PHOTOLYSIS OF ACETONE - HYDROGEN CHLORIDE MIXTURES1

ABSTRACT The photolysis of acetone - hydrogen chloride mixtiires has been investigated

a t 150" C. and a t room temperature. A strong suppression of ethane forl~latiorl with a corresponding large increase in the formation of ~ilethane results from additions of relatively very small amotrnts of hydrogen chloride to acetone. The importance of the reactions

(1) CHI + HCI --+ CHI + C1 and

(7) CI + CHICOCHI -+ HCI + CH?COCH, has been demonstrated. The collisiorl yield of reaction (1) a t 28O C. is 2 X 10-1, and, therefore, the upper lirrlit for EI is 5 . 1 kcal. per mole. The effects observed a t 150" and 28" indicate that, on the assumption of a zero activation energy and a steric factor of unity for combination of rrlethyl radicals, El = 2 .1 & 1 lical. per mole, and PI is approximately 7 X

IN'TRODUCTION

Methyl radicals produced in the pl~otolysis of acetone have been widely used in the investigation of hydrogen abstraction reactions. Recently these investigations have been extended to halide substituted hydrocarbons. Raal and Steacie (12) have determined the activation energies and the steric factors of the reactions of methyl radicals with a number of halogenated methanes in the temperature range 110 to 220' C. The following mechanism of methane and ethane formation is proposed:

(1) CH3 + RII -+ CH4 + R, (2 ) CH3 + CH3 --+ CzHs, (3 ) CH3 + CH3COCH3 --+ CH4 + CH,COCH3.

For experimental conditions where the concentrations of acetone and RH can be regarded constant, and assuming a time independent concentration of nlethyl radicals, it can be shown that

I l l

Experimental measurenlents of the rates of formation of methane (RCH,) and ethane ( R y s ) in the absence and presence of RH are used to evaluate k3/k2h and k l / kg respectively. Arrhenius plots of such determinations over a temperature range give the values of (E3 - Ez) and (El - + Ez). Since E2 is approximately zero, these values are to a good approximation identical

the activation energies of the respective elementary reactions (I) and (3) . For the series of reactions of nlethyl radicals with CH3C1, CH2Cl2, and

CHCI3, Raal and Steacie thus find a substantial increase in the ease of removal of hydrogen as the number of chloriile atoms in the rllolecule increases. This

1 Manuscript received October 64, 1952. Contribution from the Division of Chemistry, National Research Cor~ncil, Ottawa, Canada.

Issued as N.R.C. No. 2899. 2 National Research Cozr~lcil of Canada Postdoctorate Fellow, 1951-52.

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result is not unexpected in view of the other observations on the reactivity of these compounds. The authors have, however, also observed a very pro- nounced effect of carbon tetrachloride on the rates of formation of metharle and ethane in the photolysis of acetone. The effect is analogous to that p1-o- duced by chloroform. Relatively small amounts of carbon tetrachloride added to acetone suppress very strongly formation of ethane and the formation of methane is correspondingly increased. While in the case of partially halo- genated methanes the increased methane formation is easily explained by hydrogen abstraction, the corresponding- effect due to additions of carboil tetrachloride is entirely unexpected. The apparent anomaly introduces, there- fore, an element of uncertainty into the above results. The authors' suggestion that the observed anomaly may be due to sinall amounts of hydrogen chloride, present as an impurity or formed during the photolysis, makes it necessar-y to obtain some quantitative information on the effect produced by additions of varying ainounts of hydrogen chloride to acetone.

In a recent paper A. S. Kenyon (9) reports formation of hydrogen chloride as a product in the photolysis of acetone in the presence of sec-butyl chloride. At the same time ethane formation is completely suppressed even a t 25". The author ascribes the latter fact to a'rapid reaction of methyl radicals with sec-butyl chloride. The possibility of an effect of the presence of hydrogen chloride is not considered.

Williams and Ogg (17) have investigated the rate of reaction of ~nethyl radicals, produced by photolysis of methyl iodide, with hydrogen chloride. They find the activation energy of this reaction only about 1 . 6 ltcal. per inole greater than that of the rapidly occurring reaction with hydrogen iodide, although they suggest that in the former case a reaction of "hot" n~ethlrl radicals is of importance.

