1289
Ind. Eng. Chem. Res. 1991,30,1289-1293 Sahimi, M. On the Determination of Transport Properties of Disordered Systems. Chem. Eng. Commun. 1988,64,177-195. Sahimi, M.; Gavalas, G. R.; Tsotsis, T. T. Statistical and Continuum Models of Fluid-Solid Reactions in Porous Media. Chem. Eng. Sci. 1990,45, 1443-1502. Sahimi, M.; Hughes, B. D.; Scriven, L. E.; Davis, H. T. Real-Space Renormalization and Effective-Medium Approximation to the Percolation Conduction Problem. Phys. Reo. E 1983,28,307-311. Schopper, J. R. A Theoretical Investigation of the Formation Factor/Permeability/PorodlityRelationship Using a Network Model. Geophys. Prospect. 1966,14, 301-341. Sing, K. S. W. The Use of Gas Adsorption for the Characterization of Porous Solids. In Fundamentals of Adsorption; Proceedings of the Third International Conference on Fundamentals of Ad-
sorption, May 7-12,1989, Sonthofen, Bavaria, Federal Republic of Germany; Mersmann, A. B., Scholl, S. E., Ma.; Engineering Foundation: New York, 1990. Stauffer, D. Introduction to Percolation Theory; Taylor & Francis: London, 1985. Stinchcombe,R. B.; Watson, B. P. Renormalization Group Approach for Percolation Conductivity. J.Phys. C: Solid State Phys. 1976, 9,3221-3247.
Wakao, N.; Smith, J. M. Diffusion in Catalyst Pellets. Chem. Eng. Sci. 1962, 17, 825-834.
Received for reuiew May 14, 1990 Revised manuscript receioed November 19, 1990 Accepted December 2, 1990
Kinetics of the Conversion of Calcium Sulfate to Ammonium Sulfate Using Ammonium Carbonate Aqueous Solution Elamin M. Elkanzi* and Mohamed
F.Chalabi
Department of Chemical and Petroleum Engineering, U.A.E. University, P.O.Box 17555, Al-Ain, U.A.E.
A two-step procedure is described for measuring the intrinsic rate of the slurry reaction between calcium sulfate and ammonium carbonate. In the first step, experiments were conducted following the shrinking core model concept to get estimates of the kinetic parameters. These estimates were then used in a mathematical model to predict an experimental program under which the conditions of chemical reaction control is likely. The second step is to measure the intrinsic kinetic parameters under these favorable conditions. A simple power law model fitted to the experimental data of the second step predicted an activation energy of 72 kJ/mol and a pseudo-first-order reaction with respect to ammonium carbonate. The procedure may prove useful in measuring the intrinsic kinetics of fluidaolid heterogeneous reactions and may reduce the effort put on experimental work to establish favorable conditions for chemical reaction control. Introduction Ammonium sulfate for use as a fertilizer is generally manufactured from anhydrous ammonia and strong sulfuric acid. An alternative process originally developed in Germany as long ago as 1909 (Higson, 1951)is now in use. It is based on combining ammonia and carbon dioxide to produce ammonium carbonate, which then reacted with gypsum or anhydrite in a three-phase slurry reactor to yield ammonium sulfate and calcium carbonate. However, little work is reported regarding the kinetics of the reactions involved (Ganz et al., 1959; Chalabi et al., 1975). This may be attributed to the complexity resulting from the influence of physical processes that must take place in order for the overall reaction to proceed. In the threephase gas-liquid-solid system involved here, the overall rate of reaction at any instant of time may be controlled by one or more of the following steps: (1) gas-liquid mass transfer; (2) chemical reaction between ammonia solution and carbon dioxide to form ammonium carbonate; (3) transfer of ammonium carbonate solution through the liquid-solid film; (4)diffusion of ammonium carbonate solution through the calcium carbonate product layer; ( 5 ) chemical reaction at the surface of unreacted calcium sulphate core. As these steps take place in series, it is evident that the slowest step would be rate controlling. Andrew (1954),in an investigation on carbon dioxide absorption by partially carbonated ammonia solutions, reported that the rate of carbon dioxide absorption is controlled by the reaction between carbon dioxide and ammonia. Chalabi et al. (1975)reported that ammonium carbonate was present in
* Author to whom correspondence should be addressed. 0888-588519112630- 1289$02.50/ 0
appreciable amounts even at high conversions of calcium sulfate when the overall reaction was carried out in a fluidized bed reactor. These findings suggest that the first two steps may not control the overall reaction. The purpose of this paper is to report a kinetic study for the reaction described by the equation CaS04.2H20(s)+ (NH4)&03(aq) CaC03(s) + (NH4)2S04(aq)+ 2H20(1) (1) A two-step procedure is to be investigated in order to measure the intrinsic rate of the above reaction. The first step is to establish reaction conditions under which chemical reaction control is likely. These reaction conditions, namely, calcium sulfate particle size, calcium sulfate slurry concentration, ammonium carbonate initial concentration, rate of mixing, and reaction temperatures, are predicted by using a mathematical model based on the relative importance of the resistances offered by steps 3-5 above. The kinetic parameters of the model are estimated experimentally by using the shrinking core model concept (Levenspiel, 1972). The second step is to measure the intrinsic kinetics of the reaction by using an experimental program predicted from the results of the first step.
