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The Gas Phase Photolysis of Acetone at 3130 A in the Presence of Hydrogen Bromide. A Study of the Primary Photochemical Decomposition Processes of Acetone
by Carl W. Larson and H. Edward O’Neal’ San Die00 State College, San Diego, California
(Received December $7, 1966)
The effect of hydrogen bromide on the photochemical excitation of acetone a t 3150 A is described. Two primary decomposition processes are proposed : (1) spontaneous decomposition from upper vibrational energy levels of the acetone triplet state and (2) decomposition of thermally equilibrated triplet acetone molecules in the low-pressure (or falloff) region. Chemical quenching of the acetone triplet state by hydrogen bromide sometimes leading to the formation of isopropyl alcohol is also proposed. Quantitative estimates from quantum yield data are shown to give for the spontaneous decomposition (RRK) rate constant @(e) = veAst’R[(e - e O ) / e ] a - ’ ) , veAst/R = 10’5.’ sec-’, co = 17 kcal/mole, and s = 18; for the low-pressure thermal decomposition rate constant, 3 k d = 1012.2sX 10-9.95/e M-I sec-l; and for the triplet trapping rate constant, kT = X M-‘ sec-’. A mechanism for the formation of the unusual products observed (e.g., isopropyl alcohol, isopropyl acetate, and isopropyl bromide) is presented. The negligible yields of acetaldehyde and the regeneration of hydrogen bromide are also rationalized by semiquantitative calculations.
Introduction In an earlier study of the gas phase photolysis (3130 A) of acetone in the presence of hydrogen bromide, Steacie% reported that very small pressures of hydrogen bromide ( i e . , less than 0.1 mm at 25”) were sufficient to reduce ethane quantum yields to negligible proportions, trap almost all methyl radicals as methane, and reduce quantum yields of carbon monoxide to very low values. As in the acetone-hydrogen chloride photochemical system,2b hydrogen bromide (like HCI) was not used up in product formation but rather was regenerated in the course of the reaction. Although product analysis was confined to “noncondensable” gases (CHI, C2Hs, and CO), Steacie assumed trapping of the acetyl radicals by hydrogen bromide to produce acetaldehyde. The corresponding reaction of acetyl radicals with hydrogen iodide has since been well established.8 Initial results in these laboratories on the products of the acetonehydrogen bromide photochemical system confirmed Steacie’s findings with
regard to the very effective trapping of methyl radicals by hydrogen bromide and the low CO and C2He quantum yields. However, the “condensable” products of the photolysis contained almost no acetaldehyde. Surprisingly, the major condensable products were isopropyl alcohol, isopropyl acetate, and isopropyl bromide. In addition, the decomposition quantum yields of acetone were found to be strongly dependent on hydrogen bromide pressure and on temperature. These unusual results stimulated the further moderately extensive study of the acetone-hydrogen bromide photochemical system reported here. Although the nature of the secondary reactions was of considerable interest and has been fairly well illucidated, the most important results of this study (1) On leave of absence at Stanford Research Institute, Menlo Park, Calif. (2) (a) E. W.R. Steacie, Can. J . Chem., 33, 383 (1955); (b) R. J. Cvetanovi6 and E. W. R. Steacie, ibid., 31, 158 (1953). (3) E. O’Neal and S. W. Benson, J . Chem. Phys., 36, 2196 (1962).
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concern the nature of the primary processes in the acetone photolysis. This is a subject which has received considerable a t t e n t i ~ n but , ~ nevertheless is one which has remained to a large degree both qualitatively and quantitatively unresolved.
