H. A. SWAIN, JR., L. S . SILBERT,
2562
radical concentration. If then CH4 comes largely from decomposition of n-propyl, and chain hydrogen from the decomposition of isopropyl, both the very low values of kq and the normal values of k5, obtained by calculations based on the total radical concentration, are explained. (2) The corrected values for kd appear to support the observation of Kerr and Trotman-Dickenson that
[CONTRIBUTIOSFROM
THE
AND
J. G. MILLER
Vol. 86
a t temperatures above 300" kd falls well below an extrapolation of data obtained a t lower temperatures, leading to an apparently lower value of E 4 . I t seems unlikely that the true value of El is as low as 22 or 25 kcal., and the low values observed probably arise from a systematic decline in the pre-exponential term with rising temperature. Speculation as to the cause of such behavior is not warranted here.
DEPARTMENT O F CHEMISTRY, UNIVERSITY
OF
PENNSYLVANIA, PHILADELPHIA 4,
PENKSYLVANIA]
The Heats of Combustion of Aliphatic Long Chain Peroxyacids, &Butyl Peroxyesters, and Related Acids and Esters' BY HOWARD A . SWAIN, J R . , LEONARD S.SILBERT, AND
JOHN
G . MILLER^
RECEIVED JANUARY 29, 1964 T h e heats of combustion of peroxylauric, peroxymyristic, peroxypalmitic, and peroxystearic acids, myristic, palmitic, and stearic acids, t-butyl peroxycaprate, peroxylaurate, and peroxymyristate, and t-butyl caprylate, caprate, iaurate, and myristate have been measured. By use of values or estimates of values of the heats of fusion and vaporization, the results have been converted t o the 25" standard-state heats of combustion for the gaseous substances for ail the peroxyesters and esters and for peroxylauric acid. T h e application of the thermochemical group additivity rule to the heats of formation of organic peroxides has been examined. -4higher order approximation to the general law of disproportionations may be needed for some of the peroxides. Calculated dissociation energies of I-butyl peroxylaurate and peroxylauric acid are compared with enthalpies of activation observed for the free radical decompositions of these substances.
Introduction The understanding of many of the properties of organic peroxides will be aided greatly by the thermodynamic information that can be obtained by direct calorimetric determination of the heats of combustion of those substances. Unfortunately, such determinations have been generally impractical owing to the instability and lack of purity of the peroxides so that, for the great part, less direct means have .been used to estimate those thermodynamic quantities. Recently, improved methods have been for the preparation of long chain aliphatic peroxyacids and t-butyl peroxyesters in pure form. With these highly purified, relatively stable substances available, it has been possible to determine for the first time their heats of combustion. In order to obtain auxiliary data for the calculation of the heats of some other reactions, the heats of combustion of the parent fatty acids and their t-butyl esters have also been measured. Experimental Materials.-The long chain fatty acids used in the combustion measurements and as starting materials for the preparation of the esters, peroxyacids, and peroxyesters were obtained from the Eastman Kodak Co. and, after fractional distillation at 4-7 mm., samples converted t o their methyl esters were examined for homologs by gas-liquid chromatography. Methyl esters of several of the acids were also tractionated. No homologs or other impurities were present in any case. The fatty acids whose heats of combustion were t o be determined were then treated by recrystallization procedures designed t o convert them into their C forms.s T h e final identification and (1) A report of work done under contract with t h e U. S. Department of Agriculture and authorized by t h e Research and Marketing Act of 1946. T h e contract is being supel-vised by t h e Eastern Ctilization Research a n d Development 1)ivision of t h e Agricultural Research Service. Presented a t the 14.5th riational Meeting of the American Chemical Society, New Y o r k , N Y , S e p t . , 1963 ( 2 ) To a h o r n inquii-ies are to be directed ( 3 ) W. E. Parker, C. Ricciuti, C I, Ogg, and D . Swern, J . A m . Chem. SOL., 11, 4087 1195.5) ( 4 ) I,, S. Silbert and D Swern. rbid , 81, 2364 (1959). ( 5 ) I, S . Silbert, E . Siegel, and D. Sweun, J . O r g . Chem., 17, 1336 (1962).