Both in the older and the recent literature a somewhat controversial attitude with respect to the importance of the reaction between irlethyl radicals and hydrogen chloride is found. At the same time it can be conceived that in the photolysis of acetone in the presence of chlorine containing compounds hydrogen chloride might be formed in the process or could be introduced in varying amounts as an impurity. I t was, therefore, considered of interest to investigate separately its effect. A largely qualitative preliminary experiment, conducted in collaboration with Dr. F. A. Raal, showed that the addition of a relatively small amount of hydrogen chloride to acetone photolyzed a t 150" suppressed ethane formation alnlost coinpletely and the arnount of methane produced was much increased. 4 quantitative investigation of this effect necessitated some arrangement for the iiltroduction of known very small quantities of hydrogen chloride into the reaction system. Such an investigation, conducted a t 150" and at room temperature, is reported in this paper.

EXPERIMEN'l'AL A pparatz~s

The apparatus was of the same general design as used previously by Raal and Steacie (12). Alterations were made: ( I ) to allo\\r introduction of knou-n

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small cluantities of ll~.drogen chloride and to ensure thorough mixing with acetone; (2) to eliminate mercury vapor as niuch as possible from the reaction cell. The former was achieved by a system of calibrated volumes and by ulti- mately freezing over the two reactants into a 300 ml. calibrated bulb and, on expansion, allowing sufficient time for iilterdiffusion; the mixed reagents were subsequently expanded into the reaction cell. The latter was accomplished bj. measuring the pressures of vapors of the reactants contained in large accessory volumes, which were filled sim~~ltaneously with the reaction system. The accessory volulnes were subsequently isolated from the reaction system and opened to mercury manometers. A correction for i-he slight increase in volume (amounting to 1 to 27,) was applied to the measured pressure. Re- peated calibration and testing of this arrangement ensured satisfactory accuracy.

The light from a Hanovia S-500 medium pressure arc was roughly collimated bjr a highly polished aluminum cylinder. For inost deterininations the light beam was passed through a Corning 9-54 filter, and in a few determinations a Corning 0-53 filter was used instead. In the former case the light beam con- tained the wave lengths greater than about 2200 A, in the latter greater than about 2800 A. The light beam filled the cylindrical silica reaction cell 120 ~ n l . in volun~e.

'The reaction cell was contained in an aluminum blocli furnace. The heating current was supplied from a Sorensen voltage regulator and the temperature \\-as maintained to within about 1".

The products were analyzed for methane, ethane, and carbon nionoxide. The analysis system was of the conventional type. Carbon monoxide was determined b\r conlbustion over copper oxide a t 215'. In a number of cases simultaneous determillations were made by inass spectrometer an a 1 y s ~ s - ' as a check. Ethane was separated a t -1GS0. At that temperature no hydrogen chloride was recovered from acetone, as was evident from determinations ~vllei-e large amounts of hydrogen chloride were added to acetone and tlle ethane fraction \\-as negligibl?. small.

Acetone was a Mallinclcrodt Analytical Reagent product, refluxed over potassium permanganate, dried over Drierite, and distilled. I t %\-as degassed in the usual manner by bulb to bulb distillation in vacuo. Hydrogen chloride \vas a Matheson Compan\- product. I t was degassed and further p~lrified by several bulb to bulb distillations i n vucuo (drj. ice - acetone to liquid air). The deuterated acetone (87% de) was made by Dr. I,. C. I2eitcll in this labora- tor!.

'I'hc photol~.sis of acetoile in the presence of hydrogen chloride was investi- gated a t 150" and a t room temperature. Most determinations were carried out with a light beam of wave lengths greater than about 2200 A and only a few determinations were made with wave lengths greater than about 2800 A. In all cases addition of hydrogen cllloride led to an increase in the a ~ n o u ~ l t of

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CI'E1':INOI~ZC .-!iYD ST'EACZE: .4 CETONE - HYDROGEN CIZLORZDE PHOTOLI'SZS ] (jl

methane formed with a corresponding decrease in ethane formation. Fig. 1 shows the quantum yields of methane and ethane a t 150' and X > 2200 A for -varying initial h?-drogen chloride concentrations, the other experimental factors being kept approximately constant.

