prediction of azeotrope formation based on infrared spectral data in

PREDICTION OF AZEOTROPE FORMATION BASED ON INFRARED SPECTRAL DATA IN BINARY SOLUTIONS CONTAINING METHANOL K0 IC H I

R0

, Department ofzndustrial Chemistry, Kyoto University, Kyoto, Japan N D H I D E K 0 S H I R A I , Department of Industrial Chemistry, Shinshu

NA KA N IS H I

S H0G0 I C H I N0S E A

University, Nagano, Japan

A correlation between the activity coefficient of methanol in binary mixtures and the shift of OH-stretching infrared spectra of methanol in these liquids is presented. Based on this correlation, a boiling pointinfrared shift diagram has been prepared, which can predict the formation or absence of azeotropes in binary solutions containing methanol at atmospheric pressure.

data are important for distillation processes involving multicomponent systems. Much effort has been devoted to studies on the thermodynamic aspects of azeotropic and extractive distillations. Malesinski (1965) and the Polish school have recently developed a statistical thermodynamic treatment of azeotropy based primarily on the regular solution theory (Hildebrand and Scott, 1962). Extensive information as to whether or not a given binary or multicomponent system forms a n azeotrope are available from Horsley's compilations (1952, 1962). Horsley (1952) has also presented a useful graphical method for predicting azeotrope formation range in binary systems of particular types such as alcohol-hydrocarbon and alcohol-ester. However, no general method for the prediction of azeotrope formation for a given binary system is avail able. The condition for azeotrope formation is determined first by the boiling point difference between the components and secondly by the degree of deviation from the ideal solution law-for example, the boiling point of ethyl acetate (77.1' C.) is very close to that of ethanol (78.3' C.). T h e former forms a n azeotrope with methanol a t atmospheric pressure, while ethanol-methanol is a nonazeotropic system. This difference comes from the fact that the ethyl acetate-methanol system shows a positive deviation from the ideal solution law and that the mixtures of aliphatic alcohols are almost ideal. I t is thus concluded that the boiling point difference between the two components, which is necessary to avoid azeotrope formation, becomes larger with deviation from the ideal solution law in binary systems. Figure 1 illustrates this situation. Recently, we have studied the vapor-liquid equilibrium of binary systems containing methanol (Sakanishi, 1965 ; Nakanishi et al., 1967a, c) and pointed out that the activity coefficient of various organic liquids in excess of methanol has a close connection with the infrared spectroscopic data (Nakanishi, 1965). If this correlation is valid for a large variety of binary methanol solutions, spectroscopic information, such as the hydrogen bond shift, A,,, can be used as a measure of the deviation from the ideal solution law. I n this paper, this correlation is discussed in detail and, based on it, a method for predicting azeotrope formation in binary systems containing methanol or other aliphatic alcohols is presented.

Downloaded by STOCKHOLM UNIV on September 1, 2015 | http://pubs.acs.org Publication Date: August 1, 1968 | doi: 10.1021/i160027a006

A

ZEOTROPIC

Activity Coefficient Correlation

Mixtures of alcohols with nonpolar organic liquids, such as the aliphatic hydrocarbons, show large positive deviations from the ideal solution law, while mixtures of two alcohols are nearly

ideal. Moreover, recent studies by Copp and Everett (1953), Copp and Findlay (1960), and Nakanishi et al. ( 1 9 6 7 ~ )indicate that some aliphatic amine-alcohol solutions show a negative deviation from Raoult's law.

v Phase Separation I

0 MOLE FRACTION

Figure 1. Schematic diagram of relation between azeotrope formation and nonideality of binary solutions

These facts may be interpreted qualitatively as follows: T h e alcohol molecules tend to associate and form polymer-like networks in their liquid states. I n such associated clusters, each molecule is connected by hydrogen-bonding bridges where the hydroxyl group in the alcohol molecule acts as both proton donor and proton acceptor. Since the proton-accepting ability of the hydroxyl group is very strong, the hydrogen bonds between alcohol molecules cannot be broken when they are mixed with hydrocarbons or similar compounds of weak proton-accepting ability. Ample spectroscopic evidence has indicated that the methanol molecules are mostly present as the associated species even in a very dilute region. I n such a case, the excess entropy has a large negative value, and the excess free energy is positive. Phase separation occurs when such a n effect is marked, VOL. 7

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Table 1.

