232
Ind. Eng. Chem. Prod. Res. Dev. 1980, 19, 232-237
Kinetics of Acetylene Hydrochlorination H. Subbaraman Shankar" and John B. Agnew Department of Chemical Engineering, Monash University, Clayton, Victoria, 3 168, Australia
The results of an experimental study of the kinetics of acetylene hydrochlorination over mercuric chloride catalyst on activated carbon supports are presented. Measurements were made over the temperature range 100-240 O C . For temperatures less than 140 O C good agreement was obtained with the results of previous workers; a Langmuir adsorption model provided a good description of the data, the activation energy for the reaction step being in the range 10-1 1 kcal/mol. An Arrhenius plot of overall reaction rate between 75 and 240 O C showed an abrupt change of slope at 135 OC, the overall activation energy changing from 9.5 to 2.4 kcal/mol. This discontinuity could not be ascribed to mass transport effects alone, but appeared to be due to a change in reaction mechanism.
Introduction Vinyl chloride was first produced commercially by catalytic hydrochlorination of acetylene. As ethylene became plentiful in the early 1950's it quickly displaced acetylene as the major feedstock, so that today direct chlorination and oxychlorination of ethylene are the principal processes employed commercially. Mixed-gas and acetylene processes are still used in a number of countries around the world, however. McPherson et al. (19791, in a detailed survey of the vinyl chloride industry, mention the recent development of a new crude oil cracking process which produces a substantial yield of acetylene along with ethylene. This process, they suggest, could be advantageous under certain economic and geographic conditions in providing a mixed feedstock for vinyl chloride production. This paper is concerned with the kinetics of the hydrochlorination of acetylene over a catalyst consisting of mercuric chloride supported on activated carbon. Previous studies of this reaction have been reported by Wesselhoft et al. (1959), Gel'bshtein et al. (1963, 19721, and Weston and Agnew (1973) (see Table I). However, most of the data were limited to temperatures below 140 OC. Above this temperature, deactivation is known to occur, the rate of deactivation increasing markedly with temperature (Shankar, 1976). As commercial reactors are known to operate with hot spots between 140 and 200 "C+, a detailed study of the reaction which includes this highertemperature region was carried out and is now reported. Experimental Section Equipment. A continuous stirred-tank catalytic reactor (CSTCR) was used for experimental rate determinations. This consisted of a cylindrical stainless steel reactor vessel 7.4 cm i.d. and 18 cm long flanged at both ends. A Magnedrive stirrer was fitted to the top lid. Four wire-mesh catalyst baskets (11 X 1.8 x' 0.7 cm) were fitted to the stirrer shaft (9.5 mm diameter) and propellers were attached above and below the baskets to improve gas mixing. The bottom of the reactor vessel was fitted with a bursting disk of nickel-plated mild steel (rupture pressure 550 kPa). Separate heating elements were wound on the reactor body and top and bottom flanges, with power inputs regulated individually. The temperature inside the reactor could be measured at a number of points by thermocouples
* Address correspondence to this author a t the Department of Chemical Engineering, Indian Institute of Technology, Bombay 400076, India. 0196-4321/80/1219-0232$01 .OO/O
inserted in thermowells fitted to the lid. The whole assembly was insulated. Materials. Hydrogen chloride of 99.9% purity was obtained from Matheson Gases Inc. Commercial grade acetylene was used; this was treated in a purification train containing water, silica gel, and activated carbon to remove impurities before entering the reactor. Vinyl chloride of 99.9% purity was obtained from B. F. Goodrich (Australia) Ltd. for calibration purposes. Tsurumicoal activated carbon was used as a catalyst support; this was obtained from Kogyo Ltd., Japan, through I.C.I. Australia Ltd. Three different sizes were used, viz. 4 X 4 mm cylinders (TCI), 2 X 2 mm cylinders (TC2), and 0.25-0.42 mm granules (TC3). Catalyst Preparation. The catalyst was prepared by first dissolving the required quantity of mercuric chloride in distilled water. The required amount of carbon support was then added and stirring