estimation of the polymorphic forms present was always carried out after pelletizing the acid. T h e third- and fifth-order long spacingss on the Cu Ka: X-ray diffraction powder photographs were measured for this purpose. Stearic acid was recrystallized successfully from olefin-free, dry, freshly-distilled petroleum ether a t room temperature and then kept a t 50" for 48 hr. t o eliminate completely all traces of the A form.6 Palmitic acid was converted satisfactorily by recrystallization from acetic acid. Myristic acid was recrystallized from its own pure melt and also from acetic acid. In the case of lauric acid, neither the recrystallization from the melt nor t h a t from various solvents a t different temperatures gave reproducible polymorph content and, furthermore, the content changed on standing. Similar difficulty with lauric acid is indicated in studies made by Davies and Malpass.' The peroxyacids were prepared from the f a t t y acids by the method of Parker and co-workers3 and Silbert and co-workers.5 T h e peroxy oxygen content of these compounds determined by iodometric titration3j5was between 99.4 and 100.57, of theoretical, with half of the batches having values in the range 99.9 t o lOO.17,. Pelletizing had no effect on the peroxy oxygen content. The long spacing of the peroxypalmitic acid, 37.2 A., was the same as found by Swern and co-workers.8 I t was unaffected by pelletizing. S o evidence for polymorphism within the evencarbon peroxyacid series below 18 carbon atoms has been noted t o date, although recent evidence suggests a polymorphic difference between the odd and even ~ e r i e s .Only ~ the even ca; on members are considered in this investigation. Both the t-butyl esters and the t-butyl peroxyesters were prepared by pyridine acylation methods.4,10 T h e esters after fractional distillation in vacuo showed no impurities in gas-liquid chromatography. The peroxyesters were purified by passage through silica gel and elution with petroleum ether containing 4 t o 59" of diethyl ether, the best fractions being chosen by iodometric titration." The peroxy oxygen content of the liAi.i1 prod(6) (a) K . S. Markley, " F a t t y Acids," P a r t I , Interscience Publishers, ( b ) E. von Sydow, Arkiw K ~ m i9 , 231 Inc., New York, I C Y . , 1960: (1956). (7) M. Davies and V . E . Malpass, J . Chem Soc., 1048 (1961). ( 8 ) D . Swern, I,. P. Witnauer, C. R . E d d y , and W . E . Parker, J . A m , Chem. Soc., I T , 5537 (195,5). (9) I.. S Silbert, E . Siegel. and T) A . I.utz, Abstl-acts. 146th National Meeting of t h e American Chemical Society, Denver, Colo.. J a n . 20-23, 1964, paper no. 54, p. 221). (10) I
ZW, cal
13 5 15 0 17 4 13 0 15 0 19 8 18 0
-0
-AH,',
kcal mole-'
kcal.
7 7
2089 8 2098 2 2090 5 2096 2 2098 3 - .6 2090 3 - .6 2093 7 - 6 Mean 2093 9 f 2 3
-
-
tion of the mean standard deviations of all the quantities from which the separate -AH," values were computed. The Heat of Fusion of Peroxylauric Acid.-Because of the rapid decomposition of the molten pure acid, no effort was made to determine the heat of fusion directly Instead, the AHru.value was obtained by measuring the freezing points of solutions of peroxylauric acid with hexadecane or methyl laurate. Six solutions were measured, two with hexadecane and four with methyl laurate, the mole fractions, N ,of the peroxylauiic acid lying in the range 0.9657 to 0.9913. Philadelphia differential thermometers, supplied by the Precision Scientific Co., were used to read the temperatures. Plots of -In AT against l / T , where T was the freezing point (point of separation of peroxylauric acid), showed t h a t all six points lay close to the same straight line revealing remarkably ideal behavior from which could be calculated. It will be recalled t h a t peroxycarboxylic acids are monomeric in benzene ~ o l u t i o n . ' ~X least-squares fit of the data t o the equation
-In
N
=
(AHfuS/I?) [(l/TnJ)- (1/T) ]
gave AHru. = 11.07 f 0.18 kcal. mole-' and the melting point value 49.39 f 0.16" for the peroxylauric acid.
Results The standard state heats of combustion a t 25' are given in Table 11. TABLE I1 STANDARD STATEHEATSOF COMBUSTION AT 25" FOR A SERIES OF PEROXYACIDS, ACIDS,PEROXYESTERS, A N D ESTERS R
No. of samples measured
-AHo', kcal mole - 1
Solid peroxyacids, RCOOH
I1
0 n-CllH23 6 Peroxylauric n-ClsH~7 7 Perox ymyristic n-Cl&I 6 Peroxypalmitic n-C17H36 9 Peroxystearic Solid acids ( C form) RCOH
1785.8 f 2 . 0 2093.9 f 2 . 3 2406 2 f 2 . 2 271i.i i 2 7
I!
