- Email: [email protected]
Thermal decarboxylation of acetate. Part I. The kinetics and mechanism of reaction in aqueous solution
Thermal decarboxylation of acetate. Part I. The kinetics and mechanism of reaction in aqueous solution DONALD
A. PALMER and S. E. DRUMMOND
Chemistry Division. Oak Ridge National Laboratory, Oak Ridge, TN 3783 1. U.S.A. (Received Uc&hr 16, 1985; accepted in revised.form February 14, 1986) Abstract-In an effort to understand the kinetics of the thermal decarboxylation of acetate and the role of catalysis, a series of laboratory experiments were conducted to measure the rate constants for the decomposition of acetate (acetic acid and sodium acetate) in the presence of titanium, silica. stainless steel, gold, and magnetite. Activation energies for decarboxylation of acetic acid and acetate ion range from about 8 kcal mol-’ in stainless steel vessels to 69 kcal mol-’ in silica tubes. Extrapolated rate constants at 100°C for acetic acid differ by more than fourteen orders of magnitude between the experiments conducted in stainless steel and the catalytically least active titanium vessels. Gold and titanium were the least active catalysts for the acetic acid substrate, while stainless steel, silica, and magnetite showed marked catalytic effects. Methane and carbon dioxide were the predominant reaction products of most of these ex~~rnent~ although mass spectrometric analyses of the gas phase revealed concentrations of carbon monoxide and hydrocarbons (apparent mass range from 29 to 56) amounting to as much as 55 mole percent of the total volatile products, depending on the catalyst. The reactions were generally first order in acetic acid or acetate ion, except for those involving the acid over silica and magnetite which were zero order. These results and the observed effects of variations in surface area are rationalized in terms of changes in the mode of surface catalysis. The mechanistic ~ignment is simplified by the existence of three unique straight lines on an isokinetic plot (Le., activation enthalpy versusactivation entropy) which lit all the respective first- and zeroorder reactions. The results described here provide the nucleus for the discussion in Part II of the role of acetate in the primary migration of methane and the transportation of metals in hydrothermal solutions. INTRODUCTION RECENTLY, KHARAKA ei al. (1983) demonstrated ex~~rnen~ly for the first time that acetic acid will readily decompose to carbon dioxide and methane in
the range of temperatures common to sedimentary basins. However, these kinetic data reveal an abnormally low activation energy (8.1 kcal mol-‘) for a reaction involving the breaking of a Karen-carbon bond and suggest that acetic acid decarboxylation was catalyzed in these experiments. Our experimental results confirm that surface catalysis is predominant under their conditions and demonstrate that many complexities exist in the decarboxylation of acetate that may not have been recognized by mARAK.4 et al. These new data suggest that acetate may play an impo~ant
role in the
primary mi~tion of natural gas and the insertion of metals in hydrothermal systems. Knowledge of the high temperature behavior of aqueous acetate solutions may also have important applications in problems associated with metal transport and corrosion in commercial steam generators. EXPERIMENTAL Kinetic experiments were carried out in a variety of vessels over a temperature range of 300 to 440°C. The vessel materials were as follows: 32 1stainless steel, gold (a gold bag suspended in a stainless steel vessel), titanium (commercially pure), titanium alloy (6% Al, 4% V), fused quartz and Pyrex tubes. Prior to each experiment the vessels were “conditioned” to ensure mp~u~bility. The 32 1stainless steel vessel was heated in air for several hours at 300°C to oxidize the surface. The gold bags were annealed in a Bunsen flame. The silica and quartz tubes (0.9 cm O.D., 0.3 cm I.D., 20 cm length) were washed for 30 minutes in a 5% HF solution. The titanium
vessels were thoroughly oxidized by heating them to incandescent red in a Bunsen flame. As a fmal preparation, all of these vessels were rinsed in distilled water and dried. In each experiment, the solutions were injected into the vessel at room tern~mtu~ and heated to the experimental temperature in a tube furnace. Before injection the titanium vessels were evacuated to (50 microns of pressure, whereas the gold bag and silica tubes were purged with argon to eliminate air. Neither of these measures were taken during the loading of the 321 stainless steel vessel (I.D. 3.5 em, internal length 28 cm) with acetic acid. The catalytic effects of various materials were measured in titanium vessels. The pure titanium and titanium alloy vessels gave identical results and hence could be used interchangeably. However, it was found that preheating the titanium alloy vessels to 400°C for ca. 24 hours while containing glacial acetic acid substantially pass&d the surfaces with respect to acetic acid decarboxylation. Three experiments were completed at 440.420, and 400°C before the surface began to return to its original reactivity. These experiments are referred to as taking place in “treated titanium.” Two experiments (numbered 16 and 171,involving M. 1m acetic acid, were carried out in the pure titanium vessels which contained one and ten tubes of 316 stainless steel (0.48 cm C&D.,0.24 cm I.D., and 10.0 cm long), respectively. The tubing had been previou~y cleaned with con~entmted HCI and was completely immersed in the acetic acid solution during the experiment. In experiment 18. a well oxidized 10.2 cm length of 1.6 mm O.D., 0.8 mm I.D. 3 16 stainless steel tube was immersed in ca. 1 m acetic acid solution. In order to assess the surface area effect on the corresponding reaction of acetate ions, experiments 3 I and 32 were conducted in a 316 stainless steel vessel (4.5 cm I.D.. I 7.8 cm internal length) containing eighteen 3 16 stainless steel tubes (0.62 cm O.D., 0.28 cm I.D., and 14 cm long). Most of the experiments with acetic acid solutions were monitored by titrating s&&on samples with a standard NaOH solution to the phenophthalein end point using a microburette with a readability of one part per thousand. This technique gave an accurate measure of the amount of acetic acid remaining and proved to be the most expedient method of fol813
814
D:A. Palmer and S. E. Drummond
lowing the decarboxylation reaction with time. Some of the initial experiments were monitored by measuring the pressure increaseamociatedwith the production of methane and carbon dioxide. It should be noted that this technique is normally subject to larger uncertainties which are not due to a lack of sensitivity in the pressure measurement, but rather to the difficulty in estimating both the solubihties of the product gases and the volume of tbe vapor phase. Calculating the latter is particularly arduous for experiments with the tlexible gold bag. The solubilities of CH, and CO2 were obtained from DRUMMOND (1981). The extent of the reaction of sodium acetate was determined exclusively by potentiometric titration. This involved titrating a sample, typically 1 cm3, to pH 2.00 with 1 m ha with rapid stirring to liberate bicarbonate as CO2 . The samples were then diluted to cu. 10 cm3 and backtitratcdwithstandardlmNaOHtoapHof11.OOtodetermine the total acetate concentration. These results were corrected for the dissociation of water and were reproducible to within 1%.Finally, a few scattered measurements of total acetate by ion chromatography were taken to confirm the accuracy of the other methods. At the termination of each experiment, the gas phase was also sampled at temperature and subsequently analyxed by mass spectrometry. The temperature of all these experiments was generallycontroUedto within 0.5”C. Stock solutions were prepared from 99.9% glacial acetic acid (Mallinckrodt, AR grade) or AR grade sodium acetate (F&herScientific) and deionized, doubly distilled water. Fine grainad,synthetic magnetite (Aha Producta 99.5%purity) was used without further purification. However, an X-ray powder diffraction pattern of this material failed to reveal any impurities.
RESULTS The decarboxylation of aqueous acetic acid can be represented by the general equation: [CHFOOH],
-
CH, + CO2
+ {CO + H2 + CnH2n+2 + - - . }
(1)
A summary of the pertinent experimental conditions is presented in Table 1. The measured molalities of acetate as a function of time ate given in Table 2 and the resulting rate constants are recorded in Table 3. Figures 1 and 2 show plots representative of the firstand zero-order reactions involving sodium acetate in stainless steel (Expr. 31 and 32) and acetic acid over magnetite (Expr. 20 and 2 1), respectively. KHAMKA et al. (1983) repotted a first-order rate constant for the decarboxylation of cu. 0.52 m acetic acid at 3OO’C in a stainless steel vessel of 8. I7 X lo-’ s-‘, approximately twice the value obtained from experiments 16 and 17 (Table 3). Extrapolation of their data to 349 and 359°C resulted in values of 1.42 X lo-’ and 1.58 X lo-’ s-t, respectively, which are conversely about half the value of the rate constants derived from experiments 18 and 19 in Table 2. Considering the great variability observed in the decarboxylation rates with both the composition and condition of the exposed surfaces, agreement of this order is quite reasonable. Moreover, this agreement is despite the large differences in the surface areas of iron exposed in each case, and in turn demonstrates that the kinetics are una&cted by the presence of titanium in experiments 16through IKKHARAKA et al. also reported olxrving high concentrations of iron during the initial stages of decarboxylation resulting in the formation of copious
magnetite crystals on the vessel walls by the end of the experiment. Experiments 16 to 19 duplicated the same behavior. The results of experiments with sodium acetate and mixtures of sodium acetate/acetic acid are presented in Table 3. Note that although f&tARAKA et al. (1983) did not observe any reaction of a 0.52 m solution of sodium acetate in stainless steel at 300°C after 1430 hours, this is consistent with the high activation enthalpy determined in this study at higher temperatures for this process which would predict only 1.4% reaction under their conditions. It is particularly relevant to the later discussion that in experiment 32 a significant amount of recrystallimtion of the titanium oxide surface layer was observed. Furthermore, neither the acetic acid solutions in experiments 3 to 11, nor those involving exclusively sodium acetate (i.e., 26-28) had any observable effect on the titanium surface. As mentioned earlier, the gas phase was sampled in situ whenever possible at the end of each experiment and analyzed by mass spectrometry. The results are given in Table 4, where CH represents hydrocarbons other than methane with the exception of mass 56. Some of the masses listed in parentheses are presumably fragments of larger hydrocarbons cleaved in the mass spectrometer. The liquid phase was only analyzed in one experiment (20) for nonvolatile organic mole cules other than acetic acid. A large array of organic molecules wem observed in the gas chromatograph of this solution with the dominant class of compounds being the C& alkenes. It is noteworthy that clear visual evidence was obtained in several experiments for the formation of a liquid organic phase after the reaction vessel had been cooled to room temperature. Experiments (4,6,7,8) over the temperature range 340-4OO”C in “untreated” titanium show that there is a discontinuity in an Arrhenius plot of the rate constants at or near the critical point (cu. 372°C for 1 m acetic acid). Extrapolated rate constants from data below the critical point are about 20% greater than those extrapolated from above. This phenomenon may reflect the inherent difference between the gas- and liquidphase kinetics. In order to test the possibility that decarboxylation occurs homogeneously in the liquid phase by interacting with metal ions released from the vessel walls, experiments 22 and 23 were carried out. In the former, added ferrous ion had no apparent influence on the rate which is compatible with the extrapolated rates observed previously in titanium vessels (Expr. 3 and 6). Similarly, chromium(II1) ions, which could be leached out of stainless steel by acetic acid solutions, had no measurable effect on the tate of decarboxylation (cf: Expr. 3 and 23, Table 3). Conceivably other metal ions in solution such as nickel@) could play a catalytic role, but the consensus of evidence points to a dominant heterogeneous path for decarboxylation. A typical Arrhenius plot of the rate constants for the decarboxylation of sodium acetate over titanium (Expr. 26-28) is presented in Fig. 3. The complete set
815
Acetate decarboxylation. Part I Table 1.
Summary of Experimental Conditions Prevailing in the Kinetic Runs
Expr. Temp.
Vessel Material
Volume of Vessel
Surface1 Area
Weight Initial Of Halality Solution
Method*
Acetic Acid :
357.2 359.0
3 4 5 6 7 8 9 10 11 12 13 14 15 lb
359.0 359.0 359.0 340.0 380.0 400.4 400.4 421.8 440.0 340.0 359.0 340.0 359.0 300.0
I7
3no.0
1R
359.0
19 20
349.4 359.0
gold bag gold bag titanium titanium titanium titanium titanium titanium treated titanium treated titanium treated titanium silica silica silica silica titanium (1 ss tube) titanium (10 ss tubes) titanium (1 ss tube) stainless steel titanium (magnetite, 59) titanium (mgnetite. 59) titanium (F&l,, O.OOlL) titanium (chromic nitrate, 0.0001 m)
166 178
64 68
64 35 64 64 64 64 64 64 64 1 1 1 1 35
80 63 40 49 109 109 109 109 109 15 18 15 18 23(ss)
25 15 15 25 15 12 12 12 12 0.5 Il.5 0.5 0.5 15
35
229(ss)
170
150 250 64
b(ss) 209 6x106 (ma9.j 6x106
t1tr.