Relatively small amounts of hydrogen chloride (less than 1 mm. added to 100 mm. of acetone) were sufficient to suppress ethane forination alrnost con~pletely; the formation of methane was correspondingly increased and approached a t 150' a value close to twice the amount of carbon morloxide produced. The yield of carbon monoxide was not- or possibly was only slightly affected.

I t was found convenient to investigate the influence of different experi- mental conditions by following trends in the values of k1/k2i calculated from Equation [I] after substituting in it [HClIi for [RH]. I t is, therefore, implicitly assumed that hydrogen chloride concentration does not vary with time and is equal to the concentration of the initially introduced hydrogerl chloride (HCI,). (In the follo\ving the symbol (kl/k2+)l will denote the values obtained to this approxi~nation.) The reason for this choice may be evident from the follonring considerations. The amounts of hydrogen chloride added to acetone in these experiments are much too small to account for the increased methane production if it is to be assumed that the additional methane is formed by reaction ( I ) , i.e., hydrogen abstraction from the originally introduced hydrogen chloride. A t the same time, as will be seen from the subsequent discussioi~, there is evidence that reaction (1) is iilvolved in the process. I t seems, there- fore, to be necessar!. to assuIne, as was done previously by Williains and

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I (i2 CANADIAlV JOUlWAL OF CHEMISTRY. VOL. 31

Ogg (17), that hydrogen chloride is regenerated, in this case by reactioll of chlorine atoms with acetone. If the hydrogen chloride regeneration is ~0111- plete, and since in all experi~nents large concentrations of acetone are used and the reactions are carried out to small conversions, the concentratio~ls of l~ydrogen chloride and acetone remain constant. Equation [I] can then be used for evaluation of k l / k + . If the hydrogen chloride regeneration is not complete, the effective mean hydrogen chloride concentratioil will be smaller than the initial hydrogen chloride concentration and the calculated values of k l / k 2 t will be too small. A criterion of the completeness of the hydrogen chloride regeneration is, therefore, constancy of ( k l / k ? a ) , for different experi- mental conditions.

Typical results of the experiments carried out are presented in Tables I to IV. An attempt was made to investigate trends in the values of ( k l / k $ ) l by varying one experimental factor a t a time. This, however, was achieved only approximatell-. For this reason a graphical representation of the results would show a scattering of points not necessarily representative of the actual experin~ental error. As a result of the necessity of keeping all experimental factors constant but one, in some of the experiments the amounts of the products are very small and may involve a considerable analytical error. They are believed, however, to indicate correct l~~ the trends in the values and have been, therefore, included in the tables.

Table I shows drifts in the values of ( k l / k z t ) l with increasing initial hydrogen chloride conceritrations for the photolysis of acetone a t 150'. Acetone pressure was kept a t approximately 100 mm. The first five determinations were carried

Filter 9-5L

TABLE I P N O T O L Y S ~ S OF ACETONE-HCI M I X T U ~ E S A T 150' C .

Effect of HCI concentrat~on

Filter 0-53 I I I I

Time, Run Temp., C .

[HClli , molec./cc.

X 10-I'

In molec./cc. sec. x 10-l3

--- --- - Rca I RCH. 1 Rc,H.

[Acl, molec./cc.

X l O - ~ b e c .

CzHc f i C H I CO

( k I / K P : ) L X 10"

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Cl~r 'TANOI ' IC AND STEACIE: ACETONE - HI'DROGEN CHLORIDE PHOTOI.YSIS 163

out under conditions of free access of mercury vapor into the reaction cell from the adjoining mercury manonleter. In the case of all the remaining determinations reported mercury vapor was removed from the reaction s)-stem. There was no significant difference i l l the values obtained in the two cases. With increasing hydrogen cl-lloride concentration the values of ( ~ ~ / k $ ) ~ be- come progressively greater. The material balance for the conversion of meth\-I radicals in to methane and ethane, as expressed by the ratio (GI-16 + 3 CH,)/CO, remains similar to that of acetone alone. TVith hydrogen chloride concen- trations considerably larger than those reported in Table 1 ethane was not detected by the analytical procedure used and the amount of methane formed approached a value about twice that of carbon monoxide without exceeding it. The three determinations with X > 2800 A show a similar trend.

