ZE’EV B. ALFASSIAND DAVID AT. GOLDEN
3314
Dissociation Energy of the C-H Bond in Ethanol. inetic Study of the Reaction I,
+ Ethanolla
by Ze’ev B. Alfassilband David M. Golden” Department of Thermochemistry and Chemical Kinetics, Stanford Research Institute, Menlo Park, California 9.4026 (Received M a y 30,197’9)
The reaction of iodine with ethyl alcohol has been studied in a static systeui in the temperature range 203-252’ and in the pressure range 30-750 Torr. The rate-determining step of this reaction has been found to be I 1 CH3CH20H CH,CHOH HI and k~ is given by the equation log ( h / N sec-I) = (11.5 ~ r 0.3) . - (22.7 f b).Bj/B, 0 = 2.303RT kea1 mol-’. From the activation energy and the assumption E-1 = 1 i 1 kea1 mol-1, it has been calculated that D[CH&H(OH,-H] = 93.0 k 1.0 and aHr’(CH&HOH,g) = -15.2 rt 1.0 kcal mol-‘. The data of the reaction I1 CH30Hare reinterpreted, increasing log A by 0.3.
-
+
+
+
Introduction The first values of the bond dissociation energies
(BDE’s) of C-H bonds in alcohols were measured by Whittle and coworker^^*^ using photobromination techniques. They found D[CHz(OH)-H] 5 92 kcal/ mol and D[CH&H(OH)-H] 5 90 kcal/mol. Two years later Walsh and Benson4 studied the reaction of iodine with 2-propanol and found D [(CH,),C(OH)-K] = 90.3 f 1.1kcal/mol. The similarity of the strengths of the tertiary C-R bonds in isobutane and in 2-propanol led them to suggest that “the inductive effect of hydroxyl group is the same as that of a methyl group with respect to homolytic bond breaking a t the a-carbon position.’’ This suggestion would lead to D [CHz(OH)--H] = 98 keal/mol and D[CH3CH(OH)-H] = 95 kcal/mol. Cruickshank and Benson5 have measured the BDE of the C-€€ bond in methanol and found D [CIIz(OH)-H] 2 95.9.5p6 These discrepancies lead to the present situation in which the C-H bond strengths in alcohols are not too well known. I n order to resolve these discrepancies, more data are needed. The kinet-. ics of gas-phase iodine atom reactions have been found to be a reliable murce of bond dissociation energies! This work deals with the reaction of iodine with ethanol in order to measureD(CH&H(OH)-HI.
E x ~ e ~ i m e $~ tea~~ t i ~ ~ Iodine (Jlallinkrodt) was resublimed and degassed at liquid nitrogen temperature before use. Ethanol (Kossville, Gold Shield) and acetaldehyde (n4atheson Coleman and Bell) were distilled and found to be more than 99.9% pure by gas chromatography (F and M, Model 810)) wing a 25-ft column of 20% Carbowax, 20 AI, on Chromosorb, MT. AW 60-80 mesh. The Cary 15 spwtrophotometer modified for use with a heated reaction vessel has been described previously.’ Iodine was introduced into the reaction vessel at the desired pressure, first, and the spectrophotometer was The Journal of Physical Chemistry, Vol. 76, N o . 23, 197,9
set to record the absorbance as a function of time at a wavelength between 502 and 440 mp, according to the I, pressure. The wavelength was chosen such that the initial absorption was 1.7-1.9 OD, except in one experiment where the amount of 1, was too low and the absorption a t the peak was 0.95 OD. I n some of the experiments the absorptions a t 270 and 235 my were also measured in order to measure the amount of H I formed. I n order to minimize surface effects, the vessel was coated with Teflon by polymerization of 150-200 Torr of C2F4 catalyzed by 1-5 Torr of tert-butyl peroxide a t 185’.8 Initial rates were measured by following the rate of disappearance of 5-10% of the original iodine absorption.
