Solubility of Aromatic Hydrocarbon Solids in Pyridine and Thiophene

When the binary parameter, /,2, determined from the lowest solubility datum point temperature is used in the Scatchard-Hildebrand equation, the activi...
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Ind. Eng. Chem. Fundam. 1983, 22, 46-51

46

Solubility of Aromatic Hydrocarbon Solids in Pyridine and Thiophene Peter B. Chol and Edward McLaughlln' Deparfment of Chemical Engineering, Louisiana State University, Baton Rouge, Louisiana 70803

Solubilities of eight aromatic hydrocarbon solids, biphenyl, naphthalene, fluorene, phenanthrene, acenaphthene, fluoranthene, pyrene, and 0 -terphenyl, have been determined in two pure solvents, pyridine and thiophene, from around room temperature to temperatures near the melting point of the solids. The solubility data of the first five solids are correlated well by the Scatchard-Hildebrand regular solution theory if solubility parameters of the components in each binary system are evaluated at the melting temperatures of the solutes rather than at 298 K, the conventional datum temperature. In this case activity coefficients of the solutes are predicted on average to within 3%. When the binary parameter, /,2, determined from the lowest solubility datum point temperature is used in the Scatchard-Hildebrand equation, the activity coefficients are predicted to within an average of 0.6%.

Introduction The solubility of solids in liquids is expressed by eq 1

when there is no phase transition in the solid phase. When a phase transition takes place between the system temperature T and the melting point T,, the solubility equation for temperatures below that of the phase transition must include the effect of the transition. The result for a first-order transition as given, for example, by Weimer and Prausnitz (17) is

naphthalene, fluorene, phenanthrene, and acenaphthene, such physical properties are available, and accordingly solubility measurements on these five solids enable evaluation of y2 in solution with any solvent. For correlation of the activity coefficients, the Scatchard-Hildebrand regular solution model, eq 6, will be used as it has been found to be the most useful for these systems. On this model, the activity coefficient is given by

where the solubility parameter, 6, is defined as 6 = (aEV/V')'/2

(7)

Equation 6 predicts y2 from pure-component properties only. The accuracy of prediction by the Scatchard-Hildebrand regular solution theory is improved if an adjustable constant 112 is added to the equation giving

If the phase transition is of the X type, a correction term, A, given by eq 3

(8)

This equation requires one experimental datum point for evaluation of the binary parameter l12.

(4) must be added to eq 1, giving

T, + X + In y2 (5) R T where T, and Tbspecify the temperature range over which the transition takes place. CPBin eq 3 is the base line of the specific heat during the phase transition. A h P 2 and A S P z in eq 4 stand for the overall contribution of enthalpy and entropy changes caused by the specific heat anomaly from the beginning to the end of the transition. By using eq 1 and either eq 2 or 5, one can evaluate activity coefficients of the solute, y2, if solubility data ( X 2 vs. 5") are available, assuming that other physical constants in the equations are known. For the solids, biphenyl, Ac,f2

-In -

Experimental Work Biphenyl and naphthalene were purchased from Eastman Kodak, phenanthrene, fluorene, and pyrene were purchased from Eastern Chemical, and acenaphthene was purchased from British Drug H o w , o-terphenyl from ICN pharmaceutical, Inc., and fluoranthene from the Matheson Co., Inc. The samples, except acenaphthene and fluoranthene, were purified further by activated alumina chromatography using toluene as eluant. After recrystallization, the toluene was removed in vacuo. Fluorene was further purified by batch distillation to remove lowboiling impurities and acenaphthene was zone refined. The final samples were analyzed by GC mass spectrometry and melting points of the samples measured. The results were as follows: biphenyl, 99.67 wt % (342.6 K); naphthalene, 99.21 w t 5% (352.8 K); fluorene, 97.85 wt % (387.6 K); phenanthrene, 98.75 w t 70 (372.8 K); pyrene, 99.09 w t % (422.7 K); o-terphenyl, 99.63 wt % (328.7 K); fluoranthene, 98.75 wt % (383.0 K); and acenaphthene, 99.2 wt % (366.5 K). Pyridine and thiophene solvents were "gold label" quality Aldrich products and were used without further purification.

