blake1968 aa

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Phys. Org. ¹¹⁵³

The Thermal Decomposition

of

Acetic Acid

By P. G . Blake" and G. E. Jackson, Chemistry Department, University College, Cardiff CFI 3 N R

The homogeneous gas-phase thermal decomposition of acetic acid has been studied in a static system between 460 and 595°C and at pressures up to 1 atmosphere. A second-order decomposition to keten and water occurs simultaneously with a first-order process yielding methane and carbon dioxide, the respective Arrhenius equations being k = 1 OS'45exp(-34,200/RT) l.mole-l sec.-l and k ⁼1 011'12exp(-58,500/RT) sec.-I. The decarboxylation is similar to that pteviously reported in flow work at 770-920" but the dehydration processes are quite different.

In the static system it closely resembles dehydration in formic and propionic acids.

THE only reported kinetic study of the homogeneous thermal decomposition of acetic acid is by Bamford and Dewar,l who used a flow system at 770-920". They found that decomposition was represented by reactions (1) and (2). Both processes were first-order, the

CH,CO,H CH,CO

+

^{H 2 0} ⁽¹⁾

CH,C02H

--+

CH,

+

^CO, ^(2,)

respective Arrhenius equations being k = 1012*95exp- (-67,5W/RT) sec.-l, and k = 1011*90exp(-62,000/RT) sec.-l. Studies on formic and propionic acids in static systems a t around 500" indicate t h a t homogeneous dehydration is a second-order reaction in both cases, with activation energies about half of t h a t shown b y acetic acid. Decarboxylation is first-order in each case. The homogeneous decomposition of acetic acid has therefore been investigated in detail in a static system; dehydration kinetics closely resemble those of its homologues, whereas decarboxylation is similar in flow and static apparatus.

EXPERIMENTAL

Apparatus and Procedure.-The static apparatus used has been described previously.3 Unpacked and packed silica vessels had surface-to-volume ratios of 0.88 and 9-69 cm.-1 respectively. Vessels were prepared for use by pyrolysjng acetic acid or isobutene in them a t 600" until reproducible rates of acid decomposition were obtained.

Rates measured in different prepared vessels were identical under the same conditions.

Decomposition did not normally exceed 5y0, so complic- ations due t o further reactions of keten were minimised, and rate constants were obtained from initial rates. Unreacted acid and products were separated into two fractions by C. H. Bamford and M. J. S. Dewar, J . C h e w . SOC., 1949, 2877.

P. G. Blake and Sir C. Hinshelwood, Proc. Roy. SOC., 1960, A , 255, 444.

distillation a t -78". Of the volatile fraction, CO, CO,, CH,, C,H,, and C,H, were measured by gas-solid chromato- graphy (g.s.c.) on activated silica gel a t 25", propadiene on silica gel a t 120°, and keten by i.r. absorption a t 2150 cm.-l.

The involatile fraction contained water and unchanged acetic acid; the former was measured by gas-liquid chro- matography, and the acid by titration with alkali. Un- successful attempts to detect acetic anhydride employed

U . V . absorption a t 253 my and i.r. analysis.

Materials.-Glacial acetic acid (moisture content 0.01- 0.1 yo) was fractionally distilled under dry conditions and stored in a desiccator. Gas-chromatographic analysis using large injections of acid demonstrated that the water content did not exceed 0.02yo. The acid was degassed before use by repeatedly freezing, pumping, and thawing. Keten was prepared by pyrolysing acetic anhydride in a flow system a t 400°, and purified by low-temperature distillation. Other gases used were generally of C.P. grade and obtained from cylinders.

RESULTS AND DISCUSSION

Acetic acid has been decomposed in the temperature range 460-595" at pressures ranging from 10 to 730 mm.

The major products are keten, water, carbon dioxide, and methane, and a typical product-time curve is shown in Figure 1. Very much smaller amounts of carbon monoxide, propadiene, and ethylene, and a trace of ethane, are also formed. The minor products are secondary in nature, the ratio of CO to CH,, for example, falling t o zero a t the beginning of the reaction, and presumably stem from slight decomposition of keten according to reactions (3) and (4). Pressure-time plots

2CH2C0

--+

2CO

+

^C,H, (3) 2CH,CO CH,:C:CH,

+

^CO, ⁽⁴⁾

P. G. Blake and I<. J . Hole, J . Chem. SOC. ( B ) , 1966, 577.

J. R . Young, J . Chem. SOC., 1958, 2909; W. B. Guenther and W. D. Walters, J . A m e r . C?aern. SOC., 1959, 81, 1310.

Published on 01 January 1968. Downloaded by University of Western Ontario on 31/10/2014 06:33:38.

