INITIATOR
EFFICIENCY OF 2,2’-AZOBIsISOBUTYRONITRILE
predicted and the measured AF*zas.z may be due to the solvent. Only studies of Dure amides were used to develop the correlation. A solvent effect upon E, in dimethylamides has been r e c o g n i ~ e dbut , ~ ~a ~quantita~~ tive relationship between E’, or AF* and solvent properties has not yet been developed. It would be of value to measure the activation parameters of all dimethylamides in the same solvent. However, it would be necessary to find a solvent, preferably inert, which remained liauid from about 225 to 395°K. It is likely that the other amides in Table I will also have relatively low values for E, and AF*z98.2, primarily because of their very low COakSCenCe temperatures
2699 when compared with similar but not as highly substituted amides.
Acknowledgments. We would like to thank Lester Isbrandt for synthesizing amides IV and VI-IX. Chang Y. Chang, Gary Martinie, and Michael Schafer have also assisted a t various stages of this project. We wish to acknowledge the support of the National Science Foundation through summer research institutes in which L. I. and R. E. D. participated. (13) J. C. Woodbrey and M. T. Rogers, J . Amer. Chem. Soc., 84, 13 (1962). . . (14) A. G.Whittaker and s. Siegel, J. Chem. Phys., 42, 3320 (1965).
Photo- and Thermal Initiator Efficiency of 2,2’-Azobisisobutyronitrileat 25 by R. D. Burkhart and J. C. Merrilll .Department of Chemistry, University of Nevada, Reno, Nevada
(Received January 15, 1969)
TJsing thiols as free-radical scavengers and exceedingly long reaction times, the fraction of kinetically free radicals produced in the thermal decomposition of 2,2’-azobisisobutyronitrile(ABN) has been measured in benzene and cyclohexane a t 25’. Using a chain reaction as a radical counter, a similar fraction is obtained for the photodecomposition of ABN thereby providing, for the first time, a comparison of these efficiencies a t a common temperature. The photoprocess is found to be somewhat more efficient in both solvents but both photo- and thermal decompositions are less efficient in cyclohexane than in benzene.
Introduction Peroxides or azo compounds are widely used as initiators of free-radical reactions, but they rarely function with 100% efficiency; therefore independent experimentation is needed to measure actual rates of initiation. The techniques usually used include addition of known amounts of a radical inhibitor followed by a measurement of the inhibition period or by the use of a radical chain reaction for which the overall rate constant is known and which, therefore, can function to measure the frequency of the initiation process. Another method involves the use of a scavenger such as iodine, a thiol, or a moIecule containing an unpaired electron so that radicals produced in a kinetically free state can be counted by the rate of disappearance of the scavengermZ I n the present study the last of the above methods has been used to measure the fraction of kinetically free radicals produced in the thermal decomposition of 2,2’azobisisobutyronitrile (ABN) a t 25”. Interest in the reaction is an outgrowth of kinetic studies being carried out in these laboratories on the radical chain reaction between triethyl phosphite (TEP) and various thiols.