0196-4313/83/1022-0046$01.50/00 1983 American Chemical Society

Ind. Eng. Chem. Fundam., Vol. 22, No. 1, 1983

Following purification, mixtures of known composition were sealed in glass ampules and the solubilities determined by procedures similar to those of McLaughlin and Zainal(9-11). This essentially consists of determining the temperature at which the last trace of solid disappears while the temperature of the bath, in which the samples were rotated at 0.25 rps, was raised at a rate of 0.1 K every 1200 s. The solubility data taken in this manner are tabulated in column 2, Table I, for naphthalene, fluorene, phenanthrene, and acenaphthene and in Table I1 for the remaining three solids. For reasons that will be clear later biphenyllpyridine data are in Table I and biphenyl/ thiophene data in Table 11.

Correlation of Activity Coefficients Activity coefficients for biphenyl, naphthalene, fluorene, and acenaphthene were evaluated for all data points by using eq 1. For phenanthrene, eq 1was used at solution temperatures above the end of the phase transition while eq 5 was used at temperatures below the end of the phase transition. hsf, and AC,fz for the evaluation of the activity coefficients of the solutes considered are given in Table 111. To satisfy the condition X 2 1while T,/T 1, we used the observed melting temperatures of the samples instead of the literature values given in Table 111. ASf, and AC,f? of biphenyl in the table yielded negative values for In y2 in solutions with thiophene. Solubility data for biphenyl in pyridine at 323.7 and 331.7 K also showed a negative deviation when Asf, and AC,f2 given in Table I11 were used. To check this negative deviation, we reevaluated activity coefficients of biphenyl in both pyridine and thiophene with the new values of and AC,f, calculated by using the available VLE data for the biphenyl-benzene system (1, 2, 6). To calculate the new values of Asf, and AC,f2, we fitted isothermal activity coefficients of benzene in the biphenyl-benzene system, which were given by Guggenheim (6),to the Van Laar equation at four temperatures of 303, 313, 323, and 333 K and back-calculated activity coefficients of biphenyl at saturation compositions at these temperatures. These values were then fitted to eq 1 by a least-squares method to yield 53.411 X lo3 and 33.835 X lo3 J mol-' K-' for ASf, and AC,fz, respectively, which were used to calculate y2 for biphenyl in both pyridine and thiophene solutions. The new values improve the result for biphenyl in pyridine but not for biphenyl in thiophene. To confirm the negative deviation of biphenyl in thiophene solution, one needs vapor-liquid equilibrium data of the binary system. As these are not available at the present time, the system biphenyl-thiophene is excluded from the correlation work of this study. The Scatchard-Hildebrand regular solution model (eq 6) needs the physical properties of the pure components such as energy of vaporization Ai?P and liquid molar volume v' for evaluation of the solubility parameter 6. AE?' and v' are readily available for pyridine and thiophene at any particular choice of the measured solution temperatures since they exist as liquids at these conditions. To correspond, AEv and v' of the solutes would have to be found in their subcooled liquid state because they do not normally exist as liquids at the solution temperatures. Physical properties of solids at such conditions are rarely given in the literature and this causes difficulty in evaluating the solute solubility parameters. The assumption in the regular solution theory that ASE = 0, however, allows evaluation of y2 if solubility parameters 6, and are evaluated at any fixed convenient temperature at which all parameters on the right-hand side of eq 6 can be determined. The most convenient tem-

-

-

47

perature with the minimum extrapolation is obviously the melting point of the solutes, where no difficulty is encountered in finding AE" and v'. AE?' was calculated by subtracting RT from Ahv which was estimated by using the Clausius-Clapeyronequation. Antoine equations given by Reid et al. (13)and Dean (3) were extrapolated to the melting points of solutes for use in the Clausiusxlapeyron equation. Solubility parameters and liquid volumes of the solutes and solvents, evaluated at the melting point of the solute, are given in Table IV. We call this the "floating datum point method" as it contrasts with the normal choice, which involves a fixed temperature for all systems, usually 298 K. When these values were used, activity coefficients of the solutes were predicted by using eq 6 and are given in column 4 in Table I. The deviation of the values from the experimental results, defined by eq 9, is A = 1001[(Y2)exptl-

(Y2)predl /(Yl)exptll

(9)

given in column 5. Similarly, eq 8 was used to calculate activity coefficients with the binary parameter 112 evaluated at the lowest temperature at which the solubility was measured for each system. The value of y2 and the calculated deviation are given in columns 6 and 7 in Table I as well as l12.