-
Mathematical Model As far as is known, the effects of mass-transfer limitations on the course of reaction 1 have not yet been investigated. Previous experimental investigations (Ganz et al., 1959; Chalabi et al., 1975) have shown that the calcium sulfate conversion increases with decreasing size of calcium sulfate particles and with decreasing slurry concentrations. Therefore it seems a reasonable assumption that the chemical reaction occurs mainly at the surface of the available calcium sulfate. As the reaction proceeds, 0 1991 American Chemical Society
1290 Ind. Eng. Chem. Res., Vol. 30, No. 6, 1991
the calcium sulfate will be used up, but the particles may become coated with solid calcium carbonate product. Higson (1951) provides a photomicrograph of reaction magma in which crystals of calcium sulfate and aggregates of deposited calcium carbonate are clearly shown. This evidence and the above-mentioned findings would suggest that the overall reaction can be visualized to take place via the following sequence: (a) mass transfer of aqueous ammonium carbonate from the bulk liquid phase to the external surface of the calcium sulfate particle; (b) diffusion of aqueous ammonium carbonate through the deposited calcium carbonate product layer; (c) chemical reaction at the core surface. The relative importance of each of these steps will change as the reaction proceeds. The model equations describing the rates of the steps at one position in the reactor are given as
(-rb = k,ap(CAL - ) ,C
step a:
(2)
step b: this is modeling the shrinking core model for a first-order reaction: d/dr (P(dCA)/dr)= 0
(-rA) = upkCAc
step c:
rc C r C dp/2
r = rc
(3a)
(4)
Considering the progress of reaction from fresh solid spherical particle to acompletely converted particle in an isothermal, irreversible fmborder reaction, the average rate of reaction of the fluid reactant (A = (NH,),CO,) based on unit surface area of particle may be obtained by combining eqs 2-4 as follows: (-r~)av
= CAL/P/(~+ J (dp/4De) + (3/k)l
(5)
Since the shrinking core model envisages the overall process as composed of three resistances in series, the concentration difference of the transferring component (NH1)2C03over each of these steps will be inversely proportional to the resistance to transfer in that step. From eq 5 we have ACA/CAL= [l/kn + (dp/4De)l/[l/k, + dp/4De + 3 / k ] (6) where ACA is the difference in concentration of (NH4)&03 between the liquid bulk ( C,) and that at the gypsum core of a particle (CAc). Reduction of the mass-transfer resistances to low values allows chemical reaction to be rate controlling. Equation 6 is used to make estimates of the relative importance of the resistances to mass transfer, diffusion through product layer, and chemical reaction. Although none of these resistances could be predicted with confidence, it was thought that the effort might lead to a worthwhile reduction in experimental work, before conditions were established where the reaction was truly the rate-determining step.
Estimation of Model Parameters The liquid-film mass-transfer coefficient may be estimated in the case of a mechanically agitated reactor from correlations reviewed by Pandit et a1 (1986): (ha/ U ) N s $ 7 6 ( d p U p ~ / p ~ ) o 3.04 .6
(7)
when 17 C N b C 270 and 13.7 C N h C 3200, or from
k,/ U = 2.165/ (Nb)O.*(Ns,) 2/3 when N b C 17.