0
Myristic Palmitic Stearic
n-CtJL7 9 2073.6 f 2 . 3 n-ClJ%l 6 2384 7 f 2 . 1 7 2693.9 f 3 . 2 n-CirH35 Liquid t-butyl peroxyesters, RCOOC4Hg
I/
0 Perox ycaprate Peroxylaurate Peroxymyristate
n-CgHlg n-Cdh n-C13Hz7
7 7 4
2108.6 f 2 . 1 2421.4 f 2 4 2732.4 f 2 . 5
Liquid t-butyl esters, RCOC4H9
I1
T h e error estimated for the mean value of -AH," was computed as a standard deviation value for the whole set by propaga( 1 2 ) J. Coops, R S.Jessup. and K . van Xes in "Experimental Thermochemistry," F. I) Rossini, E d . , Interscience Publishers, Inc , New Y o r k , 1\; Y , 1956, C h a p t e r 3 . (13) W N Hubhal-d, I). W S c o t t , and G . Waddington. ref. 12, C h a p t e r 5 , see also E J Prosen. C h a p t e r 6
Caprylate Caprate Laurate Myristate
n-CTH n-CgH19 n-CllHnn n-ClaHni
0 8
i
4
1781.4 f 3 0 2093 2 f 1 9 2402 4 f 2 9 2716.4 f 3 0
(14) J R . Rittenhouse: W. Lobuenz. D S w e r n , and J C Miller, J Chem S o c . , €40, 4850 (1958).
Am
H . A. SWAIN, J R . , L. S. SILBERT, A N D J. C . MILLER
2564
A useful formulation of the values of the heats of combustion of such substances may be written -AHc" = cy Pc, where p is the increment per carbon atom in the acyl group, c being the number of those carbon atoms. This formulation serves as a check on the self-consistency of the data and also permits useful correlations of the AH," values of one type of substance with those of another. Table I11 gives the 25" values of cy and @ obtained by formulation of the values in Table I1 for each of the types of compounds. Here and in all the subsequent cases, the errors shown for cy and p are the standard errors of the least-squares fit of the data to the formula
+
TABLE I11 THEV A L U E S
OF a A S D
6
FOR THE D I F F E R E N T
TYPES OF
SUBSTANCES IS TABLE I1 p . kcal mole-' per Type
Solid fatty acid Solid peroxyacid Liquid t-butyl ester Liquid f-butyl peroxyester
0 ,kcal
-97 -80 +535 +549
mole -1
1 1 4 4
f4 3 f4 9 =k 3 4 f3 1
acyl carbon a t o m
155.1 f 0 . 3 155.4 f . 3 1.55 71 i 31 155.95 f 26
The 25" values of -AHc" of only three of the individual substances listed in Table I1 are available elsewhere for comparison. These are the three fatty acids, myristic, palmitic, and stearic, the most accurate measurements made of them to date being those of Adriaanse,15 who obtained the values for those acids in their resolidified stabilized forms (C forms). His values were 2073.8, 2384.6, and 2695.8 kcal. mole-', respectively, which agree with ours within the limits of error stated in Table 11. Adriaanse also determined the heats of fusion of those fatty acids with good accuracy. If we apply those heats of fusion to our fatty acid values we obtain a = --1(il.9 f 3.9 and /3 = 156.2 f 0.2 for -AHc" of the liquid substances a t 25". Adriaanse found cy = -102.82 and p = 156.26 for the whole series of fatty acids from c' = 5 to 20 a t 25" in the liquid state. The value l5G.%Gkcal. mole-' for the increment per CH? group appears also to fit the standard heats of combustion of liquid normal alcohols,lfi acids,I6 and esters, and liquid normal parafTinslSa t 23". The 25" heats of vaporization, AH,,,, may be taken from the work of S p i z z i ~ h i n o 'on ~ these three f a t t y acids by use of her extrapolations of accurate vapor pressure measurements which were carried down to the melting points of the acids. Application of these AH,,, values to the -AHco(l) values gives cy = -85.4 3 , 2 and 9 = 157.2 f 0.2 for -AHc"(g) for these fatty acids a t % S o , where the vapor pressures are so low (well below 1 m m . ) that the vapors must be very little associated. These results may be compared with the values which are obtained on applying to our -AHc" (s) values the formulation of the heats of sublimation. A H s u b , determined by Davies and Alalpass': A H s u h (kcal. mole-]) = 12.7 1.5Oc, for the even-carbon acids from Clo to Cna. The effect of temperature upon the heat of sublimation is not considered in their formulation which was based on values of AHsu),determined from sublimation pressure varia-
*
+
(1.5) S Adi-iaanse, Ilissei tdtion. l'iee I'niversity of A m s t e r d a m , 1460 ilfi! (8) J €1 S . G,-een, ( ' h r m i i i d Il.ondon!. 1213 (lB60). ( b ) J H. S. 'i Kco (I.ondrrn), 16, 1.18 11061) pizzichino. J K p c h C ~ n t r e.Ya[i, Rech. S c i . L a b Brllrvue ( P a r i s ) , So 3 4 [IWYV
Vol. 86
