摘要:

282 OXIDATION OF TRIMETHYLETHYLENE THE CATALYZED OXIDATION OF TRIMETHYL- ETHYLENE IN SOLUTION BY C. E. H. BAWN, A. A. PENNINGTON AND C. F. H. TIPPER Received 1 s t Mavch, 1951 The oxidation of trimethylethylene by molecular oxygen occurs very slowly at low temperature. Soluble cobalt salts are effective catalysts for the reaction and the kinetics of oxidation (in the range 25-45"C) have been determined in glacial acetic acid which is a good solvent for both hy-drocarbon and catalyst. The oxidation rate obeys the law - dO,/dt ot [RHI2[Cat]* at catalyst concentration greater than 2.0 x I O - ~ molell. The initial product of the reaction is a hydroperoxide which forms an unsaturated aldehyde as a product of the reaction. At an early stage of the oxidation the hydroperoxide is present at an equilibrium concentration given by The mechanism of the reaction has been discussed in terms of the existing theories of the oxidation of unsaturated hydrocarbons.[ROOH] ot [RHIa. Recent studies of the interaction of molecular oxygen with unsaturated hydrocarbons have shown that the reactions exhibit certain common features in that (a) a hydroperoxide appears as the initial product of the reaction, and (b) the oxidation is a radical chain process, the rate of which may be accelerated by ultra-violet light, added peroxides or the peroxidic products of the oxidation and by certain heavy metal ions or their co- ordinated complexes. lBolland and Gee, Trans. Furaday SOC., 1946, 0, 230; Bolland, Trans. Faraday SOC., 1950, 46, 358; Idem, i b i d . , Quart. Rev., 1949, 3, I ; George, Rideal and Robertson, Proc.Roy. Soc. A , 1946, 185, 288, 308, 337 ; Bamford and Dewar, Proc. Roy. SOC. A , 1949, 188, 252 ; Robertson and Waters, J . Gem. Soc., 1948, 1574, 1578, 1585 ; T r w s . Faraday Soc., 1946, +?, 201.C. BAWN, A. PENNINGTON AND C. TIPPER 283 On account of solubility difficulties, measurements of the homogeneous heavy metal catalyzed oxidation has been limited to the metallic soaps, naphthenates and resinates, often of ill-defined composition, and in only a few instances have kinetic studies been made. The action of the heavy metal salt in catalyzing the oxidation is not clear although it is usually associated with its effect in promoting the formation and decomposition of the intermediary hydroperoxide.The present work deals with the kinetics of the cobalt ion catalyzed oxidation of trimethylethylene in acetic acid solution, a medium in which both the hydrocarbon and metallic salt were readily soluble. Experimental The course of the oxidation was followed by measuring the volume of oxygen absorbed a t constant pressure by trimethylethylene solutions in acetic acid. The concentrations of the hydrocarbon used were in the range 0.5 and 1-20 mole/l, and the catalyst concentration between I O - ~ and I O - ~ mole cobalt acetate per litre. The reaction vessel of approximately 50 ml. capacity was thermostated (.+ 0.03' C) and mechanically sbaken a t a rate of 7 times per sec. It was estab- lished that a threefold decrease in the shaking rate was without effect on the fastest rate of oxygen absorption and thus tbe measured values are real reaction rates.The apparatus and experimental method was similar to that previously used in measurements of the rate of oxidation of acetaldehydee3 Purification of Materials .-Amylene was stirred with its own volume of I : I concentrated sulphuric acid and water for several hours a t oo C. Under these conditions the trimethylethylene and z methyl-I-butene in the mixture dissolved to give tert.-amyl hydrogen sulphate and any pentene-e present was not attacked.* The acid layer was separated, and an equal volume of water added and the mixture ot trimethylethylene and z-methyl-I-butene distilled off. The mixture was dried and fractionally distilled under dry nitrogen through a zo-plate column.The middle fraction, b.p. 38-3 f 0.1' C at 760 mm., which was used for kinetic measurements, was stored under nitrogen and had a per- oxide content less than 0-01 %. Oxidation measurements of the product ob- tained from Kahlbaum and B.D.H. amylene agreed closely. Acetic acid and cobalt acetate were dried and purified as described by Bawn and Williamson.3 Both the acetic acid and cobalt acetate solutions used were free from aldehydes and peroxides and contained less than 0-01 % water as determined by the Karl Fischer method as modified by Smith, Bryant and Mit~hell.~ All reactants were stored in specially designed flasks with burette attachment under dry nitrogen. Oxygen taken from a gas cylinder was freed from water and carbon dioxide by passage through a trap cooled with liquid nitrogen.Cylinder nitrogen was purified from oxygen over heated Cu + CuO, followed by passage through a cooled trap. PEROXIDES were determined by the ferrous thiocyanate method of Young, Vogt and Neiuland.6 The ferric compound formed was determined photometrically, the photometer being calibrated by the use of hydrogen peroxide solutions. Reaction Products .-A complete analysis of the primary and secondary products of the oxidation has not yet been possible and attention so far has been directed t o the determination of the hydroperoxide, aldehydes, water and gases which are formed. The rate of hydroperoxide build-up is discussed later in this paper and it may be noted here that at catalyst concentrations greater than z x I O - ~ mole/l. the hydroperoxide rapidly reaches a stationary concentra- tion which corresponds to a small percentage of the total oxygen absorbed. Even in the initial phases of the oxidation a t the above catalyst concentration a relatively small percentage of the oxygen absorbed is present as hydroper- oxide.At much lower concentration of catalysts the hydroperoxide formation is very much greater and in the initial stages the oxygen absorbed corresponds to the almost exclusive production of hydroperoxide as observed in the uncatalj zed thermal oxidation of some mono-olefins, I : 4-dienes and hydroaromatics, and in the heavy metal catalyzed oxidation of tetralin,l and acetaldehyde a George, Rideal and Robertson, ref. (I) ; Robertson and Waters, ref. ( I ) . Bawn and Williamson, Trans.Furaday SOC. (in press). Norris and Joubert, J . Amer. Chem. SOC., 1927, 48, 873. J . Amer. Chem. SOC., 1939, 61, 2407. Ind. Eng. Chem. (Anal.), r936, 8, 198,284 OXIDATION OF TRIMETHYLETHYLENE and ben~aldehyde.~ This observation shows that the hydroperoxide is the primary product and the low hydroperoxide concentration a t high cobalt ion concentration is most probably a result of its rapid decomposition by the cobalt catalyst. An attempt was made t o isolate the hydroperoxide from the products of the room tempeIature uncatalyzed oxidation of 500 ml. of trimethyl- ethylene. By careful fractionation the initial hydrocarbon which contained 0.5 yo hydroperoxide was concentrated t o a colourless liquid containing 24 yo peroxide. This solution could be concentrated by distillation under pressure without decomposition and contained very small amounts of aldehydes and ketones , indicating that the peroxide was a relatively high boiling stable liquid.The exact structure of the hydroperoxide is unknown but as shown below one of the main products of the reaction is an unsaturated aldehyde, indicating a probable structure CH3\C_C<H CH/ CH,OOH The aldehydic products were separated from the reaction mixture as the 2 : 4-dinitrophenyl hydrazones. Chromatographic analysis of a benzene solu- tion using an alumina column gave two bands. The constituent of the main band on recrystallization from acetic acid had a m.p. 178.5' C and its absorption 260 300 340 380 I FIG. I. spectrum (Fig. I) showed maxima a t 3810, 2810 and 2540 A.These results are in very close agreement with data reported by Jones and Braude 7 for the 2 : 4-dinitrophenyl hydrazone of C H 3 \ G C / , viz., m.p. 179" C ; abs. max. CH/ \CHO 3810, 2810 and 2560 A respectively. In order t o ascertain if any gaseous products were formed the gases above the reaction mixture were analyzed a t various stages of the oxidations for carbon monoxide, carbon dioxide, hydrogen and the lower hydrocabons. In a typical experiment in which a solution containing 1.9 x IO-, mole trimethylethylene had absorbed 204 ml. (9.1 x I O - ~ mole) of oxygen the residual gases contained 3-5 yo and 1.0 yo of C 0 2 and CO respectively. At lower degrees of oxygen ab- sorption (30-40 mi.) no oxides of carbon were detected. These gases are thus products of an advanced stage of the oxidation and are formed in small amounts in any case.Titration with the Fischer reagent showed that water was produced in small amounts ; for instance, 15 ml. of a solution containing 1-26 mole/l. trimethyl- ethylene after absorption of 195 ml. 0, contained 5 x IO-* mole/l. of water corresponding t o 8-6 yo of the oxygen absorbed. The water concentration during the early stages of the oxidation was such that the inhibiting effect of water, discussed later, would be negligible. J , Chem. SOG., 1945, 498. H No hydrogen, methane or ethane were formed.C. BAWN, A. PENNINGTON AND C. TIPPER 285 Kinetic Measurements .-Trimethylethylene and its solutions in acetic acid are oxidized very slowly in air or oxygen in the temperature range 20-5oo C and samples exposed to air were found to contain less than 0.5 yo peroxide after one year.In the presence of a soluble cobalt salt (less than I mole yo of the hydrocarbon concentration) oxygen absorption occurs at a measurable rate above 25' C and the general form of the typical absorption curves is shown in Fig. 2. After a short period of acceleration the rate of absorption showed a linear de- pendence on time over a considerable part of the reaction and this was followed FIG. 2.-Rate of oxygen absorption with cobaltous and cobaltic acetates. by a gradual falling off in rate. The accelerating part was relatively short and the steady state was attained when about 0.01-0.02 mole oxygen per mole hydro- carbon bad been absorbed. Results were reproducible within z yo for any given set of conditions.The slope of the linear part of the curve was taken as a measure of tbe rate of oxidation. DEPENDENCE OF RATE ON HYDROCARBON AND CATALYST CONCENTRATION. -The effect of variation of the hydrocarbon concentration a t 40-45' C for three different catalyst concentrations is shown as a log-log plot in Fig. 3. The slope showed a constant value of 2.0 and thus the rate a t constant catalyst concentra- tion was proportional to [hydrocarbon]*.286 OXIDATION OF TRIMETHYLETHYLENE Fig. 4 shows the dependence of the reaction rate on catalyst concentration for two typical hydrocarbon concentrations. In all cases the slope of the log-log plot was 0.50 f 0-02 thus corresponding to a rate proporticnal to [catalyst]&. EFFECT OF OXYGEN PRESSURE.-A~ constant concentration of trimethyl- ethylene and catalyst the rate of oxygen absorption, as summarized in Table I, was independent of the oxygen pressure over the range 560-910 mm.Other experiments a t higher hydrocarbon and catalyst concentration in which dry air at atmospheric pressure was used showed that the rates were unchanged by sub- stituting air for oxygen. It may tberefore be concluded that the rate was in- dependent of oxygen pressure down to about 150 mm. TABLE I.-VARIATION OF RATE WITH OXYGEN PRESSURE AT 40'45°C [Hydrocarbon] = 1-12 mole/l. [Cobaltous acetate] = 1-04 x IO-% mole/l. pressure Steady Rate of 0% Absorption (mole/l. hr. (N.T.P.)) 559 610 664 763 864 911 Total Cobalt Salt 26.1 x I O - ~ 26.1 x I O - ~ 27.8 x I O - ~ 28.4 x I O - ~ 29.0 x I O - ~ 27-8 x I O - ~ Rate of O2 Absorption (mole /l.hr.) ACTIVATION ENERGY.-The plot of the logarithm of the rate of oxygen ab- sorption against 1/37 for a series of a catalyst concentrations is shown in Fig. 5 . In all cases the linear relationship corresponded to an overall activation of 15,000 cal., independent of the catalyst concentration. EFFECT OF THE INITIAL ADDITION OF COBALTIC ACETATE AND THE PER- OXIDIC PRoDucT.-Since it was observed that a t higher catalyst concentration a colour change occurred which may be a result of the oxidation of the catalyst, to the higher valence state, measurements were made in which cobaltic acetate prepared by the electrolytic oxidation of the cobaltous salt, was used as catalyst. As may be seen from curve 11, Fig.2, the cobaltic salt caused an immediate and increased rate of oxygen absorption but after less than 0-01 mole oxygen/mole hydrocarbon had been absorbed a very pronounced slowing down of the reac- tion occurred and from then onwards the reacticn followed a course similar to that observed with the cobaltous catalyst. In fact, after the above short initial period, as the iesults in Table I1 show, the steady state rate was close to that expected for the corresponding concentration of initially added cobaltous acetate. TABLE II.-COMPARISON OF COBALTOUS AND COBALTIC ACETATE [Hydrocarbon] = 1-12. Temp. = 40-5"C 1'45 X 10-' 2'9 x 10-3 4'3 x 10-8 6.6 x I O - ~ 8.8 x I O ~ 11'2 x 10-8 14'4 x I O - ~ 17.6 x 103 21.9 x 104 25'4 X IOs Concentration (g.-atom Call.) Cobaltous Cobaltic 9-2 x I O - ~ 15.3 x I O - ~ 19.5 x I O - ~ 20.5 x I O - ~ 24.6 x 10-3 Visual observation showed that depth of the green colour characteristic of the cobaltic ion complex slowly disappeared during the initial reaction phase, the solution becoming brown and finally showing the pink colour of the cobaltous ion with the lowest catalyst concentrations.In view of the remarkable oxidizing properties of the cobaltic ion towards organic substances i t was of interest to examine its behaviour towards trimethyl- ethylene and some preliminary observations carried out a t our suggestion byC. BAWN, A. PENNINGTON AND C. TIPPER 287 Dr. A. G. White showed that this hydrocarbon reacted very rapidly with dilute solutions of cobaltic sulphate in 4N H,SO, a t room temperature in the absence of oxygen.Solutions N/5o with respect t o cobaltic ion were reduced t o the cobaltous state within 10 min. at 20° C. By analogy with previous observations it may be concluded that the cobaltic ion reacts readily with the trimethyl- ethylene molecule. FIG. 4.-Dependence of rate on catalyst concentration. I ' I 3-20 3-30 FIG. 5.-Temperature dependence of the overall reaction. The effect of the initial addition of the peroxidic product of the reaction was studied in the following way. The catalyst was added t o the hydrocarbon solution in contact with oxygen and the mixture left for a definite time during which hydroperoxide was formed. On carrying out the rate of oxygen absorp- tion on these mixtures in the usual way it was found that the accelerating period was either considerably reduced or eliminated.With sufficient added hydro- peroxide the initial reaction rate was faster than the steady rate but in all cases the system rapidly settled down to steady state attained without the addition.288 OXIDATION O F TRIMETHYLETHYLENE EFFECT OF WATER.-Small amounts of added water had a marked inhibiting effect on the reaction, and not only lengthened the period of acceleration to the steady rate but also caused a considerable decrease in the latter. The variation of the rate of oxygen uptake with water concentration a t constant catalyst con- centration is sbown in Fig. 6. The plot of the reciprocal of rate against the water concentration was linear. A similar inverse dependence of the rate on water concentration has previously been observed in the cobalt acetate catalyzed oxidation of acetaldehyde in acetic acid solution.FIG. 6. HYDROPEROXIDE FORMATION.-Determination of the hydroperoxide forma- tion during the course of the reaction showed that its concentration rapidly reached a constant value, and as will be seen from Fig. 7 the amount of hydro- peroxide with a catalyst concentration > 2 x I O - ~ mole/l., corresponded to a small fraction of the oxygen absorbed. Furthermore, equilibrium hydroperoxide concentration was proportional to the initial [hydrocarbon] and independent of the catalyst concentration (2.21 - 6.63 x I O - ~ mole/l. cobaltous acetate) as shown by the results in Table 111, TABLE III.-VARIATION OF EQUILIBRIUM HYDROPEROXIDE CONCENTRATION WITH HYDROCARBON AND CATALYST CONCENTRATION (All concentrations molep.) Hydrocarbon 1-10 0.84 0.56 0.84 0.56 0.84 0.56 1-12 1-12 Cobaltous Acetate x 1 0 4 4'70 4'42 4'42 2'21 2'2 I 2'2 I 6.63 6.63 6.63 Oxygen Absorbed 10-3 3 1'4 16.5 6-1 13'3 7'0 3'5 5 1-0 31-8 13-1 H ydroperox ide 10-8 2.05 1-24 0.5 I 2.33 1-33 0.45 2'39 1.41 0.50 [Hydrocarbon] 2 [Hydroperoxide] x I02 -- 5'9 5'7 6.0 5'4 5'4 6.0 4'8 5'0 6.2C.BAWN, A. PENNINGTON AND C. TIPPER 289 At much lower catalyst concentrations the above relationships were no longer true. Thus with [hydrocarbon] = 1-68 mole/l. and [catalyst] = 0.56 x 10-3 and 1-12 x I O - ~ mole/l. the rate of oxygen absorption was 10 yo and 50 yo of that required by the rate law dO,/dt = [RH]* [Catla. Moreover, under these conditions the equilibrium concentration of hydroperoxide was established very slowly and about 75 yo of the oxygen absorbed was present as hydroperoxide.At low catalyst concentration therefore another mechanism is operative and the kinetic relationships in this region have not yet been determined. FIG. 7.-Variation of reaction rate with added water. COBALTOUS-COBALTIC RATIO.-At high initial c a d y s t concentration a noticeable colour change of the solution was observed during the course of the reaction, and the pink colour characteristic of the cobaltous state slowly dis- appears with the appearance of a brown colour. In no instance was the green colour of the complexes of the cobaltic ion in acetic acid observed. Independent experiments with cobaltous-cobaltic mixture showed the brown coloration was due t o the presence of both the cobaltous and cobaltic ions in the solution.Determination by a spectrophotometric method of the cobaltous-cobaltic ratio in the reaction medium when the equilibrium hydroperoxide concentration bad been attained showed that the highest concentration of catalyst used, viz., I x I O - ~ mole/l. cobaltous acetate, g yo of the cobalt was in the higher valency state but that a t concentrations below 7'1 x 10-a mole/l. all the catalyst was in the lower valency state throughout the course of the reaction. This result is in direct contrast with the corresponding cobalt ion catalyzed oxidations of acetaldehyde and benzaldehyde 8 where the catalyst is maintained in the higher valency state during the oxidation.Discussion Extensive investigations of the oxidation of non-conjugated unsaturated hydrocarbons, including mono-olefinic, I : 4-dienic and hydroaromatic 8 Bawn and Jolley (unpublished). K290 OXIDATION OF TRIMETHYLETHYLENE molecules 1 have shown that reaction obeys the chain mechanism initially formulated by Gee and Bolland,' viz., 1 k 8 k8 Initiation R + h v + 2ROOH -+ ROOR + R + 0, +- ROB ROOH + h v + or RO, R~ Propagation RO, + RH +ROOH + R non-radical k4 products " R + R O , + RO, + RO, + IZ6 Termination where R is the radical obtained by extraction of hydrogen from the hydro- carbon RH, and RO, is the corresponding peroxidic radical. At high oxygen pressures (> 40 mm.) this scheme leads to the rate equation, * (1) k d0,jdt = Ri* 3 [RH], .k6' where Ri is the rate of the initiation reaction. In the metal ion catalyzed oxidation there is no reason to assume any other propagation reaction and this adequately accounts for the initial formation of the hydroperoxide. Any modification of the kinetic laws may therefore be attributed to the different nature of the initiation and possibly also the termination reactions. The initiation reaction may be the direct reaction between the cobaltic ion and the hydrocarbon, viz. (The cobalt ion exists as a complex in acetic acid and here as elsewhere in the paper the simple designation is used although this may need modifica- tion when the structure of the ions has been worked out. Preliminary results reported above show that reaction (2) occurs readily at room temperature with trimethylethylene, the higher valence state of the ion being formed by peroxidic oxidation.Inserting the rate of initiation, reaction ( 2 ) in the Gee-Bolland scheme leads to Co3+ + RH R + H+ + Co2+. . * (2) d(O,)/dt = k, $ [Co3+]* [RH]', . (3) which is not in agreement with the experimental relationship. It is of interest to note that eqn. (3) is obeyed for the cobalt ion catalyzed oxida- tion of benzaldehyde in glacial acetic acid.8 The build-up of the hydroperoxide to an equilibrium value demands the occurrence of a reaction causing decomposition of the peroxides. The experimental results indicate that this is the principal function of the metallic ion catalyst which may decompose the hydroperoxide according to the reactions ROOH + Co8+ -r RO + OH- + Co*+ (or RO- + OH + Me2+) (4) These reactions appear to be more probable than (3) as the initiation step, the initial small concentration of hydroperoxide necessary being formed by thermal oxidation.The metal ion may also take part in termination of the chain by an electron transfer reaction with the radical RO, (or RO), viz. The observed dependence of the oxidation rate on the square root of the catalyst concentration, however, indicates that the controlling reaction in termination is a radical-radical interaction. Bawn and White, J . Chem. SOC., 1951, 331, 339, 344, ROOH + CO'+ -f RO, + Co8+ + H+. . - ( 5 ) RO, + Co2+ RO,- + Cog+. (6)C. BAWN, A. PENNINGTON AND C. TIPPER 291 Using the initiation reaction, (4) and (5), together with the propagation and termination reactions of the Gee-Bolland mechanism, give a rate expression of the form (7) which in conjunction with the experimental rela tionship for the equilibrium hydroperoxide concentration at high catalyst concentration, viz. gives the observed oxidation law dO,/dt cc [RH] [ROOHI* [Cat],, [Cat] = [Co8+] or [CoS+], [ROOH] cc [RH], . * (8) - dO,/dt cc [RH]*[Cat]*. The suggested reaction scheme, however, is not consistent with the relationship (8) and leads to a peroxide dependence of the form [ROOH] cc [RHI2/[Cat], [Cat] = [Co*+] or [Cos+]. The interpretation of (8) is not yet clear and this result seems to be con- nected with several factors. (i) At low catalyst concentration, (8) no longer holds. (ii) The cobaltous-cobaltic ion ratio in the reacting systems depends on the initial hydrocarbon concentration. (iii) The catalyst functions as a complex ion, the structure of which is unknown. (iv) The mechanism of the inhibiting action of small addition of water ; this may be due to its effect by substitution in the co-ordination shell of the cobalt ion. The understanding of the complete kinetics may await the answers to these questions. We are indebted to the D.S.I.R. for financial aid to one of us (C. F. H. T.). Department of Iraorganic and Physical Chemistry, University of Liverpool.