0.987 1.054 1.064 0.0988 9.037 1.064 1.064 1.064 0.987 0.987 0.987 lb.39 16.65 1 .Ob4 1.068 0.987
pres. titr. titr. titr. titr. titr. titr. titr. titr. titr. titr. titr. titr. titr. titr.
15
0.987
titr.
60
1.054
pres.
100 25
1.054 0.989
pres. titr.
25
0.985
titr.
35
47(ma9.)
15
0.987
titr.
64
80
25
0.988
titr.
165 175 60
121 122 26
64 305 236
75 76 189 905
32 30 140 130
1.003 1.003 1.003 1.003 1.003 1.003 1.003
tjtr. titr. titr. titr. titr. titr. titr.
236
904
122
1.003
titr.
15
0.5401 0.540
titr.
31
1.097
titr.
64
21
340.0
22
300.0
23
359.9
24 25 26
359.0 381.0 340.0
gold cup gold CUP titanium
299 299 64
:; 29 30
381.0 359.0 359.0 359.0
31
381.0
titanium, titanium stainless steel stainless steel (1R ss tubes) stainless steel (18 5s tubes)
Sodium Acetate
Acetic Acid/Sodium Acetate 32
359.0
titaniwn
35
55
n-Butyric Acid 33
359.0
titanium
64
100
1
Surface Area represents the approximate area exposed of the dominant catalytic ccinponent prior to smnpling.
*Methods: titr. refers to titration against a standard NaOH solution; pres. refers to monitoring the pressure increase. 3
Inltlal twlalities refer here to concentrations at room temperature. i.e.. without correction for loss of water to vapor phase at temperature.
of these plots is represented in Fig. 4 for the zero- and first-order reactions listed in Table 5 incorporating those temperatures of immediate geological interest (see Part II). The activation parameters in Table 5 were determined according to the Eyring-Polanyi relationship, (MOORE and PEARSON, I98 1) Eqn. (2), and are shown together with the extrapolated values of the rate constants at 100°C. k = &T/h)
exp(--AH*/RT)
exp(ASIR).
(2)
In Eqn. (2) k is the observed rate constant; k, is the Boltzman constant; T is the temperature, K; h is Planck’s constant; AH* is the enthalpy of activation; AS* is the entropy of activation; and R is the universal
gas constant. The previously published results for decarboxylation over silica included in Table 5 refer to supercritical acetic acid in the absence of water. The relationships between the activation parameters are shown in Fig. 5. DKXUSSION One of the original goals of this endeavor was to establish the minimum, or least catalyzed, rate of decarboxylation of acetate. This rate would then form the basis for studying the effects of catalysis and, in addition, would provide the ultimate lower kinetic boundary for a natural gas migration model, as well as the maximum lifetime for acetate as a potential li-
D. A. Palmer and S. E. Drummond
816 Table 2.
Hours Hours
mlality
klality
Hours
Expr.
19.5 79.5 170.5
0.987 0.976 0.963
240.5 337.5 418.3
Hours
1.040 1.040 LO47 1.045 1.043 1.046 1.035 1.040 1.033 1.035 1.040 1.035 1.032
399.5 427.5 471.5 :~~*x 595:o 620.0 644.0 7b1.0 784.5 808.0 835.0 907.5
1.031 1.030 1.022 1.018 1.013 1.011 1.007 1.006 0.998 0.994 0.994 0.989 0.981
307.5 580.0 680.0
0.914 0.908
0.969 0.947 0,929 0.912
1.042 1.043 1.017 0.992
527.0 954.5 981.0 999.5 K8 11Ob:O 1125.0 1152.0 1169.5 1222.5 1240.0
0.980 0.971 0.974 0.974 0.961 0.963 0.962 0.961 0.959 0.954 0.945 0.945
1339.0 1510.0 1511.0 1637.0
0.895 0.872 0.867 0.865
Expr. 4 67.0 166.0 233.0
0.0855 0.0766 0.0713
328.5 428.3
72.0 159.0 165.0 242.5
9.846 9.395 9.366 9.280
333.0 333.n 405.5 501.0
0.0637 0.068§
501.0 569.0
0.0531 0.0481
556.5 775.5 914.0
8.631 8.097 7.965
Expr. 5 9.023 9.023 8.687 a.525 Expr. 6 159.0 663.0
1170.5 1698.0
1.018 0.984
2173.5 2653.5
I.059 1.062 1.053
1.041
153.5
1.023 1.012 0.998 0.978
:::*: 3so:o Expr,
140.5 209.5
1.010 0.967
312.0 405.0
417.5 541.5 684.5
0.962 0.935 0.901
0.878 0.841
379.5 503.0
0.835 0.816
477.5 573.0
0.955 0.949
89.5
0.980
0.563 0.213
122.0
fxpr.
Expr.
965.5 1274.0
312.5 601.0
313.0 601.0
0.967 0.943
22.5 49.0 67.0 75.5
1.054 0.978 0.956 0.948
115.0
965.0 1274.0
k!
114:o 123.3
gin 96.5 112.0
ii-;: a:934 0.922 0.904 0.892 0.887 0.883
114.5 146.0 163.0 168.0 184.0 2oR.o 234.0 304.5 328.0 333.5
71.o 94.0
0.729 0.690
143.0 167.3
16.0 22.3 39.5 t?: ;:I:
1.043 0.961
Expr.
Expr.
c:
95.0 118.5 142.0 163.0 264.5
.
a:::
0.913 0.882
94.0
0.988
189.5
16.49 16.4S1 16.M 16.20
333.0 333.0 it:*:*
15.52l 15.79 15.6g1 15.26
1.0
0.891
1.097 1.062 1.023 0.944 0.875
1.277 If *oOz-t:; 23:s
27.5
::z .
136.5
15.52
136.5
0.105
1613.5
0.777
1613.0
0.809
139.0 170.5 235.3 259.0
0.879 0.862 0.809 0.801
zz-8 373:o 384.0 399.0 408.0 42b.0 449.0 471.5
0.692 0.684 0.684 0.670 0.670 0.657 0.645 0.636 0.629
236.0 303.5
0.368 0.223
333.0 426.0 496.5 576.3
0.485 0.307 0.242 0.182
:::“33.
0.977 0.973
16 0.881 0.830
0.908 0.857
0.931 0.918 0.907 0.898
0.875 0.850 t-iii 0:82O 0.801 0.778 0.729 0.712 0.699 20 0.600 0.542 21 0.704 0.690 0.664 0.652 0.566
23 0.963
359.0
0.950
:fF: 140:5 164.5
oh79
42.5 51.0 66.5
0.830 0.796 0.744 0.691
192.5 261.0 286.5 308.5
Xii 0:513 0.485
75.0 75.5 98.5
0,434 0.417 0.301
618.0 721.0
0.742 0.703
253.5 301.5
0.643 0.584
25 0.717 0.632 0.489
Expr. 26 525.0 648.0
15.47 lb.39
22.0 219.0 382.5
1.037 0.959 0.871
:::-: . Expr.
Expr. 13 93.0 122.0
1.0031 0.160
Expr. 24
Expr. 12 43.5 68.5 168.0 193.0
355.0 525.0
15
0.210
0.991 0.990
f$t: . Expr.
2255 51:5 81.0
0.841
0.631 0.506
Expr. 22 0.987 0.995
Expr.
it:::.
14
Expr. 19
0.798 0.763
Expr. 11 0.980 0.955
wr11ty
Expr. 18
20.5 643.0 808.5
Hours
Expr. 17
0.843 0.797
Expr. 10 ::*:0
93.0 121.5
312.0 312.0
Expr. 9 68.3 309.5
312.0 355.0
17.0 23.5 40.5
8 0.918 0.875
1,060 0.925
0.955 0.933
Expr. 7 I!$ If:5 86.5
lklrlfty
43.5 193.0
0.886
E!cpr. :z 331:5 498.5
Hours
Expr.
Expr. 2 87.3 120.0 161.5 167.0 213.0 237.0 265.0 283.0 304.0 312,5 336.0 352.0 386.5
Rolality
rnolality
1
0.953 0.941 0.928
tonttnued
Table 2.
Smmry of Kinetic Data for the Decarboxylatlffn reaction
136.6
15.51
14.0 43.5 73.0
:z ‘
0.837 0.792 27
z:: .
817
Acetate decarhoxylation. Part I Table
Hours
2.
HOUPS
Molality
Continued
Holality
Expr.
1;:: 26.5
Expr.
3:::
Expr. 2.5 18.0 49.5
1.024 0.966 0.860
71.5 116.0 138.5
::: 3.5 5.5 7.5 8.0 11.0
1.044 0.998 0.925 0.909 0.907 0.822 0.766
11.5 23.0 23.5 27.5 28.0 32.5
0.0 28.0 46.5
65.0 70.5 102.5
2.0 25.5 91.0
1
These
1.027 1.035 0.972
samples
138.0 193.0 265.5
were
taken
2The molallties in Expr. acetic acid content.
162.5 186.5 213.0
0.563 0.519 0.482
47.5 48.0 54.0 94.5 95.0 121.0
0.355 0.297 0.321 0.122 0.129 0.0845
139.0 140.0
0.029 0.018
457.0 672.0 935.5
0.922 0.888 0.856
33 0.959 0.963 0.946
from
0.513
32 0.120 0.094 0.018
Expr.
202.5
31 0.630 0.655 0.550 0.541 0.491 0.370
Expr.
0.182 0.103 0.095
30 0.784 0.655 0.602
Expr.
142.0 211.0 213.5
29 0.871 0.615
85.5 178.0
1.106 1.082
Holality
28 0.611 0.484 0.321
41.5 66.0 98.0
1.045 O.R28 0.738
Hours
quartz
stainless steel. In acidic solution, iron is readily soluble as evidenced by the high levels observed initially in acetic acid solutions during this study and reported by KHARAKA et al. (1983). e.g., 329 (300°C) and 5600 (200°C) mg L-‘. In contrast. iron concentrations in the sodium acetate solutions were < 1 mg L-’ and decarboxylation was found to be orders of magnitude slower. An explanation of this phenomenon is that the dissolution, or oxidation. of iron is mainly promoted by the acetic acid content of the solution due to the formation of [Fe&]aq which is predicted to be the dominant iron acetate complex at these conditions (DRUMMONDand PALMER, 1985) and is independent of the acetate anion concentration. The act of dissolution produces active sites on the walls of the vessel, such that the overall rate of decarboxylation is governed by the rates of adsorption/desorption of acetic acid molecules onto these sites and the rate of the subsequent decomposition of the adsorbed species. The implicit assumption being that decomposition. the actual breaking of the C-C bond, occurs predominantly at the active sites on a given surface. This reaction sequence is consistent with the earlier observation reported by WARAKA et al. ( 1983) that in an equimolal acetic acid/sodium acetate solution in stainless steel only the acetic acid component reacted during the course of the experiment at a rate which was first order in acetic acid, rather than total acetate. In other words, the total number of active sites generated is related to the solubility of iron which, at these experimental con-
tubes.
32 refer exclusively
Tahle 3. Constants
to the
gand in a metal transport model. The rate constants shown in Table 3 establish that the slowest reaction of aqueous acetic acid is in “treated” titanium. The high activation enthalpy and near-zero entropy (Table 5) indicate minimal catalytic influence, similar to the corresponding values reported at higher temperatures and low pressures for the same process in a nonaqueous environment (BAMFORD and DEWAR, 1949; BLAKE and JACKSON, 1969). Indeed, BLAKE and JACKSON ( 1969) demonstrated that free radicals were not formed under their conditions and concluded that the reaction is therefore homogeneous, which in this case implies uncatalyzed. The apparent indifference exhibited by the rate constants for acetic acid decomposition to either the presence of added ferrous chloride (Expr. 22), or the varying concentrations of ferrous ion resulting from the initial dissolution of stainless steel and the subsequent precipitation of magnetite in Expr. 19 provides strong evidence for a heterogeneous mechanism, which is controlled by reaction sites at a surface, at least for those reactions carried out in stainless steel. Consider first the implications of the pH effect on the heterogeneous mechanism for decarboxylation in
Expr. No.
