J. Phys. Chem. 1995,99, 11883-11888
11883
Theoretical Study of the Thermal Decomposition of Acetic Acid: Decarboxylation Versus Dehydration Minh Tho Nguyen," Debasis Sengupta, Greet Raspoet, and Luc G. Vanquickenborne Department of Chemistry, University of Leuven, Celestijnenlaan 200F, B-3001 Leuven, Belgium Received: March 20, 1995; In Final Form: May 25, 1995@
The (C2II402) potential energy surface related to the thermal decomposition of acetic acid has been reinvestigated using ab initio MO calculations. Thermochemical parameters have been estimated using approximate QCISD(TC)/6-3 11++G(d,p) -t- ZPE calculations based on MP2/6-3 lG(d,p) geometries. In contrast to a previous theoretical study, our results show that the decarboxylation of acetic acid (CH3COOH C02 C h ) requires an activation energy similar to that of its dehydration (CH3COOH H2CCO H2O). It is also confirmed that the direct dehydration is less favored than the two-step process involving the 1,l-ethenediol intermediate. Higher level calculations predict that the latter is 115 f 10 kJ mol-' higher in energy than acetic acid with a heat of formation at 0 K, AH& (H2C=C(OH)2) = -303 f 10 kJ mol-'. The classical barrier heights (AH# at 0 K in W mol-') have been estimated as follows: 301 for decarboxylation, 317 for direct dehydration, 31 1 for 1,3-H shift, and 312 for dehydration of ethenediol. The ratio of the rate constants k(dehydration)/k(decarboxylation), calculated using a quantum RRK approach, is about 2 to 9 in the temperature range 1300-1800 K, in agreement with experiment. Thus, there is no need to invoke a bimolecular water-catalyzed decarboxylation mechanism as proposed in an earlier theoretical study.
-
-
+
Introduction An interesting feature of the gas phase thermal decomposition
of simple alkanoic acids (RCOOH) is that both decarboxylation and dehydration processes could occur. In contrast to the case of formic acid in which the decarboxylation (HCOOH C02 H2) is a rather minor process, acetic acid has been shown to decompose homogenously and unimolecularly via two competing channels:'-3
-
+
Bamford and Dewar' reported in 1949 the f i s t kinetic study of the acetic acid decomposition in a flow system at temperatures of 770-920 "C. The first-order rate constants were derived for both channels, namely
for decarboxylation and
k2 = 1012.95exp(-282.4 kJ mol-'/RT)
s-'
for dehydration. When the decomposition was reexamined by Blake and Jackson2 in 1968 in both batch and flow systems, they found that the C02 elimination (1) is a first-order reaction throughout a wide range of temperatures (530-1950 K) whereas the rate constant associated with the H20 elimination (2) is of secondorder below 600 K (k2 = 109.52exp(-164.4 kJ mol-'lRT)l mol-' s-'), but the Arrhenius plot curves upward at temperatures greater than 600 K and the rate constant order falls toward firstorder (kz = 10'2.45 exp(-271.5 kJ mol-'lRT) s-'). The dehydration channel seems thus to be sensitive to temperature. @
Abstract published in Advance ACS Abstracts, July 1, 1995.