Tablc I1 shows the analogous experiments a t room temperature. I t is

TABLE I1

PHOTOLYSIS OF ACETONE-HCI MIXTURES A T ROO11 1.IShlPEKATURE E f f e c t o f HCI concentration

I I I I 1 In rnoGc./cc. sec. I

evident that even a t room temperature the effect of small hydrogen chloride concentrations is very pronounced. The general trend in the values of (kl/k2f)l is similar to that a t 150°.

The rates of formation of ethane and methane vary with light intensity, as shown in Table 111. I t was found that in the photolysis of acetone -hydrogen chloride mixtures a t 130° the quantum yield of carbon monoxide formation was close to unity, as in the case of acetone alone. The data on the light absorption in Table 111 are calculated on the assumption that the quantum yield of carbon monoxide formation is unity. The trend in the ratio of RCH,/RCzHG is qualitatively the same as in the case of acetone alone and the values of (kl/kz':;)l become greater as the amount of light absorbed is reduced. The light intensity in these experiments varied as a result of (u) progressive decline in the output of the light source and (b) regulation of the current through the mercury arc. 'The effects in the two cases were approximately the same.

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CANADIAN J0UR;V:IL O F CIIEJIISTRY. 1.01,. .;I

PH~~~oI . \ - s IS OF :\cI~ToKI':-HCI h1lXTURES (150' C., FILTER 9-54) Effect of light intensity

I I 1 I I I I

Run Light In nlolec./cc. sec.

Temp., [HCl];, [Ac], absorbed, " C. molec./cc. molec./cc. q~ranta/sec. x lo9

X L O - 1 6 ~ 1 0 - ~ 9 1 0 - ~ I Rco R C ~ . R C , H .

The lime of phololysis i n all runs was 6400 sec. T h f u7nount of lifiht absorbed w zs cnlczrlaled 071

Ikc assun~phion lhnt the quiznl7~nz yield of carbon ,n!~noxida .forrnnliot~ was z~ni ty .

The effect of varying time of photol>.sis is shown in Table IV. There is a distinct decrease in the ratio R C I - I , / R ~ 2 ~ G with increased time of photolysis and the values of (kl/k2f)l become smaller.

PHO.I.OLYSIS OP ACETONE-HCI MIXTURES (FII.TISK 9-51) Effect of the time of irradiation

I I I I I I I

., Time, [HCl];, sec. molec./cc. molcc./cc. 1 * ~ X 10-16 1 L::'18 1

DISCUSSIOK

The expe;imental results on the photol>-sis of acetone - hydrogen chloride mixtures can be summarized in the following manner: ( a ) there is a very pro- nounced effect on addition of relatively small quantities of hydrogen chloride both a t 150" and a t room temperature: ethane formation is strongly sup- pressed and the formation of methane is correspondingly increased; (b) the effect is approxilnately the same a t X > 2200 A and X > 2800 A; (c) the yield of carbon lnonoxide and the material balance for the conversion of methyl radicals into methane and ethane are approximatel-). the same as in the case oE acetone alone; (d ) the products of the photolysis contain no detectable quantities of h? drogen. There is also a decrease in the values of (kl/k$)l with (e) time of photolysis, ( I ) decreasing hydrogen chloride concentration, and (g) increasing light absorption. At the same time, (IL) no significant difference in (kl/k24)1 was observed when acetone pressure was reduced from 100 to 30 mm. (This experiment was performed a t 150" and X > 2200 A.)

I t can be concluded fro111 (e) that hydrogen chloride is a t least partially

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CVEI'ANOI'IC d;VD STEACIE: ACETOIVE - HYDROGEN CHLORIDE PIIOTOLYSIS 165

consunled in the reaction. At the same t i~ne , i t is evident from the material balance for ~nethane formation that much less than one molecule of hydrogen chloride is used up whenever an additional ~nolecule of methane is formed. This implies either that for the most part hydrogen chloride participates in the reaction only as a second body and is not itself chemically changed, or, if it does undergo a change, that it must be rapidly regenerated. I t is difficult to conceive of a plausible mechanism for the first alternative co~npatible with all the experimental observations.