Temp. “C
First and Zero-Order Rate for the Decarbowlation of Acetate
First-Order
Zero-Order “S
: 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 1R 19 20 21 22 23 24 25 2’; 28 29 30 31 32 33
359.0 357.2 359.0 359.0 359.0 340.0 400.4 380.0 400.4
4.28 2.51 3.89 3.08 7.19 1.55 6.64 1.52 4.72
x x x x x x x x
421.8 440.0 340.0 359.n 340.0 359.0 300.0 300.0 359.0 349.4 359.0
1.99 5.03
x 1o_7 x 10
340.0 300.0 359.0 359.0 381.0 340.0 359.0 381.0 359.0 359.0 381.0 359.0 359.0
5.03 3.83 2.A9 2.94
co.3 3.2 7.13 4.15 1.57 5.09 3.08 1.00 1.02 6.30 6.75 4.20
x x x x
x x x x x x x x x x x x
k-l
1O-8 1O-8 10-8 lo-’ 10-8 10-8 10:; lo_, 1O_8
10-8 1O-8 10-l 10-7
10-8 lo-’ 10-7 10-6 10-7 10-7 lI+ 10-6 Ml-6 1O-6 1O-6 1O-8
7.97 2.31 6.39 3.02
x 10:; x x lo_, 1O_6 x 10
6.15
x
3.18
x 1O_7 10
-7
D. A. Palmerand S. E. Drummond
0.8 0.6
s”
-0.8 -LO
0
40
20
80
60
100
TIME(h) FE. 1. First-order plot for the reaction of sodium acetate
in stainlesssteel at 38 I"C(Expr.3 1).
ditions, is controlled largely by the acetic acid content of the solution. A direct consequence of this hypothesis is that the hrst-order rate constants should be independent of surface area, given that there is sufficient solid to satisfy the amount of acetic acid present. All the first-order reactions studied here support this mechanism as the results show no evidence of a surtkce area dependence. This mechanism can be formulated Table 4.
Expr. No.
"Active Surface'
Tcnp. OC
gold goldb tjtan. titan. titan. titan. titan. silica quartz silica silica silica 5s 5s 55
nag. nag. titan.
gold gold titan. titan. titan.
359.0 357.2 359.0 359.0 340.0 400.4 400.4 340.0 340.0 359.0 340.0 359.0 300.0 300.0 349.4 359.0 340.0 300.0
Mstribution
Curation of Expr. h
in terms of a general rate law based on the following reaction sequence, namely, S (solid) + HAc +
Of Gas Pbasc Products1
CM4
co*
ketic
kid
co
H2
820 1240 1676 1255 265 651
41.6 22.5 20.3 10.3 31.5 10.3
42.8 48.9 47.6 71.3 60.0 47.2
6.6 25.3 7.0 5.4 2.5 4.5
3.1 1.1 3.9 0.9 0.4 5.2
:: 333 137 525 137 1274 1274 316 766 577 1613
35.7 0.9 9.1 15.1 22.4 32.2 2.R 5.3 49.1 lb.3 5.2 3.2
42.3 52.4 15.5 54.3 28.5 35.2 15.6 46.3 45.1 63.2 63.0 34.4
4.5 7.8 20.0 7.2 4.6 5.1 1.1 4.9 4.3 2.3 3.8 53.8
6.1 0.3 0.8 0.3 0.3 CO.01 75.4 30.9 1.5 4.0 1.5 1.4
Sodim
ketate
309 143 789
60.3 80.4 60.6
18.3 34.6 17.3
3:4 0.2
::: 21.3
titan.
305 214 203 213 121 167
7A.4 79.2 79.5 85.9 76.2 37.3
12.9 lb.7 20.0 12.6 19.6 59.9
0.1 0.1 0.1 0.1
:*: 0:5 1.2 4.0 2.6
titan.
359.0
940
19.3
21.6
5s OS
n-Butyrlc
%mallred vapor are
7.0
CH(masses)
&
0.5 1.0 2.2 3:4 0.3 0.8 2.9 10.3 2.3 3.7 3.3 0:5 2:b 015
5.4(29.44.55) 1.3(48-55.57) 18.0(29.41.55) 12.1 2:::(29.44.55) 10.6(29.44.55) 27.7(29.41,55) 44.3129.41.55) 20.8(29.41.55) 40.5(29.41.55) 24.2(29.41.55) 5.1f29.44.55) 5.4(29,44,55) 10.7(29.41.55) 26.5 b.bt29.41.55)
kid 52.0
to 100%. 1 .c ., the coexisting gases such as k. not included.
Acetic acldlsodium acetate mixture.
HAC& 2 products
1
S (solid) + AC-+A&
359.0 381.0 340.0 359.0 3Rl.O 359.0 359.0 381.0 359.0
55
2
RG. 2.Zero-order plots for the reaction of acetic acid over commercialmagnetiteat (a) 340°C (Expr.21) and (b) 359°C (Expr.20).
112, and water
2 products
(3) (4)
Acetate decarboxylation. Part 1
either K, > K3 (k:1 = k$/k-3) or k2 > k4 or both. It is relevant that the activation parameters associated with the thermal decomposition of acetic acid preadsorhed on a Fe( 100) surface (BENZIGER and MADIX. 1980), AH* = 32.5 kcal mol-’ and AS* = 2.1 cal K-’ mol-‘. fall well outside the error limits for the first-order straight line correlation in Fig. 5. As these parameters reflect the expected values for the liz step in Eqn. 3. this failure to correlate indicates that adso~t~on/desorption the~~ynamics associated with 4.Vtdo contribute to the activation parameters in Fig. 5. It should also be mentioned that the remarkable linearity exhibited by the three sets of data presented in Fig. 5 (namely. for first-order in acetic acid, first-order in acetate. and zero-order in acetic acid) is indicative of a common mechanism prevailing in each case (WILKINS. 1974).
-NOc)
C
-6.0
-6.5
-7.0 l.50
160
l.55
l.65
1000/T FIG. 3. An Arrhenius plot describing the temperature dependence of the reaction of sodium acetate in titanium (Expr. 26-28).
Making the reasonable assumption that the activated sites, or adsorbed acetic acid, reach a low, steady-state concentration leads to the general rate law: Rate = d[HAc]/df = -{k’&,/(k_,
+ k2)}[H/lc]
819
(5)
where k’, = k,S, with the corresponding expression for the rate of acetate ion decomposition. Given the validity of this rate law, the measured first-order rate constants in Table 3 are really composite values of the form k’,k2/(k_, + k2) with k2 most likely representing the slower, rate determining step, such that this constant term reduces to K~k~(=k~k~~k_,). For acetic acid to react substantially faster than acetate ions, with the corresponding large difference in activation enthalpy,
The following discussion provides an explanation of the kinetic behavior observed in titanium. The sol-
ubility of titanium oxide is less affected by acid than by complexation with acetate ions, and hence the rate constant verszts pH profile is a complex function of the acetic acid and acetate concentrations (~6 Expr. 3. 27 and 32. Table 3): keeping in mind that the dissolution of titanium molecules is associated with the formation of active sites presumably by exposing titanium metal. The maximum solubility of titanium oxide corresponds to the highest concentration of titanium acetate complexes, and is presumably reached near equimolal ratios of acetic acid and acetate ions (Expr. 32). This is reflected by the highest rate constant observed at this condition, namely 6.75 X 10m6s-’ {
FIG. 4. A summary of the temperature dependencies of the reactions presented in Table 5 extrapolated to in&de temperatures of geological interest. where the legendis as folfows:(a) sodium acetate over magnetite: (b) 16.5 m acetic acid in s&a; (c) 1.06 m acetic acid in silica; (d) acetic acid in stainless steel, KHARAKA et al.: (et suoercritical acetic acid in titanium: (f) acetic acid in tj~niurn; (g) sodium acetate in ti~nium: (h) so&&n acetate in stainless steef; (i) sodium acetate in gold; (k) acetic acid in treated titanium. The units of k are m s-’ for zero order and s-’ for the first-order ease. The logarithmic half-life (tre) scale is relevant to the zero-order plots only for an acetate concentration of 1.39 m.
820
D. A. Palmer and S. E. Drummond Table 5.
Expr. nos.
ktlration
Temp. Range
Farametcrs for the lkcarboxylation
k( lM)‘Y)
“Active Surface"
ms
s-1
-1
of ketate
AH* kca!1 no1
As*_ 1 cal _K1 nO1
Acetic Acid 3413-359
titan.
5.5 x 10-17
35.9
-37.1
titan. tr. titan silica silica SiliCS1 silica2 si ica3 ss 1 mag.
9.8 5.9 x 10-17 10-23
34.2 ::*; 61:4 55.8 60.3 63.7
20.21
400-440 380-400 340-359 340-359 494-585 642-700 767-916 200-300 340-359
2:::
-10.9 -40.5 -19.5 11.8 -14.2 -0.2 -5.2 -79.6 -51.9
24.25 26-28 30.31
359.381 34ll-381 359-381
gold titan. ss
1.4 x 10-22
64.5 56.8 66.8
+13.9 1.4 *la.2
9111 :*: 12.13 14.15
1.2 x 10-16 3.2 x 1O-21 1.2 7.0 2.1 1.8
x x x x
10-23 10-28 10-26 10-9
8.0 x 10-10
Sodiun Acetate
'(Blake and Jackson, 1968). no water vapor added; static technique with p(acetic acid) - 200 m in unpacked silica tubes. '(Blake and Jackson, 1969). no water vapor a&dad; flow technique through "mature" silica tubes. '(Banford and C&war. 1949). no water vapor added; flow technique through carbon coatad silica tubes. 4(Kharaka et al., 1983).
and acetate concentrations, the formation of Ti(Ac)S and higher titanium acetate complexes must be evoked in accordance with the following general equation: Ti02 + 2HAc + H2 + nAc- = Ti(Ac)L
+ 2H20 (6)
E 0
FIG. 5. Isokinetic plot for the kinetics of decarboxylation of acetate in aqueous solution. The symbols are defined in Fig4,wiUlthcadditionthat99aadIXIrrp*lcntthereaaion of anhydrous acetic acid in silica taken from BLAKE and JACKSON (1968) and ( 19691, and BAMFORD and DEWAR (19491,mspectively. The solid line is a linear least-uquates fit of the first-orderacetic acid reactions, the chaindashed line apt&s to the tirst-order reactions of acetate ion, while the dashed line fits the xere-order results.
