SOLVENT DESIGN FOR CHEMICAL REACTIONS K .
F .
W O N G
A N D
C.
A .
E C K E R T
Department of Chemistry and Chemical Engineering, University of Illinois, Urbana, Ill. 61801 Thermodynamic analysis can be used with the transition state theory to predict solvent effects on the rates of chemical reactions. This technique may be applied to the prediction of the optimum solvent for a reaction; often a properly chosen solvent mixture may be better for a given reaction than any pure solvent. Also, this mixture is applied to sets of competing reactions. Examples are presented in terms of Diels-Alder reactions.
CHEMICAL reactor design generally includes consideration of many interrelated factors, such as temperature, pressure, and heat and mass transport. However, for the rational design of a liquid-phase reaction one other factor can be of prime importance-choice of the optimum solvent system. Kinetic solvent effects on reaction rates can be enormous; variations in rate due to solvent of the order of a factor of 100 to 1000 are rather common, and in some cases factors as high as lo9 have been observed (Cram et al., 1961). A method of predicting such effects would provide a very powerful tool for the design engineer. Recent reviews on various aspects of solvent effects on reaction rates are available (Amis, 1966; Baekelmans et al., 1964). Empirical correlations of solvent effects and chemical reactivities using the linear free energy relationships have been well discussed by Wells (1963). Frisch et al. (1962) have developed a corresponding states treatment using the critical parameters, which gives a fair correlation, a t least qualitatively, of the relative rates of several types of reactions in noninteracting solvents. This work demonstrates a method for predicting the relative rates of chemical reactions in various solvents and solvent mixtures, based on the transition state theory and thermodynamic properties. This has the advantage of reducing a problem in kinetics to a more tractable one in thermodynamics, since both the theoretical correlations and the experimental data are better for thermodynamic properties than for rate processes. Examples of this technique are presented in applications to the prediction of kinetic solvent effects in pure solvents and solvent mixtures, and for the case of competing reactions. The two limitations of this method are that it does not predict absolute reactions rates, but only relative rates, and, although it may be capable of extension to more complex situations, thus far it has been applied only to elementary reactions with well-defined mechanisms. Transition State Theory
The transition state approach defines an intermediate species, M, in an elementary reaction, called the activated complex
A +B
- M
product
The assumption of quasi-equilibrium between the reactants and the activated complex leads to an expression for the rate of reaction in a thermodynamically nonideal system, given by Bransted (1922) and Bjerrum (1924) 568
I & E C PROCESS D E S I G N A N D DEVELOPMENT
where h, is the rate constant in the thermodynamically ideal reference system, the y’s are the activity coefficients of the individual species, and the [ A ] ,[ B ]are the concentrations of A and B. I n the ideal system all the activity coefficients are unity. The apparent rate constant in the real system is given by
(3) The validity of this result depends on the postulation of quasi-equilibrium between transition state and reactants. Marcus (1967) has shown that an equivalent result may be obtained using the adiabatic assumption. Three assumptions are inherent in the derivation of transition state theory. Briefly, the motion of the nuclei is assumed to be adiabatic and to obey the laws of classical mechanics. Deviation from classical mechanics is usually small and requires only small corrections. This is the BornOppenheimer approximation in quantum mechanics. In addition, a Maxwell-Boltzmann distribution is necessary. Clearly the transition state theory fails in the case of fast reaction. As a rule of thumb, it is valid only if the activation energy of the reaction is greater than 5 kcal. per mole. These assumptions are discussed by Kondratev (1964). The applicability of Equation 3 was first demonstrated in the study of ionic reactions in dilute aqueous solutions using the Debye-Huckel treatment. I t has been applied to gas-phase reactions (Eckert and Boudart, 1963; Mills and Eckert, 1968) in terms of valid equations of state. For liquid-phase reactions a number of different approaches can be used (Eckert, 1967) for obtaining the activity coefficients in Equation 3-various expressions for excess Gibbs energies of mixtures-or when available, experimental data are preferable. For any application, the units in Equation 3 must be consistent. For example, in the Debye-Huckel treatment, activity coefficients are expressed in volumetric concentration units (moles per liter), and so are the rate constants. For most other expressions for activity coefficients, the concentration units are the dimensionless mole fraction, so we use consistent rate constants in mole fraction units, k,. This point has often been overlooked. For reaction in dilute solution, Equation 3 is valid, with k and k ,