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1154

for CO, curve slightly upwards owing to reaction (4), but methane is a negligibly minor product of keten. Initial rates of formation of CH, and CO, are substantially equal, the latter being slightly the greater in some cases.

Since CO, and CH, together frequently constitute less than 5% of the total products, the difference is not thought to be significant. Water and keten pressures match fairly well but reproducibility in measuring the former is relatively poor.

The ratio A$/Aacid is 0.95 & 0.05, giving a stoicheio- metry of 1-95 & 0.05. Aacid is calculated from the differences in the titrations of acid in the vessel removed a t time equals zero and a t the end of the given run. The width of the limits is due to Aacid being a small difference between large quantities in the early stages of decomposition. In fact, Ap/Aacid remains constant up

15

r

10

-

E

H,O and

s

0 ) L

CO, and CH,

%+---%--%

0 3 0 6 0

Time (sec.)

FIGURE 1 Major product-time curves for pyrolysis of 200 mm. acetic acid at 585"

keten; 0 H,O; x CO, and CH,

to 20% decomposition despite appreciable secondary reaction, since decomposition and polymerisation of keten occur together without appreciable pressure change around 500-550". Well over 90% of the acid decomposed was normally accounted for by the products.

The results show that the overall decomposition reactions of acetic acid are correctly represented by (1) and (2).

Kinetics.-Representative rate constants for acid dehydration and decarboxylation are in Table 1. Similar

TABLE 1

Rate constants for dehydration and decarboxylation of acetic acid

Temp.

585.0"

565.0 535.0 535.0 535-0 535-0 493.5 460.0 ("c)

$acid

(mm.1 200 200 200 50 200 400 200 400

Vessel Unpacked Packed Unpacked Unpacked Packed Unpacked Unpacked Unpacked

10k (H,O) (1. mole-l sec.-l)

3.52 2.42 0.99 1.00 1.06 1-05 0-342 0.124

1 0 5 ~ ( ~ 0 , ) (sec.-l)

7.33 6.50 1-42 1.50 1.97 1-21 0.13 0.033

J. Chem.

^{S O C .}

(B), 1968

results are obtained in packed and unpacked vessels, except that rates are slightly higher in the former a t low pressures, owing principally to an increase in the rate of decarboxylation. Even this is relatively small (50% for an 1100% increase in surface-to-volume ratio), and thus both reactions are substantially unaffected by changes in this ratio.

Decomposition to keten and water is second-order with respect to acid pressure. Activation energies were measured at nominal pressures of 50, 200, and 400 mm., and, corrected to constant concentration, are 35.7, 33.4, and 33.6 kcal. molep1 respectively, with a mean standard deviation of 1.1 kcal. The mean Arrhenius equation is k = 108.45e~p( -34,20O/RT) 1. molep1 sec.-l. Decarb- oxylation follows first-order kinetics with activation energies of 60.1, 58.6, and 56.8 & 2.1 kcal. mole-l at 50,200, and 400 mm. pressure, and a mean Arrhenius equation k = 1011*12exp( -58,50O/RT) sec.-l. Here, k increases as the pressure falls to the lowest values measured, indicating again a slight surface contribution.

Rates were measured in the presence of an excess of propene (Table 2) in order to test for free-radical chains.

TABLE 2

Effect of inhibitor and added gases on acetic acid decomposition a t 493.5"

Added gas Propene ...

Propene ...

Benzene ...

C,F1, ...

CH, ...

C,H, ...

co

² ...