All of this work has utilized ABN as initiatorJaj4and at the outset, rates of initiation in the photochemical process were determined by measuring rates of polymerization of methyl methacrylate. It was decided to make another evaluation of the initiator efficiency in this reaction for several reasons. First, it has been noted6 that not all radical scavengers perform their function with equal efficiency. Also, the evaluation of polymerization rates relied on gravimetric analyses of polymer formed, and a more accurate and precise analytical method was desired. A rather straightforward solution to these problems would be to evaluate the initiator efficiency directly, using a thiol as scavenger. The potentiometric titration of thiols using (1) NDEA Predootoral Fellow, 1966-1969. (2) G. M. Burnett and W. H. Melville, “Technique of Orgmic Chemistry,” Vol. VIII, Part 11, S. L. Frieas, E. S. Lewis, and A. Weissberger, Ed., Interscience Publishers, Inc., New York, N. Y . , 1963,pp 1109-1110. (3) R. D.Burkhart, J. Phys. Chem., 70, 605 (1966). (4) R. D.Burkhart, J . Amer. Chem. SOC.,90, 273 (1968). (5) (a) J. C. Bevington, Trans. Faraday SOC.,51, 1392 (1955); (b) J. C. Bevington, H. Bradbury, and G. M. Burnett, J. Pol2/mer Sci., 12, 469 (1954). Volume 75, Number 8 August 1969
R. D, BURKHART AND J. C. MERRILL
2700
mercuric ion can be carried out with good precision and accuracy and disulfides do not interfere with the analyFurthermore, the scavenger reaction, involving abstraction of a thiol hydrogen atom, is the same as the assumed initiation step for the thiol-TEP reaction. Unfortunately, the extinction coefficient of ABN a t 366 mp7 is so small that the necessary experiments would be too lengthy to be practical using photodecomposition. It was decided, therefore, to measure the efficiency of the thermal process which, even though more time consuming, had the virtue of not requiring equipment needed for other work and made possible an indirect determination of the efficiency of the photoprocess. In spite of the extensive investigations made on ABN decompositions,s it is interesting that the efficiency for thermal decomposition has never been measured at 25". Also, the efficiencies of photo- and thermal processes have never been measured at a common temperature.
Experimental Section
the thermal decomposition of ABN at 25", utilized the same type of inverted Y vessel described above. I n these experiments the initiator solution was kept separate from the solution containing thiol and TEP during degassing. The solutions were brought to thermal equilibrium in the dark a t 25" before being mixed. Zero time was taken at the instant of mixing. Here again a replicate of the reaction solution was prepared and analyzed to give a zero-time thiol concentration. I n all reactions the thiol analyses were carried out by potentiometric titration using mercuric ion.G Light intensities were measured using the potassium ferrioxalate actin0rneter.O The monochromaticity of the light transmitted by the Corning No. 5840 filter and the Pyrex reaction cell was checked by preparing a transmission curve for this combination using a Cary 14 spectrophotometer. It was found that less than 1% of light above 3960 and below 3130 is transmitted. Taking into account the relative intensities of the emission lines from the mercury arc, it is estimated that greater than 98% of the quanta counted by actinometry result from the 3660-A emission.
Purification of the thiols and ABN (Aldrich Chemical Co.) and of triethyl phosphite and benzene (Eastman Results and Discussion Organic Chemicals) has been described p r e v i o u ~ l y . ~ ? ~ Table I gives a summary of the scavenger experiments Cyclohexane (Eastman) was extracted several times carried out in benzene and in cyclohexane. Values of with concentrated H2SOa,was washed with dilute base f ' were obtained from the relation and with water, and was then dried over anhydrous sodium sulfate. It was then distilled twice through a [RSH],, - [RSH], = 2f'[ABNIo[l - exp(--lcdt)l (1) 30-in. Vigreux column; center fractions only were retained on each distillation. Table I : Values off' for the Thermal Decomposition of The apparatus used to carry out photochemical reacABN in Benzene and Cyclohexane at 25' tions has also been described b e f ~ r e . ~The vessels Reaction used for thermal reactions were in the shape of an [RSHlo [RSHlt [ABNlo time inverted Y with 15-mm Pyrex tubing forming the two x IO*, x IO*, x 102, x 10-8, arms branching off from the central 20-mm tube. A M M M sec f' 24/40 ground-glass joint was sealed to the end of the Benzene central tube so that the vessel could be attached to the 6.05 0.56 2.95 3.28 3.79" vacuum system by way of an adapter and stopcock. 