replaced by k , and k x o , respectively. For a bimolecular reaction with volumetric concentrations in moles per cubic centimeter, these quantities differ from each other by a constant factor, the molar volume of the solvent. Thus by Equation 3, the prediction of relative rates for a simple reaction becomes a thermodynamic problem of accounting for nonidealities with a suitable solution theory or sufficient thermodynamic data, or for the transition state, a technique for estimating its activity in solution. For demonstration here we choose to use the regular solution theory of Hildebrand and Scott (1962, 1964). Although this approach is not the only possible one, it has certain advantages. I t is simple, requiring only purecomponent data, and it is readily extended to multicomponent mixtures. I t has the limitation of applying strictly only to solutions of nonionic, nonpolar, or slightly polar molecules. Applying this technique to kinetics was first proposed by Glasstone et al. (1941). According to the regular solution theory, the activity coefficient of component i in a multicomponent mixture is given in terms of its liquid molar volume, u,, solubility parameter, 6, (the square root of the cohesive energy density), and the volume-fraction average solubility parameter of the solvent, 6.
RT In
yi = ~ ~ (6 ,6)'
(4)
Combining Equations 3 and 4, the relative rate expression for an elementary reaction has been derived (Wong, 1967)
Equation 5 is valid under conditions where the assumption that there is no excess entropy of mixing is justified. For dilute solution the average solubility parameter can be approximated by the solubility parameter of the solvent. Then Equation 5 reduces to that of Frost and Pearson (1961). The values of the molar volume and solubility parameter for the reactants in Equation 5 are readily available purecomponent properties. However, the properties of the transition state may also be obtained from independent measurements. Its volume is available from the effect of pressure on the reaction rate in the liquid phasei.e., volume of activation data. For the solubility parameter. we combine this volume with the energy of vaporization, found from a thermodynamic cycle similar to that suggested by Harris and Prausnitz (1969),
A,
+
B, + M I
The energy of vaporization of the complex is equivalent to the liquid-phase dissociation of M to species A and B, vaporization of A and B, and the gas-phase recombination to give M. To a good approximation the activation energies of the dissociation and the recombination are equal. The 6.w is given by
Prediction of Solvent Effects
The Diels-Alder reaction provides a good class of reactions to use as examples. It consists of the addition of a compound containing a double or triple bond (usually activated by additional unsaturation in the a , 6-position) to the 1,4-positions of a conjugated diene system, with the formation of a six-membered hydroaromatic ring. The simplest example is the addition of butadiene (diene) to ethylene (dienophile) to form cyclohexene. A number of such reactions have been studied and the Sesults are available. I n general, there are few side reactions and the reactants and products are not too highly polar. The Diels-Alder condensation of isoprene with maleic anhydride to yield 4-methylcyclohexene-1,2-dicarboxylic
anhydride has been measured by Dewar (as quoted by Frisch et al., 1962) in 10 solvents, Grieger (1969) in five solvents, and Snyder (1968) in two solvents. For this reaction, the volume of the activated complex has been measured by Grieger and Eckert (1969), and its solubility parameter has been estimated by Equation 6. I n general, in applying Equation 5 it is convenient to eliminate h, by choosing one solvent as a reference. The rate in benzene is taken as the standard rate; all other rates are relative to that. The experimental results are given in Table I and compared with the predictions of regular solution theory (Equation 5) in Figure 1. Except for the highly polar solvent, nitromethane, the theory is in good agreement with the data. Table I. Reaction of Isoprene and Maleic Anhydride
Reactants and Complex Isoprene Maleic anhydride Transition state
Molar Volume, Cc. 101 74" 145O
Solubzlity Parameter, (Cal Cc.)"