Vapor-liquid Equilibrium and Infrared Spectral Data for Binary Methanol Solutions

Component Methanol Vapor Liquid Cyclohexane Benzene Toluene Chloroform Carbon tetrachloride Chlorobenzene Isopropyl ether p-Dioxane Anisole Acetone Ethyl acetate n-Butyraldehyde Acetic acid Nitromethane N,N’-Dimethylacetamide Acetonitrile Pyridine n-Butylamine Diethylamine Triethylamine

Vapor-Liquid Equilibrium Dataa log Y1° log 7 2 0

0.88 0.72

0.80 0.92

0.85 0.56

0.89 0.62

0.70

0.81

0.45

0.43 -0.05c

0.55

0.62

-0.03 -0.10 -0.09 +0,37

0 . 45c +O ,03 -0.48 -0.44 +O .32

Infrared Spectral Data VOHi

cm .--I

A v ~

-’

cm .

3685t 3365t 3647 3620 3614 3644 3643

0 320 38 65 71 41 42

3448 3504 3565 3524 3550

232 181 120 161 135

3364 3614

321 71

3427 3552 3322

258 133 363

3226

459

a Data taken from Nakanishi et al. (7967a, e, 7968) and Shirai and Nakanishi (7965). 7 refers to methanol. * Data taken from Cole et al. (7965). Nakanishi et al. ( 79676).

as in the case of methanol-n-hexane a t room temperature. A different situation occurs when two kinds of alcohols are mixed. I n this case, the hydroxyl groups of other kinds of alcohols act as proton acceptors of similar strength, and therefore, the two components mix ideally and only a very small amount of heat evolves. Thus the thermodynamic properties of alcohol mixtures are nearly ideal. The hydrogen bonding OH-N is comparable with or even stronger than that of OH-0 in alcohols. I n some cases, especially in the presence of oxygen, alcohol will react with amine. Such a strong interaction results in a negative deviation from the ideal solution law. From the above considerations, one can easily understand the situation when the alcohols are mixed with other organic liquids between these two extreme cases. The proton-accepting ability of a large majority of the functional groups in these molecules is stronger than that of the aliphatic hydrocarbons, but evidently weaker than that of alcohol itself. Thus these molecules have the ability to dissociate the hydrogen-bonded clusters incompletely. A relatively small positive deviation from the ideal solution law should then be expected. The ester-alcohol, or ether-alcohol, mixture is an example. We now t r y . to find a possible quantitative correlation between any parameter expressing the degree of the deviation from the ideal solution law and the strength of a hydrogen bond in these systems. The hydrogen bond energy between OH-0 in alcohol is estimated to be 5 to 8 kcal. per mole and those for alcohol and various proton acceptors decrease with decreasing proton-accepting ability. But the hydrogen bond energy will still be larger than the London dispersion energy, as in the hydrocarbon group. Thus the contribution of the hydrogen bond interaction to the total intermolecular force is predominant, and the degree of the hydrogen bond strength should be primarily responsible for determining the magnitude of the deviation from the ideal solution law, if the size and shape of the two components do not differ appreciably. T h e concept 382

