Ind. Eng. Chem. Res. 2000, 39, 843-845
843
Acid-Catalyzed 2-Furaldehyde (Furfural) Decomposition Kinetics Ian C. Rose, Norman Epstein,* and A. Paul Watkinson† Department of Chemical and Bio-Resource Engineering, The University of British Columbia, Vancouver, British Columbia, Canada V6T 1Z4
Experiments on acid- (HCl-) catalyzed thermal decomposition of 2-furaldehyde in a dilute aqueous solution were performed in a stainless steel batch reactor over a temperature range of 130-170 °C. First-order reaction kinetics best described the experimental results. The rate of reaction increased with temperature and yielded a kinetic activation energy of 48 kJ/mol, somewhat lower than the value previously reported in the literature. Introduction 2-Furaldehyde is a decomposition product of sugar and is present in most foods in various concentrations. Detection of 2-furaldehyde is therefore an indication of the deterioration of food products as a result of the effects of heat and/or storage. In an acidic (H2SO4), aqueous solution of 2-furaldehyde, Williams and Dunlop1 measured the concentration of 2-furaldehyde over time at several fixed temperatures (150-210 °C). They also made one such measurement using HCl instead of H2SO4 at 160 °C. The following kinetic equation was written for the H2SO4 runs:
d[2-furaldehyde] ) -k′[2-furaldehyde][H+][H2O] dt (1) However, the concentration of acid and water remained essentially constant, and therefore a pseudo-first-order reaction was assumed:
d[2-furaldehyde] ) -k[2-furaldehyde] dt
(2)
The following Arrhenius-type equation correlated the experimental results over the temperature range 150210 °C, 7
k ) (1.396 × 10 )e
-83.6/RT
(3)
where the rate constant k is in min-1, the universal gas constant R in kJ/mol‚K, and the absolute temperature T in K. The present study was undertaken to evaluate the kinetic rate constant for a series of temperatures and thereby the activation energy, using HCl instead of H2SO4 as the acidic catalyst. Instead of sealed tubes submerged in heating oil and the Hughes-Acree2 method of analysis used by Williams and Dunlop,1 the present study uses a stirred cell reactor and ultraviolet spectrophotometric analysis. Experimental Procedure and Analysis The isothermal decomposition of 1 wt % 2-furaldehyde in 0.1 N HCl was studied using the batch reactor shown * To whom correspondence should be addressed. † E-mail:
[email protected]. Phone: (604) 822-3238. Fax: (604) 822-6003.
Figure 1. Schematic diagram of the stirred cell reactor apparatus.3
in Figure 1. A 2.3-L vessel constructed from 316 stainless steel consisted of a stirrer, a K-type thermocouple, and a pressure relief valve. A 3.2-mm i.d. sample line was used with a needle valve to remove samples for analysis. Samples were collected in 20-mL disposable vials, placed immediately into an ice bath, and stored in a dark cupboard until analysis. Prior to an experiment, this sample line was used to purge nitrogen through the reactor for 5 min to eliminate air from the system and therefore minimize autoxidation.4 The reactor was mounted into a stabilized mineral oil bath which contained two 500-W immersion heaters capable of heating the bath to approximately 200 °C and a stirrer. An experiment was initiated by heating the oil bath to 10-20 °C above the desired temperature. This process would take approximately 2 h, during which time the reactor was charged with the test fluid (1.4 L) and purged with nitrogen. The reactor was then pressurized with nitrogen to above the vapor pressure of the solution and immersed into the bath, marking time zero
10.1021/ie990550+ CCC: $19.00 © 2000 American Chemical Society Published on Web 02/04/2000
844
Ind. Eng. Chem. Res., Vol. 39, No. 3, 2000
Figure 2. Kinetic description of 2-furaldehyde decomposition at T ) 169 °C.