Results The initial rates of the reactions of iodine with organic compounds have been foundBto obey the equation d(lz) - k , ~ ( 1 (RH) ) = kobsdK,z1’2(12)1’2(RH) (1) dt where R H is the organic compound and XI, is the equilibrium constant for the dissociation of iodine molecules (1) (a) This work was supported, in part, by Grant No. AP 00353-08 from the Environmental Protection Agency. (b) Pofitdoctoral Research Associate. (2) (a) E. Buckley and E. Whittle, Trans. Faraday Sac., 58, 529 (1962) ; (b) E. Buckley and E. Whittle, ibid., 58,536 (1962). (3) A. M. Tarr and E. Whittle, Trans. Faraday SOC.,60,2039 (1964). (4) R . Walsh and S. W. Benson, J . Amer. Chem. Soc., 88, 3480 (1966). (5) F. R. Cruickshank and S. W. Benson, J . Phvs. Chem., 73, 733 (1969). (6) D. M. Golden and 9. W. Benson, Chem. Rev., 69, 125 (1969). (7) D. M. Golden, R.Walsh, and 8 . W. Benson, J . Amer. Chem. SOC., 87,4053 (1965); R. Walsh and S. W. Benson, ibid., 88,3480 (1966). (8) R. K. Sally, D. M. Golden, and S. W. Benson, I n k J . Chem. Kinet., 2 , 3 8 1 (1070).
BONDDISSOCIATION ENERGY OF THE C-H BONDIN ETHANOL
12
+M
XI,
21
+M
2oo
3315
3 Packed Vessel (Teflonized)
A
The results for the reaction of iodine with ethanol given in 'Table I show that this reaction is governed by the same rate law, since sixfold changes in iodine pressure and more than tenfold changes in ethanol pressure give the same kobsd (calculated according to eq 1) within experimental error. The dependence of k o b s d on the temperature in the range 203.2-252.0" is given by log (kohsd/l. mol-l sec-l) = (11.5 f 0 . 3 ) - (22.7 f 0.6)/0, where e equals 2.303RT kcal/mol, and the quoted errors are standard deviations. This Arrhenius plot is shown in Figure 1.
a Packed Vessel (Unteflonized)
Table I : The Observed Rate Constants for the Reaction of Iodine with Ethanol: k&d = ( -d(Iz)/dl)i,itiai/(I)(C2H6 OH) Run no.
Temp, O C
(Id, Torr
1 2
203.0 203.0 203.2 203.2 213.5 213 5 2,13.5 213.5 225
21.0 24.1 22.8 20.1.
3 4
5 6
7 8 9 10 la 12 13 14
15
I
236.5
2.62 5.6 16.7 21 .e 24.5 3.9
5.71 11.8
18
30
21.8
17.8
16 17 19 20 21 22 23 24 25 26 27 28 29
5.85 15.3 20.4
252
212.8" 244a
215b 217b
17.1 25.6 8.65 13.4 18.6 22.4 19.5 20.9 96.6 18.4 22.4 14.5 18.7
689 425 743 689 535 362 122.5 629 579 540 527 38 27.4 603.5 633 494 367 96 447 354 169 195 277 303 536 185 280.5 146 519.5 221
1 . 6 6 X lo4 1 . 6 6 X lo4 1.64 x 104 1 . 6 4 X lo4 1.089 X lo4
7.03
x
103
4 . 5 9 X los
2.74 X loa
1.10 x 104 3.55
x
103
1 . 0 2 5 X lo4 9 . 4 0 X loa
a Measurements done in packed vessel. in uncoated packed vessel.
1.185 1.20 1.09 1.12 1.55 1.71 1.55 1.84 2.84 2.81 2.85 3.05 2.78 3.13 4.72 4.25 4.44 5.55 5.00 9.9 10.9 10.8
11.1 2.01 1.66 7.15 6.30 7.75 1.76 2.45
Measurements done
1O~IT
Figure -C
1.
Arrhenius plot for the reaction 1
CHsCHOH
+ HI.