I&EC FUNDAMENTALS

of “group interaction” in solution proposed by Redlich (1959) and developed by Shell’s group (Black et al., 1963; Pierotti et al., 1959; Wilson and Deal, 1962) is essentially the same as the above. Their treatment was mainly concerned with mixtures of hydrocarbons with hydrocarbons, alcohols, and water. Weimer and Prausnitz (1965) have also extended this idea by using the solubility parameter concept (Hildebrand and Scott, 1962). A recent comprehensive review by Pimentel and McClellan (1960) shows that the strength of the hydrogen bond is most conveniently correlated by the so-called “hydrogen bond shift” in infrared absorption spectra. Thus, it is not unreasonable to expect a correlation between infrared spectral data and the degree of deviation from the ideal solution law. Sweeny and Rose (1963) have suggested the possibility of such a correlation in their paper on the theoretical calculation of vapor-liquid equilibrium relation based on the Barker model (1952). A possible correlation of binary methanol solutions follows. As a measure of the deviation from the ideal solution law, we use log y o , the logarithm of the limiting value of the activity coefficient of liquids in a large excess of methanol, obtained from the atmospheric vapor-liquid equilibrium data. By adopting this quantity, one can consider an isothermal value at 64.5’ C. (the normal boiling point of methanol), although it is obtained by extrapolation from the isobaric data. More than 30 sets of vapor-liquid equilibrium data are available in the literature, and the authors have recently obtained additional data (Nakanishi, 1965; Nakanishi et al., 1967a, c, 1968). Table I lists the values of log y o estimated from our measurements. The quantity which is adopted as a partner of log y o in the present correlation is A,,, which is arbitrarily defined as the frequency of the OH absorption peak of methanol in the vapor phase minus the frequency of the OH absorption peak due to the complex between methanol and the acceptor. This is different from the conventional definition by which the frequency shift is measured by the deviation from the free OH absorption peak in dilute carbon tetrachloride solution. The frequency of the absorption peak in the vapor phase is taken as 3685 cm.?, according to Cole et al. (1965). Many AvS data are available in the literature and additional data have been obtained by us for 16 liquids. The experimental part of the infrared study is briefly described below. The values of AvS are also given in Table I. Before correlating log y o with A,,, the contribution of the size factor, log y s , to the activity coefficient is evaluated and subtracted from log y o , by using the Flory-Huggins equation (Flory, 1942), log ys = log n f (1

- n)/2.303

(1 1

where n is the molar volume ratio, taking methanol as monomer. Whenever possible, n was calculated from the density data a t 25’ C. The same procedure had been used by Redlich et al. (1959) and by Wilson and Deal (1962). The magnitude of this correction term depends solely on the molar volume ratio, n, and cannot be neglected if n is larger than 1.2. The uncorrected log yo’s are plotted against log y8 in Figure 2 and the corrected values ( = log y o - log ys) are plotted in Figure 3. The estimated values of log y o are subjected both to experimental errors and to uncertainty in extrapolation, while different Avs values were often reported, probably because of broad absorption of the association bands. Therefore possible error limits are indicated in Figures 2 and 3. I n view of the diversity of molecular interactions in alcoholic solutions, the correlation is satisfactory. Except for some minor cases, al-

I 4- 1.0

-

I

I

I 1 1 1

I

I

1

1

\CHLOROBENZENE -

\

$AN ISOLE

-

E HYLACETATE

-!-0.5

6 0

ACE1ON I TEK&~%E

i. A

2- B U T A ~ N I

I CHLOROFORM

ro

-

0,

0

0

TINE-

ALCOHOLS


0 I R data estimated

DIETHYLAMINE

- 0.5 -

n-BUTYLAMINE

40

60 Figure 2.

I50 1cm-I)

100

80

200

300

6

400

500

Limiting activity coefficient correlation

log y o V.I Av. log y o values taken from Table I. Amer et al. (1956), Ballard and van Winkle (1952), Bittrich and Fleircher (1963), Burke et al. ( 1 964), Hill and van Winkle ( 1 9521, Hipkin and Myers (19541, Nagata (19621, Padgitt ef al. ( 1 942), Privott et al. (1 9661, Riur et al. ( 1 959) Av8 data taken from Table 1. Baitinger et al. ( 1 964), Brandmuller and Seevogel (1 9641, Cole et 01. (19651, Findlay and Kidman (1965), Gramrtad (1962), Krueger and Mettee (1964), Mitra (1962), Ungnade et 01. (1 964)

I

1

I

1

1

1

I

1

1

0

40

\ \ \

METHANOL

60

80

150

100

A

Figure 3.