ISSN:0366-9033    

DOI:10.1039/DF9511000282

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

C. BAWN, A. PENNINGTON AND C. TIPPER 291 THE REACTIONS OF LIQUID ETHYL BENZENE WITH OXYGEN IN GLASS VESSELS BY G. M. HENDERSON Received 1st February, 1951 Pure ethyl benzene, in clean glass systems, reacts with oxygen at 130OC first t o give a very slow radical reaction. This is replaced by a faster wall- catalyzed reaction. Both these reactions produce only the hydroperoxide of the hydrocarbon. The wall-catalyst apparently consists of stabilized free radicals derived from ethyl benzene. The amount of the wall-catalyst laid down from any one sample of ethylbenzene at any one temperature is a function of the original volume. The stabilized free radicals show a remarkable degree of stability. General experience with industrial liquid-phase air oxidations suggested that an experimental investigation of the oxidation of an isolated methylene group at atmospheric pressure, in glass at 1 3 0 O C both with and without catalysts, would aid in understanding what happens in the induction period of the more complex catalyzed technical oxidations which are usually necessarily carried out under pressure in metal equipment, and on materials presenting several places of attack per molecule.Ethyl benzene was chosen as it presented an isolated methylene linkage, open to oxidation, between the relatively inert phenyl and methyl groups. In the preliminary investigation, with which we are alone concerned here, no catalysts were used and the attempt was made to exclude all292 REACTIONS O F ETHYL BENZENE foreign ions known to be oxidation catalysts. Carefully purified ethyl benzene was used throughout and, although different specimens showed the familiar scatter of results of this field under any fixed experimental conditions, large bulked samples gave reproducible results.Consequently in making comparative experiments samples were withdrawn from the same stock sample as far as possible at the same time (Fig. 3). After the experimental work discussed in this paper was finished, a recent patent disclosed that commercial isopropyl benzene, even after careful fractionation, contains traces of styrene and its derivatives. Unless these are removed by special chemical means, isopropyl benzene will react only sluggishly with oxygen to form hydroperoxides. As the purification employed in our experiments involved only the sequence fractionation, washes with sodium bisulphite and sodium bicarbonate solutions, water washing and refractionation in a column of 30 theoretical plates, this fact must be borne in mind.Jacketted oxidizers in Pyrex glass were constructed whose inner cylindrical reaction tube was 5 cm. wide by 30 cm. deep. Sintered glass discs 2-5 cm. in diameter were set on the end of flared-out narrow tubes forming a spoon. The porosity could be varied from G1 (coarse) to G, (fine) and were adapted from standard Pyrex Gooch crucibles. These discs were set with their faces upwards, at 45" to the vertical. The shallow spoon on the end of the air line under these circumstances was completely self-draining on commencing air flow. Various jacketting liquids or vapours could be used to adjust the reaction temperature.The reactor was " blacked out " for dark reactions (as in the bulk of the present experi- ments) by swathing in layers of aluminium foil. Porosity of the disc and air flow rate were shown to have little effect provided the system was saturated with oxygen and a flow of 20 l./hr. of air was fixed as a standard. Gas analysis on oxidations using air showed that the exit gas contained 20 yo oxygen on average. The ascend- ing stream of bubbles from the sintered disc at the foot of the reactor gave excellent agitation by an air lift effect, bubbles being swirled below the level of the injection area. Cooled portions of the liquid contents were withdrawn at any. time for analysis by means of a built-in water jacketted pipette with a capillary stem dipping to the foot of the reactor.Analytical samples were then repipetted from this cooled zone and the residue re- turned to the reactor. The most important experimental feature is that the various Quickfit joints must be dry or unlubricated. The whole apparatus before each experiment must be thoroughly cleaned with a strong metal-free oxidizing agent ; concentrated hot nitric acid or mixtures of nitric and hydrochloric acid were used with equal success. The acid was then removed with dis- tilled water and the apparatus finally baked at 130-150" C in an electric oven for at least I hr. If this cleaning was not carried out properly, a series of repeat experiments showed a steady decay in rate each time. A few experiments suggested hot concentrated sulphuric acid was also effective, but it was much less convenient and more suspect of metallic contamination.One or two experiments were carried out in spherical reactors of different volumes so as to vary the ratio, wetted wall area/volume. It proved difficult to fill the reactors with glass rods or tubes without disturbing the even distribution of gas bubbles in the system and creating stagnant enclosures of liquid. The results obtained over several years have not fallen easily into the current theories and a very tentative hypothesis is now advanced which seems to cope with some of the experimental facts. The rate of reaction in all cases was followed by titrating samples of the reactor contents for hydroperoxide by the Kokatur and Jelling a method and the hydroperoxide content was expressed as ml.of 0.1 N sodium Kokatur and Jelling, J . Amer. Chem. Soc., 1941, 63, 1432. 1 Distillers Company Ltd., Brit Pat., 630,286.G. M. HENDERSON 29 3 thiosulphate finally consumed by I ml. of reactor contents. The sum- marized findings are : (I) The specific rate of hydroperoxide formation at 130’ C in the pure thermal reactions depends primarily on the volume of the charge being oxidized from any one batch of purified ethyl benzene (Fig. 4). ( 2 ) Dilution of ethyl benzene with up to 33 yo of pure diphenyl before oxidation does not alter this rate dependence on the original volume of the solution (Fig. 6). ( 3 ) The rate of the thermal reaction, once established, is independent of rate of air flow, bubble size or even the partial pressure of oxygen. The solutions in our experimental range were always fully saturated with oxygen. (4) Once the steady rate for a given volume of ethyl benzene has been established, exposure to u.-v.light leads to a rise in rate, which is mostly retained on removal of the exciting source of radiation (Fig. 10). ( 5 ) Oncc the steady rate for a given volume has been established (0) addition of fresh ethyl benzene, say by doubling the volume, does not alter the rate (Fig. 8) ; ( b ) removal of a major portion of the charge, say by reducing the volume to one-half, greatly enhances the specific rate to somewhat over twice the original figure (Fig. 9). (6) If a volume of ethyl benzene is oxidized at its steady rate, and the reactor contents are wholly removed from the reactor, two effects may be subsequently observed : ( a ) If the reactor is cleaned with methanol, water, and then baked at 130’ C an attempted repeat experiment with the previous volume of ethyl benzene or larger volume of ethyl benzene, only reproduces the steady rate of the previous experiment.Some factor of the original rate has been imposed on the walls as a “reaction memory ”. The rate shows a slight but variable decay according to the severity of the solvent washing. If this subsequent experi- ment takes the form of using a smaller volume than the original experiment, higher than normal rates may be experienced. (b) If the walls are cleaned with nitric acid following the standard procedure, the now clean walls are without effect and the specific rates observed initially depend once more only on the initial volume of the ethyl benzene charged to the oxidizer; the other two variables (the temperature and provenance of the samples) being kept constant.(7) If a volume of ethyl benzene is oxidized at its steady rate and is then poured into a new clean reactor, the reaction proceeds at the same steady rate in the new reactor, in spite of depositing rate influencing sur- faces on the walls of both reactors (Fig. 7). (8) Very pure and fresh samples of ethyl benzene show an induction period during which there is a slow but steady rate of oxidation which is then replaced by a faster and steady rate. This induction period is fleeting and can only be observed under favourable conditions. This suggests that the reaction commonly observed is catalyzed in some way, even with attempted exclusion of catalyst ions (Fig.3). (9) In all these reactions, hydroperoxide is the only detectable reaction product, the formation of water characteristic of the metal-catalyzed reaction has never been observed without the deliberate addition of traces of metal catalyst such as cobalt naphthenate. (10) I t is noticeable that the light-catalyzed reaction shows signs of sensitivity to oxygen partial pressure, apparently being poisoned by high partial pressures of 0, at temperatures of IOOO C. From these facts i t seems that the reaction generates a limited quantity of material on the walls of the glass reactor which acts as a catalyst. This294 REACTIONS OF ETHYL BENZENE is presumably organic in nature.This bears some relation to the original volume of the system and is one factor in influencing the specific rate. Furthermore, this catalyst is able to short circuit the original catalyst- forming reaction, which it is reasonable to suppose to be a homogeneous chain reaction by analogy with other hydrocarbon systems. Since the system is at all times saturated with oxygen, the two reactions cannot compete over the system : hydrocarbon + dissolved oxygen. We are forced to assume that there is a slow equilibrium absorption of some of the oxygen into a more closely bound form slow RH2 + 0, [RH,-++O,] the oxygen here being held as an activated complex with the methylene bridge group under attack. We then suppose that this may decompose in at least two ways, back to dissolved oxygen or most infrequently into a hydrocarbon free radical.We may now set down the slow chain reaction of the induction period : very slow CHAIN INITIATION. [RH, c --f O,] - RH' + H02 CHAIN PROPAGATION. RH' + 0, + RHO2' RHO,'+ RH, -+ RHOOH + RH'. CHAIN TERMINATION (at least in part). RH' + wall -+ RH- wall+. I t is now supposed that at least part of the captured radicals on the walls are stabilized by withdrawing electrons from the glass and are then sufficiently stable to form long-lived complexes. It is suggested that a wall charged with such stabilized free radicals will lead to a facile re- arrangement on its surface of the ethyl benzene-oxygen complex leading to the peroxide and a regenerated ion or radical with only a movement of charges.One further supposes that this surface reaction is much faster than the original chain-initiating step ; it thus greatly reduces the stationary concentration of the active complex and hence practically suppresses the original classical chain reaction in solution which led to the setting-up of this wall process. The wall reaction may. then be regarded as a chain reaction carried out on the surface of the wall by a fixed quota of stabil- ized free radicals, or as a wall reaction catalyzed by a highly efficient organic ion catalyst supported on glass. The wall process resists simple organic solvents, water and dry baking for brief periods with only slight decay but is completely destroyed by rezgents such as hot concentrated nitric acid, We can now fit the hypothesis to the experimental facts if we suppose that the totaI charge of wall catalyst is a function of the original volume Vo.Total catalyst oc fV,. We now suppose the system circulates over the surface, the catalyst coming in contact with succeeding units of volume. The system may be best treated as the pouring of a volume V , of ethyl benzene through a catalyst bed containing a quantity of catalyst fV, in unit time, the amount of catalyst alone controlling the efficiency of the reaction. If we double the volume, we suppose we exactly double the efficiency of the catalyst bed and again we have a volume 2 V 0 treated as pouring through a bed of catalyst 2fV0 in unit time, the catalyst bed now being of twice the efficiency. This gives us a rate dependence as found on varying the initial volume ; furthermore it matches the movement of specific rates &-hen the volume is subsequently doubled or halved, The invigoration This is illustrated in Fig.I and 2 . The wall coating is quite invisible.G. M. HENDERSON 29 5 of the reaction by u.-v. irradiation, leading to a permanent enhancement of the reaction rate is explicable by supposing the superimposed photo- initiated chain reaction is able to generate free radicals anew by an in- dependent route which generates a further quantity of catalyst. FIG. RELATIVELY RH- CHAIN TERUINATION t Slow chain reaction in brief induction period. PHOOH WROPL ROXlDE ph phase I. RH, + Q, 1"'"" CHAIN REACTION COATED WALLS RHOOH HYDROPEROX1 DE FIG. 2.-Phase I J . Final fast wall reaction governed by stabilized radicals.XPIJ. 0 AND 8) E.B. PURIFIED E.B I 2 3 4 5 FIG. 3.-Effect of provenance on specific rate of oxidation of ethyl benzene. Conditions : 130' C, 150 g. ethyl benzene.296 REACTIONS OF ETHYL BENZENE The experimental data are appended in the form of graphs showing rates of reaction under different conditions. The rates are expressed as the build-up in time of the peroxide titre of I ml. of the reactor content in ml. 0.1 N sodium thiosulphate. t5 2 3 4 5 I '- I I - I I L FIG. +-Specific rates of peroxide formation in relation to volume of charge. Conditions : varying E.B. charges, 130' C. t5 5 2 T I- 2 3 4 5 FIG. 5.-Effect of temperature on thermal reaction a t 100' C and 130' C on same bulked sample ethyl benzene, Conditions : 150 g.charges of E.B. 150 9. ETHW- BENZENE COMPARED WITH EQUAL VOLUME SOLUTION OF DIPHENYL IN ETHYL BENZWE AND 13OoC. (SAME STOCK) AT too0c. FIG. 6.-Comparison of pure ethyl benzene with a mixture of diphenyl (I mole) and ethyl benzene (2 moles), equal volumes of solution being compared.G. M. HENDERSON 297 It is suggested that this type of oxidation probably only appears under favourable circumstances with particular hydrocarbons, but nevertheless may explain the variable results of past workers in this field. 4 4 w " @= DISC IMMERSED 8.5 Cm. IN NARROW CYLINDRICAL REACTOR. A= DISC IMMERSED 443Cm. IN SPHERICAL REACTOR n r 4 vl L E -3 TIME SCALE ADJUSTED. REACTOR, FIG. 7.-Effect of transferring charge of oxidizing ethyl benzene to new clean reactor of different shape.Conditions : 150 g. ethyl benzene, 130° C. HOURS 3 3 4 3 FIG. 8.-Effect oi doubling ethyl benzene charge during oxidation (temperature 1 3 0 O C). FIG. g.-Effect of removing half the charge of oxidizing ethyl benzene during an oxidation. Conditions : 500 ml. ethyl benzene in I litre round flask in oil bath at 130' C. It may be best to visualise this reaction as an extension of the original homogeneous chain reaction which has colonized the reactor walls. To draw another analogy, it is the developed and fixed latent image on the walls of the stationary state in the earlier homogeneous reaction.298 OXIDATION OF LUBRICATING OILS We had at one time supposed that this type of wall-process might also take place on the molecules of oil soluble cobalt catalysts used in catalyzed air oxidations of hydrocarbons, thus giving the striking colour charge (blue +green) of cobaltous salts to cobaltic complexes when the rapid oxidation commences after an induction period. The recent ex- periments of Bawn and co-workers 8 demonstrate that hydroperoxides alone can account for this phenomenon, ... w t5 5 I FIG. 10.-Effect of period of exposure t o u.-v. lamp on a reaction. Conditions : I 50 g. ethyl benzene, I 30° C. However, for metal-catalyzed oxidations, the slow growth of hydro- peroxide during the induction period may well arise by such a wall-process which will then in turn be suppressed by the hydroperoxide activated metal ions, whose reactivity is such, that the concentration of dissolved oxygen in many catalyzed reactions is nearly reduced to zero once the rapid catalyzed reaction is running. Imfierial Chemical Industries Limited, Research Laboratories, Hexagon House, Blackley, Manchester, 9. Private communication.

ISSN:0366-9033    

DOI:10.1039/DF9511000291

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

OXIDATION OF LUBRICATING OILS IRON AND COPPER AS CATALYSTS I N THE OXIDATION OF HYDROCARBON LUBRICATING OILS BY J. H. T. BROOK AND J. B. MATTHEWS Received 26th January, 1951 The kinetics of the catalyzed oxidation of a lubricating oil in the liquid phase have been studied, using the rate of absorption of oxygen as a measure of the reaction velocity. The most important catalysts which come in contact with oils used for the lubrication of internal combustion engine crankcases being iron, copper and " crankcase catalyst ", i t is shown that the catalytic activity of all these is probably due to the dissolved iron and copper compounds formed by corrosion. A chain reaction mechanism for the catalyzed oxidation is proposed, and it is found that dissolved copper is a good chain-initiation catalyst, while dissolved iron is a good chain-branching catalyst. With increasing concentra- tions of dissolved copper the overall rate of oxidation reaches a maximum, and i t is suggested that this is due to micelIe formation of the dissolved copper soap.J.H. T. BROOK AND J. B. MATTHEWS 299 Whilst many catalyzed oxidation tests have been devised for pre- dicting the relative performance of hydrocarbon oils used for the lubrication of the crankcases of internal combustion engines, only Larsen and co-workers have made any systematic examination of the properties of the catalysts which are most likely to be encountered in the engine. Thus, Larsen, Thorpe and Armfield have shown that lubricating oils are mixtures of hydrocarbons stabilized by weak natural inhibitors, and Larsen and Armfield * have investigated the behaviour of several hydro- carbon oils under a wide range of conditions using lead, iron and corJper as catalysts, without, however, analyzing their results kinetically. Larsen, Armfield and Whitneya reached the conclusion that the oil- insoluble material present in used engine oils, which in the dry state was described by them as “ crankcase catalyst ”, was the most active catalyst present under service operating conditions.In addition George and Robertson 4, have investigated the catalyzed oxidation of some hydro- carbons, with special emphasis on tetralin, and presented reaction schemes in general agreement with those proposed by BollandB for the oxidation of olefins. The investigation described now was undertaken in order to examine in more detail certain parts of the field which are important with respect t o lubricating oils, and in particular to determine whether the simple kinetics found applicable to the reactions of pure hydrocarbons could be applied to complex mixtures such as lubricating oils.In addition to crankcase catalyst, metallic iron and copper are generally accepted as being the most important catalysts for the oxidation of crank- case lubricating oils in the field. It is now shown that all three catalysts owe their activity to the production of dissolved iron and copper com- pounds. Semenov’s concept of the degenerate chain-branching reaction has been applied to describe the kinetics of the liquid-phase oxidation of a typical lubricating oil in the presence of dissolved copper and iron soaps, and of these catalysts the former is found to be more effective for the chain-initiation step than for the chain-branching steps, whereas the reverse holds for the latter.Experimental Apparatus.-A circulatory oxygen absorption apparatus (Fig. I) of the type described by Dornte 7 was employed for all the measurements except when an inert solvent was used. The apparatus differed from Dornte’s in that glass valves were used in the circulating pump, a capillary flow meter was in- serted to measure oxygen flow rates, larger reaction vessels were used holding IOO g. of sample, the condensers did not return volatiles to the flask, a solid CO, cooled trap was provided, and the absorber train consisted of tubes filled with calcium chloride, activated charcoal, Hopcalite and soda asbestos.Auto- matic recording of oxygen uptake was provided. A static oxidation apparatus was employed for the measurements involving an inert solvent. This apparatus 8 consisted of a IOO ml. conical flask, con- nected to a gas burette and manometer, and gave the same induction periods for oxidation inhibited oils as the circulatory apparatus. Oil samples were degassed and saturated with oxygen. Five g. samples of oil were weighed into the flask and the apparatus was then flushed with oxygen prior to the com- mencement of each experiment. Whilst it could not be shown that a maximum oxidation rate existed in the circulatory oxygen absorption apparatus, the rates obtained with reaction vessels of varying dimensions and gas diffuser pore-size were constant within 1 Larsen, Thorpe and Armfield, Ind. Eng.Chem., 1942, 34, 183. Larsen and Armfield, I n d . Eng. Chem., 1943, 35, 581. Larsen, Armfield and Whitney, SOC. Auto Eng. J., 1943, 51(T), 310, 4 George and Robertson, Trans. Faraday SOC., 1946, 42, 217. 5 George and Robertson, J . Inst. Petrol., 1946, 32, 382. 6 Bolland, Proc. Roy. SOC. A , 1946, 186, 218. 7 Dornte, Ind. Eng. Chem., 1936, 28, 26. Reaven, Irving and Thompson, J . Inst. Petrol, 1951, 37, 25300 OXIDATION O F LUBRICATING OILS the limits of experimental error. However, the rate of oxidation depended slightly on the ratio of the oxygen flow rate to the sample weight. This was taken to indicate that the oxygen flow rate was not rate determining, but that variations in the degree of removal of reaction products affected the rate slightly.Similarly, attention had to be given to the absorption train, and the combination used was found to give the highest reaction rates. The repeatability of the measurements was variable, usually being about f 5 yo and about f 3 yo for the uninhibited and inhibited reactions respectively. n COLD TRAP FIG. I .-Circulatory oxygen absorption apparatus. Materials .-OIL.-A typical solvent refined paraffinic lubricating oil Carbon group analysis DILUENT.-B.D.H. diphenyl, twice recrystallized from 60-80 petroleum INHIBITOR.-The calcium phenate derivative of the iso-octyl phenol-formal- CATALYSTS.-Cupric stearate, recrystallized, (cu 9.9 yo : theory 10 yo). Ferric stearate, recrystallized, (Fe 7-7 yo : FeSt, requires Fe 6-18 yo, and Copper wire, soft commercial, 18 S.W.G.Reduced iron powder (specific surface of the order of 7000 sq. cm./g.). Crankcase catalyst, as described by Larsen, Armfield and W h i t n e ~ , ~ essenti- ally the oil-insoluble material which accumulates in internal combustion engine crankcase oils. r ] l o ~ ~ ~ 121.2 cs., pk5 0.879, sulphur content 0-42 % w. (Leendertse Q) : aromaticity 4.9 yo, naphthenicity 27-1 yo, paraffinicity 68 yo. spirit. dehyde resinlo hence this material is probably a basic stearate). Results Solid Catalysts.-The rates of oxidation of the oil catalyzed by copper metal (Fig. 2) proved t o be insensitive to the surface area of the catalyst over the surface area range 0.25 crn.,/g.oil to I cm.2/g. oil, and copper was found to corrode during the oxidation to an appreciable extent, giving, for example, 50 p.p.m. of dissolved copper after the absorption of 20 ml. of oxygen per g. oil, with I cm., of copper exposed per g. oil. Assuming that the amount of corrosion is proportional to the area of the catalyst surface and the amount of oxygen absorbed, it was found possible to correlate the results obtained with copper metal with results obtained using dissolved copper in the form of the stearate. Thus, from the above corrosion data one would expect to find, with I cm.2 of copper surface per g . oil, that the dissolved copper contents after absorption of I ml. O,/g. were 2-5 p.p.m., of z ml. O,/g., 5 p.p.m., of 4 ml. O,/g., 10 p.p.m., etc. It may be assumed that over the oxygen absorption range, 1-5 ml. O,/g.to 3 ml. O,/g., the average dissolved copper content was 5 p.p.m., and from the measured oxidation rate for this concentration of dissolved copper, the time taken in covering this oxygen absorption range can be determined. These various time intervals corresponding to various oxygen absorption intervals may be summed, and the results obtained in this way are presented as plot points in Fig. 2, showing good agreement between observed and computed oxygen absorptions. In the inhibited reaction (M/5o inhibitor concentration), copper metal again corroded slightly to give dissolved copper. Difficulties in the exact analysis of about 5 p.p.m. of copper in oil prevented the determination of the corrosion rate, and therefore an exact comparison of data on induction periods obtained 9 Leendertse (in press).l o U.S. Pat. 2,280,419, granted to the TJnion Oil Co. of California.J. H. T. BROOK AND J. B. MATTHEWS 301 with solid and dissolved copper could not be made. However, the induction period with I cm.2 of copper per g. oil was 870 min., while the average dissolved copper concentration appeared to be of the order of 4-5 p.p.m. and the induction period found with 5 p.p.m. of dissolved copper added as copper stearate was 830 min. From this it follows that, when metallic copper is used, dissolved copper is probably the active catalyst in the inhibited reaction, as well as in the uninhibited reaction. With iron powder as catalyst, corrosion was again observed, together with a very high degree of autocatalysis, the reaction obeying Semenov’s equation l1 dV/dt = A exp ($t), where A is the initial rate of oxidation, r j is the net chain-branching coefficient and V is the volume of oxygen absorbed. The initial rate A was very erratic, and it was not possible to apply the same reasoning as in the case of the copper catalyst in order to determine whether the catalytic activity was entirely due t o dissolved iron.In the inhibited reaction, with MI50 concentration of inhibitor and I g. of iron, giving about 70 cm.2 of iron surface per g. oil, the induction period was over 1200 min., showing that iron surfaces must be very weak catalytically. FIG. 2.-Influence of metallic copper upon the rate of ab- sorption of oxygen a t 1 6 0 O C ; ( I ) 0.10 cm.2 of Cu per g.oil; ( 2 ) 0.25 cm.2 of Cu per g. oil; (3) 0.5 cm.2 of Cu per g. oil ; (4) 1-0 cm.2 of Cu per g. oil. 0 and represent calcul- ated absorptions for 0.5 cm.2 and 1.0 cm.2 Cu per g. oil respectively. D so 100 I50 200 TIME FOR ABSORPTION MINUTES Larsen’s “ crankcase catalyst ” gave results with the uninhibited oil closely resembling those obtained with ferric stearate, taking I yo of the “ crankcase catalyst ” as being equivalent to 20 p.p.m. of dissolved iron. The high degree of dispersion of this catalyst in oxidized oils prevented accurate measurement of their dissolved iron contents, but the values obtained were of the right order of magnitude to account for the catalytic activity of this material. Since no soluble iron could be extracted from the catalyst by oils under non-oxidizing conditions, i t is probable that a corrosion mechanism is involved in the mode of action of this catalyst also. Soluble Catalysts .-Uninhibited Reactions.-Copper stearate exerts a maximum catalytic effect in this oil at about 20 p.p.m.of copper. It is possible that the reaction velocities are proportional to some simple power of the catalyst concentration a t low catalyst concentrations (of the order of magnitude of 2 p.p.m. or less), but the results a t higher catalyst concentrations (data a t 150’C are given in Fig. 3) are not easily interpreted. The apparent “ saturation ” of the inhibited oil by soluble copper with regard to catalytic activity affords an explanation for the copper metal catalyzed reaction being insensitive to the catalyst surface area, when the ratio of surface area to quantity of oil is greater than 0-25 cm.2 per g.oil. The oxidation rate curves do not fit Semenov’s equation for a reaction accelerating to a constant velocity, and are partially described by the parameters presented in Table I. l1 Sernenov, Chemical Kinetics and Chain Reactions (Clarendon Press, Oxford, 1935).302 OXIDATION OF LUBRICATING OILS The initial oxidation rates are not easily measured in the circulatory ap- paratus, since there are temperature changes in various parts of the gas flow lines on starting the oxygen flow. The reactions in the presence of dissolved iron obey Semenov's equation Catalyst Initial Rate Concentration ml.-~.~roo g. min. p.p.m. Cu 5 2'4 1 0 4'0 20 6-7 40 8.0 particularly well a t the higher catalyst concentrations as illustrated in Fig.4 where log V is plotted against time, but a t low catalyst concentrations the re- action rates fall away from the theoretical values as the reactions proceed. I500 - 0 d -- ItKQ E: 2 a 500 f 2 0" E 0 0 50 I00 IS0 200 TIME FOR ABSORPTION MINUTES FIG. 3.-Influence of dissolved cupric stearate on the rate of absorption of oxygen a t 150O C. Values for A , the initial rate of oxidation, fit an equation of the form A = k[Fe] + A , , and are thus not directly proportional to the catalyst concentration. On the other hand the values of 6, a more reliably determined constant, are approxim- ately proportional to the iron concentration (Fig. 5). and their temperature- dependence, given in Fig.6, leads t o an activation energy of 36 kcal./mole for the chain-blanching reaction. TABLE I Steady State Rate ml. 02/Ioo g. min. 5'6 8.25 10-5 9'1 The addition of soluble copper to the soluble iron catalyzed reaction gives a highly autocatalytic reaction with higher rates than are obtained with iron alone. Some of these rates are thought to be too high to be measured in the apparatus under the standard conditions without giving rise to a diffusion controlled re- action. Some values for A and 4 in Semenov's expression, where applicable, are given in Table 11. Inhibited Reactions .-In investigating the inhibited reaction, the in- hibitor was added a t a single concentration (M/50), and the measured induction period was taken to be the time required for an amount of free radicals to be produced equivalent to the amount of inhibitor originally present.The reciprocal of the induction period ~ / t , can be considered to be proportional t o the chain initiation rate. Experiments at low partial pressures of oxygen were conducted by filling the circulating system with appropriate oxygen +J. H. T. BROOK AND J. B. MATTHEWS 303 FIG. +-Influence of dis- solved ferric stearate on the rate of absorption of oxygen a t 150'C. 10 0 so 100 I so TIHE FOR ABSORPTION MINUTES FIG. 6.-Activation energy plot for chain-branching reaction in the oxidation catalyzed by 50 p.p.m. Fe as ferric stearate. FIG. 5.-Variation of 4 with concentration of dissolved ferric stearate a t 150' C. 0 and A represent values obtained with different samples of oil.-00235 00240 00245 00250 IIT304 OXIDATION O F LUBRICATING OILS nitrogen mixtures, and inserting a length of capillary tubing between the burette (which must necessarily be full of oxygen) and the circulating system to prevent mixing of the gases. As all the available inert diluents are in- conveniently volatile a t I 50' C for measurements in the circulatory apparatus, the static oxidation apparatus, in which volatility is less important, was used to obtain the dependence of the induction period upon hydrocarbon concentra- tion, diphenyl being chosen as the diluent. As will be seen from Fig. 7, 8 and g , the data for the copper catalyzed reaction can be represented by the expression, w = k [catalyst] [OJ* [RHIa, in which w is the chain-initiation rate and k is a constant, provided that the copper concentration is not greater than 30 p.p.m.The activation energy plots Simple kinetics were found for the reactions. ~~ 0 10 20 ppm Cu 30 49 so 20 40 ppm Fe 60 80 I I#) FIG. 7.-Influence of cupric and ferric stearates on the rate of chain initiation a t 150' C in the presence of M/50 concentration of inhibitor. are not good (Fig. IO), but lead to a value of about 30 kcal./mole for the activa- tion energy of the chain-initiation reaction. The chain-initiation rates are additive. For example, the value of w for 20 p.p.m. of copper in this system is 0.0038 min.-1, and for 50 p.p.m. of iron is 0.0018 min.-l. Experimentally, the value obtained from 20 p.p.m. Cu + 50 p.p.m. Fe is 0.0057 min.-l, in good agreement with the value for the sum of the separate rates, which is 0.0056 min.-l.TABLE I1 Fe p.p.m. I00 50 5 0 20 20 I 0 20 20 I 0 Cu p.p.m. A ml./roo g. min. 2.I* 1'3 1-5 * 1'1 * 0.8 4'2 2.4 1.0 1 2'1 4 min.-1 0.162 * 0.064 0.087 * 0.030 0.032 * 0.018 0.061 0.052 0.0 I 8 * Values obtained on a second sample of oil.J. H. T. BROOK AND J. B. MATTHEWS 305 FIG. 8.-Influence of partial goo01 pressure !of oxygen on the I rate of chain initiation in the 2 copper- and iron-catalyzed oxidations a t 15oO C in the presence of M/50 concen- tration of inhibitor ; ( r ) 20 p.p.m. of Cu as cupric stearate ; ( 2 ) 50 p.p.m. of Fe as ferric stearate. 0 05 I 0 F2 ( p o l PRESSURE, ATMOSPHERES) FIG. 10.-Activation energy plot for the chain-initiation reaction in the presence of 20 p.p.m.Cu as cupric stearate and 50 p.p.m. Fe as ferric stearate. FIG. 9.-Variation of chain- initiation reaction velocity a t 150' C with concentra- tion of oil in diphenyl. Catalyst: 20 p.p.m. Cu as cupric stearate. I 0 00235 ,00240 a I IT 45306 OXIDATION OF LUBRICATING OILS Discussion A simple reaction scheme may be postulated to represent the oxidation as a chain reaction with degenerate branching, including the following steps, in which the symbols have the usual significance and X represents a molecule of natural inhibitor.1 catalyst R H - R- W l - (1) R-+O2 -+ R0,- - (2) R- + X -+ inert products k4 - - (4) R 0 , - + R H --f R O O H + R - } kp : - (3) catalyst ROOH -* R- (or RO-) ROOH + inert products k6 * - (6) It is of little importance in the development of the final equations whether the radical in (4) is R- or R0,--, whether the radical produced in ( 5 ) is R- or RO--, etc.Following Semenov, and regarding the concentra- tions of radicals as being small and in equilibrium with each other, we can Put Or, since d [R-]/dt = W, - k4 [R-] + k5 [ROOH] * 0. dV/dt = kp [R-1, then dV/dt = (wl + K s [ROOH]) kp/k4. . . * (1) Now, - (ks + k6) [ROOHI. * a ( 2 ) d [ROOH]/dt = W ~ V + { K ~ ( v - I ) - k6) [ROOH] . * ( 3 ) Also v, the chain length, may be written for K,/k4. d[ROOH] dV dt dt Substitution in (2) of the value for dV/dt from eqn. (I) gives Integration of (3), with the condition that [ROOH] = o at t = o gives --- {exp [k5(v - I) - k6]t - I}. Wl V [ROOH] = K 5 ( v - I) - k ; Substituting for [ROOH] in eqn.(I) shows that This equation can be simplified under the two following conditions. Firstly, if and v the chain length is large then - I) k 6 , dV/dt = A exp ($t), . - ( 5 ) in which and $ = k,v. A = w1v = initial (unbranched) rate of oxidation, Secondly, when k d v - I) < K 6 , then dV/dt = A + B{I - exp (- $t)}, . ( 6 ) where and 4 = k6 - K ~ ( v - I). Eqn. (6) represents a reaction accelerating to a constant velocity. The oxidation of the uninhibited oil catalyzed by the higher concentra- tion of iron, and by iron and copper together, follows eqn. ( 5 ) . The experimental values for A (Table 11) may not in fact be values of wp, since they do not show the same dependence upon the catalyst concentra- tion as do the values for w, obtained from measurements on the inhibitedJ.H. T. BROOK AND J. B. MATTHEWS 307 system (Fig. 7), v being, according to the reaction scheme, independent of catalyst concentration and reaction rate. This d8ficulty in identifying A with w, is in all probability due to the nature of the oil. Thus, traces of unstable compounds or possibly of naturally occurring strong in- hibitors in the oil, might displace the reaction velocity curves in Fig. 4 along the time axis sufficiently to prevent this identification, since the values of A are derived from the intercepts of these curves on the volume axis. The chain-branching coe'fficient 4 which is given by the slopes of the curves in Fig. 4 is not affected by such considerations, and, since from Fig. 5 the value of 4 is proportional to [Fe], it may be identified with k,v.It will be noted from the few results in Table I1 that the addition of copper to iron in general increases A much more than 4. On a quali- tative basis, therefore, the data lead to the conclusion that copper is a good chain-initiation catalyst but a less effective catalyst for chain branching, Additional evidence in support of this conclusion is provided by a comparison of the shapes of the reaction velocity curves for the iron catalyzed and copper catalyzed reactions, since the former show a high degree of autocatalysis whereas the latter do not. The only example in the present work of a reaction closely following eqn. (6) is the oxidation catalyzed by 5 p.p.m. of iron, the data for which are given in Fig. 4. With copper catalysts, the reactions accelerate to a constant velocity without, however, obeying eqn.(6), as a result possibly of the value of the term {K5(u - I) - K,} being such that the conditions imposed in deriving eqn. (5) and (6) are not valid. The maxima in the rates of reactions catalyzed by copper stearate have been commented on before,2 and were thought by George and Robertson 4 to be due to the catalyst participating in a chain-termination reaction, mainly on the grounds that the addition of copper stearate reduces the rate of oxidation of long chain paraffins and tetralin catalyzed by cobalt stearate and ferric stearate respectively. However, as shown in Table 11, the addition of copper increases the already rapid rates of the iron catalyzed oxidations. Furthermore, even in the presence of a strong inhibitor, a maximum in the rate of oxidation is reached at ap- proximately the same concentration of copper as found in the uninhibited reaction. For these reasons it appears to be necessary to seek an alter- native explanation to that advanced by George and Robertson for the " saturation " effect observed with copper. The molecular weights of copper and ferric stearates have recently been measured ebullioscopically in cyclohexane in these laboratories, and were found to be 5000 for copper stearate in the concentration range zoo p.p.m. Cu to 800 p.p.m. Cu, and 1600 for ferric stearate in the concentration range IOO p.p.m. Fe t o 400 p.p.m. Fe. At these concentrations, therefore, copper stearate is highly associ- ated and ferric stearate is present largely as a dimer. Hence the possi- bility suggests itself that the maxima in the oxidation rates obtained with copper are due to micelle fortnation of the copper soap. This may be expected to occur at different copper concentrations in different systems. The authors wish to thank Mr. J. L. Dawson and Mr. D. K. Gilmour for experimental assistance and the Shell Refining and Marketing Co. Ltd., for permission to publish. Thornton Research Centre, P.O. Box No. I, Chester.