6.03 0.54 1.89 4.09 2.84" 6.11 0.54 2 96b 1.77 0.500 Scavenger experiments were carried out by separately 6.04 0.54 0.623 1 . 4 0 0. 94Bb preparing solutions of ABN and of the thiol. Five Av 0.54 milliliters of each of these solutions was then placed in Std dev 0 . 0 1 each arm o f the reaction vessel. I n these experiments Cyclohexane one of the arms of the vessel was constricted near the 6.10 0.49 0.967* 0.755 1.01 branch of the Y so that it could be pulled off later. 6.24 0.50 0.439 2 . 23b 2.13 The separated solutions were subjected to several 6.56 0.52 1.13 0.627 1,2ab freeze-evacuate-thaw cycles to remove oxygen. They Av 0.50 were then mixed in the arm having the constriction, and Std dev 0.01 the mixture was immediately frozen. The arm was 1-Octadecanethiol. ' 1-Butanethiol. then carefully pulled off under vacuum by heating at the constriction. The solutions were kept in the dark in a 'Onstant temperature water bath at 25" for periods On (6) J. S. Fritz and T. A. Palmer, Anal. Chem., 33, 94 (1961). the order of 10 weeks. Another 5-ml portion of the (7) p. Smith and A. M. Rosenberg, J. Amer. Chem. SOC.,81, 2037 I
tration. Reactions between TEP and the thiol, initiated by The Journal of Physical Chemistry
(9) C. G. (1958).
Hatchard and C. A. Parker, Proc. Roy. SOC.,A235, 518
INITIATOR
2701
EFFICIENCY O F 2,2'-AZOBISISOBUTYRONITRILE
and assumed a 1: 1 stoichiometry. The decomposition rate constant for ABN a t 25" was calculated from the Arrhenius parameters obtained by Van Hook and TobolslrylO and is found to be 4.05 X sec-'. For present purposes, the quantity of interest is f'kd which sec-l in benzene and has the average value 2.20 X 2.03 X lov8 sec-1 in cyclohexane. Only Walling and Kurkovll and Carlsson and Ingold12 have made measurements off' at temperatures approaching those used here. The former workers found a value of 0.22 (at 29.5'), and the latter authors concluded that the best value was 0.50 (at 30'). We have found no tendency for f' to vary with either changing thiol or ABN concentration. Also, 1-octadecanethiol and 1-butanethiol apparently have the same scavenger efficiency, and the effect of changing solvent is of small but measurable magnitude. Reaction times, in units of days, varied from 69.8 to 75.8 corresponding to ABN conversions of 21.5 to 23.3%. Conversions with respect to thiol varied from 4.5 to 40.2%. Neither of these variations seemed to have an appreciable effect onf'. The next series of experiments involved measurement of the rates of reaction between 1-butanethiol and TEP using thermal initiation by ABN. Previous workSs4has shown that when [TEP]/ [RSH] is sufficiently large the rate of thiol disappearance obeys the relation -d [RSH]/dt
=
lep/k~1~z[RSH]Ri"2
(2)
where, for thermal initiation Ri = f'kd[ABN]. The over-all mechanism which has been found consistent with this rate law is ABN ---t 2A iY2 (3)
+
+ RS RS + P(0Et)n +R + SP(0Et)a R + RSH -% RH + RS A,
+ RSH
---f
AH
(4)
Table 11: Vrtlues of kp/kt'/Z for the Thermally Initiated 1-Butanethiol-TEP Reaction in Benzene at 25""
2.59 3.27 2.34 2.17 0.621 1.06 0.302
14.0 19.5 23.8 16.6 4.73 12.4 3.55
2.60 3.03 0 358 2.62 0.750 1.67 0.477
x
k p/ k $ / a
(M-1
102,
M
I
aeo-1)1/9
1.78 1.94 2.02 2.18 1.83 1.97 2.02 Av 1.96 Std dev 0.13
is taken to be 2.20 X lo-* sec-l.
a f'kd
The photochemical experiments utilized the 366-mp mercury emission, and since ABN is the only component of the reaction mixture which absorbs light a t this wavelength, it is assumed that kp/kt'/' is independent of the mode of initiation. For the photo process Ri = 2.30310& [ABN] where 1 0 is the incident light intensity, 1 is the optical pathlength, B is the decadic extinction coefficient (equal to 9.7 at 366 mp'), and @ is the fraction of radicals produced per quantum of light absorbed. The results of experiments utilizing photoinitiation of the thiol-TEP reaction are summarized in Table 111.
