Electron spin resonance studies of gas-phase oxidation reactions. The hydrogen-oxygen system at atmospheric pressure

COMBbSTIONAYD FLAME 23, 47-55 (1974)

47

E;l'ectron Spin Resc)nance Studies of Gas-Phase Oxidation Reactions. The Hydrogen-Oxygen System at Atmospheric Pressure AKISS AGKPO and LOUI,~-RENI~ SOCHET* Laboratcire de Citl ~tique et Chimie de la Combustion, et Cent,,e de Spectrochimie. Universit~ des Sciences et Techniques de Lille. B.P. 36-59650 Villeneuue D'Ascq, France

An appar;!Etus is described in which radicals can be identified by ESR and molecular products by polar. ography c r gas chromatography. The radicals are removed from the reaction mixture thr mgh a micro~eak and condensed in a matrix on a cold fintgar. From a study of the system, 1 H2 (1 D2) - 1 02 ,i N2 at 5,:1.6°Cand a pressure of I aim, in reactors with and without a B203 coal ag, evidence has have been pro:lueed for the presence of HO.z and DO~ radicals. Following the earlier suggestion that the iormation of H2Oa occurs by bimolecular reaction of H02, a mechanism is studied and compared with the el
/~though it is l,,enerally believed that free radical chain processes are important in combustion reactions, th~1 existe~ ee of ra-icals has seldom been demonst rated dilectly. Knowledge of t he identity of the radicals and of their cop centrati::)n throughout the course of the ]leaction would give us a more detailed understanding o f th." reaetic:a mechanism. Several methods of s~udying radic:ds have been employed. In parllienlar, mass spe{ trometry has made useful ad,~ances i~ this fie/d, and the same is true of emis~:ion and adsorptJ,on spectroscopy. Nevertheless, progress has b,.'.er~ slow because of difficulties in interpretation an:., above all, because of the high concentration of stable 'species interfering with the !ipectra c f the sh.rt-lived radicals present in very low cone entratioa. On the other h;~tnd, ESR is more selective. Several stud;,es have )een made using low-pressure flames, and r.~ce~ fly a few papers [1,2] have reported t t e reautt~ of experiments with flames at *Please address correspoudence to this author.

atmospheric pressure. In slow combustion reactions, since the radical concentrations are very low, it was found necessary to condense the radicals on a cold surface [3, 4], This paper describes an example o f the ESR technique and shows the advantages o f the simultaneous spectroscopic and kinetic study of gasphase oxidation reactions. 1. Experimental The apparatus is shown in Fig. ! .~r,d comprises an ordinary flow system operating at atmospheric pressure from which a sample is taken thrc,ugh a microleak into a l~w-pressure (0.1-1 Torr) region. The reaction vessd (Ve) was a silica cylinder, diameter 3.6 cm and volume 95 cm a . Experiments were performed with both an untreated reaction vessel and ol~e whose surface had been coated with B~Oa. The flows of reactants were controlled by needle valves (V) and the individual flows measured by rotameters (F). The total flow was measured by the soap-bubble meter (F6). The reactants were preheated (Pr) before entering the reactor and on leaving the reactor the Copyright © 19"/4 by The Combustion Institute Published by American Elsevier Publishing Company, Inc.

48

AKISSI AGKPOand LOUIS-RENE,~;OCHET

,

02

H2N2

•

Qh

Fig. I. Schematicof appara$.usused for the simultaneouskinetic study by ESR, chromatographyand pola~ographyof gas-phaseoxidation reactions. products were passed through a.trap at - 80°C to remove water and hydrogen peroxide before being analyzed by gztschromatograpiay(Ch). Hydrogen peroxide was condensed m another trap at the same temperature and analyzed~iby pclarography. A sample of the reactants w~s taker, through a silica leak (diameter 20-100 #) to a silica cold finger cooled with liquid nitrogen, the end of which was situated in the resonant cavity of the ESR spectrometer situated be~:ween the poles of the electromagnet (H). A measured flow (F7) of a matrix material (usuall~ CO;~) entered the lowpressure region through a side arm. The total flow through the low-pressure "egion was measared by Fs and the pre-.sure measured by the pirani ~auge G, For a given contact time (5), after the reactant concentrations as shown by the Beckman Analysar (E) had become steady, a sampl.* was injected from the loop into the chromatograph. Than, by turning Ts fire mixture was made to pass via tube a i:lstead of tube b. The condensable products were trapped over a 15-rain period in two traps in series (P) for subsequent polmrographic analysis. Wl-~.nthe flow of reactants through the leak it, to

