlock (Swansea) (communicated) Horrex and Miles postulate the reaction (e) PhCHCH,Ph -f PhCH=CHPh + H-(128 - Q) kcal. I among the reactions resultant upon the production of PhCHCH,Ph I radicals by the attack of benzyl radicals upon dibenzyl.This reaction is endothermic to the extent of about 58 kcal. being taken as 70 kcal. and therefore an activation energy of at least 58 kcal. will be required for this reaction. This activation energy is greater than that observed for the primary reaction. From the data given in Table I and assuming a 26 Szwarc J . Chem. Physics 1950 18 1660. 27 Szwarc and Taylor Trans. Furaduy SOC. (in press). 28 Szwarc J . Chem. Physics 1948 16 128. 29 Horrex and Miles this Discussion. GENERAL DISCUSSION I 230 normal value for the temperature independent factor ( 1 0 ~ ~ sec.-l) i t can be shown that for the lower temperature data (630-700’ C) decomposition of the PhCHCH,Ph radical will take place to the extent of only 10-50 yo.Similar coksiderations apply to reaction (i). Reaction ( f ) 2PhCHCH,Ph -+ PhCH=CHPh + PhCH,CH,Ph + (29- 128) kcal. on the other hand is exothermic and should have a much lower activation energy. In contrast to reaction (e) i t will produce only one molecule of stilbene per two PhCHCH,Ph radicals. Therefore on the basis of the authors’ reaction mkchanism a variation of the toluene/stilbene and toluene/styrene ratios should be obtained with variation in temperature. Investigation of the pyrolysis of mixtures of benzyl iodide and dibenzyl a t about 5 0 0 O C should therefore lead to the virtual elimination of re- actions (e) and (i) together with their dependent reactions (g) ( h ) and ( j ) and thus lead to toluene and stilbene as the sole reaction products in the ratio of 2/1.Dr. C. Horrex and Dr. S . E. Miles (St. Andrew) (communicated) In reply to Dr. Gowenlock we wish to point out that our analyses showed no significant variation in the composition of the products with the tem- perature of reaction. When introducing our paper we drew attention to the endothermicities of reactions (e) and (i) and pointed out that these decompositions would have to have normal temperature-independent factors of about 1 0 1 ~ in order to be significant in our conditions. We consider that the uncertainty in the value of Q makes i t unprofitable to pursue such calculations at present since an error of 4 kcal. alters the estimated rate of (e) or (i) at 630’ C by a factor of 10.Since the reactions appear to be feasible and a homogeneous chain sequence improbable we prefer to await the result of further work which one of us is carrying out. It must be noted that although the primary decomposition is not a heterogeneous reaction the experimental conditions favour access to the wall and we cannot at present assess the part i t may play in radical decompositions. The extent of a homogeneous disproportionation re- action which requires the encounter of two radicals is very dependent on the concentrations of the latter and hence on the rate constants of (d) (e) and ( 2 ) . It may be significant to note that the early work which reported stilbene and toluene as main products was done with much higher concentrations than we employed ; this would result in decreased accessibility of the wall to the radicals and reactions (e) or (i) which might occur there would be less prominent.Investigations in progress include work on the reaction of benzyl radicals with other molecules. With reference to the points made by Dr. Bateman we have stated that we evaluated the best straight line by use of the high partial pressure data since in our opinion the analytical precision in the other experiments was impaired by the small amounts of products available to us. We consider this procedure legitimate and a closer examination of our Fig. 7 shows that the spread of points in the data used is small (with the noted exception of one point which by any test must be regarded as a faulty experiment).The point concerning the revision of the length of the central C-C bond in dibenzyl has also been drawn to our attention by Prof. Cox. The amount of the shortening and its relationship to the strength of that link are of importance and in summarizing Szwarc’s arguments we may have given the impression that the strengthening is dependent OR the value for this shortening. Actually any “predicted ” value for the dissociation energy of the central bond in dibenzyl depends on the evalu- ation of the heat of formation of the benzyl radical. This has been done in several independent ways a summary of which has been given recently.30 30 Quart. Rev. 1951 5 42. 23 1 GENERAL DISCUSSION The agreement shown by these methods appears to us to be convincing evidence for the higher value for the resonance energy of the benzyl radical.Dr. W. A. Waters (Oxfovd) said In connection with the paper by Horrex and Miles on the pyrolysis of dibenzyl I should like to draw attention to the rather different conclusions of Dr. A. F. Bickel and myself 31 concerning the free benzyl radical which we prepared by de- composing w w'-azotoluene in boiling decalin solution at about 200' C. We found that even at this low temperature the benzyl radicals did not just recombine to form dibenzyl but underwent about 35 yo dispropor- tionation t o toluene and stilbene. We did not detect either benzene or styrene which may be secondary pyrolysis products of stilbene. Though our results could perhaps be attributed to a very rapid de- hydrogenation of dibenzyl by benzyl radicals i t was significant that the decalin solvent was not dehydrogenated at all.The alternative ex- planation would therefore require both an abnormally high lability of the C-H groups of dibenzyl and a very low probability of recombina- tion of benzyl radicals. Unlike Szwarc we see no fundamental objection to postulating the simple reaction . 2Ph. CH . + Ph . CH + Ph . CH. It is interesting also to note that the activation energy for the vapour phase dissociation of dibenzyl accords with the experimental data given by Ziegler 3 2 for liquid phase dissociations of many similar compounds some of which give radicals that undergo disproportionation at temper- atures as low as 1 5 0 O C.Dr. M. Szwarc (Manchester) (communicated) Dr. Waters suggests the possibility of disproportionation of benzyl radicals i.e. D(C6H5 . CH-H)-D(C,H The activation energy of this reaction should certainly be less than 5 kcal./mole if reaction ( I ) is to compete successfully with dimerization. This follows from the approximate equality of A factors for both dis- proportionation and dimerization (since the activated complexes for both reactions are very alike and in the liquid phase the deactivation of " hot " djbenzyl molecules cannot be the rate determining step). The activation energy of a process must be a t least equal to its endothermicity i.e. . CH2-H) < 5 kcal./mole D(C6H . CH-H) < 82.5 kcal./mole. CH3). On the other hand D(C6H5 .CH,-H) +D(C,H5 . CH-H) +D(C,H,-CH) +D(C-H) = 3AHf(H) + f AHf(C6H6) -AHf(C6H5. Taking D(C6H5. CH,-H) = 77-5 kcal./mole D(C,H5. CH-H) < 82.5 kcal/mole D(C-H) = 80 kcal. ; AH,(H) = 52 kcal./mole ; AH,(C) = L ( L being the heat of sublimation of carbon) and finally AH,(C6H5) ~3.68 kcal. /mole (the latter value being derived from a plausible assumption D(C6H5-H) m IOO kcal./mole) one derives D(CsH,-CH) > L - 28 kcal./mole. The heat of sublimation of carbon is probably not less than 136 kcal./mole and therefore D(C6H,-CH) > 108 kcal. /mole. The latter value appears to be much too high. It might be reasonable if the activation energy of (I) is about 25 kcal./mole but then this re- action could not be observed.3l Rec. truu. cbim. 1950 69 316. 32 Annalen 1942 551 161 ; 1950 567 134. GENERAL DISCUSSION 232 It seems to me that the results of Bickel and Waters could easily be interpreted on similar lines to those of Horrex and Miles. The reason why benzyl radicals do not dehydrogenate decalin seems to be simple the C-H bond dissociation energy in decalin is higher than D(CQH6. CHZ-H) and thus the activation energy for dehydrogenation is too high. Dr. C . Eorrex (St. Andrews) (communicated) In reply to Dr. Waters I think the products obtained from the interaction of benzyl radicals depend on the experimental conditions. The conditions in solution and in dilute gaseous systems are distinctly different and even in our work with gaseous systems we have observed differences in the proportions of the products when the only substantial change has been the pressure of the inert carrier gas.Dr. Ruth Lapage when using toluene as an acceptor for methyl radicals at 500 to 600" C in a flow technique with 3 to 8 mm. nitrogen found the main solid product was dibenzyl with a small amount of stilbene. With 600 mm. nitrogen however the product was stilbene. Mr. C. B. Cowan has obtained similar data with phenyl radicals and toluene at 600" to 700" C. In the toluene pyrolysis Mr. J. 0. McCrae and Mr. R. B. Cundall examined the solid products from an experiment a t 850°C when using the pyrolysis technique as published by Dr. Szwarc. By use of u.-v. absorption spectroscopy' they' found the solids contained dibenzyl stilbene and anthracene in the ratio 380 7 I.Thus Dr Szwarc is substantially correct in speaking of dimerization only under his conditions. It was also found that using 1-5 mm. toluene in 5 mm. nitrogen carrier gas gave even smaller amounts of stilbene and anthracene at 850" C. In toluene pyrolysis the benzyl radicals are formed at temperatures where dibenzyl would be rapidly decomposed and dimerization probably occurs at the exit of the furnace. Since the gases are cooled quickly there will be little opportunity' for attack by the radicals on the first fractions of the dibenzyl formed. In the work of Lapage and Cowan the temperatures required should permit dibenzyl to be formed within the furnace and remain unpFolyzed. If stilbene was formed from di- benzyl by radical attack it seems that higher inert gas pressures materially assist in the formation of the dibenzyl.If benzyl radicals disproportionate I would have expected the reaction to be rapid at the temperatures of toluene pyrolysis and more stilbene produced than has been found. Dr. Miles used w a'-azotoluene as well as benzyl iodide to provide benzyl radicals when examining the attack of these radicals on dibenzyl. The products in both cases were the same and similar to those obtained from dibenzyl alone on pyrolysis. I would expect the abstraction of a hydrogen atom from decalin by a benzyl radical to be a distinctly endothermic process and the data given by Dr. Kooyman in his Tables I and I1 support the idea that hydrogen abstraction from dibenzyl is an easier process.With reference to Ziegler's data I would like to add that his tem- perature independent factors are generally high. Prof. S . W. Benson (California) said The data of Horrex and Miles seem to indicate a steric factor which is unusually low for a unimolecular decomposition involving the rupture of a single bond and an activation energy lower than that which would be expected for a C-C bond if the proper length of the bond were employed. In addition the dependence of the rate on the surface/volume ratio shows a strong temperature effect. Thus the ratio of rates in a packed to unpacked flask fall regularly from a factor of 3 at 635O C to about 1-4 at 720° C . The authors make the assumption that the ratio of heterogeneous to homogeneous reaction is proportional to the surface/volume ratio.This is by no means necessarily the case and must certainly depend on the rate and energy of adsorption of reactant on the surface as well as on the overall GENERAL DISCUSSION 233 mechanism. In particular the dependence may vary from complete independence to direct proportionality according to the conditions. Further the pressure range investigated (about 0-1 mm. Hg reactant) is that in which wall collisions proceed a t a rate comparable to intermolec- ular collisions. In the light of these uncertainties i t would seem premature to accept with any certainty the value of 48 kcal./mole and the low frequency factor for the unimolecular homogeneous decomposition of dibenzyl. Dr.C. Horrex and Dr. S. E. Miles (St. Andrews) (communicated) Prof. Benson is in error in stating that in our conditions of 0.1 mm. Hg reactant pressure the wall collisions are about as frequent as intermolec- ular ccl'isions. He appears to be ignoring the presence of the carrier gas which impedes diffusion to the wall considerably. Calculation of the average number of collisions made by a molecule diffusing to the wall 33 gives about 105 for the pressures and reaction vessel used. By considering his comments on surface effects in conjunction with details given in his previous contribution on this topic we note that the effects of changes in surface/volume ratios are difficult to interpret where chains are initiated and end at the walls. In our case we see no reasonable evidence for a chain mechanism and consider the relatively small influence of the large change in surface implies little heterogeneity in the primary reaction.This is not in conflict with the possibility (mentioned in reply to Dr. Gowenlock) that some radicals produced in secondary reactions may reach the walls and undergo reaction there. Prof. Benson's assumption that the expected bond dissociation energy depends on the length assigned to the central C-C bond has been dealt with above. We agree of course that the temperature independent factor is unusually low in our experi- mental conditions. It seems to us that it is desirable to have present an acceptor other than dibenzyl for the removal of the benzyl radicals ; data on the thermal stability of potentially suitab!e substances is being completed.In view of findings on the pyrolysis of methyl iodide in this pressure range given later we are also checking to see if the high pressure limiting rate has not been attained. Dr. C. Horrex and Dr. Ruth Lapage (St. Andrews) (communicated) Some points disclosed by a detailed examination of the pyrolysis of methyl iodide seem to be of interest in connection with the kinetic difficulties in this pyrolysis work a t low pressures and particularly since C-H and other bond dissociation energies have been derived from published work on iodides. Butler and Polanyi 34 report five experiments on the pyrolysis of CH31 at 493'-495" C and by calculating first order constants and using log k (sec.-l) = 13 - E/4*57T they derived E = 54 kcal.With this value assigned to the C-I bond dissociation energy they deduced a C-H value €or CH which proved in agreement with later determinations by other methods. We have used the same techniques and conditions and found that their quoted values of k can only be realized a t temper- atures about rooo C higher. The use of the correct temperature in the above equation would give a value of E of about 60 kcal. and this seems too high for the C-I dissociation energy. The validity' of assuming first-order behaviour and a 1013 factor under such experimental con- ditions was therefore examined. With 0.07 to 0.3 mm. CH,I in 3 mm. nitrogen carrier gas the kinetics approached second order and changed towards first order above I mm.of CH31. In these conditions the re- action was mainly' homogeneous and not inhibited by iodine the free methyl radicals could be trapped by toluene but in its absence they were decomposed quantitatively' a t the wall to methane and carbon. The rate of reaction was increased by increase in pressure of the carrier gas. s3 Bursian and Sorokin 2. $lay&. Chew. B 1939 12 247. ,* Butler and Polanyi Trans. Faraday SOC. 1943 39 19. GENERAL DISCUSSION g mm. N Goo mm. N z mm. C,H 3 mm. N 1-1-3 mm. 3 mm. N 0.3 mm. 1-2 mm. CH,I CH,I 1-1.3 mm. CH31 0.0066 0.028 CH,I 0.0056 A A na rp A A 234 By using 600 mm. Nz z mm. toluene and 1-2 mm. CH,I we obtained rate constants showing a temperature variation given by k (sec.-l) = 2.6 x 1olS exp -547oo/RT.The strong influence of experimental conditions on the rates of decomposi- tion can be seen from the following values of k calculated on a first order basis. Conditions K (sec.-l) at M~ I + 0.2 MM. 1,) - FIG. 5. 0.0016 Soo0 K The presence of a radical acceptor was essential in experiments at high total pressures since with the decreased accessibility of the wall to the radicals a recombination reaction became important. An analvsis of our data has beeimade which considers activation of the CHJ molecules as occur- ring by collisions with CH31 or N mole- cules the latter having about one-fifth the efficiency of the former. Applying the con- ventional test for data on second- to first- order transitions of plotting the reciprocal of a first-order constant against the recip- rocal of the concentrations (appropriately weighted) we obtain Fig.5 The experi- mental point on the vertical axis clearly represents the high pressure limiting rate for the reaction. Details of this work are being prepared for publication. These findings while not altering the C-I bond dissociation energy make the initial derivation of the 54 kcal. value seem rather fortuitous. Similar troubles may exist with other substances in the Butler- Polanyi conditions although one might expect the fall off in first-order constants for more comdicated molecules to occur a t lower pressures than for CHJ. A number of C-H bond energy values which rest on the original iodide pyrolysis determinations may' be uncertain for these reasons.A 2.0 53 &A A A /;o Dr. Horrex and J. 0. McCrae (St. Andrews) said Dr. Kooyman may be interested in our preliminary results on the pyrolysis of diphenyl- methane which show that the products are essentially tetraphenylethane fluorene and hydrogen. The technique of investigation has been similar to that used for dibenzyl (see this Discussion). We assume the primary process is the breaking of the methylenic C-H link and that fluorene production may proceed through the planar diphenylmethyl radical radical. The calculation of the rate of this primary process via the hydrogen production (allowing for fluorene formation) yields a rate constant k = 10'~ exp (- 73ooo/RT).We are extending the range of experimental variables but suggest that this energy of activation can be considered as the C-H bond dissociation energy Prof. 6. W . Benson (California) said There seems to be some question in the work of Rowley and Steiner of the role played by each of the three possible simultaneous reactions cis plus cis ; cis plus trans and trans plus trans (butadiene). W e can write a general rate expression hobs = a2k + z a ( ~ - a)kd + (I - a)'LKtl where a is the fraction of the cis form and K, K,t and Kg are the specific rate constants for the individual reactions of the species indicated. If i t GENERAL DISCUSSION 23 5 is assumed that the active complex is the same for all of these cases then it can be easily shown that the observed activation energy is not dependent on the extent to which the different species react since a = e -AE/RT ; E, = Ell - zAE = E,# - AE and the temperature dependence of the 01 exactly cancels the different temperature dependence o f the different rate constants.In this case the above expression reduces to kobs = ktt = a'ku. If however there is more than one possible complex for the system and the steric factors are different for the different isomers then we should expect a different contribution at different temperatures from the different species. Thus i t may be quite possible that a cyclic and linear complex both exist having different activation energies and different frequency factors. In this case the system becomes much more complicated and alternative explanations must be considered for interpreting the data.Dr. M. Magat (Paris) said There is one point that strikes me as not being taken enough into account in all theoretical discussion of butadiene cyclization reaction. It is the fact pointed out by Aston Szasz Wooley and Brickwedde 35 that the trans configuration is 2.5 kcal./mole more stable than the cis configuration. On the other hand at least one of the two molecules must have the cis configuration in order to make the cyclic complex possible while the linear complex can be realized with any one of the two. Mr. B. Eisler and Dr. A. Wasserman (London) (communicated) Experiments have been carried out which enable a comparison to be made of the activation energies E of Diels-Alder diene associations of the following type (1) Butadiene + dienophile -+ adduct Cyclopentadiene + dienophile + adduct .* (11) The difference between E and EII can be accounted for,ss if it is assumed that the formation of the transition state of (I) involves the conversion of trans-butadiene into the cis form. In the reaction discussed by Steiner (14 2 Butadiene -+ vinylcyclohexene 2 Cyclopentadiene -+ dicyclopentadiene and collaborators. . * the dienophile of (I) is butadiene. A comparison of the E values of (Ia) and (IIa) indicates that both butadiene molecules are converted into the cis form before the transition state of the reaction (Ia) is taken up. . . . (IIa) This is due firstly to steric requirements and secondly to effects which operate generally in Diels-Alder reactions and which bring about the formation of relatively closely packed transition Dr.H. Steiner (Petrocarbon Ltd. Manchester) (communicated) I n reply to the remarks by Benson Magat and Wasserman a statistical rate calculation of the butadiene dimerization using the values for the thermodynamic functions of butadiene obtained by Aston Szasz Woolley and Brickwedde,s* may give an answer to some of the questions raised. Since the interconversion of trans- into cis-butadiene is endothermic by only 2-5 kcal. /mole whereas the activation energy of the dimerization is some 25 kcal./mole the initial state which has to be assumed in such a calculation is most likely' an equilibrium mixture of trans- and cis- butadiene at the appropriate temperature.It may be that the deviations from the Arrhenius function which we have observed are connected with the gradual shift of the cis-trans-butadiene equilibrium particularly at low temperatures. Unfortunately our statistical calculations were carried out using approximate values for the partition function of buta- diene only. I hope to recalculate this prablem using the data of Aston 85 J . Chem. Physics 1946 14 67. 36 Detailed considerations and experimental results to be reported elsewhere. 37 J . Chem. Soc. 1935. 833 1512a8 1936 432 ; Trans. Faraday SOC. 1936 J Chem. Physics 1946 14 67. 35 841. D. ROWLEY AND H. STEINER 213 GENERAL DISCUSSION 11. HYDROCARBON REACTIONS A.THERMAL REACTIONS Map I point out rather an interesting correlation between results quoted in papers by Sir Cyril Hinshelwood and Prof. Glockler ? The latter shows that the C-H and C-C bond energies in n-paraffin chains increase in passing from the centre towards the ends ; the former states that in the thermal chain fission of n-paraffins the probability of C-C fission (and Dr. P. Torkington (Brit. Rayon Res. Assoc.) (communicated) 214 GENERAL DISCUSSION hence also of C-H fission accompanying H-transfer) also increases to-wards the ends. I suggested earlier in this Discussion that Prof. Glockler’s results might be interpreted as evidence for the existence of molecular orbitals (overall but not necessarily delocalized in the Sense of implying electron mobility) favouring the chain ends.The correlation here implies that the presumably higher terminal electron density is associated with greater activity i.e. electron mobility an interesting negative analogy with true bimolecular reactions. Regarding the concentration of energy via vibrational coupling I should like to draw attention to another previous remark of mine following Prof. Ubbelohde’s paper. The envisaged concentration of energy could surely never occur with “ perfect ” coupling ; it seems that if the occur-rence of energy concentration by vibrational coupling is accepted as proved then so also must the occurrence of two or more collisions pro-viding the energy sources. Transfer from two or more such sources by coupling might well lead to a concentration of energy favourable to re-action by’ normal interference.It seems quite probable that a long-chain molecule might be involved in several simultaneous collisions but against fundamental principles for a concentration of energy to arise from a single collision. Finally is it possible for a little more non-mathematical commonsense to make a little more headway against a little more formalism ? How does symmetry lead to electro delocalization 9 Why because one can write functions 4 = ,Zai$i because of symmetry does an electronic system become mobile ? Perhaps too much n-bond theory has preceded a study’ of normal molecules. In an unsymmetrically-substituted benzene de-rivative there is no symmetry but still the same delocalization as in the parent compound. Dr.G. Gee (Welwyn Garden City) said I should like to outline a possible explanation of the observed thermal decomposition products of paraffins. Considering first radical mechanisms to a first approxima-tion the thermal dissociation of a long paraffin chain will occur randomly at any link. The resulting radicals will then tend to degrade by the facile process of losing a series of ethylene molecules from the end, and this process will continue until a methyl or ethyl radical remains. Thus primary dissociation of long chains will result mainly in the produc-tion of ethylene and ethyl and methyl radicals and the ratio of these primary products will not be markedly temperature dependent. The most probable secondary reaction of a methyl or ethyl radical appears to be the dehydrogenation of a paraffin molecule a process which again will occur nearly randomly along the chain (3) and will be followed by dissociation of the resulting radical (4) CH3* + CH,(CH,)nCH + CH + CH,(CHa),*CH(CHa),-,,CH ( 3 ) CH3(CH2).*CH(CH2),-,,CH3 + CH,(CH,),CH=CH,+CH,(CH,),-, (4) The situation now becomes more complex in that the olefin produced will be more reactive toward further methyl and ethyl radicals than is the original paraffin.Without seeking to follow o u t these processes in detail i t can be seen that a further series of degradative processes will follow from the secondary radicals in which the principal final products will again be ethylene methane and ethane. Thus i t is easy to see why these products predominate and also why their relative yields are largely independent of conditions.It also follows from these considerations that in presence of nitric oxide the main hydrocarbon product to be expected from radical dissoci-ation is ethylene since nitric oxide would compete more successfully with The point is obvious and not worth developing. CH,(CHz)nCH + CH3(CHa)z + CH,(CHa)L - - ( 1 ) CH3(CH,),* + CH3(CH2)g2 + CH*CH2 * (2 GENERAL DISCUSSION 215 (3) than with ( 2 ) . The fact that methane and ethane are still produced in presence of nitric oxide therefore argues in favour of Stubbs and Hinshelwood’s contention that a molecular reaction occurs. The mechanism of this reaction does not appear very clear but it may be plausibly sug-gested that ethylene could be eliminated from the chain at any point through the formation of a cyclobutane type of activated complex : CH,-R CH,R I -+ I + CHeCH,+R-R .- ( 5 ) CH,-R CH,-.R, where R R are alkyl radicals or hydrogen. It is not a t first sight clear why this reaction should lead to methane and ethane rather than to larger paraffins but it appears likely that increase in the size of the radicals R and R would lead to steric difficulties. If this is indeed the case each successive step in the degradation of a long chain (by removal of ethylene) will lead to a more readily dissociable paraffin and-as in the radical decomposition-the principal final products will be ethylene methane and ethane. Dr. F. J. Stubbs and Prof. C. N. Hinshelwood (Oxford) (co.mununi-cated) In connection with Dr. Gee’s interesting suggestions we may state for comparison the views which we have formed about the reason for the similarity of products from a chain reaction and a molecular process.For the typical case of n-hexane the molecular process would be CH,CH,CH,CH,CH,CH -+ CH + CH CHCH,CH,CH, where El and El’ are the activation energies of the two reactions. Initia- CH,CH,CH,CH,CH,CH 3 X + Y (1) (1’) (2) CH;+CH,CH,CH,CH,CH,CH +- CH4+’CH,CH2CH,CH,CH,CH, (2’) C,H; +CH,CH,CH,CH,CH,CH -f C,H,+’CH,CH,CH,CH,CH,CH, where e, e, etc. are the respective activation energies. El CH,CH,CH,CH,CH,CH -z C,H + CH CHCH,CH, El’ The various steps in the chain reaction would be tion { X+CH,CH,CH,CH,CH,CH + XH+’CH,CH,CH,CH,CH,CH, DCH,CH2CH2CH,CH,CH -+ CH;+CH CHCH,CH,CH, *CH,CH,CH,CH,CH,CH -+ C,Hi+CH CHCH,CH, With ling chains the ielative proportions of methane and ethane formed will depend only on the repeating steps ( I ) and (1’) in the above scheme and their sequels ( 2 ) and (2’).The nature of the initiating re-action is unimportant for this purpose. For the products from the molec-ular process and chain reaction to be in the same proportions all that is necessary is that (&-El’) should equal (el-el’). This seems quite likely since the alternative steps in the molecular process are closely parallel with the steps (I) and (1’) respectively the two positions of rupture being in the same relative position in the radical CH,CH,CH,CH,CH,CH; as in the molecule CH,CH,CH,CH,CH,CH,. Dr. M. Szwarc (Manchester l7niversit-y) said Stubbs and Hinshelwood proposed in their paper an interesting mechanism for a unimolecular decomposition.I would like to examine some of the conclusions which, it seems to me follow from their assumptions. Let us consider a system composed of two components viz. an inert gas X and a reacting gas M which may decompose unimolecularly. Furthermore let u s choose the experimental conditions in such a way that the energy supplied to molecules M arises from collisions between M and X. According to Hinshelwood and Stubhs we can observe two kinds of collisions. (if One type of collisions between M and X lead to the formation of energized molecules M* i.e. molecular species which contain an amount of energy at least sufficient for their decompositio 2 16 GENERAL DISCUSSION but distributed unfavourably amongst various internal degrees of freedom.This process we represent by the equation X + M + M* + X (collision of the 1st type). The decomposition of energized molecules requires a reshuffling of their energy which has to be accumulated eventually in the reacting centre, i.e. M* +- MS + Products, MS denoting an " activated " molecule i.e. a species which contains at least the necessary amount of energy localized in the reacting centre. Although collisions of the first type are frequent the probability for the energy redistribution is low and consequently the reaction caused by these collisions attains a first-order character at comparatively low pressure of the gas X. Plotting the rate of this decomposition as a function of partial pressure of gas X we obtain the curve I shown in Fig.I . (ii) Collisions of the second type lead directly to the formation of I ' activated " species MS X + M -t MS + X (collisions of the 2nd type). Since the chance of the occurrence of such a collision is very low and the probability for the subsequent decomposition of Mt is very high, the reaction caused by collisions of the second type retains the second-order character up to comparatively high pressures of the gas X. Plotting the rate of the latter reactions as a function of pressure of the gas X we obtain the curve 2 shown in Fig. I. FIG. I. Fig. I illustrates the features of the decomposition claimed by Hinshelwood and Stubbs. There is a comparatively low pressure region (C-D) in which the first type of collision predominates and in which the reaction obeys first-order kinetics.There is a comparatively high pressure region (E-F) in which collisions of the second type are predominant and conseqfiently in this region the reaction obeys second-order kinetics. However in addition to these two regions there is a very low pressue region (A-B) in which both types of collision call for a decomposition of the second order and finally there is a very high pressure region (G-H) in which both types of collision demand first-order kinetics of decomposition. If we plot therefore the order of reaction as a function of the partial pressure of gas X we obtain a curve represented in Fig. 2. I would like to ask Prof. Hinshelwood if he agrees that this behaviour of decomposition is a consequence of his assumption.There is however another point which I consider to be important. If the reaction takes place in a very high pressure range say G-H then the first type of collision must lead to an equilibrium between the energized molecules M* and the ordinary molecules M i.e., X + M + M* + X (collisions of the 1st type) GENERAL DISCUSSION 217 I n the same pressure region collisions of the second type must lead to an equilibrium between the " activated " molecules MS and the ordinary molecules M i.e. X + M $ MS + X (collisions of the 2nd type). One has to realize that any type of " activated " molecules MZ can be produced either by collision of the 1st type via M* or directly by collision of the 2nd type. If the latter modes of activation leads to the state of equilibrium between ordinary molecules M and activated species MZ , then i t would be expected that the activation through collisions of the first type should also lead to an equilibrium between these two species (this follows from the fact that both types of collisions occur in the same gas phase).Consequently the following scheme should be valid for a reaction resulting from collisions of the first type : M + X + M* + X Jr MZ -f Products, and the equilibrium concentration of MZ should be the same as in the case of collisions of the 2nd type. That however would contradict the as-sumptions on which Fig. z was obtained. I Px FIG. 2. One can argue that the collisions of the second type do not lead to an equilibrium concentration of Mi since their rate of decomposition is too high being of the order of 1013 (MS) sec.