Experimental Section Reagent grade acetone obtained from Matheson Coleman and Bell was distilled several times under vacuum and stored over calcium chloride. Gasliquid partition chromatography (glpc) indicated no low-boiling (Le., bp = ('~)-'/[1
+ kiC/%d(M)]
We will show in a future communication that the phosphorescence data of Heicklen and Noyes' and of Luckey and Noye@ are consistent with an activation energy for the triplet decomposition between 8 and 10 kcal/ mole and are, therefore, in good agreement with the value of 9.45 kcal/mole accepted here. From the d the above data, one obtains k T = ratio k ~ / ~ k and 109.25*0.65 x 10(-0.35*1-1)/0 &,-' sec-l. Both rate constants are reasonable. The A factor of the decomposition reaction is greater than collision frequencies, which is in accord with expectation for a unimolecular reaction in the low-pressure region. Error limits in the Arrhenius parameters of the trapping reaction are rather large; however, the over-all rate compares very favorably with that measured for the methyl plus hydrogen bromide reaction. It is clear, by reference to the activation energy obtained for the decomposition reaction, that the activation energy for triplet trapping by HBr must be small (ie., ET 0-1.5 kcal/mole). A higher activation energy would result in an unreasonably high A factor for k ~ . As a compromise and by analogy with reaction 1, we suggest a most reasonable value for ' k ~ ,N109.6 x IO-'.'/@ M-' sec-1. Quantitative estimates of the rates of spontaneous decomposition from high vibrational levels of the acetone triplet state relative to collisional deactivation can be obtained from the limiting decomposition quantum yields. With reference to the mechanism
(XII)
From (XII), using the relation kic/'kd(M) = (1
-
(19) G. W. Luckey and W. A. Noyes, Jr., J . Chem. Phys., 19, 227 (1951).
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and with the assumption of steady-state conditions, one obtains for the quantum yield of spontaneous decomposition (XIII)
4.4
-a -
Thus
4.0
xu m
3.6 3.2 2.8
Values obtained for k(e)/(ZM) at the various temperatures are listed in column 6 of Table V. For a collision diameter of 5 A and a total pressure of 70 mm, the ratio at 44" corresponds to a spontaneous decomposition rate of k(e) = 3.5 X lo7 sec-l and to a lifetime of the vibronically excited species of about 3 X lo4 sec. One can see that the spontaneous decomposition rate constants are strongly temperature dependent. This is a most interesting result because it implies that quite small variations in total energy content are sufficient to cause large variations in decomposition rates. The RRK expression for a spontaneous rate constant is k(e)
=
v e A S t E/ R (Eoy )s-l
3 .O
I / T x i03-
3.2
3.4
OK-'
Figure 6. Arrhenius plot of the pseudo-unimolecular rate constant for triplet decomposition obtained from phosphorescence lifetime data of Groh, Luckey, and Noyes.17
REACTION COORDINATE
Figure 7. Energy diagram relating t'he triplet state to the ground electronic state of acetone.
For our system, the following identifications can be made (see Figure 7)
temperature range of 44 to 126'). Rebbert and Aus~ O O S in , ~ a~ study (at 3130 A) of the triplet state quenching of perdeuterioacetone phosphorescence by azoe = (H"-- H " 0 ) y i b hv €0 - D(CHTCOCH~)~OKalkanes, in agreement with our finding, have observed strongly temperature dependent limiting quantum In the above, ( H " - H " 0 ) v i b is the average vibrational yields of carbon monoxide. The magnitude of Rebbert energy content of acetone in the ground electronic and Ausloos' observed temperature dependence was state, hv is the energy of the light absorbed, eo is the very similar to ours, although their limiting decomcritical energy for dissociat,ion from the triplet state position yields were lower. Thus, at 30 and 133", in its high-pressure region and D(CH~-COCH~)O.K Rebbert and Ausloos found $Id* 'v 0.025 and 50.16, is the bond dissociation energy of acetone at 0°K. respectively (an increase of a factor of 6.4). If reasonIt is clear, since (H - HoO)vit, depends on temperature, able estimates of e, eo, s, (H" - H"o)vir,,and D(CH3that t and therefore k(e) should be both temperature COCH3) are made, one can show that the variations and wavelength dependent. Kirk and Porterm in a of 4 d * with temperature as observed in this study are study of the photooxidation of acetone at 2800 of the correct order of magnitude. For this purpose, and 3130 A have observed the latter. Their quantum the following values have been selected: (1) s 'v 18. yield data were consistent with &* 5 0.16 at 3130 The six C-H stretches have been excluded. (2) 1.0 at 2800 A. The magnitude of the A and &* D(CH3-COCH3) = 83 kcal/mole. This follows from temperature dependence is far more difficult to assess. the recent heats of formation of LWt"(CH3) = 34 kcal/ Heicklen and Noyes,' using biacetyl as triplet quencher, molez2 and of AHf"(CH&O) = -3.0 k ~ a l / m o l e . ~ ~ ~ ~ have reported a temperature-independent limiting decomposition quantum yield of &* = 0.22 f 0.06 over (20) A. D. Kirk and G. B. Porter, J . Phys. Chem., 66, 556 (1962). (21) R. E. Rebbert and P. Ausloos, J . Am. Chem. SOC.,87, 1847 the temperature range 40-70" at 3130 A. Their error (1965). limits are large enough, however, to accommodate the (22) D. M. Golden, R. Walsh, and S. W. Benson, (bid., 87, 4053 temperature dependence of &* that was obtained in (1965). this study (i.e., +d* varied from 0.08 to 0.32 over the (23) J. A. Kerr and J. G. Calvert, J . Phys. Chem., 69, 1022 (1965).