Table 111: Evaluation of 9, a t 25' in Benzene for the Photodecomposition of ABN Utilizing the 1-Butanethiol-TEP Reaction"
(5)
[RSHlo
M
ITEPlo, M
[ABNlo x 106, M
1.63 2.19 1.23 1.21 1.92
0.158 0.174 0.179 0.152 0.162
2.68 2.68 2.68 2.68 2.68
x
(6)
2R "t, products
(7) A possible dimerization involving thiyl radicals to form products such as disulfides is inconsistent with the observed rate law. The point of these experiments was to evaluate kp/ktl/' so that subsequent rate experiments utilizing photoinitiation could be used to find photochemical initiation rates. Table I1 summarizes the results of rate experiments utilizing thermal initiation, The average error in kp/ktl/' found here is about 5% where TEP and ABN concentrations are varied about sixfold and the thiol concentration about tenfold. Apparently, therefore, eq 2 is valid either for photo- or thermal initiation, and this finding certainly enhances the credibility of the mechanism proposeda for the reaction. Also, we have here the unusual situation that the scavenger reaction and the supposed initiation process for the chain reaction are identical, and given by eq 4. Thus, whether or not variability in scavenger efficiency exists is irrelevant for the present system.
[ABN lo
M
[TEPlo x IO? M
[RSHlo x 102,
10%
Av Stddev a
0
0.41 0.38 0.32 0.38 0.35 0.37 f0.03
Calculations amume kp/kt'/2 = 1.96.
Smith and Rosenberg' have already found that the quantum yield for photodecomposition of ABN is 0.47. Thus, the fraction of radicals produced per decomposition which subsequently initiate the chain reaction is f = 0.78. This is probably a much larger efficiency than one might have predicted for this reaction and is (10) J. P.Van Hook and A. V. Tobolsky, J. Amer. Chem Soc., 80, 779 (1968). (11) C. Walling and V. P. Kurkov, ibid., 89, 4895 (1967). (12) D.J. Carlsson and K. U. Ingold, ibid., 89, 4885 (1967). Volume 73, Number 8 August 1969
2702
R. D. BURKHART AND J. C. MERRILL
over a factor of 2 larger than that obtained previouslya using methyl methacrylate as the scavenger. Although one might suspect that this large value off results from a photoinitiation process which does not involve ABN, it has been found3 that no detectable reaction occurs when a degassed solution containing thiol and TEP is irradiated in the absence of ABN.I3 It should be pointed out that Rij defined in terms of the rate of production of initiator radicals, is d[A.]/dt = 2Ri and kt is defined by
(8)
-d [R. ]/dt = 2kt[R*l2
(9)
Experiments similar t o those described above utilizing benzene as the solvent were also carried out in cyclohexane. The results of thermally and photochemically initiated reactions are summarized in Tables IV and V.
Table IV: Values of kp/kt'/z for the Thermally Initiated Thiol-TEP Reaction in Cyclohexane at 25'
[RSHlo" x
106,
[ABNlo
x
[TEPlo,
M
M
4.64 4.64 3.22 3.22
0.708 0.708 0.708 0.708
104,
M
kp/ktJ2b
1.38 6.90 3.02 0.91
21.1 27.8 22.2 28.3 24.9 ~k3.2
Av Std dev a
1-Butanethiol.
a
Assuming
f'kd
=
2.03 X IO-s see-'.
_.