Discussion of the Results The Scatchard-Hildebrand regular solution model (eq 6) predicted yzwith an average deviation of 2.8% for 54 data points as shown in column 4 in Table I. The result signifies that solubility parameters can be evaluated at the melting temperature of the solute, thus eliminating the difficulties in obtaining solubility parameters of the solute components in their subcooled liquid state. There is a general trend, however, in that the deviation becomes larger as the system temperature decreases. Prediction of activity coefficient by this method should therefore be made with caution at system temperatures far from the melting point of the solute. Equation 8, which includes the binary parameter 112, improves the accuracy of prediction to an average deviation of about 0.6% as illustrated in column 6 of Table I. In this case, however, one datum point is necessary to evaluate the binary parameter l12. While no prescription is available for determining Z12 from first principles, it is useful to try to correlate the present values from solubility studies. For thiophene 112 is so close to zero that nothing useful follows from the values. From the data in pyridine, however, the results conform roughly to a group contribution pattern. This can be seen if a benzene ring is assigned a value of 0.0018 from the biphenyl result and a -CH2- in a ring arrangement a value of 0.0030 from the fluorene result. With these numbers, calculated and experimental lI2 values as given in Table V show reasonable agreement. The number of systems is however too limited to have confidence in the values and additionally, as the values would only apply to these particular solvents, such predictions would have limited applicability. Generalization of Solid Solubility Data I t was shown by McLaughlin and Zainal (9-11) that their solubility data for solids of the same series of aromatic hydrocarbons could be generalized for each of the solvents benzene, carbon tetrachloride, and cyclohexane when -log X 2 was plotted as a function of T,/T, showing that each solid had the same solubility in the same solvent at the same reduced reciprocal temperature T,/ T. The solubility data in Tables I and I1 are plotted in the same fashion in Figures 1and 2 for pyridine and thiophene solutions. The solid lines in the figures, which are obtained by a least-

48

Ind. Eng. Chem. Fundam., Vol. 22, No. 1, 1983

Table I. Solubility and Activity Coefficients at Saturation of Aromatic Hydrocarbon Solids in Pyridine and Thiophene

T/K

ix1 ) expt1

iY 2)exptl

(Y 2 )pred

297.3 300.9 307.1 312.4 323.7 331.7

0.3742 0.4085 0.4743 0.5327 0.6884 0.8170

Biphenyl in Pyridine 1.054 1.013 1.011 1.042 1.030 1.008 1.027 1.006 1.005 1.002 1.000 1.000

297.6 3 25.3 333.2 337.2 344.4

0.3032 0.5748 0.6808 0.7422 0.8570

1.061 1.034 1.022 1.012 1.003

Naphthalene in Pyridine 1.006

1.001 1.001 1.000

1.000

A I%

A 1%

( Y 2 )pred

3.8 2.9 2.1 2.0 0.3 0.0

1.054 1.045 1.032 1.023 1.008 1.001 l , , = 0.0036

5.1 3.2 2.0 1.1 0.3

1.061 1.016 1.008 1.005 1.001

0.3 0.2 0.3 0.3 0.1

1.7 1.4 0.7 0.2

I,, = 0.0036 311.5 3 27.1 340.2 349.0 359.2

0.1979 0.2936 0.4002 0.4909 0.6127

1.153 1.111 1.075 1.042

1.011

Fluorene in Pyridine 1.003 1.002 1.001 1.001 1.000

13.0 9.8 6.8 3.9 1.1

1.153 1.096 1.057 1.036 1.017

1.3 1.6 0.6 0.6

I , , = 0.0066 299.8 307.7 314.3 316.6 323.4 342.8 349.6 355.6 361.0 366.5

0.2459 0.3011 0.3513 0.3690 0.4283 0.6170 0.6961 0.7651 0.8349 0.9111

Phenanthrene in Pyridine 1.087 1.004 1.057 1.003 1.002 1.042 1.040 1.002 1.025 1.002 1.001 1.020 1.012 1.000 1.012 1.000 1.007 1.000 1.002 1.000