(8)
The effective diffusivity of (NH,),CO, solution through ciolid CaC03 may be estimated from the molecular diffuriivity using a solid tortuosity factor (Satterfield, 1980; Epstein, 1989). The value of the reaction velocity constant has not been reported in literature. Estimates of this parameter have to be determined experimentally. Here the application of the Shrinking core model to the reaction would yield an estimate of the apparent reaction rate constant provided that the assumptions and conditions embodied in the model are approached. With these parameters being established, eq 6 becomes a function of calcium sulfate particle size and of speeds of mixing and of temperature. A plot of ACA/C, versus calcium sulfate particle size, with temperature and with speed of mixing as parameters, would yield a family of curves from which the conditions under which chemical reaction control is likely may be predicted.
Experimental Section The first set of experiments were conducted to establish an estimate for the apparent reaction rate constant under conditions approaching those embodied in the shrinking core model. The reaction conditions selected were as follows: (1)Calcium sulfate slurry concentration was of 5% by weight (Ganz et al., 1959; Chalabi et al., 1975). (2) Ammonium carbonate aqueous concentration was kept constant at ita saturation value at each temperature. In this manner the calcium sulfate conversion would be a function only of its slurry concentration. This would ensure that the time for complete conversion and the apparent reaction rate constant be constant for constant temeprature, particle density, and diameter. The assumption of pseudo-first-order kinetics embodied in the shrinking core model would also be approached. (3) Calcium sulfate particle size was kept below 50 pm. (4) The temperature was varied between 10 and 20 "C in steps of 5 "C. These temperatures were lower than those used in previous work. (5) Speed of mixing (impeller speed) was kept constant at its maximum value of 2200 min-'. The second set of experiments were conducted to establish the intrinsic rate expression for the reaction under conditions predicted by the mathematical model of eq 6. The independent parameters investigated include ammonium carbonate concentration, temperature, and conversion time. Apparatus. All experiments were conducted in a l-L three-necked round-bottom flask supplied with a stirrer, thermocouple ports, a sampling port fitted with a microfilter, and suction pump. Agitation of the reactor contents was carried out with a pitched-blade turbine having three blades. The turbine was driven by a variable-speed motor. The reactor temperature was maintained constant by a constant-temperature bath provided with a thermoelectric controller. The reactor was operated in a batch manner. Materials. The calcium sulfate used in this study was Analar quality microfine powder. The ammonium carbonate used was of the commercial type (NH4HC03+ NH2COONH4)genral reagent quality. Deionized water was used. Other chemicals used as reagents in analysis were all of the AnalaR type. Procedure. For each experimental run, a weighed amount of ammonium carbonate was charged to the reactor containing 500 mL of deionized water. Stirring of the reactor content, at 36.6 revolutions/s, was then started. When the desired temperature had been reached, calcium sulfate already cooled to the reaction temperature was added as required for the reaction to start. It is important
Ind. Eng. Chem. Res., Vol. 30,No. 6,1991 1291
to charge first the ammonium carbonate because of its highly negative heat of solution. The course of the reaction was followed by means of 1-mL clear-liquid samples that were analyzed for ammonium sulfate content through the determination of the sulfate radical. The analytical method devised was based on measuring light absorption by a turbid suspension caused by reacting barium chloride with the already treated sample containing ammonium sulfate. For this purpose a Pye Unicam Model SP8-400double-beam UVvis spectrophotometer was used. The analytical method is quite sensitive, and ammonium sulfate concentrations as low as 1 X lo4 M can be detected. The accuracy of the method depends upon the sample being dilute and therefore the original 1-mL sample goes through a series of dilutions. Corrections to the analytical results were made on account of the sulfate content of dissolved calcium sulfate, of the withdrawal of liquid by sampling, and of dilution due to the generation of water by reaction. In the first set of experiments the data were taken in the temperature range 10-20 O C , ammonium carbonate concentrations corresponding to the saturation value at each temperature, and calcium sulfate concentration of 0.05 kg/L of water. In the second set of experiments the data were taken under the conditions predicted by the mathematical model, namely, the temperature range and calcium sulfate particle size. Ammonium carbonate initial concentrations ranged from 0.1 to 0.2 M and calcium sulfate was present in exceas. Different initial concentrations of ammonium carbonate were used so as to measure the kinetic parameters over a wider range of concentrations.