ISSN:0366-9033    

DOI:10.1039/DF9511000298

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

GENERAL DISCUSSION 308 GENERAL DISCUSSION Prof. W. A. Noyes Jr. (Rochester) said Recent experiments by Mr. Martin in our laboratory indicate that methyl radicals from the photo- lysis of mercury dimethyl upon reaction with oxygen at low pressures give CO and CO,. Not enough results were available when I left to allow calculation of rate constants although the rates are in approximate agreement with the acetone results. Some uncertainty exists concerning the nature of the primary process in mercury dimethyl so that many more experiments may be necessary (only two had been made before I left) before quantitative conclusions can be drawn. Other results have been obtained with acetone-oxygen mixtures by Messrs. Marcotte and Durbetaki at intensities about 10 times those pre- viously used both in the presence and absence of mercury vapour.It is concluded that mercury vapour does not affect the results by more than the experimental error. At these high intensities in the absence of oxygen at 120Othe acetyl radical seems to be sufficiently stable to permit appreci- able reaction of the type . CH + COCH = CH,COCH and possibly - (1) - (2) Thus a t this temperature the yield of CO + CO may be as much as 2COCH = (COCH,) . six times the yield of CO in the absence of oxygen. At zooo however the rate of decomposition of COCH is sufficiently rapid so that the results at these higher intensities agree with those previously found. These experiments are being performed with a view to permitting an analysis of products other than CO CO, CH and C,H,.Mr. Marcotte has also made experiments at 225OC. In general the results are similar to those obtained at lower temperatures but i t is evident is setting in. Thus CO + CO may have yields of 3-6 or more. This fact that at this temperature one has reached a point that chain propagation may limit the temperature range over which i t will be possible to deter- mine the activation energy for HCO decomposition. The fact that the CO is less sensitive to oxygen pressure and is higher than would be ex- pected from the low temperature results may be due to hydrogen abstrac- tion by radicals other than methyl. Dr. Finkelstein has preliminary results on the decomposition of diethyl ketone in the presence of oxygen.At 25°C the yield of CO falls as the pressure of oxygen increases indicating that the COC,H radical reacts with oxygen to give something other than CO at this temperature. In this respect i t seems to resemble the acetyl radical. However the quantum yield of oxygen disappearance is still rising at the highest oxygen pressure used (0.18 mm.) and has reached a value of 6. Complete oxidation of I molecule of ketone to CO and H,O would require 7-5 molecules of oxygen. Since small amounts of C hydrocarbons are found either more than I molecule of ketone disappears per quantum (doubtful a t 25OC) or the oxidation is proceeding quite completely. Further conclusions are not warranted at this time. At 200' C the yield of CO + CO seems to reach a constant value of 3 to 3.2 with increase in oxygen pressure the CO being about 2.3 and the CO about 0.8.However the quantum yield of oxygen disappearance has reached 7 at 0.78 mm. oxygen pressure and shows no sign of reaching a constant value. C hydrocarbons are still formed to the extent of about 0.3 mole per quantum even under these conditions. The total yield of CO + CO is about the same as for acetone thus in- It would be premature to give a mechanism based on these results. dicating that either a CH radical or a C,H radical yields either CO or CO when i t reacts with oxygen. However since the CO yield falls with increase in oxygen pressure for acetone and rises under similar circum- GENERAL DISCUSSION equation Two primary mechanisms suggested themselves namely * * CH,I-+ h9 CH +- I CH + 0 -f CH,O + OH e + HCOOH +- HCOOH.+ HCO + OH* occurs in preference to HCO + RH + HCOH + R HCO -f H + CO. or 3 09 stances for diethyl ketone the explanation for the two phenomena must have important differences. Dr. R. Spence (Harwell) said I would like to recall some experiments by Dr. Bates and myself on the oxidation of methyl radicals produced by the photodissociation of gaseous methyl iodide a t room temperatures. The products formaldehyde methylal water and iodine were identified and determined quantitatively and the reaction was shown to follow the qCH,I + 2 0 % -f (HCHO) + (CH,O),CH + H,O + P I . (1) - ( 2 ) and CH + 0 -+ CH,O,. The absence of any inert gas effect which might have been expected with mechanism (2) seemed to favour the hydroxyl mechanism (I) but Blaedel Ogg and Leighton failed to detect the OH radical by its ab- sorption spectrum and therefore proposed that a mechanism of type (2) is operative.Prof. Noyes’s results at high temperatures suggest that the fonnyl radical is produced from the interaction of methyl radicals and oxygen but the above results indicate fairly clearly that i t is not formed at room temperatures. This discussion has drawn attention to the fact that after twenty years we are still unable to give precise answers to such fundamental questions as how long such entities as 0 and CH can stay together on collision and what type of stabilizing collisions are subsequently required. Dr. N. Uri (Munchester) said The HCO radical was postulated by Marcotte and Noyes as an intermediate product in the photolysis of acetone in the presence of 0,.It appears to me very important to learn more about this radical which may also play an important role in the process of photo-synthesis. I consider that formaldehyde may be formed from formic acid by an electron transfer i.e. HCO + R . H 3 HCOH + R where the bond R H has a relatively small bond dissociation energy. On the other hand i t is postulated that HCO decomposes into H + CO. In this connection i t would be interesting to have more data on the formation of formaldehyde in the experiments discussed by Prof. Noyes so as to ascertain under what conditions Prof. A. R. Ubbelohde (Belfast) said It may be useful to reconsider to what extent the recombination of two radicals as in CH + CH -f C,H H+O,+HO - (B) requires a ternary collision to stabilize the product.Why do the energy- rich products not fly apart again almost at once as they would in simple atomic collisions such as H + 0 -f HO ? In the theory put forward in 1935 for the formation of peroxide radicals * R + 0 -f R02* (A) (1) . (2) . . R02* + RH + ROOH + R . I attributed the temporary stabilization of the RO complex to “ quantum smudging ” notable for R > C,H, but more rigorous conditions should apply when R is small as in reaction (B) and especially (A). Bates and Spence J . Amer. Chem. SOC. 1932 53 1689. Blaedel Ogg and Leighton J . Amer. Chem. SOC. 1942 64 2499.3 10 GENERAL DISCUSSION Prof. S . W. Benson (California) said Prof. Ubbelohde has raised a question with regard to the allowability of two-body recombinations such as have been employed by the present authors. I believe that this question can be answered with considerable more certainty than has heretofore been the case. A recent theory of Dr. N. B. Slater has shown that to a first approximation the rate of unimolecular decomposition of an activated complex with energy E in excess of a critical energy E* needed for decomposition is given by k(E) = (1 - f>"-l where Y lies in the range of normal frequencies (1012-1014 sec.-l) and n is the number of normal co-ordinates required for the description of the energy molecule decomposition.- E* + kT The so average that for any n v and with E* energy we can in calculate excess of the E* mean has life t ( c ) = I / K ( E ) . If we use mean values such as E*/kT - 40 ; v = 1013 sec.-l we find (with Slater) that t(E) w 10-l3 sec. if n = I ; t ( ~ ) E I O - ~ sec. if for n = I) are very much larger for low temperatures. Thus for (except e*/kT n = - 5 140 and at t ( ~ ) room E 10-4 temperature sec. if n = and 10. if we These set mean n = 5 life-times the mean ethane life of an excited molecule will be I O - ~ sec. These results imply that for the reverse reactions the recombinations of free radicals complexes may be formed having mean lives sufficiently long to assure even at very low pressures quite efficient stabilization by collisional deactivation.This will be true for the high-temperature pyrolysis reactions and will be more pronounced for photochemical re- actions where temperatures are considerably lower. A reaction to form the radical HO is one that is close to the permissible limits. Here n < 3 ; E* > 45 kcal. and the average complex formed at 130O C has a mean life t ( ~ ) < v(kT/E*)2 w 3 x I O - ~ sec. which is of the order of collision times a t S.T.P. At temperatures of 500' C to 600° C or higher this mean life is much shorter and i t is probable that most of the HO complexes decompose before they can be stabilized. In the more complex methyl radical the theory gives a reasonable accounting of the high collision efficiency which has been observed for their recombination.It will not take many degrees of freedom in a com- plex with such a high critical energy to give i t a very long mean life. Finally we may say that recombinations will manifest their three-body character when the complex has a low critical energy the temperature is high and the combining radicals are not complex. Dr. C. A. McDowell (Liverpool University) (communicated) I should like to mention that Dr. Thomas and I have found that in the presence of a large excess of oxygen the acetyl radical is quite stable up to about zoo0 c. Dr. A. D. Walsh (Leeds) said Were any analyses for formaldehyde carried out by Noyes and Marcotte ? A t the higher temperatures normally used for the study of methane oxidation it was difficult not to suppose the reaction of CH radicals and 0 to produce HCHO took place.Dr. Peter Gray (Cambridge) (communicated) I am very interested to read of these experiments on the peroxides and in particular of the role of the alkoxy radical RO. Peroxides are encountered as an important feature of some of the chemiluminescent (cool-flame) processes in hydro- carbon oxidation and the alkoxy radicals are presumably also present. Such RO radicals are also important in the thermal decompositions of the alkyl nitrates and nitrites. The initial step is the fission of the molecule to give the alkoxy radical and the appropriate oxide of nitrogen. Both these processes may be accompanied by a luminescence.41 Methyl McDonell and Thomas J .Chem. Soc. 1949 2208 2217. Gray and Y o f f e Proc. Roy. SOC. A 1949 200 114 fi Gray xg51 unpublished work. GENERAL DISCUSSION nitrite gives nitric oxide which may not be expected to oxidize the CH ,O- . Thus this chemiluminescence observed with nitrates nitrites and in the oxidation by nitrogen dioxide of the simple alcohols may be a property of the alkoxy radical itself. Furthermore from experiments on the nitrates values may be ob- tained under favourable circumstances for the heats of formation of the alkoxy radicals. The heats of formation of the other species involved and the energy required to break the 0-N bond are needed the latter may be identified with the energy of activation derived from measure- ments of initial rates.Such values may be employed both as a check on other measurements and as a means of calculating provisional values of heats of formation of peroxides of which the activation energies of decomposition but not the heats of combustion are known. The heats of formation of the methoxy and ethoxy radicals may be obtained from the following data on nitrate and nitrite esters. (i) CH,ONO,(g) = CH,O + NO ; Eact METHOXY COMPOUNDS Qj = 29.4 (ii) CH,ONO(g) = CH,O + NO ; Eact = 39-5 kcal. - 8.1 Qf(CH30) = - 2 , - 20.9 Qj(CH,O) = = 36'4 + 1.2 8 , - 8.1 Qf(C,H,O) = + z'o, The formation of each radical from its elements is numerically small and errors in any of the primary quantities appear magnified.The values (- 0.4 f 1.2) kcal./mole and (2 f 2) kcal./mole might be assigned to the CH,O and C,H,O radicals respectively. Qf = 16-7 ETHOXY COMPOUNDS CZH~ONO,(~) = C&',50 + NO ; Eact Qj = 33.4 Sir Alfred Egerton (ImperiaE College) said I can allay Dr. Stevenson's anxiety about the measurements of heat of formation of peroxide which are mentioned in the preliminary draft of the paper by myself Emte and Minkoff. The values found though the subject of careful measure- ments are unconvincing; for instance the heats of formation of the peroxides are in most cases greater than those of the corresponding more stable substances. It is not easy to determine accurately the heats of combustion of these explosive peroxides. We intend to redetermine them and to withdraw for the present the last section of our paper referring to the heats of formation of peroxides.Dr. A. G. Gaydon (Imperial College) said I should like to point out that the values for the dissociation energies of H, 0 and OH taken from my book are for oo K not 298OK. Fortunately the errors about 0.8 kcal./mole nearly cancel out for the equations as used. Dr. A. D. Jenkins (Courtaulds Ltd Maidenhead) (communicated) Vaughan ei al. report that vapour-phase pyrolysis of methyl-tert.-amyl peroxide (I) gives an 8 % yield of methyl ethyl ether which they ascribe to vapour phase combination of a methoxy radical with an ethyl radical. The scheme of decomposition of this peroxide will be substantially CH / CH,-O-O-G-CH -+ CHSO + O-C-CH = 39-5 kcal.CH / Gray and Yoffe J. Chem. SOC. 1950 3180. 312 GENERAL DISCUSSION bond Instead of the combination reaction suggested above i t is quite possible that the ether could be formed by reaction between the ethyl radical and the substrate by attack at the -0-O- CH3 CH3 CZHS + CH300C-CH3 -+ CH30CzHs + -0-C-CH3 / \ C2HS CZHS C2HS / \ Whilst no data are available for estimating the activation energy of this process from experimental results one can make a comparison with E for the alternative hydrogen abstraction reaction CH C,H + CH,OOC-CH / -+ C,H + radical Thus the general reaction \ in the following way. In the former reaction an -0-0- bond is broken to form an 0-C bond so that the exothermicity will be of the order (go-55)=35 kcal.and we can represent this process by the diagram FIG. I. In the second case the reaction is nearly thermoneutral and the bond broken is a C-H bond. Thus the corresponding diagram will be Fig 2. The crucial factors are the shapes and positions of the two repulsion curves but if they are not very different in the two cases the activation energy of the former process will almost certainly be lower than t h a t of the latter and therefore less than approx. 13 kcal. If the decomposition of this peroxide followed a chain type of reaction the chains would be very short if 8 % of the products resulted from the termination process so that the methyl ethyl ether may be at least partly formed by the reaction proposed here.R + R’OOR” -+ R’OR + R”0 or R”0R + R’O may occur in diaIkyl peroxide decompositions where R results from R’O -f R + CH,O or R”Q 3 R + CH,O. GENERAL DISCUSSION 313 It may be noted that nitrates or nitrites might undergo a similar reaction even more readily since the 0-N bond has a dissociation energy of the order 38 kcal. and the reaction will be even more exothermic than the corresponding peroxide reaction. Dr. L. Bateman and Mrs. Hilda Hughes (Welwyn Garden City) (communicated) The decomposition of cyclohexene hydroperoxide in solution in the temperature range 60-100’ C exhibits rather different features to those reported by Bell ei! al. and also by other workers for other hydroperoxides.718 Our investigations in this field are still pro- ceeding but certain results appear to warrant consideration here.In benzene solution the reaction is almost exactly second order. The non-formation of diphenyl and phenol indicates that the solvent is inert under our conditions. Cyclohexenone and R0,-double bond addition compounds appear to be the main organic products. Each mole of peroxide decomposing liberates nearly 0.5 mole of water and much less oxygen (- 0.1 mole). The Arrhenius activation energy is 26 kcal./mole. We conclude that the measured rate refers to the one-step bimolecular decomposition RO2-H + HO-OR + RO2- + H2O + RO-. - (1) Radical Radical t C C H5 H5 + H - - - - Radical I \ C H~ + CN 00ccc>15 /C H . . 3 Inactive products ’C H i FIG.2. In cyclohexene solution the decomposition proceeds faster (roughly five-fold) and the order is reduced to 1-7-1.8. These characteristics the isolation of cyclohexenol as the major product,O and a linear relationship between [R02H] /rate and I /[RO,H] are consistent with a chain reaction comprising (I) as initiation step and RO- + R’H + ROH + R’- R’- + ROJ3 + R’OH + RO- zRO-1 2R’- i ’ . - (4 * (3) * * (4) (5) (6) wheze in this case R and R’ are identical and the reasonable assumption 10 is made that kz = A&,. In the presence of stearic acid the reaction is greatly accelerated and becomeF of the first-order with respect to the peroxide. These facts correlate strikingly with the catalytic action of the hydro- peroxide on the oxidation of the parent olefin with molecular oxygen R’- + RO- Farkas and Passaglia J .Amer. Chem. SOC. 1950 72 3333. Kharasch Fono and Nudenberg. J . Org. Chem. 1951 16 113. Farmer and Sundralingam J . Chem. SOC. 1942 121. lo Cf. Bateman Gee Morris and Watson this Discussion. GENERAL DISCUSSION * roxidation = const. d([RO2HI2) 3 I4 Normally the relationship is obeyed ; in the presence of stearic acid this changes to (7) . (8) . Reaction (3) has not been considered by Bell et al. and an indication of its existence in their work viz. the formation of methanol4 from CMe,O,D and Me radicals has been otherwise explained. However we believe the following experiment provides strong evidence for its occur- rence. Cyclohexene hydroperoxide was decomposed in the substituted I 4-diene ethyl linoleate (R’H).R’ is thus a mesomeric pentadienyl radical which if formed will react preferentially to give a conjugated I 3-diene derivative. In fact conjugated diene units were produced they were located only in an ethyl hydroxy-linoleate and the formation of this substance was greatly in excess of that expected to result either directly or indirectly via radical-radical interaction. The alternative reaction . H j ( 7 . H !n R’- R’- + ROZH -+ RO2- + R’H has possibly been shown to occur by Bell et al. at 195’ in the gas phase reaction of Me and Et radicals with CMe,O,D. It should be noted how- ever that even in these systems the reaction is roughly 10 kcal. endo- thermic and this figure is increased to about 30 kcal.when R’H is an olefin and allylic resonance energy has to be supplied. At lower temper- atures and especially in unsaturated solvents (8) must be quite unim- portant compared with ( 3 ) which is about 30 kcal. exothermic (if Do-o = 40 kcal.). Reaction (3) is envisaged as a radical substitution reaction (SRz-to extend the Hughes-Ingold terminology) formally analogous to (I) cf. (-y with RO,-I-H + O-:-OR + 0-\--OR ui t-)! where the arrows denote one-electron displacements. The corresponding ionic processes involving heterolytic bond scission can obviously be re- garded likewise. We do not suggest that the decomposition mechanism now advanced is necessarily applicable to other hydroperoxides under different con- ditions-it is quite evident that analogy can be especially misleading in this field.Thus the effect of increasing temperature may be particularly critical in facilitating a unimolecular S,I relative to an S,Z decomposi- t ‘on because the increasing tendency of the comparatively weak 0-0 bond to rupture will be roughly paralleled by a decrease in the molecular association in the hydroperoxide (evidenced by hydrogen bonding) which is undoubtedly precursory to reaction (I). Nevertheless we do suggest that the relevant conclusions of Bell et al. require careful reconsideration in particular the postulation of unimolecular dissociation of tert.-butyl hydroperoxide in solution at temperatures not greatly above 100’ C. Not only are the rate data at 150’ C given in Table I1 of these authors’ paper consistent with our scheme (they correspond to a reaction order of 1-55) but Bolland and Morris 11 have made the very significant observation that this hydroperoxide in common with numerous primary secondary and tertiary allylic hydroperoxides catalyzes olefin oxidations in accordance with eqn.(7). Dr. W. E. Vaughan (EmeryvilZe California) (communicaded) The questions raised by Dr. Bateman and Mrs. Hughes concerning our paper seem to u s to be well answered by the experimental evidence presented namely the isolation of methane-d and ethane-d and of ethyl tert.-butyl l1 See Quart. Rev. 1949 3 I. GENERAL DISCUSSION 31 5 peroxide from the appropriate reactions. These actual product isolations in substantial yields provide definite proof for the reactions R' + (CH,),COOD -+ RD + (CH,),COO* R.+ (CH,),COO' -f (CH,),COOR . and and they are not otherwise explained by the mechanism proposed by Dr. Bateman and Mrs. Hughes. R* + - CHdH-CH2-CHSH- . It is our opinion that reaction (I) is not endothermic as claimed by Bateman and Hughes but rather is exothermic to the extent of about 20 kcal./mole when R is methyl. No quantitative evidence is available but qualj tative comparison of ted-butyl hydroperoxide as a hydrogen atom donor with cyclohexene and formaldehyde leads us to the conclusion that Do-,(ROOH) is of the order of 80 kcal./mole rather than 105 or 118 kcal./mole as observed in certain alcohols and in water. Further proof for the chain decomposition we have proposed may be obtained from a comparison of the rates of decomposition a t 140~ of tert.- butyl hydroperoxide in a reactive solvent n-octane and in an inert medium chlorobenzene.In the latter solvent reaction (3) cannot intervene R'* + ROOH -f R'OH + RO' because decomposition leads to quantitative yields of tert.-butyl alcohol and oxygen. Moreover in this solvent the rate is some 40 times that in n-octane. (A) . Bateman and Hughes have acknowledged the difficulties in drawing analogies between decompositions of various peroxides under differing conditions. However the decomposition of cyclohexene hydroperoxide in an unsaturated solvent can also be interpreted by our mechanism which has the evidence of actual isolation of a stable peroxide intermediate.For ethyl linoleate this intermediate would correspond to (B) ; -+ RH + -CH=CHL:CHLCHCCH- OOR ROO' + (A) -+ -CHdH-CHdH-CH- I (B) 0. I (B) -+ KO' + -CH=CH-CHeCH-CH- - (1) ' (4 * (3) OH (4) (5) (6) (7) (C) (C) + ROOH (or RH) -+ ROO' (or R.) + -CH=CH-CH=CH-CH- The allylic stabilization of (A) would increase the possibility for the postulated association reaction (2) at the expense of metathetical reactions such as (I) or ( 3 ) . I The decomposition of cyclohexene hydroperoxide in cyclohexene can also be interpreted by a strictly analogous mechanism. However we think that in an inert solvent the reactions of the hydroperoxide group are greatly complicated by the presence of reactive methylene groups and the double bond in cyclohexene hydroperoxide.Although our in- formation offers nothing either pro or con concerning reaction (3) we see no objection to it on purely theoretical grounds and have in fact indicated that the processes proposed in our paper do not exclude other possible competitive transformations which may be proceeding simul- taneously and undetected. Dr. W. A. Waters (Oxford) said Whilst we all realize the great value of the kinetic studies of olefin autoxidation carried out by members of the British Rubber Producers' Research Laboratory I would suggest that there will be little value in increasing the precision and detail of GENERAL DISCUSSION their present types of kinetic work beyond the stage which it has now reached.Kinetic studies of chain reactions give velocity equations which are dependent very much on one particular chain-ending process e.g. 2R' 3 products or R' + ROa' +- products (dimers often unspecified). When however organic chemists start looking for these diagnostic compounds in autoxidized mixtures they seldom find them and sometimes can show that the postulated dimers are far too reactive to persist in the final oxidation product. For example in our current study of the effects of phenols on benzaldehyde autoxida- tion 12 Mr. C. Wickham-Jones and I deduced kinetically that the chain- termination product in mixtures containing +-cresol was a dimer of a mesomeric phenoxy radical (ArO*)2. We have failed as yet to isolate any of the known dimers of the tolyloxy radical from our reaction product and have found that the material which is eventually formed must be more complex.Moreover by adding the known dimers separately to our reacting system we have shown that they have a comparable reactivity to the p-cresol and so could not persist throughout the whole autoxidation. With 2 6-xylen-1-01 which gives similar kinetics the obtainable reaction product is not the phenol dimer but its further oxidation product. the corresponding diphenoquinone ; even this is an autoxidation retarder which therefore must be converted to still another product. One important point that kinetic investigators should remember is that the chain-terminating reaction which will be revealed by their reaction velocity measurements is the one which most rapidly removes active free radicals from the system.None of a whole series of slower radical-removing processes or of secondary reactions involving products which initially are present in very low concentration may be able to influence the reaction velocity to a discernible extent unless the experi- mental conditions are especially designed so as to detect one of them. " Initial velocity I' measurements in particular are of little help in the diagnosis of the chemical identities of products of chain termination. In reactions involving hydroperoxides it is important to remember that if the system becomes appreciably acidic or basic then an ionic decomposition of the hydroperoxide may set in to give products which are often quite different to that of its homolytic fission.The homolytic fission seems to proceed + 'OH RlR2R3G-O-OH -+ RlR,R3C-O' L RlR2C==0 + 'R3 where the fission rules for the breakdown of the ROO radical are in general those given by Walsh.lS The acid-base catalyzed reaction however proceeds RlR2R3C-0-0-H ways. + H+ -* where R is the group that most easily carries with it an electron pair. These different products influence the subsequent reactions in different la J . Chem. SOC. 1951 812. l3 Trans. Faraday SOC. 1946 42 269. GENERAL DISCUSSION 317 Dr. A. J. Harding (Cambridge) said A good approximation to the kinetics of hydrocarbon oxidation may be derived without postulating intermediates of definite types such as peroxide or aldehyde.It is only necessary to assume that the chains consist of links which are alternately reacting with oxygen and with the hydrocarbon that the chains are terminated by destruction of the radicals which would otherwise react with hydrocarbon and that the chains are initiated by the breakdown of an intermediate which they themselves produce. The use of special features such as the ratio pmax/Apmax becomes essential if the finer points of the kinetics are to be evaluated. One of these points is the type of degenerate branching. Dr. Mulcahy has described the type in which the reaction fails to reach an infinite rate (i.e. ignition) because the re- actants are consumed before the intermediate concentration becomes sufficiently great.Another type of degenerate branching has been suggested for the oxidation of hydrocarbons at higher temperatures. In this second type the intermediate (e.g. aldehyde or peroxide) which generates chains is itself oxidized by the radicals i t produces. Since the production of intermediate is proportional to the first power of its concentration and the destruction is proportional to the second power a limiting concentration of intermediate is reached. This means that even if no reactants are consumed in the process the reaction attains a limiting finite velocity. A distinction between the two types of degenerate branching is readily made. In the first case addition of intermediate at the beginning of the reaction will give an increase of maximum rate as well as a decrease of induction period.In the second addition of intermediate will reduce the induction period but will produce no change in the reaction velocity until so much intermediate has been added that the induction period has been completely eliminated. The oxidation of ethylene at temperatures in the region of 400" C with formaldehyde as intermediate shows the behaviour expected of the second t41~e.l~ A derivation of the variation of pmax/Apmax with hydrocarbon con- centration for the second type of degenerate branching leads to a relation- ship approximately the same as that found by Dr. Mulcahy for the first type. This relationship does not therefore provide a clue to the type of degenerate branching although i t does give valuable support to the thesis that degenerate branching is taking place.l4 Dr. Peter Gray (Cambridge) (communicated) May the fact that the propylene oxidation studied by Mulcahy appears to occur through a peroxide which may react in such a way as not to produce active radicals be associated with the characteristics of propylene as a chain-breaking agent ? (In this case the constant C might include a term corresponding to the propylene concentration.) Dr. M. F. R. Mulcahy (Melbourne) (communicated) Dr. Gray's sug- gestion would require the chain-terminating effect of propylene to be specific towards the radicals produced immediately on decomposition of the peroxide. The participation of the hydrocarbon in a chain-terminating reaction in the main cycle would affect the second rather than the third term in eqn.(4) (tending to make B independent of [RH] under conditions of excess oxygen). However i t now appears that the significance of the value of C derived from the experimental results is somewhat complex. Some very recent experiments carried out in this laboratory by Mr. Ridge with propylene at 288" C have shown that the intercept of the pmax/Apmax against [RH] curve on the pmax/Apmax axis is dependent on the surface conditions and may become positive. It seems therefore that the simple interpretation of this intercept (C) given in my paper is in need of some amplifications. 1 4 Harding Thesis (Cambridge 1948). Norrish X V I Int. Coll. C.N.R.S. (Paris 1948) p. 16. Harding and Norrish (in preparation).GENERAL DISCUSSION Dr. E. J. Harris (London) (communicated) In 1935 l5 i t was suggested that compounds known to promote the branched chain reaction in the slow oxidation of hydrocarbons such as ethyl nitrite and ethyl peroxide did so by providing -OR radicals and the converse process a condensa- tion of -OR to peroxide seemed not unlikely. In 1938 Neumann and Tutakin l6 showed that small additions of diethyl peroxide to butane + oxygen mixtures would cause the appearance of a flame similar to the cool flame seen during the oxidation reaction. when induced in hydrocarbon + oxygen mixtures. The result flame does not inevitably promote hydrocarbon oxidation even quent reaction. This is explicable on the basis of radical formation. 318 To put the hypothesis of the intervention of peroxide on a frrmer basis two approaches are possible namely the physical and chemical investiga- tion of the slow oxidation and the study of the peroxides themselves under appropriate conditions.In 1936 Sir Alfred Egerton suggested the latter method as a profitable one. Before mentioning a few relevant properties of peroxides i t will be useful to state the analytical evidence concerning peroxides isolated from oxidation reactions. At 320-270' propane and butane can be made to yleld considerable quantities of hydrogen peroxide making up 10 yo of the total condensate. The H,O combines with two molecules of form- aldehyde one of the other main products and in one experiment l7 0.5 g. of the compound was isolated.If however the reaction is carried out when the walls of the vessel are slightly contaminated with alkaline material the reaction no longer leads to the formation of the peroxide though all the other products are nearly the same. This and other evidence suggested that the H,O,-formaldehyde compound was only formed after condensation. The fact that H,OB can be isolated indicates that radicals -OH or -0,H intervene in the slow oxidation the sensi- tivity to the state of the wall may mean that hydrogen peroxide is only formed when collisions involving the wall permit it. The analytical problem of showing the presence of an alkyl peroxide is of course com- plicated by the presence of a large amount of H202. Results of the study of the alkyl peroxides 18 19 2o and their influence on slow oxidation reactions are consistent with the following (a) They may be formed in traces as by-products but chemical evidence is lacking.(b) The alkyl hydrogen peroxides like H,O, are very sensitive to the state of the surface being decomposed by traces of alkali. They will ignite giving a diffuse blue luminescence when introduced with oxygen into a vessel at 250-3ooo. The luminosity resembles that sometimes seen in the hydrocarbon oxidation and is general rather than flame-like. The dialkyl peroxides will decompose with a blue flame in certain conditions of temperature pressure and gas mixture. Below the critical limits they decompose homogeneously. The artificial blue depends upon the surface and the temperature.( d ) Traces of the peroxides reduce the induction period of the slow oxidation of hydrocarbons without altering the rate of the subse- The small quantities of peroxide which are needed to produce visible effects make i t not surprising that they have not been isolated. Only 0.007 cm. of dipropyl peroxide will give a flash with air at 270° and 0-01 cm. 16 Egerton Smith and Ubbelohde PhiE. Trans. Roy. SOC. 1935 243 433. 16 Neumann and Tutakin Acta physicochim. 1938 9 861. l7 Harris and Egerton Chem. Rev. 1937 21 287. 1s Harris and Egerton Proc. Roy. SOC. A 1938 168 I. l9 Harris Proc. Roy. SOC. A 1939 173 126. 20 H a r k Proc. Roy. SOC. A 1940 175 254. GENERAL DISCUSSION 3 I 9 of ethyl hydrogen peroxide can be seen to luminesce when oxygen is present.In spite of the explosive properties of the peroxides the alkyl hydrogen peroxides like H,O,,21 can survive quite high temperatures if the vessel has a low surface/volume ratio and particularly if a carrier gas is present. Thus at 180' only 5 yo ethyl hydrogen peroxide was decomposed when passed with N through a tube at 180' (contact time 4 sec.). A t 320° z yo survived. Under somewhat similar conditions ( I sec. contact time) 85 yo H,O sometimes survived. Post-war work on hydrocarbon oxidation (e.g. Hinshelwood 22 Mulcahy 23) appears to be in accord with the scheme outlined by Egerton and Harris 24 in 1938 in which i t was proposed that the initial step was peroxide forma- tion followed by splitting to give radicals -OR -OH and -O,H which carry on the main reaction.Analogous schemes were outlined by Ubbel~hde,,~ Pease 28 and Lewis and v Elbe.,' Egerton and Harris however did expressly suggest (cf. their Table 11) the reaction z -OH -f H,O as a chain-terminating one and this together with a reaction of -0,H with a hydrogen compound would explain hydrogen peroxide formation. Organic peroxides as such were only invoked in the initiation of the reaction and this agrees with the fact that surfaces active in decomposing the alkyl hydrogen peroxides (e.g. salt-coated ones which under the experimental conditions become alkaline) are those which give rise to long induction periods for the slow oxidation reaction. Dr. G. J. Minkoff (ImperiaE College) said I should like to raise two points in connection with Dr.Mulcahy's paper. First with regard to the evidence of Badin which is quoted in support of the formation of hydroperoxides in oxidation Badin observed a line at 11.4 p in the infra- red spectrum of tert.-BuOOH; he concluded that since H,O also has a line at 11*4p this line is characteristic of the -0OH grouping. This is unsound for several reasons; in many observations of the infra-red spectrum of freshly prepared ten!.-BuOOH I have only once found a line at 11-4p and that was in an impure sample ; the position of the line was otherwise found to be at 11.6~. Di-tert.-butyl peroxide does have a line at 11*4p so that clearly the line cannot be characteristic of -0OH ; in fact as was pointed out before i t may be connected with both the -04- group and with the tevt.