-l.This seems to be quite correct. Indeed collisions of the second type can never be of any significance and if they have to maintain a first-order character of decomposition a pressure of the order of 105 atm. would be required. It seems that the best method of approaching the problem is t o sum up aZZ the contributions arising from all possible collisions (appropriately weighed). The reaction would be of the second order a t low pressures and would attain the first-order character at sufficiently high pressures In the latter case the activation energy of decomposition would be given by the activation energy of the process M +M$. Dr. F. J. Stubbs and Prof.C. N. Hinshelwood (Oxford) (communic-ated) Given the hypothesis that the nitric-oxide inhibited reaction is a definite molecular reaction (which as we have explained there is a good prima facie case for accepting) the changes of order referred to by Dr. Szwarc are the basic experimental facts upon which our views about the kinetically composite nature of the reaction are founded. The tentative interpretation which we have advanced is not in any way built upon a preconceived theory and may be modified in the light of further experi-ments now in progress. In the meantime we would prefer not to regard the facts about the order as deductions from a theory as Dr. Szwar 218 GENERAL DISCUSSION implies but as observations calling for explanation. What we have suggested is that the normal theory requires some degree of amplification.Prof. S. W. Benson (California) (conznzunicated) The data presented by Ingold Stubbs and Hinshelwood and quoted here seem to indicate that even in the " fully inhibited " pyrolysis there is a surface effect which increases with decreasing temperature and is virtually independent of pressure. It seems to me that until this effect is further elucidated i t is difficult to accept the discussion and rate constants quoted as anything but tentative. There still remains the possibility that there is a heterogeneous reaction initiated on the walls which starts chains which also end at the wall. Rice and Herzfeld 1 have recently shown that under such conditions i t is possible that depending on the mechanism the rate may show any de-pendence on the surface volume ratio between complete independence and direct proportionality.Even in a direct decomposition to molecular products i t is conceivable that heterogeneous decomposition competes with homogeneous decom-position. In such case the activation energy of the two would be ex-pected to be different and since the relative amount of each would be expected to show a monotonic if not simple pressure dependence there would be observed an effect of pressure on experimental activation energy. For n-pentane i t is difficult to believe that when the experimental activation energy is found to vary between 80 and g o kcal./mole in the pressure range 50-150 mm. Hg (inhibited reaction) t h a t the free radical mechanism of splitting of the C-C bond which has an activation energy of 80-85 kcal.is not proceeding at a competitive rate. Finally it is difficult to reconcile the observed decrease of activation energy with increasing pressure with any reasonable behaviour of mole-cules. In the Rice-Ramsperger-Kassel theories of the unimolecular decomposition the activation energy is supposed to decrease at the lower pressures where the higher energy states can no longer make their com-plete contribution to the overall rate. The explanation offered by the authors to the effect that there may be a single mode which requires a very long time to reach activation even when the energy is present in other modes seems completely ad hoc and unsupported by any other experimental or theoretical evidence.It also leans upon an interpreta-tion of a molecular collision that seems in immediate conflict with the law of microscopic reversibility. That is there seems to be an implicit assump-tion that an active molecule can only lose energy on collision and never gain it. If there is a mode which is only slowly activated then the chances are that i t is just as slowly deactivated and collisions do not necessarily play any decisive role in the time of decomposition. The example of ethane where the kinetics seem to go through a second to first order transition while the activation energy changes but little (or possibly is higher a t the lower pressures) seems a striking example of the above-mentioned contradiction. I doubt very much that these pyrolyses will be very tractable theoreticallv until the nature of the surface efiects and the inhibition axe better understood.Dr. F . J. Stubbs and Prof. C. N. Hinshelwood (0,vford) (conz~lzunic-nted) Chain reactions which are initiated by a decomposition are not at all likely to start and stop at a wall. If the molecule dissociates on the wall radicals are likely t o remain adsorbed. In any case chain reactions staxting at a wall should be subject to the usual kinds of inhibition just like any other chain reactions. This assumption of Prof. Benson seems to us to be improbable. In pentane we have always believed the chain reaction normally to proceed a t a " competitive rate " but we state a prima facie case for concluding that We are unable to understand his fourth paragraph.1 Symposium on Abnormalities in Reaction Kinetics Amer. Chem. SOC., Minn. hfinn. Spring 1950 GENERAL DISCUSSION CH3-CH3 CH3. CH,-CH3 CH3. CH,-CH,. CH3 CH . CH,-CH . CH . CH, 219 80 CH CH-H I00 80 60 < 60 1 slightly ( 8 0 CH . CH . CH,-H 1 due to second- CH CH . CH,-CH3 J order effects CH CH . CH,-CH . CH3 ’lightly this competing reaction has in the conditions of the experiments been suppressed. In his penultimate paragraph we do not recognize our views very clearly and as to the last paragraph we can assure him that our approach continres to be essentially experimental and that all the details of these admittedly rather puzzling phenomena are under examination still and from several new points of view. Dr. L. Bateman (Welwyn Garden City) (communicated) The form and relative position of the paraffin and olefin decomposition rate curves in Fig.I of Stubbs and Hinshelwood’s paper are just as would have been predicted on the basis of A H changes on breaking the most easily broken bonds so that some reference to this mode of primary bond scission is perhaps of interest. The AH changes (in round figures) for the C,-C, compounds are summarized in the following Table. Ethylene should be I I much more stable than ethane propane and propylene should be of similar stability and the higher members of the two series should exhibit a difference reflecting the partial utilization of allylic resonance energy in the bond-breaking process. These features are evident experimentally, but ethane and I-butene appear to be relatively more stable than might be expected and may serve to illustrate how secondary processes affect the overall decomposition.The common factor is the primary formation of methyl radicals and not higher alkyl radicals capable of ready decom-position into smaller fragments. Hence in ethane recombination will be specially favoured since the only alternative reaction is the rather difficult hydrogen exchange Me + C,H + MeH + C,H, In I-butene the corresponding reaction to the latter (E N 10-15 kcal.). Me + CH CH . CH,Me -f MeH + CH CH . CH . Me ( E - 5 kcal.) 1 is much more facile but the resulting substituted allyl radical like that initially produced will be far less active than an alkyl radical in inducing further decomposition.An important aspect of these considerations is that all A1-olefins are required to yield simple allyl radicals as primary fission products. The reactivity of these radicals under the decomposition conditions their ultimate products and possibly the estimation of their extent of forma-tion might well be stcdied experimentally by carrying out co-decomposi-tions with I 5-hesadiene as a foreign source of allyl radicals. Informa-tion on the third point would establish the contribution of allylic bond primary scission to the evidently composite overall decompositi m. Dr. R. G. Partington (Oxford) (communicated) In their paper Stubbs and Hinshelwood state that Steacie and Folkins suppnrt the view that the reaction taking place In the presence of sufficient nitric oxide to reduce the rate to a minimum (the “ residual reaction ”) is “ the primary process of what in the absence of inhibitors would be a chain reaction ”.This is, I think a mistaken interpretation of Steacie and Folkins’ conclusions, Steacie and Folkins Can. J . Res. B 1940 18 I 2 20 GENERAL DISCUSSION which read '' the addition of nitric oxide then merely diminishes the chain length without completely suppressing the chains ". Had the suggestion been made that all the initidly formed radicals reacted with nitric oxide one would have to consider the possibility of the formation of such a molecule as (I) in the following scheme e.g., CH,-CH,-CH,-CH,-CH3 = CH,+CH,-CH2-CH2-CH3 CH3+N0 = CH3N0 CH3-CH2-CH,-CH,+ NO = CHs-CH~-CH~-CH2-NO (1) which might well participate in further reaction (starting by its decom-position) in a similar way to the large radical itself.A further paper by Steacie and Folkins 8 is of interest in this connection. It is concerned with the decomposition of n-butane initiated by free radicals from ethylene oxide and the authors conclude that their results " cast some suspicion upon the idea that maximum inhibition by nitric oxide in all cases corresponds to complete suppression of chains *'. The whole problem of the action of nitric oxide as ail inhibitor (and the foregoing remarks were largely included to justify the opinion that the problem still exists) seems to depend on the relative rates of the two reactions viz., reaction of radicals with the hydrocarbons e.g., and reaction of radicals with the inhibitor e.g., A C R + RH = R,H + R .* (1) R + NO = R,NO. . * (4 Reactions of the type (I) have recently been studied and similar specific attention to reactions of the type ( z ) , would undoubtedly be of the greatest value in helping to solve the above problem and might be less subject to ambiguity than experiments on a system in which reactions of this type are among the very many taking place. There are many possible sources and convenient methods of producing the appropriate radicals either thermally or photochemically and in connection with the possibility of making such a study some experiments of mine on the chemical analysis of small concentrations of nitric oxide in gas samples of the size usually handled in work of this character might be of value.A gas sample (say 50ml. of gas at N.T.P.) containing say Q to I yo of nitric oxide is enclosed in the holder A (Fig. 3 ) . This is then con-nected to a flask B containing air and small amount of an aqueous solution of sodium or potassium hydroxide free of nitrite. On opening the wide-bore stop-cock C and gently shaking the reactions FIG. 3 . 2NO + 0 = 2N0 . * (3) (NO + NO,) + 2NaOH = 2NaN0 + H,O . - (4) take place. Reaction ( 3 ) is slow and the conversion to nitrite is almost complete. The solution of nitrite can then be analyzed with great accuracy by the standard colorimetric method^.^ The conversion factor NO -f nitrite is reproducible for a given nitric oxide concentration and its value could be determined over the desired range of concentrations by making up mixtures.Can. J . Res. B 1939 17 gg. 4 Trotman-Dickenson and Steacie J . Amer. Ckem. SOC. 1950 72 2310 ; 6 See for example Vogel A Text Book of Quantitative Inorganic Analysis J . Chem. Physics 1950 18 1097 and following papers. (19391 P. 700 GENERAL DISCUSSION 22 I Dr. A. F. Trotman-Dickenson (Manchester) said I am not clear why Anderson Davison and Burton have disregarded the recent work of Dodd in selecting 8.6 kcal. as the best available value for the activation energy of the reaction CH + CH,CHO -f CH + CH,CO. Dodd has shown that the results of all previous workers (including those of Rollefson and Grahame) may be interpeted as giving a value for the activation energy of 10.7 f 0.5 kcal. This value makes the agree-ment between the acetaldehyde and acetone results much worse (15.8 and 13.2 kcal.respectively). Dr. Steacie and I recently reviewed the data on the reaction CH + H +CH4 + H and concluded that the best work (three separate determinations) gave El = 8.8 kcal. steric factor = 4 f 2 x I O - ~ (oR2 = 2-8 A ucas = 3.5 A) and K / K t at 182O C = 50 x 10-1 molecules-* cm.812 sec.-4. The results of Anderson Davison and Burton give K,/kt2 at 182O C = I x 10-1, molecules-* ~ m . ~ / e set.-* if the same collision diameters were used as above. No satisfactory reasons for these serious discrepancies have been given. Prof. M. Burton (Notre Dame) (communicated) Mr. Trotman-Dickenson mLst not have realized that the full paper by Dodd did not reach our laboratory until after the paper by Anderson Davison and myself was submitted for publication.Otherwise we would certainly have commented on it. Our actual work was concerred with establishment of activation energy differences. involved also assumptions of reliability of E values of Grahame and Rollefson for and of Trotman-Dickenson and Steacie for Since the agreement of values of El based on the two separate assumptions was good one conclusion was that the two sets of results were consistent. The more recent value of Dodd (E20 = 10.7 f 0.5 kcal.) is simply not reconcilable with the older value (8.6 kcal.) of Grahame and Rollefson. It is consistent with the results of Trotman-Dickenson and Steacie only if it is assumed that our measurements of E differences are inaccurate to about 2 kcal.Never-theless the new evidence presented by us does support a high value of El and Dodd’s results emphasize that fact. Regarding the relative merits of the results of Dodd and of Grahame and Rollefson we have no firm opinion. However we may note that E, >El* is a new idea. The reverse order of activation energies ac-cepted prior to the work of Dodd is consistent with the usual notion that, since in reaction (20) a formyl H is involved while in reaction (18) it is a methyl H reaction (20) probably has a lower E,. Mr. Trotman-Dickenson’s remark concerning discrepancies between values of kl/kl24 from old experimental data and from calculation involving our El and s emphasizes the dangers inherent in acceptance of old cal-culations based on experiments which may contain innate error e.g.the general difficulty of interpretation of experiments of involved mechanism. Prof. W. A. Noyes Jr. (Rochester) said It is interesting to note in the paper by Prof. Burton and co-workers that for the reaction Calculation of E for CH,+H + CH4+H . - (1) CH + CH,CHO -+ CH + CH,CO . - (20) CH + (CH,),CO -+ CH + residue. * (18) We are cognizant of the difficulties of this work. CH + H = CH + H . (1) Dodd J . Chem. Physics 1950 18 234 ; Trans. Faraday SOC. 1951 47 56. 7 Trotman-Dickenson and Steacie J . Physic. Chem. (in press) 222 GENERAL DISCUSSION the activation energy is 13-14 kcal. and the steric factor about I O - ~ . These results are based on studies in which reaction (I) could compete with reactions of the type The activation energies for (2) depend on R but generally fall in the region of 8-10 kcal.(e.g. if R is CH,COCH,- at about 9-7 kcal.). Thus if steric factors for (I) and (2) were identical (2) would proceed about 102 times as f a s t as (I) at equal pressures of H and of RH at temperatures of 200-300° C. A ratio of rates as high as 102 would invalidate most conclusions due to attendant experimental difficulties. The very fact of successful com-petition under experimental conditions used by most authors indicates quite strongly that steric factors for (2) are usually 10-1000 times lower than for (I). This fact and other related facts raises questions about “ hot ” radical conclusions. If a radical separated from a parent molecule by absorption of energy much greater than that necessary to break the bond retains energy solely as kinetic energy i t could hardly undergo 103-105 collisions without being reduced essentially to thermal equilibrium with its sur-roundings.One must conclude either that hot radical effects would be unobservable or that the so-called steric factor is a very pronounced function of kinetic energy. Theory concerning this matter is very vague and qualitative although one is led to predict an increase in steric factor with increase in kinetic energy. An examination of published data on hot radical effects for reactions of the type of (2) reveals that positive conclusions are not warranted. Uncertainties exist either due to scatter in the data themselves or to effects other than hot radicals which afford equally satisfactory explanations.This is true particularly in those cases such as CH,COCH3 and Hg(CH,), where the full nature and yield of the primary photochemical process have not been elucidated as a function of temperature and other variables. CH + RH = CH + R. . * (2) A word should be said also about the reaction CH + CH = CpH6. . (3) All workers seem agreed that activation energies for radical combination reactions of the type of (3) are very low. Less agreement is found con-cerning the steric factor but the safest conclusion seems to be that i t is high and probably near unity. The necessity for a third body has not been clearly shown but an examination of the data a t sufficiently low pressures will indicate a trend which might show a third-body effect, More data on this point will be found in a forthcoming article by Dr.A. J. C. Nicholson as well as in a recent article by Gomer and myself.* A third body may be introduced in several ways but i t is not satisfactory to introduce a mere triple collision. One way would be in which case the Under conditions reduces to ks 1 4 CH + X c Z CH,X . (4) CH,X + CH = C,H6 + X * ( 5 ) rate of ethane formation would be of high radical concentrations (high intensities) eqn. (6) an equation which is satisfactory for part of the ethane formation during photolysis of mercury dimethyl but which is not satisfactory under con-ditions so far studied in acetone. The trend in acetone at low pressures may indicate the necessity for a three-body collision.J . Amer. Chem. SOL 1950 71 101 GENERAL DISCUSSION 223 The most satisfactory way to obtain the right rate equation for the majority of experimental conditions is that used by Gomer and others, viz. the reversible formation of an intermediate complex from two methyls, the intermediate complex being stabilized to ethane by collisions. At sufficiently high pressures the rate-controlling step is the formation of the intermediate complex a step whose rate depends only on (CH3)2. One comes to the conclusion therefore that the nature of the third body may determine the way in which i t acts in agreement with the classical experiments of Rabinowitch and Wood on the recombination of bromine atoms. Under most conditions the rate of ethane formation is given by in which h may be a complex constant which indicates nevertheless the rate of (3) to be very high.Dr. P. Torkington (Brit. Rayon Res. Assoc.) (corutmuvzicated) Is the steric factor temperature-dependent ? If so its variation might possibly be related to the proportion of molecules (in this case radicals) undergoing a mode of vibration more favourable to reaction than other modes. Thus, in the reaction CH3- + H, it might be thought that approach along the line C . . . H-H coinciding with opening of the CH " umbrella " at the correct stage (the hydrogen molecule lying on the opposite side to the hydrogen atoms of the methyl group) would have a high probability of successful reaction. I n this case then the favourable vibration is the symmetrical breathing frequency of the methyl group.The argument is not essentially altered if the group is planar as has been fairly recently suggested though the favouring would not be so noticeable. Possibly the principle could be applied generally. As regards feasibility with a frequency of 1500 cm.-l and amplitude of the order 10-g cm. the hydrogen atoms in the breathing mode of a methyl group have a mean velocity of the order 105 cm./sec. ; the root-mean-square velocity of hydrogen molecules at 2 5 O C is about 2 x 105 cm./sec. The example shows that vibrations might co-operate. Dr. George Porter (Cambridge) (communicated) The high value for the bimolecular rate constant of methyl radical combination at room temperature of about 5 x 1013 cm.a mole-I sec.-l quoted by Sztvarc and others in the discussion now seems well established having been obtained by several workers using the sector method There appears a discrepancy, however with the results of mirror experiments which cannot be dis-regarded on the grounds that radical removal under these conditions is known to be a wall reaction.Whatever the mechanism by which the radicals disappear it is the rate-determining one and thus sets an upper limit for the homogeneous bimolecular rate constant. If the concentration of radicals is low as in the photochemical experi-ments of Norrish and Porter O the rate constant is found to be high and no discrepancy appears but as the concentration is increased so is Kmax in-creased for the radical lifetime does not decrease proportionately.Thus Forsyth's values 10 give kmax = 6 x 1012 cm.3 g . mole-1 sec.-l and although other workers have not recorded concentration measurements specifically i t seems certain from a consideration of mirror removal times that even higher concentrations were involved in the experiments of Rice Johnston and Evering l1 and of Paneth Hofeditz and Wunsch l2 who also increased the life-time to 0.1 sec. by decreasing the rate of the wall reaction. As the rate-determining reaction is heterogeneous i t appears that the velocity of the bimolecular gas phase reaction determined in this way is consider-ably less than that corresponding to unit collision efficiency. Norrish and Porter Faraday SOC. Discussions 1947 2 97. Forsyth Trans. Faruday Soc. 1941 37 312. RQG = k(CH,)' * (8) l1 Rice Johnston and Evering J.Amer. Chem. Soc. 1932 54 3529. la Paneth Hofeditz and Wunsch J. Chem. Soc. 1935 372 GENERAL DISCUSSION The value of kmax is obtained from t* and the concentration the latter being determined from the weight of mirror removed or metal alkyl formed in unit time which can be measured to within 10 yo without great difficulty. The half-life is also a reasonably accurate determination and the well-known difficulties of the method such as mirror poisoning and irreproducible wall conditions can produce only a scatter and not a general trend of this kind. It appears therefore that this discrepancy must be considered seriously unless a reason can be given for doubting the basic assumptions of the mechanism of mirror removal. The bimolecular rate constant must eventually decrease at low pressures when the collision life becomes greater than the lifetime of the collision complex.A possible explanation therefore lies in the lower pressure of the mirror experiments which is about I mm. whereas the pressure in the sector experiments was usually much higher and always above 5 mm. Dr. E. J. Bowen (Oxford) said The importance of the amount of delocalization energy in the products of a dissociation affecting the ease with which it occurs is also shown by work carried out by Miss Rohatgi on the photochemical reaction of anthracene with liquid chlorinated hydrocarbons. The following quantum efficiencies were found : In light absorbed by anthracene the reaction is AX + RC1 -+ ARC1 (derivative of dihydroanthracene).Carbon tetrachloride CC . 0'4 Pentachlorethane CHCl,CCl . . 0.27 Tetrachlorethylene CCl,=CCl . . 0.07 Tetrachlorethane CHCl ,CHCl . 0.018 Trichlorethylene CHC1-CC1 a . . 0.008 Ethylene dichloride CHCl=CHCl . . 0.0076 Chloroform CHCl . 0.0032 High quantum efficiencies appear only where the CCl radical might be involved. Prof. S. W. Benson (California) said It seems rather strange that the Br atoms liberated in the initial step do not undergo addition and hydrogen abstraction reactions with the cyclohexene. I wonder if Dr. Robb has any information on this point. Dr. E. C. Kooyman (Amsterdam) (communicated) It seems likely that the main differences between the mechanisms of the addition of bromotrichloromethane and that of carbon tetrachloride to olefins arise from the far greater reactivity of the bromine atom in CBrCl as compared with that of the chlorine atoms in CCl,.Kharasch and Friedlander la found bromotrichloromethane to react rapidly with styrene at 50°C under the influence of ultra-violet light. However little reaction occurred at 20' C ; at this temperature styrene was found to inhibit the addition of CBrCl to other olefins. These facts were interpreted by Kharasch and Friedlander as resulting from the low reactivity of the benzyl type radical formed by preferential addition of CCl to the styrene double bond : -CCl + CH,=CH-Ph -+ CC1 ,CH,CH-Ph. These data suggest the abstraction of a bromine atom from CBrCl, to be a slow reaction at low temperatures when the attacking radical is re1 atively stable.In detailed analyses of the reaction between cyclohexene CCl and benzoyl peroxide at 78" C Kooyman and Farenhorst l4 found 60-70 ?& of 3 3-dicyclohexenyl on the basis of peroxide decomposed. In our in-vestigation on a-methylenic reactivity l5 the " retardation constant '' for cyclohexene was found to be 11 x I O - ~ ; in view of the value 0 - 2 reported l 3 Kharasch and Friedlander J . Org. Chem. 1949 14 239. l4 Kooyman and Farenhorst Rec. trav. clzim. 1951 (in press). 15 Kooyman this Discussion GENERAL DISCUSSION 225 by Kharas3i and Friedlander and by Kharasch and Sage l6 for the ratio of the addition rates of CC1 radicals to the double bond in cyclohexene and n-octene respectively this suggests the abstraction of a-hydrogen atoms in cyclohexene by trichloromethyl radicals to proceed at a rate which is not much smaller than the rate of addition.Finally attention is called to the experiments of Kharasch and Fried-lander with respect to the overall kinetics of the photochemical addition of bromotrichloromethane to various olefins including cyclohexene at 10-50OC. The authors found the reaction to be zero order with respect to olefin. Olefin determinations were made by direct titration with bromide-bromate solution rather than by dilatometry. On the bases of these arguments it is suggested that the rate constants reported by Melville and his co-workers may contain contributions from cyclohexenyl radicals rather than to apply to the trichloromethyl radicals only. Thus termination might consist both of dimerization of trichloro-methyl radicals and of dimerization of cyclohexenyl radicals.The in-fluence of the latter will be of course more pronounced a t lower halide/ olefin ratios. Prof. H. W. Melville Dr. J. C. Robb and Mr. Tutton (Birminghum) (communicated) The kinetic evidence presented in our paper entirely pre-cludes atatck by trichloro-methyl radicals on the a-methylenic hydrogen atoms to form chloroform and a stable cyclohexenyl radical as suggested by Dr. Kooyman since under all our conditions the rate of reaction is very accurately proportional to (rate of initiation) 4 and also to the concentration of olefin. If such a reaction did occur exclusively as a termination CC1 + CgHlo +CHCl + CBHg . . . . k', the rate of reaction would be given by which is of course not the case.If on the other hand both processes, i.e. reactions given by k' and R were operative then the intensity ex-ponent would lie between 0.5 and I. The fact that i t is so accurately 0.5 supports the original kinetic scheme suggested in our paper. Since the discussion an experiment has been done in which cyclo-hexene in presence of benzoyl peroxide as catalyst with excess of carbon tetrachloride was placed in a sealed dilatometer tube. By irradiating with light of suitable wavelength the photo-catalyzed reaction was studied and in agreement with the work reported by Dr. Kooyman this was shown to be dependent directly on the light intensity thus establishing that under his conditions termination is indeed first order with respect to radical concentration.The rate of reaction under these conditions is very much less than when CC1,Br is used. This then raises the point of the reason for the different mechanism under the two different experimental conditions but further experiments are necessary in order to establish the nature of the differences in the mechanism involved and more precise information is required regarding the velocity constants for all possible steps in the reaction. Prof. H. W. Melville Dr. J. C. Robb and Mr. Tutton (Birmingham) said Since the paper on the reaction of trichloromethyl radicals with cyclohexene was submitted further values have been obtained for the proposed kinetic steis. These are listed below. R (30" C) k (40" C) 1-5 x I O ~ J , k (5oOC) 1-90 x I O ~ I Eovers1l = E - &Es = 4-5 kcal./mole.2kS 2k6 2-5 x IO* I f 1.3 x 1 0 2 1. mole-1 sec.-l 2.0 x 10' 1. mole-1 sec.-l. 16 Kharasch and Sage J . Org. Chern. 1949 14 537. 226 GENERAL DISCUSSION This gives a value of 4 = 5.5 where 2k6 = +V2k4 . 2k5. This low value is in accord with those recorded elsewhere in the dis-cussion by Bateman Gee Morris and Watson. Further attempts to obtain a value for the energy of activation of the termination step by means of experiments conducted using the technique of intermittent illumination at various temperatures has resulted in failure to detect any activation energy for this step. It should be pointed out, however that this technique is not sufficiently sensitive or accurate to measure small activation energies less than 3 or 4 kcal.over the rather small temperature range normally available in experiments of this kind. Dr. M. Magat (Paris) said I would like to mention some additional evidence concerning two points raised by Dr. Kooyman in his paper. First concerning so-called " stable " i.e. non-dimerizable radicals, some caution is advisable. Dr. Chapiro in our laboratory has for instance, observed that the reputedly stable radical I I-diphenyl-2-picrylhydrazyl reacts with double bonds of polymerizable vinyl compounds forming molecules of a molecular weight higher than that of the radical dimer. We are now investigating the kinetics of this process and the nature of the molecules formed. Secondly concerning the reaction of radicals with poly cyclic com-pounds we have investigated the effect of addition of cancerogenic hydro-carbons and their homologues on the rate of thermal polymerization of styrene at 37" C.If the usually present traces of peroxides are destroyed by a preheating under vacuum at this temperature for 2-3 weeks all these cancerogenic compounds slow down the polymerization. It can be shown by persistent fluorescence of the polymer that these compounds do enter the chain. It is remarkable that the absorption and fluorescent spectra are not significantly shifted. The slowing-down efficiency decreases in the order 3 4-benzpyrene, eo-methylcholantrene phenantrene a-methylanthracene I 2 5 6-di-benzanthracene. Chrysene and pyrene have no effect at a1l.l' Dr. K. S . Pitzer (Washington D.C.) said Prof.Kistiakowsky has reported in his paper some exceptionally fine experimental work and the minor point of interpretation which I am about to raise does not in any way detract from the principal results of the work. In the text just after Table IV i t is stated that the presently unknown corrections for anharmonicity in the vibrations would raise the calculated heat capacity of ethane more than that of ethylene. I would urge caution in accepting this estimate as certain although i t may well be correct. While i t is true that ethane has more vibrations in a given frequency range than ethylene i t could be that the effect of anharmonicity in the torsional vibrations overshadows all others at moderate temperatures. The present treatment as a restricted internal rotation accounts for the anharmonicity in the torsional motion in ethane while the more highly restricted torsion of ethylene is presently treated as a harmonic oscillator.Thus i t seems possible that the anharmonicity correction for the torsional motion in ethylene might be large enough to make the total for ethylene exceed that for ethane. The same statements can also be applied to propane and propylene. Prof. E. A. Guggenheim (Reading) (comr~zunicated) It j s clear that the measurements of Kistiakowsky and Nickle on the ethane-ethylene equilibrium are appreciably more accurate and more reliable than any previous measurements. It is particularly satisfactory to notice how well these results agree with the most recent theoretical values. It is perhaps, not entirely without interest to compare the experimental values with theoretical values calculated before the experiments.The following l7 Bodme and Magat Comfit. rend. 1951,232 1657 GENERAL DISCUSSION 227 Table gives a comparison between experimental measurements made in 1942 and here reported and the values calculated according to a formula la published in 1941 based on Kistiakowskfs own experimental value for the heat of hydrogenation a t 82’ C. EQUILIBRIUM VALUES OF @CIQB[2/p,-,Hb IN ATM. 380’ C 4-04 f 0.17 x 10-6 Temperature . 