Solvents Diisopropyl ether
7.0
n-Hexane n-Butyl chloride Ethyl acetate Benzene m-Dichlorobenzene Chlorobenzene 0-Dichlorobenzene Anisole o-Dimethoxybenzene Phenylisoc yanide Nitrobenzene
7.3 8.45 9.04 9.15 9.2 9.5 9.55 9.8 9.8 10.0 10.0
Dichloromethane Kitromethane
10.2 12.6
Sol ub 11it) Parameter, (Cal CCI" 7.45 13 11.5 Rate Constant Relative to Benzene, log ( k , k , 1 at 30 3" C -0.745 -0.648' -0.E195~ -0.443' -0.333' 0' 0.170' 0.100' 0.480' 0.069' 0.093' 0.233 0.419 ' O.44Od 0.316' 0.498' 0.573'
' Extrapolated from density in liquid range. ' Data of Grieger (1969). Data of Deuar as quoted by Frisch et al. (1962). a Data of Snyder ( 1 968).
VOL. 8 NO. 4 OCTOBER 1 9 6 9
569
',O
/
t
Table II. Dimerization of Cyclopentadiene
/
0-Dichlorobenzene Nitrobenzene
0.2
i
Dichloromethane
Soluents
-
Ethyl Acetate [r
-0.4
0 n-Butyl Chloride
-0.6c
A !,
n -Hexane
/&i-isopropyl
20
40
60
Ether
80
1
0 Data of Dewor 0 Dota of Grieger A Data of Snyder -
100 120
140 160 180 200
Cohesive Energy Density (8') cal /cc
Figure 1. Reaction of isoprene-maleic anhydride in various solvents Comporison with prediction of regulor solutlon theory
Another example is the liquid-phase Diels-Alder dimerization of cyclopentadiene. This reaction has been studied both in the pure state and in various solvents by Wassermann (1952), Harkness et al. (1937), Schmid et al. (1948), and more recently Shuikin and Naryshkina (1957). The volume of activation has also been measured by Raistrick et al. (1939). The data are shown in Table 11, and the comparison with the regular solution method is presented in Figure 2. The dimerization in the pure liquid has been done by several authors (Harkness et al., 1937; Schmid et al., 1948; Shuikin and Naryshkina, 1957; Wassermann, 1952). The deviations among the various authors and consequently the disagreement with the predictions of the regular solution theory are within the experimental error estimated by Wassermann (1952). As Harkness et al (1937) pointed out, the reaction in the pure liquid is complicated by polymerization, especially in the initial stage of the reaction. The scattering in the case of carbon tetrachloride further shows the extent of uncertainty of the available kinetic data. The cohesive energy density has often been used with moderate success as a correlating parameter for kinetic solvent effects, not only for Diels-Alder reactions but for many other types of reactions. I n most cases the reactants and solvents are so polar that the success of the correlation must be considered surprising. For example, Stefani (1968) has shown such a result for the solvent effect on the rates of two Menschutkin reactions-the reaction of an alkyl halide with a tertiary amine (both highly polar) to form an ionic quaternary ammonium salt. For the reaction of pyridine with methyl iodide in 13 solvents ranging in solubility parameter from 7 to 11 (cal./cc.)' (including hydrocarbons, ethers, organic halides, and nitro compounds), Stefani found a good linearity of the logarithm of the rate constant with respect to the square root of internal pressure (equivalent to the solubility parameter). Equation 5 has been applied to these data, using the value of 110 cc. per mole for 570
C yclopentadiene Transition state
Solubility Parameter, iCal. Cc.)' '
83.1 135O
8.76 9.75
Solubility Parameter, (Cal. cc.y
Rate Constant Relative to Benzene, W h y k, 1 at 25. c.