tion over short temperature ranges different for each member of the series. The range for the Clo acid was 16.8-27.3' and for C22 was 71.GO-78.30°. On the questionable assumption that their formula gives the AHsubvalues a t 25' for the three solid fatty acids, we obtain a = -84.4 f 4.3 and = 156.6 0.3. These extrapolations of both the Spizzichino and the Davies and Malpass data to 23" are rather hazardous because of the temperature effects one might expect on the vaporization and sublimation of such highly associated substances. n'ormal behavior would be expected for the esters which are little associated. Spiz~ichino~'showed strikingly the additivity of group contributions18l 9 to the vapor pressure by the exact equality of the vapor pressures of the three esters, ethyl stearate, n-octadecyl acetate, and n-decyl decanoate, a t fourteen temperatures over the range 68 to 125". She has made accurate measurements of the vapor pressures of the methyl esters of a series of higher fatty acids a t temperatures suficiently low that by interpolations or short extrapolations we may estimate accurately the 23" values of AHvapfor the methyl esters of the same molecular weights as the esters in Table 11. If a slight correction is made for the effect of the configuration of the t-butyl group, these values may be used to convert the heats of combustion of the esters to the gaseous state values. The correction was made by first estimating the difference between the normal boiling points of the t-butyl esters and the methyl esters of the same molecular weight by use of the KinneyId group contributions. These values were then used to calculate the correction by use of the empirical relation established by Scott?O for AH,,, a t 23" in terms of Tb. The corrections were only 1.1 kcal. mole-' for the t-butyl laurate and niyristate and 1.2 for the other two esters. When these estimates of AHvapare applied to the -4Hco(l) values in Table 11, the results give cy = 543.4 =k 3.5 and 9 = 133.8 =t0.3 for -AHc"(g) for the four t-butyl esters a t 23". With the t-butyl peroxyesters, AHvapcan be taken the same as for the parent t-butyl esters on the basis of the observations of Egerton, Emte, and MinkofP1 that in going from the parent compound to a peroxide the change in heat of vaporization due to the introduction of the extra oxygen atom is quite small except for peroxyacids. This gives cy = 537.6 f 3.1 and p = 157.1 f 0.3 for -AHco(g) a t 25" for the t-butyl peroxyesters in Table 11. While conversion of the heats of combustion for the solid fatty acids to the standard gaseous state values a t 23" was complicated by the strong intermolecular H-bonding in the condensed states a t that temperature, no such complication appears to exist for the liquid peroxyacids. Here the strong intramolecular H-boiidZ2 prevents intermolecular bonding. Egerton, Emte, and Minkoff2l measured the vapor pressures of liquid peroxyacetic, peroxypropionic, and peroxybutyric acids
*
( I S ) C . R K i n n e y , .I. A m C h r m Soc , 60, 3032 (1938) ( l e ) K J Laidler, C a n . .I, C h r v a , 34, 626 (1956). ( 2 0 ) J H Hildehrand a n d R I. S c o t t , " T h e Solubility of Nonelectrolytes," 31-d E d , Reinhold Publishing Coi-p , Yew Y o r k , X Y , 1850, pp. 427-128 ( 2 1 ) A C. E g e r t o n , K' E m t e , a n d G. J Minkoff, D i s c i i s s i o m Faraday SOL, 10, 278 (1951). (22) For a listing of t h e evidence lo]- t h e s t r o n g intramolecular hydrogen bonding, see ref 11
July 5 , 1964
HEATSOF COMBUSTION OF 1-BUTYL PEROXYESTERS
over a wide temperature range (0 to 110 or 120') and found that plots of log p against l / T were linear over the whole range, indicating rather constant values of AHva, over this range. The vapor pressures were higher than those of the corresponding fatty acids a t the same temperature. The normal boiling points as given by the vapor pressure determinations, which were carried up to those boiling points, were 110, 120, and 126O, respectively. From the data available, a comparison of the standard state heats of combustion of gaseous lauric and gaseous peroxylauric acid may be made. For the lauric acid, one may take the accurate 25" value of - A H c o ( s ) determined for the C form by Adriaanse,Is 1763.1 f 0.2 kcal. mole-', and add the heat of sublimation, AHsu, = 30.7 kcal. mole-', estimated by the smoothing formula