450’ c Calculated 1941 . . 5-2 x 10-4 4’3 x 10-6 Measured 1942 . . 5-13 f 0.13 x 10-4 Dr. L. Bateman (Welwyn Garden City) (commu.nicated) Horrex and Miles’s reference to the bond shortening in dibenzyl requires revision.Cruickshankl9 has re-analyzed Jeffrey’s data and finds that they are actually indicative of only half the contraction quoted. Moreover, recent isomerization equilibrium measurements by Dr. J. I. Cunneen and myself 20 fail completely to reveal any chemical effect (AG differences < 0.1 kcal.) which could be attributed to bond shortening in I 5-dienes. I suggest therefore that the agreement claimed between the experi-mental and ‘ I calculated ” bond dissociation energies needs critical reconsideration. Three uncertainties are apparent ((i) the correctness of the observed activation energy absolutely and as regards its identity with the energy of primary bond scission ; (ii) the appropriateness of using CH,-CH, rather than say CH . CH,-CH . CH as the “ non-resonance ” refer-ence compound for P h .CH,-CH,. P h ; (iii) the magnitude of the resonance energy of the benzyl radical. Horrex and Miles present con-vincing evidence for the approximate validity of their conclusions con-cerning (i) and this receives support from Bolland and Orr’s 21 investigation into the analogous breakdown of aliphatic I 5-dienes. I question, however whether f I kcal. is a fair estimate of the overall uncertainv in the activation energy in view of the complexity of the decomposition process the analytical intricacies and the spread of the points in Fig. 7. Any error associated with (ii) is undoubtedly small but will be such as to tend to reduce the “ standard ” CC-bond energy (by about 1-2 kcal.). The third factor is thus left as the main source of uncertainty.has derived the value of 24.5 kcal. by subtracting the activation energy for the pyrolytic decomposition of toluene (this being identified with the energy required to break an acyclic CH-bond) from the CH-bond dissoci-ation energy in methane. In fact there is sufficient uncertainty in this derivation for the formerly accepted value of 19 kcal. still to be tenable. Szw-arc reports the activation energy as 77-5 f 1.3 kcal. but statistical regression analysis of all his tabulated data (for uniform surface condi-tions) leads to the result E = 80.0 f 4-2 kcal. (95 yo limit). Further, apart from the small uncertainty in D, in methane (101 f I kcal.) i t would seem more legitimate to compare with an alkane €We where D, in the Me group is undoubtedly lower,2s rather than with the sym-metrical methane molecule.For the present therefore i t would seem preferable to calculate the dissociation energy of dibenzyl simply as (84 - 2 x 19) = 46 kca1.-in satisfactory agreement with the thermochemical and kinetic data-rather than as (85 + 11 - z x 24.5) = 47 kcal. in which a large other-wise unrecognized bond energy term has to be invoked in order to com-pensate for a higher resonance energy which is not definitely specified experimentally. Szwarc l 8 Trans. Favaday Soc. 1940 37 272 l9 Cruickshank Acta Cvyst. 1949 2 65. a1 Bolland and On I.R.I. Tram. 1943 21 133. 22 Szwarc J . Chem. Physics 1948 16 138. J Chem. Physics 1944 IS 479. To be published shortly. Stevenson J . Chem. Physics 1942 10 291 ; Anderson and Van Artsdalen 228 GENERAL DISCUSSION Dr.M. Szwarc (Munchester) (communicated) Dr Bateman's remarks enables me to clear up some details of calculation of activation energy in the pyrolysis of toluene.24 The technique elaborated in this investiga-tion makes i t possible to measure the rate of pyrolysis down to about 0.01 yo. This seems to be the lower limit permissible in this technique, and indeed runs 89 88 87 and 94 (crosses in Fig. 4) show that losses of products start to be noticeable when the limiting case of 0.01 yo is reached. Therefore these four results were ignored in calculating activation energy and were omitted in the graph which was presented in the paper but they were included in the Table giving all the results. The value of 80 kcal./mole obtained by Dr.Bateman results from including these four runs; their omission leads to an activation energy of 77.5 kcal./mole as quoted in the original paper. I take this oppor-tunity to include in Fig. 4 the results of pyrolysis of toluene obtained in 1948 by Dr. J. S . Roberts and in 1949 by Mr. J. Murawski. Pyrolysis of toluene. A Roberts (1948) E = 77'5 kcal./mole. x runs No. 89 88 87 94. Murawski (1949) 0 szxTm2 (1947) FIG. 4. In his contribution Dr. Bateman deals with the possible connection existing between the length and the dissociation energy of the central C-C bond in the molecule of dibenzyl. This problem has been discussed previously,26 and here I would like to clarify further certain points which need additional emphasis.(i) The dissociation energy of a bond linking two atoms depends not only on the nature of the two atoms but i t is also greatly influenced by the molecular environment of the bond in question (e.g. Table I in ref. (2)). (ii) The factors which influence the magnitudes of bond dissociation energies can be divided into two groups those connected with the struc-ture of the undissociated molecule and those due to the nature of radicals produced on dissociation. On the whole i t is not possible to ascertain which factor and to what extent is responsible for the observed changes in bond dissociation energies (see however ref. (3) in which an exceptional case is discussed). (iii) Further treatment of the problem of variations in bond dissoci-ation energies requires therefore the introduction of certain simplifying t4Szwarc J.Chem. Physics 1948 16 128. 25 Szwarc Faraday SOC. Discussions 1947 2 39 GENERAL DISCUSSION 229 assumptions. It has been assumed tentatively 26 that the variations in the C-H bond dissociation energies in molecules of the type R . H are due entirely to the factors arising from the nature of the radical R (i.e. i t has been assumed that these variations reflect the changes in the stabilities of various radicals R). Taking the value of D(CH,-H) as the point of reference it is possible zo to build up a system of experimental resonance energies " for various radicals R defined as D(CH,-H) -D(R-H). (iv) It follows from the examination of heats of formation of the relevant compounds that variations in other bond dissociation energies cannot be accounted for by the experimental resonance energies only.For example, D(CH,-CH,) - D(R-R,) -+ [D(CH,-H) - D(R-H)] + [D(CH,-H) - D(R,-H)]. I would like to emphasize that this inequality arises from thermochemical data only i.e. the value of [D(CH,-H) - D(R-H)] + [D(CH,-H) - D(R,-H)] - [D(CH,-CH,) - D(R-R,)] is independent of the values of dissociation energies used in this expression. (v) I t has been assumed that the above-mentioned value the " strengthening effect ",26 is related to the length of the relevant R-R, bond. For example its value for dibenzyl has been calculated as about 11 kcal./mole and i t has been suggested 2 5 ~ 26 that this value is related to the shortening of the central C-C bond in the molecule of dibenzyl.Since, as has been said above the value of II kcal./mole is independent of the value of the C-H bond dissociation energy in toluene the uncertainty (iij) quoted in the communication by Dr. Bateman is irrelevant €or the problem of the C-C bond shortening in the molecule of dibenzyl. (vi) The relation between the C-H bond dissociation energy in toluene and the C-C bond dissociation energy in dibenzyl is given by the heats of formation of toluene dibenzyl and the H atom i.e. ZD(CeH5. CH2-H) - D(C6H5. CH,-CH,. C6H5) = ZAH,(H) -ZAHf(C6H5 . CH,) +AHI(C6HQ . CH . CH . C6H5), = 106 kcal. /mole. Hence the value of D(C6H5. CH,-H) = 77-5 kcal./mole,28 requires D(C6H5. CH,-CH . C6H,) to be 49 kcal./mole while the value of the latter C-C bond dissociation energy is experimentally estimated by Horrex and Miles 29 as 48 kcal./mole.If D(C,H5 . CH,-H) = 80 kcal./mole as Dr. Bateman suggests then D(C6H5 CH,-CH . C6H5) must be 54 kcal./mole i.e. the values D(C,H . CH,-H) = 80 kcal./mole and D(C6H5 . CH,-CH . C6H5) = 46 kcal./mole suggested by Dr. Bateman are incompatible. Dr. B. G . Gowenlock (Swansea) (communicated) Horrex and Miles postulate the reaction (e), I PhCHCH,Ph -f PhCH=CHPh + H-(128 - Q) kcal., among the reactions resultant upon the production of PhCHCH,Ph radicals by the attack of benzyl radicals upon dibenzyl. This reaction is endothermic to the extent of about 58 kcal. being taken as 70 kcal., and therefore an activation energy of at least 58 kcal. will be required for this reaction. This activation energy is greater than that observed for the primary reaction.From the data given in Table I and assuming a I 26 Szwarc J . Chem. Physics 1950 18 1660. 27 Szwarc and Taylor Trans. Furaduy SOC. (in press). 28 Szwarc J . Chem. Physics 1948 16 128. 29 Horrex and Miles this Discussion 230 GENERAL DISCUSSION normal value for the temperature independent factor ( 1 0 ~ ~ sec.-l) i t can be shown that for the lower temperature data (630-700’ C) decomposition of the PhCHCH,Ph radical will take place to the extent of only 10-50 yo. Similar coksiderations apply to reaction (i). Reaction ( f ) , 2PhCHCH,Ph -+ PhCH=CHPh + PhCH,CH,Ph + (29- 128) kcal., I on the other hand is exothermic and should have a much lower activation energy. In contrast to reaction (e) i t will produce only one molecule of stilbene per two PhCHCH,Ph radicals.Therefore on the basis of the authors’ reaction mkchanism a variation of the toluene/stilbene and toluene/styrene ratios should be obtained with variation in temperature. Investigation of the pyrolysis of mixtures of benzyl iodide and dibenzyl a t about 5 0 0 O C should therefore lead to the virtual elimination of re-actions (e) and (i) together with their dependent reactions (g) ( h ) and ( j ) , and thus lead to toluene and stilbene as the sole reaction products in the ratio of 2/1. Dr. C. Horrex and Dr. S . E. Miles (St. Andrew) (communicated) : In reply to Dr. Gowenlock we wish to point out that our analyses showed no significant variation in the composition of the products with the tem-perature of reaction.When introducing our paper we drew attention to the endothermicities of reactions (e) and (i) and pointed out that these decompositions would have to have normal temperature-independent factors of about 1 0 1 ~ in order to be significant in our conditions. We consider that the uncertainty in the value of Q makes i t unprofitable to pursue such calculations at present since an error of 4 kcal. alters the estimated rate of (e) or (i) at 630’ C by a factor of 10. Since the reactions appear to be feasible and a homogeneous chain sequence improbable we prefer to await the result of further work which one of us is carrying out. It must be noted that although the primary decomposition is not a heterogeneous reaction the experimental conditions favour access to the wall and we cannot at present assess the part i t may play in radical decompositions.The extent of a homogeneous disproportionation re-action which requires the encounter of two radicals is very dependent on the concentrations of the latter and hence on the rate constants of (d) (e) and ( 2 ) . It may be significant to note that the early work which reported stilbene and toluene as main products was done with much higher concentrations than we employed ; this would result in decreased accessibility of the wall to the radicals and reactions (e) or (i) which might occur there would be less prominent. Investigations in progress include work on the reaction of benzyl radicals with other molecules. With reference to the points made by Dr.Bateman we have stated that we evaluated the best straight line by use of the high partial pressure data since in our opinion the analytical precision in the other experiments was impaired by the small amounts of products available to us. We consider this procedure legitimate and a closer examination of our Fig. 7 shows that the spread of points in the data used is small (with the noted exception of one point which by any test must be regarded as a faulty experiment). The point concerning the revision of the length of the central C-C bond in dibenzyl has also been drawn to our attention by Prof. Cox. The amount of the shortening and its relationship to the strength of that link are of importance and in summarizing Szwarc’s arguments we may have given the impression that the strengthening is dependent OR the value for this shortening.Actually any “predicted ” value for the dissociation energy of the central bond in dibenzyl depends on the evalu-ation of the heat of formation of the benzyl radical. This has been done in several independent ways a summary of which has been given recently.30 30 Quart. Rev. 1951 5 42 GENERAL DISCUSSION 23 1 The agreement shown by these methods appears to us to be convincing evidence for the higher value for the resonance energy of the benzyl radical. Dr. W. A. Waters (Oxfovd) said In connection with the paper by Horrex and Miles on the pyrolysis of dibenzyl I should like to draw attention to the rather different conclusions of Dr. A. F. Bickel and myself 31 concerning the free benzyl radical which we prepared by de-composing w w'-azotoluene in boiling decalin solution at about 200' C.We found that even at this low temperature the benzyl radicals did not just recombine to form dibenzyl but underwent about 35 yo dispropor-tionation t o toluene and stilbene. We did not detect either benzene or styrene which may be secondary pyrolysis products of stilbene. Though our results could perhaps be attributed to a very rapid de-hydrogenation of dibenzyl by benzyl radicals i t was significant that the decalin solvent was not dehydrogenated at all. The alternative ex-planation would therefore require both an abnormally high lability of the C-H groups of dibenzyl and a very low probability of recombina-tion of benzyl radicals. Unlike Szwarc we see no fundamental objection to postulating the simple reaction .2Ph. CH . + Ph . CH + Ph . CH. It is interesting also to note that the activation energy for the vapour phase dissociation of dibenzyl accords with the experimental data given by Ziegler 3 2 for liquid phase dissociations of many similar compounds, some of which give radicals that undergo disproportionation at temper-atures as low as 1 5 0 O C. Dr. M. Szwarc (Manchester) (communicated) Dr. Waters suggests the possibility of disproportionation of benzyl radicals i.e. The activation energy of this reaction should certainly be less than 5 kcal./mole if reaction ( I ) is to compete successfully with dimerization. This follows from the approximate equality of A factors for both dis-proportionation and dimerization (since the activated complexes for both reactions are very alike and in the liquid phase the deactivation of " hot " djbenzyl molecules cannot be the rate determining step).The activation energy of a process must be a t least equal to its endothermicity i.e. D(C6H5 . CH-H)-D(C,H . CH2-H) < 5 kcal./mole, D(C6H . CH-H) < 82.5 kcal./mole. On the other hand D(C6H5 . CH,-H) +D(C,H5 . CH-H) +D(C,H,-CH) +D(C-H) Taking D(C6H5. CH,-H) = 77-5 kcal./mole D(C,H5. CH-H) < 82.5 kcal/mole D(C-H) = 80 kcal. ; AH,(H) = 52 kcal./mole ; AH,(C) = L ( L being the heat of sublimation of carbon) and finally AH,(C6H5) ~3.68 kcal. /mole (the latter value being derived from a plausible assumption D(C6H5-H) m IOO kcal./mole) one derives : D(CsH,-CH) > L - 28 kcal./mole.The heat of sublimation of carbon is probably not less than 136 kcal./mole and therefore, D(C6H,-CH) > 108 kcal. /mole. The latter value appears to be much too high. It might be reasonable if the activation energy of (I) is about 25 kcal./mole but then this re-action could not be observed. = 3AHf(H) + f AHf(C6H6) -AHf(C6H5. CH3). 3l Rec. truu. cbim. 1950 69 316. 32 Annalen 1942 551 161 ; 1950 567 134 232 GENERAL DISCUSSION It seems to me that the results of Bickel and Waters could easily be interpreted on similar lines to those of Horrex and Miles. The reason why benzyl radicals do not dehydrogenate decalin seems to be simple : the C-H bond dissociation energy in decalin is higher than and thus the activation energy for dehydrogenation is too high.Dr. C . Eorrex (St. Andrews) (communicated) In reply to Dr. Waters I think the products obtained from the interaction of benzyl radicals depend on the experimental conditions. The conditions in solution and in dilute gaseous systems are distinctly different and even in our work with gaseous systems we have observed differences in the proportions of the products when the only substantial change has been the pressure of the inert carrier gas. Dr. Ruth Lapage when using toluene as an acceptor for methyl radicals at 500 to 600" C in a flow technique with 3 to 8 mm. nitrogen found the main solid product was dibenzyl with a small amount of stilbene. With 600 mm. nitrogen however the product was stilbene. Mr. C. B. Cowan has obtained similar data with phenyl radicals and toluene at 600" to 700" C.In the toluene pyrolysis Mr. J. 0. McCrae and Mr. R. B. Cundall examined the solid products from an experiment a t 850°C when using the pyrolysis technique as published by Dr. Szwarc. By use of u.-v. absorption spectroscopy' they' found the solids contained dibenzyl stilbene and anthracene in the ratio 380 7 I. Thus Dr Szwarc is substantially correct in speaking of dimerization only under his conditions. It was also found that using 1-5 mm. toluene in 5 mm. nitrogen carrier gas gave even smaller amounts of stilbene and anthracene at 850" C. In toluene pyrolysis the benzyl radicals are formed at temperatures where dibenzyl would be rapidly decomposed and dimerization probably occurs at the exit of the furnace.Since the gases are cooled quickly there will be little opportunity' for attack by the radicals on the first fractions of the dibenzyl formed. In the work of Lapage and Cowan the temperatures required should permit dibenzyl to be formed within the furnace and remain unpFolyzed. If stilbene was formed from di-benzyl by radical attack it seems that higher inert gas pressures materially assist in the formation of the dibenzyl. If benzyl radicals disproportionate I would have expected the reaction to be rapid at the temperatures of toluene pyrolysis and more stilbene produced than has been found. Dr. Miles used w a'-azotoluene as well as benzyl iodide to provide benzyl radicals when examining the attack of these radicals on dibenzyl. The products in both cases were the same and similar to those obtained from dibenzyl alone on pyrolysis.I would expect the abstraction of a hydrogen atom from decalin by a benzyl radical to be a distinctly endothermic process and the data given by Dr. Kooyman in his Tables I and I1 support the idea that hydrogen abstraction from dibenzyl is an easier process. With reference to Ziegler's data I would like to add that his tem-perature independent factors are generally high. Prof. S . W. Benson (California) said The data of Horrex and Miles seem to indicate a steric factor which is unusually low for a unimolecular decomposition involving the rupture of a single bond and an activation energy lower than that which would be expected for a C-C bond if the proper length of the bond were employed.In addition the dependence of the rate on the surface/volume ratio shows a strong temperature effect. Thus the ratio of rates in a packed to unpacked flask fall regularly from a factor of 3 at 635O C to about 1-4 at 720° C . The authors make the assumption that the ratio of heterogeneous to homogeneous reaction is proportional to the surface/volume ratio. This is by no means necessarily the case and must certainly depend on the rate and energy of adsorption of reactant on the surface as well as on the overall D(CQH6. CHZ-H) GENERAL DISCUSSION 233 mechanism. In particular the dependence may vary from complete independence to direct proportionality according to the conditions. Further the pressure range investigated (about 0-1 mm. Hg reactant) is that in which wall collisions proceed a t a rate comparable to intermolec-ular collisions.In the light of these uncertainties i t would seem premature to accept with any certainty the value of 48 kcal./mole and the low frequency factor for the unimolecular homogeneous decomposition of dibenzyl. Dr. C. Horrex and Dr. S. E. Miles (St. Andrews) (communicated) : Prof. Benson is in error in stating that in our conditions of 0.1 mm. Hg reactant pressure the wall collisions are about as frequent as intermolec-ular ccl'isions. He appears to be ignoring the presence of the carrier gas which impedes diffusion to the wall considerably. Calculation of the average number of collisions made by a molecule diffusing to the wall 33 gives about 105 for the pressures and reaction vessel used.By considering his comments on surface effects in conjunction with details given in his previous contribution on this topic we note that the effects of changes in surface/volume ratios are difficult to interpret where chains are initiated and end at the walls. In our case we see no reasonable evidence for a chain mechanism and consider the relatively small influence of the large change in surface implies little heterogeneity in the primary reaction. This is not in conflict with the possibility (mentioned in reply to Dr. Gowenlock) that some radicals produced in secondary reactions may reach the walls and undergo reaction there. Prof. Benson's assumption that the expected bond dissociation energy depends on the length assigned to the central C-C bond has been dealt with above.We agree of course, that the temperature independent factor is unusually low in our experi-mental conditions. It seems to us that it is desirable to have present an acceptor other than dibenzyl for the removal of the benzyl radicals ; data on the thermal stability of potentially suitab!e substances is being completed. In view of findings on the pyrolysis of methyl iodide in this pressure range given later we are also checking to see if the high pressure limiting rate has not been attained. Dr. C. Horrex and Dr. Ruth Lapage (St. Andrews) (communicated) : Some points disclosed by a detailed examination of the pyrolysis of methyl iodide seem to be of interest in connection with the kinetic difficulties in this pyrolysis work a t low pressures and particularly since C-H and other bond dissociation energies have been derived from published work on iodides.Butler and Polanyi 34 report five experiments on the pyrolysis of CH31 at 493'-495" C and by calculating first order constants and using log k (sec.-l) = 13 - E/4*57T they derived E = 54 kcal. With this value assigned to the C-I bond dissociation energy they deduced a C-H value €or CH which proved in agreement with later determinations by other methods. We have used the same techniques and conditions and found that their quoted values of k can only be realized a t temper-atures about rooo C higher. The use of the correct temperature in the above equation would give a value of E of about 60 kcal. and this seems too high for the C-I dissociation energy.The validity' of assuming first-order behaviour and a 1013 factor under such experimental con-ditions was therefore examined. With 0.07 to 0.3 mm. CH,I in 3 mm. nitrogen carrier gas the kinetics approached second order and changed towards first order above I mm. of CH31. In these conditions the re-action was mainly' homogeneous and not inhibited by iodine the free methyl radicals could be trapped by toluene but in its absence they were decomposed quantitatively' a t the wall to methane and carbon. The rate of reaction was increased by increase in pressure of the carrier gas. s3 Bursian and Sorokin 2. $lay&. Chew. B 1939 12 247. ,* Butler and Polanyi Trans. Faraday SOC. 1943 39 19 234 GENERAL DISCUSSION By using 600 mm. Nz z mm. toluene and 1-2 mm.CH,I we obtained rate constants showing a temperature variation given by The strong influence of experimental conditions on the rates of decomposi-tion can be seen from the following values of k calculated on a first order basis. Conditions Goo mm. N g mm. N 3 mm. N 3 mm. N, k (sec.-l) = 2.6 x 1olS exp -547oo/RT. z mm. C,H 1-1.3 mm. 1-1-3 mm. 0.3 mm. 1-2 mm. CH,I CH31 CH,I CH,I K (sec.-l) at 0.028 0.0066 0.0056 0.0016 Soo0 K The presence of a radical acceptor was essential in experiments at high total pressures since with the decreased accessibility of the wall to the radicals a recombination reaction became important. An analvsis of our data has beeimade which considers activation of the CHJ molecules as occur-ring by collisions with CH31 or N mole-cules the latter having about one-fifth the efficiency of the former.Applying the con-ventional test for data on second- to first-A order transitions of plotting the reciprocal of a first-order constant against the recip-rocal of the concentrations (appropriately weighted) we obtain Fig. 5 The experi-mental point on the vertical axis clearly represents the high pressure limiting rate for the reaction. Details of this work are being prepared for publication. These findings while not altering the C-I bond dissociation energy make the initial derivation of the 54 kcal. value seem &A rather fortuitous. Similar troubles may exist with other substances in the Butler-A Polanyi conditions although one might expect the fall off in first-order constants for more comdicated molecules to occur a t A A na 53 rp A A A M~ I + 0.2 MM.1,) - lower pressures than for CHJ. A number /;o 2.0 of C-H bond energy values which rest on the original iodide pyrolysis determinations may' be uncertain for these reasons. Dr. Horrex and J. 0. McCrae (St. Andrews) said Dr. Kooyman may be interested in our preliminary results on the pyrolysis of diphenyl-methane which show that the products are essentially tetraphenylethane, fluorene and hydrogen. The technique of investigation has been similar to that used for dibenzyl (see this Discussion). We assume the primary process is the breaking of the methylenic C-H link and that fluorene production may proceed through the planar diphenylmethyl radical radical. The calculation of the rate of this primary process via the hydrogen production (allowing for fluorene formation) yields a rate constant k = 10'~ exp (- 73ooo/RT).We are extending the range of experimental variables but suggest that this energy of activation can be considered as the C-H bond dissociation energy Prof. 6. W . Benson (California) said There seems to be some question in the work of Rowley and Steiner of the role played by each of the three possible simultaneous reactions cis plus cis ; cis plus trans and trans plus trans (butadiene). hobs = a2k + z a ( ~ - a)kd + (I - a)'LKtl, where a is the fraction of the cis form and K, K,t and Kg are the specific rate constants for the individual reactions of the species indicated. If i t FIG. 5. W e can write a general rate expression GENERAL DISCUSSION 23 5 is assumed that the active complex is the same for all of these cases then it can be easily shown that the observed activation energy is not dependent on the extent to which the different species react since a = e -AE/RT ; E, = Ell - zAE = E,# - AE and the temperature dependence of the 01 exactly cancels the different temperature dependence o f the different rate constants.In this case the above expression reduces to If however there is more than one possible complex for the system and the steric factors are different for the different isomers then we should expect a different contribution at different temperatures from the different species. Thus i t may be quite possible that a cyclic and linear complex both exist having different activation energies and different frequency factors. In this case the system becomes much more complicated and alternative explanations must be considered for interpreting the data. Dr. M. Magat (Paris) said There is one point that strikes me as not being taken enough into account in all theoretical discussion of butadiene cyclization reaction. It is the fact pointed out by Aston Szasz Wooley and Brickwedde 35 that the trans configuration is 2.5 kcal./mole more stable than the cis configuration. On the other hand at least one of the two molecules must have the cis configuration in order to make the cyclic complex possible while the linear complex can be realized with any one of the two. Mr. B. Eisler and Dr. A. Wasserman (London) (communicated) : Experiments have been carried out which enable a comparison to be made of the activation energies E of Diels-Alder diene associations of the following type : kobs = ktt = a'ku. Butadiene + dienophile -+ adduct . * (1) Cyclopentadiene + dienophile + adduct . * (11) The difference between E and EII can be accounted for,ss if it is assumed that the formation of the transition state of (I) involves the conversion of trans-butadiene into the cis form. In the reaction discussed by Steiner the dienophile of (I) is butadiene. A comparison of the E values of (Ia) and (IIa) indicates that both butadiene molecules are converted into the cis form before the transition state of the reaction (Ia) is taken up. 2 Cyclopentadiene -+ dicyclopentadiene . . (IIa) This is due firstly to steric requirements and secondly to effects which operate generally in Diels-Alder reactions and which bring about the formation of relatively closely packed transition Dr. H. Steiner (Petrocarbon Ltd. Manchester) (communicated) I n reply to the remarks by Benson Magat and Wasserman a statistical rate calculation of the butadiene dimerization using the values for the thermodynamic functions of butadiene obtained by Aston Szasz Woolley and Brickwedde,s* may give an answer to some of the questions raised. Since the interconversion of trans- into cis-butadiene is endothermic by only 2-5 kcal. /mole whereas the activation energy of the dimerization is some 25 kcal./mole the initial state which has to be assumed in such a calculation is most likely' an equilibrium mixture of trans- and cis-butadiene at the appropriate temperature. It may be that the deviations from the Arrhenius function which we have observed are connected with the gradual shift of the cis-trans-butadiene equilibrium particularly at low temperatures. Unfortunately our statistical calculations were carried out using approximate values for the partition function of buta-diene only. I hope to recalculate this prablem using the data of Aston and collaborators. 2 Butadiene -+ vinylcyclohexene . (14 85 J . Chem. Physics 1946 14 67. 36 Detailed considerations and experimental results to be reported elsewhere. 37 J . Chem. Soc. 1935. 833 1512a8 1936 432 ; Trans. Faraday SOC. 1936 35 841. J Chem. Physics 1946 14 67