Qhtgmethani
rn-DichlorobenzeneD Phenylisocyanide Chlorobenzene g-o Dimethoxybenzene Benzene Anisole
0
Reactant and Complex
Molar Volume. Cc.
I & E C PROCESS D E S I G N A N D D E V E L O P M E N T
a-Dicyclopentadiene Tetralin Carbon tetrachloride
8.35 8.59 8.6
C yclopentadiene
8.76
Toluene Benzene Nitrobenzene Carbon disulfide
8.9 9.15 10.0 10.0
0.07O -0.675' 0.llC -0.94' 0.11' -0.25d -0.59' -0.65' -0.214O On
0.504d -0.192'
a Data of Raistrick et al. (1939). Data of Schmid et al. (1948). 'Data of Harkness et al. (1937). dData of Wassermann (1952). 'Data of Shuikin and iVaryshkina (1957).
the volume of the transition state, from the volume of activation data of Hartmann et al. (1965). Stefani's correlation corresponds to a solubility parameter of the Menschutkin complex of about 14 (cal./cc.)' '. Similarly, Equation 5 has been applied to compare with Stefani's correlation for the triethylamine-ethyl iodide reaction in a similar range of solvents. Using a molar volume of the transition state of 190 cc. per mole (Kondo, Tojima, and Tokura, 19671, the correlation corresponds to a solubility parameter for this Menschutkin complex of about 13.5 (cal./cc.)' '. Since the transition state for the Menschutkin reaction is known t o be nonionic but highly polar, the two values found for the solubility parameter of the complexes are qualitatively reasonable. However, the rate constants should be based on mole fraction to be strictly correct. Design of Solvent Mixtures
The greatest utility of applying thermodynamics to the analysis of kinetic solvent effects lies in the optimum design of the solvent. In many cases, as evident from Figures 1 and 2, the optimum solvent would be merely that solvent with the lowest or highest possible cohesive energy density. However, such is not always the case. Recently on the basis of the Gibbs-Duhem equation, it has been shown that in any ternary system one component must exhibit an extremum in its activity coefficient (Grieger and Eckert, 1967). Since the relative rate, h,/ h r , is proportional to the ratio of the individual activity coefficients by Equation 3, the rate of reaction may also show an extremum. In many cases then, a properly chosen mixture of solvents may maximize or minimize the rate of reaction better than any individual solvent alone. For the case of reaction in a binary solvent mixture where the regular solution behavior is followed and the reactants are dilute, Equation 5 can be used with the average solubility parameter given by
1.2
l
1
'
i
"
'
1.0
Nitrobenzene 0
I
lI
l
l
i
Prediction of
0.6
0.4 Pure Liquid
Toluene
-0.2
Carbon Disulfide Pure Liquid
-0 4
a Carbon Tetrachloride
-1,o
1 0,2
0 Data of Wasserman
A Data of Schmid
V
Data of Harkness
0
20
40
60
100 120 140
80
Cohesive Energy Density ( 8 ' )
160 180 200 col./cc
Figure 2. Relative rate of dimerization of cyclopentadiene in various solvents
where @ is volume fraction and the subscripts 1 and 2 refer to the two solvents. Then the extremum of-the ratio of the rate constants k x / k x o with respect to 6 is given by -
6 extremum
-
+ U B ~ B- U M ~ M
U A ~ A
(UA f U B
- UM)
1
I
0.6
0.8
1
I .o
Volume F r a c t i o n Neapentane
0 Data of Shuikin
- 1.6
I
0.4
(8)
The denominator in Equation 8 is no more than the negative of the volume of activation for the reaction. Further differentiation indicates that whether the >extremum given is for a maximum or minimum depends on the sign of the volume of activation. For a Diels-Alder reaction where the volume of activation is generally negative, the 6exrremum given is for a minimum in the rate. The Diels-Alder dimerization of isoprene provides an example of this phenomenon. Calculations based on regular solution theory, assuming a value of 6~ = 7.1 (cal./cc.)' ', show that a mixture of about 3 0 ' ~ (by volume) of neopentane with carbon disulfide will give a minimum in the rate constant. The rate in pure CSS is about 6% faster than in the mixture and in pure neopentane more
than 35% faster. The results are shown in Figure 3. Although the effect in this case is not terribly large, it illustrates the extremum possible in a mixture. For more polar mixtures (for which regular solution theory is less applicable) the effect would be much greater. For example, experimental results (Szmant and Roman, 1966) for the rate of the Wolff-Kishner reaction of benzophenone hydrazone show a maximum in the rate in a mixture by butyl carbitol and dimethyl sulfoxide. This maximum rate appears to be a t least two orders of magnitude faster
Figure 3. Calculated rate of isoprene dimerization in neopentane-carbon disulfide mixtures at 0" C.