of Davies and Malpass7 for the C form. This AH,,, value should be correct for this purpose since the range a t which their formula applies to lauric acid is 22.7 to 41.4'. This gives -AHco(g) for lauric acid as 1793.8 kcal. mole-'. For peroxylauric acid, we have measured AHfusand may estimate the 25' value of AHvapby use of the Scott equation.20 To do this, we have used the normal boiling points of the three peroxyacids mentioned above to determine the peroxycarboxylic group contribution to the boiling point by the Kinney procedurela and have used this to estimate the boiling point of the peroxylauric acid. This gives AHva, = 16.9 f 0.3 kcal. mole-' a t 25'. Applying AH,,, and AHvap to the -AHCo(s) value for peroxylauric acid in Table 11, we obtain - AHco(g)= 1813.8 kcal. mole-'. The error in this estimate is about 1 3 kcal. mole-', being much greater than the corresponding error for the lauric acid. The following equation can now be written. 0
II
+
C I I H Z ~ C O H ( ~ '/*Odg) ) --+O...H
li I
C l l H d 2 0 0 ( g ) ; AH" = 20.0 f 3.0 kcal. mole-'
(1)
Discussion The results given above may be discussed from the rtandpoint of their relation to other thermochemical values, especially those determined by other workers for peroxy compounds. The discussion is given below in two parts. Group Additivity Calculations.-Benson and in their systematization of the laws of additivity of molecular properties have put out a general limiting law which shows clearly the superior accuracy of calculations based on the additivity of group properties as compared with bond or atomic properties. They were also able to show that the additivity of group properties gives estimates of heats of formation, AHfo, with high accuracy in general, whereas the result of use of bond values is one-fifth as accurate and the additivity of atomic values fails badly for enthalpy calculations. In their analysis of the data available for AHfo of organic peroxides, they found strong signs of inaccuracies in the determinations of some of the best known values as shown by gross deviations from group additivity. The values determined here for the esters and peroxyesters meet the self-consistency test t h a t Benson (23) S W. Benson and J H . Buss, J C h r m . P h y s . , 39, 546 (1958)
2565
and Buss suggest when not all of the group contribution values are available for calculation of the A H f O values themselves. Thus, by using their group contribution [C-(C)z(H)z] = -4.95 kcal. mole-' and the A H f o values,24-68.3171 kcal. mole-' for HzO(1) and -94.0518 kcal. mole-' for COz(g), i t can be shown that the p term for the - A H c o ( g ) formulation a t 25' should be 157.42 kcal. mole-' per carbon atom in the acyl group, which is close to the values we have found for the esters and peroxyesters. Benson and Buss questioned the values of AH," reported by Jaffe, Prosen, and SzwarcZ5 for the diacyl peroxides ( C H ~ C O Z )(C2H6C02)zr ~, and (n-C3H7C02)2on the basis of the large increments in AHt" for the series. As an example of the use of the general limiting law for disproportionations, Benson and noted that on the basis of the group additivity approximation A H o should be zero for the reaction HzOr(g) t-CqHQOOC?H&) Zt-CiHgOOH(g), but that the literature values of the heats of formation for the peroxides involved give A H o = 12.6 kcal., a discrepancy far beyond any others they found in use of group additivity. They used this to emphasize the inconsistency of the data for the peroxides, especially for the hydroperoxides. In this connection, i t should be noted that causes do exist for failure of group additivity in the case of some peroxides owing to interaction of the groups with each other across the 0-0 nucleus. For example, it is known t h a t in cumene hydroperoxide the H atom of the OH group forms a hydrogen bond with the benzene ring across the 0-0 nucleus the strength of the bond being 0.9 kcal.26 A similar interaction across the 0-0nucleus occurs in l-phenylcyclohexyl hydroperoxide,27 A stronger interaction occurs in the case of peroxyacids, as shown by the strong intramolecular hydrogen bond formed by the H atom of the OH group with the carbonyl group across the 0-0nucleus.22This means that the group contribution sum given by A H o for reaction 1
-
[O-(CO)(O)]
+ [O-(O)(H)] - [ 0 - ( C O ) ( H ) ]
+
=
20.0 =t3.0 kcal.