ISSN:0366-9033    

DOI:10.1039/DF9511000213

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

11 HYDROCARBON REACTIONS B. OXIDATION REACTIONS THE REACTIONS OF RADICALS FROM ACETONE WITH OXYGEN BY FRANK B. MARCOTTE AND W. ALBERT NOYES, JR. Received I 8th January, 195 I The reactions of radicals formed during photochemical decomposition of acetone with oxygen have been discussed. At temperatures up to 2ooOC the results are relatively simple, and a mechanism to explain them is advanced. The activation energy for reaction of methyl radicals with oxygen is low and about the same as for the combination of methyl radicals to give ethane. The steric factor is low, probably about I O - ~ to 10-5. Either carbon monoxide or carbon dioxide results indirectly from the re- action of methyl radicals with oxygen. It is suggested that HCO is an inter- mediate. A tentative activation energy f a - dissociation of HCO is given.The primary process during photolysis of acetone a t 25' C is discussed. The mechanism of the photochemical decomposition of acetone has received much attention, and the principal steps seem to be well sup- ported by the evidence.l Remaining doubts centre chiefly about the details of the primary process.2 The radicals produced photochemically in acetone are CH, and CH,CO. When methane is produced by the abstraction of hydrogen atoms from acetone by methyl radicals, CH,COCH, must also be formed.8 Hence a study of the photochemical behaviour of mixtures of acetone and oxygen may be expected to vield valuable information both about the reactions of these radicals with oxygen and also, possibly, about the behaviour of any excited molecules which may be produced during the primary process.Experimental The light source was a Hanovia Alpine burner collimated by a quartz lens and filtered through I mm. of Pyrex glass. The effective wavelengths were 3130 and 3020 A since longer wavelengths are not absorbed by acetone vapour, and shorter wavelengths are transmitted very little by Pyrex. After exposure of acetone-oxygen mixtures, condensation was effected by liquid nitrogen ( - 196' C) and residual gases (02, CO, CH,, traces of C,H, if any) were removed by a Toepler pump. These were passed into a mixture of copper and copper oxide in a furnace at a temperature of about 220° C. The oxygen is removed as CuO, and the carbon monoxide is oxidized to carbon dioxide by this treatment. Ethane was never found in sufficient quantity to be determined quantitatively when oxygen was present. Carbon dioxide was separated from the methane by condensation with liquid nitrogen.Reference should be made to Noyes, Jr., and Dorfman, J . Chem. Physics, 1948, 16, 788, for a discussion of this problem. An extensive review of earlier work has been published by Davis, Jr., Chem. Rev., 1947, 40, 201, so that the literature references in the present article will be limited t o those pertinent to the discussion. See Noyes, Jr., J . Physic. Chem., 1951, 54, 000. A more recent study of fluorescence is in press, Luckey and Noyes, Jr., J . Chem. Physics, 1951, 19, 227. Rice, Rodowskas and Lewis, J . Amer. Chem. SOC., 1934, 56, 2497 ; Allen, J . -4~ner. Chem. SOC., 1941, 63, 708.The methane is not affected. 236F. B. MARCOTTE AND W. A. NOYES, JR. 23 7 Oxygen Disappearance The analyses were, therefore, for the following : oxygen consumption, methane, carbon monoxide, ethane, carbon dioxide. The latter was removed a t - 150' C. The methane fraction never showed an indication of a measurable amount of ethane. At temperatures from 120' to 200' C i t was assumed that one molecule of carbon monoxide is formed per quantum absorbed by oxygen-free acetone.4* At 25OC a value of 0-1 for this quantum yield was used,5 and admittedly quantum yields at this tem- perature are less accurate than a t the higher temperatures. Since i t was desired t o obtain quantum yields relative to quantum yields of methane formation, and since no methane is formed except a t low oxygen pressures, it was necessary t o replenish the oxygen to prevent its complete disappearance during an experiment.This was accomplished by a doser which was essentially a small Toepler pump which added a definite amount of oxygen when desired. Thus the oxygen pressure fluctuated to some extent during a run, but once the approximate rate had been determined, the oxygen could be added so as to maintain its pressure nearly constant. The oxygen pressures given in the table represent mean values about which there were fluctuations of about 10 yo a t low oxygen pressures and less a t high. Since oxygen would be depleted most rapidly in the illuminated zone, it was necessary to ensure thorough and continuous mixing. This was accomplished by inserting a small glass propeller with a piece of steel sealed into the top.The propeller was rotated by a permanent magnet attached to the shaft of a motor. The acetone was used as an internal actinometer. Thus no metal surface was in contact with the gases. co + co, Results The results are divided into those obtained at temperatures from 120' to 2ooOC as shown in Table I, and those obtained a t lower temperatures, mostly a t 25OC, as shown in Table 11. The phenomena appear t o be quite different under these two temperature conditions, and this is due mainly t o the nature of the primary process in acetone. TABLE I.-AVERAGE QUANTUM YIELDS AT 120-200' C I20 150 I75 200 4-0 f 0-2 4-0 f 0.2 3'9 f 0'2 3'9 f 0'4 - 3-0 f 0.1 3.2 f 0.2 3.1 & 0-2 - I- - 3'5 3'9 3'6 (Note.-Values for CO + CO, as well as for oxygen disappearance are low either a t very low or very high oxygen pressures. The averages are for oxygen pressures from about 30 to about 300 x I O - ~ mm.The values for CO + CO, + CH, have little quantitative significance except t o indicate that the products can be accounted for only if more than one acetone molecule disappears per quantum absorbed.) The salient facts which may be obtained from the data a t 120' to 200° C are as follows : (i) No ethane is formed even when pressures as low as 0.01 mm. of oxygen are present. (ii) Methane is formed only a t low oxygen pressures, but the oxygen pressure necessary to prevent methane formation increases markedly as the temperature increases. (iii) The quantum yield of oxygen consumption is approximately four and varies extraordinarily little with experimental conditions, although departures from this figure greater than experimental error are noted.Leermakers, J . Amer. Chem. SOC., 1934, 56, 1899. Herr and Noyes, Jr., J . Amer. Chem. SOC., 1940, 62, 2052.238 RADICALS FROM ACETONE and passes through a maximum of about three. pressures i t may fall below unity. falls, and the sum of these two yields is very roughly constant a t three. (iv) The carbon monoxide quantum yield is a function of oxygen pressure At sufficiently high oxygen (v) The carbon dioxide quantum yield rises as that of carbon monoxide At 25' C the following facts are apparent : (i) The carbon monoxide yield is lower in the presence than in the absence of oxygen but is nevertheless nearly constant as the oxygen pressure varies.(ii) The carbon dioxide yield is higher than the carbon monoxide yield. (iii) Neither methane nor ethane is formed if oxygen pressure exceeds 0.02 mm. (the lowest pressure studied). (iv) The number of molecules of oxygen used per quantum is much less than unity and is approximately constant and independent of the oxygen pressure. Acetone Pressure (m.1 131'3 - - - Acetone Pressure (m.1 Oxygen Pressure 1 co Oxygen Disappearance (mm. x Id) - 20-4 0.56 0.044 50'4 0.48 0.048 123 0.58 0-040 245 0.40 0.032 TABLE II.-QUANTUM YIELDS AT 2ooOC Oxygen Pressure (mm. x 103) 43 38 53 74 98 165 511 1064 2072 Oxygen Disappearance 4'5 3'4 3'4 4'1 3'8 4'7 3'7 (3'1) . (2'2) Quantum Yields co -- 3'3 2.6 2'4 2'5 2'4 1'4 1-05 2'2 1-2 COa - 0.46 0.59 0.80 0.91 0.80 1'3 1-7 1-8 0'74 0.80 0.59 0'43 0.28 0.23 <0-05 0-00 0.00 Additional results are being obtained on the reactions in the system diethyl ketone + oxygen.These cannot be reported in detail a t present, but a t room temperature the carbon monoxide yield is lower in the presence than in the absence of oxygen. Thus neither ethyl nor propionyl radicals appear, upon reaction with oxygen at room temperature, to give carbon monoxide. The oxygen reacts sufficiently rapidly with propionyl radicals to prevent their dissociation into carbon monoxide and ethyl radicak6 The results will not be presented in detail, but parts important to the dis- cussion will be given. The trends in the data are illustrated in Table I1 which gives results a t 203O c.The data obtained a t 25O C are given in Table 111. Table I gives a summary of certain values. TABLE III.-QUANTUM YIELDS AT 25OC (Gc0 is assumed to be 0-1 in the absence of oxygen) I I Quantum Yields .- 0.20 0.18 0.26 0.16 (Note.-No methane and ethane were formed in any of these experiments.) 6 These results will be reported in detail later by Dr. A. Finkelstein.F. B. MARCOTE AND W. A. NOYES, JR. 239 Discussion A minimum yield of one molecule of carbon monoxide per quantum would be expected if the primary process has an efficiency of roo yo and if the acetyl radical always dissociates. These conditions seem to be met at 150, 175, and zooo C, but at high oxygen pressures at 120' C the acetyl radical does not appear to dissociate completely, and at 25' C neither condition is met.At zooo C the carbon monoxide yield approaches unity asymptotically with increase in oxygen pressure. The maximum yield of about three for carbon monoxide implies that reaction of methyl radicals with oxygen yields carbon monoxide under conditions of low oxygen pressure. However, as the oxygen pressure increases, carbon dioxide increases, and carbon monoxide decreases, although the sum of the two is approximately constant. Methane formation does not occur at the expense of carbon monoxide and carbon dioxide, but the last column in Table I indicates that too many carbon atoms are found among the products in such a case to be due to decomposition of a single molecule of acetone. Any mechanism for these reactions must account, therefore, for a reaction of methyl radicals with oxygen which may by alternative paths lead either to carbon monoxide or to carbon dioxide.This suggests an intermediate, such as HCO, which may either dissociate or react with oxygen. The quantum yield of oxygen disappearance of about four indicates that the product of the reaction of a methyl radical with oxygen also consumes one molecule of oxygen but that the product of this second reaction does not consume oxygen. It is possible to suggest a mechanism which will explain these facts, and it is given by the following equations : Reaction Quantum Yield CH3COCH3 +hv=CH3 + COCH, 4 (1) COCH,=CO +CH3 a+ (2) COCH +O, = CO, + (CH ,O) ( 1 - 4 4 (3) CH, +CH3COCH3=CH, +CH,COCH, ( 1 + 4rS4 (4) CH,COCH, +0,=CH3C02H +HCO P+a)fl+ ( 5 ) CH3 3-0, = HCO +H,O ( 1 + 4 ( I - - B ) + (6) HCO=H+CO [ ( I + 4 ( I -B)4+(1+4/34IY (7) HCO +Oo= CO, +OH (8) H +02 =HO, (9) The fates of OH, CH30 and HO, are not specified.None of these would be expected to react with oxygen, but the first two might form H,O and CH30H to some extent by hydrogen abstraction. Some slight reaction of this type may account for CO + COO yields slightly greater than three in a few experiments. Peroxides would probably react with mercury during the analysis, and this is suggested as the final fate of the products of OH and HO,. [(I +a) (1 -PI$+ ( 1 + a o B # l ( I - Y ) [( 1 + 4 ( 1 -B) 4 + ( 1 + 4B41 Y Hence @(lo = a$ + (1 + .)Y$. ' (10) A maximum value of @go = 3 could be obtained only if t$ = I (primary quantum yield}, a = I (fraction of acetyl radicals which dissociate), and y = I (fraction of formyl radicals which dissociate).This would be obtained at low oxygen pressures, but it should be noted that does not depend on /3, the fraction of methyl radicals which form methane. The evidence is quite strong that # = I in the absence of ~xygen,~. and the results lend strong supporting evidence to the statement that oxygen does not inhibit the primary dissociation. Also @co, = 24 - (1 + a)r+, * (11) @GO + @C@ = (2 + a>$. ' ( 1 4 which will attain a maximum value of 2 only if $ = I and Y = 0. This will have the value 3 if 4 = I, a = I. The data are not extensive240 Temp. OC I20 150 I75 200 RADICALS FROM ACETONE (16) makes the second assump- ks lk4 tion but not the first.Table IV shows average values of k , / k , at four temperatures. A plot of the logarithm of the ratio in Table IV against I/T gives a reasonably good 4-6(f 1.1) x 104 2-0(fo.02) x 104 I - I ( ~ 0.2) x 104 5.8(f 0.8) x 103 enough at IZOO C to establish this conclusion, but at 150, 175 and zooo C the conclusion appears warranted. This will have the value of 4 if a = I , + = I . These conclusions seem warranted unless the oxygen pressure is several hundred microns or more. It is of interest now to calculate the relative rates of reactions (6) and (4). 002 = (3 + a)+. * - (13) This may be done from the following relationship : (1 + a)B+ k4(CH3COCH3) /3 . ' * (14) (1 + - rw - 'B(O2) = 1 - 8 - The evaluation of /3 in terms of quantum yields may be obtained from the following combined with (13) @C& = (1 + 4 8 4 - (15) p = - @CHa Hence and eqn.(14) becomes 0 0 , - 24'F. B. MARCOTTE AND W. A. NOYES, J R . other reactions occur also. Let us assume tentatively that all of the carbon monoxide comes from dissociation of CH,CO and that all carbon dioxide comes from reaction of that radical with oxygen. Under such conditions The number of methyl radicals formed would be 2 CO + CO, = 0.28. If each methyl and each acetyl consumed one molecule of oxygen, the quantum yield of oxygen disappearance should be 0.28 + 0.20 = 0.48 which is well within experimental error of the average value found of 0-5 I. The nature of the products formed by reaction of methyl radicals with oxygen molecules at room temperature is not certain. Bates and Spence 9 suggest whereas Blaedel, Ogg, and Leighton l o suggest Reaction (20) requires two methyls per oxygen, and on this assumption the quantum yield of oxygen disappearance would be 0.38. Thus while the data do not lead to a definite conclusion, the disappearance of one 0, per radical is favoured over two.If the correct primary yield at these low oxygen pressures is 0.24, the question arises as to whether it is the same in the absence of oxygen. Oxygen quenches the blue fluorescence of acetone quite strongly., Activ- ated molecules mainly dissociate at temperatures over 100" C leaving a residual fluorescence similar to that found at 25°C in the presence of several mm. of oxygen. Thus oxygen removes by some process many of the activated molecules which could fluoresce.If this removal is a true deactivation rather than a predissociation induced by collision, the primary process might have a considerably higher e'fficiency in the absence than in the presence of oxygen. Finally, certain conclusions about the formyl radical are possible if the mechanism for the reaction is correct. By reference to the mechanism it can be shown that 4 = CO + CO, - 0.24. . - (18) CH, + 0, = H2C0 + OH, . - (19) CH302 + CH, = CH,OH + CH,O. * (20) * (21) 2 - @GO* @co, --- - - * k,(HCO) MHCOW*) - 1 - Y @co - 1 =- @co* ' * where (21) is valid if a = I whereas (22) is valid only if a = 4 = I. This condition is apparently valid at 175" C and at 200°C except, possibly, at high oxygen pressures. Unfortunately the data are not extensive enough to permit adequate test of (21) and (m), but values have been calculated at 175°C and at zoo" C. These give k,/kB (av.) = 1-6 f 0.5 x 1015 at 1 7 5 ~ C and 3-1 & 1-2 x 1015 at 200" C in molecules x cm.-3. Hence E, - E, - 12 kcal. ; but this figure is purely tentative. Further information about this matter should be obtained. Values for E, in the literature show little consistency.ll The authors wish to express their appreciation to the Office of Naval Research, United States Navy, for support of this work through Contract N6onr-241, Task I, with the Department of Chemistry, University of Rochester. We are also indebted to several staff members and graduate students for helpful discussions. University of Rochester, New Ydrk. Bates and Spence, J . Amer. Chem. SOC., 1931, 53, 381, 1689. lo Blaedel, Ogg and Leighton. J . Amsr. Chem. SOC., 1942, 64, 2499. l1 Cf. Steacie, Atomic and Free Radical Reactions (Reinhold Publishing Cor- poration, New York, 1946), p. 362.