was continued for 24 h. The distribution of mercuric chloride between activated carbon and water is such that almost all of the salt adsorbs on carbon. The carbon containing mercuric chloride was removed by filtration and dried at 105 "C. The filtrate was analyzed for mercury using an atomic absorption spectrophotometer to determine, by difference, the quantity of mercuric chloride adsorbed on carbon. Catalyst was prepared separately for each run. Gas Analysis. An F & M Model 700 gas chromatograph fitted with a flame-ionization detector was used; 20% w/w oxidipropionitrile supported on 80-120 mesh celite packed in 120 cm long, 6.3 mm stainless tubing was used to achieve the desired separation. Mixtures of acetylene and vinyl chloride required for calibration purposes were prepared by mercury displacement. Hydrogen chloride was removed from the product gas stream prior to analysis in an absorber containing potassium hydroxide pellets. Procedure. Reaction runs were carried out by operating the reactor at temperatures between 100 and 240 OC, pressure of 1 atm, and HCl inlet mole fraction in the HC1-C2H2 feed in the range 0.10 to 0.74. Concentration of mercuric chloride on the activated carbon supports was varied between 4.0 and 12.0% w/w, with 10% being used for the majority of runs. Details are given in Table 11. The mixing characteristics of the reactor had been determined previously by Weston (1970) using a standard tracer technique. Additional tests were carried out by Shankar (1976) at lower gas flow rates to confirm good mixing above 1200 rpm impeller speed. A speed of 2500 rpm was used for all reaction rate determinations apart from a few check runs. 0 1980 American Chemical Society
Ind. Eng. Chem. Prod. Res. Dev., Vol. 19, No. 2, 1980
233
Table 1. Experimentad Kinetics Studies of Acetylene Hydrochlorination (Mercuric Chloride/Activated Carbon Catalyst)
--
author Wesselhoft ( 1 9 8 9 )
reactor
T , "C
p , atm
diff'l
75-125
1-4
reaction rate, mol h-' (g of HgCl,).'
model no.
~ ~ K H K APPAH
IA
(1+ KHPH + KVpV)(1 + KAPA) Wessellioft (1959)
diff'l
75-125
1-4
IB
krKHKApHpA
(1 + KHPH) (1 + K A P A )
Shankar ( 1 9 7 6 )
CSTCR
100- 240
1-1.7
Shankar ( 1 9 7 6 )
CSTCR
100-240
1-1.7
~ J PH PH A (1+ KHPH + KAPA)
IIA IIB
krKHPHPA
(1 + KHPH) recycle
Gel'hshtein (19163) Weston (1973) Weston (197 3)
CSTCR CSTCR
75-140 75-140
I11
krKApHpA
1
100-180
(1 + K A P A + KHPH) k r P H h P A a P V V exP ( - . E / R T ) k r P H h P A a exp ( - E / R T )
1.7-2.4 1.7-2.4
IV V
Table 11. Operating Conditions for Reaction Rate Determinations run no. 10 15 16 17 19 20 21 22 25 27
support TC 1 TC1 TC1 TC2 TC1 TC1 TC1 TC 1 TC3 TC3
g of HgC1, 1.79 1.70 1.81 1.19 1.24 1.00 0.65 1.25 1.00 0.85
28 29 30 31 32 35 36
TC3 TC1 TC2 TC2 TC2 TC2 TC2
0.67 1.02 1.02 1.02 1.02 1.0 1.0
conditions
T , "C 100,200 100 100 140, 240 100, 140 100,210 100,210 100
% w/w
10 8.5 12 10 6 10 4 10 10 10
{;E;:
10 10 10 10 10 10 10
a = exp(-kdt)
(1)
where eXp(-Ed/Rr)
P , atm 1.0 1,o 1.0 1.0 1.0 1.0 1.0 1.0 1.0, 1.68 1.0, 1 . 6 8 1.0 1.0 1.0 1.0
;;;
140, 240 100, 180 1 0 0 , 240 100,210 140,240 100, 180 100,210
The time taken tO bring the reactor up to operating temperature from room temperature was about 2 h in order to ensure uniformity of heating. For operating temperatures above 150 "C this procedure resulted in some loss of activity which had to be accounted for in expressing reaction rates in terms of fresh catalyst. A deactivation study was carried out (Shankar, 1976) to determine appropriate correction factors, and this showed that deactivation could be described by the simple experimental decay relationship
k d := k d o
_______
catalyst
-_
(2)
Using determined d u e s of kdO and Ed and the measured temperature-time record for heat-up it was possible to correct the measured rates to apply to fresh catalyst. Results Detailed rate data are given by Shankar (1976). Temperatures below 140 "C. Figure 1 is a plot of p A / r against l / p Hat a temperature of 100 "C, for different mercuric chloride concentrations and supported sizes. Reaction rates expressed per unit mass of mercuric chloride are seen to be essentially independent of mercuric chloride loading on carbon up to 1 2 % w/w. Rates were also independent of carbon granule size. The results fall reasonably within the range predicted by earlier studies. The data at 140 "C showed similar agreement when plotted in this manner.