I

0

0

*

1

200

300

400

0

Us (cm-')

Limiting activity coefficient correlation log y o - log y e vb. A u ~

VOL. 7

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383

most all the points fall in a linear band, the center of which passes through the line log y o = 0 at Av, = 320 cm.-1 This value of Av, corresponds to the frequency shift of the OH association band in liquid methanol. (Actually this value is based on the polymeric band in CC14 solution and is dependent on the solvent used. But this dependence may be small enough to be negligible in the logarithmic plot.) Although both plots give comparable results, the corrected log y o us. AvS correlation seems to be better when one considers the larger AvS range. Most of the Av, data were obtained a t room temperature, while log y o values are referred to 64.5' C. However, the temperature dependence of the frequency of OH stretching vibration is usually very small for both monomeric and hydrogen-bonded species (Finch and Lippincott, 1957; Liddel and Becker, 1956), although their integrated intensities change appreciably.


Azeotrope Diagram

The correlation suggested in Figures 2 and 3 indicates that Av, can be used as a measure of the degree of deviation from the ideal solution law. Whether or not a given binary solution containing methanol forms an azeotrope can be predicted by combining the boiling temperature difference between the components and the Avs data. If the boiling temperature of each liquid, t, is plotted against log Avs, one can fix their positions on a t - log Av, diagram, as in Figure 4. O n Figure 4, the position of methanol is t = 64.5' C. and AvS = 320 cm.-' The positions of other liquids which form an azeotrope with methanol are given by a circle and those which do not form an azeotrope are given by a triangle. As is easily seen, there is a distinct regularity in the distribution of all the liquids for which both azeotropic and infrared information is available; one can divide the azeotropic and nonazeotropic regions clearly on the figure. A set of suggested straight lines for this division may be given by the equation,

t (" C.) = +78.2 log Avs (cm.-I) t (" C.) = -78.2 log Av, (cm.-I)

-

131.0

+ 260.0

The thermodynamic significance of this divided line is discussed in the Appendix. As far as binary methanol solutions are concerned, the prediction of azeotrope formation by the lines from Equation 2 is perfect. Although the prediction by Figure 4 and Equation 2 is satisfactory, it covers only some 30 binary systems and it is desirable to extend its scope of application. For this purpose infrared spectral data are needed. However, we might suggest the following extensions for the use of Figure 4. Homologous Series Correlation. T h e value of AI), is characteristic of the kind of proton-accepting group and almost independent of the size and shape of the hydrocarbon residue in the molecule. O n e can easily verify this from the data on nitriles (Mitra, 1962), nitroalkanes (Baitinger et al., 1964), and halides (Krueger and Mettee, 1964). Then one can safely assume that the AvS value is approximately constant in a homologous series. According to this assumption, any suggested line for a particular compound group is perpendicular in Figure 4. As an example, the line for saturated aliphatic monoesters, RCOOR', where R and R ' are paraffin hydrocarbon residues, is given as a dashed line on Figure 4. This is based on the value of Au, = 135 cm.-', for ethyl acetate. Using this line and the boiling point data, one can fix the position for such esters. The azeotropic data with methanol are available (Horsley, 1952) and the prediction by the lines of 384

I&EC FUNDAMENTALS

Equation 2 is correct except for one among 14 esters. Similar prediction will be possible for other homologs. Azeotropic Temperature. Since no quantitative information as to the composition dependence of the activity coefficients is deduced from the present correlation, it is difficult to predict the azeotropic composition and temperature on Figure 4. However, a n empirical estimation can be made by utilizing Horsley's figure (Horsley, 1952). According to Horsley, the boiling point difference between low boiling component and azeotrope, 6, may be roughly proportional to the absolute value of the difference in the boiling points of two components, IAI. (Strictly speaking, the 6 - / A / relation has a small curvature, but it can be neglected in this treatment.) The maximum value of 6 for given combinations of methanol with a series of homologous compounds is nearly equal to /Almax/16, where lAImsx is the temperature range in which an azeotrope is formed. Then the 6 value for any azeotropic systems of methanol binaries can be given by