the low-pressure chamber and the flow of matrix gas had become steady, liquid nitrogen was placed in the cold finger and the spectrum recorded as a function of time. At n ,~ve~lmoment the flow of matrix gas was stopped by o~'eration f T7 and the reactor e,,acuated through Tq. Under these eondition~ the spectrum did not change with time, and could be recorded and studied more e .silly. The concentration of radicals was determined by c~mparL, u with the spectrum given under the same conditions at 77°K by 5/~l of a freshly prepared solution of diphenylpicrythydrazylof lmown strength,. This apparatus is a development of that described earlier [3-8] and enables the nature of the trapped radicals to be studied and the relationship between the concentrations of radieat and molecular species to be determined [5, 9]. The work described here concerns the oxidation of a mixture of H2 and air at 546°C and 1-atm ' pressure, 2, Spectroscopic Analysis 2.1. Detection and Characterizationof the HO~ Rad~a! The spectrum given by the products trapped in the

ESR STUDIES OF THE HYDROGEN-OXYGENSYSTEM

g=2,0028

49 other radicals appears to be negligible, in agreement with the kinetic study (see Table 3).

i

9~" 'g, |

2.2 Detection and Charae~erizatio~of the DO= Radical To confirm the assignment of the spectrum observed in the oxidation of hydrogen to the HO2 radical, these experiments have been repeated under the same conditions using deuterium. As remarked in paper [15], the concentration of the molecular species (02, D202) and tl~,e kinetics are the same, but the observed spectrum is quite different (see Fig. 3). It is characterized by

HO"2

U 9,;

91;

9=2j0028 1 =91

Fif~. 2. ESR spectrum of HO2 radical. /

matrix (see Fig. 2) was the same for all contact times studied. The two peaks (~) are due to gas-phase molecular oxygen. The spectrum attributed to the HO2 radicals [10] is characterized by anisotropy and hyperfine structure due to the interaction of the odd electron with the single H-atom (I = V2) of ~ e radical. The measured parameters are gll = 2.0348,g~. = 2.0057,a, = 15 g,a± = 12 g. These values are in good agrezment with those obtained by others. Using the rotating cryostat method [11 ], the following parameters were measured at liquid nitrogen temperature in a matrix of HI/H20 gll = 2.035 l, gz = 2.0039, a = 11.5 g. By phntolysis of HI, foliowed by reaction with oxygen and trapping at 4°K in an argon matrix, other authors [12, 13] found gll = 2.0393, gl = 2.004.4, all = t3.5 g, a± = 8.6 g. The small differences are ascribable to variations in the tem~'¢rature of trapping' ; in the matrix material. If a splitting of about 508 g is assigned to the free H-atom, the splitting of about 13 g for HO~ radical implies that about 2% of the unpaired spin density is centered at the H-atom. Under similar conditions of temperature and pressure, the spectrum of the HO2 radical has already been observed [3, 14, 15] and has shown evidence of a reactivity quite different from that shown in the explosion peninsula. In this system the concentration of

Do'2

t

2o__.~g

/

J

Fig. 3. ESR spectrum of DO2 radical a very marked anisotropy without apparent hypert'me structure. We assign this spectrum to the DO~ radical. The mea~.ured parameters are gll = 2.0368,g t = 2.0088; they are approximately the same as those obtained in a study [13] of DO2 mdical~ formed in the pEotolysis of a D20~/D20 glass at 77°K (gz = 2.0344,g2 = 2.0087,ga = 2.0031)'. -£he hyperfine stracture d~"~ to the D atom ( I = 1) is very feeble z.nd is or perceptible as a very small shoulder, in part J, on the peakg 1. I f a splitting of about 508 ~, =.,J 78 g is ascribed to free atoms of H and D and a splitting