than the rate in either pure solvent. Kondo and Tokura (1964, 1967) have applied the transition state theory and a solution model to correlate some Menschutkin reactions in mixed solvents. Competing Reactions
The same type of analysis can be applied to calculation of the relative rates of two competing reactions. Consider two bimolecular reactions,
A~ + B 3 - i . ~ ~ A?+ B 3~~ where k l and hz are the rate constants for the two reactions, AI and Al are the reactants competing for B, and M I and MS are the corresponding activated complexes. The variation of the relative rate from solvent to solvent is given by
where ( h l / h 2 ) 0again refers t o the ideal reference system defined previously. If sufficient thermodynamic data are known, the relative rates of the two competing reactions can be calculated. Otherwise the regular solution theory can be applied
An example of a pair of competing reactions is the Diels-Alder condensation of methyl acrylate with cyclopentadiene, which yields two products.
VOL. 8 N O . 4 OCTOBER 1 9 6 9
571
The separate reactions, through separate transition states, to form the two possible stereoisomeric products, constitute a pair of competing reactions. Since both reactants are the same for both reactions, the two terms in A drop out of Equation 10 and it becomes
Table 111. Reaction of Cyclopentadisne and Methyl Acrylate
Reactants and Complex
Molar Volume,
cc.
Cyclopentadiene Methyl acrylate Complex
ks In - = ik,)
83.1 90.3 150"
Relative Rates in Various Solvents
I t is reasonable to assume that the molar volumes of the two isomeric transition states are equal, u . ~ , = u.~,. Then, Equation 11 becomes
or the natural logarithm of the relative rate constants ( k c / k x ) is linear in 6. The solvent effect on this reaction has been studied by Berson et ul. (1962), and Stefani (1968) has empirically found the ratio (h,\ihx) for these data to be a linear function of the square root of the cohesive energy density of the solvent. Equation 12 has been applied to the data of Berson et ul. (1962), and the solubility parameter difference for the two activated complexes, (&.M, - 6.M,), adjusted to give the best fit. Since the solvents used were generally fairly polar, the agreement between theory and experimental data, shown in Figure 4, must be considered remarkably good. The parameters used for this comparison are given in Table
Triethylamine !,2-Dimethoxyethane 1,2-Dichloroethane Pyridine Dimethylformamide Acetonitrile Ethanol Nitromethane a
Solub zlzty Parameter, (Cal cci
Log ( k l k x ) at 30" c n
7.85 8.8 9.8 10.7 11.2 11.8 12.2 12.6
0.445 0.543 0 600 0.595 0.620 0.692 0.718 0.680
Estimated frum volume of activation for Diels-Alder reactions. Data of Berson et al. 11962)
111. This best fit is achieved by a rather large value of (6.M, - 6 ~ v , ) of -2.44 (cal./cc.)'*. However, such a value is entirely reasonable; the endo form of the transition state (as of the product) is expected to have a larger permanent dipole moment than the eno form, and thus has a much higher heat of vaporization and a higher solubility parameter. Since in this case a ratio of the rate constants is involved, the concentration units used can be arbitrary as long as the reaction is run in dilute solution.