1.0 1.0 1.0 1.0
T
.
XHI
0.10 -0.7 4 0.1 2-0.60 0.10-0.37 0.12-0.28 0.12-0.30 0.12-0.17 0.12-0.39 0.1 2-0.17 0.1 2-0.39 0.12-0.39 0.1 2-0.39 0.12-0.39 0.1 2-0.39 0.12-0.39 0.12-0.39 0.12-0.39 0.12-0.39 0.12-0.39
iO0.C
12
WESTON (19701
1
'tvv: Y 2 9 .
234
Ind. Eng. Chern. Prod. Res. Dev., Vol. 19, No. 2 , 1980
Table 111. Percentage Deviation of the Present Data from Predictions of Previous Rate Equations Weston ( 1 9 7 0 )
T, "C 100 140
Wesselhoft ( 1 9 5 9 )
___
u2, %
u 3 ,%
u2,%
u 3 ,%
7 4
7.6 12
-18
21
Table IV. Results of Regression Analysis for the Present Data, Using Model IA at 1 0 0 ° C
k, KH KA KV 0 1
3: u4
1 . 7 5 i 0.2 4.3 t 0.4 0.20 i 0.03 4 . 8 t 0.4 93 9 .o 1.c 0.9
at 1 4 0 " C 7 . 6 ?- 0 . 4 3.4 t 0.3 0.18 i 0.03 3.2 i 0.15 94 10.0 1.8 0.9
a procedure. A statistical comparison is given in Table 111. In the present study the suitability of several models shown in Table I alone is examined. This is because there are ample data which fit earlier models quite well and the present measurements were not designed to postulate new models because suitable means of simultaneous and independent determination of adsorption parameters were not available. The main concern was to seek a simple model which describes the present data and earlier data sets reasonably. As models shown in Table I are nonlinear, parameter estimation posed some problems. To circumvent this, a linearized estimation using appropriate transformation was first used to obtain an initial estimate for a subsequent nonlinear search. In cases where the linearized estimation led to unrealistic parameter values, the power transformation weighting method was used to obtain acceptable initial estimates (Box and Hill, 1974). The present data and those of the earlier data sets were each examined against each of the models shown in Table I and the parameter estimates and the fit were examined statistically (Draper and Smith, 1967). All models provided a reasonable description of the data sets (Shankar, 1976), but Langmuir-Hinshelwood-type models IA and IB were distinctly superior because of the consistency of the parameter estimates and the overall fit with respect to each of the data sets. The results of regression for model IA are shown in Tables IV and V. An activation energy of 10-11 kcal/mol is indicated by all the data sets for the rate constant K,. The values of KH, KA, and Kv also agree reasonably well. The value of KH is much larger than KA, indicating that HC1 is more strongly adsorbed than acetylene. Quantitatively, therefore, there is good agreement despite the different total pressures at which the data were obtained. In addition to the correction for deactivation occurring during heat-up, correction was also applied for intraparticle diffusional resistance by calculating effectiveness factors using the method of Weisz and Prater (1954). Values are given in Table VI. Temperature above 140 "C. Corrected rate data were plotted against reciprocal temperature (semilog coordinates) to determine the overall activation energy, as shown in Figure 2. Also plotted are values obtained from Wesselhoft's and Weston's studies. An abrupt change in slope is observed at about 140 "C. Up to this temperature there is good agreement between the three studies, with an activation energy of around 9.5 kcal/mol indicated. The higher value of 11.2 kcal/mol shown in Table V was obtained when only the rate con-
Ind. Eng. Chem. Prod. Res. Dev., Vol. 19, No. 2, 1980 235
Table VI. Estimated Values of Effectiveness Factor for Different Suuuortsi effectiveness factor temp,"C 180 210
240 240
l.OOL
TC1 0.73-0.82 0.65-0.7:? 0.45-0.513
' ,.