.......

a~m~ 3

/

100

50

~b

-180

-150 '

-121') ' ~

T ° c :'

Fig. 4. Decay of H02 radical diJting "pulse" warming of the matrlLx.

of about 13 g is observed for the EIO~ rad:a~l, a splitting of about 2 g must be expected for the DO~ radi,gal. 2.3. Effect of Matrix TemperJttureon the Stmcf,tre and Concentration of Radicals In order to study the effect of ~ e te:uperature of the matrix on the structure and co~c.'ntration of trapped raaicals, an apparatus was c,mstructed which could be operated at tempe,a~ures between that of Liquid nitrogen and that of th~ laboratory. Gaseous nitrogen was passed through a c,;~ cooled in liquid nitrogen and then through a therm, statically controlled heater. This cooled nitrogen was passed through the cold finger and the temperatare of the matrix was mcssured by a copperconstantan thermocouple immersed in the matrix. Radicals were then collected for a known time with the cold finger at the temperature of liquid nitrogen and alter the react~,:,n and the flow of C02 had been stopped, th ,° . ,.itrix was isolated while the reaction vessel wa~ ' vacuated (or a slow stream of nitrogen was pa:;sea). A temperature gradient was established between the liquid rdtrogen in the cold tinge~ and the temperature of the matrix (- 165°C). At this temperatufa, neve.-theless, the radicals are stabile and there was no detectable change in the spectrum with

1~o

.... 2 1s0~r=

Fig. 5. Concentration profiles for HO2, H202, arid O2 during oxidation of the mixture 1 H2 - 1 02 - 4 N2 at atmospheric pressure and T = 546°C in a 95 cc silica vessel. time. The effect of changing temperature was studied as f,Alows: the matrix was warmed to A T for 5 rain and then recooled to - 165°C, then reheated to A T '° ( A T ' > A T ) and cooled to - 165°C and so on. At temperatures above - 165°C, the observed spectra did not show ~my special features, except that the line width was slightly enhanced. FollowirJg heating and cooling to - 165°C, the ~pectra were the same as shown in Fig. 2; the signal height h decreased and neaiIy all the radicals disappeared by quadratic recombination at T = - 120°C (Fig. 4). 3. Kinetic Study 3.1. Effect of Contact Time and Surface Treatmem The concentration of HO~ radicals, the consumption of oxygen, and the formation of hydrogen peroxide have been determined over a range of contact times in treated and untreated vessels. Under our experimental conditions, hydrogen peroxide was the only compound that lead to degenerate chaL'~-branching. The experimental remits are shown in Figs. 5 and 6. In an untreated y actor (Fig. 5), the H:O2 and HO~ radical concentrations grow exponentially in the early stages and then reach maxima; thereafter, the H202 falls more quickly than the HO~. The maxima are reached in both cases at very low oxygen consumption (2%).

ESR STUDIES OF THE HYDRGGEN-OXYGENSYSTEM H202 I'106

.o~ !~_9

,.v,ov,~

O~4I03 "

T,5 ~.6 *¢ a2o3 2C

~

lS

S0

100

~S0.~i

2

Fig. 6. Concentration prof'flesfor HO2, H202, and 02 during the slow oxidation of the mixture 1 H2 - 1 02 4 N2 at atmospheric pressure and T = 546°C in a 95 ce B203 silica vessel. In a freshly B203-coated reactor (Fig. 6), in which the destruction of radicals and H202 at the wall is known to be small, the above observations remain true. However, the maximum concentrations of HO2 and of H202 are greater (see Table 2), the induction period (def'med as the time from the start of the reaction to the maximum rate) is slightly less, the fall in 1-[202 concentration is less rapid and the HO~ c~ncentration remains practically constant. Our experimental results are in good agreement with ~hose obtained by Shakhnazarian et al. [ 15, 27j by a similar technique. These experimental facts lead us to the following conclusions: (a) Treatment with B:O3.erthances the cc,neentration of H202 and HO2 radicals in agreement with the usual behavior of this type of coating [16]. (b) I"I202 appears as an intermediate in trace amounts. Practically all the H2 consumed is found as water. (c) The ratio ffI202)/(HO~)~sabout 102,pro riding experimental justif':ation for the use of Semenov's partial stationary concentration principle in kinetic calculations [17]. (d) Whatever the nature of the vessel surface, H202 and HO2 reach maximum concentrations for a very small consumption (2 and 4%) of the initial products. The maxima cannot be attributed