+
In a subsequent paper, Mackie and Doolan3reported in 1984 a kinetic study in a single-pulse shock tube in which acetic acid was diluted in argon. These authors concluded that over the temperature range 1300- 1950 K, both fragmentation channels 1 and 2 are first-order reactions having nearly equal rates, the measured activation energies amounting to about 27 1.5-295.0 kJ mol-' for (1) and 282.4-295.0 kJ mol-' for (2). More recently, Ruelle4 carried out a theoretical study using molecular orbital calculations at the MP4/6-31G level and obtained the classical barrier heights of 376.6 kJ mol-' for (1) and 325.5 kJ mol-' for (2), in clear disagreement with earlier experimental results. To reconcile his theoretical results with the experimental ones, Ruelle4 proposed that in a batch system the dehydration occurs first; a water molecule produced in that step will then exert a catalytic effect on the decarboxylation. Such a catalysis which takes place through a six-membered transition structure reduces substantially the barrier height for decarboxylation. However, Nguyen and Ruelle5 have subsequently pointed out that the transition structures reported in ref 4 do not correspond to the first-order saddle points and suggested, on the other hand, that the dehydration could better be rationalized in terms of a two-step process involving the intermediacy of 1,l-ethenediol (eq 3), similar to the hydration reaction of ketene? H3CCOOH
-
H,C=C(OH),
-
H,C=C=O
+H20
(3)
Moreover, the calculations carried out by Ruelle4 had some severe limitations. The atomic basis set employed is quite small without polarization functions, and the zero-point vibrational corrections have not been taken into account. Recent theoretical studies' confirmed that the favored gas phase dehydration of acetic acid is in fact a two-step process (3), but the competition between both decarboxylation and dehydration channels and a possible change of mechanism with respect to the temperature have not been examined. In view of the existing conflicting results, we have carried out a theoretical study with the hope of addressing the experimental facts and uncertainties. In the
0022-365419512099- 11883$09.0010 0 1995 American Chemical Society
Nguyen et al.
11884 J. Phys. Chem., Vol. 99, No. 31, 1995 first part of the present work, we have reexplored the potential energy surface related to the unimolecular rearrangements 1,2, and 3 of acetic acid using uniform and reliable ab initio molecular orbital calculations. In the second part, the rate constants of different transformations have been analyzed in detail making use of a quantum statistical method and the data obtained from ab initio calculations. Our treatment provides, thus, a consistent picture which could rationalize several seemingly conflicting observations.
Ab Initio MO Calculations All calculations were carried out using a local version of the GAUSSIAN 92 set of programs.8 Four atomic basis sets, labeled 6-31G(d,p), 6-31 l++G(d,p), 6-311++G(2d,2p), and 6-311++G(2df,2p) were employed in the present study. Complete structural optimizations were initially performed using the Hartree-Fock method and the 6-31G(d,p) basis, and the structures located were characterized by harmonic vibrational wave numbers at this level. Geometrical parameters of the equilibrium and transition state structures were subsequently reoptimized using the second-order Moller-Plesset perturbation theory with the same basis set. Thermochemical parameters were then determined by the electronic energies computed at the MP2/6-3 1G (d,p)-optimized geometries and the quadratic configuration interaction (($21) and full fourth-order perturbation (MP4SDTQ) methods. Due to the limited computational resource, we have estimated the QCISD(T)/6-3 1l++G(d,p) electronic energies by assuming that the contributions of the triple substitutions (TC) are the same with both 6-31G(d,p) and 6-311++G(d,p) basis sets. In other words, we have used relationships 4 and 5 :
E(QCISD/6-31G(d,p)) E(QCISD(TC)/6-31 l++G(d,p)) = E(QCISD/6-31 l++G(d,p))
la
1
1.353 0, H\+33%/ ,H 1367 ~/12*\2' 2.138 H'
H '
2
TS l / l a
;
0
' . : 1.895 c'1.188
l.099,c'..!.363 H- - --- - . ~ 4 & 8 . 2 1*583i.283
TS 1/3 (C,)
TS 1/1
TS 1/2
TS 114
(4)
+ TC (5)
The selected geometrical parameters for acetic acid and relevant stationary points are given in Figure 1. The total and relative energies calculated through various ab initio methods are summarized in Tables 1 and 2. Schematic energy profiles illustrating the unimolecular decomposition of acetic acid are presented in Figure 2 which, for the purpose of comparison, also includes a selection of possible homolytic fragmentations. The relative energies of different fragments are compared with experimental values in Table 3. To determine the relative energy of 1,l-ethenediol,calculationsusing the larger 6-3 11 f G (2d,2p) and 6-31 l++G(2df,2p) basis sets have also been carried out; the obtained results are recorded in Table 4. The geometry of acetic acid 1 is well established; the performances of the HF and MP2 methods with various basis sets have been intensively examined in earlier s t u d i e ~ . ~We ,'~ note that the MP2/6-31G(d,p) structural parameters of 1 are very close to the experimental values determined by microwave ~pectroscopy.~ A comparison can also be made by considering its rotational constants as follows.