-butyl group.The line which Badin found may well have belonged to H,O,. The other point is connected with the attempt t o draw up a mathe- matical expression for the rate of pressure change. This is based on the proportionality observed between peroxide concentration and the rate of pressure change. In the derivation in terms of initial concentrations the tacit assumption is made that P is the hydroperoxide of the original hydrocarbon. However this assumption may not be correct because in the degradation of a long cha'n paraffin several steps must occur all prob- ably involving peroxides. In the calculations made by Bolland this ob- jection does not arise since only one peroxide is involved. Another complicating factor is that the different peroxides formed (i.e.primary secondary peracids etc.) will react to different extents with the potassium iodide reagent; thus the maximum peroxide concentration may be masked by the lack of reaction with KI of some of the peroxide present. Dr. W. A. Waters (Oxford) (communicated) Sir Cyril Hinshelwood's review of the influence of substituents on the rate of hydrocarbon oxida- tion can be given an alternative interpretation which is equally consistent ar Harris Trans. Faraday Soc. 1.948 9 764. 2 Hinshelwood Faraday SOC. Dascussaons 1947. z8 Mulcahy Trans. Furuduy Soc. 1949. 45 575. ,4 Egerton and Harris Proc. I8me Congr. Chim. Ind. (Nancy 1938). 2s Ubbelohde and Egerton Proc. Boy. SOC. A 1935 152 354. 26Pease J .Amer. Chem. SOL 1929 51 1839 et seq. 27 Lewis and v. Elbe J . Amer. Chem. SOC. 1937 59 976. 320 GENERAL DISCUSSION with the theories of general polarity. The substituents which he finds to increase oxidation rates are also those which promote the attack of methyl radicals on C-H groups. Kharasch and his colleagues for in- stance have reported that methyl radicals from the decomposition of diacetyl peroxide preferentially attack C-H bonds vicinal to C-C1 and also CH-CO- CH-0- CH-CO-OMe but do not attack CH groups. From studies of tert.-butyl peroxide it now appears that R-0' radicals (e.g. Me,C-O.) have not quite the reactivity of alkyl radicals such as methyl and are even more selective in the same sense in their reactivities towards C-H groups.28 I would therefore suggest that the reaction which immediately follows Sir Cyril's chain initiation process R-0-0-X + R-0' + '0-X i.e.R-0' + H-C- -+ R-0-H / \ (iii) R-M + M 3 R-M-M' groups to build up high polymers. + 'G- / \ may be much slower and much more selective than the subsequent stages of the gas-phase oxidation involving R' and R-0-0. radicals. It may well be rate-determining for the whole breakdown of the oxidized compound. A similar state of affairs is well known in polymerization chemistry where the sequence (ii) R' + Monomer -+ Hydrocarbon type radical R-M' ; (i) Catalyst (eg. benzoyl peroxide) -+ Radicals R' ; etc. (fast chain growth) results in a polymerization kinetically dependent upon the rates of both processes (i) and (ii) as for instance in the benzoyl peroxide catalyzed polymerizations of both styrene and vinyl acetate.In polymer chemistry again the stability of CH in comparison with A H z - is strongly marked as for instance in the ready chain transfer to -CHz-CH=CH- and the converse tendency of compounds with CH,-C=C In connection with later states of hydrocarbon oxidation i t may be noted that in the Dyson Perrins Laboratory we have recently shown that in the liquid phase the breakdown (1) Alkyl-CO' + AllrJtl' + CO occurs quite readily a t temperatures as low as SO-IOO~ C. This is con- firmatory evidence for the view that in hydrocarbon oxidation at elevated temperatures only the initial stages of the reaction sequence are rate- determining.Dr. A. D. Walsh (Leeds) said It is helpful in considering the kinetics of oxidation reactions of hydrocarbons in the gas phase to write the step that produces the alkyl peroxide radical in the form R + 0 = RO,*. Commonly the * is omitted but its inclusion serves to remind us that in the first instance the RO radical is energy-rich. This is important for at least the following reasons. (i) It makes it less likely that the reverse reaction to (I) will be neglected in considering the full kinetic scheme. (ii) The excess energy is presumably distributed over various vibrational degrees of freedom. The radical is therefore to be thought of as vigorously twisting turning and generally distorting itself far more than does a " normal " radical.This means that subsequent reactions which axe rather more complicated than are usually found in gas-phase oxidation processes are rendered more pIausible than they would otherwise be. Cf. also Farmer and Moore J . Ckem. Soc. 1951 131. 321 GENERAL DISCUSSION Those gas phase chain reaction steps which are best substantiated commonly' fall into one of 4 classes 29 ( a ) those in which one linkage is broken e.g. CH,CH3 = CH + CH CH + NO = CH,NO* . - (2) ( b ) those in which one linkage is formed e.g. (I) above the reverse of (2) and * (3) (G) those in which one linkage is formed and one is broken e.g. abstraction of an atom (particularly H) from a molecule by a radical as in * OH + RH = H,O + R; If this is re-written as . RCH,CHO + 0 = RCHO + HCHO CH + 0 = HCHO + OH.CH3 + 0 = [C,HO,"] = HCHO + OH. (4) ( d ) energy transfer reactions e.g. reactions involved in the quenching of fluorescence A * + B = A + B * . As illustrations of these four types it is instructive to look at the paper by Bell Raley Rust Seubold and Vaughan. The great merit of that paper is that some experimental evidence for each step postulated has been obtained. Consideration of these steps shows that they all con- form (or in one case can readily be slightly altered to conform) to the above types. - ( 5 ) In other words the steps occurring most commonly appear to be the simplest. This should not be surprising for in postulating a reaction step i t is vital to consider just how the reactants could be converted into the supposed products; and such consideration makes it difficult not to conclude that more complicated steps than (a) t o (a) are likely to have unfavourable steric factors or activation energies.30 If a postulated gas- phase reaction step does not conform to one of the 4 types one cannot say it is impossible-but one can and should demand that the evidence for it be unusually strong unless it is to be dismissed as mere speculation.This is incidentally a serious criticism of many of the reaction steps postulated by certain authors. To take but one example the supposed reaction involves the simultaneous breaking of two linkages and the forming of two linkages. Unless particularly strong evidence in support of this was forthcoming-and such has certainly not yet been produced-this postulated step can be taken as very unplausible.In the special case of a reaction such as (I) however i t is not unplaus- ible to suppose a subsequent reaction for RO,* more complicated than types (a) to (d). An example is the reaction it is split into simpler steps and the fact that the second involves the breaking of two linkages and the forming of one does not appear un- plausible. Another example might be a somewhat complicated isomer- ization of RO,*. (iii) Even if the reaction supposed subsequent to (I) conforms to (a) to (d) inserting the * in (I) helps one to remember that the activation energy' for the following reaction need not be as great as would be the case for " normal " RO,.Some or all of the activation energy may be provided by the energy produced by (I). The step RO + RH = ROOH + R - (6) 29 Cf. Chamberlain and Walsh Rev. Inst. Frangais du Pe'trole 1949 4 307. 3 0 Cf. Ubbelohde Rev. Inst. Frangais du Pe'trole 1949 4 315. 31 Norrish this Discussion. L GENERAL DISCUSSION HO,* + NO = EIO + NO -/- 0 322 although probably exothermic when RH is an olefin is probably endo- thermic when RH is a paraffin.32 The activation energy required for (6) in gas phase reactions however need not be as great as this endothermicity. The idea is closely similar to that of (ii). The energy of formation of RO, carried over into a subsequent reaction reduces any improbability of that subsequent step. To realize this is to recognize that the process of gas-phase peroxide formation may’ be subject to inhibition by inert gases.The greater the dilution of the reactants the greater the chance that collisions (and reactions of type (a)) will take so much energy from RO,* that the probability of (6) is seriously reduced. Peroxide formation is involved in cool flame phenomena. To test whether inhibition by inert gases of cool-flame formation can occur however one needs to choose conditions carefully’. At temperatures near the lower temperature limit of spontaneous cool-flame formation the dominant chain-ending reaction is usually a surface destruction of radicals and the inhibition may be swamped by a greater promotion viz. the hindering of radical diffusion to the walls. Along the upper temperature limit of cool flame formation however the dominant chain-ending reaction is a gas phase one (in most cases the decomposition of a bulky alkyl radical to an olefin and a small alkyl radical 33 3ii).At such temperatures for diethyl ether (Fig. 3 of ref. (33)) and for propane (Fig. 11 of ref. (35)) a small inhibition by’ inert gases is observed. It seems probable that this inhibition represents the expected effect of reduction of the energy content of RO,* though of course i t has to be remembered that cool-flame formation3(j and pro- pagationa7 involve a thermal factor and inert gases may also inhibit purely by virtue of their thermal properties. Finally i t is important to be clear about the nature of the excess energy carried by RO,*.This is surely vibrational in nature. For the analogous HO,* radical however the following reaction has been supposed to occur 38 the HO,* radical transferring the whole of its energy to the NO,. Such a complete transfer of energy seems only likely for electronic energy. Yet the effects of inert gases (M) on the second limit of the €-I + 0 ex- plosion are usually supposed to require the transfer of energy from HO,* (i.e. H + 0,) to M ; and the different gases that are effective make i t difficult to suppose other than vibrational energy is concerned. Some confusion seems therefore to exist. Dr. C. A. McDowell (Liverpool) (communicated) I was interested to read Dr. Walsh’s remarks on the existence of RO,* radicals i.e. RO radicals endowed with excess energy.Similar views were expressed by Dr. Thomas and myself in our paper on the inhibitory effect of nitrogen peroxide on the gas-phase oxidation of acetaldehyde. 39 In considering possible subsequent reactions of RO,* radicals one has to be extremely careful for the excess vibrational energy with which these radicals are endowed may or may not persist throughout the numerous collisions which these molecules may make before they undergo reaction. If the excess of vibrational energy does not persist then ob- viously little is to be gained by ma’ntaining the above notation. Numerous 32 Walsh J . Chem. SOC. 1948 339. 33 Chamberlain and Walsh 3rd S y m p . Combustion Flame and Explosion Phe- nomena (Williams and Wilkins Baltimore 1949). 34 Walsh in course of publication.35 Malherbe and Walsh Trans. Favaday SOC. 1950 46 835. 36 Malherbe and Walsh Trans. Faraday SOC. 1950 46 824. 37 Spence and Townend 3rd S y m p . Combustion Flame and Explosion Phe- nomena (Williams and Wilkins Baltimore 1949). 38 Dainton and Norrish Proc. Roy. Soc. A 1941. 177 395. 3% McDowell and Thomas J . Chem. Soc. 1950 1462. GENERAL DISCUSSION 323 cases are however known where vibrational energy does persist through- out a time interval during which a molecule may make as many as 5 x 104 collisions before one quantum of vibrational energy is dissipated. Theoretically there is no reason why vibrational energy should not be transferred from one molecule to another provided their vibrational levels are sufficiently close together to permit resonance to occur.Such cases are well known and it should perhaps be pointed out that it is not necessary that all the excess vibrational energy should be transferred from one molecule to another ; in fact it seems probable from elementary theoretical considerations that the excess energy is more likely to be transferred in small amounts in successive quanta. One other type of transfer which must be borne in mind is the transfer of vibrational energy from one molecule into translational energy of another. This is most easily under- stood in the case of the transfer of vibrational energy from vibrationally excited molecules to rare gas atoms. This latter type of transfer probably also occurs with light molecules. RCH,CHO + O=RCHO + H,CO was originally postulated as a source Prof.R. G. W. Norrish. (Cambridge) (communicated) The reaction of excited formaldehyde to account for the luminosity of the cool flame. In view of the very low luminosity observed (ca. I quantum per 10 mole- cules of hydrocarbon reacting) its probability in comparison with the other radical reactions would be very low in accordance with Dr. Walsh’s expressed view. However I have as will be observed in my paper dissociated myself from this reaction except as a possible component in the dark blue flame and even then the suggestion must not be taken too literally for the reaction may occur in two stages RCH,CHO + 0 = RCH CHO + OH RCH . CHO + OH = RCHO + H,CO. In any case I am not fully able to agree with limitations which Dr.Walsh would lay down for chain processes. Not enough is yet known about the possible configurations in the transition state. I agree with his remarks about RO,*. They conform to the views we have already expressed with reference to HO ; see for example the footnote in my paper or Axford and Norrish on the oxidation of formal- deh~de.4~ Prof. A. R. Ubbelohde and Mr. Small (Belfast) (communicated) The suggestion that RO,* is in the first instance “ energy-rich ” and is more effective in reaction ( 2 ) above if i t can use this energy before i t is randomized by collisions etc. has an interesting experimental corollary. It has recently been observed 41 that molecular hydrogen inhibits certain reaction chains in both aldehyde and hydrocarbon oxidations much more markedly than molecular nitrogen.Tentatively this may be attributed to the efficiency with which hydrogen can effect the transfer between internal and translational molecular energy in collisions. RO,* + H + H,* + RO,. In (pseudo) unimolecular reactions the special efficiency of hydrogen collisions in the reverse process translational + internal is well known and from the principle of microscopic reversibility collision process (3) would also be expected. Dr. A. D. Walsh (Leeds) (communicated) Prof. Sir Cyril Hinshelwood refers to my suggestion that bonds between strongly electronegative elements should be strengthened by electron-repelling groups. As originally put forward this was based upon the possibility of increas- ing the overlap of the two atomic orbitals concerned in a single bond 40 Axford and Norrish Proc.Roy. SOC. A 1948 192 518. 41 Small and Ubbelohde J . Chem. Soc. 1950 723 and unpublished results. L* 324 GENERAL DISCUSSION between the elements. It is perhaps worth pointing out however that the suggestion could also be based upon the possibility of reducing lone pair- lone pair repulsion between the 0 atoms by attaching electron-repelling groups. The interaction between lone pair electrons on a halogen atom and electrons of a neighbouring system is greatest for the most electro- negative halogen (F) and least for I.43 A similar statement applies to interaction involving lone pair electrons on a Group VI eIeme11t.4~ The interaction between lone pairs on the adjacent 0 atoms of an -0-O- group is expected to be repulsive (witness e.g.the skew nature of the H,O molecule). Attaching electron-repelling groups to an 0-0 group will reduce the effective electronegativity’ of each 0 atom and so be expected to reduce the lone pair-lone pair repulsion i.e. be expected to increase the 0-0 bond strength. Dr. N. Uri (Manchester) said In the primary step relating to the oxidation of hydrocarbons as postulated by’ Prof. Norrish viz. RH + 0 -+ R + HO the endothermicity and activation energy are likely to be of the order of 40-70 kcal. if the dissociation energy of the HO radical into H + O Z ( D ~ o 2 ) is as low as 36 kcal. This latter value is obtained 4 4 9 45 from the electron affinity (in solution) of the HO radical which is in turn evaluated by inter- polation from the energies corresponding to the absorption maxima of various ferric ion pairs.While some unrecognized error may cause our quoted value of D H O ~ to be a few kcal. low it would be difficult to account from hydrocarbon chemistry for a value of D H O ~ as high as 65 k~a1.,4~ unless i t also assumed that the activation energy of the first step of the oxidation of hydrocarbons is much larger than its endothermicity. Dr. A. J. B. Robertson (King’s College London) said The heat evolved in the formation of HO may be determined by the method outlined by’ Stevenson in this Discussion. I find that hydrogen peroxide is dissoci- ated by electron impact to give the HO,+ ion with a small probability.The appearance potential is 16.1 eV. The ionization potential of H,O is found to be 12.1 elr. The ionization of O, HO and H,O very prob- ably involves the removal of a weakly anti-bonding or non-bonding electron located approximately on an oxygen atom. In these circum- stances we may provisionally estimate the ionization potential of the HO free radical as 12.2 eV intermediate between that of 0 and H,O, and i t may be less. This assumption gives 46 kcal. for the heat evolved in the formation of HO from H and O, subject to an experimental uncer- tainty a t present of about g kcal. If the ionization potential of HO is lower than 12-2 eV the formation of HO is even less exothermic. Dr. R. Spence (Harwell) said Prof. Norrish has raised many inter- esting points.I have always thought that the phenomena of com- bustion are so numerous and varied that general mechanisms should not be pressed too far unless all the well-established facts have been taken into account. For instance I do not think that the isolation of aldehydes as products of the reaction is necessarily evidence that they occur as essential intermediates. This was one of the important lessons to be learned from the early work of Bone who tried to establish the hydroxylation theory of combustion by the identification and separation of a set of products which were regarded as intermediates. Formalde- hyde for instance is a product of the oxidation of acetylene at 3 o 0 - 3 ~ 0 ~ C.4’ It is known that its rate of reaction with oxygen at these temperatures la See Baker and Hopkins J .Chem. SOC. 1949 1089. 43 Unpublished work of Baker and Barrett. 44 Evans and Uri Trans. Faraday SOC. 1949 45 224. 45 Evans Hush and Uri (in the course of publication). 46 Walsh J . Chem. SOC. 1948 331. 47 Spence and Kjstiakowsky J . Amer. Chem. SOC. 1930 52 4846. 48 Spence J. Chem. SOC. 1936 652. GENERAL DISCUSSION 32 5 is quite slow even at moderate partial pressures. Thus ordinary formaldehyde cannot be an intermediate in the combustion of acetylene ; i t can only be an intermediate if i t reacts whilst still in the excited state and i t is important as Dr. Walsh has mentioned earlier to recognize this in the mechanism. Then as regards the OH radical I have referred earlier in this dis- cussion to the work of Dr.Bates and myself 49 on the reaction between methyl radicals and oxygen and to the possibility of the formation of OH radicals or of CH,O in the primary step. More recently Blaedel Ogg and Leighton 5 O failed to detect the OH radical in this reaction by optical methods so that if present its concentration must be very small. Another piece of evidence which is not easy to fit into a simple radical chain theory was obtained by my wife 51 when working in Prof. Townend's laboratory. The propagation of a cool flame through a vertical tube at room temperature containing acetaldehyde ether hexane or heptane and oxygen or air is determined purely by thermal considerations. Free radical chain effects usually associated with the vessel wall and with the addition of inert gases were not observed.The addition of methyl radicals or of iodine produced no measurable effect ; the only indication that a chain mechanism might be operating was the inhibitory effect of additions of NO,. Another experimental result which is not easy to reconcile with a mechanism involving oxygen atoms is the homogeneous slow combustion of formaldehyde in reaction vessels of only I mm. diam.48 In this case the chain carriers must be relatively unaffected by the wall. The surface oxidation only predominates when the vessel is packed with powdered glass. The " dark blue " flame mentioned by Prof. Norrish is to be identified with the " blue" flame of Townend and co-workers 51 which follows in the path of the " cool " flame.Prof. R. G. W. Norrish (Cambridge) (communicated) There can be no question that aldehydes are formed as intermediates in hydrocarbon oxidation because they are readily observed the question a t issue is whether they are to be identified as the moderately stable intermediate Iesponsible for the delayed branching. I have given the reasons in my paper for supposing that they are particularly in the second paragraph I do not think that more weighty evidence has yet been produced for any other intermediate product playing this role. It must be remembered that the process of degenerate branching also takes place at temperatures where peroxides are completely unstable and unobservable and that Semenov's hypothesis demands a finite measureable concentration.The possible origin of aldehydes via a transitional peroxide is not ruled out on this account e.g. RCH + 0 -+ [RCH,O] + RCHO + OH as I have indicated in my paper. I do not feel that the remarks about the oxidation of acetylene are relevant. My paper did not concern this reaction but i t may be men- tioned that the rate of oxidation of formaldehyde at 286" C is appreciable (for partial pressures of formaldehyde and oxygen of IOO mm. in a vessel of 23-6 mm. diam. and volume 80 cm. the rate of reaction as measured by pressure change is 1-9 mm./rnir~.).~ Further i t is highly probable that the formaldehyde product in the acetylene oxidation is excited and thus more highly reactive. However I do not wish to commit myself to any view concerning acetylene oxidation in the present remarks.49 Bates and Spence J . Amer. Chem. SOC. 1932 53 1689. 50 Blaedel Ogg and Leighton J . Amer. Chem. SOC. 1942 64 2499. 51 Kate Spence Thesis (Leeds 1945). 52 Axford and Norrish Proc. Roy. SOC. A 1948 192 518. 326 GENERAL DISCUSSION As shown by Dr. Reagh and myself,5s the surface effect in the slow oxidation of hydrocarbons only becomes apparent at a certain limiting diameter of about 5-10 mm. The theory of slow branching requires that the reaction shall be entirely inhibited a t a finite diameter when the net branching factor becomes less than zero by the increase of the surface deactivation. This was found to be the case for methane ethane propane acetylene ethylene and propylene.We should not expect the cool flame to be affected until a limiting diameter of tube was reached. Finally with regard to the oxidation of formaldehyde this shows none of the criteria of a degenerate branched reaction and the results have been explained by us in terns of a straight chain mechanism. We should not expect the dramatic results observed with hydrocarbons on decreasing the diameter. There is however a competition in chain ending between the volume and surface reactions in which the surface effect only becomes predominant a t very small diameters ( < I mm.) corresponding to packed vessels. Nothing here appears to me to be inconsistent with our inter- pretation of the oxidation of formaldehyde as a straight chain reaction. With regard to the “ dark blue ” flame i t will be seen that I have drawn attention to the similar flames in acetaldehyde observed by Townend and his co-workers in my paper.Dr. G . Porter (Cambridge) (communicated) As far as I am aware no short-lived free radical has ever been detected by its absorption spectrum in a chemical reaction at normal temperatures nor in an ordinary photo- chemical reaction despite numerous attempts with systems where radical mechanisms are known to be operative. I think it important to bear this in mind when assessing negative evidence of the kind mentioned by Dr. Spence though the workers quoted used a particularly sensitive method and their results certainly suggest a short lifetime of I O - ~ sec. or less of the OH radical if it was present.By using the method of flash photolpsis54 very high instantaneous concentrations of intermediates can be obtained and I have been able to observe the OH radical in absorption at high intensity in two systems which almost certainly involve the reaction between methyl and oxygen viz. the photochemical oxidation of acetone and the chlorine photo- sensitized oxidation of methane. Unfortunately even this positive result does not enable one to conclude that OH is formed in the primary reaction of methyl with oxygen rather than in the subsequent steps though i t is hoped that studies of the kinetics of OH appearance will make this dis- tinction possible. Dr. L. Bateman (Welzuyn Garden C i t y ) (communicated) Norrish’s contention that peroxides do not absorb light and are therefore photo- chemically inert in the wavelength range 3000-4000 A is misleading and his arguments as presented have not the force claimed.Admittedly peroxides ranging from about 0.5-10 at 3000 to 0-01-1 at 3650 A but such absorption is weak E of typical saturated and allylic hydro- and di- Gee and I have presented quantitative Fvidence 55 that the very strong catalytic action of irradiation a t 3650 A on the autoxidation of liquid olefins originates in the photolysis of the hydroperoxide. This catalysis certainly extends to wavelengths as high as 4000 A and possibly higher.56 The quantum efficiency of the photolysis at 3650 is of the order of 0.1. However simple carbonyl compounds are even less absorbing in this spectral region (for acetone E~~~~ N o - o o ~ ) while photolysis is of com- parable efficiency to the peroxide (Norrish Crane and Saltmarsh 5 7 report + = 0-2 at 3150 A).53 Norrish and Reagh Proc. Roy. SOC. A 1940 176 429. 54 Porter Proc. Roy. SOC. A 1950 200 284. 55 Bateman and Gee Proc. Boy. SOC. A 1948 195 376. 56 Bateman Trans. Faraday Soc. 1946 42 266. 57 Norrish Crane and Saltmarsh J . CAem. SUG 1934 1456. 327 GENERAL DISCUSSION Prof. R. G. W. Norrish (Cambridge) (communicated) I cannot sub- scribe to the cogency of Dr. Bateman’s criticism with acetaldehyde for example at 3400 the extinction coefficient is 1.0 whereas for tert.-butyl hydroperoxide and also for hydrogen peroxide the extinction coefficient is 0.1. In addition formaldehyde has an absorption coefficient of 0.07 a t 3560 while the peroxides have fallen to negligible values.I would point out that the peroxide with which he was concerned namely cyclo- hexene hydr~peroxide,~~ is one that would be expected to have consider- able absorption owing to the unsaturated nature of the compound. The combined effect of the double bond and the peroxide group will un- doubtedly cause light to be absorbed more strongly at longer wavelengths than in the compound where the peroxide group is the only chromophore. Add to this the fact that the intermediate concentration of aldehydes in the oxidation of methane and ethylene and in higher hydrocarbons at the higher oxidation temperatures is of the order of millimetres while the concentration of peroxides is undetectable analytically and I think i t must be agreed that the photochemical effect is to be ascribed to the aldehyde.In addition Mr. Booth in this laboratory has recently been studying the photolysis of tert.-butyl peroxide using a solution of 2 yo by volume in n-hexane. He finds no photolysis whatsoever when all wave- lengths below 3350 are filtered out. but considerable decomposition when only those wavelengths above 3000 are removed. Dr. C. F. Cullis (Imperial College) (communicated) Prof. Norrish has suggested that the large structural effect encountered in the oxida- tion of the normal paraffins is simply attributable to the greater number of points of attack in the longer molecules Thus he argues that n-octane is more readily oxidized than n-pentane since there are more CH groups in the former molecule and the removal of hydrogen atoms from such groups is considerably easier than from terminal methyl groups.On this view the increase in oxidation rate with chain length should be roughly proportional to the number of CH2 groups whereas in fact the variation is much more marked.59 In any case an extension of this argu- ment fails to explain qualitatively the facts relating to the oxidation of branched chain paraffins. It is well known in hydrocarbon chemistry that the order of reactivity of C-H bonds towards radical attack decreases in the order 3” > zo > IO and Rice has shown that at 3ooOC the chances of attack at primary secondary and tertiary carbon atoms are in the ratio I 3 33.60 For the isomeric hexanes for example i t is possible to calculate the relative probabilities of hydrogen atom removal from the molecules concerned by multiplying the numbers of such atoms attached to primary secondary and tertiary carbon atoms by Rice’s factors and summing over the whole molecule.If the ease of removal of a hydrogen atom from a hydrocarbon molecule is the main factor controlling the ease of oxidation as Prof. Norrish’s argument implies the figures in the final column but one of the following Table should be measures of the oxidizability of the compounds concerned. On this basis therefore the order of ease of combustion of the isomeric hexanes would be V > I1 = I11 > I > IV whereas the experimentally determined relative oxidation rates (given in the final column of the Table) show that the order is I > I1 > I11 > IV > V.59 Prof.R. G. W. Norrish (Cawbridge) (communicated) My remarks about the relative ease of oxidation if hydrocarbons are intended to indicate that other things being equal the ease of oxidation per carbon atom will increase from methane and ethane which will have values 58 Bateman and Gee PYOC. Roy. SOC. A 1948 195 376. Cullis and Hinshelwood Faraday Soc. Discussions 1947 2 I 17 ; Cullis and Mulcahy Revue l’lnst. Franpais Pe‘trole 1949 4 283. 60 Rice and Rice The Aliphatic Free Radicals (Johns Hopkins I935) p. 100. 328 characteristic of the CH,.group to a value characteristic of the CH group for an infinitely long chain.Such comparisons as are made by Dr. Cullis in his criticism of my statement appear to me to have little relevance because the bases of comparison are so doubtful. Firstly his comparison of the rates of oxidation of hydrocarbons is based on the rates of pressure changes the reactants being at comparable pressures in all cases. This does not take account of the great increase in " carbon concentration " as we pass to higher hydrocarbons which would have the effect of weighting his results strongly in favour of the higher hydrocarbons. Secondly the comparisons could not all be made at the same temperature and the accumulation of intermediate products greatly' depends on the tem- perature as well as the nature of the hydrocarbon. For example in Compound I CH .CH . CH . CH . CH . CH (I) CH CH,. C H . CH . CH . CH CH CH,. CH,. CH . CH,. CH CH CH, C . CH,. CH I CH CH,CH I I I I CH,. CH . C H . CH one experiment of Dr. Patnaik the oxidation of butane at 270' C using a butane oxygen ratio of 1-5 and a total pressure of butane of 76.1 mm. intermediate products such as aldehydes and alcohols accounted for over 25 mm. in a total pressure change of 60 mm. It is not legitimate to com- pare the rates of oxidation of hydrocarbons by pressure changes in view of this great variability of end-products and the only certain way is to measure the rate of disappearance of hydrocarbon by analytical technique. This has not yet been done. Finally Rice's figures refer to the relative reactivity of CH, CH and CH respectively with alkyl groups whereas in comparing the rates of oxidation one is concerned with their reactivity towards oxygen or OH.There is no reason why there should be any close quantitative parallelism in the two cases which will undoubtedly be differently conditioned by temperature and steric hindrance. Such calculations as have been made on this basis can have little quantitative value until much more data have been accumulated and until some more uniform basis of comparison of reaction rates has been devised. Dr. D. W. G. Style (London) said When illuminated with Schumann ultra-violet light diethyl peroxide and ethyl nitrate emits as a fluorescence the same band system which extends from approximately 3300 %L to 5000 A.The emiss:ons from methyl chloroformate and methyl nitrite are also identical but different from that from the first two substances. 78 GENERAL DISCUSSION TABLE I1 Relative Probabilities of H Atom Removal Total telative telative Proba- Oxida- ilitr of tion Rate From IO From 2' I Atom From 3" C Atoms C Atoms Z A toms :emoval 1580 4 x 1 30 8 x 3 54 560 (11) 9 x 1 4 x 3 60 54 9 x 1 4 x 3 (111) I 2 18 I 2 x I (IV) 2 x 3 - I I 2 x I (V) GENERAL DISCUSSION 3 29 I t is probable that the emitters of the two band systems are respectively C,H,O. and CH,O-. Have these spectra been observed with other sources particularly cool flames ? Dr. Peter Gray (Cambridge) (communicated) It is very interesting to see the spectra of molecules containing the alkoxyl group shown by Dr.Style and to learn that the emitter is likely to be the alkoxy radical itself. It is hoped soon to examine a system in which the methoxy (CH,O-) group thermally produced may be an emitter viz. the lumin- escent thermal decomposition of methyl nitrate. The initial decom- position of methyl nitrate gives nitrogen dioxide and methoxy groups and is accompanied by a blue luminescence. The same behaviour is displayed by methyl nitrite. The light may be due either to CH,O- or to excited formaldehyde CH,O* and now both spectra are known i t may be possible to identify the source of this luminescence. Sir Alfred Egerton (Imperial College) (communicated) The emission bands which Dr. Styles has found and ascribed to the C,H,O radical in the spectrum of glow in diethylperoxide are interesting.When Dr. Harris was working on the decomposition of diethylperoxide in my laboratory' we looked for bands in the glow but a t that time had not the means for photographs of long exposure. It is possible these bands have some relation to the pronounced absorption bands obtained in the com- bustion of hydrocarbons which Pidgeon and I reported in 1933. Dr. C. A. McDowell (Liverpool Univevsity) (communicated) With re- gard to the oxidation of acetaldehyde Dr. Thomas Mr. Farmer and I have shown fairly conclusively by inhibition studies G 1 ~ 6 2 that the thermal reaction u p to about 140'C is a chain reaction in which two radicals the acetyl and peracetyl radicals play the predominant part.Dr. G. J. Minkoff (Imperial College) said The heats of combustion presented in our paper clearly have some peculiarities since the values of (-0-0-) which can be derived are of the order of 100-140 kcal. We intend to repeat thf measurements a t some future time but have mentioned the results already obtained as the work was done most care- fully and as the values are self-consistent within themselves (cf. Qc for members of series and for isomers). We have also recorded the refractive indices and melting points which we obtained so as to have as many published data on this subject as possible. These should be of use in determining the purities of future samples though we would stress the great advantages of determining the infra-red spectra of the peroxides during the preparations ; as the standard methods of purification (e.g.fractional distillation) are often inapplicable i t is a great help to know what particular impurities have to be removed. A brief point is that the latent heats of evaporation show high Trouton coefficients for all the hydroperoxides; this agrees with infra-red studies of the OH fre- quencies in suggesting that the compounds are strongly associated. DeteIminations of the heat of formation and of D(-0-0-) should thus if possible make allowance for the heat of dimerization etc. involved. Finally I should like to explain the use of the terms " acetyl propionyl and butyryl hydroperoxides " to describe the corresponding peracids. Mr. Everett and I have measured the (classical) dissociation constants in aqueous solutions of these compounds and of H,O, Me0,H and EtO,H and have found that the peracids are intermediate in acid strength between the corresponding hydroperoxides and carboxylic acids.In fact the increase over the hydroperoxide can be shown to be approx- imately accounted for by inductive and polarizability effects. We there- fore suggest naming the peracids " acyl hydroperoxides " to avoid the implication of a great acid strength sometimes inferred from the term McDowell and Thomas J . Chem. SOC. 1949,2208 2217 ; 1950 1462 ; Trans. Faraday SOC. 1950 46 1030. G 2 McDowell and Farmer unpublished work. GENERAL DISCUSSION 3 30 peracid (cf. perchloric and permanganic acids). Further if the diacyl peroxides are referred to as such there will be less chance of confusion with the corresponding acyl hydroperoxides (cf.dibenzyl peroxide benzoyl peroxide and perbenzoic acid). Dr. W. E. Vaughan (Emeryville California) (communicated) In the course of our own studies on organic peroxides we have obtained certain physical data on tert.