51 to reactant consumption, but are probably due to radical-radical [18] interaction or reaction between radicals and the branching intermediate [5, 10]. 3.2. lnfluenee of Surface Coating on the Third Explcsion Limit "l'~e • third explosion iim~t was determined in both B~O3-treated and -untreated vessels. For a constant time of 60 sec, the temperature was increased in 2 ° stepz starting at 546°C. In both cases, the limit was reached at 575°C, clearly above that observed in a static system and close to that previously recorded in a flow system [19]. The displacement of the explosion limit toward lower temperatures in a static system is probably due to changes in temperature occurring on adrrdssion of reactants to the reaction vessel [20]. Unexpectedly, treatment of the vessel with B203 does not affect either the maximum rate (Figs. 5 and 6) or the position of the limit, although it noticeably increases the concentration of both HO2 radicals and H202. In this connection, it may be remarked that the third limit is ,.'ssentially thermal [21 ]. 4. Study of the Mechanism and Comparison with Experimental Resulis The mechanism of hydrogen-oxygen reaction has been examined thoroughly by many workers and especially by Baldwin an~t Walker [22, 23]. The mechanism generally used, ineludas the 0-11 reactions shown in Table 1. The rate cons]:ants of most of the elementary reactions are known and we have used the recommended rate constant published by Baulch, Drys. dale, and lloyd [24]. Tl'.e value of k9 is the value determined by Foner and Hudson [25] and more recently by Paukert and Johnston [26]. Table 1 also !ncludes th.~ preexponential factors, activation energies of these reactions, together with the calculated values of the rate constants at 546°C and atmospheric pressure. The value ofk~ has been calculated from ]rl ~,~e)'the composition of the mixture and the efiieiencias (e0 relative to he,Iinm of N20x = 2), O2(~! = 2), and H : ( a = 5) as third bodies [24]. The value of ka is obtained from k3 ( N ) ' the composition of the mixture and by assuming that the relatltve effic~encies are the ~ m e in reactions (1) and (3).

AKISSI A G K P O

52

and LOUIS-RENESOCHET

TABLE 1 Preexponential Factors (A), Aetiva:tionEnergies0E), and Rate Commnts in the Oxidation of a Mixture of 1 H2-1. O2-- 4 N2 at 5460C and a Pxessnteof 1 arm (ct,~eentxation in moles 1-I and time in S ) A H2+O 2

o ~ HO2+H"

1011

H.+02+ M

t ~ HO2+M

HO2 + H2

2 > H2024- H' 3 • ~ 2OH ,'M

i"6"109 2.5 • 1.6" 11)9 9.6" 11)9 1.17- 11)14 1-25 - 1.17 - 1014 2.2" 11)10 1 " 11)10 3.2" 1011 2.8- 11)'0

H202 + M OH'+H2 ~

H20+H"

OH'+ H202 ~

H' +H202 - - ~6 H02 +H20 ~ H202

Ho~

H20 ÷ I'I02

H~O+OH + H202 +OH

E (M = He)

(M = mixture) (M = N2) (M = mixture)

k

56a

1.2.10 -'I

-1

7.3- 109

24

4.1 • l0 s

45.5

102

5.1 1.8 9 32.7

9.3 • 10s 3.3. 109 1.2 • 109 6 • 101

16.8 9.4

1.8.109 7.7. I06 5.4 • I07

s-,,

HO~ + HO~ ~ H202 + O2 H ' + O 2 t-2~O OH+O O"+ H 2 ~ OH + H

2.2" 11)11 1,7" 10 l e

aE=AH=S6k~ It has been proposed that H202 formation takes place by bimolecular reaction of HO2 [19, 22, 7.31. This hypothesis has been directly vefLfi~,d [I0] by showing that the concer~tration of H0~ was suffieiently important, at least in a B20s coated vessel for rite rate of reaction (9) to be much greater ~ a n the rate of reactiom (2) and evidently (6'). 4.1 Kinetic Analysis of the Mecha~m Ignoring, therefore, reaction (6'), the equations describing the changes of active centers and miti',d and final products concentrations with time ~ a y be written and simplified by application of Semenov's principle of partial stationary concentrations: d(OH)/dt = d(O)[dt ~ d fH)/dt = d(HO2)/dt =: 0 and d(H202)]dt ~ O.