is probably not applicable to AHf' calculations in general. The extent by which the separate parts of that group sum are affected by the internal hydrogen bond cannot be estimated from data available a t present, but i t is obvious that a higher order approximation is needed. Whether one can simply make a constant correction for the hydrogen bond is not certain since the hydrogen bond may weaken the 0-0 bond. Gig u h e and Olmosz8estimated that the internal hydrogen bond energy is roughly G kcal. mole-' in peroxyformic acid and 7 kcal. mole-' in peroxyacetic acid. They and others? have suggested that the instability of the peroxyacids is indicative of a weak 0-0bond caused by the chelation. The peroxyacids studied here are much (24) F. D. Rossini, 11. D W a g m a n , W. H. Evans. S.I.evine, and I Jaffe, "Selected Values o f Chemical Thermodynamic P-operties," National Bui.eau o f S t a n d a r d s Circular 500, U S. Government Printing Office, W a s h i n g t o n , D . C . , 1952 ( 2 . 5 ) L. Jaffe, E . J Prosen, and M Szwal-c, J . C h r m P h y s , 27, 416 (1957). ( 2 6 ) V. V Zharkov and N . K . Rudnevskii, Opl i S p r k / i . o s k o p r y n . 7 (No. 6 ) . 497 (1959) (27) H . K w a t t and R T Keen, J . A m ChPm . S O L , 81, 943 (19.59) (28) P. A Gigueie and A . U' Olmos, Con J C h m , 3 0 , 821 (19J21
H. A. SWAIN,J R . , L. S.SILBERT, AND J. G. MILLER
2566
more readily decomposed thermallyz9than are the pero x y e s t e r ~and ~ ~ have weaker 0-0 bonds as determined polarographically. lo Dipole moment studies of the peroxyacids14 and of the peroxye~ters~l indicate that the hydrogen bond in the peroxyacids twists the configuration about the 0-0 bond greatly away from the stable skew configuration. This twisting would bring about a repulsion between the lone-pair s-orbitals of the peroxy oxygen atoms that would, therefore, weaken the 0-0 bond. Bond Dissociation Energies.-In the reversible reaction Ei XY,X.
+Y. E,
where the enthalpies of formation (25') of the gaseous species involved are AHf(XY), AHf(X.), and A H f (Y .), the relation A H = AHf(X.) AHr (Y . ) - AHf(XU) is the basis of the kinetic method of determining bond dissociation energies developed by Szwarc.3 2 He set AH = E1 - Ez, the difference between the activation energies of forward and reverse reactions. The bond dissociation energy is then equivalent to the energy of activation for decomposition of the X-Y bond when recombination of the radical fragments involves no activation energy (E? = 0). This procedure has been used extensively by Gray and Williams33 for determining the thermochemistry of alkoxy1 radicals. The Szwarc procedure may be applied as a test comparison of the calculated and experimental energies of activation for the free-radical decomposition of t-butyl peroxylaurate and peroxylauric acid. Any serious discrepancies would suggest either higher order additivity effects in the heats of formation used or unexpected decomposition phenomena related to the activation energies. The heat of formation calculated from the heat of combustion of t-butyl peroxylaurate a t 25' is -155.1 kcal. mole-'. The heat of formation of the lauroyloxy radical is estimated34to be -96.4 kcal. mole-'. Using AHf(t-butoxy radical) = -25 kcal. per the derived gaseous dissociation energy of 33.7 kcal. mole-' compares to 35 kcal. mole-' obtained from solution studiesso of the thermal decomposition of t-butyl peroxyesters in chlorobenzene, nitrobenzene, and diphenyl ether. A slightly greater value of the activation energy is expected in the solution reaction because of caged recombination of the radicals produced.35 This recombination also has the effect of making the observed frequency factor slightly higher than normal, i.e., yielding a positive value for the entropy of activation, an effect which is in accord with the results obtained for the peroxyesters in solution.30 It appears t h a t the