ISSN:0366-9033    

DOI:10.1039/DF9511000236

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

REACTIONS OF FREE RADICALS ASSOCIATED WITH LOW TEMPERATURE OXIDATION OF PARAFFINS BY E. R. BELL, J. H. RALEY, F. F. RUST, F. H. SEUBOLD AND W. E. VAUGHAN Received 5th February, 1951 An attem2t has b-en made to elucidate the elementary, free radical reactions which oscur in the low temperature oxidation of paraffin hydrocarbons. The problem has b-en approached through a study of the reactions of radical and mdecular sp3cies which are known or postulated to be intermediates in the oxidations. Foi this purpose a variety of aliphatic peroxides has been used both to provide the free radicals desired and to allow examination of the be- haviours of the peroxides themselves. An oxidation mechanism which describes the reaction sequence from the initial union of an alkyl radical and oxygen t o the formation of the primary stable molecular species has been postulated, and experimental evidence for each step has been obtained.Of particular import- ance in this scheme are the processes by which alkylperoxy radicals are con- verted to alkoxy radicals which are the immediate precursors of the primary stable oxidation products. The detailed mechanism of the low temperature oxidation of paraffin hydrocarbons has been a matter of speculation for many years.l With this material as a basis, we have attacked the problem by a study of the reactions of both known and postulated peroxidic intermediates. From this study a self-consistent mechanism has been evolved, and each step has been verified experimentally. The major contribution of this work is an explanation of the processes involved in the conversion of alkyl- peroxy radicals, through alkoxy radicals, to the primary stable products of the oxidation.The further transformations of these primary products may well take place in an analogous manner, but these possibilities are not within the scope of the present study. This picture of the low temperature oxidation process is illustrated in Fig. IA and IB, each step of which will be discussed in more detail in later sections of this paper. In summary, however, alkyl radicals pro- duced in the initiation process (Step I) react with oxygen quantitatively to yield alkylperoxy radicals (Step 2). These may then be converted to alkoxy radicals by (i) abstraction of a hydrogen atom to form a hydro- peroxide (Step 3A), followed by decomposition to alkoxy and hydroxyl (Step 4A) ; (ii) reaction with one another to yield alkoxy radicals and oxygen directly (Step 3B) ; or (iii) reaction with an alkyl radical (possible only a t very low oxygen concentrations) to yield a dialkyl peroxide (Step 3C) which will decompose to alkoxy radicals (Step 4C).Attack on a hydroperoxide to regenerate an alkylperoxy radical (step 3A’) has also been demonstrated. The alkoxy radicals are the immediate pre- cursors of the first stable oxidation products. By decomposition or free radical attack aldehydes or ketones are formed ; by hydrogen atom abstraction alcohols are produced ; and by association with alkyl radicals ethers may be formed in small amounts. It should be emphasized that the conclusions have been based on a study of the reactions of postulated intermediates, rather than on 1 For pertinent references see Lewis and Elbe, Combustion, Flames and Ex- plosions of Gases (Cambridge, 1938) ; Jost, Explosion and Combustion Processes in Gases (McGraw-Hill, 1946) ; Faraduy SOG.Discussions, 1946. 242BELL, RALEY, RUST, SEUBOLD AND VAUGHAN 243 observation of actual hydrocarbon oxidation, and demonstrate the value of this method of approach to the solution of mechanistic problems. Step 2. Formation of Alkylperoxy Radicals by the Reaction of Alkyl Radicals with Oxygen.-The present investigation has been coil- cerned solely with the propagation steps of the oxidation chain and no attempt has been made to specify the mode of initiation. It is assumed that the reaction begins by the production of an alkyl free radical by some means (Step I ) , and the union of this radical with molecular oxygen is the step which will introduce this discussion.The combination of an alkyl radical and oxygen to form an alkyl- peroxy radical is now generally accepted. Some information on the rate of this bimolecular process for the methyl radical is available from earlier studies, and additional data are reported below. A recalculation of data on the photo-oxidation of methyl iodide 3 has shown that methyl radicals combine with oxygen at about 1/8ooth of the rate at which they react with iodine. Also, the activation e n e r a for the oxygen combination has been estimated as I or z k ~ a l . ~ In the pressnt study, observations on the pyrolysis of di-tert.-butyl peroxide, a source of free methyl, in the presence of oxygen afford a comparison with the rates of formation of ethane and methane, the normal pyrolysis p r o d ~ c t s .~ As shown in Table I, in all experiments in which the available oxygen was not ex- hausted, hydrocarbon formation after the oxygen was admitted was vanishingly small. In view of this suppression, it is concluded that the only significant reaction of methyl under these conditions is a union with oxygen. CH,O, + CH, + CH,OH + H,CO is Considered less likely since, to conform to the data, a very high specific rate, relative to other methyl radical reactions, would have to be ascribed to this process. TABLE I.-EFFECT OF OXYGEN ON C2H, AND CH, FORMATION IN THE A possible alternative, the occurrence of the reaction DECOMPOSITION OF DI-tett.-BUTYL PEROXIDE Oxygen remaining .- Total CH, found as CH, + C,H, Total CH, released * . . CH, released before O2 added * . . 1 5 9 - 8 O 147.20 122.1° (m.1 (mm.1 (mm.) 34 I7 I 2 I 0 340 310 238 284 81 28 19 I 84 24 I9 2 < 0'2 251 I 22 Step 3. Formation of Alkyl Hydroperoxides by the Reaction of Alkylperoxy Radicals with Hydrogen Atom Donors .-This reaction is well known from experiments carried out in both the liquid and gas phases. In the former category typical examples are the oxidations of tetralin 6 and cumene,' in which the hydrocarbon itself serves as the hydrogen atom donor : Blaedel, Ogg and Leighton, J. Amer. Chern. SOC., 1942, 64, 2500. Bates and Spence, J. Amer. Chena. SOC., 1931, 53, 1689.Van Tiggelen, Ann. Mines Belg., 1941, 43, 117-44. 6 Raley, Porter, Rust and Vaughan, J. Amer. Chem. SOC., 1951, 73, 15. 6 Hartmann and Seiberth, Helv. cAim. Acta, 1932, 15, 1390. 7 Hock and Lang, Ber., 1944, 77, 257.244 OXIDATION O F PARAFFINS The oxidation of isobutane,s in which hydrogen bromide performs the function of donor, is a pertinent example of hydroperoxide formation in the gas phase : CH3 CH, 1 I I CH, I CH, CH,COO* + HBr -+ CH,COOH + Br' In the three cases cited the hydroperoxides may be isolated in nearly quantitative yield. (Part I of z Parts) X H i STEP 1 8. STEP4A STEP 2 R a R02H + R- a SfEP4C 0 2 32 1 Ror STEP 38 R02. t ' + 0 2 2 ROD ROD FIG. IB.-A generalized mechanism of the low temperature oxidation of paraffin Formation of Alkylperoxy Radicals by the Reaction of Alkyl Hydroperoxides with Free Radicals.-Although the abstraction of the hydrogen atom bonded to oxygen in a hydroperoxide has been postulated in the literature,@ the reaction was first verified in the vapour phase reaction of tert.-butyl deuteroperoxide with di-tert.-butyl and di- tert.-amyl peroxides, carried out in a flow system at 1 g 5 O .l ~ Methyl and ethyl radicals, derived respectively from the two dialkyl peroxides, attacked the deuteroperoxide to yield deuteromethane and deutero- ethane : CH, . (or C,H,-) + (CH,) ,COOD -+ CH,D (or C,H,D) + (CH,) ,COO. As will be explained in a subsequent section, methoxy and ethoxy radicals were also produced in the respective rcactions. Methanol4 was isolated from the reaction of di-tert.-butyl peroxide with tert.-butyl deuteroperoxide, hydrocarbons.Step 3A'. * Bell, Dickey, Raley, Rust and Vaughan, Ind. Eng. Chem., 1949, 41, 2597. Robertson and Waters, Trans. Faraduy SOC., 1946, 0, 201-210. l o Seubold, Rust and Vaughan, J . Amer. Chem. SOC., I951,73, 18.BELL, RALEY, RUST, SEUBOLD AND VAUGHAN 245 presumably by the attack of methoxy radicals on the deuteroperoxide : Although it is considered unlikely, exchange in the vapour phase between the deuteroperoxide and methanol cannot be excluded as a source of methanol-d. Formation of Dialkyl Peroxides by the Combination of Alkylperoxy and Alkyl Radicals.-As mentioned in Step 3A’, the re- action of tert.-butyl deuteroperoxide with di-tert.-amyl peroxide in the (Part I1 of z Parts) CH,O’ + (CH,) ,COOD -+ CH,OD + (CH,) ,COO.Step 3C. If RO* is o primoty alkoxy radicol. R‘C H 2 0. STEP R& CH 2~ R’CH~OH+R~ If R0.k o secondory olkoxy rodicol. R:CHO- STEP RltR’CHO R;CHOH+R* R;CO+RH R;CO+ROH Sy y;;P RbGO+Ri R;COH+R FIG. IA.-A generalized mechanism of the low temperature oxidation of paraffin hydrocarbons. gas phase at 195” yielded ethane-d and tert.-butylperoxy radicals. In addition, a small yield of ethyl tert.-butyl peroxide was also obtained. To determine the importance of this process, the same reaction was carried out at I ~ o ” , at which point 65 yo of the input dialkyl peroxide decomposed in the 2-min. residence time. On the assumptions that the rate of decomposition of ethyl tert.-butyl peroxide was the same as that di-tert.-amyl peroxide, and that every successful collision of an ethyl with a tert.-butylperoxy radical yielded ethyl tert.-butyl peroxide, the maximum yield of the mixed peroxide would be 23 yo of the input di- tert.-amyl peroxide.This calculation assumed that residence times of the two dialkyl peroxides were the same. Since, in reality, the ethyl derivative is formed throughout the reactor, the yield should be somewhat higher. Actua.lly, a 30 yo yield was obtained, indicating that the most important mode of disappearance of alkylperoxy radicals in the absence of oxygen was combination with alkyl radicals : (CH,) ,COO* + C&,* -+ (CH,),COOC,H,.246 OXIDATION O F PARAFFINS The decomposition of the mixed dialkyl peroxide has been reported l1 and the data will be summarized subsequently. The general reaction has also been confirmed by the isolation of methyl tert.-butyl peroxide in low yield from the reaction of tert.-butyl hydroper- oxide with di-tert.-butyl peroxide at 195~.The Formation of Alkoxy Radicals by the Interaction of Alkylperoxy Radicals.-In order to simulate more closely the con- ditions of an oxidation, oxygen more than sufficient to react with all alkyl radicals produced was included in the tert. -butyl hydroperoxide + di-tert.-amyl peroxide system, which, as shown in the preceding section, gives in the absence of oxygen a 30 yo yield of ethyl tert.-butyl peroxide at 180'. While those products typical of alkoxy radical reactions, e.g. ethanol, methanol, and minor amounts of acetic and formic acids and carbon monoxide (see Table 111), were isolated in good yields, no trace of either ethyl tevt.-butyl ar diethyl peroxide could be detected.I n addition, no normal products of the reactions of methyl or ethyl radicals, such as methane, ethane, or butane, could be found. From the evidence presented for Step 2, it would appear that all alkyl radicals were converted to peroxy radicals and that the latter reacted according to the following overall equation ROO. + R'OO. -+ RO. + R'O. + 0, to yield the corresponding alkoxy radicals. The same conclusion was drawn from a study of the oxidation of methyl radicals.s Further evidence for this transformation was derived from the decomposition of tert.-butyl hydroperoxide in various solvents.la The pertinent example is the case in which chlorobenzene served as the solvent.It is apparently highly resistant to radical attack and a very rapid chain decomposition of the hydroperoxide ensued at 140°, yielding tert.-butyl alcohol and oxygen almost quantitatively by the mechanism : Step 3B. (CH,) ,COOH -+ (CH,) ,Coo + 'OH (CH,) ,COO + (CH,) ,COOH -4 (CH,) ,CO HO' } 2(CH3) ,COO. + 2(CH3) ,COO + 0,. The mechanistic details of peroxy radical interaction cannot as yet be specified. The Thermal Dissociation of terf.-Butyl Hydroperoxide.- The decomposition of this hydroperoxide has been studied qualitatively in the gas-phase and quantitatively in the liquid phase.12 In the former case it was passed through an open Pyrex reactor with a molar excess of cyclohexene as a hydrogen atom donor at 260'. In addition to the ex- pected products (acetone, tert.-butyl alcohol, methane, methanol, carbon monoxide and water), a 13 yo yield of cyclohexanol (based on decomposed peroxide) was found.The most reasonable process by which this alcohol could be formed is a dissociation of the hydroperoxide at the 0-0 bond followed by addition of the hydroxyl radical to the olefin : Step 4A. (CH,),COOH + (CH,),CO* + HO. The possibility of unimolecular fission of the 0-0 bond was con- firmed by a study of the kinetics of the decomposition in n-octane in the range 150-180'. The reaction proved to be a combination of uni- molecular and chain processes, the observed " first order I' rate constant l1 Rust, Seubold and Vaughan, J . Amer. Chew. SOC., 1950, 72, 338. IZ Unpublished data from this Laboratory.BELL, RALEY, RUST, SEUBOLD AND VAUGHAN 247 increasing markedly with the initial concentration of the peroxide.The values of these constants at the various temperatures and concentrations employed are presented in Table 11. The true values of the first order rate constants were obtained by extrapolation to zero concentration on a TABLE II.-RATE CONSTANTS FOR THE DECOMPOSITION OF tert.-BUTYL HYDROPEROXIDE I N %?-OCTANE T, O K 423'43 423'39 423.15 422'98 433'08 433'19 433.13 433'54 433'01 442'92 442'97 443'14 442.92 453.01 452'96 453'03 422'96 45"96 452'99 Ik 0.04 - 10skI (sec.-1) extrap. 0.8, 2'57 6-99 18.2 plot of hobs against [ROOH],. From these extrapolated ra.te constants the activation energy for the process was calculated to be 39.0 f 0.6 kcsl./mole, in good agreement with the value calculated from thermo- chemical data as follows : z tert.-C,H, + zH + 2 ~ s o - C ~ H ~ , H,O(z) H2 + 0.50, H, + 2H 2 0 + 0, z iso-C4Hl, + 130, + 8C0, + 1oH,O(Z) gH,O(Z) + 8C0, + (tert.-C,H,O), + 11-50, (tert.-C4HgO), -+ z tert.-C,H,O z tert.-C4Hg + 2 0 -+ z tert.-C4Hg0 Thus Do-o (tert.-C,H,O) = 84.0 kcal. /mole tert.-C,H, + 0 + tert.-C,H,O H f O - t O H 0, + 2 0 ~ s o - C ~ H ~ , -+ tert.-C,H, + H 4CO, + 5H,O(Z) --f iso-C4Hl, + 6.50, 5-50, + tert.-C,H,OOH 3 4C02 + 5H,O(Z) tert.-C4HgOOH 3 tert.-C,H,O + OH Thus Do+ (tert.-C4HgOOH) = 38.5 kcal./mole. A EL=, kcal./mole - 172-l~ + 103.2 la - I 368.5 l5 - II7.2l4 + 67'5 + 1279.9 IS + 39'1 l7 - 168.0 - 84.0 - 101-l~ + 86*13 + 117.2 l4 + 6 8 4 ~ 3 ~ ~ - 664.0 13 + 38.5 l3 Baughan and Polanyi, Trans.Faruday Soc., 1943, 39, 19. l4 Gaydon, Dissociation Energies (Chapman and Hall Ltd., London, 1947). l6 Calculated from heat of combustion (Rossini, J . Res. Nut. Bur. Stand., l* Calculated from heat of combustion (Raley, Rust and Vaughan, J . Amer. 1935, 15, 357). Chem. SOC., 1948, 70, 88). 17 See ref. (16).OXIDATION OF PARAFFINS From the experimental activation energy and the various extrapolated rate constants, the entropy of activation for the unimolecular fission of the 0-0 bond is calculated to be 9-7 f 0.3 cal./mole deg. Interpretation of this entropy change in comparison with that observed for di-tert.-butyl peroxide, 14.5 cal. /mole deg.,ls is straightforward. The greater rigidity of the dialkyl peroxide is evident from a study of models, so that a greater increase in entropy on dissociation would be expected.Considerable interaction between the tert.-butyl group and the peroxidic hydrogen atom in the hydroperoxide is indicated, however, by the relatively high value of the entropy term observed for that decomposition. It is much more difficult to explain the coincidence of the energies of activation for the two decompositions when it is considered that the 0-0 bond dissociation energy in hydrogen peroxide is 52 kcal. /mole? Probably at least two opposing factors are operative, one inductive, the other perhaps steric. No definite answer to the problem can be made at this time. In summary, the overall mechanism of the decomposition of tert.- butyl hydroperoxide may be portrayed as : (CH,),COOH -+ (CH,),CO* + HO- (CH,),CO* --f (CH,),CO + CH,.X- + (CH,),COOH + XH -+- (CH,),COO. Re + (CH,),COO- --f (CH,),COOR (CH,) ,COOR -+ (CH,) ,CO- + ROO 2(CH3),COO* + 2(CH3),C0. + 0, wherein X = HO-, RO-, (CH,),CO*, CH,. and R = CH,* or a radical derived from the solvent. Further transformations of the alkoxy radicals will be discussed in a subsequent section. The Formation of Alkoxy Radicals by the Thermal De- composition of Dialkyl Peroxides .-This reaction was early proposed 2 o but unequivocal proof awaited the study of the decomposition of di- tert.-butyl peroxide in various solvents capable of donating a hydrogen atom. At 1 2 5 O , for example, in cumene,21 tri-n-butylaminc,21 and benzal- dehyde,22 tert.-butyl alcohol is formed in yields of 80, 95 and IOO yo re- spectively, the remainder being accounted for as acetone. (CH,),CO- + RH + (CH,) ,COH + Re.The radicals formed by attack on the solvent produced a dimer in the case of cumene and gave more complex association products in the other two solvents. Step 5. The Reactions of Alkoxy Radicals.-A study of the vapour phase decomposition of a series of alkyl tert.-butyl peroxides at 195' in the presence of a molar excess of cyclohexene as a hydrogen atom donor provided a means of comparing the stabilities of the alkoxy radicals generated.1' The results of this survey are summarized in Table 111, which shows the order of decreasing stability to be Step 4C. (CH,) ,COOC(CH,) , + 2(CH3) ,COO CH,O > C2H,0 > n-C,H,O > iso-C,H,O > iso-C,H,O = tert.-C,H,O The high yields of tert.-butyl alcohol in the decompositions of the n- and isobutyl and tert.-butyl peroxides suggest the possibility of intramolec- ular reaction to yield the appropriate butyraldehyde and tert.-butyl alcohol without the formation of the corresponding free butoxy radicals.The additional tert.-butyl alcohol is not due to the presence of a relatively l8 See ref. (16). l9 GiguCre, Can. J . Res., 1950, 28, I?. 2o Rieche, AZkyZperoxide und Ozonzde (Theodor Steinkopff, Leipzig, 1931). 21 Raley, Rust and Vaughan, J . Amer. Chem. SOC., 1948, 70, 1336. 22 Rust, Seubold and Vaughan, J . Amer. Ckem. SOC., 1948, 70, 3258. P- 30.BELL, RALEY, RUST, SEUBOLD AND VAUGHAN 249 high concmtration of butyraldehyde, since decomposition of ethyl tert.- butyl percxide with a.molar excess of isobutyraldehyde at 195" gave only an 11.2 yo yield of tert.-butyl alcohol. TABLE III.-REACTIONS OF (CH,) ,COOR WITH CYCLO~EXENE AT 1g5O RO Yield in moles/roo moles of Peroxide Reacted Products Derived from RO By H Abstraction 76 CH,OH 65 CH,CH,OH 30 CH,CH,CH,CH,OH 19 (CH,),CHOH 6 (CH,),CHCH20H 8 (CH,),COH By H Loss 8 CH,O 5 co 8 CH,CHO scot 10 CH,CH,CH,CH( 16 CO t 11 (CH,),CO 11 (CH,),CHCHO I4 CO t By De- composition - * 10 CH,O I C O t 26 CH,O 3COt 35 CH,CHO 23 cot 61 CH,O 6COt 94 (CH3)2CC Products Derived from t&.-BuO By Sta- bilization (CHa)&OH 1 0 I2 2 0 6 22 8 By Decom- position (CHa)&O S2 80 69 82 67 94 * Decomposition of CH,O* according to CH,O* -+ CH,O + H* cannot exceed 0-2 mole/Ioo moles of peroxide reacted, the amount of hydrogen produced. t Carbon monoxide is formed by decomposition of aldehydes following free radical attack. The assumed distribution is based on the stability of formaldehyde as observed in the decomposition of ethyl tert.-butyl peroxide with no additive present . If the relative stability of an alkoxy radical is defined as the ratio of the moles of alcohol produced to the sum of that quantity and the moles of decomposition products, the alkoxy radicals tested are ranked a3 follows : CH,O', 1-00 ; C,H,O' 0.86 ; n-C,H,O*, 0.51 ; iso-C,H,O', 0.25 ; iso-C,H,O' 0.08 ; tert.-C,H,O', 0.08. In addition to decomposing and participating in hydrogen atom transfer reactions, certain alkoxy radicals may associate with . copresent alkyl radicals to form ethers. Thus, a small (8 yo) yield of methyl ethyl ether was prcduced by the union of methoxy and ethyl radical: as supplied by the vapour phase decomposition of methyl tert.-amyl peroxide.23 Still lower yields would be expected with the other, less stable, alkoxy radicals. The authors are pleased to acknowledge the assistance of Messrs. D. 0. Collamer, Jr., and L. M. Porter on portions of the experimental work. Shell Development Company, California. Emeryville, 8, 23 Raley and Collamer (in press).