o" "= 2k~ (m) (H~%)

-n)

,~+(H~)

+ k6 (H202) + 2klo(O2) k4 (H2) (I-l). (2) If the ,'ate of initiation is neglected a relation between (H), (HO2), and (H202) is obtained from d(rt)lat = 0

(H)

1

(02) [k, @1) - 2k,o] [2ks (M) (H202) + k2 (H2) (H02)].

It can be shown that k2 (H2) (1-IO2) < < 2k~ (M) (H:O2) in our conditions and the. ",~re

{a] Estimationof the Relative Concentrari~,nof Active Centers From d(O)/dt = 0 we obtain

(o) = ~ n (rI2) (~.

p.,ession (2) if it is assumed, as it is justified in om conditions, that k, (H2) > > ks (H202).

(1)

Introducing (1) in d(OH)ldt = 0 lezds to the ex-

2k3(M) (H)= - - (t12o~) (02) [k, CM)- 2k, o]

(3)

By summation of equations d(X)/dt = 0 (where X = H, O, OH, HG2), and from Eq. (3) we obtain a

ESR STUDIESOF THE HYDROGEN-OXYGENSYSTEM simple relation between the two main active centers if surface destruction of He2 is neglected.

53

{b} The Maximum Concentrationof Hydrogen Perexide The rate of formation of hydrogen peroxide is given by the expression

a ( H z O 2 ) = k 2 ( H 2 ) ( H O 2 ) - ka(H202)(M) m

dt

/

1

~/2

(4) -

-

From these equations, knowing the experimentally determined maximum concentration of Hv02, (H3 O2)M, the maximum concentration of H e [ , (He2) M , may be estimated. The values calculated in this way are plott¢d on Table 2 and are in relatively good agreement ~ith those deter. mined by ESR. This implies that the value taken ;or ke which ts obtained from experiments at room temperature is correct and that therefore the activation ene,'gy E9 is close to zero. TABLE 2 Estimation of (He[) at the MaximumRate of Reaction from Eq. (4) and the Experimental Values of (H202)M. Comparison with Experiment~ Results. Type of coating

(H202)M exp. moles 1"~

(HO2)Mexp. moles 1-1

3-10-6

2"10-8

6"I0 -8

20"10-6

9"10-a

14"10-8

Without coating With B203 coating

k s ( H z O z ) ( O H ) + kg(HO2) 2 k ~ ( H 2 0 2 ) - k6(H2G,)(H).

If the values for (OH), (HOe), and (H) given by (1).-(4) are substituted in (5) we obtain a relation which ~how that when consumpt;on of reactants is negleett d, the concent:ation of hydrogen peroxide reaches a maximum (llzO2)M. From (I)-(4), similar results are also obtained for the other active centers and the overall rate of reaction. I f it is supposed that in a B2 03-treated reactor, the heterogeneous decomposition of H2 O: is ~egligible, the maximum concentrations of hydrogen peroxide are given by the following equation

(6). s12 + B(H2 ' ,92)M 3/2 _ C(H202)~2 = b A(H2Oz)M

(HO2)Mca]c. moles 1-1

(6) where,

2k3ksk0__ B = ~ j A=klk,~(H2)(Oa)'

The con.:aatretions of other radicals 131ative to (H2 O2)M = 1 must be calculated from the ,'elations (I)-(4) in a B20:-treated vessel. The results are shown in Table 3. The concentration of HO~ is thus very much greater than those of the otLer radicals and their possible contributions to the ESR .:. 'etra may be neglected. TABLE 3 Estimati.~n of the Relative Concentration of Active C~aters from Eqs. (1)-(4) and the Experimental Value (H202)M in a B203 Treated Vessel. X

H202

He2

(X)/(I'1202)

1

7"10-3

H

10~

0

OH

10 -6

10-.6

(5)

2k~k,o

-

k0 + k,(M)] Lkl(O2) ka(H2)J'

(k3(ra)~'/~ t -~/

This equation has been resolved by successive approximations and the results are shown on Table 4.