+
B (MHz) C (MHz)
\ 1510 lo 4C-i 1.366 'g 115.0 0 H'
-
TC = E(QCISD(T)/6-31G(d,p))
A (MHz)
H
.so1 /0 ~5'3 K 111.9 ,H 331 H \
calcd 11210 9506 5312
exptl 11335 9419 5325
The relative errors in the observed rotational constantsg are small overall ( 5 1%). The 1,l-ethenediol could exist in two
TS 214 Figure 1. Selected MP2/6-3 lG(d,p) optimized geometries for the C~H402structures considered. Bond angles are given in angstroms and bond lengths in degrees. TS X / Y stands for a transition structure connecting both equilibrium structures X and Y.
distinct conformations, but 2 is the most stable conformer.6 There is a significant difference between both C-0 bond distances of 2 (0.014 A). This is probably due to the existence of a weak intramolecular hydrogen bridge. Otherwise, both the C=C distance and the longer of the two C-0 distances are quite similar to the corresponding values in ethenol." It is useful to note that the rotational constants of 2 based on MP2/ 6-31G(d,p) geometries are: A = 11 018 MHz, B = 10 107 MHz, and C = 527 1 MHz. Hence, only the B constant of 2 is clearly distinguishable from that of acetic acid. Regarding the relative energies, we first examine the energies of the fragments for which experimental heats of formation are available.'1-13 As seen in Table 3, the agreement between the QCISD(TC)/6-31 l++G(d,p) and experimental results is satisfactory, in view of the moderate size of the basis set employed. The absolute errors range from 1 to 18 kJ mol-'; the largest errors concem the HOC0 and C H F O radicals where the degree of spin contamination in UHF references is relatively high. In addition, uncertainty in the experimental estimates is also The energy difference between acetic acid 1 and
J. Phys. Chem., Vol. 99, No. 31, 1995 11885
Thermal Decomposition of Acetic Acid
TABLE 1: Calculated TotaP (hartree) and Zero-Point Vibrational (ZPE, kJ mol-') Energies for the (C21&02)Structures Considered 6-31G(d,p) structureb
M"
H&-COOH 1 H3C-COOH l a HzC=C(OH)2 2
-0.469 -0.458 -0.415 -0.488 -0.398 -0.446 -0.415 -0.346 -0.338 -0.342 -0.341
c&+ c02 3
+
HzCCO H20 4 TS: llla
TS: 111 TS: la/3 TS: 114 TS: 112 TS: U4
6-3 11++G(d,p)
QCISD 33 57 63 22 74 73 41 68 51 04 77
-0.478 -0.468 -0.427 -0.492 -0.408 -0.457 -0.419 -0.346 -0.338 -0.347 -0.347
93 45 03 35 53 24 82 22 92 42 16
QCISD(T)
MP4SDQ
MP4SDTQ
-0.497 -0.486 -0.445 -0.512 -0.426 -0.475 -0.440 -0.369 -0.362 -0.368 -0.368
-0.589 -0.580 -0.542 -0.603 -0.530 -0.569 -0.528 -0.462 -0.456 -0.462 -0.461
-0.618 -0.609 -0.569 -0.561 -0.638 -0.598 -0.562 -0.501 -0.495 -0.496 -0.495
13 70 13 44 93 28 78 62 27 43 78
66 21 16 09 59 41 34 39 29 99 73
Core orbitals are frozen except in MP2(F) calculations. The values reported here are E(tota1) optimized geometries. Zero-point energies from HF/6-31G(d,p) and scaled by 0.9.