Conclusions l
!
l
I
I
l
l
o,61 ,/
xi2 -
l
I , 2 -Richloroethane 0
0 Pyridine lformamide
I,2-DimethoxyethaneO
pl
-
-I
0,5
W
0,3r
0 Roto of Berson
5
6
7
8
9
IO
II
12 3I,
14
15
S Solubility Parameter ( c o l / c c )
Figure 4. Competing reactions of cyclopentadiene and methyl acrylate Comparison with regular solution theory
572
Nomenclature h, = rate constant, time-' k , = rate constant in ideal solution, time-' R = gas constant, cal./mole-"K. T = absolute temperature, K. u = liquid molar volume, cc./mole y = activity coefficient 6 = solubility parameter, (cal./cc.)' ' SUBSCRIPTS O
5,5
I
".L
Thermodynamic analysis can be used in conjunction with the transition state theory to predict kinetic solvent effects on elementary reactions. Although experimental thermodynamic data would be preferable, they are generally not available, so some theory of solutions must be applied. Here, regular solution theory has been used to predict the rates of Diels-Alder reactions, to design an optimum solvent mixture, and to compare the rates of competing reactions. I n theory, this type of analysis could be extended to complex sequences of elementary reactions, often encountered in reactor design. The success of such an extension would, of course, depend on knowledge of the detailed mechanism for each step and on the accuracy with which activity coefficients are known, or could be estimated.
I & E C PROCESS D E S I G N A N D DEVELOPMENT
A, B g 1 M
= reactants = gas phase = liquid phase
= transition state
Literature Cited
Amis, E. S., “Solvent Effects on Reaction Rates and Mechanisms,” Academic Press, New York, 1966. Baekelmans, P., Gielen, M., Nasielski, J., Ind. Chim. Belge 29, 1265 (1964). Berson, J. A., Hamlet, Z., Mueller, W. A., J . Am. Chem. SOC.84, 297 (1962). Bjerrum, N., 2. Phys. Chem. (Leipzig) 108, 82 (1924). Brdnsted, J. M., 2. Phys. Chem. (Leipzig) 102, 169 (1922). Cram, D. J., Rukborn, B., Kingsbury, C. A., Haberfield, P., J . Am. Chem. SOC. 83, 3684 (1961). Eckert, C. A., Ind. Eng. Chem. 59 (9), 20 (1967). Eckert, C. A., Boudart, M., Chem. Eng. Sci. 18, 144 (1963). Frisch, H . L., Bak, T. A., Webster, E . R., J . Phys. Chem. 66, 2101 (1962). Frost, A. A., Pearson, R. G., “Kinetics and Mechanism,” 2nd ed., Wiley, New York, 1961. Glasstone, S., Laidler, K. J., Eyring, H., “Theory of Rate Processes,” McGraw-Hill, New York, 1941. Grieger, R. A., personal communication, 1969. Grieger, R. A., Eckert, C. A., A.1.Ch.E. J . , in press, 1969. Grieger, R. A., Eckert, C. A., IND.ENG. CHEM.PROCESS DESIGNDEVELOP. 6, 250 (1967). Harkness, J. B., Kistiakowsky, G. B., Mears, W. H., J . Chem. Phys. 5, 682 (1937). Harris, H. G., Prausnitz, J. M., Ind. Eng. Chern., Fundarnentals 8, 180 (1969). Hartmann, H., Kelm, H., Rinck, G., 2. Phys. Chern. (Frankfurt) 44, 335 (1965). Hildebrand, J. H., Scott, R. L., “Regular Solutions,” Prentice-Hall, Englewood Cliffs, Fi.J., 1962.