(c) TheMaximurn.RateofReactton The rate of reaction is equal, for instance, to the rate of removal of oxygen, and if we consider all the reactions (I)-(1 I) a general expression is obtained:

AKISSIA G K P O and LOUk3-REN ~" SOCHET TABLE4 Observedand CalculatedPa~amelersat the Maximum]~ateof Reaction for the Mixtme 1 il2-1 02 - 4 N2 at 546°C and AtmosphericPsessusein a BaOa-CoatedSilicaVessel

Experimentalresults with B203 coating Kineticanalysis Computertreatment Computertreatment k I = k ~ / 2 , k s =k3/$

~M

(HO2)M× 108

(H202)M X 106

S

moles I-II

moles l-I

~TJM

moles l"I sec-~"

× 106

30 19

9 14 12

29 iI 15

6 36 28

29

8

17

13

57

4

7

4

Computertreatment ~1o = ~l;ll =0 kt = klL/2, k3 = k3/3

_ d(02) = ~ (0=) (M) (S) -/c9 ~ 0 = )+" dt

+ klo(02) (I~. (7) By substRutingthe concentration of radicals by their values (1)-(4) this equation becomes d(O:O

. ~.. [kt (~) + 2k~o] ,= ,~,

(8)

From this equation, knowing the previously ealcuiated m.~rSrnumconcentrations of hydrogen peroxide (H20z )m, it is possible to calculate the maximwn rate of reaction ( - d ( O 2 ) / d O M as can be seen in Table 4. 4.2. Comp¢~erTfearmentof the Mechanismand Discussion The preceding kinetic analysis of the mechanisrn has been made on the assumption that the con sumption of reactants is negligible. Fwthennore, this treatra.=nt does not ~ve any iufo~maticn on ~ e tkae of reaction. By computer treatment it is possible to ob rain the variation of all products a~td radicals against thn~e. The computer treatment =rod kinetic analysis give similar results; the values givenby tht,• former are only slightly lower because of the cons~mption of reactants. With the ;recommended rate constants [24], it can be seen in Table 4 that the m~x/rnum rate of oxygen consumption is too great and the time

(~M) to obtain this maximum too short. Best values for these parameters am obtained if the values of ks and kt are diminished by a Cactor of 3 and 2. These two variations =:e within the limit of the suggerted error [24] for cal-uiated r-te constant (AE$ = -+ 2 keel, A/~'1 = e It keel). Nevertheless, it should be noted that the .,~,,,e effects can be obtained ff the measured remparature inside the vessel is higher than "theaverage temperature because of the exothermicity of the reaction which is not negligible [28] near the third limit. To test the importance of the direct ctu,inbranching reactions in the mechanism in ou, conditions (atmospheric pressure), rite branching reactions (10) and (11) have been neglected. We obsarve 'that the induction period is longer and both the rate of reaction and the conc~ntration of active centers lower. Conclusions tm apparatus has been developed which enables the simtdtaneous detection and estimation of rLiical products l:y ESR and molecular products by polarography or gas chromatography. In the oxidation of H2 (or D2) close to the tIfird limit we have obtained evidence of relativaiy large amounts o:f HO~ (,or DO~). It seems ,.hat H2 02 is formed raainly by bimoleeular recombination of HO~, Evidence is also presented concerning the role of

surface coating. Using valuesof the rate constantsfrom the literature, the induction period and the concenh~-