-butyl hydroperoxide and di-tert.-butyl peroxide pounds were better than gg yo pure we do not wish to claim superior which differ from those just presented. Although we believe our com- accuracy for our data but wish rather to present them for the record in the hope that the differences may be resolved in the future. tert.-Butyl hydroperoxide 1.4010 5'5 I 15-0 log10 P (mm.1 I 1-5 kcal./mole 65 4.2 kcal./mole +alp . K p . "C . B.p. "C . Vapour pressure tained Di-lert.-butyl Penoxide 1.3890 - 40.0 I 11'0 9.6 kcal. /mole 12 75.0 kcal. /mole 0.0005 . Heat of vaporization . . Heat of combustion . Dr. M. Magat (Paris) said Leadbeater 63 has recently prepared in our laboratory some carefully purified peroxides and measured their refractive indices and Raman spectra. The following results were ob- Diethyl peroxide nb5 = 1.3720 f 0-0005 Di-a-hydroxydiethyl peroxide nh6 = 1.4265 Ethyl-a-hydroxyhydroperoxide n:4 = 1.4150 f 0-0010. For the diethyl peroxide the agreement with the refractive index given by Egerton et al. is excellent. The characteristic Raman frequency for the 0-0 bond found in all peroxides investigated by Leadbeater as well as in dibenzylperoxide and in H,O is located at 880 -& 3 cm.-l.I would like also to call attention to the danger of working with very pure crystallized peroxides even with those classified as relatively stable a spontaneous or induced rupture of the crystal has led to an extremely serious explosion when monohydroxydiethyl peroxide was purified by' crystallization. Dr. N. S . Wooding (Courtaulds Ltd. Coventry) (communicated) Cobalt acetate in concentrations of IO- to 10-3 M has been shown to be an effective catalyst for the oxidation of trimethylethylene in solution by gaseous oxygen. Recently i t has been found that the autoxidation of cellulose under alkaline conditions was also catalyzed by cobalt acetate.The effect of other metal ions was investigated and some were found to catalyze while others retarded the autoxidation. Such behaviour has been reported elsewhere.65 However manganese salts were found to behave both as catalysts and as retarders depending upon the concentra- tion of salt used. A possible explanation of this effect has been suggested in terms of the mechanism of metal ion catalysis postulated in this paper and else~here.~42 O 5 I would like to ask Prof. Eawn if he has found any evidence for negative catalysis by metallic cations in the autoxidation 63 Leadbeater Cow@. rend. 1950 230 829. 64 Entwistle Cole and ?Tooding Textile Res. J. 1948 19 527 609. 65 George and Robertson Tram. Faraday SOC. 1946 42 217. GENERAL DISCUSSION 33 1 of trimethylethylene or any of the other systems he has investigated since such phenomena if the suggested explanation is correct should be observable in other systems under the appropriate conditions.Prof. C. E. H. Bawn (Liverpool) (communicated) Under the experi- mental conditions so far studied the condition in which the rate of oxida- tion is independent of the catalyst concentration as observed by George and Robertson with saturated hydrocarbons has not been attained. There is no evidence at present that the catalyst terminates chains by a process such as Solvent 258 (-4 I <IO % decrease'in I conc. in 3 hr. 66 RO + Co++ -+ R0,- + Co+++. Dr. C. F. H. Tipper (Edinburgh) said It is nearly always assumed when considering the oxidation of hydrocarbons in solution that the hydroperoxides initially formed decompose to give free radicals.However under certain conditions for example in solvents of high dielectric constant or especially in the presence of acids as Dr. Waters has stressed heterolytic fission of the 0-0 bond may occur to give ions. The evidence for this has so far been mainly' organic,66 and so I would like to report the results of some kinetic measurements on the decomposition of decalin hydroperoxide in various solvents. The overall decomposition was found to be first order in all cases at any rate at low concentrations ( < I O - ~ mole/l.). The rates at 130' C and overall activation energies of decomposition in different solvents are shown below.TABLE I11 First Order Constant Time of Half Cnange Chlorobenzene . . o-Dichlorobenzene . Ethylene glycol . . @in.-1) Very slow 0*0105 0.002 68 0.0332 0.359 30,250 29,800 16,700 22,140 21 2 I Acetic Acetic acid acidlwater i 40% H,O by volume J . The addition of water to the ethylene glycol or the acetic acid thus presumably greatly increasing the dielectric constant of the solvent increased the rate of decomposition considerably. In the case of the acetic acid addition of 2 yo by volume of water had little effect but further addition up to about 30 yo by volume caused a large increase in the rate. Above 30 yo increase in the water content had no effect. Also in chlorobenzene solution a t I 15.5' C no decomposition was detectable over a period of hours but if 1-5 mole yo of acetic acid was added the time of half decomposition fell to 64 min.These results would seem to show that in the first three solvents de- composition of the hydroperoxide molecules into radicals is taking place possibly followed by a chain decomposition but that with a change in conditions an ionic decomposition can occur very readily'. Water and organic acids are very often stable products of oxidation of hydrocarbons in solution and thus as the reaction proceeds it is possible that ionic decomposition of the hydroperoxide formed might become important. Prof. Bawn (Liverpool) said IXr. S. F. Mellish and I have observed that the stable free radical aa-diphenyl p-picryl hy'drazyl reacts rapidly with the radicals of the type RO.and this provides a simple and convenient method for studying the rates of dissociation of peroxides into radicals 66 For example Robertson and V'aters J . Chem. Soc. 1948 1577. Bartlett and Cotman J . Amer. Chem. SOC. 1950 72 3095. Kharasch Fono and Nudenberg J . Org. Chem. Igjo 15 748; 1951 16 113 128. GENERAL DISCUSSION 3 32 (ROOR -f RO. + RO.). The vividly coloured radical which gives stable solutions in a wide range of organic solvents undergoes a sharp colour change on reaction with radicals which may be measured in a simple colorimeter. This method has been used to measure the rate of dissocj- ation of polymerization in;t'ators such as peroxide and azonitriles. Dr. W. A. Waters (Oxford) said A similar instance to the decomposi- tion of decalin hydroperoxide quoted by Dr.Tipper is that of tetralin hydroperoxide which was reported some time ago by Robertson and This afforded a good example of a reaction which appeared from kinetic study to be much more simple than i t really was. The results reported by Dr. Henderson show the extent to which very minor products or impurities can influence the course of autoxida- t:on. In the autoxidation of a related compound dibenzyl ether (Ph . CH,),O which has recently' been studied by Mr. Wickham-Jones and myself we have been able to show that irregularities in the oxidation are due to the formation of a trace of a phenolic by-product. Dibenzyl ether gives a stable peroxide and in the main the uncatalyzed rate of oxygen absorption is independent of the concentration of the peroxide formed.Homolytic dissociation of the peroxide to give more chain- starting free radicals does not therefore play a major role in determining the oxygen uptake rate. However the autoxidation is self-retarding and we have been able to show that as the peroxide of dibenzyl ether decomposes there is gradually formed just enough phenolic material to give a positive indophenol reaction. We ascribe the gradual retardation of the autoxidation to the formation of this phenol and have noted that when our reaction vessel was packed with chopped glass wool there was less peroxide decomposition less formation of phenol and less retardation of the autoxidation. Here the decisive factor seems to be surface catalysis of the mode of the secondary reaction-the peroxide decomposition- yet i t significantly influences the whole autoxidation process.Sir Alfred Egerton (Imperial College) (communicated) With refer- ence to Dr. Henderson's paper i t is well known that reproducible results are not obtained in hydrocarbon oxidation until the surface of the vessel has been conditioned by previous experiments ; the wall catalyst obtained in the liquid-phase oxidaton of ethylbenzene seems to be in line with this effect. Prof. J. P. Wibaut (Amsterdam) said In collaboration with Dr. A. Strang 6 8 we have carried out an investigation into the oxidation in the liquid phase by molecular oxygen of a number of normal alkanes with 8 to zz carbon atoms and of some branched octanes.When cobalt stearate is used as a catalyst the oxidation proceeds at a measurable rate. It has been found that the first stage of the oxidation reaction consists in the formation of a hydroperoxide. z 5-Dimethylhexane is slowly oxidized by molecular oxygen a t zoo C and a crystalline dihydro- peroxide melting a t 106.5~ C is formed CH3 CH CH3 CH I I + 2 0 2 4 H,C-C-CH2-CH2-C-CH3 I H I H3C-C-CHz-CH2-C-CH3 H I 0 0 I 1 0 0 I H Ir The presence of small quantities of peroxides can also be detected in samples of other saturated hydrocarbons which have been kept for a long time for instance in 3 4-dimethylhexane z 5-dimethylhexane 3-methylheptane n-nonane n-hexadecane methylcyclohexane.The primarily formed peroxide is decomposed under the influence of the cobalt 67 J . Chem. Soc. 1948 1578. 68 W'ibaut and Strang Proc. Kon. Neder. Akad. Wet. B 1951 54 ( z ) 101. 333 GENERAL DISCUSSION ions ; the radicals thus formed start a chain mechanism so that the oxida- tion reaction proceeds ROOH + Co++ -+ RO' + OH- + Co+++ ROOH ~2 ROO- + H+ Co+++ + ROO- -+ Co++ + ROO. We ascertained what products are formed by the oxidation of 2 5- dimethylhexane ; with reference to these reaction products a reaction scheme can be drawn up. The characteristic feature of the reaction scheme for the catalytic oxidation in the liquid phase of z 5-dimethyl- hexane and of other branched hydrocarbons is that the chain mechanism is initiated by an alkoxy radical.We ascertained the maximum rate of oxidation of normal alkanes with 8 g 10 12 14 16 18 20 and 22 carbon atoms. The experiments were carried out at 110.4" C and with 0.112 mmoles cobalt stearate per 61.7 mmoles hydrocarbon. There is a linear relation between the number of carbon atoms and the maximum rate of oxidation from C,,H, to C Z z H 4 6 . This can be explained by assuming that all the secondary carbon atoms have an equal chance of reaction. Some branched alkanes oxidized much more easily than the isomers with normal chains ; we ascertained the maximum rate of oxidation at 78.1" C measured in ml. 0 per 61.4 mmole hydrocarbon per hour (catalyst 70 mg. cobalt stearate) to be . 11.0 - 35'0 3-methylheptane z 5-dimethylhexane z 2 4-trimethy'lpentane .0-0 a-methylheptane 3 4-dimethy'lhexane . 3-methyl-3-ethylpentane . 0-0 . . 3'0 The oxidation begins at a tertiary C-H bond from which a hydro- peroxide group is formed. A considerable quantity' of acetone is formed during the oxidation of 2-methylheptane methylethylketone being formed during the oxidation of 3-methylheptane. The fact that z z 4-trimethylpentane (iso-octane) is not oxidizable under the conditions chosen can be explained by steric hindrance. The quaternary carbon atom which has three methyl groups screens the tertiary carbon atom to such a degree that an oxygen atom cannot approach the tertiary hydrogen atom to within the distance pertaining in the transition state. Dr. M. F. R.Mulcahy (Melbourne) (counrtzunicated) Several con- tributors 89-7 have discussed the liquid-phase oxidation of hydrocarbons catalyzed by decomposing peroxide or metallic catalysts in terms of the propagation mechanism ROZ- + RH + ROOH + R-. It may perhaps be of interest to recall that there is evidence that . . 5.0 Roy. SOC. A 1946 183 337. R- + 0 + ROZ- in the absence of catalysts (and of light) formation of hydroperoxide may occur by' a different mechanism. This was shown by the work of George and Robertson 73 on the " thermal " oxidation of tetralin. A similar result has recently been found by Mr. Watt in this laboratory with the uncatalyzed oxidation of benzaldehyde,' the kinetic behaviour being analogous to that of the tetra'in reaction.In the presence of benzoy'l peroxide however a reaction of the type shown above is initiated and is (additionally) superimposed on the uncatalyzed reaction. 69 Bateman Gee Morris and Watson this Discussion. 7O Bau-n Pennington and Tipper this Discussion. 71 Brook and Matthews this Discussion. 72 Mulcahy this Discussion. 73 George and Robertson Proc. Roy. SOC. A 1946 183 309 ; George Proc. 74 Mulcahy and Watt Nature (in press). GENERAL DISCUSSION 334 Dr. N. Uri (Munchester) said The liquid-phase reactions in which a wall effect is observed are few and therefore remarkable. There are ody four such cases known to me (a) the experiments relating to the oxidation of liquid ethyl benzene described by G. M. Henderson in this Discussion; (b) my own findings in the catalytic decomposition of hydrogen peroxide under certain experimental conditions ; 75 (G) the observations made by Dain and Kachan 7 G in the photochemical oxidation of water by ceric ions.It is not unlikely that the wall effect in this case leads to a recombination of OH radicals and a subsequent instantaneous decomposition of hydrogen peroxide by’ ceric ion. (d) In our work on the photo-initiated free radical polymerization of vinyl compounds in aqueous solution77 we made the observation that distance < rmm. from the wall the polymer is produced (by’ recombina- under conditions when practically all the free radicals are formed a t a tion of active endings) exclusively on the wall; none appears to be SUS- pended in solution.No doubt these experiments require some co-ordination and in this connection i t is interesting to note that all these reactions involve free radicals as intermediates and the process effected by the wall is con- sidered to be a termination process which may in some cases lead to a new type of chain reaction as reported by Henderson. In the oxidation of lubricating oils the effect of iron or copper as single catalysts was studied by Brook and Matthews. It is well known that in the catalytic decomposition of hydrogen peroxide the joint action of iron and copper as co-catalysts is much more than additive. It would therefore be interesting to study those effects in the oxidation of hydrocarbons particularly if peroxides are postulated as intermediate products.Dr. G . M. Henderson (Blackley) (communicated) The effects of walls on liquid-phase chain reactions are fairly’ well indicated in past work and to restrict it to two examples on oxidations excluding all references to solid catalysts the most relevant papers are those of Stephens 78 working with cyclohexene and Medvedev and Podyapol~kaya,~~ working with tetra- lin where many effects akin to the present work were noted and the effects of walls mentioned only with very different explanations. The novel feature of ethyl benzene to which we wished to draw attention as a convenient expet imental medium is that the two hydroperoxide forming mechanisms seem more clearly distinguishable and mutually exclusive. With very fresh and pure ethyl benzene we presume that the original wall termination process is without further visible effect but that in an older sample or after sufficient of the initial peroxide has further decom- posed some of the breakdown products possibly acids act as haptens (to borrow a term from immunology) and a new wall termination process arises which gives rise to visible effects on the rate of reaction.One can generalize that in other similar reactions that the two possible re- actions might then run concurrently or that the second stage might never even arise. Sir Alfred Egerton (Imperial College) (communicated) This paper refers particularly to the inhibition of the oxidation of lubricating oils I would like to draw the authors’ attention to a paper by Hanson and 75 Uri J .Physic. Chem. 1949 53 1070. 76 Dain and Kachan A C.S. Abstr. 1949 43 7349. 77 Unpublished observations. 78 Stephens J . Amer. Chew. Soc. 1936 58 219. 7g Medvedev and Podyapolskaya J . Physic. Chew. U.S.S.R. 1939 12 79. tiENERAL DISCUSSION 335 myself 8o on " Nitrogen oxides in internal combustion engine gases " in which the promoting action of the nitrogen oxides on the oxidation of lubricating oils was investigated and a subsequent paper entitled " In- fluence of catalysis on oil oxidation " 81 iron oxide was found to inhibit the nitrogen oxide catalyzed reaction. Dr. J. B. Matthews and Mr. J. H. T. Brook (Thornton) (communi- cated) In connection with Dr. Uri's suggestion that the study of mixed iron and copper catalysts would be of interest in the oxidation of hydro- carbons] results have been given in the paper showing that in the inhibited reaction the two catalysts are additive in their effect on the initial rate of reaction.In the uninhibited reaction] however the measured values of t$ are greater than the added effects of the individual catalysts thus lending support to the inclusion of peroxides as intermediate products. Values of t$ for iron and mixed iron-copper catalysts are given in Table 11 I 0 c s WEiCHT FRACTION OF HYDROCARBON FIG. I. of the paper and values of 4 for copper can be calculated from the data in Table I together with values of the intercepts with the oxygen absorbed axis of the asymptotes to the curves in Fig. 3 using the equation = ( A + B)t - (B/4) to describe the asymptote.Mr. J. H. T. Brook ( T ~ C Y M ~ O O ~ Z ) said I would like to add some further information on the kinetics of the iron-catalyzed inhibited oxidation. Using 50 p.p.m. of iron added as ferric stearate at 1 5 0 O C and using diphenyl as the inert solvent the dependency of the reaction upon the oil concentration was found to be of the form (as Fig. I above) [RHI cc 220 + 330 [RH]' No immediate explanation of the difference in kinetics between the iron and copper-catalyzed reactions is apparent. ' 0 PYOC. ROY. SOC. A 1937 153 90. 81 Symposium on Engine Wear Inst. Mech. Eng. 1937. 308 GENERAL DISCUSSION GENERAL DISCUSSION Prof. W. A. Noyes Jr. (Rochester) said Recent experiments by Mr. Martin in our laboratory indicate that methyl radicals from the photo-lysis of mercury dimethyl upon reaction with oxygen at low pressures give CO and CO,.Not enough results were available when I left to allow calculation of rate constants although the rates are in approximate agreement with the acetone results. Some uncertainty exists concerning the nature of the primary process in mercury dimethyl so that many more experiments may be necessary (only two had been made before I left) before quantitative conclusions can be drawn. Other results have been obtained with acetone-oxygen mixtures by Messrs. Marcotte and Durbetaki at intensities about 10 times those pre-viously used both in the presence and absence of mercury vapour. It is concluded that mercury vapour does not affect the results by more than the experimental error.At these high intensities in the absence of oxygen at 120Othe acetyl radical seems to be sufficiently stable to permit appreci-able reaction of the type and possibly CH + COCH = CH,COCH . - (1) 2COCH = (COCH,) . - (2) Thus a t this temperature the yield of CO + CO may be as much as six times the yield of CO in the absence of oxygen. At zooo however, the rate of decomposition of COCH is sufficiently rapid so that the results at these higher intensities agree with those previously found. These experiments are being performed with a view to permitting an analysis of products other than CO CO, CH and C,H,. In general the results are similar to those obtained at lower temperatures but i t is evident that at this temperature one has reached a point that chain propagation is setting in.This fact may limit the temperature range over which i t will be possible to deter-mine the activation energy for HCO decomposition. The fact that the CO is less sensitive to oxygen pressure and is higher than would be ex-pected from the low temperature results may be due to hydrogen abstrac-tion by radicals other than methyl. Dr. Finkelstein has preliminary results on the decomposition of diethyl ketone in the presence of oxygen. At 25°C the yield of CO falls as the pressure of oxygen increases indicating that the COC,H radical reacts with oxygen to give something other than CO at this temperature. In this respect i t seems to resemble the acetyl radical. However the quantum yield of oxygen disappearance is still rising at the highest oxygen pressure used (0.18 mm.) and has reached a value of 6.Complete oxidation of I molecule of ketone to CO and H,O would require 7-5 molecules of oxygen. Since small amounts of C hydrocarbons are found either more than I molecule of ketone disappears per quantum (doubtful a t 25OC) or the oxidation is proceeding quite completely. Further conclusions are not warranted at this time. At 200' C the yield of CO + CO seems to reach a constant value of 3 to 3.2 with increase in oxygen pressure the CO being about 2.3 and the CO about 0.8. However the quantum yield of oxygen disappearance has reached 7 at 0.78 mm. oxygen pressure and shows no sign of reaching a constant value. C hydrocarbons are still formed to the extent of about 0.3 mole per quantum even under these conditions.It would be premature to give a mechanism based on these results. The total yield of CO + CO is about the same as for acetone thus in-dicating that either a CH radical or a C,H radical yields either CO or CO when i t reacts with oxygen. However since the CO yield falls with increase in oxygen pressure for acetone and rises under similar circum-Mr. Marcotte has also made experiments at 225OC. Thus CO + CO may have yields of 3-6 or more GENERAL DISCUSSION 3 09 stances for diethyl ketone the explanation for the two phenomena must have important differences. Dr. R. Spence (Harwell) said I would like to recall some experiments by Dr. Bates and myself on the oxidation of methyl radicals produced by the photodissociation of gaseous methyl iodide a t room temperatures.The products formaldehyde methylal water and iodine were identified and determined quantitatively and the reaction was shown to follow the equation qCH,I + 2 0 % -f (HCHO) + (CH,O),CH + H,O + P I . h9 Two primary mechanisms suggested themselves namely, CH,I-+ CH +- I CH + 0 -f CH,O + OH * * (1) and CH + 0 -+ CH,O,. - ( 2 ) The absence of any inert gas effect which might have been expected with mechanism (2) seemed to favour the hydroxyl mechanism (I) but Blaedel Ogg and Leighton failed to detect the OH radical by its ab-sorption spectrum and therefore proposed that a mechanism of type (2) is operative. Prof. Noyes’s results at high temperatures suggest that the fonnyl radical is produced from the interaction of methyl radicals and oxygen but the above results indicate fairly clearly that i t is not formed at room temperatures.This discussion has drawn attention to the fact that after twenty years we are still unable to give precise answers to such fundamental questions as how long such entities as 0 and CH can stay together on collision and what type of stabilizing collisions are subsequently required. Dr. N. Uri (Munchester) said The HCO radical was postulated by Marcotte and Noyes as an intermediate product in the photolysis of acetone in the presence of 0,. It appears to me very important to learn more about this radical which may also play an important role in the process of photo-synthesis. I consider that formaldehyde may be formed from formic acid by an electron transfer i.e.e + HCOOH +- HCOOH. + HCO + OH* HCO + R . H 3 HCOH + R, where the bond R H has a relatively small bond dissociation energy. On the other hand i t is postulated that HCO decomposes into H + CO. In this connection i t would be interesting to have more data on the formation of formaldehyde in the experiments discussed by Prof. Noyes, so as to ascertain under what conditions HCO + RH + HCOH + R occurs in preference to HCO -f H + CO. Prof. A. R. Ubbelohde (Belfast) said It may be useful to reconsider to what extent the recombination of two radicals as in H+O,+HO . (A) or CH + CH -f C,H . - (B) requires a ternary collision to stabilize the product. Why do the energy-rich products not fly apart again almost at once as they would in simple atomic collisions such as H + 0 -f HO ? In the theory put forward in 1935 for the formation of peroxide radicals R + 0 -f R02* * (1) R02* + RH + ROOH + R .. (2) I attributed the temporary stabilization of the RO complex to “ quantum smudging ” notable for R > C,H, but more rigorous conditions should apply when R is small as in reaction (B) and especially (A). Bates and Spence J . Amer. Chem. SOC. 1932 53 1689. Blaedel Ogg and Leighton J . Amer. Chem. SOC. 1942 64 2499 3 10 GENERAL DISCUSSION Prof. S . W. Benson (California) said Prof. Ubbelohde has raised a question with regard to the allowability of two-body recombinations such as have been employed by the present authors. I believe that this question can be answered with considerable more certainty than has heretofore been the case.A recent theory of Dr. N. B. Slater has shown that to a first approximation the rate of unimolecular decomposition of an activated complex with energy E in excess of a critical energy E* needed for decomposition is given by k(E) = (1 - f>"-l, where Y lies in the range of normal frequencies (1012-1014 sec.-l) and n is the number of normal co-ordinates required for the description of the decomposition. The average molecule with energy in excess of E* has energy - E* + kT so that for any n v and E* we can calculate the mean life t ( c ) = I / K ( E ) . If we use mean values such as E*/kT - 40 ; v = 1013 sec.-l we find (with Slater) that t(E) w 10-l3 sec. if n = I ; t ( ~ ) E I O - ~ sec. if n = 5 and t ( ~ ) E 10-4 sec.if n = 10. These mean life-times (except for n = I) are very much larger for low temperatures. Thus for ethane e*/kT - 140 at room temperature and if we set n = 5 the mean life of an excited molecule will be I O - ~ sec. These results imply that for the reverse reactions the recombinations of free radicals complexes may be formed having mean lives sufficiently long to assure even at very low pressures quite efficient stabilization by collisional deactivation. This will be true for the high-temperature pyrolysis reactions and will be more pronounced for photochemical re-actions where temperatures are considerably lower. A reaction to form the radical HO is one that is close to the permissible limits. Here n < 3 ; E* > 45 kcal.and the average complex formed at 130O C has a mean life t ( ~ ) < v(kT/E*)2 w 3 x I O - ~ sec. which is of the order of collision times a t S.T.P. At temperatures of 500' C to 600° C or higher this mean life is much shorter and i t is probable that most of the HO complexes decompose before they can be stabilized. In the more complex methyl radical the theory gives a reasonable accounting of the high collision efficiency which has been observed for their recombination. It will not take many degrees of freedom in a com-plex with such a high critical energy to give i t a very long mean life. Finally we may say that recombinations will manifest their three-body character when the complex has a low critical energy the temperature is high and the combining radicals are not complex.Dr. C. A. McDowell (Liverpool University) (communicated) I should like to mention that Dr. Thomas and I have found that in the presence of a large excess of oxygen the acetyl radical is quite stable up to about zoo0 c. Dr. A. D. Walsh (Leeds) said Were any analyses for formaldehyde carried out by Noyes and Marcotte ? A t the higher temperatures normally used for the study of methane oxidation it was difficult not to suppose the reaction of CH radicals and 0 to produce HCHO took place. Dr. Peter Gray (Cambridge) (communicated) I am very interested to read of these experiments on the peroxides and in particular of the role of the alkoxy radical RO. Peroxides are encountered as an important feature of some of the chemiluminescent (cool-flame) processes in hydro-carbon oxidation and the alkoxy radicals are presumably also present.Such RO radicals are also important in the thermal decompositions of the alkyl nitrates and nitrites. The initial step is the fission of the molecule to give the alkoxy radical and the appropriate oxide of nitrogen. Both these processes may be accompanied by a luminescence.41 Methyl McDonell and Thomas J . Chem. Soc. 1949 2208 2217. Gray and Y o f f e Proc. Roy. SOC. A 1949 200 114, fi Gray xg51 unpublished work GENERAL DISCUSSION nitrite gives nitric oxide which may not be expected to oxidize the CH ,O- . Thus this chemiluminescence observed with nitrates nitrites and in the oxidation by nitrogen dioxide of the simple alcohols may be a property of the alkoxy radical itself.Furthermore from experiments on the nitrates values may be ob-tained under favourable circumstances for the heats of formation of the alkoxy radicals. The heats of formation of the other species involved and the energy required to break the 0-N bond are needed the latter may be identified with the energy of activation derived from measure-ments of initial rates. Such values may be employed both as a check on other measurements and as a means of calculating provisional values of heats of formation of peroxides of which the activation energies of decomposition but not the heats of combustion are known. The heats of formation of the methoxy and ethoxy radicals may be obtained from the following data on nitrate and nitrite esters. METHOXY COMPOUNDS (i) CH,ONO,(g) = CH,O + NO ; Eact = 39-5 kcal.Qj = 29.4 - 8.1 Qf(CH30) = - 2 ,, (ii) CH,ONO(g) = CH,O + NO ; Eact = 36'4 8 Qf = 16-7 - 20.9 Qj(CH,O) = + 1.2 ,, CZH~ONO,(~) = C&',50 + NO ; Eact Qj = 33.4 - 8.1 Qf(C,H,O) = + z'o,, The formation of each radical from its elements is numerically small and errors in any of the primary quantities appear magnified. The values (- 0.4 f 1.2) kcal./mole and (2 f 2) kcal./mole might be assigned to the CH,O and C,H,O radicals respectively. Sir Alfred Egerton (ImperiaE College) said I can allay Dr. Stevenson's anxiety about the measurements of heat of formation of peroxide which are mentioned in the preliminary draft of the paper by myself Emte and Minkoff. The values found though the subject of careful measure-ments are unconvincing; for instance the heats of formation of the peroxides are in most cases greater than those of the corresponding more stable substances.It is not easy to determine accurately the heats of combustion of these explosive peroxides. We intend to redetermine them and to withdraw for the present the last section of our paper referring to the heats of formation of peroxides. Dr. A. G. Gaydon (Imperial College) said I should like to point out that the values for the dissociation energies of H, 0 and OH taken from my book are for oo K not 298OK. Fortunately the errors about 0.8 kcal./mole nearly cancel out for the equations as used. Dr. A. D. Jenkins (Courtaulds Ltd Maidenhead) (communicated) : Vaughan ei al. report that vapour-phase pyrolysis of methyl-tert.-amyl peroxide (I) gives an 8 % yield of methyl ethyl ether which they ascribe to vapour phase combination of a methoxy radical with an ethyl radical.The scheme of decomposition of this peroxide will be substantially ETHOXY COMPOUNDS = 39-5 kcal. CH, / CH, / CH,-O-O-G-CH -+ CHSO + O-C-CH, Gray and Yoffe J. Chem. SOC. 1950 3180 312 GENERAL DISCUSSION Instead of the combination reaction suggested above i t is quite possible that the ether could be formed by reaction between the ethyl radical and the substrate by attack at the -0-O- bond : CH3 / \ CH3 / \ CZHS + CH300C-CH3 -+ CH30CzHs + -0-C-CH3 C2HS CZHS Whilst no data are available for estimating the activation energy of this process from experimental results one can make a comparison with E for the alternative hydrogen abstraction reaction, CH, / \ C,H + CH,OOC-CH -+ C,H + radical, C2HS in the following way.In the former reaction an -0-0- bond is broken to form an 0-C bond so that the exothermicity will be of the order (go-55)=35 kcal. and we can represent this process by the diagram : FIG. I. In the second case the reaction is nearly thermoneutral and the bond broken is a C-H bond. The crucial factors are the shapes and positions of the two repulsion curves but if they are not very different in the two cases the activation energy of the former process will almost certainly be lower than t h a t of the latter and therefore less than approx. 13 kcal. If the decomposition of this peroxide followed a chain type of reaction the chains would be very short if 8 % of the products resulted from the termination process so that the methyl ethyl ether may be at least partly formed by the reaction proposed here.Thus the corresponding diagram will be Fig 2. Thus the general reaction R + R’OOR” -+ R’OR + R”0 or R”0R + R’O, R’O -f R + CH,O or R”Q 3 R + CH,O. may occur in diaIkyl peroxide decompositions where R results fro GENERAL DISCUSSION 313 It may be noted that nitrates or nitrites might undergo a similar reaction even more readily since the 0-N bond has a dissociation energy of the order 38 kcal. and the reaction will be even more exothermic than the corresponding peroxide reaction. Dr. L. Bateman and Mrs. Hilda Hughes (Welwyn Garden City) (communicated) The decomposition of cyclohexene hydroperoxide in solution in the temperature range 60-100’ C exhibits rather different features to those reported by Bell ei! al.and also by other workers for other hydroperoxides.718 Our investigations in this field are still pro-ceeding but certain results appear to warrant consideration here. The non-formation of diphenyl and phenol indicates that the solvent is inert under our conditions. Cyclohexenone and R0,-double bond addition compounds appear to be the main organic products. Each mole of peroxide decomposing liberates nearly 0.5 mole of water and much less oxygen (- 0.1 mole). The Arrhenius activation energy is 26 kcal./mole. We conclude that the measured rate refers to the one-step bimolecular decomposition, In benzene solution the reaction is almost exactly second order.RO2-H + HO-OR + RO2- + H2O + RO-. - (1) Radical t C H5 + H - - - -Radical /C H C H~ + CN 00ccc>15 I \ ’C H i FIG. 2. In cyclohexene solution the decomposition proceeds faster (roughly five-fold) and the order is reduced to 1-7-1.8. These characteristics the isolation of cyclohexenol as the major product,O and a linear relationship between [R02H] /rate and I /[RO,H] are consistent with a chain reaction comprising (I) as initiation step and RO- + R’H + ROH + R’- - (4 R’- + ROJ3 + R’OH + RO- . * (3) zRO-1 . * (4) R’- + RO- 3 Inactive products (5) 2R’- i ’ . * (6) wheze in this case R and R’ are identical and the reasonable assumption 10 is made that kz = A&,. In the presence of stearic acid the reaction is greatly accelerated and becomeF of the first-order with respect to the peroxide.These facts correlate strikingly with the catalytic action of the hydro-peroxide on the oxidation of the parent olefin with molecular oxygen, Farkas and Passaglia J . Amer. Chem. SOC. 1950 72 3333. Kharasch Fono and Nudenberg. J . Org. Chem. 1951 16 113. Farmer and Sundralingam J . Chem. SOC. 1942 121. lo Cf. Bateman Gee Morris and Watson this Discussion 3 I4 GENERAL DISCUSSION Normally the relationship is obeyed ; in the presence of stearic acid this changes to roxidation = const. d([RO2HI2) . * (7) Reaction (3) has not been considered by Bell et al. and an indication of its existence in their work viz. the formation of methanol4 from CMe,O,D and Me radicals has been otherwise explained.However we believe the following experiment provides strong evidence for its occur-rence. Cyclohexene hydroperoxide was decomposed in the substituted I 4-diene ethyl linoleate (R’H). R’ is thus a mesomeric pentadienyl radical which if formed will react preferentially to give a conjugated I 3-diene derivative. In fact conjugated diene units were produced, they were located only in an ethyl hydroxy-linoleate and the formation of this substance was greatly in excess of that expected to result either directly or indirectly via radical-radical interaction. The alternative reaction R’- + ROZH -+ RO2- + R’H . . (8) has possibly been shown to occur by Bell et al. at 195’ in the gas phase reaction of Me and Et radicals with CMe,O,D. It should be noted how-ever that even in these systems the reaction is roughly 10 kcal.endo-thermic and this figure is increased to about 30 kcal. when R’H is an olefin and allylic resonance energy has to be supplied. At lower temper-atures and especially in unsaturated solvents (8) must be quite unim-portant compared with ( 3 ) which is about 30 kcal. exothermic (if Do-o = 40 kcal.). Reaction (3) is envisaged as a radical substitution reaction (SRz-to extend the Hughes-Ingold terminology) formally analogous to (I) cf. H j ( 7 . (-y H !n t-)! ui R’- + O-:-OR with RO,-I-H + 0-\--OR, where the arrows denote one-electron displacements. The corresponding ionic processes involving heterolytic bond scission can obviously be re-garded likewise. We do not suggest that the decomposition mechanism now advanced is necessarily applicable to other hydroperoxides under different con-ditions-it is quite evident that analogy can be especially misleading in this field.Thus the effect of increasing temperature may be particularly critical in facilitating a unimolecular S,I relative to an S,Z decomposi-t ‘on because the increasing tendency of the comparatively weak 0-0 bond to rupture will be roughly paralleled by a decrease in the molecular association in the hydroperoxide (evidenced by hydrogen bonding) which is undoubtedly precursory to reaction (I). Nevertheless we do suggest that the relevant conclusions of Bell et al. require careful reconsideration, in particular the postulation of unimolecular dissociation of tert.