Table V : Evaluation of 0 a t 25" in Cyclohexane for the Photodecomposition of ABN [RSHlo'
x
108,
[ABNlo
[TEPlo,
x
106,
M
M
M
1.96 2.90 2.71 2.62 2.95
0.944 0.708 0 708 0.944 0.708
1.24 0.62 2.48 0.55 1.48
I
ipb
Av Stddev
' i-Butanethiol.
' Assuming ICp/kt'/2
=
0.22 0.25 0.29 0.36 0.34 0.29 f0.05
24.9.
It will be noted that there is a considerable influence of solvent on the over-all rate constant for the chain reaction, in fact, more than a factor of 10 increase. Assuming 0.47 to be the correct quantum yield for ABN decomposition in cyclohexane, one finds that the fraction of radicals which initiate chains per decomposition is 0.62 in this solvent. Although the precision is not as good in cyclohexane, it is still possible to conclude The Journal of Physical Chemistry
that both the thermal and photochemical efficiencies are smaller in cyclohexane and that in both solvents the photochemical efficiency is larger than the thermal. Since cyclohexane has a higher viscosity than benzene (0.89 cP vs. 0.61 CPat 25"), the lower efficiency in the former solvent seems reasonable and is in general agreement with the findings of Booth and Noyesl4 in their study of iodine atom reactions in various solvents. The larger efficiency found for processes involving light absorption is somewhat unusual for decompositions of azo compounds although there have apparently been no direct comparisons made under the same conditions as those used here. Nelsen and Bartlett,lB however, have investigated the thermal and photochemical decomposition of azocumene, and from their data it is possible to estimate values off' and f at 25". They are 0-70and 0.62, respectively; that is, the thermal efficiency is slightly greater than the photochemical. Data of Hammond and Foxls on the decomposition of ethyl 2,2'-azobisisobutyrate (EAB) in CCI, and chlorobenzene suggest that thermal and photochemical efficiencies are probably not too different for that system at 25". It has been inferred from these results that both processes involve decomposition from a singlet state. The fact that sensitized photolyses of both azocumene and EAB using triplet sensitizers yield the same f values as are found by direct phot~lysis'~~" has been interpreted to mean that the spin relaxation process is fast compared to cage recombination. If this interpretation is correct, then the similarity in thermal and photochemical efficiencies gives no information about the possible intermediacy of a triplet state in the direct photoprocess. It would be tempting to attribute the greater efficiency for the photoprocess in ABN decompositions to intersystem crossing prior to dissociation. Bartlett and Engel, l8 however, have found that the sensitized decomposition of azo-2-methyl-2-propane is mainly due to the transfer of singlet excitation, and it is quite possible that a similar situation holds for ABN. A practical result of the present experiments is that it is now possible to obtain more reliable quantitative data for the thiol-TEP reaction. Using results from Tables I11 and V it is found that kD/ktl/' should be revised downward from 0.81 t o 0.53 (M-I sec-l)'/* for the CYtoluenethiol-TEP reaction and from 2.24 to 1.46 (M-I sec-') 'Iz for the 1-pentanethiol-TEP r e a ~ t i o n . ~Sirnilady, kt values involving benzyl and pentyl radicals are decreased somewhat to 1.8 X los and 4.3 X los 1 M - I
(13) This test was more recently repeated with the same result a8 that reported in ref 3. (14) D. Booth and R. M. Noyes, J . Amer. Chem. SOC.,82, 1868 (1960). (16) S. F. Nelsen and P. D. Bartlett, ibid., 88, 143 (1966). (16) G.S. Hammond and J. R. Fox, ibid., 86, 1918 (1964). (17) J. R.Fox and G. S. Hammond, ibid., 86,4031 (1964). (18) P.D.Bartlett and P. S. Engel, ibid., 90, 2960 (1968).