L
2iO
. 1,
TC2
TC3 0.99-1.0 0.99-1.0 0.985-0.99
0.92-0.95 0.48-0.55 0.48-0.58 TE HPE RAT URE ' C I40 4 00
180
75
/ I
'=,:
\
E: 9 5 %
A
WESTON
11970)
WESSELHOFT
E T AL I 1 9 5 9 1
> Y
\
L 0.0 1 1.9
1 -I
2.1
23
2.3 L'/T)
IO'
,
2.7 K
2.9
3.0
-'
Figure 2. Arrhenius plot.
stant k, was considered. The lower value incorporates the opposing effects of the temperature dependences of k,, kH, K A ,and Kv. The indicated activation energy for the high-temperature branch (7' > 140 "C) is 2.4 kcal/mol, which is in reasonable agreemenit with Weston's value of 2.0 kcal/mol based on limited data. Discussion Interphase heat and mass transport resistances, and intraparticle heat transfer resistance, were checked using established criteria and were found to be negligible (Shankar, 1976). Besides, the high-temperature branch of Figure 2 is seen to lie well below the line of slope Eo/2 for complete diffusicnal control. Hence the break in the Arrhenius plot cannot be due to simple diffusional effects alone. Similar discontinuities have been observed in a number of other systems despite independent evidence for the apparent absence of transport effects (Shankar, 1976). Two possible explanations can be offered for the discontinuity observed here. (a) Catalyst-Phase Distribution. The extent of internal diffusional effects in the case of supported catalysts is likely to depend on the catalyst particle size and on the distribution of the catalyst phase. In the case of supported metal catalysts, these factors can be determined by selective chemisorption and electron microscopic methods. In the present case, however, adsorbates specific to mercuric chloride could not be found. BET nitrogen adsorption measurements indicated that the surface areas were essentially the same ((920m2/g) for activated carbons and those containing mercuric chloride (Shankar, 1976). IR and X-ray diffraction measurements to determine the chemical nature of mercuric chloride on the surface were inconclusive. However, mercury particles could be detected on a scanning electron microscope; mercury particle size was estimated to be 2500 A by this technique. Considerable loss of resolution resulted at higher magnifications and therefore it was not possible to scan for mercury in small pores. Gel'bshtein et al. (1963) reported unsuc-
cessful results from electron microscopy. It has been shown by Kipling (1965) that adsorption of solutes from solution can be used for estimating catalyst phase areas if a clear and identical plateau region can be discerned in the adsorption isotherm for more than one solvent. Isotherms at 25 "C from solution of HgClz in water and methanol were found to obey the Freundlich isotherm but no plateau region was observed (Shankar, 1976). Mercury in mercuric chloride is essentially covalent and has a covalent radius of 3.5 A (Sanderson, 1966). For monolayer coverage, a 10% w/w HgC12 would occupy approximately 96 m2 of the carbon surface (920 m2/g). The value of 2500 A obtained for mercury particle size from scanning electron microscopy indicates that mercuric chloride is probably present as clumps. Doraiswamy (1978) has pointed out that discontinuities can occur when the catalyst phase is distributed uniformly on the internal surface of supports exhibiting bi-disperse pore structure; activation energy in such situations can fall to Eo/4. In comparing the results of this study with this value it should be remembered that reaction rates plotted in Figure 2 had already been corrected for intraparticle diffusion. Since efforts to determine location, dispersion, and particle size of mercuric chloride were inconclusive, no definite conclusion could be reached for this possible explanation. (b) Change in Reaction Mechanism. Mechanism I. It has been shown by Bond (1972) that Arrhenius plots of observed reaction rates can exhibit discontinuities if the surface coverage of one of the adsorbed species participating in the reaction remains constant over one temperature range and changes over another. The observed behavior in the present case could arise from acetylene surface coverage changing in this manner. Mechanism 11. Gel'bshtein et al. (1963) showed that HC1 is several times more strongly adsorbed than acetylene. The results of the present regression analysis also confirm this. Adsorption is an exothermic process. Therefore, as temperature rises the surface concentrations of adsorbed acetylene and hydrogen chloride decrease. Above 140 "C the acetylene surface coverage becomes so low that reaction no longer proceeds by the Larigmuir mechanism proposed by Wesselhoft et al. (1959) but predominantly by collision between adsorbed HCl and gasphase acetylene with a clearly different activation energy. Since quantitative evidence for the applicability of these mechanisms was scarce, kinetic modelling above 140 "C presented some problems. Weston (1970) fitted an empirical power-law expression for his limited data above 140 "C. In this study, the data sets have been examined against Langmuir adsorption models I and I1 shown in Table I. In fitting the data to models IA and IB it was assumed that mechanism I applies, while for models IIA and IIB, mechanism I1 was assumed. The nonlinear regression analysis indicated that the change of slope occurs at 135 "C, which was thus used for demarcation between data sets from a particular source. The results of regression together with mean and root-mean-square deviations are summarized in Table VII. It can be observed from Table VI1 that the parameter values obtained for the data from different sources show reasonable agreement. To illustrate this, model IA parameter values are shown in Table VIII. It is clear that reasonable agreement exists although k, values for the Gel'bshtein (1972) data are relatively large; the problems associated with this were covered earlier.