Since (AImax for any liquids can be read on Figure 4, one can easily calculate 6 and the azeotropic temperature. The accuracy of this empirical calculation was examined for 22 azeotropes on Figure 4. The average error was =t1.bo C. Azeotropic Composition. A similar argument can be applied to the estimation of the azeotropic composition, c, in weight per cent of methanol. Again, according to Horsley (1952), the c us. IAI relation seems to have a general form. However, the position of the azeotrope on the boiling point curve is fairly sensitive to the excess properties of binary methanol solutions; no simple relationship can be expected. Generally speaking, c decreases as the boiling point of the second component decreases a t fixed IAL(mar and its average value at A = 0 is about 20 weight %. To get a rough estimate of the value of c, one can use this empirical relation with the result subject to an average error of ='=loweight %. O t h e r Alcohols. A survey of the literature indicates that Avs data are available for the d-methanol, phenol, pyrrole, or thioalcohol system but, except for 2-methyl-2-propanol (Cole et al., 1965), no extensive data have been published for aliphatic alcohols other than methanol and not enough data are available on the Av, of other alcohols. According to Cole et al. (1965), the absorption peak due to the 2-methyl-2-propanol-solvent interaction shifts slightly to a lower frequency in comparison with the corresponding ones due to methanolsolvent and the Av, for liquid 2-methyl-2-propanol (the OH absorption peak in the vapor phase taken as reference) is smaller than that for methanol. These facts prevent a straightforward application of Figure 4 to other alcohol systems. However, if we assume that the position of all the liquids remains unchanged and that the Av8 value of all the aliphatic alcohols is 320 crn.?, one can use the same procedure as for methanol by a translational migration of the common point of the lines given by Equation 2 to each position for other alcohols. As an example, suggested lines for ethanol are given as chain-dotted lines in Figure 4. The prediction by these lines was tested for 26 ethanol binaries for which the azeotropic data are available (Horsley, 1952, 1962). Only one case among 26 was in error. A similar test was attempted for propanols and butanols. Although the results were satisfactory, better agreement seemed to be obtained when the slope of the divided lines was somewhat smaller than that given by Equation 2. A brief discussion on this point is given in the Appendix.

I40

I20

A HAC A 0 N

IO0


n

c)

e

80

t

60

4c

A VS (ern-') Figure 4.

Diagram for prediction of azeotrope formation in binary methanol solutions

Discussion

I n spite of the complicated molecular interactions in alcoholic solutions, the correlation between log y o and AvS is promising. Before accepting this as a general rule, however, careful consideration should be given to the experimental limitations, physical meaning of both quantities, and so forth. A detailed discussion will be postponed until more experimental evidence is obtained. AvS is not a true measure of the basicity of each acceptor group. I t seems better to use other quantities, such as the equilibrium constant or free energy of complex formation between methanol and liquid. But many troublesome complications occur, owing to the presence of many different species in the systems. The absorption band owing to methanolsolvent interaction widens with increasing AvS. Moreover, the complex band sometimes has two or more peaks. No detailed experimental analysis has been made on the solvent effects with bifunctional molecules. Even where AvS can be determined precisely, it is sometimes dependent on the acceptor

concentration in CCla. There is no satisfactory theoretical explanation for this phenomenon. There are some experimental limitations in the infrared measurements-for example, the interaction between alcohols and water cannot be studied extensively by the infrared technique. I n the case of amines, the N H absorption bands overlap those of the hydrogen-bonded O H band. The latter obstacle may be eliminated by measuring the hydrogen bond shift in deuteromethanol, MeOD (Gordy, 1941 ; Searles and Tamres, 1951), and by utilizing Bellamy’s rule (Bellamy et al., 1958). I n view of these limitations, it will be necessary eventually to use other data such as the N M R chemical shift as a measure of the hydrogen bond strength. In correlating the limiting activity coefficients with AvS, we have introduced a correction for an entropy contribution owing to the difference in size of the molecules. However, there is no assurance that the value calculated by Equation 1 gives an accurate correction. I n fact, Redlich (1959) reported that this correction is too small to account for the entropy contribuVOL. 7