7.6 5.1 3.8 3.6 2.2 1.8 1.1 1.2 0.7 0.2

1.087 1.067 1.052 1.047 1.035 1.012 1.007 1.004 1.002 1.000

0.9 0.9 0.6 0.9 0.7 0.5 0.7 0.4 0.2

I , , = 0.0041 306.7 320.0 332.9 337.5 343.7

0.2102 0.3166 0.4592 0.5191 0.6110

Acenaphthene in Pyridine 1.242 1.009 1.158 1.006 1.079 1.003 1.058 1.002 1.032 1.001

0.3588 0.5011 0.5510 0.7372 0.8053 0.9167

Naphthalene in Thiophene 1.023 1.014 1.022 1.007 1.021 1.005 1.011 1.002 1.000 1.004 1.003 1.000

18.8 13.1 7.0 5.3 3.0

1.242 1.145 1.072 1.052 1.030

1.1 0.6 0.6 0.2

I , , = 0.0101 303.2 318.2 322.7 336.8 341.1 348.1

0.8 1.4 1.5 0.9 0.4 0.3

1.023 1.012 1.009 1.003 1.001 1.000

I,, 303.6 321.0 335.5 350.2

357.5

299.4 304.9 310.0 321.2 340.9 348.3 355.0

Fluorene in Thiophene 1.052 1.034 1.020

0.1844 0.2762 0.3802 0.5146 0.5954

1.017 1.031 1.027 1.017 1.008

0.2379 0.2742 0.3146 0.4016 0.5907 0.6757 0.7541

Phenanthrene in Thiophene 1.114 1.112 1.092 1.094 1.063 1.077 1.046 1.050 1.035 1.017 1.022 1.009 1.018 1.005

1.010

1.006

3.4 0.3 0.6 0.6 0.2

0.4 0.2 1.3 0.4 1.7 1.2 1.2

0.9 1.1 0.7 0.3 0.3 =

0.0008

1.017 2.0 1.9 1.4 0.6

1.011 1.007 1.003 1.002 I,,

-0.0016

1.114 1.095 1.078 1.050 1.017 1.010 1.005 I , , = 0.0001

0.2 1.4 0.4 1.7 1.1 1.2

Ind. Eng. Chem. Fundam., Vol. 22,

No. 1, 1983 49

Table I (Continued eq 6

eq 8 -_ ( Y 2 )pred A /% ~-

I _ _

( X ?Iexptl

T/K

--

(Y ? Ipred A /% Acenaphthene in Thiophene 1.024 1.021 0.3 1.021 1.014 0.6 1.014 1.009 0.5 1.009 1.006 0.3 1.002 1.002 0.0

(7z)exptl

I _

0.2584 0.3542 0.4407 0.5224 0.6428

307.3 319.6 328.4 335.7 344.8

1.024 1.016 1.010 1.007 1.003 l , , = 0.0002

0.5 0.4 0.2 0.1

Table 11. Solubility of Fluoranthene, Pyrene, and o-Terphenyl in Pyridine and Thiophene and Biphenyl in Thiophene _ _ I _

T /K

fluoranthene

x,

pyrene

T/K

o-terphenyl

X? T/K

X? T/K

fluoranthene

x2

T/K

pyrene

X? T/K

o-terphenyl

x2

T/K

biphenyl

Table 111. -

x,

In Pyridine 328.4 340.1 0.3537 0.4545 326.7 340.5 0.1899 0.2561 303.6 308.5 0.6061 0.6753

354.1 0.6071 355.3 0.3474 317.9 0.8140

360.0 0.6729 365.8 0.4201 320.8 0.8618

299.5 0.1754 307.0 0.1217 298.6 0.5745 295.9 0.3851

In Thiophene 316.5 326.7 0.2652 0.3354 323.7 335.9 0.1764 0.2323 303.4 311.4 0.6296 0.7300 298.4 309.7 0.4071 0.5218

342.5 0.4752 350.2 0.3123 317.6 0.6158 317.0 0.6060

357.1 0.6355 362.5 0.3934 321.6 0.8758 329.5 0.7653

Physical Properties of Five Aromatic Hydrocarbon Solids

1 0 - ~ ~ h f , / ( ~ AS~J(J mol-') mol-' K-' I

Tm/K

Reference 1 4 .

341.3a 353.3b 387.9' 372.4c 366.6'

Reference 15.