The important initial step in the dissolution process associated with Eqn. (6) can be envisaged as adsotption of HAc on the surface in accordance with Eon. (3). thus fulfilling the same requirements for a first-order reaction as laid down for stainless steel. The high concentration of titanium in solution would lead to the gradual deposition of Ti4 during the experiment as the concentration of acetic acid declines, resulting in the observed ‘YecryWliz.ation of TiO*” in Expr. 32 described earlier. Raising the initial concentration of acetic acid to 9.8 m (Expr. 5) increases the titanium complex concentration, which according to Eqn. (6) is propottional to [HAc]*[Ac-I”, giving rise to increased TiOz solubility compared to that at 1 m. Thus the number of potential active sites is larger and the rate of decarboxylation increases accordingly. Conversely, lowering the acetic acid concentration in Expr. 4 to about 0.1 m should decrease the concentration of titanium acetate in solution giving rise to a slower rate. However, the reaction is in fact faster than either the 9.8 or the 1 m solutions. As experiment 4 shows a good first-order linear dependence them is no reason to doubt the authenticity of this result and it remains an enigma. Gold is less soluble than titanium and iron in acetic acid solutions and hence fewer active sites are generated for adsorption and decarboxylation. Moreover dissolution of a surface gold atom merely exposes more inert gold atoms, whereas removal of relatively inert oxide layers for titanium and iron reveals highly active metal underneath which can readily be imagined as providing a more teactive surface for acetate to adsorb and react Thus the rates of reaction in the presence of gold tend to be closer to the background homoge-
Acetate decarboxylation. Part 1 neous rate. It also follows that if the oxide layers on titanium and iron remain intact, as presumably is the case in the presence of sodium acetate, as well as for acetic acid in treated titanium, the rates of reaction will approach that observed in gold. Comparing the rate constants extrapolated to 100°C (Table 5) for sodium acetate on all three surfaces and acetic acid on treated titanium, which must have a thicker than usual oxide layer protecting the reactive titanium metal, shows that they are all approximately similar. This leads to the conclusion that in weakly catalyzed environments acetic acid and acetate anions have similar “inherent” reactivities with the activation enthalpy reflecting mainly the energy required to break the C-C bond, within the weakly bound substrate. Having established a mechanistic framework to rationalize the pH dependence for acetate decarboxylation kinetics, it now remains to be seen how this hypothesis can be used to explain the observed changes in reaction order. First consider the somewhat analogous dehydrogenation of dry formic acid in the gas phase (HCOOH - Hz + CO*) studied in detail by MARS et al. ( 1963). The order of this reaction is known to vary from zero at low temperatures and high pressures to first at high temperatures and low pressures. This tendency was correlated with surface adsorption of formate being complete for the zero-order case, thereby ensuring that the rate-determining step involves the decomposition of the adsorbed intermediate and is consequently not sensitive to the bulk formate content. In the case of a first-order reaction, active sites are still available on the surface and the rate of adsorption, which is directly proportional to the formate content, controls the process. Accordingly, the activation energy varied from 15.8 kcal mol-’ for zero order to 6.2 kcal mol-’ for the first-order adsorption process. If the same criteria operate in the solution phase, increasing the surface area to concentration ratio for a given surface material should only accelerate the zero-order reaction, until it eventually leads to a change in mechanism from zero to first order. It is tempting to draw parallels between the kinetics of formic and acetic acid because there is ample evidence for strong adsorption of the latter (and acetate ion) on iron oxide (KURIACOSEand JEWUR, 1977; BENZIGERand MADIX, I980), titanium oxide (GONZALEZ et al., 1978) as well as other transition metal surfaces (YINGand MADIX, 1979; BARTEAU et al., 198 I). From the earlier discussion of the origin of the first-order acetate dependence gleaned from Eqn. 3, it is apparent that the analogy to the formate system is only partly correct in this case. However, from a consideration of the conditions which must prevail in this scheme to promote zero-order kinetics, it will be seen that the formate analogy does present a plausible explanation. For the reactions of acetic acid and acetate ions in stainless steel, all of which have been determined to be first order, no significant change in rate at a given temperature was observed despite substantial changes
821
in the ratio of surface area to moles of total acetate in solution. For example, in Expr. 16 and 17 the ratio varies from 0.99 to 9.9 cm* mmol-‘, while the rate constants at 300°C are quite similar at 5.03 X lo-’ and 3.83 X lo-* s-‘, respectively. Extrapolation of the rate constant for Expr. 18 to 300°C yields a value of 7.8 X lo-’ s-’ for a ratio of 0.072 cm* mmol-‘. indicating that this 138-fold change in area had no significant effect on the rate. Similarly. the rate constants for experiments 29 and 30 are not influenced by a 4.8-fold change in surface area. Thus in all these cases. and most likely for the first-order reactions in general, there is sufficient surface available to promote enough active sites to satisfy either the acetic acid or acetate concentrations in the solution. Here the analogy with the formic acid mechanism ends, with the first-order dependence arising from the formation of active sites coupled with the chemical decomposition of the adsorbed intermediate, rather than being simply due to a diffusion and adsorption of the reactant molecules to the surface. The high activation enthalpies observed for the reactions on gold and titanium. for example, would be incompatible with a diffusion/adsorption rate-determining process, cf: 6.2 kcal mall’ for formic acid. Attempts to reduce the surface area ratio to even lower values in order to promote an adsorption limiting rate process requires containing the acetate solution in a vessel of some other, more inert material. Unfortunately. at least at the temperatures where the reaction rate can be measured over a reasonable length of time, no vessel material could be found which did not impose its own catalytic effect thereby dominating that of the smaller surface of interest. Hence the postulate that the first-order processes can be converted to surfacearea dependent, zero-order paths could not be tested. Both zero-order reactions have unique features which set them apart and neither has yet been shown to be surface-area dependent. Firstly, the reaction of acetic acid over magnetite is unquestionably zero order as demonstrated by the clearly linear plots in Fig. 2. Nevertheless, in the corresponding reaction in stainless steel (Expr. 19) where a large amount of magnetite was gradually formed during the experiment, the reaction rate neither increased in proportion to the magnetite buildup, nor deviated from a slower, strictly first-order dependence. Therefore, it must be concluded that minor unidentified impurities in the commercial magnetite used in Expr. 20 and 2 I are responsible for the zero-order kinetics. This is consistent with the fact that the surface area of this fine-grained magnetite is orders of magnitude larger than found in any of the other experiments and should therefore promote a first-order reaction if Fe304 sites were involved. The second example of zero-order kinetics was found for those experiments conducted in silica tubes where it had been initially hoped that essentially uncatalyzed rates may have been observed. The silica tubes containing the one-molal solutions were strongly etched and substantial induction periods of about I30 h (Expr.
D. A. Palmer and S. E. Drummond
a22
14) and 50 h (Expr. 15) were observed. In the 16.5 m solutions the dissolution of silica was less apmfent and no corresponding induction period was seen. The dissociation constant for acetic acid at 359°C is approximately 1.8 X lo-* (FISHER and BARNES, 1972) while that of silicic acid is about 1.9 X lo-” (BUSEY and MESMER, 1977) and hence the solubility of silica in these experiments is mainly a function of the water activity rather than proton activity. The generation of active surface sites would consequently no longer be dependent on the acetic acid concentration. Assuming that these sites are all immediately occupied by acetic acid molecules, then the observed rate of decomposition would be independent of the total acetic acid concentration and zero-order kinetics would result. In the case of the reaction over magnetite, the number of sites is limited by the amount of impurities in the solid and these must also be considered to be completely occupied by acetic acid molecules. This rationale is similar to that proposed for the formic acid analog. An alternate explanation for the zero-order kinetics, at least over silica, can be found in the following reaction sequence: (SiOr)d
+ 2HrO $
active site + H4c 2
active site + (H&iO,),
(7)
(MC)& 2 products
(8)
Solving the rate expression for this reaction sequence with the assumption that the number of active sites is at a low, steady-state level and that ks > k, leads to, d[I-ZAc]/dt = -kI . However, from a study of the kinetics of silica dissolution, RIMSTIDTand BARNES ( 1980) reported an activation energy for this process of only 11.7 kcal mol-‘, far less than the observed AH* (AH* = E. - RT) listed in Table 5 for the zero-order decarboxylation reactions. Not only is this mechanism incompatible with the results, but a linear isokinetic plot would not be expected for the different solids, magnetite impurities and silica, if k, is rate determining Thus, until experimental evidence to the contrary becomes available, the zero-order behavior is best explained by the concept of a limited surface area on which all the active sites are occupied. CONCLUSIONS The salient points raised in this discussion are summarized as follows: (I) The decarboxylation kinetics can be classified into two groups, namely: those which are first order in acetic acid or acetate concentration and occur in the presence of stainless steel, titanium, and gold and show no surface area dependence; and those which are zero order, occur over silica and synthetic magnetite, and may, or may not, be surface-area dependent. The data for each group can be unified by remarkably linear isokinetic relationships (Fig 5) which enable the kinetic boundaries of acetate decomposition to be considered in terms of the activation parameters, AH* and AS*,
rather than the individual rate constants. In other words, one set of AH? ASS values describe the rates as a function of temperature for a given order. (2) Based on the current experimental results, it would appear that both acetic acid and acetate ions possess similar inherent (uncatalyxed) reactivities. (3) The mechanism favored to rationalize the firstorder results is derived from a reaction sequence where surface adsorption is a function of the solubility of the solid catalytic phase and depends on the concentration of acetic acid, and in specific cases acetate, in solution. For zero-order kinetics the solubility is not dependent on the acetic acid concentration and the limited number of surface active sites are presumed to be completely covered by adsorbed acetate. (4) The principal products of acetic acid decomposition are generally CH., and COr, but varying amounts of CO, Hr, and volatile hydrocarbons were also observed. Their relative abundances appear related to the nature ofthe active surface but not the reactivity of the surface. (5) The result for Expr. 33 involving butyric acid indicates that higher carboxylic acids than acetic acid may possess similar reactivities and exhibit similar behavior although the range of products is undoubtedly wider. Acknowfedgemenrs-F. W. Dickson provided initial technical assistance. Y. K. Kharaka released data from his decarboxylation experiments prior to publication that helped guide our early etforts. The research benefitted substantially from consultation with M. L. Poutsma and from several analyses performed in his laboratory. For all their assistance, we gratefully acknowlec&e our debt. Financial support for this work was provided by the Offtce of Basic Energy Sciences. U.S. Depmtment of Energy, under contract DE-ACtX-84OR2I400 with the Martin Marietta Energy Systems. Editorial handling: R. P. Philp
REFERENCES BAMPORDC. H. and DEWARM. J. S. (1949) The thermal decomposition of acetic acid. J. Chem. Sot. (B). 2871-2882. BARYWJ M. A., BOWKERM. and MADIX R. J. (198 I) Formation and decomposition of acetate intermediates on the Ag( I IO) surface. J. Cataf. 67, 1I8- 128. BENZIGER J. B. and MADIXR. J. (1980) Reactions and rcaction intermediateson iron surkes. II. Hydrocarbons and carboxylic acids. J. Catal. 65.49-58. BLAKEP. G. and JACKSONG. E. ( 1968) The thermal decomposition of acetic acid. J. Chem. Sot. (B). 1153-l 155. BLAKEP. G. and JACKSONG. E. (1969) High- and low-temperature mechanisms in the thermal decomposition of acetic acid. J. Chem. Sot. (B). 94-96. BUSEYR. H. and MESMERR. E. (1977) Ionization equilibria of sihcic add and polysilicate formation in aqueous;odium chloride solutions to 300°C. Inore. C/tern. 16.2444-2450. DRUMMOND S. E. ( 198I) Boiling and mixing of iydro&rmal fluids: Chemical efkcts on mineral precipitation. Ph.D. dissertation, The Pennsylvania State University, 380p. DRUMMOND S. E. and PALMER D. A. (1985) Formation constants for aqueous ferrousacetate compkxes horn magnetite solubility measurements from 100” to 250°C and 250 to 1250 bars (abstr.]. Geol. Sot. Amer., Absiracts wifh Programs 17,561.
Acetate decarhoxylation. Part I
823
catalytic decomposition of formic acid. In Advances in CaR. and BARNESH. L. (1972) The ion-product confalysis (eds. D. D. ELEY. H. PINES and P. B. WEISZ). pp. stant of water to 350°C. J. Phys. Chem. 76,90-99. 35-l 13. Academic Press. GONZALEZF., MUNUERAG. and PRIETOJ. A. (1978) MechMOOREJ. W. and PEARSDN R. G. (198 I) Kinetics and Mechanism of ketonization of acetic acid on anatase TiOr suranisms. 3rd ed.. Wiley, 178~. faces. J. Chem. Sot.. Faraday Trans. 74, 15 17-l 529. KHARAKA Y. K., CAROTHERSW. W. and ROSENBAUER RIMSTIDT J. D. and BARNES H. L. (1980) The kinetics of silica-water reactions. Geochim. Cosmochim. Acra 44, 1683R. J. (1983) Thermal decarhoxylation of acetic acid: Im1699. plications for origin of natural gas. Geochim. Cosmochim. WILKINS R. G. ( 1974) TheStudy qfKinefics and Mechanisms Acta 47, 397-422. qf Reactions cf Transition Melal Complexes. pp. I OO- IO1. KURLKXXE J. C. and JEWURS. S. (1977) Studies on the surface Allyn and Bacon. Inc. interaction of acetic acid on iron oxide. J. Caral. SO, 330YING D. H. S. and MADIX R. J. (1979) Thermal desorption 341. study of the acetic acid decomposition on clean Ni/Cu( 110) MARS P.. SCHOLTENJ. J. F. and ZWIETERINGP. (1963) The alloy surfaces. J. Carol. 60.44 l-45 I FISHER J.