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385

tion in the mixtures of nonpolar liquids. There should be some other entropy contribution probably owing to the mutual orientation of molecules; it will be fairly large in the hydrogen bond systems. Thus y o should consist of the following main contributions:

+

AGe = xixz[A

+


contribution owing to the difference in size = Equation 1)

(4)

Other contributions may be small in comparison with the above three. Triethylamine, isopropyl ether, and substituted benzene derivatives deviate appreciably from the present correlation. The presence of a compact and bulky hydrocarbon residue adjacent to the proton-accepting group seems to be responsible for this deviation. The present correlation intends to connect two quantities at different temperatures-Le., AvS at room temperature and log y o a t 64.5' C. To correct for this discrepancy and evaluate each term in Equation 4 separately, it is desirable to use heat of mixing data. The authors are currently studying this problem, with the hope of obtaining a correlation more accurate than the present one. Experimental

The hydrogen bond shifts of OH frequency of methanol in 15 organic liquids were measured in CC14 solution and are listed in Table I. They were obtained in the following manner. Solvents. T h e sample liquids used for infrared spectroscopic measurement were purified by the procedure used for the recent vapor-liquid equilibrium study (Nakanishi et al., 1967a, c). Precautions were taken to prevent moisture contamination of carbon tetrachloride and other liquids. Apparatus. T h e instrument used to obtain the infrared spectra was a Nihon Bunko DS-301 spectrophotometer with NaCl optics. T h e wave number was reproducible within h 1 . 5 crn.-' The cells used were matched pairs, of silica, having caps and a thickness of 10 mm. They were transparent in the wave number ranges between 3000 and 4000 cm . Infrared Measurement. All the sample solutions were prepared by weighing into a narrow-necked volumetric flask and making successive dilutions with pipets. They always contained 0.005 mole of methanol in 1 liter of carbon tetrachloride, to which three different amounts of solvent were added. Some solvents absorb near the absorption bands of methanol, and it was necessary to keep their concentration exactly the same in the reference and sample solutions. T h e absorption was recorded with a scanning speed of 10 cm.-' per minute. For the determination of absorbance, each solution was run two or three times under the same conditions. The listed values represent the average. T h e temperature of the sample solution was 23' f 0.5' C. during the measurement. Appendix

Statistical thermodynamic treatment of the azeotropic phenomena is given by Prigogine and Defay (1954), Malesinski (1965), and Yoshimoto (1961). Yoshitnoto has shown that the condition for azeotrope formation is expressed by the following equation,

- (A + B)/ASzO < Ti' - Tz' < + ( A

- B)/ASiO

(AI)

where T o is the boiling point of the pure component, ASo is the entropy of vaporization of the pure component, and A and B are the coefficients of the expression for the excess free energy of mixing, AGe. 386

l&EC FUNDAMENTALS

XZ)

XZ)'

...]

(A2)

If the solution is regular, B is equal to zero and

A

=

2.303 R T l o g y '

('43)

The correlation shown in Figures 2 and 3 may be written as

(energy contribution owing to the hydrogen bond) (entropy contribution owing to orientation) (entropy

yo =

+ B ( X I- + C(XI -

log 72' = a log Av8

+b

('44)

where a and b are constants and su& 1 refers to methanol and 2 to any proton-accepting molecule. Now we can select the ordinate of Avs to make the value of b zero. Then, by neglecting the difference between AS1' and ASz', Equation A1 can be rewritten as

- (2.303 a R T log AvS)/ASo < T ' M ~ O-H Tzo < (2.303 a R T log Av,)/AS'

(A5)

Therefore the equation for the divided lines on the azeotrope diagram is given by

\AT1 = K log Av8 where A T =

- Tz'

TohlleO~

(Ab)

and K = (2.303 aRToMeoa)/ASo.