18.659' 18.23Sb 19.591' 16,474' 21.476'

' Reference 5.

,T ,

334.8 0.8698

--

--

I _ -

biphenyl naphthalene fluorene phenanthrene acenaphthene a

309.2 0.2282 304.3 0.1119 302.0 0.5880

54.671 51.623 50.493 44.254 58.573

AC f l ~ ( ~

mot1

~ O - ~ A ~ P , ~ ASP,'/(J / ( J

K-11

36.274' 8.901 1.444' 12.586' 14.855'

mol-' I

mol-' K-I 1

1.307

3.835

Ahf?, A S f ? , and A C D f 2 are values a t the triple point.

Table IV. Solubility Parameters and Molar Liquid Volumes of Solids and Liquids a t the Melting Points of the Solids __-----___solute pyridine thiophene T,(obsd)/K biphenyl naphthalene phenanthrene fluorene acenaphthene

342.6 352.8 372.8 387.6 366.5

10-462 / (J m-3)1'2

1.9304' 1.9396a 1.977 2' 1.850Sg 1.8930g

io6v1?/

( m 3 mol-')

155.16b 130.86e 168.0Ejf 163.7 149.8'

__ IO-~S~

~ o - ~bv/ ' ,

(J m-3)1'2 (m3mol-')

2.0259 1.9938 1.9377 1.8841 1.9519

84.87 85.86 87.31 89.36 87.19

--

i~-~b,ci

io6vlId/

( J m-3)1'z ( m 3mol-')

1.8759 1.8426 1.7759 1.7259 1.7974

83.78 84.87 87.14 88.96 86.38

a Enthalpy of vaporization was evaluated by using the Clausius-Clapeyron equation with Antoine equation constants given by Reid e t al. ( 1 3 ) . Reference 15. ' Enthalpy of vaporization data given by Dreisbach ( 4 ) . Reference 7u. Antoine equation constants e References 3, 13, and 1 6 . f ReFeren.ce 7b. Extrapolated by using eq 12-3.2 of ref 1 3 . given by Dean ( 3 ) . Reference 1 2 . Reference 8.

squares fit of the data, subject to passing through the point (O,l), can be expressed approximately by eq 10, which is -log

xz = -

(lo)

obtained from eq 1 by dropping the smaller terms. The

magnitude of AS' in eq 10 can be used to compare nonideality in the different solvents. The line for the pyridine solutions in Figure 1, for example, yields 50.803 J mol-' K-' for AS', while the line for the thiophene solutions in Figure 2 yields 49.819 J mol-' K-l. McLaughlin and Zainal (9-11) obtained 57.778,66.570, and 101.739 J mol-' K-l for AS' in solutions with benzene, carbon tetrachloride, and

50

Ind. Eng.

Chem. Fundam., Vol.

22, No. 1, 1983

Table V. Comparison of Calculated and Experimental Values of I , , for Pyridine Solutions _I____._____

Phenanthrene A Fluorene 0 Acenaphthene Fluoranthene v 0 - Terphenyl v Pyrene

I _ _ _ -

biphenyl

naphthalene fluorene

a

o Biphenyl A Naphthalene

01

calcd

exptl

0.0

phenanthrene acenaph thene Fitted values.

0.0036 0.0036 0.0066

0.0036a 0.0036 0.0066a

0.2

0.0041

0.0054

0.3

0.0101

0.0096 04

00

o Biphenyl

a Naphthalene Phenanthrene A Fluorene

o Acenaphthene m Fluoranthene D 0-Terphenyl ' I Pyrene

0 2

RJ

X

04

0.9

0 5

I

11

0.9

10

\

\

1.01

I

I

\I-

12

14

~

I O

.o

1.2

0.8 0'7

I I

1.2

\

0.8

I .3

( T m ) obs. T

Figure 1. Generalization of solid solubilities in pyridine.