Rssults and Discussion For the process design, it is important to obtain the kinetic expression for the reaction of eq 1. This reaction is the most important and effective one occurring during the ammonium sulfate production process. A two-step procedure is followed in this study: (a) determination of the conditions under which the reaction is chemically controlled and (b) measurement of the intrinsic kinetic parameters under the conditions determined in step (a). Results from the kinetic analysis of each step are discussed separately. Reaction Conditions. The conversions for calcium sulfate with time, at temperatures of 10,15, and 20 OC and ammonium carbonate concentration of 1.38,1.60, and 1.85 kmol/m3 (saturation value at each temperature, respectively) were obtained. These experimental data were analyzed on the basis of the shrinking core model. Levenspiel (1972) showed that when the rate is controlled by the chemical reaction at the core surface, then for both shrinking and constant-size particles the solid reactant conversion is given by t/T = 1 - (1 - x)1/3
(9)
where T
= pBdp/2bkCfi
(10)
The experimental data were plotted on the basis of eq 9 for all three temperatures as illustrated in Figure 1. The values of T were obtained from the slopes of the lines in Figure 1. The value of k can then be calculated from eq 10. It was assumed that the size distribution of calcium sulfate particles remains constant and an average value for dp of 30 pm was used. The molar density of calcium sulfate is 24.6 kmol/m3. The value of b can be obtained from eq 1. The saturation concentration of ammonium carbonate,
.55
-
45-
D
1 :291K
&
T
.2edn
T
281
. 0
2
6
4
8 REACTION TI=,
IO Minules
n
12
14
16
Figure 1. Shrinking core model.
-2.5-
3-
. P
-1.5
-
3.4
1.12
1.47
1.45
3 50
1.52
1.55
1 I 103 n" 1
Figure 2. Arrheniua plot of the shrinking core model.
C, at any temperature was kept nearly constant by using excess carbonate. Thus all the parameters in eq 10 are constant at each temperature, and hence k would change with temperature in approximately the same ratio as the reciprocal of T. A plot of In 1/r against the reciprocal of absolute temperature is illustrated in Figure 2. The slope of the line indicates an apparent activation energy of 70 kJ/mol for the reaction of eq 1. The apparent reaction rate constant can be represented by k = 1.54 X (11) Since the stirring speed was kept constant at 36.6 revolutions/s, the value of k, can be estimated from eq 7. An order of magnitude value for the effective diffusivity De of 3 X lo-" m2/s is estimated from the molecular diffusivity by using a solid tortuosity factor. Substitution of these values in eq 6 allows the calculation of the fractional change in concentration with the particle diameter. Figure 3 illustrates some of the results as a plot of ACA/Cfi vs calcium sulfate particle diameter with temperature as a parameter. As expected, the fractional change in concentration decreases with decreasing particle size at each temperature. Figure 3 indicates that with particles of less than 50 pm and with temperature of 10 OC or lower, the mass-transfer limitations on the rate of reaction would be less than 2%. Therefore it seemed likely that with a stirring speed of 36.6 revolutions/s, with temperatures of 10 "C or lower, and with particle sizes below 50 pm, the reaction of eq 1 would be effectively controlled by the speed of the chemical reaction. Besides the effect of the particle size on mass-transfer limitations, the effect of
1292 h d . Eng. Chem. Res., Vol. 30, No. 6, 1991
1.5-
(5)
T
i
293 K 1.25-
5
( 4 ) _1
20
1 -
,751
13)
(2) o 1
0 PARTICLE
DIAMETER,
20
40
60
80
r m
Figure 3. Fractional change in concentrationsv8 particle diameter (es 6).