The energetically possible que~lching reaction ( la ) CH3COCH3* + HC1+ CH4 + CH3C0 + C1

would have to be followed b>. CH3C0 + CHB + CO. Evidently, it cannot explain complete suppression of ethane formation unless it is assu~ned that either of the reactions

(1) CH3 + HC1 +CH, + C1 or ( I b ) CH3CO + HCI + CH4 + CO + C1 also takes place and is very fast. I;urtherrnore, 110 kinetic evidence can be derived from the experimental results that reaction ( la ) plays an important role. I t may, therefore, in the present treatment be disregarded. Since a t 150" CH3C0 decomposes rapidly into CH3 and CO, reaction (1) must be of importance. The following schenze of reactions would then seem to be sufficient for a satisfactory explanation of the experimental results:

(I) CH3COCH3 + h~ + CH3 + CH3CO (11) CH3CO +CH3 + CO

(111) CHXCO + CHSCO + (CHSCO)?

(1) CH3 + HCl + CH4 + C1 (2) CH3 + cI33 + C2H 6

(3) CHa + CH3COCHB + CH4 + CH2COCH3 ( 4 CH3 + CHZCOCH 3 -+ CrH,COCH3 ( 5 ) 2 CH3COCH3 + (CH,COCH,), ( 6 ) CHTCOCHj + CH 3 + CHZCO

(7) C1 + CH3COCIHj + HCl + CH2COCH3

(8) { c"l" +S + (HC1.S) + Y + (C1.Y)

where (8) is some process by which part of the hydrogen chloride is consumed. Reaction (6) is probably not important below 200". -1 proof that hq,drogen is abstracted from h\,dl-ogen chloride to form methane

might be furnished by photol>.sis of deuterated acetone in the presence of relatively large quantities of ll\di-ogen chloride when the process is carried to a small conversion. M'e thus find that in the photolysis of 100 mm. of deuterated acetone (87% ds) and 30 rnm. of hydrogen chloride the methane produced consists predon~inantly of CD3H, both a t 150" and a t room tempera- ture. An eventual sponta~leous exchange of H and D between acetone, and hydrogen chloride, even in the case of the most u~lfavorable redistribution, cannot explain the predominant formation of CD3H.

The possibility of an appreciable direct photolysis of h>,drogt~z chloride

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166 CANADIAN JOURNAL OF CIIEMISTRY. I.OL. 31

under conditions of the present experiment seems to be excluded. The lollg wave length limit of the hydrogen chloride absorption continuum is about 2200 A (7; 10; 16) and with the filters and the small hydrogen chloride concen- trations used no appreciable absorption by hydrogen chloride would be expected. This is substantiated by the lack of hydrogen in the reaction products and the similarity of the observed effects with alternate use of filters 9-54 and 0-53.

The mechanism proposed for the photolysis assumes a partial consun~ption of hydrogen chloride. A number of chemical processes could explain the de- crease in the effective hydrogen chloride concentration, although purely physical phenomena may be also involved to some extent. .An assunlption that for higher hydrogen chloride concentrations [HCI] = Const. = [HCl]i should, however, be a good approximation, regardless of the nature of these processes.

An approximate treatment of the mechanism of hydrogen chloride con- sumption can be attempted by assuming that hydrogen chloride concentration is diminished during the process predominantly through a partial removal of free chlorine atoms. A direct abstraction of a chlorine atom from hydrogen chloride is highly unlikely. Results (f) and (g), stated in the summary a t the beginning of this section, represent a decrease i l l the values oi ( ~ l / k $ ) ~ with an increase in the methyl radical concentration. For simplicit!-, therefore, only the following possibilities will be considered:

(8a) CH3 + Cl -+ CH3Cl (8b) CHXCO + Cl -+ CH3COCI

CO + Cl -+COCl (8c) / COCl - - CO + Cl

\ CH3 + COCI -+ CH3COCI . Reactions (8c) are suggested as a possibility on the basis of the investigations

~ f ' ~ h o t o l y s i s and thermal decoillposition of phosgene ( 3 ) . If (8c) plays an important part, additions of relatively small quantities of carbon monoxide should inhibit the reaction very strongly. The analytical accuracy in such experiments is solnewhat decreased through the presence of additional quanti- ties of carbon monoxide. Nevertheless, it is evident from Table V that (8c) is

