KINETICS OF
THE
DECOMPOSITION OF OXALICACIDIN GLYCERINE
1729
Kinetics of the Decomposition of Oxalic Acid in Glycerine
by M. A. Haleem and Peter E. Yankwich Noyes Laboratory of Chemistry, University of Illinois, Urbana, Illinois (Received January 4, 1966)
The decomposition of oxalic acid in 96% glycerine solution has been studied over the range 88-121’ in a flow system, The nonstoichiometry observed by Clark at somewhat higher temperatures (and in a closed system) was not found, but it was discovered that the simple first-order character of the reaction altered suddenly at a degree of reaction which increased with increasing temperature. Formic acid was found to accelerate the decomposition; the effect of small amounts of added water was found to be very small, a difference from the behavior of the molten dihydrate. For the first-order region, the observed Arrhenius parameters are E = 28.4 f 0.2 kcal. mole-l and log ( A , sec.-l) = 12.3 0.1; these and other activation parameters are in good agreement with those obtained for solvents expected to be kinetically inert toward oxalic acid and from the gas phase reaction.
*
Introduction Clark1 has published the results of an investigation of the kinetics of the decomposition of oxalic acid in 95% glycerine between 124 and 159’; activation parameters were calculated, apparently for the temperature range 141-159’. In that study, the author observed, “the decomposition of oxalic acid in glycerol is not stoichiometric a t temperatures below 150’,” the apparent yield of carbon dioxide product being 61% a t 123.6’, 83y0 a t 139.5’, 95q7b at 150.1’, etc. Further, he suggested that competition with decomposition of a slower reaction such as ester formation between the solute and solvent might be responsible for the effect. During work preparatory to a study of the carbon isotope fractionation effects in the oxalic acid decomposition in glycerine, we undertook to use Clark’s results to guide procedure development; our early runs were a t 121’, just below a temperature a t which he had observed about 60% yield of carbon dioxide. Since the previous experiments were carried out in a closed system and ours (vide infra) employed a flow apparatus, it seemed possible that the differing observations could be due to a carbon dioxide solubility effect in his experiments. However, the apparent thermal effect associated with the low yields observed by Clark is about -25 kcal. mole-l, while the solubility data2s3 correspond to a heat of solution almost twice as large. The study reported here was undertaken primarily
to secure kinetics information useful to the planning of later isotope effects work, but it became important to check possible “method dependence” of the rate results.
Experimental Reagents. Fisher Certified analytical reagent materials were used. Anhydrous oxalic acid was heated a t 110’ for 1 hr., then stored in vacuo over magnesium perchlorate. The 96% glycerine used as solvent and a small sample of 96% formic acid were used as received. Where required, deionized water was used. Apparatus. The kinetics apparatus was a flow system operating on dry, tank nitrogen gas a t an inlet pressure of ca. 60 em. maintained by house vacuum; the nitrogen flow rate was 50-75 ~ mmin.-l, . ~ and the volume to be swept approximately 500 ~ m . ~The . nitrogen stream bubbled through a glycerine solution of oxalic acid contained in a reaction vessel thermostated to 10.05’, thence through a condenser cooled by water a t 15’, and finally through a series of traps a t -78 and - 195’ which were part of a high-vacuum gas-handling apparatus of conventional design. Procedure. Glycerine (50 ml.) was placed in the reactor, and 90.00 mg. of oxalic acid was weighed into a small glass dish which was set onto the end of a ~
~~
(1) L.W.Clark, J . Am. Chem. SOC.,77, 6191 (1955). (2) G.Just, 2. physik. Chem., 37, 342 (1901). (3) A. yon Hammel, ibid., 90, 121 (1915).
Volume 69, Number 6
M a y 1966
M. A. HALEEM AND PETER E. YANKWICH
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movable probe inside the reaction vessel. After assembly of the system and placement of the reactor in the thermostat, nitrogen flow was started, and 30 min. was allowed for thermal equilibrium to be established. The reaction was started by tipping the oxalic acid into the solvent (solution being assisted by agitation of the reactor) and was quenched by addition of 50 ml. of ice water through the condenser tube. The reaction vessel was then lifted out of the thermostat, nitrogen flow increased to ea. 150 min.-', and sweeping continued for 20 min. Carbon dioxide product in the traps was puritied carefully by distillation, then estimated manometrically. Each kinetics run thus yielded just one a (degree of decomposition)-t (time) data pair; a plot of -In (1 - a) vs. t was used to guide the collection of data at each temperature.
Results Between 10 and 23 decomposition runs were carried out at each of four temperatures: 121.0, 112.0, 100.0) and 88.0'. The last data pair at 121.0' corresponded to a = 0.827; at no point at this temperature was there significant deviation from first-order dependence of the rate on oxalic acid concentration. At 112.0' similar observations were made up to ea. a = 0.70, at which point the rate of decomposition increased. This same anomalous acceleration was found at ea. Q! = 0.40 at 100.0' and at ea. a = 0.15 at 88.0'. The values of the apparent first-order rate constant, kl, and other information concerning the welkbehaved region af a at each temperature are collected in Table I ; the kl were obtained from least-squares analysis of all data pairs in that region. The various activation parameters are shown in Table 11.