A. PALMER and S. E. DRUMMOND
Chemistry Division. Oak Ridge National Laboratory, Oak Ridge, TN 3783 1. U.S.A. (Received Uc&hr 16, 1985; accepted in revised.form February 14, 1986) Abstract-In an effort to understand the kinetics of the thermal decarboxylation of acetate and the role of catalysis, a series of laboratory experiments were conducted to measure the rate constants for the decomposition of acetate (acetic acid and sodium acetate) in the presence of titanium, silica. stainless steel, gold, and magnetite. Activation energies for decarboxylation of acetic acid and acetate ion range from about 8 kcal mol-’ in stainless steel vessels to 69 kcal mol-’ in silica tubes. Extrapolated rate constants at 100°C for acetic acid differ by more than fourteen orders of magnitude between the experiments conducted in stainless steel and the catalytically least active titanium vessels. Gold and titanium were the least active catalysts for the acetic acid substrate, while stainless steel, silica, and magnetite showed marked catalytic effects. Methane and carbon dioxide were the predominant reaction products of most of these ex~~rnent~ although mass spectrometric analyses of the gas phase revealed concentrations of carbon monoxide and hydrocarbons (apparent mass range from 29 to 56) amounting to as much as 55 mole percent of the total volatile products, depending on the catalyst. The reactions were generally first order in acetic acid or acetate ion, except for those involving the acid over silica and magnetite which were zero order. These results and the observed effects of variations in surface area are rationalized in terms of changes in the mode of surface catalysis. The mechanistic ~ignment is simplified by the existence of three unique straight lines on an isokinetic plot (Le., activation enthalpy versusactivation entropy) which lit all the respective first- and zeroorder reactions. The results described here provide the nucleus for the discussion in Part II of the role of acetate in the primary migration of methane and the transportation of metals in hydrothermal solutions. INTRODUCTION RECENTLY, KHARAKA ei al. (1983) demonstrated ex~~rnen~ly for the first time that acetic acid will readily decompose to carbon dioxide and methane in
the range of temperatures common to sedimentary basins. However, these kinetic data reveal an abnormally low activation energy (8.1 kcal mol-‘) for a reaction involving the breaking of a Karen-carbon bond and suggest that acetic acid decarboxylation was catalyzed in these experiments. Our experimental results confirm that surface catalysis is predominant under their conditions and demonstrate that many complexities exist in the decarboxylation of acetate that may not have been recognized by mARAK.4 et al. These new data suggest that acetate may play an impo~ant
role in the
primary mi~tion of natural gas and the insertion of metals in hydrothermal systems. Knowledge of the high temperature behavior of aqueous acetate solutions may also have important applications in problems associated with metal transport and corrosion in commercial steam generators. EXPERIMENTAL Kinetic experiments were carried out in a variety of vessels over a temperature range of 300 to 440°C. The vessel materials were as follows: 32 1stainless steel, gold (a gold bag suspended in a stainless steel vessel), titanium (commercially pure), titanium alloy (6% Al, 4% V), fused quartz and Pyrex tubes. Prior to each experiment the vessels were “conditioned” to ensure mp~u~bility. The 32 1stainless steel vessel was heated in air for several hours at 300°C to oxidize the surface. The gold bags were annealed in a Bunsen flame. The silica and quartz tubes (0.9 cm O.D., 0.3 cm I.D., 20 cm length) were washed for 30 minutes in a 5% HF solution. The titanium
vessels were thoroughly oxidized by heating them to incandescent red in a Bunsen flame. As a fmal preparation, all of these vessels were rinsed in distilled water and dried. In each experiment, the solutions were injected into the vessel at room tern~mtu~ and heated to the experimental temperature in a tube furnace. Before injection the titanium vessels were evacuated to (50 microns of pressure, whereas the gold bag and silica tubes were purged with argon to eliminate air. Neither of these measures were taken during the loading of the 321 stainless steel vessel (I.D. 3.5 em, internal length 28 cm) with acetic acid. The catalytic effects of various materials were measured in titanium vessels. The pure titanium and titanium alloy vessels gave identical results and hence could be used interchangeably. However, it was found that preheating the titanium alloy vessels to 400°C for ca. 24 hours while containing glacial acetic acid substantially pass&d the surfaces with respect to acetic acid decarboxylation. Three experiments were completed at 440.420, and 400°C before the surface began to return to its original reactivity. These experiments are referred to as taking place in “treated titanium.” Two experiments (numbered 16 and 171,involving M. 1m acetic acid, were carried out in the pure titanium vessels which contained one and ten tubes of 316 stainless steel (0.48 cm C&D.,0.24 cm I.D., and 10.0 cm long), respectively. The tubing had been previou~y cleaned with con~entmted HCI and was completely immersed in the acetic acid solution during the experiment. In experiment 18. a well oxidized 10.2 cm length of 1.6 mm O.D., 0.8 mm I.D. 3 16 stainless steel tube was immersed in ca. 1 m acetic acid solution. In order to assess the surface area effect on the corresponding reaction of acetate ions, experiments 3 I and 32 were conducted in a 316 stainless steel vessel (4.5 cm I.D.. I 7.8 cm internal length) containing eighteen 3 16 stainless steel tubes (0.62 cm O.D., 0.28 cm I.D., and 14 cm long). Most of the experiments with acetic acid solutions were monitored by titrating s&&on samples with a standard NaOH solution to the phenophthalein end point using a microburette with a readability of one part per thousand. This technique gave an accurate measure of the amount of acetic acid remaining and proved to be the most expedient method of fol813
814
D:A. Palmer and S. E. Drummond
lowing the decarboxylation reaction with time. Some of the initial experiments were monitored by measuring the pressure increaseamociatedwith the production of methane and carbon dioxide. It should be noted that this technique is normally subject to larger uncertainties which are not due to a lack of sensitivity in the pressure measurement, but rather to the difficulty in estimating both the solubihties of the product gases and the volume of tbe vapor phase. Calculating the latter is particularly arduous for experiments with the tlexible gold bag. The solubilities of CH, and CO2 were obtained from DRUMMOND (1981). The extent of the reaction of sodium acetate was determined exclusively by potentiometric titration. This involved titrating a sample, typically 1 cm3, to pH 2.00 with 1 m ha with rapid stirring to liberate bicarbonate as CO2 . The samples were then diluted to cu. 10 cm3 and backtitratcdwithstandardlmNaOHtoapHof11.OOtodetermine the total acetate concentration. These results were corrected for the dissociation of water and were reproducible to within 1%.Finally, a few scattered measurements of total acetate by ion chromatography were taken to confirm the accuracy of the other methods. At the termination of each experiment, the gas phase was also sampled at temperature and subsequently analyxed by mass spectrometry. The temperature of all these experiments was generallycontroUedto within 0.5”C. Stock solutions were prepared from 99.9% glacial acetic acid (Mallinckrodt, AR grade) or AR grade sodium acetate (F&herScientific) and deionized, doubly distilled water. Fine grainad,synthetic magnetite (Aha Producta 99.5%purity) was used without further purification. However, an X-ray powder diffraction pattern of this material failed to reveal any impurities.
RESULTS The decarboxylation of aqueous acetic acid can be represented by the general equation: [CHFOOH],
-
CH, + CO2
+ {CO + H2 + CnH2n+2 + - - . }
(1)
A summary of the pertinent experimental conditions is presented in Table 1. The measured molalities of acetate as a function of time ate given in Table 2 and the resulting rate constants are recorded in Table 3. Figures 1 and 2 show plots representative of the firstand zero-order reactions involving sodium acetate in stainless steel (Expr. 31 and 32) and acetic acid over magnetite (Expr. 20 and 2 1), respectively. KHAMKA et al. (1983) repotted a first-order rate constant for the decarboxylation of cu. 0.52 m acetic acid at 3OO’C in a stainless steel vessel of 8. I7 X lo-’ s-‘, approximately twice the value obtained from experiments 16 and 17 (Table 3). Extrapolation of their data to 349 and 359°C resulted in values of 1.42 X lo-’ and 1.58 X lo-’ s-t, respectively, which are conversely about half the value of the rate constants derived from experiments 18 and 19 in Table 2. Considering the great variability observed in the decarboxylation rates with both the composition and condition of the exposed surfaces, agreement of this order is quite reasonable. Moreover, this agreement is despite the large differences in the surface areas of iron exposed in each case, and in turn demonstrates that the kinetics are una&cted by the presence of titanium in experiments 16through IKKHARAKA et al. also reported olxrving high concentrations of iron during the initial stages of decarboxylation resulting in the formation of copious
magnetite crystals on the vessel walls by the end of the experiment. Experiments 16 to 19 duplicated the same behavior. The results of experiments with sodium acetate and mixtures of sodium acetate/acetic acid are presented in Table 3. Note that although f&tARAKA et al. (1983) did not observe any reaction of a 0.52 m solution of sodium acetate in stainless steel at 300°C after 1430 hours, this is consistent with the high activation enthalpy determined in this study at higher temperatures for this process which would predict only 1.4% reaction under their conditions. It is particularly relevant to the later discussion that in experiment 32 a significant amount of recrystallimtion of the titanium oxide surface layer was observed. Furthermore, neither the acetic acid solutions in experiments 3 to 11, nor those involving exclusively sodium acetate (i.e., 26-28) had any observable effect on the titanium surface. As mentioned earlier, the gas phase was sampled in situ whenever possible at the end of each experiment and analyzed by mass spectrometry. The results are given in Table 4, where CH represents hydrocarbons other than methane with the exception of mass 56. Some of the masses listed in parentheses are presumably fragments of larger hydrocarbons cleaved in the mass spectrometer. The liquid phase was only analyzed in one experiment (20) for nonvolatile organic mole cules other than acetic acid. A large array of organic molecules wem observed in the gas chromatograph of this solution with the dominant class of compounds being the C& alkenes. It is noteworthy that clear visual evidence was obtained in several experiments for the formation of a liquid organic phase after the reaction vessel had been cooled to room temperature. Experiments (4,6,7,8) over the temperature range 340-4OO”C in “untreated” titanium show that there is a discontinuity in an Arrhenius plot of the rate constants at or near the critical point (cu. 372°C for 1 m acetic acid). Extrapolated rate constants from data below the critical point are about 20% greater than those extrapolated from above. This phenomenon may reflect the inherent difference between the gas- and liquidphase kinetics. In order to test the possibility that decarboxylation occurs homogeneously in the liquid phase by interacting with metal ions released from the vessel walls, experiments 22 and 23 were carried out. In the former, added ferrous ion had no apparent influence on the rate which is compatible with the extrapolated rates observed previously in titanium vessels (Expr. 3 and 6). Similarly, chromium(II1) ions, which could be leached out of stainless steel by acetic acid solutions, had no measurable effect on the tate of decarboxylation (cf: Expr. 3 and 23, Table 3). Conceivably other metal ions in solution such as nickel@) could play a catalytic role, but the consensus of evidence points to a dominant heterogeneous path for decarboxylation. A typical Arrhenius plot of the rate constants for the decarboxylation of sodium acetate over titanium (Expr. 26-28) is presented in Fig. 3. The complete set
815
Acetate decarboxylation. Part I Table 1.
Summary of Experimental Conditions Prevailing in the Kinetic Runs
Expr. Temp.
Vessel Material
Volume of Vessel
Surface1 Area
Weight Initial Of Halality Solution
Method*
Acetic Acid :
357.2 359.0
3 4 5 6 7 8 9 10 11 12 13 14 15 lb
359.0 359.0 359.0 340.0 380.0 400.4 400.4 421.8 440.0 340.0 359.0 340.0 359.0 300.0
I7
3no.0
1R
359.0
19 20
349.4 359.0
gold bag gold bag titanium titanium titanium titanium titanium titanium treated titanium treated titanium treated titanium silica silica silica silica titanium (1 ss tube) titanium (10 ss tubes) titanium (1 ss tube) stainless steel titanium (magnetite, 59) titanium (mgnetite. 59) titanium (F&l,, O.OOlL) titanium (chromic nitrate, 0.0001 m)
166 178
64 68
64 35 64 64 64 64 64 64 64 1 1 1 1 35
80 63 40 49 109 109 109 109 109 15 18 15 18 23(ss)
25 15 15 25 15 12 12 12 12 0.5 Il.5 0.5 0.5 15
35
229(ss)
170
150 250 64
b(ss) 209 6x106 (ma9.j 6x106
t1tr.