QCISD
75
-0.590 -0.580 -0.543 -0.530 -0.606 -0.569 -0.528 -0.461 -0.455 -0.463 -0.462
40 21 10 01 05 83 67 80 22 49
+ 228 hartree.
ZPE'
20 75 09 63 16 96 69 78 59 69 15
158 157 156 142 135 153 145 136 135 144 143
Based on MP2/6-31G(d.p)
TABLE 2: Calculated Relative EnereiesD (kJ mol-') for C2I-LO2 Structures Considered ~
6-31G(d,p) structure
MP2(F)
QCISD
H$-COOH 1 H$2-COOH l a H2C=C(OH)2 2 CH4 C02 3 HzCCO HzO 4
0 27 140 -65 163 54 130 300 320 320 320
27 135 -50 163 52 143 326 345 331 331
+
+
TS: llla TS: 111
TS: la/3
TS: 114 TS: 112 TS: U4 a
6-31 l++G(d,p) QCISD(T)
MP4SDQ
MP4SDTQ
0 26 136 -55 162 52 136 312 33 1 324 322
0 24 124 -50 133 48 149 312 327 319 321
0
0
24 129 -66 129 49 135 285 300 308 309
24 123 -57 134 48 149 315 330 318 321
0
QCISD
QCISD(TC)
0 24 123 -53 134 49 142 30 1 317 311 312
All relative energies are corrected for zero-point vibrational energies.
E
kJ/mo1
500 H3C-COO + H k 3 c c o OH H2CCOOH + H
45 1
+
400
H3C + HOC0
300
200 H20 + H2CCO
4
100 0
1
1
'
H3CCOOH
CH4 + C02
134
-53
3
-loa Figure 2. Schematic potential energy profiles (at the QCISD(TC)/6-31l++G(d,p)
ethenediol 2 is of particular interest as they represent another simple enol-ketone pair. For the sake of comparison, the energy difference between the prototypical pair, namely ethenol and acetaldehyde, is also listed in Table 4. It is apparent that, in both cases, extension of the one-electron basis functions and improvement of the correlation treatment tend to reduce the
+ ZPE level) for unimolecular reactions of acetic acid.
energy difference. It is pleasant to note that, in the ethenolacetaldehyde case, the higher the theoretical level, the better the theory-experiment agreement. Our best estimate for the ethenol-acetaldehyde energy gap of 44 kJ mol-' compares well with the experimental14 value of 41 f 8 kJ mol-' determined in the gas phase by mass spectrometry. A previous theoretical
Nguyen et al.
11886 J. Phys. Chem., Vol. 99, No. 31, 1995 TABLE 3: Comparison of Calculated and Experimental Fragments Relative Enerdes (in kJ mol-') of the CZH~OZ fragments
calcd"
H3CCOOH CH4 COL, HzCCO H2O H3C HOCO H2CCOOH H H3CCO OH H3CCOO H
-53 134 365 399 420 45 1
exDtlb
0
+ + + + + +
0 -42 135 377 400 439 455
OH
+
At the QCISD(TC)/6-31 l++G(d,p) ZPE level. Based on heats of formation at 0 K (AHZo) taken from ref. 12.