Table I : Rate of Decomposition of Oxalic Acid in 96% Glycerine
O C .
ac (end of wellbehaved region)
No. of data pairs
Av. ki (sec.-1) X 106
88.0 100.0 112.0 121.0
0.15 0.40 0.70 (>0.83)
10 17 9 10
1.33 f 0.03 4.62 f 0.03 15.3 f 0 . 2 36.7 f 0.9
Run temp.,
Several additional experiments were carried out at 100.0' to test the effect on the rate of added water. Clark' found the rate of decomposition of oxalic acid dihydrate to be affected drastically by the addition of equal amounts of water, while Bredig and Lichty4 reported even larger effects when oxalic acid was disThe Journal of Phveieal ChsmiStry
Table II: Calculated Activation Parameters
d Log (A, sec-1) AH*b AS*d
AF**
kl (sec.-l) X 1@, at 140'
Earlier resultsa
This research
27.2 11.8 26.4 -7.8 29.6 251
28.4 f 0.2c 12.3 f 0 . 1 27.6 f 0 . 2 -4.9 f 0 . 4 29.6 f 0 . 3 197
Errors in this column are all a See ref. 1. Kcal. mole-'. standard deviations evaluated by a least-squares procedure. Cal. mole-' OK.-'.
solved in concentrated sulfuric acid. To test for such an effect in 96% glycerine solution, three runs (a = 0.16, 0.20, and 0.24) were carried out with 0.50 (27.7 mmoles) of water added. The value of kl (set.-') X lo5 found in these experiments was 4.58 f 0.08. In three other experiments at 100.0' (a = 0.17, 0.31, and 0.43), 0.040 cm.8 of 96% formic acid (1.02 mmoles) was added to the glycerine solvent. The value of kl (sec. -I) X lo6observed in these experiments was 5.17 f 0.05, definitely higher than that shown in Table I. Unrecorded in Table I are the results of experiments for which the degree of decomposition was greater than the limit (m) of the well-behaved region; at 100.0', six such results were obtained. The following data pairs are a, kl (sec.-l) X 106 for each such run: 0.458, 4.86; 0.501, 5.27; 0.534, 5.31; 0.547, 5.28; 0.563,5.30; and0.599,5.64.
Discussion All of the activation parameters derived from this research (Table 11) are in good agreement with those obtained by Clark from experiments at temperatures where normal yields of carbon dioxide were observed; fortuitously, all of the earlier kinetics work was carried out at temperatures sufficiently high that the apparent acceleration of the reaction was not observed. The small effect of added water which was observed suggests that any large influence of water on the rate in glycerine solution must occur a t higher glycerine concentrations than 96% in consonance with Bredig and Lichty's findings for concentrated sulfuric acid s01vent.~ The acceleration produced at low CY by the addition of 1 mmole of formic acid is similar to that observed beyond a = 0.40 when only product formic acid is present. The accelerative effect of added sulfuric acid on the rate of oxalic acid decomposition in aque(4) G. Bredig and D. M. Lichty, 2. Elektrochm., 12, 469 (1906).
KINETICSOF THE DECOMPOSITION OF OXALIC ACIDIN GLYCERINE
ous solutions was noted by Dinglinger and Scbroer6 and is about twice as large as that of formic acid on glycerine solutions of oxalic acid noted above. It is not likely that the relative acidities of formic and sulfuric acids in glycerine are sufficiently similar that the formic acid effect is simple acid catalysis. Carbon dioxide is rapidly removed from the solution by the nitrogen sweep, but it seems certain that formic acid product is neither rapidly volatilized from the solution (its Duclaux constant from glycerine must be quite low) nor appreciably decomposed to carbon dioxide and hydrogen under the reaction conditions. Interestingly, a plot of In (formic/oxalic)c vs. 1/T is linear, with an associated enthalpy of about 29 kcal. mole-'. These considerations suggest a model mechanism which has mathematical properties in consonance with the few experimental observations available: formic acid is slowly swept from the solution (the rate increasing with temperature), but that remaining is complexed by the solvent6; as the temperature increases, larger concentrations of formic acid (larger CY) would be required to yield a given concentration of complex. When a critical concentration of complex is reached, a second-order reaction between complex
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and oxalic acid contributes to the observed rate of decomposition. The activation parameters in Table I1 are similar to those observed for oxalic acid decomposition in the gas phase' and in kinetically neutral solvents like dioxane and triethyl phosphate and suggest a simple intramolecular mechanism for the firsborder reaction. In further exploration of this decomposition, carbon isotope effect studies and experimentation on the apparent formic acid catalysis are under way in this laboratory.
Acknowledgments. We are indebted to our colleagues Professors D. Y. Curtin and S. G. Smith for illuminating discussion of the formic acid catalysis. This research was supported by the U. S. Atomic Energy Commission; (5) A. Dinglinger and E. Schroer, 2. physik. Chem., A179, 401 (1937). (6) Perhaps esterification actually ocours, aa proposed for oxalio acid by Clark.' The acceleration could not be due to a formic acidoxalic acid complex because the effect of its formation would reach a maximum at a = 0.6, independent of temperature. (7) G.Lapidus, D. Barton, and P. E. Yankwich, J. Phys. Chem.,68, 1863 (1964).
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