'TABLE V

PHOTOLYSIS 01; ACI~T~N~C-HCI MlXTUKIC5 (FILTER 9-34) Effect of carbon monoxide concentration

I I I I I I I

74 28.3 7200 0.384 0.44 3.20 0.309 1 0.190 0.55 27.7 7200 0.136 0 3 8 6 0.139 0.73 :B 1 7 . 1 7200 1 6:ut 1 8::: 1 :::: 1 o a 4 1 o o o r 1 1.3

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CVETANOI'IC d M D STEACIE: ACETONE - HYDROGEN CHLORIDE PHOI'OLYSIS 167

of little or no importance both a t 150° and 25O C. The amounts of carbon monoxide added are in excess of the quantities norn,ally produced in the photolysis under conditions where (k~/k$)l shows a n~nrlced decrease.

An exact mathematical treatment of the proposed mechanism, taking into account reactions (8a) and (8b), is a t present not possible. Chlorine atom and hydrogen chloride concentrations are mutually related in a very complex way. An assumption of a steady state concentration of CH3 and Cl implies that there is a mean hydrogen chloride concentration which is srnaller than IHCli and remains constant during the photolysis. I t can be shown that the value of kI/k$ to this new approxi~nation (denoted by (kI/k2+)2) is'

Rcrr, k a 1 1 2 - TZ [Acl

k2 [HCI], - 6

where

1 + Rc H Rco k' --=A + k" -- [Acl [Acl

and

t is the time of photolysis. Equation [2] is in qualitative agreement with the experimental observations.

At l5Oo kII is very large and as a further simplification k" can be neglected. A sinlultaneous plot of quantum yields of methane and ethane for~nation a t 150° and X > 2200 A for varying initial hydrogen chloride concentrations, other experimental factors being approxinlately constant, is shown in Fig. I . Acetone pressure in all cases was close to 100 mm. By plotting also the re- maining experimental data a t 150° and X > 2200 A it is found that all the points either fall on the pair of curves (quantum yields of methane andethane respectively) or can be brought thereon through a sinlple shift in hydrogen chloride concentration. Thus there is an indication that kI/k2* = f([IHCl]i - A). Equation [2] is a very sinlplified expression of this functional dependence and an accurate agreement with the experimental results cannot be expected. By taking k' = 4.2 X 101° and k" = 0 the drifts observed a t 150° and X > 2200 A largely disappear and the mean value of (kI/k2*)2 is about 4 .5 X lop9. However, there remains a large scatter of the values. Also, this treatment fails with the results obtained a t room temperature, though the values of kl/k2f can be mutually approached.

Drifts in the values of (~ , /kI? i )~ make an accurate evaluation of the activation energy of reaction (I) impossible. On the basis of the proposed ~nechanis~n extrapolations to t = 0, l a = 0, or to large hydrogen chloride concentrations should supply the true values of kI/k$. Limitations in the accuracy of the analytical procedure make these extrapolations unreliable. I t seems possible,

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however, to obtain an approximate estimate of the activation energy of the reaction from the data in Tables I and 11. With increasing hydrogen chloride concentrations (kl/kzt)l tends to a value of about 4 . 5 X lop9 a t 150" and 1 . 3 X a t 28" C. From the ratio of the two values, and taking into account the ~4 factor in the pre-exponential term, 2.1 kcal. per mole is obtained for El - + Ez. Assuming mean collision diameters of 3 . 5 A and 1 . 3 A for CH3 and hydrogen chloride respectively, P~/P$ is found to be 6 . 7 X With unity as a plausible value for Pz (6), the value of PI becoines 6 . 7 X In view of the uncertainty,of the extrapolation to high hydrogen chloride concen- trations, an uncertainty of + 1 Ircal. per mole in the above value of El - + Ez would seem to be a sufficiently conservative estimate. This introduces a factor of 3 . 3 as the correspondiilg uncertainty in the value of P~/P?.