236 Ind. Eng. Chem. Prod. Res. Dev., Vol. 19, No. 2, 1980 Table VIII. Comparison of Parameter Values of Model I A for the Different Data Sets
T , 'C
data source
k,, mol K H , h-' g" atm-'
KA,
atm"
Kv,
atm-.' .I__
100
140
present work Weston (1970) Wesselhoft (1959) Gel'bsht ein (1972) present work Weston
Y
T o
0
+I
180
(1970) Gel'bshtein (1972) present work
Weston
1.75 2.0
4.35 4.35
0.195 0.195
4.78 4.0
1.31
2.67
0.46
3.8
5.2
9.66
0.23
-
7.6 8.6
3.4 3.4
0.176 0.176
3.17 4.0
22.6
7.0
0.10
-
25.4 29.0
3.80 2.80
0.077 0.077
2.25 4.0
78.1
5.30
0.037
-
(1970)
Gel'bshtein ( 197 2 )
m
0 2 +I
F?
c90 O
N
0 t I
i
4
e
35 ga
U
&
t i
N"?
0 0 0 +I
+I
+I
i
The value of KAin model IA for present data varies from 0.176 atm-' at 140 "C to 0.03 atm-' at 240 "C. Therefore the term (1 + KA P A ) for experimental conditions becomes virtually equal to unity. As a result, the magnitude of the parameter k, for model IIA becomes approximately equal to kJA in model IA. Therefore the fit of the model I1A will be much the same as model IA. The same argument applies to model IB. Gel'bshtein et al. (1963) quoted a value of 4-5 kcal/mol for the heat of adsorption of HC1 and 2 4 kcal/mol for acetylene over the temperature range 100-140 "C and equilibrium pressures up to 300 mmHg. Although much of the present work was confined to temperatures above 140 "C, the values of heats of adsorption obtained from regression analysis are of the same order as stated above. Both models I and I1 fitted the data of the present investigation to within 15%, after accounting for internal diffusional effects by the method of Weisz and Prater. Conclusions (1) Good agreement exists between reaction rate data from this study and earlier investigations for temperatures up to 180 "C. (2) Kinetic models I and I1 both fit the rate data adequately for temperatures up to 240 "C. (3) The abrupt change in apparent activation energy at 135 " C is not due to internal diffusional limitations alone. (4) The reaction above 135 "C appears to proceed predominantly by collision between adsorbed HCl and gasphase acetylene, while below 135 "C the reaction occurs between adsorbed HC1 and adsorbed acetylene. Acknowledgment H.S.S. wishes to thank Monash University for providing financial assistance to carry out this study. Nomenclature a = activity E = observed activation energy, kcal mol-' Eo = true activation energy, kcal mol-' E,, = activation energy for deactivation, kcal mo1-l El, E 2 , E , = heats of adsorption, kcal mol-' F = critical value of F distribution KH = exp(x2 + E 2 / R T ) ,atm-' K A = exp(x, + E,/RT), atm-I Kv = exp(x, +, E , / R T ) , atm-' kd = deactivation velocity constant, h-l kdo = preexponential factor defined in eq 5 . h-' k , = exp(xl - E,iRT), reaction velocity constant (units dependent on model) P = reactor total pressure, atm
Ind. Eng. Chem. Prod. Res. Dev. 1980, 19, 237-241
pH, P A , p v = partial pressure of HCl, CzH2,and vinyl chloride.