ISSN:0366-9033    

DOI:10.1039/DF9511000242

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

THE VELOCITY COEFFICIENTS OF THE CHAIN PROPAGATION AND TERMINATION REACTIONS I N OLEFIN OXIDATIONS IN LIQUID SYSTEMS BY L. BATEMAN, G. GEE, A. L, MORRIS AND W. F. WATSON Received 19th February, 1951 The purpose of this new investigation of the kinetics of the liquid phase oxidation of olefins was twofold : ( I ) to study more completely the relative importance of the three termination reactions ; (2) to try by improved experi- mental techniques to obtain more reliable estimates of the absolute velocity coefficients. It is shown that the cross termination process, between unlike radicals, is relatively more favoured in the more reactive olefins. In contrast to co- polymerizations,ll where only substituted-alkyl radicals are involved, no value of 4 very much greater than I has been found, and in one example a fractional value was actually observed. New estimates of the velocity coefficients still leave much to be desired in accuracy.They show little variation between the magnitude of the different radical-radical reactions, and in this respect oxygen behaves as a radical. Non-conjugated olefins in the liquid phase a t moderate temperatures absorb oxygen to form hydroperoxides by the following radical chain reaction mechanism : Initiation : Production of R- or R0,- radicals y i . (I) Propagation : R- + 0, -+ R0,- A , * (4 R0,- + RH -+ RO,H + R- k3 - (3) Termination : k4 * (4) k , - ( 5 ) R- + RO,- --f Non-initiating K6 * ( 6 ) 2R- --f 1 products, 2R0,- --f where RH represents the olefin, R- the radical obtained on removal of an a-methylenic hydrogen atom H, vi is the rate of initiation, and the A’s are the respective velocity coefficients of the reactions indicated.There are two uncertainties in the earlier kinetic analysis : (i) the absolute magnitudc of ri, i.e. for example, its relationship to the overall rate of decomposition of a free radical producing catalyst such as benzoyl per- oxide and (ii) the validity of the assumption made hitherto that k,% = k&,. This paper is mainly concerned with investigating (ii) and with evaluating the velocity coefficients on a more general kinetic basis. The absolute values deduced for these coefficients necessarily depend on (i), and here certain inconsistencies remain to be resolved. Our results assume that in benzoyl peroxide catalyzed oxidations each molecule of peroxide de- composing initiates one chain., Stationary State Relationships for the Generalized Kinetic Scheme.-- A complete kinetic analysis of the set of reactions listed above leads to a complex and intractable rate equation, which, however, can be greatly simplified by introdu5nz the single assumption, valid if the chain length is long, that ~,rR-l[0,1 = A3[RHIrKO,-l * .(1) Bolland, Quart. Rev., 1949, 3, I. Bolland and ten Have, Tram. Faradny SOC., 1947, 43, 201. 250BATEMAN, GEE, MORRIS AND WATSON 25 ' The rate of oxidation Y can then be expressed in terms of three composite constants defined by 1 . - * ( 2 ) A = k2k,-* B = k,kB-*, 4 = k4-* K,k,-* The rcsult is Y = Y ~ { I + z ~ A - ' B [RH][O,]-l + A-2B2 [RH]2[02]-2)-*, . ( 3 ) where the limiting rate a t sufficiently high oxygen pressures is given by yrn = B [RHIY,~ * (4) Y = Y ~ { I + A-'B [RH][02]-1)-1 .* (5) If 4 = I, as hitherto assumed, eqn. (3) reduces to In this case, a plot of y-l against [02]-' will be linear, and the constants A and B are obtainable from the slope and intercept. Data consistent with this relationship have been published % 4, 5 and used to evaluate A and B. More critical consideration of eqn. (3) shows that a plot of y-l against [0,]-1 will be concave to the latter axis if 4 < I and convex if 4 > I, but that so long as 4 does not differ greatly from unity, it would not be easy experimentally to observe the curvature by measurements in the pressure range (Q= 5 mm.) ncrmally employed in earlier work. It is also readily sten that i f eqn.( 5 ) is wrongly employed when 4 $. I the constant deduced as A is actually more nearly equal to k2k5-1k6*. Values pre- viously listed 4 s 5, 8 for A&,-* are therefore open to suspicion, and we shall in fact show below that some of them are in need of considerable revision. Determination of 4 Materials .-The preparation and purification of e thy1 linoleate, ethyl linolenate, methyl oleate and z : 6 : 11 : 15-tetramethylhexadeca-z : 6 : 10 : 14- tetraene (digeranyl) have been indicated elsewhere . 7 ~ 8 Hexadec-Iene, ng 1.4419, was obtained by fractionating a commercial sample through a 15-plate column. Phytene, ng 1-4515, was prepared by reducing phytol in liquid ammonia with sodium and alcohol. All specimens were examined spectro- scopically ; phytene was thus found to contain 3 yo of its less reactive Al-isomer, and the hexadecene contained z yo of its A2-isomer.Immediately before use, any adventitious peroxidic or other polar impurities were removed by running petrol solutions through a 15 cm. alumina column. Azobis(isobutyronitri1e) was recrystallized from methanol. Benzoyl per- oxide was purified by several precipitations from chloroform solution by methanol. These catalysts were introduced into the reaction vessel in acetone solution, the solvent then being evaporated off in a stream of nitrogen. Apparatus .-The constant pressure gas burette system described by Bolland 9 was used with minor modifications. At very low pressures, the con- stant pressure control lacked sensitivity and manual control using the octoil (octyl phthalate) manometer as indicator was adopted.Procedure .-A series of rate measurements ( r ) at oxygen pressures varying from 0.4 to 50 mm. were made on the same sample of olefin plus catalyst. The series was interspersed with measurements (r,) a t a standard pressure (generally 20 mm.) in order that any change in rt due to autocatalysis or catalyst decom- position could be allowed for. A plot of r,/r against [O,]-l was extrapolated to L0,J-l = o to obtain r/va. Then, according to eqn. (3), a plot of {(roo /r)2- I}[OJ Bolland, Trans. Faraday SOC., 1948, 44, 669. Bolland and ten Have, Trans. Favaday SOC., 1949, 45, 93. Bolland, Trans. Faraday SOC., rg50,46, 358. Bateman, Bolland and Gee, Trans. Faraday SOC. (in press).Bolland, Proc. Boy. SOC. A , 1946, 196, 21s. 6 Bamford and Dewar, Proc. Roy. SOC. A , 1949, 198, 252. * Bateman and Gee, Trans. Faraday SOC. (in press).252 OLEFIN OXIDATIONS against [0,]-1 should then be a straight line of slope A -zB2[RHJe and ordinate intercept z#A -IB[RH]. From these two quantities, and making use of eqn. (4)' we can evaluate the three constants A , B and 4. In the above equations [O,] represents the concentration of oxygen in the liquid, and is related to the gaseous oxygen pressure p by [O,] = j3p - IKs-l, 1 (6) where is the solubility of oxygen in the olefin in moles 1. -I (mm. Hg) -l and k, a shaking constant. The latter was determined for the present experimental conditions as described by Bateman, Bolland and Gee,8 by a method independent of 4, and eqn.(6) then used to calculate [O,] from the measured p . In the early part of this work, benzoyl peroxide was used as the initiating agent, but evidence was found of a variation in its efficiency with oxygen pressure apparent a t low pressures. This, together with uncertainty concerning the amount of carbon dioxide evolved per mole decomposed were soon found t o present serious difficulties. Azobis(isobutyronitri1e) was therefore used sub- sequently to obtain all the results reported in this section of the paper. The small values of I encountered a t the lower pressures necessitated corrections for the amount of nitrogen evolved by the catalyst; these were as much as 0-1 I a t the lowest pressure. Nitrogen evolution, which is stoichiometric, was measured in ethyl linoleate solution in vacuo and its rate a t 45" was found to agree with published data There was no evidence to suggest that dilution of the gaseous oxygen by the liberated nitrogen introduced any spurious effects.On the basis that each mole of cata- lyst decomposing initiates one chain (and this is certainly too high an estimate), the chain lengths a t the lowest oxygen pressure varied from g to 40 for the east and most active olefin studied, respectively. to within I yo, i.e. k = 1.06 x 10-6 sec.-l. 3'3 2.5 3.1 I* 0.3 Results Fig. I shows plots of I,/Y against [0,]-1 for ethyl linoleate catalyzed by ( a ) azobis(isobutyronitri1e) and ( b ) benzoyl peroxide. Fig. 2 gives sirnilax plots for phytene, methyl oleate and hexadec-I-ene. The marked curvature of some of these plots shows that 4 differs irom unity. In Fig.3 and 4 are given the deiived plots of { ( Y ~ / Y ) ~ - 1}[0,] against [OJ-l. A satisfactory linear relation- ship evidently exists in the nitrile catalyzed systems-and this applies also to the other olefins examined (Table 1)-and serves to establish that eqn. ( I ) accurately describes the dependence of I on [O,]. The departure from linearity at low values of LO,] when benzoyl peroxide is used as catalyst (Fig. 3b) un- doubtedly reflects the variations in catalytic efficiency referred to above. TABLE I 1.3 1.6 0.9 0.40 0.53 Olefin Ethyl linolenate Ethyl linoleate Digeranyl . Phytene . Methyl oleate . Unsaturation Pattern ~. CH : CH. CH,. CH : CH . CH,. CH : CH . 1 . CH : C H . CH,. CH : CH .. CMe : CH . CH, . CH, . CMe : CH . . CH, . CMe : CH . CH, . . CH,. C H : CH. CH,. , ~~~ ~~ * On the basis of the Y ~ / Y against [O,]-l plot. Values of $ and K,K,-& for five olefins a t 45O C are given in Table I. It proved impossible to determine these quantities for hexadec-I-ene because its rate of oxidation is insensitive to diminishing oxygen pressure above 0.5 mm. (Fig. 2). Comparison of this behaviour with that of ethyl linoleate, for example, affords a striking demonstration of the influence of olefinic structure. The pressure insensitivity of the oxidation rate of the former is readily understandable in lo Matheson, Auer, Bevilacqua and Hart, J. Amer. Chem. Sot., 1949, 71, 2610.BATEMAN, GEE, MORRIS AND WATSON 253 view of the low reactivity of RH and correspondingly high reactivity of R-, these two factors operating in the same direction.The validity of the original assumption that k, = k,+k,+ is obviously justified to a first approximation by the magnitude of the 4's given in Table I. Notwithstanding this, a definite trend in the values is apparent, viz., the more reactive the olefin (at high oxygen pressures), the higher the 4. FIG. I. 0 Mekhyl deale x Phyfene A Hexadec-I-ene 45 "c FIG. 2. Non- Stationary State Relationships To proceed further, it is necessary to make observations of the photo- chemical oxidation during the non-stationary period when the light has just been switched on or off. Such measurements can lead to the evalua- tion of two further composite constants : These, together with A , B and 4 permit the five individual velocity con- stants to be calculated.C = k , / k , ; D c ks/k6. . - (7)254 OLEFIN OXIDATIONS The photochemical after-effect has been discussed by Bateman and (compare Bamford and Dewar 6, who define the " decay intercept " Gee I d by where YD is the rate of the stationary dark reaction and t is the time FIG. 4. measured from the instant of switching off the light (cf. Fig. 5 ) . a " growth intercept " I, is defined by Similarly Ig E ~ , " ( Y L - r)dt, . - (9) where YL is the rate of the stationary light reaction, and t the time of illumination.BATEMAN, GEE, MORRIS AND WATSON 25 5 By an extension of the method of Bateman and Gee to the more general kinetic system discussed above, these intercepts may be evaluated as D [RH] + A-,B2C [RH]2[0,]-1 where a = I + z+A-'B [RH][O,]-l + A - 2 B 2 [RHl2[O,]-2 .(12) For reasons already advanced 0 and further discussed below, the experi- mental observations are best analyzed as the difference of these two intercepts, i.e. We have therefore used this equation to evaluate a. Tne denominator oi the expressim for a (eqn. (12)) will be recognized as the main term of eqn. ( 3 ) , and is entirely determined from stationary state data. This quantity, multiplied by- a is, according to eqn. (I z ) , a linear function of [O,]-l, the slope and intercept determining C and D respectively. Bateman, Bolland and Gee have measured the photo- chemical pre- and after-effects in oxidizing ethyl linoleate and digeranyl at an oxygen pressure of I mm.in an attempt to derive k,R,-l. However, serious difficulties were encountered in analyzing the experimental data owing to [O,] being appreciably different in the dark and light periods. It was not found possible to make due allow- ance for this difference at the same time as considering a reaction mechanism involving more than one termination process. The simplifying, but FIG. 5. crude, assumption that the rate-determining propagation step was reaction (2) and that termination occurred solely by reaction (4) under the con- ditions prevailing was therefore introduced, and it was recognized that the final results could only be roughly quantitative. In this paper, we have adopted the alternative approach, viz., experiments have been re- stricted to pressure regions where, with high shaking efficiency, it was permissible to neglect the effect on Y of changes in [O,] during the non- stationary period.Eqn: (13) is only applicable when this assumption is valid. A considerable improvement in precision has thus been effected, but the numerical unceratinty in the derived values still remains rather high. In particular, since measurements do not extend to the pressure region where reaction (4) is the main terminating reaction, our estimates of k z and h, are only approximate. l1 See Mayo and Walling, Chem. Rev., 1950, 46, 272 ; Burnett, Quart. Rev., 1950, 4, 322.256 OLEFIN OXIDATIONS Determination of k3/k6 and kz/kp Apparatus.-This is shown in Fig. 6 and was used for all the experiments over the oxygen pressure range covered (- 5-700 mm.).It incorporates several improvements over that used previously : 6 (i) the reaction vessel, glass spiral and manometers were totally immersed in the same thermostat (maintained to f 0 * 0 0 5 O C) ; (ii) the spiral was sufficiently flexible to permit a shaking frequency of up to GU. 1000 per min. ; (iii) octoil was replaced by lower viscosity silicone fluid in the constant volume manometer and (iv) the shape of the reaction vessel gave improved shaking efficiency, while the wide inset neck prevented small liquid drops momentarily sealing the connecting tube and causing jerky manometer movements. Irradiation was provided by a B.T.H. stabilized 125 W, bare quartz mercury arc, focused on the reaction vessel by a quartz condensing lens through the Pyrex thermostat window.A manually-operated shutter served to interrupt the light. Shaking was effected by a potentiometer-controlled, A.C. stirrer motor, the shaft being rigidly coupled to the stem of the reaction vessel, and was measured t o about f 20 rev-jmin. by a Whidbourne stroboscope. 3 ml. samples of olefin were used. E L - A FIG, 6.-A is the pear-shaped reaction vessel (- 23 ml.) with glass ring below the neck, B a volume-compensating bulb, C and D, silicone and mercury manometers, respectively. E connects to an electrolytic cell controlled by the make-and-break I;. LL indicates the thermostat water level. The decay and growth intercepts were measured as divisions of a graticule in the cathetometer trained on one limb of the silicone manometer ( IOO div.N 0.4 mm.). Appropriate calibration to permit these quantities to be expressed in absolute units was effected by directly comparing steady rates of oxidation on the silicone and mercury manometers. Movement of the meniscus over one graticule division corresponded to an oxygen uptake of 4-20 x 10-6 mole 1.-1. Under the best conditions, the intercepts could be defined to f 0.1 div. and I,,’ are not identical with Id and I, as defined above, but are equivalently related to them,e viz., Disturbing FaCtOrS.-THE EXPERIMENTALLY OBSERVED INTERCEPTS, 1,’ I,’ = I , + (’L - YD)(kr-l + Km-l) . - (14) I,’ = I, + (‘L - *Jg)(k,-l + Km-l), (15) . where K, is a constant expressing the shaking efficiency and k , the speed of mano- meter response. Since Km-l mas found to be negligible compared with k,-1 in the earlier work 8 and substitution of silicone fluid for octoil ensures more rapid manometer response it need not be further considered.The term (rL - r ~ ) K , - l measures the difference in [O,] produced by changing the reaction rate from rD to rL, and i t is obvious from eqn. (14) and (15) that its effect can in principle be eliminated by using (I,‘ - I,’). It must, however, be stressed that if ( r ~ - r ~ ) k , - l is large compared with I , or I,, the experimental uncertainty in the derived constants will necessarily be magnified. Unfortun- ately, although our investigations have been confined to conditions whereBATEMAN, GEE, MORRIS AND WATSON 257 (rL - YD)k8-l is small compared with [O,], i t is in most instances large com- pared with I,.Two examples may be quoted from experiments on the oxidaxion of ethyl linoleate at 550 mm. pressure with a shaking speed of 650 per niin. At 15" C, Id', I,', and (rL - rD)k,-l were respectively 27, 19 and 16 ( x 10-6 moIe/l.), a t 45" C the corresponding figures were 24, 17 and 12. EVALUATION OF k,/k,, AND k,/k,.-The non-stationary state data required for solving eqn. (13) for ethyl linoleate and digeranyl * are presented in Tables I1 and 111, respectively. Also required are values of k & J , k , k , 3 and 4 from stationary state measurements: these are given in Table IV. Values of the TABLE II.-Id' AND I,' FOR ETHYL LINOLEATE AT 25°C Liquid vol., 3 ml. ; shaking speed, 650 per min. Ethyllinoleate . Digeranyl . . Po* mm. 1.21 x IO-* 1.07 XIIO-~ xOBa mole/l.1.6 x 103 0.9 x 103 I95 31 16 9'7 6.9 5'4 2.3 x 10, 1.3 x 103 I 6-0 16-8 15.1 12-6 13.5 9'3 9'3 7'3 7'3 5'5 5'9 12-6 12.6 10-9 8-4 8-4 4'2 4'2 2'5 2'5 1'7 1'7 22-8 20.9 17-2 16-8 13-7 I 4-0 10.6 21'0 10'2 ;:; 4'74 3.00 3-20 2-03 1-89 1-14 0.98 0'77 0'74 0'53 0.62 6-1 5'2 5'5 4'3 4'9 4'0 3'8 3'5 3'5 2.6 2'9 TABLE III.-ld' AND I,' FOR DIGERANYL AT 25' C Liquid vol., 3 ml. ; shaking speed, 650 per min. PO2 mm. 71 1 292 37 I4 7'7 4'9 I mole L-1 - I 15.1 I 6.0 19'4 I 8-5 14'3 14'3 10.5 10.9 8.4 13'9 12'2 12'1 15'9 I 6.0 11'0 11'2 8.0 9.0 6.1 I 0.4 TABLE IV 1-65 1-72 2'34 1.60 1.61 0.95 0.95 0.60 I .03 2'22 7'7 9.0 9.1 8.4 8.4 7'3 6.4 6.5 5'3 4'9 I 2.5 3'1 * See text. "Since each propagation cycle here results in the incorporation of two molecules of oxygen,4 all velocity coefficients for digeranyl are half those cal- culated by the equations given.I258 OLEFIN OXIDATIONS last two quantities at 25’ have not yet been measured, but adequate correction to the data at 45’ is possible : the correction is in fact small and its effect on the values of k J k , and k3jk, almost negligible. The resulting plots of eqn. (13) show somewhat scattered points (Fig. 7), and statistical regression analyses have therefore been made to ascertain the best linear correlations. These were kindly carried out by Mr. G. E. Blackwell, who reports that highly significant correlation coefficients are found (probability level being better than 0-1 %). From the slopes and intercepts of the lines drawn the following results are obtained : Digeranyl .ka/k4 1’4 rt 0.3 106k,/k, 0.36 & 0.05 Ethyl linoleate . k2/k4 0.6 f 0.3 106k,/k, 2.3 f 0.5 r Ethyl linoleate . Digeranyl . . L 4 I@ke Idka I d k s x 10-6 x 10-6 I d k s 4 x 10-6 x 10-6 -I____ _____- 9 60 20 50 30 I 3 0-9 9 9 / CF X- -D/gerany J €lhyJ /inoleare 25 “c x 0;5 /i 0 /i5 ZiU 2.5, FIG. 7. Values and Significance of the Velocity Coefficients The individual k’s derived from the measured composite quantities are given in Table V; the uncertainty’ in the k,/k4 and k,/k, values do not justify more than onc significant figure being quoted. When com- pared with earlier results,% considerable discrepancies are apparent, particularly in the absolute values of k, and k,, and these naturally raise the question of the quantitative significance of such measure- ments generally.The results now given are believed to be reliable, but we wish to stress most strongly the inherent limitations of the present technique owing to the smallness of the photochemical pre- and after- effects and the relatively large effects due to delayed oxygen diffusion. It is now clear that at oxygen pressures much below 5 mm. significant TABLE V.-VELOCITY COEFFICIENTS AT 2 5 O C (mole-1 1. sec.-1)BATEMAN, GEE, MORRIS AND WATSON 259 quantitative estimates are unobtainable, and even at high pressures it is doubtful whether the individual K's are correct to f 50 "/b in the most favourable cases. In less reactive olefins, e.g. methyl oleate and phytene, we know this uncertainty to be increased several fold. The main conclusion reached previously * that increased oxidizability (at constant ri) reflects an increase in K,, with K , remaining substantially constant, is essentially confirmed, although some difference in K , appears to be indicated for ethyl linoleate and digeranyl. What is surprising is the higher k4 value of ethyl linoleate, since energetic considerations require that the reactivity of R- be the inverse of RH, as is indeed consistent with other oxidation features. However, the activation energy of radical- radical interaction is very small and unquestionably steric factors could become dominant and effectively confer special stability on the R-radical from digeranyl.' This work forms part of a programme of fundamental research under- taken by the Board of the British Rubber Producers' Research Association. 48 Tewin Road, Welwyn Garden City, Herts.

ISSN:0366-9033    

DOI:10.1039/DF9511000250

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

BATEMAN, GEE, MORRIS AND WATSON 259 THE KINETICS OF OXIDATION OF HYDRO- CARBONS IN THE GAS PHASE A THEORY OF THE LOW-TEMPERATURE MECHANISM BY M. F. R. MULCAHY Received 2nd January, I 95 I It is suggested that the mechanism of oxidation of hydrocarbons in the gas phase in the region of 200-300° C is basically similar t o that shown by Bolland and others to occur in the oxidation of more reactive hydrocarbons in the liquid phase at lower temperatures. The kinetic consequences of such a mechanism are worked out for the conditions which occur in gas phase experiments and shown to be in agreement with the experimental results. Introduction.-Evidence has accumulated that in the oxidation of many organic compounds at not too elevated temperature, the first product to be formed is a hydroperoxide,* viz.RH + O8 + ROOH. . * (1) The evidence is most clear in the cases in which the molecule contains a weak C-H bond (e.g. aldehydic, s, 6m * or a to a double bond or conjug- ated system 1, 8s 6 ) , presumably since the reaction proceeds at measureable speed at a sufficiently low temperature for the resultant hydroperoxide to remain undecomposed. On the other hand, such low-temperature * For examples, see ref. (I)-(8). 1 Farmer and Sutton, J . Chem Soc., 1942, 139. George, Rideal and Robertson, Proc. Roy. SOC. A , 1946, 185, 288. Almquist and Branch, J . Amer. Chem. Soc., 1932, 54, 2293. McDowell and Thomas, J . Chem. SOL, 1949, 2208. 6 Wittig and Pieper, Liebig’s Ann., 1941, 546, 142. 6 Bolland, Proc. Roy. SOC. A , 1946, 186, 218. 7 Badin, J .Amer. Chew. Soc., 1950, 72, 1550. Bowen and Tietz, J . Chem. SOL, 1930, 234.260 LOW TEMPERATURE MECHANISM oxidatjons are catalyzed by decomposing peroxides 6* sp l o p loa and, in suitable circumstances, may become autocatalytic, hydroperoxide being regenerated faster than it is decomposed. There is ample evidence that these reactions proceed by free radical chain mechanisms. The experiments of Cullis and Hinshelwood,ll of Egerton and Young,12 and others have shown that during the gas-phase oxidation of hydro- carbons at 200-300° C the peroxide content of the mixture rises to a maximum and falls away again, the maximum peroxide concentration being reached close to the time at which the maximum rate of pressure change occurs.11 Cullis has obtained evidence that during my one reaction (with hexane) the rate of pressure change is proportional to the instantaneous con- centration of peroxide.It is the purpose of this communication to examine to what extent the type of kinetic mechanism shown by Bolland %I4 and others to be operative in reactions proceeding in the liquid phase near room temper- ature may also apply to the oxidation of hydrocarbons in the gas phase in the region of 200-300° C. Probably the three most outstanding phenomena associated with the low temperature oxidation of hydrocarbons in the gas phase are (a) the autocatalytic nature of the reaction, (b) the steep dependence of the maximum rate of pressure change on the initial partial pressure of hydro- carbon combined with independence of the oxygen pressure (both under conditions of excess oxygen) l1 and (c) a large effect of the structure of the hydrocarbon on the maximum rate which, for example, increases about fivefold with each successive member of the paraffin series.1l If then a mechanism is to be considered valid for these conditions it is necessary at least to show that it provides an explanation of these basic facts.The Reaction Mechanism.-If, by some means, the radical R (or RO,) is produced (at the rate given by 4), then hydroperoxide is formed by the following reactions : R- + 0 2 -+ ROZ- . (2) ROZ- + RH -t RO2H + R- . * (3) RO2- + ROZ- -+ X . * (4) R-++- - t X . - ( 5 ) Chain termination may occur either by bhnolecular recombination of radicals, viz. { RO2 + R -tx . * (6) A or by some reaction first order with respect to the radicals, e.g.by collision with the wall or with some other (inhibiting) molecule ROS--+X . * (4’) B {Ib 4 x . * (5’) Chain termination of type A leads (for long chains) to the following expression : l4 * (1) --= dP21 d[ROOHI - --- d[PI %d* [RHI [Ozl dt dt - dt - KZK4) [O,] + K3k5) [RH]’ Chain termination of type B yields 9 George, Proc. Roy. Soc. A , 1946, 185, 337. 10 Bolland, Trans. Furaduy Soc., 1950, 46, 358. 10a Robertson and Waters, Trans. Faruday SOC., 1946, 42, 201. 11 Cullis, Hinshelwood, Mulcahy and Partington, Dzscusszon Fuvaduy Soc., 12 Egerton and Young, Trans. Faraduy SOC., 1948, 44, 755. 13 Cullis, Thesis (Oxford, 1947). 1947, 2, 111. 14 Bolland, Quart. Rev., 1949, 3, I.M. F. R. MULCAHY 26 I Except for the exponent of r j l both expressions are of the same form with regard to the concentrations of oxygen and hydrocarbon, and for our present purposes need not be distinguished.However, it may be remarked that, for a gas phase reaction, first-order chain termination (B) seems the more probable. In uncatalyzed reactions 4 is a function of [RH] and [O,]. Where peroxide decomposition is the only initiating reaction 4 is a function of [PI. Derivation of an Expression for the Maximum Rate in Terms of Initial Concentrations.-We shall take as our fundamental assumption that the rate of pressure rise is given by p = 22 = K[P], . (111) dt where K is a numerical constant. In the system under discussion in which peroxide is both decomposing and being regenerated at comparable rates, we have for the rate of change of peroxide concentration : dB = A + B[P] - C[P].. dt A is the rate of production of peroxide in the absence of peroxide, and is of the form (see eqn. (11)). This gives the rate at the very beginning of the reaction and in a favourable case may approximate to the rate during the in- duction period.l% lo B[P] gives the rate of peroxide production produced by autocatalysis. B is of the form This term is taken proportional to the first power of the peroxide con- centration, without deciding whether this results from unimolecular radical-producing decomposition of the peroxide followed by first-order termination reactions (eqn. (11)) or bimolecular decomposition (which occurs in solution in some cases14) followed by second-order radical recombination (eqn.(I)). * C is a constant which allows for the possibility that the peroxide may also decompose (unimolecularly) in a manner which does not give rise to active radicals. We may now write where K is a constant including the speci,fic rate of chain initiation by per- oxide breakdown. We shall not deal further with the very early part of the reaction, but shall consider the (usual) case in which the rate during the induction period is very small compared with the rate subsequently attained. This means that from the end of the induction period, the contribution of A to the rise of peroxide concentration may be neglected. We have therefore from (IV) ( V I ) l5 Mulcahy, Trans. Faraday SOC., 1949, 45, 575. 16 Ridge and Mulcahy (to be published).* The molecularity of the initial act of decomposition of hydroperoxides in the homogeneous gas phase is not known. has found that when hydroperoxides are caused to decompose per se, the reaction occurs largely a t the walls of the vessel. Harris Harris, PYOC. Roy, SOC. A , 1939, 173, 126.262 LOW TEMPERATURE MECHANISM The solution of eqn. (VI) does not follow immediately, since B is time-dependent because of the diminution of [RHI and [O,] as the reaction proceeds. A simplification may be introduced, however, by considering the case in which one of the reactants is in excess so that its change in concentration during the reaction may’ be neglected. Taking [O,] to be in sufficient excess that [O,] > Y’[RH] in eqn. (V),* and that this condition remains until the end of the reaction, we have In this case, therefore, the eventual falling-off in the rate is due to de- pletion of RH.Now in the ideal case, RH disappears from the system exclusively by reaction leading to peroxide : RH + 0, -+ ROOH. (The rise in pressure may be attributed to breakdown products from ROOH.) B = K[RH]. . . . (VII) Hence, at any time t [RH] = [RH], - Dr[P]dt, . (VIII) where [RH], is the initial concentration of RH and D a numerical con- stant which will not be far from unity. We now have from (VI), (VII) and (VIII) Setting (K[RH], - C) = ,!? we have from (111), This gives, on transformation and integration, - (XI) dAP KD p = dt = PAP - -(Ap), + const . 2k Since p is small where Ap is small it is clear that the integration constant is small compared with PAP.Moreover, in order to make the solution continuous with that which obtains for conditions towards the end of the induction period, it is also necessary that the integration constant should be small. We shall take it to be zero. However, though this step is not unreasonable, it can only be strictly justified a posteriori. Integration of eqn. (XI) then gives Ap = (I + tanh . (XII) KD where t, is a constant of integration. At t = t , we have (XIII) and from (X), ( $ ) , = o . . (XIV) Thus from (XIV) it follows that t , is to be identified with the time at which the maximum rate Pmax is reached. Hence, substituting (XIII) in (XI) gives where k’ is a numerical constant. * It is to be remarked from (I) and (11) that Y’ is very probably < I since k , is almost certainly > k,, and k , and k, probably of the same order of magnitude.M.F. R. MULCAHY 263 A very simple result may be had from (XIII) and (XV), viz. . (XVI) . (XVIU) Pmax fl Apmax 2' Pmsx I Apmax 2 -=- i.e. -- --(K[RH]o - C), . Apmax being the pressure increment from the beginning of the reaction to the point at which Pmax occurs. Comparison with Experiment.-The slow beginning and eventua.1 acceleration of the reaction are accounted for in terms of a slow produc- tion of hydroperoxide from the pure reagents followed by auto-catalysis due to free-radical-producing breakdown of the hydroperoxide, as set out above. This general explanation coincides with that previously proposed by Hinshelwood.17. 11 THE RATIO Pmax/Apmax.-Eqn. (XVIa) predicts a linear relationship between Pmax/A$max and the initial concentration of hydrocarbon, pro- vided an excess of oxygen is present.In Fig. I Pmax/Apmax derived from experiments with different initial pressures of butane with 250 mm. FIG. I .-Variation of hx/APmax with initial pressure of hydrocarbon. (Initial oxygen pressure constant.) oxygen at 263OC and with different initial pressures of propylene with 400 mm. oxygen at 298OC l6 respectively, is plotted against the initial partial pressure of hydrocarbon. In both cases a linear relationship is found. It is interesting to observe that with butane the straight line passes through the origin, i.e. C = o in eqn. (XVIa). In terms of the present theory (eqn. (IV)) this means that the decomposition of butane hydroperoxide under these conditions occurs only by a mechanism (one or more) which produces radicals capabIe of initiating more peroxidc formation.On the other hand, with propylene C 9 0, which indicates that the peroxide formed from propylene decomposes in two ways, one of which does not produce active radicals. It is to be observed that K in eqn. (XVIzz) is independent of oxygen concentration when the latter is in excess. Hence Pmax/Apmax should also be independent of oxygen pressure under these conditions. Fig. z shows that this is so for butane (50 mm.) + oxygen mixtures at 258.5" C when the partial pressure of oxygen is greater than - zoo mm. Measure- ments of the final pressure reached at the end of the reaction indicate that three molecules of oxygen are required for complete reaction with one of butane ; hence in these experiments the condition of excess oxygen Mulcahy, Faraday SOC.Discussion, 1947, 2, 128. 17 Hinshelwood, J . Chem. Soc., 1948, 531. lS Mulcahy, Thesis (Oxford, 1948).264 LOW TEMPERATURE MECHANISM is not fulfilled until at least 150mm. are present. Below this pressure PmaxlApmax rises steeply. A similar result was found with 80 mm. butane and various oxygen'pressures at 263O C (cp. ref. (IS)).* THE MAXIMUM RATE.-h cases where C = 0, the theory indicates that, in the presence of excess oxygen, Pmax should increase with the square of the initial hydrocarbon pressure and where C .c: o with a power be- tween the second and first. In both cases Pmax should be independent of oxygen pressure. The latter result has been observed in many cases.lln 20* 2 l s 2 2 ~ 25 A steep rise in pmax with [RH], is invariably observed.In several cases 1% 20, 21, 2 2 the exponent of [RH], has been found to be approximately 2 (1.7 in the butane experiments and ca. 2 in the propylene experiments in Fig. I). FIG. 2 .-Variation of pmax/APmaxwith initial pressure of oxygen. (Initial butane pressure constant.) THE STRUCTURAL EFFECT.-For equal pressures of different hydro- carbons with excess oxygen under the same conditions and C = o in from eqn. (XV). Taking the case of unimolecular chain termination, reference to eqn. (11) shows that K = k,"k,/k,' where k" is the specific rate of radical production from peroxide decom- position. Assuming that all peroxides produce the same number of active radicals for each molecule decomposed, this may be taken to be directly proportional to the unimolecular rate constant K, of peroxide decomposition : k,' is unlikely to be very specific, consequently the effect of the structure of the hydrocarbon on Pmax is to be attributed mainly to variations in KIA,.The increase in K with increasing chain length of the straight chain paraffins and its decrease with branching in the chain must then be ascribed to corresponding variations in A,, since k , would be expected to alter little with increasing chain length and to increase in those cases where branching involves introduction of tertiary C-H groups into the molecule. A similar conclusion has been arrived at by Hinshelwood on somewhat different grounds.ll9 l7 Ceteris paribus, a variation of 2 kcal./ * There is some indication that Pmax/Afimsf reaches a maximum when the 2o Cullis, Trans.Faraday Soc., 1949, 45, 709. 21 Cullis, Hinshelwood and Mulcahy, Proc. Boy. SOC. A , 1949, 196, 160. 22 Pease, J . Amer. Chem. SOG., 1938, 60, 2244. 23 Day and Pease, J. Amer. Chem. SOC., 1941, 63, 914. each case pmax = const. K ROOH -+ RO- + -OH. initial oxygen and butane pressures are approximately equal.M. F. R. MULCAHY 265 per mole in the dissociation energy of the 0-0 bond would produce (at 250' C) a seven-fold change in Pmax ; 5 kcal./mole difference would change the rate by a factor of 120. Conclusions.-In brief, the theory postulates that the mechanism of oxidation of hydrocarbons in the gas phase at 200-300' C is basically similar to that by which more reactive hydrocarbons are oxidized in the liquid phase at lower 14, 32 Hydroperoxide is formed during the induction period and thereafter catalyzes its own formation.This occurs by the production of radicals incidental to the decomposition of the peroxide. The rate is prevented from accelerating indefinitely by the consumption of one or both of the reagents; in the case con- sidered this is due to depletion of the hydrocarbon concentration. The rise in pressure during the reaction is attributed to the formation of products of peroxide decomposition and their subsequent (non-rate- controlling) oxidation. It is clear that the rate equation developed from this mechanism (eqn. (IX)) and its solution ((XV) and (XVI)) can only be approximate, since no account is taken of the possibility of sensitization of the reaction by radicals produced during the oxidation of degradation products or of inhibition by substances formed during the reaction. In view of the complexity of the reaction products, which usually include both formalde- hyde and higher aldehydes, it is most probable that both effects occur to some extent.However, the theory accounts well for the very different effects of initial hydrocarbon and oxygen pressures on the maximum rate. Furthermore the prediction of the correct relation between the initial concentrations and Pmax/Apmax is particularly striking. Again, the strong dependence of Pmax on the structure of the hydrocarbon is quali- tatively in accordance with the considerable range of stability shown among different peroxides (for example, cf. ref. (16a), (24)-(31)). It is evident, therefore, that the proposed reaction mechanism, which is itself a natural extension of a mechanism well substantiated for liquid phase reactions, provides a. reasonable correlation of several of the main facts of the kinetics of the low-temperature gas-phase oxidation. There are various other points at which the ideas and conclusions given above may be tested experimentally. Such experiments are at present in progress. The author is very much indebted to Dr. J. K. Mackenzie for providing him with the solution of the differential eqn. (IX). This work forms part of the general programme of research of the Division of Tribo- physics, Commonwealth Scientific and Industrial Research Organization. C.S.I. R.O., Division of Tribophysics, The University, Melbourne, A ustrali a. 24Harris and Egerton, Proc. Roy. SOC. A , 1938, 168, I. 25 Raley, Rust and Vaughan, J . Amer. Chem. SOC., 1948, 70, 88. 26 Milas and Surgenor, J . Amer. Che.m. SOC., 1946, 68, 205. 27 Medwedew and Alexejewa, Ber., 1932, 65, 131. 28 Robertson and Waters, J. Chem. Soc., 1948, 1578. 29 Farkas and Passaglia, J . Amer. Chem. SOC., 1950, 72, 3333. 30 Nozaki and Bartlett, J . Amer. Chem. SOC., 1948, 68, 1686. 31 Redington, J. Polymer Sci., 1948, 3, 503. 32 Ramford and Dewar, Proc. Boy. SOC. A , 1949, 198, 252.