+ CHiCHzOH
surface catalyzed the disappearance of 1%. I n order to check if the difference in the surface effects is caused by the different alcohols or by the fact that our surfaces have been Teflonized, the Teflon coating was decomposed by heating to 400" (the decomposition is verified by the decrease of the absorbance (scattering) of light in the reaction vessel). Two measurements were done in the un-Teflonized packed vessel. The results of these measurements (cf. Table I and Figure 1) show that a quartz surface does not catalyze the reaction between Izand ethanol either.
Discussion Following the mechanism found for methanol,6 the scheme of reactions in the case of ethanol will be
I * + CII3CH20H
1
CH36MOH
+ HI
2
+ IzJ'CH3CMIOH 4- I CH,CHIOH A H I + CH,CHO CH3CH0 + I HI + C H 3 6 0 CH.360 + I2I_ CH3COI + I
CH3CH:OH
4
The possibiliLy of a heterogeneous contribution to the disappearance of I2was checked by measuring the rate of the reaction in a packed vessel in which the surface to volume ratio is about 14 times larger than in the usual vessel. The results of these measurements given in Table T and in Figure 1 show that the disappearance of 1, is homogeneous. I n the study of the reaction of IZ with methanol5 it was found that the
5
I n the case of methanol, it was assumed that, since reactions 3 and 4 are much faster than reaction 1, for every molecule of CH30H which reacts, two molecules of 1 2 will be consumed, and hence kobsd = 2kl. This assumption has to be considered carefully since, although the rate constant lc4 is larger than Jcl, the conThe Journal of Physical Chemistry, VoE. 16,No. $8,1978
ZE'EV B. ALFASSIAND DAVID M. GOLDEN
3316 centration of the alcohol is usually much higher than that of the aldehyde. The rate constant for the abstraction of hydrogen from acetaldehyde by an iodine atom was measured previously a t one temperatureg and the Arrhenius parameters were chosen by assuming a reasonable A factor. I n order to be sure of our comparison of Jil and k4, we measured kq for one of the temperatures for which kobsd was measured (225"); the results are given in Table IE, Comparison with the results of kobsd
about SOO), every molecule of formaldehyde formed would have reacted, and they did not consider the ratio of the concentrations of the aldehyde and alcohol. The rate constant of the reaction I f CEZOH has been found to be log (kl/l. mol-' sec-l) = 11.5 - 26/0 and of the reaction of iodine atoms with formaldehyde'O log k4 = 10.9 - 17.4/0, and hence, k4/kl = Hl-o.6+8.6/B. For runs 11-13 (the numbers are the numbers of the runs in the original paper5), 0 = 2,64 and k4/k1 = 102.". A similar expression to the one derived previously for ethanol will, in this case, be
~~
Table 11: The Reaction of Iz with Acetaldehyde at 225"
19.4 20.3
42.8 24.5
1.71 X 9.26
x
10-3
7.03 X lo3 7.03 x 103
1.98 X IO3 1.90
x
108
which is either kl (if reaction 3 or 4 do not occur), or twice kl (if these reactions are completed) shows that lc1l/k1is between 67 and 134. The initial rates are measured up to 10% change in the (IJO concentration and thus the concentration of the acetaldehyde, which is equal or less than the amount of the iodine consumed, is given by (CHLXIO) 5 O.l(Iz)o. The ratio of the rakes of reactions 4 an rate 4 5 h rate 1 kl
O.f(12)0
>(: --___
-< 13.4
(C2H60H)o
(12)o
(CzH8H)o