The present estimate of the activatioil energy of the reaction (1) is based on the values of rate constants a t 150" and 28" C. In the photolysis of acetone alone two distinct values of the activation energy of hydrogen abstractioil by methyl radicals have been reported: 9 . 7 kcal. per mole bet\veen approxin~ately 120" and 300" (15) and 6 . I kcal. per mole between 120" and 25" (5). There is still no complete explanation of this apparent anomaly. This discrepailcy, however, should not affect the estimated value of the activation energy of reaction (1) because in the estrapolztion of (kl/klq)l to large hydrogen chloride concentrations the coiltribution of the reaction (3) becon~es negligibly small. At the same time, Equation [:I] should be valid regardless of the discrepancy in the value of E3 - $ E2 provided methane and ethane are formed by a free radical mechanism. Dorfman and Noyes (5) have shown that the experimental data on photolysis of acetone are fully coinpatible \\lit11 the free radical mechanism both a t 25" and a t higher temperatures. Bellson and Falterman (1) have recently investigated the photolysis of ecl~~in~oleculai- mixtures of ordinary and deuterated acetone and found that both a t 35" and a t higher temperatures C ~ D G , CD3CH3, and C2H6 are formed in the ratio of 1: 2: 1, as would be ex- pected in the case of a free radical mechanism. In analogous experiments we find tha t both a t 25" and 150" several times as much CD3I-I is formed as CD4. I t would seem, therefore, that a t rooin temperature as well as a t higher temperatures metharle and ethane are formed by a free radical mechanism. Nevertheless, it is possible that an eventual explanation of the curvature of the Arrhenius plot for acetone photolysis a t lower temperatures ma)- necessitate

. a revision of the conclusions suggested here. The value of the steric factor obtained for reaction ( I ) seems to be in

agreement with the values previously observed for similar t).pes of reactions. I t is also of interest that froin the experimental results of Williams and Ogg (17) a value of 2 . 4 kcal. per mole is obtained for El if it is assumed that the reaction of ~nethyl radicals with inolecular iodine requires no activatioil energy. These authors, however, suggest that under the conditions of their experiments (large hydrogen chloride concentrations) a reaction of "hot" methyl radicals with hydrogen chloride is of importance. On the other hand, in the experiments reported in the present paper, the hydrogen chloride concentrations used were very small and, consequently, any "hot" methyl

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CI'ETANOI~IC dXr) SI'EAC'IE; .ACETONE - HYDROGEN CHLOILIIIE PHOTOLYSIS l(j9

radicals were thermalized by collisions with acetoile before reacting with hydrogen chloride, so that a 1-eaction of "hot" methyl radicals is unlikely to be inlportant.

The lour value of the activation energy obtained for reaction (1) may appear a t first sight sorne\vhat surprising in view of the value of more than 9 kcal. for the reaction of methyl radicals with hydrogen (11). A consideration of bond strengths shows that reaction (1) is nearly thermoneutral or only slight])- endothermic. El would then be expected to be comparable \vith the activation energy of the reverse reaction

(9) C1 + CH4 + HC1 + CH,. For the comparable set of reactions (which are also approximately therrno- neutral)

(10) H + HCI +Hz + C1 and (11) Cl + H, + I-ICl + H Bodenstein (2) estimates collision yields of the same order of magnitude, lou4. Tamura (14), using the method of thermal analysis of reaction velocity (8), has investigated reaction (I)) a t 19' C. and obtained a collision yield of 2 .4 X and, consequently, Eg < 6 . 2 kcal. per mole. The maximum value of 6 .2 kcal. per mole corresponds to a steric factor of unity; for a steric factor of lop3 E9 reduces to 2 .1 kcal. per mole. 'Talting EZ = 0, Pz = 1, and mean molecular diameters of 3 .5 A and 1 . 3 A for CHI and 11)-drogen chloride respectively, we obtain for reaction ( I ) a t 28' a collision yield of 2 . 0 X 10-4, which leads to a value of El < 5.1 lccal. per mole. 'l'hei-e is, therefore, an approximate agreement between the collision yields of reactions (1) and (I)). The consistency of the value obtained for El and Tainura's value for Eg is thus seen to be uncertain to the extent of the uncertainty of the steric factor of the latter reaction. Unfort~inatelj-, most reactions involving chlorii~e atoms have been studied a t a single temperature. Rodebush and I<lingell~oefer (13) have investigated reaction (11) a t 25' and 0' C. Chlorine a to~ns were pro- duced by electrodeless discharge and their concentration ~neasured by a gauge similar to the Wrede gauge. From the experimental measurements reported and the uncertainty suggested by the authors, i t would seem that PII is between lo-? and unity. However, in view of the relatively small temperature range employed, we do not feel justilied in drawing- any definite coi~ciusions relevant to the present work.