+ +
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GASPHASE PHOTOLYSIS OF ACETONE
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-
(3) eo N 17 kcal/mole. The activation energy observed for the triplet decomposition, 3Ed 9.45 kcal/ mole, is for the low-pressure (or falloff) region. In the classical RRK theory, the activation energy in the highpressure region could be larger than this by (s - 3/2) X RT 'v 11-12 kcal/mole. By analogy with molecules of similar size (e.g., cyclobutane or cyclopropane) a 7-8 kcal energy difference between the high- and lowpressure activation energies is not unreasonable. (4) (H"- H " 0 ) v i b = 1.74 (44"); 2.37 (96"); 2.76 (126"); 3.12 (150"). (5) hy = 3150 A in this study 1 : 90.76 kcal/mole. Rotational and translational contributions to the enthalpy have been subtracted from the heat contents calculated and tabulated for acetone by Pennington and Kobe.24 The remainders are the numbers above and represent the energy above zero point which reside in the vibrational modes of the ground-state molecule. Using these numbers, ratios of the spontaneous rate constants referred to the rate constant at 96" have been calculated. The results are
These numbers are to be compared to those in parentheses which were obtained from the experimental k(e)/(ZiLI) ratios given in column 6 of Table V. The agreement is good considering the experimental errors involved and the assumptions made. The predicted variation is an increase of a factor of 4 from 44 to 150" while the observed temperature dependence is stronger showing an increase of a factor of 6.3. Very slight adjustments in the suggested variables could easily reproduce the observed variation [e.g., a 0.2-0.5 kcal increase in D(CHrCOCH3) would be sufficient]. It should be mentioned that the approach presented does not account for the sizable increase of the acetone extinction coefficient with temperature. Such an increase implies that the effective (H" - HoO)vib energy carried to the excited state by absorption should have a stronger temperature dependence than the values used. Corrections of (H" - H " 0 ) v i b for this effect would be in the required direction to account for the stronger temperature dependence observed for k(e). The general trend of the data is quite clear. Spontaneous decomposition rates are temperature dependent and the magnitudes of their dependence are reasonable. We will show in a following communication that the parameters used to calculate the
spontaneous decomposition rate constant ratios are also reasonably consistent with wavelength and pressure studies of the acetone phosphorescence. This latter fact has influenced our choice of s and eo, which admittedly do seem somewhat too high. The same numbers for E, eo, and s can be used to obtain an estimate of the frequency term in the spontaneous decomposition rate constant. Thus
If a recombination rate for CH3 and CH3C0 radicals of about 109.5M-' sec-' is assumed, along with entropies of 47 and 64.6 eu, respectively, an A factor of about 1016J sec-' for the thermal decomposition of acetone is obtained. This is higher than the spontaneous decomposition A factor by nearly a factor of 30. Thus, there appears to be an appreciable entropy increase (in addition to electronic) in going from the acetone ground state to the first triplet state. The low A factor for spontaneous decomposition is between those normally observed for methyl radical fissions from molecules (-1016-1s sec-l) and those observed from radicals (-lOI3-l4 sec-l) and therefore is not unreasonable. Some of the apparent discrepancies in the absolute values observed for the limiting decomposition quantum yields by Heicklen (&I* 0.20 between 40 and 70") and by Rebbert and Ausloos (&* N 0.25 and 2 0 . 1 6 a t 30 and 133") can be rationalized. The magnitude of k ( e ) is extremely sensitive to the exact wavelength of the light absorbed. Although Heicklen used chemical filters similar to our own, he operated with a medium-pressure mercury arc which would have had less pressure broadening than our high-pressure arcs. It is quite possible, then, that his radiation was distributed closer to the 3130-A mercury line. Since 3130A radiation would produce excited-state acetone molecules more energetic than ours (by about 0.7 kcal/ mole) Heicklen's limiting decomposition yields at 40" should be more analogous to ours a t 96", as observed. Rebbert and Ausloos used a third-order interference filter with radiation a t 3130 A, but performed their quenching studies with perdeuterioacetone pressures about 2.5 times larger than ours. A calculation of the expected limiting decomposition quantum yields in their system, using the suggested values for k ( ~ and ) for 2, gives r$d* (30") 0.034 and +d* (133") 0.09.25 The fact that acceptable agreement is obtained when the total pressures of the two studies were so
-
-
-
(24) R. E. Pennington and K. A. Kobe, (1957).