This value of 13.4 is a high upper limit including the assumption that k 0 b s d = 2kl and that the amount of acetaldehyde present is equal to the amount of reacted iodine. These are contradictory assumptions and a morc reasonable relakion is obtained by using the value 67 for the ratio of the rate constants. Thus
The results of Table I show that at 225" the ratio (12)0/(C2HaOH)o has beell varied in the range of 7.9 X lo-' to 4.5 X without any noticeable change in kohsd. For the cases of higher values of this ratio (7.9 X IO-' and 4.7 X almost all the acetaldehyde formed would have had to react and thus kobsd = 2k1, while for the rung with low (12)o/(C2H50H)o ratios only a very small fraction of the acetaldehyde would have reacted with iodine and kohsd = IC'. The fact that in both cases we get the same result for kobsd shows that kobad = kl in all the runs. This same conclusion can be reached for the rne1,hanol" experiments when the data are reevaluatpd. I n this latter study, the authors5 amumed that, since the reaction of I with formaldehyde (reaction 4 in their paper) is much faster than the reaction with methanol (the ratio of the rate constants is The Journal
of
Physical Chemistry, Val. 76, N o . 25, 1972
Thus, for runs 12 and 13 where the ratio of the pressures is about unity, reaction 4 would be much faster than reaction 1, and hence kobsd =: 2kI, since the overall reaction would be 212 f C€130R 4 CO f 4HI. Whereas for run 11, where the ratio (12)o/(CH30H)o= 1.5 X the rate of reaction 4 would be smaller than that of reaction 1 and kobsd < 21c1. The solution of the rate equations (Appendix A) shows that in run 11 for 10% consumption of I, only about 28% of the newly formed formaldehyde would have reacted, thus kobsd = l.28kl, while in run 12, 98.5% of the formaldehyde would have been consumed by 1 2 , and thus kobsd = 2k1. Yet the experimental results of kobsd for runs 41-13 are the same, within the experimental error, meaning that in both cases k o b s d = kl, and reaction 4 does not occur. I, are found for runs The same results for CH,OH 15-17. I n run 16, after consumption of 10% of the iodine, only 18% of the newly formed HCNO will have reacted, while in run 17, 78y0will have reacted, meaning that in run 16, Jiobsd = l , l S k l , and in run 17, kobsd = 1.78kl; yet the experimental results arc the same, leading to the conclusion that no iodine a t all has been consumed by the aldehyde. The rea8on that no iodine has been consumed by the aldehyde, although in the reaction of iodine with pure aldehyde iodine is consumed, is either that no aldehyde is formed or that the iodine consumed by the aldehyde is regenerated. Rollefson and Faull" found that in the reaction of iodine with acetaldehyde and propionaldehyde the I, disappeared in the first stcp, but at the end of the reaction the original amount of iodine is regenerated. These studies were done at higher temperatures (>300"), and the mechanism for this regeneration, proposed by O'Neal and Benson,12 does not predict the same phenomena a t 225". l\Iore than that, both Walsh and B e n ~ o n ,at~ 20S0, and the present authors, a t 225", found that the absorption of iodine
+
(9) R . Walsh and S. W. Benson, J . Phys. Chem., 70,3751 (1966). (10) R. Walsh and S.W. Benson, J . Amer. Chem. Sac., 8 8 , 4570 (1966). (11) G. K. Rollefson and R. F. Faull, i b i d . , 59,625 (1937). (12) E. O'Keal and S.W. Benson, J . Chem. Phys., 40,302 (1964).