-butyl hydroperoxide in solution at temperatures not greatly above 100’ C.Not only are the rate data at 150’ C given in Table I1 of these authors’ paper consistent with our scheme (they correspond to a reaction order of 1-55), but Bolland and Morris 11 have made the very significant observation that this hydroperoxide in common with numerous primary secondary and tertiary allylic hydroperoxides catalyzes olefin oxidations in accordance with eqn. (7). Dr. W. E. Vaughan (EmeryvilZe California) (communicaded) The questions raised by Dr. Bateman and Mrs. Hughes concerning our paper seem to u s to be well answered by the experimental evidence presented, namely the isolation of methane-d and ethane-d and of ethyl tert.-butyl l1 See Quart. Rev. 1949 3 I GENERAL DISCUSSION 31 5 peroxide from the appropriate reactions.in substantial yields provide definite proof for the reactions These actual product isolations R' + (CH,),COOD -+ RD + (CH,),COO* . - (1) and R. + (CH,),COO' -f (CH,),COOR . ' (4 and they are not otherwise explained by the mechanism proposed by Dr. Bateman and Mrs. Hughes. It is our opinion that reaction (I) is not endothermic as claimed by Bateman and Hughes but rather is exothermic to the extent of about 20 kcal./mole when R is methyl. No quantitative evidence is available, but qualj tative comparison of ted-butyl hydroperoxide as a hydrogen atom donor with cyclohexene and formaldehyde leads us to the conclusion that Do-,(ROOH) is of the order of 80 kcal./mole rather than 105 or 118 kcal./mole as observed in certain alcohols and in water.Further proof for the chain decomposition we have proposed may be obtained from a comparison of the rates of decomposition a t 140~ of tert.-butyl hydroperoxide in a reactive solvent n-octane and in an inert medium chlorobenzene. In the latter solvent reaction (3) cannot intervene because decomposition leads to quantitative yields of tert.-butyl alcohol and oxygen. Moreover in this solvent the rate is some 40 times that in n-octane. Bateman and Hughes have acknowledged the difficulties in drawing analogies between decompositions of various peroxides under differing conditions. However the decomposition of cyclohexene hydroperoxide in an unsaturated solvent can also be interpreted by our mechanism which has the evidence of actual isolation of a stable peroxide intermediate.For ethyl linoleate this intermediate would correspond to (B) ; R'* + ROOH -f R'OH + RO' . * (3) R* + - CHdH-CH2-CHSH- -+ RH + -CH=CHL:CHLCHCCH-(A) (4) OOR (5) I ROO' + (A) -+ -CHdH-CHdH-CH-0. (B) (6) I (B) -+ KO' + -CH=CH-CHeCH-CH-OH (C) I (C) + ROOH (or RH) -+ ROO' (or R.) + -CH=CH-CH=CH-CH- (7) The allylic stabilization of (A) would increase the possibility for the postulated association reaction (2) at the expense of metathetical reactions such as (I) or ( 3 ) . The decomposition of cyclohexene hydroperoxide in cyclohexene can also be interpreted by a strictly analogous mechanism. However we think that in an inert solvent the reactions of the hydroperoxide group are greatly complicated by the presence of reactive methylene groups and the double bond in cyclohexene hydroperoxide.Although our in-formation offers nothing either pro or con concerning reaction (3) we see no objection to it on purely theoretical grounds and have in fact, indicated that the processes proposed in our paper do not exclude other possible competitive transformations which may be proceeding simul-taneously and undetected. Dr. W. A. Waters (Oxford) said Whilst we all realize the great value of the kinetic studies of olefin autoxidation carried out by members of the British Rubber Producers' Research Laboratory I would suggest that there will be little value in increasing the precision and detail o GENERAL DISCUSSION their present types of kinetic work beyond the stage which it has now reached.Kinetic studies of chain reactions give velocity equations which are dependent very much on one particular chain-ending process e.g. 2R' 3 products or R' + ROa' +- products (dimers often unspecified). When however organic chemists start looking for these diagnostic compounds in autoxidized mixtures they seldom find them and sometimes can show that the postulated dimers are far too reactive to persist in the final oxidation product. For example, in our current study of the effects of phenols on benzaldehyde autoxida-tion 12 Mr. C. Wickham-Jones and I deduced kinetically that the chain-termination product in mixtures containing +-cresol was a dimer of a mesomeric phenoxy radical (ArO*)2. We have failed as yet to isolate any of the known dimers of the tolyloxy radical from our reaction product and have found that the material which is eventually formed must be more complex.Moreover by adding the known dimers separately to our reacting system we have shown that they have a comparable reactivity to the p-cresol and so could not persist throughout the whole autoxidation. With 2 6-xylen-1-01 which gives similar kinetics the obtainable reaction product is not the phenol dimer but its further oxidation product. the corresponding diphenoquinone ; even this is an autoxidation retarder which therefore must be converted to still another product. One important point that kinetic investigators should remember is that the chain-terminating reaction which will be revealed by their reaction velocity measurements is the one which most rapidly removes active free radicals from the system.None of a whole series of slower radical-removing processes or of secondary reactions involving products which initially are present in very low concentration may be able to influence the reaction velocity to a discernible extent unless the experi-mental conditions are especially designed so as to detect one of them. " Initial velocity I' measurements in particular are of little help in the diagnosis of the chemical identities of products of chain termination. In reactions involving hydroperoxides it is important to remember that if the system becomes appreciably acidic or basic then an ionic decomposition of the hydroperoxide may set in to give products which are often quite different to that of its homolytic fission.The homolytic fission seems to proceed RlR2R3G-O-OH -+ RlR,R3C-O' + 'OH L RlR2C==0 + 'R3, where the fission rules for the breakdown of the ROO radical are in general those given by Walsh.lS The acid-base catalyzed reaction however, proceeds RlR2R3C-0-0-H + H+ -* where R is the group that most easily carries with it an electron pair. ways. These different products influence the subsequent reactions in different la J . Chem. SOC. 1951 812. l3 Trans. Faraday SOC. 1946 42 269 GENERAL DISCUSSION 317 Dr. A. J. Harding (Cambridge) said A good approximation to the kinetics of hydrocarbon oxidation may be derived without postulating intermediates of definite types such as peroxide or aldehyde. It is only necessary to assume that the chains consist of links which are alternately reacting with oxygen and with the hydrocarbon that the chains are terminated by destruction of the radicals which would otherwise react with hydrocarbon and that the chains are initiated by the breakdown of an intermediate which they themselves produce.The use of special features such as the ratio pmax/Apmax becomes essential if the finer points of the kinetics are to be evaluated. One of these points is the type of degenerate branching. Dr. Mulcahy has described the type in which the reaction fails to reach an infinite rate (i.e. ignition) because the re-actants are consumed before the intermediate concentration becomes sufficiently great. Another type of degenerate branching has been suggested for the oxidation of hydrocarbons at higher temperatures.In this second type the intermediate (e.g. aldehyde or peroxide) which generates chains is itself oxidized by the radicals i t produces. Since the production of intermediate is proportional to the first power of its concentration and the destruction is proportional to the second power a limiting concentration of intermediate is reached. This means that even if no reactants are consumed in the process the reaction attains a limiting finite velocity. A distinction between the two types of degenerate branching is readily made. In the first case addition of intermediate at the beginning of the reaction will give an increase of maximum rate as well as a decrease of induction period. In the second addition of intermediate will reduce the induction period but will produce no change in the reaction velocity until so much intermediate has been added that the induction period has been completely eliminated.The oxidation of ethylene at temperatures in the region of 400" C with formaldehyde as intermediate shows the behaviour expected of the second t41~e.l~ A derivation of the variation of pmax/Apmax with hydrocarbon con-centration for the second type of degenerate branching leads to a relation-ship approximately the same as that found by Dr. Mulcahy for the first type. This relationship does not therefore provide a clue to the type of degenerate branching although i t does give valuable support to the thesis that degenerate branching is taking place. l4 Dr. Peter Gray (Cambridge) (communicated) May the fact that the propylene oxidation studied by Mulcahy appears to occur through a peroxide which may react in such a way as not to produce active radicals be associated with the characteristics of propylene as a chain-breaking agent ? (In this case the constant C might include a term corresponding to the propylene concentration.) Dr.M. F. R. Mulcahy (Melbourne) (communicated) Dr. Gray's sug-gestion would require the chain-terminating effect of propylene to be specific towards the radicals produced immediately on decomposition of the peroxide. The participation of the hydrocarbon in a chain-terminating reaction in the main cycle would affect the second rather than the third term in eqn. (4) (tending to make B independent of [RH] under conditions of excess oxygen).However i t now appears that the significance of the value of C derived from the experimental results is somewhat complex. Some very recent experiments carried out in this laboratory by Mr. Ridge with propylene at 288" C have shown that the intercept of the pmax/Apmax against [RH] curve on the pmax/Apmax axis is dependent on the surface conditions and may become positive. It seems therefore that the simple interpretation of this intercept (C) given in my paper is in need of some amplifications. 1 4 Harding Thesis (Cambridge 1948). Norrish X V I Int. Coll. C.N.R.S. (Paris 1948) p. 16. Harding and Norrish (in preparation) 318 GENERAL DISCUSSION Dr. E. J. Harris (London) (communicated) In 1935 l5 i t was suggested that compounds known to promote the branched chain reaction in the slow oxidation of hydrocarbons such as ethyl nitrite and ethyl peroxide, did so by providing -OR radicals and the converse process a condensa-tion of -OR to peroxide seemed not unlikely.In 1938 Neumann and Tutakin l6 showed that small additions of diethyl peroxide to butane + oxygen mixtures would cause the appearance of a flame similar to the cool flame seen during the oxidation reaction. To put the hypothesis of the intervention of peroxide on a frrmer basis two approaches are possible namely the physical and chemical investiga-tion of the slow oxidation and the study of the peroxides themselves under appropriate conditions. In 1936 Sir Alfred Egerton suggested the latter method as a profitable one.Before mentioning a few relevant properties of peroxides i t will be useful to state the analytical evidence concerning peroxides isolated from oxidation reactions. At 320-270' propane and butane can be made to yleld considerable quantities of hydrogen peroxide making up 10 yo of the total condensate. The H,O combines with two molecules of form-aldehyde one of the other main products and in one experiment l7 0.5 g. of the compound was isolated. If however the reaction is carried out when the walls of the vessel are slightly contaminated with alkaline material the reaction no longer leads to the formation of the peroxide, though all the other products are nearly the same. This and other evidence suggested that the H,O,-formaldehyde compound was only formed after condensation.The fact that H,OB can be isolated indicates that radicals -OH or -0,H intervene in the slow oxidation the sensi-tivity to the state of the wall may mean that hydrogen peroxide is only formed when collisions involving the wall permit it. The analytical problem of showing the presence of an alkyl peroxide is of course com-plicated by the presence of a large amount of H202. Results of the study of the alkyl peroxides 18 19 2o and their influence on slow oxidation reactions are consistent with the following : (a) They may be formed in traces as by-products but chemical evidence is lacking. (b) The alkyl hydrogen peroxides like H,O, are very sensitive to the state of the surface being decomposed by traces of alkali. They will ignite giving a diffuse blue luminescence when introduced with oxygen into a vessel at 250-3ooo.The luminosity resembles that sometimes seen in the hydrocarbon oxidation and is general rather than flame-like. The dialkyl peroxides will decompose with a blue flame in certain conditions of temperature pressure and gas mixture. Below the critical limits they decompose homogeneously. The artificial blue flame does not inevitably promote hydrocarbon oxidation even when induced in hydrocarbon + oxygen mixtures. The result depends upon the surface and the temperature. ( d ) Traces of the peroxides reduce the induction period of the slow oxidation of hydrocarbons without altering the rate of the subse-quent reaction. This is explicable on the basis of radical formation. The small quantities of peroxide which are needed to produce visible effects make i t not surprising that they have not been isolated.Only 0.007 cm. of dipropyl peroxide will give a flash with air at 270° and 0-01 cm. 16 Egerton Smith and Ubbelohde PhiE. Trans. Roy. SOC. 1935 243 433. 16 Neumann and Tutakin Acta physicochim. 1938 9 861. l7 Harris and Egerton Chem. Rev. 1937 21 287. 1s Harris and Egerton Proc. Roy. SOC. A 1938 168 I. l9 Harris Proc. Roy. SOC. A 1939 173 126. 20 H a r k Proc. Roy. SOC. A 1940 175 254 GENERAL DISCUSSION 3 I 9 of ethyl hydrogen peroxide can be seen to luminesce when oxygen is present. In spite of the explosive properties of the peroxides the alkyl hydrogen peroxides like H,O,,21 can survive quite high temperatures if the vessel has a low surface/volume ratio and particularly if a carrier gas is present.Thus at 180' only 5 yo ethyl hydrogen peroxide was decomposed when passed with N through a tube at 180' (contact time 4 sec.). A t 320°, z yo survived. Under somewhat similar conditions ( I sec. contact time), 85 yo H,O sometimes survived. Post-war work on hydrocarbon oxidation (e.g. Hinshelwood 22 Mulcahy 23) appears to be in accord with the scheme outlined by Egerton and Harris 24 in 1938 in which i t was proposed that the initial step was peroxide forma-tion followed by splitting to give radicals -OR -OH and -O,H, which carry on the main reaction. Analogous schemes were outlined by Ubbel~hde,,~ Pease 28 and Lewis and v Elbe.,' Egerton and Harris, however did expressly suggest (cf. their Table 11) the reaction z -OH -f H,O as a chain-terminating one and this together with a reaction of -0,H with a hydrogen compound would explain hydrogen peroxide formation.Organic peroxides as such were only invoked in the initiation of the reaction and this agrees with the fact that surfaces active in decomposing the alkyl hydrogen peroxides (e.g. salt-coated ones, which under the experimental conditions become alkaline) are those which give rise to long induction periods for the slow oxidation reaction. Dr. G. J. Minkoff (ImperiaE College) said I should like to raise two points in connection with Dr. Mulcahy's paper. First with regard to the evidence of Badin which is quoted in support of the formation of hydroperoxides in oxidation Badin observed a line at 11.4 p in the infra-red spectrum of tert.-BuOOH; he concluded that since H,O also has a line at 11*4p this line is characteristic of the -0OH grouping.This is unsound for several reasons; in many observations of the infra-red spectrum of freshly prepared ten!.-BuOOH I have only once found a line at 11-4p and that was in an impure sample ; the position of the line was otherwise found to be at 11.6~. Di-tert.-butyl peroxide does have a line at 11*4p so that clearly the line cannot be characteristic of -0OH ; in fact as was pointed out before i t may be connected with both the -04- group and with the tevt.-butyl group. The line which Badin found may well have belonged to H,O,. The other point is connected with the attempt t o draw up a mathe-matical expression for the rate of pressure change.This is based on the proportionality observed between peroxide concentration and the rate of pressure change. In the derivation in terms of initial concentrations the tacit assumption is made that P is the hydroperoxide of the original hydrocarbon. However this assumption may not be correct because in the degradation of a long cha'n paraffin several steps must occur all prob-ably involving peroxides. In the calculations made by Bolland this ob-jection does not arise since only one peroxide is involved. Another complicating factor is that the different peroxides formed (i.e. primary, secondary peracids etc.) will react to different extents with the potassium iodide reagent; thus the maximum peroxide concentration may be masked by the lack of reaction with KI of some of the peroxide present.Dr. W. A. Waters (Oxford) (communicated) Sir Cyril Hinshelwood's review of the influence of substituents on the rate of hydrocarbon oxida-tion can be given an alternative interpretation which is equally consistent ar Harris Trans. Faraday Soc. 1.948 9 764. 2 Hinshelwood Faraday SOC. Dascussaons 1947. z8 Mulcahy Trans. Furuduy Soc. 1949. 45 575. ,4 Egerton and Harris Proc. I8me Congr. Chim. Ind. (Nancy 1938). 2s Ubbelohde and Egerton Proc. Boy. SOC. A 1935 152 354. 26Pease J . Amer. Chem. SOL 1929 51 1839 et seq. 27 Lewis and v. Elbe J . Amer. Chem. SOC. 1937 59 976 320 GENERAL DISCUSSION with the theories of general polarity. The substituents which he finds to increase oxidation rates are also those which promote the attack of methyl radicals on C-H groups.Kharasch and his colleagues for in-stance have reported that methyl radicals from the decomposition of diacetyl peroxide preferentially attack C-H bonds vicinal to C-C1, and also CH-CO- CH-0- CH-CO-OMe but do not attack CH, groups. From studies of tert.-butyl peroxide it now appears that R-0' radicals (e.g. Me,C-O.) have not quite the reactivity of alkyl radicals such as methyl and are even more selective in the same sense in their reactivities towards C-H groups.28 I would therefore suggest that the reaction which immediately follows Sir Cyril's chain initiation process R-0-0-X + R-0' + '0-X / / \ \ i.e. R-0' + H-C- -+ R-0-H + 'G-may be much slower and much more selective than the subsequent stages of the gas-phase oxidation involving R' and R-0-0.radicals. It may well be rate-determining for the whole breakdown of the oxidized compound. A similar state of affairs is well known in polymerization chemistry where the sequence, (i) Catalyst (eg. benzoyl peroxide) -+ Radicals R' ; (ii) R' + Monomer -+ Hydrocarbon type radical R-M' ; (iii) R-M + M 3 R-M-M' etc. (fast chain growth), results in a polymerization kinetically dependent upon the rates of both processes (i) and (ii) as for instance in the benzoyl peroxide catalyzed polymerizations of both styrene and vinyl acetate. In polymer chemistry again the stability of CH in comparison with A H z - is strongly marked as for instance in the ready chain transfer to -CHz-CH=CH- and the converse tendency of compounds with CH,-C=C groups to build up high polymers.In connection with later states of hydrocarbon oxidation i t may be noted that in the Dyson Perrins Laboratory we have recently shown that in the liquid phase the breakdown occurs quite readily a t temperatures as low as SO-IOO~ C. This is con-firmatory evidence for the view that in hydrocarbon oxidation at elevated temperatures only the initial stages of the reaction sequence are rate-determining. Dr. A. D. Walsh (Leeds) said It is helpful in considering the kinetics of oxidation reactions of hydrocarbons in the gas phase to write the step that produces the alkyl peroxide radical in the form Commonly the * is omitted but its inclusion serves to remind us that in the first instance the RO radical is energy-rich.This is important for at least the following reasons. (i) It makes it less likely that the reverse reaction to (I) will be neglected in considering the full kinetic scheme. (ii) The excess energy is presumably distributed over various vibrational degrees of freedom. The radical is therefore to be thought of as vigorously twisting turning and generally distorting itself far more than does a " normal " radical. This means that subsequent reactions which axe rather more complicated than are usually found in gas-phase oxidation processes are rendered more pIausible than they would otherwise be. Cf. also Farmer and Moore J . Ckem. Soc. 1951 131. Alkyl-CO' + AllrJtl' + CO R + 0 = RO,*. (1 GENERAL DISCUSSION 321 Those gas phase chain reaction steps which are best substantiated ( a ) those in which one linkage is broken e.g.( b ) those in which one linkage is formed e.g. (I) above the reverse commonly' fall into one of 4 classes 29 : CH,CH3 = CH + CH - (2) CH + NO = CH,NO* . * (3) of (2) and (G) those in which one linkage is formed and one is broken e.g., abstraction of an atom (particularly H) from a molecule by a radical as in ( d ) energy transfer reactions e.g. reactions involved in the quenching OH + RH = H,O + R; * (4) A * + B = A + B * . . - ( 5 ) of fluorescence : As illustrations of these four types it is instructive to look at the paper by Bell Raley Rust Seubold and Vaughan. The great merit of that paper is that some experimental evidence for each step postulated has been obtained.Consideration of these steps shows that they all con-form (or in one case can readily be slightly altered to conform) to the above types. In other words the steps occurring most commonly appear to be the simplest. This should not be surprising for in postulating a reaction step, i t is vital to consider just how the reactants could be converted into the supposed products; and such consideration makes it difficult not to conclude that more complicated steps than (a) t o (a) are likely to have unfavourable steric factors or activation energies.30 If a postulated gas-phase reaction step does not conform to one of the 4 types one cannot say it is impossible-but one can and should demand that the evidence for it be unusually strong unless it is to be dismissed as mere speculation.This is incidentally a serious criticism of many of the reaction steps postulated by certain authors. To take but one example the supposed reaction RCH,CHO + 0 = RCHO + HCHO involves the simultaneous breaking of two linkages and the forming of two linkages. Unless particularly strong evidence in support of this was forthcoming-and such has certainly not yet been produced-this postulated step can be taken as very unplausible. In the special case of a reaction such as (I) however i t is not unplaus-ible to suppose a subsequent reaction for RO,* more complicated than types (a) to (d). If this is re-written as it is split into simpler steps and the fact that the second involves the breaking of two linkages and the forming of one does not appear un-plausible.Another example might be a somewhat complicated isomer-ization of RO,*. (iii) Even if the reaction supposed subsequent to (I) conforms to (a) to (d) inserting the * in (I) helps one to remember that the activation energy' for the following reaction need not be as great as would be the case for " normal " RO,. Some or all of the activation energy may be provided by the energy produced by (I). An example is the reaction CH + 0 = HCHO + OH. CH3 + 0 = [C,HO,"] = HCHO + OH. The step RO + RH = ROOH + R - (6) 29 Cf. Chamberlain and Walsh Rev. Inst. Frangais du Pe'trole 1949 4 307. 3 0 Cf. Ubbelohde Rev. Inst. Frangais du Pe'trole 1949 4 315. 31 Norrish this Discussion. 322 GENERAL DISCUSSION although probably exothermic when RH is an olefin is probably endo-thermic when RH is a paraffin.32 The activation energy required for (6) in gas phase reactions however need not be as great as this endothermicity.The energy of formation of RO, carried over into a subsequent reaction reduces any improbability of that subsequent step. To realize this is to recognize that the process of gas-phase peroxide formation may’ be subject to inhibition by inert gases. The greater the dilution of the reactants the greater the chance that collisions (and reactions of type (a)) will take so much energy from RO,* that the probability of (6) is seriously reduced. Peroxide formation is involved in cool flame phenomena. To test whether inhibition by inert gases of cool-flame formation can occur however one needs to choose conditions carefully’.At temperatures near the lower temperature limit of spontaneous cool-flame formation the dominant chain-ending reaction is usually a surface destruction of radicals and the inhibition may be swamped by a greater promotion viz. the hindering of radical diffusion to the walls. Along the upper temperature limit of cool flame formation, however the dominant chain-ending reaction is a gas phase one (in most cases the decomposition of a bulky alkyl radical to an olefin and a small alkyl radical 33 3ii). At such temperatures for diethyl ether (Fig. 3 of ref. (33)) and for propane (Fig. 11 of ref. (35)) a small inhibition by’ inert gases is observed. It seems probable that this inhibition represents the expected effect of reduction of the energy content of RO,* though of course i t has to be remembered that cool-flame formation3(j and pro-pagationa7 involve a thermal factor and inert gases may also inhibit purely by virtue of their thermal properties.Finally i t is important to be clear about the nature of the excess energy carried by RO,*. For the analogous HO,* radical however the following reaction has been supposed to occur 38 The idea is closely similar to that of (ii). This is surely vibrational in nature. HO,* + NO = EIO + NO -/- 0, the HO,* radical transferring the whole of its energy to the NO,. Such a complete transfer of energy seems only likely for electronic energy. Yet the effects of inert gases (M) on the second limit of the €-I + 0 ex-plosion are usually supposed to require the transfer of energy from HO,* (i.e.H + 0,) to M ; and the different gases that are effective make i t difficult to suppose other than vibrational energy is concerned. Some confusion seems therefore to exist. Dr. C. A. McDowell (Liverpool) (communicated) I was interested to read Dr. Walsh’s remarks on the existence of RO,* radicals i.e. RO, radicals endowed with excess energy. Similar views were expressed by Dr. Thomas and myself in our paper on the inhibitory effect of nitrogen peroxide on the gas-phase oxidation of acetaldehyde. 39 In considering possible subsequent reactions of RO,* radicals one has to be extremely careful for the excess vibrational energy with which these radicals are endowed may or may not persist throughout the numerous collisions which these molecules may make before they undergo reaction.If the excess of vibrational energy does not persist then ob-viously little is to be gained by ma’ntaining the above notation. Numerous 32 Walsh J . Chem. SOC. 1948 339. 33 Chamberlain and Walsh 3rd S y m p . Combustion Flame and Explosion Phe-34 Walsh in course of publication. 35 Malherbe and Walsh Trans. Favaday SOC. 1950 46 835. 36 Malherbe and Walsh Trans. Faraday SOC. 1950 46 824. 37 Spence and Townend 3rd S y m p . Combustion Flame and Explosion Phe-38 Dainton and Norrish Proc. Roy. Soc. A 1941. 177 395. 3% McDowell and Thomas J . Chem. Soc. 1950 1462. nomena (Williams and Wilkins Baltimore 1949). nomena (Williams and Wilkins Baltimore 1949) GENERAL DISCUSSION 323 cases are however known where vibrational energy does persist through-out a time interval during which a molecule may make as many as 5 x 104 collisions before one quantum of vibrational energy is dissipated.Theoretically there is no reason why vibrational energy should not be transferred from one molecule to another provided their vibrational levels are sufficiently close together to permit resonance to occur. Such cases are well known and it should perhaps be pointed out that it is not necessary that all the excess vibrational energy should be transferred from one molecule to another ; in fact it seems probable from elementary theoretical considerations that the excess energy is more likely to be transferred in small amounts in successive quanta. One other type of transfer which must be borne in mind is the transfer of vibrational energy from one molecule into translational energy of another.This is most easily under-stood in the case of the transfer of vibrational energy from vibrationally excited molecules to rare gas atoms. This latter type of transfer probably also occurs with light molecules. Prof. R. G. W. Norrish. (Cambridge) (communicated) The reaction RCH,CHO + O=RCHO + H,CO was originally postulated as a source of excited formaldehyde to account for the luminosity of the cool flame. In view of the very low luminosity observed (ca. I quantum per 10 mole-cules of hydrocarbon reacting) its probability in comparison with the other radical reactions would be very low in accordance with Dr. Walsh’s expressed view. However I have as will be observed in my paper, dissociated myself from this reaction except as a possible component in the dark blue flame and even then the suggestion must not be taken too literally for the reaction may occur in two stages : RCH,CHO + 0 = RCH CHO + OH RCH .CHO + OH = RCHO + H,CO. In any case I am not fully able to agree with limitations which Dr. Walsh would lay down for chain processes. Not enough is yet known about the possible configurations in the transition state. They conform to the views we have already expressed with reference to HO ; see for example the footnote in my paper or Axford and Norrish on the oxidation of formal-deh~de.4~ Prof. A. R. Ubbelohde and Mr. Small (Belfast) (communicated) : The suggestion that RO,* is in the first instance “ energy-rich ” and is more effective in reaction ( 2 ) above if i t can use this energy before i t is randomized by collisions etc.has an interesting experimental corollary. It has recently been observed 41 that molecular hydrogen inhibits certain reaction chains in both aldehyde and hydrocarbon oxidations much more markedly than molecular nitrogen. Tentatively this may be attributed to the efficiency with which hydrogen can effect the transfer between internal and translational molecular energy in collisions. I agree with his remarks about RO,*. RO,* + H + H,* + RO,. In (pseudo) unimolecular reactions the special efficiency of hydrogen collisions in the reverse process translational + internal is well known, and from the principle of microscopic reversibility collision process (3) would also be expected.Dr. A. D. Walsh (Leeds) (communicated) Prof. Sir Cyril Hinshelwood refers to my suggestion that bonds between strongly electronegative elements should be strengthened by electron-repelling groups. As originally put forward this was based upon the possibility of increas-ing the overlap of the two atomic orbitals concerned in a single bond 40 Axford and Norrish Proc. Roy. SOC. A 1948 192 518. 41 Small and Ubbelohde J . Chem. Soc. 1950 723 and unpublished results. L 324 GENERAL DISCUSSION between the elements. It is perhaps worth pointing out however that the suggestion could also be based upon the possibility of reducing lone pair-lone pair repulsion between the 0 atoms by attaching electron-repelling groups.The interaction between lone pair electrons on a halogen atom and electrons of a neighbouring system is greatest for the most electro-negative halogen (F) and least for I.43 A similar statement applies to interaction involving lone pair electrons on a Group VI eIeme11t.4~ The interaction between lone pairs on the adjacent 0 atoms of an -0-O-group is expected to be repulsive (witness e.g. the skew nature of the H,O, molecule). Attaching electron-repelling groups to an 0-0 group will reduce the effective electronegativity’ of each 0 atom and so be expected to reduce the lone pair-lone pair repulsion i.e. be expected to increase the 0-0 bond strength. Dr. N. Uri (Manchester) said In the primary step relating to the oxidation of hydrocarbons as postulated by’ Prof.Norrish viz., RH + 0 -+ R + HO, the endothermicity and activation energy are likely to be of the order of 40-70 kcal. if the dissociation energy of the HO radical into H + O Z ( D ~ o 2 ) is as low as 36 kcal. This latter value is obtained 4 4 9 45 from the electron affinity (in solution) of the HO radical which is in turn evaluated by inter-polation from the energies corresponding to the absorption maxima of various ferric ion pairs. While some unrecognized error may cause our quoted value of D H O ~ to be a few kcal. low it would be difficult to account from hydrocarbon chemistry for a value of D H O ~ as high as 65 k~a1.,4~ unless i t also assumed that the activation energy of the first step of the oxidation of hydrocarbons is much larger than its endothermicity.Dr. A. J. B. Robertson (King’s College London) said The heat evolved in the formation of HO may be determined by the method outlined by’ Stevenson in this Discussion. I find that hydrogen peroxide is dissoci-ated by electron impact to give the HO,+ ion with a small probability. The appearance potential is 16.1 eV. The ionization potential of H,O, is found to be 12.1 elr. The ionization of O, HO and H,O very prob-ably involves the removal of a weakly anti-bonding or non-bonding electron located approximately on an oxygen atom. In these circum-stances we may provisionally estimate the ionization potential of the HO, free radical as 12.2 eV intermediate between that of 0 and H,O, and i t may be less. This assumption gives 46 kcal. for the heat evolved in the formation of HO from H and O, subject to an experimental uncer-tainty a t present of about g kcal.If the ionization potential of HO is lower than 12-2 eV the formation of HO is even less exothermic. Dr. R. Spence (Harwell) said Prof. Norrish has raised many inter-esting points. I have always thought that the phenomena of com-bustion are so numerous and varied that general mechanisms should not be pressed too far unless all the well-established facts have been taken into account. For instance I do not think that the isolation of aldehydes as products of the reaction is necessarily evidence that they occur as essential intermediates. This was one of the important lessons to be learned from the early work of Bone who tried to establish the hydroxylation theory of combustion by the identification and separation of a set of products which were regarded as intermediates.Formalde-hyde for instance is a product of the oxidation of acetylene at 3 o 0 - 3 ~ 0 ~ C.4’ It is known that its rate of reaction with oxygen at these temperatures la See Baker and Hopkins J . Chem. SOC. 1949 1089. 43 Unpublished work of Baker and Barrett. 44 Evans and Uri Trans. Faraday SOC. 1949 45 224. 