0.987 1.054 1.064 0.0988 9.037 1.064 1.064 1.064 0.987 0.987 0.987 lb.39 16.65 1 .Ob4 1.068 0.987
pres. titr. titr. titr. titr. titr. titr. titr. titr. titr. titr. titr. titr. titr. titr.
15
0.987
titr.
60
1.054
pres.
100 25
1.054 0.989
pres. titr.
25
0.985
titr.
35
47(ma9.)
15
0.987
titr.
64
80
25
0.988
titr.
165 175 60
121 122 26
64 305 236
75 76 189 905
32 30 140 130
1.003 1.003 1.003 1.003 1.003 1.003 1.003
tjtr. titr. titr. titr. titr. titr. titr.
236
904
122
1.003
titr.
15
0.5401 0.540
titr.
31
1.097
titr.
64
21
340.0
22
300.0
23
359.9
24 25 26
359.0 381.0 340.0
gold cup gold CUP titanium
299 299 64
:; 29 30
381.0 359.0 359.0 359.0
31
381.0
titanium, titanium stainless steel stainless steel (1R ss tubes) stainless steel (18 5s tubes)
Sodium Acetate
Acetic Acid/Sodium Acetate 32
359.0
titaniwn
35
55
n-Butyric Acid 33
359.0
titanium
64
100
1
Surface Area represents the approximate area exposed of the dominant catalytic ccinponent prior to smnpling.
*Methods: titr. refers to titration against a standard NaOH solution; pres. refers to monitoring the pressure increase. 3
Inltlal twlalities refer here to concentrations at room temperature. i.e.. without correction for loss of water to vapor phase at temperature.
of these plots is represented in Fig. 4 for the zero- and first-order reactions listed in Table 5 incorporating those temperatures of immediate geological interest (see Part II). The activation parameters in Table 5 were determined according to the Eyring-Polanyi relationship, (MOORE and PEARSON, I98 1) Eqn. (2), and are shown together with the extrapolated values of the rate constants at 100°C. k = &T/h)
exp(--AH*/RT)
exp(ASIR).
(2)
In Eqn. (2) k is the observed rate constant; k, is the Boltzman constant; T is the temperature, K; h is Planck’s constant; AH* is the enthalpy of activation; AS* is the entropy of activation; and R is the universal
gas constant. The previously published results for decarboxylation over silica included in Table 5 refer to supercritical acetic acid in the absence of water. The relationships between the activation parameters are shown in Fig. 5. DKXUSSION One of the original goals of this endeavor was to establish the minimum, or least catalyzed, rate of decarboxylation of acetate. This rate would then form the basis for studying the effects of catalysis and, in addition, would provide the ultimate lower kinetic boundary for a natural gas migration model, as well as the maximum lifetime for acetate as a potential li-
D. A. Palmer and S. E. Drummond
816 Table 2.
Hours Hours
mlality
klality
Hours
Expr.
19.5 79.5 170.5
0.987 0.976 0.963
240.5 337.5 418.3
Hours
1.040 1.040 LO47 1.045 1.043 1.046 1.035 1.040 1.033 1.035 1.040 1.035 1.032
399.5 427.5 471.5 :~~*x 595:o 620.0 644.0 7b1.0 784.5 808.0 835.0 907.5
1.031 1.030 1.022 1.018 1.013 1.011 1.007 1.006 0.998 0.994 0.994 0.989 0.981
307.5 580.0 680.0
0.914 0.908
0.969 0.947 0,929 0.912
1.042 1.043 1.017 0.992
527.0 954.5 981.0 999.5 K8 11Ob:O 1125.0 1152.0 1169.5 1222.5 1240.0
0.980 0.971 0.974 0.974 0.961 0.963 0.962 0.961 0.959 0.954 0.945 0.945
1339.0 1510.0 1511.0 1637.0
0.895 0.872 0.867 0.865
Expr. 4 67.0 166.0 233.0
0.0855 0.0766 0.0713
328.5 428.3
72.0 159.0 165.0 242.5
9.846 9.395 9.366 9.280
333.0 333.n 405.5 501.0
0.0637 0.068§
501.0 569.0
0.0531 0.0481
556.5 775.5 914.0
8.631 8.097 7.965
Expr. 5 9.023 9.023 8.687 a.525 Expr. 6 159.0 663.0
1170.5 1698.0
1.018 0.984
2173.5 2653.5
I.059 1.062 1.053
1.041
153.5
1.023 1.012 0.998 0.978
:::*: 3so:o Expr,
140.5 209.5
1.010 0.967
312.0 405.0
417.5 541.5 684.5
0.962 0.935 0.901
0.878 0.841
379.5 503.0
0.835 0.816
477.5 573.0
0.955 0.949
89.5
0.980
0.563 0.213
122.0
fxpr.
Expr.
965.5 1274.0
312.5 601.0
313.0 601.0
0.967 0.943
22.5 49.0 67.0 75.5
1.054 0.978 0.956 0.948
115.0
965.0 1274.0
k!
114:o 123.3
gin 96.5 112.0
ii-;: a:934 0.922 0.904 0.892 0.887 0.883
114.5 146.0 163.0 168.0 184.0 2oR.o 234.0 304.5 328.0 333.5
71.o 94.0
0.729 0.690
143.0 167.3
16.0 22.3 39.5 t?: ;:I:
1.043 0.961
Expr.
Expr.
c:
95.0 118.5 142.0 163.0 264.5
.
a:::
0.913 0.882
94.0
0.988
189.5
16.49 16.4S1 16.M 16.20
333.0 333.0 it:*:*
15.52l 15.79 15.6g1 15.26
1.0
0.891
1.097 1.062 1.023 0.944 0.875
1.277 If *oOz-t:; 23:s
27.5
::z .
136.5
15.52
136.5
0.105
1613.5
0.777
1613.0
0.809
139.0 170.5 235.3 259.0
0.879 0.862 0.809 0.801
zz-8 373:o 384.0 399.0 408.0 42b.0 449.0 471.5
0.692 0.684 0.684 0.670 0.670 0.657 0.645 0.636 0.629
236.0 303.5
0.368 0.223
333.0 426.0 496.5 576.3
0.485 0.307 0.242 0.182
:::“33.
0.977 0.973
16 0.881 0.830
0.908 0.857
0.931 0.918 0.907 0.898
0.875 0.850 t-iii 0:82O 0.801 0.778 0.729 0.712 0.699 20 0.600 0.542 21 0.704 0.690 0.664 0.652 0.566
23 0.963
359.0
0.950
:fF: 140:5 164.5
oh79
42.5 51.0 66.5
0.830 0.796 0.744 0.691
192.5 261.0 286.5 308.5
Xii 0:513 0.485
75.0 75.5 98.5
0,434 0.417 0.301
618.0 721.0
0.742 0.703
253.5 301.5
0.643 0.584
25 0.717 0.632 0.489
Expr. 26 525.0 648.0
15.47 lb.39
22.0 219.0 382.5
1.037 0.959 0.871
:::-: . Expr.
Expr. 13 93.0 122.0
1.0031 0.160
Expr. 24
Expr. 12 43.5 68.5 168.0 193.0
355.0 525.0
15
0.210
0.991 0.990
f$t: . Expr.
2255 51:5 81.0
0.841
0.631 0.506
Expr. 22 0.987 0.995
Expr.
it:::.
14
Expr. 19
0.798 0.763
Expr. 11 0.980 0.955
wr11ty
Expr. 18
20.5 643.0 808.5
Hours
Expr. 17
0.843 0.797
Expr. 10 ::*:0
93.0 121.5
312.0 312.0
Expr. 9 68.3 309.5
312.0 355.0
17.0 23.5 40.5
8 0.918 0.875
1,060 0.925
0.955 0.933
Expr. 7 I!$ If:5 86.5
lklrlfty
43.5 193.0
0.886
E!cpr. :z 331:5 498.5
Hours
Expr.
Expr. 2 87.3 120.0 161.5 167.0 213.0 237.0 265.0 283.0 304.0 312,5 336.0 352.0 386.5
Rolality
rnolality
1
0.953 0.941 0.928
tonttnued
Table 2.
Smmry of Kinetic Data for the Decarboxylatlffn reaction
136.6
15.51
14.0 43.5 73.0
:z ‘
0.837 0.792 27
z:: .
817
Acetate decarhoxylation. Part I Table
Hours
2.
HOUPS
Molality
Continued
Holality
Expr.
1;:: 26.5
Expr.
3:::
Expr. 2.5 18.0 49.5
1.024 0.966 0.860
71.5 116.0 138.5
::: 3.5 5.5 7.5 8.0 11.0
1.044 0.998 0.925 0.909 0.907 0.822 0.766
11.5 23.0 23.5 27.5 28.0 32.5
0.0 28.0 46.5
65.0 70.5 102.5
2.0 25.5 91.0
1
These
1.027 1.035 0.972
samples
138.0 193.0 265.5
were
taken
2The molallties in Expr. acetic acid content.
162.5 186.5 213.0
0.563 0.519 0.482
47.5 48.0 54.0 94.5 95.0 121.0
0.355 0.297 0.321 0.122 0.129 0.0845
139.0 140.0
0.029 0.018
457.0 672.0 935.5
0.922 0.888 0.856
33 0.959 0.963 0.946
from
0.513
32 0.120 0.094 0.018
Expr.
202.5
31 0.630 0.655 0.550 0.541 0.491 0.370
Expr.
0.182 0.103 0.095
30 0.784 0.655 0.602
Expr.
142.0 211.0 213.5
29 0.871 0.615
85.5 178.0
1.106 1.082
Holality
28 0.611 0.484 0.321
41.5 66.0 98.0
1.045 O.R28 0.738
Hours
quartz
stainless steel. In acidic solution, iron is readily soluble as evidenced by the high levels observed initially in acetic acid solutions during this study and reported by KHARAKA et al. (1983). e.g., 329 (300°C) and 5600 (200°C) mg L-‘. In contrast. iron concentrations in the sodium acetate solutions were < 1 mg L-’ and decarboxylation was found to be orders of magnitude slower. An explanation of this phenomenon is that the dissolution, or oxidation. of iron is mainly promoted by the acetic acid content of the solution due to the formation of [Fe&]aq which is predicted to be the dominant iron acetate complex at these conditions (DRUMMONDand PALMER, 1985) and is independent of the acetate anion concentration. The act of dissolution produces active sites on the walls of the vessel, such that the overall rate of decarboxylation is governed by the rates of adsorption/desorption of acetic acid molecules onto these sites and the rate of the subsequent decomposition of the adsorbed species. The implicit assumption being that decomposition. the actual breaking of the C-C bond, occurs predominantly at the active sites on a given surface. This reaction sequence is consistent with the earlier observation reported by WARAKA et al. ( 1983) that in an equimolal acetic acid/sodium acetate solution in stainless steel only the acetic acid component reacted during the course of the experiment at a rate which was first order in acetic acid, rather than total acetate. In other words, the total number of active sites generated is related to the solubility of iron which, at these experimental con-
tubes.
32 refer exclusively
Tahle 3. Constants
to the
gand in a metal transport model. The rate constants shown in Table 3 establish that the slowest reaction of aqueous acetic acid is in “treated” titanium. The high activation enthalpy and near-zero entropy (Table 5) indicate minimal catalytic influence, similar to the corresponding values reported at higher temperatures and low pressures for the same process in a nonaqueous environment (BAMFORD and DEWAR, 1949; BLAKE and JACKSON, 1969). Indeed, BLAKE and JACKSON ( 1969) demonstrated that free radicals were not formed under their conditions and concluded that the reaction is therefore homogeneous, which in this case implies uncatalyzed. The apparent indifference exhibited by the rate constants for acetic acid decomposition to either the presence of added ferrous chloride (Expr. 22), or the varying concentrations of ferrous ion resulting from the initial dissolution of stainless steel and the subsequent precipitation of magnetite in Expr. 19 provides strong evidence for a heterogeneous mechanism, which is controlled by reaction sites at a surface, at least for those reactions carried out in stainless steel. Consider first the implications of the pH effect on the heterogeneous mechanism for decarboxylation in
Expr. No.