BONDDISSOCIATION ENERGY OF THE C-H BONDIN ETHANOL
3317
reaches an equilibrium value and does not return to the This value is larger than the value found by Tarr and original absorption, at least in the time range of our Whittle3 by photobroniination studies ( 5 9 0 kcal experiments (2- 4 hr). Thus the remaining explanamol-l), but it is close to the value predicted by Golden tion is that the iodohydrin is not in a steady-state conand Benson as a result of both bromination and iodinacentration, and tEx concentration of the aldehyde is tion studies6 (-92 kcal mol-’). An interesting obserlo’wer than that calcula,tcd by assuming a steady state vation is that the iodination method gives higher reof the iodohydrin. This must mean that the iodohysults for the BDE of C-H bonds in alcohols than the , ~ the drin is more stable than previously a s s ~ m e dand photobromination method. A possible good way to lower limit of k3 which can be deduced from these recheck this discrepancy might be through a study of the sults is compared with a theoretical calculation in kinetics of thermal bromination, which may eliminate Appendix 11. The fact that CQ has been found to be some of the problems connected with photosensitizaone of the products5 and the existence of an accomtion. panying pressure increase do not contradict the conThis value of the bond dissociation energy can be ciusion concerning lack of formation of aldehyde in the used to calculate the heat of formation of the radical initial steps. The iodohydrin reaches its steady-state CH&HQH. concentration after the initial steps and then the aldehyde is formed. This sort of behavior has been observed in the reaction of iodine with methyl ethyl AHf” (H) - AHfO(CH3CHzOH) ketone,13 where in the initial step 2-iodo-3-butanone is formed, and only later does this intermediate decomThe heats of formation of a hydrogen atom and of the pose t o give methyl vinyl ketone by elimination of hyethyl alcohol molecule a t 298°K are 52.1” and --56.118 drogen iodide. kcal mol-I, respectively; thus the heat of formation of I n the case of ethyl alcohol, there is still another posthe hydroxyethyl radical a t 298°K is -15.2 i 1.0 sibility, namely, that the iodohydrin decomposes to kcal mol-’. water and vinyl iodide (instead of the hydrogen iodide and acetaldehyde), This possibility does not seem to Appendix A be favored by the calculation of its homogeneous rate Solution of the Rate Equations for the Reaction of constant (Appendix XI), but it might be the predomiIz with Methanol nant reaction as a result of heterogeneous catalysis. Some support for this possibility is the fact that in two I CH30H cH,OH HI runs carried out a t 271”, after the absorption of the iodine goes to a minimum value and remains a t this c H 2 0 H Iz CHJOH 4- 1 value for some time, it starts to increase again. A CHJOH --%CHzO HI mass spectrometric analysis of the mixtures shows, besides C,H5Q1-I, b,, and HI, the presence of water, CH~O I cao HI ethane, and ethylene. Ethane and ethylene have previously been found to be the products of the reaction of Assuming steady state for eH&H and CIIJOH gives vinyl iodide and hydrogen iodide. l4 The conclusion that in the case of the reaction of I2 (CHJOH),, = ?(I) (CH,OH) k3 \pith methanol, knbsd = kl and not kl = kobsd/2, as was postulated before, leads for that system to log kl = (13) R. K. Solly, D. M. Golden, and S. W. Benson, I n t . J . Chem. 11.8 - 26/9. Kinet., 2,381 (1970). The Bond Dissociation Energy, D(CH3CH(OH)-H). (14) A . 8. Rodgers, D. M. Golden, and S. W. Benson, unpublished The reaction of hydrogen iodide with radicals is usually results. (15) The values used for heat capacity are: I, Cp’298 = 5.0; C,’KOZ= assumed (and in Some cases has been proven6) to have 5.0 gibbs mol-‘ (S. W. Benson, “Thermochemical Kinetics,” Wiley, an activation energy of 1 f 1 kcal/mol, and thus New York, N. Y., 1968, p 195); CzHsOH, C P o m = 15.64; Cp’SOZ = 22.80; HI, cp0298 = 6.97, C,”SOZ= 7.11 (D. R. Stull, E. F. Westrum, AH(1) = I& - E-1 = 22.7 - 1 = 21.7 kcal/mol Jr., and G. C. Sinke, “The Chemical Thermodynamics of Organic Compounds,” Wiley, New York, N. Y., 1969, pp 228, 423); CHa‘This enthalpy is ior 502°K, the average temperature at CHOH, CPo298= 14.0, CPo~o2= 20.3 calculated by deduction from which the reaction has been studied. Since the average the value of ethyl alcohol, the values for the three niissing vibrations heat capacity is small, it can be neglected and AHzSs(l) (H-C-H scissors, C-C-H rocking, and C-H stretching) and for the hindered rotation (S. W. Benson, “Thermochemical Kinetics,” Wiley, = 21.7 kcal r n ~ l - ~ . The ’ ~ enthalpy of reaction is the New York, N. Y., 1968, p 44). AACpo = (0.42 - 0.34) /2 = 0.04 eu. difference between the BDE of the broken bond D(CH3(16) B. deB Darwent, “Bond Dissociation Energies in Simple Molecules,” NSRDS-NBS 31 (1970). CH(OH)-H) and that of the newly formed bond, (17) S. W. Benson, “Thermochemical Kinetics,” JViley, New York, D(H-1) thus
+
+
+ A
+ + 2- +
J
D(CH&H(OIf)-H)
+ D(H-I) = + 71.316 = 93.0 kcal mol-l
21.7 21.7
I=
N. Y., 1968, p 195. (18) D. R. Stull, E. F. Westrum, Jr., and C . Sinke, “The Chemical. Thermodynamics of Organic Compounds, Wiley, New York, N. Y . , 1969, p 423.