P (mm.1 115

0 210 0 225 0 405 408 407 167 615

jbacid (mm.) 105k(acid) (sec.-1)

53 5.7

53 5.2

185 12.5

186 13-5

114 8.2

114 8.9

55 5.1

52 5.3

52 5.9

51 6.1

51 5.7

Propene affects neither the overall rate nor the individual processes. The amount of CH, found is, of course, in- creased, since propene itself is decomposing slowly. The rate of this latter decomposition is not affected by the acetic acid. Both dehydration and decarboxylation thus appear t o be molecular processes a t this temperature, and the absence of chains, and of surface depend- ence, indicate that they are also homogeneous. At pressures above 400 mm. when the rate is largely due to dehydration, the overall order falls with increasing pressure. This effect is shown in Figure 2, where the overall first-order rate ' constant ' is plotted against pressure, at some of the temperatures studied. The form of the curves, and the effect of temperature on them, suggested that dehydration might be a unimolecular process, being studied mainly in the second-order region. Accordingly, the effect of various gases on the rate was studied (Table 2). Either there was no effect or a very slight increase, and, unless the efficiencies of added gases in transferring energy are improbably

A. F. Trotman-Dickenson, ' Gas Kinetics,' Butterworths, London, 1955, p. 84.

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Phys. Org. 1155

i t seems unlikely t h a t the reaction is unimolecular. A pre-exponential factor of 106.01 sec.-l is also abnormally low. Dehydration of formic acid departs only slightly from second-order over the same pressure range, and that of propionic acid remains second-order up to 200 mm., which was the highest pressure studied. The Arrhenius parameters are similar in the three acids :

k = 107'46exp( -28,50O/RT) 1. mole-l sec.-l (formic) k = 105*45exp( -34,20O/RT) 1. mole-l sec.-l (acetic) k = 10**76exp( -35,15O/RT) 1. molep1 sec.-l (propionic) All three are homogeneous molecular processes, and the results suggest t h a t a similar mechanism operates in

301-

n y v 2 0

0 )

e9

t

S'@

S8Sa

/"

0

0 2 0 0 400 6 0 0 Acid pressure (mm.)

' constants ' for acid decomposition FIGURE 2 Effect of acid pressure on first-order rate

0 Unpacked vessel; 0 packed vessel

each case. This might involve the formation of a six-centre transition state from two acid molecules, fol- lowed b y elimination of water and the rapid decomposition of anhydride a t these temperatures6 Acetic anhydride could not be detected among the products,

&'I. Szwarc and J. Murawski, Trans. Faraday Soc., 1951, 47, W. C . Child and A. J. Hay, J . A m e r . Chem. SOC., 1964, 86, 269.

182.

but the equilibrium concentration at 530" and 200 mm.

initial acid pressure is only 0.3 mm.' Alternative six- centre transition states from two acid molecules are

r

^Me

1

l o

4 c l

^*\ o

]

^CH2CO ^3. ^MeC02H

L H

possible, which might decompose directly to keten, water, and acid. However, a t 34 kcal. mole-l, the activation energy is only 4 kcal. greater than the heat of reaction, making the direct process seem less likely. Activation energies for dehydration of the three acids fall with increasing electron-attracting power of the ' alkyl ' group (Taft ^C* C,H, -0.10; CH, 0.00; H +0.49). The trend is consistent with the weakening of the O-H bond in the six-centre transition state.

It seems that a completely different mechanism pre- dominates a t the much higher temperatures of the flow work, since dehydration has an activation energy of 67 kcal. mole-l and is first-order at all pressures. The bimolecular process has a relatively low A-factor and activation energy, and therefore gives way a t high temperatures to the unimolecular decomposition, with its high activation energy and normal A-factor. The marked increase in the Arrhenius constants is consistent with a change of mechanism from six-centre bimolecular t o four-centre unimolecular transition states. Acetic acid resembles acetamide, which appears to undergo unimolecular decomposition in a flow systems at 700", when k ⁼1015~70exp(-73,400/RT) sec.-l, but t o be second-order a t 500" in static a p p a r a t ~ s , ~ k = 10s*59exp- (-36,20O/RT) 1. mole-l sec.-l.

The kinetics of decomposition to methane and carbon dioxide under flow and static conditions are very similar, a difference in activation energy of 3.5 kcal. molep1 being very small when the widely different temperature ranges are taken into account.

G. E. J . thanks I.C.I. Fibres Limited for a maintenance

[S/752 Received, M a y 29th, 19681

* ^M.Hunt, J. A. Kerr, and A. F. Trotman-Dickenson, J . Chem. SOC., 1965, 5074.

J. Aspen, A. Maccoll, and R. A. Ross, Trans. Faraday Soc., 1968, 64, 965.

grant.