45 Evans Hush and Uri (in the course of publication). 46 Walsh J . Chem. SOC. 1948 331. 47 Spence and Kjstiakowsky J . Amer. Chem. SOC. 1930 52 4846. 48 Spence J. Chem. SOC. 1936 652 GENERAL DISCUSSION 32 5 is quite slow even at moderate partial pressures. Thus ordinary formaldehyde cannot be an intermediate in the combustion of acetylene ; i t can only be an intermediate if i t reacts whilst still in the excited state and i t is important as Dr.Walsh has mentioned earlier to recognize this in the mechanism. Then as regards the OH radical I have referred earlier in this dis-cussion to the work of Dr. Bates and myself 49 on the reaction between methyl radicals and oxygen and to the possibility of the formation of OH radicals or of CH,O in the primary step. More recently Blaedel Ogg and Leighton 5 O failed to detect the OH radical in this reaction by optical methods so that if present its concentration must be very small. Another piece of evidence which is not easy to fit into a simple radical chain theory was obtained by my wife 51 when working in Prof. Townend's laboratory. The propagation of a cool flame through a vertical tube at room temperature containing acetaldehyde ether hexane or heptane and oxygen or air is determined purely by thermal considerations.Free radical chain effects usually associated with the vessel wall and with the addition of inert gases were not observed. The addition of methyl radicals or of iodine produced no measurable effect ; the only indication that a chain mechanism might be operating was the inhibitory effect of additions of NO,. Another experimental result which is not easy to reconcile with a mechanism involving oxygen atoms is the homogeneous slow combustion of formaldehyde in reaction vessels of only I mm. diam.48 In this case, the chain carriers must be relatively unaffected by the wall. The surface oxidation only predominates when the vessel is packed with powdered glass.The " dark blue " flame mentioned by Prof. Norrish is to be identified with the " blue" flame of Townend and co-workers 51 which follows in the path of the " cool " flame. Prof. R. G. W. Norrish (Cambridge) (communicated) There can be no question that aldehydes are formed as intermediates in hydrocarbon oxidation because they are readily observed the question a t issue is whether they are to be identified as the moderately stable intermediate Iesponsible for the delayed branching. I have given the reasons in my paper for supposing that they are particularly in the second paragraph : I do not think that more weighty evidence has yet been produced for any other intermediate product playing this role.It must be remembered that the process of degenerate branching also takes place at temperatures where peroxides are completely unstable and unobservable and that Semenov's hypothesis demands a finite measureable concentration. The possible origin of aldehydes via a transitional peroxide is not ruled out on this account e.g., RCH + 0 -+ [RCH,O] + RCHO + OH as I have indicated in my paper. I do not feel that the remarks about the oxidation of acetylene are relevant. My paper did not concern this reaction but i t may be men-tioned that the rate of oxidation of formaldehyde at 286" C is appreciable (for partial pressures of formaldehyde and oxygen of IOO mm. in a vessel of 23-6 mm. diam. and volume 80 cm. the rate of reaction as measured by pressure change is 1-9 mm./rnir~.).~ Further i t is highly probable that the formaldehyde product in the acetylene oxidation is excited and thus more highly reactive.However I do not wish to commit myself to any view concerning acetylene oxidation in the present remarks. 49 Bates and Spence J . Amer. Chem. SOC. 1932 53 1689. 50 Blaedel Ogg and Leighton J . Amer. Chem. SOC. 1942 64 2499. 51 Kate Spence Thesis (Leeds 1945). 52 Axford and Norrish Proc. Roy. SOC. A 1948 192 518 326 GENERAL DISCUSSION As shown by Dr. Reagh and myself,5s the surface effect in the slow oxidation of hydrocarbons only becomes apparent at a certain limiting diameter of about 5-10 mm. The theory of slow branching requires that the reaction shall be entirely inhibited a t a finite diameter when the net branching factor becomes less than zero by the increase of the surface deactivation.This was found to be the case for methane ethane propane, acetylene ethylene and propylene. We should not expect the cool flame to be affected until a limiting diameter of tube was reached. Finally, with regard to the oxidation of formaldehyde this shows none of the criteria of a degenerate branched reaction and the results have been explained by us in terns of a straight chain mechanism. We should not expect the dramatic results observed with hydrocarbons on decreasing the diameter. There is however a competition in chain ending between the volume and surface reactions in which the surface effect only becomes predominant a t very small diameters ( < I mm.) corresponding to packed vessels.Nothing here appears to me to be inconsistent with our inter-pretation of the oxidation of formaldehyde as a straight chain reaction. With regard to the “ dark blue ” flame i t will be seen that I have drawn attention to the similar flames in acetaldehyde observed by Townend and his co-workers in my paper. Dr. G . Porter (Cambridge) (communicated) As far as I am aware, no short-lived free radical has ever been detected by its absorption spectrum in a chemical reaction at normal temperatures nor in an ordinary photo-chemical reaction despite numerous attempts with systems where radical mechanisms are known to be operative. I think it important to bear this in mind when assessing negative evidence of the kind mentioned by Dr. Spence though the workers quoted used a particularly sensitive method and their results certainly suggest a short lifetime of I O - ~ sec.or less of the OH radical if it was present. By using the method of flash photolpsis54 very high instantaneous concentrations of intermediates can be obtained and I have been able to observe the OH radical in absorption at high intensity in two systems which almost certainly involve the reaction between methyl and oxygen, viz. the photochemical oxidation of acetone and the chlorine photo-sensitized oxidation of methane. Unfortunately even this positive result does not enable one to conclude that OH is formed in the primary reaction of methyl with oxygen rather than in the subsequent steps though i t is hoped that studies of the kinetics of OH appearance will make this dis-tinction possible.Dr. L. Bateman (Welzuyn Garden C i t y ) (communicated) Norrish’s contention that peroxides do not absorb light and are therefore photo-chemically inert in the wavelength range 3000-4000 A is misleading and his arguments as presented have not the force claimed. Admittedly such absorption is weak E of typical saturated and allylic hydro- and di-peroxides ranging from about 0.5-10 at 3000 to 0-01-1 at 3650 A but Gee and I have presented quantitative Fvidence 55 that the very strong catalytic action of irradiation a t 3650 A on the autoxidation of liquid olefins originates in the photolysis of the hydroperoxide. This catalysis certainly extends to wavelengths as high as 4000 A and possibly higher.56 The quantum efficiency of the photolysis at 3650 is of the order of 0.1.However simple carbonyl compounds are even less absorbing in this spectral region (for acetone E~~~~ N o - o o ~ ) while photolysis is of com-parable efficiency to the peroxide (Norrish Crane and Saltmarsh 5 7 report + = 0-2 at 3150 A). 53 Norrish and Reagh Proc. Roy. SOC. A 1940 176 429. 54 Porter Proc. Roy. SOC. A 1950 200 284. 55 Bateman and Gee Proc. Boy. SOC. A 1948 195 376. 56 Bateman Trans. Faraday Soc. 1946 42 266. 57 Norrish Crane and Saltmarsh J . CAem. SUG 1934 1456 GENERAL DISCUSSION 327 Prof. R. G. W. Norrish (Cambridge) (communicated) I cannot sub-scribe to the cogency of Dr. Bateman’s criticism with acetaldehyde for example at 3400 the extinction coefficient is 1.0 whereas for tert.-butyl hydroperoxide and also for hydrogen peroxide the extinction coefficient is 0.1.In addition formaldehyde has an absorption coefficient of 0.07 a t 3560 while the peroxides have fallen to negligible values. I would point out that the peroxide with which he was concerned namely cyclo-hexene hydr~peroxide,~~ is one that would be expected to have consider-able absorption owing to the unsaturated nature of the compound. The combined effect of the double bond and the peroxide group will un-doubtedly cause light to be absorbed more strongly at longer wavelengths than in the compound where the peroxide group is the only chromophore. Add to this the fact that the intermediate concentration of aldehydes in the oxidation of methane and ethylene and in higher hydrocarbons at the higher oxidation temperatures is of the order of millimetres while the concentration of peroxides is undetectable analytically and I think i t must be agreed that the photochemical effect is to be ascribed to the aldehyde.In addition Mr. Booth in this laboratory has recently been studying the photolysis of tert.-butyl peroxide using a solution of 2 yo by volume in n-hexane. He finds no photolysis whatsoever when all wave-lengths below 3350 are filtered out. but considerable decomposition when only those wavelengths above 3000 Dr. C. F. Cullis (Imperial College) (communicated) Prof. Norrish has suggested that the large structural effect encountered in the oxida-tion of the normal paraffins is simply attributable to the greater number of points of attack in the longer molecules Thus he argues that n-octane is more readily oxidized than n-pentane since there are more CH groups in the former molecule and the removal of hydrogen atoms from such groups is considerably easier than from terminal methyl groups.On this view the increase in oxidation rate with chain length should be roughly proportional to the number of CH2 groups whereas in fact the variation is much more marked.59 In any case an extension of this argu-ment fails to explain qualitatively the facts relating to the oxidation of branched chain paraffins. It is well known in hydrocarbon chemistry that the order of reactivity of C-H bonds towards radical attack decreases in the order 3” > zo > IO, and Rice has shown that at 3ooOC the chances of attack at primary, secondary and tertiary carbon atoms are in the ratio I 3 33.60 For the isomeric hexanes for example i t is possible to calculate the relative probabilities of hydrogen atom removal from the molecules concerned by multiplying the numbers of such atoms attached to primary secondary and tertiary carbon atoms by Rice’s factors and summing over the whole molecule.If the ease of removal of a hydrogen atom from a hydrocarbon molecule is the main factor controlling the ease of oxidation as Prof. Norrish’s argument implies the figures in the final column but one of the following Table should be measures of the oxidizability of the compounds concerned. On this basis therefore the order of ease of combustion of the isomeric hexanes would be V > I1 = I11 > I > IV whereas the experimentally determined relative oxidation rates (given in the final column of the Table) show that the order is I > I1 > I11 > IV > V.59 Prof.R. G. W. Norrish (Cawbridge) (communicated) My remarks about the relative ease of oxidation if hydrocarbons are intended to indicate that other things being equal the ease of oxidation per carbon atom will increase from methane and ethane which will have values are removed. 58 Bateman and Gee PYOC. Roy. SOC. A 1948 195 376. Cullis and Hinshelwood Faraday Soc. Discussions 1947 2 I 17 ; Cullis and Mulcahy Revue l’lnst. Franpais Pe‘trole 1949 4 283. 60 Rice and Rice The Aliphatic Free Radicals (Johns Hopkins I935) p. 100 328 GENERAL DISCUSSION characteristic of the CH,.group to a value characteristic of the CH group for an infinitely long chain.Such comparisons as are made by Dr. Cullis in his criticism of my statement appear to me to have little relevance because the bases of comparison are so doubtful. Firstly his comparison of the rates of oxidation of hydrocarbons is based on the rates of pressure changes the reactants being at comparable pressures in all cases. This does not take account of the great increase in " carbon concentration " as we pass to higher hydrocarbons which would have the effect of weighting his results strongly in favour of the higher hydrocarbons. Secondly the comparisons could not all be made at the same temperature and the accumulation of intermediate products greatly' depends on the tem-perature as well as the nature of the hydrocarbon.For example in TABLE I1 Compound CH . CH . CH . CH . CH . CH (I) CH, I (11) (111) (IV) CH,. C H . CH . CH . CH, CH I CH,. CH,. CH . CH,. CH, CH, I I CH, I I CH, C . CH,. CH, CH,CH, CH,. CH . C H . CH (V) Relative Probabilities of H Atom Removal From IO C Atoms 4 x 1 9 x 1 9 x 1 I 2 x I I 2 x I From 2' C Atoms 8 x 3 4 x 3 4 x 3 2 x 3 -From 3" Z A toms Total telative Proba-ilitr of I Atom :emoval 30 54 54 18 78 telative Oxida-tion Rate 1580 560 60 I 2 I one experiment of Dr. Patnaik the oxidation of butane at 270' C using a butane oxygen ratio of 1-5 and a total pressure of butane of 76.1 mm., intermediate products such as aldehydes and alcohols accounted for over 25 mm.in a total pressure change of 60 mm. It is not legitimate to com-pare the rates of oxidation of hydrocarbons by pressure changes in view of this great variability of end-products and the only certain way is to measure the rate of disappearance of hydrocarbon by analytical technique. This has not yet been done. Finally Rice's figures refer to the relative reactivity of CH, CH and CH respectively with alkyl groups whereas in comparing the rates of oxidation one is concerned with their reactivity towards oxygen or OH. There is no reason why there should be any close quantitative parallelism in the two cases which will undoubtedly be differently conditioned by temperature and steric hindrance. Such calculations as have been made on this basis can have little quantitative value until much more data have been accumulated and until some more uniform basis of comparison of reaction rates has been devised.Dr. D. W. G. Style (London) said When illuminated with Schumann ultra-violet light diethyl peroxide and ethyl nitrate emits as a fluorescence the same band system which extends from approximately 3300 %L to 5000 A. The emiss:ons from methyl chloroformate and methyl nitrite are also identical but different from that from the first two substances GENERAL DISCUSSION 3 29 I t is probable that the emitters of the two band systems are respectively C,H,O. and CH,O-. Have these spectra been observed with other sources, particularly cool flames ? Dr. Peter Gray (Cambridge) (communicated) It is very interesting to see the spectra of molecules containing the alkoxyl group shown by Dr.Style and to learn that the emitter is likely to be the alkoxy radical itself. It is hoped soon to examine a system in which the methoxy (CH,O-) group thermally produced may be an emitter viz. the lumin-escent thermal decomposition of methyl nitrate. The initial decom-position of methyl nitrate gives nitrogen dioxide and methoxy groups and is accompanied by a blue luminescence. The same behaviour is displayed by methyl nitrite. The light may be due either to CH,O-or to excited formaldehyde CH,O* and now both spectra are known i t may be possible to identify the source of this luminescence. Sir Alfred Egerton (Imperial College) (communicated) The emission bands which Dr.Styles has found and ascribed to the C,H,O radical in the spectrum of glow in diethylperoxide are interesting. When Dr. Harris was working on the decomposition of diethylperoxide in my laboratory' we looked for bands in the glow but a t that time had not the means for photographs of long exposure. It is possible these bands have some relation to the pronounced absorption bands obtained in the com-bustion of hydrocarbons which Pidgeon and I reported in 1933. Dr. C. A. McDowell (Liverpool Univevsity) (communicated) With re-gard to the oxidation of acetaldehyde Dr. Thomas Mr. Farmer and I have shown fairly conclusively by inhibition studies G 1 ~ 6 2 that the thermal reaction u p to about 140'C is a chain reaction in which two radicals, the acetyl and peracetyl radicals play the predominant part.Dr. G. J. Minkoff (Imperial College) said The heats of combustion presented in our paper clearly have some peculiarities since the values of (-0-0-) which can be derived are of the order of 100-140 kcal. We intend to repeat thf measurements a t some future time but have mentioned the results already obtained as the work was done most care-fully and as the values are self-consistent within themselves (cf. Qc for members of series and for isomers). We have also recorded the refractive indices and melting points which we obtained so as to have as many published data on this subject as possible. These should be of use in determining the purities of future samples though we would stress the great advantages of determining the infra-red spectra of the peroxides during the preparations ; as the standard methods of purification (e.g.fractional distillation) are often inapplicable i t is a great help to know what particular impurities have to be removed. A brief point is that the latent heats of evaporation show high Trouton coefficients for all the hydroperoxides; this agrees with infra-red studies of the OH fre-quencies in suggesting that the compounds are strongly associated. DeteIminations of the heat of formation and of D(-0-0-) should thus if possible make allowance for the heat of dimerization etc., involved. Finally I should like to explain the use of the terms " acetyl, propionyl and butyryl hydroperoxides " to describe the corresponding peracids. Mr. Everett and I have measured the (classical) dissociation constants in aqueous solutions of these compounds and of H,O, Me0,H and EtO,H and have found that the peracids are intermediate in acid strength between the corresponding hydroperoxides and carboxylic acids.In fact the increase over the hydroperoxide can be shown to be approx-imately accounted for by inductive and polarizability effects. We there-fore suggest naming the peracids " acyl hydroperoxides " to avoid the implication of a great acid strength sometimes inferred from the term McDowell and Thomas J . Chem. SOC. 1949,2208 2217 ; 1950 1462 ; Trans. Faraday SOC. 1950 46 1030. G 2 McDowell and Farmer unpublished work 3 30 GENERAL DISCUSSION peracid (cf. perchloric and permanganic acids). Further if the diacyl peroxides are referred to as such there will be less chance of confusion with the corresponding acyl hydroperoxides (cf.dibenzyl peroxide benzoyl peroxide and perbenzoic acid). Dr. W. E. Vaughan (Emeryville California) (communicated) In the course of our own studies on organic peroxides we have obtained certain physical data on tert.-butyl hydroperoxide and di-tert.-butyl peroxide which differ from those just presented. Although we believe our com-pounds were better than gg yo pure we do not wish to claim superior accuracy for our data but wish rather to present them for the record in the hope that the differences may be resolved in the future. +alp . B.p. "C . Vapour pressure . K p . "C . Heat of vaporization . . Heat of combustion . tert.-Butyl hydroperoxide 1.4010 5'5 I 15-0 log10 P (mm.1 I 1-5 kcal./mole 65 4.2 kcal./mole Di-lert.-butyl Penoxide 1.3890 - 40.0 I 11'0 9.6 kcal. /mole 12 75.0 kcal. /mole Dr. M. Magat (Paris) said Leadbeater 63 has recently prepared in our laboratory some carefully purified peroxides and measured their refractive indices and Raman spectra. The following results were ob-tained : Diethyl peroxide nb5 = 1.3720 f 0-0005 Di-a-hydroxydiethyl peroxide nh6 = 1.4265 Ethyl-a-hydroxyhydroperoxide n:4 = 1.4150 f 0-0010. For the diethyl peroxide the agreement with the refractive index given by Egerton et al. is excellent. The characteristic Raman frequency for the 0-0 bond found in all peroxides investigated by Leadbeater as well as in dibenzylperoxide and in H,O is located at 880 -& 3 cm.-l.I would like also to call attention to the danger of working with very pure crystallized peroxides even with those classified as relatively stable : a spontaneous or induced rupture of the crystal has led to an extremely serious explosion when monohydroxydiethyl peroxide was purified by' crystallization. Dr. N. S . Wooding (Courtaulds Ltd. Coventry) (communicated) : Cobalt acetate in concentrations of IO- to 10-3 M has been shown to be an effective catalyst for the oxidation of trimethylethylene in solution by gaseous oxygen. Recently i t has been found that the autoxidation of cellulose under alkaline conditions was also catalyzed by cobalt acetate. The effect of other metal ions was investigated and some were found to catalyze while others retarded the autoxidation.Such behaviour has been reported elsewhere.65 However manganese salts were found to behave both as catalysts and as retarders depending upon the concentra-tion of salt used. A possible explanation of this effect has been suggested in terms of the mechanism of metal ion catalysis postulated in this paper and else~here.~42 O 5 I would like to ask Prof. Eawn if he has found any evidence for negative catalysis by metallic cations in the autoxidation 0.0005 63 Leadbeater Cow@. rend. 1950 230 829. 64 Entwistle Cole and ?Tooding Textile Res. J. 1948 19 527 609. 65 George and Robertson Tram. Faraday SOC. 1946 42 217 GENERAL DISCUSSION First Order Constant @in.-1) 33 1 Time of Half Cnange (-4 of trimethylethylene or any of the other systems he has investigated, since such phenomena if the suggested explanation is correct should be observable in other systems under the appropriate conditions.Prof. C. E. H. Bawn (Liverpool) (communicated) Under the experi-mental conditions so far studied the condition in which the rate of oxida-tion is independent of the catalyst concentration as observed by George and Robertson with saturated hydrocarbons has not been attained. There is no evidence at present that the catalyst terminates chains by a process such as RO + Co++ -+ R0,- + Co+++. Dr. C. F. H. Tipper (Edinburgh) said It is nearly always assumed when considering the oxidation of hydrocarbons in solution that the hydroperoxides initially formed decompose to give free radicals.However, under certain conditions for example in solvents of high dielectric constant or especially in the presence of acids as Dr. Waters has stressed, heterolytic fission of the 0-0 bond may occur to give ions. The evidence for this has so far been mainly' organic,66 and so I would like to report the results of some kinetic measurements on the decomposition of decalin hydroperoxide in various solvents. The overall decomposition was found to be first order in all cases, at any rate at low concentrations ( < I O - ~ mole/l.). The rates at 130' C and overall activation energies of decomposition in different solvents are shown below. TABLE I11 Chlorobenzene . . o-Dichlorobenzene . Acetic acid . Acetic acidlwater i 40% H,O by volume J Ethylene glycol .. Solvent Very slow <IO % decrease'in conc. in 3 hr. 0.002 68 258 30,250 0*0105 66 29,800 0.0332 21 16,700 0.359 2 22,140 I I I The addition of water to the ethylene glycol or the acetic acid thus presumably greatly increasing the dielectric constant of the solvent, increased the rate of decomposition considerably. In the case of the acetic acid addition of 2 yo by volume of water had little effect but further addition up to about 30 yo by volume caused a large increase in the rate. Above 30 yo increase in the water content had no effect. Also in chlorobenzene solution a t I 15.5' C no decomposition was detectable over a period of hours but if 1-5 mole yo of acetic acid was added the time of half decomposition fell to 64 min.These results would seem to show that in the first three solvents de-composition of the hydroperoxide molecules into radicals is taking place, possibly followed by a chain decomposition but that with a change in conditions an ionic decomposition can occur very readily'. Water and organic acids are very often stable products of oxidation of hydrocarbons in solution and thus as the reaction proceeds it is possible that ionic decomposition of the hydroperoxide formed might become important. Prof. Bawn (Liverpool) said IXr. S. F. Mellish and I have observed that the stable free radical aa-diphenyl p-picryl hy'drazyl reacts rapidly with the radicals of the type RO. and this provides a simple and convenient method for studying the rates of dissociation of peroxides into radicals Bartlett and Cotman J .Amer. Chem. SOC. 1950 72 3095. Kharasch Fono and Nudenberg J . Org. Chem. Igjo 15 748; 1951 16 113 128. 66 For example Robertson and V'aters J . Chem. Soc. 1948 1577 3 32 GENERAL DISCUSSION (ROOR -f RO. + RO.). The vividly coloured radical which gives stable solutions in a wide range of organic solvents undergoes a sharp colour change on reaction with radicals which may be measured in a simple colorimeter. This method has been used to measure the rate of dissocj-ation of polymerization in;t'ators such as peroxide and azonitriles. Dr. W. A. Waters (Oxford) said A similar instance to the decomposi-tion of decalin hydroperoxide quoted by Dr. Tipper is that of tetralin hydroperoxide which was reported some time ago by Robertson and This afforded a good example of a reaction which appeared from kinetic study to be much more simple than i t really was.The results reported by Dr. Henderson show the extent to which very minor products or impurities can influence the course of autoxida-t:on. In the autoxidation of a related compound dibenzyl ether (Ph . CH,),O which has recently' been studied by Mr. Wickham-Jones and myself we have been able to show that irregularities in the oxidation are due to the formation of a trace of a phenolic by-product. Dibenzyl ether gives a stable peroxide and in the main the uncatalyzed rate of oxygen absorption is independent of the concentration of the peroxide formed. Homolytic dissociation of the peroxide to give more chain-starting free radicals does not therefore play a major role in determining the oxygen uptake rate.However the autoxidation is self-retarding, and we have been able to show that as the peroxide of dibenzyl ether decomposes there is gradually formed just enough phenolic material to give a positive indophenol reaction. We ascribe the gradual retardation of the autoxidation to the formation of this phenol and have noted that when our reaction vessel was packed with chopped glass wool there was less peroxide decomposition less formation of phenol and less retardation of the autoxidation. Here the decisive factor seems to be surface catalysis of the mode of the secondary reaction-the peroxide decomposition-yet i t significantly influences the whole autoxidation process. Sir Alfred Egerton (Imperial College) (communicated) With refer-ence to Dr.Henderson's paper i t is well known that reproducible results are not obtained in hydrocarbon oxidation until the surface of the vessel has been conditioned by previous experiments ; the wall catalyst obtained in the liquid-phase oxidaton of ethylbenzene seems to be in line with this effect. Prof. J. P. Wibaut (Amsterdam) said In collaboration with Dr. A. Strang 6 8 we have carried out an investigation into the oxidation in the liquid phase by molecular oxygen of a number of normal alkanes with 8 to zz carbon atoms and of some branched octanes. When cobalt stearate is used as a catalyst the oxidation proceeds at a measurable rate. It has been found that the first stage of the oxidation reaction consists in the formation of a hydroperoxide.z 5-Dimethylhexane is slowly oxidized by molecular oxygen a t zoo C and a crystalline dihydro-peroxide melting a t 106.5~ C is formed : CH CH3 CH CH3 I I 1 I I I I H Ir H3C-C-CHz-CH2-C-CH3 + 2 0 2 4 H,C-C-CH2-CH2-C-CH3 0 0 0 0 I H H The presence of small quantities of peroxides can also be detected in samples of other saturated hydrocarbons which have been kept for a long time for instance in 3 4-dimethylhexane z 5-dimethylhexane, 3-methylheptane n-nonane n-hexadecane methylcyclohexane. The primarily formed peroxide is decomposed under the influence of the cobalt 67 J . Chem. Soc. 1948 1578. 68 W'ibaut and Strang Proc. Kon. Neder. Akad. Wet. B 1951 54 ( z ) 101 GENERAL DISCUSSION 333 ions ; the radicals thus formed start a chain mechanism so that the oxida-tion reaction proceeds : ROOH + Co++ -+ RO' + OH- + Co+++ ROOH ~2 ROO- + H+ Co+++ + ROO- -+ Co++ + ROO.We ascertained what products are formed by the oxidation of 2 5-dimethylhexane ; with reference to these reaction products a reaction scheme can be drawn up. The characteristic feature of the reaction scheme for the catalytic oxidation in the liquid phase of z 5-dimethyl-hexane and of other branched hydrocarbons is that the chain mechanism is initiated by an alkoxy radical. We ascertained the maximum rate of oxidation of normal alkanes with 8 g 10 12 14 16 18 20 and 22 carbon atoms. The experiments were carried out at 110.4" C and with 0.112 mmoles cobalt stearate per 61.7 mmoles hydrocarbon.There is a linear relation between the number of carbon atoms and the maximum rate of oxidation from C,,H, to C Z z H 4 6 . This can be explained by assuming that all the secondary carbon atoms have an equal chance of reaction. Some branched alkanes oxidized much more easily than the isomers with normal chains ; we ascertained the maximum rate of oxidation at 78.1" C measured in ml. 0 per 61.4 mmole hydrocarbon per hour (catalyst 70 mg. cobalt stearate) to be : a-methylheptane . . 5.0 3-methylheptane . . 3'0 3 4-dimethy'lhexane . . 11.0 z 5-dimethylhexane - 35'0 3-methyl-3-ethylpentane . 0-0 z 2 4-trimethy'lpentane . 0-0 The oxidation begins at a tertiary C-H bond from which a hydro-peroxide group is formed. A considerable quantity' of acetone is formed during the oxidation of 2-methylheptane methylethylketone being formed during the oxidation of 3-methylheptane.The fact that z z 4-trimethylpentane (iso-octane) is not oxidizable under the conditions chosen can be explained by steric hindrance. The quaternary carbon atom which has three methyl groups screens the tertiary carbon atom to such a degree that an oxygen atom cannot approach the tertiary hydrogen atom to within the distance pertaining in the transition state. Dr. M. F. R. Mulcahy (Melbourne) (counrtzunicated) Several con-tributors 89-7 have discussed the liquid-phase oxidation of hydrocarbons catalyzed by decomposing peroxide or metallic catalysts in terms of the propagation mechanism : R- + 0 + ROZ-ROZ- + RH + ROOH + R-. It may perhaps be of interest to recall that there is evidence that in the absence of catalysts (and of light) formation of hydroperoxide may occur by' a different mechanism.This was shown by the work of George and Robertson 73 on the " thermal " oxidation of tetralin. A similar result has recently been found by Mr. Watt in this laboratory with the uncatalyzed oxidation of benzaldehyde,' the kinetic behaviour being analogous to that of the tetra'in reaction. In the presence of benzoy'l peroxide however a reaction of the type shown above is initiated and is (additionally) superimposed on the uncatalyzed reaction. 69 Bateman Gee Morris and Watson this Discussion. 7O Bau-n Pennington and Tipper this Discussion. 71 Brook and Matthews this Discussion. 72 Mulcahy this Discussion. 73 George and Robertson Proc.Roy. SOC. A 1946 183 309 ; George Proc. 74 Mulcahy and Watt Nature (in press). Roy. SOC. A 1946 183 337 334 GENERAL DISCUSSION Dr. N. Uri (Munchester) said The liquid-phase reactions in which a wall effect is observed are few and therefore remarkable. There are ody four such cases known to me : (a) the experiments relating to the oxidation of liquid ethyl benzene described by G. M. Henderson in this Discussion; (b) my own findings in the catalytic decomposition of hydrogen peroxide under certain experimental conditions ; 75 (G) the observations made by Dain and Kachan 7 G in the photochemical oxidation of water by ceric ions. It is not unlikely that the wall effect in this case leads to a recombination of OH radicals and a subsequent instantaneous decomposition of hydrogen peroxide by’ ceric ion.(d) In our work on the photo-initiated free radical polymerization of vinyl compounds in aqueous solution77 we made the observation that under conditions when practically all the free radicals are formed a t a distance < rmm. from the wall the polymer is produced (by’ recombina-tion of active endings) exclusively on the wall; none appears to be SUS-pended in solution. No doubt these experiments require some co-ordination and in this connection i t is interesting to note that all these reactions involve free radicals as intermediates and the process effected by the wall is con-sidered to be a termination process which may in some cases lead to a new type of chain reaction as reported by Henderson.In the oxidation of lubricating oils the effect of iron or copper as single catalysts was studied by Brook and Matthews. It is well known that in the catalytic decomposition of hydrogen peroxide the joint action of iron and copper as co-catalysts is much more than additive. It would therefore be interesting to study those effects in the oxidation of hydrocarbons particularly if peroxides are postulated as intermediate products. Dr. G . M. Henderson (Blackley) (communicated) The effects of walls on liquid-phase chain reactions are fairly’ well indicated in past work and to restrict it to two examples on oxidations excluding all references to solid catalysts the most relevant papers are those of Stephens 78 working with cyclohexene and Medvedev and Podyapol~kaya,~~ working with tetra-lin where many effects akin to the present work were noted and the effects of walls mentioned only with very different explanations. The novel feature of ethyl benzene to which we wished to draw attention as a convenient expet imental medium is that the two hydroperoxide forming mechanisms seem more clearly distinguishable and mutually exclusive. With very fresh and pure ethyl benzene we presume that the original wall termination process is without further visible effect but that in an older sample or after sufficient of the initial peroxide has further decom-posed some of the breakdown products possibly acids act as haptens (to borrow a term from immunology) and a new wall termination process arises which gives rise to visible effects on the rate of reaction. One can generalize that in other similar reactions that the two possible re-actions might then run concurrently or that the second stage might never even arise. Sir Alfred Egerton (Imperial College) (communicated) This paper refers particularly to the inhibition of the oxidation of lubricating oils, I would like to draw the authors’ attention to a paper by Hanson and 75 Uri J . Physic. Chem. 1949 53 1070. 76 Dain and Kachan A C.S. Abstr. 1949 43 7349. 77 Unpublished observations. 78 Stephens J . Amer. Chew. Soc. 1936 58 219. 7g Medvedev and Podyapolskaya J . Physic. Chew. U.S.S.R. 1939 12 79 tiENERAL DISCUSSION 335 myself 8o on " Nitrogen oxides in internal combustion engine gases " in which the promoting action of the nitrogen oxides on the oxidation of lubricating oils was investigated and a subsequent paper entitled " In-fluence of catalysis on oil oxidation " 81 iron oxide was found to inhibit the nitrogen oxide catalyzed reaction. Dr. J. B. Matthews and Mr. J. H. T. Brook (Thornton) (communi-cated) In connection with Dr. Uri's suggestion that the study of mixed iron and copper catalysts would be of interest in the oxidation of hydro-carbons] results have been given in the paper showing that in the inhibited reaction the two catalysts are additive in their effect on the initial rate of reaction. In the uninhibited reaction] however the measured values of t$ are greater than the added effects of the individual catalysts thus lending support to the inclusion of peroxides as intermediate products. Values of t$ for iron and mixed iron-copper catalysts are given in Table 11 c s I 0 WEiCHT FRACTION OF HYDROCARBON FIG. I. of the paper and values of 4 for copper can be calculated from the data in Table I together with values of the intercepts with the oxygen absorbed axis of the asymptotes to the curves in Fig. 3 using the equation to describe the asymptote. Mr. J. H. T. Brook ( T ~ C Y M ~ O O ~ Z ) said I would like to add some further information on the kinetics of the iron-catalyzed inhibited oxidation. Using 50 p.p.m. of iron added as ferric stearate at 1 5 0 O C and using diphenyl as the inert solvent the dependency of the reaction upon the oil concentration was found to be of the form (as Fig. I above) : = ( A + B)t - (B/4) [RHI cc 220 + 330 [RH]' No immediate explanation of the difference in kinetics between the iron and copper-catalyzed reactions is apparent. ' 0 PYOC. ROY. SOC. A 1937 153 90. 81 Symposium on Engine Wear Inst. Mech. Eng. 1937