TABLE 4: Relative Energies (in kJ mol-') between the Pairs Acetic Acid-Ethenediol and Acetaldehyde-Ethenol at Different Levels of Theorv ~~~
level"
AE (ethenediolacetic acid)b
AE (ethenolacetaldehyde)b
139.6 124.9 128.8 122.6 119.3 122.4 116.5 120.3 115.2
62.8 52.6 56.7 52.2 47.5 51.1 45.1 48.9 44.3 (41 iz 8)'
MP2/6-3 lG(d,p) MP2/6-31 l++G(d,p) MP4SDTQ16-3 11++G(d,p) QCISD(T)/6-311++G(d,p) MP2/6-311++G(2d,2p) MP4SDTQ16-311++G(2d,2p) MP2/6-3 11++G(2df,2p) MP4SDTQ16-311++G(2df,2p) QCISD(T)/6-311++G(2df,2p) exptl
+ CH3COOH - CH2COOH + CH, CH,COOH CH,CO + OH
CH,
-
+ CH,COOH
--
CH,
+ CH,COOH CH,COO
and
+ HOCO CH,COO + CH, CH, + CO,
(7)
+
-
-
+
HpCCO*
+ H20
-
-*
HaC-C,
//u
c
OHH 1*
*
study derived a value of 47 kJ mol-' for this quantity using the G1 method." On the basis of the results of Table 4, we would suggest an energy difference of 115 f 10 kJ mol-' for the 1,lethenediol-acetic acid pair. Combination with the well established heat of formation of acetic acid, AH&o(CH3COOH)= -418 kJ mol-' at 0 K, leads to the corresponding value for ethenediol 2: AH&o(H~CC(OH)~) = -303 f 10 kJ mol-'. While the experimental gas phase energetics of 2 are not available yet, its free energy of formation in aqueous solution has recently been estimated making use of the PKa values of a series of exchange reaction^,'^ namely, AGdH2CH2CC(OH)2)aq = -272 f 8 kJ mol-'. Together with the corresponding value for acetic acid,I6 AGf (H3CCOOH),, = -394 f 2 kJ mol-', the difference in free energy of formation in aqueous solution between 2 and 1 could be evaluated at 122 kJ mol-'; this value is in the same order of magnitude of our gas phase estimate for enthalpy difference mentioned above. We now turn to the rearrangement and decomposition reactions. The geometrical parameters of the transition structures have abundantly been commented on in earlier theoretical ~ t u d i e s . ~Conceming -~ the energetics, examination of Table 2 and Figure 2 allows several important points to be noted: (a) The most facile homolytic fission of acetic acid is cleavage along the C-C bond. Nevertheless, all these dissociation limits lie at least 48 kJ mol-' higher in energy than the transition structures for unimolecular rearrangements. This fact, however, rules out the possibility of formation of C&, CO, H20, and H2CC=O via a simultaneous free radical chain mechanism which could be the following:
-
+ H,O
The results are consistent with the established fact that the C-C bond is weaker than the C - 0 single bond in the ground state of acetic acid. The C-0 bond cleavage becomes, however, the dominant path following a '(n n*)photolysis of acetic acid (CH3COOH CH3CO OH).I7 (b) The syn-anti 1 la and degenerate 1 1 isomerzations require far smaller activation energies than decarboxylation and dehydration. The TS l/1 for the 1,3-shiftof the hydrogen within the carboxyl moiety has an energy content similar to that of the product H20 H2C=C=O 4. This is of interest regarding the dehydration of labeled acetic acid such as 1* (eq 8). The fast exchange between two isotopic forms of 1* presumably leads to the formation of a mixture of H2CCO* and H2CCO with equal proportion.
Based on MF'2/6-31G(d,p) optimized geometries. All relative energies are corrected for zero-point energies. Reference 14.
CH,COOH
CH,COOH
CH,
(6)
/ObH H3C \\ 0
-
HpO*
+
HzCCO
(8)
(c) The energies of the four transition structures 1/2,1/4, la/ 3, and 2/4 are quite close to each other; the separation gap between the higher-lying and lower-lying structures amounts only to 16 kJ mol-'. Our best estimate suggests that the direct dehydration 1 4 has the largest barrier height whereas the decarboxylation has the smallest barrier. Both barriers related to the two-step dehydration through intermediate 2 have intermediate values. This result differs fundamentally with that reported by Ruelle4 who found a separation of 51 kJ mol-' in favor of the transition structure for direct dehydration 1/4. As mentioned in the Introduction, such a result arose no doubt from the use of small atomic basis functions and partly from neglect of zero-point energy corrections. (d) The TS 2/4 for dehydration of 2 is calculated to be only 1 kJ mol-' above the TS 1/2 for the 1,3-hydrogen shift converting acetic acid to ethenediol. Within the accuracy of the level of theory employed, which is about f 1 2 kJ mol-', it is rather hard to identify the rate-controlling step of the twostep dehydration. The latter could be expected to behave as a reversible process. In any case, owing to the small energy difference between the relevant TS, the competition between the decarboxylation and dehydration could better be analyzed in terms of their rate constants.