ESR STUDIES OF THE HYDROGEN-OXYGENSYSTEM tio:~t o f active centers have been calculated from the generally used mechanism which involves both direct chain branching (reactions 10-1 I ) and indirect chain branching (reaction 3) with interaction of reaction chains. I f it is taken into account the suggested errors [24] about the rate constants and the exothermicity of the reaction, these values are in relatively good agreement with experiments. Referer, ees I. Bennett, J. E., Mile, B., and Summers, R , Nature 225, 932 (1970). 2. S~moiiov, I. B., Ryabiokov, O. B., Zaichikov, V. V., C-e~henzon, Yu. M., and Gus~ak, L. A., Dokl. Akad. Nauk. 205,1138 (1972). 3. Nalbandtan, A. B., Vest. Akad. Nauk. 11, 46 (1969) 4. CHilies,M., and Soehet, L. It., C.R. Acad. Sci. Paris C271,114 (1970); Soehet, L. R., The kinetics o f chain reactions, Dunod, Paris, 1971, p. 17; Cariics~ M, and Sochet, L. IL,Z Chim. Phys. 71 (1974). 5o S",ehet~L. R.,J. Chim. Phys, 70,456 (1973). 6. Cariier, M., Th~se 3~me cycle, Lille, 1972. 7. Agkpo, A., Th*3se3~me cycle, Lille, 1973. 8. Vatdanian, I. A., Sachianl G. A.1 and Nalbandian, A. B., Comb. Flame. 17, 315 (1971). 9. Agk'po, A., and Soehot, L. R., Combustion Institute. European Symposium 1973, (F. J. Weinberg, Ed.), Academic Press, London, 1973, p. 65. 10. Agkpo, A., and Sochet, L. iL, C.R. A cad. ScL Paris C276, 631 (197: , 11. Bermett, .I.E., Mille,B., and Tt,omas, A, 11th Syrup. International on Combustion, The Combustion Institute, Pittsb,rzh. Pa., !9~7, p. 853. 12. Adrian, F. L, Coehran~ ~. L., and Bowers, V. A., J. Chem. Phys. 57, 5441 (1967). !3. Wyaxd, S. J., Smith I R. C., and Adrian, F. 3, .L Ch~n~Phys. 49, 2780 (1968).

55 14. Sachian, {3. A., Shakhnazarian, L K., and N:.flbandian, A. B.,Arm. Chlm. J. 22, 371 (1969). 15. Shakhn~.,arian, 1. IC, Danghian, T. M., Saehia,, '.;. A., and Nalbaodian, A. B.,Arm. Chirn. J. 25,543 (1972). 16. Hoare, D. E., Peacock, G. B., and RtLxton, G. R. D., Trans. Farad. Soe. 63, 2498 (1967). 17. Semanov, N. N.,Z. Fiz. Khim. 17, 187 (1943). 18. Conze, A. T., Gaillatd-Casin, F., and James, H., Bull. SOC.Cbim. 2367 (1968). 19. Lalo-Koutilsky, C., and .lames, H., Bull. 5oc. Chim. 1775 (1967). 20. Ben Aim Biquatd, J., and Diamy, A. M, Bull. Soc Chin~ 4114 (1972). 21. Barnard, J. A., arid Platts, A. G., Comb. ScL Tech. 6, 133 (1972). 22. Baldwin, R. R., and Mayor, L., 7th Symposium Interrational on Combustion, The Combustion Institute, Pittsburgh, Pa., 1959, p. 8. 23. Baldwin, R. IL, Doran, P., and Mayor, L , Sth Symposium lnfcrnational on Combustion, The Combustion Institute, Pittsburgh, Pa., 1962, p. 103; Baldwin, IL IL, Jackson, D., Valker 11. W., and Webster, S. J., Trans. FaradaySoc. 63,166~ (1967); 63,1676 (1967). 24. Baulch, D. L , Dlysdale, D. D., and Lloyd, A. C., High tempemtnre reaction rate data llniver.,;ityof Leeds 1968,1969. 25. Fonet, S. N., and Hudson, R. L.,, ~dv. Che~ 36, 42 t1962). 26. f'ankert, T. T., and Johnston, H. $.,£ Chem. Phys. 56, 2824 (1972). 27. Shakhna2atian, L K., Danghian, T. M., Saehian, G. A., and Nalbnndlan, A. B.,Arm. Chim. J. 26,182 (1973). 28. FOO,K. K., and Yang, C. H., Comb. Flame. 171223 (1971).

Received Lecember 3, 1973; revised March 2.2, l 9 74