2~
The Journal of Physical Chemistry, Vol. 76,N o . $3, 197%
ZE’EV B. ALFASSIAND DAVIDM. GOLDEN
334.8 d(CH20) dt
-
k~(CHz10H)- kr(1) (CH20)
I
+
and for the init,ial kl
(CRz0) = -(CHBOH)(l kd The total amount of formaldehyde formed (both the amount present and the amount already consumed) is given by h(CH30H) (Ut
(C&O)+,otal
and the fraction of the formaldehyde consumed by the reaction with iodine F , is given by
F = (CH20)totai - (CHzO) - = I (CWz0)total = 1
I
-
+
For IZ CBOH, run 12, where we expect the reaction of the aldehyde to have the most effect, T I= 574,2”K, K~~ = 10-10.13 M , ( I ~ )= 3.1 x 10-4 M , (ROH) = 2.9 X M , and kl = 101.65J4-l sec-’, and thus ks - 10-4.30 sec-I. For I’L C2H50N,run 13, T = 498”K, K I ~ =~ 8.07 / ~ X lo-’ M’”, (12) = 7.0 X I l l , (ROH) = 8.85 X N , and kl = 2.9 X lo1M-’ sec-l, and thus IC3 = 10-5.105ec-l. The rate constant for the H I elimination from the iodohydrin can be calculated theoretically, as suggested by Cruickshank and B e n ~ o n . ~The A factor is calculated by O’Neal and Benson’s method19 for four- and -2 eu, six-center unimolecular elimination as AS’ The activation energy is calcuand thus A = lated from the activation energy for the reverse reaction, the four-center addition of HX to the carbonyl double bond, calculated by Benson and Haugen’s method,20and adding the endothermicity of the elimination reaction. The endothermicity is calculated from the known bond dissociation energies
(CHz0)
(0) total
e-ka(I)t
- -__--
h(I>t The maximum value for k&)t can be calculated by the upper limit of the measured disappearance of iodine, ie., A(12) = O.l(I& A(I2)
kl(CHaOH)(I).t
-
H
I I
R-C-OH
--+ R-C
h(CH30H) Appendix B Estimated Upper Limit Value for k3 and Comparison to the Theoretical Estimation The upper limit value of k3 can be estimated by assuming that, in the region where the initial rate law applies (A(1,) 0.1(12))>the iodohydrin reaches its stationary-state concentration. This can be calculated by assuming that the iodohydrin steady-state concentration is equal to the amount of the iodine which has been consumed (0.1(12)). The steady-state concentration of the iodohydrin is given by kl/k3(I) (ROH) (Appendix A], where (I) and (ROH) are the concentration of the iodine atoms and the alcohol, respectively. T’huE:the upper limit value of k2 is given by
-
O.l(Is)
kl -(I)(ROH) k3
The Joimnal of Physical Chemistry, Vol. 76, N o . 23,1972
0
+ HI
T
+
and
I
c
AH = D(0-H) D(C-I) - D(C=O, 7r bond) D(H-1) and using the values D(0-H) = 104.0, 105.0, D(C-I) = 53.5, 56.0, D(C=O) = 76, 73, and D(H-I) = 71.3 (the first value is for ethanol iodohydrin and the second is for methanol iodohydrin), thus AN = 10.2, 16.7 kcal mol-’, respectively. The activation energies were calculated by using the data in penson and HauEen’s paper, and (YCHGHO = 3.60 A3, CWH~O = 3.56 A3nz1, U C I ~ ~ C H O= 2.69OlZ2/.LCH,O = 2.330,*’ TO=E 0.96 A,23~c-1= 2.14 The C-0 bond distance in the transition state has been taken as the average vflue between double and single C-0 bond lengths, 1.33 A.23 The activation energies calculated thus are for CHICHIOH and CHJOH, 31.5 and 38.7 kea1 mol-l, respectively. Thus kB for formaldehyde iodohydrin a t 574°K is lo-’.’ M-l sec-’, much higher than our upper limit estimate and the same is true for acetaldehyde