Temp. “C
First and Zero-Order Rate for the Decarbowlation of Acetate
First-Order
Zero-Order “S
: 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 1R 19 20 21 22 23 24 25 2’; 28 29 30 31 32 33
359.0 357.2 359.0 359.0 359.0 340.0 400.4 380.0 400.4
4.28 2.51 3.89 3.08 7.19 1.55 6.64 1.52 4.72
x x x x x x x x
421.8 440.0 340.0 359.n 340.0 359.0 300.0 300.0 359.0 349.4 359.0
1.99 5.03
x 1o_7 x 10
340.0 300.0 359.0 359.0 381.0 340.0 359.0 381.0 359.0 359.0 381.0 359.0 359.0
5.03 3.83 2.A9 2.94
co.3 3.2 7.13 4.15 1.57 5.09 3.08 1.00 1.02 6.30 6.75 4.20
x x x x
x x x x x x x x x x x x
k-l
1O-8 1O-8 10-8 lo-’ 10-8 10-8 10:; lo_, 1O_8
10-8 1O-8 10-l 10-7
10-8 lo-’ 10-7 10-6 10-7 10-7 lI+ 10-6 Ml-6 1O-6 1O-6 1O-8
7.97 2.31 6.39 3.02
x 10:; x x lo_, 1O_6 x 10
6.15
x
3.18
x 1O_7 10
-7
D. A. Palmerand S. E. Drummond
0.8 0.6
s”
-0.8 -LO
0
40
20
80
60
100
TIME(h) FE. 1. First-order plot for the reaction of sodium acetate
in stainlesssteel at 38 I"C(Expr.3 1).
ditions, is controlled largely by the acetic acid content of the solution. A direct consequence of this hypothesis is that the hrst-order rate constants should be independent of surface area, given that there is sufficient solid to satisfy the amount of acetic acid present. All the first-order reactions studied here support this mechanism as the results show no evidence of a surtkce area dependence. This mechanism can be formulated Table 4.
Expr. No.
"Active Surface'
Tcnp. OC
gold goldb tjtan. titan. titan. titan. titan. silica quartz silica silica silica 5s 5s 55
nag. nag. titan.
gold gold titan. titan. titan.
359.0 357.2 359.0 359.0 340.0 400.4 400.4 340.0 340.0 359.0 340.0 359.0 300.0 300.0 349.4 359.0 340.0 300.0
Mstribution
Curation of Expr. h
in terms of a general rate law based on the following reaction sequence, namely, S (solid) + HAc +
Of Gas Pbasc Products1
CM4
co*
ketic
kid
co
H2
820 1240 1676 1255 265 651
41.6 22.5 20.3 10.3 31.5 10.3
42.8 48.9 47.6 71.3 60.0 47.2
6.6 25.3 7.0 5.4 2.5 4.5
3.1 1.1 3.9 0.9 0.4 5.2
:: 333 137 525 137 1274 1274 316 766 577 1613
35.7 0.9 9.1 15.1 22.4 32.2 2.R 5.3 49.1 lb.3 5.2 3.2
42.3 52.4 15.5 54.3 28.5 35.2 15.6 46.3 45.1 63.2 63.0 34.4
4.5 7.8 20.0 7.2 4.6 5.1 1.1 4.9 4.3 2.3 3.8 53.8
6.1 0.3 0.8 0.3 0.3 CO.01 75.4 30.9 1.5 4.0 1.5 1.4
Sodim
ketate
309 143 789
60.3 80.4 60.6
18.3 34.6 17.3
3:4 0.2
::: 21.3
titan.
305 214 203 213 121 167
7A.4 79.2 79.5 85.9 76.2 37.3
12.9 lb.7 20.0 12.6 19.6 59.9
0.1 0.1 0.1 0.1
:*: 0:5 1.2 4.0 2.6
titan.
359.0
940
19.3
21.6
5s OS
n-Butyrlc
%mallred vapor are
7.0
CH(masses)
&
0.5 1.0 2.2 3:4 0.3 0.8 2.9 10.3 2.3 3.7 3.3 0:5 2:b 015
5.4(29.44.55) 1.3(48-55.57) 18.0(29.41.55) 12.1 2:::(29.44.55) 10.6(29.44.55) 27.7(29.41,55) 44.3129.41.55) 20.8(29.41.55) 40.5(29.41.55) 24.2(29.41.55) 5.1f29.44.55) 5.4(29,44,55) 10.7(29.41.55) 26.5 b.bt29.41.55)
kid 52.0
to 100%. 1 .c ., the coexisting gases such as k. not included.
Acetic acldlsodium acetate mixture.
HAC& 2 products
1
S (solid) + AC-+A&
359.0 381.0 340.0 359.0 3Rl.O 359.0 359.0 381.0 359.0
55
2
RG. 2.Zero-order plots for the reaction of acetic acid over commercialmagnetiteat (a) 340°C (Expr.21) and (b) 359°C (Expr.20).
112, and water
2 products
(3) (4)
Acetate decarboxylation. Part 1
either K, > K3 (k:1 = k$/k-3) or k2 > k4 or both. It is relevant that the activation parameters associated with the thermal decomposition of acetic acid preadsorhed on a Fe( 100) surface (BENZIGER and MADIX. 1980), AH* = 32.5 kcal mol-’ and AS* = 2.1 cal K-’ mol-‘. fall well outside the error limits for the first-order straight line correlation in Fig. 5. As these parameters reflect the expected values for the liz step in Eqn. 3. this failure to correlate indicates that adso~t~on/desorption the~~ynamics associated with 4.Vtdo contribute to the activation parameters in Fig. 5. It should also be mentioned that the remarkable linearity exhibited by the three sets of data presented in Fig. 5 (namely. for first-order in acetic acid, first-order in acetate. and zero-order in acetic acid) is indicative of a common mechanism prevailing in each case (WILKINS. 1974).
-NOc)
C
-6.0
-6.5
-7.0 l.50
160
l.55
l.65
1000/T FIG. 3. An Arrhenius plot describing the temperature dependence of the reaction of sodium acetate in titanium (Expr. 26-28).
Making the reasonable assumption that the activated sites, or adsorbed acetic acid, reach a low, steady-state concentration leads to the general rate law: Rate = d[HAc]/df = -{k’&,/(k_,
+ k2)}[H/lc]
819
(5)
where k’, = k,S, with the corresponding expression for the rate of acetate ion decomposition. Given the validity of this rate law, the measured first-order rate constants in Table 3 are really composite values of the form k’,k2/(k_, + k2) with k2 most likely representing the slower, rate determining step, such that this constant term reduces to K~k~(=k~k~~k_,). For acetic acid to react substantially faster than acetate ions, with the corresponding large difference in activation enthalpy,
The following discussion provides an explanation of the kinetic behavior observed in titanium. The sol-
ubility of titanium oxide is less affected by acid than by complexation with acetate ions, and hence the rate constant verszts pH profile is a complex function of the acetic acid and acetate concentrations (~6 Expr. 3. 27 and 32. Table 3): keeping in mind that the dissolution of titanium molecules is associated with the formation of active sites presumably by exposing titanium metal. The maximum solubility of titanium oxide corresponds to the highest concentration of titanium acetate complexes, and is presumably reached near equimolal ratios of acetic acid and acetate ions (Expr. 32). This is reflected by the highest rate constant observed at this condition, namely 6.75 X 10m6s-’ {
FIG. 4. A summary of the temperature dependencies of the reactions presented in Table 5 extrapolated to in&de temperatures of geological interest. where the legendis as folfows:(a) sodium acetate over magnetite: (b) 16.5 m acetic acid in s&a; (c) 1.06 m acetic acid in silica; (d) acetic acid in stainless steel, KHARAKA et al.: (et suoercritical acetic acid in titanium: (f) acetic acid in tj~niurn; (g) sodium acetate in ti~nium: (h) so&&n acetate in stainless steef; (i) sodium acetate in gold; (k) acetic acid in treated titanium. The units of k are m s-’ for zero order and s-’ for the first-order ease. The logarithmic half-life (tre) scale is relevant to the zero-order plots only for an acetate concentration of 1.39 m.
820
D. A. Palmer and S. E. Drummond Table 5.
Expr. nos.
ktlration
Temp. Range
Farametcrs for the lkcarboxylation
k( lM)‘Y)
“Active Surface"
ms
s-1
-1
of ketate
AH* kca!1 no1
As*_ 1 cal _K1 nO1
Acetic Acid 3413-359
titan.
5.5 x 10-17
35.9
-37.1
titan. tr. titan silica silica SiliCS1 silica2 si ica3 ss 1 mag.
9.8 5.9 x 10-17 10-23
34.2 ::*; 61:4 55.8 60.3 63.7
20.21
400-440 380-400 340-359 340-359 494-585 642-700 767-916 200-300 340-359
2:::
-10.9 -40.5 -19.5 11.8 -14.2 -0.2 -5.2 -79.6 -51.9
24.25 26-28 30.31
359.381 34ll-381 359-381
gold titan. ss
1.4 x 10-22
64.5 56.8 66.8
+13.9 1.4 *la.2
9111 :*: 12.13 14.15
1.2 x 10-16 3.2 x 1O-21 1.2 7.0 2.1 1.8
x x x x
10-23 10-28 10-26 10-9
8.0 x 10-10
Sodiun Acetate
'(Blake and Jackson, 1968). no water vapor added; static technique with p(acetic acid) - 200 m in unpacked silica tubes. '(Blake and Jackson, 1969). no water vapor a&dad; flow technique through "mature" silica tubes. '(Banford and C&war. 1949). no water vapor added; flow technique through carbon coatad silica tubes. 4(Kharaka et al., 1983).
and acetate concentrations, the formation of Ti(Ac)S and higher titanium acetate complexes must be evoked in accordance with the following general equation: Ti02 + 2HAc + H2 + nAc- = Ti(Ac)L
+ 2H20 (6)
E 0
FIG. 5. Isokinetic plot for the kinetics of decarboxylation of acetate in aqueous solution. The symbols are defined in Fig4,wiUlthcadditionthat99aadIXIrrp*lcntthereaaion of anhydrous acetic acid in silica taken from BLAKE and JACKSON (1968) and ( 19691, and BAMFORD and DEWAR (19491,mspectively. The solid line is a linear least-uquates fit of the first-orderacetic acid reactions, the chaindashed line apt&s to the tirst-order reactions of acetate ion, while the dashed line fits the xere-order results.