Figure 3. Temperature dependence of acid-catalyzed decomposition of 2-furaldehyde.
for the experiment, and the temperature and pressure were recorded. The experiment was terminated once particulate matter was clearly visible in the sample vial. At this point the samples were stored in an ice bath overnight in a dark cupboard for analysis. A decrease in 2-furaldehyde concentration over time was measured using a UV spectrophotometer at 270 nm, the measured wavelength of maximum absorbance for 2-furaldehyde. Experimental samples of 1 wt % 2-furaldehyde were diluted 500 times to produce a maximum possible sample concentration of 20 ppm. The following standards were prepared to develop a calibration curve: 20, 15, 10, 5, and 1 ppm. A linear least-squares
regression gave rise to the following calibration:
A ) 0.156 × [2-furaldehyde(ppm)] + 0.094 r2 ) 0.9999 where A is the absorbance. Results and Discussion Figure 2 shows the decrease in 2-furaldehyde concentration with time for experiment 6 at 169 °C as ln(C0/C) versus t, that is, in terms of first-order reaction kinetics. Correlation of C versus t according to zero-
Ind. Eng. Chem. Res., Vol. 39, No. 3, 2000 845 Table 1. Summary of Stirred Cell Reactor Experiments standard exp. C0 Ta deviation duration no. (g/100 mL) (°C) (°C) (min)
analysis
k (min-1)
4 5 6 lit.
UV UV UV Hughesb
1.45 × 10-3 7.79 × 10-4 2.78 × 10-3 2.16 × 10-3
0.985 1.012 0.951 1.000
150 130 169 160
6.5 2.2 2.3
290 300 180 210
a
Average temperature over the duration of the experiment. b Hughes-Acree2 method of analysis.
order kinetics or of C-1 versus t according to secondorder kinetics always yielded values of the correlation index r2 smaller than that for first-order. In the regression of the experimental data, the first three data points were ignored because these samples were taken during the initial heat-up period of the experiment. Table 1 shows the kinetic results from analysis using the UV spectrophotometer. Also included for comparison is a value of the pseudo-first-order rate constant from the one experiment in the cited study1 in which HCl was used as the catalyst. Figure 3 shows a semilog plot of the kinetic rate constant versus reciprocal temperature, k ) k0 exp(∆E/RT). These results apply to temperatures of 130170 °C. This figure also shows a comparison between the experimental and literature results. The one data point from Williams and Dunlop1 using HCl fits the data obtained in the present study. For the experimental conditions described above, the kinetic activation energy ∆E ) 48.1 kJ/mol, based only on the three runs of the present study (compared to 83.6 kJ/mol obtained by Williams and Dunlop1 using H2SO4) and the preexponential factor k0 ) 1297 min-1 (versus 1.396 × 107 min-1). It is interesting to note that the activation energy reported here is nearly half the value reported in the previous study,1 despite the fact that the experiments were performed using the same normality of HCl and H2SO4. It appears that although the same nominal
strength of acid has been used in the two studies, dissociation of H2SO4 is not driven to completion, as is the case with HCl, thereby giving rise to a higher activation energy. Acknowledgment Continuing financial support from the Natural Sciences and Engineering Research Council of Canada is gratefully acknowledged. Nomenclature A ) UV spectrophotometer absorbance at 270 nm C ) concentration of 2-furaldehyde, g/100 mL of solution C0 ) original concentration of 2-furaldehyde, g/100 mL of solution k ) pseudo-first-order rate constant, min-1 k′ ) kinetic rate constant, min-1 k0 ) Arrhenius pre-exponential factor, min-1 R ) universal gas constant, kJ/mol‚K t ) time, min T ) temperature, K ∆E ) activation energy, kJ/mol
Literature Cited (1) Williams, D.; Dunlop, A. Kinetics of Furfural Destruction in Acidic Aqueous Media. Ind. Eng. Chem. 1948, 40, 239-241. (2) Hughes, E. E.; Acree, S. F. Volumetric Estimation of 5-Bromo-2-Furoic Acid with Standard Bromate. Ind. Eng. Chem., Anal. Ed. 1934, 6, 292-293. (3) Rose, I. C. Model Investigation of Initial Fouling Rates of Protein Solutions in Heat Transfer Equipment. Ph.D. Thesis, The University of British Columbia, Vancouver, BC, Canada, April 1999. (4) Dunlop, A.; Peters, F. The Furans. ACS Monographs: Reinhold Publishing Corp.: New York, 1953; Chapters 8 and 9.
Received for review July 26, 1999 Revised manuscript received November 23, 1999 Accepted January 6, 2000 IE990550+