12C
Figure 4. Continuous reaction model. ,
I5
352
355
357
3 5 f I 10
temperature level should not be overlooked. In fact, the reaction of eq 1 is both reversible and parallel with respect to ammonium carbonate. Nowak et al. (1989) studied the kinetics of the decomposition of ammonium carbonate in aqueous solutions. They found that the reaction is second order with respect to ammonium carbonate concentration and has an activation energy of 98.9 kJ/mol. Therefore, low-temperature level and low concentration level would not favor the decomposition of ammonium carbonate. To minimize the effect of this side reaction in comparison to the main reaction of eq 1, and to make mass-transfer limitations small, the following reaction conditions may be used for the study of the intrinsic rate of the reaction of eq 1: (i) calcium sulfate concentration of 5 wt 90' (Ganz et al., 1959; Chalabi et al., 1975); (ii) ammonium carbonate aqueous concentration of 0.1-0.2 kg mol/m3; (iii) speed of mixing of 36.6 revolutions/s; (iv) temperature range of 0-10 OC; (v) calcium sulfate particle sizes below 50 pm. Intrinsic Rate Measurement. Several runs were made using ammonium carbonate initial concentrations of 0.1, 0.15, and 0.2 kmol/m3. Calcium sulfate concentration was kept constant a t 5 wt 9'O (this is in excess of its stoichiometric proportion of eq 1). This would allow the surface area to remain essentially constant during the course of the reaction and the rate to be independent of the calcium sulfate concentration. Each run was repeated in duplicate and repeated for a third time to check reproducibility. Concentration vs time results were as presented in Table I for the initial concentration of ammonium carbonate of 0.1 kg mol/m3. Having correlated the results in the absence of masstransfer limitations, it was found that the kinetic equation representing the continuous reaction model, of a firsborder
278 K
REACTION TIME,MInulcr
8
Tuble I. Experimental Results for the Reaction Starting with 0.1 ka mol/m8 (NHACO, Aqueous Solution (NH,)$O, concn, kg mol/m3 X lo2 reaction time, min T = 273 K T = 278 K T = 283 K 0 0 0 0 1.2 0.8 10 0.5 1.5 2.3 0.9 20 2.1 3.3 30 1.3 2.1 4.1 1.7 40 5.4 2.4 3.8 60 4.7 6.5 3.0 80 5.4 7.3 3.1 100 1.9 6.1 4.3 120 8.4 4.8 6.1 140 1.2 5.2 160 1.6 5.6 180
1W
i
'.
362
365
357
37
K1
Figure 5. Arrhenius plot of the continuous reaction model.
irreversible reaction, is best satisfied. The integral form of the kinetic equation is In C,/CA = k,t (12) The fitting of the above equation to the experimental data is illustrated in Figure 4. The rate constants k, a t different temperatures were calculated from the regression analysis of these data. An Arrhenius plot shown in Figure 5 gives an activation energy of 72 kJ/mol for the reaction of eq 1. From this plot the estimated frequency factor is 4.4 X lo9 s-l at an ammonium carbonate concentration of 0.1 kmol/m3 and in the temperature range 0-10 O C . At these conditions the following equation holds: k, = 4.4 X lo9 exp(-17234/Rl') (13) The rate of reaction was found to be independent of the initial concentration of ammonium carbonate in the range 0.1-0.2 kmol/m3. Substituting the value of k, from eq 13 into eq 6 and using a particle diameter of 50 pm and the previously estimated values of k, and De,the percent fractional change in concentration was less than 0.019 i for the temperature range studied. This clearly demonstrates that masstransfer limitations were negligible and that the rate of conversion is controlled by the reaction step. Conclusions The work presented in this paper applies a general method involving a two-step procedure for the measurement of intrinsic kinetics of fluid-solid heterogeneous
Ind. Eng. Chem. Res. 1991,30, 1293-1300
reactions. The method has been applied to the measurement of the intrinsic kinetics of the conversion of calcium sulfate to ammonium sulfate using aqueous "onium carbonate solution. In the first step the conditions for chemical reaction control were established. In the second step the intrinsic kinetics were measured under the conditions established in the first step. The experimental findings showed that the main advantage of this method is that it helps to predict the reaction kinetics by establishing the conditions where the chemical reaction is truly the rate-controlling step. Under the experimental conditions studied here, the kinetics of the conversion of calcium sulfate to ammonium sulfate using ammonium carbonate aqueous solution were tentatively represented by a pseudo-fmt-order model, and an activation energy of 72 kJ/mol was obtained.