cyclohexane, respectively. Therefore, it can be concluded that nonideality increases in sequential order in thiophene, pyridine, benzene, carbon tetrachloride, and cyclohexane. It can also be noticed that the differences in nonideality between solutions in pyridine and in thiophene are rather small, even though pyridine is a six-membered ring and thiophene a five-membered ring. The solvents containing a hetero atom such as N or S in their molecules form solutions conforming more closely to ideal behavior than the solvents containing only homogeneous atoms such as benzene when they dissolve aromatic hydrocarbon solids. Generalization of the solid solubility data in each solvent is due to two facts. First, the solids studied show similar entropy changes of fusion, ranging from 44 to 59 J mol-l K-l. Secondly, they have approximately the same order of magnitude of A6, the difference in solubility parameters 61and 62. As a result, solubility of a solid either in pyridine or in thiophene can be approximately predicted from Figure 1or 2 if Asf, and 6, of the new solid are similar to those of the solids studied. Conclusions The solubility study presented in this paper suggests the following results. 1. When no experimental data are available, the Scatchard-Hildebrand regular solution model (eq 6) may be safely used for the first approximation to activity coefficientswith the solubility parameters evaluated at the melting points of the solutes.

1.2 I .3 I.4 ( T m ) obs. T Figure 2. Generalization of solid solubilities in thiophene. 1.1

2. When one experimental datum point is available, the regular solution model with the binary parameter 112 (eq 8) may be used for improved accuracy of prediction. In this case, however, the data a t a lower temperature are preferred for evaluation of the binary parameters. 3. Solubility of solids in pyridine and thiophene can be generalized for each solvent as in the case of benzene, carbon tetrachloride, and cyclohexane. The generalization shows that nonideality of the aromatic hydrocarbon solids increases in sequential order in thiophene, pyridine, benzene, carbon tetrachloride, and cyclohexane solutions.

Nomenclature A C i 2 = difference of molar specific heat between those of

liquid and solid at fusion temperature

aEv = molar energy of vaporization

AhPz = molar enthalpy change of phase transition Ahv = molar enthalpy of vaporization lI2 = binary parameter R = gas constant ASE = excess entropy of mixing ASf, = molar entropy change of fusion AS12 = least-squares-fitted value ASPz = molar entropy change of phase transition T = system temperature T, = melting temperature P = temperature of first-order phase transition v' = molar liquid volume X = mole fraction Greek Letters y = activity coefficient 6 = solubility parameter

Subscripts 1 = solvent

2 = solute (solid) Registry No. Biphenyl, 92-52-4; naphthalene, 91-20-3; fluorene, 86-73-7; phenanthrene, 85-01-8;acenaphthene, 83-32-9;fluoranthene, 206-44-0; ppene, 129-00-0;o-terphenyl,84-15-1; pyridine, 110-86-1; thiophene, 110-02-1.

Ind. Eng. Chem. Fundam. 1983, 22, 51-54

Literature Cited

51

(9) McLaughlin, E.; Zainai, H. A. J . Chem. SOC. 1950, 863.

(1) Baxendale, J. H.; Eniistun, B. V.; Stern, J. fhilos. Trans. R . SOC. London, Ser. A 1051, 243, 169. (2) Everett, D. H.; Penny, M. F. R o c . R . SOC. London, Ser. A 1952, 212, 164. (3) Dean, A. "Lange's Handbook of Chemistry", 12th ed.; McGraw-Hill: New York, 1979; pp 10-37, 10-113. (4) Dreisbach, R. R. "Physical Properties of Chemical Compounds"; Amerlcan Chemlcal Society: Washington, DC, 1955; A&. Chem. Ser. No. 15. (5) Finke. H. L.; Messerly, J. F.; Lee, S. H.; Osborn, A. G.; Douslin, D. R. J. Chem. Thermodyn. 1977, 9, 937. (6) Guggenhelm, E. A. "Mixtures"; Oxford at the Clarendon Press: London, 1952; p 236. (7) (a) "Internatlonal Critical Tables"; National Research Councll, McGraw-Hill: New York, 1926; VoI. 3, p 28. (b) IbM. Vol. 7, p 88. (8) Lange, N. A.; Forker, G. M. "Handbook of Chemistry", 10th ed.; McGraw-Hill: New York, 1961; pp 402, 1289.