-
Quantum Statistical Analysis of the Decomposition Kinetics In an attempt to obtain more insight into the mechanism of the thermal decomposition of acetic acid, we have applied a quantum version of the Rice-Ramsperger-Kassel statistical theory (QRRK),IE combined with the chemical activation processI9 for the calculation of rate constants, to examine the feasibility of various decomposition pathways. The major advantage of the QRRK theory is the substantial simplification of the input data required. This method has been applied succesfully to calculate the rate constants of several recombination For details of the theory, we refer to the paper by Dean.20 The two-step decomposition of acetic acid 1
J. Phys. Chem., Vol. 99, No. 31, 1995 11887
Thermal Decomposition of Acetic Acid
SCHEME 1 1
ki(E)
k-iW
E*
I
12,
k3W
1800K
I WOK 700K
products
Bks (M)
B
SCHEME 2 ki’W
I+M-l*+M k-iW
can be represented in Scheme 1, where B* and B are the activated and stabilized forms of the intermediate or isomer; M, the bath gas, is generally taken as an inert gas; 8, the collision efficiency, was calculated from the expression suggested by Troe2,; ki(E) are the energy dependent rate constants for the individual processes calculated as a function of the internal energy E according to RRK theory.I8 The excess energy which is stored in the quantized vibrational level of B* can be utilized to cross the barrier for the formation of the products; B* can also be stabilized by collisions. These reactions involve the competition between stabilization and dissociation of the energized complex and should not be treated as single-step reactions. In other words, the apparent rate constants are more meaningful than the rate constants for the individual processes [i.e., ki(E)]. The apparent rate constants for the decarboxylation and the stabilization of the anti form of acetic acid la are defined as
Kdecarb
CH,COOH 1
-
CH,COOH 1
K,,
CH,
-
+ CO, 3
(via la)
CH,COOH la
Similary, the apparent rate constants for the dehydration and stabilization giving the 1,l-ethenediol isomer 2 are defined as
Kdehyd
CH,COOH 1
Kisomer
-
CH,C--O iH204 (via 2)
CH,COOH 1
-
CH,C(OH), 2
As we have discussed earlier, acetic acid can also decompose to ketene and water via a one-step concerted path. The calculation of the rate constant via QRRK theory for direct decomposition of acetic acid has been carried out following Scheme 2. The rate constant for the direct decomposition (KdiRct)has been calculated using the equation given in ref 18. The frequency factors (A) for the individual steps have been calculated from the equation
where e# and Q are the complete partition functions for transition state and reactant structures, respectively, k is the Boltzmann constant, T i s the temperature in degrees K, and h is Planck’s constant. Partition functions were obtained from ab initio calculated vibrational wave numbers and moments of inertia at the HF/6-31G(d,p) level; barrier heights are taken from our best estimates at the QCISD(TC)/6-31 l+G(d,p) ZPE level (Table 2, Figure 2). The apparent rate constants for the reaction are calculated as a function of temperature and pressure. The variations of both Kroland Kisomeras a function of pressure at three different temperatures (700,1O00, and 1800 K) (see Figure 3) show that the fall-off curves of Kisomer tend to shift toward the higher pressure with the increase in temperature while the
+
10
-...... ....................................Ka... .. ... .................. .. .......