Hildebrand, J. H., Scott, R. L., “Solubility of Nonelectrolytes,” 3rd ed., Dover, New York, 1964. Kondo, Y., Tojima, H., Tokura, N., Bull. Chem. Soc Japan 40, 1408 (1967). Kondo, Y., Tokura, N., Bull. Chem. SOC. Japan 37, 1148 (1964). Kondo, Y., Tokura, N., Bull. Chem. Soc. Japan 40, 1433, 1438 (1967). Kondratev, V. N., “Chemical Kinetics of Gas Reactions,” Pergamon Press, New York, 1964. Marcus, R. A., J . Chern. Phys. 46, 959 (1967). Mills, T. R., Eckert, C. A., Ind. Eng. Chem. Fundamentals 7, 327 (1968). Raistrick, B., Sapiro, R. H., Newitt, D. M., J . Chem. SOC.1939, 1760.‘ Schmid, H., Kubassa, F., Herdy, R., Monatsh. Chem. 79, 430 (1948). Shuikin, N. I., Naryshkina, T. I., Zh. Fiz. Khirn. 31, 493 (1957). Snyder, R. B., M.S. thesis, University of Illinois, Urbana, Illinois, 1968. Stefani, A. P., J . Am. Chern. SOC.90, 1694 (1968). Szmant, H. H., Roman, M. N., J . Am. Chem. SOC.88, 4034 (1966). Wassermann, A,, Monatsh. Chem. 83, 543 (1952). Wells, P. R., Chem. Rev. 63, 171 (1963). Wong, K. F., M.S. thesis, University of Illinois, Urbana, Ill., 1967. RECEIVED for review Pu’ovember 8, 1968 ACCEPTED April 14, 1969 155th Meeting, ACS, San Francisco, Calif., April 1968. Work supported financially by the National Science Foundation.
EFFECTIVENESS OF A FLUIDIZED BED IN REMOVING SUBMICRON PARTICULATE FROM A N AIR STREAM C H A R L E S
H .
B L A C K ’ A N D
R I C H A R D
W .
B O U B E L
Oregon State University, Coruallis, Ore. 97331
ONE of the difficulties associated with any filtration system is removal of the collected material from the filter without shutting down the operation. Cyclones and scrubbers simply drain the collected material from the bottom of the system. However, most filtration systems require periodic regeneration of the filter media; this usually necessitates taking that part of the filtration system “off the line.” If the elements of a filter can be moved readily, an arrangement becomes possible for cycling them continuously through a dust-laden stream and a cleaning system. This may make it possible to use certain materials, having desirable properties, which would not otherwise be feasible. The fluidized bed offers an opportunity to maintain moving elements in proper position with respect to each other and the dust stream, thus producing a suitable filter medium and the opportunity to regenerate the filter on a continuous basis. The purpose of this study was to investigate the factors contributing to removal
’ Present address, Sorthern Arizona University, Flagstaff, Arizona 86001
efficiencies of small-diameter aerosols in a bed of fluidized glass shot. Experimental Program
The effectiveness of a 2-inch diameter fluidized bed in removing airborne particulate from an air stream was investigated a t superficial gas velocities of 8.75 t o 25.0 feet per minute. Higher velocities resulted in excessive bed carryover. Bed height-to-diameter ratios were varied from 2 to 6. Concentrations of aerosol ranged from 0.03 to 8.3 mg. per cubic meter. Ambient temperature conditions prevailed, normally 20” to 30”C. The apparatus used is shown in Figures 1 and 2. Room air, after passing through a filter, entered the aerosol generating flask, where sublimated ammonium chloride particles were picked up and carried to the stirred settling chamber. The chamber entrance valve was shut off and the vacuum pump pulled air from the stirred settling chamber through the fluidized column and a SinclairPhoenix Model J M 2000 photometer. The photometer was calibrated before and after each series of measurements in accordance with manufacturer’s instructions. The fourway reversing valve allowed the column to be shunted VOL. 8 NO. 4 OCTOBER 1969
573