ISSN:0366-9033    

DOI:10.1039/DF9511000308

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

TABLES OF THE HEATS OF FORMATION AND IONIZATION POTENTIALS OF HYDROCARBON RADICALS BY M. SZWARC Received 20th July, 1951 The values of bond dissociation energies and heats of formation of radicals have been collected recently in two publications.*t The latter contains all the experimental data published up to the end of 1949, thus enabling one to assess the uncertainties involved. The work done during the last two years provides some new data and requires a slight revision of the old. The present meeting provides an opportunity to supplement these Tables t and thus to bring the data on hydrocarbon radicals up to date. The following equation enables one to compute the desired R-R, bond dissociation energy : The heats of formation used in this equation should be computed for the relevant species in the gaseous state and at oo K.However, the extra- po!ations to oo K are somehow uncertain, and it seems that the corrections which should be introduced are of the same magnitude as, or perhaps even smaller than, the experimental errors involved in the currently available data on bond dissociation energies. Consequently, the values AHf(R . R,) corresponding to 25" C have been used for calculating the heats of formation of radicals listed in Table I. Our knowledge of ionization poten tia?s of hydrocarbon radicals is very scanty. The ionization potentials of CH,, C,H5, sec.-C,H,, and tert.-C,H, have been determined by electron impact method, and the relevant data are summarized in the paper by Stevenson. There are, however, data which enable us to calculate ionization potentials of few more hydrocarbon radicals.(i) PROPYL RADICAL.-The appearance potential of C,H,+ ion pro- duced in the mass-spectrum of n-butane is given by Stevenson and Hipple 8 at 11-14 eV/molecule. D(CH, . CH, . CH,-CH,) is calculated as 3-57 eV/ molecule on the basis of D(CH, . CH, . CH,-H) = 4-3 eV/molecule (see ref. (2) and (3)). (ii) ALLYL RADICAL.-The appearance potential of C,H,+ ion pro- duced in the mass-spectrum of butene-1 is given by Stevenson s as 11-65 eV/molecule.$ D(CH, : CH . CH,-CH,) has been determined 1 0 as 2.6 eV/molecule. Therefore, Iz(CH, : CH . CH,) = 9.0, eV/molecule. This value seems to be reasonable, the ionization potential of ally1 radical is expected to be only slightly lower than that of methyl radical.This has been the writers' intention in preparing Table I. D(R-R,) =z AH,(R) + AHf(R1) - AH,(R. R1). Therefore, Iz(CH,. CH, . CH,) = 7.5, eV/molecule. * Roberts and Skinner, Trans. Faraday SOC., 1949, 45, 339. t Szwarc, Chem. Rev., 1950. 47, 75. 8 Stevenson and Hipple, J . Amer. Chem. SOC., 1942, 64, 1588. 9 Stevenson, J . Amer. Chem. SOC., 1943, 65, 299. $ The objections which Stevenson raises in his paper seems to be irrelevant. I t appears that electron impact dissociates the molecule of propylene into CH, : C . CH, ion and not into CH, : CH . CH,+ ion. Indeed, this assumption is consistent with the results of electron impact studies of iso-butene. This probkm is discussed fully by Evans and Szwarc, J . Clwn. Physics (in press). 10 Sehon and Szwarc, Proc.Roy. SOC. A , 1950, 202, 263. 336M. SZWARC 337 TABLE I.-TABLE OF C-H BOND DISSOCIATION ENERGIES AND HEATS OF FORMATION OF HYDROCARBON RADICALS Radical CH, . C*H5 * n-C,H, iso-C,H, . tert.-C4H, . C2H3 * CH, : CH . CH,' . C H , : C . CH,' . HC: C - YL-C~H, CH,' Ph- . Ph.CH,* . P-CH, . C6H4. CH,' WZ-CH,. C6H4. CH,' o-CH,. C G H d , CH,* Ph.CH- . dH3 CH, CH, P h . $-CH,. C,H, . CH* CH, '">,He . Ph ph\ ;v= -CH; m D (R-H) kcal. /mole 101-102 96-97 - 9 9 ? - 94 N I01 ? - 88 92-104 77-78 - 76 < I 2 1 104 ? 77'5 75 77'5 - 74 - 74 - 71 - 74 - 72 75 - 76 - 76 AH, f W kcal./mole 31-32 24-25 - 22 ? - 17 N 1 9 . - 4'5 52-64 29. j-30'5 - 20 - I20 ? - 72 ? 37'5 27 29.5 - 26.5 - 29 - 20 - 2 1 - 53 Experimental Data published since 1949 Pyrolysis CH, . NO, { Ph PZgi?(?2H5 Electron impact Pyrolysis Ph .CH, . C,H, Electron impact Electron impact Electron impact Pyrolysis of dibenzyl Pyrolysis PhCH (CH 3) Pyrolysis PhC(CIX,) Pyrolysis P-CH3. CeHdCH(CH3)Z Pyrolysis Ph, . CH2 Ref. N 0. I 2 3 2 4 3 4 Cottrell, Graham and Reid, Trans. Faraday SOC., 1951, 47, 584. Leigh and Szwarc (in course of publication). Stevenson, this Discussion. 3 Stevenson (private communication). 5 Horrex and Miles, this Discussion. 6 Leigh and Szwarc (in course of publication). 7 Horrex and McCrae (communication to this Discussion).338 (ii;) VINYL RADIcAL-The appearance potentials of C,H,+ and C2H,+ ions in the mass-spectrum of butene-1 make i t possible to calculate D(CH, : CH-H) and ionization potential of vinyl radical. However, these calculations lead to D(CH, : CH-CH) = g z kcal./mole, a value which seems to be too low. The data on ionization potentiah of the six hydrocarbon radicals dis- cussed above are listed in Table 11. TABLES OF HEATS OF FORMATION TABLE 11.-IONIZATION POTENTIALS OF HYDROCARBON RADICALS Radical CH, . n-C,H, . CH,:CH.CH,' : tert.-C,H, . C2H6 - ZSO-C~H, Ionization Potential eV /molecule 1 k cal. /mole 10'1 8'7 7.57 7'4s 9-05 6.9 232 I74 172 208 I59 2 00 I wish to express my appreciation to Dr. Stevenson for allowing me to use his unpublished results. Chemistry Department, University of Manchester, Manchester, 13.

ISSN:0366-9033    

DOI:10.1039/DF9511000336

出版商:RSC    

年代:1951

数据来源: RSC

摘要:

AUTHOR INDEX * Anderson, R. D., 136. Aston, John G., 73, 119. Bartindale, G. W, R., 104. Bateman, L., 219,227, 250,313, 326. Baughan, E. C., 106. Bawn, C. E. € I . , 282, 331. Bell, E. R., 242. Benson, S. W., 218, 224, 232, 234, 310. Blekkingh, J. J. A., 120. Bowen, E. J., 224. Bradley, R. S., 127. Brook, J. H. T., 298, 335. Brown, J. I<., 118. Burawoy, A., 104, 107. Burton, M., 136, 221. Cox, E. G., 127. Cullis, C. F., 327. Davies, C. N., 124. Davison, S., 136. Egerton, A. C., 278. Egerton, Sir Alfred, 311, 329, 332, 334. Eider, B., 235. Emte, W., 278. Evans, Alwyn G., 109. Evans, M. G., I. Gaydon, A. G., 108, 311. Gee, G., 214, 250. Gibbs, Julian H., 122. Glockler, Geo., 26. Gordon, Manfred, 125. Gowenlock. B. G., 108, 229. Gray, Peter, 128, 310, 317, 3-29. Guggenheim, E.A., 116, 118, 226. Hall, G. G., 18. Harding, A. J., 317. Harris, E. I., 318. Hazebroek, P., 87. Henderson, G. M., 291, 334. Hinshelwood, C. N., 129, 21.5, 217, 218, 266. Horrex, C., 187, 230, 232, 233, 234. Hughes, Hilda, 3 I 3. Jenkins, A. D., 311. Kistiakowsky, G. B., 175. Kooyman, E. C., 163, 224. Lapage, Ruth, 233. Lennard-Jones, Sir John, 9, 18, Linnett, J. W., 119. Luft, N. W., 117. McCoubrey, J. C., 94, 127. McCrae, J. O., 234. McDowell, C. A., 53, 106, IIO 116, 103, 104, 106. 310, 322, 329. Magat, M., 113, 118, 126, 226, 235. Marcotte, Frank'R., 236. Matthews, J. B., 298, 335. Melville, H. W., 154, 225. Miles, S. E., 187, 230, 233. Minkoff, G. J., 108, 278, 319, 329. Morris, A. L., 250. Mulcahy, M. F. R., 259, 317, 333. Nickle, A. Gordon, 175. Norrish, R.G. W., 269, 323, 325, Noyes, Jr., W. A., 221, 236, 308. Oosterhoff, L. J., 79, 87, 122. Partington, R. G., 219. Pennington, A. A., 282. Pitzer, Kenneth S., 66, 119, 124, 127, 226. Pople, J. A., 9. Porter, George, 108, 115, 223, 326. Raley, J. H., 242. Robb, J. C., 154, 225. Robertson, A. J. B., 324. Rowley, D., 198. Rust, F. F., 242. Schissler, D. O., 46. Seubold, F. H., 242. Sheppard, N., 118. Sheridan, J.3 120. Small, Mr., 323. Spence, R., 309, 324. Steiner, H., 112, 198, 235. Stevenson, D. P., 35, 110, 113. Stubbs, F. J., 129, 215, 217, 218. Sxyle, D. W. G., 328. Szwarc, M., 143, 215, 228, 231, 336. Thompson, S. O., 46. Tipper, C. F. H., 282, 331. Torkington, P., 104, 108, 128, 213, 223. Trotman-Dickenson, A. F., 111, 128, Turkevich, John, 46. Tutton, R. C., 154, 225. Ubbelohde, A. R., 94, 103, 124, 127, 128, 309, 323. Uri, N., 309, 324, 334. Vaughan, W. E., 242, 314, 330. Walsh, A. D., 310, 320, 323. Warren, J. W., 53, 110, 116. Wasserman, A., 235. Waters, W. A., 231, 315, 319, 332. Watson, W. F., 250. Wibaut, J. P., 332. Wooding, N. S., 330. 330. 327. 221. _ _ ~ * The references in heavy type indicate papers submitted for discussion. 339

ISSN:0366-9033    

DOI:10.1039/DF9511000339

出版商:RSC    

年代:1951

数据来源: RSC