Conclusions

The present divided lines are based on the assumptions of linear relationship between log yz' and log AvS, a regular solution, and invariance of the entropy of vaporization for all liquids. If the same argument can be extended to other alcohols, the slope of the divided lines is primarily dependent on the value of a, provided that the Avs values are common to all alcohols. Although the available log y o data for other alcohol systems are limited, it is natural to expect the log y o of nonpolar liquids in higher alcohols always to be smaller than those in methanol or ethanol, so that the value of a for these systems becomes smaller. It is obvious that this reflects on the decrease in the slope of the divided lines. Acknowledgment

The authors express their appreciation to Nobuatsu Watanabe and Hidekazu Touhara for their interest and discussions. Yoshinari Kawasaki assisted in some of the experiments. The experimental section of this paper is based in part on a thesis by Shogo Ichinose for the B. Eng. degree, Shinshu University, 1966. Nomenclature

coefficients in AGEequation constants in Equation A4 azeotropic composition, wt. yo of methanol excess free energy of mixing constant in Equation A6 molar volume ratio gas constant entropy of vaporization of liquid boiling temperature, ' C. boiling point of liquid, ' K. mole fraction limiting value of activity coefficient of liquid at infinite dilution of methanol entropy contribution to y owing to difference in size boiling point difference between two components temperature range of azeotrope formation boiling point difference between low boiling component and azeotrope frequency shift of infrared OH stretching vibration absorption, crn.-'

literature Cited

Amer, H. H., Paxton, R. R., van Winkle, M., Znd. Eng. Chem. Data Ser. 3, 224 (1956). Baitinper., W. F.. - Schlever. P. von R.. Murtv. T. S. S. R.. Robinson,-L., Tetrahbdron 26, 1635 (1964).’ Ballard, L. H., van Winkle, M., Ind. Eng. Chem. 44,208 (1952). Barker, J. A., J.Chem. Phys. 20, 1526 (1952). Bellamy, L. J., Hallam, H. E., Williams, R. L., Trans. Faraday SOC.54, 1120 (1958). Bittrich, H. J., Fleischer, W., J.Prakt. Chem. 4, 151 (1963). Black. C.. Derr. E. L.. Pauadououlos. M. N.. 2nd. Enp. Chem. 55. No: 8, 40, NO. 9, 38’(1963). Brandmuller, J., Seevogel, K., Spectrochim. Acta 20, 453 (1964). Burke, D. E., Williams, G. C., Plank, C. A., J . Chem. Eng. Data 9, 212 (1964). Cole, A. R. H., Little, L. H., Michell, A. J., Spectrochim. Acta 21, 1169 (1965). Copp, J. L., Everett, D. H., Discussions Faraday Sac. 15, 174 (1953). 56, 13 (1960). Copp, J. L., Findlay, T. J. V., Trans. Faraday SOC. Finch, J. N., Lippincott, E. R., J . Phys. Chem. 61,894 (1957). Findlay, T. J. V., Kidman, A. D., Australian J . Chem. 18, 521 (1965). Flory, P. J., J . Chem. Phys. 10, 51 (1942). Gordy, W., J.Chem. Phys. 9,215 (1941). Gramstad, T., Acta Chim. Scand. 16, 807 (1962). Hildebrand, J. H., Scott, R. L., “Regular Solutions,” PrenticeHall, Englewood Cliffs, N. J., 1962. Hill, W. D., van Winkle, M., 2nd. Eng. Chem. 44, 205 (1952). Hipkin, H., Myers, H. S., Ind. Eng. Chem. 46,2524 (1954). Horsley, L. H., Advan. Chem. Ser., No. 6, 35 (1952, 1962). Krueger, P. J., Mettee, H. D., Can. J. Chem. 42, 288 (1964). Liddel, U., Becker, E. D., J.Chem. Phys. 25,173 (1956). Malesinski, W.,“Azeotropy and Other Theoretical Problems of Vapor-Liquid Equilibrium,” Interscience, New York, 1965. Mitra, S. S., J.Chem. Phys. 36, 3286 (1962). , I