The important initial step in the dissolution process associated with Eqn. (6) can be envisaged as adsotption of HAc on the surface in accordance with Eon. (3). thus fulfilling the same requirements for a first-order reaction as laid down for stainless steel. The high concentration of titanium in solution would lead to the gradual deposition of Ti4 during the experiment as the concentration of acetic acid declines, resulting in the observed ‘YecryWliz.ation of TiO*” in Expr. 32 described earlier. Raising the initial concentration of acetic acid to 9.8 m (Expr. 5) increases the titanium complex concentration, which according to Eqn. (6) is propottional to [HAc]*[Ac-I”, giving rise to increased TiOz solubility compared to that at 1 m. Thus the number of potential active sites is larger and the rate of decarboxylation increases accordingly. Conversely, lowering the acetic acid concentration in Expr. 4 to about 0.1 m should decrease the concentration of titanium acetate in solution giving rise to a slower rate. However, the reaction is in fact faster than either the 9.8 or the 1 m solutions. As experiment 4 shows a good first-order linear dependence them is no reason to doubt the authenticity of this result and it remains an enigma. Gold is less soluble than titanium and iron in acetic acid solutions and hence fewer active sites are generated for adsorption and decarboxylation. Moreover dissolution of a surface gold atom merely exposes more inert gold atoms, whereas removal of relatively inert oxide layers for titanium and iron reveals highly active metal underneath which can readily be imagined as providing a more teactive surface for acetate to adsorb and react Thus the rates of reaction in the presence of gold tend to be closer to the background homoge-
Acetate decarboxylation. Part 1 neous rate. It also follows that if the oxide layers on titanium and iron remain intact, as presumably is the case in the presence of sodium acetate, as well as for acetic acid in treated titanium, the rates of reaction will approach that observed in gold. Comparing the rate constants extrapolated to 100°C (Table 5) for sodium acetate on all three surfaces and acetic acid on treated titanium, which must have a thicker than usual oxide layer protecting the reactive titanium metal, shows that they are all approximately similar. This leads to the conclusion that in weakly catalyzed environments acetic acid and acetate anions have similar “inherent” reactivities with the activation enthalpy reflecting mainly the energy required to break the C-C bond, within the weakly bound substrate. Having established a mechanistic framework to rationalize the pH dependence for acetate decarboxylation kinetics, it now remains to be seen how this hypothesis can be used to explain the observed changes in reaction order. First consider the somewhat analogous dehydrogenation of dry formic acid in the gas phase (HCOOH - Hz + CO*) studied in detail by MARS et al. ( 1963). The order of this reaction is known to vary from zero at low temperatures and high pressures to first at high temperatures and low pressures. This tendency was correlated with surface adsorption of formate being complete for the zero-order case, thereby ensuring that the rate-determining step involves the decomposition of the adsorbed intermediate and is consequently not sensitive to the bulk formate content. In the case of a first-order reaction, active sites are still available on the surface and the rate of adsorption, which is directly proportional to the formate content, controls the process. Accordingly, the activation energy varied from 15.8 kcal mol-’ for zero order to 6.2 kcal mol-’ for the first-order adsorption process. If the same criteria operate in the solution phase, increasing the surface area to concentration ratio for a given surface material should only accelerate the zero-order reaction, until it eventually leads to a change in mechanism from zero to first order. It is tempting to draw parallels between the kinetics of formic and acetic acid because there is ample evidence for strong adsorption of the latter (and acetate ion) on iron oxide (KURIACOSEand JEWUR, 1977; BENZIGERand MADIX, I980), titanium oxide (GONZALEZ et al., 1978) as well as other transition metal surfaces (YINGand MADIX, 1979; BARTEAU et al., 198 I). From the earlier discussion of the origin of the first-order acetate dependence gleaned from Eqn. 3, it is apparent that the analogy to the formate system is only partly correct in this case. However, from a consideration of the conditions which must prevail in this scheme to promote zero-order kinetics, it will be seen that the formate analogy does present a plausible explanation. For the reactions of acetic acid and acetate ions in stainless steel, all of which have been determined to be first order, no significant change in rate at a given temperature was observed despite substantial changes
821
in the ratio of surface area to moles of total acetate in solution. For example, in Expr. 16 and 17 the ratio varies from 0.99 to 9.9 cm* mmol-‘, while the rate constants at 300°C are quite similar at 5.03 X lo-’ and 3.83 X lo-* s-‘, respectively. Extrapolation of the rate constant for Expr. 18 to 300°C yields a value of 7.8 X lo-’ s-’ for a ratio of 0.072 cm* mmol-‘. indicating that this 138-fold change in area had no significant effect on the rate. Similarly. the rate constants for experiments 29 and 30 are not influenced by a 4.8-fold change in surface area. Thus in all these cases. and most likely for the first-order reactions in general, there is sufficient surface available to promote enough active sites to satisfy either the acetic acid or acetate concentrations in the solution. Here the analogy with the formic acid mechanism ends, with the first-order dependence arising from the formation of active sites coupled with the chemical decomposition of the adsorbed intermediate, rather than being simply due to a diffusion and adsorption of the reactant molecules to the surface. The high activation enthalpies observed for the reactions on gold and titanium. for example, would be incompatible with a diffusion/adsorption rate-determining process, cf: 6.2 kcal mall’ for formic acid. Attempts to reduce the surface area ratio to even lower values in order to promote an adsorption limiting rate process requires containing the acetate solution in a vessel of some other, more inert material. Unfortunately. at least at the temperatures where the reaction rate can be measured over a reasonable length of time, no vessel material could be found which did not impose its own catalytic effect thereby dominating that of the smaller surface of interest. Hence the postulate that the first-order processes can be converted to surfacearea dependent, zero-order paths could not be tested. Both zero-order reactions have unique features which set them apart and neither has yet been shown to be surface-area dependent. Firstly, the reaction of acetic acid over magnetite is unquestionably zero order as demonstrated by the clearly linear plots in Fig. 2. Nevertheless, in the corresponding reaction in stainless steel (Expr. 19) where a large amount of magnetite was gradually formed during the experiment, the reaction rate neither increased in proportion to the magnetite buildup, nor deviated from a slower, strictly first-order dependence. Therefore, it must be concluded that minor unidentified impurities in the commercial magnetite used in Expr. 20 and 2 I are responsible for the zero-order kinetics. This is consistent with the fact that the surface area of this fine-grained magnetite is orders of magnitude larger than found in any of the other experiments and should therefore promote a first-order reaction if Fe304 sites were involved. The second example of zero-order kinetics was found for those experiments conducted in silica tubes where it had been initially hoped that essentially uncatalyzed rates may have been observed. The silica tubes containing the one-molal solutions were strongly etched and substantial induction periods of about I30 h (Expr.
D. A. Palmer and S. E. Drummond
a22
14) and 50 h (Expr. 15) were observed. In the 16.5 m solutions the dissolution of silica was less apmfent and no corresponding induction period was seen. The dissociation constant for acetic acid at 359°C is approximately 1.8 X lo-* (FISHER and BARNES, 1972) while that of silicic acid is about 1.9 X lo-” (BUSEY and MESMER, 1977) and hence the solubility of silica in these experiments is mainly a function of the water activity rather than proton activity. The generation of active surface sites would consequently no longer be dependent on the acetic acid concentration. Assuming that these sites are all immediately occupied by acetic acid molecules, then the observed rate of decomposition would be independent of the total acetic acid concentration and zero-order kinetics would result. In the case of the reaction over magnetite, the number of sites is limited by the amount of impurities in the solid and these must also be considered to be completely occupied by acetic acid molecules. This rationale is similar to that proposed for the formic acid analog. An alternate explanation for the zero-order kinetics, at least over silica, can be found in the following reaction sequence: (SiOr)d
+ 2HrO $
active site + H4c 2
active site + (H&iO,),
(7)
(MC)& 2 products
(8)
Solving the rate expression for this reaction sequence with the assumption that the number of active sites is at a low, steady-state level and that ks > k, leads to, d[I-ZAc]/dt = -kI . However, from a study of the kinetics of silica dissolution, RIMSTIDTand BARNES ( 1980) reported an activation energy for this process of only 11.7 kcal mol-‘, far less than the observed AH* (AH* = E. - RT) listed in Table 5 for the zero-order decarboxylation reactions. Not only is this mechanism incompatible with the results, but a linear isokinetic plot would not be expected for the different solids, magnetite impurities and silica, if k, is rate determining Thus, until experimental evidence to the contrary becomes available, the zero-order behavior is best explained by the concept of a limited surface area on which all the active sites are occupied. CONCLUSIONS The salient points raised in this discussion are summarized as follows: (I) The decarboxylation kinetics can be classified into two groups, namely: those which are first order in acetic acid or acetate concentration and occur in the presence of stainless steel, titanium, and gold and show no surface area dependence; and those which are zero order, occur over silica and synthetic magnetite, and may, or may not, be surface-area dependent. The data for each group can be unified by remarkably linear isokinetic relationships (Fig 5) which enable the kinetic boundaries of acetate decomposition to be considered in terms of the activation parameters, AH* and AS*,
rather than the individual rate constants. In other words, one set of AH? ASS values describe the rates as a function of temperature for a given order. (2) Based on the current experimental results, it would appear that both acetic acid and acetate ions possess similar inherent (uncatalyxed) reactivities. (3) The mechanism favored to rationalize the firstorder results is derived from a reaction sequence where surface adsorption is a function of the solubility of the solid catalytic phase and depends on the concentration of acetic acid, and in specific cases acetate, in solution. For zero-order kinetics the solubility is not dependent on the acetic acid concentration and the limited number of surface active sites are presumed to be completely covered by adsorbed acetate. (4) The principal products of acetic acid decomposition are generally CH., and COr, but varying amounts of CO, Hr, and volatile hydrocarbons were also observed. Their relative abundances appear related to the nature ofthe active surface but not the reactivity of the surface. (5) The result for Expr. 33 involving butyric acid indicates that higher carboxylic acids than acetic acid may possess similar reactivities and exhibit similar behavior although the range of products is undoubtedly wider. Acknowfedgemenrs-F. W. Dickson provided initial technical assistance. Y. K. Kharaka released data from his decarboxylation experiments prior to publication that helped guide our early etforts. The research benefitted substantially from consultation with M. L. Poutsma and from several analyses performed in his laboratory. For all their assistance, we gratefully acknowlec&e our debt. Financial support for this work was provided by the Offtce of Basic Energy Sciences. U.S. Depmtment of Energy, under contract DE-ACtX-84OR2I400 with the Martin Marietta Energy Systems. Editorial handling: R. P. Philp
REFERENCES BAMPORDC. H. and DEWARM. J. S. (1949) The thermal decomposition of acetic acid. J. Chem. Sot. (B). 2871-2882. BARYWJ M. A., BOWKERM. and MADIX R. J. (198 I) Formation and decomposition of acetate intermediates on the Ag( I IO) surface. J. Cataf. 67, 1I8- 128. BENZIGER J. B. and MADIXR. J. (1980) Reactions and rcaction intermediateson iron surkes. II. Hydrocarbons and carboxylic acids. J. Catal. 65.49-58. BLAKEP. G. and JACKSONG. E. ( 1968) The thermal decomposition of acetic acid. J. Chem. Sot. (B). 1153-l 155. BLAKEP. G. and JACKSONG. E. (1969) High- and low-temperature mechanisms in the thermal decomposition of acetic acid. J. Chem. Sot. (B). 94-96. BUSEYR. H. and MESMERR. E. (1977) Ionization equilibria of sihcic add and polysilicate formation in aqueous;odium chloride solutions to 300°C. Inore. C/tern. 16.2444-2450. DRUMMOND S. E. ( 198I) Boiling and mixing of iydro&rmal fluids: Chemical efkcts on mineral precipitation. Ph.D. dissertation, The Pennsylvania State University, 380p. DRUMMOND S. E. and PALMER D. A. (1985) Formation constants for aqueous ferrousacetate compkxes horn magnetite solubility measurements from 100” to 250°C and 250 to 1250 bars (abstr.]. Geol. Sot. Amer., Absiracts wifh Programs 17,561.
Acetate decarhoxylation. Part I
823
catalytic decomposition of formic acid. In Advances in CaR. and BARNESH. L. (1972) The ion-product confalysis (eds. D. D. ELEY. H. PINES and P. B. WEISZ). pp. stant of water to 350°C. J. Phys. Chem. 76,90-99. 35-l 13. Academic Press. GONZALEZF., MUNUERAG. and PRIETOJ. A. (1978) MechMOOREJ. W. and PEARSDN R. G. (198 I) Kinetics and Mechanism of ketonization of acetic acid on anatase TiOr suranisms. 3rd ed.. Wiley, 178~. faces. J. Chem. Sot.. Faraday Trans. 74, 15 17-l 529. KHARAKA Y. K., CAROTHERSW. W. and ROSENBAUER RIMSTIDT J. D. and BARNES H. L. (1980) The kinetics of silica-water reactions. Geochim. Cosmochim. Acra 44, 1683R. J. (1983) Thermal decarhoxylation of acetic acid: Im1699. plications for origin of natural gas. Geochim. Cosmochim. WILKINS R. G. ( 1974) TheStudy qfKinefics and Mechanisms Acta 47, 397-422. qf Reactions cf Transition Melal Complexes. pp. I OO- IO1. KURLKXXE J. C. and JEWURS. S. (1977) Studies on the surface Allyn and Bacon. Inc. interaction of acetic acid on iron oxide. J. Caral. SO, 330YING D. H. S. and MADIX R. J. (1979) Thermal desorption 341. study of the acetic acid decomposition on clean Ni/Cu( 110) MARS P.. SCHOLTENJ. J. F. and ZWIETERINGP. (1963) The alloy surfaces. J. Carol. 60.44 l-45 I FISHER J.