(IO) McLaughlin, E.; Zainal, H. A. J . Chem. SOC. 1960, 2485. (11) McLaughiin, E.; Zainal, H. A. J . Chem. SOC. 1060, 3854. (12) McLaughlin, E.; Ubbelohde, A. R. Trans. Faraday&. 1057, 53, 628. (13) ReM, R. C.; Prausnk, J. M.; Sherwood, T. K. "The Properties of Gases and Liquids"; McGraw-Hill: New York, 1977; p 629. (14) Spaght, M. E.; Thomas, S. B.; Parks, G. S. J. f h y s . Chem. 1932, 36, 882. (15) Tlmmermans. J. "Physico-Chemical Constants of Pure Organic Compounds"; Elsevler: New York, 1950; Vol. I , pp 174, 177, 568, 1965; Voi. 11, pp 127, 130. (16) Weast, R. C. "CRC Handbook of Chemistry and Physics"; CRC Press: Boca Raton, FL, 1979; p c-381. (17) Welmer, R. F.; Prausnitz, J. M. J. Chem. Phys. 1965, 42, 3643.

Receiued for reuiew October 19, 1981 Revised manuscript received September 17, 1982 Accepted September 23, 1982

Distribution and Miscibility Limits in the System Ethanol-Water-Tri-n -butyl Phosphate-Diluent James W. Roddy" and Charles F. Coleman Chemical Technology Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee 37830

Distribution coefficients of (14C-) ethanol, DEmH, and f'H-) water, D,,, from aqueous solutions to TBPln-octane were measured at 25 O C . With undiluted TBP and >5 M aqueous ethanol, DEm is near unity; it decreases with dilution of either phase to about 0.005 at 0.03 M TBP and 0.01 M aqueous ethanol. The corresponding DHon decreases from about 0.1 to 0.00003, so that the separation factor DEmHlDwH increases from about 9 to about 160. The four-component phase diagram contains one (threedimensional)two-phase region, which shows that the maximum aqueous ethanol product concentration would be around 90 wt % in extraction with 1 M TBP and 65 wt % with undiluted TBP. Infrared spectral data in the frequency range of the hydroxyl and phosphoryl stretching modes were analyzed to assess the shifting nature of the intermolecular bonding, although they were not sufficient to identify the individual species.

Introduction As a part of the response to present and potential curtailment of petroleum supplies, significant attention is being given to the possibility of using fermentation ethanol as a fuel (Department of Energy, March 1978; Department of Energy, June 1979). A recent study by one of the present authors (Roddy, 1981) involved a series of scouting tests attempting to identify a solvent extraction system that could be substituted for the distillation step (a large energy user) in some significant portion of the separation of ethanol from fermentation broth. Although no system was found which produced as high a distribution coefficient as desired, several classes of compounds exhibited values warranting further study, with TBP systems indicated for first choice. The present investigation was undertaken to provide fundamental phase and distribution data basic to further study in this area. Distribution coefficents for both ethanol and water and the miscibility limits for the four-component system at 25 OC were determined in equilibrations of aqueous ethanol solutions (0.01 to 10 M) with TBP undiluted and diluted (3.65 to 0.03 M) with n-octane. In addition, the bonding in this system was probed by infrared spectral measurements in the frequency region of the O-H and P=O stretching modes. Experimental Section Radiotracer. The methods used for preparing and counting the tritiated water and the 14C-tagged ethanol 0196-4313/83/1022-0051$01.50/0

(a-carbon tagged) have been described previously (Roddy, 1981). Materials. The TBP was purified as discussed by Roddy and Mrochek (1966), and the methods of preparation of the other substances are described by Roddy ( 1981). Distribution Measurements. The procedures as given by Roddy (1981) were used for all equilibrations. Determination of Phase Diagrams. In addition to the direct analyses of the equilibrated phases (each equilibration establishing one tie line), the method described by Othmer et al. (1941) was used to locate the boundary of the solubility curve for the two-phase systems. A composition was selected which produced a clear solution (single phase). One of the components was then added from a micropipet until the solution became just turbid as detected by visual observation. The precision of this detection method is estimated to be better than 1%, The total amounts of the components used (obtained by weight gain of the titration vessel or by a weighing buret) represented a saturated phase and a point on the solubility curve. Successive additions of one of the other materials (usually ethanol) to produce a single phase and then of the initial component to produce turbidity again traced out the desired solubility curve. It is obvious that this method will not define tie lines. IR Spectral Procedure. Spectral measurements were made with a Beckman Model IR 4240 filterlgrating double-beam infrared spectrophotometer using matched cells @ 1983 American Chemical Society