2703
RADICAL-RADICAL REACTIONS IN DIFFERENT SOLVENTS sec-', respectively, One finds, therefore, that these recombination processes are even further removed from the diff usion-controlled limit than was previously supposed. This point is more thoroughly explored in connection with studies of the effects of solvents on
Acknowledgment is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society, for support of this research.
kt.19
(19) R. D.Burkhart, J. Phya. Chem., 73, 1741 (1969).
Acknowledgments.
Radical-Radical Reactions in Different Solvents. Propyl, Cyclohexyl, and Benzyl Radicals by R. D. Burkhart Department of Chemistry, University of Nevada, Reno, Nevada
(Received January I S , 1969)
Rate constants for radical-radical reactions involving the propyl, cyclohexyl, and benzyl radicals have been determined in both cyclohexane and benzene. Steady-state radical concentrations were monitored using the thiol-triethyl phosphite reaction, and rate constants for the reaction R . RSH --e RH RS were also obtained. Both k ~+ .R. and k ~+ .RSH when measured in cyclohexane are larger than or equal to those found in benzene. The solvent effect is greatest for the propyl radical and essentially nonexistent for the benzyl radical with cyclohexyl intermediate in behavior. The results can be rationalized by using the assumption that the alkyl radicals interact with benzene to form a complex species.
+
Introduction I n recent years there has been considerable activity in the measurement of rate constants for hydrocarbon radical recombination1 reactions in solution. Some of these studies have utilized spectroscopic methods2 making use of esr signals or uv absorption to measure the decay of radicals produced by pulsed radiolysis or flash photolysis. Rotating sector experiments have also been employedaf4in making these measurements. A rather sizable body of information has, therefore, been built up concerning radical-radical reactions in solution. Many of the available data have been obtained a t ambient temperatures, and the solvent most often used is cyclohexane, although an exception was the work on the recombination of pentyl and benzyl radicals in b e n ~ e n e . ~This study utilized rotating sector measurements carried out on the radical chain reaction between triethyl phosphite (TEP) and the appropriate thiols and yielded a kt vadue of 1.0 X lo8 M-l sec-1 for the recombination of pentyl radicals. This value seemed to agree reasonably well with results obtained in cyclohexane by other workers; however, recent studies on rates of initiation for the thiol-TEP reaction6 show that earlier estimates of this rate were probably too low. The recalculated kt for pentyl radicals in benzene was
+
found to be, 0.43 X lo9 M-l sec-I. This number is considerably smaller than rate constants for similar radicals obtained by others in cyclohexane and suggests that there is either a significant solvent effect in these recombination reactions or else there is some problem associated with the thiol-TEP monitor reaction. Because it is a natural part of a continuing study on the kinetics of radical recombinations and because of this basic disagreement with the results of others, it was decided to investigate in more detail the comparison of kt values in benzene and cyclohexane. The particular species chosen for study were propyl, cyclohexyl, and benzyl radicals representing unpaired electrons a t primary and secondary carbon atoms and a delocalized system. (1) In most of the processes to be discussed here, products may result either from combination or disproportionation reactions. In what follows, the term, recombination, shall be used to describe these reactions and the symbol kt is used for the rate constants. (2) (a) R. J. Hagemann and H . A. Schwarz, J . Phys. Chern., 71, 2694 (1967); (b) E. J. Burrell, Jr., and E'. K. Bhattacharyya, ibid., 71, 774 (1967); (c) M. C. Sauer, Jr., and I. Mani, ibid., 72, 3856 (1968); (d) S. Weiner and G . S. Hammond, J . Amer. Chem. SOC., 90, 1959 (1968). (3) (a) R. W. Fessenden, J . Phya. Chem., 68, 1508 (1964); (b) D. J. Carlsson and K. U. Ingold, J. Amer. Chem. Soc., 90, 1055 (1968). (4) R. D.Burkhart, ibid., 90, 273 (1968). (5) R.D.Burkhart and J. C. Merrill, J . Phys. Chem., 73,2698 (1969). Volume 7S,Number 8 August 1969