237
Box, G. E. P., Hill, W. J., Technometrics, 16(3), 385 (1974). Doraiswamy, L. K.. N.C.L., Poona, India, personal communication, Jan 1978. Draper, N. R., Smith. H.."Applied Regression Analysis", Wiley, New York, 1967. Gel'bshtein, A. I., Slin'ko, M., Shcheglova, G. G., Yabbnskii. G. S.,Timoshenko, V. I.. Kamenko, 8. L., Kinet. Catal. ( f n g l . Trans/.), 13, 634 (1972). Gel'bshtein, A. I., Siling. M. I., Sergeeva, G. A., Shcheglova, G. G., Khomenko, A. A., Kinet. Catal. (Engl. Trans/.), 4, 123, 262, 543 (1963), Kipling, J. J., "Adsorption from Solution of Non-electrolytes", Academic Press, London, 1965. McPherson, R. W., Starks, C. M., Fryar, G. J., Hydrocarbon Process., 75 (Mar 1979). Sanderson, J. T., "Chemical Periodicity", Reinhold. New York, 1966. Shankar, H. S.,Ph.D. Thesis, Monsah University, Clayton, Australia, 1976. Weisz, P. B., Prater, C. D., Adv. Catal., 6, 143 (1954). Wesselhoft, R. D., Woods, J. M., Smith, J. M., A I C M J . , 5(3), 361 (1959). Weston, D. F., Agnew, J. B., Indian Chem. Eng., 15(3). 37 (1973). Weston, D. F., Ph.D. Thesis, Monash university, Clayton, Australia, 1970.
atm R = gas constant, kcal mol-' K-' r = reaction rate, mol h-l (g of HgC12)F' Greek L e t t e r s p = percent w/w HgCI,, in catalyst u I 2 = percent variance explained by model u2 = percent deviation between measured and predicted data u3 = percent root variance between measured and predicted data u4 = percent deviation between sum of squares and sum of squares due to pure error L i t e r a t u r e Cited
Received f o r reuieu; August 23, 1979 Resubmitted February 13, 1980
Bond, G. C., "Principles of Catalysis' , pp 36-39, Chemical Society, London, 1972
GENERAL ARTICLES Utilization of Cellulosic Feedstock in the Production of Fuel Grade Ethanol Jackson Yu' and Steven F. Miller' Bechtel National, Inc., San Francisco, California 94 1 19
Recent interest in producing ethanol from renewable resources has focused on the use of lignocellulose as a possible feedstock. Ethanol production could become a compatible addition to integrated forest products operations. This paper outlines the status of the various process steps available for liberating fermentable sugars from the lignocellulose, for ferimenting the sugars to alcohol, and for recovering alcohol and byproducts. Process and laboratory studies associ,ated with these steps are discussed. Finally, this paper outlines some of the developmental activities which will lead toward commercialization.
The production of motor fuel quality ethanol from renewable resources has attracted great interest for extending existing petroleum based motor fuel. Fermentation alcohol has been used as a motor fuel in a number of countries and is being used currently in Brazil, South Africa, and certain midwestern states in the United States. Most of the fermentation ethanol produced today is based on sugars, grains, or cassava. Lignocellulose is being seriously considered as an alternative feedstock candidate. However, in order to use lignocellulose, fermentable sugars must be first liberated and separated from the lignin and hemicellulose fractions of the feedstock. The production of ethanol from cellulose is not new. In fact, much of the developmental work was carried out in the 1940's and 1950's. Commercial facilities in the United States and abroad have used cellulosic sugars for ethanol production. However, most current processes are energy intensive. In general glucose yield is low. Figure 1 shows in a block diagram form the basic steps involved in producing ethanol from lignocellulose. The cellulose fraction must be separated and converted to Cutter Laboratories, Inc., Berkeley, Calif. 94710. 0196-4321/80/1219-0237$01.00/0
glucose to provide substrate for the fermentation. The ethanol thus produced must be concentrated from the fermentation beer to produce anhydrous fuel quality material. There are a number of approaches to the liberation of sugars. The three most important process categories are weak acid, strong acid, and enzymatic hydrolysis. Weak acid hydrolysis processes generally use 0.5 % sulfuric acid at a relatively low temperature of 140 of 190 " C to degrade the cellulose. The Madison-Scholler process, as practiced today in the Soviet Union, yields about 50% of theoretical sugars as a 4% glucose solution along with significant degradation of glucose to undesirable byproducts. The low concentration glucose is fermented to only 2% ethanol, thus requiring energy intensive ethanol distillation and byproduct evaporation procedures. Experimental work conducted recently at a higher temperature and much shorter residence time has resulted in a higher glucose yield and reduced glucose degradation (Brenner, 1978; Bender, 1978). At about 500 " C , glucose yield was approximately 70% of theoretical and about a 20 70 post-hydrolysis glucose concentration. This final glucose concentration is accomplished with residence times in seconds rather than the hours required a t the lower
0 1980 American
Chemical Society