ISSN:0366-9033    

DOI:10.1039/DF9511000259

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

THE INFLUENCE OF SUBSTITUENTS ON THE OXIDATION OF HYDROCARBONS BY C . N . HINSHELWOOD Received 15th February, 1951 The influence of substituents on the ease of oxidation of hydrocarbons is described, and discussed in relation to theories of oxidation and electronic theories about reactivity in general. Hydrocarbon oxidation has been discussed on more than one occasion in recent years, so that a general survey of the problem will not be made in this paper which will deal rather with one special aspect of the problem, namely the influence of structure and substitution on reaction rate. In spite of various ambiguities and complexities in the situation there is general agreement that in the range of what are conventionally called low temperatures, that is from 1 5 0 O to 300° C, the oxidation of hydrocarbons takes place by a mechanism depending on the formation of a peroxide which slowly decomposes into radicals, thereby setting up branching chains.1 R-0-O-R’ (where R’ may be H) splits into R-O- and R’-0- which in general break down, e.g.by the reaction The smaller hydrocarbon radical attacks oxygen and continues the chain. The range characteristic of this so-called low temperature mechanism is some zooo to 300’ lower than that in which the pyrolytic reactions be- come conveniently measurable. The contrast is clearly due in part to the reactivity of the oxygen biradical, but since the initial rate of attack of oxygen on hydrocarbon is quite slow compared with that attained when the branching chains have developed to the maximum extent, a second major factor is the relative ease of splitting of the central bond in RO-OR’.In the Discussion on the Labile Molecule 2 it was pointed out that the oxidation reactions are enormously more sensitive than the decomposition reactions to the structure of the hydrocarbon. This fact has, in part at least, a kinetic explanation. The rate of the branching chain reaction is of the form RICHS-O- = R1 + HCHO. B(I - e--At) A according as the system is stationary or non-stationary. In either case A or A’ measures the degree of branching and is a difference of two terms, one of which includes the velocity constant for the splitting of the peroxide into radicals. Variations in this velocity constant have therefore a magnified influence on the rate. There is nothing comparable to this situation in the reactions of pyrolysis.This is one reason for the contrast, though another may well be that the -0-O- link itself is specially sensitive to the nature of the attached groups. The expression for the oxidation rate contains two factors, one pro- portional to the initial rate of attack of the oxygen molecule on the hydrocarbon (chain-initiating process) , the other dependent upon the rate of branching. In principle, both map be influenced by the structure Semenov, Chemical Kinetics and Chairt Reactions (Oxford, Clarendon Press, 1935) ; Cullis and Hinshelwood, E;araday SOC. Discussions, 1947, 2, 117. a Faraday SOC. Discussions, 1947, 2, 117; Hinshelwood, J . Chem. SOC., B(eA’l - I) A’ ’ or 194% 531. 266C. N. HINSHELWOOD 267 of the hydrocarbon, but the second is likely to be the more important, since its effect is enhanced in the manner described.We shall consider, therefore, the influence of changes in the structure of R (including the introduction of non-hydrocarbon groups as substituents) on the reactions (a) R-0-OX = RO- + XO- (b) RH + 0, = R- + H-0-0-, and with special attention to the first. The facts recently brought to light about the structural effect may be summarized as follows. I. The rate of oxidation is very much lowered by the introduction of extra methyl groups into a hydrocarbon.2 The mode of initial attack by oxygen which leads to the peroxide most actively capable of chain-branch- ing is that made on a CH, group as remote as possible from methyl. 2. If there is no place for the attack except a methyl group, as in CH,CH,, CH,COCH, or CH,OCH, there is a veyy great reduction in rate of oxidation compared with that of a corresponding compound containing a CH, group., 3.The substitution of chlorine increases oxidation rate, and acts in two ways, firstly by destroying the symmetry of any methyl group into which the substituent enters, and secondly by a direct inductive effect of the chlorine atom.4 If definite -numerical magnitudes are as- signed to the stabilizing influence of methyl groups and to the activating influence of chlorine atoms, a coherent relation appears between the oxidation rates and the calculated " stability factors '' for a whole series of hydrocarbons and chloro compounds. Acetone, it is true, oxidizes only about a quarter as fast as propane under similar conditions, but about 20 times as fast as ethane.The slowness compared with propane is due to the destruction of the CH, group on passage from CH,CH,CH, to CH,COCH,. 2-Pentanone reacts about as East as pentane and 3-pentanone several times faster. Butanone reacts about as fast as butane. In the passage from CH,CH,CH,CH, LO CH,COCH,CH, two equivalent CH, positions become reduced to a single one, and the fact that in spite of this the rate is maintained shows that the inductive effect of the carbonyl oxygen is essentially one favouring oxidation. 5. The introduction of an amino group leads to an increased oxidation rate.6 With propane and butane the effect is not very great but with ethane it is most marked, once again, presumably because the symmetry of a methyl group is destroyed and a CH, group appears in The molecule.6. Ethers are very much more easily oxidized than paraffins,' diethyl ether reacting about 2500 times as rapidly as pentane. 7. With unsubstituted hydrocarbons, chloro compounds, amino compounds and ketones the rate increases very rapidly with the lengthen- ing of the normal carbon chain, but with the ethers the increase from ethyl ether to higher ethers is relatively small, as has recently been found by Dr. Eastwood. The enormous increase caused by the 0 atom appears to have saturated the possible response of the molecule. It will be observed that the only substituent which stabilizes the mole- cule towards oxidation is the methyl group. This is an electron-repelling group, with an inductive action represented by - I in the conventional symbolism.If we regard ethers as hydrocarbons with the substituent 4. Carbonyl groups likewise tend to increase oxidizability.6 Mulcahy, Trans. Faraday SOC., 1949, 45, 537. Cullis, Hinshelwood and Mulcahy, R o c . Roy. SOC. A , 1949, 196, 160. Cullis and Smith, Trans. Faraday SOC., 1950, 46, 42. 6 Bardwell (in press). 'Malherbe and Walsh, Trans. Faraday SOC., 1950, 46, 835 ; Eastwood (unpublished work).2 68 INFLUENCE OF SUBSTITUENTS OR, then the groups which accelerate oxidation are OR, C1 and NH,, and all these are of the + I type. The moment C+-0- of the carbonyl group indicates that in ketones also a flow of electrons into the sub- stituent occurs. We may conclude that the two major influences of structure on oxidation are firstly, the great tendency of methyl groups to preserve themselves intact and, secondly, a stabilizing influence of electron access to the seat of reaction. The first influence, which wherever possible displaces the seat of reaction away from methyl groups, is probably conditioned by the symmetry of their structure and the consequent possibility of electron delocalization.The second depends upon the response to electron accession or recession of the reactions ( a ) and (b) referred to above. Walsh 8 has argued that bonds between strongly electronegative elements should be strengthened by electron-repelling groups, since these facilitate the expansion of atomic orbitals and allow increased overlap without the occurrence of nuclear repulsion.If this is so, then the stabilization of RO-OR’ by the -I groups and its weakening by +I groups would provide a satisfactory interpretation of the observed effects. Electron displacement phenomena must show saturation and this fact would explain why introduction of -0- (in diethyl ether) is not reinforced by a lengthening of the carbon chain, whereas lengthening of the chain in a simple paraffin causes progressive and marked increases in the rate of oxidation. The influence of electronic displacements on reaction (b) is, as already pointed out, less important, and is likely to be smaller in any case, since the C-H bond exists between atoms with a much less pronounced electro- chemical character than oxygen. It is surprisingly difficult to form an unambiguous opinion about the effect which electron-accession should have on the strength of this bond.There is, however, another possible line of argument about the ease of reaction which, with all due reserve, seems to be as follows. RH + 0, = R + HO,, the attacking oxygen must be changed into the biradical form under the influence of RH. In this change electrons are drawn from the O=O bond, and the very initial stages of the necessary re-distribution may well be helped by the action of a positive centre in R which would tend to attract electrons towards it. If this were so +I substituents would favour attack and -1 substituents would weaken it. Thus methyl groups, by causing a flow of electrons towards the seat of reaction, would weaken the attack and the other substituents would intensify it. It is, of course, not very difficult to formulate qualitative arguments which might seem to lead to an opposite conclusion. But in general the influence of substituents on the reactivity of organic molecules seems most profitably to be assessed in terms of the ease of approach of the attacking agent, and the initiation of the requisite reorganization. If this point of view is applicable here also, the effects of the substitution on reaction (a) and on reaction (b) are in the same direction and reinforce one another so that a general interpreta.tion of the observed changes in oxidizability becomes possible. Physical Chemical Laboratory, In the process South Parks Road, Oxford. a Walsh, J . Chem. SOC., 1948, 398.