(19) H. E. O’Neal and S. W. Benson, J . Phys. Chem., 71, 2903 (1967). (20) S. W. Benson and G. R. Haugen, J. Amer. Chem. Sac., 87,4036 (1965); J. Phys. Chem., 70, 3336 (1966). (21) The molar refraction, R , was calculated from those of n-butylaldehyde and propylaldehyde (K. G. Denbigh, Trans. Faraday SOC., 36, 936 (1940)), assuming additivity of molar refraction. The molar polarizability is calculated from the equation a = 3R/4nN, and the ratio longitudinal to total polarizability is taken to be equal to those of propane and ethylene, respectively ( 3 . 0.Hirschfelder, C. F. Curtiss, and R. B. Bird, “Molecular Theory of Gases and Liquids,” Wiley, New York, N . Y., 1954, p 950). (22) A. L. McClellan, “Tables of Experimental Dipole Moments,” W. H. Freeman, San Francisco, Calif., 1963, pp 41,64. (23) “C. R. C. Handbook of Chemistry and Physics,” 5lst ed, 19701971, p F-157.
EPOXIDATION OF ALKENESIN
GASPHASE
THE
3319
iodohydrin. The conclusion thus is that the Benson and Haugen calculation cannot be applied to fourcenter addition to a carbonyl bond, unless further modification would be done, such as enlarging the H I bond distance in the transition state more than in the transition state for additiqn to olefins (The ground sta,te bond length 0.40 A.2o). The rate constant For the elimination of water from
+
acetaldehyde iodohydrin can be calculated by the same method, using Haugen and Benson's c a l c ~ l a t i o n sfor ~~ addition of water to olefinic double bonds, which gives k = 1013.6-64.2/8, which is much smaller than the calculated rate constant for H I elimination. (24) G. R. Haugen and 9. W. Benson, Int. J . Chem. Kinet., 2, 235 (1970).
Epoxidation of Alkenes in the Gas Phase by D. J. M. Ray and D. J. Waddington* Department of Chemistry, the University of York, Heslington, York YO1 bDD, England
(Received April I d , 1972)
It is suggested that the retardation of acetaldehyde oxidation by alkenes is due to addition of peracetyl radicals to the alkene. This leads to the formation of the corresponding epoxide from the alkene. Addition of 2,3-dimethylbut-a-ene leads to long induction periods before acetaldehyde is oxidized and this may be due to interaction of peracetic acid with the alkene, leading to enhanced suppqeasion of chain branching. The addition reaction between peracetgl radicals and cis-but-2-ene is calculated as 2 X lo3 m3 mol-' sec-l a t 457 K.
I n a previous paper,' we reported that many aliphatic alkenes retard the gas-phase oxidation of acetaldehyde. By comparing the relative retardation coefficients of the alkenes and by considering the products formed from the alkenes, we suggested that the alkene reacted with acetylperoxy radicals (formed during the propagation reactions in acetaldehyde oxidation2), the resulting alkene-acetylperoxy radicals being stabilized by hyperconjugation. The addition reaction (reaction 1) is able to compete with the a1t)ernative reaction (reaction 2) However, 2-methyl-2-butene and 2,3-di~
03CCH3 GH&Oa.
+. >C=C
C-C
C-C