... ... . . .... .... . . ,
0
-2
-10
-
-20
-
-30
-
-40-
Kdmi
-50
I O O W T . Tin Kelvin
Figure 4. Calculated apparent rate constants of all unimolecular processes of acetic acid as a function of tempereature.
syn-anti isomerization of acetic acid is not particulary sensitive to the temperature; the formation of 2 by a 1,3-H shift becomes faster at higher temperatures. It is clear that the formation of la (Kml) is, as expected, the fastest process indicating the existence of a large amount of the anti form la of acetic acid at high temperatures. The anti form l a has the correct nuclear configuration for a hydrogen transfer to the methyl carbon. At all temperatures and pressures, the formation of 2 by 1,3-H shift (&mer) competes favorably with either the decarboxylation (Kdecm) or the dehydration (Kdehyd). Hence, we predict a substantial formation of 1,l-ethenediol as a primary product. Figure 4 shows the variation, with temperature, of all the apparent rate constants for the thermal decomposition of acetic acid. The pressure was kept at a value where both Krot and KiSomerreach their high-pressure limits. The apparent rate constants for the one-step concerted decomposition to ketene and water (Kdirect)have much lower values compared to that for the two-step decomposition process. This clearly confirms that the one-step dehydration of acetic acid is not a favorable pathway. Figure 4 also Suggests that both Kdecarb and Kdehyd are nearly equal with a small preference for the dehydration at dl temperatures between 1300- 1800 K. The ratio KdehydKdecar varies between 9 and 2.3 in the temperature range 1300-1800 K. In other words, both thermal decomposition modes in acetic acid are competing unimolecular processes in this temperature range. This is in good agreement with the experimental results of Mackie and Doolan3 and points out that it is not necessary to invoke a bimolecular water-catalyzed decarboxylation mechanism as previously proposed by Ruelle4 to interprete the
11888 J. Phys. Chem., Vol. 99, No. 31, 1995 experimental observation. The small preference toward the dehydration pathway may be due to the larger value of its frequency factor compared with that of the decarboxylation process; the frequency factor is of course related to the entropies. The formation of 1,l-ethenediols as intermediates during the hydration reaction of ketenes in solution has recently been e s t a b l i ~ h e d , ~while ~ - ~ in ~ some cases, the diols could be detected s p e c t r o s ~ o p i c a l l y ;in~ ~other ~ ~ ~ cases, the reversibility of the ketene hydration is the only evidence for their intermedia~y.~~ We have also calculated the rate constants for both backward and forward transformations of 2 to form acetic acid and CH2C=O H20, respectively, via the Arrhenius equation. Our calculated values for the rate constants indicate that the backward and forward reactions are equally favorable. Thus, the gas phase ketene hydration is also a reversible process. It should be noted that this reversibility has already been taken into account in the calculation of the rate constants.
+
Summary In the present paper we present a reexamination of the (C2H402) potential energy surface related to the decomposition of acetic acid making use of reliable molecular orbital methods. Calculations of the apparent rate constants for various decomposition pathways have also been carried out in the framework of the quantum version of the RRK theory. The dehydration of acetic acid giving ketene and water via a two-step mechanism is much preferred over the direct dehydration. The decarboxylation giving carbon dioxide and methane is a one-step concerted process starting from the anti conformer acetic acid. In agreement with experimental results, our calculations indicate that the rate constants of both unimolecular decomposition channels of acetic acid are nearly equal at high temperatures. The ratio k(dehydration)/k(decarboxylation) amounts to 2-9 in the temperature range 1300- 1800 K. The previously proposed mechanism invoking a water-catalyzed decarboxylation could thus be disregarded. The dehydration of acetic acid involves its 1,l-ethenediol isomer as the primary intermediate. The latter is predicted to be 115 kJ mol-’ higher in energy than acetic acid; its heat of formation at 0 K being AH&o(l,l-ethenediol) = -303 f 10 kJ mol-’. The hydration of ketene by one water molecule in the gas phase is a reversible process.
Note Added in Proof: A theoretical paper (Duan, X.; Page, M. J. Am. Chem. SOC. 1995, 117, 51 14) and an experimental paper by N. I. Butkovskaya, G. Manke, and D. W. Setser in J. Phys. Chem. (in press) have recently addressed the decomposition of acetic acid. Setser et al. have found that the probability
Nguyen et al. of water formation is approximately 2 times higher than that of COS, in agreement with our analysis.
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