~


-

Nagata, I., J . Chem. Eng. Data 7, 367 (1962). Nakanishi, K., Chem. Eng. ( J a p a n ) 10,1122 (1965). Nakanishi, K., Nakasato, K., Toba, R., Shirai, H., J . Chem. Eng. Data 12, 440 (1967a). Nakanishi, K., Shirai, H., Ichinose, S., Kyoto University, unuublished results. 1967b. Nakanishi, K., Sh&ai, H., Minamiyama, T., J . Chem. Eng. Data 12, 591 (1967r). Nakak h i , K., Shirai, H., Nakasato, K., J . Chem. Eng. Data 13, 181S (1968). Padg itt, F. L., Amis, E. S., Hughes, D. W., J. A m . Chem. Sac. 64, 122 . - 3 l (1942). Pierotti. G. J.. Deal. C. H., Derr, E. L., Ind. Enp. Chem. 51, 95 (1959). Pimentel, G. C., McClellan, A. L., “The Hydrogen Bond,” W. H. Freeman, San Francisco, Calif., 1960. Prigogine, I., Defay, R., “Chemical Thermodynamics,” Longmans Green, London, 1954. Privott, W. J., Paul, D. R., Jolls, K. R., Schoenborn, E. M., J.Chem. Eng. Data 11, 331 (1966). Redlich, O., Derr, E. L., Pierotti, G. J., J. A m . Chem. SOC. 81, 2283 (1959). Rius, A.; Otero, J. L., MaCarron,A., Chem. Eng. Sci. 10,105 (1959). Searles, S., Tamres, M., J.Am..Chem. SOC.73, 3704 (1951). Shirai, H., Nakanishi, K., Kagaku Kogaku ( J a p a n ) 29,180 (1965). Sweeny, R. F., Rose, A,, A . I . Ch. E. J.9, 390 (1963). Ungnade, H. E., Roberts, E. M., Kissinger, L. W.,J.Phys. Chem. 68.3225 (1964). Weimer, R’. F., ‘Prausnitz, J. M., Hydrocarbon Processing 44, 237 (1965). Wilson, G. M., Deal, C. H., IND.ENC.CHEM.FUNDAMENTALS 1, 20 (1962). Yoshimoto, T., Nippon Kagaku Zasshi 82,530 (1961). \ - . -

- I .

I

RECEIVED for review September 5, 1967 ACCEPTED February 19, 1968

REACTOR DESIGN AND CONTINUOUS SAMPLING CRITERIA FOR AN ULTRASONIC REACTION SCOTT FOGLER

Department of Chemical and Metallurgical Engineering, University of Michigan, Ann Arbor, Mich.

Ultrasonically induced cavitation can bring about increased rates in liquid-phase chemical reactions. If the variation in chemical yield is to remain at a minimum in an ultrasonic reaction, certain limits of the acoustic pressure must not be exceeded. If these limits are exceeded in some instances, the acoustic pressure will drop below the cavitation threshold and the reaction will cease. Design criteria for a continuous stirred tank reactor and batch reactor, along with the maximum sample size, are developed which meet the prescribed restrictions.

NE

of the usual techniques for obtaining rate data from a

0 chemical reaction is by sampling the reaction vessel con-

tents a t various time intervals after the reactants are mixed. By plotting the proper function of concentration against time, one can then determine the various reaction rate parameters for simple reactions. I n liquid-phase reactions where no ultrasonics are applied to the reacting species, the amount of sample withdrawn a t any time interval will not affect the remainder of the chemical reaction. However, when a n ultrasonic wave is applied to a chemical reaction or extraction, the amount of liquid sampled for concentration determination can have a

dramatic effect on the rate of a chemical reaction. This paper gives an approximate set of conditions under which continuous sampling of an ultrasonic reaction is permissible. These results are extended to continuous ultrasonic processing in the design consideration of a continuous stirred tank reactor

(CSTR) . Various investigators (Chen and Kalback, 1967; Ostroski and Stambaugh, 1950; Weissler et al., 1950) have applied ultrasonics to liquid-phase chemical reactions and withdrawn samples for titration continuously throughout the course of the reaction. Further search of recent literature on ultrasonic VOL. 7

NO. 3

AUGUST

1968

387

prediction of azeotrope formation based on infrared spectral data in

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