Reaction between methyl radicals and hj-drogen chloride is seen to be fast enough so that the presence of even relatively small quantities of hydrogen chloride may have a pronounced effect on the over-all course of a reaction involving methyl radicals. Actualbr, k? is probably about ten thousand times as great as k1 a t rooin temperature and the pronounced catalytic effect of hydrogen chloride is only due to the exti-emely small CH3 concentrations and to the rapid regeneration of hydrogen chloride. The action of hydrogen chloride may manifest itself in most cases solely in an increased production of methane and almost complete suppression of ethane without affecting in any other way the over-all composition of the products. Under such con- ditions the CH3 concentration is so milch suppressed that on1)- negligible

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170 CANADIAN J O U R N A L OF CHEhfISTRY. V O L . 31

quantities of CH3C1 result from reaction (Sa) and lnay not be analytically detectable. Thus, in the photolysis of mixtures of acetone and chlorinated h~-drocarbons, hydrogen chloride rnay be a product of the reaction or could be present as an impurity. The strong suppression of ethane formation could be in such cases predominantly due to a very fast abstraction of hydrogen from the hydrocarbon. At the same tiine; to a smaller or greater extent, the effect may be due to a fast abstractio~l of hydrogen via hydrogen chloride. The relative importance of the two possibilities can be ascertained from quanti- tative considerations when hydrogen chloride coilcentration is known, or by a suitable selective inhibition of the hydrogen chloride.effect (4). The signifi- cance of the catalytic effect of hydrogen chloride in the apparent carbon tetrachloride anomaly mentioned in the introduction, and in some other instances, will be discussed elsewhere in this Journal (4) on the basis of addi- tional experiments performed.

ACKNOWLEDGMENT

\Ire are indebted to Dr. F. P. Lossing and Miss Frances Gauthier for the mass spectrometer analyses.

REFERENCES

1. BENSON, S. W. and FALTERMAN, C. W. J . Chern. Phys. 20: 201. 1952. 2. BODENSTEIN, M . Trans. Faraday Soc. 27: 413. 1931. 3. BODENSTEIN, M., BRENSCHEDE, \V., and SCHUMACHER, H. J . Z. physik. Chenl. B,

40: 121. 1938. 4. CVETANOVIC, R. J . , RAAL, F. A., and STEACIE, E. W. R. Can. J . Chem. 31: 171. 1953. 5. DORFMAN, L. M. and NOSES, W. A., JR. J . Chem. Phys. 16: 557. 1948. G. DURHAM, R. \Y. and STEACIE, E. W. R. J . Chem. Phys. 20: 5S2. 1952. 5. HERZBERG, G. Molecular spectra and ~nolecular structure. I. Spectra of diatomic

molecules. 2nd ed. D. Van Nostrand Company, Inc., New York. 1950. p. 534. S. ICHIKAWA, T . Z. physik. Chem. B, 10: 299. 1930. 9. KENYON, A. S. J . Am. Chem. Soc. 74: 3372. 1952.

10. LEIFSON, S. W. Astrophys. J . 63: 73. 1926. 11. MAJURS, T. G. and STEACIE, E. W. R. J . Chem. Phys. 20: 197. 1952. 12. RAAL, F . A. and STEACIE, E. W. R. J . Chem. Phys. 20: 578. 1952. 13. RODEBUSH. W. H. and KLINGELHOEFER, W. C., JR. J . Am. Chem. Soc. 55: 130. 1933. 14. TAMURA, M. Rev. Phys. Chem. Japan, 15: 86. 1941. 15. TKOTMAN-DICKENSON, A. F. and STEACIE, E. W. R. J . Chem. Phys. 18: 1097. 1950. 16. VODAR, B. J . phys. radium, 9: 166. 1948. 17. WILLIAMS, R. R., JR. and OGG, R . A,, JR. J . Chem. Phys. 15: 696. 1947.

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