Nomenclature a, = area per unit volume, m-l b = stoichiometric coefficient (eq 10) CA = concentration of aqueous ammonium carbonate, kmol/m3 CAc= concentration of aqueous ammonium carbonate at the core surface, kmol/m3 C, = concentration in bulk liquid, kmol/m3 C, = initial concentration, kmol/m3 Ck = concentration at the surface of the solid, kmol/m3 ACA = molar concentration drop of component A, kmol/m3 De = effective diffusivity of aqueous ammonium carbonate solution through calcium carbonate ash layer, m2/s d, = particle diameter, m k = apparent reaction rate constant, m/s k, = intrinsic rate constant, s-l k, = liquid-film mass-transfer coefficient, m/s NRe = Reynolds number, defined by N R =~ dpUeL/pL, dimensionless N k = Schmidt number, dimensionless R = gas constant, cal/mol K r = particle radius, m (-rA) = rate of reaction, kmol/m3 s r, = core radius, m T = temperature, K
1293
t = reaction time, s
U = slip velocity, m/s U, = relative velocity between a particle and liquid generated
X
because of impeller action, m/s = fractional conversion
Greek Letters = liquid viscosity, kg/ms = molar density of calcium sulfate, kmol/m3 pL = liquid density, kg/m3 7 = time required for complete conversion, s Registry No. CaS04, 7778-18-9; (NH4)2C03, 506-87-6; pL pB
H4)2S04,7783-20-2.
(N-
Literature Cited Andrew, S. P. Carbon dioxide Absorption by Partially Carbonated Ammonia Solutions. Chem. Eng. Sci. 1954, 3, 279. Chalabi, M. F.; Younis, S. M. Conversion of CaS04 into (NH4)804 at Ammoniated Gypsum Slurries by Co, Gas. Presented at Chisa International Congress on Chemical Engineering, Prague, Checozlovakia, 1975. Conversion of Gypsum or Anhydrite to Ammonium Sulfate. Nitrogen 1967, No. 46. Epstein, N. On tortuosity and the tortuosity factor in flow and diffusion through porous media. Chem. Eng. Sci. 1989,44,777-779. Fertilizer Manual; No. IFDC-R-1; International Fertilizer Development Center: AL, December 1979. Ganz, S. N.; Leibovish, S. B.; Gorbman, S.I. Conversion of CaS04 into (NH4)2S04by Agitating a Mixture of Gypsum Slurry, NHS and Coz in a Variable Speed Disc Mixer. Zh. A'klad. Khim. 1959, 32,975-978.
Higson, G. I. The Manufacture of Ammonium Sulfate from Anhydrite. Chem. Ind. 1951, Sept 8,750-754. Levenspiel, 0. Chemical Reaction Engineering, 2nd ed.; Wiley New York, 1972. Nowark, P., Skrzypek, J. The Kinetics of Chemical Decomposition of Ammonium Bicarbonate and Carbonate in Aqueous Solutions. Chem. Eng. Sci. 1989,44, 2375-2377. Pandit, A. B.; Joshi, J. B. Maes and Heat Transfer Characteristics of Three Phase Sparged Reactors. Satterfiold, C. N. Heterogeneow Catalysis in Practice; McGraw-Hilk New York, 1980. Received for review May 25, 1990 Accepted December 8, 1990
Heterogeneous Photocatalytic Oxidation of Phenol with Titanium Dioxide Powders Tsong-Yang Wei and Chi-Chao Wan* Department of Chemical Engineering, National Tsing Hua University, Hsinchu, Taiwan, Republic of China
The photocatalytic oxidation of phenol in oxygenated solution with suspensions of titanium dioxide powders has been investigated. The experimental results indicate that there is an optimum value for the Ti02content. Sufficient O2supply is needed, but a too high gas flow rate induces large bubbles that eliminate the gas residence time and the gas/liquid contact interface. The initial concentration of phenol demonstrates a negative effect on the pseudo-first-order reaction rate constant. A mechanism of the phenoxide ions being adsorbed on the Ti02 surface has been proposed to account for this inverse effect. Moreover, the different pH regions affect the competitive adsorption between the phenoxide ions and OH- ions on the Ti02 surface, which in turn influences the generation of the hydroxyl radicals on the T i 0 2 surface.
Introduction Phenol is a major pollutant and under strict effluent restriction. owingto its and solubility in water,
* To whom correspondence should be addressed.
the degradation of this compound to reach the safety level in the range of 0.1-1 mPL-' is not easy. Many conventional methods have been proposed to treat phenol solutions (Devlin and Harris, 1984; Eisenhauer, 1964, 1968; Getoff, 1986; Hackman, 1976; Kharlamova and Tedoradze, 1981; Sadana and Katzer, 1984; Sharifian and Kirk, 1986;
0888-5885/91/2630-1293$02.50/00 1991 American Chemical Society