+
Vol. 86
agreement of the observed and calculated energies of activation is satisfactory in this case. The peroxyesters apparently contain a normal 0-0 bond, there being little interaction between the t-butyl group and the carbonyl In a similar calculation for peroxylauric acid, a great discrepancy results. The gaseous heat of formation a t 25' is calculated to be -134.6 kcal. mole-'. Taking AHf(H0.) = 9 kcal. per mole,33the calculated dissociation energy is 47.2 kcal. mole-' compared with 19 kcal. mole-', the experimental activation energy determined by Vorobiev and c o - ~ o r k e r for s ~ ~the free-radical decomposition of peroxylauric acid in benzene (decomposition in the presence of diphenylpicrylhydrazyl) . The 28 kcal. difference greatly exceeds all errors due to the data used or due to failure of group additivity wherever used in this evaluation. There may be, therefore, some unexpected phenomenon occurring in the process measured by Vorobiev, et al. That such a phenomenon does exist in that process is revealed by calculation of the entropy of activation. Using their velocity constants (Table 11, ref. 36), calculation shows that AH* is 19.1 1.5 kcal. mole-', which is in accord with their value of El, but also that AS* is -30.4 f 4.3 e.u. This large, negative entropy of activation was not considered by Vorobiev and his co-workers in their discussion of the speed of the reaction and the strength of the 0-0 bond relative to those of other peroxides. Such a value may be due to the participation of other particles besides the peroxylauric acid in the activated complex. In any case, the highly abnormal A S * value (low frequency factor) would rule out use of the Szwarc procedure of estimating the dissociation energy in the case of peroxylauric acid on the basis of data now available. For the same reason, the rate constants for the thermal decomposition of peroxyacetic acid in aromatic solvents3' and in the vapor phase (in a stream of toluene)38cannot be used to obtain the 0-0 bond dissociation energy. For example, we have calculated from the k-values for the decomposition reaction in liquid toluene
*
0
CHjCOOH
0 --f
II
CHaCO,
+ .OH
(20) W . E P a r k e r , L. P. Witnauer. and D . Swern, J . A m . C h e w . Soc., 80, 32;i (1958) (30) h1 T r a c h t m a n a n d J . C . Miller, :bid., 84, 4828 (1962). (:i1) F. 11 Verdei-ame and J. G . Miller, J P h y s . Chem., 66, 2185 (1962). (:32) M Szwarc. C h e m . Rrw , 47, 75 (1950). (:33) P. G r a y a n d A . Williams, i b i d . , 89, 239 (1959).
(Table 11, ref. 37) that AH* = 20.1 f 0.3 kcal. mole-' and A S * = -27.9 + 0.7 e.u., values which are substantially the same as for peroxylauric acid in benzene. For the decomposition of gaseous peroxyacetic acid in gaseous toluene,3bthe data are less accurate but indicate even smaller AH* and larger negative A S * values. The entropy values indicate need for further investigation of the mechanism of the thermal decomposition of peroxyacids. Acknowledgments.-The authors thank F. D. Verderame, M . J . Kornblau, and A. Hettinger for aid in some of the measurements. The .use of facilities a t Wilkes College is also gratefully acknowledged.
(31) This computation was derived from AHf reportedzs for t h e acetoxy radical ( - 4 3 kcal. mole-') by adding - 6 8 kcal. mole-1 for t h e fii-st - C H r inci-ement and -44 5 3 foi- t h e remaining nine methylenes (I-ef 23, p. 564). (3.5) %f I x v y , M. Steinberg, and M Szwarc, J . A m . Chem Soc , 7 6 , 5978 ( 19.54 )
(36) V . Vorobiev, D . Lefoi-t. J Sorba, a n d D Rouillal-d, Bull. soc chrm. France, 1577 (1962). (37) F. W. E v a n s a n d A . H. Sehon, C a n . J . Chem , 41, 1826 (1963). (38) C. Schmidt a n d A H Sehon, i b i d . , 41, 1819 (1963)