J. Am. Chem. Soc.,
79, 300
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widely different is particularly significant since it tends to confirm the competition between spontaneous decomposition and collisional deactivation proposed in the mechanism. This competition is also evident in the reduced spontaneous decomposition quantum yield (by a factor of 2) of run 93 in which a 36-cm pressure of SFe was added. Further studies on the effect of total pressure on both decomposition processes are needed and will soon be underway. An examination of the energetics of the tripletHBr trapping reaction provides a final point of interest. Our numbers place the energy of the acetone triplet ground vibrational state at about D(CH&O-CH3) 17 kcal, s 6 6 kcal above the acetone ground state. This roughly corresponds to the maximum observed in the acetone phosphorescence (-4500 A).19 Thus, the maximum could correspond to an 60-o emission from the triplet to the ground state. The st,rength of the A bond in acetone may be estimated (from the difference between the consecutive 0-H and C-H bond dissociation energies in isopropyl alcohol to produce acetone) at about 75 kcal. This latter value corresponds roughly to the shortest wavelength observed in the fluorescence emission of acetone (-3850 A)26 and could also correspond to an transition from the upper singlet to the ground state. The trapping reaction of the triplet is about 5.5 kcal exothermic.*' It is surprising, then, that the activation energy of the HBr trapping reaction is as low as 1 kcal since the data
The JoUTnd of Phyeical Chemistry
CARLW. LARSON AND H. EDWARD O'NEAL
of Fettis and Trotman-Dickenson28 for alkyl radical reactions with HBr would imply that the trapping reaction should be 14-17 kcal exothermic for such an activation energy. The reason for this discrepancy is not clear. It may mean that the usual Polanyi relation should not be used to compare the chemical behavior of a triplet state with that of normal free radicals.
Acknowledgment. The authors wish to express their appreciation to the National Science Foundation for financial support of this research. Many interesting and instructive discussions of this work with Professor S. W. Benson are also appreciatively acknowledged. (25) These calculations are for k(e) of acetone and not deuterated acetone. However, the behavior of the deuterated acetone should be similar. Comparisons of the differences in the six vibrational modes being lost in the dissociation (using CZHSand CzDs as analogous) lead to an estimate of D(CDs-CoCDa)o"K = 83.7 kcal/ mole. The lower vibrational modes resulting from the deuterium substitution would tend to increase the temperature dependence of (Ho- X"0)vib. Both effects are in the proper direction to produce even better agreement between calculated and experimental C$d*'S. (26) R. E. Hunt and W. A. Noyes, Jr., J. Am. Chem. SOC.,7 0 , 467 (1948); G. W. Luckey, A. B. F. Duncan, and W. A. Noyes, Jr., J. Chem. Phys., 16, 407 (1948). (27) This is estimated relative to ground-state acetone. Thus, in the trapping reaction, acetone containing 66 kcal of triplet energy forma an 0-H bond worth about 102 kcal and breaks a T bond and an H-Br bond at roughly 75 and 87.5 kcal, respectively. (28) G. C. Fettis and A. F. Trotman-Dickenson, J. Am. Chem. SOC., 81, 5260 (1959).