ISSN:0366-9033    

DOI:10.1039/DF9511000266

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

EVIDENCE RELATING TO THE COMBUSTION OF HYDROCARBONS BY R. G. w. NORRXSH Received z z n d March, 1951 It is generally agreed that the combustion of hydrocarbons occurs by way of a chain process showing delayed branching, but there have been differences of opinion about the identity of the intermediate compound essential to this kinetic conception. In the present memoir some new evidence bearing on the problem -not all of which has yet been published-is cited. The kinetics of the combustion of gaseous hydrocarbons exhibit certain common characteristics which are mechanistically explained by the conception of degenerate branching. According to this idea, which has been fully expounded by Semenoff,l the auto-catalytic character of the reactions and the long induction periods are associated with the build-up of a moderately stable intermediate compound by way of a chain reaction, and this intermediate is responsible for a delayed branching, leading to the slow multiplication of reaction centres.While there is not any serious difference of opinion about the general kinetic character of the reactions, there are more than one hypothesis as to the exact nature of the reactions involved and the relevant intermediate substances re- sponsible for the branching reactions. It is the purpose of this paper to cite some fresh evidence which nominates the intermediate aldehydes as the most generally important branching agents, and to show how the oxidation of hydrocarbons can be represented as an attack on the hydrocarbon molecules by hydroxpl radicals. It is thought that the mechanism proposed, which does not exclude the participation of intermediate peroxides at low temperatures, gives a logical interpretation of the phenomena with reasonable economy of hypothesis.The Concentration of 1ntermediates.-If AT is the average life of the molecules of the intermediate substance in seconds, and Y is the maximum rate of reaction of the hydrocarbon in mm. Hglsec., then the maximum pressure reached by the intermediate during the reaction is YAT mm., and the concentration of the intermediate must be greatest at the time of maximum rate. Now the rate of a branched chain reaction can be shown to conform to the relationship Y = A(e4t - I) and when 9 the net branching factor is greater than zero the reaction accelerates either towards ignition or degenerate explosion.In the latter case characteristic of hydrocarbon oxidation the time of development to maximum velocity is of the order of minutes and hours. And arguing on kinetic grounds, Semenoff has concluded that the average life of a chain centre before it enters into a branching reaction, is of the order of seconds or minutes. Thus for a value of AT of 60 sec., and a maximum reaction rate of 0-1 mm. of hydrocarbon oxidized per second, the maximum pressure of intermediate is YAT = 0.1 x 60 = 6 mm. oxidation indicates an intermediate responsible for delayed branching, In general it may be judged that the time scale of the rate of hydrocarbon sufficiently stable to reach pressures of the order millimetres when the reaction is proceeding isothermally at an ordinary measurable velocity.Semenoff, Chemical Kinetics and Chain Reactions (rg35), p. 68. 269270 COMBUSTION OF HYDROCARBONS With both methane and ethylene the pressure of formaldehyde pro- duced as intermediate reaches values of this order, but the concentration of hydroperoxide, if formed, is too small to be measured. Upon the basis of this calculation it is not present in sufficient concentration to allow it to function as the essential agent in the degenerate branching process. That formaldehyde IS a in these two cases is the functioning intermediate is shown also by the fact that its addition to the system reduces the period of acceleration (induction period) and in sufficient quantity completely eliminates it. Further, Dr.Harding' has found that if more formalde- hyde be added than is necessary completely to suppress the induction period, the reactions will start immediately at an enhanced rate, which will rapidly decline to the normal rate as the pressure of formaldehyde adjusts itself to the value corresponding to the stationary state proper to the reaction parameters. In such a case we mag speak of a negative induction period. The Effect of Light on the Oxidation of Methane and Ethylene.- These considerations lead to the conclusion that any influence favouring the dissociation of the formaldehyde into free radicals should have the effect of increasing the net branching factor with a consequent reduction of the time of development of the reaction and an acceleration of the maximum rate.This effect was achieved by Dr. D. Patnaik6 and the author, using a very powerful beam of ultra-violet light of wavelength A 3800-2400 A which is known to dissociate formaldehyde into H atoms and carbon monoxide. The reactions (methane at 484-5' C and ethylene at 431.5' C) responded immediately to the effect of the light, accelerating when the light was admitted and decelerating when it was excluded, and the induction period was reduced. In cases near the thermal ignition limit the slow reaction could indeed be converted to ignition by irradiation. A further exam- ination of the kinetics both in the light and the dark showed that the effect of the light was merely to augment the thermal process without altering the course of its kinetics, and that while the rate RD of the dark reactions is represented by the expression the rate in the light R, is given by where (Hy) is the pressure of the hydrocarbon, I the intensity of the light, and P the total pressure in the system, the diameter and surface being kept constant. If, therefore, we subtract the rate in the dark from that in the light we see that the augmented reaction due to irradiation follows a similar law to the dark reaction, with the intensity of light re- placing the oxygen pressure.The kinetic results are explained in detail if the delayed branching due to the formaldehyde originates from its thermal oxidation by the reaction * RD = K W Y ) (0,) + KAHy) * ( O P B, = K,(HY)(02) + K,'(HY)I + K2(HY)2(OIP + KawYT)ZIP R, - RD = KI'(Hy)I + K,'(Hy)'IP /OOH 0 HCHO+O,= [ HC ,o ] =HCCoH+ 0, AH-0, Bone and Gardner, Pvoc.Roy. SOC. A , 1936, 154, 297. 9 Foord and Norrish, Proc. Roy. Soc. A , 1936, 157, 503. 4 Harding (not yet published). 5 Norrish and Patnaik, Nature, 1949, 163, 883. * In what follows, the reaction of an aldehyde with oxygen t o form oxygen atoms will frequently be postulated. It is believed that the process is associated with the formation of a peracid which is known to occur a t low temperatures, possibly by a chain reaction involving RC radicals, as formulated by Ubbelohde 80 and by McDowell and Thomasa1 /oo- N OR. G. W. NORRISH and in the photochemical reaction from H2C0 + hv = H + HCO \ H + CO. The effect also persists in the spectral region h 3500-3800 A where the formaldehyde does not yield H atoms directly and must be ascribed to some form of induced predissociation, such as H2C0 + h u = H2C09 ‘€€,COX + X = H + HCO + X which has been already postulated by Henri to set in at elevated tem- peratures. Very similar results have been obtained by Dr.Garbatski and the author for the photochemical oxidation of ethane. It must be stressed that we have, in all the above cases, influenced thermal gas reactions photochemically in spite of the fact that none of the primary reactants show the slightest absorption of the light used. The photochemical effect is exerted through the intermediate product, a result which in itself provides confirmation of the kinetic conception of delayed branching. It is further to be noted that hydroperoxide and hydrogen peroxide do not absorb light of wavelength longer than 3000 while the photochemical effect persists up to 3800 A The kinetic requirements of both the thermal and photochemical oxidation of methane and ethylene are satisfied by a simple extension of the scheme advanced for these reactions by the author.6 For methane this takes the form Initiation and degenerate branching CH4 + O2 = CH20 + H,O H2CO + 0, = H,CO, + 0 0 + CH4 = CH, + OH OH + CH4 = CH, + H,O CH, + 0% = H2C0 + OH propagation 1 coupled formalde- H + 0, + H,CO = CO + H,O + OH CH@ + OH = CO + H,O + H H + 0, + X = HO, + X’ OH + surface + termination termination * 1 photochemical H2C0 + X + hu = H + H + CO + X } branching- H&O + hv = H + H + CO It may be said to be an academic question whether the CH,00 radical takes part.and the radical may have a transitory existence, but at the temperatures At higher temperatures of hydrocarbon oxidation the transient peracid is assumed to eliminate an oxygen atom Reaction (5) might be written : CH, + 0, --+ CH,COO H,CO + OH RC/ooH) -+ \O The net overall thermal value of the reaction RCHO -!- 0, = RCOOH + 0 is zero, and in all cases where the latter reaction is written an intermediate tiansition process of the above type is to be undeistood. 6 Norrish, Colloquium on Reactions of Inflammation and Combustion in Gases (C.N.R.S.) (Pans), 1948, 16, or Revue de L’Inst. France de Petrole, 1949, 4, 288. * Reactions (7) and (8) are only formally to be regarded as ternary reactions. The transition complex HO, formed with ca. 50 kcal. is unstable unless deactivated (reaction (8)) and reactive only if it encounters a formaldehyde molecule (re- action (7)) before it decomposes or is deactivated.272 COMBUSTION OF HYDROCARBONS of the oxidation of methane there is no evidence of the reaction for no trace of methyl hydroperoxide can be detected. Argument from Effect of Temperature on Velocity of Reaction.- Further kinetic analysis of the process of degenerate branching shows that the overall temperature dependence of the reaction, when self heating is excluded, is mainly dependent on the energy of activation of the branch- ing reaction furnished by the intermediate.Dr. Harding' and the author have shown that for ethylene the value of the energy of activation varies in a parallel way with that of formaldehyde from 2 6 kcal.at 400' C to 53 kcal. at 550" C. They conclude that two processes of degenerate branching are involved, the reaction predominant at lower temperatures giving place to H2C0 + O2 = HO, + HCO, AH =,- 50. It may be, however, that traces of olefinic peroxide are formed with ethylene at the lower temperatures (though none was ever detected in our work in spite of careful search) for there sometimes occurred a small drop in pressure during the induction period. C2HI + 0, = C2HI02 = &H,O is not excluded. The Oxidation of Higher Params.-The ignition phenomena of higher paraffins, from propane upwards and possibly ethane, fall into two groups : the higher temperature ignitions in the region above 450' C, and the cool flame phenomena in the temperature range between ca.260-400' C. This is illustrated in Fig. I, due to Newitt and Thornes.* The Higher Temperature Ignition.-The kinetics associated with the higher temperature ignition appear to follow closely the pattern of those of methane, both as regards the time scale and intermediates. They have all the characteristics of delayed branching reactions and with propane at pressures below atmospheric outside the cool flame region, Mr. Galvin finds similar dependence of the reaction rate at 395" C on partial pressures, surface and total pressure, as is found for methane, and a similar effect of irradiation. Aldehydes are easily detected at concentrations consistent with their functioning as the essential intermediate while peroxides, Q at these higher temperatures, are unstable and, if formed, their concentration would appear to be too low to satisfy the kinetic requirement of degenerate branching.Thus the high temperature oxidation of higher paraffins can be represented by a similar mechanism to that of methane, modified by reactions to account for the observed degradative oxidation of the alde- hydes to formaldehyde and the final products, carbon monoxide and water. We may write : CH,COO + CHI = CH,OOH + CH, H,CO + 0, = H2C02 + 0, AH = 0, A reaction such a s RCH,CHO + 0, = RCH,COOH + 0 Initiation and branching RCH2CH3 + 0, = RCH,CHO + H20 RCH,CH, + 0 = RCH,CH, + OH propagation RCH,CH, + 0, = RCH,CHO + OH RCH,CH3 + OH = RCH,CH, + H,O RCH,CHO, + O H = RCH, + CO + H,O , c o ~ ~ ~ d ~ RCH, + 0, = RCHO + OH gradation of aldehyde.Termination may take place a t the surface or in the gas phase by the recombination of propagating radicals. In addition, alcohols may be 7 Harding and Nomsh, Nutwe, 1949, 163, 797. 8 Newitt and Thornes, J . C!;enz. SOC., 1937, 1656. Q Neumann, A d a Physzcochzm., 1938, 9, 527.R. G. W. NORRISH 27 3 expected to result from an attack on the secondary hydrogen atoms of the hydrocarbon by hydroxyl radicals, e.g., RCH,CH3 + OH = RCHCH, + H,O RCHCH, + 0, = RCHO + CH30 CH30 + RCH,CH3 = CH,OH + RCH,CH, and Mr. Knox, in agreement with Newitt and Thornes,8 has recently found measurable quantities of alcohols in the oxidation of propane at 390° C. In addition, Mr. Galvin has been able strongly to accelerate the oxidation of propane at 3 8 1 O C (in regions of pressure and temperature outside the cool flame zone) by irradiation with light of wavelength 3000-3800 A, a part of the spectrum which is not absorbed by peroxides.CURVE POR ICNlTlON OF CJ Ha + 0, MIXTURES. AFTER NEWlTT AM THORNES. 5 Cool Flames LOW Temperature 2 Codflamas Ignltion ReTim. I I I I 200 400 600 PRESSURE rn m. 9. FIG. I . There is thus evidence that the kinetics of the high temperature oxidation of the higher hydrocarbons conform to those of the oxidation of methane and it is in accordance with the arguments advanced here to conclude that formaldehyde which is the final product of the oxidative degradation of aldehydes and which, owing to its greater stability accumu- lates to intermediate concentrations much greater than the higher alde- hydes, begins to function as a branching agent only at the higher tem- peratures.Thus, as the temperature rises toward the upper ignition limit, ,$ will increase due to the participation of the formaldehyde in the branching In this way we can explain the fact that when formaldehyde is added to a mixture of propane and oxygen it sensitizes the high temperature ignition, but does not affect the low temperature oxidation below 350° C. There does not appear to be any reason for looking further than the alde- hydes, and in particular formaldehyde, as the essential intermediates leading to branching in the reaction zone approaching the high temperature ignition. The Lower Temperature Ignition Phenomena.-The cool flame phe- nomena are illustrated by the curve for propane air mixtures, due to Newitt and Thornes (Fig. I).Between 400° and 270° C the temperature- pressure curve of ignition is associated with a region extending on the low reaction : RCHO + 0, = RC0,H + 0.274 COMBUSTION OF HYDROCARBONS pressure side of ignition within the bounds of which intense reaction associated with cool flames is observed. In closed vessels the cool flame may be present during the whole reaction, which follows the normal sinusoidal curve of pressure against time, or under other conditions of partial or total pressure, one or more flames may successively traverse the reacting mixture. Under such conditions the passage of each flame is associated with a sharp pressure pulse of very short duration imposed upon the normal pressure time curve.The pressure and temperature limits within which Newitt and Thornes observed one, two or more flames are shown in Fig. (I). Outside the cold flame region the reaction is much slower both at high and low temperature but it still follows the normal sinusoidal pressure-time curve. The temperature limits within which the cool flames occur are much the same for all the higher hydrocarbons studied, namely, about z60° to 400° C. Similar observations are also characteristic of aliphatic ethers, ketones and aldehydes. This in itself would suggest a common origin of the phenomenon and the conclusion is borne out by the fact that the spectra of all cool flames which have been observed lo, 11-acetaldehyde, diethylether, propionaldehyde and hexane- are identical with the fluorescence spectrum of formaldehyde.The fact that cool flames of this specific nature are characteristic of the oxidation of aldehydes and that aldehydes are formed in profusion as common inter- mediates in the oxidation of the paraffins and other ketones in the cool flame zone suggests that their origin is to be sought in some reaction derived from the aldehyde itself. That this reaction is a side reaction and not part of the main course of oxidation is shown by the fact that only one quantum of radiation per 1oS molecules of aldehyde reacted is emitted ; but whatever it is it must be a reaction capable of yielding sufficient energy to provide excited molecules of formaldehyde. The spectrum emitted required a maximum excitation of the order of 95 kcal. It has been suggested by the author that this is to be found in the reaction but it is now clear that there is serious objection to this.In a static system of oxidizing acetaldehyde, Newitt, Baxt and Kelkar la showed that multiple cool flames can follow one another through the mixture, and it is not easy to see how the above reaction could give rise to such periodic phenomena. It would appear that some substance derived from the alde- hyde is alternatively built up to a limiting concentration and exploded, and Neumann has offered evidence, in the case of higher hydrocarbons, that this may be a hydroperoxide or dialkyl peroxide. Recently Bardwell and Hinshelwood,13 studying the cool flame of butanone oxidation, have obtained similar periodic cool flames and observed a fluctuation of per- oxide in the medium corresponding to the passage of each flame.They further found that the induction period preceding the cool flames was greatly reduced by the addition of a small quantity of acetaldehyde, while it is unaffected by formaldehyde. In the scheme of oxidation described above for the high pressure oxidation the aldehyde has been indicated as the source of alkyl radicals in hydrocarbon combustion. Similar processes could occur in acetal- dehyde oxidation itself, e.g. the scheme : RCH2CH0 + 0 = RCHO + H,CO, AH = - 74, 3 CH, lo + o /OOH CH,CHO + 0, --f CH,C 1 0 \OH CH,CHO + 0 -+ CH,CO + OH CH,CHO + OH -+ CH, + CO + H,O CH,CO + CH, + CO lo Emeleus, J . Chem. SOL, 1926, 2948 ; 1929, 1733. l1 Kondrat’ev, 2. Physik, 1930, 63, 322. lZNewitt, Baxt and Kelkar, J Chem.SOC., 1939, 1703. 13 Bardwell and IIinshelwood, Proc. Roy. SOC. A , 1951, 205, 375R. G. FV. NORRISH 275 leads to methyl radicals. radical can have but a transient existence, owing to its instability, i.e., as is indicated by its non-appearance in the methane oxidation, there can be little doubt about its existence at low temperatures, for-hydroperoxides are built up in measurable quantities, in the cold flame region. Thus the aldehyde may be the source of peroxide formation through the alkyl radicals it generates in hydrocarbon combustion, in its own combustion, and in the combustion of ethers and other bodies in which it is found, as an intermediate. We have the possibility of the reaction of the hydrocarbon radical with oxygen taking two courses Now while at high temperature the peroxide CH, + 0, = [CHaOJ + HCHO + OH RCHO + OH RCH,OO RCH, + 0 2 < the peroxide radical coming more into the picture at lower temperatures.Two courses of propagation in hydrocarbon systems are then opened up, namely, OH + RCH, = RCHa + Ha0 ROO + RCH, = ROOH + RCH,. If, therefore, some reaction of the accumulating peroxide, subject to ignition limits, could give rise to excited formaldehyde, as suggested by Harris and Egerton,14 the periodic cool flames in aldehyde hydrocarbon and other similar oxidizing media find ready explanation, but such a conclusion must be subject to the result of experiments designed to measure the spectra emitted by the explosive, decomposition and oxida- tion of hydroperoxides. The facts require that one or both should be identical with the fluorescence of formaldehyde.It is significant that the conclusions of Bardwell and Hinshelwood, that the cool flame of butanone oxidation is only a secondary process and unconnected with the main course of oxidation, are in agreement with the conclusions to be drawn from Townend’s intensity measurements for acetaldehyde.16 It is apparent from the results of Newitt and Thornes, for propane, and Bardwell and Hinshelwood, for butanone, that intermittent or periodic cool flames are associated with a parallel periodicity in the self-heating of the oxidizing medium. Each interval between cool flames is a period of exponential acceleration of the main oxidation reaction, with consequent self-heating. Towards the end of the process the peroxide which has been accumulating reaches its ignition limit and explodes with the emission of light.Thereafter the main oxidation process, if it does not pass to com- plete thermal ignition, suddenly decelerates, the self-heating is arrested and the temperature falls, The whole process of acceleration of the reaction to the next phase of self-heating and associated cool flame is then repeated. This periodic self-heating of the medium may be explained by a fluctu- ation in the branching factor associated with a small periodic variation in the concentration of the higher aldehydes, and peroxides, in accordance with the general principle suggested by Salnikov.16 By hypothesis the aldehyde is subject to two processes of oxidation-ne leading to branching the other not, and mechanisms for these two processes have been postulated above.If the latter, which is to be identified with the oxidative degrada- tion of the higher aldehydes to formaldehyde, increases more rapidly than the former with rise of temperature, the stationary concentration of higher 1* Harris and Egerton, Nature, 1938,412, S30 ; Pvoc. Roy. Sac. A , 1938, 168, I. ’5 Topps and Townend, Trans. Favaduy Soc., 1948.42, 345. 16 Salnikov, Compt. vend. U.R.S.S., 1948. 60, 405.276 COIMBUSTION O F HYDROCARBONS aldehyde will be diminished and the branching factor may fall.* It is known that in the cool flame zone the stationary concentration of aldehyde is higher, and the reaction much more intense than in the region of slow reaction outside.Any process of self-heating can therefore, if of sufficient magnitude, carry the system outside the cool flame region to the region of slow reaction and no self-heating where the branching factor is very small. With the suppression of self-heating the temperature of the system will very rapidly fall to that characteristic of the cool flame region. There will be a short period of induction and acceleration while the aldehyde and peroxide concentration is re-established, followed by self-heating and the next cool flame. The fluctuation of aldehyde pressure during the periodic process need not be great for the branching factor is very sensitive to small changes in the concentration of intermediate. It is understandable that such fluctuations could be within the experimental error of Bardwell and Hinshelwood, though with the slower reaction of Newitt and Thornes they were observable. That the periodic fluctuations in aldehyde concentration are indeed small is in accord with the long period of induction before the first flame, and the short interval between successive flames.The first may be said to represent the time for the branching process to build up the critical aldehyde concentration from zero, while the second represents the time required to make good the small deficiency of aldehyde due to the thermal fluctuation. In cases where there is one continuous cool flame, it is to be concluded that the self-heating is insufficient to take the reaction system outside the cool flame zone. We therefore have a continuous glow emitted by the peroxide which is all the time maintained at its ignition limit.The facts are further consistent with the conclusion that final ignition of the system in the region of cool flames is conditioned by a thermal ignition of the accumulated aldehyde. A striking observation made by Mr. Galvin and the author may be mentioned here : if a mixture consisting of 160 mm. of propane and 160 mm. of oxygen is reacted at 281O C-just inside the cool flame region-the in- duction period of 5 hr. is progressively cut down to zero by admixture of acetaldehyde up to 0.86 yo. The same induction period, of 5 hr., is also reduced to 20 min. by strong irradiation with ultra-violet light. Similar results were obtained by Townend and Chamberlain 17 who were able to induce a lower ignition peninsula in the oxidation of ethane, and to reduce the time lags to ignition in the same region from hours to seconds.The conclusion would seem to be reasonable, that it is the aldehyde whose build-up is accelerated by irradiation and which functions as the essential intermediate. Thus we conclude that the formation of peroxide is secondary to the production of aldehyde. There is nothing, however, in the above conception which would ex- clude the accumulated peroxide as a source of branching at these low tem- peratures. Such a reaction as could be superimposed on the aldehyde mechanism without any material alteration of the kinetics, but since the stability of the hydroperoxide rapidly decreases with rise of temperature the contribution of this reaction to branching will rapidly wane in the same direction. The above reaction may also, through the HO,, be the source of the hydrogen peroxide which Dr.Bailey and the author have found to be * The branching factor will in fact pass through a minimum as the temperature rises, since the formaldehyde, which all observations show to be comparatively stable and not contributing to the branching proceys a t low temperatures of the cool flame zone, begins to function in this capacity as the high temperature ig- nition region is approached (methane oxidation). l8 Bailey and Norrish (not yet published). ROOH + 0, = RO, + HO, Townend and Chamberlain, Proc. Roy. SOG. A , 1936, 154, 95.R. G. W. NORRISH 277 the main peroxide isolated in the products of the combustion of hexane in the cool flame region.In this work we have made an extensive ana- lytical study of the products formed in the cold flame of hexane by a flow method in which the intensity of the cold ffame could be measured. The results are in accord with the " aldehyde " scheme outlined above, modified by the production of derivatives of hydroperoxide or peroxide radical (cyclic ethers, etc.) and consistent with attack on the hexane chain at secondary as well as primary hydrogen atoms. Thus, while all the alde- hydes from caproic aldehyde to formaldehyde have been identified as required by the aldehyde scheme, there is evidence, in addition, of such reactions as CH,CHCH,CH,CH, -+ CH,O + CH,(CH,) ,CHO CH ,OH 00 I '\ which lead to the production of alcohols, and of reactions of cracking, and disproportionation of hydrocarbon radicals which yield considerable quan- tities of olefines.The cold flame itself, at atmospheric pressures, may, under certain circumstances, become associated with a second more intense blue flame. This second flame establishes itself in the products of the first flame as the mixture is made richer in oxygen. Similar two-stage cold flames have been observed by Townend and his co-workers 18 in the oxidation of acetaldehyde and ethers, and by Newitt and Thornes 8 with propane, and it has been shown that both parts emit the spectrum of formaldehyde. We have been able to show that the intensity of emission of the first cold flame increases as the mixture becomes richer in oxygen and that simultaneously the production of aldehydes and peroxide in- creases also.But when the second flame forms the concentrations of aldehydes, peroxides and other condensable products decrease rapidly with increasing intensity of light emission, giving place to carbon monoxide and water and, in addition, a much increased yield of olefines. There seems to be little doubt that the second flame accompanies the combustion of aldehydes and other intermediates formed in the main reactions that proceed parallel with that of the first flame. Its temperature is about 430°C but it gets higher as the oxygen content is increased and may finally give place to the white flame of ignition. It is possible that at these higher temperatures its origin may be found in the direct oxidation of aldehydes by the reaction previously proposed in any case it is a stage which precedes the thermal ignition of the whole system. The results of our work on the products from the cool flame of hexane have led us to the conclusion that both the aldehyde and peroxide mechan- isms are functioning at the same time in this region. Relative Oxidizability of Hydrocarbons .-Finally we may consider the interesting facts concerning the relative rates of oxidation of hydro- carbons as we ascend the homologous series, to which attention was first drawn by Cullis and Hinshelwood.22 The fact that other things being made equal, the rate of oxidation of hydrocarbons increases towards a maximum as we ascend the homologous series can be readily understood when i t is realized that the secondary hydrogen atoms are more reactive than the primary. Thus as we ascend the series from ethane, the ratio of " primary " to I ' secondary " characteristics progressively falls towards RCH,CHO + 0 = RCHO + H,CO, AH = - 74 ; Maccormac and Townend, J . Ckem. Soc., 1940, 143. 2 O Ubbelohde, Proc. Roy. Soc. A , 1936, 152, 355. *1 McDowell and Thomas, J . Chem. Soc., 1949, 2205, 2217. 22 Cullis and Hinshelwood, Faraduy Soc. Discussions, 1947, 2, 117.27s ORGANIC PEROXIDES zero and since the primary hydrogen is less easily attackable than the secondary, we shall pass from a rate solely defined by primary hydro- gens in ethane, to one solely defined by secondary hydrogens in the case of an infinitely long hydrocarbon. Thus we should expect the oxidiza- bility per carbon atom to increase to a maximum as the length of the molecule increases. Department of Physical Chemistry, University of Cambridge, Free School Lane, Cambridge.

ISSN:0366-9033    

DOI:10.1039/DF9511000269

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

ORGANIC PEROXIDES SOME PROPERTIES OF ORGANIC PEROXIDES BY A. C. EGERTON, W. EMTE AND G. J. MINKOFF Received 22nd February, 1951 A number of alkyl bydro-, di-alkyl and acyl hydrb-peroxides have been prepared, and several of their physical properties have been studied. The refractive indices and melting points are recorded and compared with values in the literature. The vapour pressures have been measured in a static system, and the latent heats of vaporization deduced from the data. The heats of com- bustion of the peroxides were measured in a bomb calorimeter, leading t o values of the heats of formation. Comparisons are made between the heats of formation of R . 00 . R' and of R . 0 . R' ; in general, differences in heats of formation are 35-50 kcal., though acyl hydro-peroxides give results more difficult to interpret.Investigations have been made in this laboratory relating to the slow, and rapid, combustion of hydrocarbons, e.g. by Bone and Hil1,l Newitt and Thornes,S G a y d ~ n , ~ Egerton and Minkoff,4 and Egerton, Harris and Young.6 That the peroxides play a role of importance in combustion, particularly in relation to the phenomenon of " knock ID, has long been recognized (Egerton, Callendar (1927)). Investigation of their decom- position was therefore made by Harris and Egerton.6 I n spite of these investigations and many others elsewhere, there is still doubt as to the particular peroxide which is formed under different circumstances during hydrocarbon combustion. The object of the present investigation was to gain familiarity with some of the different types of simpler peroxides, and to determine some of their properties.The following have been investigated : alkyl hydroperoxides, dialkyl peroxides, peracids, monohydroxy-alkylhydroperoxides, dihydroxy-dial- kyl peroxides and monohydroxy dialkyl pkroxides. Methods of preparation have been described by Baeyer and Villiger,' by Rieche and Hitz 8 and by D'Ans and Frey ; they have been further applied by Harris and Egerton,6 Harris,lO Milas and Surgenor and particularly by Eggersgluss and Schutt (1941). The latter work was Bone and Hill, Proc. Roy. SOC. A , 1930, 129, 434. Gaydon, Proc. Roy. SOC. A , 1942, 179, 435. Egerton and Minkoff, Proc. Roy. SOC. A , 1947, 191, 145. a Newitt and Thornes, J. Ckem. SOC., 1937, 1656. 6 Egerton, Harris and Young, Trans.Faraday SOC., 1948, 44, 745. 6 Egerton and Harris, PYOC. Roy. SOC. A , 1938, 168, I. 7 Baeyer and Villiger, Bey. B, 190. 33, 2479, etc. 8 Rieche and Hitz, Ber. B, 62, 2458, etc. D'Ans and Frey, Ber. B, 1912, 45, 1848. loHarris, Proc. Roy. SOC. A , 1939, 173, 126. 11 Milas and Suigenor, J Amer. Chem. SOC., 1946, 68, 205.A. C. EGERTON, W. EMTE AND G. J. MINKOFF 279 not published, but an account of i t was available and one of the present authors (W. E.) took part in the investigation. The methods of prepara- tion will not be described in detail in this communication. The preparation of the peroxides in a state of purity is still difficult. The properties which have been measured are the vapour pressures, refractivities, melting points and heats of combustion. The properties of hydroxy-peroxides will be considered in a later paper. Experimental The Vapour Pressures, Melting Points and Refractivities .-The static method for measurement of vapour pressure was chosen.It was possible t o use a differential manometer containing mercury for di-tert.-butyl peroxide, which is comparatively stable, and even, with certain precautions, for diethyl- peroxide, but most peroxides attack mercury and a spoon gauge F (see Fig. I ) , was used with all the peroxides except di-terl.-butyl peroxide. The peroxide was placed in B and frozen with liquid air : the apparatus was then evacuated. After closing tap TI, the peroxide was distilled from B t o C and back from C to B, the middle fractions being rltained and the rest pumped off.Finally the peroxide was distilled into A and, after evacuating the spoon gauge with A cooled in liquid air, the apparatus sealed off a t D. Counter pressure was pro- vided from nitrogen in bulb E to bring the pointer back t o zero (using a mirror in conjunction with a Pointolite lamp) : 4 scale divisions were equivalent to I mm. Hg pressure difference. The vapour pIessure could be measured on the manometer to 0-1 mm. Hg. Observations were made a t six t o ten different temperatures, the mean of five a t each temperature being plotted ; the separate determinations agreed to about I yo. Up t o room temperature A was main- tained a t the required temperature by immersing the container in cold alcohol and above room temperatlie in water, both baths being weil stirred.The calibrated mercury thermometers could be read t o 0.01' C. Leakage or de- composition could be detected by freezing the peroxide and measuring the residual pressure : measurements were discarded when the presence of permanent gases was indicated. Materials .-METHYLHYDROPEROXIDE was made from dimethylsulphate and hydrogen peroxide (30 yo) ; the distillate was extracted with ether and frac- tionated.2 80 ORGANIC PEROXIDES ETHYL HYDROPEROXIDE was prepared similarly from diethylsulphate. i!ert.-BuTYL HYDROPEROXIDE was obtained by the action of 85 yo H,O, on tert.-butyl alcohol in presence of anhydrous magnesium sulphate. After remcval of alcohol, the residue was fractionated and then redistilled. DIETHYLPEROXIDE was prepared from diethyl sulpbate and 30 yo hydrogen peroxide ; the oily layer was separared from that containing most of the ethyl hydrogen peroxide, distilled and fractionated.DI-fert.-BUTYL PEROXIDE was prepared from butyl hydrogen sulphate (ob- tained via tert.-butyl alcohol) and 30 yo hydrogen peroxide. The oily layer was fractionated and redistilled. PERACIDS (acetyl hydroperoxide, propionyl hydroperoxide and butyryl- hydroperoxide) were prepared by perhydrolysis of the diboro-carboxylic acids, followed by fractionation. Some samples were also prepared from the acid anhydrides and hydrogen peroxide, but the product was generally less pure by this method. The impurities were a small peicentage of tbe corresponding acid and occasionally some ethyl ether. HYDROGEN I?EROXIDE.-T~~S was a redistillate cf a IOO yo H,O, obtained from Messrs.Laporte. Results The results of the vapour pressure measurements were plotted as log p against I/T and the equations of the straight lines obtained are given in Table I, together with the value of the latent heat of vaporization within the range of the measurements. TABLE I.-VAPOUR PRESSURES Peroxide Methyl hydroperoxide . Ethyl hydroperoxide . tart.-Butyl hydroperoxide . Diethyl peroxide . Di-tart.-butyl peroxide . Acetyl hydroperoxide . Propionyl hydroperoxide . Butyryl hydroperoxide . Hydrogen peroxide . iange of Measurements OC - 20 t o + 40 - 20 t o + go 0 to I20 - 20 to + 60 - 20 t o + I00 0 t o I10 0 to I20 0 to I20 o t o 60 L(ca1.) 9020 10190 11130 6930 7420 105 70 10320 10860 I I 300 8.38 -1972/T 8.834-2228JT 8.891-2432JT 7-140-1621 JT 8-91 I - 23 I I/T 8.623 - 2256 JT 7.356- 1517/T 8.83 -2376/T 8-63 -2469/T The slopes of the lines are approximately the same except for those of the dialkyl peroxides which are less steep.The Trouton constant (LIT,) is about normal for the dialkyl peroxides, but is high for the others. At the higher temperatures, above the range over which the equation holds, some decomposi- tion was observed with methyl hydroperoxide and hydrogen peroxide, and t o a less extent with ethyl hydroperoxide. The refractivities of these peroxides determined with a Hilger-AbbB refracto- meter are recorded in Table 11. The values in column 4 were obtained, using different preparations, subsequent to the vapour pressure determinations by one of the authors (G.J. M.). The values given in column 5 are quoted from the literature. The melting points were also determined with a thermometer (Reichenstalt calibration) and a thermocouple. Considering the difficulty of obtaining the peroxides in a pure state, there is fair agreement between the values. The freezing point of methyl hydroperoxide is higher than that given by Rieche. The vapour pressure of this peroxide also appears t o be considerably @eater than the value of 65 mm. at 3g0 C quoted by the same author. The vapour pressure of 2ert.-butyl hydroperoxide is rather lower than 15 mm. a t 37" C quoted elsewhere. The peracids are particularly difficult to obtain in a pure state and there is some doubt as to the accuracy of the values given, although they agree reasonably well with other determinations.The acetyl hydroperoxide was made by two quite different methods, and the refractivities agreed to 0.2 yo ; nevertheless, the acetic acid content was different. (Iodometry is of little use for estimation of the purity of the peroxides.) That the slopes or the vapour pressure lines are of the expected magnitude is shownA. C. EGERTON, W. EMTE AND G. J. MINKOFF 281 by comparison with those given for (a) ethyl alcohol, (b) diethyl ether and (c) tert.-butyl alcohol, viz., (a) log P = 9.205 - 2216/T ; L = 10140 cal./mole. (b) log P = 7-92 - 1543/T ; L = 7059 cal./mole. (c) log P = 9-53 - 2361/T ; L = 10800 cal./mole. TABLE II.-REFRACTIVITIES Peroxide Methyl hydroperoxide . Ethyl hydroperoxide . tert.-butyl hydroperoxide .Diethyl peroxide . Di-tert.-butyl peroxide . Acetyl hydroperoxide . Propionyl hydroperoxide . Butyryl hydroperoxide . Hydrogen peroxide . Peroxide ,I 5 O D 1.3646 1.3832 1.4027 1.3712 1.3908 1'3994 1.4101 1.4148 1.4086 ,209 D 1.3608 1.3801 1.4007 1.3888 1'3974 1.4041 1'4125 1.4058 - R = determination by Reiche. TABLE III.-MELTING POINTS ,200 D 1.3648 I 1'4009 1.3698 1'3905 1.4022 1.4057 - - Methyl hydroperoxide Ethyl hydroperoxide . tert.-Butyl hydroperoxide Diethyl peroxide. . Di-tert. -butyl per oxide Acetyl hydroperoxide Propionyl hydroperoxide Butyryl hydroperoxide Hydrogen peroxide . Melting Points C -72 to -78 - 8 t o - I 0 - 68 to - 69.5 - 18 t o - 19'5 glass 0'2 - - 13-1 - 10'0 - 1.8 Literature Value Literature Value OC - I00 t o 100.5 - I00 t o 100.5 - 13'5 - 70 - 18 - 13'3 - 10.3 - 1.7 0'1 The latent heat change due to the introduction of the extra oxygen atom in the molecule is quite small, but the vapour pressures of the peroxides are, however, lower, except for the peracids where the position is reversed.The boiling points in O C estimated by extrapolating the plots t o log p = 2.8808 are as follows : Methyl hydroperoxide . . 86 Methyl alcohol . (64'5) Ethyl hydroperoxide . 95 Ethyl alcohol . - 77 (78.3) Diethyl peroxide . . 66 Diethyl ether . * 33'5 (34'6) terl.-Butyl hydroperoxide . 133 ten!.-Butyl alcohol . 82 (82.5) Di-tert.-butyl peroxide . . 108 Di-tert.-butyl ether . Acetyl hydroperoxide . . IIO Acetic acid . . 118.5 Propionyl hydroperoxide . 120 Propionic acid . . 140 Butyryl hydroperoxide .. 126 Butyric acid . . 162-3 Hydrogen peroxide . - I57 The figures in parentheses indicate the true b.p (literature values). The values obtained from measurements of vapour pressures and extrapolation would be expected io give somewhat lower values than the true boiling points, but they are probably not more than about a degree lower. Decomposition prevents their accurate determination. The value obtained for diethyl per- oxide by Rieche, 64-5O a t 760 mm., is lower than the above estimate. The value found by direct determination of the comparatively stable di-tevt.-butyl peroxide was I O ~ O C .282 OXIDATION OF TRIMETHYLETHYLENE Heats of Combustion.-A number of determinations of the heats of combus- tion of the above mentioned peroxides have been made during the course of this work by the bomb calorimeter method and the authors are indebted to the Director of the Fuel Research Station for permission to use the apparatus a t the Station for the purpose. The results, however, show discrepancies, presumably due to the difficulty of ensuring complete combustion, and further determinations are in hand. The difference between the heats of formation of water vapour and of hydrogen peroxide vapour is 24-2 kcal. The difference, however, between the heats of formation from the atoms, of hydrogen peroxide and of water vapour is 34.8 kcal./mole. This is the energy difference due to the -0-0- group instead of the -0- in the molecule and t o the consequent change in the strength of the OH bands. One of us (G. J. M.) wishes to express his gratitude to the University of London for the award of an I.C.I. Research Fellowship. We are grateful to the Director of the Fuel Research Station, Greenwich, for permission to use the bomb calorimeter. Department of Chemical Engineering and Applied Physical Chemistry, Imperial College, s. w.7.

ISSN:0366-9033    

DOI:10.1039/DF9511000278

出版商:RSC    

年代:1951

数据来源: RSC